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QUALITATIVE
CHEMICAL ANALYSIS
A GUIDE IN CyjALITATIVE WORK, WITH DATA FOR
ANALYTICAL OPERATIONS AND LABORATORY
METHODS IN INORGANIC CHEMISTRY.
BY
ALBERT B. ^RESCOTT,
AND
OTIS c. Johnson,
PROFESSORS IN THE UNIVERSITY OK MICHIGAN.
FIFTH REVISED AND ENLARGED EDITION,
ENTIRELY REWRITTEN.
^^^^^^^
NEW YORK:
D. VAN NOSTRAND COMPANY
23 Murray and 27 Warren Sts.
IQOI
015075
Copjmghted 1901, by
D. VAN NOSTRAND COMPANY.
The Fribobwwald Company,
baltimore, md., u. s. a.
PREFACE.
In this, the fifth full revision of this manual, the text has been
rewritten and the order of statement in good part recast. The subject-
matter is enlarged by fully one-half, though but one hundred pages
have been added to the book.
It has been our aim to bring the varied resources of analysis within
reach, placing in order before the worker the leading characteristics of
elements, upon the relations of which every scheme of separation de-
pends. This is desired for the working chemist, and no less for the
working student. However limited may be the range of his work, we
would not contract his view to a single routine. It is while in the
course of qualitative analysis especially that the student is forming
his personal acquaintance with the facts of chemical change, and it is
not well that his outlook should be cut off by narrow routine at this
time.
The introductory pages upon Operations of Analysis, setting forth
some of the foundations of qualitative chemistry, consist of matter
restored and revised from the editions of 1874 and 1880. This subject-
matter, omitted in 1888, is now desired by teachers. For the portion
upon Solution and Ionization, we are indebted to Dr. Eugene C. Sulli-
van, a pupil of Professor Ostwald, now teaching qualitative analysis.
The pages upon the Periodic System have been added to afford a more
connected comparison of the elements than that undertaken in each
' group by itself, in previous editions, and referred to in the preface in
1874. The use of notation with negative bonds, in balancing equations
for changes of oxidation, introduced by one of the authors in 1880,
has been retained substantially as in the last edition. Other authors
adopt the same notation with various modifications. For the present
revision there has been a general search of literature, and authorities
are given for what is less commonly known or more deserving of further
iv PREFACE.
inquiry. . The number of citations is so large that to save room special
abbreviation is resorted to.
For convenient reference, on the part of teachers, students and
analysts using the book, the section for each element and each acid is
arranged in uniform divisions. For instance, in each section, solu-
bilities are given in paragraph 5, the action of alkalis in paragraph 6a,
the action of sulphur compounds in paragraph 6e, etc. In the para-
graph (9) for estimations it should be said, nothing more than a general
statement of methods is given, for the benefit of qualitative study, with-
out directions and specifications for quantitative work, in which, of
course, other books must be used.
The authors desire to say with the fullest appreciation that Perry
F. Trowbridge, instructor in Organic Chemistry in this University, has
performed a large amount of labor in this revision, collecting data from
original authorities, confirming their conclusions by his own experi-
ments, elaborating material, and making researches upon questions as
they have arisen.
University of Michigan,
April, 1901.
CONTENTS.
PART I.— THE PRINCIPLES OF ANALYTICAL CHEMISTRY.
PAGE
The Chbmical Elements and their Atomic Weights 1
Table of the Periodic System of the Chemical Elements 2
Discussion of the Periodic System » 3
Classification of the Metals as Bases 10
Commonly Occurring Acids 13
The Operations of Analysis 13
Solution and Ionization 20
Order of Laboratory Study 24
PART II.-THE METALS.
THE SILVER AND TIN AND COPPER GROUPS.
(FIRST AND SECOND GROUPS).
General Discussion 27
THE SILVER GROUP (FIRST GROUP).
fjead 29
Mercury 37
SUver 45
Comparisou of Certaiu Keactious of the Metals of the Silver
Group 51
Table for Analysis of the Silver or First Group 52
Directions for Analysis with Notes 53
THE TIN AND COPPER GROUP (SECOND GROUP).
THE TIN GROUP, OR SECOND GROUP, DIVISION A.
Arsenic 56
Antimony 72
Tin 82
Ck>inparisou of Certain Reactions of Arsenic, Antimony and Tin. 90
Ctold 91
Platinum 93
Molybdenum 97
THE COPPER GROUP, OR GROUP II, DIVISION B.
Bismuth 1 00
Copper 104
Cadmium 110
Comparison of Certain Reactions of Bismuth, Copper and Cad-
mium 112
Vi CONTENTS.
PAGE
The Precipitation op the Metals of the Second Group 113
Table for the Analysis of the Tin Group (Second Group, Division A). 116
Directions for Analysis with Notes 118
Tablb for Analysis op the Copper Group (Second Group, Division B). . 124
Directions for Analysis with Notes 126
RARER METALS OF THE TIN AND COPPER GROUP.
Ruthenium 139
Rhodium 180
Palladium 131
Iridium 132
Osmium 133
Tnn^ten 134
Vanadium 135
Germanium 136
Tellurium ! 137
Selenium 138
THE IRON AND ZINC GROUPS (Third and Fourth Groups) 140
THE IRON GROUP (THIRD GROUP).
Aluminum 142
Chromium 1 47
Iron 151
Table for Analysis of the Iron Group (Third Group) 160
Directions for Analysis with Notes 161
the zinc ghoup (fourth group).
Cobalt 163
Nickel 168
Mani^anese 172
Zinc 178
Comparison of Some Reactions of the Iron and Zinc Group
Bases 182
Table for the ANALY^iI8 of the Zinc Group (Foikth Group) 183
Directions for Analysis with Notes* 184
Analysis op Iron and Zinc Groups after Pkecipitation by Ammonium
Sulphide 186
Iron and Zinc Groups in Presence of Phosphates 188
Iron and Zinc Groups in Presence of Oxalates 189
Table of Separation of Iron, Zinc and Calcium Group Metals
and Phosphoric Acid by Means of Alkali Acetate and Ferric
Chloride 191
Table of Separation of Iron, Zinc and Calcium Group Metals
and Phospboric Acid by Means of Ferric Chloride and Barium
Carbonate 192
the rarer metals of the iron and zinc groups.
Cerium 193
Columbiuin (Viobinm) 193
Didymium 194
Erbium 195
CONTENTS. vii
Gallinm 195
Glacinam (Berylliam) 195
Indinm 196
lianthanam 197
Neodymium 197
Praseodymium 197
Samarium 197
Scandium 198
Tantalum 19s
Terbium 198
Thallium - 199
Thorium 199
Titanium 200
Uranium 201
Ytterbium 203
Yttrium 203
Zirconium 202
The Calcium Group (Fifth Gboup). (The Alkaline Earth Metals) 203
Barium 205
Strontium 208
Calcium '. 310
Ma^rnesium 314
Table for the Analysis of the Calcium Group (Fifth Group) 217
Direction* for Analysis with Notes 218
Separation of Barium, Strontium, and Calcium by the Use of Alcohol 330
Alkaline Earth Metals as Phosphates 330
Alkaline Earth Metals as Oxalates 330
The Alkali Group (Sixth Group) 231
Potassium 233
Sodium 336
Ammonium 339
Caesium 233
Rubidium 234
liithinm 234
Directions for Analysis with Note:* 236
PART III.— THE NON-METALS.
Balancing of Equations 338
Hydrog:en 243
Boron 245
Boric Acid 245
Carbon •. 347
Aceiic Acid 249
Citric Acid 251
Tartaric Acid 253
Carbon Monoxide 254
Oxalic Acid 255
Carl)on Dioxide (Carbonates) 259
Vlll CONTENTS,
PAOB
Cyanogen 263
Hydrocyanic Acid 263
Hydroferrocyanic Acid 267
Hydroferricyanic Acid 269
Cyanic Acid 271
Thiocyanic Acid 272
Nitrogen 273
Hydronitric Acid 274
Nitrous Oxide 275
Nitric Oxide 275
Nitrons Acid * 276
Nitrogen Peroxide 277
Nitric Acid 277
Oxygen 282
Ozone 284
Hydrogen Peroxide 285
Fluorine 288
Hydrofluoric Acid 289
Fluosilicic Acid 289
Silicon 290
Silicic Acid 290
Phosphorus 292
Phosphine 295
Hypophosphorous Acid 295
Phosphorous Acid 297
Hypophosphorio Acid 298
Phosphoric Acid 298
Sulphur 304
Hydrosnlphurio Acid 806
Thiosulphuric Acid 312
Hyposulphurous Acid 314
Dithionio Acid 314
Trlthionio Acid 315
Tetrathionio Acid 315
Pentathionic Acid 316
Table of Thionio Acids 317
Sulphurous Acid 318
Sulphuric Acid 321
Persulphuric Acid 326
Chlorine 327
Hydrochloric Acid 830
Hypochlorous Acid 837
Chlorous Acid 337
Chlorine Peroxide 338
Chloric Acid 339
Perchloric Acid 341
Bromine 842
Hydrobromic Acid 845
Hypobromous Acid 348
CONTENTS. ix
PAGE
Bromic Acid 348
Iodine 350
Hydriodic Acid 858
Iodic Acid 357
Periodic Acid 360
COMPABATIVB REACTIONS OP THE HaLOOEN COMPOUNDS 361
PART IT.— SYSTEMATIC EXAMINATIONS.
Rkmotal op organic Substances 362
Pbeliminabt Examination op Solids 363
Conversion op Solids into I^iquids 366
CoNYE^siON op Solutions into Solids 367
Treatment op a Metal or an Allot 367
Separation op Acids prom Bases 368
Table por Preliminary Examination op Solids 870
Bkhatior op Substances Beporb the Blow-Pipe 374
Tablb op the Grouping op the Metals 375
Table por the Separation op the Metals 376
Acids — First Table 378
Acids — Second Table 386
Acids — Third Table 887
Acids— Fourth Table 388
notbs on the detection op acids 389
Principles : 398
E<2nATiONS 396
Problems in Synthesis 397
Table op Solubilities 398
Reagents 403
ABBREVIATIONS.
A
A. €h.
Am.
Am. S.
Arch. Pharm.
Am. Chem.
B.
Bl.
B.J.
Oomey.
C. N.
Ch. Z.
C. r.
C. C.
D.
Fehlinff.
Fresenlus.
G. O.
Gazzetta.
Gilb.
Gmelln-Krant.
J.
C.
pr.
Soc. Ind.
Anal.
Am. Soc.
J. Pharm.
Laden bur^.
M.
Phil. M&S'
Pogg.
Proc. Roy. Soc.
Pharm. J. Trans.
Ph. C.
Tr.
Watt's.
18ft8»
* Indicates continuance to the present time.
Liebig's Annalen. 1832*
Annales de Chimie et de Physique. 1789«
American Chemical Journal. 1879*
American Journal of Science. 1818»
Analyst. 1876*
Archives der Pharmacie. 1832*
American Chemist. 1870-77.
Berichte der Deutschen Chemischen Gesellschaft.
Bulletin de la Societe Chimique. 1859*
Berzelius Jahresbericht. 1832-51.
Comey's Dictionary of Solubilities. 1896.
Chemical News. 18fi0*
Chemiker Zeituujr. 1877*
Comptes Rendus des Seances de TAcad^mie des Sciences.
Chemiscbes Centralblatt. 1830*
Dingler'8 Polytechniscbe Journal. 1 820*
Dammer's Anorganischc Chemie. 1893*
Fehling's Handbuch der Chemie. 1871*
Fresenius: Qualitative Chemical Analysis.
Graham-Otto: Lehrbuch der anorj^auischen Chemie
Gazzetta chimica italiana. 1871*
Gilbert's Annalen der Physik und Chemie. 1799-1824.
Gmelin-Kraut: Handbuch der anorganischen Chemie. 1877.
Jahresbericht uber die Fortschritte der Chemie. 1847*
Journal of the Chemical Society. 1849*
Journal fiir praktische Chemie. 1834*
Journal of the Society of Chemical Industry. 1882*
Journal of Analytical Chemistry. 1887-1893.
Journal of the American Chemical Society. 1876*
Journal de Pharmacie et de Chimie. 1809*
Handworterbuch der Chemie. 1882-1895.
Monatshefte fur Chemie. 1880*
Menschutkin. Locke^x Trandathm, 1895.
Philosophical Magazine. 1798*
PoggendorflPs Annalen der Physik und Chemie. 1824-1877.
Proceedings of the Royal Society of London. 1832*
Pharmaceutical Journal and Transactions. 1841*
Pharmaceutische Centralhalle. 1859*
Transactions of the Royal Society. 1665*
Watt's Dictionary of Chemistry. 1888.
18.S.5*
Wells' Trans., 1897.
1885.
ABBREVIATIONS.
XI
W. A. Wiedemann's Annalen. 1877*
W. A. (Beibl.) Wiedemann's Annalen Beiblatter. 1877*
Wormley. Wormley's Microcbemistry of Poisons. 1867.
Wnrtz. Dictionnaire de Chimle. 1868.
Z. Zeitschrift fur analytiscbe Cbemle. 1863.*
Z. Ch. Zeitschrift fur Chemie. 1865-1871.
Z. anorsr. Zeitschrift fur anorganiscbe Cbemie. 1891*
Z. AD^ew. Zeitschrift fur angewandte Chemie. 1888*
Z. phys. €h. Zeitschrift fur pbysicaliscbe Chemie. 1887*
PART I.
THE PBIXCIPLES OF ANALYTICAL CHEMISTRY.
§1. The Chemical Elements and their Atomic Weights.f
Name.
Aluminum .
Antimony .
Argon
Arsenic
Barium . . . . .
Bismuth ...
Boron
Bromine
Cadmium . . ,
Cassium . . . ■
Calcium . . .
Carbon
Cerium
Chlorine . . . .
Chromium . .
Cobalt
Columbium .
Copper
Erbium
Fluorine . .
Gadolinium .
Gallium . . .
Germanium .
Glucinum . . .
Gold
Helium . . . .
Hydrogen . .
Indium . - . .
Iodine
Indium
Iron
Krypton . . . .
Lanthanum .
Lead
Lithium . . .
Magnesium .
Manganese .
Mercury . . .
Molybdenum
Sym-
601.
H-l.
0-16.
Al
26.9
27.1
Sb
119.5
120.4
At
40.?
40. ?
As
74.46
76.0
Ba
186.4
187.40
Bi
206.6
208.1
B
10.9
11.0
Br
79.34
79.95
Cd
111.55
112.4
Cs
181.9
182.9
Ca
89.8
40.1
C
11.9
12.0
Ce
138.0
189.0
CI
35.18
86.46
Cr
61.7
62.1
Co
68.55
69.00
Ob
98.0
98.7
Cu
68.1
63.6
Er
164.7
166.0
P
18.9
19.06
Gd
166.8
167.0
Ga
69.6
70.0
Oe
71.9
72.6
Gl
9.0
9.1
Au
196.7
197.2
He
4.?
4.?
H
1.00
1.008
In
118.1
114.0
I
125.89
126.86
Ir
191.7
198.1
Fe
65.5
65.9
Zr
59.?
69.?
T.a
137.6
188.6
Pb
205.36
206.92
Li
6.97
7.03
Mg
24.1
24.3
Mn
61.6
65.0
Hg
198.50
200.0
Mo
96.3
9G.0
Name.
Neodymium . . .
Neon
Nickel
Nitrogen
Osmium
Oxygen
Palladium
Phosphorus . . .
Platinum
Potassium . . . .
Praseodymium .
Rhodium
Rubidium . . .
Ruthenium . . . .
Samarium . . I . .
Scandium . . . .
Selenium
Silicon
Silver
Sodium
Strontium
Sulphur
Tantalum
Tellurium . . . .
Terbium
Thallium
Thorium
Thulium
Tin
Titanium
Tungsten
Uranium
Vanadium
Xenon
Ytterbium
Yttrium
Zinc
Zirconium
Sym-
bol.
H = l.
Nd
Ne
Ni
N
Os
O
Pd
P
Pt
K
Pr
Bh
Bb
Bu
Sm
So
Se
Si
Ag
Na
Sr
S
Ta
Te
Tr
Tl
Th
Tm
Sn
Ti
W
u
V
X
Yb
Y
Zn
Zr
142.6
20.?
68.25
18.98
189.6
15.88
106.2
30.76
193.4
38.82
139.4
102.2
84.76
100.9
149.2
48.8
78.6
28.2
107.11
22.88
86.95
31.83
181.6
126.6
16a8
202.61
230.8
169.4
118.1
47.8
182.6
287.8
51.0
?
171.9
88.3
64.9
89.7
O = 10.
143.6
20.?
68.70
14.04
191.0
« 16.000
107.0
31.0
194.9
89.11
140.5
108.0
85.4
101.7
150.8
44.1
79.2
28.4
107.92
28.05
87.60
32.07
182.8
127.6?
160.
204.15
232.6
170.7
119.0
48.16
184.
239.6
51.4
?
173.2
89.0
65.4
90.4
t Elflrhth Annual Report of the Committee on Atomic Weigrhts. F. W. Clarke, J. Am. Soc.^
19(0, SS, 90.
*The atomic wei^rbts used in this book are taken O— 16.
2 TABLE OF THE PERIODIC SYSTEM OF CHEMICAL ELEMENTS, §2.
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§3. DISCI SSl()\ OF THE PERIODIC SYfH'EM. 3
§3. In this system of the chemical elements certain regular gradations
of chemical character are to be studied and held in view, to simplify the
multitude of facts observed in analysis. Passing from Li 7.03 to F 19.05
in the first Series of this system, the elements are successively less and
less of the nature to constitute bases and more and more of the nature to
form acids, as their atomic weights increase. The acid-forming elements
are electro-negative to the elements which form bases.*
But in passing from 19.05 to the next higher atomic weight, Na 23.05,
we return from the acid extreme to the basal extreme and begin another
period, in gradation through the seven Groups. There is a like return
from one extreme to the other in the steps between chlorine and potass^ium
♦ Bases are the oxygen rompoundK of the metals. Acids are compounds of elements for
the most part not metals. In the chemical union of nodium with t-hlorine, for example,
these two elements differ widely from each other in their various properties. The <-hlorine
is the opposite of the sodium in that very p<iwer by virtue of which the one combines with
the other in the making of sodium chloride, a distinct product. In the polarity of electro-
lysis the sodium is the positive element, while the chlorine is the negative element. The
IKiwer of opposite action exercise<l by the one element upon the other, in their combination
together. Is represente<l by the opiK>slte polarity of the one in relation to the other during
electrolysis. Electrolysis is an exercine of the name energy that is (»therwise manifested
io chemical union or in a chemical change. Strictly speaking, it may be said that It is only
in electrical results that a positive or u negative polarity appears. Hut the term positive
{lolarity, applied to sodium because it g<»e8 to the negative pole of a battery, is u term
which well designates the opposlteness of the chemical a<'tion of sodium in its union with
chlorine. That is to say, the metals are in general *' positive," the not-metals in general
*• negative," in the relation of the former to the latter, and this relation may be termed
one of ** polarity," whether it appear in electrolysis, in chemical <'ombluatlon, or in a
ehemicai change.
In chemical combination, the atoms of each element a<t with a '* polarity." the extent
of which may be expresseii In terms of hydrogen eciuivaleuce or ** valen<*e." The valence of
an element, when in combination with another element, may be counted as relatively
•• positive " or •* negative " to the latter. For example. In the cc»mi)ound known as hydro-
sulpharic acid, the sulphur is negative, the hydrogen positive, in the relation of one to the
other, as represented by the diagram,
in which the plus and minus signs of mathematics arc used to represent the ** positive "
and *• negative " activities of chemical elements. That is, the sulphur acts with two units
of valence, both in negative polarity. In sulphuric acid the sulphur is positive in relation
to both the oxygen and the hydroxyl, as indicated In the diagram
(HO)- -4- ] +.
(HO)-+'
That Is, the sulphur acts with six units of valence, all in positive polarity. In respect to
oj^idation and reduction, the dlfferent'e between the action of sulphur In hydrosulphurlc
add on the one hand, and In sulphurl<* a<'ld nn the other hand, is a difference equivalent to
eight units of valence, the combining extent of t'lght atoms of hydrogen. This value is In
agreement with the factors of oxidizing a^rents In volumetric analysis.
In the same sense there is a change of ** polarity " eiiulvalent to the extent of eight units
of valence. In reducing periodic acid to hydrlodic acid, In redu<*lng arsenic add to arsine, or
In re<luclng carbon tetrachloride to methane. That is. In any of the groups from IV. to
VII. there N a difference, equivalent to the combining extent of eight hydrogen units, be-
tween the negative polarity of the element in its regular combination with hydrogen, such
as IVU,, aDd its positive i)olarIty in its highest combination with oxygen, suih as XO^ (OH).
4 DlSCUSSlOy OF THE PERIODIC SYS'TE}f, §4.
and- in those between bromine and rubidium. This fact of a periodic
return in the gradation of the properties of the elements, as their atomic-
weights ascend, constitutes a periodic system. A period is termed a Series.
A Group in this system consists of the corresponding members of all the
Series, which members are found to agree in valence, so that the number
of the groups, from I. to Vll. (not in YIIL), expresses the typical
valence of the elements as grouped. Further inquiry shows that all the
properties of the elements are in relation to their atomic weights, as they
appear in the periodic system. But this system is not to be depended upon
to give information of the facts; it is rather to be used as a compact simpli-
fication of facta found independently, by the student and by the author-
ities on whom the student must depend. A full account of the Periodic
System, as far as it is understood, is left to works on General Chemistry.
§4. The remarkable position of Group VIII., made up of three series,
each of three elements near each other in atomic weight, respectively in
Series 4, G, and 10, is in central relation to the entire system. In this
group there is something of a return, from negative to positive polarity,
from higher to lower valence. Group YIII. lies between Group VII. and
Group I., that is to say in this group there is a return from negative to
positive nature, and from higher to lower valence. Moreover, the newly
discovered elements related to argon, destitute of combining value as they
are, appear to constitute a Group 0. The latest results render this position
of the argon grou]) of elements so probable that it has been placed in the
chart for convenience of study, su])ject to further conclusions. (W. Ramsay.
Br, Assoc. Adv. Sci., 180:, oOS-GOl ; B. 1808, 31, 3111. J. L. Howe, C. N.,
1899, 80, 74; 1900, 82, 15, 52. Ostwald, Gntndr. AUg. Chem., 3te Auf.,
1899, S. 45.) In comparison with the members of Group Yll. those of
Group YIII. certainly have a diminished negative polarity, and a lower
valence, the latter being easily variable. Some of the particulars are given
below under the head, " Metals in Relation to Iron.'' The most remark-
able thing about Group T'lll. is the fact that the return to Group I. from
Group YIII. is less comy)lete than the return from Grou]) YII. That is to
say, the character of copper is divided between Group YIII. and Group I..
and the same is true of silver and of gold. This relation to Group YIII.
can be traced, in some particulars, to zinc and cadmium and mercury in
Group II. For these reasons Series 4 and 5 may be studied as one lonfj
period of seventeen members, Series G and 7 as another long period and
Series 10 and 11 as a third and final long period.
§5. It is to be observed that each one of the Groups, from I. to YII., falls
in two columns, a column consisting of fhe oHrniafe elemenls in the group.
Thus, H, Li, K, Eb and Cs make up the first column of Group I. It is
among the alternate members of a group that the closer grade-relations of
§9. DISCI S810X OF THE PERIODIC SYSTEM. 5
the elements are found. The gradations represented under one column
are distinct from those under the other in the same group. The well
known alternate elements of a Group, so far as found clearly graded
together in respect to given properties, are to be studied as a Family of
elements. Again a number of elements next each other in a Series are to
be studied together, either by themselves or with an adjoining half-group.
For the studies of analytical chemistry the following given are the more:
strongly marked of the families of the well known elements.
§6. The AMlLMetal8,—U 7.03, (Na 23.05), K 39.11, Eb 85.4, Cs 132.9.
The first part and sodium of the second part of Group I. In the grada-
tion of these elements the basal power increases qualitatively with the rise
in atomic weight. The hydroxides and nearly all salts of these metals are
freely soluble in water, wherein they are unlike the ordinary metals of all
the other groups. For the most part, however, these solubilities increase
with the atomic weight of the metal, and the carbonate and orthophosphate
of lithium are but slightly soluble.
§7. Th^ Alkaline Earth Metah.—(Kg 24.3), Ca 40.1, Sr 87.60, Ba 137.40.
These metals, like those of the alkalis, form stronger bases as they have
higher atomic weights. Both in Group I. and in Group II. the member
in Series 3 (Na, Mg), though in the second set of alternate members, agrees
in many ways with the next three of the first set of alternates. The
hydroxides of these metals are not freely soluble in water but are regularly
more soluble as the atomic weight of the metal is higher. The sulphides
are freely soluble; the carbonates and orthophosphates quite insoluble.
The sulphates have a graded solubility, decreasing as the atomic -weight
is higher, an order of gradation the reverse of that of the hydroxides and
of wider range. That is, at one extreme the magnesium sulphate is freely
soluble, at the other barium sulphate is insoluble.
§8. Tlie Zinc Family.— tUg 24.3, (Al 27.1), Zn 05.4, Cd 112.4, ,
Hg 200.0. These metals, save aluminum, belong to the second alternates of
Group IL, and, like those of the corresponding half of Group L, in their
gradation they are in general less strongly basal as they rise in their atomic
weights. Aluminum, here drawn in from Group III. second half, has the
valence of the third group, and differs from the others in not forming a
sulphide. The sulphide of magnesium is soluble, the sulphides of zinc,
cadmium and mercury insoluble in water, and these throe show this grada-
tion, that the zinc sulphide is the one dissolved by dilute acid, while the
mercury sulphide is the one requiring a special strong acid to dissolve it.
both these differences being depended upon in analysis. ^lercury, sepa-
rated from cadmium by two removes in the periodic order, is but a distant
member of this family.
§9. Metals in Relation to Iron.— Cr 52.1, Mn 55.0, Fe 55.9, Ni 58.70,
6 DiaCl'SSIoy OF THE PERIODIC SYSTEM. §10.
Co 59.00. The atomic weights of these metals lie nearly together. They
all belong to one Series, the fourth, representing Groups VI. and VII.,
and make the first of the instances of three members together in one series
in Group VIII. Chromium, being in the first division of its group, could
not be expected to grade with sulphur and selenium, nor would manganese
be expected to grade with chlorine and bromine, but the disparity is strik-
ing in both cases, especially in the comparison of melting points. The
valence of both chromium and manganese appears partly exceptional to
their positions in the system but the maximum valence of each is regular.
That all of these five elements, neighbors to chlorine and bromine, are
counted as metals, is not contrary to the periodic order. Group VIII. binds
Group I. to Group VII. After Co 59.00 follow Cu 63.G and then Zn 65.4.
Indeed each of " the well-known metals related to iron " is capable of serv-
ing as either a base or an acid, by change of valence. These metals are the
special subjects of oxidation and reduction. So far they resemble their
non-metallic neighbors, the halogens. Of the ^ve, chromium and man-
ganese (nearest the halogens) form the best known acids. Nickel and
cobalt, like cop])er, have a narrower range of valence, a more limited extent
of oxidation and reduction, within which they as readily act. These
valences, in capacity of combination with other elements, not including the
most unusual valences, may l)e written in symbols as follows:
2-3-6 2-3-4-6-7 2-3-6 2-3 2-3 1-2 2
Gr , Mn , ^e , Ni , Co , Cu , Zn
On reaching zinc, 65.4, in this gradation, the capacity of oxidation and
reduction disappears. Sulphides are formed by such of these metals as act
with a valence of two (all except chromium), and these sulphides are insolu-
ble in A'ater. In the conditions of precipitation sulphides are not formed
with the metal in any valence other than two. Chromium acting as a
base with a valence of three, like aluminum whose only valence is three,
refuses to unite with sulphur. Trivalent iron in ])reci])itation by sulphides
is mainly reduced to ferrous sulphide (FeS). In chromates the chromium
valence is reduced from six to three by hydrogen sulphide acting in solu-
tion. A carbonate is not formed by chromium, this being another agree-
ment with aluminum, and the same is true of trivalent iron.
§10. The Metals not Alhilis in Gronp /., Second Part, and their Relatives
in Group VIIL—Cu 63.6, Ag 107.92, , An 197.2. In gradation these
metals are less strongly basal, and more easily reduced from their com-
pounds to the metallic state, as their atomic weights rise. This is in agree-
ment with tlie gradation among the second set of alternates in Group II..
the Zinc Family. It likewise agrees with second part of Group VII., the
halogens. These elements of Group I. are to be studied with those of
Group VIII., especially with those respectively nearest them in atomic
§12. DISCUSSION OF THE PERIODIC SYSTEM. 7
weight: Cu 63.6 with Ni 58.70 and Co 59.00, Ag 107.92 with Pd 107.0, and
An 197.2 with Pt 194.9. Those with atomic weights above that of copper
rank as " noble metals," from their resistance to oxidation and other
qualities, so ranking in higher degree as their atomic weights increase.
Their melting points (those of Pd, Ag, Au, Pt) rise in the same gradation.
By action of ammonium hydroxide upon solutions of their salts these
(seven) metals form metal ammonium compounds, all of which are soluble
in water except the compounds of platinum and gold (highest in atomic
weight). All of the seven named form sulphides insoluble in water, in
condition of precipitation. For the most part their sulphides are relatively
more stable than their oxides. Silver differs from the others in the insolu-
bility of its chloride, and agrees irregularly in this fact, one prominent in
analysis, with mercury in its lower valence, and partly with lead.
§11. The Nitrogen Family of Elements.— It 14.04, P 31.0, As 75.0,
Sb 120.4, , Bi 208.1. The entire second part of Group V., and from
the first part the Leading Element of the group. Nitrogen and phosphorus
count as non-metals, antimony and bismuth as metals, arsenic as inter-
mediate, the polarity being more positive as the atomic weight increases.
In combinations with hydrogen, like ammonia and ammonium compounds,
phosphine and phosphonium salts, and also like analogous organic bases
where carbo-hydrogen takes the place of a part or all of the hydrogen, there
is a remarkable unity of type in this family. The same is true of the com-
binations with oxygen, like nitric acid. It is in Group V. that the group
valence for oxygen begins to diverge in gradation from the group valence
for hydrogen. In ammonium compounds nitrogen exercises a valence of
five, it doubtless is true, but this total of five units is always limited in
polarity to a balance of three negative units at most. In ammonia:
N - ^ HHH. In ammonium chloride: N - * + ^ = - ' • HHHHCl. Bismuth
is a distant member, a vacancy falling between it and antimony.
Phosphorus, arsenic and antimony are in gradation with each other as
to their indifference to chemical combination and readiness of reduction to
the elemental state, these qualities intensifying with the rise in atomic
weight. In this gradation nitrogen, belonging among the other alternate
members, has no part. In its chemical indifference it stands in extreme
contrast to phosphorus.
§12. Relniion of Tin and Lead to the Nitrogen Family. — These metals
are in Group IV., each combining both as dyad and tetrad, a valence dis-
tinctly unlike the valence of the nitrogen family, which is entirely regular
for Group Y. In Series 7: Sn 119.0, Sb 120.4. In Series 11: Pb 206.92,
Bi 208.1. The metals in the first named pair are two removes from those
in the second pair, all being among the second alternate members. In their
salts tin and antimony are more easily subject to changes of valence than
8 DISCUSmON OF THE PERIODIC SYSTEM. §12.
are lead and bismuth. In further comparison, arsenic, in its deportment
as a metal, may be included, making the list: As 75.0, Sb 120.4 (Sn 119.0),
Bi 208.1, (Pb 206.92). Of these, only arsenic fonns a higher oxide soluble in
water (separation after treatment with nitric acid and evaporation). Arsenic
and antimony form gaseous hydrides, in this agreeing with phosphorus and
nitrogen, the others do not. The stability of the hydrides of N, P, As, 8b,
all in the type of ammonia, is in the ratio inverse to that of the atomic
weight. All of these metals are precipitable as hydroxides save arsenic,
all are precipitated as sulphides, and these have chemical solubilities some-
what in gradation with atomic weights, the arsenic sulphide being most
fully separable by chemical solvents. The sparing solubility of the chloride
of lead, referred to in description of silver, is approached by the insolu-
bility of the oxy-chlorides of bismuth, tin, and antimony, and this fact
must be borne in mind, when precipitation by hydrochloric acid is employed
for separation of silver and univalent mercury in analysis.
Nitrogen in its trivalent union with hydrogen, the leading element of the
group of alkali metals, constitutes an active alkali. In its prevalent union
with oxygen, the leading element of Group VI., that is with oxygen and
hydroxyl, nitrogen forms an acid which is very active though not very
stable, its decomposition being represented by its gunpowder salt. The
degree of negative polarity of nitrogen, or its capacity for acid formation,
in accordance with its place next to oxygen among the atomic weights, is
shown in that singular instable body, hydronitric acid, HN3, of decided
acid power, constituting well marked salts, such as Na Ng, in which a ring
of nitrogen alone acts as an acid radical. The first four members of the
nitrogen family agree with each other in fonning trivalent and pentavalent
anhydrides and acids, the pentavalent ones being the more stable. The
pentavalent acids are of especial interest. In nitric acid the five units of
positive valence of an atom of nitrogen are met by two atoms of oxygen
with two units each of negative valence and a unit of negative valence
of hydroxyl: H — — W^q. The same constitution is found in metaphos-
phoric acid HO P O2 , meta-arsenic acid HO As O2 , and in antimonic acid
HOSbOs* The so-called ortho acids, phosphoric and arsenic, have the
constitution (H0)3 P and (H0)3 As , respectively. Phosphoric and
arsenic acids have a remarkable likeness to each other in nearly all the
properties of all their salts, behaving alike in analysis so long as preserved
from action of reducing agents. These sharply separate arsenic, usually in
one of its trivalent forms, AsHj or As^S,, . Antimony is reduced from its^
acid even more readily than is arsenic, in accordance with the gradation
stated above.
In the solubility of its metal salts the acid of nitrogen is, again, in
§14. Djscussroy of the periodic system. 9
strong contrast with the acids of the elements of the second part, phos-
phoric and arsenic acids. Metal nitrates are generally all soluble in water.
Of the metal phos))hate8 and arsenates, that is the full metallic salts of
phosphoric and arsenic acids, in their several forms, only those of the alkali
metals dissolve in water.
§13. The Halogens.— F 19.05, CI 35.45, Br 79.95, I 126.85. The lead-
ing element of Group VII., one of its first set of alternate members.
and the three known members of the second alternates. In the halogen
family fluorine has a relation like that of nitrogen in its family, taking
part in the group gradation as to polarity, solubility of compounds and
other qualities, but standing quite by itself in respect to certain properties.
It is the most strongly electro-negative of the known elements, a fact in
accord with the relation of its atomic weight.
For the common work of analysis we may confine our study of the
halogens to chlorine, bromine, and iodine. In the order of their atomic
weights, these elements appear, respectively, in gaseous, liquid, and solid
state, under common conditions. Their hydrogen acids, HCl , HBr , and
HI, show a stability in proportion to the electro-negative polarity of the
halogen, hydriodic acid being so unstable as to suffer decomposition in the
air. In the solubility of their metal salts these acids are nearly alike, all
being soluble except the silver, univalent mercury, and lead salts, but the
iodides of divalent mercury, bismuth and divalent palladium are sparingly
soluble. Each of these halogens, most especially iodine, forms a class of
salts each containing two metals, one of the united metals being that of an
alkali, such as (KI)2 Hglg and K^ Pt Cl^ . The periodides show that iodine
atoms have a power of uniting with each other, in the molecules of salts,
a power partly shared by bromine and chlorine and probably exercised in
many complex halogen compounds. By this means two atoms of a halogen
may serve the same as one atom of oxygen, in the linkings of molecular
structure.
Of the oxygen acids of chlorine, bromine and iodine, those in which the
halogen has a valence of five are more stable than the others. These acids
are chloric, HOClOg; bromic, HOBrO^; and iodic, HOIO^. Chloric acid
agrees with nitric acid, HO N 0^ , in the fact that it forms soluble salts \nth
all the metals. Chlorates decompose more violently than nitrates: iodates
for the most part less readily than the latter. Of the oxygen acids with
a halogen valence of seven, periodic acid, HO I O3 , also (H0)5 1 , is pre-
served intact without difficulty.
§14. Th^ Relations of Sulphur.— 8 32.07. Sulphur is the first member
of a family including selenium and tellurium. It differs from oxygen
almost as much as phosphorus differs from nitrogen, and we may say more
than silicon differs from carbon. The higher valence of Group VI., exer-
10 THE CLA881FJCAT10X OF THE METALS AS BASES. §16.
cised toward oxygen, cannot be met by oxygen itself. Of the acids of
sulphur, HjS , in which sulphur has two electro-negative units of valence,
is quite unstable, while (H0)2 S Oj , in which the sulphur has six electro-
positive units of valence, is the most stable. The sulphides (salts of HjS)
of the heavier metals quite generally are insoluble in water, an important
means of separation in analysis. The sulphates (salts of H2SO4) of the
larger number of the metals are soluble in water, the exceptions being
im1)ortant to observe, those of Pb 206.92, Ba 137.40, Sr 87.60, and (with
sparing solubility) Ca 40.1. Of these sulphates, that of barium (least solu-
ble), is the one usually employed in analytical separation.
§15. The Relations of Carbon. — C 12.0. Carbon, in a central position
in respect to polarity, stands alone in its capacity for a multitude of dis-
tinct compounds with hydrogen and oxygen, with and without nitrogen,
these being the so-called organic compounds. This capacity goes with
the power of carbon atoms to unite with each other in the same mole-
cule. It appears in acetylene C2H2(HC=CH), also in oxalic acid,
(HO) OC — CO (OH). The same capacity of union of the atoms of an
element with each other, in the molecules of compounds, is exercised
by other elements in fewer instances, as by nitrogen in hydronitric acid,
by oxygen in ozone, by siil])hur in thiosulphuric acid, and by iodine
in periodides. In carbon, nitrogen, and oxygen we see a decreasing grada-
tion of this capacity, as the atomic weights ascend. Silicon, next to carbon
in Group IV., but in tlie opposite set of alternates, agrees with carbon in
the formation of many corresponding compounds, while it is entirely desti-
tute of the capacity of uniting its atoms to each other in building up
combinations.
§16. The Classification of the Metals as Bases,
The grouping of all the elements, both metals and not metals, according
to their properties as related to their atomic weights, is the object of The
Periodic System, briefly given in the foregoing pages for studies bearing
especially upon the main methods of analysis.
The ordinary grouping of the bases in the work of analysis, outlined in
the next paragraph, is done by the action of a few chemical agents, termed
"group reagents," which have been chosen from a large number of re-
agents, as being more satisfactory than others, for the use of the greater
number of analysts. This ordinary grouping, therefore, is not the only
way in which the metals can be separated, in the practice of anaMical
chemistry, nor is any one scheme of separation adopted throughout by all
authorities. The principal separations of analysis can be well understood
by gaining an acquaintance with the properties of the leading bases and acids.
.^16.
THE CLASSIFICATION OF THE METALS AS BASES.
11
in their action upon each oilier. Without this acquaintance, the analyst is
the servant of routine, and his results liahle to fallacy.
The following named are the bases of more common occurrence.
The Alkali Bases.
The sixth group,*
Potassium (Kalium), K^f
Sodium (Xatrium), Na^
Ammonium, (HH^)^
The Alkaline Earth Bases.
The fifth group.
Magnesium, Mg".
Calcium, Strontium, Barium, Ca".
Sr'^ Ba".
Not precipitated from their salts
hy any of the groxip reagents. Potas-
sium and sodium are found after re-
moving all the following named
groups. Ammonium is found by
tests of the original, this base being
added in the " group reagents."
In coinbination in potassium hy-
droxide, KOH , and in potassium
salts, such as the chloride KCl , and
the nitrate, KHO3 .
In the base, sodium hydroxide and
its salts.
Forms ammonium hydroxide,
HH4OH , representing ammonia,
HH3 , and water, and serving as the
base of ammonium salts, such as
(.N 114)0804 , ammonium sulphate.
(Precipitated by carbonates, which
fact alone ^oes not separate them
from the following named groups.)
Separated hy precipitation as a
phosphate after removing all the fol-
lowing named bases. Forms magne-
sium hydroxide, Mg(OH)o , and mag-
nesium salts, such as MgS04 .
Separated by precipitation with
Ammonium Carbonate, adding
HH4GI to keep magnesium from pre-
cipitation. Calcium carbonate, a
normal salt, CaCO.; .
• The sixth divlBion of the bases, in the order in which they are separated from each other by
precipitation with the erroup reag-ents.
tThe Roman numerals (as i) express units of valence, each equivalent to an atom of
hydrogen, in the formation of salts and other combinations.
13
CLASSJFlCATlOy OF THE METALf! A8 BASES,
§18.
The Zinc and Iron Groups.
The Zinc Oroup.
The fourth group.
Zn": zinc salts.
Mn": manganous salts.
Mn"^ : manganic salts.
Mn^: salts unstable.
Mn^: salts of manganic acid.
Mn^^: salts of permanganic acid.
Ni": nickel salts.
Co": cobaltous salts.
Co"' : cobaltic salts.
The Iron Oroup.
The third group.
Fe" : ferrous salts.
Fe'": ferric salts.
Cr'": chromic salts.
Cr^: chromates.
Al"': aluminum salts.
Metals falling with Copper and Tin.
The second group.
The Copper Group.
Division B, second group.
Mercury (Hydrargyrum).
Hg": mercuric salts.
Hg': mercurous salts.
Silver (Argcntum).
Ag': silver salts.
Lead (Plumbum).
Pb": lead salts.
Bi"': bismuth salts.
Cu": copper or cupric salts.
Cu': cuprous salts.
Cd": cadmium salts.
(Precipitated by sulphides, this
being a separation from the fore-
going, not from the following named
groups of bases.)
Separated by precipitation ivith
Ammonium Sulphide, after removal
of all the following named bases as
directed below. (The precipitates
are all sulphides.)
Separated by precipitation with
Ammonium Hydroxide, in presence
of HH4CI , after the removal of the
groups named following. (The pre-
cipitates are all hydroxides.)
Precipitated by HoS in acidulated
solution. (The precipitates are sul-
phides.)
Separated by the insolubility of
the precipitated sulphides in treat-
ment with Ammonium Sulphide.
ns.
THE OPERATIOXi^ OF AyALYSlS.
13
The Tin Oroup.
Division A, second group,
Sn": stannous salts.
Sn^: stannic salts and stannates.
Sb"^: antimonoiis compounds.
Sb^: antimonic compounds.
A»^": arsenous compounds.
A»^: arsenic compounds and arsen-
ates.
Ketalfl Precipitated as Chlorides.
The Silver Oroup.
The first group.
Separated by dissolving the pre-
cipitated sulphides with Ammonium
Sulphide.
The silver, lead, and univalent
mercury, grouped in the division last
above given. Silver and the mer-
cury of mercurous salts can be re-
moved, as chlorides, by precipitation
with hydrochloric acid. The precip-
itate of lead is not insoluble enough
to remove this metal entirely, in sep-
aration from other groups.
§17. The Acids of Certain Commonly Occurking Salts.
Name of Acid.
Name of Salt.
Formula.
ShowinfiT Hydrozyl.
Anhydrl(
Carbonic
Carbonate
H,CO.
(HO),CivO
CO,
Oxalic
Oxalate
H=C,0,
(HO),C,ivO,
C.O.
Nitric
Nitrate
HNO,
(HO)NV'O,
N,0.
Nitrous
Nitrite
HNO,
(HO)NmO
N,0.
Phosphoric (ortho)
Phosphate
H.PO,
(H0),PV0
P.O.
Metaphosphoric
Metaphosphate
HPO,
(HO)PVO,
P.O.
Pyrophosphoric
Pyrophosphate
H,P,0,
(H0),PV,0,
P.O.
Sulphuric
Sulphate
H=SO,
(HO),Svib,
SO.
Sulphurous
Sulphite
H,SO,
SO,
Hydrosulphuric
Sulphide
H,S
Hydrochloric
Chloride
HCl
Hydrobromic
Bromide
HBr
Hydriodic
Iodide
HI
Chloric
Chlorate
HCIO,
(HO)ClvO,
C1.0.
Iodic
lodate
HIO,
(HO)IVO,
1.0.
The Operations of Analysis.
§18. Chemical analysis is the determination of any or all of the compo-
nents of a given portion of matter, whether this be solid, liquid or gaseous.
A portion of matter is made up of one or more definite and distinct sub-
stances, or chemical individuals, each of which is either a "compound" or
14 THE OPERATIONS OF ANALYSIS. §19.
an " element " and is always and everywhere the same. It is required of
analysis to determine a chemical compound as a body distinct from the
chemical elements that hsve formed it. For example, the analyst may
have in hand a mixture containing sodium sulphate, Na2S04 ; sodium sul-
phite, NajSOg, and sodium thiosulphate, HaaSaOaj.but not containing any
sodium or sulphur or oxygen as these bodies are severally known to the
world and described in chemistry. In this instance the analyst in his
ordinary work does not separate the sulphur or the sodium, as elements
uncombined with oxygen, cither in qualitative or in quantitative oper-
ations. Each one of the compounds of the sulphur with the oxygen is
usually sought for and found and weighed as a chemical individual. Cer-
tain of the chemical elements, however, are frequently separated free from
all combination, as a method of determination of their compounds.
§19. The analysis of gaseous material is termed Gas Analysis; that of
mixtures of the complex compounds of carbon. Organic Analysis. An
examination of organic matter, when limited to a determination of its ulti-
mate chemical elements is styled Ultimate Organic Analysis. When it is
undertaken to determine individual carbon compounds actually existing in
organic matter, it has been spoken of as Proximate Organic Analysis. If
the same distinction were to be applied to inorganic analysis, we should
have to say that it is mostly " proximate ^' but is sometimes " ultimate ''
in its methods of operation.
§20. The term Qualitative Chemical Analysis as commonly used is con-
fined to a chemical examination of material, chiefly inorganic, in the solid
or liquid state, the inquiry being limited for the most part to well known
substances.
§21. In the methods of analysis of a mixture, it is often required to
separate individual substances from each other, but sometimes a distinct
compound can be identified and sometimes its quantity can be estimated
while it is in the presence of other bodies. Both the identification and
separation are accomplished, nearly always, by effecting changes, physical
and chemical.
Methods of analysis are as numerous as are the ways of bringing into
action the physical and chemical forces by which chemical changes are
wrought. The characteristics of any chemical individual, by which it is
distinguished and removed from others, lie in its responses to the physical
and chemical forces, including especially the chemical action of certain
well known compounds called reagents.
§22. The response toward heat and pressure fixes the melting and boiling
points, its ordinary solid or liquid or gaseous state. The operations "in
the dry way " are done over a flame or in a furnace, with or without solid
^'reagents" and with regard to oxidation. They represent some of the
§27. THE OPERATIONS OF ANALYSIS, 15
methods of metallurgical manufacture. The liquid state, whether by
fusing or by solution, is the state commonly necessary or favorable to chem-
ical change and its control.
§23. The deportment of a solid substance toward light comprises its
color and that of its solutions, as well as that of its vapor, in ordinary light,
and the bands and primary colors it exhibits in the uses of the spectroscope
(Crookes, J. C, 1889, 66, 255; Welsbach, i¥., 1885, 6, 47).
§24. The conduct of a chemical compound in electrolysis is, in various
cases, a means both of identification and of separation. Electric conduc-
tivity methods are used for establishing the presence or absence of minute
traces of substances (Kohlrausch Whitney, Z, phys. Ch., 1896, 20, 44).
Again, traces of dissolved matters too minute for other means of detection
can be revealed by the difference of eUctric potential between electrode and
solution (Ostwald, Lehrh, 2 Aufl., II, 1, 881; Behrend, Z. phys. Ch., 1893,
11, 466; Hulett, Z. phys, Ch,, 1900, 33, 611).
§26. By far the most extensive of the resources of analysis lie in the
chemical reaction of one definite and distinct substance with another, ac-
cording to the character of each, giving rise to a chemical product having
peculiarities of its own in evidence of its origin. In this way the com-
pounds are bound in regular relations to each other. Therefore it belongs
to the analyst to gain personal acquaintance with the behavior of the repre-
sentative constituent bases and acids toward each other.
§26. Operations for chemical change are commonly conducted in solu-
tion. The material for analysis is dissolved, and is treated with reagents
that are in solution. A solid or a gas is dissolved in a liquid in making a
solution. Wlien the dissolved substance is converted into one that will
not dissolve a precipitate is formed. It is necessary therefore to under-
stand the nature of solution and to give heed to its obvious limitations.
Certain facts and conclusions as to the chemical state of dissolved com-
pounds are presented under the head next following, '' Solution and Ioniza-
tion." But it must first be observed that the universal solvent, water, is
always understood to be present in somewhat indefinite proportion in opera-
tions " in the wet way." It serves as a vehicle, as such not being included
in any statement of the substances operated upon, nor formulated in equa-
tions, any more than is the material of the test tube, but often some portion
of it enters into combination or suffers decomposition, and then it must 1;(?
placed among the substances engaged in chemical change.
§27. Xo other pro])crty of substances has so great importance in analysis
and in all chemical operations, as their sohthiHty in irater. It must never
^ be forgotten that there are degrees of solubility, but there is hardly such a
Uu't as absolute solubility, or insolubility, regardless of the proportion
of the solvent. There are liquids which are miscible with each other
16 THE OPERATIOXS OF AXALYi^IS. §28.
in all proportions, but solids seldom dissolve in all proportions of the sol-
vent, neither do gases. For every solid or gas, there is a least quantity of
solvent which can dissolve it. One part of potassium hydroxide is soluble
in one-half part of water (or in any greater quantity), but not in a less
quantity of the solvent. One part of sodium chloride requires at least two
and a half parts of water to dissolve it. One part of mercuric chloride will
dissolve in two parts of water at 100 degrees, but when cooled to 15 degrees
so much of the salt recrystallizes from the solution, that it needs twelve
parts more of water at the latter temperature to keep a perfect solution.
Lead chloride dissolves in about twenty parts of hot water, about half of
the salt separating from the solution when cold. Calcium sulphate dis-
solves in about 500 times its weight of water — this dilute solution forming
one of the ordinary reagents. Barium sulphate is one of the least soluble
precij)itates obtained, requiring about 430,000 parts of water for its solution
at ordinary temperature (Hollemann, Z. pliys, Ch., 1893, 12, 131). In ordi-
nary reactions it is not appreciably soluble in water. Lead sulphate dis-
solves in about 21,000 parts of water: in many operations this solubility
may be disregarded, but in quantitative analysis the precipitate is washed
with alcohol instead of water, losing less weight with the former solvent.
These examples indicate the necessity of discriminating between degrees of
solubility. Also the solubility of a particular compound is dependent upon
the physical form of that compound (§69, 5h); e, g., amorphous magnesium
ammonium phosphate is quite soluble in water, the crystalline salt being
almost insoluble. "When a solvent has dissolved all of a substance that it
can at a particular temperature, in contact with the solid, the solution is
said to be saturated at that temperature. It frequently happens that a
saturated solution of a substance at a higher temperature may be cooled
without separation of the solid. Such a solution (at the lower temperature)
is said to be supersaturated and precipitation frequently is induced by
jarring the solution, more surely by adding a crystal of the dissolved sub-
stance.
§28. The ordinary liquid reagents are solutions in water — sulphuric acid
and carbon disulphide being exceptions. Hydrochloric acid, liquid hydro-
sulphuric acid, and ammonium hydroxide (reagents) are solutions of gases
in water; on exposure to the air these gases gradually separate from their
solutions. All these gases escape much more rapidly when their solutions
are warmed. The majority of liquid reagents are solids in aqueous solu-
tion. (See the list of Reagents.)
§29. Substances are said to dissolve in acids, or in alkalis, and this is
termed chemical solution; more definitively it is chemical action and solu-
tion, the solution being counted as a physical change. We say that cal-
cium oxide dissolves (chemically) in hydrochloric acid; that is, in the
§33. THE OPERATIOyS OF ANALT8I8, 17
reagent named hydrochloric acid, a mixture of that acid and water. The
acid unites with the calcium oxide, forming a soluble solid, which the water
<lissolves. Absolute hydrochloric acid cannot dissolve calcium oxide.
§30. Solids can be obtained, without chemical change, from their aqueous
solutions: Firstly, by evaporation of the water. This is done by a careful
application of heat. Secondly, solids can be removed from solution, with-
out chemical change, by (physical) precipitation — accomplished by modify-
ing the solvent. If a solution of potassium carbonate, or of ferrous sul-.
phate, be dropped into alcohol, a precipitate is obtained, because the salts
will not dissolve, or remain dissolved, in the mixture of alcohol and water.
But, in analysis, precipitation is more often effected by changing the dis-
solved substance instead of the solvent.
§31. Solids can be separated from their solution by precipitation due to
chemical change, to the extent that the product is insoluble in the quantity
of the solvent present. Calcium can be in part precipitated from not too
dilute solutions of its salts, by addition of sulphuric acid; but there still
remains not precipitated the amount of calcium sulphate soluble in the
water and acid present, which is enough to give an abundant precipitate
with ammonium oxalate, the precipitated sulphate being previously re-
moved by filtration.
Time and heat are required for the completion of most precipita-f
tions. If it is necessary to remove a substance, by precipitation, before
testing for another substance, the mixture should be warmed and allowed
to stand for some time, before filtration. Xcglect of these precautions often
occasions a double failure; the true indication is lost, and a false indication
is obtained.
§32. Reagents should be added in very small portions, generally drop by
drop. Often the first drop is enough. Sometimes the precipitate redis-
solves in the reagent that produced it, and this is ascertained if the reagent
be added in small portions, with observation of the result of each addition.
If it is a final test, a quantity of precipitate which is clearly visi])le is suffi-
cient, but if the precipitate is to be filtered out and dissolved, a considerable
quantity should be formed. If the precipitate is to be removed and the
filtrate tested further, the precipitation must be completed — by adding the
reagent as long as the precipitate increases, with the warmth and time
requisite in the operation; and a drop of the same reagent should be added
to the filtrate to obtain assurance that the precipitation has been completed.
It will be found, with a little ex])erience, tliat some reagents must be used
in relatively large quantities. On the contrary, the acids, sulphuric, hydro-
chloric and nitric, are required in a volume relatively very small.
§33. Certain very exact methods of identification can be conducted by
drop tests upon a black or white ground, or upon a glass slide and especially
18 THE OFERATIOyS OF ANALYSIS, §34.
with help of a microscope and with studies of crystalline fonn. Further
Kee Behrens, Z. 1891, 30, 125; and Herrnsehmidt and Capelle, Z, 1893, 32,
608.
§34. Precipitates are removed — usually by filtration, sometimes by decan-
tation. If they are to be dissolved, they must be first washed till free from
all the substances in solution. For complete precipitation some excess of
the reagent must have been used. Beside the reagent there are other dis-
solved matters, after precij)itations, some of which are indicated by the
equation written for the change. All these dissolved substances permeate
and adhere to the porous precipitate with greater or less tenacity. If they
are not wholly washed away, some portion of them will be mixed with the
dissolved precipitate. Then, the separation of substances, the only object
of the precipitation is not acc()m])lislied, while the operator, proceeding
just as though it was accom])lished, imdertakes to identify the members of
a group by reactions on a mixture of groups. The washing, on the filter,
is best completed l)y repeated additions of small portions of water — ^around
th"e filter border, from the wash bottle — allowing each portion to pass
through before another is added. The washings should be tested, from
time to time, until they are free from dissolved substances.
§36. In dissolving preci])itates — by aid of acids or other agents — use
the least possible excess of the solvent. Endeavor to obtain a solution
nearly or quite saturated, chemically. If a large excess of acid is carried
into the solution to lie o])erated upon, it usually has to be neutralized, and
the solution then becomes so greatly encumbered and diluted that reactions
become faint or ina])precial)le. Preei])itates may be dissolved on the filter,
without excess of solvent, by ])assing the same portion of the (diluted)
solvent repeatedly through the filter, following it once or twice with a few
drops of water. The mineral acids should be diluted to the extent required
in each case. For solution of small quantities of carbonates and some
other easily soluble precipitates the acids may be diluted with fifty times
their weight of water. "Washed precipitates may also be dissolved in the
test-tube, by rinsing them from the filter, through a jnincture made in its
point, with a very little water. If the filter be wetted before filtration, the
precipitate will not adhere to it so closely.
§36. When the addition of a reagent is to cause a change in the acid,
alkaline or neutral condition of the solution, the addition of sutlicient
reagent to cause the desired chan<:e should always bo governed by testing
a drop of the solution, on a glass rod, with a piece of litmus paper.
§37. AVhen substances in separate solution are brought together, an
evidence of the formation of a new substance is the appearance of a solid
in the mixture, a precipitate. A chemical change between dissolved sub-
stances — salts, acids, and bases — will be ])ractically complete when one or
§40. THE 0PER4TI0XS OF ANALYSIS. 19
more of the products of such change is a solid or a gas, not soluble in the
mixture. As an example, Calcium carbonate + Hydrochloric acid = Cal-
cium chloride + Water + Carbon dioxide (gas).
§38. In the practice of qualitative analysis, the student necessarily refers
to authority for the composition of precipitates and other products. For
example, when the solution of a carbonate is added to the solution of a
calcium salt, a precipitate is obtained; and it has been ascertained by quanti-
tative analysis that this precipitate is normal calcium carbonate, CaCO^ ,
invariably. Were there no authorized statement of the composition of this
precipitate, the student would he unable, without making a quantitative
analysis, to declare its formula or to write the equation for its production.
When the results of analytical operations are substances of unknown, uncer-
tain, or variable composition, equations cannot be given for them.
§39. The written equation represents only the substances, and the quan-
tity of each, which actually undergo the chemical change that is to be
expressed. Thus, if a reagent is used to effect complete precipitation, an
excess of it must be employed, beyond the ratio of its combining weight in
the equation. That is, if magnesium sulphate be employed to precipitate
barium chloride, the exact relative amount of magnesium sulphate indicated
by the equation: BaClg + MgSO^ = BaSO^ + HLgCh , fails to precipitate all
of the barium. The soluble sulphate must be in a slight excess. On the
other hand, to effect complete precipitation of the sulphate the barium
must be in a slight excess.
§40. By translating chemical equations into statements of proportional
parts by weight, they are prepared to serve as standard data of absolutely
pure materials, and applicable in operations of manufacture, with large or
small quantities, after making due allowance for moisture and other im-
purities, necessary excess, etc. In quantitative analysis the equation is the
constant reliance. For example, in dissolving iron by the aid of hydro-
chloric acid, we have the equation:
Fe + 2HC1 = FeCl^ + H, .
56 + 72.9 = 126.9 + 2 .
Also in precipitating ferrous chloride by sodium phosphate, we have the
equation:
Feci, + Na2HP0..12H=0 = FeHPO, + 2NaCl + 12H,0 .
126.9 + (142.1 -f 216) = 152 -f 117 .
Suppose it is desired to determine from the above:
(1) How much hydrochloric acid, strength 32 per cent, is required to
dissolve 100 parts of iron wire.
(2) What quantities of 32 per cent hydrochloric acid and iron wire are
necessary to use in preparing 100 parts of absolute ferrous chloride.
20 SOLUTION AND IONIZATION. §41.
(3) What materials and what quantities of them, may be used in prepar-
ing 100 parts of ferrous phosphate.
In practice allowance must be made for the facts that the iron wire will
not be quite pure, and that a considerable excess of the hydrochloric acid
would be necessary to the complete solution of the iron. Also that some
excess of the phosphate would be necessary to the full precipitation of the
iron. Irrespective of impurities, oxidation product and excess, the re-
quired quantities are found by the combining weights as follows:
^ f 56/72.9 = lOO/x = parts of absolute HCl for 100 parts of iron wire.
' \ 32/100 = x/y = parts of 32 per cent HCl for 100 parts of iron wire.
i 126.9/72.9 = lOO/x
32/100 = x/y = parts of 32 per cent HCl for 100 parts of PeCl, , absolute.
126.9/56 = lOO/z = parts of iron wire for 100 parts of PeCl,.
152/72,9 = lOO/x
32/100 = x/y = parts of 32 per cent HCl for 100 parts of reHP04 .
152/56 = lOO/z = parts of metallic iron for 100 pai-ts of 7eHP0«.
152/358.1 = lOO/u = parts of Na,HP04,12H,0 for 100 parts of TeHPO*.
Practice in reducing the combining numbers of the terms in an equation
to simple parts by weight, is a very instructive exercise, even in the early
part of qualitative chemistry. It enforces correct and clear ideas of the
significance of formula* and equations, and refers all chemical expressions
to the facts of quantitative work.
§41. The chief requirement in qualitative practice is an experimental
acquaintance with the chemical relations of substances, rather than the
identification of one after the other by routine methods. The acids and
bases, the oxidizing and reducing agents, are all linked together in a net-
work of relations, and the ability to identify one, as it may be presented in
any combination or mixture, depends upon acquaintance with the entire
fraternity.
§42. The full text of the book, rather than the analytical tables, should
be taken as the guide in qualitative operations, especially in those upon
known material. The tabular comparisons are commended to attention,
especially for review. In actual analysis, the tables serve mainly as an
index to the body of the work.
Solution and Ionization.
§43. The Theory of Electrolytic Dissociation, proposed by Arrhenius in
1887 (Z. phys. Ch., 1887, 1, 631), assumes that salts, acids, and bases in
water solution are present not as the intact molecule but split up into
certain components, and that the characteristics of the dissolved substance
result very largely from the extent to which this breaking down of the
§43. SOLUTION AND IONIZATION. 21
molecule has taken place. The facts upon which the theory is based are
in a word the parallelism between osmotic pressure,* electric conductivity,
and chemical activity of substances in solution.
The gas-laws (Boyle's, Gay-Lussac's, Henry's, and Dalton's) are found
to hold for dissolved substances, osmotic pressure being substituted for
gas-pressure (van 't Hoff, Z, phys. Ch., 1887, 1, 481). Avogadro's Hypoth-
esis is therefore applicable to solutions as well as to gases, and as abnormal
gas-pressure points to dissociation in the gas (NH^Cl, PCI5) so excessive
osmotic pressure is taken as indicating dissociation of the dissolved sub-
stance. The osmotic pressure is a measure of this dissociation.
Faraday gave the name ions to the components of a substance conducting
the electric current in solution. It is an observed fact that transmission
of the current by a solution is always accompanied by movement of the
ions in opposite directions (Hittorf, Pogg. 1853, 89, 177). This is quite
independent of any separations taking place at the electrodes. From this
it is concluded that the ions carry the electricity from one pole to the
other through the solution. If the ions are the carriers of electricity then
the power of a solution to conduct the current will be in proportion to their
number, that is, to the extent of dissociation of the dissolved substance.
And experiment shows that the dissociation calculated from the osmotic
pressure is identical with the dissociation calculated from the electric
conductivity.
Further, if in analysis of a substance in solution we are dealing not with
the substance in its integrity but with certain ions, then our ordinary
analytical reactions are reactions of the ions, and we may expect that where
the substance for some reason is transformed from the ionized condition
to the undivided molecule these reactions will fail. Here again the chemi-
cal activity will be proportional to the number of ions; and experiment
shows that unquestioned quantitative parallelism exists, to take the case
of acids, between (1) the characteristic acid activity — the dissolving of
metals, the influence as catalyzer on such changes as the inversion of cane-
sugar and the saponification of esters; (2) the extent of dissociation as
indicated by osmotic pressure, and (3) the extent of dissociation as indicated
by electric conductivity. The same parallelism holds for other bodies in
solution. The very active acids and bases and the neutral salts underi^o
wide dissociation in water solution, while weak acids and bases retain
almost entirely the non-dissociated condition.
The Electrolytic Dissociation Theory in its assumption of a separation
•The pressure by virtue of which a soluble substance In contact with the solvent, as common
salt In water, is enabled to rise against the force of gravity and distribute Itself uniformly
throu^iout the solvent, just as a gas by virtue of the gas-pressure occupies the entire space at
its diiposaL
22 SOLUTION AND IONIZATION. §44.
into ions groups together and gives system and meaning to these three
classes of facts, experimentally absolutely independent and up to Arrh'enius'
time without any suspected relationship. In each case the results calculated
on the assumption of such a dissociation are in quantitative agreement with
those obtained by measurement.
Corresponding in actual experience to the view that the common analyti-
cal reactions arc due to the ions rather than to the molecule as a whole, is
the analyst's practice of testing for acid radicle or basic radicle without
regard to the other component; and on the other hand, to take a specific
case, the fact that the sulphur in H^iS does not give the same precipitation
reactions as that in K,S or HgSO^ or H2SO3 or H2S2O3 . Further, HgClj in
its chemical behavior is unlike other mercuric salts and unlike other
chlorides. The mercury is not readily precipitated by alkali hydroxides
nor is the chloride readily precipitated by silver salts. In agreement with
this, its conductivity and osmotic pressure are also unlike those of the great
majority of neutral salts, both pointing to very slight dissociation into the
ions. CdClo is another neutral salt anomalous in that its conductivity and
osmotic pressure are both low. And here also for precipitation of the
chloride a considerable concentration of the reagent is necessary. Similar
instances of the parallelism referred to are numberless.
§44. The Law of Mass-Action embodies the familiar principle that the
chemical activity of a substance is proportional to its concentration. It
was first recognized, although imperfectly, by BerthoUet and was given
mathematical expression by Guldberg and Waage in 1867. The latter
investigators found it to accord well with the observed facts in some cases;
in others there were wide discrepancies which were later shown by Ar-
rhenius to disappear when the concentration, not of the reacting body as a
whole but only of that part present in the ionized condition, was taken
into consideration. "We must assume that every chemical reaction is rever-
sible, that is, that none of them proceed until the reacting substances are
completely transformed. Then by a simple process of reasoning it is found
that when equilibrium sets in the product obtained by multiplying together
the concentrations of the reacting substances will be in a certain definite
ratio to the product of the concentrations of the substances formed, con-
centration being defined as the quantity in unit volume.* For example,
in the reaction indicated by the equation CH3CO0H -\- CsH^OH =
CH3C02C2H^ + H2O , when equilibrium sets in ab = ked , in which a and b
are the concentrations of acid and alcohol respectively, c and d those of
ester and water, while k is a constant peculiar to the reaction. AVhere the
♦ The unit of quantity is the molecular weight taken in firrams (the ** mol "). Where there are
18.23 grams HCl in a liter either in solution or as gas the concentration Is H* wherid there are
72.02 grams in the same volume the concentration is 2, and so on.
§46. SOLlTIOy AA'D lONIZATIOy. 23
reaction is a dissociation, as with gaseous NH4CI , we have ab = k'c , a and b
representing the concentrations of NH3 and HCl respectively, c that of the
undecomposed NH^Cl, and k' the constant characteristic of this change.
Dissociation into ions must follow the same laws, and for the electrolytic
dissociation of acetic acid a similar equation holds, a and b in this case
standing for concentration of H and acetic ions, c for concentration of non-
dissociated acetic acid, while the constant is one governing only this par-
ticular dissociation. It is apparent from each of these equations that, if .
we add one of the products of the reaction and thus increase its concentra-
tion, the concentration of the other product must decrease in the same
proportion — the extent of the reaction will be decreased; while, on the
other hand, removing either or both of the products will tend to make the
transformation complete. This deduction is of great significance. In
making ethyl acetate from the acid and alcohol, in order to use the materials
as completely as possible, the ester is distilled off as rapidly as produced
while the water is taken up by some absorbent. Introducing gaseous NH3
or HCl diminishes the dissociation of NH^Cl by heat, and similarly adding
either H ions or acetic ions will diminish the dissociation of acetic acid.
Acetic acid is much weakened by the presence of a neutral acetate. A
ferrous solution moderately acidified with acetic acid gives no precipitate
on saturation with HjS , but on addition of sodium acetate the black FcS
is brought down. Similarly a weak base, as NH^OH , is made still less
effective by the presence of its strongly-dissociated neutral salt, as NH^Cl .
Quantitative agreement is obtained between observed effect of NH4CI on
ITH^OH as saponifying agent and that calculated from the equation:
^NH^ • ^OH' ~ '^^NH^OH (Arrhenius, Z, phys, Ch., 1887, 1, 110).
§46. The Solubility-Product. — In the saturated solution which always
remains after precipitation we have the usual dissociation equilibrium, as:
^£g • ^QY ^AeCl • "^^^ ^^^ quantity of non-dissociated substance in
a saturated solution is invariable and the right side of this equation is
therefore constant. That is, in saturated solution the product of the con-
centrations of the ions is always the same for a given substance (Xernst).
This Ostwald has called the Solubility-Product. Where the saturated solu-
tion is made by bringing the salt into contact with the solvent ^j^^ • ^ ^ny •
From such a solution precipitation will take place on addition of either a
silver salt or a chloride, for such addition largely increases the concentration
of one ion and, to restore equilibrium, the concentration of the other ion
must decrease in the same proportion, which is possible only by precipita-
tion. From this follows the old empirical rule to add an excess of the
reagent in making a precipitation. Experiments on this point give quanti-
24 ORDER OF LABORATORY STUDY, §46.
tative agreement with the theory (Xernst, Z, pJiys. Ck,, 1889, 4, 372;
Noyes, Z. phys, CA., 1890, 6, 241; 1892, 9, 603).
The Solubility-Product of the alkaline-earth carbonates is
^M " ^CO " ^ • I^ *h® solution of a neutral salt, as CaClj , Ca ions are
present in large concentration. When a substance containing CO, ions in
large concentration is added, as NasCO, , the solubility-product is exceeded
and precipitation takes place. Carbonic acid, however, is shown by con-
ductivity and osmotic pressure measurements to be but slightly disso-
ciated, that is, it contains few COs ions, and in accord with this is the
familiar fact that the alkaline earths are not precipitated by carbonic acid.
Similarly the fixed alkali hydroxides, strongly dissociated, will precipitate
alkaline-earth hydroxides, while ammonium hydroxide, shown by other
measurements to contain but few hydroxyl ions, will not.
For the metallic sulphides the solubility-product is ^jf^S"^^
The alkali sulphides as normal salts contain the S ion in large concentra-
tion and so produce precipitation even of the more soluble sulphides of
the Iron and Zinc Groups. The slightly dissociated HgS contains sufficient
S ions to reach the solubility-product of the sulphides of the Silver, Tin.
and Copper Groups, but not enough to attain to the larger solubility-
product of the Iron and Zinc Group sulphides. A strong acid, as HCl .
containing as it does H ions, one of the dissociation products of HgS , drives
back the dissociation of the H^S, so decreasing the concentration of the
S ions and making precipitation of the sulphide more difficult.
For the application of the dissociation theory to the details of analytical
work we arc indebted chiefly to Ostwald. See his " Scientific Foundations
of Analytical Chemistry " and " Outlines of General Chemistry.'^
Order of Laboratory Study.
§46. The following is a suggestive outline to be modified by the teacher
to suit the ability of the students, and the amount of time to be given to
the study :
a, A review of chemical notation and the writing of salts.
h, A study of the action of the Fixed Alkalis upon solutions of the salts
of the metals in the order of their groupings; including the action of an
excess of the reagent. The fact of the reaction should be stated; e. {j..
lead acetate + potassium hydroxide = a white precipitate readily soluble in
excess of the reagent. The text should then be consulted for the products
of the reaction (6a), and the reactions expressed in the form of equations:
2Pb(C,H,0,), -h 4K0H=Pb, 0(0H),» (white) -h 4KC,H,0, + H,0
Pb,0(OH), -h 4K0H (excess)"^ 2K,PbO, -f- SH^O
or PbCCaH.O,), -f- 4K0H (excess) = K,PbO, + 2KCaHaOa + 2HjO .
* It has been found helpful to require students to underscore all precipitates.
j46. ORDER OF LABORATORY STUDY, 25
The results should all be tabulated and then summarized in form of a
carefully worded generalization (§205, Ga).
c. Action of Ammoniiim Hydroxide (volatile alkali) upon solutions of
the salts of the metals, etc., as in {b) above; e. g., lead nitrate + ammonium
hydroxide = a white precipitate not dissolving in excess. Consult text
(§67, 6a) and write the equation :
3Pb(H0,), + 4NH«0H = 2PbO.Pb(yO,)« + 4NH«N0, + 2H,0 .
After the work has been completed in the laboratory and the results
discussed in the class room, summarize in the form of a generalized state-
ment (§207, 6a).
d. A study of the action of the Fixed Alkali Carbonates, and generaliza-
tion of the results (§206, 6a).
e. A study of the action of Ammoninm Carbonate. Summarize the re-
sults (§207, 6a).
f. A study of the solvent action of acids, HCl , HNO, , and E2SO4 , upon
the Hydroxides and Carbonates obtained by precipitation.
g. Action of Hydrosnlphuric Acid as a precipitating agent upon salts of
the metals in neutral and acid solutions.
h. The use of Ammoninm Sulphide as a reagent.
t. The solvent action of acids, HCl , HNOj, and HC2H3O2 , upon the
sulphides obtained by precipitation.
;. Action of Hydrochloric Acid and Soluble Chlorides.
Action of Hydrobromic Acid and Soluble Bromides.
Action of Hydriodic Acid and Soluble Iodides.
k. Precipitation by Soluble Sulphates, Phosphates, and Oxalates.
Z. The solvent action of Hydrochloric and Acetic Acids upon the Phos-
phates obtained by precipitation.
m. The reverse of certain of the above reactions as illustrating the
precipitation of Acids; e. g., Ammonium oxalate + calcium chloride = a
white precipitate. Consult the text (§227, 8), and write the equation:
(WHJjCA + CaClj = CaCjO, + 2NH4CI .
n. Application of the above reactions to the Grouping of the Hetals
for Analysis.
0. A study of the limit of visible precipitation with several reagents
upon a particular metal, or upon a number of metals.
p. A study of the analysis of the individual metals and acids; combining
them, and eifecting their separation and detection. The new work of
each day to be followed by the anah'sis of " unknown *' mixtures prepared
by the teacher to illustrate tlie new work and to give an instructive review
of the preceding work. The order of the study of the metals and acids
may be varied greatly. In no case should the metals of a whole group be
studied without considering the relations to the other groups.
2G ORDER OF LABORATORY STUDY, §46.
q. The study in the elass room of Oxidation and Bednotion, with work
in the laboratory to illustrate.
r. The study of problems in Synthesis involving analytical separations,
accompanied by laboratory experiments.
s. The analysis of a series of Dry " TTnknown ** Hiztnres.
f. A special study of the analysis of Phospliates, Oxalates, Borates,
Silicates, etc., and certain of the Barer Ketals.
u, Tlie analysis of mixtures in solution, illustrating Oxidation and
Beduction.
r. A study of Electrolysis as a means of detection in qualitative analysis.
PART II.-THE METALS.
THE SILVEH AND TIN AND COPPER GROXJPS.
(First and Second Groups.)
§47, The Silver group (first group) includes the metals whose chlorides
are insoluble in water and which are precipitated from solutions ^pon the
addition of hydrochloric acid or soluble chlorides : Pb, Hg', Ag .
The Tin and Copper group (second group) includes those metals whose
sulphides are precipitated by hydrosulphuric acid from solutions acid with
dilute hydrochloric acid, and whose chlorides (soluble in water for the
most part) are not precipitated by hydrochloric acid or soluble chlorides.
Lead*
?b
206.92
Germanium
Ge
72.5
Mercury
Hg
200.0
Iridium
Ir
193.1
Silver
Ag
107.9*i
Osmium
Os
191.0
Arsenic
Ajb
75.0
Palladium
Pd
107.0
Antimony
Sb
120.4
Rhodium
Bh
103.0
Tin
Sn
119.0
Kuthenium
Bu
101.7
Gold
Au
197.2
Selenium
6e
79.2
Platinum
Pt
194.9
Tellurium
Te
127.5?
Molybdenum
Mo
96.0
Tung^sten
W
184.
Bismuth
Bi
208.1
Vanadium
V
51.4
Copper
Cadmium
Cu
Cd
63.6
112.4
§48. Owing to the partial solubility of lead chloride in water, it is never
completely precipitated in the first group; hence it must also be tested
for in the second group. Monovalent mercury belongs to the first group
and divalent mercury to the second. Silver, then, is the only exclusively
first-group metal.
§49. The metals included in these groups are less strongly electro-
positive than those of the other groups. Only bismuth, antimony, tin,
and molybdenum decompose water, and these only slowly and at high
temperatures. The oxides of silver, mercury, gold, platinum, and palla-
dium are decomposed below a red heat. Copper, lead, and tin tarnish by
*In this list of the metals of the SUvcr, Tin and Ck)pper Groups the more common, those In
the first column, are arranged in the order of their discussion and separation in analysis. The
rare metals are arranged in alphabetic order, but are discussed in order of their relations to
each other, beginning at i 104.
28 OEXERAL DISCUSSION. §50.
oxidation in the air. In general, these metals do not dissolve in acids
with evolution of hydrogen, or do so with difficulty. Nitric acid is the
best solvent for all, except antimony and tin, which are rapidly oxidized
by it. Concerning the separation and detection of the metals of these
groups by electrolysis, see Schmucker, Z. anorg., 1894, 6, 199, and Cohen,
J., Soc. Ind., 1891, 10, 327.
§60. Mercury, arsenic, antimony, and tin form, each two stable classes
of salts. Therefore, the lower oxides, chlorides, etc., of these metals act
as reducing agents; and their higher oxides, chlorides, etc., as oxidizing
agents, each to the extent of its chemical force. Arsenic, antimony, tin,
molybdenum, and several of the rare metals of these groups enter into
acidulous radicles, which form stable salts. Arsenic, selenium and tellu-
rium are metalloids rather than metals. Arsenic, antimony, and bismutli
belong to the Nitrogen Series of Elements.
§61. A large proportion of the compounds of these metals are inBoluble
in water. Of the oxides or hydroxides, only the acids of arsenic are
soluble in water. The only insoluble chlorides, bromides, and iodides are
in these groups. The sulphides, carbonates, oxalates, phosphates, borates,
and cyanogen compounds are insoluble. Most of the so-called soluble
compounds of bismuth, antimony, and tin, and some of those of mercury,
dissolve only in acidulated water, being decomposed by pure water, with
formation of insoluble basic salts.
§62. Among the many soluble double salts of the metals of these groups
are especially to be mentioned the double iodides with KI and the iodides
of Pb , Hg , Ag , Bi and Cd . Platinum forms a large number of stable
double chlorides, soluble and insoluble; and gold forms double chlorides,
cyanides, etc.
§53. The oxides of arsenic act as acid anhydrides and form soluble salts
with the alkalis; oxides of antimony, tin, and lead, are soluble in the fixed
alkalis; oxides of silver, copper, and cadmium, in ammonium hydroxide.
Metallic lead, like zinc, dissolves in the fixed alkalis with evolution of
hydrogen.
§64. The solubility of certain sulphides in the alkali sulphides forming
sulpho salts or double sulphides, separates the metals of the second group
into two divisions. .4 (tin group) — As , Sb , Sn , Oe , Au , Ir , Ho , Ft , Se ,
Te , W . and V ; sulphides soluble in yellow ammonium sulphide; and B
(copper group)— Hg , Pb , Bi , Cu , Cd , Os , Pd , Rh , and Ru ; sulphides
not soluble in yellow ammonium sulphide.
§55. Mercury, antimony, silver, and gold do not form hydroxides. The
oxides of gold are very unstable.
§56. The metals of thei=e groups are all easily reduced to the metallic
state by ignition on charcoal. Except mercury and arsenic, which vaporize
§67, 4. LEAD, 29
readily, and certain rarer metals diflScultly fusible, the reduced metals melt
to metallic grains on the charcoal.
The Silveb Group (First Group).
Lead, Heronry (Mercurosum), Silver.
§67. lead (Plumbum) Pb = 206.92 . Valence two and four.
1. Tropwi^eB.— Specific gravity, 11.37 (Reich, J. pr.^ 1859, 78, 328). Melting point,
327.69® (Callendar and Griffiths, C. N., 1891, 63, 2). It begins to vaporize at a
red heat and boils at a white heat. Vaporization is said to take place at 360^^
(Demarcay, C. r., 1882, 95, 183). It can be distilled in vacuo (SchuUer, B., 1883,
16. 1312).
Pure lead is almost white, soft, malleable, very slightly ductile, tarnishes in
the air from formation of a film of oxide. The presence of traces of most of
the other metals makes the lead sensibly harder. It is a poor conductor of heat
and electricity. It forms alloys with most metals; lead and tin in various pro-
portions form solder and i>ewter; lead and arsenic form shot metal; lead and
antimony, type metal; lead, bismuth, tin and silver form a fusible alloy melting
as low as 45® ; bell metal consists of tin, copper, lead and zinc.
2. Occurrence. — ^It is rarely found native (Chapman, PhiL Mag,y 1866, (4), 31,
176) ; its most abundant ore is galena, PbS ; it also occurs as cerussite, PbCO, ;
anglesite, PbS04; pyromorphite, 3Pb,P,08 + PbClj; krokoite, PbCrO^: and
also in many minerals in combination with arsenic, antimony, etc. The
United States produces more lead than any other country. Spain produces
about one-fourth the world's supply.
3. Preparation. — From galena (a) It is roasted in the air, forming variable
quantities of PbS04 , PbO , and PbS ; then the air is excluded and the tempera-
ture raised, and the sulphur of the sulphide reduces both the PbO and the
PbSO^ , SO, being formed: PbSO^ + PbS = 2Pb + 2S0, . 2PbO -j- PbS = 3Pb +
SOj . (6) Similar to the first except that some form of carbon is used to aid
in the reduction, (c) It is reduced by fusing with metallic iron: PbS -|- Fe =
Pb 4- PeS . Frequently these methods are combined or varied according to
the other ingredients of the ore.
4. Oxides.— Lead forms four oxides, PbjO , PbO , PbO, , and Pb.O* . Lead
suboxide (Pb,0) is little known: it is the black powder formed when PbCjO* is
heated to 300**, air being excluded. Lead oxide (litharge, or massicot) is formed
by intensely igniting in the air Pb , Pb,0 , PbO, , Pb.O* , Pb(OH), , PbCO, ,
PbC.Of , or Pb(K0,)2 . It has a yellowish-white color, melts at a red heat, and
is volatile at a white heat.
Trilead tetroxide (red lead or minium), Pb^O^ , is formed by heating PbO
to a dull-red heat with full access of air for several hours. Strong, non-reduc-
ing acids, such as HNO, , H3SO4 , HCIO, , etc., convert it into a lead salt and
PbO, (a). But concentrated hot H^SO^ converts the whole into PbSOi , oxygen
being evolved (6). But with the dilute acid and reducing agents, such as
C,Hb(OH), , C«H„06,HX,0, , H,C«H«0<,, Zn , Al , Cd . Mg , As, Pb . etc.,
it is all reduced to the dyad lead without evolution of oxygen (f)» (<^)» and (e).
Hydra cids iisually reduce the lead and are themselves oxidized (f).
(a) Pb,04 + 2H,S04 (dilute) = PbO, + 2PbS04 + 2TL,0
(6) 2Pb,04 + 6H2SO« (concentrated and hot) = 6PbS0« -f- 6H,0 + 0,
(c) Pb,0, + H,C,04 4- 6HN0, = 3Pb(N0,), + 4H,0 + 200,
(d) 10Pb,O* 4- As, + 30H,SO, = SOPbSO, 4- 4H,AsO* -f 24H,0
(e) Pb.O, + Zn + 4H2SO, = 3PbS0, + ZnSO* + 4H,0
(/) Pb.O, 4- 8HC1 = 3PbCl, + 01, 4- 4H,0
The valence of Pb.O, is best explained by the theory that it is a union of the
dyad and tetrad (Pb" and Pbiv) , Pb.O* = 2PbO 4- PbivO, .
30 LEAD. §57, 5a.
Lead dioi^ide or peroxide, PbO, , is formed: (1) by fusion of PbO with KCIO,
or KNO, ; (2) by fusing PbgO^ with KOH : (3) by treating" any compound of
Pb*' with CI, Br, K,Fe(CN)« , KMnO^ , or HjO, in presence of KOH; (4) by
treating Pb^O^ with non-reducing acids:
Pb,0, -f 4HN0, = PbO, + 2Pb(N0,), + 2H3O.
Ignition forms first PbaO^ and above a red heat PbO, oxygen being given off.
It dissolves in acids on same conditions as Pbs04 . Very strong solution of
potassium hydroxide, in large excess, dissolves it. with formation of " potassium
plumbate," KaPbO, . Lead dioxide is a powerful oxidizing agent, one of the
strongest known. Digested with ammonium hydroxide, it forms lead nitrate
and water. Triturated with one-sixth of sulphur, or tartaric acid, or sugar,
it takes fire; with phosphorus, it detonates.
5. Solubilities. — a. — Met^^ — Nitric acid is the proper solvent for metallic lead,
the lead nitrate formed is readily soluble in water but insoluble in concentrated
nitric acid*: hence if the conct.itrated acid be used to dissolve the lead, a
white residue of lead nitrate will be left which dissolves on the addition of
water. Dilute sulphuric acid is without action, the concentrated acid is almost
without action in the cold (Calvert and Johnson, J, C, 1863, 16, 66), but the hot
concentrated acid slowly changes the metal to the sulphate with evolution of
sulphur dioxide, a portion of the salt being dissolved in the acid, precipitating
on the addition of water. Hydrochloric acid very slowly dissolves the metal
(more rapidly when warmed), evolving hydrogen; the chloride formed dissolves
in the acid in quantities depending upon conditions of temperature and con-
centration (r). The halogens readily attack the metal forming the correspond-
ing haloid salts. Alloys of lead are best dissolved by first treating with nitric
acid, if a white residue is left it is washed with water and, if not dissolved, it
is then treated with hydrochloric acid, in which it will usually be soluble.
Water used for drinking or cooking purposes should not be allowed to stand
in lead pipes. Pure water free from air is without action upon pure lead, but
water containing air and carbon dioxide very slowly attacks lead, forming the
hydroxide and basic carbonate. This action is promoted by the presence of
salts, as ammonium nitrate, nitrite, chloride, etc.: the action seems to be
hindered by the presence of sulphates.
h,—()j'idc8.—'Lead oxide, litharge, PbO , and the hydroxides, 2PbO.H,0:
SPbO.H.O, are readily dissolved or transposed by acids forming the correspond-
ing salts, i. ۥ., PbO -h H2SO4 == PbSO* -f H^O . The oxide and hydroxide are
soluble in about 7000 parts of water, to which they impart an alkaline reaction.
Thev are soluble in the fixed alkalis forming plumbites; soluble in certain salts
as NH,C1, CaCla, and SrCL (Andre, T. r., 1883, 96, 435; 1887, 104, 359); very
soluble in lead acetate, forming basic lead acetate.
IJead dioxide, PbOa , lead peroxide, is insoluble in water or nitric acid: it is
dissolved by the halogen hydracids with liberation of the halogen and reduction
of the lead forming a dyad salt: PbO^ -f- 4HC1 = PbCI, -f CI, + 2H2O; it is
attacked by hot concentrated sulphuric acid, forming the sulphate and liberat-
ing oxygen; it is soluble in glacial acetic acid forming Pb(C2H,03)4 , unstable
(Hutchinson and Pollard, J. C, 1896, 69, 212). Some of the salts of the tetrad
lead seem to be formed when the peroxide is treated with certain acids in the
cold. Thev are. however, very unstable, being decomposed to the dyad salt
upon warming (Fischer, ,/. r.,*1879, 35, 282; Nickels, A, Ch., 1867, (4), 10, 328).
The peroxide is slowlv soluble in the fixed alkali hvdroxides forming plum-
bates, i. e., PbO, + 2K0H = K^PbO, + H.O .
Trilead tetroxide, PbaO^ , red lead, mitiium, is insoluble in water, is at-
tacked by nearly all acids in the cold forming the corresponding dyad lead
salt and lead peroxide, PbOj . Upon further treatment with the acids using
heat the lead peroxide is decomposed as described above. The presence of
many reducing agents, as alcohol, oxalic acid, hydrogen peroxide, etc., greatly
* The solubility of a salt is lessoned by the presence of another substance having an ion in
common with it i^fi). In some cases, as with Pbl, and KI, this is offset in conoentrated aolution
by formation of a complex comjx)und.
S57, OC. LEAD. 31
facilitates the solution of red lead or lead peroxide in acids, (. e., nitric acid
does not dissolve lead peroxide, but if a few drops of alcohol be added the
solution is readily obtained upon warming, leaving the lead as the soluble
nitrate, which greatly facilitates the further analysis.
c. — Salts. — The carbonate, borate, cyanide, ferrocyanide, phosphate, sul-
phide, sulphite, iodate, chromate, and tannate are insoluble in water.
The sulphate is soluble in about 21,000 parts of water at 18° (Ivohlrausch
and Rose, Z. phys, Ch., 1893, 12, 241), the presence of HNO., or HCl in-
creases its solubility in water; it is insoluble in alcohol even when quite
dilute; sparingly soluble in concentrated H2SO4, from whicli solution it is
precipitated by the addition of water or alcohol; less soluble in dilute HgSO^
than in water; soluble in 682 parts 10 per cent HCl, in 35 parts 31.5 per
cent (Rodwell, J. C, 1862, 15, 59); transposed and dissolved by excess of
HCl , HBr , or HI forming the corresponding haloid salt ; insoluble in
HP (Ditte, A. Ch., 1878, (5), 14, 190); soluble in ammonium sulphate,
nitrate, acetate, tartrate and citrate, and from these solutions not readily
precipitated by ammonium hydroxide or sulphate (Fleischer, J. C, 1876,
29, 190; Woehler, .4., 1840, 34, 235). The sulphate is almost completely
transposed to the nitrate by standing several days with cold concentrated
nitric acid (Rodwell, Z. c). The oxalate is sparingly soluble in water, insol-
uble in alcohol ; the ferricyanide is very slightly soluble in cold water, more
soluble in hot water; the chloride is soluble in 85 parts water at 20° and in
32 parts at 80° (Ditte, C. r., 1881, 92, 718); the bromide is soluble in 16G
parts water at 10°, in about 45 parts at 80°; the iodide is soluble in 1235
parts water at ordinary temperature, and in 194 ])arts at 100° (Denot, J,
pr,, 1834, 1, 425). The* chloride is less soluble in dilute HCl or H.SO4 than
in water, but is more soluble in the concentrated acids (Ditte, /. c.) ; HNO,
increases the solubility of the chloride more and more as the HNO., is
stronger. The chloride is less soluble in a solution of NaCl than in water
(Field, J. C, 1873, 26, 575): soluble in NH.Cl —90 grams dissolving in 200
grams NH.Cl with 200 cc. water (Andre, C. r., 1893, 96, 435). The chloride,
bromide, and iodide are insoluble in alcohol. The iodide is moderately
soluble in solutions of alkali iodides; it is decomposed by ether. The
basic acetates are permanently soluble if carbonic acid is strictly excluded.
The basic nitrates are but slightly soluble in water, and are precipitatinl
on adding solutions of KNO3 to a solution of basic lead acetate.
The relative insolubility of PbCl, in cold water or in dilute HCl makes
it possible to precipitate the most of the lead (by means of HCl) from
solutions containing also the other metals of the Silver Group; while its
solubility in hot water is the means of its separation from the other
< hlorides of that group (§61). The lead is separated and identified in
the second group as the insoluble sulphate. (§96).
32 LEAD. §57, ().
6. Beaotions. a. — ^Fixed alkali hydroxides precipitate, from solutions of
lead salts, basic lead hydroxide (i), Pb20(0H)2 (Schaffner, A., 1844, 61, 175),
white, soluble * in excess of the reagent as plumbite (2) (distinction from
silver, mercury, bismuth, copper, and cadmium). The normal lead hy-
droxide, Pb(0H)2 , may be formed by adding a solution of a lead salt to
a solution of a fixed alkali hydroxide.
(i) 2Pb(N0,), 4- 4K0H = Pb,.0(OH), + 4KN0, + H,0
(2) Pb,0(OH), + 4K0H = 2K,PbO, + 3H,0 .
Ammonium hydroxide precipitates white basic salts, insoluble in water
and in excess of the reagent (distinction from silver, copper, and cad-
mium); with the chloride the precipitate, insoluble in water, is
PbClj.PbO.HjO (Wood and Bordeu, C. N., 1885, 62, 43); with the nitrate
2PbO.Pb(N03)2 (I^-y 2, 2, 358). With the acetate, in solutions of ordinary
strength, excess of ammonium hydroxide (free from carbonate) gives no
precipitate, the soluble tribasic acetate being formed.
Alkali carbonates precipitate lead basic carbonate, white, the composition
varying with the conditions of precipitation. With excess of the reagent
and in hot concentrated solutions the precipitate consists chiefly of
Pb,(0H)2(C03)2 . Precipitation in the cold approaches more nearly to the
normal carbonate (Lefort, Pharm. J., 1885, (3), 15, 26). Solutions of lead
salts when boiled with freshly precipitated barium carbonate are com-
pletely precipitated. Carbon dioxide precipitates the basic acetate but
not completely.
ft. — Oxalic acid and alkali oxalates precipitate lead oxalate, PbCjO^, while y
from solutions of lead salts, soluble in nitric acid, insoluble in acetic acid.
A solution of lead acetate preci})itates a large number — and a solution of
lead subacetate a still larger number — of organic acids, color substances,
resins, gums, and neutral principles. Indeed it is a rule, with few excep-
tions, that lead subacetate removes organic acids (not formic, acetic,
butyric, valeric, or lactic). Tannic acid precipitates solutions of lead
acetate, and of the nitrate incompletely, as yellow-gray lead tannate,
soluble in acids.
Soluble cyanides precipitate lead cyanide, Pb(CN)2 , white, sparingly soluble
in a large excess of the reagent and reprecipitated on boiling*. Potassium ferro-x^
cyanide precipitates lead ferrocyanide, Pb^FeCCN), , white, insoluble in wat^r
or dilute acids. Potassium ferricyanide precipitates from solutions not too
dilute lead ferricyanide, Pbs(Fe(CN)o)2 , white, sparingly soluble in water,
soluble in nitric acid. Solutions of lead salts are precipitated by potassium
sulphocyanate as lead sulphocyanate, PbCCNS)^ , white, soluble in excess of the
reagent and in nitric acid.
c. — Lead nitrate is readily soluble in water, and dissolves the oxide to form
the basic nitrate, which may also be formed by precipitating lead acetate with
♦Nparlyall the salts of lead are soluble in the flxed alkali hydroxides, PbS forminfir almoflt
the only notable exception.
§67, 6e. LEAD, 33
potasaium nitrate. The solubility of lead nitrate is greatly increased by the
presence of the nitrates of the alkalis and of the alkaline earths, a complex
compound being formed (Le Blanc and Noyes, Z. phy8, Ch., 1890, 6, 385).
d.— The higher oxides of lead are all reduced by hypophosphoroua acid, lead
phosphate being formed. Lead phosphite, PbHPO, , white, is formed by
nearly neutralizing phosphorous acid with lead carbonate or precipitating
J7a,HP0, with Pb(NO,)a (Amat, C. r., 1890, 110, 901). Sodium phosphate,
KatHFO« , precipitates from solutions of lead acetate the tribasic lead phosphate,
Pb.CPO*), , white, insoluble in the acetic acid which is set free (D., 2, 2, 562):
3Fb(C,H,0,), + 2Ka,HF0« = Pb.CPOJ, + 4NaC2H,0, + 2HC,H,0,. The same
precipitate is formed when sodium phosphate is added to lead nitrate, soluble
in nitric acid, insoluble in acetic acid. Lead phosphate is also precipitated
upon the addition of phosphoric acid to solutions of lead acetate or lead nitrate.
The pyrophosphate, PbaP.OT , white, amorphous, is formed by precipitating a
lead solution with TSiB.^^tO^ , soluble in excess of the precipitant, in nitric acid,
and in potassium hydroxide; insoluble in ammonium hydroxide and in acetic
acid (Gerhardt, A. CK, 1849, (3), 25, 305). The metaphosphate, Pb(PO,), ,
white, crystalline, is obtained by the action of NaPO, upon PbCNO,), in exces^.
e. — ^Hydrosnlphnrio acid and the soluble sulphides precipitate — from
neutral, acid, or alkaline solutions of lead salts — lead sulphide^ PbS,
brownish black, insoluble in dilute acids, in alkali hydroxides, carbonates,
or sulphides. Freshly precipitated CdS, lUnS, FeS, CoS, and NiS also
give the same precipitate. Hydrosulphuric acid and the soluble sulphides
transpose all freshly precipitated lead salts to lead sulphide.* Moder-
ately dilute nitric acid— 15 to 20 per cent — dissolves lead sulphide with
separation of sulphur (i), some of the sulphur, especially if the nitric acid
be concentrated, is oxidized to sulphuric acid, which precipitates a portion
of the lead (^), unless the nitric acid be sufficiently concentrated to hold
that amount of lead sulphate in solution. The oxidation of sulphur always
occurs when nitric acid acts upon sulphides, and in degree proportional
to the strength of acid, temperature, and duration of contact.
(!) 6PbS + 16HN0, = 6Pb(N0,), + 3S, + 4N0 + 8H,0
(2) 3PbS + SHNO, = SPbSO, + 8N0 + 4H,0
In solutions too strongly acidulated, especially with hydrochloric acid,
either no precipitation takes place, or a brick-red double salt, PbaSClg ,
* The oondition for equilibrium is that a certain ratio of concentration exist between the loos,
in the case of FbBOf between the 8 ions and the 8O4 ions. These concentrations are the same
as those in a solution obtained by digesting the two salts, PbS04 and PbS, together in water.
Pb90« dissolves more freely than PbS. and for equilibrium therefore *gQ ,, must be corres-
pondincriy firroater than ^^„. But adding H,S or a soluble sulphide to PbS04 gives just the
opposite of this condition, and transformation accordingly results* increasing the S04^' con-
centration by formation of soluble sulphate and deicreasing the S'^ concentration by precipita-
tion of PbS, until the equiUbrium-ratio is produced or, if the quantity of PbS04 present is in-
sufScient for this, until all the PbS04 has been transformed to sulphide. On the other hand,
treatment of PbS with a very large excess of HsS04 will cause the reverse action, S ions going
into solution until the same equilibrium results as before.
The general principle is then that unless a constituent of the more soluble substance is in
great preponderance in the solution the least poluble of two or more possible products will
slways be formed. This principle determines the direction in which a reaction takes place ;
AfCn + KI - Agl + KCl I CaS04 + Ka.CO, = CaCO, + Na,S04 ($4^).
34 LEAD. §67, iif.
is formed, the precipitation being incomplete. In neutral solutions con-
taining 100,000 parts of water lead is revealed as the sulphide; a test
which is much more delicate than the formation of the sulphate.
Ferric chloride decomposes lead sulphide, forming lead chloride, ferrous
chloride and sulphur. The reaction takes place in the cold and rapidly when
warmed (Gabba, C. C, 1889, 667).
When galena, PbS , is pulverized with fused KHSO« , HjS is evolved (Jan-
uettaz, J. C, 1874, 27, 188).
Lead thiosulphate, PbS^O, , white, is precipitated by adding sodium thiosul-
phate to solutions of lead salts: the precipitate is readily dissolved in an excess
of the reagent, forming the double salt, PbS20„2Naa820, (Lenz, A., 1841, 40,
94); on boiling, all the lead is slowly precipitated as sulphide (Vohl, J.., 1855,
96, 237).
Sodium sulphite precipitates lead sulphite, PbSOg , white, less soluble in
water than the sulphate, slightly soluble in sulphurous acid; decomposed
by sulphuric, nitric, liydrochloric, and hydrosulphuric acids and by alkah
sulphides; not decomposed by cold ])ho8phoric and acetic acids.
Sulphuric acid and soluble sulphates precipitate from neutral or acid
solutions, lead sulphate, PbSO^ , white, not readily changed or permanently
dissolved by acids, except hydrosulphuric acid, yet slightly soluble in
strong acids (5c). Soluble in the fixed alkalis and in most ammonium
salts, especially the acetate, tartrate, and citrate (Woehler, A,, 1840, 34,
235). Soluble in warm sodium thiosulphate solution, in hot solution
decomposed, lead sulj)hide, insoluble in thiosul])hate, being formed (dis-
tinction and separation from barium sulphate, which does not dissolve in
thiosulphates).
The test for lead as a 8ul])liate is from five to ten times less delicate
than that with hydrosulphuric acid; but lead is quantitatively separated
as a sulphate, by precipitation with sulphuric acid in the presence of
alcohol, and washing with alcohol. When heated with potassium chromate
transposition takes place and yellow lead chromate is formed (j^). Excess
of potassium iodide transposes lead sulphate (f), a distinction of lead from
barium. Repeated washing of lead sulphate with a solution of sodium
chloride com])letely transposes the lead to the chloride (Matthey, J. C,
1879, 36, 124). See footnote on previous page.
f. — Hydrocliloric acid and soluble chlorides precipitate, from solutions
not too dilute, lend chloride, PbCL , white. This reaction constitutes lead
a member of the FIEST GROUP— as it also is of the second. The solu-
bility of the precipitate is such (5r) that the filtrate obtained in the cold
gives marked reactions with hydrosulphuric acid, sulphuric acid, chro-
mates, etc.; and that it can be quite accurately separated from silver
chloride and mercurous chloride by much hot water. Also, small propor-
tions of lead escape detection in the first group, while its removal is
necessarily accomplished in the second group.
i^67, 7. LEAD. 35
Hydrobromic acid and soluble bromides precipitate lead bromide, PbBrg ,
white, somewhat less solublu in water than the chloride (oc); soluble in
excess of concentrated potassium bromide, as 2KBr.PbBr2 , which is decom-
posed and PbBr2 precipitated by dilution with water.
Hydriodio acid and soluble iodides precipitate lead iodide, Pbig , bright
yellow and crystalline, much less soluble in water than the chloride or
bromide (oc); soluble in hot moderately concentrated nitric acid and in
solution of the fixed alkalis; soluble in excess of the alkali iodides, by
forming double iodides, JQPblo with small excess of KI , and 4EI.Pbl2
with greater excess of KI ; these double iodides are decomposed by addi-
tion of water with precipitation of the lead iodide. Lead iodide is not
precipitated in presence of sodium citrate; alkali acetates also hold it in
solution to some extent, so that it is less perfectly precipitated from the
acetate than from the nitrate. Freshly precipitated lead peroxide, PbOa ,
prives free iodine when treated with potassium iodide (Ditte, C. r., 1881,
93, 64 and 67).
In detecting lead as an iodide in solutions of the chloride by precipita-
tion with potassium iodide and recrystallization of the yellow precipitate
from hot water, care must be taken that the potassium iodide be not
added in excess to form the soluble double iodides.
ff, — ^Arsenous acid does not precipitate neutral solutions of lead salts; from
alkaline solutions or with soluble arsenites a bulky white precipitate of lead
arsenite is formed, insoluble in water, but readily soluble in all acids and in the
fixed alkali hydroxides. Arsenic acid and soluble arsenates precipitate lead
arsenate, white, from neutral or alkaline solutions of lead salts, soluble in the
fixed alkali hydroxides and in nitric acid, insoluble in acetic acid. For the
composition of the arsenites and arsenates of lead see (D., 2, 2. 565). Hot
potassium stannite (SnClj in solution by KOH) gives with lead salts or lead
hydroxide a black precipitate of metallic lead.
h. — Chromic acid and soluble chromates — both K2Cr04 and K^Ct.Oj —
precipitate lead cliromate, PbCr04 , yellow, soluble in the fixed alkali
hydroxides (distinction from bismuth), insoluble in excess of chromic acid
(distinction from barium), insoluble in ammonium hydroxide (distinction
from silver), decomposed by moderately concentrated nitric and hydro-
chloric acids, insoluble in acetic acid.
7. Ignition. — Lead salts when fused in a porcelain crucible with sodium
carbonate are converted into lead oxide, PbO {a). After fusion and diges-
tion with warm water, the aqueous solution is tested for acids, and the
residue for bases after dissolving in nitric or acetic acid. If charcoal (or
some organic compounds as sugar, tartrates, etc.) be present, metallic lead
is formed (ft); and with excess of charcoal the acid radicle may also be
changed (c). If the fusion with sodium carbonate is made on a piece of
charcoal, instead of in a crucible, using the reducing flame of the blow-
36 LEAD, §57, 8.
pipe, globules of metallic lead are produced and at the same time the
charcoal is covered with a yellow incrustation of lead oxide, PbO .
(o) PbCl, + Na^CO, = 2NaCl + PbO + CO,
(6) 2PbS04 + 2Na,C0. + C = 2Pb -f 2Na,S0, + 3CO,
(r) 2PbSO, + 2Na2CO, + 5C = 2Pb + 2Na,S + 7C0,
8. Detection. — Lead is precipitated, incompletely, from its solutions by
HCl as PbClg ; separated from AgCl and HgCl by hot water, and confirmed
by HjS , JS2SO4 , KjCrO^ , and KI . It is separated (in the second group)
from As, Sb, Sn, etc., by non-solubility of the sulphide in (1114)28,;
from HgS by HNO3 ; from Bi , Cu , and Cd by precipitation with dilute
sulphuric acid. Insoluble compounds arc transposed by an alkali sulphide,
being then treated as lead in the second group, or they are examined by
ignition as described in (7).
9. Eatimation. — («) As an oxide into which it is converted by ignition (if a
carbonate or nitrate), or by precipitation and subsequent ignition. (6) As a
sulphate. Add to the solution twice its volume of alcohol, precipitate with
H3SO4 , and after washing with alcohol ignite and weigh, (r) It is converted
into an acetate, or sodium acetate is added to the solution, then precipitated
with KaCTjOT , and after drying at 100°, weighed as PbCrO* . (d) It is con-
verted into PbS , free sulphur added, and after ignition in hydrogen gas
weighed as PbS . (f) The lead is precipitated with standardized sodium iodate
and the excess of io<late is determined by retitration. Lead iodate is less
soluble in water than lead sulphate (Cameron, J. C, 1879, 36, 484). (f) In
presence of bismuth, ignite the halogen compound, or convert into a sulphide
and ignite in a current of bromine. The haloid salts of bismuth sublime upon
ignition (Steen, Z. amjew,^ 1895, 530). (</) Gas roluinetric method. Precipitate as
a chromate, filter, wash and transfer to an azotometer with dilute sulphuric
acid and estimate the amount of chromium by the volume of oxygen set free
by hydrogen peroxide (Baumann, Z. anpew., 1891, 329).
10. Oxidation. — Metallic lead precipitates the free metals from solutions
of Hg , Ag , An , Pt , Bi , and Cu . Lead as a dyad is oxidized to the
tetrad as stated in (4), also electrolytically in separation from Cu (Nissen-
son, Z. angew.y 1893, 646). PV^ is reduced to Pb° in presence of dilute
H2SO4 by nascent hydrogen, and by all metals capable of producing nascent
hydrogen (such as Al , Zn , Sn , Mg , Fe), and to Pb" by soluble compounds
of Hg', Sn", Sb'", As'", (AsH,, gas), Cu', Fe", Cr"', Mn", Mn'", Mn^^
Mn^i. Also by H^Q.O, , HNOo , H3PO2 , H3PO., , P , SO^ , HjS , HCl , HBr ,
HI, HCN, HCNS, H4Fe(CN)«, glycerine, tartaric acid, sugar, urea, and
very many other organic compounds. In many cases the reduction to
Pb" or to Pb° takes place in presence of KOH . The freshly precipitated
peroxide oxidizes ammonia, NH3 , to nitrite and nitrate in the course of a
few hours (Pollacci, Arch. Pharm., 1886, 224, 176).
From lead solutions Zn , Mg , Al , Co , and Cd precipitate metallic lead.
568, 5a. MERCURY. 37
§68. Mercury (Hydrargyrum) Hg = 200.0 . Valence one and two.
1. Properties.— Spfcf/?c gravity, liquid, 13.5953 (Volkmann, W. A,, 1881. 13, 209) ;
solid, 14.1932 (Mallet, Proi\ R. ISoc., 1877, 26, 71). Melting (freezing) point, —38.85*^
(Mallet, Phil. Mag., 1877, (5), 4, 145). Boiling point, 357.33° at 760 mm. (Ramsay
and Young, J, C, 1885, 47, 657). It is the only metal which is a liquid at
3rdinary temperatures, white when pure, with a slightly bluish tinge, and
having a brilliant silvery lustre. The precipitated or finely divided mercury
appears as a dark gray powder. Mercury may be " extinguished " or " dead-
ened," i. e., reduced to the finely divided state, by shaking with sugar, grease,
zhalk, turpentine, ether, etc. It is slightly volatile even at — 13** (Kegnault,
C. r., 1881, 93, 308); is not oxidized by air or oxygen at ordinary temperature
(Shenstone and Cundall, J. C, 1887, 51, 619). The solid metal is composed of
octahedrgil and needle-shaped crj'stals, is very ductile and is easily cut with a
knife. Owing to its very strong cohesive property it forms a convex surface
with glass, etc. It is a good conductor of electricity, and forms amalgams with
Al, Ba, Bi, Cd, Cs , Ca , Cr , Co, Cu , Au , Fe, Pb , Mg, Mn, Ki , Os ,
Pd, Pt, K,.Ag, Ka , Tl, Sn , and Zn . An amalgam containing about 30
per cent of copper is used for filling teeth (Dudley, Proc, Am, Asac. for Adi\ of
8rt., 1889, 145).
2. Occorrence. — The principal ore of mercury is cinnabar, HgS , red, found in
California, Illyria, Spain, China, the Ural, and some other localities. The free
metal is sometimes found in small globules in rocks containing the ore. It is
also found amalgamated with gold and silver, and as mercuric iodide and
mercurous chloride.
3. Preparation. — (a) The ore is roasted with regulated supply of air: HgS -|-
Oj = lELg -\- SO,, (ft) Lime is added to the ore, which is then distilled:
4Hg8 -h 4CaO = 3CaS + CaSO^ + 4Hg . (c) The ore is heated with iron
(smithy scales): Hg , FeS , and SO, are produced. The mercury is usually con-
densed in a trough of water. Commercial mercury is freed from dirt and other
impurities by pressing through leather or by passing through a cone of writ-
ing paper having a small hole in the apex. For the separation of mercury
from small quantities of Pb , Sn , Zn , and Ag without distilling, see Briihl (B.,
1879, 12, 204), Meyer (B., 1879, 12, 437), and Crafts (B/.» 1888, (2), 49, 856).
4. Oxides.— Mercury forms two oxides, Hg.O and HgO . Mercurous oxide,
HgoO , is a black powder formed by the action of fixed alkalis on mercurous
salts. It is converted by gentle heat into Hg and lELgO and by a higher (red)
heat, to Hg and O. Mercuric oj'ide, HgO, is made (/) by keeping Hg at its
boiling point for a month or longer in a flask filled with air; (2) by heating
HgNO, or HgfKOs), with about an equal weight of metallic mercury:
Hg(N08)j -f 3Hg = 4H£^ + 2N0; {3) by precipitating mercuric salts with
KOH or NaOH . Made by (/) and (2) it is red, by (S) yellow. On heating it
changes tp vermillion red, then black, and on cooling regains its original color.
A red heat decomposes it completely into Hg and O . Mercurj' forms no
hydroxides.
5. Solubilities. — a. — Metal. — Unaffected by treatment with alkalis. The most
effective solvent of mercury is nitric acid. It dissolves readily in the dilute
acid hot or cold; with the strong acid, heat is soon generated; and with con-
siderable quantities of material, the action acquires an explosive violence. At
ordinary temperatures, nitric acid, when applied in excess, produces normal
mercuric nitrate, but when the mercury is in excess, mercurous nitrate is
formed; in all cases, chiefly nitric oxide gas is generated. Both mercurous and
mercuric nitrates require a little free nitric acid to hold them in solution.
This free nitric acid gradually oxidizes mercurosum to mercuricum, making a
clear solution of Hg(N0a)2 , if there is sufficient HNO, present, otherwise a
basic mercuric nitrate may precipitate. A solution of mercurous nitrate may
be kept free from mercuric nitrate by placing some metallic mercury in the
bottle containing it; still after standing some weeks a basic mercurous nitrate
crystallizes out, which a fresh supply of nitric acid will dissolve. Sulphur
attacks mercury even in the barometric vacuum, forming HgS (Schrotter,
J. C, 1873, 26, 476). H3SO4 concentrated at 25** has no action on Hg (Pitman,
S^ JiERCCRY. §68, 56.
.7. A»i. ^w-. IfSti. 20. 1C«C*)- ^JTh the hot concratnted acid SO, is evolved and
Hg.SO, is fomw^ if Hg be in irreat exc«*; HgSO. if the H^SO. be in excess.
Hi ±rcit-±tjoric acid irsis at 2ti0= is without action iBerthelot, ^1. C*., 1856, (3), 46»
J.H£ i: tiio lie acid >p.. gr., 1^«. Railev and Fowler iJ. C, ISS^, 53, 759) say that
err hx crcirfLlonc acid eas in pre*<-nct- of **\\sezi and mercury, at ordinary' tem-
;»-niTLrt frr rhnrr weeks. form> H^.OCl. withoTit evoln'tion of hydrogen:
'-^S^ — iHCl — O; = H^.OCl^^-O . Hydrc«hroni:-c and hydriodic acids, g^ses,
i»^<"L tr-£.;i; nrrczrv. rvolrf H. and fi-rm respectively H^Br and H^I (Ber-
Vr'"'" -' ' '', H;. dTvts:.:-.h-3Tic arid. Htv sras. at l*yj- does not attack dry Hg
i:i*T"b*K'. \' . H.v-:r«-t>.:ly.hiiric acid, in solution, and alkali sulphides form
S^ - *- -"tt"' ■^. - r« TTjIr.-- ar.^ i •^ine. dry or mo:>?. attack the metal: merourous
t^r.^ hT*r f-.rr:.- if -L- n:rT\-jry W in rxc<»ss^ mercuric salts if the halogens be
1. — ^*.-id'*.— Mercarons oxide :<i in^-Oubje It. water -^r alkalis. Hydrochloric
tx'jjt fi-m:* HgCl: -'^^pburio acid fonns Hg.SO, . changed by boiling with
*x^•r«^ :-f t: :i -: HgSO.: nitrii a<id fomis HgVO, . changed by excess of acid
"tr- H^ XO; .. . Merauric oxide i< sol-ble in acSc*. ir.soluble in alkalis, soluble
-z. i' . II -r. '.',** 7iar:s water <B:neaT:. •\ «•_ I^^^, 41. 50&>- It is decomposed
:t h'\.hl cr::T-:iT-s*f.:.rm:nxr HgCl* .Mia^be. .1. •'!., l>4f, (5>. 5, 177). soluble in
IHC . fr:~ wh, ch =^.! :-5on XS.OH i*r«^ip:tates HH«CUrH{rH,Cl +
^fH.HgC: T^-r. '. '.. >-i. 112. -T.i'j. s^.rjble in KI, forming iKI^Hgl,
•^ti^ /. •_ :-:i:.25, i^ZK
• — i-*7i. — Mfr:~rT f-.rTi:> rv-- trt-:" ir^rkt-: .la??^ of salts — mercurous,
zlili-^lItI.:- i::i •-.^'-.-^K;. iiTilvi:! — ^r:-.»>: Tr—r: iit^vjs o^iiipcumds are per-
n-L^e*: iz. 'JL- iir. '*:-: :.r\ .h^zijcv' '-y >^Tvrf^j: .>xidi2iDg agents to
iiriTi-zr., .':z:t« :zLlf. Thv Icr.UT srv- >>n:rTha: m»:»re stable, but are
.mjutt*: "■- i^i-j rvi-riiii: ajr:::>. ::rsT :-> njrrv-r->-> i:x»ni pounds and then
: ZLr'.Ll:: li rurr !'• . <>>-:: r.s -: !r.vrv-urT sahs reilden litmus.
