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Full text of "A System Of Inorganic Chemistry"


m<OU 168494 



By the same Author 


Fnr the Tbt- <>f S( liuols and C.lli'ir^ 
SHIM" I'D^ n S\i>, I , ui iiiterlt ll^ i d, "'/'> 


TNO KG AN I C (MI K M 1ST 1 1 Y 

WILL LA M It AM SAY, IMi.l)., K.Il.S 

J. & A. C H U IJ C H 

11. NL\V JU IiTJNu'I o\ S'lK' 
1 Si) J 

I L 


FOR more than twenty years, the compounds of carbon have 
been classified in a rational manner ; and the relations 
between the different groups of compounds and between the 
individual members of the same groups have been placed in 
a clear light. It is, doubtless, owing to the brilliant origin- 
ators of this method of classification Kekule, llofmami, 
Wurtz, Frankland, and others too numerous to mention, but 
whose names occupy a prominent place in the history of our 
science that the domain of organic chemistry has been so 
systematically and successfully enlarged, and that it presents 
an aspect of orderly arrangement which can scarcely l)e 

This has unfortunately not been the fate of the chemistry 
of the other elements. Nearly twenty-five years have 
elapsed since the discovery by Newlands, Mendel6eff, and 
Meyer of the periodic arrangement of the elements ; and, in 
spite of the obvious guide to a similar classification which it 
furnishes, no systematic text-book has been written in English 
\v1th the periodic arrangement of the elements as a basis. 

The reasons for this neglect have probably been that the 
ancient and arbitrary line of demarcation between the non- 
metals and the metals has been adhered to ; that too great 
importance (from the standpoint of pure chemistry) has been 
assigned to the distinction between acid hydroxides and 
basic hydroxides (acids and bases), which has tended to 
obscure the fact that they belong essentially to the same 
class of compounds, viz., the hydroxides ; and that the 
chemistry of text-books has almost always been influenced 
by commercial considerations. The first of these reasons 


has often led, among other anomalies, to the separation of such 
closely allied elements as boron arid aluminium, antimony 
and bismuth, silicon and tin ; the second reason has often led 
to the ignoring of the double halides, except in a few special 
instances, and to the neglect of compounds such as double 
oxides of the sesquioxides of the iron group, in which these 
oxides play an acidic part ; while, for the third reason, those 
methods of preparing compounds which are of commercial 
importance are usually given, while other methods, as im- 
portant from a scientific point of view, are often ignored ; the 
borides, nitrides, &c., have been almost completely neglected 
since the time of Berzelius ; and the less easily obtained 
elements and compounds have been dismissed with scant 
iiotice because of their rarity ; whereas they should obviously 
be considered as important as the commoner ones in any 
treatise on scientific chemistry. 

The methods of classification adopted in this book are, as 
nearly as the difference of subject will permit, those which 
have led to the systematic arrangement of the carbon com- 
pounds. After a short historical preface, the elements are 
considered in their order; next their compounds with the 
halogens, including the double halides ; the oxides, sulphides, 
Selenides/and tellurides follow next, double oxides, such as 
sulphates, for example, being considered among the com- 
pounds of the simple oxides with the oxides of other ele- 
ments ; a few chapters are then occupied with the boridep, 
carbides, and silicides, and the nitrides, phosphides, arsen- 
ides, and antimonides ; and in these the organo-metallic coni- 
pounds, the double compounds of ammonia, and the cyanides 
are considered; while a short account is given of alloys and 
amalgams. The chemistry of the rare earths, which must at 
present be relegated to a suspense account, is treated along 
with spectrum analysis in a special chapter ; and the 
systematic portion of the book concludes with an account 
of the periodic table. 

The periodic arrangement has been departed from in two 
instances : the elements chromium, iron, manganese, cobalt, 
and nickel have been taken after those of the aluminium 


group ; and the elements copper, silver, gold, and mercury 
have been grouped together and considered after the other 
elements. It appeared to me that the analogies of these 
elements would have been obseuired, had the periodic arrange- 
ment been strictly adhered to. 

It has been thought desirable, instead of treating of 
processes of manufacture under the heading of the re- 
spective elements or compounds, to defer a description of 
them to the end of the book, and to group them under 
special headings, those compounds being considered together 
which are generally manufactured under one roof. In 
describing manufactures, chemical principles have been con- 
sidered, rather than the apparatus by means of which the 
manufactures are carried on. The student, having acquired 
the requisite acquaintance with facts, is now better able to 
appreciate these principles. 

The physical aspects of chemistry have generally been 
kept in the background, and introduced only when necessary 
to explain modern theories. I hold that a student should 
have a fair knowledge of a wide ran go of facts before he 
proceeds to the study of physical chemistry, which, indeed, 
is a science in itself. But short tables of the more important 
physical properties of elements, and of the simpler com- 
pounds, have been introduced for purposes of reference. 

It may be asked if such a system is easily grasped by the 
student, and if it is convenient for the teacher. To this 
question I can reply that, having used it for four years, I am 
'perfectly satisfied with the results. For the student, memory 
work is lightened ; for the teacher, the long tedious descrip- 
tion of metals and their salts is avoided ; and I have found 
that the student's interest is retained, owing to the fact that 
all the "fire-works" are not displayed at the beginning of 
the course, but are distributed pretty evenly throughout. 

It need hardly be mentioned that the teacher is not 
required to tpach, nor the student to remember, all the facts 
as they are here set forth. It is necessary to make a 
judicious selection. But it is of advantage to have the list 
fairly complete for purposes of reference. It should be stated 

that, in the case of compounds of questionable existence, 
they have received the benefit of the doubt. It is at least 
well that they should be known, in order that their existei>ce 
may be brought to the test of renewed experiment. 

References to original memoirs have been given where 
important theoretical points are involved: or where doubt 
exists ; and an attempt has been made to guide the reading 
of students. As a rule, references to recent papers are 
given ; the older references may be found in one of the 
chemical dictionaries. 


January, 1891. 




CHAPTER I. Introductory and Historical .. .. .. .. .. 1 

CHAPTER II. Historical 14 


CHAPTER III. Group 1. Hydrogen, lithium, sodium, potassium, 

rubidium, and caesium . . . . . . . . . . . . 25 

Group 2. Beryllium, or glucinum, calcium, strontium, barium . . 31 

Group 3. Magnesium, zinc, cadmium . . . . . . . . 33 

Group 4. Boron, scandium, (yttrium), lanthanum, (ytterbium) .. 35 

Group 5. Aluminium, gallium, indium, thallium , . . . , . 37 

CHAPTER IV. Group 6. Chromium, iron, manganese, cobalt, nickel . . 40 

Group 7. Carbon, titanium, zirconium, cerium, thorium . . . . 43 

Group 8. Silicon, germanium, tin, (terbium), lead.. .. .. 40 

CHAPTER V. Group 9. Nitrogen, vanadium, niobium, (didymium), 

tantalum . . . . . . , . . . . . . . . . f>$ 

Group 10. Phosphorus, arsenic, antimony, (erbium), bismuth . . 5G 
Group 11. (Oxygen, chromium). Molybdenum, tungsten, uran- 
ium . . GO 

Group 12. Oxygen, sulphur, selenium, tellurium . . . . . . 61 

Appendix. Air . . . . . . , . , , . , . . . . 70 


CHAPTER VI. Group 13. Fluorine, chlorine, bromine, iodine . . . . 72 

Groups 14 and 15. Ruthenium, rhodium, palladium, osmium. 

iridium, platinum . . . . . . . . . . . . . 77 

Group 16. Copper, silver, gold, mercury . , . . . , . . 79 

General remarks on the elements , . . . . . . . . . B3 


CHAPTER VII. Compounds and mixtures ; nomenclature . . . . 88 
The states of matter ; Boyle's law ; Gay-Lussac's law ; Avogadro's 

law 91 

Methods of determining the densities of gases . . . . . . 97 


CHAPTER VIII. Hydrogen fluoride, chloride, bromide, and iodide . . 104 
Halides of lithium, sodium, potassium, rubidium, caesium, and 

ammonium . . . . . . . . . . . . . . . .115 

CHAPTER IX. Halides of beryllium, calcium, strontium, and barium .. 120 

Halides of magnesium, zinc, and cadmium . k .. .. ..* 123 

Molecular formulae ; specific heats of elemeuts . . . . . . 126 

CHAPTER X. Halidea of boron, scandium, (yttrium), and lanthanum, 

(ytterbium) 131 

Halides of aluminium, gallium, indium, and thallium . . . . 133 

Halides of chromium, iron, manganese, cobalt, and nickel. . . . 137 

CHAPTER XI. Halides of carbon, titanium, zirconium, cerium, and 

thorium , 144 

Halides of silicon, germanium, tin, (terbium), and lead .. .. 148 

CHAPTER XII. Halides of nitrogen, vanadium, niobium, tantalum . . 157 

Halides of phosphorus, arsenic, antimony,, (erbium), and bismuth. . 160 

Halides of molybdenum, tungsten, and uranium . . . . . . 164 

Halides of sulphur, selenium, and tellurium, . . . . . . . 166 

CHAPTER XIII. Compounds of the halogens with each other . . . . 169 

Halides of ruthenium, rhodium, and palladium 170 

Halides of osmium, iridium, and platinum . . . . . . . . 172 

Halides of copper, silver, gold, and mercury. . . . . . .. 174 

CHAPTER XIV. Eeview of the halides ; their sources, preparation, and 
properties j their combinations and their reactions j also their 

molecular formulee . . 181 



CHAPTER XV. Compounds of oxygen, sulphur, selenium, and tellurium ^ 

with hydrogen . . . . . . . . . , . . ., 191 

Physical properties of water . . . . . .. . . . . 199 

Compounds of water with halides . . . . . . . 203 

CHAPTER XVI. Classification of oxides . . 205 

The dualistic theory 207 

Constitutional and rational formulae 208 

Oxides, sulphides, &c., of lithium, sodium, potassium, rubidium, 

caesium, and ammonium .. .. .. ..211 

Hydroxides and hydrosulphides . . 214 

CHAPTER XVII. Oxides, sulphides, and selenides of beryllium, calcium, 

strontium, and barium 218 

Hydroxides and hydrosulphides . . . . 222 

Oxides, sulphides, selenides, and tellurides of magnesium, zinc, 

and cadmium j . . 225 



Hydroxides and hydrosulphides . . . . . . . . . . 229 

Double oxides (zincates) j oxyhalides.. .. .. .. .. 229 

CHAPTER XYIII. Oxides and sulphides of boron, scandium, (yttrium), 

* lanthanum, (and ytterbium) . . . . . . . , . . 232 

Double oxides (boracic acid and borates) . . . . . . . . 233 

Oxyhalides 236 

Oxides, sulphides, and selenides of aluminium, gallium, indium, 

and thallium 237 

Hydroxides and double oxides (aluminates, &c.) . . . . . . 239 

Double sulphides and oxyhalides . . . . . . . . . . 242 

CHAPTER XIX. Monoxides, monosulphides, monoselenides, and mono- 

tellurides of chromium, iron, manganese, cobalt, and nickel . . 243 

Dihydi>xides ; double sulphides . . k . . . . . . . 246 

Sesquioxides, sesquisulphides, and sesquiaelenides . . . . . . 248 

Trihydroxides 251 

Double oxides (spinels).. .. .. .. .. . . .. 253 

Double sulphides and oxyhalides . . . . . . . . . . 256 

Dioxides and disulphides . . . . . . . . . . . 258 

Hydrated dioxides and double oxides (manganites) . . . . . . 260 

Trioxides 261 

Double oxides (chromates, ferrates, and manganates) .. .. 262 

Perchromates and permanganates . . . . . . . . . . 266 

Oxyhalides 208 

CHAPTER XX. Monoxides and monosulphides of carbon, titanium, 

zirconium, cerium, and thorium . . , . . . . . 270 

Sesquioxides and sesquisulphides . . . . 273 

Dioxides and disulphides . . . . . . . . . . . . 274 

Compounds with water and with hydrogen sulphide . . . . 283 

Carbonates, tilanates, zirconates, thorates; carbon oxysulphide; 

oxysulphocarbonates and sulphocarbonates . . . . . . 284 

Oxyhalides 292 

CHAPTER XXI. Monoxides, monosulphides, monoselenides, and mono- 

tellurides of silicon, germanium, tin, and lead . . . . . . 294 

Hydroxides ; compounds with oxides and with halides . . . . 297 

Sesquioxides and sesquisulphides . . . . . . . . . . 299 

Dioxides, disulphides, diselemdes, and ditelluride , . , . . . 300 
Compounds with water and oxides : silicates, stannates, and 

plumbates , 303 

Sulphostannates . . . . . , . . . . 316 

Oxyhalides 317 

CHAPTER XXII. Oxides and sulphides of nitrogen, vanadium, niobium, 

and tantalum . . . , . . 319 

Compounds of pentoxides with water and oxides ; nitric, yanadic, 
niobicy and tantalic acids : nitrates, vanadates, niobates, and 

tantalates 322 

Oxyhalides 331 



Tetroxides or dioxides : tetrasulphide or disulphide . . . . 333 

Compounds with oxides and sulphides ; hypoyanadates and hj po- 

sulphovanadates . . . . . . . . . . . . . . 335 

Compounds with hahdes. Trioxides . . . . . . . . . . 336 

Nitrites and vanadites . . . . . . . . . . . .*. 337 

Compounds with halides . . . . . . . . . . . . 340 

Nitric oxide ; vanadium monoxide . . . . . . . . . 341 

Nitrogen sulphide and selemde. Nitrosulpludes . . . . . . 343 

Nitrous oxide ; hyponitrites . . . . . . . . . . . . 343 

CHAPTER XXIII. Oxides, sulphides, selenides, and tellurides of phos- 
phorus, arsenic, antimony, and bismuth. . . . . . _ . 346 

Compounds of the pentoxides arid pentasulphides ; orthophos- 

phates, orthoarsenates, and orthoantimonates, &c. .. .. 352 

Pyrophosphates, pyrarsenates, and pyrantimonates . . %. ... 363 

Metaphosphates, metarsenates, and metantimonates . . . . 369 

CiiAi'TEB XXIV. Ilypophosphoric acid .. .. .. .. .. 373 

Compounds of trioxides and trisulphides ; phosphites, arsemleb, 

and antimomtes ; their sulphur analogues . . . . . . 374 

11} pophosplutes . . . . . . . . . . 380 

Compounds of oxides and sulphides with halides . . . . . . 332 

CHAPTER XXV. Ozone (oxide of oxygen) . . . . . . . . . . 387 

Oxides and sulphides of molybdenum, tungsten, and uranium . . 392 
Hydroxides ; molybdates, tungstates, and uranates ; sulphur 

analogues . . . . . . . . . . . . . . . . 396 

Peruranates, persulphomolybdatcs . . . . . . . . . . 405 

Compounds with huxidea . . . . . . . . . . . . 406 

CHAPTER XXVI. Oxides of sulphur, selenium, and tellurium . . . . 409 

Sulphuric, selenic, and telluric acids . . . . . . . . . . 414 

Sulphates, selenates, and tellurates . . . . . . . . . . 419 

Anhydrosulphurio acid and anhydrosulphates . . . . . . 432 

CHAPTER XXVII. Sulphurous, selenious, and tellurous acids . . . . 43 

Sulphites, selenites, and tellurites * 436 

Compounds of oxides with halides j sulphury 1 chloride ; chloro- 

sulphonic acid . . . . . . . . . . . . . . 440 

Other acids of sulphur and selenium . . . . . . . . . . 443 

Thiosulphates . . . . 444 

Seleniosulphates. . . . . . . . . . 447 

Hyposulphurous acid and hyposulphites . . . . . . . . 447 

Dithionic acid and dithionates , . . . . . . . . . . . 448 

Trithionic acid and trithionates 449 

Seleniotrithionic acid $ tetrathionic acid . . . . . . . . 450 

Pentathionic acid 451 

Hexathionic acid. Constitution of the acids of sulphur and 

selenium . . . . . . . , . . . . . . . . 452 

Nitrososulphates 455 

Compounds of sulphur, selenium, and tellurium with each other . . 455 



CHAPTER XXVIII. Oxides of chlorine, bromine, and iodine . . . . 459 

Hypochiorites, hypobromites, and hypoiodites . . . . . . 4G1 

Chlorous acid and chlorites . . . . . . . . . . . . 464 

Chlorates, bromates, and iodates . . . . . . . . . . 464 

Perchlorates and periodates . . . . . . . . . . . . 469 

CHAPTER XXIX. Oxides*, sulphides, and selcnides of rhodium, ruthen- 
ium, and palladium .. .. .. . . .. .. 476 

Hydroxides 478 

Sulphopalladites ; ruthenates and perruthenates .. .. .. 479 

Oxides, sulphides, and selenides of osmium, iridium, and platinum 4-80 

Hydroxides 182 

Osmites and platinates . . . . . . . . , . . . . . 483 

Platmonitrites ; platinochlorosulphites ; platimcarbonyl com- 
pounds ; dichloroplahmphosphonic acid, . . . , . . . 485 
Oxides, sulphides, selenidcs, and tellurides of copper, silver, gold, 

and mercury . . . , . . . . . . . . . . 487 

Hydroxides . . . . . . . . . . . . . . . . 41)1 

Aurates. Double sulphides. Oxy- and sulpho-hahdes . . . . 492 

Concluding remarks on the oxides, sulphide*, &c ; classification of 

oxides . . , . . . . . . . . . . . . , 494 

Constitutional formulae ; oxyhalidca and double Imlides . . . . 495 


CHAPTER XXX. The borides ; hydrogen borido .. .. .. .. 497 

Magnesium, aluminium, manganese, silver, and iron borides . 198 

The carbides and silicides ; methane, or marsh -gas . . . . . . t-98 

Hydrogen sihcide , . . . . . . . . . . . . 500 

Ethane ; silicoethane . . , . . . . . . . . . 501 

Double compounds of ethyl and methyl , " organo-metallic " com- 
pounds . . . . . . . . . . . . . . . . 502 

'Ethylene 507 

Acetylene. . . . . . . . . . . . . . . 508 

Carbides and silicides of iron, &c. . . . . . . . 510 



CHASTER XXXI. Hydrogen nitrides, phosphides, arsenide, and anti- 

monide ; ammonia, hydrazine, hydrazoic acid, Ac. . . . . 512 

Salts of phosphonium .. 517 

Hydroxylamine . . . . . . 523 



Amido-corapounds or amines . . . . . . . . . 524 

Salts of the amines . . . . . . . . . . . . . . 525 

Chromamme salts . . . . . . . . . . . . . . 526 

Cobaltamine salts 528 

Methylaraine, &c. ; the phosphines and arsines . . . . . * 532 

Carbamide 532 

Stlicamines, titanamine, and zirconamine salts . . . . . . 533 

Amides of phosphorus ; phosphamio acids . . . . . . . 534 

Su^phaminos (sulphamic acids) .. .. .. .. .. 536 

Amjnes of rhodium, ruthenium, and palladium .. .. .. 537 

Osmamines, iridamines, and platinammes . . . . . . . . 539 

Cupramines, argentamines, auramines, and mercuramines . . . . 545 

CHAPTER XXXII. The nitrides, phosphides, arsenides, and antimonidcs. 550 

Cyanogen (carbon nitride) and its compounds . . .% . . 558 

Ferrocyanides and ferricyanides, and analogous compounds .. 562 

Platino- and platini-cyanides, and similar compounds . . . . 570 

Constitution of cyanides . . . . . . . . , . . . 572 


CHAPTER XXXIII. Alloys. Hydrides 575 

Alloys of lithium, sodium, potassium, &c. . . . . . . . . 577 

calcium, barium, magnesium, zinc, &c . . . . . . 578 

aluminium, chromium, iron, &c. . . , . . . . . 581 

tin and lead . . . . . . . . . . . . . . 585 

antimony and bismuth . . . . . . . . . . 587 

the palladium and platinum metals . . . . . . 588 

copper, silver, gold, and mercury . . . . . . . . 589 


CHAPTER XXXIV. Spectrum analysis 591 C 

Spectroscopy applied to the determination of atomic weights . . 598 
The rare earths ; the didyrnium group j the erbium group ; the 

yttrium group . . . . . . . , . . . . . . 602 

Solar and stellar spectra . . . . . . . . . . . . 606 

CHAPTER XXXV. The atomic and molecular weights of elements, and 

the molecular weights of compounds . . . . . , . . 611 

The specific heats of elements and compounds . . . . . . 617 

The law of replacement 619 

Isomorphism . .. . . .. . . . . .. 620 

The complexity of molecules .. .. .. .. .. ., 621 

The monatomic nature of mercury gas . . . . . . , . 624 

CHAPTER XXXVI. The periodic arrangement of the elements . . .. 627 

Numerical relations between atomic weights. . . . . , . . 629 


Eelations between atomic weights and physical properties of 

elements 633 

Comparison of the elements and their compounds . . . . . . 634 

Prediction of undiscovered elements . . . . . , . . . . 639 


CHAPTER XXXVII. Processes of manufacture 642 

Combustion ; fuels . . . . . . . . . . . . . . 642 

CHAPTER XXXVIII. Commercial preparation of the elements . . .. 651 

Manufacture of sodium. . . . . . . . . . . . 651 

magnesium, zinc, and aluminium . . . . . . 652 

iron and steel. . ., .. .. .. .. 653 

nickel 658 

tin and lead 659 

,, antimony . . . . . . . . . . . . 660 

bismuth and copper . . . . . . . . . . 661 

silver 662 

gold 663 

mercury . . 664 

phosphorus , . . . . . . 665 

CHAPTER XXXIX. Utilisation of sulphur dioxide 667 

Manufacture of sulphuric acid. . . . . . . . . . 667 

alkali 670 

Preparation of sodium sulphate .. .. . .. .. 672 

The Lcblanc soda-process . . . . ' 673 

Manufacture of caustic soda . . . . . . . . . . 675 

Utilisation of tank-waste . . . . . . . . 676 

Manufacture of chlorine . . . . . . . . . 678 

bleaching powder . . . . . . . . . . 681 

potassium chlorate . . . . . . . 682 

The ammonia- soda process . . . . . . 683 



THK first object of the Science of Chemistry is to ascertain the 
composition of the various things which we see around us. Thus, 
among familiar objects are air, water, rocks and stones, earth, the 
bark, wood, and leaves of plants, the flesh, fat, and bones of animals, 
and so on. Of what do these things consist ? 

The second is to ask, Can such things be made artificially, and, 
if so, by what methods ? Attempts to answer these questions 
have led to the discovery of many different kinds of matter, some 
of which have as yet resisted nil efforts to split them up into still 
simpler forms. Such ultimate kinds of matter are termed elements. 
But other kinds of matter can often be produced when two or 
more of the simpler forms or elements are brought together ; the 
elements are then said to comlitie, and the new substances resulting 
from their combination are called compounds. 

The third object of the Science of Chemistry is the correct 
classification of the elements and of their compounds ; those sub- 
stances which are produced in a similar manner, or which act in a 
.rim-far manner when treated similarly, being placed in the same 

The fourth, inquiry relates to the changes which diifercnt forms 
of matter undergo when they unite with each other, or when they 
split into simpler forms. 

Fifthly, the conditions of change are themselves compared 
with each other and classified ; and thus general laws are being 
deduced, applicable to all such changes. 

Lastly, the Science of Chemistry and the sister Science of 
Physics join in speculations regarding the nature and structure of 
matter, in the hope that it may ultimately be possible to account 
for its various forms, the changes which they undergo, and the 
relations existing between them. 



To answer questions such as these, it is obvious that experi- 
ments must be made. Each form of matter must be separately 
exposed to different conditions ; heated, for example ; or placed 
under the influence of an electric current; or brought together 
with other kinds of matter ; before it is possible to know what ft will 
do. Now, the ancient philosophers did not perceive this necessity ; 
nor indeed were they much concerned in making tlio inquiry. 
Those nations which have left behind them a record of their 
thoughts, the ancient inhabitants of India, Egypt, Greece, and 
Rome, devoted their attention, if they aspired to be learned, to 
oratory, to history, or to poetry. Their only scientific pursuits 
were politics, ethics, and mathematics. Distinction was to be 
gained in the forum, in the temple, or in the battlefield ; not in 
wresting secrets from Nature. The practice of such of the arts as 
wore then known was in the hands of slaves and the lower classes 
of the people, who were content to transmit their methods from 
lather to son, and whose achievements were unchronicled. The 
citizens of the State, the wealthy and the leisured, despised these 
low-class arts; and, indeed, it was taught by Socrates and his fol- 
lowers that it was foolish to abandon the study of those things 
which more nearly concern man, for that of things external 
to him. It was generally believed that by the exercise of pure 
thought, without careful observation and experiment, a man 
could know best the true nature of the objects external to him. 
Thus Plato says in the 7th Book of the " Republic," " We shall 
pursue Astronomy with the help of problems, just as we pursue 
Geometry; but, if it is our design to become really acquainted with 
Astronomy, we shall let the heavenly bodies alone." Elsewhere 
he states that, even if we were to ascertain these things, we could 
neither alter the course of the stars, nor apply our knowledge so as 
to benefit mankind. And in "Tirnanis," Plato remarks, " God only 
has the knowledge and the power which are able to combine many 
things into one, and to dissolve the one into the many. But no man 
either is or ever will be able to accomplish either the one or the 
other operation." 

It was impossible, with such a mental attitude towards science, 
for any accurate knowledge to exist, or for any probable theories 
to be devised. Yet, as it is interesting to know something of the 
old ideas concerning matter and its nature, a short sketch will be 
given here. 

The origin of the world was for the ancient philosophers of 
Egypt, India, and Greece, as it is for ourselves, a subject of the 
greatest interest ; and in attempting to frame some theory to explain 


the Creation it was necessary to speculate on the nature of matter. 
The various aspects of matter which we see around us were sup- 
posed by Empedocles (492 B.C.), and later by Aristotle (384 iu.), 
to be modifications of one fundamental original material, occurring 
in various forms, the difference between which was caused by the 
"assumption of certain " elements," or as we should now namo them 
" properties." This original material was imagined by Empedocles 
to consist of small particles, which he termed atoms, or "indi- 
visibles," because they were in his view the ultimate particles into 
which matter could be divided. Plato imagined such atoms to 
have the form of triangles of different sizes, equilateral, isosceles, 
or scalene ; and ascribed the " perfect ion " or " i in perfection " of 
matter to \^ due to the form of its ultimate partieleH. But such 
particles were modified by the "elements" earth, water, air, and 
Jire ; that is, they assumed a solid, liquid, aeriform, or flaming- 
nature, according to the element which predominated in them 
Along with this view, a certain confusion of thought arose which 
led to the conception that earth, water, air, and tire were actually 
present in, and constituents of, matter, and that all the elements 
originated in one, supposed by Thales (t)00 B.C.) to bo water, and 
by Anaximenes (about 550 u.c.) to be air or tire. The well- 
known poem of Lucretius, De rcrum natitrd, is a transcript of 
these views of the atomic constitution of the universe. But such 
speculations were wholly without a basis of fact, and led to no 
new knowledge. These ideas, in all probability, were originally 
derived from India, where the four elements already men- 
tioned were associated with a fifth and sixth, ether and 
consciousness, as appears from the teaching of Buddha. The 
notion that matter was one in kind, modified by certain attributes, 
developed the belief that by changing the attributes, the matter 
itself would be transmuted. Thus Tinwms is made to say by 
Plato: " In the first place, that which we are now calling water, 
when congealed, becomes stone and earth, as our sight seems to 
show us [here he refers probably to rock-crystal, a transparent, 
hard material, which was supposed to be petrified ice] ; and this 
same element, when melted and dispersed, passes into vapour and 
fire. Air again, when burnt up, becomes fire, and again tire, when 
condensed and extinguished, passes once more into the form of air ; 
and once more air, when collected and condensed, produces cloud 
and vapour ; and from these, when still more compressed, comrn 
flowing water ; and from water come earth and stones once more ; 
and thus generation seems to be transmitted from one to the other 
in a circle.** Here the elements are evidently conceived in their 

* B 2 


concrete sense ; but he goes on to say that certain matter in a state 
of change partakes of the nature of fire to some extent, and to some 
extent of the nature of the other elements. . 

Guided by such considerations, Aristotle gave precision to these 
speculations by his system of "contraries." The properties shared 
by all matter in varying proportions were " hotness," '* coldness," 
" moistness," and "dryness." Thus air was hot and moist; fire, 
hot and dry; water, cold and moist; and earth, cold and dry. By 
imparting heat to water it becomes steam, that is, air, hot and 
moist; by taking away its moisture it becomes earth, that is, ice, 
cold and dry. 

The "Timnpus" of Plato, which has been quoted several times 
here, was held in high esteem by the great school of learning which 
existed at Alexandria during the first centuries of our era. It was 
here, in all likelihood, that the second great era of chemicnl theory 
began. Based on such ideas regarding the constitution of matter, 
attempts wero made to change one substance into another, arid 
above all to transmute the baser metals into gold. The attempt 
was called Alchemy (the Arabic prefix al signifying "the"), and 
from that word, which probably means the dark or secret art, is 
derived our modern name Chemistry. 

In order to realise the atfitute of mind which led to the belief 
in the possibility of the transmutation of mrfah, as the change of 
one metal into another is called, wo mnst note that it was supposed 
that the apparent change of one form of matter into another in- 
volved the destruction of the first form, and the creation of the 
second ; the properties of the matter were changed, and hence the 
matter itself was supposed to be changed ; no attempt was made, 
so far as we know, to compare the weights (or masses) of the 
matter before and after the change had taken place. Pure sub- 
stances, moreover, were almost unknown, and the separation of an 
impurity from a compound in many cases entirely altered its pro- 
perties. Now, the Arabs, who conquered Egypt in the 7th century, 
and transmitted their knowledge to posterity, possessed a theoiy 
of which we learn in the writings of Qeber, an Arabian alchemist 
of the 8th century, and in which we can trace a germ of the 
modern views concerning matter, inasmuch as we find here the 
first dawn of a conception of a chemical compound, in the modern 
sense of the word. 

Geber, and probably his predecessors the Alexandrians, re- 
garded the metals as alloys of mercury and sulphur in varying 
proportions. Now-a-days, mercury is the name of a metal which 
possesses definite unalterable properties ; nor does sulphur vary, 


but is always a distinct substance capable of certain changes, 
though radically the same throughout these changes. But Geher 
held that the mercury and the sulphur each varied in kind and in 
properties, and were not what we should now term definite 
chemical individuals. His views inny best bo learned from his 
own words : "It in folly to attempt to extract one substance from 
another wlilc.lt Joes not contain if. But, as all metals consist of mer- 
cury and sulphur, it is possible to add to one what is wanting, or 
to take from another what is in excess." Yet he did not discard 
the older elements, earth, water, air, and fire, but appears to have 
regarded them as more re-mote constituents of mutter, while mer- 
cury and sulphur were the proximate constituents. The mercury 
was suppoMjd to impart to metals their brilliancy, their malleability, 
and their fusibility ; while the sulphur which they contained ren- 
dered them alterable by fire, which changes many metals to earthy 
powders. Tn his writings also we find the first allusion to a con- 
nection between the curing of disease and the transmutation of 
metals, in his illustration, '' Bring me the six leper.*, and I will 
heal them," referring to the conversion of six of the metals then 
known into gold, the seventh. 

It is not wonderful that the alchemists should have been led 
into error by attempts to transmute the metals into gold, for 
their properties are radically changed by the presence of mere 
traces of foreign bodies. 

Thus the presence of a minute trace of lead or arsenic, for 
example, renders gold exceedingly brittle, and alters its colour; 
the presence of a very small quantity of carbon in iron renders it 
elastic, or if more be present, hard and brittle; a small amount of 
arsenic in copper colours it white, and lowers enormously its power 
of conducting electricity. These changes, which are still un- 
explained, received much more attention in the early days of 
chemistry than of recent years ; but it is to be hoped that they 
wilF again be exhaustively studied. 

For many centuries after Geber's time, although numerous 
compounds were discovered in the search for gold, no new 
development of theory can be noticed. The attitude of the 
Roman Church was hostile to the progress of any knowlege of 
nature. All learning was in the hands of the priests, and the 
study of the ancient writers was discouraged or forbidden, as not 
only useless in itself, but as tending to distract the mind from the 
higher studies of Divine things. When permitted on sufferance it 
was with the avov\ed object of combatting disbelief with its own 
weapons. Roger Bacon (1214 1291; even, who published 


several works on alchemy, wrote that that science which is not 
prosecuted with a view of defending the Christian faith leads " to 
the darkness of hell." Yet after the conquest of Spain by -the 
-Arabs, in the beginning of the 8th century, the study of medicine 
of mathematics, and of optics slowly grew in the west of Europe 
The works of many of the Greek philosophers were known onl 
through Arabic translations. It was not in the nature of the 
Arabs to originate new theories ; they merely preserved those 
which they had. 

In the 15th century, Basil Valentine, a Benedictine friar, 
added one to the two supposed constituents of metals. This 
was a principle of fixity ; something which resisted the action 
of heat without volatilising into gas. Valentine "Permed this 
principle " salt." Again, it is not to be understood that any 
particular salt is referred to ; yet, in the minds of many of Basil 
Valentine's disciples, the earthy residue obtained by calcining 
metals such as lead in the air was regarded as the same in essence, 
from whatever source it had been derived. And this theory was 
extended to include all matter ; it is well described in the words of 
Paracelsus (born 1493, died 1541). "In all things four elements 
are mingled with each other ; among those four, only one is fixed 
and perfect ; in that element lies the true ' quintessence.' The 
other elements are imperfect; yet any one of them is able to 
tinge and qualify the ouiers, according to its nature. Thus in 
some the element water predominates ; in some fire ; in some 
earth ; in others air. In order to separate the predominating 
element as salt, sulphur, and mercury, each must be broken and 
destroyed by solution, and calcination, or by such means." 
" There are various minerals in which the elements are not ?o 
firmly locked up as in the metals, and which can be split into theiij 
three principles: salt, the fixed element; sulphur, fiery andoily ; 
and mercury, aeriform and watery." ( Wunsch Huttlein, Erfurt, 
1738, p. 27.) 

The language of the mediaeval alchemists is most obscure. 
Not only did they confuse substances now known to be perfectly 
distinct ; not only was their nomenclature ambiguous ; but a 
spirit of mysticism pervaded their writings which led them to 
believe that it would have been impious to reveal to the common 
people the processes with which they were acquainted. Chemical 
elements and compounds form in the pages of their books proces- 
sions of kings and queens, bridegrooms and brides, lions and 
dragons, eagles and swans ; gold is the Sun ; silver the Moon ; this 
king and queen^, Apollo and Diana, are devoured, on their bridal 


eve, by Saturn, (lead), a dragon and serpent which has for ages 
slept in his rocky cavern. Pluto enters with exceeding heat, 
expelling the dragon as an eagle with scorched wings, and leaving 
the royal pair reposing on a bed, white as the mountain snows. 
Such is Basil Valentine's account of the refining of gold in the 
second of his " Twelve keys to unlock the door leading to the 
ancient stone '*; among innumerable descriptions of the kind it is 
one of the few in which the actual processes can be followed under 
their mystical disguise. We meet with fanciful analogies between 
the Divine Trinity ; the human body, soul, and spirit ; and the 
trio of salt, sulphur, and mercury, in which religion, medicine, and 
chemistry are mingled in inextricable entanglement. Yet such 
analogies erved a good purpose ; they led to the combatting of 
disease, not as before with charms and incantations, and remedies 
of a disgusting and fantastic nature, but by the administration of 
chemical substances as drugs. In spite of all their false theories, 
connecting certain of the organisms of the bodies with the stars, 
and these again with the metals, the compounds of which were 
supposed to act on those organisms with which they were so 
fancifully related, true progress was made by the only way in 
which progress is possible by experiment and deduction ; and the 
virtues of antimony, of mercury, and of other remedies were 
gradually discovered. This change in the ultimate goal of chemical 
research was begun by Basil Valentine ; its chief advocate, how- 
ever, was Paracelsus, who boldly announced tha,t " The true scope 
of chemistry is not to make gold, but to prepare medicines." Yet 
in his writings there are numerous receipts for the preparation of 
the alcuhest, or universal solvent; and of that magic elixir, capable 
not only of converting baser metals into gold, but of conferring on 
its fortunate possessor long life and eternal youth. 

Although no advance in chemical theory was made by the 
school of alchemists represented by Valentine and Paracelsus, yet 
the indefatigable labours of these men and of their disciples 
enriched chemistry by the discovery of many new compounds, 
and laid a foundation of facts for chemists of a later age. 

The era of modern chemistry opens with Robert Boyle 
(1626 1691). In his works, which are very voluminous, we meet 
with no traces of the spirit of mysticism which prevailed up to his 
time, but he manifests that aspect of rational inquiry which is 
typical of modern science. His most important work on chemistry 
is named " The Sceptical Chymist, or considerations upon the 
experiments usually produced in favour of the four elements, and 
of the three chymical principles of the mixed bodies." In it he 


defines the word element : " The words element and principle arc 
here used as equivalent terms : and signify those primitive and 
simple bodies of which the mixed ones are said to bo composed, 
and into which they are ultimately resolved. 'Tis said that a 
piece of green wood, by burning, discovers the four element^ of 
which mixed bodies are composed, the fire appearing in the flame 
by its own light; the smoke ascending and readily turning into air, 
as a river mixes with the sea ; the water, in its own form, boiling 
out at the end of the stick, and the ashes remaining for the element 

of earth But there are many bodies from whence 

it seems impossible to extract four elements by fire, and which of 
these can be obtained from gold by any degree of fire whatsoever ?" 
He then proceeds to consider in detail the supposed evidence for 
the existence of the Aristotelian elements, and of the principles, 
salt, sulphur, and mercury ; and he finally shows that they cannot 
be considered as elements, using the word in the modern sense of 
the constituents of bodies ; and he incidentally points out that 
compounds, as a rule, do not resemble the elements of which they 
are composed. 

Boyle thus successfully combatted the ancient doctrines of the 
alchemists ; and although a belief in alchemy lingered on into the 
last century, and has even had a few disciples in our own day, yet 
the formation of the learned societies of Florence (1657), London 
(1662), Paris (1666), and Vienna, and the open interchange of 
ideas among men who submitted every doubtful point to the test of 
experiment, did much to destroy the veil of mysticism which had 
surrounded the labours of the ancient alchemists, and has ulti- 
mately proved the correctness of Boyle's views. 

The phenomena of burning and combustion played a great part 
in the theories of the mediaeval alchemists, and were held to sub- 
stantiate their views. For when a candle burns, it disappears * v the 
solid is changed into air and flame ; a transmutation has taken 
place, proving the identity of the elements. Many metals, when 
heated in air, are converted into earthy powders, differing entirely 
from their originals in properties. Paracelsus seems to have 
imagined that in certain similar cases, for example when iron 
pyrites lies exposed to air, "the old demogorgon," as he calls this 
compound of iron and sulphur, " absorbs the universal salt, 
whereby it is converted into a greyish crystalline powder." But 
no consistent theory had been advanced to account for the pheno- 
mena of combustion. John Mayow (16451679), indeed, a con- 
temporary of Boyle's, a medical graduate of Oxford, who practised 
at Bath, had he lived, would in all probability have advanced the 


knowledge of chemical theory to the stage which it reached more 
than a centmy after his death. During his short life, he antici- 
p^ted most of the deductions of Lavoisier, who, as we shall shortly 
see, effected a revolution in the science. Guided by Boyle's 
researches, and with a rare faculty of devising experiments admi- 
rably adapted to decide the points at issue, Mayow pointed out 
that atmospheric air consists of two kinds, one capable of sup- 
porting combustion and life, which he named " spiriius igno-aerifs" 
and another, devoid of these properties. He concluded from 
his experiments that this " spiritus igno-aerius^ or to give 
it Lavoisier's name, oxygen, was a constituent of nitre or salt- 
petre, and was also contained in nitric acid ; that when oxygen 
combines ,with other bodies, such as metals, it increases their 
weight ; that it is the common constituent of acids, sulphuric acid 
being its compound with sulphur, and nitric acid its compound 
with the inactive constituent of air, now known as nitrogen ; and 
he also devised a method of estimating oxygen by mixing with it 
one of the compounds of nitrogen and oxygon, nitric oxid<* 9 a process 
which was afterwards largely employed, and which has been re- 
cently revived. Lastly, he showed the function of oxygen in acid 
fermentation, and, in his " Tractatus yuinque medico -physici," in 
which his investigations and conclusions are recorded, he showed 
very clearly the part played by oxygen in restoring venous blood 
to the arterial state, and in maintaining animal heat. 

But Mayow was alone in his work ; his early death cut short 
his researches ; and his contemporaries and successors did not recog- 
nise their merit. Stephen Hales, for example, though he pre- 
pared in an impure state carbonic acid, nitrogen, hydrogen, and 
oxygen gases, and also marsh-gas, regarded them all as modifica- 
tions of air, not as distinct gaseous substances. He ascribed to 
atmospheric air " a chaotic nature," inasmuch as it was found to 
be Qndowed with so many different properties. 

But in spite of Mayow's correct surmises regarding the nature 
of combustion, the opinion which chemists generally held was that 
when a body was burnt something escaped from it, viz., fire or 
heat. For although it was well known that combustible sub" 
stances do not continue to burn in a confined space, this was 
attributed not to the exclusion of air, but to the prevention of the 
escape of flame. And in spite of its having been noticed by Boyle 
and others that metals gain in weight by being calcined, yet no 
special attention was paid to the fact. So long indeed as what we 
now know to be different kinds of gases were assumed to be only 
common air containing impurities, it was impossible to account 


for the apparent loss of weight which many combustible substances 
suffer when burnt. 

A consistent though erroneous theory of combustion, which 
served to unite in one group such apparently different processes as 
the burning of a candle and the conversion of a metal into" a 
" calx " or earthy powder when it was heated in air, was first pro- 
pounded by Stahl (16601734). Stahl taught that when a 
substance burns, it loses something; this he called "phlogiston " 
(from 0Xo77To*, inflammable), which signified the common con- 
stituent of all combustible bodies. This theory, however, dates 
from before Stahl's time ; phlogiston is identical with the " terra- 
pinguis" of Becher (16351682), and the idea that combustible 
bodies lost a fiery matter, a " sulphur," is even older tha* Becher. 
The more readily a substance burns, according to Stahl, the more 
phlogiston it contains. A substance containing much phlogiston 
was carbon, or charcoal. And when a metal had lost its phlogiston 
and had become a " calx," it was possible to restore the lost 
phlogiston by heating it with charcoal, which would yield up to 
the calx its phlogiston, again converting it into the metallic state. 

But in the meantime the progress of chemistry was furthered 
by the discovery of many new gases, and the conviction spread 
not only that gases were not impure atmospheric air, but that 
matter was capable of existence in three forms, solid, liquid, and 
gaseous. Black (1728 ,1*799) was the first clearly to show 
(probably about 1755) that carbonic acid gas (carbon dioxide) or 
' fixed air " was radically distinct from ordinary air, inasmuch 
as it could combine with or be "fixed" by lime, magnesia, and 
the caustic alkalies, potash and soda. As acids have this pro- 
perty, Keir suggested that it belonged to the class of acids, and 
Berg man n (1735 1784), following Priestley's suggestion that it 
was a constituent of air, named it "aerial acid." It is the rst 
substance which was named " gas " (from geist, equivalent to gust) ; 
the name is due to Van Helmont (1577 1644), who had noticed 
that it could be obtained by heating limestone. 

The merit of Black's work consists in his having shown that, 
whereas limestone lost a definite weight by being calcined, its 
weight is exactly restored if the lime resulting from its calcination 
is reunited with carbonic acid gas. This was the first success- 
ful chemical experiment dealing with quantities. A complete 
investigation of " inflammable air," or, as it is now named, 
hydrogen gas, is due to Henry Cavendish (1731 1810). It is 
the gaseous substance produced when metals such as iron, tin, or 
zinc are treated with acids. Cavendish, in 1766, proved tho 


identity of the substance from whichever source it was prepared, 
and examined its properties. He found it to be exceedingly light, 
aod to burn very readily ; and it was supposed by some to be the 
long sought *' phlogiston " of Stahl. Cavendish also discovered 
that its product of combustion was water. 

But the chemistry of gases, or as it was then termed " pneu- 
matic chemistry," was most advanced by the researches of Joseph 
Priestley (1733 1804). He was the first to devise a convenient 
method of collecting gases over water or mercury, and his plan is 
still used in our own day. To him is due the discovery of most 
of the gaseous substances now known, especially of oxygen gas, on 
August 1st, 1774, which he named " vital air," owing to its 
property *of supporting life, or a dephlogisticated air,'' because it 
was the most ardent supporter of combustion, though not itself 
combustible. The discovery of oxygen was made independently 
and almost simultaneously by Scheele, a Swede (1742 1786), to 
whom we also owe the discovery of chlorine. 

These discoveries prepared the way for the grand generalisation 
of Lavoisier. Black, Cavendish, Priestley, and Scheele were all 
adherents of Stahl's phlogistic theory. But Lavoisier (1743 
1794), having been shown the method of preparing oxygen by 
Priestley, who paid him a visit in the autumn of 1774, saw the 
grand importance of the discovery, and made the great generalisa- 
tion that, when bodies burn, they combine with this constituent of 
air, to which he gave the name oxygen. This discovery laid the 
foundation of the present science of chemistry ; the time was now 
ripe ; and in a very complete series of researches Lavoisier showed 
first : that water cannot be converted into earth by boiling, but 
that it merely dissolves some of the constituents of the glass vessel 
in which it is boiled, leaving the dissolved matter as a residue 
after it has evaporated ; second, that when tin is heated with air 
in a closed vessel, although it is changed into a whitish-grey calx, 
yet the combined weight of the vessel and the tin remains 
unchanged ; thus showing that nothing has escaped from the tin 
or been lost from the vessel ; and that on opening the vessel air 
entered, so that the whole apparatus increased in weight ; and 
that this increase in weight was practically equal to the increase 
in weight of the tin due to its conversion into " calx." From this 
experiment he drew the correct conclusion that the gain in weight 
of the tin was due to its absorption of one of the constituents of 
air. Thirdly, he repeated this experiment, substituting the metal 
mercury for tin ; the red powder produced, when heated strongly, 
yielded up the absorbed gas, identical with, the " vital" or 


" dephlogisticated " air of Priestley, to which Lavoisier gave the 
name oxygen. Fourthly, he showed that organic matters yield, 
when burnt, carbonic acid and water; and that carbonic acid, 
identical with Black's " 6xed " air, is produced by the combustion 
of carbon or charcoal. His views are stated by himself as 
follows : 

1. Bodies burn only in pure air. 

2. This air is used up during combustion, and the gain in 
weight of the body burned is equal to the loss of weight of the air. 

3. The combustible body is generally converted into an acid by 
its union with pure air; but the metals are converted into 
calces or earthy matters. 

To this last statement is due the name " oxygen," cr " pro- 
ducer of acids." Up to that date, acid (from acetum, vinegar) was 
the name applied to substances with a sour taste, which acted 
on calces, producing crystalline substances, termed salts. Many 
attempts have since been made to give precision to the conception 
of the word acid ; but, however convenient the colloquial use of the 
word, it has ceased to have a definite chemical signification. It 
was soon after shown that bodies may possess the defined pro- 
perties of an acid and yet contain no oxygen. 

The discoveries of Lavoisier werf3 owing in great degree to two 
fundamental conceptions, with regard to which he held the firmest 
convictions : first, that heat v;-is not a substance capable of entering 
and escaping from bodies like a chemical element, but a condition 
of matter ; and that its gain or loss implied no gain or loss of 
weight ; and second, that matter was indestructible and uncreat- 
able ; and that the true measure of its quantity was its mass, or 
weight ; hence the weight of a compound body must equal the 
sum of the weights of its constituents. 

It was many years before Lavoisier's views gained complete 
acceptance amongst chemists ; but the discovery of Cavendish in 
1784 85, that the only product of the combustion of hydrogen 
was water, showed the true relations of that important substance 
to oxygen, and explained many difficulties. 

To Lavoisier, too, belongs the merit of having invented a 
systematic nomenclature, which is still retained in its main 
features ; its convenience and general applicability did much to 
promote the acceptance of the theory on which it was based. 

We have traced the gradual evolution of the science of 
chemistry from the earliest speculations of the Greek philosophers 
to the end of last century. With this century opens a new era, 
which will form the subject of the next chapter. 


Note. The chief works on the history of chemistry are Kopp's Geschichie 
der Chemie, 1843-47 ; Entwickelung der Chemie in der neveren Zeit, 1873 ; 
Thomson's History of Chemistry , 1830; Meyer's Geschichfe der Chemie, 
Leipzig, 1889 : the last is specially to be recommended. For short sketches of 
the subject, see also Muir's Heroes of Chemistry, and Picton'a The Story of 



As most of the common substances which we see around us contain 
oxygen, their composition could not be- determined before it had 
been shown by Lavoisier that the phlogistic theory was untenable, 
and before the phenomena of oxidation had received tjieir true 
explanation Lavoisier himself showed the true nature of sulphuric 
acid,* viz., that it was a compound of sulphur and oxygen, and not 
a constituent of sulphur, deprived of phlogiston ; and also of 
carbonic acid,t that it was an oxide of carbon, and not carbon 
deprived of phlogiston. These and similar discoveries of Lavoi- 
sier's pointed the way to others, and numerous attempts were 
made to discover the composition of substances, or to analyse 
them (fii>rtXt/<rf<?). And from the time of Lavoisier's enunciation 
of the true nature of combustion, to the beginning of the 19th 
century, many analyses were made, and confirmed in many cases 
also by synthesis, that is placing together (<r?;i/0e<m) the con- 
stituents of the compounds, so as to reproduce the compound 
which had been analysed. 

At that time very few accurate methods of analysis were 
known. The qualitative composition of compounds was as a rule 
not difficult to ascertain ; but the proportions in which the con- 
stituents were contained in the compounds analysed, or their 
quantitative composition, were not accurately determined, and the 
results of the same experimenter often varied among themselves. 
It is therefore not to be wondered at that two views were held 
regarding the composition of compounds : one, of which Berthollet 
(1748 1822) was the author, and which is set forth in his Essai 
de Statique Chinnque (1803) ; he regarded every compound as 
variable in composition, or, if in some cases its composition was 
found to be constant, attributed such constancy to the fact that it 
had been submitted to precisely similar conditions during its 

* The name "sulphuric acid*' used to be, but is not now, applied to the 
compound of sulphur and oxygen referred to. According to present nomencla- 
ture, the acid contains in addition the elements of water. 

f See former note. 


preparation at successive times. JBerthollet held that the propor- 
tion in which elements existed in a compound depended on the 
relative amounts of the elements present during the change which 
led to their combination, and on other conditions such as tempe- 
imure. The other and contrary view, that the same substance had 
always the same composition, was defended by Proust (1755 
1826), and the dispute, which was eagerly watched by all chemists, 
lasted from 1799 to 1808. 

But the question had already been decided by Richter 
(1762 1807). The law of "constant proportions," as it is 
termed, was announced by Richter in the involved language 
of the phlogistic theory in papers which appeared between 
1792 anfl, 1794. Stated in ordinary language, his discovery 
is as follows : If two acids, A! and A 2 , combine with two 
bases, B! and B^, to form compounds, AiB 1? A 2 B b AJl^, and 
A-.B 2 , the proportion by weight between A] and A a in the first 
two compounds is the s;nie as that between A! and A 2 in tho 
second pair if the weight of B! and also of B 2 is the same in both 
cases. Or, to take a particular case : If 80 grams of sulphuric 
acid* combine with 62 grams of soda,* or with 94 grams of 
potash ;* and if 108 grams of nitric acid* likewise combine with 
62 grams of soda ; then 108 grams of nitric acid will combine 
with 94 grams of potash. Therefore 94 grams of potash are said 
to be equivalent (or of equal value) to 62 grams of soda in their 
power of combining with acid ; and 80 grams of sulphuric acid 
are equivalent to 108 grams of nitric acid in their power of com- 
bining with base. Richter determined and tabulated a number of 
such u equivalent weights." And Proust went still further. In. 
1799 1801, he showed that tin forms two compounds with oxygen, 
in which the proportion of oxygen varies not gradually but 
suddenly ; and that iron forms two similar compounds with 
sulphur; but here he stopped. The discovery of the reason of 
definite proportions is due to Dalton ; it gave a new impetus to tho 
study of chemistry, and has been, in its results, perhaps the most 
fruitful speculation of any known to science. 

John Dalton was born in 1766, at Eaglesfield, in Cumberland. 
In his younger days he was a schoolmaster at Kendal ; he went to 
Manchester in 1793 as Lecturer on Mathematics and Natural Philo- 
sophy in the New College, and afterwards acted as a private 
mathematical and chemical tutor in Manchester, giving occasional 

* These names are used in their old sense of the combinations of the ele- 
ments sulphur, nitrogen, sodium, and potassium with oxygen. See previous 


lectures in the larger towns of England and Scotland. He inves- 
tigated the relations between the temperature and pressure of 
liquids, the expansion of gases by heat, the solubility of gases in 
liquids, and other similar subjects ; but his discoveries in chemical 
theory were those which conferred on him a world- wide fame,*and 
have exercised a lasting influence on the science. 

It was the habit of the analysts of that time, as it IP now, to 
state their results in parts per 100. Thus Proust gives the 
following analyses of the compounds of copper and tin with 
oxygen : 

" Suboxide " Protoxide " Suboxide " Protoxide 
of Copper." of Copper." of Tin." of Tin." 

Metal 80-2 BO 87 78 '4 

Oxygon 13'8 20 13 21 '6 

100 -0 100 100 100 

It is obvious that, from inspection of tho above numbers, no 
simple relation between tho amounts of oxygen in the lower and 
higher oxides of copper, and in the lower and higher oxides of tin, 
is evident ; yet, if Proust had calculated the ratios, he might have 
guessed thnt the proportion of oxygen to copper in the second 
oxide is nearly double that in the first, viz., 13'8 : 21*5; and 
similarly with tin, 13 : 24. But still the analyses are not accurate 
enough to render this proportion self-evident, even if thus stated. 

It was during an investigation of two compounds of carbon 
with hydrogen, viz., marsh gas and olefiant gas, or, as they are 
now named, methane and ethylene, and two compounds of carbon 
with oxygen, carbonic oxide and carbonic acid, or, as the latter 
gas is now called, carbonic anhydride, that Dalton was led to 
investigate the subject. He found that, if he reckoned the carbon 
in each the same, then marsh gas contains just twice as much 
hydrogen as olefianb gas ; and carbonic acid just twice as much 
oxygon as carbonic oxide. He then considered the proportions of 
hydrogen and oxygen in water, and of hydrogen and nitrogen in 
ammonia, and having found, first, that when two elements comline 
with each other, they do so in constant proportions ly weight, and 
second, that when two elements, A and B, form more than one com- 
pound with each other, they comline in simple multiple proportions, 
he deduced the following laws to account for these facts : 

1. Each element consists of precisely similar atoms of constant 

2. Chemical compounds consist of complex " atoms,"* which are 

* As the expression " complex atom " is a contradictory one, it was after- 
wards replaced by the word " molecule," or * httlo mrss" of atoms. 


produced &?/ tie union of the atoms of the constituent elements in 
simple numerical ratios.* 

An example will render these statements clear. Olefiant gas 
consists of six parts of carbon by weight united with one part of 
hydrogen; marsh gas of six parts of carbon united with two parts 
of hydrogen. Similarly, carbonic oxide contains six parts of 
carbon and eight parts of oxygen; and ''carbonic a^id," six parts 
of carbon and 16 of oxygen. The following table shows the 
relations : 

Olefiant Gas. Ratio. Marsh Gas. Ratio. 

Carbon .... 85 '71 per cent. G 75 per cent. 6 

Hydrogen. . 14 '28 1 25 2 

Carbonic Oxide. Ratio. " Carbonic Acid." Ratio. 

Carbon 42 -80 per cent. 6 27 '27 per cent 6 

Oxygen .. . 57 '14 8 72 '72 16 

It is again evident here that no obvious relation exists between 
the amounts of hydrogen in marsh gas and olefiant gas, unless 
they are compared with a uniform weight of carbon. From thpsp 
results Dalton concluded that olefiant gas consists of one atom of 
carbon united to one atom of hydrogen, and marsh gas of one 
atom of carbon united to two atoms of hydrogen ; and, similarly, 
that carbonic oxide is composed of one atom of carbon and one of 
oxygen, and carbonic acid of one atom of carbon and two of 

It necessarily follows from this conception that the atom of 
carbon is six times as heavy as the atom of hydrogen, and that the 
relative weights of the atoms of carbon and oxygen are as 6 to 8. 

Extending these observations to water, the only compound of 
hydrogen and oxygen then known, the following relation was 
determined : 

Water. Ratio. 

Hydrogen 11*11 per cent. 1 

Oxygen 88 '88 8 

Hence Dalton concluded that water is a compound of one atom 
of hydrogen with one atom of oxygen, and that the atom of oxygen 
is eight times as heavy as the atom of hydrogen, thus bearing out 
the conclusions of his former analyses. 

Dalton then proceeded to determine the relative weights of the 
atoms of other elements by similar methods. His numbers are far 

* Dalton's New System of Chemical Philosophy, 1808; Thomson's Chemutry 
1807 j also edition 1810, Vol. Ill, p. 441. 



from accurate, and indeed, in the above tables, the actual numbers 
found by him have not been stated, in order to avoid confusion. 
He next arranged a number of compounds of the elements in 
classes, according to the number of atoms contained in each class. 
Thus if only one compound of two elements was known, Dalton 
assumed it to contain one atom of each element, and named it a 
binary compound, " unless some cause appear to the contrary." 
If two compounds were known, they were represented as A + B, 
and as A -f- 2B ; the latter was named a ternary compound, 
because it contained three atoms ; and so on with quaternary, &c. 
Thus he regarded water as a binary compound, in which one atom 
of hydrogen weighing 1, and one atom of oxygen weighing 8 rela- 
tively to the hydrogen were united. Ammonia, a compound of 
nitrogen and hydrogen, was regarded as also composed of one atom 
of hydrogen weighing 1 and one atom of nitrogen weighing 4|. 
Thus he constructed a table of atomic weights ; and to render his 
theory more tangible, he assigned to each element a symbol ; thus 
oxygen was Q? hydrogen 0, nitrogen , sulphur , and so on; 
and the symbols of the metals consisted of circles circumscribed 
round the initial letter of the name of the metal ; thus stood 
for iron, for zinc, and so on. These symbols also stood for the 
relative weights of the atoms ; hence Q0 denoted water, QCD am - 
mooia, QQ olefiant gas, O0O marsh gas, and so with others. 

Now it is evident that Dalton here made a great assumption, in- 
asmuch as he had no sure basis to guide him in assigning such 
atomic weights. Let us consider his results from another point of 
view, and we shall see that another set of atomic weights might 
with equal justice have been adopted. 

Turning back to the table on p. 17, it is seen that Dalton 
assumed that the four substances, marsh gas, olefiant gas, carbonic 
oxide, and " carbonic acid " each contained one atom of carbon,. But 
it is equally justifiable to assume that each of the first pair contains 
one atom of hydrogen, and each of the second pair one atom of 
oxygen. We should then have the ratio : 

Olefiant Gas. Ratio. Marsh Gas. Ratio. 

Carbon 85 71 per cent. 6 75'0 per cent. 3 

Hydrogen... 14'28 1 25'0 H 1 

Carbonic Oxide. Ratio. " Carbonic Acid." Ratio. 

Carbon 42 86 per cent. 6 27 '27 per cent. 3 

Oxygen. .. 57'14 8 7272 8 

The smallest amount of carbon in combination is now found 


to weigh three times as much as the hydrogen ; i.e. the atomic 
weight of carbon is 3. And the first body would then consist of 
2 atoms of carbon and 1 of hydrogen ; while the second, marsh 
gas, would contain 1 atom of carbon and 1 of hydrogen. Similarly, 
carbonic oxide might be composed of 2 atoms of carbon and 1 of 
oxygen, while " carbonic acid " might consist of 1 atom of carbon 
and 1 of oxygen. 

Dalton himself was quite aware of this difficulty, as is seen by 
his remarks in the appendix to his second volume, published in 
1827. He therefore contented himself by assuming those numbers 
to be the correct atomic weights which give the simplest propor- 
tions between the numbers of atoms contained in all the known 
compounds \>f the elements. But Dalton did not possess the 
analytical skill necessary to determine the composition of the 
compounds from which such deductions were to be made. In 1808, 
Wollaston published an account of accurate experiments on the 
carbonates and oxalates of sodium and potassium, in which he 
showed that the ratio of carbonic acid or oxalic acid in one (the 
" subcarboiiate " or " suboxnlate ") to the sodium or potassium was 
half that which it bore in the other (the " supercarbonate " or 
" superoxalate "). The work of determining the composition of 
compounds was, however, chiefly undertaken by Berzelius, pro- 
fessor of chemistry, medicine, and pharmacy in Stockholm 
(1779 1848). The aim of this great chemist was to forward the 
work which had been suggested by Dalton, and, by preparing 
numerous compounds and analysing them, to determine the ratios 
of the weights of their atoms. His industry was untiring, and the 
number of new compounds prepared and analysed by him almost 
incredible. But it is obvious that for the reasons stated it is impos- 
sjble, even by comparing all the compounds which one element forms 
with others, to determine which compound contains only 1 atom of 
that element. What Dalton and Berzelius really determined was 
the equivalents of the elements, that is, the proportion by weight 
in which they are capable ot combining with or replacing 1 
part by weight of hydrogen ; they had no data sufficient to enable 
them to determine what multiple of the equivalent is the true 
atomic weight. In subsequent chapters the various reasons in' 
favour of the atomic weights at present assigned to the elements 
will be discussed. We must leave the historical part of the 
subject at this point, and proceed to discuss the facts of the science, 
and to arrange the various compounds in an orderly manner. 

Assuming, then, that, for reasons to be given hereafter, the 
relative weights of the atoms are represented by the^ numbers used 

c 2 


in this- book, the question arises, what element should be made 
the standard of comparison ? Dalton having found that, of all the 
elements investigated by him, a smaller weight of hydrogen 
entered into combination than of any other element, assigned the 
weight 1 to the atom of that element, and arranged the other 
atomic weights accordingly. Thus, according to him, the weight 
of an atom of oxygen was 8 times that o an atom of hydrogen, 
because water, which he supposed to consist of 1 atom of each, 
was found on analysis to contain 1 part by weight of hydrogen 
combined with 8 parts by weight of oxygen. And so with the 
other elements. There are reasons which will follow in their place 
(p. 202) for believing that a number between 15 87 and 16*00 (01 
double the number assigned by Dalton) represents Ihe relative 
weight of an atom of oxygen referred to hydrogen as unity. Bui 
it happens that the equivalents of most of the elements have 
been determined by synthesising or analysing their compounds 
with oxygen, or with oxygen and some other element. Hence il 
appears advisable to accept the atomic weight of oxygen as 16, 
and to refer the weights of the other elements to that scale 
Until the ratio between tbe atomic weights of hydrogen and 
oxygen is satisfactorily determined, this appears the best course tc 
pursue ; for then the accepted atomic weights of the majority oi 
the elements need not 'be altered to suit any proposed alteration in 
tho ratio of the accepted itomic Weights of hydrogen and oxygen 
Moreover this plan has the great advantage that many of the 
atomic weights are whole numbers, and are therefore more easily 
remembered. It should here be noticed that if the ratio between the 
atomic weights of hydrogen and oxygen is really 1 to 15 96, ther 
by placing the atomic weight of oxygen equal to 16, that o 
hydrogen is no longer 1, but 1-0025, for 15*96 : 16 :: 1 : 1*0025 
A very remarkable relation between the atomic weights f the 
elements and their chemical and physical properties was pointec 
out by Mr. J. A. R. Newlands in 1864,* and this relation has beei 
further studied by Professors Mendel^efff and Lothar Meyer*] 
It is briefly this. If the elements be arranged in the order of thei] 
atomic weights in seven double columns, those elements whicl: 
resemble each other fall in the same column. It is on this principlt 
that the elements and their compounds are classified in this text 
book. Such an arrangement is termed a periodic arrangement 

Chem. News, July 30th, 1864; August, 1865 ; March, 1866 j also On th 
Discovery of the Periodic Law, Spon, 1884. 
t Annalen, Suppl., 8, 133 (1869). 
J Annalen, Sispl., 7, 354. 


and the following table is named the periodic table. The letters, 
such as H, Li, <fcc., are abbreviations for the names of the elements ; 
they are termed symbols ; and they also represent the numbers 
which precede or follow them. Thus represents not merely 
oxygfcn, but 16 parts by weight of oxygen ; CaO represents not 
merely a compound of calcium and oxygen, but of 40 08 parts by 
weight of calcium, and 16 parts by weight of oxygon ; CaCU 
represents a compound of 40 parts by weight of calcium with 
2 x 35*46 parts by weight of chlorine. Such a representation of 
compounds by the symbols of the elements which they contain is 
termed a formula. 

While most of the elements are represented by the initial letters 
of their English names, some of the symbols require explanation. 
The following is a list : 

Na, Natrium (connected with the word nitre) Sodium. 

K, Kahum (from alkali, an Arabic name) . Potassium. 

Cu, Cuprum (Latin) Copper. 

Ag, Argentum (Latin) -Silver. 

Au, Aurum (Latin) Gold. 

Hg, Hydrargyrum (Greek = water-silver) Mercury. 

Sn, Stannum (Latin) Tin. 

Ph, Plumbum (Latin) . Lead. 

Sb, Stibium (Latin) Antimony. 

W, Wolfram, a mineral containing Tungsten . . . Tungsten 

Fe, Ferrum (Latin) Iron. 

Note. For this portion of chemical history, Wurtz's History of the Atomic 
Theory, London, 1880, may be consulted ; also Cook's The New Chemistry ; 
and the works previously referred to. 







OO fa 


% 8 is 



s w 


O lO" rH 




< H csf 

fl S 



rH W cv 



s w 





C5 o 10 <~) Q j-^ o 

rH CO lO QO O CM lO 

S S 

rH rH 




^ ' 

(Z4 r3 cv. cv. 



00 ia H 



p o 



CO 01 CH ib CO 

rH rH 

rH rH 






o o a s? 



* -? S 





"^ j rH \O -^ O rH 
rH CO vO Jt> OS CM Tf 

rH rH 

CO 00 
rH rH 




fe k rS * 



2 s 

^ 5 fi 



V 5 10 10 

i4 _ 




s s 5 g g s 




rH rH 

rH rH 



O H N O 



^1 l^J rH 




rH l> -f O OS ^f CM 






rH CM ^ l> 00 rH <* 

i 1 rH 

rH rH 




AA W ? ^ 

W O2 r*1 r^ 




Jsf N S 


kO lO lO 




Oi ** O W3 JO CM Jt> 

CM ^* CO 00 rH CO 
rH rH 

00 CM 

rH rH 





<C gj u. * 




^ ^ ^ 


* - 



Jt* ^ o> cb o oo co 





"*"" vi ^ 

rH rH 




a M S 3 



H S 

'S - 

03 rj 


8 3 


cv. 1 



Table of Atomic Weights of Elements (0 = 16.). 


, Al. 



... Ni 

58 6 


. Sb 



. .. Nb 



. As 


, .. N 



. Ba 

137 -00 


... OB 

191 -3 


. Be 



... O 



. Bi 

208 -10 

Palladium. . . . 

.. . Pd 

106 -35 

. B 


Phosphorus.. . 

.. . P 

31 03 


. Br 


. .. Pt 



. Cd 



39 -It 


. Cs 



. . . Prd 



. Ca 

40 -08 


. .. Bh 

103 () 


. C 


Rubidium .... 

... Bb 

85 5 

Cerium. . , 

. Ce 


Ruthenium . . . 

. .. Bu 

101 -65 


. Cl 


... So 



. Cr 


. .. So 

79 -0 


. Co 



. .. Si 

28 -33 


. Cu 



... Ag 

107 -930 


. Er 



. .. Na 



. F 


Strontium.. . . 




. Ga 

69 9 


.. . S 

32 06 


. oe 


Tantalum . . . . 


182 5 


. Au. 

197 '22 


. . Te 

125 ? 


. H 

1 to 1 0082 Terbium 

. . Tb 

162 ? 

Indium , 

,. In 

113 7 


. . Tl 

204 -2 


. I 

126 85 


. .. Th 



,. Ir 



.. . Sn 



. Fe 


Titanium.. . . 

.. Ti 


Lanthanum . . . 

. La 



. . W 



. Pb 



.... 17 





Vanadium . . . 

.... V 

51 '4 

Magnesium .... 

.. MfiT 


Ytterbium. . . 

. . . . Yb 



.. Mu 



,.,. Y 






. ,. Zn 

65 -3 

Molybdenum. . 

. Mo 





Neojlymium . . . 

.. Ndi 



Liquid, Gas. 

2fote. In this table recent determinations have been incorporated with the 
mean results given by Clarke (" Constants of Nature," Part V, 1882). It is to 
he understood that the last digit of the figures given may vary within one or 
two units. Thus zirconium = 90 means that the atomic weight is not certain, 
and may be 89'5 or 90'5 ; thallium = 204 2 leaves it uncertain whether the 
true weight is 204'1 or 204*3 ; and so on. Where a query (?) is appended, it 
is to be understood that the weight given may be one or more units wron#. 
The standard works on the subject are by Clarke, mentioned above ; by Lothar 
Meyer and Seubert, Die Atomyewichte der Elemente ; arid, as a mode] of 
research, by Staa, Recherches sw les Rapports rectproyues de9 Poids alvmiyues, 
Brussels, 1860. 

Table of Metric Weights and Measures 

Measures of Length. 

I metre = 10 decimetres = 100 centimetres = 1000 millimetres. 

1 metre = 1 '09363 yard == 3 '28090 feet = 39 '37079 inches. 

Log metres + '038870 1 = log yards; + '5159930 = log feet, 4 + 1 '5951743 
= log inches. 

Log n yards + '9611296 = log metres; log n feet + 0'484007l = log deci- 
metre; log n inches + '40 18257 ~ log centimetres. 

Measures of Capacity. 

\ cubic metre = 1000 litres = 1,000,000 cubic centimetres = 1,000,000,000 cubic 

1 litre = (U'02705 cubic inches = '035317 cubic foot = 1 '76077 pints 

0-22097 gallon, 
log / litres + 1' 7855223 = log cubic inches; + 2 -5479838 = log cubic feet; 

+ 0'2457026 = logpinta; + 1-3443333 = log gallons. 
Log / cubic inches -f- 1 '2144774 log cubic centimetres. 
Log n cubic feet 4- 1*4520162 = log litres. 
Log n gallons + '6556667 log litres. 

Measures of Weight \ 
1 gram ~ weight of 1 cubic centimetre of water at 4 . 

1 kilogram * 1000 grams = 100,000 centigrams = 1,000,000 milligrams. 

1 kilogram = 2 -2046213 Ibs. ; = 35' 273941 oz. - 15432-35 grains. 

].og n kilos. + '3433340 log Ibs.j log n grams -I- 1' 5474540 log oz.; 

+ 1 '1884323 log grains. 
Log n Ibs. + 1*6566660 = log kilograms; log n grains + 2 '8115677 loggranw. 





THE elements, it has been seen, when arranged in the order of 
their atomic weights, fall into certain groups. The various 
members of these groups resemble each other in their physical and 
chemical properties, and it is therefore advisable to consider the 
members of each group in connection with each other. They 
possess certain properties in common, while exhibiting individual 
peculiarities. In the following chapters, an account will be given 
of the sources of the elements, whether they occur "free," or 
"native," that is, as elements, or whether combined with other 
elements in the form of compounds ; of their properties ; and of 
the methods of their preparation. ; but fuller details will in some 
cases be given under the heading of the compounds from which 
they are prepared. 

In the main, the order of the periodic table will be followed ; 
butj as it is still under investigation, and the position of all the 
elements cannot be regarded as finally settled, certain elements 
will be grouped together which do not occur near each other in the 

GROUP I. Hydrogen, Lithium, Sodium, Potassium, 
Rubidium, Caesium. 

Sources. Hydrogen occurs free in the neighbourhood of 
volcanoes, owing probably to the decomposition of its compound 
with sulphur, hydrogen sulphide, by the hot lava through which 
it issues. It has also been proved by the evidence of the spectro- 



scope (see chap. XXXV), to exist as element in the atmosphere of 
the sun, in certain fixed stars, in nebulae, and in comets. It has 
been found associated with iron and nickel in many meteorites. t ln 
combination with | oxygen, it occurs in water (hence its name 
from vdtap, water, *fewaw, I produce) in the sea, lakes, riverS, in 
the atmosphere, in many minerals ; in all organised matter, animal 
and vegetable. Tt is thus one of the most widely distributed and 
abundant of elements. 

Preparation. 1. By heating its compounds with boron, 
carbon, silicon, nitrogen, phosphorus, arsenic, antimony, sulphur, 
selenium, tellurium, iodine, or palladium to a red heat ; or with 
oxygen, chlorine, or bromine to a white heat (see these com- 

2. By the decomposition of its compounds dissolved in water 
by an electric current (see p. 62). 

3. By displacing it from these compounds by means of certain 
metals. The most usual methods of preparation are by the action 
(a) of sodium on water (oxide of hydrogen, see p. 192) ; (6) of 
iron on gaseous water at a red heat (see p. 255) ; or (c) of zinc on 
dilute sulphuric or hydrochloric acid (see pp. 415, 112). 

Method (a). Ajar is filled with water, covered with a glass plate, and in- 
verted in a trough of water as shown in figure 1. A piece of the metal sodium 

FTO. 1. 

not larger than a pea is placed in a spoon made of wire gauze, which is passed 
quickly under the water beneath the jar, when the hydrogen evolved passes in 
bubbles into the jar. The sodium melts, moves about, and displaces hydrogen 
from the water. Other small fragments are successively introduced into the 
spoon until the jar is full. (Note. Large pieces must not be used, else an ex- 
plosion may ensue.) 

Method (b). A piece of iron gas-pipe, of |-inch bore, is filled loosely with 
iron turnings, and closed by stoppers made of asbestos cardboard moistened 


with water, and moulded round glass tubes, placed as shown in figure 2. The 
iron tube is then heated in a gas furnace, and the water in the flask is boiled. 
The iron combines with the oxygen of the steam, setting free the hydrogen, 
which may be collected in a jar as shown in the figure. 

FIG. 2. 

Method (c). A flask or bottle, as shown in figure 3, is provided with a cork 
and delivery tube. Some granulated zinc, prepared by pouring melted zinc into 
water, is placed in the flask A ; and a mixture of one volume of hydrochloric acid 
and four volumes of water, or of one volume of oil of vitriol (sulphuric acid),* 
and eight volumes of water, is poured through the funnel B. Bubbles begin to 
appear on the surface of the zinc, and the liquid effervesces. A few minutes 
must be allowed, so that the hydrogen may displace the air from the bottle. It 
can then be collected in jars. The zinc displaces the hydrogen from its com- 

FIG. 3. 

pound with chlorine in hydrochloric acid, or from its compound with 
sulphur and oxygen in sulphuric acid. The substances produced are named zinc 
chloride, or zinc sulphate, according as one or other acid has been used. 

* If sulphuric acid be used, sulphur dioxide and hydrogen sulphide are 
produced if the proportion of water be not a large one, Chem. Soc., 53, 54. 


Properties. A colourless, odourless gas ; the lightest of all 
known bodies. As it is nearly fourteen and a half times as light 
as air, it may be poared upwards from one jar into another ; on if 
a light jar or beaker be suspended mouth downwards from the 
arm of a balance, and counterpoised, and hydrogen be poured into 
it from below, that arm of the balance rises, the heavier air being 
replaced by the lighter hydrogen. Balloons used to befilled with 
it ; but coal gas is now employed. It burns in airj^bmbining 
with oxygen to form water, and when mixed with air (about 
2 times its volume) the resulting mixture is explosive (see p. 
192). It is sparingly soluble in water; 100 volumes of water 
absorb 1*93 volumes of hydrogen gas. It is not poisonous, but 
cannot be respired for any long time, as ttie oxygen of the air, 
which is necessary for the support of life, is thereby excluded. 
Owing to the rate at which it conveys sound, speaking with 
hydrogen gives a curious shrill tone to the voice. It has never 
been condensed to the liquid or solid states. Cailletet, and also 
Pictet, who claim to have condensed it by cooling it to a very low 
temperature,* and at the same time strongly compressing it, had 
m their hands impure gas. Its critical temperature, above which 
it cannot appear as liquid, is probably not above 230. 

It unites directly with the halogens ; with oxygen and with 
sulphur ; also with carbon at a very high temperature j and with 
potassium and sodium. It is absorbed by certain metals, notably 
by palladium, which can be made to take up 900 times its own 
volume (see p. 576). From this its density and its specific heat in 
the solid state have been calculated. f 

Lithium, sodium, potassium, rubidium, and caesium are 
always found in combination with chlorine, or with oxygen and 
the oxides of other elements such as silicon, carbon, boron, sulphur,, 
phosphorus, &c. ; they never occur free. They are named " m/stals 
of the alkalies." 

Sources. Lithium occurs as silicate in lepidolite and petallite ; 
as phosphate in triphylline; as chloride in many mineral waters, 
especially in the Wheal-Clifford Spring, near Eedruth, in Cornwall ; 
in sea- water ; and in many soils, whence it is absorbed by plants, 
tobacco-ash, for example, containing about 0*4 per cent. Its com- 
pounds are usually prepared from lepidolite. 

Sodium forms, in combination with chlorine, common salt, or 
sodium chloride ; it occurs in deposits in Chili and Peru as nitrate 
and iodate, in which its oxide is combined with the oxides of nitrogen 
* Comptes rend., 98, 304. 
f Ibid., 78, 968 ; also Phil. Mag. (4), 47, 324. See Palladium. 


and iodine respectively ; as sulphate in mineral springs ( Glauber's 
salts) ; as silicate in soda-felspar or allite ; and as fluoride along 
with, aluminium fluoride in cryolite; as borate in certain American 
lakes. It is obtained as carbonate by incinerating sea-plants. 

Potassium, is found as chloride in mineral deposits at Stass- 
furth, in N. Germany; the mineral is termed sylvin; as nitrate 
(saltpetre, nitre), forming an incrustation on the soil in countries 
where rain seldom falls ; and as silicate in many rocks, chiefly in 
potash felspar and mica. It is abundant and very widely dis- 
tributed, being a constituent of every soil. It remains as car- 
bonate on burning to ash all kinds of wood, hence its name, 
from " pot-ash." 

Rubidium and Caesium are widely distributed, but occur in 
small amount. They are contained in lepidolite, along with lithium 
and potassium, as silicates ; also in castor and pollux, two rare 
minerals, found in the Isle of Elba. They also occur in some 
mineral waters, particularly in a spring at Diirkheim, in the 
Bavarian Palatinate, from which they were first extracted by 
Bunsen, their discoverer, in 1860. They are widely distributed in 
the soil, and are absorbed by some plants to a considerable extent. 
Thus the ash of beetroot contains 1*75 per mille of rubidium. 

Preparation. These metals are prepared : 1. By passing a 
current of electricity through their fused hydroxides, chlorides, 
or cyanides. It was in this way that Davy,* in 1807, obtained 
potassium and sodium from their hydroxides, which up to that 
date had not been decomposed ; the electrolysis of lithium chloride 
is still the only method of preparing lithium ; and Setterberg,f in 
1881, prepared considerable quantities of rubidium and ceesium by 
electrolysing a fused mixture of their cyanides with cyanide of 
barium, using as poles strips of aluminium. 

Tc prepare lithium, which may serve as a type of this kind of operation, 
about 30 grams of lithium chloride are melted over a Bunsen flame in a nickel 
crucible ; when the chloride is quite fused, a piece of gas carbon (the sticks of 
a Jablochkoff candle answer well), is connected with the positive pole of four or 
six Bunsen or Grove cells ; and a knitting-needle, passing through the hole in the 
stem of a tobacco pipe, made into a shallow cup at its broken end, is connected 
with the negative pole ; these are dipped in the fused chloride ; and when a bead 
of lithium as large as a small pea has collected on the negative electrode, the fused 
chloride is allowed to cool, and the bead plunged into rock oil. The bead of 
lithium is then scraped off with a knife, and the process repeated, until a suffi- 
cient quantity has been collected. 

* Phil. Trans., 1808, 1 ; 1809, 39 ; 1810, 16. 
t Annalen, 211, 100. 


2. Sodium, potassium, and rubidium may be prepared by dis- 
tilling the hydroxides with carbon.* The carbon unites with the 
oxygen of the hydroxide, while the hydrogen ,is liberated and 
comes off as gas. (A hydroxide, it should be here explained, is a 
compound of oxygen with hydrogen and with a metal.) "The 
industrial preparation of sodium is thus carried out (see Chapter 
XXXVIII, p. 651). 

Properties. These elements are all white metals, so soft at 
the ordinary temperature that they can be cut with a knife, but 
brittle at low temperatures ; they are malleable, and may be 
squeezed into wire by forcing them through a small hole by means 
of a screw-press ; they may be welded by pressing clean surfaces 
together ; they melt at moderate temperatures, and ai*6 all com- 
paratively volatile; hence, lithium excepted, they may be distilled at 
a bright red heat from a malleable iron tube or retort. They are 
all lighter than water; lithium, indeed, is the lightest solid known. 
Each imparts its special colour to a Bunsen or spirit flame ; thus 
compounds of lithium give a splendid crimson light ; of sodium 
a yellow light; potassium compounds colour the flame violet; 
rubidium red, hence the name of the metal (from rubidus) ; and 
coesium blue (ccesius). (See Spectrum Analysis, Chapter XXXV.) 
Potassium vapour is green, and sodium vapour, violet. These ele- 
ments crystallise, when m,elted and cooled, in the dimetric system. 

They all combine readily with the elements chlorine, bromine, 
iodine (these elements are termed the "halogens"), oxygen, 
sulphur, phosphorus, &c., with evolution of light and heat; and 
they all decompose water at the ordinary temperature, liberating 
hydrogen (see p. 26). 

Physical Properties. 
Mass of 1 cub. cent. 

Lithium . . 
Sodium. . . . 

Solid. Liquid. 

.. 0-62 0'63f 0-025J 
0-59 ? 
0-985 ? 


TT 1 


Below -230 


0-865 ? 




1-50 ? 



Caesium . . . 

1-88 ? 



26 27 

* Castner, Chem. News, 54, 218. 

f Deduced from the mass of 1 c.c. of its alloy with palladium. 
J At 0, under a pressure of 275 atmospheres j deduced from the density of 
a mixture of 1 volume of hydrogen with 8 rols. of carbonic anhydride. 


Boiling- Specific Atomic Molecular 

point. Hoat. Weight. Weight. 

Hydrogen Below -230 (Gas) 2 -411 1 '0025 ? 2 '0 

(Solich 5 88 

Lithium ? 941 7 '02 

Sodium 742 '293 23 04 23 '04 

Potassium 667 0'1G6 89 '14 39 '14 

Rubidium ? ? 85 '5 

Caesium ? ? 132 '9 

Giiour II. Beryllium or Glucinum, Calcium, 
Strontium, Barium. 

These metals, like those of the previous group, always occur in 
nature in -combination, never in the metallic state. They are 
found combined with silicon and oxygen, as silicates ; with carbon 
and oxygen, as carbonates ; with sulphur and oxygen, as sul- 
phates ; and with phosphorus and oxygen, as phosphates. Calcium 
is also associated with fluorine and with chlorjne. They are named 
" metals of the alkaline earths.'* 

Sources. Beryllium is a somewhat rare element. Its most 
common sources are : beryl, a silicate of beryllium and aluminium, 
a pale greenish- white mineral, which, when transparent, and of a 
pale sea-green colour, is natned aquamarine ; and when bright 
green, emerald (the green colour is due to the element chromium) ; 
phenacite, also a silicate of beryllium ; and chry so beryl, a compound 
of the oxides of beryllium and aluminium. 

Calcium is one of the most abundant elements. Its carbonate 
when pure and crystalline is named Iceland-spar or calc-spar ; 
earthy and less pure varieties are limestone, chalk, and marble. When 
associated with magnesium carbonate, the mineral is named dolo- 
mite. Calcium sulphate is named gypsum, selenite, and anhydrite, 
according to its state of aggregation. Its phosphate, in which it 
is Combined with phosphorus and oxygen, is named phosphorite or 
apatite. The fluoride is named fluor- or Derby shire- spar ; and its 
chloride is a constituent of sea- water and many mineral waters. 
Most natural water contains hydrogen calcium carbonate (bi- 
carbonate) in solution. 

Strontium, like calcium, occurs as carbonate, in strontianite, 
and as sulphate in celestine. Its name recalls the source in 
which it was first found Strontian, a village of Argyllshire, in 

Barium also occurs as carbonate, witherite; and as sulphate, 
barytes or heavy -spar, so named from its high specific gravity. 
Henc'e the name of the metal, from /k/>v, heavy. 


Preparation. Beryllium, the chloride of which volatilises 
at a red heat, may for that reason be prepared from that compound 
by passing its vapour over fused sodium contained in an iron boaj;.* 
The sodium combines with the chlorine, which leaves the metal as 
such. Sodium reacts with cold water, while beryllium does not ; 
hence the sodium may be removed by treatment with water. 

Barium, strontium, and calciumf are best prepared by passing 
a current of electricity through their respective chlorides, fu^ed in 
a porcelain crucible over a blowpipe, using a carbon rod (see 
lithium) as one electrode, and an iron wire as the other. Solutions 
of the metals in mercury are -easily made by electrolysing strong 
solutions of the chlorides of the metals, using mercury as the 
negative electrode. Barium amalgam crystallises out of the 
mercury; it may be collected, and after washing it with .cold 
water and drying it, the mercury can be distilled off in a vacuum, 
leaving the barium as a yellowish-white metallic powder, still, 
however, containing mercury. Another method of preparing an 
amalgam of mercury and barium (alloys of mercury are termed 
" amalgams ") is to shake up sodium amalgam with a strong 
solution of barium chloride. The sodium combines with the 
chlorine, leaving the barium in the mercury. Amalgams of 
strontium and calcium cannot be made in this manner. 

Properties. Beryllium and calcium are white metals ; the 
other two have a yellovv^tinge. They melt at a bright red heat, 
oxidising in presence of air. Calcium and beryllium are brittle ; 
strontium and barium malleable. They are all heavier than water. 
The compounds of the last three impart characteristic colours to a 
Bunsen flame, and have well-marked spectra (see Chapter XXXY). 
The chloride of calcium tinges the flame brick-red ; of strontium, 
bright crimson-red like lithium ; and of barium, pale-green. The 
metals have not been volatilised. They unite readily with the 
halogens, with oxygen and sulphur, and with phosphorus. 
Beryllium does not decompose water unless boiled with it ; the 
others act on it at the ordinary temperature, with evolution of 

Physical Properties. 

Mass of 1 cub. Melting- Specific Atomic 

cent, solid. point. Heat. Weight. 

Beryllium. ... 1 '85 at 20 Eed heat Variable 9 *1 

(above 1230). (See Appendix) . 

Strontium .... 


Bright red 

? 87-5 
? 137 '00 

* Chem. Ifiws, 42, 297. f Annalen, 183, 367. 


Appendix. The specific heat of beryllium varies greatly with the tempera- 
ture. The following results were found by Ilumpidge.* 

Temperature .. ' 100 200 306 400 500 

Specific Heat .. 0-375G 0-4702 0*5420 0-5910 0'6172 0-6206 

GROUP III. Magnesium, Zinc, Cadmium. 

Sources. These three metals are never found native. All 
three occur as carbonate and as silicate; and the two latter as 
sulphide. Magnesium sulphide is decomposed by water ; hence 
its non-occurrence in nature. Magnesium occurs also in consider- 
able quantity as sulphate {Epsom salts), in sea- water, and in many 
mineral springs. Its native compounds are named as follows : 
Magnesium carbonate, magnesite ; double carbonate of magnesium 
and calcium, dolomite ; it occurs in great rock masses in the range 
of hills in the Italian Tyrol named the Dolomites. There are 
many silicates of magnesium and other metals. Among the more 
important are talc, steatite or soap-stone (French chalk) , serpentine, 
and meerschaum. Augite, hornblende, asbestos, olivine, and biotite 
(a variety of mica) are also rich in magnesium (see Silicates, p. 313), 
Carnallite, a chloride of magnesium and potassium, is found at 
Stassfurth. The commercial sources of metals are named "ores.'* 
The ores of zinc are : Calamine, zinc carbonate ; silicious cala- 
mine, the silicate ; and blende, or " Black Jack," the sulphide. 
Cadmium always accompanies zinc ; the only pure mineral con- 
taining it is greenocJcite, cadmium sulphide. The name magnesium 
is derived from the town of Magnesia, in Asia Minor. Its oxide 
is sometimes called magnesia alba, from its white colour. The 
word zinc is perhaps connected with the German equivalent 
for tin, Zinn. " Cadmium " is adopted from the name given by 
Pliny to the sublimate found in brass-founders' furnaces (cadmia 
fornaeum) . 

Preparation. Magnesium is prepared like beryllium ; dried 
carnallite, a double chloride of magnesium and potassium com- 
bined with water, is heated with sodium. The sodium unites 
with the chlorine, removing it from the magnesium, which is set 

The mixture is heated in large iron crucibles to a high temperature. When 
the reaction is over, the crucible is allowed to cool, and the contents chiselled 
out. Small globules of magnesium are disseminated throughout the fused 
mass, and at the bottom of the crucible is a mass of magnesium embedded in 

* Proc. Roy. Soc., 39, 1. 



flux, as the fused chlorides are termed. The salt with the globules of mag- 
nesium is transferred to a crucible, A, the bottom of which is perforated, as shown 
in the figure, and a tube, B, passes through the bottom, reaching up to near tbe 
top of the crucible. The lid is then luted on (i.e., fastened on by clay), the 

top of the tube having been closed by a wooden plug. When the temperature 
rises to bright redness, the magnesium rises in vapour, and distils down the 
centre tube, condensing on the lower portion, whence it drops into heavy oil. 
Hence the old term for tlr^ process " distillatio per dewensum " 

Zinc is produced by distilling its oxide with coke (carbon) in 
clay cylinders. The carbon unites with the oxygen, setting free 
the zinc, which distils over. 

The old English method of extracting zinc from its oxide used to be carried 
out in apparatiis like that employed in making magnesium. The roasted zinc 
ore, consisting of oxide of zinc, was mixed with coke or anthracite coal (carbon), 
and placed in clay crucibles, similar in construction to the iron one sViown in 
Fig. 4. On raising the temperature to bright redness, the zinc distils over, and 
drops through the tube which passes through the bottom of the furnace. 

The Belgian process, which is now all but universally adopted, consists in 
distilling the zinc ore with coke from clay cylinders, arranged in tiers. The 
zinc condenses in conical tubes of cast iron or iron plate, which fit the mouths 
of the cylinders, and are made tight at the joint by a luting of clay. When 
the operation is over, these tubes are removed, and the zinc, which forms a 
crust adhering to their interiors, is chiselled off. 

Cadmium accompanies zinc, and as it boils at a lower tempe- 
rature, the first portions wliich distil over contain it. 

Properties. These three metals are all white. Zinc, however, 
has a bluish tinge, and cadmium a yellow tinge. Of the three, 

BORON. 35 

magnesium is the hardest, and cadmium the softest ; it may be cut 
with a knife, but with difficulty. Magnesium and zinc are malle- 
able and ductile at a moderately high temperature (zinc at 120), 
but are brittle at the ordinary temperature. Zinc is also brittle at 
200* and may be easily powdered in a hot iron mortar. Those 
metals may all be distilled, cadmium most easily, and magnesium 
at the highest temperature. They are all heavier than water. 
They combine directly with the halogens ; they burn when heated 
in air, combining with its oxygen. Magnesium gives out a brilliant 
white light, and it is prepared in the form of ribbon, wire, or dust 
for signalling, pyrotechnic, and photograph ical purposes. Zinc 
burns with a light blue -green flame, and cadmium with a dull 
flame; they* tarnish very slowly in air. They also unite directly 
with sulphur, phosphorus, &c. When boiled with water, mag- 
nesium and zinc slowly decompose it, hydrogen being evolved. 
Cadmium is without action on water except at a red heat. 

Physical Properties. 

Muss of 1 c c Density, Melting- 
Solid. H = 1. point. 

Magnesium 1-743 ? 700800 

Zinc 7'15 34-5 412 

Cadmium 8 -6 52'15 315 

Boiling-point Specific Atomic Molecular 

at 760 mm. Heat. Weight. Weight. 

Magnesium About 1000 '250 24 '30 24 '30 

Zmc 930 to 942 '095 65 '43 65 '43 

Cadmium About 770 -056 112 '1 112 '1 

OKOUP IV. Boron, Scandium, Yttrium,* Lantha- 
num,* Ytterbium.* 

These elements are never found in the free state. They all 
exist in nature in combination with oxygen, and other oxides. 

Sources. Boron issues from the earth as hydroxide, or 
boracic acid, along with steam in the neighbourhood of vol- 
canoes. The hydroxide also occurs as sassolite; its other sources 
are tincal or native borax, in which its oxide is combined with 
oxide of sodium and with water (the beds of certain dried up 
American lakes contain enormous quantities of borax) ; boracite, 
boron oxide with magnesium oxide and chloride ; boronatrocalcite, 

* It is doubtful if these metals belong to this group. 

* D 2 


boron oxide, calcium oxide, and sodium, oxide ; and datolite, boron 
and silicon oxides with calcium oxide. 

The remaining elements of this group are usually associated 
with cerium, didymium, erbium, terbium, samarium, &c., as 
oxides, in combination with oxides of silicon, niobium, tantalum, 
titanium, and other elements. The minerals containing them 
are named euxenite, orthite, columbite, gadolinite, yttrotantalite, 
samarsltite, and cerite. They have been found chiefly afc Areiidal 
and Hittero, in Norway, and in Connecticut, U.S. 

Preparation. Boron is obtained by heating with metallic 
sodium the compound which its fluoride forms with potassium 
fluoride ; or by heating its oxide with potassium, or better, with 
magnesium dust. The fluorine or oxygen combines with the potas- 
sium or magnesium, leaving the boron in the free state. It was 
by the latter method that it was first prepared in 1808 by Gay- 
Lussac and Thenard, and later by Deville and Wohler.* 

Metallic scandium has not been prepared. 

Yttrium was prepared in an impure state, mixed with erbium, 
by Berzelius, by the action of potassium on the impure chloride. 
It was a greyish-black lustrous powder. 

Lanthanum has been prepared by passing a current of 
electricity through its fused chloride (see Lithium). 

Metallic ytterbium has not been obtained. 

Properties. Boro-^is a brown amorphous (i.e., non-crystalline) 
powder, which has not been melted even at a white heat. It is 
insoluble in all solvents which do not act on it chemically. It was 
for long supposed possible to crystallise it from molten aluminium ; 
the resulting black crystals, however, are not pure boron, but a 
compound of boron and aluminium. Yellow crystals, obtained by 
Wohler and Deville, and also supposed by them to be pure boron, 
consist of a compound of boron, carbon, and aluminium. Tl\e mass 
of 1 c.c. of pure boron has not been determined. Boron combines 
with the oxygen and nitrogen of air, burning to oxide and nitride. 
It is one of the few elements which combine directly with nitrogen. 
It is also attacked by chlorine and by bromine. Lanthanum is 
the only one of these elements which has been prepared in a 
compact state. It resembles iron in colour ; is hard, malleable, 
and ductile. It melts at a^lower temperature than silver (below 
1000), and burns with great brilliancy when heated in air; its 
specific gravity is 6*05 at the ordinary temperature. 

The specific heat of boronf undergoes a remarkable change as the tempera- 
ture is raised. The following results were obtained by Weber :- 

* Annalv (3), 52, 63. .f Phil. .Mag. (4), 49, 161, 276. 


Temperature .. 40 +27 77 126 177 233 

Specific Heat .. 0*1915 0*2382 0-2737 0'3069 0-3378 0'366:i 

, In this it resembles beryllium, carbon, and silicon. 

GROUP V. Aluminium, Gallium, Indium, Thallium. 

These elements are found only in combination. The sources of 
aluminium are its oxide, corundum ; when coloured blue, probably 
by cobalt, it forms the precious stone the sapphire, and when red, 
coloured by chromium, the ruby. Associated with iron oxide, it is 
named emery. Silicate of aluminium is a constituent of many 
rocks ; it exists in felspar, hornblende, mica, and numerous other 
minerals. .China clay or kaolin is a slightly impure silicate of 
aluminium (see Silicates). The mineral cryolite, found in Green- 
land, is a fluoride of aluminium and sodium. The sulphide of 
aluminium is decomposed by water ; hence its non-occurrence in 

The other three elements of this group occur as sulphides. 
Gallium and indium are found in extremely minute amount in 
some zinc ores ; thallium is contained in some specimens of iron 
pyrites (disulphide of iron) and copper pyrites. Zinc sulphide, 01 
blende, from the Pyrenees, contains about 0*002 per cent, ol 
gallium ; the zinc ores from Freiberg, in Saxony, about 0*05 to O'l 
per cent, of indium. 

Preparation. Aluminium is prepared : 

1. By passing the vapour of its chloride over heated sodium ;* 
the sodium unites with the chlorine, while the aluminium remains 
in the metallic state. 

2. By heating its oxide mixed with carbon to an enormously 
high temperature in the electric arc.f The oxide is thus decom- 
posed, and the carbon unites with the oxygen, while the metal iy 
left This process is better adapted for preparing the alloys ui 
aluminium than the metal itself. 

3. By heating with metallic sodium cryolite, the double 
fluoride of aluminium and sodium, previously fused with salt.J 

Gal Hum is prepared by passing a current of electricity 
through a solution of its oxide in caustic potash. 

Indium 1 1 may be obtained by passing a stream of hydrogen 

* Wohler, Annalen, 37, 66 ; Deville, Annales (3), 43, 5, and 46, 415. The 
literature on this subject is now very large, 
f Chem. News, 1889, 211, 225, 241. 
J Brit. Asscn., 1889. 
Comptes rend., 82, 1098 ; 83, 636. 
)| J. praJft. Chem., 1863, 89, 441 ; 02, 480 j 04, 1 ; 05, 414; 102, 273. 


gas over its oxide heated to a high temperature ; the hydrogen 
combines with the oxygen, producing water, and the indium is- 
left; or by heating its oxide with sodium; or by removing chlorine 
from indium chloride by placing metallic zinc in a solution of 
that substance. 

Thallium* is most easily obtained by heating its chloride to a 
red heat with potassium cyanide, a compound of carbon, nitrogen, 
and potassium. The potassium removes the chlorine, forming 
potassium chloride ; cyanogen, a compound of carbon and nitrogen, 
escapes as gas ; and thallium remains behind as fused metal. 

Properties. Aluminium, gallium, and indium are tin-white 
metals, while thallium has a duller lustre, resembling that of lead* 
These metals are moderately malleable and ductile. Indium and 
thallium are soft, and may be cut with a knife ; aluminium and 
gallium are hard. Thallium and its salts impart a magnificent 
green colour to the flame of a Bunsen's burner ; indium burns with 
a violet light ; aluminium and gallium do not volatilise sufficiently 
easily to colour the flame. 

All these elements unite readily with oxygen at a red heat; 
aluminium and thallium become tarnished in air at the ordinary 
temperature. They also combine directly with the halogens and 
with sulphur. They are not acted on by water at the ordinary 
temperature, but decompose it at higher temperatures, combining 
with its oxygen. 

Aluminium is contained in alum, hence its name ; gallium was 
discovered in 1875 by the French chemist, Lecoq de Boisbaudran, 
and patriotically named after Gaul ; indium derives its name from 
the blue line in its spectrum (from " indigo ") ; and thallium was- 
named by its discoverer Crookes, from 0a\\o'v, a green twig, in 
allusion to the green colour it imparts to the flame. 

Of these elements aluminium is the only one which has fcund 
a commercial use; the barrels of opera glasses, telescopes, and 
optical instruments are made of it ; and, alloyed with copper, it is 
employed for cheap jewellery, under the name of " aluminium 
bronze." Of recent years its manufacture has been greatly in- 
creased, and in the near future it will rank as one of the commoner 

The metals beryllium, magnesium, zinc, cadmium, lanthanum, didymium, 
cerium, and aluminium used to be classified together as "metals of the earths; " 
the so-called earths being their oxides, which are insoluble in water, and hence 
hare not an alkaline reaction like those of calcium, strontium, and barium. 

* CJiem. News, 3, 193* 303 ; froc. Roy. Soc., 12, 150. 


Physical Properties. 

Mass of 1 c.c. Melting- Specific Atomic Molecular 

Solid. point. Heat. Weight. Weight. 

Aluminium.. 2*583 at 4 About 700 '2253 from to 27*01 27 '01 

Gallium .... 5 -91 at 23 29 '5 Solid '079 from 69*9 69 '9 

12 to 23 
Liquid 0-080 from 

106 to 119 
Indium .... 7 '42 at 16 -8 176 '0565 to 0' 0574 113 '7 

Thallium.... 11-9 290 0'0336 204*2 204*2to 



The equations expressing the preparation of tho foregoing elements are as 
follows : 

Hydrogen. (\) 2H 2 O = 2H 2 + O 2 . 

(2) 2H 2 + 2Na = 2NaOH + H 2 

(3) 4H 2 + 3Fe = Fe 3 4 + 4H 2 . 

(4) H 2 S0 4 + Zn = ZnS0 4 -I- H 2 . 
Lithium, $c. 2LiCl = 2Li + C1 2 . 

Sodium and Potasnum.-~2N&OIl + 2C = 2Na + 2CO + H 2 . 
2KOH + 2C = 2K + 2CO + H 2 . " 
.Beryllium. BcCl 2 + 2Na = Be + 2NaCl. 
Calcium, Strontium, and Jlarium. BaClo = Ba + C1 2 . 
Magnesium. MgCl 2 .KCl + 2Na = Mg + 2NaCl + XC1. 
Zinc, Cadmium. ZnO + C = Zn + CO. 
CdO 1- C = Cd + CO. 
Boron. (1) BC1,, + 3Na = B + 3NaCL 

(2) KF.BF 3 + 3Na = B + KF + 3NaF ; 

(3) B 2 3 +' 3Mg = 2B + 3MgO. 
Aluminium. (\) A1C1 3 + 3Na = Al + 3NaCl.' 

(2) A1 2 3 + 3C = 2A1 + 3CO. 

(3) AlF 3 .3NaF + 3Na = Al + GNaF. 
lallium. 2Ga20 3 = 2Gb, + 30 2 . 

^ndium. (1) In 2 3 + 3H 2 = 2In + 3H 2 O. 

(2) 2InCl 3 + 3 Zn = 2In + 3ZnCl 2 . 
Thalliwn. 2T1C1 3 + 6KCN = 2T1 + GKC1 + 3(CN) 2 . 




GROUP VI. Chromium, Iron, Manganese, Cobalt, 


The elements of this group are not, generally speaking, asso- 
ciated in the periodic table, yet they closely resemble each other ; 
and it is convenient to consider them together. 

Sources. They invariably occur in combination with oxygen, 
when of terrestrial origin. Certain meteorites, however, consist 
largely of metallic iron and nickel with a little cobalt and a trace of 
hydrogen. Common proportions are 90 per cent, of iron, 9 per 
cent, of nickel, and 1 per cent, or less of cobalt. 

The chief ore of chromium is chrome iron ore, or chromite ; it 
is a compound of oxygen with chromium and iron (see Chromium, 
oxides, p. 254). It is found in Silesia, Asia Minor, Hungary, 
Norway, and N. America. The green colour of the emerald and 
serpentine is due to traces of chromium. 

Compounds of iron are very numerous in nature. Its oxides,, 
when found native, are named : Hcematite, of which varieties 
are termed specular iron ore, kidney ore, and titaniferous ore 
(these occur largely in Cumberland, also in the south of Spain) ; 
combined with water, gothite, brown iron ore, bog iron ore, and 
ake ore, the latter of which are named from their sources : 
they are found in Northamptonshire, the Forest of Dean, and 
Glamorganshire ; magnetic iron ore, magnetite or loadstone, an 
oxide of a different composition (see p. 255) : it does not 
occur largely in England, but is worked in Sweden ; the largest 
deposit of iron ore in the world consists of magnetite : it occurs in 
Southern Lapland, but is as yet inaccessible. Spathic ore, or car- 
bonate of iron, is a white crystalline substance when pure, but is 
usually interstratified and mixed with clay or shale, when it is 


termed "clay-band" or "black-band." Spathic ores occur in 
Durham, Cornwall, Devon, and Somerset ; clay iron-stone in the 
coaj-measures in Staffordshire, Shropshire, Yorkshire, Derbyshire, 
Denbigh, and South Wales ; while black-band is mined largely in 
the Clyde basin, in Scotland. 

Iron occurs in combination with sulphur as pyrites ; it is very 
widely distributed ; perhaps the largest sources are in the south 
of Spain. At Bio Tinto this ore is worked, not for the iron 
which it contains, but for its copper (about 3 per cent.) and its 
sulphur. Iron is also a constituent of most rocks and soils : it is 
cue of the most abundant as well as one of the most widely dis- 
tributed of elements. 

Manganese is nearly always found associated with iron, in 
combination with oxygen. Its most important source is pyrolusite 
or black oxide. Other manganese minerals are braunite and haus- 
mannite, also oxides ; manganite, psilomelane, and wad, compounds 
of oxides and water ; manganese-spar, the carbonate ; it also occurs 
in combination with silicon and oxygen as silicate, and with sulphur 
as sulphide. 

Cobalt and nickel are almost invariably associated. AvS 
already mentioned, they accompany iron in some meteorites in the 
state of metals. Cobalt occurs as smaltite or tin- white- cobalt, in 
combination with arsenic; and as glance-cobalt, in combination 
with arsenic and sulphur. 

The chief ore of nickel is the oxide and the double silicate 
of nickel and magnesium, large quantities of which are now im- 
ported from New Caledonia, a French convict settlement north-east 
of Australia. It is found on the continent of Europe chiefly as 
the arsenide, a compound of nickel and arsenic named Kupfer- 
nickel or copper-nickel, from its red colour resembling copper ; it is 
also called niccolite. The sulphide, or capillary pyrites, also occurs 

Preparation. These metals in an impure state may all be 
prepared by reducing (i.e., removing oxygen from) their oxides by 
means of carbon. Iron and nickel are prepared for commercial 
purposes ; alloys of iron and manganese, and iron and chromium 
are also produced ; and nickel is often deposited by means of an 
electric current on the surface of other metals, which are then said 
to be nickel-plated. 

Chromium, in the pure state, has been prepared by removing 
chlorine from its chloride, by means of metallic zinc or magnesium.* 
The chloride is mixed with potassium and sodium chlorides, and 
* Annalen, 111, 117. 


heated with metallic zinc to the boiling-point of the latter metal 
(about 940). An alloy of zinc and chromium remains, from which 
the zino may be removed by treatment with nitric acid; ,the 
chromium remains as a pale-grey crystalline powder. It has also 
been prepared by decomposing by electricity its chloride in* con- 
centrated solution. It then deposits in brittle scales with the 
lustre of metallic iron. 

Iron, in a state of purity, is hardly known. It has been 
prepared by reducing its oxide by means of hydrogen at a red heat, 
and heating the resulting greyish-black powder, which consists of 
pure iron in a state of fine division, to whiteness in a porcelain 
crucible under a layer of fused calcium fluoride in the oxyhydrogen 
flame.* It does not fuse, but agglomerates to a sintereid mass. It 
may also be deposited electrically from solution. Ordinary iron 
contains small quantities of several elements, notably carbon and 
silicon, which completely alter its properties, and it must, there- 
fore, be considered as a compound. A description of the metallurgy 
of iron is therefore deferred to Chapter XXXVI. 

Manganese, like iron, is almost unknown in a pure state ; 
when produced by the aid of carbon, it combines with that element 
and acquires peculiar properties. Its metallurgy will be considered 
along with that of iron. It has recently, however, been prepared 
in a pure coherent state by heating to redness with magnesium 
dust a mixture of manganese dichloride with potassium chloride. f 

Nickel is prepared in a manner exactly similar to that by which 
iron is made. Impure nickel can be prepared by heating its oxide 
with charcoal ; the pure metal is obtained by electrolysis. The 
same remarks apply to cobalt. 

Properties. These elements are all greyish- white, with metallic 
lustre, like iron. Manganese and cobalt have a redd ish- tinge ; 
nickel is whiter than iron, but not so white as silver. Tlj^ey all 
melt at a very high temperature, so high, indeed, that it is reached 
only by means of the oxyhydrogen blowpipe. The addition of a 
small amount of carbon, as has been remarked, profoundly modifies 
their properties ; and, indeed, the pure elements are almost un- 
known in a compact state, owing to the difficulty of melting them 
into a compact mass in any vessel capable of withstanding the 
requisite temperature, and not attacked by the metal. The figures 
in the following table refer, for the most part, to such impure 

They all combine with oxygen, on exposure to moist air, but are 
* Troost, Sull. Soc. Chim. (2), 9, 250. 
f G-latzel, Ser. Deutsch. CJhem. #<?*., 22, 2857. 


permanent in dry air ; they unite directly with the halogens ; with 
sulphur, selenium, and tellurium ; with phosphorus, arsenic, arid 
antimony ; with carbon, silicon, and titanium ; and they form 
alloys with each other and with many other metals. Iron and 
iiicke*! also absorb hydrogen gas to a small extent. 

Physical Properties. 

Specific .Atomic Molecular 

Mass of 1 c.c. Solid. Heat. Weight. Weight. 

Chromium . . 7 '3 ; G '81 (at 25) .... Not determined 52 3 ? 

Iron 8 -00 (at 10) pure '112 (impure) 56 '02 ? 

8 '14 (at 15 -5) electro- 

Manganese.. * 7 '39 at 22 0-122 550 55 '0 

Nickel About 9-0 0'109 58'6 ? 

Cobalt About 9'0 0'107 58*7 ? 

GROUP VII. Carbon, Titanium, Zirconium, 
Cerium,* Thorium. 

Of these elements, carbon is the only one found in the free 
state. The others are always found combined with oxygen, and 
usually with silicon and oxygen as silicates. 

The native forms of carbon are the diamond, carbonado, and 
graphite, black-lead, or plumbago. Diamonds are found in situ 
in pegmatite, or graphic granite, near Bellary, in the Nizam, 
India, and also in an aqueous magnesian breccia in S. Africa. It 
is probable that they have been formed simultaneously with these 
rocks ; the conditions of their formation are unknown. Diamond- 
fields, or districts which yield diamonds, occur in Brazil, India, the 
Cape, California, Borneo, and the Ural Mountains. 

Cdrbonado, a variety of carbon found in the Soap Mountains of 
Bahla, is a reddish-grey, porous substance ; it is evidently closely 
allied with diamond. 

Graphite occurs in nests of trap in the clay slate at Borrowdale, 
Cumberland, and is also found in certain coal-measures, e.g., at 
New Brunswick. 

Such different forms of an element are said to be allot ropic, n 
word which signifies "different forms." 

Carbon also occurs in combination with oxygen (th.3 atmo- 
sphere contains about 0'04 per cent, by weight of carbon dioxide) , 
and its dioxide, with the oxides of various metals; the most 

* It is doubtful if cerium belongs to this group of elements. 


important of the carbonates are those of calcium, of magnesium, 
and of irou (q.v.). 

Along with hydrogen, oxygen, and nitrogen, it is a constituent 
of all organised matter ; coal, which consists of ancient yegetable 
matter, agglomerated by pressure and decomposed by heat,- contains 
a large percentage of carbon, anthracite, for instance, containing 
over 90 per cent. 

Titanium occurs only in combination with oxygen, as rutile, 
anatase, and brookite ; and associated with oxides of iron, as titani- 
ferous iron ; with oxide of calcium in perowsJcite ; and with the 
oxides of silicon and calcium in sphene. 

Zirconium is found as oxide, in combination with oxide of 
silicon in zircon, and in other rare minerals. r 

The chief source of cerium is cerite, a compound of oxide of 
cerium with oxide of silicon and with water ; and it occurs asso- 
ciated with oxides of niobium, tantalum, lanthanum, didymium, &c., 
in orthite, euxenite, and gadolinite, and other very rare minerals. 

Thorium occurs in thorite as oxide, in combination with oxide 
of silicon and with water ; it also occurs in euxenite, &c., along 
with cerium. 

Preparation. Carbon is produced by the decomposition by 
heat of its compounds with hydrogen, sulphur, and nitrogen ; at a 
very high temperature its oxide is also decomposed. It may also 
be produced by withdrawing chlorine from any of its chlorides by 
means of metallic sodium, or oxygen from its oxides by metallic 
potassium. Chlorine also removes hydrogen at a red heat from 
its compounds with that element, setting free carbon in the form 
of soot. It is best prepared in a pure state by the first of these 
processes. Sugar and starch consist of carbon in union with hydro- 
gen and oxygen. On heating these bodies out of contact with air, a 
large portion of the carbon which they contain remains in the state 
of clement. It is advisable, in order to remove hydrogen com- 
pletely, to heat to redness in a current of chlorine. The deposit on 
the upper surface of the interior of retorts during the manufacture 
of coal-gas by the distillation of coal is named gas- carbon, and is 
nearly pure. It is thus produced by the decomposition of hydro- 
carbons (compounds of hydrogen with carbon) by heat. Various 
impure forms of carbon are prepared for industrial purposes. Wood 
charcoal is obtained by heating wood to redness in absence of air. 
This used to be the work of " charcoal burners," and the manufac- 
ture still survives in Epping Forest. Faggots of wood are piled 
into a tightly packed heap, covered over with turf, and set on fire, 
a limited quantity of air being admitted to support combustion. 


Most of the wood is thus charred ; and when smoke ceases to be 
emitted more turf is heaped on, so as to extinguish the fire. 
The mass of charcoal is then allowed to cool, and when cold, the 
covering of turf is removed, and the billets of charcoal unpiled. 
The wood yields about 34 per cent, of its weight of charcoal. In 
the present day, oak or beech wood is distilled from iron retorts, 
for the production of acetic acid, or viuegar ; the retorts are heated 
with coal, and the charcoal remains in the same form as the logs 
which are put into the retort. The charcoal made in this way is 
used chiefly by iron -founders to mix with sand in making moulds 
for castings. Charcoal for gunpowder is made from willow, dog- 
wood, or alder. 

Coke, th residue on distilling coal, is also impure carbon. The 
coke forms from 40 to 75 per cent, of the weight of the coal. 
Coke is largely used as a fuel, especially in iron smelting. 

Bone or animal charcoal, or bone-Hack, produced by distilling 
bones, contains about 10 per cent, of carbon, the remainder chiefly 
consisting of the mineral constituents of bones, calcium phosphate 
and carbonate. It is used for decolorising solutions of impure 
sugar, which are filtered through the bone-black, ground to a 
coarse powder. Its decolorising properties are much increased by 
dissolving out the calcium compounds by washing it with hydro- 
chloric acid. 

Lamp-black, chiefly used for printers' ink, is prepared by burning 
certain compounds of carbon and hydrogen, especially one con- 
jstituent of coal-tar oil, named naphthalene. The hydrogen and a 
portion of the carbon burn, while the greater portion of the carbon 
is carried away as smoke, and condensed in long flues. 

Titanium,* like carbon, may be produced by passing the 
vapour of its chloride over heated sodium,, when the sodium 
removes the chlorine as sodium chloride, leaving the element, with 
whioh sodium does not appear to form a stable compound. It may 
also be produced by projecting into a red hot crucible potassium, 
cut into small pieces, along with potassium titanifluoride (a com- 
pound of titanium, potassium, and fluorine) ; the fluorine is removed 
by the potassium as potassium fluoride, which is soluble in water, 
and may be separated from the titanium by treatment with water, 
in which titanium is insoluble, and with which it does not react in 
the cold. 

Zirconium, f like titanium, may be produced by heating potas- 

* Wohler, Annales (3), 29, 181. 

f Berzelius, Pogg. Ann., 4, 124, and 8, 186. Troost, Comptes rend., 61, 


slum zirconifluoride with potassium, or magnesium. The metal 
aluminium also withdraws fluorine from this compound and the 
zirconium dissolves in the metal, crystallising from it when it 
cools. The aluminium is removed by treatment with dilute hydro- 
chloric acid, in which zirconium is insoluble. 

Cerium* has been prepared by electrolysing cerium chloride, 
covered with a layer of ammonium chloride, contained in a porous 
earthenware cell, placed inside a non-porous crucible, filled with a 
mixture of sodium and potassium chlorides. The whole arrange- 
ment is heated to redness, so as to melt the compounds. On 
passing an electric current, the cerium deposits on the negative 
electrode, which is made of iron, inserted through the stem of a 
clay pipe to prevent its oxidation by the hot air. r 

Thoriumf has also, like carbon and titanium, been prepared 
by heating with sodium in an iron crucible potassium thorifluoride, 
covered with a layer of common salt. 

Properties. Carbon, as already mentioned, exists in several 
different forms, each of which has distinct properties. It is there- 
fore said to display allotropy. 

The diamond is transparent, crystalline, the hardest of all 
known substances ; it is nearly pure carbon. It is usually colour- 
less, but is occasionally coloured green, brown, or black by mineral 
matter. It was found to be combustible by the Florentine 
Academicians, in the !Mi century, who succeeded in burning it 
by concentrating the sun's rays on it by means of a large lens. 
Lavoisier discovered its identity with carbon. It can be converted 
into a coke-like substance when exposed to the intense heat of the 
electric arc. It is used as a gem, also for rock boring, and for cut- 
ting glass. Its dust is employed in cutting and polishing precious 
stones, and in cutting other diamonds. 

The weight of diamonds is measured in carats^ (1 carat 
= 0'205 gi'am, or 3'1C5 grains). The value, however, is not pro- 
portional to the weight, but to approximately the square of the 
weight. Among the most remarkable diamonds one of the largest 
belongs to the Nizam of Hyderabad, and weighs 277 carats ; the 
Crown of Russia possesses another, of a somewhat yellow colour, 
weighing 194 carats; the Koh-i-Noor, or " mountain of light," 
belonging to the British Crown, weighs 106 carats. It was 

* Hillebrand and Norton, Pogg. Ann., 155, 633 ; 166, 466. 
f Nilson, BericUe, 1882, 2519 and 2537; 1883, 153. 

{ Kerat (Arab.), supposed to be derived from rati, the Indian name for the 
seeds of Abrus precatorius. 


originally much larger, but was reduced in weight by cutting. 
The cutting of diamonds is intended to display their great power 
of refracting light. The two forms in which diamonds are cut are 
that of the brilliant, which fig. 5 represents, and the table or 
rosette form, shown in fig. 6. The former is the most valuable. 

FIG. 5. Fia. 6. 

Carbonado is a very hard substance, which is also used for rock 
boring. It has been noticed as a constituent of some meteorites. 

Graphite is a blackish-grey, lustrous, soft substance, chiefly 
used for making very refractory crucibles, when mixed with clay ; 
also for fine iron-castings, and for lead pencils. 

No attempt to produce diamonds artificially has succeeded, 
except perhaps one in which carbon was kept in contact with a 
large quantity of melted silver ; the carbon appears to be slightly 
soluble in the fused metal, and microscopic crystals which sepa- 
rated out on cooling were said to possess the properties of the 

Graphite may be made artificially by dissolving carbon in 
molten iron, which dissolves 1 or 2 per cent When the metal is 
slowly cooled, part of the carbon separates in this form. It has 
also been prepared by heating amorphous carbon to an extremely 
high temperature by passing an electric current of high potential 
through a rod of carbon, and thus heating it to brilliant incan- 
descence. The temperature at which the change is produced is 
unknown, but is enormously high. 

Carbon is infusible at the ordinary pressure. It is volatile in 
the electric arc, which when formed, as is usually the case, by 
passing an electric current between two rods of gas- carbon, always 
possesses the temperature at which carbon volatilises. What that 
temperature is has not yet been ascertained. 

The diamond is a non-conductor of electricity, like indeed all 
transparent bodies ; but the other forms of carbon conduct, though 
not so well as metals. 

Carbon in one form or another, especially as coal, is the source 
of all the heat and energy practically utilised by mankind. To 
utilise this energy stored in carbon it is burned in air, uniting with 
its oxygen. Charcoal unites very slowly with oxygen at the 


ordinary temperature, but rapidly at a red heat. The other forms 
of carbon also burn, but slowly, when heated to redness in oxygen. 
At a red heat carbon deprives most other oxides of their oxygen, 
and is therefore used in extracting metals from their oxides. It 
unites directly with sulphur when heated to 'redness in sulphur 
vapour ; and with hydrogen at the temperature of the electric arc, 
to form acetylene. It does not appear to combine directly with the 
other elements, although many compounds have been prepared by 
indirect methods. 

Animal charcoal and, to a less degree, wood charcoal, owing 
probably to their cellular nature and to the great amount of 
surface which they possess, have the power of condensing and 
absorbing gases. The amount of absorption is in the* same order 
to the condensibilities of the gases, those gases which are condensed 
to liquids by the smallest lowering of temperature being absorbed 
in greatest amount. Thus 1 volume of boxwood charcoal absorbs 
90 volumes of ammonia-gas, which condenses to a liquid at 36, 
whereas it absorbs only 1*75 vols. of hydrogen, which is probably 
not liquid at a temperature of 230. 

Titanium is a dark grey powder, like iron which has been 
reduced from its oxide by hydrogen. It has not been fused. It 
unites directly with oxygen and with chlorine, burning when 
heated in these gases. It has also the rare property of uniting 
directly with nitrogen when heated in that gas. It decomposes 
water at 100, combining with its oxygen, and liberating 

Zirconium forms brittle lustrous scales. It fuses at a very 
high temperature, and does not combine with oxygen at a red 
heat; but at a white heat it burns to oxide. It unites directly 
with chlorine. 

Thorium is an iron-grey powder, which burns brilliantly 
when heated in air, forming oxide. Like titanium and zirconium, 
it is attacked by chlorine, burning brilliantly in the gas ; also by 
bromine and iodine. It unites directly with sulphur. 

Physical Properties. 

Mass of 1 c.c. 


Atomic Molecular 



Weight. Weight. 

Carbon (Diamond) .. 

3-514 at 18 . . 

(see below) . . 

12 -00 ? 

(Graphite) . . 

2-25 at ? . . 

(see below) . . 


(Charcoal) .. 

About 1-8 at ? 





48-13 ? 

4 '15 at ? . . 


90*0 ? 


11 -1 at 17. . 


232-4 ? 



Note. The specific heat of carbon increases very rapidly \\ith rise of tem- 
perature (of. beryllium and boron).* 

Temperature . . 
Sp. heat 
diamond . 



+ 10 

+ 33 

+ 58 

+ 8(> 

Graphite . . 





Temperature . 
Sp. heat 
Diamond . 

+ 140 

+ 206 

+ 24-7 

-t (500 

+ 800 

-f 1000 

Graphite . . 







It is noticeable that, although at low temperatures the specific heut of the 
diamond differs considerably from that of graphite, yet at high temperatures 
they nearly coincide. 

GROUP VIII. Silicon, Germanium, Tin, Terbium (?), 


The first of these elements, silicon, closely resembles titanium ; 
it is a blackish, lustrous substance. Germanium and tin are white 
metals, with bright lustre; lead is of a greyer hue. Terbium, as 
element, has not been prepared in a pure state. 

Sources. The element silicon, next to oxygen is the most 
widely distributed and abundant of the elements on the surface of 
the globe, forming about 25 per cent, of its total weight. It 
never occurs in the free state, being attacked by oxygen : hence 
it exists only as oxide (silica), alone ; or in combination with other 
oxides, as silicates. As such, it is contained in a vast number 
of minerals. Some of the most typical of these are described 
under the heading silica (p. 800). The more commonly occurring 
forms of silica are quartz, sandstone, and flint ; pure crystallised 
silica is named rock-crystal, bog diamond, or Iritfh diamond; agate, 
chalcedony, opal, &c., are other forms. Granite, trap, basalt, 
porphyry, schist, and clay are rocks entirely composed of silicates. 
The name is derived from silex, flint. Germanium,! an element 
patriotically named, has been recently discovered by Winkler in 
a mineral termed argyrodite found in the Himmelsfiirst mine, near 
Freiberg. It contains 6 or 7 j)er cent. (?) of sulphide of ger- 
manium. Euxenite is also said to contain a trace of germanium, 
about O'l per cent. 

Tin is a moderately abundant element, although not widely dis- 
tributed. The chief mines are in Cornwall ; it also occurs in the 
Erzgebirge, in Saxony and Bohemia, in the Malay Peninsula, and 

* Weber, Fogg. Ann., 154, 367. 

f Winkler, Berichte, 19, 210; J. praTct. Chem. (2), 34 M 177. 


in Peru. Large deposits of tin-ore have recently been discovered 
in Australia and in Borneo. It occurs as oxide, in tin-stone or 
c (twite rite, and as sulphide, a comparatively rare mineral. It is 
never found as n, metal, owing to its tendency to oxidise. 

Terbium, the connection of which with this group of elements 
is open to question, is associated with yttrium (<.#.), and is con- 
tained in the same minerals as yttrium. 

Lead, like tin, is always found in combination, chiefly with 
sulphur in galena, a very widely distributed ore, found in the Isle 
of Man, in Cornwall, in Derbyshire, in the south of Lanarkshire, 
and in many foreign countries. Other ores of smaller impor- 
tance are the carbonate or cerussite, the sulphate, phosphate, and 
arsenate. * 

Preparation. Silicon,* like titanium, is produced by with- 
drawing chlorine from its chloride by passing the vapour of the 
Litter over red hot sodium ; or by removing fluorine from its double 
compound with fluorine and sodium, sodium silicifluoride, by 
means of metallic sodium; the metal zinc may also be used 
alorg with sodium to withdraw the fluorine, when the silicon 
crystallises from the zinc, which may be removed by dissolving 
it in weak hydrochloric acid. Germanium, tin, and lead are 
reduced from their oxides by heating 111 hydrogen or with carbon. 
Terbium has not been prepared. For the preparation of tin and 
lead on the large scale, the chapter on the oxides and sulphides of 
these metals must be consulted (p. 296, and also Chap. XXXVIII) ; 
for their metallurgy involves somewhat intricate operations. 

Properties. Silicon lacks metallic lustre, and is therefore 
usually classed with the non-metals. It is a blackish brown 
powder, which, when crystallised from zinc or aluminium, 
separates either in black lustrous tablets resembling graphite, or in 
brilliant hard iron-grey prisms. It fuses at a high temperature, 
and may be cast into rods. It is contained in cast iron, probably 
however in combination with the iron. The crystalline variety 
conducts electricity. When heated in oxygen, chlorine, bromine, 
or sulphur gas, silicon combines with these elements. The crystal- 
line variety can be dissolved only by fusion with caustic potash 
(see Silicates, p. 310). 

Germanium is a white metal, somewhat resembling antimony. 
It is very brittle and can be readily powdered. It may bo melted 
under a layer of borax, which prevents oxidation ; it is, however, not 
very easily oxidised. It melts at a bright red heat. It combines 

* Deville and Caron, Annales (3), 67, 435. 



directly witb oxygen, sulphur, and the halogens when heated in 
the vapour of these elements. 

Tin is a lustrous white metal resembling silver. It is very 
soft and malleable, and may be hammered into foil (tin foil), but 
its wire has little tenacity. Up to 100 its malleability increases ; 
but, like zinc, it becomes brittle at higher temperatures, and may 
be powdered at 200. Its fracture is crystalline. It melts at a 
low temperature. 

Although not oxidised at the ordinary temperature, it burns in 
air with a white flame when strongly heated ; it also unites directly 
with the halogens and with sulphur. It forms alloys with many 
other metals which find commercial use. It is also largely used 
in tinning i^on (see Alloys, p. 58. J). An allotropie form of tin is 
produced when tin is cooled to a low temperature, or when it is kept 
fora long time ; it is greyish-red, and exceedingly brittle. When 
heated to 50 for some hours it is reconverted into ordinary tin.* 

Lead has a greyer shade than tin. It is soft, and may be cut 
with a knife. It may be hammered into foil, and drawn into wire, 
which however has little tenacity. It is easily fused, and volati- 
lises at a white heat. 

Lead combines directly with oxygen at a high temperature, 
forming " dross " ; although not affected by dry oxygon, moist 
atmospheric air soon tarnishes it. When heated with the halogens 
or with sulphur, it combines directly with them. It is used 
largely for pipes, for covering roofs, for bullets, shot, &c. ; and 
its various alloys find a very wide application. 

Mass of Ice. 


Graphitoidal 2 2 at ? ____ 
Adamantine 2 -48 at ? . 

J'hyaical Properties. 


Specif! (3 

Atomic Molecular 
Weight. Weight. 

(see below) 28 '33 


5 '47 at 20 '4 

Tin (solid) ____ 

(liquid) .. 

Lead (solid) . . 

7 '29 at 13. . 
7 '18 at 226 
6 '99 at 226 
5 '8 to 6 '0.. 

11 '35 at 14. . 

11 -Oat 325. . 
(liquid).. 10-65 .. 

About 1100 
About 90U 



00758 723 
(100 to 440) 

'0562 119 1 


206 -93 

206 '93 

The specific heat of silicon, like that of beryllium, boron, and carbon, varies 

* Fritsche, Phil. Mag. (4), 38, 207. 

F t 


greatly with the temperature, and attains approximate constancy only at high 
temperatures. The following determinations were made by Weber.* 

Temp - 40 +22 +57 + 86 + 129 + 184 + 232 

Sp. heat... 0*1360 0'1697 0'1833 O'lOOl 0'1964 02011 0'2029 


Equations expressing the preparation of elements of Groups VT, YIT, and 

Chromium.-~2CrC]s + 3Zn = 2Cr + 3ZnCL> 
Iron. FeO + H 2 Fe -f IF 2 O. 
Manganese. MnO 4- C = Mn + CO. 
Carbon. CC1 4 + 4Na = C + 4NaCl. 
Titanium. 2KF.TiF 4 + 4K - Ti + GKF. 
Cerium 2CeCl 3 = 2Co + 3C1 2 . 
Silicon 2NaF.SiF 4 + 4Na = Si + GNaF. 
Germanium. Ge0 2 + 2H 2 = Go + 2IT 2 O. 
T/w.SnO 2 + 2C - Sn + 2CO. 
Lead. PbO + C = Pb + CO. 

Pogg. Ann , 154, 367. 



GHOUP IX, Nitrogen, Vanadium, Niobium (or 
Columbium), Didymium* (?), Tantalum. 

THE first element of this group, like the first of the seventh group, 
does not outwardly resemble the remaining ones. It is a colourless 
gas, whereas the others are solids with metallic lustre. It exists 
tree, like carbon, while the others occur only as oxides, because 
they readily combine with oxygen. 

Sources. Nitrogen forms nearly four-fifths of the volume 
as well as of the weight of air. It occurs also in small amount in 
air as ammonia, in which it is combined with hydrogen. Am- 
monia also exists in the soil, being carried down by the rain, and 
yields its nitrogen to plants, which use it as food, assimilating it 
by means of their roots. Nitrogen is essential to the life of plants 
and animals, and is a constituent of the albuminous matters of 
which they largely consist. Coal, the relic of a former vegetation, 
also contains nitrogen in combination with carbon, hydrogen, and 
oxygen. Lastly, it occurs in combination with oxygen and sodium, 
and with oxygen and potassium, in sodium and potassium nitrates, 
which encrust the surface of the soil of dry countries. They are 
exported from India, and from S. America. Nitrogen has no 
great tendency to combine with other elements ; hence it chiefly 
occurs in the free state. The spectroscope has also revealed its 
presence in some nebulae. 

Vanadium is a comparatively rare element. It is found in 

* It is doubtful whether didymiuin belongs to this group of elements. It, 
appears to be a mixture, not a simple substance. See p. 002. > 


combination with oxygen, along with lead, copper, and zinc oxides, 
as vanadates of these metals. A crystalline incrustation on the 
Keuper Sandstone, at Alderley Edge, in Cheshire, in which vana- 
dium is associated with phosphorus and copper, named mpttra- 
mite, is one of its chief sources. 

Niobium, tantalum, and didymium are associated with rare 
metals, such as yttrium, cerium, lanthanum, &c., in euxenite and 
similar minerals. The two former are also found in combination 
with iron and manganese in niobite and tantalite, minerals found 
in the United States and in Gi*eenlaiid. 

Preparation. Nitrogen is usually prepared by removing 
oxygen from air, which consists mainly of these two gases. By 
heating ammonia, its compound with hydrogen, to a red heat, it is 
decomposed into its constituents ; but the hydrogen is not easily 
separated from the iiitrogen ; hence the plan usually adopted is to 
decompose ammonia by the action of chlorine, or by oxygen at a 
red heat, both of which unite with the hydrogen, liberating 
nitrogen. Perhaps the best method is one in which the oxygen of 
the air is made to combine with the hydrogen of the ammonia; 
the nitrogen of both air and ammonia is thus collected. The appa- 
ratus is shown in the accompanying figure. 

A gas-holder, A, is connected with a JJ-tube, B > filled with weak 
sulphuric acid, which in its turn communicates by means of 
indiarubber tubing witl^a tube of hard glass, c, containing bright 
copper turnings. The other end ol the hard-glass tube is joined 

Fia. 7. 

to a wash^bottle, half full of strong ammonia solution. The copper 
is heated ]fco bright redness, and the water in the gas-holder is 
allowed to'estaDe, a current of air being thus drawn through the 


ammonia- solution. The gaseous ammonia is carried along with 
the air over the red-hot copper. The oxygen of the air unites in 
presence of the red-hot copper with the hydrogen of the am- 
monia, forming water, and the nitrogen of the air along with 
the nitrogen from the ammonia both pass on. The sulphuric 
acid in the (J-tube serves to retain excess of ammonia, and pure 
nitrogen is the product. 

Nitrogen may also be prepared by leaving air in contact with 
any absorbent for oxygen ; for example, phosphorus ; or a solution 
of cuprous chloride in ammonia; or potassium pyrogalhite. 

Vanadium* has been prepared by withdrawing chlorine from 
one of its compounds with that clement by the action of hydrogen 
a red heat. The utmost precautions must be taken to exclude* 
oxygen and moisture, as vanadium is at once oxidised at a red heat 
by these substances. As it attacks porcelain, it must be heated in 
a platinum boat placed in a porcelain tube, during the passage of 
the hydrogen. The method of preparation of niobiumf is similar 
to that of vanadium. 

Tantalum, is said to have been prepared by a method similar 
to that which yields silicon, viz., by heating with metallic sodium 
its compound with potassium and fluorine. 

Didymium has been made in the same; manner as cerium 
(V/.y.). The substance thus called is certainly a mixture of at 
least two metals (see p. (50.5). 

Properties. Nitrogen iV a colourless, odourless, tasteless gas, 
somewhat lighter than air It condenses to a colourless liquid at 
the very low temperature 19M,J and solidifies to white flakes 
at 214, when cooled by boiling oxygen. The liquid is lighter 
than water. 

The only elements with which it combines easily and directly 
at a red heat are magnesium, boron, titanium, and vanadium. At 
a, higher temperature, that of the electric spark, it combines with 
hydrogen and with oxygen ; indeed combination between oxygen 
and nitrogen may be caused by burning magnesium in air, when 
the constituents of air unite to form peroxide of nitrogen ; and 
this gas is suddenly cooled by escaping away from the source of 
heat, and therefore remains undecom posed. 

Vanadium is a white substance with metallic lustre. It does 
not combine with oxygen at the ordinary temperature, nor even 

* Roscoe, Proc. Roy. Roc., 18, 37 and 310. 
t Roscoe, Chem. j\Vm, 37, 25. 
Comptes rend., 100, 350. 


at 100, but it takes fire spontaneously and burns in chlorine. It 
is unaltered by water, except at high temperatures. 

Niobium forms an irridescent steel -grey powder. . 

Didymium is a white metal with a faint yellow tinge. 

Tantalum is said to be a black powder, but it is doubtful 
whether it has been isolated. 

Physical Properties. 

Mass of 1 c.c 

A ._^ Density 

Liquid. Gas 11 = 1. 

-89 at - 194 -4 '001258 14 '08 

Nitrogen .... 



89 at 

Vanadium . 
Niobium . . 

5 87 at 15 
. . 7 0(> at 155 

'tantalum . . 

... 10 -5 ? 

Melting- Boiling- Specific Atomic Molecular 

point. point. Heat. "Weight Weight. 

Nitrogen -214 -It) 14 2138 11-03 2800 

Vanadium . . Not at bright ? 51 t p 

red heat 

Niobium . . Very high ? 94 v 

Didymium .. ., 0150 j p d f } !^ 1 ' 

J \ 1 rdi 143 6 J 

Tantalum.... - ? 182' 5 ? 

The critical temperature i r i' nitrogen is 1 10, and the critical prew>urr 35 
atmospheres f Its vapour-pressures are as l'ollo\\s . 

Pressure in atmospheres .. 35 31 17 1 Verj low 

Temperature - L M) - L IS 2 -1GO 5 - 19 1 t" -213 

Under a pressure of about 4000 atmospheres, nitrogen has the density 
at ordinary temperature, eomj)ared ^ith water. 

(Hour X.Phosphorus, Arsenic, Antimony, 
Erbium,J Bismuth. 

Owing to the great tendency of phosphorus to unite with 
oxygen, t is always found in combination with it. Arsenic, too, 
is seldom found native ; it also is easily oxidised. Antimony 
and bismuth ai*e botli found native. Erbium is always asso- 

* Neodjnnum and praseodymium, two bodies into which did\mium lias been 
re* o I red. 

t Comptes rend., 09, 133, 184; 100, 350. 

1 It is doubtful whether erbium belongs to this group. 


ciated with cerium, lanthanum, yttrium, &c. Phosphorus, arsenic, 
and antimony display allotropy. 

{Sources. Phosphorus occurs chiefly in combination with 
oxygen and calcium, as calcium phosphate, in minerals named apatite, 
in which it is associated with fluorine ; ph wphorite, an earthy 
variety; and in coprolites. It is also found as phosphate of alu- 
minium, or wavellite in large deposits; lead and copper phosphates 
also occur native. It is a constituent of all soils, though in 
minute amount. From them it is absorbed by plants, and is 
hence a constituent of all vegetable matter, especially seeds. 
Through plants it is assimilated by animals, and forms a con- 
stituent of the bones (about 58 per cent.) and the nerves. Ignited 
bones consist mainly of calcium phosphate. 

Arsenic occurs most abundantly in combination with iron as 
arsenical iron, and with nickel and cobalt as kupfnr-nichel and 
smaltife ; also with iron and sulphur in arsenical pyrites and 
minpickeL With sulphur it forms realgar and orpiment ; and it is 
also found combined with oxygen and metals as arsonatcs. It is 
sometimes found native. 

Antimony is rarely found native ; its most abundant ore is 
stibnite, or antimony sulphide ; it also occurs as antimony ochre or 
oxide; and it is associated in various minerals with sulphur and 
lead, mercury, copper, silver, &c. 

Bismuth is a comparatively rare metal, and nearly always 
occurs native. It is sometimes associated with tellurium. 

Erbium accompanies yttrium, cerium, &c. ((/.v.). It is 
extremely rare. 

Preparation. Phosphorus was originally obtained by 
Brandt by distilling dried and charred urine at a white heat. 
The carbon resulting from the decomposition of the animal 
matter deprived the sodium phosphate of its oxygen, and phos- 
phorus distilled over. It is still made by a somewhat similar 
process. Metaphosphoric acid, a compound of phosphorus with 
oxygen and hydrogen, is distilled with powdered coke or charcoal 
from clay retorts. The carbon deprives this substance of its 
oxygen, and phosphorus, hydrogen, arid oxide of carbon pass over in 
the gaseous state. The hydrogen and carbonic oxide gases escape, 
while the phosphorus is condensed and falls into warm water. For 
a detailed description of the process see Chap. XXXVIII. 

Arsenic is produced by distilling mispickfl, when the arsenic, 
which is very volatile, distils over, leaving the sulphur arid iron 
behind as ferrous sulphide. 

Antimony is prepared by heating its sulphide (stibnite) with 


scrap iron. The iron withdraws the sulphur, and the antimony 
separates in the metallic state. It is not sufficiently volatile to be 
conveniently distilled, but it flows down, forming a layer below 
the sulphide of iron. These operations must all be conducted in 
absence of air, for phosphorus, arsenic, and antimony readily corn- 
bine with oxygen. 

Bismuth is freed from earthy impurities by melting it in a 
crucible, when it sinks to the bottom. Arsenic, antimony, and 
bismuth may also be obtained by heating their oxides in a current 
of hydrogen. Krbium has not been prepared. 

Properties. Phosphorus exists in two allotropic forms. 
The variety longest known, culled yellow or ordinary phosphorus, 
is a yellowish-white, waxy substance, possessing a strong disagree- 
able smell. It has a great tendency to combine with oxygen even 
at ordinary temperatures, and shines in the dark owing to slow 
oxidation ; hence the name phosphorus (from 0v, light, and 
06/jeti/, to bear). It must, therefore, be kept under \\atcr. It is 
easily fusible, but when melted in air it takes tire and burns. It 
also catches fire when rubbed on a rough surface, owing to the 
heat produced by friction. Hence its use for lucifer matches. ]t 
is soluble in carbon disulphide, a liquid compound of carbon and 
sulphur, and may be obtained in crystals by the slow evaporation 
of the disuiplude ; this solution is named " Greek tire." When 
the solvent evaporates, the phosphorus is left in a tinely divided 
state, and is spontaneously inflammable. It is also soluble in 
alcohol, ether, olive oil, turpentine, benzene, and in certain of its 
own compounds, such as chloride and oxyrhloride of phosphorus. 
It is easily distilled at a moderate temperature (290). Its vapour 
is yellow. 

When heated to 240 for some time in a closed vessel in 
absence of oxygen, it changes to a red variety, named, red, or 
amorphous, phosphorus. The change may be brought about more 
quickly by a higher temperature, or by addition of a trace of iodine 
to the phosphorus. It is also produced under water on exposure 
of the yellow variety to light. But red phosphorus, when heated, 
also changes back to yellow phosphorus. Such a change, which 
can take place in two directions, is termed a limited reaction. Red 
phosphorus is insoluble in all ordinary solvents ; hence it may be 
purified from yellow phosphorus by digestion with carbon disul- 
phide. It does not combine with oxygen at the ordinary tempera- 
ture, nor perhaps at any temperature, for it burns in air only when 
made so hot that the change into the yellow variety begins. The 
colour of allotropic phosphorus varies, according to the tempera- 


ture at which it is formed. If prepared at 260 it is deep red, and 
has a glassy fracture ; at 440 J it is orange, ami has a granular 
fracture; at 550 it is violet-grey; it fuses at 580, and solidities 
to red crystals, which have a ruby-red fracture.* It is necessary 
to exclude air and to heat the phosphorus under pressure to pro- 
duce these changes. It dissolves m melted lead, and separates on 
cooling in nearly black crystals. 

Yellow phosphorus combines directly and very readily with 
oxygen and the halogens. It also unites with sulphur, selenium, 
and tellurium, and with most metals. lied phosphorus combines 
directly with the halogens. Neither variety unites directly with 
nitrogen or with hydrogen. Yellow phosphorus is poisonous, doses 
of 1 grain and upw.mls producing fatal effects, but in simill doses 
it is a powerful remedy for nervous disorders. Yellow phosphorus 
is a noii-eonductor of electricity, but red phosphorus conducts. 

Arsenic is a very brittle steel-grey substance with nu'tallic 
lustre on freshly broken surfaces, When heated, it sublimes with- 
out melting, and condenses partly in crystals, partly in ?i black 
amorphous (i.e., non-crystal line) state. Tt may, however, be 
melted under great pressure. The amorphous variety is rendered 
crystalline by heating it to .*Wo .f Jt readily combines with 
oxygen, and hence loses its lustre on exposure to moist air It 
biifiis when heated in air, spreading a garlic-like smell. It unites 
directly with oxygen, with the halogens, and wif h most other ele- 

Antimony, like phosphorus and arsenic, also exists in two 
forms Ordinary antimony is a bluish- white metal, vory brittle 
and crystalline. It is not oxidised by air at ordinary tempera- 
tures, and does not tarnish on exposure. AHotropic antimony J is 
obtained by electrolysing a strong solution of antimony chloride. 
A greyish deposit is formed on tho negative polo, which has the 
remarkable property of exploding when struck. Its specific 
gravity is considerably less than that of the ordinary variety. It 
is said, however, to contain hydrogen. Antimony unites directly 
with all elements, except nitrogen and carbon. 

Bismuth is a greyish- white metal with a red tingo, also very 
brittle and crystalline. The conductivity for electricity of the 
three elements arsenic, antimony, and bismuth rises in the order 
given. Bismuth is the most dianiagnetic of the elements. 

Erbium, has not been isolated. 

* Compte* rend., 78, 748. 

t Ibid., 96, 497 and 1314. 

1 Gore, Chem. Soc. 7 M 16, 365 ; Bottger, J. pralct. Chem., 73, 484 j 1O7, 43. 



Physical Properties. 
Mass of 1 c.c. 

Phosphorus, yellow . . 


II = 1. 

65 -Oat 1040 

45 '4 at 1700 

Arsenic, amorphous . . 
,, crystalline . . 
Antimony, ordinary . . 
,, explosive . . 

4- 7 at 14 
5-73 at 14 
6-67 at 155 
5 -7 to 5-8 
9 -8 at 13-5 

Solid. Liquid. 

1 -83 at 10 1 -75 at 40 

1 -49 at 278 
2-15 to 2 -3 

(at 0) 


I 77 -5 at 1736 
C'5? ri55 'I at 1572 

1 141 -2 at 1640 

44 4 



10-0 24G -2 at 1700 2G8 3 


Phosphorus, yellow . . 


Arsenic, amorphous . . 

,, crystalline . 

Antimony, ordinary . 

,, explosive 

Boiling- Specific 
point. Heat. 
o 7H ,, r 0-017401 ,, , 
^ 7H 6 \0-01895J di Ud 
. 0-0170 


1300 0508 

1610" 0-0308 


62 -06 to 124 -12 

75 '09 150 -18 to 300 36 
120 30 120-3 to ? 
208-1 208-1 to ? 

(Jiioup XL (Oxygcjji, Chromium). Molybdenum, 
Tungsten, Uranium. 

The resemblance between oxygen and the other four members 
of this group is a slight one. It is advisable to consider oxygen 
along with the three elements sulphur, selenium, and tellurium, 
with which it displays much greater analogy. 

The elements molybdenum, tungsten, and uranium present 
some analogy with chromium, both in their properties as well as 
in the compounds which they form. But chromium is best con- 
sidered along with aluminium, iron, and manganese. 

Sources. The chief source of molybdenum is the sulphide, 
molybdenum glance, or molybdenite, and widfeuite, a compound of 
molybdenum, oxygen, and lead. These are rare minerals ; an 
alloy of lead and molybdenum has also been found native in the 
State of Utah. 

Tungsten occurs in ivolfram, combined with oxygen, iron, and 
manganese; and in scheelite, with oxygen and calcium 

Uranium is chiefly found as pitchblende, in combination with 


Preparation. Molybdenum* is prepared by heating its 
chloride to bright redness in a tube through which a stream of 
hydrogen gas is passed. The hydrogen unites with the chlorine, 
forming gaseous hydrogen chloride, leaving the non-volatile 
molybclenum. It may also be obtained by heating its oxide with 

Tungstenf can be prepared in a similar manner, or from its 
oxide by the action of hydrogen ; the hydrogen removes the 
oxygen as water, which passes off as gas, whilo the metal remains. 

Uranium J is best got from its chloride by heating it with 
metallic sodium in an iron crucible. The sodium and chlorine 
unite, forming common salt, while the uranium, which does not 
unite with sodium, sinks to the bottom of the crucible, being 
heavier than the fused salt. 

Properties. These elements all possess metallic lustre, 
Molybdenum is a brittle silver-white substance, exceedingly 
hard. It fuses at a high temperature. Tungsten is a steel -grey 
crystalline powder, which fuses at a white hent. Uranium is a 
black powder which is fusible to a grey metallic button of great 

These metals do not combine with oxygen at the ordinary tem- 
perature, but are converted into chlorides when thrown into chlorine 
gas in the state of powder. At a high temperature they burn in 
air, forming oxides. They also unite with sulphur at a red heat. 
They are unchanged by water at the ordinary temperature. 

Physical Properties. 

Mass of 1 c.c 










Molybdenum. , 


White heat.. 




Tungsten.. . . 

19 2 

While heat. . 

OH 34 




(at 12). 

Ur/mium . . . 

18 7 

Red heat . . . 




(at 4). 

GROUP XII. Oxygen, Sulphur, Selenium, Tellurium. 

These elements all occur native, as well as in combination. 
The first is a gas ; the other three are solids at the ordinary 
temperature, and are often associated with each other. 

* Debray, Compte* rend , 46, 1098. 

t Roscoe, Chem. Soc J., 10, 286. 

J Pehgot, Annales (3), 6, 53 ; 12, 549. n 


Sources. Oxygen is the most abundant and widely distri- 
buted of the elements, forming, as has been estimated, 50 per cent, 
of the earth's crust. About one-fifth of the weight as well as of the 
volume of air consists of oxygen, the remaining four-fifths being 
nitrogen, with which the oxygen is mixed. It constitutes 'eight- 
ninths of the weight of water, and is found in union with every 
element in nature, except fluorine, chlorine, bromine, platinum 
and its analogues, and gold, silver, and mercury. Many compounds 
into which it enters have been already mentioned as sources of the 

Sulphur occurs native in the neighbourhood of volcanoes, and 
coats the surface of the soil in districts of volcanic activity. It is 
chiefly mined in Italy and Sicily. It also occurs m combination 
with iron as iron pyrites, and with iron and copper as copper 
pyrites ; with lead as galena, with zinc as blende, with mercury as 
c>nnabar. It also occurs in union with oxygen and a metal, e.g., 
in the sulphates of calcium, magnesium, sodium, &c. Its principal 
sources are native sulphur; and copper pyrites, of which large 
mines exist in the South of Spain. It exists also in certain 
volatile oils, such as oil of mustard, oil of garlic, &c. 

Selenium in small quantities almost invariably accompanies 
sulphur ; both native and in its compounds. It is also, but rarely, 
found in combination with lead and copper ; and with nickel, 
silver, molybdenum, &c. 

Tellurium is found in the free state, and also in combination 
with bismuth, silver, lead, and gold. It is a very rare element. 

Preparation. There is no convenient method of separating 
nitrogen from air; hence pure Oxygen, unlike pure nitrogen, cannot 
be directly prepared from that source. Owing to its tendency to 
unite with almost all elements, it cannot well be prepired by dis- 
placing it from any one of its compounds. The only elements 
capable of displacing it appear to be fluorine and chlorine, for 
almost all other elements combine directly with it. It must 
therefore be prepared by heating certain of its compounds with 
other elements certain oxides and double oxides or salts ; or by 
the electrolysis of certain of its compounds, e.g., water. The 
methods of preparing it may be grouped under three heads : 

1. The electrolysis of water, or of fused oxides or hydroxides, i.e., 
oxides of hydrogen and another element. Water, however, is a 
non-conductor of electricity when pure, and it is necessary, in 
order to make it conduct, to dissolve in it some substances with 
which it reacts. In practice, the operation is conducted as follows : 
A U'tuke, ^ * s filled vriili dilute sulphuric acid. Through the 



lower end of each of tLese tubes is sealed a piece of platinum 
wire, connected each with a slip of platinum foil, and the pieces 
of wire projecting outside are connected by copper wires to the 
poles of a battery of four Bunsen's or Groves' or bichrome elements 
(two are sufficient, but the operation is more rapid with four cells). 
The oxygen is evolved from the electrode connected with the car- 

- ',"" .: "- . - 

bon or platinum plate ; the gns issuing from the other electrode is 
hydrogen. After the current has passed for some time, the tube o 
is partly filled with oxygen gas, and the hydrogen in the tube h 
occupies about twice the volume of the oxygen. On opening the 
stopcock o carefully, the characteristic property possessed by 
oxygen of rekindling a glowing piece of wood may be shown by 
allowing the escaping gas to play on it; and the hydrogen may be 
set on fire as it escapes from the tube h. 

2. The heating of certain oxides. All compounds with oxygen 
of the metals of the platinum group ; of gold, silver, or mercury ; of 
the chlorine group of elements; of the higher oxides. of nitrogen , the 
higher oxides of the chromium group of elements (e.g., chromium 
trioxide, chromates, potassium ferrate, manganate or permanga- 
nate, manganese dioxide, nickel and cobalt sesquioxides) ; of the 
calcium group of elements, and of lead; all these part with oxygen 



at a bright red heat, and in many cases at a lower temperature. 
The action of sulphuric acid on the higher oxides also yields oxygen 
(see Manganese Dioxide, and Chromate of Potassium) . 

Three typical examples are chosen : 

(a). The action of heat on mercuric oxide. The apparatus is 
shown in fig. 9. A is a tube of combustion glass, which is more 
difficult of fusion than ordinary glass, sealed at one end, and closed 
at the other end with a perforated indiarubber cork through which 
a bent glass tube is inserted. This tube dips below the surface 
of the water in a glass trough, K, and its open end bends upwards, 

FIG. 9. 

so as to deliver gas into an inverted jar, D, full of water. The 
hard glass tube contains some mercuric oxide. Heat is applied 
with a Bunsen's burne'vB, care being taken to wave about the 
flame at first, so as to heat the glass tube gradually ; else it is apt 
to crack. After allowing some bubbles to escape, so as to ensure 
the expulsion of air from the tube, the glass jar is placed above 
the exit tube, and the gas is collected. The mercury collects in 
the depression C. It was by this means that Priestley, one of the 
discoverers of oxygen, first prepared it in 1774. He named it 
dephlogisticated air (see p. 11). 

(#). The action of heat on potassium chlorate. This body is a 
compound of potassium, chlorine, and oxygen. The oxygen is 
wholly expelled, potassium chloride, a compound of chlorine and 
potassium, remaining behind. The chlorate is heated in a hard 
glass flask, by aid of a Bunsen burner (see Potassium Cltlorate, 
p. 466). The salt melts and begins to froth, owing to the evolution 
of oxygen. If some manganese dioxide be mixed with the chlorate, 
the gas is evolved at a lower temperature, but is not so pure 
(see Perchlorates ; also Oxides of Manganese) * 

(c). The action of heat on bnrium dioxide. Barium forms two 
oxides, one, the monoxide, containing less oxygen than the second, 
* Chem. Soc. J. t 51, 274. 



the dioxide. When the monoxide, a grey porous solid, is heated to 
dull redness in contact with dry air, it absorbs oxygen from the 
air, producing the dioxide ; the absorption is increased by pressure. 
On decreasing the pressure, the dioxide formed is decomposed : 
the oxygen may be pumped off by means of an air-pump and 

FIG. 10. 

forced into iron or steel bottles. This process is now carried out on 
a large scale, and indeed is the only method by which oxygen is 
made commercially. The barium dioxide is contained in vertical 
iron tubes, which are heated with gas from a Siomens's " pro- 
ducer," the temperature being carefully regulated. 

3. By displacement. The gaseous element fluorine, which 
has only recently been prepared, at once acts on water, combining 
with its hydrogen, and liberating its oxygen (see Ozone, p. 387). 
Chlorine and steam at a red heat react in a similar manner, 
hydrogen cbloride and oxygen being produced. Chlorine gas 
also slowly acts on water exposed to sunlight, liberating oxygen. 
The halogens expel oxygen from certain oxides at a red heat; 
e.g'., from the oxides of lead, bismuth, zinc, &c. None of these are 
practical methods of preparing the gas (see Oxides of Manganese, 
Chlorine, Silver, and Lead ; also Hi/pochlorites). 

Sulphur. Sulphur may be prepared (1) by heating certain 
sulphides, e.g., those of gold and platinum, which part with their 
sulphur, leaving the metal; or by heating hydrogen sulphide, which 
splits into sulphur and hydrogen ; and (2) by heating certain per- 
sulphides (compounds of metals with sulphur which form more 
than one sulphide), the most important of which is iron pyrites. As 
sulphur combines directly with most other elements, there are few 
methods of preparing it by displacing it from its compounds ; yet 
chlorine, bromine, or iodine dissolved in water combines with the 


hydrogen of hydrogen sulphide in preference to the sulphur, so that 
the element is liberated (see also Poly sulphides of Sodium). 

The elements selenium and tellurium are most readily.pre- 
pared by displacement. The compounds which they form^with 
oxygen are decomposed by sulphur dioxide, which absorbs their 
oxygen, itself changing to sulphur trioxide, and liberating the 
selenium or tellurium (see Selenium and Tellurium Dioxides). 
Their compounds with hydrogen, dissolved in water, arc decom- 
posed by atmospheric oxygen, the element falling to the bottom of 
the solution. 

An important source of sulphur is native sulphur, of which the 
chief impurity is earthy matter. The modern method of extraction 
is to melt it under water in a boiler by forcing in steam until the 
pressure rises to 25 Ibs. on the square inch. The temperature of 
the water is thus raised to over 115, the melting point of sulphur. 
The melted sulphur is run off through a stop-cock in the side of 
the boiler, and when cold a fresh charge of impure sulphur is 
introduced, and the operation repeated. Sulphur is usually brought 
into commerce in the form of sticks cast in wooden moulds, and is 
in this form named roll sulphur. 

Sulphur is a by-product in the manufacture of alkali, being 
obtained from iron or copper pyrites (see Chapter XXXIX). 

Selenium is best obtained from certain residues in the manu- 
facture of sulphuric acid, by treating them with nitric acid, and 
then precipitating the selenium with sulphur dioxide. 

Tellurium may be purified by distilling native tellurium at a 
red heat in a current of hydrogen gas. It is precipitated from its 
solutions by metallic zinc. 

Properties. Oxygen is a colourless, odourless, tasteless gas, 
somewhat heavier than atmospheric air. It is very sparingly 
soluble in water; 100 volumes of water at 4 dissolve 3'7 volumes 
of oxygen. It has been condensed to a colourless transparent 
liquid by application of great pressure at a very low temperature, 
and when still further cooled, it freezes to a white crystalline solid. 
It was discovered independently by Priestley and by Scheele (see 
p. 11) in 1774 and 1775; it had previously, however, been recog- 
nised as a distinct "air" or gas by Mayow, about 1675 (see p. 9). 
Its true nature was made public by Lavoisier, as has already been 
noticed, although Mayow anticipated him in most of his con- 
clusions. Its name, " acid-producer " (ofvs <yewaw), was invented 
by Lavoisier. 

Oxygen combines directly with all elements except the halogens, 
gold, and some metals of the platinum group. Silver and mercury, 


although they do not readily combine directly with oxygen, can be 
made to unite under pressure. Many elements, such as carbon, 
sulphur, and certain metals, do not unite with oxygen except when, 
heated ; others, such as sodium, phosphorus, &c., combine at the 
ordinary temperature. The word " combustion " usually signifies 
union with oxygen, with evolution of light. All substances which 
burn in air burn with increased brilliancy in oxygen gas. It is 
respirable; in its usual dilute state in air, it is when breathed 
absorbed by the blood, and serves to oxidise the carbon and. 
hydrogen in the body, thereby generating animal heat ; if breathed 
in a pure state, however, oxidation takes place with too gremt 
rapidity, and acute febrile symptoms are produced after a short 
time, followed by death unless the animal is allowed to respire air. 
The respiration of fishes is sustained by the small amount of oxygon 
dissolved in the water in which they exist. 

When an electric discharge is passed through oxygen, or when 
the element is liberated by the action of fluorine on water, a portion 
of it is changed to an allotropic form, which from its strong smell 
has been named ozone (ofci*/, to smell). This substance will be 
considered as an oxide of oxygen, and is described on p. 387. 

The remaining three elements of this group, sulphur, selCr 
nium, and tellurium, form a well-marked series. They show 
progression in their atomic weights : thus S = 32, Se = 79, 
Te = 125. The atomic weight of selenium is nearly the mean of 
those of sulphur and tellurium. They show a gradation of colour: 
sulphur is yellow, selenium red, and tellurium metallic. Sulphur 
is practically a non-conductor of electricity, selenium conducts 
when exposed to light, and tellurium is a conductor. No allo- 
tropic form of tellurium is known. Selenium is known to exist in 
three forms : amorphous, which changes to crystalline when fused 
and kerjt for some time at 210; this crystalline variety is insoluble 
in carbon disulphide ; the amorphous variety, produced by precipi- 
tating selenium with sulphur dioxide, is a bright-red powder, soluble 
in carbon disulphide, from which it deposits on evaporation in dark 
red crystals. Sulphur crystallises in two distinct forms : rhombic 
crystals (fig. 11), which are found native, and which may be arti- 

FIGK 11. FIG. 12. , 

F 2 


ficially produced by crystallising sulphur from carbon disulphide ; 
and monoclinic needles (fig. 12), which may be prepared by melting 
sulphur, allowing it to cool till the surface has solidified, breaking 
the solid surface layer, and pouring out the liquid. The interior of 
the mass is filled with crystals. The monoclinic form also deposits 
from a solution of sulphur in ether or in benzene. The monoclinic 
form is less stable than the rhombic ; and the crystals, which are 
clear, transparent, and brownish-yellow, soon become opaque on 
standing, changing spontaneous!}' into a mass of minute rhombic 
octohedra. This change is accompanied with evolution of heat. 
Other crystalline forms have recently been obtained. 

Selenium or sulphur, when distilled into a large chamber* 
condenses in part as a fine powder, named flowers of aulphur, or of 
selenium. This is really a mixture of two varieties, one of which 
is insoluble in carbon disulphide, the other soluble. 

Again, in the state of liquid, sulphur exhibits allotropy. It 
melts at 115 to a clear, pale yellow, mubile liquid. At 200 it 
turns brown and viscous. When the first variety is poured into 
water, it at once solidifies to ordinary brittle crystalline sulphur, 
soluble in carbon disulphide. The viscous variety, however, if 
poured into water, changes to a curious elastic, indiarubber-like 
substance, insoluble in carbon disulphide, which only slowly regains 
its former condition. Between 400 and 446, its boiling point, 
sulphur again become" "mobile, still remaining brown. A variety 
of sulphur soluble in water has recently been discovered.* Sulphur 
produced by precipitation has a white colour, and its mixture with 
water is known as milk of sulphur. 

In the gaseous state also, sulphur displays allotropy. Its 
density at low temperatures implies a high molecular weight, but 
at high temperatures the molecule is simpler and weighs less 
(see p< 614). 

These elements unite directly with oxygen, burning in the air 
when heated; they also combine with each other, with the halogens 
(the compounds with bromine and with iodine are ill-defined), and 
with all other elements except nitrogen, when heated in contact 
with .them. They are without action on water at the ordinary 

Chem. Soc. J., 63, 283. 



Physical Properties. 

Mass of 1 c.c. 



Sulphur (rhombic) * . 2 *07 at 

(monoclinic) 1 * 98 

(plastic) l-95atO 

Selenium, crystallised 4*4 

from fusion 

Selenium, crystallised 4 -8 at 15 
from solution 

Selenium, amorphous 4 '3 

Tellurium 6 *23 at 

Liquid. Gas. 

1-24 at -200 0-001429 
(at and 
760 mm.) 


H = 1. 


32 -27 
(at 1040) 


(at 1040) 

(at 1440) 


Sulphur (rhombic) . . 


Selenium, crystallised 

from fusion 
Selenium, crystallised 

from solution 
Selenium, amorphous 

Melting- Boiling- 
point. point. 
Below -212 -186 
115 446 








16 '0 


217 W 


0-0746 79 '0 

About 100 

Below Bright red 0'0483 125*0 
redness heat (crystalline) 











Vapour Pressures of Oxygen at different Temperatures.* 

-118-8 (crit.) -121-6 -125 -6 -129'0 

Temperature .... 
Pressure, atmos. . . 

Temperature . . . 
Pressure, atmos. 

50-8 (crit.) 


-146-8 -155-6 
13-7 8-23 






-181-5 -190 -192-71 -196-2 -198-7 -200-4 -211-5 
Pressure, mm. 

740 160 71 50 20 20 9 

These low temperatures were produced by the evaporation of liquid ethylenc 
under reduced pressure. The mass of 1 c.c. of oxygen at its boiling point, 
181-4, under a pressure of 742'1 mm. was found to be T124 grams. 

* Comptes red. t 98, 982; 100, 350, 979; 103, 1010. 


Vapour Densities of Sulphur, Selenium, and Tellurium. These are dis- 
cussed on p. 614. 

Appendix. Air. Air is not a chemical compound, but a 
mere mixture of nitrogen and oxygen gases. That this is the case 
is shown by the following considerations : (1.) There is no heat 
change on mixing oxygen and nitrogen gases ; when a compound 
is formed, heat is usually evolved. (2.) The density of air is the 
mean of the densities of the constituent gases ; its refractive index 
for light is also the mean of those of oxygen and nitrogen taken in 
the proportions in which they occur in air ; and so with other 
physical properties. Such properties, possessed by a compound, 
are not the mean of those of its constituents. (3.) Oxygen is 
more soluble in water than nitrogen. On saturating water with 
air, oxygen dissolves in greater amount than, nitrogen ; and the 
gas evolved from the water when it is heated contains a larger 
proportion of oxygen than does air. (4.) There is no simple rela- 
tion between the atomic proportions of the nitrogen and oxygen in 
the air. Such a relation would be characteristic of a compound. 
The actual composition by weight is, approximately, nitrogen 
= 77 per cent. ; oxygen = 23 per cent. Dividing these numbers 
by the atomic weights of nitrogen and oxygen respectively, 14 and 
16, we obtain the quotients 5*55 arid 1'44, representing the relative 
number of atoms of nitrogen and oxygen in air. The simplest 
proportion between thes'? numbers is 3'85 to 1 ; although the ratio 
approximates to the ratio 4:1, yet it is far from being a simple 
one, as it would be, were air a compound. A substance of the 
formula N^O would contain 77'8 per cent, of nitrogen and 22*2 per 
cent, of oxygen. 

Air contains, in addition to nitrogen and oxygen, water-vapour 
(about 0*84 per cent, by weight, or 1*4 per cent, by volume, on the 
average), carbon dioxide, from 0*049 to 0'033 per cent, by volume, 
and a few parts of ammonia per million. Subtracting these from 
air, tLe ratio of oxygen to nitrogen by volume approximates to 
79*04 volumes of nitrogen, and 20'96 volumes of oxygen. But 
its composition varies slightly in different places and at different 
times, although the action of air currents and winds tends to 
keep it nearly constant. 

Air has been liquefied by cooling to 192 ; but, as oxygen and 
nitrogen have not the same boiling points, the less volatile oxygen 
doubtless liquefies first. 

Air is analysed (1) by mixing a known volume with a known 
volume of hydrogen, and exploding the mixture. The oxygen is 
withdrawn as water, and the residual nitrogen is measured. 

AIR. 71 

2. By passing air deprived of carbon dioxide and moisture over 
ignited copper. The oxygen unites with the copper, forming 
oxide ; and its amount is ascertained by the gain in weight of the 
copper; the nitrogen passes on into an empty globe, previously 
weighed. The gain in weight of the globe gives the weight of 


Equations expressing the preparation of elements of Groups IX, X, XT, and 

Nitrogen. WS^s - N 2 + 3H 2 . 

2NH 3 + 3C1 2 = N 3 4- 3HCL 
4NH 3 4- 30 2 = 2N 2 4- 6H 2 0. 
Vanadium. 2VC1 3 + 3H 2 = 2V + 6HC1. 
Phosphorus. 4HPO 3 + 120 = P 4 + 2H 2 + 120O. 
Arsenic. FeSAs = As + FeS. 
Antimony. Sbfiz 4- 3Fe = 2Sb -I- 3FeS. 
Molybdenum. MoCl 4 4- 2H 2 = Mo 4- 4HC1. 
Tungsten. W0 3 + 3H 2 = W 4- 3II 2 O. 
Uranium. UC1 4 + 4Na = U 4- 4NaCL 
Oxygen. 2H 2 O = O 2 4- 2iI 2 . 

2HgO == O 2 4- 2Hg. 

2KC1O 3 = 30 2 4- 2KC1. 

2BaO 2 = 2 4- 2BaO. 

H 2 + C1 2 = O 4- 2HC1. 
SuIpkur.2An 2 S 3 = 3S 2 4- 4Au. 

2FeS 2 - S 2 4- 2FeS. 

2H 2 S 4- 2 = S 2 4- 2H 2 0. 
Selenium. Se0 3 4- 2S0 2 4- 2H 2 = Se 4- 2H 2 SO i . 




GROUP XIII. Fluorine, Chlorine, Bromine, Iodine. 

The elements fluorine and manganese present little, if any, 
analogy. Hence, just as oxygen is best classified along with 
sulphur, selenium, and tellurium, presenting little, if any, analogy 
with chromium, so with the elements of this group, manganese and 
fluorine having little or nothing in common. Both chromium and 
manganese, it will be remembered, are most conveniently classed 
with iron, nickel, arid cobalt. 

The halogens, as these elements are called, are strikingly like 
each other. They have all low boiling and melting points ; and 
they all combine readilv, with other elements, oxygen and nitrogen 
excepted. They all are found in combination ; free iodine has 
been found in the water from Woodhall Spa, near Lincoln.* 

Sources. Fluorine occurs in nature, combined with calcium, 
influor spar or Derbyshire spar ; in cryolite, combined with sodium 
and aluminium ; and sometimes in apatite, a compound of phos- 
phorus, oxygen, and calcium calcium phosphate, combined with 
calcium fluoride. It occurs in small quantity also in the enamel 
of the teeth and in the bones : it is very widely distributed in 
nature, although not very abundant. 

Chlorine, bromine, and iodine are all contained in sea- 
water, in combination with sodium, potassium, and magnesium. 
Chlorine also occurs in rock salt, of which large mines exist in 
Cheshire, and in the neighbourhood of the Tyne, in Northumber- 
land. Certain rare ores of silver and mercury contain these metals 
in combination with chlorine, bromine, and especially with iodine. 
The chief source of bromine and iodine is sea-weed, which 
absorbs the compounds of these elements from sea-water.f Iodine 
is also largely obtained from Chili saltpetre, or sodium nitrate, 

Chem News, 64, 300. f Dingl, polyt. J., 126, 85. 



which contains it in small amount in combination with oxygen and 
sodium as sodium iodate.* 

Preparation. 1. By electrolysis. This is the only method of 
preparing fluorine ; liquid compounds, or solutions of compounds 
in water, of the other halogens also yield the elements by this 
process. The preparation of lithium by the electrolysis of its fused 
chloride (see p. 29) affords an example of the application of elec- 
trolysis to a fused compound of chlorine. The gas is evolved at the 
positive or carbon pole, the metal being deposited on the negative or 
zinc pole. Concentrated solutions in water of chlorides, bromides, 
or iodides yield these elements on electrolysis, for such solutions 
conduct electricity, and the halogens, with exception of fluorine, 
are not readily acted on by water. Thus hydrogen chloride dis- 
solved in water (hydrochloric acid) splits into chlorine and hydro- 
gen gases when electrolysed. The poles should consist of gas 
carbon, or platinum, all other substances being attacked, more or 

FIG. 13. 

less, by the chlorine produced. Fluorine, however, cannot be 
liberated in presence of water, for it immediately acts on it, liberat- 
ing oxygen as ozone. Hence, it is prepared by electrolysing in a 

* Dingl. polyt. J. t 253, 48. 


U-tube made of an alloy of platinum and iridium, which is but 
slightly attacked, a solution of potassium fluoride in dry hydro- 
fluoric acid cooled to a low temperature. It is necessary to use 
such a solution, inasmuch as pure hydrogen fluoride does not conduct 
electricity, and unless the liquid conduct, electricity cannot pass, and 
electrolysis cannot take place. The apparatus used by M. Moissan,* 
who has recently isolated this element, is shown in fig, 13. 

2. By displacement. No element appears capable of displacing 
fluorine from its compounds without combining with it. But 
chlorine, bromine, and iodine are usually prepared by displacing 
them from their compounds with potassium, sodium, magne- 
sium, &c., by means of oxygen. The oxygen, however, is not 
employed in the gaseous state. At a red heat, indeed, such dis- 
placement is possible. The action of oxygen, for instance, on red- 
hot magnesium chloride yields chlorine, while a double compound 
of chlorine, oxygen, and magnesium (oxychloride) remains behind : 
again, Deacon's process for producing chlorine, which depends on 
the interaction of the oxygen of the air and hydrogen chloride at 
330 in presence of bricks coated with dry copper chloride, yields 
chlorine and water as products. The usual method, however, of pre- 
paring halogens consists in acting on hydrogen chloride dissolved 
in water (hydrochloric acid) with a peroxide. The peroxide 
yields a portion of its oxygen to the hydrogen chloride, forming 
water and chlorine. "The remaining hydrogen chloride sub- 
sequently reacts with the oxide. The peroxide generally used 
is manganese dioxide ; but peroxides of lead, barium, &c., potas- 
sium permanganate, bichromate, and other peroxides may also be 
employed. When a mixture of chloride, bromide, and iodide of 
potassium or sodium is treated so as to liberate the halogens, 
the iodine is liberated first, then the bromine, and lastly the 


3. By Ji eating compounds of the elements. Most of the com- 
pounds of the halogens are remarkably stable, and although some, 
such as hydrogen chloride, may be decomposed by exposure to an 
exceedingly high temperature, yet re. combination takes place on 
cooling, so that the halogen cannot be isolated. Fluorine, how- 
ever, is said to have been prepared by heating cerium or lead tetra- 
fluoride.f The chloride, bromide, and iodide of nitrogen are ex- 
tremely explosive bodies, at once decomposing into their elements 
when warmed or when exposed to shock ; the higher chlorides 
and bromides of phosphorus and arsenic, when heated, yield lower 

* Comptes rendu*, 102, 1543 j 103, 202 and 256. 
t Berichte, 14, 1944. 



compounds and the halogens; compounds of halogens with 
oxygen are also unstable, and are resolved with explosion into 
their, elements when heated ; compounds of the halogens with 
each pther are also easily decomposed by heat. The halogen 
compounds of gold, platinum, &c., are decomposed into their 
elements by heat. This type of reaction, however, does not afford 
a practical method of preparation. 

Preparation of chlorine. About 30 grams of powdered man- 
ganese dioxide are placed in a flask closed with a double bored 
cork ; through one hole passes a tube communicating with a wash- 
bottle full of water; through the other a thistle funnel passes. 
Strong hydrochloric acid (solution of hydrogen chloride in water) 

FIG. 14. 

is added, and gentle heat is applied. The gas issues from the exit 
tube of the wash-bottle, and may be collected over warm water, in 
which it is less soluble than in cold ; or, better, by downward dis- 
placement, for it is heavier than air. The latter arrangement is 
shown in the figure. To show the tendency of chlorine to com- 
bine with other elements, powdered antimony may be thrown into 
a jar containing it ; the metal will burn. A candle will be found 
to burn in chlorine with a sooty flame ; the hydrogen combines, 
but the carbon is liberated as soot. A solution of chlorine in 
water acts as a bleaching agent : a coloured rag dipped in such a 
solution is soon bleached ; the chlorine combines with the hydro- 
gen of the water, liberating oxygen, which oxidisee 


substances to colourless ones. Lastly, some chlorine-water, as a 
solution of chlorine in water is termed, added to a solution of a 
bromide or iodide, e.g., potassium bromide or iodide, liberates these 
elements. Similarly, bromine-water, added to an iodide, liberates 

Properties. In the gaseous state these elements have all a 
strong disagreeable smell ; that of fluorine, however, is the smell 
of ozone, for it acts on the moisture in the nose, liberating ozone. 
Fluorine appears to be colourless, clllorine is greenish-yellow, 
bromine deep red, and iodine violet. The names %Xw/)os, 
yellowish -green, ppfe/mo?, a stench, and coet^t, violet, refer to these 
properties. As it is impossible to confine fluorine in any vessel 
which ifc does not attack, no attempt to liquefy it Was been made. 
Chlorine may be condensed to a greenish-yellow liquid, boiling at 
a very low temperature ; it solidifies to a solid of the same colour.* 
Bromine is at ordinary temperatures a deep brownish-red liquid, 
freezing to a blackish-red solid ; and iodine at ordinary tempera- 
tures is a bluish-black lustrous solid, melting to a brownish-black 
liquid. It sublimes readily. 

Chlorine, bromine, and iodine dissolve in carbon disulphide and 
tetrachloride, in alcohol and ether, and also in water. One volume 
of water at absorbs 2'58 volumes of chlorine ; and at 15, 
2*36 volumes. Bromine is soluble in 30 times its weight of water 
at 10 ; iodine is ve/y sparingly soluble in pure water. The 
presence of chlorides, bromides, and iodides in the water greatly 
increases the solubility of the halogens: it is possible that the 
solubility of chlorine and bromine in water depends on their 
interaction with the water. Chlorine and bromine combine with 
water to form crystalline hydrates. Bromine and iodine form 
compounds with starch; the former has an orange colour, the 
latter is deep blue. The compound of iodine with starclj is used 
as a delicate test both for iodine and for starch. 

These elements combine directly with each other, and at a 
high temperature, or when moist, with all others except carbon, 
nitrogen, and oxygen. f The only solid elements which withstand 
their action even partially are carbon, and iridium, or better, its 
alloy with platiaum. Fluorine attacks glass and porcelain, but the 
other halogens are without action on these substances, and may be 
exposed in glass or porcelain vessels to a high temperature. 

All these elements tend to combine with hydrogen, whether 
free or in combination, hence their irritating action on the 

* Monatsh. Chem., 5, 127. t Chem. Soc. 7., 43, 153. 


organism, which chiefly consists of compounds of carbon, hydrogen, 
and oxygen. They produce catarrh of the mucous membranes 
when, breathed. 

No allotropic modifications of these elements are known. 

Physical Properties. 

Mass of Ice. 

f * ^ Density, Melting- 
Solid. Liquid. H = 1. point. 

Fluorine ? ? 18 *3 at 15 . . ? 

Chlorine ? 1 '33 at 15 '5 35 '4 at 200 Below - 102 

Bromine ? 3 '18 at .. 80 "0 at 228 -7 '05 

Iodine 4 '95 4 '00 at m. p. 128 '85 at 445 114 15 C 








35-46 70-92 

58 -75 

0-0843 solid 






Iodine 184 '35 '0541 .... 126 '85 126 '85 253 '7 

Vapour-densities of Chlorine, Bromine, and Iodine. These are considered 
on p. 616. 

GROUPS XIV AND XV. Ruthenium, Rhodium, Pal- 
ladium; Osmium, Iridium, Platinum. 

These metals are always associated. They fall into two groups 
of three, members of the first of which have atomic weights of 
about 105, and of tlie second, about 193. They are always found 
native, or in combination with each other. They are very difficult 
of attack by any process : even chlorine or oxygen at a red heat 
produces little effect ; hence their occurrence in the free state. 

' Sources. (a). Metallic particles, consisting chiefly of plati- 
num and palladium, but containing small quantities of the other 
metals, occur as flattened grains in the sand of certain rivers in 
Brazil, Mexico, California, and on the west side of the Ural 

(b). Metallic particles, chiefly consisting of osmium and 
iridium, and named osmiridium, occur along with the grains of 
platinum. The complete separation of these metals is a matter of 
great difficulty (see p. 475).* 8,000 kilos, of platinum ore were 
exported from the Ural Mountains in 1881. 

* Consult Annales (3), 66, 1 and 385; also Chem. News, 39, 175. 



Properties. These elements are all white, with a greyish 
tinge, and possess strong metallic lustre. They melt only at a 
very high temperature ; in practice they are fused by meana of a 
blowpipe flame of oxygen and hydrogen in crucibles of linpe, on 
which they are without action (see fig. 31, p. 194). Owing to 
their ability to resist oxidation, an alloy of 90 per cent, of plati- 
num and 10 per cent, of iridiam is used for crucibles, retorts 
for evaporating oil of vitriol, &c., and for standards of length (e g., 
the French standard metre). The alloy of osmium and iridium, 
owing to its extreme hardness, is employed in tipping gold pens, 
and as bearings for very delicate wheel work. These alloys are 
very costly, which somewhat limits their use. The metals can be 

Platinum and palladium form compounds with hydrogen, in 
which the last element appears to play the part of a metal in an 
alloy (see Alloys, p. 576). 

The name platinum, signifying "little silver," was given to the 
metal by the Spaniards* The name rhodium refers to the red 
colour of its salts. The other names are fanciful, except osmium, 
so called from O<T/I^, a smell, in allusion to the strong odour of its 
volatile oxide. 

Allotropic forms of iridium, ruthenium, and rhodium have been 
prepared by fusing the metals with zinc or lead, and subsequently 
dissolving out the zi'xic or lead with an acid.* The iridium, 
ruthenium, or rhodium is left as a black powder which explodes 
on gently warming, being converted into the ordinary form of the 
metal. Osmium, iridium, and platinum are the heaviest substances 
known, being more than 21 times as heavy as water. 

Physical Properties. 

Mass of 1 c.c. Melting- Specific Atomic Molecular 

Solid. point. Heat. Weight. Weight. 

Ruthenium 12 -26 at 0. . -0611 101 '66 ? 

Rhodium 12'lat? .. 0'0580 103 '0 ? 

Palladium 11 '4 at 225 0-0593 106 '35 ? 

10 '8 (liquid) 

Osmium 22'48at?.. 0'0311 191 '3 ? 

Iridium 22 '42 at 17 '5 - 0-0326 193-0 ? 

Platinum 21 '50 at 17 '6 1700? 0'0324 194' 3 ? 

18 '91 (liquid) 

* Debray, Comptes rend., 90, 1195. 


GROUP XVI. Copper, Silver, Gold, Mercury. 

Of these elements copper, silver, and gold probably belong to 
the same group : owing to considerable resemblance which mer- 
cury bears to them in its compounds, it is convenient to include it 
in the group. 

Sources. All these metals are found native, for all can resist 
the action of oxygen at the ordinary temperature. All occur, 
besides, in combination with sulphur and with arsenic. The chief 
ore of copper is copper pyrites, in which it is combined with iron 
sulphide and sulphur ; and other important ores are the oxide, 
cuprite, or red copper ore, and the sulphide, copper-glance ; besides 
these, it is found in two forms in combination with carbon and 
oxygen as carbonate, viz., malachite and azurite. 

Silver is mostly found native. But silver-glance or sulphide, 
pyrargyrite, proustite, and silver-copper-glance, in which it is associ- 
ated with sulphur, antimony, arsenic, and copper, are also impor- 
tant, and it also occurs in combination with the halogens. Th6 
chloride is named horn-silver. 

Gold chiefly occurs native, forming veins and nuggets in 
quartz-rock; but it also accompanies copper and silver as arsenide 
and sulphide ; and is sometimes associated with tellurium and 
bismuth. The chief mines are in California, Australia, and the 
Cape ; but it is now mined in Wales, and it is found in upper 
Lanarkshire, in the Leadhills. 

Mercury sometimes occurs free, but its most important ore is 
cinnabar, the sulphide, of which large mines are worked in Austria, 
Spain, China, and California. 

Preparation. The preparation of copper from ores in which 
it is not associated with sulphur is simple. The ore is powdered 
anji heated with some form of carbon. The carbon unites with the 
oxygen, forming gaseous carbon monoxide, and the copper fuses, 
and owing to its greater specific gravity settles at the bottom of 
the furnace. Copper oxide does not decompose by heat alone ; 
but when heated in an atmosphere of hydrogen it is " reduced," 
the hydrogen uniting with the oxygen to form water. 

If iu union with sulphur, one of two methods may be adopted : 
(1.) The sulphide of copper is roasted in air, whereby it absorbs 
oxygen, and is converted into sulphate of copper. This is some- 
times brought about by leaving the copper ore lying exposed to air 
for years. The sulphate of copper is treated with water, which 
dissolves it ; and on addition of scrap-iron, the iron replaces the 



copper in its compound with snlphur and oxygen, forming sulphate 
of iron, and the copper is precipitated as a metallic sponge. It is 
then collected, dried, and smelted. This is called the "wet" 
process of extraction. The latter part of this process may be shown 
on a small scale by dipping into a solution of copper sulphate a 
piece of bright iron, such as the blade of a knife ; it will soon 
become covered with a deposit of metallic copper. (2.) The dry 
process consists in roasting the ore : the iron contained in it com- 
bines with oxygen, the copper remaining as sulphide. The oxide 
of iron is made to unite with sand, or silica, forming a "slag," and 
by repeating this process several times the copper is finally 
obtained as sulphide. The sulphide is roasted ; both copper and 
sulphur are oxidised, and a reaction occurs whereby copper sepa- 
rates in the metallic state ; the oxygen of the copper oxide unites 
with the sulphur of the copper sulphide, forming sulphur dioxide, 
a gas, which escapes, while metallic copper remains. It is melted 
and brought to market (see Chapter XXXVIII). 

Copper chloride loses its chlorine when heated in hydrogen ; 
hydrogen chloride is formed, and the metal remains. 

Mercury is easily separated from the sulphur with which it ia 
combined in cinnabar, by roasting in specially constructed fur- 
naces ; the oxygen of the air unites with the sulphur, forming 
gaseous sulphur dioxide, and the mercury passes in the form of 
gas through a series rf cold chambers or earthenware pots, in 
which it condenses to the metallic state, while the sulphur dioxide 
escapes. This process may be illustrated by heating in a hard 
glass tube some powdered cinnabar and aspirating over it a 

FIG. 15. 


current of air. The metallic mercury will condense in the cold 
part of the tube in small globules, while the sulphur dioxide gas 
will be carried on into the aspirator (see fig. 15). 

^prcury can also be prepared by heating its oxide (see p. 491) 
although its compounds with the halogens also split into mercury 
and halogen when heated, yet they recombine on cooling ; hence 
the metal cannot be prepared by this method unless hydrogen, or 
some other metal, e.g., iron, is present to combine with the halogen. 

Mercury may be purified from iron, zinc, lead, and other metals 
accompanying it by distillation, preferably at a low pressure; and 
by drawing a slow stream of air for several days through an 
inclined tube containing the impure metal. 

Silver is extracted from its ores, in which it exists chiefly as 
sulphide, by roasting the ore with common salt, which is a com- 
pound of sodium and chlorine. The change is represented as 
follows : 

an. f Sodium , Sodium sulphide. 

balt \ChlormoI3II3 ~^ 

r Silver . . . ^ ::z::::> Silver chloride. 

Silver sulphide i g i v, ^^ 

The silver and chlorine combine, and the sulphur and sodium. 
Such a reaction is termed a " double decomposition." The next 
stage in the process is to mix the mass with water, and to add 
scrap-iron, rotating the mixture in wooden barrels. The chlorine 
and iron combine, the silver separating as a spongy mass. 
Mercury is added to dissolve the silver ; and after renewed rota- 
tion of the barrels, the mercury is drawn ofr, dried, and distilled. 
The volatile mercury distils away, leaving the much less volatile 
silver behind. 

SilveV is also largely extracted from lead ores and from copper 
ore's (see Chapter XXXVIII).* 

The process of extracting gold from gold quartz is a mechani- 
cal one, for the most part. The ore is stamped to fine powder in 
mills for the purpose, and washed with water. The fine powder 
is made to run over a runnel of copper, coated with mercury ; the 
sand is carried on, but the grains ot gold unite with the mercury, 
and are retained. The mercury is then squeezed through chamois- 
leather : the alloy (or amalgam) of gold and mercury is very 
sparingly soluble in mercury; hence it remains almost entirely 
behind. The mercury is then distilled off, and, along with that 

* For the preparation of pure silver, see Stas, Annalen, Suppl. 4, 168. 



portion which had passed through the chamois-leather, used for 
re-coating the copper plates. 

When the gold exists mixed with sulphides, these are roasted 
to remove sulphur and arsenic, which unite with the oxyoren of 
the air and volatilise away. The residue containing the gold is 
heated under pressure with chlorine- water, when the gold unites 
with the chlorine, going into solution as chloride of gold. The 
gold is then precipitated on metallic copper. 

The preparation of mercury, silver, and gold from the chlorides 
may he shown, (a) by placing a piece of bright copper in a solution 
of mercuric chloride; (b) by laying on the top of a bead of fused 
silver chloride a piece of zinc and adding a little hydrochloric 
acid ; (c) by placing a slip of clean copper foil in* a solution of 
chloride of gold. 

Properties. Copper is a red metal ; silver, brilliant white ; 
gold, yellow ; and mercury, white with a faint grey tinge. Mercury 
is liquid at the ordinary temperature, but freezes at 40 to a 
malleable solid. Silver is the most ductile of the remaining three 
metals ; gold is the softest, and the most malleable. Gold and 
silver leaf, used for gilding and silvering , are made by beating the 
metals into leaves with wooden mallets : when thin they are pro- 
tected from the direct blow of the mallet by layers of gold-beaters' 
skin. Copper may also be beaten or rolled into foil and leaf. 
Gold leaf transmits ft ,'oen light ; silver leaf, blue light ; and copper 
leaf, bluish or pink light. All of those metals conduct electricity 
well. Placing silver equal to 100 at 0, copper has a conductivity 
of 77-43 at 18 '8, and gold of 55'19 at 22 7; mercury follows with 
a conductivity of 1'63 at 22 8. 

Silver can be distilled at a white heat. Its vapour is bluish- 
purple. It has the peculiarity of dissolving oxygen (about 
22 times its volume) when liquid and discharging it during solidi- 
fication (''spitting"). 

Mercury distils about 358. Its vapour is colourless. 

Copper is rendered brittle by slow cooling, and is softened by 
rapid cooling. The properties of all these elements are very 
singularly modified by the presence of traces of others. Thus the 
smallest tr,ace of arsenic renders gold exceedingly brittle ; a trace 
of phosphorus in copper greatly increases its tensile strength ; a 
minute traco of zinc in mercury completely modifies its surface 

For the composition of coins, jewellers' metal, &c., see Alloys 
(p. 587). 

Of these elements, none is oxidised on exposure to air, but 


copper at a red heat, mercury at the temperature of ebullition, 
and silver heated in air under a pressure of several atmospheres 
unite wit'h oxygen. Gold does not combine directly with oxygon. 
The oxides of the last three are easily decomposed by heat. These 
metals all unite directly with sulphur, selenium, and tellurium, and 
with arsenic ; with chlorine, bromine, and iodine, and with most 
metals. They do not unite directly with nitrogen. 

Physical Properties. 
Mass of 1 c.c. 



..... 8 90 
(at 0) 

Silver 10-57 

G-old 19-29 

Mercury 34 -19 



Copper ? 

Silver White heat 

G-old ? 

Mercury 358 2 


H = 1. 











(at 0) 











107 -93 


032 fc 

197 22 

197 22 


200 2 



General Remarks on the Elements. 

(1.) Classification. It has been customary to divide tho 
elements into two classes : the metals, those which are opaque and 
which exhibit metallic lustre ; such elements are more or less good 
conductors of electricity and heat: and the non-metals, comprising 
the remaining elements. Such a division tends to obscure the rela- 
tions between them ; it is, so far as we know, an arbitrary division, 
and is sanctioned, only by long custom. Other objections which 
might be taken to this division are that a number of elements, such 
as titanium, arsenic, and tellurium, are difficult to classify, being* 
sometimes considered as metals, sometimes as non-metals : and a 
still more important objection is that certain elements can exist in 
both forms. Thus silicon, phosphorus, and carbon exist in com- 
pact crystalline forms, with dull metallic lustre, and are then 
conductors of electricity ; while they also exist in forms incapable 
of conducting, and without metallic lustre. Such reasons have 
led to the abandonment of this classification here. Still the name 
metal has generally been applied in this book to those, elements 

a 2 


which are usually ranked as such ; though it is to be understood in 
a loose, colloquial sense. 

It- will, however, be convenient to trive a list of non-metals, so 
that the old classification may be understood. , 

List of non-metals.' Hydrogen (?), boron, carbon, silicon, 
titanium (?), zirconium (?), nitrogen, phosphorus, vanadium (?), 
arsenic (?), antimony (?), oxygen, sulphur, selenium, tellurium (?), 
fluorine, chlorine, bromine, iodine. 

The sign (?) denotes that these elements are sometimes in- 
cluded in, sometimes excluded from, the class of non-metals. The 
remaining elements are classed as metals. 

(2.) Sources. As a general rule, those elements are found 
native which are unaffected by oxygen and moisture in air at the 
ordinary temperature. Ttms carbon, nitrogen, sulphur, selenium, 
tellurium, the platinum group of metals, and copper, silver, 
mercury, and gold are among these. It is curious that hydrogen 
is not found native to any great extent, for it fulfils these condi- 
tions. There appears no reason why air should not contain small 
traces of hydrogen, unless, indeed, its molecular motion may carry 
it out of the sphere of the earth's attraction.* 

Those compounds of elements with the halogens which are not 
decomposed by water as a rule exist native. Among these are 
chlorides, bromides, and iodides of sodium, and potassium ; of 
silver, lead, and moi>ury. From the abundance of oxygen, and 
the tendency which most elements have to combine with it, the 
oxides and double oxides are the most widely spread compounds : 
for example, the silicates, carbonates, phosphates, nitrates, &c. 
The sulphides rank next in order of distribution ; only those 
stable 111 presence of air and water, however, occur abundantly. 
It is indeed probable that the mass of the earth consists largely 
of sulphides ; for the specific gravity of our globe has b&en found 
by astronomical measurements to be 5^ times that of water, while 
the average specific gravity of the crust cannot well exceed 3. 
It appears not unlikely that the greater density is caused by 
the presence of the denser sulphides in the interior; and the 
prevalence of sulphur in volcanic districts, where the interior of 
the earth is in a state of disturbance, would support this supposi- 
tion. Some few elements occur in combination with arsenic alone, 
or with arsenic and sulphur, 

3. Preparation. It will have been noticed that there are 

* A similar theory would account for the absence of an atmosphere on the 


three general methods of preparing elements from their com- 
pounds. These are 

(ot.) Electrolysis of a liquid compound of the element or 
of a solution of a solid compound in water. It is question- 
able whether solids or gases can be electrolysed ; at all events, the 
constituents cannot be conveniently collected ; hence the limitation 
to the liquid state. It appears probable that no perfectly pure 
compound is capable of conducting electricity ; those at least, such 
as pure water, hydrogen chloride, &c., which can be obtain eel 
nearly pure do not appear to do so. A liquid mixture, however, .is 
almost always an electrolyte, i.p., capable of yielding its elements 
under the influence of a current of electricity. In many cases no 
easily fusible compound of the element required is known, or it 
is difficult of preparation, or it does not conduct; in other discs 
the liberated element acts upon water, forming an oxide and 
liberating hydrogen ; hence the method is somewhat limited. 

(?>.) Heating a compound of the element required. It is 
almost certain that all compounds, it' heated to a sufficiently high 
temperature, would decompose into their elements. But, unless 
one of the elements possesses a much lower boiling-point than the 
others with which it is combined, separation cannot be effected, ns 
a rule, for in most instances recombination occurs on cooling. It, 
is owing to the great difference in volatility of mercury and 
oxygen that the latter can be prepared by heating mercuric oxide ; 
on similar grounds, chlorine can be prepared from gold chloride ; 
or sulphur, by heating platinum sulphide or hydrogen sulphide. 
In many cases only a portion of one of the combined elements is 
evolved as gas, as, for instance, oxygen from manganese, barium, or 
lead dioxides, or from chromium trioxide. 

(c.) By displacing one element from a compound by 
the action of another. This method is very largely used. 
Tte agents of displacement, however, are limited in number. It 
is obviously essential to the success of the process that the element 
used as a displacer shall not combine with the one to be displaced ; 
or, if it do so combine, that it shall be easily expelled from its com- 
bination by heat; or that it shall combine much more, readily 
with one of the elements in the compound acted on than with the 

Thus no metal will displace phosphorus or oxyen from their 
compound, phosphorus pentoxide, because all metals combine with 
phosphorus and oxygen. Again, aluminium may be prepared by 
removing chlorine from its chloride by the action of sodium ; for 
the compound or alloy of aluminium and sodium which is doubt- 


Jess produced is easily decomposed by heat into sodium, which 
volatilises away, and aluminium, which remains non- volatile at 
the temperature employed. And lastly, sulphur may be produced 
by the action of an insufficient quantity of oxygen on it$ com- 
pound with hydrogen ; for hydrogen combines so much more 
readily with oxygen than sulphur does that water is formed, 
little of the sulphur combining with the oxygen; and, as another 
instance, carbon is liberated from its compounds with hydrogen 
when they burn in chlorine gas, because at the temperature of 
reaction the chlorides of carbon are decomposed. 
In practice the following methods are used : 

1. The action of carbon (coal, charcoal) on the oxide of the 
element, or on its compound with oxygen and hydrogen (hydr- 
oxide), at a red heat. The most important elements thus prepared 
are : 

Hydrogen, potassium, rubidium ; zinc, cadmium ; impure 
chromium, iron, manganese, nickel and cobalt (these elements 
combine with a small quantity of the carbon employed in their 
liberation) ; germanium, tin, lead; phosphorus, arsenic, antimony, 
bismuth ; molybdenum, tungsten j copper. In many cases this is 
in reality the action of carbon monoxide on the oxide of the ele- 
ment : the carbon monoxide unites with the oxygen combined with 
the element, forming carbon dioxide, and the element is liberated. 

2. The action of .hydrogen on the oxide of the element 
required at a red heat. Elements which may be thus prepared 
are : Indium, thallium, tin, lead ; nitrogen, arsenic, antimony, 
bismuth, tungsten ; iron, nickel, cobalt, and copper. 

3. The action of hydrogen on the chloride of the element 
at a red heat. Examples : Vanadium, niobium, arsenic, antimony, 
bismuth, and others. 

4. The action of sodium or potassium, or of ^Jnc, on 
the fused Chloride, double chloride, or double fluoride of the 
element required. Examples : Magnesium, boron, aluminium, 
yttrium, carbon, titanium, zirconium, thorium, tantalum, chrom- 
ium, uranium. 

6. The action of another element on the solution of a 
C0mpound of the element required. Examples : Iodine may be 
prepared by the action of chlorine or bromine on iodide of potas- 
sium ; bromine, by the action of chlorine on potassium bromide ; 
sulphur, selenium, or tellurium, by the action of atmospheric oxygen 
on a solution of their compounds with hydrogen ; copper, by the 
action of iron on a solution of copper chloride or sulphate ; mer- 
cury or silver, by the action of copper on a solution of mercuric or 


silver nitrates ; gallium, by the action of zinc on a solution of 
gallium chloride, and many others. 

properties. -The elements, like other forms of matter, exist in 
the three states of gas, liquid, and solid. Those gaseous at the 
ordinary temperature are hydrogen, nitrogen, oxygen, fluorine, 
and chlorine. Two are liquid, viz., bromine and mercury; the 
remainder are solid. 

. The mass of one cubic centimetre varies from 0*0000896 gram 
in the case of hydrogen gas to 22*48 grams in the case of osmium. 
The variation of this constant with atomic weight will be considered 
in Chapter XXXVI. 

The atomic weights of the elements vary from 1 (hydrogen) to 
240, (uranium) ; and their specific heats from 5'4 (hydrogen alloyed 
with palladium) to 0'0277 (uranium). It will subsequently be 
shown that the product of the two is usually a constant number. 

It cannot be doubted that many elements remain to be dis- 
covered. On referring to the periodic table on p. 23, it will be 
seen that many atomic weights are accompanied by queries (?). 
Within the last few years several such gaps have been filled ; 
notably thallium (Crookes), gallium (Lecoq de Boisbaudran), 
scandium (Oleve), and germanium (Winckler). But this subject 
will be fully considered in a later chapter. 

Equations expressing the preparation of elements of Groups XIII and XVI. 

Fluorine. 2HF = H 2 + F 2 . 

Chlorine MnO 2 + 4HC1 = CI 2 4- MnCl 2 + 2II 2 O. 

Bromine. 2KBr + C1 2 = Br 2 + 2KC1. 

Iodine. 2KL + Br 2 = I 2 + 2KBr. 

Copper. CuO + C = Cu + CO. 

f CuS + 2O 2 = CuSO 4 . 

\ CuS0 4 + Ife = Cu + FeS0 4 . 
CuS + 2CuO = 3Cu + SO 2 . 
Mercury.HgS -f O 2 Hg + S0 2 . 
o/ J A g* s + 2NaCl =. 2AgCl 

Silver. 4- Fe = 2Ag - 





Compounds and Mixtures. 

Elements are said to winltinr when on bringing them together a 
new substance is produced, differing from its constituents nnd pos- 
sessing properties which, as a rule, are not the mean of their pro- 
perties. Such combination is always attended with a rise or fall of 
temperature, or " heat change ; " and, as heat is a form of energy, 
or power of doing work, elements either gain or lose energy by 
combination with each other. It appears that direct combination 
is always attended with loss of energy, heat being evolved. This 
is illustrated by the combustion of oarbon in oxygen, of antimony 
in chlorine, and by many other instances ; and the evolution of 
heat in many such cases is so great as to raise the substance to the 
temperature of incandescence, so that it emits light. 

Two or more elements may, however, be mingled without 
sensible evolution of heat. They are then said to constitute a 
mixture. Atmospheric air is an instance in point. On mixing its 
constituent gases, oxygen and nitrogen, no heat change takes place. 
But if electric sparks be passed through the air, its constituents 
are raised to a high temperature and combine ; the product is an 
oxide of nitrogen, possessing a brown-red colour and a strong smell. 
Certain mixtures are thus definitely distinguished from compounds. 
But in many cases it is difficult to affirm positively that an element 
is or is not combined. Some metals mix freely with others, as, for 
example, tin and lead ; but there is no way of absolutely testing 
whether or not they are combined. Another instance is that of a 
solution of c^ ln rint3 in bromine. In such cases, however, the pro- 


perties of tlie mixture are apparently the mean of those of its 

The best criterion of a compound is its definite composition. 
this are associated definite physical properties, such as con- 
stancy of melting point, of boiling point, and of crystalline form. 
An amorphous condition, i.e., lack of crystalline form, almost 
always accompanies indefinite composition ; but, on the other Land, 
a substance may possess a definite crystalline form (as, for example, 
many silicates), and yet have an indefinite composition. Such 
bodies are, however, usually mixtures of compounds with each 

Nomenclature. Chemical nomenclature in its present form 
was mainly devised by Lavoisier, and, although extended, its prin- 
ciple has not been materially modified since his time. Hut even lie 
was constrained to adopt certain expressions which had been in use 
from a very early date, such as " base;" " acid," and " salt." These 
terms are incapable of accurate definition, and must therefore be 
used loosely. Et may be said generally that the word bane is applied 
to the oxides of certain elements, either alone or in combination with 
hydrogen oxide (water) ; the word acid, the oxide of hydrogen and 
certain other elements, not usually those of which tho oxides are 
called bases ; and the word salt, a body produced by the interaction 
of a basic oxide with an acid oxide. The words xa// and acid, however, 
are frequently applied to substances containing no oxygen, such UH 
sodium chloride, or hydrochloric acid. Tn fact, no rule can be 
given, and the words must be employed in a vague sense, custom 
alone determining their use. 

A compound formed by the union of two elements retains the 
names of both, one of them, however, acquiring the termination 
" ide." It is a matter of indifference which receives that ending; but, 
as most compounds which have been investigated contain one of 
ten or twelve elements, the names of these are commonly modified. 
Thus we speak of oxides, sulphides, selenides, tellurides, fluorides, 
chlorides, bromides, iodides, nitrides, phosphides, arsenides, borid^s, 
carbides, and silicides ; also of hydrides. The Greek numeral pre- 
fixes mono-, di-, iri-, tetra-, penta-, and the Latin one sesqni.-, signi- 
fying respectively, one, two, three, four, five, and one-and-a-half, 
are employed, when required, to denote the relative numbers of 
atoms in the compound. 

Many compounds of fluorides, chlorides, bromides, and iodides 
with each other, and of oxides and sulphides with chlorides, &c., 
are known. These have generally been named double chlorides, 
oxy chlorides, or basic chlorides, sulphochloridos, &$. Another 


nomenclature is sometimes used. It is as follows ; an example 
will render it plain. Platinum forms two compounds with chlorine, 
one containing twice as much chlorine as the other, proportionately 
to the metal. The one containing least chlorine is named platinows 
chloride; that containing most, platim'c chloride. Each of these 
forms a compound with potassium chloride ; the first is named 
potassium platinoclilori.de, platino- being contracted from platinums ; 
the second potassium platirn'ohloride, the word platim- standing for 
platin/r. So with ferrous and ferric, phosphorous and phos- 
phor/V, <fec. 

The double oxides have names which do not show that they 
contain oxygen. Thus compounds of oxides of chlorine and of a 
metal are named hypucliloritcs (hypo = below), chlof tfes, chlorates 
or perclilovates (per = over, from hyper), according to the amount 
of oxygen in combination with the chlorine ; so also with com- 
pounds of nitrogen, phosphorus, sulphur, gold, &c., &c. 

In the case of a few common substances, such as water (hydrogen 
monoxide), ammonia (hydrogen nitride), vitriol (sulphuric acid or 
hydrogen sulphate), old and familiar names have been retained. 
These are fortunately in many cases becoming obsolete. 

The Elements 

Will be considered in the following order: 

1. Compounds of the halogens fluorine, chlorine, brom- 

ine, and iodine with the elements, arranged in 
groups according to the periodic table, including 
double compounds. 

2. Compounds of oxygen, sulphur, selenium, and tellur- 

ium with the elements; including oxychlorides, 
sulphochlorides, &c., and double oxides and sul- 
phides, usually called hydroxides, hydrosulphides, 
acids, and salts. 

3. Borides, carbides, silicides. 

4. Nitrides, phosphides, arsenides, and antimonides. 

Double compounds. 

5. Alloys and amalgams. 

Before proceeding with the consideration of the halogen com- 
pounds it is necessary, in order to understand the relations between 
these substances, to study tho methods of expressing chemical chancre, 
and some of the reasons for assigning definite atomic weights to 
the elements. This involves a knowledge of the nature of gases, 
and their bqjiaviour as regards temperature and pressure. 


The States of Matter. 

Matter is known in three states : the solid, the liquid, and the 

Solids. Solids are peculiar in possessing form ; they have 
rigidity, enabling them to keep their shape. It is believed that 
minute particles of which all matter consists, which aie named 
molecules, are so closely packed together in solids as to attract 
each other powerfully, and to possess very little freedom of motion. 
Such particles possess symmetrical arrangement in crystals ; but 
are heaped together at random in amorphous solids. Solids 
generally expand when their temperature is raised, but only to a 
small degree. At a sufficiently high temporature, they either 
melt, volatilise without melting, or decompose. They are very 
slightly compressible. 

Liquids. Liquids differ from solids in not possessing form, 
and from gases by possessing a surface. The condition of the 
liquid matter at the surface differs from that in the interior, and 
the surface is under a lateral strain, named surface,- tens ion. A 
drop, for example, behaves as if it wore covered with a stretched 
skin or film. The molecules of which liquids consist possess 
greater freedom of motion than do those of solids; so that they 
move about, continually gliding past each other, and hence a liquid 
has no fixity of form. On raising the temperature of a liquid, this 
motion increases. The motion of the molecules of a liquid is termed 
diffusion or osmosis. When liquids are cooled they generally con- 
tract, and at a sufficiently low temperature they freeze, or turn to 
8hds; on raising their temperature they expand, and at a suffi- 
ciently high temperature the) r volatilise, changing into gas. 
Vapour is continually being evolved from the surface of a liquid, 
and if. the liquid be in a closed vessel the pressure which its 
vapour exerts can be measured. This pressure is termed its 
vapour-pressure. The vapour- pressure increases with rise of tem- 
perature ; and when it exceeds the pressure of the atmosphere the 
liquid boils and changes wholly into gas, if heat be supplied in 
sufficient amount. 

Gases. Gases or vapours have neither form nor surface. A 
solid or a liquid in changing into vapour acquires a greatly increased 
volume ; thus the gas of water occupies about 1700 times the space 
occ.upied by its own weight of liquid water at the same tempera- 
ture and at the same pressure, viz., 100 and 7b'0 mm. pressure. 
While solids and liquids are but slightly altered in volume by 
alteration of pressure and temperature, the volumes^of gases are 


greatly changed. The molecules of gases are evidently much 
more distant from one another than those of solids or liquids, and 
therefore possess much greater freedom for motion, or free ptith. 
They occupy but a small portion of the space which they inhabit. 
And while the molecules of solids and of liquids are so near each 
other as to exercise great attraction on one another, those of gases 
are so far apart that the attraction is barely sensible. Hence 
gases exhibit simple relations to temperature and pressure. 

Relation between the volume of a gas and the pressure 
to which it is exposed. Boyle's law. The temperature of 
a gas being kept constant, its volume varies inversely as 
the pressure to which it is exposed. This law was discovered 
by Robert Boyle in 1660. 

The barometer. It has been remarked in Chapter II that gases 
have weight; the weight of a given quantity of matter is not 
changed by change of state : thus a pound of water weighs a 
pound, whether it bo ice, water, or steam. Air, which is a mixture 
of nitrogen and oxygon gases, therefore possesses weight ; and, the 
longer or higher a column of air, the greater its weight. A column 
of air reaching to the upper confines of the atmosphere and rest- 
ing on the earth at the level of the sea, of 1 square centimetre 
in section, weighs on the average 1033 grams ; or, if 1 square 
inch in section, about 1C Ibs. ; but 1033 grams is the weight of 
a column of mercury at of 1 square centimetre in sectional 
area and 760 millimetres in length; and 1C Ibs. is the approxi- 
mate weight of a mercury column 1 square inch in sectional area 
and approximately 30 inches long; or of a column of water about 
33 feet in length, also 1 square inch in sectional area. Hence, if it 
were possible to support the end of such a column of air on one 
pan of a balance, and to place on the other pan a column of 
mercury 760 millimetres in length, removing the pressure of the 
air from its upper surface (else the weight of both air and mer- 
cury would press on the other pan), the two columns would 
balance. Such an operation is actually performed in construct- 
ing a barometer. The air is removed from the upper portion 
of a glass tube, the lower end of which is open and dips in mer- 
cury ; and the mercury rises in the tube until it balances a column 
of air of equal sectional area to the tube, rising, in order that it 
may do BO, to a height of 700 millimetres. If the weight of the 
atmosphere increases, owing to its cooling, or to its compression, 
the column of mercury rises proportionately, so as to balance it ; 
and, conversely, when the weight of the atmosphere decreases, the 



balancing column is shorter. The pressure of the atmosphere 
might be expressed in units of weight for a given sectional area, 
say,, 1 square centimetre; it might be, and indeed sometimes is, 
measured in fractions or multiples of 1033 grams, jnst as it is 
the custom for engineers to express the steam pressure in a boiler, 
which is closely analogous, in pounds 011 the square inch of 
boiler surface; but it is commonly expressed as equal to the pres- 
sure of 760 millimetres of mercury, or of a column of greater or 
less length, according as the weight of the atmosphere vaiies. 

All gases contained in vessels communicating with the atmo- 
sphere are therefore under this pressure; hence it must be allowed 
for in ascertaining the relation between the volume of a gas and 
the pressure. That Boyle's law is approximately true can be 
proved by the following experiment: A U" tuOc ' as shown in 
fig. 16, about 50 centimetres in length, contains air in its closed 


limb. The amount of air is adjusted so that when the mercury 
is level in both limbs it occupies a volume represented by 
273 millimetres -f a number of millimetres equal to the tem- 
perature of the day. Thus, for example, if the temperature of 
the surrounding atmosphere is 15 C., the length of the column 
of air enclosed should be 288 millimetres. The reason of this 
adjustment will appear later. Now mercury is poured into the 


open limb of the JJ-tube so as nearly to fill it; the difference in 
level of the mercury in the open and in the closed limb is read 
off. The air in the closed limb will be compressed by the weight 
of the mercury in the open limb, the column being equal in length 
to the difference in level of the mercury in the open and in the 
closed limb. For example : 

Distance from top of tube to surface of mercury, 

that in both limbs being at the same level. . 288 mm. 

Level of mercury in open limb, after filling it. . ,, 

Level of mercury in closed limb, after filling 

open limb 223 

Difference in level between mercury in closed , 

and open limbs 223 

The initial volume of the gas was 288 X x cubic centimetres. 

After compression the volume decreased to 223 X x cubic 

The initial pressure was that of the atmosphere, say, 760 milli- 

The final pressure on the gas was that of the atmosphere, 
760 millimetres + 223 millimetres = 983 millimetres of mercury. 

But 1)83 : 760 :: 288a> : 223*, nenrly. 

Hence tho volume, of fye gas decreases proportionately to tlip in- 
crease of pressure at a temperature of 15. 

If pi, ^) 2 , V}j and ?'> represent the pressures and volumes respec- 
tively before and nf ter alteration, then 

P^VI = p-ii\, provided temperature be kept constant. 

Similar expoiiments may be performed, decreasing the amount 
of mercury in the JJ-tube, ^7 running out mercury through the 
stopcock, nnd so reducing the pressure on the gas; and ]V>vle's 
law may thus be proved true for such small alterations of pres- 
sure. When the pressure is very great, it ceases to hold : gases 
become more compressible up to a certain point, and then less 
compressible with greater rise of pressure. 

. Gay-Lussac's law. The volume of a gas increases one 
two hundred and seventy-third of its volume at (0-00367) 
for each rise of 1 C., provided pressure remain constant. 
Thus 1 cubic centimetre of air or other gas measured at becomes, 
when heated to l,l f <f or 1-00367 c.c.; at 2, l f y ,or 1 + (0'00367 

* The barometer should be read at the time, and its height substituted here. 
In the above instance it is supposed to be at its normal height. 


X 2) ; at 100, lijg, or 1 + (0-00367 x 100). This can lie illus- 
trated by means of the apparatus used for demonstrating Boyle's 
la w^ with a slight addition to allow of an alteration in the tem- 
pera^ure of the gas. The closed limb of the JJ-tubo is sur- 
rounded by a jacket or mantle of glass, the lower part of which 
is closed by an indiarubber cork, perforated to allow the limb 
of the (J-tube to pass through. The liquid in the bulb is 
water. It is boiled by a flame, and the steam jackets the 


(J-tube, raising its temperature to 100, provided the atmo- 
spheric pressure is 76*0 millimetres. If the pressure does not 
differ much from the normal one, the difference in temperature* 
may be neglected. At 15, supposed to be the atmospheric tem- 
perature of the day, the air in the closed limb of the JJ-tubo 
is adjusted so as to occupy 288 millimetres of the tube's 
length, measured from the top downwards, the mercury in both 
limbs being level. On boiling the water so as to raise the tem- 
perature of the gas to 100, the gas will expand, pushing down 1 


the mercury in its own limb, and raising it in the other. When 
the level is stationary, mercury is run out of the JJ-tube, 8O as 
to restore equal level in both limbs. The gas will then occupy 
373 millimetres of the length of the \J-tube. Thus : 6 

Initial volume of gas at 15, at atmospheric pressure, 288.2 c.c. 
Final volume of gas at 100, and at 373a? 

Expansion, 373^ 288^ = S^x c c. 
Rise of temperature, 100 15 = 85. 

The expansion is thus seen to be proportional to the rise of 

It is obvious that by cooling the gas to 0, by surrounding the 
tube with melting ice, the volume would contract from 2S8.r c c. to 
273, c.c. This may also be proved experimentally. It may also 
be .shown that Boyle's law holds equally well at the temperature 
100 as at 15 by means of this apparatus. 

Such an instrument might be, and indeed with altered con- 
struction is, used as a thermometer. It would be convenient to 
place the number 273 at the level of the mercury when ice sur- 
rounds the tube ; then the expansion of the gas and the tempera- 
ture will march pari pas^u. The zero of such a scale will mani- 
festly bo at the top of the tube ; the degrees are ordinary 
Centigrade degrees, the interval of temperature between the 
melting-point of ice ard- the boiling-point of water under normal 
pressure being 100; but on this scale the former is marked 273 
and the latter 373. Such a scale is termed the absolute scale] 
and the temperature 273 C. is equal to absolute. 

As this behaviour with respect to pressure arid temperature is 
common, speaking approximately, to all gases, it may be con- 
jectured that they possess some property in common, as the cause 
of their similar changes. This property was discovered yi 1811 
by Avogadro, and is known as 

Avogadro's law. Equal volumes of gases, under the 
same pressure, and at the same temperature, contain equal 
numbers of molecules. It must be noted that this state- 
ment postulates nothing as regards the actual si'/e of the gaseous 
molecules; it merely asserts that, temperature and pressure being 
constant, a detmite number of molecules of one gas, say, hydrogen, 
inhabit the same space as the same number of molecules of any 
other gas, say, oxygen or chlorine. The actual number of mole- 
cules is, of course, unknown ; and, although attempts to estimate it 
have been made, they do not concern us here. JYoin the known 


laws of expansion of gases, and their relation towards pressure, it 
is possible to compare the weights of equal volumes of different 
gases* and so to compare the relative weights of the molecules of 
whicH they are composed ; for it is obvious that if the weight of 
n molecules of, say, oxygen is 16 times that of n molecules of 
hydrogen at some temperature and pressure the same for both, the 
weight of 1 molecule of oxygen is 16 times that of 1 molecule of 

Ifc is, therefore, exceedingly import-ant to be able to compare 
the relative weights of gases, inasmuch as it affords a simple 
means of comparing the relative weights of their molecules. 
The term drnsity is applied to the weight of a gas relative to 
hydrogen, the density of which is arbitrarily placed = 1. Some- 
times air is chosen as the unit of comparison The absurdity of 
this is evident; for it has been repeatedly shown that the composi- 
tion, and hence the density, of air, which is a chance mixture of tho 
gases oxygen and nitrogen, is not uniform, but varies within small 
limits. The variation, however, is so smalt as to be within the 
usual errors of experiment in determining the density of gasos ; 
hence, for practical purposes, as air is about 14*47 times as heavy as 
hydrogen, densities compared with air may be converted to the 
hydrogen standard by multiplying the number expressing thorn by 
14'47. The density of a gas which exists as a liquid at ordinary 
atmospheric temperature is termed a vapour-density ; there is no 
real distinction between the words gas and vapour. 

Methods of determining the Density of Gases. 

1. When the substance is a gas at the temperature of the 
atmosphere. Two globes of nearly equal capacity (half a litre to 
five, litres, and which should have as nearly as possible the same 
weight), provided with tight-fitting stop-cocks, are pumped empty, 
first by means of a water-pump, and finally with a Sprengel's or 
other mercury-pump ; the stop-cocks are then closed, and they 
are suspended one from each arm of a balance, as shown in fig. 18, 
and if not quite equal in weight, counterpoised by addition of 
weights to one or other pan. The gas to be weighed, is thru 
admitted from a gas-holder into one of the globes, care being taken 
to dry it, by passing it slowly through JJ*^ u ^ e8 filled with strong 
sulphuric acid or phosphorus pentoxide, which has a great ten- 
dency to combine with water, and so removes it from the gas. If 
the gas be soluble in water, it may be passed straight from the 

' H 



generating flask through the drying tubes 'into the empty globe, 
the stop-cock of the latter being opened slowly so as to ensure the 
gas being thoroughly dried. It is again suspended from the hook of 
the balance pan, and after some hours, the amount of gas whi 3h has 
entered is weighed. The volume of the globe is then ascertained 

FIG. 18. 

by filling it completely with water and weighing it. The difference 
between the weight of the vacuous globe and the globe full of 
water gives the weight of water tilling the globe. It is sufficient 
for the present purpose to consider that 1 gram of water occupies 
1 cubic centimetre, though for accurate determinations the true 
volume of the water at 4 must be calculated. Here, also, the 
expansion of the globe between and atmospheric temperature, 
and also its diminution of volume when empty of air, c'ue to the 
presence of the atmosphere, have been neglected. 
We have accordingly the data. 

Weight of globe full of gas at T temp., and 

P mm. pressure W 2 grams. 

Weight of empty globe W! 

Weight of V cub. centimetres of gas W grams. 

From this the volume of 1 litre of the gas at and 760 milli- 
metres pressure can be calculated thus : 

(a.) To ascertain the volume of the gas at 760 millimetres 


Law. The volume is iftiiersely as the pressure. Hence, 
V 760 = V P x P/7GO. 

To ascertain the volume, corrected for pressure, at 0C. 
Law. The volume of a gas increases by 0036*7 of its volume at 
for each rise of 1. Hence, 

v _ V P x P/760 

V and 750- 

(c.) This volume of gas weighed W grams. To find the weight 
of 1 litre : W 1000 co = 1000 W/V. aud 7fio mm 

(^.) From Regnault's very accurate experiments wo learn that 
1000 cubic centimetres of hydrogen weigh 0'089G gram. Hence, 
the density of the gas = W l()00( .e /0'0896. 

The relative weight of a molecule of hydrogen is taken aw 2, 
for reasons which will afterwards ho considered (p. 109), Hence, 
the relative weight of a molecule, or the molecular wc-ight of the 
gas = 2W,ooo cr /0089f>, or is equal to twice its density. 

2. When the substance becomes gaseous at a tempera- 
ture higher than that of the atmosphere. Ono oi the follow- 
ing methods may be employed. 

(a.) Dumas' Method. This method differs from the method 
already described only in one particular, viz., in the manner of 
filling the globe. The globe usually has a capacity of 250 to 
500 cubic centimetres. About 10 cubic centimetres of the liquid 
or solid, of which the density in the gaseous state is required, is 
introduced into the globe by warming it gently BO as to expel air, 
and dipping the thin neck of the globe into the liquid ; or by 
introducing the solid into the globe before its neck is drawn out. 
It is then placed in a bath of some liquid or vapour, depending on 
the temperature required. If the boiling-point is below 100, 
water may be used ; if between 100 and 250, olive or castor-oil ; 
and vapour-baths, such as that of boiling mercury (358), or 
sulphur (444), or phosphorus pentasulphide, or stannous chloride, 
or even the vapours of boiling cadmium or zinc, may be used for 
.higher temperatures, but with the last two the globe must be a 
porcelain one, for glass softens at about 700. The liquid or solid 
begins to evaporate, and its vapour displaces the air from the 
globe. As soon as vapour ceases to escape, the drawn-out end of 
the neck of the globe is sealed by means of a hand- blowpipe, or of 
an oxyhydrogen blowpipe, if a porcelain globe is wed (see fig. 19). 
The globe is then removed, allowed to cool, cleaned, and weighed, 
balancing it, as before, by a similar globe hung from the pther pan 

H 2 



of the balance. The calculations are performed exactly as before, 
but the expansion of the globe must here be allowed for ; if of 
glass, it may be calculated as V + (0*0000250 ; it is assumed that 
the gas will remain a gas when cooled to 0. It would b more 

FIG. 19. 

rational to compare the weight of the gas with that of a,n equal 
volume of hydrogen at the same temperature and pressure as 
those of the vapour at the time of sealing the globe, but the end 
result is the same whichever method of calculation be used. 

(b.) Hofmaim's Method, modified. The principle of this 
method is to ascertain the volume of a known weight of the gas. The 
apparatus consists of a graduated tube of the form shown in fig. 20. 
The tube is filled with mercury and inverted into a glass basin 
containing mercury, and after the jacketing tube has been put on, 
the apparatus is clamped in a vertical position. The graduated 
tube passes through a wide hole in an indiarubber cork fitting the 
jacket ; but as this cork is apt to be attacked by the boiling 
liquid, a little mercury is poured in, so as to cover and protect 
it. The substance is weighed out in a small bulb, and pushed 
under the open end of the tube, so that it floats up to the surface of 
the mercury in the closed end. The temperature is then raised by 
boiling tLe liquid, which must be pure, in the bulb of the jacket. 



FIG. 20. 

The following is a list of convenient substances, with their vcspe 
tive boiling points under a pressure of 760 millimetres : 

T. A. 

Carbon disulphkle. 40' 2 25 

Alcohol 78 3 M 

Chlorobenzene. ... 132*1 25 

Bromobenzene. . . . 156' 1 20 


Aniline ]84'5 

Ciiinoline 237'5 J 

Bromonapkthalene 280'4 


The column, A, represents the average difference in pressure in 
millimetres per degree at about the pressure 70*0 millimetres. 
Thus, if the height of the barometer is 740 millimetres, t,e , 
20 millimetres less than 760, the temperature of the carbon di- 
sulphide vapour will be not 46*2, but 40*2 * Jths of 1 = 45'4. 
The mercury in the tube will be displaced by the vapour, and 
will enter the glass basin in which the tube stands. 'J'he volume 


of the vapour is then read off, if the tube is a graduated one ; if 
not, the level of the mercury in the tube is read on the graduated 
scale, and also the level of the top of the tube. The volume may 
be afterwards determined, by inverting the tube, and filling" it to 
the required height with water from a burette. The pressure is 
that of the atmosphere, diminished by the length of the column of 
mercury in the tube. But mercury itself, when heated, expands, 
arid a correction must be introduced, because at the length of 
the mercury column would be less. Again, the gas in the tube 
consists partly of mercury vapour ; its pressure must be calculated 
and subtracted.* But neglecting these corrections, the plan of cal- 
culation is as follows : 

A certain volume of gas, V, has been produced from a known 
weight of liquid or solid, W. This gas is at the temperature of 
the jacketing vapour, and under atmospheric pressure diminished 
by the length of the column of mercury, equal to the distance 
between the level of mercury in the glass basin and that in the 
tube. The weight of an equal volume of hydrogen at the same 
temperature and pressure is calculated, and the weight of the 
vapour is divided by the weight of the hydrogen. The quotient is 
the density. 

(c.) Victor Meyer's Method. In this case not mercury but 
air (or some other gas) is displaced ; and the volume of a known 
weight of the vapour is deduced from that of the displaced gas, or 
air. A cylindrical bulb, c (fig. 21), with a long stem, fr, closed by 
a cork at its upper extremity, as shown in the figure, is heated to 
some constant temperature by an oil- or vapour-bath, as already 
described. The air expands while the temperature is rising, and 
issues through the side tube, d, escaping in bubbles through the 
water in the trough. When bubbles cease to rise the temperature 
is assumed to be constant. The tube is quickly uncorked, a small 
tube, full of the liquid or solid whose vapour-density is sought, is 
dropped in, falling on sand, placed at the bottom of the cylinder, 
so as to avoid breaking it. The cork is then rapidly replaced. 
The substance turns to gas, and expels air from the cylindrical 
bulb. This air is cooled in passing up the stem and through 

* The following data are available for this calculation : 

Temperature 46 78 132 156 184 237 280 

Expansion of 1 c.c. of 

mercury between 

and f 1 '0083 1 -0141 1 '0240 1 '0285 1 '0338 1 -0438 1 '0521 

Vapour pressure of 

mercury, in mm. . . '1 1 '0 3 '0 10 '0 52 -5 157'0 



the water; it is collected in a graduated tube. Its volume ia 
equal to that of the vapour, supposing the latter to have been 
cooled to the atmospheric temperature, and to have withstood the 
process without condensing. We have then a given volume of 
air at atmospheric temperature and pressure corresponding to that 
of the vapour ; and also the weight of substance which has pro- 
duced the vapour by which the air has been expelled. From these 
data it will be seen the density of the vapour may be calculated. 

Such are the available means of ascertaining the weights of 
one litre of various gases and their densities. The processes have 
been described in some detail, because such determinations have 
the utmost chemical importance. The deductions to be drawn 
irom them will appear in the next chapter. 

. 21. 




Hydrogen Fluoride, Chloride, Bromide, and 
Iodide, * 

Only one compound of each of these elements with hydrogen 
is known. 

Sources. Hydrogen chloride is present in the atmosphere in 
the neighbourhood of volcanoes ; it has been doubtless formed by 
the action of steam on certain chlorides, easily decomposed by 
water into oxide of the element and hydrogen chloride. The 
others do not exist free. 

Preparation. 1. By direct union, (a.) Hydrogen Fluo- 
ride. During the preparation of fluorine by Moissan, by the 
electrolysis of hydrogf^ potassium fluoride, hydiogen was liberated 
from the negative, and fluorine from the positive pole (see p. 73). 
When a bubble of hydrogen escaped round the bend of the 
JJ-tube, and mixed with the fluorine, an explosion was heard, 
showing that these two elements unite at the ordinary tempera- 
ture, and in the dark. 

(6.) Hydrogen Chloride. Equal volumes of hydrogen and 
chlorine gas unite directly on exposure to violet light, or on 
application of heat. This may be shown as follows : 

A tube of the form shown in fig. 22 is employed. The stop- 
cock in the middle divides it equally into two halves. The stop- 
cock in the middle being shut, one side is filled with dry chlorine 
by downward displacement, a capillary tube serving to conduct 
the chlorine gas to the lower closed end, as shown in the figure. 
The stop-cock is then closed. The other half of the tube is then 
filled with dry hydrogen by upward displacement, for hydrogen is 
lighter, though chlorine is heavier, than air. The tube is then 
placed in a dark place, for example, a close fitting drawer, for 
some hours, the stop-cock in the middle being opened. The two 
gases will mix, but will not combine. It is then placed for an 
instant in direct sunlight, or, if that is not available, illumined by 



burning a piece of magnesium ribbon witbin a few inches of it. A 
flash ,vill be seen inside the tube, showing that combination has 

taken place, and the green colour of the chlorine will disappear. 
It is safer, however, to expose the tube for some hours to diffuse 
daylight. One end of the tube is now dipped in mercury, and the 
lower stop-cock is opened. The mercury does not enter the tube, 
showing that the hydrogen chloride retains the same volume as its 
constituents; it does not act on mercury. The stop-cock is again 
cldsed, and the lower end of the tube is now dipped in water, and 
the stop-cock again opened. The water rushes in, and completely 
fills the tube, provided both compartments were exactly equal, and 
that all air was displaced on filling it with chlorine and hydrogen. 
Chlorine is sparingly soluble in water, hydrogen nearly insoluble. 
Hence a gas has been produced by the combination of equal 
volumes of hydrogen and chlorine, which occupies the same 
volume as its two constituents, but which differs from them in 

A jet of hydrogen gas may be burned in a jar of chlorine. The 
hydrogen is lit, and, while burning in the air, a jar of chlorine is 
brought under it, and raised so that the jet dips into fche chlorine. 



The hydrogen continues to burn, but with a greenish-white flame. 
Fumes are produced. 

(c.) Hydrogen Bromide. Hydrogen and bromine do jaot 
combine so readily as hydrogen and fluorine or as hydrogen and 
chlorine. Their direct combination may be shown as follows : A 
bulb tube is connected with an apparatus for generating hydrogen. 
A few cubic centimetres of bromine are placed in the bulb ; the 
hydrogen passes over the bromine, and carries some with it as gas. 

FIG. 23. 

The hydrogen is lit, and burns, combining partly with the oxygen 
of the air, partly with the bromine. The hydrogen bromide 
formed unites with the water-vapour forming a white cloud of 
small liquid particles. It is owing to the formation of a similar 
compound with water that fumes are produced when hydrogen 
burns in chlorine. 

A practical plan of preparing hydrogen bromide is to pass the 
mixture of hydrogen and bromine, prepared as described, through 
a glass tube containing a spiral coil of platinum wire, heated to 
redness by an electric current. The uncombined bromine is 
absorbed by passing the resulting gas through a tube filled with 
powdered antimony. 

(d.) Hydrogen and Iodine may be made to combine directly by 
heating them together in a sealed tube to 440 for many days. Com- 
plete combination does not take place, however long the mixture 
is heated, and about one quarter of the hydrogen and one quarter 
of the iodine remain uncombined. 

2. By the Action of the Halogen on most Compounds of 
Hydrogen. t Instances. (a.) On water. A solution of chlorine 


gas in water exposed to sunlight yields oxygen and hydrogen 
chloride ; if chlorine and water-gas be led through a red-hot tube, 
some of the water-gas reacts with the chlorine, yielding hydrogen 
chloride and oxygen. (6.) On hydrogen sulphide, dissolved in 
water. The products are sulphur and the hydrogen compound of 
the halogen. This is a convenient method of preparing hydrogen 
iodide. Sulphuietted hydrogen gas (see p. 196) is passed through 
water in which iodine is suspended. The liquid becomes milky, 
owing to separation of sulphur, and the colour gradually dis- 
appears, owing to the union of the iodine with the hydrogen of 
the hydrogen sulphide. When the reaction is over, the sulphur is 
separated by t filtration, and the liquid distilled. It is, however, 
impossible to separate hydrogen iodide from its solution in water 
by distillation. The aqueous solution is termed hydriodic acid, 
(c.) Chlorine, bromine, and iodine act on ammonia, yielding nitro- 
gen and the compound of the halogen with hydrogen. Nitrogen 
combines with the halogen, if the latter is in excess, yielding very 
explosive bodies (see p. 158). (cZ.) Chlorine and bromine act on 
hydrocarbons (carbides of hydrogen) giving compounds of carbon 
with both chlorine (or bromine) and hydrogen, and the haloid 
acid. Generally it may be stated that almost all compounds of 
hydrogen are decomposed by the halogens, yielding a haloid com- 
pound of the element, and h}drogen chloride, bromide, or iodide. 

3. By the Action of Water, or of Double Oxides of 
Hydrogen and some other Element on Compounds of the 
Halogens, a. Action of Water. The halogen compounds of 
boron, silicon, titanium, phosphorus, sulphur, selenium, and 
tellurium, are at once decomposed by cold water. Hence the 
halogen added to water in which one of these elements is sus- 
pended, combines with part of the hydrogen of the water, the 
remaining hydrogen and oxygen combining with the element (see 
these haloid compounds, p. 188). Instances: (a.) This is a 
practical method of preparing hydrogen bromide. The bromine is 
added very gradually to phosphorus, lying in water in a retort. 
Phosphorus bromide is produced, and decomposed by the water, 
forming phosphorous and phosphoric acids, and hydrogen bromide. 
After all the phosphorus has disappeared, the liquid is distilled. 
The solution of hydrogen bromide in water is named hydrobromic 

There is little doubt that all soluble chlorides, bromides, and 
iodides are decomposed by excess of water, forming the hydroxide 
of the metal and hydrogen chloride, bromide, or iodide. But in 
most cases there is no available method of separating the hydr- 


oxide from the hydrogen halide, for, on evaporation, the reverse 
reaction takes place, and water alone escapes. Yet, at a high 
temperature, magnesium chloride and some other chlorides react 
with water-gas, giving an oxy-chloride and hydrogen chloride. 4 

(6.) This is a recently patented method of manufacturing 
hydrogen chloride, and promises to be successful. Steam is led over 
magnesium chloride, heated in tubes ; hydrogen chloride is evolved, 
and a compound of magnesium oxide and chloride remains.* 

b. Action of Hydroxides. The hydroxides which react in 
this manner are termed acids. Generally stated, the hydrogen 
halides can be prepared by the action of any hydroxide which does 
not react with them. Phosphoric, sulphuric, and selenic acids are 

c. This is the common method of preparing hydrogen 
fluoride. The fluoride generally employed is calcium fluoride, or 
fluor-spar, which occurs native ; it is treated with sulphuric acid 
in leaden vessels, and the gas evolved is condensed in a worm of 
lead and stored in leaden or gutta-percha bottles. It acts on 
silica, which is a large constituent of glass and porcelain ; hence 
the use of lead, which is but slightly attacked. On a small scale, 
platinum vessels and potassium fluoride answer better. 

d. This is also the best method of preparing hydrogen 
chloride. On a small scale, about 50 grams of sodium chloride 
(common salt) are placed in a retort, and covered with a mixture 
of equal volumes of sulphuric acid and water. On applying a 
gentle heat the hydrogen chloride comes over in the gaseous state. 
It may bo led into water; the solution is called hydrochloric acid. 

On a large scale, the operation is conducted in circular furnaces 
with a revolving bed. The salt and sulphuric acid are introduced 
from above, and fall on to the middle of a revolving plate of iron 
covered with fire-clay, which forms the bed of the furnace. The 
product, sodium sulphate, or " salt-cake," is raked by mechanical 
means towards the circumference of the plate, and drops through 
traps for the purpose. The hydrogen chloride is led up brick towers 
filled with small lumps of coke, kept moist with water from above. 
The water dissolves the hydrogen chloride, which is sent to market 
in carboys. 

e. As both hydrogen bromide and iodide react with and 
decompose sulphuric acid (see p. Ill), bromine or iodine being 
liberated, phosphoric acid must be used for their preparation. 
The method of operation is similar to that of preparing hydrogen 

t * Soc. Chem. Industry, 1837, 775. 


4. Heating Compounds of the'Hydrogen Halide with the 
Haloid Compounds of other Elements. Such compounds 
always decompose when heated. In practice, this method is 
employed for the preparation of pure hydrogen fluoride. Its 
compound with potassium fluoride, after being dried, is heated to 
redness in a platinum retort, and the hydrogen fluoride which 
distils over is condensed by passing through a platinum tube 
surrounded with a freezing mixture, and collected in a platinum 
bottle. The preparation of pure hydrogen fluoride is exceedingly 
dangerous, owing to its great corrosive action. 

Before considering the properties of these bodies, the nature 
of the changes which have been described, and the method of 
representing these changes, must be discussed. 

Atoms and Molecules. 

It was stated in last chapter that equal volumes of gases contain 
equal numbers of molecules. Now, it has been shown that equal 
volumes of hydrogen and chlorine unite to form hydrogen chloride. 
It might be concluded that such a compound consists of 1 molecule 
of chlorine in union with 1 molecule of hydrogen ; but the follow- 
ing considerations will show that such a supposition is inconsistent 
with Avogadro's law. The actual facts are that 1 0025 gram 
of hydrogen, occupying at standard tempera! uro and pressure 
11*16 litres, combines with 35*46 grams of chlorine, also occupying 
11*16 litres, and that the volume of the product is 11' 16 X 2, or 
22*32 litres. We de not know the actual number of molecules of 
hydrogen, or of chlorine, in 11*16 litres of these gases ; let us cnll 
it n. Then n molecules of hydrogen, on this supposition, unite 
with n molecules of chlorine, and as chemical combination has 
occurred, n molecules of hydrogen chloride are formed. But the 
volume of the hydrogen chloride is 22*32 litres; hence n molecules 
of hydrogen chloride would thus occupy (11*16 X 2) litres, instead 
of 11*16 ; or the requirements of Avogadro's law would not be 
complied with, inasmuch as there would be only half as many 
molecules in a given volume of hydrogen chloride as in the same 
volume of hydrogen or of chlorine. But there is no reason to 
suppose that hydrogen chloride does not fulfil Avogadro's law ; 
its expansion by rise of temperature and behaviour as regards 
pressure are practically the same as those of hydrogen and 
chlorine, hence the conclusion is evidently false. The accepted 
explanation is as follows : 


A molecnle of hydrogen, or a molecule of chlorine, is not a 
simple thing"; it consists of two portions in combination with 
each other; these portions are named atoms. When chlorine^and 
hydrogen combine to form hydrogen chloride, their double atoms 
or molecules split, each atom of hydrogen leaving its neighbour 
atom, and uniting to an atom of chlorine, which has also parted 
with its neighbour atom. The original arrangement may be 
represented thus : 

and the final arrangement, thus : 

In 11 '16 litres of hydrogen chloride there is, therefore, the 
same number of molecules as in an equal volume of hydrogen or 
of chlorine; but whereas tho hydrogen chloride molecules contain 
an atom of each element, those of hydrogen contain two atoms of 
hydrogen, and those of chlorine contain two atoms of chlorine. 

Symbols are employed to express such changes. The expression 
of the change is termed an equation ; and the above change is 
written thus : 

JT 2 + Ck = 2ZTC7. 

Where the small numeral follows the letter, it signifies the 
number of atoms in the molecule, as H 2 , Cl z ; where a large 
numeral precedes the formula, it signifies the number of molecules; 
thus, 2HOI. 211 would mean two uncombined atoms of hydrogen ; 
Ua signifies two atoms combined into a molecule. Atoms of 
hydrogen have not been obtained uncombined with each other ; 
atoms of chlorine, however, exist uncombined, or in the free state, 
at a sufficiently high temperature. 

Such an equation expresses the following facts : 

1. That 22*32 litres of hydrogen react with 22'32 litres of 
chlorine, producing 44'64 litres of hydrogen chloride ; and 

2. That 2-005 grams of hydrogen react with 70'92 grams of 
chlorine, forming 72*925 grams of hydrogen chloride. 

It is obvious that 22*32 litres of hydrogen chloride weigh 
72'925/2 grams ; and as 22'32 litres of hydrogen weigh 2'005 grams, 
hydrogen chloride is 18'231 times as heavy as hydrogen. This 
has been found to be the case by direct experiment. Hence the 
molecular weight of hydrogen chloride = 36'4625 is twice its 
density compared with hydrogen. 

Such formulae as HCl, If,, Ck, apply only to gases. In this 
book the symbols for gaseous elements and compounds are prin f ed 


in italics ; those for liquids in ordinary type ; and those for solids 
in bold type. It is still doubtfnl whether liquids and solids possess 
sucli simple formulae ; it is the author's opinion that in many 
case's they do ; but there are certainly many cases in which they 
possess more complex formulae. There is, however, as yet no 
method of determining with certainty the degree of complexity ; 
hence, the simplest formulae are employed. Liquid hydrogen 
chloride may have the formula HC1 ; or it may have the formula 
(HCl)w; but what the value of n is, there is no means of deter- 

The reactions, whereby the halides of hydrogen are prepared, 
are represented thus : 

la. H% + F = 277Fathigh temperatures (see p. 115). 

b. H. 2 + cL = 2HCI. 

c. JL 2 + L 2 = 2HT. 

2a. 2ir 2 O + 2C1 2 = 4,1101 + O 2 , or 2H 2 + 2C7 2 = 4HCI + O 2 . 

b. HoS + I 2 + Aq = 2IIT.Aq + S. (Aq = a^wa, water). 
Hydrogen sulphide. Ilydriodic acid. 

c. 2H 3 N Aq + 3C/ 2 = GIICl.Aq + N> 2 . 

d. CI^^ + C1 2 = CH^Cl + HCl. 
Methane. Cliloro- 


3a. 2P + 5Br,.Aq + 8H 2 O = 2n 3 PO 4 Aq + lOHBr.Aq. 
Phosphoric acid. 

b. 2M&C1 2 + HO = MgCl 2 MO + 2HCI. 

Magnesium Magnesium 

chloride. oxychloride. 

c. CaF 2 -f H 2 SO 4 + CaSO 4 + 2HF. 
Calcium Sulphuric Calcium 
flubnde, acid. sulphate. 

d. NaCl -H H 2 S0 4 = NaHSO 4 + HCl. 
Sodium Sulphuric Sodium hydrogen 

chloride. acid. sulphate. 

NaCl + NaHS0 4 = N^SC^ + HCl. 

Sodium Sodium hydrogen Sodium, sulphate, 
chloride. sulphate. 

e. NaBr + H 3 FO 4 = NaH 2 PO 4 + HSr. 

Sodium Phosphoric Dihydrogen sodium 
bromide. acid. phosphate. 

The action of hydrogen bromide or iodide on hot sulphuric acid is repre- 
sented thus : 

H 2 SO 4 + 2HSr (or 2HI) = SO 2 + 2II 2 O -f J0r 2 (or J 2 ). 


4. KF.HF = KP + HF. 

Hydrogen potassium Potassium 

fluoride. fluoride. 

Properties. Hydrogen fluoride is a colourless very volatile 
liquid, boiling at about 19 under atmospheric pressure ; hydrogen 
chloride, bromide, and iodide are all colourless gases. Hydrogen 
fluoride is fearfully corrosive ; a drop on the skin produces a 
painful sore, and several deaths have occurred through inhaling 
its vapour. The other three gases are suffocating, but do not 
produce permanent injury when breathed diluted with air. They 
condense to liquids at low temperatures. They are exceedingly 
soluble in water, in all probability forming compounds which mix 
with excess of water or of the halide. One volume f water at 
dissolves about 500 times its volume of hydrogen chloride ; the 
solution is about 1*21 times heavier than water, and contains 
42 per cent, of its weight of the gas. On cooling a strong solution 
of hydrogen chloride in water to 18, and passing into the cold 
liquid more hydrogen chloride, crystals of the formula HC1.2H 2 O* 
separate out. It is probable that, at the ordinary temperature, this 
compound exists in an aqueous solution of hydrogen chloride, and 
is decomposed into its constituents to an increasing extent with 
rise of temperature. Hydrogen fluoride, bromide, and iodide are 
also exceedingly soluble in water, and their solutions probably 
contain similar hydrates. The corresponding compound of 
hydrogen bromide, HBi '^MI 2 0, has been prepared; it melts at 
11. The solutions of these compounds are termed hydrofluoric, 
hydrochloric, hydrobromic, and hydriodic acids. When saturated, 
they are colourless fuming liquids ; they possess an exceedingly 
sour taste, and are very corrosive ; they change the blue colour of 
litmus (a substance prepared from a lichen named Iccanora tinctoria, 
and itself the calcium salt of a very weak acid) to red, owing to 
the liberation of the red-coloured acid. This is the usual test for 
an acid. 

The great solubility of hydrogen ehloride may be illustrated by help of the 
apparatus shown in the figure. (Fig 24.) 

The lower flask is filled with water coloured blue with litmus ; the upper flask 
is filled with hydrogen chloride by downward displacement, and inverted over 
the lower flask. The stopcock is then opened, establishing communication 
between the two flasks. By blowing through the tube, a little water is forced 
up into the hydrogen chloride. It immediately dissolves, producing a partial 
vacuum in the upper flask ; and the pressure of the atmosphere causes a 
fountain of water to enter it. The blue colour of the litmus is at the same time 
changed to red. 

Comptet rendvs, 86, 279. 



All elements are attacked and dissolved by these acids, 
hydrogen being liberated, while the halogen combines with the 
eleinemt, with the exception of : Silver, gold, mercury ; boron 
(attacked by hydrofluoric acid), carbon; silicon, zirconium (both 
attacked by hydrofluoric acid), lead; nitrogen, vanadium, phos- 

phorns, arsenic, antimony, bismuth ; molybdenum ; oxygon, 
sulphur, selenium, tellurium, and the elements of the platinum 
group. Mercury and lead are attacked by strong hydriodic acid; 
moist hydrogen chloride, bromide, and iodide are decomposed by 
light in presence of oxygen.* The first two are not decomposed 
when dry ; dry hydriodic acid, however, yields water and iodine. 

Use's. Hydrofluoric acid is employed for etching on glass. 
The glass is protected by a coating of beeswax, and a pattern is 
drawn on the wax. The article is then dipped in the strong acid, 
and the pattern remains after removing the wax, the glass 
appearing frosted where the acid has attacked it. Hydrochloric 
acid is used for many purposes, one of the chief of which is the 
manufacture of chlorine and the chlorides of metals. 

* Chem. Soc., 61, 800. 


Proofs of the Volume-Composition of the Halides 

of Hydrogen. 

It has already been shown that hydrogen chloride con- 
sists of equal volumes of hydrogen and chlorine united with- 
out contraction ; it may be shown to contain its own volume of 
hydrogen by the following experiment: A U-tube, as shown in 
fig. 25, is filled with mercury, which is then displaced in the 

FIG. 25. 

closed limb by gaseous hydrogen chloride. The level of the 
mercury is made equnl Jn the two limbs, and the position marked. 
The open limb is then filled with liquid sodium amalgam (an alloy 
of mercury and sodium containing about 2 per cent, of sodium) 
and closed with the thumb. The tube is then inverted, so as to 
bring the gas into contact with the sodium amalgam. The sodium 
reacts with the hydrogen chloride, liberating hydrogen, thus : 

2IICI + 2Na = 2NaCl + H z . 

The hydrogen is then again transferred into the closed limb by 
inclining the tube, and the levels again equalised ; it will be seen to 
occupy half the volume originally occupied by the hydrogen chloride. 

That hydrogen bromide and iodide yield half their volume of 
hydrogen when similarly treated has also been proved. Hydrogen 
fluoride has been synthesised by heating silver fluoride with 
hydrogen gas. The product occupied at 100 twice the volume of 
the hydrogen employed for its formation. The equation is 2AgF -f 
H* = 2HF -h 2Ag, silver being set free as metal. 

It is argued that hydrogen fluoride, bromide, and iodide possess 
respectively the formulae HF, HBr, and HI, from these experiments*, 



from their densities, and from their similarity to hydrogen 
chloride. Recent experiments have, however, shown that at low 
temperatures gaseous hydrogen fluoride has a greater molecular 
weigtil than that expressed hy the formula JTF; but the- actual 
degree of complexity is not yet certain (see below). 

Physical Properties. 
Mass of 1 c.o. 



Gas. IT - 1. 

Hydrogen fluoride . . 


988 at 12 

7 See below See below 

Hydrogen chloride . . 



0001633* 18'23 

Hydrogen bromide . . 



0' 003620* 40-47* 

Hydrogen iodide .... 



0-005727* 63 92* 



Specific Molecular 



Heat. Weight. 

Hydrogen fluoride. . 



? 20 (soo below) 

Hydrogen chloride . 



0-]?(H(gas) 3'46 

Hydrogen bromide . 



? 80 95 

Hydrogen iodide. . . 



? 127 '85 

Heat of formation 77., 


C7 8 = 27/rv + 440K. 

7??- 2 -= 2777?r *- 242K. 

7, - 27/7 + OK at about 184. 

Note. Molecular weight of hudroqen Jlnor*de.\ Tlie vapour-dcnHity of 
hydrogen fluoride increases with fall of temperature, implying the nwormtion of 
molecules of 77F to form (HF)n (The value of w appear* to be k) The 
highest density was found at atmospheric prefrmre, nnd at 26 1, to be 25 59, 
implying the molecular weight of 51 '18 Tins corrcwpouds to a mixture of 
81 -2 i per cent molecules of 77 4 F 4 and 18 70 per cent, of moleculcfl of JfF At 
100 and above, the density is normal, and corresponds to the formula IFF. 

Compounds of the Halogens with Lithium, 
Sodium, Potassium, Rubidium, and Caesium 
- (Ammonium). 

Sources. Sodium fluoride occmrs native in Greenland in com- 
bination with aluminium fl uoride, as cryolite, 3NaF.AlF 3 . Lithium, 
sodium, and potassium chlorides, bromides, and iodides occur in sea- 
water; sodium chloride in by far the greatest amount about .75 
per cent. ; and also in many mineral springs. That at Diirkheirn, 
in the Bavarian Palatinate, is comparatively rich in ccosium and 
rubidium chlorides, and was the source from which Bunsen and 

* These numbers are calculated. 

t Thorpe, Chem. Soc., 63, 765 j 65, 163. 

I 2 


Kirchhoff extracted these elements for the first time.* TheWheal 
Clifford spring, in Cornwall, is specially rich in lithium chloride. 
Sodium chloride also occurs as rock-salt in mines, in various parts 
of the world ; the largest in Britain are in Cheshire, but recently 
other deposits have been discovered near the Tyne. Very large 
deposits of potassium chloride occur at Stassfurth, near Magdeburg, 
in N". Germany. It also occurs in kelp, the ash of fucus palmatus, 
species of seaweed. The ash of the beetroot contains about 
0*17 per cent, of rubidium chloride. 

Preparation. 1. By direct union of the elements. This 
takes place with great loss of energy (i.e., evolution of heat) ; the 
elements take fire and burn in chlorine gas. Perfecljy dry chlorine, 
bromine, or iodine, however, does not act on sodium in the cold.f 
A subchloride of a purple colour is said to be produced by the 
action of chlorine on metallic potassium. 

2. By double decomposition, (a.) Action of the halogen 
acids on the oxides, hydroxides, or carbonates of the metals ; 
in the first two cases, the hydrogen of the halogen acid unites with 
the oxygen of the oxide, or the hydroxyl (a name applied to the 
group OH) of the hydroxides ; in the third case, carbon dioxide 
and water are liberated. Examples of this action are given in 
the following equations : 

KOH + HP = KP + H-OH. 
Na,O + 2HCI = 2NaCl + H,0. 
Li 2 CG 3 + 2HI = 2LiI + H a O + CO Z . 

These reactions also occur in solution. 

On adding to a solution of a hydroxide, containing an unknown 
quantity of the hydroxide, a solution of a hydrogen halide, the 
completion of the reaction, or the "point of neutralisation," may 
be ascertained by the addition of litmus, or of phenol-phthalein, 
to the hydroxide ; the former gives a blue, the latter a cherry- 
red colour with these hydroxides ; when the colour is on the point 
of changing to brick-red, with litmus, or being discharged entirely, 
with pheiiol-phthalem, the reaction is complete, and there is no 
excess either of acid, or of alkali, as such hydroxides are named. 

If carbonates be used, the solution must be boiled during the 
addition of acid, so as to expel carbon dioxide gas, else it will 
produce a colour change. 

(6.) By certain other "double decompositions;" thus 
sodium chloride is obtained as a by-product in the manufacture 

* Poffff. Ann., 110, 161 ; 113, 337 ; 119, 1 ; Annales (3), 64, 290. 
t JBtrichte, 6, 1518 ; Chem. Soc., 43, 155. 


of potassium nitrate from sodium nitrate and potassium chlor- 
ide : 

I KCl.Aq + NaN0 3 .Aq = KN0 3 .Aq + NaCl. 

The sodium chloride, being much less soluble in water than 
potassium nitrate, separates in crystals on evaporation. The 
sulphides and hydrosulphides of these metals also yield halides 
on treatment with halogen acids. 

3. By heating compounds of these metals with oxygen 
and with the halogens, e.g., chlorates, iodatcs, &c. (see p. 4G6). 

4. Compounds of ammonium with the halogens are pre- 
pared by addition of the halogen acid to a solution of ammonia in 
water. Direct combination ensues, thus : 

NH 3 .Aq + HCl.Aq == NH 4 Cl.Aq. 

Ammonium, NH 4 . 

The group of elements to which the name ammonium has been 
given exhibits the greatest similarity to metals of the sodium group, 
and is usually classed along with them. It has never been isolated 
(see, however, pp. 577, 578). But ammonia, consisting of one atom 
of nitrogen and three atoms of hydrogen, JV7/ 3 (see p. 512), has 
the power of union with acids (as well as with oxides and double 
oxides); compounds of ammonium with the halogens differ from 
those of sodium and the other metals by splitting up when 
heated into ammonia and the hydrogen halide. 

The union of ammonia with a halide of hydrogen may be 
illustrated by placing a jar filled with ammonia gas (sec p. 512) 
mouth to mouth over a jar of hydrogen chloride, both being 
covered Vith glass plates; when the plates are withdrawn, dense 
white fumes of ammonium chloride are seen ; they gradually settle 
in the lower jar as a white powder. 

The decomposition of this compound by heat may be shown by 
applying heat to a fragment in a platinum basin ; it will volatilize 
completely, being decomposed into its constituents ammonia, NH 3 , 
and hydrogen chloride, HCl ; they unite when cooled by the air, 
forming dense white fumes. 

Special methods of extraction and preparation. Owing to their im- 
portance, the following compounds require consideration: Common salt, or 
sodium chloride, is produced by the evaporation of sea- water in " salt pans," 
. shallow ponds exposed to the air. To promote evaporation, the salt water is 
sometimes allowed to trickle over ledges, running into gutters wltich lead it to 


ihe ponds. When a portion of the water has been thus removed, it is boiled 
down in shallow iron pans. Rapid evaporation produces fine-grained salt, such 
as is used for the 'table j slow evaporation causes the salt to separate m, larger 
crystals; it is used for curing fish, &c. 

In Cheshire, water is run into the mines, and the brine is pumped up and 
evaporated. In cold climates, the salt is sometimes extracted from sea- water by 
freezing j the ice which separates is nearly pure, while the salt remains dissolved 
in the last portions of water. 

Potassium bromide and iodide are prepared (a) by the action of bromine 
or iodine on a solution of potassium carbonate ; the water is removed by evapora- 
tion, and the residue is heated to redness (see p. 467) ; or (6) by treating iron 
filings with bromine or iodine, producing ferrous bromide or iodide, to which a 
solution of potassium carbonate is then added. The resulting ferrous carbonate is 
insoluble in water ; it is removed by filtration, and the filtrat^ is evaporated to 
dryness. The equations are : 

Fe + Br 2 + Aq FeBr 2 .Aq ; 

FeBr^Aq + K 2 C0 3 .Aq = 2JOr.Aq + FeCO 3 . 

The equation for the preparation of potas&ium iodide is similar. 

Properties. These substances are all white solids, crystallis- 
ing in the cubical system, with the exception of caesium chloride, 
which crystallises in rhombohedra. They are all soluble in water ; 
lithium chloride, sodium bromide, and sodium and potassium 
iodides are also soluble in alcohol. 

100 grams of water dissolve at the ordinary temperature (about 15) 

Fluoride. Chloride. Bromide. Iodide. 

Lithium trace 

Sodium 4 36 88 373 

Potassium 33 65 143 

Ammonium 37 72 

grams of these salts. 

They all form double compounds with water, e.g., NaC1.2H 2 O, 
crystallising at a low temperature. They melt at a red heat and 
volatilise at a bright red heat ; ammonium chloride dissociates at 
339, under ordinary atmospheric pressure, into hydrogen chloride 
and ammonia ; the other compounds of ammonium behave similarly. 

Melting-points. Mass of 1 c.c. 

R Cl. B^ L Cl. Br. ~L 

Lithium 801 598 547 446 2 -29 2 "00 3-10 3 '48 

Sodium 902 772 708 628 2 '56 2 '16 308 3*65 

Potassium .... 789 734 699 634 2 '10 1 '98 2 -60 3 '01 

Kubidium.,.. 753 710 683 642 3 '20 2 '80 3 '36 3 '57 

Caesium ? 4'00 4-46 4*54 

Ammonium - ? 1 '52 2 '46 2 '44 


The vapour densities of the following compounds have recently 
been determined at about 1200 by V. Meyer's method : 

Fond ........ KI = 184-1 ; RbCl = 139-4 ; Rbl = 221-6 

Calculated .... KI = 166'0 ; RbCl = 121-0 ; Rbl = 212*3 

Found ......... CsCl = 179-2; Csl = 267; 

Calculated ...... CsCl = 168-4; Csl = 259'8* 

These numbers represent molecular weights, i.e., vapour- 
densities multiplied by two. 

It may be concluded from analogy that the other halides, in 
the gaseous state, have also simple formulas, such as NaCl ; at 
present we know nothing about the molecular weights of these 
bodies in the liquid or solid state. 

Double compounds. 1. With halogens Iodine unites directly with 
potassium iodide in aqueous or alcoholic solution, and forms daj-k lustrous prisms, 
possessing the formula KI 3 .f The mass of 1 c.c. is 3*50 grams at 15. It melts 
at 45. Chlorine and bromine are more soluble in solutions of chlorides and 
bromides than in pure water, owing probably to the formation of similar 
compounds, which are partially dissociated at the ordinary temperature. 
Ammonium tri-iodide and tribromide, (NH 4 )I 3 and (NH 4 )Br 3 , have been 
prepared by a similar method, and are closely analogous. 

2. With hydrogren halides. Potassium fluoride unites with hydrogen 
fluoride in three proportions, forming (a) KF.HF, (b) KF.2HF, and (c) 
KF.3HF.J They are all stable in dry air, but decompose when heated into 
potassium and hydrogen fluorides. No doubt, similar compounds of the other 
halogen salts would prove stable at low temperatures. 

For compounds of the formula 4NH 3 .HC1, and 7NH 3 .HC1, see p. 525. 

Heats of formation 

Li + 

Na + 
Na + 
Na + 
K + 

K + 
K + 




= LiCl 4- 
= NaCl + 
= NaBr -t- 
- Nal + 
- KC1 + 
= KBr + 
= KI + 




Aq = 
Aq = 
Aq = 
Aq = 
Aq = 
Aq = 
Aq = 

+ 84K. 
-51 -IK. 

* Scott, Brit. Assn., 1887, 668 ; Proc. Roy. Soc. Edin., 14. 

t Chem. Soc., 31, 249 ; 33, 397 ; Berichte, 14, 2398. 