Xui^ : :lr -.i!> : n-rvur}- .:>:• r.rhtr ::.-y.uVc :n water, or require the
: Tz^z.' :i :rv i.i 1 :> k-.-tp :h-.:v. :r. > !u::. r,. 'vir.^ ^i-evomposed by water
i: L i-rr-Lii. :.jT\-r : •liluiior.. yr-.- .;. -:^::rr a Sasio salt and leaving an
b j'. T,Lz -JL -:*.-:.". Hemux>iis /./.^r:>:. " r^:r.:*ir. a::«i i«>dide are insolu-
- ..-_ Ti.: r: :!•. >u>. '..:..:•. :> > .. " • *•■. ''•- .Mns ov-l and 3«>l parts hot
vi>7. «. l-^T Lz ::::::. -tri: ;u: i W,:.kv::r^:tT. J.. l>4e, 41, 319). The
ii.t^iir u.r !."••:-: :'::i >./■.-. --^'u* :':::^> ^-.s :>.i >::".y-hi:c. Meniurous nitrate
.* •utI-^ItI^ r- '.M* '• ::: -r.^Ttr. Or. >:Ar- Mr.^ :: jmiduallv changes to
.zi:'T. -IT.' ::_ : ri : T . : > v , r. : > *. ' y : >. ^ - t^ >< :: . > ->: : rw nieniury. bu t i f free
:i.'!r:Tr- - tt-^ -: ■ ■». y::o:-: :' \->:. r.:eTv*j^r^^u> nitrate gradually
: ni.-. Menric Mr::-. i> << 13"-. :r. !•> y<£r:> ^: o^x-i water and 3 parts
• Tw Zjl-w :t ILags^JL :^:c. r*»;u:r**:iai ^^b^re "rijs^^-vc^Lrjrrnt? :t * slifkUj-ioniaed substance
irr» ir««**!ir -111- *>r:iRai>* 5a;i._ ::rzi *: t^ t».\:.^c:?^ .: :ii\-«» ».-«>f fnvkQ^ty kmiied. Such a
.:.*/nr. --J r .^i :...t7 ^ H«C1.^ w^^c: B|cO ;* ^r•..:^: -:isv> cvsKk^*^ wiih KO solution H^ and
••-i -ni;-!!:** - - .cni -J^ :=»;c-^i:;asi.vi4:'-*i HcCl^ .^«' ir^ WL aad Ot v^sch unite with water, im-
^ % r— r.if - -J., i. . i'n: c i *zr':c:f Alfcfcl.r* w*.-, ?;r- KHr acii Kl ac? -f-nM aw^re strongly. H^O.
».• :ii 1UZ1 -^ m - u- r"^:.- ijt:rcip*.^i<c:-'*c \\; .:'* «l.::«v >y water aoi iKCi Us etcy rpducibility a
r.-.i^ -Ai^. '.r T-l^ ?»7ufc;r^ :if A A*:. r:-,-c*-s * i-er? a ..^r:v-di5SK.via9e«i B« eoBpoand results.
i.- •: r*-- T* B^SO, , i»ji.M.v* vi-.r...V. "^ .-^.cTji^. *&i xxlxSe of ■« aad A^ owing to the
<ani»* raiL-^. -nt* Be' ^^'o:^ - '^-^ Mr,^"^\> i:.-«t.vdt««l r.'rtte JWvonftstimr the afaready slight
'...-?- .-uir.. n r -j** ■ai-r'.-sirx'' iJLv-'.U »^^ - I^*? tfju'-re .*=;f il«Cl» t*> ^ve Bany of the pre-
•■n'":ar.i a- -"'UTnira.* -.-ruai^.V v.:i o:ier AV.:si* »rtv';ir.vr *fclw » oi jv^one due to the same
jJ58, 6a. MEIiCVRY. 39
warm water; the bromide is soluble in 94 parts water at 9° and 4-5 parts
at 100**, decomposed by warm nitric or sulphuric acids; the iodide is
soluble in about 25,000 parts water (Bourgoin, A. Ch, 1884 (6), 3, 4?9),
soluble in Na^S^Os (Eder and Ulen, J C, 1882, 42, 80G), and in many
alkali salt^, forming double salts. Normal mercuric sulphate is decom-
posed by water into a soluble acid sulphate and the basic sulphate, HgSO^ ,
^HgO , which is practically insoluble (soluble in 43,4T8 parts water at
16°, Cameron, Analyst, 1880, 144). The normal nitrate is deliquescent,
very soluble in a small amount of water, but more water precipitates the
nearly insoluble basic nitrate, BHg^.N^Os > changed by repeated washing
into the oxide, HgO (Millon, A, Ch., 1846 (3), 18, 361). The basic nitrate
is soluble in dilute nitric acid. The cyanide is soluble in eight parts water
at 15**. The acetate is readily soluble, the chromate and citrate sparingly,
and the sulphide, iodide, iodate, basic carbonate, oxalate, phosphate, arse-
nate, arsenite, ferrocyanide, and tartrate are insoluble in water.
6. Eeaotions. a. — Fixed alkali hydroxides precipitate, from solutions of
mercnrons salts, mercurous oxide, HgoO , black, insoluble in alkalis, readily
transposed by acids; from solutions of mercuric salts, the alkali, added
short of saturation, precipitates reddish-brown basic salts, when added in
excess, the orange-yellow mercuric oxide. HgO , is precipitated. If the
solution of mercuric salt be strongly acid no precipitate will be obtained
owing to the solubility of the mercuric oxide in the alkali salt formed; or,
in the language of the Dissociation Theory, owing to the slight dissocia-
tion of the soluble mercuric salt (§46). Ammonium hydroxide and car-
bonate precipitate from solutions of mercurous salts mixtures of mercury
and mercuric ammonium compounds: The same is true of the action of
ammonium hydroxide on insoluble mercurous salts : 2HgGl -f- 2NH^0H
= Hg + HHzHgCl -f 2H2O -f HH.Cl ; eHgNOg + eNH.OH = 3Hg -f
(MH-^HOOsHgO -f 4ira:,N0, -f 5H,0 : 4Hg,S0, -+: 8NH,0H = 4Hg -f
{liIgK^)^s6^.21iIeO -f 3(NH,)2S0, + GH.O ; or uniting the salt in dif-
ferent manner, 4HgCl -f 4NH,0H := 2Hg + Hg,NCl.NH,Cl + 2NH^C1
-f 4H2O . Examination with a microscope reveals the presence of Hg° .
The mercuric ammonium precipitate dissolves in a saturated solution of
(]IE4)oS04 containing ammonium hydroxide and can thus be separated
from the Hg (Francois, J. PJiarm,, 1897 (6), 5, 388; Turi, Gazzetta, 1893,
23, ii, 231; Pesci, Gazzetta, 1891, 21, ii, 569; Barfoed, J. pr.. 1889, (2), 39,
201). With mercuric salts ammonium hydroxide produces " white precipi-
tate," recognizable in very dilute solutions; that with cold neutral solu-
tions of mercuric chloride being mercurammonium chloride, (NHoHg)Cl ,
also called nitrogen dihydrogen mercuric chloride (a); if the solution be
hot and excess of ammonium hydroxide be added, dimercurammonium
chloride, also called nitrogen dimercuric chloride (h) is formed. Treat-
40 MERCURY. §68, G&.
ing with fixed alkali hydroxide until no more ammonia is evolved changes
the former compound to the latter (Pesci, I c). The precipitates are
easily soluble in hydrochloric acid, slightly soluble in strong ammonium
hydroxide, and more or less soluble in ammonium salts, especially am-
moniimi nitrate and carbonate (Johnson, C. iV., 1889, 69, 234). A soluble
combination of ammonium chloride with mercuric chloride, 2NH4CI.
HgClj, or ammonium mercuric chloride, called "sal alembroth," is not
precipitated by ammonium hydroxide, but potassium hydroxide precipi-
tates therefrom the white mercurammonium chloride, (HH3)2HgCl2 (c) :
(a) HgCla + 2NH,0H = NH.HgCl + NH.Cl + 2H,0
(6) 2HgCla + 4NH,0H = NHg,Cl + 3NH,C1 -f 4H,0
(c) 2NH,Cl.HgCl, 4- 2K0H = (NHOaHgCl, + 2KC1 + 2H,0
A solution of HgClj in KI with an excess of KOH (Nessler's Reagent) is
precipitated by NH^OH (or by ammonium salts), as NHgjI (§207, 6i).
Fixed alkali carbonates precipitate from merciirous salts an unstable mer-
eurous carbonate, HgzCO, , gray, blackening to basic carbonate and oxide when
heated. Carbonates of barium, strontium, calcium and magnesium precipitate
mercurous carbonate in the cold. Mercuric salts are precipitated as red-brotcn
basic salts, which, by excess of the reagent with heat, are converted into the
yellow mercuric oxide. The basic salt formed with mercuric chloride is an oxy-
chloride, Hg^l2.(Hg^)2 , >* or <; with mercuric nitrate, a basic carbonate,
(H£rO),HgC03 . Barium carbonate precipitates a basic salt in the cold, from
the nitrate, but not from the chloride.
b, — Oxalic acid and soluble oxalates precipitate from solutions of mercurous
salts mercurous oxalate, 'Rg2^204 , white, slightly soluble in nitric acid; from
solutions of mercuric salts, except HgCl, , mercuric oxalate, HgCjO^ , white,
easily soluble in hydrochloric acid, difficultly soluble in nitric acid. A solution
of HgCI, boiled in^the sunlight with (NH4)3C.04 gives HgCl and COj .
Hydrocyanic acid and alkali cyanides decompose mercurous salts into me-
tallic mercury, a gray precipitate, and mercuric cyanide, which remains in
solution. Mercuric salts are not precipitated, since the cyanide is readily
soluble in water. Soluble ferrocyanides form with mereurosum a white gela-
tinous precipitate, soon turning bluish green: with mercuricum a white precipi-
tate, becoming blue on standing. Soluble ferricyanides form with mercurous
salts a yellowish green precipitate; with mercuric salts a green precipitate,
soluble in hydrochloric acid. Potassium thiocyanate precipitates mercurous
thiocyanate, HgCNS , white, from solutions of mercurous salts (Clans, J. pr.,
1838," 15, 406); from solutions of mercuric .salts, mercuric thiocyanate,
HgCCNS), , soluble in hot water (Philipp, Z. Ch., 1867, 553).
r. — Nitric acid never acts as a precipitant of mercury salts, the salts being
more soluble in strong nitric acid than in water or the dilute acid; also nitric
acid dissolves all insoluble salts of mercury except HgS , which is insoluble in
the hot acid (sp. (jr, 1.42) (Howe, Am., 1887, 8, 75). HgCl is slowly dissolved by
nitric acid on boiling. All mercurous salts are oxidized to mercuric salts by
I'xcess of nitric acid.
</, — H3rpophosphorous acid reduces mercuric salts to Hg°, but the presence of
hydrogen peroxide causes the formation of HgCl from Hg^, and is of value
as a quantitative method for estimation of mercury (Vanino and Treubert, B.,
1S97, 30, 1099).
Phosphoric acid and alkali phosphates precipitate, from mercurous salt>..
mercurous phosphate, HgaPO^ , white, if the reagent be in excess; but if HgKO,
be in excess, HgjPO^.HgN'Oa , white, with a yellowish tinge. Mercurous pha<-
phate is soluble in dilute HNO, , insoluble in HsPO^ , From mercuric nitratv.
§58, 6e. MERCURY. 41
mercuric phosphate, Hg^CPOJ,, white, is precipitated, soluble in HNO, , HCl ,
and ammonium salts, insoluble in HsFO^ . Phosphoric acid does not precipitate
HgCl, , and Ka-^HPO^ does not precipitate the white HgsCFO^), from HgCl, ,
but on standing a portion of the mercury separates as a dark brown pre-
cipitate (Haack, J. C, 1891, 60, 400; 1892, 62, 530).
e. — ^Hydrosulphnrio acid and soluble sulphides, precipitate from mcr-
cnrous salts, mercuric sulphide, HgS , black, and mercury , gray. Mercurous
sulphide, HgjS, does not exist at ordinary temperatures. According to
Antony and Sestini {Oazzetta, 189-t, 24, i, 193), it is formed at — 10° by
the action of H2S on HgCl , decomposing at 0° into HgS and Hg . From
meronric salts there is formed, first, a white precipitate, soluble in acids
and excess of the mercuric salts, on further additions of the reagent, the
precipitate becomes yellow-orange, then brown, and finally black. This
progressive variation of color is characteristic of mercury. The final and
stable black precipitate is mercuric sulphide, HgS ; the lighter colored
precipitates consist of unions of the original mercuric salt with mercuric
t^ulphide, as HgCl2.HgS , the proportion of HgS being greater with the
.darker precipitates. When sublimed and triturated, the black mercuric
sulphide is converted to the red (vermillion), without chemical change.
Mercuric sulphide is insoluble in dilute HNO3 (distinction from all other
metallic sulphides); insoluble in HCl (Field, J. C, 18G0, 12, 158); soluble in
chlorine (nitro-hydrochloric acid) ; insoluble in (MH4)2S except when KOH
or NaOH be present (Volhard, .4., 1891, 255, 252); soluble in K^S (Ditte,
C. r., 1884, 98, 1271), more readily if KOH be present (separation from
Pb , Ag , Bi , and Cu) (Polstorff and Billow, Arch, Pharm., 1891, 229, 292).
It is soluble in KoCS., (one part S , two parts CSj , and 23 parts KOH , sp,
gr, 1.13) (separation from Pb, Cu, and Bi); reprecipitated as sulphide by
HCl (Rosenbladt, Z., 1887, 26, 15).
Mercurous nitrate forms with sodium thiosulphate a grayish black precipi-
tate, part of the mercury remaining in solution. Mercurous chloride forms
metallic mercury and some mercury salt in solution as double salt (Schnauss,
/. C., 1876, 29, 342). Mercuric chloride added to sodium thiosulphate forms a
white precipitate, which blackens on standing: if the mercuric chloride be
added in excess a bright yellow precipitate is formed, which blackens when
boiled with water, nitric acid or sulphuric acid, but does not dissolve or
blacken on boilinpr with hydrochloric acid. Sodium thiosulphate added to
mercuric chloride forms a white precipitate, which blackens on standinjjf or on
adding excess of thiosulphate, but if excess of thiosulphate be rapidly added to
HgCla no precipitate is formed; boiling or long standing produces the black
precipitate. Mercuric salts are not completely precipitated by sodium thio-
sulphate. The black precipitate is HgS.
Sulphurous acid and soluble sulphites form from mercurous solutions a
black precipitate (Divers and Shimidzu, ./. C, 1886, 49, 567). Mercuric nitrate
with sulphurous acid forms slowly a flocculent white precipitate soluble in
nitric acid. The precipitate and solution contain mercurosum as evidenced by
HCl. Mercuric nitrate with soluble sulphites forms a voluminous white pre-
cipitate, soluble in HNO, and containing mercurosum. Mercuric chlonde is
not precipitated by sulphurous acid or sulphites in the cold, but is reduced, by
boiling with sulphurous acid, to HgCl and then to Hg° .
42 MEiiC(in\ §68, G/.
Sulphuric acid and soluble Bulpliates precipitate from mercurous solu-
tions not too dilute, mercurous sulphate, Hg2S04 , white, decomposed by
boiling water, sparingly soluble in cold water (5c), soluble in nitric acid
and blackened by alkalis. Mercuric salts are not precipitated by sulphuric
acid or sulphates. For action of H2SO4 on UgClj see next paragraph and
(§269, 8, footnote).
f. — Hydrochloric acid and soluble chlorides precipitate from solutions of
mercurous salts, mercurous chloride, HgCl , " Calomel," white, insoluble iu
water, slowly soluble in hot concentrated HCl . Boiling nitric acid slowly
dissolves it, forming Hg(N03). and HgCL ; dissolved by chlorine or nitro-
hydrochloric acid to HgClj ; soluble in Hg(N03)2 (^^ footnote) (Dreschsel,
J. C, 1882, 42, 18). This precipitation of mercurous salts by hydro-
chloric acid is a sharp separation from mercuric salts and places mer-
curous mercury in the First (Silver) Group of Metals. Mercuric salts
are not precipitated by hydrochloric acid or soluble chlorides, unless the
mercuric solution is more concentrated than possible for a mercuric
chloride solution under the same conditions, i. e., a strong solution of
Hg(N03)2 gives a precipitate of Hg^L, on addition of HCl , soluble on
addition of water. Mercuric chloride is not decomposed by sulphuric
acid. A compound HgCL.HoSO^ is formed which sublimes undecom-
posed. The same compound is formed when HgSO^ is treated with HCl
and distilled (Ditte, .4. Ch., 1879, (5), 17, 120).
Hydrobromic acid and soluble bromides precipitate, from solutions of
mercurous salts, mercurous bromide, HgBr, yellowish white, insoluble in
water, alcohol, and dilute nitric acid; from concentrated solutions of
mercuric salts, mercuric bromide, HgBio , white, decomposed by concen-
trated nitric acid. Mercuric l)romido is soluble in excess of mercuric salts
(5& footnote), or in excess of the precipitant; hence, unless added in
suitable proportions, no precipitate will be produced. Sulphuric acid does
not transpose HgBr^ but forms compounds exactly analogous to those
with HgClo . Excess of concentrated H0SO4 gives some Br with HgiBr. .
Hydriodic acid and soluble iodides precipitate from solutions of mer-
curous salts, mercurous iodide, Hgl, greenish yellow — "the green iodide
of mercury '^ — nearly insoluble in water, insoluble in alcohol (distinction
from mercuric iodide), soluble in mercurous and mercuric nitrates; decom-
posed by soluble iodides with formation of Hg and Hgig , the latter being
dissolved as a double salt with the soluble iodide: 2HgI -f- 2KI = Hg -j-
HgIo.2KI . Mercurous chloride is transposed by HI or EI to form Hgl ,
excess of the reagent reacts according to the above equation (D., 2, 2, 867).
Ammonium hydroxide in the cold decomposes Hgl into Hg and Hgl,
(Francois, J. Pharm., 1897, (6), 5, 388).
Mercuric salts are precipitated as mercuric iodide, Hgis , first reddish-
§58, 7. MERCURY. 43
yellow then red, soluble in 24,814 parts of water at 17.5° (Bourgoin, A, Ch.,
1884, (6), 3, 429), soluble in concentrated nitric and hydrochloric acids;
quickly soluble in solutions of the iodides of all the more positive metals,
i. e. in excess of its precipitant, by formation of soluble double iodides; a?^
(Sl)^gl.j^ variable to KIHgl^ . A hot concentrated solution of potas-
sium iodide dissolves 3HgI^ for every 2KI. The first crystals from this
solution are Klftglo . These are decomposed by pure water, and require
a little alkali iodide for perfect solution, but they are soluble in alcohol
and ether. A solution of dipotassium mercuric tetraiodide, EoHgl^ =
(KI)2Hgl2 (sometimes designated the iodo-hydrargyrate of potassium), is
precipitated by ammonium hydroxide as mercurammonium iodide, 'SKgJ.
(Nessler's test), and by the alkaloids (Mayer's reagent).
Potassiiixn broinate precipitates, from solutions of mercurous nitrate, mer-
curous bromate, HgBrOs , white, soluble in excess of mercurous nitrate and
in nitric acid; from solutions of mercuric nitrate, mercuric bromate, HgCBrO,)..
soluble in nitric acid, hydrochloric acid, and in excess of mercuric nitrate,
soluble in 650 parts of cold and 64 parts of hot water (Rammelsberg, Pogg.^ 1842,
55, 79). No precipitate is formed when potassium bromate is added to mercuric
chloride (56, footnote). Iodic acid and soluble iodates precipitate solutions
of mercurous salts as mercurous iodafr^ HglO, , white with yellowish tint, solu-
ble with difficulty in dilute nitric acid, readil^^ soluble in HCl by oxidation to
mercuric salt. Mercuric nitrate is precipitated as mercuric iodate, HgCIO,), ,
white, soluble in HCl , insoluble in HNO, and H2SO4 , soluble in NH^Cl , trans-
posed and then dissolved by KI . Mercuric chloride is not precipitated bv
DO, (5b, footnote) (Cameron, C. .V., 1876, 33, 253).
g, — Arsenous acid or arsenites form a white precipitate with mercurous
nitrate, soluble in HNO, (Simon, Pogg., 1837, 40, 442). Merciiric nitrate is
precipitated by a solution of arsenous acid: the precipitate is soluble in HNO,
(/)., 2, 2, 920). Arsenic acid or Na.HAsO^ precipitates from mercurous nitrate
SHgsAsO^.HgrNOs.H.O , light yellow if the HgNO, be in excess (D., 2, 2, 921):
dark red HgjAsO^ if the arsenate be in excess. HgaAsOf is changed by cold
HCl to HgCl and H.AsO, , by boiling with HCl to Hgo , HgCl, , and HgAsO^ :
and is soluble unchanged in cold HNO, , insoluble in water and acetic acid
(Simon, Pogg., 1837, 41, 424). Arsenic acid and soluble arsenates precipitate
from mercuric nitrate, Hg3(As04)2 , white, soluble in HNO, and HCl , slightly
soluble in water. Arsenic acid and potassium arsenate do not precipitate
mercuric chloride from its solutions.
Stannous chloride precipitates solutions of mercuric salts (by reduction),
as mercurous chloride, white; or if the stannous chloride be in excess,
as metallic mercury (a valuable final test for mercuric salts) (10).
h, — ^Soluble cliroinates precipitate from mercurous solutions mcrcurouH
chromaie, Hg^CrO^ , brick-red, insoluble in water, readily transposed by HCl to
HgCl and H^CrO^ , soluble with difficulty in HNO, without oxidation (Richter,
B., 1882, 15, 1489). Mercuric nitrate is precipitated by soluble chromates as a
light yellow precipitate, rapidly turninp" dark brown, easily soluble in dilute
acids and in "HgCl... Mercuric chloride forms a precipitate with normal chro-
mates, but not with K.CtzOj .
7. Ignition. — Mercury from all its compounds is volatilized by heat as
the undecomposed salt or as the free metal. Mercurous chloride (Debray,
44 MERCURY, §58, 8-
J. C, 1877, 31, 47) and bromide and mercuric chloride and iodide sublime
(in glass tubes) uudecomposed — the sublimate condensing (in the cold part
of the tube) without change. Most other compounds of mercury are
decomposed by vaporization, and give a sublimate of metallic mercury
(mixed with sulphur, if from the sulphide, etc.). All compounds of mer-
cury, dry and intimately mixed with dry sodium carbonate, and heated in
a glass tube closed at one end, give a sublimate of metallic mercury as a
gray mirror coat on the inner surface of the cold part of the tube. Under
the magnifier, the coating is seen to consist of globules, and by gently
rubbing with a glass rod or a wire, globules visible to the unaided eye are
obtained.
8. Detection. — Mercury in the mercurou^ condition belongs to the first
GROUP (silver group), and is completely precipitated by HCl . It is iden-
tified by the action of ammonium hydroxide, changing the white precipi-
tate of mercurous chloride to the black precipitate of metallic mercury
and nitrogen dihydrogen mercuric chloride (a delicate and characteristic
test for Hg'). Mercury in the mercuric condition belongs to the second
GROUP (tin and copper group), and is separated from all other metals of
that group by the non-solubility of the sulphide in (NH^)2S, and in dilute
HNO3 . The sulphide is dissolved in nitrohydrochloric acid, and the pres-
ence of mercury confirmed by the precipitation of Hg° on a copper vrire, or
by the reduction to HgCl or Hg° by SnCL .
9. Estimation. — (a) As metallic mercury. The mercury is reduced by means
of CaO in a combustion-tube at a red heat in a current of COj . The sublimed
mercury is condensed in a flask of water, and, after decanting^ the water, dried
in a bell-jar over sulphuric acid without application of heat. The mercury may
also be reduced from its solution bj^ SnCl. (or H,POs at 100**) and dried as
above, (h) As mercurous chloride. It is first reduced to Hg' by H,PO, (Uslar,
Z., 1805, 34, 391), which must not be heated above 60°, otherwise metallic mer-
cury will be formed; and after precipitation by HCl and drying on a weighed
filter at 100°, it is weighed as HgCl . Or enough HCl is added to combine with
the mercury, then the Hg" is reduced to Hg' by FeSO^ in presence of NaOH:
2HgO H- 2FeO + 3H3O = HgaO -f 2re(0H),. H.SO, is added, which causes the
formation of HgCl , which is dried on a weighed filter at 100®. (c) As HgS .
It is precipitated by HjS, and weighed in same manner as the chloride. Any
free sulphur mixed with the precipitate should be removed by CS, . (d) As
HgO , by heating the nitrate in a bulb-tube in a current of dry air not hot
enough to decompose the HgO. (e) Volumetrically, by NaaSjO.; from the
nitrate the precipitate is ijellotc, from the chloride it is whHe:
3Hg(N0,), + 2Na,S,03 -h 2H,0 =: Hg.S,(NO.), + 2Na.,SO, -f 4HKO,
3HgCl, -f 2Na,S,03 -f 2H,0 = Hg.SXl, -f 2Na,S0, -f 4HC1 .
(f) Volumetrically, HgCL is reduced to HgjO by FeSO^ in presence of KOH ,
and after acidulating with H^SOf the excess of FeSO^ is determined by KsCrjOr
or KMnO^ (Jiiptner, C. r., 1882, 727). (g) By iodine. It is converted into HgCl
and then dissolved in a gnuluated solution of I dissolved in KI: 2HgCl -|- 6KI -h
I, = 2K2HgI, -h 2KC1 . The excess of iodine is determined by Na^S^O, . (h)
The measured solution of HgCL is added to a graduated solution of KI:
4KI -f HgCL = K,Hgl4 -f- 2KCi . The instant the amount of HgCl, shown
in the equation is exceeded a red precipitate of Hgl, appears, (i) Volumetric^
§69,2. SILVER. 45
by adding a few drops of ammonium hydroxide to HgCl, and then titrating
with standard KCN , the ammonium hydroxide precipitate disappears when the
mercury becomes Hg(CN), (Hannay, /. C, 1873, 26, 570; Tjison, J. C, 1877, 32,
679). (/) Electrolytically, by obtaining the mercury as HgNO, , Hg(NO,), ,.
or HgsSOA and precipitating as Hg° on platinum by the electric current.
Mercuric chloride cannot be used, as it is partly reduced to HgCl , and that
is not readily reduced to Hg® by the electric current (Hannay, I. c),
10. Oxidation. — Free mercury (Hg°) precipitates Ag, Au, and Pt from
their solutions, and reduces mercuric salts to mercurous salts (Hada, J. C,
1896, 69, 1667). Potassium permanganate in the cold oxidizes the metal
to Hg^O , when hot to HgO (Kirchmann, J. C, 1873, 26, 476). Mercur}^
and mercurous salts are oxidized to mercuric salts by Br , CI , I , HNO3 ,
H2SO4 (concentrated and hot), and HCIO3 .
Eeducing agents, as Fb , Sn , Sn", Bi , Cu^ Cu', Cd , Al , Fe , Co , Zn ,
ThS Mg, H3FO2, HaFOe and H2SO3, precipitate, from the solutions of
mercuric and mercurous nitrates, dark-gray Hg** ; from solution of mer-
curic chloride, or in presence of chlorides, first the white, HgCl , then gray
Hg®. Strong acidulation with nitric acid interferes with the reduction,.
and heating promotes it.
The reducing agent most frequently employed is stannous chloride:
2HgGl, + SnCla = 2HgCl + SnCl*
2HgCl + SnCla = 2Hg + SnCl*
or HgCl, + SnCl, = Hg + SnCl*
also 2Hg(N0,)a +'SnCl, = 2HgCl -f Sn(NO,)*
A clean strip of copper, placed in a slightly acid solution of a salt of mer-
cury, becomes coated with metallic mercury, and when gently rubbed
with cloth or paper presents the tin-white lustre of the metal, the coating
being driven off by heat; 2HgN03 + Cu = 2Hg + Cu(N03)2 . Formic acid
reduces mercuric to mercurous chloride, and in the cold does not affect
further reduction. Dry mercuric chloride, moistened with alcohol, is
reduced by metallic iron, a bright strip of which is corroded soon after
immersion into the powder tested (a delicate distinction from mercurous
chloride).
§59. Silver (Argentum) Ag = 107.92 . Monovalent.
1. Properties. — Specific gravity 10.512 heated in vacuo (Dumas, C. A'., 1878. 37,
82). Melting point, 960.7*» (Heycock and Neville, ./. C, 1895, 67, 1024). Does not
appreciably vaporize at 1567° (V. and C. Meyer, B., 1879, 12, 1428). It is the
whitest of metals, harder than pold and softer than copper. Silver is hardened
by copper; United States silver coin contains 90 per cent silver and 10 per cent
copper. In malleability and ductility it is inferior only to gold; and as a con-
ductor of heat and electricity it exceeds all other metals.
2. Occorrence. — Found in a free state in United States, Mexico, Peru, Siberia,
etc.; more frequently in combination. Its most important ores are argentite or
» Held, C. N,, 1866, 12, 242 ; « Heumann, /. C, 1875, 28, 182.
46 SILVER, §59, 3.
silver glance, AgjS , pyrargyrite, AgsSbSs , and horn silver, AgCl ; it is fre-
quently found in paying quantities in galena, PbS , and copper pyrites, and
in man3' other ores.
3. Frepaxation. — {a) It is alloj'ed with lead by fusion and the lead separated
by oxidation, (b) It is amalgamated with mercury and the mercury separated
by distillation, (c) It is brought into solution and the metal precipitated by
copper, (d) It is very easily reduced from the oxide or carbonate by heat
alone, and from all its compounds by ignition with hydrogen, carbon, carbon
monoxide and organic compounds.
4. Oxides. — Silver oxide, Ag.O , argentic oxide, is formed by the action of
alkali hydroxides on silver salts or by heating the carbonate to 200**. It is a
brown powder, a strong oxidizing agent, decomposed at .'iOO° into metallic silver
and oxygen. Concerning the existence of argentous oxide, Ag40 , and silver
peroxide, Ag.O-. , and their properties, see Muthmann {B., 1887, 20, 983); Pford-
ten {H., 1887, 80, 1458) and Bailey ((\ A'., 1887, 55, 263).
5. Solubilities. — a. — MeUiL — The fixed alkalis do not act upon silver, hence
silver crucibles are used instead of platinum for fusion with caustic alkalis.
Ammonium hydroxide dissolves finely divided silver, no action if air be excluded.
Acetic acid is' without action (Lea, Am. »S., 1892, 144, 444). Nitric acid is the
ordinary solvent for silver, most effective when about 50 per cent, the dilute
acid free from nitrous acid has little or no action (Lea, I.e.); silver nitrate is
formed and nitrogen peroxide is the chief product of the reduction of the
nitric acid (Higley and Davis, Am., 1897, 18, 587). Silver is not oxidized by
water or air at any temperature; it is attacked by phosphorus or by substances
ea.sily liberating phosphorus; it is tarnished in contact with hydrosulphuric
acid, soluble sulphides, and many organic compounds containing sulphur;
except that pure dry hydrosulphuric acid is without action upon pure dry silver
(Cabell, r. .v., 1884, 50, 208). Dilute sulphuric acid slowly dissolves finely
divided silver (Lea, /. c), a sulphate is formed and, with the" hot concentrated
acid, sulphur dioxide is evolved. Hydrochloric acid, sp. gr., 1.20, is without action
upon pure silver, but the metal is readily attacked by chlorine, bromine or
iodine. //.—O.Wr/r.— Silver oxide, Ag.O, soluble in 3000 parts of water, com-
bines with nearly all acids, except COj , forming \he corresponding salts. The
hydroxide is not known.
c. — Salts. — Silver forms a greater number of insoluble salts than any
other known metal, though in this respect mercury and lead are quite
similar. The nitrate is very soluble in water, 100 parts HoO dissolving
227.3 parts AgNO^ at 19.5°, soluble in glycerol, and sparingly soluble in
alcohol and ether. The chlorate dissolves in about ten parts cold vrater:
the acetate in 100 parts; the sulphate in about 200 parts cold water and
SS parts at 100°, and is more soluble in nitric or sulphuric acid than in
water; the borate, thiosulphate, and citrate are sparingly soluble in water.
The oxalate, tartrate, carl)onate, cyanide, ferrocyanide, ferricyanide, phos-
phate, sulphide, sulphite, chloride, bromide, iodide, iodate, arsenite, arse-
nate, and chromate are insoluble In water.
The chloride is soluble in 244 parts HCI , but its solubility is very much
lessened by the presence of mercurous chloride (lluyssen and Varenne, BL,
1881, 36, 5). If a solution of silver nitrate be dropped into concentrated
hydrochloric acid no precipitate appears until one half per cent of tht*^
HCI becomes AgCI (Pierre, J. C, 1872, 25, 123). Concentrated nitric acic'^
upon long continued boiling scarcely attacks AgCl (Thorpe, J. C, 1872, 2S ^
453); sulphuric acid, sp. gr. 1.84, completely transposes even the fuse*.1
§69, 6&. SILVER. 47
chloride on long boiling (Sauer, J. C, 1874, 27, 335). Silver chloride is
also soluble in ammonium hydroxide and carbonate; in sodium chloride
forming a double salt; in a concentrated solution of mercuric nitrate
(§68, 1; §68, 56 footnote); and in many other metallic chlorides and
alkali salts to a greater or less extent. All the salts of silver which are
insoluble in water are soluble in ammonium hydroxide, except the sulphide
and iodide; in ammonium carbonate, except the bromide, iodide, and
sulphide, the bromide very slightly soluble; in cold dilute nitric acid,
except the chloride, bromide, bromate, iodide, iodate, cyanide, and thio-
cyanate; in a solution of potassium cyanide (and by many other cyanides)
except the sulphide; and in alkali thiosulphates almost without exception.
6. Beactions. a, — The fixed alkali hydroxides precipitate from solu-
tions of silver salts (in absence of citrates), silver oxide, AgoO , grayish
brown, insoluble in excess of the reagents; soluble in acids, alkali cyanides,
and thiosulphates; somewhat soluble in ammonium salts. Most silver
salts are transposed on boiling with the fixed alkalis, except the iodide,
which is not thus transposed (Vogel, J. C, 1871, 24, 313).
Ammoninm hydroxide, in neutral solutions of silver salts, forms the
same precipitate, Ag.^0 , very easily dissolving in excess, by formation of
ammonium silver oxide, NH^AgO : AgNOg + 2NH^0H — NH^AgO -f
HE^HO, + HjO (Prescott, J. Am, Soc, 1880, 2, 32). Tn solutions con^
taining much free acid, all precipitation is prevented by the ammonium
salt formed.
Alkali carbonates precipitate silver carbonate, AgoCO, , white or yellow-
ish white, very slightly soluble in water and in the fixed alkali carbonates,
readily soluble in ammonium hydroxide and carbonate, transposed by
inorganic acids forming the corresponding salts. Carbon dioxide does
not transpose silver salts.
5.— Oxalic acid and soluble oxalates precipitate silver oxalate, Ag.C^O^ , white,
slightly soluble in water, soluble with difficulty in dilute nitric or sulphuric
acids, Veadily soluble in ammonium hydroxide. When heated it decomposes
with detonation, forming metallic silver.
Potassium cyanide precipitates from neiftral or slightly acid solutions
*i7rer cyanide, AgCN , white, quickly soluble in excess of the reagent as
silver potassium cyanide, AgCN.KCN . Hydrocyanic acid precipitates
''^lutions of silver salts but the precipitate does not dissolve m excess of
tlip reagent. Silver cyanide is transposed by H^SO^ or HCl and is soluble
'" <immonium hydroxide and carlmnate (Schneider, J. pr., 1868, 104, 83).
The ready solubility of nearly all silver compounds in potassium cyanide
^'^^) affords a means of separating silver from many minerals.
. ^otassium ferrocyanide precipitates silver ferrneyanide. Ag4Fe(CN)«, yellow-
'•"li white, soluble with difficulty in ammonium hydroxide and carbonate:
48 SILVER. §59, Gc.
metallic silver separates on boiling and a ferricyanide is formed. The ferro-
cyanide is not decomposed by hydrochloric acid, but it is changed to the
ferricyanide by nitric acid. Exposure to the air gives it a blue tinge. Potas-
sium ferricyanide precipitates tfUrer fenicjjanide, Ag3Fe(CN)c , reddish yellow,
readily soluble in ammonium hj'droxide and carbonate. Potassium thiocyanate
gives sUcer thiocyanate, AgCNS , white, soluble in ammonium hydroxide and
carbonate, insoluble in dilute acids. Concentrated sulphuric acid w^th the aid
of heat dissolves silver thiocyanate when some free silver nitrate is present. This
may be used as a separation from silver chloride, which is transposed b^* hot
concentrated sulphuric acid only on long-continued boiling (5r). To effect this
separation a little silver nitrate should be added to the silver precipitates and
then concentrated sulphuric acid and heat. To avoid danger of decomposition
of the chloride the mixture should not be heated above 200**. The pure silver
thiocyanate (silver nitrate being absent) is decomposed by hot concentrated
sulphuric acid with formation of a black precipitate containing silver.
c. — Silver nitrate is soluble in 500 parts of concentrated nitric acid (Schultz.
Z. Ch„ 1869, 531), and is precipitated from its concentrated water solutions by
the addition of concentrated nitric acid. d. — Disodium phosphate precipitates
silver phosphate, AggPO^ , yellow, soluble in dilute nitric acid, in phosphoric
acid, and in ammonium hydroxide and carbonate; but little soluble in dilute
acetic acid. Sodium pyrophosphate precipitates silver pyrophosphate, white, same
solubilities as the orthophosphate.
e. — Hydrosulphuric acid and soluble sulphides precipitate from neutral
acid or alkaline solutions silver sulphide, AgS , black, soluble in moderately
concentrated nitric acid (distinction from mercury), insoluble in potassium
cyanide (distinction from copper), insoluble in alkali sulphides (distinction
from arsenic, antimony, and tin). Certain insoluble sulphides fonn silver
sulphide from solutions of silver nitrate,* e. g., cupric sulphide gives silver
sulphide, cuprous sulphide gives silver sulphide and metallic silver, in
both cases cupric nitrate resulting (Schneider, J. C, 18T5, 28, 133 and
G12).
Thiosulphates precipitate silver thiosulphate, Ag.S^Og , white, unstable,
readily soluble in excess of the precipitant, by formation of double thiosul-
phates: with excess of sodium thiosulphate Na4Ag2(S20,), is formed (Cohen,
J. f'., 189(), 70, ii, 167). Silver thiosulphate turns black on standing or heating:
Ag'^S.O^ H- HjO = AgjS + H.SO^ . Sulphurous acid and soluble sulphites
precipitate silver sulphite^ AgsSOg , white, readily soluble in excess of alkali
sulphite or in dilute nitric acid: on boiling precipitated as metallic silver with
formation of sulphuric acid. Sulphuric acid and soluble sulphates precipitate
silver sulphate, Ag^SO^ , white, from concentrated solutions of the nitrate or
chlorate: sparingly soluble in water, quite soluble in concentrated sulphuric
acid.
f. — Hydrochloric acid and soluble chlorides precipitate silver chloride,
AgCl , white, curdy; separated on shaking the solution; turning ^aolet to
bro\vn on exposure to the light; fusible without decomposition; verj'
easily soluble in ammonium hydroxide as ammonio silver chloride,
(NH3)3(Ag:Cl)2 (Jarry, C. r., 1897, 124, 288). If mercurous chloride be
present with silver chloride the solubility in ammonium hydroxide is
♦ AgaS Is one of the least soluble of the sulphides. See $57, 6e, footnote.
§59, 7. SILVER. 49
greatly lessened, in fact a great excess of mercurons chloride may entirely
prevent the solution of silver chloride in ammonium hydroxide by forming
metallic silver. Silver chloride is quite soluble in a solution of mercuric
nitrate, which, if present in large excess, may entirely prevent the pre-
cipitation of the silver chloride by hydrochloric acid. The precipitation
by hydrochloric acid (in absence of a great excess of Hg(N0a)2) is the most
delicate of the ordinary tests for silver, being recognized in 250,000 parts
of water. As mercuric salts are not at all precipitated by HCl and lead
salts only imperfectly, silver is the only metal which belongs exclusively
to the FIRST OR SILVER GROUP OF BASES (§16).
Hydrobroxnic acid and soluble bromides precipitate sillier bromide^ AgBr «
"white, with a slight yellowish tint; but slightly soluble in excess of alkali
bromides, and much less easily soluble in ammonium hydroxide than silver
chloride. If silver nitrate be added to a bromide containing an excess of am-
monium hydroxide, the precipitate which first forms readily dissolves on shak-
ing; no solution is obtained with the iodide.
Hydriodic acid and soluble iodides precipitate silver iodide, Agl , pale yellow,
soluble in excess of the concentrated reagents by formation of double iodides,
as KTAgI , which are decomposed by dilution with much water. The precipi-
tate dissolves in 26,000 parts of ten per cent ammonium hydroxide: not at all in
a five i)er cent solution (Longi, Qazzctta, 1883, 13, 87). It is insoluble in dilute
acids, but is decomposed by hot concentrated nitric or sulphuric acids.
Silver bromate formed by adding potassium bromate to silver nitrate is soluble
in about 600 parts water and in 320.4 parts nitric acid (sp. gr.^ 1.21) at 25°, and
readily soluble in ammonium hydroxide. Silver iodaie formed in manner simi-"
lar to the bromate is soluble in about 28,000 parts water and in 1044.3 parts
nitric acid (»p. gr,, 1.21) at 25**, and readily soluble in ammonium hydroxide
(Longi, I. c).
ff. — Soluble arsenites precipitate silver arsenite, Ag,AsO, , yellow, very readily
soluble in dilute acids and in ammonium hydroxide. Soluble arsenates precipi-
tate ifilver arsenate, AgaAs04 , red-brown, soluble in ammonium hydroxide,
nitric acid, arsenic acid, and almost insoluble in acetic acid.
A solution of alkali stannite — as EsSnOg — precipitates metallic silver
from solutions of silver salts. A solution of silver nitrate in a great
excess of ammonium hydroxide constitutes a very delicate reagent to
detect the presence of tin in the stannous condition in the presence of fixed
alkalis; antimony does not interfere if a great excess of ammonium hy-
droxide be present.
A. — Cbromates and dichromates, as K3Cr04 and K^CrsOr , precipitate silver
chromale, AgjCrO* , dull-red, sparingly soluble in water and in dilute nitric
acid, soluble in ammonium hydroxide.
7 Xgniitioii. — ^Silver nitrate melts undecomposed at 218**, at a red heat it is
decomposed into Ag** , O, N, and NO (Fischer, Po(jg., 1848, 74, 120). Silver
chloride fuses at 451°, the bromide at 427°, and the iodide at 527°. On charcoal
with sodium carbonate, silver is reduced from all its compounds by the blow-
pipe, attested by a bright malleable globule. Lead and zinc, and elements more
volatile, may be separated from silver by their gradual volatilization under
the blow-pipe, or in the assay furnace (see Cupellation in works on the assay
of the precious metals).
50 SILVER. §59, S.
8. Detection. — Silver is identilied by ity precipitation with hydrochloric
acid, the insolubility of the precipitate in hot water, and its solubility in
aninionium hydroxide, with repreeipitation on rendering add with nitric
acid (§61).
0. Estimation. — (a) As metallic silver, into which i , " — •
ip-nition if it is the oxide or carbonate, or by ignit
chloride, bromide, iodide or sulphide (Vogel, •/. (\, l^
precipitated as AgGl , and after i|>:inting- to incipient Ti
converted into Ag.S by HjS , and uci|rh(»d yfter dryi
in case of an acid that niif^ht liberate free sulphur,
solution of KAg'(CN). is formed, precipitate with HK
1(»0°, weipfh as AgCN . {€) Volumetrically, by adding
NaCl until a precipitate is no Unifier formed. This n
the measured silver solution to the frraduated NaCl 80
drops of KnCrOf , until the red ])recipitate bejrins to fo
add a ^'raduated solution of ammonium thiocyanate, co
until thr red color ceases to disa])]>ear. (y) Add the i
to a standard solution of KCN until a permanent white
10. Oxidation.— Metallie silver precipitates gol
their solutions, reduces cuprie chloride to (jupro
chloride to mercurous chloride, and ])erinanganateB
Silver is precipitated from its solutions by: Pb,
Sb , SbH, , Sn , Sn", Bi , Cu , Cu'\ Cd , Te , 7e , Pc
.PS PH, , H,PO, , H,S03 , SiH,\ H,0./', and H (very
Tn alkaline mixture silver is also reduced by Hg
Mn". An anial<]fam of mercury and tin reduces h
silver in the wet way, tbe silver amal^^amates witl
tin becomes Sn^^ (Laur, (\ r., lSS->, 95, 3S).
Ferrous sulphate in the cold incompletely reduces sll*
ferric salt formed is reduced and the silver dissolved (L-
reduction of silver by certain orj?anic reaji^ents, the meti
silver coatins^ or mirror upon the inner surface of the
vessel. Usually a sli^-htly ammoniaeal solution of sil
allowed to stand some time with the reapent: such as
of cloves or cassia, formic acid, aldehyde, chloral, ta _, ,.,.v. vjeiine
warminp facilitates the result. If a jrood mirror is desired, great care must be
taken to free the inner surface of the jrlass from all organic impurities by
careful washing with ether, chloroform, etc. In these deoxidations, generally
the nitric acid radical of the silver nitrate is not decomposed, but nitric acid iV
left: 4AgN0, -f -*H,0 = 4Ag + 4HN0, + O, .
Light acts upon nearly all salts <>f silver when mixed with gelatine or other
organic substances used in preiiarinjir ])hotographic plates, etc. It is quite
probable that the silver is reduced to metallic silver or argentous oxide, Ag^O ,
or both: but the action is not well understood. The nitrate in crystal or pure
water solution, the phos])hate, bromide, iodide and cyanide are not decomposed
by liirht alone: but light greatly hastens their decomposition by organic sub-
stanees, or other reducing agents, as of solution of silver nitrate in rain water,
or written as an ink upon fabrics. Silver is the base of most indelible inks.
' I^a, Am. S., 1892, IJi, 4«. 2 A, 2, 2, 759. ' Skey, C. X.. 1871, 2S, 232. * Senderons, C. r., IW,
lOJ, 17r). ft D.. 2, 1, 4r)f.. « Uiejfler, J. C, 1896, 70, li, 471. ^ PoUet, B., 1874, 7, 656 ; Schwaiwnbach
and Kritschewsky. Z., 1886. 25, 374 ; Cooke, C. A'., 1888. 5S, 103. • Mlllon. Am. S^ 1863, 86. 417.
§60. COifPARISOy OP REACTI0X8 OF MKTALS OF THE aiLVEH OROLJ'. 51
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TABLE FOR ANALYSIS OF THE SILVER OR FIRST GROUP.
§61.
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§63, 6e. DIHECTIONS FOR ANALYSIS WITH NOTES. 53
Directions for the Analysis of the Metals of the First Group.
§62. Hanipolation. — To the solution add hydrochloric acid (whenever
directions call for the addition of a reagent it is to be used reagent
strength unless otherwise stated) drop by drop (§32) until no further
precipitate is formed and the solution is distinctly acid to litmus (§36).
The precipitate will consist of the chlorides of Pb , Hg', and Ag , e, g., ^
Pb(H03)2 + 2HC1 = PbCl2 -f 2HNO3 . Shake thoroughly and allow to
stand a few moments before filtering; if the solution is warm it should
be cooled to the temperature of the room. Decant the solution and
precipitate upon a filter paper previously wetted. (§35) with water and
wash two or three times with cold water or until the filtrate is not strongly
acid to litmus. The washings with cold water should be added to the
first filtrate and the whole marked and set aside to be tested for the
metals of the remaining groups (§16).
§63. Xotes, — 1. Failure to obtain a precipitate upon the addition of HCl to
an acid reaction is proof of the absence of Hg' and Ag , but a solution of a
lead salt may be present* of such a degrree of dilution that the lead chloride
formed will be soluble in the dilute acid (§57, 5c).
2. The solution should not be strong-ly acid with nitric acid, as it forms
nitrohydrochloric acid with the hydrochloric acid, causing oxidation of the
Kg' (§58, 5c). Lead chloride is also more soluble in nitric acid than in dilute
hydrochloric acid (§57, 5c). By a study of the solubilities of the silver group
metals it will be seen that HaSO^ , HCl , HBr or HI cannot be used in prepar-
ing" a solution for analysis when these metals are present.
3. A great excess of acid is to be avoided, as it may interfere with the reac-
tion in Group II. (§57, 6e). Complete precipitation should be assured by
testing the filtrate with a drop of HCl , when no further precipitation should
occur (§32). If a white precipitate is .formed by adding a drop of HCl to
the filtrate it is evident that the precipitation was not complete and more
HCl should be added and the group separation repeated.
4. The presence of a slight exce3s of dilute acid does not aid or hinder the
precipitation of the Hg' or Ag, but as PbClj is less soluble in dilute HCl
than in water, a moderate excess of the acid causes a more complete precipita-
tion of that metal in the first group.
5. Concentrated HCl dissolves the chlorides of the first group quite appre-
ciably (§59, 5c).
6. Hydrochloric acid added to certain solutions may cause a precipitate
when none of the first group metals are present. Some of the more important
conditions are mentioned:
a. A concentrated solution of BaCl, is precipitated without change by the
addition of HCl, readily soluble in water (§186, 5c).
b. An acid solution of Sb , Bi , or Sn , with some other acid than HCl ,
and saturated with water as far as possible without precipitation, on the
addition of HCl, precipitates the oxychloride of the corresponding metal
(§76, (if). These precipitates are readily soluble in an excess of the HCl . It
must, however, be remembered that a trace of AgCl will also be dissolved by
an excess of HCl (§59, 5c).
c Solutions of metallic oxides in the alkali hydroxides are precipitated when
neutralized with acids, e. {;., K^ZnOj -f 2HC1 ^ Zn(0H)2 -f- 2KC1 .
d. The sulphides of As , Sb , Sn , Au , Pt , Mo (Ir , W , Ge , V , Se and Te)
in solution with the alkali polysulphides are reprecipitated together with
sulphur on the addition of HCl (§69, 6c).
e. Soluble polysulphides and thiosulphates give a precipitate of sulohur,
white, with HCl <§256, 3a).
54 DIRECTIONS FOR AXALYSIS WITH NOTES. §63, 6f.
f. Certain soluble double cyanides, as Ni(CN)2.2KCN , are precipitated
as insoluble cyanides, Ni(CN)j , on the addition of HCl (§133, 66).
g. Solutions of silicates (§249, 4), borates, tungstates, molybdates; also
benzoates, salicylates, urates, and certain other organic salts, are precipitated
by acidulation with HCl, many of the precipitates being soluble on further
addition of the acid.
h. Acidulation with HCl may induce changes of oxidation or reduction,
%vhich in certain mixtures may result in precipitation: for example, Cu" salts
with KGNS in ammoniacal solution (§77, 66); mixture of solutions of KI and
KIO, (§280, 6, fi, 7), etc.
7. If the precipitate, obtained by the addition of HCl to the solution, is
colored or does not give further reactions wliich are conclusive and perfectly
satisfactory in every respect, it should be separated* by filtration, and treated
as a solid substance taken for examination (see conversion of solids into
liquids, §301).
8. Compounds of the first group metals insoluble in water or acids are trans-
posed to sulphides by digestion with an alkali sulphide. The lead and silver
sulphides thus formed are readil3' soluble in hot dilute nitric acid. The mer-
curous compounds are changed to mercuric sulphide (§58, or; and 6r), a second
group mercury compound insoluble in HNO, .
9. If but one metal of the first group be present, the action of NH4OH
determines which it is; FbCl, does not change color or dissolve; HgCl blackens;
and AgCl dissolves (§60).
§64. Manipulation. — The precipitate (white) on the filter should now
be washed once or twice with hot water. The first hot water should be
poured upon the precipitate a second time. This hot filtrate is divided
into four portions and each portion tested separately for lead with the
following reagents, KSO^ , HoS , KXt^O^ , and KI (§57, 6 e, h, and f) :
PbCl, H- H3SO, = PbSO, (white) + 2HC1
PbCl, + H^S = PbS (black) + 2HC1
2PbCl, H- K,Cr,0, + H,0 = 2PbCr0, (yellow) + 2KC1 + 2HC1
PbCl, H- 2KI = Pbl^ (yellow) + 2Ka
The yellow precipitate with potassium iodide (the KI must not be used
in great excess (§57, 5c)) should be allowed to settle, the liquid decanted,
and the precipitate redissolved in hot water, to a colorless solution which
upon cooling deposits beautiful yellow crystalline scales of Pbia (charac-
teristic of lead).
§65. Notes. — 1. Lead is never completely precipitated in the first group
(§57, 6f). The presence of a moderate excess of dilute HCl and the cooling of
the solution both favor the precipitation.
2. Lead can be completely separated from the second group metals by sul-
phuric acid applied to the original solution (§57, 6^, §95 and §98), but that
would necessitate a regrouping of the metals; as, Ba , Sr , and Ca would also
be precipitated (Zettnow, Z., 1867, 6, 438).
3. In order to precipitate the lead chloride, not removed in the first group, in
the second group with H^S , the solutions must not be strongly acid, either
the excess of HCl should be removed by evaporation or the solution should be
diluted (§57, Gc, and §81, 3, 5 and 9).
4. If the lead chloride is not all washed out with hot water it is changed to
an insoluble basic salt (white) by the NH4OH , part remaining on the filter
and part carried through mechanically which causes turbidity to the am-
monium hydroxide solution of the AgCl and makes necessary the filtration
of that solution before the addition of HNO, , otherwise it does not interfere.
5. The precipitation of lead as the sulphide while not characteristic of lead.
5568, 3. DlRECTlOyS FOR ANALYSIS WITH NOTES. 55
is exceedingly delicate, much more so than the formation of the white PbSOj
(§57, 5c). In extremely dilute solutions no precipitate occurs, merely a brown
coloration to the solution. The presence of free acid lessens the delicacy of
the test.
fi. PbCrO* is blackened by alkali sulphides and dissolved by the fixed alkalis
(important distinction from BaCrOt); the solubility in the tixed alkalis is also
^n important distinction from bismuth chromate (§76, 6/i).
7. Other tests for lead by reduction on charcoal before the blow-pipe, or in
the wet way by Zn, should not be omitted (§57, 7 and 10). If to a solution of
lead salt nearly neutral a strip of zinc be added, the lead will soon be deposited
on the zinc as a spongy mass.
§68. Uanipulation. — The white precipitate remaining on the filter after
washing with hot water consists of HgCl and AgCl, with usually some
PbCls which was not removed. To this precipitate NH^OH , one or two cc.
is added and allowed to pass through the filter into a clean test-tube.
An instantaneous blackening of the precipitate is conclusive evidence of
the presence of mercurosum; 2HgCl + 2NH4OH = Hg + NH.^HgCl -f
ira.Cl + 2JLJQ .
The AgCl is dissolved by the NH.OH : 2AgCl + 3NH^0H = 3NH3 .
2AgCl + 3H2O , and is found in the filtrate; its presence being confirmed
by its reprecipitation on rendering the solution acid with HNO3 : 3NH3 .
•2AgCl + 3HHO3 = 2AgCl + 3NH,N0., .
§67. A'ofc«.---Merciiry.— 1. The black precipitate on the filter, caused by the
addition of NH4OH to the HgCl may be examined under the microscope for
the detection of globules of Hg°, or the precipitate may be digested with
concentrated solution of (1)^4)5804, which dissolves the NH2HgCl , leaving
the Hg*" unattacked (§58, 6fi).
2. If the original solution contains no interfering metals, the distinctive
reactions of mercurous salts with iodides, chromates and phosphates should be
obtained (§58, 6e, h and d),
3. The precipitation with HCl and blackening with NH4OH is conclusive evi-
"dence of the presence of mercury in the mercurous condition; should further
•confirmation be desired, the black precipitate may be dissolved in nitro-
hydrochloric acid, the excess of acid removed by evaporation and the free
metal obtained as a coating on a copper wire, by immersing the freshly
polished wire in the solution of HgCL (§58, 10).
4. Mercury has but few soluble mercurous compounds, and in preparing
solutions of the insoluble compounds for anal^-sis, oxidizing agents are usually
employed and the mercury is then found entirely in the second group as a
fiulphide (§96 and §97).
5. Additional proof may be obtained by mixing a portion of the black residue
with sodium carbonate, drying and heating in a glass tube (read §58, 7, also
197, 7).
§68. Silver. — 1. The presence of a large excess of Hg(N0s)2 prevents the
precipitation of AgCl from solutions of silver salts by HCI (§59, he). In this
case the metals should be precipitated by H.,.S and the well-washed precipitate
digested with hot dilute HNO^ . The silver is dissolved as AgNOj . while the
mercury is unattacked: 6Ag,S + IfiHNO, = 12AgN08 -f 38, + 4N0 4- 8H0O .
After evaporation of the excess of HNO, the solution may be treated with
HCl as an original solution.
2. A small amount of AgCl with a large amount of HgCl is not dissolved by
HH 4OH , but is reduced to Ag** by the Hg** formed by the addition of the
HH4OH to the HgCl (§58, 6<i, §59. 10 and §60).
3. If Hg' be present and Ag is not detected, the black precipitate on the
56 ARSEyiC. §88, 4.
filter should be digested for some time with (NHJaS , washed, and boiled with
hot dilute nitric acid. The Ag , if any be present, is dissolved and separated
from the HgS :
NH.HgCJl + (NHJ.S H- 2H2O = HgS + NH,C1 -f 2NH,0H
Hg + (NHJ,S, = HgS + (NHJaS..,
4. If only a trace of silver be present, its detection by adding* HNO, to the
NH4OH solution of the chloride may fail, unless the excess of the NH4OH be
first removed by evaporation (because of the solubility of the Ag^l in the
ammonium salt, §59, 5c),
5. As a further test for silver, the chloride, precipitated by the nitric acid,
may be reduced to the metal by zinc; by adding to the ammoniacal solution
a few drops of potassium stannite (§71, 6a and 8); by warming with grape
sugar in alkaline mixture. In all cases the well-washed grayish black metal
may be dissolved in nitric acid as AgNO, .
6. To identify the acid of silver salts which are insoluble in HNO,(AgCl,
AgBr , Agl), (i) Add metallic zinc and a drop of H3SO4 ; when the silver is all
reduced test for the acid in the filtrate. (2) Fuse with NajCO, , add water,
and test the filtrate for acids. {S) Add HjS , or an alkali sulphide, digest
warm for a few minutes, filter and test filtrate for acids. (4) Boil with KOH
or NaOH (free from HCl), and test the filtrate in the same manner. It must
not be overlooked that by the first three methods, and not by the last,
bromates and iodates are reduced to bromides and iodides (§257, 6B).
The Tin and Copper Group (Second Group).
Arsenic, Antimony, Tin, Gold, Platinum, Molybdenum, Mercury, Lead,
Bismuth, Copper, Cadmium (Ruthenium, Rhodium, Palladium, Iridium,
Osmium, Tungsten, Vanadium, Germanium, Tellurium, Selenium).
The Tin Group (Second Group, Division A).
Arsenic, Antimony, Tin, Gold, Platinum, Molybdenum (Iridium, Tungs-
ten, Vanadium, Germanium, Selenium, Tellurium).
§69. Arsenic. As = 75.0. Valence three and five.
1. Properties. — SpeeifU* gravity, pure crystalline 5.727 at 14°; amorphous 4.716
(Bettendorlf, A., 1867, 144, 110). Melting pmnt, at dull red heat, under pressure
in sealed tube (Landolt, J., 1859, 182); between the melting point of antimony
and silver (Mallet, C. y„ 1872, 26, 97). VoUitilizcs in an atmosphere of coal gas
without melting at 450** (Conechy, C. N., 1S80, 41, 189). Vapor density (H = 1).
147.2 (Deville and Froost, C. r., 1863, 56, 891); therefore the molecule is assumed
to contain four atoms (AS4). At a white heat the vapor density is less, but
the dissociation is not low enough to indicate As, (Mensching and V. Meyer.
B,, 1887, 20, 1833). Arsenic exists in two forms, crystalline and amorphous.
The crystalline arsenic is steel-gray with a metallic luster, brittle and easily
pulverizable; forms beautiful rhombic crystals on sublimation with slow-
condensation. For ductility, malleability, etc., see D., 2, 1, 161. Amorphous
arsenic is grayish black, of less specific gravity than the crystalline; long
heating changes it to the crystalline form (Engel, C. r., 1883, 96, 1314). The
vapor of arsenic is citron-yellow (Le Roux, 0. r., 1860, 51, 171), with an oppres-
sive and poisonous alliaceous odor. It is slowly oxidized in moist (not in dry)
air at ordinary temperature; when heated in the air, it bums with a bluish
flame and becomes the white arsenous anhydride, ASjO, . The burning metal
evolves a strong garlic odor, not noticed when the pure arsenous anhydride is
sublimed. In its physical properties arsenic is a metal, but its failure to act
as a base with oxyacids classes it chemically with the non-metallic elements
(Adie, ,/. (7., 1889, 55, 157; Stavenhagen, Z, angew., 1893, 283). Its chief use as a
metal is in mixing with lead for making shot.
§69» 5b. ARSENIC. 57
2. Occurrence. — Arsenic is very widely distribnted geographically. Found
native; as As,0,; as an alloy with other metals, e.g., FeAs, , NlAs , CoNlAsx;
as realgar, AStS, ; orpiment, AsjS, ; arsenical pyrites, FeAsFeS, ; as an arsenate
in cobalt bloom, Go,(As04)2 ; and in a great variety of minerals. Most sulphide
ores of zinc and iron contain arsenic, hence arsenic is frequently found in
these metals and in sulphuric acid made from the sulphur, and also in the
products made therefrom.
3. Preparation. — (i) Reduced from its oxide by ignition with carbon; 2A830»
-f 3C =: AJI4 + SCO, . (2) From arsenical .pyrites, FeAsFeS, , by simple igni-
tion, air being excluded; 4(FeA8.FoS,), = 8F0S + As* . (3) From orpiment,
ASjS, , by fusion with sodium carbonate and potassium cyanide; 2A82SS -|-
6Ka,CO, + 6KCK = As* + 6Na,S + 6KCN0 + 6C0, .
4. Oxides. — ^Arsenic forms two oxides: arsenous oxide or anhydride, AB2O,
(BUtz, Z. phys. Ch., 1896, 19, 385; C. C, 1896, 793), and arsenic oxide or anhydride,
ASjOft . Arsenous oxide, AszO, (ichite arsenic, arsenous anhydride, arsenous acid,
arsenic trioxide}, is usually prepared by burning arsenic; it may also be prepared
by heating arsenic in sulphuric acid till SO, is evolved, or by decomposing
AsCla with HzO . It sublimes easily on gradually heating, forming beautiful
octahedral and tetrahedral crystals. On suddenly heating under pressure it
melts, and on cooling forms the opaque arsenic glass. It is very poisonous,
usually producing violent vomiting. One hundred fifty milligrams are con-
sidered a fatal dose for an adult. No acids (hydroxides) of arsenous anhydride
(oxide) have been isolated; but its solutions with bases form salts, arsenites,
as if derived from the meta, ortho, and pyro arsenous acids. The alkali
arsenites are usually meta compounds; the arsenites of the alkaline earths and
heavy metals are usually ortho compounds (D., 2, 1, 170).
Arsenic pentoxide, AsaO^ (arsenic anhydride, arsenic oxide), is formed by heat-
ing arsenic acid, H^AsO* (Berzelius, A. Ch., 1819, 11, 225). It is a white
amorphous mass, melts at a dull red heat, is slowly deliquescent, combining
with water to form HgAsO* . The pentoxide, ASjOs , forms three acids or
hydroxides; meta -arsenic acid^ HAsO, =: AsOjCOH); ortho-arsenic acid,
HsAsO* =: AsO(OH),; and pyro-arsenic acid, H^AsjO, ^ As30,(OH)4; each
of these forming a distinct class of arsenates with bases. Ortho-arsenic acid is
formed by adding water to arsenic anhydride, AsjOs -f 3H3O = 2H3AsO« ,
or by oxidizing arsenic or arsenic anhydride with nitric acid. Pyro-arsenic
acid is formed by heating the ortho acid to between 140° and 180**: 2H3ASO4 =
H4AS2O7 -f H2O . The meta acid is formed by heating the ortho or pyro acid
to 206°: H,AsO« = HAsO, -f H2O (D., I. c).
5. Solubilities. — a. — Metal. — Arsenic is insoluble in pure water. It is readily
attackefd by dry chlorine and bromine upon contact and by iodine with the aid
of heat. Arsenous chloride, bromide and iodide are formed. It combines
with sulphur, forming from AS2S2 to AS2SS , depending upon the proportion of
sulphur present (Gelis, A. Ch., 1873, (4), 30, 114). Chlorine and bromine in
presence of water oxidize it, first to arsenous then to arsenic acid (Millon,
A. Ch., 1842, (3), 6, 101): As* + lOCl, + I6H2O = 4H,A804 -f 20HC1 . It is not
attacked by concentrated hydrochloric acid at ordinary temperature and but
slowly by the hot acid in presence of air forming AS3O, , then AsCl, ; nitric
acid readily oxidizes it first to AsjO, then to HsAsO* ; upon fusion with KNO,
it becomes KsAsO*; readily soluble as HjAsO* by nitrohydrochloric acid;
sulphuric acid, dilute and cold, is without action; with heat and the more con-
centrated acid ASxO, is formed and the sulphuric acid is reduced to SO2 .
Ammonium hydroxide is without action (Guenez, C. r., 1892, 114, 1186). Hot
solution of potassium or sodium hydroxide dissolves it as arsenite: As* 4-
4KOH -f 4HaO = 4KAsO, -f 6H2 .
6. — Oxides. — Arsnwus oxide exists in two forms, crystalline and amorphous, the
solubilities of which differ considerably (§27). At ordinary temperature 100
parts of water dissolve 3.7 parts of the amorphous and 1.7 parts of the crystal-
line, several hours being necessary to effect the solution. 100 parts of boiling
water dissolve 11.46 parts of the amorphous and 10.14 parts of the crystalline
oxide in three hours (Winkler, /. pr., 1885, (2), 31, 247). The presence' of acids
greatly increases the solubility in water (Sehultz-Sellac, B., 1871, 4, 109).
Arsenous oxide is readily soluble in alkali hydroxides or carl3onates to arsenites
68 ARSEXIC. §69, 5c.
(Clayton, C. N», 1891, 64, 27). Arsenic 'penioxide, ASsOg , is deliquescent, soluble
in water forming HsAsO« . The meta and pyro acids are easily soluble in
water forming the ortho acid (Kopp, A. Ch., 1856, (3), 48, 106).
c. — Salts. — Arsenic does not act as a base with oxyacids, but its oxides combine
with the metallic oxides to form two classes of salts, arsenites and arsenates.
Arsenites of the alkalis are soluble in water, all others are insoluble or only
partially so; all are easily soluble in acids. AUcali arsenates, and acid arsenates
of the alkaline earths, are soluble in water; all are soluble in mineral acids,
including H,A804 (LeFevre, C, r.,.1889, 108, 1058). See also under the respec-
tive metals.
Arsenous sulphidCf ABoS., , is insoluble in water when prepared in the
dry way ; when prepared in the moist way it may be transformed into the
soluble colloidal * form by treatment with pure water, from which solu-
tions it is precipitated by solutions of most inorganic salts (Schulze, J. pr.,
1882 (2), 25, 431). The presence of acids or solutions of salts prevents
the solubility of A82S., in water. Boiling water slowly decomposes the
sulphide forming ABoO., and HoS (Field, C. N., 1861, 3, 115; Wand, Arch.
Phar., 1873, 203, 296). It is completely decomposed by gaseous HCl form-
ing AbCI,^ (Piloty and Stock, B., 1897, 30, 1649)r, very slightly decomposed
by hot concentrated acid (Field, l. c). Chlorine water and nitric acid
decompose it readily with formation of H3A8O4; with sulphuric acid
ASoO., and SO^ are formed (Hose, Pogg., 1837, 42, 53G). The alkali hy-
droxides or carbonates dissolve it readily with formation of BAsOj and
KAsS. (K = K, Na and NHJ (/)., 2, 1, 183); soluble in alkali sulphides
and poly-sulphides forming E^ABjS- , and RAsSo (Berzelius, Pogg., 1826,
7, 137; Xilsson, J. C; 1872, 25, 599).
Arsetiic sulphide, Ab^S.^ , is insoluble in water; soluble in HCl gas, as
AbCL ; insoluble in dilute HCl , soluble in HNO3 or chlorine water, as
H3ASO4 ; soluble in alkali hydroxides and carbonates, as E^AsS^ and
E3ABO3S : A8,S, + 6NH,0H = (NHJaAsS, + (NH,)3A803S + 3HoO (Mc-
Cay, Ch. Z., 1891, 16, 476); soluble in alkali sulphides, as TL^AbS^ (Nilsson,
J.pr., 1876 (2), 14, 171).
Arsenous chloride, bromide and iodide (AsClj , AsBr, , Asl,) are decomposed
by small amounts of water into the corresponding oxyhalogen compounds,
AsOCl , etc. A further addition of water decomposes these compounds into
arsenous oxide and the halogen acids.
6. Keactions. — a. — The alkali hydroxides and carbonates unite with arsenous
and arsenic oxides (acids), the latter with evolution of carbon dioxide, forming
soluble alkali arsenites and drsenates. These alkali salts are chiefly meta arse-
nites and ortho arsenates (Bloxam, J. C, 1862, 15, 281; Graham, Pogg., 1834, 32.
47).
•Colloids l8 a name grivea by Graham to a class of glue-like bodies in distinction to the crystal-
loids, which have a well-deflncd solid form. The colloids are indefinitely soluble in water,
giving the little-understood ** pseudo-solutions,*' which stand midway between the mechanical
suspension or emuls'on and the true solution. Gelatine, starch, the metallic sulphides, silicic
acid, and the hydroxides of iron and aluminum are some of the substances that may take on the
colloid form. The colloid solutions arc as a rule broken up by addition of an acid or a neutral
salt.
§69, 6e. ARSKXIC. 59
b, — Oxalic acid does not reduce arsenic acid* (Nay lor and Braithwaite, Plmrm,
J. Trans., 1883, (3), 13, 464). Potassium ferricyanide in alkaline solution oxi-
dizes arsenous compounds to arsenic compounds, very rapid h' when g-ently
warmed, c. Nitric acid readily oxidizes all other compounds of arsenic to
arsenic acid. d. Hypophosphites in presence of concentrated hydrochloric acid
reduces all oxycompounds of arsenic to the metallic state. 0.00001 gram oi
arsenic may be detected by boiling with 10 cc. strong hydrochloric acid and 0.2
gram calcium hvpophosphite (Engel.and Bernard, C. r., 1896, 122, :J90; Thiele
and Loot C, T., "l890, 1, 877 and 1078; and Hager. J. C, 1874, 27, 868).
e. — ^Hydrosnlphnric acid preoipitates the lemon-yellow arsenmis sulphide,
A82S3 , from acidulated solutions of arsenous acid. The precipitate forms
in presence of concentrated hydrochloric acid. Citric acid and other
organic compounds hinder the formation of the precipitate, hut do not
wholly prevent it if strong hydrochloric acid be present. Nitric acid
should not he present in strong excess as it decomposes hydrosulphuric
acid, with precipitation of sulf)hur.
In aqueous solutions of arsenous acid the sulphide forms more as a
yellow color than as a precipitate, being soluble to quite an extent in pure
water, especially when boiled (5r) : ASoS, + SHjO = ASjO., + SHjS . This
has been given as a method of separating arsenous sulphide from all other
heavy metal sulphides (Clermont and Frommel, J. C, 1879, 38, 13). The
precipitate is not formed in solutions of the arsenites except upon acidu-
lation. Alkali sulphides produce and, by further addition, dissolve the
precipitate {5c):
As,0, + 3(NHJ,S -f 3HjO = As^S, -f 6NH4OH
A8,S, + 2(NH,),S = {KR^),As,S, or As^S, + (NH4)2S = 2NH4ASS,
Arsenous sulphide is also soluble in alkali hydroxides and carbonates,
forming arsenites and thioarsenites (5r). The thioarsenites are precipi-
tated by acids forming ASjS,, : (NHJ^ASjS, + 4HC1 = ABjSg + 2HoS +
4ira[^Cl or 2NH,A8S, + 2HC1 = As.Sg + H^S + 2NH^C1 .
The soluhility of the sulphides of arsenic in yellow ammonium sulphide
separates arsenic with antimony and tin from the other more common
metals of the second group; and the solubility in ammonium carbonate
effects an approximate separation from antimony and tin (Eager, J. C,
1885, 48, 838). Arsenous sulphide is soluble in solutions of alkali sul-
phites containing free sulphurous acid (separation from antimony and
tin): 4AS2S3 + 32KHS0, = 8E:A80, + 12K,S20., + 3So + USO, + IGH^O.
It may also be separated from antimony and tin by boiling with strong
hydrochloric acid, the ASjS., remaining practically insoluble; the sulphides
of antimony and tin being dissolved. It is easily dissolved by strong
♦ Patrouillard {Pharm. J. Trans., 1883, (3), 18, 882) claims the reduction of A»t to A»"' by oxalic
acid ; and Hagrer (C. C, 1882, 090) reports a microscopic tost for arsenic by reduction to metallic
arsenic on boiling with oxalic and sulphuric acids. Experiments in the authors' laboratory fail
to conftrm these results.
60 ARSEmc, §69, 6f.
nitric acid, and by free chlorine or nitrohydrochloric acid, as arsenic acid :
6A82S3 + 20HHO, + 8H2O = 12H3A80^ + QS^ + 20N0 ; gAs^Sg + lOCU
+ I6H2O = 4H,A804 + 3S2 + 20HC1 . Usually a portion of the sulphur
is oxidized to sulphuric acid, completely if the nitric acid or chlorine be in
great excess and heat be applied: As^Sg + I4CI2 + 2OH2O = 2H3ASO, +
3H2SO4 + 28HC1 .
Arsenic pentasulfhide, ASgSg , is formed by passing H^S for a long time
into a solution of alkali arsenate and then adding acid (McCay, Am,, 1891,
12, 547); by saturating a solution of arsenic acid with HjS and placing, in
stoppered bottle, in boiling water for one hour; or by passing a rapid
stream of H.S into an HCl solution of H^AsO^ (Bunsen, A., 1878, 192, 305;
Brauner and' Tomicek, J. C, 1888, 53, 146); 2H3A8O4 + 5H2S + xHCl =
A82S5 + 8H2O + xHCl . Carbon disulphide extracts no sulphur from the
precipitate, indicating the absence of free sulphur. The presence of
FcClg or heating the solution does not reduce the A82S5 to A8283 . If there
be a small amount of HCl and the HjS be passed in slowly about 15 per
cent of A82S3 is formed: 2H3A8O4 + SHgS + xHCl = ASjSj + S2 +
8H0O + xHCl . If NH^Cl be present more A82S3 is formed. According
to Thiele (C C, 1890, 1, 877), arsenic acid cold treated with a slow stream
of HjS gives arsenous sulphide, while the hot acid with a rapid stream of
the gas gives the pentasulphide. Arsenic sulphide has the same solubili-
ties as arsenous sulphide. When distilled with hydrochloric acid gas
arsenous chloride is formed (A8CI5 is not known to exist). The solutions
in the alkali hydroxides, carbonates and sulphides form arsenates and
thioarsenates (5r). Ammonium sulphide added to a neutral or alkaline
solution of arsenic acid forms arsenic sulphide which remains in solution
as ammonium thioarsenate (5c). The addition of acid at once forms
arsenic sulphide, not arsenous sulphide and sulphur. The reaction is
much more rapid than with hydrosulphuric acid and is facilitated by
warming.
Arsitie, A8H3 , does not combine with hydrosulphuric acid until heated
to 230°, while stibine, SbH, , combines at the ordinary temperature (Brunn,
R, 1889, 22, 3202).
Acidulated solutions of arsenic boiled with thiosulphates form arsenous
sulphide (separation from Sb and Sn) (Lesser, Z., 1888, 27, 218). Arsenic
may be removed from sulphuric acid by boiling with barium thiosulphate
and no foreign material is introduced into the acid: As^O^ + 3BaSo03 =
As.Ss + 3BaS04 ; 2H;,As04 + 5Na.S203 = AS2S3 + SNa^SO^ + S. + 3HoO.
(Thorn, /. C, 1870, 29, 517; Wagner, Dingl, 1875, 218, 321). "
Sulphurous acid readily reduces arsenic acid to arsenous acid : H3A8O4 +
H^SO., = H3ASO3 + HoSb, (Woehler, A., 1839, 30, 224).
f, — The arsenic from all arsenical compounds treated with concentrated
§69, 6». ARSENIC. 61
hydrochlorio acid and then distilled in a current of hydrochloric acid gas,
passes into the distillate as arsenous chloride, AsCl, . Nearly all of the
arsenic will be carried over in the first 50 cc. of the distillate. This is a
very accurate quantitative separation of arsenic from antimony and tin
and from other non- volatile organic and inorganic material. The AsClj
passes over at 132°, condenses with HCl and may be tested with SnCl^
{g)y or, after decomposition with water (oc) by the usual tests for arsenous
acid (Huf Schmidt, B., 1884, 17, 2245; Beckurts, Arch. Pharm., 1884, 222,
684; Piloty and Stock, B., 1897, 30, 1649).
Eydrobromic acid in dilute solutions is without action upon the acids
of arsenic. The concentrated acid reduces arsenic acid to arsenous acid :
H^O^ + 2HBr = H3A8O3 + Br^ + E,fi. Hydriodic acid reduces
arsenic acid to arsenous acid with liberation of iodine. This is a method
of detecting As^ in the presence of As'". 0.0001 gram of HjAsO^ may be
detected in the presence of one gram of As^Og : 2H3ASO4 + 4HI = ASoO^
+ 21, + 5H2O (Naylor, J. C, 1880, 38, 421).
Chloric and bromic acids oxidize arsenous compounds to arsenic acid with
formation of the corresponding hydracid: SASaO. + 2HBrO, + OHjO =
i>H,A80« + 2HBr . Iodic acid oxidizes arsenous compounds to arsenic acid
with liberation of iodine: SAs^O. + 4HI0, + ISHjO = 10HsAsO4 + 2lj .
g. — Stannous chloride, SnClj , reduces all compounds of arsenic from their
hot concentrated hydrochloric acid solutions, as llocculent, black-brown, metal-
loidal arsenic, containing three or four per cent of tin. The arsenic, in solution
with the concentrated hydrochloric acid, actn as arsenous chloride: 4AsCl, +
5SnCl, = A84 -f BSnCl* . The hydrochloric acid should be 25 to 33 per cent; if
not over 15 to 20 per cent, the reaction is slow and imperfect.
In a wide test-tube place 0.1 to 0.2 srram of the (oxidized) solid or solution
to be tested, add about 1 gram of ttodium chloride, and 2 or ^ cc. of sulphuric
acid, then about 1 gram of crj'stallized stantious chloride; agitate, and heat to
boiling several times, and set aside for a few minutes. Traces of arsenic give
only a brown color; notable proportions give the flocculent precipitate. A
dark gray precipitate may be due to mercury (§58, 6y), capable of being gath-
ered into globules. If a precipitate or a darkening occurs, obtain conclusive
evidence whether it contains arsenic or not, as follows: Dilute the mixture
with ten to fifteen volumes of about 12 per cent hydrochloric acid; set aside,
decant; gather the precipitate in a wet filter, wash it with a mixture of h^'dro-
chloric acid and alcohol, then with alcohol, then with a little ether, and dry in
a warm place. A portion of this dry precipitate is now dropped into a small
hard-glass tube, drawn out and closed at one end, and heated in the flame;
arsenic is identified by its mirror (7), easily distinguished from mercury
(§58, 7). Antimony is not reduced by stannous chloride; other reducible
metals give no mirror in the reduction-tube. Small proportions of organic
material impair the delicacy of this reaction, but do not prevent it. It is
especially applicable to the hydrochloric acid distillate, obtained in separation
of arsenic, according to f.
h. — Chromates boiled with arsenitcs and sodium bicarbonate give chromium
arsenate (Tarugi, J. C, 1896, 70, ii, 340 and 390).
i. — ICagnesium salts with ammonium chloride and ammonium hydroxide
pre cipit ate from solutions of arsenates, magnesium ammonium- arsenate,
KgNH«A804 , white, easily soluble in acids. The reagents should be first
mixed together, and used in a clear solution (** magnesia mixture*^) to make
sure that enough ammonium salt is present to prevent the precipitation of
magnesium hydroxide by the ammonium hydroxide. The crystalline precipi-
€2 ARSENIC. §69, 6;.
tate forms slowly but completely. Compare with the corresponding magnesium
ammonium phosphate (§189, 6d). Maynvnium arsenite is insoluble in water, but
is soluble'in ammonium hydroxide and in ammonium chloride (distinction from
arsenates).
j. — Silver nitrate solution precipitates From neutral solutions of arsenites, or
ammonio-silver nitrate * precipitates from a water solution of arsenous oxide,
silver Hi'senite, AgsAsO, , yellow, readily soluble in dilute acids or in ammonium
hydroxide (§59, 6fr). Neutral solutions of arnenates are precipitated as ailier
arsenate, Ag,As04 , reddish brown, hiving the same solubilities as the arsenite.
A*.— Copper sulphate solution precipitates from neutral solutions of arsenites.
or ammonio-copper sulphate (prepared in the same manner as the ammonio-
silver oxide described above) precipitates from water solutions of arsenous
oxide, the green eopper arsenite, CuHAsOs (Scheele*s green), soluble in ammo-
nium hydroxide and in dilute acids. Copper acetate, in boiling solution, pre-
cipitates the (jreen eopper aceto-arsenite (CuOA8sOs)aCu(CxH,02)2 (Schweinfurt
green), soluble in ammonium hydroxide and in acids. Both these salts are
often designated as Paris green (§77, Cmj). Copper sulphate with excess of freft
alkali is reduced to cuprous oxide with formation of alkali arsenate (10).
K.AsO, -f 2CuS04 -f 4K0H = K.AsO^ -f 2K2SO, + Cu^O -f 2H5O . Solutions
of arsenates are precipitated by copper sulphate as copper arsenate, CuHAsOt ,
greenish blue, the solubilities and conditions of precipitation being the same
as for the arsenites.
/. — Ferric salts precipitate from arsenites, and freshly precipitated ferric
hydroxide (used as an antidote, Wormlej*, 246), forms with arsenous oxide,
variable basic ferric arsenites, scarcely soluble in acetic acid, soluble in hydro-*
chloric acid. Water slowly and sparingly dissolves from the precipitate the
arsenous anhydride: but a large excess of the ferric hydroxide holds nearly all
the arsenic insoluble. To some extent the basic ferric arsenites are trans-
posed into basic ferrous arsenates, insoluble in water, in accordance with the
red\icing power of arsenous oxide. In the presence of alkali acetates, arsenic
acid, or acidulated solutions of arsenates, are precipitated by ferric salts as
ferric arsenate. FeAsO^ , vellowish white, insoluble in acetic acid (compare
§126, 0(/).
wi. — Ammonium molylxlate, (NH4)...Mo04, in nitric acid solution, when slightly
warmed with a solution of arsenic acid or of arsenates gives a yellow precipi-
tate of ammonium arseno^nolffbdate, of variable composition. No precipitate is
formed with As"'. This precipitate is verj- similar in appearance and proper-
ties to the ammonium phospho-molybdate: except the latter precipitates com-
pletely in the cold.
(i'. Special Reactions, a. — Marsh's Test. — Arsenic, from all of its solu-
ble eompoundv^, is reduced by the action of dilute sulphuric or hydrochloric
acid on uhq, forming at first metallic arsenic and then arsenous hydride,
AsH,, gaseous: As.O, + GZn + (>H.SO, = ^AsH, + GZnSO, + STLJ) ;
H,A80, + 4Zn + 4H,S0, = A8H3 -f IZnSO^ -f 4H2O . The arsenic is
precipitateil with the other metals of the second group by hydrogen
sulphide, separated with antimony, tin (gold, platinum and molybdenum)
by yellow anunoninm sulphide. This solution is precipitated by dilute
hydrochloric acid and the mixed sulphides, well trashed,^ are dissolved in
hydrochloric acid using as small an amount of potassinm chlorate crystals
as possible. The solution is boiled (till it does not bleach litmus paper)
* Prepared by addinir ammonium hydroxide to a solution of silver nitrate tUl the precipitate
at AT9t prtxiuced is ntarly aU redissolved.
t If the ammonium salts are not thoroughly removed by washing there is dancer of the for-
mation of the very explosi\-e chlorile of nitnj^ren i5«a»% l> when the precipitate is treated
with hydrochloric acid and |K>ta$«ium chlorate.
§88,6'a. ARSSEXic. 63
to remove excess of chlorine and is then ready for the Karth apparatus.
This apparatus consists of a strong Erlenmeyer flask of about 125 cc.
capacity fitted with a two hole rubber stopper. Through one hole is passed
a thistle (safety) tube, reaching nearly to the bottom of the flask ; in the
other is fitted a three-inch Marchand calcium chloride tube, which projects
just through the stopper and is filled with glass-wool and granular calcium
chloride to dry the gases generated in the flask. To the other end of
the Marchand tube is fitted, with a small cork or rubber stopper, a piece
of hard glass tubing of six mm. diameter and one foot long. This tube
should be constricted one-half, for about two inches, beginning at the
middle of the tube and extending toward the end not fastened to the
calcium chloride tube. The outer end of the tube should also be con-
stricted to about one mm. inner diameter. A short piece of rubber tubing
should connect this constricted end with a piece of ordinary glass tubing,
dipping into a test tube about two-thirds filled with a two per cent solu-
tion of silver nitrate. The rubber tubing should make a close joint with
the constricted end of the hard glass tube, and yet not fit so snug but that
it can be easily removed.
From 10 to 20 grams of granulated zinc * are placed in the flask with
sufficient water to cover the end of the thistle tube. Four or five cubic
centimeters of reagent sodium carbonate are added and the stopper
tightly fitted to the fiask. Dilute sulphuric acid (one^ of acid to three of
water) should now be added, very carefully at first,t until a moderafe
evolution of hydrogen is obtained.
The hydrogen should be allowed to bubble through the silver nitrate
for about five minutes. There should be no appreciable blackening of
the solution (§59, 10), thus proving the absence of arsenic from the zinc
and the sulphuric acid. The purity of the reagents having boon estab-
lished the solution containing the arsenic may be added in small amounts
at a time through the thistle tube. If arsenic be present there will be
almost immediate blackening of the silver nitrate solution.
6AgN0. -f AsH. -f 3H2O = r,Ag + H,AsO, -|- GHNO,
The hard glass tube should now be heated J to redness by a fiame from
• The zinc and aU the reagents should be absolutely free from arsenic. If the zinc bo strictly
chemioaUy pure It will be but slowly attacked by the acid. It should be platinized (S210« *n) or
should oontain traces of iron. Hote {A. CTi., 18S4, (6), 3, 141) removes arsenic from zinc by ad<lin^
anhydrous KffCla to the molten metal, AsCl, is evolved. The zinc purified in this way i?
readily attacked by acids.
t The acid first added decomposes the alkaU carbonate forminer carbon dioxide which rapidly
displaces the air and grreatly lessens the dangrer of explosion when the gtLB is ignited. I f too
much acid be added before the carbonate is decomposed violent frothing may take place and
the liquid contents of the fiask forced into the calcium chloride tube.
X Before heating the tube or igniting the gas, a towel should be wrapped around the fiask to
insure safety in case of an explosion due to the imperfect removal of the air ; or the tube con-
necting the hard glass tube with the Marchand tube should be of larger size and provided with
a plug of wire gauze (made of 10 or 20 circles of gauze lAie size of the tube).
64 ARSENIC. §69, 6'6.
a Bunsen burner provided with a flame spreader. The flame should be
applied to the tube between the ealciimi chloride tube and the constricted
portion. The tube should be supported to prevent sagging in case the
glass softens, and it is customary to wrap a few turns of wire gauze around
the portion of the tube receiving the heat. The heat of the flame decom-
poses^ the arsine and a mirror of metallic arsenic is deposited in the con-
stricted portion of the tube just beyond the heated portion. This may
be tested as described under c 1. When a sufficient mirror has been
obtained the flame is withdrawn, and, removing the rubber tube, the
escaping gas * is ignited.
h. ArsenoTis Hydride (arsine), AsHg , burns when a stream of it is ignited
where it enters the air, and explodes when its mixture with air is ignited.
It burns with a somewhat luminous and slightly bluish flame (distinction
from hydrogen); the hydrogen being first oxidized, and the liberated
arsenic becoming incandescent, and then undergoing oxidation; the vapors
of water and arsenous anhydride passing into the air: 2ASH3 -|- SOj =
^^'2^:i + 3HaO . If present in considerable quantity a white powder may
be observed settling on a piece of black paper placed beneath the flame.
If the cold surface of a porcelain dish be brought in contact with the
flame the oxidation is prevented and lustrous black or brownish-black
spots of metallic arsenic are deposited on the porcelain surface; 4ASH3 +
3O2 = As^ -f 6H0O . A number of spots should be obtained and all the
tests for metallic arsenic applied. The arsenic in the silver nitrate solu-
tion is present as arsenous acid and can be detected by the usual tests (6e)
by first removing the excess of silver nitrate with dilute hydrochloric acid
or calcium chloride.
To generate arsine, maprnesiiim or iron t may be used, instead of zinc, and
hydrochloric acid instead of sulphuric acid. Arsine cannot be formed in the
presence of oxidizing agents as the halogens, nitric acid, chlorates, hypo-
chlorites, etc. Arsinuretted hydrogen (arsine) may also be produced from
arsermuft compounds by nascent hydrogen generated in alkaline solution. Sodium
amalgam,^ zinc (or "zinc and magnesium) and potassium hydroxide or alumi-
num and potassium hydroxide may be used as the reducing agent. There is
no reaction with AsV , or with compounds of antimony (§70, 6/); hence when
* Arsine is an exceedingly poisonous gas, the inhalation of the unmixed gas being quickly
fatal. Its dissemination in the air of the laboratory, even in the small portions vrhich are not
appreciably poisonous, should be avoided. Furthermore, as It is recognized or determined, in
its various analytical reactions, only by its decomposition, to permit it to escape undecomposed
is so far to fail in the object of its production. The evolved gas should be constantly run into
silver nitrate solution, or kept burning.
t According to Thiolo {C. C, 1890, 1, 877) arsenic may be separated from antimony In the Marsh
test by using electrolytic ally deposited iron instead of zinc. Stibine is not evolved. According
to Sautermeister (Analyst, 1891, 218) arsine is not produced when hydrochloric acid act^ upon
iron containing arsenic, but if several grams of zinc be added a very small amount of arsenic in
the iron may bo detected.
t Sodium amalgam is conveniently prepared by adding (in small pieces at a time") one part of
sodium to ci<rht parts (by weight) of dry mercury warmed on the water bath. When cold the
amalgam becomes solid and is easily broken. It should be presen-ed in well stoppered bottles.
^69, G'e. AR8EXIC, 65
the arsenic is present in the triad condition (Asv may be reduced to As'" by
SOa) the use of one of the above reajjents serves admirably for the detection
of arsenic in the presence of antimony. This experiment may be made in a
test-tube, the arsenic being detected by covering the tube with a piece of filter
paper moistened with silver nitrate. It is very difficult to drive over the last
traces of the arsenic and therefore the method is not satisfactory for quanti-
tative work (Uager, J. C, 1885, 48, 838; Johnson, C. A'., 1878, 38, 301; and Clark.
J. C, 1893, 63. 884).
If ferrous sulphide contains metallic iron and arsenic, arsine may be gen-
erated with the hydrogen sulphide. It cannot be removed by washing the
gases with hydrochloric acid (Otto, B„ 1883, 16, 2947).
Arsine does not combine with hydrogen sulphide until heated to 230**, while
fitibine, SbH, , combines at ordinary temperature (method of separation)
(Bnmn, B., 1889, 22, 3202; Myers, J, C, 1871, 24, 889). As dry hydrogen sul-
phide is without action upon dry iodine, it may be freed from arsine by passing
the mixture of the dried gases through a tube filled with glass wool inter-
spersed with dry iodine. AsH. + 31^ = Asl, -f 3HI (Jacobson, B., 1887, 20.
1999). Arsenous hydride is decomposed by passing through a tube heated to
redness (mirror in March test) 4AsH, = A84 + 6H, . Nitric acid oxidizes it
to arsenic acid, 3 AsH, + 8HN0, = 3HsAs04 -f SNO + 4H2O; and may be used
instead of silver nitrate to efPect a separation of arsine and stibine in the
Marsh test. The nitric acid solution is evaporated to dryness and the residue
thoroughly washed with water. Test the solution for arsenic with silver
nitrate and ammonium hydroxide (Ag,As04 , reddish brown precipitate, 6/).
Dissolve the residue in hydrochloric or nitrohydrochloric acid and test for
antimony with hydrogen sulphide (Ansell, J. C.,*1853, 5, 210).
c. — Comparison of the mirrors and spots obtained with arsenic and anti-
mony. — 1. Both the mirror and spots obtained in the Marsh test ex]iil)it
the properties of elemental arsenic (5a). The reactions of these deposits
having analytical interest are such as distinguish arsenic from antimony.
Arsexic Mirror. Antimony Mirror.
Deposited beyond the flame; the Deposited before or on both sides
gas not being decomposed much be- of the flame; the gas being decom-
low a red heat. posed considerably below a red heat.
Volatilizes in absence of air at The mirror melts to minute glob-
450® (1), allowing the mirror to be ules at 433°, and is then driven at
driven along the tube; it does not a red heat,
melt.
By vaporization in the stream of The vapor has no odor,
gas, escapes with a garlic odor.
By slow vaporization in a cur- By vaporization in a current of
rent of air a deposit of octahedral air, a white amorphous coating is
and tetrahedral crystals is obtained, obtained; insoluble in water, soluble
forming a white coating soluble in in hydrochloric acid, and giving rc-
water and giving the reactions for actions for antimonous oxide.
arsenous oxide.
66
ARSEMC.
§68, ii'c.
The heated mirror combines with
hydrogen sulphide, forming the
lemon-yellow arsenous sulphide,
which, being volatile, is driven to
the cooler portion of the tube.
The dry sulphide is not readily
attacked by dry hydrochloric acid
gas (6/).
Arsenic Spots.
Of a steel gray to black lustre.
Volatile by oxidation to arsenous
oxide at 218°.
Dissolve in hypochlorite.*
Wanned with a drop of ammon-
ium sulphide form yellow spots,
soluble in ammonium carbonate, in-
soluble in hydrochloric acid (6^).
With a drop of hot nitric acid,
dissolve clear. The clear solution,
with a drop of solution of silver
nitrate, when treated with vapor of
ammonia, gives a brick-red precipi-
tate.
The solution gives a yellow pre-
cipitate when warmed with a drop
of ammonium molybdate.
With vapor of iodine, color yel-
low, by formation of arsenous
iodide, readily volatile when heated.
The heated mirror combines with
hydrogen sulphide forming the
orange antimonous sulphide, which
is not readily volatile.
The sulphide is readily decom-
posed by dry hydrochloric acid gas,
forming antimonous chloride which
is volatile, and may be driven over
the unattacked arsenous sulphide.
Antimony Spots.
Of a velvety brown to black sur-
face.
Volatile, by oxidation to anti-
monous oxide, at a red heat.
Do not dissolve in hypochlorite.
Warmed with ammonium sul-
phide, form orange-yellow spots, in-
soluble in ammonium carbonate,
soluble in hydrochloric acid (§70.
Ge).
With a drop of hot dilute nitric
acid, turn white. The white fleck,
by action of nitric acid treated with
silver nitrate and vapor of ammo-
nia, gives no color until warmed
with a drop of ammonium hydrox-
ide, then gives a black precipitate*.
With the white fleck no further
action on addition of ammonium
molybdate.
With vapor of iodine, color more
or less carmine-red, by formation
of antimonous iodide, not readily
volatile by heat.
* The hypochlorite reagent, usually NaClO, decomposes in the air and Ught on standing.
It should instantly and perfectly bleach litmus paper (not redden it). It dissolves arsenic by
oxidation to arsenic acid. ASf + lOlWaClO + 6H3O ^ 4H,As04 + lONaCl.
§69, e'd. ARSENIC, 67
2. To the spot obtained on the porcelain surface, add a few drops of
nitric acid and heat; then add a drop of ammonium molybdate. A yellow
precipitate indicates arsenic. Antimony may give a white precipitate
with the nitric acid, but gives no further change with the ammonium
molybdate (Deniges, C, r., 1890, 111, 824).
3. Oxidize the arsenic spot with nitric acid and evaporate to dryness.
Add a drop of silver nitrate or ammonio-silver nitrate (6;). A reddish-
brown precipitate indicates arsenic.
4. After the formation of the mirror in Marsh's test the generating
flask may be disconnected and a stream of dry hydrogen sulphide passed
over the heated mirror. If the mirror consists of both arsenic and anti-
mony, the sulphides of both these metals will be formed, and as the
arsenous sulphide is volatile when heated, it will be deposited in the cooler
portion of the tube. The sulphides being thus separated can readily be
distinguished by the color. If now a current of dry hydrochloric acid
gas be substituted for the hydrogen sulphide the antimonous sulphide
will be decomposed to the white antimonous chloride which volatilizes and
may be driven past the unchanged arsenous sulphide (5c).
5. The tube containing the mirror is cut so as to leave about two inches
on each side of the mirror and left open at both ends. Incline the tube
and beginning at the lower edge of the mirror gently heat, driving the
mirror along the tube. The mirror will disappear and if much arsenic
be present a white powder will be seen forming a ring just above the
heated portion of the tube. This powder consists of crystals of arsenous
oxide, and should be carefully examined under the microscope and iden-
tified by their crystalline form (Wormley, 270).
r>. The crystals of arsenous oxide obtained above are dissolved in water
and treated with ammonio-silver nitrate forming the yellow silver arse-
nite (fi;): or with ammonio-copper sulphate forming the green copper
arsenite ((^k) (Wormley, 259). Any other test for arsenous oxide may be
applied as desired.
7. Magnesia mixture (Gi) is added to the solution of the mirror or spots
in nitric acid. A white crystalline precipitate of magnesium ammonium
arsenate, HgNH^AsO^ , is formed (Wormley, 316).
d^ — ^Beinsch's Test. — If a solution of arsenic be boiled with hydrochloric acid
and a strip of bripht copper foil, the arsenic is deposited on the copper as a
gray film. Hager (C. C, 1886, OSO) recommends the use of brass foil instead of
copper foil. When a large amount of arsenic is present the coating of arsenic
operates from the copper in scales. The film docs not consist of p\ire metallic
arsenic, but appears to be an alloy of arsenic and copper. Arsenous compounds
are reduced much more readily than arsenic compounds. The hydrochloric
acid should compose at least one-tenth the volume of the solution. The arsenic
is not deposited if the acid is not j)resent. This serves as one of the most
KBtisfactory methods of determining the presence or absence of arsenic in
G8 ARSENIC, §69, 6V.
hydrochloric acid. Dilute the concentrated acid with five parts of water and
boil with a thin strip of brig-ht copper foil. A trace of arsenic if present will
soon appear on the foil. For further identification of the deposit, wash the
foil with distilled water, dry, and heat in a hard glass tube, as for the oxida-
tion of the arsenic mirror (GV% 5). The crj'stals may be identified by the mic-
roscope and by any other tests for arsenous oxide. It is important that the
surface of the copper should be bright. This is obtained by rubbing the sur-
face of the foil with a file, a piece of pumice or sand-paper just before usinsr.
The copper should not contain arsenic, but if it does contain a small amount
no film will be deposited due to its presence unless agents are present which
cause partial solution of the foil. If a strip of the foil, upon boiling with
hydrochloric acid for ten minutes, shows no dimming of the brightness of
the copper surface; the purity of both acid and copper may be relied upon for
the most exact work. Antimony, mercury, silver, bismuth, platinum, palladium
and gold are deposited upon copper when boiled with hj'drochloric acid. Under
certain conditions most of these deposits may closely resemble that of arsenic.
Of these metals mercury is the only one that forms a sublimate when heated
in the reduction tube (7), and this is readily distinguished from arsenic by
examination under the microscope. Antimony may be volatilized as an amor-
phous powder at a very high heat. Organic material may sometimes give a
deposit on the copper which also yields a sublimate, but thfs is amorphous and
does not show the octahedral crvstals when examined under the microscope
(Wormley, 209 and ff.: Clark, J. C\ 1893, 63, 880).
f. — Detection in Case of Poisoning. — Arsenic in its various compounds is
largely used as a poison for bugs, rodents, etc., and frequently cases arise of
-accidental arsenical poisoning. It is also iised for intentional ^ioisoning, chiefly
suicidal. It is usually taken in the form of arsenous oxide (white arsenic), or
" Fowler's Solution '* (a solution of the oxide in alkali carbonate). One hun-
dred fifty to two hundred milligrams (two to three grains) are usually suflRcient
to produce death. Violent vomiting is a usual symptom and death occurs in
from three to six hours. In cases of suspected poisoning vomiting should be
induced as soon as possible by using an emetic followed by demulcent drinks,
or the stomach should be emptied by a stomach pump. Freshly prepared ferric
hydroxide is the u.sual antidote, of which twenty-five to fifty grams (one to
two ounces) may be given. The antidote may be prepared by adding magnesia
(magnesium oxide), ammonium hydroxide, or cooking soda (sodium bicarbo-
nate) to ferric chloride or muriate tincture of iron: straining in a clean piece
of muslin, and washing several times. If magnesia be used it is not necessary
to wash, as the magnesium chloride formed is helpful rather than injurious.
A portion of the ferric hydroxide oxidizes some of the arsenous compound,
being itself reduced to the ferrous condition, and forming an insoluble ferrous
arsenate. When the ferric oxide is in excess the ferrous arsenate does not
appear to be acted upon by the acids of the stomach. Of course it will be seen
that the ferric hydroxide will have no efPect upon the arsenic which has
entered into the circulation.
It frequently' becomes necessary for the chemist to analyze portions of sus-
pected food, contents of the stomach, urine; or, if death has ensued, portions
of the stomach, intestines, liver, or other parts of the body. At first a careful
examination should be made of the material at hand for solid white particles,
that would indicate arsenous oxide. If particles be found they can at once be
identified by the usual tests. Liquid food or liquid contents of the stomach
should be boiled with dilute hydrochloric acid, filtered and washed and the
filtrate precipitated with hydrogen sulphide, etc. When solid food or portions
of tissue are to be analyzed, it is necessary first to destroy the organic miaterial.
Several methods have been proposed:
(1) Method of Fresenius and Babo. — The tissue is cut in small pieces and
about an equal weight of pure hydrochloric acid added to this, enough water
should be added to form a thin paste and dilute the hydrochloric acid five or
six times. The mass is heated on the water bath and crystals of potassium
chlorate added in small amounts at a time with stirring until a clear yellow
liquid is obtained containing a very small amount of solid particles. The
heating is continued until there is no odor of chlorine, but concentration shouM
§69, i. ARSENIC. 69
be avoided by the addition of water. The solution should be cooled and filtered;
the arsenic now being present in the filtrate as arsenic acid. This solution
should be treated with sodium bisulphite or sulphur dioxide to reduce the
arsenic acid to arsenous acid and then the arsenic may be precipitated with
hydrogen sulphide. It is advisable to pass the hydrogen sulphide through the
warm liquid for twenty-four hours to insure complete precipitation. A yel-
lowish precipitate of organic matter will usually ])e obtained even if arsenic
be absent. The precipitate should be filtered, washed, and then dissolved in
dilute ammonium hydroxide, which separates it from other sulphides of the
silver, tin and copper groups, that may be present. A portion at least of the*
precipitated organic matter will dissolve in the ammonium hydroxide. The
filtrate should be acidulated with hydrochloric acid, filtered and washed.'
Dissolve the precipitate in concentrated nitric acid and evaporate to dryness.
Kedissolve in a small amount of water, add a drop of nitric acid, filter and test
the filtrate by Marsh's test or any of the other tests for arsenic.
(2) Hydrochloric acid diluted alone may be used for the disintegration of
the soft animal tissues. The solution will usually be dark colored and viscous
and not at all suited for further treatment with hydrogen sulphide; but may
be at once subjected to the Reinsch test (6'd).
(3) ICethod of Danger and Flandin. — The tissue may be destroyed by heat-
ing in a porcelain dish with about one-fourth its weight of concentrated sul-
phuric acid. When the mass becomes dry and carbonaceous it is cooled,
treated with concentrated nitric acid and evaporated to dryness. Moisten with
ijrater, add nitric acid, and again evaporate to dryness: and repeat until the
niass 18 colorle'ss. Dissolve in a small amount of water and test for arsenic by
the usual tests. This method is objectionable if chlorides are present as the
volatile arsenous chloride will be formed.
(4) Method by distillation with hydrochloric acid. The finely divided tissue
IS treated, in a retort, with its own weight of concentrated hydrochloric acid
and distilled on the sand bath. Salt and sulphuric acid may be used instead of
hydrochloric acid. A receiver containing a small amount of water is connected
to the retort and the mass distilled nearly to dryness. If preferred, gaseous
hydrochloric acid may be conducted into the retort during the process of dis-
tillation, in which case all the arsenic (even from arsenous sulphide (5r)) will
be carried over in the first 100 cc. of the distillate. The receiver contains the
arsenic, a great excess of hydrochloric acid and a small amount of organic
nnatter. To a portion of this solution the Reinsch test may be applied at once
and other portions may be diluted and tested with hydrogen sulphide or the
solution may at once be tested in the Marsh apparatus.
For more detailed instructions concerning the detection and estimation of
arsenic in organic matter, special works on Toxicology and Legal Medicine
must be consulted. The following are valuable works on this subject: Micro-
Chemistry of Poisons, Wormley: Medical Jurisprudence, Taylor: A System of
Legal Medicine, Hamilton: Ermittelung von Giften, Dragendorff; Poisons,
Taylor; etc.
7. Ignition. — Metallic arsenic is obtained by igniting any compomifl
containing arsenic with potassium carbonate and charcoal,* or with potas-
sium cyanide:
2A8,0. -f 6KCN = AS4 + 6KCN0
2A8,S, + 6KCN = As, + 6KCNS
2A8,S, + 6Na,C0, + 6KCN = As^ + 6Na,S + 6KCN0 + 6C0, .
4H, AsO, + 5C = As, + SCO, + GH^O
♦ A very suitable carbon for the reductioD of arsenic is obtaioed by Ignitinj? an alkali tartrate
in abflenoe of air to oomplete carbonization.
72 ASTIMOyY. §70, 1.
the use of the copper-iron wire couple for the detection of small quantities
of arsenic by reduction to the elemental state. O.OOOOO'J'5 grams may be
detected. In solution As^ is reduced to As'" by HgPOy, HjS, H^SOj,
NaaS^Oa (6e), HCl , HBr , HI (6/), HCNS , etc. As^ and As'" arc reduced
to As~^"H3 by nascent hydrogen generated by the action of Zn and dilute
H^SO^ , or, in general, by any metal and acid which will give a ready
generation of hydrogen, as Zn, Sn, Fe, Mg, etc., and H^SO^ and HCl
(Draper, Dinghy 1872, 204, 320). As'" is reduced to As-^'H, by nascent
hydrogen generated in alkaline solution as, Al and EOH, Zn and HOH,
sodium amalgam, etc. (separation from antimony) (Davy, Ph, C, 187G,
17, 275; Johnson, C. N., 1878, 38, 301).
§70. Antimony (Stibium) Sb = 120.4. Valence three and five.
1. Properties.— »S*peri/fc gravity, 0.697 (Schroeder, J., 1859, 12). Melting point,
432° (Ledebur, Wild. Beibl., 1881, 650). Boiling point, between 1090** and 1450"
(Camelley and Williams, J. C, 1879, 35, 566). Its molecular weight is unknown,
as its vapor density has not been takf*n. Antimony is a lustrous, silver white,
brittle and readily pulverizable metal. It is but little tarnished in dry air and
oxidizes slowly in moist air, forming* a blackish gray mixture of antimony and
antimonous oxide. At a red heat it burns in the air or in oxygen with incan-
descence, forming white inodorous (distinction from arsenic) vapors of anti-
monous oxide.
2. Occurrence. — Native in considerable quantities in northern Queensland,
Australia (Mac Ivor, C. 2S\, 1888, 57, 64); as stibnite, Sb2S8; as valentinite, Sb«Ox:
in very many minerals usually combined with other metals as a double sulphide
(Campbell, Phil. Mag., 1860, (4), 20, 304; 21, 318).
3. Preparation. — (a) The sulphide is converted into the oxide by roasting in
the air, and then reduced by fusion with coal or charcoal. (//) The sulphide is
fused with charcoal and sodium carbonate: 2SboS, + 6Na3CO, -f 3C = 4Sb -f-
6NajS + 9CO3 . (r) It is reduced by metallic iron: SbzS, -f 3Fe = 2Sb -f- 3FeS .
(d) To separate it from other metals with which it is frequently combined
requires a special process according to the nature of the ore (Dexter, J. pr.^
1839, 18, 449; Pfeifer, A., 1881, 209, 161).
4. Oxides. — Antimony forms three oxides, Sb.O, , SbzO^ , and Sb.Os . (a)
Antimonous oxide, SbjO, , is formed (/) by the action of dilute nitric acid upon
Sb°; {2) by precipitating SbCl, with NajCO, or NH4OH: (,*f) by dissolving Sb*"
in concentrated HoSO« and precipitating with Na^COa: {If) by burning antimony
at a red heat in air or oxygen; (.5) by heating Sb304 or SbjOs to 800** (Baubigny,
C, r., 1897, 124, 499, and 560). It is a white powder, turning yellow upon heat-
ing and white again upon cooling; melts at a full red heat, becoming crystalline
upon cooling: slightly soluble in water, fairly soluble in glycerine (5^>). Anti-
monous oxide sometimes acts as an acid, SbsO, + 2NaOH = 2NaSb02 -f H;0:
but more commonly as a base. Ortho and pyro antimonous acids are known
in the free state. The meta compound exists only in its salts (/)., 2, 1. 19S).
{h) Diantimony tetroxide, Sb204 , is formed by heating Sb** , SbjS, , Sb..O, ,
or Sb.jOa in the air at a dull red heat for a long time. The antimony in this
compound is probably not a tetrad, but a chemical union of the triad and
pentad: 2Sb304 = 2Sb'"Sbv04 = SbjOs.Sb.Os . It is found native as antimony
ochre, {e) Antimonic oxide, SbaOs , is formed by treating Sb** , Sb-O, or
Sb304 with concentrated nitric acid. When heated to 300° it loses oxygen,
forming Sb204 (Geuther, J. />r., 1871, (2), 4, 438). It is a citron-yellow powder,
insoluble in water but reddening moist blue litmus paper. Antimonic acid
exists in the three * forms, analogous to the arsenic and phosphoric acids,
• Beilstein and Blaese (C. C, 1889, 803i have prepared a number of antlmoDates and coDclude
that the acid is always the meta, H SbOa .
§70, 56. AXTIMOXT. 7a
t. f., ortho, meta and pyro (Geuther, /. r., and Conrad, C. .V., 1879, 40, 198). The
ortho acid, HsSbO« is formed by the decomposition of the pentachloride with
water and washing" until the chloride is all removed (Conrad, /. c, and Dau-
brawa, 4., 1877, 186, 110). The most of the antimonates formed in the wet way
by precipitation from the acid solution of antimonic chloride are the ortho
antimonates. By heating the ortho acid to 200** the meta acid, HSbO, , is
formed. Strong ignition of Sb^Og with potassium nitrate and extraction with
water gives the potassium metantimonate, KSbO, , and by adding nitric acid
to a solution of this salt the free acid is formed. The ortho acid dried at 100**
gives the pyro acid: 2HsSb04 = H^Sb.OT + HjO (Conrad, 1. c), which upon
further heating to 200° gives the meta acid. The pyroantimonic acid forms
two series of salts, IC^SbsOr and ICsHsSbaOr . The sodium salt N'a.HjSbzOr
is insoluble in water and is formed in the quantitative estimation of antimony
(9), and also in a method for the detection of sodium (§206, 6^;). For the latter
the soluble potassium salt KsH^SbsOT is used as the reagent. It is prepared
by fusing antimonic acid with a large excess of potassium hydroxide: then
dissolving, filtering, evaporating and digesting hot, in syrupy solution, with a
large excess of potassium hydroxide, best in a silver dish, decanting the
alkaline liquor, and stirring the residue to granulate, dry. This reagent must
be kept dry, and dissolved when required for use; inasmuch as, in solution, it
changes to the tetra potassium pyroantimonate, K^SbsO? , which does not
precipitate sodium. The reagent is, of course, not applicable in acid solutions.
The reaction is as follows: K^S.Bb.O^ + 2NaCl = Na^HaSb^O^ + 2KC1 .
The ortho acid, H,Sb04 , is sparingly soluble in water, easily soluble in KOH,
but insoluble in NaOH. The meta acid, HSbO, , is sparingly soluble in water,
easily soluble in both the fixed alkalis: the pyro acid, H^SbsOr , is sparingly
(more easily than the meta) soluble in water; the normal fixed alkali salts,
H^SbaOr , are soluble in water, also the acid potassium salt, KnH.SboO, , but
not the corresponding sodium salt, Na.H-..Sbs07 .
5. Solubilities. — «. — Metal. — Antimony is attacked but not dissolved by nitric
acid, forming Sb^O, (a) or SbjO^ (/>), depending upon the amount and degree
of concentration of the acid: it is slowly dissolved by hot concentrated sulphuric
acid, evolving SO, and forming Sb2(S04)g (c); it is insoluble in HCl out of con-
tact with the air, but the presence of moist air causes the oxidation of a small
amount of the metal to Sb^O, , which is dissolved in the acid without evolution
of hydrogen (Ditte and Metzner, .4. C/i., 1S96, (6), 29, 389).
The best solvent for antimony is nitric acid, followed by hydrochloric acid or
nitrohydrochloric acid containing only a small amount of nitric acid. Anti-
raonous chloride, SbCl, , is at first formed (r/), but if sufficient nitric acid be
present this is rapidly changed to antimonic chloride, SbCl^^ (e). If, however,
too much nitric acid be present, the corresponding oxides (not readily soluble
in nitric acid) are precipitated (6r). The halogens readily attack the metal
forming at first the corresponding trihalogen compounds (d). Chlorine and
bromine (gas) unite with the production of light, and if the halogen be in
excess, the pentad chloride (r) or bromide is formed (Berthelot and Petit, A. Ch.y
1891, (6), 18, 65). The pentiodide, Sbl^ , does not appear to exist (Mac Ivor,
J. C 1876, 29, 328).
(a) 2Sb -f 2HN0, = Sb,0, + 2N0 + H,0
(6) 6Sb + lOHNO, = 3Sb,05 + lONO + 5H,0
(r) 2Sb + (>H,SO, = Sb,(SO,), + 3S0, + GH^O
(d) 2Sb + 3CL = 2SbCl.
(e) SbCl, -f CI3 = SbCl,
6. — Oxides. — .\ntimonou8 oxide, Sb.O, , is soluble in 55,000 parts of water at
15« and in 10,000 parts at 100° (Schulze, ./. /V., 1883, (2), 27, 320); insoluble in
alcohol; soluble in hydrochloric («), sulphuric and tartaric (h) acids with
formation of the corresponding salts. The dry ignited oxide is scarcely at all
soluble in nitric acid; the moist, freshly precipitated oxide, on the other hand,
dissolves readily in the ailute or concentrated acid, be it hot or cold. Under
certain conditions of concentration a portion of the antimony precipitates out
upon standing as a white crystalline precipitate. It is soluble in the fixed
74 AXTIMOXY. §70, OC.
alkali hydroxides with formation of metantimonites (c) (Terreil, A. Ch.^ 1866,
(4), 7, 350). Fixed alkali carbonates dissolve a small amount of the oxide with
the probable formation of some antimonite (d) (Schneider, Pogg,^ 1859, 108, 407).
It is fairly soluble in glycerine (Kohler, Ditigl., 1885, 258, 520).
(a) Sb,0, -f 6HC1 = 2SbCl, + 3HaO
(h) Sb,0. + H,C,H,0, = (SbO),C,H,Oe + H,0
(c) Sb,0, + 2K0H = 2KSbO, + H,0
(d) Sb,0, + NaaCO, = 2NaSbOa + CO,
Antimony tetroxide, Sb304 , is insoluble in water, slowly dissolved by hot
concentrated hydrochloric acid. Antimonic oxide, SbjOs , is insoluble in water;
soluble in hydrochloric and tartaric acids without reduction; hydriodic acid
dissolves it as antimonous iodide with liberation of iodine (6f); slowly soluble
in concentrated fixed alkalis; soluble in alkaline solution of glycerine (Kohler,
J. C, 1886, 50, 428). The hydrated oxides of antimony (acids) have essentially
the same solubilities as the oxides (4).
f. — »S'<f/^8.^Antimonous chloride, SbCl, , is very (leHquescent, decomposed by
pure water, forming a basic salt; soluble in water strongly acidulated with an
inorganic acid, or tartaric, citric, or oxalic acids (66), but not when acidulated
with acetic acid; it is also soluble in concentrated solutions of the chlorides of
the alkalis and of the alkaline earths (Atkinson, C. A\, 1883, 47, 175). The
bromide and iodide are dcliqueacefit and require moderately concentrated acid to
keep them in solution. The sulphate, Sb3(S04)a , dissolves in moderately con-
centrated sulphuric acid. Antimonous tartrate and the potassium antimonous
tartrate (tartar-emetic) are soluble in water without acidulation; the latter is
soluble in glycerine and insoluble in alcohol. The trichloride, bromide and
iodide are soluble in hot CS,; the chloride and bromide are soluble in alcohol
without decomposition, but the iodide is partially decomposed by alcohol or
ether (Mac Ivor, ,/. f., 1876, 29, 328).