J Comptes rendus, 106, 547. 

For an explanation of " K," see p. 127. 






Beryllium, Calcium, Strontium, and Barium 

Sources. Calcium fluoride, or fluor-spar, CaP 2 . This 
beautiful mineral, crystallising in cubes, sometimes showing octa- 
hedral modifications, occurs in granite and porphyry rocks, 
especially where the veins border other strata. It forms the 
gangue of the lead-veins which intersect the coal-formations of 
Northumberland, Cumberland, Durham, and Yorkshire; it is 
abundant in Derbyshire and also in Cornwall, where the veins 
intersect much older rocks. A large vein occurs in Jefferson Co., 
New York State, in ^vanular limestone. It often possesses a pink, 
amethyst, or green colour, from the presence of certain metallic 

Calcium chloride is a constituent of all natural waters, and 
exists in small amount in sea-water. Traces of the chlorides of 
strontium and barium are also found in some mineral waters. 

Preparation. The methods of preparation are similar to those 
of the halides of the alkali-metals. 

1. By direct union of the elements. The metals of this 
group are so difficult to prepare that the method is impractic- 

2. By double decomposition. (a.) The action of the haloid 
acid on the oxides, hydroxides, sulphides, or hydrosulphides, or on 
double oxides, such as carbonates, silicates, &c. This method 
serves for the production of the chlorides, bromides, and iodides ; 
not well for the fluorides, for the fluorides of calcium, strontium, 
and barium, are insoluble in water, and the hydroxide or carbonate 
becomes coated over with the insoluble fluoride, and action ceases. 
The reactions may be typified by the following equations : 


BeO + 2HCI = Bed, + H 2 0. 
Ca(OH), + 2HCI = CaCl, + 2H- 2 0. 
SrCO 3 4- 2HCI = SrCl 3 -f H 3 + 00 8 . 
BaS -f 2HCI = BaCla -f 

These reactions occur both in solution and with the dry 

This process is practically made use of in preparing strontium 
and barium chlorides, from their carbonates and sulphides. 

(fe.) The fluorides of calcium, strontium, and barium, being 
insoluble in water, may be precipitated by adding a soluble fluoride, 
such as potassium fluoride, to a soluble salt of one of these metals, 
such as calcium chloride, barium iodide, &c. The reaction is, for 
example : 

CaCl 2 .Aq + 2KF.Aq = CaF, -f- 2KCl.Aq. 

Potassium chloride is soluble in water, and may be separated 
from the insoluble calcium fluoride by filtration. 

Doubtless similar reactions occur on mixing soluble compounds 
of the other halogens with soluble compounds of these metals ; thus 
it may be supposed that 

2KI.Aq + BaCl 2 .Aq = 2KCl.Aq + BaI,.Aq. 

But as all the compounds concerned in the change are soluble 
in water, they cannot be separated. It is probable that such 
changes are only partial ; i.e., not all the potassium iodide is con- 
verted into chloride, nor all the barium chloride converted into 
iodide, but that after mixture the solution contains all four com- 

This method of " double decomposition,'* i.e., reciprocal ex- 
change, is also practically applied in the preparation of strontium 
and baaium chlorides on a large scale. The chief sources of these 
metals are the sulphates of strontium and barium (see p. 422). 
These substances are heated to redness with calcium chloride, 
when the calcium transfers its chlorine to the strontium or barium, 
itself being converted into sulphate, thus : 

BaSO 4 + CaCl 2 = BaCl 2 -f- CaSO 4 . 

On treatment with water the insoluble calcium sulphate re- 
mains, while the soluble strontium or barium chloride dissolves, 
and may be purified by crystallisation from water. 

Properties. Beryllium fluoride has not been prepared free 
from water ; on attempting to dry the gummy mass obtained by 
its evaporation it reacts with the water (see below). 


The fluorides of calcium, strontium, and barium are white crys- 
talline powders, insoluble in water. 

The remaining halides of this group are all white solids, soluble 
in water. They unite with water, forming crystalline compounds. 
Among these are BeCl 2 .2H 2 O; CaCL.6H,O; SrCl 2 .3H 2 O ; 
BaBr 2 .2H 2 O; arid BaI 2 .7H 2 O. The only one of the halides 
which has been volatilised is beryllium chloride, which becomes 
vapour somewhat below 520 under ordinary pressure. At higher 
temperatures (812) it has the vapour- density 40*42, implying the 
molecular weight 80'02.* Tho compounds of beryllium have a 
sweet, disagreeable taste ; the soluble compounds of the other 
elements are saline and burning. 

Uses. Calcium fluoride is employed as a flux, or material to 
be added to metals to make them flow (fluo) when they are being 
fused. It probably acts by dissolving a film of oxide encrusting 
the globules, and thereby causes the metallic surfaces to come 
in contact and unite. It is also a source of hydrogen fluoride 
(see p. 106). Calcium chloride is employed on a small scale for 
drying gases, and liquid compounds of carbon ; it has a great 
tendency to unite with water, hence it deliquesces on exposure to 
moist air, attracting so much moisture as to dissolve. 

Some of these substances react with water ; hence beryllium 
halides, calcium chloride, bromide, and iodide, and strontium and 
barium bromides and iodides, cannot be prepared pure in an an- 
hydrous state by evaporating their solutions. The reaction is a 
partial one. With calcium bromide, for instance, it is : 

CaBr 2 + H 2 = CaO + 2HBr. 

But the calcium bromide and oxide unite, forming various 
oxybromides, which remain, while a portion of the hydrogen 
bromide escapes. 

Physical Properties. 
Melting-points. Mass of 1 c.c. 

Beryllium .... 

F. Cl. Br. I. 
? 600 600 ? 
902 719 676 631 
902 825 630 507 
908 860 812 ? 

F. 01. Br. 
? ? ? 
3-14 2-20 3-32 
4-21 3-05 3-98 
4-83 3-82 4-23 




Strontium .... 

Nilson and Fetterssen, Comptes rendus, 98, 988. 


Heats of formation : 

Ca + Clt - Oa01 2 + 1698K + Aq - + 174K. 

Ca + Br 2 = CaBr 2 -r 140DK + Aq = + 256K. 

Ca + I 2 ^ CaI 2 + 1073K + Aq = + 277K. 

Sr + CU = SrClj + 1846K + Aq - + 111K. 

Sr + Br 2 SrBr 2 + 1577K + Aq = + 161K. 

Ba + C7 2 = Bad; + 1917K + Aq = + 21K. 

Ba + Br 2 = BaBr + 1700K -f Aq = -t- 50K. 

Double compounds. The scare all prepared by direct addition. Among 
them may be mentioned : BeF 2 .2KF, BeCL.SKCl, and similar compounds 
with sodium and ammonium chlorides, and BaF^BaCL,. The solubility of 
barium and strontium fluorides in hydrofluoric acid is probably due to the 
formation of double compounds with hydrogen fluoride. 

Magnesium, Zinc, and Cadmium Halides. 

Sources. Magnesium chloride, bromide, and iodide are con- 
tained in sea-water, and in many mineral springs. Carnallite, 
MgCl 2 .KC1.6H 2 0, occurs in large quantities at Stassfurth, and is a 
valuable source of magnesium and potassium compounds. 

Preparation. 1. By direct union. The halogens unite 
with these metals directly, even in the cold, to produce halides. 
In presence of water, solutions are obtained. 

2. By the action of the halogen acid on the metal hy- 
drogen is evolved, and the halide of the metal is formed. 

3. By double decomposition. (a.) By the action of the 
halogen acid 011 the oxides, hydroxides, sulphides, and on some 
double oxides, such as carbonates, borates, &c. This process yields 
solutions of the halides (except in the case of magnesium fluoride, 
which is insoluble in water). But the water cannot be removed 
completely by heat, for it reacts with the chlorides, forming oxy- 
chlorides. The double chlorides with ammonium chloride, how- 
ever, are unacted on when evaporated with water, hence anhydrous 
magnesium chloride may be produced by heating the compound, 
MgCl 2 .2NH 4 Cl, to redness ; the ammonium chloride sublimes (see 
p. 117), leaving the anhydrous magnesium chloride. It can also 
be prepared by heating the aqueous chloride in a current of hydro- 
gen chloride. Similar methods would probably succeed with the 
bromides and iodides. 

(6.) Other methods of double decomposition may be sometimes 
employed; e.g., MgS0 4 .Aq -f BaCl 2 .Aq = MgCl 2 .Aq + BaSO 4 . 
Barium sulphate is insoluble, and may be removed by filtration. 
Another method, which succeeds on a large scale, is to heat, under 


pressure, magnesium carbonate with a solution of calcium chloride ; 
the equation 

MgCO 3 + CaCl 2 .Aq = MgCl z .Aq 4- CaCO 3 

represents the reaction, the insoluble calcium carbonate being 
removed by nitration. 

Typical Equations 

1. Zn + Ck = ZnCl,. 

2. Cd + 2HI.Aq = CdI 2 .Aq + H z . 

3. MgO + 2HBr.Aq = MgBr 2 .Aq + H 2 O. 
ZnS -f 2HCl.Aq = ZnCl 2 .Aq + H 9 S. 

CdCO 3 + 2HF.Aq = CdF 2 .Aq + H 2 + CO.. 

Properties. With the exception of magnesium fluoride, the 
hfilides of these metals are soluble in water. They are white and 
crystalline. The fluorides excepted, they are all volatile and are 
decomposed at a red heat by atmospheric oxygen, yielding the 
halogens and oxyhalides. This has been proposed as an effec- 
tive method of manufacturing chlorine. They also react with 
water at a red heat ; the products are oxyhalide and hydrogen 
halide. This method is in operation for the preparation of 
hydrogen chloride ; the equation has been given on p 111. They 
all unite with water, forming crystalline compounds ; for example, 
MgCl 2 .6H 2 O ; MgBr 2 .3H 2 O ; ZnF 2 .4H 2 O ; ZnCl 2 .H 2 O ; 
Cd01 2 .2H 2 O; CdBr 2 .H 2 O ; CdI 2 crystallises as such from water. 
Zinc chloride has prvh a strong tendency to combine with water 
as to be able to withdraw the elements, hydrogen and oxygen, 
from compounds in which they do not exist as water ; thus it 
chars wood and destroys the skin ; it is therefore used in surgery 
as a caustic. They all, except magnesium fluoride, attract mois- 
ture from moist air, and deliquesce. 

Uses. Magnesium chloride is employed as a disinfectant, and 
is also used fraudulently for " weighting " flannel and cotton goods. 
Zinc chloride is also employed as a disinfectant under the name 
of " Burnett's Disinfecting Fluid. " Cadmium bromide and iodide 
are used in photography. 

Physical Properties. 
Mass of 1 c.c. solid. Melting-point. 


Magnesium . 

F. 01. Br. I. 
2-86 2-18 ? ? 

4 -60 2 -75 3-64 4'7 
6-00 3-62 4*8 5*7 












F. 01. 

? ? 

? 680 

P f86l 
? t954 







Oadmium . . 


Another variety of cadmium iodide is known, with the specific gravity 
4'6 or 4 7 ; it has a brownish colour, whereas the usual variety is white. It is 
converted at 50 into the usual modification.* 

Heats of formation. Mg> + C1 2 = Mg-CL, + 1510K + Aq = -t 359K. 

Zn + CL 2 = 2nCl 2 + 972K + Aq = +150K. 

Zn + Hr 2 = ZnBr + 760IC + Aq = +150K. 

Zn + I 2 = ZnI 2 + 492K + Aq = + 113K. 

Cd + (7/ 2 = CdCl 2 -f 932K + Aq = + 30K. 

Cd + Br 2 = CdBr 2 -f 952K + Aq = + 4'4K. 

Cd + I 2 = CdI 2 + 488K + Aq = - 9 'OK. 

Molecular weights. The vapour-densities of zinc chloride and of cad- 
mium chloride, bromide, and iodide nearly correspond to the formula ZnCl% and 
CtfC7 2 ,f there n slight dissociation at the temperatures employed (898 and 
1200) ; cadmium iodide undergoes considerable dissociation at the higher 

Double compounds. 1. "With hydrogen halides. 

2ZnCL HC1.2H>O and ZnCl 2 .HCl 2H 2 O 

are produced in crystals by saturating a concentrated aqueous solution of zinc 
chloride with hydrogen chloride. They decompose on rise of temperature.^ 
2. With halides of the alkali metals. 

(a.) Fluorides. MgF 2 NaF j ZnF 2 2KF. 

(I ) Chlorides Mg-Cl 2 NaCl.H 2 O ; Mg-Cl, KC1 6H 2 O. 

ZnCl 2 NH^Cl ; ZnCl 2 2NH 4 C1 ; ZnCl 2 3NH 4 C1; ZnClo 2KC1 ; 
ZnCl 2 .2NaC1.3H,O ; 2Cd01 2 .2KCl H 2 O ; CdCl 2 2NaCl. 
3H 2 O; CdCl 2 .2NH 4 Cl.H 2 O; CdCl 2 .4NH 4 Cl ; OdCL. 

(c.) Bromides. CdBr 2 .KBr.H,O ; 2CdBr 2 .2NaBr 5H 2 O ; 2CdBr 2 .2NH 4 Br. 
H 2 O; CdBr 2 4KBr; CdBr 2 .4NH 4 Br. 

(d.) Iodides. Zn.I 2 KI ; ZnI 2 .2NH 4 I. 

CdI 2 KI H^O ; CdI 2 .2NaI GH 2 O ; CdI 2 2KI.2H,O ; CdI 2 . 
2NH 4 I.2H 2 O. 

3. With calcium, strontium, and barium halides 

2CdCl 2 .CaCl 2 7H 2 O ; CdCl 2 .SrCl 2 .7H 2 O ; 
CdCl 2 .BaCl 2 .4H 2 O ; CdCl J .2CaCl J 2H 2 O. 
2ZnBr 2 .BaBr 2 ; CdBr 2 BaBr 2 2H 2 O 
2ZnI 2 .BaI 2 ; 2CdI 2 .ZnI 2 .8H 2 O ; 2CdI 2 BaI 2 . 

4. With each other MgrCl 2 .ZnCl 2 .6H 2 O ; MgrCl 2 .2CdClo 12H 2 O. 

These are some of the numerous compounds which have been prepared. 
The ratios between the numbers of atoms of chlorine in the constituents ap- 
pear to be : 2 :1; 2:2; 2:3; 2:4; and 4 : 1. 

* Amer. Chem. Jour., 5, 235. 

t Brit. Assn., 1887, 668; Serichte, 12, 1195. 

I Compt. rend., 102, 1068. 


As examples vre may select : 2 : 1 ; Mg:Cl 2 .NaCl ; 2CdCl2 CaCL ; 
2 . 2 CdClj2NaCl ; CdBr 2 .BaBr., ; 2 : 3 ZnCl 2 3NH 4 C1 ; 

2 : 4 CdCl 2 4KC1 ; CdCl 2 .2CaCl 2 ; 4 : 1 2ZnCl 2 HC1. 

These bodies are all prepared by direct addition, concentrated aqueous8olu- 
tions of their constituents being added to one another. 

Concluding remarks on these groups. Molecular for- 
mula*. It lias been seen that whereas the metals of the alkalies 
combine with the halogens in the ratio 1:1, as a rule, e.g., 
NaCl, those of the beryllium and magnesium groups display the 
ratio 2:1, as for example, CaCl 2 , BeCl 2 . The inquiry no ay bere 
be made : How is this known to be the case ? To take a specific 
instance : We know, from the densities of gaseous HCl, HBr, 
KBr, Rbl, &c., that these compounds contain an atom of each 
element ; the vapour-density of zinc chloride has been found to 
correspond to the molecular weight 136*37 ; now subtracting 
35*46 x 2, corresponding to the weight of two atoms of chlorine, 
the remainder, 65 '45, is the relative weight of an atom of zinc, 
provided the compound contains only one atom of zinc. But how is 
this known ? Might not its formula be Z-u^Cl^ ? In which case 
65'45 would represent the relative weight of two atoms of zinc, 
and 32' 72 that of one. And if such a question may be asked in 
the case of zinc, where we know the molecular weight of one of 
its compounds in the gaseous state, the uncertainty in the case of 
barium would appear to be much greater, for in this instance no 
compound has ever been gasified. 

The answer to this question is to be found (1) in a study of 
the specific heats of these elements, and (2) in their position in 
the periodic table. These will now be considered in their order. 

1. Specific Heats of Elements. 

The data for these have been given in the tables of j^Jiysical 
properties appended to the description of the groups of elements. 

The specific heat of a body is defined as the amount of 
heat required to raise the temperature through 1, compared 
with the amount of heat required to raise the temperature 
of an equal weight of water through 1. Or, as water is 
chosen as unit of weight as well as of specific heat, specific heat 
may be defined as the amount of heat required to raise the tem- 
perature of 1 gram of a body through 1. Bat the specific heat 
of water is not constant; more heat is required to raise a gram of 
water from 99 to 100, than from to 1. Hence the unit is now 
generally accepted to be the hundredth part of the heat required 
to raise the temoerature of 1 gram of water from to 100. This 


happens nearly to coincide with the value of I heat unit at tbe 
temperature 18. Such a heat unit is termed a calory, and its 
abbreviated symbol is c. Where large amounts of beat are in 
question a unit of 100 calories is often used, and is represented by 
the letter K. This unit is convenient in expressing heat changes 
which take place during chemical action. 

In 1819, a simple relation was discovered by Dulong and Petit 
to exist between the amount of beat required to raise the tempera- 
ture of 1 gram of each of the following thirteen elements through 
1 : copper, gold, iron, lead, nickel, platinum, sulphur, tin, zinc, 
bismuth, cobalt, silver, and tellurium. 

Dulong and Petit's law. The specific heats of the ele- 
ments are inversely proportional to their atomic weights, 
approximately, or 

(Sp. Ht.) A x (At. Wt.) A = (Sp. Ht.) B x (At. Wt.) B . 

Now the product of the specific heat of an element, or heat re- 
quired to raise the temperature of 1 gram of the element through 1 
into its atomic weight, is termed its atomic heat. For instance, the 
atomic weight of sodium is 23, and its specific heat 0*293 ; and 
the atomic weight of lithium is 7, and its specific heat 0*941. The 
product of the first pair, 23 X 0*293 = 6*74 calories, represents 
the amount of heat necessary to raise the temperature of 23 grams 
of sodium through 1 ; and the product of the second pair, 
7 X 0*941 = 6*59 calories, is similarly the amount of heat re- 
quired to raise the temperature of 7 grams of lithium through 1. 
But 23 and 7 are the relative weights of the atoms of sodium 
and lithium ; and to raise these relative weights expressed in 
grams through 1 requires 6'74 and 6'59 calories respectively ; 
these numbers are approximately equal. Hence the conclusion 
from this and similar instances, that the atomic heats of the 
elements are approximately equal. 

This law is not without apparent exceptions, as, for example, 
in the cases of beryllium, boron, carbon, and silicon, but it holds 
closely enough to be a valuable guide in selecting the true 
atomic weights. It appears also to apply only to solids. As 
regards the real meaning of this law, we have at present no 
knowledge. We can form no probable conception of the change 
in the motion or position of the atoms in a molecule due to their 
rise of temperature; but it is a valuable empirical adjunct for the 
purpose mentioned. 

The product of atomic weight and specific heat, in the instances 
given, is approximately 6*5 ; in other cases it falls as low as 5*5. 


It may be stated then, that this product is approximately a con- 
stant, not differing much from the number 6. Hence 6/specific 
heat of any element should approximately equal its atomic weight ; 
and conversely 6/atomic weight, should give an approximation to 
its specific heat. 

As the atomic weight of hydrogen is 1, its atomic heat should 
be 6, and should be identical with its specific heat. Solid hydro- 
gen, however, has never been prepared. But it forms a solid 
alloy with palladium ; and as the specific heat of an alloy is the 
mean of those of its constituents, that of solid hydrogen has been 
indirectly determined. It has been found equal to 5*88, a suffi- 
ciently close approximation to 6. f 

To return now to the atomic weights of members of the beryl- 
lium and magnesium groups ; the following table gives their 
atomic heats : 

Name. Weight Specific Heat. Atomic Heat. 

Beryllium 9'1 x G'6206 (at 500)* = 5'65 

Calcium 40'08 x 0'167 = 6'69 

Strontium.. 87'5 x ? = ? 

Barium 137'00 X ? = ? 

Magnesium . . . 24-30 x 0*250 = 6*07 

Zinc 65-43 x 0*095 = 6*22 

Cadmium . . . . lljj'l x 0'056 = 6'28 

At 100 the specific heat of beryllium is 0*4702 ; its atomic 
heat is therefore 4*28. It was for long doubtful whether beryllium 


had not the atomic weight 13*65, i.e., 3 x ^-~- ; the formula of its 


chloride would then have been BeCU, and its atomic heat 
0'4702 X 13*65 = 6*42, agreeing with those of many other 
elements; but its vapour-density decided the question. A sub- 
stance of the formula BeCl s should have had the vapour-density 
{13*65 + (3 X 35-46)} 2 = 60*01. Actual experiment gave 40'42 
(see p. 122), hence its molecular weight is 80*84 (9*1 + (2 x 35*46) 
= 80'02).f 

The atomic weights of calcium, magnesium, zinc, and cadmium 
given in the table, correspond, it will be seen, with the usual 
atomic heat. 

2. The similarity of the metals calcium, strontium, and barium, 
and of their compounds, lead to the inference that they belong to 
the same group of elements, hence they find their position in the 

* So p. 88. f Compte* rend., 98, 988. 



periodic table. The atomic weights are deduced from this simi- 
larity, and from their position in the table (see p. 22). 

For these reasons it is concluded that the general formula of 
the halides of this group of elements is MX 8 , where M stands for 
metal and X for halogen ; that of the members of the lithium 
group is MX. Lithium and its congeners are termed HiunaJ or 
inonovalent elements in these compounds; beryllium, magnesium, 
and elements of their groups, are termed dyad or divalent in their 
compounds. But it has been amply shown that Valencia ns flio 
property of acting as a monad, a dyad, a triad element is termed, is 
not a constant quality of any element ; nor in such compounds as 
KI<5, or in the double halides mentioned, can we tell how the 
atoms are hel?l together, whether the metal attracts halogen, or 
halogen attracts halogen, or both attracts both. \Vc arc at present 
without any satisfactory theory to account for such compounds, and 
must, in the meantime, simply accept the fnct of their existence. 

The specific heats of some elements may be simply detei mined \\ith fair 
approximation by the (< method of mixture," and l)uloit and iVtit's \\\\\ may 
be easilr illustrated. A cylindrical can of thin sheet brHs nerves us u calori- 
meter (Mir 2(>) It should have a capacity of about 300 culm < < nhim I r8. 
Huving placed in it 200 cubic centimetres of wafer, the tempi i, it nn- <if the 
water is accurately ascertained \>\ .1 d< IK ,ite thei IMOHM l< r, ^udmih-d m trnllin 
of a degree. Three small hcun^plieri > <! /me, tin mid Icid < ,K li \M i iL, r lnu^ 100 
grams, airo suspended in a h.itli nt hoilin^ \\,it<r l>\ Hun \\ins The ^inc is 

quickly lifted out and dropped into the calorimeter, the wntiT is stirred with 
the thermometer or with a npecial ntirrer, IKS bliown \\\ 1li< figure, and its tem- 
perature ascertained. Similarly, the amount of heat given up to fre)i supplies 
of cold water by the other two metals, tin and lead, i* found. r l b< ir specific 
Heats may be calculated as follows : 


Rise of temperature of the water x 200 = heat given up to the water by 
100 grams of metal in cooling from 100 to the final temperature of the water. 
Hence, if t - t' rise of temperature, then (t - f) 200 = (100 - t)x, where x = 
capacity for heat of the metal ; and a?/100 = specific heat of the metal. 

This experimental illustration, rough as it is, yields fairly good results, 
probahly because the errors neutralise each other in part. The sources of error 
are (1) Hot water is carried over by the metal into the calorimeter ; (2) heat 
is lost by the metal during its transit ; (3) no allowance is made for the capacity 
for heat of the metal of the calorimeter ; and (4) no correction is made for the 
loss of heat of the calorimeter by radiation. 



Boron, Scandium, Yttrium, Lanthanum, and 
Ytterbium Halides. 

Of these elements boron is the only one the halides of which 
are well known. 

Sources. None of the haloid compounds of these elements 
exist in nature. 

Preparation. 1. By direct union. Boron burns when 
heated in chlorine gas, producing the chloride BCk; the bromide 
may also be prepared by passing bromine vapour through a tube 
in which amorphous boron is heated to redness. The iodide is 

2. By the simultaneous action of chlorine or bromine 
and carbon (charcoal) on the oxide at a bright red heat. 
The carbon withdraws the oxygen, producing carbon monoxide, 
while the halogen unites with the boron ; thus : 

B 2 8 + 3C + 3(% = 2BCI, + SCO. 

An intimate mixture of sugar-charcoal, oil, and boron oxide is 
made into balls, and ignited to carbonise the oil out of contact with 
the air. They are then heated to bright redness in an atmosphere 
of halogen. 

Carbon monoxide is a gas, very difficult to condense; boron 
chloride and bromide are liquids at the ordinary temperature ; 
hence by leading the products through a freezing-mixture, the 
halide condenses. The halides of the other elements may be simi- 
larly prepared ; but as they are solids, difficult to volatilise, they 
remain mixed with the surplus carbon. 

3. By double decomposition. (a.) The action of the halo. 
g6n acid on the oxides or hydroxides. This is the usual ijiethod of 



preparing boron fluoride. The hydrogen fluoride is prepared from 
calcium fluoride and sulphuric acid (see p. 108), and while being 
formed acts on boron oxide contained in the mixture. 

The first action is 3CaP 2 + 3H 2 S0 4 = 3CaSO 4 + 6Hf; and 
the second B,O 3 + 6HF = 2BF 3 + 3H 2 O. The water produced 
would decompose the boron fluoride, were it not that it combines 
with the sulphuric acid (see p. 415), and it is thus withdrawn 
from the action. The other hydrogen halides have no action on 
boron trioxide. With other oxides of the group, and with the 
hydroxides, aqueous solutions of the halogen acids yield halides. 

(6.) Boron chloride may be produced by heating together 
phosphorus pentachloride, PCl a , and boron trioxide ( in sealed tubes 
to 150. The equation 6PC1 8 + 5B,O 3 = 3P 2 O 6 + 105C7, ex- 
presses the change. 

Properties. Boron fluoride is a colourless gas very soluble in 
water (1059 volumes at 0). Boron chloride and bromide are 
volatile colourless liquids, the former boiling at 18*23, the latter 
at 90' 5 ; they react at once with water, forming the hydroxide and 
hydrogen halide, thus :BCl 3 + 3H 2 = B(OH) 3 4- 3HCL Boron 
fluoride has such a tendency to combine with water that it with- 
draws hydrogen and oxygen from carbon compounds containing 
them, liberating carbon, and in this respect resembling zinc 
chloride. It also reacts with water ; the first stage of the reaction 
is 2#F 3 + 3H a O = B 2 O 8 .6HF.* On heating the solution, BF* and 
H t O are evolved, and the compound HB0 2 .3HF named fluoboric 
acid remains (see p. 236). On dilution with water, boron hydroxide 
deposits and hydroborofluoric acid is formed, thus : 

4(HBO 3 .3HF) = B(OH) 8 + 3HF.BF 3 + 5H,O.f 
The halides of the other elements of this group are white crys- 
talline substances soluble in water, and decomposed on evaporation 
with water. They are not easily volatile, hence they may be pro- 
duced anhydrous by evaporation with ammonium chloride, as 
anhydrous magnesium chloride is prepared (see p. 123). Yttrium 
iodide is unstable in moist air. 

Heat of formation. B + Cl* - BC7 3 + 1040K. 

Double halides. The double halides of boron fluoride only have been 
studied. It was mentioned above that on heating a solution of boron fluoride, 
some fluoride escapes, but some reacts with the water, giving HF.BF 8 , named 
hydroborofluoric acid. It is also produced by dissolving boron oxide, B^O*, in 
hydrofluoric acid. It is known only in aqueous solution, for on concentration 

* Basaroi*, Comptes rend., 78 1698. 

f C^nsidei able doubt exists regarding these changes (see p. 236). 


hydrogen fluoride is evolved, while boron hydroxide, B(OH) 8 , remains in solu- 
tion, thus: HF.BF, + 3H,O B(OH) 8 + 4JHTF. Compounds with other 
fluorides can also be produced by direct union of boron fluoride with the 
fluorides of these elements ; but such compounds are also formed by the action 
of hydroborofluoric acid on the oxides, hydroxides, or carbonates of the metals. 
They are almost all soluble in water and crystalline. The potassium compound 
has the formula KF.BF 3 ; the barium compound BaF. 2 .2BF 3 .H 2 O. The fine 
compound may be prepared by the action of the hydrogen compound on metallic 
zinc, when hydrogen is evolved, thus: 2HF.BF 3 .Aq 4-Zn ZnF 2 .2BF s + H^. 
These bodies are commonly termed salts of hydroborofluoric acid or boro- 

Aluminidm, Gallium, Indium, and Thallium 

Sources. The only important compound found native is alu- 
ninium fluoride, which, in combination with sodium fluoride, 
forms the white crystalline mineral cryolite, 3NaF.AlF 8 . 

Formation. These elements combine with the halogens in 
several proportions, as seen in the following table: 

Fluorine. Chlorine. Bromine. Iodine. 

Aluminium. A1F 2 *; A1F 8 AlCl, AlBr 8 All,,. 

Gallium ... P G-aF 3 Q-aCl* ; GaOl a f P GaBr 8 P G-alg. 

Indium... ? InF 8 InCl; India ; ? InBr* ? InI 8 . 


Thallium.. T1F; T1F 3 T101; T1CL, ; TlBr ; TlBr 3 ; Til; T1I 8 . 

T101 8 TlBr, 

Preparation. 1. By direct union. The compounds of the 
general formula MX 3 are formed in this way. 

2. By replacement. A1X 3 , GaX 3 , and InX 3 , are produced 
by dissolving the respective metals in the haloid acid; hydrogen 
is evolved; thallium dissolves very slowly, being protected by a 
layer of sparingly soluble halide, forming a thallous salt, T1X. 
By heating indium in dry hydrogen chloride, however, InOl 8 IH 

3. The lower chlorides, GaCl 2 , InCl 2 , and InCl, have been 
produced by heating the higher chlorides with the respective 

4. By double decomposition. (a.) Solution of the respective 
oxides, hydroxides, or sulphides in the haloid acid. 

* Only known in the compound 2NaF. AlF a . 
f Compteg rend., 93, 294 and 329. 


4- GHCl.Aq = 2AlCl 3 .Aq + 3H a O ; 

4- 2HCl.Aq = 2TlCl.Aq + H Z ; 
T1,O 3 + GHCl.Aq = 2TlCl 3 .Aq + 3H a O ; 
In a S a + GHCl.Aq = 2InCl 3 .Aq + 3# a #. 

Thallous carbonate dissolves in baloid acids, giving thallons salts. 
(i.) By precipitation. The chloride, bromide, and iodide of thal- 
lium being nearly insoluble in water, may be prepared by treating 
a soluble compound, e.g., the nitrate, T1N0 3 , with a soluble halide ; 

TINOa.Aq + KI.Aq = Til + KNO 3 .Aq. 

(c.) Aluminium chloride and bromide, like th corresponding 
halides of boron, may be produced by passing chlorine over a mix- 
ture of the oxide and charcoal heated to redness ; or by passing the 
vapour of carbon tetrachloride, CC1 4 , over red-hot alumina. The 
equations are : 

A1,O 3 + 3O + 3C7, = 2AICI, + 300; and 
2Al a O 3 + 3(70/4 == 4AICI* + 300,. 

Properties. MX 3 . These compounds, with the exception of 
InI 3 , which is yellow, T1F 3 , green (?), TlBr 3 , yellow, and T1I 3 , 
red, are colourless crystals ; they are all soluble in water. They 
melt and sublime at comparatively low temperatures. They 
crystallise from water with water of crystallisation. Their solu- 
tions, when evapor,rtt)d, decompose, halogen acid being liberated, 
and an oxy halide being left. The anhydrous halides all attract 
atmospheric moisture. 

MX). Gallium and indium dichlorides are white ; that of 
thallium pale yellow, as also its di bromide. They are attacked by 
water, indium and gallium dichlorides apparently decomposing 
into mono- and trichlorides, thus : ' 

2InCl, + Aq = InCl + InCU.Aq. 

The monochloride in contact with water deposits the metal, 
trichloride remaining in solution, thus : 

SInCI + Aq = InCl 3 .Aq + 2In. 

MX. InCl* is reddish-yellow, and is decomposed by water 
(see above). TIP, T1C1, and TIBr, are white crystalline bodies ; 
TU is yellow. The fluoride is the most soluble, the iodide almost 
insoluble in cold water. They all crystallise from solution in hoi 
water, and do not react with it on evaporation. 

Ckem. Soc. t 63, 820. 


crucible in a current of hydrogen. It is a white insoluble substance, evolving 
hydrogen on treatment with hydrochloric acid.* 

3 .Derivatives of MX. The only representative known is T1F.HF, which 
is protluced by direct addition. It resembles its potassium analogue, KF.HF 
(see p. 119), in being decomposed by heat. 

It has been shown that the compound TILjKI may equally well be produced 
from Til and KI 3 .f We cannot therefore regard it as necessarily composed of 
thalh'c iodide and potassium iodide ; it may equally well be viewed as a compound 
of potassium triiodide, BLI 3 , and t hallow* iodide. Til. In fact we have to confess 
our complete ignorance of the manner of combination of the atoms in the mole- 
cule. It might therefore be better to write the formula KT1I 4 , thus committing 
ourselves to neither view j but simplicity of arrangement is certainly aided by 
the method adopted. 

Chromium, Iron, Manganese, Cobalt, and Nickel 


Sources. None of these compounds is found native except 
ferric chloride, FezCU, which sometimes occurs in the waters of 
volcanic districts. 

These elements, generally speaking, combine with the halogens 
in two proportions, as shown in the following table : 

Fluorine. Chlorine. Bromine. Iodine. 

Chromium ... CrF 3 . CrCl 2 ; Cr01 3 . CrBr 2 j CrBr 3 . CrI 3 . 

Iron FeF 2 ; FeF 3 . FeCl 2 ; FeCl s . FeBr 2 ; FeBr 3 . FeI 2 ; FeI 3 . 

Manganese... MnF 2 ; MnF 3 . MnCl 2 ; MnCl 8 t MnBr 2 ; MnI 2 ; 
MnF 4 . 

Cobalt CoF 2 ; CoCl 2 ; CoCl 3 J CoBr 2 ; CoI 2 j 

Nickel NiF 2 ; NiCl 2 ; NiBr 2 ; 

Manganese forms a tetrachloride, stable in ethereal solution ; chromium a 
hexafluoride, CrF 6 . 

Preparation. 1. By direct union. Chromium and iron 
form dihalides, if the halogen be not in excess ; and trihalides 
with excess of halogen; manganese, nickel, and cobalt, form 
only dihalides. 

2. By the action of the halogen acid on the metals with 
or without presence of water. In all cases the dihalide ia 
formed, thus : 

Fe + 2HC1 = Fed, + H 2 . 

3. By double decomposition. The action of the halogen 

* Chem. News, 59, 75. 

t Johnson, Chem. Soc. t 33, 183. 

X Known only in solution. 



Physical Properties. 

. Mass of 1 c.c. solid. 



F 3 . 01,. Br 3 . I,.' If. 01. Br. 1. F. 01. , Br. I. 

Boron (liquid. ? 1'35 2'69 ? ? ? ? 18'2 90'5 


Aluminium . . 3'1 ? 2'54 2'63 ? ?* 93 125 ? * 264 350 

GhOlium .... ? 236 ? ? ? 75'5 ? ? P 220 ? ? 

at 80 

Indium ..... ? ? ? ? 

Thallium,TlX ? 7'0 7'54 7'8 

Ga01 2 : m.-p., 164; b.-p,, c. 635. 

? * ? ? ? * ? ? 
? 427' 458 439 719 ? ? 800 

1 Heats of formation : 

1. Al + 3C7 --= AlOls -*- 1610K + Aq - 768K. 
Al + 3Br AlBr 3 + 1197K + Aq - 853K. 
Al + 81 = A1I 8 + 704K + Aq = 890K. 

2. Tl + Br = TlBr + 413K. 
Tl + I = TU + 302K. 

Double halides. Of these, only the compounds of aluminium and thallium 
seem to have been prepared. They are all obtained by direct addition, some- 
times, however, being prepared in presence of water, sometimes by fusion. 


1. Derivatives of 1013. 

AlIVSNaF. A1P 3 .2KF. AlCl 3 .NaCl f Similar iodides are 

A1F 3 .3KF. AlP 3 .2KaF. AlBr s .KBr I said to exist. 

T1C1 3 .3NH 4 C1. T1C1 8 .2K01. TlBr 3 .NH 4 Br. 

T1C1 3 .3T1C1. T1C1 3 .T1C1. 

TlBr 3 .3TlBr. TlBr 3 .TlBr. 

T1I 3 .KI. 

Besides these are known : T1I (V 5TU ; 2TlBr 3 .3KBr ; and its analogue, 
2T1I 3 .3KI; also 4AlF 3 .MgrF 2 .NaF, a mineral named ralstonite. The most 
important of these is the mineral cryolite, AlF 3 .3NaF, which is mined at 
Evigtok, in West Greenland, where it forms a deposit 80 x 300 feet in deptli 
and length. It is used as a source of fluorine, of pure alumina, and of caustic 

2. Derivatives of MX2. The compound AlF 2 .2NaF, belonging to this 
group, is an interesting one, inasmuch as it is the only one in which aluminium 
is combined with two atoms of a halogen, or, more comprehensively, the only 
one in which aluminium functions as a dyad (see p. 129). It has recently been 
prepared by heating cryolite with metallic aluminium to redness, in an iron 

* Sublimes without fusing. 


crucible in a current of hydrogen. It is a white insoluble substance, evolving 
hydrogen on treatment with hydrochloric acid.* 

3 .Derivatives of MX. The only representative known is T1F.HF, which 
is protluced by direct addition. It resembles its potassium analogue, KF.HF 
(see p. 119), in being decomposed by heat. 

It has been shown that the compound TILjKI may equally well be produced 
from Til and KI 3 .f We cannot therefore regard it as necessarily composed of 
thalh'c iodide and potassium iodide ; it may equally well be viewed as a compound 
of potassium triiodide, BLI 3 , and t hallow* iodide. Til. In fact we have to confess 
our complete ignorance of the manner of combination of the atoms in the mole- 
cule. It might therefore be better to write the formula KT1I 4 , thus committing 
ourselves to neither view j but simplicity of arrangement is certainly aided by 
the method adopted. 

Chromium, Iron, Manganese, Cobalt, and Nickel 


Sources. None of these compounds is found native except 
ferric chloride, FezCU, which sometimes occurs in the waters of 
volcanic districts. 

These elements, generally speaking, combine with the halogens 
in two proportions, as shown in the following table : 

Fluorine. Chlorine. Bromine. Iodine. 

Chromium ... CrF 3 . CrCl 2 ; Cr01 3 . CrBr 2 j CrBr 3 . CrI 3 . 

Iron FeF 2 ; FeF 3 . FeCl 2 ; FeCl s . FeBr 2 ; FeBr 3 . FeI 2 ; FeI 3 . 

Manganese... MnF 2 ; MnF 3 . MnCl 2 ; MnCl 8 t MnBr 2 ; MnI 2 ; 
MnF 4 . 

Cobalt CoF 2 ; CoCl 2 ; CoCl 3 J CoBr 2 ; CoI 2 j 

Nickel NiF 2 ; NiCl 2 ; NiBr 2 ; 

Manganese forms a tetrachloride, stable in ethereal solution ; chromium a 
hexafluoride, CrF 6 . 

Preparation. 1. By direct union. Chromium and iron 
form dihalides, if the halogen be not in excess ; and trihalides 
with excess of halogen; manganese, nickel, and cobalt, form 
only dihalides. 

2. By the action of the halogen acid on the metals with 
or without presence of water. In all cases the dihalide ia 
formed, thus : 

Fe + 2HC1 = Fed, + H 2 . 

3. By double decomposition. The action of the halogen 

* Chem. News, 59, 75. 

t Johnson, Chem. Soc. t 33, 183. 

X Known only in solution. * 


acid on the oxide, hydroxide, sulphide, carbonate, sulphite, 
&C. With oxides, sulphides, &o., in which the metal acts as a 

dyad, the dihalides are formed, thus : 


FeO + 2HCl.Aq = FeClj.Aq + H a O. 
Mn(OH) 2 + 2HCl.Aq = MnCl 2 .Aq + H 2 0. 
NiS + 2HCl.Aq = NiCl 2 .Aq 4- H 2 S. 
CoCO 3 + 2HCl.Aq = CoClo.Aq + 00 2 + H 2 0. 

If the sesquioxide, dry or hydrated (hydroxide), be employed, 
the trihalides are produced when capable of existence ; if not, the 
halogen is evolved, thus : 

Fe 2 O 3 + GHCl.Aq = 2FeCl 3 .Aq + 3H 2 O. 
Cr(OH) 3 .Aq + SHCl.Aq = CrCl 3 .Aq + 3H 2 0. 
Ni a O 3 + 6HCl.Aq = 2JSTiCl 2 .Aq + 3H 2 O + C1 2 . 
Mn 2 O 8 + GHBr.Aq = 2MnBr 2 .Aq + 3H 2 + Br 2 . 

With a higher oxide of the metal, or a double oxide containing 
such a higher oxide, the highest halide capable of existence at the 
temperature of action is produced, and the halogen is liberated : 
thus, if the solution be cold, 

2MnO 3 + 4HCl.Aq = 2MnCl 3 .Aq + 4H 2 + CZ 2 ; but if hot, 
\ MnOa -f 4HCl.Aq = MnCl a .Aq + 2H 2 O + Cl*. 
Similarly, 2OO 3 .Aq^ 12HI.Aq = 2CrI 3 .Aq + 6H 2 O -f 3I 2 ; and 
K 2 O 2 O 7 .Aq(= K 2 0.2CrO 3 ) H- 14HCl.Aq = 2KCl.Aq + 

2CrCl 3 .Aq -f 7H 2 + 3Z 2 . 

Also, 2KMn0 4 .Aq(= K 2 O.Mn a O 7 ) + IGHCl.Aq = 2KCl.Aq -f 
2MnCl a .Aq + 8H 2 

These last methods, involving the use of higher oxides, are the 
practical methods of preparing the elements chlorine,, bromine, 
and iodine (see p. 75). Fluorine cannot be thus liberated. Hydro- 
gen fluoride either is without action, or it liberates oxygen as ozone, 
or (in the case of manganese dioxide or of chromium trioxide), 
higher fluorides are produced (see p. 142). 

4. With chromium alone, the action of hydrogen at a low 
red heat on the trihalide produces the dihalide, thus : 

2Cr01 3 -f H> = 2CrCl a + 2HCL 

This is best carried out practically by heating a mixture of 
chromic chloride and ammonium chloride to bright redness in a 
porcelain retort. 

On tr^itment with hydrogen at a red heat, the other chlorides 


are reduced to metal; as is that of chromium at a high tem- 

5. By the action of the halogen on a red-hot mixture of 
the oxide and carbon. This method is specially used for pre- 
paring the trihalides of chromium, for the metal is difficult to 
prepare. The halide volatilises, and is thus separated from the 
excess of carbon. 

Properties. Dihalides. These compounds, if anhydrous, 
crystallise in lustrous scales. Their colours are : 

Chromium. Iron. Manganese. Nickel. Cobalt. 

Fluoride.. ? White ? ? ? 

Chloride.. White White Rose Yellow Blue. 

Bromide . White Yellowish Pale-red Yellow Green. 

Iodide... ? Grey White? Dark, metallic Black, lustrous. 

They are all deliquescent, and dissolve in water, heat being 
evolved by the union. They also dissolve in alcohol. They 
crystallise from such solutions, with more or less water of 
crystallisation. They cannot be dried, for they react with water, 
giving oxyhalides. The colours of these compounds with water 

Chromium. Iron. Manganese. Nickel. Cobalt. 

Fluoride ? Colourless Amethyst Green Kose. 

Chloride Blue Blue-green Rose Green Pink. 

Bromide Blue Green Bed Green Bed. 

Iodide ? Green White Green Green. 

Manganous fluoride is insoluble in water, but dissolves in 
aqueous hydrofluoric acid, doubtless forming a double fluoride. 
Almost all these compounds are soluble in alcohol ; manganous 
chloride dissolves with a green colour. The halides of nickel and 
of cobalt undergo a curious change on concentration, or on addition 
of halogen acid ; those of nickel turn yellow ; those of cobalt blue, 
'or green. This is probably due to the formation of the anhydrous 
chloride. The solutions are used as " sympathetic inks." 

When the paper on which they are traced as ink is warmed, a change of 
colour takes place. A very curious effect may be produced by combination of 
ordinary water-colours with such sympathetic inks; a landscape, cleverly 
painted, may be made to show a transition from a winter to a summer scene 
when held before the fire. 

The chromous and ferrous halides, on exposure to air, combine 
with its oxygen, forming chromic or ferric oxyhalides (see p. 257). 
Their solutions, especially those of the chromium halides, rapidly 
absorb oxygen; the oxidation being accompanied by# change of 

Chromium . . 

F. Cl. Br. I. 
? 2-75 ? ?" 
P 2'53 ? ? 

Manganese. . 

? 2-48 ? ? 
? 2-94 ? ? 
2-86 2-56 ? ? 


colour to green, in the case of chromium, and to brown-yellow, in 
the case of iron. Such substances are said to have power of 
" reduction," meaning that they tend to absorb oxygen from 
bodies capable of parting with it, they themselves being 
"oxidised." In presence of halogen acid, such a reaction as this 
occurs : 2FeCl -f 2HC1 + O = 2FeCl 3 + H 2 O ; the oxygen being 
derived from the air, or from any substance capable of yielding it. 
Hence, chromous and ferrous halides are converted into chromic 
or ferric halides, by the action of the halogen in presence of 

Physical Properties. 

Mass of 1 c.c. Melting-points. Boiling-points. 

Unknown. Unknown. 

Hydrated : N1C1 2 .4H 2 O, 2'01 ; FeCl 2 .4H,O, T93 ; CoCl 2 .6H 2 O, 1 84. 

Heats of formation : 

Cr + C/ 2 = OrCl 2 + ? + Aq = ? 
Fe + C7 2 = FeCl 2 4- 821K + Aq = 179K. 
Mm + Cj - Mn01 2 + 1120K + Aq = 160K. 
Ni + C1 2 -^iCLj + 745K + Aq - 192K. 
Co + Cl* = CoOl 2 + 765K + Aq - 183K. 

Double compounds of the dihalides. One hydrochloride is known, viz., 
2H01.3Cr01 2 .13H 2 O ; and crystals, too unstable to be collected, have also been 
obtained bypassing hydrogen cbloride into a cold solution of cobaltous chloride. 
The other double salts may be divided into two groups, of which instances are 
FeF 2 .2KF, FeCl 2 .2KC1.2H 2 O, MnCl 2 2NH 4 Clj also NiCl 2 .NH 4 Cl, and 
MnCl 2 .NH 4 Cl. 

Not many such compounds have been prepared. 

Trihalides. The anhydrous trihalides also form lustrous 
scales. Their colours are 

Chromium. Iron. 

Fluoride Dark green Pale yellow. 

Chloride Pale violet Black. 

Bromide Dark olive green Black. ? 

Iodide ? Black. 

Chromic chloride, after sublimation', is insoluble in cold water, 
but dissolves after long boiling. If prepared by drying the 
hydrated chloride in a current of hydrogen chloride, it is soluble ; 
as soon as i has been sublimed, it is insoluble. The presence of a' 


trace of chromous chloride causes the insoluble variety to dissolve 
at once. The other halides are deliquescent, and readily soluble 
in water. They also, like the dihalicles, react with water, forming 
oxyllalides (see p. 257). 

The trihalides of manganese and cobalt are unknown in the 
anhydrous state. 

The aqueous solutions have different colours, owing, no doubt, to the 
presence in solution of a compound with water. They are 

Chromium. Iron. Manganese. Cobalt. 

Fluoride .... Green Colourless Buby ? 

Chloride . . . Q-reen Yellow Brown-yellow Brown 

Bromide .... Green Brown-red ? ? 

Iodide Green Brown ? ? 

Chromic chloride exists in two modifications, green and violet. 
The green solution has possibly a more complex molecule than 
the violet one. The violet modification is produced from the 
violet sulphate (see p. 426) by double decomposition with barium 
chloride, thus, Cr 2 3S0 4 .Aq + 3BaCl*.Aq = 2Cr01 3 .Aq + 3BaSO 4 ; 
or by dissolving the grey modification of the hydroxide (see p. 252 ) 
in hydrochloric acid. These chlorides probably alt react with 
water, giving oxychlorides. That of manganese, indeed, if much 
water be added, gives a precipitate of sesquioxide, thus : 

2MnCl 3 .Aq + 3H 2 = Mn 2 O 3 .Aq + CHCl.Aq. 

Manganic fluoride, when heated with water, gives off oxygen, 
and hydrogen fluoride, thus : 2MnF 3 .Aq-f H 2 O = 2MnF 2 .Aq + 2HF 
-f 02. Manganese and cobalt trichlorides are very unstable, 
evolving chlorine at the ordinary temperature, thus : 2MnCl 8 .Aq 
= 2Mn,01 2 .Aq -f C7 2 - Ferric chloride is more stable, but it may be 
reduced or deprived of chlorine by means of nascent hydrogen, 
*&.e., hydrogen in process of formation. Hydrogen gas may be 
passed through a solution of ferric chloride without action; but if 
the hydrogen be prepared in a solution of ferric chloride by the 
action of zinc and hydrochloric acid for example (see p. 27), 
the ferric chloride is changed to ferrous chloride, thus : 
FeCl 3 .Aq + H = FeCl 2 .Aq + HCl.Aq. It is supposed, with great 
probability, that the hydrogen is liberated in the atomic condition. 
In presence of ferric chloride it unites with chlorine ; but if no 
reducible substance is present, it combines with itself to form 
molecular hydrogen, H^ which is then without action. Chromic 
4 chloride cannot be easily reduced in aqueous solution. 


Physical Properties. 
Mass of 1 o.c. solid. Melting-point. Boiling-point. 

01. Br. I. F. 01. Br. I. F. Cl. Br. I. 
Chromium .... ? 2'76 ? ? 1 

-, , 
Iron .......... ? 2-80 P ? - TTntnOWn - 

.Hea of formation : 

Fe + C7 8 - Fe01 3 + 961K ; + Aq = FeCl 8 .Aq 4 633K. 

Double compounds of the trihalides. These are made by direct addition 
and helong to the following four types : 

l; CrBr 8 .KBr ; CrI 3 .KI ; FeF 3 .KF. 

These are stable in presence of excess of the hydrogen-halide, but decom- 
pose with water. 

2. CrF 3 .2KF; FeF 8 .2KF; Fe01 3 .2XCl; Fe01 3 2NH 4 C1 ; 

MnF 3 .2KF ; MnF 3 .2NH 4 F; MnF 3 .2NaF; Mn 
3. CrF 8 .3XF. 
4. 2FeI 3 FeI 2 j 2MnF 3 .MnF 2 . 

The green modifications of chromic halides do not form double compounds. 
They are possibly combinations of molecules of the chromium halides with 
each other. 

Higher halides. Manganese tetrafluoride, MnF 4l is produced 
by treating manganese^jiioxide with aqueous hydrogen fluoride, 
thus : 

MnO 3 -I- 4HF.Aq = MnF 4 .Aq + 2H 2 0. 

It is soluble in alconol and in ether. Its aqueous solution, when 
warmed, decomposes, depositing the dioxide, MnO 2 .Aq. On ad- 
dition of a solution of potassium fluoride it forms the double 
compound, 2KF.MnP 4 , as a rose-coloured precipitate. 

Manganese dioxide, suspended in ether, and saturated with 
hydrogen chloride, gives a green solution of MnCl^ 

Chromium hexafluoride, CrF 6 , is produced by the action of 
hydrogen fluoride on chromium trioxide, Cr0 3 , in presence of 
anhydro-sulphuric acid to absorb the resulting water, thus : 

Cr0 3 + 6HF + 3H 8 S 2 7 = CrF fl + 6H 2 S0 4 . 

It is a fuming volatile liquid, of a blood-red colour, which 
attacks ttilicon oxide, and hence cannot fye kept in glass vessels.* 
General remarks. The elements of this group combine with 

* This substance is also said to be an oxyfluoride of the formula CrO a F 3 
(Gazzetta chimica italiana, 16, 218). 


halogens in four different proportions, thus : MX a , MX 3 , MX 4 , and 
MX<j. The higher members are most stable with chromium, 
and ^fche lower ones most stable with nickel. The molecular 
formulae of these bodies have given rise to much dispute. 
Chromium dichloride appears to exist partly as CrCl^ partly 
as Cr^Cli, in the gaseous state at 1600 ; at 1400-1500, ferrous 
chloride possesses the simpler formal, FeClj. Chromic chloride, 
above its volatilising-point, about 1060, has the formula, CrCk ; 
ferric chloride, at temperatures below 620, is Fe^Ch ;* but/ as tem- 
perature rises, these complex molecules dissociate, and at 750 and 
upwards, its density shows it to have the formula, FeCl*.^ The 
molecular weights of the double compounds of these halides are 
unknown, but it appears probable that they possess the simpler 
formulae given them. 

The formulas of these compounds are deduced 

1. From the simplicity of the ratios of metal and halogen: 
viz., 1:2; 1:3; 1:4; and 1 : 6. 

2. From the vapour-densities. 

3. From the atomic heat of the metals. These are : 

Or. Fe. Mn. Ni. Co. 

? 6-27 6-69 6-43 6'31 

* Comptes rend., 107, 301. 

f Zeitschr. Ptys. Chem., 2, 659; Chem. Soc., 53, 814. 




Carbon, Titanium, Zirconium, Ceriiim, and 
Thorium Halides. 

The halides of carbon differ from those of the remaining elements 
of this group, in being more numerous, and in being insoluble in 
water. It appears advisable, in the present state of our know- 
ledge, to include cerium in this group, although its halides do not 
closely resemble those of the other elements of the group. 

Sources. None of these halides occur native, except fluocerite, 
to which Berzelius gave the formula CeF 3 , and tyscmite, 4CeP 3 , 
3LaF 3 . 

These elements form the following compounds with the 

Fluorine. Chlorine. Bromine. Iodine. 

Carbon ____ CF 4 CC1 4 2 Cl fl ; C 2 C1 4 , &c. CBr 4 ; C 2 Br 6 j C 2 Br 4 . CI 4 . 

Titanium.. TiF^j TiF 4 T1C1_>; Ti^; TiCl 4 TiBr 4 TiI 4 . 

Zirconium . ZrP 4 ZrOl 4 ZrBr 4 * ? 

Cerium.... CeF 3 j*CeF 4 *CeCl 3 CeBr 8 * CeI 8 * 

Thorium . . ThP 4 ThC^ ThBr 4 * ThI 4 *. 

Preparation. 1. By direct union. Carbon does not com- 
bine directly with halogens, except with fluorine. The other 
elements are converted into those compounds which contain the 
largest amount of halogen. 

2. By the action of the halogen on a red-hot mixture of the 
OXide With charcoal. By this means, TiCl 4 , TiBr 4 , Zr01 4 , and 
ThCli have been prepared. The preparation of chloride of 
titanium may serve as a type of the rest : 

TiO, + 20 + 2CZ, = TiCl* + 2(70.f 

* These have been obtained only in combination with water. 
f Chew. Soc., 47, 119; Comptea rend., 104, 111; 106, 1074. Carbon tetra- 
cbloride mayj>e substituted for free carbon and free chlorine. 


CeCl 3 has also been prepared by passing a mixture of carbon 
monoxide, CO, and chlorine over the ignited oxide ; and Ti01 4 , by 
the action of CC1 4 on ignited TiO^. 

3! By the action of the halogen on the hydride or 
sulphide of the element. This is the method by which 
carbon tetrachloride, CCli, is commercially prepared. The 
disulphide (see p. 282), mixed with chlorine, is passed through 
a tube fi lied with pumice-stone and heated to redness. The 
chlorine combines with both carbon and sulphur, thus : 

os. + 30^ = cch + &OZ,. 

The chloride of sulphur is afterwards decomposed by the action 
of lime- water* (see p. 167), and the carbon tetrachloride purified 
by distillation 

Methane or marsh gas ^hydrogen carbide), OH* (see p. 560), 
is also converted by the prolonged action of chlorine into the tetra- 
chloride, thus : 

+ 4CZ2 = CCh + 4HCL 

There are, however, three intermediate stages,, GH 3 Cl, CH.Cl^, 
and CHC1 3 . 

Similarly, O 2 H 6 can be converted into C 2 C1 6 , through the 
following stages : 

CtH 6 Cl; C 2 H 4 C1 2 ; C 3 H 3 C] 3 ; C 2 H 2 C1 4 ; C 2 HC1 8 , and 0,01*. . 

4. By the action of the hydrogen halide on the element. 
By this method TiCl 3 , ZrF 4 , CeF 3 , CeCl 3 , CeBr a , CeI 8 , aud ThCl 4 , 
have been produced in solution. Hydrogen is evolved. 

5.* By the action of heat on CC1 4 other chlorides are pro* 
duced, thus : 20Ck = 2 OZ 8 + Cl* ; 2CCZ 4 = 2 0/ 4 + 20Z, ; 600? 4 
= 6 OZ fl -f 12C7 a > Special names are given to these bodies, viz.', 
CC1 4 , tetrachlorcimethane ; C 2 C1 6 , hex achlore thane ; C 2 C1 4 , tetra- 
chlorethylene ; C 6 C1 6 , hexachlorobenzene. 

6. By the action of hydrogen at a red heat on titanium 
tetrachloride or tetrafluoride they yield the trifluoride or tri- 
chloride. The dichloride is produced by the farther action of 
hydrogen on the trichloride. 

7. Double decomposition. (a.) The action of the 
hydrogen halide on the oxide or hydroxide of the element. 
All the fluorides, except that of carbon, have been thus prepared 
in solution; also solutions of ZrCl 4 , ZrBr 4 , CeCl 3 , CeBr 3 , CeI 8 , 
ThCl 4 , ThBr 4 > and ThI 4 . These substances, in solution, react with 
water on evaporation. Cerium chloride has been dried in the same 
manner as magnesium chloride, viz., by preparing the dyuble salt 


with ammonium chloride, and, after drying it, igniting it to remove 
ammonium chloride; also by passing a mixture of chlorine and 
carbon monoxide over the sesquioxide at a red heat. It is probable 
that the others could be obtained anhydrous in a similar manner. 

(&.) This process is applied to the preparation of carbon 
bromide and iodide from the tetrachloride. A mixture of alu- 
minium bromide or iodide and carbon tetrachloride, all diluted 
with carbon disulphide, yields carbon tetrabroraide or iodide on 
heating ; carbon tetrafluoride, CF 4 , is produced by heating silver 
fluoride, AgF, in a sealed tube with carbon tetrachloride. Cerous 
fluoride, CeP 3 , which is an insoluble white substance, is also pre- 
pared by this general method by the interaction between solutions 
of sodium fluoride and cerium chloride, thus ; CeCl 3 . Aq -f 3NaF. Aq 
= 2CeF 3 .H 2 O + SNaCl.Aq. 

Properties. The tetrahalides are all volatile at compara- 
tively low temperatures. Carbon tetrafluoride is a gas ; carbon 
tetrachloride, bromide, and iodide, titanium tetrachloride, and 
tetrachlorethylene are colourless liquids; hexachlorethane, zir- 
conium chloride, cerium trichloride, and thorium chloride are 
colourless solids,' which can be sublimed. Titanium dichloride is 
a black powder,* which rapidly decomposes water, with evolution 
of hydrogen, combining with the oxygen to form an oxychloride. 
Titanium trifluoride and trichloride consist of violet scales, 
soluble in water with a violet colour. Titanium tetrabromide 
is a red liquid; and the tetriodide forms brown needle-shaped 
crystals, deric fluoride is not known in the anhydrous state. 
Combined with water as CeP 4 .H 2 O, it is a brown irwOnKJa, 
powder, produced by treating the hydrated dioxide with< aqueous 
hydrofluoric acid. It is doubtful whether the si^stance described 
as thorium fluoride is not in reality an oxyfluoriu] e ThOF 2 . 

Carbon tetriodide decomposes when heated, or \vhen exposed to 
air. With the exception of the carbon compounds.* cerium tetra- 
fluoride, and possibly thorium fluoride, these substances are deli- 
quescent, and soluble in water, probably reacting wit\h it to form 
oryhalides ; this change, certainly takes place on evaporation, in 
some cases an oxyhalide, in others the oxide, being] produced. 
Carbon tetrabromide occurs as an impurity in commercial bromine. 

Friedel and Ghierin, Annale* (5), 7, 24. 


Physical Properties of Bodies of the Formula 

Melting-points. Boiling-points. 

Mass of 1 c.c. solid 
or liquid. 

F. CL. Br. I. F. Ci. Br. I. F. Cl. Br. I. 

Carbon.., 1 '632 3 '42 4 '34 ? ? 91 100* ? 76' 7 J 189 '5 

atO at 14 at 20 

Titanium. ? 1 '761 2 '6 ? ? ? 39 150 ? 136'4 230 360 C * 


Zirconium. ? ? ? ? ? ? ? white ? P ? 


Cerium ..? ??????? ? 

Thorium . ? ? ? ? ? ? ? ? ? ? ? ? 

Of the other halides : 

Mass of 1 c.c. 

CoCl 6 1-62 

C 2 "Br 6 ? 

C 2 C1 4 l-65atO 

C 2 Br 4 ? 

TiF 3 ? 

Ti 2 Cl 6 ? 


187 rf 








CeCl 3 . 

not at bright redness 

Heats of formation. The following only have been determined : ' 

C + 2CL = CCl< + 210K. 
2C + 2Cl] = C 2 C7 4 - 12K. 

The vapour-densities of many of these compounds have been 
determined, and it may be safely concluded that, in the gaseous 
state, most of them possess the molecular formulae given above. 

Double halides. These are for the most part produced by mixing solutions 
of the two halides and crystallisation. Those of carbon are produced by sub- 
stitution of chlorine for bromine, or by addition of bromine to a chloride (e.g., 
C2C1 4 + Br 2 = C 2 Cl 4 Br 2 ), or of chlorine to a bromide. 

Carbon compounds. CCl 3 Br; a liquid boiling at 104'3. CCl^Br 3 boils at 
a higher temperature. C 2 Cl 4 Br^ exists in two forms, isomeric with each other, 
one produced by direct addition of bromine to C 2 C1 4 ; the other by the action of 
bromine on C 2 HC1 6 . There are also known : C 2 Br 4 CL ; C 2 Br 3 Cl ; and C 2 Br 2 Cl 2 . 
These bodies have vapour-densittes corresponding with the formulae given. 

The other halides combine in varying amount with halides of other elements. 
As instancesj the following compounds may be given : 

8 : 1. 2ThCl 4 .KC1.18H 2 O. 

6 : 1. 3Ti01 4 2PH 4 C1. 

4 : 1. ZrF 4 .KF; ThP 4 .KP. 

4 : 2.TiP 4 .2HP ; TiP 4 .2KF ; 

TiF 4 .NiF 2 ; ZrF,'.2KF; 

TliF 4 .2XF. 
8 : 3. 2CeF 4 .3K7. 

TiF 4 .2NH 4 F; TiF 4 .CaF 2 ; TiF 4 .OaF 2 ; 
ZrF 4 .MnF 3 ; ZrOl 4 .2NaOl; 

Melts with decomposition. 


4 : 3. Ti01 4 .3NH 4 01; ZrP 4 3KP; 2ZrF 4 .3CuF>. 

4 : 4. ZrF 4 .2ZnP 2 ; ZrP 4 20dP 2 ; ZrP 4 .2MnP 2 ; ZrP 4 .2NiF 2 ; 

2ZrP 4 .2KP.NiP 2 . 

4 . 6, TiF 4 2FeF j; TiCl 4 .6NH 4 Cl, 
4 : 8. ThCl 4 8NH 4 C1. 
3 : 3. TiF 3 .3NH 4 F. 

These halides arc able to combine with others in many proportions. The 
products are crystalline substances often combined with water, sometimes anhy- 
drous. As regards their molecular weights, nothing is known ; hence the 
simplest possible formulae have been assigned to them. 

Halides of Silicon, Germanium, Tin," Terbium, 
and Lead. 

It has been already remarked as doubtful whether terbium 
belongs to this group of elements. Tbese bodies, like those of the 
last group, show a decrease of volatility witb increase of the 
atomic weight of the metallic element. 

Sources. The only native halide is lead chloride, PbCl 2 , 
which was found in the crater of Vesuvius, after the eruption of 
1822. A chloride and carbonate of lead also occurs native, though 
rarely, as corneous lead ; its formula is PbCO 3 .PbCl 2 . 

The following compounds are known : 

Fluorine. Chlorine. Bromine. Iodine 

Silicon ..... Si 2 F 6 ; SiF 4 . Si 2 Cl 4 ; Si 2 01 6 ; SiCl 4 . Si^Br 6 ; SiBr 4 . SiI 2 ; Si 2 I 6 ; SiI 4 . 

Fluorine. Chlorine. 

Silicon ..... Si 2 F 6 ; SiF 4 . Si 2 Cl 4 ; Si 2 01 6 ; SiCl 4 . . 

Germanium. ? GeF 4 . GeCV aeC! 4 . ? GeI 4 . 

Tin ........ SnF 2 ; SnF 4 .* SnCl, ; nCl 4 . SnBr 2 ; SnBr 4 . SnI 2 ; SnI 4 . 

Terbium.... TbCl 3 ?* 

Lead ....... PbF 2 . PbCl, ; PbCl 4 ?* PbBr 2 . PbI 2 . 

Preparation. 1. By direct union. These elements readily 
combine with the halogens, when they are heated together, forming 
the compounds containing the greatest amount of halogen. 

Silicon takes fire in fluorine gas, burning to silicon fluoride. 

This is the only method of preparing silicon tetriodide, SiI 4 . 

2. By the action of the halogen on a red-hot mixture of 
the oxide with charcoal (see p. 131). This. is the most con- 
venient method of preparing silicon tetrachloride and tetrabromide. 
It is necessary to take the utmost precaution to exclude moisture 
by' scrupulously drying the halogen ; for the chloride and bromide 
are instantly decomposed 'by water. The silicon chloride or 

* Not knowji in the anhydrous state. 


bromide is condensed in a (J-tube, cooled by a freezing mixture. 
The equation is : SiO 2 + 2C + 2C7 2 = SiCl* + 200. 

3. By the action of the hydrogen halide on the element. 
By this means germanium fluoride and tin dichloride, bromide, 
and iodide may be conveniently prepared. Silicon fluoride may also 
be formed thus. Hydrogen gas is in every case evolved. It is 
believed that hydrogen chloride, at a red heat, converts germanium 
into the dichloride, GeCl 2 . 

The usual method of preparing stannic chloride, which bears a 
close analogy to the action of a haloid acid on the element, is by 
distilling a mixture of granulated tin with mercuric chloride. The 
stannic chloride distils over, leaving the mercury in combination 
with the excess of tin, thus : 

2HgCl 2 + Sn = 2Hg + 8nCl*. 

4. By double decomposition. (a.) This is the usual and 
easiest method of preparing the halid.es of lead, a solution of the 
nitrate or the acetate of lead being treated with a solution of any 
soluble halide, for example, with the nitrate, Pb(N0 3 ) 2 .Aq -f 2KF.Aq 
= PbP 2 + 2KN0 3 .Aq ; and with the acetate, Pb(C 2 H 3 2 ),. Aq + 
2HCl.Aq = Pb01 2 + 2C,H 4 O 2 .Aq. 

(6.) The action of the hydrogen halide on the oxide or 
hydroxide of the element. Silicon tetrafluoride, the halides of 
tin, and terbium chloride have been thus produced. The oxides of 
lead are attacked superficially by the halogen acids ; but, the halides 
of lead being sparingly soluble, a coating of halide is formed, which 
renders the action slow. By alternately boiling lead oxide with 
the halogen acid, and with water, in order to dissolve this coating, 
complete conversion into halide may be accomplished. 

Lead dioxide, thus treated with solutions of hydrogen chloride, 
bromide,* or iodide, undergoes the following reactions, half the 
halogen being liberated : 

Pb0 2 + 4HCl.Aq = PbCl a + 2H 2 O + Cl z + Aq. 

Hydrogen fluoride is without action on lead dioxide. 

5. By the action of the element at a red heat on the 
tetrahalide the disilicon hexahalide has been prepared, thus : 

6SiCl 4 + 2Si = 4Si*Cl 6 . 

As eiamples of these methods of preparation, the following instances may be 
chosen : 

1. Tin, melted in a deflagrating spoon, and plunged into ajar of chlorine 
gas, burns to the tetrachloride. 



2. A mixture of silica and carbon, made into a paste with starch, and 
moulded into balls, and then strongly ignited, is heated in a porcelain tube 
bj means of a Fletcher's tube-furnace, provided with a blast, in a current 
of chlorine, perfectly dried by passing through tubes filled with phosphorus 

Fra. 27. 

The silicon chloride produced must bo condensed in a 
in a freezing-mixture. The preparation is not easy, and is not well adapted for 
a lecture experiment. 

3. Tin, granulated by pouring the melted metal into water, is boiled in a 
flask with strong hydrochloric acid, a few pieces of platinum-foil being added to 
form a galvanic couple and assist solution. It sldwly dissolves, forming 
tannous chloride. 

4. Silicon tetraflnoride may be prepared by heating in a glass flask a 
mixture of equal parts of fine sand and powdered fluorspar with excess of sul- 
phuric acid. The hydrogen fluoride liberated attacks the sand, forming water, 
which unites with the sulphuric acid, and hence does not exercise a decomposing 
action on the silicon fluoride. The latter escapes as a colourless gas. It may 
be made to react with water, by causing the exit-tube to dip into a little 
mercury in a beaker, the beaker being filled up with water. The mercury is 
required, else the exit-tube would be soon blocked by deposition of silicon 
hydroxide (or silicic acid), resulting from the decomposition of the fluoride (see 
p. 163). 

The action of lead dioxide on the halides of hydrogen may be easily shown 
by warming in a test-tube a few grams with some hydriodic acid. Violet fumes 
of iodine escape, and the dioxide is converted into yellow iodide. 

5. The formation of the halides of lead may be shown, as in 4a. 

Properties. Tetrahalides. These compounds boil at com- 
paratively low temperatures. Silicon tetrafluoride is a colourless 
gas at ordinary temperatures, the chloride and bromide are volatile 
liquids ; and the iodide a white solid. Germanium chloride* is a 
colourless volatile liquid ; and tin tetrachloride is also mobile and 
colourless, boiling at a somewhat higher temperature. Germanium 

J.prakt. Chem. (2), 84, 177. 


bromide and fluoride do not appear to have been prepared ; the 
iodide is a yellow solid, giving a yellow vapour. It dissociates 
somewhat below 658. Tin tetrafluoride has not been obtained in 
the anhydrous condition ; the bromide forms volatile white crystals, 
and the iodide is yellowish-red, and also volatile. All these sub- 
stances react with water, forming oxides, or oxyhalides; hence, 
being volatile, they all fume in the air. The vapour-densities of 
most of them have been determined, and correspond to the simple 
formulae MX4. 

Bodies of the formula M 3 X 6 . These are only known to exist 
as compounds of silicon. The fluoride, Si 2 F fl (?), is a white powder 
(probably an oxyfluoride). The iodide, Si 2 I 6 , produced by the 
action of finely-divided silver on the tetriodide, is separated from 
the excess of silver by solution in carbon disulphide, from which 
it deposits in colourless prisms. By warming it with mercuric 
chloride it is converted into the corresponding chloride, Si 2 Cl 6 , 
which is a colourless mobile liquid. The corresponding bromide 
is produced by shaking a solution of the iodide with bromine dis- 
solved in carbon disulphide, and removing the iodine by agitation 
with mercury. It forms white crystals. A determination of the 
vapour- density of the chloride, Si^Cl^ showed it to possess the 
molecular weight corresponding to that formula.* 

Dihalides. Silicon dichloride is a liquid, which has not yet 
been obtained pure ; the di-iodide remains as an orange-coloured 
residue on distillation of the compound Si 2 I 6 , which splits into the 
tetriodide and di-iodide, thus :Si 2 I 6 = SiI 4 + SiI 2 . It is in- 
soluble in all known solvents, and is decomposed by water. 

Germanium dichloride is a colourless liquid. Its formula is 
as yet uncertain, and it may possibly be GeHCl 3 , for it has not 
been analysed. 

Tin (^fluoride has not been obtained anhydrous. It crystallises 
from water in small opaque prisms. The dichloride crystallised 
from water is known as " tin-salt.*' On evaporation of its solution, 
a portion reacts with water, forming oxychloride and hydrogen 
chloride. The excess of water evaporates along with the hydrogen 
chloride. On raising the temperature the undecomposed stannous 
chloride distils over, leaving the oxychloride. It forms a white 
lustrous crystalline mass. With a large quantity of water it gives 
a precipitate of oxychloride, SnCl 2 .SnO.2H 2 O. Its solution is a 
powerful reducing agent, ior it tends to take chlorine from 
hydrogen chloride or oxygen from water, liberating hydrogen, 

* Annale* (4), 9, 5; 19, 334; 28, 430; 27, 416; (5), 19, 390. 


when there is any substance present with which the hydrogen can 
combine. The clibromideis similar to the dichloride. The di-iodide 
is a dark-red mass ; its iodine is replaced by oxygen when, it is 
heated in air. 

Lead di fluoride, dichloride, and dibromide are white solids, 
sparingly soluble in boiling water and crystallising therefrom in 
long needles. The iodida is yellow and crystallises in golden- 
yellow spangles. 

From the vapour-density of stannous chloride it would appear 
that these bodies in the state of gas have, at temperatures not far 
removed above their boiling-points, the double formula, e.g., 
Sn 2 Cl4* ; but that, as the temperature rises, the cornplex molecule 
dissociates into two simpler ones, viz., SnCl 2 (see N 2 O4, p. 833). 
Lead chloride appears to dissociate before its volatilises, for its 
density corresponds to the simple formula PbCl 2 .t 

Silicon . ... 






,ss of 1 c.c, liquid. 



1 -524 2 
at 18 
2 -379 

Br. I. F. Cl. Br. 

823 ? -102^? -12 

? ? ? ? ? 

? 4-orr> ? ? 30 

at 11 


F. Cl. 

? 57-6 

? 86 
? 114 







Hexahalides : Si 2 Cl 6 , sp. gr. T58 at ; m.-p. -1; b.-p. 146148. 
Si 2 Br 6 , b.-p. about 240. Si 2 I 6 , m.-p. about 250, with decomposition. 
Dihalides : Sp. gr. : SnCl a ?. SnBr 2 , 5'117 at 17. SnI 2 P. 
M.-p.: 2493. 215-5 316. 

B.-p.: 601. 620 ? 

Sp. gr. : PbF 2 ,~8'24 at 2 ; PbCl 2 , 5*80 at 15 } PbBr 2 , 6'60 at 7 '5 ; PbI 2 , 6'06 


M.-p. : PbUJ 2 , 498; PbBr 2 , 499; PbI 2 , 383. 
B.-p. : PbCl 2 , 900? ; PbBr 2 , above 861 j PbI 2 > 861954. 

* Zeitschr. phys. Chem., 2, 184. The author differs entirely from the con- 
cluding words of this memoir regarding the non-existence of Sn 2 Cl 4 in the state 
of gas. 

\ Brit. Assn., 1887, 668. 

J Volatilises without melting. This behaviour is explained as follows : The 
boiling-point of a liquid is dependent on the pressure. By lowering the pressure, 
the boiling-point is lowered, whereas the melting-point is almost unaffected by 
small alteration of pressure. It is evident that by a sufficient reduction of 
pressure the boiling-point may be lowered till it occurs at a temperature below 
the melting-point. Such bodies as silicon fluoride, hexachlore thane, C 2 C1 6 , and 
many others are in this condition under ordinary atmospheric pressure. By in- 
creasing the pressure, BO as to raise their boiling-points, they can be melted. 


Heats of formation : 

Sn + C1 2 - SnCL, 4- 808K ; + Aq = 811K. 

Sn +2CL 2 = Sn01 4 + 1273K; + Aq - 299K. 

The last number implies decomposition when solution takes place 

Pb + C1 3 - PbCl 2 + 828K; + Aq - -68K. 
Pb 4- Br 2 = PbBr^-f 645K ; + Aq = -100K(?). 
Pb + I 2 = PbI 2 " + 398K; + Aq = -160K(?). 

Double halides. Silicon, like carbon, forms double halides, 
of which the molecular weights have been determined in many 
cases. For example, by the action of bromine on the compound 
SiHCl 3 , named silicon chloroform (see p. 501), three chloro- 
bromides have been obtained : one has the formula SiCljBr, the 
second, SiCl 2 Br 2 , and the third, SiGlBiv* They are all liquids : 
the first boiling at 80, the second at about 100, and the third at 
140 141. There appear to be similar chlorobromides of tin, 
which, however, are not stable in the gaseous state. 

The tetrahalides form numerous double salts. Those of silicon tetrafluorule 
have been most carefully studied ; they are named silicifluorides. G-ermam* 
fluorides and stannifluondes have also been prepared. The following is a list 
the more important ones : 

SiF 4 .2HF.Aq; SiF 4 .2KF; SiF 4 .BaF 2 ; &c. 

GeF 4 .2KF. 

SnF 4 .2KF; SnF 4 .BaF 2 . 
SnCl 4 .2HCl; SnCl 4 .2KCl, SnCV^CaCLjj SnCl 4 .BaCl 2 . SnOl 4 .2NH 4 Cl. 

SnBr 4 .2HBrj SnBr 4 .2NaBr ; SnBr 4 .MgBr 2 , and others. 
PbCl 4 .2HCl.Aq(?) ; PbCl 4 .9NaCl; PbCl 4 .lGCaCL 2 . 

The compound SnCl 4 2NH 4 C1 is known as " pink salt," being used as a 
means of fixing pink dyes. 

These compounds are mostly prepared by direct addition ; but those of 
silicon may also be produced by the action of 8iF 4 2HF.Aq on the oxides, 
hydroxides^ or carbonates of the metals. When silicon fluoride is passed into 
water the following reaction takes place (see Boro fluorides, p. 132) 
3SiF 4 -4- 3H 2 O + Aq = H 2 SiO 3 + 2H^SiF 6 .Aq. The gelatinous precipitate 
formed when silicon tetrafluoride ia passed through water consists of silicic 
acid, H 2 3iO 3 ; the aqueous solution contains the body H 2 SiF 6 , hydrosilicifluoric 
acid ; its formula is deduced from that of its salts, as it decomposes on evapora- 
tion into hydrofluoric acid and silicon fluoride, a portion of which reacts with 
the water to form more silicic acid. 

The more important compounds of this acid are potassium silicifluoride, 
K 2 SiF 6 , which is one of the few sparingly soluble salts of potassium ; it is used 
as a source of silicon (see p. 50) ; and the barium salt, which is insoluble in 
water, the corresponding salts of strontium and calcium being soluble. This is 
utilised as a method of separating barium from these metals. 

Chem. Soc. t 61, 590. 


Ceesium stannichloride, SnCl 4 2CsCl, being nearly insoluble in water, may 
be separated as such from the corresponding compounds of sodium, potassium, 
and rubidium. 

All these double salts crystallise in the same form, and are therefore termed 

Many double halides are known of the dihalides of tin and lead. None have 
been gasified ; hence their molecular weights are unknown ; the simpler formulae 
are therefore given as a rule. 

Compounds containing- two different halogens : 

SnOlI; PbPCl; PbClBr; PbBrl. 

3PbBr 3 .PbI 3 ; 6PbBr;.PbI 3 . 

Compounds with the halides of other elements : 

2:1. SnCl 2 .HCl; SnCl c KCl ; SnBr 2 .KBr ; SnBr 2 .NH 4 Br; SnI. 2 .KI-; 

SnI,.NH 4 I; PbI 2 .XI. 
2 : 2. SnCl 2 .2XCl; SnCL 2 .BaCl 2 ; SnBr 2 .2NH 4 Br; SnI 2 .2KI; 

PbI 3 .2HI; PbI 2 .2KI; PbI 2 .2NH 4 Cl. 

Many more complex ratios have also been noticed among the lead 
halides, e.g. : 2 : 3, PbI 2 .3NH 4 Cl ; 2 : 4, PbI 2 .4KI ; 2 : 6, PbBr 2 .6NH 4 Br ; 
2:7, PbBr 2 .7NH 4 Br; 2:9, PbCl 2 .9NH 4 Cl ; 2:10, PbCl 2 .10NH 4 Cl, and 
others still more complex. These last bodies possess the qualifications usually 
attributed to definite chemical compounds, viz., definite crystalline form, 
coupled with constant composition. 

The formulae of these halides of the carbon and silicon groups 
have been determined : 

1. Prom the vapour-densities of many of the compounds, 
and from the analogy of those of which the vapour-densities have 
not been determined with those in which that constant is known. 

2. By the method of replacement. It is argued, for instance, 
that the formula of the compound CCl 3 Br implies the existence of 
four atoms of chlorine in the compound CCh, inasmuch as one- 
fourth of the total amount of chlorine it contains has been replaced 
by bromine. In this case, and in that of the similar silicon com- 
pounds, SiCl 8 Br, SiCl 2 Br 2 , and SiClBr 3 , this view is confirmed by 
the vapour-densities of the bodies. But there is no means of 
ascertaining whether such a body as SnClI possesses that formula 
or the formula SnCl 3 .SnL, for it has never been gasified. Indeed, 
judging from the vapour-density of Sn^CU, the latter formula 
would appear the more probable ; and no simpler formula than 
2PbCl 2 PbI 3 is possible in the case of the tetrachlorodiiodide of 

3. The atomic heats of carbon and silicon present special 
anomalies. It has been shown by Weber,* however, that, like 

Poaa. Ann.. 164. 367. 


those of beryllium and boron, they approach constancy at high 
temperatures, and become approximately normal. They are as 
follows : 






Sp. ht. 


Sp. ht. 


Sp. ht. 




+ 22. 




+ 10. 












The atomic weights of carbon and silicon have been deduced 
from their atomic heats at 1000 and 232 respectively, which are, 
for carbon, 5'608, and for silicon, 5'671. 

A few words must be added as to the views which are held re- 
garding the nature of the atomic combination in the compounds 
C 2 C1 6 , Si 2 Cl 6 , C 2 C1 4 , Sn 2 Cl 4 , and analogous bodies. These views are 
based on the behaviour of the compound of carbon and hydrogen 
named ethane, C 2 ff^ which is analogous to C 2 Cle, and which, 
indeed, can be converted into the latter by the continuous action 
of chlorine, whereby all the hydrogen atoms are successively re- 
placed by an equal number of atoms of the halogen. The compound 
CH 3 I, named iodomethane, when acted on by sodium, loses its 
iodine, sodium iodide being produced. But the vapour-density of 
the resulting gas shows it to possess not the formula CH 3 , but the 
double formula CH* CH 3y or (7 2 JBT 6 . This is also borne out by the 
fact that the hydrogen in ethane, C 2 H 6 , may be replaced by 
chlorine in sixths, giving C Z H(>CI, monochlorethane, C 2 H 4 C1 2 , 
dichlorethane, &c. 

It is argued that the group CH 3 may be regarded as replacing 
an atom of chlorine in CH 3 Cl, or of iodine in CH 8 I, and that the 
compounds CH 3 C1 and CH 3 CH 3 are in that sense analogous. 
Hence the formula of C^Cla may be written CC1 3 CC1 3 ; and of 
Si 2 Cl<j, SiCl 3 SiCl 3 . 'And by similar reasoning it is argued that 
the compound C 2 C1 4 may be regarded as composed of two separate 
portions, viz., CCl 2 ZlCCl 2 , the two horizontal lines expressing the 
hypothesis that the group CC1 2 replaces two atoms of chlorine in 
the compound CC1 4 . And the vapour-density of these compounds 
C 2 C1 6 and C 2 C1 4 , and of their hydrogen analogues, CJIs and C 2 H 4 , 
even at the highest temperatures to which they can be submitted 
without decomposition, shows that they still possess the formulae 
given. On the other hand, there can be no doubt that stannous 


chloride, Sw 2 C7 4 , at a sufficiently high temperature, has a vapour- 
density corresponding to the simpler formula SnG'k. The conclu- 
sion appears, therefore, to follow, that, if it were possible to subject 
(7-2(7/4 to a sufficiently high temperature without inducing decom- 
position, it, too, would possess the formula CCk, Si a O/ e , when 
heated, splits into SiClt and SiOl<i\ it is, therefore, extremely 
improbable that any member of this group, at any temperature, 
will be found to have the formula MX 3 , for more stable forms of 
union exist. But, in the chromium group, chlorides of both 
the general formula? MC1 2 and MCI.} are known ; and these appear 
capable of existence in the two molecular states, MCI* arid JfC/a, and 
M^Ck and i/ 2 OZ4, respectively ; it will be rcmetubei^d that, in the 
chromium group, chlorides of the general formula MC1 4 are ex- 
ceedingly unstable, the only representative definitely known being 
MnF 4 , and that only in aqueous solution. Hence the stability of 
compounds with the simpler molecular form MI'V 




Halides of Nitrogen, Vanadium, Niobium, Tan- 
talum (Neodymium, see p. 605). 

Again it is to be noticed that the compounds of nitrogen, the 
first element of this group, differ considerably from those of the 
other members. While the halogen compounds of nitrogen are 
exceedingly explosive, those of the other elements are stable, 
though decomposed by water. For these reasons none of them are 
found in nature. The following table shows tho compounds 
known : 

Fluorine. Chlorine. Bromine. Iodine. 

Nitrogen... NC1 3 . Nflr 3 ? NI 3 . 

Vanadium.. YF 4 ?* VCL^; VC1 3 ; VC1 4 . VBr 3 . VI 4 ?* 

Niobium NbF 5 * NbCl 3 ; NbCl 6 . NbBr 6 . 

Tantalum .. TaF 6 .* TaCl 6 . TaBr 6 . 

The iodides of niobium and tantalum, though probably capable of existence, 
have not been prepared. 

Preparation. 1. By direct union. Nitrogen will not com- 
bine directly with the halogens. Vanadium tetrachloride and tri- 
bromide are prepared by passing the vapour of the halogen over 
the heated element; and tantalum pentachloride has also been thus 

2. By the action of the halogen on a red-hot mixture of 
the oxide with charcoal. This is the method of preparation of 
niobium and of tantalum pentachloride and pentabromide. Vana- 
dium oxytrichloride, VOC1 3 (see p. 332), when passed over red-hot 
charcoal along with chlorine, also yields the trichloride. 

* Known only in solution. 


3. By heating a higher halide. Vanadium tetrachloride 
by distillation alone splits up into chlorine and the trichloride ; along 
with hydrogen, the dichloride is formed ; and niobium pen tfc chlo- 
ride, passed through a red-hot tube, yields the trichloride and 

4. By the action of the halogen on a compound of the 
element. Ammonia (hydrogen nitride, NH 3 ) on treatment with 
excess of chlorine, bromine, orjiodine, yields exceedingly explosive 
bodies. The method of preparation is as follows : 

A flat leaden dish, in which a smaller thick dish is placed, is filled with a 
strong solution of ammonium chloride. A small jar, of about 200 cubic centi- 
metres capacity, provided with a neck, is placed in the solution, standing on 
the smaller leaden dish. The neck is closed with a cork, through which a 
tube passes, which is connected with an apparatus for generating chlorine by 
means of a short piece of india-rubber tubing, on which a clip is placed. The 
solution of ammonium chloride is drawn up into the jar by suction, and when 
the jar is full the clip is closed. The chlorine apparatus is then connected, 
and by opening the clip the jar is quickly filled with chlorine. The chlorine is 
absorbed by the solution, while oily drops collect on the surface, and sink, col- 
lecting in the leaden dish. Air is then admitted by disconnecting the chlorine 
apparatus and opening the clip, and the jar is removed. These drops, when 
touched with an oiled feather tied to the end of a long stick, explode with the 
greatest violence, shooting a column of water into the air and flattening the 
leaden vessel. 

Recent analysis* has shown that the hydrogen of the ammonia 
is replaced by stages, exactly as in the case of the hydrogen of 
methane, CH* (see p. 145). By passing chlorine for half an hour 
into water in which these drops are suspended, the trichloride is 
finally formed. The equations are these : 

NH 4 Cl.Aq + Ck = 2HCl.Aq + NH 2 C1. 
NH 4 Cl.Aq + 2CZ 2 = SHCl.Aq + NHC1 2 . 
NHC1 2 -f Aq + Cl* = HCl.Aq + NC1 3 . 

The corresponding bromine compounds have been little investi- 
gated, but are made by treating ammonia with excess of bromine. 
Aqueous ammonia reacts with iodine dissolved in alcohol giving 
NI S ; but with a weaker solution of ammonia NHI 2 is produced. 

The action of chlorine or bromine on vanadium nitride, VN, 
at a red heat gives the trichloride or di bromide, and nitrogen. 

The oxygen in niobium oxy trichloride, NbOCl 3 , is replaced by 
chlorine when its vapour mixdd with chlorine is passed through a 
red-hot tube. 

5. By double decomposition. Action of the hydrogen 
halide on the oxide of the element. Vanadium tetroxide dis- 

* Berichte, 21, 751. 


solves in hydrofluoric acid, yielding a blue solution, which on evapora- 
tion deposits green crystals. This oxide is also soluble in the other 
haloid acids, giving similar solutions. The pentoxide, when boiled 
with hydrochloric acid, yields chlorine. Tantalum pentoxide, if 
hydrated, likewise dissolves in hydrofluoric acid, and the solution 
on evaporation is said to evolve the fluoride, leaving a residue of 
oxyfluoride. Niobium pentoxide is also soluble in hydrofluoric 

Properties. The halides of nitrogen are exceedingly explo- 
sive, and the preparation of more than a drop or two of the chloride 
and bromide is attended by great danger. They are oily yellow 
liquids, insolvble in water, which slowly decompose when left in 
contact with water or solution of ammonia. The iodide is a brown, 
ish black powder, of which it is also advisable to prepare only a 
few decigrams at a time. They explode on contact with an oiled 
feather, or indeed by the slightest impact, and often without any 
apparent cause. The pure chloride has been heated to 90 without 
decomposition, but at 95 a violent explosion occurred. 

Vanadium dichloride forms apple-green crystalline plates. The 
element may be obtained from it by heating it to redness in a 
current of very carefully dried hydrogen. Vanadium trichloride 
closely resembles chromium trichloride in appearance. When 
heated in air, its chlorine is replaced by oxygen, and the pentoxide 
is formed by further absorption of oxygen. The tribromide is a 
greyish-black amorphous mass ; it is very unstable. The tetra- 
shloride is a reddish-brown volatile liquid, soluble in water with a 
jlue colour. 

Niobium and tantalum pentafluorides form colourless solutions. 
Niobium trichloride closely resembles iodine in appearance ; it is 
unaffected by water. Niobium and tantalum pentachlorides form 
yellow volatile crystals ; the bromides are similar in appearance, 
but of a darker colour. 

Physical Properties. 

Mass of 1 c.c. Melting-point. Boiling-point. 

Dichloride. Vanadium . . 3'23 c.c. at 18 ? ? 

Trichlorides. Nitrogen . . . T65 ? Above 90. 

Vanadium . . 3'00 Decomposes. Decomposes. 

Niobium. ... ? P ? 

Tetrachloride. Vanadium . . T858 at 0. Below - 18. 154*. 

Pentachlorides. Niobium ... ? 194 240'5. 

Tantalum .. ? 211 242. 

The properties of the other halides hare not been determined. > 


Double halides. Although double halidesof the oxyfluorides of vanadium, 
niobium, and tantalum have been studied (see p. 336), the tantalifluorides are 
the only compounds of any of the halides of this group with the halides of other 
elements. They have all the general formula TaF 5 .2MF. They are produced 
by, direct union of the respective fluorides in aqueous solution, and crystallise 
Well. They are soluble in water. The following have been prepared : 

f TaF 6 .2NH 4 Fj TaF 6 .2KFj TaF 6 .2NaFj TaF 5 .CuF 2 j and TaF 6 .ZnF 2 . 

Halides of Phosphorus, Arsenic, Antimony, 
(Erbium), and Bismuth. 

Sources. None of these compounds are found in nature. 
The halogen compounds known arc given in ''the following 
table : 

Fluorine. Chlorine. Bromine. Iodine. 

Phosphorus.. PF 3 ; PF 5 . PC1 3 ; PC1 5 . PBr 3 ; PBr 6 . P 2 I 4 ; PI,. 

Arsenic AsF 3 . As01 3 . AsBr 3 . As a I 4 ; AsI 3 . 

Antimony SbF 3 ; SbF 6 . SbCl 8 ; SbCl 6 . SbBr 3 . SbI 3 . 

(Erbium).... ErF 3 .* ErCl s * ErBr 3 .* 

Bismuth BiF 3 . Bi 2 Cl 4 ; BiCl 3 . Bi 2 Br 2 ; BiBr 3 . BiI 3 . 

Preparation. 1. By direct union. All these bodies are 
best prepared thus, except the fluorine compounds ; phosphorus, 
arsenic, and antimony take fire spontaneously in fluorine and chlo- 
rine gases, and all combine with the halogens with great evolution 
of heat. With excess of halogen, the higher halide is formed 
where they exist; with excess of the other element,' the lower 

As examples of this method of formation, the following types may be 
chosen : 

(a.) A retort is half filled with dry sand, and on it are placed a few pieces of 
phosphorus. Dry chlorine is led into the retort, so as to impinge on the phos- 

FIG. 28. 
* Known only in aqueous solution. 


phorus frorh the generating flask. The phosphorus burns with a greenish flame, 
and the liquid chloride distils over, and may be condensed in the receiver. It is 
purified by distillation. 

(bj} A little powdered antimony is thrown into a jar of dry chlorine. It 
burns with scintillation, and on standing, the fumes condense to crystals of the 

(c.) A mixture is made of 1 volume of bromine and 3 volumes of carbon disul- 
phide and placed in a flask. Powdered antimony is added in small quantities 
at a time, and the flask is warmed gently with continuous shaking over a water- 
bath, taking care to have no flame near, for fear the disulphide vapour should 
inflame It is then allowed to cool, when the tribromide separates in crystals. 

2. By heating a higher halide. Phosphorus and antimony 
pentachloridea yield the trichloride when heated ; phosphorus 
pentabromide behaves similarly. Phosphorus pentafluoride, on 
the other hand, is stable, showing no decomposition, even at the 
high temperature of the electric spark. The decomposition of phos- 
phorus pentachloride may be well seen by heating it in a flask ; 
its vapour has a greenish-yellow colour, due to the presence of free 
chlorine. Phosphorus tri-iodide gives off iodine when heated. 

3. By heating a lower halide. Bismuth dichloride at 
330, and the dibromide at a temperature not much above that of 
the atmosphere decompose into bismuth and the trihalide. 

4. By the action of the halogen on a compound of the 
element. This method is not generally employed ; yet hydrogen 
phosphide, arsenide, and antimonide are at once acted on by 
chlorine or bromine, yielding hydrogen halide, and the halide of 
the element. It is certain that all other compounds, except per- 
haps the oxides, would behave similarly. 

5. By double decomposition. (a.) The action of the 
hydrogen halide on the oxide or sulphide of the element. 
The oxides of phosphorus are not attacked. But the oxides and 
sulphides.of the other elements yield the respective halides, either 
when heated in a current of the hydrogen halide, or when treated 
with a halogen acid. In presence of a great excess of water, 
the halides are decomposed; hence the acids must not be too 

Arsenious fluoride is prepared by heating together arsenions 
oxide, As 4 O fi , fluorspar, CaF 2 , .and sulphuric acid, H 2 S0 4 .* The 
hydrogen fluoride liberated by the action of the sulphuric acid on 
the calcium fluoride (see p. 108) attacks the arsenious oxide, pro- 
ducing arsenious fluoride and water, which combines with the 
excess of sulphuric acid, thus : 

As 2 Q 3 + 6HF -f 3H 2 S0 4 = 3(H 2 S0 4 .H 2 O) 4- 2AsF B . 

Compte* rend., 99, 874. 



The chloride may be similarly prepared from sulphuric acid, 
sodium chloride, and arsenious oxide. Ifc may also be obtained by 
distilling arsenic with mercuric chloride, thus : 

2As + 3HgCl a = ZAsCl* + 3Hg. 

(6.) Phosphorus trichloride or peutachloride reacts with arsenic 
trifluoride, yielding the trifluoride* or pentafluoridef of phos- 
phorus, thus : 

PCI + AsF 3 = P-F 3 + AsCJ 3 ; 
and 3PC1 6 + 5AsF 3 = 3PF 6 + 5AsCl 3 . 

Properties. The pentafluorides of phosphorus and arsenic are 
gases at the ordinary temperature ; the trichloride and tribromide 
of phosphorus, the trifluoride and trichloride of arsenic, and the 
pentachloride of antimony are colourless liquids, fuming in air, 
owing to their reacting with the water-vapour which it contains ; 
the remaining trifluorides, chlorides, and bromides are colourless 
crystalline solids. Phosphorus di-iodide forms orange-coloured, 
and tri-iodide, red, crystals. Arsenic di-iodide, produced by 
melting arsenic with iodine in theoretical proportion, forms a dark 
cherry-red mass, which crystallises from carbon disnlphide in 
prisms, and which decomposes into arsenic and the tri-iodide on 
addition of water. The tri-iodide forms red tablets. Antimony 
tri-iodide exists in three forms : when crystallised from carbon di- 
sulphide, it forms red hexagonal crystals ; when sublimed below 
114, yellow trimetric crystals; and from its solution in carbon 
disulphide exposed to sunlight, in monoclinic crystals. The last 
variety is converted into the hexagonal modification at 125. 
Bismuth tri-iodide forms a greyish mass with metallic lustre. 

Phosphorus pentachloride and pentabromide are yellowish 
crystalline solids ; antimony pentachioride is a colourless fuming 
liquid. These three substances dissociate when heated into the 
trihalides and two atoms of the halide; hence, their vapour- 
densities do not correspond to their formulee. In excess of tri- 
chloride, however, the decomposition of phosphorus pentachloride is 
prevented, and it volatilises as PC1 6 .$ Bismuth dichloride, on the 
other hand, decomposes into bismuth and the trichloride when 
heated, thus : 

3Bi 2 Cl 4 = 4BiCl 3 + 2Bi. 

These substances are all, with the exception of bismuth tri- 
fluoride, deliquescent, attracting water, and reacting with it to 

* Comptes rend., 99, 655; 100, 272. 

f Proc. Roy. Soc. t 25, 122; Comptes rend., 101, 1496. 

j Comptes rend., 70, 601. . 


form oxyhalides or hydroxides (acids). They are all soluble in 
carbon disulphide, benzene, &c. The erbium halides form colour- 
less solutions. 

Physical Properties. 
Mass of Ice. solid 


or liquid. 




Cl. Br. I. 

F. Cl. 



Phosphorus . ? 

1-613 2-923 ? 

? ? 



at at 

Arsenic 2 -66 

2 205 3-66 4-39 

? -18 




at at 15 at 13 

Antimony . . ? 

3-064 4-148 4 8,5* 

292 73 '2 



at 26 at 23 at 24 

Bismuth ... 5 '32 

4-56 5-4 5-64 

? 227 



at 20 

at 11 at 20 at 20 


F. Cl. 



Phosphorus . 

76 -0 

172 -9 


60-4 130 '2 



Antimony . . . 

? 223-5 

275 -4 


Bismuth .... 

? 427-439 



SbCl 5 : mass of 1 c.c. 2 '346 at 20; m.-p. -6. PC1 5 : m.-p. 148, under 
increased pressure ; volatilises at 148. 

Bi. 2 Cl 4 : m.-p. 176. Bi 2 Br 4 : m.-p. 202 (uncorr.). 

Heat of formation : 

P + 3CI = PC1 3 + 755K. 
P + 5 CZ = PC1 5 + 1050K. 

P + 21 = PI 2 
As + 3CT = AsCl 3 
As + 31 = AsI 3 
Sb + 3C7 = SbCl 3 


P + 3Br = PBr 3 + 448K. 

P + 5Br = PBr 5 + 591K. 

P -H 31 = PI 3 + 109K. 

As + 3Br = AsBr 3 + 449K. 

Sb -f 5CI = SbCL + 1049K. 

Bi + ZCl = BiCl 3 -r 906K. 

The vapour-densities of phosphorus tri- and penta- fluorides, tri- 
and penta-chlorides, tribromide, and tri-iodide have been deter- 
mined ; also those of the trihalides of arsenic and of antimony 
and bismuth trichlorides ; their molecular weights are represented 
by the formula given. Diphosphorus tetriodide has a vapour- 
density corresponding to the formula P-^Ii; moreover, the 
analogous compounds of bismuth easily decompose into the tri- 
halide and metal ; hence, the more complex formulas have been 
chosen, although the choice is not justified by any absolute proof 
in the case of bismuth. 

* Hexagonal ; monoclinic : 4 '77 at 22. 
f Decomposed. 


Double halides. 1. Compounds containing- two halogens. These are 
known only in the case of phosphorus. They are, for the most part, made by 
adding the halogen to the halide dissolved in carbon disulphide, and crystal- 
lising from that solvent. Their molecular weights are unknown. 
They are/ as follows : * 

PF 3 Br 2 , PC1 5 IC1. PCl 2 Br 7 . PCl 3 Br 8 . 

P01 3 Br 3 . PCl,Br 4 . 

PCl 4 Br. PCl 2 Br 5 , 

The following compounds with the halides of other elements are 
known : 

f PC1 5 FeCl 3 . PCl 6 .SnCl 4 . PCl 5 .SbCl 6 . PCl 5 .3HgrCl 2 . 

Phosphorus I PC1 5 .A1C1 3 . PCl 6 .SeCl 4 . PC1 6 .TTC1 5 . 
pentahalides ] PCl 6 .CrCl 3 . PC1 5 .WC1 4 . 
lPCl 5 .AsCl 3 . PCl 5 .MoCl 4 . 

Arsenic pentahalides .... AsF 5 .XF. AsF 6 .2XF. 

. . rSbF 6 .:NaF. SbF 6 .2KF. SbCl 5 .SCl 4 . 

p'Sides { * *^"=W- SbCl,SeCl 4 . 

Trihalides. Phosphorus. PC1 3 AuCL 

Antimony. SbF,, KF. SbF 3 .2KF. SbF 3 .3NaF. 2SbI 3 .8KI. 

2SbCl 3 .HC1.2H 2 0. SbI 3 ,KI. SbCl,.2NH 4 Cl. SbCl 3 .3KCl. 
SbI 3 .NH 4 I. SbCl 3 BaCla, SbCl 3 .3KBr.f 
SbI 3 .BaI 2 . SbBr 3 .3KCl f 

Bismuth.... 2BiCl 3 .NH 4 Cl. BiI 3 .HI. BiCl 3 .2NaCl. Bi01 3 .3NH 4 Cl. 

BiCl 3 .4NH 4 Cl. 

2BiI 3 .3NaI. BiF^HF. BiI 3 .BaI 2 . BiF 3 .3HF. 
2BiCl 3 .HC1.3H 2 O. BiI 3 XI. 

These are some of the compounds known. It will be noticed that the ratios 
of the number of atoms of halogen in the two components vary between 6 : 1 
and 3 : 4. All these substances react with water, producing oiyhalides 
(seep. 385). 

Halides of Molybdenum, Tungsten, and Uranium. 

These bodies present some analogy with the halides of chro- 
mium, which, indeed, in the periodic table, falls in this group. 
Sources. None of these halides occurs native. 
The following is a list of the known compounds : 

Fluorine. Chlorine. 

Molybdenum MoF 3 ;t MoF 4 ;J MoF 6 . MoOl 2 ; MoCl 3 ; MoCl 4 ; MoCl 5 
Tungsten ..WF 2 t WF 6 . WCLj; WC1 4 W01 6 ; WC1 6 
Uranium,... UF 4 UC1 3 ; UC1 4 ; U01 5 

* Chem. Soo., 49, 815. 

t These bodies are identical, although prepared by direct addition. 

J Known only in solution. 


Bromine. Iodine. 

Molybdenum . . . MoBr 2 j MoBr ;{ ; MoBr 4 MoI 2 ;* MoI 4 .* 

Tungsten *WBr 2 WBr 4 ; WBr 5 . WI 2 . 

Uranium UBr 4 UI 2 . 

Preparation. 1. By direct union. It is important to avoid 
the presence of air and water-vapour, else oxyhalides are 
obtained. This process yields molybdenum and uranium penta- 
chlorides, tungsten hexachloride, molybdenum tetrabromide, and 
tungsten pentabromide, all of which are volatile. 

2. By the action of the halogen on a mixture of the 
oxide and charcoal. By this means, molybdenum tribromide 
and uranium pentaehloride have been prepared ; it is doubtless 
adapted for the production of auy of the higher halides. 

3. By heating the higher halides. Molybdenum tri- and 
tetra-bromides, when distilled, undergo decomposition into bromine 
and the dibromide, MoBr 2 . Tungsten hexachloride, between 360 
and 440, dissociates into pentaehloride and free chlorine. In. 
other cases, the distillation of a halide yields a mixture of two 
halides ; for example, molybdenum trichloride, sublimed in dry 
carbon dioxide, splits into the di- and tetra-chlorides, thus : 
2Ifo(7/ 3 = MoClz + MoCli. And the tetrachloride is also unstable 
when distilled, giving tri- and penta-chlorides, 2MoCh = MoCl$ -f 

4. By the action of hydrogen on the heated halide. 
Molybdenum pentaehloride yields hydrogen chloride and the tri- 
choride at 250. Tungsten hexachloride and uranium penta- 
ehloride also yield a mixture of lower chlorides when thus 

5. By double decomposition. The action of a halide on 
the oxide. (a.} The fluorides are all thus prepared from the cor- 
responding oxides by the action of aqueous hydrofluoric acid. 
Solutions of many of the other halides may also be prepared thus. 

(6.) Tungsten hexachloride is produced by heating in a sealed 
tube to 200 a mixture of tungsten trioxide and phosphoric 

Properties. 1. Dihalides. Molybdenum dihalides when pre- 
pared in the dry way are insoluble in water , but when obtained 
from the oxides they form brown or purple solutions. The di- 
chloride is a sulphur-yellow powder. Tungsten dichloride is a 
loose grey powder ; the fluoride forms a yellow solution. 

2. Trihalides. Molybdenum trichloride is a red powder like 

* Known only in solution, 


amorphous phosphorus; the tribromide forms dark needles ; both 
are insoluble in water. Uranium trichloride is dark brown. 

3. Tetrahalides. Molybdenum tetrafluoride forms a red solu- 
iion ; and uranium tetrafluoride an insoluble green powder. 
Molybdenum tetrachloride is a volatile brown substance ; that of 
tungsten a greyish- brown crystalline powder; while uranium 
tetrachloride forms magnificent dark green octohedra, and yields 
a red vapour. The tetvabromid.es form brown or black crystals. 
These compounds are deliquescent, and soluble in water. 

4. Pentahalides. Molybdenum pentachloride is a black sub- 
stance yielding a brown-red vapour ; those of tungsten and 
uranium consist of black needles. 

5. Hexahalide. Tungsten hexachloride volatilises in bluish - 
blaok needles, resembling iodine. 

Many of these compounds require further investigation. As 
has been seen, they are very numerous, and their reactions have 
by no means been exhaustively studied. 

Physical Properties. 

The mass of 1 cubic centimetre lias not been determined for any one of these 
halides. The following melting- and boiling-points and vapour-densities have 
been determined : 

MoCl 5 . M.-p , 194 ; b.-p., 268 ; vap.-dens. at 350, 13G'0 to 1 37'9. Calc. 136'5. 

WC1 5 . M.-p., 248 ; b -p , 275 6 ; vap -dens, at 360, 182*8. Calc. 180 65. 

WC1 6 . M,-p., 275 b.p., ? ; vap -dens, at 300, 190 9. Calc. 208 65. 
UC1 & dissociates when its vapour, mixed with carbon dioxide, is heated. 
Dissociation begins at 120 and is complete at 235. 

The heats of combination are undetermined. 

Double Halides. These again have been very little studied Rome com- 
pounds of molybdenum containing two halogens are known, e /; , 2MoCL MoBr,, 
2MoBr 2 MoI 2 , &c., and one compound of the formula 2MoClj MoBr 2 ,KBr ; a 
compound of the tetrachloride lias also been prepared, viz , IiMoOl 4 .2KCl. No 
similar compounds of tungsten have been prepared, and only one of uranium, 
viz., XJF 4 .K3T. These bodies much require investigation. 

The atomic weights of these elements have been determined 
from the equivalents, and by the vapour-densities given above. 

Halides of Oxygen, Sulphur, Selenium, and 

The halides of oxygen are best considered as oxides of the 
1 ->gens, q. v. (p. 459). Those of the three other elements form a 
we marked group. None of them occurs in nature. Thoy are as 
follows : 


Fluorine. Chlorine. Bromine. Iodine. 

Sulphur. ... ? S 2 C1 2 ; SC1 2 ; SC1 4 . ? S 2 I 2 ? 

Selenium... ? Se 2 Cl 2 ; SeCl 4 . Se 2 Br 2 ; S6Br 4 . ? 

Tellurium.. ? TeCl 2 ; TeCl 4 . TeBis ; TeBr 4 . TeI 2 ; Tel.,? 

Preparation. 1. By direct union. This is the general 
method of preparing these bodies. Sulphur is said to burn in 
fluorine. When chlorine is led over sulphur contained in a retort 
it grows warm, and di sulphur dichloride is formed ; by keeping 
sulphur in excess it is the only product. Diselenium dichloride is 
similarly produced, and may be obtained fairly pure by distillation 
in presence of selenium. But it dissociates to some extent daring 
distillation, v^ith formation of tetrachloride and free selenium. 
Sulphur dichloride, produced by saturating S^Ch with chlorine, is 
stable up to nearly 20, but above that temperature dissociation 
proceeds rapidly, so that at 120 it has nearly all decomposed into 
the compound S 2 C1 2 and chlorine. The tetrachloride is still 
more unstable; at 22 it can exist, but at 4-6 it has wholly 
split up into dichloride and chlorine. It will be remembered that 
it forms a double chloride with antimony trichloride, which 
crystallises and has a definite composition, 2SbCl 3 .3SCl 4 . 

Selenium tetrachloride is freed from accompanying dichloride 
by washing it with carbon disulphide, in which it is sparingly 
soluble. It may be volatilised without decomposition. 

Sulphur and bromine mix in all proportions with evolution of 
heat, but no definite compound has been isolated. It is not im- 
probable that the resulting liquid is a mixture of the compounds 
S 4 Br 2 and SBr 4 with excess of uncombined sulphur and bromine. 

Sulphur and iodine, and selenium and iodine, mix in all propor- 
tions when melted together, but no products of definite composition 
have been isolated. Tellurium di-iodide is similarly prepared ; the 
excess of iodine is volatilised away by gentle heat. 

2. By double decomposition. Sulphur and selenium 
fluorides are said to have been prepared by distilling a mixture of 
dry lead fluoride and sulphur or selenium. They have not been 
analysed. Tellurium dioxide dissolves in hydrofluoric acid, but no 
definite compound has been isolated. With hydriodic acid tellu- 
rium yields the tetriodide as a soft black powder. 

Properties. The chlorides of sulphur are yellow-brown oily 
liquids decomposed by water with separation of sulphur, thus : 

2S 2 C1 8 + 2H,O 4- Aq = H 2 S0 3 .Aq + 4HCl.Aq + 3S. 
2SC1, + 2H 2 O + Aq = H 2 SO 3 .Aq + 4HCl.Aq + S. 
SC1 4 + 2H,O + Aq = H 2 SO,.Aq + 4flCl.Aq. 


Diselenium dichloride and dibromide are dark-brown liquids, 
which., when vaporised, dissociate partially into free selenium, and 
tetrahalide. The tetrachloride and tetrabromide form yellow 
crystals. Tellurium dichloride is a black amorphous solid, melt- 
ing to a black liquid, and giving a yellow vapour. The di- 
bromide forms black needles, and the diiodide black flocks. 
The tetrachloride is a yellow crystalline mass, melting to a 
yellow liquid ; it is volatile without decomposition. The tetra- 
bromide sublimes in pale-yellow needles, which melt to a red 
liquid. The tetraiodide is a black powder. All these bodies are 
decomposed by water. 

Physical Properties. 

The following determinations have been made of the mass of 1 c.c. of 
these compounds : 

S 2 C1 2 : 1*709 grams at 0. Se 2 Cl 2 : 2'90f> grams at 17'5 
S 2 Br 2 : 2628 at 4. Se 2 Br 2 ; 3 604 at 15, 

The other known constants are as follows : 

Melting and boiling points : S 2 C1 2 : b.-p. 138 ; TeCl 2 : m.-p. 175, b.-p. 324 ; 
TeI 2 : m.-p. 160 ; TeCJ 4 : m.-p. 209 C , b.-p. 380. The vapour-densities of S 2 C] 2 , 
SeCl 4 , SeCl 3 Br, TeCl 2 , and TeCl 4 have been determined, and are normal, cor- 
responding to the formulae given. 

Heats of combination : 

2S + 2CI = 8 2 C1 2 + 143K. 

2Se + 2CI = Se 2 Cl 2 + 222 K ; Se + 4C7 SeCl 4 + 351K. 

Te + 4,01 TeCl 4 + 774K. 

It is thus seen that the more stable compounds are formed with greatest 
evolution of heat. 

Double halides. 1, SeClBr 3 , SeCl 2 Br 2 , and SeCl 3 Br, have been prepared 
by addition. They are yellowish powders.* 

2. By acting with chlorine on sulphides, the following bodies have been 
obtained : 

SC1 4 .2A1C1 3 ; 3SCl 4 .2SbCl 3 ; 2SCl 4 .SnCl 4 ; and 2SC1 4 TiCl 4 . 

3. By mixing: aqueous solutions of the constituent halides, tellurium 
halides combine thus : 

TeF 4 .Fj 2TeF 4 .BaF 2 . TeCl 4 .2KClj TeCl 4 .2AlCl 3 . ToBr 4 .2KBr; TeI 4 .2KT. 

These compounds form reddish crystals. Few attempts have been made to pre- 
pare' double halides. 

* Chem. Soc., 45, 70. 




Compounds of the Halogen Elements with each 


These compounds have no great stability. Fluorides of chlorine 
and bromine are unknown. Iodine is said by Moissan to unite 
with fluorine when exposed to it, and to be a colourless fuming 
liquid. Chlorine and bromine mix, but yield no definite com- 
pound; similarly, iodine dissolves in bromine, but separates on 
distillation. No attempts are recorded of cooling mixtures of these 
elements, but it is highly probable that evidence of combination 
would be obtained if the experiment were made. The only com- 
pounds investigated are the chlorides of iodine. They do not 
occur in nature. They are two in number, IC1, of which two 
modifications exist, and IC1 3 . 

Preparation. 1. By direct union. Iodine heated in chlo- 
rine yields the monochloride with iodine in excess ; with excess of 
chlorine, the trichloride. 

2. By displacement and subsequent combination. This 
is accomplished by heating a mixture of iodine and potassium 
chlorate, KC1O 3 . This body decomposes thus: K 2 O.C1 2 O 5 + I 2 = 
K 2 O -f 50 -h 2ICZ. Subsidiary reactions take place, thus : 
K,O.C1 2 O 6 = 2KC1 + 60 ; KC1 + 40 = KC1O 4 ; KI + 30 = 
KIO 3 , perchlorate and iodate of potassium being simultaneously 
formed. The reaction is a violent one, and the iodine monochloride 
distils over very rapidly; hence the arrangements for condensing it 
must be complete. 

3. By double decomposition. The trichloride is thus 
formed by treating iodine pentoxide with dry hydrogen chloride, 
thus:I 3 O fi + 10HCI = 5H 2 + 2CI* 4- 2JC7 3 . The higher 
chloride, 2I 3 C1 6 , presumably formed for an instant, is unstable, and 
decomposes, liberating chlorine. 


Properties. Monochloride, IC1. The liquid product, if 
cooled to 25, solidifies in long dark-red needles, melting at 
27*2. This is the a-modification. The /3-modification is sometimes 
obtained as dark-red plates, melting at 13'9, on crystallising the 
liquid between 4-5 and 10. On cooling it below 12 it 
changes into the a form.* 

The trichloride forms yellow needles, melting under pressure 
at 101. The monochloride is only slightly decomposed at 80 ; , 
boiling with partial dissociation between 102 and 106 ; whereas 
the trichloride dissociates when gasified. 

Heats of combination. 

I -f a = IC1 + 68K. 
I -f C/3 = IC1 3 + 215K. 

Both of these bodies react with water, forming iodic acid, HI0 3 , 
hydrogen chloride, and free iodine. Among the products a yellow 
body of the formula ICl.HClf is said to exist, soluble in ether. 

Halides of Ruthenium, Rhodium, and Palladium. 

Sources. These substances do not occur native. 
The following compounds are known : 

Fluorine. Chlorine. Bromine. Iodine. 

Ruthenium.. BuCl 2 ; RuCl^ (RuCl 4 ) % 

Rhodium ... BhCl 3 

Palladium... PdF 2 . PdCl 2 PdCl 4 . PdBr 4 . PdI 2 . 

Preparation. 1. By direct union. The respective metals, 
heated in chlorine, yield RuCL, RuCl <? and RhCl 3 . 

2. By the action of hydrogen chloride on the metal. 
The presence of nitric acid, HNOj, is necessary to furnisji oxygen, 
with which the hydrogen of the hydrogen chloride may combine. 
By this means, PdCl 2 is formed; with excess of nitric acid the 
product is PdCl 4 . 

3. By heating a higher halide. PdCl 4 loses chlorine, 
yielding PdCL. 

4. By removing chlorine from a higher chloride. 
A solution of RhCl 3 , on treatment with hydrogen sulphide, yields 
RhCl 2 . 

5. By double decomposition. (a.) The action of the hydro- 

* Rec. trav. chim., 7, 152. 
f Compt. Rend., 84, 389. 
J Known only in combination with KOI. 


gen halide on the hydrated oxide, in presence of water. This is 
the method of preparing PdF a (from PdO), RuCl 3 (from 
Ru 2 O^H 2 0), and RuCl 4 .2KCl (from Ru0 2 .rcH,0), in presence of 
KC1 ; also RhCl 3 (from Rh 2 3 ,wH 2 0) ; and PdBr 4 . 

(&.) By double decomposition. On adding a solution of 
an iodide to that of a soluble compound of palladium, e.</., the 
nitrate, Pd(NO a ) 2 , palladous iodide, PdI 2 , is precipitated in a 
gelatinous form. 

Properties. Ruthenium dichloride remains as a black crys- 
talline powder when chlorine is passed over ruthenium, while the 
trichloride volatilises. The trichloride, prepared in the wet 
way, is a yellgw- brown crystalline substance. On passing hydro- 
gen sulphide through its solution it is converted into ruthenious 
chloride, thus : 2RuCl 3 .Aq + H^S = 2RuCl 2 .Aq + 2HC1 + S. 
The dichloride forms a blue solution. 

Rhodium trichloride, prepared in the dry way, is a reddish- 
brown insoluble body. Prepared from the hydrated oxide, it forms 
a red solution. 

Palladium difluorido, obtained by evaporating palladous 
nitrate, Pd(N0 3 ) 2 , with hydrogen fluoride, forms colourless 
soluble crystals. The dichloride fuses to a black mass. The 
letrachloride and tetrabromide are said to form dark-brown 
solutions. It is probable that they are really compounds with 
hydrogen chloride and bromide, PdCl 4 .2HCl and PdBr 4 .2HBr. 
The di-iodide is a black gelatinous precipitate, drying to a black 
powder, and decomposing into its elements at 300 360. The 
only one of these compounds which finds practical application is 
palladium di-iodide, which is insoluble, the corresponding chlorides 
and bromide being soluble. It is therefore used as a means of sepa- 
rating iodine from the other halogens. The physical constants and 
molecular weights of these bodies are unknown. 

Double halides. Palladium fluoride is said to form double compounds 
with fluorides of potassium, sodium, and ammonium. The following compounds 
of the other halides have been prepared : 

PdCl 2 .2KCl. RuCl 3 .2KCl. RuCl 4 .2KCl. 

RhCl 3 .2KCl. PdCl 4 .2KCl. 

RhCl 3 .3KCl. PdBr 4 .2KBr. 


Halides of Osmium, Iridium, and Platinum. 

None of these compounds is found in nature. 

The following is a list of the known compounds : 

Fluorine. Chlorine. Bromine. Iodine. 

Osmium... OsCl. 2 ; OsCl 3 ; OsCl 4 . 

Indium... IrCl 2 ?i IrCl 3 ; IrCl 4 . IrBr 3 ; IrBr 4 . IrI 2 ; Irl a ; IrI 4 . 

Platinum.. PtF 4 PtCl 2 ; PtCl 4 . PtBr 2 ; PtBr 4 . PtI 2 ; PtI 4 

Preparation. 1. By direct action of the halogen on the 
metal at a red heat, osmium dichloride, trichloride, and tetra- 
chloride, iridium trichloride, and platinum tetrafluoride and tetra- 
chloride have been prepared. The double chlorides are in many 
cases produced by the action of chlorine, at a red heat, on a mixture 
of chloride of potassium, &c., with the metal. Platinum tetra- 
bromide is formed by the action of bromine and hydrobromic acid 
on spongy platinum at 180 in a sealed tube. 

2. By the action of nitro-hydrochloric acid on the metal. 
The action between nitric and hydrochloric acids generates free 
chlorine, thus : 

HN0 3 + SHCl.Aq =r 2H 2 O.Aq + NOCl + 0Z a . 

Metallic iridium and platinum dissolve in aqua regia, as the 
mixture is called, with formation of the double compounds of 
hydrogen chloride \\ith tetrachlorides. Platinum tetrachloride, 
PtCl 4 .4H 2 0, is produced by dissolving the calculated amount of 
platinic oxide in this solution. Similarly, a mixture of nitric and 
hydrobromic acids yields the tetrabromides, in solution. 

3. By heating higher halides. Iridium trichloride and tri- 
bromide have been obtained from the tetrachloride and tetrabroui- 
ide by heat. Platinic chloride (i.e., the tetrachloride) yields the 
dichloride at 440. There is little doubt that in every case the 
application of heat to a tetrahalide would be followed by the forma- 
tion of a lower halide ; but in many cases it appears to be difficult 
to avoid complete loss of halogen and reduction to metal, 

4. Double decomposition. (a.) The action of a halogen 
acid on the corresponding oxide or hydroxide of the metal. 
Osmium dichloride in solution has been thus prepared from 
osmium monoxide, OsO j similarly, iridium trichloride is pro- 
duced from Ir a O 3 . 

(b.) Iridium tetriodide is produced by mixing solutions of tho 
tetrachloride and potassium iodide. On mixing it with ammonium 


iodide, the tetraiodide is probably formed at first, but it loses 
iodine, yielding the tri-iodide. Platinum tetrafluoride is produced 
by adding silver fluoride to platinum tetrachloride, filtering from 
the precipitated silver chloride, and evaporating the solution. 
Platinum di- and tetra-iodides are formed on addition of potas- 
sium iodide to the di- and tetra-chlorides. Iridium tetrabromide 
may be similarly produced by the action of potassium bromide on 
the tetrachloride. 

5. By reduction of a higher halide. Various reducing 
agents may be used to prepare a lower from a h'gher halide. The 
one commonly used is sulphurous acid, which absorbs oxygen from 
water, libera^ng hydrogen, which combines with a portion of the 
halogen. By this means osmium di- and tri-chlorides and iridium 
di-iodide are produced. The last reaction is as follows : 

ir! 4 .Aq + H 2 O -f H 2 S0 3 .Aq = IrI 2 .Aq + H 2 SO 4 .Aq + 2HI.Aq. 

Properties. Most of these bodies are non- crystalline powders. 
Iridium trichloride, tetrachloride, tri-iodide, and tetra-iodide are 
black powders. Osmium dichloride is blue-black. It is very 
unstable, but its compound with chloride of potassium is more 
permanent. The trichloride is known only in solution. The 
tetrachloride is a red mass. Iridium dichloride is an olive-green, 
and the di-iodide a brown, powder. The tribrornide forms olive- 
green crystals. Platinum tetrafluoride is a buff-yellow crystalline 
deliquescent mass. The tetrachloride forms orange-brown crystals 
containing water. The tetrabromide is a non-deliquescent black 
mass, soluble with brown colour. The dichloride and dibromide 
are greenish-brown masses. These substances are all easily decom- 
posed by heat. The following are soluble in water: OsCl 2 , 
dark-violet; OsCl 3 , green; OsCl 4 , red. IrCl 3 and IrCl 4 are de- 
liquescent ; PtP 4 , yellow ; PtCl 2 , oransro ; PtCl 4 , orange-brown ; 
IrBr 3 , olive-green ; IrBr 4 , red. Osmium tetrachloride decomposes 
on addition of much water. 

Double halides. These bodies are, as a rule, crystalline in this group, and 
are more stable than the simple halides. The following is a list : 



OsClo.wKF. IrCLz.nKOl. 

Pt01 2 .KCl. 

PtCl 4 .2K01. 

OsCl 3 .3KCl. IrCl 2 .3KCL 

PtCl 2 .2HCl. 

PtOl 4 .2NH 4 Ol. 

OsCl 4 .2KCl. Ir01 3 .3AgrCl. 

PtCl 2 .2KCl. 

PtCl 4 .BaCL. 

OsCl 4 .2NaCl. IrCl 4 .2KCL 

3PtCl 2 .2AlCl 3 . 

PtCl 4 .A101 8 . 

OsCl 4 .2AffCl. 

2PtCl 2 .Sn01 4 . 

PtCl 4 .FeCl 3 . 

PtCl 4 .SnCl 4 . 

PtCl 4 .Se01 4 . 


Bromides and Iodides. 

PtBr 2 .2KBr. IrBr 3 .3HBr. IrBr 4 .2KBr. Ir01 4 .NH 4 I. 
PtBr2.CuBr 2 . IrBr 3 .3KBr. PtBr 4 .2HBr. IrI<.2NH 4 I. 
IrI 2 .2NH 4 I. IrI 3 .3KI. PtBr 4 .2KBr. PtI 4 .2KV 

IrI 3 .3AffI. PtBr 4 .BaBr 2 . 
Also, PtOl 4 .PtI 4 , or PtCl 2 I 2 is known. 

A compound of platinum dichloride with phosphorus trichloride is formed 
by heating to 250 spongy platinum with phosphorus pentachloride ; its formula 
is PtCl 2 .PCl 3 . The resulting crystals melt at 170, and are soluble in carbon 
tetrachioride and in chloroform. It combines with chlorine to form the double 
compound PtCl a .PCl 4 . 

The most important of these compounds are P^Cl^HCI, pro- 
duced by direct addition, and the corresponding potassium and 
ammonium compounds, produced by double decomposition, thus : 

PtCl 4 .2HCl.Aq + 2KCl.Aq = PtCl A .2KCl + 2HCl.Aq. 

These compounds are yellow crystalline powders, sparingly soluble 
in water, and nearly insoluble in a mixture of alcohol and ether. 
As the similar sodium platinichloride dissolves in these solvents, 
potassium and ammonium are usually separated from sodium by 
precipitation as platiniehlorides, and weighed as such. The 
ammonium salt at a red heat yields spongy platinum as a porous 
grey metallic mass. Alt these compounds, indeed, lose halogen 
when heated, leaving a mixture of the metal of the platinum 
group with the halide of the conjoined metal. 

The mass of 1 c.c. of platinum dichloride is 0'87 gram at 11. The mass of 
1 c.c. of many of the platinichlorides has also been determined, but with these 
exceptions the physical constants are unknown. 

Halides of Copper, Silver, Gold, and Mercury. 

These elements resemble each other in their monohalides. The 
monochlorides, bromides, and iodides are all insoluble in water. 
They have a certain analogy with the compounds of the palladium 
and platinum groups, and in their formula) correspond with those of 
the elements of the potassium group, in which the first three 
members are classed. Mercury, in the periodic table, is the last 
element in the magnesium group, which it resembles in the formulae* 
of its dihalide compounds. 

Sources. Silver chloride, AgCl, occurs native as horn silver, 
or kerar^yrite 9 in waxy translucent masses. Bromargyriie is the 


name of native silver bromide, a lustrous yellow or greenish 
mineral. Chlorobromid.es of silver of the formulae 3AgCl.AgBr, 
3Aggi.2AgBr, SAgBr.AgOl, 5AgCL4AgBr, and 3AgCl.AgBr 
also \>ccur native. Native iodide or iodargyrite is also found in 
yellow-green masses. AgCl.AgBr.AgI has also been found 

Mercurous chloride, HgCl, or horn quicksilver, accompanies 
cinnabar, HgS, occurring in dirty-white crystals. 

The following halides are known : 

Fluorine. Chlorine. Bromine. Iodine. 

Copper.. Cu 2 F 2 ; CuF 2 . Cu 2 Cl 2 ; CuCl 2 Cu 2 Br 2 ; OuBr 2 . Cu 2 I 2 ; CuI 2 . 

bilver .. Agr^. Ag-Cl. AgBr. Agl. 

Mercury. HgF; HgrF 2 . Hg-Cl; HgrCl 2 . Hg-Br; HgrBr. 2 . HgrI; H?I 2 . 

Gold... AuCl;AuCl 2 ; AuBr Aul 

AuCl 3 . AuBr,. 

Preparation. 1. By direct union. Fluorine, chlorine, bro- 
mine, and iodine attack these elements when finely divided in the 
cold, but the action is promoted by heat. In this way cuprous 
and cupric chlorides and bromides, and cuprous iodide have been 
prepared, the monohalide being formed in presence of a small 
amount of halogen ; but the dihalide with excess of halogen. Silver 
chloride, bromide, and iodide, inercurous and mercuric chloride 
and iodide and mercurous bromide, and gold dichloride, AuCl 2 
or Au 3 Cl4, and the corresponding bromide, AllBr^ or Au^Br 4 , have 
also been thus obtained. 

The higher halides are often prepared by the action of a 
mixture of nitric and hydrochloric, or nitric and hydrobromic, 
acids on the elements (see p. 172). The free halogen attacks the 
metal, forming the halide. Thus mercuric chloride, HgCL, cupric 
chloride, JuCL, and auric chloride and bromide, AuCl 3 and AuBr 3 , 
are produced in solution by this means: 3Hg + GHCl.Aq -f 
2HN0 3 .Aq = 3HgCl 2 .Aq + 4H.O + 2.VO; Au + 3HBrAq + 
HN0 3 .Aq = AuBr 3 .Aq 4- 2H 2 + NO. 

2. By the action of the halide of hydrogen on the metal. 
A solution of hydrogen iodide dissolves silver, forming the 
double halide, Agl.HL Hydrochloric acid dissolves copper in 
presence of air : Cu -}- + 2HCl.Aq = CuCl 2 .Aq -f H,0. 

3. By heating a higher halide. Oupric chloride and 
bromide, CuCl 2 and CuBr 2 , when heated, yield cuprous halide, 
Cu 2 Cl> and CuJBr 3 ; and cupric iodide decomposes spontaneously 
into cuprous iodide, Cu 2 I 2 , and iodine. Aurous chloride is pro- 
duced at 185 from auric chloride, and auric bromide yields 


aurons bromide at 115. Auric iodide decomposes spontaneously 
into anrons iodide and iodine. 

4. By the action of the metal on the higher halicte. A 
solution of cupric chloride in hydrochloric acid, when shaken with 
scraps of metallic copper, is converted into the dichloride, thus : 
Cu -f CuCl 2 .nHCLAq = Cu 2 Cl 3 .wHCl.Aq. Mercuric chloride or 
bromide triturated with mercury yields mercurous chloride or 

5. By double decomposition. (a.) By the action of the 
halogen acid on the oxide or carbonate of the metal. 
All these compounds may be thus prepared. It is, however, not 
convenient for the preparation of insoluble con? pounds, inas- 
much as the oxides, being insoluble, become coated over with a 
film of the insoluble halide and protected from the further action 
of the halogen acid. The following compounds have been pre- 
pared thus: CuiF 2 , CuF 2 , Cu 2 01 2 , CuCl,, CnBr,, AgF, HgF (by 
the action of HF on Hg 2 0) ; HgF 2 , HgCl a , HgBr 2 , Aul (from 
Au 2 O 3 and HI, thus : Au 2 O 3 + 6HI = 2 Aul + 3H,O + I 2 , 
the auric iodide, AuI 3 , decomposing at the moment of its forma- 

(6.) Other cases of preparation by double decomposition : 

Cu 2 Cl 2 . This is the best method of preparation. A strong 
solution of copper sulphate, CuS0 4 , and sodium chloride, 
NaCl, in equivalent proportions, is saturated with sulphur 
dioxide. The sulphur dioxide liberates hydrogen from 
water, itself forming sulphuric acid ; and the nascent 
hydrogen removes chlorine from cupric chloride, produced 
by the interaction of copper sulphate and sodium chloride, 
precipitating cuprous chloride, thus : 

CuS0 4 .Aq + 2NaCl.Aq = CuCl,.Aq + Na 2 S0 4 .Aq; and 

2CuCl 2 .Aq + 2H,0 + SO,.Aq = CuCl 2 + H,S0 4 .Aq + 


Hg 2 Cl 2 may be similarly prepared from mercuric chloride, 
HgCl 2 , and sulphur dioxide. 

CuJ 2 . Copper sulphate, or any other soluble salt of copper, 
reacts with potassium iodide, giving in very dilute solution 
a blue solution of cupric iodide; in strong solution the 
cupric iodide decomposes into cuprous iodide and free 
iodine. The reactions are as follows: 

CuSO 4 .Aq + 2KLAq = CuI 2 .Aq -f K a S0 4 .Aq; and 
2CuI 2 .Aq = Cu*I 2 + If + Aq. 


AgCl, AgBr, and Agl. These are prepared by adding a soluble 
'salt of silver, e.g., the nitrate, to the required halide of 
t hydrogen, or to any other soluble halide, thus : 

AgNO 3 .Aq + KI.Aq = Agl + KNO 3 .Aq. 

AuF 3 ? An attempt to prepare auric fluoride by adding silver 
fluoride to auric chloride resulted in the precipitation of 
auric oxide, Au 2 s , through the action of water on the 
fluoride, thus : 

2AuF 3 + 3H 2 O = Au 2 O 3 + 6HF., 

AuI 3 . Auric iodide is formed by the addition of auric chloride 
to potassium iodide, thus : 

AuCls.Aq + 4KI.Aq = AuT 3 .KI.Aq + SKCl.Aq. 

The double iodide is decomposed on addition of more auric 
chloride, with precipitation of. auric iodide : 

3KI.AuIa.Aq + Au01 8 .Aq = 4AuI 3 + SKCl.Aq. 

HgF. Mercurous chloride, digested with silver fluoride, yields 
mercurous fluoride and silver chloride, thus : 

AgF.Aq + HgCl = AgCl + HgF.Aq. 

HgCl, HgBr, and Hgl. By precipitation. Mercurous 
nitrate, Hg(N0 3 ), and a soluble halide yield mercurous 
halide and a soluble nitrate, e.g., HgNO 3 .Aq + NaCl.Aq 
= HgCl -f- NaNOa.Aq. Another method of preparing 
HgCl is to sublime mercurous sulphate, Hg 2 SO 4 , with salt, 
NaCl : 

Hg 2 SO 4 + 2NaCl == 2HgCl + Na 2 SQ 4 . 

l? 1 . Mercuric sulphate, HgSO 4 , and salt yield mercuric 
chloride on sublimation ; hence its name corrosive sublimate. 
HgI 2 . Mercuric iodide, being insoluble, is precipitated by 
addition of mercuric chloride to potassium iodide. The 
sesquiiodide, HgI 2 .HgI, is similarly precipitated from a 
mixture of mercurous and mercuric nitrates by potassium 

Properties. These substances are all solid. The cuprous and 
mercurous, and the silver and aurous compounds are all insoluble 
in water, but dissolVe in concentrated halogen acids ; mercurous 
and . anrou talides are decomposed when boiled with acids. 
Cuprous fluoride is a red powder, fusing to a black mass ; when 


prepared by precipitation it is white. The chloride is also white, 
but is affected by light, which turns it dirty violet; it appears to lose 
chlorine. The bromide is greenish-brown, and the iodide brownish- 
white. Silver fluoride is a white soluble mass ; the chloride is white, 
but turns purple on exposure to light. This is said to be owing to 
the formation of asubchloride, Ag a Cl, inasmuch as the purple sub- 
stance is not dissolved by nitric acid, in which silver itself is soluble. 
The bromide is pale-yellow, and the iodide darker yellow. These 
substances are used to detect and estimate the halogens, for they 
are almost absolutely insoluble in water. They melt to horny 
masses. Mercurous fluoride is a light-yellow crystalline powder, 
partly decomposed on boiling with water, and decomposed by 
heat. The chloride, the common name for which is calomel, is 
dirty white in colour, and also partially decomposes when volati- 
lised, but its constituents recombine on cooling ; hence it can 
be sublimed. It condenses as a fibrous, translucent, very heavy 
solid. It is quite insoluble in water. The bromide is also a 
fibrous yellow mass. The iodide is a greenish-yellow powder, 
sparingly soluble in water. 

Aurous chloride, AuCl, is white, insoluble in water, but decom- 
posed on boiling with water into gold, and auric chloride, ATlCl 3 . 
The bromide is also insoluble in water, and yellowish-grey in colour. 
It is decomposed by hydrobromic acid, thus : 

SAuBr + HBr.Aq = AuBr 3 .HBr.Aq + 2Au. 

Aurous iodide is an insoluble yellow powder, soluble in hydriodic 

The higher halides are all soluble in water. Those of mercury 
and cupric chloride are also soluble in alcohol and in ether. 

Cupric fluoride forms sparingly soluble blue crystals ; mercuric 
fluoride is a white crystalline mass. 

Cupric chloride is a brownish-yellow deliquescent powder ; it 
dissolves in water with a blue colour, and deposits blue crystals 
of CuCl 2 .2H 2 O. The bromide consists of iron-black crystals, 
soluble in water with a brown colour. 

Gold dichloride,* AUgCl^ is regarded as a compound of AuCl 3 
with AuCl. Its molecular weight, however, is unknown. It is a 
hard dark-red substance, decomposed by water into AuCl 3 and 
AuCl. The trichloride, AuCl 3 , forms dark-red crystals, and is 
soluble in water, alcohol, and ether. The dibromide is a black 
substance, which reacts with water like the corresponding 
chloride, yielding monobromide and tribromide, The latter is 
* J. praM. Chem. (2), 37, 105. 


dark- brown and dissolves in water, alcohol, and ether. Auric 
iodide, AuI 3 , is a dark-green precipitate, decomposing spontane- 
ously % into aurous iodide and iodine. 

Mercuric chloride, or corrosive sublimate, is a white crystalline 
substance ; 100 parts of water dissolve 7*4 parts at 20 ; 100 parts 
of alcohol dissolve 40 parts at the ordinary temperature. The 
bromide crystallises in soft white laminae. The iodide is a 
scarlet powder, sparingly soluble in water, more soluble in alcohol 
and ether. It crystallises from aqueous potassium iodide in red 
octahedra. When sublimed, it condenses in yellow prisms, which, 
when rubbed, suddenly change into red octahedra. 

Physical Properties. 
Mass of 1 c.c. Melting-point. Boiling-point. 

F. 01. Br. I. F. 01. Br. I. F. 01. Br. I. 

Copper. ? (ous) 4'72 5'70 908 434 504 601 ? 954f 861 759 

Silver.. 5-505 6'215 5'67 ? 451 427 527 ? ? ? White 

atO at 17 . heat. 

Gold. . . ? ? ? ? 250* 115* * * * * * 

M __f 6-56 (ous) 7'31 7'64 ? 405 290 ? 400 J? 310 

r Ury I5'45(ic) 573 6-30 130* 288 244 238 ? 303 319 339 

Double halides. Cupric fluoride is said to combine with the fluorides of 
the alkaline metals to form black compounds. The following compounds of the 
other halides have been prepared: 

Cu 2 Cl 2 .4HCl. HffCl 2 .KCl. CuCl 2 ,2HOl. AuOl 3 .KCl. 

2HffCl 2 .CaCl 2 . CuCl 2 .2KCl. AuCl 3 .NaCl. 

HgrBr 2 .KBr. CuCl^NB^Cl. 2AuCl 3 .CaCl 2 . 

A?G1.NH 4 O1. 2HffBr 2 .SrBr 2 . HgrOl 2 .2NH 4 Cl. 2AuOl a .ZnCl 2 . 

AgOLKGl. HgrI 2 .KI. HgrI 2 .2NH 4 I. AuBr 3 .HBr. 

Affl.HI. H*C1 2 .NH 4 C1. HgCl 2 .2KCl. AuBr 3 .XBr. 

Ag-I.KX Hfflo.Hgrl. H?I 2 .2KI. AuIy.KI. 

Ag-I.2KI. 2HgrI 2 .BaI 2 . HfirI 2 .HfirGl 2 . 

2H?Gl.SrGl 2 . 2H*Ol 2 .HgI 2 . 
2Hg-Cl.SCl 2 . 

Besides these, 2Hg01 2 .K:Cl, 3HgCl 2 ,Mg01 2 , and SHgCL.CaCla are known, in 
which the mercuric chloride bears a larger ratio to the other chloride than in 
the tabulated examples. The name aurichlorides (sometimes, but incorrectly, 
" chloraurates ") has been applied to the compounds of auric chloride. The 
compound CujI 2 .HgI 2 is a red body, and has the curious property of turning 
black when heated. It has been used as a means of indicating whether the 
axles of engines become superheated. The compound HgCl 2 .2NH 4 Cl has been 

* Decomposes. 

f Between 954 and 1032 3 ; CuCl 2 , 498; Cu 2 Br a , 861-964. 
J Sublimer between 400 and 500 without melting. 

N 2 


known since the times of the alchemist, and was termed by them tal alemoroth. 
All these bodies are prepared by direct addition. Those of silver are decom-. 
posed on dilution, giving precipitates of halides. The compound HgCl 2 .SnCl 2 
is produced by subliming an alloy of tin and mercury with mercurous chlpride. 

The molecular weights of some of these compounds have been 
determined. The density of cuprous chloride, Cu 2 Cl 2 , was found to 
be 102*0, while the calculated number for tha.t formula is 106*86.* 
Silver chloride gave a density corresponding to the molecular 
weight 160'8, instead of the theoretical one, 143'39, for the 
formula AgCl.f As regards mercurous chloride, it 'is most pro- 
bable that the molecular weight is that equivalent to the formula 
HgCl. It is not difficult to vaporise mercurous t chloride ; the 
difficulty has been to ascertain whether it decomposes, in the state 
of gas, into mercuric chloride, HgCl 2 , and mercury, or is stable. 
In each case the density found corresponded to the formula 
HgCl, not to the formula Hg 2 Cl 2 . The actual number was 
231'8; the calculated molecular weight, 235'4. The density was 
determined in presence of an atmosphere of mercuric chloride, and 
under these circumstances little or no dissociation takes place. J 

The molecular weights of the remaining halides are unknown, 
but the formulae have been made to accord with those of which the 
value has been ascertained. 

* Zerichte, 2, 1116. 

t Proc. Roy. Soc. JSdin., vol. 14 

J Gazzetta, 1881, 341 ; Chem. Soc. Abs., 42, 466. 





Having concluded the description of the compounds of the 
halogens with other elements, and with each other, it may be here 
advisable to give a summary of their leading features. This will 
be done in the same order as that observed in the special descrip- 
tion of each class of compounds, viz., their sources, their prepara- 
tion, and their properties. 

1. Sources. If a compound occur free in nature, it must 
either be unacted on by substances around it at the temperature 
at which it exists, or must have only an ephemeral existence. 
The two most important and widely spread agents are the oxygen 
of the air and water. It must, therefore, be able to resist the 
combined action of both of these substances. 

As an instance of a compound produced under certain unusual 
circumstances, hydrogen chloride may be named. It is found in 
the air and water in the neighbourhood of volcanoes ; but, although 
not altered by air or water, it soon is dissolved by the rain, and 
reacts with the constituents of the soil, forming chlorides of cal- 
cium, sodium, potassium, &c., which ultimately find their way into 
the sea, being carried down by rivers. It is, therefore, only found 
in the locality where it is formed before it has been exposed to 
those influences. Ferric chloride, Fe 2 01 6 , occurs under similar 

The chlorides, bromides, and iodides of lithium, sodium, potas- 
sium, calcium, and magnesium are all soluble in water. It is net 
improbable that they are partially decomposed by solution ; thus, 
for example, NaCl + H 2 = NaHO + HC1. But when such a 
solution is evaporated, the reaction, if there is one, occurs, in the 
inverse sense, and the water evaporates, leaving the chloride. By 
the evaporation of inland lakes, such as the Dead Sea, these salts 
are deposited. Such has doubtless been the case where mines of 


known since the^t ; and at Stassfurth, in N. Germany, the layers of 
All these bodiind in the order of their solubility, the least soluble 
posed on ne lowest layers, 
is P r du >l a ble salts, such as fluorspar (calcium fluoride), cryolite 

rminium sodium fluoride, AlP 3 .3NaF), silver chloride, bromide, 
frA iodide, lead chloride, <fec., which are not attacked by water or 
oxygen, are also found in nature. 

Preparation. The general methods of preparation may be 
summed up as follows : 

1. Direct union. The halidcs may, as a rule, be thus pro 
pared. Fluorine appears to act on all elements, oxygen and 
nitrogen excepted, at the ordinary temperature. The metals 
iridium and platinum are, perhaps, the least affected of any in 
the cold ; hence the use of an alloy of these metals in forming the 
vessel in which fluorine was isolated by electrolysis. Chlorine, 
when dry and cold, appears not to attack some metals, such as 
sodium and zinc, which are readily acted on when hot ; but, as a 
rule, the elements combine with chlorine, bromine, and iodine 
when heated in contact with them. Those which do not combine, 
even at a red heat, are carbon, nitrogen, and oxygen. 

2. Replacement. Action of a compound of the halogen on the 
element ; or action of the halogen on a compound of the element. 
The most common instance of the first method is the action of the 
halide of hydrogen on a metal. A list of the elements not thus 
attacked is given on p. 112. But there are many other processes 
involving similar reactions, where the method is not used as a 
means of preparing a halide, but of liberating the element with 
which the halogen was in combination. The elements magnesium, 
boron, aluminium, silicon, and others are prepared by the action 
of sodium or potassium on their halides, which, of course, results in 
the formation of sodium or potassium halides. The action of the 
halogen on a compound of the element, of which the halide is 
required, is also a method not frequently employed ; for, owing to 
the fact that there are few elements which do not combine with 
the halogen, a mixture of two halides is thus obtained, which are 
often not easily separated. An instance of its application, how- 
ever, is found in the preparation of hydrogen iodide, by the action 
of iodine and water on hydrogen sulphide ; and of carbon tetra- 
chloride, by the action of chlorine on carbon disulphide. The 


perhaps, the most usual method of preparing compounds of the 
halogens. As a rule, the resulting halide must be gaseous or solid, 
or wtfler or hydrogen sulphide must be the product of the action. 
Instances of such action are very numerous. Among them may 
be mentioned the action of sulphuric or phosphoric acid on halides 
of the metals, whereby the hydrogen halide is formed ; the 
action of the halides of boron, silicon, phosphorus, &c., on water ; 
the action of a halide of hydrogen on oxides, hydroxides, sulphides, 
or carbonates of the metals ; the action of calcium chloride on 
barium sulphate at a red heat ; the precipitation of calcium 
fluoride ; the preparation of magnesium chloride ; of boron fluoride : 
boron chloricfcj ; and many other cases. The method is almost 
universally applicable ; but it does not yield halides of nitrogen or 
of oxygen. 

A special method, applicable to the preparation of aluminium 
chloride, is the action of the vapour of carbon tetrachloride on the 
red-hot oxide. The simultaneous action of carbon and chlorine on 
the oxides of silicon, boron, &c., at a red heat can hardly be 
considered double decomposition, inasmuch as the chlorine and 
carbon are not combined, but it is difficult to classify such actions 
elsewhere, unless they be regarded as cases of direct union. 

To distinguish the halogens when all three may be present, the 
mixture is distilled with strong sulphuric acid and potassium di- 
chromate. If chlorine be present, the volatile chromyl chloride, 
CrOCl 2 , is produced, and distils over. If the distillate contains 
chlorine, chromium will be found therein. To detect bromine and 
iodine in presence of each other, chlorine-water is gradually added 
to the solution of their sodium or potassium salts, and the liquid 
is shaken with carbon disulphide or chloroform, which do not mix 
with water. If iodine be present, a violet solution is obtained; if 
bromine he also present, further addition of chlorine-water will 
destroy the violet colour of the chloroform or carbon disulphide, 
and it will be replaced by an orange-red colour. 

4. If two or more halides exist, the compound containing most 
halogen may almost always be prepared by heating the one con- 
taining less with the required halogen. Thus, iron diohloride 
yields the trichloride when heated in chlorine ; mercurous is con- 
verted into mercuric chloride ; stannous into stannic, &c. 

5. By heating the higher halide, in certain cases, the 
halogen is evolved; and the lower halide is left. Thus, thallic 
chloride, T1C1 3 , yields thallous chloride, T1C1, when heated ; and 
auric chloride similarly gives aurous chloride, two atoms of chlorine 
being lost. 


Sometimes, but rarely, the lower halide decomposes into the 
element and the higher halide. This is the case with bismuth 
dichloride, BiCl 3 . It is sometimes necessary to heat in contact 
with some element capable of combining with the halogen. For 
example, aluminous sodium fluoride, AlP 2 .2NaF, is prepared 
by heating cryolite with metallic aluminium; the compounds 
GaCl 2 , In01 2 , and InCl, by heating GaCl 3 and InCl 3 , with 
gallium and indium respectively ; disilicon hexachloride is similarly 
prepared from the tetrachloride ; and chromous chloride, CrCl 2 , 
results from the action of hydrogen at a red heat on CrCl 3 ; 
the lower chlorides of titanium, molybdenum, and tungsten are 
also prepared thus. ' 

Sometimes the removal of halogen from the higher halide may 
be accomplished in solution. Thus, the familiar operation of 
" reducing " ferric chloride in solution by means of the hydrogen 
generated from zinc and hydrochloric acid, or by sulphur dioxide, 
or by stannous chloride, falls under this head ; also the formation of 
mercurous from mercuric chloride, and that of osmium di- and tri- 
chlorides, and iridium di-iodide. Hydrogen sulphide is also used as 
a reducing agent for ferric halides, for rhodium trichloride, &c. 

Properties. (a.) Physical properties : Colour. The colour 
of objects is due to their absorbing light rays of certain wave- 
lengths in the visible part of the spectrum. It is to be noticed 
that the iodides of those metals which form white fluorides, 
chlorides, and bromides often are yellow or red ; as examples, the 
cases of thallium, silver, mercury, &c., may be noticed. In general, 
those halides with higher molecular weights towards the end of 
the periodic table display colour. But substances which appear 
colourless to our eyes have the power of absorbing vibrations of 
wave-lengths which do not affect our sight, and to eyes sensitive 
to other scales of vibration than ours such bodies woufd appear 
coloured. It may also be generally stated that halides containing 
a large proportion of halogen display colour when those containing 
less are colourless. 

Form. The halides are almost without exception crystalline, 
but up to the present their crystalline form has not yet been 
connected with their chemical nature (see Isomorphism, 
Chap. XXXV). 

State of aggregation. Compared with the oxides and sulphides, 
the halides may generally be said to be easily fusible and volatile. 
This is probably due to their simplicity of structure and low 
molecular weight. The fluorides, however, have, as a rule, greater 
complexity than the chlorides, bromides, and iodides. For example, 


hydrogen fluoride is known to have a more complex molecule than 
hydrogen chloride, even in the gaseous state (see p. 115) ; and the 
non-volatility of many fluorides, compared with the volatility of 
the corresponding chlorides, would lead to the inference that their 
molecules are complex. Some fluorides, however, such as those 
of boron and silicon, have undoubtedly simple formula ; and it is 
to be remarked that these bodies are very stable. The comparative 
insolubility of many fluorides, e.g., those of calcium, strontium, 
barium, magnesium, tin, &c., may also point to complex molecular 
structure ; and further evidence may be derived from the fact that 
the fluorides form double compounds more easily than the other 

The solubility of a compound, however, may perhaps partly 
depend on its chemical action on the solvent, though probably not 
invariably. It certainly appears to be connected with simplicity 
of molecular structure, implying low molecular weight. 

The mass of one cubic centimetre of the halides also shows regu- 
larity. The iodides are, as a rule, specifically heavier than the 
bromides ; the bromides than the chlorides ; the chlorides, how- 
ever, are not always heavier than the fluorides ; but, again, this 
may depend on molecular complexity, contraction always occurring 
when chemical union occurs, even between molecules of the same 
kind. It is also to bo noticed that, in each group of elements, the 
halides of those which possess the highest atomic weights are 
specifically heavier than the earlier members of each series. 

(b.) Chemical properties. Some halides, when heated, 
decompose into their elements, or into lower halides and halogen. 
It is probable, indeed, that at a sufficiently high temperature all 
chemical compounds would decompose thus. In certain cases, for 
example, the halides of nitrogen, oxygen, and carbon, when the 
elements itre once apart, they do not again combine. The halides of 
oxygen and nitrogen are formed, not, as usual, with evolution of 
heat, but with absorption, and suqh compounds are always readily 
decomposed. Those of nitrogen and of oxygen are exceedingly 
explosive, and cannot be produced by direct union. Other halides, 
such as those of gold, platinum, &c., decompose into their elements 
when Jieated, but if kept in contact the elements would again 
recombine. But, as the metallic element is volatile only at a very 
high temperature, the halogen, which is easily volatile, distils 
away, leaving-the metal. Other halides, such as the higher ones 
of selenium, phosphorus, and antimony, are also decomposed, out 
the lower halide is not so different in volatility from the halogen 
itself ; hence, the two are difficult to separate. When a compound 


decomposes into constituents which reunite on cooling, it is said 
to dissociate. The term decomposition includes dissociation, but 
may be employed in the stricter sense of splitting up without 
recombination. There is a temperature of decomposition peculiar 
to each compound, at which, if recombination does not occur, after 
sufficient time, all the compound would be decomposed; whereas, 
if recombination is possible, a state of balance is maintained, the 
relative proportions of the constituents depending on the tempera- 
ture, on the pressure, and on the relative amounts of the con- 
stituents. Excess of one constituent prevents decomposition. 
Thus, phosphorus pentachloride is stable in the gaseous form in 
presence of excess of chlorine or of phosphorus tlichloride, and 
mercurous chloride can exist as gas in presence of mercuric 
chloride. These statements probably also apply to bodies in solution. 

The halogens are capable of replacing each other. Here, 
again, the relative amounts have a great influence on the result. 
Bromine replaces iodine from its compounds with elements of the 
potassium, calcium, and magnesium groups dissolved in water; 
and chlorine replaces bromine and iodine. But a current of 
bromine vapour led over hot potassium chloride results in the 
formation of potassium bromide. Again, on digesting precipitated 
silver chloride with bromine- water, silver bromide is formed ; and 
iodine, under similar circumstances, replaces both chlorine and 
bromine. Yet, on heating silver iodide in a current of chlorine or 
bromine, the iodine is expelled, and replaced by chlorine or 
bromine. In these cases,, the mass of the halogen acting on the 
halide has the effect of reversing the process which takes place 
in presence of water. 

Combinations. The halides of the elements in most cases com- 
bine with water to form crystalline compounds containing water 
of crystallisation. It is sometimes, but not always, possible to 
expel such water by heat ; in many cases, the water reacts with 
the halide, forming hydroxide, oxide, or oxyhalide. The crystalline 
form is altered by the presence of the water, and when several 
hydrates exist, they have usually different crystalline forms. The 
lower the temperature, the greater the amount of water with 
which the substance will combine. A halide crystallising without 
water at the ordinary temperature sometimes forms a hydrate at 
low temperatures, as is the case with sodium chloride. The 
remarkable change of colour of some halides, e.g., those of nickel, 
cobalt, iron, <frc., when hydrated appears to point to some profound 
modification in molecular structure by hydration ; and the per- 
sistence ^ *his colour in dilute solution leads to the inference 


that the hydrate exists dissolved in the water. It has been 
pointed out that compounds of halides with hydrogen halides 
invariably contain two molecules of water of crystallisation for 
every molecule of hydrogen halide present. 

Double halides. The halides in almost all cases, as has been 
seen, combine with each other, forming double compounds. These 
are usually prepared by mixing solutions of the two halides of 
which it is desired to form a compound, and evaporating the mix- 
ture, best at the ordinary temperature, for a low temperature is 
favourable to combination. The compounds with halites of hydro- 
gen are generally, but not always, called acids. In many cases 
they are exceedingly unstable, and mere removal from the presence 
of a strong solution of the halogen acid is sufficient to decompose 
them, the hydrogen halide escaping as gas. They usually crystallise 
with water, if, indeed, they can be obtained crystalline ; the anhy- 
drous compounds are rare. Of the four halogens fluorine is most 
prone to form double compounds. This is probably connected with 
the tendency of its compounds to polymerise^ i.e., the tendency for 
several molecules to enter into combination with each other. It is 
probable, indeed, that there is no difference in kind between com- 
pounds of two molecules of the same halide, such as Pe 2 Cl 6 (which 
may be regarded as a compound with each other of two molecules 
of FeCl 3 ), and compounds produced by the union of the halides of 
two different elements, such as Pt01 4 .2KCl, SbCl 6 .SCl 4 , &c. ; such 
bodies, however, exhibit very different degrees of stability, certain 
of them withstanding a fairly high temperature without decompo- 
sition, so far as can be ascertained, while others exist only at a low 
temperature. If one of the halide constituents of a double halide 
is easily decomposed by heat, it is usually rendered more stable by 
combination ; although on heating such a double halidc, the more 
easily decomposable halide is decomposed, while the more stable 
one resists decomposition. An ins-tance is given above; SbCl 5 .SCl4 
is stable at the ordinary temperature, while SC1 4 can exist only 
below 22 ; but on heating the double halide chlorine is evolved, 
while the stable chloride of sulphur, S 2 C1 2 , is formed, the anti- 
mony pentachloride remaining unaffected. Similarly the other 
double halide mentioned above, PtCl 4 .2KCl, when heated, decom- 
poses , a mixture of metallic platinum and potassium chloride being 
left, while chlorine is evolved. Here, again, the comparatively 
unstable platinum tetrachloride is decomposed, the stable potassium 
chloride resisting decomposition. It is said that solution in watei 
decomposes such double halides into their constituent halides. 
But it appears more likely that the degree of decomposition 


depends on the relative proportion of water and doable balide, 
and on the temperature of the solutions ; and that such a solution 
really contains in many cases both the double halide and the, two 
simple halides. With increase of solvent, or with rise of tem- 
perature, it is probable that the relative amount of the double 
halide decreases, while that of the single halides increases. These 
are matters, however, still involved in considerably obscurity. 

Action of water. The action of water on many of the halides is 
to decompose them, hydrogen halide and the oxyhalide or hydrox- 
ide of the element being produced. The following halides are 
known to be thus decomposed by water : (a.) At the ordinary tem- 
perature : Halides of boron, silicon, zirconium, germanium ; tetra- 
halides of tin ; halides of phosphorus,, arsenic, antimony, bismuth, 
vanadium, niobium, tantalum, molybdenum, tungsten, uranium, 
sulphur, selenium, and tellurium. In certain cases the halide is 
not decomposed in presence of great excess of hydrogen halide, 
even although water be present, possibly owing to the formation of 
a double halide of the element and hydrogen. This is known to be 
the case with the fluorides of boron and of silicon, which form the 
compounds BF 3 .HP, and SiF 4 .2HF, which are stable even in pre- 
sence of a large amount of water. Arsenic, antimony, and bismuth 
trihalides dissolve in excess of halogen acid, probably forming 
similar stable compounds. (6.) At a red heat, most of the halides 
react with water-gas to form the oxides, those of lithium, sodium, 
potassium, rubidium, and caesium excepted. 

It is, however, not improbable that, as has been already stated, 
solutions of all halides in water are partially decomposed by the 
water, sodium chloride, for example, reacting to form sodium 
hydroxide and hydrogen chloride, thus : JSTaCl + H 2 O = NaOH 
-|- HC1 ; and so with other chlorides. The degree of this decom- 
position depends, no doubt, largely on the relative aftnounts of 
water and halide, and on the temperature, and varies for each salt. 
The presence of a second halide appears in many cases to retard or 
diminish giich decomposition, and to render salts stable in solution 
which would decompose or react with water in their absence. 

Action of hydroxides. Halides which are not decomposed by 
water, so that their constituents can be separated, and which^are not 
re-formed on alteration of temperature, dilution, &c., can in most 
cases be decomposed by a soluble hydroxide. Thus sodium or 
potassium hydroxides react with almost all 1 halides producing 
hydroxides, that is, oxides in combination with water. Ammonia 
dissolved in water has in most cases a similar action, the solution 
acting as if it were hydroxide of ammonium, NH^H. In 


many instances, particularly if the element belongs td the class 
generally termed " non-metals," the hydroxide produced com- 
bines with the reacting hydroxide, forming a double oxide, or 
salt, and water. Oxides such as these are termed " acid-forming 
oxides," or "chlorous" oxides; those which have less tendency 
to such combination being named " basic " or " basylous " 
oxides. The following instances will exemplify what has been 
stated : 

The action of potassium hydroxide on cupric chloride is to 
form potassium chloride and cupric hydroxide, thus : 

CuCl 2 .Aq + 2KOH.Aq = Cu(OH) 2 + 2KCl.Aq. 

Cupric hydroxide may be viewed as a distinct individual, 
or as a compound of cupric oxide, CuO, with water. This 
point will be discussed later. A great excess of caustic potash, 
KHO, develops the slight power of combination of copper oxide, 
which dissolves with a blue colour, forming, no doubt, some com- 
pound such as CuO.K 2 O, or Cu(OK) 2 . Such a compound is 
certainly formed by the action of zinc chloride, ZnCl 2 , on caustic 
potash, KOH, the body Zn(OK) 2 being produced. But this kind 
of change is the usual and normal one of the chlorides of those 
elements whose halides are decomposed by water ; thus phosphorous 
chloride at once gives with water phosphorous acid, H 3 P0 3 , or 
P(OH) 3 (?), and with caustic potash, KOH, potassium phosphite, 
the caustic potash reacting thus with the phosphorous acid : 

2KOH.Aq + H 3 P0 3 .Aq = HK 2 P0 3 .Aq + 2H 2 0. 

As the hydroxides when heated are as a rule transformed into 
oxides with loss of water, this forms one of the most convenient 
methods of preparing hydroxides and oxides, as will soon appear. 

The fdrmulse of the halides are, as a rule, undoubtedly simple. 
It has already been remarked that we do not know with certainty 
the formulae of liquids and of solids, inasmuch as their molecular 
complexity is unknown. But it is probable that mere change of 
physical state from gas to liquid, or from liquid io solid, is not 
necessarily accompanied by chemical aggregation. Thus, if the 
formuja of hydrogen chloride as gas is HOI, and if no sign of 
aggregation is seen on its approaching its temperature of lique- 
faction ; that is, if its contraction on cooling runs pari passu with 
that of .hydrogen, fhere would appear to be no good reason to 
suppose that merely because it has liquefied its formula is thereby 
rendered more complex; but where, as in the case of hydrogen 
fluoride, distinct signs of molecular aggregation are to be noticed 


as the temperature falls, no doubt can be entertained as regards 
the fact that the molecular structure is complex in liquid hydrogen 
fluoride; but that it begins to occur before the liquid sta^te is 
reached would appear to negative the supposition that it is directly 
connected with change of state. In the present state of our know- 
ledge, therefore, it may be concluded that the formula possessed 
by a halide in the gaseous state also represents its molecular 
weight in the liquid condition, although there may well be examples 
of aggregation beginning in the liquid or solid states with fall of 
temperature, which are not to be detected by determination of the 
density of the gas. A full discussion of this point is better reserved 
until the oxides and sulphides have been studied^ for there is 
strong ground for the belief that their molecular structure is 

In every case, however, where the molecular complexity of a 
compound is unknown, the simplest formula have been adopted. 

These formulse are deducible : 

1. From the results of analysis, which yields the equi- 

valents of the elements, but gives no information 
as regards their atomic weights. 

2. By the law of simplicity, as applied by Dalton and 


3. By use of Avogadro's law, that equal volumes of gases 

contain equal numbers of molecules: the chief 
method of investigation being the method de- 
pending on the vapour-densities of compounds. 

4. From the atomic heats of the elements (Dulong and 

Petit's law). 

Other methods will be considered in a subsequent chapter. 


Detection and Estimation of the Halogens. Fluorine is detected by 
heating the suspected fluoride with strong sulphuric acid, and trying if the gas 
evolved will etch glass, i.e., will produce silicon fluoride. Chlorine, bromine, 
and iodine, when in combination, are detected by adding to a solution of the 
suspected compound in nitric acid a solution of silver nitrate. A chloride gives 
a white precipitate ; a bromide, a yellowish precipitate ; an iodide, a yellow 
precipitate. These may be further distinguished by addition of excess of 
aqueous ammonia. Silver chloride easily dissolves j the bromide is sparingly 
soluble ; and the iodide insoluble. 





The elements oxygen, sulphur, selenium, and tellurium, like 
the elements fluorine., chlorine, bromine, and iodine, combine 
readily with other elements, and many of their compounds have 
been carefully studied. Like the halogens, these four elements 
bear a marked resemblance to each other, oxygen being the 
analogue of fluorine, while the other three elements correspond 
more or less closely to chlorine, bromine, and iodine. The pre- 
vious arrangement of matter will be adhered to ; but additional 
paragraphs must be added, describing the double compounds of 
the elements of this group with those of the halogens and with 
each other. 

Compounds of Oxygen, Sulphur, Selenium, and 
Tellurium with Hydrogen. 

Hydrogen oxides, sulphides, selenide, and telluride ; H 2 ; 
H 2 2 ; H 2 S-, H 2 S 3 ; H 2 Se; H.Te. 

Sources. Water, H 3 0, is the most widely distributed of com- 
pounds, and occurs in larger proportion in nature than any other. 
It forms the sea, lakes, and rivers ; as ice it caps the tops of high 
mountains, and covers the land in the neighbourhood of the North 
and Soath Poles ; in the form of small liquid particles it forms 
clouds, fogs, and mist ; its vapour is always present in the atmo- 
sphere in greater or less amount, and is known as " humidity." It 
is a constituent of many minerals, and of all organised beings-, 
vegetable and animal, forming from 70 to 95 per cent, of their 
'weight. It is conjectured, from the appearance of the planets 


Mars and Venus, that their atmospheres contain wafer-vapour, 
and that their land is intersected by seas. It has not been proved 
to exist in the Moon, and it probably does not exist as such^in the 
Sun. V 

Hydrogen dioxide, H 2 O 2 , is present in minute amount in rain 
and snow, and in all natural waters,* and, being a body prone to 
give up oxygen, probably plays an important part in oxidising 
dead vegetable and animal matter. It appears to be produced 
by the evaporation or exposure to light of water in which oxygen 
gas is dissolved. 

Hydrogen sulphide, T 2 $, escapes from fissures in the earth 
in volcanic districts, and is a constituent of many mineral springs ; 
such waters are termed li hepatic," and are used as a cure for 
diseases of the skin. The wells at Harrogate are much fre- 
quented on this account. It is not widely spread, being slowly 
oxidised on exposure to air. Hydrogen selenide and telluride do 
not occur in nature. 

Preparation. 1. By direct union. (/i.) Water. A mix- 
ture of hydrogen and oxygen gases in the proportion of two 
volumes of the former to one volume of the latter explodes 
violently when heated to its igniting point at the ordinary pres- 
sure, forming water. The fact that by the union of hydrogen 
with oxygen water is the sole product was first proved by 
Cavendish, though its true nature was first determined by 

The combination may be easily shown by filling a strong soda-water bottle 
two-thirds full of hydrogen and one-third with oxygen, and after wrapping it 
in a cloth, for fear of the glass being shattered to fragments by the explosion, 
applying a lighted taper to the mouth. A violent explosion will occur, owing to 
the sudden expansion of the water-gas caused by the heat evolved by the union 
of its constituents. 

The quantitative relations between the volumes of the gases and their pro- 
duct, water-gas, may be shown in a more instructive manner as follows : 

A is a strong U"tube, of about 15 mm. in internal diameter, with platinum 
wires sealed through its upper end, surrounded by a jacketing tube, B, in the 
bulb of which water is boiled. A is filled with dry mercury, and placed in 
position in a mercury trough. A mixture, obtained by electrolysing water (see 
below) , of oxygen and hydrogen in the approximate proportions of two volumes 
of hydrogen to one volume of oxygen is introduced into the tube A, so^as to fill 
it about one-third. The water is then boiled so as to jacket the inner tube, A, 
with steam. The mixed gases expand, and when the temperature has become con- 
stant the mercury is run out by opening the stop-cock C until it is level in both 
limbs of the LHube. The level of the gases in then marked by an india-rubber 
i - L _ _ _ . .. . __ , . . 

Schone, Serichte, 7, 1693. 



ring, and mercury is again allowed to flow out so as to reduce the pressure on the 
gas. A spark from an induction coil is then caused to pass between the 
platinujn wires sealed through the glass. The gases are heated to their 
temperature of ignition j the portions thus heated unite, and the heat evolved 
by the union raises the neighbouring portions to their ignition-point. An 
explosion takes place, but owing to the increased volume of the gas, it is not 
so violent as it would be at atmospheric pressure and ordinary temperature. 

FIG 2<J 

The gases after combination contract, and, to bring them back to atmospheric 
pressure, mercury is poured into the open limb of the IJ-tube until it stands at 
equal height in both limbs. The volume of the water-gas is seen to be about 
bwo-thirds of that of the mixed gases before combination ; three volumes have 
become A wo. This experiment is adapted only as an illustration ; it is inaccu- 
rate owing to the non-introduction of various corrections; for example, a 
mixture of oxygen and hydrogen, prepared by electrolysis, contains ozone 
(see p 387), and hence occupies too small a volume; and some water- vapour 
condenses on the glass, and hence possesses a smaller volume than it ought to 

Ojryhydrogen blowpipe. By forcing mixed hydrogen and ox\gen gases 
through a narrow tube and setting them on fire, a pointed ilame ^p 


of a very high temperature. But the rate of explosion of a mixture of these gases 
is very rapid, and there is great danger of the explosion travelling back through 
the narrow tube and inflaming the mixture. Hence a special form of blowpipe 
must be employed. The temperature of ignition of the mixed gases is a big" 1 ! one ; 
probably 000 to 700 at the ordinary pressure. By cooling the gases below this 
temperature they will not ignite. The cooling is effected by passing the mixed 
gases through a tube filled with copper gauze, or packed with fragments of 
copper wire The explosion cannot travel back through such a tube, for the 
flame is extinguished owing to its giving up its heat to the copper, which is a 
good conductor of heat. The danger of explosion can be thus avoided. An 
almost equally hot flame, however, may be produced without danger by urging 
oxygen under pressure through a flame of hydrogen gas by a blowpipe of 
the form bhown m fig. 30. The temperature of such a flame is estimated at 

FIG. 30. 

2200 to 2400. The very infusible metal, platinum, can be melted, and even 
boiled when thus heated; silica can be melted and drawn into threads like 
glass ; and the stem of a pipe, which is composed of aluminium silicate, can be 
softened and bent. With such a flame the hardest glass (combustion glass) can 
be worked as easily as ordinary glass; and when directed on a piece of lime or 
of zirconium oxide, a dazzling light is emitted, the solid being raised to the tem- 
perature of brilliant incandescence. Coal-gas, which contains about 50 per 
cent, of hydrogen, is usually substituted for hydrogen in such experiments 
the temperature, though not quite so high, is still high enough for practical 
purposes. The applications of such a blowpipe are the fusion of platinum and 
iridium, and the production of the lime-light, or, as it is named from its dis- 
coverer, Captain Drummond, the " Drummond" light (Fig. 30). The crucible 
shown in Fig. 31 is made of lime, which is almost the only material capable of 
withstanding such a high temperature without softening. In it such metals as 
platinum, iridium, &c., can be melted. a 

Hydrogen dioxide, H 2 2 , is also formed in small amount 
when water is evaporated; it exists in very minute quantity in 
all natural waters, and is apparently produced by the action of 
heat and light on water containing oxygen in solution. 

Hydrogen sulphide. Hydrogen burns in sulphur vapour, 


with formation of monosulphide, H Z S. This may be shown by 
boiling sulphur in a flask, and introducing a jet of burning hydro- 
gen into the vapour ; the hydrogen continues to burn feebly in the 
sulphur gas. Selenium and tellurium also unite directly with 
hydrogen at about 500. 

2. By replacement. (a.) Action of hydrogen on an oxide. 
This process has already been alluded to as a means of obtaining 
the elements indium, iron, germanium, tin, and lead, nitrogen, 
arsenic, antimony, and bismuth, tungsten, the metals of the plati- 
num group, and copper, silver, mercury, and gold The method 
consists in heating the solid oxide to redness in a tube through 
which a current of hydrogen is passing, when the hydrogen unites 
with the oxygen of the oxide, forming water, and the reduced 
element is left. In some cases higher oxides, such as manga.riese 
dioxide, chromium trioxide, &c., are reduced not to the state of 
element, but only to lower oxides. Similar experiments on sulph- 
ides, selenides, and tellurides have not been thus carried out, but 
would doubtless prove efficient in many cases. 

(&.) Action of oxygen, sulphur, <fcc., on a compound of 
hydrogen. All compounds of hydrogen, excepting hydrogen 
fluoride, are thus decomposed by oxygen. This is the principle of 
Deacon's chlorine process (p. 74), and of the manufacture of 
lampblack (p. 45) ; while a useful method of preparing hydrogen 
sulphide consists in heating a mixture of paraffin wax (a mixture 
of compounds of carbon and hydrogen) with sulphur. The sul- 
phur replaces some of the hydrogen, which combines with excess 
of sulphur to form hydrogen sulphide. Similarly, by heating 
selenium with colophene, hydrogen selenide is continuously 

3. By double decomposition. Water is produced by in- 
numerable interactions of this kind. For example, when many 
oxides, hydroxides, carbonates, silicates, &c., are treated with 
halogen acids, halides are formed together with water. This is 
also the usual and only available method of manufacturing hydro- 
gen dioxide.* For this purpose barium dioxide is dissolved in 
dilute hydrochloric acid until the acid is nearly neutralised. 
Dilute^ baryta- water is then added to bhe filtered and cooled 
solution in order to precipitate foreign oxides and silica, which are 
often present as impurities in commercial barium dioxide. The 
solution, again filtered, is again treated with a strong solution of 
barium hydroxide, which throws down a precipitate of hydrated 

* Th^nard, Annales (2), 0, 441; 10, 114, and 335; 11, 208. BericUe, 7, 
73 j Annalen, 192, 257. 



barium peroxide. This precipitate is filtered and washed until free 
from hydrogen chloride. It is then added to dilute sulphuric acid 
(I part H a S0 4 to 5 parts H 2 0) with constant stirring, until the^acid 
is nearly neutralised. The precipitated barium sulphate, which is 
practically insoluble in water, is then removed by nitration, and 
the small trace of sulphuric acid remaining is precipitated by 
careful addition of dilute baryta-water. The slight precipitate is 
allowed to settle, and the clear liquid decanted and evaporated in 
a vacuum over strong sulphuric acid. The equations are as fol- 
lows : 

BaO 2 + 2HCl.Aq = BaCl 2 .Aq + H 2 2 .Aq t ; 
Ba(OH) 2 .Aq + H 2 2 .Aq = BaO 2 .8H 2 O + Aq ; 
BaO..8H 2 O 4- H,SO 4 .Aq = BaSO 4 4- H 2 O 2 .Aq. 

Hydrogen sulphide, selenide, and telluride are also usually 
prepared by double decomposition. Sulphide of iron, FeS, is 
treated with dilute sulphuric acid; or sulphide of antimony, Sb^S 3 , 
or stlenide of zinc or telluride of magnesium, ZnSe ov MgTe, 
are treated with hydrochloric acid, thus : 

PeS + H 2 S0 4 .Aq = FeS0 4 .Aq + H 2 S ; 
Sb 2 S 3 -f GHCl.Aq = ->SbCl,.Aq 
ZnSe + H 2 S0 4 .Aq = ZnS0 4 .Aq -f 

Hydrogen sulphide, prepared from crude ferrous sulphide con- 
taining metallic iron, obtained by heating together iron and 
sulphur, always contains hydrogen. The pure gas may be pro- 
duced from antimony sulphide. Many other sulphides are simi- 
larly attacked ; among those which resist the action of acids 
(dilute sulphuric or hydrochloric) are the sulphides of tin, lead, 
arsenic, bismuth, platinum, &(\, copper, silver, mercury, and 
gold. Certain sulphides and hydrosulphides are decomposed by 
water alone ; among these are sulphides of magnesium, alumi- 
nium, boron, silicon, phosphorus, chlorine, &c. The heating of a 
solution of magnesium hydrosulphide to 100 causes such a reac- 
tion : Mg(SH) 2 .Aq = Kg(OH), -f H.,S + Aq. This method 
yields pure hydrogen sulphide. The selenide and telluride could 
doubtless be similarly prepared. The gases are best collected by 
downward displacement. 

Hydrogen trisulphide is prepared in an impure state by 
pouring into cold hydrochloric acid a solution of sodium poly- 
sulphide. The resulting yellow oil does nob correspond to the 
formula H 2 Sa, for it contains sulphur in solution. An orange- 


coloured compound with the alkaloid strychnine is, however, 
known, which on treatment with strong sulphuric acid yields 
colourless drops of the trisulphide, H 2 S 3 . No persulphides of 
selenium or tellurium are known. 

Properties. Water is a liquid at ordinary temperatures, 
colourless in thin layers, but blue when a white light is passed 
through a stratum 6 feet long contained in a blackened tube. Ice, 
when seen in thick masses, has also a bluish-green colour. The 
vapour of water also appears to be blue. Hydrogen sulphide, 
selenide, and telluride are colourless gases ; the first has been 
condensed to a clear liquid, and frozen to a colourless solid. 
Water, when^ure, possesses no smell or taste; hydrogen sulphide 
has the smell of rotten eggs, being produced by the decomposition 
of the albumen of eggs, which contains sulphur; the odour of 
hydrogen selenide and telluride is not so offensive as that of the 
sulphide, but they produce exceedingly disagreeable nervous effects. 
The sulphide, selenide, and telluride are exceedingly poisonous ; 
when breathed undiluted with air, instant insensibility is produced. 
Hydrogen dioxide is a colourless viscid liquid, miscible in all pro- 
portions with water. It has a faint pungent, smell, and a sharp 
metallic taste. Hydrogen trisulphide has a pungent smell, and is 
insoluble in water. 

These compounds are of very different degrees of stability. 
While water decomposes only at a very high temperature that of 
melted platinum, for example into its elements, hydrogen sulphide 
is resolved into hydrogen and sulphur at a low red heat, and 
hydrogen selenide and telluride slowly decompose at the ordinary 

The dissociation of water may be shown by passing steam through a tube 
containing a spiral of platinum wire heated to whiteness by an electric current. 
The hydrogen and oxygen produced by the dissociation mix with the steam, 
and are cooled below the temperature of ignition ; and a test-tube full of explo- 
sive gas may thus easily be collected. The dissociation of sulphuretted hy- 
drogen may be shown by passing the gas through a red-hot glass tube, when 
sulphur deposits on the cool part of the tube. 

Hydrogen dioxide and hydrogen trisulphide are very unstable 
bodies^ The former, even at 18 or 20, begins to decompose 
into water and oxygen. It thus dilutes itself, and in dilute 
solution it is mor/e stable. On warming even a very dilute solu- 
tion, however, it decomposes, bubbles of oxygen being evolved. 
Many substances of a porous consistency cause this decomposition 
to take place at the ordinary temperature ; and it reacts with 
certain oxides and peroxides, depriving them of oxygen, while it 


also loses oxygen. Silver oxide, manganese dioxide, and potas- 
sium permanganate have an action of this nature. With 
silver oxide, for example, the action is shown by the equation 
Ag 2 O + H 2 2 .Aq = 2Ag + H 2 O.Aq + 2 . The tendency of the 
oxygen of the silver oxide to combine with one atom of the oxygen 
of the dioxide so as to form a molecule of oxygen, 2 , causes the 
change to take place. Hydrogen dioxide cannot be vaporised 
appreciably without decomposition, but the fact of its possessing 
a smell points to its being able to exist for some time as gas.* 

Hydrogen trisulphidef when heated, at once splits up into 
sulphur and hydrogen sulphide. This decomposition occurs spon- 
taneously when hydrogen trisulphide is kept in a sealed tube, and 
pressure rises until the resulting hydrogen sulphide is liquefied, 
solid sulpltnr separating out. 

Many instances have already been given of the decomposition 
of water by elements. Some, such as sodium and calcium, decom- 
pose it at the ordinary temperature ; others, such as magne- 
sium, iron, copper, carbon, phosphorus, &c., act on it at a high 
temperature. In all such cases hydrogen is evolved, while the 
element combines with the oxygen ; the resulting oxide often itself 
combines with the excess of water, forming a hydroxide or an acid. 
Sulphuretted, seleniuretted, and telluretted hydrogen are similarly 
decomposed, yielding sulphides, selenides, and tellurides of the 
elements, with evolution of hydrogen. But when fluorine or chlo- 
rine acts on water, oxygen is evolved. 

Hydrogen sulphide, selenide, and telluride are soluble in water, 
but their solutions soon decompose on exposure to air. A solution 
of the first is largely employed as a reagent in qualitative and 
quantitative analysis. 

The presence of water can be detected and estimated by heat- 
ing the substance containing it in a current of dry air, and leading 
the current through a weighed tube containing dry calcium chlor- 
ide, phosphorus pentoxide, or strong sulphuric acid, all of which 
bodies are hygroscopic. The amount of water preseLt is deter- 
mined by weighing the absorbing tube a second time. Hydrogen 
dioxide may be detected}; by adding to the liquid containing it a 
little ether, and one drop of a solution of potassium bichromate ; 
on shaking, the ether is tinged blue, if dioxide be present, by a 
compound of chromium of the formula Cr0 3 .H 2 O2, produced by 
the union of the hydrogen dioxide with the chromium trioxide, 

* Comptes rend., 100, 57. 

t Comptes rendus, 60, 1095 ; Chem. Soc., 27, 857. 

J Annales (3), 20, 364. 


Cr0 3 , of the bichromate. Another very delicate test is freshly pre- 
pared titanium hydroxide, with which the peroxide gives a yellow 
colour. Hydrogen sulphide is recognised by its smell and its 
blackening a piece of paper soaked in a solution of lead acetate ; 
black sulphide of lead is formed. 

Physical properties of water. As water, owing to its abun- 
dance, and the ease with which it can be purified, serves as the 
standard substance for many physical constants, a somewhat detailed 
description of its physical properties is necessary. 

(a.) Mass of 1 cubic centimetre. The mass of 1 cubic 
centimetre of water at 4 is accepted as the unit of weight, 1 gram. 
Ice is specifically lighter than water. 1 cubic centimetre of ice at 
weighs 0'917 gram ; hence ice floats in water with about 9/10 ths 
of its bulk submerged. 

(&.) Expansion. Water, unlike other liquids, has a point of 
maximum density at 4; when cooled below that temperature, or 
warmed above it, it expands. It is possible to cool water a few 
degrees below without its freezing ; it continues to expand on 
fall of temperature, instead of contracting as all other known sub- 
stances do. 

(c.) Vapour-pressures. At 100 Centigrade, 80 Reaumur, or 
212 Fahrenheit, water- vapour exerts a pressure equal to that of 
760 millimetres of mercury ; it is then at its boiling-point under 
normal atmospheric pressure. With decrease of temperature its 
vapour-pressure decreases, and at its vapour-pressure is equal to 
that of 4'6 millimetres of mercury. When pressure is reduced by 
pumping out air, its temperature falls, that portion of water 
which evaporates withdrawing heat from the remainder, until at 
a pressure of 4*6 millimetres its temperature is O c - On still 
further reducing pressure, its temperature falls still lower, but it 
is difficult to prevent freezing. It is, however, possible to lower 
temperature to 5 or 7 without freezing. Ice has also a 
vapour-pressure. At it is equal to that of water at the same 
temperature, viz., 4*6 millimetres ; on reducing the pressure still 
further, the temperature of the ice falls by evaporation, exactly as 
with water, owing to its cooling itself by evolving vapour ; if heat 
be communicated to the ice, it does not raise the temperature 
of the ice, provided the pressure does not rise, but is entirely 
expended in evaporating the ice, which passes directly from the 
state of solid to that of vapour. The vapour-pressures of water 
are as follows : 

T. 0. 10. 20. 30. 40. 60. 60. 70. 
P. mm. .. 4-60 9'16 17 '40 31'55 64*91 91'98 148-79^23309 



T. 80. 90. 100. 110. 120. 130. 140. 
mm... 354-64 525 '45 760 '0 1075'4 1484 2019 2694 




mm. .. . 

. . 4652 



190. 200. 
9403 11625 










(d.) Specific heat. The amount of heat required to raise the 
temperature of 1 gram of water through 1 is termed a calory. 
But the specific heat of water, like that of other substances, is not 
a constant; hence the hundredth part of the heat required to raise 
the temperature of a gram of water from to 100 is generally 
accepted as the value of a calory. This amount* is practically 
coincident with the amount required to raise the temperature of 
1 gram from 18 to 19. A unit of 100 calories is employed in 
this book under the symbol K. It is better adapted to express 
large amounts of heat, such as are evolved or absorbed during 
chemical reactions. The specific heat of ice between 78 and 
is 0*474 calory per degree ; that of water-gas at constant volume is 
0-4805 calory. 

(e.) Heat of fusion of ice. To melt 1 gram of ice, 80 calories 
are absorbed ; hence to melt 18 grams (or 1 gram-molecule) of ice 
requires 14'4 K at atmospheric pressure. 

(/.) Heat of evaporation of water. To evaporate 1 gram of 
water at 100 into steam of that temperature requires an absorp- 
tion of 537 calories; hence to evaporate ]8 grams, or 1 gram- 
molecule requires (537 X 18)/100 = 96'66 K. To convert 1 gram 
of water at into steam at t requires an absorption of heat of 
(606-5 + 0-3050 calories. 

(0.) Volumes of saturated steam. From direct measure- 
ments the following numbers have been obtained : 

Temperature 140. 150. 160. 170. 180. 190. 

Vol. of 1 gram ; c.c.. . 506 -0 392 '4 307 "9 246 '4 197 '1 160 '9 

Temperature 200. 210. 220. 230. 240. 250. 

Vol. of 1 gram j c.c.. . 129'8 108 '7 89 '2 73 '8 62'1 52-1 

Physical properties of water, hydrogen sulphide, hydrogen selenide, 

and hydrogen telluride. 

Mass of 1 c.c. Melting-point. 

H 2 0. H 2 8. H 2 Se. H 2 Te. H 2 O. H 2 S. H 2 Se. IL,Te. 

Solid.... 0-917 ? ? ? -85 ? ? 

at at 760 mm. 

Liquid.. TOO 1-19 ? ? 
at 4 at ? 


H 2 O. H 2 S. H 2 Se. H 2 Te. 

Liquid 100 ? ? ? 


Heats of combination : 

2H + = H.O + 68iK; -f + Aq = H 2 2 .Aq -231K. 
2H + S = H,S + 47K; H 2 S + Aq = H 2 S.Aq + 46K. 
2# + Se = JBT 8 flfe - 111K. 

Proofs of the composition of the oxide, sulphide, selen- 
ide, and telluride of hydrogen. We have seen that two 
volumes of hydrogen and one volume of oxygen unite to form two 
volumes of wtiter-gas. An experiment has also been described on 
p. 62, whereby it is shown that when water is electrolysed, it 
decomposes into two volumes of hydrogen and one volume of 
oxygen approximately. From Avogadro's law it may therefore be 
concluded that the reaction occurs between 2 molecules of hydro- 
gen and 1 molecule of oxygen, 2 molecules of water-gas being 
formed, thus : 

2H Z -f 2 = 2H,0, 

or in gram-molecules, 4 grams of hydrogen, occupying 11'16 X 4 
= 44 64 litres, unite with 32 grams of oxygen, occupying 
11 '16 x 2 = 22'32 litres, to form 44'u'4 litres of water-gas weigh- 
ing 36 grams. Hence, as the weight of 11*16 litres of hydrogen is 
1 gram, water-gas under similar conditions of pressure and tem- 
perature weighs 36/4 = 9 times as much as hydrogen. Its 
molecular weight is therefore 18 ; that is, a molecule of water-gas 
weighs 18 times as much as an atom of hydrogen. 

Similarly the weight of 22 32 litres of hydrogen sulphide is 
34 grams, and its specific gravity 17 ; arid the specific gravities of 
hydrogenselenide and telluride have been found equal to 40*5 and 
64*3 respectively, giving molecular weights of 81 and 128*6. 

The fact that hydrogen sulpliide contains approximately its own volume of 
hydrogen may be shown by heating in a tube, by means of a spiral of platinum 
wire traversed by a current, a known volume of hydrogen sulphide. The gas is 
decomposed into h j drogen and sulphur, and on opening the tube under water 
no contraction takes place. 

The exact quantitative composition of water has been the 
subject of numerous researches, and is even now by no means certain. 
The processes for ascertainirg the composition may be grouped in 
two divisions : (1) Determination of the relative weights of 
oxygen and hydrogen gases, and of the exact proportions 


by volume in which they combine; and (2), Synthesis of 
water by passing a known weight of hydrogen over a 
weighed quantity of red-hot copper oxide, CuO, and f /esti- 
mating its loss of weight, the weight of the water produced 
being also determined. 

1. By the second method, Erdmann and Marchand,* in 1842, 
established the ratio between the weights of hydrogen and oxygen 
in water as 2 : 16. 

2. In the same year, Dumas also obtained the ratio 2 : 16, 
and therefore the ratio between the atomic weights of hydrogen 
and oxygen of 1 : 16.f 

3. Stas, in 1867, determined the ratio between the atomic 
weight of silver, and the molecular weights of ammonium chloride 
and bromide, by precipitating the chlorine and bromine contained 
in weighed quantities of these compounds by silver nitrate pro- 

< duced from pure silver. As he had previously determined the 
ratios of the atomic weights of silver, chlorine, bromine, and 
nitrogen to oxygen (these numbers are given on p. 23), the 
ratio of hydrogen to oxygen could be calculated. He found 
H : O : : 1 : 15-885.J 

4. Regnault, in 1847, found the relative densities of hydrogen 
and oxygen 1 : 15*964. Applying a correction overlooked by him, 
but necessary on account of the decrease of the volume of the 
vacuous globe, owing to the external pressure of the atmosphere, 
the ratio is reduced to ]5'939. 

5. Scott, in 1887~8,|| redetermined the ratios between the 
volumes of hydrogen and oxygen combining with one another, and 
found it to be = 1, H = 1*994; applyiug this correction to 
Regnault's results, the ratio 1 : 16*01 is obtained. 

6. Van der Plaats, in 1886, found the ratio 1 : 15*05 by oxi- 
dising a known volume of hydrogen. 

7. Lord Eayleigh, in 1888 and 1889,^ found the ratio 1 : 15*89, 
from the relative weights of the gases. 

8. Cooke and Richards, in 1888,** by weighing the water 

* J. pr. Chem., 26, 461. 

f Annales (3), 8, 189. 

J Recherches sur les Rapports reciproques des Poids atomiques, Brussels, 
1860. , 

Relations des Experiences, Paris, 1847, 151. 

11 Proc. Roy. Soc., 42, 396 ; Brit. Assn. Rep., 1888, 631. Scott has since 
found the ratio to exceed 1 : 2. 

1 Proc. Roy. Soc., 43, 356. 

** Amur. Chem. Jour., 10, 81. 


produced by the combustion of known weights of hydrogen, ob- 
tained the number 15'869. Lastly, 

Sl'Keiser, in 1888,* weighed hydrogen in combination with 
palladium, and after combining it with oxygen, found the ratio 
1 : 15-949. 

These numbers vary between 15 '869 and 16 '01 ; their difference 
amounts to nearly 1 per cent., and the question cannot be regarded 
as settled. Hence, as remarked on p. 20, seeing that most atomic 
weights have been determined by the analysis of oxides, it is 
advisable to assume as the basis of atomic weight, = 16, leaving 
the exact ratio between hydrogen and oxygen to the tost of further 
experiment. f 

Compounds of water with halides. The compounds of 
water with halides are very numerous. The water thus com- 
bined is generally termed " water of crystallisation," and com- 
pounds containing water are said to be "hydrated." To give a 
complete list of such compounds would occupy too much space. 
In some instances, the amount of water has been stated in the 
formulae given. The same salt may crystallise with several different 
amounts of water; thus, ferric chloride, Fe 2 Cl 6 , forms the hydrates, 
Fe^Cl 6 .10H 2 O and Fe^Cl 6 .5H 2 O ; calcium chloride combines with 
water in the proportions CaCl 2 .6H^O, and 2H 2 O ; and so with 
other halides. It may generally be stated that the lower the 
temperature, the larger the amount of water of crystallisation 
with which the halide will combine. The halides of hydrogen 
also form compounds with water (see p. 112), which are partially 
decomposed at the ordinary temperature ; but when distilled, an 
acid of a definite strength always comes over ; the relative 
amounts of halide and water depend, however, on the pressure. 

Some double halides are unstable, and are not known in a solid 
state unless combined with water. This is particularly the case 
with the double halides of hydrogen with those of other elements. 
The compounds SiF 4 .2HF, PtCl 4 .2HCl, and many others, are 
unknown except in combination with water. Their formulae are 
deduced from those of their salts, i.e., from compounds such as 
SiF 4 .2KF, PtCl 4 .2KCl, &c., which can be dried. Such hydro- 
chlorides appear to be unstable unless for every molecule of 
hydrogen chloride two molecules of water are present. 

This water tends to leave the substance with which it is com- 
bined, evaporating into the air. Its vapour, therefore, exerts a 
definite pressure. If the pressure of the water- vapour in the air 

* JSerichte, 20, 2323. 


be equal to or but little greater than that of the water of crystal- 
lisation, evaporation is balanced by assimilation of water, and no 
change occurs. If, however, it be greater, the compound turn&wet, 
and is said to " deliquesce ; " such substances are termed " hygro- 
scopic ; " if less, the compound loses water, turns opaque and 
lustreless, and is said to " effloresce/' Water of crystallisation is 
usually expelled by heating to 100, but a much higher tem- 
perature is often required. 

Compounds of hydrogen sulphide, selenide, and telluride with 
the halides are unknown. 

Compound of hydrogen sulphide with water, Crystals 
of the compound H a S.7H 2 O are deposited when a saturated solu- 
tion of hydrogen sulphide in water, under a pressure slightly 
higher than that of the atmosphere, is cooled to 0. 




The Oxides, Sulphides, Selenides, and Tellurides. 

Like the halogens, oxygen, sulphur, selenium, and tellurium form 
many double compounds. But (and this is especially true of the 
double oxides) such compounds have been usually placed in a differ- 
ent class, and viewed in a different manner from the doable 
halides. Many of the double halides are decomposed into their 
constituent single halides on treatment with water ; but there is 
no obvious sign of decomposition with most of the double oxides. 
Water, also, is an oxide, and enters into combination with other 
oxides, as, indeed, it does with halides ; but it is often expelled 
only at a high temperature, and, in one or two cases, cannot 
apparently be expelled at any temperature short of that of the 
electric arc, in which the constituent oxide is itself decomposed 
into oxygen and element. But, besides such firmly bound water, 
some oxides crystallise with water, and such " water of crystal- 
lisation " is expelled with more or less readiness at a moderate tem- 
perature, as it is from the double halides also united with water 
of crystallisation. Some double sulphides, selenides, and tellurides 
are also known, but they, unlike the double oxides, are often 
unstable in presence of water, tending, indeed, to react with the 
water in. which they are dissolved, forming hydrogen sulphide and 
an oxide. The sulphides, moreover, do not, as a rule, form stable 
compounds with hyclrogen sulphide, and the few compounds which 
exist have been little investigated. 

Classification of oxides. The oxides of the commoner 
elements have long been divided into two classes ; those of the one 
class chiefly consist of the oxides of elements of low atomic weight, 


with some marked exceptions, and have been termed acids or acid- 
forming oxides; elements forming such oxides are generally termed 
non-metals ; while those of the other class which yield compounds 
with acid oxides have been termed bases, or basic oxides. Examples 
of the first class are : B 2 O 3 , boron oxide ; SiO 2 , silica, or silicon 
dioxide; 00 2 , carbon dioxide; $0 2 and SO 3 , sulphur di- and tri- 
oxides ; and of the second, Na^O, sodium oxide ; CaO, calcium 
oxide ; A1 2 O 3 , aluminium oxide ; Pe 2 O 3 , iron sesquioxide, &c. In 
certain cases, an oxide may belong to both of these classes, as, for 
example, A1 2 O 3 , which combines with basic oxides, on the one 
hand, to form compounds such as A1 2 O 3 .K 2 O, or KA1O 2 ; and, on 
the other, with acid oxides, such as SO 3 , to form such^ompounds as 
A1 2 O 3 .3SO 3 , or A1 2 3SO 4 . And with some elements, which combine 
with oxygen in seveial proportions, basic properties are displayed 
by those oxides containing least oxygen, as, for example, Cr 2 O 3 ; 
while the higher oxides show acid properties, for instance, CrO 3 . 

The Dualistic Theory. Snch properties led Lavoisier to 
assign the nomenclature to bodies which he did, and suggested to 
Davy* the theory of "dualism," as it was subsequently termed by 
Berzelius, its great expositor.f Inasmuch as an oxide, decomposed 
by the electric current, yields up its oxygen at the positive pole, and 
the other constituent element at the negative pole of the battery, 
Berzelius supposed that the atoms of oxygen were negatively, and 
the atoms of the element with which it is in combination, positively 
electrified. When combinations of such oxides are electrolysed, it 
was supposed by Berzelius that they also decompose in like manner, 
the electro-negative constituent of the double oxide being attracted 
to the positive pole, and the electro-positive constituent to the 
negative pole of the battery. Thus, as examples, the oxides 

+ -4- + + 

FeO, BaO, S0 3 , CO 2 , 

were supposed to be constituted of electro-positive and electro- 
negative atoms respectively, while the compounds 

+ + 

BaO.S0 3 , and FeO.C0 2 , 

were likewise imagined to consist of groups of atoms, .which, 
taken as a whole, themselves displayed positive or negative electri- 
fication. On these grounds, he explained the dualistic theory,, 
namely, that every chemical compound is composed of two con- 

* Phil. Trans., 1807, 1. 

f Schwaigger's Jour., 6, 119. 


stituents, one electro-negative and one electro-positive, in combin- 
ation with each other. 

J3ut among the reasons which led to the abandonment of this 
view, two are of special importance. First, many compounds 
exist, especially of the element carbon, which cannot be repre- 
sented on the dualistic system.* Such compounds are, for example, 
CCl 3 Br, C 2 H 5 C1, and numerous others, the molecular weights of 
which are established by their vapour-densities ; hence they have 
not such formulae as 3CCl 4 .CBr 4 , 5C 2 H 6 .C2Cl b , &c. Second, on 
electrolysis of solutions of compounds, such as Na 2 SO 4 , or 
Na 2 O.S0 3 , the basic oxide does not accumulate at the negative, 
and the acid oxide at the positive pole, but the compound splits 
into the element sodium and the group SO*, neither of which are 
stable in the presence of water, but react with it, sodium com- 
bining with its oxygen and half its hydrogen, liberating the 
other half ; while the group, S0 4 , parts with a fourth of its 
oxygen, remaining as S0 3 . It cannot, therefore, be supposed 
that compounds such as sodium sulphate, Na^O.SOs, really consist 
of two distinct portions Na 2 and S0 3 ; but its molecule exists as 
a complete individual, Na^SOi. In further support of the second 
argument, it has also been adduced that a similar compound, 
PbS0 4 , lead sulphate, may be produced by the following methods : 
union of PbO and S0 3 ; union of Pb0 2 and S0 2 ; and union of 
PbS with 40. 

The first argument is termed the argument from substitu- 
tion ; it was suggested by the French chemist, Dumas, and by the 
Swiss chemists, Laurent and Gerhardt, and its development has 
led to the classification of the compounds of carbon, and to the 
discovery of an enormous number of new bodies. t 

This view of the constitution of chemical compounds has also 
been extended to include compounds other than those of carbon, 
and compounds of which the molecular weight is absolutely un- 
known. Thus, sodium monoxide has a composition most simply 
expressible by the formula Na^O. This oxide unites with water 
with great readiness, producing the compound Na^O.H 2 O. But 
the same compound may be produced by the action of the metal 
sodium on water ; the equation is 

2Na + 2H 2 = 2NaHO + H 2 . 

An atom of so'dium expels and replaces an atom of hydrogen 
from water. The secondary action of the union of two atoms of 

* Dumas, Annales (2), 56, 113 and 140. 

f References are not introduced, as they refer almost exclusively to the 
compounds of carbon. See B. v. Meyer's Qeschichte der Chemie, Leipsig, 1889. 


hydrogen to form a molecule afc once occurs, and ordinary hydrogen 
is evolved. The formula NaHO is identical with the formula 
Na^O.HaO, so far as concerns the expression of the composition of 
the body, for Na 2 O.H 2 = 2NaHO; but, as the further action of 
sodium on fused NaHO is to yield Na0 and hydrogen, thus : 

2JSTaHO + 2Na = Na,O + JBT 2 , 

the reactions are adduced as a proof that water contains two atoms 
of hydrogen, inasmuch as the hydrogen can be replaced by sodium 
in two stages, the series of compounds being 
HA NaHO, Na a O. 

Again, many chlorides on treatment with water exchange their 
chlorine for oxygen. Thus, PC1 5 with a small quantity of water 
forms PC1 3 0, thus : 

PC1 6 + H 2 = PC1 3 + 2HCZ, 

one atom of oxygen taking the place of two atoms of chlorine ; 
and that PC1 3 is really the formula of the compound is proved 
by its vapour- density ; it is not 3PC1 5 .P 2 5? which would express 
the same percentage composition. And so with many other 

Constitutional or rational formulae. The analogy between 
the halides and the hydroxides, as bodies such as NaOH are 
termed (the word being a contraction for " hydrogen-oxides"), is 
also a close one. Thus we have NaCl, NaOH ; CaCl 2 , Ca(OH) 2 ; 
SiCl*, Si(OH) 4 , and so on ; and although no hydroxide is volatile 
enough at high temperatures, or indeed, as a rule, stable enough to 
make it possible to determine its molecular weight by means of its 
vapour- density, the analogy is an instructive one. The molecule 
of chlorine, moreover, Cl Zj finds its analogue in hydrogen peroxide, 
or dihydroxyl, (OH) 2 . 

The action of halides of hydrogen on the hydroxides can also 
be well represented on the scheme of replacement. Thus we have 
NaOH + HCl = NaCl + H.OH ; sodium hydroxide being con- 
verted into sodium chloride, while hydrogen chloride is changed 
to hydrogen hydroxide or water; and so with Ca(OH) a -f 2HCI 
= CaCl 3 + 2H.OH. 

An example of the reverse action, viz., replacement of chlorine 
by hydroxyl, is given in the action of water on phosphorus trichlo-< 
ride, PCI 3 , thus : 

Cl H.OH fOH H.Cl 

Cl + HOH = P< OH -f H.Cl. 
l H.OH [OH H.OI 


(See, however, p. 375). Certain oxy chlorides of known molecular 
weight undergo similar changes, for instance : 

SO / C1 4- H - OH <50 / OH -i- 

bU HCl + H.OH = b *\OH H0l ; 

arid so with many other examples. Such formulae as those given 
above are termed constitutional or rational formula, in contradis- 
tinction to empirical formulae, such as H 3 P0 3 , H 2 S0 4 , by which the 
percentage composition of the body only is expressed, and not the 
possible functions which it may exhibit. 

The action of such compounds on hydroxides may also be 
similarly represented. Thus, the formation of sodium sulphate by 
the action of sulphuric acid on sodium hydroxide is represented 
empirically : 

H,S0 4 + 2STaHO = NazSO, + 2H 2 O. 
Its rational representation is : 

NaOH _ .ONa , H.OH 


In both instances, however, the exchange of hydrogen for sodium 
and of sodium for hydrogen is obvious. The name " sodoxyl " may 
be given to the group (ON~a), and it may be supposed to exist in 
combination with itself in sodium peroxide (ONa) 2 , or 

Intermediate compounds are also known, such as 

chlorosulphonic acid, half chloride, half hydroxide; and S0 2 

sodium hydrogen sulphate, only half of the hydrogen being ex- 
pelled by sodium (see p, 421). 

This nfethod of representation has evidently great advantages ; 
it permits an insight, if only a limited one, into the constitution 
of such double oxides and chlorides; and it has been almost 
universally adopted, save among certain French chemists. It has 
been founded largely on the behaviour of compounds of carbon, 
the constitution of which is elucidated in a similar manner and 
in a much more extended degree. 

Molecular Compounds. The universal acceptance of this 
^system, however, has not been wholly good. There are many com- 
pounds which cannot be thus classified, and which have conse- 
quently been relegated to the position of so-called " molecular " 
compounds. Such is the case with the double halides described 
in previous chapters. The name " molecular " has been applied to 



all double compounds the formation of which cannot be represented 
by the device of replacement, and it has been attempted to draw a 
distinction between " atomic " compounds, such as the siifiple 
halides, and compounds such as those represented above, and 
" molecular'' compounds. Thus, NaCl, CaCl 2 , FeClj, CC1 4 , PF 6 , 
are regarded as atomic compounds, the halogen being in direct com- 
bination with its neighbour element; and such elements are 
termed monad, dyad, triad, tetrad, or pentad, according as they 
combine with one, two, three, four, or five atoms of halogen. And 
compounds such as S0 2 C1,, SO.(OH) 2 , SO.(ONa),, POC1 3 , PO(OH) 3 , 
<fcc., are also regarded as atomic compounds, inasmuch as they 
fulfil the required condition of replacement. But c6mpounds like 
BF 3 .HF, AlF 3 .3NaF, FeCl 3 .2KCl, and of double oxides with each 
other, such as MgS0 4 .K 2 S0 4 (although the latter compounds may 
often be represented as formed by replacement) have been regarded 
as molecular or addition compounds. The water which often 
accompanies crystalline salts, commonly called water of crystalli- 
sation, has also been regarded as molecularly combined. 

Now it is questionable whether it .is permissible to arbitrarily 
divide compounds into two classes without sufficient reason. And 
there is justice in the view that a uniform system of representation 
should be adopted. Yet, as we know nothing of the true internal 
arrangement of atoms in a molecule, any systems which contribute 
towards classification of like compounds, and representation of 
like changes which they undergo, may be made use of in arrang- 
ing compound bodies. The method of representing compounds 
constitutionally often 1 serves a useful purpose, and likewise thu 
method of representation of compounds as addition-products. There 
is advantage to be gained by representing sodium sulphate as 

S0 2 (OH) 2 , inasmuch as its analogy with S0 2 C1 2 and S0 2 <5L 

is thereby brought out : and there is also advantage in repre- 
senting it as S0 3 .H 2 0, inasmuch as reactions- occur in which the 
group S0 3 remains unaltered, while the group H 2 is affected. 
For example, on distillation with phosphorus pentoxide, the com- 
pound S0 3 is liberated as such, while the water combines with the 
phosphoric oxide. Both systems of representation will therefore 
be employed as occasion offers. 

With these preliminary remarks, which apply mutatis mutandis 
to the sulphides, selenides, and tellurides, we proceed to the con- 
sideration of the compounds of elements of the sodium group. 


Compounds of Oxygen, Sulphur, Selenium, and 
tellurium, with Lithium, Sodium, Potassium, 
Rubidium, Caesium, and Ammonium. 

The following table gives a list of these compounds : 

Lithium ---- 





Oxygen. Sulphur. Selenium. 

Li 2 O 2 P Li 2 S P ? 

Na 2 O; Na,O 2 * Na^S; Na 2 S 2 ; Na 2 S 3 . NaJSe. 
Na 2 S 4 ; Na 2 S 5 .t 

K 2 S; K 2 S 2 ; K 2 Sj. K 2 Se. KL 2 Te p 

K,S 4 ; K 2 S 6 .f 

Bb 2 S? Bb 2 Se? 

Cs 2 SP Cs 2 SeP 

(NH 4 ) 2 S; S 2 ; S 3 ; S 4 ; ? 
S 6 ; andS 7 t 

K 2 O; K 2 O 2 . 
K 2 3 ; K 2 4 . 
Bb 2 O? 
Cs 2 OP 

Bb. 2 Te? 
Cs 2 Te? 


It will be seen that the compounds of potassium, sodium, and 
ammonium alone have been investigated with any degree of com- 

Sources. None of these compounds occurs free in nature ; 
the monoxides of the type M 2 O occur in combination with other 
oxides, especially with C0 3 , SiO 2 , N 2 O 6 , and S0 3 , as carbonates, 
silicates, nitrates, and sulphates. 

Preparation. 1. By direct union. The monoxides are pro- 
duced when thin slices of the metals are exposed to dry oxygen. 
At higher temperatures higher oxides are formed when the metals 
are heated in oxygen or nitrous oxide, N Z ; this process yields 
KiO 2 , Na^O 2 , and higher oxides. The formula of lithium monoxide 
is conjectural; the monoxides of potassium and sodium have been 

A mixture of sulphides is produced on heating the metals 
with sulphur, unless excess of sulphur is used, when the penta- 
sulphides are formed. 

Ammonium monosulphide is produced by the union of am- 
monia and hydrogen sulphide at a temperature not higher than 
-18, thus: 

2NH, + H,S = (NH^S. 

2. By expelling or withdrawing an element from a com- 
pound.Sodium and potassium monoxides have been produced 

* Chem. Soc., 14, 267; 30, 565. 
t Pogg. Ann., 131, 380, 
j J. prakt. Chem., 24, 460. 

P 2 


by heating the hydroxides NaOH and KOH with the metal, 
thus : 

4- 2Na = 2Na,O + H 2 . * 

The higher oxides of potassium are formed on exposing the 
dioxide to moist air ; a portion of the potassium is converted into 
hydroxide, and the remainder stays in combination with oxygen 
as trioxide and tetroxide, thus : 

3K 2 O 2 + 2H,0 = 2KOH -f H 2 + 2KA; 
2K 2 O 2 + 2H 2 = 2KOH + H 2 -f K 2 O 4 . 

The hydrosulphides on exposure to air yield thepoly sulphides, 
the hydrogen uniting with atmospheric oxygen, thus : 

2KSH -f = K 2 S 2 + H 2 0. 

The sulphates, selenates, and tellurates, when heated, do not 
lose oxygen as the chlorates, bromates, and iodates do, leaving 
sulphide, selenide, or telluride as the halogen- compounds leave 
halide ; but if hydrogen or carbon is present, oxygen is lost at a 
red heat, thus 

Na,SO 4 -f 4ff 2 = 4H 2 4- Na 2 S; 
or Na 2 SO 4 -f 40 = 40O + Na,S. 

The action of heat on ammonium pentasulphide, (NH 4 ) 2 S 5 , 
yields ammonium mono- and heptasulphides, thus : 3(NH 4 ) 2 S 6 = 
2(NH 4 ) 2 S 7 + (NH^S. The sulphide being unstable at tempera- 
tures above 18, decomposes into hydrosulphide and ammonia, 
thus : 

S = NH> -f NH 4 SH. 

3. By double decomposition. Hydrogen sulphide passed 
over fused sodium chloride produces monosulphide, H^S -f 2NaCl 
= Na 2 8 -f 2HCI. The sulphides of potassium have also been 
produced by double decomposition ; the trisulphide by exposing 
red hot potassium carbonate to the vapour of carbon disulphide, 
thus : 

2K.CO3 -f 3C& = 2K 2 S 3 -f 4CO + CO,. 

And the tetrasulphide by similar treatment of the sulphate : 

200 -f S0 2 (i>). 

The existence of this compound is doubtful. 

By distillation of ammonium chloride with a sulphide of potas- 
sium, the corresponding ammonium sulphide is produced, e.g., 


K 2 S 2 + 2NH 4 C1 = 2KC1 + (NH,\S Z . In this manner (NH 4 ) 2 S 2 , 
(NH,),S 3 , (NH 4 )>Si, and (NH 4 ) 2 S 5 have been prepared. 

v Liver of sulphur " or " hepar sulphuris," a substance which 
has been long known, is produced by fusing 4 gram molecules of 
potassium carbonate with 10 gram atoms of sulphur, thus : 

4K 2 CO a + 108 = K 2 SO 4 + 3K 2 S 3 + 4(70,. 

It is a mixture of sulphate and trisulphide. 

Properties. The monoxides, so far as they have been pro- 
pared, are white or grey solids. Lithium monoxide is said to be 
non-volatile at a white heat ; the others melt with difficultly and 
volatilise at a^ very high temperature. Ammonium monoxide is 
incapable of existence, decomposing at once into ammonia and 

Sodium dioxide is a white, and potassium dioxide a brownibh- 
yellow solid. Potassium trioxide is lemon-yellow, and the tetroxide 
sulphur-yellow ; both fuse to orange-red liquids, turning black with 
rise of temperature, but returning to yellow on solidification. 

The sulphides of potassium, sodium, and ammonium are all 
yellow or brownish-yellow solids which have a peculiar " hepatic " 
smell. Ammonium heptasulphido is a deep-red substance, vola- 
tilising without dissociation at 300. With acids, the polysulphides 
give off hydrogen sulphide, while sulphur separates as a white 
emulsion (milk of sulphur). 

Potassium selenide is a greyish or brownish mass ; the telluride 
is a brittle substance with metallic lustre. Both are soluble in 
water and deposit selenium or tellurium on exposure to air. 

All these substances are soluble in water, in all probability 
combining with it. The union of the monoxides with water taken 
place with great evolution of heat, and the water cannot be ex- 
pelled on. ignition (see Hydroxides, below). But water may be 
expelled from solutions of sodium dioxide, and of sodium and potas- 
sium monoselenides and disulphides, the anhydrous salts being 
left on evaporation. Hydrated sulphides are known of the formula) 
K 2 S.2H 2 O, K 2 S.5H 2 O, aNajS.QHA Na 2 S.5H 2 O, and sulphide, 
selenide, and telluride of sodium with 9H 2 0. 

Little is known of the physical properties of these substances. The 
following data, however, are approximate : 

Mass of 1 c.c.LijO, 2 '102 at 15 j Na 2 0, 2'805 ; KaO, 2'656j Na a S, 
2-471; K 2 S,2-130. 

Volatility. Li 2 O has not been volatilised ; K 2 O volatilises at a red heat ; 
melts at a red heat and volatilises with difficultly. The sulphides appear 
to be difficultly volatile ; potassium peutasulphide meltd at a red heat. 


Heats of formation 

2Na + - Na 2 + 804K; + H 2 O - 2NaOH + 352K; + Aq = 198K. 

2Na + S = Na 2 S + 870K ; + Aq = Na 2 S.Aq + 150K. 

K^O has not been investigated. 

2K + S = K 2 S + 1012K; + Aq = K 2 S.Aq + 100K. 

None 'of these substances has been gasified ; their molecular 
weights are therefore unknown. 

Double Coonpounds. Double oxides of potassium are known of the 
formula K 2 O 2 .K 2 O ; K 2 O 2 SB^O j and K 2 O 2 .3K 2 O. These are bluish solids 
produced by heating potassium in oxygen or nitrous oxide. They melt to deep 
red liquids. 

Hydroxides, HydrosulpMdes, Hydroselenides, 
and Hydrotellurides. 

These names are given to compounds of the oxides with water, 
or of the sulphides, &c., with hydrogen sulphide, selenide, or 
telluride. 'None of these compounds occurs free in nature. The 
double selenides and tellurides have not been investigated. 

Monoxides and Monosulphides ; Monohydrates and MonMiilphydrates. 

H 2 O. Li 2 O.H 2 O; Na 2 O.H 2 O; K 2 <XH 2 O; Bb 2 O.H 2 O; Cs 2 O.H 2 O, 

H 2 S. Na 2 S.H 2 S. K 2 S.H 2 S. (NH 4 ) 2 S.H 2 S 

folyhydrates and Polysulpliydrate?. 

Na 2 0.5H 2 0. K 2 O.5H 2 O 
Na 2 Q.8H 2 O. 

Na 2 S.5H 2 0. K 2 S.2K20. 

Na 2 S.9H 2 O. K 2 S.5H 2 O. 
. 12H 2 O, 

Hydrated Polyoxides and Poly sulphides. 

Na 2 S 2 .5H 2 0. Na^SE^O. Na 2 S 4 .8H 2 O. 
Na 2 O 2 .8H 2 0. K 2 S 4 .2H 2 0. 

Preparation. 1. By direct addition. All of these substances 
may be thus prepared. As has been remarked, it is still an open 
question whether the formula of sodium hydroxide is NaOH, one 
atom of sodium replacing one atom of hydrogen in water ; or 
Na 2 O.H 2 O, which may be viewed as an additive product. If the* 
second view be chosen, the, analogy with the halides is concealed, 
and the substances should be named hydrates: if the first, the 
compounds with more molecules of water are difficult to classify ; 


and there appears no good reason for preferring one method of 
representation to another. These remarks apply also to the 
sulpjiicjes. Similar compounds of the selenides and tellurides have 
not been investigated. 

The compound Na 2 O.5H^O is prepared by crystallising a solu- 
tion of Na 2 O.H 2 O from alcohol containing 2 per cent, of water ; 
fche similar compound of potassium separates from water; this 
compound, when treated with metallic sodium, gives a liquid alloy 
of potassium and sodium. 

The hydrate, Na,O.8H 2 O, crystallises from water. 

The compound, Na 2 O 2 .8ILO, crystallises from water, and when 
dried over sulphuric acid it Iqses water, and has then the formula ' 
Na,0 2 2H 2 0.* 

The hydrates of fche mono- and poly sulphides are all obtained 
by crystallising them from water. In most cases the water may 
be evaporated by heat, leaving the anhydrous sulphides. 

Ammonium hydrosulphide, NHJ3S, is produced by direct 
addition of ammonia to hydrogen sulphide above 18. 

2. By double decomposition. The hydrates are prepared 
by (a) the action of barium hydroxide on the sulphate, thus : 

Li 2 S0 4 .Aq + Ba(OH) 2 .Aq = 2LiOH.Aq + BaSO 4 ; 
the barium sulphate being insoluble, it may be separated by filtra- 
tion; (/>), the action of calcium hydroxide on the carbonate 

Na 2 C0 3 .Aq + Ca(OH) 2 .Aq = 2NaOH.Aq 4- CaCO 3 ;* 
or (c) by the action of silver hydroxide on the chloride, bromide, 
or iodide 

KOl.Aq + AgOH = KOH.Aq + AgCl. 

The second method (6) has been long made use of in cauticising 
soda or potash, i.e., in converting the carbonate into the hydroxide, 
named caustic soda or caustic potash ; a solution of the carbonate 
is boiled with milk of lime (i.e., calcium hydroxide stirred up 
with water) in an iron, nickel, or silver vessel, for vessels of other 
metals or of glass or china are attacked by the soluble hydroxide. 

Potassium hydrosulphide has been prepared by passing a 
stream of hydrogen sulphide over red-hot potassium hydroxide or 
carbonate, thus : 

KOH + jET 2 S = KSH + JT 2 0; 

K 2 O.C0 2 + 2H 2 S = 2K 2 S.H 2 S 

various other methods of preparing caustic soda and caustic 

potash (NaOH and KOH) have been employed on a manufacturing 

scale. The most important of these, which yields a mixture of 


hydroxide and carbonate, is the Leblanc process. The principle of 
this process is the simultaneous action of calcium oxide and carbon 
on sodium sulphate. The reaction may be conceived to take jiiace 
in two stages, which, however, are not separated in practice : 

Na 2 S0 4 + 20 = Na 2 S + 2(70 2 ; and 
Na 2 S -h CaO = Na 2 O + CaS. 

The product is termed " black-ash." On treatment with lukewarm 
water in tanks, the hydroxide dissolves and the calcium sulphide 
remains insoluble. 

If the mixture were boiled, the hydroxide of sodium would 
react with the calcium sulphide, reversing the second of these 
equations, thus :2NaOH.Aq + CaS = 2NaSH.Aq f Ca(OH) 2 .Aq. 
But the solution is separated from the solid as quickly as possible 
and concentrated by evaporation. During evaporation chloride 
and sulphate of sodium contained as impurities separate out ; they 
are " fished " out with perforated ladles, and hence are termed 
"fished salts;" while the solution is concentrated, freed from 
carbonate by addition of lime, and finally evaporated in hemi- 
spherical iron pofcs till fused caustic soda, NaOH, remains. It is 
then run into iron drums and brought to market. The principle 
of the manufacture of caustic potash is similar. 

Properties. The hydroxides of the metals lithium, sodium, 
potaslium, rubidium, and cessiurn of the general formula MOH 
have been termed " caustic " lithia, soda, &c., owing to their 
corrosive and solvent properties (KQ/W, T burn). They are all 
white soluble solids melting at a red heat and volatilising at a 
white heat. When dissolved in water, great heat is evolved owing 
to combination. When fused, they attat-k glass and porcelain, dis- 
solving the silica of the glass and the silica and alumina of the 
porcelain ; they act on metals, converting them into oxides, with 
exception of nickel, iron, silver, and gold. Csssium hydroxide is 
most, and lithium hydroxide least volatile. 

Sodium and potassium hydroxides usually contain as impuri- 
ties sulphates, carbonates, and chlorides. A partial purification 
may be effected by treatment with absolute alcohol in which the 
hydroxides dissolve, while the salts are insoluble. The clear solu- 
tion is decanted from the undissolved salts, the alcohol is removed 
by distillation, and the residue fused. 

If absolutely pure hydroxides are required, however, they are 
best prepared from the metal^by throwing small pieces into water, 
and subsequently evaporating the solution of hydroxide in a silver 


The hydrosulphides are white crystalline bodies, which fuse to 
black liquids, but turn white again on solidification. They may 
be Obtained in solution by saturating solutions of the hydroxide 
with hydrogen sulphide, thus : 

NaHO.Aq 4- H 2 S = NaHS.Aq + H 2 0. 

The hydroxides volatilise as such when heated ; but bydrosulphides 
lose hydrogen sulphide, and leave the sulphides. 

Ammonium hydrosulphide dissociates into ammonia and hydro- 
gen sulphide at 50, and above, and on cooling, the constituents 
re-unite.* It forms colourless crystals. 

Appendix. Manufacture of sodium and potassium. An indication of the 
method of preparing these metals was given on p. 30. As they are now pre- 
pared from the hydroxides, by a process devised by Mr. Castner,f a short 
sketch of the manufacture is here appended. 

The following reaction takes place at a red heat between carbon and the 
hydroxide :6NaOH -h 2C = 2Na2C0 3 + 3// 2 + 2Na. But if carbon is 
heated with caustic soda, the hydroxide melts, and the carbon, which is lighter 
than the soda, floats to fche surface, and is for the most part unacted on. Hence 
it is necessary to weight the carbon so as to cause it to sink, or else to add some 
substance to prevent the caustic alkali fusing completely, so that the carbon 
may remain mixed with it. The old plan consisted in adding lime j but the 
temperature at which the metal distilled off was rendered so high that the yield 
was small, and the destruction of the wrought-iron tubes used as stills was 
enormous. The new method is to heat a mixture of pitch and finely-divided 
iron (spongy iron) to redness. Compounds of hydrogen and carbon distil off, 
and an intimate mixture of iron and carbon is left in a porous state. This 
mixture is introduced along with caustic soda into cast-iron crucibles provided 
with tight lids, from each of which a tube conv-eys the metallic vapour to the 
condensers, which themselves are tubes about 5 inches in diameter and 3 feet 
long, and which are placed in a eloping position so that the melted metal runs 
down into a small pot through a hole about 20 inches from the nozzle. The 
crucibles are heated by means of gas to about 1000 ; and when the distillation 
is over, in atjout an hour and a quarter, the crucible is lowered in the furnace, 
so as to separate it from the lid which is stationary; it is then withdrawn, emptied, 
recharged while still hot, and replaced. It is next lifted by hydraulic power 
till it again meets its lid, and the operation again commences. The mixture of 
sodium carbonate and spongy iron emptied from the crucible after each distil- 
lation is treated with water, the iron is recharged with carbon, and the sodium 
carbonate is converted by means of lime into caustic soda to be used in a 
subsequent operation. 

The tnetals potassium and rubidium can be similarly prepared j but lithium 
and caesium must be obtained by electrolysis. 

* Engel and Moitessier, Comptes rend., 88, 1353. 
t Chem. News, 54, 218. 







Oxides, Sulphides, Selenides, and Tellurides of 
Beryllium, Calcium, Strontium, and Barium. 

The compounds of beryllium differ from those of the other 
metals ; those of calcium, strontium, and barium strongly resemble 
each other. 

Sources. These compounds are never found free ; but the 
oxides occur in combination with carbon dioxide, silica, and sulphur 
trioxide, as carbonates, silicates, and sulphates. 

List. Oxygen. Sulphur. Selenium. Tellurium. 

Berjllium.. BeO. BeS. (?) BeSe. ? 

Calcium.... CaO; CaO 2 . OaS; CaS 2 ; CaS 5 .* CaSe. ? 

Strontium.. SrO; SrO 2 . SrS; SrS 4 . SrSe. ? 

Barium BaO; BaO 2 . BaS; BaS 3 ; BaS 5 .f BaSe. ? 

Preparation. 1. By direct union. All of these metals 
readily oxidise when exposed to air, and burn when heated in air 
or oxygen, producing monoxides. They would also in c all proba- 
bility combine with sulphur, selenium, and tellurium. 

Barium dioxide is produced when the monoxide is heated to 
450 in a current of pure dry air; the polysulphides of these 
metals are also formed when the hydrosulphides are boiled with 
sulphur, thus: Ca(SH) 2 . Aq + S = CaS 2 .Aq + H 2 S-> and similarly 
with others; also by heating the monosulphides with sulphur. 

2. By heating hydroxides, nitrates, or carbonates. 

These compounds may be viewed as compounds of the oxides^ 
with oxides of hydrogen, nitrogen, carbon, or iodine, thus : 

* Chem. Soc., 47, 478. 
t Ibid., 49, 369. 


CaO.H 2 O; CaO.N 2 O 5 ; CaO.CO 2 . At a red or white heat, the 
water, nitrogen pentoxide (which splits into lower oxides of 
nitrogen . and oxygen), or carbon dioxide, are evolved as gas, 
while the non-volatile oxide of the metal remains. The Joss of 
water takes place readily with beryllium hydroxide ; slowly, 
beginning at 100, or even lower, with calcium hydroxide, and at 
a very high temperature with strontium and barium hydrox- 
ides. The loss of N.,0 8 takes place at a red heat in all cases. 
This method is adopted as the only practical one in preparing 
barium oxide, which is now made on a large scale. To ensure 
thorough expulsion of oxides of nitrogen, the partially decomposed 
oxide is headed in a vacuum.* Beryllium carbonate is 'decom- 
posed at low redness ; calcium carbonate begins to decompose 
below 400 ; and provided the carbon dioxide be removed by a 
current of air or steam, so that recombination cannot take place, the 
decomposition, if sufficient time be given, is complete at that 

The decomposition of calcium carbonate (limestone) by 
heat, termed " lime-burning " is carried out in " lime-kilns," 
towers open above, with a door below, into which alternate layers 
of lime and coal are introduced from above. The coal is set on 
fire, and the " burnt " or " quick " lime is withdrawn below, after 
all carbon dioxide has been expelled, and when cold. Strontium 
and barium oxides may also be produced from their carbonates, but 
at a higher temperature ; it is well to mix them with a little coal, 
which reduces the carbon dioxide to monoxide, so that no recom- 
bination takes place. 

Calcium sulphide is similarly formed by heating calcium hydro- 
sulphide, Ca(SH) 2 = CaS.H 2 S, in a current of hydrogen sul- 
phide. Strontium and barium sulphides could no doubt be 
obtained in an analogous manner. 

The monoxides of calcium, strontium, and barium are also 
obtainable by heating the dioxides to 450 under reduced pressure, 
or to a higher temperature. This process is made use of in pro- 
ducing oxygen on a large scale (see p. 65). 

Calcium dioxide is also said to be produced in small amount 
when the carbonate is heated to low redness. The hydrated 
diox'dCs may be dried by moderate heat. 

3. By double decomposition. Monoxides. Barium mon- 
%xide is prepared by heating together barium sulphide and copper 
or zinc oxide. On treatment with water, barium hydroxide goes 
into solution. 

* Boussingault, Annales (5), 19, 464, 


Sulphides. The hydroxides, when heated in a current of 
hydrogen sulphide yield the monosulphides, thus : 

Ca(OH) 2 4- H Z S = CaS 

4. By removing oxygen from the sulphates, selenates, or 
selenites by heating to redness with carbon or carbon monoxides; 
ttie sulphides or seleuides are left. The sulphides of calcium, 
strontium, and barium are thus prepared. The selenides are 
similarly prepared by heating selenates or selenites to dull redness 
in a current of hydrogen. It is in this way that barium com- 
pounds are produced from the insoluble sulphate, which is mixed 
with bituminous coal and heated to redness. The Sulphide thus 
produced is converted into the chloride by treatment with hydro- 
chloric acid, or into the oxide, by heating with copper or zinc oxide. 
The soluble hydroxide is produced on treatment with water. 

Properties. Monoxides. These are white powders, or hard, 
white, or greyish-white masses. They all unite with water with 
evolution of much heat. Beryllium oxide forms the least stable, 
and barium oxide the most stable compound. Beryllium oxide is 
eaid to volatilise at a high temperature ; calcium oxide melts 
only in the electric arc, while strontium and barium oxides melt 
at a white heat. The oxides are crystalline when prepared by 
heating the nitrates in covered porcelain crucibles. Beryllium 
oxide crystallises from its solution in fused sulphate of beryllium 
and potassium, or in fused boron oxide. 

The dioxides are white substances, which evolve oxygen when 
heated, calcium dioxide most readily, barium dioxide at a bright- 
red heat ; barium dioxide is said to fuse before evolving oxygen (?). 
They dissolve in water with moderate ease, forming compounds. 

The monosulphides are white amorphous powders, very 
sparingly soluble in water, but reacting with it (see below). 

The monoselenides are also white, sparingly soluble powders, 
which turn red on exposure to air, owing to the expulsion of 
selenium by oxygen; the monosulphides turn yellow, owing to 
the formation of polysulphides. The tellurides have not been 

The polysulphides are yellow solids, soluble in ^water. 
Barium monosulphide, when heated in a current of steam, decom- 
poses it, hydrogen being evolved, and barium sulphate remaining^ 
The impure monosulphides, produced by heating the powdered 
carbonates with sulphur, or the sulphate with carbon, possess the 
curious property of remaining luminous in the dark, after having 
been exposed to light. Such substances used to be 


phori. The calcium compound used to be known as " Canton's 
phosphor us, " and the barium compound as " Bolognian phos- 
phorus." The modern u luminous paint " owes its property to this 

All these oxides are converted into chlorides when heated in a 
current of chlorine. 

Uses. Calcium oxide (lime) when heated to whiteness in the 
oxy-hydrogen flame evolves a brilliant light (Drummond's light) ; 
barium oxide and dioxide are employed in the commercial manu- 
facture of oxygen. 

Physical Properties. The melting and boiling-points of these bodies are 
unknown. * 

Mass of one cubic centimetre 

BeO. CaO. SrO. BaO. BaO 2 . 

318 at 14 3-25 4'75 5'72 496 

Heats of formation: 

Oa + O = CaO -f 1310K ; + H 2 O = 155K ; + Aq = 30K. 
Sr + = SrO + 1284K ; + H 2 O = 177K; + Aq = 116K. 
Ba + O = BaO + 1242 (?)K ; V H 2 O = 233K; + Aq = 122K. 
BaO + O = BaO 2 + 172K; + H 2 O 2 = 102K. 
Ca + Q = CaS + 869K. 
Sr + S = SrS + 974K. 
Ba + S = BaS + 983 (?)K. 

Double Compounds. 

(a.) With water, &C. The following bodies are known : 

Dioxide with 
Dioxides hydrogen 

Oxides with water. with water. dioxide. 

Beryllium.. *3BeO.10H 2 O. *2BeO.3H 2 O. 

*BeO.4H 2 O. BeO.H 2 O. 

Calcium..* CaO.H 2 O. = Ca(HO) 2 . 
Strontium,. SrO.H 2 O = Sr(OH) 2 . 

Sr0.9H 2 = Sr(OH) 2 .8H 2 0. 
Barium. . . . BaO.EUO = Ba(OH) 2 . 

BaO.9H 2 O = Ba(OH) 2 .8H 2 O. 

CaO 2 .8H 2 O. 

SrO 2 .8H 2 O. 

BaO 2 .8H 2 O. BaO 2 .H 2 2 . 

Beryllium . 
CalciuA . . 

Barium .... 

Sulphides with water. 
BeS (?)H 2 O. 
CaS.H 2 - Ca(SH)(OH). 
CaS.4H 2 - Oa(SH)(OH).3H 2 0. 
SrS.H 2 = Sr(SH)(OH). ? 
BaS.HoO - Ba(SH)(OH)P 

Sulphides with 
hydrogen sulphide. 
CaS.H 2 S = Ca(SH) 2 .f 

SrS.H 2 S = Sr(SH) 2 ? 
BaS.H 2 S Ba(SH) 2 ? 

* The existence of these compounds is doubtful, 
t Chem. Soc., 45, 271 and 696. 


Sulphide with water and hydrogen sulphide 

Calcium CaS.H 2 S.6H;O = Ca(SH) 2 .6H 2 O. 

Hydrated poly sulphides. Ca 2 S 2 .3H 2 O ; SrS 4 .6H 2 O, and others. 

Preparation. Hydrated oxides, and hydroxides. 1. By 
direct addition. All of these oxides unite with water directly ; 
beryllium oxide shows least tendency ; calcium oxide unites with 
great evolution of heat ; the water is at first absorbed, and then 
the lumps of lime grow so hot as to evolve clouds of steam, and 
break up into a bulky white powder. This is the familiar opera- 
tion of "slaking lime." The product is termed "slaked lime.'* 
Barium oxide unites with water with so great an evolution of heat 
as to turn red hot when thrown into water. Calcium hydroxide is 
sparingly soluble in water, and the solubility diminishes with rise 
of temperature. At 15, 1 gram of calcium oxide dissolves in 
779 grams of water ; at 20, in 791 grams ; and at 95, in 1650 grams. 
It would thus appear that calcium hydroxide loses water when 
heated even in contact with water, and hence shows no tendency 
towards further hydration. Strontium and barium hydroxides, on 
the other hand, dissolve to some extent in hot water, and on 
cooling, crystals of Sr(OH) 2 .8H 2 O, or Ba(OH) 2 .8H 2 O separate. 
At 15, 1 gram of barium hydroxide dissolves in about 
20 grams of water ; and at 100, in 2 grams. Strontium hydroxide 
is less soluble. Calcium hydroxide, Ca(OH) 2 , separates in crystals 
when its solution is evaporated in vacuo. The hydrated peroxides 
are also formed by dissolving the peroxides in water and crystal- 
lising. The compound Ba,(X.H 2 O 2 , separates from a solution of 
Ba0 2 in H 2 O a containing water. A possible, though improbable, 
view of the constitution of the compound Ca(SH)(OH).3H}O is 
that it consists of CaO.H 3 S.4H 2 O. It is produced ty passing 
sulphuretted hydrogen into a paste of calcium hydroxide and 

2. By double decomposition. -(a.) By addition of a soluble 
hydroxide (e.g., of lithium, sodium, potassium, &c., or ammonia 
and water) to a soluble compound of beryllium, calcium, strontium, 
or barium, thus : 

CaCla.Aq + 2KOH.Aq = Ca(OH) 2 + 2KCl.Aq. 

No doubt this change -always takes place to a greater or less 
extent. But as strontium and barium hydroxides are fairly soluble 
in water, they separate only when the solution is a concentrated 
one.' With beryllium, the hydroxide produced by heating any 


oluble salt, such as the chloride, sulphate, or nitrate, with potas- 
sium hydroxide, thus : BeCL.Aq + 2KOH.Aq = Be(OH) 2 .2H 2 O 
-f 2KCl.Aq, redissolves in excess.of the potassium hydroxide, doubt- 
less producing a soluble double oxide of beryllium and potassium ; 
but the solution of this substance, when boiled, decomposes into 
beryllium hydroxide, Be(OH) 2 , which precipitates, and potassium 
hydroxide, which remains in solution. 

Solutions of strontium and barium hydroxides give precipitates 
with soluble salts of beryllium and calcium, owing to the greater 
insolubility of the hydroxides of the latter metals. 

The hydrated peroxides may be similarly produced by addition 
of some dioxicfe, such as hydrogen or sodium dioxide, to a solution 
of the hydroxide of the metal, thus : 

Ca(OH) 2 .Aq + H 2 2 .Aq = CaO 2 .8H 2 O + Aq. 

As they are sparingly soluble they are precipitated. 

(&.) By the action of hydrogen sulphide on the hydroxides, 
the hydrated sulphides are formed, and in presence of excess of 
hydrogen sulphide the sulphydrated sulphides. With calcium, for 
example, the action is as follows : 

Ca(OH) 2 .Aq + H 2 S = Ca(SH)(OH).Aq + H 2 ; or 
CaO.H 2 O.Aq + H,S = CaS.Aq + H 2 ; 

and further, 

Ca(SH)(OH).Aq + R.,8 = Ca(SH) 3 .Aq -f HA 
If the solutions are strong and cold, the substances 

Ca(SH)(OH)3H 2 O (= CaS.4H 3 O) and Ca(SHV6H 2 O 

(= CaS.HaS.6HaO) 
separate in crystals. 

The calcium compounds are the only ones which have been 
carefully investigated as regards their behaviour with hydrogen 
sulphide ; similar compounds no doubt exist with beryllium, 
strontium and barium, and also with hydrogen selenide and 

The hydrosulphide, Ca(SH) 2 ; when heated with water (as it 
cannoi? be obtained free from the six molecules of water with which 
it crystallises, this water reacts), gives off hydrogen sulphide, and 
the hydroxy-hydrosulphide rjemains, thus : 

Ca(SH) 2 .Aq. + H 2 = Ca(SH)(OH).Aq + H 2 S. 
The hydrosulphide, when treated with sulphur, evolves hydrogen 


sulphide with formation of a polysulphide. Such polysulphides 
fire known only in solution. 

Properties. The hydroxides are white powders ; that of 
beryllium is insoluble in water, but dissolves in a solution of 
ammonium carbonate, and is reprecipitated on boiling., This 
reaction serves to separate it from aluminium hydroxide, which is 
insoluble in aqueous ammonium carbonate. The hydroxide of 
calcium is sparingly soluble in water (see p. 222), that of strontium 
more soluble, and barium hydroxide easily soluble. The hydrates 
of strontium and barium, Sr(OH) 2 .8H 2 O, and Ba(OH) 2 .8H 2 O, are 
white crystalline bodies, rapidly turning opaque on exposure to 
air, owing to absorption of carbon dioxide. When Cheated to 75, 
7 molecules of water are lost, and the eighth only at a red heat. 
From this it would appear that the compound BaO.2H 2 O is not 
much inferior in stability to BaO.H 2 O, and that the formula 
Ba(OH) 2 for the latter does not express any exceptionally stable 
form of combination between water and oxide. 

The hydrated dioxides are crystalline powders, which may be 
dried in vacuo to the dioxides. That of barium, indeed, may be 
heated to over 300, without loss of oxygen. 

The hydrosulphides are very unstable bodies, capable of exist- 
ence only when cooled by ice in presence of hydrogen sulphide. 
When placed in water at the ordinary temperature, hydrogen sul- 
phide is evolved, and the hydroxy-hydrosulphide, 

Ca(SH)(OH).3H 2 0, 
is left. 

There appear to be various compounds of oxides and sulphides 
of these metals (the existence of which, however, requires further 
proof), e.g., 2CaO.Ca,S 2 , 3CaO.CaS 2 , 3CaO.CaS 3 , &c., in com- 
bination with water. 

On boiling solutions of the hydroxides, calcium, strontium, or 
barium, with sulphur, polysulphides are formed, together with 
thiosulphates, thus : 

3Ca(OH) 2 .Aq + 2S -f nS = CaS 2 3 .Aq + 2CaSn/ 2 . 

The polysulphide formed depends on the amount of sulphur pre- 
sent. A deep yellow solution is obtained from which tne thio- 
eulphate separates in crystals. 

(The slaking of lime, the precipitation o'f calcium hydroxide with sodium 
hydroxide, the crystallisation of barium hydroxide from a hot solution ; the 
preparation of calcium sulphide by the action of hydrogen sulphide or calcium 
hydroxide; the formation of polysulphidea of calcium on boiling "milk of 


lime" with sulphur; and the precipitation of "milk of sulphur" on addition 
of sulphuric acid to the orange solution form suitable lecture experiments.) 

(<?.) Double compounds with halides. These are few in number, 
BeCUBeO, is said to be obtained on evaporating an aqueous solution of beryl- 
lium chloride. CaCl 2 .3CaO.15H 2 O is prepared by boiling calcium hydroxide 
in a solution of calcium chloride, and filtering while hot ; BaOl2.BaO.5HjO, 
BaBr.JBaO.5H 2 O, and BaI 2 .BaO.5H 2 O are similarly prepared. 

There appear also to be indications of similar calcium and strontium 
compounds, SrClo.SrO.OH^O having been prepared. 

It is possible to regard these compounds as hydroxy chlorides, 
thus : Ba<QTT.2H 2 O, &c. Although somewhat similar formula 

could be constructed for more complex compounds, as, for example, 
Cl Ca O Ca O Ca O Ca C1.15H 2 O ; yet, inasmuch 
as similar double halides exist in number, which cannot in reason 
be similarly represented, it appears advisable, in the present state 
of our knowledge, to adhere to the simpler and older methods of 
representation. * 

Oxides, Sulphides, Selenides, and Tellurides of 
Magnesium, Zinc, and Cadmium. 

As many of these compounds are unaffected by air and carbon 
dioxide, and do not react or combine with water, they occur 

Sources. Magnesium oxide occurs as periclase ; also, in com- 
bination with water, Mg(OH) 2 , or magnesium hydroxide, as 
brucite, in white rhombohedra. It also occurs in combination with 
carbon dioxide, silicon dioxide, &c. Zinc oxide, ZnO, is named 
zincite or red zinc ore ; it is red owing to its containing ferric oxide 
in small quantity ; it is also found in combination with oxides of 
iron and manganese as franklinite. Zinc sulphide occurs as blende, 
associated with many other sulphides, both in crystalline and in 
sedimentary rocks. It is the chief ore of zinc. It has usually a 
black colour, but is white when pure. Cadmium sulphide, CdS, 
occurs as the rare mineral greenockite. Zinc oxide also occurs in 
combination with carbon dioxide and with silica. 

kist. Oxygen. Sulphur. Selenium. Tellurium. 

Magnesium.. MgrO. M*S. MgSeP BtarTeP 

Zinc ZnO; ZnO 2 P ZnS; ZnS 5 P ZnSe. ZnTeP 

Cadmium.,.. CdO; CdO 2 P CdS. CdSe. CdTe. 

Preparation.!. By direct union. These elements all burn 
in oxygen, or when heated to a high temperature in air. Magne- 
sium burns with .a brilliant white flame, but if the supply of uir is 


limited, the nitride, Mg 3 N 2 , is simultaneously produced. The 
metal is sold in the form of thin ribbon for purposes of signalling, 
photographing dark chambers, &c. ; and in fine dust, for signalling. 
A little powder, when thrown into a flame, gives a brilliant flash 
of light. Zinc burns with a green flame, giving off filmy clouds of 
oxide. Cadmium also burns to a brown oxide. 

The sulphides are also produced by throwing sulphur on to 
the red-hot metals. Zinc and cadmium do not readily combine with 
selenium ; if the metal be fused with selenium, the latter distils 
off, leaving the metal coated with a crust of selenide. But with 
tellurium, tellurides are produced, the boiling-point of that 
element being higher. ? 

2. By heating a compound. The hydroxides, carbonates, 
nitrates, or sulphates of these metals, when heated, leave the oxide. 
The hydroxides and carbonates are decomposed at a low red heat ; 
the nitrates and sulphates require a higher temperature. 

3. By double decomposition. Sulphides of these metals are 
produced by heating the oxides in a current of hydrogen sulphide 
or carbon disulphide, thus : 

MgO + H 2 8 = MgS + JT 2 0; and 
2MgO + CS t = 2MgS + 00* 

Zinc and cadmium selenides have been similarly prepared. 

Inasmuch as the sulphides, selenides, and tellurides of zinc and 
cadmium are insoluble in water, they may be produced by precipi- 
tation, viz., by passing a current of hydrogen sulphide through a 
solution of a soluble salt of the metals ; thus : 

ZnS0 4 .Aq + H Z S = ZnS + H 2 S0 4 .Aq. 

There appear good grounds for believing that this reaction 
gives not a sulphide such as ZnS, but a hydros ulphide, I3nS.wH 2 S. 
The body produced contains more sulphur than corresponds to the 
formula ZnS, and gives off hydrogen sulphide on heating. The 
precipitate produced as above is soluble iu many acids; hence, to 
ensure thorough precipitation, the acid must be neutralised by an 
alkali, e.g., by soda or ammonia. Acetic acid, however, has no 
solvent action ; hence precipitation is complete from a solution of 
zinc acetate. Cadmium sulphide, prepared in a similar manner, is 
also probably a hydrosulphide. It is, unlike zinc sulphide, in- 
soluble in dilute acids ; but dissolves in moderately strong hydro- 
chloric acid. 

Magnesium sulphide cannot be thus prepared ; if the hydr- 
oxide J8 employed the hydrosulphide is produced. 


The selenides and tellurides of zinc and cadmium may be 
similarly obtained. 

Zinc and cadmium peroxides, and probably also magnesium 
peroxide, are formed by addition of hydrogen dioxide to the 
hydroxides. They appear to be compounds of dioxide with mon- 
oxide in proportions as yet unascertained. The pentasulpbide of 
zinc is produced when a zinc salt is treated with a solution of 
potassium pentasulphide. 

Properties. Magnesium and zinc oxides and sulphides are 
white; cadmium oxide brown, and its sulphide yellow. When 
prepared by the union of the metal with oxygen, magnesium oxide 
is dense, anA has the specific gravity 3*6. Magnesia usta, or 
calcined magnesia, is a very loose white powder produced by 
gently glowing the hydroxycarbonate, known as magnesia alba. 
When produced from the native carbonate, magnesite, it is dense 
and hard, and is made use of as a lining for the interior of 
Bessemer converters. It is known as " basic lining." It is very 
sparingly soluble in water, 50,000 parts of water dissolving only 
one parb of oxide ; it probably dissolves as hydroxide. It unites 
slowly with water, when it has not been strongly ignited ; and also 
attracts carbon dioxide from the air, if moist. It is soluble in all 

Zinc oxide is also a soft white powder. When produced by 
burning zinc, it is sometimes named " lana philosopbica," on 
account of its woolly texture. When heated it turns yellow, but 
its white colour returns on cooling. It is insoluble in and does 
not combine directly with water, nor does it unite with carbon 

Cadmium oxide is a soft brown powder. 

None of these bodies are easily volatilised, nor do they melt 

Magnesium sulphide reacts with water, giving hydroxyhydro- 
sulphide (?) or hydroxide and hydrosulphide. It is an amorphous 
pinkish body, infusible, and burning when heated in air to oxide 
and sulphur dioxide. 

Zinc sulphide, as blende, forms compact masses of various 
colour^ due to impurities; it is usually black, and is known to 
miners as " black-jack." It is translucent and crystalline. When 
tf roasted " or heated in air, it changes to oxide and sulphur 
dioxide. Prepared artificially, by precipitation and subsequent 
heating, it forms a white infusible powder. It is employed as a 
pigment under the name of " zinc-white." Its " covering power " 
is not so great as that of white lead (see Carbonates, p. 289), 


but it has the advantage of not turning black on exposure to 
hydrogen sulphide as white lead does, zinc sulphide being white. 
Cadmium sulphide, as greenockite, occurs in yellow transparent 
crystals; prepared by precipitation, it is a yellow powder, and 
is used as an artist's colour, under the name of " cadmium 
yellow," or "jaune brillant" It is not permanent, being easily 
oxidised by moist air. When heated to redness it turns first 
brownish, then carmine-red. It fuses at a white heat, and crys- 
tallises in scales on cooling. 

The oxygen of these oxides is displaced at a red heat by 

The peroxides of zinc, magnesium, and cadmivm are white 
powders. They do not contain enough oxygen to correspond to 
the formulae Mg0 2 , &c., and are either mixtures or compounds of 
higher oxides with the monoxides. 

Zinc pentasulphide is a flesh-coloured precipitate, which, on 
treatment with hydrochloric acid, dissolves with effervescence of 
hydrogen sulphide, sulphur being deposited. 

Zinc selenide, ZnSe, is a yellow amorphous powder, which 
changes into yellow crystals when heated in a current of hydrogen. 
Cadmium selenide forms deep reddish-black crystals. The amorph- 
ous telluride has metallic lustre, but forms a red powder. When 
heated in hydrogen it forms ruby-red crystals; cadmium telluride 
is also a metallic-looking substance giving black crystals. These 
bodies are probably decomposed by hydrogen into the elements, 
which recombine in the cooler part of the tube. It is improbable 
that they are volatile as compounds. 

Physical Properties. 
Mass of one cubic centimetre : 

Magnesium. . 

Oxygen. Sulphur. 
.. 3-636* ? 

S 7R n.t. 1 K 4. -n?i 

Selenium. Tellurium. 

5 -4 at 15 6 '34 at 15 
6 -8 at 15 6 '2 at 15 

+ H 2 O - 50K. 
+ H 2 O - -26K. 
+ H 2 O - - 98K. 


Cadmium .... 8 -11 4 -5 
(crystalline) (precipitated) 
Heats of formation : 
Mgr + O = Mg-O + 1440K j 
Zn + O = ZnO + 853K ; 
Cd + O - OdO + 755K ; 
Mar + S - M*S + 776K. 
Zn -h 8 - ZuS + 396K. 
Cd -h S = OdS + 824K. 

* The density increases on calcination j magnesia produced by igniting 
carbonate has the density 3 '19 at 0. 


Double compounds. (a.) With water: hydrates or hydr- 
oxides. The mineral brucite, MgO.H 2 O, or Mg(OH) 2 , occurs 
native, usually in masses of serpentine. It crystallises in rhombo- 
hedra. Magnesium oxide, when prepared from the nitrate or 
carbonate at a low red heat, unites with water, forming a solid 
translucent substance harder than marble. After being heated to 
whiteness, it loses the property of combination with water. Zinc 
and cadmium oxides do not combine with water directly. 

Soluble salts of magnesium, zinc, and cadmium, on treatment 
with hydroxides of sodium, potassium, or barium give gelatinous 
precipitates of the hydrates. Ammonia in water (equivalent to am- 
monium hydroxide) also produces precipitates, but rcdissolves 
them if added in excess. Magnesium hydroxide does not react with 
excess of sodium or potassium hydroxides, whereas zinc and cad- 
mium hydroxides are soluble in excess of the precipitant, forming 
double compounds (see infra). 

Crystals of ZnO.H 2 O and of CdO.H 2 O are produced after some 
time by placing a stick of zinc or cadmium in aqueous ammonia, in 
contact with iron, lead, or copper. The zinc compound forms 
rhombic prisms, of 2'68 specific gravity. And octahedral crystals 
of ZnO.2H 2 O have been formed by allowing a solution of Zn0 2 K 2 
to stand for some months. The following bodies are thus 
known : 

MgO.H 2 O = Mg(OH) 2 ; ZnO.H 2 O = Zn(OH) 2 ; 
CdO.H 2 O = Cd(OH) 2 ; ZnO.2H 2 O. 

(6.) With hydrogen sulphide. Zinc and cadmium sulphides 
do not appear to combine with hydrogen sulphide. But if a stream 
of that gas is led through water in which magnesium oxide or 
carbonate is suspended, a soluble compound is formed, which has 
not been Obtained solid, but which is supposed to have the formula 
MgS.H 2 S = Mg(SH) 2 , and to be magnesium hydrosulphide. 
When gently warmed, this solution evolves hydrogen sulphide, 
thus: Mg(SH) 2 .Aq + 2H 2 = Mg(OH) 2 .Aq + 2H*S. This 
solution dissolves sulphur with a yellow colour, and may then 
contain poly sulphides of magnesium. 

TJje selenides and tellurides have not been investigated. 

(c.) Compounds of oxides with oxides. White crystals of 
,ZnO.K*O and ZnO.Na.0 [= Zn(OK) 2 , and Zn(ONa) 2 ] 
separate from solutions of zinc hydroxide in caustic alkali. 
Metallic zinc dissolves in boiling caustic potash or soda, with evo- 
lution of hydrogen, thus: Zn + 2NaOH.Aq=Zn(ONa) 3 .Aq-hI/2. 
A similar cadmium compound is formed by dissolving cadmium 


oxide in fused potassium hydroxide. On treating a solution of 
zinc hydroxide in caustic soda with alcohol, the compound 
ZnQ,Na 2 O.8H 2 O is thrown down in crystals. These bodies cor- 
respond to the hydroxides, the hydrogen being wholly or partially 
replaced by sodium or potassium. 

(d.) Compounds of sulphides with sulphides. Zinc 
sulphide is said to be wholly dissolved when added to a solution 
of sodium sulphide containing a weight of sulphur equal to that 
contained in the zinc sulphide. The inference is that the 
compound ZnS.Na 2 S is produced. Cadmium sulphide is also 
sparingly soluble in excess of alkaline sulphides. 

Cadmium sulphide is supposed to polymerise whun boiled with 
acids or with sodium sulphide; and the sulphide produced by 
treating with hydrogen sulphide cadmium hydroxide which has 
been boiled with water is vermilion-coloured. Cadmium sul- 
phide may also be obtained dissolved in water by washing the 
precipitated sulphide thoroughly, and treatment with solution of 
hydrogen sulphide. A yellow solution is produced, which coagu- 
lates on treatment with weak solutions of salts, especially those of 

(e.) Compounds of sulphides with oxides, Magnesium 
oxide heated in a mixture of carbon dioxide and disulphide is 
converted into MgO.MgS. The corresponding zinc compound has 
been prepared by heating zinc sulphate, ZnSO 4 , in hydrogen ; and 
the cadmium compound, CdO.CdS.H 2 O, is thrown down as a red 
precipitate when hydrogen sulphide is passed through a boiling 
solution of a cadmium salt. The compound 4ZnO.ZnS has been 
found in zinc furnaces. 

(/.) Compounds of oxides with halides. The following 
" basic " halides have been prepared by the reaction of water at a 
high temperature on the halides : ' 

; MgrCl 2 .5MgrO ; MgrCL 2 .9MgO ; 
ZnClj.SZnO ; ZnCl s .6ZnO ; ZnCl 2 .9ZnO. 
OdBr 2 .CdO. 

These bodies crystallise with varying amounts of water j thus crystals of 
MgClj.5MgO have been obtained with 17, 14, 8, and 6H 2 0. Zinc oxychlorides 
possess the property of dissolving silk, but not wool or cotton, and their 
solutions are employed as a means of separating the constituents of mixed 
fabrics. The zinc oxychlorides are used by dentists as a stopping for teeth. * 



Physical Properties. 
Mass of 1 cubic centimetre : 



(OH) 2 .. - 

3-16 3-32* 





5'32 5'72* 


S ...... 






2 -683 -05 








The asterisked higher numbers usually refer to the crystallised varieties, 
but are sometimes the results of different experimenters. 

Heats of formation : 
Ca + O = CaO 

Sr + 
Ba + 


+ 1310K ; 
+ 1284K ; 

O = 

= BaO + 1242K ; 
M*+ = MffO + 1440K; 

Zn + = ZnO + 853K ; 
Cd + 

Ca + S = CaS + 896K; 

Sr + S = 8rS + 974K. 

Ba + S = BaS + 983K j 

+ H 2 O Ca(OH) 2 + 155K. 
+ H 2 - Sr(OH) 2 + 177K. 
+ H 2 = Ba(OH) 2 + 223K. 
+ H 2 = M(OH) 2 + 50K. 
+ H 2 O = Zn(OH) 2 - 26K. 
+ H 2 O = Cd(OH) 2 + 657K. 

Mff + S =- MffS + 776K. 
Zn + S = ZnS + 396K, 
Cd + S = CdS + 324K. 




Oxides and Sulphides of Boron, Scandium, 
Yttrium, Lanthanum, and Ytterbium. 

Of these, boron oxide and sulphide, and the oxides of the 
remaining elements of the group have alone been investigated. 
The selenides and tellurides are unknown. 

Sources. These compounds do not occur native. Boron 
oxide is found in combination with water, as B 2 O 3 .3H 2 O, as 
sassolite; with sodium oxide as borax, 2B 2 O3.Na2O.10H 2 O ; with, 
magnesium oxide and chloride as boracite, 8B 2 O 3 .6MgO.MgCl 8 ; 
and with silicon and calcium oxides as datolite, 

3SiO ? .B 3 O 3 .2CaO.H 2 O. 

Scandium, yttrium, and ytterbium oxides are found in combination 
with silica in gadolinite, and with niobium and tantalum oxides in 
yttrotantalite, samarskite^ and euxenite; while lantharium oxide 
accompanies cerium and didymium oxide in cerite, in combination 
with silica. 

List. Boron. Scandium. Yttrium. Lanthanum. Ytterbium. 
Oxygen.. B 2 O 3 . SoaO 3 . Y 2 O 3 ; Y 4 O 9 . La 2 O 3 ; La 4 O 9 . Yb 2 3 
Sulphur , B 2 S 8 . 

Preparation. 1. By direct combination. Boron burns in 
oxygen or nitric oxide, NO. Yttrium is also oxidised when, 
heated in air, and lanthanum becomes covered with a steel-blue 
film. When strongly heated it takes fire and burns. The other 
elements of this group have not been prepared. Boron unites 
with sulphur at a white heat. 


2. By heating the hydroxides, &c. This is the usual method 
of preparation. These substances part with water at a red heat, 
leaving the oxides. The oxalates, carbonates, and nitrates of 
scandium, yttrium, lanthanum, and ytterbium also yield the 
oxides when heated to redness. 

3. By double decomposition. Boron oxide mixed with 
Carbon, and heated to redness in a stream of carbon disulphide gas, 

'yields the sulphide. 

Properties. Boron trioxide, B 2 O 3 , is a non-volatile glass, 
melting to a viscid liquid at a red heat. It reacts with and 
dissolves in alcohol and in water. When fused with the oxides of 
metals they*are dissolved, forming borates, i.e., double oxides of 
boron and the metal. The sulphide, B s Sj, is a whitish-yellow 
substance, volatile when heated in a stream of hydrogen sulphide, 
and melting at a red heat. It is decomposed by water, yielding 
boracic acid and hydrogen sulphide. 