The pentachloride, SbCl, , is a liquid, very readily combining with a small
amount of water to form crystals containing one or four molecules of water.
The addition of more water decomposes the salt forming the basic salt; if,
however, a few drops of HCl have been added first, any desired amount of
water (if added at one time) may be added without causing a precipitation of
the basic salt. If after acidulation water be added slowly, the basic salt will
soon be precipitated.
Antimonous sulphide, Sb^S., , is readily soluble in KjS , and on evapora-
tion large yellow transparent crystals of K^SbaS. are obtained (a) (Ditte,
C, r., 1886, 102, 168 and 212). It is soluble in moderately concentrated
HCl with evolution of HoS (h); slowly decomposed by boiling with water
into Sb.O,, and H.,S (c); and on boiling with NH^Cl into ShCl^ and (NHJ.^S
(de Clermont, T. r., 1879, 88, 972). Dilute HaSO^ is almost without action,
dilute HNO3 gives SbaOg (d). Sparingly soluble in hot NH^OH solution,
soluble in the fixed alkalis (on fusion or boiling) (e); insoluble in (liUjaCO^
(distinction from arsenic); insoluble in the fixed alkali carbonates in the
cold but on warming they effect complete solution (/) (distinction from
tin); very sparingly soluble in normal ammonium sulphide; readily soluble
in yellow ammonium sulphide with oxidation (g) (6f). The pentasulphide,
SboS,^ , is insoluble in water; soluble in the alkali sulphides (fe), and in the
fixed alkali carbonates and hydroxides; insoluble in ammonium carbonate
and sparingly soluble in ammonium h3^droxide, more readily when warmed
{D., 2, 1, 217). On boiling with water it slowly decomposes into SbjO, ,
|70, 5d. A\TIM0\Y. 75
H,S and S (Mitscherlich, J. pr., 1840, 19, 455). Hydrochloric acid on
wanning dissolves it as SbCly {\):
{a) Sb,S, + 2K,S = K«Sb,SB
(6) Sb^S, + 6EC1 = 2SbCl, + 3H,S
(c) Sb,S, 4- 3H,0 = Sb^O, + 3H,S
(d) 2Sb,S, + 4HN0. = 2Sb,0, -f- -^S, + 4N0 + 2H,0
(c) 2Sb,S, 4- 4K0H = rKSbS, + KSbO, + 2H,0
(D 2Sb,S, + 2Na,C0, = 3NaSbS5 + NaSbO, -f- 200,
(y) 2Sb,S, + C(NH,),S, = 4(NH,)aSbS4 + S,
(*) Sb,S, + 3(NH,)3S = 2(NH4),SbS,
(i) Sb,Ss + 6H01 = 2Sb01, 4- 3H,S + S,
A — Water* — With the exception of the compounds of antimony with
some organic acids, as tartaric and citric, all salts of antimony are decom-
posed by pure WATER. For this reason it will be seen that water is a very
important reagent in the analysis of antimony salts. The salts with
inorganic acids all require the presence of some free acid (not acetic) to
keep them in solution. If the acid be tartaric the further addition of
water causes no precipitation of the antimony salt. Water decomposes
the inorganic acid solutions precipitating the basic salt, setting more acid
free which dissolves a portion of the basic salt. The addition of more
water causes a further precipitation and at the same time dilutes the acid
so that upon the addition of a sufficient amount of water a nearly com-
plete precipitation may be obtained. If the precipitate of the basic salt be
washed with water the acid is gradually displaced, leaving finally the anti-
mony as oxide.
With solutions of antimonons chloride the basic salt precipitated is
white antimonous oxychloride, Sb4Cl,0.., , " Powder of Algaroth," soluble
in tartaric acid (distinction from bismuth, §76, 5</) (Mac Ivor, C. N., 18T.">,
-32, 229), 4SbCl3 + 5H,0 = Sb^Cl^O^ + lOHCl . The basic salt repeatedly
washed with water is slowly (rapidly if alkali carbonate be used) changed
to the oxide, Sb.O., (Malaguti, J. pr., 1835, 6, 253), Sb.CLO, + H,0 =
2Sb.A + 2HC1. With antimonic chloride, SbCl, , the basic salt is
SbOCl^ ; SbCa, + H^O = SbOCl, + 2HC1 (Williams, C, N.. 1871, 24, 224).
Solutions of the tartrates of antimony and of antimony and potassium
are not precipitated on the addition of water; and antimonous chloride
*The acidity of water solutions of certain salts having a weak base and the alkalinity of
others containing a weak acid is due to a partial decomi>osition (hydrolysis) of the salt by the
ions of the water, B« and OH', forming again the original acid and base. lVa,CO,, for instance,
it split up into the weak non-di88<K^iate<l H.CO, and the strongly-dissociated NaOH, whose
OH ions give the "alkaline reaction." FeCl, in water forms soluble colloid F«(OH)3, which
may be separated by dialysis from the free HCl resulting or precipitated by addition of a
neutral salt, as HaCl, to the dilute solution ; KCN gives alkaline KOII nnrl non-<lissociated
HCK, readily det<!!cted by its odor. In other cases precipitation is caused, as in the treatment
of bismuth or antimony solutions with wator or ti\ heating 'N'lisZnOa solution, hydrolysis in
general being Increased bj* raising the temperature. The action of water on soap belongs to
this class.
76 AXTlMoyY. §70, 6a,
dissolved in excess of tartaric or citric acid solution is not precipitated on
addition of water.
6. Beactions. — a, — The alkali hydroxides and carbonates precipitate front
acidulated solutions of inorganic antimonpus salts, antimonous oxide,* SbgO,
((i) (Rose, Pogg., 1825, 3, 441), white, bulky, readily becoming crystalline on
boiling; sparingly soluble in water (56), readily soluble in excess of the fixed
alkalis^ forming a metantimonite (ft) (Terrell, A. Ch,, 1866, (4), 7, 350); slowly
soluble in a strong excess of a hot solution of the fixed alkali carbonate (c>
(distinction from tin); insoluble in ammonium hydroxide or ammonium car-
bonate. The freshly precipitated oxide is readily soluble in acids (not in acetic
acid). If the alkaline solution of the antimony be carefully neutralized with
an acid (not tartaric or citric) the oxide is precipitated (d) and at once dissolved
by further addition of acid. The presence of tartaric or citric acids prevents
the precipitation of the oxide by means of the alkalis or alkali carbonates.
The solutions of antimonous oxide by alkalis is due to combination with 'them,
acting as a feebly acidulous anhydride and forming antimonites, which are
found to be monobasic^ so far as capable of isolation. Sodium antimonite,
NaSbO, , is the most stable and the least soluble in water: potassium anti-
monite, KSbO, , is freely soluble in dilute potassium hydroxide solution, but
decomposed by pure water. By long standing (24 hours), a portion of the
antimonous oxide deposits from the alkaline solution, and the presence of alkali
hydrogen carbonates causes a nearly complete separation of that oxide (c).
(a) 2SbCla + 6K0H = Sb,0, -f 6KC1 -f 3HaO
2SbCl, 4- 3Na,C0, = Sb,0, -f 6NaCl + 3C0,
(6) SbjOa -f 2K0H = 2KSb02 + H,0
or SbCl, + 4K0H = KSbO, + 3KC1 + 2H,0
(c) Sb,0. 4- Na,CO, = 2NaSbO, -f- CO,
(d) 2KSbO, + 2HC1 = Sb^Os + 2KC1 + H,0
(e) 2NaSbO, + 2NaHC0. = Sb,0. -f 2Na,C0. + H,0
Antimonlc salts are precipitated under the same conditions as the antimonous
salts. The freshly formed precipitate is the orthoantimonic acid, HgSbO^ =
SbO(OH), = Sb2d«,3H20 («) (Conrad. C. N., 1879, 40, 198); insoluble in am-
monium hydroxide or carbonate: soluble, more readily upon warming, in
excess of the fixed alkali hydroxides and carbonates as metantimonate (b).
(a) SbCl, -f 5K0H = SbO(OH). + 5KC1 + TL,0
(b) SbO(OH)8 + KOH = KSbO, + 2H,0
h. — The freshly precipitated antimonous oxide is soluble in oxalic acid, but
(in absence of tartaric acid) the antimony soon slowly but completely separates
out as a white crystalline precipitate; unless an alkali oxalate be present, when
the soluble double oxalate is formed. The precipitate of antimony oxalate
dissolves upon the further addition of hydrochloric acid. Freshly precipitated
antimonic oxide dissolves readily in oxalic acid and does not separate out upon
standing. Acetic acid precipitates the solutions of antimony salts if tartaric
acid be absent. Potassium cyanide gives a white precipitate with antimonous
salts soluble in excess of the cyanides.
With potassium ferrocyanide antimonous chloride (not tartrate) gives a
white precipitate, f,ohible in hydrochloric acid (distinction from tin), or fixed
alkali hydroxides (Warren, C .V., 1888, 67, 124). Potassium fcrricyanide is
reduced' to ferrocyanide by antimonous salts in alkaline solution (Baumann,
Z. anoew,, 1892, 117).
r.._From the solutions of the fixed alkali antimonites or antimonates the
oxides or hydrated oxides (acids) are precipitated upon neutralization with
nitric acid (or other inorganic acids) ; the freshly formed precipitates readily
* Men'chiitkin (pajre 185) says the precipitate formed by the action of alkalis upon antimonous
salts is the meta acid, HSbO..
§70, 6e. AXTiMoyr, 77
dissolving in an excess of the acid. Antimonous nitrate is rery unstable and
the antimonic nitrate is not known to exist. It is quite probable that these
solutions in nitric acid are merely solutions of some of the hydrated oxides
(acids).
d. Compounds of antimony with the acids of phosphorus are not known,
(KaaHPO^ does not precipitate antimony salts, separation from tin, §71, 6</).
e. Hydrogen sulphide precipitates, from acid * solutions of antimonous
salts, antimonous sulphide (a), SbjSj, orange-red; in neutral solutions
(tartrates) the precipitation is incomplete. In strong fixed alkali solu-
tions (6a) the precipitation is prevented, or rather the sulphide first
formed (6) is at once dissolved in the excess of the fixed alkali (c), sparingly
in HH^OH. The alkali sulphides give the same precipitate sparingly
soluble in normal ammonium sulphide, readily soluble in the fixed alkali
sulphides (d) and in yellow ammonium sulphide (e). Antimonous sulphide
is slowly decomposed by boiling water (f) ; insoluble in ammonium carbon-
ate (distinction from As); slowly soluble in boiling solution of the fixed
alkali carbonates (g) (distinction from Sn) ; soluble in hot moderately con-
centrated hydrochloric acid (h) (distinction from arsenic). The alkaline
solutions of antimonous sulphide are oxidized upon standing by the oxygen
of the air or rapidly in the presence of sulphur (e); from the alkaline solu-
tions hydrochloric acid precipitates the antimony as trisulphide, penta
sulphide or a mixture of these, depending upon the degree of oxidation (i)^
(a) 2SbCl, -h 3H,S = Sb,S, -f 6HC1
(6) 2KSbO, + 3HaS = Sb^S, -f 2K0H -f- 2H,0
(c) 2Sb,S, 4- 4K0H = 3KSbS, -f- KSbO, -f 2H,0
(d) Sb,S. + K^S = 2KSbS,
(c) 2Sb,S3 -h 6(NH,),S, = 4(NH,)3SbS4 + S,
(f) SbaS, + 3H,0 = Sb,0, -h 3H,S
(g) 2Sb,S. -h 2K2CO, = 3KSbS, -f- KSbO, -f- 2C0,
(h) Sb,S. + 6HC1 = 2SbCl, + 3H,S
(0 3KSbS, -h KSbO, + 4HC1 = 2Sb,S, + 4KC1 + 2H,0
OP 2(NrH,),SbS4 + 6HC1 = Sb.S, -f- 6NH4CI + SH^S
Hydrosniphnrio acid f and alkali sulphides precipitate (under like condi-
tions as for antimonous salts), from solutions of antimonic salts, antimonic
sulphide, SbjSg , orange, having the same solubilities as the tri-sulphide.
The alkaline solution of the sulphide consists chiefly of the ortho-thioanti-
monate instead of the meta, as in antimonous compounds. Sb^S^ -|- 3K.S
= 2K3SbS, ; 4Sb2S5 + 18K0H = SK^SbS^ + SKSbO, + OH^O . When
dissolved in HCl the penta-sulphide is reduced to SbCla with liberation
of sulphur, SbjSs -f 6HC1 = 2SbCl5 -f 3HoS + S^ .
•According to LoAiton (J, C, 1888, 54, 992) the precipitation takes place in the presence of
quite strong' hydrochloric acid (one to one) separation from tin, which is precipitated only when
three or more parts of water are present to one of the acid.
t In order to precipitate pure antimonic sulphide, the solution of the antimonic salt must be
cold, and the hydrof^cn sulphide added rapidly. If the solution be warmed or the hydrogen
sulphide added slowly more or less antimonous sulphide is precipitated (BOsek, J, C, 1895, 67»
515).
78 ANTIMONY. §70, 6/.
All salts of antimony when wanned with sodium thloBalphate, NasSsO, ,
are precipitated as the sulphide (separation of arsenic and antimony). 2SbCl,
4- .SNa^SjOs + 3H3O = Sb.S, + SNa^SO^ + 6HC1 . Sulphurous acid reduces
antimonic salts to antimonous salts (Knorre. Z. angeic, 1888, 155). Sulphates of
antimony are not prepared by precipitation, but by boiling the oxides with
strong sulphuric acid. They dissolve only in very strongly acidulated water.
/. — Antimony occurs most frequently for analysis as the chlorides, it is
therefore important that the student familiarize himself with the deport-
ment of these salts with the various reagents, used in qualitative analysis.
The most important of the properties have been discussed under 5a, h. c, d.
Hydrocliloric acid, or any other inorganic acid, carefully added to a solu-
tion of antimony salts in the fixed alkalis will precipitate the correspond-
ing oxide or hydrated oxide, soluble upon further addition of the acid.
Potassiuin iodide added to antimonous chloride solution, not too strongly
acid, gives a yellow precipitate of antimonous iodide, soluble in hydro-
chloric acid. The precipitation does not take place in the presence of
tartaric or oxalic acids. Hydriodic acid (or potassium iodide in acidu-
lated solutions) added to solutions of antimonic salts causes a reduction
of the antimony to an anfimonoiis salt with liberation of iodine (distinc-
tion from Sn'^:" SbCl^ + 2HI = SbClg -f 2HC1 -f I^ . The iodine may be
detected by heating and obtaining the violet vapors, or by adding carbon
disulphide and shaking. It should be remembered that the solution to
be tested must be acid, for in alkaline solutions the reverse action takes
place, iodine oxidizing antimonous salts to antimonic salts: SbCl, +
8K0H + I2 = K^SbO, -f 2KI -f 3KC1 + 4HoO (Weller, A,, 1882, 213,
364). Also the absence of other oxidizing agents which liberate iodine
from hydriodic acid must be assured.
[/. — If antimony and arsenic compounds occurring together are strongly
oxidized with nitric acid there is danger that the insoluble precipitate of anti-
monic oxide may contain arsenic, as antimonic arsenate, in.sohible (Menschut-
kin). Stannous chloride reduces antimonic compounds to the antimonous
condition, but in no case caiKses a precipitation of the metal (distinction from
arsenic).
h. — Antimonous salts in acid, neutral or alkaline solution, rapidly reduce
solutions of chromates to chromic compounds. Acid solutions of antimonous
salts reduce solutions of manganates and permanganates to manganous salts:
with alkaline solutions to manganese dioxide. These reactions are capable of
quantitative application in absence of other reducing agents. The antimony is
oxidized to the antimonic condition (9 and 10).
I. — An antimonous compound when evaporated on a water bath with an
ammoniacal solution of silver nitrate gives a black precipitate (Bunsen, .4.,
1855, 106, 1). A solution of an antimonous compound in fixed alkali when
treated with a solution of silver nitrate gives a heavy black precipitate of
metallic silver, insoluble in ammonium hydroxide, and thus separated from the
precipitated silver oxide. If instead of a water solution of silver nitrate, a
solution with great excess of ammonium hydroxide (one to sixteen) be added,
no precipitation occurs in the cold (distinction from Sn"); nor upon heating
until the excess of ammonia has been driven off. Antimonates with silver
nitrate give a white precipit^ite of silver antimonate, soluble in ammonium
hj'droxide.
§70, 6;. AXTi^foyr, 79
j. — Stibine. — By the action of zinc and sulphuric or hydrochloric acid all
compounds of antimony are first reduced to the metallic state. The
formation of stibine is a secondary reaction and requires the moderately
rapid generation of hydrogen in acid solution. If a few drops of a solu-
tion of an antimony salt, acidulated with hydrochloric acid, be placed
upon a platinum foil and a small piece of zinc be added, the antimony is
immediately deposited as a black stain or coating adhering firmly to the
platinum; 2SbCll, + 3Zn = 2Sb + SZnClj . In this test tin, if present,
deposits as a loose spongy mass, while arsenic, if present, does not adhere
so firmly to the platinum as the antimony. In the presence of arsenic
this test should be applied with caution under a hood as a portion of the
arsenic is almost immediately evolved as arsine (§69, 6'&).
If hydrogen be generated more abundantly than in the operation above
mentioned, by zinc and dilute sulphuric or hydrochloric acid, the gaseous
antimony hydride, stihitie, SbHj , is obtained for examination. For com-
parison with arsine and details of manipulation see " Marsh's Test " under
arsenic (§69, 6'a) :
Sb,0, -h BZn 4- 6H,S04 = 6ZnS04 + 3H,0 + 2SbH,
SbCla -h 3Zn -h 3HC1 = aZnCla + SbH.
Stibine is a coloriess, odorless gas, not nearly so poisonous as arsine. It
bums with a luminous and faintly bluish-green flame, dissipating vapors
of antimonous oxide and of water (a); or depositing antimony on cold
porcelain held in the flame, as a lusterless brownish-black spot (&). The
gas is also decomposed by passing through a small glass tube heated to
low redness (c), forming a lustrous ring or mirror in the tube. The stibine
is decomposed more readily by heat than the arsine and the mirror is
deposited on both sides of the heated portion of the glass tube. The spots
and mirror of antimony are compared with those of arsenic in §69, G'c.
The antimony in stibine is deposited as the metal when the gas is passed
into a concentrated solution of fixed alkali hydroxide or when it is passed
through a U tube filled with solid caustic potash or soda-lime (distinction
and separation from arsenic).
(a) 2SbH, -h 30, = Sb,0, + 3H,0
(ft) 4SbH. + 30^ = 4Sb + 6HaO
(c) 2SbH, = 2Sb + 3H,
When the antimony hydride (stibine) is passed into a solution of silver
nitrate, the silver is reduced, heaving the antimony with the silver, as
antimonous argentide, SbAg^^ , a black precipitate, distinction from arsenic,
which enters into solution (S69, (>'</ and b); SbH^ + ^AgNOg = SbAgg +
3HN0, . The precipitate should bo filtered and washed free from unde-
oomposed silver salt (and arsonous acid, if that be present), and dissolved
with dilute hydrochloric acid (HCl docs not dissolve uncombined anti-
80 AXTIMOXY. §70, 7.
mony, 5a) : SbAgg + 6HC1 = SbClg -f 3AgCl + SHg . The solution con-
sists of antimonous chloride, leaving silver chloride as a precipitate.
However, in the excess of hydrochloric acid used a small portion of the
fiilver chloride may be dissolved (§69, 5c), interfering with the final test
for the antimony. If this be the case the silver should be removed by a
drop of potassium iodide (8).
Rtibine is not evolved by the action of strong KOH upon zinc or aluminum,
nor by sodium amalgam in neutral or alkaline solution (distinction from triad
arsenic); the antimony is precipitated as the metal (Fleitmann, J, C, 1852, 4,
329). Stibine is slowly oxidized by sulphur to SbsS, in the sunlight at ordinary
temperature and rapidly when the sulphur (in a U tube mixed with glass wool)
is heated to 100°. The reaction takes place according to the following equation:
2SbH3 + ?.S., = Sb.S, + 3H,S (Jones, J, C, 1876, 29, 645).
7. Ignition. — By ignition in the absence of reducing agents, antimonic acid
and anhydride are reduced to antimonous antimonate, SbsOstSbsO;^ or Sb^Ot
(Sb'"Sbv04), a compound unchanged at a dull red heat, but when heated to
800° this oxide is further reduced to antimonous oxide (4b).
The antimonates of the fixed alkali metals are not vaporized or decomposed
when ignited in the absence of reducing agents; hence, bj' fusion in the crucible
with sodium carbonate and oxidizing agents, i. e., with sodium nitrate and car-
bonate, the compounds of antimony are converted into non-volatile sodium
pyroantimonate, NafSbaOf , and arsenic compounds if present are at the same
time changed to sodium orthoarsenate, Na^AsOf . If now the fused mass be
digested and disintegrated in cold water and filtered, the antimonate is sepa-
rated as a residue, Na^HsSbjOT (4c), while the arsenate remains in solution
with the excess of alkali. The operation is mucji more satisfactory when the
arsenic and antimony are previou.sly fully oxidized — as by digestion with nitric
acid — as the oxidation by fusion in the crucible is not effected soon enough to
retain all the arsenic or antimony which maj' be in the state of lower oxides,
sulphides, etc. If compounds of tin are present in the operation — and If the
fusion is not done with excess of heat, so as to convert sodium nitrite to caustic
soda and form the soluble sodium stannate — the tin will be left as stannic oxide,
SnO, , in the residue with the NajH.SboO, , But if sodium hydroxide is added
in the operation, the tin is separated as stannate in solution with the arsenic
(Meyer, J. C, 1849, 1, 388).
All compounds of antimony are completely reduced in the dry way on char-
coal with sbdium carbonate, more rapidly with potassium cj'anide; the metal
fusing to a brittle globule. The reduced metal rapidly oxidizes, the white
antimonous oxide rising in fumes, and making a crystalline deposit on the
support. If now ammonium sulphide be added to this white sublimate, an
orange precipitate is a sure indication of the presence of antimony (Johnstone,
C. N., 1883, 68, 296). The same white oxide is formed on heating antimony or
its sulphides in a gla-ss tube, through w^hich air is allowed to pass.
8. Detection. — Antimony is precipitated, from the solution acidulated
with hydrochloric acid, in the second group by liydrosiilphuric acid as the
sulphide (6e). By its solution in yellow ammonium sulphide * it is sepa-
rated from Hg , Pb , Bi , Cu , and Cd . In the Marsh apparatus the anti-
mony is precipitated on the Zn as the metal, a portion being still further
reduced to stibine. By passing the gases, sfihine and arsine, into AgNO..
solution, the antimony is precipitated as SbAg^^ , antimony argentide, sepa-
* Antimony an sulphide solution in potassium sulphide may bo detected electrolytically, beinff
deposited as Sb**. Delicate to one part in 1,500,000 (Kohn, J. Soc. hid., 1891, 10, 827).
^70, 10. AxmioxY. 81
lating it from the arsenic which is oxidized and passes into solution as
arsenous acid. The SbAgg is dissolved in HCl and the presence of the
antimony is confirmed by the precipitation of the orange colored sulphide
with HjS . Study text at 6 and §84 to §89. For distinction between Sb^
and SV" see §89,>.
9. Estimation. — (1) Tartaric acid and water are added to SbCl, , which is
then precipitated by H2S as SbjS, , and after washing on a weighed filter it is
dried at 100** and weighed. If from any cause the precipitate contains free
sulphur, it is separated by heating in CO, . (2) Antimonous oxide, sulphide,
or any oxysalt of antimony is first boiled with fuming nitric acid, which con-
verts it into SbjO. , and then by ignition it is reduced to Sb^O^ , and weighed
as such. (3) The trichloride is precipitated by gallic acid, and weighed after
drying at 100**. (4) In the presence of tin and lead oxidize the hydrochloric
acid solution of the salts with KClOa (the tin must be present as Sniv) and
distil in a current of HCl . The stannic and antimony chlorides are volatile
(separation from lead). To the distillate add metallic iron, obtaining stannous
chloride and metallic antimony; filter and wash (separation from tin). Fuse
the precipitate with sodium nitrate and sodium carbonate, digest the fused
mass with cold water, filter, wash, dry and w^eigh as NasHzShjOr (7) (Tookev,
J. C, 1862, 15, 462; and Thiele, A., 1894, 263, 361). (5) For estimation of anti-
TOonv and separation from arsenic and tin bv the use of oxalic acid, see Lessen
(Z., 1888, 27, 218) and Clarke (C. .V., 1870, 21, 124). (6) VolumetricaUy. The
antimony compound is converted into stibine (6j) and the gas passed into
standard silver nitrate solution. The solution is filtered and the excess of
silver nitrate is titrated with standard sodium chloride. If arsenic be present
it must also be estimated (§69, 9 (/o)), and the true amount of antimony-
present computed from the two determinations (Ilouzeau, J, C, 1873, 26, 407).
(7) Sb"^ is oxidized to Sbv in presence of NaHCO, by a standard solution of
iodine. The end of the reaction is shown bv the blue color given to starch.
(8) Sb'" is oxidized to Sbv in presence of H^C.H^O. by KMnO« . (9) Sb'" is
oxidized to Sbv by KjCroOr , and the excess of KgCraO/used is determined by
a standard solution of FeSO« , K8Fe(CN)« being used to show the end of the
reaction. (10) The antimony as the triad salt is treated with an excess of
standard K|Fe(CN)«; the excess of which is estimated in a gas apparatus with
H,0, (Baumann, Z. angew,, 1892, 117).
10. Oxidation. — Stibine, SbH, , is decomposed by heat alone into anti-
mony and hydrogen (6;). By burning in the air it is oxidized to SboO;.
and HjO . Passed into a solution of silver nitrate, SbAgs is produced, or
passed into a solution of antimonous chloride or potassium hydroxide,
sp. gr. 1.25, metallic antimony is produced. Excess of chlorine, bromine,
or nitric acid in presence of water oxidizes it to Sb^; but if the SbH, be in
excess metallic antimony is precipitated. With excess of iodine in pres-
ence of water Sb'" is produced; if the stibine be in excess metallic anti-
mony. Metallic antimony is oxidized by nitric acid, chlorine or bromine
.to Sb'" or Sb^, depending upon the amount of these reagents and the
temperature. Iodine oxidizes the metal to Sb'" only, except in alkaline
mixtures when Sb^ is formed.
Antimonous compounds are oxidized to antimonic compounds by CI ,
Br , HHO3 , KjCraO^ , and EMnO^ ; by silver oxide in presence of the fixed
alkalis (6i); by gold chloride in hydrochloric acid solution, gold being
82 TIX. §71, 1.
deposited as a yellow precipitate (§73, 10). The antimony is precipitated
as SbjOg unless sufficient acid be present to dissolve the oxide : 4AnCl3 -|-
aSbjOg + GHjO = 4Ati + SSb^Os + 12HC1 .
Antimonic compounds are reduced to antimonous compounds by HI (6/)
and by SnCls (§69 and §71, 10); the antimony not being further reduced
(distinction from As). Antimonic and antimonous compounds are reduced
to the metallic state by Pb , Sn , Bi , Cu , Cd , Fe , Zn , and Mg ; but in
the presence of dilute acids and metals which evolve hydrogen the anti-
mony is still further reduced to stibine. Iron in the presence of platinum
(iron platinum wire couple) precipitates the antimony from acid solutions
as Sb°; 0.000012 grams can be detected (Rideal, C. N,, 1885, 61, 292).
Sodium amalgam with dilute sulphuric acid evolves stibine from all
antimony solutions (Van Bylert, B., 1890, 23, 2968) but the generation
of hydrogen in alkaline solution, i. e., Zn -(- EOH, causes the reduction
of the antimony salt to the metal only, in no case evolving stibine.
§71. Tin (Stannum). Sn = 119.0. Valence two and four.
1. Properties.— S^/>e<'i'/?<? gravity, 7.293 (Rammelsberg, J5., 1870, 3, 724); meltvng
point, 231.68° (Calleiidar and (Jriffiths, C. A\, 1891, 63, 2). Boils between 1450^
and 1000° (Carnelley and Williams, J. C, 1879, 35, 566). Does not distill in a
vacuum at a red heat (Schuller, J., 1884, 1550). Tin is a silver white metal, does
not tarnish readily in pure air. At a red heat it decomposes steam with evolu-
tion of hydrogen: at a white heat it burns in the air with a dazzling white
light, forming SnOa . It is softer than gold and harder than lead, can readily
be hammered or rolled into thin sheets (tinfoil); at 100° it can be drawn into
wire and at 200° can be pulverized. Tin possesses a strong tendency to crystal-
line structure, and when bar or block tin is bent a marked decrepitation
" Zinngeschrei " (Levol, A. C/?., 1859. (3), 56, 110) is noticed, due to the friction
of the crystals. Block tin exposed to severe cold (winter of 1867-68, at St.
Petersburg, — 30°) crumbles to a grayish powder (Fritsche, B., 1869, 2, 112).
This same property of crumbling is noticed in samples of tin that have been
preserved several hundred years (Schertel, J. pr., 1879, 2, 19, 322). Tin forms
alloys with many nietnls. Bronze consists of copper and tin, brass frequently
contains from two to five per cent of tin, solder consi.sts of lead and tin. All
the easily fusible metals as Wood's metal, etc., contain tin. For many refer-
ences concerning tin alloys, see Watts (IV, 720).
2. Occurrence. — The chief ore of tin is cassiterite or tinstone, a nearly pure
crystallized dioxide, found in England, Australia, Malay Peninsula, United
States, etc. (/)., 2, 1, 643). Tin pyrites, impure SnS., , is found in small quanti-
ties in various tin veins.
3. Preparation. — The reducing agent emjjloyed is carbon. The impure ore,
SnO, , is first roasted, which removes some of the arsenic as As.O, , and some
of the sulphur as SOa . Then, by washing, the soluble and some of the in-
soluble impurities are washed away, the heavier SnO, remaining. It is then
fused with powdered coal, lime being introduced to form a fusible slag with
the earthy impurities. It is refined by repeated fusion. Strictly pure tin is
best made bj^ treating the refined tin with HNO, , and then reducing the oxide
thus formed by fusion with charcoal; or by reducing the purified chloride.
4. Oxides and Hydroxides. — Tin forms two stable oxides and corresponding
classes of salts; stannous oxide, SnO , black or blue black, and stannic oxide,
SnOj , white; the latter acts both as a base, in stannic salts, and as an anhy-
dride, in stannates. Stannous oxide is formed (1) by precipitating SnCl, with
KaCO, , washing with boiled water in absence of air, drying at 80° or lower;
then dehydrating by heating in an atmosphere of hydrogen or carbon dioxide
§71, 5b. TIN, 83
(Long^, C. C, 1886, 34); (2) by melting a mixture of SnCl, and NasCO. with
stirring until it becomes black, and removing the NaCl by washing (Sandal,
Phil. Mag., 1838, (3), 12, 216; Bottger, A., 1839, 29, 87). Stannous hydroxide,
Sn(OH),*, white to yellowish white, is formed by adding alkalis or alkali
carbonates to stannous chloride, washing and drying at a low temperature
(Ditte, A. Ch., 1882. (6), 27, 145).
Stannic oxide exists in two forms, crystalline and amorphous. The native
tinstone is nearly pure crystalline SnO, . For preparation see Bourgeois (C r.,
1887, 104, 231) and Levy and Bourgeois (C. r., 1882, 94, 1365). Amorphous SnO,
is formed (!) by heating tin in the air to a white heat; (2) stannic salts are
precipitated by alkali carbonates, the precipitate washed and ignited; (3) tin
is oxidized by nitric acid; (4) tin filings are ignited in a retort with HgQ
(D., 8, 1, 647). Stannic h}fdrox\de or stannic acid exists in two forms: (i) Nor-
mal stannic acid, SnOCOH), = H^sSnO, , is formed when a solution of stannic
chloride is precipitated by barium or calcium carbonate (Frelng, Pogg,, 1842, 55^
519); if an alkali carbonate be used some alkali stannate is also formed. {2)
Metastannic acid, HioS^sOis* , is formed by decomposition of tin with nitric
acid (Hay, C. N„ 1870, 22, 298; Scott, C. N,, 1870, 22, 322); insoluble in acids but
changed on standing with acids to normal stannic acid, which is readily soluble
in acids (56). It is also formed when stannic chloride is boiled in concen-
trated solution with most of the alkali salts: SSnCl* -f- 20Na2SO4 + 15H2O =
H,«S]isOie 4- 20NaCl + 20NaHSO4 , or according to Fresenius (16th edition),
271: SnCl^ + 4Na,S04 + 4H2O = Sn(OH), + 4NaCl + 4NaHS0, .
5. Solubilities. — a, — MctaL — Tin dissolves in hydrochloric arid slowly when the
acid is dilute and cold, but rapidly when hot and concentrated, stannous
chloride and hydrogen being produced {n)\ in dilute sulphuric acid, slowly, with
separation of hydrogen (6), (not at all even in hot acid if more dilute than
HaSO^.GHaO (Ditte, A. Ch., (5), 27, 145); in hot concentrated sulphuric acid,
rapidly, with separation of sulphurous anhj'dride and sulphur (r); nitric acid
rapidly converts it into metastannic acid, insoluble in acids ((/): very dilute
nitric acid dissolves it without evolution of gas as stannous nitrate and am-
monium nitrate (e) (Maumene, BL, (2), 35, 59S): nitro-hydrochloric acid dis-
solves tin easily as stannic chloride (f), potassium hydroxide solution dissolves
it very slowly, and by atmospheric oxidation (g): or, at high temperatures,
with evolution of hydrogen (h). Bromine vapors readily attack melted tin
with formation of SnBr4 , colorless crystals, melting point 30° (Carnelley and
O'Shea, J. C, 1878, 33, 55).
((I) Sn -f 2HC1 = SnCl^ + H,
(h) Sn -h H2SO, = SnSO* -f H,
(c) Sn 4- 2H2S04 = SnSO, + 2H,0 -f SO,
and then 4SnS04 -f 2S0, -f 4H3SO4 = 4Sn(S04), + S, + 4H2O
((f) 15Sn + 20HNO, + 5H,0 = 3H»oSn,0,5 -f 20NO
(6) 4Sn + IOHNO3 = 4Sn(NO,)2 + 3H,0 -f NH^NO,
(f) Sn + 2CI2 = SnCl,
iu) 2Sn + 4K0H + 02= 2K2SnO, + 2H2O
(h) Sn + 2K0H = K^SnO, + H,
h. — Oxides. — Stannous oxide \?i insoluble in water, soluble in acids (Ditte, A. Ch.,
1882, (5), 27, 145; Weber, ,/. T., Ifs82, 42, 1266), oxidized by nitric acid when
heated, forming the insoluble metastannic acid. Stannous hiidroxide is readily
soluble in all the solvents of the oxide, and is also readily soluble in fixed
alkali hydroxides. Siannic o.ridr, SnO. . is insoluble in water; soluble with
difficult^' in alkalis; insoluble in acids except in concentrated H.SO4 (D., 2, 1,
648). Sulphur forms SnSj and SOo; chlorine forms SnCl^ (Weber, Pogg., l«r>l,
112, 619). Normal stannic acid, H..SnO, , freshly precipitated, is soluble in
fixed alkali hydroxides and in acids (Ditte, C. v., 1SS7, 104, 172); insoluble in
water and changed by hot nitric acid to the insoluble metastannic acid.
Metastannic acid, HioSn^O^s , is insoluble in water and acids, HCl changes it to
*Acoordinfir to other authorities Sn(OH)a dors not exist, but a hydnited oxide is formed,
SnO.SiKOH). (Graham-Otto, -4. 2, ISrJT: A, 2, 1, 6r»7; Gmelln-Kraut, 3, 107).
84 Tiy. §71, 5r.
metastannic chloride insoluble in the acid, but soluble in water after removal
of the acid; soluble in the fixed alkalis as metcistannates, which are soluble in
water and precipitated by acids. Metastannic acid in contact witt HCl is
gradually changed to stannic acid (Barfoed, J, pr,, 1867, 101, 368).
c. — Salts. — The sulphides and phosphates of tin are insoluble in water, also
stannous oxychloride: stannous sulphate,* bromide and iodide; and stannic
chloride and bromide dissolve in pure water with little or no decomposition
(Personne, C, r., 1862, 54, 216; and Camelley and O'Shea, J, O., 1878, 33, 55).
Stannous chloride is soluble in less than two parts of water (Engel, A, Ch., 1891,
(6), 17, 347); but more water decomposes it, unless a strong excess of acid be
present: 2SnCl2 + HjO = SnO.SnCl^ + 2HC1 . Pure stannic chloride is a
liquid; sp, gr., 2.2; boiling point, 144**; solidifies at —33** (Besson, C. r., 1889, 109,
940). A small amount of water added to the liquid combines with heat to form
crystals of SnCl^.BHsO , which are readily soluble in excess of water (D., 2, 1,
662). Stannic chloride is not readily decomposed on boiling with water. The
nitrates of tin are very easily decomposed by water and require free acid to
keep them in solution (Weber, J. pr,, 1882, (2), 26, 121; Montemartini, Qazzettn,
1892, 22, 384). Stannic iodide is readily soluble in water (Schneider, Pogg,, 1866.
127, 624). Stannic sulphate is easily soluble in water, but is decomposed by a
large excess (Ditte, C. r., 1887, 104, 171). Stannous and stannic chloride, and
stannic iodide are soluble in alcohol. Stannous nitrate and stannic sulphate,
and bromide are deliqueHcent. Stannous sulphide is insoluble in water, soluble
in HCl with formation of HjS; decomposed by HNOg with oxidation to meta-
stannic acid; insoluble in solution of the normal alkali sulphides, but soluble
in the polysulphides with oxidation to a stannic comjKDund (6e). Stannic sul-
phide is soluble in HCl , with evolution of HjS; and in solutions of the alkali
sulphides.
6. Reactions. — a. Alkali hydroxides and carbonates precipitate from
solutions of stannous salts, stannous hydroxide, Sn(0H)2 (4), white, readily
soluble in excess of the fixed alkali hydroxides, insoluble in water, am-
monium hydroxide and the alkali carbonates (distinction from antimony).
It is also precipitated by barium carbonate in the cold (Schaflfner, A., 1844,
61, 174).
SnCl, -f 2K0H = Sn(OH), + 2KC1
Sn(OH)j -f 2K0H = K,SnO, + 2H,0
SnCl, -f 4K0H = K^SnO, -\- 2KC1 -f 2H,0
SnCla -\- Na^CO, -f H,0 = Sn(OH), + 2NaCl + CO,
By gently heating the solution of potassium stannite, KsSnOj , crystalline
stannous oxide, SnO , is formed. By rapid boiling of a strong potassium
hydroxide solution of stannous hydroxide part of the tin is oxidized and
the remainder precipitated as metallic tin; 2K2Sn02 + HoO = Sn +
K^SnOs + 2K0H . The reaction proceeds more rapidly upon the addition
of a little tartaric acid. Stannic salts are precipitated by alkali hydroxides
and carbonates as stanjiic acid, HoSnOa soluble in excess of the fixed alkali
hydroxides, insoluble in ammonium hydroxide and the alkali carbonates
(Ditte, A. Ch., 1897 (6), 30, 282).
SnCl, + 4K0H = H.SnO, + 4KC1 + H,0
HjSnO, -f 2K0H = K^SnO, -f 2H,0
SnCl, + OKOH = K.SnO, + 4KC1 + 3H,0
SnCl, + 2Na2CO, + H^O = H^SnO, + 4NaCl + 2C0,
* Stannous sulphate Is decomposed by an excess of cold water forming S8n8O4.4SnO.8H,0|
ASd by a small amount of hot water forming SnS04.28nO (Ditte, A. Ch., 1883, (5), 37, 161).
§71, 6e. TIS. 85
Metastannic salts are precipitated as metastannic acid soluble in potassium
hydroxide not too concentrated, not readily soluble in sodium hydroxide,
insoluble in ammonium hydroxide and the alkali carbonates.
h, — Oxalic acid forms a white crystalline precipitate with a nearly neutral
solution of stannous chloride, soluble in hydrochloric acid, not readily soluble
in ammonium chloride. If a nearly neutral solution of stannous chloride be
added drop by drop to a solution of ammonium oxalate, the white precipitate
which forms at once dissolves in the excess of the ammonium oxalate. Stannic
chloride is not precipitated by oxalic acid or ammonium oxalate (Hausmann
and Loewenthal, A., 1854, 89, 104).
Potassiiixii cyanide precipitates both stannous and stannic salts, white, in-
soluble in excess of the cyanides. Fotassiam ferrocyanide precipitates from
stannous chloride solution stannous ferrocyanide^ Sn^FeCCN). , white, insoluble
in water, soluble in hot concentrated hydrochloric acid. Stannic chloride is
precipitated as a greenish white gelatinous precipitate, soluble in hot hydro-
chloric acid, but reprecipitated upon cooling (distinction from antimony)
(Wyrouboff, A, Vh., 1876, (5), 8, 458). Potassium ferricyanide precipitates from
solutions of stannous chloride, stannous ferricyanide, Sn,(Pe(CN)B)j , white,
readily soluble in hydrochloric acid. On warming, the ferricyanide is reduced
to ferrocyanide with oxidation of the tin. No precipitate is formed by the
ferricyanide with stannic chloride.
e. — The nitrates of tin are not stable. Stannous nitrate is deliquescent and
soon decomposes on standing exposed to the air. Stannous salts when heated
with nitric acid are precipitated as SnO,; but if stannous chloride be warmed
with a mixture of equal parts of nitric and hydrochloric acids, stannic chloride
and ammonium chloride are formed (Kestner^ A, Ch,, 1860, (3), 68, 471).
d, — Hypophosphorons acid does not form a precipitate with stannous or
stannic chlorides, nor are these salts reduced when boiled with the acid. Sodium
liypophosphite forms a white precipitate with stannous chloride, soluble in
excess of hydrochloric acid: no precipitate is formed with stannic chloride.
Phosphoric acid and soluble phosphates precipitate from solutions of stannous
salts, not too strongly acid, stannous phosphate, white, of variable composition,
soluble in some acids and KOH; insoluble in water (Lenssen, A., 1860, 114,
113). With stannic chloride a white gelatinous precipitate is formed, soluble
in HCl and KOH , insoluble in HNO, and HCjH.O, . If the stannic chloride be
dissolved in excess of NaOH before the addition of Na2HP04 and the mixture
then acidulated with nitric acid, the tin is completely precipitated as stannic
phosphate (separation from antimony). However, the precipitate always car-
ries a little antimony (Bomemann, Z. angew., 1899, 635).
e, Hydrosnlpharic acid and soluble sulphides precipitate from solutions
of stannous salts dark brown hydrated stannous sulpliide, SnS (a), insol-
uble in dilute, soluble in moderately concentrated HCl (6). It is readily
dissolved with oxidation by alkali supersulphides, the yellow sulphides,
forming thiostannates (c) ; from which acids precipitate the yellow stannic
sulphide (d). The normal, colorless alkali sulphides scarcely dissolve any
stannous sulphide at ordinary temperature, compare (§69, 6e and §70, Ge),
but hot concentrated X^S dissolves SnS forming EoSnSs and Sn (e) (Ditto,
C. r., 1882, 94, 1419; Baubigny, J. C, 1883, 44, 22). Potassium and
sodium hydroxides dissolve it as stannites and thiostannites (/), from
which acids precipitate again the b^o^vn stannous sulphide ((/). Am-
monium hydroxide and the alkali carbonates do not dissolve it (distinction
from arsenic, §69, 6^). The insolubility in fixed alkali carbonates is a
86 TIN, §71, 6f.
distinction from antimony (§70, 6e). Nitrohydrochloric acid (free chlorine)
dissolves it as stannic chloride, with residual sulphur (h). Nitric acid
oxidizes it to raetastannic acid without solution (i) (separation from
arsenic, §69, 6e).
(a) SnCla + H,S = SnS -f 2HC1
(ft) SnS + 2HC1 = SnCl, -f H,S
(c) SnS + (NH,),S, = (NHO,SnS,
(d) (NHJjSnS, + 2HC1 = SnS, + 2NH,C1 + H,S
(c) 2SnS + K,S = K,SnS, + Sn
(f) 2SnS + 4K0H = K,SnO, + K,SnS, + 2H,0
(17) (K,SnOa + K,SnS,) -f 4HC1 = 2SnS + 4KC1 + 2H,0
(/*) 2SnS + 4C1, = 2SnCl, + S,
(0 30SnS + 40HNO, + 10H,0 = GHioSn^Oj^ + 40NO + 15S,
Solutions of stannic salts are precipitated as stannic sulphidey SnSo .
hydrated, yellow, having much the same solubilities as those given for
stannous sulphide, with this difference, that stannic sulphide is moderately
soluble in normal, coloriess, alkali sulphides. The following equations
illustrate the most important reactions:
SnCl* -f 2HjS = SnSj + 4HC1
SnS, + 4HC1 = SnCl, + 2H,S
SnS, 4- (NH,),S= (NHJ,SnS,
2SnS, 4- 2(NH,),S, = 2(NH,),SnS. + S,
3SnS, + 6K0H = K^SnO, + 2K,SnS, + 3H,0
(K.SnO, -f 2K,SnS,) + 6HC1 = aSnS, + 6KC1 + 3H,0
SnS, + 2C1, = SnCl^ + S,
15SnS, -f 20HNO, -f 5H,0 = SHjoSn.O.a + 15S, + 20NO
Sodium thiosnlphate does not form a precipitate with the chlorides of tin
(separation from As and Sb) (Lesser, Z., 1888, 27, 218). Sulphurous acid and
sodium sulphite precipitate from stannous chloride solution not too strongly
acid, stannous sulphite, SnSO, , white, readily' soluble in HCl . When warmed in
the presence of hydrochloric acid, si?lphur dioxide acts as an oxidizing agent
upon the stannous salt. A precipitate of SneOmS, or SnS, is formed, or H,S
is evolved and SnCl^ formed, depending upon the amount of HCl present.
GSnCl, + 2S0a + GHaO == Sn^OjoSa + 12HC1
eSnCl, -f 2S0, + 8HC1 = SnS, + 5SnCl4 + 4H,0
3SnCl, + SO2 + 6HC1 = 3SnCl, + H,S + 2H,0
Stannic chloride does not give a precipitate with sulphurous acid or sodium
sulphite.
The sulphates of tin are formed by dissolving the freshly precipitated
hydroxides in sulphuric acid and evaporating at a gentle heat. They cannot be
formed by precipitation and are decomposed bv water (Ditte, A, Ch., 1882, (5).
27, 145).
f. — Fotassium iodide added to a concentrated water solution of stannous chlo-
ride forms first a yellow precipitate soluble in excess of the SnCl, . Further
addition of KI gives a yellow preeipitale rapidly turning to dark orange needle-
like crystals, often forming in rosette-like clusters. If a drop of the stannous
chloride solution be added to an excess of potassium iodide the yellow precipi-
tate is formed, which remains permanent unless a further quantity of stannous
chloride be added when the orange precipitate is formed. The orange precipi-
tate is probably SnI., , and is soluble in HCl , KOH , and CaHjOH , soluble in
large excess of KI and sparingly soluble in HjO with some decomposition.
§71, 7.. TIN. 87
The yellow precipitate is probably a double salt of stannous iodide and potas-
sium iodide, and has about the same solubilities as the orange precipitate
(Personne, J., 1862, 171; BouUay, A, Ch., 1827, (2), 34, 372). Potassium iodide in
concentrated solution precipitates stannic iodide^ yellow, from very concentrated
-water solutions of stannic chloride. The precipitate is readily soluble in water
to a colorless solution (Schneider, J„ 1866, 229). llydriodic acid does not give
free I with Sniv , distinction from Sbv and Asv (Harroun. J, C, 1882, 42, 661).
The chlorates, bromates and iodates of tin have not been thoroughly studied
{Watts, 1, 539, III., 22; /)., 2, 1, 675). Stannous chlorate appears to be formed
-w'hen potassium chlorate is added to a concentrated water solution of stannous
chloride; it dissolves on addition of HCl, and nearly all dissolves in excess of
-water. With KBrO, , bromine is liberated, and with KIO, iodine is liberated.
Potassium chlorate, bromate and iodate all form precipitates with stannic
chloride, soluble in HCl without liberation of the halogen.
g. — StazmoiiB arsenate, 2SnO.As205 , a voluminous flocculent precipitate is
formed by adding a solution of SnClj to a concentrated acetic acid solution of
XsAsOf , decomposed by heating to As , ASjO, and SnOj (Lenssen, A,, 1860, 114,
115). Stannic arsenate, 2Sn02.As205 , a white gelatinous precipitate is formed
by adding HNO, to a mixture of NajSnO, and NajAsO^ (Haeffely, J., 1855, 395).
With antimony, tin acts as a base, forming stannous and stannic antimonites
and antimonates (Lenssen, I. c).
ft.— Jf potassium chromate be dropped into a hydrochloric acid solution of
stannous chloride there is immediate reduction of chromium with formation
of a dirty brown precipitate. If stannous chloride be carefully added to potas-
sium chromate in excess, an abundant yellowish precipitate is obtained without
much apparent reduction of the chromium. Potassium chromate added to
stannic chloride gives an abundance of bright yellow precipitate soluble in
excess of SnCl^ , insoluble in HoO , soluble with difficulty in HCl . KjCrjOt
also gives a precipitate with SnClj and SnClf (Leykauf, J. pr., 1840, 19, 127).
i. An ammoniacal solution of silver nitrate is reduced to metallic silver
by a solution of potassium stannite. The reagent (silver nitrate solution
one part, to ammonium hydroxide sixteen parts) serves as a delicate test
for the presence of Sn" in solution in KOH. The addition of KOH in
excess to an unknown solution removes all heavy metals except Pb , Sb ,
8n 9 Al ^ Cr 9 and Zn ; of these tin only precipitates metallic silver from the
strongly ammoniacal solution in the cold. Antimonous and arsenous
compounds give the black precipitate of metallic silver if the solution be
boiled.
;. A solution of mercuric chloride, HgClj , reacts with stannous
chloride solution, forming SnCl4 and a precipitate of HgCl (white) or Hg°,
gray, depending upon the relative amounts present (§68, 6g),
Jc. Stannous salts react with (jrE^)^l/Lo0^y giving a blue-colored
solution of the lower oxides of molybdenum, constituting a delicate test
for Sn" (§76, Gg),
7. Ignition. — Before the blow-pipe, on charcoal, with sodium carbonate, and
more readily by addition of potassium cyanide, tin is reduced to malleable
lustrous globules — brought to view (if minute, under a magnifier) by repeated
trituration of the mass with water, and decantation of the lighter particles.
A little of the white incrustation of stannic oxide will collect on the charcoal
near the mass, and, by persistence of the flame on the globules, the same coat-
ing forms upon them. This coating, or oxide of tin, moistened with solution of
cobalt nitrate, and again ignited strongly, becomes of a blue-green color. SnOt
fused with KCN gives metallic tin (Bloxam, J, C, 1865, 18, 97),
88 TIN. §71, 8.
8. Detection. — Tin is precipitated, from the solution acidulated with
hydrochloric acid, in the second group by hydrosulphuric acid, as the sul-
phide (6e). By its solution in yellow ammonium sulphide it is separated
from the Copper Group (Hg , Pb , Bi , Cn , and Cd). By the reaction in
the Marsh apparatus the tin is reduced to the metal and is not dissolved
as long as zinc is still present. The residue Sn (Zn , Sb , Au , and Pt) in
the Marsh apparatus is warmed with hydrochloric acid, which dissolves
the Sn as SnCla . This is detected by its reducing action on HgCl, , giving
a white precipitate of HgCl or a gray one of Hg** (6;).
A short test for the detection of tin in the stannous condition, or after
its reduction to that condition, consists in treating the solution with an
excess of cold KOH (separation of Pb , Sn , Sb , Al , Cr , and Zn , from
all other heavy metals) ; and adding to this solution, filtered if necessary,
a solution of AgNO, in a great excess of NH^OH (one part AgNOg to sixtf en
parts NH4OH). A brown-black precipitate of metallic silver indicates
that tin was present in the stannous condition (6t). Consult also §90
and §92.
9. Estimation. — (1) Gravimetrically. It Is converted into SnO, , and after
ignition weighed. (2) Yohimetrically. To SnCl, add KNaC^H^O, and NaHCO, ,
then some starch solution and a graduated solution of iodine, until a pernaa-
nent blue coloration appears. (3) To SnCl, add slight excess of PeCl, , and
determine the amount of FeCls formed, by a graduated solution of KMnOt .
(4) By electrolytic deposition from a solution of the double oxalate, rendered
slightly acid with oxalic acid.
10. Oxidation. — Metallic tin reduces solutions of Ag , Hg , Bi , Cn , Pt ,
and Au , to the metallic state. Sn" is oxidized to Sn^ by free HNO2 ,
HNOgS H3Fe(CN)e , H^SO, and H^SO^ (if hot), CI , HCIO , HCIO, , HCIO, ,
Br , HBr03 , 1 », and HIO3 . Also bv Pb" (in alkaline solution only), Pb^^ ,
Ag'^ Hg', Hg", As^ As"' (in presence of HCl), Sb^ Mo^i, Bi'", On',
Pd(N08)2 , Pt^^ S Fe'", Fe^^ Cr^i, Co'", Ni'", and Mn^+n. Chlorine, bromine
and iodine act more vigorously in alkaline than in acid mixtures. The
above mentioned metallic forms oxidize Sn" in both acid and alkaline
mixtures.
Stannous chloride is one of the most convenient and efficient of the
ordinary discriminative deoxidizing agents for operations in the wet way.
As stannic chloride is soluble in the solvents of stannous chloride no
' precipitate of tin is made by its reducing action; but many other metals
are so precipitated by reduction to insoluble forms, and are thus identified
in analysis, e. g., mercuric chloride is reduced from solution, first to white
mercurous chloride, and then to gray mercury (detection of mercury);
silver nitrate, to brown-black silver (detection of tin); all soluble com-
» Kestner, A, Ch,, 1880, (8), 58, 471. « Ditte, A. CTi., 1882. (6), 37, 146. » Thomas, C. r., 1896, IM,
1639. * DItte, C. r., 1882, 94, 1114.
§71,10. Tiy. 89
pounds of arsenic in strong HCl (detection of arsenic); bismuth salts, to
metallic bismuth (in alkaline mixture §76, 6/7); and ferric salts, to
ferrous salts, left in solution, much used in volumetric analysis of iron
(9, and §126, 6^ and 9); auric chloride is reduced to the metal by stannous
chloride, forming a colored precipitate varying from brown to reddish-
brown or purple-red according to the amount of stannic chloride present.
This finely divided precipitate of gold is called ^* Purple of Cassius " (Max
ITuUer, J. pr., 1884, 30, 252).
Solutions of Sn^ and Sn" are reduced to the metallic state by Cd , Al ,
Zn , and Mg . According to Bideal (C. N., 1885, 61, 292) 0.00003 grams
of tin in solution may be detected as the metal by reduction, using the
gold zinc wire couple. Stannic salts are reduced to stannous salts by
metallic tin, copper or iron (Allen, /, C, 1872, 26, 274).
90
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§78, 5a. OOLD, 91
§73. Gold (Aurum) An = 197.2. Valence one and three.
1. Properties.— fif|Jeoi/lc gravUy, 19.30 to 19.34 (Rose, Pogy., 1848, 75, 403). Melt-
ing point, 1061.7** (Heycock and Neville, J. C, 1895, 67, 189). It is a yellow metal,
that from different parts of the world varying slightly in color; the presence of
very small traces of other metuls also aifects the color. It is softer than silver
and harder than tin; possesses but little elasticity or metallic ring. It is the most
malleable and ductile of all metals; one gram can be drawn into a wire 2000
metres long. The presence of other metals diminishes the ductility. It may be
rolled into sheets 0.0001 mm. thick. At a very high heat it vaporizes (Deville
and Debray, A. Ch., 1859, (3), 56, 429). It is a good conductor of electricity,
«qual to copper, not so good as silver. It has a high coeflBcient of expansion
and cannot be moulded into forms but must be stamped. On account of its
softness, gold is seldom used absolutely pure, but is hardened by being alloyed
-with other metals, as Ag , Cu , etc.
2. Occxirrence. — Gold is usually found native, but never perfectly pure, being
always alloyed with silver, and occasionally also wnth other metals. It is found
as gold-dust in alluvial sand, sometimes in nuggets, and sometimes disseminated
in veins of quartz.
3. Preparation. — (i) Washing. "Which consists in treating the well-powdered
ore with a stream of water, the heavy gold settling to the bottom. (2) Amalga-
mation. Which consists in dissolving the gold in mercury and then separating
it from the latter by distillation. (S) By fusing with metallic lead, which dis-
solves the gold, the liquid alloy settling to the bottom of the slag. The gold is
afterward separated from the lead by cupellation. The silver is separated from
the gold by dissolving it in nitiic or sulphuric acid. Or the whole is dissolved
in nitrohydrochloric acid, and the gold precipitated in the metallic state by
some reducing agent; ferrous sulphate being usually employed. Another
method is to pass chlorine into the melted alloy. The silver chloride rises to
the surface, while the chlorides of Zn , Bi , Sb , and As (if present) are vola-
tilized, and the pure gold remains beneath. A layer of fused borax upon the
surface prevents the silver chloride from volatilizing. (Jf) By treatment with a
solution of KCN . (.5) By amalgamation with mercury and electrolysis at the
same time.
4. Oxides and Hydroxides. — AurouR oxide, AUjO , is very unstable, heating to
about 250° decomposes it into the metal and oxygen. The hydroxide is pre-
pared by reducing the double bromide with SOo in ice-cold solution; heating to
200** changes it to the oxide - (Kriiss, A., 1886, 237, 274). Auric hydroxide,
Aa(OH), , is prepared by precipitation from the chloride solution with MgO
(Kriiss, /. c). It is a yellow to brown powder, changing to the oxide upon dry-
ing at 100°. Heating to 250° gives the metal and oxygen.
5. Solubilities. — a, — Metal. — CJold is not at all tarnished or in any way acted
upon by water at any temperature, or by hj'drosulphuric acid. Neither nitric
nor hydrochloric acid attacks it under any conditions; but it is rapidly attacked
by chlorine (as gas or in water solution), dissolving promptly in nitrohydro-
chloric acid, as auric chloride, AuCl^; by bromine, dissolving in bromine water,
as auric bromide, AiiBr,; and by iodine; dissolving when finely divided in hydri-
odic acid by aid of the air and potassium iodide, as potassium auric iodide,
KIAiiI.: 4An + 12HI + 4KI + "O, = 4KIAuI, + 6H.0 . Potassium cyanide
jiolution. with aid of the air, dissolves precipitated gold as potassium auro'
eyanide, KA.u(CN).: 4Au + 8KCN + O, + 2H,0 = 4KAu(CN)2 -f- 4K0H .
Gold is separated, from its alloys with silver and base metals, by solution in
nitric acid; the gold being left as a black-brown powder — together with
platinum and oxides of antimony and tin. When the gold-silver or gold-copper
has not over 20 per cent gold, nitric acid of 20 per cent disintegrates the alloy,
and effects the separation; w-hen the gold is over 25 per cent, silver or lead
(three parts) must be added, by fusion, to the alloy before solution. (If gold-
silver alloy contains 60 per cent or more of silver^ it is silver color; if 30 per
cent silver, a light brass color; if 2 per cent silver, it is brass color.)
If gold and other metals are obtained in solution by nitrohydrochloric acid,
leaving most of the silver as a residue, the noble metals can be precipitated by
zinc or ferrous sulphate, and the precipitate of gold, silver, etc., treated with
92 OOLD. §73, 56.
•
nitric acid, which will now dissolve out any proportion of silver not less than
15 per cent, to 85 per cent of gold, and dissolve the baser metals. Concentrated
sulphuric acid dissolves silver, and leaves gold.
6. — The oxides and hydroxides of gold are insoluble in water, soluble in acids.
e. — The salts of the oxyacids are not stable, being decomposed by hot water.
Gold sulphide is insoluble in water or acids, '^except nitrohydrochloric acid,
soluble in alkali sulphides. Aurous salts are decomposed by water, forming
An° and Au'". Auric chloride is deliquesceni; both the chloride and bromide
are readily soluble in water. The iodide is decomposed by water, forming
aurous iodide. The double chlorides, bromides, iodides and cyanides are soluble
in water.
6. Beactions. a. The fixed alkali hydroxides and carbonates in excess
do not precipitate AnCl, solutions, as a soluble aurate, KAnO, , readily
forms; but upon boiling and neutralizing the excess of alkali, Au(OH);,
is precipitated. Ammoninm hydroxide precipitates from concentrated
solutions a reddish-yellow ammonium aurate, (NH8)2Au203 , " fulminating
gold." b. Oxalic acid reduces gold chloride from solutions, slowly (nitric
acid should be absent and the presence of ammonium oxalate is advan-
tageous), but completely. The gold separates in metallic flakes or forms
a mirror on the side of the test-tube. gAuCl, + SHjCgO^ = 2Au + 600.
-f- 6HC1 . As platinum, palladium, and other second group metals are
not reduced by oxalic acid, this method of removal of gold should be
employed upon the original solution before the precipitation of the second
gronp metals as sulphides. Potassium gold cyanide, KCN.An(CN)3 , is
formed when a neutral solution of AnClg is added to a hot saturated
solution of KCN . It is very soluble in water and by heating above 200*^
it is decomposed into CN and ECN.AnCN , which latter product is formed
when gold is dissolved in KCN in the presence of air (5a). c. A solution
of AuClg is precipitated as An° by a solution of ENO^ . d. Sodinm
pyrophosphate forms with AUGI3 a double salt which has found application
in gold plating, e. Hydrosulphuric acid precipitates from gold chloride
solution, hot or cold, gold sulphide, variable from AUjS to AUjS, , brown,
insoluble in acids, hot or cold, except in nitrohydrochloric acid, in which
it readily dissolves; soluble in alkali sulphides to a thio-salt. Alkali
sulphites precipitate gold chloride solution as double sulphite, L e.
Ano(S03)3.(NHj2S08.6NH3 + 3H2O . Upon boiling the sulphite acts as
a reducing agent, giving metallic gold.
f. Potassium iodide, added in small portions to solution of auric chloride
(so that the latter is constantly in excess where the two salts are in
contact), and when equivalent proportions have been reached, gives a yel-
low precipitate of aurous iodide, Aul , insoluble in water, soluble in large
excess of the reagent; the precipitate accompanied with separation of free
iodine, brown, which is quickly soluble in small excess of the reagent as a
colored solution: AuClj + 4KI = Aul + 3KC1 + Ij with KI . But, on
gradually adding auric chloride to solution of potassium iodide, so that the
§74, 1. PLATINUM. 95
latter is in excess at the point of chemical change, there is first a dark-
green solution of potassio-auric iodide, KTA11I3 ; then a dark-green precipi-
tate of auric iodide, Anlg , very unstable, decomposed in pure water, more
quickly by boiling; changed in the air to the yellow aureus iodide.
g. Stannous chloride gives a purple precipitate containing the oxides of
tin with the gold, " purple of Cassius " insoluble in acids.
h. FerrouB sulphate is the most common reagent for the detection of
gold, reducing all gold salts to the metallic state; AuClg -f SFeSO^ =:
Aii + Fe,(S0,)3 + FeCl3.
7. Ignitioii. — Gold is reduced from many of its compounds by lig:ht, and from
all of them by heat — its separation in the dry way being readily effected by
faaion with such reagents as will make the material fusible. Very small pro-
portions are collected in alloy with lead, by fusion; after which the lead is
vaporized in " cupellation " (§59, 7).
8. Detection. — In the dry way gold is detected by fusion of the mineral
matter with lead, to the formation of a " button " which is then ignited
to drive oflE the lead, leaving the gold and silver behind as the metals.
In the wet way the material, if not in solution, is digested with nitro-
hydrochloric acid which dissolves all the gold. The excess of acid is re-
moved by evaporation and the gold is precipitated by oxalic acid or ferrous
sulphate, and identified by its color and insolubility in acids. If the
gold be not removed from the original solution it is precipitated in
Group II. by HjS , passes into Division A (tin group) by (1^4)28 , and may
be detected in the flask of the Marsh apparatus by the usual methods.
9. Xstimation. — Gold it always weighed in the metallic state, to which form
St is reduced: (!) By ignition alone if it is a salt containing no fixed acid; if in
an ore, by mixing w^ith lead and fusion to an alloy, and final removal of the
lead by ignition at a white heat in presence of air. (2) By adding to the solu-
tion some reducing agent, usually PeSO^ , H2C2O4 , chloral hydrate, or some
easily oxidized metal, such as Zn , Cd , or Mg . (S) Gold is also estimated volu-
metrically by H2C2O4 and the excess of HjCaO^ used, determined by XMnO^ .
10. Oxidation. — Gold is reduced to the metalljc state by very many
reducing agents, among which may be mentioned the following: Pb , Ag^,.
Eg, Hg', Sn, Sn", As, As'", AsH,, Sb, Sb'", SVE,, Bi, Cu, Cu',
Pd, Pt, Te, Fe, Fe", Al, Co, Ni, Cr"', Zn, Mg, JL^Cfi,, HNO., P^
E3PO2 , EsPO, , PHg , H2SO3 , and a great number of organic substances.
§74. Platinum. Pt = 194.9 . Valence two and four.
1. TropeMeB.— Specific gravity at 17.6°, 21.48 (Deville and Debray. C. r., I860,
50, 1038). Melting point, 1775** (Violle, C. r., 1879, 89, 702). Pure platinum is a
tin-white metal, softer than silver, hardened by the presence of other metals,
especially iridium, which it frequently contains. It is surpassed in ductility
and malleability only by An and Ag . Plntinum black is the finely divided
metal, a black powder, obtained by reducing an alkaline solution of the platinous
salt with alcohol (Low, B,, 1S90, 23, 289) ; platinum sponge, a gray spongy mass^
94 PLATINUM. §74, 2
by ignition of the platinum ammonium double chloride; platinized asbestos
(usually 10 per cent Pt), the metal in finely divided form deposited by reduction,
from the salt upon asbestos. These finely divided forms of platinum have great
power of condensation of gases, and by their presence alone bring about a num-
ber of important chemical reactions (catalytic reaction); e. g,^ a current of
hydrogen mixed with air ignites when passed over platinum black, also
hydrogen and chlorine unite. SO, unites with O to form SO,; alcohol is oxi-
dized to acetic acid, formic and oxalic acids to CO3 , As'" to AsV , etc.
2. Occurrence. — Found in nature only in the metallic state, generally alloyed
with palladium, iridium, osmium, rhodium, ruthenium, etc. The Ural Moun-
tains furnish the largest sui>ply of platinum.
3. Preparation. — Usually by the wet method. The finely divided ore is treated
with nitrohydrochloric acid until the platinum is all dissolved. The filtrate is
then treated with lime water to a slightly acid reaction; this removes the
greater part of the Fe , Cu , Ir , Bh , and a portion of the Pd . The filtrate is
now evaporated to dryness, ignited and washed with water and hydrochloric
acid. This gives a commercial platinum which is melted with six times its
weight of lead and the finely divided alloy digested with dilute HNO, , which
dissolves out the Pb , Cu , Pd , and Bh . The black powder which remains is
dissolved in nitrohydrochloric acid, the Pb remaining, removed with H.SO4 ,
and the Pt precipitated with NH4CI . The precipitate contains a little rhodium,
which is removed by gently igniting the mass with potassium and ammonium
di-sulphate, and exhausting with water, which dissolves out the rhodium
sulphate (§105, 7). In the laboratory the platinum residues are boiled with
SOH or K2CO1 and reduced with alcohol. The fine black pow^der is filtered,
washed with water and hydrochloric acid and ignited.
4. Oxides and Hydroxides. — Platinum forms two oxides, PtO and PtO, .
Platinous hydroxide is formed by treating a dilute solution of platinous potas-
sium chloride with NaOH and boiling (Jorgensen, J. pr., 1877, (2), 16, 344).
A black powder easily soluble in HCl or HBr , reduced by formic acid to Tt° ,
gentle heating changes it to the oxide PtO. PUiHnic hydroxide, Pt(OH)t , is
formed by treating a solution of HoPtCla with Na,'COa in excess, evaporating
to dryness, washing with water and then with acetic acid. It is a red-brown
powder, soluble in NaOH , HCl, HNO,, and H.SO,; insoluble in HCaHjO, .
Gentle heating changes it to the oxide PtOa (Topsoe, B„ 1870, 3, 462).
5. Solubilities. — a— Metal. — Platinum is not affected by air or water, at any
temperature; is not sensibly tarnished by hydrosulphuric acid gas or solution;
and is not attacked at any temperature by nitric acid, hydrochloric acid or
sulphuric acid, but dissolves in nitrohydrochloric acid (to platinic chloride)
less readily than gold. h. — Oxides and hydroxides. — See 4. e.— Salts. — Platinum
forms two classes of salts (both haloid and oxy), platinous and platinic. The
oxysalts are not stable. None of the platinous salts are permanently soluble in
pure water. The chloride is soluble in dilute hydrochloric acid and the sul-
phate in dilute sulphuric acid. Platinic chloride, PtCl4 , and bromide, all the
platinicyanides (as PbPt(CN)n), and the platinocyanides of the metals of the
alkalis and alkaline earths (as KiPtiCN)^), are soluble in water. The platinous
and platinic nitrates are soluble in water, but easily decomposed by it, with the
precipitation of basic salts. The larger number of the metaUo-platinic chlorides
or " chloroplatinates " are soluble in water, including those with sodium
[Na.PtClo or (NaCl)..PtCl4], barium, strontium, magnesium, zinc, aluminum,
copper; and those with potassium, and ammonium, are sparingly soluble in
water, and owe their analytical importance as complete precipitates to their
insolubility in alcohol. Of the metallo-platinous chlorides (the "chloroplatinites")
— those with sodium [Na2PtCl4], and barium, are soluble; zinc, potassium and
ammonium, sparingly' soluble; lead and silver, insoluble in water. Platinic
sulphate, Pt(S04)2 , is soluble in water.
6. Beactions. — a. — Platinous chloride, PtCL , is precipitated by KOH as
Pt(OH), , soluble in excess of the reagent to KjPtO, , potassimn platinite, which
solution is reduced by alcohol to " platinum black " (1). Platinic chloride.
PtCl* , a brown-red solid, soluble in alcohol and water, forms with KOH or
IIH4OH , not too dilute, a yellow crystalline precipitate of an alkali (K or NH*)
platinum chloride, c. g., K,PtClc , sparingly soluble in water, soluble in excess
§74, 7, 1. PLATINUM. 95
of the alkalis and reprecipitated by hydrochloric acid. K,COa and (NH4),C0,
give the same precipitate, insoluble in excess of the reag'ent. A more complete
precipitation of the K or NH4 is obtained by the use of the chlorides. The
sodium platinum chloride, NatPtCl. , is very soluble in water and is not formed
by precipitation with sodium salts, h. — Oxalic acid does not reduce platinum
salts (distinction from gold). A solution of chloral hydrate precipitates pla-
tinum from its solutions. Platinous and platinic salts form with cyanides a
great number of double salts, c. — See 5r. d. — Hypophosphorous acid reduces
platinum salts to metallic platinum. Phosphates do not precipitate platinum
salts.
e. Eydrosnlphuric acid precipitates solutions of the platinous salts as
the black sulphide, PtS , insoluble in acids, sparingly soluble in water and
in alkali sulphides; platinic salts are precipitated as platinic sulphide,
PtS,, black; slowly soluble in alkali sulphides (Ribau, C. r., 1877, 86, 283),
insoluble in acids except nitrohydrochloric. Sulphur dioxide decolors a
solution of platinum chloride giving a compound which does not respond
to the usual reagents for platinum and requires long boiling with HCl for
the removal of the SO2 (Birnbaum, /!., 1871, 169, 116).
f. The chlorides of potassium and ammonium are estimated quantita-
tively by precipitation from their concentrated solutions with a solution
of platinic chloride. Fotassium iodide colors a solution of platinum
chloride brown-red and precipitates the black platinic iodide, Ptl4 , excess
of the B3 forming Es^^n > brown, sparingly soluble {5c). g. Stannous
chloride does not precipitate the platinum from platinic chloride (distinc-
tion from gold), but reduces it to platinous chloride.
h. Ferrous sulphate solution on boiling with a platinum chloride solu-
tion precipitates the platinum as the metal, the presence of acids hinders
the reduction.
7. Ignition. — All platinum compounds upon ignition are reduced to the
metal. Owing to the high point of fusibility of the metal and to the
difficulty with which it is attacked by most chemicals, platinum has
an extended use in the chemical laboratory for evaporating dishes, cruci-
bles, foil, wire, etc. Ix the use of platinum apparatus without
rXNECESSARY INJURY IT SHOULD BE REMEMBERED:
(1) That free chlorine and bromine attack platinum at ordinary tem-
peratures (forming platinic chloride, bromide); and free sulphur, phos-
phorus, arsenic, selenium, and iodine, attack ignited platinum (forming
platinous sulphide, platinic phosphide, platinum-arsenic alloy, platinic
selenide, iodide). Hence, the fusion of sulphides, sulphates, and phos-
phates, with reducing agents, is detrimental or fatal to platinum crucibles.
The ignition of organic substances containing phosphates acts as free
phosphorus, in a slight degree.
The heating of ferric chloride, and the fusion of bromides, and iodides,
act to some extent on platinum.
96 PLATINUM. §74, 7, i2,
(2) The alJcali hydroxides (not tjieir carbonates) and the alkaline earths,
especially baryta and lithia, with ignited platinum in the air, gradually
corrode platinum (by formation of platinites: 2Pt + 2BaO -f- Oj =
2BaFtO, . Silver crucibles are recommended for fusion with alkali
hydroxides.
(3) All metals which may he reduced in ths fusion — especially compounds
of lead, bismuth, tin, and other metals easily reduced and melted — and all
metallic compounds with reducing agents (including even alkalis and earths)
form fusible alloys with ignited platinum. Mercury, lead, bismuth, tin,
antimony, zinc, etc., are liable to be rapidly reduced, and immediately to
melt away platium in contact with them.
{Jf) Silica with charcoal (by formation of silicide of platinum) corrodes
ignited platinum, though very slowly. Therefore, platinum crucibles
should not be supported on charcoal in the furnace, but in a bed of mag-
nesia, in an outer crucible of clay. Over the flame, the best support is the
triangle of platinum wire.
(5) The tarnish of the gas-flame increases far more rapidly upon the
already tarnished surface of platinum — going on to corrosion and crack-
ing. The surface should be kept polished — preferably by gentle rubbing
with moist sea-sand (the grains of which are perfectly rounded, and do not
scratch the metal). Platinum surfaces are also cleansed by fusing horax
upon them, and by digestion with nitric acid.
8. Detection. — Platinum is identified by the appearance of the reduced
metal; by its insolubility in HCl or HNO3 and solubility in HNO3 -f HCl ;
and by its formation of precipitates with ammonium and potassium
chlorides. It is separated from gold by boiling with oxalic acid and am-
monium oxalate, which precipitate the gold, leaving the platinum in solu-
tion. The filtrate from the gold should be evaporated, ignited, and the
residue examined and after proving insolubility in HCl or HNO3 , dissolved
in nitroliydrochloric acid and the presence of platinum confirmed witli
ITH4GI . If the gold and platinum have been precipitated in the second
group with H.^S and dissolved with (1^4)28, they may be separated from
As , Sb , and Sn by dissolving the reprecipitated sulphides in HCl + ECIO3 ,
evaporating to remove the chlorine and boiling after adding EOH in ex-
cess, with chloral hydrate, which precipitates the Au and Pt , leaving the
As , Sb , and Sn in solution. The An and Pt may then be dissolved in
HNO3 + HCl and separated as directed above. FeSO^ may be use to pre-
cipitate Au and Pt , separating them from As , Sb , and Sn .
9. Estimation. — Platinum is invariably weighed in the metallic state. It is
brought to this condition: (i) By simple ignition; (2) by precipitation as
(NHO,PtCl, , KaPtCla , or PtS, and ignition; (3) by reduction, using Zn , Mg,
or FeSOf .
10. Oxidation. — Solutions of platinum are reduced to the metallic state by the
§76, 6c. MOLYBDENUM, 97
following" metals: Pb , Ag , Hg, Sn (Sn" to Pt" only), Bl, Cu , Cd , Zn .
JFe , Pe'' , Co , and Ni . Very many organic substances reduce platinum
^compounds to the metallic state.
§76. Molybdennm. T/Lo = 96.0 . Valence two, three, four and six.
1. Properties.— Sp«?i/kr ffravity, 8.56 (Loughlien, Am. S., 18G8, (2), 45, 131).
Pure molybdenum appears not to have been melted; when heated to a very
high heat in a graphite crucible it takes up carbon and melts. It is a silver-
white, hard, brittle metal, not oxidized in the air or water at ordinary tem-
peratures. Upon heating in the air it becomes brown, then blue, and finally
burns to the white MoO, . Heated to a red heat in contact with steam, h
forms first a blue oxide, then MoO, .
2. Occurrence. — Not found native, but occurs chieflj' as molybdenite, M0S2;
as an oxide in molybdenum ochre, MoO.; and as wulfenite, PbMoO^ .
3. Preparation. — (i) By heating the oxide, sulphide or chloride in a current
of oxygen free hydrogen (von der Pfordten, B., 1884, 17, 732; Rogers and
>ntchell, J, Am, Soc, 1900, 22, 350); (2) by heating with C and Na,CO«; (3) by
heating KoO, with KCN (Loughlien, I.e.).
4. Oxides and Hydroxides. — Molybdous hydroxide, HoO.xHjO , is formed when
molybdous chloride or nitrate is precipitated with alkali hydroxides or carbon-
ates^ dark brown becoming blue in the air by oxidation. Mo (OH), , black,
turning red-brown by oxidation in the air, is formed by treating M0CI3 with
JSOH: also by electrolysis of ammonium molybdate (Smith, B,, 1880, 13, 751).
By heating the hydroxide in a vacuum MO2O, is obtained as a black mass,
insoluble in acids. MoO^ , a dark bluish mass, insoluble in KOH or HCl , is
formed by igniting a mixture of ammonium molybdate, potassium carbonate
and boric acid, and exhausting the fused mass with water (Muthmann, A., 1887,
-238, 114). Molybdic anhydride (acid), MoO, , white, occurs in nature; it is
obtained by the ignition of the lower oxidized compounds in the air or in the
presence of oxidizing agents.
5. Solubilities. — Molybdenum is readily soluble in nitric acid with oxidation
to MoO| , evolving NO; in hot concentrated sulphuric acid, evolving SO, . The
various lower oxides of molybdeniim are soluble in acids forming corresponding
salts, not very stable, oxidizing on exposure, to molybdic acid and molybdates;
on the other hand, reducing agents reduce molybdates to the lower forms of
molybdenum salts, nearly all of which are colored brown to reddish brown or
violet. The salts of molybdenum are nearly all soluble in water. Molybdic
anhydride, MoO, , white, is sparingly soluble in water and possesses basic
properties towards stronger acids, dissolving in them to form salts. The
chlorides and the sulphates are soluble in water (Schulz-Sellack, 5., 1871, 4, 14);
the nitrates in dilute nitric acid. The anhydride MoO, combines with the
alkalis to form molybdates, soluble in water. Molybdates of the other metals
are insoluble in water. Solutions of the alkali molybdates are decomposed by
acids forming, MoO, , which dissolves in excess of the acids.
G. Heactions. — a. — The dyad, triad and tetrad molybdenum salts are precipi-
tated by the alkali hydroxides and carbonates, forming the corresponding
hydroxides, insoluble in excess of the precipitant. These hydroxides oxidize
in the air to a blue molybdenum molybdate. 6. — A solution of a molybdate
acidulated with hydrochloric acid gives no red color with KCNS (distinction
from Pe'"); but if Zn be added, reduction to a lower oxide of molybdenum
takes place and an intense red color is produced. Phosphoric acid does not
destroy the color (difference from ferric thiocyanate). Upon shaking with
ether the sulphocyanate is dissolved in the ether, transferring the red color
to the ether layer. In molybdic acid solutions, acidulated with hydrochloric
acid, potassium ferrocyanide gives a reddish brown precipitate. An alkaline
solution of molybdates is colored a deep red to brown by a solution of tannic
acid. c. — See 5.
98 MOLYBDENUM, §75, 6d.
d. — Tribasic phosphoric acid and its salts precipitate, from strong nitric
acid solutions of ammonium molybdate,* somewhat slowly, more rapidly
on warming, ammonium phospho-molyhdatey yellow, of variable composition,
soluble in ammonium hydroxide and other alkalis, sparingly soluble in
excess of the phosphate. Hydrochloric acid may be used instead of nitric.
The sodium phospho-molyhdate is soluble in water, and precipitates am-
monium from its salts; also, it precipitates the alkaloids — for which reac-
tion it has some importance as a reagent. f Arsenic acid and arsenates
give the same reaction ; ammonium arseno-molyhdate being formed (g).
e. — Neutral or alkaline solutions of molybdates are colored yellow to
brown by hydrosulphuric acid but are not precipitated. From the acid
solutions a small amount of the hydrogen sulphide gives no precipitate
but colors the solution blue; with more hydrosulphuric acid the brovvm or
red-brown precipitate, 1/LoS.^ , molyhdenum trisulphidey is obtained after
some time. The precipitate is soluble in ammonium sulphide, better when
hot and not too concentrated, as ammonium thiomolybdate, (NH^)2MoS^ ,
from wdiich acids precipitate the trisulphide (Berzelius, Pogg.y 1826, 7,
429), soluble in nitric acid, insoluble in boiling solution of oxalic acid
(separation from stannic sulphide).
If Na,SaO, he added to a solution of nmmoniuin molj'bdate, slightly acid,
a blue i)recipitate and bine-colored solution is obtained. If the solution be
more strongly acid, a red brown precipitate is obtained. An acid solution of a
molybdate treated with hypophosphorous and sulphurous acids gives an in-
tense bluish green precipitate or color, depending upon the amount of molyb-
denum present.
f. — Halogen compounds not important in analysis of molybdenum.
17.— Arsenic acid and arsenates form, with a nitric acid solution of ammonium
molybdate, a yellow precipitate of ammonium ar.snw-moIybdat€^ in appearance
and reactions not to be distinguished from the ammonium phospho-molybdate:
except the precipitation does not take place until the solutions are slightly
warmed, while with phosphates the precipitation begins even in the cold.
Stannous salts give with (NH4);Mo04 a blue solution of the lower oxides of
molybdenum (a delicate test for Sn") (Longstaff, C. A"., 1899, 70, 282).
*.— Solutions of the alkali molybdates are soluble in water and precipitate
solutions of nearly all other metallic salts, forming molybdates of the corre-
sponding metals, insoluble in water, e.tj., K2M0O4 -f PbCNO,), = PbMoO* -f
2KN0, .
• The reagent ammonium molylxiate, iNH*^, B[o04. is prepared by dissolving molybdic acid,
MoO, (lOO^rams), in ammonium hydn>xide (250 co. sp. jrr. 0.90 with 250 cc. water) cooHng, and
t^lowly pouring this sohition into well cooled fairly concentrated nitric acid (750 cc sp. gr. 1.42
with 750 cc. water) with constant stirring.
t Sodium Phwi)?if>-m«W?/Mfltr— Sonnenschcin's reagent for acid solutionsof alkaloids— is pre-
pared as follows : The yellow preinpitate formed on mixing acid solutions of ammonium molyb-
date and sodium phosphate^— tho ammonium phospho-molybdate— is well washed, suspended in
water, and heate<l with s^niium carl>onate until completely dissolve*!. The solution is evapor-
ated to dryness, and the ivsiduo gently ignited till all ammonia is expelled, sodium being sub-
stituted for ammonium. I f blackening occurs, from reiluctionof molybdenum, the residue is
moisteneti with nitric acid, and heated again. It is then dissolved with water and nitric acid
to strong acidulation : the solution Inking made ten parts to one part of residue. It must be
kept from contact with vajwr of ammonia, both during the preparation and when preserved
for use.
§76, 10. MOLYBDENUM. 99
7. Ignition. — With microcosmic salt, in the outer blow-pipe flame, all com-
pounds of molybdenum give a bead which is greenish while hot, and colorless
on cooling*; in the inner flame, a clear green bead. With borax, in the outer
flame, a ^ad, yellow while hot, and colorless on cooling; in the inner flame, a
brown bead, opaque if strongly saturated (molybdous oxide). On charcoal,
in the outer flame, molybdic anhydride is vaporized as a white incrustation; in
the inner flame (better with sodium carbonate), metallic molybdenum is
obtained as a gray powder, separated from the mass by lixiviation. Dry molyb-
dates, heated on platinum foil with concentrated sulphuric acid to vaporiza-
tion of the latter form, on cooling in the air, a blue mass.
8. Detection. — In the ordinary process of analysis, molybdenum appears
in Division A (tin group) of the second group with As , Sb , Sn , Au , and
Pt . The solution remaining in the Marsh apparatus is decanted from
the residue (Sn , Sb ^ Au , Pt and excess of Zn) and heated with concen-
trated HNO3 , the molybdenum is oxidized to molybdic acid. This solution,
evaporated to dryness, dissolved in ammonium hydroxide and poured into
moderately concentrated HCl forms a solution of ammonium molybdate
which may be identified by the many precipitation and reduction tests
(G by Cy dy By l, ctc, 7, aud 9). If the molybdenum be present as a molybdate
it may be precipitated from its nitric acid solution by NasHFO^ , washed,
dissolved in ammonium hydroxide, the phosphate removed by magnesia
mixture (§189, 6a), and the filtrate evaporated to crystallization (Maschke^
Z.y 1873, 12, 380). The crystals may be tested by the various reduction
tests for molybdenum.
9. Estixiiation. — (1) Molybdic anhydride and ammonium molybdate may be
reduced to the dioxide by heating in a current of hydrog-en pas. The heat
must not be permitted to rise above dull redness. Or the temperature may
rise to a white heat, which reduces it to the metallic state, in which form it is
weigrhed. (2) Lead acetate is added to the alkali molybdate, the precipitate
washed in hot water, and after i|?nition weighed as FbMoO^ . (S) Volumet-
rically. The molybdic acid is treated with zinc and HCl , which converts it into
MoCla . This is converted into molybdic acid again by standard solution of
potassium permanganate.
10. Oxidation. — Xteducin^ agents convert molybdic acid either into the hlue
intermediate oxides, or, by further deoxidation, into the black molybdous oxide,
MoO . In the (hydrochloric) nrcid solutions of molybdic acid, the blue or black
oxide formed by reduction, will be held in solution with a bhie or brown color.
Nitric acidulation is, of course, inconnpatible with the reduction. Certain
reducing agents act as follows:
Ferrous salts (in the hydrochloric acid solution) give the hlue oxide solution.
Cane sug^r, in the feebly acid boiling sohition, forms the hlue color — seen
better after dilution: a delicate test. Stannous chloride forms first the hlue,
then the frroirn, or the fireenifih brown to black-brown, solution of both the
intermediate oxide and the molybdous oxide. Zinc, with HCl or H3SO4 , gives
the ft/t/e, then green, then hroirn color, by progressive reduction. Formic and
oxalic acids do not react. A solution of 1 milligram of sodium (or ammonium)
molybdate in 1 cc. of concentrated sulphuric acid (about 1 part to 1840 parts) is
in use as Froehde's Reagent for alkaloids. The molybdenum in this solution,
which must be freshly prepared for use each time, is reduced by very many
organic substances: and with a large number of alkaloids, it gives distinctive
colors, blue, red, brown and yellow.
^l^V^^l^
100 BISMUTH. §76, 1.
The Copper Group (Second Group, Division B).
Uercury (Herctiricuin), Lead, Bismuth, Copper, Gadminm (Buthenium,
Rhodium, Palladium, Osmium).
§76. Bismuth, Bi = 208.1 . Valence three and five.
1. Properties.—iSfpecf/lo gravity, 9.7474 (Classen, B., 1890, 23, 938); melting point,
269.22 (Callendar and Griffiths, 0. iV., 1891, 63, 2); it vaporizes at 1700° and the
density of the vapor shows that the molecule Bi has begun to dissociate (Biltz
and V. Meyer, B., 1889, 22, 725). It is a hard, brittle, reddish-white, lustrous
metal; forming beautiful rhombohaedral crystals when a partially cooled mass
is broken into and the still molten mass decanted. Alloys of bismuth with
other metals give compounds of remarkably low melting points, e. g,, an alloy
of: BI two, Sn one, and Pb one part by weight melts at 93.7°; and an alloy ot:
Bi fifteen, Pb eight, Sn four, and Cd three parts by weight melts at 6S°
" Wood's Metal."
2. Occurrence. — It is a comparatively rare metal, not very widely distributed,
usually found native. It is found in greatest quantities in Saxony; also found
in Bohemia, France. England and South America. As mineralogical varieties
it occurs as bismuth ochre (Bi^Og), bismuthite (4Bi,0,.3C02.4H30), bismuth
glance (Bl^S,), etc.
3. Preparation. — The rock containing bismuth, usually with large amounts
of cobalt, etc., is roasted to remove sulphur and arsenic, which is nearly
always present. The mass is then fused with charcoal. The molten bismuth
settles to the bottom below the layer of cobalt. The cobalt becomes solid
while the bismuth is still molten, and the two are separated mechanically.
The metal is further purified by melting with XKO, or KCN .
4. Oxides. — Bismuth trioxide, BiaO, , is formed by heating the metal in the
presence of air, or by igniting the hydroxide; it is a pale citron-yellow powder.
The hydroxide, Bi(0H)8 , white, is formed by precipitating a solution of a salt
of bismuth with an alkali hydroxide. If bismuth chloride is used the hydroxide
formed always contains some oxychloride, BiOCI (Strohmeyer, Pogg., 1832, 26,
649). The meta hydroxide, BiO(OH) , is formed upon drying the orthohydroxide
at 100° (Arppe, Pogg., 1845, 64, 237). Bismuth pentoxide, BljO. , is formed by
igniting Bi(OH), with excess of KOH or NaOH in presence of the air, and
washing the cooled mass repeatedly with cold dilute nitric acid (Strohmeyer,
1. r.); or by treating Bi(OH)s with three per cent H.Oj in strong alkaline solu-
tion (Hasebrock, /?., 1887, 20, 213). It is a heavy dark brown powder. At 150 "•
it gives off O, and at the temperature of boiling mercury becomes Bi^O, . It
is decomposed in the cold by HCl with evolution of chlorine. Bimnuthic acid.
HBiO, , or more probably Bi.Os.H^O , is formed upon conducting a rapid
current of chlorine into Bi(OH)a suspended in concentrated KOH solution.
It is a beautiful scarlet red powder w^hich at 120° gives off its w^ater, becoming
BijO, (Muir, J. C, 1876, 29, 144; Muir and Carnegie, J. C, 1887, 51, 86). It is
doubtful if any alkali salt of bismuthic acid exists, although mixtures of KBiO,
and HBiO, are claimed by Hoffmann (.4., 1884. 223. 110), and Andr^ (C. r., 1891,
113, 860). The so-called bismuth t^troxide, "BUO^ , is probably a mixture of the
trioxide and pentoxide.
5. Solubilities. — n.—.l/f/cr?.— Metallic bismuth is insoluble in hydrochloric
acid *: soluble in warm concentrated sulphuric acid with evolution of sulphur
dioxide: readily soluble in nitric acid and in nitrohydrochloric acid. It burns
in chlorine with production of light: it combines vvith bromine, but more slowly
than antimony; it combines readily upon fusing together with I , S , Se , Te .
As , and Sb , besides the many metals with which it combines to form com-
* A trace of bismuth can always be found In solution when the metal is boiled with hydro-
chloric acid, but no more than when the metal has been boiled with pure water (Dltte and
Hetzner, A. Ch., 1896, (6), 99, 889).
§76, ea. BISMUTH, 101
mercial alloys (1). The halogen derivatives of pentad bismuth are not known
(Muir, J, C, 1876, 29, 144). ly— Oxides and hydrojr ides. —Bismuth oxide, BijO, ,
and the hydroxides, Bi(0H)3 and BiO(OH), are soluble in hydrochloric, nitric
and sulphuric acids; insoluble in water and the alkali hydroxides or carbonates.
The presence of glycerol prevents the precipitation of bismuth hydroxides
from solutions of its salts by the alkalis.* Bismuth pentoxide, BizOg , is solu-
ble in HCl , HBr , and HI with evolution of the corresponding halogen and
formation of the triad salt^ Nitric and sulphuric acids in the cold have but
little or no action; when hot the triad bismuth salt is formed with evolution
of oxygen.
c. — Salts. — Most of the salts of bismuth are insoluble in water. The
chloride, bromide, iodide, nitrate, and sulphate are soluble in water acidu-
lated with their respective acid, or with other acids forming " soluble ''
bismuth salts. Pure water decomposes the most of the solutions of bis-
muth salts forming corresponding oxy-salts (§70, od footnote).
The chloride, bromide and sulphate are deliquescent.
d. — ^Watcr. — A solution of bismuth chloride in water acidulated with
hydrochloric acid is precipitated on further dilution with water, bismuth
oxy-chloride, BiOCl being formed; e. g., BiClg -f B.fi = BiOCl -f 2HC1 ,
insoluble in tari^aric acid (distinction from antimony, §70, r)d). The hydro-
chloric acid set free serves to hold a portion of the bismuth in solution.
The presence of acetic, citric, and other organic acids prevents the pre-
cipitation of solutions of bismuth salts upon further dilution with water.
The washing of the precipitated oxy-salt with pure water removes more of
the acid forming a salt still more basic.
Bi(NO,), + H,0 = BiONO, + 2HNO3
12BiONO. + H2O = 6Bi,0.,5Na05 + 2HN0.
This is prevented by the presence of one part ammonium nitrate to five
hundred parts water (Lowe, J. pr., 1858, 74, 341).
Bismuth nitrate crj'stallizes with ten molecules of water, Bi(N03)3.
IOH2O . It is decomposed by a small amount of water forming the basic
nitrate, BiONOs ; this is soluble in dilute nitric acid, when further dilution
with water to any extent is possible without precipitation of the basic
salt, but a drop of hydrochloric acid or a chloride causes a precipitate of
the oxychloride in the diluted solution. The bromide is readily decom-
posed by water to BiOBr ; the iodide is stable to cold water, but is decom-
posed by hot water to Bid (Schneider, A. Ch, 1857 (3), 60, 488); the
normal sulphate very readily absorbs water to form Bi2(S04)3.3HoO , whicli
is decomposed by more water to Bi203.S03 .
6. Beaotions. a. — The alkali hydroxides precipitate from solutions of
bismuth salts bismuth hydroxide, Bi(0H)3 , white; insoluble in excess of
the fixed alkalis (distinction from Sb and Sn), insoluble in ammonium
• Lowe (C. N^ 1882, 45, 206) disaolves the hydroxides of copper and bismuth in grlycerol, adds
gluoofle and firently warms. The copper is completely precipitated and separated from the bis-
muth. Upon bolUnff the flltrato for some time the bismuth Is completely precipitated as the
metaL
102 BISMUTH, §76, 6b.
hydroxide (distinction from Cu and Cd). The hydroxide is converted by
boiling into the oxide, BijOj , yellowish white. The precipitation is pre-
vented by the presence of tartaric acid, citric acid, glycerol, and certain
other organic substances (Kohler, J, C, 188G, 60, 428).
The alkali carbonates precipitate basic bismuth carbonate, Bi^Og.CO, , white,
insoluble in excess of the reagent. Freshly precipitated barium carbonate
forms the same precipitate without heating.
6.— Oxalic acid and soluble oxalates precipitate bismuth oxalate^ BUiCiO^), ,
white, soluble in moderately dilute acids. Potassium cyanide forms a white
crystalline precipitate insoluble in excess of the reagent but soluble in nitric
or hydrochloric acid. Potassium ferrocyanide forms a yellowish white pre-
cipitate, potassium ferricyanide a brownish yellow, both soluble in hydrochloric
acid.
0. — The action of nitric acid upon bismuth and its salts is fully explained
under (5). d. — Metallic bismuth is precipitated when bismuth salts are warmed
with hypophosphorous acid (separation from Zn and Cd) (Muthraann and
Mawron, Z., 1874, 13, 209). From solutions of bismuth nitrate (5d) phosphoric
acid and soluble phosphates precipitate bismuth phosphate, BiPO^ , white,
readily soluble in HCl; from solutions of the chloride, diluted as much as pos-
sible without precipitation, phosphoric acid gives no precipitate, but the pre-
cipitate of the phosphate (soluble in HCl) is obtained with soluble phosphates.
e, — ^HydrosulphtLric acid and sulphides precipitate hismuth sulphide,
BijSg , black, insoluble in dilute acids and in alkali hydroxides; insoluble in
alkali sulphides (distinction from the metals of the tin group) and in alkali
cyanides (distinction from copper). It is soluble by moderately concen-
trated nitric acid (distinction from mercury), the sulphur mostly remain-
ing free.
Sodium thiosulphate when warmed with solutions of bismuth salts precipitates
bismuth sulphide. Sulphuric acid does not precipitate solutions of bismuth
chloride or nitrate. Potassium sulphate gives a precipitate with solutions of
both, that with the chloride being apparently caused by the dilution of the
solution.
f, — Hydrochloric acid and soluble chlorides form a precipitate of bis-
muth oxy-chloride, BiOCl , in solutions of bismuth nitrate not containing
too much free nitric acid. This makes it possible for bismuth to be precipi-
tated with the silver group salts (§63, 6h). The precipitate is readily
dissolved on addition of more hydrochloric or nitric acid (distinction from
the silver group chlorides).
Hydrobromic acid and soluble bromides do not precipitate solutions of bis-
muth chloride, but do precipitate solutions of the nitrate, forming the oxy-
broraide, BiOBr , white. The presence of potassium bromide prevents the pre-
cipitation of a bismuth chloride solution by water and also dissolves the oxy-
chloride which has been precipitated by the addition of water.
Hydriodic acid and soluble iodides precipitate from solutions of bismuth
salts, unless strongly acid, bismuth iodide, black or brownish gray crystals,
quite readily soluble in excess of the reagent * or in strong HCl without warm-
♦ Bismuth iodide dissolves in solution of potassium iodide witb an intense yellow color, deli-
cate to one-millionth (Stone J. Soc. Chem, I«4., 1887, 6. 416). The potassium iodide solution of
bismuth iodide is used as Dragendorff's reagent to detect the presence of an alkaloid. Leger
(!??„ 1888, 50, 91) uses cinchonine and potassium iodide to prove the presence of bismuth. Del-
Jeate to one-nve hundred thousandth. Other metals must be removed.
§76, 0. BISMUTH, 103
ing. It is reprecipitated on diluting" the solution with water. Bismuth iodide
is scarcely at all decomposed by washing with cold water, but on boiling with
-water it is decomposed into bismuth oxy-iodide, BiOI , red, insoluble in KI ,
soluble in HCl , and in HI (Gott and Muir, J, C„ 1888, 53, 137).
Chloric acid dissolves bismuth hydroxide, but the compound decomposes upon
evaporation (Wachter, A., 1844, 52, 233). Potassium bromate and iodate both
precipitate solutions of bismuth nitrate. The iodate formed is scarcely soluble,
the bromate easily soluble in HNOg .
g. — Potassium or sodium stannite hot, when added in excess to bismuth
solutions^ cause a black precipitate, from reduction to metallic bismuth, a
very delicate reaction.* The stannite is made, when wanted, by adding
to a stannous chloride solution, in a test-tube, enough sodium or potas-
sium hydroxide to redissolve the precipitate at first formed: 2BiCl3 +
aZ^SnOs + 6K0H = 2Bi + 6KC1 + SKgSnOa + SHjO (Vanino and Treu-
bert, S., 1898, 31, 1113).
h. — Solutions of bismuth salts, nearly neutral, poured into a hot solution of
potassium bichromate precipitates the orange red chromate, (B10),Cr207; but
if poured into, a cold solution of the neutral chromate a citron-yellow precipi-
tate, .3BiaO,.2CrO, , is formed. These precipitates are soluble in moderately
concentrated acids, insoluble in fixed alkalis (distinction from Pb). The pre-
cipitate with XsCraOr is used in the quantitative determination of bismuth (9).
7. Ignition. — On charcoal, with sodium carbonate, before the blow-pipe, bis-
muth is readily reduced from all its compounds. The globule is easily fusible,
brittle (distinction from lead), and gradually oxidizable under the fiame, form-
ing an incrustation (BijO,), orange-yellow while hot, lemon-yellow when cold,
the edges bluish-white when cold. The incrustation disappears, or is driven
by the reducing flame, without giving color to the outer flame. With borax
or microcosmic salt, bismuth gives beads, faintly yellowish when hot, colorless
when cold.
A mixture of equal parts cuprous iodide and sulphur forms an excellent
reagent for the detection of bismuth in minerals by the iise of the blow-pipe.
The reagent mixed with the unknown is fused on charcoal or on a piece of
aluminum sheet. A red sublimate indicates bismuth. Mercurj' gives a mix-
ture of red and yellow sublimates (Hutchings, C. ^^, 1877, 36, 249).
Bismuth chloride may be sublimed at the temperature of boiling sulphur;
recommended as a separation from lead (Remniler, 5., 1891, 24, 3554).
8. Detection. — Bismuth is precipitated from its solutions by H2S form-
ing BijSs . By its insolubility in (JiflJoS^and solubility in hot dilute
HHO3 it is separated with Pb, Cn, and Cd from the remaining metals
of the tin and copper group. Dilute HgSO^ removes the lead and
ITH4OH precipitates the bismuth as Bi(0H)3 , leaving the Cu and Cd in
solution. The presence of the bismuth is confirmed by the action of a
hot solution of KjSnOj on the white precipitate of Bi(0H)3 , giving metallic
bismuth (6g) or by dissolving the Bi(0H)3 in HCl and its precipitation as
BiOCl upon dilution with water (od),
9. Estimation. — (1) As metallic bismuth formed by fusion with potassium
cyanide. (2) As BijO, formed by ignition of bismuth salts of organic acids, or
of the salts of volatile inorganic oxy acids. (3) By precipitation by H,S , and
♦ For a modification of this test seo MuIr {J, C, 1877, 32, 45).
104 COPPER, §76,10.
after drying at 100**. weighingr as BlaS, . (If) By precipitation by KtCr,OT , and
after drying at 120°, weighing as (B10),Cr20T . (o) Volumetrically. By precipi-
tation with KjCrzOr . Dissolve the chromate in dilute acid, transfer to an
azotometer and reduce the chromate with hydrogen peroxide (Baumann, Z.
angetCj 1891, 331). (6*) By precipitation as a phosphate with standard sodium
phosphate; dilution to definite volume and determination of the excess of
phosphate in an aliquot part with uranium acetate (Muir, J. C, 1877, 32, 674).
10. Oxidation. — Metallic bismuth reduces salts of Hg, Ag, Pt, and
An to the metallic state. Bismuth is precipitated as free metal from its
solutions by Pb , Sn , Cu , Cd , Fc , Al , Zn , Mg , and HH^POa {6d), All
salts of bismuth are oxidized to BiaOg by CI or H2O2 in strong alkaline
mixture (Hasebrock, 5., 1887, 20, 213; Schiif, A. Ch., 1861 (3), 63, 474).
All compounds of bismuth arie reduced to the metal by potassium stannite
EjSnOs (Gg), Bismuth chloride or bromide heated in a current of hydro-
gen is partially reduced to the free metal (Muir, J, C, 1876, 29, 144).
It is precipitated as free metal upon warming in alkaline mixture with
grape sugar (56).
§77. Copper (Cuprum) Cn = 63.6 . Valence one and two.
1. Properties.— ;Sfperi/f<? gravity, electrolytic, 8.914; melted, 8.921; natural crys-
tals, 8.94; rolled and hammered sheet, 8.952 to 8.958 (Marchand and Scheerer,
J. pr,, 1866, 97, 193). Melting point, 1080.5 (Heycock and Neville, J. C, 1895, 67,
190). A red metal, but thin sheets transmit a gnreenish-blue light, and it also
shows the same grreeuish-blue tint when in a molten condition. Of the metals
in ordinary use, only gold and silver exceed it in malleability. In ductility it
is inferior to iron and cannot be so readily drawn into exceedingly fine wire.
Although it ranks next to iron in tenacity, its wire bears about half the weight
which an iron wire of the same size would support. As a conductor of heat it
is surpassed only by gold. Next to silver it is the best conductor of electricity.
Dry air has no action upon it; in moist air it becomes coated with a film of
oxide which protects it from further action of air or of water. It forms a
number of very important alloys with other metals; bronze (copper and tin)„
brass (copper and zinc with sometimes small amounts of lead or tin), German
silver (copper, nickel and zinc).
2. Occurrence. — Copper is found native in various parts of the world, and
especially in the region of Lake Superior. It is found chiefly as sulphides in
enormous quantities in Montana, Colorado, Chili and Spain; as a carbonate in
Arizona. It is very widely distributed and occurs in various other forms.
Copper pyrites is CuPeSj; copper glance, Cu,S; g^reen malachite, Cu3(OH)2CO,;
blue malachite, Cu,(OH)a(CO,),; red copper ore, CUjO: and tenorite, CuO .
3. Preparation. — For the details of the various methods of copper-smelting
and refining, the works on metallurgy should be consulted. In the laboratory
pure copper may be produced (/) by electrolysis; (2) reduction by ignition in
, hydrogen gas; (S) reduction of the oxide by ignition with carbon, carbon
monoxide, illuminating gas, or other forms of carbon; (4) reduction of the
oxide by K or Na at a temperature a little above the melting point of these
motals; (.7) reduction by fusion with potassium cyanide: CuO -j- KCK = Cu +
KCNO . For its reduction in the wet waj% see 10.
4. Oxides and Hydroxides. — Cuprous oxide (CUjO), red, is found native: it is
prenared: (/) bv reducing CuO by means of grape-sugar in alkaline mixture;
(2J by igniting duO with metallic copper: (S) by treating an ammoniacal cupric
nolutlon with metallic copper; then adding KOH and drying. Cuprous hydrox-
<fff, CuOH , brownish yellow, is formed by precipitating cuprous salts with
KOH or NaOH . Cupric oxide, CuO . black, is formed by igniting the hydroxide.
§77, 5c. COPPER. 105
carbonate, sulphate, nitrate and some other cnpric salts in the air: or by
heating the metal in a current of air. Cupric hydivxide, Cu(OH), , is formed
by precipitating cupric salts with KOH or NaOH . It is stated by Rose {Poyg.^
1863, 120, 1) that tetracupric monoxide, (CU4O , is formed by treating a cupric
salt with KOH and a quantity of K^SnO, insufficient to reduce it to the metallic
state. A perojriJe of copper; CuO. , is supposed to be formed by treating
Cn(OH), with H,Oj at 0° (Kriiss, B„ 1S84, 17, 2593).
5. Solubilities. — a. — J/e/«/.— Copper does not readily dissolve in acids with
evolution of hydrogen: it dissolves most readily in nitric acid chiefly with
evolution of nitric oxide* 3Cu + 8HN0, = 3Cu(N0a), + 4H.,0 + 2N0 (Freer
and Higley, Am., 1899, 21, 377); also in hot concentrated sulphuric acid, with
evolution of sulphurous anhydride: Cu -f 2H3SO4 = CUSO4 -\- 2H,0 + SO, . If
dry hydrochloric acid gas be passed over heated copper, CuCl is formed with
evolution of hydrogen (Weltzien, A. Ch., 1865, (4), 6, 487). A saturated solution
of hydrochloric acid at 15** dissolves copper as CuCl with evolution of hydrogen.
The action is very rapid if the copper be first immersed in a platinum chloride
solution. Heat favors the reaction and the presence of IOH2O to one HCl pre-
vents the action (Engel, C. r., 1895, 121, 528). Ilydrobromic acid concentrated
acts slowly in the cold and rapidly when warmed, forming CuBr, , with evolu-
tion of hydrogen. Cold hydriodic acid, in absence of iodine, is without action
(Mensel, B., 1870, 3, 123). Ammonium sulphide, (NHJ^S, colorless, acts upon
copper turnings with evolution of hydrogen, forming CUjS (Heumann, J. C,
1873, 26, 1105).
&. — Oxides. — Cuprous oxide and hydroxide are insoluble in water, soluble
in hydrochloric acid with formation of cuprous chloride, white, unstable,
readily oxidized by the air to colored cupric salts. Cnpric oxide, black,
and hydroxide, blue, are insoluble in water, soluble in dilute acids; in a
mixture of equal parts glycerine and sodium hydroxide, sp. gr. 1.20 (sepa-
ration from Cd) (Donath, J. C, 1879, 36, 178), in a mixture of tartrates
and fixed alkalis (but precipitated as CUjO l)y heating with glucose) (sepa-
ration from Cd and Zn) (Warren, C. iV., 1891, 63, 193); insoluble in
ammonium hydroxide in absence of ammonium salts (Maumene, J. C,
1882, 42, 1266).
r. — Salts, — All salts of copper, except the sulphides, are soluble in am-
monium hydroxide. All cuprous salts are insoluble in water, soluble in
hydrochloric acid and reprecipitated upon addition of water. They are
readily oxidized to cupric salts on exposure to moist air. Cuprous chloride
and bromide are soluble in ammonium chloride solution (^lohr, J. C,
1874, 27, 1099). Cupric salts, in crystals or solution, have a green or
bine color; the chloride (2 aq.) in solution is emerald-green when concen-
trated, light blue when dilute; the sulphate (5 aq.) is "blue vitriol."
Anhydrous cupric salts are white. The crystallized chloride and chlorate
are deliqnescent ; the sulphate, permanent; the acetate, efflorescent.
Cupric basic carbonate, oxalate, phosphate, borate, arsenite, sulphide,
cyanide, ferrocyanide, ferricyanide, and tartrate are insoluble in water.
The ammonio salts, the potassium and sodium cyanides, and the potassium
and sodium tartrate, are soluble in water. In alcohol the sulphate and
acetate are insoluble; the chloride and nitrate, soluble. Ether dissolves
the chloride.
106 COPPER. §77, 6a.
6. Beactions. — a. — Fixed alkali hydroxides precipitate acid solutions of
cuprous chloride, first as the white cuprous chloride, changing with more of
the alkali to the yellow cuprous hydroxide, insoluble in excess. AmnLonium
hydroxide and carbonate precipitate and redissolve the hydroxide to a color-
less solution, which turns blue on exposure. The .colorless ammoniacal solution
is precipitated by potassium hydroxide. Fixed alkali carbonates precipitate
the yellow cuprous carbonate, CUjCO, .
Fixed alkalis — ^ZOH — added to saturation in solutions of cupric salts,
precipitate cupric hydroxide, Cn(0H)2 , deep blue, insoluble in excess unless
concentrated (Locw, Z., 1870, 9, 463), soluble in ammonium hydroxide (if
too much fixed alkali is not present), very soluble in acids, and changed,
by standing, to the black compound, Cii30o(0H)2; by boiling, to CuO .
K tartaric acid, citric acid, grape-sugar, milk-sugar, or certain other
organic substances are present, the precipitate either does not form at all,
or redissolves in excess of the fixed alkali to a blue solution. The alkaline
tartrate solution may be boiled without change; in presence of glucose,
the application of heat causes the precipitation of the yellow cuprous
oxide. Alkali hydroxides, short of saturation, form insoluble basic salts,
of a lighter blue than the hydroxide.
Ammonium hydroxide added short of saturation precipitates the pale
blue basic salts ; added just to saturation, the deep blue hydroxide (in both
cases like the fixed alkalis); added to supersaturation, the precipitate dis-
solves to an intensely deep blue solution (separation from bismuth). The
blue solution is a cuprammonium compound, not formed unless ammonium
salts be present. It has been isolated as CuSO^. (1^8)4 (§77, 06). The deep
blue solution probably consists of this compound in a hydrated condition,
i. c. Cu(0H)2.2NH,0H.(irH,)2S0, ; or (NH,)4Cu(0H)4S0, . Other salts
than the sulphate form the corresponding compounds: CnClj -|- -ilTH^OH
= Cu(0H)o.2irH40H.2irH4Cl . The blue color with ammonium hydroxide
is a good test for the presence of copper in all but traces (one to 25,000),
its sensitiveness is diminished by the presence of iron (Wagner, Z., 1881,
20, 351). Ammonium carbonate, like ammonium hydroxide, precipitates
and redissolves to a blue solution. Carbonates of fixed alkali metals — ^as
K2CO3 — precipitate the greenish-blue, basic carbonate, Cii2(0H)2C03 , of
variable composition, according to conditions, and converted by boiling to
the black, basic hydroxide and finally to the black oxide. Barium carbon-
ate precipitates completely, on boiling, a basic carbonate.
From the blue ammoniacal solutions a concentrated solution of a fixed
alkali precipitates the blue hydroxide, changed on boiling to the black
oxide, CnO .
h. — Oxalates, cyanides, ferrocyanides, ferricyanides and thiocyanates pre-
cipitate their respective ciiprous salts from cuprous solutions not too strongly
acid. The ferricyanide is brownish-red, the others are white. The thiocyanate
is used to separate copper from palladium (Wohler, A, Ch,, 1867, (4), 10, 610);
and also from cadmium. In solutions of cupric salts, oxalates precipitate cupric
§77, 6«. COPPER. lor 5
oxalatey CnG,04 , bluish-white, insoluble in acetic acid, and formed from mineral
acid salts of copper by oxalic acid added with alkali acetates.
Potassinm cyanide forms the yellowish-green cupric cyanide, Cii(CN)2 ,
soluble in excess of the reagent with formation of the double cyanide,
2ZCH.Cu(CN)2 , unstable, changing in whole or in part to cuprous cyanide.
The potassium cyanide also dissolves cupric oxide, hydroxide, carbonate,
sulphide, etc., changing rapidly to cuprous cyanide in solution in the
alkali cyanide. This explains why hydrogen sulphide does not precipitate
solutions of copper salts in potassium cyanide, used as a separation from
cadmium. Potassium ferrocyanide precipitates cupric ferrocyanide,
Cii^e(CH)«, reddish-brown, insoluble in acids, decomposed by alkalis; a
very delicate test for copper (1 to 200,000); forming in highly dilute solu-
tions a reddish coloration (Wagner, Z., 1881, 20, 351). Potassium fern-
cyanide precipitates cupric ferricyanide, Cu3(re(CN)o)2 , yellowish-green,
insoluble in hydrochloric acid.
Potassium, thiocyanate, with cupric salts, forms a mixed precipitate of
cuprous thiocyanate, .white, and a black precipitate of cupric thiocyanate,
which gradually changes to the white cuprous compound, soluble in NH4OH;
in the presence of hypophosphorous or sulphurous acid the cuprous thiocyanate
is precipitated at once (distinction from cadmium and zinc) (Hutchinson, J. C,
1880, 38, 748). Ammonium, benzoate (10 per cent solution) precipitates copper
salts completely from solutions slightly acidified (separation from cadmium)
(Gucci, B., 1884, 17, 2659).
If to a solution of cupric salt slightly acidulated with hydrochloric acid, an
excess of a solution of nitroso-B-naphthol in 50 per cent acetic acid be added,
the copper will be completely precipitated on allowing to stand a short time
(separation from Pb , Cd , Hg , Mn , and Zn) (Knorre, /?., 1887, 20, 283).
Potassium, xanthate gives with very dilute solutions of copper salt a yellow
coloration; according to Wagner (/. r.) one part copper in 900,000 parts water
may be detected.
c. — Nitric add rapidly oxidizes cuprous salts to cupric salts, d. — A solution
of cupric sulphate slightly acidulated with hydrochloric acid is precipitated as
cuprous chloride by sodium hypophosphite (Cavazzi, Oazzetta^ 1886, 16, 167); if
the slightly acidulated copper salt solution be boiled with an excess of the
hvpophosphite the copper is completely precipitated as the metal. Sodium
pn osph ate» Na2HP04 , gives a bluish-w-hite precipitate of copper phosphate,
CUHPO4 , if the reagent be in excess and CUg(P04)2 if the copper salt be in
excess. Sodium, pyrophosphate precipitates cupric salts, but not if tartrates
or thiosiilphates be present (separation from cadmium) (Vortmann, 5., 1888,
21, 1103).
€. — Cuprous salts (obtained by treating cupric salts wuth SnCL) when boiled
with precipitated sulphur deposit the copper as CUaS (separation from cad-
mium) (Orlowski, J, T., 1882, 42, 1232). Cuprous salts are precipitated or trans-
posed by hydrosulphuric acid or soluble sulphides, forming cuprous sulphide,*
Cu^S , black, possessing the same solubilities as cupric sulphide.
With cupric salts HgS gives CuS, black (with some CUgS), produced
alike in acid solutions (distinction from iron, manganese, cobalt, nickel)
• Freshly precipitated cuprous sulphide transposes silver nitrate forming sUvor sulphide,
metalUo silver and cupric nitrate ; with cupric sulphide, silver sulphide and cupric nitrate are
formed (Schneider, Pogg,^ 1874, 152, 471). Freshly precipitated sulphides of F«, Co, Zn, Cd,
Pb, Bl, Sn'', and Sn>^, when boiled with CuCl in presence of NaCl give Cii,8 and chloride of
the metal: with CnCl,, CuS and a chloride of the metal are formed, except that SnS gives
C«,B, CiaCl and 8n>^ (Raschig, II., 1884, 17, d97).
108 COPPER, §77, 6f.
and in alkaline solutions (distinction from arsenic, antimony, tin). — Solu-
tions containing only the one-hundrcd-thousandth of copper salt are
colored brownish by the reagent. The precipitate, CuS , is easily soluble
by nitric acid (distinction from mercuric sulphide); with difhculty soluble
by strong hydrochloric acid (distinction from antimony); insoluble in hot
dilute sulphuric acid (distinction from cadmium) ; insoluble in fixed alkali
sulphides, and but slightly soluble in ammonium sulphide (distinction
from arsenic, antimony, tin); soluble in solution of potassium cyanide
(distinction from lead, bismuth, cadmium, mercury).
Concerning the formation of a colloidal cupric sulphide, see Spring (B., 1883,
16, 1142). According' to Brauner (C. A'., 1896, 74, 99) cupric salts with excess
of hj'drogen sulphide always yield a very appreciable amount of cuprous
sulphide. See also Ditte (C. r., 1884, 98, 1492). Solutions of cupric salts are
reduced to cuprous salts by boiling* with sulphurous acid (Kohner. (\ (\, 1880.
813). Sodium thiosulphate added to hot solutions of copper salts gives a black
precipitate of cuprous sulphide. In solutions acidulated with hydrochloric
acid, this is a separation from cadmium (Vortmaun, J/., 1888, 9, 165).
/. — Hydrobromic acid added to cupric solutions and concentrated by
evaporation gives a rose-red color. Delicate to 0.001 m. g. (Endemann
and Prochazka, C. N., 1880, 42, 8). Of the common metals only iron
interferes. Potassium bromide and sulphuric acid may be used instead
of hydrobromic acid.
Hydriodic acid and soluble iodides precipitate, from concentrated solu-
tions of copper salts, cuprous iodide, Cul , white, colored dark brown by the
iodine separated in the reaction * (a). The iodine dissolves with color in
excess of the reagent, or dircolvcs colorless on adding ferrous sulphate or
soluble sulphites, by entering into combination. Cuprous iodide dissolves
in thiosulphates (with combination).
The cuprous iodide is precipitated, free from iodine, and more com-
pletely, by adding reducing agents with iodides; as, Na^.SO:( , HoSOg ,
reso/(&).*
(a) 2CuS0. + 4X1 = 2CuI + la -f 2K,S0,
(6) 2CuS0, -f 2KI + 2FeS0, = 2CuI -h K,SO, -h l!e,(SOJ,
2CuS0. + 4KI -h H,SO, + H,0 = 2CuI + 2X^80. -|- H^SO, + 2HI
/;. — Arsenites. as KAsO* , or arsenous acid with just sufficient alkali hydrox-
ide to neutralize it, precipitate from solutions of cupric salts (not the acetate)
the green copper arscnitc, chiefly CuHAsOa (Scheele's green, ** 1 aris green''),
readily soluble in acids and in ammonium hydroxide, decomposed by strong
potassium hydroxide solution. From cupric acetate, arsenites precipitate, on
boiling, copper aceto-arMetiUe, (CuOAs,08)sCu(C.^HsO,)., , Schweinfurt green or
Imperial g-reen, " Paris green," dissolved by ammonium hydroxide and by
acids, decomposed by fixed alkalis.
Soluble arsenates precipitate from solutions of cupric snlts cupric arsenate,
bluish-green, readily soluble in acids and in ammonium hydroxide.
h. — Potassium bichromate does not precipitate solutions of cupric salts:
* The precipitation is incomplete unlosa the free iodine, one of the products of the reaction, is
removed by means of a reducinfir agent (S^4).
|77, 10.
COFPEB.
10t>
normal potas^um chromrtte forms a browni shared precipitate, loltible in um-
tuonuim hydroxide to a grreeti eolution, soluble In dilute acidis.
7. Ignition* — l|fnjtioti with soiUum carliOTiale ou tiiarecia.1 leaves mtMallic
copper ia liuely divided grains. The partielt-ts txre iruthered by triturathiM" i)ie
ehareoal iiiaj>8'in a Bmall tnortflr» with Thtr repeated addition and deeantatioii
of water until the copper Bubside** clean* It is recognized by its color* and
its softness under the knife. Copper readily dissoUej^, from itis eonipound.s in
beads of Ixirax and of mjcroco^mio salt. In the onter flame of the blow-pipe*
The beads are green whiltf bot» tind hlur when cold. In the inner ilnmc the
borax be^d becomes colorleiis when hot: the niicrorosmic i^alt tarns dark green
when hot, both leaving' a redd is^h -brown tint when cold (Chi;0) (helped bv add-
ing" tin), Corapoiindf^^ of eopper» heated in the iniier tlame, ealoi the outer flame
green. Addition of hydrochloric acid int^reaees the delicacy of the reaction,
giving a greeniah-blne color to the flanie.
8, Betcction,— Copper m precipitated from its solutions by H.S , form-
ing CnS. By its in^^nhibility iii (NHJ.S^and soluliUity in hot diiuto HHO;,
it is 6<?parated with Pb , Bi , and Cd from the Temaining metalg of the tin
and copper group. Dibit t^ H-^SO^ with C.H,.OH removes the lead and
ammonium hydroxide precipitates the bismuth as Bi(OH)^, leaving the
Cu and Cd in solution. The presence of the Cm m indicated by the blue
color of the ammoniaeal eoliition, by its precipitation as the brown ferro-
cvanide after acidulation with HCl {Gh)\ and by its reduction to Cu° with
Fe^, from its neutral or acidulated solutions (10). Study the text on
reactions (6) and |102 and §103.
9- liBtimation.^ — (/) It is precipitated on platinum by the **lectric current or
by means of zinc* the excesR of y.mc may be dissolved b^^ dilute Tbydrochloric
acid* {ij It ifi converted into CuO and weiffhed after ignition, or the oxide is
reduced to the metal in an atmosphere of hydroaren and weij^'^hed as such,
{$) It may be precipitated either by H.,S or Na^S^O^ , ond^ after addhijEr free
sulphur and ij^-uitiD^ in hydrog-en ^im^ weighed as euprotis i^ulphideH, or it may
be precipitated by KOKS In presence of H,SOa or H^PO. . and, after adding" s\
ig-nited tti M aud wei|rhed a« Cu^S . Cu,0 , CuO , Cu(NO,):,CiiCO| , CnSO, ,
and many other cuprie salts, are converted into CuiS by luldiiip- S and igrniting-
In hydrog-cn gn^. ()> By adtiin^ KI to tlie eupric suit and titr^itini^ the liber-
ated I by Na,,3;0j; not iiermi^sible with acid radieals whieh oxidize HI,
(5) By precipitation as Cnl ^nd welfrhinR^ after drying at 150° (Browning,
Am. S*. lSf»3 [31, 46, 280)* (6) By titratiujr in eoueeuirated HBr , using a
solution of SnClj in eonceutratrd HCl: the end reaction is sharper than with
SnClj alone (Etard aud Lebeau, T. r„ 1B90, 110. 40H), (7) By titration with
Na^S. Zinc does not interfere (Borntrag-er, Z. timfew-, 18^3, ."in), {^) By
reduction with SO, and precipitation with excess of standard HH^CIfS: dilu-
tion to definite volume and titration of the excess of NH^CNS in sin allijuot
part, with AgNO, (Volhurd, A., lS7y, 100, .>1). (9) Small amounts are treated
with un excess of KH|OH and estimated col ori metrically by comparing with.
fctandard tubes,
10. Oxidation.— Sol utiouis of Cu" and Cu' are reduced to tbe metallic
etatc by Zn , Cd , Sn , Al , Bb , Fe , Co , Ni , Bi , Mg *, P , and in presence
of SOH by K.SnO. . A briirht strip of iron in solution of eupric salts
addulated with hydrochloric acid, receives a bright copper coating, reeo^-
nizahle from solutions in 120,000 parts of water. With a zinc-platinum
• Warren, CiV^., 1886,71,92. |
L^rffc
110 CADMIUM. §78, 1.
couple the copper is precipitated on the platinum and its presence can be
confirmed by the use of H2SO4 , concentrated, and KBr , an intense violet
color is obtained (Creste, J, C, 1877, 31, 803). Cu" is reduced to Cu' by
Cu** (Boettger, J. C, 1878, 34, 113), by SnClg in presence of HCl, in
presence of KOH by A82O3 and grape sugar, by HI , and by SOj . Metallic
copper is oxidized to Cu" by solutions of Hg", Hg', Ag', Pt^^, and Au'",
these salts being reduced to the metallic state. Ferric iron is reduced to
the ferrous condition (Hunt, Am. S., 1870, 99, 153). Copper is also oxi-
dized by many acids.
§78. Cadmium. Cd = 112.4 . Valence two.
1. Properties.— iSpc«7?o gravity , liquid, 7.989; cooled, 8.67; hammered, 8.6944.
Melting point, 320.68** (Callendar and Griffiths, C. A^, 1891, 63, 2). Boiling point,
Tea** to 772° (Carnelley and Williams, J. C, 1878, 33, 284). &]}€ciflc heat is 0.0567.
Vapor density (H = 1), 55.8 (Deville and Troost, A. Ch., 1860, (3), 58, 257). From
these data the gaseous molecule of cadmium is seen to consist of one atom
(Richt^r, Anorg. Chern'., 1893, 363). It is a white crystalline metal, soft, but
harder than tin or zinc; more tenacious than tin; malleable and very ductile,
can easily be rolled out into foil or drawn into fine wire, but at SO** it is brittle.
Upon bending it gives the same creaking sound as tin. It maj' be completely
distilled in a current of hydrogen above 800°, forming silver white crystal's
(Kammerer, B., 1874, 7, 1724). Only slightly tarnished by air and water at
orjiinary temperatures. WTien ignited burns to CdO . When heated it com-
bines directly with CI , Br , I , F , S , Se , and Te . It forms many useful alloys
having low melting-points.
2. Occurrence. — Found ns greenockite (CdS) in Greenland, Scotland and Penn-
sylvania; also to the extent of one to three per cent in many zinc ores.
3. Preparation. — Reduced by carbon and separated from zinc (approximately)
by distillation, the cadmium being more volatile. It may be reduced by fusion
with H , CO , or coal gas.
4. Oxide and Hydroxide. — Cadmium forms but one oxide, CdO , either by
burning the metal in air or by ignition of the hydroxide, carbonate, nitrate,
oxalate, etc. It is a brownish-yellow powder, absorbs CO« from the air, becom-
ing white (Gmelin-Kraut, 3, 64). The hydroxide 'C&{OH.). is formed by the
action of the fixed alkalis upon the soluble cadmium salts; it absorbs CO, from
the air.
5. Solubilities. — a. — Metal. — Cadmium dissolves slowly in hot, moderately
dilute hydrochloric or sulphuric acid with evolution of hydrogen; much more
readily in nitric acid with generation of nitrogen oxides. It is soluble in
ammonium nitrate without evolution of gas; cadmium nitrate and ammonium
nitrite are formed (Morin, C. 7*., 1SS5, 100, 1497). 6.— The oxide and hydroxide
are insoluble in water and the fixed alkalis, soluble in ammonium hj^droxide,
readily soluble in acids forming salts; soluble in a cold mixture of fixed alkali
and alkali tartrate, reprecipitated upon boiling (distinction from copper)
(Behal, J. Phann., 1885, (5), 11, 553). c— Salts.— The sulphide, carbonate,
oxalate, phosphate, cyanide, ferrocyanide and ferricyanide are insoluble (§27)
in water, soluble in hydrochloric and nitric acids, and soluble in KH«OH .
except CdS . The chloride and bromide are deliquescent, the iodide is perma-
nent; they are soluble in water and alcohol.
6. Ecactions. a.— The fixed alkali hydroxides— in absence of tartaric
and citric acids, and certain other organic substances — ^precipitate, from
solutions of cadmium salts, cadmium hydroxide, Cd(0H)2 , white, insoluble
5578, 6». CADMIUM. Ill
in excess of the reagents (distinction from tin and zinc). Ammonium
hydroxide forms the same precipitate which dissolves in excess. If the
concentrated cadmium salts be dissolved in excess of ammonium hydroxide
with gentle heat and the solution then cooled, crystals of the salt, with
variable amounts of ammonia, are obtained; e. g,, CdCIjClTHa)., ,
Cd80,(15rH3),, Cd(N03)2(irH3)e (Andre, C. r., 1887, 104, 908 and 987;
Kwasnik, Arch, Pharm,, 1891, 229, 569). The fixed alkali carbonates pre-
cipitate cadmium carbonate, CdCOg , white, insoluble in excess of the
reagent, ammonium carbonate forms the same precipitate dissolving in
excess. Barium carbonate, in the cold, completely precipitates cadmium
salts as the carbonate.
6. — Oxalic acid and oxalates precipitate cadmium oxalate, white, soluble in
mineral acids and ammonium hydroxide. Fotassiom cyanide precipitates
cadmium cyanide, white, soluble in excess of the reapent as Cd(CN)3.2KCN:
ferrocyanides form a white precipitate; ferricyanides a yellow precipitate,
both soluble in hydrochloric acid, and in ammonium hj'droxide. Fotassium
salphocyanate does not precipitate cadmium salts (distinction from copper).
Cadmium salts in presence of tartaric acid are not precipitated by fixed alkali
hydroxides in the cold; on boilinpr, cadmium oxide is precipitated (separation
from copper and zinc) (Aubel and Ramdohr, A. Ch., 1858, (3), 62, 109).
c. — ^Nitric acid dissolves all the known compounds of cadmium, d. — Soluble
phosphates precipitate cadmium phosphate, white, readily soluble in acids.
Sodium pyrophosphate precipitates cadmium salts, soluble in excess and in
mineral acids, not in dilute acetic. The reaction is not hindered by the pres-
ence of tartrates or of thiosulphates (separation from Cu) (Vortmann, B., 1888,
21, 1104).
e. — ^Hydrogen sulphide and soluble sulphides precipitate, from solutions
neutral, alkaline, or not too strongly acid, cadmium sulphide, yellow;
insoluble in excess of the precipitant (Fresenius, Z,, 1881, 20, 236), in
ammonium hydroxide, or in cyanides (distinction from copper) ; soluble in
hot dilute sulphuric acid and in a saturated solution of sodium chloride *
(distinction from copper) (Cushman, Am., 1896, 17, 379).
Sodium thiosulphate, NajSjO, , does not precipitate solutions of cadmium
salts (Follenius, Z., 1874, 13, 438), but in excess of this reagent, ammonium
salts being absent, sodium carbonate completely precipitates the cadmium as
carbonate (distinction from copper) (Wells, C. A'., 1891, 64, 204). Cadmium
salts with excess of sodium thiosulphate are not precipitated upon boiling
with hydrochloric acid (distinction from copper) (Orlowski, J. C, 1S82, 42, 1232).
f. — ^The non-precipitation by iodides is a distinction from copper, g, — Soluble
arsenites and arsenates precipitate the corresponding cadmium salts, readily
soluble in acids and in ammonium hydroxide, h. — Alkali chromates precipitate
yellow cadmium chromate from concentrated solutions only, and soluble on
addition of water.
/. — A solution of copper and cadmium salts, verj' dilute, when allowed to
spread iipon a filter paper or porous porcelain plrte, gives a ring of the cad-
mium salt beyond that of the copper ^It, easily detected by hydrogen sulphide
(Bagley, J, C 1878, 33, 304).
*Otrinflr to the formation of incomplotcly-dissociat-od OdCl,. Cdl, Is Btill less dissociated
and accorrlingly CdS dissolves more roadily in HI than in HCl and much more readily than in
IINOa of the same concentration. On the other hand, of course, precipitation of the sulphide
takes place with more rlitllculty from the iodide than frr^m the other salts.
112
REACTIONS OF BISMUTH, COPPER AND CADMIUM, §78, 7.
7. Ignition. — On charcoal, with sodium carbonate, cadmium salts are reduced
before the blow-pipe to the metal, and usually vaporized and reoxidized nearly
as fast as reduced, thereby forming a characteristic brown incrustation (CdO).
This is volatile by reduction only, being driven with the reducing flame. Cad-
mium oxide colors the borax bead yellowish while hot, colorless when cold:
microcosmic salt, the same. If fused with a bead of KaS, a yellow precipitate
of CdS is obtained (distinction from zinc) (Chapman, J, C, 1877, 31, 490).
8. Deteetion. — Cadmium is precipitated from its solutions by HjS form-
ing CdS. By its insolubility in (SK^)^^^ and solubility in hot dilute HNO^^
it is separated with Pb , Bi , and Cu from the remaining metals of the
second group. Dilute H0SO4 with C2H5OH removes the lead and NH^OH
precipitates the bismuth as Bi(0H)3 , leaving the Cu and Cd in solution.
If copper be present, KCN is added until the solution becomes colorless,
when the Cd is detected by the formation of the yellow CdS with HoS .
If Cu be absent the yellow CdS is obtained at once from the ammoniacal
solution with H2S . See also 6i.
9. Estimation. — (i) It is converted into, and after ignition weighed as an
oxide. (2) Converted into, and after drying at 100°, weighed as CdS. (3) Pre-
cipitated as CdCsO^ and titrated by KMnO^. (4) Electrolytically from a slightly
ammoniacal solution of the sulphate or from the oxalate rendered acid with
oxalic acid, (o) Separated from copper by KI; the I removed by heating: the
excess of KI removed bv KNO, and H2SO4; the cadmium precipitated bv
Na^CO, and ignited to Cd'O (Browning, Am, S,, 1893, 146, 280). (6) By adding
a slight excess of H2SO4 to the oxide or salt, and evap)oration first on the water
bath and then on the sand bath, weighed as CdSO^ (Follenius, Z., 1874, 13, 277).
10. Oxidation. — Metallic cadmium precipitates the free metals from
solutions of Au , Pt , Ag , Hg , Bi , Cu , Pb , Sn , and Co ; and is itself
reduced by Zn , Mg , and Al .
§79. Comparison of Certain Reactions of Bismuth, Copper, and Cadmium.
Taken in Solutions of their Chlorides, Nitrates^ Svlphates, or Acetates.
KOH or NaOH, in
excess
NH4OH, in excess
Dilution of satu-
rated solutions. .
Iodides
Sulphides
Iron or zinc
Glucose, KOH, and
heat
K^SnO, + KOH..
Bi
Bi(OH)„ white.
Bi(OH)„ white.
BiOCl, white (§76,
5d).
Partial precipita-
tion in solutions
not very strongly
acid (§76, 6f).
Bi^S,, black, in-
soluble in KCN.
Bi, spongy precipi-
tate.
Bi, black.
Bi, black.
Cu
Cu(0H)2, dark
blue.
Blue solution.
Precipitation of
Cul, with libera-
tion of iodine
(§77, 6f).
Cu^S and CuS,
black, soluble in
KCN.
Cu, bright coating
(§77, 10).
CUjO, yellow (§77,
5ft).
Cu, precipitated
metal.
Cd
Cd(OH)„ white.
Colorless solution.
CdS, yellow, insol-
uble in KCN.
Cd, gray sponge
with zinc, no ac-
tion with iron.
§81,-4. PRECIPITATION OF METALS OF SECOND GROUP. 113
Systematic Analysis of the Metals of the Tin and Copper Group.
The precipitation of the metals of the second group (Tin and Copper
Group) hy hydrosulphuric acid, and their separation into Division A (Tin
Group) and Division B (Copper Group). See §312.
§80. Manipulation. — The filtrate from Group I. (§62), or the original
solution, if the metals of the silver group be absent, is rendered acid with
a few drops of HCl , warmed and saturated with hydrosulphuric acid gas.
2H,A804 -h xHCl + 5H,S = As,S, + xHCl + 8H,0
•or 2H,As04 + xHCl + 5H,S =As,S, + xHCl + S, + SHaO
SnCl^ 4- 2H,S = SnS, + 4HC1
SnCl, + H,S = SnS + 2HC1
2Bi(N0,), + 3H3S = Bi,S, + 6HN0,
CdSO* + H,S = CdS + H,S04
The precipitate, after being allowed to settle a few minutes, is filtered and
thoroughly washed with hot water containing a little HCl . A portion of
the filtrate diluted with water is again tested with H2S to insure complete
precipitation (§81, 2), and if necessary the whole of the filtrate is diluted
and again precipitated. The filtrate containing no metals of the second
group is set aside to be tested for the remaining metals (§128).
§81. Notes, — 1, Hydrosulphuric acid gas should be used in precipitating the
metals of the second group. It should be generated in a Kipp apparatus,
using ferrous sulphide, FeS , and dilute commercial sulphuric acid (1-12).
Commercial hydrochloric acid riay be iised instead of sulphuric. The gas
should be passed throiigh a wash bottle containing water to remove any acid
that may be carried over mechanically. It should always be conducted through
a capillary tube into the solution to be analyzed. Less gas is required and the
solution is less liable to be thrown from the test tube by the excess of unab-
sorbed gas.
2. In testing the filtrate for complete precipitation, instead of the gas, a cold
saturated water solution of the gas may well be employed. This dilutes the
solution at the same time. In treating the unknow^n solution with HjS or in
making a saturated water solution of the gas, it should be passed into the
liquid until, upon shaking the test tube or bottle capped with the thumb, there
is no formation of a partial vacuum due to the further absorption of the gas by
the liquid.
3. H2S is decomposed by HNO, or HNO, + HCl (nitrohydrochloric acid)
(§257, 6^), hence these acids must not be present in excess. If these acids
were used in preparing the solutions for analysis, they must be removed by
evaporation. Sulphuric acidulation is not objectionable to precipitation w^ith
HjS , but could not be used until absence of the metals of the calcium group
(Group V.) had been assured.
4. The precipitation of the silver group has left the solution acid with HCl ^
and prepares the solution for precipitation with HjS , if other acids are not
present in excess and if too much HCl was not employed. The presence of a
great excess of HCl does not prevent the precipitation of arsenic (§69, 6e), but
does hinder or entirely prevent the precipitation of the other metals of this
group, especially tin, lead (§67, 6e), cadmium and bismuth. The solution must
be acid or traces of Co , Ni and Zn (§135, 6c) will be precipitated. No instruc-
tions can be given as to the exact amount of HCl to be employed. About one
part of HCl to 25 of the solution should be present to prevent the precipitation
114 PRECIPITATION OF METALS OF SECOND OROUP. §81, 5.
of Zn , and it is seldom advisable to use more than one part of HCl to ten of
the solution ♦ (this refers to the reagent HCl , §324).
o. The precipitation takes place better from the warm solutions than from
the cold (§31); hence it is directed to warm the solution before passing in the
HaS , and before Altering heat again nearly to boiling. If arsenic be present^
the solution should be kept at nearly the boiling point, and the gas passed
into the solution for several minutes (§69, 6c).
6. The precipitated sulphides of the metals of the tin and copper group
(second group) present a variety of colors, which aid materially in the further
analysis of the group. CdS , SnSj , ASjSa and ASjSj are lemon-yellow; Sb^Ss.
and SbsS, are orange; SnS , HgS , PbS , BisS, , CUjS and CuS are black to
brownish-black. If too much HCl be present, lead salts frequently precipitate
a red double salt of lead chloride and lead sulphide (§57, 6e). Mercuric chloride
at first forms a white precipitate of HgCl2.2HgS , changing from yellow to red,
and finally to black with more H3S , due to the gradual conversion to Hg^
(§58, 6e).
7. Addition of water to the solution before passing in HjS may cause the
precipitation of the oxychlorides of Sb , Sn or Bi (5d; §70, §71 and §76). These
should not be redissolved by the addition of tnore HCl, as they are readily
transposed to the corresponding sulphides by HaS , and the excess of acid
necessary to their resolution may prevent the precipitation of cadmium or
cause the formation of the red precipitate with lead chloride.
8. Arsenic when present as arsenic acid is precipitated exceedingly slowly
from its cold solutions, and tardily even from the hot solutions. Frequently
the other metals of the group may be completely precipitated and removed by
filtration, w^hen a further treatment with HsS causes a precipitation of the
arsenic as As.Si^ from the hot solution. This slow formation of a yellow pre-
cipitate is often a very sure indication of the presence of pentad arsenic (§69,
6'e,i).
9. The presence of a strong oxidizing agent as HNO, , K.Ct.Oj , FeCl, , etc.,
causes with HjS the formation of a -white precipitate of sulphur (§125. 6c) ^
which is often mistaken as indicating the presence of a second group metal.
If the original solution be dark colored, it is advisable to warm with hydro-
chloric acid and alcohol (§125, 6f and 10) to effect reduction of a possible higher
oxidized form of Cr or Mn before the precipitation with HjS , thus avoiding
the unnecessary precipitation of sulphur.
10. Complete precipitation of the metals of the second group with H,S may
fail: (1) from incomplete saturation with the gas (§81, 2): (2) from the pres-
ence of too much HCl (§81, 4); (S) from the presence of much pentad arsenic
(§69, fif). The first cause of error may be avoided by careful observance of the
directions in note (2). To prevent the second cause of error a portion of the
filtrate, after the removal of the precipitate by filtration, should be largely
diluted with water (10 volumes) and HaS (gas or saturated water solution)
again added. In case a further precipitate is obtained, the whole of the filtrate
Fhould be diluted and again precipitated with HaS . This should be repeated
until the absence of second group metals is assured. If a slow formation of a
yellow precipitate indicating Asv is observed, HjS should be passed into the
•Addition of a strong acid, containing H Ions in large quantity, diminishes the already slight
dissociation of the H^S ($44), thus decreasing In number the S ions, whose concentration multi-
plied by that of the metal Ions must equal the solubility-product of the sulphide In question,
before precipitation can take place. Precipitation of some of the sulphides of the Tin and
Copper Group may be entirely prevented in this way.
It frequently happens that addition of water alone will cause precipitation of these sulphides
from a strongly acid solution which has been saturated with H,S. This appears strange in view
of the fact that the acid which prevented precipitation and the acid which Anally produced it
were both diluted by the added water In the same proportion. But as a matter of fact dilution
does not have the same effect on a strong acid as on a weak one. Dissociation is always in-
creased by dilution, but in much greater ratio in the case of a weakly-diaaooiated body as H,8
than whore the dissociation of the substance Is already practically complete, as in the case of
the strong acid. Dilution in the case mentioned increases the relative concentration of the S ,
ions and so the solubility-product is reached and precipitation results.
§8854. PRECIPITATIOy OF METALS OF SECOND GROIP. 115
hot sohition for fully 30 minutes (Note J) or the solution should be treated
with SO, or some other agent for the reduction of Asv to As'" (§69, 10).
§82. Kanipnlation. — After the precipitate has been well washed with
hot water the point of the filter is pierced with a small stirring rod and
the precipitate washed into a test-tube, using as small an amount of water
as possible. Yellow ammonium sulphide (1^4)28, (§83, 2) is then added
and the precipitate digested for several minutes with warming:
A8,S, + 2(NH,),S, = (NH,),A8,S. + S,
SnS + (NHJ2S2 = (NHJaSnSa
2SnS2 + 2(NH,),S2 = 2(NHJaSnS, + S,
2Sb,S, -h 6(NH,),S, = 4(NH,).SbS, + S,
2MoS, -h 2(NHJ,S, = 2(NHJ,MoS, + S,
The precipitate is then filtered and washed once or twice with a small
amount of (1^4)28, , and then with hot water. The filtrate consisting of
solutions of the sulphides of As, 8b, 8n, An, Pt, Ho {Or, Ir, Se, Te,
Wy V), constitutes the Tin Oroup (Division A of the second group). The
precipitate remaining upon the filter, consisting of the sulphides of Hg ,
Pb, Bi, Cu, Cd (O5, Pd, Bh, and Ru), constitutes the Copper Oroup
(Division B of the second group, §95).
§83. Notes, — 1. The precipitate of the sulphides of the tin and copper group
must be thoroughly washed with hot wa^r (preferably containing HjS and
about one per cent of reagent HCl to prevent the formation of soluble colloidal
sulphides ($69, 5e), to insure the removal of the metals of the iron and zinc
groups, which would be precipitated on the addition of the ammonium sulphide
(5144).
2. Yellow ammonium sulphide, (NH4)2Sx. forms upon allowing the normal
sulphide, (NH4)2S , to stand for sometime, or it may be prepared for imme-
diate use by adding sulphur to the freshly prepared normal sulphide (§257. 4).
For arsenic sulphides the normal ammonium sulphide may be employed, but
the sulphides of antimony are soluble with difficulty, and stannous sulphide is
scarcely at all soluble in that reagent; while they are all readily soluble in the
yellow polysulphide (6e; §69, §70 and §71).
.3. Cupric sulphide, CuS , is sparingly soluble in the yellow ammonium sul-
phide and will give a grayish-black precipitate upon acidulation with HCl .
The sulphides of the tin group are soluble in the fixed alkali sulphides, KoS
and NajS; cupric sulphide is insoluble in these sulphides. Mercuric sulphide,
however, is much more soluble in fixed alkali sulphides than cupric sulphide is
in the (KH4)3Sx. If copper be present and mercury be absent, it is recom-
mended to use K.S or Na.S instead of (NH4)2Sx for the separation of the
second group of sulphides into divisions A (tin group) and B (copper group).
But if Hg^ be present, the (NHJ.Sx should be used, and the presence or
absence of traces of copper be determined from a portion of the filtrate from
the silver group before the addition of H,S (§103).
4. The sulphides dissolve more readily in the (NH4)2Sx when the solution is
warmed. An excess of the reagent is to be avoided, as the acidulation of the
ssolution causes the precipitation of sulphur (§256, 3), which may obscure the
precipitates of the sulphides present.
116
TABLE FOR THE AXALYt^HS OF THE TIX OliOVP,
§84.
09
CO
ace
u
V
X
a:
c
O
c
C
pa
1 2 ^
a « js
1; s:
: II
: I
lis
'St- .
^ i a
CC IS
C OS 0^
■^ p— •'^
a; "d ac |H
S 4; C
o Jf o S
.5 S -S fe ^
o s
1^
— ? CS
O -SC 13
2 w = a;
."tr i; « :::
— ^ o 2,
r3 fc- 3 o
•3 = .b
2
c
c
o
p
CC IT"
r
be C3
.5 u
2 c
o ^
2^
C -S ,5 ^
o
s
o
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c
fi
1
2 a 4* - "^
Is :
C3 \0
p.
o i: c^
^©9
. N
t!t ^« .S -E
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^ = =
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7: o
-as"!
p CC
a. ^
o c
Si
^« i
g-ag
es
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s
I
w
S d
jgis .
OS'S
IS?
o
M
o
ft
r oc o c5«c-a o t
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1
P4
O >w"*- to* (k
-SI c-d-E 5
.S^ « < « o £ .?.
CO
1i
L-^ ?" - iT-
c ^
« = -= ci-t: > s:
u ^ 5 F -r t: ^
o g.
18
o a, ^^
*' /i. e *^ -.
o ^' c s i: c.
P c i fc t- •= ^
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to*
5 - s
^ c
o —
6 ..? ^
C « y' .^
r >- „ 3/ cs ♦; "»-b.
r5xcC:cP<P.c
s
.kJzS^'':
-^ r tt >» c = c
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2 i c « = ^ ^
'q ^ X S K ♦- —
§84.
TABLE FOR THE AXALYSfS OF THE TIX GROUP,
iir
or
0) — " w
U 0(9 c
-5 -.«
f^l
!«
es s
^'
lie
c'fi Si's
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ce
^ O t^ o *" ^
'3
« « <^
C' M :?
TT — «M d O S "^
9J
^ a; fc- s es -
2:^ C C 5 O
^ ^« ^ c> ^^ ~
g>^^« -^ CS ♦- *- ■♦J -
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fc «3 ? C O
^-^ 2 "To •s^'ing'fcG
O
118 DIRECTIONS FOR ANALYSIS WITH NOTES, §86.
§85. Manipulation.— The solution of the sulphides in {irE^)^S^ is care-
fully acidulated with hydrochloric acid:
2(NH,),Sa + 4HC1 = 4NH,a + S, + 2H,S
(NH4)4As,S» + 4HC1 = As^S, + 4NH4CI + 2H,S
2(NHJ,SbS, + 6HC1 = SbjS^ + 6NH,a + 3H,S
(NH4),SiiS. + 2HC1 = SnS, + 2NH4CI + H,S
The precipitate obtained when the metals of the tin group are present,
is usually yellow or orange-yellow and is easily distinguished from a pre-
cipitate of sulphur alone (SnS and M0S3 are brownish-black). It should
be well washed with hot water and then dissolved in hot HCl using small
fragments of KCIO3 (§69, Ge) to aid in the solution:
2As,S, -h lOCl, -f I6H2O = 4H,A80, -f 20HC1 + 3S,
SnS, + 4HC1 = SnCl, + 2H,S
PtS, + 2Cl,= PtCl, -f S,
The solution is boiled (to insure removal of the chlorine (§69, 10) until it
no longer bleaches litmus paper.
§86. NoteiK. — 1. If the precipitate obtained is white, it probably consists of
sulphur alone and indicates absence of more than traces of the metals belong-
ing to this group (GeS, is white, §111, 6).
2. Care should be taken not to use too much HCl in precipitating the sul-
phides from the (NH4)2Sx solution, as some of the sulphides (especially SnSj)
are quite soluble in concentrated HCl .
3, It will be noticed (§85) that the low^er sulphides of Sb and Sn are oxidized
by the (NH4)2Sx . and are precipitated by the HCl as the higher sulphides
SbaSft and SnS^ respectively. This fact may be most readily observed by the
precipitation of a solution of SnCl, with HjS , giving a brown precipitate of
SnS , then dissolving this precipitate in (NH4)2Sx and reprecipitating with HCl
as the orange-colored SnS, .
J. Hot reagent HCl (§324) dissolves the sulphides of tin quite readily
without reduction; the sulphides of antimony, slowly forming SbCl, only; and
the sulphides of arsenic practically not at all, or at most only traces. The
sulphides of Au and Pt are not soluble in HCl . MoS, is soluble in hot con-
centrated HCl . The relative solubility of these sulphides in HCl is used by
some chemists as the basis of a separation of As from Sb and Sn (§69, 6r. also
bottom of next note, .7).
o. The sulphides of arsenic are readily soluble in ammonium carbonate (§69.
5r) and are thus separated from the sulphides of Sb and Sn, which are prac-
tically insoluble. The following table suggests a method of analysis based
upon this property of these sulphides.
§86,5. DIRECTIONS FOR ANALYSIS WITH NOTES.
Digest with Bolution of ammonium carbonate and filter.
119
Residue: SnS, , 8bsS» , (S) .
Dissolve in hot hydrochloxic acid (5<*, §70
and §71).
Solation: SnCl4 , SbCl, .
Treat with zinc and hydrochloric acid in
Marsh*a apparatus (§69, 6'a).
Deposit: Sn, (Sb) .
Dissolve by hydro-
chloric acid.
Solution: SnCl, .
(Residue, Sb .)
Test by ammoniacal
silVer nitrate and
bv mercuric chlo-
ride (§71, 6i and j).
Oas: SbH, .
(Test the spots,
§69, 6V, i.)
Receive the gas in
solution of silver
nitrate. Dissolve the
precipitate (SbAgs)
(§70. 6/), and test
by H,S (§87 and
§89).
Solution:
(NHJ.AsS, + (NHJ.AsO,
and
(NH,)4As,S. + (NHJ.As^O, .
Precipitate by hydrochloric acid;
filter; wash the precipitate and
dissolve it by chlorine gener-
ated from a minute fragment of
potassium chlorate and a little
hydrochloric acid (§69, 5c).
Expel all free chlorine (note 9,
and §69, 10).
Solution: H.AsO^ .
Apply Marsh's Test, as directed in
§69, 6'a, testing the spots (§69,
6V); receiving the gas in solu-
tion of silver nitrate, and test-
ing the resulting solution (§87).
Examine the original solution, as
indicated in §88, i.
The plan above given may be varied by aepnralinp antimony «»?(/ tin by ammo-
hImw carttonate in fully oxidized solution, as follows: The Sb^S^ and SnSj are
dissolved by nitrohydrochloric acid, to obtain the antimony as pyroantimonic
acid. The solution is then treated with excess of animonium carbonate^ in a
vessel wide enough to allow the ctrrbonic acid to escape without waste of the
solution.
The soluble diammoniuir dihydrogen pyroantimonate, (NH4)2H2Sb30T . is
formed. Meanwhile the SnCl^ is fully precipitated as HjSnG, (§71, Gw), and
may be filtered out from the solution of pyroantimonate.
The liability of failure, in this mode of separating antimony and tin, lies in
the non-formation of pyroantimonic acid by nitrohydrochloric acid. The ordi-
nary antimonic acid forms a less soluble ammonium salt, but this acid is not
so likely to occur in obtaining the solution with nitrohydrochloric as anti-
mont/us chloride, SbCl, . Excess of ammonium carbonate does not redissolve
the SbjO, which it precipitates from SbCl, , as stated in §70, 0^/.
The above plan may also be varied as follows: After removal of the arsenic
sulphide with (KH4)2COs , the residue is dissolved* in strong HCl , not iising
KC1O3 or HNOs . The solution consists of SnCl^ and SbCl, . Divide in two
portions: (1) Add Sn on platinum foil. A black precipitate indicates Sb° .
(2) Add iron wire, obtaining Sb° and Sn"; filter and test the filtrate for Sn by
HgCl, (Pieszczek. Arch. Pharm., 1S91, 229, 667).
6. The siilphides of As , Sb and Sn are all decomposed by concentrated nitric
acid, which furnishes a basis of en excellent separation of the arsenic from the
antimony and tin (Vaughan, American Chemi,Ht, 1875, 6, 41). The sulphides
reprecipitated from the (NHJjSx solution by HCl are well washed, transferred,
to an evaporating dish, heated with concentrated HNO, until brown fumes are
no longer evolved, and then evaporated to dryness, using sufficient heat to
expel the HNO, and the H.SO4 formed by the action of the HNO, upon the S .
The heating should be done on the sand bath. The cooled residue is digested
for a few minutes with hot water, the arsenic passing into solution as HjAsO^ ,
and the antimony and tin remaining as residue of SbjO^ and SnOj . The pres-
ence of arsenic may be confirmed by the reactions with AgNO, (§69, 6/), CUSO4
(§69, 6k) by the Marsh test (§69, 6'a), or by precipitation with magnesia mix-
120 DIRECTIONS FOR ANALYSIS WITH NOTES, §86, 7.
ture (§69, 6i). A portion of the residue may be tested in the Marsh apparatus
for the Sb (§70, Oy), another portion may be reduced and dissolved in an open
dish with Zn and HCl (not allowable if As be present, §71, 10), and the result-
ing: SnClj identified by the reaction with HgCl, (§71, 6i).
7. The precipitated sulphides must be thoroughly washed to insure the
removal of the ammonium salts, since in their presence the dangerously ex-
plosive nitrogen chloride (§268, 1) could be formed when the sulphides were
dissolved in HCl with the aid of XClOs .
8. Instead of chlorine (HCl + KCIO,), nitrohydrochloric acid may be em-
ployed, but it is liable to cau&e the formation of a white precipitate of SbjO^
and SnOa .
9. The chlorine should all be removed, as the metals cannot be reduced by
the Zn and H2SO4 in the Marsh apparatus in the presence of powerful oxidizing
agents as CI . This would also require evaporation to expel the HNO, , if
nitrohydrochloric acid were used to effect solution.
10. Hydrogen peroxide, HsO, , decomposes the sulphides of arsenic and anti-
mony with oxidation. The arsenic will appear in the solution, the antimony
remaining as a white precipitate of the oxide (a sharp separation) (Luzzato,
Arch. Pharm., 1886, 224, 772).
§87. Manipulation. — The solution of the metals of the tin group is
then ready to be transferred to the Marsh apparatus (the directions for
the use of the Marsh apparatus are given under arsenic (§69, 6'a), and
should be carefully studied and observed. They will not be repeated
here). Only a portion of the solution should be used in the Marsh appar-
atus, the remainder being reserved for other tests. The gas evolved from
the Marsh apparatus is passed into a solution of silver nitrate, which by
its oxidizing action effects a good separation between the arsenic and
antimony (§89,2):
AsH, -f CAgNO, -f 3HaO = H,AsO, + 6Ag -f- 6HN0,
SbH, -h 3AgN0, = SbAg, + 3HN0,
The hard glass tube of the Marsh apparatus is heated while the gas is
being generated, a mirror of arsenic and antimony being deposited, due
to the decomposition of the gases (§69, 6'c) : 2SbH3 = 2Sb + SH^ . The
ignited gas is brought in contact with a cold porcelain surface for the
production of the arsenic and antimony spots (§69, 6'b). Failure to obtain
mirror, spots, or a black precipitate in the AgNO, is proof of the absence
of both arsenic and antimony. The black precipitate obtained in the
silver nitrate solution is separated by filtration, washed and reserved to be
tested for antimony. The filtrate is treated with HCl, or a metallic
chloride, as CaClj or NaCl , to remove the excess of silver and, after evapor-
ation to a small volume, is precipitated with HgS . A lemon-yellow pre-
cipitate indicates arsenic. The black precipitate from the silver nitrate
solution is dissolved in hot reagent HCl : SbAgs -f 6HC1 = SbCl, -f
3AgCl . The excess of acid is removed by evaporation, a little water is
added (§70, 5rf and §59, 5c) and the AgCl removed by filtration. The
filtrate is divided into two portions. To one portion HgS is added; an
orange precipitate indicates antimony. The HjS may give a black precipi-
tate of AgjS from the AgCl held in solution by the HCl . If this be the
§89,4. DlRECTIOyS FOR ANALYSIS WITH NOTES. 121
case, to the other portion one or two drops of KI are added and the
solution filtered. This filtrate is now tested for the orange precipitate
with H^S.
The mirror obtained in the hard glass tube should be examined as
directed in the text, especially by oxidation and microscopic examination
(§69, &c 6). The spots should be tested with NaClO and by the other tests
as given in the text (§69, 6'c 1),
§88. Notes, — ^Arsenic. — 1, All compounds of arsenic are reduced to arsine hj
the Zn and H^SOf in the Marsh apparatus. Hence if strong oxidizing agents
are absent, the original solution or powder may be used directly in the Marsh
apparatus for the detection of arsenic; but sulphides should not be present.
2, The burning arsine forms As^Og , which may be collected as a heavy white
powder on a piece of black paper placed under the flame. Antimony will also
deposit a similar heavy white powder.
3. The arsine evolved is not decomposed (faint traces decomposed) upon
passing through a drying tube containing soda lime or through a solution of
KOH (distinction and separation from antimony).
jf. Arsenites and arsenates are distinguished from each other by the following
reactions: (a) Arsenous acid solution acidulated with HCl is precipitated in the
cold instantly by H^S; arsenic acid under similar conditions is precipitated
exceedingly slowly (§69, 6c). (b) Neutral solutions of arsenites give a yellow
precipitate with AgNO,; neutral solutions of arsenates give a brick-red pre-
cipitate. Both precipitates are soluble in acids or in ammonium hydroxide
(559, 6g), (c) Magnesia mixture precipitates arsenic acid as white magnesium
am moni um arsenate, MgNH4A804; no precipitate with arsenous acid (§189, (^ff).
id) HI gives iree iodine with arsenic acid; not wnth arsenous acid (§69, i)f).
(f) Alkaline solutions of arsenous acid are immediately oxidized to the pentad
arsenic compounds by iodine (§69, 10). (f) Potassium permanganate is imme-
diately decolored by solutions of arsenous acid or arsenites; no reaction with
arsenates (§69, 10).
§89. Notes. — Antixnony. — /. If antimony be present in considerable amount,
it (in the form of the sulphide) is most readily separated from arsenic by
boiling with strong HCl (solution of the antimony sulphide, (§70, Of)); or by
digesting with (NH4),C0g or NH4OH (solution of the arsenic (§69, 5c)).
2. For the detection of traces of antimony, the most certain test is in its
volatilization as stibine in the Marsh apparatus and precipitation as SbAg, ,
antimony argentide, with AgNO,; this is a good separation from arsenic and
tin, and after filtration it remains to dissolve the SbAg, in concentrated HCl
and identify the Sb as the orange precipitate of Sb...Ss . The formation of the
black precipitate in the AgNO, solution must not be taken as evidence of the
presence of antimony, as arsine gives a black precipitate of metallic silver with
AgNO, . A trace of antimony may be found in the filtrate from the SbAg, ^
hence a slight yellow-orange precipitate from this solution must not be taken
as evidence of arsenic without further examination (§69, 7).
3. SbjS, is precipitated from solutions quite strongly acid with HCl i i. e., in
the presence of equal parts of the concentrated acid («/). gr. 1.20). Tin is not
precipitated as sulphide if there be present more than one part of the con-
centrated acid to three of the solution (§70, 6f). This is a convenient method
of separation. The addition of one volume of concentrated HCl to two volumes
of the solution imder examination before passing in the HjS will ])revent the
precipitation of the tin while allowing the complete precipitation of the anti-
mony.
4. If the sulphides of As , Sb and Sn are evaporated to dryness with con-
centrated HNO,; the residue strongly fused with Na^CO, and NaOH: and the
cooled mass disintegrated with cold water, the filtrate will contain the arsenic
as sodium arsenate, Na8A804 , and the tin as sodium stannate, NaoSnO^; while
the antimony remains as a residue of sodium pyroantimonate, NajHsSbsOf
(§70, 7).
122 DIRECTIONS FOR ANALYSIS WITH NOTES. §89, 5.
5. Stibine is evolved much more slowly than arsine in the Marsh apparatus,
and some metallic antimony will nearly always be found in the flask with the
tin (§70, 6;).
6. If organic acids, as tartaric or citric, be present, they should be removed
by careful ignition with KzCO, as preliminary to the preparation of the sub-
stance for analysis, since they hinder the complete precipitation of the anti-
mony with HjS (§70, 6c).
7. Antimonic compounds are reduced to the antimonous condition by HI with
liberation of iodine (§70, 6^ and 10). Chromates oxidize antimonous salts to
antimonic salts with formation of green chromic salts (§70, 6h). KMnO^ also
oxidizes antimonous salts to antimonic salts, a manganous salt being formed
in acid solution (§70, 6h). No reaction with antimonic salts. Antimonous
salts reduce gold chloride; antimonic salts do not (§73, 10).
§90. Manipulation. — The contents of the generator of the Marsh appar-
atus should be filtered and washed. The filtrate, if colorless, may be
rejected (absence of Mo). A colored filtrate, blue to green-brown or black,
indicates the probable presence of some of the lower forms of molybdenum.
The solution should be evaporated to dryness with an excess of HNO3 ,
which oxidizes the molybdenum to molybdic acid, M0O3 . The residue
is dissolved in NH^OH (the zinc salt present does not interfere) and poured
into moderately concentrated nitric or hydrochloric acid (§75, 6rf footnote).
This solution is tested for molybdenum by Na2HP04 . The original solu-
tion should also be examined for the presence of molybdenum as molybdic
acid or molybdate (§75, Gd).
The residue from the generator of the Marsh apparatus may contain
Sb , Sn , Au , and Pt with an excess of Zn . It should be dissolved as
much as possible in HCI . Sb , Au , and Pt are insoluble (§70, 5a). The
Sn passes into solution as SnCI^ and gives a gray or white precipitate with
HgCl2 , depending on amoimt of the latter present (§71, 6/) :
SnClj -f- HgClj = SnCl4 + Hg
SnCl, + 2HgCl2 = 2HgCl + SnCl^
The presence of Sn" should always be confirmed by its action in fixed
alkali solution upon an ammoniacal solution of AgNOs, giving Ag°
(§71, 6t).
Au and Pt may be detected in the residue, but it is preferable to precipi-
tate them from a portion of the original solution by boiling with ferrous
sulphate (6^, §§73 and 74). Both metals are precipitated. They are then
dissolved in nitro-hydrochloric acid and evaporated to dry^ness with am-
monium chloride on the water bath. The residue is treated with alcohol
which dissolves the double chloride of gold and ammonium, leaving the
platinum double salt as a precipitate, which is changed to the metal upon
ignition. The alcoholic solution is evaporated, taken up with water and
the gold precipitated by treating with FCSO4 (§73, 6h), by boiling with
oxalic acid (§73, 66), or by treating with a mixture of SnClj and SnCl^
(Cassius' purple) (§73, 6g),
If a portion of the original solution, free from ENO3 , be boiled with
§94,5. DIRECTIONS FOR ANALYSIS WITH NOTES. 123
oxalic acid the gold is completely precipitated as the metal, separation
jfrom the platinum which is not precipitated (§74, 6&).
591. — Notes, — Molybdenum. — /. In the regular course of analysis, molyb-
denum remains in the flask of the Marsh apparatus as a dark colored solution,
the Zn and H2SO4 acting as a reducing agent upon the molybdic acid.
2. If the molybdenum be present in solution as molybdic acid or a molybdate,
it may be separated in the acid solution from the other metals by phosphoric
acid in presence of ammonium salts, forming the ammonium phosphomolyb-
date; insoluble in acids, but soluble in ammonium hydroxide (§75, 6d).
3, In ammoniacal solution of a phosphoraolybdate, magnesium salts precipi-
tate the phosphoric acid, leaving the molybdenum as ammonium molybdate in
solution, which may be evaporated to crystallization (method of recovering
ammonium molybdate from the ammonium phosphomolybdate residues).
§92. Tin. — /. Tin requires the presence of much less HCl to prevent its pre-
cipitation by H3S than arsenic or antimony (§89, 3).
2. The yellow ammonium sulphide (NH4)3Sx must be used to effect solution
if tin (Sn") be present, SnS being practically insoluble in the normal am-
monium sulphide (§71, 5c).
5. Tin in the stannous condition, dissolved in the fixed alkalis (stannites),
readily precipitates metallic silver black from solutions of silver salts. An
arsenite (hot) or an antiraonite in solution of the fixed alkalis produces the
same result, but not if^ the silver salt be dissolved in a great excess of ammo-
nium hydroxide (§70, (ii). This reaction also detects stannous salts in the
presence of stannic salts.
.|. Tin in the Marsh apparatus is reduced to the metal, and then by solution
of the residue in HCl , forms SnCL , which may be detected by the reduction
of HgClz to Hg^Gl or Hg** (§71, 6;), and by the action in fixed alkali solution
upon the strong ammoniacal solution of silver oxide (§71, 6i).
5. If the Zn in the Marsh apparatus is completely dissolved, the Sn must be
looked for in the solution, which in this case must not be rejected. The tin
remains us the metal as long as zinc is present (§135, 10).
6. The presence of the tin may be confirmed by its action as a powerful
reducing agent (§71, 10). If it be present as Sniv , these tests must be made
after reduction in the Marsh apparatus or in an open dish with zinc and HCl.
§93. Gold. — i. Gold will usually be met with in combination with other metals
as alloys, and is separated from most other metals by its insolubility in all
acids except nitrohydrochloric acid.
2. If more than 25 per cent of gold be present in an alloy, as with silver,
the other metal is not removed by nitric acid (§73, rya). Either nitrohydro-
chloric acid must be used or the alloy fused with about ten times its weight of
silver or lead, and this alloy dissolved in nitric acid when the gold remains
behind.
3. If the presence of gold is suspected in the solution, it should be precipi-
tated with FeSO^ before proceeding with the usual method of analysis.
4. If gold be present (in the usual method of analysis) it will remain as a
metallic residue in the Mnr.sh apparatus, insoluble in HCl and may be identi-
fied by the reactions for Au° .
5. The reactions of gold chloride with the chlorides of tin forming Cassius'
purple (§73. Of/) is one of the most characteristic tests for gold.
§84. Platinum. — /. Notes / to .J under gold apply equally well for platinum,
except that it is necessary to hoil with FeS04 to insure complete precipitation
of the platinum.
2. Oxalic acid is the best reagent for the separation of gold from platinum
(§73, 6ft).
3. The most important problems in the analysis of platinum consist in its
separation from the other metals of the platinum ores (§74, 3).
324
TABLE FOR AXALT8IS OF THE COPPER GROUP.
§95.
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§95.
TABLE FOR ANALYSIS OF THE COPPER OROUP.
125
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126 DIRECTIONS FOR AXALT8IS WITH NOTES. §96.
§96. Uanipulation. — The well washed residue after digesting the pre-
cipitated sulphides of the second group (the Tin and Copper Group) in
(1^4)28, may contain any of the metals of the Copper Group, and in
addition frequently contains sulphur, formed by the action of the HoS
upon oxidizing agents : 4FCCI3 + 2HoS = 4FCCI2 + 4HC1 + 83 . Pierce
the point of the filter with a small stirring rod and, with as little water as
possible, wash the precipitate into a test-tube, beaker, or small casserole.
Sufficient reagent nitric acid (§324) should be added to make about
one part of the acid to two parts of water and the mixture boiled vigor-
ously for two or three minutes: *
2Bi,S, -h 16HN0, = 4Bi(N0,), -f- 4N0 -f- 8H3O -f- 3S,
6CdS -f- 16HN0, = GCd(NO,), -f 4N0 -f SH^O + 3S,
Mercuric sulphide is una t tacked (§58, Ge) and remains as a black pre-
cipitate together with some sulphur as a yellow to brown-black precipitate.
The precipitate is filtered and washed with a small amount of hot water.
The filtrate is set aside to be tested later, and the black residue on the
filter is dissolved in nitro-hydrochloric acid : 2Hg8 + 2Clo = 2HgCl2 + S. .
This solution is boiled to expel all chlorine and the presence of mercury
determined by reduction to HgCl or Hg° by means of 8nCl2 (§58, Gg) :
HgCl2 + 8nCl2 = Hg + 8nCl, , 2HgCl2 + 8nCl2 = 2HgCl + SnCi, ; or
by the deposition of a mercury film on a strip of bright copper wire
(§58, 10): HgClj + Cu =z Hg + CuCL . Confirm further by bringing in
contact with iodine in a covered dish: Hg -f L = Hgl^ (Jannaesch, Z.
anorg., 1896, 12, 143). The mercury may also be detected by using
NH4OH and KI as the reverse of the Nessler's test (§207, 6/r) (delicate
1 to 31,000) (Klein, Arch. Pharm., 1889, 227, 73).
§97. Xotcs.—l, The concentration of HNO, (1-2) is necessary for the solution
of the sulphides of Fb , Bi , Cu and Cd , and may also dissolve traces of HgS .
However, the concentrated HNO:, {up. ffr,, 1.42) dissolves scarcely more than
traces of HgS (§58, 6e). Lonjr-continued boilinpr of HgS with concentrated
HNOs changes a portion of the HgS to Hg(NOa),.HgS , a white precipitate,
insoluble in HNO3 .
^. In the use of nitrohydroohloric acid to dissolve the HgS , the HCl should
be used in excess to insure the decomposition of the nitric acid, which would
interfere with the reduction tests with SnCL and Cu** . One part of HNO^
to three parts HCl g^ives about sufficient HCl to decompose all the HNO, .
hence in this reaction a little more than that proportion of HCl should be
used.
3, A small amount of black residue left after boiling the sulphides with
HNO, may consist entirely of sulphur, which can best be determined by
burning the residue on a platinum foil and noting the appearance of the
flame, the odor, and the disappearance of the residue. The residue of sulphur
frequently possesses the property of elasticity (§256, 1).
4, Boiling the sulphides of the copper group with HNO, will always oxidize
♦ If preferred the precipitate on the filter may be washed with the boilinfir hot nitric acid of
the above mentioned Btrengrth, pouringr the same acid back upon the precipitate, reheating each
time, until no further action takes place.
§99,5. DIREGTIOXS FOR ANALYSIS WITH NOTES. 127
a trace at least of sulphur to E^SOf (§256, 6B, 2), which will form FbS04 if any
lead be present:
S, + 4HN0, = 2H2SO« + 4N0
3PbS + 8HN0, = aPbSO^ + 4H3O + 8N0
If the boiling be not continued too persistently, the amount of PbS04 formed
is soluble in the HNO, present (§57, 5o), and does not at all remain behind
with the HgS .
5. If the Sb and Sn are not removed, through an insufficiency of (NHJjSx
they will appear as a white precipitate mixed with the black precipitate of
HgS , due to the fact that HNO, decomposes the sulphides of Sb and Sn ,
forming the insoluble SbjOs and SnOz :
6Sb,S, + 20HNO, = 6Sb,0, -f OS^ -f 20NO -f- IOH3O
6. Traces of mercury may be detected by using a tin-gold voltaic couple.
The Hg deposits on the Au , and can be sublimed and identified with iodine
vapor. Arsenic gives similar results (Lefort, C. r., 1880, 90, 141).
7. Mercury may quickly be detected from all of its compounds by ignition
in a hard glass tube with*^ fusion mixture (Na^CO, -f K,CO,) (§58, 7), and then
adding a few drops of HNO, (concentrated) and a small crystal of KI . Upon
warming the iodine sublimes and combines with the sublimate of Hg , forming
the scarlet red Hgln . As and Sb both give colored compounds with iodine, de-
composed by HNO." (Johnstone, C. N„ 1889, 59, 221).
§98. Manipulation. — To the filtrate containing the nitric acid solution
of the sulphides of Pb , Bi , Cu , and Cd , should be added about two cc. of
concentrated H2SO4 and the mixture evaporated on a sand bath or over
the naked flame in a casserole or evaporating dish until the fumes of
H2SO4 are given off:
Pb(NO,), -f H2SO, = PbSO, -h 2HNO3
Cu(N03), -h H2SO, = CuSO, + 2HNO3
About 20 cc. of 50 per cent alcohol should be added to the well cooled
mixture and the whole transferred to a small glass beaker. Upon giving
the beaker a rotatory motion the heavy preci])itate of PbSO^ will collect
in the center of the beaker, and its presence even in very rmall amounts
may be observed. The filtrate from the PbS04 should be decanted through
a wet filter, and the PbSO^ in the beaker may be further identified by its
transference into the yellow chromate with KgCrO^ or into the yellow iodide
withKI (67, 6/ and 70.
§09. Notes, — 1. In analysis, if lead was absent in the silver group, it is
advantageous to test only a portion of the nitric acid solution with H0SO4 for
lead, and if that metal be not present, the above step may be omitterl with
the remainder of the solution and the student may proceed at once to look
for Bi , Cu and Cd . Tf, however, lead is present, the whole of the solution
must be treated with H2SO4 .
2. The nitric acid should be removed by the evaporation, as PbSO^ is quite
appreciabh' soluble in HNO, (§57, *>r)^
.i. The H3SO4 should be present in some excess, as PbSO^ is less soluble in
dilute H2SO4 than in pure water (§57, 5e).
jj. Alcohol should be present, as it greatly decreases the solubility of PbSO,
in water or in dilute H2SO4 (§57, 5r, 6e).
.5. Too much alcohol must not be added, as sulphates of the other metals
prv'sent are also less soluble in alcohol than in water (§77, 5c). These sul-
phates, if precipitated by the alcohol, are readily dissolved on dilution with
water.
132 IRIDIUM. §106, :.
color is produced, or the precipitate separates after warming. At a red heat,
the precipitate is decomposed.
Palladous nitrate gives most of the above reactions; no precipitate with
ammonia, and a less complete piecipitate with iodides.
7. Ignition. — Nearly all the palladium compounds are reduced by heat, before
the blow-pipe, to a ** sponge." If this be held in the inner flame of an alcohol
lamp, it absorbs carbon at a heat below redness: if then removed from the
flame, it glows vividly in the air, till the carbon is all burnt away (distinction
from platinum).
8. Detection. — Palladium is precipitated with the second group metals by H2S.
not dissolved by (NH4)2Sx (separation from the tin group). It is distingiiished
from mercury by its precipitation as a cyanide with mercuric cyanide. It is
precipitated from quite dilute solutions by KI (distinction from Bi and Cd);
an excess of the KI dissolves the black palladous iodide, Pdlj , to a dark brown
solution. XCNS does not precipitate palladium salts, not even after the addi-
tion of SO2 (separation from Cu). The addition of H2SO4 and alcohol separates
lead from palladium. The presence of the metal should be further confirmed
by reduction and study of the properties of the " sponge " obtained.
9. Estimation. — (/) As metallic palladium, to which state it is reduced by
mercuric cyanide or potassium formate, and ignition, first in the air and
then in hydrogen gas. (2) As XsPdClfl . Evaporate the solution of palladic
chloride with potassium chloride and nitric acid to dryness, and treat the mass
when cold with alcohol, in which the double salt is insoluble. Collect on a
weighed filter, dry at 100**, and weigh.
10. Oxidation. — Palladium is reduced as a dark-colored precipitate, from all
compounds in solution, by sulphurous acid, stannous chloride, phosphorus, and
all the metals which precipitate silver (§59, 10). Ferrous sulphate reduces
palladium from its nitrate, not from its chloride. Alcohol, at boiling heat,
reduces it; oxalic acid does not (distinction from gold §73, Ob),
§107. Iridinm. Ir = 193.1 . Usual valence three and four.
1. TropertieB.— Specific gravity, 22.421 (Deville and Debray, C. r., 1875, 81, 839).
Melting point, 1950** (Violle, O. r., 1879, 89, 702). When reduced by hydrogen it
is a gray powder, which by pressing and igniting at a white heat changes to a
metallic mass capable of takir.g a polish. It is used mostly as an alloy with
platinum, forming a very hard, durable material for standard weights and
measures. A platinum-iridium dish containing 25 to 30 per cent iridium is not
attacked bN' nitrohydrochloric acid.
2. Occurrence. — Found in platinum ores, usually as an alloy with platinum
or osmium.
3. Preparation. — The platinum residues are mixed with Pb and PbO and
heated at a red heat for one-half hour, then treated with acids. The residue
contains the iridium as osmium-iridium or platinum-iridium with other plat-
inum metals. This residue is mixed with NaCl in a glass tube and heated to
a red heat in a current of chlorine. Much of the osmium passes over as the
volatile perosmic acid, and is condensed. The d(mble sodium chlorides of Ir ,
Os , Rh , Pt , Pd and Ru are dissolved in water filtered and, when boiling hot,
decomposed by H.S . The iridium is reduced from the tetrad to the triad, but
is not precipitated until after all the other metals. By stopping the current of
HaS just as the brown iridium sulphide begins to form, a complete separation
can be made by filtration. By rf»crystallization the pure sodium double salt,
6NaC1.2lrCl3 -f 24H3O , is obtained, which is changed to the tetrad ammonium
double salt, (NH4)2lrCl, , by the addition of NH4CI and oxidation with chlorine
(Wiihler, Pogg., 1834, 31, 161). This upon ignition gives the pure metal as
iridium sponge. Or, the double sodium salt is ignited with sodium carbonate
exhausted with water and reduced by ignition in a current of hydrogen, leav-
ing the metal as a fine gray powder (See also §106, 3).
4. Oxides and Hydroxides. — Iridium forms two series of oxides and hydrox-
ides, the metal acting as a triad and tetrad respectively. IrOj is formed by
§108, 5. OSMIUM, 133
ig'niting the metal in the air at a bright red heat, henoe the scaling of platinum
dishes which contain iridium. The hydroxide, Ir(0H)4 » is formed by boiling a
solution of the trichloride, IrCl, , in a fixed alkali hydroxide or carbonate.
Careful addition of KOH to IrClj in a vessel full of liquid and closed to exclude
air gives Ir(OH), , easily oxidized to Ir(0H)4 (Claus, J. pr„ 1846, 39, 104).
5. Solubilities. — Freshly precipitated iridium may be d^ -solved in nitrohydro-
<jhloric acid. The ignited metal is insoluble in all acids. Its proper solvent is
chlorine. Iridium trichloride, IrCl, , is soluble in water and forms with the
alkali chlorides double chlorides, soluble in water, insoluble in alcohol. The
tetrachloride with sodium chloride, NasIrCla , is formed when the platiniinr*
residues mixed with NaCl are heated in a current of chlorine. It is soluble in:
•water. The corresponding ammonium salt may be formed from the sodium:
salt by precipitation from the concentrated solution with NH4CI , a reddish-
brown precipitate, soluble in 20 parts of water (Vauquelin, A, Ch.^ 1806, 50, 150
and 225). Tlie potassium double salt is sparingly soluble in water.
6. Beactions. — Fixed alkali hydroxides or carbonates precipitate from boil-
ing solutions of iridium chloride, IrCl, or IrGl4 , iridium hi/droTide, Ir(0H)4 ,
dark blue, insoluble in all acids except HCl . Potassium nitrite added to a hot
solution of iridium salts gives, first a yellow color and finally a yellow precipi-
tate, insoluble in water or acids. Hydrogen sulphide reduces irCl4 to IrCl, ,
and then precipitates the trisulphide, Ir-jS, , brown, soluble in alkali sulphides.
7. Ignition. — When iridium is fused with potassium acid sulphate it is oxid-
ized, but does not go into solution (difference from rhodium, §105, 7). Ignition
on charcoal reduces all iridium compounds to the metal. Fusion in the air
with sodium hydroxide or with sodium nitrate causes oxidation of the metal,
the iridium oxide formed being partially soluble in the fixed alkali.
8. Detection. — See 3 and 0.
9. Estimation. — It is converted into the oxide by igniting with KNO, and
then reduced by ignition in an atmosphere of hydrogen.
10. Oxidation. — Formic acid (from hot solution), zinc and HaSOf or HCl
reduce iridium compounds to the metal. SnClj , FeSO^ and H3C2O4 reduce
tetrad iridium to triad, but do not further reduce (separation from gold, §73,
6^, h and h).
§108. Ofanium. Os = 191.0 . Valence two to eight.
1. Properties. — ffpecific grnrity, 22.477, the heaviest of all bodies (Deville and
Debray, C. r., 1876, 82, 1076). In the absence of air it may be heated above the
vaporization point of Ft without melting or oxidizing. In presence of air,
when heated a little above the melting point of Zn , it bums to the volatile
poisonous perosmic acid, OSO4 . In compact form it is very hard, cutting glass,
and possesses a metallic lustre, with a bluish color rescmlDling Zn.
2. Occurrence. — Always present in the residues of the platinum ores, in com-
bination with iridium.
3. Preparation. — The iridium osmium alloy or other Os containing material
is finely divided and distilled in a current of chlorine or with nitrohydrochloric
acid, the osmium passes into the receiver containing KOH . Wy repeated
additions of HNO3 and further distillation, the osmium may all be driven into
the receiver. The distillate is treated with HCl and Hg and the amalgam
ignited in a current of hydrogen (Berzelius, Pogij,y 1829, 15, 208).
4. Oxides,— Osmium forms five different oxides, OsO , OSjO, , OsOj , OsO, ,
O8O4 . The first three are bases, the salts of which have been but little
studied; OsO, forms salts with bases, and OsO* acts rather as an indifferent
peroxide. Perosmic acid, OsO* , exists as white glistening needles, melting
under 100**, sparingly soluble in water, its solution having a very penetrating
odor, resembling that of chlorine. The fumes of the acid are very poisonous,
and cause inflammation of the eyes. HjS is recommended as an antidote
(Clauss, A„ 1847. 63, 355).
5. Solubilities. — The metal in compact condition is not at all attacked by any
acid. The precipitated metal is slowly dissolved by nitrohydrochloric or fum-
134 TUN08TEN. §108, 6.
ing nitric acid. By heating the metal in a current of chlorine a mixture of
OsCls and O8CI4 is formed. They are both unstable.
6. Reactions. — Perosmic acid, O8O4 , when boiled with alkalis, is reduced to
osmates, as XjOsO^ . A solution of perosmic acid decolors indigo, oxidizes
alcohol to aldehyde, and liberates iodine from potassium iodide. In the pres-
ence of a strong mineral acid, HaS precipitates osmium sulphide, OsS« , brown-
ish black (Glaus, J, pr., 1860, 70, 28); insoluble in alkali hydroxides, carbonates
or sulphides.
7. Ignition. — Osmium when heated on a piece of platinum foil gives an in-
tensely luminous flame of short duration. By holding the foil in the reducing
flame and then again in the oxidizing flame, the luminosity may be repeated.
If a mixture of the metal or of the sulphide and potassium chloride be heated
in a current of chlorine, a double salt of potassium osmic chloride is formed,
sparingly soluble in cold water, more readily in hot water. Alcohol precipitates
it from its solutions as a red crystalline powder.
8. Detection. — By the intensely luminous flame when ignited on a platinum
foil; by oxidation and distillation as perosmic acid and identification by odor,
action on indigo and on potassium iodide.
9. Estimation. — It is weighed as the metal (see 3).
10. Oxidation. — OSO4 is reduced to OsO, by ferrous sulphate. Zn and many
other metals in presence of strong acids precipitate the metal. The metal is
also obtained from all osmium compounds by ignition in a current of hydrogen.
§109. Tungsten (Wolframium). W = 184 . Valence two to six.
1. Properties.— Spefi/fc gravity, 19.129 (Roscoe, A., 1872, 162, 359). A tin-white
or steel-gray metal, brittle, harder than agate. That precipitated from acid
solutions is a velvet-black powder. Non-magnetic. Stable in the air at ordi-
nary temperature; burning at a high temperature, it decomposes steam at a
red heat.
2. Occurrence, — Tungsten does not occur in nature in large amounts, nor is
it widely disseminated. The moat common tungsten minerals are scheelite,
CaWO^ , and wolframite, FeW04 and MnWO* , in variable proportions. It never
occurs native.
3. Preparation. — By reduction of WO, in H at a red heat (Zettnow, Pogg,^
1860, 111, 16); by ignition of WO, and Na under NaCl . Tungstic r.cid of
commerce is prepared by igniting for several hours: 100 parts NajCO, , ignited;
150 parts finely ground wolframite; and 15 parts NaNO, . The cooled mass is
exhausted with water and the filtrate poured into hot, moderately concentrated
HCl (Franz, J. />r., 1871, (2), 4, 238).
4. Oxides. — WO3 is obtained as a brown powder by decomposing WCI4 with
water (Roscoe, I.e.). WO, is a lemon-yellow, soft powder, insoluble in water
or acids. It is formed by ignition of the metal, lower oxides or decomposable
salts in the air. The blue tungsten oxides are compounds between WOa and
WO,.
5. Solubilities.— The metal is scarcely at all attacked by HCl or H,S04 , slowly
by HNO, or nitrohydrochloric, slowly soluble in alkalis. The halogens com-
bine directly upon heating. WO, is readily soluble on heating with HCl and
H3SO4 to a red color. It is also soluble in KOH with red color, evolving
hydrogen. Both the acid and alkaline solutions deposit the blue oxide on
standing (von der Pfordten, A., 1884, 222, 158). WO, is insoluble in water or
acids, not even soluble in hot concentrated H3SO4 . Soluble in KOH , K.CO,
and NH4OH . In an atmosphere of COj it reacts with the chlorides of Ca ,
Mg, Co, Ni and Fe (not with those of Pb , Ag , K and Na), e.g., MCI, -|-
2 WO, = MWO4 -h WO2CI2 . Heated with chlorine, WO2CI, is formed, and also
WCI4 , decomposed by water. S , HjS or HgS form WS, on heating with WO, .
Soluble alkali tungstates are formed by fusion of the acid, WO, , with the
alkali metal carbonates, more slowly by boiling with the carbonates. Acids
form, from solutions of the alkali tungstates, a white precipitate of the
hydrated acid turning yellow on boiling, insoluble in excess of the acids (dis-
§110,5. VANADIUM. 135
tinction from MoO,), soluble in NH4OH. Phosphoric acid changes tungstic
acid to the metatungstic acid, which is soluble in water and not precipitated
by other acids. Long boiling of the solution of metatungstic acid causes the
precipitation of tungstic acid. Fusion of WO, with XHSO4 gives a compound
of potassium tungstate and tungstic acid, not readily soluble in water but very
readily soluble in (NH4)aC0, (distinction from silica, §249, 5).
6. Beactions. — Solutions of salts of Ba , Ca , Eb , Ag and Hg produce w-hite
precipitates with solutions of alkali tungstates. H^S precipitates WB.
from acid solutions, the sulphide dissolving readily in (NH4)2S , forming a
thiotungstate (NH4)2WS4 . The tungstates, like the molybdates, form complex
compounds with phosphoric acid, i. f., phosphomolybdates and phosphotung-
states, which react very similarly with ammonium salts and with organic bases
(§75, Gd), K4re(CN)e gives with tungstates (in presence of acids) a deep
brownish-red fluid, forming after some time a precipitate of the same color.
Solution of tannic acid gives a brown color or precipitate.
7. Ignition. — With NaPO, , WO, dissolves, on fusion, to a clear or yellowish
bead in the oxidizing flame; in the reducing flame it has a blue color, changing
to red on addition of FeS04 . Heated on charcoal in presence of NaaCO, with
the blow-pipe, using the reducing flame, the metal is obtained.
8. Detection. — If a tungstate be fused with Na::C03 , the mass warmed with
water and the water then absorbed with strips of filter paper, the tungsten
may be detected by moistening the strip with HCl and warming, obtaining the
yellow color of WO,; and the blue color of a lower oxide by moistening with
SnCl, and warming. (NH4)3S does not color the paper, even after adding HCl ,
but on warming a blue or green color is obtained.
9. Estimation. — It is converted into WO, and weighed as such after ignition.
10. Oxidation. — WO, gives with SnCl, , or Zn in presence of HCl or H2SO4 ,
a beautiful blue color, due to the formation of oxides between WO, and WO, ,
blue oxides of tungsten (delicate and characteristic).
§110. Vanadium. V = 51.4 . Valence two to five.
1. Proi>ertie8. — 'Specific gravity, 5.5. A grayish non-magnetic powder; slowly
oxidized in the air, rapidly on ignition with formation of VjO, . It forms with
chlorine the dark brown tetrachloride.
2. Occurrence. — It is often found in iron and copper ores and in some clays
and rare minerals, e.g„ vanadinite, SPbjVaOs -f PbCl,; volborthite, (Cu.Ca)8V,0g;
mottramite, (Cu.Pb)5VaOio.2H20; etc.
3. Preparation. — The vanadium ores are treated chiefly for the preparation of
ammonium vanadate and vanadic acid. The ores are fused with KNO, , form-
ing potassium vanadate. This is precipitated with Pb or Ba salts and then
decomposed with H2SO4 . The vanadic acid is neutralized with NH4OH and
precipitated with NH4CI , in which it is insoluble. This upon ignition gives
VjO, pure (Wohler, A., 1851, 78, 125). The metal is prepared from the dichlo-
ride, VCl, , by long-continued ignition in a current of hydrogen.
4. Oxides. — Vanadium forms four oxides: VO , gray; V3O, , black; VO, , dark
blue; and VjOs , dark red to orange red.
5. Solubilities. — Vanadium is not attacked by dilute HCl or H3SO4; concen-
trated H2SO4 gives a greenish-yellow solution: HNO, a blue solution. VO dis-
solves in acids to a blue solution with evolution of hydrogen. VjO, dissolves
in dilute HCl to a dark greenish-black solution. Chlorine forms with VjO, ,
VOCl, and VjO, . VO3 dissolves in acids to a blue solution, from which solu-
tions NajCO, gives a precipitate of V203(OH)4 -+- 5H3O , grayish-white mass,
losing 4H3O at 100** and turning black, soluble in acids and alkalis. V2O5
exists in several modifications with different solubilities in water, the red
modification being soluble in 125 parts of water at 20° (Bitte, C. r., 1880, 101,
698). Vanadic acid forms three series of salts, ortho, meta and pyro, analogous
to the phosphates. Most salts are the metavanadates. The ortho compounds
are quite unstable, readily changed to the meta and pyro compounds. Alkali
vanadates are soluble in water, the ammonium vanadate least soluble and not
at all in NH4CI .
136 GERMAN IV M. §110, G-
6. Beactions. — Solutions *of vanadic acid produce brown precipitates with
alkalis, soluble in excess to a yellowish-brown color. Potassium ferrocyanide
gives a gfreen precipitate, insoluble in acids. Tannic acid gives a blue-black
solution, which is said to make a desirable ink. Ammonium sulphide precipi-
tates VaSft , brown, soluble with some difficulty in excess of the reagent to a
reddish-brown thio salt. From this solution acids reprecipitate the brown
vanadic sulphide, V^Sb .
If to a solution of a vanadate, neutral or alkaline, solid NH4CI be added, the
vanadium is completely precipitated as NH4VOS , ammonium metavanadate,
crystalline, colorless, insoluble in NH4CI solution; upon ignition in air or oxy-
gen, pure vanadic oxide, V,Ob , is obtained.
7. Ignition. — Borax gives with vanadium compounds in the outer flame a
colorless bead, yellow if much vanadium be present; in the inner flame a green
bead, or brown when vanadium is present in large quantities and hot, becoming
green upon cooling. All the lower oxides of vanadium ignited in air or
oxygen give VaO» .
8. Detection. — Vanadium will almost always be found as a vanadate (2) and
is detected by the reactions used in its purification (3); also by the reactions
with reducing agents, forming the colored lower oxidized compounds (10).
9. Estimation. — (i) It is precipitated as basic lead vanadate and dried at
100°. (2) It is precipitated as ammonium vanadate, KH4VO8 , in strong
KH4CI solution, ignited to the oxide VaO. , and weighed.
10. Oxidation. — ^Zn , in solutions of vanadates with dilute H^SO^ , reduces the
vanadium to the tetrad, a green to blue solution, then greenish-blue to green,
the triad, and finally to lavender blue, the dyad. H3S reduces vanadates to the
tetrad with separation of sulphur. Oxalic acid and sulphurous acid also reduce
vanadates to the tetrad, the solution becoming blue.
§111. Germanium. Ge -= 72.5 . Valence two and four.
1. Properties.— »Sf/)ert/fo gravity, 5.469 at 20.4°; melting point, 900° (Winkler,
J, pr,y 1886, (2), 34, 177). A gray-white crystalline metal. Fused under borax
it gives a grayish-white regulus with a metallic lustre. It is stable in the air,
volatilized at a high heat (Meyer, B., 1887, 20, 497), and is easily pulverized.
It bums in oxygen to form germanic oxide, GeO, .
2. Occurrence. — It is found in small quantity in argyrodite, a sulphide of
silver and germanium, SAgjS -f- GeS, , a silver ore from Freiburg, Saxony.
It is also found in euxenite from Sweden (Kriiss, C. (7., 1888, 75).
3. Preparation. — It is formed by reduction of the oxide, GeO, , with H , C
or Mg (Winkler, B., 1891, 24, 891); also by reduction of the sulphide in H .
4. Oxides. — It forms two oxides, GeO and GeO, . To prepare pure GeO^ , the
mineral argj-rodite is pulverized and intimately mixed with equal weights of
KaaCO, and S and heated to a good full ignition. The mass must be added
carefully to prevent foaming. The fused mass is exhausted with HjO , the
germanium going into solution as a thiosalt. With a decided excess of H2SO4 ,
the sulphide is completely precipitated. The precipitate is now dissolved in
KOH , the sulphides of Ag , On and Pb remaining undissolved. By adding to
the KOH solution H2SO4 not quite to neutralization, the As and Sb sulphides
are precipitated on boiling, while the GeS remains in solution with some
A83S,; HjS is carefully added to the solution until the As^S, is all precipitated,
then the filtrate is made strongly acid with H2SO4 , and the solution evaporated
till SO, fumes escape. The mass is dissolved in hot water, and upon cooling
GeOa crystallizes out (Winkler, I, c).
5. Solubilities. — Germanium is insoluble in HCl , soluble in nitrohydrochloric
acid as GeCl^ , and oxidized with HNO, to GeO, . Hot concentrated HjSO^
evolves SO, and forms 06(804)3 . Insoluble in KOH solution but dissolves
with incandescence in fused KOH. It unites directly with CI, Br and I
(Winkler, /. c). Germanic oxide, GeO, , is a white powder, very sparingly
soluble in water or acids. Fused with fixed alkali hydroxides or carbonates it
is converted into compounds soluble in water. GeCl* is a liquid, boiling at 84°;
SU8,5,
TELURHM,
137
St \% decompo&etl by water. If a soliittcin of the oxide in excess of HCl \m
fiiiporated to dryness the Oe U till vokitUizeiJ. GeS, m solabjc! in '222 pfirta
water* in alkali Hulphitle^^ and hydroxides; insoluble in HCl or H^SO, » whieh
preeipitate it from it^ solutions; soluble in nitrohydroehlorio aeitl with &eptira-
r»uu *if sulphnr. Xiirie oxide changes it to G«Oj with sepunition of sulphnr,
6. B«actioB». — Germanium jsait^ give aimoai no chynu*Trrisrif n.^yrtions willi
th*^ various ren|rents. HjS pri'L-ipitates gerinanle sulphide* GeSj * white, from
solulions of the aults quite strongly acid* The sulphide m soluble In ammonium
Mutphide, forming* a tMo suit, iXiun phiciug- G« in division A of the second ^roup.
7, I^ition, — Heated tj**fore the blowpipe in the redneing- flame withonl an
alkali ut* flux the metal is formetl, and at the same time a white coating of
the oxide. It forms a colorless bead with borax.
8- Detection.— In the niineral. argyrodite, by heating- in an atmosphere of
HjS or lllnminatJiipr ff^s, an oran^^e-yellow sublimate is obtained, whieb may be
examined under the mieroscope and in the wet way (Paushofer* V. f\, 1888»
9, Estimation.— It is converted into the sulphide, 0«Sa , and then heated
viith HNOj and weighed hk tHQ, .
10, Oxidation, — Zn in ncid solutions of Qe salts pt»edpitateR the metal as a
dark lirowu slime. If GeS^ is healed in a current of H, deS ii at first formed
with H,S, finally Qe^
§112* TeUurinm. Te = 127,5? Valence two, four and poagibly six,
1, Troperti^B.—Sprriflr fjmrttt/, a.244r> (Borsielius, Pngff., \B3A, 32* 1 and r»77).
Mdtlmj jmittt, 4.V2' (Carnelley and Williams, J, C* 18S0, 37* 125). Te is t^rystal-
line* silver white, brittle* stable in the air and in boiling- water; heated in the
nir, i! burns w*ith a i^reenish flame. In its (general properties and reactions it
s»tiitidsi closely related to S and Be (2).
2. Occtirrence.'-rti few^ plftees and in small quantities in Germany* Mexico,
Bolt via, Tnited Stsites and .la pun. Some of the minerals are: tellurite, TeOj*
tetrmlymite, 231^X6^31,^^: ferroteJlLrite* FeTeOi , etc. It also occurs native*"
X Preparation.— ( rj Fusion with ulknli rarbunate and C, whicli converts it
Into a tellurfde, as Na.Te; thi?n solution in (air free) wuter. the air being
exelndi*d us much as po^ssildt, and the filtrate precipitated by passing air
through the solution. Tlic Te is precipitated as a priiy metaliic" powder, con-
taining w'lat Se may have been preseni. (ij Conversion into TeCI^ by distilla*
lion in n current of chlorine, deeompositiim of the chloride with water to
H;TeO, and precipitatirm of tht* Te with KHSO, . iTj From lead chamber
rcaie by digestion with Na.COa und KCN * foritiin|^ KCKTe * The decanted
i^olntion is acidified with HNO^ and thr" Te precipitati-d with H^S (Schimose^
r, K*, I'SH-I* 49, 157). (j) For purification of the commercial Te , see Brauner
iM., iSB% 10, 411) and Schimosc (T. A\, l^S4, 49, 25, and 1S85* 51, 199).
4. Oxides and Hydroxides.— TeO is j^nid to be formed by heatinj:^ TeSO. in a
vacnuni abovr ISfP: TeSO, ^ TeO -f SO. (Divers and Schi'mose, C. X., IHKi, 47,
221)* TeO, forms when Te U buru'^Hl in the air* and when TeCI^ i«s decomposed
by boiling" water. It t?^ a white crystalline solid, sparing-ly soluble in H^O ,
more sohd^Ie in acids from which solutions wiitcr causes a %vhite precijiitate of
TeOi or HiTeO, . H,.TeO„ is formed when a HNO^ solution of Te is inmiediritely
pt>iired into eold watt r. warming to 40^* changes it to TeO^ , HgTeO^ is made
hy fusiuR- TeO. with KNOjj , treating the E^TeO^ so obtained with soluble lead
or bariviiTi salt and decomposing this salt w^ith H.SO^ or H^S * colorless crystals,
insolulde in alcohol or ether-alcohol (separation from HhSO^). It can be
recrvstalllzed from water and upon heating- forms TeO, (Clarke, Am. S., 1S77,
114,' 281; 1S78, 110* 401).
5, Solubilities.— Te is insoluble in HCl: HlfOa and nitro\vdrochlone acids
oxidi^ce it to H,TeO,: in H.SO* it becomes H.TeO, with evolution of SO, (lliljrer,
J,, 1S74, 171, 211): soluble in warm concentrated solution of KGN', from which
Roluti*oi HCl precipit.ites all the Te . H,TeO, is fairly stduhh- in water, red-
dens moist litmus paper and easily decomiioscii into TeO; and H4O . Acid solu-
138 SELENIUM. §112, G.
tions of TeO, are precipitated upon addition of water or upon standing. TeO,
and HaTeO. form soluble alkali salts with the alkalis from which solutions of
the other metallic salts precipitate the respective tellurites. HsTeO^ is soluble
in water, acids and alkalis; alkali carbonates form acid tellurates, less soluble
than the corresponding normal salts. Solutions of the alkali tellurates form
insoluble tellurates with soluble salts of the other metals, c. at., KsTeO. 4-
BaCl, = BaTeO, + 2KC1 .
6. Beactions. — Tellurium is classed with second group metals because of its
precipitation from solutions of tellurites and tellurates by H,S . The precipi-
tate is not a sulphide, but is T« mixed with varying proportions of S , for CS,
removes nearly all the sulphur (Becker, A., 1876, 180, 257). In appearance the
precipitate of Te with HsS very much resembles SnS, and is very soluble in
(NHJ.S.
At a high temperature Te and H unite directly, forming HjTe (Brauner, ^f.,
1889. 10, 446). HjTe is best prepared by heating together Te and Fe or Zn and
decomposing these tellurides with HCl (analogous to the corresponding reac-
tions with sulphur, §257, 4). A colorless gas, odor similar to HsS , bums with
a blue flame, fairly soluble in water and is precipitated as Te° from its solution
by the oxygen of Ihe air. HjTe precipitates solutions of metallic salts very
similarly to HjS and H^Se .
7. Ignition. — Te combines on ignition with most metals to form tellurides.
TeO, ignited, decomposes into TeO, and O . All lower Te compounds ignited
with XNO, give XjTe04 . All Te compounds give on charcoal with the blow-
pipe a white powder, which colors the reduction flame green and disappears.
Heated in an open glass tube, Te compounds give a sublimate of TeOj , which
melts upon heating. Te compounds fused with XCN in a current of hydrogen
form potassium tellurocyanate, XCNTe; soluble in water but precipitated by a
current of air as Te** (distinction and separation from Se). Heated with Na^COa
on charcoal Te compounds give NazTe , which blackens silver with formation
of AgzTe .
8. Detection. — By reduction to Te** and solution in cold concentrated H.SO«
to a purplish-red solution (characteristic). Separated from Se by fusion with
XCN in a current of hydrogen and precipitation from the solution by a current
of air.
9. Estimation. — The Te compound is heated in a current of CI , TeCl* being
sublimed. This is decomposed by water to TeOz , which is reduced to Te" by
SO, and weighed as such after drying at 100** .
10. Oxidation. — Hydrogen at a high temperature reduces Te compounds to
HzTe . H.S reduces Te compounds to Te° mixed with S . Fusion with XNO,
oxidizes all Te compounds to X2TCO4 . SO2 reduces Te compounds to Te"* .
SnClj and Zn in acid solutions give with Te compounds a black precipitate
of Te° . Te compounds warmed with dextrose in alkaline solution are reduced
to Te** . Tellurates boiled with HCl evolve chlorine and are reduced to H^TeOj .
which precipitates as TeOj on adding water if too much HCl be not present
(distinction from Se).
§113. Selenium. Se = 79.2 . Valence two and four, possibly six.
1. "Properties,— Specific gravity^ of the red variety, 4.259; of the black variety,
4.796 (Schaffgotsch, J. pr., 1848, 43, 308). It begins to soften between 40'» and
50°; it is half fluid at about 100°, but is not completely molten until 250''
(Draper and Moss, C\ A'., 1876, 33, 1). The molten Se does not become com-
pletely solid until cooled to 50°. Selenium with tellurium is closely related to
sulphur, and like sulphur exists in amorphous forms (§256, 1). The precipi-
tated Se is red. The brown or brown-black powder obtained by quickly cool-
ing from the molten state is insoluble in CS-, . Boiling point, 676° to 683 =*
(Carnelley and Williams, C. N., 18;t9, 39, 286).
2. Occurrence. — Tn no place abundantly: never native. It is found in com-
bination with minerals in the Hartz Mountains, Sweden, Argentine Republic and
Mexico (Billandot, C, N,, 1882, 46, 60). It occurs in very small quantities vnih.
some sulphides of Fe , Cu and Zn .
§118, 10. SELENIUM. 139
3. Preparatioxi. — In the lead chambers of the HsS04 works it is found as a
red deposit with some S , ASjO, , Sb^O, , PbS04 , etc. The scale is washed with
water and digested with XCN solution at 80° to 100°, until the red color entirely
disappears. The filtrate is then treated with HCl , which precipitates the Se .
It is further purified by oxidation to BeO^ . sublimed and then reduced with
SO, (Nilson, B., 1874, 7, 1719).
4. Oxides and Hydroxides. — HoSeO, is prepared by oxidizing Se with HiSTO, ,
or nitrohydrochloric acid. H.SeO, evaporated to dryness gives HjO and SeO, ,
crystalline. SeO, is also formed by burning Se in air or oxygen; it has an
odor similar to decaying radish. It sublimes at about 200° as a yellow vapor,
condensing to white needles on cooling. SeO, is not known. H3Se04 , pure,
is a white crystalline mass, melting at 58°. H3Se04.H:20 is crystalline at — 38°,
and if recrystallized melts at 25°. The selenic acid usually obtained is a thick
oily liquid, resembling HjSO^ and containing about 95 per cent H3Se04 . It is
obtained by fusing Se or SeOj with XNO, and precipitation of the X2Se04 with
soluble salts of "Bk , Pb , Ca or Cu and decomposing the washed precipitates,
suspended in water, with H2SO4 or H3S .
5. Solubilities. — Se dissolves in cold concentrated H3SO4 to a green colored
solution without oxidation (dilution with water precipitates the Se); if the
solution be warmed SO2 is evolved and the green color disappears (dilution
with water gives precipitate), the Se being oxidized to SeO, . HNO. and nitro-
hydrochloric acid oxidize it to SeOj . Selenous oxide, SeOj , is soluble in water
in all proportions, forming HsSeO, . The selenites and selenates of the alkaline
earths are insoluble and may be formed by adding a solution of the metal to
an alkali selenite or selenate, e. g„ NajSeO. + BaClj =: BaSeO, + 2NaCl . Many
of the selenites are soluble in excess of HjSeO, . Selenates are less stable
than selenites. BaSeO^ is soluble in HCl (distinction and separation from
BaS04) and upon long-continued boiling is reduced to BaSeO, .
6. Beactions. — Selenous acid precipitates with HjS a mixture of Se and S ,
lemon yellow, bright red upon heating (Divers and Shimose, C. ^., 1885, 61,
199). This mixture is soluble in (NH4)5S, hence in qualitative analysis Se is
classed among the metals of division A, second group, while because of its
general properties it belongs with sulphur. When Se and H are heated to-^
gether they begin to combine directly at 250°, forming H3Se (Ditte, 0. r., 1872,
74, 980); which in practically all its reactions is similar to H:.S . H.Se is also
formed by treating K.Se . FeSe , etc., with dilute HCl or H.SO^ ; HNO, gives
HjSeO, with selenides. H^Se is a colorless gas. odor similar to HjS but more
penetrating. It is more poisonous than HnS , burns when ignited, combines
slowly but completely with Hg° , evolving hydrogen. It dissolves in water to a
greater extent than H2S , reacting acid and depositing red flakes of Se on
standing. It precipitates the selenides of the metals having almost the same
' solubilities as the corresponding sulphides (von Reeb, J, Pharm., 1869, (4), 9,
173). With soluble sulphites H.Se gives a precipitate of a mixture of Se and S .
7. Igpiition. — When Se or compounds of Se are fused with KCN in a current
of hydrogen, potassium selenocyanate, KCNSe , is formed. Long boiling with
HCl separates the Se , but this does not take place on exposure of the solution
to the air (separation from tellurium). Selenium compounds heated on char-
coal with Na.CO, are changed to NajSe , which yields a black stain with Ag°
and H2Se with dilute acids.
8. Detection. — If in solution as selenites it is precipitated with H2S (soluble
in (NH4)2S); oxidized to SeO, and obtained as the white needles by sublima-
tion, and reduced from its solution in water to the red Se° by SO, . If present
as selenides, decomposed by HCl or H3SO4 , forming HjSe , which is conducted
into water and the Se° precipitated by passing air or oxygen through the solu-
tion.
9. Estimation.— Oxidized to selenic acid and precipitated as BaSeO^ and
weighed as such. If BaS04 be present the precipitate is reduced in H , and
the resulting BaSeO, separated by solution in HCl . Selenides are heated in a
current of chlorine in a hard glass tube, being converted into SeCl4 , which
vaporizes and is decomposed in water; continued chlorination of the water
solution forms H2Se04 .
10. Oxidation.— 86° is oxidized to. SeOa by HNO, , nitrohydrochloric acid,
140 THE IRON AXD ZINC GROUPS. §114.
H2SO4 hot concentrated, by heating in nir or oxygren, etc. HaSeO, is oxidized
to H.SeOt by continued chlorination, and by fusion with KNO, . H2Se04 is
reduced to HaSeO, by boiling with HCl . S6j reduces selenous compounds to
the red Se° , even in H.SO4 solutions (distinction from tellurium) (Keller,
J. Am. Sor„ 1900, 22, 241). H.S forms a precipitate of Se mixed with S . SnCl,
precipitates Se** from HCl or HaSO^ solutions of selenous compounds.
The Ibox and Zinc Gboups (Third and Fourth Groups).
§114. The Metals of the Earths and the more Electro-Positive of the
Heavy Metals.
Aluminum Al = 27.1 ] Lanthanum La = i;tS.G
Chromium Cr = 52.1 Neodymium Nd = 14:{.r>
Iron Fe = 66.9 , Praseodymium Pr = I40.r>
Cobalt Co = 59.00 I Samarium Sm = 150.3
Nickel Ni = 58.70 Scandium So = 44.1
Manganese Mn = 55.0 Tantalum Ta = 1S2.8
Zinc Zn = 65.4 Terbium Tr = KiO.
Cerium Ce = i:j9.0 Thallium Tl =204.15
Colitmbium Cb = 93.7 j Thorium Th =232.0
Erbium £ =166.0 I Titanium Tl = 48.i:>
Gallium Ga = 70.0
Glucinum Gl = 9.1
Indium In =114.0
Uranium IT = 239.6
Ytterbium Yb = 173.2
Yttrium Y = 89.0
Zirconium Zr = 90.4
§116. The metals above named gradually oxidize at their surfaces in
the air, and their oxides are not decomposed by heat alone. Zinc, iron,
cobalt, nickel, and, with more difficulty, manganese, chromium, and most
of the other metals of the groups, are reduced from their oxides by igni-
tion at white heat with charcoal. They are all reduced from oxides by
the alkali metals. Iron is gradually changed from ferrous to ferric
combinations by contact A\ith the air. Chromium and manganese are
oxidized from bases to acid radicals by ignition with an active supply of
oxygen in presence of alkalis; these acid radicals acting as strong oxidizing
agents.
§116. The oxides and hydroxides of these metals are insoluble in water
and they are precipitated from all their salts by alkalis. In the case of
zinc, the precipitate redissolves in all the alkalis; the aluminum hydroxide
redissolves in the fixed alkalis, but very slightly in ammonium hydroxide;
the precipitate of chromium redissolves in cold solution of fixed alkalis,
precipitating again on boiling; the hydroxides of cobalt and nickel dissolve
in ammonium hydroxide. The oxide of chromium after ignition is insol-
uble in acids; the oxides of aluminum and iron are soluble with difficulty.
The presence of tartaric acid, citric acid, sugar, and some other orgfanic
substances, prevents the precipitation of bases of these groups by alkalis.
§117. Ammonium salts, as NH^Cl , dissolve moderate quantities of the
§120. THE IRON AND ZINC 0R0UP8, 141
hydroxides of manganese, zinc, cobalt, nickel, and ferrous hydroxide; but,
so far from dissolving the hydroxide of aluminum, they lessen its slight
solubility in ammonium hydroxide.
§118. It thus appears that ammoninm hydroxide, with ammonium
chloride, the latter necessary on account of magnesium (§189, Ga), man-
ganese (§134, 6a)y and aluminum, will fully precipitate only aluminum,
ohromimn, and ferrioum of the important metals above named. These
metals therefore constitute the THIRD GROUP (§127), and the reagent
of this group is AMMONIUM HYDROXIDE in the presence of AM-
MONIUM CHLORIDE. Since aluminum, chromium, and ferrioum are
precipitated by ammonium hydroxide in the presence of ammonium
chloride (Fe" by its previous oxidation with HNO3 is present as Fe'")
constituting the THIRD GROUP; the remaining of the most important
metals — cobalt, nickel, manganese, and zinc — constitute the FOURTH
GROUP (§137). They are precipitated by the group reagent, AMMON-
IUM SULPHIDE or HYDROSULPHURIC ACID in an AMMONIACAL
SOLUTION. Some chemists do not make this classification of these
metals, but precipitate them all as one group with ammonium sulphide
(§144), from neutral or ammoniacal solutions. The sulphides of Fe , Co ,
Hi , Mn , and Zn are not formed in presence of dilute acids, which acids keep
them in solution during the second group precipitation; but are insoluble
in water, which enables them to be precipitated by alkali sulphides, and
separated from the fifth and sixth groups. The other two metals, Al and
Cr, do not form sulphides, in the wet way, but are precipitated as hy-
droxides by the alkali sulphides.
§119. Hydrosulphurio acid scarcely precipitates the metals of these
groups, unless it be from some of their acetates (§135, 6e), owing to the
solubility of the sulphides in the acids, which would be set free in their
formation. Thus, this change cannot occur — FeClo + HjS = FeS -U
2HC1 — ^because the two products would decompose each other. Therefore
when it is desired to precipitate the metals as sulphides, neutralized
hydrosulphuric acid — an alkali sulphide— is used in neutral or alkaline
solution; or, what is equivalent, hydrosulphuric acid gas is passed into the
strongly ammoniacal solution,
§120. As most of the chemically normal salts of heavy metals have an
acid reaction to test-paper, we can only assure ourselves of the requisite
neutrality by adding sufficient ammonium hydroxide, which itself precipi-
tates the larger number of the bases, as we have just seen (§116). But
the resulting precipitate of hydroxide, as Fe(0H)2 , is immediately changed
to sulphide, FeS, by subsequent addition of ammonium sulphide; as the
student may observe, by the change in the color of the precipitate.
Ferric and manganic salts are reduced to ferrous and manganous salts^
142 ALUMINUM, §121.
by hydrosulphuric acid, in solution, with a precipitation of sulphur, anil
the corresponding reaction occurs with chromates.
§121. Soluble carbonates precipitate all the metals of these groups, in
accordance with the general statement for bases not alkali (§206, 6a).
With aluminum and chromium, the precipitates dissolve sparingly in ex-
cess of potassium or sodium carbonate; with Co , Ni and Zn , the precipitate
dissolves in excess of (iiiiJsCOs . In the case of ferrous and manganous
salts, the precipitates are normal carbonates; with zinc, cobalt, and nickel
salts, they are basic carbonates; while with ferric, aluminum, and chrom-
ium salts, the precipitates are hydroxides. Barium carbonate precipitates
Al , Cr'" and Fc'", which, in the cold and from salts not sulphates, is a
separation from the fourth group metals.
§122. Soluble phosphates precipitate these as they do other non-alkali
bases. The acid solutions of phosphates of the metals of the third and
fourth groups are precipitated by neutralization. Phosphates of Co , Si ,
and Zn are redissolved by excess of NH^OH , and those of Al , Cr , and Zn
by excess of the fixed alkalis. The recently precipitated phosphates of all
the metals of these groups which form sulphides, are transformed to sul-
phides by ammonium sulphide, due to the fact that the sulphide is less
soluble than the phosphate: TeSPO^ + (NH^aS = FcS + (NHJjHPO, .
Hence, the only phosphates which may occur in a sulphide precipitate are
those of Al , Cr , Ba , Sr , Ca , and Mg .
§123. The metals of the third and fourth groups are not easily reduced
from their compounds to the metallic state by ignition before the blow-
pipe, even on charcoal, except zinc, which then vaporizes. Three of them,
however — iron, cobalt, and nickel — are reducible to magnetic oxides. The
larger number of them give characteristic colors to beads of borax and of
microcosmic salt, fused on a loop of platinum wire -before the blow-pipe.
None of them color the flame or give spectra, unless vaporized by a higher
temperature than that of a Bunsen burner (spark spectra).
The Iron Group (Third Group).
Alnminnm, Chrominm, Iron.
§124, Alnminnm. Al = 27.1 . Valence three.
1. Properties.— iSf peci/f^ gravitv, 2.583 (Mallet, C. N., 1882, 46, 178). Melting
point, 654.5° (Heycock and Neville, J, C, 1695, 67, 187). It is a tin-white metal
(the powder is gray), odorless and tasteless, very ductile and malleable, about
as hard as silver. It has not been vaporized, impurities increase the melting-
point, when molten it possesses great fluidity. As a conductor of heat it is a
little better than tin and about two-thirds as good as silver. It conducts
electricity about one-half as well as copper (Voggendorf, Pogg., 1856, 97, 643).
about one-third as well as silver (Matthiessen, Pogg., 1858, 103, 428), and about
eight times better than iron. Commercial aluminum is never pure, containing
§124, 6a. ALUMINUM. 143
small amounts of silicon and iron, and sometimes Cu and Pb , with 96 to 90
per cent aluminum (Hampe, A., 1876, 183, 78). It is used for cooking utensils,
canteens and other military equipments, boats, small weights, measures,
articles of ornament and scientific instruments; a^ an alloy with copper
(aluminum bronze) it finds extensive application.
2. Occurrence. — Not found free in nature. Is found in corundum, ruby and
sapphire, as nearly pure AljO,; in diaspore (AlOOH); in bauxite (AlaO(OH)4);
in felspar (XsAl^SiOe); in cryolite (NasAlFe). As a silicate in all clays and in
very many minerals. It is widely distributed, constituting about one-twelfth
of the earth's crust.
3. Preparation.— (i) By electrolysis of the fused NaAlCl^ . (2) By fusion of
cryolite or the chloride with Na or X . (3) By heating NaAlCl« with zinc, with
•which it forms an alloy from which the zinc is driven off by a white heat.
(-#) By fusion of the chloride with potassium cyanide, (o) By fusing AljSs
\¥ith iron. A great many new methods have been patented. See Dammer, 3,
79.
4. Oxide and Hydroxides.— AljOg is formed by heating the hydroxide,
nitrate, acetate or other organic salt, difficultly soluble in acids after ignition,
but may be dissolved after fusion with KHSO4 or NajCO, . Al(OH), is
formed when aluminum salts are precipitated with cold ammonium hydroxide.
A1,0(0H)4 is formed if the precipitation is made at 100°.
5. Solubilities. — a. — Metal, — Pure aluminum scarcely oxidizes at all in dry or
moist air; the electrolytically deposited powder oxidizes gradually in the air.
Powdered or leaf aluminum when boiled with water evolves hydrogen, forming
the hydroxide. It is attacked by the halogens forming the corresponding
halides (Gustavson, 5/., 1881, (2), 36, 556). Dilute sulphuric acid attacks it
slowly, evolving hydrogen (Ditte, C, r., 1890, 110, 573); the hot concentrated
acid dissolves it readily with evolution of SO^ . Nitric acid, dilute or con-
centrated, attacks it very slowly (Deville, A. Ch., 1855, (3), 43, 14; Montemartini^
Oazzetta, 1892, 22, 397; Ditte, /. c, 782). Hydrochloric acid, dilute or concen-
trated, dissolves it readily with evolution of hydrogen; also attacked readily
by fixed alkalis, sparingly by NH«OH (Gottig, B., 1896, 29, 1671), evolving
hydrogen with formation of an aluminate: 2A1 -f 2K0H 4- 2H2O = 2KA102 +
3Ha . It is attacked by fixed alkali carbonates (/)., 3, 87). When ignited with
sodium carbonate, aluminum oxide is formed, sodium is vaporized and a small
amount of aluminum nitride produced (Mallet, J. f., 1876, 30, 349). Fused
XOH is decomposed by aluminum at very high temperature, the potassium
being vaporized (Deville, ./., 1857, 1.V2). It is not at all attacked by cold four
per cent acetic acid (vinegar) even in presence of NaCl , and when boiled for
14 hours with the above mixture a square meter of surface (weighing 24.7426
grams) lost but 0.047 grams (one i)art in 526).
6. — Oxide and hydroxide. — The oxide is insoluble in water, and when not
too strongly ignited dissolves readily in dilute acids and in fixed alkalis.
Corundum, crystallized ALOs , is insoluble in acids, but is rendered soluble
by fusion in fixed alkali carbonates or sulphates. The hydroxide Al(OH),
is insoluble in water, readily soluble in acids and in fixed alkalis, sparingly
.soluble in ammonium hydroxide, the solubility, however, being much
decreased by the presence of ammonium salts, r. — Salts. — Aluminum phos-
phate is the most important of the aluminum salts, insoluble in water. The
normal acetate is soluble, the basic acetate insoluble in water (separation
from Or and the fourth group). The chloride is deliquescent. The double
sulphates of aluminum and the alkali metals (alums) are soluble and readily
melt in their water of crystallization, becoming anhydrous. Anhydrous
aluminum sulphate is insoluble in water (Persoz, A. Ch., 1859, (3), 56, 102).
Solutions of normal salts of aluminum have an acid reaction.
6. Beactions. a.— The alkali hydroxides and carbonates* precipitate
aluminum hydroxide (7), A1(0H)3 (4), grayish-white, gelatinous insoluhle
• According to Lanjjlois (A. Ch., ]8ft6. (3), 48. 502) the precipitate with alkali carbonates always
contains CO,. He assitms the formula 3(Al,Oa CO,) + 5i Al,<>,.8tl,0).
144 ALUMINUM, §124, 6&.
in water, soluble in excess of the fixed alkali hydroxides* (2) (Prescott,
./. Am. Soc, 1880, 2, 27; Ditte, A. Ch., 1897 (6), 30, 266), sparingly soluble
in the fixed alkali carbonates and in ammonium hydroxide but much less
so if ammonium salts be present. The solution of fixed alkali aluminate
is precipitated as aluminum hydroxide by careful neutralization of the
alkali with acids including hydrosulphuric (5), and carbonic, as basic
hydroxide, by adding excess of ammonium chloride {J^) (distinction from
zinc which is precipitated by a small amount of NH^Cl , but redissolves on
adding an excess) (Lowe, Z., 1865, 4, 350). The excess of potassium
hydroxide liberates ammonia forming potassium chloride, thus reducing
the amount of fixed alkali present. The precipitate is more compact and
washes more readily than the gelatinous normal hydroxide. Barium car-
bonate, on digestion in the cold for some time completely precipitates
aluminum salts as the hydroxide (5) mixed with a little basic salt. (See
§126, Ga.) The presence of citric, oxalic, or tartaric acid greatly hinders
the precipitation of aluminum hydroxide, and an excess may entirely pre-
vent its precipitation by the formation of a soluble double salt, e, g.,
Hlk!L{CJSLfi^2 • Other organic substances, as sugar, pieces of filter paper,
"etc., hinder the precipitation. To obtain complete precipitation all or-
ganic substances should be decomposed.
(1) AlCl. -f 3K0H = Al(OH), -f- 3KC1
2A1C1, + 3X,C0, -f 3HaO = 2A1(0H), -f- 6KC1 -f- 3C0,
{2) Al(OH), -h KOH = KAIO, -f- 2H,0
or AlCl, -h 4K0H = KAIO, -|- 3KC1 -f- 2H,0
(3) 2KA10, + H,S -h 2H,0 = 2A1(0H). -f- K3S
(.» 2KAIO2 -h 2NH4CI -h HaO = AUOCOH), -f 2KC1 + 2NH,
(o) 2A1C1, -h 3BaC0, -h 3H3O = 2A1(0H), -f- 3BaCl, -f- 3C0,
&. — Oxalates do not precipitate aluminum salts. The acetate of alum-
inum is decomposed upon boiling, forming the insoluble basic acetate
(separation of iron and aluminum from the fourth group) : A1(C2H302)3 +
H2O = A1(C2H[302)20H[ -f HC2H3O2 . The basic acetate is best formed as
follows: To the solution of aluminum salt add a little sodium or am-
monium carbonate, as much as can be added without leaving a precipitate
on stirring, then add excess of sodium or ammonium acetate, and boil for
some time, when the precipitation at length becomes very nearly complete.
Phenyl hydrazine, CqH^NHNHo , completely precipitates aluminum as
the hydroxide from the neutral solution of its salts (complete separation
of aluminum and chromium from iron which should be in the ferrous
condition) (Hess and Campbell, J. Am. Soc, 1899, 21, 776).
•A solution of barium hydroxide may be used to dissolve the A KOH), in separating fr«m
Fe(OH), and CMOH),; espeoiaUy valuable in detecting the presence of smaU amount** of
Aluminum when the reagents NaOH and KOH contain aluminum (Neumann. M.^ ISM, 15 f: .
§124, Ge. ALUMINUM, 145
c. — ^Nitric acid is a very poor solvent for metallic aluminum, but a good
solvent for the oxide and hydroxide. The metal dissolves in a solution of the
normal aluminum nitrate, evolving hydrogen and forming the basic nitrate
A1^0»(N0,), (Ditte, C. r., 1890, 110, 782).
d. — ^Alkali phosphates precipitate aluminum phosphate, AIPO4, white,
insoluble in water and acetic acid, soluble in mineral acids, and in the
fixed alkalis (separation from FePO^) (Grueber, Z, angew,, 1896, 741).
A separation of Al and PO4 may be effected by dissolving in hydrochloric
acid adding tartaric acid and then ammonium hydroxide, and digesting
some time with magnesia mixture (magnesium sulphate to which sufficient
ammonium chloride has been added so that no precipitate is obtained
when rendered strongly alkaline with ammonium hydroxide). The filtrate
contains nearly all of the aluminum. The same method may be employed
with Fc'" and PO4 . See also 7.
e, — The sulphide of aluminum cannot be prepared in the wet way, that
prepared in the dry way being decomposed by water (Curie, C. iV., 1873,
28, 307). Hydrosulphuric acid does not precipitate aluminum from acid
or neutral solutions; from its solutions in the fixed alkalis it is precipitated
as the hydroxide on addition of sufficient hydrosulphuric acid to neutralize
the fixed alkali (distinction from zinc which is rapidly precipitated from
its alkaline solutions, as the sulphide). The alkali sulphides precipitate
aluminum from its solutions, as the hydroxide; from acid or neutral solu-
tion HjS is evolved: 2AICI3 + SCNHJ^S + GH^O = 2A1(0H)3 + 6NH,C1
+ 3HoS, from solutions in the fixed alkalis ammonia is evolved, fixed
alkali sulphide being formed: 2K&10o + (NH4)oS + 2H2O = 2A1(0H)8 +
KjS + 2NH3 .
Sodium thiosulphate precipitates, from aluminum salts, in neutral solutions,
aluminum hvdroxide ivith free sulphur and liberation of sulphurous anhydride:
2Ala(S0J, + GNa^SjO, -h 6H,0 = 4A1(0H), -f 3Sj -f GNa.SO^ + GSO^ . A
sniiall amount of sodium tetrathionate is formed and also some hydrosulphuric
acid (Vortmann, B., 1889, 22, 2307). Sodium sulphite also precipitates alu-
minum hydroxide, with liberation of sulphur dioxide: 2AlCls -f SNa^SO, -f
SHjO = 2A1(0H)3 -h ONaCl + aSO^ . Neither of the above reagents precipi-
tate iron salts, thus effecting* a separation of aluminum (and chromium) from
iron.
Aluminum, chromium and ferric sulphates crystallize with the sulphates
of the alkali metals, forming a class of compounds, alums, of which the
potassium aluminum compound is perhaps best known, KA1(S04)2.12H20 ,
common alum. These compounds melt in their water of crystallization,
becoming anhydrous upon further heating. The freshly ignited alum is
only sparingly soluble in cold water, but upon standing becomes readily
soluble, dissolving in less than one part of hot wate^*. The alums are usu-
ally less soluble than their constituent sulphates and may be precipitated
by adding a saturated solution of alkali sulphate to a very concentrated so-
lution of Al , Cr'" , or Fe'" sulphate.
146 ALUMINUM. §124, 6/:
/. — Aluminum chloride is a very powerful dehydrating" agent and is much
used in organic chemistry as a halogen carrier. An impure aluminum chlorate,
mixture of KCIO, and AlsCSO^), , is much used in calico printing (Schlum-
berger, DinyL, 1873, 207, 63). g. — Aluminum salts are precipitated by solu-
tions of alkali arsenites and arsenates, but not by arsenous or arsenic acids,
/i. — Potassium chromate forms a yellow gelatinous precipitate, potassium
bichromate gives no precipitate with aluminum salts, i. — Solution of borax
precipitates an acid aluminum borate, quickly changed to aluminum hydroxide.
7. Ignition. — Compounds of aluminum are not reduced to the metal, but
most of them are changed to the oxide, by ignition on charcoal. If now this
residue is moistened with solution of cobaltous nitrate, and again strongly
ignited, it assumes a blue color. This test is conclusive only with infusible
compounds, and applies only in absence of colored oxides. Aluminum com-
pounds ignited on charcoal in presence of sulphur are changed to AI3S, (Buch-
erer, Z, amjew,, 1892, 483).
To separate Al irom PO4 , fuse the precipitate or powdered substance with
1*4 parts finely divided silica and 6 parts dried sodium carbonate in a platinum
crucible, for half an hour. Digest the mass for some time in water: add
ammonium carbonate in excess, filter and wash. The residue consists of
aluminum sodium silicate; the solution contains the PO4 , as sodium phosphate.
The Al can be obtained from the residue bj' dissolving it in hydrochloric acid,
evaporating to dryness to render the silica insoluble. Treat with hydrochloric
acid and filter; the filtrate containing aluminum chloride.
8. Detection. — After the removal of the first two groups it is precipi-
tated with Cr and Fe"' as the hydroxide, A1(0H)3 , by NH^OH in the pres-
ence of NH4CI . It is separated from Fe(0H[)3 and Cr(0H)3 by boiling
with KOH . From the filtrate acidulated with HCl it is precipitated as
hydroxide with {iAA^)^^^^ ; or it is precipitated from the KOH solution
by an excess of NH4CI (6a).
9. Estimation. — Aluminum is usually weighed as the oxide, after ignition.
It is separated from zinc as a basic acetate; from chromium by oxidizing the
latter to chromic acid, by boiling with potassium chlorate and nitric acid, or
by fusing with KNO, and Na-jCOg , or by action of CI or Br in presence of
KOH . and after acidulating with HCl precipitating the aluminum with am-
monium hydroxide. It may be separated from iron by boiling with KOH (6a).
by Na.S..O, (Ge), or by phenylhydrazine (Ob). It is separated from iron by
conversion into the oleate and dipsolving the oleate of iron (Fe'" or Fe") in
petroleum (Borntraeger, Z., 1893, 32, 187). It is sometimes precipitated and
weighed as the i)hosphate.