The oxides of scandium, yttrium, lanthanum, and ytterbium 
are white powders, insoluble in water, and soluble with difficulty in 
acids. They do not react with alkaline hydroxides, nor do they 
fuse in the oxyhydrogen flame. The peroxides of yttrium and 
lanthanum are also white powders, which part with the excess of 
oxygen when heated. 

Mass of 1 cubic centimetre : B 2 O 3 , 1'85 grams at 14'4 ; Sc 2 3 , 3*8 grains ; 
Y 2 3 , 5 03 grains at 22 ; La 2 O 3 , 6'5 grams at 17 ; Yb 2 O 3 , 9*2 grams. 
Heat of formation :B 2 + 3O = B 2 O 3 + 3172K; + Aq = 180K. 

Double compounds. (a.) With water. Preparation. 
Boron trioxide dissolves in water with evolution of heat, com- 
bining with it to form the compound B a O 3 .3H. 2 O, or H 3 BO 3 , 
commonly called boracic acid. The same compound can also be 
prepared .by addition of sulphuric acid to a solution of borax or 
some other borate in water, when the sodium of the borax is 
replaced by hydrogen, thus : 

NaaB 4 7 .Aq + H 3 S0 4 .Aq + 5H 2 = 4H 3 BO 3 + Ka 8 S0 4 .Aq. 

The boracic acid separates in pearly- white scales, which have 

a bitterish cooling taste. Boracic acid is also obtainable by the 

action of moist air on boron ; also by boiling boron with nitro- 

.hydrochloric acid, when it unites simultaneously with oxygen and 


The hydrated oxides of scandium, ytterbium, lanthanum, and 
didymium, are produced, like those of magnesium, by adding 
hydroxide or any soluble hydroxide to solutions of the 


chlorides, or any other soluhle compounds of the metals. They 
are insoluble in and do not combine with these hydroxides to 
form compounds undecomposed by water. 

Boracic acid is a natural product, obtained in volcanic 
districts, especially in Tuscany, and in the Lipari Islands. The 
native form is named sassolite. Steam containing vapour of 
boracic acid issues from jets in the ground called soffioni. The 
steam from these jets is made to blow into artificial basins or' 
lagoni, where the boracic acid condenses along with the steam. 
The solution is concentrated by causing it to flow over long sheets 
of lead, heated by the waste steam of the soffioni. It finally runs 
into crystallising tanks, where the boracic acid separates out on 
cooling. The crude product contains about 76 per cent, of boracic 
acid; it is purified by recrystallisation. Other compounds of 
boron trioxide with water are produced by heating H 3 BO 3 ; these 
are B*O 3 .H 2 O and 2B,O 3 .H 2 O. The first remains on heating to 
100; the second is left at 160; while at 270 the compound 
8B 2 O 3 .H 2 O is said to remain. 

Properties. Boracic acid, H 3 BO 3 (B 2 O 3 .3H 2 O) crystallises in 
nacreous laminae; the other compounds are glassy substances. 
The hydrates of scandium, &c., are white gelatinous precipitates. 
Their exact composition has not been ascertained. Boracic acid is 
volatile with steam ; and it reacts also with ethyl and especially 
with methyl alcohol, forming volatile compounds. It is estimated 
by distilling with sulphuric acid and methyl alcohol ; the distillate 
is evaporated to dryness with a known weight of lime. It is used 
as an antiseptic, and is employed as a preservative of milk, fish, 
&c. A flame held in the steam evolved from a boiling solution is 
tinged green ; if alcohol be present, it burns with a green flame. 
This constitutes the usual qualitative test for boron. 

(&.) With hydrogen sulphide. None of these possible com- 
pounds has been investigated. 

(c.) Compounds of oxides with oxides. No compounds of 
scandia, &c., are known with the oxides of elements preceding 
them in the periodic table. They combine with sulphur trioxide, 
forming sulphates, colourless crystalline bodies ; with nitric pent- 
oxide, forming nitrates, &c. These compounds are considered later. 

Boron trioxide combines with other oxides when they are 
heated together. The resulting compounds are termed borates. 
The most important of these is borax, sodium borate. The follow- 
ing is a list of the more typical of these compounds ; in this classi- 
fication the combined water has not been included, as there is no 
evidence that it replaces either oxide of boron or oxide of the com- 



bined metal. The ratios are very numerous and complex. The 
metal, in the following table, has been considered analogous to 
calcium oxide, CaO, and has been termed MO in the heading. It 
would correspond to M 2 3 , or to M 2 0. The amount of water in 
the salts which have been prepared has been placed in brackets ; 
if another classification is adopted (see Silicates, p. 308), it often 
becomes an integral portion of the formula. The question of these 
formulae will be treated of further on, under silicates, phosphates, 
&c. The ratio given is that of the oxygen in the boron trioxide to 
the oxygen in the metallic oxide, the water, as before stated, being 

Ratio 2 


(2B 2 O 3 .15MO). 

,, 2 


(2B 2 O 3 .12MO). 

,, 2 


(2B 2 3 .9MO). 

>, 1 


(2B 2 3 .6MO). 

,, 6 


(2B 2 O 3 .5MO). 



(2B 2 O 3 .4MO). 

,, 2 


(2B 2 O 3 .3MO). 



(2B 2 3 .2MO). 

4 : 1 (4B 2 O 3 .3MO). 

5 : 1 (5B 2 O 3 .3MO). 

6 : 1 (2B 2 3 .MO). 

9^: 1 (3B 2 O 3 .MO). 
12:1 (4B 2 O 3 .MO). 

15 : 1 (5B 2 O 3 .MO). 
18 : 1 (6B 2 O 3 .MO). 

2B 2 3 

2B 2 O 3 .4Al2O 3 (6H 2 O ; also anhydrous). 
2B 2 3 .3A1 2 3 .(7H 2 0). 
B 2 O :{ 3Na 2 O; B 2 O 3 .3CaO.CaCl 2 ; 

B 2 3 .3CdO(3H 2 O). 
2B 2 O 3 .5BaO. 

B 2 O 3 .2BaO; B 2 O 3 .2MgO. 

2B 2 3 .aCaO ; 2B 2 3 .3SrO ; 2B 2 O 3 .8CoO(4H 2 0). 
B 2 O 3 Na 2 O(3H 2 O, also 4H 2 O) ; K 2 O ; CaO(2H 2 O, 
also anhydrous) ; SrO ; BaO (10H 2 O, also H 2 O) 
MgO(4H 2 O, also 8H 2 O) ; CdO ; 

3B 2 O 3 Fe 2 O 3 (3H 2 O) j 

B 2 3 .NiO(2H 2 0); PbO(H 2 O); Agr 2 b(H 2 O); 
also B 2 3 .PbO.PbCl2(H 2 0). 
4B 2 3 .3A8T 2 0. 
5B 2 3 .3SrO(7H 2 0). 

2B 2 O 3 .Li2O.5H 2 O ; 2B 2 O 3 .Na 2 O(10H 2 O, borax ; 
6H 2 O ; 5H 2 O ; also 3H 2 O). 
K20(5H 2 0) ; (NH 4 ) 2 0(3H 2 0, also 4H 2 O) } 
BaO.(H 2 O) ; SrO(4H 2 O, also anhydrous) ; 
BaO(5H 2 O, also anhydrous) ; PbO(4H 2 O). 
3B 2 O 3 .Li2O.6H 2 O ; 3B 2 O 3 BLjO (8H 2 O) ; 
BaO(14H 2 0); M*0(8H 2 0). 
4B 2 3 .Li 2 0.10H20 ; 4B 2 O 3 .Na2O(10H 2 O) ; 
(NH 4 ) 2 0(6H 2 0) ; CaO(9H 2 0)j SrO(6H 2 O)j 

6B 2 3 .K20(9H 2 0)j (NH 4 ) 2 0(9HaO) j 
M*0.(18H 2 0). 

T6is list comprises nearly all the known berates. They are prepared by one 
of three methods : (1.) By mixing a solution of boracic acid with the hy 
droxide or carbonate of the metal, evaporating, and crystallising. This method 
applies only to the borates of the elements of the sodium group ; their hy- 
"droxides and carbonates, as also their borates, are soluble. (2.) By heating the 
oxide or carbonate, or even the nitrate or sulphate, of the metal with boron 
trioxide to a high temperature. The mass often crystallises on cooliug. The 


berates of many oxides such as those of copper, nickel, &c., are coloured. Few 
of them have been analysed. (3.) By adding a soluble borate such as sodium 
borate to a soluble salt of the metal. A precipitate is formed with all elements 
except those of the sodium group. These precipitates, when washed with 
water, are decomposed, the boracic acid being washed out, and the hydroxide of 
the metal remaining behind. They are thus unstable compounds, largely or 
wholly decomposed by water. 

The compounds containing water are almost always crystalline ; those pro- 
duced by fusion are also often crystalline, but are sometimes, like glass, amor- 
phous j those produced by precipitation are of doubtful existence, inasmuch 
as a mixture of hydroxide and borate might on analysis give numbers which 
would lead to a definite formula. 

The most important of these bodies is borax. It occurs as an 
incrnstration on the soil of districts in Central Asia, and is known 
as tincal ; it is found most abundantly, however, in lakes in 
California, 450 miles S.E. of San Francisco, the most impor- 
tant of which is 12 miles in length and 8 miles broad ; the greater 
part of " Borax Lake " is dry, and the surface is charged with 
borax, common salt, sodium and magnesium sulphates, and salts 
of ammonium. These salts are collected and purified by reerystal- 
lisation. A solution of borax dissolves many substances insoluble 
in water, such as stearic acid, resins, arsenious oxide, &c. It is 
chiefly employed for glazing porcelain and for soldering metals ; 
the film of oxide coating the heated metal dissolves in melted 
borax, and clean surfaces of the metal can thus be brought in 
contact. It has also considerable antiseptic and detergent 

(d.) Double compounds of sulphides, selenides, and tellu- 
rides are unknown, also (e.) compounds of sulphides and 

(/.) Compounds of oxides with halides. The only com- 
pounds which have been prepared are the doiible fluorides and oxides 
of boron and metals, and an oxychloride. Boron trioxide dissolves 
in hydrofluoric acid, and the solution, when concentrated by stand- 
ing over sulphuric acid, is a syrup, which contains B 2 3 and HF in 
the ratio B 8 3 .6HF.H 2 O; it has been named fluoboric acid. The 
same liquid is obtained by saturating water with boron fluoride, 
BF 3 , and distilling. The existence of this body as a definite sub- 
stance appears to be questionable. It is decomposed by water into 
boracic and hydrofluoric acids.* 

The oxychloride, BOC1, is produced by heating to 150 a mix- 
rare of B 2 O 3 and 2BC1 8 . It is a fuming liquid. With water it 
/ields boracic and hydrochloric acids. 

Badsarow, Comptes rend., 78, 1698. 


Oxides, Sulphides, Selenides, and Tellurides of 
Aluminium, Gallium, Indium, and Thallium. 

These are as follows : 

Oxygen. Sulphur. Selenium, Tellurium. 

Aluminium A1 2 O 3 . Al^Sa- P ? 

Gallium QaO(?) ; Ga 2 O 3 . Qa^S, (?), ? ? 

Indium In 4 O 3 ? ; In 2 O 3 . In 2 S 3 . ? ? 

Thallium T1 2 O ; T1 2 O 3 ; (T10 2 )*. T1 2 S ; T1 2 S 3 . Tl 2 Se. ? 

Sources. Alurainiutn oxide, A1 2 O 3 , occurs native in a pure 
state as GO; undum ; contaminated with ferric oxide as emery ; 
coloured blue by cobalt oxide as sapphire; coloured red by chromium 
oxide as ruby; coloured purple by manganese sesquioxide, as ame- 
thyst; and yellow by ferric oxide, as topaz. It also occurs in com- 
bination with water, with silica, and with other oxides (see below ; 
Silicates, p. 303; and Spinels, p. 241). Gallium and indium sulph- 
ides accompany zinc in some blendes ; and thallium is found in the 
" flue-dust " of pyrites burners, being contained in certain samples 
of iron pyrites, KeS 2 . 

Preparation. 1. By direct union. The metals all oxidise 
when heated in air, but not very readily. Fused aluminium 
becomes coated with a film of its oxide, A1 2 O 3 ; gallium, too, 
oxidises only on its surface, even when strongly heated ; indium 
forms a film of pale-yellow In,O H , and thallium becomes covered 
with a layer of a mixture of TL 2 O and T1 2 O 3 . The sulphides and 
selenides may also be prepared by direct union ; Tl 2 S a can be 
prepared only thus. 

2. By heating compounds. (1.) The hydrates, when heated, 
yield the oxides. Aluminium hydrate loses all its water at 360 ; 
indium hydrate at 655 ; and thallium hydrate at 230. (2.) The 
compound of indium sulphide, In 2 S 3 , with hydrogen sulphide 
loses hydrogen sulphide when heated. (3.) Aluminium oxide has 
been prepared by heating potash alum, KJEJO^Al^SO^g, to white- 
ness; a mixture of potassium sulphate and alumina remains, 
sulphuric anhydride escaping ; the potassium sulphate is dissolved 
out with water, leaving the alumina. (4.) Gallium oxide has 
been* prepared by heating the nitrate. 

3. By double decomposition. Gallium sulphide, Ga 2 S 3 , is 
produced by addition of a soluble sulphide (ammonium sulphide 
has been used) to a soluble salt of gallium ; indium sulphide, 
In,S 3 , is precipitated by hydrogen sulphide, Solutions of thallous 

* In combination. 


salts, such as T1NO 3 , or T1 2 S0 4 , give with hydrogen sulphide a 
precipitate of Tl^S. If a thallic salt be used, it is first reduced to 
a thallous salt by the hydrogen sulphide, with separation of 
sulphur, and thallous sulphide is then precipitated, thus : 

TlCl 3 .Aq + H 2 S = T1C1 4- 2HCl.Aq + S 
2T101.Aq + H 2 S = T1 2 S + 2HCl.Aq. 

When carbon disulphide gas is passed over red-hot alumina, 
some of the oxide is converted into sulphide. A similar action 
takes place with hydrogen sulphide. Indium sulphide, In 2 S 3 , may 
be produced in scales like mosaic gold, by fusion of indium 
trioxide, Ill^Oa, with sodium carbonate and sulphur!* No doubt 
sodium sulphate is formed at the expense of the oxygen of the 
indium oxide, and the indium combines with the excess of 

4. By the action of heat, in a current of hydrogen, gallium 
trioxide gives a bluish-grey sublimate, supposed to be monoxide ; 
and indium trioxide, IHjO 3 , similarly treated, gives a mixture of 
oxides, one of which is said to have the formula In,O 3 . It is 
probably a mixture or a compound of In 2 O with In,O 3 . When 
thallic oxide, T1 2 O 3 , is heated to 360 it begins to lose oxygen, 
giving the compound 3T1 4 O 2 .T1 2 O, which is perfectly stable up to 
565; at higher temperatures, up to 815, T1 2 volatilises away; 
and the residue T1 2 O 3 is stable in presence of air above that 
temperature. The monoxide, T1 2 O, when heated in air is partially 
oxidised to T1 2 O 3 . 3 

By removing oxygen from thallous sulphate, ^I t **Oi, thallous 
sulphide is left. This action is analogous to the l&s of oxygen 
which sodium, and barium sulphates, &c., suffer when heated in 
hydrogen or with carbon. In the case of thallium, the sulphate is 
heated with potassium cyanide, KCN, which is doubtless con- 
verted into cyanate, KCNO. 

5. Special methods. Crystalline alumina has been produced 
by fusing the amorphous variety in the oxyhydrogen flame ; by 
heating the oxide along with aqueous hydrochloric acid to 350 in 
a sealed tube ; and by melting together at a white heat aluminium 
oxide with lead monoxide (litharge), or with barium fluoride. 
The last two processes yield artificial corundum ; and if a trace of 
cobalt oxide or chromium oxide be added, artificial sapphires or 
rabies are produced.* 

Properties. Trioxides. Aluminium and gallium trioxides 
are white powders, or friable masses ; indium trioxide has a tinge 
Compt. r*nd. t 104, 737. 


of yellow, especially when hot ; and thallium trioxide is a brown 
powder. Crystalline aluminium trioxide is exceedingly hard, and 
is insoluble in acids. The amorphous trioxide, when ignited, 
appears also to alter its structure, probably polymerising (i.e., 
several molecules unite to one), for it is then almost unattacked 
by -acids. It can still be dissolved, however, by boiling sulphuric 
or strong hydrochloric acid, forming the sulphate or chloride ; the 
crystalline variety is totally insoluble. All the trioxides are 
without action on water. 

Trisulphld.es, &e. Aluminium tri sulphide forms yellow 
crystals, which turn dark when heated ; the selenide and telluride 
are black non- volatile powders ; gallium trisulphide has not been 
described ; indium trisulphide is a brown powder, or gold-coloured 
scales ; and thallium trisulphide, a black, ropy substance, brittle 
below 12. Aluminium sulphide is decomposed by water, giving 
the hydrate and hydrogen sulphide, thus : 

A1 2 S 3 + 3H 2 O.Aq = Al,O 3 .rcH 2 O + 3H,8. 

The other three are unchanged by water, but decompose when 
boiled with acids. 

Other oxides and sulphides. There are no lower oxides or 
sulphides of aluminium ; the lower oxide of gallium, produced by 
heating the trioxide in hydrogen, is a bluish-grey substance. The 
lower oxides of indium are black powders. 

Thallium monoxide is a reddish-black substance, melting about 
300, and is volatile between 585 and 800. When heated with 
sulphur, the oxygen is replaced by sulphur. It combines directly 
with water, forming the hydrate, T1 2 O.H 8 O, and absorbs carbon 
dioxide from moist air. It has thus some resemblance to the 
hydroxides of the metals of the sodium group. Thallous sulphide, 
when precipitated, forms greyish or brownish flocks ; from a hot, 
slightly acid solution it comes down in blue-black crystals. It 
may be fused to a black lustrous mass like plumbago. The 
selenide is a black crystalline body. 

Physical Properties. 

1. Mass of 1 cubic centimetre : ALjOj, : 3'98 grams at 14 ; In 2 O 3 : 7*18; 

TLjS: 8'0. 

2. Melting-point: TljO 3 : 759. 

3. Heats of formation : 2A1 + 38 = A1 2 S 3 + 1224K. 

2T1 + O - TljO + 423K; + H S O - 33K. 
2T1 + 8 - TLzS + 197K. 

Double compounds. (a.) With water: hydrates or hydr- 
oxides. The hydrated trioxides are produced by addition of a 


soluble hydroxide, such as that of sodium, potassium, or barium, 
or even of thallium (T10H), to solutions of soluble salts of the 
metals. A solution of ammonia in water acts in a similar manner, 
as if it contained ammonium hydroxide, NHi.OH. The reaction 
is as follows, e.g., with aluminium: 

Al 2 Cl 6 .Aq + CKOH.Aq = Al^wH.O + GKCl.Aq. 

Excess of precipitant (except ammonia) dissolves the hydrates 
of aluminium and gallium ; gallium hydroxide is soluble even in 
solution of ammonia. Solution takes place owing to the formation 
of soluble double compounds (see below). 

Aluminium hydroxide may also be produced tyy passing a 
current of carbon dioxide into a solution of potassium alnminate 
(A1 2 O 3 K 2 O). Potassium carbonate is formed, and the hydrate of 
aluminium precipitated. Aluminium sulphide, A1 2 S 3 , reacts with 
water, giving the hydrate and hydrogen sulphide. Hence, when 
solution of ammonium sulphide is added to a soluble aluminium 
compound, the hydrate is precipitated, whilst sulphuretted 
hydrogen is evolved. 

The sulphides are not known to form compounds with water. 

Thallium monoxide, Tl a O, dissolves in water, and on cooling, 
or. on evaporation, the solution deposits yellow needles of 
TLO.H 2 O = 2T1OH. Its solution absorbs carbon dioxide from 
the air. Aluminium hydrate, prepared by precipitation, forms 
gelatinous flocks, and when dried at ordinary temperature in 
air, has approximately the formula A1 2 O.5H 2 O. This is a 
hard, horny mass ; when heated it gives up its water. Up 
to 65 the loss is rapid, and at that temperature the hydrate 
has approximately the formula Al a O 3 ,3H^O. The rate of loss 
of water diminishes as the temperature rises to 150, and 
increases up to 160, diminishing again up to 200. 4 The com- 
position of the hydrate between 160 and 200 is nearest the 
formula A1*O,.'2H 8 O. From 200 to 250 the rate of loss of water 
is rapid, but is much slower between 250 and 290, and here the 
formula approximates to AliO 3 .H 2 O. Complete dehydration does 
not occur till 850 is reached. It is probable that there are many 
hydrates of alumina, but that no one is stable over any consider- 
able range of temperature. 

The action of excess of water, however, on aluminium 
amalgam yields a crystalline hydrate of the formula A1(OH) 3 
s AljO,.3H,O. 

Three natural hydrates are known, gibbsite, A1<O 3 .3H 2 O, 
bauxite, Al a O 3 .2HtfO, and diaspore, Al 2 O a .H 2 O. Artificial crystals 


of gibbshe have been produced by the slow action of the carbonic 
acid of the air on a solution of aluminate of potassium ; and by 
boiling aluminium acetate or hydroxide for a long time with water, 
the dihydrate is said to be precipitated. 

Indium hydrate is a gelatinous white precipitate, which when 
air-dried has approximately the formula In 2 Oj.6H 2 O. When 
heated, it loses water gradually up to 150. The rate of loss then 
increases to 160, again to slacken. The composition between 150 
and 160 nearly corresponds to the formula In 2 O 3 .3H 3 O. It is not 
dehydrated till 655 ; and there are no signs of other hydrates. 

Air-dried hydrate of thallium has the formula Tl 2 O 3 .H,jO. At 
higher temperatures it is dehydrated. 

(?>.) With hydrogen sulphide. Indium sulphide, In,S>, 
when precipitated from soluble compounds of indium with 
hydrogen sulphide, has a deep yellow colour. It can be dried in 
air, but when heated it evolves hydrogen sulphide, leaving the 
sulphide. It is probably a compound of the nature of a hydrate : 
In 2 S 3 .MH,S. The white precipitate produced by ammonium 
sulphide with salts of indium is also probably of this nature. It 
is soluble in solution of ammonium sulphide. 

(c.) Compounds of oxides with oxides. On adding a 
solution of potassium hydroxide to aluminium hydrate, complete 
solution occurs when the ratio of the alumina to the potash 
is as A1 2 O 3 : K 2 O ; the same compound is precipitated when a 
solution in excess of hydroxide is mixed with alcohol, in 
which caustic potash is soluble, but not the aluminate. It has 
the formula A1 2 O 3 .K 2 O = 2KA1O 2 . A similar sodium compound 
has been prepared. The compounds Al 2 O 3 .2Na 2 O and Al a O3.3Na 2 
are also said to have been prepared. By dissolving hydrate of 
alumina in solution of barium hydroxide and evaporating, crystals 
of Al 2 O 3 .BdO.6H 2 O, Al 2 O 3 .2BaO.5H a O, and ALO 3 .3BaO.llH 2 O 
are successively deposited.* These bodies may be compared with 
the borates. 

The mineral named spinel is a compound of alumina with 
magnesium oxide, Al 2 O 3 ,MgO. It crystallises in octahedra, and 
has been prepared artificially. Gahnite is a similar compound 
with zinc oxide of the formula Al 2 O 3 .ZnO, and chrysoberyl with 
beryllium oxide Al 2 O 3 .BeO. 

Two compounds with barium oxide and chloride are also known, 
viz., Al 2 O 3 .BaO.BaCl2 and Al 2 O 3 .BaO.3BaCl 2 . 

Gallium oxide would, no doubt, enter into similar combinations, 
but these have not been investigated. 

* Berichte, 14, 2151 j J. prakt. Chem., 26, 385, 474 j Cfam. News, 42, 29. 



A higher oxide of thallium in combination with barium oxide is 
produced bypassing a rapid current of chlorine through potassium 
hydroxide, in which thallic hydrate is suspended. The solution 
turns violet, and with barium nitrate gives a violet precipitate 
which contains the oxide T1O 2 .* 

(d.) Compounds of sulphides with sulphides. Indium sul- 
phide forms with potassium and sodium sulphides red crystalline 
compounds of the formulae In,S 3 .K^S, and InJSa.NaaS. A silver 
compound of similar formula is produced on addition of silver 
nitrate to their solutions. Thallic sulphide, T1 2 S 3 , also unites with 
thallous sulphide, T1 3 S, giving black crystalline bodies. 

(e.) No compounds of oxides with sulphides ark known. 

(/.) Compounds with halides. On evaporating an aqueous 
solution of aluminium chloride, it is probable that oxy chlorides are 
produced, inasmuch as hydrogen chloride is evolved. On repeated 
evaporation, all aluminium remains as hydroxide. Similar com- 
pounds, but somewhat indefinite, have been produced by the action 
of aluminium chloride on aluminium in presence of air. Gallium 
chloride, on addition of wa.ter, gives a white precipitate of oxy- 
chloride, the formula of which is unknown. 

Uses. The chief use of alumina is as a mordant. When a salt 
of aluminium in solution is boiled in contact with animal or vege- 
table fibre, it splits into acid, and hydrate of alumina, the latter 
depositing on the fibre. The fibre has the power of absorbing 
and "fixing" colouring matters, when boiled with their solutions. 
If the colouring matter be dissolved in water along with a salt of 
aluminium, and the solution be boiled, the hydrated alumina often 
comes down in combination with the colour, giving a " lake/* 
Such lakes are made use of as paints. 

Physical Properties. 
Mass of 1 cubic centimetre : 

B. Sc. Y. La. Yb. Al. Oa. In. Tl. 

1-85 3-8 5-0 6-5 9'2 3 '90 4 7'18 

OH 1-49 - 2'89f 

S ______ __ _ _ 8-00^ 

Heats of formation. 

2B + 3O = B 2 O 3 + 2 x 1586K ; + 3H 2 O = 2B(OH) 3 + 2 x 60K. 
2A1 -I- 3O + 3H 2 O = 2A1(OH) 8 + 2 x 1945K. 

2T1 + = T1 2 + 423K; + H 2 O = 2T1OH + 33Kj + Aq = - 32K. 
2T1 + 3O + 3H 2 O = 2T1 2 (OH) 8 + 2 x 432K. 
2A1 + 3S = A1 2 S 3 + 2 x 612K. 
2T1 + 8 - T1 2 S + 2 x 98-5 K. 
* Gazzetta, 17, 450. 

t Gibbsite, A1(OH) S . Diaspore, AIO(OH) = 3-4. 
t T1 3 S. 








Oxides, Sulphides, Selenides, and Tellurides of 

Chromium, Iron, Manganese, Cobalt, and 


These compounds may be divided into five well-defined groups : 
(1) the monoxides, monosulphides, &c., such as FeO, PeS, 
&c. ; (2) the sesquioxides, sequisulphides, &c., for example, 
Fe 2 O 3 , Fe 2 S 3 , &c. ; (3) the dioxides, such as MnO 2 ; (4) the 
trioxides, of which CrO 3 is an instance ; and (5) the heptoxides, 
of which compounds are known in the case of manganese, Mn 2 O 7 . 
The double compounds will be considered in connection with each 
group. As these bodies or their compounds are very numerous, it 
is advisable to consider them in the order of the above groups. It 
may be noticed generally that in formula, preparation, and proper- 
ties, the monoxides, &c., show a certain resemblance to those of 
magnesium, zinc, and cadmium ; while the sesquioxides, &c., are 
comparable with those of aluminium. The trioxides find their 
closest analogues in the sulphur group ; and the compounds of 
manganese heptoxide have the same crystalline form as the per- 

I. Monoxides, monosulphides, monoselenides, and mono- 
tellurides : 

List. Oxygen. Sulphur. Selenium. Tellurium. 

Chromium ... CrS. CrSe. ? 

Iron FeO. (Fe 2 S)FeS. FeSe. FeTeP 

JManganese . . . MnO. MnS. ? ? 

Cobalt OoO. CoS. CoSe. ? 

Nickel NiO. (NisS)NiS. NiSe. ? 

Sources. CrO is said to exist in combination with Cr^O 3 in 
some chrome ores* FeO exists ID combination with C0 t as carV*^ 

B 2 


nate in spathic iron ore, and with Fe 2 O 3 in magnetite. MnO has 
been found native. It forms crystals which, reflect green, and 
transmit red light. NiO is also found native. PeS is sometimes 
found in meteoric iron, in combination with dinickel sulphide, 
Ni;S, as 2FeS.Ni 2 S. Mariganous sulphide, MnS, occurs as man- 
yanese blende, or alabandine, in iron-black lustrous cubes or octa- 
hedra. Native cobalt sulphide is known as syepoorite. It forms 
steel-grey to yellow crystals, and is used by Indian gold workers 
to give a rose-colour in burnishing gold. Nickel sulphide, NiS, 
occurs in nature as long brass-yellow needles, and is named capillary 
pyrites, or millerite. Nickel oxide, NiO, along with magnesium 
oxide, occurs as a silicate in the ore from New Caledonia. The ore 
contains about 18 per cent, of nickel oxide. 

Preparation. 1. By direct union. Higher oxides are pro- 
duced by the union of chromium, iron, manganese, and cobalt with 
oxygen ; but nickel burns to NiO. Iron, manganese, cobalt, and 
nickel unite directly with sulphur, selenium, and probably tellu-' 
num, forming monosulphides, <fcc. 

The union with sulphur may be illustrated by heating an intimate mixture 
of iron filings with sulphur in a test tube j the mixture glows throughout, and 
is convert ed into ferrous sulphide. 

2. By heating double compounds. Iron, manganese, cobalt, 
and nickel oxides may be obtained by heating the oxalates, thus : 

FeC 2 O 4 = FeO + GO + CO*. 

Manganese, cobalt, and nickel monoxides are produced when their 
carbonates or hydroxides are heated, thus : MnCO 3 = MnO -f 
00 2 ; Ni(OH)i = NiO + H 2 0. Air must be excluded, except 
in the case of nickel. Nickel monoxide alone is produced on 
igniting the nitrate; with the other metals higher oxides are 
formed. We here see a proof of the comparative 1 stability of 
the higher oxides ; those of chromium being most, and tho^e of 
nickel least ^stable. 

3. By reducing a higher oxide or sulphide. Iron sesqui- 
oxide, Fe^Oa, heated in a mixture of carbon monoxide and dioxide, 
such as is produced by the action of sulphuric acid or oxalic acid, 
is reduced to the monoxide. It is also produced in a crystalline 
form by heating iron to redness in a current of carbon 'dioxide ; 
and by heating the sesquioxide, Fe 2 O 3 , in hydrogen ; between he 
temperatures 330 and 440 magnetic oxide, Fe 3 O 4 , is produced ; 
bat from 500 to 600 the product is FeO. Ab still higher tern- 
peratures metallic iron is formed.* The higher oxides of cobalt 

* Chem. Soc., 38, 1, 606; 37, 790. 


and nickel lose oxygen when heated alone, the former at a white 
heat, the latter at a red heat. 

Chromium monosulphide and raonoselenide, CrS and CrSe,have 
been produced by heating the sesquisulphide or selenide to redness 
in hydrogen. Ferric sulphate, Fe 2 {SO 4 ) 3 , heated in hydrogen is 
said to give Fe 9 S ; and ferrous sulphate, FeSO 4 , heated in sulphur 
vapour, Fe 2 S. As both these bodies are strongly magnetic, there 
appears reason to suspect that they contain metallic iron; they 
are blackish- grey powders. When heated with carbon, FeSO 4 is 
said to yield FeS ; cobalt sulphate -behaves similarly. 

Ferrous sulphide, FeS, is produced by heating to redness the 
disulphide, Iron pyrites, FeS 2 , or magnetic pyrites, Fe,S 3 .3FeS ; 
sulphur volatilises ; it may also be formed by heating pyrites, 
FeS 2 , with metallic iron. Cobalt sulphate, CoSO 4 , heated in 
hydrogen, gives an oxysulphide, CoO.CoS(see below) ; but nickel 
sulphate yields dinickel sulphide, Ni 2 S. 

4. By double decomposition. Manganous oxide is most 
easily prepared by heating the di chloride, MnCl 2 , with sodium 
carbonate, Na 2 CO 3 , and a little ammonium chloride. The reaction 
is as follows, MnCl 2 + Na 2 CO 3 = MnO 4- 2NaCl + 00* The 
oxide is really formed by decomposition of the carbonate produced 
by double decomposition. The fused mass is deprived of sodium 
chloride by treatment with water. Higher oxides of iron or man- 
ganese, when heated with sulphur to a high temperature, yield the 
monosulphide ; the sulphur combining with, as well as replacing 
oxygen. Thus Fe 2 O 3 and Fe^, Mn 2 O 3 and MnO 2 yield mono- 
sulphides, and sulphur dioxide ; both reduction and double de- 
composition proceed simultaneously. Manganese dioxide is also 
converted into sulphide when it is heated in vapour of carbon 
disulphide, the carbon removing oxygen while manganese and 
sulphur unite. Cobalt and nickel sulphides have also been pro- 
duced by heating the oxides in a current of hydrogen sulphide or 
sulphur gas. All monosulphides (and probably also monoselenides 
and tellurides) are precipitated on adding to a soluble chromous, 
ferrous, manganous, cobalt, or nickel compound a soluble sulphide 
(selenide or telluride). Ammonium sulphide is commonly em- 
ployed. Manganese, cobalt, and nickel sulphides are also precipi- 
tated from solutions of their acetates by hydrogen sulphide. The 
typical equations are : 

FeS0 4 .Aq + (NH 4 ),S. Aq = FeS + (NH 4 ) 2 S0 4 .Aq ; 
Mn(C 2 H 3 2 ) 2 .Aq + H 2 S = MnS + 2C 2 H 4 O 2 .Aq. 

Properties. Ferrous oxide is a black amorphous powder, 


pyrophoric, i.e., igniting and glowing like tinder on exposure to an 
it decomposes water, slowly at the ordinary temperature, quickly 
on boiling, liberating hydrogen. When prepared by the action of 
carbon dioxide on metallic iron, it forms small black lustrous 
crystals. Manganous oxide is a greyish-green powder, melting about 
1500 to a green mass. When heated to redness in a current of 
hydrogen chloride it is converted into transparent emerald-green 
octahedra. Cobalt monoxide is an olive-green, and nickel monoxide 
* greyish-green, powder. The latter has been obtained in crystals 
by fusing a mixture of nickel sulphate and potassium sulphate ; 
sulphur trioxide and its decomposition products, sulphur dioxide 
and oxygen escape, and crystals of nickel oxide remain*disseminated 
through the potassium sulphate, the latter of which can be re- 
moved by solution in water. These bodies are all insoluble in 
water, and are not easily attacked by acids. 

Ohromous, ferrous, cobaltous, and nickelous sulphides when 
prepared by precipitation are black flocculent masses ; manganous 
sulphide, similarly obtained, is pale yellowish-pink. Very finely 
.divided iron sulphide is green when suspended in water. Pink 
manganous sulphide when heated in a sealed tube with yellow 
ammonium sulphide (polysulphide) changes to green, owing prob- 
ably to some molecular change. When prepared in the dry way, 
chromous and cobalt sulphides are grey, ferrous aad nickel sul- 
phides brass-yellow, and the native form of manganous sulphide 
iron-black. They all exhibit dull metallic lustre. Manganous 
sulphide changes to yellow-green hexagonal prisms when heated 
to redness in a current of hydrogen sulphide. The selenides are 
white, yellow, or grey bodies, also with dull metallic lustre. All 
these substances are insoluble in water ; they react with acids, 
giving, for example, with hydrochloric or sulphuric acids, the 
chloride or sulphate of the metal and hydrogen sulphide.' Hydro- 
chloric acid, if dilute, does not attack nickel or cobalt sulphides 
unless it is heated. The action of dilute hydrochloric or sulphuric 
acid on ferrous sulphide is the usual method of preparing hydrogen 

A wet mixture of iron filings, sulphur, and ammonium chloride 
turns hot, owing to the combination of the iron and sulphur. Such 
a mixture is employed in cementing iron, for example in the*' con- 
struction of submarine piers for bridges. The sulphides can all 
be fused at a white heat. Dinickel sulphide, Ni 2 S, can even be 
melted in glass vessels. 

Double compounds. (a.) With water. Hydrates or hy- 
droxides. These substances are prepared, as usual, by the 


action of a soluble hydroxide or of ammonia dissolved in water 
on some soluble componnd of the metal, e.gr., the chloride or 
sulphate. With chromium, iron, and, in a less degree, man- 
ganese and cobalt, great care must be taken to exclude oxygen ; 
the water in which the precipitant is dissolved must be boiled 
in vacuo, to remove dissolved oxygen, and the precipitation, 
nitration, &c., conducted in an atmosphere of hydrogen. Chro- 
mous compounds are best prepared from the acetate, which is 
made by the action of nascent hydrogen from zinc and hydro- 
chloric acid on a solution of chromium trichloride. On adding 
potassium acetate, chromous acetate is precipitated as a red 
powder. Oe. treatment with potassium hydroxide, it yields 
chromous hydrate, 2CrO.H 2 0, as a substance yellow when wet, 
turning brown when dried.* When boiled with water, hydrogen 
is evolved, and chromic hydrate is produced. The water which it 
contains cannot be removed by heat, for the reaction takes place 
2CrO.H 2 O = Cr 2 O 3 H- # 2 . 

Ferrous hydrate, PeO.H 2 O (?), is a white precipitate, which 
becomes much denser on standing in a solution of potassium 
hydroxide. It is sparingly soluble in water (1 in 150,000). It 
absorbs carbon dioxide from air ; and when dry it turns hot and 
oxidises on exposure to air. The wet hydrate, on atmospheric 
'oxidation, turns first green, then rust coloured. 

Manganous hydrate is also white, and turns brown on exposure 
to air. It is said to contain 24 per cent, of water, and hence must 
have approximately the formula 3MnO.4H 2 O. It can be produced 
by boiling manganous sulphide, MnS, with caustic potash. Co- 
baltous hydrate is a dingy- red powder, prepared by boiling a solu- 
tion of cobalt dichloride with caustic potash, and collecting and 
drying the precipitate. In the cold, a blue oxychloride is precipi- 
tated. Tke hydrate of nickel, NiO.2H 2 O, occurs native, in small 
emerald-green prisms ; andNiO.H 2 O = Ni(OH) 2 isan apple-green 
precipitate. By leaving a solution of nickel carbonate in excess 
of ammonia to crystallise, this hy (irate separates m green 

It appears probable that the precipitated sulphides of these 
metals are in reality compounds of the sulphides with water. 

() No hydrosulphides are known ; and (c.) No double oxides 

, (d.} Compounds of Sulphides with Sulphides : 

2FeS.KoS ; obtained by igniting FeJS^KaS in hydrogen. 

3MnS.K 2 S j dark red scales, -produced by heating together man- 

* Annales (5), 25, 416; Comptes rend., 92, 792, 1051. 


ganese sulphate, MnSO 4 , with potassium sulphide and 
carbon, and dissolring out the excess of potassium eul- 
phide witji water. 
3MnS Na 2 S. Light red needles, similarly prepared. Also 

Ni8.2FeS. A double sulphide of nickel and iron, named pentlandite, 

which forms bronze-yellow crystals. 
FeS.2ZnS, known as chnstophite, occurs native j also CoS.CuS, carro* 

(e.) Compound of oxide and sulphide. CoS CoO, a dark grey sintered mass, 

produced by heating cobalt sulphate, CoSO 4 , in hydrogen. 
(f). Compounds with halides. Chromous chloride is said to give a light grey 
oxychloride ; and cobalt chloride heated with water a greenish-blue 
oxy chloride. Similar bodies, green and insoluble, are* produced, wh'en 
nickel chloride or iodide is heated with nickel hydroxide. Their 
formulse are unknown. 

II. Sesquioxides, sesquisulphid.es, sesquiselenides (the 
tellurides have not been investigated). 

List. Oxygen. Sulphur. Solenium 

Chromium ........ Cr 2 O 3 . Cr 2 S 3 . Cr 2 Se 3 . 

Iron .............. Pe 2 O 8 . Fe 2 S 3 . Pe 2 Se 3 . 

Manganese ....... Mn 2 O 3 . 

Cobalt ......... CooO 3 . Co 2 S 3 . 

Nickel ............ Ni 2 O 3 . 

Sources. Chromium sesquioxide exists in combination 
"with ferrous oxide in chrome iron ore or chromite, the chief source 
of chromium. It occurs in veins in serpentine rock. As clirome- 
ochre it forms a yellow-green earthy deposit, which is found in 
Shetland. Iron sesquioxide is very widely distributed, and 
occurs as red hcematite or specular ore in large deposits in Cumber- 
land and Lancashire in early formations ; in carboniferoub strata 
as brown hcematite or Umonite in the Forest of Dean, fn Glamor- 
ganshire, or associated with oolitic rocks as the earthy hsematite 
of Northamptonshire and Lincolnshire. More recent deposits of 
Umonite occur as log -iron-ore in Ireland and North Germany. 
Magnetic ore, or magnetite, Fe a O 3 .FeO, is also very widely dis- 
tributed. It is, perhaps, the purest form of iron ore, and occurs 
as sand in Sweden. From it the celebrated Swedish iron isjnade. 
Magnetic pyrites, Fe 2 S 3 .FeS, and 2Fe 2 S 3 .FeS, and copper pyrites, 
Fe 2 S 3 .Cu 2 S, are made use of as sources of sulphur. Manganese 
sesquioxide, Mn 2 O 3 , occurs as braunite, and hydrated, lAn 2 O 3 .H 2 O, 
as grey maganese ore. Wad is admixture of oxides of manganese, 
probably consisting largely of Mn 2 O 3 . In combination with MnO, 
it forms hausmannite, Mn 2 O a .MnO (see Spinels). Cobalt and 


nickel sesquioxides do not occur native, but COjS 3 .CoS is known 
as linnceite. 

Preparation. 1. By direct union. Chromium, heated in 
air, forms Cr^O 3 ; but iron, manganese, and cobalt burn to com- 
pounds of the sesquioxides and monoxides, depending on the tem- 
perature. A steel watch-spring set on fire by being tipped with 
burning sulphur, burns in oxygen with brilliant scintillations to 
Pe 2 O 3 .FeO, or magnetic oxide, which fuses and drops from the 

FIG. 32. 


This forms a telling experiment, and illustrates well the direct union of 
metals of this group with oxygen. The jar in which the combustion takes 
place should stand in an iron tray, or in a plate full of water, for the fused 
oxide is certain to crack any glass tray on which it falls. 

Iron filings, heated to dull redness in a current of sulphur gas, 
forms Fe^a ; and the corresponding selenide, Pe 2 Se 3 , has been 
similarly made. 

2. By reducing a higher compound. Chromium trioxide, 
CrO 3 , when strongly ignited, loses oxygen, forming the sesqui- 
oxide. Compounds of the trioxide, such as mercurous chromate, 
Hg,CrO 4 , 'ammonium dichromate, (NH 4 ) 2 Cr^O 7 , and others also 
yield the sesquioxide on ignition. Chromates, such as bichrome, 
KoCr-jO?, at a white heat give neutral chromate, chromium 
sesquioxide, and oxygen, thus : 2K 2 Cr 2 O 7 = 2K 2 CrO 4 -f Cr,6 3 
+ 30. Manganese dioxide, at a dull-red heat, likewise loses 
oxygen, giving Mn 2 O 3 . 

S.JBy oxidising a lower compound. Ferrous sulphate, 
FeSO 4 , when distilled for the manufacture of anhydrosulphuric 
rcid (see p. 433), leaves a residue of sesquioxide. It may be sup- 
posed that the ferrous oxide decomposes water-gas, arising from 
water still combined with the ferrous sulphate, producing the 
sesquioxide. Ferrous carbonate heated gently in air yields ferrous 
oxide, PeO, which unites with oxygen, forming the 


It is also produced by Keating ferrous sulphate with a little nitre, 
KNO 3 , to supply oxygen. Ferrous oxalate, FeC 2 O 4 , yields the 
monoxide on ignition, and in air the sesquioxide is produced. The 
lower oxides of manganese, MnO and Mn 3 O 4 , when heated in 
oxvgen give the sesquioxide, when the pressure of oxygen is 
greater than 0*26 of an atmosphere. As the pressure of the 
oxygen in ordinary air is approximately one-fifth of an atmo- 
sphere, such oxidation does not occur in air, unless it be com- 
pressed. The nitrates of these metals, when heated, yield 
the sesquioxides. This is a case of simultaneous decomposition 
and oxidation. The nitrate is decomposed into monoxide and 
nitric pentoxide, thus : Fe(NO.,) 2 = PeO -f N 2 6 ; bat the 
pentoxide parts with its oxygen, being itself converted into lower 
oxides of nitrogen, NO and NO* thus : 2FeO + N^0 & = Fe 2 O 3 
4- 2AT0 2 ; and 6FeO + JV 2 6 = 3Fe 2 O 3 + ZNO. And similarly 
with the other metals. 

4. By the action of heat on a compound. The hydrates of 
these metals when heated leave the oxides. Ferric hydrate, when 
boiled for a long time in water, is ultimately dehydrated, and dry 
ferric oxide settles out. The nitrates and sulphates, &c., are also 
decomposed by heat, and also the borates. The excess of boracic 
acid is removed by weak hydrochloric acid. 

5. By double decomposition. Ferric oxide is produced in 
a crystalline form when ferric chloride and lime are heated to 
redness, or when ferrous sulphate and sodium chloride are heated 
together in air. The ferrous oxide is oxidised by the air, and crys- 
tallises from the salt. The sulphides are generally prepared by 
double decomposition. Chromium sesquisulphide is obtained when 
chromium trioxide is heated to whiteness in a current of carbon 
disulphide gas ; heated in sulphur gas or in hydrogen sulphide, it 
suffers no change ; but the chloride is converted into uulphide or 
selenide by hydrogen sulphide or selenide at a red heat, and the 
hydrate, when heated to 440 in sulphur gas, or to a higher tem- 
peratare in selenium vapour, yields the sulphide or selenide. 
Cobaltic hydrate gently heated in hydrogen sulphide also gives 
cobalt sesquisulphide. Nickel sesquisulphide is unknown. 

Properties. Oxides. Chromium sesquioxide is an amorph- 
ous green powder; when crystalline it forms green tablet^, or, if 
produced at a high temperature, brown crystals. The amorphous 
variety, if it has not been exposed to a high temperature during its 
formation, becomes incandescent when gently heated, no doubt 
owing to polymerisation, several molecules uniting to form one. 
It is then practically insoluble in acids. This behaviour is 


also seen with aluminium, manganese, and iron sesquioxides. The 
crystalline varieties of chromium oxide are produced in presence 
of chlorine, or by some solvent for the oxide. Thus chromium 
oxychloride, CrO z Cl 2 , when passed through a red-hot tube, de- 
posits crystalline oxide ; similarly, potassium dichromate, heated 
in chlorine, gives a mixture of crystalline oxide and potassium 
chloride, the excess of oxygen being expelled. Iron sesqui- 
OXide may be obtained in crystals by fusing the amorphous 
variety with calcium chloride, or by heating it in a current of 
hydrogen chloride. It would appear that in such cases the 
volatile chloride is formed ; and that it is decomposed by oxygen, 
yielding oxidte, which is deposited in crystals. Crystalline 
varieties of the sesquioxides of cobalt and nickel, owing to their 
easy decomposition, have not been obtained. That of manganese 
has not been prepared artificially. Amorphous ferric sesquioxide 
is brown- red or red, according to the method of preparation. If 
prepared from ignited ferrous sulphate it has a fine colour, and is 
used as a paint, under the name " Venetian red." It is also used 
under the name of " rouge " for grinding and polishing glass objects, 
such as the lenses of telescopes, &c., and as " crocus " to produce 
shades from purple-red to yellow, according to the amount, on 
porcelain, in combination with silica. The crystalline variety is 
black. When native, as specular ore, it forms very lustrous 
rhombohedra ; another crystalline variety, martite, occurs in octa- 
hedra ; while hcematite consists of kidney-shaped (botryoidal) 
masses, with a radiated crystalline structure. Manganese 
sesquioxide, when amorphous, is a black powder ; as braimite it 
forms brownish-black lustrous quadratic pyramids. Cobalt 
sesquioxide, prepared by heating the hydrated compound to 
600 700, is a black powder, as is also nickel sesquioxide. 

These bodies show different degrees of stability. While 
chromium sesquioxide can be fused at a white heat without change, 
iron sesquioxide is converted into Fe 3 O 4 , and at a bright-red heat, 
manganic sesquioxide gives Mn 3 O 4 . Cobalt and nickel sesqui- 
oxide lose oxygen at a dull-red heat, giving Co 3 O 4 and NiO 
respectively. Cobaltic oxide, as borate, is made use of as a black 
pigment in enamel painting. Chromium sesquisulphide and 
sesquielenide form brilliant black plates; iron sulphide and 
selenide are yellowish- grey with metallic lustre ; and cobalt 
sesquisulphide forms a dark iron-grey mass. 

Double compounds. (a.) With water : hydrates or hydr- 
oxides. These are produced as usual from a soluble salt of the 
hydroxide. Those of cobalt and nickel are formed by the 


action of an alkaline solution of sodium or potassium hypo- 
chlorite on a salt of the metal. Hydrated monoxide is produced 
and further oxidised by the hypochlorite, thus : 

2CoO//?H 2 O + NaClO.Aq = Co 2 O 3 .rcH 2 O + NaCl.Aq. 

Cobalt is more easily oxidised than nickel, for chlorine water 
converts the hydrated monoxide into the-sesquioxide, thus : 

2CoO.rcH 2 O -f Cl 2 .Aq + H 3 = Co 2 O 3 .^H 2 O + 2HCl.Aq. 

Hydrated chromium sesquioxide is dissolved by excess of cold 
caustic potash or soda, but is precipitated on warming (see 
below) . 

There are two varieties of chromic salts, which are respectively 
green and violet. Both varieties give with alkalis a grey-green 
precipitate. By varying the conditions, the following hydrates 
have been prepared : 

Cr 2 O 3 .9H 2 O. Grey- violet powder. 

Cr a O 3 .7H 2 O. Greyish-green; soluble in alkali with violet 

Cr 2 O 3 .6H 2 O. Green, gelatinous, drying to a hard black mass. 

Cr 2 O 3 .5H 2 O. Similar to last. 

Cr^O 3 .4HjO. Green ; by boiling chromic chloride and caustic 

Cr 2 O 3 .2H 2 O. Guignet's or Pannetier's green ; produced by heat- 
ing bichrome, K 2 Cr 2 O 7 , and borax. Oxygen is lost, and a borate 
of chromium and alkali is formed. On treatment with water, the 
borate is decomposed, leaving the hydrate. This body is a fine 
green pigment. 

These hydrates dissolve in cold acids, giving violet salts, the 
solutions of which turn green when warmed, most probably owing 
to the formation of a basic salt. Thus chromic sulphate, 
Cr 2 (S0 4 ) 3 .Aq (or Cr 2 3 .3SOj.Aq), when warmed, is supposed to 
give Cr;O.(S0 4 )2.Aq (or Cr 2 Oj.2S0 3 .Aq), losing the elements of 
sulphuric anhydride. 

No native form of chromium hydrate is known. 
Ferric hydrates are found native. Brown or yellow clay iron 
ore is supposed to be the trihydrate, Fe 2 O 3 .3H 2 O or F % e(OH) 3 ; 
xanthosiderite is Fe 2 O 3 .2H 2 O, or Fe 2 O(OH) 4 ; and gb'thite or needle 
iron ore, Fe^O 3 .H 2 O or FeO.(OH). Limonite is 2Fe^O 3 .3HjO ; 
and turgite, 2Fe 2 O 3 .H 2 O. Precipitated hydrate, dried in air, 
possesses the approximate formula Fe 2 O 3 .5H 2 O; when heated, 
water is gradually lost, no sign of formation of intermediate 
hvdrates being found. It is probable that there are many 


hydrates, each of which is stable within a very limited range of 
temperature; hence, on drying, indefinite mixtures are produced.* 
By prolonged boiling in water, the hydrate Fe 2 O 3 .H a O is produced, 
and after a long time the precipitate consists of anhydrous sesqui- 
oxide ; it appears, therefore, that the hydrate may lose water even 
in presence of. great excess of water at 100. Hydrate of iron is 
used as a mordant (see aluminium hydrate, p. 242). It produces 
stains of " iron-mould " on linen ; these can be removed by oxalic 
acid, and a little metallic tin to reduce the iron from sesquioxide 
to monoxide, which is more easily soluble. 

Hydrated manganese sesquioxide occurs native as manganite 
or grey manganese ore ; its formula is Mn 2 Oj.H 2 O. Wad, a 
mixture of oxides of manganese, probably contains some other 
hydrates. Both ferrous and manganous hydrates, suspended in 
water, when shaken with oxygen or air, are converted inlo 
hydrated sesquioxides. That of iron is rust-brown, and of man- 
ganese dark-brown. 

Hydrated sesquioxides of cobalt and nickel are black 
precipitates. That of nickel is said to have the formula 

It is probable that the sesquisulphides, produced by precipita- 
tion, are also hydrated. A green flocculent precipitate is pro- 
duced by addition of a polysulphide of ammonium, (NH 4 ) 2 S W 
(yellow sulphide), to a solution of ferric chloride, to which a small 
quantity of chlorine water or solution of bleaching powder has 
been added. With excess, it is oxidised and dissolved. This 
green precipitate is soluble in ammonia with a green colour, 
possibly giving a double sulphide. Its formula is said to be 
2Pe 2 S 3 .3H 2 O. A cobaltic salt gives, with hydrogen sulphide, a 
dark- grey precipitate of cobalt sesquisulphide, also probably 
hydrated. * No similar nickel compound has been prepared, 
(fc.) No compounds with hydrogen sulphide are known, 
(c.) Compounds of oxides with oxides. As has been stated, 
hydrated chromium sesquioxide dissolves in cold solutions of tha 
hydroxides of potassium of sodium, but is reprecipitated on 
wanning. This behaviour is so far analogous to that of alumi- 
nium hydrate ; the double oxide of aluminium and alkali-metal, 
however, is more stable than that formed by chromium, for its 
golution can be boiled without change. The other hydrates of 
this group are insoluble in alkalis. 

The Spinels. Compounds of these sesquioxides with mon- 
oxides of dyad metals form a very important group of minerals, 
* Chem. Soc., 53, 50. 


crystallising in octahedra, or in rhombic dodecahedra, named 
spinels, the name spinel being generally applied, but being 
specially applicable to the oxide of aluminium and magnesium, 
AUO 3 .MgO. The following is a, list : 

Or 2 O 3 .PeO ; chromite, or chrome-iron ore. A1 2 O 3 ZnO; gohnite. 

Fe 2 O 3 .FeO ; magnetite, or magnetic iron ore. ALO 3 .FeO ; zeilanite. 

Fe 2 O,q.MgrO; magnesio-ferrite. Al 2 O 3 .BeO ; chrysdberyl. 

Fe 2 O 3 .ZnO ; franJclinite. Mn 2 O 3 .ZnO ; hetaerolite. 

AloO 3 .MgrO; spinel. Mn 2 O 8 .MnO; haicsmannite. 

Besides these, Cr 2 O 3 .ZnO, Cr 2 O 3 .CrO, Cr 2 O 3 .MnO, Fe 2 3 .CaO, Co 2 O 3 .CoO, 
and Ni 2 3 .NiO have been made artificially.* 

Chromous hydrate, Cr(OH) 2 , made by addition of caustic 
soda to chromium dichloride, when exposed to air, changes to a 
s nnfp- coloured powder, of the formula Cr^Oj.CrO. It has not 
been obtained crystalline. When iron wire and lime are heated to 
whiteness in presence of air, black crystals of Pe 2 O 3 .CaO are pro- 
duced; the same compound is formed by strongly igniting a 
mixture of hematite and chalk. FranJclinite, Fe 2 O 3 .ZnO, has 
been produced by strongly igniting a mixture of iron sesquioxide, 
zinc sulphate, and sodium sulphate. The zinc oxide remaining 
after decomposition of the sulphate combines with the oxide of 
iron. The sodium sulphate may act as a solvent. Iron, man. 
ganese, and cobalt sesquioxides lose oxygen, the first at a white 
heat, the second at bright redness, the last at a dull-red heat, 
giving these complex oxides. That of iron is the important magnetic 
iron ore, occurring largely in Sweden. Manganoso-manganic 
oxide is a reddish-brown powder, which turns black when heated, 
but recovers its red colour on cooling. Cobaltoso-cobaltic and 
nickeloso-nickelie oxides form grey octahedra with metallic 
lustre. That of cobalt may be produced by heating frhe nitrate 
or oxalate to redness, and boiling the residue with hydrochloric 
acid; and that of nickel by heating nickel dichloride, NiCl 2 , 
to 350 400 in a current of moist oxygen. Manganese dichloride, 
on exposure to moist air, is also changed into the crystalline oxide ; 
and it may also be produced by heating manganous oxide, MnO, 
to redness in water-gas. 

These bodies are also known in a hydrated condition. 

The snuff - coloured powder, obtained as described, from chrom* 
ous oxide in air, is probably hydrated. A dingy green hydrate of 
ferroso- ferric oxide is produced by oxidation of ferrous hydrate ; 
and black hydrates are precipitated by addition of an alkali to a 

Comptes rend., 104, 580. 


mixture in molecular proportions of a ferrous and ferric salt, 
thus : 

FeCl 2 .Aq + 2FeCl 3 .Aq + BKOH.Aq = SKCl.Aq + Fe 3 O 4 .wH 2 O. 

Like the anhydrous oxide, Fe 3 O 4 , these hydrates are magnetic. 
A solution of manganoso- manganic oxide in phosphoric acid gives 
a brown precipitate with potash, doubtless of hydrate. 

A few other double compounds are known, in which the sesqui- 
oxide and protoxide are present in different ratios. Thus, by 
addition of ammonia solution to a solution of a mixture of calcium 
chloride and chromium trichloride, the body Cr 2 O 3 .2CaO is pre- 
cipitated. A Somewhat similar compound, but containing calcium 
chloride in addition, of the formula Fe 2 O 3 .2CaO.CaCl 2 , crystal- 
lises from a solution of iron sesquioxide and lime in fused calcium 
chloride, in shining black prisms. And lastly, by heating a 
mixture of hydrated ferric oxide, potassium carbonate, and potas- 
sium chloride, till the latter is volatilised, ferric oxide, in combina- 
tion with a small quantity of water and potassium oxide, remains 
as transparent red-brown crystals. 

" Smithy scales " are produced by heating iron to redness in 
air. Two layers are formed ; the outer layer has approximately 
the composition Pe 3 O4 ; the inner layer forms a blackish-grey, 
porous, brittle mass, and has the formula Pe 2 O 3 .6PeO. Ferroso- 
ferric oxide is produced also when iron burns in oxygen, when 
iron is heated in water-gas, or when the monoxide is heated in a 
current of hydrogen chloride. 

It is possible to take two views of the constitution of these 
oxides ; the first is that the sesquioxides are chemical individuals, 
derived from the corresponding trichlorides ; and there appears 
little doubt that this is the case with chromium and iron sesqui- 
oxides, Cr 2 Q 3 and Pe 2 O 3 , being easily derived from and convertible 
into Cr^Cl 6 and PesCle respectively. Similarly, their compounds 
with the protoxides would justly have 'the formulas Cr 2 O 3 .CrO and 
Pe 2 O 3 .PeO. But in the case of manganese there appears to be 
some evidence of the existence of two bodies of like formula, but 
of different properties, implying different constitution. There is 
little doubt that the fact that such bodies become much more 
dense find insoluble in acids on ignition, sometimes, indeed, them- 
sglves evolving heat when gently warmed, is due to polymerisa- 
tion, i.e., the association of several simple molecules to form a 
more complex one. But the evidence as regards manganese 
sesquioxide points to a different cause. That body may be regarded 
as either a chemical individual, Mn^O 3 , and the derived manganoso- 


manganic oxide as Mn 2 O 3 .MnO ; or it may be conceived to be 
MnO 2 .MnO, a compound of dioxide and monoxide, or a manganite 
of manganese ; and the substance, Mn 3 O 4 , might be MnO 2 2MnO. 
Now manganese sesquioxide, when treated with dilute nitric acid, 
gives a solution of manganous nitrate, Mn(N0 3 ) 2 , and a residue of 
MnOo. With sulphuric acid oxygen is evolved, and manganous 
sulphate, MnSO 4 , dissolves. The acetate, phosphate, &c., of 
Mn 2 3 can, however, be prepared ; and it is very unlikely that 
such bodies are mixtures of manganous salts and salts of manga- 
nese dioxide ; salts of the latter being almost unknown. On 
addition of alkali to such salts a brown precipitate is produced, 
soluble in acids with formation of salts of the sesquicxide ; whereas 
the hydrated sesquioxide, Mn(OH) 3 , produced by oxidation in air 
of manganons hydrate is split by nitric acid into manganous 
nitrate and insoluble hydrated dioxide, MnO 2 .nH 2 O ; and it is 
insoluble in dilute sulphuric acid. These facts would lead us to 
conclude that two bodies of the formula Mn 2 O 3 exist, one of 
which, however, has the constitution MnO 2 .MnO. The oxides 
would well repay study in this direction. 

(d.) Compounds of oxides with sulphides. Iron sesqui- 
oxide, heated in sulphur gas, gives the compound Pe 2 O 3 .3Pe 2 S3. 
No other compounds of this nature have been prepared in this 

(e.) Compounds of sulphides with sulphides. The follow- 
ing is a list : * 

Cr 2 S 3 .Na 2 S: brick-red powder. Cr 2 S 3 MnS: chocolate-coloured powder. 

CrS 3 .CrS: grey-black powder. Fe^S 3 .Cu s S. Copper-pyrites. 

Cr 2 S 3 .ZnS : dark brown powder. Co 2 S 3 .CoS. Linuceite. 

Cr 2 S 3 .FeS: Daubreelite; black. Ni 2 S.j.NiS. Berychite. 

Iii these compounds, as in the spinels, one metal may replace 
another without reference to atomic weight. If any one molecule 
be considered, it of course possesses a definite formula, such as 
Co 2 S 3 .CoS. But the mineral named nickel-linnce-ite contains some 
Ni 2 S 3 .NiS ; or, perhaps, Co 2 S 3 .NiS, along with the former. The 
atomic ratio of metal to sulphur is a constant one, but as these 
bodies have the same crystalline form, and as their molecules 
occupy approximately the same volume, they can replace one 
another in any crystal. The usual way of denoting such replace- 
ment in any proportion is. to write the formula, for example, of 
nickel-linnceite, thus : (Co,Ni) d S 4 . 

?n. Akad. Ser. (2), 81, 531 j Monatsh. f. Chem., 3, 266. 


The same peculiarity is noticeable in the spinels, where alu- 
minium, chromium, iron, and manganese may replace each other 
as sesquioxides, and beryllium, magnesium, zinc, &c., as monoxides. 
This will be again referred to in treating of the silicates. 

The double sulphides which have been prepared artificially 
have been obtained by passing hydrogen sulphide over a heated 
mixture of the hydrates of the respective metals ; thus, a mixture 
of chromic hydrate and zinc hydrate thus treated, gives a mass 
which, when boiled with hydrochloric acid, leaves a dark brown 
powder of the formula Cr 2 S 3 .ZnS. 

More complex biilphidesof iron are found native, and are generally 
termed magnetic pyrites. They have the formulas FeJ3 3 .3MS, 
Fe>S 3 4MS, Fe.Sj.SMS, and Fe 2 S 3 .6MS, M representing iron, 
cobalt, or nickel. They form yellow crystals with metallic lustre, 
Copper pyrites, barnhardiite, and chalcopyrrhotite are sitnilar 
bodies, containing copper, and have respet lively the formula) 
Fe,S 3 .Cu,S, Fe 2 S a .2Cu 2 S, and Fe,S 3 .2CuS.FeS. Purple copper ore 
is a similar compound of uncertain formula. By fusing iron with 
sulphur and potassium carbonate, purple-brown needles, of the 
formula KFeS. 2 , are formed. By ignition in hydrogen it yields 
2FeS.K 2 S. 

(/.) Compounds with halides. These bodies, as usual, are 
formed either by evaporating or heating an aqueous solution of the 
trichlorides, or by heating a mixture of chloride and hydrate. 

The following have been prepared : 

Cr,0 3 .4CrCl 3 8H 2 , and 3H 2 O. 
Cr 2 O 3 2CrCl 3 2H 2 O. 
Cr 2 O J .CrCl J 3H 2 O. 

The corresponding compounds of iron are not so definite. 
Weak solutions of ferric chloride, when heated, give 1, soluble 
ferric hydrate and hydrogen, chloride, which recombine slowly on 

2. From stronger solutions mixtures of oxychlorides separate. 

3. At high temperatures the ferric hydrate loses water, and 
ferric oxide is deposited. 

Dafk red plates, of the formula 9Fe 2 O 3 .FeCl 3 , separate from a 
strong solution of ferric hydrate in ferric chloride on evaporation 

in vacuo. 

Oxychlorides are also produced when solutions of ferrous 
chloride are exposed to air. 

Oxychlorides of manganese, cobalt, and nickel are unknown. 



III. Dioxides and disulphides. 

List. Chromium. Iron. Manganese. Cobalt. Nickel. 
Oxygen..,. CrO 2 . MnO 2 . (CoO 2 ).* (NiO 2 ).* 

Sulphur ... PeS 2 . NiS 2 . 

Sources. Manganese dioxide, or pyrolusiie (from nap, fire, 
and Xveti/, to loose, referring to its action in removing the green 
and brown tints of glass coloured by iron, owing to the comple- 
mentary action of its purple colour), is one of the chief ores of 
manganese. It is an iron-black or grey mineral, very hard, and 
somewhat brittle, with fibrous texture. It is largely employed for 
making chlorine. 

Nodules containing manganese dioxide are of common occur- 
rence on the sea-bottom ; they have been dredged from the bed of 
the Pacific and Atlantic Oceans, and are found in the Firth of 

Iron pyrites or wiundiCy FeS 2 , is a golden-yellow mineral 
crystallising in cubes. It is very hard and brittle, and was 
formerly used as a m-eans of striking fire, whence its name. 
Marcasite is a whitish mineral with metallic lustre, of the same 
formula, crystallising in the trimetric system. Both of these 
minerals occur in slate, coal, shale, &c. They oxidise on exposure 
to moist air, and furnish the sulphuric acid necessary for alum in 
alum shale. They are used as a source of sulphur. 

Preparation. 1. By direct union. Hydrated chromium 
sesquioxide, Cr 2 O 3 .t/H 2 O, heated in air to 200, is oxidised to the 
hydrated dioxide, CrO 2 .H 2 O ; the hydrated compound is dried at 
253. Iron and sulphur combine below redness to form FeS 2 ; and 
lower sulphides of iron unite with sulphur when gently heated in 
a current of hydrogen sulphide. 

2. By heating a compound. The hydrated dioxides can be 
dried at 200 250, yielding the dioxides. Chromium nitrate, 
when heated, yields the dioxide, and manganese dioxide is produced 
by heating manganous nitrate, or manganous carbonate and potas- 
sium chlorate. The oxygen is derived from the nitric anhydride, 
or from the potassium chlorate. 

3. By double decomposition. Oxides of iron, heated in 
hydrogen sulphide to above 100, are converted into disulphide; 
and an alkaline poly sulphide reacts with ferrous chloride or 
sulphate at 180, yielding disulphide. Nickel disulphide is 

* In combination, as 2CoOj.CoO and 2NiO 2 .NiO. See also Cobalt-amines. 


produced by heating a mixture of nickel carbonate, potassium 
carbonate, and sulphur to dull redness. 

Properties. Chromium dioxide is a black powder, giving 
off oxgen at 350. It is insoluble in water, but soluble in acids, 
and reprecipitated from its solution as hydrate by ammonia. 
Manganese dioxide is a black powder when prepared artificially. 
It dissolves in strong sulphuric acid, yielding a yellow sulphate, 
MnO 2 .2SO 3 . On diluting this solution, the hydrated dioxide is 
precipitated. Iron disulphide when prepared artificially is a 
black powder, or sometimes yellow cubes like the native form, 
insoluble in acids ; and nickel disulphide is a steel-grey powder. 
Double compounds, (a.) With water. Hydrated chromium 
dioxide is produced, as before mentioned, by the spontaneous oxida- 
tion of the hydrated sesquioxide at 200; and also by reducing 
chromium trioxide or its compounds. Thus, by passing a current 
of nitric oxide, NO, through a dilute solution of potassium 
dichromate, K 2 Cr 2 O 7 , it is deprived of part of its oxygen, and 
gives a flocculent brown precipitate of the hydrated dioxide. The 
reduction may be effected by ammonia, as when a solution of am- 
monium dichromate, (NH 4 ) 2 Cr 2 O 7 .Aq, is boiled, the oxygen going 
to oxidise the hydrogen of the ammonia ; or by means of a chromic 
compound, e.g., by heating together a solution of chromium tri- 
chloride, CrCl 3 , with potassium dichromate, K 2 Cr 3 O 7 or K 2 O.2Cr0 3 ; 
chromium hydrate may be supposed to be formed by the action of 
water on chromium trichloride, thus : 2CrCl 3 .Aq + 3H 2 = 
O 2 O 3 .Aq + GHCl.Aq; and the hydrate then acts on the trioxide 
combined with potassium oxide in the dichromate, thus: 
Cr 2 3 .Aq + CrG 3 .Aq = 3O0 2 .nH 2 0. The complete equation is : 
4CrCl 3 .Aq + 5H,0 + K,Cr 2 7 .Aq = 2KCl.Aq + 6CrO 2 .nH 2 O + 
lOHCl.Aq. Heat alone expels oxygen from chromium trioxide, 
but the resulting substance is said to be 3CrO 2 .Cr 2 O s . Oxalic 
acid, H 2 C 2 O 4 , or alcohol may also be used to effect the reduction. 

It is still a question whether this body is not a chromate of 
chromium, Cr0 3 .Or 2 O 3 . Against this view, it may be stated that 
while chromates, when distilled with sodium chloride and strong 
sulphuric acid, give chromyl dichloride, Cr0 2 Cl 2 (see p. 268), 
this substance does not do so ; and that it dissolves in acids as a 
whole'and is reprecipitated by alkalis, as it would be, were it a 
definite individual. Yet, on boiling with alkalies, hydrated 
chromium sesquioxide is precipitated, and the trioxide combines 
with the alkali, forming a chromate. 

The compounds MnO 2 .2H 2 O, MnO 2 .H 2 O, 2MnO 2 .H 2 O, 
3MnO 2 .H 2 O and 4MnO 2 .H 2 O are known. They are all 

3 2 


brownish-black or black powders. The last of these is produced 
by treating Mn 3 O 4 or Mn 2 O 3 with strong nitric acid, whence 
the conclusion that these bodies are compounds of MnO 2 with 
2MnO and MnO respectively. The monohydrate, MnO 2 .H 2 O, is 
formed by the spontaneous decomposition of a solution of potassium 
permanganate, KMn0 4 or K 2 O.Mn 2 O 7 , or by the action of chlorine 
on manganous carbonate suspended in water. The compound 
2MnO 2 .H 2 O is precipitated by addition of potassium hypochlorite 
to a manganous salt in presence of excess of ferric chloride ; and 
the compound SMnO^.H^O by evaporating a solution of manganous 
bromate. The dihydrate, MnO 2 .21LO, is precipitated on addition 
of water to the sulphate Mn0 2 .2SO 3 ; the existence of this 
sulphate appears to lend support to the theory that the dioxide is 
a chemical individual, and not a manganate of manganese, 
MnO j.MnO. It need hardly be pointed out that the molecular 
weights of all these bodies are unknown. 

(fe.) Double oxides. Several manganese compounds are 
known, viz : MnO 2 .MnO, MnO,.CaO, 2(MnO>).K,O, 
2(MnO 2 ).CaO. These substances are formed by the action of 
air on (1) warm hydrated manganese monoxide precipitated from 
the dichloride MnCl 2 by its equivalent cf calcium hydrate ; (2) by 
the same process, twice the equivalent of lime being added, 
thus : MnCl 2 .Aq 4- 2Ca(OH) 2 + = MnO 2 .CaO + CaCl 2 .Aq 
-h 2H 2 O, and (3) by the action of manganese dichloride on the 
former compounds, thus : 

2(MnO 2 .CaO) + MnCl.Aq = 2(MnO 2 ).CaO + Mn(OH), + 

CaCl 2 Aq. 

These bodies are all hydrated, but the amount of com- 
bined water is unknown. Their formation is the principle of 
u Weldon's manganese-recovery process " whereby manganese 
dioxide which has been used for the manufacture of chlorine, and 
converted into dichloride, is restored to the state of dioxide, and 
thereby again rendered available for preparing chlorine (see p. 75). 
Such bodies as MnO^.CaO are termed manganites. The com- 
pound 2(MnOj).K 2 O is a black powder; others containing less 
oxide, e.g., 12(MnO,).Na 2 O, 5(MnO 2 ).Na,O, &c., are produced by 
heating manganous chloride with sodium hydrate and "Sodium 

Compounds of the formula 5 (MnO 2 )M"O, where M" stands for 
calcium, strontium, zinc, or lead, may be produced by heating 

* Bull. Soc. Chim. (2), 30, 110; Dingl. polyt. Jour., 129, 51 j Chem. Soc. J., 
37, 22; 591 j Compt. rend., 101, 167; 103, 261. 


chlorides of these metals with potassium permanganate. They 
form black crystals. At higher temperatures, 2(MnO 2 ).M"O and, 
at still higher, MnO 2 .M"O are produced. 

Similar cobalt and nickel compounds, 2(CoO 2 ).CoO and 
7CoO 2 4CoO (with water of hydratioii from 4H,0 to H,O), also 
3NiO 2 .5NiO.9H 2 O are produced by adding sodium hypochlorite, 
NaCIO, to a mixture of the hydrate of cobalt or nickel and excess 
of soda. A cobalt compound of the formula 

3(2CoO 2 .3CoO).K 2 O.3H 2 O 

is produced by heating the monoxide with caustic potash in 
presence of ir. No doubt, double compounds wilh other metals 
could be prepared. 

(c.) Oxyhalides. An oxy fluoride of the formula MnO 2 .MnP 4 
or MnOF 2 is said to be produced by adding manganese tetra- 
chloride to a boiling solution of potassium fluoride. It is a rose- 
coloured powder, and combines with potassium fluoride, forming the 
compound MnOP 2 .2KP. The trifluoride is said to yield similar 
double salts, e.g., Mn 2 P 4 O.4KP. These bodies are produced by 
treating potassium permanganate, KMn0 4 , with aqueous hydrogen 

IV. Trioxides. (a.) The only trioxides known in the free state 
are chromium trioxide, or chromic anhydride, CrO 3 , and man- 
ganese trioxide, MnO 3 . Iron trioxide exists in combination 
with potassium monoxide in potassium ferrate, and that of man- 
ganese in potassium manganate. 

Preparation. By double decomposition. Chromyl fluoride 
(see p. 268), led into a crucible slightly damp and loosely covered 
with damp paper, reacts with the water, depositing long needles 
of the trioxide ; thus : 

Cr0 2 F 2 + E,0 = CrO 3 

By the action of sulphuric acid on a chromate, a sulphate and 
chromic anhydride are formed. On pouring 1 volume of a saturated 
solution of potassium dichromate, K 2 Cr 2 O 7 , into 1 J volumes of strong 
sulphuric acid, long needles of chromic anhydride hydride deposit on 
cooling. They are difficult to free from sulphuric acid ; and the 
presgnt method of preparing the trioxide for commercial use is b) 
adding to strontium chromate exactly enough sulphuric acid tc 
precipitate the strontium as sulphate, to decant the solution o 
trioxide from the insoluble sulphate, and to evaporate to drynessJ 

Manganese trioxide is obtained by dropping a solution o 

* The author has tried the process, but without success. 


potassium permanganate in strong sulphuric acid on to sodium 
bicarbonate ; MnO 3 is liberated, and is carried on in the solid state 
by the carbon dioxide. 

Properties. Chromium trioxide forms a red crystalline 
powder, a mass of loose woolly crystals, or scarlet crystals. It 
melts at 190, and begins to decompose at 250, losing oxygen. It 
is soluble in water,, and the solution contains chromic acid, 
CrO 3 .H 2 0, or H 2 Cr0 4 ,* or H 2 Cr 2 O 7 . Its compounds with other oxides 
are called chromates. The blue solution obtained by shaking a 
dilute solution of chromium trioxide with hydrogen dioxide, and 
extracting with ether, is said to be a compound of the formula 
Cr0 3 .H 2 2 . On evaporation of the ether, it remains a$ a blue oil.f 

Manganese trioxide is a reddish, amorphous, deliquescent sub- 
stance, unstable at the ordinary temperature. 

(6.) Compounds with other oxides. Chromates, ferrates, 
and manganates. Of these the chromates are the most stable, and 
have been best investigated. They may be divided in to four classes : 

1. Basic chromates ; those in which the number of atoms of 
oxygen in the base exceeds one-third of that in the chromic anhy- 
dride. These compounds are orange, red, or brown in colour. 
They are produced by double decomposition, a solution of a soluble 
chromate, such as potassium chromate, K 2 CrO 4 , being added to a 
soluble salt of the metal ; in such a case, uncombined chromic 
anhydride exists in solution; or by digesting a chromate, such as 
PbCrO 4 = PbO.CrOjwith alkali, or with excess of base. They 
are as follows : 

Ratio, 3 : 9 : CrO 3 .3Bi 2 3 . Ratio, 6 : 9 : 2CrO 3 .3Bi 2 O 3 . 
3:4 . CrO 3 .4ZnO,3H 2 O : CrO, } 4CuO. 
3:3: CrO 3 Al 2 O 3 ; CrO 3 .Cr 2 O 3 (?)(this body is CrO 2 ) ; CrO 3 .Fe 2 O 3 ; 

CrO 3 .3NiO.3H 2 O; CrO 3 .Bi 2 O 3 ; CrO 3 .3CuO; CrO 3 .3HgO. 
6:5 : 2CrO 3 .5NiO.12H 2 O. Ratio, 21 9 : 7CrO 3 .3Bi 2 O 3 . 
3:2 : Cr0 3 .2ZnO.H 2 O ; CrO 3 .2CdO.H 2 O ; CrO 3 .2MnO.2H 2 j 

CrO 3 .2CoO.2H 2 O ; CrO 3 .2NiO.6H 2 O ; CrO 3 .2PbO; 

CrO 3 .2CuO; CrO 3 .2HgrO ; CrO 3 .2Hgr 2 O; 

2CrO 3 .3CuO.K 2 O.3H 2 O. 
6 : 3 : 2Cr0 3 .3PbO ; 2CrO 3 .Bi 2 O 3 ; 2CrO 3 .CuO.2PbO. 

These compounds are orange, red, or brown powders, and are 
insoluble in water, or nearly so ; they dissolve in acids, being* con- 
verted into chromates containing a larger proportion of trioxide^ 
of chromium. The most important of them is the chromate 
CrO 8 .2PbO; it is named "chrome-red " or " Persian-red. " It is pro- 

* Comptes rend., 98, 1581. 
f Ibid., 97, 96. 


duced by addition of lead oxide to the monoplumbic chromate, 
PbCrO 4 , or CrCX.PbO; or with a purer shade by heating that 
body with potassium nitrate ; the potassium oxide withdraws 
chromic anhydride, and on washing with water, excess of potas- 
sium nitrate and chromate are withdrawn and the basic chromate 
is left as a red powder. Cloth on which a precipitate of yellow 
PbCrO 4 has been formed, may be changed to a brown-red by 
plunging it into a bath of boiling milk of lime (CH,(OH) 2 .Aq), 
which withdraws half the chromic anhydride. 2CrO 3 .3PbO 
occurs in scarlet crystals as melanochwite ; and 2CrO 3 CuO.2PbO 
as a yellowish-brown mineral named vauquelinite. 

2. The second class of chromates is often termed " neutral." 
This name was originially applied to those substances incapable of 
affecting the colour of litmus. But most of these chromates are 
insoluble ; moreover, the typical " neutral " chromate of potassium, 
K 2 Cr04(Cr03.K 2 0) has an alkaline reaction arid turns red litmus 
blue. It is better therefore to discard the misleading name. 
The oxygen of the chromic; anhydride bears to that of the base 
the ratio 3:1. The following is a list : 

Eatio 3-1. 

CrO.,.H 2 O (chromic acid) ; CrO 3 .Li 2 O.2H 2 O; CrO.j.NasO.lOHaO (crystal- 
lised above 30, this body is anhydrous) ; CrO 3 .K 2 O ; CrO 3 .K(NH 4 )O 
(= KNH 4 CrO 4 ); CrOjMgrO.Tl^O ; CrO (J .CaO.4H 2 O; CrO 3 .SrO ; 
CrO.,.BaO; CrO 3 .Tl 2 O; CrO^.PbO; CrO^.CuO; CrO 3 .Agr 2 O; Cr 

Of these, the hydrogen, lithium, sodium, potassium, magnesium, 
calcium, copper, and mercuric compounds are soluble in water. 
Hydrogen chromate, HoCrO 4 , is produced by dissolving chro- 
mium trioxide in water, and cooling with melting ice. It forms 
small red deliquescent crystals, which readily part with water. 
Potassium chromate, K^CrO4, has a light-yellow colour, and a 
bitter, cooling taste ; it is exceedingly poisonous ; it is insoluble in 
alcohol, but soluble in water (100 grams of water at 15 dissolve 
48*3 grams of chromate). Lt melts at a low red heat, and crystal- 
lises in double hexagonal pyramids. Strontium chromate, 
SrCrO 4 , is sparingly soluble. It is the one from which chromic 
anhydride is now made commercially by addition of sulphuric 
acid. * It is found to be the only available chromate from which 
the chromium trioxide is completely expelled by its equivalent of 
sulphur trioxide, by the action of sulphuric acid ; hence its use. 
Barium chromate, BaCrO 4 ,is an insoluble yellow powder, used as 
a pigment under the name, u yellow ultramarine." Lead chromate, 
PbCrOi, is found native, as red lead-ore or crocomte. It crystal- 


lises in monoclinic prisms. It is a translucent yellow body, and 
occurs in decomposed granite or gneiss. Prepared by addition of 
potassium chromate or dichromato to a soluble salt of lead, it is a 
yellow powder, and is known as " chrome-yellow " and used as a 
pigment. It fuses to a brown liquid, and solidifies to a brown- 
yellow mass. It is made use of in estimating carbon and hydrogen 
in carbon compounds. It is practically insoluble in acids, but 
dissolves easily in potassium hydrate, forming chromate and plum- 
bite of potassium. Silver chromate, Ag 2 CrO 4 is a deep-red 
precipitate, crystalline in structure ; the individual crystals trans- 
mit green light. 

3. Bichromates. These bodies are often called " acid " 
chromates, and their solutions have an acid reaction with litmus. 
They are produced by adding some acid, e.g., chromic acid, or more 
of fcen nitric acid to the monochromates. They are as follows : 

KatioGrl. 2CrO 3 .Li 2 O; 2CrOj.Na2O.2H2O; 2CrO 3 .K 2 O ; 2CrO 3 .(NH 4 ) 2 O ; 
2CrO 3 .CaO.3H 2 O; 2CrO 3 .BaO ; 2CrO 3 .TLO ; 2CrOj.PbO ; 
2CrO 3 .Agr 2 O. 

The most important of these is potassium dichromate, or 
" bichrome," which is prepared on a manufacturing scale. It 
is produced by acidifying the monochromate, K 2 Cr0 4 , with sulph- 
uric acid, thus: 2K 2 Cr0 4 .Aq + H 2 S0 4 .Aq = K 2 S0 4 .Aq + 
K 2 Cr 2 7 . Aq ~h H 2 0. It forms deep orange-red tables or prisms. It 
is insoluble in alcohol, but soluble in water (100 grams dissolve at 
20 12'4 grams of bichrome) . It melts at a dull red heat, and decom- 
poses at a white heat into potassium chromate, chromium sesqui- 
oxide, and oxygen. It is affected by light, and has the curious 
property of rendering gelatine impregnated with it insoluble in 
water after exposure to light, and it thus finds an application 
in photography. It is largely used as an oxidising agent, and for 
making chrome-yellow, &c. 

The dichromates are decomposed by much water, excepting 
those of sodium, potassium, and ammonium. 

The name anhydrochromates is sometimes applied to these bodies, 
the view being taken that they are compounds of monochromate 
and anhydride, thus : K 2 CrO 4 .CrO 3 . 

4. Polychromates; tri-, tetra-, &c. 

Ratio 9 :2. 3CrO 3 .2ZnO, soluble, crystalline; 3CrO.<.2Tl 2 O. 
Batio9:l. 3CrO 3 .K 2 O; 3CrO 3 .(NH 4 ) 2 O; 3CrO 3 .Tl 2 O. % 

These bodies are deep-red crystals, formed on crystallising the dicbromates 
from strong nitric acid. 

Batio 12 : 1. 4CrO 3 .X2O, similarly prepared. The polychromates decom- 
pose on treatment with much water. 


Ferrates. Of these, only the potassium, sodium, and barium 
salts are known. Their formulae are supposed to be Fe0 3 .K 2 ; 
Fe0 3 .Na 2 O j and Fe0 3 .BaO ; but the potassium and sodium salts 
are stable only in presence of a large excess of alkali, and the 
barium salt has not been analysed. The ratio of oxygen to iron in 
the iron trioxide has been determined ; hence the deduction of 
the formula, Fe0 3 . 

Sodium or potassium ferrate is formed by heating iron-filings 
and sodium or potassium nitrate to dull redness ; by igniting iron 
sesquioxide with sodium or potassium hydrates in an open crucible, 
better with addition of sodium or potassium nitrate ; by passing 
chlorine thrcTugh a very strong solution of sodium or potassium 
hydrates in which ferric hydrate is suspended ; the ferrate, being 
insoluble in the strong alkali, is precipitated as a black powder ; 
and by electrolysing a strong solution of potash or soda with iron 
poles ; the ferrate crystallises on the positive pole. The produc- 
tion of ferrate may be shown as a lecture experiment by adding 
a few lumps of potassium hydrate to some solution of ferric chloride, 
and adding bromine and warming. 

The potassium ferrate may be dried on a porous plate ; it 
cannot be filtered through paper, as it at once loses oxygen. It 
forms a fine cherry-red solution, but it soon decomposes with loss 
of oxygen. Barium ferrate is a purple precipitate produced by 
adding a solution of barium hydroxide to the solution of potas- 
sium ferrate. The ferrates at once lose oxygen on addition of 
an acid. 

Manganates. Of these, only the sodium, potassium, calcium, 
and barium salts are known. They arc prepared by heating man- 
ganese dioxide with sodium or potassium hydroxides or carbonates ; 
manganate and a lower oxide of manganese are formed ; or nitrate 
or chlorate* of calcium or barium with manganese dioxide. The 
yield may be increased by adding sodium or potassium nitrate to 
the hydroxides. On treatment with cold water, they form a deep 
green solution, and when it is evaporated in a vacuum, crystals 
are deposited. These crystals have the formula K 2 MnO 4 = 
MnO 3 .K>O; the barium, calcium, and sodium manganates are 
supposed to have similar formulas. On leaving a strong solution 
of pofassium manganate exposed to air, crystals of dimanganate 
tuve been formed, 2MnO3.K2O.H2O, the carbon dioxide of the air 
having withdrawn half the potash. 

Potassium manganate is stable only in presence of excess of 
alkali, and is decomposed by pure water with formation of per- 
manganate and dioxide, thus : 


3(MnO,.K,0) + 2H 2 + Aq = Mn 2 O 7 .K 2 O.Aq -f 4KOH.Aq 4- 
MnO 2 .nH 2 O. 

Owing to this change of colour from green to purple, the old 
name for potassium manganate was u mineral chameleon." 

Manganate of barium is known as " baryta-green." Potassium 
manganate having been produced by gradually adding manganese 
dioxide to a fused mixture of two parts of potassium hydrate and one 
part of potassium nitrate, the cooled mass is treated with water 
and filtered. On addition of barium nitrate to the filtrate, a violet 
precipitate of barium manganafce is produced, which is heated to 
redness with solid barium hydrate till it assumes a bright-green 
colour. It is then treated with water to remove barium hydrate. 
The green colour is in all cases probably due to basic man- 

Perchromates and permanganates. These bodies are com- 
pounds of oxides with the heptoxides of chromium or manganese, 
Cr 2 O 7 , or Mn 2 7 . Those of chromium are very unstable, if, indeed, 
they are capable of existence. If hydrogen dioxide, H 2 2 , be 
added to a solution of chromic acid, or of potassium chromate and 
sulphuric acid, a dark-brown colour is produced. On shaking the 
solution with ether, the upper layer of ether has a fine blue colour; 
on evaporation at 20, a deep indigo-blue oily liquid is left; this 
is possibly perchromic anhydride, or chromium heptoxide, Cr 2 O 7 , 
but is also said to be a compound of the formula CrO 3 .H 2 0>. 
Its salts are unknown. This reaction affords a very delicate 
test both for chromium trioxide and for hydrogen dioxide (sec 
p. 197). 

Potassium permanganate, Mn ? .O 7 .K,O, or KMnO 4 , is pro- 
duced by acidifying potassium manganate. It may be supposed 
that the manganic acid, Mn0 3 .H 2 0, decomposes at the^moment of 
its liberation, yielding manganese dioxide and permanganic acid, 
thus : 3Mn0 3 .H 2 O.Aq = Mn 2 7 .H 2 0. Aq -f- MnO 2 .?/H 2 O. The same 
change is produced by boiling a solution of potassium manganate, 
or by treating sodium manganate with magnesium sulphate, thus : 
3(Mn0 3 .lSra 2 0)Aq + 2(S0 3 .MgO)Aq + 2H 2 O = Mn 2 O 7 .Na 2 O.Aq + 
2(SO. } .Na 2 0)Aq -f 2(MgO.H 2 O) -f MnO 2 .wH 2 O; magnesium man- 
ganate being unstable. Manganate may also be converted into 
permanganate without separation of dioxide by means of chlorine, 
thus: 2(MnO 3 .K 2 O)Aq + C7 a = 2KCl.Aq -f Mn 2 7 .K 2 O.Aq. % 

The following permanganates are known : 

Mn 2 O 7 .H 2 O.Aq, or HMnO 4 .Aq; Mn 2 O 7 .K 3 O, or KMnO 4 ; 
NH 4 MnO 4 ; Ba(MnO 4 ) 2 ; Pb(MnO 4 ) 2 ; and AgMnO*. 


The barium salt is made by the action of carbon dioxide on 
barium manganate ; and from it the free acid, HMn0 4 , maybe 
separated by addition of monohydrated sulphuric acid, H 2 SO 4 .H a O. 
It forms a greenish -yellow solution, and deposits slowly a dark, 
reddish-brown liquid, not solidifying at 20; it is said to be 
manganese heptoxide, or permanganic anhydride, Mn 3 7 . It is 
non-volatile.* This liquid dissolves in strong sulphuric acid with 
a yellow-green colour ; it explodes when strongly heated. The 
yellow-green solution contains (Mn0 3 ) 2 S04. On adding water the 
colour changes to violet that of permanganic acid. 

The silver and lead salts are formed by adding soluble salts 
of silver or ?ead to potassium permanganate. They are dark- 
coloured precipitates. The ammonium salt is made by mixing 
the silver salt with ammonium chloride. Potassium permanganate, 
with excess of potassium hydrate, turns green with formation 
of manganate, oxygen being evolved, thus : 2KMn0 4 .Aq -f- 
2KOH.Aq = 2K 2 MnO 4 .Aq + H 2 -f 0. 

Potassium permanganate forms dark-red, almost black, crystals, 
with greenish reflection ; its solution is sold as a disinfectant under 
the name of " Condy's fluid," and has a splendid purple colour. 
The dichromate and permanganate of potassium are used as means 
of oxidising substances in presence of water. Bichrome does not 
readily part with its oxygen, even to an easily oxidisable body, 
unless an acid be present ; when it does, chromium dioxide is pro- 
duced. Thus : 

2Cr0 3 .K 2 O.Aq + H 2 = 2CrO 2 .nH 2 O + 2KOH.Aq + 20, 
Potassium permanganate acts similarly, thus : 

Mn 2 O 7 .K 2 O.Aq + H 2 = 2MnO 2 .nH a O + 2KOH.Aq + 30. 

In presence 'of an acid (usually sulphuric acid) a salt of chromium 
or manganese is produced, thus : 2CrO 3 = Cr 2 3 -f 30 ; and 
Cr a O 3 + 3H 2 S0 4 = Cr,(S0 4 ) 3 + 3H,O. AlsoMn 2 7 = 2MnO + 
5O ; and MnO -f H,S0 4 = MnSO 4 + H 2 0. The complete equa- 
tions are : 

K 2 O 2 7 .Aq + 4H 2 S0 4 = K 2 S0 4 .Aq + Cr 2 (S0 4 ) 3 .Aq + 4H 2 + 30; 
and 2KMnO 4 .Aq + 3H 2 S0 4 .Aq = K 2 S0 4 .Aq + 2MnS0 4 .Aq 
+ 3H 2 + 50. 


The oxygen, being in the nascent or atomic state, is available for 
oxidation of compounds of carbon, &c. 

* See Chem. Soc., 53, 175. 


Compounds of oxides with halides. These are as fol- 
low: Cr0 2 Cli, chromyl dichloride; Cr0 2 F 2 , chromyl difluoride, and, 
possibly, Mn0 2 Cl 2 ,manganyl chloride. They are formed by distilling 
a mixture of sodium chloride or fluoride, potassium dichromate or 
permanganate, and strong sulphuric acid. The reaction takes 
place between the liberated chromium irioxide or manganese 
heptoxide and the hydrogen halide, the sulphuric acid combining 
with the water produced, which would otheiwise decompose the 
chromyl or manganyl halide, thus : 

Cr0 3 + 2HC1 + H,S0 4 = CrO,Cl 2 + H,SO 4 .H 2 O. 

Chromyl dichloride may indeed be obtained by the t direct action 
of dry hydrogen chloride on pure chromium trioxide. Hydrogen 
bromide and iodide are decomposed with liberation of bromine or 
iodine. Chromyl chloride is a deep-red liquid, closely resembling 
bromine in appearance; it boils at 118, and gives a deep-red 
vapour. It mixes in all proportions with carbon disulphide and 
with chloroform. The manganese compound is said to be a purple 
vapour, condensing at a very low temperature ; but it requires re- 
investigation. Chromyl fluoride may be made by a similar pro- 
cess. Chromyl chloride reacts with water, forming chromium 
trioxide or chromic acid, thus : 

.01 , H.OH _ .OH , HC1 

As its vapour- density shows it to have the formula Cr0 2 CL>, it is 
concluded that chromic acid is analogously constituted, and may 
be represented by the structural formula CrO a (OH) 2 , and chro- 
mates as Cr(X(OM') 2 . It is obvious that an intermediate com- 
pound between Cr0 2 Cl 3 and Cr0 2 (OM') 2 should exist of the 

formula Cr0 2 <[p| . Such a body is known, and is termed a 

chloro-chromate. The potassium salt is produced by saturating a 
hot solution of dichromafe of potassium with hydrogen chloride 
and leaving it to crystallise. Flat rectangular prisms of the 

compound CrO^^i are deposited ; on treatment with water they 

decompose. The mercuric salt is also known. Compounds have 
also been prepared of the formulae 2CrO 3 .KF and 2CrO 3 .lTH 4 F ; 
they are produced by adding aqueous hydrofluoric acid to potas- 
sium or ammonium dichromates ; they may be constituted thus : 


On heating chromyl dichloride to 180190 in a sealed tube 
chlorine separates, and the compound Cr 3 Cl 2 O fl remains as a black 
powder. Its constitution may be thus represented : 

Cl CrO, Cr0 2 CrO a Cl. 

The corresponding potassium salt is produced by saturating potas- 
sium chlorochromate with ammonia, and the ammonium salt by 
saturating a solution of chromyl dichloride in chloroform with 
ammonia. Their constitutional formula may be : 

KO CrO, Cr0 2 CrCV-OK ; and 

(NH 4 ) 0- 00, Cr0 2 - Cr 2 (NH 4 ) . 

Such constitutional formula) will be further referred to in treating 
of silicates, phosphates, and sulphates. 

No compounds of bromine or iodine analogous to chromyl chlo- 
ride are known ; bromine or iodine are invariably liberated. Tho 
volatility of the chlorine compound serves to identify chlorine in 
presence of bromine or iodine ; on distilling a mixture of halides 
with bichrome and sulphuric acid, if chromium is found in the 
distillate, the presence of chlorine in the mixture is proved. 

Physical Properties. 1. Weight of I cubic centimetre. 

MnO, 5-1 ; CoO, 5-6 ; NiO, 5'6 ; 6'8 (crystallised). 

FeS, 4-8; MnS, 4'0 ; NiS, 5'6. 

Cr 2 Oj, 6'2 (crys.); Fe 2 O.j, 5'3 (native); Mn 2 3 (braunite), 4'75; Co 2 O 3 , 48. 

Cr 2 S 3 , 4-1 ; Fe 2 S 3 , 4'3 ; Co 2 S 3 , 4'8. 

Ni 2 3 , 4-8. 

Cr 5 O 9 , 4-0 ; Fe a O 4 5'12 (magnetite) ; Mn d O 4 , 4'85 (native) ; Co.>O 4 , 6'3. 

MnO 2 , 4 83 ; 

MnS 2 , 346j FeS 2 , 5'04 (pyrites); 4'8 (marcasite). 

CrO.,, 2'8. 

Heats of formation. 

Cr 2 O 3 + *3O = 2CrO 3 + 143K. Mn -f O + 2H 2 O = Mn(OH) 

+ 948K. 
Fe + O + H 2 O = Fe(OH) 2 + 683K. Mn + 2O + H 2 O = MnO 2 .H 2 O 

+ 1164K. 
2Fe + 30 + 3H 2 O = Fe(OH) 3 + 1912K. Mn + S + wF 2 O = MnS.wH 2 O 

+ 444K. 
3Fe + 40 = Fe 3 O 4 + 2647K. Co + + H 2 = Co(OH)o 4- 


F + S + IT 2 O =- FeS.wH 2 O + 238K. 

2Co(OH) 2 + + H 2 O = Co 2 3 .3H 2 O Co + S + H 2 = CoS.H,O 
+ 223K + 197K. 

Ni + O + H 2 = Ni(OH) 2 + 


2Ni(OH) 2 + + H 2 = Ni 2 3 .3H 2 O Ni + S + H 2 O = NiS.wH 2 O 
+ 13IT. + 174K. 




Oxides, Selenides, and Tellurides of Carbon, 
Titanium, Zirconium, Cerium, and Thorium. 

This group gives representatives of monoxides, sesqui oxides, 
dioxides, and peroxides. The monoxides show little tendency 
towards combination ; the dioxides form compounds with the 
oxides of other elements, which are named carbonates, titanates, 
and zirconates. Some similar compounds of the sulphides have 
also been prepared. 

Carbon. Titanium. Zirconium. 

Oxygen.... CO; CO 2 . TiO ;* Ti 2 O 3 ; TiO 2 ; TiO 3 . ZrO 2 ; Zr 2 O 6 . 
Sulphur ... OS ; CS 2 . TiS ; Ti 2 S 3 ; TiS 2 . ZrS 2 P 

Cerium. Thorium. 

Oxygen Ce 2 O 3 ; CeO 2 ; CeO. } . ThO 2 ; T^O;. 

Sulphur Ce 2 S 3 ; ThS 2 . 

1. Monoxides and monosulphides (selenides and tellurides 
have not been prepared). 

Sources. Carbon monoxide is produced by the decay of 
organic matter, and by the incomplete combustion of fuel. 

Preparation. By direct union. Carbon is said to combine 
with oxygen to form monoxide ; it appears more likely that the 
dioxide is first formed, and by its contact with red-hot carbon is 
converted into monoxide, thus COz + C = 200. 

2. By replacement. Steam, led over white-hot carbon, yields 
a mixture of hydrogen and carbon monoxide. This mixture is weK- 
adapted for heating purposes, and is commercially termed " water- 
gas." It is frequently employed in driving gas-engines. Carbon 

* As hydrate, Ti(OH) 2 . 


withdraws oxygen from sodium sulphate, NauS0 4 , forming mon- 
oxide and sodium sulphide. Carbon withdraws oxygen from 
many oxides, carbon monoxide being formed. 

3. By reduction. Zinc or copper withdraws oxygen from 
carbon dioxide, producing monoxide ; heating a mixture of mag- 
nesium carbonate and zinc dust is an available method of pre- 
paration. Carbon monosulphide is deposited from carbon di- 
sulphide, after long exposure to light ; and titanium monosulphide 
is produced by the action of hydrogen on the red-hot disulphide. 

4. By decomposition of a compound. The oxide C 2 O 3 
appears to be incapable of existence, but oxalic acid, C 2 O 4 H 2 , may 
be viewed a$i its compound with water. On depriving oxalic acid 
ot water by the action of concentrated sulphuric acid, a mixture of 
carbon monoxide and dioxide is evolved, thus 

C 2 O 4 H 2 + H 2 S0 4 = CO + CO, + H 2 S0 4 .H 2 0. 

Similarly, if the elements of water are withdrawn from formic acid, 
C0 2 H 2 , by strong sulphuric acid, carbon monoxide is produced. 
This is by far the most convenient, though not the cheapest, method 
of preparation, and yields perfectly pure monoxide. 

5. By double decomposition. Hydrocyanic acid, HCN (see 
p. 559), liberated in presence of fairly strong sulphuric acid, takes 
up water, forming carbon monoxide, and ammonia which combines 
with the sulphuric acid, thus 

HON" + H 2 S0 4 .H 3 = CO + (NH 4 )HS0 4 . 

The hydrocyanic acid is conveniently produced from potassium 

Properties. Carbon monoxide is a colourless gas at ordi- 
nary temperatures; it condenses to a liquid at 190, and the 
white solid produced by its evaporation melts at 199. Its critical 
temperature is about 139*5; and its critical pressure is 3|>'5 
atmospheres. It is soluble in alcohol ; 100 volumes dissolve about 
20 volumes ; but it is very sparingly soluble in water, 100 volumes 
dissolving only 3 volumes of the gas. It has a faint smell, but no 
taste. It is poisonous, forming a compound with the hemoglobin 
of the blood which gives a spectrum closely resembling that of 
oxyLtemoglobin ; but, while the latter is at once altered by ammo- 
qium sulphide, the spectrum due to carbon monoxide lasts after 
such treatment for several days. It is absorbed by potassium 
(see below), and by compounds of silver, and gold ; also by cuprous 
chloride. When left long in contact with potassium or sodium 
hydroxide it combines, forming formate of potassium or sodium. 


Carlson monosulphide is a red powder, sparingly soluble in 
carbon disulphide and in ether ; it dissolves in solution of potassium 
hydrate, and is reprecipitated by acids ; it decomposes at 200 into 
carbon and sulphur. It is probably a polymeride of CS. Tita- 
nium monosulphide is a black insoluble substance, decomposed 
only by fusion with alkalis. 

Compounds with water. It is sometimes stated that carbon 
monoxide is the anhydride of formic acid, CO 2 H 2 , and if their 
formulse alone be considered such might bo the case. But there 


can be no doubt that formic acid has the constitution H C OH, 

and that it is partly a carbide of hydrogen, and is derived from 
tetrad carbon. The true acid derived from carbon monoxide is 
unknown ; its formula should be HO C OH. Hence carbon 
monoxide reacts slowly with potassium hydroxide, a molecular re- 
arrangement being effected in order to produce potassium formate, 

II C OK. The explosive grey compound produced by direct 
combination of carbon monoxide with potassium, which has the 
formula K,,C M is probably also partly a carbide of potassium. 

Ti(OH) 2 is said to be produced by the action of sodium amalgam 
on the tetrachloride, TiOU, in presence of water. Titanium di- 
chloride, TiCl 2 , decomposes water, giving a mixture of trichloride 
and sesquioxide. "No compounds of these monoxides with oxides 
or sulphides are known. 

Compounds with chlorides. Carbon monoxide combines 
directly with platiiious chloride, to form the body PtCl 2 2CO, 
with platinic chloride to form PtCli.SCO, and with cuprous 
chloride to form Cu 2 Cl 2 .2CO. These are insoluble crystalline 
compounds. The last is formed when carbon monoxide is shaken 
with a solution of cuprous chloride in hydrochloric acid, and is 
used as a means of separating carbon monoxide from other gases 
with which it may be mixed. 

A compound of the formula TiO.TiCl 3 is also known ; it is 
produced by the action of oxygen on titanium tetrachloride, TiCl 4 , 
at a red heat. 2CeO.CeCl 2 is formed by the action of steam and 
nitrogen ou a mixture of cerium and sodium chlorides ; it forms 
silvery scales. - 

II. Sesquioxides and sesquisulphides. Carbon sesqui- 
oxide is unknown; its compound with water is oxalic acid. 
Carbon sesquisulphide is said to be produced by the action of 
sodium amalgam on the disulphide ; it is a red-brown powder. 


Titanium sesquioxide is formed when the dioxide is heated in 
hydrogen, or during the preparation of the trichloride (see page 145) 
due to the action of air. It forms copper-coloured crystals, and 
has the same crystalline form as specular iron, Pe 2 O 3 . Titanium 
sesquisulphide is produced by the action of a moist mixture of 
hydrogen sulphide and carbon disulphide on the dioxide, TiO 2 , at 
a bright red heat. It is a black powder. Cerium sesquioxide is 
produced by heating the oxalate in a current of hydrogen. It is a 
grey solid, reacting with acids forming salts. The sesqui- 
sulphide,* produced by the action of dry hydrogen sulphide on 
red hot cerium dioxide, or by passing that gas over a fused 
mixture of cerium trichloride and sodium chloride, is a crystalline 
vermilion or black compound, according to the temperature. It 
is slowly decomposed by warm water. Similar compounds of 
zirconium and thorium have not been prepared. 

Compounds with water. Oxalic acid may be regarded as 
the hydrate of the unknown carbon sesquioxide. It has the 
formula C 2 O4H 2 , and not CO 2 H, as can be shown by the following 
synthesis : Ethylene is known to possess the formula 2 JJ 4 from 
its vapour-density. On bringing ethylene and bromine together, 
direct addition takes place, and ethylene dibromide, C 2 H 4 Br 2 , is 
formed. This body, on treatment with silver hydroxide, exchanges 
bromine for hydroxyl, thus : C 3 H 4 Br 2 + 2AgOH = C 2 H 4 (OH) 3 
+ 2AgBr. Glycol, as the substance C 2 H 4 (OH) 2 is named, on 
oxidation yields oxalic acid, thus : C 2 H 4 (OH) 2 + 4O = C 2 O 2 (OH) 2 
+ 2H 2 O. It is therefore concluded that oxalic acid contains 
two atoms of carbon. Its constitutional formula is written 

| , and it would thus appear that the atom of carbon is 


here capable of combining with four monads, and is a tetrad. 
As carbon tetrachloride possesses the formula CC1 4 , and carbon 
hexachloride is C 2 C1 6 (see p. 155), it is seen that two atoms 
of carbon possess the property of combining with each other. 
Now, in contrasting this with the behaviour of members of 
the previous group, such as iron, it must be remembered that 
ferric chloride possesses the formula FeCl 3 , as shown by its 
vapouv- density at high temperatures. At low temperatures, its 
formula is Fe 2 Cl 6 , and it has been supposed that iron at low tem- 
peratures, like carbon under almost all circumstances, is a tetrad. 
The hydroxide, Fe 2 O 3 .H 2 O, has probably a high molecular weight, 
for the sesquioxide, Fe 2 O 3 , has the power of combining with a 

* Comptes rend., 100, 1461. 


large number of molecules of other oxides, and presumably com- 
bines with itself to form considerable molecular aggregates. But 
ignoring this, the formula of this hydroxide may be O=Fe OH, 
or it may have a constitution analogous to that of oxalic acid, viz., 

case * ne a t ms f i ron would 

be tetrad. At present there is no means of deciding the point, 
though the opinion of chemists favours the triad nature of the 
atom of iron. 

The study of oxalic acid and its compounds belongs to the 
domain of Organic Chemistry; and these are so numerous that 
their formation and relationship would occupy too large a space in 
such a book as this. 

Titanium tetrachloride on treatment with metallic copper or 
silver in a state of fine division yields the hexachloride, Ti 2 Cl 6 ; 
and on addition of an alkali, to its solution in water, a brown pre- 
cipitate of the hydroxide, Ti 2 O 3 .3H 2 O, is produced. It is soluble in 
acids giving violet salts. 

Hydrated cerium sesquioxide is formed by addition of an 
alkali to a solution of the trichloride. It rapidly oxidises on 
exposure to air. 

Compounds with halides. On treating trichloride of titanium 
with a little water, the body Ti 2 2 Cl 2 is produced. Supposing 
it to be constituted like oxalic acid, its formula would be 

Qj^Ti Ti<^,. It is also formed by the action of a mixture of 

hydrogen and titanium tetrachloride on the red hot dioxide, thus : 
TiClt + H, + TiO 2 = 2HC1 + Ti 2 O 2 Cl 2 . It forms reddish- 
brown laminee. 

A compound of the formula Ce 2 O 3 .Ce 4 Cl(5 is produced by the 
action of sodium hydroxide, and subsequently of water on the tri- 
chloride, or of a mixture of steam and nitrogen on the trichloride. 
It is an insoluble dark purple powder. 

HI. Dioxides. Sources. All of these dioxides, that of cerium 
excepted, are found native. Carbon dioxide occurs in air. 
Ordinary country air contains somewhat under 4 volumes per 
10,000 of air ; in cities, owing to its evolution from chimneyn and 
from respiration, it is present in somewhat higher amount, and in fogs 
may amount to 6 volumes. It issues from the ground in volcanic 
districts. The " Grotto del Cane," near Naples, is well known in 
this respect ; the gas in the depression in the ground contains from 
60 to 70 per cent, of carbon dioxide. It is a frequent constituent of 


mineral waters, and is present in small quantity in all natural 
water, including sea- water. It is the source from which plants de- 
rive their carbon, and is produced by the decay of all organic matter. 
Some specimens of quartz contain cavities filled with liquid carbon 
dioxide. In combination with other oxides, especially with lime, 
as carbonate, it forms a great portion of the earth's crust. 
Titanium dioxide seems native in dimetric prisms, as rutile, in 
granite, gneiss, or mica slate ; also as anatase, in acute rhombo- 
hedra ; and as brookite in trimetric crystals. Zirconium dioxide 
occurs in combination with silica as zircon, or hyacinth, ZrO 2 .SiO 2 , 
and as malacone, in some granites. Thorium dioxide occurs as 
thorite, 3(ThO 2 .SiO 2 ).4H 2 O, and is also combined with niobic and 
tantalic pentoxides in euxenite. 

Preparation. In considering the methods of preparation of 
these compounds it must be remembered that carbon dioxide is a gas, 
while the dioxides of the other elements are non- volatile solids. 

1. By direct union. The elements all burn in oxygen, forming 
dioxides > with exception of cerium. In presence of excess of the 
element, carbon forms monoxide, and titanium forms sesquioxide. 
Cerium yields, not dioxide, but sesquioxide. Compounds of carbon 
also burn, giving carbon dioxide. Carbon unites with sulphur at a 
red heat, forming disulphide ; but it does not combine directly with 
selenium or tellurium ; and zirconium and thorium also form 
disulphides when heated in sulphur gas. The selenides and tellurides 
of the other elements have not been prepared. 

The combustion of carbon in oxygen may be shown by heating a piece of 
charcoal to redness in a Bunsen's flame, and plunging it into oxygen gas, as shown 
in fig. 33. The charcoal continues to burn brightly, and the product is carbon 

Fio. 33. 

T 2 


dioxide. The combustion of a diamond may also be shown, as in fig. 34, by 
wrapping up a fragment of diamond in a small spiral of thin platinum wire 
connected with two stout copper wires which pass through an mdiarubber cork 
closing the end of a wide test-tube. The test-tube is filled with oxygen, and 
by means of an electric current from four Bunsen's cells, the thin platinum wire 

FlGK 34. 

ia heated to whiteness. The diamond is thus raised to its point of ignition, 
and on discontinuing the current it continues to glow until it is finally totally 
consumed. That carbon dioxide is the product of combustion may be shown by 
shaking the contents of the tube with a little baryta-water (Ba(OH) 2 .Aq), when 
a white precipitate of barium carbonate, BaCO 3 , is formed, Another instructive 

FIG. 35. 


experiment is devised to show that the volume of carbon dioxide produced by the 
union with carbon of a known volume of oxygen is equal to that of the oxygen. 
The oxygen is contained in the bulb, and confined over mercury. The carbon 
is wrapped in a piece of platinum wire, and, as in the case of the diamond, 
heated to its point of ignition. The gas expands at first, of course, owing to 
its temperature being raised, but on cooling, the mercury in the two limbs of 
the (J-tube returns to its original level, showing that the volume of gas is the 
same after it has been converted into carbon dioxide. (See fig. 35.) 

Carbon also withdraws oxygen from its compounds with other 
elements, combining with it, a mixture of monoxide and dioxide 
being usually formed. Carbon heated to bright redness in steam 
gives a mixture of monoxide, dioxide, and hydrogen (" water-gas ") 
There is littfe doubt that the oxides of the other elements of this 
group could be similarly formed. 

Compounds of carbon with hydrogen and oxygen also burn in 
oxygen forming dioxide. Thus when a candle, consisting chiefly of 
carbon and hydrogen, burns, both its carbon and hydrogen unite 
with oxygen. The union takes place more rapidly in oxygen gas 
than in air, but the total amount of heat evolved is the same which- 
ever be employed. But owing to the greater rapidity of combina- 
tion, the temperature ib higher during combustion in oxygen than 
in air. The oxidation of the blood of animals is also a slow com- 
bustion, taking place in the capillary bloodvessels, the oxygen being 
derived from the inspired air. 

2. By union of a lower oxide or sulphide with oxygen 
or sulphur. Carbon monoxide burns in air or oxygen to form 
dioxide. A mixture of the two gases explodes on passing a spark, 
provided they are moist. No explosion takes places when they are 
dry, although combination occurs in the space through which the 
spark passes. ' Carbon monoxide also withdraws oxygen from 
oxides of many other elements, such as those of iron, copper, &c., 
to form dioxide. When heated to whiteness with steam, a portion 
is converted into dioxide. Titanium sesquioxide and sesquisulphide 
readily unite with oxygen or sulphur, forming dioxide or disul- 

3. By the action of heat on a compound. All carbonates, 
those of lithium, sodium, potassium, rubidium, and coesium ex- 
cepted, lose carbon dioxide when heated. Barium carbonate re- 
quires a white heat ; strontium carbonate a bright red heat, and 
calcium carbonate a red heat. These carbonates decompose more 
readily if heated in a current of some indifferent gas, such as air or 
steam. Compounds of the other dioxides have not been thus 
decomposed, owing to the non-volatility of the dioxides. But 
the sulphocarbonates, like the carbonates, are decomposed by 


heat into sulphides and carbon disulphide. Calcium compounds, 
for example, decompose thus 

CaCO 3 = CaO + 00 2 ; CaCS 3 = CaS + CS 2 . 

The dioxides of titaninm, zirconium, cerium, and thorium are 
produced by heating their hydrates or sulphates, and that of thorium 
by heating its oxalate. 

4. By displacement. This method, as a rule, yields the 
hydrates ; but as carbonic acid (the hydrate of carbon dioxide) is 
very unstable, it is produced thus : for example, a carbonate, 
treated with sulphuric acid, yields a sulphate, carbon dioxide, and 
water : Na 2 C0 3 .Aq + H 2 S0 4 .Aq = Na 2 S0 4 .Aq + V0 2 + H 2 O ; 
or the reaction may be thus written : CC^.Na^O + S0 3 .H 2 O = 
S0 3 .Na 2 O + C0 2 + H 2 0. There is no tendency to form a compound 
between carbon dioxide and sulphur trioxide. 

In actual practice, carbon dioxide is prepared on a large scale 
by burning carbon in air, or by treating calcium carbonate with 
sulphuric or hydrochloric acid. When the last acid is used, some 
spray of hydrogen chloride is apt to be carried over with the carbon 
dioxide, hence it is advisable to wash it by leading it through a 
solution of hydrogen sodium carbonate. If sulphuric acid is employed, 
the calcium carbonate must be in the state of fine powder, else it 
becomes coated with an insoluble layer of sulphate which hinders 
further action. It is by this method that carbon dioxide is usually 
made in the manufacture of " aerated water." 

Cerium tetrafluoride, when heated in air, loses fluorine, and 
yields the dioxide. This is probably due to the moisture in the 
air, forming hydrogen fluoride, and would come under the next 

5. By double decomposition. Carbon disulphide has been 
produced by heating carbon tetrachloride to 200 with phosphorus 
pentasulphide ; substituting selenium for sulphur, a liquid was 
produced containing about 2 per cent, of diselenide. 

Special method. Carbon dioxide is produced by the de- 
composition of grape-sugar, C 6 Hi 2 O 6 , under the action of the yeast 
ferment (Saccharomyces cerevisice), when ethyl alcohol, C 2 H 5 .OH, 
and carbon dioxide are the chief products. 

The starch contained in grain is converted during the process of '* malting" 
or incipient germination, during which the grain is kept warm and moist on 
the "mal ting-floors," into grape-sugar, by aid of the ferment diastase, con- 
tained in the grain. The growth is then stopped by heating the malt ; it is 
crushed, and is known as " grist j" it is transferred to the "mash -tun," a 
large cask or vat, where it is treated with warm water. The solution of grape- 


sugar thus obtained is called the "wort;" it is mixed with yeast, and left to 
ferment, when the change already mentioned takes place. The carbon dioxide 
fills the vat and escapes into the air. The equation is C 6 II 12 6 = 2C 2 H 6 .OH 
* 2C0 2 . 

Properties. At the ordinary temperature carbon dioxide is a 
gas. Its boiling point under normal pressure is about 79. It 
melting point is nearly the same as its boiling point ; it is given as 
78*5, hence the liquid easily freezes by its own evaporation. It 
may be condensed to a liquid at a pressure of about 36 atmospheres 
at 10. The gas is colourless, has a faint sweetish smell and taste, 
and is much^ heavier than air, hence it is best collected by down- 
ward displacement. Its great density (22, compared to air=14'47) 
permits of its being poured from one vessel to another without 
much loss. Its density is easily shown by pouring it into a light 
beaker, suspended from the beam of a balance, and counterpoised. 

FIG. 36. 

Carbon dioxide supports the combustion of the elements potas- 
sium? sodium, and magnesium. They deprive it of a portion of its 
oxygen, forming oxides and carbon monoxide, as well as some free 
carbon ; the oxide then unites with excess of carbon dioxide, forming 
a carbonate. Carbon may also be said to burn in carbon dioxide, inas- 
much as when the dioxide is led over red hot carbon, the monoxide 
is formed : but because the heat evolved by this reaction is com- 


paratively small, the carbon is not thereby kept at its temperature 
of incandescence, and action ceases, unless a supply of heat be 
added from without. When carbon burns in oxygen, therefore, the 
whole of the oxygen is not converted into carbon dioxide ; the action 
ceases when the dioxide formed bears a certain proportion to the 
total gas present ; the reverse action then tends to begin. Hence 
a candle, burning in air, goes out when the carbon dioxide formed 
reaches a certain proportion of the total gas ; and for the same 
reason, an animal dies when breathing a confined atmosphere, loog 
before it has completely deprived it of oxygen. A man can breathe, 
however, for some time in an atmosphere in which a candle refuses 
to burn, as was shown by the late Dr. Angus Smith.' Carbon di- 
oxide is decomposed by the green colouring matter of plants in 
sunshine ; the exact nature of this decomposition is not known ; 
there are grounds for supposing that it consists in a reaction 
occurring between carbon dioxide and water, as follows : 

C0 2 4- H 2 = H 2 CO + O 2 . 

The substance H 2 CO is named formic aldehyde, and it has been 
recently shown to be easily transformable into a kind of sugar, 
CeHizOe, named formose. There may be some connection between 
this transformation and the formation of sugar in plants. The 
carbon dioxide is absorbed by the stomata or " small mouths " in 
the under surface of the leaves of plants, and oxygen is evolved. 
This may be experimentally shown by placing some blades of grass 
in a jar of water inverted over a trough. The oxygen gas collects 
in the upper portion of the jar during several days' exposure to 
sunlight, and may be recognised by the usual tests. 

Liquid carbon dioxide is heavier than water, and does not mix 
with it. It is a non-conductor of electricity. Above the temperature 
30*9, the critical point of carbon dioxide, the gas cannot be made to 
assume the liquid state by compression. The solid dioxide is a 
loose white powder, like snow, produced by allowing the liquid to 
escape into a thin flannel bag; the liquid absorbs heat during its 
conversion into gas, and a portion solidifies, owing to its being 
thus cooled. A mixture of solid carbon dioxide with ether gives 
a temperature of 100. 

It has been recently shown that carbon and carbonic oxide do 
not unite with perfectly dry oxygen, unless they be kept exposed 
to a very high temperature. The presence of water or some 
other compound containing hydrogen being necessary, it is sup- 
posed that the carbon or carbonic oxide reacts with the water 
liberating hydrogen, thus CO + H Z = CO Z + 2flT, or C + H 2 



= CO -f 2J5T, and that the hydrogen then unites with the oxygen, 
to form water, which is again acted on.* 

The test for carbon dioxide is its combination with calcium or 
barium oxide, when shaken with a solution of the respective hydr- 
oxide, to form carbonate, in either case a white powder, which 
effervesces with acids. 

The presence of carbon dioxide in expired air may be demonstrated by the 
arrangement shown in the figure : 

The air entering the lungs passes through lime-water in the bottle on the 
right hand side ; as ordinary air contains only 4 volumes of carbon dioxide 
in 10,000, a turbidity is not seen for some time. The exhaled air passes 
through the lime-water in the left hand bottle and soon turns it turbid. 

The amount of carbon dioxide in atmospheric air may be esti- 
mated comparatively by measuring the amount required to produce 
incipient turbidity in baryta water. 

The little apparatus is shown in fig. 38. The indiarubber ball is squeezed, 
the air escaping through the opening. The opening is then closed with tlie 
finger, and, on allowing the ball to expand, air is drawn through the baryta 

FIG. 38. 

* Dixon, Chem. Soc., 49, 94. 


water. On removing the finger the ball is again squeezed empty, and air is again 
drawn through the baryta water. Having found the number of charges 
of the ball which must pass through the baryta water to produce a turbidity 
with ordinary air, it may be assumed with fair correctness that the normal 
amount is present, viz., 4 volumes in 10,000. On applying the same test to 
vitiated air, fewer charges are required, and the amount of carbon dioxide may 
be calculated by simple proportion. 

Carbon dioxide rapidly combines with the hydroxides of sodium 
and potassium, as well as with those of calcium and barium. The 
method of absorbing it from gaseous mixtures is to shake them 
with a strong solution of potassium hydroxide. It may also easily 
be absorbed by passing it through a solid mixture of Hydroxides of 
calcium and sodium, commonly termed " soda-lime." 

Carbon disulphide is a limpid colourless liquid, heavier than 
water and not mixible with it, melting at 110 and boiling at 
46*04. In the crude state it contains hydrogen sulphide and dis- 
agreeably-smelling sulphur compounds. It may be purified from 
hydrogen sulphide by shaking it with a solution of potassium 
permanganate, which oxidises that impurity, and from sulphur- 
compounds and sulphur by shaking it with mercuric chloride and 
mercury and distilling it. When pure it has a not unpleasant 
ethereal odour. Its vapour is very poisonous when breathed. Its 
vapour ignites very easily when mixed with air (at 149), hence it 
must be kept away from a flame and distilled by aid of a water- 
bath. It is decomposed by light, acquiring thereby a disagreeable 
smell. It is slightly soluble in water. Its vapour explodes when 
exposed to the shock of decomposing fulminate of mercury, being 
resolved into the elements carbon and sulphur. It is formed with 
absorption of heat, hence its instability ; heat is evolved when it is 
exploded. It mixes easily with alcohol, ether, and oils, and is used 
for extracting oils and fats from acids, animal refuse, wool, &c., 
and as a solvent for sulphur chloride in vulcanising caoutchouc. 

It unites with sulphides, giving sulphocarbonates (see below), 
and when passed through a hot tube with chlorine it yields 
sulphur chloride (S 2 C1 2 ) and carbon tetrachloride (see p. 145). 

In preparing the pure dioxides of titanium, zirconium, cerium, and 
thorium, the chief difficulty is the separation from the oxides of other elecnents, 
especially from silica. The process is, fusion with a mixture of potassium and 
sodium carbonates (fusion-mixture), which yields in each case silicate, titanate, 
zirconate, or thorate of the alkaline metals, and the oxides of the other metals, 
if these are present. In the case of titanium, hydrogen fluoride is added to the 
solution of the fused mass in water, and the titanium thrown down as double 
fluoride of titanium and potassium, TiF 4 .2KF. These crystals are afterwards 
dissolved in water, and on addition of ammonia the titanium is thrown down as 


hydrate. With zirconium, the fused mass, consisting of silicate and zirconate 
of sodium, and potassium, is mixed with excess of hydrochloric acid, and evapo- 
rated to dryness. This gives a mixture of silica and oxychlorides of zirconium. 
On treatment with hydrochloric acid, the silica, not being thus converted 
into chloride, does not dissolve, hut the zirconium dissolves as chloride, 
along with iron, &c. The solution is boiled with thiosulphate of sodium, which 
precipitates the zirconium, leaving the iron in solution. On ignition ot the 
thiosulphate of zirconium, pure zirconia, ZrO 2 , is left. 

Cerium is similarly separated,* but it is precipitated as oxalate, and on 
ignition the oxide Ce 2 O 3 is left. Thorium is precipitated as oxalate, from its 
solution in hydrochloric acid, after separation of silica ; and from a solution of 
the oxalate in hydrochloric acid by a strong solution of potassium sulphate, 
with which it combines, forming a double sulphate of thorium and potassium 
see p. 428) .f It also yields an insoluble thiosulphate. 

Titanium dioxide, native as rutile, forms reddish -brown 
crystals ; artificially prepared it is a reddish-brown powder. It is 
insoluble in water and does not react with acids, except with 
strong sulphuric acid or fused bisulphates. 

It melts in the oxyhydrogen flame. It has been artificially 
crystallised by passing vapours of titanium tetrachloride and 
steam through a red-hot tube. 

Zirconia, or zirconium dioxide, is a white powder ; it is 
obtained in small quadratic prisms by crystallisation from fused 

Cerium dioxide is a pale-yellow insoluble substance, which 
also crystallises from fused borax in tesseral crystals. On boiling 
with hydrochloric acid, chlorine is evolved, and the trichloride is 
produced, CeCl 3 . With sulphuric acid it also dissolves, the sul- 
phate Ce 2 (S04) 3 .Ce( SO^ being formediWitb evolution of oxygen. 
It is soluble in nitric acid. 

Thorium dioxide, or thoria, is a white powder, separating 
from its solution in borax in transparent quadratic crystals. 

Compounds with water and hydrogen sulphide. Carbon 
dioxide as gas dissolves to some extent in water; 100 volumes 
of water at 20 dissolve about 90 volumes, and at 15 about 
100 volumes. The solution has a pleasant sharp taste, and is 
usually called " soda-water." The carbon dioxide is, however, 
forced in under a pressure of several atmospheres. The gas 
escajlks quickly if the pressure is decreased immediately; but 
after some days or weeks it appears to have entered into combina- 
tion to some extent with the water, and does not then escape so 

* For details regarding cerium compounds see Brauner, Chem. Soc., 47, 
879 ; references to other papers are given, 
f Cleve, Bull. Soc. CMm. (2), 21, 115. 


readily. The solution turns litmus solution claret coloured. It acts 
on zinc, iron, and magnesium, forming carbonates and liberating 
hydrogen. Itfprobably consists of a weak solution of carbonic acid, 
H 2 C0 3 , with carbon dioxide uncombined but mixed with the water. 

Carbon disulphide does not unite directly with hydrogen sul- 
phide, but sulphocarbonic acid, as the compound is named, 
H 2 CS 3 , is produced on addition of weak hydrochloric acid to a 
solution of a sulphocarbonate, e.g., Na 2 CS 3 (see below). It is a 
dark-yellow oil, with a pungent odour, and on rise of temperature 
it rapidly decomposes into carbon disulphide, CS 2 , and hydrogen 
sulphide, H Z 8. 

Many hydrates of titanium dioxide have beSn described, 
but the data regarding them are as a rule contradictory. On 
heating titanic hydrate thrown down from its chloride by an alkali 
it loses water gradually, with rise of temperature, and shows no 
sign of any definite hydrates. It is probable that there are many, 
and that no one is stable over any large range of temperature. 

The hydrates of zirconium dioxide appear also to be numer- 
ous. The only sudden break in drying the hydrate precipitated 
from the chloride by ammonia is at 400. On reaching this tem- 
perature the body suddenly turns incandescent, and all water is 
expelled. It has then become difficult to dissolve in acids, and it 
is believed that sudden polymerisation has occurred, many mole- 
cules of Zr0 2 having united to form one complex molecule. 

Cerium hydrate at 600 has the formula CeO 2 .2H 2 O. At 
lower temperatures it is brownish-yellow, but at that temperature 
and above it is bright yellow ; as it dries further, its colour changes 
to a salmon-pink. It is produced by the action of sodium hypo- 
chlorite on Ce 2 O 3 . 

Thorium hydrate is a gelatinous mass ; it probably resembles 
titanium hydrate. 