10. Oxidation. — Aluminum reduces solutions of Pb , Ag , Hg *, Sn , Bi
(incompletely), Cu f, Cd , Co , Ni , Zn J and Gl (in alkaline mixture only),
Te, Se, Au, and Pt, to the metallic state; ferric salts to ferrous salts;
As and Sb with HCl become respectively AsH, and SbHg ^vith alkalis As'"
is reduced to AsHy . As^ is unchanged (§69, C/ft and 10), and Sb'" and
Sb^ become Sb°. Aluminum salts are not reduced to the metallic statt'
by any other compounds at ordinary temperature; by fusion with K or Na
metallic aluminum is obtained, much better, however, by the aid of the
electric current,
• Klandy, C. C, 1898, 201 ; WIsliconus, B. 1«95. 28, 1338. t Tommasi, Bl., 1882, (2). 37, ua.
t Flavitsky, B., 1878, 6, 105 ; Zimmerman, Z., 1888, 27, 61.
§125, 5c. CHROMIUM. 147
§126. Chromium. Cr = 52.1 . Valence two, three and six.
1. Properties.— Speci/k? gravity, 6.81 (Woehler, A., 1859, 111, 231). Melts with
greater difficulty than platinum (CJlatzel, B„ 1890, 23, 3127). A grayish-white
crystalline metal. The hardness of steel is greatly increased by the presence
of less than one per cent of chromium. It is non-magnetic (Woehler, i. c). It
burns to the oxide CPaO, when heated to 200° to '600° in the air (Moissan, C. r.,
1879, 88, 180).
2. Occurrence. — Not found native. It is found in several minerals. Chrome-
ironstone or chromite (FeOCr^O,) is the chief ore of chromium, and is usually
employed in the manufacture of chromium compounds. Chromite and also a
double sulphide of iron and chromium, FeCr2S4 , are found in many meteors.
3. Preparation. — (i) By electrolysis of the chloride. (2) By fusing the
chloride with potassium or sodium. (5) By ignition of the oxide with carbon.
(4) By fusing CrCl, with Zn , Cd or Mg , using KCl and NaCl as a flux, and
removing the excess of the Zn , Cd or Mg by dissolving in nitric acid, which
does not dissolve metallic chromium. (5) By ignition of the oxide with alu-
minum (Goldschmidt, J.., 1898, 301, 19).
4. Oxides and Hydroxides. — ChromQUS oxide, CrO , has not been isolated. The
corresponding hydroxide, Cr(OH)a , is made by treating CrClz with KOH .
Chromic oxide, CrjO, , is made by a great variety of methods, among which are
fusing the nitrate, or higher or lower oxides and hydroxides in the air; heating
mercurous chromate, or the dichromates of the alkalis:
4Hg2CrO« = 2CraO, + 8Hg -j- 50^
(NHOaCraO, = Cr,0, + N, + 4H,0
4X,Cr,07 = 2Cr208 -j- 4K,Cr04 -\- 30,
In the last the K3Cr04 may be separated by water. After heating to redness,^
Cr,0, is insoluble in acids. Chromic hydroxide, Cr(0H)8 , is precipitated by
adding NH4OH to chromic solutions. That formed by precipitating with KOH
or NaOH retains traces of the alkali, not easily removed by washing.
Chromium trioxide or chromic anhydride, CrO, , is formed as brown-red
needles upon addition of concentrated sulphuric acid to a concentrated solution
of "KiCTiOf', to be freed from sulphuric acid it must be recrystallized from
water, in which it is readily soluble, or treated with the necessary amount of
BaCrO^ (Moissan, A, Ch,, 1*885, (6), 5, 568). It is also prepared by transposi-
tion of BaCrO* with HNO, or H3SO4; PbCrO^ with H2SO4; and Ag2Cr04 with
HCl; etc. It melts at about 170° (Moissan, /. r.), decomposing at higher tem-
perature into CraOj and O . It is used in dyeing silk and wool, but not
cotton fabrics. It is a powerful oxidizing agent, being reduced to chromic
oxide. The existence of chromic acid, HsCrOf , is disputed (Moissan, /. r.;
Field. C. N., 1892, 66, 153; and Ostwald, Zeit, phys. Ch., 1888, 2, 78). Two
series of salts are formed as if derived from chromic acid, H.CrO^ , and
dichromic acid, "K^Ct^O^ . The salts are quite stable and find an extended
application in analytical chemistry (G7i, §57, §59, §186, etc.).
5. Solubilities. — a. — Metal. — Chromium is not at all oxidized by water or
moist air at 100°. Heated above 200° it is oxidized to CroOa , rapidly in pres-
ence of KOH . It is soluble in HCl or dilute H.SO^: insoluble in concentrated
H^SO* or in HNO, , dilute or concentrated. Chlorine or bromine attack it
with formation of the corresponding halides (Woehlei*, I. c.; Ufer, A., 1S59. 112,
'102). 6. — Oxides and Hydroxides. — Chromic oxide, Cr^Os , is insoluble in water,
slowly soluble in acids, but not at all if previously ignited (Traube, A., 1848,
66. 88); the hydroxide is insoluble in water, soluble in acids, sparingly soluble
in ammonium hydroxide, soluble in fixed alkalis to chromites, reprecipitated
again upon boiling. The presence of other metallic hydroxides, as iron, etc.,
hinders the solution in fixed alkalis. Chromic anhydride, CrOg , is very soluble
in water, soluble in reducing acids to chromic salts.
r. — Salts, — Chromic sulphide is not formed in the wet way, being
decomposed by water; the phosphate is insoluble in water. The chloride
148 CHROMIUM, §126, Ga.
exists in two modifications; a deliquescent soluble chloride, which al^o
forms a soluble basic chloride (Ordway, Am, S,, 1858 (">), 26, 202);
and a violet sublimed chromic chloride absolutely insoluble in water,
hot or cold, or in dilute or concentrated acids, the presence of a very
small amount of chromous or stannous chloride at once renders this modi-
fication soluble in water (Peligot, A, Ch., 1846 (3), 16, 208); the bromide
and sulphate also exist in soluble and insoluble modifications; the nitrate
and also the basic nitrates are readily soluble in water (Ordway, 1. c).
There are many double salts, the sulphates of chromium and the alkali
metals, chrome alum, forming salts similar to the corresponding aluminum
compoimds. There are two modifications of solutions of chromium salts,
one having a green color and the other violet to red, the tints are modified
somewhat by the degree of the concentration. All normal chromic salts
in solution have an acid reaction, being partially hydrolized.
6. Eeactions.* a. — Alkali hydroxides and carbonates precipitate solu-
tions of chromic salts, as chromium hydroxide, gelatinous, gray-green or
gray-blue according to the variety of solution from which it is obtained
(5c), insoluble in water, soluble in acids; soluble in excess of the fixed
alkalis to chromites: Cr(0H)3 + KOH = ECrOo + 2H2O ; the chromium
is completely reprecipitated on long boiling (distinction from aluminum),
or on heating with an excess of ammonium chloride. The presence of
ferric hydroxide and some other compounds greatly hinders the solution
in fixed alkalis, hence chromium cannot be separated from iron by excess
of fixed alkali. Chromium hydroxide is slightly soluble in excess of colrl
ammonium hydroxide to a violet solution, completely reprecipitated on
boiling. The precipitate formed with the alkali carbonates is almost
entirely free from carbonate: 2CrCl3 -f SNaoCO.^ -f 3HoO = 2Cr(0H)3 +
GNaCl + 3C0.^ . Barium carbonate precipitates chromium from its solu-
tions (better from the chloride) as a hydroxide with some basic salt, the
precipitation being complete after long digestion in the cold (separation
from the fourth group). For removal of excess of reagent, add H2SO4
and the filtrate will contain the chromium as a sulphate.
Alkali dichromates are changed to normal chromates by alkali hydrox-
ides or carbonates.
ft. — Chromium forms no basic acetate and remains in solution when the
basic acetates of aluminum and ferric iron are formed (6&» §124 and §126).
Potassium cyanide precipitates chromium hydroxide. Oxalates and ferro-
cyanides cause no precipitate. H2Cr04 is reduced to chromic compounds
•Chromous salts aro very unstable, they are great reducing agents* oxidizing rapidly when
exposed to the air. Thoy are almost never met with In analysis. Chromous chloride, Cr€l,, is
formed when the metal is heated in contact with hydrochloric acid gas (tJfer, I. c ); also by re-
duction of CrCIa with hydrogen in a heated tube (Moberg, J, j)r., 1848, 44. 822). Precipitates are
formed in its solutions by the allcali hydroxides, carb<mates, sulphides, etc. fMoissan, BL, 1«*2
(2/, 37. 296).
§125, 6ft. CHROMIUM. 149
by K,re(CN')e and XCNS. r. — Nitrites or nitrates are without action upon
chromium salts in the wet way, but upon fusion in presence of nitrites or
nitrates and alkali carbonate a ehromate is formed (separation from Fe and
Al). d, — ^Hypophosphorous acid reduces chromates to chromic salts. Soluble
phosphates, as Na2HP04 , precipitate chromic phosphate, CrPO^ , insoluble in
acetic acid, decomposed by boiling- with XOH , leaving the phosphate in solu-
tion (Kammerer, J, C, 1874, 27, 1005).
e. — Hydrosulphuric acid is without action upon neutral or acid solutions
of chromium salts, chromites as ECrOs are precipitated as chromium'
hydroxide; 2KCr02 + H^S + 2H2O = 2Cr(0H)3 + K^S. The hexad
chromium of chromates is reduced to the triad condition with liberation
of sulphur, in neutral or alkaline solutions, chromium hydroxide being
formed: 2K^CTfi^ + SR^S = 4Cr(0H)3 +2X38 + SSo + 2H2O ; in acid
solutions a chromium salt is formed (10). Alkali sulphides precipitate
chromium salts as the hydroxide liberating HgS :
2CrCl, -h 3(NH,),S -f- 6HaO = 2Cr(0H), + 6NH4CI + 3H,S
Chromates are reduced and precipitated as chromium hydroxide with sepa-
ration of sulphur: 4K.CtO^ + 6{NH^)2S + 4H2O = 4Cr(0H)3 + 8K0H
-f- 3S2 + I2NH3 • Soluble sulphites and thiosulphates reduce chromates
in acid solution (Donath, /. C, 1879, 36, 401; Longi, Gazzetta, 1896, 26,
ii, 119).
f. — Hydrochloric acid reduces chromates to chromic chloride on boiling,
with evolution of chlorine: 2K2Cr04 + 16HC1 = 2CrCl3 + 4KC1 + SClg +
8H2O ; more readily without evolution of chlorine in presence of other
easily oxidized agents, as alcohol, oxalic acid, etc.: K2Cr^07 + 8HC1 +
aC^HsOH = 2KC1 + 2CrCl3 + SCjH^O (acetaldehyde) + 7H2O . If the
dry ehromate be heated with sulphuric acid and a chloride (transposable
by sulphuric acid) (§269, 5), brown fumes of chromium dioxydichloride
are evolved: K.,CTfi^ + 4NaCl + 3H2SO4 = 2Cr02Cl2 + K2SO4 + 2Na2S04
-f 3H2O (§269, Sd) (Moissan, Bl, 1885 (2), 43, 6). To obtain a quantity of
Cr02Cl2, Thorpe (J, C, 1868, 21, 514) recommends 10 parts of NaCl and
12 parts KjCrjOT fused together and distilled with 30 parts of HoSO^ .
Hydrobromic acid reduces chromates to chromic bromide with evolution
of bromine; hydriodic acid to chromic iodide with evolution of iodine.
In the presence of hydrochloric or sulphuric acids all the bromine or
iodine is set free. KXrfi^ + 6HI + 4H2SO4 = KjSO^ + Cr>(S0,)3 +
3I2 + 7H2O . Hydriodic acid acts most readily upon chromates, the
hydrochloric least readily. Chromic hydroxide and chromic salts, when
boiled with chloric or bromic acids, or potassium chlorate or bromate and
nitric, sulphuric or phosphoric acids, become chromic acid.
g, — Soluble arsenites and arsenates form corresponding' salts with chromic
salts. Chromates in acid solution are instantly reduced to chromic salts by
arsenites or arsenous acid. Chromic acid boiled with arsenous acid in excess
^ves CrAsO, (Neville, J. C, 1877, 31, 283).
h, — Potassium ehromate colors an acid solution of chromic salt brown-yellow;
150 CHROMIUM. §126, 7.
on addition of ammonium hydroxide, a precipitate of the same color is obtained,
chromic chromate (Maus, Pogq,, 1827, 9, 127). The alkali metals form two
classes of chromates: yellow normal chromates and reddish dichromates
(Schulernd, J, C, 1879, 36, 298). The chromates of the alkalis, and those of
magnesium, calcium, zinc and copper are soluble; those of strontium, mercury
(H^') are sparingly soluble; and those of barium, manganese, bismuth, mer-
cury (Hg'), silver and lead are insoluble in water. Alkali chromates or
dichromates are precipitated as normal chromates (in some cases as dichro-
mates) (Preis and Kayman, B., 1880, 13, 340) by solutions of silver, lead, mer-
cury (Hg') and barium salts. Silver chromate is dark red, soluble in nitric
acid and ammonium hydroxide (§69, 6/i); lead chromate is yellow, transposed
with difficulty by nitric acid (Duvillier, A, Ch,, 1873, (4), 30,*^ 212), insoluble in
acetic acid (§57, 6h); barium chromate, yellow, is soluble in hydrochloric and
nitric acids, sparingly soluble in chromic acid (§186, 6h),
7. Ignition. — Chromic oxide, chromic salts and chromates dissolve in beads
of microcosmic salt, and of borax, before the blow-pipe, in both reducing and
oxidizing flames, with a yellowish -green tint while hot, becoming emerald
green when cold. By ignition on charcoal the carbon deoxidizes chromic
anhydride, CrO. , free or combined, and a green mass, CrjO, , is left. When
chromium compounds are fused with an alkali carbonate, and a nitrite, nitrate,
chlorate, bromate or iodate, an alkali chromate is formed, soluble in water
(distinction from Al and Fe).
8. Detection. — If present as chromate (solution red or yellow), it is
reduced by HCl and alcohol. Precipitated with Fe'" and Al, after the
removal of the metals of the first and second groups, by NH4OH in pres-
ence of NH4CI. Boiling with KOH separates the Al and leaves the Cr
with the Fe , as hydroxides. The precipitate is fused on a platinum foil
with NajCOs ^^^ ^^0.3 which oxidizes the Cr to an alkali chromate, soluble
in water (separation from the Fe). The Cr is identified after acidulation
with HC2H3O2 bv the formation of the vellow lead chromate, using
Pb(C,H,0,),.
9. Estimation. — Chromium is usually estimated gravimetrically (/) as the
oxide. It is brought into this form either by precipitation as a hydroxide (fid)
and ig-nition or, in many cases, by simple ignition (4). (2) As chromate, it may
be precipitated with barium chloride, dried and weighed as such; or in acetic
acid solution it may be precipitated as PbCrO* by Pb(C2H80,)3 , dried and
weighed. Volumetrieally, as a chromate (if present as chromic salt it may be
oxidized to a chromate). (3) By titration with a standard solution of ferrous
sulphate. (4) By liberation of iodine from hj^driodic acid (dg) and measuring
the amount of iodine liberated with standard sodium thiosulphate solution.
10. Oxidation. — Chromous compounds are very strong reducing agents,
changing HgClj to HgCl , CuSO^ to Cu°, SnClj to Sn**, etc. Chromic com-
pounds are oxidized to chromates by chlorates (Giacomelli, UOrosi, ISO^.
18, 48; Storer, Am. S., 1869,98,190) (6/), Na^O^, MnO^ (Marchal and Wier-
nick, Z. angew.y 1891, 511), and PbOjj in acid solution; in alkaline mixture,
by reducing PbOg to PbO , Ag,0 to Ag°, Hg,0 and HgO to Hg**, CuO to
CUgO , EMnO, and EoMnO^ to MnOo (Donath and Jeller, C. C, 1887, 151);
by CI, Br, and I, forming the corresponding halide; and by HjOj*
*Tho use of H,0, in alkaUne solution is proposed by Rigrgs (Am. S., 1894, 148, 409) in the sepa-
ration of Al, Fe and Cr. 100 cc. water, 10 cc. H,0,, and ono gram of NaOH are added to tlie
freshly precipitated hydroxides and dig-ested until effervescence ceases. Filter off the precipi-
tate of ferric hydroxide, acidify the filtrate with acetic acid and precipitate the aluminum with
ammonium hydroxide. The chromium if present will be in the filtrate as sodium chromate.
§126, 3. IRON. 151
(Bai^ann, Z. angew.. 1891, 139). A chromate is also formed when
chromium compounds are fused with an alkali carbonate and an oxidizing
agent (7). Chromic oxide (not ignited) or chromic chloride at 440°
in a current of chlorine become CrOjClj (Moissan, Bl, 1880 (2), 34, 70).
Chromic acid and chromates are reduced to chromic compounds by
HjC^O^ (Wemer, J. C, 1888, 63, 602), K^Fe(CN)e , KCNS , HoS , (1^4)28 ,
HajSjO,, SO2, H2O2, etc. Of most common occurrence in qualitative
analysis is the action of hydrosulphuric acid and alkali sulphides; at first
sulphur is liberated, a part of which may be oxidized to sulphurous and
sulphuric acids (Parsons, C. iV., 1878, 38, 228).
2K,Cr20T + 16HC1 + 6H,S = 4CrCl, + 4KC1 -f 38, + 14H,0
12H3CrO* + 38, = 4Cr,02Cr04 + 680, -f- 12H,0
2H,CrO« 4- 380, = Cr,(804), + 2H,0
Wliile HjOa in alkaline solution oxidizes Cr"' to Cr^, in acid solution the
reverse * action takes place: 2H2CrO^ + 3H2SO4 + SHjOa = Cr2(S0j8 +
3O3 + 8H2O (Baumann, 1. c).
§126. Iron (Ferrum). Fe = 65.9 . Usual valence two and three.
1. Properties. — Specific gravity^ variable, depending upon the purity and
methods of preparation. 7.85 at 16*» (Caron, C, r., 1870, 70, 1263), 8.139
(Chandler-Roberts, C. N., 1875, 31, 137). Melting point, cast iron, 1100° to 1300°;
steel, 1300° to 1600°; wrought iron, 1800° to 2200°. The pure metal melts at
1804° (Carnelley, B., 1880, 13, 441). Pure iron is silver-white, capable of taking
a remarkably fine polish; it is among the most ductile of metals, in this
property being approached by nickel and cobalt (§73, 1); it is the hardest of
the ductile metals (Calvert and Johnston, Dingl., 1859, 152, 129), and in tenacity
it is only surpassed by cobalt and nickel (§132, 1). It softens at a red heat
and may be welded at a white heat. Finely divided iron burns in the air when
i^ited; that made by reduction in hydrogen may ignite spontaneously when
exposed to the air. Steel for tempering purposes contains 0.3 to 1.5 per cent of
carbon, cast iron from 1.7 to 4.6 per cent, and wrought iron less than 0.2 per
cent. Pure iron is attracted by the magnet, but does not retain its magnetism.
Permanent magnets are made of steel. Iron forms two classes of oxides,
hydroxides and salts: ferrous, in which the metal acts as a dyad; and ferric, in
which the metal acts as a triad. The ferrous compounds are changed to ferric
by moist air and by oxidizing agents in general; while ferric compounds are
readily reduced to ferrous compounds by very many reducing agents. Ferric
compounds are much more stable than the corresponding ferrous compounds.
2. Occurrence. — Native iron is rarely found except in meteorites. The chief
ores of iron are red hematite or specular iron ore (FejO,), brown hematite
(2Fe,Os.3B:30), magnetic iron ore (FejO^), iron pyrites (FeS,), spathic iron
ore (FeCOj), clay iron-stone (FeCO, vdth clay), black band (FeCO. mixed with
bituminous matter).
3. Preparation. — Pure iron is not usually found in the market. It is made:
(/) by electrolysis; (2) by heating its purified salts with hydrogen; {3) by
heating the purified salts with some form of carbon; (4) in metallurgy iron is
made from the ores, and the reducing agents are coal, coke, charcoal and
natural gas.
•With a ohromate in acid solution H.O, at first gives a deep blue solution (probably of per-
chromio acid, HOrOf) a very delicate reaction, followed by the reduction to a chromic salt.
152 IRON. §126, 4.
4. Oxides and Hydroxides.— Fcrro«« oxide, FeO , is made from FeaO, by heat-
ing it to 300° in an atmosphere of hydrogen; also by heating FezCsO^ to 160**,
air being excluded. It takes fire spontaneously in the air, oxidizing to FejO, .
Ferrous hydroxide, Fe(0H)2 , is formed by precipitating ferrous salts with KOH
or NaOH , perfectly white when pure, but usually green from partial oxidation.
Ferric oxide, FejO, , is formed by heating FeO , Fe(0H)2 » or any ferrous salt
consisting of a volatile or organic acid in the air; more rapidly by heating
Fe(OH). , Fe(NO,). , or Fe^CSO*), . Ferric hydroxide is formed by precipitat-
ing cold dilute ferric salts with alkalis or alkali carbonates, and drying at 100*.
If KOH or NaOH is used, the precipitate requires longer w^ashing than when
NH4OH is employed. By increasing the temperature and concentration of the
solutions, the following definite compounds may be formed: FeO (OH) ,
Fe,0(0H)4 , Fe^OsCOH)^, Fe40,(0H)e, FegOsCOH)^ . Fe^O^ is slowly formed
by heating FeO or Fe^Os to a white heat. Its corresponding hydroxide may be
made by precipitation: FeCl^ -f 2FeCl, + 8NH4OH = Fes(0H)8 -f 8NH4CI .
Fe,(0H)8 when heated to 90° forms Fe.O^ . The black color and magnetic
properties show that it is a chemical salt and not a mechanical mixture of FeO
and FeaO, . Fe'" acts as an acid towards the Fe"; this oxide, Fe^O^ , or
FeFejO^ , maj^ be called ferrous ferrite. Other ferrites have been formed, €, (/.»
calcium ferrite, CaFe^O^; MgFejO^ and BaFe^O^ (List, B., 1878, 11, 1512): zinc
ferrite, ZnFejO^ . Compare potassium aluminate, KAIO, (§124, 6a), and potas-
sium chromite, KCrO, (§125, 6a). Ferric acid, HjFeO^ , and its anhydride,
FeOs , have not been isolated. Potassium ferrate, XaFeO^ , is made (1) by elec-
trolysis; (2) by heating iron-filings, FeO or FejO, , to a red heat with KNO,;
(3) by heating Fe(OH)a with potassium peroxide K.O2; (-^) by passing CI or Br
into a solution of 5 parts of KOH in 8 parts of water in which Fe(OH), is
suspended; the temperature should be not above 50**. It has a purple color; is
a' strong oxidizing agent. It slowly decomposes on standing: 4X3Fe04 -|-
IOH2O = 8K0H -f 4Fe(0H), -f- 30, . With barium salts it precipitates a
stable barium ferrate, BaFeO^ .
5. Solubilities. — a. — Metal. — Iron dissolves, in hydrochloric acid and in dilute
sulphuric acid, to ferrous salts, with liberation of hydrogen (a): concentrated
cold H^tSO^ has no action, but if hot, SOj is evolved and a ferric salt formed (h):
in moderately dilute nitric acid, with heat, to ferric nitrate, liberating chiefly
nitric oxide (c); in cold dilute nitric acid, forming ferrous nitrate with pro-
duction of ammonium nitrate (d), of nitrous oxide (e), or of hydrogen f)
(Langlois, A. Ch., 1856, [3], 48, 502).
(a) Fe + H2SO4 = FeS04 -j- H,
(ft) 2Fe + 6H,S04 = Fe,(S04), + 3S0, -f- 6H2O
(c) Fe + 4HN0, = Fe(NO,), -f- NO + SH^O
id) 4Fe -f lOHNO, = 4Fe(N0,)a + NH4NO, -f 3HaO
(e) 4Fe + lOHNO, = 4Fe(N0,), + N^O + 5H,0
if) Fe -\- 2HN0, = Fe(NO,), -f H,
In dissolving the iron of commerce in hydrochloric acid, the carbon which it
always contains, so far as combined in the carbide of iron, will pass off in
gaseous hydrocarbons (Campbell, Aw., 1896, 18, 836), and so far as uncombined
will remain undissolved, as graphitic carbon. The metal is attacked by moist
air, forming chiefly 2Fe208.3H20 , iron rust. WTien hot iron is hammered, scale
oxide, Fe,0,.6FeO, is formed. Cold concentrated HNO, forms passive iron.
h. — Oxides and hydroxides. — Ferrous oxide and hydroxide unite with acids
with rapid increase in temperature, forming ferrous salts, always mixed with
more or less ferric salts. The ferrous salts are much more readily prepared
by the action of dilute acids upon the metal, or upon FeCO, or FeiS . FejO^ ,
treated with an insufficient amount of HCl, forms FeClj and FeaO,; treated with
HCl sufficient for complete solution, a mixture of FeCI. and FeCls is obtained,
which, when treated with excess of ammonium hydroxide and dried at 100"*
again exhibits the magnetic properties of the original. Ferric oxide, Fe.©., , dis-
solves in acids, quite slowly if the temperature of preparation of the oxide has
been high. Mitscherlich (J. pr., 1860, 81, 110) recommends warm digestion with
ten parts of a mixture of sulphuric acid and water (8-3). If the oxide be
§126, 6a. IRON. 153
headed with alkalis or alkali oarbonales, it then dissolves much more readily in
acids. Ferric hydroitide^ Fe(0H)8 , is insoluble in water (for a soluble colloidal
ferric hydroxide, see Sabanejeff, C. C, 1891, i, 11), readily soluble in acids to
ferric salts. Freshly precipitated ferric hydroxide readily dissolves in ferric
chloride and in chromium chloride, not in aluminum chloride. A solution of
ferric hydroxide in ferric chloride is soluble in water after evaporation to dry-
ness if not more than ten parts of FeaO, are present to one of the FeCla (Be-
champ, A. Ch„ 1859, (3), 66, 306)
c— Salts. — Ferrom salts, in crystals and in solution, have a light green
color. Solutions of the salts have a slight acid reaction toward litmus.
The sulphate TeSO^.lK^O y is efflorescent; the chloride, bromide, iodide,
and citrate are deliquescent Solutions of all ferrous salts are unstable,
gradually changing to basic ferric salts, more or less insoluble in water.
The carbonate, hydroxide, phosphate, borate, oxalate, cyanide, fcrro-
cyanide, ferricyanide, tartrate, and tannate are insoluble in water.
Ferric salts in solution have a brownish-yellow color, redden litmus and
color the skin yellow. The chloride, bromide, nitrate, and sulphate are
deliquescent. The ferrocyanide, tannate, borate, phosphate, basic acetate,
and sulphite are insoluble in water; the sulphate is soluble in alcohol
(separation from ferrous sulphate). Ferric chloride is soluble in ether
saturated with hydrochloric acid, separation from aluminum (Gooch and
Havens, Am. S., 1896, 162, 416). Solutions of ferric salts, when boilod,
frequently precipitate a large portion of the iron as basic salt, especially
if other soluble salts are present (Fritsche, Z. angew,, 1888, 227; Pickering,
J. C, 1880, 37, 807) (§70, hd footnote).
6. Eeactions. a. — The alkali hydroxides precipitate ferrous hydroxide,
Fe(0H[)2 , white if pure, but seldom obtained sufficiently free from ferric
hydroxide to be clear white, and quickly changing, in the air, to ferroso-
ferric hydroxide, of a dirty-green to black color, then to ferric hydroxide
(4), of a reddish-brown color. The fixed alkalis adhere to this precipitate.
Ammonium chloride or sulphate, sugar, and many organic acids, to a slight
extent, dissolve the ferrous hydroxide or prevent its formation (§§116 and
117). The soluble carbonates precipitate, from purely ferrous solutions,
ferrous carbonate, FcCOj , white if pure, but soon changing, in the air, to
the reddish-brown ferric hydroxide.
Solutions of ferric salts are precipitated by the alkali hydroxides and
carbonates as ferric hydroxide, re(0H)3 , variable to re203.H20 — reO(OH) —
reddish-brown insoluble in excess of the reagents (distinction from alumi-
num and chromium which are soluble in excess of the fixed alkali hy-
droxides and from cobalt, nickel and zinc which are soluble in ammonium
hydroxide). Salts of the fixed alkalis adhere to. this precipitate with great
tenacity and the precipitate obtained from the use of the fixed alkali
carbonates invariably contains traces of a carbonate. Freshly precipitated
barinm carbonate completely precipitates ferric salts in the cold as ferric
154 IRON. §126,66.
hydroxide (separation of ferric iron, with aluminum and chromium, from
ferrous iron, cobalt, nickel, manganese, and zinc; 2FeCl3 + 3BaC0, +
3H2O = 2Fe(0H)8 + SBaCla + SCOj). The mixture should be allowed to
stand several hours (chromium precipitates more slowly than aluminum
or iron), and, sulphates must be absent, as freshly precipitated barium
carbonate reacts with solutions of the sulphates of the fourth group; e. g.,
NiS04 + BaCOj = NiCOg + BaSO^ . The reaction takes place most read-
ily if the metals be present as chlorides. If the precipitate obtained be
treated with an excess of dilute sulphuric acid the ferric hydroxide dis-
solves, leaving the excess of barium as the insoluble sulphate. Freshly
precipitated carbonates of Ca , Hg , Hn , Zn , and Cu react similar to the
barium carbonate.
h, — Oxalic acid and soluble oxalates precipitate from solutions of ferrous
salts, ferrous oxalate, FeC204 , yellowish-white, crystalline, sparingly soluble in
hot water, soluble in HCl , HNO, and H2SO4 ; ferric salts are not precipitated
by oxalates except as reduction to ferrous oxalate takes place.
The acetates, as NaCsHsOj , form in solutions of ferric salts a dull red *
solution of ferric acetate, Fe(C2H30o)3 , which upon boiling is decomposed
and precipitated as basic ferric acetate of variable composition (separation
of iron and aluminum from phosphoric acid (d), chromium, and the metals
of the fourth group). The red colored ferric acetate solution is not
decolored by mercuric chloride (distinction from Fe(CNS)3). The basic
precipitates are soluble in HCl, HNO3 and H2SO4 and are transposed by
alkali hydroxides.
Tannic acid precipitates concentrated solutions of ferrous salts: ferric salts
are ijrecipitated as blue-black ferric tannate {the basis of common ink), insoluble
in water or acetic acid, very soluble in excess of tannic acid. Ferric salts are
completely precipitated by anunonium succinate from hot solutions (Young.
«/. C. 1880, 37, 074). Both ferrous and ferric salts (not nitrates) slightly acid
are completely precipitated by a solution of nitroso B. naphthol (separation
from aluminum and chromium) (Knorre, B,, 1887, 20, 283; Menicke, Z, anfjeir.,
1888, 5). If the Fe'" be in excess of the PO4 the phosphate will all be pre-
cipitated. Hydrochloric acid should be absent, i. e., excess of NaCJSaO:; should
be added (Kiiorre, Z. angetc., 1893, 267).
Potassium cyanide gives with solutions of ferrous salts a yellowish-red pre-
cipitate, which dissolves in excess of the reagent to potassium ferrocyanide.
K4Fe(CN)a: with solutions of ferric salts, ferric hydroxide is precipitated with
evolution of hydrocyanic acid (equation (a), page 156).
Potassium ferrocyanide precipitates ferrous salts as potassium ferrou<
ferrocyanide (&), K2FeFe(CN)« , (Everitt's salt), bluish-white, insoluble in
* Meconic acid and formic acid form red solutions with ferric salts ; benzoic acid grives a flesh
colored precipitate ; phenol, creosote, salljrenin, and other hydroxy aromatic derivatives grive
a blue to violet color. Morphine gives a blue color. The following is recommended as a very
satisfactory test for a trace of iron In copper sulphate. Dissolve one gram of the CaSO« in five
cc. of water, add five cc. of a ten per cent, rtherial solution of salicylic add. If the layer of
contact assumes a \iolet color iron Is present (Grigge, Z., 1806, 34, 460).
!, 66. IRON, 155
acids, transposed by alkalis (c). This is converted into Prussian blue
(see below), gradually by exposure to the air, immediately by oxidizing
agents (d). With ferric salts, ferric ferrocyanide (e), 'Ee^{Te(ClS)^)^ , Prus-
sian blue, is formed, insoluble in acids, decomposed by alkalis (/). If the
reagent be added in strong excess the precipitate is partially dissolved to
a blue liquid. Strong acids should not be present as they color the re-
agent blue. In neutral solutions diluted to one in 500,000 the iron may be
detected (Wagner, Z., 1881, 20, 350). The ferrocyanides are transposed
by KOH and decomposed by fusion with NaNO^ and Na^CO.^ , the iron being
obtained as FCjOg (Koningh, Z. angew,, 1898, 463). Potassium ferri-
cyanide precipitates from dilute solutions of ferrous salts ferrous ferri-
cyanide (g), Fe3(Fe(CN)e)2 (Tumbuirs blue), dark blue, insoluble in acids,
transposed by alkali hydroxides (h): with ferric salts no precipitate is
obtained, but the solution is colored brown or green (t). This is a very
important reagent for the detection of the presence of even traces of
ferrous salts in the presence of ferric salts. As iron is so readily oxidized
or reduced by various reagents the original solution should always Be
tested. The solutions should also be sufficiently diluted to allow the
detection of the precipitate of the ferrous ferricyanide in the presence of
the dark colored liquid due to the presence of ferric salts. If no precipi-
tate be obtained (indicating absence of ferrous iron) a drop of stannous
chloride or some other strong reducing agent constitutes a delicate test
for ferric salts and reconfirms the previous absence of ferrous salts.
Potassium thiocyanate gives no reaction with ferrous salts; with ferric
salts the blood red ferric thiocyanate, Fe(CNS)3 (solution),* is formed (;).
This constitutes an exceedingly delicate test for iron in the ferric condi-
tion (the original solution should always be tested). According to Wagner
(Z., 1881, 20, 350) one part of iron, as ferric salt, may be detected m
1,600,000 parts of water. The red salt of ferric thiocyanate is freely
soluble in water, alcohol, and ether; it is extracted by ether from aqueous
solutions and thus concentrated, increasing the delicacy of the test (Natan-
son. A., 1864, 130, 246). The red color of the liquid is destroyed by
mercuric chloride (^), also by phosphates, borates, acetates, oxalates, tar-
trates, racemates, malates, citrates, succinates, and the acids of these s^lts.
Nitric and chloric acids give red color with potassium thiocyanate, re-
moved by heat.
• The quantity of non-dissociated FetCJiS), , to which the color Is due, is Increased by an ex-
cess of either of the products of the dissociation. The test for iron is therefore more delicate
if considerable KCNS is added. The decoloration by HgCl, is due to the breakingr up of the
Fe(ClV8)a to form Hg(CN8)a which is even less dissociated in water silution than HgCl, .
156 iRoy. §126, 6c:.
(fl) Feci, + 3KCN -h 3H,0 = Fe(OH), + 3KC1 + 3HCN
(6) FefiO* + B:,Fe(CN)« = K^eFe(CN)e + K,SO,
(c) K^FeFeCCN). + 2K0H = Fe(OH), + K^Fe^CN),
(d) 4K^eFe(CN)« + O, + 4HC1 = Fe,(Fe(CN)«), + K^FeCCN). + 4KC1 + ^5,0
(e) 4FeCl, -f 3K,Fe(CN), = Fe,(Fe(CN),), + 12KC1
(0 Fe,(Fe(CN).)s + 12K0H = 4Fe(0H), + aSyPeCCN).
{g) SFeSO, + 2K,Fe(CN)« =Fe,(Fe(CN),), + 3K3SO4
(/r) Fe,(Fe(C?N)o), + CKOH = 3Fe(0H)= + 2K,Fe(CN),
(/) FeCl, + K,Fe(CN)e = FeFe(CN)e + 3KC1
(/) FeCl, 4- SKCNS = Fe(CNS), + 3KC1
(A-) 2Fe(CNS), + 3HgCla = 3Hg(CNS), + 2FeCl,
c. — Nitric acid readily oxidizes all ferrous salts to ferric salts, the reac-
tion being hastened by the aid of heat. As the iron is reduced to the
ferrous condition in the precipitation of the metals of the second group
with hydrosulphuric acid, the oxidation with nitric acid is necessary to
insure the precipitation of all the iron as hydroxide in the third group
(6a and §117).
d. — Hypophosphorons acid reduces ferric salts to ferrous salts. From
solutions of ferrous salts, alkali phosphates, as Na2HP04, precipitate
secondary ferrous phosphate, FeHPO^ , mixed with the tertiary salt,
Fe8(P04).^ , white to bluish white, soluble in mineral acids. By the addi-
tion of an alkali acetate, the precipitate consists of the tertiary phosphate
alone: SFeSO^ + 2Na2HP0, + 2NaC2H30o = Yt^{VO,). + aNa^SO, +
2HC2H3O2 . Ferric salts are precipitated as ferric phosphate, FePO^ ,
scarcely at all soluble in acetic acid, but readily soluble in hydrochloric,
nitric and sulphuric acids.* Hence ferric salts which are not acetates
are precipitated by phosphoric acid with co-operation of alkali acetates:
FeClg + H3PO, + 3NaCoH302 = FePO, + 3NaCl + 3HC2H3O, . If phos-
phates of the fourth <^oup and the alkaline earths be present they are
precipitated with the third group metals by ammonium hydroxide in the
usual course of analysis (§146 and ff.); phosphates of Co, Ni, and Zn beinjj:
redissolved by the excess of ammonium hydroxide. To prevent this gen-
eral precipitation with the metals of the third group, when phosphates
are present, the acid solution (after removal of the second group by hydro-
gen sulphide and the expulsion of the gas by boiling) is treated with an
• Equilibrium requires that a weak acid, as phosphoric, be present for the most part as the
non-dissociated molecule. But FePOf , as any neutral salt, is dissociated, so far as it dissolves
In water, into its ions, as is also the strong hydrochloric acid. Bringing these together will re-
sult In the union of the H ion of the acid and the PO4 ion to non-dissociated H,P<>4, thus
maintaining the equilibrium for H3P04 , but disturbing that between solid and dissolved
PeP04 , which requires a certain concentration of PO4 Ions. To restore the latter more PePO^
dissolves, only to react with the 11 ions as before, and this process continues until the H ions
of the hydrochloric acid are reduced to such small quantity as to be in equilibrium with the
PO4 ions or, if the HCl is in excess, until the FePO^ is entirely dissolved. This process talces
place whenever a strong acid dissolves the ^alt of a weak one. It is analogous to the solution
of a base in an acid, forming non-dissociated water.
§126, 6c. IROy. 157
excess of sodium acetate and ferric chloride is added drop by drop, until
a red color indicates complete precipitation of the phosphate and forma-
tion of ferric acetate. The mixture is then boiled and filtered hot.
Evidently another portion of the solution must be tested for iron. All
of the phosphoric acid present is thus precipitated and separated from
the metals of the remaining groups. Care should be taken to avoid an
excess of the ferric chloride as the ferric phosphate is somewhat soluble
in ferric acetate solution. The alkali hydroxides transpose ferric phos-
phate (freshly precipitated), forming ferric hydroxide and alkali phosphate^
The transposition is not complete in the cold. With fixed alkali hydroxide
aluminum phosphate is dissolved, thus effecting a separation from chrom-
ium and iron. Ferric phosphate warmed with ammonium sulphide forms
ferrous sulphide, ammonium phosphate and sulphur: 4PeP04 + 6(1^4)28
= 4FcS + 4(NHJ3P0, + 82 .
e. — Hydrosulphuric acid is without action upon ferrous salts in acid or
neutral solutions, except a slight precipitate is formed with neutral fer-
rous acetate. Alkali sulphides and H28 in alkaline mixture, form ferrous
sulphide, Fe8, black, insoluble in excess of the reagent, readily soluble in
dilute acids with evolution of hydrogen sulphide. The moist precipitate
is slowly converted, in the air, to ferrous sulphate and finally to basic
ferric sulphate, Fe20(804)2 . Ferric salts are reduced to ferrous salts with
liberation of sulphur by H28 (i), or soluble sulphides, the latter at once
reacting to precipitate ferrous sulphide {2) :
(i) 4reCla -h 2H,S = 4reCl, -f- 4HC1 -f- S,
(2) 4FeCl, -f 6(NH,)aS = 4PeS -|- 12NH,C1 -h S,
After the removal of the metals of the second group by H28, the iron
present will always be in the ferrous condition (it will therefore be neces-
sary to test the original solution to find the condition of the iron at the
beginning of the analysis). The excess of H28 should be removed by
boiling and the iron oxidized by carefully adding nitric acid drop by drop
and boiling until the solution assumes a pale straw color (6&). If this be
done the iron will be completely precipitated in the third group by the
ammonium hydroxide (6a).
Ferrous sulphite is but little soluble in pure water, easily soluble in excess of
sulphurous acidy to a colorless solution. The moist salt oxidizes rapidly oji
exposure to the air (Fordos and Gelis, J. Pharm,^ 1843, (3), 4, 333). Ferric
sulphite is only known as a red solution formed by the action of SO., upon
freshly precipitated Fe(OH), , rapidly reduced to the ferrous condition accord-
ing to the following equation: Fe,(S0.,)3 = FeSO, + FeSA (Gelis, C. C, 1862,
896). Ferrous thlosulphnte, FeS,0,, is formed, together with some FeS and FeSOa,
by the action of SO, upon Fe** (Fordos and Gelis, /. c). Ferric salts are reduced
by sodium thiosulphate to ferrous salts in neutral solutions with formation of
sodium tetrathionate: 2Fe€l, + 2Na.S,0, = 2FeCl3 -f 2iraCl -h NaJSA (Fordos
and Gelis, T. r., 1842, 15, 920); in acid solutions sulphuric acid and sulphur are
formed: iFeQ, + 2NaaSA + 2H,0 = 4FeCl, + 4NaCl + ^JH^SO^ + S, (Men-
158 IRON. §126, 6/.
schutkin, 78). Ferric iron is precipitated as basic nitrate by the addition of a
solution of aznmoniuzn sulphate to a solution of iron in KNO, evaporated to
dryness and taken up with water (separation from aluminum) (Beilstein and
Luther, C. C, 1891, i, 809).
/. — Chlorides and bromides of both ferrous and ferric iron are formed
but only ferrous iodide exists. Ferric salts are reduced to ferrous salt?
by hydriodic acid with liberation of iodine.
g, — Soluble arsenites and arsenates precipitate solutions of ferrous and ferric
salts, forming the corresponding* arsenites and arsenates. Basic ferric arsenite.
4Fe20,.As3'03 4* SBCjO , is formed when an excess of ferric hydroxide is added
to arsenous acid. It is insoluble in acetic acid. It is formed when moist
ferric hydroxide is given as an antidote in case of arsenic poisoning (§69, 0/
and 6'e; D., 3, 352).
h. Ferrous salts are rapidly oxidized to ferric salts by solutions of chro-
mates, the chromium being reduced to the triad condition (9 and 10).
With ferric salts potassium chromate forms a reddish-brown precipitate.
i. — Zinc oxide precipitates solutions of Fe'" , Al , Cr"' and Cu completely and
Pb partially, effecting a separation of these metals from Mn , Co and Ni
(Meineke, Z. angeur,, 1888, 258).
7. Ignition. — The larger number of iron salts are decomposed, as solids, by
heat; FeCl, vaporizes partly decomposed, at a very little above 100**. Igni-
tion in the air changes ferrous compounds, and ignition on charcoal or by
reducing flame changes ferric compounds to the magnetic oxide, which is
attracted to the magnet. Ferrous oxalate ignited in absence of air gives FeO .
Ferric oxide ignited in a current of hydrogen gives Fe^O^ from 330° to 440°, FeO
from 500° to 600°, and Fe° above 600° (Moissan, A. Ch., 1880, (5), 21, 199).
In the outer flame, the borax bead, when moderately saturated with any
compound of iron^ acquires a reddish color while hot, fading and becoming
light yellmc when cold, or colorless, if feebly saturated. The same bead, held
persistently in the reducing flame, becomes colorless unless strongly saturated,
when it shows the pale green color of ferrous compounds. The reactions with
znlcrocosmlc salt are less distinct, but similar. Cobalt, nickel, chromium and
copper conceal the reaction of iron in the bead.
Ferric compounds, heated briefly in a blue borax bead holding a very little
cupric oxide, leave the bead blue; ferrous compounds so treated change the
blue bead to red — the color of cuprous oxide.
8. Detection. — After removal of the first two groups the iron (now in
the ferrous condition) is oxidized by HNO^ and then precipitated in pres-
ence of NH4CI with Al and Cr'" by an excess of NH^OH , The Al is re-
moved by boiling with excess of EOH . If more than traces of Fe be
present it is detected in presence of the Cr(0H)3 , by dissolving in HCl
and obtaining the blood-red solution with ECNS . In case Cr be present
in great excess the Cr(OH),^ and Fe(0H)3 are fused on a platinum foil with
NajCOs and KSO^ , oxidizing the Cr to a chromate soluble in water. After
filtering, the precipitate of FCjO., is dissolved in HCl and tested with KCNS.
The original solution must be tested to determine whether the iron wa-*
present in the ferrous or ferric condition. A portion of the original
solution acidified with HCl gives blood red color with ECNS if Fe'" is
§126, 10. IRON. 15»
present, no color for the Fe". Another portion gives a blue precipitate
with K3Fc(CN)e if Fe'' is present, only a brown or green color for the
Fc'" (66).
9. Estimation.— -(i) After oxidation to Fe'" , if necessary, it is precipitated
with NH4OH , dried, ignited to a dull-red heat and weighed as FezO, . (2) By
precipitation with nitroso-^naphthol in slightly acid solution (Knorre, ^., 1887,
20, 283). Volumetrically: (S) As ferrous iron, by titration with a standard
solution of EMnOt-. IDFeSO* -f 2KMn04 + 8H^0« = SFe^CSO,), + K,S04 +
2MnS04 H- 8HsO . (4) By titration with a standard solution of XsCraOr , using
a solution of KaFoCCN"), as an external indicator: 6FeS04 + KsCTzOt + THjSO* =
3Fe2(S04), + X2S04 H- Cr, (804)8 + THgO . (5) As ferric iron, by titration with
a standard solution of Na^jO,, using XCNS as an indicator: 2FeCl3 + 2N&2S3O,
= 2FeCl3 + Na2S40« + 2N&C1 . A few drops of a solution of C11SO4 are added,
which seems to hasten the reaction and gives more accurate results; or use
excess of the NaaSjO, and titrate back with standard iodine (Crafts, J. C, 1873,
26, 1162). (6) The iron as ferric salt is treated with an excess of a standard
SnClj solution, the excess of the SnCl, being determined by a standard solution
of iodine in potassium iodide: 2FeCl8 + SnCl, = 2F6CL + SnCl4 . (7) Potas-
sium iodide is added to the nearly neutral ferric chloride; the flask is stoppered
and warmed to 40°. The iodine set free is titrated by standard NaaSzO^
(very accurate for small amounts of iron). (8) When present in traces it is
determined colorimetrically as Fe(GNS)s in etherial solution (Lunge, Z. angetc^
1894, 669).
10. Oxidatio^. — Metallic iron precipitates the free metals from solu-
tions of Au , Pt , Ag , Hg , Bi , and Cu (separation from Cd).
Solutions of Fe" are changed to Fe'" solutions by treating with solutions
of Au , Ag , Cr^, Mn^, Mn^, and HjOj . In presence of some dilute
acid, such as HjSO^ or H3PO4 by PbOj, PbgO^, MngO^, MnOa, Hn^Oa,
CO2O3 , NijOa . The following acids also oxidize Fe" to Fe'", HNO2 , HNOg ,
HCIO , HCIO2 , HCIO, , H.SO, (if concentrated and hot), HBrO , HBrOs
HIO3 , also Br , CI . Br and CI in presence of KOH changes Fe" and Fe'"
to XzTeO^^ . Barium ferrate is the most stable of the ferrates ; they are
strong oxidizers, acting upon nitrites, tartrates, glycerol, alcohol, ether>
ammonia, etc. (Eosell, J. Am. Soc, 1895, 17, 760).
Fe'" is reduced to Fe" by solutions of Sn", Cu', H3PO2 , H3PO3 , H2S ,
H2SO3 , Na2S203 , and HI . Also by nascent hydrogen, or by any of the
metals which produce hydrogen when treated with acids, including Pb,
As , Sb , Sn , Bi , Cu *, Cd , Fe , Al , Co , Ni , Zn , and Mg f.
• Carnegie, J. C, 1888, 58, 468. t Warren, C. N., 1889, 60, 187.
160
AXALTSIS OF THE IRON OROCP.
§127.
§127. Table for Analysis of the Iron or Third Group (Phosphates
and Oxalates being absent). See §312.
To the clear filtrate from the Second Group, in which HjS will cause no pre-
cipitate (§80), and freed from HjS by boiling, add a few drops of Nitric
Acid and boil an instant (to oxidize ferrosum*). Immediately add
Ammonium Chloride (§134, 5b; §189, 56) and an excess (§135, 6a) of
Ammonium Hydroxide (§116). If there is a precipitate, filter and wash.
Precipitate: Al(OH), , Cr(OH), , Fe(OH). .
Pierce the point of the filter, and with a little water wash the precipitate
into a casserole or evaporating" dish; add a few drops of Potassium or
Sodium Hydroxide and boil for several minutes. If a residue remains, filter
and wash.
Besidue: Cr(OH), , Fe(OH), .
Fuse a portion of the residue on a platinum foil
with potassium nitrate and sodium carbonate,
cool, digest in warm water and filter (§125, 7).
Besidue: FszO, .
Dissolve the residue in
HCl and test for iron
with potassium thio-
cyanate (§126, 66).
If the residue after re-
moval of the aluminum
does not indicate an ex-
cess of Cr by its green
color, it may be dis-
solved in HCl and test-
ed for the blood-red
color with KCNS .
Iron being found, to de-
termine whether it is
ferric or ferrous, or
botht, in the original
solution, test the latter,
after acidulating with
hydrochloric acid, with
KCNS for ferricum,
and with K^Ye(CN)c for
ferrosum (§126, 66).
Study §136, §128, §129,
§130 and §131.
Solution: Na2Cr04,
K,CrO, (NaaCO,) .
Acidify with 'HCJS.fli and
precipitate the chro-
mium as lead chromate
(yellow) with a solu-
tion of lead acetate
(§57, 6/1).
If the original solution
contains a chromate it
will be j-ellow (normal
chromate), or red (acid
chromate), and will
give the reactions for
chroma tcs with
ThiC^.O^, , BaCl, ,
etc. (§125, 6h). If the
chromium is present as
a chromic salt,Crj(SOJ„
the solution will have
a green or bluish-green
color and will give the
general reactions as de-
scribed at §125, 6.
Chromates should be re-
duced by boiling with
HCl and CjH^OH be-
fore proceeding with
the regular course of
analysis (§125, 60-
Study §136, §128, §129,
§130, §131.
Solution: KALO, .
Make the solution slight-
ly acid with hydro-
chloric acid, and then
add ammonium car-
bonate. A precipitate
is Al(OH), .
The same result is ob-
tained with nearly
equal certainty by add-
ing an excess of NH^Cl
to the alkaline solution
(§124, 6a; §130).
Lead and antimony give
similar results if
(through carelessness)
they have not been
removed (§131, 6).
Study §136, §128, §129,
§131, 6, and §124, 6.
* In the filtrate from the Second Oronp iron is necessarily in the ferrous condition (186 6*-).
•+ Ferrous salts, which have been kept in the air, are never wholly free from ferric compound?.
§129,8. DIRECTIONS FOR ANALYSIS WITH NOTES, 161
Directions for the Analysis of the Metals of the Third Group.
§128. Manipulation. — Boil the filtrate from the second group (§80) to
expel the HjS and then oxidize any ferrous iron that may be present by
the addition of a few drops of HNO3 , continuing the boiling to a clear
straw-colored solution (§126, Gr):
3FeS0« -f 4HN0, = Fe.CSO*), + FeCNO,), + NO + 2H,0
Add to the solution about one-half its volume of NH4CI (5&, §§134 and
189) and warm and then add NH4OH in a decided excess (§135, 6a):
KLgCU + NH.Cl -f NH,OH = NH^MgCls -f NH4OH
Fe^CSOJ, -f 6NH,0H = 2Fe(0H), + SCNHJ^SO,
Z11SO4 + 4NH,0H = (NH,).JZnO, + (NH,),SO, -f 2H,0
Heat nearly to boiling for a moment, filter, and wash with hot water.
Notice that the filtrate has a strong odor of ammonium hydroxide and
set aside to be tested for the metals of the succeeding groups (§138).
§129. Notes. — (1) If the HjS is not all expelled, it becomes oxidized by the
HNO, with deposition of a milky precipitate of sulphur (§257, 6^), which
tends to obscure the reactions following: OHjS -f 4HN0g = SS^ -f 4N0 + SHjO.
Also any HjS not decomposed by the HNOg would cause a precipitate of the
sulphides of the fourth group upon the addition of the NH«OH: H^S -f NiClj +
2NH,0H = NiS 4- 2NH,C1 + 2H,0 .
(2) Any iron that may have been present in the original solution in the
ferric condition is reduced to the ferrous condition by the H3S (§126, Ge):
4FeGl, -f 2H8S = 4FeCl2 + S^ + 4HC1 . The ferrous hydroxide is not com-
pletely insoluble in the ammonium salts present (§117), and hence unless the
oxidation with the HNOg be complete, some of the iron will be found in the
next group.
(5) If considerable iron be present the solution becomes nearly black upon
addition of nitric acid, due to the combination of the nitric oxide with the
ferrous iron (§241, 8a). Therefore the boiling, and addition of HNO, , a drop
or two at a time, must be continued until the solution assumes a bright straw
color.
(-J) If nitric acid be added in excess there is danger that Mn will be oxid-
ized to the triad or tetrad condition then it is precipitated with iron in the
third group (§134, 6fl). The careful addition of the nitric acid (avoiding an
excess) prevents this oxidation of the manganese.
(J) Ammonium hydroxide precipitates a portion of Mn (§134, 6«) and Mg
(§189. 6a)y but these hydroxides are soluble in NH4GI (5c, §§134 and 189);
hence if that reagent be added in excess the Mn (§134, 6«) and Mg are not at
all precipitated by the NH«OH:
2MnCl, -f 2NH,0H = Mn(OH), -f (NH,),MnCl,
Mn(OH), -f 4NH,C1 = (NHJ^MnCl^ + 2NH,0H
2MgCl, 4- 2NH,0H = KLgiOB.), + NH.MgCl, + NH.Cl
Mg(OH), + 3NH,C1 = NH,MgCl, + 2NH,0H
(6) Ammonium chloride lessens the solubility of A1(0H)8 in the NH4OH
solution and effects an almost quantitative precipitation of that metal (§117).
(7) NH4OH precipitates solutions of Co , Ni and Zn , but these precipitates
are readily soluble in an excess of the NH4OH (§116). To insure the presence
of an excess of 1I'H40H the odor should be noted after shaking the test tube
and after the solution has been heated.
(/?) The precipitates of the hydroxides of Al , Cr and Fe'" filter much more
rapidly if the precipitation takes place from a hot solution (§124, 4 and 6a).
162 DIKECl'IOy& FOR ANALYSIS WITH NOTES, §129, 9,
(9) In the presence of chromium the filtrate from the third group is usually
of a slight violet color, due to the solution of a trace of chromium hydroxide
in the NH^OH (§125, 6fi). Boiling the solution to remove excess of ammonia
prevents this.
(10) A small portion of the filtrate of the second group after the removal of
the H,S by boiling should be tested for the presence of phosphates by am-
monium molybdate (§75, 6d). If phosphates are found to be present, the
method of analysis of the succeeding groups must be considerably modified.
These modifications arc fully discussed under §145 to §153.
§130. Manipulation. — The well washed precipitates of Al , Cr , and Fe'"
hydroxides are transferred to a small casserole or evaporating dish by
piercing the point of the filter and washing the precipitate from the filter
with as small an amount of water as possible; and then boiled for a
minute or two with an excess of NaOH :
Al(OH), + NaOH = NaAlO, -f 2H,0
Cr(OH), -f NaOH = NaCrO, -f 2H,0 (in the cold)
NaCrO, + 2HaO = Cr(OH)s -f NaOH (upon boiling)
The alkaline liquid is filtered (§131, 1) (the filtrate is reserved
to be tested for aluminum), and the remaining precipitate fused on a
platinum foil with a mixture of equal parts of KNO3 and NajCOs : 2Cr(0H)3
+ 2KNO3 + Na^COs = K^CrO, + Na^CrO, + 2N0 + CO^ + 3H,0
(§125, 7). The fused mass is then dissolved in water, filtered, rendered
acid with acetic acid and tested for chromium with Pb(CoH30o)2 , a yellow
precipitate at this point being sufficient evidence of the presence of
chromium : Na.CrO^ + KoCrO^ + 2Pb(C2H30o)2 = 2PbCrO^ + 2NaC2H30,
+ 2KC2H3O2 (§67, Gh).
The residue of the fused mass not soluble in water should be washed
with hot water and then dissolved in HCl : FCoOs + 6HC1 = 2FeCl3 -f
3H2O , and tested for iron with KCNS : FeClg + 3KCNS = Fe(CNS)3 -f-
3KC1.
If iron has been found to be present, the original solution acidulated
with HCl (or a few drops of the filtrate from the first group) should be
tested with ECNS for the presence of ferric iron (§126, 6&) and with
K3Fe(CN)o for the dark blue precipitate of Fe3(Fe(CN)g)2 indicating the
presence of ferrous iron (§126, 6&): 3FeS0^ + 2K^'Ee{ClSt)^ = Fe3(Fe(CN)o)2
+ 3X280^ .
The alkaline filtrate obtained after boiling the precipitated hydrox-
ides with NaOH, is slightly acidulated with HCl: KAIO2 + 4HC1 =
AICI3 + KCl + 2H2O , and then precipitated with (NHJoCOa , a white
gelatinous precipitate being evidence of the presence of aluminum:
2AICI3 + 3(NH,)2C03 + 3H2O = 2A1(0H)3 + 6NH,C1 + 3CO2 . Or an
excess of NH4CI may be added directly to the alkaline filtrate, giving the
white gelatinous precipitate of aluminum oxide-hydroxide: 2EAIO2 +
2NH,C1 + H2O = Al20(0H), + 2KC1 + 2NH3 (§124, 6a).
§132, 1. COBALT, 1(>3
§131. Notes, — (1) Chromium hydroxide when precipitated from solutions of
pure chromic salts by NaOH is readily soluble in an excess of the cold reagent
(§125, 6a); but in presence of ammonium salts or of ferric hydroxide the
chromium hydroxide is not completely soluble in a cold solution of the fixed
alkali. This prevents the use of the cold fixed alkali as a means of separation
of Cr and Al from Fe'" . The student is therefore directed to boil the mixture
of these three hydroxides with NaOH , thus precipitating the whole of the
chromium and effecting" a quantitative separation of Cr and Fe'*' from Al . If
the alkaline liquid is too concentrated to filter, it must be diluted with water.
(2) Unless the precipitate of the hydroxides is a very dark green, due to
the presence of a large amount of chromium, a portion of the precipitate should
be dissolved in HCl and tested with KCNS for the presence of iron. The
presence of a moderate amount of chromium does not interfere.
(3) In the absence of chromium the presence of more than traces of iron
gives a brown color to the ammonium hydroxide precipitate (§126, 6fl), alu-
minum hydroxide being a white gelatinous precipitate.
(.)) If the fused mass has a green color, manganese (§134, 7) is evidently
pres ent in large quantities and was not completely separated by the NH4GI
and NH«OH (§134, 6a). By dissolving the fused mass in water and carefully
warming with HCl, the manganate, X2Mn04 , may be reduced (a) (§134, 5c)
Avithout effecting a reduction of the chromate, which may be precipitated as
BaCr04 by BaClt after neutralization with NH4OH . Or the fused mass may
be warmed with hydrochloric acid and alcohol, effecting complete reduction (?>),
and this solution again precipitated with NH4OH , which will prevent more
than traces of the manganese from being precipitated with the third group
hydroxides. If again upon fusion with XNO, and X2GO, a green mass is
obtained, the operation should be repeated:
(a) K,MnO« -f 8HC1 = MnCl, -h 2KC1 + 2C1, + -IH^O
(6) 2K,Cr04 + lOHCl -f 3C,H,0 = 2CrCl, + 4KC1 -f 3C2H4O -f ^H^O
(5) The presence of chromium as chromic salts is usually indicated by the
preen or bluish-green color of the original solution. Chromium as chromntes
(red or yellow) should be reduced to chromic salts by boiling with HCl and
CsHflO before proceeding with the regular group separations (§125, Ge and f).
H.S will effect this reduction but gives also a precipitate of sulphur which
should be avoided when convenient to do so: ^K.Cr.Or -|- 16HC1 -h 'iH^S =
4CrCl, + 4KC1 -f 3S2 -f I4H2O .
(6) Too much stress cannot })e laid upon the necessity for removing all the
metals of one group before testing the filtrate for the metals of the next
succeeding group. If through lack of sufficient HjS or too much HCl , lead or
antimonj' are not completely removed in the second group, they will give all
the reactions for aluminum' (§57, Oa, and §70, Ca); hence as a safeguard it is
advised to test the white precipitate, indicating aluminum, with HoS . A
black or orange precipitate is evidence of unsatisfactory work and the student
should repeat his analysis.
(7) The presence of a trace of white precipitate in the final test for aluminum
may be due to the presence of that metal in the fixed alkali (§124, 6rt, footnote),
or it may be caused by the use of too concentrated fixed alkali, which may
dissolve silica from the glass of the test tubes or remove it from the filter
paper (§249, 5).
The Zixc Group (Fourth Group).
Cobalt, Nickel, Manganese, and Zino.
§132. Cobalt. Co = 59.0(1 . Usual valence two and three.
1. Properties. — Fipcciflc prantjt^ powder from the oxide reduced by hydrogen,
mean of five samples, 8.957 (Rammelsberg, Pogg., 1849, 78, 93); melting point,
1500** (Pictet, C. r., 1879, 88, 1317). Cobalt is similar to iron in appearance, is
164 COBALT. §132, 3.
harder than Fe or Ni . It is malleable, very ductile and most tenacious of any
metal, the wire being about twice as strong as iron wire (Deville, A, Ch., 185t*,
(3), 46, 202). The fine powder oxidizes in the air quite rapidly and may even
take fire spontaneously; in a compact mass it is but little tarnished in moist air.
At a white heat it burns rapidly to GOsO^ . It is attracted by the magnet and
can be made magnetic, retaining (unlike steel) its magnetism at a white heat.
2. Occurrezice. — Cobalt does not occur in a free state, except in meteoric
iron. It is found in linnaeite (C0JS4); skutterudite (CoAs,); speiss cobalt
(CoNiFeAsJ; glance cobalt (GoFeAsSj); wad (Go.MnO,2Mii03 + 4H3O): etc.
3. Preparation. — (/) By electrolysis of the chloride. (2) By heating with
potassium or sodium. (S) By heating any of the oxides, hydroxides or the
chloride in hj-drogen gas. {^) By fusion of the oxalate under powdered glass.
(5) Also reduced by carbon in various ways.
4. Oxides and Hydroxides. — Cobaltous oxide, CoO , is made (/) by heating
any of its oxides or hydroxides in hydrogen to (not above) 350°; (i) by ignition
of Go(OH)3 or CoGO, , air being excluded; (3) by heating GOsO^ to redness in
a stream of CO3 (Russell, J. C, 1863, 16, 51); (-)) by heating any of the higher
oxides to a white heat (Moissan, A. Ch., 1880, (5), 21, 242). Cobaltous hydrox'uir
is made from cobaltous salts by precipitation with fixed alkalis; oxidizes if
exposed to the air (6rt). The most stable oxide is the cohaltoso-i'obaltic ( 00,04)
tricohalt letroxide; it is made by heating any of the oxides or hydroxides, the
carbonate, oxalate or nitrate to a dull-red heat in the air or in oxygen gas.
Several oxide-hydroxides are known, r. </., Go,Oa(OH)4, GosO(OH)fl, Co'sd,(OH);.
Cohaliic oxide, Co.O, , is made by heating the nitrate just hot enough for de-
composition, but not hot enough to form Go,04 . Cohaltic hydroxide, Go(OH)3 .
is made by treating any cobaltous salt with 01 , HOlO , Br or I in presence of
a fixed alkali or alkali carbonate. It dissolves in HOI with evolution of chlo-
rine, in H3SO4 with evolution of oxygen, forming a cobaltous salt. CoO, has
not yet been isolated, but McConnell and Hanes {J. C, 1897, 71, 584) have
shown that it exists as HoCoO, and in certain cobaltites.
5. Solubilities. — a. — Metal. — Slowly soluble on warming in dilute HOI or
H3SO4 , more rapidly in HNOs , not oxidized on exposure to the air or when
heated in contact with alkalis. Like iron, it mav exist in a passive form
(Nickles, ./. pr., 1854, 61, 168; St. Edme, C. r., 1889, 109, 304). With the halogens
it forms cobaltous compounds (Hartley. ./. C. 1874, 27, 501). h. — Oxides and
hydroxides. — Cobaltous oxide (gray-green) and hydroxide (rose-red) are in-
soluble in water; soluble in acids, in ammonium hydroxide, and in concentrated
solutions of the fixed alkalis when heated (Zimmerman, .4., 1886, 232, 324):
the various higher oxides and hydroxides are insoluble in ammonium hydroxide
or chloride (separation from nickelous hydroxide after treating with iodine
in alkaline mixture) (Donath, Z., 1881, 20, 386), and are decomposed by acids,
evolving oxygen with non-reducing acids, or a halogen from the halogen acids,
and forming cobaltous salts. Oo.Of is said to be soluble in acids with great diffi-
culty (Gibbs and Geuth, Am. /?., 1857, (2), 23, 257). c.—^alts.—Coha\i forms two
classes of salts: eolxtltous, derived from OoO , and cohaltic, from Oo.Os . The
latter salts are quite unstable, decomposing in most cases at ordinary tem-
peratures, forming cobaltous salts. The cobaltous salts show a remarkable
variation of color. The crystallized salts with their water of cr>'stallization
are pink; the anhydrous salts are lilac-blue. In dilute solution the salts are
pink, but most of them are blue when concentrated or in presence of strong
acid. A dilute solution of the chloride spreads colorless upon white paper,
turning blue upon heating and colorless again upon cooling, used as " sympa-
thetic ink."
Cobaltous nitrate and acetate are deliquescent', chloride, hygroscopic; sulphate,
efflorescent. The chloride vaporizes, imdecomposed, at a high temperature.
The carbonate, sulphide, phosphate, borate, oxalate, cyanide, ferrocyanide
and ferricyanide are insoluble in water. The potassium-cobaltous oxide is in-
soluble; the ammonio-cobaltous oxide, and the double cyanides of cobalt and the
alkali metals, soluble in water. Alcohol dissolves the chloride and nitratip;
ether dissolves the chloride, sparinglv, more so if the ether be saturated with
HOI gas (separation from Ni) (Pinerfia, C r.. 1897, 124, 862). Most of the
salts insoluble in water form soluble compounds with ammonium hj'droxide.
|132, 66. COBALT, 165
6. Beactions. a. — The fixed alkali hydroxides precipitate, from solu-
tions of cobaltous salts, blue basic salts, which absorb oxygen from the air
and turn olive green, as cobaltoso-cobaltic hydroxide; or if boiled before
oxidation in the air, become rose-red, as cobaltous hydroxide, Co(0H)2 .
The cobaltous hydroxide is not soluble in excess of the reagent, but is
somewhat soluble in a hot concentrated solution of EOH (distinction from
Hi) (Beichel, Z,, 1880, 19, 468). Freshly precipitated Pb(0H)2 , Zn(0H)2 ,
and HgO precipitate Co(OH)2 from solutions of cobaltous salts at 100°.
Ammonium hydroxide causes the same precipitate as the fixed alkalis;
incomplete, even at first, because of the ammonium salt formed in the
reaction, and soluble in excess of the reagent to a solution which turns
brown in the air by combination with oxygen, and is not precipitated by
potassium hydroxide. The reaction of the precipitate with ammonium
salts forms soluble double salts (as with magnesium) ; the reaction of the
precipitate with ammonium hydroxide produces, in different conditions,
■dififerent soluble compounds noted for their bright colors, as (NH3)4CoCl2 ,
<NH3)eCoCl2, (NH3),CoCl3 , etc.
Alkali carbonates precipitate cobaltous basic-carbonate, 00,^05(003)3 ,
peach-red, which when boiled loses carbomic anhydride and acquires a
violet, or, if the reagent be in excess, a blue color. The precipitate is
soluble in ammonium carbonate and very slightly soluble in fixed alkali
carbonates. Oarbonates of Ba , Sr , Oa , or Mg do not precipitate cobaltous
chloride or nitrate in the cold (separation from Fe'", Al, and Or"'), but
by prolonged boiling they precipitate them completely. However, if a
solution of a cobaltous salt be treated with chlorine, a cobaltic salt is
formed (5a), which is precipitated in the cold on digestion with Ba003
{distinction from Ni).
6. — Oxalic acid and oxalates precipitate reddish-white cohaUous oxalate,
CoCa04 , soluble in mineral acids and in ammonium hydroxide.
Alkali cyanides — as EON — precipitate the brownish-white cobaltous
cyanide, 0o(0N)2 , soluble in hydrochloric acid, not in acetic or in hydro-
cyanic acid, soluble in excess of the reagent, as double cyanides of cobalt
and alkali metals — (K0N)20o(0N)2 — potassium cobaltous cyanide, the solu-
tion having a bro\\Ti color: OoOlj + 2K0N = 0o(0N)2 + 2K01 . Then
Co(OH)2 + 2K0N = (K0H)20o(0N)2 • Dilute acids, without digestion,
reprecipitate cobaltous cyanide from this solution (the same as with Ni) :
(K0N)20o(0N)2 + 2H01 = 0o(0N)2 + 2H0N + 2K01 . But if the solu-
tion, with excess of the alkali cyanide and with a drop or two of hydro-
chloric acid,* insuring free HON , be now digested hot for some time, the
* Moore (C. 2V., 1887, 56, 8) adds erlacial pbosphorid acid to the neutral solutions of cobalt and
nickeU until the precipitate first formed be^ns to redlssolve ; then he adds KCN and boils,
continulnfir the boiling and addition of KCIC until KOH fails to give a precipitate. He then
warms with excess of bromine in presence of KOH, whereupon the nickel is completely pre-
cipitated leavlufir the «obalt In solution. See Also Hambly (C N., 1893, 60, 280).
166 COBALT, §182, Sc.
cobaltous cyanide is oxidized and converted into alkali cobalticyanide — ^as
K8Co(CN)5 — corresponding to ferricyanides, hut hiving no corresponding
nickel compound:
4Co(CN)a + 4HGK + O, = 4Co(Cir), (cobaltic cyanide) + 2HaO
Go(CN), + 3KCN = K,Co(CN)« (potassium cobalticyanide).
In the latter solution acids cause no precipitate {important distinction from
nickel, whose solution remains (KCN)2Ni(CN)2 , and after digestion as
above is precipitated with acids). The potassium cobalticyanide solution,
after removal of the Ni , may be precipitated with HgNO, (Gibbs, J. C,
1874, 27, 92). The oxidation of the cobalt may be hastened by the pres-
ence of chromic acid, which is reduced to trivalent chromium compound:
6Co(ClS02 + 24KCN + 2Cr03 + SH^O = 6K3Co(CN)e + Cr.Oa + 6K0H
(McCulloch, C. iV., 1889, 69, 51).
FerrocTanideSy as K^FeCCN)^ , precipitate cohaltous fenocyanide, COsFeCCN). ,
gray-green, insoluble in acids. Ferricyanides, as KsFe(GN)a , precipitate cobalt-
ous ferricyanide, Co,(Fe(Cir),), , brownish-red, insoluble in acids. But a more
distinctive test is made by adding ammonium chlo-ride and hydroxide^ with the
ferricyanide, when a blood-red color is obtained, in evidence of cobalt (distinc-
tion from nickel). Potassium xanthate forms a green precipitate in neutral or
slightly acid solutions of cobalt .salts (§133, Gb),
Nitroso-/9-naphthol completely precipitates solutions of Cu , Fe , and Co ;
Ag , Sn , and Bi salts are partially precipitated; and Pb , Eg , As , Sb , Cd ,
Al, Cr, Mn, Ni, Zn, Ca, Mg, and Gl remain in solution (Burgass, Z.
angew,y 1896, 59()). In analysis for the separation of cohalt and nickel it is
recommended to proceed as follows : The metals preferably as sulphates or
chlorides are acidulated with hydrochloric acid and treated with a hot
solution of nitroso-/?-naphthol in 50 per cent acetic acid, until the whole
of the cobalt is precipitated. The brick-red precipitate is then washed with
cold HCl , then with hot 12 per cent HCl , and finally with water. The
separation is quantitative. The precipitate may be ignited in air to the
oxide or with oxalic acid in an atmosphere of hydrogen and weighed as
the metal. For qualitative purposes the cobalt in the precipitate may be
identified by the color of the borax bead (7). The nickel in the filtrate
may be precipitated by hydrosulphuric acid and identified by the usual
tests (Knorre, 5., 1887, 20, 283 and Z, angew., 1893, 264).
c, — Potassium nitrite forms with both cobaltous and nickelous salts the
double nitrites, Co (N02)2.2KN02 and Ni(N02)2.2KN02 , soluble. The nickel
compound is very stable, but if the cobalt compound, strongly acidulated
with acetic acid, be warmed and allowed to stand for some time, preferably
twenty-four hours; the cobalt is completely precipitated as the yellow
crystalline potassium cobaltic nitrite, Co(N02)3.3KN02 (separation from
Ni): C0CI2 + 6KN0. + HC2H3O2 + HNO2 = Co(N02)s.3KN02 + 2KC1 +
KC2H3O2 + H2O + NO.
§132, 7. COBALT. 167
d.-^PhosphateSy as yaaHP04 , precipitate cobaltous salts as the reddish
cohalious phosphate, G0HPO4 , soluble in acids and in ammonium hydroxide.
Sodium pyrophospliate forms a gelatinous precipitate with solutions of cobalt
salts* soluble in excess of the reagent. The addition of acetic acid causes a
precipitation of the cobalt even in the presence of tartrates (separation from
TSfi , but not from Mn or Fe) (Vortmann, B., 1888, 21, 1103). If a solution of
cobaltous salt be treated with a saturated solution of ammonium phospliate
and hydrochloric acid, and when hot treated with an excess of ammonium
hydroxide, a bluish precipitate of C0NH4FO4 will appear on stirring (separa-
tion from nickel ♦) (Clark, C. N., 1883, 48, 262; Hope, J. Soc, Ind., 1890, 9, 375).
e. — ^HydroBTilphnric acid, with normal cobaltous salts, gradually and
imperfectly precipitates the black cobalt sulphide, CoS ; from cobalt acetate,
the precipitation is more prompt, and is complete; but in presence of
mineral acids, as in the second group precipitation, no precipitate is made.
Immediate precipitation takes place with hydrosulphuric acid acting upon
solutions of cobaltous salts in ammonium hydroxide. When formed, the
precipitate is scarcely at all soluble in dilute hydrochloric acid or in acetic
acid ; slowly soluble in moderately concentrated hydrochloric acid ; readily
soluble in nitric acid; and most easily in nitrohydrochloric acid. By
exposure to the air, the recent cobaltous sulphide is gradually oxidized to
cobalt sulphate, soluble, as occurs with iron sulphide (§126, 6e). Alkali
sulphides precipitate immediately and perfectly the black cobaltous sul-
phide, described above, insoluble in excess of the reagent. When cobaltous
salts are boiled with sodium thiosulpliate a portion of the cobalt is precipi-
tated as the black sulphide.
f. — The higher oxides of cobalt and cobaltic salts are reduced by warming
with halogen acids, liberating the corresponding halogens (HCl does not reduce
the cobalt in X,Co(Cll')«).
g, — Soluble arsenites and arsenates precipitate cobaltous salts, forming the
corresponding cobalt arsenites or arsenates, bluish-white, soluble in ammonium
hydroxide or in acids, including arsenic acid. h. — Soluble chromates precipi-
tate cobaltous chromate, yellowish-brown, soluble in ammonium hydroxide and
in acids, including chromic acid. No precipitate is formed with potassium
dichromate. i. — KMn04 added to an ammoniacal solution of cobaltous salts
oxidizes the cobalt and prevents its precipitation by XOH (separation from
ITi) (Delvaux, C. r., 1881, 92, 723).
/.—Cobaltous salts in ammoniacal solution, warmed with HzO, and then
rendered acid with acetic acid, are precipitated by ammonium molybdate
(separation from Ni) (Carnot, C. r., 1889, 109, 109).
7. Ignition. — In the bead of borax, and in that of microcosmic salt, with
oxidizing and with reducing flames, cobalt gives an intense blue color.
The blue bead of copper changes to brown in the reducing flame. If
strongly saturated, the bead may appear black from intensity of color, but
will give a blue powder. This important test is most delicate with the
borax bead. Manganese, copper, nickel, or iron interfere somewhat. By
igfnition, with sodium carbonate on charcoal or with the reducing flame,
* Krauss (Z., 1801, 80, 227) gives a good review of the most important methods for the separa-
tion of cobalt and nickel.
168 NICKEL, §182, S.
compounds of cobalt are reduced to the metal (magnetic). Cobaltous
oxide dissolves in melted glass and in other vitreous substances, coloring
the mass blue — used to cut oflf the light of yellow flames (§205, 7). The
black cobaltoso-cobaltic oxide, COgO^ , as left by ignition of cobaltous oxide
or nitrate, combines or mixes, by ignition, with zinc oxide from zinc com-
pounds to form a green mass, with aluminum compounds a blue, and with
magnesium compounds a pink mass.
8. Detection. — After removal of the metals of the first three groups
cobalt is precipitated by HgS in ammoniacal solution with Ni , Mn and Zn .
The sulphides are digested with cold dilute HCl which dissolves the Hn
and Zn . The borax bead test (7) is now made upon the remaining black
precipitate, and if Ni be not present in great excess the characteristic blue
bead is obtained. If the nickel be present in such quantities as to obscure
the blue borax bead the sulphides are dissolved in hot cone. HCl , using a
few drops of HNO3 . The solution is heated to decompose all the nitric
acid and, after dilution, the cobalt is precipitated with nitroso-/9-naphthol,
according to directions given in 6&, and further identified by the bead test.
9. Estimation. — (i) As metallic cobalt, all compounds that may be reduced
by ignition in hydrogen gas, e. g., CoCl, , Co (NO.) 3 , GoCO, , and all oxides and
hydroxides. {2) As CoO , all soluble cobalt salts, all salts whose acids are
expelled or destroyed by ignition, all oxides and hydroxides. The salt is con-
verted into Co(OH)j by precipitation with a fixed alkali, and ignited in a
stream of CO;^ . The carbonate and nitrate may be ignited directly in CO^ ,
and organic salts are first ignited in the air until the carbon is oxidized, and
then again ignited in CO2 . (3) After converting into a sulphate it is ignited
at a dull-red heat and weighed as a sulphate. (-)) After converting into the
oxalate, titrated with KMn04 . (.5) In presence of nickel, it is oxidized in
alkaline solution by HjOj , KI and HCl are added, and the liberated iodine
titrated with sodium thiosulphate (Fischer, C. C, 1889, 116). (6) Electroly-
tically. (7) Separated from nickel by iiitroso-/?-iiaph.thol, and after ignition
in hydrogen weighed as the metal (6ft).
10. Oxidation. — Co'' is oxidized to Co'" in presence of a fixed alkali by
PbOj, CI, KCIO, Br, KBrO, I and HjOj*; in presence of acetic acid by
KlSrOj (6c). Co'" is reduced to Co" by H^C^O, , H3PO2 , H^S , H^SO^ , HCl ,
HBr , and HI . Metallic cobalt is precipitated from solution of CoCL by
Zn , Cd , and Hg .
§133. Nickel. Ni = 58.70 . Usual valence two and three.
1. Properties.— fifp^i7?c gravity, 8.9 (Schroeder, Pogg., 1859, 106, 226). Melting
poini, 1450*' (Pictet, C. r., 1879, 88, 1317). It is a hard white metal, capable of
taking a high polish; malleable, ductile and very tenacious, forming wire
stronger than iron but not quite so strong as cobalt (§132, 1). It does not
oxidize in dry or moist air at ordinary temperatures. It is magnetic but loses
its magnetism like steel on heating to redness (Gangain, C. r., 1876, 83, 661).
It burns with inoafidescence when heated in O , CI , Br or S . It is much
• Durrant, C. JT., 1897, 75, 43w
§138, Ga. NICKEL. 169
used in plating other metals, in making coins of small denominations, in
hardening armor plate, projectiles, etc. The pretence of small amounts of
phosphorus or arsenic renders it much more fusible, without destroying its
ductility; a larger amount makes it brittle.
2. Occurrence. — Nickel almost always occurs in nature together with cobalt.
It is found as millerite, NiS,; as nickel blende, NiS; as iron nickel blende,
NiPeS; as cobalt nickel pyrites, (NiCoEe),S4 , etc.
3. Preparation.— (i) By electrolysis. (2) By heating in a stream of hydrogen.
The oxide is reduced in this manner at 270° (W. Muller, Pogy., 1869, 136, 51).
(3) By fusing the oxalate under powdered glass (CO, being given off).
(4) Reduction by igniting in CO . (5) Reduction by fusing with carbon in a
variety of methods. (6) By heating the carbonyl,* Ni(C0)4 to 200°.
4. Oxides and Hydroxides. — Xii^kclous oxide is formed when the carbonate,
nitrate, or any of its oxides or hydroxides are strongly ignited. Nickelouif
hudrojride is formed by precipitation of nickelous salts with fixed alkalis.
Nickelie oxide, NijO. , is made from NiCO, , Ni(NO,),.or NiO by heating in the
air not quite to redness, with constant stirring. It is changed to NiO at a red
heat. Xickelic hydroxide, Ni(OH)s , is formed by treating nickelous salts
first with a fixed alkali hydroxide or carbonate and then with CI , NaClO , Br
or NaBrO (not formed by iodine), a black powder forming no corresponding
salts (Campbell and Trowbridge, J. Anal., 1893, 7, 301). A trinickelic tetroxide,
NisO^ , magnetic (corresponding to COgO^ , "Fe^O^ , Mn,04 and PbaOJ, is formed,
according to Baubigny (C. r., 1878, 87, 1082), by heating NiClz in oxygen gas
at from 350° to 440°; and by heating Ni,0, in hydrogen at 190° (Moissan, A. Ch.,
1880, (5), 21, 199).
5. Solubilities. — a. — Metal. — Hydrochloric or sulphuric acid, dilute or con-
centrated, attacks nickel but slowly (Tissier, C. r., 1860, 50, 106); dilute nitric
acid dissolves it readily, while towards concentrated nitric acid it acts very
similar to passive iron (Deville, C, r., 1854, 38, 284). It is not attacked when
heated in contact with the alkali hydroxides or carbonates, h. — Oxides and
hydroxides, — Nickelous oxide and hydroxide are insoluble in water or fixed
alkalis, soluble in ammonium hydroxide and in acids. Nickelie oxides and
hydroxides are dissolved by acids with reduction to nickelous salts, with halogen
acids the corresponding halogens are liberated. The moist nickelie hydroxide,
formed by the action of CI , Br , etc., in alkaline solution, after washing with
hot water liberates free iodine from potassium iodide (distinction from cobalt).
Nickelie hydroxide when treated with dilute sulphuric acid forms NiSO^ ,
oxj'gen being evolved. With nitric acid the action is similar, distinction from
cobaltic hydroxide, which requires a more concentrated acid to effect a similar
reduction, e. — Salts. — The salts of nickel have a delicate green color in crystals
and in solution; when anhydrous, they are yellow. The nitrate and chloride
are deliquescent or efflorescent, according to the hygrometric state of the
atmosphere; the acetate is efflorescent. The chloride vaporizes at high tem-
peratures.
The carbonate, sulphide, phosphate, borate, oxalate, cyanide, ferrocyanide
and ferricyanide are insoluble; the double cyanides of nickel and alkali
metals, soluble in water. The chloride is soluble in alcohol, and the nitrate in
dilute alcohol. Most salts of nickel form soluble compounds by action of
ammonium hydroxide.
6. Reactions, a. — Alkali hydroxides precipitate solutions of nickel
salts as nickel hydroxide, Ni(0H)2 , pale green, not oxidized by exposure to
the air (§132, 6a), insoluble in excess of the fixed alkalis (distinction from
zinc), soluble in ammonium hydroxide or ammonium salts, formins: a
greenish-blue to violet-blue solution. Excess of fixed alkali hydroxide
•Nickel carbonyl Is prepared by heating the nickel ore in a current of CO. It is a Uquid, sp.
gT^ 13U6^ boiling at 43* and freezing at —26**. When hented to 200* it is decomposed into Nl and
CO (Berthelot, C. r., 1891, lia, 1843; 113, 679; Mond, J. Soc, Ltd., 1892, 11, 750).
170 NICKEL. §188, 66.
will slowly precipitate nickel hydroxide from the ammoniacal solutions
(distinction from cobalt). Alkali carbonates precipitate green hasi^
nickelotis carbonate, Ni5(0H)8(C08)2 (composition not constant), soluble in
ammonium hydroxide or ammonium salts, with blue or greenish-blue color.
Carbonates of Ba, Sr, Ca, and Hg are without action on nickelous
chloride or nitrate in the cold (distinction from Fe'", Al , and Cr"'), but
on boiling precipitate the whole of the nickel.
ft. — Oxalic acid and oxalates precipitate, very slowly but almost completely,
after twenty-four hours, nickel oxalate, green. Alkali cyanides, as KCN , pre-
cipitate tiirkel cyanide, NiCCN), , yellowish-green, insoluble in hydrocyanic
acid, and in cold dilute hydrochloric acid; dissolving in excess of the cyanide,
by formation of soluble double cyanides, as potassium nickel cyanide
(KCN)2Ni(CN)2 . The equation of the change corresponds exactly to that for
cobalt (§132, 66); and the solution of double cyanide is reprecipitated as
Ni(CN), by a careful addition of acids (like cobalt); but hot digestion, with
the liberated hydrocyanic acid, forms no compound corresponding to cobalti-
cyanides, and does not prevent precipitation by acids (distinction from cobalt).
It will be observed that excess of hydrochloric or sulphuric acid will dissolve
the precipitate of Ni(CN)3 . Eerrocyanides, as K4Fe(CN)0 , precipitate a
greenish-white nickel ferrocyanide, Ni3Ee(CN)0 , insoluble in acids, soluble in
ammonium hydroxide, decomposed by fixed alkalis. Eerricyanides precipitate
greenish-yellow nickel ferricynnide, insoluble in acids, soluble in ammonium
hydroxide to a green solution (§132, 6b). A solution of nitroferricyanide
precipitates solutions of cobalt and nickel salts, the latter being soluble in
dilute ammonium hydroxide (CavalU, Oazzetta, 1897, 27, ii, 95).
A solution of potassium xanthate precipitates neutral solutions of nickel and
cobalt, the former being soluble in ammonium hydroxide (distinction), from
which solution it is precipitated by (NHJjS (Phipson, C. .V.. 1877, 36, 150).
The xanthate also precipitates nickel in alkaline solution in presence of
Na4P207 (a separation from Ee'") (Campbell and Andrews, J, Am, Soc., 1895.
17, 125).
Xickel saltfi are not precipitated by an acetic acid solution of nitroso-,?-
naphthol (separation from cobalt) (Knorre, fi., 1885, 18, 702).
c. — Potassium nitrite in presence of acetic acid does not oxidize nickelous
compounds (distinction from cobalt), d, — Sodium phosphate, Na3HP04 , pre-
cipitates nickel phosphate, 'Nii{'PO^)i , greenish-white.
e, — Hydrosulphnric acid precipitates from neutral solutions of nickel
salts a portion of the nickel as yiickel sulphide, black (Baubigny, C. r., 1882,
94, 1183; 95, 34). The precipitation takes place slowly, and from nickel-
ous acetate is complete. In the presence of mineral acids no precipita-
tion takes place. Alkali sulphides precipitate the whole of the nickel,
as the black sulphide. Although precipitation is prevented by free acids,
the precipitate, once formed, is nearly insoluble in acetic or in dilute
hydrochloric acids; slowly dissolved by concentrated hydrochloric acid,
readily by nitric or nitro-hydrochloric.
Nickel mlphide, NiS, is partially soluble in yellow ammonium sulphide,*
from which brown-colored solution it is precipitated (gray^ black mixed with
• Hare (J. Am. Soc., 1895, 17, 537) adds tartaric acid to the solutloiis of nickel and cobalt, and an
excess of sodium hydroxide. He then passes in H^S. The cobalt is completely precipitated
while the nickel remains in solution, and can be precipitated upon acidulating^ the filtrate.
§138, 10. NICKEL, 171
sulphur) on addition of acetic acid (distinction from cobalt). Freshly pre-
cipitated nickel sulphide is soluble in KCN and reprecipitated as NiCCN), on
adding- HCl or HaS04 (separation from cobalt) (Guyard, BL, 1876, (2), 25, 509).
When nickel salts are boiled with a solution of NajSjOs , a portion of the nickel
is precipitated as the black sulphide.
/. — The halogen acids reduce the higher oxides of nickel to nickelouft
salts with liberation of the corresponding halogen. Potassium iodide
added to freshly precipitated nickelic hydroxide gives free iodine (distinc-
tion from cobalt).
g. — Nickel salts are precipitated by arsenltes and arsenates^ white or green-
ish-white, soluble in acids, including arsenic acid. h. — Potassium chromate
precipitates basic nickel chromate, yellow, soluble in acids, including chromic
acid (Schmidt, A., 1870, 156, 19). XjCrsOr forms no precipitate.
7. Ignition.— Nickel compounds dissolve clear in the borax bead, giving with
the oxidizing ilame a purple-red or violet color while hot, becoming yellowish-
brown when cold: with the reducing flame, fading to a turbid gray, from
reduced metallic nickel, and Anally becoming colorless. The addition of any
potassium salt, as potassium nitrate, causes the borax bead to take a dark
purple or blue color, clearest in the oxidizing flame. With micTOCOsmlc salt,
nickel gives a reddish-brown bead, cooling to a pale reddish-yellow, the colors
being' alike in both flames. Hence, with this reagent, in the reducing flame,
the color of nickel may be recognized in presence of iron rfhd manganese, which
are colorless in the reducing flame; but cobalt effectually obscures the bead
test for nickel. The yellow-red of copper in the reducing flame, persisting in
beads of microcosmic salt, also masks the bead test for nickel. By Ignition
with sodium carbonate on charcoal, compounds of nickel are reduced to the
metal, slightly attracted by the magnet,
8. Detection. — We proceed exactly as with cobalt for the nitroso-/?-
naphthol precipitation. The Ni remains in the filtrate and can be precipi-
tated with HjS (after neutralizing with NH^OH), and its presence con-
firmed by the usual tests. Or dissolve the sulphides of Ni and Co in
HHO3 , evaporate nearly to dryness, add an excess of KOH or NaoCOs ,
boil, add bromine water and boil to complete oxidation of the Co and Ni ,
filter, wash thoroughly with hot water and add hot solution of KI to the
precipitate on the filter paper. Free iodine (test with CSo) is evidence of
the presence of nickel.
9. Eatimation.^**— (/) Nickel hydroxide, oxide, carbonate or nitrate is ignited
at a white heat and weighed as NiO . (i) It is converted into the sulphate and
deposited on platinum as the free metal by the electric current. (3) Volu-
metrically. By titration in a slightly alkaline solution with KCN , using a
small amount of freshly precipitated Agl as an indicator (Campbell and
Andrews, J, Am, Soc„ 1895, 17, 127).
10. Oxidation. — ^Ni" is changed to Ni'" in presence of fixed alkalis by
CI , NaClO , Br , and NaBrO (not by I , distinction from cobalt, Donath,
B.y 1879, 12, 1868). Ni'" is reduced to Ni" by all non-reducing acids with
evolution of oxygen; by reducing acids, HjCgO^ is oxidized to COg, HNO,
* Gonial (Z. angew,, 1808, 177) grlves a summary of the methods proposed for the volumetrio*.
estimation of nickel.
172 MANGANESE, §134, 1.
to HHO3 , HgPOa to HgPO^ , H2S to S , H2SO3 to HaSO^ , HCl to CI , HBr to
Br , HI to I , HCNS to HCN and H^SO, , H,Fc(CN)e to HaFeCCH), . Hi"
is reduced to the metal by finely divided Zn ^ Cd ^ and Sn .
§134. Manganese. Hn = 55.0 . Valence two^ three^ four, six and
seven.
1. Properties.— Spcci/kf gravity, 7.138 to 7.206 (Brunner, Pogg,, 1857, 101, 264);
imlting point, at a high white heat (blue heat) (DeviUe, A. Ch., 1856, (3), 46,
199); volatilizes at the highest heat of the blast furnace (Jordan, C r., 1878,
86, 1874). It is a brittle metal, having the general appearance of cast iron,
non-magnetic, takes a high polish. According to Deville it has a reddish
appearance. It is readily oxidized, decomposing water at but little above the
ordinarj"^ temperature (Deville, /. c). It is used largely as ferromanganese in
the manufacture of Bessemer steel.
Oxides and hydroxides of manganese exist as dyad, triad and tetrad; the
salts exist most commonly as the dyad with some unstable triad and tetrad
salts; as an acid it is a hexad in mangunates and a heptad in permanganates.
2. Occurrence. — Not found native. It accompanies nearly all iron ores. Its
chief ore is pyrolusite, MnO, . It is also found as braunite, lCn,0,; hausman-
nite, Mn.O^; manganite, MnO(OH); manganese spar, MnCO,; manganese
blende, Ifl^S; and as a constituent of many other minerals.
3. Preparation. — (/) By electrolysis of the chloride. (2) By reduction with
metallic sodium or magnesium (Glatzel, B., 1889, 22, 2857). (S) By reduction
with some form of carbon. It has not been rexluced by hydrogen. (4) By
ignition with aluminum (Goldschmidt, A., 1898, 301, 19).
4. Oxides and Hydroxides. — (a) Manganous oxide, MnO , represents the only
base capable of forming stable manganese salts. It is formed (i) by simple
ignition of Mn(OH)j , MnCO, or MnC^O^ , air being excluded; (2) by ignition
of any of the higher oxides of manganese with hydrogen in a closed tube
(Moissan, A, Ch., 1880, (5), 21, 199). If prepared at as low a temperature as
practicable, it is a dark gray or greenish-gray powder, and oxidizes quickly
in the air to MngO^ . If prepaied at a higher heat it is more stable. Man-
ganous hydroxide, Mn(0H)3 , is formed from manganous salts by precipita-
tion with alkalis.' It quickly oxidizes in the air. forming MiiO(OH), thus
changing from white to brown. (^) Manganic oxide, MiiaO, , is formed by
heating any of the oxides or hydroxides to a red heat in oxygen gas or in air
(Schnieder, Pogg., 1859, 107, 605). Manganic oxide-hydroxide, MnO (OH) , is
formed (1) by oxidation of Mn(0H)3 in the air; (2) by treating MnO, with
concentrated H.SO4 at a temperature of about 130°, forming Mn2(S04)a and
then adding water: Mn,(S04), + 4HaO = 2Mn0(0H) + 3H,S04 (Carius. A.,
1856, 98, 63). (c) Trimanganese tetroxide, Mn.O^ , is formed when any of the
higher or lower oxides of manganese or any manganese salts with a volatile
acid are heated in the air to a white heat (Wright and Luff, B., 1878, 11, 2145).
The corresponding hydroxide would be Mn3(OH)8; this has not been isolated.
A corresponding oxide-hvdroxide is formed by adding freshly formed and
moist MnO, to an excess *of MnCl, containing ITH^Cl (Otto, .4., 1855, 93, 372).
(d) Manganese peroxide, MnO, , is formed (/) bv heating Mn(NO,). to 200**
(Gorgeu, C. r., 1879, 88, 796); (2) by heating MnCO, with KCIO, to 300°; (3) by
boiling any manganous salt with concentrated HNO, and KClOs. A correspond-
ing hydroxide, MnCOH)^ , has not been isolated. Several other hydroxides.
e.g., MnO(OH)3, Mn.Os^OH),, Mn,04(OH)4 etc., have been produced. The
chief use of manganese dioxide is in the preparation of chlorine or bromine.
(e) Manganates. — Manganic acid, HjMnOt , is not known in a free state. The
corresponding salt, K^MnO^ , is formed when any form of manganese is fused
with KOH or K,CO, (/) in the air, oxygen being absorbed: or (2) with KNO,
or KCIO. , NO or KCl being formed. A manganate of the alkali metals is
soluble in water, icith gradual decom position into manganese dioxide and per-
manganates: 3K,Mn04 -f 2H,0 = 2KMn04 + MnO, -f 4K0H . Free alkali
l,5e.
^ANQANESB,
ITS
^^13, and free acids and boilingr promote, this change. Manganates Imve
een Cf/lnr. whieh turns to the red of permanganates dnrin^ the derumpc^si-
iiwu inevitiibie in solution. This is* the nsiial method of manufactuniig KMnO^.
(f) Permanganic acid is not in use as an acid, but is repreaeuteU b.v tht- pcr-
manga nates, as KMnO, . The permanganic acid radical is at once deconiptiHcd
by addition of tiot H^SO| to a ^joJid permuni^anate (/), but in water solution
this decomposition does not at once take place, except by contact with oxidiz-
able subetanceti. The ox;idrdn|T powor of permanganates extends to a great
number of substances, possesses different charactfri^,tics in acid and in «lka-
line solutions* iind acts in many rawes bo rapidly as* to be violently explosive.
The react jons with ferrous sallu (2) and with oxalic acrd (J) are much used in
volumetric analysis.
(/) 4KMnO, + 2H,S0, = 2K,B0, ^ 4MbO, + m, + 2Mfi
and 2MnO, + 2n,B0^ = SMnSO* + 3H,0 + 0,
R4KMnO, + GH.SO, = iMnSO. + 2K,S0, + 50, + GH,0
(2) KMnO, + :iFeCl: + S^Cl = MnCl, -j- KCl + 5FeCl, + 4H,0
(3) 2KMiiO, + 5H,C,0, + GHCl = 2MnCl, + 2KC1 + SH,0 + lOCO,
. S^lubUities* — (L — jl/effJ^.^ManR^anese dissolves readily in dilute acids to
form manga nous -salts. Concentrated H^SO, di^solvi'sj it only ou warming, SO,
being evolved. It combines readily with chlorine and bromine* h. — fjxkh'if
and hjfdrfixidrM, — All oxides and hydroxides of manganese an* insoluble in
water. Thry are soluhk>, upon warming, in hydrochloric acid, forming man-
gmnons chloride; the higher oxides and hydroxides being reduced with evolU'
tion of chlorine (commercial method of preparation of chlorine). Instead of
hydrochloric acid, sulphuric ac'd and a chloride may be employed (HBr and
HI act similarly to, and more readily than HCl). In the cold, hydrochloric
acid djKsolves JfnO;^ to a greenish -brown solution, containingj prohHl>ly. BtnCl,
or MnCli , unstable* giving chlorine when warmed and forming MnO. when
strongly diluted with water (Pickering, J. t\, 1870, 35, 654; Nicklcs, A. Vh.^
1865, (4), 5, 161). Xitrie and sulphurie acids dissolve manganous oxide and
hydroxide to manganous saUs ^langanese dioxide (or hyd rated oxide) is
insoluble in nitric acid, dilute or concentrated; concentrated sulphuric acid
with heat deromposes it, evolving oxygen and forming manganous sulphate;
2lfnO, + *2H,S0, =r 2MnS0» + 2H3O -f O, . Manganons hydroxide is insoluble
in the alkalis but mhihle '^n solutions of nrnmonium 90IU.
^mc, — Sails, — ^Manganous sulphide, carbonate, phosphate, oxalate, borate,
and sulphite are insoluble in water, readily Boluble in diluto acids, Man*
^nic lalts are somewhat unstable compounds, of a reddish-bro\^Ti or
piiq»le-red cqIqt^ becomiog paler and of lighter tint in reduction to the
nianganons combination. ICnOl^ and MnSO^ are deliquescenL Man-
ganic cMoride., MnCl^ , cxista only in solution, which is reduced to
ICitCl^ by boiling, also by evaporation to a solid. Mongnnic sulphate^
Mn.(SOJ;, 3 is sohible in dilute sniphnric acid, but is reduced to ICnSO^ by
^be attempt to dissolve it in water alone; potasBinm manganic sulphate
Hbd other mangamr alums arc al^o decomposed by water, xilkali mnnfian-
ateg and permantianatm are aolnble in watepj the former rapidly changing
to manganese dioxide and permanganate, which is much more stable in
iolotioH. In presence of reducing agents both manganates and perman-
ganateB are reduced to lower forms,
KjMnO, + 8HC1 = MnCl, + 2KC1 + 2Cl, + 4H,0
SKMnO, + 3MnS0, + 2H,0 = 3MnO, + K,BO, + SH.SO,
174 MAXOSNESE. §134, 6a.
Concentrated HjSO^ in the cold digsolves EMnO^ , forming (111103)2804
(a sulphate of the heptad manganese : 2Elin04 + SHgSO^ = (111103)2804 +
2EH8O4 + 2H2O (Franke, J. pr,, 1887, 86, 31). If heat be applied oxygen
is evolved and the manganese is reduced to the dyad (4/).
6. Reactions, a. — The fixed alkali hydroxides precipitate from solu-
tions of manganous salts, manganous hydroxide, Mn(0H)2, white, soon
turning brown in the air by oxidation to manganic hydroxide, MnO(OH) .
The precipitate is insoluble in excess of the alkalis; but, before oxidation,
is soluble in excess of ammonium salts with formation of a double am-
monium manganese compound * {!). Ammoninm hydroxide precipitates
one half of the manganese as the hydroxide from solutions of manganous
salts, the other half being held in solution as a double salt by the am-
monium salt formed (2) {Dammer, 3, 237). The presence of excess of
ammonium salt prevents the precipitation of the manganese by ammonium
hydroxide (S) (separation of manganese from the metals of the third
group) (Pickering, J. C, 1879, 35, G72; Langbein, Z., 1887, 26, 731).
Mtinganic hydroxide, MnO(OH), is insoluble in the alkalis or in ammonium
salts. It gradually precipitates, completely on exposure to the air, as
a dark brown precipitate from solutions of manganous hydroxide in am-
monium salts. Alkali carbonates precipitate manganotis carbonate, MnCO., ,
white, oxidized in the air to the brown manganic hydroxide, and before
oxidation, somewhat soluble in ammonium chloride. Strong ammonium
hydroxide gradually reduces a solution of potassium permanganate to
manganese dioxide (10&).
(/) Mxi(OH), + 4NH,C1 = MiiCl,.2NH,Cl -j- 2NH,0H
(2) 2MxiS0« + 2NH,0H = MnSO^.CNHJ^SO^ -f Mn(OH),
(3) MnCl, -j- 2NH4CI = MnCl,.2NH,Cl or (NH4),MnCl,
b, — Oxalic acid and alkaline oxalates precipitate Tnanganoxis oxalate.
soluble in mineral acids not too dilute. All compounds of manganese of
a higher degree of oxidation are reduced to the manganous condition on
warming with oxalic acid, or oxalates in presence of some mineral acid:
2KMnO, + 5H2C2O, -f 3H28O, = K28O, + 2Mn80, + IOCO2 -f 8H2O .
Soluble cyanides, as KCN , precipitate manganous cyanide, Mn(CN)2 , white,
but darkening in the air; soluble in excess of the precipitant by formation of
double cyanides, as Mii(CN)2.2KCN . This solution, exposed to the air, pro-
duces manganicyanides (analogous to ferricyanides), with oxidation of the
• It has been questioned whether the solubility of llIn(OH)s in ammonium salts is due to com-
bination between the two. As has been already stated, the Law of Mass- Action causes that
reaction to take place which leads to the formation of a slightly dissociated substance. Thus
Pe(OH)3 dissolves in HCl and Aa^Os in KnOH because in each case water, a non-dissociated
substance, results; and FeS and A ■28.1 dissolve In HCI and NaHS respectively because the
Uttle-dissociated H^S is a product. Similarly, BTHfCl with Mn(OH)s gives opportunity for the
formation of NH4OH, a compound of small dissociation-constant. Solution due to this caufe
can take place only with hydroxides ha%'ing a comparatively large solubility-product '§45).
See Ostwald on the solubility of AIip(OH)s, ♦• WissenschaftUche Grundlagen der analytischen
Chemie," 2d ed., p. 183.
§134, 6c. MAXQANESE. 17?
manganese: 12(Mn(CN)a.2KCN) -f 30, + 2H,0 = 8K,Mn(CN). -f 4MnO(OH),
Pe'" and Mn" may be separated by treating a solution of the two metals with
a strong excess of KCN and then with iodine. The manganese is precipitated
as MnOa and the iron remains in solution (Beilstein and Jawein, B., 1879, 12,
1528). Ferrocyanides piecipitate white manyanous ferrocyanide, MnaFe{CN)e ,
soluble in hydrochloric acid. Eerrlcyanides precipitate brown mamjanous ferri-
cyanide, Mn8(Fe(CN)e)a , insoluble in acids (separation, with Co and Ni , from
Zn) (Tarugi, Oazzetta,, 1895, 25, ii, 478). If an alkali or alkali carbonate be
present, potassium ferricyanide oxidizes manganous compounds to manganese
dioxide, the ferricyanide being reduced to ferrocyanide. Potassium ferro^
cyanide reduces manganates and permanganates to manganous compounds.
c. — ^Nitric acid is of value in analysis of manganese compounds in that
it, as a non-reducing acid, acts readily with oxidizing agents, as PbOa,
KCIO3 , etc., to oxidize manganous compounds to manganese dioxide or to
permanganic acid. Reducing agents as HCl, etc., should be absent*
Sulphuric acid may be used instead of nitric acid.
2Mn(NO,)3 + 5PbO, -f 6HN0, = 2HIIX1O4 + 5Pb(N0a)j + 2H,0
5MnS0, 4- 2KC10, -f H,SO, -f 4H,0 = 5MnO, + K,SO, + CI, + 5H,S04
In using PbOg and HNO3 to detect manganese, the compound should first
be reduced with hydrochloric acid, precipitated with potassium hydroxide
and this precipitate dissolved in nitric acid, as HnOj is not all oxidized
by PbOj and HNO3 (Koninck, Z. angew., 1889, 4).
d, — Hypophosphorous acid reduces all higher forms of manganese to the
manganous condition. Alkali phosphates, as NajHPO^ , precipitate, from
neutral solutions of manganous salts, normal m<tnganous phosphate^ Mn,(P04)3 ,
white, slightly soluble in water, and soluble in dilute acids. It turns brown in
the air. The manganous hydrogen phosphate, MnHFG^ , is more soluble in
water, and is obtained by crystallization from a mixture of manganous sul-
phate acidulated with acetic acid and disodiura phosphate, Na^HPOi , added
till a precipitate begins to form. From the ammonium-manganese solution,
freshly formed (6«), phosphates precipitate all the manganese as manganous
ammonium phosphate.
e. — Hydrosnlphuric acid precipitates manganous acetate but imperfectly,
and not in presence of acetic acid, and does not precipitate other salts, as
manganous sulphide is soluble in very dilute acids, even acetic acid.
Ammoninm sulphide precipitates from neutral solutions, and forms from
the recent hydroxide of mixtures made alkaline, the flesh -colored w^n-
ganous sulphide^ MnS . Acetic acid, acting on the precipitated sulphide?,
separates manganese from cobalt and nickel, and from the greater part of
zinc. All the higher oxidized forms of manganese (in solution or freshlv
precipitated) are reduced to the manganous condition, with separation of
sulphur (10), by hydrosulphuric acid or soluble sulphides: 4Elin04 +
14(NH,)oS + VSRJd = 4MnS + 4K0H + 28NH^0H -f 5So . The green
manganous svlphide, MnS , crystalline, anhydrous, is formed by the action
of HoS on a hot ammoniacal manganous solution not containing an exce.'^s
of ammonium salts (Meineke, Z, angew,, 1888, 3).
Soluble sulphites precipitate from solutions of manganous salts, manganous
sulphite, MxiSO» , white, insoluble in water, soluble in acids (Gorgeu, C. r.,
l*-8.'^, 96, 341). Solutions of manganates or permanganates are tawsi^d^ftl^V^
176 AiAyoANESt:. §134,6/.
reduced to the flocculent brown-black manganese dioxide by solutions of
sodium sulphite or sodium thiosulphate; if acids be present, the reduction is
complete to manganous salts.
/. — HCl, HBr, and HI readily reduce the higher compounds of man-
ganese to manganous salts with evolution of the corresponding halogen.
When manganese dioxide is dissolved in concentrated HCl without heat,
the dark brownish colored solution is said to consist of manganese tetra-
chloride, HnCl^, which deposits MnO^ on dilution with water and on
warming decomposes into manganous chloride and chlorine (56) (Picker-
ing, J, C, 1879, 35, 654). Potassium iodide instantly reduces a solution
of potassium permanganate, forming manganese dioxide and an iodate
{distinction from chloride and bromide). Potassium chlorate or bromate
when boiled with concentrated nitric or sulphuric acids and manganous
•compounds forms manganese dioxide (c),
g. — Soluble arsenites precipitate manganous arsenitc, and arsenates precipitate
manganous arsenate^ insoluble in water, soluble in acids. Arsenous acid and
arsenites reduce solutions of mangauates or permanganates, forming a brown
flocculent precipitate; or a colorless solution if warmed in presence of a
mineral acid, h, — Normal potassium chromate precipitates manganous salts,
brown, soluble in acids and in ammonium hydroxide; no precipitate is formed
with potassium dichromate. ♦. — Soluble manganates and permanganates pre-
cipitate manganous salts as manganese dioxide, being themselves reduced to
the same form; SMnSO* -j- 2KMn04 + 2H3O = SMnO, -f K2SO4 + 2H,S04 .
7. Ignition with alkali and oxidizing agents, forming a bright green mass
of alkaline manganate, constitutes a delicate and convenient test for man-
ganese, in any combination. A small portion of precipitate or fine powder
is taken. If the manganese forms but a small part of a mixture to bo
tested, it is better to submit the substance to the systematic course of
analysis, and apply this test to the precipitate by alkali, in the fourth
group. A convenient form of the tost is by ignition on platinum foil with
potassium or sodium nitrate and sodium carbonate (a). Ignition, by an
oxidizing flame, on platinum foil, with potassium hydroxide, effects the
same result, less quickly and perfectly (&). Ignition by the oxidizing flame
of the blow-pipe, in a bead of sodium carbonate, on the loop of platinum
wire, also gives the green color (c),
(a) 3Mn(OH)2 -f 4KN0, + Na,CO« =
2K,MnO« + Na,Mn04 -f 4N0 + CO, + 3H,0
(ft) Mn(OH), -f 2K0H 4-0, = K^MnO^ -f 2H.0
(c) Mn(OH), -f Na.CO, -f O, = Na^nO^ -f H^O + CO,
With beads of borax and microcosmic salt, before the outer blow-pipe flame,
manganese colors the bead violet while hot, and amethyst-red when cold. The
color is due to the formation of manganic oxide^ the coloring material of the
amethyst and other minerals, and is slowly destroyed by application of the
inner flame, which reduces the mangranic to manganous oxide.
8. Detection. — After the removal of the metals of the first three groups
^the third group in the presence of NH^Cl in excess, 56 and 60), the Mn
§134, 10&. MANGANESE. 177
with Co , Hi and Zn is precipitated in the ammoniacal solution by HoS .
By digestion in cold dilute HCl the sulphides of Mn and Zn are dissolved,
and after boiling to remove the HgS , Un is precipitated as the hydroxide
by excess of KOH , which dissolves the Zn . The precipitate of the man-
ganese is dissolved in HNO3 and boiled with more HNO3 and an excess of
PbOj . A violet-colored solution is evidence of the presence of manganese.
9. EstimatioxL — (i) By converting into l[n,04 (4r), and weighing as such,
(i) By precipitating as MX1NH4PO4 , and after ignition weighing as M1I3P2O7 .
(3) By treating the neutral manganous salt with a solution of KM11O4 of
known strength (6J). If some ZnSO^ is added the action is more satisfactory
(Wright and Menke, J, C, 1880, 37, 42). (4) By boiling the manganous com-
pound with PbOj and HNO, , and comparing the color with a permanganate
solution of known strength (Peters, C, N,, 1876, 33, 35). (5) The manganous
compound is oxidized to MnO, by boiling with KClOs and HNO, . This is
then reduced by an excess of standard H2O2 , HjCjO^ or EeSO^ , and the excess
of the reagent estimated by the usual methods. (6) MnO, , obtained as in (t5),
is treated with HxCjOf and the evolved COj measured or weighed. (7) M11O2 ,
obtained as in (o), is boiled with HCl and the evolved CI estimated.
10. Oxidation.— (a) Mn" is oxidized to Mn'" in alkaline mixture on
exposure to the air; to Mn^^ in neutral solution by KjMnO^ and KMn04 ,
in alkaline mLxture by CI , Br , I , K3Fe(CN)e , KCIO , KBrO , HoOoS etc. ;
in acid solution by boiling with concentrated HNO3 or H0SO4 , and KCIO,
or EBrOs . Mn^""° is oxidized to Mn^' by fusion with an alkali and an
oxidizing agent, or by fusion with KCIO3 alone (Boettger, Z., 1872, 11,
433). Mn^-«» is oxidized to'Mn^" by warming with PbOo or PbgO^ and
HNOs or H2SO4 . The higher oxide of lead should be in excess and reduc-
ing agents should be absent as they delay the reaction ; hence in analysis
the manganese, should be precipitated as the hydroxide or sulphide, fil-
tered, washed, and then dissolved in HNO3 or H2SO4 , and boiled with the
higher oxide of lead (6c). A solution of potassium manganate decomposes
into potassium permanganate and manganese dioxide on standing, more
rapidly on warming or dilution with water, (h) All compounds of man-
ganese having a higher degree of oxidation than the dyad, (Mn"+") are
reduced to the dyad (Mn") by H^C.O^ , HH^PO^ , H^SS K.S , H.SO, , H.O,^
<in neutral or alkaline solution to Mn'^), HCl , HBr , HI , HCNS , Hg', Sn",
As'", Sb"', Cn', Fc", Cr", Cr"', etc. ; the reducing agents becoming respec-
tively COj , P^, S° to S^ (depending upon the temperature, concentration,
and the agent used in excess), CI , Br , I , HCN and S^i, Hgr", Sn^^^ ^v^ g^v^
Cn'', Fc'", and Cr^. Mn^+^ is reduced to Mn^^ (or Mn'") by H \ AsH3^
«bH3», PH,», Ha^SOg*, NaoS^Og*, NH,OH« (slowly), Mn", etc. KMnO, is
reduced to EoMnO^ on boiling with concentrated KOH : ^EXnO^ -f 4K0H
= 4K2MnO^ + 2H2O + 0. (Rammelsberg, J5., 1875, 8, 232).
1 Klein, Arch, Fharm.^ 1889, 227, 77; Jannacscb and von Cloedt, Z. annrg., 1895, 10, 398 and 410;
Gamot, C. r., 1888, 107, 997 and 1150.
'QtrDot, Bl., 1889, (3), 1. 277 ; Oorgeu, C. r., 1890, 110, 958. > Jones, J. C, 1878, 33, 96. * Hoenig
and Zatzck, IT., 1883, 4. 738 ; Glaeecr, If., 1835, 6, 339.
178 zjyc, §136,1.
§136. Zinc. Zn = 65.4 . Valence two.
1. Properties.— iSfped/Tr gravity, 7.142 (Spring, B., 1883, 16, 2723). Melting
point, 418.5*' to 419.35** (Heycock and Neville, J. C, 1895, 67, 185). Boiling pointy
940° (Violette, C. r., 1882, 94, 720). It is a bluish-white metal, retaining its
lustre in dry air, but slightly tarnished in moist air or in water. When heated
to the boiling point with abundant excess of air it burns with a bluish-white
flame to zinc oxide. Zinc dust mixed with sulphur is ignited bj*" percussion
(Schwarz, B., 1882, 15, 2505). At ordinary temperature it breaks with a coarse
crystalline fracture. It is more malleable at 100° to 150° than at other tem-
peratures, and at that temperature may be drawn into wire or rolled into
sheets. At 205° it is so brittle that it may be easily powdered in a mortar.