Compounds with Oxides and Sulphides Carbon- 
ates, Titanates, Zirconates, Thorates Carbon 
Oxysulphide, Oxysulphocarbonates, and Sul- 
phocarbonates. % 

These compounds maybe divided into two classes : (1) 
compounds, those in which the ratio of the number of oxygen 
atoms in the dioxide to that of the oxide of the metal is as 
2:1; and (2) basic compounds, those in which the ratio is 
less than 2:1; no acid compounds, those in which the ratio is 


greater than 2:1, are known. The normal compounds are most 

But before considering these bodies it is advisable to describe 
carbon oxysulphide, of which the formula is COS* as shown by 
its gaseous density. This body, therefore, cannot be regarded as a 
compound of carbon dioxide and carbon disulphide, CO 2 .CS 2 , but 
as carbon dioxide, of which one atom of oxygen is replaced by sul- 
phur. It may be produced by leading a mixture of carbon dioxide 
and carbon disulphide gases through a tube filled with platinum 
black, i.e., finely-divided platinum, or by the union of carbon 
monoxide with sulphur. But it is most easily produced by the 
reaction between sulphocyamde of hydrogen and water. The 
compound KCNS, on treatment with sulphuric acid, yields the 
acid HCNS. If the sulphuric acid be moderately strong and 
warm, it combines with the ammonia produced by the decom- 
position of the acid, thus : HCNS -f H 2 = NH 3 + COS. Car- 
bon oxysulphide is a not infrequent constituent of mineral springs, 
but, as a rule, it has for the most part reacted with water to form. 
carbon dioxide and hydrogen sulphide, thus : COS + H 2 O = 
COz H- #2$. It is a colourless gas, without odour or taste when 
pure, somewhat soluble in water, and combustible to dioxides of 
carbon and sulphur. It is hardly affected by aqueous potash, but 
is easily absorbed by an alcoholic solution. Its physiological 
effects resemble those of nitrous oxide.f 

There are thus three bodies, all of which form compounds with 
oxides and sulphides, viz. : 00 2 , carbon dioxide ; COS, carbon 
oxysulphide ; and CS 2 , carbon disulphide. 

Compounds of Carbon Dioxide with. Oxides. 

1. Normal carbonates. Ratio of oxygen in carbon dioxide 
to oxygen in combined oxide, 2:1. 

The following is a list of the known compounds : 

Simple carbonates: Li 2 CO 3 ; Na^COs with 15, 10, 8, 7, 6, 5, 2, and 
1H 2 ; K 2 C0 3 with 2H 2 O and H 2 O ; Rb 2 CO 3 .H 2 O ; Cs 2 CO 3 ; 
(NH 4 ) 2 C0 3 .H 2 0. 

Complex carbonates : HNaCO ;i ; H 2 Na 4 (CO 3 ) 3 .3H 2 O ; HKCO 3 .HjO ; 

X HRbC0 3 ; HCsCOg ; HNH 4 CO 3 ; H 2 (NH 4 )4(CO 3 ) 3 .2H 2 O. 

' These carbonates are all made by the action of carbon dioxide on 
a solution of hydroxide of the metal, thus : 2NaHO.Aq -f CO.. = 
q + H 2 0. 

* Than, Annalen, Suppl. 1, 236. 
t J. prakt. Chem. (2), 36, 64. 


Of these, lithium carbonate, Li 2 CO 3 , occurs in mineral waters ; 
it is sparingly soluble in water (about I '4 grams in 100 of 
water at 20), and may be produced by addition of a concen- 
trated solution of sodium carbonate to a soluble salt of lithium. 
For the preparation of sodium and potassium carbonates, see 
p. 671. 

Sodium carbonate is a constituent of certain " soda-lakes "in 
Egypt and Hungary ; it also occurs in volcanic springs. 

The ordinary name for the carbonate Na 2 CO 3 is soda-ash; for 
the crystalline salt, Na^COa.lOI^O, soda crystals or " washing- 
soda;" and for hydrogen sodium carbonate HNaCO 3 , bicarbonate, 
or " baking-soda." The latter is produced by treating the normal 
carbonate (crystals) with carbon dioxide, thus : Na 2 CO 3 -f CO Z 4- 
H 2 O = 2NaHCO 3 . Hydrogen sodium carbonate is less soluble 
than sodium carbonate. 

Carbonate of sodium melts at about 818, and of potassium 
at about 830. On heating hydrogen sodium carbonate it loses 
water and carbon dioxide, and yields sodium carbonate, thus: 
2HNaCO 3 = H Z + C0 2 + NaCO B . The simple carbon- 
ates, except those of ammonium (see p. 533), volatilise un- 
changed at a bright red or white heat, and are not decomposed 
into carbon dioxide and metallic oxide. It is probable that a car- 
bonate of sodium and potassium also exists, of the formula 
NaKCO 3 ; a mixture of the two is named "fusion mixture," and 
is used in the decomposition of silicates, &c. It has a much lower 
melting point than either of the pure salts. The compound 
H 2 Na4(CO3)3.3H 2 O occurs native, and is known as trona or urao ; 
in old times it used to be an important source of soda. These bodies 
have all an alkaline, cooling taste ; the ammonium compound 
smells of ammonia, owing to its decomposing, on exposure, into 
ammonia, carbon dioxide, and water. Hydrogen ammonium car- 
bonate is found in guano deposits. 

Simple carbonates : BeCO 3 ,4H 2 O ; CaCO 3 , also 5H 2 O ; SrCO 3 ; 

BaCO 3 ; M?CO 3 ; also 3H 2 O and 5H 2 O ; ZnCO 3 ; CdCO 3 . 
Complex carbonates : H 2 Ca(CO 3 ) 2 .Aq (?) j Na 2 Oa(CO 3 ) 2 .5H 2 O j 

H 2 Mg(C0 3 ) 2 .Aqj Na 2 Mfir(C0 3 ) 2 ; HKMff(CO 3 ) 2 .4H 2 O j 

(NH 4 ) 2 Mgr(C0 3 ) 2 .4H 2 0; H 2 K 8 Zn 6 (CO 3 ) u .7H 2 O : Na 6 Zn 8 (OO 3 ) n .8H 2 O. 
Na 6 Zn 8 (CO 3 ) 11 .8H 2 O. % 

The&e carbonates are all white solids. They are decom* 
posed by heat (see the respective oxides), barium carbonate 
requiring the highest temperature. The following are found 
native : Calcium carbonate, CaC0 3 , as calcspar or Iceland-spar, 
in hexagonal rhombohedra ; as arrayonite in trimetric rhombic 


prisms ; as marble, limestone, chalk ; a constituent of shells, 
bones, &c. It may be produced in the form of calcspar by 
crystallisation from a mixture of fused sodium and potassium 
chlorides ; by precipitation from solution below 30 ; and as 
arragonite by precipitation above 90.* Between these tempera- 
tares mixtures of microscopic crystals of the two are precipitated 
by addition of sodium or ammonium carbonate to a solution of a 
soluble salt of calcium. When heated to redness in a closed iron 
tube, calcium carbonate fuses, and then yields a crystalline mass 
resembling marble. The carbonates of calcium, strontium, and 
barium are formed by direct union of oxide with carbon dioxide ; 
the union is^attended with great evolution of heat, causing the 
oxide to become incandescent ; the product with lime has the 
formula CO 2 .2CaO. The compound Na 2 Ca(CO 3 ) 2 .5H 2 O is named 
gaylussite ; strontium carbonate, SrCO 3 , is found native as stron- 
tianite ; and barium carbonate, BaCO 3 , as witherite. MgCO 3 is* 
magnesite, and a double carbonate of calcium and magnesium, in 
which indefinite amounts of both metals are present, is dolomite ; 
it forms large mountain ranges, named " The Dolomites," in 
northern Italy. Zinc carbonate, ZnCO 3 , occurs native as calamine, 
and is accompanied by cadmium carbonate, CdCO 3 . 

The so-called acid carbonates, e.g., H 2 Ca(COj) 2 , H 2 Mg(C0 3 ) 2 , 
and similar compounds of barium, strontium, &c., have not been 
isolated. Their existence is assumed because the normal car- 
bonates dissolve freely in a solution of carbonic acid. On warm- 
ing the solution, they are decomposed with evolution of carbon 
dioxide and precipitation of the simple carbonates. 

Carbonates of boron and scandium are unknown ; of this group 
only Y 2 (CO 3 ) 3 .12H 2 O, and 2H 2 O ; and La a (CO 3 ) 2 , found native 
as lanthanite, are known. The existence of a carbonate of alu- 
minium is doubtful ; carbonate of gallium is unknown. Ijl;(CO 2 ) 3 , 
however, has been prepared. These are insoluble white bodies, 
which lose carbon dioxide when heated, leaving the oxides. 

Thallium forms no thallic carbonate, but thallous carbonate, 
T1 2 CO 3 , is produced by precipitation. There is some evidence of 
a hydrogen thallium carbonate, HT1(CO 3 ). 

Chromic, ferric, and manganic carbonates are unknown. On 
addition of a soluble carbonate to their soluble salts, e.g., chlorides, 
the hydrates are precipitated, and carbon dioxide escapes, 
thus : 

2CrCl 3 .Aq + SNa^COa.Aq = Cr 2 3 .Aq + GNaCl.Aq + 3C0 3 . 
* Comptes rend., 92, 189. 


The carbonates derived from the monoxides of these metals are 
as follows : 

Simple carbonates : CrOO 3 ; FeCO 3 ; MnGO 3 ; CoOO 3 ; NiCO 3 . 
Complex carbonates : HKCo(CO 3 ) 2 .2H 2 O ; H 2 NaCo(GO 3 ) 3 .4H 2 O ; 
Na 2 Co(00 3 )2.10H 2 0; HKNi(CO 3 ) 2 .4H 2 O ; K<>Ni(CO 3 ) 2 .4H 2 O ; 

Na2NiCo(CO 3 ) 3 .10H 2 O. 

Of these, chromous carbonate is produced by mixing a solu- 
tion of chromous chloride with sodium carbonate ; FeCO 3 is found 
native, and named spathic iron ore or siderite ; in an impure state, 
mixed with clay or shale, it is termed clay-land or black-band, and 
forms one of the most important ores of iron. WhejL pure it is a 
whitish crystalline rock. It is soluble in water containing carbon 
dioxide ; such a solution may contain hydrogen ferrous carbonate, 
H 2 Fe(C0 3 ) 2 , which, however, has not been isolated. It is in this 
form a constituent of iron springs, and, on exposure to air, it loses 
carbon dioxide, and the iron oxidises to ferric hydrate, and deposits 
on the bed of the stream. A hydrated carbonate, FeCO 3 .H 2 O, 
also occurs native. Manganese carbonate, MnCOj, occurs native 
as manganese spar. 

No carbonates of titanium or zirconium are known. 

Cerium hydrate, however, on exposure to air, absorbs carbon 
dioxide, yielding Ce 2 (CO 3 ) 3 .9H 2 O. Silicon and germanium do 
not yield carbonates, but tin forms a basic carbonate (see below). 
Lead carbonate, PbCO 3 , occurs native, and is known as cerussite. 
Lead oxide sometimes replaces calcium oxide in native calcium car- 
bonate, to the extent of 3 or 4 per cent. ; the compound is called 
plumbocalcite. Plumbo-arragonite has also been found native at 
the lead hills in Lanarkshire. A chlorocarbonate, of the formula 
PbCO 3 .PbCl 2 , may be produced by boiling lead carbonate and 
chloride together in water. It is an insoluble white substance. 
It also occurs native as corneous lead. When heated, it loses 
carbon dioxide, leaving PbO.PbCl 2 . 

Carbonates of nitrogen, vanadium, niobium, and tantalum are 
unknown, and also carbonates of phosphorus, arsenic, and anti- 
mony; a basic carbonate of bismuth has been prepared (see 

Carbonates of molybdenum and tungsten do not appear to Vxist, 
but several double carbonates of uranyl (U0 2 ) (see p. 40?) 
have been prepared. These are Na 4 (UO 2 )(CO) 3 , K 4 (UO 2 )(CO 3 ) 3 , 
and (NH 4 ) 4 (UO 2 XCO 3 )3. A calcium compound occurs native ; its 
formula is Co^UOy^CO^.lOH^O. It is seen that the group U0 2 , 
or nranyl, acts like a dyad metal. 


Normal carbonate of copper is unknown. The only known 
normal compound has the formula Na2Cu(CO 3 ) 2 .6H,O. Silver 
carbonate, Ag 2 CO 3 , is a yellowish- white powder, produced by 
precipitation; it loses carbon dioxide at 200 ; KAgCO 3 is formed 
if HKCO 3 be used ; it is white. Mercurous carbonate, Hg 2 CO 3 , 
is a very unstable brown precipitate. Carbonates of gold and of 
the metals of the palladium and platinum groups are too unstable 
to exist. 

Considering *these carbonates as a whole, it may be noticed 
(i) that with exception of those of the sodium group of metals all 
are decomposed by heat into oxide and carbon dioxide ; (2) those 
of the sodium, calcium, smd magnesium groups, and thallous and 
cerium carbonates are formed by direct union of the hydroxides 
and carbon dioxide ; (3) that the oxides, except those of the 
Hodium group, do not combine directly with carbon dioxide ; cal- 
cium oxide, however, begins to combine at 415 ; and (4) that the 
carbonates of calcium, strontium, barium, and silver, are the only 
normal ones produced by precipitation by addition of a soluble 
carbonate to a soluble salt of the metals. In all other cases, basic 
carbonates are precipitated. These will now be considered. 

2. Basic carbonates.^These bodies contain a greater proportion of the oxide 
of the metal than is represented by the ratio given before. The oxygen of the 
metallic oxide bears a larger ratio to that of the carbon dioxide than 1 : 2. Their 
formulae are most conveniently stated as addition-formulae ; the relations then 
appear most clearly. They are unknown in the sodium group of elements. 

CO 2 .5BeO.5H. 2 O ; (jCOs.SE^O^BeO ; CO 2 .2CaO.H 2 O ; CO 2 .2CaO (produced 
by heating CaC0 3 j by heating CaO in CO 2 , the mass turning incandescent 
during union ; or by exposure of Ca(OH) 2 to air) ; 3CO 2 .4CaO, also produced 
by direct union; CO 2 .2SrO ; CO 2 .2BaO ; 4CO 2 .5MgO (precipitated hot); 
3GO 2 .4MgrO (native ; Tiydromagnesite) CO 2 .2ZnO.H 2 O ; OO 2 .3ZnO.3H 2 O 
(native; zinc bloom) ; CO 2 .5ZnO.6H 2 O ; 20O 2 .3ZnO.H 2 O ; 2C0 2 .5ZnO.5H 2 O ; 
4CO 2 .5ZnO.H 2 O; 4OO 2 .9ZnO.6H 2 O ; CO 2 .3CdO. These substances (the 
cadmium, strontium, and barium compounds excepted) are produced by pre- 
cipitation under various conditions of temperature and dilution. 

2CO 2 .4ThOt>.8H 2 O appears to exist, but there are no corresponding com- 
pounds of tin or lead. CO 2 .2SnO, however, is thrown down on addition of 
sodium carbonate to stannous chloride, SnCl 2 . as a white precipitate. 

The substance known as white-lead is probably a mixture of basic carbonates 
of lead* Seen under the microscope, it consists of small spherical masses, each 
of yniich is opaque and reflects white light. Hence its use as a paint. 
It possesses great " covering power," owing to its not transmitting light. 
It is produced by the action of acetic acid, carbon dioxide and water on 
metallic lead; similar basic carbonates, which however, have not the same 
opaque quality, are produced by precipitation. The following have been 
analysed : 


2C0 2 .3PbO.H 2 j 300 2 .4PbO.H 2 ; 

and 5CO 2 .8PbO.3H2O. 

Ferrous salts, on treatment with a soluble carbonate, give a white precipitate 
of presumably basic carbonate. This precipitate rapidly turns green, absorbing 
oxygen : it has not been analysed. With manganese, cobalt and nickel the 
following compounds are known : CO 2 30oO.3H 2 O; 2CO2.5CoO.4H.jO; 
CO 2 .3NiO.6H 2 O (found native, and named emerald nickel) ; 2CO 2 .5NiO.7H 2 O. 

Copper and mercury also form basic carbonates. CO 2 2CuOH 2 O occurs 
native as malachite, a beautiful green mineral, and 2CO 2 3CuO.H 2 O, as 
azurite, which has a splendid blue colour. By precipitation CO 2 6CuO and 
CO 2 .8CuO 5H 2 O are formed as light blue precipitates. Mercuric salts with 
soluble carbonates give a reddish precipitate of CO 2 4Hg*O. 

Titanates, zirconates, and thorates. These have been litole investigated. 
The compounds which have been prepared are : 

; M*TiO 3 ; FeTiO 3 ; CaTiO 3 ; and ZnTiO 3 ; also 
Ti0 2 2ZnOj 2Ti0 2 .32nOj 5TiO 2 .4ZnO. 

The titanates of sodium and potassium are yellowish, fibrous masses, pro- 
duced by heating titanic oxide with excess of carbonate of sodium or potassium. 
On treatment with water they decompose, a sparingly soluble (acid ?) salt 
being precipitated, while a (basic ?) salt remains in solution. Obviously these 
compounds have little stability. Magnesium and iron titanates are produced 
by heating titanium oxide with magnesium chloride, or with a mixture of ferrous 
fluoride and sodium chloride. The iron titanate forms long thin steel-grey 
needles. It is formed native as ilmenite : it is isomorphous with and crystal- 
lises along with iron sesquioxide. The compounds TiO 2 2MgrO and TiO 2 2FeO 
are similarly prepared. Calcium titanate, CaTiO 3 , occurs native as perowsJcite. 

By igniting together zirconia and sodium carbonate, the compound Na 2 ZrOj 
is formed. It is decomposed by water into zirconium hydrate and sodium 
hydrate. A larger amount of carbonate yields ZrO 2 2Ka 2 O ; it is also decom- 
posed by water, and deposits hexagonal crystals of a salt of the formula 
8ZrO 2 Na 2 O 12H 2 O. Magnesium zircon ate has also been prepared by fusing 
zirconium dioxide and magnesium oxide in presence of ammonium chloride. It 
is a powder consisting of transparent crystals. 

Although thorium dioxide dissolves in alkalies, and probably unites with 
oxides, no thorates have been analysed. 

Compounds of sulphides with sulphides. These bodies 
have been investigated only in the compounds of carbon. They are 
named sulphocarbonates or thiocarbonates, from the Greek word for 
sulphur, Qeiov. They are produced by the action of carbon disul- 
phide on sulphides, which is analogous to that of carbon dioxide 
on oxides. Those which have been prepared and analysed 
are : 

NaCS 3 ; KsCSs; (NH 4 ) 2 CS 3 ; MfirCS 3 ; CaCS 3 ; SrCS 3 ; BaCS ;i . 

Precipitates are produced by potassium sulphocarbonate in 
solutions of zinc, cadmium, chromium, iron, manganese, cobalt, 


nickel, tin, lead, bismuth, platinum, silver, gold, and mercury. 
These require further investigation. 

Potassium sulphocarbonate consists of yellow deliquescent 
crystals ; it is formed by digesting an aqueous or alcoholic solution 
of potassium sulphide, K 2 S, with carbon disulphide ; the crystals 
contain water, which is expelled at 80, leaving a brown- red solid. 
On heating it, potassium trisulphide, K 2 S 3 , remains, mixed with 
carbon. The ammonium salt is produced along with ammonium 
sulphocyanide, by digesting carbon disulphide with alcoholic 
ammonia, thus : 2CS> + 4NH 3 = (NH 4 ) 2 CS 3 4- NH 4 CNS. It 
forms yellow crystals, insoluble in alcohol, but soluble in water. 
The calcium and barium salts are prepared like the potassium salt. 
Milk of lime and carbon disulphide give an orange-red basic salt, 
CaCS 3 .2CaO.8H,O ; at 30 it melts to a red liquid, from which 
CaCS 3 .3CaO.10H 2 O separates. The action of carbon oxysulphidcr 
on sulphides requires investigation. 

Compounds of oxides with halides. It has already been 
stated that carbon monoxide and chlorine combine directly ; the 
product is carbonyl chloride, or carbon oxy chloride, GOGk. Its 
vapour-density shows it to have that formula, and not to be a 
compound of CO 2 and CCU. It is produced on exposing a mixture 
of the two gases to sunlight, hence its old name, phosgene gas, a 
gas produced by light (0u>s). It is more easily prepared by passing 
carbon monoxide through hot antimony pentachloride, SbCl 5 , which 
loses two atoms of chlorine ; or by passing a mixture of the gases 
through a tube filled with hot animal charcoal. It condenses to a 
liquid boiling at 8'4.* When treated with water it produces carbon 
dioxide and hydrogen chloride. Assuming the carbon dioxide to 
remain in combination with water, as carbonic acid, the change 
may be thus represented : 

H.OH _ .OH HC1 


Light is thus thrown on the constitution of carbonic acid. 
It appears to consist of carbon monoxide in combination 
with hydroxyl ; and the normal carbonates may be similarly 

represented ; for example, sodium carbonate as ^O^oNa' 


hydrogen sodium carbonate as CO<Q^ a ; calcium carbonate as 

/-\^>Ca j basic copper carbonate as CO<^ /~v~p U >0, each atom 

* For sulphochlorides, see P. Klason, Berichte, 20, 2376. 

u 2 


of copper being half oxide, half carbonate. The more complex basic 
carbonates may also be similarly represented ; e.g., basic lead car- 

CO< OPb \O 
bonate may be written QQ<^ Q>Pb. But such complicated 

formula are not confirmed by any other considerations, and should 
be sparingly used ; moreover, it is impossible to represent the 
various amounts of water in combination with such compounds in 
any way but by simple addition. 

The oxychlorides of titanium have recently been investigated, 
and their formulae appear capable of similar modes of expression. 
Titanium tetrachloride may be supposed to react wittf water form- 
ing the hydroxide Ti(OH) 4 , which, however, appears to be unstable 
(see p. 284). The corresponding carbon hydroxide, C(OH)4, is 
Certainly incapable of existence, but if, instead of hydrogen, it 
contain certain hydrocarbon groups, such as ethyl, C 2 H 8 , it 
becomes stable. For example, the body C(OCv.H 5 ) 4 is known, and 
is named ethyl orthocarbonate, the name orthocarbonic acid being 
applied to the unknown C(OH) 4 . If water containing hydrogen 
chloride in solution (36 per cent. HC1) be mixed with titanium 
chloride, a violent reaction occurs, and a yellow, spongy, very 
deliquescent mass is produced, which has the formula Ti(OH)Cl,. 
It is tolerably stable in aqueous solution. On adding titanium 
tetrachloride to very cold water in theoretical amount, the dihy- 
droxydichloride, Ti(OH) 2 Cl 2 , is produced. It is a yellow deli- 
quescent substance ; and may also be mixed with water. On 
exposing the di- or tri-chloride to moist air for some time, the 
trihydroxymonochloride is formed, Ti(OH) 3 Cl. It has been 
obtained in a crystalline form. It is insoluble in water, but soluble 
in weak hydrochloric acid. We have thus the series : 

TiCl 4 ; Ti(OH)Cl 3 ; Ti(OH) 3 Cl 2 ; Ti(OH) 3 Cl; and Ti(OH) 4 . 

All of these compounds, when heated alone, evolve titanium 
tetrachloride or hydrogen chloride, leaving a residue of dioxide. 

Oxychloride of zirconium, ZrO01 2 , separates in tetragonal 
crystals from a hydrochloric solution of the oxychloride in water ; 
a similar bromide is known. 

A higher oxide of titanium is produced on treating titanium hydrate^ ^ith 
hydrogen dioxide.* It is a yellow substance, the formula of which approxi- 
mates to TiO 3 .3H 2 O. It appears to form compounds with Ti0 2 in the ratios 
4TiO 2 .Ti0 3j 3Ti0 2 .TiO 8 ; 2TiO 2 .TiO 3 ; and TiO 2 .TiO 3 . 

* Chem. Soc., 49, 150, 484. 


Certain fluorine derivatives of this body in combination are also known. 
They are as follows : 

Ti0 2 F 2 .2HF; Ti0 2 F 2 2KP; Ti0 2 F 2 .BaF 2 ; 2TiO 2 F 2 .3BaF 2 ; TiOF 4 ,BaF 2 ; 
and Ti0 2 F 2 .3NH 4 F. 

Attempts to prepare similar zirconium compounds yielded Zr 2 & %H 2 0, as a 
white precipitate ;* and cerium trioxide has been thus prepared as an orange- 
red precipitate.f Thorium yields an oxide of the formula Th 2 7 by similar 

Physical Properties. 
Mass of 1 cubic centimetre : 

C. Ti. Zr. Ce. Th. 

Monoxides. ...... ? 

Dioxides 1'2 1/6J 4 25 5'85 6'93 7 09 10'22 

Ilydrated dioxides 

Monosulphides . . . 1*66 5'1|| 

Bisulphides 1'29 (0) 829 

Carbonates. Li. Na. K. Ca. Sr. Ba. Mg. Zn. Cd. 
1-79 2-4 21 29 3'6 4'3 3'0 44 43 

Tl. Pb. Mn. Fe. Ag. 
7'2 6-5 36 3-8 6*0 

Heats of formation : 

C + = CO + 290K; CO + = C0 2 + 6SOK; C + 20 = C0 2 + 

CO + CT 2 = COC1 2 + 261K; C + + S - COS + 370K; C + 2S = 

CS<i - 260K. 
2NaOH.Aq + COo Na 2 CO a .Aq + 261K; 2KOH.Aq + C0 2 = K 2 CO 3 .Aq 

+ 261K. 
CaO + C0 2 = CaC0 3 + 308K (?) ; SrO 4- C0 2 = SrCO 3 -I- 958K ; 

BaO + C0 2 = BaC0 3 + 622K. 
AgTjO + C0 2 => Aff 2 C0 3 + 200K; PbO + C0 2 = Pb00 3 - 744K. 

* Berichte, 15, 2599. 

t Comptes rend., 100, 605. 

t Solid. 

Artificial; Rulile, 4'42j BrooTcite, 3'89-4'22j Anatase, 375-4'06. 

I! Ce 2 S 3 . 





Oxides, Sulphides, Selenides, and Tellurides of 
Silicon, Germanium, Tin, and Lead. 

The formula of the compounds of this group of elements 
resemble those of carbon and titanium. There are monoxides, 
sesquioxides, dioxides, and intermediate combinations, but no per- 
oxides have been prepared. But a noticeable difference is that the 
monoxides, except ihat of silicon, form compounds ; the com- 
pounds of the dioxides are very numerous ; and we again meet 
with resemblances between the first number of the previous group 
with that of this group; i.e., between compounds of carbon and 
silicon, as we do between those of beryllium and magnesium, and 
of boron and aluminium. 

I. Monoxides, monosulphides, selenides, and tellurides. 

Silicon. Germanium. Tin. Lead. 

Oxygen SiO (?). GteO.* SnO. PbO. 

Sulphur SiS. O-eS. SnS. PbS. 

Selenium ? ? SnSe. PbSe. 

Tellurium PtaTe. 

Sources. Lead monoxide occurs as lead ochre, a yellow earthy 
mineral found sparingly among lead ores. The sulphide is the 
chief ore of lead. It is named galena. It occurs in crystals 
derived from the cubical system, usually rhombic dodecahedra. 
It has a very distinct cubical cleavage, and forms leaden -colpured 
masses with brilliant metallic lustre. It is found in the Isfc\of 
Man, at the lead hills in Lanarkshire, in Cornwall, in the raoun- 
tain limestone of Derbyshire, and in the lower silurian strata of 
Cardiganshire and Montgomeryshire. It is also found in combina- 

* J.prakt. Chem. (2), 84, 177 j 36, 177; Chem. Centralbl., 1887, 329. 


tion with the sulphides of arsenic, antimony, and copper. Lead 
selenide occurs as clausthalite, and the telluride as altaite. 

Preparation. 1. Direct union. The only monoxide obtain- 
able thus is that of lead. It is prepared as massicot by heating 
lead in a reverberatory furnace to dull redness, taking care that 
the resulting oxide shall not fuse, and raking it away as it is 
formed. If the oxide fuses, it forms litharge. The monosulphides 
of tin and lead are also produced directly, by melting the metal 
and adding sulphur. In the case of lead, the mixture becomes 
incandescent owing to the heat liberated during combination. 
Lead sele'hide is similarly prepared. 

2. By heating a compound. Germanous, stannous, and lead 
hydrates, heated in a current of carbon dioxide, lose water, leaving 
the monoxides. If heated in hydrogen, the temperature must not 
exceed 80, else reduction to metal takes place. The dehydration 
of stannous hydrate takes place on boiling water in which it is 
suspended, the condition being the absence of ammonia. Lead 
hydrate suspended in water loses water on exposure to sunlight 
forming crystalline monoxide. Tin oxalate and lead oxalate, 
carbonate, or nitrate, when heated, yield monoxides. 

3. By reducing a higher compound. Silicon monoxide is said 
to have been formed as one of the products of heating silica in the 
Cowles' electric furnace, which is lined with carbon. No doubt it 
would be possible to prepare germanium and tin monoxides from 
the dioxides by careful heating in hydrogen gas ; but the reduc- 
tion is apt to go too far, and to produce metal. Lead dioxide and 
its compounds, when strongly heated, yield monoxides. 

Silicon disulphide, when heated to whiteness, loses sulphur, 
and yields monosulphide ; and germanium disulphide is reduced to 
monosulphide by heating in hydrogen. 

Stannic sulphide, SnS 2 , loses sulphur at a red heat, forming 
monosulphide ; also the sesquisulphide, Sn 2 S 3 . 

4. By double decomposition. Tin and lead monoxides are 
produced by heating their corresponding chlorides, SnCl 2 or PbCl 2 , 
with sodium carbonate. It may be supposed that the carbonates 
first formed are decomposed, leaving the monoxides. 

The sulphides are produced by heating the oxides in vapour of 
carton disulphide, or in the case of germanium, tin, and lead, by 
tr^mting a solution of a salt of the metal or of the hydroxide in 
potassium hydroxide with hydrogen sulphide or some other soluble 

Stannous selenide is best prepared by the action of selenium on 
hot stannoas chloride. 


Properties. Silicon monoxide is said to be an amorphous 
greenish powder ; those of germanium and of tin blackish powders. 
Tin monoxide may be obtained crystalline by heating a mixture of 
the hydroxide and acetate to 133 ; and of a vermilion colour by 
evaporating a solution of ammonium chloride in which the hydrate 
is suspended. Lead monoxide, in the form of massicot , is lemon- 
yellow; it may be prepared pure by strongly heating lead car- 
bonate or nitrate ; and in the form of litharge as a yellowish-red 
laminated mass of crystals. A red variety is produced by heating 
the hydroxide to 110. It, too, can be obtained crystalline by fusion 
with caustic potash ; it separates out in cubes on slow cooling ; if it is 
allowed to deposit from an aqueous solution of potassitfm hydroxide 
it separates first in yellow laminae, and afterwards in red scales. 

Of these oxides, lead oxide is the only one soluble in water ; it 
requires 7000 times its weight of water for solution. 

The monoxides of silicon, germanium, and tin appear to have 
very high melting points ; lead oxide melts at a red heat. 

Silicon monosulphide is a volatile yellow body; that of 
germanium, when obtained by precipitation, forms a reddish- 
brown amorphous powder ; but when prepared in the dry way 
it consists of thin plates, transparent and transmitting red light ; 
but grey, opaque, and exhibiting metallic lustre by reflected light. 
Its vapour- density is normal, corresponding to the formula GeS. 
It volatilises easily. 

Tin monosulphide is a leaden-grey crystalline substance, 
exhibiting metallic lustre. It has also been prepared by electro- 
lysis of a solution in alkaline sulphide, and then forms metallic- 
looking cubes. The precipitated variety is brown and amorphous, 
and is sparingly soluble in alkaline sulphides. It dissolves in and 
crystallises from fused staimous chloride, SnCl*. The selenide 
forms steel-grey prisms. 

The appearance of lead sulphide as galena has been already 
described. When heated it melts, and volatilises a6 a high tem- 
perature. Prepared by precipitation, it is a black amorphous 
powder if the solution be cold ; and if warm, greyish and crys- 

Lead sulphide and oxide react together when heated, yielding 
metallic lead and sulphur dioxide, thus V N 

PbS + 2PbO = 3Pb + 80 2 . 

This reaction is made use of in the extraction of lead from its ores. 
The sulphide when roasted is converted partially into the oxide ; 
and on raising the temperature, metallic lead is produced. 


The selenide is a grey porous mass when artificially prepared ; 
native as clausthallite it forms leaden grey crystals with metallic 

Compounds of the monoxides, <fcc. (a.) With water. 
Silicon monoxide has not been obtained in combination with water. 
The hydrate of germanium monoxide has not been analysed ; it is 
a white precipitate produced on boiling germanium dichloride with 
caustic potash. That of tin monoxide is produced by adding 
sodium carbonate to a solution of tin dichloride ; this precipitate 
is also said to consist of a basic carbonate of the formula 
CO 2 .2Snfcl (see p. 289). 

Hydrate V lead monoxide, prepared by precipitation and dried 
in air, has the formula 2PbO.H 2 O ; and after standing for some 
weeks over sulphuric acid, so as further to dry it, its formula is 
3PbO.H 2 O. The first of these hydrates forms microscopic crystals, 
and the second, lustrous octahedra. 

A mixture of lead hydrate and basic carbonate is produced on 
exposing metallic lead to the action of water and air. Water alone 
has no effect on lead, nor has oxygen ; but together they attack it, 
and as the metal lead is commonly used for water-pipes, the slight 
solubility of the oxide is apt to cause it to contaminate the water. 
It is found that the presence of sulphates, carbonates, and chlorides 
stops this action. 

(6.) Compounds with oxides. No compounds have been pre- 
pared with silicon or germanium monoxides ; but hydrated tin and 
lead monoxides are soluble in sodium, potassium, calcium and 
barium hydroxides, no doubt forming compounds. The calcium 
compound is said to form sparingly soluble needles. A yellow 
precipitate of the formula 2PbO.Ag 2 O is produced by adding 
caustic potash to a mixture of two soluble lead and silver salts ; it 
is probably analogous to the compounds with the former oxides. 
On boiling a solution of stannous hydrate in caustic potash, 
metallic tin is deposited, and a stannate (see below) is formed. 

(c.) Compounds of sulphides with sulphides. Mono- 
sulphides of silicon, germanium, tin, and lead are insoluble in 
solutions of monosulphides of the alkalies, and no compounds 
are known. Compounds of lead sulphide with those of arsenic 
and sAtimony occur in nature, and will be described later on. 

\d.) Compounds with halides. No compounds of silicon or 
germanium monoxides with halides are known ; but stannous 
chloride, SnCl 2 , if dissolved in much water, deposits a white pow- 
der of the formula SnO.SnCl 2 .2H 2 O, according to the equation 
2SnC ! a .Aq + 3H a O = SnO.SnCl 2 .2H 8 O + 2HCl.Aq. 


The same compound is produced by the action of atmospheric 
oxygen on a solution of stannous chloride 

3SnCl 2 .Aq + = SnCl 4 .Aq + SnO.SnCl 2 .2H 2 O. 

The decomposition may be prevented by addition of a soluble 
chloride, such as hydrogen or ammonium chloride, which forms a 
double salt with stannous chloride not decomposed by air, and not 
altered by water (see p. 154). 

The oxyhalides of lead are pretty numerous. A fusible oxy- 
fluoride is produced by heating together fluoride and oxide. Five 
oxychlorides are known, viz. : j* 

PbO.SPbCLj, a laminar pearl-grey substance, obtained by fusion 
of oxide with chloride, and treatment with water. 

PbO.PbCL, found native as matlocJcite in yellowish translucent 
crystals ; and produced by fusing together lead chloride and carbon- 
ate, or by heating lead chloride in air. It is manufactured as a 
pigment by adding to a hot solution of lead chloride, lime water 
in quantity sufficient to remove half the chlorine as calcium chloride. 
It has a white colour, and good covering power. 

2PbO.PbCl 2 , a mineral known as mendipite, forming white, 
translucent crystals. 

3PbO.PbCl } , prepared by fusion ; or by adding a solution of 
sodium chloride to lead oxide. It is a yellow substance, and is 
used as a pigment under the name of Turner's yellow. 

5PbO.PbCl 2 , produced by fusion, is a deep yellow powder, and 
7PbO.PbCl 2 , prepared by heating together litharge and ammonium 
chloride, forms cubical crystals. It is a fine yellow pigment, and is 
known as Cassel yellow. 

Two oxybromides have also been produced, by decomposition of 
the double bromide of lead and ammonium (see p. 154) with 
water, viz., PbO.PbBr 2 .H 2 O, and 2(PbO.PbBr 2 )3H 2 O. The same 
compounds are produced by the action of atmospheric oxygen on 
fused lead bromide, PbBr 2 , but anhydrous. Oxy iodides of the 
formulse PbO.PbI 2 , 2PbO.PbI 2 .H 2 O, 3PbO.PbI 2 .2H a O, and 
5PbO.PbI 2 , are produced by similar reactions. The first of these, 
when prepared by the action of hydrated lead peroxide on potas- 
sium iodide in contact with air, combines with the potassium 
carbonate produced by the action of the carbon dioxide of tltoair 
on the resulting potassium hydroxide, giving compounds of The 
formulse PbO.PbI 2 .K 2 CO 3 2H 2 O, 2(PbO.PbI ? )3K 2 CO 3 .2H 2 O, and 
PbI 2 .2KI.K>CO 3 .3H 2 O ; and by mixing together potassium 
iodide and lead carbonate, the compound PbO.PbI 2 .CO 3 is pro- 


It appears possible also to obtain mixed halides ; one of these 
produced by the action of lead oxide on zinc chloride has the 
formula PbO.ZnCl,. 

II. Sesquioxides and sesquisulphides. Of these compounds, 
hydrated sesquioxides of silicon and tin, sesquioxide of lead, and tin 
sesquisulphide are the only representatives. Their formulae are 
Si 2 O 3 .H 2 O, Sn2O 3 nH,O, Pb 2 O 3 , and Sn 2 S,. The first, Si 2 O 3 .H 2 O, 
from its analogy to the corresponding carbon compound, oxalic acid, 
is sometimes named silico-oxalic acid. The constitution of oxalic 
acid has been noticed on p. 273, and it is probable that the 
analogousS^ilicon compound is similarly constituted. It is pro- 
duced by the* action of ice-cold water on silicon hexachloride ; and 
its formation may be represented graphically thus : 

m TT 

gi< ^0 2HC1 

Cl H.OH = 

X C1 H.OH ^ D HC1 


\ C1 


four atoms of chlorine being replaced by two atoms of oxygen, 
and two by hydroxyl (OH)'. It is a white mass ; but unlike oxalic 
acid the remaining hydrogen of the hydroxyl cannot be replaced 
by metals. It is, therefore, said to be " devoid of acid properties." 
When treated with any soluble hydroxide, it gives a silicate with 
evolution of hydrogen. The compound is, however, of considerable 
interest, inasmuch as it displays the analogy between silicon and 

Hydrated sesquioxide of tin is said to be produced by boiling 
together hydrated ferric sesquioxide, Fe^O 3 .wH 2 O, and stannous 
chloride, SnCl 2 . It is a slimy grey precipitate. 

Lead sesquioxide, Pb 2 O.<, is produced by the action of sodium 
hypochlorite, NaOCl.Aq, on salts of lead, or on a solution of lead 
hydrate in caustic soda ; and also by the action of alkalies on a 
solution of red lead in acetic acid. The last action will be noticed 
below, in treating of red lead. The sesquioxide is a reddish-yellow 
insoluble powder ; it dissolves for a moment in hydrochloric acid, 
but Almost at once chlorine is evolved, and the dichloride precipi- 
tated. No double compounds of sesquioxides are known. 

Tin sesquisulphide is produced by heating three parts of the 
monosulphide with one part of sulphur to dull redness. It has a 
greyish- yellow metallic lustre, and at high temperatures decom- 
poses into monosulphide and sulphur. 


III. Dioxides, disulphides, diselenides, and ditellurides. 

List. Oxygen. Sulphur. Selenium. Tellurium. 

Silicon Si0 2 SiS 2 SiSe 2 ? SiTe 2 

Oermanium GeO 2 GeS. 2 ? ? 

Tin SnO 2 SnS 2 SnSe 2 P 

Lead PbO 9 

These are the most stable compounds with silicon, gejmanium, 
and tin ; lead dioxide, however, easily loses oxygen. , * 

Sources. Silicon dioxide occurs native in hexagonal prisms, 
capped by hexagonal pyramids, as ruck-crystal, bog-diamond, or Irish 
diamond. When coloured yellow or orange by sesquioxide of iron 
it is named cairngorm ; it also also occurs with an amethyst 
colour due to manganese sesquioxide ; and of a rose-red colour (rose- 
quarts) . It is very hard, easily scratching glass. It frequently con- 
tains small cavities, filled with liquid carbon dioxide, often contain- 
ing a minute cubical crystal of sodium chloride. Quartz is a name 
applied to less perfectly crystalline silica, and usually occurs in 
white translucent masses. When perfectly transparent it is used 
for the lenses of spec-tacles, being harder and less easily scratched 
than glass. It is cut into slices by a copper disc, moistened with 
emery and oil, then ground and polished. Flint and chert are 
forms of silica found embedded in chalk, or older limestones, and 
are due to the siliceous spicules of sponges, now extinct. It has 
usually a dull grey-brown colour, owing probably to its containing 
some free carbon, derived doubtless from the animal matter of the 
shell-fish, the remains of which constitute the chalk, for it turns 
white on ignition. Chalcedony is a variety of quartz, not display- 
ing definite crystalline structure, but showing a fibro-radial struc- 
ture, and occurring in kidney-shaped, translucent masses. Varieties 
of chalcedony are named agate, hornstone, onyx, carnelian, catseye, 
chrysoprase, &c. Sandstone consists mainly of water- or air-rolled 
grains of quartz, bound together by a little lime. 

Silica also occurs in combination with many other oxides, as 
silicates. With water, it occurs as opal, an amorphous translucent 
substance, which has been deposited in thin layers. This ^re- 
duces in some specimens a brilliant play of colours, owing to the 
refraction and interference of the light which it reflects. Opal 
is soluble in a hot solution of potassium hydrate ; it is thus dis- 
tinguished from quartz. The other silicates will be considered 


Germanium disulphide, in combination with silver sulphide, 
forms the mineral argyrodite, found in the Himmelsfiirst mine at 
Freiberg. It is almost the only mineral in which germanium has 
been found. 

Tin dioxide, named cassiterite, or tinstone, is the only important 
source of tin. It occurs in veins, traversing the primitive granite 
and slate of Cornwall ; it is also exported from Melbourne. It 
forms translucent white, grey, or brownish quadratic crystals. Its 
crystalline form is the same as that of anatase, one of the forms of 
titanium 'dioxide. 

Stanixte sulphide, SnS 2 , occurs in combination with sul- 
phides of iron and copper, and is named tin pyrites. 

Preparation. 1. By direct union. Silicon dioxide, or silica, 
is formed when silicon burns in air or oxygen. Germanium 
dioxide and stannic oxides are similarly produced. The oxides thus 
prepared are amorphous . Lead unites with oxygen to form mon- 
oxide, PbO, as already mentioned. The highest stage of oxida- 
tion produced directly is that of red lead, P'o 3 04 = PbO 2 .2PbO. 
Stannous oxide also unites directly with oxygen to produce the 

2. By decomposing a double compound. All these oxides 
remain on heating to redness their various hydrates ; germanium 
dioxide has also been prepared from its sulphate, Ge(SO 4 ) 2 , which 
loses sulphur trioxide at a red heat. 

3. Prom lower compounds. Lead monoxide heated to dull 
redness with potassium chlorate is oxidised to the dioxide. The 
potassium chloride and excess of chlorate are dissolved out by 
water. It is also formed by fusing lead monoxide with potassium 
hydroxide. Hydrogen is evolved, and potassium plumbate is pro- 
duced; on treatment with water the dioxide remains in black 
hexagonal tables. 

Tin disulphide and diselenide are prepared by a somewhat 
curious method. When tin and sulphur are heated together, the 
sesquisulphide is the highest sulphide formed. But if ammonium 
or mercuric chloride be heated in a glass retort with the mixture 
of tin and sulphur the disulphide is produced. It is supposed 
that a double chloride of tin and ammonium, or of tin and mercury, 
is fo*st formed, and that this reacts with sulphur, thus : 
2(anCl 2 .2NH 4 Cl) + 2S = SnS 2 + SnCl 4 .2NH 4 Cl + 2NH 4 C1. 
Diselenide of tin is produced by the action of iodine on the 
monoselenide, in presence of carbon disulphide, thus : 2SnSe + 
2 I a = SnI 4 H- SnSe 2 ; at the same time some selenium is liberated, 
according to the equation, SnSe + 2I a = Snl* -f- Se. The tin 


tetriodide dissolves in the carbon disulphide, leaving the di- 
selenide, which is insoluble. 

4. By double decomposition. Tin dioxide is produced in a 
crystalline form by passing the vapours of stannic chloride, SnCl 4 , 
and water through a red-hot tube. The crystals produced are of the 
same form as brookite (Ti0 2 ) : quadratic crystals are formed by 
the action of hydrogen chloride on the red-hot dioxide. The di- 
sulphides of silicon and germanium are both produced by double 
decomposition. To prepare the former, silica, or a mixture of 
carbon and silica, is exposed at a white heat to the/action of 
carbon disulphide; the monosulphide is simultaneously produced, 
probably owing to the decomposition of the disulphide. The 
disulphides of germanium and of tin are precipitated from solutions 
of the dioxides by hydrogen sulphide. Tin disulphide is also 
produced by passing a mixture of hydrogen sulphide and gaseous 
tin tetrachloride through a tube heated to dull redness. 

Properties. The properties of native silica have been already 
described. It fuses at a white heat in the oxyhydrogen flame to a 
glassy mass, which can be drawn into threads. In this form it 
furnishes one of the best insulators for electricity, and has been 
used to suspend the needles of galvanometers. Such threads have 
great tenacity and are very elastic. Even when moist they do not 
conduct. Amorphous silica, produced by heating the hydrate, is 
a loose white powder ; it is said to volatilise when heated to 
whiteness in water-vapour, resembling boron oxide in this respect. 
"While the crystalline form is not attacked by solutions of potas- 
sium or sodium hydroxide, the amorphous variety dissolves 
slowly. Crystalline silica is attacked only by hydrofluoric acid. 

Germanium dioxide is a dense, gritty, white powder, 
sparingly soluble in water, and crystallising from it in small 
rhombohedra. Its solubility is : 1 gram of QeO 2 dissolves in 
247*1 grams of water at 20, and in 95*3 grams at 100. 

Tin dioxide is a white or yellowish powder, insoluble in 
water. It turns dark yellow when heated, but again becomes 
white on cooling. Under the name " putty powder " it finds 
commercial use in polishing stone, glass, steel, &c. 

Lead dioxide is a soft brown powder, insoluble in water; 
when heated to redness it loses oxygen, leaving a residues of 

Silicon disulphide forms long white volatile needles. It is 
remarkable that the oxide is so non- volatile, while the sulphide 
can be sublimed; it leads to the supposition that while the 
sulphide has the formula assigned to it, SiS 2 , the formula of the 


oxide, as we know it, is really a high multiple of Si0 2 . And on 
comparing the silicon and carbon compounds, this conclusion is 
strengthened. For while the boiling-points of carbon dioxide, 
disulphide, and tetrachloride are respectively 78'5, 46, and 76*7, 
an ascending series, we have with silicon, the dioxide melting at 
a white heat, the sulphide easily volatile, and the chloride boiling 
at 58. The order of volatility is reversed. And as it is found 
with compounds of carbon and hydrogen, that the more complex 
the molecule, the higher the boiling-point, we may conclude that 
the non-yolatility of silica is due to its molecular complexity. 
There is at^iresent, however, no means of ascertaining how many 
molecules of Si0 2 are contained in the complex molecule of ordi- 
nary silica, the formula of which must therefore be written 

Germanium disulphide is described as a white precipitate, 
sparingly soluble in 221 '9 parts of water, and also soluble in 
sulphides. It appears not to decompose on solution ; but silicon 
disulphide reacts with water, forming hydrogen sulphide and a 
hydrate of silica. 

Tin disulphide, prepared by precipitation, is a buff- yellow 
powder, insoluble in water. When obtained by the dry process 
it forms golden-yellow scales, and is named " mosaic gold.'* The 
diselenide is a red-brown crystalline powder. 

Double compounds. It is convenient to group these as 
follows : (a) Compounds of the oxides with water and oxides ; 
(/?) oxyhalides and sul phohalides ; (c) compounds of sulphides 
with other sulphides ; and (d) oxysulphides. 

(a.) Compounds with water and oxides. l4ie most im- 
portant of these compounds are the silicic acids and the silicates ; 
allied to them are the stannates and plumbates, and there appears 
to be indications of the existence of germanates. 

General Remarks on the Silicates. The ratios between the 
oxygen of the silica and the oxygen of the metallic oxides com- 
bined with it are very numerous. The silicates form a very large 
portion of the crust of the earth, and they have very varied com- 
position. Among the native silicates the ratio of oxygen in silica 
to that in oxide of metal may vary for monad and dyad metals, 
such^as potassium or calcium, between 2 : 4 and 4:1; or to take 
hypothetical specific instances, between Si0 2 .4K 2 0, or Si0 2 .4CaO 
and 2Si0 2 .K 2 0, or 2Si0 2 .CaO ; and for silicates of triad metals, such 
as aluminium or iron, between 2 : 6, as in Si0 2 .2Al 2 3 , and 12 to 
3, as in 6Si0 2 .Al 2 3 . It must be remembered that the native 
silicates have almost always been formed in a matrix containing 


compounds of many elements ; hence it is rare to find among the 
silicates pure compounds such as those of which the formulae have 
been given above. For instance, it is not unusual to find a silicate 
containing both the metals, potassium and calcium, as oxides com- 
bined with silica ; or the oxides of the metals, iron and aluminium, 
or of cilcium and aluminium, and that not in atomic proportion; 
for we may have a silicate of aluminium containing only a trace of 
iron, or a silicate of calcium containing only a trace of magnesium 
or ferrous iron, the crystalline form of which does not differ from 
that of the pure silicate. It is not to be conceived thgt in such 
instances any given molecule has not, as is usual amona^ompounds, 
a perfectly definite formula ; but it would appear that it is possible 
for an apparently homogeneous crystal to be made up of molecules 
of silicate of aluminium and silicate of iron, or of silicate of mag- 
nesium and silicate of calcium in juxtaposition ; so that, to take a 
suppositions case, a crystal containing 1000 molecules might 
consist of 999 molecules of magnesium silicate and one molecule 
of calcium silicate, or of 998 molecules of magnesium silicate and 
two molecules of calcium silicate, and so on ; oxides of magnesium 
and calcium being mutually replaceable in any proportion what- 
ever. And similarly with the compounds of silica with the sesqui- 
oxides of iron and aluminium. But all oxides are not capable of 
mutually replacing each other. While beryllium, calcium, mag- 
nesium, iron, manganese, nickel, cobalt, sodium, and potassium 
monoxides mutually replace one another, and while the sesqui- 
oxides of aluminium, iron, manganese, chromium, &c., are also 
mutually replaceable, it is found that the place of a monoxide is 
not taken by a sesquioxide, and vice versd. But a silicate may 
contain at once a mixture of sesquioxides and a mixture of mon- 
oxides in combination with silica. 

To deduce the formula of a natural silicate from its percentage 
composition is a problem of some difficulty. It is solved by ascer- 
taining the ratio of all the oxygen combined with dyad metals, 
whatever they may be, to that combined with triad metals, and 
to that contained in the silica. An example will render this clear. 

On analysis, a specimen of the felspar named orthoclase (which is 
essentially a silicate of aluminium and potassium, but which may 
contain iron sesquioxide replacing alumina, and sodium, magne^pm, 
and calcium oxides replacing potassium oxide) gave the following 
numbers : 

SiO 2 . A1 2 8 . CaO. K a O. Na^O. 

65-69 17'97 T34 13'99 I'Ol = 100-00 per cent, 


These constituents contain oxygen in the following propor- 
tions : 

32 48 16 16 16 

Cu-33 102-02 60-u8 94r28 62*09, 

the denominators being- the molecular weights of the oxides, and 
the numerators the oxygen contained in these weights. The ratios 
are, therefore, as folloAvs : 

SiO 2 . A1 2 3 . CaO. K 2 O. 

65-69 >y32 17-97 X 48 1*34 x 16 13-99 x 16 1*01 X 16 

60'33'T 10202 56*08 94-28 62*09 ' 

or 34-84 + 8'45 + 0*38 + 2'37 + 0'26 = 46'30 per cent, of 

We have, therefore, the ratio : ^ 

Oxygen in silica. Oxygen in alumina. Oxygen in lime, potash, and soda. 
34 84 : 8-45 0'38 + 2*37 + 26 = 8'01* 

or 12 : 3 : 1, nearly. 

Hence the formula is 6SiO 2 .Al..Oj.M 2 O, where M stands for 
calcium, potassium, or sodium. It is usually written thus : 
6Si02.Al 2 O 3 .(Ca, K 2 , Na 2 )0, the comma between those symbols en- 
closed within brackets signifying that they are mutually replaceable 
in any proportions. Had iron sesquioxide been present, the oxygen 
contained in it would have been added to that of the alumina, and 
the formula would then have been written, 

6Si0 2 (Al,Fe) 2 3 (Ca, K 2 , Na 2 )O. 

As with the borates, chromates, and carbonates, there are 
two methods of representing the formulae of the silicates. The 
first method is to consider them as addition compounds of 
silica with other oxides, and the formula of orthoclase, as written 
above, is constructed on that principle. It must, however, clearly 
be understood that, inasmuch as we know almost nothing regard- 
ing the internal constitution of such compounds, we can only guess 
at their structure from analogy with the hydrocarbons and their 
derivatives. j 

">he method of writing given above does not imply that the 
compound contains as such the molecular group Si0 2 united with 

* The calculated ratio of oxygen in the above compound is 

SiO 2 . A1 2 O 3 . M 2 O. 

3439 : 869 : 2'87. 


molecular groups A1 2 3 and K 3 0. It is merely a method of show- 
ing the proportions of ingredients which the compound contains 
in an orderly manner, and is better than if we were to write the 
formula, Al 2 K 2 Si60 16 . 

The second method starts from the fact that in such compounds 
silicon is a tetrad element ; that analogous to its compounds 
with fluorine or chlorine, SiF 4 or SiCl4, the typical silicic acid 
has the formula Si(OH) 4 . This substance is named orthosilicic 
acid. Its salts may be supposed to be formed by replacing the 
hydrogen of the hydroxyl groups by metals ; thus the potassium salt 

has the formula Si(OK) 4 , the calcium salt, Si0 4 Ca IT 2 , or 

\ Q >Ca 

and the aluminium salt, 3(Si0 4 ) lv Al 4 m . These are the same as 
Si0 2 .2K 2 0, Si0 2 .2CaO, and 3Si0 2 .2Al 2 3 , and are named ortho- 

Silicates of the formula, SiO 2 .K 2 0, SiO 2 .CaO, &c., are also 
known, and in them the oxygen of the silica beais the ratio to that 
of the oxide as 2 : 1. These may be supposed to be derived from 
the hydroxide Si0 2 .H 3 O, which is named metasilicic acid, and 
which may be regarded as orthosilioic acid deprived of a mole- 


cule of water ; its constitution may be represented Si\OH, and its 


^ ^ 

potassium and calcium salts as Si\~OK, and SKT CL ~ 

N OK \ >Ca. 

It will be remembered that an analogy was drawn between 
chromyl dichloride Cr0 2 Cl 2 , and chromic acid CrO 2 (OH) 2 (see 

p. 268), and it was pointed out that the substance 

might be regarded as partaking of tLe nature both of the 

dichloride and of potassium chromate, Cr0 2 < o -g., being, in fact, 

an intermediate stage. We should expect, therefore, intermediate 
compounds between silicon tetrachloride, SiCl 4 , and silicon tetra- 
hydroxide, Si(OH) 4 . Only one such body is known, viz., SiCl 3 .SH, 
in which hydrosulphuryl replaces hydroxyl. But derivatives of the 
elements of this group are known, which represent similar com- 
pounds connected with metasilicic acid, SiO(OH) 2 . Although the 
corresponding chloride SiOCl 2 is unknown, yet it is represented 


by GeOClj ; and although SixrOH is also unknown, it finds a re- 



preservative in the compound of tin, Sn\~ OH. This method of repre- 


seiitation, which may be termed the method of substitution, is, there- 
fore, justified. 

But we may go still further Hitherto we have been dealing 
with compounds containing only one atom of silicon. It is, how- 
ever, conceivable that two molecules of orthosilicic acid may form 
an anhydride, water being lost between them, thus : 

/ OH 

Sl %)H 
-H Z = (1) Q ; and further (2) 

and (3) >0 

The compound (1) is termed disilicic acid; (2) is the first, and 
(3) the second anhydride of disilicic acid. A representative of 
(1) is serpentine, 2SiO 2 .3MgO ; wollastonite, 2SiO 2 ,2CaO, may be 
a representative of (2), although its formula may equally well be 
the simpler one, SiO 2 .CaO, or SiO(O 2 )Ca; and okenite, 2SiO 2 .CaO, 
may represent the calcium salt of (3). 

A chlorine-derivative of (1), however, is known, viz., Si 2 OCl 6 , 
in which all hydroxyl is replaced by chlorine. That it possesses 
that simple formula is shown by its vapour-density. 

In a similar manner, a trisilicic acid may be derived from 
three molecules of orthosilicic acid, by loss of two molecules of 
water ; it in its turn will yield three anhydrosilicic acids ; and a 
tetrasilicic acid may be supposed to exist, of which four anhydro- 
acids are theoretically capable of existence. Of this tetrasilicic 
acid three chlorine-derivatives have been prepared of the formula, 
Si 4 0,Clio, SiiOiCle, and Si 4 O 6 Cl6, corresponding to the respective 
acids, Si 4 3 (OH) 10 , Si 4 4 (OH) 8 , and Si 4 O 6 (OH) 6 , as shown by their 
vapour-densities. The first is tetrasilic acid itself ; the second and 
third its first and second anhydrides respectively. Salts of even 
more condensed silicic acids may exist. 

Many silicates are known, containing more base than that 

x 2 


contained in orthosilicates, in which the ratio is Si0 2 .2M''0. For 
example, collyrite has the formula Si0 2 .2Al 3 3 , the ratio of oxygen 
in the silica to that in the oxide being 2 : 6. Such silicates are 
termed basic. Their formulae may be written in an analogous 
manner, on the supposition that the metal exists partly as oxide, 
partly as silicate. Thus the above compound may be represented 
thus : 


AtzO ; 
X) AlziO .j 

each atom of aluminium being one-third ortho-silicate, and two- 
thirds oxide. And so with other instances. 

These remarks must be held to apply also to the titanates, 
zirconates, stannates, and plumbates ; but similar compounds of 
tin and lead are not numerous, 

One point must still be noticed before proceeding with a 
description of the silicates, viz., the question as to whether or not 
water, occurring in combination with a silicate or stannate, should 
be included in the formula. For example, by including water, a 
compound of the formula Si0 2 .CaO.H 2 may be represented as 

/o H 

an orthosilicate, SiXV>,^>Ca, or, excluding the water, as a meta- 

silicate, Si;rO^Q .H 2 0, the water being regarded as water of 


crystallisation. There is no rule for guidance in discriminating 
water of crystallisation from combined water ; and indeed we have 
no reason to regard water of crystallisation as combined in any 
other fashion than other oxides. At present, however, no satis- 
factory theory has been devised whereby water of crystallisa- 
tion can be rendered a part of the formula, like the molecule of 
water in the first example given above ; and in the present state of 
our knowledge the only course is to exercise discretion as regards 
its inclusion or exclusion. 


Silicates, Stannates, and Plumbates. 

SiO 2 .2H 2 O (?) - Si(OH) 4 ;' SiO 2 .H 2 O - SiO(OH) 2 . aeO 2 .H 2 O. 
Sn0 2 4H 2 (?) - Sn(OH) 4 ; SnO^O (?) - SnO(OH) 2 ; 3SnO,.H 2 ; 

7SnO 2 .2H 2 O ; 5SnO 2 .5H2O. 
Pb0 2 H 2 j 3Pb0 2 .H 2 0. 

These compounds are very indefinite. On addition of dilate 
hydrochloric acid to a dilute solution of sodium or potassium sili- 
cate, na precipitate is produced. Placing this solution in a 
dialyser- a circular frame, like a tambourine, covered with parch- 
ment or parchment paper or bladder, and floated on water the 
crystalline sodium chloride passes through the diaphragm, while 
the colloid (glue-like) non-crystalline silicic acid remains behind 
for the most part. It was suggested by Graham, the discoverer of 
this method of separation, that the molecules of the colloid body are 
much more complicated and larger than those of the crystalline 
substance, and hence pass much more slowly through the very 
minute pores of the dialyser. To such passage Graham gave the 
name osmosis, and the general phenomenon is termed diffusion. 
Recent researches appear to confirm this view, and to show that 
the molecules of colloid bodies are very complex. It is supposed 
that the silicic hydrate thus remaining soluble is orthosilicic acid, 
Si(OH)4, inasmuch as it is produced from an orthosilicate. To 
obtain it pure, the water outside the dialyser must be frequently 
renewed. A solution of silicic acid containing 5 per cent, of Si0 2 
may thus be prepared ; and by placing it in a dry atmosphere o*er 
sulphuric acid, it is slowly concentrated until it reaches a strength 
of 14 per cent. 

It forms a limpid colourless liquid, with a feeble acid reaction. 
When warmed, it gelatinises ; this change is retarded by the 
presence of a small amount of hydrochloric acid, or of caustic soda 
or potash ; but is furthered by the presence of a carbonate. 

Similar results were obtained from a stannate mixed with 
dilute hydrochloric acid, and also from a titanate. The solutions 
have similar properties. 

It is supposed that the gelatinous substance produced from 
orthosilicic acid is metasilicic acid, SiO(OH) 2 . When dried for 
several months over strong sulphuric acid, it corresponds with 
that formula. This hydrate is also supposed to be produced when 
a halide of silicon reacts with water. A convenient method of 
preparation consists in leading silicon fluoride into water (see 
p. 153). It is said to have been obtained in crystals of the 


formula SiO(OH) 2 .3H 2 O, by the action of hydrochloric acid on a 
siliceous limestone. 

On drying precipitated silica for five months over sulphuric 
acid, it had the approximate formula 3SiO 2 .4H 2 O. When the 
temperature was raised, it lost water gradually, but no evidence of 
any definite hydrate was obtained ; no point could be found at 
which a small rise of temperature did not produce a further loss of 
water. The same remarks apply to stannic hydrate. But about 
360, the substance, which had the composition 3SnO 2 .H. 2 O, and a 
dirty brown colour, displayed a sudden change of colou/* to pale 
yellow, and had then the formula 7SnO;.2H 2 O. /* 

When metallic tin is treated with strong nitric acid, it is 
oxidised ; copious red fumes are evolved, and a white powder is 
produced. While ordinary hydrate, prepared by precipitation, is 
sol able in acids, this white substance is not ; and after drying in a 
vacuum or at 100 it has the formula 5SnO 2 .5H 2 O (see below). 

Hydrated lead peroxide, dried in air, has the formula 

On further heatinor, water and then oxygen are lost. 

By passing a current of electricity between two lead platos, 
dipping in dilute sulphuric acid, hydrated peroxide of lead is 
formed on the positive, while hydrogen is evolved at the negative 
plate. This hydrate has the formula PbO 2 .H 2 O. Such an 
arrangement gives a very powerful current, lead peroxide being 
very strongly electro-positive ; and it is made use of for " storage 

SiO 2 .2Li 2 O; SiO,.Li 2 O ; 5SiO 2 .Li 2 O. SiO 2 .2Na 2 O(?) ; 

5SiO 2 .2Na 2 O; 4SiO 2 Na 2 O. SiO 2 .2K. 2 O (?) ; SiO^O ; 9SiO 2 .2K 2 O, 

or perhaps 4SiO 2 .Kl 2 O. 
Sn0 2 .Naa0.3, 8, 9, and 10H 2 O : SnO 2 .K 2 O.3H 2 O ; 2SnO 2 .(NH 4 ) 2 O.wH 2 O. 

PbO 2 .K 2 O.3H 2 O, and others. 

When silica is fused with a carbonate or hydroxide of lithium, 
sodium, or potassium, a glass is formed of indefinite composition, 
depending on the proportions taken. The lithium glass, however, 
dissolves in fused -lithium chloride, and crystallises out on cooling. 
The lithium chloride withdraws lithia from the silicate, forming 
oxychloride ; and by keeping the mass fused for different times, 
the three compounds given above are formed. 

Soluble glass, orwater-yla8s,is a silicate of sodium or potassium. 
It is prepared as mentioned ; or by heating silica (quartz, flint, 


sand, &c.) with solution of caustic soda or potash under pressure. 
The proportion of silica and potash usually corresponds with the 
formula 4SiO 2 .K 2 O ; on treating the solution with alcohol, a sub- 
stance of that formula is thrown down; it is suggested that the 
more probable formula is 9SiO 2 . k 2K 2 O. It is probably a mixture. 
If the sodium silicate be saturated with silica, 4SiO a .Na 2 O, is 

Soluble glass is a syrupy liquid, obtained by dissolving the 
product of fusion in water, or by evaporating the solution of silica 
in alkaline hydroxide. It is used to form artificial stone ; for it 
reacts with calcium hydrate or carbonate, giving insoluble calcium 
silicate, which may be used to bind together large amounts of sand 
into a coherent mass. It is also employed in mural painting ; and 
it is added to cheap soaps. 

Hydrated silica dissolves to some extent in solution of am- 
monia, but no solid compound has been obtained. 

Decomposition of silicates. The usual method of decomposing insoluble 
silicates is by fusing them with a mixture of sodium and potassium carbonates, 
named " fusion-mixture." Carbonates or oxides of the metals remain, and the 
silica combines with the sodium and potassium oxides, forming a mixture of 
silicates. This mixture is now treated with water, when the silicates of the 
alkalies pass into solution, and may be removed from the insoluble oxides by 
filtration. But as it is usually required to separate the silica, it is more common 
to add hydrochloric acid, which, if the solution be strong, precipitates gelatinous 
silicic acid, and converts tho oxides of the metals into chlorides. On evapora- 
tion to dryness and heating, the silicic acid is decomposed into water and silica, 
and after re-evaporation with a little hydrochloric acid, it is m?oh\bJo in dilute 
hydrochloric acid, which dissolves the chlorides of the instate, and thoy may then 
be separated by filtration. On ignition, the silicp, remains as a wbte loose 
powder, and if required it may be weighed. 

The corresponding stannates are prepared by fusing tin 
dioxide with hydroxide, sulphide, nitrate, or chloride of sodium or 
potassium ; or by heating metallic tin with a mixture of hydroxide 
and nitrate, from the latter of which it derives its oxygen. On 
treatment with water the mass dissolves, and on evaporation 
deposits crystals containing 3, 8, 9, or 10 molecules of water, 
according to circumstances. The salt with 3H 2 may also be 
precipitated by adding alcohol. 

Stannic hydrate is soluble in ammonia, forming a jelly, in 
which the ratio of Su0 2 to ammonia corresponds with the formula 
2Sn0 2 .(NH 4 ) 2 O. 

Metastannates. By boiling the product of the action of nitric 
acid on tin, 5SnO 2 .5H 2 O, with sodium or potassium hydroxide, a 


solution is obtained, from which, if caustic soda be used, granular 
white crystals deposit on cooling, of the formula 
5SnO 2 .Na 2 O.4H 2 O. 

If potash be used, a similar compound, 5SnO,>.K 2 O.4!H 2 O, is pre- 
cipitated by addition of excess of potash, in which it is insoluble. 
It is a gummy non-crystalline substance. Both of these com- 
pounds are decomposed by boiling water into alkali and meta- 
stannic acid. It is the fact that one-fifth of the water is replaced 
by sodium or potassium oxide, which leads to the formula 
5SnO 2 .5H 8 O for metastannic acid, instead of SnO 2 .H^O, which 
would more simply represent its percentage composition. 

On mixing metastannic acid, dissolved in hydrochloric acid, with 
caustic potash, until the precipitate at first produced redissolves, 
and then adding alcohol, a precipitate of the formula 

7SnO 2 .K 2 O.3H 2 O 

is produced. There appear also to be other analogous substances. 
Plumbates. By fusing lead dioxide with excess of caustic 
potash, it dissolves ; the solution of the product, in a little water, 
deposits octahedral crystals of the formula PbO 2 .K 2 O.3H 2 O 
analogous to the stannate. By fusing litharge with potassium 
hydroxide, the compounds PbO 2 .K 4 O and 3PbO 2 ,2K 2 O, are formed 
with absorption of oxygen. These salts are decomposed, on treat- 
ment with water, into potassium hydroxide and hydrated lead 
dioxide ; they are stable only in presence of excess of alkali. 

SiO 2 .2BeO (pkenacite, beryl, emerald) ; SiO 2 .CaO (wollastonite) ; 

28iO 2 .CaO.2H 2 O (okenite) ; 3SiO 2 .2CaO.H 2 O (gyrobteYi 
SnO,.CaO, also 5H 2 O ; 2SnO 2 .: SrO.10H 2 O ; SnO 2 .2BaO.10H 2 O.* 

These silicates are found native ; they are well crystallised 
minerals. By adding to solutions of calcium, strontium, or barium 
chlorides a solution of sodium or potassium silicate, white curdy 
insoluble precipitates are produced of the respective silicates, the 
composition of which is analogous to that of the alkaline silicate 
from which they are produced. 

Of the native silicates, phenacite is an orthosilicate ; wollastonite 
probably a metasilicate ; dkenite a salt of disilicic acid, Si 2 0(OH) 6 ; 
and gyrolite, of the second anhydride of trisilicic acid, Si 3 4 (OH) 4 . 
And with the stannates, we have barium orthostannate ; calcium 
metastanuate (rejecting the water) ; and the strontium salt of 
distannic acid. 

* Comptes rend., 96, 701. 


Two compounds are known, the first occurring native, a 
titanate and silicate of calcium, named sphene ; and the second, of 
similar crystalline form (monoclinic prisms) produced by heating 
a mixture of silica and tin dioxide with excess of calcium chloride 
to bright redness for eight hours. These bodies are derived from 
a compound analogous to the second anhydride of disilicic acid. 
Their formulae are probably 

Ca<>Si<g>Ti<>Ca, and Ca<>Si<>Sn<>Ca, 

Similar silico-zirconates occur native. 

Ordinary Mortar is made by mixing sand with slaked lime. The rapid 
setting of the mortar is, however, not due to the combination of the calcium 
and silicon oxides, but to the formation of the compound CO 2 .2CaO, by absorp- 
tion of carbon dioxide from the air. But, after the lapse of years, combination 
of the silica does take place, and very old mortars contain much calcium 

Hydraulic mortars, as those mortars are named which " set" under water, 
on the other hand, cannot be produced from anhydrous silica. A mixture of 
precipitate silica or of crushed opal and lime hardens under water ; but the best 
hydraulic mortars are made from hydrated silicates of alumina. The celebrated 
pozzolana of Naples is such a substance. When mixed with lime, there is 
formed a silicate of aluminium and calcium, which is rapidly produced, and 
perfectly insoluble in water. Tufa, pumice, and clay-slate form similar insoluble 
mortars. Marl, a mixture of clay and calcium carbonate, after ignition, sets 
when moistened. This is probably in the first instance due to hydration, and 
subsequently to the formation of a silicate of aluminium and calcium. 

SiO. 2 2(Mgr, Fe)O (chrysolite, oHmne) ; SiO 2 .(Mgr, Fe", Mn", Ca)O (augite 
and hornblende ; these differ in crystalline form, but are both monoclmic) ; 
2SiO 2 .3(Mgr, Fe")O.2H 2 O (serpentine, sometimes containing Na 2 , Mn", and 
Ni"); 3SiO 2 .2MgO.2H 2 O,or4H 2 O (meerschaum)-, 5SiO 2 4M#O (talc, contains 
a little water. SiO 2 .2ZnO.HO (siliceous calamme) \ SiO 2 .2ZnO (willemite). 
2SnO 2 .3ZnO.10H 2 <X 

These silicates are all found native and, as a rule, crystalline. 
Chrysolite and willemite are orthosilicates ; siliceous calamine 
possibly a basic metasilicate of the formula SiO.(OZnOH) a ; 
augite and hornblende are metasilicates, but one is probably a 
polymeride of the other, possibly a derivative of the disilicic acid, 


r, like sphene, with which, however, neither is 

isomorphous. Serpentine is a derivative of disilicic acid, and 
meerschaum and talc of tri- and penta-silicic acids respectively. 

The silicates of boron, aluminium, ferric iron, &c., are very 


numerous, and it is here impossible to do more than give a sketch 
of their nature. 

Datolite has the formula Si0 2 .B 2 O 3 .CaO ; and botryolite contains, 
in addition, two molecules of water. They are doubtless derived 
silicates of boron and calcium, and may be constituted thus : 


BCK Q ./0 CaOH 
\ > bl <OH. , 

XenoUte is aluminium orthosilicate, 3SiO 2 .2Al 2 3 . A number 
of minerals, including fibrolite, topaz, muscovite, paragonite and 
eucryptite (varieties of mica), dumortierite, grossularite (a lime 
alumina garnet), prehnite, and natrolite (or soda mesotype), may be 
simply derived from it ; the following structural formulae show 
their relations (Clarke) : 

Alr-Si0 4 =Al 

Xenolite. Fibrolifce. Topaz. 

Dumortierite. Muscovite. Paragonite. 

^Si0 4 ~ " 

Eucryptite. Natrolite. Grossularite. 

/OH _ ,BO Z 

Prehnite. Kaolin. 

Kaolin or china-clay, is, it will be seen, partly hydrate of 
aluminium. R in the last formula may be calcium, iron (dyad), 
magnesium, sodium, or potassium, or generally a mixture. 

The metasilicates may be similarly represented. Among these 
are pyropJiyllite, 4SiOi.AlaO3.HaO and spodumene, jade, and leucite 
containing lithium, sodium, and potassium respectively. They may 


all be represented by the formulae Al-^V,-^ 3 , where B stands 
r * o LUs iv 

for H, Li, Ma, or K. 


There are at least two other silicic acids, 2SiOi.H 2 0, or 
Si 2 3 (OH) 2 , the second anhydride of disilicic acid, and 3SiO z .2HA 
or Si,0 4 (OH) 4 , the second anhydride of trisilicic acid, which yield 
salts. Petalite, (Si 2 3 ) 2 Al.Li, is a salt of the first, and the felspars albite, 


orthoclase, Si 8 4 <\ of the second. 

By trebling these formulas we obtain groups analogous to those of 
the orth^silicates ; and this shows a striking analogy between 
these and other minerals, otherwise difficult to classify. Thus 

/Si0 4 = ySi 3 8 ^ 

analogous to Alx-SiO 4 z=Al, we have Al\-Si 8 O=Al, The calcium 

X Al N Si 3 8 ~Al 

salt corresponding to the first formula and the sodium salt of the 
second are respectively the minerals 

xSi0 4 = Ca 3 ~Si0 4V /Si 3 O 8 =Na 3 

Al , and 

, ~ 38 
4 =Al Al=SiO x Si 3 8 =Al 

Anorthite. Albite, 

/Si/) 8 =(Al(OM) 2 )3 

Labrador ite. 

If potassium rep 1 aces the sodium of albite, the mineral is orthoclase, 
or potash felspar.* 

Lastly, an instructive analogy is pointed out by Clarke, which 
promises to throw light on a curious compound of a brilliant bine 
colour, found native, and named lapis lazuli, which is now manu- 
factured in large quantities as a paint under the name of ultra- 
marine, by heating together sodium sulphate, sulphur, felspar, and 
some carbon compound such as resin. The mineral sodalite has 
the formula 

4SiO,.4Al 8 0,.2Na 9 O.NaCl. 

Ultramarine maybe represented as the sodo- sulphury 1 (SN"a) com- 
pound of which sodalite is a chloride ; and the analogy is streng- 
thened inasmuch as the constitution of another mineral, nosean, 
closely allied to ultramarine, is thereby represented. The formulas 

* F. W. Clarke, Amer. Jour, of Science, Nov., 1886 j Aug., 1887 : Amer. 
Chem. ,7.,10, 120 j 38,384. 


Al SiO<=Al. Al 

N C1 S0 4 Na. S Na. 

Sodalite. Nosean. Ultramarine. 

Such are some of the attempts which have been made to 
classify these complex silicates. Whether they are justified or 
not, if they serve to connect together bodies resembling each other, 
and to point the way to new researches, they have their fise. 

A few other silicates have still to be mentioned. 

SiOo.2MnO (tephroite) ; SiO 2 .MnO (rhodonite) ; SiO 2 .CuO.2H 2 O (chryso- 
colla) ; 3SiO 2 .2Ce 2 O 3 (cerite) j 3SiO 2 .2(Y, Ce, Fe, Mn, &c.) 2 O 3 (gadolinite) ; 
3(SiO 2 .ThO 2 )4H2O (thorite)^ 

After what has been said, these may easily be grouped in their 
respective classes. Other stannates have also been prepared, for 

SnO 2 .KiO.5H 2 O ; SnO 2 .CuO.6H 2 O ; SnO 2 .CuO.4H2<> } 
SnO 2 .CuO(NH 4 ) 2 O.2H 2 O ; SnO 2 .Aff 2 O. 

The germanates have not been investigated. But as dioxide of 
germanium is soluble in excess of caustic alkali, they are, without 
doubt, capable of existence. It has lately been announced that 
germanium exists in small amount in euxenite ; and it is present, 
no doubt, in the form of a germanate. 

Double Compounds of the Sulphides and Selenides. 

Those of tin alone have been investigated. 

Stannous sulphide, SnS, when treated with a /very strong 
solution of potassium sulphide, K 3 S, dissolves ; while tin precipi- 
tates in the metallic state. The equation is 

2SnS + K 2 S.Aq = SnS a .K 2 S.Aq + Sn. 
By further action, hydrogen is evolved, thus 
Sn + 3K 2 S.Aq 4- 4H 2 = SnS 2 .K 2 S.Aq +4KOH.Aq + 2# 2 . 

The same compound is also produced by warming stannous 
sulphide with the polysulphide of an alkali, e.g., K 2 S 5 .Aq, or 
(NH 4 ) 2 S 6 .Aq; the monosulphide is then converted into the di- 
snlphide which dissolves in the solution of sulphide ; or, more 
simply still, tin disulphide may be dissolved in a solution of potas- 
sium sulphide. 

The hydrogen salt of sulphostannic acid, SnS 2 .H 2 S, or 
SnS(SH) (it will be noticed that this is a meta-acid), is produced 
on adding an acid to a sulphostannate, as a yellow precipitate, which 


becomes dark-coloured on exposure to air. The following salts 
exist ; they are all prepared thus, and are soluble in water : 

SnS(SNa) 2 .3H 2 0, also 2H 2 O ; SnS(SK) 2 .10H 2 O, also 3H 2 O ; 
3SnS 2 .(NH 4 ) 2 S.6H 2 O; SnS(S 2 )Ba.l4H 2 O ; SnS(S 2 )Sr.l2H 2 O j 
and SnS(So)Ba.8H 2 O. 

SnSe(SeK) 2 .3H 2 O has been analysed ; and a mixed compound 
obtained by digesting potassium sulphide with tin and selenium 
has the formula SnSe 2 .K 2 S.3H 2 O.* It would appear that two iso- 

merides might here exist, viz., SnSe<|^, and SnS<|^ ; but 

they have not been identified. 

A native sulphostannate of copper, iron, and zinc is known as 
tin-pyrites. Its formula is SnS 2 .(Cu 2 , Fe, Zn)S. It is also a 

(6.) Compounds with halides. These are, as a rule, difficult 
to prepare, for almost all are acted on by water. No compound of 
the formula SiOCl 2 is known. The corresponding germanium 
compound, GeOCl 2 , is produced by distilling germanium tetra- 
chloride in contact with air. It is a colourless, faming, oily liquid 
boiling above 100. By passing a mixture of silicon tetrachloride 
and hydrogen sulphide through a red-hot tube the substance 
SiCl 3 .SH is formed. It boils at 196. The sulphochloride, 
SiSCl 2 , is said to be formed by the same process ; probably the 
former compound dissociates at a high temperature into hydrogen 
chloride and the latter. 

Silicon tetrachloride, SiCl 4 , led over fragments of felspar con- 
tained in a white-hot porcelain tube, deprives the felspar of 
oxygen, and yields the oxychloride (Si01 3 ) 2 0. It is a liquid boiling 
at 136 139 U ; and this compound, passed through a hot glass tube 
along with oxygen, yields a liquid from which, on fractionation, the 
following compounds have been isolated : Si 2 OCl 6 (136 139) ; 
Si40 3 Cl 10 (152154); SiACl 8 (198202). These substances 
give vapour- densities which lead to the formulae ascribed. A fourth 
is formed which, on analysis, gives the numbers for Si 4 O ft Cl 6 ; the 
molecular weight of such a body should be 405 ; that deduced from 
its vapour-density was 614 ; its formula is therefore doubtful. It 
boiled at a very high temperature.f 

A somewhat analogous body to the compound SiCl 3 (SH) is the 
substituted orthostannic acid produced by the action of water on 

stannic chloride. Its formula is SnO<^ . It may be regarded 

* Comptes rend., 95, 641. 

f Troost and Hautefeuille, Annales (5), 7, 453. 


as metastannic acid, with one hydroxyl-group replaced by chlorine. 
On treatment with ammonia it yields the salt SnO<, 4 . 

In conclusion, a set of curious compounds of carbon and silicon 
with oxygen and sulphur may be mentioned, which require further 
investigation.* The first of these is a greenish- white mass pro- 
duced by the action of carbon dioxide on white-hot silicon. Its 
formula is SiCO. Vapours of hydrocarbons passed over silicon, 
heated in a porcelain tube, yield a bottle-green substance of tbe 
formula SiC0 2 , the oxygen being derived from the tube. ( * By sub- 
stituting a mixture of carbon dioxide and hydrogen the substance 
Si 2 CsO is produced ; and by the action of silicon and carbon at a 
white heat on porcelain, a body of the formula Si 2 C 3 2 is formed. 
No clue has been obtained regarding the constitution of these 

Here also may be mentioned a very remarkable compound of 
carbon monoxide with nickel, produced by passing that gas over 
hot finely-divided nickel, and condensing by means of a freezing 
mixture. It has the empirical formula Ni(CO) 4 , and is a colour- 
less liquid, boiling about 45. Its vapour density corresponds 
with the formula given. It deposits metallic nickel when heated 
to 180, as a brilliant mirror, j 

Physical Properties. 

Weights of 1 cubic centimetre : 

Si. Ge. Sn. Pb. Si. Ge. Sn. Pb. 

O.. 2-89 - 6-06-6 8-74J 9'29 2 .. 2'65|| 6'7 8'9 

S .. 5-0 7'5 S 2 . -* 4-6 6-31T 

g e ,. _ _ 6-2 8'1 Be,.. 5-1 

Te.. 6-5 8-1 Te,., 

Heats of formation : 

Si + 20 + Aq = 
Sn + 20 Sn0 2 

Si0 2 .Aq + 1779K (?). 
+ 1400K (?) 

Sn + O SnO ., 

+ 700K(P). 

Sn + + H 2 - 
Pb + PbO 

SnOgBL, + 681K. 
' + 503K. 

Pb + s PbS .. 

+ 184K. 

* Comptes rend., 93, 1508. f C*^^; 
|| Rock crjstal at 10 : Tridymite, 2 

n. Soc., 57, 749. J Red. Yellow. 
3 -, fused to glass, 2 '22. f Pb 2 H a . 






Oxides, Sulphides, Selenides, and Tellurides of 
Nitrogen, Vanadium, Niobium, and Tantalum. 

These are very numerous. The compounds of nitrogen are not 
formed by direct union, for heat is absorbed during their formation 
and they therefore are easily decomposed. Those of vanadium 
niobium, and tantalum, on the other hand, are very stable. 

List of Oxides. 

Nitrogen. Vanadium. Niobium. Tantalum. 

Monoxides N$O 

Dioxides NO VO* NbO 

Trioxides N 2 O 3 "V^Oj 

Tetroxides N0 2 ; N 2 O 4 VO 2 * NbO 2 TaO 2 

Pentoxides N 2 O 3 V 2 O 5 Nb 2 O 6 Ta 2 O 6 

Hexoxido N 2 O 6 

List of sulphides, selenides, and tellurides: 

NS; NSe; VS 2 ; V 2 S 6 ; TaS 2 (?). 

Sources. None of these bodies occurs native. The pentoxides 
occur in combination with the oxides of metals in the nitrates, 
vanadates, niobates, and tantalates, which will be described later. 
Among the most important are nitrates of sodium and of potas- 
sium, named respectively Chili saltpetre and saltpetre or nitre; 
vanadinite, a vanadate and chloride of lead ; pyrochlore, a niobate 
of calcium, cerium, &c. ; euxenite, a niobate and tantalate of 
cerium, yttrium, &c. ; and tantalite, a tantalate of iron and man- 

Preparation. The starting-point for the preparation of all 

* As the molecular weight of these bodies is unknown their simplest formulae 
are given. t 


the oxides of the members of this group is the compounds of the 
pentoxides with other oxides. For nitrogen oxides, the nitrates of 
potassium and sodium; for vanadium oxides, the vanadate of lead; 
for the oxides of niobium and tantalum, the niobates and tantal- 
ates of yttrium, lanthanum, iron, manganese, &c. On treatment 
of these compounds with strong sulphuric acid, hydrates of the 
pentoxides are set free. This may be regarded as the displace- 
ment of an oxide by another oxide, viz., S0 3 . As nitric acid, 
N 3 O 5 .H 2 O, or as its vapour-density shows us, HNO 3 , is a liquid, vola- 
tile without decomposition, it can be distilled away from/ the solid 
sulphate of sodium or potassium ; the vanadate of lead, on treat- 
ment with sulphuric acid, or, better, on fusion with hydrogen 
potassium sulphate, HKSOi, is decomposed, lead sulphate, which 
is insoluble in water, being left behind ; and on treatment with 
water vanadate of potassium is dissolved, from which strong nitric 
acid sets free vauadic acid, Y 2 O 6 .H 2 O, as a reddish precipitate. 
The pentoxides of niobium and tantalum are also produced by fusing 
the ores with hydrogen potassium sulphate, and after cooling, 
boiling the fused mass with water ; the iron, yttrium, &c., all go 
into solution as sulphates, and the pentoxides remain as insoluble 

We shall begin reversing the usual order with the pent- 
oxides, because they form the sources of the lower oxides. 

Pentoxides. Nitrogen pentoxide is produced by the action 
on nitric acid of phosphoric anhydride, P 2 O 5 , a body which has a 
great tendency to combine with water, and which, therefore, with- 
draws it from nitric acid. The acid cannot be dehydrated by heat 
alone, for the pentoxide easily decomposes into the teteoxide, losing 
oxygen. Phosphorus pentoxide is gradually added to ice-cold, 
pure nitric acid, and the syrupy liquid is distilled at a low tem- 
perature. The liquid distillate consists of two layers, the upper 
one being the pentoxide, mixed with a little of the compound 
2N^O 6 .H..O ; the lower consisting of the latter compound. The 
upper layer is separated, and cooled with a freezing mixture, when 
the pentoxide deposits in crystals. The equation is : 

P 2 5 = 2HPO 3 + N 2 O 6 . 

This substance may also be prepared by heating silver nitrate, 
AgNO 3 , to 58 68 in a current of perfectly dry chlorine. This reac- 
tion should yield a hexoxide, N 2 6 , thus, AgNO 3 -f OZ 2 = 2AgCl 
4- N 3 O 6 ; but the hexoxide is unstable, and decomposes at the moment 
of liberation into pentoxide and oxygen. The hexoxide is said, 
however, to be produced by passing an electric discharge through 


a mixture of nitrogen and oxygen at 23, and to form a volatile 
crystalline powder. 

Another method, which appears to act well, is to pass a mixture 
of nitric peroxide, NO 2 , and chlorine over dry silver nitrate at 
6070. The equation is A T 0, + Cl + AgNO 3 = AgCl + N 2 6 . 

The pentoxide must be condensed in a (J-tube, surrounded by a 
freezing mixture ; and the most scrupulous care must be taken to 
exclude moisture, by drying the apparatus and materials perfectly 
before use, and by preventing the access of moist air. 

Vansfdium, niobium, and tantalum pentoxides are produced 
(1) by burning the elements in air, or by the oxidation of the lower 
oxides when heated in air. (^) By heating their hydroxides 
(arids), or in the casa of vanadium by heating ammonium vanadate, 
2NH 4 VO, = 2NH, + H,0 + V 2 O a ; or by heating a solution of 
the oxide VO 2 in strong sulphuric acid ; the first reaction is the 
formation of the sulphate, V 2 O.,. k 2SO 3 .H 2 O, a portion of the sul- 
phuric acid losing oxygen to oxidise the tetroxide to pentoxide, 
thus 2VO 2 + H^SO 4 = V 2 O 6 -f H 2 O + SO Z ; the pentoxide then 
forms the above sulphate ; V 2 O 5 + *H 2 SO 4 = V 2 O a .2SO, + 2H a O. 
The sulphate is decomposed on further ignition into V 2 O 6 and SO*. 
(3) By the action of water on the pentahalide or oxyhalides. This 
yields the hydroxides, from which the oxides are obtainable. 

Properties. Nitric pentoxide forms brilliant, colourless, 
transparent rhombic prisms; it melts at 30, and boils about 45. 
It is very unstable, forming nitric peroxide with loss of oxygen, but 
can be preserved for several days at 10 in a dry atmosphere. It 
hisses when dropped into water, forming hydrated nitric acid. 

Vanadium psntoxide is a reddish-yellow solid, sparingly 
soluble in water, to which it imparts a yellow colour. The solution 
is tasteless, but has an acid reaction. It melts when heated to 
redness, and on solidifying it turns incandescent, probably display- 
ing the phenomenon of supervision. 

Niobium pentoxide is a white insoluble solid, turning yellow 
when heated, but regaining its whiteness on cooling. It has been 
fused at a white heat. After ignition it is insoluble in acids. 

Tantalum pentoxide is also a white insoluble powder, which 
has not been fused. It is also insoluble in acids. 

Vanadium is the only element of which a pentasulphide is 
known. It is pioduced by adding ammonium sulphide to a solution 
of the pentoxide, and precipitating with hydrochloric acid. It is 
a brown precipitate, which turns black on drying. It is soluble in 
sulphides of sodium and potassium, forming suJphovanadates (see 


Compounds with water and oxides. Of these oxides that 
of nitrogen is the only one which readily dissolves in water, forming 
a compound. That of vanadium is slightly soluble ; but the pent- 
oxides of niobium and tantalum do not combine with water. The 
name "acid "is applied to the hydrates of these oxides, because 
the hydrogen of the combined water is replaceable by metals, when 
the compound in solution is treated with hydroxides of the metals, 
or heated with the carbonates. These acids are as follows : 

2N 2 6 .H 2 0; N 2 8 .H 2 0, orHN0 3 ; V 2 O 6 .H 2 0, or HV0 3 ; 


(this body contains another molecule of water, which is easily ex- 
pelled by heat, and which is therefore not regarded as essential to 
its composition) ; Nb 2 O fi .nH 2 O, and Ta^Os.wHgO, the value of n being 

There are two classes of nitrates, the ordinary nitrates, and 
the basic nitrates ; and many classes of vanadates, niobates, and 

Nitric acid and nitrates. Preparation. The method of 
preparation of nitric acid is by distillation of sodium or potassium 
nitrate with excess of sulphuric acid. The reaction is as follows 

KNO 3 4- H 2 SO 4 = HKSO 4 + HM) 3 . 

It would appear as if economy of sulphuric acid might be attained 
by using the proportions 2KN0 3 + H 2 S0 4 = K 2 S0 4 + 2HNO 3 ; 
but at the temperature at which hydrogen potassium sulphate 
attacks a nitrate, nitric acid is largely decomposed. On the small 
scale, the distillation is carried out in a glass retort (see Pig. 39) ; 
on the large scale in one of iron. The iron is protected by a film 

of ferroso-ferric oxide, Fe 3 O 4 , which is at once formed on the 
surface, and on which nitric acid is without action. The worm of 
the condenser and the receivers are usually made of stoneware. 

Nitric acid is also produced along with nitrous acid by the 
action of water on nitric peroxide, N^O* or N0 3 , thus N 2 4 + H 2 


= HN0 3 4- HNO 2 ; also by heating a solution of nitrous acid, 
SHNOa = HNO 3 -f H 2 O 4- 2NO. 

When prepared by distillation it usually has a yellow colour, 
owing to its containing peroxide, NO 2 , in solution. This substance 
is easily volatile, and may be removed by passing a current of air 
through the acid for some hours. 

Properties. Nitric acid, when pure, is a colourless liquid, 
fuming slightly in the air, being somewhat volatile at the ordinary 
temperature. It freezes at 55. and boils at 86, partially de- 
cotnposrrtg into tetroxide, N 2 4 , oxygen and water, a weaker acid 
remaining behind. It is completely decomposed when heated in a 
sealed tube to 256. Its density corresponds with the formula 
HNOi. It absorbs water from the air, forming, no doubt, a 
hydrate, which, however, has not been isolated, although it is 
stated to have the formula 2HN0 3 .3H 2 O, or N 2 O 5 .5H 2 0. 

The hydrate 2N 2 O 5 .H 2 O is produced during the preparation of 
nitric anhydride, N 2 5 , by use of phosphorus pentoxide. It is the 
lower layer, into which the distillate separates, and it crystallises 
out when cooled by a freezing mixture ; and it can also be prepared 
by adding nitric anhydride to nitric acid. At the ordinary tem- 
perature it is a liquid, but it turns solid at about 5. 

A solution of vanadium pentoxide in water perhaps contains 
the compound V 2 5 3H 2 O, or H 3 VO 4 ; but the hydrate best known 
is V 2 O 5 .H 2 O, or HVO 3 , corresponding to nitric acid. This sub- 
stance is a brown-red powder, prepared by adding nitric acid to a 
solution of one of its salts, e g., V 2 O 5 .K 2 O, or KVO 3 . It is also 
formed by heating a mixture of solutions of copper sulphate with 
vanadic acid and a large excess of ammonium chloride to 75. The 
copper vanadate decomposes, depositing golden-yellow scales of 
metavanadic acid. It contains a molecule of water in addiiioti, 
V 3 O 5 .2H 2 O, but as the second molecule is lost when it is dried over 
strong sulphuric acid, it must be very loosely combined. It is also 
produced by the action of water or vanadium pentachloride, or 
oxychloride, VOC1 3 . It is a reddish-yellow powder or golden- 
yellow scales ; it is very hygroscopic. 

Niobic and tantalic acids are precipitated as white powders 
on adding hydrochloric acid to a solution of sodium or potassium 
niobate or tantalate ; or by the action of water on the penta- 
chloride of niobium or tantalum. When heated they lose water, 
and leave the pentoxide. 

Nitrates, vanadates, niobates, tantalates. These salts are 
all produced by the action of nitric, vanadic, niobic or tantalic 
acid, in presence of water, on hydrates, oxides, or carbonates, or 


by fusion of the pentoxides of the three last with hydrates or 
carbonates of lithium, sodium, potassium, &c. The following 
equations may be taken as typical : 

HNO 3 .Aq -f KOH.Aq - KNO 3 .Aq -f H 2 O; 
2HN0 3 .Aq + CuO = Cu(NO 3 ) 2 .Aq + H 2 O. 
2HNO 3 Aq -f CaCO., - Ca(NO 3 ) 2 .Aq + H 2 O + C0 2 . 
V 2 O 5 + 2KOH Aq = 2KVO 3 .Aq -f H 2 O ; 
V 2 6 + SNa-zCOa = 2Na 3 VO 4 

These equations are rendered more simple by the older method of 
representation, thus : 

N 2 O 6 .Aq -f K 2 O.Aq = N 2 O 6 .K 2 O.Aq j 
N 2 O 5 .Aq + CuO = N 2 O 5 CuO.Aq. 
N 2 O 6 Aq -f CO 2 CaO = N 2 O 5 .CaO Aq + C0 2 . 
V 2 5 + ILjO H 2 V 2 5 K^O t- H 2 : 
V 2 5 + 3C0 2 .Na 2 = V 2 O 5 .3Na 2 O + 3C0 2 . 

All nitrates are soluble, and hence cannot be produced by precipita,' 
tion, unless the solution be a concentrated one. It is possible to 
prepare certain nitrates, however, such as those of lead, silver, 
and barium, by addition of much nitric acid to a soluble salt of 
such metals; for the nitrates produced are sparingly soluble in 
nitric acid. Thus : 

BaCl 2 -f 2HNO 3 = Ba(NO,) 2 + 2HC1 ; 

Ag 2 S0 4 + 2HN0 3 = 2AgN0 3 4- H 2 S0 4 . 

The nitrate of barium or silver is precipitated as a crystalline 
powder. Many yanadates, niobates, and tantalates are produced 
by precipitation, e.g., those of lead and silver. 

Nitrates, vanadates, niobates, and tantalates. 

below 10, 2L1N0 3 3H 2 ; NaNO 3 ; KKO d ; KNO 3 .2HNO., 

melting at 3. 

KbNO 3 : BbNO 3 .5HNO 3 . CeNO 3 . NH 4 NO 3 ; NH 4 NO 3 2HNO 3 : 
NH 4 N0 3 .HN0 3 . 

These are all white, soluble salts. Those containing excess of 
nitric acid are made by mixture and cooling. With the exception 
of ammonium nitrate, the action of heat on which is peculiar, and 
will be fully treated of later in this chapter, these salts when 
heated to bright redness fuse and give off oxygen, forming at first 
the corresponding nitrites MNO 2 ; at very high temperatures they 
give off nitrogen and oxygen, and leave oxides and peroxides. They 


cannot be strongly ignited even in gold, silver, or platinum vessels 
without attacking them, forming oxides. 

Two of them, sodium and potassium nitrates, are commercially 
important. Sodium nitrate, named " Chili saltpetre," does not 
occur in Chili, hut forms immense beds, several feet thick, at 
Tarapaca, in Northern Peru. Its local name is " caliche." Its 
crystalline form is nearly cubic, hut in reality it forms very obtuse 
rhombohedra ; it is often erroneously named " cubic saltpetre." 
One gram of the salt dissolves m 1'4 grams of water at 15 ; it is 
also soluble in alcohol. It is largely used as a manure. 

Potassium nitrate, also called " nitre " or " saltpetre," is 
present in most soils, being especially abundant in chalk or marl. 
It also occurs in the leaves of many plants, especially in those of 
the tobacco-plant. It is found as an efflorescence on the soil of 
hot countries, being formed by the action of a ferment or ammonia 
in presence of bases, the ammonia being derived from decomposing 
animal or vegetable matter.* The nitrate ferment is a minute 
organism similar to those which produce fermentation. Nitrifica- 
tion, as the process of transformation of ammonia into nitric acid 
is called, goes on beneath the surface of the soil, the necessary 
conditions being apparently presence of air and absence of light. 
It ceases and does not recommence if the soil be kept for some 
time at 100, the organism being destroyed ; but after exposure to 
the atmosphere fresh germs enter, and it again proceeds. The 
manufacture of nitre by this process has been carried out for ages 
in Upper India; stable-manure and limestone are exposed to air 
for several months, and the resulting nitrate of calcium is con- 
verted into nitrate of potassium by treatment with potassium 
carbonate or sulphate ; the soluble potassium nitrate is easily 
separated from the insoluble calcium carbonate or sulphate. The 
process is also carried out in France and elsewhere. 

Potassium nitrate is now largely prepared from the Peruvian 
sodium nitrate by mixing the latter with potassium chloride. 
Sodium chloride is formed, and, as it is much less soluble in hot 
water than potassium nitrate, it separates out on evaporation. The 
mother-liquor, after removal of most of the salt, deposits crystals 
of nitre. 

Potassium nitrate crystallises in two forms : in trimetric prisms, 
and in rhombohedra, like calcspar. It has a cooling taste ; it is 
soluble in 3 parts of water at 18, but insoluble in alcohol. It 
melts at 339.t 

* Ckem. Soc , 35, 454. 

t For lists of melting points, gee Carnclley and Williams, Chem. Soc., 33, 279. 


Ammonium nitrate is prepared by mixing nitric acid and 
amnjonia, and evaporating till the water is expelled. It dissolves 
in half its weight of water at 18, and is also soluble in alcohol. It 
melts at 108, and can be distilled at 180, splitting into nitric acid 
and ammonia, \\hich recombine on cooling (?). At a higher tem- 
perature it decomposes into nitrous oxide and water. It is formed 
in solution by the action of dilute nitric acid on some metals, espe- 
cially on tin. 

Orthovanadates : Na 3 VO 4 ; K 3 VO 4 ; also with 3H 2 O and 2H 2 O. Pjro- 
Tanadate : V 2 O 5 .2K 2 O.3H 2 O. Metavanadates : LiVO 3 ; NaVO 3 ; also with 
2^0; KVO.,2H 2 0; NH 4 .VO 3 . Acid Vanadates : 2V,O 5 .Li 2 O ; also with 
9H 2 O j 2V 2 O 6 Na 2 O j aVgOfi.KgO 3, 4, and 7H 2 O ; 2V 2 O 6 (NH 4 ) 2 O 4H 2 O ; 
3V 2 O 5 .2Na 2 O 9H 2 O ; 3V 2 O 5 .K, 2 O 6H 2 O (insoluble) ; 3V 2 O 5 Na 2 O 9H 2 O ; 
3V 2 O 6 .(NH 4 ) 2 O.6H 2 O ; 4V 2 O d .6Na 2 O. 

The orthovanadates are produced by fusing vanadium pent- 
oxide with carbonates in theoretical proportions. With sodium 
carbonate the pyrovanadate, Na 4 V. 2 O 7 , is formed first. They are 
soluble in water, but decompose slowly at the ordinary tempera- 
ture, rapidly on warming, into sodium or potassium hydroxides 
and pyrovanadates or metavanadates. They are yellowish solids. 

The metavanadates are white, soluble, earthy bodies which, 
on acidifying with acetic acid, turn orange, and on evaporation 
deposit orange-yellow crystals of the acid vanadates. Ammonium 
metavanadate is produced by addition of ammonia in excess to 
vanadic acid; the acid vanadate, 2V 2 O5.(NH4) 2 O, by saturating 
ammonia with vanadic acid ; and on acidifying with acetic acid 
the body 3V 2 O 5 .(NH 4 ) 2 O is produced. 

Niobates. ?Nb 2 O 5 4^0 1GH 2 O ; 7Nb,O 5 .3K 2 O 32H 2 O ; 

2Nb 2 O 5 .3K2O.13H 2 O ; 3Nb 2 O 5 K. 2 O.5H 2 O ; 4Nb 2 O 5 .2K 2 O.llH 2 O ; 
Nb 2 O 5 2K 2 0.11H 2 O ; Nb 2 O 5 Na 2 O.6H 2 O ; 3Nb 2 O 5 .2Na 2 O.9H 2 O. 

The first of these is obtained by fusion of niobic pentoxide with 
potassium carbonate solution in water, and crystallisation. ; the 
solution also deposits crystals of the second compound ; and the 
third is formed by addition of potassium hydroxide to one of the 
former, and crystallisation. Tho fourth is produced by boiling 
potassium fluoxyniobate, NbOF 3 .2KF.H 2 O, with hydrogen potas- 
sium carbonate ; it is nearly insoluble in water. These compounds 
are all white, and crystallise. The sodium salts are easily decom- 
posed by water into hydrated niobic pentoxide and sodium 

Tantalates. 3Ta 2 O 5 .4Na 2 O.25H 2 O ; 3Ta 2 O 5 .4K 2 O.16H 2 O ; Ta 2 O 5 .Na 2 O j 

On fusing tantalum pentoxide with excess of caustic potash or 
cda, and washing out excess of the alkali with alcohol, the salts of 
the formula 3Ta 3 O 5 .4M 2 O remain as crystalline powders. Their 
solutions, when warmed, deposit the other salts of the formula 
Ta 2 O 3 .M^O as insoluble precipitates. 

Be(N0 3 ) 2 .3H 2 0; BeiOH^NOa.H^O ; Ca(NO 3 ) 2 .4H 2 O j Sr(NO 3 ) 2 .5H 2 O ; 
Ba(N0 3 ) 2 . 

These are also white soluble salts. The basic beryllium 
nitrate is produced by digesting a solution of the ordinary nitrate 
with beryllium hydrate. Calcium nitrate often occurs as an 
efflorescence on caverns frequented by bats and birds, and in stables, 
&c., where animal matter decomposes in presence of calcium 
carbonate. It is easily soluble in water, and in alcohol, and may 
be fused without decomposition. Strontium nitrate is also an 
easily soluble salt ; it is used to produce red fire in pyrotechny. 
Barium nitrate is one of the important salts of barium. It is 
formed by dissolving barium sulphide (q.v.) or carbonate in dilute 
nitric acid, or on account of its sparing solubility (1 part in 11*7 
of water at 20) by addition of potassium nitrate to a strong 
solution of barium chloride. It is insoluble in strong nitric acid 
and also in alcohol. These nitrates, when heated, yield nitrites, 
and then oxides at a bright red heat. 

Ca(V0 3 ) 2 ; Sr(V0 3 ) 2 ; Ba(VOj 2 .H 2 O ; 2Ca 2 V 2 O 7 .5H 2 O ; Ba 2 V 2 O 7 ; 
2V 2 O 6 .CaO ; 2V 2 O 5 BaO ; 2V 2 O 5 .3BaO.19H 2 O. 

The three first are yellowish-white gelatinous precipitates 
formed by adding ammonium metavanadate to soluble salts of the 
metals; the three last are orange-coloured, and are produced by 
acidifying the former with acetic acid. The other vanadates are 
insoluble and are formed on adding to a soluble salt of the metal 
potassium orthovanadate. They have not been analysed. The 
pyrovanadates are produced by precipitation. 

Nb 2 5 .2CaOj Nb 2 O 5 .CaO. 

These are prepared by fusing niobium pentoxide with calcium 
chloride, or with calcium fluoride and potassium chloride. They 
are insoluble. Other niobates and tantalates are formed as in- 
soluble precipitates on adding a soluble niobate or tantalate to a 
soluble salt of calcium, strontium, or barium. They have not been 


Mgr(NO 3 ) 2 6H 2 ; Zn(NO 3 ) 2 .6H 2 O ; Cd(NO 3 ) 2 .4H 2 O Basic nitrates : 
3N 2 5 .4Zn0.3H 2 ; N 2 O 5 .2ZnO.3H 2 O ; N 2 O 5 2CdO 3H 2 O ; and N 2 O 5 .8ZnO. 

Magnesium, zinc, and cadmium nitrates are white deliquescent 
crystals, soluble in alcohol. The basic nitrates of zinc, produce^ 
by digesting the ordinary nitrate with zinc hydrate, are non- 
crystalline soluble masses, Nb 2 O 5 .4MgO and Nb 2 O 5 3MgO are 
also known. 

Sc(N0 3 ) 3 ; Y(N0 3 ) 3 4H 2 and 12H 2 O; La(NO d ) 3 .6H 2 O and 4H 2 O. 
These are colourless soluble deliquescent salts. A crystalline 
vanadate of boron has been produced by fusion. 

A1(N0 3 ) 3 .9H 2 ; Ga(N0 3 ) 3 ; In(NO 3 ) 3 3H 2 O ; T1NO 3 ; T1NO 3 3HNO 3 
Aluminium nitrate is deliquesent ; when digested with hy- 

droxide, or when heated, it forms basic salts, similar to those of 

zinc. Indium also forms basic nitrates. 

Thallous nitrate is insoluble in alcohol ; the acid salt crystal- 

lises from strong nitric acid. All these salts are colourless. 

T1 4 V 2 7 ; T1VO 3 ; also 4V 2 O 5 TL,O ; 5V 2 O 5 .6T1 2 O ; 7V 2 O 5 .6T1 2 O. 
The first of these, prepared by fusion, is a red substance ; the 
second is precipitated by addition of or^ovanadate of sodium, 
NaaVCh to a thallous salt, as a yellowish powder; and the third 
is produced by fusion ; it forms dark scales. The pyrovanadato 
is formed with liberation of alkali ; if it be warmed with water, 
more and more alkali goes into solution, and the other acid vana- 
dates are produced. 

Cr(N0 3 ) 3 .9H 2 0; Fe(NO 3 ) 3 .9HoO. Basic salts, 2N 2 O ft Cr 2 O 3 ; 3N 2 O 5 .2Cr 2 O :i ; 
also many basic ferric salts. Two basic salts of chromium should be here 

/N0 3 /NOT 

included, viz. : CrOH, and Cr^-NO 3 , produced by treating the compound 

NCI \ci 

Cr(OH) 2 Cl with nitric acid. 

Chromium nitrate is a violet crystalline substance; and the 
ferric salt lavender- blue ; both are very soluble. The basic salts 
of chromium are green ; those of iron orange-yellow. 

Pe(N0 3 ) 2 .6H 2 ; Mn(NO 3 ) 2 6H 2 O, and 3H 2 O ; Co(NO,) 2 .6H 2 ; 
Ni(N0 3 ) 2 .6H 2 0. Basic salt : N 2 O 5 .6CoO.5H 2 O. 

The solution of ferrous nitrate must be evaporated in the cold ; 
when heated, oxygen from the nitric group oxidises the iron to 
ferric nitrate, and a basic substance is formed. The ordinary 
nitrate of iron is green ; that of manganese, pink ; that of cobalt, 
red ; and of nickel, grass-green. They are all soluble in alcohol, 


Basic cobalt nitrate is produced as a blue precipitate by adding a 
solution of ammonia to the normal nitrate ; and nickel nitrate, 
similarly treated, gives a green basic salt. 

,^ Hydrated titanium and zirconium dioxides are soluble in nitric 
acid ; but on warming the solution of the former, the hydrate 
separates out. Zirconium nitrate can be evaporated to dryness ; it 
leaves a gummy mass. Cerium sesquioxide dissolves in nitric acid, 
and on evaporation a crystalline mass of Ce(NO : ,) 3 .6H 2 is left. 
The dioxide also forms an orange-yellow solution in nitric acid. 

Thorium nitrate, Th(NO.Oi, is a crystalline salt ; it also forms 
a double salt with potassium nitrate Th(NO.<)4.KNO,. 

Silica, recently precipitated, is sparingly soluble in nitric acid. 
The nitrate of germanium has not been prepared ; that of tin, 
Sn(NO 3 ) 4 , is obtained by dissolving stannic hydrate, Sn(OH) 4 , in 
dilute nitric acid ; on rise of temperature it easily decomposes into 
metastannic acid, 5SnO 2 .5H 2 O, and nitric peroxide, N0 2 . If am- 
monium nitrate be present, the decomposition does not occur, 
probably because it forms a double salt. 

Stannous nitrate, Sn(NO) 2 , is produced by dissolving tin in 
dilute cold nitric acid. It also is easily decomposed when heated, 
giving metastannic acid. Lead nitrate, Fb(NO 3 ) 2 , forms octa- 
hedra ; when crystallised below 16, it contains 2H 2 ; it is in- 
soluble in alcohol. By digesting it with lead hydrate, or by adding 
ammonia to ordinary lead nitrate, the basic salts, N.Os 2PbO H 2 O 
(= NO, - Pb - OH) ; N 3 (V2PbO ; 2N a O 5 .3PbO.3H 2 O ; and 
3N,O 5 .10PbO.4H 2 O, are formed. The last three are nearly in- 
soluble in water. 

Two vanadates of lead are found native, viz., Pb(VO 3 ) 2 , lead 
metavanadate, or dechenite, and Pb 2 V 2 O 7 , lead pyrovaiiadate, or 
desdoizitc. Lead orthovanadate, Pb}(VO 4 ) 2 , has also been prepared ; 
it is a yellow precipitate. An orange-coloured acid salt is also 
produced on treating one of these vanadates with acetic acid ; it 
has the formula 2V 2 O 5 .PbO. The mineral vanadinite, 
3Pb}(VO 4 ) 3 .PbClo, is a compound of lead orthovanadate and 

Nitrates, vanadates, &c., of members of the vanadium group do 
not appear to exist. The compound nitric peroxide, N 2 4 , has been 
viewed as nitrate of nitrosyl, NO, thus, NO(NO 3 ) ; but of the 
justice of this view there is no proof. A nitrate of the oxide V 3 O 4 
appears to exist ; and Y 2 O 6 is soluble in acids, but the hydrates of 
tantalum and niobium pentoxides are insoluble in nitiic acid. 

Similarly, although the oxides of phosphorus and arsenic dis- 
solve in nitric acid, no compound has been isolated. But with 


antimony, N 2 O 5 ,Sb 4 O 6 , has been prepared ; and the pentoxide, 
Sb 2 O 5 , is slightly soluble in nitric acid, 

Bismuth nitrate, Bi(NO 3 ) 3 .5H a O, is a well crystallised salt. 
On treatment with water it gives a mixture of three salt*/, 
each of which may, however, be prepared fairly pure by careful 
attention to temperature and dilution. These are N 2 O .Bi 2 O 3 .H>O 
= 2{BiO(NO 3 )}.H 2 O ; 2N 3 O 5 .Bi 2 O 3 H 2 O = Bi(OH)(NO 3 J a ; and 
N 2 O 6 .2Bi 2 O 3 .H 2 O These basic nitrates are insoluble in water. 

Molybdenum trioxide is soluble in nitric acid ; so too is oxide of 
tungsten, but no compounds are known. Uranium fornts yellow 
nitrates of uranyl of the formulas UO 2 (NO 3 ) 2 3H 2 O and BH 2 O. 

Samarskite consists chiefly of niobates of uranyl, iron, and 

A nitrate of tellurium of the formula NjOa.^TeO^ is produced 
on dissolving tellurium in nitric acid, and evaporating. 

Rhodium oxide is soluble in nitric acid, but the nitrate is 
unstable. But on adding sodium nitrate the stable double salt 
Rh(NOj) 3 .NaNO.5 may be obtained in crystals. Palladium nitrate, 
Pd(NOj) 8 is easily prepared by dissolving palladium monoxide, or 
the metal, in nitric acid. It is a brown compound ; and on evapo- 
ration a basic salt is produced. 

Osmium oxide is also soluble in nitric acid. Platinic nitrate, 
Pt(NO 3 ) 4 , is unstable, but as with rb odium the addition of potas- 
sium nitrate yields a stable double salt of the formula 
Pt(NO d ) 4 .KN0 3 . 

Cu(N0 3 ) 2 .3H 2 5 Cu 3 (V0 4 ) 2 H 2 0; V 2 O 5 .4CuO.3H 2 O, also H 2 O. The 
latter is possibly VO.(OCu.OH).(O 2 )Cu. It is found native, and named 

Cu(N0 3 ) 2 NH 4 NO a . 

Copper nitrate is a soluble blue salt, crystallising well. It 
is the source of copper oxide for the analysis of organic sub- 
stances, for, like almost all the nitrates, it yields the oxide ori 
ignition. The vanadates are brown substances. 

AgrNO 3 .NH 4 NO ,; 2AffNO 3 .Pb(NO 3 ) a . 
4 ; Agr 4 V 2 7 ; Ag>VO 3 . 
The first nitrate is an important substance. Great use is made 
of it in photography, electro typing, &c., and under its old name 
" lunar caustic" (luna = silver), it is employed as a caustic, being 
cast into sticks for medical use. It is a white easily fusible salt 
(m. p. 218) ; it is soluble in about its own weight of cold water, 
and iii about four times its weight of alcohol. It crystallises with 
eodiuai and lithium, to form double salts like those of potassium 


and ammonium, but not in molecular proportions. A number of 
double nitrates and halides are known ; e.g., 

AffN0 3 AgrCl; AgrN0 3 .AgBr ; A*NO 3 AffI ; 2AgrNO 3 .AgrI ; 

3 .Pb (N0 3 ) 2 .2Ag-I ; 2AffNO 3 .Pb(NO 3 ) 2 .2AgrI. 

These are sparingly soluble salts prepared by mixture. 

The mercurous nitrates are numerous, many basic compounds 
being known. They are as follows : 

t HffNO 3 .H 2 O; 3N 2 O 5 .4Hgr 2 O.H 2 O ; 3N 2 O 5 .5Hgr 2 O 2H 2 O ; N 2 O 6 .2Hg 2 O 2H 2 O. 

Others are said to have been obtained, but their existence is 

questionable. Mercurous nitrate is formed by digesting mercury 
with cold dilute nitric acid. The basic nitrates are produced by the 
action of water on the ordinary salt. Double salts with strontium, 
barium, and lead nitrates are also known, of formula? such as 
3N 2 O ft .2PbO.2Hg 2 O. All these salts are crystalline and soluble. 

By dissolving mercuric oxide, HgO, in excess of nitric acid 
and evaporating, crystals of the salt 2Hg(NO 3 ) 2 .H,O, are deposited. 
Crystals with 8H 2 O may also be produced by cooling the solution. 
These crystals, when fused, deposit a basic salt, N 2 O 5 .2HgO.3H 2 O ; 
and with water they yield N 2 O 5 .3HgO.H 2 O. Like silver nitrate, 
mercuric nitrate combines with mercury halides, forming colour- 
less crystalline compounds, e.g., Hg(NO 3 ) 2 .HgI 2 ; 2Hg(NO 3 ) 2 .HgI 2 ; 
Hg(NO 3 ) 3 .2HgI 2 ; and 2Hg(NO 3 ) 2 .3HgI 2 . These are all decom- 
posed by water. The compound 2Hg(NO 3 ) 2 .4AgI.H,O is also 

Oxide of gold dissolves in nitric acid, but the solution decom- 
poses spontaneously at the ordinary temperature, again depositing 
gold oxide. 

Compounds of vanadium pentasulphide. This body is 
soluble in sulphides of the alkalies. On adding alcohol to its 
solution in potassium sulphide, a scarlet precipitate is produced, 
consisting of potassium sulphovanadate ; it lias probably tho 
formula V 2 S 5 .K 2 S = KVS 3 , and is a meta-compound. A solution 
of this substance gives brown precipitates with soluble salts of 
other elements, but the formulae of the compounds are unknown. 

Compounds containing halogens. 
VOF 3 ; YOC1 3 ; VOBr 3 ; NbOF 3 ; NbOCl 3 ; NbOBr 3 ; TaOF 3 . 

No corresponding nitrogen compound is known, although a 
mixture of nitrosyl chloride, NOC1, and chlorine reacts with water 
as. if it consisted of 


Vanadyl trifluoride is known in combination (see below). 

Vanadyl chloride, VOC1 3 , is produced by heating the oxide, 
V 2 2 (or VO P) in a current of chlorine, when direct union ensues. 
The higher oxides, mixed with carbon and heated in. chlorine, a.\>o 
yield it. The bromide, VOBr 3 , is similarly prepared; also by pass- 
ing bromine over the heated trioxide, V^O 3 , but no corresponding 
iodide seems capable of existence. 

Niobium oxyfluoride (niobyl fluoride), oxychloride, and 
OXybromide are volatile white crystalline bodies. The chloride 
has a vapour density corresponding to the formula NbOQl 3 . They 
are prepared along with the pentahalides by the action of chlorine 
on a mixture of the pentoxide with charcoal at a bright red heat. 
The trichloride at a red heat decomposes carbon dioxide, pro- 
ducing carbon monoxide and the oxychloride. Tantalum oxy- 
fluoride is produced by the action of air or water- vapour on the 
pentafluoride ; the oxychloiide and oxy bromide could probably be 
similarly produced. 

Vanadyl chloride is a golden-yellow liquid, boiling at 127. Its 
density leads to the usual formula. The oxybromide forms a dark 
red liquid boiling at about 130 under a pressure of 100 mrns., bufc 
it is decomposed at/ 180 into the dibromide VOBr,. 

These oxyhalides form the following compounds with the 
htiKdes of other elements : 

Fluoxyvanadates : 

V 2 O 5 .2VOF.,.6KF.2H 2 O ; V 2 2 2VOF 3 6NH 4 F 2H 2 O ; 
V 2 O 5 .2VOF 3 .12NH 4 F. 

Vanadoxy fluorides : 

2VOF 3 .3KHF 2 ; 2VOF 3 .3NH 4 HF 2 ; 2VOF 3 .ZnF 2 Zn0.14H 2 O. 

Nioboxyfluorides : 

NbOF 3 2KF H 2 O ; NbOF 3 2NH 4 F H 2 O ; NbOF 3 .3KF - NbOF 3 .3NH 4 F ; 
3NbOF 3 5KF.H 2 O ; 3NbOF d .5NH 4 F.H 2 O ; NbOF 3 3KF.HF j 
3NbOF 3 4KF.2H 2 ; NbOF 3 NH 4 F ; NbOF,, NbF 5 .3NH 4 F j 
NbOF 3 ZnF 2 .6H 2 O. 

Tantaloxyfluoride : TaOF 3 .3NH 4 F. 

These bodies are produced by direct union. They are crystal- 
line salts. The tantatoxyfluorides react with water, forming 
hydrated tantalum pentoxide and tan talifluorides, such as TaF 6 2KF, 
hence they have been little investigated. The corresponding 
chlorides and bromides of these elements are also easily decom- 
posed by water, hence their derivatives have not been prepared. 


Tetroxides, or dioxides. These are as follows : 

N0 2 and N 2 4 j VO 2 or V 2 O 4 ; NbO 2 or Nb 2 O 4 ; TaO 2 or Ta^ ; 

N VS 2 or V 2 S 4 . 

The formula of nitric peroxide, as this substance is usually 
called, depends on the temperature. In the liquid state it is a 
tetroxide, N 2 O 4 . The gas, at the lowest possible temperature, 
also approximates to this formula: but on raising the temperature, 
dissociation ensues, the extent of dissociation depending on the 
temperature and pressure, until, at 140, at atmospheric pressure, 
the more complex molecules of N 2 4 are entirely resolved into mole- 
cules of N0 2 . At higher temperatures the compound NO disso- 
ciates in its turn into NO and 0, and at 620 the gas contains no 
molecules of peroxide. On cooling, recombination takes place, 
and the phenomena are reversed. It is possible to trace these 
changes by the alteration of colour of the gas ; N 2 4 is an almost 
colourless substance when solid; NO 2 is dark reddish-black ; and 
a mixture of NO and O is also colourless. On heating a tube of 
hard glass filled with the gas, it turns dark at first, and then 
lightens in colour, turning nearly colourless at the temperature at 
which the glass begins to soften. As we have the two substances, 
one of which is a polijmeride of the other, it is convenient to give 
them different names. The first, N0 2 , we shall call nitric peroxide, 
reserving the name tetroxide for the compound N 2 4 . 

Alternative formulae have been ascribed to the oxides of vana- 
dium, niobium, and tantalum. They are non-volatile solids, and 
nothing is kno^n regarding their molecular complexity. 

Preparation.!. By the union of the lower oxides with 
oxygen. Nitrous oxide, N 2 O, does not combine directly with 
oxygen; but nitric oxide, NO, mixed with half its volume of 
oxygen, at once combines, forming a mixture of peroxide and 
tetroxide. Nitrogen trioxide, N 2 3 , which is a blue liquid, is also 
slowly converted into peroxide and tetroxide when kept in presence 
of oxygen or air. 

Vanadium tetroxide is formed when the trioxide is heated 
in air ; but on prolonged heating it is oxidised to the pent- 

2. By depriving a higher oxide of oxygen. It has been 
already remarked that nitrogen pentoxide decomposes spon- 
taneously into peroxide and oxygen. Nitric acid is more stable ; 
but when its vapour is led through a red-hot tube, a large propor- 
tion is decomposed. It is more convenient, however, to deprive 
nitric acid of oxygen by distilling it with arsenious anhydride. 


The reaction is :4N 2 6 .H 2 O 4- As 4 O fl = 4>N 2 4 + 2As 2 5 + 4H 2 0. 
The water, however, reacts with the tetroxide, thus : ;JN 2 4 -f 
2H 2 = 4HN0 3 + 2 NO ; and a mixture of tetroxide, peroxide, 
and nitric oxide is produced. On condensing the product, thcy*d 
combine to form trioxide, thus, NO* + NO = N 2 3 . Hence, in 
order to remove water from the sphere of action, a considerable 
quantity of strong sulphuric acid or phosphorus pentoxide is added. 
The product is then pure peroxide and tetroxide. To remove 
nitric acid, some of which is apt to distil over, the liquid is again 
distilled, with addition of a little more arsenic trioxide Mid phos- 
phorus pentoxide. 

The tetroxide may also bo formed by the action of nitric pent- 
oxide on the trioxide. The blue liquid containing trioxide may be 
rendered orange by addition of a mixture of nitric acid and phos- 
phoric anhydride, which must contain N 2 8 . 

When a nitrate is heated, it decomposes into an oxide and oxides 
of nitrogen. If the pentoxide were not so unstable, one would 
expect that it would be formed, but, as a rule, the peroxide resolved 
by heat into nitric oxide and oxygen is produced by its decomposi- 
tion. On cooling the resulting gases they re-combine to form 
tetroxide and peroxide. The most convenient nitrate to employ is 
that of lead. The equation is : 

Pb(NO 3 ) 2 = PbO + 2NO, + 0. 

Metallic tin may also be used to withdraw oxygen from nitric 
acid. The equation is : 

Sn + 4HN0 3 = SnO, + 2# 2 4 + 2H 2 0. 

Nitric and nitrous oxides, NO and N 2 0, are, however, produce*} 
simultaneously. The nitric acid must be strong and somewhat 
warm. It will be remembered that the tin is oxidised to meta- 
stannic acid, 5SnO 2 .5H 2 O. 

The compound, V0 2 C1, decomposes when heated in carbon 
dioxide into VO 2 and chlorine. 

Niobium pentoxide is reduced to tetroxide by heating it to 
whiteness in hydrogen, and tantalum pentoxide when heated to 
whiteness in a crucible lined with carbon loses oxygen, leaving the 

Properties. Nitrogen tetroxide is a colourless solid below 
10*14. At that temperature it melts, but the liquid has a pale 
straw colour, owing to incipient dissociation. As the temperature 
rises its colour changes to yellow and then orange-red j it boils at 


22, giving off a brown-red gas, which consists largely of the per- 
oxide. The peroxide is not known in the solid form, but the 
liquid tetroxide apparently contains some, judging from its colour. 
T^e liquid compound is heavier than water (1*45 at 15). It 
reacts with ice-cold water, forming nitrous and nitric acids, 
N 2 O 4 + H 2 = HNO 3 + HNO 2 ; and at higher temperatures 
forming nitric acid and nitric oxide, 3N 2 O4 + 2H 2 O = 4HNO 3 -h 
2NO. It dissolves in strong nitric acid, forming the red fuming 
acid often employed for oxidation of sulphides, &c. ; and in sul- 
phuric auid, giving salts of nitrosyl, NO (see sulphates). It acts 
violently on cork and indiarubber, hence, in preparing it, all the 
joints should be of sealed glass. 

Vanadium tetroxide, V^O 4 or VO 2 , is a dark green amorphous 
powder, insoluble in water, but soluble in hydroxides of sodium 
and potassium, forming hypovanadates, and in acids, forming salts 
of vanadyl (VO). 

Niobium tetroxide is a dense black insoluble powder, which 
On ignition in air yields the pentoxide ; and tantalum tetroxide 
is a dark substance, which acquires metallic lustre under the bur- 

Tetrasulphides of vanadium, V 2 S 4 , and tantalum, Ta 2 S 4 (?) 
are known. The first, produced by heating the tetroxide in a 
stream of hydrogen sulphide, is a black powder, insoluble in water, 
alkalies, or alkaline sulphides ; the second, which may be an oxy- 
sulphide, is produced by heating tantalum pentoxide in vapour of 
carbon disulphide or tantalum pentachloride in hydrogen sulphide. 
It is a black powder, which when burnished acquires a brass-yellow 

Compounds with oxides and sulphides. Nitric peroxide 
does not combine with water, but is decomposed (see above). It 
combines, however, with lead oxide, producing a compound of the 
formula PbN 2 O 6 , which may be a salt of the hypothetical acid, 

H 3 N 2 6 , or may be a double nitrite and nitrate of lead, Pb^MQ*' 

Similar compounds, but containing more lead oxide, are produced 
by heating lead nitrate with metallic lead. 

Vanadium tetroxide dissolves in alkalies, forming hypo- 
vanadates. On addition of a hydroxide to its solution in hydro- 
chloric or sulphuric acids, its hydrate, V 2 O 4 .7H 2 O, is thrown down 
as an amorphous black precipitate, which may be viewed as 
h yd rated hydrogen hypovanadate. An arbitrary division is usually 
drawn between the compounds called hypovanadates and those 
termed vanadyl salts. They are here considered as chemically 


similar ; both contain vanadium tetroxide in combination with 
other oxides. They are as follows : 

2V 2 4 .K 2 0.7H 2 ; 2V204.Na20.7H20; 2V 2 O 4 .(NH 4 ) 2 O.3H 2 Oj 2V 2 O 4 .BaO^ 
2V 2 O 4 PbO ; 2V 2 O 4 . Aff 2 O. 

These are termed hypovanadates. There are also V2O4.3SOj.4HoO 
and 15H 2 O; V 2 O^SO 3 .7H 2 O and 10H 2 O. These are termed 
vanadyl sulphates, and will be considered among the sulphates. 

Potassium hypovanadate, 2V a O 4 .K2O.7H 2 O, forms dark brown 
crystals, soluble in water, but nearly insoluble in caustic potash, 
and quite insoluble in alcohol. The sodium salt is similar. The 
barium, lead, and silver salts are brown or black, and are produced 
by precipitation. 

Hydrated vanadium tetrasulphide is precipitated on addi- 
tion of an acid to a solution of the tetroxide in sulphides of the 
alkalies. It is a brown powder. It dissolves in sulphides of the 
alkalies, forming the hyposulphovanadates. These have been 
little studied ; they are black solids dissolving with a brown 
colour. Those of the alkalies are soluble, and give precipitates 
with solutions of the metals. 

Compounds with halides. Compounds of the formulae 
NOCl a and N0 2 C1, though it has been stated that they are formed 
by various reactions, have been proved to consist of solutions of 
chlorine in nitrosyl chloride, NOC1, or in nitrogen tetroxide. Nor 
is any compound known of the formula NOC1 2 . 

But vanadium oxytrichloride, VOC1 3 , when heated to 400 with 
metallic zinc is converted into VOC1., a liglit green crystalline 
solid, deliquescent, and soluble in alkalies. The corresponding 
bromide is a yellow-brown deliquescent solid, produced by heating 
the tribromide to 180. The corresponding fluoride, VOF a , is 
known in combination with ammonium fluoride, in the blue mono- 
clinic crystals of VOF 2 .2NH 4 F, produced by adding hydrogen 
ammonium fluoride, HNH 4 F 2 , to a solution of tetroxide, V 2 4 , in 
hydrofluoric acid. 

Trioxides, N 2 O 3 ; V 2 O 3 . Preparation. Nitrogen trioxide, 

or nitrous anhydride, is produced by the union of nitric peroxide, 
NO 2 , with nitric oxide, NO. It is apparently formed by all 
reactions involving these products ; but as it cannot exist in the 
gaseous state, it is formed only on cooling the mixture of its 
products of decomposition.* Such a gaseous mixture is liberated 
on treating a nitrite with sulphuric acid, thus : 

* Chem. Sac., 47, 187. 


2KNO 2 + H 2 SO 4 = K2SO4 + H 2 + NO* + NO; 
or by adding water to hydrogen nitrosyl sulphate : 
% 2H(NO)S0 4 + Aq = 2H 2 S0 4 .Aq + N0 l + NO. 

When fairly pure it is a mobile blue liquid, stable only at a very 
low temperature. It does not solidify even on cooling it to about 
90. If warmed, it decomposes into its constituents ; and as more 
nitric oxide escapes than peroxide, the colour of the remaining 
portion Changes to green, and subsequently to dirty red : for the 
colour of the remaining peroxide is changed by that of the blue 
trioxide. It is also formed by the action of a small quantity of 
water on nitrogen tetroxide, thus : 2N 2 4 -f H 2 O = 2HNO 3 + 
N 2 3 ; and this is one of the easiest methods of preparing it.. 
A mixture of nitric oxide, NO, and oxygen, even if the oxygen 
be in excess, combines to some extent to form the trioxide when 
cooled in a freezing mixture. 

Vanadium trioxide is produced by heating vanadium pent- 
oxide in a current of hydrogen or with carbon. It is also formed 
when V 2 O 2 is heated gently in air. It is a black insoluble powder, 
possessing semi-metallic lustre. It is insoluble in acids. When 
heated to redness in air, it glows and burns to the pentoxide. It 
is a conductor of electricity. Like nitrogen trioxide, it combines 
with oxides, forming the vanadites. 

No trioxide of niobium or tantalum has been prepared. 

Compounds with oxides. Nitrites and vanadites. It is 
probable that two sets of nitrites exist, having the same formulae 
but different constitution ; these may be regarded as derivatives of 

two hypothetical nitrous acids, HN<^Q, and HO N=O. 

It is probable that the silver and mercury salts are derivatives 
of the first, and the potassium and calcium salts of the second. 
The reason for this view is as follows : 

The compound of carbon, hydrogen, and iodine, known as 
methyl iodide, has the formula CH 3 I. When heated with silver 
nitrite in a sealed tube, silver iodide is produced, along with the 
compound CH 3 .N0 2 , named nitromethane. Now, on exposing this 
liquid to the action of nascent hydrogen, produced, for example, by 
the action of tin on hydrochloric acid, the following reaction 
occurs : 

CH 3 .N0 2 + 6H = CH 3 .NH 2 + 2H a O ; 

the oxygen is replaced by hydrogen, forming the compound 



(CH J ).!N'H 7 , analogous to ammonia, NH 3 ; and it is argued that the 
nitrogen and the carbon must be combined with each other. 

On heating methyl iodide, CH 3 I, with potassium nitrite, on the 
other hand, a compound of the same formula is produced, vjf ., 
CH 3 .N0 2 , along with potassium iodide. But this body, which is 
named methyl nitrite, differs entirely in properties from its 
isomeride, nitromethane. And on treatment with nascent hydro- 
gen, this reaction takes place : 

CH 3 .N0 8 + 6H = CH 8 .OH + NH 3 + H 2 0. , 

The body CH 3 .OH is named methyl alcohol, and it is certain 
that carbon and oxygen are here combined. Hence the formula 
CH 3 O NO is attributed to it, and KONO to the nitrite from 
which it is derived ; whereas silver nitrite has apparently the 
formula Ag N0 2 . 

These conclusions are confirmed by a study of the action of 
caustic potash on these bodies. For while nitromebhane reacts 
thus : 

CH 3 .N0 2 + KOH = CH a .KN0 2 + H 2 O, 

methyl nitrite is decomposed, thus : 

CH 3 .0]SrO + KOH = CH 3 .OH + KONO, 

the original potassium nitrite being reproduced. 

While, therefore, silver nitrite should probably be regarded as 
a nitride of silver and oxygen, and should be considered among 
the nitrides, and potassium nitrite as a derivative of nitrous 
anhydride, yet we do not know which bodies to place in one class 
and which in the other ; and as we are not sure whether some of 
the compounds named nitrites are not mixtures of both com- 
pounds, it is more convenient to include both varieties at present 
in one class.* 

Preparation. The nitrites are prepared : 1. By reducing the 
nitrates. This is best done by fusing them with metallic lead. 
For instance, three parts of potassium nitrate fused with two parts 
of metallic lead with constant stirring yield potassium nitrite and 
lead monoxide, thus : 

KNO 3 + Pb = K]srO 2 + PbO. 

Potassium sulphite may also be employed as a reducing agent. 
2. By the action of a mixture of NO 2 and NO on 

Chem> Soc. t 47, 203, 205, 631. 


hydroxides. Those reactions which produce such mixtures in 
correct proportions are to be preferred. An example is 

NO + NO, + 2KOH.Aq = 2KNO 2 .Aq + H 2 O. 

3. By passing a mixture of oxygen and ammonia over 
heated platinum black (finely divided platinum), ammonium 
nitrite is formed, thus : 

2NH A + 30 = NH A NO 2 + H 2 O. 

The nitrites of lead and silver are nearly insoluble, whereas 
the nitrates are very soluble salts ; hence, on adding to a nitrite a 
soluble salt of one of these metals (nitrates), the respective 
nitrites are precipitated. They may be converted into other 
nitrites by digestion with a soluble chloride in the case of silver, or 
a sulphate in the case of lead. 

List Of Nitrites. The following have been prepared : 
NaNO 2 j KNO 2 ; NH 4 NO 2 .H 2 O. 

White deliquescent salts. That of sodium is soluble in alcohol. 
The ammonium salt is produced by addition of nitrous anhydride, 
N 2 O 3 , to ammonia, keeping it cold ; or by mixing solutions of lead 
nitrite and ammonium sulphate, filtering off insoluble lead sul- 
phate, and evaporating in a vacuum to crystallisation. When 
heated, even in solution, it undergoes the curious decomposition 
NH 4 N0 2 = N z + 2H 2 0. 

This forms a convenient method of preparing pure nitrogen. 
It may be carried out more conveniently by heating a mixture of 
potassium nitrite and ammonium chloride, best after addition of 
copper sulphate. 

The corresponding vanadites have not been analysed. They 
are produced by dissolving vanadium trioxide in alkalies. They 
are red when 'hydrated, but green when anhydrous. 

Ca(N0 2 ) 2 H 2 Oj Sr(N0 2 ) 2 .H 2 Oj Ba(NO 2 ) 2 .H 3 O ; Ba(NO 2 ) 2 .KNO 2 .H 2 O. 

These salts may be formed by heating a nitra.te of one of these 
metals, dissolving the product in water, and, in order to separate 
oxide, passing carbon dioxide to remove it as carbonate. The fil- 
trate is evaporated and crystallised. Calcium nitrite is insoluble 
in alcohol. These are all soluble white salts. 

Mgr(N0 2 ) 2 .3H 2 and 2H 2 O ; Zn(NO 2 ) 2 .3H 2 O ; Od(irO ft ) f .H i O. Also basic' 
salts :N 2 O 3 .2ZnO, and N 2 O 3 .2CdO; and double salts, Cd(NO 3 ) 2 .2XNO 2l 
and Cd(N0 2 ) 2 .4KN0 2 . 

These are all white soluble salts. 

Nitrites of chromium and iron have not been investigated. 


Manganous nitrite is a pink deliquescent salt ; that of cobalt is 
rose-coloured, and of nickel green. 

The double nitrates of the last two metals are better known. 
They are as follows : 

3(Co(NO 2 ) 2 .2KN0 2 ).H 2 O, also with other amounts of water. 
Kl(NO 3 ) 3 .Oa(N > O2) 2 .KN > O2 ; also similar strontium and barium salts. 

These contain the metals as dyads, and are derivatives of CoO, 
and NiO. 

2Co(NO 2 ) 3 .4NaNO 2 .H 2 O ; also 6NaNO 2 .H >O ; 2Co(NO 2 ) 3 .4KoO. 

These compounds are produced by boiling a cobalt salt with 
acetic acid and nitrite of sodium or potassium. The cobalt is 
here triad, as in Co 2 O 3 . Nickel forms no corresponding com- 
pounds, and as the double nitrite of cobalt and potassium is nearly 
insoluble in water, its formation is used as a means of separating 
cobalt from nickel. Jt has a bright yellow colour, and is therefore 
used as a pigment. 

The following compounds of lead are known: 
Pb(NO 2 ) 2 ; N 2 3 .2PbO.HjO, and 3H 2 O ; N 2 O 3 .3PbO.H 2 O ; N 2 O 3 .4PbO.H2O. 

The last three are yellow bodies, and are made by boiling 
a solution of lead nitrate with metallic lead; the first, by 
passing a current of carbon dioxide through one of the latter 
suspended in water ; the excess of lead oxide is removed as car- 
bonate. When lead nitrate solution is boiled with lead, a double 

nitrate and nitrite is also formed. Its formula is 4Pb< j 

a basic salt is also produced, viz., N 2 O 3 .N 2 O 6 .9PbO.3H 2 O. The first 
of these has been viewed as a salt of the anhydride N 2 4 ; as N 2 O 4 .PbO 
(see p. 335) ; but the formula given is more probably correct. 

Copper nitrite, Cu(NO 2 ) 2 is an apple- green crystalline salt; 
and silver nitrite, AgNO^, forms long needle-shaped pale-yellow 
crystals, sparingly soluble in cold water. 

Some interesting double nitrites of platinum have been pre- 
pared (see pp. 485 and 344*). 

Compounds with halides. - NOC1 ;* VOC1. The first of 
these bodies has the molecular weight given by the formula. It 
is prepared (1) by passing a mixture of nitric peroxide and chlorine 
through a red-hot tube. The nitric peroxide is doubtless dis- 
sociated into nitric oxide and oxygen, and the former combines 
with the chlorine. It is also produced by direct combination of 
nitric oxide with chlorine at a red heat. (2) By the action of salt 
(NaCl) on hydrogen nitrosyl sulphate, H(ETO)S0 4 , produced by 
* Chem. Soe., 27, 630 ; 49, 222, 


saturating strong sulphuric acid with NO* and NO, thus: 
H(NO)S0 4 + Nad = HNaSO 4 + NOCL (3) Along with free 
chlorine, by heating a mixture of hydrochloric and nitric acids, 
thtfc : 3HC1 + HN0 3 = 2H 2 O -f NOGl + Cl* ; and probably by 
the action of hydrogen chloride on nitrogen tetroxide, which may 
be regarded as nitrate of nitrosyl, NO(NO 3 ), thug : 

HC1 -f NO(N0 3 ) = NOC1 -f- HN0 3 . 

A mixture of nitric and hydrochloric acids has heen long known 
under thefnanie "aqua regia." Owing to the nascent chlorine, it 
has the property of dissolving gold and platinum, converting them 
into chlorides. It is a powerful oxidising agent, the chlorine re- 
acting with water forming nascent oxygen and hydrogen chloride. 

The corresponding vanadosyl chloride, VO01, is a brown 
powder formed by heating the trichloride, VOC1 3 , to redness in a 
current of hydrogen. At the same time the compound V 2 O 2 C1 is 
formed as a heavy shining powder like mosaic gold, and also the 
oxide V 2 O 2 or VO. With vanadium we have thus the series, 
VOC1 3 , VOC1 2 , VOC1, and V.O.OL 

Nitric oxide and vanadium dioxide, NO and V 2 O 2 . The 
first of. those is often erroneously named nitrogen dioxide. Its 
formula, however, even at 100, is NO, as shown by its vapour- 
density. No tendency towards increased density has been noticed ; 
the gas contracts paripassu with hydrogen. The molecular weight 
of the vanadium compound is unknown, but as it is derived from 
V 2 O,C1, it is possibly V 2 O 2 . 

Nitric oxide is produced in an impure state by the action of 
nitric acid on certain metals. It is probable that the normal action 
of nitric acid is similar to that of other acids ; that a nitrate is 
produced with liberation of hydrogen. But nascent hydrogen 
(i.e., hydrogen in the state of being liberated, when it consists in 
all probability of single un combined atoms) cannot exist in 
presence of nitric acid, but deprives it of oxygen. In theory, the 
following reactions are possible : 

3 + M" = M(N0 3 ) a -f- 2H 

5H a O. 

The conditions determining the prevalence of any one of these 
reactions are temperature, presence of water, and of the products of 
reaction. But the oxides of nitrogen produced may themselves 
react with water or with nitric acid. For example, if N a 4 be 


liberated in presence of water, the reaction described on p. 337 
will take place, and a mixture of nitric oxide and nitric acid will 
be produced. But some peroxide may escape along with nitric 
oxide. The gases NO, N Z 0, and nitrogen, not being affectedT by 
water, will be liberated as such, if formed. 

Nitric acid diluted with its own volume of water, acts on 
copper at 15 and on aluminium at 60 65, producing a mixture 
containing 98 and 97 per cent, respectively of nitric oxide, along 
with a small amount of -nitrous oxide and nitrogen. With silver, 
acid of the same strength at 15 gives 31 per cent of nitric oxide 
and 60 per cent, of nitrous oxide, N 2 0, while iron with nitric acid 
of any dilution, gives chiefly nitric oxide (from 86 to 91 per cent.).* 

The action of nitric acid on copper therefore forms the most 
convenient method of preparing nitric oxide. The equation is : 
3Cu + 8HN0 3 .Aq = 3Cu(N0 3 ) 2 .Aq + 4H 2 O + 2NO. To prepare 
the pure compound, this gas is passed through a strong cold solu- 
tion of ferrous sulphate, FeS0 4 , with which nitric oxide combinesf 
(see p. 428) . On warming the solution, the compound is decom- 
posed, and pure nitric oxide is liberated. It is a colourless, nearly 
insoluble gas, which, when mixed with air or oxygen, gives red 
fumes of nitric peroxide. It condenses to a colourless liquid at 
11 under a pressure of 104 atmospheres. Under normal 
pressure, it boils at 153*6, and begins to solidify when the 
pressure is reduced to 138 mms. at 167. It does not support 
combustion, but like other gases containing oxygen, it is 
decomposed at a high temperature, and thus glowing charcoal 
or phosphorus burn in it. With the vapour of carbon disulphide 
it forms a mixture which, when set on fire, burns rapidly with a 
brilliant blue-white flame. When mixed with hydrogen, it can 
be exploded by a powerful spark. 

The corresponding oxide of vanadium, V 2 O 2 , may be formed 
by the action of potassium on a higher oxide of vanadium, and 
used to be considered to be metallic vanadium. It is also pro- 
duced when a mixture of vanadyl trichloride, VOC1 3 , and hydrogen 
are passed through a tube full of red-hot charcoal. It is a light- 
grey powder with metallic lustre, difficult of fusion, and insoluble 
in water and acids. When heated in air, it burns to higher oxides. 

It may be produced in solution by reducing a solution of 
vanadium pentoxide in sulphuric acid by means of zinc. Such a 
solution has a lavender colour, and is one of the most powerful 
reducing agfents known. 

* Chem. 800., 28, 828; 32, 52. 
t Qompt. rend. } ; 89, 410. 


Nitrogen sulphide and selenide, NS and NSe. The first 
is produced by the action of ammonia on sulphur chloride dissolved 
in carbon disulphide, thus :8NH 3 + 3S 2 C1 2 = 6NH 4 C1 + 2NS + 
4 A; the ammonium chloride, being insoluble in carbon disulphide, 
is removed by filtration, and the ca,rbon disulphide on evaporation 
deposits nitrogen sulphide in yellow rhombic prisms. The corre- 
sponding selenium compound, produced, however, from selenium 
tetrachloride, is an amorphous, orange-coloured, insoluble substance. 
Both of these bodies explode by percussion. 

When mixed with chloroform and treated with chlorine, sulphur- 
yellow crystals of the formula NSC1 are deposited, analogous to 
nitrosyl chloride, NOC1. A second chloride (NS) 3 C1 is also formed ; 
it deposits in copper- coloured needles. 

Nitroso-sulphides. A curious set of compounds of nitric 
oxide with sulphide of iron and of a metal has been produced* 
by dropping a solution of ferric chloride into a mixture of solutions 
of potassium nitrite and ammonium sulphide, when black crystals of 
Fe 3 S4(NO)4.H 2 S are deposited. When the solution of these crystals 
is heated with caustic soda, they yield large black crystals of the 
compound Fe 2 S 3 (NO) 2 .3Na 2 S; and with an acid, a black precipitate 
of " nitrososulphide of iron/' Fe 2 S 3 (NO) 2 separates. The first 
compound, heated to 100 with sodium sulphide, deposits red prisms 
of the body Pe 2 S 3 (NO) 2 .Na 2 S.H 2 O. 

The constitution of these bodies is unknown ; but they appear 
to be related to the nitroferricyanides (see p. 566). It is sug- 
gested that a corresponding amido- compound has the formula 
Pe(NO 2 ).SNH 2 , and the last nitroso- sulphide may be analogously 
represented Pe(SNa).SNO. 

Nitrous oxide, N^O, is produced (]) by the action of metals 
on nitric acid. Zinc and pure nitric acid at 15 yield a mixture 
consisting of 1 per cent, of nitric oxide, 78 per cent, of nitrous 
oxide, and 21 per cent, of nitrogen. Nickel and cobalt, too, with 
acid diluted with its own volume of water, yield a mixture contain- 
ing about 80 per cent. ; and tin, at ordinary temperatures, furnishes 
a mixture containing from 67 to 85 per cent., with acids of all con- 
centrations. (2) The simplest method of preparation is to heat 
ammonium nitrate to above 185, when it decomposes like the nitrite, 
thus : NH,NO ? = N 2 + 2# 2 0. (3) Nitrous oxide is also formed 
by the action of an acid or an hyponitrite (see below). 

* BericAte, 15, 2600. 


Nitrous oxide, or hyponitrous anhydride, as it is sometimes 
named, is a colourless gas, possessing a faint sweetish smell and 
taste. It is somewhat soluble in water, and is best collected over 
hot water, or by downward displacement. When exposed to^i* 
sudden shock, as, for instance, the detonation of a fulminate, it 
explodes into its constituents ; this is a property common to bodies 
produced with absorption of heat. It is condensed by pressure to a 
liquid, boiling at 88 to 92 q , and when the liquid is evaporated 
by a current of air some of it freezes to a white solid, melting at 
99. Its most striking property is its action on the r nervous 
system when breathed, which has gained for it the name " laughing- 
gas." When pure, it produces insensibility, and is used as an 
anaesthetic in minor surgical operations and in dentistry ; but when 
diluted with air it causes excitement and intoxication. It easily 
decomposes when heated, hence a candle burning brightly con- 
tinues to burn more brightly in the gas. But if the candle is 
burning feebly it is extinguished. 

Compounds with oxides. Hyponitrites.* A solution of 
potassium nitrate or nitrite, exposed to nascent hydrogen generated 
from sodium amalgam (an alloy of sodium and mercury) loses oxygen, 
and potassium hyponitrite, KNO, is produced, about 15 per cent. 
of the nitrate or nitrite suffering change. The same compound is 
formed by fusing iron filings with potassium nitrate. The sodium 
salt forms white, needle-shaped crystals, and has the formula 
NaNO.3H 2 O. With silver nitrate, in presence of acetic acid, the 
silver salt is precipitated ; it is a pale yellow body, of the formula 
AgNO. On addition of hydrochloric acid to the silver salt sus- 
pended in water, the acid, presumably HNO, is liberated. Ifc 
reduces potassium permanganate: and on standing, decomposes 
into water and nitrous oxide. No other salts have been analysed ; 
but a solution of the sodium salt gives precipitates with solubl 
alts of most metals, almost all of which are insoluble in acetic acic 

We have thus a series of oxides and acids of nitrogen, vanadiun 
niobium, and tantalum : 

, nitrous oxide or hyponitrous anhydride. HNO acid. 
NO, nitric oxide. 

N 2 O 3 , nitrogen trioxide or nitrous anhydride. HK0 2 acid. 
W 2 O 4 , NO}, nitrogen tetroxide and peroxide. 
."NgOf, nitrogen pent/oxide or nitric anhydride. HNO 3 acid. 

HjNA, acid. 
nitrogen hexoxide. 

* Divers, Proc. Roy- Soc. t 19, 425 ; 33, 401 j Chem. Soe., 45, 78; 47, 361. 
t The hexoxide has been formed by passing sparks through a mixture of 

Similarly : 

Nb 2 6 

H 2 V 4 O tf (?) - - 

fH 3 V0 4 - 

Nb 2 6 .H 2 Ta 2 5 .H 2 0. 

Physical Properties* 

Mass of 1 cubic centimetre. 

Monoxides ..... See below. 

Dioxides ....... 

Trioxides ....... 

Tetroxides ...... 1-49 at 

Pentoxides ..... 

Vanadium. Niobium. Tantalum. 

3'64 at 20 
472 at 16 - 

3*5 at 20 4'37-4-53 7'35-8'Ol 

Mass of 1 c.c. N 2 0. 

Temp. -20-6 -11'6 -5*6 -2'2 +6'6 4-11-7 + 19'8 +237. 
Mass. 1-002 0-952 0'930 0'912 0-849 O'SIO 0'758 0'698 
NS, 2-22 at 15 ; VS 2 , 47 at 21 ; V&, 3'0. 

HN0 3 , 1-552, at 15. 2N 2 & Jl20, 1'642 at 18. VOC1 2 , 2'88 at 13 j 

VOC1 3 , 1-865 at 0. 

Heats of Combination. 

2J\r + Q ^ 2 o . 180K. 22V + 30 4- Aq = N 2 8 .Aq - 68K. 
N + = ^0 - 215K. N + 20 NO* - 77K. 
2^0 2 = JV 2 4 + 129K. 2N + 40 = ^ 2 4 - 26K. 
2N 4- 50 = N 2 5 + 131K; + Aq = 2HN0 3 .Aq -f 167K. 
N 2 4 = N 2 4 - 31K. N 2 8 = N 2 6 - 83K. 

Specific heat of gaseous N 2 4 or N0 2 .* 

T f26-5 f 277 / 28-9 / 29-0 f 292 / 27'6 

lem P ..... \667 \103'1 \ 150-6 \198-5 1.2531 \289-5 
Spec.|heat 0747 0-663 0-513 0'395 0'319 0-298. 

oxygen and nitrogen, cooled to 23. From volumetric measurements the 
compound produced a volatile crystalline powder is declared to have the 
formula N0 3 (Comptes. rend , 94, 1306). 

* This great change is due to absorption of heat in the conversion of N 2 4 
into N0 2 (Compt. rend., 64, 237). 




xides, Sulphides, Selenides, and Tellurides of 
Phosphorus, Arsenic, Antimony, and Bismuth. 

List of Oxides, Sulphides, Selenides, and Tellurides. 

P 4 0. _ 

P,0 6 . As 4 6 . Sb 4 6 . Bi 4 6 . 
P 2 4 . Sb 2 4 . Bi 2 4 . 

P 2 O 5 . As 2 O 4 . Sb 2 O 5 . Bi 2 O 5 . 

P 4 S 3 . - _ _ 

As 2 S 2 . Bi 2 S 2 . 

Sb 2 S 3 . Bi 2 S 3 . 

P 2 S 4 . - 

P 2 S 5 . As 2 S 6 . Sb 2 S & . 

Selenides and tellurides. P 2 Se 6 ; AsSeS 2 ; As 2 Te 2 ; As 2 Te 3 ; Sb 2 Se 3 ; 
SbTe; Sb 2 Te 3 ; Bi 2 Se a ; Bi 3 Te; Bi,Te 2 ; 

Sources. Pentoxide of phosphorus occurs in combination 
with oxides of metals, especially calcium and aluminium, as apatite, 
phosphorite, wavellite, &c. Arsenious oxide, A8 4 8 , is found as 
arsenite, or arsenic bloom ; and Sb 4 O 6 as antimony bloom in trimetric 
prisms, and as senarmontite in regular octahedra. The oxide, Sb 2 O 4 , 
is named antimony ochre. 

The sulphides As^a (realgar) As 2 S 3 (orpiment), Sb 2 S 3 (stibnite), 
and BLS 3 (bismuthine) also occur native ; as well as in combination 
with many other sulphides. 

Preparation. 1. By direct union. When phosphorus is 
burned in excess of air or oxygen, the pentoxide is formed. 
Arsenic and bismuth burn to trioxides ; and antimony to trioxide 
and tetroxide. In a limited supply of air, and at moderately high 
temperature, phosphorus gives P 4 O, P 2 O*, and P 2 O 5 ; by careful 
regulation of air a considerable amount of P 4 O 6 is produced, even 
as much as 50 .per cent., the other oxide being mainly P 2 O 6 . 

The process of preparing phosphorus pentoxide is to drop pieces of dry 
phosphorus through a tube passing through a cork closing the neck of a glass 


balloon, while a current of air, dried by passing through a U-tube filled with 
pumice-stone moistened with sulphuric acid, is blown in. The fumes are 

Fia. 40. 

condensed partly in the balloon, partly in the bottle communicating with it by 
a wide-mouthed tube. 

By the glowing of phosphorus in dry air the pentoxide is the only product. 

Arsenious oxide, As 4 O 6 , is usually produced by condensing 
in brick chambers the fumes resulting from the roasting in muffles 
of arsenical ores of tin, cobalt, and nickel, or arsenical pyrites. To 
purify it, the condensed product is sublimed in cast-iron pots. 

By limiting the supply of air, antimony burns to Sb^Oe, but* 
with free access of air, to Sb 2 O 4 . 

The sulphides, selenides, and tellurides of all these elements are 
produced by direct union. 

2. By decomposition of other oxides. Phosphorus tetr- 
oxide, P 2 O 4 ,* is produced by distilling in a vacuum the product of 
the combustion of phosphorus in a slow current of air. Bright 
orthorhombic crystals sublime, of the formula P2O 4 , arising from 
the decomposition of the phosphorous oxide, thus : 

7P 4 6 = 10P 2 O, + 2P 4 0. 

Arsenic pentoxide loses oxygen, forming trioxide at a dull red 
heat ; antimony pentoxide yields tetroxide at temperatures above 
275 ; and bismuth pentoxide, heated to 250, is converted into 

* Chem. Soc., 49, 833, 


tetroxide, and over 305 into trioxide. No known rise of tempera- 
ture, however great, deprives phosphorus pentoxide of oxygen. 

3. By oxidation in the wet way. This method in reality 
yields hydrates or acids. The usual oxidising agents are a mix^ 
ture of nitric and hydrochloric acids (aqua regia, see p. 341), or 
caustic potash and chlorine or bromine. Water cannot be expelled 
by heat from phosphoric acid, P 2 O 6 .H 2 O = HPO 3 ; but arsenic 
acid, As 2 O 6 .H 2 O, is dehydrated at a dull red heat, antimonic acid, 
Sb 2 O 6 .H 2 O, by heating not above 275, and hydrated bismuth 
pentoxide at 120. - 

4. By decomposition of compounds. The hydrates, as 
remarked above, lose water; and the nitrates and sulphates of 
antimony and bismuth decompose, when strongly heated, leaving 
trioxides. Phosphoryl chloride, POC1 3 , when heated with metallic 
zinc, yields zinc chloride and tetraphosphorns oxide, P 4 O ; and Ihe 
same body is formed by heating phosphoryl chloride with phos- 
phorus, thus : 

POC1 3 + P 4 = PC1 3 4- P^CX 

5. By double decomposition. As a rule, this process yields 
the hydroxides or acids, for example : PC1 3 4- 3H 2 O = H 3 PO 3 -f- 
3HC1 ; POC1 3 + 3H 2 O = H 3 PO 4 + 3HC1 ; 2SbOCl + 2KOH.Aq 
= Sb 2 3 .Aq + 2KCl.Aq + H 2 O ; 2BiCl 3 + GKOH.Aq = 
Bi 2 O 3 .H 2 O + GKClAq + 2H 2 O ; and, with the exception of the 
compounds of phosphorus, 'these yield oxides when heated. It 
forms, however, the usual method of preparing the sulphides, 
excepting those of phosphorus : e.g., 2AsCl 3 .Aq + 3H%S = As 2 S a 
+ GHCl.Aq ; 2SbCl 3 .Aq + 3H 2 S = Sb 2 S 3 + GHCl.Aq, &c. 

Properties. P 4 O is a light red or orange powder resembling 
red phosphorus, for which it was formerly taken ; when prepared 
by oxidation of phosphorus, it possesses reducing properties ; but 
when by depriving POC1 3 of chlorine, it does not reduce salts of 
mercury, silver, or gold. 

Phosphorous oxide or anhydride, P 4 O 6 > forms feathery 
crystals, melting at 22*5, and boiling at 173'3. It is decomposed 
by heat thus : 

2P 4 6 = 3P 2 4 + P 2 . 

It is slowly attacked by cold water, with formation of phosphorous 
acid, H 3 PO 3 , and immediately and with violence by hot water. It 
is luminous in the dark in presence of oxygen at a less pressure 
than that of the air ; and when heated gently in air, it burns to 
P 2 6 . It also burns in chlorine, forming POC1 3 and P0 2 01. 

The tetroxide forms orthorhombic crystals. It is soluble iu 


water, giving a mixture of phosphorous and phosphoric acids, 
thus : P 2 O 4 + 3H 2 = H 3 P0 4 + H 3 P0 3 . It is, therefore, sup. 
posed to have the formula P 2 4 or PO(P0 3 ) ; it would then be 
n^med phosphoryl metaphosphate. But of this there is no other 

The pentoxide or phosphoric anhydride is a snow-white 
powder, volatile below redness. It has a great tendency to com- 
bine with water, and is, therefore, used as a dehydrating agent, 
e.g., in the preparation of nitrogen pentoxide and sulphur tri- 
oxide. * When heated with carbon, it yields carbon monoxide and 

Arsenious oxide or anhydride, sometimes called arsenic 
trioxide, exists in three forms. When condensed at high tem- 
peratures, it is an amorphous porcelain-like mass ; its specific 
gravity is then 374. When cooled quickly^ or when it crystallises 
Irom solution, it furms colourless regular octahedra, the specific 
gravity of which is nearly the same, viz,, 3' 70. But when crys- 
tallised at low temperatures, or when it separates from its saturated 
solution in caustic potash, it forms rhombic crystals of the specific 
gravity 4*25. 

Arsenious oxide is sparingly soluble in water (vitreous, 4 in 
100 ; crystalline, 1*2 or 1*3 parts in 100 of water). It does not 
combine with water, but crystallises out from its solution in the 
anhydrous state. It is sparingly soluble in alcohol. Its vapour- 
density at a white heat corresponds to the formula As 4 6 .* It 
sublimes without fusion, but when heated under pressure it can be 

It is both an oxidising and a reducing agent, tending with 
certain oxides nitric acid, chromic acid, &c., to remove their 
oxygen, while it is itself reduced by carbon, phosphorus, sodium, &c. 
It is exceedingly poisonous ; less than 0'4 gram has been known to 
cause death ; but by continually increasing doses, the system may 
become inured to as much as 0'2 gram at a time. The antidote is 
a mixture of hydrated ferric oxide and magnesium chloride, pro- 
duced by adding magnesium oxide or carbonate in excess to tri- 
chloride of iron ; such a mixture forms an insoluble arsenite of 
iron, while the magnesium chloride and oxide act as a purgative. 

Arsenic pentoxide is a white mass, dissolving in water to 
produce arsenic acid. It is poisonous, but is not so deadly as the 

Antimonious oxide is found native in trimetric prisms as 
antimony -bloom, and in regular octahedra as senarmontile. It is a 
* Berichte, 12, 1112. 


white powder, turning yellow when heated, but white again on- 
cooling. It melts at a red heat, and volatilises at 1550. Its 
vapour-density points to the formula Sb 4 6 , like arsenious oxide.* 
It is insoluble in, and does not combine with water. One of th^> 
best solvents is a solution of tartrate of hydrogen and potassium 
(cream of tartar) HKC 4 H 4 6 .Aq ; it forms the potassium salt of 
the acid Sb(OH)C4H 4 6 , a substituted antimonious acid. 

Antimony tetroxide, Sb 2 O 4 , also occurs native as antimony 
ochre. It is a white powder when cold, and yellow when hot. It 
bas not been melted or volatilised. It is possibly metantimonate 
of antimonyl, SbO(SbO 3 ). 

The pentoxide, Sb 2 O 5 , is an insoluble lemon-coloured powder. 

Bismuth dioxide, Bi 2 O 2 , is a black crystalline powder, ob- 
tained by the reduction with tin dichloride of the trioxide sus- 
pended in alkali ; it must be dried out of contact with air. On 
treatment with acid, it gives a salt of the oxide Bi 2 O 3 , and a pre- 
cipitate of metallic bismuth. It oxidises at 180. 

The trioxide, Bi 2 O 3 , is a yellow- white solid, which crystallises 
from fused potassium hydroxide. No compound with an oxide is 
known, but it is not impossible that such a hot solution contains 
an easily decomposible bismuthite. 

The tetroxide, Bi 2 O 4 , is a brown-yellow solid, produced by 
treating the trioxide suspended in a cold solution of potash with 
chlorine; and the pentoxide, Bi 2 O A , is a red powder, similarly 
prepared, the solution of potash being kept boiling during passage 
of chlorine. The pentoxide combines with water, forming the 
hydrate Bi 2 O 6 .H 2 O. 

As hydrogen sulphide has no action on a solution of a phos- 
phate, the sulphides of phosphorus are prepared by direct 
union. There appear to be only three deQnite compounds.! Phos- 
phorus and sulphur may be melted together, but combination takes 
place only above 130". Owing to the great violence of the action 
and the inflammability of phosphorus in presence of air, a large 
quantity of sand is added to the melted mixture, and the retort is 
filled with carbon dioxide. If phosphorus is in excess, the com- 
pound produced is P 4 S 3 . This substance is reddish-yellow, melts 
at 167, and boils constantly about 380. If sulphur is in excess, 
the pentasulphide, P 8 S 5 , is formed, melting at 210 and boiling at 
519. Phosphorus and sulphur both dissolve in these compounds, 
but apparently without altering them. On heating a solution of 
the body P 4 S 8 ia carbon disulphide, however, with sulphur, yellow 
* Serichte, 12, 1282. 
t Bull. Soc. Chim.j 41, 433 j Comptes rend., 108,1386. 


crystals of th.6 compound P 2 S 4 are deposited ; and intermediate 
indistinct crystals are said to have been obtained of the formula 

The selenides of phosphorus are somewhat doubtful in com- 
position. The bodies P 4 Se, P 2 Se, P 2 Se 3> and P,Se 5 , are said to 
have been prepared, but, except perhaps the last, they are probably 
mixtures of compounds analogous to the sulphides. Phosphorus 
and tellurium apparently mix in all proportions \ no definite com- 
pounds have been isolated. 

Arsenic disulphide, ASsS 2 , is found native as realgar, in mono- 
clinic prisms. It is a reddish-orange body, and may be produced 
by heating arsenic and sulphur together in the right proportions. 
The trisulphide, A 8283, similarly produced, occurs native in 
trimetric prisms as orpiment ; it forms translucent lemon-yellow 
crystals. Prepared by double decomposition, it is a yellow powder, 
which is easily melted and volatilised. When hydrogen sulphide 
is passed into an aqueous solution of the trioxide no precipitate is 
produced, but the solution turns yellow. The substance in solution 
is probably a hydrate and hydrosulphide ; on addition of hydro- 
chloric acid, the trisulphide, As^S 3 , or more probably its compound 
with hydrogen sulphide, is thrown down. It is soluble in solu- 
tions of hydroxide or hydrosulphide of sodium or potassium, 
forming oxysulpharsenites and sulpharsenites (see below). The 
pentasulphide is an easily fusible yellow powder ; it is formed by 
direct union ; by addition of an acid to a sulpharsenate ; and by 
the action of a rapid current of hydrogen sulphide on a solution 
of arsenic acid. It is easily soluble in solutions of sulphides of 
the alkalies, forming sulpharsenates (see below). The action of 
a slow current of hydrogen sulphide on a solution of arsenic pent- 
oxide is first to reduce it, thus : 

As 2 6 . Aq -f 2H Z S = As 2 3 .Aq + 2H 2 -f 2S ; 

and then to precipitate the trisulphide.* 

Selenides of arsenic have not been prepared ; but two double 
sulphoselenides have been obtained by direct union, viz., Aa>SeS 2 , 
and As 2 SSe 2 . They are red bodies ; the latter may be distilled 
unchanged. The tellurides, also directly prepared, have the 
formulas ASaTe-j and As 2 Te 3 . 

Antimony trisulphide, Sb 2 S 3 , occurs native in trimetric grey 

metallic-looking or in orange-coloured prisms, as stibnite. It can be 

prepared by direct union, or by the action of hydrogen sulphide on 

a soluble salt of antimony. The former method yields crystals } 

* Bunsen, Annalen t 192> 30&; Brauner, Chem. Soc. y 53, 145. 


the latter, an orange-red powder, which, until dried, appears to he 
a hydrosulphide ; it dries to a hrown powder. * It turns grey at 
200220, and melts easily. The selenide, Sb 2 Se 3 , is a greyish, 
metallic-looking solid, produced by direct union; the telluride, 
SbTe, is iron-grey; and Sb 2 Te 3 , silver- white. The penta- 
Slllphide is not produced by direct union, but by decomposition of 
a sulphantimonate (see below) by an acid. It is a dark orange- 
coloured powder. The pentaselenide is a -brown precipitate, 
similarly prepared. 

Bismuth trisulphide, Bi 2 S 3 , is fouud in nature as bismutJi- 
glance, or bismuthine, in rhombic crystals, with steel -grey metallic 
lustre. A body of similar appearance is prepared by direct union, 
which becomes crystalline when heated with an alkaline sulphide. 
The brown-black precipitate, obtained by passing hydrogen sulphide 
through an acid solution of bismuth nitrate or chloride, is a com- 
pound of bismuth sulphide with water and hydrogen sulphide. 
The action of hydrogen sulphide on an alkaline solution of bismuth 
trioxide is said to yield the disulphide, Bi 2 S 2 , in combination with 
water. The triselenide is a black lustrous powder, similarly pre- 
pared; and the telluride is indefinite. The mineral telluric bismuth, 
Bi2S 3 .2Bi 2 Te,, occurs native. 

Compounds with Water and Oxides ; with Hydro- 
gen Sulphide and Sulphides ; with Selenides ; 
and with Tellurides. 

The constitution of the acids derived from the pent- 
oxides, pentasulphides, &c., of phosphorus, arsenic, and 
antimony. Phosphorus, it will be remembered, forms two 
chlorides, PC1 3 and PC1 5 (see p. 160). When the pentachloride is 
treated with a small quantity of water, an oxychloride, of the 
formula POC1 3 is produced (see below). The equation is : 

PC1 6 + H 2 = POC1 3 + 2HCI. 

It is probable that this oxychloride, which corresponds to those of 
vanadium, VOC1 3 , and niobium, NbOCl 8 , and to tantalum oxy- 
fluoride, TaOF 3 , is in reality the decomposition product of a 
dihydroxy trichloride, P(OH)2C1 3 , the reaction taking place thus : 

PC1 5 + 2H 8 = P(OH) 2 C1 3 + 2HC1; 
but that body beiug unstable forms an anhydride, thus : 
P(OH) S C1, = H a O + POC1 8 . 


The action of water on phosphoryl chloride, POC1 3 , is to yield 
orthophosphoric acid, PO.(OH) 3 , thus: 

POC1 3 + 3H 2 = PO(OH) 3 + 3HOL 
We have thus the series : 

01 xCl HO V /Cl .01 /OH 

>P^C1 ; >P^C1 ; 0=P^C1 ; and 0=Pf OH. 

C1 X X C1 HO X X C1 X C1 \)H 

The density of the vapour of phosphoryl chloride, POC1 3 , shows 
it to havfe the molecular weight corresponding to that formula ; 
and the fact that the hydrogen in orthophosphoric acid is replace- 
able in three stages by such a metal as potassium is a strong 
argument in favour of the analogy between phosphoryl chloride 
and phosphoryl hydroxide, or phosphoric acid ; such phosphates 
are : 

PO(OH) 2 OK; PO(OH)(OK) 3 ; and PO(OK) 3 . 

It would thus appear that phosphoric hydroxide, or the true 
orthophosphoric acid, should possess the formula P(OH) 5 ; but 
of this body, the first anhydride, PO(OH) 3 , is the one to which the 
name orthophosphoric acid is applied. 

By heating the first anhydride, PO(OH) 3 , the elements of water 
are expelled, and the second anhydride, metaphosphoric acid, 
PO^OHj, is produced, thus: 

PO(OH) 3 = H Z + PO 2 (OH). 

This substance is usually a monobasic acid, that is, its hydrogen is 
replaceable in one stage; hence its formula (see, however, p. 369), 
The analogous compound PO 2 01 has also been prepared. 

But intermediate between PO(OH) 3 and PO 2 (OH), there 
exists an acid of the formula H 4 P 2 O 7 , named pyrophosphoric acid. 
And corresponding to this hydroxide, P 2 O 3 (OH) 4 , a chloride, 
P 2 3 C1 4 , exists, which, however, has not been gasified, inasmuch as 
it decomposes. But arguing from the relation of the chloride 
POC1 8 to the acid PO(OH) 3 , the analogy of pyrophosphoric acid 
to pyrophosphoryl chloride appears justified, for its hydrogen is 
replaceable in fourths. And just as in the case of the silicic acids, 
acids are derived from two molecules of the ortho-acid with loss of 

0=p OH 

/ ^OH 
water, so here. We have therefore the series Ov ^TJ and 


the second member of which has the same compo- 



sition as metaphosphoric acid, but is a polymeride. Salts of this 
acid are called dimetaphosphates ; the acid is dibasic. Salts of the 
unknown acid, H 6 P 4 13 , are also known. Such an acid would be 
the fourth anhydride of tetraphosphoric acid, HuPiOn, also /in- 
known. And salts of the hypothetical acid, H 12 P 10 31 , are also 
known, which would be similarly derived. There are also tri-, 
tetra-, and hexa-metaphosphates, apparently corresponding to 
condensed acids. 

Such compounds can be also represented as formed by union 
of phosphoric anhydride with oxides. We have, for example, the 
series : 

P 2 O 6 .M 2 O = 2PO 2 (OM), monometaphosphates. 
P 2 O 5 .2M 2 O = P 2 O 3 (OM) 4 , pyrophosphates. 
P 2 O 5 .3M 2 O = 2PO(OM) 3 , orthophosphates. 

2P 2 O 5 .2M 2 O = 2P 2 O 4 (OM) 2 , dimetaphosphates. 

2P 2 O 6 .3M 2 O = P 4 O 7 (OM) 6 , a-phosphates. 

3P 2 O 6 .3M 2 O = 2P 3 O 6 (OM) 3 , trimetaphosphates. 

4)P 2 O5.4M 2 O = 2P 4 O 8 (OM) 4 , tetrametaphosp hates. 

6P 2 O 6 .6M 2 O = P 10 O 19 (OM)i 2 , |S-phosphates. 

6P 2 O 6 .6M 2 O = PeOj^OM)^ hexametaphosphates. 

Such compounds are, as a rule, soluble in water without decom- 
position. The sodium salts like - and /?-, however, named " Fleit- 
inann and Henneberg's phosphates," are decomposed by much hot 
water into mixtures of other salts. But the corresponding pyro- 
and meta-arsenates are converted into ortho-arsenates on treat- 
ment with water, unless they happen to be insoluble. For example, 
the ortho-arsenate, Na^HAsOi, is a well-known body ; on ignition, 
it loses water and yields Na^ASsO?, corresponding to the pyro- 
phosphate, NaJP 2 O 7 ; but on treatment with water, while sodium 
pyrophosphate dissolves as such, sodium pyro-arsenate reacts with 
the water, thus: NaiAs 2 7 + H 2 = 2Na2HAs0 4 . No ortho- 
antimonates are known except that of hydrogen, SbO(OH) a ; 
some pyroantimonates and many metantimonates have been pre- 
pared, and these have the formulae M 4 Sb 2 O 7 and MSbO 3 .* The 
nydrate of bismuth, Bi 2 O 5 .H 2 O, is analogous to a meta-acid ; it 
appears to be incapable of combination with other oxides. 

Compounds analogous to the orthophosphates have been pre- 
pared, in which the oxygen of the phosphate is partially replaced 
by sulphur, such as K 3 PSO 3 , Na3PS 2 O 2 , and possibly Na 3 PS 3 O. 
These bodies are termed thiophosphates or sulphophosphates. 
With selenium, compounds analogous to the pyrophosphates hava 

* Ifc is unreasonable to name compounds of the general formula M 4 Sb 3 O 7 
" metantimonates," as is usually done. These bodies have here been named 
systematically " pyroantimonates." 


been prepared, e.g., K 4 P 2 Se T . Orthothioarsenic acid, AsS(SH) 3 , is 
said to have been prepared ; and ortbo-, pyro-, and meta-thioarsenates 
are known. Similarly, orthothioantimonates are known : but no 
pyifc- or meta-derivatives have been prepared, nor are there any 

Double compounds of the pentoxides, &c.; phosphates 
and similar compounds. Ortho-acids. 

Orthophosphoric acid is formed by the oxidation of phos- 
phorus with boiling nitric acid, best in presence of a little iodine; 
by treating an orthophosphate with some acid which forms an in- 
soluble compound with the metal ; and by the action of a penta- 
halide or an oxytrihalide on water. If the first method be employed, 
the first product is phosphorous acid. The nitric acid should have 
the specific gravity 1'2, and should be employed in considerable 
excess ; and at the last, stronger acid may be employed to oxidise 
the phosphorous to phosphoric acid. The second method is the one 
employed on a large scale ; calcium orthophosphate, Ca^PO^, is 
mixed with sulphuric acid, and the precipitated calcium sulphate 
removed by subsidence. The equation is : 

Ca3(PO 4 ) 3 + 3H 2 SO 4 .Aq = 3CaSO 4 + 2H 3 P0 4 .Aq. 

It is common to use calcined bones or apatite (see p. 358) as 
the source of calcium phosphate. The third method is the most 
convenient for preparing phosphoric acid in the laboratory, and it 
may be coupled with the preparation of hydriodic acid. Bed 
phosphorus and iodine in the proportions equivalent to the formula 
I'ls are placed in a retort ; excess of water is added, and the mix- 
ture is distilled. Water distils over first, and then an aqueous 
solution of hydrogen iodide, while phosphoric acid remains in the 
retort. It is advisable then to evaporate the viscid residue with 
nitric acid. 

Orthophosphoric acid is also produced by dissolving phosphorus 
pentoxide in cold water, and boiling the solution of the resulting 
metaphosphoric acid ; and also by oxidation with nitric acid of 
hypophosphorous, phosphorous, and hypophosphoric acids. 

By spontaneous evaporation of its aqueous solution, it crystal- 
lises in long colourless prisms, melting at 41 '75, and has the 
formula H 3 PO 4 . From, the mother liquor of these crystals fresh 
crystals deposit on cooling, of the formula 2H 3 PO 4 .H 2 O ; these 
melt at about 27. Commercial phosphoric acid is a mixture of 
these two compounds. 

The solution of phosphoric acid is very sour ; the acid may be 

2 A 2 


heated to 160 without alteration, but at 212 it is largely con- 
verted into pyrophosphoric acid. 

By similar processes orthoarsenic acid is produced. The most 
convenient plan is to boil elementary arsenic, or arsenious oxide, rith 
nitric acid, or to pass chlorine through water with which powdered 
arsenious oxide is mixed. The solution is evaporated to dryness, 
and heated for some time to 100 ; water is then added, and on 
spontaneous evaporation the hydrated acid 2H 3 AsO4.H 2 O deposits 
in small needle-shaped crystals; and on heating to 150 ortho- 
arsenic acid, H 3 AsO 4 , remains. 

Orthoantimonic acid has been produced by treating potas- 
sium metantimonate, KSbO 3 , with nitric acid. It forms an in- 
soluble white precipitate. The usual product of this action, 
however, is metantimonic acid, HSbO 3 . 

The only corresponding sulphur compound is orthosulph- 
arsenic acid, H 3 AsS 4 , which is precipitated by addition of hydro- 
chloric acid to a solution of sodium sulpharsenate, Na 3 AsS 4 .Aq. 
Thiophosphates, on similar treatment, give ofE hydrogen sulphide, 
and yield phosphates. 

List of Orthophosphates and Orthoarsenates. The follow- 
ing have been prepared : 

Simple salts : 

2Li 3 P0 4 .H 2 0; Na 3 P0 4 .12H 2 ; K 3 PO 4 ; (NH 4 ) 3 PO 4 . 
2Li 3 As0 4 H 2 0; Na 3 As0 4 .12H 2 0; K 3 AsO 4 ; (NH 4 ) 3 As0 4 3H 2 O. 

Mixed salts : 

H2&iPO 4 ; H 2 NaPO 4 .H 2 O; H 2 KPO 4 ; H 2 (NH 4 )PO 4 . 
SHoLiAsO^H.O ; H 2 NaAs0 4 .H 2 0, and 2H 2 O ; 

H 2 XAs0 4 ; H 2 (NH 4 )As0 4 . 
HNa2P0 4 .12 and 7H 2 O ; H(NH 4 ) 2 PO 4 . 
HNa : Ae0 4 .12 and 7H 2 O ; HK 2 AsO 4 ; H(NH 4 ) 2 AsO 4 . 
(Li,Na) 3 PO 4 ; HNaKPO 4 ; HNa(NH,)PO 4 .4H. 2 O j 

Na(NH 4 ) 2 PO 4 .4H 2 O; HNaKAsO 4 .7H 2 O. 
Na 3 P0 4 .2NaF. 

These bodies are all white salts. They are prepared by the 
action of hydroxide or carbonate of lithium, sodium, or potassium, 
or of ammonia, on phosphoric or arsenic acid. The simple 
saltg are produced only by the action of hydroxide, if in solution, 
for carbonic acid decomposes them, giving a carbonate and a 
phosphate or arsenate containing an atom of hydrogen. But the 
carbonates ignited with the theoretical amount of phosphoric acid 
yield simple phosphates. The phosphates containing one and two 
atoms of hydrogen, however, cannot be made by fusion. 

Hydrogen di-lithium .phosphate has not been obtained pure. 


On adding hydrogen disodium phosphate to a concentrated solution 
of a soluble lithium salt, a precipitate is produced of the formula 
(Li 3 HPO 4 .LiH 2 PO 4 )H 2 O. It is a sparingly soluble salt (1 in 
20^ parts of water). The other salts are easily soluble. 

Hydrogen disodium phosphate, HNgi2POt.12H.iO, is the 
ordinary commercial " phosphate of soda;" the corresponding 
arssnate is also a commercial product ; they crystallise in mono- 
clinic prisms. The salt HNa(NH 4 )PO 4 .4H 2 O is known as " micro- 
cosmic salt," because it occurs in urine ; the human organism 
used to tie known as the " microcosm.*' It is used as a blowpipe 
reagent (see Metaphosphoric Acid). 

The following thiophospliates are similar in composition : * 

Na 3 P0 3 S 12H 2 ; Na { PO 2 S 2 .llH 2 O ; (NH 4 ) 3 PO 2 S 2 .2H 2 O. 

Salts of pota.ssium have been obtained in solution; and also 
sodium tritliiophosphattt, Na } POS 3 . These bodies are produced 
by the action of sodium hydroxide on powdered phosphorus penta- 
sulphide. They are unstable, especially the trithiophosphate, 
which decomposes when fche solution is heated to 50 ; a tempera- 
ture of 90 destroys the dithiophosphates, and they are precipitated 
by addition of alcohol to their aqueous solutions. They resemble 
the phosphates in appearance. 

Analogous oxythioarsenates have been made by dissolving 
arsenious oxide in a solution of sodium sulphide. They are 
separated by fractional crystallisation. Their formulas are : 
Na 3 AsO,S 12H 2 ; Na^HAsOjS 8H 2 O ; and Na 3 AsO 3 S 2H 2 O. 

Analogous to these are the thioarsenates, 2Na 9 AsS 4 .I5H 2 O ; 
K,AsS 4 ; (NH 4 ) 3 AsS 4 ; and Na 3 (NH 4 ) 3 (AsS 4 ) 2 . They are pro- 
duced along with pyro- and mefca-thioarsenates by digesting arsenic 
pentasulphide, As 2 S 5 , with alkaline sulphides, and evaporating 
the solution until crystals separate ; or by dissolving arsenic tri- 
sulphide, As 2 S 3 , in the solution of a polysulphide. They may also 
be produced by fusion. If arsenic pentasulphide be dissolved in 
solution of sodium or potassium hydroxide, a mixture of aryenate 
and thioarsenate is produced. They form yellowish crystals, and 
are very soluble in water. The solution of arsenic sulphide in 
ammonium polysulphide, a process used in qualitative analysis in 
order to separate sulphide of arsenic from sulphides of copper, lead, 
bismuth, mercury, and cadmium, depends on the formation of 
these bodies. The sulphides of antimony and of tin form similar 
compounds, and may be separated iu the same manner from sul- 
phides of lead, copper, &c. 

* J. prakt. Chem. (2), 31, 93. 


The thioantimonates maybe classed with the preceding salts. 
The following are known : 

Na 3 SbS 4 .9H20 j 2X 3 SbS 4 .9H 2 O. 

They are prepared by boiling a mixture of caustic alkali, sul- 
phur, and antimony trisulphide ; they form yellowish crystals. 
The sodium salt has been long known as u Schlippe's salt." The 
compound Na 3 SbSe 4 .9H 2 O forms orange-red tetrahedra ; it is pro- 
duced by fusing together sodium carbonate, antimony triselenide, 
sulphur, and carbon. Trisodium trithioseleno-antimonato, 

Na 3 SbS 3 Se.9H.O, 

is formed by boiling the tetrathioantimonate with selenium. It 
forms yellow crystals. 

Simple salts : 

/Be,(P0 4 ) 2 .7H 2 Oj Ca^PO^a; Ba.(PO 4 \ H 2 O. 

I - Ca,(As0 4 ) 2 ; Ba 4 (AsO 4 ) 8 . 

Mixed salts : 

f HOaPO 4 .4, 3, and 2H 2 O ; HSrPO 4 ; 
I HCaAsO 4 j HSr AsO 4 ; 

Ca 3 (PO 4 ) 2 2CaHPO 4 ; Ba 3 (PO 4 ) 2 .2BaHPO 4 . 

f Ca(H 2 P0 4 ) 2 ; Ba(H 2 P0 4 ) 2 j 

I Ba(H 2 As0 4 ) 2 ; 

fLiCaPO 4 ; KCaPO 4 ; NaSrPO 4 ; KSrPO 4 ; NaBaPO 4 ; 

I KBaFO 4 ; NaSrAsO 4 ; 2NH 4 CaAsO 4 H 2 O, also of Ba ; 

H 2 (NH 4 ) 2 Ca(As0 4 ) 2 ; also of Ba. 
3Ca 3 (PO 4 ) 2 .CaF 2 (apatite) ; 3Ca s (PO 4 ) 2 .CaCl 2 (apatite). 

CaCLj IL>O. 

The simple salts are produced by addition of the chloride of 
the metal to trisodium phosphate or arsenate. They are inso- 
luble white powders. The salts containing an atom of hydrogen 
are also insoluble, and are similarly precipitated with hydrogen 
disodium phosphate or arsenate. By boiling with water these are 
decomposed, giving the insoluble simple phosphate, while the 
soluble salt containing one atom of hydrogen goes into solution, 
The simple salt may also be precipitated by addition of excess o1 
ammonia, or of caustic soda or potash, to the mono- or di-hydrogen 
salts. These compounds are soluble in acids, the soluble di-hydrk 
salts being formed ; but are reprecipitated as simple salts on addi 
tion of alkaline hydroxide. 

Calcium phosphate is the chief mineral constituent of bones 
bone-ash, or calcined bones, contains about 93 per cent, ol 


Ca 3 (P0 4 )2. It is also widely distributed in soil. When fonnd 
native in combination with calcium chloride or fluoride, it is 
known as phosphorite, or apatite (see above) ; the chlorine and 
fluorine are mutually replaceable. Coprolites consist of the remains 
of the excreta of extinct animals, and are found in the Lias. They 
contain from 80 to 90 per cent, of phosphates. These bodies are 
largely used for artificial manure. 

To render the tricalcium phosphate soluble, so that its phosphorus may be 
easily assimilated by plants, it is treated with sulphuric acid in sufficient amount- 
to convert it into monocalciuin phosphate, thus: Ca 3 (P0 4 ) 2 + 2H 2 8Oj = 
Ca(H 2 P0 4 ) 2 + 2CaSO 4 . 

The mixture of monocaleium phosphate and sulphate is applied to the soil, 
usually mixed with organic matter containing nitrogen. The old plan of allow- 
ing land periodically to lie fallow had the effect of promoting a similar decom- 
position by aid of the carbon dioxide of the air. It appears that one part of 
tricalcium phosphate dissolves as monocaleium phosphate in from 12,000 to 
100,000 parts of water saturated with carbon dioxide. At the same time the 
carbon dioxide decomposes silicates, rendering their potash available for the use 
of plants ; and nitrogen in the form of ammonia collects on the soil, being 
brought down by rain. In the modern system of agriculture, artificial manure 
is applied to the soil, containing these substances in a soluble form ; the phos- 
phorus as monocaleium phosphate, the potash as chloride or carbonate, and the 
nitrogen as salts of ammonia, or as sodium nitrate ; or in the form of animal 
matter, from which ammonia is formed by putrefaction, such as manure, guano, 
dried blood, &c. 

Calcium arsenate, CaHAsO*, is found native as pharmacolite. 

The double phosphates and arsenates are produced by mixture. 
A arseriato-chloride, corresponding to apatite, has been produced 

The monothiophosphates of calcium, strontium, and barium 
are all insoluble white precipitates ; the dithiophosphates of 
strontium arid barium, and the trithiophosphate of barium, are 
stlbo insoluble. 

Thioarsenates of beryllium and of strontium have been pre- 
pared, but not analysed ; those of calcium and barium have the 
formulae Ca 3 (AsS 4 ) 3 , and Ba 3 (AsS 4 ) 2 ; they are insoluble yellow 
precipitates, produced by adding alcohol to the product of the 
actiou of hydrogen sulphide on HBaAsO*. The resulting thio- 
arsenate, HBa(AsS4), decomposes thus: 

3HBaAsS 4 .Aq = Ba 3 (AsS 4 ) 2 + BaAsS 3 .Aq -f J^S; 

the metathioarsenate remains dissolved. The corresponding thio- 
antimonate, Ba 3 (SbS 4 ) 2 , has also been obtained from the corre- 
sponding sodium salt by precipitation. 


Simple salts : 

M8T 3 (P0 4 ) 2 j Zn 3 (P0 4 ) 2 .5H 4 ; Cd 3 (PO 4 ) 2 ; 
Mgr 3 (As0 4 ) 29 Zn 3 (As0 4 ) 2 .3H 2 0; Cd 2 (AsO 4 ) 2 .3H 2 O ; 
Mixed salts: 

HMgrPO 4 7H 2 O ; HZnPO 4 .H 2 O ; Zn(H 2 PO 4 ) 2 .2H 2 O. 
HMffAs0 4 .7H 2 ; HZnAs0 4 ; H 2 Cd 5 (AsO 4 ) 4 4H 2 O. 
NaMgrP0 4 ; KMgPO 4 ; H 2 (NH 4 ) 2 Mg(PO 4 ) 2 3H 2 O j 

NH 4 MgrPO 4 6H 2 O ; NH 4 ZnPO 4 .2H 2 O. 
KMgrAsO 4 ; NaMg-AsO 4 .6H 2 O. 

(wagnerite) = PO 

^OMff F 

Similar arsenates have been prepared artificially. 

Trimagnesium orthophosphate is a constituent of the ash of 
seeds, especially of wheat. It and the corresponding arsenate are in- 
soluble in water. The other salts are produced by precipitation, and 
are sparingly soluble. The most important are ammonium mag- 
nesium phosphate and arsenate. The former is a constituent 
of certain urinary calculi, and is formed by the puti-efaction of 
urine, and separates in crystals. Both of these salts are very 
sparingly soluble in water (about 1 in 13,000), and are used in the 
estimation of magnesium, and of phosphoric and arsenic acids. 
They are produced by adding a solution of magnesium chloride 
and ammonium chloride, commonly called " magnesia mixture," 
along with ammonia, to a soluble phosphate or arsenate. On igni- 
tion they leave a residue of pyrophosphate or pyroarsenate. 

Thioarsenates of magnesium, zinc, and cadmium have also been 
prepared : they are soluble crystalline salts 

BPO 4 ; 2YPO 4 .5H 2 O (xenotime) ; LaPO 4 . 
YAs0 4 

Boron phosphate is an insoluble white substance produced by 
heating boron hydrate with orthophosphoric acid. Yttrium phos- 
phate and arsenate and lanthanum phosphate are white gelatinous 
precipitates produced by double decomposition. Yttrium phos- 
phate occurs native, and that of lanthanum occurs in several rare 

A1P0 4 .3 and 4H 2 O ; AlAsO 4 .2H 2 O. 

Phosphates of aluminium and hydrogen: A1(H 2 PO 4 ) 3 and 

Basic phosphates of aluminium : 6A1P0 4 A1 2 3 .18H 2 O ; 
4A1P0 4 A120 8 .12H 2 ; P 2 O 5 .2A1,O 3 8, 6, and 5H 2 O. 

Thallous phosphates :^-Tl 3 PO 4 j HTl2PO 4 .H 2 O j H 2 T1PO 4 . 

Aluminium phosphate, produced by precipitation, is a white 
bulky precipitate, closely resembling hydrated alumina, from which 
it is difficult to distinguish and to separate. The arsenate closely 


resembles the phosphate. The compound A1PO 4 .4H 2 O occurs 
native as gibbsite ; it is also produced on boiling a solution of 
hydrogen aluminium phosphate. The first basic phosphate is pro- 
duped by adding ammonia to a solution of the orthophosphate in 
hydrochloric acid ; the second is wavellite. The third, with 5H 3 0, 
is turquoise, which owes its blue colour to a trace of copper ; with 
6H 2 it is peganite, and with 8H 2 O it forms crystals of fischerite. 

Thallic arsenate is a flocculent, insoluble precipitate; the 
thallous phosphates are nearly insoluble, and separate from dilute 
solutions in crystals. 

CrP0 4 7, 6, 5, and 3H 2 O j FePO 4 . 

The chromic salt exists in two forms : the violet modification, 
with 7H 3 0, which is soluble and crystalline, and is produced by 
treating a solution of violet chromic chloride with silver phos- 
phate ; and the green modification, precipitated by addition of a 
soluble phosphate to a green chromium salt. The violet variety, 
when heated, changes into the green one; and the green precipi- 
tate becomes violet and crystalline on standing. Ferric phosphaio 
is a white precipitate produced in a neutral solution of a ferric salt 
by hydrogen disodium phosphate, or by exposing ferrous phosphate 
to air. Arsenates give similar precipitates with chromium and 
iron salts. 

Iron also forms the following double phosphates with hydro- 
gen : 

Fe(H 2 P0 4 ) 3 ; FeH a (PO 4 ) 2 ; Fe 6 H 3 (PO 4 ) 7 ; Fe 8 H 3 (PO 4 ) 9 ; and Fe 4 H 3 (PO 4 ) 5 . 
Also the basic phosphates P 2 O 5 2Fe 2 O 3 .12H 2 O (cacoxem) ; 5H 2 O (dufren'de 
or green iron, ore) ; and 12H 2 O or 18H 2 O (delvauxite) . 

Basic ferric phosphate is also a frequent constituent of bog- 
iron ore. Manganic and cobaltic orthophosphates and arsenates 
are unknown. 

Simple salts : 

Cr 3 (P0 4 ) 2 (?) ; Fe 3 (P0 4 ) 2 .8H 2 ; Mn 3 (PO 4 ) 2 7H 2 ; Co 3 (P0 4 ) 2 .8H 2 O ; 
Ni 3 (P0 4 ) 2 .7H 2 ; Fe. t (As0 4 ) 2 ; Co 3 (AsO 4 ) 2 .8H 2 O a 
(cobalt- and nickel -bloorn) . 

Mixed salts : 

Fe(H 2 PO 4 ) 2 .2H 2 O ; Mn(H 2 PO 4 ) 2 .2H 2 O ; 

(NH 4 )Fe(P0 4 ).H 2 ; (NH 4 )Mn(PO 4 ).H 2 O. 
AJso the arsenates, Co (H 2 AsO 4 ) 2 ; Mn(H 2 AsO 4 ) 2 ; MnHAsO 4 . 
And the minerals childrenite, a phosphate of aluminium, iron, and man- 
ganese : tripltte, (Fe,Mn) 3 (PO 4 ) 2 , and triphylline, (Li2,Mfir,Fe,Mn) 3 (PO 4 )2. 

Chromous phosphate is a blue precipitate ; ferrous phosphate 
is white and insoluble ; it occurs native as vivianite or blue iron 
earth; the hydrogen manganous salts and the double ammonium 


salts are obtained by mixture and crystallisation, e.g., Mn 3 (PO 4 ) 2 -f- 
H 3 PO 4 .Aq = 3HMnPO 4 + Aq ; Mn 3 (PO 4 ) 2 + (NH 4 ) 3 PO 4 .Aq = 
3NH 4 MnPO 4 + Aq. The cobalt salt is reddish-blue, and the nickel 
salt light-green. Arsenates of cobalt and nickel occur native ; 
cobalt-bloom forms red, and nickel-bloom green, crystals. 

Elements of the carbon-group form no normal phosphates. 
Carbon phosphate is unknown ; titanium forms the compound 
Ti 2 Na(PO 4 ) 3 when titanium dioxide is fused with hydrogen sodium 
ammonium phosphate (microcosmic salt) ; and sodium thorium 
phosphate, Th 2 Na(PO 4 ) 3 , is similarly prepared ; zirconium salts by 
precipitation give the basic phosphate (ZrO) 3 (PO 4 ) 2 ; thorium 
phosphate is a white precipitate ; cerous phosphate, CePO 4 , occurs 
native as cryptolite and phosphocerite ; prepared artificially, it forms 
a white precipitate. Arsenates of titanium, zirconium, and thorium 
have been prepared ; also cerous arsenate, CeAsO* (?) and sulph- 
arsenate, CeAsS 4 (?), which require investigation. 

SiH 2 (PO 4 ) 2 .3H 2 O is deposited from a solution of silica in 
phosphoric acid kept at 125 for several days. It is soluble in, 
and decomposed by contact with, water. Germanium phosphate 
has not been prepared ; a basic phosphate of tin, P 2 O 8 .2SnO 2 .10H 2 O, 
is deposited on treatment of tin dioxide (metastannic acid) with 
phosphoric acid; this compound is insoluble in nitric acid, and is 
therefore used in separating phosphoric acid from solutions con- 
taining it. Corresponding ar senates are unknown. By fusing 
stannic oxide with borax and microcosmic salt, crystals of 
the formula Na 2 Sn(PO 4 ) 3 are produced. With microcosmic 
salt alone, the body NaSn 2 (PO 4 ) 3 is formed in microscopic 

Stannous phosphato-chloride, Sn 3 (PO 4 ) 2 .SnCl 2 , is precipitated 
by adding a solution of ordinary sodium phosphate to excess of tin 
dichloride ; but with excess of sodium phosphate the precipitate 
has the formula Sn (PO 4 ) 2 .2SnHPO 4 .3H 2 O. The arsenates, 
similarly produced, are said to have the formulae 2HSnAsO 4 .H^O 

and C1 |>As0 4 .H 2 0. 

Lead orthophosphate, Pb 3 ( r O 4 ) 2 , produced by precipitation, is 
a white amorphous substance, fusible, and crystallising on cooling. 
By adding phosphoric acid to a dilute boiling solution of lead 
nitrate, the compound Pb(H 2 PQ 4 ) 2 is thrown down in sparkling 
white laminee. In the cold, a phosphato-nitrate, of the formula 
Pb(NO) 3 .Pb 3 (P0 4 ) 3 .2H,O, is precipitated. It is decomposed by 


boiling water. By employing a boiling solution of lead chloride 
and excess of sodium phosphate, the compound 
Pb 3 (P0 4 ) 2 .PbCl 2 .H 2 

ia precipitated. With excess j of lead chloride the precipitate con- 
sists of 2Pb %J (PO 4 ) 2 .PbClo (?). Pyromorphite, another phosphato- 
chloride, 3Pb 3 (PO 4 ) 3 .PbCl 2 , occurs native in hexagonal prisms, 
usually of a green colour. The corresponding arsenate, 

3Pb 3 (AsO 4 ),.PbCl 2 , 

is also found in nature, and is named mimetesite. Crystals in 
which arsenic and phosphorus replace each other partially are 
common. The arsenates Pb 3 (AsO 4 )j and HPbAsO 4 have been 
produced by precipitation, and also the sulpharseriate, Pb a AsS 4 . 

(VO)P0 4 .7H 2 ; (VO)As0 4 .7H 2 ; (VO) 2 H,(PO 4 ) 3 .3H,O. 
These are the simpler phosphates and arsenates of elements of 
the nitrogen group. They are brilliant yellow or red crystals. It 
is to be noticed that these bodies may equally well be conceived as 
vanadates of phosphoryl and arsenyl, thus : 

(PO)VO 4 .7H 2 O ; (AsO)V0 4 .7H 2 O , and (PO.OH) 3 (VO 4 ) 2 3H 2 O. 

Tantalum pentoxide, dissolved in hydrochloric acid, forms a jelly 
with phosphoric acid, due probably to a combination between them. 

A curious compound of the formula 4MgHPO 4 .NO 2 is produced 
by boiling magnesium pyrophosphate with strong nitric acid, and 
heating it in a paraffin-bath until it ceases to emit fumes. It is 
a crystalline whitish-yellow powder, which gives off nitric per- 
oxide when strongly heated. 

The vapour-density, and consequently the molecular weight, of 
phosphorus pentoxide is unknown. If its formula be P*O 6 , it 
may perhaps be regarded as phosphoryl phosphate, (PO)PO 4 , 
O=P~O 3 ^P=O ; and arsenic pentoxide and the other pentoxides 
might be similarly regarded. 

Many very complicated compounds of the pentoxides with 
each other have recently been discovered. Among these are 

P 2 5 .V 2 O 6 .(NH 4 ) 2 O.H 2 O ; 4P 3 O 3 .6V 2 5 .3K,O.21H,0 ; 
P 2 O 5 .20V 3 O ft .69H 2 ; 5As 2 O 6 .8V 2 O s ,27H 3 O. 

Some also contain vanadium dioxide, for example, 

2P 2 O 6 .V0 2 .18V 2 O 6 .7(NH 4 ) 2 O.50H 2 0. 

Compounds of arsenious and arsenic oxides are also known ; 
thus : 
2ASjO6.3ASaOj.HaO ; ASaO 6 .2As 4 O 3 .H 2 O j and ASaO 6 .ASaOj.H 2 O. 


They are produced by partial oxidation of arsenioiiB oxide, 
As 4 O t; , by nitric acid, and are definite crystalline bodies.* The 
bismuth phosphate corresponding to the last of these, BiPO 4 = 
P 2 O 6 .Bi 2 O 3 is produced by precipitation. The corresponding 
arsenate, BiAsO 4 .H 2 O is a yellowish-white precipitate ; they may, 
however, equally well be regarded as metaphosphate and met- 
arsenate of bismuthyl, (BiO)POj and (BiO) AsO^. 

The compounds with the elements molybdenum and tungsten 
are exceedingly complicated. Molybdenum trioxide, MoO 3 , and 
tungsten trioxide, WO d , combine with phosphorus tri- and pent- 
oxides, and with many other oxides ; these compounds will be 
described among the oxides of molybdenum and tungsten. The only 
one to be mentioned here is ammonium phosphomolybdate, which 
is produced by adding ammonium molybdate to any warm solu- 
tion containing an orthophosphate. It is a bright yellow precipitate, 
insoluble in nitric acid, and is used as a test for phosphoric acid. 
Several compounds of uranyl, (U0 2 ), are known. The normal salt 
has not been prepared, but double salts are known, for example, 

(UO 2 ) 3 (PO 4 ) 2 .2 (UO >)HPO 4 .H 2 O, 

which is formed by precipitation as a light yellow powder. By 
digestion with phosphoric acid, the salts (UO 2 )HPO 4 , and 
(UO 2 )(H 2 PO 4 ) 2 are formed ; corresponding arsenates have been 
prepared. Uranyl sodium salts, (UO 2 )NaPO 4 and (UO 2 )NaAsO 4 , 
are produced by addition of sodium phosphate in excess. The 
calcium salt, (UO 2 ) 2 Ca(PO 4 ) 2 .8H 2 O, is found native as uranite ; 
and a similar copper salt, (UO 2 ) 2 Cu(PO 4 ) 2 8H 2 O, occurs as chalco- 

Phosphates and arsenates of the palladium and platinum 
groups of metals require investigation. No compound has been 
analysed (except H^Rh(PO 4 ) 2 .H 2 O), although salts of these metals 
give precipitates with phosphates and arsenatos. Compounds of 
gold are unstable. 

Copper orthophosphate, Cu^PO^, is a blue-green precipitate ; 
or, when prepared by heating the pyrophosphate with water, 
yellowish-green crystals with 3H 2 O. The salt HCuPO* is also a 
blue-green precipitate. Many basic compounds occur native, e.g., 

P 2 O 6 .4CuO.H a O, 2H.O, and 3H 2 O ; P,O 6 .5CuO.2H 2 O and 
3H 2 O ; and P 2 O 6 .6CuO.3H 3 O. 

The last is the most important, and is named pliospTiochdlcite. 
* Comptes rend., 100, 1221. 


The arsenates, Cu^AsO^ and H 2 Cu 2 (AsO 4 ) .H 2 O, are green 
and blue powders respectively. 

Silver phosphate, Ag,PO 4 , is a yellow precipitate, produced by 
adding any soluble phosphate to a solution of silver nitrate. It is 
used as a test for phosphoric acid. Hydrogen disilver phosphate, 
HAg 2 PO 4 , produced by digesting the former with phosphoric acid, 
forms colourless crystals; it is at once decomposed by water into 
Ag.jPO 4 and H 3 P0 4 . The arsenate, Ag,AsO 4 , is a red precipitate. 
It is formed by adding an arsenate to a solution of silver nitrate, 
and cautiously adding ammonia. It serves as a test for arsenic 
acid, and distinguishes it from arsenious acid. 

Mercurous phosphate, Hg,,PO 4 , and mercuric phosphate, 
Hg i (PO 4 ) i , are white crystalline powders. A. phosp ha to- nitrate, 
HgjPO^HgNOi.H.O, is also known. The arsenate Hg^HAsO 4 is 
an orange precipitate. 

Pyro- compounds. Pyrophosphoric acid, H 4 P 2 O 7 = 
P-zO^OH)!, is produced by heating orthophosphoric acid to 215. 
The change begins at 160, but is not complete at 215, for the mass 
still contains unchanged orthophosphoric acid. If a higher tem- 
perature be employed, meta phosphoric acid begins to be formed, 
Similarly, pyroarsenic acid is formed by heating the ortho-acid 
to 140 160. Pyroantimonic acid, unlike the corresponding 
acids of phosphorus and arsenic, is produced by the action of 
water on the pentachloride. When SbCl 5 is mixed with a little 
water, crystals of the formula SbCl 5 .4H 2 O are deposited. 
Addition of more water to the cold solution of this body produces 
the insoluble oxychloride, SbOCl, ; on warming this antimonyl 
chloride with much water, the sparingly soluble pyroantimonic 
acid, H 4 Sb 2 O 7 .2H 2 O is formed. The water of crystallisation may 
be expelled at 100. No corresponding compound of bismuth is 

These bodies may also be prepared by replacing some metal 
such as lead, in the pyro-salts, by hydrogen, by the action of 
hydrogen sulphide, thus : 

Pb,P 2 O 7 + 2H 2 S + Aq = H 4 P 3 O 7 .Aq + 2PbS. 

The lead pyrophosphate is insoluble, and is suspended in water. 
Pyroarsenic acid, however, cannot be thus prepared, for it reacts 
with hydrogen sulphide, giving arsenic pentasulphide. But as 
pyrantimonic acid is sparingly soluble, it is precipitated on adding 
an acid to a solution of a pyroantimonate ; e.g., 

q + 4HCl.Aq = HJ3b 2 O 7 + 4KCLAq. 


On standing, even in contact with water, it loses water, changing 
to HSbO 3 , thus : 

H 4 Sb 2 O 7 = 2HSbO 3 + H 2 0. 

No pyrosulpho- or pyroselenio-acids are known. 

Pyrophosphoric acid is usually a soft colourless glass-like 
body ; it has, however, been obtained in opaque indistinct crystals. 
Pyroarsenic acid forms hard shining crystals ; it unites with water 
at once, giving out heat, and forming a solution of orthoarsenic 

Pyroantimonic acid is a white powder, soluble in a large 
quantity of water, from which it is precipitated by addition of 

Pyrophosphates, &c. The pyrophosphates and pyroarsenates 
are produced by heating the mono-hydrogen or mono-ammonium 
orthophosphates to redness, thus : 

2HNa 2 PO 4 = Na 4 P 2 O 7 + JBT 2 0; 
2NH 4 MgPO, = Mg,P 2 O 7 + 2NH, + H 2 0. 

The pyroarsenates require investigation. It is possible that on 

treatment with water dimetallic orthoarsenates are again formed, 

but this has not been proved. The pyroantimonates are produced 

by heating the metantimonates with water, or with an oxide, 

- thus : 

2MSbO 3 + M 2 = M 4 Sb 2 7 , and 2MSb0 3 + H 2 = M 2 H 2 Sb 2 O 7 . 

The pyrophosphates may also be produced by action of pyro- 
phosphoric acid on oxides, hydroxides, or carbonates. 

Pyrothioarsenates are the salts usually produced by dissolving 
arsenic pentasulphide in solutions of soluble sulphides, or hydrosul- 
phides, or the trisulphide in solutions of polysulphides ; or by the 
action of hydrogen sulphide on solutions of the arsenates ; or by 
fusing the sulphides of arsenic and metal together. Many are 
insoluble, and are precipitated on addition of the sodium salt to a 
solution of a compound of the element. On treatment with 
alcohol, they are often decomposed into orthothioarsenates, which 
are precipitated, while the meta-salts dissolve. 

List of Pyrophosphates, &o, 
Simple salts : 

Na 4 P 2 7 .10H 2 O ; X 4 P 2 O 7 .3H 2 ; (NH 4 ) 4 P 3 O 7 . 

Mixed salts : 

H 2 X S P 2 7 ; 2HK 2 (NH 4 )P 2 O 7 .Hc,O ; 


The pyrophosphates are produced by addition of a hydr- 
oxide or carbonate to the acid ; many of them are precipitated 
by alcohol. They are white deliquescent salts, and they are not 
altered by boiling with water; but, on boiling with acids, they 
combine with water, forming orthophosphates. The double salts 
are produced by mixture and crystallisation. On heating dihydro- 
gen monosodium orthophosphate, H^NaPO^, to 200, it loses water, 
giving dihydrogen disodium pyrophosphate, thus : 2H 2 NaPO 4 = 
H 2 Na 2 P 2 O 7 + H 2 0. Potassium pyroantiraonate, K 4 Sb 2 O 7 , is pro- 
duced by fusing the metantimonate, KSbO 3 , with caustic potash, 
and subsequent crystallisation from water. Dihydrogen dipotas- 
sium pyrantimonate is formed, along with potassium hydroxide, by 
warming the tetrapotassium salt with water. The corresponding 
sodium salt is very sparingly soluble in water it is one of the 
few nearly insoluble salts of sodium and the formation of a pre- 
cipitate in a solution free from other metals on addition of a solu- 
tion of the potassium salt indicates the presence of sodium, owing 
to the formation of 

Simple salts : 

Be2P 2 O 7 .5H 2 O ; Ca2P 2 O 7 .4H 2 O ; Sr 2 P 2 O 7 .H 2 O ; Ba2P 2 O 7 .H 2 O ; Ca 2 As,S r ; 

and others. 
Mixed salts : 

Na 2 CaF 2 O-.4H 2 O j and insoluble white pyrantimonates. 

Hydrogen pyrophosphate gives no precipitate with the 
chlorides of these metals; but with sodium pyrophosphate 
these pyrophosphates are precipitated. The calcium salt fuses to 
a transparent glass, which may be substituted for ordinary glass 
for many purposes. 
Simple salts : 

Mff2P 2 7 3H 2 ; 2ZH2P 2 7 .H 2 and 10H 2 O ; Cd 2 P 2 O 7 .2H 2 O ; M* 2 As 2 S 7 . 
Mixed salts : 

Na 2 ZnP 2 O 7 , also with 4H 2 O ; Na 2 CdP 2 O 7 . 

The anhydrous magnesium pyrophosphate is left as a white 
caked mass on. igniting ammonium magnesium orthophospbate, 
NH 4 MgPO 4 . These anhydrous salts are soluble in sulphurous 
acid, and crystallise from the solution on evaporation. The double 
salts crystallise from solutions of oxides in sodium metaphosphate. 
The sulpharsenate of magnesium is a very soluble yellow salt, also 
soluble in alcohol. 

Pyrophosphates, <fec., of the boron group of elements have not 
been prepared. 

Al4(PaO 7 )8.10H 2 O is a white precipitate, differing from the 


orthophosphate by its solubility in ammonia. The salts of gallium, 
indium, and thallium have not been prepared. The double salt, 
NaAlP 2 O 7 , crystallises from a solution of A1 2 O 3 in fused sodium 

Pyrophosphate of carbon is unknown ; titanium, zirconium, and 
tin pyrophosphates, TiP 2 O 7f ZrP 2 O 7 , and SnP 2 O 7 , are prepared by 
dissolving the dioxides in fused orthophosphoric acid. 

Silicon pyrophosphate,* SiP 2 O 7 , crystallises in octahedra from 
a solution of silica in fused metaphosphoric acid, and lead pyro- 
phosphate, Pb 2 P 2 O 7 .H 2 O, produced by precipitation, is a bulky 
white powder. Pb 2 As 2 S 7 is also known. 

Cr 4 (P 2 7 ) 8 ; Fe 4 (P 2 O r ) s .9H 2 0.-2Na 4 P 2 7 .Fe 4 (P 2 7 ) 3 .7H 2 ; Fe 4 (As 2 S 7 ) 3 . 

These salts are produced by precipitation ; that of chromium 
is green, and those of iron nearly white. They are soluble in 
excess of sodium pyrophosphate, and doubtless form salts like 
the double salt of iron of which the formula is given above. Ammo- 
nium sulphide does not precipitate chromium or iron from solu- 
tions of these double salts. NaCrP 2 O 7 crystallises from a solution 
of chromium sesquioxide in sodium metaphosphate. 

Mn 2 P 2 O 7 .3H 2 O ; Co 2 P 2 O 7 ; Ni^O^HgO; NaNHiMnP 2 O 7 .3H 2 O; 

Fe 2 As 2 S 7 ; Mn 2 As 2 S ; ; and Co.2As 2 S 7 are produced by precipitation. 

Of the nitrogen and phosphorus groups, the only pyrophos- 
phate known is that of bismuth, Bi 4 (P 2 O 7 ) 3 , which is a white pre- 
cipitate. It crystallises from a solution of bismuth trioxide in 
fused sodium metaphosphate. But hydrogen sodium pyrophoB- t 
phate dissolves antimony trioxide. The pyrophosphates of elements 
of the palladium and platinum groups have not been prepared. 

Cupric pyrophosphate, Cu 2 P 2 O 7 .H 2 O, is a greenish- white 
powder produced by precipitation. Silver pyrophosphate, Ag 4 P 2 O 7 , 
is a white curdy precipitate. Its formation serves to distinguish 
pyrophosphates from orthophosphates, which give a yellow pre- 
cipitate of Ag 3 PO 4 with silver nitrate. A double pyrophosphate of 
gold and sodium, of the formula 2NaJ? 2 O 7 .Au 4 (P 2 O 7 ) 3 .H 2 O, is 
formed by mixing gold trichloride with sodium pyrophosphate and 
evaporation ; the sodium chloride separates in crystals, leaving the 
above salt. Mercuric pyrophosphate, Hg 2 P 2 O 7 , and mercurous 
pyrophosphate, Hg t P 2 O 7 , are white precipitates. 

It is to be noticed that while there are many double pyrophos- 
phates in which the two atoms of hydrogen of pyrophosphoric acid 
are replaced by one metal, and two by another, such as H 2 NaaP 2 O 7 , 

Comptes rend., 96, 1052 ; 99, 789 j 102, 1017. 


Na2CaP 2 O 7 , &c., there are few in which the hydrogen is replaced 
in fourths. Yet instances are known, for example, NaNH 4 MnP 2 O 7 , 
HK 2 NH 4 P 2 O 7 , and one or two others. The conclusion is therefore 
justified that, inasmuch as such compounds are known, there are 
four atoms of hydrogen in hydrogen pyrophosphate. With the 
pyrothioarsenates and pyroantimonates, such double compounds 
are unknown : the only double salts being those of the pyroanti- 
monates of hydrogen and a metal such as H 2 Na 2 Sb 2 O 7 . 

Meta-compounds. Metaphosphates, etc. It cannot be said 
with certainty that more than one metaphosphoric acid is known, 
although, as mentioned on p. 354, there are grounds for infer- 
ring the existence of at least five sets of metaphosphates : mono-, 
di-, tri-, tetra-, and hexa-metaphosphates, derived from condensed 
acids.* When phosphoric anhydride is dissolved in cold water, and 
the resulting solution evaporated, or when orthophosphoric acid is 
heated above 213, a transparent glassy soluble substance remains, 
the simplest formula of which is HPO 3 . The same body is pro- 
duced by (1) heating microcosmic salt to redness, when sodium 
metaphosphateis produced, thus: HNaNH 4 PO 4 = NaPO 3 -f H 2 
-f NH 3 ; (2) dissolving this metaphosphate in water, and adding lead 
mtrate, when lead metaphosphate is formed, thus: 2NaP0 3 .Aq -f 
Pb(N0 3 ) 2 .Aq = 2NaNO 3 .Aq + Pb(PO 3 ) 2 ; and (3) suspending the 
insoluble lead metaphosphate in water, and passing through the 
liquid a current of hydrogen sulphide, when lead sulphide and 
hydrogen metaphosphate are produced, thus : Pb(PO 3 ) 2 4- Aq + 
#2$ = 2HP0 3 .Aq -f PbS. On evaporating the filtered liquid 
to dryness, the same glassy soluble body is obtained. It is 
probably a hexametaphosphoric acid, for it forms salts in which 
one-sixth of the hydrogen is replaceable. 

But it has been noticed that during the preliminary stage of 
phosphorus manufacture, in evaporating orthophosphoric acid with 
charcoal or coke, and igniting the residue, the black powder of 
carbon and metaphosphoric acid gives up nothing to water ; an 
insoluble variety is in fact produced. This variety differs therefore 
from the other, and is possibly monometaphosphoric acid, for that 
body gives insoluble salts. 

On boiling metaphosphoric acid with water, orthophosphoric 
acid is formed, thus : HPO 8 . Aq 4- H 2 O = H 3 PO 4 . Aq. The meta- 
acid, when added to a solution of albumen (white of egg) in water, 
coagulates it, producing a curdy precipitate; the silver salt is 
white, and is not produced on adding silver mtrate to a solution of 

See also Zeitschr.f. phytik. Chem., 6, 122. 

2 B 


metaphosphoric acid ; and it gives no yellow precipitate when 
warmed with ammonium molybdate and nitric acid. But, after 
boiling with water, the resulting orthophosphoric acid does not 
coagulate albumen, gives a yellow precipitate of silver ortjjo- 
phosphate, Ag 3 PO 4 , with silver nitrate, and a bright yellow 
precipitate with ammonium molybdate. The two acids are there- 
fore obviously distinct bodies. They are distinguished from 
pyrophosphoric acid by the fact that silver pyrophosphate is white 
and curdy. 

Metaphosphoric acid is volatile at a high temperature, but it 
does not lose water to give phosphorus pentoxide. 

Metarsenic acid, HAsO 3 , is likewise produced by heating 
ortho- or pyroarsenic acid to 200 206. It is a white nacreous 
substance sparingly soluble in cold water ; but its solution exhibits 
no properties differing from those of a solution of orthoarsenic 
acid, and it appears, therefore, to combine with water to form the 
latter body. The metarsenates, too, are only known as solids; 
they may be obtained from the appropriate hydrogen or ammonium 
orthoarsenates, e.g., HNaNH 4 AsO 4 = NaAsO 3 + H 2 + NH*-, 
but on treatment with water they combine, forming dihydrogen 
metallic ortho-arsenates. 

The metathioarsenates are produced by the action of alcohol 
on solutions of the pyrothioarsenates, thus : 

K 4 As 2 S 7 .Aq + Ale = K 3 AsS 4 + KAsS 3 .Aq.Alc. 

The orthosulpharsenate is precipitated, while the metasulph- 
arsenate remains in solution. The acid is unknown. They have 
been little investigated. 

Metantimonic acid, HSbO 3 , results from the spontaneous 
decomposition 'of H 4 Sb 2 07 dissolved in water; it is also produced 
when ihe pyro-acid is heated, or when a metantimonate is treated 
with an acid. It is also formed by the action of nitric acid on 
antimony. It is a soft white sparingly soluble powder. This 
compound and its salts are usually inconsistently named " anti- 
monic acid " and " antimonates." Hydrated pentoxide of bismuth, 
Bi 2 O 5 .H 2 O (see p. 350) may be classed here. 

(a.) Hexametaphosphates. These are the salts prepared by 
the usual methods from ordinary metaphosphoric acid : Na 6 P 6 O l8 ; 
(NH 4 ) 6 P 6 O 18 ; Na-aCa^PeO^; Ag 6 P 6 O I8 ; and others. 

The sodium salt is produced by strongly igniting dihydrogen 
sodium orthophosphate until it fuses, and then rapidly cooling the 
fused mass. It is an amorphous colourless deliquescent glass, 
easily soluble in water and in alcohol. It gives gelatinous preci- 


pitates with salts of most metals; its hexa-basic character is 
deduced from the formulae of double salts such as the one given 
above, Na2Ca 6 P 6 Oi 8 . The ammonium salt is prod uced by saturating 
orcjgnary metaphosphoric acid with ammonia, and evaporating. 

(6.) Tetrametaphosphates. Lead oxide, heated with excess 
of phosphoric acid, yields large transparent prisms of an insoluble 
salt. The salt is powdered, and digested with sodium sulphide ; 
lead sulphide and sodium tetrametaphosphate are formed. It is 
diluted with much water, and filtered. On adding alcohol, an 
elastic ropy mass, like caoutchouc, is precipitated. Its solution in 
water gives ropy precipitates with salts of other metals. Its 
tetra-basicity is inferred from the existence of double salts snch as 
Na 2 Cu"P 4 O 12 . 

(c.) Trimetaphosphates. When a considerable mass of 
sodium metaphosphate is slowly cooled, the mass acquires a 
beautiful crystalline structure ; and on treatment with warm 
water the solution separates into two layers, the larger stratum, 
containing the crystalline, and the smaller the ordinary vitreous, 
salt. The solution of the crystalline variety gives crystalline 
precipitates with salts of many metals, the silver salt, for example, 
depositing in crystals of the formula Ag^O^.H^O. The sodium 
salt deposits in large crystals of the formula Na 3 P<O 9 .6H 2 O. Its 
tri-basicity is inferred from formula such as 2NaBaP 3 O 9 .H u O. 
The salts of this acid uniformly crystallise well. 

(d.) Dimetaphosphates. By heating copper oxide, CuO, 
with a slight excess of phosphoric acid to 350, an insoluble 
crystalline powder is formed. On digestion with sulphides of 
sodium, potassium, &c., the corresponding dimetaphosphates are 
formed, and separate in crystals on addition of alcohol. Double 
salts are produced by mixture, such as NaNH 4 P 2 O h .H 2 O ; 
NaKP 2 O 6 .H 2 O; NaAgP 2 O 6 , &c. These salts, like the trimeta- 
phosphates, are crystalline bodies sparingly soluble in water. 

(e.) Monometaphosphates. These bodies are insoluble in 
water. They are produced by igniting together the oxides and 
phosphoric acid in molecular proportions ; or, by adding excess of 
phosphoric acid to solutions of nitrates or sulphates, evaporating, 
and heating the residues to 350 or upwards. They are crystalline 
and anhydrous, and form no double salts ; even the salts of the 
alkalies are nearly insoluble in water. The solution of the potassium 
salt in acetic acid gives precipitates with salts of barium, lead, and 

Metantimonates. These salts are produced by fusing anti- 
mony or its trioxide with nitrates, or the acid HSbO 3 with 

2 B 2 


carbonates ; or by double decomposition from the potassium salt, 
KSb0 3 .Aq. The chief compounds are : 

LiSb0 3 ; 2NaSbO 3 .7H 2 O ; NaSbO 3 .8H 2 O ; KSb0 3 ; also 2XSb0 3 .5H 2 O and 
SHsO; NH 4 Sb0 3 .2H20 ; Oa(SbO 3 ) 2 ; Sr(SbO 8 ) 2 .(H 2 O; Ba(SbO 3 ) 2 .5H J O ; 
Mff(Sb0 3 ) 2 .12H2O j Zn(Sb0 3 ) 2 ; Co(SbO 3 ) 2 ; Ni(SbO 3 ) 2 .6H 2 O ; 
8n(SbO 3 ) 2 .2H 2 O ; Pb(SbO 8 ) 2 ; Cu(SbO 3 ) 2 ; Hgr(SbO 3 ) 2 . 

All these salts, with exception of the lithium, sodium, 
potassium, and ammonium salts, are sparingly soluble in water, and 
crystalline. The compounds 2NaSbO 3 .7H 2 O and 2KSbO 3 .5 and 
3H 2 O are gummy, and may possibly be derived from a poly- 
metantimonic acid. When boiled with water they are decomposed, 
giving a residue of 3Sb 2 O 6 .K 2 O.10H 2 O. 

" Naples yellow " is a basic antimonate of lead, produced by 
heating 2 parts of lead nitrate, 1 part of tartar- emetic, and 
4 parts of common salt to such a temperature that the salt fuses ; 
the mass is then treated with water, which dissolves the salt, 
leaving the " Naples yellow " in the form of a fine yellow powder. 
Another basic antimonate of lead occurs native as bleinerite, 
Sb 2 8 ,3Pb0.4H 2 O. 

Certain complex phosphates have been prepared by fusing 
tetrasodium pyrophosphate with metaphosphate in the proportion 
Na 4 P 2 7 to 2NaP0 3 . The product is soluble without decomposi- 
tion in a small quantity of hot water, and crystallises from the 
solution ; but it is decomposed by much water. With solutions of 
salts of the metals, it gives precipitates ; the silver salt, for example, 
has the formula Ag 6 P 4 O l3 . Another salt has been produced by 
fusing together the same constituents in the proportion Na 4 P 2 7 
to 8NaP0 3 . The resulting salt is very sparingly soluble ; the silver 
salt derived from it has the formula Agi 2 P 10 O 31 . These phosphates 
go by the name of Fleitmann and Henneberg, their discoverers. 




Oxides, Sulphides, Selenides, and Tellurides of 
Phosphorus, Arsenic, Antimony, and Bismuth, 

Compounds of tetroxides. It has been already stated that 
the oxide P 2 O 4 , when treated with water, gives a mixture of phos- 
phorous and phosphoric acids, thus : P 2 O 4 + 3H 2 + Aq = 
H 3 P0 4 . Aq -f HjPOa.Aq. It is therefore concluded to be a phosphite 
of phosphoryl, thus : (PO)'"(PO 3 ) . But a tetrabasic acid is known, 
of the formula P 2 O 4 .2H 2 O = P 2 O 2 (OH) 4 , which forms distinct, 
salts, and possesses properties differing from those of such a mix- 
ture. The sodium salt, P 2 O 2 (ONa) 4 , is converted by bromine and 
water into dihydrogen disodinm pyrophosphate, and, as the acid 
has no marked reducing properties, it may possibly have the con- 

0=P(OH) 2 0=P(OH) 2 

I , that of pyrophosphoric acid being ]>0 

0=P(OH) 2 0=P(OH) a . 

Hypophosphoric acid,* as the acid P 2 O 2 (OH) 4 is called, is 
produced along with orthophosphoric and phosphorous acids, by 
the oxidation of phosphorus exposed to water and air. About one- 
sixteenth of the phosphorus is converted into hypophosphoric acid. 
On addition of sodium carbonate, the dihydrogen disodium salt 
separates out, owing to its sparing solubility in water. To prepare 
the pure acid, the barium salt is treated with the theoretical 

* Annalen, 87, 322 ; 104, 23 ; BericUe, 16, 749 j Comptes rend., 101, 1058 ; 
108, 110. 


amount of sulphuric acid ; insoluble barium sulphate is formed, 
and the acid remains in solution. On evaporation in a vacuum, 
the acid H 4 P 2 O 6 .2H 2 O separates out in large rectangular tables, 
melting at about 62. On standing in a dry vacuum, tfeese 
crystals lose water, and gradually change to needles of the pure 
aoid H 4 P 2 O 6 . This body, at 70, suddenly decomposes into phos- 
phorous and metaphosphoric acids : H 4 P 2 O 6 = H 3 PO 3 + HPO 3 . 
The following salts are known : 

Na 4 P 2 O 6 .10H 2 O ; K 4 P 2 O 6 .5H 2 O ; (NH 4 ) 4 P 2 O 6 .H 2 O ; Mgr 2 P 2 O 6 .12H 2 O ; 
Ca 2 P 2 O 6 .2H 2 O ; Ba 2 P 2 O 6 j Pb^Oe ; and Ag- 4 P 2 O 6 ; and the double salts 
H 3 NaP 2 6 .2H 2 ; H^a^O,,; HNa 3 P 2 O 6 .9H 2 ; H 3 Na 6 (P 2 O 6 ) 2 .20H 2 O ; 
H 3 KP 2 6 j H 2 K 2 P 2 6 3H 2 j HK,P 2 O 6 .3H 2 O j H 3 (NH 4 )P 2 O 6 ; 
H 2 (NH 4 ) 2 P 2 6 ; H 2 M*P 2 6 .4H 2 ; H 2 CaP 2 O 6 .6H 2 O ; H 2 BaP 2 O 6 .2H 2 O. 

With lithium salts, sodium hypophosphate gives a white pre- 

The tetra-metallic salts of the alkalis are easily soluble in 
water ; the dihydrogen disodium salt is sparingly soluble, and is 
used to separate the acid from its mixture with phosphorous and 
orthophosphoric acid. The dibarium salt is produced by precipi- 
tation ; it is nearly insoluble in water, as are most of the other 
salts. When the salts are heated they give products of decomposi- 
tion of phosphorous acid (hydrogen phosphide and metaphosphate) 
and metaphosphate of the metal. 

The silver salt may be prepared directly by dissolving 6 grams 
of silver nitrate in 100 grams of nitric acid diluted with 100 grams 
of water, and while it is kept hot on a water-bath adding 
8 or 9 grams of phosphorus. The mixture must be cooled as soon 
as the violent evolution of gas ceases, and, on standing, tetrargentic 
hypophosphate crystallises out. The silver salt is not reduced to 
metallic silver on boiling, as is silver phosphite ; and the sodium 
salt does not reduce salts of mercury, gold, or platinum. 

No similar compounds of arsenic are known ; but antimony 
tetroxide, when fnsed with potassium hydroxide or carbonate, 
yields a mass from which cold water extracts excess of alkali ; 
the residue, dissolved in boiling water and evaporated to dryness 
gives a yellow non-crystalline mass which has the composition 
SbaO^.ILO. On treatment with hydrochloric acid, it is con- 
verted into 2Sb 2 O 4 .K 2 O ; and excess of acid liberates the com- 
pound Sb 2 O 4 .H 2 O. 

Compounds of trioxides and trisulphides : Constitution 
of the acids, hydroxides, and salts derived from the trioxides 
and trisulphides of phosphorus, arsenic, antimony, and bis- 


muth. It will be remembered that phosphoryl chloride, POC1 S , 
on treatment with water, yields orthophosphoric acid, PO(OH) 8 , 
and it may be supposed that phosphorus trichloride, PC1 3 , yields 
a |imilar acid, P(OH) 3 . Such an acid ought to be tribasic, like 
orthophosphoric acid, and should yield three double salts, e.g., 
P(OHXONa), P(OH)(ONa) 2 , and PO(0]Sra) 3 . But the last of 
these is formed only when the second is mixed with great excess of 
a strong solution of sodium hydroxide, and left for some time ; it 
is then thrown down on addition of alcohol. It appears not im- 
probable, therefore, that a change has taken place during this 


time, and that the compound == P < C(OTT\ has changed to 

PO(H) 3 . And it is also to be noticed that when water acts on phos- 
phorus trichloride, some orthophosporic acid and free phosphorus 
are formed; this might take place during the change of P(OH) 3 to 
its isomeride 0=P(OH) 2 H. Moreover, an acid is known, named 
ethyl-phosphinic acid (produced by the oxidation of the compound" 
ethyl -phosphine), analogous to hydrogen phosphide (see p. 532), 
which is certainly dibasic, and in which the phosphorus is doubt- 
less in direct union with carbon. The formulas are : 


H X C 2 H 6 C 2 H 5 

Hydrogen phosphide. Ethyl pliosphine. Eth.yl-pliosph.mic acid. 

There are therefore good reasons for believing that, although 
two phosphorous acids might exist, the one known is O=P(OH) 2 H, 
and not P(OH)3. The isomerism is analogous to that of the two 
nitrous acids (see p. 337), O=N OH, and 2 =N H. The anhy- 
dride of the acid O=P(OH) 2 H would be therefore not P 2 O 3 , but 
2 PH, an unknown substance. As with orthophosphoric acid pyro- 
phosphates are known, so pyrophosphites exist, e.g., Na 4 P 2 O s . 
Such substances also find representatives among the arsenites and 
thioarsenites, all these series of salts being known, viz., MAsO 2 , 
and MAsS 2 , metarsenites and thioarsenites ; M 4 As 2 O ft and M 4 As 2 S ft , 
pyroarsenites and thioarsenites ; and M 3 As0 3 and M 3 AsS 3 , ortho- 
arbenites and thioarsenites. The corresponding metaphosphites are 
unknown. A few antimonites and sulphantimonites have also been 

Phosphorous acid, etc. 

H 3 PO 3 j H 4 Sb 2 O 5 = Sb 2 O 3 .2H 2 O ; H a SbO 3 = Sb 2 O 3 .3H 2 O. 

To prepare crystalline phosphorous acid, H 3 P(>3, a current of 
dry air is passed through phosphorus trichloride heated to 60 and 


passed into water cooled to 0. When the water is saturated, the 
crystals which separate are washed with ice-cold water, and dried 
in a vacuum. It is also slowly formed by union of the anhydride 
with water; or along with orfchophosphoric acid by the action, of 
water on the tetroxide ; or along with phosphoric and hypophos- 
phoric acids by the oxidation of phosphorus in air, in contact 
with water. Phosphorus also abstracts oxygen from a solution of 
copper sulphate, depositing copper, thus : 3CuS0 4 .Aq + 6H 2 O 
+ 2P = 3H 2 S0 4 .Aq + 2H 3 PO 3 .Aq + 3Cu. The sulphuric acid 
may be removed as barium sulphate by cautious addition of solu- 
tion of barium hydroxide. 

Pyroantimonious acid, H 4 Sb 2 O 6 , is produced by addition of 
copper sulphate to a solution of antimony trisulphide in caustic 
potash. Copper sulphide is formed, and potassium antimonite ; 
and on addition of an acid to the filtered liquid, the antimonite is 
decomposed, pyroantimonious acid being precipitated. 

Orthoantimonious acid is formed by the spontaneous de- 
composition of the peculiar compound acid of which tartar- emetic 
is the potassium salt. This acid is liberated from the barium salt 
corresponding to tartar- emetic, by the action of sulphuric acid, and 
has the formula (C 4 H 4 O 6 )"Sb.OH. With water it yields Sb(OH) 3 , 
and tartaric acid, C 4 H 6 O a . Aq. From this it would appear that tartar- 
emetic is not, as hitherto supposed, a tartrate of potassium and anti- 
monyl, K(SbO)C 4 H 4 O 6 , but a tartaro-antimonite (C 4 H 4 6 )"Sb.OK, 
two hydroxyl groups of antimonious acid, Sb(OH) 3 , being replaced 
by the dyad group (C 4 H 4 6 ). 

Phosphorous acid forms deliquescent white crystals, melting 
at 74. When heated it decomposes into hydrogen phosphide and 
phosphate : 

4H 3 PO a = 3H 3 PO 4 

Zinc and iron dissolve in it, and the liberated gas is hydrogen 
phosphide ; this action is somewhat similar to that of nitric acid 
on certain metals, whereby ammonia is produced. It is a powerful 
reducing agent, tending to combine with oxygen to form ortho- 
phosphoric acid ; hence, when added to solutions of salts of silver, 
gold, and mercury, the metals are deposited. It also reduces 
sulphurous acid to hydrogen sulphide, thus : 

3H 3 P0 3 .Aq + H 2 S0 3 .Aq = 3H 2 P0 4 .Aq + E Z 8. 

The antimonious acids are white powders, insoluble in water, 
but soluble in hot solutions of hydroxides of sodium and potassium, 


forming antimonites. The corresponding hydroxides of bismuth 
have no acid properties. The three hydrates, Bi(OH) 3 , Bi 2 O(OH) 4 , 
and BiO(OH), are all known. They are produced by heating 
solutions of bismuth salts with potash or ammonia. 

Phosphites. NagPOg is the only trimetallic phosphite known. 
It is produced by addition of a large excess of a strong solution of 
sodium hydroxide to disodium phosphite, HNa 2 PO 3 .Aq, and after 
two hours adding alcohol. The trisodium salt settles down as a 
viscid syrup, which is stirred with alcohol, and finally dried in a 
vacuum over sulphuric acid. 

; 2HNa(HPO 3 ).5H 2 O ; 
HK(HP0 3 ) ; and 2H 4 Na2(HP03)3H 2 ; and H 4 K2(HPO 8 ) 3 . 

These bodies form soluble crystals, and are produced by addition 
of phosphorous acid to hydroxides or carbonates. 

Ca(HP0 3 ).H 2 j 2Sr(HP0 3 ).H 2 ; 2Ba(HPO 3 VH 2 O j 
H,Ba (HPO 3 ) 2 .H 2 O. 

White sparingly soluble salts. 

j Od(HP0 3 ) (?) ; 2Zn(HPO 3 ).5H 2 O. 

These and an ammonium magnesium phosphite are produced by 
precipitation. They are white, crystalline, and sparingly soluble. 

Phosphites of aluminium, chromium, and iron have been pre- 
pared, but not analysed. They are sparingly soluble precipitates. 

Mn(HPO 2 ) j Co(HPO 2 ).2H 2 O, and Ni(HPO 3 ).3H 2 O. 

Coloured precipitates. 

Sn(HPO 3 ) and phosphites of tin dioxide and of titanium have 
also been prepared ; they are white precipitates. Pb(HPO 3 ) is also 
white, and is formed by precipitation. It is nearly insoluble. When 
digested with ammonia, the basic phosphate, P 2 O 3 .4PbO.2H 2 O, is 
produced. Bismuth phosphite is a white precipitate ; and copper 
phosphite, Cu(HPO 3 ).2H 2 O, forms sparingly soluble blue crystals ; 
when boiled, metallic copper is precipitated. 

All these phosphates decompose when heated, evolving -hydro- 
gen and a little hydrogen phosphide, and leaving a phosphate. 

It is stated that when the compound 2HNa(HPO 3 ).5H 2 O is 
heated to 160 it loses six molecules of water, forming a pyrophos- 
phite, NaJHgPaOfi.* Data concerning the phosphites are exceed- 
ingly meagre, and the whole series of salts requires reinvewtigation. 

* Comptes rend., 106, 1400. 


Some oxythiophosphites* have been prepared by the action of 
a solution of sodium hydroxide on phosphorus trisulphide (pre- 
sumably P 4 S 3 ). Hydrogen, mixed with hydrogen phosphide, is 
evolved, and on evaporation crystals are deposited of the comppsi- 
tion Na 4 P 2 O 3 S 2 .6H 2 O, analogous to a pyrophosphite. With sodium 
hydrosulphide, NaSH, hydrogen phosphide and sulphide are evolved, 
and the solution, evaporated in a vacuum, deposits crystals of 
Na 4 P 2 OS 4 .6H 2 O. These crystals lose hydrogen sulphide at the 
ordinary temperature, probably forming the salt previously men- 
tioned. With ammonium hydrosulphide, crystals of the formula 
(NH 4 ) 4 P 2 S 6 .3H 2 O are deposited, which, when dried at 100 in a 
current of hydrogen sulphide lose hydrogen sulphide, giving the 
compound (NH 4 ) 4 P 2 O 2 S 3 .2H 2 O. 

From the mother liquor of these crystals the compound 
(NH 4 ) 4 P 2 O 3 S 2 .2H 2 O, analogous to the potassium salt has been 
obtained. Solutions of these salts when boiled lose hydrogen sul- 
phide, and yield phosphites. 

Arsenites and thioarsenites. KAsO 2 and NH 4 AsO 2 are 
white soluble salts, produced by dissolving arsenious oxide (As 4 O 6 ) 
in caustic potash or ammonia. They are apparently metarsenites. 
By similarly treating arsenic trisulphide with potassium sulphide, 
either by solution or by fusion, the corresponding thioarsenite, 
KAsS 2 , is produced. It decomposes when treated with warm water. 

By adding alcohol to a solution of a large amount of arsenic 
trioxide in caustic potash, the pyroarsenite, H 3 KAS2O 5 , is produced. 
When digested with caustic potash, the salt K 4 As2O 6 is formed, and 
may be precipitated with alcohol. A similar ammonium salt is 
produced by direct addition, (NH 4 ) 4 AS2O 6 . The sodium salts are 
all very soluble, and have not been isolated. The corresponding 
pyrothioarsenites are unknown; but orthothioarsenites of potas- 
sium and ammonium, K 3 AsS 3 and (NH 4 )3AsS 3 , are precipitated on 
adding alcohol to a solution of arsenic trisulphide in excess of 
colourless ammonium sulphide. 

Oa(AsO 2 ) 2 j Ca^AsoO^ Ca 3 (AsO 3 ) 2 ; Sr(AsO 2 } 2 ; Ba(AsO 2 ) 2 ; 

H 4 Ba(As0 3 ) 2 . 

These are white sparingly soluble salts, produced by addition 
of arsenious oxide to the hydroxides, or arsenites of potassium or 
ammonium to salts of the metals. 

Corresponding to these are Ca 3 (AsS 3 )3.15H 2 O ; B 0^8285 ; and 
; they are soluble substances precipitated by alcohol. 

* Comptes rend., 93, 489 j 98, 45. 


Mg 3 (AsO 3 ) 2 ; MgHAsO 3 ; Mg 2 Aa 2 O 6 ; Mg 2 As2S 6 ; and 

are produced by double decomposition. 

Arsenites of the boron and alu minium groups have not been 

Various basic arsenites of iron are known. These are insoluble, 
and are produced by addition of a ferric salt and an alkali to 
solutions of arsenious oxide, and for this reason a mixture of ferric 
hydrate and magnesia is employed as an antidote in cases of arsenical 
poisoning. Among these are FeAsO 3 .Fe 2 O 3 ; 2FeAsO 3 .Pe 2 O 3 .7H 2 O, 
and 5H 2 O. 

Ferrous pyroarsenite, Fe 2 As2O 6 , is a greenish precipitate ; 
Mn 3 H 6 (AsO 3 ) 4 .H 2 O and Co a H 6 (AsO 3 ) 4 .H 2 O are rose-red precipi- 
tates ; the corresponding nickel salt, Ni.}H 6 (AsOa) 4 .H 2 O, is a 
greenish- white precipitate which yields Ni 3 (AsO 4 ) 2 on ignition. 
The sulpharsenites of these metals are all pyro-derivatives, viz., 
Pe 2 As 2 S 6 , Mn 2 As 2 S 6 , Co a As 2 S 6 , and Ni 2 As 2 S 6 . 

Stannous and stannic arsenites and sulpharsenites have been 
prepared, but not analysed. The three lead arsenites, Pb(AsO 2 ) 2 , 
Pb 2 As 2 O 5 , and Pb 3 (AsO 3 ) 2 , are all white precipitates. The com- 
pound Pb(AsS 2 ) 2 is a mineral named sartorite; Pb 2 As 2 S 6 is 
named dufrenoysite, and Pb 3 (AsS 3 ) 2 , guittermannite. All these are 
crystals with metallic lustre, and occur native. 

The arsenite of hydrogen and copper, HCuAsO 3 , is obtained 
by adding to a solution of copper sulphate a solution of potassium 
arsenite, a solution of arsenious oxide, and a small amount of 
ammonia. It is a fine green powder, and is named, from its dis- 
coverer, " Scheele's green." The arsenite Cu(AsO 2 ) 2 is produced 
by digesting copper carbonate with arsenious oxide and water. 

Copper sulpharsenite, CU2As 2 S 6 , is formed by precipitation; 
and some minerals exist which appear to be compounds of copper 
sulpharsenite and sulphide, e.c/., Cu 4 AsS 4 , julianite, Cu 6 As 4 S 9 , bin- 
mte, and Cu b ASoS 7 , tennantite. 

Silver arsenite, Ag 3 AsO 3 , is a yellow precipitate produced by 
adding to silver nitrate a solution of arsenious oxide in ammonia. 
It is soluble in excess of ammonia. It serves, along with Scheele's 
green, as a distinctive test between arsenious and arsenic oxides ; 
it will be remembered that copper arsenate is blue, and silver 
arsenate red. The corresponding sulpharsenite, Ag 3 AsS 3 , occurs 
native as proustite; and the mineral scanthoconate, Ag 9 A8 3 S 10 
appears to be a double sulpharsenite and sulpharsenate of silver. 

Only two antimonites are known, viz., NaSbO 2 .3H 2 O, which 
forms octahedra, and is obtained by dissolving antimonious oxide, 
(Sb 4 O 6 ) in caustic soda ; and an acid compound, NaSbO 2 .2HSbO 2 , 


similarly prepared. The corresponding thioantimonite,* 
NaSbS 2 , separates on addition of alcohol to a solution of Sb 2 S 3 in 
sodium hydroxide ; and copper-coloured crystals of 2NaSbS 2 .H 2 O 
deposit from a concentrated solution of the same substances. Many 
sulphantimonites occur native ; among them are Fe(SbS 2 ) 2 , 
berthierite ; Pb(SbS 2 ) 2 , zirikenite; Pb 2 Sb 2 S 6 , jamesonite; Pb 3 Sb 2 S 6 , 
boulangerite ; Pb 4 Sb 2 S 7 , meneghinite ; Pb 6 Sb 2 S 8 , geocronite; CuSbS 2 , 
chalcostibite ; Cl^Sh^, guejarite; CuPbSbS 3 , bournonite; Ag 3 SbS 3 , 
pyrargyrite ; AgSbS 2 , miargyrite ; Ag 5 SbS 4 , stephanite ; Ag 9 SbS 6 , 
polybasite ; and Hg(SbS 2 ) 2 , living stonite. Besides these, similar com- 
pounds of bismuth are known, e.g., AgBiS 2 , silver bismuth glance ; 
Pb(BiS 2 ) 2 , galenobismuihite ; Pb 2 Bi 2 S 6 , cosalite ; Pb fl Bi 2 S 9 , beeger- 
ite ; CuBiS 2 , emplectite ; Cu 3 BiS 3 , wittichenite ; and others. These 
double sulphides of bismuth have not been made artificially ; but 
the compound KBiS 2 , produced by fusing bismuth with sulphur 
and sodium carbonate, forms steel-grey shining needles. 

Hypophosphites.f Hydrogen hypophosphite, H 3 P0 2 , is a 
monobasic acid ; and it is therefore concluded that its constitution is 
somewhat analogous to that of phosphorous acid, inasmuch as it 
may be regarded as a hydroxyl-derivative of an oxidised hydrogen 
phosphide, thus, 0=P(OH)H 2 . It is only the hydrogen of the 
hydroxyl which can be replaced by metals. The anhydride of 
such an acid would not be the oxide P 2 O, but the unknown com- 
pound 0=PH 2 O PH 2 =O = H 4 P 2 3 . Such a body might be 
expected to be devoid of acid properties. 

Hypophosphorous acid, H 3 PO 2 , is produced by decomposing 
a solution of the barium salt, Ba(H 2 P0 2 ) 2 , with its equivalent of 
sulphuric acid. The dilute solution is boiled down, and finally 
evaporated at 105, the temperature being gradually raised to 130. 
It is then cooled to 0, and on shaking it crystallises. It melts at 
17*4. When heated, it decomposes into phosphoric acid and 
hydrogen phosphide, thus : 

2H 3 PO 2 = H 3 PO 4 + PH*. 

It yields salts on neutralisation with hydroxides or oxides. 
But sodium, potassium, and barium hypophosphites are easily pre- 
pared by boiling phosphorus with their hydroxides. The hydrogen 
phosphide which is evolved is spontaneously inflammable, owing 
to its containing a trace of liquid hydride, P 2 H 4 . The reaction is : 
4P + SKOH.Aq + 3H 2 = PH 3 + 3KH 2 P0 3 .Aq. 

* See also Ditte, Comptes rend., 102, 168, for pyrothioantimonites. 
f Rammelsberg, CKem. 8oc., 26, 1. 


It is from the barium salt, thus prepared, that the acid is 


LiH2P0 2 .H 2 j NaH 2 P0 2 .H 2 ; KH 2 PO 2 ; (NH 4 )H 2 PO 2 , 

These salts are white crystalline bodies, produced as described. 
Those containing water may be rendered anhydrous at 200. They 
decompose, when more strongly heated, as follows : 

5NaH 2 PO 2 = Na 4 P 2 O 7 + NaPO 3 + 2PH 3 + 2H Z . 

The ammonium salt undergoes a different change, thus : 
7NH 4 H.PO 2 = H 4 P 2 O 7 + 2HPO 3 + H Z + 7NH, 

Ca(H 2 P0 2 ) 2 j Sr(H2P0 2 ) 2 .H 2 0; 
White soluble salts. When heated they decompose, thus : 
7Sr(H 2 PO 2 ) 2 = 8Sr a P a O 7 + Sr(PO 3 ) 2 -f 6Pff 3 + H % + 4# 2 . 

Mg(H 2 P0 2 ) 2 .6H 2 ; Zn(H 2 P0 2 ) 2 6H 2 O ; and Cd(H 2 PO 2 ) 2 . 

These are also soluble crystalline salts, which can be dried at 
200. When heated they decompose, thus : 

5Zn(H 2 PO 2 ) 2 = 2Zn 2 P 2 O 7 4- Zn(PO 3 ) 2 + 4PJff 3 -f 4Bi. 

Aluminium and chromium hypophosphites are gumniy solids ; 
the ferric salt is a white sparingly soluble powder (?). 

Fe(H 2 P0 2 ).6H 2 Oj MnpEJPOjJs.HjO ; Co(H 2 PO 2 ) 2 .6H 2 O ; Ni(H 2 PO 2 ) 2 .6H 2 O. 

The ferrous salt has been prepared by dissolving iron in the 

acid ; the others by neutralisation. They can be dried at 200. They 

are all crystalline and soluble. They change thus, when heated : 

6Co(H 2 PO 2 ) 3 = 4Co(PO 3 ) 2 -f 2CoP -f 2PT 3 -f 9JGT 8 . 

Pb(H2PO 2 ) 2 is crystalline and sparingly soluble ; when heated, 
it decomposes, thus : 

9Pb(H 2 PO 2 ) 2 = 4Pb a P a O 7 -f Pb(PO 3 ) 2 -f 8Ptf 3 -f 2fl,0 -f 4ff a . 

Thallous hypophosphite, T1H 2 PO 2 , forms soluble white crystals. 
It decomposes like the sodium salt when heated. The uranyl salt, 
UO 2 (H 2 PO 2 ) 2 .H 2 O, is a sparingly soluble yellow crystalline salt. 

Like the phosphites, the hypophosphites possess great power 
of reduction ; the reaction, for example, with silver nitrate, is 
Ba(H 2 P0 2 ) 2 .Aq + 6AgN0 8 .Aq + 4H a O = 2H 8 P0 4 .Aq + 4HNO,.Aq 
-f Ba(N0 3 ).Aq + 6Ag+ H. The free hydrogen further reduces the 


nitric acid. With solutions of cupric salts, cuprous salts are first 
produced, and then a reddish precipitate of copper hydride, CuH, 
is formed. HypophosphorouR acid also withdraws oxygen from 
sulphur dioxide, liberating sulphur. 

Double compounds with halogens. With phosphorus, 
compounds of the type POC1 3 are best known. 

Arsenic forms only one compound of this nature, viz., 
AsOF 3 .KF.H 2 O ; its characteristic compound is AsOCl ; and 
antimony and bismuth resemble arsenic; the compound SbOCl 3 
is, however, known. 

S ; POC1 3 ; PSC1 3 ; POBr 3 ; PSBr a ; POCLjBr ; PSCl 2 Br. 
SbOCl 3 ; SbSCl 3 . 

These compounds have the formulae assigned, inasmuch as their 
vapour- densities have, in almost all cases, been determined. Their 
constitution is, without doubt, analogous to 0=PEEC1 3 ; and it will 
be remembered that when treated with water or alkalies they give 
rise to orthophosphoric or orthothiophosphoric acid. The corre- 
sponding antimony compound, SbOCl 3 , on treatment with water, 
yields the more stable pyrantimonic acid, H4Sb 2 O 7 .2H 2 O. 

POF 3 , phosphoryl trifluoride, is a colourless gas, liquid at 
50, or at 16 under a pressure of 15 atmospheres. By evapora- 
tion of the liquid a portion solidifies to a snow-like solid. It is 
produced by direct combination of phosphorus trifluoride and 
oxygen, which takes place with explosion on passing a spark 
through the mixed gases. It is more easily produced by distilling 
powdered cryolite, AlF.3NaF, with P 2 5 . 

PSF& the corresponding sulphur compound, is a spontaneously 
inflammable gas, liquefying under pressure of 10*3 atmospheres at 
13*8. It is best prepared by heating a mixture of phosphorus 
pentasulphide and lead fluoride, thus : 

P 2 S 6 + 5PbF 2 = 5PbS + 2PF 5 ; and 3PF 5 + P 2 S 6 = 

The density of this gas shows that, like the other similar com- 
pounds, it cannot be regarded as the compound ^PF 5 .P Z S^ but as 
the simpler body PSF^. 

POC1 3 , phosphoryl trichloride, is produced by the action of 
water on the pentachloride, thus : 

PC1 6 + H 2 = POC1 3 + 2HCI (see p. 353). 

But, as it quickly reacts with water, it is convenient to use eom- 
bined water, in the form of boracic acid, H 3 BO 3 , in its formation. 
It is easily obtained by distilling a mixture of phosphorus penta- 


chloride and boracic acid in theoretical proportions. It can also 
be produced by heating together the pentachloride and pentoxide 
of phosphorus, thus : 

3PC1 5 + P 2 O 5 = 5POC1 3 , 

It is a colourless liquid, heavier than water (1*7), boiling at 110. 
It fumes in the air, forming phosphoric and hydrochloric acids. 
It may be solidified, and melts at 2'5. It combines with some 
other chlorides, forming, for example : 

POC1 3 BC1 3 ; a white solid, melting at 11. 

POC1 3 A1C1 3 ; a white solid, melting and boiling without decomposition (?). 

POCl 3 .Mg:Cl 2 ; a white solid, decomposed when heated. 

POC1 3 ZnCL; white rhombic crystals. 

POC1 3 SnCl 4 ; a liquid, boiling at 180. 

It also forms gelatinous compounds with sodium and potassium 

PSC1 3 , sulphophosphoryl trichloride, is also a colourless 
fuming liquid, heavier than water, boiling at 124'25. It could 
doubtless be prepared by heating phosphorus pentachloride with 
pentasulphide ; but it is more readily obtained by the action of 
hydrogen sulphide on phosphoryl trichloride, thus : POC1 3 -f 
J9T 2 $ = PSC1 3 H- H 2 ; or by distilling phosphoric chloride with 
antimony trisulphide, thus : 

6PC1 5 + 5Sb. 2 S 3 = 3P 2 S 5 + 10SbCl 3 ; and 
3PC1 5 + P 2 S 5 = 5PSC1 3 . 

The easiest method of preparation is to distil phosphorus with 
sulphur chloride, S 2 C1 2 ; the reaction is : 

2P + 3S 3 Cl a = 48 + 2PSC1 3 . 

A compound of this body, with sulphur dichloride, PSC1 3 .SC1 2 , 
is produced by the action of sulphur on phosphoric pentachloride. 
It is a colourless liquid, boiling at 100. Its molecular weight has 
not been determined, 

POBr 3 and PSBr 3 are crystalline solids, similarly prepared. 
The former melts at 45 and boils at 195 ; the latter is yellow, and 
cannot be distilled without partial decomposition. POCl 2 Br and 
PSCl 2 Br are also known. 

Analogous compounds of arsenic are unknown ; but two com- 
pounds of arsenyl fluoride, of the formulae AsOP 3 .KF.H 2 O and 
AsOF 3 .AsF 6 .4KF.3H 2 O, have been prepared by treating arsenate 
of potassium with much hydrofluoric acid. They are colourless 
crystalline bodies. 


Antimonyl trichloride, SbO01 3 , is produced by the action 
of a trace of water on the pentachloride, SbCl 6 . Another state- 
ment is that this body has the formula SbCl 6 .H 2 ; it might well 
be SbOCl 3 .2HCl. It crystallises from chloroform. The corre- 
sponding compound SbSCl 3 forms white crystals ; it is produced 
by the action of hydrogen sulphide on the pentachloride. It is 
said to decompose when heated into SbCl 3 and S 2 C1 2 (?). 

Pyroplxosphoryl chloride, P 2 3 Cl4, corresponding to pyro- 
phosphoric acid, P 2 3 (OH) 4 , has been produced by the action of 
nitrogen tetroxide, N 2 C>4, on phosphorus trichloride. It is a colour- 
less liquid, boiling between 210 and 215. On treatment with 
water it yields orthophosphoric acid. Pyrosulphophosphoryl 
bromide is produced by dissolving phosphorus trisulphide, P 2 S 3 (a 
mixture ?), in carbon disulphide, and adding the necessary quantity 
of bromine. The carbon disulphide is distilled off, and the residue 
is extracted with ether, which dissolves the compound P 2 S 3 Br 4 . A 
light yellow oily liquid remains on evaporating the ether. When 
heated with phosphorus pentabromide, this substance yields the 
orthosulphophosphoryl bromide, PSBr 3 ; and when distilled aloue, 
the compound P 2 SBr 6 , which may be regarded as corresponding to 
the unknown thiodiphosphoric acid,P 2 S(OH) 6 , with hydroxyl re- 
placed by bromine (see p. 353). 

This substance is a white solid, melting at 5 to a yellow 

Metaphosphoryl chloride, PO 2 C1, is said to have been 
obtained by heating in a sealed tube for six hours a mixture of 
phosphorus pentoxide and phosphoryl trichloride, thus : 

P 2 O 6 + POC1 3 = 3P0 2 C1. 

It is a viscid colourless substance. The corresponding sulpho- 
phosphoryl bromide, PS 2 Br, is the insoluble residue after dissolving 
out the compound P 2 S 3 Br 4 with ether (see above). The analogous 
metantimonyl chloride, SbO 2 Cl, is produced by the action of much 
water on SbCl 6 . 

No chlorine derivatives of the oxide or sulphide, P 2 3 or P 2 S 3 , 
are known. But the bismuth haloid compounds and almost all 
those of arsenic and antimony are thus composed. 

Arsenosyl chloride, or arsenyl monochloride, AsOCl, is 
a hard white translucent fuming solid, formed by the action of a 
small amount of water on arsenic trichloride. It forms the follow- 
ing compounds : 

AsOCl.As 2 O 3 ; AsOCl.NH 4 Cl ; and AsOCl.H^O. 


The last of these may be viewed as orthoarsenious acid, with one 
hydroxyl group replaced by chlorine, thus : As(OH) 2 Cl. Arsen- 
osyl bromide, AsOBr, is a brown waxy solid similarly prepared. 
It forms the compound 2AsOBr.3H 2 O, which may perhaps be con- 
ceived as As 2 O(OH) 2 Br 2 .2H 2 O, a derivative of pyroarsenious acid, 
two hydroxyl groups being replaced by bromine. 

By similarly treating arsenic tri-iodide with water, the com- 
pound AsOI.As 4 O 6 crystallises out in thin plates. 

Sulpharsenosyl iodide, AsSI, is said to be formed by the action 
of iodine on arsenic trisulphide; and on addition of powdered 
arsenic to a solution of sulphur and bromine in carbon disulphide, 
the compound AsSBr.SBr 2 separates in dark-red crystals. 