Zinc finds an extended use in laboratories for the generation of hydrogen.
It is molded in sticks or granulated by pouring the molten metal into cold
water. The pure metal is not suitable for the generation of hydrogen, as the
reaction with acids proceeds too slowly (Weeren, B., 1891, 24, 1785). Com-
mercial impurities render the metal readily soluble in acids, or the pure metal
may be treated with a dilute solution of platinum chloride (twenty milligrams
PtCl4 per litre). Metallic platinum is deposited upon the zinc: PtCl* -f 2Zn =
Pt -f 2ZnCl, .
2. Occurrence. — It is found as calamine (ZnCO,), as zinc-blende (ZnS): also
associated with other metals in numerous ores.
3. Preparation. — The process usually employed consists of two operations:
(1) Roasting: in case of the carbonate the action is: ZnCO, = 25nO -f- CO-: if it
is a sulphide, 2ZnS -f 30, = 2ZnO -f 280, . (2) Reduction with distillation:
after mixing the ZnO with one-half its weight of powdered coal» it is distilled
at a white heat. Its usual impurities are As, Cd , Pb , Cu , Fe and Sn . It is
purified by repeated distillation, each time rejecting the first portion, which
contains the more volatile As and Cd , and the last which contains the less
volatile Pb , Cu , Fe and Sn . Strictly chemically pure zinc is best prepared
from the carbonate which has been purified by precipitation.
4. Oxide and Hydroxide. — Zinc oxide (ZnO) is made by igniting in the air
either metallic zinc, its hydroxide, carbonate, nitrate, oxalate, or any of its
organic oxysalts. Zinc hydroxide. Zn(0H)2 , is made from solutions of zinc
salts by precipitation with fixed alkalis (6a).
5. Solubilities. — (a) Metal, — Pure zinc dissolves very slowly in acids or alkalis,
unless in contact with copper, platinum or some less positive metal (Baker.
J, C, 1885, 47, 349). The metallic impurities in ordinary zinc enable it to
dissolve easily with acids or alkali hydroxides. In contact with iron, it is
quite rapidly oxidized in water containing air, but not dissolved by water
unless by aid of certain salts. It dissolves in dilute hydrochloric, sulphuric *
and acetic acids (/), and in the aqueous alkalis (2), with evolution o*f hj^drogen:
in very dilute nitric acid, without evolution of gas (S); in moderately dilute
cold nitric acid, mostly with evolution of nitrous oxide (i); and, in somewhat
less dilute nitric acid, chiefly with evolution of nitric oxide (5). Concentrated
nitric acid dissolves zinc but slightly, the nitrate being very sparingly soluble
in nitric acid (Montemartini, Ofizzetta, 1802, 22, 277). Hot concentrated sul-
phuric acid dissolves it with evolution of sulphur dioxide (6).
(1) Zn + H3SO, = ZnSO. + H,
(2) Zn 4- 2K0H = K^ZnO, + H,
(3) 4Zn + lOHNO, = 4Zn(N0,), + NH4NO, -f 3H,0
{J,) 4Zn 4- lOHNO. = 4Zn(N03), + N,0 + 5H,0
(J) 3Zn -j- 8HN0, = 3Zn(N0,), -f 2N0 -f 4H,0
(G) Zn + 2H2SO, = ZnSO^ -h SO, -f 2H,0
(h) (Iridr and Hydroxide,— All the agents which dissolve the metal, dissolve also
its oxide and hydroxide.
•Muir and Robbs, O. A'., 1882, 45, 69.
§136, 6e. , 7Ayc, 179
(c) Salts, — The chloride, bromide, iodide, chlorate, nitrate (6aq), and
acetate (7aq) are deliquescent; the sulphate (7aq) is efflorescent. The
chloride is readily soluble in alcohol in all proportions (Kremers, Pogg,^
1862, 115, 360). The sulphide, basic carbonate, phosphate, arsenate,
oxalate, and ferrocyanide are insoluble in water; the sulphite is sparingly
soluble. The ferrocyanide is insoluble in hydrochloric acid (Fahlberg, Z,,
1874, 13, 380). The sulphide is almost insoluble in dilute acetic acid (sepa-
ration from MnS). All zinc salts are soluble in EOH and NaOH except
zinc sulphide, and all in NH^OH except ZnS and Zn2Fc(CH)e .
6. Beactions. a. — The alkali hydroxides precipitate zinc hydroxide,
Zn(0H)2 9 white, soluble in excess of the precipitant forming an alkali
zincate:
ZnCl, -f 2K0H = Zn(OH)a + 2KC1
Zn(OH), + 2K0H = K^ZnO, + 2H3O
ZnCl, + 4NH4OH = (NH4),ZiiO,+ 2NH,C1 -f 2H,0 •
The precipitate of zinc hydroxide dissolves more readily in excess of the
alkalis at ordinary temperature than when heated. Unless a strong excess
of the alkali be present, boiling causes a precipitation of ?inc oxide, more
readily from the solution in ammonium hydroxide than in the fixed
alkalis. The presence of other metals — as iron or manganese— makes
necessary the use of much more alkali to effect solution. An alkali solu-
tion as dilute as tenth Normal does not dissolve zinc hydroxide, no matter
how great an excess be added (Prescott, J. Am. Soc, 1880, 2, 29).
Alkali carbonates precipitate the basic carbonate, Zn5(0H)e(C03)2 , white,
soluble in ammonium carbonate, readily in alkali hydroxides (Kraut, Z.
anorg., 1896, 13, 1). Carbonates of Ba , Sr , Ca , and Hg have no action
at ordinary temperatures (separation from Fe'", Al , and Cr'"), but upon
boiling precipitate the whole of the zinc.
ft. — Alkali cyanides, as KCN , precipitate zinc cyanide, Zxi(CN)a , white,
soluble in excess of the precipitant. Alkali ferrocyanides, as K4Fe(CN)« ,
precipitate zinc ferrocyanide, Zii2Fe(CN)e , white (5c). Alkali ferrlcyanides,
as K8Fe(CN)« , precipitate zinc ferricyanide, Zn8(Fe(CN)e)a , yellowish, c. —
See ,5r. <2.— Sodium phosphate, Na2HP04 , precipitates zinc phosphate^ soluble
in alkali hydroxides and in nearly all acids.
e. — Hydrosnlphuric acid precipitates a part of the zinc from neutral
polutions of its salts with mineral acids, and the whole from the acetate;
also from other salts of zinc, by addition of alkali acetates or monochlor-
acetic acid, in small excess (separation from Mn , Co , Ni , and Fe) (Berg,
* Ostwald incUnes to the view that the solubility in NH4OH Is duo to the formation of a
•omplex ammonium-zino Ion (Scicntiflo Foundations, p. 151 ; see also second German edition, p.
147). The fact that NH4CI precipitates Zn(OH), from its solution in fixed alkali, and on further
tddition redlssolves it and also that NH4CI hinders precipitation by heat from the ammoniacal
H).utlon of the hydroxide speaks against the assumption that solution in the latter case arises
from the formation of a zincate.
180 Z/iST. §135, 6f.
Z., 1886, 25, 512): ZnCl^ + 2KC^ILfi^ + H^S = ZnS + 2KC1 +
2HC2H3O2 .* That is : Zinc sulphide is not entirely soluble in dilute acids,
though much more soluble in mineral acids than in acetic acid. The
precipitate is white when pure. Alkali sulphides completely precipitate
zinc as sulphide, both from its salts with acids and from its soluble com-
binations with alkalis.
Concentrated solutions of sodium sulphite precipitate solutions of zinc salts
as basic zinc sulphite; or if the solutions be too dilute for immediate precipita-
tion, boiling will cause the immediate formation of the bulky white precipittte
of the basic sulphite (Seubert, Arch. Pharm., 1891, 229, 316). ^— If a hot con-
centrated zinc chloride solution be treated with ammonium hydroxide until
a precipitate begins to form, a basic chloride, 2ZnCl2.9ZnO , will separate out
upon cooling as a white precipitate (Habermann, M., 1884, 5, 432).
y. — Zinc salts are precipitated by solutions of alkali arsenites and arsenates,
forming respectively zinc arsenite or arsenate, white, gelatinous, readily solu-
ble in alkalis and acids, including arsenic acids, h. — Normal potassium cliro-
mate forms, with solutions of zinc salts, a yellow precipitate readily soluble
in alkalis and acids, including chromic acid. No precipitate is formed with
7. Ig^tion. — With sodium carbonate, on charcoal, before the blow-pipe, com-
pounds of zinc are reduced to the metallic state. The metal is vaporized, and
then oxidized in the air, and deposited as a non-volatile coating, yellow when
hot and white when cold. If this coating, or zinc oxide otherwise prepared,
be moistened with solution of cobalt nitrate and again ignited, it assumes a
green color (Bloxam, ./. C, 1865, 18, 98). With borax or microcosmic salt, zinc
compounds give a bead which, if strongly saturated, is yellowish when hot,
and opaque white when cold.
8. Detection. — After the removal of the first three groups, the Zn i^
precipitated with Co , Ni and Mn from the ammoniacal solutions by H-S .
Digestion of the precipitated sulphides with cold dilute HCl dissolves the
Mn and Zn as chlorides. The solution is thoroughly boiled to expel the
HgS and the zinc changed to "SeL^ZnO^ by an excess of NaOH , which precipi-
tates the manganese as the hydroxide. From the alkaline filtrate HgS gives
a white or grayish-white precipitate — evidence of the presence of Zn .
9. Estimation. — (/) Zinc is weighed as an oxide, into which form it is
brought by simple ignition if combined with a volatile inorganic oxyacid,
otherwise it should be changed to a carbonate and then ignited. (2) It is
converted into a sulphide, and after adding powdered sulphur it is ignited in
a stream of hydrogen or hydrogen sulphide, and weighed as a sulphide (Kiinzel,
Z., 1863, 2, 373). (3) It may be converted into Z11NH4FO4 , and, after drying:
at 100°, weighed. Ignition converts it into Zn^PjO, , with slight loss of zinc.
(4) Volumetrically, by converting into ZiiiFe(CN)e and titrating with potas-
sium permanganate or by using EeCls acidulated with HCsHsO, as external
• In the equation for acetic acid, ab = kc, a and b, the concentrations of the H and C,IIsOa
Ions respectively, are small, c is large, and k, the so-called " dissociation-constant," to which
the strengrth of the acid is proportional, is very small. But addition of the fuUy-dissociated
sodium acetate to the Likewise completely-Ionized hydrochloric acid gives a solution containing
the Ions In very large concentration and practically none of the non-dissociated acetic acid.
To restore equilibrium the II ions of the HCl unite with the aoetio ions of the sodium acetate,
leaving Na and CI ions in the solution. The displacement of a weak add from its salt by a
strong one lies then not so much in an attraction of the strong acid by the base aa in the ten-
dency of the weak acid to form the non-ionized molecule.
§136,10. ZlUfC. 181
indicator (Voigt, Z. angew, 1889, 307). (J) By precipitation as Zn,(Fe(CN)a)j »
treating the precipitate with potassium iodide and titrating the liberated iodine
(Mohr, Dingl., 1858, 48, 115). ((>) By titration in hydrochloric acid solution
with KfFeCCN), , using a uranium salt as an indicator (Fahlberg, Z., 1874, 13,
379; KoniDck and Prost, Z. angew., 1896, 568). (7) By titration in alkaline
solution with NOjS , using a copper salt as an indicator. (8) The zinc is pre-
cipitated as ZnNH4As04 , the precipitate decomposed with HI and the liber-
ated iodine titrated with standard Na,S,0, (Meade, J. Am, 8oc., 1900, 22, 353).
10. Oxidation. — Metallic zinc precipitates the free metal from solutionsf
of Cd , Sn , Pb , Cn , Bi , Hg , Ag , Pt , An , As , Sb , Te , In , Fe S Co ,
ITi, Pd, Eh, It, and Os (Gmelin-Kraut, HandbucJi, 1875, 8, 6). Zinc
with copper (zinc-copper couple, used in water analysis) reduces nitrates
and nitrites to ammonia, chlorates to chlorides, iodates to iodides, fern-
cyanides to ferrocyanides, etc. (Thorpe, J. 0., 1873, 26, 541). Solutions
of chromates are reduced to chromic salts, ferric salts to ferrous salts,
and compounds of manganese having more than two bonds are reduced to
the dyad in presence of some non-reducing acid. Zinc is precipitated as
the metal from acetic solutions by Mg (Warren, C. 2V., 1895, 71, 92).
The oxide is reduced to the metal by heating in a current of hydrogen
(Deville, A. Ch^ 1855 (3), 43, 477).
1 Davles, J, C, 1875, 28, 311.
182
REACTIONS OF IROIf AND ZINC GROUP BASES.
§138,
§137,
TABLE FOR ANALYSIS OF THE ZINC GROUP.
183
§137. Table fob Analysis of the Zinc Gboup (Foubth Gboup)
(Phosphates and Oxalates being absent).
Into the clear ammoniacal filtrate from the Third Group pass HYDBOSUL-
PHURIC ACID GAS, and if a precipitate appears, warm until it subsides.
Filter and wash with a one per cent solution of NH4CI . Test filtrate, in
which H2S gives no precipitate for the Fifth Oroup.)
Precipitate: CoS , NiS , MnS , ZnS .
Treat on the filter with cold dilute Hydrochloric Acid.
Besidue: CoS, NiS'' (black).
For Cobalt:
Dissolve in nitro
hydrochloric
acid, evaporate
and add NaHCO,
and HgO,; warm
g-ently and filter.
A green color to
the filtrate indi-
cates cobalt
(gi40>.
Test the black resi-
due with the
borax bead (blue
color charncteris-
tic of cobalt,
5132, 7).
If fiuflHcient nickel
be present to ob-
scure the bine
bead (JS133, 7>,
disisolTe the aul-
pbides in nitro-
hydrochloric add.
evaporate and add
an excess of nl-
troso-/? -naphthol
in acetic acid so-
lution (5132, afe):
filter, wflfjh and
tesi the brick-red
precipitate, with
the borax bead.
5136,4138,5130,
S140,§141,S144,
5145 and ff.
ForKickelj
Bisfiolve the snl-
phidea in nitro-
hydrochloric
acid, evaporate
and add an CK-
ceaa of nitroso-/if-
naphtbol in acet-
ic solution to re-
move the cobalt
§132, 6f»). Filter
and add to til*
trate ammonium
hydroxide till al-
kaline, filter and
to the filtrate
add HjS. A black
precipitate, HIS^
indicates nickel.
Or: Dissolve the
CoS and ms,
add excess of
hot SOH and
Br, boil, filter,
wash (until fil-
trate g'lvea no
precipitatt* with
AgNO,), add so-
lution of hot KI
and test the fil-
trate with CSj.
If free iodine ap-
pears^ nickel ia
present(il33,6f).
Stud5^ the text at
|133, f>fl. ft, c and
f; R132, fift and e:
S136,§13S, §139,
gl40, fl4t.§144,
5145 and if.
SDlution: HnCl, , ZiiCl,(H3S,HCl),
Boil the sfthition thnrouffhlu to remove the
H2S , cool, and add a decided excess
of potassium or sodium hydroxide and
digest without warming (§135, Cct),
Filter and wash.
Precipitate T
M[n(OH),*
Dissolve in nitric
acid and boil
with an exeens of
PbO, and HNO«.
Violet Bolutioiv
(HMnOO indi-
cates mj^iuganese
(chsracteristic-
reaction, §134,
Dark-colored orig-
inal Bohttlons iu"
djcaling an alka-
li salt of mang-a-
neso should be
reduced by
warming with
HCl before pro-
ceeding- with the
analysis (§134,
fiC and 6/).
Confirm by study
of tbetext, S134,
7, 1136, §138,
§139, 1142. §143,
§144, gl45 and
Solution:
Teat for ztnc by
adding H,S. A
white precipitate
(ZnS) indicates
zfnc
Study the text at
§135, 6rt and e,'
B36, §138.n39,
§142, n 43. §144,
§145 and /f.
•Small portions of cobalt and nickel sulphides may be dissolved by the cold dilute HCl, and
will be proclpltated with the Bf n'OH). . These traces will not interfere with the further tests
for manganese.
184 DIRECTIONS FOR ANALYSIS WITH NOTES. §138.
Directions fob the Analysis of the Metals of the Fourth Group.
§138. Kanipnlation. — Into the warm strongly ammoniacal filtrate from
the third group (§128), HjS gas is passed until complete precipitation is
obtained :
MnCl2.2NH,Cl -f 2NH4OH + H,S = MnS + 4NH4CI + 2H,0
(NHO^ZnO, -f 2H,S = ZnS + (NH,),S + 2H,0
The solution is warmed until the precipitate subsides, allowed to stand
for a few minutes, and is then filtered and the precipitate washed with
hot water containing about one per cent of NH4CI (§139, 2). The filtrate
should be again tested with HgS and if complete precipitation has been
obtained it is set aside to be tested for the metals of the succeeding groups
(§191). The well washed precipitate of the sulphides of Co , Ni , Mn , and
Zn is digested on the filter or in a test-tube with cold dilute HCl (one part
of reagent HCl to four of water) : MnS + 2HC1 = MnClj + HjS . The
black precipitate remaining undissolved contains the sulphides of Co and
BR, the filtrate contains Mn and Zn as chlorides with an excess of HCl
and the HjS which has not escaped as the gas.
§139. Notes. — (1) Instead of passing* the H2S into the ammoniacal solution, a
freshly prepared solution of ammonium sulphide may be used. The yellow
ammonium sulphide, (NH4)3Sx, should not be employed to precipitate the
metals of the fourth group, as nickel sulphide is quite appreciably soluble in
that reagent (§133, 6e).
(2) The sulphides of the fourth group, especially MnS and ZnS , should not
be washed with pure water, as they may be changed to the colloidal sulphides,
soluble in water. The presence of a small amount of NH4CI prevents this, and
does not in any way interfere with the analysis of the succeeding groups.
(3) If the precipitates are to be treated on the filter with the dilute HCl,
the acid solution should be poured on the precipitate three or four times. For
digestion in a test tube, the point of the filter is pierced and the precipitate
washed into the test tube with as little water as possible.
(4) The sulphides of Co and Ni are not entirely insoluble in the cold dilute
HCl , and traces of them may usually be detected in the precipitate for Mn
(§137, footnote).
(5) Dilute acetic acid readily dissolves MnS but scarcely attacks ZnS (§135.
6c). If desired, dilute acetic may be used, first removing the Mn and then
adding dilute HCl to dissolve the Zn .
(6) If large amounts of iron are present, a portion of the Mn will always
appear in the third group (§134, 6a), and is detected by the green color of the
fused mass when testing for Cr: 3Mn(OH)2 + 4KN0, -f- NazCO. = 2KoMnO« -f
Na^MnO^ -f- 4N0 + COj -f 3H2O . Too much HNO, in the oxidation of the
iron favors this precipitation of Mn with Fe'" due to the oxidation of the Mn to
the triad or tetrad combination.
§140. Manipulation. — The hlack precipitate of cobalt and nickel sul-
phides should first be tested with the borax bead (§141, S) for the blue
bead of cobalt (delicate and characteristic but obscured by the presence
of an excess of nickel (§132, 7)). The sulphides are then dissolved in hot
HCl , using a few drops of HNO3 (§141, 1), and boiled to expel excess of
HNOa : 6C0S + 12HC1 + 4HNO3 = 6C0CI0 + 38^ + 4lSrO + 8H0O .
Divide the solution into three portions: To one portion of the solution
§142. DIRECTIONS FOR ANALYSIS WITH NOTES. 185
add an excess (§142, 2) of nitroso- /S-Naphthol, filter, and wash with hot
water and then with hot HCl (§132, 65). Test the red precipitate with
the borax bead for cobalt, llender the filtrate ammoniacal, filter again
and test this last filtrate with H.S for the black precipitate of NiS (§133,.
i'}b and e). To another portion of the solution add NaHCOs ^^ excess,
then add HsOj^ warm and filter, a green color to the filtrate indicates
cobalt (§132, 10). The third portion of the solution is boiled with an
excess of NaOH , bromine water (10, §§132 and 133) is added and the solu-
tion is again boiled. The black precipitate of the higher hydroxides
(§141,-4) of Co and Ni is thoroughly washed with hot water and then
treated on the filter with hot solution of KI (§133, 6/), catching this last
filtrate in a test-tube containing CSg (§141, 6), Free iodine is evidence of
the presence of nickel.
§141. Notes.— (1) HNOs interferes with the nitroso- /?-naphthol reaction that
follows the solution of the sulphides of Co and Ni » hence an excess is to be
avoided. A crystal of KClOs may be used instead of HNO, .
(2) If an insufficient amount of nitroso- )8-naphthol has been used a portion
of the cobalt may be in the filtrate and will give the black precipitate for
nickel. The filtrate must be tested with the reagent to insure complete
removal of the cobalt.
(.?) Test with the borax bead as follows: Make a small loop on the end of a
platinum wire, dip this loop when hot into powdered borax, and heat the
adhering mass in the flame until a uniform transparent glassy bead is obtained.
Repeat until a bead the size of a kernel of wheat has been made. Bring this
hot bead into contact with the precipitate or solution to be tested and fuse
again in the burner flame. Allow the bead to cool and notice the appearance.
A deep blue indicates cobalt, obscured, however, by a large excess of nickel.
(4) The nickel and cobalt may also be oxidized for the KI test as follows:
Add five or ten drops of bromine to the solution to be tested in a beaker,
warm on a water bath under the hood until the bromine is nearly all expelled,
then add rapidly an excess of a hot saturated solution of Na^COs . The black
precipitate so obtained will filter rapidly.
(o) The test for nickel by adding KI to the mixed higher oxides of cobalt
ailH nickel is characteristic of nickel and is also a very delicate test. Fully
nine-tenths of the cobalt salts sold for chemically pure, show the presence of
nickel by this test.
(6) In the reaction of nickelic hydroxide with potassium iodide some potas-
sium iodate is formed and a greater amount of free iodine will be obtained if
a drop of hydrochloric acid be added to the filtrate: KIO, + 5KI + 6HC1 =
31, + 6KC1 + 3H,0
(7) If the sulphides of Ni and Co be digested with yellow ammonium sul-
phide, a portion of the NiS will be dissolved (§133, 6c) and may be reprecipi-
tated as a gray precipitate (black with free sulphur) upon acidulating the
filtrate with acetic acid. It is not a delicate test.
§142. Manipulation. — The solution of the sulphides of manganese and
zinc in cold dilute hydrochloric acid is boiled thoroughly to insure the
removal of the hydrosulphuric acid (§143, 1), cooled (§135, 6a), and then
treated with an excess of sodium hydroxide. The zinc forms the soluble
zincate, IfasZuOj , while the manganese is precipitated as the hydroxide,
white, rapidly turning brown by oxidation :
MnCl, + 2NaOH = Mn(OH), + 2KC1
ZnCl, + 4NaOH = Na,ZnO, + 2NaCl + 2H,0
186 ANALYSIS OF IRON AND ZINC GROUPS, §143, 1.
Filter and test the filtrate with HjS , a white or grayish-white precipitate
indicates zinc (characteristic). Dissolve the well washed precipitate of
Mn(0H)2 in nitric acid and boil with an excess of lead peroxide, adding
more nitric acid. A violet color to the nitric acid solution indicates the
presence of manganese (very delicate and characteristic) :
2Mn(0H), + 5PbO, + lOHNO. = 2HMiiO« + 5Pb(N0.), + 6HaO
§143. Notes. — 1. If the H,S is not completely removed the Zn will be pre-
cipitated as the sulphide upon adding the NaOH , and will not be separated
from the manganese: ZnCl, + HjS + 2NaOH = ZnS + 2NaCl + 2H,0 .
2. Frequently the precipitate of zinc sulphide is dark gray or almost black.
This is usually due to the presence of traces of other sulphides. If iron has not
been all removed, through failure to oxidize completely with the nitric acid,
it may appear as a precipitate with the manganese, and also as a black precipi"
tate with the zinc sulphide.
3. Small amounts of Co and Ni are frequently dissolved by the cold dilute
HCl and will appear with the precipitate of Mii(OH)a . They do not interfere
with the final test for manganese.
4. The precipitate of Mii(OH)j must be washed to remove all the chloride,
as the manganese will not be oxidized to permanganic acid until the chloride
is completely oxidized to chlorine.
5. Instead of PbO^ , red lead, Pb804 , is frequently employed with the nitric
acid to oxidize the manganese to permanganic acid:
2Mn(OH)2 4- SPb.O^ + 30HNO, = 2HMn04 + 15Pb(N0.)j + 16H,0
6. It is very difficult to procuie PbOj or Pb804 which does not contain traces
of manganese. The student should always boil the lead oxides with nitric acid,
and if a violet-colored solution is formed, this should be decanted and the
operation repeated until the solution is perfectly colorless after the black
precipitate of PbOj has subsided. Then the unknown solution in HNO, may
be added and the boiling repeated to test for the manganese.
7. The student is not advised to apply the permanganate test to the original
substances. All reducing agents interfere, and MnO. frequently fails to give
permanganic acid when boiled with PbO, and HNO. until after reduction
(§134. 6c).
Analysis of Iron and Zinc Groups after Precipitation by Ammonium
Sulphide.
§144. It is preferred by some to precipitate the metals of the third
and fourth groups together, by m^ans of ammonium snlphide; using
ammoninm chloride to prevent the precipitation of magnesium (§189, bh
and 6a), and to insure the complete precipitation of the aluminum as the
hydroxide §124, 6a). In the manipulation for this method of separation,
the HgS is not removed from the second group-filtrate, nor is nitric acid
used to oxidize any iron that may be present. To the second group filtrafe
(§80), warmed, an excess of NH^Cl is added (§189, 5c), then NH^OH till
strongly alkaline, and, paying no attention to any precipitate that may be
formed {6a, §§124, 125 and 126), normal ammonium sulphide is added (or
what is equivalent HoS is passed into the alkaline mixture). Aluminum
and chromium are precipitated as the hydroxides, the remaining metals as
the sulphides. The following table illustrates a plan of separation of the
ammonium sulphide precipitates of the third and fourth group metals,
phosphates being absent :
§144.
ANALYSIS OF IRON AND ZINC GROUP.
187
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188 IROy AXD ZINC GROUPS. §145.
§145. The presence of phosphates greatly complicates the work of the
analysis of the metals of the third, fourth, and fifth groups. The phos-
phates of the alkali metals are soluble, those of the other metals insoluble
in water. As the solutions for precipitation of first and second group
metals are acid; phosphates remain in solution and do not in any way
interfere with the analysis for the metals of those groups; i. e., silver
phosphate in nitric acid solution is readily transposed by HCl ; copper
phosphate in acid solution is readily transposed by H2S ; etc.
§146. When the filtrate from the second group is rendered strongly
ammoniacal (§128) the phosphates of all the metals present, except those
of the alkalis, are precipitated. Phosphates of cobalt, nickel and zinc are
redissolved by an excess of ammonium hydroxide. Freshly precipitated
ferric phosphate is transposed by the alkali hydroxides (incompletely in
the cold). The phosphates of Al , Cr , and Zn are soluble in the fixed
alkalis, the solution of chromium phosphate is decomposed by boiling,
precipitating Cr(0H)3 and leaving the alkali phosphates in solution.
§147. In analysis a portion of the filtrate from the second group (after
the removal of the HjS) (§128) should be tested for phosphoric acid with
ammonium molybdate (§75, 6d), If phosphates are present the usual
methods of analysis for third, fourth, and fifth groups must be modified.
Several methods have been recommended :
§148. First. — To the filtrate from the second group, HgS, being re-
moved (§128), an excess of the reagent ammonium molybdate is added,
the mixture set aside in a warm place for several hours, until the yellow
ammonium phospho-molybdate has completely formed and settled
(§75, Gd). Filter and evaporate nearly to dryness to remove the nitric acid.
Take up with water and a little hydrochloric acid if necessary to obtain a
clear solution, and remove the excess of molybdenum with HjS (§75, 6e).
From this point proceed by the usual methods of analysis (§§127, 128
and ff.).
§149. Second. — Precipitation of the phosphate as ferric phosphate in
acetic acid solution. This method of separation rests upon the fact that
the phosphates of the fourth group and of the alkaline earths are soluble,
and the phosphates of Al , Cr"' and Fe'", insoluble in acetic acid.
To the filtrate from the second group, freed from H.^S by boiling (128),
and nearly neutralized with Na^COs , an excess of NaCoH^Os is added and
then FeClg solution, drop by drop, as long as a precipitate is formed.
Care must be taken to avoid an excess of FeClg , as the ferric phosphate
is soluble in a solution of ferric acetate. As soon as the phosphate is all
precipitated the blood-red ferric acetate is formed at once, indicating the
presence of a sufficient amount of FeClg . The mixture should be boiled
§161. IRON AND ZINC GROUPS. 181>
to precipitate the ferric acetate as basic ferric acetate (§126, 66) and at
once fQtered.
Upon the addition of the sodium acetate the aluminum and chromium
are precipitated as phosphates, provided there be sufficient phosphate
present to combine with them; if not the whole of the phosphate will be
precipitated and the first drop of FeClg will give a red solution showing
the addition of that reagent to be unnecessary. •
By the above method of manipulation any iron present in the original
solution is in the ferrous condition and does not react to precipitate the
phosphate, as ferrous phosphate is soluble in acetic acid. If the iron has
been previously oxidized with nitric acid it will react with the phosphate
upon the addition of the sodium acetate ; but if there be more iron present
than necessary to combine with the phosphate, the red ferric acetate solu-
tion will be formed with the excess of the iron and render the precipita-
tion of the phosphate incomplete. In this case the previous oxidation of
the iron is detrimental.
If alkaline eari;h salts are present in quantity more than sufficient to
combine with the phosphoric acid radical, not all of these metals will be
precipitated with the third group metals upon the addition of ammonium
hydroxide. The table (§162) illustrates the separation of the metals in
presence of the phosphates by the use of FeClg in acetic acid solution.
§150. Third. — A method of separation of the third group metals with
phosphates from the remaining metals is based upon the action of freshly
precipitated barium carbonate. Solutions of Al , Cr"', and Fe'" are pre-
cipitated as the hydroxides by digestion in the cold with freshly precipi-
tated BaCOj {Qa, §§124, 126 and 126): 2AICI3 + SBaCOg + SH^O =
2A1(0H)3 + SBaClj + SCOg . Solutions of the chlorides or nitrates of
the fourth group and of the alkaline earths are not transposed by cold
digestion with BaCOg . Sulphates of the fourth group are transposed by
freshly precipitated BaCOs in the cold: C0SO4 + BaCOs = BaSO^ -f-
CoCO, , etc.; and must not be present in this method of separation
(§126, ea).
If an excess of ferric chloride be present the phosphates will all be
precipitated as ferric phosphate and the Al, Cr'" and excess of Fe'" as
the hydroxides upon the digestion with BaCO, . The table (§163) gives
an illustration of the use of the BaCOg in effecting the separation.
It should be observed that presence or absence of FeClg or of BaCOs ^^
the sample must be fully determined before their addition as reagents.
§161. Oxalates do not interfere with the usual course of analysis of the
first two groups of metals; with the other metals oxalates interfere very
much the same as phosphates. They, however, with other interfering
190 IRON AND ZINC 0R0UP8. §161.
organic matter, can readily be removed by ignition. If the presence of
an oxalate has been established (§§188, 6h and 227, 8), the second group
filtrate should be evaporated to dryness, moistened with concentrated
HNO3 and gently ignited. The residue, dissolved in HCl, is then ready
for the usual process of analysis. For the analysis in presence of silicates
and borates the student is referred to the text under those elements
<§§249,8and4i21,8).
§168.
IROy, ZINC AND CALCIUM GROUP METALS.
191
.4 «»
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192
IRON, ZINC AND CALCIUM GROUP METALS.
§153.
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§166. . CERIUM— COLUMBIUM. 193
The Rarer Metals of the Iron and Zinc Groups.
€erinm, Coltunbitun (Niobium), Didymium, Erbium, Oallium, Olucinum
(Beryllium), Indium, Lanthanum, Neodymium, Praseodymium, Sama-
rium, Scandium, Tantalum, Terbium, Thallium, Thorium,
Titanium, Uranium, Ytterbium, Yttrium, Zirconium.
§164. Cerium. Ce = 139.0 . Valence three and four.
Specific gravity, 6.628. Melts higher than Sb and lower than Ag (Hillebrandt
and Norton, Pofjy., 1875, 156, 466). Cerium is a comparatively rare metal, never
found native; it is found in many minerals in Sweden, especially in cerite,
which is chiefly a silicate of Ce , La, Ne , Pr , Al and Fe; also found in a
brick-making clay near Frankfurt, Germany (Strohecker, J. pr., 1886, (2), 33,
133 and 260). It was first described in 1803 by Klaproth, but in 1839 Mosander
showed the supposedly pure cerium oxide to consist of oxides of at least three
metals: Ce , La, D (Ne and Pr) (Pogg., 1842, 56, 503). The metal is obtained
from the chloride, CeCl, , by electrolysis or by heating with sodium. It is a
steel-gray, lustrous, malleable, ductile metal; fairly stable in air under ordinary
conditions. When heated in air it burns with incandescence. It burns in CI ,
Br and in vapor of I , S and P . Soluble in acids. Two oxides are known,
COaOs and CeOj , forming two classes of salts, cerous and eerie, the latter being '
less stable. Ignition in air or oxygen changes Ce.O, to CeOa . CejO, is white
or grayish-white, soluble in acids and formed by igniting Ce^^CO,). , Ce3(C30«),
or CeOa in an atmosphere of hydrogen. Cerous salts are white and form color-
less solutions in water. Ceric oxide, CeO^ , is yellowish-white, orange-yellow
when hot, soluble in acids with difficulty; the hydroxide dissolves readily.
Ceric salts are yellow or red, forming yellow solutions. Ceric hydroxide,
Ce(OH)« , dissolves in HCl with evolution of chlorine, forming colorless cerous
■chloride. Sulphurous acid decolorizes solutions of ceric salts, forming cerous
salts. Fixed alkali hydroxides and ammonium sulphide precipitate, from
solutions of cerous salts, the white cerous hydroxide, turning yellow by absorp-
tion of oxygen, with formation of ceric hydroxide. The precipitate is in-
soluble in excess of the fixed alkalis (distinction from Al and Gl). The pre-
<!ipitation is hindered by the presence of tartaric acid (distinction from
yttrium). Ammonium hydroxide precipitates a basic salt. Alkali carbonates
precipitate cerous carbonate, soluble in excess of the fixed alkali carbonates.
Oxalic acid forms cerous oxalate, white, from moderately acid solutions, soluble
in hot (NH4)2C204 , but reprecipitated on dilution with cold water. A con-
centrated solution of K3SO4 forms the double sulphate^ K3Ce(S04), , white,
sparingly soluble in water, insoluble in K3SO4 solution (distinction from Gl).
NajSjO, does not precipitate cerium salts. BaCO, does not precipitate cerous
salts in the cold, but precipitates them completely on boiling. Ceric salts are
completely precipitated by BaCO, in the cold. Alkali hypochlorites precipitate
cerous salts as the yellow ceric hydroxide. If cerous nitrate be boiled with
PbOj and HNOg , ceric nitrate, a deep yellow solution is formed {delicate test
for cerium). Cerium givas no absorption spectrum, but the spark spectrum
shows several brilliant lines.
§155. Columbium (Niobium). Cb = 93.7 . Valence five. •
Columbium usually occurs with tantalum in such minerals as columbite and
tantalite; it is also found in tantalum free minerals as euxenite, pyrochlor, etc.
The metal is prepared by passing the penta-chloride mixed with hydrogen
repeatedly through a hot tube. It is a steel-gray lustrous metal, specific
gravity, 7.06 at 15.5**. By ignition in the air it burns readily to the pentoxide.
Not attacked by chlorine in the cold, but when warmed combines readily,
forming CbCl, . The metal is not soluble in hydrochloric, nitric or nitrohydro*
194 DIDYMIUM. §156.
chloric acids, but is readily soluble in hot concentrated sulphuric acid, forming*
a colorless solution (Koscoe, C. A'., 1878, 37, 25). It forms several oxides, CbO ,
CbO, and CbsOt . Columbic acid (anhydride) Cb,Ot , is a white powder, yellow
when hot (distinction from tantalum); it is obtained by ignition of the low^er
oxides, or by decomposition of solutions of the salts by water or alkalis and
igniting. CbOj , black, is prepared by strongly igniting CbjOs in a current of
hydrogen. Cb^OB , not too strongly ignited, is soluble in acids, from which
solutions NH4OH and (NH4)tS precipitate colttmbic acid containing some am-
monia. By mixing Cb^Oa with charcoal and heating in a current of chlorine, a
mixture of CbOCls and CbCls is obtained. CbGla is a yellow crystalline solid
(needles), melting at 194** and distilling at 240.5** (Deville and Troost, C. r., 1867.
64, 294). Upon treating the chloride with water, it is partially decomposed
to columbic acid, a large portion remaining in solution and not precipitated
by HjSOf (distinction from tantalum). Cb-^Oo not previously ignited dissolves
in HF; which solution when mixed with KF , the HF being in excess, gives
a double fluoride, 2KF.CbF5; if the HF be not in excess, a double oxy-fluoride
is obtained, 2KF.CbOF. (Kruess and Nilson, B,, 1887, 20, 1676). The potassium
columbium fluoride is much more soluble than either the corresponding tita-
nium or tantalum compounds. Fusion of columbic acid with the alkalis gives
the columbates, the potassium salt being quite soluble in water and in potas-
sium hydroxide; the sodium salt is only soluble in water after removal of the
excess of the sodium hydroxide. From a solution of potassium columbate,
sodium hydroxide precipitates, almost completely, sodium columbate. Carbon
dioxide precipitates columbic acid from solutions of columbates. Soluble salts
of Ba , Ca and Mg form white bulky precipitates with a solution of potassium
columbate. AgNO, gives a yellowish-white precipitate, CuSO« a green pre-
cipitate. CbgOs in presence of HCl or H2SO4 gives a blue to broicn color with
Sn or Zn, due to partial reduction of the Cb (distinction from tantalum).
Fused with sodium meta-phosphate, columbic acid gives in the inner flame a
violet to blue bead; a red bead by addition of FeSO« .
§156. Didyiniiim = |?f^y^'^"'
^ I Fraseodymii
Nd = 143.6 . Valence three.
Praseodymium. Pr =-- 140.5 . Valence three.
Specific gravity, G.544. Melts with greater difliculty than Ce or La . Present
in cerite in Sweden and in monazite sand from Brazil. Didymium was reported
about 1840 by Mosander, having been separated from cerium and lanthanum.
In 1885 Welsbach (J/., 1885, 6, 477) separated didymium salts into two distinct
salts, neodymium and praseodymium. By the absorption spectrum bands
other chemists are of the opinion that the so-called didymium consists of a
group of elements, nine or more (Kruess and Nilson, B., 1887, 20, 2166; Kreuss,
A., 1892, 265, 1). Concerning the separation of didymium compounds, see
Dennis and Chamot (./. Am. Soc, 1897, 19, 709). By repeated fractionation of
the nitrate (several thousand times) Welsbach obtained a pale green salt and
a rose-colored salt, which gave different spectra but which, united, gave the
spectrum of didymium. Didymium oxide absorbs water to form the hydroxide,
which absorbs CO3 from the air, but does not react alkaline to litmus. The
salts are soluble in water to a reddish solution. The saturated sulphate solu-
tion does not deposit crystals until heated to Imling; while lanthanum sulphate
precipitates from the saturated solution at 30°. Fixed alkalis precipitate the
hydroxide: NH,OH , a basic salt: insoluble in excess of the reagents. Alkali
carbonates form a bulky precipitate, insoluble in excess of the reagent, barium
carbonate precipitates slowly but completely. Precipitation by alkalis is pre-
vented by tartaric acid. Oxalic acid precipitates didymium salts completely,
soluble with diflRculty in HCl . The double potassium sulphate forms much
more slowly and less completely than with cerium. The salts give a distinct
and i'haraei eristic absorption spectrum. Consult Jones {Am., 1898, 20, 345),
Scheie (Z. anorg., 1898, 17, 319), Boudard (C. r., 1898, 126, 900), Demarcay
(C. r., 1898, 126, 1039), and Brauner (C. .V., 1898, 77, 161).
§169, ERBIVM-^ALLIIM—GLUCINUM. 195
§167. Erbium. Er = 166.0 . Valence three.
Erbium metal has not been prepared. As oxide or earth it is described by
Cleve (C. r., 1880, 91, 381) as that yttrium earth the most beautiful rose
colored. It forms a charavi^riHtic absorption spectrum, and a spark spectrum
with sharp lines in the orange and green. This earth has not been thoroughly
studied and quite probably consists of the oxides of several metals (Boisbau-
dran, C. r., 1886, 102, 1003; Soret, C. r., 1880, 91, 378; Crookes, C. 2V^., 1886, 54,
13). The oxide gives upon ignition an intense green light; it is not fusible or
volatile.
§168. Oalliiiin. (Ja = 70.0 . Valence three.
Specific graritif; the solid, at 23** to 24.5*', 5.935 to 5.956; the melted, at 24.7**,
6.069. Melting point, 30.15^*; frequently may be cooled to O'* without again be-
coming solid. It is a grayish-white metal, crystallizing in octahaedra or in
broad plates. It is quite brittle and gives a bluish-gray mark on paper. It
gfives a very weak and fugitive flame spectrum; the spark spectrum shows two
beautiful violet lines. When heated in the air or in oxygen it is but slightly
oxidized; does not vaporize at a white heat; soluble in acids and alkalis;
attacked by the halogens (with iodine only upon warming). In the Periodic
System it is the Ekaaluminum of Mendelejeff, who described the general prop-
erties before the metal was discovered (C r., 1875, 81, 969). It occurs in zinc
blende (black) from Bensberg on the Rhine; in brown blende from the
Pyrenees; and in some American zinc blendes (Cornwall, Ch. Z., 1880, 4, 443).
It is prepared by electrolysis after previous purification of the ore by chemical
methods. 4300 kilos of the Bensberg ore gave 55 kilos of pure gallium (Bois-
baudran and Jungfleisch, C, r., 1878, 86, 475). The oxide, GhtjOa , is a white
powder obtained by igniting the nitrate. After strong ignition it is
insoluble in acids or alkalis. It is easily attacked on fusion with KOH
or KHSO4 . The alkalis and the alkali carbonateB precipitate the salts
as the hydroxide, perceptibly soluble in fixed alkali carbonates, more easily
in ammonium hydroxide and in ammonium carbonate, atid very readily in
the fixed alkalis. Tartrates hinder the precipitation of the hydroxide. The
salts of gallium are colorless and for the most part soluble in water. The
neutral solutions upon warming precipitate a basic salt, dissolving again upon
cooling. Excess of zinc forms a basic zinc salt which precipitates the gallium
as oxide or basic salt. BaCO, precipitates gallium salts in the cold. K,Fe(CN)«
gives a precipitate, insoluble in HCl , noticeable in very dilute solutions
(1-175,000). H2S does not precipitate gallium salts from solutions acid with
mineral acids; from the acetate or in presence of ammonium acetate the irhite
fftUphide, Gb.Sz , is precipitated; (NH4)2S precipitates the sulphide. Gallium
chloride, GaCl, , is a colorless salt, melting at 75** and volatilizing at 215° to
220**. The vapor density indicates the molecule to be Gk^Cl« , which decomposes
to GaCl, at about 400" (Friedel and Kraft, C. r., 1888, 107, 306). Upon evaporat-
ing a solution of the chloride on a water bath the salt is perceptibly volatil-
ized, not so if H^SOt be present. Gallium sulphate forms with ammonium
sulphate an alum. For separation from other metals, see Boisbaudran, C. r.,
1882, 95, 410, 503, 1192, 1332.
§159. Olncinum (Beryllium). 01 = 9.1 . Valence two.
Specific gravity, 1.85 (Humpidge, Proc. Roy. Soc, 1871, 39, 1). Melting point,
below 1000 (Debray, A. Ch., 1855, (3), 44, 5). It is a white malleable metal,
obtainable in hexagonable crystals (Nilson and Pettersson, B., 1878, 11, 381
and 906). It was first discovered in 1797 by Vauquelin from beryl. It is
stable in the air, does not decompose steam at a red heat, and at red heat is
scarcely attacked by oxygen or sulphur. It is a strongly positive element,
196 IXDIUM. §160.
in greneral properties between ahiminum and the alkaline earths: as lithium
is between the alkaline earths and the alkali metals. It should be classed
with the alkaline earths. It is found in ehrysoberyl, Ol(A102)2 , in phenakite,
GLSiOf , and in some other silicates. It is prepared by heating the chloride.
GlGl;^ , with Na in a closed iron crucible (Nilson and Pettersson, /. c); or bv
heating the oxide, GIO , with Mg (Winkler, /?., 1890. 23, 120). The oxide, G10\
is obtained by igniting the hydroxide. It is a white infusible powder, soluble
in acids and in fixed alkalis. The hydroxide is prepared by precipitating the
salts with NH^OH , soluble in the fixe^ alkaUn and in ammonium carbonate,
concentrated: precipitated on dilution and boiling (distinction and sejjaration
from Al). The metal is soluble in acids except that when in the compact
form it is scarcely attacked by HNOa . The hydroxide is soluble on continued
boiling with NH4CI , forming GIGL . The more common salts of gliicinum are
soluble in water to a solution having a sweetish taste. The carbonate and
phosphate are insoluble, the oxalate and sulphate soluble, the existence of a
sulphide is doubtful. Solutions of glucinum salts are precipitated by the
alkalis, the precipitate being soluble in excess of the fixed alkalis. The alkali
carbonates precipitate the carbonate, soluble in concentrated ammonium car-
bonate, reprecipitated on diluting, boiling and adding an excess of NH4OH
(Joy, Am. ^'., 1863, (2), 36, 8:5). The salts are not precipitated by H^S , but are
precipitated by (NH4)2S as the hydroxide. BaGO, does not precipitate Gl salts
in the cold, but precipitates them upon boiling. GlCl:; melts at about 600"
and sublimes at a w^hite heat, forming white needles. The oxide has not been
melted or sublimed. Gl usually occurs as a silicate with aluminum. The
mass is fused with alkali carbonate, acidified with HCl and the Al and 01
chlorides filtered from the SiOa . An excess of ammonium carbonate precipi-
tates both metals, but redissolves the Gl . After repeating this separation
several times pure glucinum hydroxide, Gl(OH)a , is obtained upon boiling off
the ammonia. The hydroxide thus obtained is ignited and weighed as the
oxide.
§160. Indium. In = 114.0 . Valence three.
specific granty^ 7.11 to 7.28 at 20.4°. Meltinfl point, 17G°. Indium was discov-
ered in Freiberg zinc blende by Reich and Richter (J. pr., 1863, 89, 441: 90, 175:
1864, 93, 480), by use of the spectroscope. It is found chiefly as sulphide, never
native, in the Freiberg blende to the extent of about 0.1 per cent. It is found
in a few other places, but in much smaller amounts (Boettger, J. pr., 1S66, 98,
26). In the preparation of indium the Freiberg zinc is dissolved in HCl or
S2SO4 , leaving an excess of the zinc. When no more hydrogen is evolved, the
mass is digested for a day or more with the excess of Zn , whereby the indium
is obtained as a precipitate with Pb , Cu . Cd , Sn , As, Fe and Zn . This
precipitate is dissolved in nitric acid and evaporated with sulphuric acid: then
taken up with water separating from lead. The solution is precipitated with
NH4OH , which precipitates the In and Fe; this precipitate is dissolved in
HCl and boiled for some time with NaHSC, . The indium sulphite is obtained
as a fine crystalline powder, which is treated with HNO, and H2SO4 , forming
indium sulphate, from which the metal is precipitated by zinc (Bayer, J.., ISTl.
158, 372: Boettger, J. pr., 1869, 107, 39: Winkler, J. pr., 1867, 102, 276). Indium
is a grayish-white metal, very soft, makes a good mark on paper, is ductile,
easily fusible, less volatile than Zn or Cd. It is less electro positive than Zn
or Cd and hence it is precipitated from its solutions by both these elements.
In the air or in water it is rather more stable than zinc. Heated in the air it
burns with a violet flame and brown smoke, forming the oxide, In.O, , Indium
does not decompose water at 100°. At a red heat it combines with sulphur
and the halogens. By ignition with charcoal or in a current of hydrogen it is
reduced to the metal from its compounds. It is soluble in HCl and HJ30«.
evolving H: in HNO, , evolving NO . In the reactions of its salts indium
deports itself quite similar to Fe'" and Al . Its most characteristic property is
its spectrum: two lines, an indium a, intense blue, and an indium iS , less
intense violet (Schroetter, J, pr„ 1865, 95, 441). In^O, is brown when hot.
§184. LANTHAyiM-SAMARIUM, 197
light yellow when cold, slowly soluble in cold acids, rapidly when heated.
Indium salts are precipitated by the alkalis as Iii(OH), , soluble in excess of
the fixed alkalis, reprecipitated by boiling" or treating with NH4GI . Tartrates
prevent the precipitation by alkalis. Alkali carbonates precipitate the indium
carbonate, soluble in ammonium carbonate, but reprecipitated on boiling.
^aCOa carbonate precipitates the indium completely as a basic salt (separation
from Co , Ni , Mn , Zn and Fe")- Phosphates form white precipitates from
neutral solutions. HjS precipitates frOni neutral solutions, or solutions acid
with acetic acid, yellow indium sulphide. In alkaline solutions HjS , or in
neutral solutions {'NMi)3 , forms a white precipitate containing In^Ss . Yellow '
XnA boiled with (NH4)2Sx becomes white and is partly dissolved. Upon cool-
ing the solution a bulky white precipitate separates out. KtFe(CN)a gives a
white precipitate: K,Ct04 gives a yellow precipitate; EaCrzOr , K^lPe(CN)t and
ISCNB do not form precipitates.
§161. Lanthannin. La = 138.6 . Valence three.
Specific gravity, 6.163. Melts somewhat higher than Ce . In general appear-
ance and properties very similar to Ce . It is prepared almost exclusively from
cerite. By treating the mineral with an insufficient quantity of HNO3 , a
solution rich in La may be obtained. The cerium is precipitated from the
solution by alkali hypochlorite. The filtrate is converted into the sulphate and
separated from Ne and Pr sulphates by fractional crystallization, the latter
being more soluble (Holzman, J. pr., 1858, 75, 346). Fractional precipitation
with NHtOH is also used to separate I*a from Ne and Pr , the latter precipitat-
ing first (Cleve, Bl., 1874, 21, 196; 1883, 39, 287). The metal is prepared from
the chloride, LaCl, , by electrolysis or by ignition with potassium. The igni-
tion point of I*a is higher than that of Ce; it is also not so readily attacked
by HNO3 . In cold water La is slowly attacked, but in hot water the- action
is violent (Winkler, B., 1890, 23, 787). The oxide, La^O, , is a white powder,
readily soluble in acids; with water it forms the hydroxide, La(OH)s , which
reacts alkaline towards litmus and absorbs CO3 from the air. La(OH), is
soluble in a solution of NH4CI (similar to Mg(0H)2). The salts are colorless.
XL2SO4 and H2C2O4 form precipitates with lanthanum salts as with cerium salts.
Fixed alkalis precipitate lanthanum salts as La(0H)3 , white, insoluble in
excess of the reagent and not changing color on exposure to the air (distinc-
tion from Ce). Alkali carbonates precipitate La2(COj)3 , insoluble in excess.
BaCOe precipitates the salts completely in the cold. NH4OH precipitates basic
salts. HjS forms no precipitate; (NH4)2S precipitates the hydroxide. Lantha-
num gives a number of characteristic lines in the spark spectrum (Bettendorf,
A., 1889, 256, 159).
§162. Ncodymium. Nd = 143.6 . See Didymium (§156).
§163. Praseodymium. Pr = 140.5 . See Didymium (§156).
§164. Samarium. Sm = 150.3 . Valence three.
Samarium was found in 1879 by Boisbaudran from didymium earths by its
peculiar spectrum (C. r., 1879, 88, 323). According to Crookes (C. r., 1886, 102,
1464), it consists of at least two elements and is found in all yttrium earths.
Its salts are light yellow, giving an absorption spectrum of six bands (Kruess,
jB., 1887, 20, 2144). In its chemical properties it is more similar to Nd and Pr
than to Y . It is separated from Nd and Pr by the fractional precipitation of
the hydroxide, basic nitrate, oxalate and sulphate; which separate before the
corresponding Nd and Pr compounds.
198 SCAyDIUM—TANTALUM—TERBIUM, §165.
§165. Scanditun. Sc = 44.1 . Valence three.
It is found in euxenite and gadolinite with yttrium. Its name comes from
Scandinavia, where it was first found. It is separated from ytterbium, with
which it is always closely associated, by heating the nitrates; the basic scan-
dium nitrate being precipitated before the ytterbium basic nitrate, or by
precipitating as the double potassium, sulphate, the corresponding ytterbium
salt remaining in solution. The oxide, ScjO, , is a white flocculent infusible
powder, readily soluble in warm acids. The solutions of the salts show no
absorption bands in the spectrum. The spark spectrum of the chloride gives
over 100 bright lines (Thalen, C. r., 1880, 91, 45). Solutions of the salts taste
sweet and have an astringent action. The alkalis precipitate the hydroxide,
a wrhite bulky precipitate, insoluble in excess of the precipitant. Tartrates
hinder the precipitation in the cold, but not upon heating. Na^CO, gives a
bulky white precipitate, soluble in excess of the reagent, H^ is without
action, but (NHt)S precipitates the hydroxide. K2SO4 precipitates the double
scandium sulphate, SKaSO^JSCaCSOi), , soluble in water but not in a saturated
K.BO4 solution.
§166. Tantalum. Ta = 182.8 . Valence five.
Tantalum occurs in tantalite and columbite, silicates, nearly always ac-
companied by columbium. It is prepared by heating the tantalum alkali
fluoride with K or Na in a well-covered crucible (Rose, Pogg., 1856, 89, 65). It
is a black or iron-gray powder with a metallic lustre. Specific gravity^ 10.78.
Heated in the air it burns with incandescence to form Ta»Os . It is insoluble
in acids except HF , in which it dissolves with evolution of H . Upon ignition
in a current of chlorine, TaClj , volatile, is formed. Solution of alkalis has
no action, upon fusion with the fixed alkalis an alkali tantalate is formed.
Ta^Os is a white infusible powder, specific gravity, 8.01 (Marignac, A, Ch,, 1866,
(4), 9, 254). The oxide fused with fixed alkalis gives also an alkali tantalate,
M'TaO, . When KOH is used, the fused mass is soluble in water. When NaOH
is used, water removes the excess of alkali, leaving the NaTaO, as a white
residue, which dissolves in pure water, but not in NaOH solution. Tantalum
chloride is a yellow solid, melting at 211.3° and boiling at 241.6°, with 75.']
mm. atmospheric pressure (Deville and Troost, C. r., 1867, 64, 294). It is com-
pletely decomposed by water, forming the hydrated acid, 2HTaOi.S[30 =
'H,4Tb..jOj . The freshly precipitated acid is solu ble in acids and reprecipitated
by NH4OH . The acid is readily soluble in HT , which solution with KF forms
a characteristic double salt, 2KF.TaP5 , crystallizing in fine needles, insoluble in
water slightly acidulated with HF (distinction and separation from colum-
bium). A solution of alkali tantalate gives with HCl a precipitate of tnntalic
acid, soluble in excess of the HCl . From this solution NH4OH or (NHJjS
precipitates tantalic acid; H0SO4 precipitates tantalic sulphate. Tartaric acid
prevents the precipitation with NH4OH and (NH4)aS . A solution of tantalic
acid gives no coloration with zinc (distinction from CTb). Solutions of alkali
tantalates form tantalic acid with CO2 . The acid fused with sodium meta-
phosphate gives a colorless bead (distinction from SiOj), which does not become
blood-red upon adding FeSO« and heating in the inner flame (distinction from
titanium).
§167. Terbium. Tr = 160. Valence three.
The terbium compounds are very similar to the yttrium compounds. The
salts are colorless and give no absorption spectrum. The double potassium
terbium sulphate has about the same solubilities as the corresponding cerium
compound, and so the terbium is frequently precipitated with cerium com-
pounds. Terbia, TPjOg , is the darkest colored of the yttrium earths, soluble
§169. THALLIUM—THORIUM. 199
in acids and sets ITHa free from ammonium salts. The hydroxide is a
gelatinous precipitate which absorbs CO, from the air. It is quite probable
that terbia is a mixture of rare earths (Boisbaudran, C, r., 1886, 102, 153, 395,
483 and 899).
§168. Thallium. Tl = 204.15 . Valence one and three.
Thallium was discovered by Crookes by means of the spectroscope in 1861,
in selenium residues of the H^SO^ factory at Tilkerode in the Hartz Mountains,
Germany (C. A'., 1861, 3, 193, 303; 1863, 7, 290; 1863, 8, 159, 195, 219, 231, 243,
255 and 279). It is found widely distributed in many varieties of iron and
copper pyrites, but in large proportions it is only found in Crookesite in
Sweden. This mineral contains as high as 18.55 per cent Tl (Nordenskjoeld,
J.., 1867, 144, 127). It is prepared by reduction from its solutions with Zn or
Al; by electrolysis; by precipitation with KI , and then reduction by Za or Al
or by electrolysis. Specific gravity, 11.777 to 11.9 (Werther, J. pr.y 1863, 89, 189).
Melting point, 290° (Lamy, C. r., 1862, 54, 1255). It is a bluish-white metal,
softer than lead, malleable and ductile; tarnishes rapidly in the air; may be
preserved under water, which it does not decompose below a red heat; soluble
in HaSO^ and HNO, , in HCl with great difficulty; combines directly with
CI , Br , I , P , S , Se , and precipitates from their solutions Cu , Ag , Hg ,
Au and Pb in the metallic state. As a monad its compounds are stable, and
not easily oxidized; as a triad it is easily reduced to the univalent condition.
ThaUious oxide, Ti^O , is black; on contact with water it forms an hydroxide,
TIOH , freely soluble in water and in alcohol, to colorless solutions. The car-
bonate is soluble in about 20 parts of water; the sulphate and phosphate are
soluble; the chloride very sparingly soluble; the iodide insoluble in water.
Sydrochloric acid precipitates, from solutions not very dilute, thallious
chloride, TlCl , white, and unalterable in the air. As a silver-group precipitate,
thallious chloride dissolves enough in hot water to give the light yellow pre-
cipitate of iodide, TU , on adding a drop of potassium iodide solution, the
precipitate being slightly soluble in excess of the reagent. HaS precipitates
the acetate, but not the acidified solutions of its other salts. (N1S.^)..S pre-
cipitates TI3S , which, on exposing to the air, soon oxidizes to sulphate.
Perrocyanides give a yellow precipitate, Tl^PeCCN),; phosphomolybdic acid a
yellow precipitate; and potassium permanganate a red-brown precipitate, con-
sisting in part of TlaO, . Chromates precipitate yellow normal chromate; and
platinic chloride, pale orange, thallious platinic chloride, TljPtCl«. Thallium
compounds readily impart an intense green color to the flame, ^ind one emerald-
green line to the spectrum (the most delicate test). The flame-color and
spectrum, from small quantities, are somewhat evanescent, owing to rapid
vaporization. Thallic oxide, TljOa , dark violet, is insoluble in water; the
hydroxide, an oxyhydroxide, TIO(OH), is brown and gelatinous. This hydrox-
ide is precipitated from thallic salts by the caustic alkalis, and not dissolved
be excess. Chlorides and brom.ides do not precipitate thallic solutions; iodides
precipitate Til with I. Sulphides and HjS precipitate thallious sulphide, with
sulphur. Thallic oxide, suspended in solution of p>otassium hydroxide, and
treated with chlorine, develops an intense violet-red color. Thallic chloride
and sulphate are reduced to thallious salts by boiling their water solutions.
§169. Thorium. Th = 232.6 . Valence four.
Thorium is a rare elemept foimd in thorite (a silicate), orangite and some
other minerals. It was described by Berzelius in 1828 {Pogg., 1829, 16, 385),
who also prepared the metal by reduction of the potassium thorium fluoride
ivith potassium. The metal is a gray powder; «peci/fc gravity, 11.000; stable in
air at ordinary temperature, but igniting when heated; attacked by vapors of
CI , Br , I and S . Sparingly soluble in dilute acids, easily soluble in concen-
trated acids; insoluble in the alkalis (Nilson, B„ 1882, 15, 2519 and 2537; Kruess
200 TITAXIUM. §170.
and Xilfion, B., 1S87, 20, 1665). Thorium forms one oxide, ThOj , upon ignition
of the oxalate. It is a snow-white powder, not easily soluble in acids if highly
ignited (Cleve, */., 1874, 261). The hydroxide, ThCOH), , is formed by precipita-
tion of the salts by the alkalis. It is a white, heavy, gelatinous precipitate*
drying to a hard glassy mass. The chloride, ThCl^ , and the nitrate, Tli(KOa)« ,
are deliquescent. The chloride is a white body melting at a white heat and then
subliming in beautiful white needles (Kruess and Xilson, Lc). The sulphate
is soluble in five parts of cold water. The carltonnte, oxalate and phosphate are
insoluble in water; the oxalate is scarcely soluble in dilute mineral acids.
Alkali hydroxides or sulphides precipitate thorium hydroxide, Tli(OH)«,
insoluble in excess of the reagent. Tartaric and citric acids hinder the pre-
cipitation. Alkali carbonates precipitate the basic carbonate, soluble in ex-
cess, if the reagent be concentrated. The solution in (NH4).COs readily repre-
cipitates upon warming. Ba.CO, precipitates thorium salts completely. Oxalic
acid and oxalates form a white precipitate (distinction from Al and 01), not
soluble in oxalic acid or in dilute mineral acids: soluble in hot concentrated
('XfH.t)iC/}t and not reprecipitated on cooling and diluting (distinction from
Ce and La). A saturated solution of K3SO4 slowly but completely precipitates
a solution of Th(S04)2 , forming potassium thorium sulphate; insoluble in a
saturated K^0« solution, sparinglj^ soluble in cold water, readily soluble in
hot water. HF precipitates ThF, , insoluble in excess, gelatinous, becoming
crystalline on standing. Boiling freshly precipitated Th(0H)4 with KF in
presence of HF forms K3ThFo.4H;,0 , a heavy fine white precipitate almost
insoluble in water. The distinf/uishinfj reactions of thorium are the precipitation
with oxalates and with K3SO4 , and failure to form a soluble compound on
fusion with Ka,CO| (distinction from SiOj and TiO^).
§170. Titanium. Ti = 48.15 . Valence three and four.
Titanium is found quite widely distributed as rutile, brookite, anatase.
titanite, titaniferous iron, FeTiOj , and in many soils and clays. Never found
native. It is prepared bj' heating the fluoride or chloride with K or Na . It
is a dark gray powder, which shows distinctly metallic when magnified. Heated
in the air it burns with an unusually brilliant incandescence; sifted into the
fiame it burns with a blinding brilliance. Chlorine in the cold is without action,
when heated it combines with vivid incandescence. It decomposes water at
100°. It is soluble in acids, with evolution of hydrogen, forming titanous
chloride. At a higher temperature it combines directly with Br and I. It is
almost the only metal that combines directly with nitroyen when heated in the
air (Woehler and Deville, A., 1857, 103, 230; Merz, ./. pr., 1866, 99, 157). The
most common oxide of titanium is the dioxide, TiOa , analogous to COj and SiO..
It occurs more or less pure in nature as rutile, brookite and anatase; it is
formed by ignition of the hydrated titanic acid or of ammonium titanate
(Woehler, J., 1849, 268). Ignition of TiOa in dry hydrogen gives Ti,0, , an
amorphous black powder, dissolving in H-^O^ to a violet-colored solution (Ebel-
men, A. Ch., 1847, (3), 20, 392). TiO is formed when TiO, is ignited with Mg;
2TiO, 4- Mg = TiO + MgTiO, (Winkler, B., 1890, 23, 2660). Other oxides have
been reported. Titanic acid, TiOj , is a white powder, melts somewhat easier
than SiOa , soluble in the alkalis unless previously strongly ignited. Mixed
with charcoal and heated in a current of chlorine TiCl* is formed. The
bromide is formed in a similar manner. TiOj acts as a base, forming a series
of stable salts; also as an acid, forming titanates. TiCl* is a colorless liquid,
fuming in the air; it boils at 136.41° (Thorpe, J. C, 1880, 37, 329); it is de-
composed by Wfitrr, forming titanic acid, which remains in solution in the HCl
present. Solutions of most of the titanic salts, when boiled, deposit the
insoluble meta-titanic acid. HF dissolves all forms of titanic acid; if the
solution be evaporated in presence of H3SO4 no TiF* is volatilized (distinction
from SiFJ. When evaporated with HF alone, TiF< is volatilized. The double
potassium titanium fluoride, K^TiFe , formed by fusing TiO- with acid KF , is
sparingly soluble in water (96 parts), readily soluble in HCl . Solutions of
titanic salts in water or acid solutions of titanic acid are precipitated by
il7L
IRAXllM,
201
alkali Jiydroxides, carbonates ami HHlphhi^if t\^ the hifdrftfvil tltonh* tivuh insolu-
ble ill t'xet^ss of tin' pr<Hn[)itiiiits unrj in tmmionUmi suits. BaCO, givtrs iht' same
l>rt?cipitate* K«Fe(CN)^ iriVL'.s a rt-ddi^h^jellovv preripititte; E^FeiCN), u yelluw
precipitate, Ka^HPO^ precipitates the titanUiiii (ihnmt rmnpkiviih even in the
presence of st rtm^' HCl . An acid solution of TIO^ when treated with Sn or
ZiL irives a imk' hluv to ikiki tmktrtitktn to the solution^ due to h partial rfdttftioH
of the titaiiiiini to the triad condition. These colon-d ^lolntions are preeipit filed
by alkali hydroxides, carbonates and sulphides, h:,S ]s withont action. The
solution rerlTices Fe'" to Fe" , Cu" lo Cil^ , and Kails of Hg , Ag" and Au to the
metallic etiite; tht* titiiniiim becninhip' iiffain the tetrad. The red net ion by Sn
or Zn takei? plaee in pres^enee of HF {distinction from eolumbic acid). Titanium
cotnponnclts fused in the flame with luirro^^nstnic salt give in the rrdueitipr flame
ai yellow bead when hot* cooliiiir to reddish and violet (red net ion of the tUu-
nium). With Fe80, in tht* redueing tlame a hiood'ird it€fid ib obtained.
§171t TTraMium. V = 239, li . Valence four and sis.
f.fM^i^e fjmnlih l^.nsTj (Zinimermritin. ,1m 1^82, 213, 2S5). Melts at a bright
red heat (r'eli?»ot, A. Ch., iStiD, (4). 17. 3ftS). Found in vnrion8 minerols; its
ehief ore Is p»reh-bh*iide, which contaia^ from 40 to UO per cent of U,0, ,
Prepared by fusing UCl* with K or Na (Zimmerinann, A., iss:i, 216. i: ls80,
232* 27;^), It has the color of nickel, hard, but softer than steeU iiialloable,
permau€*iit In the air and watt*r at ordinury tempemtnreB; when ignited biirus
wiUi incandf&cenve to U,0,: unites direetly with CI , Br , I and S w hen heated;
soluble in HCl , HJSO^ and b lowly in HNO, . IrutmuJi nj^kU. VQ, , formed by
igtiiting' the higher oxides in carbon or bydrog-en, is a brown powder, £ooq
tnrtiinir yellow b^^ absorption of oxypen from the air. VfmitntH hudnhriiif' Is
formed by precipitating nranous salts with alkali j^, Vrttnw luidt, XIO, , Is
formed by heating- nrank nitrate cBUtionsly to 25^, and npon iVnition in the
|ir both this and other uraniani oxides^ hydroxides and araninta oxy salts with
'VolntHe ttcid^ are convened into UnO, ^ UO^SUOa . rrfialnm acts as a base in
two eb^Kes of saltw, ttrfUtnuA and nratiifl jvolts. rmnous siilts are green and give
|,'reen solutions, from which alkab's precipitiite nranous hydroxide, insohihle hi
excesjH of the ^dknli: alkali earlxmatcs precipitate ITfOff)* ♦ soluble in
<KS,hCO,; with B0CO3 the precipitation is* complete even in the cokb H3S is
wilhont action; (NH^^JS givcfe a dark-Virown precipitate: K4F&(CN)fl i^rives a
reddish-brown precij>ltate. In their action towiird oxidizing* and reducing'
Mprenffi nranous and nranyl (nranicj salts resemble closely ferrous and ferric
iiults: oranous salts are even more easily oxidized than lerrouss salts* r. 1?., by
exposure to the air, by WNO^. CI, HClOs . Br ♦ KMmO^ , etc. Gold, silver and
pliitinutii palts are rcdaeed to the free nictnl. The hexad nmniam (TJVi) acts
as a hnse. but nsiiiilly forniFi biL^ie siilts. never normal: we have U'OjNOf.lt ,
Tjot trfNOft! UO.JBO/, not UfSO j, . These basic salts were formerly ealled
nranic salts, but at present (UOj'' ik reg-arded as tk basic radical and called
nmtiitU and ita salte are called nranyl ^alta, e. |f „ ITO^fClj urariyl chloride,
(TTOj.tPO,), nranyl orthophosphatc. Solutions of nranyl salts are yellow:
KOH and NaOH |*ive a yellow prcei]ittate, nranates, ii.V ,Qr &nd Na^TI^Oj ,
in^^olnlde in cxcmsk. Alkali carbonates jrivc a yellow precipitate, soluble in
excess: BaCO^ and CaCO, prive UOn . HJ3 does not precipitate the Mranbitti,
hilt slowly rednees nranyl salts to uranous salts (Formanek. A., IKliO, 25t, 115)*
(KHJ.S Ifivcs a dark-brown precipitate. K.FeiCN'Jfl R-ives a rcddi^li-bro',vn
precipitate. Used in the analysis and sepamtjon of uranium compounds
rFresenlns and Hintsc, Z. fittfjetr.. 1^5, 502). Sodium phosphate i^ives a yellow
precipitate. The hexad nrinunm act:^ as an acid toward some stronger 'liases.
Thus we have KIT.O. and Na.-U'.OT , formed l>y precipitating nranyl salts with
KOH and NaOH: compare the similar salts of the hcxad chrominm, K,Cr,OT
and Na,Cr.OT . Other oxides of uraT>inm are described, but arc doubtless com-
hinatlon*4 of UO, and TJO, . Zti , Cd ♦ Sn , Pb . Co , Cu , Fe , and fcrrons salts
rediirc uranyl sjifts fo nrai>ot]s salts, Sokitions of Sn , Pt , Au , Cu . Hg' and
A|f are rcditeed to the metal by metallic ;iraninm (Zimmprmafin, !, c.b For
iiiethod of recovery of waste nrimium compounds, see Lnnbe (Z. fini/€ii\f 1SS%
TmS),
202 YTTERBIUM— YTTRWM—ZIRCOXIUM. §172.
§172. Ytterbium. Tb = 173.2 . Valence three.
Obtained as an earth by Marignac (C. r., 1878, 87, 578) from a gadolinite
earth; by Delafontaine (C. r., 1878, 87, 933) from sipylite found at Amherst, Va.
Nilson (B., 1879, 12, 550; 1880, 13, 1433) describes its preparation from euxenite
and its separation from Sc . It has the lowest bacisity of the yttrium earths.
The double potassium ytterbium sulphate is easily soluble in water and in
potassium sulphate. The oxalate forms a white crystalline precipitate, in-
soluble in water and in dilute acids. The salts are colorless and g^ve no
absorption spectrum. For the spark spectrum see Welsbach (Jf., 1884, 5, 1).
The oxide, Yb.O, , is a white powder, slowiy soluble in cold acids, read i|y u pon
warming". The hitdroxidc forms a pfelatinous precipitate, insoluble in XhBt«OH
but soluble in KOH . It absorbs COj from the air. The nitrate melts in its
water of crystallization and is very soluble in water.
§173. Yttrium. Y = 89.0 . Valence three.
Yttrium is one of the numerous rare metals found in the gadolinite mineral
at Ytterby, near Stockholm, Sweden: also found in Colorado (Hidden and
Mackintosh, Am. 5>'., 1889, 38. 474). The metal has been prepared by electro-
lysis of the chloride: also by heating the oxide, YjO, , with Mg (Winkler, B..
1890, 23, 787). The study of these rare earths is by no means complete. It is
also claimed that they have not yet been obtained pure, but that the so-called
pure oxides really consist of a mixture of oxides of from five to twenty ele-
ments (Crookes, V. A"., 1887, 55, 107, 119 and 131). The most of these rare
earths do not give an absorption spectrum, but give characteristic spark spectra:
and it is largely by this means that the supposedly pure oxides have l>een
shown to be mixtures of the oxides of several closelv related elements (Wels-
bach, J/., 1883, 4, fi41: Dennis and Chamot, ./. Am. »SVk\, 1897, 19. 799). Y^ttrium
salts are precipitated by the alkalis and by the alkali sulphides as the
hydroxklc, YCOH), , a white bulky precipitate, insoluble in the excess of the
reagents (distinction from Gl). The oxide and hiidroxide are readily soluble
in acids; boiling with NH4CI causes solution of the hydroxide as the chloride.
The alkali carbonates precipitate the carbonate Y2(C0,)s , soluble in a large
excess of the reagents. If the solution in ammcmium carbonate be boiled, the
hydroxide is precipitated. Soluble oxalates precipitate yttrium salts as the
white oxalate (distinction from Al and Gl): soluble with some difficulty in
HCl . The double sulphate with potassium is soluble in water and in potassium
sulphate (distinction from thorium, zirconium and the cerite metals). BaCO,
forms no precipitate in the cold (distinction from Al , Fe'" , Cr"' , Th , Ce .
La, Nd and Pr). Hydrofluoric acid precipitates the gelatinous fluoride, YF, ,
insoluble in water and in HF . The precipitation of yttrium salts is nor
hindered by the presence of tartaric acid (distinction from Al , Gl , Th and
Zr). The analysis of yttrium usually consists in its detection and separation
in gadolinite (silicate of Y , Gl , Fe , Mn , Ce and La). Fuse with alkali car-
bonate, decompose with HCl , and filter from the SiOa . Neutralize the filtrate
and precipitate the Y, La and Ce as oxalates with {'$rH.4).CnOt . Ignite the
precipitate and dissolve in HCl . Precipitate the La and Ce as the double
potassium sulphates, and from the filtrate precipitate the yttrium as the
hydroxide with NH4OH . Ignite and weigh as the oxide. In order to effect
complete separations the operations should be repeated several times.
§174. Zirconium. Zr = 90.4 . Valence four .
Zirconium is a rare metal found in various minerals, chiefly in zircon, a
silicate: never found native. The metal was first prepared by Berzelius in
1824 by fusion of the potassium zirconium fluoride with potassium (Popg.^ 1S25,
4, 117). Also prepared by electrolysis of the chloride (Becquerel, A. Ch., 1831,
48. 337). The metal exists in three modifications: crystalline, graphitoidal and
amorphous. The amorphous zirconium is a velvet-black pow^der, burning when
§175. TEE CALCIUM GROUP. 203
heated in the air. Acids attack it slowly even when hot, except EP , which
dissolves it in the cold. It forms but one oxide, ZrOj , analogous to SiOa and
TiOa . ZrOa is prepared from the mineral zircon by fusion with a fixed
alkali. Digestion in water removes the most of the silicate, leaving the
alkali zirconate as a sandy powder. Digestion with HCl precipitates the last of
the SiOj and dissolves the zirconate. The solution is neutralized, strongly
diluted and boiled; whereupon the zirconium precipitates as the basic chloride
free from iron. Or the zirconium may be precipitated by a saturated solution
of KjSO^ , and after resolution in acids precipitated by NH4OH and ignited
to ZrO, (Berlin, J, pr„ 1853, 58, 145; Roerdam., C. C, 1889, 533). ZrO^ is a white
infusible powder, giving out an intense white light when heated: it shows no
lines in the spectrum. It is much used with other rare earths, La^Oa , Y,0, ,
etc., to form the mantles used in the WeUbach gas-burners (Drossbach, C. C,
1891, 772; Welsbach, J., 1887, 2670; C. A'., 1887, 55, 192). The 0J:ide (or hydroxide
precipitated hot) dissolves with difficulty in acids to form salts. The hydroxide^
ZrO(OH)2 , precipitated in the cold dissolves readily in acids. As an acid^
zirconium hydroxide, ZrO(OH)2 = H^ZrO, , forms zirconates, decomposed by
acids. As a l)ase it forms zirconium salts with acids. The sulphate is easily
soluble in water, crystallizing from solution with 4H2O . The phosphate is
insoluble in water, formed by precipitation of zirconium salts by Na^HPO^ or
HsPO^ . The silicate, ZrOj.SiO. , is found in nature as the mineral zircon,
usually containing traces of iron. Zirconium chloride is formed when a current
of chlorine is passed over heated ZrO, , mixed with charcoal. It is a white
solid, may be sublimed, is soluble in water. vSolutions of zirconium salts are
precipitated as the hydroxide, ZrO(OH)2 , by alkali hydroxides and sulphides,
a white flocculent precipitate, insoluble in excess of the reagents, insoluble in
19^4 CI solution (difference from Gl). Tartaric acid prevents the precipitation.
Alkali carbonates precipitate basic zirconium carbonate, white, soluble in
excess of KHCOg or (NH4)2C03: boiling precipitates a gelatinous hydroxide
from the latter solution. BaCOj does not precipitate zirconium salts com-
pletely, even on boiling. The precipitates of the hydroxide and carbonate are
soluble in acids. Oxalic acid and oxalates precipitate zirconium oxalate, solu-
ble in excess of oxalic acid on warming, and soluble in the cold in (NH4)2C204
(difference from thorium); soluble in HCl. A saturated solution of K3SO4
precipitates the double potassium zitroiiium sulphate, white, insoluble in excess
of the reagent if precipitated cold, soluble in excess of HCl; if precipitated
hot, almost absolutely insoluble in water or HCl (distinction from Th and Ce).
Zirconium salts are precipitated on warming with NajSjO, (separation from
Y , Nd and Pr). Solution of H.O, completely precipitates zirconium salts.
Tumeric paper moistened with a solution of zirconium salt and HCl is colored
orange upon drying (boric acid gives the same reaction) (Brush, »/. pr,, 1854,
62, 7). HF does not precipitate zirconium solutions, as zirconium fluoride,
ZrF4 , is soluble in water and in HF (distinction from Th and Y).
The Calcium GRorr (Fifth Group).
(The Alkaline Earth Metals.)
Barium. Ba = 137.40 . Calcium. Ca = 40.1 .
Strontium. Sr — 87.60 . Magnesium. Mg = 24.3 .
§175. Like the alkali metals, Ba, Sr, and Ca oxidize rapidly in the air
at ordinary temperatures — ^forming alkaline earths — and decompose water,
forming hydroxides with evolution of heat. Mg oxidizes rapidly in the air
when ignited, decomposes water at 100°, and its oxide — in physical proper-
ties farther removed from Ba , Sr , and Ca than these oxides are from each
204 THE CALCIUM OROUP. §176.
other— slowly unites with water without sensible production of heat. As
compounds, these metals are not easily oxidized beyond their quantivalence
as dyads, and they require very strong reducing agents to restore them
to the elemental state.
§176. In basic power, Ba is the strongest of the four, Sr somewhat
stronger than Ca, and Mg much weaker than the other three. It will be
observed that the solubility of their hydroxides varies in the same decreas-
ing gradation, which is also that of their atomic weights; while the
solubility of their sulphates varies in a reverse order, as follows :
§177. The hydroxide of Ba dissolves in about 30 parts of water; that of
Sr, in 100 parts; of Ca, in 800 parts; and of Mg, in 100,000 parts. The
sulphate of Ba is not appreciably soluble in water (429,700 parts at 18.4°;
Hollemann, Z. phys. Ch., 1893, 12, 131); that of Sr dissolves in 10,000
parts ; of Ca , in 500 parts ; of Mg , in 3 parts. To the extent in which thev
dissolve in water, alkaline earths render their solutions caustic to the
taste and touch, and alkaline to test-papers and phenolphthalein.
§178. The carbonates of the alkaline earths are not entirely insoluble
in pure water: BaCOs is soluble in 45,566 parts at 24.2° (Hollemann,
ZeiL phys. Ch., 1893, 12, 125); SrCOg in 90,909 parts at 18° (Kohlrausch
and Rose, Zeit. phys, Ch,, 1893, 12, 241); CaCO., in 80,040 parts at 23.8°
(Hollemann, /. c); MgCO, in 9,434 parts (Chevaiet, Z., 1869, 8, 91). The
presence of NH^OH and (NH4)2C03 lessens the solubility of the carbonates
of Ba , Sr , and Ca , while their solubility is increased by the presence of
NH4CI . MgCOa is soluble in ammonium carbonate and in ammonium
chloride, so much so that in prcvsence of an abundance of the latter it is
not at all precipitated by the former, i. e, (UHjoCOg does not precipitate a
solution of MgClo as the NH^Cl formed holds the Mg in solution.
§179. These metals may be all precipitated as phosphates in presence
of ammonium salts, but their further separation for identification or esti-
mation would be attended with ditTiculty (§145 and ff.).
§180. The oxalates of Ba, Sr, and Mg nre sparingly soluble in water,
calcium oxalate insoluble. Barium chromate is insoluble in water (§§27
and 186, 5r), strontium chromate sparingly soluble, and calcium and mag-
nesium chromates freely soluble.
§181. In qualitative analysis, the group-separation of the fifth-group
metals is effected, after removal of the first four groups of bases, by
])recipitation with carbonate in presence of ammonium chloride, after
which magnesium is precipitated from the filtrate, as phosphate.
§182. The hydroxides of Ba, Sr, and Ca, in their saturated solutions^
necessarily dilute, precipitate solutions of salts of the metals of the first
four groups and of Mg , as hydroxides. In turn, the fixed alkalis precipi-
tate, from solutions of Ba , Sr , Ca , and Mg , so much of the hydroxides
§186, 4. BARIUM 20S
of these metals as does not dissolve in the water present *; but ammonium
hydroxide precipitates only Mg , and this but in part, owing to the solubility
of Hg(0H)2 in ammonium salts.
§183. Solutions containing Ba , Sr , Ca , and Mg , with phosphoric, oxalic,
boric, or arsenic acid, necessarily have the acid reaction, as occurs in dis-
solving phosphates, oxalates, etc., with acids; such solutions are precipi-
fated hy ammonium hydroxide or by any agent which neutralizes the solu-
tion, and, consequently, we have precipitates of this kind in the third
group (§145 and //.):
CaCl, + H,PO, + 2NH,0H = CaHPO« + 2NH4CI + 2H,0
CaH,(PO,), + 2NH4OH = CaHPO« + (NH4),HP0, + 2H,0 .
If excess of the ammonium hydroxide be added the precipitate is Ca8(P04)2.
In the case of a magnesium salt the precipitate is HgNH^FO^ .
§184. The carbonates of the alkaline earth metals are dissociated by
heat, leaving metallic oxides and carbonic anhydride. This occurs with
difficulty in the case of Ba .
§185. Compounds of Ba , Sr , and Ca (preferably with HCl) impart char-
acteristic colors to the non-luminous flame, and readily present well-defined
spectra.
§186. Barium. Ba = 137.40 . Valence two.
1. Properties.— ^pfci/?c gmvity, 3.75 (Kern, C. 3^., 1875, 31, 243) ; melting point,
above that of cast iron (Frey, A., 1876, 183, 368). It is a white metal, stable in
dry air, but readily oxidized in moist air or in water at ordinary temperature,
hydrog-en being evolved and barium hydroxide formed. It is malleable and
ductile (Kern, /. c).
2. Occurrence. — Barium can never occur in nature as the metal or oxide, or
hydroxide near the earth's surface, as the metal oxidizes so readily, and the
•xide and hydroxide are so basic, absorbing acids readily from the air. Its
roost common forms of occurrence are heavy spar, BaSO^ , and witherite,
BaCO. .
3. Preparation. — {1) By electrolysis of the chloride fused or moistened with
strong HCl . (2) By electrolysis of the carbonate, sulphate, etc., mixed with
Hg and Hg^ , and then distilling the amalgam. (3) By heating the oxide or
various salts with sodium or potassium and extracting the metal formed with
mercury, then separating by distillation of the amalgam.
4. Oxides and Hydroxides. — The oxide, BaO , is formed by the action of heat
upon the hydroxide, carbonate, nitrate, oxalate, and all its organic salts. The
corresponding hydroxide. BaCOH), , is made by treating the oxide with water.
The peroxide, BaO, , is made by heating the oxide almost to redness in oxygen,
or air which has been freed from carbon dioxide: by heating the oxide with
potassium chlorate (Liebig, Pogy., 1832, 26, 172) or cupric oxide (Wanklyn. B.,
1874, 7, 1029). It is used as a source of oxygen, which it gives off at a white
heat, BaO remaining; also in the manufacture of hydrogen peroxide, HaO, ,
which is formed by treating it with dilute acids: BaO, + 2HC1 = BaClj +
H,0,.
* The presence of an excess of fixed alkali renders these hydroxides much less soluble, the
hJgh coDO«>ntration of the hydroxyl ions, one of the factors of the solubility product, diminish-
ingr the other factor. (H^^)*
206 BARIUM, §186, 5a.
5. Solubilities. — a. — Metal. — Metallic barium is readily soluble in acids with
evolution of hydrogen. 6. — Oxides and hydroxides, — Barium oxide is acted upon
by water with evolution of heat and formation of the hydroxide, which is
soluble in about 30 parts of cold water and in its own weight of hot water
(Rosenstheil and Ruehlmann, J., 1870, 314). Barium peroxide, BaO, , is very
sparingly soluble in water (Schone, A., 1877, 192, 257); soluble in acids with
formation of H^Oj .
c, — Salts, — Most of the soluble salts of barium are permanent; the
acetate is efflorescent. The chloride, bromide, bromate, iodide, sulphide.,
ferrocyanide, nitrate, h37)ophosphite, chlorate, acetate, and phenylsul-
phate, are freely soluble in water; the carbonate, sulphate, sulphite,
cliromate* phosphite, phosphate, oxalate, iodate, and silico-fluoride, are
insoluble in water. The sulphate is perceptibly soluble in strong HCl.
The chloride is almost insoluble in strong hydrochloric acid (separation
from Ca and Mg) (Mar, Am, S,, 1892, 143, 521); likewise the nitrate in
strong hydrochloric and nitric acids. The chloride and nitrate are insolu-
ble in alcohol.
6. Eeactions. a, — The fixed alkali hydroxides precipitate only con-
centrated solutions of barium salts (5&). No precipitate is formed with
ammonium hydroxide (§45). The alkali carbonates precipitate barium
carbonate, BaCOg , white. The precipitation is promoted by heat and
by ammonium hydroxide, but is made slightly incomplete by the presence
of ammonium salts (Vogol, /. /?r., 183(), 7, 455).
Barium Carbonate — BaCOs — is a valuable reagent for special purposes,
chiefly for separation of third and fourth group metals. It is used in the
form of the moist precipitate, which must be thoroughly washed. It is
best precipitated from boiling solutions of barium chloride and sodium or
ammonium carbonate, washed once or t^nce by decantation, then by filtra-
tion, till the washings no longer precipitate solution of silver nitrate.
Mixed with water to consistence of cream, it may be preserved for some
time in stoppered bottles, being shaken whenever required for use. When
dissolved in hydrochloric acid, and fully precipitated by sulphuric acid,
the filtrate must yield no fixed residue. This reagent removes sulphuric
acid (radical) from all sulphates in solution to which it is added {e): NEsSO^
+ BaCOg = BaSO^ + NaoCOg . When salts of non-alkali metals are so
decomposed, of course, they are left insoluble, as carbonates or hydroxides,
nothing remaining in solution:
FeSOf -h BaCO, = BaS04 + FeCO,
Fe,(SOJ, -f SBaCO, + SH.O = 3BaS0« + 2Fe(0H), + 3C0,
The chlorides of the third group, except Fe" , are decomposed by barium
carbonate; while the metals of the fourth group (zinc, manganese, cobalt,
nickel), are not precipitated from their chlorides by this reagent. Tartaric
• Kohlrausch and Rose, Z. phyn, Ch„ 1803, 13* 241 : Schweitzer, Z., 1890, 39* 414.
§186, 7. BARIUM. 207
acid, citric acid, sugar, and other organic substances, hinder or prevent
the decomposition by barium carbonate.
b, — ^Aznmoniuin oxalate precipitates barium oxalate, BaCs04 , from solutions
of barium salts, sparingly soluble in water, more soluble in presence of am-
monium chloride; soluble in oxalic and acetic acids (Souchay and Lenssen, A.,
1856, 99, 36).
c. — Solutions of barium salts are precipitated by the addition of concentrated
nitric acid (5c). d, — Soluble phosphates, full metallic, or two-thirds metallic,
as NajHFOf , precipitate barium phosphate, white, consisting of BaHF04
when the reagent is two-thirds metallic, and Bas(F04)2 when the reagent is
full metallic. Soluble phosphites precipitate barium salts, hypophosphites do
not. e. — Barium sulphide is not formed in the wet way, hence hydrosulphuric
acid and soluble sulphides are without action upon barium salts. Soluble
sulphites precipitate solutions of barium salts as barium sulphite, BaSO, , in-
soluble in water but soluble in hydrochloric acid (distinction from sulphates).
Sulphuric acid, H2SO4, and all soluble sulphates, precipitate barium
sulphate (BaSO^), white, slightly soluble in hot concentrated sulphuric
acid. Immediate precipitation by the (dilute §188, 5c) saturated solution
of calcium sulphate distinguishes Ba from Sr (and of course from Ca); but
precipitation by the (very dilute §187, oc) solution of strontium sulphate
is a more certain test between Ba and Sr . BaSO^ is not transposed by
solutions of alkali carbonates (distinction from Sr and Ca , §188, 6a foot-
note).
f. — Solutions of iodates, as NalO, , precipitate, from barium solutions not
very dilute, barium iodate, Ba(I0,)2 , white, soluble in 600 parts of hot or
1746 parts of cold water (distinction from the other alkaline earth metals).
t;. — Neutral or nmmoniacal solutions of arsenous acid do not precipitate barium
salts (distinction from calcium). Soluble arsenates precipitate solutions of
barium salts, soluble in acids, including arsenic acid.
h, — Soluble chromates, as K^.CrO^ , precipitate solutions of barium salts
as barium chromate, BaCrO^ , yellow; almost insoluble in water (separa-
tion from calcium and from strontium except in concentrated solutions),
sparingly soluble in acetic acid, moderately soluble in chromic acid and
readily soluble in hydrochloric and nitric acids. Bichromates, as KaCrjO^ ,.
precipitate solutions of barium salts (better from the acetate) as the
normal chromate (very accurate separation from strontium and calcium)
(Grittner, Z, angew,, 1892, 73).
/. — Fluosilicic acid, HzSiP, ♦ precipitates white, crystalline barium fluo-
silicate, BaSiPe » slightly soluble in water (1-4000), not soluble in alcohol
(distinction from strontium and calcium). If an equal volume of alcohol be
added the precipitation is complete, sulphuric acid not giving a precipitate in
the filtrate (Fresenius, Z., 1890, 29, 143).
7. Ignition. — The volatile salts of barium as the chloride or nitrate impart a
yellowish-green color to the flame of the Bunsen burner, appearing blue when
viewed through a green glass. The spectrum of barium is readily distinguished
from the spectra of other metals by the green bands Baa, /? and y . Barium
carbonate is very stable when heated, requiring a very high heat to decompose
it into BaO and CO, .
208 STROXTIU^i §186, 8.
8. Detection. — In the filtrate from the fourth group, barium is precipi-
tated with strontium and calcium as the carbonate by ammonium car-
bonate. The white precipitate (well washed) is dissolved in acetic acid
and the barium precipitated with KJ^T^^i ^s BaCrO^ which separates it
from strontium and calcium. The barium is further identified by the
non-Holubility of the chromate in acetic acid, the solubility in hydrochloric
acid, and precipitation from this solution by sulphuric acid. It may also
be confirmed by the color of the flame with any of the volatile salts (7)
(not the sulphate).
9. EBtimation. — Barium 18 weighed as a sulphate (Fresenius and Hurtz, Z.
antfcw., ISOO, 2y.i), carbonate or fluosilicate (BaSiF«). It is separated from
Ktrontiiim and calcium: (/) By digesting the mixed sulphates at ordinary tem-
peratures for 12 hours with ammonium carbonate. The calcium and strontium
arc thus converted into carbonates, which are separated from the barium
tnilphate by dissolving in hydrochloric acid. (2) By hydrofluosilicic acid.
(.i) By repeated precipitation as the chromate in an acetate solution.
It is separated from calcium by the solution of the nitrate of the latter in
amyl alcohol (§188, 9). The hydroxide and carbonates are also determined by
alkalimetry. Volumetrically it is precipitated as the chromate, thoroughly
washed, dissolved in dilute *HC1 and the Crvi determined by HaO, (Baumann.
Z. anffPiP., 1891, 3:j1).
10. Oxidation. — Barium compounds are reduced to the metal when heated
with Na or K (3). BaO, oxidizes MnCl, to HnjO. (Spring and Lucion, BL,
1890, (3), 3, 4).
§187. Strontium. Sr = 87.60 . Valence two.
1. Properties.— *"^/>rr//?r gwnfy, 2.4 (Franz, J, pr., 1809, 107, 254). Melts at a
moderate red heat and is not volatile when heated to a full red. It is a " brass-
yellow ** metal, malleable and ductile. It oxidizes rapidly when exposed to
the air, and when heated in the air burns, as does barium, with intense
illumination (l«>anz. I.e.).
2. Occurrence. — Strontium occurs chiefly in strontianite, SrCO, , and in
celestlne. SrSO^ .
3. Preparation. — First isolated in 1808 by Davy by electrolysis of the hydrox-
ide (7';v/«.s. Roj/al Noc, 345). It is made by electrolysis of the chloride (Frey.
A., 187(), 183, 307): by heating a saturated solution of SrClj with sodium
amalgam and distilling off the mercury (Franz, /. c): by heating the oxide with
powdered magnesium the metal is obtained mixed with Mg^ (Winkler, B., 1890,
23, 12:.).
4. Oxides and Hydroxides. — Strontium oxide, SrO . is formed by igniting the
hydroxide, carl>onate (greater heat required than with calcium carbonate),
nitrate and all organic strontium salts. The hydroxide, Sr(0H)2 , is formed
by the action of water on the oxide. The peroxide, SrOj.SHjO , is made by pre-
cipitating the hydroxide with H..O5: at 100° this loses water and becomes SrO^ .
a white powder, melting at a red heat, used in bleaching works (Conroy.
J. Snr. Awf/., 1S92, 11, S12).
5. Solubilities, — a. — ]fvtal. — Strontium decomposes water at ordinary tem-
perature (Winkler, /. f.), it is soluble in acids with evolution of hydrogen.
h. — O.rif/r.f and hi/iirnxidra. — The oxide, SrO, is soluble in about 100 parts water
at ordinary temperature, and in al>out five parts of boiling water forming the
hydroxide* (Seheibler. Xriir Zeitftchrift fitr Ruehmzurler, 1R81, 49, 257). The
peroxitle is scarcely soluble in water or in ammonium hydroxide, soluble in
iicids and in ammonium chloride.
§187, G^. STRONTIUM. 200
c. — Salts, — The chloride is slightly deliquescent; crystals of the nitrate
and acetate effloresce. The chloride is soluble, the nitrate insoluble in
absolute alcohol. The nitrate is insoluble in boiling amyl alcohol (§188,
oc). The sulphate is very sparingly soluble in water (1-10,090 at 20.1°)
(HoHemann, Z. phys, Ch., 1893, 12, 131); yet sufficiently soluble to allow
its use as a reagent to detect the presence of traces of barium. Less soluble
in water containing ammonium salts, sodium sulphate, or sulphuric acid
than in pure water; quite appreciably soluble in HCl or HNO3 ; insoluble
in alcohol. Strontium fluosilicate is soluble in water (distinction from
barium). The chromate is soluble in 831.8 parts water at 15° (Fresenius,
Z., 1890, 29, 419); soluble in many acids including chromic acid; and more
soluble in water containing ammonium salts than in pure water.
6. Ecactions. a, — The fixed alkalis precipitate strontium salts when
not too dilute, as the hydroxide, Sr(0H)2 , less soluble than the barium
hydroxide. No precipitate with ammoniuin hydroxide. The alkali car-
bonates precipitate solutions of strontium salts as the carbonate. Stron-
tium sulphate is completely transposed on boiling with a fixed alkali car-
bonate (distinction from barium, §188, 6a footnote).
ft. — Oxalic acid and oxalates precipitate strontium oxalate, insoluble in
water, soluble in hydrochloric acid (Souchay and Lenssen, A., 1857, 102, 35).
e. — The solubility of strontium salts is diminished by the presence of con-
centrated nitric acid, but less so than barium salts, d. — In deportment with
phosphates, strontium is not to be distinguished from barium.
e. — See Ge, §§186 and 188. Sulphuric acid and sulphates (including
CaS04) precipitate solutions of strontium salts as the sulphate, unless
the solution is diluted beyond the limit of the solubility of the precipitate
(or). A solution of strontium sulphate is used to detect the presence of
traces of barium (distinction from strontium and calcium). In dilute
solutions the precipitate of strontium sulphate forms very slowly, aided
by boiling or by the presence of alcohol, prevented by the presence of
hydrochloric or nitric acids (5r). It is almost insoluble in a solution of
ammonium sulphate (separation from calcium).
f. — The halldes of strontium are all soluble in water and have no application
in the analysis of strontium salts. Strong hydrochloric acid dissolves stron-
tium sulphate, but in general diminishes the solubility of strontium salts in
water, g. — Neutral solutions of arsenltes do not precipitate strontium salts,
the addition of ammonium hydroxide causes a precipitation of a portion of the ,
strontium. Arsenate of strontium resembles the corresponding barium salt.!
Alkaline arsenates do not precipitate strontium from solution of the sulphate
(distinction from calcium, §188, 6^7).
h, — Xormal cliroinates precipitate strontium chromate from solutions
not too dilute (5c), soluble in acids. In absence of barium, strontium
may be separated from calcium by adding to the nearly neutral solutionj^
a solution of K^CrO^ plus one-third volume of alcohol. The calcium
210 CALCIUM, §187, 6i.
chromate is about 100 times as soluble as the strontium ehroraiate (Fre-
senius and Rubbert, Z., 1891, 30, 672). No precipitate is formed with
potassium bichromate (separation from barium).
i. — ^Fluosilicic acid does not precipitate strontium salts even from quite
concentrated solutions, as the strontium fluosilicate is fairly soluble in cold
water and more so in the presence of hydrochloric acid (Fresenius, Z., 1890,
29, 143).
7. Ignltioii. — Volatile strontium compounds color the flame crimson. In pres-
ence of barium the crimson color appears at the moment when the substance
(moistened with hydrochloric acid, if a non-volatile compound) is first brought
into the flame. The paler, yellowish-red flame of calcium is liable to be mis-
taken for the strontium flame. The spectrum of strontium is characterized
by eight bright bands; namely, six red, one orange and one blue. The orange
line Sr a, at the red end of the spectrum; the two red lines, Sr /3 and Sr r,
and the blue line, Sr 6 , are the most important.
8. Detection. — Strontium is precipitated with barium and calcium from
the filtrate of the fourth group by ammonium carbonate. The well washed
precipitate of the carbonates is dissolved in acetic acid and the barium
removed by KoCFsO. . The strontium and calcium are separated from the
excess of chromate by reprecipitation with (H 114)^003 . The precipitate is
again dissolved in HC^HsO^ and from a portion of the solution the stron-
tium is detected by a solution of CaSO^ (6e). The flame test (7) is of value
in the identification of strontium.
y. E8tim.ation. — Strontium is weighed as a sulphate or a carbonate. The
hydroxide and carbonate may be determined by alkalimetry. It is separated
from calcium: (/) By the insolubility of its sulphate in ammonium sulphate.
(2) By boiling the nitrates with amyl alcohol (§188, 9). {S) By treating the
nitrates with equal volume of absolute alcohol and ether (§188, 9). For
separation from barium see §186, 9.
§188. Calcium. Ca = 40.1 . Valence two.
1. TTOTpertieB.— Specific iimrity, 1.6 to 1.8 (Caron, C. r., 1800, 50, 547). Melting
p(Ant^ at red heat (Matthiessen, A., 1855, 93, 284). A white metal having very
much the appearance of aluminum, is neither ductile nor malleable (Frey, A.,
1876, 183, 367). In dry air it is quite stable, in moist air it burns with
incandescence, as it does also with the halogens. It dissolves in mercury, form-
ing an amalgam.
2. Occurrence. — Found in the mineral kingdom as a carbonate in marble,
limestone, chalk and arragonite: as a sulphate in gypsum, selenite, alabaster,
etc.; as a fluoride in fluor-spar; as a phosphate in apatite, phosphorite, etc.
It is found as a phosphate in bones; in egg-shells and oyster-shells as a car-
bonate. It is found in nearly all spring and river waters.
3. Preparation. — (/) By ignition of the iodide with sodium in closed retorts
(Dumas, C. r., 1858, 47, 575). (2) By fusion of a mixture of 300 parts fused
CaClj , 400 i)arts granulated zinc and 100 parts Na until zinc vapor is given
off. From the CaZn alloy thus obtained the zinc is removed by distillation in
a graphite crucible (Caron, /. c). (3) By electrolysis of the chloride (Frey, I. r.).
(4) Bv reducing the oxide, hydroxide or carbonate with magnesium (Winkler.
B., 1800, 23, 122 and 2642). ' ^
4. Oxides and Hydroxides. — The oxide, CaO , is a strong base, non-fusible,
non-volatile; it is formed by oxidation of the metal in air; by ignition of the
§188, 5c. CALCIUM. 211
hydroxide, the carbonate (limestone), nitrate, and all organic calcium salts.
The corresponding hydroxide, Ca(OH), (slaked lime), is made by treating the
oxide with wate^. Its usefulness when combined with sand, making mortar,
is too well known to need any description here. The peroxide, Ca02.8H30 , is
made by adding hj'drogen peroxide or sodium peroxide to the hydroxide;
Ca(OH), -f H2O, = CaO, + 2H,0 (Conroy, J. Soc, Ind., 1892, 11, 808)^ Drying
at 130° removes all the water, leaving a white powder, CaOj , which at a red
heat loses h^lf its oxygen (Schoene, i.., 1877, 192, 257). It cannot be made by
heating the oxide in oxygen or with potassium chlorate (§186, 4).
5. Solubilities. — a, — Metal, — Calcium is soluble in acids with evolution of
hydrogen; it decomposes water, evolving hydrogen and forming Ca(OH)s .
h. — Oxide and hydroxide. — CaO combines with dilute acids forming cor-
responding salts, it absorbs COg from the air becoming CaCOs .* In moist
air it becomes Ca(0H)2 , the reaction takes place rapidly and with increase
of volume and generation of much heat in presence of abundance of
water. The hydroxide, Ca(OH).j , is soluble in acids, being capable of
titration with standard acids. It is much less soluble in water than
barium or strontium hydroxides (Lamy, C. r., 1878, 86, 333); in 806 parts
at 19.5° (Paresi and Rotondi, B., 1874, 7, 817); and in 1712 parts at 100°
(Lamy, /. c). The solubility decreases with increase of temperature. In
saturated solutions one part of the oxide is found in 744 parts of water
at 15° (Lamy, Z. c). A clear solution of the hydroxide in water is lime
water (absorbs COj forming CaCO..), the hydroxide in suspension to a
greater or less creamy consistency is milk of lime.
c. — Salts. — The chloride, bromide, iodide, nitrate, and chlorate are
deliquescent; the acetate is efflorescent.
The carbonate, oxalate, and phosphate are insoluble in water. The
chloride, iodide, and nitrate are soluble in alcohol. The nitrate is soluble
in 1.87 parts of equal volumes of ether and alcohol (Fresenius, Z., 1893,
32, 191); readily soluble in "boiling amyl alcohol (Browning, Am. S., 1892,
143, 53 and 314) (separation from barium and strontium). The carbonate
is soluble in water saturated with carbonic acid (as also are barium, stron-
tium, and magnesium carbonates), giving hardness to water. The oxalate
is insoluble in acetic acid, soluble in hydrochloric and nitric acids. The
sulphate is soluble in about 500 parts of water f at ordinary temperature,
the solubility not varying much in hot water until above 100° when the
solubility rapidly decreases. Its solubility in most alkali salts is greater
than in pure water. iVmmonium sulphate (1-4) requires 287 parts for the
solution of one part of CaSO^ (Fresenius, Z., 1891, 30, 593) (separation
from Ba and Sr). Readily soluble in a solution of NasSgOa (separation
from barium sulphate) (Diehl, /. pr., 1860, 79, 430). It is sy)luble in GO
parts hydrochloric acid, G. 12 per cent at 25°, and in 21 parts of the same
* Dry C«0 does not absorb dry CO, or SO, beJow 850". (Veley, J. C, 1893, 6S, 821).
r Goldhammer, C. C, 1888, 708; Droeze, B., 18T7. 10, 330; Bolsbaudran, A. Hi.. 1874. (5), S, 477
Kohlrausch and Rose, Z. phyn. CTi., 1893, 12, 241 ; Raupenstrauch. 3f., 1886, 6« 668).
212 CALCIUM, §188, (ya,
acid at 103° (Lunge, /. Soc. Ind.y 1895 14, 31). The chromate is soluble
in 214.3 parts water at 14° (Siewert, /., 1862, 149); in dilute alcohol it is
rather more soluble (Fresenius, I. c, page 672); very readily soluble in
acids including chromic acid.
6. Eeactions. a. — The fixed alkali hydroxides precipitate solutions of
calcium salts not having a degree of dilution beyond the solubility of the
calcium hydroxide formed (5ft), t. e, potassium hydroxide will form a
precipitate with calcium sulphate since the sulphate requires less water
for its solution than the hydroxide (5& and c) ; also the calcium hydroxide
is less soluble in the alkaline solution than in pure water. Ammoniiun
hydroxide does not precipitate calcium salts. The alkali carbonates pre-
cipitate calcium carbonate, CaCOs , insoluble in water free from carbon
dioxide, decomposed by acids. Calcium sulphate is completely trans-
posed upon digestion with an alkali carbonate * (distinction from barium).
Calcium hydroxide, Ca(0H)2 , is used as a reagent for the detection of
carbon dioxide (5& and §228, 8).
h. — Alkali oxalates, as (NH4)2Co04 , precipitate cakium oxalate, CaCoO^ .
from even dilute solutions of calcium salts. The precipitate is scarcely at
all soluble in aceiic or oxalic acids (separation of oxalic from phosphoric
acid (§315), but is soluble in hydrochhric and nitric acids. The pre-
cipitation is hastened by presence of ammonium hydroxide. Formed
slowly, from very dilute solutions, the precipitate is crystalline, octahedral.
If Sr or Ba are possibly present in the solution to be tested (qualitatively),
an alkali sulphate must first be added, and after (li(jestin(j a few minutes,
if a precipitate appears, SrSO^ , BaS04 , or, if the solution was concentrated,
perhaps CaSO^ , it is filtered out, and the oxalate then added to the filtrate.
If a mixture of the salts of barium, strontium, and calcium in nf?utral or
alkaline solution be treated with a mixture of (NH4)2S04 and (NH J2C2O4 ,
the barium and strontium are precipitated as sulphates and the calcium as
the oxalate; separated from the barium and strontium on addition of
hydrochloric acid (Sidersky, Z., 1883, 22, 10; Bozomoletz, Z?., 1884, 17,
1058). A solution of calcium chloride is used as a reagent for the detec-
tion of oxalic acid (§227, 8).
In solutions of calcium salts containing a strong" excess of ammonium
chloride, potassium ferrocyanide precipitates the calcium (distinction from
barium and strontium) (Baubigny, Bl, 1895, (3), 13, 326).
• Here experiment shows that for equilibrium the SO4 Ions must be present in solution In large
excess of CO, ions. With strontium also an excess of SO4 Ions Is required, although not so
£rreat as In the case of calcium. For barium, however, equilibrium demands that the oonoen-
tratlon of CO3 Ions exceed that of 8O4. This condition Is already fulfilled when an alkali car-
bonate Is added to BaS04 and therefore no changrc takes place In this case, while in the others
the sulphate is transformed Into carbonate. It Is important to notice that the relative or ab-
solute quantities of solid carbonate and sulphate present do not affect the equilibrium, which
Is determined solely by the substances in solution (§57, 6e, footnote).
§188, 9. CALCIUM. 213
c. — See 5c, d, — By the action of alkali phosphates, solutions of calcium are
not distinguished from solutions of barium or strontium.
e. — Pure sodium sulphide, Na^S , gives an abundant precipitate with calcium
salts; even with CaSO^ . The precipitate is CaCOH),: CaClz + aNa^S + 2TLJ0 =
CaCOBf). + 2NaCl + 2NaHS . The acid sulphide, NaHS , does not precipitate
calcium salts (Pelouze, A, Ch., 1866, (4), 7, 172). Alkali sulphites precipitate
calcium sulphite, nearly insoluble in water, soluble in hydrochloric, nitric or
sulphurous acid; barium and strontium salts act similarly.
Sulphuric acid and soluble sulphates precipitate calcium salts as CaSO^ ,'
distinguished from barium by its solubility in water and in hydrochloric;
acid; from barium and strontium by its solubility in ammonium sulphate
{5c). A water solution of calcium sulphate is used to detect strontium
after barium has been removed as a chromate. Obviously a solution of
strontium sulphate will not precipitate calcium salts.
f. — Calcium chloride, fused, is much used as a drying agent for solids, liquids
and gases. Chlorinated lime, calcium hypochlorite, Ca(C10)j (Kingzett, J. C,
1875, 28, 404), is much used as a bleaching agent and as a disinfectant, g. —
Neutral or ammoniacal solutions of arsenites form a precipitate with calcium
salts (distinction from barium). A solution of calcium salts including solu-
tions of calcium sulphate in ammoniacal solution is precipitated by arsenic
Acid as CaNH4A804 (distinction from strontium after the addition of sulphuric
acid) (Bloxam, C. N., 1886, 54, 16).
h. — Normal chrom.ates, as K2Cr04 , precipitate solutions of calcium salts as
<?alcium chromate, CaCr04 , yellow, provided the Folution be not too dilute (5c).
The precipitate is readily soluble in acids and is not formed with acid chro-
mates as K^CrxOr (separation from barium), i. — Fluosilicic acid does not
precipitate calcium salts even in the presence of equal parts of alcohol (separa-
tion from barium).
7. Ignition. — Calcium sulphate, CaS04.2H20 , gypsum , loses its water of
crystallization at 80° and becomes the anhydrous sulphate, CaSOf , plaster of
Paris; which on being moistened forms the crystalline CaS04.2H20 , expands
and " sets." Calcium carbonate, Umcstoney when heated (burned) loses carbon
dioxide and becomes lime, CaO .
Compounds of calcium, preferably the chloride, render the flam.e yellowish
red. The presence of strontium or barium obscures this reaction, but a mixture
containing calcium and barium, moistened with hydrochloric acid, gives the
calcium color on its first introduction to the flame. The spectrum of calcium
is distinguished by the bright green line, Ca /?, and the intensely bright
orange line. Ca a, near the red end of the spectrum.
8. Detection. — Calcium is separated in analysis from the metals of the
other groups and from barium, with strontium, as described at §187, 8.
A portion of the solution of strontium and calcium acetate is boiled with
potassium sulphate; after standing for some time (ten minutes), the filtrate
is tested with ammonium oxalate. A white precipitate insoluble in the
acetic acid present, but soluble in hydrochloric acid is evidence of th(^
presence of calcium. The flame test (7) is confirmatory.
9. Estimation. — Calcium is weighed as an oxide, carbonate, or sulphate. The
carbonate is obtained by precipitating as oxalate, and gently igniting the dried
precipitate; higher ignition changes the carbonate to the oxide. The sulphate
is precipitated in a mixture of two parts of alcohol to one of the solution. The
hydroxide and carbonate may be determined by alkalimetry. Calcium may be
^separated from barium and strontium by the solution of its nitrate in amyl
214 MAOyESIVM. §189, 1.
alcohol (5c). The best method of separation from strontium is to treat the
nitrates with a mixture of equal volumes of alcohol and ether. The calcium
nitrate dissolves, but not more than one part in 60,000 of the strontium is
found in the solution ($195). In the presence of iron, aluminum and phos-
phoric acid, calcium is best precipitated as an oxalate in the presence of citric
acid (Passon, Z. angeto,, 1898, 776). See also 9, §186 and §187.
§189. Magnesium. Mg = 24.3 . Valence two.
1. Properties.— -Spfci/fr gravity, 1.75 (Deville and Caron, A. Ch., 1863, (3), 07,
346); melting poitU, a little belov^ 800**, does not appear to be volatile (Meyer,
B., 1887, 20, 497). A white, hard, malleable and ductile metal; not acted upon
by water or alkalis at ordinary temperature and only slig^htly at 100® (Ballo,
B,, 1883, 16, 694). When heated in air or in oxypfen it burns with incandescence
to Mg^ . It combines directly when heated in contact with N , P , As , S
and CI . It forms alloys with Hg^ and Sn , forming compounds which decom-
pose water.
2. Occurrence. — Mapfnesite, MgCO.; dolomite, GaMgCCO,)^; brucite, Mg^(OH)j;
epsom salts, Mg'S04.7HsO; and combined with other metals in a g^reat variety
of minerals.
3. Preparation. — (/) By electrolysis of the chloride or sulphate (Bunsen, A.,
1852, 82, 137). (2) By ignition of the chloride with sodium or potassium
(Wohler, .4., 1857, 101. 562). (3) Mg,Fe(CN)e is ignited with Na.GO, , and
this product ignited with zinc (Lanterbronn, German Patent No. 39,915).
4. Oxide and Hydroxide. — Only one oxide of magnesium, Mg^ , is known
with certainty. Formed by burning the metal in the air, and by action of
heat upon the hydroxide, carbonate, nitrate, sulphate, oxalate and other mag-
nesium salts decomposed by heat. The corresponding hydroxide, Mg(OH), ,
is formed by precipitating magnesium salts with the fixed alkalis.
5. Solnbilities. — a. — Metal. — Magnesium is soluble in acids including
carbonic acid, evolving hydrogen: Mg + COo + HoO = MgCO^ -f S-
(Ballo, B,, 1882, 15, 3003): it is also attacked by the acid alkali carbonates,
as NaHCOg , to form MgCO., , NaoCO., and H (Ballo, I. r.). Soluble in
ammonium salts: Mg + 3NH,C1 = NH.MgClg + 2NH, +. H, . With
the halogens it acts tardily (Wanklyn and Chapman, J. C, ISGG, 19, 141).
b, — Oxide and hydroxide. — Insoluble in water, soluble in acids. Mg(OH)>
is soluble in 111,111 parts of water at 18° (Kohlrausch an<l Rose, Zeit,
pJiys, Ch,, 1893, 12, 241). In contact with water the oxide is slowly
changed to the hydroxide, Mg(0H)2 , and absorbs COo from the air. Sol-
uble in ammonium salts:* Mg(0H)2 + 3NH^C1 = NH^MgClj +
2NH4OH. c. — Salts, — The chloride, bromide, iodide, chlorate, nitrate,
and acetate (4 aq) are deliquescent \ the sulphate (7 aq) slightly efflorescent.
The carbonate, phosphate, borate, arsenite, and arsenate are insoluble in
water; the sulphite, oxalate, and chromate soluble; the tartrate sparingly
soluble. The carbonate is soluble; the phosphate, arsenite, and arsenate
are insoluble in excess of ammonium salts.
6. Reactions, a, — The fixed alkali hydroxides and the hydroxides of
barium, strontium and calcium precipitate magnesium hydroxide, Mg(0H)2^
• The conditions here are the same as In the case of Mn(OH)s, %\^A^ 6a, footnote.
§189, 7. MAGNESIUM, 215
white, gelatinous, from solutions of magnesium salts; insoluble in excess
of the reagent but readily soluble in ammonium salts : Mg(0H)2 + SNH^Cl
= MgCL^.NH4Cl + 2NH4OH . With ammonium hydroxide but half of the
magnesium is precipitated, the remainder being held in solution by the
ammonium salt formed in the reaction: 21l[gS04 + 2NH4OH = Mg(0H)2
+ (NHj2Mg(S0j2 (Rheineck, Dingl, 1871, 202, 268). The fixed
alkali carbonates precipitate basic magnesium carbonate, Mg4(0H)2-
(003)3 , variable to 1ILg^(0R)^{C0^\ : 4MgS04 + 4Na2C08 + H^O =
lIg4(0H)2(C03)3 + 4Na2S04 + CO2 . If the above reaction takes place in
the cold the carbon dioxide combines with a portion of the magnesium
<?arbonate to form a soluble acid magnesium carbonate: SMgSO^ +
oNajCOg + 2H2O = Mg4(0H)2(C03)3 + MgH2(C03)2 + SNagSO^. On
boiling, the acid carbonate is decomposed with escape of COo . Ammonium
<:arbonate does not precipitate magnesium salts, as a soluble double salt is
at once formed. Acid fixed alkali carbonates, as NaHCOs , do not precipi-
tate magnesium salts in the cold ; but upon boiling, CO2 is evolved and the
carbonate is precipitated (Engel, A. Ch,, 1886, (6), 7, 260).
h. — Soluble oxalates do nt)t precipitate solutions of magnesium salts, as they
form soluble double oxalates. If to the solution of double oxalates, preferably
magTiesium ammonium oxalate, an equal volume of 80 per cent acetic acid be
■added, the magnesium is precipitated as the oxalate (separation from potas-
sium or sodium (Classen, Z„ 1879, 18, 373).
d. — Alkali phosphates — as Na2HF04 — precipitate magnesium phosphate,
MgHFO^ , if the solution be not very dilute. But even in very dilute
solutions, by the further addition of ammonium hydroxide (and NH^Cl),
a crystalline precipitate is slowly formed, magnesium ammonium phosphate
— HLgSt&^O^ . Stirring with a glass rod against the side of the test-tube
promotes the precipitation. The addition of ammonium chloride, in this
test, prevents formation of any precipitate of magnesium hydroxide (5&).
The precipitate dissolves in 13,497 parts of water at 23° (Ebermayer,
J. pr., 1853, 60, 41); almost absolutely insoluble in water containing^
ammonium hydroxide and ammonium chloride (Kubel, Z,, 1869, 8, 125).
e. — Magnesium sulphide is decomposed by water, and magnesium salts are
not precipitated .by hydrosulphuric acid or ammoniuin sulphide; but Mg^ -f
HjO (1-10) absorbs H2S , forming in solution MgHsSz , which readily gives
off HJS upon boiling (a very satisfactory method of preparing H^S absolutely
arsenic free) (Divers and Shmidzu, J, C, 1884, 45, 699). Normal sodium or
potassium sulphide precipitates solutions of magnesium salts as the hydroxide
with formation of an acid alkali sulphide: MgS04 -\- 2Na.S -\- 2HoO = MgCOH).
-f Na,S04 + 2NaHS (Pelouze, A. Ch., 1866, (4), 7, 172). Sulphuric acid and
soluble sulphates do not precipitate solutions of magnesium salts (distinction
from Ba , Sr and Ca).
A — Magnesium chloride, in solution, evaporated on the water bath evolves
hydrochloric acid (7). g, — Soluble arsenates precipitate magnesium salts in
deportment similar to the corresponding phosphates.
7. Ignition. — Magnesium ammonium phosphate when ignited loses ammonia
216 MAOyESIlM, §189, 8.
and water, and becomes the pyrophosphate: 2Mg^NH4P04 = Mg,P,OT -+- H,0 -+-
2NH3 . The carbonate loses CO, and becomes Mg^ . In dry air mag^nesium
chloride may be ignited without decomposition, but in the presence of steam
MgO and HCl are formed: MgCl, -f HjO = MgO + 2HC1; a technical method
for preparing HCl (Heumann, A., 1877, 184, 227).
8. Detection. — If sufficient ammonium salts have been used, the mag-
nesium will be in the filtrate from the precipitated carbonates of barium,
strontium and calcium. From a portion of this filtrate the magnesium is
precipitated as the white magnesium ammonium-phosphate, 'ilLgSILJfO^ ,,
by NajHPO^ .
9. Estimation. — After removal of other non-alkali metals, magnesium is pre-
cipitated as MgNH4P04 , then changed by ignition to MgaPaOr (magnesium
pyrophosphate) and weighed as such. Separated as MgCl, from KCl and KaCl
by solution in amyl alcohol, evaporated with H3SO4 and weighed as MgSO^
(Riggs, Am. aS., 1892, 44, 103). It is estimated volumetrically by precipitation
as MgNH4P04 , drying at about 50** until all free NH4OH is removed. An
excess of standard acid is then added and at once titrated back with standard
fixed alkali, using methyl orange as an indicator (Handy, J. Am. iS*oc., 1900, 22,
31).
10. Oxidation. — Magnesium is a powerful reducer; ignited with the
oxides or carbonates of the following elements magnesium oxide is formed
and the corresponding element is liberated : Ag , Hg , Pt , Sn *, B , Al ,
Th, Ct, Si, Pb, Pt, As, Sb, Bi, Cr, Mo, Mn, Fe, Co, Ni, Cu,
Cd , Zn , Gl , Ba , Sr , Ca , Rb , K , Na , and Li . In some cases the reaction
takes place \nth explosive violence. From their corresponding salts in
neutral solution Mg precipitates Se , Te , As , Sb , Bi , Sn , Zn f , Cd , Pb,
Tl , Th , Cu , Ag , Mn f, Fe f, Co , Ni , An , Pt , and Pd (Scheibler, B.,
1870, 3, 295; Villiers and Borg, C. n, 1893, 116, 1524).
♦ Winkler, B., 1800, 23, 44, 120 and 773 ; 1801, 24, 802.
t Kern, C. N , 1876, 83, 112 and 236.
t Seubert and Schmidt, A., 1802, 267, 218.
«o
AXALT8I8 OF THE CALCIUM OROVP.
217
218 DIRECTIONS FOR ANALYSIS WITH NOTES. §191.
Directions fob Analysis of the Metals of the Calcium Gboup.
(The Alkaline Earths.)
§191. Manipulation. — To the filtrate from the fourth group in which
HgS (§192, i) gives no precipitate (§138) add NH^OH and ammonium
carbonate as long as a precipitate is formed : BaClj + (^^^4)2^08 = BaCO,
+ 2NH4CI . Digest with warming, filter and wash. The filtrate should
be tested again with ammonium carbonate and if no precipitate is formed
it is set aside to be tested for magnesium and the alkali metals (§§193
and 211).
The well washed white precipitate is dissolved in acetic acid, using as
little as possible: SrCOg + 2HC2H3O2 = Sr(C2H302)2 + CO2 + HjO .
To a small portion of the acetic acid solution add a drop of K^Ct^Oj ;
if a precipitate — BaCrO^ — is obtained, the K^CToOj must be added to the
whole solution : 2Ba(C2H302) + K^Ct^O^ + IL^O = 2BaCr04 + 2KC2H3O2
+ 2HC2H3O2 . Filter, wash the precipitate, dissolve it in HCl and pre-
cipitate the barium as barium sulphate, with a drop of sulphuric acid.
To the filtrate from the barium chromate add NH4OH and (NHJ2CO, ,
warm, filter, and wash. Dissolve the white precipitates of SrCO, and
CaC03 in acetic acid and divide the solution into two portions.
Portion 1. — For Strontium. — With a platinum wire obtain the flame
test, crimson for strontium; calcium interferes (7, §§187, 188 and 205).
Add a solution of calcium sulphate and boil ; set aside for about ten min-
utes. A precipitate — SrSO^ — indicates strontium. This SrSO^ may be
moistened with HCl and the crimson flame test obtained.
Portion 2. — For Calcium, — Add a solution of potassium sulphate, boil,
and set aside for ten minutes. Filter (to remove any strontium that may
be present; also a portion of the calcium may be precipitated, §188, 6e)
and add ammonium oxalate to the filtrate. Dissolve the precipitate in
HCl . A white precipitate — CaCoO^ — insoluble in acetic acid by its forma-
tion in that solution, and soluble in HCl is proof of the presence of calcium.
§102. Note8.—]. The failure of (NHJ.S (or HjS in presence of NH«OH) to
form a precipitate with solutions of the alkaline earths and of the alkalis,
marks a sharp separation of these metals from the metals of the preceding
groups.
2. Do not boil after the addition of ammonium carbonate, as this will drive
off ammonium hydroxide and carbonate, increasing the solubility of the CaCO,
(note 3 and §178).
3. The precipitation of barium, strontium and calcium by ammonium car-
bonate in the presence of ammonium chloride, is not as complete as would be
desirable in very delicate anal^'ses. The carbonates of barium, strontium and
calcium are all slightly soluble in ammonium chloride solution; and w^hile the
prescribed addition of ammonium hydroxide, and excess of ammonium car-
bonate, greatly reduces the solubility of the precipitated carbonates, yet even
with these the precipitation is not absolute, though more nearly so with
strontium than with barium and calcium. Thus, in quantitative analyses, if
§194,^. DIRECTIONS FOR ANALYSIS WITH NOTES. 219
barium and calcium are precipitated as carbonates, it must be done in the
absence of ammonium chloride or sulphate, and the precipitate washed with
water containing ammonium hydroxide.
4. If barium be absent, as evidenced by the failure to obtain a precipitate
w^ith KsCrxOr , the solution may at once be divided into two portions to test for
strontium and calcium.
5. With care the reprecipitation by ammonium carbonate, for the separa-
tion from the excess of K^Ct^O, , may be neglected and the filtrate from the
barium, yellow, at once divided into two portions and tested for Sr and Ca .
Reprecipitation always causes the loss of some of the metals, due to the solu-
bility of the carbonates in the ammonium acetate formed. On the other hand,
traces may escape observation in the yellow chromate solution.
6. Before reprecipitation with (NH4)2C08 , an excess of ammonium hydroxide
should be added to prevent the liberation of COj when the ammonium car-
bonate is added.
7. Strontium sulphate is so sparingly soluble in water (§187, 5c) that its
precipitation by GaS04 (or other sulphates in absence of Ca) is sufficiently
delicate to detect very small amounts of that metal. However, it is sufficiently
soluble in water to serve as a valuable reagent to detect the presence of traces
of barium. Obviously SrS04 will not precipitate solutions of calcium salts.
Solutions of strontium and barium salts (except SrS04) are all precipitated
by CaS04. The presence of excess of calcium salts lessens the delicacy of the
precipitation of strontium salts by calcium sulphate.
8. — In very dilute solutions the eulphates of the alkaline earths are not
precipitated rapidly. Time should be allowed for the complete precipitation.
Boiling and evaporation facilitates the reaction.
9, It should be noticed that the test for calcium as an oxalate is made upon
that portion of the calcium not removed by K3SO4: or in other words upon a
solution of GaSO* (1-500). A solution of SrSOf (1-10,000) may be present but
i s no t precipitated by (NH4)2C304 . The presence of a great excess of
(NH4)sS04 prevents the precipitation of traces of calcium salts by (NH4)2C,04 .
§193. Manipulation. — To a portion of the filtrate from the carbonates
of Ba , Sr , and Ca add a drop or two of (NH4)2S04 and then a few drops of
(NH4)2C204 ; filter if a precipitate is obtained and test the filtrate for Mg
with NagHPO^ . A white precipitate — llLgI^K^FO^ — is evidence of the
presence of magnesium. The other portion of the filtrate from the car-
bonates of Ba , Sr , and Ca is reserved to be tested for the alkali metals
(§211)-
§194. Notes, — i. By some, magnesium is classed in the last or alkali group
instead of in the alkaline earth group. It is not precipitated by the (NH4),C0j,
yet in the general properties of its salts it is so closely related to Ba , Sr and
Ca , that it is much better regarded as a subdivision of that group than as
belonging to the alkali group (§175 and ff.).
2. Traces of Ba , Sr and Ca may remain in solution after adding (NH4),COs
and warming; due to the solvent action of the ammonium salts present. To
prevent these traces giving a test for magnesium with Na2HP04 , a drop or
two of (NH4)2S04 is added to remove barium or strontium and a few drops of
(NH4)2C204 to remove calcium. The precipitate (if any forms) is removed by
filtration, before the Na2HF04 is added.
3, The precipitate of MgNH4F04 does not always form rapidly if only small
amounts of Mg are present, and the solution should be allowed to stand.
Rubbing the sides of the test tube with a glass stirring rod promotes the pre-
cipitation.
Ji. The precipitation of Mg as MgNH4F04 is fairly delicate (1-71,492) (Kissel,
Z., 1869, 8, 173): but not at all cMmcteristic^ as the phosphates of nearly all the
metals are white and insoluble in water. Hence the reliability of this test for
220
SEPARATION OF BARIUM, STRONTIUM AND CALCIUM. §194,5.
magnesium depends upon the rigid exclusion of the other metals (not alkalis)
by the previous processes of anal^'sis.
J. Lithium phosphate is not readily soluble in water or ammonium salts and
may give a test for magnesium. See §210, Gd.
§196. The unlike solubilities in alcohol, of the chlorides and nitrates of
barium, strontium and calcium enable us to separate them quite closely by
absolute alcohol, and approximately by " strong alcohol," as follows :
Dissolve the carbonate precipitate in HCl , evaporate to dryness on the
water-bath, rub the residue to a fine powder in the evaporating dish, and
digest it with alcohol. Filter through a small filter, and wash with alcohol
(5c, §§186, 187 and 188).
Besidue: BaCl, .
Dissolve in water, test
with GaSO« , SrSO« ,
KaCr^O, , etc.
Filtrate: SrCl, and GaCl, .
Evaporate to dryness, dissolve in water, change to
nitrates by precipitating with (NH4)2COs , wash-
ing and "dissolving in HNO, . Evaporate the
nitrates to dryness, powder, digest with alcohol,*
filter and wash with alcohol (or digest and wash
with equal volumes of alcohol and ether).
Besidue: Sr(NO,), .
Precipitation by CaS04
in water solution;
flame test, etc.
Filtrate: Ca(NO,), .
Precipitation by HsSOf
in alcohol solution, by
(NH4)2Ca04. etc.
Or, the alcoholic filtrate of SrClg and CaClg may be precipitated with (a
drop of) sulphuric acid, the precipitate filtered out and digested with
solution of (NHJjSO^ and a little NH^OH . Residue, SrS04 . Solution
contains CaSO^ , precipitable by oxalates.
§196. If the alkaline earth metals are present in the original material
as phosphates, or in mixtures such that the treatment for solution will
bring them in contact with phosphoric acid; the process of analysis must
be modified. One of the methods given under analysis of third and fourth
group metals in presence of phosphates (§146 and //.) must be employed.
§197. The presence of oxalates will also interfere, necessitating the
evaporation and ignition to decompose the oxalic acid (§151).
* Instead of alcohol the residue of the nitrates may be boiled with amyl alcohol. Calcium
liitrate is dissolved making a complete separation from the strontium nitrate ($18S« 5c).
§200. THE ALKALI GROUP. 221
The Alkali Gboup (Sixth Group).
Potassium. K = 39.11. Caesium. Cs = 132.9.
Sodium. Na = 23.05. Rubidium. Eb = 85.4.
Ammonium. (NH4)'. Lithium. li = 7.03.
§198. The metals of the alkalis are highly combustible, oxidizing quickly
in the air, displacing the hydrogen of water even more rapidly than zinc
or iron displaces the hydrogen of acids, and displacing non-alkali metals
from their oxides and salts. As elements they are very strong reducing
agents, while their compounds are very stable, and not liable to either re-
duction or oxidation by ordinary means. The five metals, Cs, Eb, K,
Ha, Li, present a gradation of electro-positive or basic power, caesium
being strongest, and the others decreasing in the order of their atomic
weights, lithium decomposing water with less violence than the others.
Their specific gravities decrease,* their fusing points rise, and as carbon-
ates their solubilities lessen, in the same order. In solubility of the phos-
phate, also, lithium approaches the character of an alkaline earth.
Ammonium is the basal radical of ammonium salts, and as such has
many of the characteristics of an alkali metal. The water solution of the
gas ammonia, NH3 (an anhydride), from analogy is supposed to contain
ammonium hydroxide, NH4OH, known as the volatile alkali. Potassium
and sodium hydroxides are the fixed alkalis in common use.
§199. The alkalis are very soluble in water, and all the important salts
of the alkali metals (including NH4) are soluble in water, not excepting their
carbonates, phosphates (except lithium), and silicates; while all other
metals form hydroxides or oxides, either insoluble or sparingly soluble, and
carbonates, phosphates, silicates, and certain other salts quite insoluble in
water.
Their compounds being nearly all soluble, the alkali metals are not pre-
cipitated by ordinary reagents, and, with few exceptions, their salts do not
precipitate each other. In analysis, they are mostly separated from other
metals by non-precipitation.
§200. In accordance with the insolubility in water of the non-alkali
hydroxides and oxides, the alkali hydroxides precipitate all non-alkali metals,
except that ammonium hydroxide does not precipitate barium, strontium,
and calcium. These precipitates are hydroxides, except those of mercury,
silver, and antimony. But certain of the non-alkali hydroxides and
oxides, though insoluble in water, dissolve in solutions of alkalis; hence,
when added in excess, the alkalis redissolve ths precipitates they at first pro-
duce with salts of certain metals, viz. : the hydroxides of Fb , Sn , Sb (oxide),.
• Except thofle of potassium (0.875) and sodium (0.9786).
223 POTASSIUM. §201.
Zn, Al, and Cr dissolve in the fixed alkalis; and oxide of Ag and hy-
droxides of CiL ^ Cd , Zn ^ Co 9 and Ni dissolve in the volatile alkali.
§201. Solutions of the alkalis are caustic to the taste and touch, and
turn red litmns blue; also, the carbonates, acid carbonates, normal and
dibasic phosphates, and some other salts of the alkali metals, give the
"alkaline reaction" with test papers. Sodiuoi nitroferricyanide, with
hydrogen sulphide, gives a delicate reaction for the alkali hydroxides
(§207, 66).
§202. The hydroxides and normal carbonates of the alkali metals are not
decomposed by heat alone (as are those of other metals), and these metals
form the only acid carbonates obtained in the solid state.
§203. The fixed alkalis, likewise many of their salts, melt on platinum
foil in the flame, and slowly vaporize at a bright red heat. All salts of
ammonium, by a careful evaporation of their solutions on platinum foil,
may be obtained in a solid residue, which rapidly vaporizes, wholly or
partly, below a red heat (distinction from fixed alkali metals).
§204. The hydroxides of the fixed alkali metals, and those of their salts
most volatile at a red heat, preferably their chlorides, impart strongly
characteristic colors to a non-luminous flame, and give well-defined spectra
with the spectroscope.
§206. Potassium. K =z 39.11 . Valence one.
1. Properties.— iS'prrt^r aravity, 0.875 at 1?>° (Hanmhaiier, 5., 1873, 6, 655).
Meltiny pmtU, (V2A° (Hagen, C. C'., 1883, 129). lioilimj point, 719° to '731° (Car-
nelley and Williams, B., 1879, 12, 13(30); GG7° (Perman, J, C, 1889, 55, 328).
Silver-white metal with a bluish tinge. At ordinary temperature of a wax-like
consistency, ductile and malleable; at 0° it is brittle. It is harder than Na
and is scratched by Li , Pb , Ca and Sr . The glowing vapor is a very beautiful
intense violet (I>udley, Am., 1892, 14, 185). It is next to caesium and rubidium,
the most electro-positive of all metals, remains unchanged in dry air, oxidizes
rapidly in moist air, and decomposes water with great violence, evolving
hydrogen, burning with a violet flame. At a red heat CO and COj are
decomposed, at a white heat the reverse action takes place. Liquid chlorine
does not attack dry potr.ssium (Gautier and Charpy, C. r., 1891, 113, 597). Acids
attack it violenth', evolving hydrogen.
2. Occurrence. — Very widely distributed as a portion of many silicates. In
sea water in small amount as KCl . In numerous combinations in the large
salt deposits, especially at Stassfurt; e. (/., camalHte, KCl.MgCls + GHjO;
kainite, K2SO4.MgSO4.MgCl, -f 6H3O , etc. As an important constituent of
many plants — grape, potato, sugar-beet, tobacco, fumaria, rumex, oxalis, etc.
3. Preparation. — (J) By reduction of the carbonate with carbon. {2) By
electrolysis of the hydroxide (Horning and Kasemeyer, B., 1889, 22, 277c:
Castner, B., 1892, 25, 179c). (S) By reduction of K^CO, or KOH with iron car-
bide: 6K0H + 2FeC2 = 6K -f 2Fe -\- 2C0 -f 200, -\- 3H, (Castner, C. N., 1886.
54, 218). (Jf) By reduction of the carbonate or hydroxide with Fe or Mg
(Winkler, i?., 1890, 23, 44).
4. Oxides and Hydroxide. — Potassium oa^Wc,* K,0 , is prepared by carefully
* The existence of the oxides M^,0 of K, Nm. and Rb is disputed (Erdmann and Koethner, A^
1896, 204, 55).
§206, 66. POTASSIUM. 223
heating potassium with the necessary amount of oxygen (air) (Kuhnemann,
('. C, 1863, 491); also by heating K^O* with a mixture of K and Ag (Beketoff,
(*. C, 1881, 643). It is a hard, gray mass, melting above a red heat. Water
changes it to KOH with generation of much heat. Potassium hydroxide^ KOH»
is formed by treating K or K^O with water; by boiling a solution of KjCO,
wMth Ba, Sr or Ca oxides; by heating KsCO, with FejOs to a red heat and
decomposing the potassium ferrate with water (Ellershausen, C. C, 1891, (1),
1047; (2), 399). Pure water-free KOH is a white, hard, brittle mass, melting
at a red heat. It dissolves in water with generation of much heat. Potassium
superoxide^ K2O4 , is formed when K is heated in contact with abundance of air
(Harcourt, J, C, 1862, 14, 267); also by bringing K in contact with KNO,
heated until it begins to evolve O (Bolton, C. A'., 1886, 53, 289). It is an amor-
phous powder of the color of lead chromate. Upon ignition in a silver dish
oxygen is evolved and K3O and AgjO formed (Harcourt, ?. c). Moist air or
water decomposes it with evolution of oxygen. It is a powerful oxidizing
agent, oxidizing S** to Svi , P<» to Pv , K , As , Sb , Sn , Zn , Cu , Fe , Ag and Pt
to the oxides (Bolton, /. c; Brodie, Proc, Roy. Soc., 1863, 12, 209).
5. Solubilities. — K and K^O dissolve in water with violent action, forming
KOH , which reacts with all acids forming soluble salts. Potassium dissolves
in alcohol, forming potassium alcoholate and hydrogen.
Potassium platinum chloride, acid tartrate, silico-fluoride, picratc, phos-
phomolybdate, perchlorate, and chlorate are only sparingly soluble in
cold water, and nearly insoluble in alcohol. The carbonate and sulphate
are insoluble in alcohol.
6. Beactions. a. — Potassium and sodium hydroxides are very strong
bases, fixed alkalis, and precipitate solutions of the smalts of all the other
metals (except Cs , Rb , and Li), as oxides or hydroxides. These precipi-
tates are quite insoluble in water, except the hydroxides of Ba, Sr, and
Ca . Excess of the reagent causes a resolution with the precipitates of
Pb , Sb , Sn , Al , Cr , and Zn , forming double oxides as, K^PbOo , potas-
sium plumbite, etc. Potassium carbonate is deliquescent, strongly alkaline.
and precipitates solutions of the salts of the metals (except Cs , Eb , Na ,
and li), forming normal carbonates with Ag , Hg', Cd , Fe", Mn', Ba , Sr ,
and Ca ; oxide with Sb ; hydroxide with Sn , Fe'", Al , Cr"' and Co'"; basic
salt with Hg", and a basic carbonate with the other metals.
&.— The potassium salts of HCN , H,Fe(CN)« , H3Fe(CN)e , and HCNS
find extended application in the detection and estimation of many of the
heavy metals.
Tartaric acid, 13.^0^3.^^^^ or more readily sodium hydrogen tartrate,
'StSRJSLfif^y precipitates, from solutions sufficiently concentrated, potas-
sium hydrogen tartrate, KHC^H^O^, granular-crystalline. If the solution
be alkaline, tartaric acid should be added to strong acid reaction. The
test must be made in absence of non-alkali bases. The precipitate is in-
creased by agitation, and by addition of alcohol. It is dissolved by fifteen
parts of boiling water or eighty-nine parts water at 25°, by mineral acids,
by solution of borax, and by alkalis, which form the more soluble normal
tartrate, K2C4H4OS , but not by acetic acid, or at all by alcohol of fifty
per cent
•224 POTASSIUM. §205, 6r.
Picric acid, C«H2(N02)30H , precipitates, from solutions not very dilute,
the yellow, crj'stalline potassium picratey C^^(SO^)^QiK , insoluble in alco-
hol, by help of which it is formed in dilute solutions. The dried precipi-
tate detonates strongly when heated.
c. — If a neutral Bohition ot a potassium salt be added to a solution of cobaltie
nitrite,* a precipitate of the double salt potassium cobaltie nitrite, K,Co(NOa), ,
will be formed. In concentrated solutions the precipitate forms immediately,
dilute solutions should be allowed to stand for some time; sparingly soluble in
water, insoluble in alcohol and in a solution of potassium salts, hence the
precipitation is more valuable as a separation of cobalt from nickel than as a
test for potassium (§132, 6o).
Potassium nitrate is not found abundantly in nature, but is formed by the
decomposition of nitrogenous organic substances in contact with potassium
salts, *' saltpeter plantations "; or by treating a hot solution of NaNO, with
KCl (/>., 2, 2, 72). it finds extended application in the manufacture of gun-
powder. (/.—See §206, erf.
c.— Potassinm sulphide may be taken as a type of the soluble sulphide?
which precipitates solutions of the metals of the first four groups as
sulphides except: Hg' becomes HgS and Hg°, Fe'" becomes FeS and 8,
and Al and Cr form hydroxides. The sulphides of arsenic, antimony and
tin dissolve in an excess of the reagent, more rapidly if the alkali sulphide
contain an excess of sulphur. For the general action of HoS or soluble
sulphides as a reducing agent see the respective metals. Potassium sul-
phate is used to precipitate barium, strontium, and load. It almost alway?
occurs in nature as double salt with magnesium, KoS04.MgS04.MgCl2 +
(JH.O , kainite, and is used in the manufacture of KA1(S0 J. ^ KXOj and
£0H . As a type of a soluble sulphate it precipitates solutions of lead,
mercurosum, barium, strontium, and calcium; calcium and mercurosum
incompletely.
/. — Potassium chloride precipitates the metals of the first group, acting
thus as a type of the soluble chlorides. It is much used with sodium
nitrate in the preparation of potassium nitrate for the manufacture of
gunpowder, in the preparation of K2CO3 , KOH , and also as a fertilizer.
Potassium bromide as a type of the soluble bromides precipitates solutions
of Pb , Ag , and Hg (Hg" incompletely). Potassium iodide finds extended
use in analytical chemistry in that it forms many soluble double iodide?:
it is also extensively used in medicine. As a typo of a soluble iodide it
precipitates solutions of the salts of Pb , Ag, Hg , and Cu'. Cu" salt^
are precipitated as Cul with liberation of iodine. Fe"' salts are merely
reduced to Fe" salts with liberation of iodine. Arsenic acid is merely
reduced to arsonous acid with liberation of iodine.
• One cc. of cobaltous nitrate solution and three cc. of acetic acid are added to five cc. of a ten
per cent solution of sodium nitrite. This srives a yellowish solution having an odor of nitxoas
jusid.
1,7. POTASSIUM, 225
Potassium chlorate is used as a source of oxygen and as an oxidizing agent
in acid solutions. Sodium perchlorate, NaClOf , precipitates from solutions of
potassium salts potaitsium peichloratey KCIO4 , sparingly soluble in water and
almost insoluble in strong alcohol (Kreider, Z. anory., 1895, 9, 342). Potassium
iodate is used as a reagent in the detection of barium as Ba(I0s)2 . g. — The
oxides of arsenic act as acid anhydrides toward KOH and form stable soluble
potassium salts, arsenites and arsenates, which react with the salts of nearly
all the heavy metals, h. — Potassium chromate and dichromate are both exten-
sively used as reagents, especially in the analysis of Ag , Pb and Ba salts.
i. — Flnosilicic acid, HsSiFg, precipitates from a neutral or slightly
acid solution of potassium salts, potassium fluosilicate (silico-fluoride),
ZjSiFe, soluble in 833.1 parts of water at 17.5*^; in 104.8 parts at 100°;
and in 327 parts of 9.6 per cent HCl at 14° (Stolba, J. pr., 1868; 103, 396\
The precipitate is white, very nearly transparent.
;. — Platinic Chloride, PtCl4 , added to neutral or acid solutions not too
dilute, with hydrochloric acid if the compound be not a chloride, precipi-
tates potassium platinic chloride, (KCl)oPtCl4 , crystalline, yellow. Non-
alkali bases also precipitate this reagent, and if present must be removed
before this test. The precipitate is soluble in 19 parts of boiling water,
or 111 parts of water at 10°. Minute proportions are detected by evapor-
ating the solution with the reagent nearly to dryness, on the water-bath,
and then dissolving in alcohol; the yellow crystalline precipitate, octahe-
dral, remains undissolved, and may be identified under the microscope.
k. — An alcoholic solution of BiCl, in excess of Na^SjO, gives a yellow pre-
cipitate with solutions of potassium salts (Pauly, C. C, 1887, 553). I. — Gold
chloride added to sodium and potassium chloride forms double salts, e. g,,
KCl.AuCl, + 2H2O . If these salts are dried at 100° to 110** to remove water
and acids, the sodium salt is soluble in ether (separation from potassium)
(Fasbender, 0. C, 1894, 1, 409).
7. Ignition. — Ignited potassium hydroxide or potassium carbonate is a
valuable desiccating agent for use in desiccators or in liquids. A mixture
of molecular proportions of K2CO3 and NEoCOs melts at a lower tempera-
ture than either of the constituents, and is frequently employed in fusion
for the transposition of insoluble metallic compounds : BaS04 + K2CO3 =
BaCOg + K2SO4 .
Potassium compounds color the flame violet. A little of the solid
substance, or residue by evaporation, moistened with hydrochloric acid,
is brought on a platinum wire into a non-luminous flame. The wire
should be previously washed with HCl , and held in the flame to insure .
the absence of potassium. The presence of very small quantities of
sodium enables its yellow flame completely to obscure the violet of potas-
sium; but owing to the greater volatility of the latter metal, flashes of
violet are sometimes seen on the first introduction of the wire, or at the
border of the flame, or in its base, even when enough sodium is present
to conceal the violet at full heat. The interposition of a blue glass, or
226 SODIUM. §206,8.
prism filled with indigo solution, sufficiently thick, entirely cuts off the
yellow light of sodium, and enables the potassium flame to be seen. The
red rays of the lithium flame are also intercepted by the blue glass or
indigo prism, a thicker stratum being required than for sodium. If
organic substances are present, giving luminosity to the flame, they must
be removed by ignition. Certain non-alkali bases interfere with the
examination. Silicates may be fused with pure gypsum, giving vapor of
potassium sulphate. Bloxam (J, C, 1865, 18, 229) recommends to fuse
insoluble alkali compounds with a mixture of sulphur, one part, and
barium nitrate, six parts: cool, dissolve in water, remove the barium with
NH^OH and (NHJ^COs and test for the alkalis as usual.
The volatile potassium compounds, when placed in the flame, give a
widely-extended continuous spectrum, containing two characteristic lines;
one line, K a, situated in the outermost red, and a second line, K ?, far in
the violet rays at the other end of the spectrum.
8. Detection. — Potassium is usually identified by the violet blue color
which most of its salts impart to the Bunsen flame (7). Sodiimi inter-
feres but the intervention of a cobalt glass (§132, 7) or a solution of
indigo cuts out the yellow color of the sodium flame and allows the violet
of the potassium to be seen. Some of the heavy metals interfere, hence
the test should be made after the removal of the heavy metals (§§211
and 212).
Potassium may be precipitated as the platinichloride {(Sj)', as the per-
chloratc (6/); as the silico-fluoride ifii)\ as the acid tartrate (66); etc.
Certain of these reactions are much used for the quantitative estimation
(9) of potassium but are seldom used for its detection qualitatively.
9. Estimation. — (/) Potassium is converted into the sulphate or phosphate
and weighed as such. {2) It is precipitated and weighed as the double chloride
with platinum. (.?) If present as KOH or K^CO, it is titrated with standard
acid (Kippenberprer, Z. amjeir,, 1S94, 495). (J) It is precipitated with H.SiF«
and strong alcohol. {5) Indirectly when mixed with sodium, by converting'
into the chlorides and weighing as such: then determining the amount of
chlorine and calculating the relative amounts of the alkalis. ((*)"It is pre-
cipitated as the bitartrate in presence of alcohol and, after filtration and
solution in hot water, titrated with deci-normal KOH. (7) By precipitation as
the perchlorate, KCIO, (Wense, Z. mujew,, 1892, 233; Caspari, Z. angew., 1893, GS).
10. Oxidation. — Potassium is a very powerful reducing agent, its affinity
for oxygen at temperatures not too high is greater than that of any other
element except Cs and Rb . For oxidizing action of K^O^ see 4.
§206. Sodium. Na = 23.05 . Valence one.
1. Properties.— ;f?/)PCi/?c gravity, 0.9735 at 13.5° (Banmhauer, B., 1873, 6, 605);
0.7414 at the boiling point (Ramsay, B., 1880, 13, 2145). Meimg poini, 97.6°
(Hag-en, B., 1883, 16, 1668). B(Hling point, 742° (Perman, C. iV., 1889, 59, 237).
§206, 6d. SODlUAf. 227
A silver- white metal with a strong metallic lustre. At ordinary temperatures
it is softer than Li or Pb, and can be pressed together between the fingers;
at — 20** it is quite hard; at 0** very ductile. It oxidizes rapidly in moist air
and must be kept under benzol or kerosene. It decomposes water violently
even at ordinary temperatures, evolving hydrogen, which frequently ignites
from the heat of the reaction: 2Na + 2H3O = 2NaOH -f Hj . It burns, when
heated to a red heat, with a yellow flame. Pure dry Na is scarcely at all
attacked by dry HCl (Cohen, C. N., 1886, 54, 17).
2. Occurrence. — Never occurs free in nature, but in its various combinations
one of the most widely' diffused metals. There is no mineral known in which
its presence has not been detected. It occurs in all waters mostly as the
chloride from traces in drinking waters to a nearly saturated solution in some
mineral waters and in the sea water. It is found in enormous deposits as rock
salt, NaCl; as Chili saltpeter, NaNOj; in lesser quantities as carbonate, borate,
Kulphate, etc.
3. Preparation. — (/) By igniting the carbonate or hj'droxide with carbon;
(2) by igniting the hydroxide with metallic iron; (3) by electrolysis of the
hydroxide; (4) by gently heating the carbonate with Mg .
"4. Oxides and Hydroxides. — Sodium oxide, Na^O , is formed by burning
sodium in oxygen or in air and heating again with Na to decompose the Na^O,
(§205, 4, footnote). Sodium hydroxide, NaOH , is formed by dissolving the
metal or the oxide in water (Rosenfeld, J. pr.y 1893, (2), 48, 599); by treating
a solution of sodium carbonate with lime; by fusion of NaNO, with CaCO, ,
CaO and NaoCOs are formed and the mass is then exhausted with water; by
igniting NaaCOs with FejO, , forming sodium ferrate, which is then* decom-
posed with hot water into NaOH and Fe(OH), (Solvay, C. C, 1887, 829). It is
a white, opaque, brittle crystalline body, melting under a red heat. The
fused mass has a sp. gr. of 2.13 (Filhol, A.Ch., 1847, (3), 21, 415). It has a very
])Owerful affinity for water, gradually absorbing water from CaCla (Muller-
Erzbach, fi., 1878, 11, 409). It is soluble in about 0.47 part of water according
to Bineau (C. r., 1855, 41, 509).
Sodium i)eroxide, NajOj , is formed by heating sodium in CO.2 free air or
oxygen (Prud'homme, C. C, 1893, (1), 199). It reacts as HjO, , partly reducing
and partly oxidizing. It may be fused without decomposition. Water decom-
poses it partially into NaOH and HjOj .
5. Solnbilities. — Sodium and sodium oxide dissolve in water, forming
the hydroxide, the former with evolution of hydrogen. In acids the
corresponding sodium salts are formed, all soluble in water except sodium
pyroantimonate, which is almost insoluble in water, and the fliiosilicate
sparingly soluble.
The nitrate and chlorate are deliquescent. The carbonate (10 aq), sul-
phate (10 aq), sulphite (8 aq), phosphate (12 aq), and the acetate (3 aq) are
efflorescent.
6. Beactions. a. — As reagents sodium hydroxide and carbonates act in
«11 respects like the corresponding potassium compounds, which see.
h, — By the greater solubility of the picrate and acid tartrate of sodium, that
metal is separated from ])otassium (§205, 6ft). c.^^odlum nitrate occurs in
nature in large quantities as Chili saltpeter, used as a fertilizer, for the manu-
facture of nitric acid, with KCl for making KNO, , etc.
d. — Sodium phosphate, JStLMTO^, is much used as a reagent in the
precipitation and estimation of Pb , Mn , Ba , Sr , Ca , and Mg . The
phosphates of all metals except the alkalis are insoluble in water (lithium
phosphate is only sparingly soluble (§210, 5r), soluble in acids). Solu-
228 SODIUM. §206, 6e.
tions of alkali phosphates precipitate solutions of all other metallic salt?
as phosphates (secondary, tertiary or basic) except : HgCL; precipitates as
a basic chloride (§68, Gd), and antimony as oxide or oxychloride (§70, 6rf).
€, A //♦ h. — As reagents the sodium salts react similar to the corresponding'
potassium salts, which see. i. — ^Sodium fluosilicate is soluble in 153.3 parts
HaO at 17.5° and in 40.G6 parts at 100° (Stolba, Z., 1872, 11, 199); hence is not
precipitated by fluosilicic acid except from very concentrated solutions
(separation from K). /.—Sodium platinic chloride, (NaGl)aPtCl4 , crystallizes
from its concentrated solutions in ml prisms, or prismatic needles (distinction
from potassium or ammonium). A drop of the solution to be tested is slightly
acidified with hydrochloric acid from the point of a jflass rod on a slip of glass,
treated with two drops of solution of platinic chloride, left a short time for
spontaneous evaporation and crystallization, and observed under the micro-
scope.
k. — Solution of potassium pyroantimonate, K^HzShzO^ , produces in
neutral or alkaline solutions of sodium salts a slow-forming, white, crv'stal-
line precipitate, NaoH^SboO. , almost insoluble in cold water. The reagent
must be carefully prepared and dissolved when required, as it is not per-
manent iu solution (§70, 4r).
7. Ignition.— Sodium bicarbonate, NaHCO., , loses HoO and CO2 at 125'
becoming NaoCO^ , no further decomposition till 400° when a very small
amount of NaOH is formed (Kirsling, Z. angew., 1889, 332).
Sodium compounds color the flame intensely yellow, the color being
scarcely affected by potassium (at full heat), but modified to orange-reil
by much lithium, and readily interce])ted by blue glass. Infusible com-
pounds may be ignited with calcium sulj)hate. The test is interfered witli
by some non-alkali bases, which should be removed (§§211 and 212).
The spectrum of sodium consists of a single broad band at the D line iu
the yellow of the solar spectrum separable into two bands, D^ and D^/,^}'
prisms of higher refractive power.
The amount of sodium in the atmosphere, and in the larger number of
substances designed to be " chemically pure '' is sufficient to give a dis-
tinct but evanescent yellow color to the flame and spectrum.
8. Detection. — Sodium is usually detected by the color of the flame,
yellow, in absence of the heavy metals. In the usual process of analysis
the presence or absence of sodium is determined in the presence 01
magnesium (as NajHPO^ is the usual reagent for the detection of mag-
nesium, it is evident that the presence or absence of the sodium must ^
determined before the addition of that reagent); and as that metal give^
a yellowish color to the flame it must be removed if small quantities 01
sodium are to be detected. For this purpose the filtrate from Ba , Sr and
Ca is evaporated to dryness and gently ignited to expel all ammonium
salts; then taken up with a small amount of water and the magnesium
precipitated as the hydroxide with a solution of barium hydroxide. After
§207, 5. AMMONIUM, 229
filtration the barium is removed by (NH4)2C03 or H2SO4 and the filtrate
tested for sodium by the flame or by the pyroantimonate test (6A;).
9. Estimation. — (i) If present as hydroxide or carbonate, by titration with
standard acid (Lunge, Z. anycir., 1897, 41). (2) By converting into the chloride
or sulphate and weighing as such. (3) In presence of potassium by converting
into the chloride, weighing as such, then estimating the amount of chlorine
with AgNO, and computing the amounts of K and Na. (4 It is precipitated
by KjHsSb^OT and dried and weighed as Na^H^hiO^ .
10. Oxidation. — Sodium ranks with potassium as a very powerful re-
ducing agent. It is not quite so violent in its reaction, and being much
cheaper is almost universally used instead of potassium. Sodium peroxide .
may act both as a reducing and oxidizing agent. The action is similar to
H2O2 in alkaline solution, which see (§244, 6).
§207. Ammoninm. (NHJ'. Valence one.
1. Properties.— ;Srppct/?r gravity of NH, gas, 0.589 (Fehling, 1, 384); of the
liquid, 0.6234 at 0° (JoHy, A., 1861, 117, 181). The liquid boils at —33.7°, at
O** the liquid has a tension of 4.8 atmospheres (Bunsen, Pogg., 1839, 46, 95).
Liquid ammonia is a colorless mobile liquid, burns in air when heated or in
oxygen without being previously heated. At ordinary temperature it is a gas
-with very penetrating odor. It burns with a greenish-yellow flame, and com-
Tjines energetically with acids to form salts, the radical NH^ being monovalent
and acting in many respects similar to K and Na . At 0** one volume of water
absorbs 1049.6 volumes of the gas; at 15**, 727.22 volumes (Carius, A., 1856, 09,
144). One g^ram of water, pressure 760 mm. and temperature 0**, absorbs
0.899 gram of NH3: with temperature 16**, 0.578 gram (Sims, A., 1861, 118, 345).
2. Occurrence. — Free ammonia does not occur in nature. Various ammonium
salts occur widely distributed: in rain water, in many mineral waters, in almost
all plants, among the products of the decay or decomposition of nitrogenous
organic bodies, etc.
3. Preparation. — It is obtained from the reduction of nitrates or nitrites by
nascent hydrogen in alkaline solution, €. g,, 8A1 + 5K0H -|- 3KN0, -|- 2H2O =
8KAIO3 + 3NH,; by the reduction with the hydrogen of the zinc-copper couple;
by boiling organic compounds containiiig nitrogen with KMn04 in strong
alkaline solution (as in water analysis): also by the oxidation of nitrogen in
organic bodies with strong sulphuric (Kjeldahl method of nitrogen determina-
tion). It is prepared on a larger scale by heating an ammonium salt with lime
(or some other strong base). Nearly all the ammonium hydroxide and am-
monium salts of commerce are obtained as a by-product in the production of
illuminating gas by the destructive distillation of coal.
4. Hydroxide. — Ammonium hydroxide, NH4OH, is made by passing
ammonia, NH3 , into water. The gas is absorbed by the water with great
avidity, and a strongly alkaline solution is produced. A solution having
a sp. gr. of 0.90 at 15® contains 28.33 per cent of NH3 (Lunge and Wiemik,
Z. angew., 1889, 183).
5. Solubilities. — Ammonia, NHg , and all ammonium salts are soluble in
water. Ammonia dissolves less readily in a strong solution of potassium
hydroxide than in water. The carbonate (acid), and phosphate are efflores-
cent. The nitrate and acetate are deliquescent^ the sulphate slightly deli-
quescent.
230 AMMONIUM. §207, 6a.
6. Beaotions. a, — The fixed alkali hydroxides and carbonates liberate
ammonia, NHg , from all ammonium salts, in the cold and more rapidly
upon heating. Ammonium hydroxide, volatile alkali, colors litmus blue,
neutralizes acids, forming salts, and precipitates solutions of the metals of
the first four groups, manganese and magnesium salts imperfectly; due to
the solubility of the hydroxide formed, in the ammonium salt produced
by the reaction, and with these metals if excess of ammonium salts be
present no precipitate will be formed by the NH^OH . The precipitate is
a hydroxide except: with Ag and Sb it is an oxide, with mercury a sub-
stituted ammonium salt and with lead a basic salt (see below, h and I),
With salts of Ag , Cu , Cd , Co , Ni , and Zn the precipitate redissolves in
excess of the reagent. Ammonium carbonate, (NNJoCOj, , is unstable and
used only in solution. It is formed by adding ammonium hydroxide to a
solution of the acid carbonate of commerce. It precipitates solutions of
all the non-alkali metals, chiefly as carbonates except magnesium salts
which are not at all precipitated, as a soluble double salt is at once formed
(separation of Ba , Sr , and Ca from Mg). With salts of Ag , Cu , Cd , Co ,
Ni , and Zn , the precipitate is redissolved by an excess of the ammonium
carbonate.
6. — Dilute solutions of picric acid with ammonium hydroxide form in-
tensely colored yellow solutions, a precipitate of ammonium picrate is
formed if the solutions are quite concentrated. Tartaric acid precipitates
ammonium salts very^ closely resembling the precipitate of potassium acid
tartrate. The ammonium salt is more soluble in water than the potas-
sium salt and does not leave K2CO., upon ignition. Sodium nitroferri-
cyanide, Na2Fe(N0)(CN)5 , added to a mixture of NH^OH and HoS
[(NH4)2S] gives a very intense purple color, characteristic of alkali
sulphides and the manipulation may Ik? modified so as to give a very deli-
cate test for the presence of an alkali hydroxide or of hydrosulphuric acid.
In no case, however, can the HjS be directly added to the sodium nitro-
ferricyanide as it causes oxidation of the sulphur. To test for ammonia
the gas should be liberated by KOH and distilled into a solution of HoS ;
and this solution added to the Na,Je(NO)(CN), .
c, — Ammonium nitrite, NH4NO2 , is used in the preparation of nitrogen
(§235, :0; ammonium nitrate in the preparation of nitrous oxide, N2O .
" laupfhing gas *' (§237). d, — Ammonium phosphate, as a reagent, acts
similarly to sodium phosphate. When sodium phosphate, NasHFOf , is used to
precipitate metals in the presence of ammonium hydroxide, a double phosphate
of the metal and ammonium is frequently formed as MnNH^FOf, Mg'NH^FO^,
etc. Ry some chemists micro<osmic salt,"NaJi'H4HP04 , is preferred to sodium
phosphate, Na..HPO, , as a reagent.
(\ — When aninionium hydroxide is saturated with H2S , ammonium .'^ul'
phide, (1^^)28 , is formed. Complete saturation is indicated by the failure
§207, ek. AMMOXllM. 231
to precipitate magnesium salts, that is, NH4OH precipitates magnesium
salts while (NHJjS does not. Freshly prepared ammonium sulphide is
colorless, but upon standing becomes yellow with loss of ammonia and
formation of the poly-sulphides, (1^4)28,. The yellow poly-sulphide
may also be formed by dissolving sulphur in the normal ammonium sul-
phide. As a precipitant ammonium sulphide acts similarly to the fixed
alkali sulphides. The sulphides of Sb'" and 8n" are with great difficulty
soluble in the normal ammonium sulphide, but readily soluble in the
poly-sulphide. Nickel sulphide, NiS, is insoluble in normal ammonium
sulphide but is sparingly soluble in the yellow poly-sulphide (distinction
from cobalt). (NH4)28 gives a rich purple color with sodium nitroferri-
cyanide (h), Ammonitun sulphate as a precipitating reagent acts similar
to all soluble sulphates (§206, 6e), A 25 per cent solution of (11^114)2804
is used to dissolve Ca804 (§188, 5c) (distinction from Ba and Sr).
f. — ^Ammonitun chloride is much used as a reagent. It prevents pre-
cipitation of the salts of Mn by the NH4OH , and is of special value in the
precipitation of the third group as hydroxides and the fourth group as
sulphides by preventing the formation of soluble colloidal compounds.
The solubility of the precipitates of the carbonates of the fifth group is
slightly increased by the presence of ammonium chloride ; i. e., very dilute
solutions- of barium chloride are not precipitated by ammonium carbonate
in presence of a large excess of ammonium chloride. The salts of mag-
nesium are not precipitated by the alkalis or by the alkali carbonates in
presence of ammonium chloride. The solubility of A1(0H)3 is diminished
by the presence of NH4CI (§124, 6a, and §117).
17, h. — Similar as reagents to the corresponding* potassium salts, i. — ^Fluo-
silicic acid, HaSiF, , does not precipitate ammonium salts, the ammonium
fluosillcate being" very soluble in water (distinction from potassium). /. — Plat-
inum chloride, PtCl* , forms with ammonium salts the yellow double ammonium
platinum chloride, (NH4)2PtCl, , very closelj' resembling" the potassium salt
with the same reagent, but upon ignition only the spongy metallic platinum
is left, i. e., no chloride of the alkali metal, as KCl .
Jc. — ^A solution of potassium mercuric iodide, 'SLflgl^, containing also
potassium hydroxide — Nessler's test * — produces a brown precipitate of
nitrogen dimercuric iodide, "SILgJ. , dimercur-ammonium iodide (§68, Qa),
soluble by excess of KI and by HCl ; not soluble by KBt (distinction from
HgO):
NH, + 2HgI, = NHgJ 4- 3HI
KH4OH -h 2K,HgI« + 3K0H = NHg,I + 7KI + 4HaO
* This reagent may be prepared as foUows : To a solution of mercuric chloride add solution
€f potMsium iodide tUl the precipitate is nearly all redlssolved ; then add solution of potassium
hydxozide sufBcient to liberate ammonia from ammonium salts ; leave until the liquid becomes
4sle«r, and deoant from any remaining sediment.
232 AlIMONIUM. §207, «,
This very delicate test is applicable to ammonium hydroxide or salts;
traces forming only a yellow to brown coloration. The potassium mercuric
iodide, " Meyers Reagent," alone, precipitates the alkaloids from neutral
or acid solutions, but does not precipitate ammonium salts from neutral
or acid solutions. Ammonium hydroxide in alcoholic solution does not
give a precipitate with Xessler's reagent, but from this solution a precipi-
tate is formed with HgClj (De Koninck, Z., 1893, 32, 188).
/.—Mercuric chloride, HgClg , forms, in solutions of ammonium hy-
droxide or ammonium carbonate, the " white precipitate " of nitrogen
dihydrogen mercuric chloride, NHsHgCl, or mercur-ammonium chloride.
If the ammonium is in a salt, not carbonate, it is changed to the carbonate
and precipitated, by addition of mercuric chloride and potassium carbonate
previously mixed in solutions (with pure water), so dilute as not to precipi-
tate each other (yellow). This test is intensely delicate, revealing the
presence of ammonia derived from the air by water and many substances
(Wittstein, Arch. Pharm,, 1873, 203, 327).
m. — Add a sman quantity of recently precipitated and weU-washed silver
chloride, and, if it does not dissolve after ag-itation, then add a little potassium
hydroxide solution. The solution of the AgCl , before the addition of the ttxed
alkali, indicates free ammonia; after the addition of the fixed alkali, ammonium
salt. (Applicable in absence of thiosulphates, iodides, bromides and sulpho-
cyanates.)
n. — Sodium phosphoxnolybdate (§75, firf) precipitates ammonium from neutraV
or acid solutions; also precipitates the alkaloids, even from very dilute solu-
tions, and, from concentrated solutions, likewise precipitates K , Rb and C^
(all the fixed alkalis except Na and Li).
7. Ignition. — Heat vaporizes the carbonate, and the haloid salts of ar^r-
monium, undecomposed (dissociated but reuniting" upon cooling); decompos«£^:
the nitrate with formation of nitrous oxide and water, and the phosphate ar-^^
borate with evolution of ammonia. NH, heated to 780** or hig-her is dissociat^^
into N and H (Ramsay and Young, J, C, 1884, 45, 88).
8. Detection. — As ammonium hydroxide and chloride arc nsed in tl^B
regular process of analysis, the original solution must be tested for tF'
presence or absence of ammonium compounds. The hydroxide of tTT
carbonate may be detected by the odor (1) ; the action on red litmus pap^^=
suspended in the test-tube above the heated solution; the blue color ii^c:
parted to paper wet with copper sulphate; the blackening of mercuro— ■*
nitrate paper; and if in considerable quantity, the white vapors wh— -^^
brought into contact with the vapors of volatile acids. In combinati-^'^^
as salts the gas is liberated by the fixed alkali hydroxides or carbonatzr^ij
(oxides or hydroxides of Ba , Sr , or Ca may be used) and distilled in ^o
Nessler's reagent, or collected in water and the test with HgClg (0?) appli^
or any of the tests for ammonium hydroxide.
9. Estimation. — Ammonium salts are usually estimated by distillation into a
standard acid, from a solution made alkaline with KOH , and titration of the
excess of the acid with a standard NH«OH solution, using tincture of cochineal
§5808, 5. CAUSIIM, 233
as an indicator. It may be converted into the chloride and precipitated by
PtCl4 and weighed as the double platinum salt.
10. Oxidation. — Ammonium salts in solution, treated with chlorine gas, gen-
erate the unstable and violently explosive "nitrogen chloride" (NCI,?) («).
The same product is liable to arise from solid ammonium salts treated with
chlorine. Gaseous ammonia, and ammonium hj'droxide, with chlorine gas,
generate free nitrogen (&), a little ammonium chlorate being formed if the
ammonia is in excess. Hypochlorites or hypobromiies (or chlorine or bromine
dissolved in aqueous alkali, so as to leave an alkaline reaction) liberate, from
dissolved ammonium salts, all of their nitrogen (as shown in the second equa-
tion of b); the measure of the nitrogen gas being a means of quantitative
estimation of ammonium. With iodine, ammonium iodide and the explosive
iodamides (c) are produced; or under certain conditions an iodate (rf). Ammo-
nium hydroxide is liable to atmospheric oxidation to ammonium nitrite and
nitrate. Permanganates oxidize to nitrate (e) (Wanklyn and Gamgee, J. C,
1868, 21, 29). In presence of Cu the O of the air oxidizes the nitrogen of
ammonia to a nitrite (f) (Berthelot and Saint-Gilles, A, Ch,, 1864, (4), 1, 381).
Ammonia is somewhat readily produced from nitric acid by strong reducing
agents (g). It is formed with carbonic anhydride, in a water solution of
cyanic acid, and, more slowly, in a water solution of hydrocyanic acid. It is
generated, by fixed alkalis, in boiling solution of cyanides (h); also in boiling
solutions of albuminoids and other nitrogenous organic compounds, this forma-
tion being hastened and increased by addition of permanganate (Wanklyn's
process). Fusion with fixed alkalis transforms all the nitrogen of organic
bodies into ammonia.
(fl) NH,C1 -f 3C1, = NCI, 4- 4HC1
(6) 8NH, + 3CI2 = 6NH,C1 + N»
2NH,C1 + 3C1, = 8HC1 -f N,
(0) 2NH, 4- la = NH,I + NH,I f
(d) 6NH4OH + 31, = 5NH,I -h NH4IO, + 3H,0
(€) 6NH,0H + 8HMnO, = 3NH,N0, + 8MnO(OH), + 5H,0
(f) 12Cu -f 2NH, -h 90, = l2CuO + 2HN0, + 2H,0
(g) 3HN0, -I- 8A1 4- 8K0H = 8KA10, -|- 3NH, 4" H^O
(h) HON 4- KOH 4- H,0 = NH, 4- KCHO, (formate).
§208. Caesium. Cs =: 132.9 . Valence one.
1. Properties.— Spref/lc gravity, 1.88 at 15" (Setterberg, A., 1882, 211, 100).
Melting point, between 26° and 27**. It is quite similar to the other alkali
metals; silver-white, ductile, very soft at ordinary temperature. It burns
rapidly when heated in the air, and takes fire when thrown on water. It may
be kept under petroleum. It is the most strongly electro-positive of all metals.
2. Occurrence. — Widely distributed but in small quantities: as caesium
aluminum silicate (mineral castor and pollux) (Pisani, C. r., 1864, 58, 715); in
many mineral springs (Miller, C. A"., 1864, 10, 181); in the ash of certain plants,
tobacco, tea, etc.
3. Preparation.— By electrolysis of a mixture of CsCN with Ba(CN)2; by
ignition of CsOH with Al in a nickel retort (Beketoff, C. C, 1891, (2), 450).
4. Oxide and Hydroxide. — An oxide has not yet been prepared. The
hydroxide, CsOH, is a grayish-white solid, very deliquescent, absorbs CO, from
the air; dissolves in water with generation of much heat, forming a strongly
caustic solution.
5. Solubilities. — Caesium dissolves with great energy in water, acids or
alcohol, liberating hydrogen and forming the hydroxide, salts or alcoholate
resi>ectively. The hydroxide is soluble in water and alcohol. The salts are
all quite readily soluble. The double platinum chloride, CSzPtCl, , and the
acid tartrate, CsHC^H^Oq , being least soluble and used in preparation of the
salts free from the other alkali metals.
234 RIBIDIIM— LITHIUM. §208,6.
6. Beactons. — In all its reactions similar to the other fixed alkalis.
7. Ignition. — Caesium salts color the non-luminous flame violet. The spec-
trum g-ives two sharply defined lines, Cs a and Cs /i, in the blue and a third
faint line in the orange-red Cs > , also several faint lines in the yellow and
green. With the spectroscope three parts of CsCl may be detected in presence
of 300,000 to 400,000 parts KCl or NaCl; and one part in presence of 1,500,000
parts LiCl (Bunsen, Pogg,, 1875, 155, 633).
8. Detection. — By the spectroscope (7 and §210, 7).
9. Estimation. — (!) As the double platinum chloride; (2) as the chloride with
BbCl , estimation of the amount of CI and calculation of the relative amounts
of the metals: (J) as the sulphate obtained from ignition of the acid tartrate
and treatment with H^SO^ (Bunsen, Pogg., 1863, 119, 1).
Bnbidimn. Eb = 85.4 . Valence one.
1. Properties.— *Spen7fc gravity, 1.52 (Bunsen, A., 1863, 125, 367). Melting
point, 38.5**; at — 10** soft as wax. A lustrous silver-white metal with a tinge of
yellow, oxidizes rapidly in the air, developing much heat and soon igniting.
Volatile as a blue vapor below a red heat. The metal does not keep well
under petroleum, but is best preserved in an atmosphere of hydrogen. Next
to caesium it is the most electro-positive of all metals.
2. Occurrence. — Widely distributed in small quantities, usually with caesium,
and frequently with the other alkali metals, always in combination. None of
the alkali metals can occur free in nature.
3. Prepe ration. — From the mother liquor obtained in the preparation of Li
salts (Heintz, J. pr,, 1862, 87, 310): (/) By ignition of the acid tartrate with
charcoal: (2) electrolysis of the chloride; (S) by ignition with Mg or Al
(Winkler, B., 1S90, 23, 51; Beketoff, B., 1888, 21, c, 424).
4. Oxide and Hydroxide. — The oxide BbsO has not been with certainty pre-
pared. The hydroxide, BbOH , is formed when the metal is decomposed by
water; also through the action of Ba(0H)2 upon BbzSO^ . It is a gray-white,
brittle mass, melting under a red heat.
5. Solubilities. — The metal dissolves in cold water, in acids and in alcohol
with great energy, evolving hydrogen. The hydroxide is readily soluble in
t\'ater with generation of heat. The salts are all quite readily soluble. The
acid tartrate is about eight times less soluble than the corresponding Cs salt.
Among the less soluble salts are to be mentioned the perchlorate, the fluosili-
cate, the double platinum chloride, the silicotungstate, the picrate, and the
phosphomolybdate. The alum is less soluble than the corresponding potassium
alum.
6. Reactions. — Similar to the other fixed alkalis.
7. Ignition. — The salts give a violet color to the flame. The spectrum gi^^^
two characteristic lines in the violet, Bb a and Bb /?; two less intensive in ^^^
outer red, Bb >^and BbJ; a fifth Bbf in the orange; and many faint lines in the
orange, yellow and green. As small a quantity as 0.0000002 gram of BbCl can
be detected (Bunsen, I.e.).
8. Detection. — By the spectroscope (7 and §210, 7).
9. Estimation.— (/) By weighing with CsCl as the chlorides, determining the
amount of CI and calculating the proportion of the metals; (2) as the double
platinum chloride.
§210. Lithium. Li =: 7.03 . Valence one.
1. Properties.— iSfpfri/fc gravity, 0.5936, the lightest of all known solid bodies
(Bunsen and Matthiessen, A., 1855, 94, 107). Melting poitU, 180*»; does not vaporize
at a red heat. It is a silver- white metal with a grayish tinge; harder than
K or Na but softer than Pb , Ca or Sr; it is tough and may be drawn into
wire and rolled into sheets. It is more electro-positive than the alkaline earth
metals but less electro-positive than K or Na . The pure metal is quite similar
§210, 8. LITHIUM, 235
in appearance and in its chemicol properties to K and Na , but does not react
so violently as those metals. It does not ignite in the air until heated to 200°,
and then burns quietiy with a very intense white light. It also burns with
vivid incandescence in CI , Br , I , O , S and dry COj . It decomposes water
readily, forming LiOH and H , but not with combustion of the hydrogen or
ignition of the metal.
2. Occurrence. — It is a sparingly but widely distributed metal. Usually pre-
pared from lepidolite, triphylene or petalite. Traces are found in a great
many minerals, in mineral springs, and in the leaves and ashes of many plants;
e. g., coffee, tobacco and sugar-cane.
3. Preparation. — It is prepared pure only by electrolysis, usually of the
chloride. A larger yield is obtained by mixing the LiCl with NH4CI or KCl
(Giintz, C. r., 1893, 117, 732). The metal is also obtained by ignition of the
carbonate with Mg , but the metal is at once vaporized and oxidized.
4. Oxide and Hydroxide. — It forms one oxide, Li^O , by heating the metal
in oxygen or dry air; cheaper by the action of heat upon the nitrate. The
corresponding hydroxide, LiOH , is made by the action of water upon the
metal or its oxide; cheaper by heating the carbonate with calcium hydroxide.
5. Solubilities. — The metal is readily soluble in water with evolution of
hydrogen, forming the hydroxide; soluble in acids with formation of salts.
The oxide, LijO , dissolves in water, forming the hydroxide. The most of the
lithium salts are soluble in water. A number of the salts, including the
chloride and chlorate, are very deliquescent. The hydroxide, carbonate and
phosphate are less soluble in water than the corresponding compounds of the
other alkali metals. In this respect lithium shows an approach to the alkaline
earth metals. LiOH is soluble in 14.5 parts water at 20** (Dittmar, J. 80c. Ind.,
1888, 7, 730); Li^CO, in 75 parts at 20**; Li,P04 in 2539 parts pure water and
3920 parts ammoniacal water, more soluble in a solution of NH4CI than in
pure water (Mayer, A., 1856, 98, 193).
6. Reactions. — Lithium salts in general react similar to the corresponding
potassium and sodium salts. They are as a rule more fusible and more easily
decomposed upon fusion. Soluble phosphates precipitate lithium phosphate,
more soluble in NH^Cl solution than in pure water (distinction from mag-
nesium). In dilute solutions the phosphate is not precipitated until the solu-
tion is boiled. The delicacy of the test is increased by the addition of NaOH,
forming a double phosphate of Na and Li (Rammelsberg, A, Ch., 1818, (2), 7,
157). The ])hosphate dissolved in HCl is not at once precipitated by neutraliz-
ing with NH4OH (distinction from the alkaline earth metals). Nitrophenic
acid forms a yellow precipitate, not easily soluble in water.
7.' Ignition. — Compounds of lithium impart to the flame a carmine-^re(f color,
obscured by sodium, but not by small quantities of potassium compounds.
Blue glass, just thick enough to cut off the yellow light of sodium, transmits
the red light of lithium: but the latter is intercepted by a thicker part of the
blue prism, or by several plates of blue glass. The spectrum of lithium con-
sists of a bright red band, Li «, and a faint orange line, Li /3. The color
tests have an intensity intermediate between those of sodium and potassium.
8. Detection. — B// the spectroscope. — To the dry chlorides of the alkali metals
a few di'ops of HCl are added and the mass extracted with 90 per cent alcohol.
The solution contains all the rare alkalis and some Na and K . Evaporate to
dryness, dissolve in a small amount of water and precipitate with platinum
chloride. The double platinum and potassium chloride is more soluble than
the corresponding salt of Bb and Cs . Boil repeatedly with small portions of
water to remove the potassium, and frequently examine the residue by the
spectroscope as follows: Wrap a small amount of the precipitate in a moistened
filter paper, then in a platinum wire and carefully char. After charring is
complete, ignite before the spectroscope. The K spectrum grows fainter, that
of Rb and Cs appear.
Evaporate to dryness the filtrate from the precipitate of the platinum double
salts, add oxalic acid and ignite, moisten with HCl, evaporate and extract with
absolute alcohol and ether. Upon evaporation of the extract LiCl is obtained,
almost pure. Test with the spectroscope and by forming the Insoluble phos-
phate.
236 DIRECTIOXS FOR ANALYSIS WITH NOTES. §810»9.
9. Estimation. — After separation from other elements it may be weig-hed as
a sulphate, carbonate or phosphate, LigPO^ . It may also be estimated by the
comparative intensity of the lines in the spectroscope (Bell, Am,, 1886, 7, 35).
DiBECTIOXS FOR THE ANALYSIS OF THE MeTALS OF THE AlKALI GROFP
(Sixth Group).
§211. If the material is found not to contain magnesium, the clear
filtrate from the carbonates of Ba , Sr , and Ca , after testing for traces
with (NHJaSO^ and (NH4)2C204 (§193), may at once be tested for the pres-
ence of potassiam and sodiiim. If magnesium be present it should be
removed in order to test for small amounts of sodium. Potassium and
large amounts of sodium may be readily detected in the presence of mag-
nesium. It is evident that the magnesium must not be removed by the
usual reagent used to detect the presence of that element, t. e. Na^HPO^ .
It is recommended by many to use ammonium phosphate. (NHJ^HPO^ .
This reagent removes the magnesium, and permits the application of the
flame test for the fixed alkalis; but the presence of the phosphate obstructs
the gravimetric determination of the alkalis. The phosphate may be
removed by lead acetate and the excess of the lead by hydrogen sulpliide.
§212. As a better method it is directed to evaporate the filtrate con-
taining the magnesium and the alkalis to dryness, ignite gently to remove
the ammonium salts. Dissolve the residue in water and add BafOH)^ to
precipitate the magnesium as Mg(0H)2 (§§177 and 182). After filtration,
the excess of barium in the filtrate is removed by H^SO^ , and the filtrate
from the barium sulphate is ready to be tested for the fixed alkalis by the
flame test or by gravimetric methods as may be desired. The presence of
sodium obscures the flame reaction for potassium, but the introduction
of a cobalt glass (§132, 7) or an indigo prism cuts out the sodium flame
and allows the violet potassium flame to be seen. Study 6, 7, 8, and 9 of
§§205 and 206.
§213. The free use of ammonium salts during the process of analysis
makes it necessary that the testing for ammonium be done in the original
solution or in the filtrate from the Tin and Copper Group.
Add an excess of KOH or NaOH to the solution and warm s^entlv. Notice
the odor (§207, 1). Suspend a piece of moistened red litmus paper in
the test-tube; in the presence of ammonia it will be changed from red
to blue color. To detect the presence of small amounts of ammonium
salts, heat the strongly alkaline mixture nearly to boiling and pass the
evolved gas into water. Test this solution (ammonium hydroxide) with
Xessler's Reagent (§207, Qk) or by the precipitation with HgCL (§207, 6/)^
Study §207, 6, 7, 8, and 9.
§214. The rare metals of the Alkali Group: lithium, rubidium, and
§2J5. DIRECTIONS FOR ANALYSIS WITH NOTES. 237
caesium, are rarely met with in the ordinary analyses. If their presence
is suspected they are tested for and detected by the spectroscope (7, §§208,
209 and 210).
§216. Lithium, because of the insolubility of its phosphate (§210, 5c),
interferes with the detection of magnesium. If the filtrate after the
removal of barium, strontium, and calcium be evaporated to dryness and
gently ignited to remove all ammonium salts; the residue, dissolved in
water and treated with an excess of barium hydroxide, will give a precipi-
tate of the magnesium as the hydroxide, leaving the lithium in solution.
The barium hydroxide precipitate may be tested for magnesium and from
the filtrate the excess of barium hydroxide may be removed by sulphuria
acid before testing for the alkali metals.
PART III.-THE NON-METALS.
§216. Balancing Equations in Oxidation and Reduction.
Statement of Bonds in Plus and Minns Nnmbcrs,* according to chemical
polarity, positive and negative (see §3 footnote).
In the terms of this notation the plus bond is the unit of Oxidation
and the minus bond is the unit of Reduction.
A bond, that is a unit of active valence, is either a plus one or a minus
one. The formula of a molecule of hydrochloric acid is stated, H+^Cl"^.
That of water, (H+^)oO~". (The plus sign is understood when no sign is
written before the valence number.)
Plus and minus bonds are represented as positive and negative quan-
tities. In the formula of hydrochloric acid, as above, the difference
between the polarity of the hydrogen atom and that of the chlorine atom
is stated as a difference of two.
In any compound the sum of the plus bonds and the minus bonds of the
atoms forming a molecule is zero.
Free elements, not having active valence, have zero bonds in this
notation.!
The Oxidation of any element is shown by an increase, and its Reduction
by a decrease, in the sum of its bonds.
When one substance reduces another the element which is reduced
loses as many bonds as are gained by the element which is oxidized.
It is evident that, changes in valence being reciprocal in oxidation and
reduction, there is no gain or loss in tlie sum of the bonds of two elements
which act upon each other.
The use of this notation is illustrated in the following equations:
3SnCl, 4- HaSO, + 6HC1 = 3SiiCl< + H,S + 3H:,0
In this equation the three atoms of tin gain six bonds; the bonds of the
sulphur in the H2SO.., have then been diminished by six; that is, it ha^^
given up six bonds to the tin, and having only four in the first place must
now have minus two (4 -6 = -2).
*0. C. Johnson, C. 2V., 1880, 42, 61. See also Ostwald, Qrundr, (tilg. Chem,, 8te Aufl., 1899, 8. 439.
tif there Is polarity in the union of like atoms with each other in forming an elemental
molecule, the sum must be zero, as in the for^lation of the molecules of compounds.
§217, f. BALANCING OF EQUATIONS. 239
3SnCl, 4- HIO, 4- 6HC1 = 3SnCl, + HIH- 3HaO
Here also the three atoms of tin gain six bonds, and these are furnished
by the iodine of the HIO3 . It has five in the first place, and being
diminished by six, has one negative bond remaining (5 -6 = -1). [In
other words, unless we deny that iodine has five bonds in HIO3 , we must
^dmit that it has one negative bond in HI (written H'l"^).]
8HMiiO« 4- 5AsH, 4* 8H2SO4 = SHjAsO* 4- SMnSO* 4- 12HaO
In this equation eight atoms of manganese in the first member have 5G
bonds, and a like amount in the second member has only 16, losing 40,
and this 40 has been gained by the five atoms of arsenic. They now have
25, after gaining 40. They must then have had — 15 in the first place
(25 — 40 = -15). That is, the atom of arsenic in arsenous hydrido ha*
-3 bonds {Asr"1L^).
SnCl, 4- HgCl, = Hg 4- SnCl*
This equation illustrates the statement that free elements have no
bonds. The tin gains two bonds, and these two bonds are taken from the
mercury in the HgCIs .
§217. Rule for Balancing Equations.
The number of oxidation bonds which any element has is determined
by the following rules :
a. Hydrogen has always one positive bond.
b. Oxygen has always two negative bonds.
c. Free elements have no bonds.
d. The sum of the bonds of any compound is zero.
e. In salts the bond of the metal is always positive.
/. In acids and in salts the acid radical has always negative bonds.
Thus, the bond of free JPb is zero, but in PbClj the lead has two posi-
tive bonds, and each atom of chlorine has one negative bond.
In BijSg , each atom of Bi has three positive bonds (e), and each atom of
S has two negative bonds (/).
In ammonium nitrite, NH^NO^ , or H4= N — — N = , the nitrogen
of the NH4 has four negative bonds and one positive bond. The other
nitrogen, that of the acid radical NOg , has three positive bonds. Each
atom of hydrogen has one positive bond and each atom of oxygen two
negative bonds, the sum being zero : +~t — 4 + 1 + 3 — 4 = 0.
In the following salts, etc., the bond of each element is marked above,
with its proper sign, plus being understood if no sign is given. Then f ol-
240 BALANCING OP EQVATIONSi. §218,1.
lows the equation in full, the bonds of each atom being multiplied by the
number of atoms^ and all being added, the sum is ^een to be aero.
Hg^(NV0-^,),.2 -f 10 — 12 =
Bi%(SviO-",),.6 + 18 — 24 =
Ba''(MiiviiC>-^^)2.2 H- 14 — 16 =
Pe'"(Nvo-",),.3 + 15 — 18 =
As"%S-'',.6 — 6 =
If the above is understood, the rule for balancing equations is easily
explained.
The number of bonds changed in one molecule of each shows the number
of the molecules of the other which must be taken, the words each and
other referring to the oxidizing and reducing agents.
§218. A few equations will illustrate the application of the rule.
(i) 3A84 + 20HNO, -h 8H,0 = i2H,AsO« + 20NO
The arsenic in one molecule gains 20 bonds, therefore 20 molecules of HNO,
are taken. The nitrogen loses three bonds, therefore three molecules of As,
are taken.
(2) OSb -f lOHNO, = 3SbjO» + lONO -|- SHjO
The antimony gains five bonds, therefore five molecules of HNO, would be
taken, and since the nitrogen loses three bonds, three of antimony would be
taken, but since we cannot write SbjOs with an odd number of atoms of
antimony, we double the ratio and take six and ten.
{3) 3H2S -f 8HNO3 = aHjSO, + 8N0 -h 4H2O
The S in the first member has 2 negative bonds (a and d); in the second
member it has 6 positive, gaining 8 bonds: hence 8 molecules of HNO3 must
be taken. The nitrogen in the first member has five bonds, and in the second
it has two. The difference is three, therefore just three molecules of HjS
must be taken.
Further, the reaction may be explained as follows:
The sulphur in the first member has two bonds (valence of two), but nega-
tive because combined with hydrogen (two atoms) to form a definite cou^-
pound; in the second member it has six bonds (valence of six), but positive
because combined with oxj'gen (SO, or Sq — ^ -—O^' '^^^ valence of the
hj'drogen does not change and hence in the reaction one molecule of H.S
gains eight bonds. The nitrogen in the first member has five bonds (valence
of five), but positive because combined with oxygen (N.O^ or H — O — N~^);
in the second member it has two bonds, still positive because combined with
oxygen. The valence of the hydrogen and oxygen does not change, hence in
the reaction one molecule of HNO, loses three bonds. Now the number of
bonds gained by the H^S (8) must equal the bonds lost by the HNOj (•?)•
The least common multiple, twenty-four, indicates the least possible totnl
change of valence for each compound: this requires that three molecules of
H;S and eight of HNO, be taken, giving for the products three molecules of
H2SO4 and eight of NO with four of water to complete the equation.
(4) 3SbaS, -f 28HNO3 = 3Sb,0a + 9H,S0, + 28NO -f 5H,0
In this case, both the Sb and the S in the molecule gain bonds, and must be
§218, 12. BALAXCIXG OF EQUATIONS. 241
considered. It is plain (from d and e) that each atom of Sb gains 2 bonds, and
the two in the molecule will gain 4.
The S in Sb^S, has 2 negati\e bonds, and in the second member (in HjSOt)
it has 6 positive bonds, a gain of 8. The three atoms in the molecule will gain
three times eight, or 24 bonds; to this add the 4 which the Sb has gained, and
we have 28 bonds gained by one molecule of SbjSs; hence 28 molecules of HKO,
must be taken. We take 3 of SbsS, for reasons explained in the first equation.
Further explain as follows: In this case both the Sb and the S gain in
valence (oxidized). Each atom of antimony gains two bonds, a total gain of
four. Each atom of sulphur gains eight, a total gain of twenty-four; or n
gain for one molecule of SbzS, of twenty-eight bonds. As in the previous
illustration, the nitrogen loses three bonds. The least common multiple,
eighty-four, indicates that for the reaction each compound must undergo a
change of at least eighty-four bonds. This requires for the Sb^Sg three mole-
cules, and for the HNOs twenty-eight molecules. The products are as indicated
in the equation.
(5) 2Ag,A80« + llZn -h llHjSO^ = 2AsH, + 6Ag + llZnSO^ -h 8H,0
The silver loses three bonds, and the arsenic in changing from plus five to
minus three loses eight bonds; this added to the three that the silver loses
makes eleven, therefore eleven molecules of zinc are taken, and since the zinc
gains two, two molecules of silver arsenate are taken,
(6) 2MnO + 5Pb,0, -f 30HNO, = 2HMiiOt -f 15Pb(N0.), -f 14H,0
The manganese gains five bonds, therefore five molecules of PbjO^ are taken.
The three atoms of lead in one molecule of PhjO^ have in all eight bonds, but
a like amount has only six in the second member, being a loss of two, there-
fore two molecules of MnO are taken.
(7) 2MnBr, -f TPbO^ -f HHNO, = 2HMnO, + 2Br2 + 7Pb(N0,)j -f CH^O
The manganese gains five bonds and the bromine gains one, the two atoms
gaining two, adding this to the five that the manganese gains makes a total
gain of seven bonds, therefore seven of PbOj are taken. The lead loses two,
therefore two of MnBr, are taken.
(8) MnS -f 4KN0, + K,CO, , fusion = K.MnO, -f K^SO, -f 4N0 + K.CO3
The manganese gains four bonds and the sulphur eight, making twelve;
therefore twelve of KNO^ would be taken, and since the nitrogen loses three
bonds, three of MnS would be taken, but since three is to twelve as one is to
four, the latter amounts are taken.
(9) 2Cr(0H), + 3Mn(NO,)3 + SK^CO, , fusion = 2X^010, +
3K,MnO« -f 6N0 -f 5C0, -f SHjO
This is a peculiar and instructive equation. The nitrogen loses six bonds, but
since the manganese in the same molecule gains four, the total loss is only two,
therefore two of Cr(0H)3 are taken. The chromium gains three, therefore
three of Mii(N08)2 are taken.
{10) 3Ag -f 4HN0, = 3AgrN0, + NO -f 2H,0
The rule here calls for three of silver and one of nitric acid, but three more
of unreduced nitric acid are needed to combine with the silver, making four
in all.
(11) 2PeI, + 6H,S0« , cone, hot = Fe,(SO0a + 3S0, + 21, + 6H,0
The rule here calls for two of Felj and three of H^SOa , but three more of
H,S04 that are not reduced are needed to combine with the iron, making six
in all.
(12) 3HN0, 4- 8A1 + 8K0H = 3NH, + 8KA10, + HjO
The nitrogen has five bonds in HNO, . and in NH, it has minus three,
losing eight, therefore eight of aluminum are taken. The aluminum gains
three, therefore three of HNO, are taken.
242 BALAyCINO OF EQVATIOyS, §218,25.
(13) 3BiONO, + llAl + llKOH = 3Bi + 3NH, + llKAlO, + H,0
The bismuth loses three bonds and the nitrogen loses eight, therefore eleven
of aluminum are taken; the aluminum gains three, therefore three of the
BiONOa are taken.
(U) MnO, + 4HC1 = MnCl, + CI, + 2HaO
The manganese loses two bonds and the chlorine gains one, but two more of
unoxidized HCl are needed to combine with the manganese, hence four are
taken.
(13) 2CrI, + 64KOH -f 27C1, = 2K,CrO, + 6KI0« + 54KC1 + 32H,0
The chromium gains three bonds and the iodine (in the molecule) gains
twenty-four, therefore twenty-seven of CI, are taken and the CI, loses two,
therefore two of CrI, are taken.
This rule holds good in organic chemistry when all the products of the
reactions are known, as the following examples will illustrate:
CH4 C--*H^. -4+4 =
CH.Cl C-»+'H'«Cl-^ —34-14-8—1 =
CHaCl, C-» + »H^Cl-',. -24-24-2 — 2 =
CHCl. C--» + »H^Cl-',. -14-84-1 — 8 =
CC1« OCl-'4. 4—4 =
HCtHsO, H'(C,) + »-»H',0-«,. 14-8-84-8 — 4 =
C.H,0 (O^-'H^O"'. 1-64-6-2 =
C.HsO. (C,)-» + »H^O-«,. -6 + 84-8 — 6 =
C«H„Oe (a)-*'-^'H'„0-«e. -7+7 + 12—12 =
(1) CH, + 4C1, = CCl, + 4HC1
The carbon is oxidized by the chlorine from negative four to positive four,
a polarity change of eight units, hence take eight molecules of chlorine; each
molecule of chlorine loses two bonds, take two molecules of methane. Two is
to eight as one is to four.
(2) 3C,H«0 + 2K,Cr30T + 8H,S0, = 3HCA0, + 2K2SO, +
2Cr,(S04), + 11H,0
The carbon of the alcohol while possessing a valence of eight, has an oxida-
tion valence of but four (minus four bonds); in the acetic acid the two atoms
of carbon have zero bonds, that is, the combinations with negative affinity
exactly equal the combinations with positive affinity; therefore take fouV
molecules of the potassium dichromate. The two atoms of the chromium lose
six bonds, take six molecules of the alcohol. Six is to four as three to two.
Eight molecules of sulphuric acid are necessary to combine with the potassium
and the chromium.
(3) aCjH.Oa + 14HN0, = 9C0, + 14N0 + 19H,0
The three atoms of the carbon in the glycerine have minus two bonds (the
negative affinity is two more than the positive affinity), and in the CO, a like
amount has twelve bonds, a gain of fourteen. The nitrogen loses three bonds.
(//) C.,Hi,0., + 12H,S0» = GCO, + 12S0, + ISH.O
The carbon in the dextrose has zero bonds (equal positive and negative
affinity combinations) and gains twenty-four bonds, while the sulphur loses
two bonds. The lower ratio is one to twelve.
For convenience of reference tlie non-metallic elements will be de-
scribed in llie order of their atomic weights; and the acids in the order
of the decree of oxidation of the characteristic element, e. g., N before S .
HCl before HCIO , HCIO, before HCIO, , (»tc.
§219,6. HTDROOEN, 243
§219. Hydrogen. H = 1.008 . Valence one.
1. Properties. — An odorless, tasteless gas. It is the lightest body known:
One litre at 0**, 760 mm. atmospheric pressure, weighs 0.08952289 gram (one
crith); specific gravity, 0.06949 (Crafts, C. r., 1888, 106, 1662). It is used for
filling balloons; also illuminating gas, containing about 50 per cent of hydrogen,
is frequently used because it is much cheaper. It is a non-poisonous gas, but
causes death by exclusion of air. It has been liquified to a colorless trans-
parent liquid by cooling to — 220° under great pressure and then allowing to
expand rapidly (01sze\yski, C. r., 1884, 99, 133; 1885, 101, 238; Wroblewski, C. r.,
1885, 100, 979). Critical temperature^ — 234.5**; critical pressure, 20 atmospheres;
tmling point, —243.5° (Olszewski, Phil, Mag., 1895, (5), 40, 202). It diffuses
through walls of paper, porcelain, heated platinum, iron, and other metals
more than any other gas (Cailletet, C. r., 1864, 58, 327 and 1057; 1865, 60, 344;
1868, 66, 847). It is absorbed by charcoal and by many metals, especially
palladium; which, heated to 100° in an atmosphere of hydrogen and then
cooled in that atmosphere, absorbs at ordinary temperatures 982.14 volumes of
hydrogen (Graham, J, C, 1869, 22, 419). This occluded hj'drogen acts as a
strong reducing agent, reducing FeCl. to FeCla , HgCl, to Hg° , etc. It is a
better conductor of sound than air (Bender, B,, 1873, 6, 665). It conducts heat
seven times better than air or 480 times poorer than iron (Stefan, C. C, 1875,
529). It refracts light more powerfully than any other gas and about six
times more than air. It bums with a non-luminous fiame and with generation
of much heat (more than an equal weight of any other substance or mixture
of substances). Hydrogen forms two oxides: water, HjO , and hydrogen
peroxide, H3O3 (§244).
2. Occurrence.— In volcanic gases (Bunsen, Pogg., 1851, 83, 197). In pockets
of certain Stassfurt salt crystals (Precht, B., 1886, 19, 2326). As a product of
the decay of organic material, both animal and vegetable. In combination as.
water and in innumerable mineraJs (H^O and OH) and in organic compounds.
3. Formation. — (a) By the reaction of alkali metals with water, (h) By
the action of superheated steam upon heated metals or glowing coals
(§226, 4a). (c) By dissolving aluminum or certain other metals in the
fixed alkalis, (d) By the action of many metals with dilute acids (seldom
HNO3). By heating potassium formate or oxalate with KOH : "KnC^O^ +
2K0H = 2K0CO3 + H2 (Pictet, A. Ch., 1878, (0), 13, 216).
4. Preparation. — (a) By the action of dilute sulphuric acid (one to
eight) on commercial or platinized zinc * (§136, 5a). The solution must
be kept cold or traces of SO2 and HoS will be evolved, (h) By the elec-
trolysis of acidulated water.
5. Solubilities. — Water at ordinary temperature dissolves nearly two per
cent (volume) of hydrogen. Chdrcoal dissolves or absorbs fully ten times its
volume of the gas (1).
6. Beactions. — Hydrogen gas is a very indifferent body at ordinary tem-
perature, combining with no other element except as it is occluded or ab-
sorbed by palladium, platinum, iron, nickel, etc.; and in the sunlight combines
with chlorine and bromine. "Nascent hydrogen " (hydrogen at the moment
of its generation), however, is a powerful reducing agent, and under proper
• For the rapid generation of hydrogrcn the zinc should be granulated by pouring the molten
metal Into cold water. Chemically pure zinc is very slowly attacked by dilute sulphuric acid;
but the commercial zinc frequently contains sufficient impurities to insure a rapid generation
of hydrogen when treated with the dilute acid. By the addition to the granulated zinc, in a tub
of water, of a few cubic centimeters of a dilute solution of platiflum chloride ; the zinc is made
readily soluble in dilute sulphuric acid and a uniform and rapid generation of hydrogen can be
obtained.
2U HTDROGEX. §819, T.
conditions combines with O , S , Se , Te , CI , Br , I , N , P , As , Sb and Si
with comparative readiness. The reduction of salts by nascent hydrogen Id
acid or alkaline solution will not be discussed here. See under the respective
elements. It should be noted, however, that " nascent hydrogen " generated
by different methods does not possess the same reducing properties. Sodium
amalgam with acids does not give hydrogen capable of reducing silver halides;
the reduction is rapid when zinc and acids are used. Neither electrolytic
hydrogen nor that from sodium amalgam and acids reduces chlorates; while
zinc and acids reduce rapidly to chlorides. Hydrogen generated by KOH and
Al does not reduce AsV; that formed by zinc and acids gives AsH, . 8V
with sodium amalgam and acids gives Sb°; with zinc* and acids, SbH^ (Cha*
brier, C. r., 1872, 75, 484; Tomraasi, BL, 1882, (2), 38, 148).
Hydrogen occluded in metals as Pd , Pt , etc., is even more active than
" nascent hvdrogen "; often causing combination with explosive violence
(Berthelot, A. Ch., 1883, (5), 30, 719; Berliner, W. A., 1888, 35, 781). Hydrogen
absorbed by palladium precipitates Ag , Au , Pt , Pd , Cu and Hg from their
solutions; permanganates acidified are reduced to Mn"; Fe'" to Fe'*; CrVi to
Cr'''; KCIO, to KCIO; CH.CO^H to CH,CHO and C,H,OH; and C.HsKO, to
C«H,NH2 . The reactions are quantitative. Salts of Pb , Bi , Cd , As , Sb , W.
Mo , Zn , Co , Ni , Al , Ce , TJ , Bb , Cs , K , Na , Ba , Sr and Ca are not reduced
(Schwarzenbach and Kritschewsky, Z., 1886, 25, 374). In the presence of
platinum black hydrogen reduces very much as described above; also KsFe(CN)«
becomes K4Fe(CN)c; dilute HNO3 becomes NH4NO, , concentrated HNO, be-
comes HNO2; CI, Br and I combine with the hydrogen in the dark; KCIO,
and KCIO are reduced to chlorides. KCIO4 is not reduced; H3SO4 , concen-
trated, is reduced to H.SO, (Cooke, C. A'.. 1888, 58, 103).
Free hydrogen very slowly acts upon a neutral solution of silver nitrate,
precipitating traces of silver; and in concentrated solution with formation of
AgNOj; hindered by HNO, or KNO, . Solutions of Au , Pt and Cu are also
acted upon (Rusself, J, C, 1874, 27, 3; Leeds, B., 1876, 9, 1456; Reichardt,, Arch.
Pharm., 1883, 221, 585; Poleck and Thuemmel, B., 1883, 16, 2435; Senderens, Bl,
1897, (3), 15, 991). KMnO^ in acid, neutral, or alkaline solution slowly
oxidizes hydrogen. It is not at all oxidized by nitrohydrochloric acid, in
diffused daylight, CrO, , at ordinarv temperature*^ FeCl, , KjFeCCN), , HNO, .
sp. or, 1.42,* or H.SO, , up, ijr, 1.84 ('Wanklyn and Cooper, Phil. Mag,, 1890, (5).
30, 431). In some cases, when hydrogen under ordinary conditions is without
action, if subjected to great pressure a reducing action takes place; e. »/..
hydrogen at 100 atmospheres pressure precipitates Hg** from HgCl, (Loewen-
thal, J, pr,, 1860, 79, 480).
7. Ignition. — Chlorine and bromine combine with hydrogen directly \v.
the sunlight, but heat is required to effect its combination with iodin.*.
fluorine, and oxygen.
All oxides, hydroxides, nitrates, carbonates, oxalates, and organic salt •
of the following elements are reduced to the metallic or elemental state h;*
ignition in hvdrogen gas: Pb , Ag , Hg , Sn , Sb , As , Bi , Cn , Cd, Pd.
Mo , Eu , Os , Rh , Ir , Te , Se , W , Pe , Cr , Co , Ni , Zn , Tl , m) , In , V .
Compounds of aluminum, manganese, and of the fifth and sixth grouj)
metals have not been reduced by hydrogen.
8. Detection. — (a) Method of formation if known. (6) Its explosive
union witli oxygen when the mixture with air is ignited, (c) Absorption
bv palladium sponge, (d) Explosive union with chlorine in the sunlight
to form HCl . (e) Separated from most other gases by its non-absorption
by the chemical reagents used in gas analysis.
9. Estimation.— By volume measurement, almost never by weight, except
'when determined in its compounds by combustion to HjO .
,4. BORON— BORIC ACID, 245
Boron. B = 11.0 . Valence three.
Boron does not occur free in nature. It is found chiefly as borax, Na^BfOr ,
and as boric acid, H|BOt , in volcanic districts. Two varieties of the element
have been prepared, amorphous and crystalline. The former is changed to the
latter by heating- to a white heat in presence of Al and C (Woehler and Claire-
Deville, .4.., 1867, 141, 268). Elemental boron is prepared (a) by electrolysis;
(6) by fusing B,0, with Al , Na or Mg-; (r) by igniting BCl, with hydrogen;
Id) by fusing borax with red phosphorus. Specific gravity of the crystalline,
2.53 to 2.68 (Hampe, A., 1876, 183, 75); of the amorphous, 2.45. Amorphou^
boron is a greenish-brown, opaque powder, odorless, tasteless, insoluble in
water, alcohol or ether. It is a non-conductor of electricity. Heated in air or
oxygen it burns with incandescence. In air it forms BjOj and BN . It is
oxidized by molten KOH or FbCrOf , with incandescence. It is dissolved by
concentrated HNO, or H2SO4 , forming boric acid. At a red heat it decom-
poses steam. When heated it combines directly with S , CI , Br , N and many
metals. It forms BCl, with chlorine, not BCI5 . Fused with V.O^ it forms
B3O, and P; with KOH , K,BO, and H; with KxCO. , KsBO. and C . Boron
forms but one oxide, B.jO, , boric anhydride. Three hydroxides are known:
2H,BO.^ = B^Os.SHaO , orthoboric acid; 2HBO3 = B,Os.H,b , metaboric acid; and
H^B^Ot = 2BaO».HaO , pyroboric acid.
Boric acid. H3BO3 = 62.024 .
B',B'"0-"3, H-0-B~J~^
1. PropertiOB. — Boron trioxide, B,0, , boric anhydride, is a brittle vitreous
mass; «p. gr. at 12*», 1.8476 (Ditte, A, Ch., 1878, (5), 13, 67). Melting point, 577**
(Carnelley, J, C, 1878, 83, 278). It is volatile at a very high heat (Ebelemen,
A, Ch., 1848, (3), 22, 211). It has a slightly bitter taste, is hygroscopic, and
shows a marked rise in temperature on solution in water (Ditte, C. r., 1877,
S5, 1069). In some respects boron tpoxide deports itself as a weak base. It
forms a sulphide, BjS, , decomposed by water (Woehler and Deville, A. Ch.,
1858, (3), 52, 90); a sulphate, B(HS04)s (D'Arcey, J, C, 1889, 55, 155); and a
phosphate, BPO4 (Meyer, B., 1889, 22, 2919). It combines with water in three
proportions, forming the ortho, meta and pyroboric acids. Orthoboric acid is
A weak acid, its solutions reddening litmus; at 12** it has a specific gravity of
1.5172 (Ditte, I.e.): melts at 184*' to 186*^ (Carnelley, I.e.). Soluble in 25 parts
water at 20°, and in 3.4 parts at 102** (Ditte, I.e.). It is volatile in steam and
in alcohol vapor. The evaporation of the water of combination of the acid
•carries with it from ten to fifteen per cent of the acid.
2. Occurrence. — Widely distributed, but usually in very small quantities. In
the rock salt deposits at Stassfurt, Germany, as boracite, MgTBisOsoCl, (62.5
per cent B3O,). In the volcanic regions of Tuscany and the Liparic Islands as
steam saturated with boric acid.
3. Formatioii. — The anhydride is formed by burning the metal in air
or oxygen, or by heating the acids. Orthoboric acid, H3BO3 , is formed
by dissolving the oxide in water; the meta acid, HBO^ , H — — B = •
by heating the ortho acid a little above 100** (Bloxam, J. C, 1860, 12,
177); the pyroboric acid, tetraboric acid, H2B4O7, by heating the ortho
or meta acid foi»6ome time at 160° in a current of dry air (Merz, J. pr.,
1866, », 179).
4. Preparation. — (a) By evaporation of the water from the lagoons of
'Tuscany, which Are saturated with boric acid, and recrystallization
246 BORIC ACID. §821,5.
from water. (5) In Nevada the boronatrocalcite, CdL^fi^^JStL^fi^ +
I8H2O (45.6 per cent B2O3), is evaporated in lead pans with KJBO^ to a
stiflf paste; and then treated with superheated steam in iron cylinders
heated to redness. The acid passes over with the steam and is collected
in lead lined chambers (Gutzkow, Z., 1874, 13, 457). (c) Commercial
borax, NasB^O^.lOHsO , is dissolved in hot water, twelve parts, and acidi-
fied with hydrochloric acid. Upon cooling, the boric acid, H3BO3 , is ob-
tained in small scales, which are purified by recrystallization from hot
water.
5. Solubilities. — More soluble in hydrochloric acid solution or in alcohol
than in water (1). The alcoholic solution bums with a beautiful green
fiame. Quite soluble in glycerine and in most alcohols and hydrocarbons,
only sparingly in ether. The borates are insoluble in alcohol; those of
the alkalis are soluble in water to an alkaline solution. Borates of the
other metals are insoluble in water (no borate is entirely insoluble in
water) ; but are usually rendered soluble by the addition of boric acid.
6. Beactions. — Silver nitrate forms, in solutions of acid borates, a white
precipitate of silver borate, AgBOs , but normal borates form in part silver
oxide, brown. Lead acetate gives a white precipitate of lead borate,
Pb(B02)2 ; calcium chloride, in solutions not very dilute, a white precipi-
tate of calcium borate; and barium chloride, in solutions not dilute, a white
precipitate of barium borate, Ba(B02)2 • With aluminum salts, the precipi-
tate is aluminum hydroxide.
Borates are transposed with formation of boric acid, by all ordinary
acids — in some conditions even by carbonic acid.
The liberated boric acid is dissolved by alcohol, and if the alcohol solu-
tion be set on fire, it burns with a green flame.
A solution of a borate, acidulated with hydrochloric acid to a barely
perceptible acid reaction, imparts to a slip of turmeric paper half wet with
it, a dark-red color, which on drying intensifies to a characteristic red color.
7. Ignition. — Boric acid is displaced from its salts by nearly all acids
including COo ; but being non-volatile except at a very high heat, it dis-
places most other acids upon ignition.
^y heating a mixture of borax, acid sulphate of potassium, and a fluo-
ride, fused to a bead on the loop of platinum wire, in the clear flame of
the Bunsen gas-lamp, an evanescent yellowish-green color is imparted to
the flame.
Borates fused in the inner blow-pipe flame with potassium acid sulphate
give the green color to the outer flame.
If a crystal of boric acid, or a solid residue of borate previously treated
with sulphuric acid, on a porcelain surface, is played upon by the flame of
Bunsen's Burner, the green flame of boron is obtained.
§222,1. CARBON, 247
If a powdered borate (previously calcined), is moistened with sulpiiuric
acid and heated on platinum wire to expel the acid, then moistened with
glycerine and burned, the green flame appears with great distinctness.
The glycerine is only ignited, then allowed to burn by itself. Barium
does not interfere (being held as sulphate, non-volatile) ; copper should be
previously removed in the wet way. The glycerine flame gives the spec-
trum. But in all flame tests, boric acid must be liberated.
Borates (fused on platinum wire with sodium carbonate) give a char-
acteristic spectrum of four lines, equidistant from each other, and extend-
ing from Ba ^ in the green to Sr d in the blue.
Borax, Na2B407 , when ignited (as on a loop of platinum wire to form
the borax bead) with many metallic compounds, forms a coloi*ed glass,
used in the detection of certain metals (§132, 7). The fused borax forms
a solid brittle mass, borax glass, used in assaying and in soldering because
of its power of combination with metallic oxides.
8. Detection. — By conversion into the acid, if present as a salt; solution
in alcohol or glycerine and burning with the formation of the green flame
(very delicate, but copper salts should be removed by HoS and barium salts
should be removed or converted into the sulphate). Also by the red color
imparted to a strip of turmeric paper.
9. Esttmation. — Boron compounds cannot be completely precipitated from
solution by any known reagents, hence most of the methods of quantitative
determination are indirect. By adding a known quantity of NaaCO, , fusing
and weighing; then after determining the CO^ subtracting its weight and
that of the Na^O present (calculated from NajCO, first added). The differ-
ence is the weight of BjO, present. See also Will (Arch, Phurm., 1887, 225, 1101).
In the presence of glycerine, boric acid may be accurately titrated with sodium
hydroxide, using phenolphthalein as an indicator: BjO, -f 2NaOH = 2NaB0a -|-
HjO . Sodium carbonate must be absent or we get: 2B3O, -f NaaCOj =
D'asB^Or Hh CO, (Honig and Spitz, Z, angew,, 1896, 549; Joergensen, Z, angetc,^
1897, 5).
§222. Carbon. C = 12^.0 . Usual valence four.
1. Properties. — Carbon exists in three allotropic forms: two crystalline,
diamond and graphite, and amorphous as charcoal, coke, etc. Specific gravity,
diamond at 4**, 3.51835 (Baumhauer, ,/., 1873, 237); graphite, Ceylon, 2.25 to 2.20
(Brodie, A., 1860, 114, 6); wood charcoal, 1.57; gas coke, 1.88. Very small
specimens only, of diamonds have been artificially prepared, by saturating iron
with carbon at 3000**. At this temperature graphite is formed and upon cool-
ing under pressure the crystalline diamond form is obtained. This cooling
under pressure is obtained by pouring the carbon saturated iron into a soft'
iron bomb, which is cooled by water (Moisson, C. r., 1893, 116, 218). Diamond
is the hardest substance known. It is very strongly refractive towards light
(Becquerel, A. Ch,, 1877, (5), 12, 5). Fluorescence and phosphorescence of
diamonds, see Kunz (C. C., 1891, ii, 562). Ignition in an atmosphere of hydro-
gen does not effect a change; in air or oxygen it burns to COa .
Graphite is a hard, gray, metal-like, opaque solid, a good conductor of
electricity and a fairly good conductor of heat. It burns with difficulty. It
248 CARBON. §222,2.
is used in lead pencils, in black lead (plumbago) crucibles, as a lubricant for
heavy machinery, in battery plates, for the arc light carbon pencils, ete.
Amorphous carbon is black, lighter than diamond or g^raphite. It is in use
as coal, coke, charcoal, animal charcoal, etc.; all impure forms. Lamp-black
is also amorphous carbon made from burning resin, fat, wax, coal gas, etc.,
with limited supply of air. It is used as a pigment in paints, in stove-black-
ing, shoe-blacking, printers' ink, etc. Charcoal, preferably animal charcoal,, is
used for decoloring organic solutions. Charcoal absorbs many gases, hence is
valuable as a disinfectant.
Carbon forms two oxides: carbon monoxide, CO , and carbon dioxide, CO, .
2. Occurrence. — Diamonds seem first to have been found in India, especially
in the Golconda pits, where, as early as 1022, 30,000 laborers are said to hate
been employed (Walker, J., 1884, 774). Also found in other parts of Asia, in
South Africa, in Brazil, etc. (Winklehner, C, C, 1888, 192; Damour, J., 1883, 774;
Gorceix, J., 1881, 345; Smit, J., 1880, 1400). Graphite is found in Ceylon (Wal-
ther, C. C, 1890, il, 20); in California (C. A\, 1868, 17, 209); in Canada (Dawson.
Am. 8,, 1870, (2), 50, 130); in New Zealand (Mac Ivor, C. iV., 1887, 55, 125);
in Russia, Germany, Greenland, etc. Pure amorphous carbon occurs in nature
as a chief product in the decomposition of organic material, air being excluded.
Anthracite coal is relatively pure amorphous carbon.
3. Formation. — Graphite remains as a residue when pig iron is dis-
solved in acids. It forms by reducing CO with FC3O4 at 400°. Amor-
phous carbon is formed by passing CCI4 over Na in a tube heated to red-
ness (Porcher, C. N., 1881, 44, 203).
4. Preparation. — Pure graphite is prepared by heating the commercial
graphite on a water bath with KCIO3 and H^SO^ and repeatedly washing.
If it contains SiOg it should also be treated with NaF and H^SO^ . Amor-
phous carbon is prepared by heating wood, coal, or almost any organic-
matter to a very high temperature in absence of air, but when so prepared
it is never pure. Amori)hous carbon is prepared approximately pure by
heating pure cane sugar in a closed platinum crucible; then boiling in
succession with HCl , KOH , and H^jO ; then igniting to redness in an
atmosphere of chlorine, cooling in the same atmosphere.
5. Solubilities. — Insoluble in water or acids. Soluble in many molten
metals with partial combination to form carbides. When the metal is
dissolved in acids the combined carbon passes off as hydrocarbons, the
excess remaining as graphite.
6. Beactions. — Xot attacked by acids or alkalis. It slowly oxidizes to
CO2 when heated with concentrated H..SO4 and KoCroO- . Upon gently
warming graphite with KGIO.^ and BSO^, graphitic acid, C^H^O^j , is
said to be formed (Stingl, B., 1873, 6, 391). The important reactions of
carbon require the aid of high heat and are described in the next
paragraph.
7. Ignition. — Unchanged by ignition in absence of air. When strongly
ignited in air or oxygen it slowly burns to COg . If the carbon and oxygen
have been previously very thoroughly dried the action is very slow,
especially with graphite. By fusion with KNO.^ or KCIO3 carbon is oxid-
ized to CO2 . With vapors of sulphur, carbon disulphide is formed; 1. p..
§223,2. ACETIC ACID, 249
by passing sulphur vapors over hot coals in a furnace. In an atmosphere
of hydrogen with the electric spark, acetylene, CjHg, is formed. By
igniting in an atmosphere of carbon dioxide, CO2 , the whole of the carbon
iKJComes carbon monoxide: C + CO^ = 2C0 .
By simple ignition with carbon, all oxides of the elements in the follow-
ing list are reduced to the elemental state (a); and if sodium carbonate is
added, all of the salts of the same are likewise reduced (&). Cu , Bi , Cd ,
Pb , Ag , Hg , As , 8b , 8n , Pd , Mo , Ru , Os , Eh , Ir , Te , Se , W , K ,
Na, Eb, Cr, Fc, Mn, Co, Ni, Zn, Ti, Tl .
(fl) Pb.O, -f 2C = 3Pb + 2C0,
(6) 2PbCl, -f 2Na,C0, + C = 2Pb + 4NaCl + 3C0,
(c) CuO + C (excess) = Cu + CO
(d) C + 2CuO (excess) = 2Cu + CO^
With excess of carbon CO is formed (c). With excess of the oxide CO2 is
formed (d). In the reduction of iron ore, the process is conducted so as
to give some CO and some COo . To obtain some metals in the free state
(such as K and Na), special methods are adopted to exclude the air, and
to produce the high temperature needed.
All compounds of sulphur when ignited with carbon are reduced to a
sulphide: BaS04 + 2C = BaS -f 2C0o .
8. Detection. — By its appearance ; failure to react with general reagents ;
and by its combustion to CO2 with oxygen (air), or with EoCTjOy and con-
centrated H28O4 (Fritsche, A., 1896, 294, 79), then by identification with
Ca(0H)2 (§228, 6).
9. Estimation. — By combustion to CO3 and weighing after absorption in KOH
solution. See works on ultimate organic analysis.
§223. Acetic acid. JLCJL^O., = 60.032 .
H
I II
H\(C2)+'"-'"0"% , H — C — C — — H = CH3CO2H.
1. Properties. — Pure acetic acid is a colorless, crystalline, hygroscopic solid,
melting at 16.5** and boiling at 118**. Its specific gravity at 0° is 1.080. It has
a sharp, sour taste, an irritating burning effect on the skin, and a very pene-
trating odor. It burns when heated nearly to the boiling point. Vinegar
contains four to five per cent of acetic acid. The U. S. P. reagent contains 36
per cent of acetic acid, and has a specific gravity of 1.0481 at 15**. It vaporizes
from its concentrated solutions at ordinary temperatures, having the char-
acteristic odor of vinegar. It is a monobasic acid, the three remaining
hydrogen atoms (linked to carbon) cannot be replaced by metals.
2. Occurrence. — It occurs in nature in combination with alcohols in the
essential oils of many plants.
250 ACETIC ACID. §223,3.
3. FormatioiL— (a) During the decay of many organic compounds. (6)
By gently heating sodium methylate, HaOCE,, in a current of carbon
monoxide: NaOCH, + CO = CHsCO^Va (HaCAOj). (c) By boiling
methyl cyanide with acids or alkalis: CH,CH + HCl + 2H,0 = HCfifi.
+ HH^Cl. (d) By the oxidation of alcohol: 30^0 + 2X^CtJlij +
8H2SO4 = 2K28O4 + 2Cr2(80j3 + 3HC^302 + 11H,0 .
4. Iteparation. — (a) By the drj- distillation of wood. (6) By the fer-
mentation of cider, beer, wine, molasses, etc. {c) Pure acetic acid is
prepared by distilling anhydrous sodium acetate with concentrated sul-
phuric acid. The distillate solidifies upon cooling and is termed glacial
acetic acid.
5. Solnbilities. — Misciblc in all proportions in water and alcohol. The
salts of acetic acid, acetates, are all soluble in water, silver and mercurous
acetates sparingly soluble. Certain basic acetates, as Fc'", Al , etc., are
insoluble in water. Very many of the acetates are soluble in alcohol.
(). Keactions. — The stronger mineral acids transpose the acetates,
forming acetic acid. Anhydrous acetates with concentrated sulphuric
acid give pure acetic acid (4), but if the sulphuric acid be in excess and
heat be applied the mixture blackens with separation of carbon; and, at
higher temperatures, COo and SO^ are evolved.
Solution of ferric chloride fonns, with solutions of acetates, a red solu-
tion containing ferric acetate, Fe(C2H302)3 , which on boiling precipitates
brownish-red, basic ferric acetate. The red solution is not decolored bv
solution of mercuric chloride (distinction from thiocyanate): but is de-
colored by strong acidulation with sulphuric acid or hydrochloric acid (dis-
tinction from thiocyanate and from meconate). The ferric acetate is pre-
cipitated by alkali hydroxides.
If acetic acid or nn acetate be warmed with solphnric acid and a little
alcohol, the characteristic pungent and fragrant odor of ethyl acetnte or
acetic ether is obtained :
HC.HaO^ 4- C.H^OH = H,0 4- C,H5C=H,0,
Acetic acid does not act as a Reducing Agent as readily as do most of
the organic carbon compounds. It does not reduce permanganates even in
boiling solution; reduces auric chloride only in alkaline solution, and does
not reduce alkaline copper solution. It takes chlorine into combination —
slowly in ordinary light, quickly in sunlight, forming chloracetic acids.
7. Ignition. — By ignition alone, acetates blacken, with evolution of
vapor of acetone, C3H,;0 , inflammable and of an agreeable odor. By pro-
longed ignition of alkali acetates in the air, carbonates are obtained free
from charcoal. By ignition with alkali hydroxides in dry mixtures,
methane, marsh-gas, CH^ , is evolved. By ignition with alkalis and
arsenous anhydride, the poisonous and offensive vapor of cacodyl oxide
S224. CITRIC ACID. 251
is obtained. This test should be made under a hood with great caution
and with small quantities. It is a very delicate test . for acetates :
4KC2H3O2 + A82O3 = A82(CH3),0 + 2K2CO3 + 2CO2 .
8. Detection. -(a) By its odor. (&) By the formation of the fragrant
ethyl acetate upon warming with sulphuric acid and alcohol, (c) By the
formation of the red solution with ferric chloride (§126, 6b and §162).
{d) By ignition of the dry acetate alone to acetone, CHaCOCHa ; with
NaOH to methane, CH4 ; or with As^Os to cacodyl oxide, {e) As a delicate
test for formates or acetates it is directed to warm a solution of CnClg in
NaCl and add a small amount of the material under examination. Form-
ates give a blackish-gray deposit; acetates give bright green precipitate
not changed by boiling. Both precipitates are soluble in acetic acid
(Field, J, C, 1873, 26, 575).
9. Estimatioii. — Other volatile acids are separated by precipitation; sulphuric
acid is then added and the acetic acid is distilled into water and estimated
by titration with standard alkali.
§224. Citric acid. HgCeH-O^ = 192.064 .
H^C — CO2H
!
S'3(Ce)+i^^H',0-% , H — — C — CO2H
I
H.C — CO2H
Found in small quantities in the juices of many fruits. The chief commercial
source is lemon-juice. It is a colorless, crystallizable, non-volatile solid; freely
soluble in water and in alcohol.
The citrates of the metals of the alkalis are freely soluble in water; those
of iron and copper are moderately soluble; those of the alkaline earth metals
insoluble. There are many soluble double citrates formed by action of alkali
titrates upon precipitated citrates, or of alkali hydroxides upon metallic salts
in presence of citric acid. In distinction from tartrates, the solubility of the
potassium salts, non-precipitation of calcium salt in cold solution; and weaker
reducing action, are to be noted.
Solution of calcium hydroxide in excess (as by dropping the solution tested
into the reagent) gives no precipitate with citric acid or citrates in the cold
(distinction from tartaric acid), but on heating, the white calcium citrate,
'Ca^(Ca'H.:,0j)2 . is precipitated (not soluble in cold ix)tassium hydroxide solu-
tion). By filtering before boiling, the tartrate and citrate may be approxi-
mately separated. Calcium chloride also gives the same precipitate after boil-
ing. Calcium citrate is soluble in acetic acid (distinction from oxalates).
Solution of lead acetate precipitates white lead citrate, '2h:^{CJS.:,0^)^ , soluble
in ammonia. Silver nitrate gives a white precipitate of «i7rer citrate,
AgsCnHsOr , which does not blacken on boiling (distinction from tartrate).
For action of citric acid or citrates in hindering many of the usual analytical
reactions, see Spiller, J. C, 1858, 10, 110.
One part of citric acid dissolved in two parts of water, and treated with a
solution of one part of potassium acetate in two parts of water, should remain
clear after addition of an equal volume of strong alcohol (absence of oxalic
Acid and of tartaric acid and its isomers).
252 TARTARIC ACID. §SS>>1.
Citric acid does not act very readily as a redudLng agent; does not reduce
alkaline cop{>er solution, or silver solution; reduces permangranate very slowly.
C*oncentrated nitric acid produces from it, acetic and oxalic acids; and digei-
tion with mang'anese dioxide decomposes it, with formation of acetone, acrylic
and acetic acids. Citrates carbonize on ignition, with various empyreumatic
products, and with final formation of carbonates. By fused potassium hydrox-
ide, short of ignition, they are decomposed with production of oxalate and
acetate.
§225. Tartaric acid. H^C ^0, = 150.048 .
H
I II
H — — C — C~0 — H CH(0H)C02H
I or I
H— — C — C — — H CR{011)C0JBi
1. Fropertiee. — Tartaric acid is a colorless, crystalline, non-volatile solid;
freely soluble in water and in alcohol. It exists in four distinct modifications:
dextrotartaric acid, levotartaric acid, raeemic acid, and mesotartaric acid.
They differ from each other in crystalline form, in solubility, and especially
in the deportment of their solutions towards polarized light. Bacemic and
mesotartaric acids are optically inactive, but the former may be resolved into
the first two acids, optically active.
2. Occurrence. — It is found in various fruits. The chief commercial source
fs grape juice.
3. Formation. — By oxidation of dextrose, cane sugar, milk sugar, starch, etc..
with HNO, (Kiliani, .4., 18S0, 205, 17.")). By action of sodium amalgam on
oxalic ether in alcoholic solution (Debus, A., 1S73, 166, 1*24). By synthesis
from succinic acid by formation first of the dilDromsuccinic acid, H2C4Br3H:04:
then substitution of the OH group for the bromine by means of water and
silver oxide.
4. Freparation. — The crude argol deposited during the fermentation of grape
juice is recrystallized, giving the commercial cream of tartar, KHC^H^Oa .
This in hot solution is treated with powdered chalk, and the filtrate from the
precipitate thus obtained is precipitated with calcium chloride. Both pre-
cipitates are washed and decomposed by the necessary quantity of hot dilute
sulphuric acid. The tartaric acid solution is evaporated to crystallization and
purified by recrystallization (Ficinius, Arch. Fharm,, 1879, 215, 14 and 310).
0. Solubilities. — The Tartrates of the alkali bases arc soluble in water;
the normal tartrates being freely 8olu])le, the acid tartrates of potassium
and ammonium sparingly soluble. The tartrates of the alkaline earth
bases and of the non-alkaline bases, are insoluble or sparingly soluble, but
mostly dissolve in solution of tartaric acid. Most of the tartrates are
insoluble. in alcohol. There are double tartrates of heavy metals with
alkali metals, which dissolve in water. Tartar-emetic is potassium anti-
mony tartrate, KSbOC^H.O,, .
Hydrochloric, nitric, and sulphuric acids transpose the tartrates
(whether forming solutions or not). Most of the tartrates are also dis-
§225, Sd. TARTARIC ACID. 253
solved (and, if already dissolved, are not precipitated) by the alkali hy-
droxides, owing to the formation of soluble double tartrates.
The freshly precipitated oxides, hydroxides, and carbonates of the fol-
lowing metals are soluble in a solution of potassium-sodium tartrate,
Rochelle salt: 8b , Sn^^ Bi , Cn , Fe , Al , Cr , Co , Ni , Mn , and Zn ; Ba ,
Sr, Ca, and Dig to quite an extent. CdCO, is not dissolved (Warren,
C. N., 1888, 67, 223).
6. Beactions. — Solution of calcinm hydroxide, added to alkaline reac-
tion, precipitates from cold solution of tartaric acid, or of soluble tartrates,
calcium tartrate, white, CaC^H^Oe . Solution of calcium chloride with
neutral tartrates gives the same precipitate. Solution of calcium sulphate
forms a precipitate but slowly, or not at all (distinction from racemic acid).
The precipitate of calcium tartrate is soluble in cold solution of potassium
hydroxide, precipitated gelatinous on boiling, and again made soluble on
cooling (distinctions from citrate), and dissolves in acetic acid (distinction
from oxalate).