Antimony trifluoride, SbP 3 , deliquesces on exposure to moist 
air, forming the compound 3SbOF.SbF a ; and bismuth oxyfluoiidj, 
BiOP, remains as a white powder on heating the crystalline com- 
pound B1OF.2HP, obtained by the action of concentrated hydro- 
fluoric acid on bismuthous oxide, Bi 2 O 3 , 

Antimonosyl chloride, SbOCl T and bismuth oxychloride, 
BiOCl, are produced by the action of water on the trichlorides 
SbCl 3 and BiCl 3 . The former is obtained in crystals by mixing 
10 parts of the trichloride with 17 of water, and, after, allowing to 
stand for some days, filtering, and washing .the precipitate with 
ether. The corresponding bismuth compound is used as a pigment 
and cosmetic under the name " pearl-white." When heated in air, 
it changes, giving the body 3BiOC1.2Bi 2 O 3 . Many other com- 
pounds are produced by the action of water on antimony tri- 
chloride ; among these are 

SbOC1.7SbCl 3 ; 2SbOCl Sb 2 O 3 ; 20SbOC1.10Sb 2 O 3 SbCl 8 , &c. 
These bodies all dissolve in concentrated hydrochloric acid, 
giving the trichloride. 

Similar bromides are known, similarly prepared ; for instance 

2SbOBr.Sb 2 O 3 ; 20SbOBr.lOSb 2 O 3 SbBr 3 ; BiOBr ; 7BiOBr.2Bi 2 O 8 . 
SbOI i 2SbOI Sb 2 O 3 ; BiOI ; and 3BiOI.4Bi 2 O 3 . 

Compounds of sulphantimonosyl chloride are also known, pro- 
duced by the action of the trichloride on trisulphide of antimony. 
Crystals of SbSCl.SbCl 3 are produced, and on washing them with 
alcohol, 2SbSC1.3Sb 2 S 3 remains. 

Sulphantimonosyl iodide, SbSI, is the product of the action of 
antimony tri-iodide on the trisulphidp ; it is a brown-red powder ; 
when boiled with zinc oxide and water, the oxy sulphide, Sb t OS 2 , is 

2 c 


The compounds BiSCl and BiSI are similarly produced ; and 
a selenochloride, BiSeCl, is formed as steel-grey needle-shaped 
crystals on adding bismuth triselenide, Bi 2 Se 3 , to molten 

No halogen compounds derivable from or connected with h} po- 
phosphorous acid are known ; dry hydrogen iodide acts on that 
acid violently, producing phosphorous acid and phosphonium 
iodide (see p. 517), thus : 

3H 3 P0 2 + HI = 2H 3 P0 3 + PH 4 I. 

Physical Properties. 

Mass of 1 cubic centimetre 

P 2 5 , 2 387 ; As 4 O 6 (amorphous), 374 ; (crystalline) 370. 
As 2 O 5 , 40; Sb 4 O 6 (octahedral), 5 11 ; (prismatic) 5 72. 
Sb 2 6 , 3-78; Sb 2 4 , 4'07 ; Bi^g, 5'1 ; Bi 2 O 4 , 5'6 ; Bi 2 3 , 8'08. 
P 4 S :J , 2 00 ; As 2 S 2 , 3 55 ; As 2 S 3 , 3'45 ; Sb 2 S 3 , 4'22 ; (stibnite) 4'6. 

Bi 2 S 3 , 7-00. 

As 2 Se 3 , 4'75 ; Bi*8e d , 6'25 ; Sb 2 Te 3 , 6 5 j Bi^, 7'23 7'94. 
POC1 3 , 1-711 at 0: P 2 O 3 C1 4 , 1-58 at 7; Sb 4 O 6 01 2 , 5'0; BiOCl, 7'2. 
POBr 3 , 2'82; PSBr 3 , 2'87; AsSBra, 279 ; BiOBr, 670 at 20. 

Heats of formation 

P P - 1-51K. 2P + 50 = P 2 5 + 3700K + Aq = 2H 3 PO 4 .Aq + 


2P + 30 -I- Aq = 2H 3 PO 8 .Aq + 2 x 1252K. 
2P + 4- Aq = 2H 3 PO 2 .Aq + 2 x 373K. 

P + + 3CI = POOL, 4- 1460K ; P + + 3Br = POBr 3 -f- 1056K 
2 As -f5O = As 2 O 5 + 2194K; + Aq = 60K. 
2As + 3O = As 2 O 3 + 1547K; + Aq =-70K. 
2Sb 4- 50 + 3H 2 O = 2H 3 Sb0 4 + 2 x 1144K. 
2Sb + 30 = Sb 2 3 + 1660K. Sb + + Cl = SbOCl + 897K. 
2Bi . 30 -r 3H 2 O = 2H3BiO 3 4- 2 x 691K ; Bi + O + Cl - 
BiOOl + 882K. 







Ozone. It has long been known that oxygen through which 
electric sparks have been passed acquires a peculiar smell, and acts 
rapidly on mercury. This behaviour is due to the conversion of 
the oxygen into an allotropic modification, to which the name 
ozone (from o'fe*>, to smell) has been given.* In this instance 
the molecular weight is known, and consequently the formula of 
ozone, 3 ; and it appears advisable, therefore, to regard it as an 
oxide of oxygen. It is true that ordinary oxygen, which possesses 
the molecular formula 2 , might also thus be regarded ; but, inas- 
much as ozone is the only allotropic modification of an element 
(except perhaps sulphur gas at a low temperature, which may 
possess the formula $ 8 ) which is fairly stable and possesses a known 
molecular weight at the ordinary temperature, it has been given a 
prominent position. 

Sources. Ozone occurs in small amount in the atmosphere, 
especially of the country. It may be recognised by its power of 
turning red litmus paper soaked in a solution of potassium iodide 
blue, owing to the liberation of potassium hydroxide (see below). 
Country air contains at most about one seven-hundred-thousandth 
of its volume of ozone. It appears to contain more ozone in spring 
than in summer, and more in summer than in autumn or winter ; 
and it is -more abundant on rainy than on fine days. Its presence 
appears also to be favoured (in the northern hemisphere) by west 
or south-west winds ; and its existence has been shown to be 
largely dependent on the prevalence of atmospheric electricity, for 
its amount is greatly increased during and after thunderstorms. 
Its presence in country air in greater amount than in town air may 

* Schdnbein, Pogg. Ann.> 50, 616 j Andrew.s, Chem. Soc. J. t 9, 168. 

2 C 2 


be due to the fact that the oxygen evolved from plants contains 
small traces of ozone, and that in the neighbourhood of towns the 
ozone is destroyed by its action on organic particles, and on the 
sulphurous acid produced during the combustion of coal. 

Preparation. Ozone is formed, as mentioned above, by the 
passage of electric sparks through oxygen ; and it is also pro- 
duced during the oxidation by free oxygen of various substances, 
such as phosphorus in contact with water, ether vapour, benzene 
and other hydrocarbons, and also by the combustion of hydrogen ; 
by the action of sulphuric acid on barium dioxide, potassium per- 
manganate, and other substances which evolve oxygen in the cold 
on being thus treated; and, lastly, by the electrolysis of dilute 
sulphuric acid. It is never obtained pure. By the first process a 
quarter of the oxygen present has been converted into ozone, but 
by the other processes a much smaller proportion undergoes 

These processes of formation may be illustrated as follows : 

1. By slow oxidation. (a.) A few sticks of phosphorus are placed in a 
large bottle and partly covered with water. After standing for about an hour, 
the air in the bottle is aspirated through a U't UDe containing a solution of 
potassium iodide mixed with a little boiled starch. The solution will turn blue 
owing to the liberation of iodine and the formation of blue iodide of staich. 
(b.) A few drops of ether are poured into a large dry beaker, covered with a 
plate, and the beaker is shaken so as to mix the ether vapour with the air. The 
gaseous mixture is then stirred with a glass rod heated over a flame till too hot 
to touch. On pouring a little solution of potassium iodide and starch into the 
beaker and shaking, a blue colour will be produced. It has been observed that 
this reaction is also shown by the air in a bottle containing a little petroleum, 
frequently opened and shaken, especially if it has been exposed to sunshine for a 
few da) s. 

2. During: combustion. If a small jet of hydrogen be burned below a 
funnel, and the products be drawn through a solution of potassium iodide and 
starch, the blue colour is produced ; but, besides ozone, hydrogen peroxide and 
ammonium nitrite are produced, both of which have the property of liberating 
iodine from such a solution. The method of distinguishing these bodies from 
ozone is described below. 

3. Dilute sulphuric acid, electrolysed by eight Grove cells, each electrode 
consisting of six thin platinum wires, yields oxygen rich in ozone. In one ex- 
periment at 9, about one quarter of the oxygen collected was in the form of 
ozone. Persu'phuric acid is also formed in the liquid by this process. 

4. The most convenient method of producing ozone is ty the passage of the 
"silent discharge" througrh oxygen. This " silent discharge " appears to 
consist of a rain of small sparks, and is best produced between two surfaces of 
glass placed very near each other, with conducting coatings on their exterior 
surfaces. A Kuhmkorff coil or an electrical machine may be used as a source 
of the electricity of the high potential required. The apparatus, of which a 



out is given in fig. 41, serves well for the purpose, and by its means the rela- 
tion between the alteration in volume of the oxygen during conversion into 
ozone and the volume of the ozone produced may be shown. It consists of a 

FIG. 41. 

wide glass tube, standing on a foot, and constricted about 2 inches from its 
Lower end. At its lower end a paraffined cork, or a glass stopper lubricated with 
vaseline, is inserted in the tubulus, and on the opposite side to the tubulus a 
vertical tube, provided with a stopcock, ending in a (J-tube, is sealed on. A 
narrower tube, which should fit the wider one very closely, but without touching, 
is sealed through its upper end. At the top of the wide tube a gauge, like 
k hat shown in the figure, is attached by sealing. The outer tube is covered with 
tinfoil where it surrounds the inner tube. 

As ozone is destroyed by grease, the stopcock should be lubricated with vase- 
line, and no india-rubber connections should be placed in contact with ozone, 
for it at once attacks india-rubber. 

To illustrate the formation of ozone with this apparatus, a slow current of 
oxygen is passed through the tube, entering through the gauge-tube, which should 
contain no liquid. A platinum wire connected with one pole of a coil is dipped 
in dilute sulphuric acid contained in the inner tube, while the outer coating 
>f tinfoil is connected with the other pole of the coil. The U'^be having been 
illed with a solution of potassium iodide and starch, a blue colour is produced 
is soon as the current passes. The characteristic smell may be noticed before 
;he solution is poured into the (J-tube. 

Properties. A.t ordinary pressure and temperature, ozone is a 
Sfas, colourless in thin layers; but by looking through a tube 
several metres in length, filled with ozonised oxygen, it is seen to 
Liave a blue colour. When compressed, this blue colour becomes 
more apparent, and at low temperatures it increases in intensity. 
A. current of ozonised oxygen, cooled to 180 by liquid oxygen 
boiling at atmospheric pressure, deposits its ozone as a dark-t^ue 
liquid, while the oxygen passes on. The blue liquid boils at 

* Comptes rend,, 94, 1249 ; Monatshcft Chem., 8, 69. 


When heated to 250 300, ozone is reconverted into ordinary 
oxygen, and the volume of the gas is found to have permanently 
increased. Ozone liberates iodine from a solution of potassium 
iodide, forming potassium hydroxide and free iodine ; it oxid'ses 
silver and mercury, which are unaffected by ordinary oxygen at 
the atmospheric temperature, and at once converts black lead 
sulphide into white lead sulphate. It reacts with hydrogen per- 
oxide, but only slowly in presence of free acid ; the action is rapid 
in presence of an alkali ; oxygen gas is evolved. It bleaches 
indigo and other vegetable colouring matters. It is very sparingly 
soluble in water, although nearly ten times more soluble than 
oxygen. It provokes coughing, irritating the bronchial tubes. It 
is poisonous when breathed in a concentrated form ; and, curiously, 
the blood of animals killed by it is found to have the dark colour 
of venous blood ; death appears to be produced by asphyxia. 

Proof of the formula of ozone. That ozone has the formula 
O 3 is rendered probable by the following experiments : 1. Oxygen, 
when ozonised, undergoes contraction. This may be proved by 
placing the apparatus shown in fig. 41, filled with oxygen, in 
water so as to maintain a constant temperature, for on passing the 
discharge the gas would become heated, and changes of volume, 
not dependent on the conversion of oxgyen into ozone, would then 
occur. Some strong sulphuric acid, coloured with indigo, is intro- 
duced into the gauge, and the stopcock connecting the apparatus 
with the (J-tube is shut. On passing a currenc, a momentary 
expansion will take place at first, due for the most part to heating 
of the gas ; it is, however, followed by a contraction shown by the 
rise of the liquid in the gauge. Its level is observed. 

2. The apparatus is then removed from the water, and by a 
rapid shake, a small thin bulb filled with oil of turpentine contained 
in the lower part of the tube is broken. The apparatus is again 
placed in water, and -allowed to stand for a few minutes, so as to 
regain its original temperature. A further contraction will have 
taken place, amounting to. twice that originally observed. The 
(J-tube is washed out and filled with fresh iodide of potassium 
and starch ; the contents of the apparatus are then expelled 
through the U^ QDe by a current of oxygen ; no coloration is pro- 
duced, showing that the ozone has been completely removed. 

The relation of the volume of the ozone to that of the oxygen 
from which it has been produced, can be inferred from these experi- 
ments. To take a suppositious case : Suppose the total volume 
of the oxygen before electrification is 100 c.c. After partial con- 
version iiito ozone, the volume may be imagined to be reduced to 

OZONE. 391 

99 c.c. ; and after absorption of the ozone by turpentine, the con- 
traction is twice as great as the original one, and the volume is 
further reduced to 97 c.c. We have thus: 

Volume of original gas 100 c.c. 

Volume of oxygen plus ozone 99 c.c. 

Volume of oxygen after removal of ozone 97 c.c. 

Hence oxygen converted into ozone 100 97 c.c. = 3 c.c. 

Volume of that ozone 9.9 97 c.c. = 2 c.c. 

We see, therefore, that three volumes of oxygen are converted into 
two volumes of ozone. 

The density of ozone should,, therefore, be that of 3 /2 = 24. 

No direct experiments on its- density have been made ; but 
corroborative evidence is furnished by experiments in which the 
rate of diffusion of ozone mixed with oxygen was compared with 
that of chlorine mixed with oxygen.* The rate of diffusion of tWo 
gases, as shown by Graham, is inversely as the square roots of 
their respective densities. Now, the density of chlorine is 35*5, 
the square root of which is 5*96 ; and that of ozone is presumably 
24, the square root of which is 4*90 ; hence, for every 4*90 grams 
of chlorine escaping into air by diffusion, 5'96 grams of ozone 
should escape. The ratio between, these numbers is 

5-96 : 4-90 : 100 : 82 2 ; 

the rate of diffusion should be, therefore, the 100/82 of that of 
chlorine ; it was experimentally found to be the 100/84th part, a 
sufficiently close approximation. A similar set of experiments 
proved it to have nearly the same rate of diffusion as carbon di- 
oxide, C0 2 , the density of which is 22. It may be regarded, there- 
fore, as a compound of three atoms of oxygen, and it possesses the 
formula 3 , with the molecular weight 48. 

Many substances, such as potassium iodide, mercury, and silver, 
convert ozone into ordinary oxygen, but remove only the third 
atom of the oxygen of ozone. The equations are : 

2KI Aq + 3 + H 2 = 2KOH.Aq + i 2 .Aq + O z ; 
2Ag + 3 = A&O + 2 ; and Hg + 3 = HgO -f- O a . 

Other bodies, such as the dry dioxides of lead and manganese, and 
copper oxide, decompose it without themselves suffering any per- 
manent change. And in some cases it has a reducing action ; for 
example, silver oxide is reduced to metallic silver, thus : Ag 2 O -f 
3 = 2Ag -f 20 2 ; moist dioxide of lead is also reduced : 

* Soret, Annalef (4), 7, 118 j 18, 2;7. 


PbO 2 -f 3 = PbO -f 20 2 . And in alkaline or neutral solution it 
reduces hydrogen dioxide, H 2 2 .Aq + 3 = H 2 O.Aq -f- 20 2 . It is 
probable thai/ the decomposing action which silver, mercury, &c., 
have on ozone is due to a double change, for instance : 2Ag -|- O 3 
=s Ag 2 O -h 2 , and Ag 2 O -f 3 = 2Ag -f 20 2 . These changes 
may be regarded as due to the action of atomic oxygen, a body 
incapable of more than a momentary existence at the ordinary 
temperature, but one which we should suspect to display great 
chemical activity. In this connection it may be noted that at the 
moment when tbe electric discharge begins to pass through pure 
Iry oxygen, a sudden expansion occurs, too sudden to be regarded 
as due to rise of temperature; an equally sudden contraction 
Bnsues. It may be supposed that the first action of the discharge 
is to partly dissociate the ordinary oxygen, 2 , into atoms, many 
Df which then combine in groups of three, forming ozone, 3 . 

Tests for Ozone. Ozone liberates iodine from potassium 
iodide, with formation of potassium hydroxide. If, therefore, half 
[>f a strip of red litmus paper be moistened with a solution of 
potassium iodide and starch, the moist portion will become blue, 
iwing to the liberated alkali. This effect is not produced by 
nitrous acid, hydrogen peroxide, chlorine, or other substances 
which have also the power of liberating iodine from potassium 

Another tost is the power which ozone possesses of oxidising 
i thallous salt to hydrated thallic oxide. Paper moistened with a 
solution of colourless thallous hydroxide therefore changes to the 
brown tint of thallic hydrate on exposure to ozonised oxygen. 

Oxides and Sulphides of Molybdenum, Tungsten, 
and Uranium. 

No selenides or tellurides of the elements of this group have 
been prepared. The following is a list of the oxides and sul- 

List. Molybdenum. Tungsten. Uranium. 

Qiygen Mo 2 O 3 ; Mo0 2 ; M60 3 . WO 2 ; W0 8 . T7O 2 ; TJO 3 j TT0 4 . 

Sulphur MoS 2 ; MoS 3 ; MoS 4 . WS 2 ; WS 3 . tJS 3 ; US 3 (?) * 

Besides these, the following oxides are known; they may be 
regarded as compounds of the simpler ones with each other: 

* Poisibly exist in combination with, other oxides or sulphides. 


Mo 2 O 5 (in combination with water) = MoO 2 .MoO 3 ; U 2 O 6 ; 
MoA* = MoO 2 .2MoO 3 ; U 3 O 8 = UO 2 .2UO 3 . 

Sources. Molybdenum and tungsten trioxides occur native 
&a*molybdic ochre and timgstic or wolfram ochre; the former often 
coats the surface of the native sulphide as an earthy powder; 
arid the latter forms a bright yellow or yellowish-green powder, 
sometimes occurring crystallised in cubes. The oxide TJ 3 O 8 is the 
chief constituent of. pitchblende, the chief source of uranium. It is 
a hard greyish, greenish, or reddish-black mineral, sometimes 
crystallising in regular octahedra. It usually accompanies lead 
and silver ores. Molybdenum, tungsten, and uranium also occur 
as trioxides, in combination with other oxides. Among such 
compounds are wulfenite, or yellow lead ore, lead molybdate, 
PbMoO 4 ; calcium tungstate, CaWO 4 , named scheelite or tungsten 
(from the Swedish words tung, heavy, and sten, stone ; its specific 
gravity is 6) ; ferrous-manganous tungstate, (Pe,Mn)WO 4 , or wol- 
fram, the chief ore of tungsten ; and scheeletine or lead tungstate, 
PbWO4. Uranium occurs as carbonate in liebigite, Ca(UO 2 )(CO 3 ) 2 ; 
as phosphate in uranium vitriol ; also in uranite, 

2(UO 2 ).Ca(P0 4 ) 2 .8H 2 O, 

and in chalcolite, in which copper replaces calcium. 

Uranium also occurs in the r&re miner&ls samarskite, fergusonite, 
pyrochlore, euxenile, &c., in combination with oxides of niobium, 
tantalum, yttrium, and other elements. 

Disulphide of molybdenum, MoS>, occurs native as molyb- 
denite, or molybdenum glance, in soft, grey, elastic, flexible 
laminae, resembling lead in colour and lustre, and graphite in its 

Preparation. 1. By direct union. Molybdenum and tungs- 
ten, when finely divided, burn when heated in air to the trioxides, 
MoO 3 and WO 3 ; uranium yields the oxide U 3 O 8 . The trioxide 
of molybdenum is also obtained by heating the metal in water- 
vapour, or with potassium hydroxide. The sulphides, MoS 3r 
WS 2 , and US 2 , are also produced directly, by heating the finely- 
divided metals with sulphur. 

2. By heating double compounds. The oxides Mo 2 O 3 , 
MoO 2 , MoO 3 , WO 3 , TJO 2 , UO 3 , and UO 4 are left anhydrous 
when their hydrates are heated. With Mo 2 O 3 and MoO 2 , air 
must be excluded, else oxidation occurs. The hydrate of the 
oxide UO a becomes anhydrous when boiled with water. To 
dehydrate UO 4 .2H 2 O, the temperature must not be allowed to rise 
much above 100, else loss of oxygen ensues. 


Molybdates, tungstates, and uranates of ammoniam and of 
mercury leave the trioxides when heated to redness, the volatile 
bases being expelled. Uranyl nitrate, UO 2 (NO3) 2 , at 250 yields 
the trioxide. When more strongly ignited, U 3 O 8 is produce^ ; 
and at an intense heat, U 2 O 5 . 

3. By reducing a higher oxide or sulphide. Molybdenum 
sesquioxide is produced by treating the trioxide with nascent 
hydrogen from zinc and hydrochloric acid. Molybdenum and 
tungsten dioxidee are produced when the trioxides are heated to 
low redness in hydrogen; at high temperatures the oxides are 
reduced to metal. Uranium dioxide is produced by heating the 
complex oxide U 3 O 8 to whiteness with carbon, or in a current of 
hydrogen. It has also been prepared by heating the oxalate, 
(CJO 2 )C 2 O 4 , or the double oxychloride, UO2Cl2.2KCl.2H2O, to 
redness in a current of hydrogen. Molybdenum dioxide is obtained 
in a crystalline form, by fusing sodium molybdate, NaJMtoO^ 
with metallic zinc, which deprives the trioxide, MoO a , of its 
oxygen. Uranium tetroxide, UO 4 , loses oxygen when heated to 
250, leaving the trioxide, and at higher temperatures gives the 
oxide U 3 O 6 , which loses more oxygen on intense ignition, leaving 
U*0 5 . 

The higher sulphides of molybdenum and tungsten, US 4 , US 3 , 
and WS 3 , likewise lose sulphur at a red heat, yielding the 

4. By oxidation of a lower oxide. The oxides MoO 3 , WO 3 , 
and U 3 O 8 are produced when the lower oxides are ignited in air. 
The higher sulphides, however, are not formed by heating the 
lower ones with sulphur. 

5. By replacement, or by double decomposition. The 
only oxide produced by this method is uranium tetroxide ; it is 
formed when a mixture of solutions of uranyl nitrate and 
hydrogen dioxide, in presence of a large excess of sulphuric acid, 
are allowed to stand fur some weeks, thus : 

UO 2 (NO 3 ) 2 .Aq + H 2 2 .Aq = UO 4 + 2HJST0 3 .Aq. 

This is, however, a method of preparing the sulphides. Molyb- 
denum or tungsten trioxide, heated with sulphur, yields sulphur 
dioxide, and the disulphide. Bisulphide of tungsten is also pro- 
duced by heating any oxide in a current of hydrogen sulphide or 
carbon disulphide ; and uranium disulphide has been obtained by 
heating uranium tetrachloride, UCli, in a current of hydrogen 


The higher sulphides are also prepared by this method. 
Molybdenum and tungsten trisulphides are formed by addition of 
hydrogen sulphide or ammonium sulphide to the solution of a 
nuolybdate or a tungstate, and subsequent addition of an acid. 
They are then precipitated. Uranium tri sulphide is produced 
by heating the tribromide in a current of hydrogen sulphide. 
Molybdenum tetrasulphide is precipitated on addition of an acid 
to a solution of a sulphopermolybdate, such as Na 2 MoS 6 .Aq. 

Properties. Molybdenum sesquioxide was believed by 
Berzelius to be the monoxide, MoO. It is a black powder, when 
obtained by igniting the hydrate; but when produced by the 
action of nascent hydrogen from zinc and hydrochloric acid on the 
trioxide, a dark brass-yellow precipitate. Molybdenum dioxide, 
from the trioxide with hydrogen, is a dark-brown powder ; when 
prepared from sodium molybdate by fusion with zinc, it forms 
blue-violpt prisms. Dioxide of tungsten forms brilliant copper- 
red plates, insoluble in water and acids ; and uranium dioxide 
also possesses metallic lustre ; prepared from the oxalate, it is a 
cinnamon-brown powder ; but when obtained from the double 
chloride, it crystallises in lustrous octahedra. The amorphous 
form glows when heated in air, burning to U 3 O 8 . 

Molybdenum trioxide forms a light porous white mass of 
silky scales ; it melts at a red heat to a dark-yellow liquid, which, 
when cooled slowly, solidifies in needles. It is volatile in a current 
of air, but not alone ; this is perhaps due to the transient forma- 
tion of a dissociable and more volatile higher oxide. It is insoluble 
in water, but dissolves, in acids. Trioxide of tungsten is a 
lemon- or sulphur-yellow powder, turning darker when heated. 
It may be obtained in transparent trimetric tables by crystallisa- 
tion from fused borax, or in octahedra by heating it in a current 
of hydrogen chloride. It melts at a bright-red heat. Uranium 
trioxide is a buff-coloured powder. Uranium tetroxide forms 
a heavy, white, crystalline precipitate. The more complex oxide, 
Mo 3 O 8 is a blue insoluble powder; U 3 O 8 , when prepared artificially, 
is a dark-green velvety powder ; and U 2 O 6 a black powder. When 
the glaze on porcelain is mixed with the oxide U 3 O 8 , and baked, 
an intense black colour is produced, and it is conjectured that 
during firing, the oxide U 3 O 8 is converted into U 2 O 5 . 

Molybdenum disulphide, prepared artificially, is a black 
lustrous powder; disulphide of tungsten forms slender black 
needles ; and that of uranium is a greyish -black amorphous body, 
becoming crystalline at a white heat. Molybdenum trisulpbide 
is a blackish -brown powder; and that of tungsten forms black 


lumps which yield a liver-coloured powder. Both of these bodies 
dissolve in solutions of sulphides of the alkalis ; that of tungsten 
is slightly soluble in water, but is precipitated on addition of 
ammonium chloride. It becomes denser when boiled witb hydro- 
chloric acid. Uranium trisulphide is a black powder. Molyb- 
denum tetrasulphide forms dark-red flocks, drying to a dark- 
green mass with metallic lustre ; when triturated, it gives a red 

Double compounds. (a.) With water. The hydrates 
are mostly prepared by double decomposition ; those of the 
trioxides, and of uranium tetroxide, may be teamed acids, inasmuch 
as they correspond in formula to numerous double oxides. Double 
oxides corresponding to the other oxides of these elements have 
not been prepared. 

Hydrated molybdenum sesquioxide is produced by adding 
potassium hydroxide to a solution of a molybdate previously 
exposed for some time to the action of nascent hydrogen from zinc 
and hydrochloric acid, or better, sodium amalgam. It is a black 
precipitate, soluble in acids, forming dark-coloured or purple 
molybdous salts ; in dilute solution they have a brownish-red 

Hydrated molybdenum dioxide is produced by adding 
ammonia solution to a solution of molybdenum tetrachloride. It 
is a rusty-coloured precipitate, sparingly soluble in water, in 
which it dissolves to a dark red solution ; hence it should be 
washed with alcohol. Its solution gelatinises on standing. It is 
soluble in acids, forming reddish-brown solutions. 

Hydrated dioxide of tungsten is unknown; that of uranium 
is precipitated in red-brown flocks from uranous salts by addition 
of an alkali. It is soluble in acids, forming green solutions. 

Molybdenum trioxide is sparingly soluble in water (I in 
500), By dialysing it, Graham prepared a stronger solution, 
possessing a yellow colour and an acid taste. This soluble modi- 
fication has also been prepared by addition of the theoretical 
amount of sulphuric acid to barium molybdate suspended in water, 
and filtering off the precipitated barium sulphate. On evaporation 
it forms a transparent blue-green mass, which slowly dries to the 
anhydrous oxide. The hydrate, or acid, 2MoO 3 .H 2 O = H 2 Mo 2 O 7 , 
has been prepared by drying the residue for two months over 
strong sulphuric acid ; and a hydrate, or acid, MoO 3 .H 2 O, once 
separated in prisms, after long standing, from a solution of the 
magnesium salt which had been mixed with nitric acid equivalent 
to the magnesium. 


The hydrated double oxide, Mo 2 O 6 .3H 2 O. forma a bine pre- 
cipitate on mixing a solution of the dioxide with a solution of the 
trioxide in hydrochloric acid. Tt may be regarded as MoO 2 .MoO 3 , 
tnolybdyl molybdate. Similarly, the oxide Mo 3 O 8 may bo viewed 
as MoO 2 .2MoO 3 , or molybdyl dimolybdate. 

Hydrated tungsten trioxide, or tungstic acid, WO 3 .H 2 O, 
2= H 2 WO 4 , forms a yellow precipitate on adding an acid to a hot 
solution of a tungstate. It crystallises without alteration of com- 
position from hydrofluoric acid. By similar treatment of a cold 
solution, a white gelatinous precipitate of WO 3 .2H 2 O is formed. 
It is also produced when water is added to a solution of tungsten 
chloride or oxychloride. 

A soluble modification of tungstic acid, named metatungstic 
acid, is produced by action of sulphuric acid on barium tungstato 
suspended in water. On evaporation, the solution deposits yellow 
crystals of 4(WO 3 .H 2 O).31H 2 O = H 2 W 4 O 13 .31H 2 O. This hydrate 
is easily soluble in water and forms soluble salts. On heating its 
concentrated solution, ordinary tungstic acid separates out. 

Hydrated uranium trioxide, UO 3 .2H 2 O, has not been ob- 
tained pure by precipitation, for alkali is always carried down. 
But by heating a weak alcoholic solution of the nitrate, oxidation 
products of alcohol are suddenly evolved ; the hydrated oxide re- 
mains as a buff-coloured mass. It is also formed by exposing 
moist U 3 O 8 to air. When dried in vacuo it forms UOj.H.0 = 
H 2 UO 4 , a lemon-yellow powder. 

The hydrated tetroxide, UO 4 .2H 2 O, is a yellow-white powder 
obtained by mixing a solution of a uranyl salt with hydrogen di- 
oxide. On treatment with potash, uranic hydrate, UO 3 .2H 2 O, is 
precipitated, and the potassium salt of an acid dissolves which 
may be conceived to have the formula H 8 TJOi . The hydrated 
tetroxide may therefore be viewed as UO 6 .2UO 3 .6H 2 O. Attempts 
to prepare the hydrated oxide UO 6 were, however, unsuccessful ; 
on addition of hydiogen dioxide to a nitric acid solution of uraninm 
nitrate, the ratio of uranium to oxygen in the precipitate corre- 
sponded approximately to the formula U 2 O 9 . 

(fc.) No hydrosulphides are known. 

(c.) Double oxides and sulphides; salts of molybdic, 
tungstic, and uranic acids ; also of corresponding sulpho- 
acids. A compound of tungsten dioxide and sodium oxide has 
been prepared, by dissolving in fused sodium tungstate as much 
trioxide as it will take up, and heating the mixture to redness in 
hydrogen. On treatment with water, the compound 2WO 2 .Na.jO 
a remains in golden-yellow scales and cubes possessing 


metallic lustre. It cannot be prepared by direct union of the 
oxides. No similar compounds of molybdenum or uranium are 

Molybdates, tungstatas, and uranates. These are among 
the most complex of compounds known. Owing to their com- 
plexity, the formulas of many are somewhat uncertain, different 
investigators drawing different conclusions from their analytical 
data. There can. be no doubt that, of all oxides, these show most 
tendency to polymerise, especially when in union with others. It 
will be convenient, in writing the formulae of these complex bodies, 
to represent them as compounds of oxides with oxides, e>.</., 
mM 2 0.r<R0 3 , where M stands for any element, and R, for molyb- 
denum, tungsten, or uranium. As nothing is known regarding 
the constitution of these bodies, the water frequently contained in 
them will be written separately. 

5(Mo0 3 .Li20).3H 2 0; 3(Mo0 3 .Li2O).8H 2 0; Mo0 3 .Na 2 O.H !? 0, and 2H 2 ; 
2(Mo0 3 .K 2 0).H 2 ; MoO 3 .K20.5H 2 ; Mo0 3 .(NH 4 ) 2 0. 
W0 3 .Na 2 0; WO :i .K 2 O.H 2 O, 2H 2 0, and 5H 2 0; 

TJ0 3 .(NH 4 ) 2 0. 

The lithium, sodium and potassium salts are obtained by 
fusing the respective trioxides with the carbonates. The molyb- 
dates and tungstates are colourless; the uranates yellow. The 
ammonium salts crystallise from solutions of trioxides in am- 
monia ; they are precipitated by alcohol ; the neutral ammonium 
tungstate is, however, unstable, and yields crystals of 
3WO 3 .2(NH 4 ) t O.3H 2 O. Ammonium molybdate is employed as a 
reagent for orthophosphoric acid. 

The double salt 3MoO3.K 2 O.2Na2O.14H 3 O is also known, and 
is produced by mixture. Sodium tungstate, NagWO^ is used as 
a mordant in dyeing, and, as it fuses at a red heat, it is employed 
to render linen and cotton cloth uninflammable. 

; 2Mo0 3 .(NH 4 ) 2 (X 2W0 3 .Na 2 0.2H 2 O and 6H 2 O; 
and 3H 2 0. 2U0 3 .Na 2 ; 

These are crystalline salts obtained by acidifying the former, or 
by adding trioxide in theoretical proportion to their solutions. 
The uranium salts are also produced by addition of excess of 
solution of potassium hydroxide to a uranyl salt, such as the 
nitrate, U0 2 (N0 3 )2.Aq. They are light orange powders. 


7MoO 3 .3Na2O.22H 2 O; 7MoO.3K2O.4H2O; 7Mo0 3 .3(NH 4 ) 2 O.22H 2 O. 
. 7W0 8 .3Li0.19H 2 0; 

These molybdates are produced by evaporation to dry ness of 
sqlutions of the trioxide in solutions of carbonates. From the 
potassium salt the curious compound 16MoO3.6K 2 O.4H 2 O2 has 
been obtained with hydrogen dioxide. The tungstates are obtained 
by the action of carbonic acid on the former salts. 

5Mo0 3 .2(NH 4 ) 2 0.3H 2 0. 5W0 3 .2Na20.11H 2 j 5WO 3 .2Kj0.2H 2 0. 
9MoO d 4E^O.GH 2 O. 

These salts, and those which follow, are produced by acidifying 
those in which the number of molecules of the two oxides are more 
nearly equal. 

12W0 3 .5Na 2 0.28H 2 0;12W0 3 .5K 2 0.11H 2 0; 12W0 3 .5(NH 4 ) 2 O.5H 2 O and 

11H 2 0. 
12WO 3 .Na 2 O.4K2O.15H 2 O ; 12WO 3 .Na 2 O.4(NH 4 ) 2 O.12H 2 O ; and others. 

3MoO 3 .Ha 2 O.4H 2 O and 7H 2 O ; SMoOa.KjO.SHgO; 3MoO 3 .Rb 2 O.2H 2 O. 

3W0 3 .Na 2 0.4H,0 ; 

Molybdates of the following types are also known ; they are all 
produced by addition of acid to those containing less trioxide : 

4Mo0 3 .Na 2 O.6H 2 O ; 

10Mo0 3 Na 2 0.21H 2 0; 16MoO 3 .Na 2 O. 

A corresponding tetratungstate, 4WO 3 Na^O, remains insoluble 
on digesting the fused salt, 12WO 3 .5Na 2 O, with water; the salt 
3WO 3 .2Na2O dissolves. A hexuraiiate, 6UO 3 .K 2 O, is also 
known; it is produced by fusion of uranyl sulphate, UO 2 .SO 4 , 
with potassium chloride. All these bodies are crystalline, and 
apparently definite chemical compounds. 

The tetratungstates, or, as they have been termed, the 
metatungstates, form a separate class, inasmuch as the tung- 
stic acid produced from them is soluble. They are produced by 
boiling solutions of ordinary tungstates with hydrated tungstic 
acid, WO 3 .2H 2 O, or by adding phosphoric acid to a solution of a 
tungstate until the precipitate at first formed redissolves. They 
are also obtained by adding carbonates to metatungstic acid, pro- 
duced from the barium salt with sulphuric acid. They form well- 
defined colourless crystals. Those of the firs.t group have the 
formulae 4WO 3 .Na 2 O.4H,O and 10H 2 O ; 4WO 3 .K 2 O.8H 2 O ; and 
4WO 3 .(NH 4 ) 2 O 8H 2 O. 

The sulpho-compounds are less complicated. They are as 
follows : 


MoS 3 Na 2 S ; MoS 3 K>S ; MoS ? ,(NH,) 2 S. WS 3 .NaoS WS :i K,S ; 
WS 3 (NH 4 ) 2 S.~Also 2MoS 3 Na 2 S ; 2MoS3.K2S.-2WS3.Na2S;* 2WS 3 X2S. 
The analogous oxysulphides have also been prepared: MoO 2 S. (NH 4 ) 2 S. 
S. 2MoO 2 S.Na 2 S. WO 3 K 2 S.H.O. 

No similar uranium compounds are known. ' 

The sulphomolybdates and sulphotungstates are prepared by 
dissolving the trisulphides of these elements in sulphides of the 
alkalis, and crystallising ; or those of potassium by fusing 
together potassium carbonate, sulphur, carbon, and molybdenum 
or tungsten trisulphide. Potassium sulphoraolybdate forms deep- 
red prisms, which reflect green light. The sulphomolybdates 
yield deep-red solutions; the sulphotuiigstates are yellowish-red. 
The disulphomolybdates and tungstates are produced by adding 
acetic acid to the mono--salts ; they are precipitated by alcohol. 
The oxysulphomolybdates and tungstates are produced by mixture, 
or by adding the hydrosulphide of the metal to a solution of a 
molybdate or tungstate, and evaporating to crystallisation. They 
form golden-red or yellow needles. 

These bodies form well crystallised double salts with potassium 
nitrate, e.g., MoSa.KjS.KNO-, and WS 3 .K 2 S.KNO 3 . 

MoO 3 .2BeO.3H 2 O j MoO 3 CaO ; M6O 3 .SrO ; MoO 3 .BaO. WO 3 CaO ; WO 3 SrO ; 
WO 3 BaO.2TJO 3 .CaO ; 2T7O 3 .SrO ; 2T7O 3 BaO. 7WO 3 .3BaO.8H 2 O j 

These molybdates and tungstates are prepared by fusing the 
chloride of the metal with molybdenum or tungsten trioxide and 
sodium chloride, or by precipitation. They are sparingly soluble 
white crystalline bodies. The uranates are reddish-yellow, and are 
produced by precipitation. 

Calcium tungstate occurs native as scheelite or tungsten in 
white quadratic pyramids, associated with tin-stone and apatite. 
The mineral is insoluble in water. 

: 4W0 3 .CaO ; 4WO 3 .SrO.8H 2 O ; 4WO 3 ,BaO.9H 2 O, 

are soluble salts, prepared by dissolving the carbonate in the acid, 

MoSg.CaS; MoS 3 SrS ; MoS 3 BaS. WS 3 .CaS : WS 3 .SrS ; WS 3 .BaS. 
3MoS 3 CaS ; 3MoS 3 SrS ; 3MoS 3 .BaS. 

The trisulphomolybdates are produced by boiling the trisulphide 
with solutions of the sulphides ; and the monosulphomolybdates 
deposit from the mother liquor. They are dark-red substances. 
The sulphotungstates are produced by treating the tungstates with 
hydrogen sulphide. 

* Annalen, 232, 244. 


MoO 3 .Mg-O.5H 2 O ; Mo0 3 ,ZnO ; MoO s .CdO.~- WO 3 .MffO ; WO 3 .ZnO ; 
WO 3 .CdO -2TTO 3 .MgrO; 2TJO 3 .ZnO. Aho7WO 3 .2MffO.(NH 4 ) 2 O.10H 2 O; 
12WO 3 .3MffO.2(NH 4 ) 2 0.24H 2 O ; 7WO 3 .ZnO.(NH 4 ) 2 O.3H 2 O. 

JThese molybdates and tungstates are produced by fusing together 
the chloride of the metal with sodium molybdate and chloride. They 
form colourless crystals. The uranates are produced by igniting 
the doable acetate of uranyl and the metal. They are not crystal- 
line. The double ammonium salts are obtained by mixture. 

Metatung-states : 4WO3.M9O.8HsO ; 4WO :i .ZnO.10H 2 O ; 
4WO 8 .CdO.10H 2 O. 

These are all colourless crystalline salts, and are prepared from 
the carbonates. 

The sulpho molybdates of zinc and cadmium are dark-bijown 
precipitates. The neutral magnesium salt is soluble, as is also the 
yellow sulphotungstate. The sulphotungstates of zinc and 
cadmium are sparingly soluble yellow bodies. 

Simple boron and yttrium molybdates have not been prepared. 
Boron tungstate is also unknown ; but double compounds of 
WO 3 , B 2 O 3 , an oxide, and water are very numerous. They are 
soluble colourless salts, crystallising well. Owing to their high 
molecular weights, too great confidence must not be placed in the 
formulae given ; but they appear to belong to the following classes : 

10WO 3 .B 2 O 3 .2BaO.20H 2 O. 

9WO 3 .B 2 O 3 .Na 2 O.3H 2 O. 

9WO 3 .B 2 O 3 .2Ba0.20H 2 O ; 9WO 3 .B 2 O 3 .2CdO.15H 2 O. 

14WO 3 .B 2 O 3 .3K 2 O.22H 2 O ; (the barium and silver salts are also known) 

12WO 3 .B 2 O 3 .4K 2 O.21H 2 O. 

7W0 3 .B 2 O 3 .Na2O.llH 2 O. 

The solution of the cadmium salt has the exceedingly high specific 
gravity 3*6. The acid corresponding to the nonotungstate has been 
prepared from the barium salt. It is a syrup, and gives insoluble 
precipitates with solutions of alkaloids, and may be used to separate 
quinine, strychnine, &c., from solutions. It may be regarded as 
boron tungstate. These bodies are all prepared by mixture. 

Aluminium tungstate has the formula 7WO 3 .A1 2 O3.9H 3 O ; it is 
obtained by precipitation. 3WO 3 .Y 2 O<.6H 2 O is also known. Salts 
of gallium and indium have not been prepared. MoO 3 .Tl 2 O is a 
crystalline powder. 

Mo0 3 .FeO; MoO 3 .M:nO.H 2 O ; MoO 3 .CoO; MoO 3 .NiO.~ 
WO 3 .PeO; W0 8 .MnO; WO 3 .OoO; WO 3 NiO; WO 3 .(Fe,Mn)O. 

Sulphomolybdates and sulpho tungstates of similar formula 


have also been prepared. They are produced by precipitation, 

or by fusion of the trioxide with chloride of the metal and common 
salt. Ferrous manganous tungstate is wolf > aw, the chief ore of 
tungsten. It is a hard dark-grey or brownish mineral, asso- 
ciated with tin-ores and galena. 7WO 3 .3MnO.llH 2 O and 
7WO 3 .3NiO.14H 2 O are produced by precipitation. 

The metatungstates known are 4"WO 3 MnO.lOH 2 O ; 4WO 3 .CoO.9H 2 O ; and 
4WO 3 Ni0.8H 2 O. They are soluble. 4MoO 3 FejO d 7H 2 O j also double salts of 
ammonium with chromium and with ferric iron of the general formula) 
10Mo0 3 M 2 3 .3K 2 O.6H 2 O, where M may be aluminium, chromium, ferric iron, 
or triad, manganese. A manganic salt is also known of the formula 

16MoO 3 .Mn 2 3 .5K 2 O.12H 2 O. 

3W0 3 .Or 2 3 .13H 2 ; 7W0 3 .Cr 2 O 3 .9H 2 O j 5WO 3 .Fe 2 O 3 .5(NH 4 ) 2 O.H 2 O; 

the double salt is soluble. 

The uranates have been little studied. Double molybdates 
and chromates have been prepared, of which an example is 
MoO 3 .CrO 3 .K 2 O.Mg0.2H 2 0. 

These oxides do not combine with oxides of carbon ; but with 
titanium dioxide, compounds similar to these with boron oxide have 
been prepared. Among them are 12WO 3 .TiO 2 .4K 2 O and 
10WO 3 .TiO 2 .4K 2 O. Zirconium, cerium, and thorium compounds 
appear to exist, but have not been investigated. 

The silicomolybdates and tungstates are also numerous. 
Silicomolybdic acid has the formula 12MoO 4 .SiO 2 .13H 2 O ; it 
forms fine yellow crystals. It gives precipitates with salts of 
rubidium and caesium, affording a means of separating these metals 
from sodium and potassium. The corresponding tungstates of 
silicon aud their derivatives have the formulae 

12W0 3 .Si0 2 .4H 2 0.wAq and 10WO 3 SiO 2 .4H 2 O.wAq. 

There are two isomorides having the first formula. The potassium 
salt of the first is produced by boiling gelatinous silica (SiO.(OH} 2 ) 
with di tungstate of potassium. This yields a precipitate with 
mercurous nitrate, from which the acid may be liberated with 
hydrochloric acid. The salts are produced by its action on carbon- 
ates. The ammonium salt of silicodecitungatic acid, 

10WO 3 .SiO 2 .4(NH,) ? O, 

is produced similarly from ammonium ditungstate. This acid also 
yields numerous salts. It is unstable, and, on evaporation, is con- 
yerted into the isom-eric acid of the first formula, which has been 



named decitungstosilicic acid. These acids aud their salts as a 
rule crystallise well. 

The salts of tin have not bpen carefully examined. 

Lead molybdate, MoO 3 .PbO, occurs native as wulfenite or 
yellow lead ore. It is a heavy orange-yellow mineral, occurring in 
veins of limestone. It may also be obtained as a white precipitate, 
or in crystals by fusing sodium molybdate with lead chloride and 
common salt. 

Lead tungstate, WO 3 PbO, also occurs native as scheeletine, in 
quadratic crystals isomorphous with the molybdate. It has a 
greenish or brown colour. It nan be prepared artificially like the 
molybdate, and is then white. The salt 7WO 3 3PbO.10H 3 O is 
produced by precipitation. The raetatungstate, 4WO 3 PbO.6H 2 O, 
crystallises in needles. 2UOj.PbO is yellowish-red and insoluble. 
MoSj PbS and WS 3 .PbS are dark coloured precipitates. 

The oxides of the elements of the vanadium and phosphorus 
groups form exceedingly complex compounds with the trioxides of 
molybdenum and tungsten, and with the oxides of other elements.* 
To these names vanadimolybdates, vanaditungstates, &c., are 
applied, the number of molecules of trioxide being denoted by a 
numerical prefix. The chief compounds are as follows ; they are 
produced by mixture, and are well -crystallised bodies : 

5WOj.V 2 6 .4(NH 4 ) 2 0.l3H 0. 

6MoO 3 . V 2 5 . 2 (NH 4 ) 2 0. 5H 2 O. 

6MoO,.P 2 5 .Aq. 

6WO 3 .As 2 O a .3K2O.3B[ 2 O. 

6WO 3 . As 2 O 3 .4K 2 O. 2H 2 O. 
10WO d .V 2 5 .22H 2 0. 
1 IW0 3 .2P 2 O .6(NH 4 ) 2 O.42H 2 O. 
!GMo0 3 .P 2 6 .3 

. As 2 O 5 .30H 2 O . 
20Mo0 3 .P 2 5 .Na 2 0.23H 2 O. 
20MoO 3 .P 2 O 5 .2BaO.24H.,O. 

20Mx>0 3 JP 2 6 .7K20. 28H 2 0. 
20WO { .P 2 5 .6BaO.48H 2 O. 


16WO 3 .P 2 O 5 .CaO.5H 2 O. 
16WO J .P 2 5 .4K 2 0.2H 2 O. 
16WO.,.P 2 5 .6(NH 4 ) 2 0.10H 2 0. 
16W0 3 .As 2 6 .6Ar 2 0.l1 H 2 0. 
18Mo0 3 .V 2 5 .8(NH 4 ) 2 0. 15H 2 O. 

22WO a .P 2 O 5 .4BaO.32H 2 O. 
24Mo0 3 .P 2 6 .62H: 2 0. 

. GH 2 O. 
24WO.,.P 2 O, i .3K 2 O.21H 2 O. 

24MoO 3 .6P 2 O.6(NH 4 ) 2 O.7H 2 O. 

It is to be noticed that the ratios of the molecules of trioxide 

* Wolcott G-ibbs, Amer. Jour. Sci. (3), 14, 61 ; Amer. Chem. Jour., 2, 217, 
281 ; 5, , 361, 391 ; 7, 209, 313, 392 ; Chem. tf'ws, 45, 29, 60, 60 ; 48, 135. 

2 D 2 


to pentoxide are 5 : 1, 6 : 1, 7 : 1, 8 : 1, 10 : 1, 1(5 : 1, 18 : 1, 
20 : 1, 22 : 1, and 24 : 1 ; that the number of molecules of the 
alkaline oxide varies from 1 to 8 ; and that phosphorus trioxide 
and monoxide (the hypothetical anhydride of hypophosphoijpus 
acid) appear also to be capable of union with these trioxides. 

Still more complex bodies have been prepared, containing two 
or more pentoxides of different elements ; or a pentoxide of one 
and a trioxide of another element, for example : 

14Mo0 3 .8V20 8 .P20 6 .8(NH 4 ) 2 0.50H 2 ; 48Mo0 3 . V 2 O 6 .2P 2 O 6 .7(NH 4 ) 2 O.30H 2 O ; 
60WO 8 .V 2 O 6 .3P 2 O 6 aO(NH4) 2 O.6H 2 O ; 16WO 3 .3V 2 O 6 .P 2 O 6 .5 (NH 4 ) 2 O.37H 2 O. 

Apparently arsenic, antimonic, niobic, and tantalic oxides, and 
the trioxides of boron, phosphorus, vanadium, arsenic, and anti- 
mony, are capable of forming similar compounds. A quaternary 
compound has even been obtained of the formula 

60WOa.8P2O5.V2O5.VO2 18BaO.150H 2 O, 

Some complex uranium compounds occur native, resembling to 
some extent those mentioned above. They are : 

Frdfferite .. .. 3TJO 3 .A8 2 O 5 .12H 2 O ; 

Walpurgin .. .. 3UO 3 .5Bi 2 O 3 .2A8 2 O 6 .10H 2 O ; 

Zeunerite .. . . 2TJO 3 .As 2 O 5 .CuO.8HO ; 

Uranospinite . . . . 2UO a .As 2 O s .BaO.8H 2 O, and 

UranosphcBrite .. .. yO3.Bi2O3.H2O. 

They are yellowish or green crystalline minerals. 

Indications also appear to exist of complex molybdotungstates, 
but they have not been investigated. 

Uranyl tungstate, WO 3 .UO^.H. 2 O, is a brown precipitate. 
Triple compounds have also been prepared of the trioxides of 
molybdenum or tungsten, one of the oxides of sulphur, and the 
oxide of an alkaline metal, but at present there are no precise data 
as to their formulae. 

Oxides of the platinum group of metals also form similar com- 
pounds. Among the few which have been prepared are : j 

10MoO 3 .PtO 2 .4Na2O.29H 2 O and lOWOg.PtO^^O.OHaO. 

They are analogous to the titani- and silici-decimolybdates and 
tungstate s. 

' The compounds of copper are : 

3MoO 3 .4CuO.5H 2 O ; the meta tungstate, 4WO 3 .CuO.llH 2 O ; the sulplio- 
molybdate, MoS 3 .CuS j and the sulphotungstate, WS 3 .CuS. 

They are obtained by precipitation. 


The silver salts are : 

MoO 3 .Agr 2 O; 2WO 3 .Agr 2 O; the metatungstate, 4WO5.AffaO.3H2O; the 
uranate, 2T7O 3 .Aff 2 O ; and the sulphomolybdate and sulphotuugstate, 
MoS 3 .A* 2 S, and WS 3 .A* 2 S. 

They are all insoluble, except the metatungstate. The action 
of hydrogen at the ordinary temperature on silver molybdate 
or tungstate is said to produce sub-argentous salts, containing 
the oxide Ag 4 0, but in the light of recent researches this action 
is improbable. 

The following mercury compounds have been prepared : 

MoO 3 .Hff 2 O; 2MoO. 1 .Hgr 2 O ; WO 3 .H* 2 O ; 2W0 3 .3H*O; 3WO 3 .2H*O; the 

metatungstate, 4WO.,.H 2 O. >5H 2 O ; and the sulpho-compounds, MoSj.H8r 2 S, 
MoS 3 .HgrS, WS 3 .fiT2S and WS^.H^S. 

These are all produced by precipitation ; even the metatung- 
state is insoluble. Mercurous tungstate, WO 3 .Hg>O, is completely 
insoluble in water, and on ignition, leaves tungsten trioxide ; hence 
tungsten trioxide is usually separated from other metals and 
estimated by precipitation with mercurous nitrate. 

The tungstates and molybdates generally resemble the sul- 
phates in their formula); and these might with reason be written 
from analogy M 2 Mo0 4 and M 2 W0 4 ; and a few uranates appear 
also to possess similar formulae. Salts analogous to anhydro- or 
di-sulphates are also known, such as M 2 Mo 2 7 = 2MoO.M 2 O and 
M 2 W 3 7 = 2WO 3 .M 2 ; the uranates, as a rule, are thus constituted. 
Bat as nothing is known of the constitution of the- more complex 
salts, which, as has been seen, are very numerous, the provisional 
method of writing the formulae of the oxides separately has uni- 
formly been adopted. 

Peruranates. It has been stated that the solution of a uranyl 
salt yields a white compound of the formula UO^HaOn on treat- 
ment with hydrogen dioxide. This compound when mixed with 
solution of potassium hydroxide gives a precipitate of the hydrated 
trioxide, UO 3 2H 2 O, while the salt UO.2K 2 O.10H 2 O, goes into 
solution, and may be separated as a yellow or orange precipitate 
on addition of alcohol. It may also be produced by adding hydro- 
gen dioxide to a solution of hydrated uranium trioxide in caustic 
potash. The sodium salt, similarly prepared, has the formula, 
UO 6 .2Na^O.8H 2 O ; and by using a smaller amount of alkalis 
hydroxide, the compound UO 6 .UO 3 .Na2O.6H 2 O is formed, and 
separates on addition of alcohol. The analogous ammonium com- 
pound has also been prepared. These compounds would lead to 
the inference that an oxide of the formula UO 8 is capable of 
existence ; but it has been suggested, apparently on insufficient 


evidence, that they are in reality compounds of uranium tetroxide 
with peroxides of the metals, thus : UO^K.CMOHaO ; 
UO4.2Na2O2.8H2O ; and 2UO4.Na2O2.GH2O. They readily part 
with oxygen, forming uranates. Similar permolybdates and per- 
tungstates are said to be capable of existence at low temperatures. 

Persulphomolybdates. Potassium dimolybdate on treatment 
in solution with hydrogen sulphide yields a mixture of potassium 
feulphomolybdate, MoS 3 K 2 S, and molybdenum trisulphide, MoS 3 . 
Such a mixture, when boiled with water for some hours, gives off 
hydrogen sulphide, and forms a copious precipitate ;> it is collected 
and washed with water until the washings give a red precipitate of 
MoS 4 with hydrochloric acid. Water extracts potassium persul- 
phomolybdate, MoS 4 .K 2 S from the residue, leaving the disulphide, 
Mo8 2 . On treatment with hydrochloric acid, the tetrasulphide is 
precipitated, and from it the salts may be obtained by treatment 
with sulphides. The alkali and ammonium salts are soluble with 
a red colour ; they yield precipitates, usually red or reddish-brown, 
with soluble salts of the metals. The magnesium salt is an in* 
soluble red precipitate. 

(d.) Compounds with halides. No simple oxyfluorides of 
molybdenum arc known. But by dissolving molybdates in hydro- 
fluoric acid and evaporating tho solutions, compounds isomorphous 
with siannifluorides, SnP 4 .2MF.H 2 O, and titani- and zirconi- 
fluorides of corresponding formulas are produced. Molybdoxy- 
fluorides of the geneial formula MoO 2 F 2 .2MF.H 2 O have been 
prepared with potassium, sodium, ammonium, and thallium ; of the 
formula MoO 2 F 2 .2MF.2H 2 O with rubidium and ammonium ; and 
with 6H 2 O with zinc, cadmium, cobalt, and nickel. 

Tungstoxyfluorides have been similarly prepared ; also one of 
the formula WO 3 .3NH4F. They are isoraorphous with the former 
salts. The oxjrfluoride itself is known with uranium, UO 2 F 2 . It 
is a white substance produced by evaporating a solution of the 
trioxide in hydrofluoric acid ; and has been obtained in crystals 
by subliming the tetrafluoride, UF 4 in air. It also forms double 
salts on mixture ; for example, UO 2 F 2 .NaF.4H.O ; UO 2 F 2 .3KF ; 
UO 2 F 2 .5KF ; and 2UO 2 F 2 .3KF.2H 2 O. They are crystalline yellow 
bodies. The salt UO 2 F 2 .KF is a yellow crystalline precipitate 
obtained by adding a solution of potassium fluoride to uranyl 
nitrate, UO,(NO 3 )2.Aq. 

Oxychlorides and oxybromides of all these elements are known, 
viz. : MoO 2 CI 2 , WO 2 C1 2 , UO 2 C1 2 ; and MoO 2 Br 2 , WO 2 Br a , and 
UO 2 Br 8 . They are all produced by the action of the halogen on 
the heated dioxides or by heating the trioxides in a current of 


hydrogen chloride or bromide. They may also be formed by passing 
the halogen over a hot mixture of the trioxide with charcoal ; 
and one, MoO 2 Br a , has been prepared by heating a mixture of the 
trioxide with boron trioxide and potassium bromide : MoO 3 + 
B 3 b 3 4- 2KBr = 2KBO 3 + MoO.Br,. Molybdyl dichloride, 
MoO 2 Cl 2 , forms square reddish-yellow plates ; it volatilises without 
fusion. The bromide also volatilises in crystalline yellow scales* 
They are soluble in water, alcohol, and ether. Tungstyl dichloride, 
WO 2 C1 2 forms lemon-yellow scales ; and the bromide consists of 
scales like mosaic gold. They decompose when heated. Uranyl 
dichloride, UO 2 C1 2 , is a yellow crystalline fusible body, volatile with 
difficulty ; the bromide forms yellow needles. An oxyiodide is said 
to have been made. 

Molybdyl and uranyl dichlorides form compounds with water, 
MoO 2 Cl 2 .H 2 O and UO 2 C1 2 .H 2 O. The first of these is a white 
crystalline substance, very volatile in a current of hydrogen 
chloride ; it is produced, along with the anhydrous body, by 
passing hydrogen chloride over molybdenum trioxide at 150 200. 
Uranyl dichloride unites with chlorides of the alkalies, forming 
bodies, such as UO 2 C1 2 .2KC1.2H 2 O, similar to the fluorides ; the 
corresponding bromide also gives salts, e.g., UO 2 Br 2 .2KBr.7H 2 O. 

Molybdenum and tungsten also form other oxyhalides, MoOClj, 
WOCli, and WOBr 4 . These may b'e named molybdanosyl and 
tungstosyl tetrachlorides, respectively. The first is produced, 
along with molybdyl dichloride, by the action of chlorine on a 
heated mixture of the trioxide with charcoal. It forms green 
easily fusible crystals, which melt and sublime below 100 ; it is 
soluble in alcohol and in ether. The corresponding tungsten 
compound is produced when tungstyl chloride is quickly heated 
above 140. It forms red transparent needles ; it melts at 210*4, 
and boils at 227'5 U . Its vapour density corresponds with the 
formula WOCli. The bromide is similarly prepared by heating 
tungstyl dibromide ; it forms light-brown woolly needles. 

Molybdenum forms some other oxyhalides. The action of 
chlorine on a mixture of molybdenum trioxide and carbon gives, 
besides the compounds already mentioned, two others : Mo 2 O 3 Cl 6 , 
which forms dark violet crystals, ruby-red by reflected light, and 
volatile without decomposition ; and Mo 4 O 5 Cli , forming large 
blackish- brown crystals, volatile in a current of hydrogen. The 
first points to a dimolybdic acid, Mo 2 3 (OH), but the second is a 
derivative of a lower oxide. 

Molybdous bromide, MoBr 2 , on treatment with alkali,' yields a 
solution from which carbonic or acetic acid throws down the 


hydroxybromide, Mo 3 Br 4 (OH) 2 , as a yellow sparingly soluble 
precipitate. This body acts as a base, yielding a crystalline 
sulphate, Mo 3 Br 4 .SO 4 , chromate, Mo 3 Br 4 .CrO 4 , molybdate, 
Mo 3 Br 4 .MoO 4 , oxalate, Mo 3 BrvC,O 4 , and phosphate, 
Mo 3 Br 4 .2PO(OH) 2 . f 

The oxy chlorides, such as MoO 2 Cl 2 and MoOCl 4 , point to 
hydroxides like MoO 2 (OH) 2 and MoO(OH) 4 ; these are known 
with all three elements and are the respective acids. 

No sulphohalides'are known. 

Physical Properties. 

Mass of one cubic centimetre : 

MoO 2 ,6'44 grams at 16. MoO 3 , 4 39 grams at 21. MoSa, 4'44 4'59 grains. 
WO 2 , 1211 grams. WO 3 , 7'23 grains at 17. WS. 2 , G 26 grams at 20. 
TJO 2 , 10-15 grams; U a O 8 , 7'19 grams ; TJ0 3 , 51)2-: 2G grams. 

The heats of formation are unknown. 






These compounds are most conveniently divided into the two 
classes: (1) the oxides and their compounds; and (2) the 
compounds of sulphur, selenium, and tellurium with each 

The following is a list of the oxides : 

Sulphur. Selenium. Tellurium. 

S 2 O 3 *; S0 2 ; SO 3 ; S,0 7 f. SeO 3 . TeO 2 ; TeO 3 . 

Besides these, the double oxides SSeOi, STeO 3 , and SeTeO 3 are 
known, analogous to the oxide S 2 Oj. 

Sources. Sulphur dioxide is the only one of these compounds 
occurring native. It is present in the air in the neighbourhood of 
volcanoes, being produced by the combustion of sulphur, and also 
in the air of towns, where its presence is due to the combustion of 
coal, which almost always contains small quantities of iron pyrites. 
Air in the neighbourhood of furnaces where sulphides are roasted 
also contains this gas. It is very injurious to vegetation, and the 
prevention of its presence in the atmosphere in large quantities 
should engage the attention of manufacturers. 

Sulphur trioxide exists in abundance in combination with other 
oxides in sea- water, or on the earth's surface, as sulphates, and 
selenium trioxide has been found native in combination with lead 
oxide. The more important of the natural sulphates are Glauber'* 
salt, or sodium sulphate, Na2SO 4 .10H 2 O, which is contained in 
sea- water and in many mineral waters, and when solid, in efflor- 
escent crusts, is named thenardite ; glaserite, K 2 SO 4 , in sea-water 
and spring- water ; schonite, K 2 Mg( k SO 4 ) 2 .6H 2 O, anhydrite, CaSO 4 , 
and gypsum, CaSO 4 .2H 2 O ; celestin, SrSO 4 ; heavy spar, or 
barytes, BaSO 4 ; Epsom salt, MgSO 4 .7H 2 O ; feather alum, 

* Pop \ Ann., 156, 531. 

f Comptes rend., 86, 20, 277 ; 90, 269. 


A1(SO4) 3 .18H 2 O ; alum ston6, A1KSO 4 .2A1(OH) 3 ; copperas, or 
green vitriol ', FeSO 4 .7H 2 O; cobalt vitriol, CoSO 4 .7H 2 O; angltsite, 
PbSO 4 ; lanarkite, a double carbonate and sulphate of lead, and 
leadhillite, PbSO 4 .PbO ; and blue vitriol, CuSO 4 .5H 2 O. Lad 
selenate, PbSeO 4 , has also been found native. 

Preparation. 1. By direct union. Sulphur, selenium, and 
tellurium burn with a faint blue flame when heated in air, forming 
the dioxides. Heated in oxygen, the flame of burning sulphur is 
much more brilliant, and of a fine lilac colour. Its combustion 
forms a telling experiment. About 3 or 4 per cent, of the product 
of the combustion consists of sulphur trioxide, S0 3 . The sulphides 
of many metals, when roasted in air, give the oxide of the metal 
and sulphur dioxide. Iron pyrites containing from 2 to 4 per cent, 
of copper is made use of in its commercial preparation, the copper 
being extracted from the residue. The sulphur dioxide is em- 
ployed directly in the preparation of sulphuric acid. It is also 
a by-product in the roasting of zinc sulphide, in the smelting 
of lead ores (see p. 429), and in various other metallurgical 

2. By oxidation of a lower oxide. Sulphur trioxide is 
thus prepared on a commercial scale. In the laboratory it may 
be prepared by the following method : 

A dry mixture of gaseous sulphur dioxide and oxygen, the 
dioxide being made to bubble through the wash- bottle containing 
strong sulphuric acid twice as quickly as the oxygen, is led through 
a tube of hard glass, heated to redness, filled with asbestos, pre- 
viously coated with metallic platinum by moistening it witt 
platinum tetraohloride, and igniting it. Under the influence of the 
finely-divided platinum, the sulphur dioxide and the oxygen com 
bine, and the sulphur trioxide produced is condensed in a flask. 
To obtain the pure trioxide, water must be rigorously excluded, 
and corks should not be exposed to its action, for they are at once 

By passing an electric discharge of high potential through a 
mixture of perfectly dry sulphur dioxide and oxygen, combination 
takes place between 4 vols. of sulphur dioxide and 3 vols. of 
oxygen to form persulphuric anhydride or disulphur heptoxide, 
S 2 7 . 

3. By reducing a higher oxide. -The trioxides of sulphur 
and of tellurium, at a red heat, decompose into the dioxides and 
oxygen. The vapour of sulphuric or selenic acid also, at a red 
heat, gives water, sulphur or selenium dioxide, and oxygen, and it 
is by this method that a mixture in the requisite proportion of 


sulphur dioxide and oxygen is obtained on a large scale for the 
manufacture of sulphur trioxide. The sulphuric acid is decom- 
posed by causing it to flow on to red-hot bricks ; and the mixed 
gfises are dried by passage upwards through a tower filled with 
coke, kept moist by strong sulphuric acid.* The mixture is then 
passed over asbestos coated with platinum, as on the small scale 
(see previous page). 

The reduction may also be effected by chemical agency. On 
heating sulphur and sulphuric acid, the dioxide and water are the 
sole products, thus : 2(S0 3 .H 2 0) + S = 30 2 + 2H 2 O. Carbon 
may be used in the form of charcoal ; in this case a mixture of 
carbon dioxide and sulphur dioxide is produced, from which it is 
not easy to separate the carbon dioxide :2(S0 3 .H 2 0) -f C = 
2$0 3 + COz + 2H 2 0. Almost all metals, when heated with strong 
sulphuric acid, yield sulphur dioxide, a sulphate and sulphide of the 
metal, hydrogen sulphide, free sulphur, and water. For example, 
with copper, the metal most frequently employed in the form of 
foil or turnings in the ordinary laboratory process for preparing 
sulphur dioxide :f 

(1.) Cu + 2H 2 SO 4 = CuSO 4 t- tfO 2 + 2H 2 
(2.) 4Cu + 5H 2 SO 4 = 4CuSO 4 + H 2 S + 4If 2 O ; 
(3.) 3Cu + 4H 2 SO 4 = 3CuS0 4 + CuS + 4H 2 O ; 
(4.) 4Cu + 4IJ 2 SO 4 = 3CuSO 4 + Cu 2 S + 4H 2 O; and 
(5.) O 2 + 2H^8 2H 2 O + 38. 

Reaction (1) is that which predominates ; but the other reactions 
doubtless take place, for the products are found in the residue. 

It is probable that these reactions are due to the action of hot 
nascent or atomic hydrogen on. sulphuric acid. Thus, the equa- 
tions may also be written : 

(1 ) Cu + H 2 S0 4 = CuS0 4 + 2H- t H 2 SO 4 + 2H = 2H 2 O + #0 2 ; 

(2.) H 2 SO 4 + 8AT = 4H 2 O + H^S ; 

(3.) CuS0 4 + J/ 2 tf = CuS + H 2 SO 4 ; and 

(4.) 20uSO 4 + 10 R = Cu 2 S + H 2 SO 4 + 4H 2 O. 

The metals osmium, iridium, platinum, and gold are the only 
ones which withstand the action of boiling sulphuric acid ; but 
strong acid may be evaporated in iron pans, for the iron becomes 
protected by a coating of sulphate, which is insoluble in oil of 

Gold is, however, attacked by selenic acid ; the acid is reduced 
by it and other metals to the dioxide. Selenic acid is also converted 

Ditffl. polyt. J., 218, 128. 
f Chem. Soc., 33, 112. 


into selenious acid, with evolution of chlorine, by boiling it with 
hydrochloric acid, thus : 

H 2 Se0 4 + 2HCl.Aq = H 2 Se0 3 .Aq + H 2 O + Ck. 

The oxides S 2 3 , SSe0 3 , STeO 3 , and SeTeO 3 are also formed 
by reduction. They are produced by dissolving sulphur in fused 
sulphur trioxide ; selenium in sulphur trioxide ; and tellurium in 
strong selenic acid. 

4. By heating compounds. Both sulphites and sulphates, 
and probably also selenites, selenates, tellurites, and tellurates, when 
hnated to a high temperature decompose, leaving the oxide of the 
metal with which the oxide of sulphur, selenium, or tellurium was 
combined. But the dioxides are usually produced, for the tem- 
perature at which decomposition occurs is almost always so high 
as to partially, at least, decompose the trioxides. Anhydrosul- 
phates, such as Na 2 S 2 O 7 , however, give off half their trioxide when 
heated, leaving the monosulphate, Na 2 SO 4 . The compounds 
with water, however, in every case, except that of selenic acid, are 
decomposed by heat, yielding the respective oxide. 

Thus a solution of sulphur dioxide in water, presumably con- 
taining sulphurous acid, H 2 S0 3 , loses the oxide when boiled; 
selenious and tellurous oxides remain on evaporating their aqueous 
solutions ; the latter, indeed, separates out on warming its solution 
to 40; sulphuric acid, H 2 S0 4 , when gasified has the density 44'5, 
proving it to have split into its constituent oxides, which, however, 
recombine on cooling; the trioxide is prepared, moreover, by dis- 
tilling anhydrosnlphuric acid, H 2 S 2 O 7 , which decomposes thus : 
H 2 S 2 O 7 = H 2 SO 4 + 80 Z ; and tellurium trioxide is produced by 
heating the hydrate to a temperature below redness. 

5. By double decomposition. As this process is usually 
carried out in the presence of water, the hydrates (acids) are the 
usual products. 

6. By displacementThis is a convenient method of pre- 
paring sulphur trioxide. Strong sulphuric acid is mixed with 
phosphoric anhydride, care being taken to keep the acid cold 
during mixing. It is then distilled, when the trioxide passes 
over, the phosphoric anhydride having abstracted water from the 
sulphuric acid, thus : 

H 2 SO 4 + P2O = 3 -f 2HP0 3 . 

A sulphate also, when strongly ignited with silicon dioxide, 
or with phosphorus pentoxide, yields sulphur trioxide, or its pro- 
ducts of decomposition, the dioxide and oxygen. This process 


finds practical application in the manufacture of glass, where 
silica in the form of sand is heated with sodium sulphate, lime, 
and carbon. The addition of carbon causes the conversion of the 
sulphate into sulphite ; the silica replaces the sulphur dioxide at a 
lower temperature than it would replace the trioxide of the 
sulphate. A double silicate of sodium and calcium ia thus 
formed, which constitutes one variety of glass. The method, it 
will be seen, is not available for the preparation of the oxides of 

Properties. Sulphur dioxide is a gas at the ordinary tem- 
perature, but it may be easily condensed to a liquid by passing it 
first through a tube filled with calcium chloride, to dry it, and 
then through a leaden worm cooled by a mixture of salt and 
crushed ice. 

It boils at 8 under normal pressure, and melts at about 
79. The liquid oxide is mobile and colourless, and heavier than 
water (1*45). It forms a white crystalline solid when sufficiently 
cooled by its own evaporation. The gas has the familiar smell of 
burning sulphur; it is irrespirable ; it supports the combustion of 
potassium, tin, and iron, which combine both with its oxygen and 
its sulphur. It is readily soluble in water; one volume of water 
absorbs about fifty times its volume of the gas at the ordinary 
temperature, probably with formation of sulphurous acid, H 2 S0 3 . 
Hence it cannot be collected over water ; but, as its density is 
high (32), it is easy to collect it in a jar by downward displace- 

Selenium dioxide is a white solid, volatilising to a yellow 
vapour without melting, at a heat somewhat below redness, and 
condensing in white quadrangular needles. Its vapour has a 
sharp acid odour. It is soluble in water, producing selenious 

Tellurium dioxide is a white solid, sometimes crystallising 
in octahedra. It melts to a deep-yellow liquid, and at a high 
temperature it volatilises. It is sparingly soluble in water, and 
does not appear to form the acid. 

Sulphur trioxide crystallises in long colourless prisms, 
arranged in feathery groups ; it somewhat resembles asbestos. It 
melts at 15, and boils at 46, producing dense white fumes with 
the moisture of the air. Its molecular weight, as shown by its 
vapour density, is 80. It unites with water with great violence, 
hissing like a red-hot iron. It is made in considerable quantity, 
being used in the manufacture of alizarine or turkey-red, and 
other artificial dyes. 


When this body is kept for some time at a temperature'. 
below 25, it changes into another modification which crystallises 
in thin needles. When heated above 50 it gradually liquefies, 
and changes into the first modification. It is distinguished 
from the first modification by the difficulty with which it dfs- 
solves in H^SO^ and by its crystallising out of the solution un- 

There appear to be indications of the existence of an oxide 
S 2 5 ; for sulphur trioxide dissolves the dioxide in large amount, 
and the solution is stable up to 5. 

No attempts to prepare selenium trioxide have succeeded. 
The acid, when heated, decomposes into selenium dioxide, oxygen, 
and water. Selenious anhydride is the only product of the action 
of oxygen, even in the state of ozone, on selenium. 

Tellurium trioxide is an orange-yellow insoluble substance, 
which does not dissolve even in hydrochloric or nitric acid. 
When strongly heated, it loses oxygen, producing tellurium 

Sulphur, selenium, and tellurium dissolved in pure melted 
sulphur trioxide give respectively blue, green, and red substances. 
The sulphnr and tellurium compounds have been isolated, and 
have been shown to have the formulae S 2 O 3 and STeO 3 ; it is pre- 
sumed that the others are similarly constituted. Selenium and 
tellurium also dissolve in concentrated selenic acid, doubtless form- 
ing similar compounds. The sulphur compound is insoluble in 
perfectly pure sulphuric anhydride, and may be separated from it 
by decantation. It decomposes, on exposure to air at ordinary 
temperatures, into sulphur dioxide and sulphur. It dissolves in 
strong snip baric acid, and, on diluting the acid, it is decomposed. 
The tellurium compound appears to exist in two modifications, a 
red one, and a buff -coloured, obtained by heating the red variety 
to 90. 

Persulphuric anhydride, S 2 O 7 , at the ordinary temperature 
forma an oily liquid ; when cooled to 0, it solidifies in long thin 
transparent flexible needles. It sublimes easily, and decomposes 
spontaneously on standing for a few days. It dissolves in strong 
sulphuric acid ; it is immediately decomposed by heat. 

A, Compounds with water and oxides; acids and salts 
of sulphur, selenium, and tellurium. 1 .Compounds of th 
trioxides; sulphuric, selenic, and telluric acids; sulphates, 
selenates, and tellurates. 

The trioxide of sulphur dissolves in water with evolution of 
great heat, forming various hydrates, according to the relative 


proportion of oxide and water. The following have been 
isolated : 

SO 3 .5H 2 O; S0 3 .3H 2 0; SO 3 .2H 2 O ; S0 3 .H 2 = H 2 S0 4 ; 
c 2S0 3 .H 2 = H 2 S 2 O 7 . 

Those containing less water than ordinary sulphuric acid are 
more conveniently produced by dissolving sulphur trioxide in the 
ordinary acid ; those containing more, by pouring ordinary sul- 
phuric acid into water. Salts have been produced corresponding 
to the acids H 2 S0 4 and H 2 S 2 7 ; they are named sulphates, and 
pyrosulphates or anhydrosulphates respectively. 

On boiling a solution of sulphuric acid in water, the water 
evaporates, and the acid becomes more and more concentrated, 
until it acquires nearly the composition expressed by the formula ' 
H 2 S0 4 ; on further heating, this compound dissociates into trioxide, 
or anhydride, $0 3 , and water, both of which evaporate together. 

Some of these hydrates may be dismissed in a few words. The 
hydrate, SO 3 .5H 2 O = H 2 S0 4 4H 2 0, crystallises out on cooling sul- 
phuric acid containing the correct amount of water to a very low 
temperature. It melts at 25. H 3 S0 4 .2H 2 is the point of 
maximum contraction of sulphuric acid and water, but has not 
been obtained in a solid state; and H 2 SO 4 .H 2 O is also obtained 
by cooling a mixture in the correct proportion. It melts at 8. 
The coi responding selenic acid, H 2 SeO 4 .H 3 O, melts at 25. The 
monohydrate requires particular attention. 

Sulphuric acid, "oil of vitriol," H 2 S0 4 . Sulphur trioxide, 
as has been mentioned, is decompo&ed by heat, and hence it cannot 
be produced in quantity by the combustion of sulphur in air or 
oxygen, for the temperature of burning sulphur is higher than that 
at which the trioxide decomposes. Hence an indirect method of 
preparation must be chosen. It can be prepared in aqueous 
solution by oxidising sulphur; for example, when boiled with 
nitric acid, that acid parts with its oxygen, oxidising the sulphur 
to sulphuric acid, while oxides of nitrogen are liberated. Sulphur 
may also be oxidised on treatment with chlorine and water, 
thus : 

g + 3C1 2 + 4H 2 = H 2 S0 4 + 6HC1. 

It will be remembered that the halogen acids may be prepared 
by the action of the halogens in presence of water on hydrogen 
sulphide (see p. 106) ; and, similarly, an aqueous solution of sulphur 
dioxide is oxidised to sulphuric acid by their action. Chromic 
adid ahd other oxidising agents also effect such oxidation. 


But such processes are too expensive to be used in manufacture. 
The main outlines of the process actually in use are given here ; 
the details and the connection of this with other manufactures 
will be described later (see p. 667). 

Sulphur dioxide at once attacks nitrogen peroxide, NO 2 . With- 
out discussing intermediate products, which will be afterwards 
considered, the final reaction, in presence of water at least, is 
S0 9 + N0 2 + H 2 = SO 3 .H 2 4- NO. In presence of air, as 
has been seen on p. 333, nitric oxide is oxidised to a mixture of 
peroxide and tetroxide, N0 2 and N 2 0^ These gases again part with 
their oxygen when brought in contact with a fresh supply of sulphur 
dioxide. In theory, then, a small amount of nitrogen dioxide is 
capable of converting an indefinite amount of sulphur dioxide, in 
presence of oxygen and water, into sulphuric acid. The nitrogen 
dioxide required for this process is derived from nitric acid, pre- 
pared in the usual manner, i.e., from sodium nitrate and sulphuric 
acid. On bringing it into contact with sulphur dioxide, it is 
reduced, and gives an effective mixture of oxides of nitrogen. 
This process may be illustrated by the following experiment : 
D is a flask containing copper turnings and strong sulphuric 

Fro 42. 

acid, from which, on applying heat, sulphur dioxide is generated. 
B is a similar flask containing copper turnings and dilute nitric 
acid, and yields a supply of nitric oxide when warmed. E is a 
flask containing water. The delivery tubes of these flasks all 
enter the large balloon, A, through a large perforated cork; a 
glass tube passes to the bottom of the globe through a fourth hole 
in the cork, and serves as an exit tube for any excess of gas. 
Nitric oxide is first parsed into the globe. It unites with the 


oxygen of the air, forming a mixture of the dioxide and peroxide, 
which are at once noticeable as red fumes. Sulphur dioxide is 
passed in next, and reacts with the peroxide ; it will be noticed 
tha^ the sides of the globe soon become covered with radiating 
crystals. These are described later ; they consist of hydrogen 
nitrosyl sulphate, SO 2 (OH)(ONO), and are known as "chamber 
crystals.'* Steam is then passed into the globe by boiling the water 
in the flask, E. The crystals disappear and the liquid which 
collects in the globe is dilute sulphuric acid. It may be concen- 
trated by evaporation in a porcelain or platinum basin, till its 
strength is little below that indicated by the formula H^SOj. 

Selenic acid* may be prepared, like sulphuric acid, by the 
action of chlorine water on selenium, or, better, on selenious acid ; 
but on concentration, the selenic acid is reduced by the hydro- 
chloric acid with evolution of chlorine. A better plan is to 
saturate a solution of selenions acid with chlorine gas, thereby 
converting that acid into selenic acid ; ' to saturate the mixed 
selenic and hydrochloric acids with copper carbonate, forming a 
mixture of copper selenate and chloride ; to evaporate to dryness, 
and extract with alcohol, which dissolves the copper chloride, 
leaving the selenate ; and, finally, to dissolve the selenate in water, 
and liberate the selenic acid by precipitating the copper as sul- 
phide by a current of hydrogen sulphide. After filtering off the 
copper sulphide, the selenic acid is concentrated by evaporation. 
It can be obtained nearly anhydrous by evaporation in a vacuum 
at 180. The acid has then the formula H-^SeO^ A higher tem- 
perature decomposes it into selenium dioxide, water, and oxygen. 
One other hydrate of selenium trioxidoiias been prepared by cool- 
ing a solution of the acid of the requisite strength to 32. It 
has the formula H 2 Se0 4 .H 2 0, and melts at 25. Attempts to 
prepare other hydrates in the solid state have not been successful. 

Telluric acid is produced in solution by treating the barium 
salt (obtained by heating tellurium with barium nitrate) sus- 
pended in water, with the requisite amount of sulphuric acid, and, 
after filtration, concentrating the acid by evaporation. Colourless 
hexagonal prisms of the formula H 2 TeO 4 .2H 2 O separate out on 
cooling. It loses its water a little above 100, leaving the acid 
HjjTeO^ as a white solid. 

These acids are also produced by the action of water on the 
chlorides, SO 2 C1 2 , Se0 2 Cl 2 , and Te0 2 Cl 2 . 

Sulphuric and selenic acids are dense, viscid, colourless 

* Proc. Roy. Soo. t 4tf, 13. 

2 E 


liquids, exceedingly corrosive, inasmuch as they abstract the 
elements of water from many organic substances containing carbon, 
hydrogen, and oxygen. A piece of wood placed in strong sul- 
phuric acid is blackened and charred, and siii^ar placed in contact 
with it is converted into a tumefied mass of impure carbon. Pure 
anhydrous sulphuric acid, H^SO 4 , is, however, a solid, melting at 
10*5, and bolenic acid, H 2 3e(X, melts at 58 J . The presence of a 
mere trace of water greatly lowers the melting points of these 

The hydrate of telluric acid, H 2 TeO 4 .2H 2 O, dissolves slowly, 
but to a considerable extent in water. The anhydrous acid, 
H>TeO 4 , can be dissolved only by prolonged boiling with water. 

These acids cannot be said to boil, in the purely physical 
sense of the word At the ordinary temperature, sulphuric acid, 
if perfectly pure, gives off sulphur trioxide, hence the only method 
of obtaining an acid precisely corresponding to the formula H 2 S0 4 , 
is to add sulphuric anh) dride to ordinary oil of vitriol. When 
concentrated by evaporation as far as possible, the acid contains 
about 98 per ctnt. of H 2 SO 4 . On further heating to 827, this acid 
dissociates with apparent ebullition into water and trioxide, which 
recombine on cooling, forming an acid of the same composition. 
By taking advantage of the different rates of diffusion of water- 
^as and sulphuric anhydride, which possess respectively the 
densities 9 and 40, and whose ratio of diffusion is therefore as 
v/40 : v/9, a much stronger acid has been obtained. Acid con- 

G. 43. 


taining 5 per cent, of water was boiled in a flask, while a gentle 
current of air passed downwards through* 'a tube, sealed on to 
the bottom of the other flask ; after an hour, the composition of 
thj remaining acid was approximately 60 per cent, of H 2 SO 4 , and 
40 per cent, of S0 3 . This process of concentration is not applied 
on a laige scale. 

Pure selenic acid begins to decompose into dioxide, oxygen, and 
water at about 200. On distilling dilute selenic acid, water passes 
over up to 205; a little dilute acid then begins to distil over, and, 
at 260, white fumes appear, containing a little trioxide, but for 
the most part consisting of selenium dioxide. 

Telluric acid is non- volatile, and parts with its water below a red 
heat, leaving the anhydride, TeO 3 . 

A great rise of temperature is produced by the action of water 
on sulphuric and selenic acids, due to their combination with it to 
form hydrates. 

The specific gravity of ordinary sulphuric acid is approximately 
1'84 at 15 ; that of selenic acid, 2*61 at 15 ; and of telluric acid, 
3-42 at 18-8. 

Sulphates, selenates, and tellurates. These salts are ob- 
tained by the action of the acids on aqueous solutions of the 
hydroxides or carbonates of the metals; by the action of the con- 
centrated acids at a high temperature on most metals, with evolu- 
' tion of the dioxides ; by the action of aqueous solutions of the 
acids on many of the metals themselves, on the oxides, or hydroxides, 
and on some of the sulphides ; and by heating a mixture of the 
acid and a halide, nitrate, or acetate of a metal, or, in short, with 
any salt containing a volatile or decomposable oxide. Thus, for 
example : 

H 2 SO 4 Aq + 2KOELAq = K 2 S0 4 .A'q + 2II 2 Oj 
TeO 3 + Na 2 CO 3 .Aq = NagTeO^.Aq + C0 2 ; 
H 2 SO 4 Aq -4- Zn = ZriSO 4 Aq + 11^ 
H 2 SeO 4 Aq + CtiO = CuS e O 4 Aq \ H 2 O ; 
H 2 SO 4 .Aq + FeS = FeSO 4 .Aq + ff 2 S ; 
H 2 S0 4 4 2NaCl = Na 2 SO 4 + 2HCI ; 
H 2 SO 4 Aq + CaSO, = CaSO 4 + SO 2 + Aqj 
. H 2 SO 4 + CaSiO., = CaSO. + H 2 SiO 3 . 

The salts of calcium, strontium, barium, and lead are insoluble, 
or nearly insoluble, and may therefore be produced by addition of 
a soluble sulphate, selenate, or tellurate, to the solution of a soluble 
salt of one of these metals, thus : 

CaCl 2 .Aq + Na,SO 4 .Aq = CaSO, + 2NaCl.Aq ; 
Pb(N0 3 ) 2 .Aq + K a Se0 4 .Aq = PbSO 4 + 2KN0 3 Aq. 

2 B 2 



The other sulphates and selenates are soluble in water. Many 
of the tellurates are insoluble, and may be produced by precipita- 
tion. The sulphates are also formed by the oxidation of sul- 
phides by boiling with nitric acid ; by the action of chlorine 
water ; or by the action of air. 

Li 2 S0 4 .H 2 0; N02S0 4 .7, and 10H 2 O; K^; Rb 2 SO 4 ; Cs 2 SO 4 ; (NH 4 ) 2 SO 4 . 
Na 2 Se0 4 .10H 2 Oj KgSeO., ; (NH 4 ) 2 SeO 4 . K^TeO^ELjO; (NH 4 ) 2 TeO 4 . 

Lithium sulphate crystallises in flat tables, easily soluble in 
water and alcohol. Sodium sulphate occurs anhydrous as thenar- 
dite ; and when crystallised with 10H 2 O it is known as Glauber's 
salt. It is prepared in immense quantity from common salt and 
sulphuric acid, as a preliminary to the manufacture of sodium car- 
bonate, and is then termed " salt-cake." It is also produced by 
passing a mixture of steam, air, and sulphur dioxide through 
sodium chloride, heated to dull redness (Hargreave's process). It 
is obtained as a residue in the preparation of nitric and of acetic 
acid ; of ammonium chloride *, and of common salt by the evapora- 
tion of sea-water. It crystallises in an anhydrous stato from water 
at 40 in rhombic octahedra. It is insoluble in alcohol, but very 
soluble in water ; 100 parts of water dissolve 12 parts at 0, and 
48 parts at 18. It crystallises with 10H 2 O in large, colourless, 
nionoclinic prisms. Crystals with 7H 2 O are deposited below 18. 
On raising the temperature of a saturated solution above 33, the 
anhydrous salt deposits, hence it appears to possess a lower solu- 
bility at high than at low temperatures. This apparent abnor- 
mality is doubtless explained by the dissociation of the solution 
of the decahydrate, Na-jSOi.lOHjO, as the temperature rises. 
Sodium selenate is isomorphous with and closely resembles the 
sulphate. The tellurate has not been carefully examined. 

Potassium sulphate crystallises from the aqueous extract of 
kelp (burned seaweed), in trimetric prisms or pyramids. It is 
among the least soluble of the potassium salts, 100 parts of water 
dissolving 8*36 parts at 0. It is insoluble in alcohol. Both 
sodium and potassium, sulphates have a saline 'bitter taste, and a 
purgative action. Potassium selenate is produced by fusing 
selenium or potassium selenite with nitre, and crystallisation from 
water; it resembles the sulphate. The tellurate forms rhombic 
crystals ; they deliquesce in air, becoming converted by carbon 
dioxide and. water into carbonate and ditellurate. Rubidium and 
caesium sulphates resemble that of potassium, but are much more 
soluble in water. . Ammonium sulphate, selenate, and tellurate are 
isomorphous with potassium sulphate, but are more soluble. The 


sulphate, when heated, decomposes above 280, yielding ammonia 
water, and nitrogen, and a sublimate of hydrogen ammonium 
sulphate ; the selenate gives, first, hydrogen ammonium selenate, 
anc^ then selenium, its dioxide, water, and nitrogen. These salts, 
with the exception of lithium sulphate, are all insoluble in alcohol. 

Double salts :HLiSO 4 ; HNaSO 4 ; HKSO 4 ; H(NH 4 )SO 4 . HKSe0 4 ; 

H(NH 4 )Se0 4 . HNaTe0 4 ; 2HKTeO 4 .3H 2 O.-HNa,(SO 4 ) 2 ; HjNa(8O 4 ), ; 
HK t (S0 4 ) 2 ; H 3 K^0 4 ) 2 ; LiKS0 4 ; NaK 3 (S0 4 ) 2 ; H 2 K 4 (SO 4 ) 3 ; Li 4 (NH 4 ) 2 (SO 4 ) 3 ; 
Li,K 4 (S0 4 ) 3 , NaK,(S0 4 ) 3 ; HK 3 (TeO 4 ) 2 . 

These substances are white crystalline bodies, very soluble in 
water, and also, as a rule, in alcohol. They are produced by 
mixture and crystallisation. Bisulphate of potassium, as hydrogen 
potassium sulphate is generally named, is used in decomposing 
various minerals, which are for that purpose reduced to fine 
powder, mixed with* the salt, and fused. When carefully heated 
it loses water and yields the anhydrosulphate, or true disulphate, 
KaSgO?. Sodium tripotassium sulphate is technically named plate- 
salt, from its crystallising in hexagonal plates ; it deposits on 
cooling an aqueous extract of kelp. 

The existence of the more complex double sulphates leads to 
the conclusion that the molecular formulae of the ordinary sul- 
phates are not so simple as they are usually written. Such 
formulae as HzK^SC^ and NaK 6 (S0 4 ) a , lead to the conclusion 
that the formula of potassium sulphate is probably at least 
K 6 (S0 4 )3. Double salts with other acids are also known; e.g., 
K 2 SO 4 .HN0 3 and K a S04.H 3 P0 4 separate from solutions of potas- 
sium sulphate in nitric or phosphoric acid. They are, however, 
decomposed by water. The existence of such salts would also 
favour the supposition of greater complexity of molecule. 

BeS0 4 .2, 4, and 6H 2 O; CaSO 4 2H 2 O ; 2CaSO 4 .H 2 O ; SrS0 4 . BaSO 4 . 
BeSeO 4 .4H 2 O ; CaSeO 4 .2H 2 O ; SrSeO 4 ; BaSeO 4 . 
CaTeO 4 ; SrTeO 4 ; BaTeO 4 .3H 2 O. 

With the exception of beryllium sulphate, which is soluble, 
all these compounds may be prepared by precipitation. Beryllium 
sulphate forms quadratic octahedra ; it is insoluble in alcohol but 
very soluble in water. On evaporation with beryllium carbonate, 
it yields gummy basic salts of the formal 

SO 3 .2BeO.3H 2 O; SO 3 .3BeO.4H 2 O ; SO 3 .6BeO.3H 2 O. 

Calcium sulphate occurs abundantly in the native form in salt 
mines. When anhydrous, it forms trimetric prisms, and is named 
anhydrite ; and with t^o molecules of water it is gypsum ; indivi- 


dual varieties of gypsum are named selenite, alabaster, and satin- 
spar ; sblenite forms transparent colourless monoclinic crystals; 
the massive variety is alabaster; and satin-spar is fibrous. When 
gypsum is heated and ground it forms " plaster of Paris,'' a 
material much employed in taking casts, and as a cement. The 
dihydrated calcium sulphate becomes anhydrous and falls to a 
powder when heated ; on mixing -the powder with water, a pasty 
mass is produced with which casts may be taken. After a few 
minutes it hardens, expanding slightly at the same time, and 
forms a fine white material. Plaster of Paris, mixer) with a 
saturated solution of potassium sulphate, gives a paste which 
solidifies more rapidly than ordinary plaster of Paris, and has a 
nacreous lustre ; for certain purposes this mixture is to be 
preferred to the ordinary one. A double salt K 2 Ca(SO 4 )2.H 2 O, is 
produced. Hydrated calcium sulphate is very sparingly soluble 
in water, and is more soluble in cold than in hot water (1 in 420 
at^ 20). This is probably due- to the solution containing the 
dihydrated compound, which loses water, becoming insoluble as 
the temperature rises. It is much more soluble in weak hydro- 
chloric or nitric acid; or in presence of common salt, or of sodium 
thiosulphate. Its solubility in the last affords a method of sepa- 
rating calcium from barium. Calcium sulphate melts at a red 
heat. The selenate closely resembles the sulphate in preparation 
and properties, and is isonaorphous with 1 it. It is reduced to the 
selenite, however, when boiled with hydrochloric acid, chlorine 
being evolved. 

Calcium tellurate is a white precipitate, soluble in hot water. 

Strontium sulphate, SrS04, occurs native as cwlestin in tri- 
metric crystals. It is soluble in about 7,000 parts of cold water ; 
it fuses at a bright red heat. The selenate resembles it. 

Barium sulphate occurs as heavy-spar or barytes, in largfe 
quantity ; it forms trimetric crystals. A solution of barium 
chloride or nitrate is the common reagent for sulphuric acid. On 
adding it to a sulphate, a dense white precipitate is produced, 
practically insoluble in water and acids. Its insolubility serves to 
distinguish it from most other bodies of similar appearance. In 
estimating sulphuric acid, it is always weighed in the form of 
barium sulphate. It is unaltered by ignition ; when heated with 
charcoal or coke, however, it yields barium sulphide ; and this is 
the usual process of preparing compounds of barium, since the 
sulphide dissolves in acids. Barium sulphate reacts to a limited 
extent when boiled with a solution of sodium carbonate ; a 
portion is converted into carbonate, thus : 


BaSO 4 4- STaaCOg.Aq = BaCO, -f Na 2 S0 4 .Aq. 

But the reaction is incomplete. It is only after removal of the 
sodium sulphate and replacement by fresh sodium carbonate that 
further decomposition takes place. On fusion with excess of sodium 
or potassium carbonate, however, it, is completely converted into 
carbonate. Barium sulphate has been used as a pigment under the 
name perminent white ; it has too little body, and hence it is 
generally mixed with white-lead or zinc- white (ZnS). JJarium 
seienate closely resembles the sulphate, but it is decomposed on 
boiling with hydrochloric acid, selenious acid and chlorine being 
formed. This serves to distinguish it, and to separate it from 
the sulphate. The tellurate is fairly soluble in warm water. 

Double sulphates, seienate s, and tellurates. K 2 Be(S0 4 ) 2 ; ;HjCa(S0 4 ) 2 ; 
H 2 Sr(SO 4 ) 2 ; H 2 Ba(S0 4 ) 2 , also with 2H 2 O ; H fi Ca(SO 4 ) 4 ; Na 2 Ca(SO 4 ) 2 ; 
Na 2 Ba(S0 4 ) 2 ; Na 4 Ca(S0 4 ) 3 .2H 2 O; K 2 Ca 2 (SO 4 ) 3 .3H 2 O ; H 2 Ba(TeO 4 ) 2 .2H 2 O. 

Hydrogen calcium and barium sulphates are crystalline bodies 
produced by dissolving the ordinary sulphates in strong sulphuric 
acid and crystallising. They are decomposed by water. The 
tellurate is soluble in water. The double salts are prepared by 
digesting the simple salts with sodium or potassium sulphate ; 
that of calcium crystallises out with 2H 2 0, but at a higher 
temperature loses water, and is then identical with the mineral 
glauberite, crystallising in rhombic prisms. 

MgSO 4 .7, 6, and 1H 2 O; ZnSO 4 .7, 6, 5, 2, and 1H 2 O; CdSO 4 .4H 2 O, 
also 1H 2 O, and 3CdSO 4 8H 2 O. 3VtgSeO 4 .7H 2 O ; ZnSeO 4 7, 6, and 2H 2 O , 
CdSeO 4 .2H 2 O. MgrTe0 4 ; CdTe0 4 . 

These salts are all easily soluble in water, except magnesium 
and cadmium tellurates, which are produced by precipitation from 
concentrated solutions. Magnesium sulphate, as Epsom salt. 
MgSO 4 7H 2 O, and as kieserite, MgSO 4 .H 2 O, occurs native in 
caves in magnesium limestone and in the salt-beds of Stassfurth. 
It is a frequent constituent of mineral waters. It has a bitter 
taste and a purgative action. It is made on the large scale by 
treating dolomite, a carbonate of magnesia and calcium, or serpentine, 
a hydrated silicate, with sulphuric acid. The hepta-hydrated sul- 
phates and selenates of magnesium and zinc are isomorphous, and 
crystallise in four-sided right rhombic prisms. When heated to 
150 they lose 6H 2 0, but retain the seventh molecule even at 200. 
Anhydrous magnesjum sulphate melts at a red heat ; the cadmium 
and zinc salts lose trioxide, leaving the oxides. These salts are 
insoluble in alcohol. Zinc sulphate, digested with hydroxide, 


yields several basic sulphates : SO 3 .2ZnO ; SO 3 .4ZnO.10 and 
2H 2 O; SO 3 .62nO.10H 2 O; and SO 3 .8ZnO.2H 2 (X Wifch excep- 
tion of the first, they are crystalline bodies. 

Mixed salts. < 

H 2 Mgr(S0 4 ) 2 ; H 6 Mgr(S0 4 > ) 4 ; Na2Mff(S0 4 ) 2 .6H 2 0; 
K2Mff(SO 4 ) 2 .6H 2 O , KoCa^Mff (SO 4 ) 4 .2H 2 O ; 

Na2Zn(S0 4 ) 2 .4H 2 O ; (NH 4 ) 2 Zn(SO 4 ) 2 .6H 2 O ; 
Mg-Zn(S0 4 ) 2 14H 2 ; Na 2 Cd(SO 4 ) 2 .6H 2 O, 

and other similar salts. These are all soluble, and are prepared by 

B 2 (S0 4 ) 3 .H 2 : Sc 2 (S0 4 ) 3 .6H 2 ; Y 2 (SO 4 ) 3 .8H 2 O; La 2 (SO 4 ) 3 .9H 2 O. 

Seleuates and tellurates have not been prepared. Boron sulphate 
is a white mass, produced by evaporating boron trioxide with sul- 
phuric acid.* It is decomposed by water. The other salts of the 
group are white and crystalline. 

Double salts. 

H 3 B(S0 4 ) 3 ; (NH 4 )Sc(S0 4 ) 2 ; K 4 Sc 2 (SO 4 ) 6 ; Na 3 Sc(SO 4 ) 3 .6H 2 O ; 

j Na 3 Y(S0 4 ) 3 .2H 2 0; (NH 4 )LtHSO 4 ) 2 4H 2 O; K 3 La(SO 4 ) 3 . 

I'hese salts are sparingly soluble, and are produced by mixture. 

Al 2 (S0 4 ) 3 .18H 2 Oj Oa 2 (S0 4 ) 3 ; In 2 (SO 4 ) 3 .9H 2 O ; T1 2 (SO 4 ) 3 .7H 2 O. 
Al 2 (Se0 4 ) 3 .wH 2 0; and the thallous salts T1 2 SO 4 ; HT1S0 4 ; and Tl 2 SeO 4 . 

Sulphate of aluminium, containing 18H 2 0, occurs native as 
alunogen, or feather alum ; it forms delicate fibrous masses or 
crnsts. It is known in commerce as " concentrated alum/' and 
is prepared by heating finely ground clay with strong sulphuric 
acid until tjie latter begins to volatilise. After lying some days, 
it is treated with water ; the solution is freed from iron by pre- 
cipitating it as ferrocyanide, or by addition of certain peroxides, 
such as those of lead or manganese ; it is then evaporated to 
dryness and fused. It crystallises with difficulty, being exceed- 
ingly soluble (1 in 2 parts of water) ; its crystallisation may be 
furthered by addition of alcohol, in which it is insoluble. Basic 
salts are known, produced by the action of hydrated alumina on 
the ordinary aulphate, by incomplete precipitation with ammonia, 
or by the action of zinc on a solution of ordinary sulphate. These 
are said to have the formulae 3SO 3 .2A1 2 O 3 ; 3SO 3 .3A1 2 O 3 .9H 2 O 
(occurring native as aluminite) ; 3SO 3 .4A1 2 O 3 .36H 2 O ; and 
3SO 3 .5A1 2 O 3 .20H 2 O. The selenate closely resembles the sulphate, 

* J. Prakt. Chem. (2), 38, 118. 


and yields corresponding basic salts. The tellurate is a white pre- 
cipitate. Gallium sulphate, Ga^SO^g, is very soluble, and crystal- 
lises in nacreous scales ; indium sulphate has been obtained only as 
a gummy mass ; and thallic sulphate forms thin colourless laminae, 
which are decomposed by water into the hydrated trioxide and sul- 
phuric acid. 

Thallous sulphate and selenate crystallise in anhydrous rhom- 
bic prisms isomorphous with potassium sulphate. They are soluble 
in water. They establish a link between the aluminium and the 
potassium groups. 

Double salts. The alums. These bodies are very numerous. 
They all crystallise in regular octahedra, are soluble in water, and 
have the general formula M / M'"RO 4 .12H 2 O, where M' stands for 
lithium, sodium, potassium, rubidium, caesium, ammonium, thal- 
lium (as a thallous compound), or silver; M'" for aluminium, 
gallium, indium, chromium, ferric iron, manganic manganese, or 
cobaltic cobalt;* and R for sulphur or selenium. Tellurium 
alums do not seem to have been prepared. The number of possible 
different alums is therefore 96 ; of these some 25 have been pre- 
pared. Alums containing aluminium, gallium, and indium are 
colourless ; chromium alums are very deep purple ; iron alums, 
pink ; and manganese alums, brownish-red. As they are all iso- 
morphous, they crystallise together. For example, an alum con- 
taining aluminium and potassium placed as. a nucleus in a solution 
of chromium alum becomes covered with a regular deposit of the 
latter, and a coating of iron alum may be deposited on the 

Alums are prepared by mixing solutions of the sulphates or 
selenates of the two metals, and crystallising. The most important 
are potassium aluminium sulphate, and ammonium alumi- 
nium sulphate, KA1(SO 4 ) 2 .12H 2 O, and NHiAKSO^ 1'2H 2 O. 
Ammonium alum, which also occurs native as tclierniigite,^ prepared 
by mixture ; 100 parts of water dissolve 5*22 parts at 0, and 421*9 
parts at 100. Potassium alum is prepared on a very large scale 
by calcining aluminous schists, which are essentially impure sili- 
cate of aluminium containing quantities of iron pyrites and car- 
bonaceous matter. The pyrites on ignition forms ferrous sulphate, 
FeSO^ and free sulphuric acid. The ignited mineral is methodi- 
cally extracted with water, and the liquors are concentrated in 
leaden pans, giving an acid solution of aluminium sulphate con- 
taining ferrous or ferric sulphates. To this liquor, a concentrated 

* Proc. Roy. Soc. Ed., 183, 203, 


solution of potassium chloride is added. It is preferable to tlie 
sulphate, for it forms, with the iron sulphate, uncrystnllisable 
ferric chloride along 1 with potassium sulphate. After settling, it 
is run into coolers to "crystallise. The confused crystals which 
separate are washed, drained, dissolved in fresh water, and re- 
crystallised in casks. It is sometimes freed from iron before the 
second crystallisation by one of the methods already described 
(p. 424). 

The chief nse of alnm is as a mordant in dyeing; the sulphato 
and acetate of aluminium are used for the same purpose. When 
cloth or any mineral or vegetable fibre is boiled in such a solution, 
it becomes impregnated with hydrated alumina ; and when treated 
with a dye, a triple combination appears to take place between the 
fibre, the alumina, and the colouring matter. 

Some basic sulphates of aluminium occur native. These are 
alunite, 4SO 3 .KjO^Al 2 O H 3EL>O, found at Tolfa, near Civita 
Vecchra, at Solfatara, near Naples, and at Puy de Garcy, in the 
Auvergne. It forms rhomboheclral crystals, and is used for the 
manufacture of Roman alum, which has been prepared from it 
from very early times. When it is calcined at a moderate heat, 
the hydrated alumina loses watei*, and on lixivtation, alum dis- 
solves, and may bo crystallised as usual . The basic sulphate, lowigite, 
4SO. V K 2 O.3AL 2 O 3 .H 2 O, is also a natural product. 

The difference of solubility of potassium alum from that of ru- 
bidium and caesium alums has afforded a means of separating from 
each other these elements, which almost always occur together. 
Rubidium and ccesium alums are insoluble in a cold saturated solu- 
tion of potassium alum ; hence, on concentrating such a mixture, 
the first portions of the crystals consist chiefly of the former. 
Caesium alum is likewise insoluble in a saturated solution of 
rubidium alum, and may be separated from the latter in a similar 

Mn(SO 4 ) 2 . Produced by dissolving potassium permanganate, 
KMnO 4 , in a mixture of 500 grams of sulphuric acid arid 150 of 
water. It is a yellow substance, which deposits a basic sulphate 
as a black powder of the formula MnO.SOi. 

Cr a (SO) 3 .15 and 5 H 2 O; Fe 2 (SO 4 ) 3 .9H,O; Mn,(SO 4 ),. There 
are two hydrated varieties of chromium sulphate, a green and a 
violet. The green salt is produced when the sulphate is pro- 
duced by the ordinary methods above 50, or by heating the violet 
variety to that temperature ; it is soluble in alcohol. The violet 
variety is produced in the cold ; it is also formed when the green 
modification is allowed to stand. It is precipitated by alcohol, 


and crystallises best from a mixture of alcohol and water. On 
heating either variety with excess of sulphuric acid to above 190, 
a light yellow mass of anhydrous sulphate is obtained, insoluble in 
waj;er, and with difficulty in acids. Several basic salts are known, 
produced by digesting a solution of the ordinary salt with 
chromium hydrate, or by incomplete precipitation. Among theso 
are 2SO 3 .Cr 3 O 3 ; 2SO 3 .3Cr 2 O 3 ; and 3SO 3 . 2CrO d . They are 
insoluble and amorphous. Ferric sulphate, Fe 2 (SO 4 ) 3 .9H 2 O, 
seems native as coquimbite. It is produced by oxidising ferrous 
sulphate with nitric acid in presence of strong sulphuric acid : 
2FeSO 4 + H 2 S0 4 + O = Fe,(SO 4 ) 3 4- H 2 O. It forms small pink 
scales, and is very difficult to dissolve in water. Manganic sulphate 
is a non-crystalline green substance produced by heating the 
hjdrated dioxide with sulphuric acid. Many basic sulphates of 
iron and manganese are known, which resemble those of chromium. 
The double salts of these oxides, or alums, have already been 
noticed. A sulphato-nitrate of chromium, Cr 2 (SO 4 )(NO<)4is pro- 
duced by dissolving the hydrated basic sulphate, Cr 2 (SOi)(OH) 4 . 
in strong nitric acid. The salt Cr 2 (SO 4 )2(NO 3 )2 is also known. 

CrSO 4 .Aq; FeSO 4 .7, 5, 3, 2, and 1H 2 O ; MnSO 4 .7, 6, 5, 4, and 2H 2 O; 

CoSO 4 .7, 6, and 4H 2 O ; NiSO 4 .7 and 6H 2 O. 
PeSeO 4 .7 and 5H 2 O ; CoSeO 4 .7H 2 O; Ni 2 SeO 4 .7 and 6H 2 O. 
FeTeO 4 ; MnTeO 4 ; CoTeO 4 ; NiTeO 4 . 

Chromous sulphate has been obtained as a blue solution, by 
dissolving the metal in dilute acid. Like all chromous salts, it 
has powerful reducing properties. Ferrous sulphate occurs native 
as green vitriol or copperas, produced by the atmospheric oxidation 
of iron pyrites. It usually crystallises with 7H 2 0, in light-green 
monoclinic crystals, which absorb oxygen slowly in moist air, 
forming a basic ferric sulphate (said to be 2(SO 3 .Fe,O-j).H,O), 
but in dry air they are permanent. When heated to redness it 
evolves sulphur dioxide, and a basic sulphate remains, which, on 
further heating, leaves a residue of ferric oxide, and yields a dis- 
tillate of sulphur trioxide. This residue is named rouge, and used 
to be known as " colcot har vitrioli," or " caput mortuum" ; it is used 
as a pigment. Ferrous sulphate has been obtained with different 
amounts of water, according to the temperature at which it is 
crystalised ; the hydrates with 3 and 2H 2 are formed in presence 
of sulphuric acid. That with 1H 3 is produced by drying the 
salt at 114; the last molecule is retained at 280, and is some- 
times termed " water of constitution.*' Ferrous sulphate absorbs 
nitrio oxide (see p. 342) ; but the composition of the resulting 


compound depends on the pressure and temperature, varying from 
3FeSO 4 .2NO to 6PeS0 4 .2NO. Manganous sulphate is a pink 
salt ; cobalt sulphate rose-red, and. nickel sulphate grass-green. 
The anhydrous salts are colourless. The hydrated sulphates, of 
these metals, containing the same number of molecules of water of 
crystallisation are isomorphous with each other; those with 
5H 2 resemble copper sulphate, CuSO 4 .5H 2 O, in crystalline 

The selenates closely resemble the sulphates ; the tellurates are 
insoluble precipitates. PeTeO 4 occurs native, and has been named 

A large number of double salts of the general formula, 
M' 2 EO4.M"R0 4 .6H 2 O, are known, where M' stands for Li, Na, K, 
Rb, Cs, IT and NH 4 ; M", for Mg, Zn, Cd, Cr", Fe", Mn", Co", 
Ni", and Cu" ; and R for S or Se. They all crystallise in mono- 
clinic crystals, and are isomorphous with each other. They are 
produced by mixture. The double salts of hydrogen, 

H 2 Mn(SO 4 ) 2 , and H 6 Mn(SO 4 ) 4 
have also been prepared. 

Sulphate of carbon is unknown. Both the monoxide and 
dioxide of carbon are insoluble in sulphuric acid. 

Ti 2 (S0 4 ) 3 ; Ce 2 (S0 4 ) 3 .5, 6, 8, 9, and 12H 2 O. 
Double salts: Ce 2 (SO 4 ) 3 . 2K 2 SO 4 . 2H 2 O; Ce 2 (SO 4 ) 3 .5K2SeO 4 , and others. 

The titanous sulphate is violet ; the cerous salts colourless. 

Ti(SO 4 ) 2 ; Zr(SO 4 ) 2 ; Ce(SO 4 ) 2 .4H 2 O ; Th(SO 4 ) 2 .4H 2 O. Also double salts, 
such as KoTi(S0 4 ) 3 ; (NH 4 ) 6 Ce(SO 4 V4H 2 O ; K 4 Th(SO 4 ) 4 .2H 2 O. 

The cerium salt is yellow ; the others colourless. Cerium also 
forms a double salt, containing the metal in two states of oxida- 
tion ; it is called ceroso-ceric sulphate. It has a brown-red colour 
and the formula 2Ce(SO 4 ) 2 .Ce 2 (SO 4 ) 3 .25H 2 O. These bodies, 
especially titanium, zirconium, and cerium, also yield basic sul- 
phates. The formation of titanium sulphate serves as a means 
of separating titanium from silica. The mixture is fused with 
hydrogen potassium sulphate, dissolved in water, and filtered from 
silica; on boiling with water the titanium sulphate is decomposed 
into hydrate and sulphuric acid. 

Silica is insoluble in sulphuric acid ; and germanium does not 
appear to form a sulphate. 


8nSO 4 ; PbSO 4 ; PbSeO 4 ; PbTeO 4 . Double salts : K 2 Sn(S0 4 ) 2 ; 
(NH 4 ) 2 Pb(SO 4 ) 2 . 

^Stannous snlphate is colourless and crystalline. The double 
salts are obtained by mixture. Lead sulphate occurs native in 
trimetric crystals as anglewte, isomorphous with those of heavy 
spar (barium sulphate). The crystalline variety may be obtained 
by fusing lead chloride with potassium sulphate The selenate 
has also been found native. As lead sulphate and selenate are 
nearly insoluble, they may be produced by precipitation; they 
form dense white powders, more easily dissolved by water than by 
the dilute acid ; but they are soluble to a small extent in strong 
acids. They dissolve in larger quantity in solutions of sulphate, 
nitrate, acetate, or tartrate of ammonium, and easily in caustic 
alkali, and in thiosulphates. Lead sulphate also dissolves in sul- 
phuric acid ; the solution deposits crystals of H 2 Pb(SO4) 2 .H 2 O. 
These bodies melt at a red heat. 

Lead sulphate, heated with the sulphide, as in lead smelting, 
yields metallic lead and sulphur dioxide, thus : PbSO 4 -|-PbS = 
2Pb + 20 2 ; or the oxide and metal : 2PbSO 4 + PbS = 30 2 + 
2PbO + Pb. 

Lead tellurate is also a white precipitate, but is more easily 
soluble in water. Basic sulphates and selenates of tin and lead 
have also been prepared; stannic hydrate dissolves in sul- 
phuric acid, but stannic sulphate is an indefinite non-crystalline 

Compounds of nitrogen and vanadium usually contain the 
nitrosyl, or vanadyl groups. Compounds of the pentoxides with 
sulphuric anhydride are, however, known. The compound, 
SO 3 .NoO 6 .4H 2 SO 4 , a white crystalline body, is produced by cool- 
ing a mixture of sulphur trioxide and nitric acid; it is at once 
decomposed by water, and, when heated, evolves red fumes, yielding 
a sublimate supposed to be SO 3 N 2 O 3 . This would be nitrosyl 
sulphate, SO 2 (ONO) 2 , to be alluded to later. The first may be 
viewed as a compound of nitryl sulphate, S0 4 (N0 2 ) 2 with sul- 
phuric acid. The compound, 2(SO 3 ).N 2 O 5 , is also known. It is a 
snowy crystalline mass, produced by the action of induction sparks 
on a mixture of sulphur dioxide, oxygen and nitrogen ; it may be 
regarded as nitryl anhydrosulphate, S 2 7 (N0 2 )4. 

Vanadyl sulphate, 3SO 3 .V 2 O 5 , is prepared by dissolving 
vanadium pentoxide in cold sulphuric acid, and expelling excess 
of sulphuric acid by heat. It may be regarded as (VO)"' 2 (SO 4 ) 3 . 


It is red and crystalline. During evaporation, the green compound 
of V 2 O 4 , 2SO 3 .V 2 O 4 = (VO)" 2 (SO 4 ) 2 separates as a crust. By 
1) eating the first compound to the temperature of melting lead, 
the basic sulphate, (VO) 2 O.(SO 4 ) 2 , is obtained as a red crystalline 
mass. A double sulphate of the formula, 2803.1^0 .V 2 O 6 .6Bf 2 O, 
is also known. 

These bodies are mostly derivatives of the pentoxides of 
nitrogen and vanadium. Niobium and tantalum are said also to 
form sulphates, but these compounds have not been investigated. 

Nitrosyl sulphate, (NO) 2 S0 4 , may be the substance alluded to on 
the previous page. Hydrogen nitrosyl sulphate, H(NO)SO 4 , is 
better known, and is produced by the action of nitrogen trioxide 
on sulphuric acid, thus : N 2 3 + 2H 2 S0 4 = 2H(NO)SO 4 + H 2 O. 
Excess of sulphuric acid must be present to combine with the water. 
The same substance is produced by the action of sulphur dioxide 
on nitric acid, or by passing the vapours from a heated mixture of 
nitric and hydrochloric acids (nitrosyl chloride and chlorine, see 
p. 341) through strong sulphuric acid. It forms long, thin, trans- 
parent crystals melting at 85-87. It is the substance known as 
" chamber crystals,'* and its solution in sulphuric acid is produced 
in the " Gay-Lussac tower," in which the escaping gases from the 
vitriol chambers are brought in contact with strong sulphuric 
acid. On treatment with water, it is at once decomposed into oxides 
of nitrogen (NO and NOz + NzO*) ; this change takes place in the 
" Glover tower," where the sulphuric acid containing hydrogen 
nitrosyl sulphate is diluted ; the oxides of nitrogen are liberated, 
and again pass into the chambers (see p 416). (See also nitrosyl 
anhydrosulphate, p. 434). 

(PO)'" 2 (SO 4 ) 3 , phosphoryl sulphate, is produced by mixture ; it 
forms thin transparent scales, and is decomposed at 30, and by 
water; the corresponding compounds of arsenic, antimony, and 
bismuth are unknown, the groups (AsO)', (SbO)' and (BiO)' 
tending, as a rule, to replace only one atom of hydrogen. 

By dissolving arsenious oxide, As 4 O 6 , in sulphuric acid* of 
different concentrations, which must not, however, be more dilute 
than corresponds with the formula, H 2 SO 4 .H 2 O, various white 
crystalline sulphates of a,rsenic have been obtained. They appear 
to have the formula? 8(SO 3 ).As2O 3 ; 4(SO 3 ).As,O 3 ; 3(SO 3 ).As 2 O 3 (?) ; 
2(SO 3 ).As 2 O 3 ; and SO 3 .As 2 O 3 , The body, 3(SO 3 ).Aa,O 3 , would 
correspond to AS2(SO 4 ) 3 ; 2(SO 3 )As 2 O 3 may be written 
SO 4 As O As SO 4 ; and SO 3 Aa 2 O 3 may represent arsenosyl 
bulpLate, (AsO/ 2 SO 4 , corresponding in formula to nitrosyl .sul- 


phate. These bodies are all decomposed by water, and are all 
very unstable. 

The sulphates of antimony are similar but more stable. The 
compounds, 4(SO 3 ).Sb 2 O 3 , 3(SO 3 ).Sb 2 O 3 , 2(SO 3 ).Sb 2 O 3 , and 
SO 3 .Sb 2 O 3 , have been prepared. The normal salt, Sb 2 (SO 4 ) 3 = 
3(SO 3 ).Sb 2 Oj, is produced by boiling antimony with strong sul- 
phuric acid. It crystallises in needles. 

With bismuth, the compounds, 3(SOj).BiOj,2(SO 3 ).Bi 2 O3, and 
SO 3 .Bi 2 O 3 , are known. Bismuth dissolves in hot, strong sul- 
phuric acid, with evolution of sulphur dioxide forming the first; it 
is decomposed by water, giving the third. Double salts with 
hydrogen, HBi(SO 4 ) 2 .H,O ; with ammonium,NH 4 Bi(SO 4 ) 2 .4H 2 O ; 
and with potassium. K,jBi(SO4) 3 are also known. The selenates 
and tellurates have scarcely been examined. Bismuth tellurate, 
however, has been found native. Its formula is TeOj.Bi 2 O,j ; it 
has been named inontanite. 

Hydrated molybdenum sesquioxide forms a dark-coloured solu- 
tion with sulphuric acid, which may contain Mo 2 (S0 4 ) 3 . The di- 
oxide gives a red solution, supposed to contain Mo(S0 4 ) 2 . 

Uranou3 sulphate, U(SO 4 ) 2 .4 and 8H 2 O, forms green crystals, 
and is produced bv dissolving hydrated uranium dioxide in 
sulphuric acid. A basic sulphate, SOs.UOs.ttHaO, is also known ; 
and also the double salt K 2 U(SO 4 ) 3 .H,O. They are green, 
soluble bodies. 

MoO 2 (SO 4 ) and UO 2 (SO 4 ), molybdyl and uranyl sulphates, 
are yellow crystalline bodies, obtained from the hydrated trjoxides. 
This sulphate of molybdenum, when boiled with water, decomposes, 
depositing the hydrated oxide, 5(MoO.j) ETO. Double salts of 
uranyl sulphate are known. e.g>, H,(UO 2 )(SO 4 ) 2 , and 

K 2 (UO 2 )(SO 4 ) 2 .2H,O 
The selenates and tellurates are little known. 

Tellurium dioxide dissolves in hot dilute sulphuric acid, and 
deposits crystals uf SO^.2TeO 2 . It is decomposed by warm water. 

Ru(SO 4 ) 2 ; B,h2(SO 4 ) 3 .12H 2 O; PdSO 4 .2H 2 O. Also KBh(SO 4 ) 2 . 
Ruthenium and rhodium sulphates are orange-brown and red 
solutions, drying respectively to a yellow-brown amorphous mass, 
and to a brick-red powder; they are produced by oxidatio-n of the 
sulphide. Palladium dissolves in sulphuric acid, mixed with a 
little nitric acid ; the solution, when evaporated, deposits brown 

OsSO 4 j Os(S0 4 ) 2 ; IrS0 4 ; IrO.S0 4 ; PtS0 4 ; Pt(SO 4 ) . 


These are all yellow syrups, drying to brown non-crystalline 
masses; they are all produced by oxidising the respective sulphides 
with nitric acid, with the exception of platinous sulphate, PtSO4, 
which is produced when the chloride, PtCl 2 , is dissolved in 
sulphuric acid. 

Afir 2 SO 4 ; HA*S0 4 ; H 3 Agr(SO 4 ) 2 .H 2 O; H 6 Ag 2 (SO 4 ) 4 ; Hgr 2 SO 4 ; Hgr 2 SeO 4 ; 
Agv>TeO 4 ; cuprous and wurous sulphates are unknown. Auric sulphate, 
however, can be prepared in solution by dissolving auric oxide in dilute 
acid. It decomposes on standing. 

Sulphates of silver and mercury are sparingly soluble white 
salts, produced by precipitation, or by dissolving the metals in 
sulphuric acid. The silver salt is isomorphous with anhydrous 
sodium sulphate. The tellurate is a dark-yellow powder. It has 
been found native, and named magnolite. 

CuSO 4 .5H 2 O ; CuSeO 4 .5H 2 O ; HgrSO 4 . Basic salts : 8O 3 .2CuO.H 2 ; 
SO 3 .3CuO.3H 2 O; SO 3 .4CuO.3H 2 O; SO (J .3HgrO. Double salts: Those 
of copper belong to the class M 2 / M' / (SO 4 ) 2 .6H 2 O ; those of mercury re- 
semble 3K2Hff(S0 4 ) 2 .2H 2 0. Also HgSO 4 .HgI 2 ; 2HgSO 4 .H^S. 

Copper sulphate, or blue vitriol, is produced on a large scale by 
the spontaneous oxidation of copper pyrites, or by the action of 
air on ignited cuprous sulphide, Cu^S, whereby cupric oxide is 
produced at the same time. It crystallises with water in large 
blue monoclinic prisms, isomorphous with ferrous sulphate of the 
same degree of hydration. Indeed, copper sulphate, if present in 
excess in a solution containing ferrous sulphate, induces the latter 
to adopt its crystalline form ; and, similarly, ferrous, zinc, mag- 
nesium, or nickel sulphate in excess, causes copper sulphate to 
assume their special form. When heated to 100, CuSO4.5H 2 O 
loses four molecules of water ; the last molecule is retained 
up to 200, and is regarded as " water of constitution." It is 
easily soluble in water, but insoluble in alcohol. The tetra- 
basic salt occurs native as brochantite. The selenate closely re- 
sembles the sulphate. Mercuric sulphate is decomposed by water 
into a soluble acid salt, 3SOj.HgO.nH.jO, and the basic salt, 
SOj.SHgO, a lemon -yellow powder, which used to be called tur- 
peth mineral. The compound, 2HgSO 4 .HgS, is precipitated by 
the action of a moderate quantity of hydrogen sulphide on a solu- 
tion of the sulphate. It is a white precipitate. 

Anhydro- or pyrosulphuric acid, H,S 2 O 7 . This substance 
is, as will appear hereafter, an analogue of pyrophosphoric acid, 


inasmuch as it may be regarded as constituted of two molecules of 
sulphuric acid, minus a molecule of water, thus . 

HO (SO 8 ) (SO,) OH. 

But it cannot be prepared by heating ordinary sulphuric 
acid, for that acid, as already remarked, distils as a whole. 
It may be obtained by dissolving sulphur trioxide in ordinary 
sulphuric acid, thus : H a SO 4 4- SO., = H 2 S 2 O 7 . The old method 
of preparation, which gained for this acid the name "Nord- 
hausen sulphuric acid," is still carried out at Nordhauseii in 
Saxony; it consists in distilling partially dried ferrous sulphate 
from ' tube-shaped retorts of very refractory fire-clay. The 
products are sulphur dioxide and anhydrosulphuric acid, while 
ferric oxide of a tine red colour remains in the retort, and is made 
use of as a pigment under the name of "Venetian ivd " or 
" rouge." This method of manufacture is a very ancient one. 
When ferrous sulphate, FeSO 4 .7H 2 O, is- dried, it lopes six mole- 
cules of water, retaining the seventh. On distilling the mono- 
hydrated salt, sulphur dioxide and water are evolved first, leaving 
basic ferric sulphate, thus : 

2FeSO 4 .H 2 O = S0 3 + 11,0 + Fe a O,(SO 4 )(= SO 3 .Fe,O.O- 
The sulphur dioxide escapes ; the temperature is then raised, 
when sulphur trioxide distils over, arid combines with the water, 
leaving iron sesquioxide in the retort. 

Anhydrosulphuric acid is a white solid, crystallising in needles, 
and melting at 35. It gives off sulphur trioxide when heated. It 
hisses when dropped into water, evolving groat heat. 

It is probable that still more condensed sulphuric acids are 
formed when more sulphur trioxide is added to sulphuric acid; 
but they have been little investigated. Corresponding compounds 
of selenium and tellurium are unknown. 

Pyrosulphates and polytellurates. The pyrosulphates are 
produced (1) by the action of pyrosulphuric acid on the oxides; 
(2) in a few cases by heating the double salts of hydrogen and a 
metal ; arid (3) by the action of sulphur trioxide on the normal 
sulphate. The following salts are known : 

NaoS 2 O 7 ; K 2 S 2 O- ; Ba. 2 S 2 O 7 ; AgvS.,O 7 ; also the double salt HKS 2 O 7 . 

The sodium and potassium salts may be prepared by all these 
methods. They are crystalline salts, which combine with water, 
forming hydrogen metallic sulpha! es. Hydrogen potassium pyrosul- 
phate crystallises from a solution of the anhydrosulphate in strong 
sulphuric acid; the other salts are best prepared by method (3). 

2 F 


Nitrosyl anhydrosulphate, S 2 7 (NO) 2 , is produced as a white 
crystalline substance by the action of sulphur dioxide on nitric 
peroxide. It is at once decomposed by water into sulphuric acid 
and the products of decomposition of nitrous anhydride. 

Several poly sulphates of arsenic, antimony, &c., have already 
been described among the sulphates. 

Di- and tetra- tell a rates are also known. The ditellurates prob- 
ably correspond to the anhydrosulphates ; and the tetratellurates 
are produced by the action of water on the monotellurates. The 
formulae of the following have been ascertained : 

EVTe. 2 7 ; (NH^TeoO, ; PbTe,O 7 ; Ag- 2 Te 2 O 7 ; also 4TeO 3 .K_O; 
4TeO a .(NH 4 ) 2 0; 4TeO 3 .BaO; 4TeO 3 .PbO; and 4TeO 3 .Agr 2 O. 

These bodies are more soluble than the ordinary tellurates. 




Compounds with Water and with Oxides (continued): 
(2) Compounds of the Dioxides ; Sulphurous, Selenioua, 
and Tellurous Acids; Sulphites, Selenites, and Tel- 

Sulphurous, selenious, and tellurous acids, in aqueous solu- 
tion, are produced either by direct combination of the anhydridc-H 
with water, or by displacement. 

Water absorbs at 15 about 45 times its volume of sulphur 
dioxide ; and on cooling the solution sevei^al definite hydrates have 
been obtained. 

By passing a current of the gas through a solution cooled to 
6, white crystals, fusing at 4, of the formula H 2 SO 3 .8H 3 O, 
were produced. By a similar process, crystals melting at 14, of