Tartaric acid prevents the precipitation by fixed alkalis of solutions of
the baits of the following metals : Al , Bi , Co , Ni , Cr , Cn , Fe , Pb , Pt ,
and Zn (Grothe, J. /w., 1864, 92, 175).
Silver nitrate precipitates, from solutions of normal tartrates, stiver
tartrate, Ag^CJlfiQ , white, becoming black when boiled. If the precipi-
tate is filtered, washed, dissolved from the filter by dilute ammonium
hydroxide into a clean test-tube, left for a quarter of an hour on the
water-bath, the silver is reduced as a mirror coating on the glass (§59, 10b),
distinction from citric acid. Free tartaric acid does not reduce silver
salts. Permanganate is reduced quickly by alkaline solution of tartrates
(distinction from citrates), precipitating manganese dioxide, brown. Free
tartaric acid acts but slowly on the permanganate. Alkaline copper tar-
trate, Fehling^s solution (§77, 6b), resists reduction in boiling solution.
Chromates are reduced by tartaric acid, the solution turning green. The
oxidized products, both with permanganate and chromate, are formic
acid, carbonic anhydride, and water.
7. Ignition. — On ignition, tartaric acid or tartrates evolve the odor of
hurnt sugar, separating carbon, and becoming finally converted to carbon-
ates. — Strong sulphuric acid also blackens tartrates, on warming. Melted
potassinm hydroxide, below ignition, produces acetate and oxalate. The
fixed alkali tartrates ignited in absence of air give an alkali carbonate and
finely divided carbon. The mixture serves as an admirable flux for the
reduction tests for arsenic (§69, 7).
8. Detection. — (a) By the odor of burnt sugar when ignited, (h) By
the deportment of the calcium salt with cold and hot KOH (6). (c) By the
formation of the silver mirror (§59, 10b). (d) By its action as an alkali
254 CARBON MONOXIDE. §2tt|9.
tartrate in preventing precipitation of the solutions of the heavy metak
by the fixed alkalis. To test citric acid for. the presence of tartaric acid,
add about one cc. of ammonium molybdate solution to about one gram
of the citric acid; then two or three drops of sulphuric acid and warm
on the water-bath. The presence of 0.1 per cent or more of tartaric acid
gives a blue color to the solution (Crismer, BL, 1891, (3), 6, 23).
9. Estimation.— See Philippe (Z., 1890. 29, 577); Haas (C. C, 1888, 1045);
Heidenhain (Z., 1888, 27. 681).
Carbon monoxide. CO = 28.0 .
CO-", C = .
1. Properties. — Carbon monoxide, rarlxmic oxide, formic anhvdride, CO , is a
colorless, tasteless gas. i^itevifiv yrarityj 0.9678. By maintaining' a pressure of
200 to 300 atmospheres at — 136° and then reducing the pressure to 50 atmos-
pheres the gas becomes a colorless transparent liquid (Wroblewski and Ols-
zewski, A. Ch., 1884 (6), 1, 128). It is. when inhaled, a virulent poison, abstract-
ing oxygen from the blood and combining with the haemoglobin. It bums in
the air with a pale blue flame to COa , but does not support combustion.
Mixed with air in suitable proportions, it explodes upon ignition. It unites
with chlorine in the sunlight to form phosgene, COCl, .
2. Occiirrenee. — In combination as formic acid in ants and in nettles.
3. Formation. — (a) By the incomplete combustion of coal, charcoal or
organic material, (h) Yrom the reduction of metallic oxides in the bla^^t
furnace with excess of charcoal : Tt^O^^ -f 3C = 2Fe -f 3C0 . (r) By
heating sodium sulphate with excess of charcoal (LeBlanc's soda process):
Va.SO^ + 4C = Na.^S + 4C0 . See also Grimm and Ramdohr (.4., 18o<),
98, 127).
4. Preparation. — (a) By passing steam over charcoal at a white heat
(water gas): H^O + C = CO + Hj (Naumann and Pistor, B., 1885, 18,
164). (b) By passing COo over red hot charcoal, (r) By heating
X,Fe(CN)e with concentrated H.SO^: K,Fe(CN)o + GH.SO^ + GH^O =
2KoS0, + 3(NH,)2S04 + FeSO, + 6C0 . With dilute acid HCN is formed.
(d) By heating a formate with concentrated sulphuric acid: 2KCH0o +
HoSO^ = K2SO4 + 2C0 + 2H.0 . (e) By heating an oxalate with con-
centrated sulphuric acid: KoCjO^ + 2Ho80^ = KjSO^ + HjSO^.HoO +
CO + CO2 .
5. Solubilities. — It is not absorbed by KOH or Ca(0H)2 (distinction
from CO2). It is absorbed by charcoal, cuprous chloride, and by several
metals, e. g., K , Ag , and Au .
6. Beactions. — It is an energetic reducing agent. Combines with moist
fixed alkalis to form a formate (Froelich and Geuther, A,, 1880, 202, 31T).
In the sunlight it combines directly with chlorine or bromine. It is
oxidized to COo by warming with K^CraOy and concentrated H^SO^ ; also
§227,46. OXALIC ACID. 255
by palladium sponge saturated with hydrogen, and in presence of oxygen
and water (Remsen and Reiser, B,y 1884, 17, 83). A solution of PdCIg is
reduced to Pd by CO .
7. Ignition. — When heated to redness with Na or K, carbon and an
alkali carbonate are formed. Upon ignition of metallic oxides in an
atmosphere of CO a reduction of the metal takes place, so far as observed
the same as when the corresponding metallic forms are ignited with char-
coal (Rodwell, J. C, 1863, 16, 44).
8. Detection. — In distinction from COj by its failure to be absorbed by
KOH or Ca(0H)2 . By its combustion to COg and detection as such. By
its combination with hot concentrated EOH to form a formate. It is
detected in the blood by the absorption spectrum (Vogel, i?., 1878, 11,
235).
9. Estimation. — The measured volume of the gas is brought in contact with
a solution of cuprous chloride in hydrochloric acid which absorbs the CO
(Thomas, (7. JV., 1878, 37, 6).
§227. Oxalic acid. H^CA = 90.016.
II II CO,H
H'2(C2)+«0-"',,H — — C — C — — H or |
CO2H
1. Properties. — Absolute oxalic acid, HjCsO^ , is a white, amorphous solid,
which may be sublimed at 150** with only partial decomposition: HjCjO^ =
CO, -f- CO -f H2O . Crystallized oxalic acid, "EL^Q^O^^ZISL^O , effloresces" very
slowly in warm, dry air, and melts in its water of crystallization at 08**: at
which temperature the liquid soon evaporates to the absolute acid. Oxaliq
anhydride is not formed.
2. Occurrence. — Found in many plants in a free state or as an oxalate. In
sorrel it is found as KHC3O4; in rhubarb as CaCjO^ . As ferrous oxalate in
lignite deposits; as ammonium oxalate in guano.
3. Formation. — (a) By decomposition of cyanogen with water, am-
monium oxalate being one of the products. (6) By the oxidation of
glycol with nitric acid, (c) By heating potassium formate above 400°
(Merz and Wcith, B., 1882, 15, 1507). {d) By passing CO^ over a mixture
of sodium and sand at 360° (Drechsel, Bl, 1868, 10, 131).
4. Preparation. — (a) By action of nitric acid sp. gr. 1.38 upon sawdust,
starch, or sugar. By the continued action of concentrated nitric acid,
after the sugar is all oxidized to oxalic acid, the latter is farther oxidized
to CO2 . {h) By heating sawdust with KOH or NaOH . Hydrogen is
evolved, the cellulose, C^HioO.^ , >^eing converted into oxalic acid. Under
certain conditions, additional products are formed. It is also foimed in
the oxidation of a great many organic compounds.
258 OXALIC ACID. §227, BS.
is more rapid in the presence of a fixed alkali, an alkali chloride and
carbonate being formed. HCIO3 forms COj and vaiying proportions of
CI and HCl . A high degree of heat and excess of oxalic acid favoring
the production of HCl (Calvert and Davies, J. C, 1850, 2, 193).
6. — Bromine decomposes oxalic acid in alkaline mixture, fonning a
bromide and a carbonate. In acid mixture a similar reaction takes place
if a hot saturated solution of oxalic acid be used in excess. With HBrO-. .
bromine and COj are formed; with excess of oxalic acid and heat hydro-
bromic acid is formed.
7. — HIO3 forms COj and I . With mixtures of chlorates, bromates, and
iodates, the chlorate is first decomposed, then the bromate, and finally the
iodate (Guyard, J. C, 1879, 36, 593).
7. Ignition. — The oxalates are all dissociated on ignition. Those of
the metals of the alkalis and alkaline earths are resolved at an incipient
red heat, into carbonates and carbon monoxide (a) — a higher temperature
decomposing the alkaline earth carbonates. The oxalates of metals, whose
carbonates are easily decomposed, but whose oxides are stable, are re-
solved into oxides, carbonic anhydride, and carbon monoxide (b). The
oxalates of metals, whose oxides are decomposed by heat, leave the metal,
and give off carbonic anhydride (c). As an example of the latter class,
silver oxalate, when heated before the blow-pipe, decomposes explosively,
with a sudden puffing sound — a test for oxalates:
(a) CaC^O, = CaCOa + CO
(/>) ZnC^O^ = ZnO + CO, + CO
(c) Aff,C,0, = 2Ag + 2CO3
8. Detection. — (a) By warming with concentrated sulphuric acid after
decomposition of carbonates with dilute sulphuric acid ; showing the pres-
ence of COa by absorption in Ca(OH)o or in a solution of BaCL alkaline
with KOH ; and showing the presence of CO by its combustibility, {b) In
solution by precipitation in neutral, alkaline, or acetic acid solution by
calcium chloride, and solubility of the precipitate in dilute hydrochloric
acid. Frey (Z., 1894, 33, 533), recommends the formation of a zone of
precipitation. To the HCl solution containing BaCl^ and CaCl^ he ad(i>
carefully a solution of NaCoH^Oo and watches the zone of contact.
9. Estimation. — {a) It is precipitated as CaCaO^; after washing, the C& is
determined by §188, 9, from which the oxalic acid is calculated, {h) By the
amount of KMnOi which it will reduce, (r) By measuring" the amount of CO;
evolved when it is oxidized in any convenient manner, usually by MnO. .
(d) By the amount of gold it reduces from AuGls .
, 4. CARBON DIOXIDE. 259
§228. Carbon dioxide. CO2 = 44.0 .
(Carbonic anhydride.)
Carbonic acid (hypothetical). H2CO3 = 62.016 .
C^0-"2 and H'^C^Q-^, = C = and H — — C — — H.
1. Properties.— The specific gravity of the gas CO, is 1.52897 (Crafts, C. r., 1888,
106, 162) ; of the liquid at —34°, 1.057 (Cailletet and Mathias, C. r., 1886, 102,
1202); of the solid (hammered), slightly under 1.2 (Landolt, B., 1884, 17, 309).
Critical temperature, 30.92° (Andrews, Trans. Roy. Soc., 1869, 159, 583; 1876, 166,
21). It is a heavy colorless gas; which at low temperatures, +3°, and high
pressure, 79 atmospheres, may be condensed to a clear mobile liquid; and upon
further cooling this becomes a snow-like mass. Liquid COj is more compres-
sible than other liquids (Natterer, J,, 1851, 59). It diffuses through porous
plates more rapidly than oxygen (Graham, C. A^, 1863, 8, 79). Non-combustible
and a non-supporter of combustion, except that K , Na and Mg" burn in the gas
forming an oxide of the metal and free carbon. It is used in chemical fire
engines. Non-poisonous but causes suffocation (drowning) by exclusion of air.
It is taken internally without injury in soda water, etc.
Liquid CO2 is insoluble in water which swims on the surface. It mixes with
alcohol and ether. It dissolves iodine but does not dissolve phosphorus or
sulphur; it is without action upon K or Na . A spirit thermometer immersed
in the liquid registers —75° (Thilorier, J. in-., 1834, 3, 109). Solid CO, at 767.3
mm. barometric pressure meltft at —77.94° (Regnault, A. Ch., 1849, (3), 26, 257).
When the solid is mixed with ether it gives a temperature of — 98.3°.
2. Occurrence. — In a free state in the air, about 0.04 per cent. Found in
g-reat abundance in the form of carbonates in the earth's crust; e. </., limestone,
marble, magnesite, dolomite, etc.
3. Formation. — (a) By burning wood, coal, etc., in the air. (6) By
turning CO . (c) By the reduction of many metallic oxides upon ignition
.vith charcoal, (d) During fermentation or decay of organic material.
e) By the reaction between acids and carbonates.
Liquid CO2 is made by compressing the gas with pumps at a reduced
temperature.
Solid CO2 is made by allowing the liquid to escape freely, into woolen
3ags and then compressing in wooden moulds (Landolt, L c).
4. Preparation. — CaCOg (chalk or marble) in small lumps is treated with
lydrochloric acid in a Kipp's gas generating apparatus. The gas is passed
rhrough a solution of NaHCOs to remove any HCl that may be carried
:>ver, and then dried by passing through a tube filled with fused CaCL .
It is also prepared on a large scale for making the liquid COj, and for
use in sugar factories by the ignition of limestone : CaCOg = CaO -|- COo .
Preparation of Carbonates.— Na^COs is made by converting NaCl into
Ka^SO^ , by treating it with H2SO4 ; then by long ignition with coal and
calcium carbonate, impure sodium carbonate is formed (Leblanc's process).
Na,SO, -h 4C + CaCO; = CaS -f 4C0 + Na.CO,
260 CARBOy DIOXIDE. §S2S, 5.
It is separated bv lixiviation with water, and farther purified. The
other method, known as the ammonia, or Solvays process, consists in pitt-
ing HH3 and CO, into a concentrated solution of NaCl (a). The HaHCO,
is converted into Na^COs by heat, and the evolved COj used over again (h).
The NH4CI is warmed with MgO (r), and the HH3 which is given off is
used over again. The MgCl, is strongly heated (d) and the KgO is used
over again, and the evolved gas sold as hydrochloric acid. This oontiill-
0X18 process has nearly superseded the Leblanc process.
(fl) Naa + NH, + H,0 + CO, = NaHCO, + NH«C1
(b) 2NaHC0, -h heat = Na^^CO, -f CO, + H,0
(r) 2KH,Cn + MgO = MgCl, + 2NH, + H,0
(d) MgrCl, -h H,0 -h heat = MgO -f 2Ha
The other carbonates are mostly made from the sodium salt (6),
5. Solubilities. — CO, is soluble in water, forming the hypothetical
H0CO3 , which reacts acid towards litmus. At 15** one volume of water
absorbs 1.002 volumes of the gas (Bunsen, A., 1855, 93, 1). It is rapidly
absorbed by hydroxides of the alkalis and of the alkaline earths, forming
normal or acid carbonates: KOH + CO, = KHCO3 or 2K0H + CO, =
K2CO3 + HmO . The carbonates of the alkalis are soluble in water (aci<]
alkali carbonates are less soluble than the normal carbonates), othi-r
carbonates are insoluble in water or only sparingly soluble. The presence
of some other salts, especially ammonium salts, increases the solubility 01
carbonates, notably magnesium carbonate (§189, 5c). Many of the car-
bonates are soluble in water saturated with CO, ; forming acid carbonate?
of variable composition. Boiling removes the excess of CO2 , causing pre-
cipitation of the carbonate.
G. Beactions. — Dry carbon dioxide does not unite with dry calcium
oxide at ordinary temperature (Birnbaum and Maher, B., 1879, 12, 1547;
Scheibler, 7^., 188(>, 19, 19T3). Also at 0° no reaction takes place between
dry CO, and dry Na^O , but at 400° combination takes place with incan-
descence (Bekctoff, /}/., 1880, (2), 34, 327).
Carbonates of the fixed alkalis precipitate solutions of all other metallic
salts: with antimony the precipitate is an oxide; with tin, aluminum,
chromium, and ferricum it is an hydroxide; with silver, mercurosum,
cadmium, ferrosum, manganese, barium, strontium, and calcium it is a nor-
nuil carbonate; with other metals a basic carbonate, except that mercuric
chloride forms an oxychloride. Carbonic acid is completely displaced by
strong acids, for example, from all carbonates, by HCl , HClOsHBr , HBrO,.*
HI , HIO3 , H0C2O, , HNO3 , H3PO, , H,SO, , and even by HjS , completely
from carbonates of the first four groups, incompletely from those of the
fifth and sixth groups (Xandin and Montholon, C. r., 1876, 33, 58).
Ammonium carbonate precipitates solutions of all the non-alkali metals.
:§228, 6. CARBON DIOXIDE, 261
-chiefly as carbonates; except magnesium salts which are not at all pre-
cipitated, a soluble double salt being at once formed (separation of barium,
strontium, and calcium from magnesium). With salts of silver, copper,
cadmium, cobalt, nickel, and zinc the precipitate is redissolved by an
excess of the ammonium carbonate.
The decomposition of carbonates by acids is usually attended by marked
effervescence of gaseous COg which reddens moist litmus paper : NaoCO, +
H,SO, = Na^SO, + H^O + CO^ .
With normal carbonates in cold solution, slight additions of acid (short
of a saturation of half the base) do not cause effervescence, because acid
carbonate is formed: 2Na2C03 + HjSO^ = NaoSO^ + 2NaHC03 ; and
-when there is much free alkali present (as in testing caustic alkalis for
slight admixtures of carbonate), perhaps no effervescence is obtained.
By the time all the alkali is saturated with acid, there is enough water
present to dissolve the little quantity of gas set free. But if the car-
T)onate solution is added drop hy drop to the acid, so that the latter is con-
istantly in excess, even slight traces of carbonate give notable effervescence.
The effervescence of carbonic acid gas, COo , is distinguished from that of
HoS or SO2 by the gas being odorless, from that of NgOg by its being color-
less and odorless ; from all others by the cfforvoscence being proportionally
more forcible. It should be remembered, however, that COo is evolved
(with CO) on adding strong sulphuric acid to oxalates or to cyanates.
On passing the gas, COg , into solution of calcium hydroxide (a); or of
barium hydroxide (b); or into solutions of calcium or barium chloride,
containing much ammonium hydroxide (c), or into ammoniacal solution
of lead acetate (d), a white precipitate or turbidity of insoluble carbonate
is obtained. The precipitate may be obtained by decanting the gas (one-
half heavier than air) from the test-tube in which it is liberated into a
{wide) test-tube, containing the solution to be precipitated ; but the opera-
tion requires a little perseverance, with repeated generation of the gas,
owing to the difficulty of displacing the air by pouring into so narrow a
vessel. The result is controlled better by generating the gas in a large
test-tube, having a stopper bearing a narrow delivery-tube, so bent as to
be turned down into the solution to be precipitated,
(a) CO, + Ca(0H)2 = CaCO, + H^O
(6) CO, + Ba(OH), = BaCO, + H,0
(c) CO, + CaClj + 2NH,0H = CaCO, + 2NH4CI -f H^O
id) CO, -f- Pb^OCCH.O,)^ = PbCO, + PbCCH.O^),
The solutions of calcium and barium hydroxides furnish more delicate
tests for carbonic anhydride than the ammoniacal solutions of calcium and
barium chlorides, but less delicate than lead basic acetate solution. The
latter is so rapidly precipitated by atmospheric carbonic anhydride, that
2Q2 CARBON DIOXIDE,
it cannot be preserved in bottles partly full and frequently opened, and
cannot be diluted clear, unless with recently boiled water.
Solutions of the acid carbonates effervesce^ with escape of CO, , on boiling
or heating, thus :
2XHC0, = K,CO, -f H,0 + CO, . (Gradually, at 100*'.)
2KaHC0. = Na,CO. + H,0 + CO, . (Gradually, at 70**; rapidly at 90** to 100*.)
2NH4HCO, = (NHJ,CO, + H,0 -f CO, . (Begins to evolve CO, at 36^)
(NHJ,H,(CO.), = 2(NH,)aC0. + H,0 + CO, . (Begins at 49^)
7. Ignition. — On ignition, the normal carbonates of the metak of the
fixed alkalis are not decomposed; the carbonates of barium and strontium
are dissociated slowly, at white heat, calcium carbonate at a full red heat,
forming the oxide and COo . At a lower temperature, ignition changes
all other carbonates to the oxide and CO2 , except that the carbonates of
silver at 250°, mercury, and some of the rarer metals are reduced to the
metallic state, CO2 and oxygen being evolved. Stannous and ferrous
oxides ignited in an atmosphere of CO2 are changed to SnOj and Fc.O,,
respectively, with evolution of CO (Wagner, Z., 1879, 18, 559).
8. Detection. — Carbonates are detected : (a) By the sudden eiTervescence
when treated with dilute acids, (b) By the precipitate which this gas
forms with solutions of Ca(0H)2 , Ba(OH:). , or PboO(C.H30o)2 . If but a
small amount of carbonate be present, the mixture must be warmed to
drive the CO^ over into the reagent (6). A non-volatile acid as HoSO^ or
H3PO4 should be used, as a volatile acid might pass over with the CO2 and
prevent the formation of a precipitate, (c) Phonolphthalein detects the
normal carbonate in solution of the bicarbonate (very delicate). Sodium
bicarbonate fails to give a precipitate with magnesium sulphate (distinc-
tion from Na^CO.,) (Patein, J. Pharm., 1892, (5), 25, 448).
To detect free carbonic acid in ])rescnce of bicarbonates, a solution of
1 part of rosolic acid in 500 parts of 80 per cent alcohol may be employed,
to which barium hydroxide has been added until it begins to acquire a
red tinge. If 0.5 oc. of this rosolic acid solution be added to about 50 oc.
of the water to be tested — spring water, for instance — the liquid will be
colorless, or at most faintly yellowish if it contains free carbonic aci<l.
whereas, if there be no free carbonic acid, but only double salts, it will
bo red (Pettenkofer, IHngl, 1875, 217, 158).
Salzer (Z., 1881, 20, 227) gives a test for free carbonic acid or bicar-
bonates in presence of carbonates, founded on the fact that the Nessler
ammonia reaction (§207, (U*) does not take place in presence of free car-
l)onic acid or bicarbonates. This reaction is also used to detect the presence
of fixed alkali hydroxides in the fixed alkali carbonates. In presence of a
fixed alkali hydroxide a broAvn precipitate is obtained (Dobbin, J, Soc. Ind,y
1888, 7, 829).
§230, 1. CYANOGEN— HYDROCYANIC ACID, 263
9. ISstimation. — (a) By decomposition of a weighed sample with acids and
determining the COj by loss of weight, after taking into consideration the
gain in weight due to the acid used. (6) By decomposition of the weighed
sample and collection of the CO, in a weighed KOH solution, (c) By decom-
position with an excess of a standard acid, boiling to expel the CO, and
titrating the excess of acid, (d) Sodium bicarbonate may be estimated by
titration with sodium hydroxide: KaHCO. + NaOH = Na^CO, -f HjO . The
first excess of sodium hydroxide beyond the reaction gives a brown precipi-
tate with silver nitrate (Lunge, Z. angew,, 1897, 169; Bohlig, Arch. Pharm,, 1888,
226, 541).
§229. Cyanogen. CN = 26.04 .
N=C — C=N.
A colorless, intensely poisonous gas; specific in'avity, 1.8064 (Gay-Lussac, Oilb,,
1816, 53, 145). The molecular weight shows the molecule to be CaN, . At
ordinary atmospheric pressure it liquifies at — 22" (Drion, J., 1860, 41); at 20®
under four atmospheres pressure (Hofmann, B., 1870, 3, 658). The gas has
an odor of bitter almonds and burns with a red color to the flame forming
COj and N . When cooled to about the freezing point of mercury it solidifies
to a crystalline ice-like mass (Hofmann, /. c). Critical temperuture, 124** (De-
war, C. N,, 1885, 51, 27). The liquid is colorless, mobile and a non-conductor
of electricity. It occurs in the gas from the coke ovens (Bunsen and Playfair,
«/. pr., 1847, 42, 145). It is prepared: (a) By heating the cyanides of mercury,
silver or gold: Hg(CN)2 = Hg + C^N, . (h) By the dry distillation of am-
monium oxalate: (11^^)20.0^ = 4H2O -f CsN, . (r) By fusing KCN with
HgClj: 2KCN -h HgClj = Hg -f 2KC1 -f CjN, . (d) By heating a solution of
CUSO4 with KCN . Half of the CN is evolved and CuCN is formed. If the
CuCN be heated with FeCl, or MnO, and HCsHsO, , the remainder of the
CN is obtained. The gas is purified by absorption with aniline; oxygen,
nitrogen and carbon dioxide are not absorbed (Jacquemin, A, Ch,, 1886, (6), 6,
140). It combines with CI , Br , I , S , P , and with many of the metals,
reacting very much like the halogens. It dissolves in water, alcohol and
€ther; but gradually decomposes with formation of ammonium oxalate and
carbonate (Vauquelin, A, Ch., 1823, 22, 132; Buif and Hofmann, A., 1860, 113,
129). At 500* it combines with hydrogen to form HCN (Berthelot, B/., 1880,
(2), 33, 2). With Zn it forms Zn(CN)2 , rapidly at lOO**. With HCl and abso-
lute alcohol it forms oxalic ether, which shows cyanogen to be the nitrile of
oxalic acid (Pinner and Klein, B., 1878, 11, 1481). With solution of KOH,
KCN and XCNO are formed: C^N, -f 2K0H = KCN -f XCNO -f H,0 . Com-
pare the reaction with chlorine and KOH (§270).
§230. Hydrocyanic acid. HCN = 27.048 .
H— C = N.
1. Properties. — Hydrocyanic acid is a clear, mobile liquid, boiling at 26**. At
— 15® it freezes to a fibrous crystalline mass. Specific gravity at 19**, 0.697
(Bleekrode, Proc. Roy, 80c., 1884, 37, 339). It burns with a bluish-red flame,
forming H,0 , CO3 and N. Its index of refraction is much less than that of
water (Mascart, C, r., 1878, 86, 321). It is one of the most active poisons
known; of a very characteristic odor, somewhat resembling that of bitter
almonds. The antidote is chlorine or ammonia by inhalation. Its water
solution decomposes slowly, forming ammonium formate: scarcely at all in
the dark. It distils readily unchanged. The U. S. P. solution contains two
per cent of HCN. It is a weak acid, scarcely reddening litmus; its salts are
partially decomposed by CO, . The free acid or soluble salts when warmed
264 HTDROCTAyiC ACID.
with dilute alkalis or acids (with strong acids in the cold) becomes formic
acid and ammonia: HON -f 2HaO = HCO,H -h NH. .
2. Occurrezice. — The free acid does not occur in nature, but in combination
in the kernels of bitter almonds, peaches, apricots, plums, cherries and
quinces; the blossoms of the peach, sloe and mountain ash; the leaves of the
peach, cherry laurel and Portugal laurel; the young branches of the peach;
the stem-bark of the Portugal laurel and mountain ash: and the roots of the
last-named tree, when soaked in water for a time and then distilled, yield
hydrocyanic acid, together w^ith bitter-almond oil. Potassium C3'anide appears
in the deposits of blast furnaces for the smelting of iron ores.
3. Formation. — {a) Decomposition of amygdaline by emulsine and distilla-
tion. (6) By the action of the electric .spark on a mixture of acetylene and
nitrogen (Berthelot, J., 1874, 113). (c) By heating a mixture of cyanogen and
hydrogen (§229). (d) By the dry distillation of ammonium formate: NH«CHO,
= HON -f 2H2O . (f) By boiling or fusing many organic compounds contain-
ing nitrogen with XOH , forming KCN (Post and Huebner, B., 1872, 5, 40S).
(f) By decomposition of metallic cyanides with mineral acids. (^7) By heating
chloroform with a mixture of ammonium and potassium hydroxides (Hof-
mann. A., 1867, 144, 116).
4. Preparation. — (a) By the action of dilute sulphuric acid on potassium
ferrocyanide: 2K,Fe(CN)« + 3H5SO, = 6HCN + K,Fe,(CN)« -f SK^SO, .
(6) By action of acids upon metallic cyanides, (r) By the action of sulphuric
acid upon mercuric cyanide in the presence of metallic iron: Hg(CN)a -f Fe -h
H,SO, = 2HCN -h FeSO, -f Hg .
Metallic cyanides are prepared: (a) By the action of HCN on metallic
hydroxides. (6) By the action of soluble cyanides on metallic salts, (r) By
igniting potassium ferrocyanide: K,Fe(CN)e = 4KCN -f FeC, -f N^ . (d) By
heating potassium ferrocyanide with potassium carbonate. If prepared in
this manner it contains some cvanate: K4Fe(CN)« -f K2CO, = 5KCflT -f- XCNO
-h Fe -f CO, .
5. Solubilities. — Hydrocyanic acid is soluble in water, alcohol and ether in
all proportions. A mixture of equal parts acid and water increases in tem-
perature from 14** to 22..")°: it also increases slightly in volume (Bussv and
Buignet, A, Ch., 1865, (4), 4, 4).
The cyanides of the alkali metals, alkaline earth metals, and mercuric
cyanide, are soluble in water, barium cyanide being but sparingly soluble.
The solutions are alkaline to test-paper. The other metallic cyanides are
insoluble in water. Many of these dissolve in solutions of alkali cyanides,
by combination, as double wetallir cyanides,
Pb , Hg , As , Sb , Sn , Bi and Cd are dissolved by KCN with absorption
of oxygen. Cu , Al , Fe (by H or CO), Co, Ni , Zn and Mg with evolution
of hydrogen: 2Cu + 2KCN -f- 2HoO = 2CuCN -f 2K0H -|- H, . Iron or steel
wire are not attacked (Goyder, (\ .Y., 1894, 69. 262, 268 and 280).
6. Reactions. — There are two classes of double cyanides, both of which are
formed when a cyanide is precipitated by an alkali cyanide, and redissolved
by excess of the precipitant: HgCL + 2kCN = Hg(CN)j -+- 2KC1; and with
excess of KCN: Hg(CN)2 -h 2KCN = (KCN)2Hg(CN)2 .
Class I. Double cyanides ichivh are not affected by alkali hydroxides, but suffer
dissociiition when treated iti4h dilute acids: (KCN)2Hg(CN)3 -f- 2HC1 = Hg(CN).
-|- 2KCI -f- 2HCN . These closely resemble the double iodides (potassium
mercuric), and the double sulphides or thiosalts (§69, 5c and 6e). The most
frequently occurring of the double cyanides of this class, which dissolve in
water, are given below:
Potassium (or sodium) zinc cyanide, K2Zn(CN)4 or (KCN),Zn(CN)2 .
Potassium (or sodium) nickel cyanide, KaNi(CN)4 or (KCN)2Ni(CN), .
Potassium (or sodium) copper cyanide, K:jCu(CN)< or (KCN)jCu(CN)j .
Potassium cadmium cyanide, K2Cd(CN)4 or (KCN)2Cd(CN)s .
Potassium (sodium or ammonium) silver cyanide, KCNAgCN or KAg(CN)3 .
Potassium (or sodium) mercuric cyanide, KsHg(CN)4 or (KCN),Hg(CN), .
Potassium (or sodium) auric cyanide, KAu(CN)4 or KCNAa(CN), •
Jj230, G. HYDROCYAyiC ACID. 265
Class II. Double cyanides which, as precipitates, are transposed hy alkali hydrox-
ides, in dilute solutio^n (a), a fid are transposed, without dissociation, by dilute acids
(h). In these double cyanides, as potassium ferrous cyanide, K4Pe(CN)e , the
whole of the cyanogen appears to form a new compound radical with that metal
whose single cyanide is insoluble in water; thus, Fe(CN)a as ** ferrocyanogen,"
giving K^FeCCN)^ as *' potassium ferrocyanide " (for the potassium ferrous
cyanide). These more stable double cyanides or " ferrocyanides," etc., cor-
respond to the platinic double chlorides or ** chloroplatinates " (§74, 5c), and
the palladium double chlorides, or chloropalladiates (§106, 5c). The most
frequently occurring of the double C3'anides of this class, which are soluble in
water, are given below.
(a) Cu,Fe(CN), + 4K0H = 2Cu(0H), + K4Pe(CN)a
(6) K,Pe(CN)o + 2H,S0, = 2X,S04 -f H,Pe(CN).
2K,Fe(CN). + 3H,S0, = 8X,S0« + 2H.Fe(CN),
Alkali ferrocyanides, as K4Fe"(CN)« , potassium ferrous cyanide.
Ferricyanides, as K,Fe'"(CN)a , potassium ferric cyanide.
Cobalticyanides, as K,Co"'(CN)8 , potassium cobaltic cyanide.
Manganicyanides, as K,Mn'''(CN)« , potassium manganic cyanide.
Chromicyanides, as Ka(Cr'")(CN)e , potassium chromic cyanide.
The easily decomposed double cyanides of Class I. are, like the single cyan-
ides, intensely poisonous. The difficultly decomposed double cyanides of Class
II. are not poisonous.
The Single Cyanides are transposed by the stronger mineral acids, more
or less readily, with liberation of hydrocyanic acid, HCN, effervescing from
concentrated or hot solutions, remaining dissolved in cold and dilute solu-
tions. Mercuric cyanide furnishes HCN by action of HjS , not by other
acids. The cyanides of the alkali and alkaline earth metals are transposed
by all acids — even the carbonic acid of the air— and exhale the odor of
hydrocyanic acid. Solution of silver nitrate precipitates, from solutions
of cyanides or of hydrocyanic acid (not from mercuric cyanide) silver
njanidcy AgCN , white, insoluble in dilute nitric acid, soluble in ammonium
liydroxide, in hot ammonium carbonate, in potassium cyanide, and in
thiosulphates — uniform with silver chloride. Cold strong hydrochloric
icid decomposes it with evolution and odor of hydrocyanic acid (recogni-
tion from chloride), and when well washed, and then gently ignited, it does
not melt, but leaves metallic silver, soluble in dilute nitric acid, and pre-
3ipitable as chloride (distinction and means of separation from chloride).
Solution of mercurous nitrate, with cyanides or hydrocyanic acid, is
•esolved into metallic mercury, as a gray precipitate, and mercuric cyanide
md nitrate, in solution. Salts of copper react, as stated in §77, 66; salts
)f lead, as stated in §57, 0&.
Ferrous salts, added to saturation, precipitate from solutions of cyan-
des, not from hydrocyanic acid, ferrous cyanide, Fe(CN)2 , white, if free
from the ferric hydroxide formed by admixture of ferric salt, and, with
:he same condition, soluble in excess of the cyanide, as (with potassium
?yanide), (SXS)^t{CS)2 = 'Sk,J^t{(jS)Q , potassium ferrocyanide (a). On
266 HYDIWCYAXIC ACID. §280, 7.
acidulating this solution, it gives the blue precipitates with iron salts,
more marked with ferric salts (6) :
(a) 2KCN + FeSO, = Fe(CN), -f K,SO,
Pe(CN), + 4KCN = K,Fe(CN),
(6) 3K,Fe(CN), + 4PeCl. = Pe*(Pe(CN),), + 12KC1
This production of the blue ferric ferrocyanide is made a delicate test fon
hydrocyanic acidy as follows: A little potassium hydroxide and ferrous
sulphate are added, the mixture digested warm for a short time; then a
very little ferric chloride is added, and the whole slightly acidulated (so
as to dissolve all the ferrous and ferric hydroxides), when Prussian blue
will appear, if hydrocyanic acid was present (Link and Moeckel, Z., 1878,
17, 456).
Solution of nitroplienic acid, picric acid, QJ3Lo(JSl0^fiJLy added, in a
small quantity, to a neutralized solution of cyanides of alkali metals, on
boiling( and standing), gives a blood-red color, due to picrocyanate (as
KCgH^NjOe). This test is very delicate, but not very distinctive, as var-
ious reducing agents give red products with nitrophenic acid (Vogel,
C. N.y 1884, 50, 270).
The fixed alkali hydroxides, in boiling solution, strongly alkaline, gradu-
ally decompose the cyanides with production of ammonia and formate:
HCN + KOH + HoO = ECHO2 +NH3 . Ferrocyanides and ferricyanides
finally yield the same products. Dilute alkalis, not heated, transpose, as
by equation a, class II above.
Cyanides are strong reducing agents. The action is not so marked in
solution as in state of fusion (7). Permanganates are reduced by cyan-
ides, and cupric hydroxide in alkaline solution forms Cu'. Solutions
of cyanides on exposure to the air take up some oxygen with formation of
a cyanate : 2KCN -f Oo = 2KCN0 . Commercial potassium cyanide always
contains some potassium cyanate. By warm digestion of a cyanide with
sulphur or with yellow ammonium sulphide a thiocyanate is formed (S).
Hydrocyanic acid reduces PbO.^ , forming Pb(CN)2 and CN : PbOo + 4HCH
= Pb(CN), -f C2N, + 2H2O (Liebig, A., 1838, 25, 3). With^HCN and
HoO^ oxafnide is formed (Altfield, J. (7., 1863, 16, 94). Chlorine forms
with hydrocyanic acid a cyanogen chloride (Serullas, A, Ch,, 1828, 38,
370); with iodine the reaction is not so marked, but a similar product is
formed (Meyer, B., 1887, 20, III, 704). Concentrated sulphuric acid
decomposes all cyanides.
7. Ignition. — By fusion with fixed alkalis, cyanides and all compounds
containing cyanogen yield ammonia. In state of fusion cyanides are very
efficient reagents for reduction of metals from their oxides or sulphides
to the metallic state (§69, 7). The cyanates or thiocyanates formed in
the reaction are not readily decomposed by heat alone.
|231. HYDROFERROCYAXIC ACID. 267
8. Detection. — Cyanides may be detected: (a) By the odor of the free
iieid upon decomposition of the cyanide with acids. This test must be
applied with extreme caution as the evolved HCN or CN is very poisonous.
(h) By formation of a ferrocyanide and its reaction with ferric salts, as
described in 6. (c) The production of the red ferric ihiocyanate is a test
for hydrocyanic acid, more delicate than formation of ferrocyanide. By
warm digestion this reaction occurs: 2ECN + 82 = 2KCNS ; or:
2(NH,),S, + 4HCN = 4NH,CNS + 2H,S + S,
To the material in an evaporating-dish, add one or two drops of yellow
ammonium sulphide, and digest on the water-bath until the mixture is
colorless, and free from sulphide. Slightly acidulate with hydrochloric
acid (which should not liberate HgS), and add a drop of solution of ferric
chloride ; the blood-red solution of ferric thiocyanate will appear, if hydro-
cyanic acid was present (Link and Moeckel, Z. c).
(d) Link and Moeckel also recommend the following test for cyanides,
delicate to 1-3,000,000. Saturate a filter paper with a four per cent
alcoholic solution of guaiac; allow the alcohol to evaporate; then moisten
the paper with a one-fourth per cent solution of copper sulphate, and
allow the unknown solution to trickle over this test paper. A deep blue
color indicates the presence of a cyanide.
To detect cyanides in presence of ferri- and ferrocyanides it is directed
to add tartaric acid and, in a distilling flask, pass a current of carbon
dioxide, warming not above 00°. Test the distillate by the methods
given above. Ferro- and ferricyanides do not yield HCN under 80° (Hilger
and Tamba, Z., 1891, 30, 529; also Taylor, C. N., 1884, 50, 227).
9. Estimation. — (a) The nearly neutral solution of cyanide is titrated with
standard silver nitrate. No precipitate occurs as long" as two molecules of
alkali cyanide are present to one of silver nitrate. Soluble AgCN,XCN is
formed. As soon as the alkali cyanide is all used in the formation of the
double cyanide, the next molecule of sih'er nitrate decomposes a molecule of
the double salt, forming" two molecules of insoluble silver cyanide: giving a
white precipitate for the end reaction. Chlorides do not interfere (Liebig, A.,
1851, 77, 102). (6) By titration with a standard solution of HgCl, , applicable
in presence of cyanates and thiocyanates (Hannay, J. C, 1878, 33, 245).
§231. hydrofcrrocyanic acid. 'E^TeiCS)^ = 216.172 .
H\Fe"(CN)-'«.
Absolute hydroferrocyanic acid (§230, 6, Class II.), is a white solid, freely
soluble in water and in alcohol. The solution is strongly acid to test-paper,
and decomposes carbonates, with effervescence, and acetates. It is non-volatile,
but absorbs oxygen from the air, more rapidly when heated, evolving hydro-
cyanic acid and depositing Prussian blue: 7H4i'e(CN)e -f 0, = Fe4(re(CN)«)a
-f 2H,0 -f 24HCN .
Potassium ferrocyanide is the usual starting point in the preparation of the
free acid or any of the salts. It is prepared by fusing together in an iron
268
BTDROFERROCTANIC ACID.
S2SL
kettle nitrogenous animal matter (blood, hair, horn, hoof, etc.), cozmnerdal
potash (KOH), and scrap iron. The ferrocyanide is formed when this maas i&
digested with water. The filtrate is evaporated to crystallization (lemon-yellow
prism), soluble in four parts of water.
Hydroferrocyanic acid is formed by transposition of metallic ferrocyanidea
in solution, with strong acids («). ^Yhen the solution is heated, hydrocyanic
acid is evolved; in the case of an alkali ferrocyanide, without absorption of
oxygen (6). Potassium ferrocyanide and sulphuric acid are usually employed
for preparation of hydrocyanic acid (e) :
(a) K,Fe(CN), + 2H,S0, = 2K,S0, + H,F6(CN),
(b) 3H,Fe(CN)« + K,Fe(CN)a = 2K,FeFe(CN)a + 12HCN
(c) 2K,Fe(CN), + 3H,S0, = 3K,S0, + K,FeF6(CN)a -f 6HCN
The forrocyanides of the alkali metals, strontium, calcium and magnesium,
are freely soluble in water; of barium, sparingly soluble; of the other metals,
insoluble in water. There are double ferrocyanUfes: soluble and insoluble: that
of barium and potassium is soluble, but potassium^ calcium ferrocyanide is in-
soluble. The most of the ferrocyanides of a heavy metal and an alkali metal
are insoluble. Potassium and sodium ferrocyanides are precipitated from their
water solutions by alcohol (distinction from ferricyanides).
The soluble ferrocyanides are yellowish in solution and in crystals, white
when anhydrous. The insoluble ferrocyitnides have marked and very diverse
colors, as seen below.
as K4Fe(CN)fl , give, with soluble salts of:
Solutions of alkali ferrocyanides
Aluminum, a white precipitate,
Antimony a white "
Bismuth, a white "
Cadmium, a white "
Calcium, a white "
Chromium, no *'
Cobalt, a green, then gray "
Copper, a red-broAvn "
Gold, no
Iron (Fe"), white, then blue "
Iron (Fe'")» a deep blue "
Xead, a white "
Magnesium, a white "
a yellow-white "
Manganese, a white "
Mercury (Hg^), a white "
Mercury (Hg"), a white
Molybdenum, a brown "
Nickel, a greenish-white
Silver, a white
Tin (Sn" and Sniv), white
Uranium (uranous), brown
Uranium (uranyl), red-brown
Zinc, a white, gelatinous
Al(OH), and Fe(CN), (formed slowly).
Sb4[Fe(CN).]..25HaO.
Bi,(Fe(CN).)..
Cd2Fe(CN)o (soluble in HCl).
K,CaFe(CN). .
Co,Fe(CN)« .
Cu2Fe(CN)e .
K,FeFe(CN)o .
Fe,(Fe(CN),)..
Pb,Fe(CN). .
(NHj2MgFe(CN). (in presence of NH,OH)
KjMgFe(CN)o (only in concentrated solu-
tion).
Mn,Fe(CN)« (soluble in HCl).
Hg4Fe(CN)fl (gelatinous).
Hg,Fe(CN)« , turning to Hg(CN), and
Fe,(Fe(CN)o)., blue.
Ni,re(CN)o .
Ag,Fe(CN)« , (slowly turning blue).
(gelatinous).
UFe(CN)o .
(UO,).Fe(CN), .
Zn,Fe(CN)o .
See Wyrouboff (A, Ch., 1876 (5), 8, 444; and 1877, (5), 10, 409).
Insoluble ferrocyanides are transposed by alkalis (§230, 6, Class II.)
It will be observed (§230. 6) that fcrrncyiimdcH are fctrom combinations, while
fcrrirt/afiidnt are ferric combinations. And, although ferrocyanides are far less
easily oxidized than simple ferrous salts, being stable in the air, they are
§232. HYDROFERRICYANIC ACID. 269
nevertheless reducing agents, of moderate power: 2X4Fe(CN)« + CI, =
2K,Fe(CN)o + 2KC1 .
PbO, with sulphuric acid forms Pb" and HaFe(CN)e .
Ag' with fixed alkali forms an alkali ferricyanide and metallic silver.
Crvi with phosphoric acid, gives Cr'" and HaF^CCN)^ (Schonbein, J, pr., 1840,.
20, 145).
Co'" with phosphoric acid forms Co" and H,Fe(CN)o .
Nl'" with acetic acid gives Ni" and H,Fe(CN)fl .
MnO, with phosphoric acid gives Mn" and H,Fe(CN)e .
Mnvii forms with potassium hydroxide MnO^ and potassium ferricyanide.
With sulphuric acid, manganous sulphate and hydroferricyanic acid.
Ferricyanides when boiled with NH4OH give ferrocyanides (Playfair, J. C,
1857,9, 128).
HNOs forms first hydroferricyanic acid, then hydronitroferricyanic acid and
NO.
HNOa forms hydroferricyanic acid, and then hydronitroferricyanic acid, NO
being evolved.
CI forms first hydroferricyanic and hydrochloric acids. Excess of chlorine to
be avoided in preparation of ferricyanides.
HCIO, forms hydroferricyanic and hydrochloric acids.
Br forms hydroferricyanic and hydrobromic acids.
HBrOs forms hydroferricyanic and hydrobromic acids.
I , iodine is decolored by potassium ferrocyanide, and some potassium ferri-
cyanide and potassium iodide are formed. The action is slow and never
complete (Qmelin's Hand-book^ 7, 459).
HIO, forms hydroferricyanic acid and free iodine.
In analysis, soluble ferrocyanides are recognized by their reactions with
ferrous and ferric salts and ccpper salts (see 66, §126 and §77). Separated
from ferricyanide, by insolubility of alkali salt in alcohol.
Ferrocyanides are estimated in solution with sulphuric acid by titrating with
standard KlfnO« . Also by precipitation with CuSO^ either for gravimetric de-
termination or volumetrically, using a ferric salt as an external indicator.
§232. Hydroferricyanic acid. H:3Fc(CN)« = 215.164 .
H'3Fc'"(CN)-', .
Absolute hydroferricyanic acid, HgFeCCN), , is a non-volatile, crystallizable
solid, readily soluble in water, with a brownish color, and an acid reaction to
test-paper. It is decomposed by a slight elevation of temperature. In the
transposition of most ferricyanides, by sulphuric or other acid, the hydro-^
ferricyanic acid radical is broken up.
Potassium ferricyanide is the usual starting point in the preparation of most
ferricyanides. It is prepared by passing chlorine into a cold solution of
X«Fe(CN)« until a few drops of the liquid give a brownish color, but no pre-
cipitate with a ferric salt. The solution is evaporated to crystallization and
the salt repeatedly recrystallized from water. Large red prismatic crystals,,
very soluble in water, freely soluble in alcohol (distinction from K4F6(CN)fl).
The free acid is made by adding to a cold saturated solution of K3Fe(CN)«
three volumes of concentrated HCl and drying the precipitate which forms,
in a vacuum (Joannis, C. r., 1882, 94, 449, 541 and 725). Lustrous, brownish-
green needles, very soluble in water and alcohol, insoluble in ether.
The ferricyanides of the metals of the alkalis and alkaline earths are soluble
in water; those of most of the other metals are insoluble or sparingly soluble.
The soluble ferricyanides have a red color, both in crystals and solution; those
insoluble have different, strongly marked colors. Potassium and sodium ferri-
cyanides are but slightly, or not at all, precipitated from their water solutions
by alcohol (separation from ferrocyanides).
Ferricyanides are not easily decomposed by dilute acids; but alkali hydrox-
ides, either transpose them or decompose their radicals (§230, 6).
270 HTDROFERRICYANIC ACID. §812.
•
Solutions of metallic ferricyanides give, with soluble salts of:
AlnmimiTn, no precipitate.
Antimony, no precipitate.
Bismuth, light-brown precipitate, BiFe(CN)« , insoluble in HCl.
Cadmium, yellow precipitate, Cdr[Fe(CN)«]3 , soluble in acids and in ammo*
nium hydroxide.
Chromium, no precipitate.
Cobalt, brown-red precipitate, Coa[re(CN);]a , insoluble in acids. With ammo-
nium chloride and hydroxide, excess of ferricyanide gives a blood-red
solution, a distinction of cobalt, from nickel, manganese and zinc.
Copx>er, a yellow-green precipitate, Cu,[Pe(CN)«], , insoluble in HCl.
Gold, no precipitate.
Iron (ferrous), dark blue precipitate^ Fe8[Fe(CN)s]2 , insoluble in acids.
Iron (ferric), no precipitate^ a darkening of the liquid.
Lead, no precipitate, except in concentrated solutions (dark brown).
Manganese, brown precipitate, Mn,[Fe(CN)«]2 , insoluble in acids.
Mercury (mercurous), red-brown precipitate, turning white on standing.
Mercury (mercuric), no precipitate.
Nickel, yellow-green precipitate, Nia[Fe(CN)o]2 , insoluble in hydrochloric acid.
With ammonium chloride and hydroxide, excess of ferricyanide gives a
copper-red precipitate.
Silver, a red-brown precipitate, K^^'E^iCN) ^ , soluble in NH4OH .
Tin (stannous), white precipitate, Snj[Fe(CN),]2 , soluble in hydrochloric acid.
Tin (stannic), no precipitate.
TTranium (uranons), no precipitate.
Zinc, orange precipitate, Zn3[Fe(CN)fl]3 , soluble in HCl and in NH«OH .
Ferricyanides, ferric combinations, are capable of acting as oxidizing agents,
becoming ferrocj-anides, ferrous combinations.
4K,Fe(CN), 4- 2H2S = .lK,Fe(CN)« -f H,Fe(CN). -h S,
2K,Fe(CN)« + 2KI = 2K,Fe(CN). -f I, .
Nitric acid, or acidulated nitrite, by continued digestion in hot solution,
effects a still higher oxidation of ferricyanides, with the production, among
other products, of nitrofcrricyaniilcs or nitroprtissides (Playfair, Phil. May,, 1845.
(3), 26, 197, 271 and 348). These salts are generally held to have the composi-
tion represented by the acid H..Fe(N0)(CN)5 . Sodium nitroprusside is used as
a reagent for soluble sulphides — that is, in presence of alkali hydroxides, a
test for hydrosulphurie acid; in presence of hydrosulphuric acid, a test for
alkali hydroxides (§207, 6/)).
K3Fe(*CN)„ is reduced to K4Fe(CN)« bv Pd , Th , Mg and As, but not bv
Pb , Hg , Ag , Sb , Sn , Au , Pt , Bi , Cu , Cd , Te , Al , Fe , Co , Mn , Zn and In.
When a sheet of any metal except Au and Pt is placed in contact with a
solution of K,Fe(CN)« and FeCl, , a coating of Prussian blue is soon formed
(Boettger, J, C, 1873, 26, 473).
Pb" with potassium hydroxide forms PbO, and potassium ferrocyanide (Watt*'
Dictionary, 1889, 2,' 340).
Sn" with potassium hydroxide forms potassium stannate, KsSnO, and potas-
sium ferrocyanide {Wattti^ Dictionary, I.e.).
Cr"' forms in alkaline mixture a chroma te and a ferrocyanide (Bloxam, C. y.,
1885, 52, 109).
Mn" with potassium hydroxic'e forms MnO, and potassium ferrocyanide
(Boudault, ,/. ;)r., 1845, 36, 23).
Co" and Ni" are not oxidized.
In alkaline solutions K,Fe(CN)o oxidizes sugar, starch, alcohol, oxalic acid
and indigo (Wallace, J. C, 1855, 7, 77; Mercer, Phil. Mag., 1847, (3). 31, 126).
HNO2 and HNOs both form hydronitroferricyanic acid, HjFe(NO)(CN)B .
NO in alkaline solution becomes a nitrate (Wallace, /. c).
T in alkaline solution becomes a phosphate (Wallace, /. c).
§233. CYANIC ACID, 271
r,PO, forms H^FeCCN). and H3PO, .
H,S forms S, then H^SO^ and H^FeCCN), (Wallace, l.c).
SO, forms H.SO^ and H«Fe(CN)« .
CI decomposes ferricyanides.
HCIO, acts upon K,Pe(CN)a , forming potassium superferricyanide, XsFeCCN).
(Skraup, A„ 1877, 189, 368).
HI forms H4Fe(CN)« and I .
Ferricyanides in solution are detected by the reactions with ferrous and
ferric salts (§126, 66). Insoluble compounds are ignited (under a hood) with
a fixed alkali, giving an alkali cyanide, ferric oxide, and an oxide of the metal
in combination. Detect the alkali cyanide as directed (§230, 8). A ferri-
cyanide is estimated by reduction to ferrocyanide with KI in presence of con-
centrated HCl; the liberated iodine being titrated with standard Na.SjO, .
Or it is reduced to ferrocyanide by boiling with XOH and FeSOf , filtering,,
acidulating with HsSO^ and titrating with XMnO« .
§233. Cyanic acid. HCNO = 43.048 .
H — — C=N.
The cyanates of the alkalis and of the fourth-group metals may be made by
passing cyanogen gas into the hydroxides. The cyanates of the alkalis are
easily prepared by fusion of the cyanide with some easily reducible oxide.
CN, + 2K0H = KCNO -f KCN + H^O
KCN -h PbO = KCNO -f Pb
4KCN -f Pb,0, = 4XCN0 + 3Pb
The free acid may be obtained by heating cyanuric acid, HgCsNaO, , to
redness, better in an atmosphere of CO, . Cyanic acid is found in the dis-
tiUate. HaC.N.O, = 3HCN0 .
Absolute cyanic acid, HCNO , is a colorless liquid, giving off pungent, irri-
tating vapor, and only preserved at very low temperatures. It cannot be
formed by transposing metallic cyanates with the stronger acids in the pres-
ence of water, by which it is changed into carbonic anhydride and ammonia:
HCNO -f H3O = NHj -+- CO2 . The cyanates, therefore, when treated with
hydrochloric or sulphuric acid, effer\^esce with the escape of carbonic anhydride
(distinction from cyanides), the pungent odor of c}/anic avid being perceptible:
2KCN0 -h 2H3SO/-f 2H.0 = K^SO, -f (NH^^SO, + 200^ . The ammonia
remains in the liquid as ammonium salt, and may be detected by addition of
potassium hydroxide, with heat.
The cyanates of the metals of the alkalis and of calcium are soluble in water;
most of the others are insoluble or sparingly soluble. All the solutions
gradually decompose, with cA'olution of ammonia. Silver cifunate is sparingly
soluble in hot water, readily soluble in ammonia; soluble, with decomposition,
in dilute nitric acid (distinction from cyanide). Copper cyanate is precipitated
greenish-yellow.
Ammonium cyanate in solution changes gradually, or immediately w^hen boiled,
to urea, or carbamide, with which it is isomeric: NH4CNO ^ CO(NH..)2 . The
latter is recognized by the characteristic crystalline laminae of its nitrate.
when a few drops of the solution, on glass, are treated with a drop of nitrit*
acid. Also, solution of urea with solution of mercuric nitrate, forms a white
precipitate, CH^N^OCHgO): , not turned yellow (decomposed) by solution of
.sodium carbonate (no excess of mercuric nitrate .being taken). Solution of
urea, on boiling, is resolved into ammonium carbonate, which slowly vaporizes:
CHfNzO -f 2H2O = (NH4)2CO, . Cyanates, in the dry way, are reduced by
strong deoxidizing agents to cyanides.
For detection of a cyanate in presence of cyanides, see Schneider, B., 1895,
28, 1540.
1272 THIOCYAMC ACID.
§234. Thiocyanic acid. HCNS = 59.118 . ,
H — S — C = N.
An aqueous solution of HCNS may be obtained by treating lead thiocyanate
suspended in water with HjS , also by treating barium thiocyanate with HjSO,
in molecular proportions. The anhydrous acid is obtained by treating dry
Hg(CNS)3 with HjS . Potassium thiocyanate is formed by fusing KCN with
S . Or two parts of K^FeCCN). with one part of sulphur. Also by fusing the
cyanide or ferrocyanide of potassium with potassium thiosulphate, X2S3O,:
2KCN -f S3 = 2KCNS
K,F6(CN)« + 3S, = 4KCNS -f Fe(CNS),
4KCN + 4K,S,0, = 4KCNS -f 3X,S0« + K,S
2K4re(CN)o 4- 12X2S,0, = 12KCNS + OX^SO, -f K,S -f 2FeS
Thiocyanic acid is quite as frequently called sulphocyanic acid, and its salts
either thiocyanates or sulphocyanates. It corresponds to cyanic acid, HCNO,
oxygen being substituted for sulphur.
Absolute thiocyanic acid, HCNS , is a colorless liquid, crystallizing at 12*"
and boiling at 85°. It has a pungent, acetous odor, and reddens litmus. It is
soluble in water. The absolute acid decomposes quite rapidly at ordinary
temperatures: the dilute solution slowly; with cA'olution of carbonic anhydride,
carbon disulphide, hydrosulphuric acid, hydrocyanic acid, ammonia, and other
products.
The same products result, in greater or less degree, from transposing soluhlf
thiocyanates with strong acids; in greater degree as the acid is stronger and
heat applied; while in dilute cold solution, the most of the thiocyanic acid
remains undecomposed, giving the acetous odor. The thiocyanates, insolubh
in water, are not all readily transposed. Thiocyanates of metals, «whose sul-
phides are insoluble in certain acids, resist the action of the same acids.
The thiocyanates of the metals of the alkalis, alkaline earths; also, those of
iron (ferrous and ferric), manganese, zinc, cobalt and copper — are solnble in
watfer. ^lercuric thiocyanate, sparingly soluble: potassium mercuric thiocyanate.
more soluble. Silrer thiorj/annte is insoluble in water, insoluble in dilute nitric
acid, slowly soluble in ammonium hydroxide.
Solutions of metallic thiocyanates give, with soluble salts of:
Cobalt, very concentrated, a blue color, Co(CNS)2 . crystallizable in blue
needles, soluble in alcohol, not in carbon disulphide. The coloration is
promoted by warming, and the test is best niade in an evaporating dish.
In strictly neutral solutions, iron, nickel, zinc and manganese, do not
interfere.
Copper, if concentrated, a black crystalline precipitate, Cu(CNS), , soluble in
thiocyanate. With sulphurous acid, a white precipitate, CuCNS: also with
hvdrosulphuric acid (used to separate a thiocyanate from a chloride)
(Mann, Z., 1889, 28, 068).
Iron (ferrous), no precipitate or color.
Iron (ferric), an intensely blood-red solution of Fe(CNS)j , decolored by solu-
tion of mercuric chloride (§126, 0^, dlHtinvHon from acetic acid); decolored
by phosphoric, arsenic, oxalic and iodic acids, etc., unless with excess of
ferric salt: decolored by alkalis and by nitric acid, not by dilute hydro-
chloric acid. On introduction of metallic zinc, it evolves hydrosulphuric
acid. Ferric thiocyanate is soluble in ether, which extracts traces of it
from aqueous mixtures, rendering its color much more evident by the
concentration in the ether layer.
Xead. gradually, a yellowish crystalline precipitate, Pb(CNS)3 , changed by
I boiling to white basic salt.
Mercury (mercurous), a white precipitate, HgCNS , resolved by boiling into
Hg and Hg(CNS)3 . The mercurous thiocyanate, HgCNS , swells greatly
on iffnUion (being used in *' Pharaoh's serpents '*), with evolution of mer-
cury, nitrogen, thiocyauogen, cyanogen and sulphur dioxide.
§286, 1. NITROQEN. 273
Hercary (mercuric), in solutions not very dilute, a white precipitate,
Hg(GNS)s , somewhat soluble in excess of the thiocyanates, sparingly
soluble in water, moderately soluble in alcohol. On ignition, it swells like
the mercurous precipitate,
^tinum. Platinic chloride, giadually added to a hot, concentrated solution
of potassium thiocyanate, forms a deep-red solution of double thioeyanate of
potassium and platinum (KCNS) .FtlCNS)* , or more properly. K,Pt(CNS)«.
potassium thiocyaiwplatinale. The latter salt gives bright-colored precipi-
tates with metallic salts. The thiocyanoplatinate of lead (so formed) is
golden-colored; that of silver, orange-red.
Silver, a white precipitate, AgCNS , insoluble in water, insoluble in dilute
nitric acid, slowly soluble in ammonium hydroxide, readily soluble in excess
of potassium thiocyanate; blackens in the light; soluble in hot concentrated
H1SO4 (separation from AgCl) (Volhard, A., 1877, 190, 1).
Certain active oxidizing agents, viz., nascent chlorin'e, and nitric acid contain-
ing- nitrogen oxides, acting in hot, concentrated solution of thiocyanates, pre-
cipitate perthiocya-noffcn, H(CNS)s , of a yellow-red to rose-red color, even blue
•ometimes. It may be formed in the test for iodine, and mistaken for that
element, in starch or carbon disulphide. If boiled with solution of potassium
hydroxide, it forms thiocyanate.
Concentrated hydrochloric acid, or sulphuric acid, added in excess to water
solution of thiocyanates, causes the gradual formation of a yellow precipitate,
perthuH'!fanic acid, (HCN)^Sa , slightly soluble in hot water, from which it
<!rj'8tallizes in yellow needles. It dissolves in alcohol and in ether.
Potas.sium thiocyanate can be^ fused in closed vessels, without decomposition;
but with free access of air, it is resolved into sulphate and cyanate, with
evolution of sulphurous acid.
'When thiocyanic acid is oxidized, the flnal product, as far as the sulphur is
concerned, is always sulphuric acid or a sulphate. In many cases (in acid mix-
ture) it has been proven that the cyanogen is evolved as hydrocyanic acid.
In other cases the fame reaction is assumed as probable.
PbO, and PbjO* form Pb" and sulphuric acid, in acid mixture only (Hardow,
./. C. 1851), 11, 174).
H,A804 forms HaAsOs , hydrocyanic and sulphuric acids.
^^'" forms Co" , hydrocyanic and sulphuric acids.
Nl'" forms Ni" , hydrocyanic and sulphuric acids.
Crvi forms Cr"' , hydrocyanic and sulphuric acids.
Mn^+n forms Ifn" \ hydrocyanic and sulphuric acids. In alkaline mixture, a
cyanate and sulphate are formed (\Vurtz\s Dirt. Chim,, 3, 05).
HJO, forms sulphuric acid and nitric oxide.
HJi'O, forms sulphuric acid and nitric oxide.
CI forms at first a red compound of unknown composition, then HCI , H2SO4
and HCK are produced. In alkaline mixture a chloride and sulphate are
formed.
HCIO same as with CI .
MClOa forms sulphuric, hydrochloric and hydrocyanic acids.
Br forms Iffir and H^SO, : but with alkalis, a bromide and sulphate.
MBrOg forms HBr and HaSO^ .
HXP» /onns HsSO« and free iodine.
§235. Nitrogen. N = 14.04 . Valence one to five.
X. •^*«>P«rtt«s.— Weight of molecule, N, , 28.08. Vapor density, 14 (Jolly, W.
g^ ^^7^* ^ 536). At — 12.3.8**, under pressure of 42.1 atmospheres, it condenses
■ftJISU^^ (Sarrau, C. r., 1882, 94, 718). Boiling point. —194.4° (Olszewski, ^^'. .1.,
' Liquid nitrogen is colorless and trausjiarent. The gas is tastc-
3 colorless. Not poisonous, but kills by excluding air from the
•n or support combusion. It is very inert, not attacking
T*« simplest combinations are the following: N— '"H%. ,
T^Oi . The number of organic compounds contain-
The nitrogen In all compounds that are the
274 HTDROyjTRJC ACID. §235,2.
immediate products of vegfetable (growth has a valence of minus three and
may without change of bouds be converted into N— "'H', . This statement is
made with a limited knowledge of the facts and without, at present, having
conclusive proof; and merely predicting that future research will verify it.
2. Occurrence. — It constitutes about four-fifths of the volume of the atmos-
phere. It occurs as a nitrate in various salts and in various forms as a con-
stituent of animal and vegetable growths.
3. Formation. — (a) From the air, the oxygen being removed by red-hot
copper, the COj by potassium hydroxide, the ammonia and watet* by passing
through H2SO4 . (6) Ignition of ammonium dichromate, (NH4)aCr20T = N, +
CTaO, -f 4H3O . (c) By heating ammonium nitrate and peroxide of manganese
to about 200° (Gatehouse, C. .V., 1877, 35, 118). (r/) Ignition of NH^Cl and
K^Ct^OjI 2NH4CI -f KsCr,0, = 2KC1 + N, + Cr,0, + 4HaO . Unless the
temperature be carefully guarded traces of NO are formed, which may be
removed by passing the gases through FeSO^ . (f) Action of chlorine upon
NH,: SNHs + 301, = 6NH4OI + N, . The NH, must be kept in excess to
avoid the formation of the dangerously explosive chloride of nitrogen, NClj .
if) Removing the oxygen from the air by shaking with NH4OH and copper
turnings, (g) Burning phosphorus in air over water, (h) By passing air
through a mixture of FeS and sawdust; then through a pyrogallate solution,
and finally through concentrated H2SO4 . (i) By shaking" air with Fe(OH),
and Mn(6H)3 . {}) By passing air through an alkaline pyrogallate. (A-) By
passing air, from which CO, has been removed, mixed with hydrogen over
heated platinum black, the hydrogen having been added in just suflScient
quantity to form water with all the oxygen (Damoulin, »/., 1851, 321). (I) By
warming a concentrated solution of NH4NO2 or a mixture of KNO, and H'H^Cl:
NH4NO2 = N, + 2H2O . Potassium dichromate is added to oxidize to nitric
acid any of the oxides of nitrogen that may be formed (Gibbs, B., 1877, 1387).
(m) Bv action of potassium or sodium livpobroniite upon ammonium chloride:
3NaBrO + 2NH,C1 = N2 '-f 3NaBr -h 2HC1 -h 3H2O .
4. Preparation. — Nitrogen has been economically produced by most of the
above methods.
5. Solubilities. — Nitrogen is nearly insoluble in all known liquids.
(). Keactions. — At ordinary temperatures nitrogen is not acted upon by other
compounds. Nodules grooving on the roots of leguminous plants absorb nitro-
gen and build up nitrogenous compounds therewith.
7. Ignition. — I'nder electric influence it combines slowly with hj'drogen:
also with B , Or , Mg , Si and V .
8. Detection.— Nitrogen is more easily detected by the nature of its com-
pounds than by the properties of the liberated element.
9. Estimation. — (a) As free nitrogen by measuring the volume of the gas.
(b) By oxidation of the organic substance with hot concentrated H2SO4 , which
also converts the nitrogen into ammonium sulphate. For details, see works
on organic analysis, (c) By decomposition of the organic material with potas-
sium permanganate in strong alkaline solution, forming ammonia, (d) By
combustion of the organic compound in presence of OuO and Cu® , absorbing
the COa by KOH and determining the nitrogen by volume.
§236. Hydronitric acid (Azoimide). N3H: = 43.128 .
N
Constitution, || ^NH
N
Curtius, B., 1890, 23, 3023. A clear mobile liquid of a penetrating odor, a
very irritative effect upon the nostrils and the skin, and readily exploding
with exceeding violence. Boiling point, about 37°, Soluble in water and
alcohol. An acid of marked activity, dissolving a number of metals with
evolution of hydrogen. Its salts, the trinitrides of the metals of the alkalis
§288,6. NITROUS OXIDE— NITRIC OXIDE, 275
and the alkaline earths, are soluble in water and crystallizable (Dennis, J, Am.
>Soc., 1898, 20, 225). Potassium trinitride precipitates from thorium salts, the
hydroxide of this metal in quantitative separation from cerium, lanthanum,
neodymium and praseodymium (Dennis, J. Am, Soc, 1896, 18, 947). Hydro-
nitric acid is formed by treating ammonia with sodium, and the resulting
sodamide, NaNH, , with nitrous oxide: 2NaNHa + N,0 = NaN, + NaOH -f
NH, (Wislicenus, B,, 1892, 25, 2084).
§237. Nitrous oxide. NgO = 44.08 .
N'^O-", N — — N.
Nitrous oxide becomes a colorless liquid at 0® under pressure of three '
atmospheres (Farady, A., 1845, 56, 157). Melts at —99*' and boils at —92**
(Wills, J. C, 1874, 27, 21). It is a colorless gas with slight sweetish smell and
taste. Supports combustion. When breathed acts as an anaesthetic of short
duration; and is used in dentistry for that purpose. Decomposed by heat
completely at 900" into N and O (Meyer, Pyrochemisch. Vntersuch., 1885). Passed
over red-hot iron N and Fe^O, are formed. K and Na burn in nitrous oxide,
liberating the nitrogen. As a rule both gases and solids that burn in air burn
also in nitrous oxide. It is formed: (a) By heating ammonium nitrate in a
retort from 170*» to 260": NH4NO3 = N^O -|- 2HjO . (&) By passing NO through
solution of SO2 . (c) By action of HNO,; sp. gr., 1.42, diluted with an equal
volume of water, upon metallic zinc, (d) A mixture of five parts of SnCl, , ten
parts of HCl , ap, gr„ 1.21, and nine parts of HNO, , sp. /;r., 1.3, is heated to
boiling: 2HN0| + 4SnCla + 8HC1 = 4S11CI4 + N^O + SH^O (Campari, J, C,
1889, 55, 569).
§238. Nitric oxide. NO = 30.04 . ;^. r ./ ;
N"0-", N = .
1. Properties. — The vapor density (15) shows the molecule to be NO (Daccomo
and Meyer, B., 1887, 20, 1832). Under pressure of one atrbosphere it is
liquified at — 153.6**, and under 71.2 atmospheres at — 93.5**, and solidifies at
— 167° (Olszewski, C. r., 1877, 85, 1016). Odor and taste unknown, on account of
its immediate conversion into NO2 on exposure to the air.
2. Occurrence. — Not found free in nature.
3. Formation. — («) Reduction of nitric acid by means of ferrous sulphate
))reviously acidulated with HjSO^ . (h) Action of cold nitric acid, sp. gr., 1.2,
upon metallic copper; unless great care be used other oxides of nitrogen are
produced, (c) SO2 is passed into slightly warmed HNOg , sp. gr.^ 1.15, and
excess of SO2 removed by passing through water, (f/) According to Emich
(-1/., 1893, 18, 73), a strictly pure nitric oxide is made by treating mercury
with a mixture of nitric and sulphuric acids.
5. Solubilities. — Soluble in about ten volumes of water and in five volumes
of nitric acid, sp. gr., 1.3. One hundred volumes of H.SO^ , sp. gr., 1.84, and
1.50, dissolve 3.5 and 1.7 volumes respectively (Lunge, B., 1885, 18, 1391). A
16 per cent solution of ferrous sulphate dissolves six times its own volume of
the gas forming the '* brown ring," which is decomposed at 100°. Soluble in
CS2 and in alcohol.
6. Beactions. — When heated in nitric oxide to 450°, Ag , Hg and Al are un-
changed; filings of Cu , Fe , Cd and Zn are superficially oxidized, but lead is
completely changed to PbO: while if the metals are in an exceedingly fine
.state of division (by reduction of their oxides by hydrogen), Ni at 200° be-
comes NiO , Fe at 200° forms FeO , Cu at 200° forms CUeO; the higher oxides of
these metals not being thus produced (Sabatier and Senderens. C. r.. 1802. 114,
1429). Oxidized to KNO, bv KMnO,: KMnO^ + NO = MnO, + KNO, (Wank-
lyn and Cooper, Phil. Mug./lSlS, (5), 6, 288).
276 yiTRois ACID, §8M,L
§289. Nitrous acid. HNO^ = ^7.048 .
H'N'"0-"o,H — — N=:0.
1. Properties. — Nitrous acid is known only in solution. Made by adding
KaO, to water. It has a blue color and, owing to its tendency to dissociation
(6HNO, = 2HNO3 + 4N0 -f 2H5O), is very unstable (Fremy, C\ r., 1870. 70,61).
Nitrous anhydride is obtained when a mixture of one volume of oxygen ami
four volumes of nitric oxide are passed through a hot tube, 4X0 -h O, = 25,0-, .
It is a deep red gas, condensing to a blue liquid at 14.4° under 755 mm. pressure
(Gains, C. .V., 1883, 48, 97).
2. Occurrence. — Traces of ammonium nitrite are found in the air, in rain
water, river water and in Chili saltpeter. When found in nature it is usually
accompanied by nitrates.
3. Formation. — By action of nitric acid, sp, gr,, 1.35, upon starch or arsenous
oxide. At 70** nearly pure N^O. is obtained, which passed into cold water
forms HNO3 . Nitrites of potassium and sodium may be formed by ignition
of their nitrates (a prolonged high heat forming the oxides). Or the alkali
nitrites may be made by fusing the nitrates with finely divided iron; lead
nitrite by fusing lead nitrate with metallic lead, and sflver nitrite may be
made from these by precipitation; and from this salt many nitrites may be
made nearly pure by transi^osition; c. f/., BaCl, -f- 2AgNOs = Ba(NOj), +
2AgCl and then BaCNOJ, + ZnSO^ = ZnCNOa)^ + BaSO^ .
4. Preparation. — Same as above.
5. Solubilities. — Silver nitrite is only sparingly soluble (120 parts of cold
water). The other normal nitrites are soluble; but many basic nitrites are
insoluble.
Nascent hydrogen in presence of an alkali reduces nitrates to nitrites; e.p..
sodium amalgam, aluminum wire in hot KOH , etc. Used in excess the nascent
hydrogen reduces the nitrogen still further, forming NH3 .
G. Keactions. — .4. — With metals and their compounds. — Nitrous acid actji
sometimes as an oxidizer, sometimes as a reducer; in the former case NO is
vsually produced (under some conditions N.O , N and NH, are formed): in the
latter case nitric acid is the usual product, but sometimes NO, is produced.
1, PbOa becomes Pb" and nitric acid.
2. Hgf' becomes Hg** and nitric acid.
S, Crvi becomes Cr'" and nitric acid.
4. Co" becomes Co'" and nitric oxide. Excess of KNO, with acetic acid is
used to separate cobalt from nickel (§132, G<*).
5. Ni'" becomes Ni" and nitric acid.
G. Mn" + n becomes Mn" and nitric acid.
B. — With non-metals and their compounds. —
U H^FeCCN),, becomes fir^-t HaFefCN), and then hydronitroferricyanic acid.
Solution of indigo in sulphuric acid is bleached by nitrites.
2, Nitrites are decomposed by nitric acid.
S. HH,PO. becomes II3PO4 and NO.
-). HjS does not displace or transpose alkali nitrites, but if acetic acid Im?
added to liberate the nitrous acid, then S° and NO are produced. HxSO, be-
comes H2SO4 and chiefly NO . With excess of H.SO, , NjO or NH, is formed.
See Weber, /*o//f/., ISOC, 127, 54.1, and 18G7, 130, 277; Fremy, G. r., 1870, 70, 61.
'). HCIO3 becomes Cl° and HNO, .
6\ HBrOs becomes Br° and HNO,
7. HI becomes 1° and NO .
HIO3 becomes 1° and HNO, .
7. Ignition. — In general nitrites are changed to oxides, but with jjotassiam
and sodium nitrites a white heat is required, and with nitrites of Ag , Hg.
Au and Pt the dissociation goes a step further, the free metals being produced.
8. Detection. — (/) Formation of brown ring when a nitrite is acidulated with
acetic acid. Nitrates require a stronger acid for their transposition. (2) A
§241,4. NITROGEN PEROXIDE— NITRIC ACID, 277
mixture of a nitrite and KI liberates iodine on addition of acetic acid (nitrates
requiring a stronger acid for transposition). (S) Nitrous acid with iodic acid
liberates iodine, and nitric acid is produced. (4) Solution of potassium per-
manganate acidified with sulphuric acid is reduced by nitrites (distinction from
nitrates).
9. ElBtiznation. — Acidif}' with acetic acid, distil and titrate the distillate with
standard solution of permanganate.
§240. Nitrogen peroxide (dioxide). N02 = 46.04.
Vapor density, 23 (Ramsay, J, C, 1890, 57, 590). Melting point, —10'
(Deville and Troost, C. r., 1867, 64, 257). Boils at 21.64** (Thorpe, J, C, 188U,
37, 224). Below —10** it is a white crystalline solid. Between —10** and 21.64°
a liquid; nearly colorless at — 9**, yellow at 0°. At 21.64**, orange, growing
nearly black as the temperature rises. The gas does not support combustion
of ordinary fuels, and is poisonous when inhaled. It dissolves in water, form-
ing a greenish-blue solution containing nitrous and nitric acids. With an
aqueous solution of a fixed alkali a nitrate and nitrite are formed: 2N0, -+-
2K0H = KNO, -h KNO3 -f- HaO .
§241. Nitric acid. HNO3 = 63.048 .
II
H'ir'0-"3, H — — N = 0.
1. Properties. — Nitric anhydride, N2O5 , is a colorless solid, melting at 30°
-with partial decomposition to NOj and O, and if exposed to direct sunlight
decomposition begins at lower temperatures.
Nitric acid, HNOg , has not been perfectly isolated; that containing 90.8 per
cent of HNOg is a colorless highly corrosive liquid (Roscoe, A., 1860, 116, 211),
solidifies at 47** (Berthelot), boils at 86**, but dissociation begins at a lower
temperature and is complete at 255**: 4HN0, = 4N0a -f- 2H,0 + Og (Carius,
B.y 1871, 4, 828). If the very dilute acid be boiled, it becomes stronger, and
if a very strong acid be boiled it becomes weaker, in both cases a sp. gr, of
1.42 and boiling point of 120° is reached; the acid then contains about 70 per
cent of HNO, (Kolbe, A. C/*., 1867 (4), 10, 136). This is the acid usually
placed on the market. The reagent usually employed has a «p. gr, of 1.2
(Fresenius standard). The so-called fuming acid has a specific gravity of 1.50
to 1.52. The stronger acid should be kept in a cool dark place to avoid decom-
position.
2. Occurrence. — Found in nature as nitrates of K , Na , NH* , Ca , Mg , and
of a few other metals, the most abundant supply coming from Chili and
Bolivia as sodium nitrate, ** Chili saltpeter."
3. Formation. — (o) Oxidation of nitrogenous matter in presence of air,
moisture and an oxide or alkali; (&) by oxidation of NO, NaO, or NO, by
oxygen (or air) in presence of moisture; (c) from NH, , by passing a mixture
of NHs and oxygen through red-hot tubes.
4. Preparation. — By treating nitrates with sulphuric acid and distilling.-
Nitrates may be made: (a) By dissolving the metal in nitric acid, except
those whose metals are not attacked by that acid, e. g., Au , Pt , Al and Cr ;
and also, antimony forms SboOj , arsenic, H3ASO4 and with excess of hot
acid tin forms metastannic acid H^o^i^oOis • (^) By adding HNO3 to the
oxides, hydroxides or carbonates. All the known nitrates can be made
278 NITRIC ACID. §241,5.
in this manner, (c) By long continued boiling the chlorides of all ordi-
nary metals are completely decomposed, no chlorine remaining, except
the chlorides of Hg , Ag , An and Pt , which are not attacked, and the
chlorides of tin and antimony, which are changed to oxides. (Wurtz,
Am. S., 1858, 75, 371; Johnson, Proc. Am. Ass. ScL, 1894, 163.)
The anhydride is made: (a) By passing chlorine over silver nitrate:
4AgN03 + 2CI2 = 4AgCl + 2N2O5 + O2 . (6) By adding anhydrous P^O,
to HNO3: 2irao3 + P2O5 = 2HPO3 + NjOb .
o. Solubilities. — All normal nitrates are soluble. A few are decom-
posed by water, e. g., 81(^03)3 + HoO = BiON03 + 2HNO3 . Most
nitrates are less soluble in nitric acid than in water, e. //., Cd , Pb , Ba , etc.;
the barium nitrate being completely insoluble in ILNO3 , sp, gr., 1.42.
Nitric acid decomposes the sulphides of all ordinary metals, except
mercuric sulphide which by long continued boilin;? with the concentrated
acid becomes 2HgS.Hg(N03)2 , insoluble in the acid.
6. Reactions. .1. — With metals and their compounds. — Nitric acid is
a powerful oxidizer but unless warmed acts more slowly than chlorine.
It can never be a reducer. The following products are formed: H,
NH3 , H2NOH *, N , N2O , NO , HNO2 , NO2 . If the acid is concentrated,
in excess and hot, the product is usually entirely nitric oxide, colorless,
but changing to the red colored NO2 by coming in contact with the air.
Excess of the reducer, low temperatures and dilute solutions favor the
production of nitrogen compounds having lower valence and of hydrogen.
Nascent hydrogen usually forms NH^ , always the ultimate product if the
hydrogen be produced in alkaline mixture.
Nitric acid oxidizes all ordinary metals, (It does not act upon chro-
mium, gold or platinum.) It forms nitrates, except in the case of tin,
antimony, and arsenic, with which it forms HioSn-Oj,^ , SbgO^ , and H^AsO^.
With the respective metals it forms Hg' or Hg", Sn" or Sn"", As'" or As^',
Sb'" or Sb^, Fe" or Fe'", according to the amount of nitric acid employed.
With copper it forms cupric nitrate (never cuprous); with cobalt it forms
cobaltous nitrate.
• Hydroxylamlne, WH.OH, is formed by the reducing action of Sn and HCl ui>on NO.NjCj.
HN 0„ etc. (Ix)ssen, X., 1888, 252, 170); also by the action of H^S, SO,, K, Na, M g. Zii, and Al upon
HNO3. or by the action of H3S uiwn certain nitrates (Divers and Haf?a, G. iV., 1886, 54, 2T1 . By
action of sodium amalgam upon sodium nitrite solution, NHgOH is produced along with nitrous
oxide, free nitrogen, ammonia, sodium hyponitritv, and sodium hydroxide, the highest yield of
the hydroxylamine being obf.iined when the nitrite solution is as dilute as one in fifty, the iniX'
ture kept cold (Divers, J. C, 1S99, 75, 87 and 89-. It is a base with an alkaline reaction an«l &
strong reducing a.'irent. When pure it is a crystalline solid, odorless, molting at 33.05°, boilln^at
58°at 22 mm. pre?-.' uro ; oxidized by oxygen to IIWO, (Lobry dc liruyn, B., ISftJ, 25, 3, HO andft^-
It is a good antisc:)tic and preservative. It combines with acids to form salts : N HjOH + II( 1=
NH,OH . Htl. Hydroxylamine hydrochloride is decomposed by alkalis forming the free base*
which is decomposed by the halogens, KAIn04, KgCrjO,, BaOj and FbO,. Its solution in ether
reacts with sodium forming a white precipitate of NHgONa.
241, 7. SUTRIC ACID. 279
1. PbOj is not changed. PbaO^ is changed thus: PbgO^ + 4HN0.T =
•bO^ + 2Pb(N03)2 + 2H2O .
2. Hg' becomes H|r"-
3. Sn" becomes Sn'^. Stannous chloride and hydrochloric acid, heated
ith a nitrate, form stannic chloride, and convert nitric acid to ammonia
ivhich remains as ammonium salt). See §71, Gc.
4. Sb'" becomes Sb^, forming SbjO^ , insoluble.
5. As'" becomes As^, forming HjAsO^ .
6. Cu' becomes Cu".
7. Fc" becomes Fe'".
B. — With non-metals and their compounds.
1, Carbon (ordinary, not graphite) becomes COg if the nitric acid be
ot and concentrated.
H2C0O4 becomes COg , in hot concentrated acid.
H4Fe(CN)e becomes first H.iFe(CN)e and then hydronitrof erricyanic acid.
HCNS is oxidized, the sulphur becoming H2SO4 .
2, Xitrites are all decomposed, nitrates being formed, the nitric acid
ot being reduced. The nitrous acid liberated immediately dissociates:
iSO^ — 2N0 + HNO3 + H2O .
S. P°, PH3 , HH.PO0 and H3PO8 become H3PO4 . That is P^-" becomes
»v
4. S becomes H2SO4 .
HjS becomes first S** and then HgSO^ .
H2SO3 becomes RJ&O^ ; and in general S^-" becomes S^^.
5. HCl , nitrohydrochloric acid: 2HNO3 + 6HC1 = 2N0 + 4H2O + SClg
Koninck and Nihoul, Z. anorg,, 1890, 477). See §269, 6B2.
HCIO3 is not reduced. Chlorates are all transposed but not decom-
osed until the temperature and degree of concentration is reached that
rould dissociate the HCIO3 if the nitric acid were absent.
6. Br° is not oxidized. HBr becomes Br° and is not further oxidized,
ill bromates are transposed but the HBrOs is not decomposed until a tem-
erature and degree of concentration is reached that would cause the
issociation of the HBr03 if the nitric acid were absent.
7. 1° becomes HIO3. Very slowly unless the fuming nitric be used.
HI become first 1° ; then as above.
8. In general organic compounds are oxidized. Straw, hay, cotton, etc.,
re inflamed by the strong acid (Kraut, B,, 1881, 14, 301). For action
n starch, see Lunge, B., 1878, 11, 1229, 1641. With many organic bodies
abstitv.tion products are formed, the oxides of nitrogen taking the place
f the hydrogen.
7. Ignition.— Nitric acid is dissociated by heat: 4HN0a = 4N0, + 2H,0 + Oj,
jmplete if at 256° (Carius, B., 1871, 4, 828). No nitrates are volatile as such;
2S0 NITRIC ACID, §241,8,
ammonium nitrate is dissociated: NH4NO3 = N,0 -+- 2H,0. Some nitrates, e.g^
those of K and Na , are first changed to nitrites with evolution of oxygen only,
and at an intense white heat further changed to oxides with evolution of H,'0
as well as oxygen. As a final result of ignition the nitrates of all ordinary
metals are left as oxides, except that those of Hg* , Ag , Au and Pt are reduced
to the free metal.
A mixture of potassium nitrate and sodium carbonate in a state of fusion
is a powerful oxidizer; e.g., changing Sn" to Sniv , As'" to Asv , Sb'" to Sbv,
Pe" to Fe"' , Cr'" to Crvi , Mnvi-n to Mnvi , Svi-n to S^^ , etc.
Heated on charcoal, or with potassium cyanide, or sugar, sulphur or other
easily oxidizable substance (as in gunpowder), nitrates are reduced ^vith
deftat/ration or explosion, more or less violent. With potassium cyanide, on
platinum foil, the deflagration is especially vivid. In this reaction free nitrof/en
is evolved.
Strongly heated with excess of potassii«m hydroxide and sugar or other
carbonaceous compound, in a dry mixture, nitrates are reduced to ammonia,
which is evolved, and may be detected. In this carbonaceous mixture, the
nitrogen of nitrates reacts with alkalis, like the unoxidized nitrogen in car-
bonaceous compounds.
8. Detection.— Most of the tests for the identification of nitric acid are
made by its deoxidation, disengaging a lower oxide of nitrogen, or even,
"by complete deoxidation, forming ammonia.
If, with concentrated sulphuric acid, a bit of copper turning, or a crystal
of ferrous sulphate, is added to a concentrated solution or residue of
nitrate, the mixture gives off abundant brown vapors; the colorless mtric
oxide, NO , which is set free from the mixture, oxidizing immediately in
the air to nitrogen peroxide, NOo :
2KN0, -f 4H,S0, + 3Cu = K^SO^ + 3CuS0« + 4H,0 -f- 2N0
2KN0, + 4H,S0* 4- CFeSO, = K,SO, + 3Fe,(S0Ja -f- 4H2O -f 2N0
The three atoms of oxygen furnished hy two molecules of nitrate suffice to
oxidize three atoms of copper; so that 3CuO with SH.SO^ , may fonn
3CuS0^ and SHoO . The same three atoms of oxygen (having six bonds)
suffice to oxidize six molecules of ferrous salt into three molecule? of
ferric salt; so that GFeSO, with 3H->S0^ , can form SFCjCSO,)., and ^HjO-
Now if, by the last-named reaction, the nitric oxide is disengaged in
cold solution, with excess of ferrous salt and of sulphuric acid, instead
of passing off, the nitric oxide combines with the ferrous salt, forming a
hlack'lrown liquid, ('Et^O^)^0 , decomposed l)y heat and otherwise un-
stable: 2KNO3 + 4H2SO, + lOFeSO, = KoSO.V SFe^CSOJg + ^H^O +
2(FeS0,)2N0 .
a. — This exceedingly delicate " Brown ring " test for nitric acid or
nitrates in solution may be conducted as follows: If the solution of '^
nitrate is mixed with an equal volimie of concentrated HoSO^ , the mixture
allowed to cool and a concentrated solution of FeSO^ then cautiously added
to it, so that the fluids do not mix, the junction shows at first a ])urple,
afterwards a brown color (Fresenius, Qual. Anal., 16th ed., 387). A second
method of obtaining the same brown ring is: Take sulphuric acid to a
§241, Sh. XITRIC ACID. 281
quarter of an inch in depth in the test-tube; add without shaking a nearly
equal bulk of a solution of ferrous sulphate, cool; then add slowly of the
solution to be tested for nitric acid, slightly tapping the test-tube on the
side but not shaking it. The brown ring forms between the two layers of
the liquid. A third method often preferred is: Take ferrous sulphate
solution to half an inch in depth in the test-tube ; add two or three drops
of the liquid under examination and mix thoroughly; incline the test-tube
and add an equal volume of concentrated H2SO4 in such a way that it will
pass to the bottom and form a separate layer. Cool and let it stand a
few minutes without shaking.
h, — ^Indigo solntion. — In presence of HCl heat moderately and blue
color is destroyed. Interfering substances, HCIO3 , HIO, , BSrOs , Fc"',
Cr^i, mC^, and all that convert HCl into CI .
c. — Sodinm salicylate is added to the solution, H^SO^ is slowly added,
test-tube being inclined. Avoid shaking, keep cool for five minutes. A
yellow ring indicates ILNO3 . To increase the brilliancy of the color^
shake, cool and add to HIT4OH .
d, — Anunoninm test, — Treat the solution with KOH and Al wire, warm
until gas is evolved. Pass the gas into water containing a few drops of
Xessler's reagent. A yellowish-brown precipitate indicates HITO3 :
3HNO3 + 8A1 -f 8K0H = SNHg + 8EAIO2 + HgO . Nothing interferes
with this test, but action is delayed by CF , P and many other oxidisers.
e. — Nitrite test. — Eeduce the nitrate to nitrite by warming with Al and
KOH . At short intervals decant a portion of the solution, add a drop of
KI, acidify with HC2H3O2 and test for I with CSj . This test should
always be made in connection with (d). Other oxidisers including CP',
Br^, I^, and As^ are reduced before the reduction of the HNO3 begins:
3HNO3 -f- 2AI -h 5K0H = 3KN0, + 2KA10, -f 4H,0
2KN0, -f- 2KI + 4HC,H30, =1,-1- 4KC,H3 0, -f- 2H,0 -f- 2N0
Other means of making the nascent hydrogen are sometimes preferred;
e, g,y sodium amalgam, a mixture of Zn and Fe both finely divided and
used with excess of hot KOH , or finely divided Mg in presence of H3PO4 .
/. — Add three drops of the solution to be tested to two drops of
diphenylamine, {CJS,^)^iSlSL , dissolved in HjSO^ . A blue color indicates
a nitrate, Cl% CF, Br^ F, Mn^ii, Cr^^ Se^^ and Fe'" interfere with this
test.
g, — ^Bruclne, dissolved in concentrated sulphuric acid, treated (on a porcelain
surface) with even traces of nitrates, gives a fine deep-red color, soon paling to
reddish-yellow. If now stannous chloride, dilute solution, be added, a fine red-
violet color appears. (Chloric acid gives the same reaction.)
h, — Phenol, CeHsOH , gives a deep red-brown color with nitric acid, by for-
mation of nitrophenol (mono, di or tri), C,H4(N03)OH to C.H2(N02)sOH ,
** picric acid " or nitrophenic acid. A mixture of one part of phenol (cryst.
carbolic acid), four parts of strong sulphuric acid, and two parts of water,
282 OXYGEN, §241, 8t.
constitutes a reagent for a very delicate test for nitrates (or nitrites), a few
drops being sufficient. With unmixed nitrates the action is explosive, unless
upon very small quantities. The addition of potassium hydroxide deepens and
brightens the color. According to Sprengel {J. C, 1863, 16, 396), the some-
what similar color given by compounds of chlorine, bromine, iodine and by
organic matter may be removed by adding ammonium hydroxide without
diminishing the brightness of the color formed by the nitrates.
i. — According to Lindo (C. .V., 1888, 68, 176), resorcinal is five times more
delicate a test than phenol. Ten grammes of resorcinol are dissolved in 100 cc.
of water; one drop of this solution with one drop of a 15 per cent solution of
HCl and two drops of concentrated H.SO^ are added to 0.5 cc. of the nitrate
to be tested. Nitrous acid gives the same purple color.
/. — A little pyrogfallol is dissolved in the liquid to be tested (less than one
mg. to one cc.) and ten drops of concentrated H38O4 are dropped down the
side of the test tube so as to form two layers; at the surface of contact a
brown or yellow coloration apx>ears if nitric acid is present. One mg. of
nitric acid in one litre of potable water can thus be detected (Curtman, Arch.
Phann,, 1886, 223, 711).
9. Estimation. — (a) If the base is one capable of readily forming a silicate,
the nitrate is fused with SiO, and estimated by the difference in weight, (ft) By
treating with hot sulphuric acid, passing the distillate into BaCOa and esti-
mating the nitric acid by the amount of barium dissolved, (c) Treating with
Al and KOH and estimating the distillate as NH, . (d) Neutralizing the free
acid with ammonium hydroxide, and after evaporation and drying at 115*'.
weighing as ammonium nitrate*, (r) In presence of free H3SO4 a ferrous solu-
tion of known strength is added in excess to the nitrate and the amount of
ferrous salt remaining is determined by a standard solution of potassium
permanganate, (f) The volume of hydrogen generated by the action of potas-
sium hydroxide upon a known quantity of aluminum is measured; and the
test is then repeated under the same conditions, but in presence of the nitrate.
The difference in the volume of the hydrogen obtained represents the quantity
of NHs that has been formed.
§242. Oxygen. = IG.OOO . Usual valence two.
1. Properties. — A colorless, odorless gas: ftperifir ijrarUn, 1.10562 (Crafts, V. r.,
1888, 106, 1662). When heated it diffuses through silver tubing quite rapidly
(Troost, i\ r., 1S84, 98, 1427). It liquifies by cooling the gas under great pres-
sure and then suddenly allowing it to expand under reduced pressure. It ItolU
at — 113° under 50 atmospheres pressure: and at — 184° under one atmosphen'
pressure (Wroblewski, C. r., 1884, 98, 304 and 982). Its critical tcmj}cra1vir is
about — 118°, and the crilical pressure 50 atmospheres. Spci'iflc gravily of the
liquid at —181.4°, 1.124 (Olszewski, 3/., 1887, 8, 73). Oxygen is sparingly soluble
in water with a slight increase in the volume (Winkler, B., 1889, 22, 1764).
Slightly soluble in alcohol (Cariu.s .4., 1855, 94, 134). Molten silver absorbs
about ten volumes of oxygen, giving it up upon cooling (blossoming of silver
beads) (Levol, C. r., 1852, 35, 63). It transmits sound better than air (Bender.
B., 1873, 6, 665). It is not combustible, but supports combustion much better
than air. In an atmosphere of ox^^gen, a glowing splinter bursts into a flanie-
phosphorus burns with vivid incandescence; al.so an iron watch spring heated
with burning sulphur. It is the most negative of all the elements except
fluorine; it combines directly or indirectly with all the elements except fluorine-
with the alkali metals rapidly at ordinary temperature. The combination <^'
oxygen with elements or compounds is termed combustion or oxidation. The
temperature at which the combination takes place varies greatly: Phosphorus
at 60°; hvdrogen in air at 552°; in pure oxvgen at 530° (Mallard and Le Chate-
lier. «/.,'l883, (2), 39, 2); carbon disulphide at 149°; carbon at a red heat;
while the halogens do not combine by heat a4one.
2. Occurrence. — The rocks, clay and sand constituting the main part of the
earth's crust contain from 44 to 48 per cent of oxygen; and as water contains
242,4/. OXYGEN, . 283
5.81 per cent, it has been estimated that one-half of the crust is oxygen,
xcept in atmospheric air, which contains about 23 per cent of uucombined
sygen, it is always found combined.
3. Formation.— (a) By igniting HgO . (b) By heating KCIO, to 350°, KCIO^
produced and oxygen is evolved; at a higher temi^erature the KCIO4 becomes
;C1 . In the presence of MnO. the KCIU, is completely changed to KCl at
)0*', without forming KCIO4 , the MnO^ not being changed. Spongy platinum,
uO , FCzOs , PbOa , etc., may be substituted for MnO, (Mills and Donald, J, C-.,
J82, 41, 18; Baudrimont, Am, fe'., 1872, 103, 370). Spongy platinum, ruthenium,
lodium and indium with chlorine water or with hydrogen peroxide evolve
sygen. The spongy ruthenium acts most energetically (Schoenbein, A. Cft.,
J66, (4), 7, 103). (c) Action of heat on similar salts furnishes oxygen; e. g,,
:C10 and KCIO3 form KCl , KBrO^ forms KBr , KIO, and KIO4 form KI ,
ad KNO, forms KNO, (at a white heat K.O , NO and 6 are formed), (d) By
le action of heat on metallic oxides as shown in the equations below, (e) By
eating higher oxides or their salts with sulphuric acid. Crvi is changed to
r"' , Co"' to Co" , Ni'" to Ni" , Biv to Bi'" , Fevi to Fe"' , Pbiv to Pb'^, and
Cn'^+n to Mn"; in each case a sulphate is formed and oxygen given off:
a, 2HgO (at 500°) = 2Hg + O,
6. lOKClO, (at 350°) = 6KCIO4 -f- 4KC1 + 30, (Teed, J. C, 1887, 51, 283)
2KC10, (at red heat) = 2KC1 + . O^
2KC10, -h nMnO, (at 200°) = nMnO, + 2KC1 + 30,
c. KCIO2 = KCl -f- O,
2KBr03 = 2KBr + 20.
2KI0, = 2KI + nO^
KIO, = KI 4- 20,
2KN0, = 2KN0, -f O,
4KN0, (white heat) = 2K,0 -f- 4N0 + O,
d. 2PbaO, (white heat) = (JPbO + O^
2Sb20, (red heat) = 2Sb,04 + O^
BljOs (red heat) = Bi,0, + O,.
4CrO, (about 200°) = 2Cr,0, + 30,
4K2Cr,07 (red heat) = 2Cr,0s -h 4K2Cr04 + 30,
CFe^Oa (white heat) = 'FCaO, -f O,
3MnO, (white heat) = Mn,©, -f O,
CCo,0, (dull-red heat) = 4Co,0, -f- 0,
2Ni20, (dull-red heat) = 4Ni0 + O,
2AgoO (:)00°) = 4Ag -f O,
2Ba02 (800°) = 2BaO -f O,
e. 2K,Cr,0, + 8H=S0, = 4KCr(S04)5 + 30, + 8H,0
4KMn04 -f 6H,S0, = 2K,S0, + 4MnS04 + 50^ + 6H,0
2Pb304 4- 6H2SO, = GPbSO* 4- 6H,0 + O,
4. Preparation. — (a) By heating KCIO, to 200° in closed retorts in the pres-
ce of MnO. or Fe,03\ If KCIO, be heated alone, higher heat (350°) is
quired, and the gas is given off with explosive violence. About equal parts
the metallic oxide and KCIO. should be taken. (6) BaO heated in the air
550° becomes BaO:, . and at 800° is decomposed into BaO and O , making
eoretieally a cheap process, (c) By heating calcium plumbate. The calcium
umbate is regenerated by heating in the air (Kassner, J, C, 1894, 66, ii, 89).
) By passing snlphiiric acid over red-hot bricks: 2H2SO« = 2SO2 + 2H,0 -+- O2:
e SO2 is separated by water, and after conversion into H0SO4 (§266, 4) is
ed over again, (e) By warming a saturated solution of chloride of lime with
small amount of cobaltic oxide, freshly prepared and moist. The cobaltic
ide seems to play the same role as NO in making H3SO4 (Fleitmann, A. Ch.,
^5, (4), 5, 507). (f) The following cheap process is now employed on a large
lie. Steam is passed over sodium manganate at a dull-red heat; Mn,Ot and
284 • OZOyE, §842,5.
oxygen are formed. Then, without change of apparatus or temperature, air
instead of steam is passed over the mixture of Mn^O. and NaOH . The JblsO^
is thus again oxidized to "NeL^HLnO^ , and tree nitrogen is liberated:
4Na,MnO, -f ^H^O (dull-red heat) = SNaOH -f 2Mn,0, + »0,
8NaOH 4- 2Mn,0, + air, 3(0, + 4N,) = 4Na,MnO« + 4H,0 + 12N,
5. Solubilities.— See 1.
6. Seactions. — Pure oxygon may be breathed for a short time without injury,
A rabbit placed in pure oxygen at 24** lived for three weeks, eating voraciously
all the time, but nevertheless becoming thin. The action of oxygen at 7.2*' is
to produce narcotism and eventually death. When oxygen is cooled by a
freezing mixture it induces so intense a narcotism that operations may be
performed under its influence. Compressed oxygen is ** the most fearful poison
known." The pure gas at a pressure of 3.5 atmospheres, or air at a pressure
of 22 atmospheres, produces violent convulsions, simulating those of strychnia
poisoning, ultimately causing death. The arterial blood in these cases is found
to contain about twice the quantity of its normal oxygen. Further, compressed
oxygen stops fermentation, and permanently destroys the power of yeast.
At varying temperatures oxygen combines directly with all metals except
silver, gold and platinum, and with these it may be made to combine by pre-
cipitation. It combines with all non-metals excei)t fluorine; the combination
occurring directly, at high temperatures, except with CI , Br and I , which
require the intervention of a third body.
7. Ignition. — Most elements when ignited with oxygen combine readily.
Some lower oxides combine with oxygen to form higher oxides, and certain
other oxides evolve oxygen, forming elements or lower oxides. Oxides of gold,
platinum and silver cannot be formed by igniting the metals in oxygen; they
must be formed by precipitation.
8. Detection. — Uncombincd oxygen is detected by its absorption by an alka-
line solution of pyrogallol; by the combination with indigo white to form
indigo blue; by its' combination with colorless NO to form the brown NO,; by
its combination with phosphorus, etc. It is separated from other gases by
its absorption by a solution of chromous chloride, pyrogallol or by phosphorus.
In combination in certain compounds it is liberated in whole or in part by
simple ignition; as with KCIO, , KMnO, , HgO , Au.O, , PtO, , Ag,0 , Sb,0, ,
etc. In other combinations by ignition with hydrogen, forming water.
9. E&timation. — Free oxygen is usually estimated by bringing the gases in
contact with phosphorus or with an alkaline solution of pyrogallol (CO, havinp
been previously removed), and noting the dimunition in volume. Oxygen in
combination is usually estimated by difference.
§243. Ozone. 0,== 48.000.
—
\/
Ozone was first noticed by Van Ma rum in 1785 as a peculiar smelling ga^
formed during the electric discharge; and which destroyed the lustre of
mercury. Schoenbein {Poog., 1840, 60, 61G) named the gas ozone and noticed
its powerful oxidizing x:)roperties. It is said to be an ever-present constituent
of the air, giving to the sky its blue color; present much more in the country
and near the seashore than in the air of cities (Hartley, J. C, 1881, 39, 57 and
111; Houzeau, C, r., 1872, 74, 712). Ozone is always mixed with ordinary oxygen,
partly due to dissociation of the ozone molecule, which is stable only at low
temperatures (Hautefeuille and Chappuis, C. r., 1880, 91, 522 and 815). It is
prepared by the action of the electric discharge upon oxygen (Bichat and
Guntz, C. r., 1888, 107, 344; Wills, B,, 187.3, 6, 7G9). By the oxidation of raoisl
phosphorus at ordinary temperature (Leeds, A., 1879, 198, 30; Marignac, C. r.,
1845, 20, 808). By electrolysis of dilate sulphuric acid, using lead electrodes
§244, 1. HYDROGEN PEROXIDE, 285
(Planti, C, r., 1866» 63, 181). By the action of concentrated sulphuric acid on
potassium permanganate (Schoenbein, J, pr., 1862, 86, 70 and 377). Many
readily oxidized organic substances form some ozone in the process of oxida-
tion (Belluci, B., 1879, 12, 1699). Ozone is a gas, the blue color of which can
be plainly noticed in tubes one metre long. Its odor reminds one somewhat
of chlorine and nitrogen peroxide, noticeable in one part in 500,000. It acts
upon the respiratory organs, making breathing difficult. When somewhat
concentrated it attacks the mucous membrane. It caused death to small
animals which have been made to breathe it. For further concerning the
physiological action, see Binz, C. C, 1873, 72. Its specific (jrmUy is 1.658 (Soret,
A.," 1866, 138, 4). It has been liquified to a deep-blue liquid, brniiny at — 106**
(Olszewski, J/., 1887, 8, 230). The gas is sparingly soluble in water (Carius, li.,
1873, 6, 806). It decomposes somewhat into inactive oxygen at ordinary tem-
perature, and completely when heated above 300**, with increase of volume.
A number of substances decompose ozone without themselves being changed;
e, g.y platinum black, platinum sponge, oxides of gold, silver, iron and copper,
peroxides of lead and manganese, potassium hydroxide, etc. It is one of the
most active oxidizing agents known, the presence of water being necessary.
When ozone acts as an oxidizing agent there is no change in volume: but one-
third of the oxygen entering into the reaction, inactive oxygen remaining.
Moist pzone oxidizes all metals except gold and platinum to the highest pos-
sible oxides.
Pb" becomes PbO,
Sn" becomes SnO,
Hgr' becomes Hg"
Bi'" becomes BijO,
Pd" becomes PdO,
Cr"' becomes Crvi
Fe" becomes PejO. ; in presence of KOH , K^FeOt
Hn^ becomes MnOj ; in presence of HaS04 or HNO, , Hlfn04 is formed.
Co" becomes Co'"
Ni" becomes Ni'" . With the salts of nickel and cobalt the action is slow,
rapid with the moist hydroxides.
K^PeCCN), becomes KsPe(CN).
NjOg becomes HNO, , in absence of water NO, is formed
SOj becomes H.SO^
H3S becomes S and H3O , the sulphur is then oxidized to HaSO« (PoUacci,
C. C, 1884, 484)
P and PH, become H3PO4
HCl becomes CI and HjO
HBr becomes Br and H3O
I becomes HIO3 and HIO* (Ogier, C, r., 1878, 86, 722)
HI and KI become I and H3O , then Iv
Most organic substances are decomposed; indigo is bleached much more
rapidly than by chlorine (Houzeau, C r., 1872, 75, 349).
Alcohol and ether are rapidly oxidized to aldehyde and acetic acid.
Ozone is usually detected by the liberation of iodine from potassium iodide,
potassium iodide starch paper being used. Because HNO2 and many other
substances give the same reaction, thallium hydroxide paper is preferred by
Schoene (B., 1880, 13, 1508). The paper is colored brown, but the reaction is
much less delicate than wn'th potassium iodide starch paper. It is estimated
quantitatively by passing the gas through a solution of KI rendered acid with
HjSOt , and titration of the liberated iodine: O, + 2HI = O^ + I, + HjO .
§244. Hydrogen peroxide. HoOo = 34.016 .
H — — — H.
1. Properties. — Pure hydrogen peroxide (99.1 per cent) is a colorless syrupy
liquid, boiling at 84** to 85® at 68 mm. pressure. It does not readily moisten
the containing vessel. It is volatile in the air, irritating to the skin, and
286 BTDROGEX PEROXIDE. §244^2.
reacts strongly acid to litmus. The ordinary three per cent solution can be
evaporated on the water bath until it contains about 60 per cent H,0, , losing
about one-half by volatilization. The presence of impurities causes its decom-
position witii explosive violence. Before final concentration under reduced
pressure it should be extracted with ether (Wolffenstein, B., 1894, 27, 3307).
The dilute solutions are valuable in surgery in oxidizing putrid flesh of wounds,
etc.; they are quite stable and may be preser\'ed a long time especially if acid
(Hanriott, C. r., 1885, 100, 57). The presence of alkalis decreases the stability.
Concentrated solutions evolve oxygen at 20®, and frequently explode when
heated to nearly lOO**. It contains the most oxygen of any known compound:
one-half of the oxygen being available, the other half combining with tlie
hydrogen to form water.
2. Occurrence. — In rain water and in snow (Houzeau, C. r., 1870, 70, 519).
It is also said to occur in the juices of certain plants.
3. Formation. — (a) By the electrolysis of 70 per cent H3SO4 (Richarz, W. i.,
1887, 31, 912). (6) By the action of ozone upon ether and water (Berthelot,
C r., 1878, 86, 71). (r) By the action of ozone upon dilute ammonium hydroxide
(Carius, B., 1874, 7, 1481 )\ (d) By the decomposition of various peroxides with
acids, (r) By the action of oxygen and water on palladium sponge saturated
with hydrogen (Traube, /?., 1SS3, 16, 1201). (f) By the action of moist air on
phosphorus partly immersed in water (Kingzett, J, C, 1880. 38, 3).
4. Preparation.— BaOj is decomposed by dilute H3SO4 , the BaSO^ bein?
removed by filtration. The BaO, is obtained by heating BaO in air or oxygen
to low redness. At a higher heat the BaO, is decomposed into BaO and
(Thomsen, B., 1874, 7, 73). Sodium peroxide, Na.O, , is formed by heating
sodium in air or oxygen (Harcourt, J. T., 1862, 14, 267); by adding H-0. to
NaOH solution and precipitating with alcohol. Prepared by the latter method
it contains water.
5. Solubilities It is soluble in water in all proportions: also in alcohol.
which solvent it slowly attacks. BaO, is insoluble in water, decomposed by
acids, including CO, and H,SiF, with formation of H^O, . 'NeuO. is soluble in
water with generation of much heat. It is a powerful oxidizing agent.
G. Reactions. .1.— With metals and their compoTinds. — Hydrocren
peroxide usually acts as a powerful oxidizing agent to the extent of 0110-
half its oxygen. Under certain conditions, however, it acts as a strong
reducing agent. Some substances decompose it into HoO and without
changing the substance employed, e, g,, gold, silver, platinum, manganese
dioxide, charcoal, etc. (Kwasnik, B., 1892, 25, 67). Many metals are
oxidized to the liighest oxides, e. //., Al , Fe , Mg , Tl , As , etc. Gold ami
platinum are not attacked.
1. Pb" becomes PbO. (Schoenbein, J. pr., 1862, 86, 129; Jannasch and
Lesinsky, 7?., 1893, 26, '2334).
2. AgoO becomes Ag and .
3. HgO becomes Hg and .
Jf, Au^O^ becomes An and .
5. As'" becomes As^.
G, Sn" becomes Sn^^.
7. Bi'" becomes Bi^'.
8. Cn" in alkaline solution (Fehling's solution) becomes QvlS^ (Hanriott,
Bl, 1886, (2), 46, 468).
9. Fe" becomes Fe'" (Traube, B., 1884, 17, 1062).
10. Tl' becomes Tl^Og (Schoene, 4., 1879, 196, 98).
§244,96. ijyi)JWGEN peroxide. 287
11. Cr"' becomes Cr^ in alkaline mixture (Lenssen, J. pr., 1860, 81,
278).
12. Cr^ with H2SO4 gives a blue color, HCrO^ , perchromic acid, soon
changing to green by reduction to Cr'". By passing the air or vapor
through a chromic acid solution, ozone is separated from hydrogen perox-
ide, the latter being decomposed (Engler and Wild, /?., 1896, 29, 1940).
13. Hn" in alkaline mixture becomes MnOj . In presence of KCN a
separation from Zn (Jannasch and Niedcrhofheim, B., 1891, 24, 3945;
Jannasch, Z. anorg., 1896, 12, 124 and 134).
Hii"+n ^^ii H2SO4 forms M11SO4 , oxygen being evolved both from the
H2O2 and from the Mn compound (Brodie, J. C, 1855, 7, 304; Lunge^
Z. angew.y 1890, 6).
H. BaO , SrO , and CaO become the peroxides.
16. NaOH becomes Na202.8H20 .
16. NH4OH becomes NH^NOj (Weith and Webber, £., 1874, 7, 17 and
45).
B. — With non-metals and their componnds.
1. K4Fe(CN)e becomes KgFeCCN),, (Weltzien, A., 1866, 138, 129); in
alkaline solution the reverse action takes place : 2K3Fe(CN)Q -|" 2K0H +
H2O2 = 2KJ?e{ClSl)^ + 2H2O + O2 (Baumann, Z. angew., 1892, 113).
2. O3 becomes O2 (Schoene, I. c, page 239).
8. H3PO2 becomes H3PO4 .
4. H2S and sulphides, and SO2 and sulphites, become H2SO4 or sulphates
(Classen and Bauer, B., 1883, 16, 1061).
5. CI becomes HCl (Schoene, I. c, page 254). It is a valuable reagent
for the estimation of chloride of lime : CaOCl2 + H2O3 = CaClg + HjO +
O2 (Lunge, Z. angetv., 1890, 6).
6. I becomes HI (Baumann, Z. angew.y 1891, 203 and 328). KCl , EBr ,
and KI liberate oxygen from HgOg but no halogen is set free ; except that
with commercial HoOg free iodine may always l)e obtained from KI
(Schoene, A., 1879, 196, 228; Kingzett, J. C, 1880, 37, 805).
7. Ingition. — The peroxide of barium is formed by ignitiiifc BaO to dull red-
ness; strong" ignition causes decomposition of the BaO, into BaO and O . The
])eroxide of calcium cannot be formed by ignition of lime in air or oxygen.
8. Detection. — In a dilute solution of tincture of guaiac mixed Avith malt
infusion, a blue color is obtained when HjO. is added. To the solution suj)-
posed to contain H2O2 add a few drops of lead acetate; then KI , starch, and a
little acetic acid; with H^Oj a blue color is produced (Schoenbein, /. r.; Struvo.
Z., isno, 8, 274). As confirmatory, its action on KMnO« and on K.Cr.OT should
he observed. A ten per cent solution of ammonium molybdate with equal
parts of concentrated sulphuric acid gives a characteristic deep yellow color
with HjOj (I)eniges, C. r., 1890, 110, 1007; Crismer, Bl.y 1891, (3), 6, 22). H.O.
jrives some extremely delicate color tests with the aniline bases (Ilosvay, B..
]>0.->, 28, 2029; Deniges, J. Phnrm., 1892, (5), 25, 591).
0. Estimation. — (a) By measuring the amount of oxygen liberated with MnOo
(Ilanriott, /??., 1885, (2), 43, 468). (6) By the amount of standard KMn04
288 FLUORINE. §245.
reduced, or by measuring' the volume of oxygen set free: 2KMn04 + SH^SO^ 4-
5H3O3 = K2SO4 -f- 2MnS04 + 8H2O -h 5O2 . (c) By decomposition of XI in
presence of an excess of dilute H2S04; and titration of the liberated iodine with
standard Na^S^Os . ((/) Dissolve a weighed sample of BaO, in dilute HCl , add
KaFe(CN)«; transfer to an azotometer and add XOH . The volume of oxygen
is a measure of the amount of HjO, (Baumann, L c).
§246. Fluorine. F = 19.05 . Valence one.
Since Davy's experiments in 1813, many others have attempted the isolation
of fluorine. In his zeal the unfortunate Louyet fell a victim to the poisonous
fumes which he inhaled. Faraday, Gore, Fremy, and others took up the prob-
lem in succession, but it was not ultimately solved until H. Moissan, in 1886.
produced a gas which the chemical section of the French Academy of Sciences
decided to be fluorine. Many ingenious experiments had been made in order
to obtain fluorine in a separate state, but it was found that it invariably
combined with some portion of the material of the vessel in which the opera-
tion was conducted. The most successful of the early attempts to isolate
fluorine appears to have been made, at the suggestion of Davy, in a vessel of
fluor-spar itself, which could not, of course, be supposed to be in any way
affected by it. Moissan's method was as follows: The hydrofluoric acid having
been very carefully obtained pure, a little potassium hydrofluoride was dis-
solved in it to improve its conducting power, and it was subjected to the action
of the electric current in a U tube of platinum, down the liml>s of which the
electrodes were in.serted; the negative electrode was a rod of platinum, and
the positive was made of an alloy of platinum with 10 per cent of iridium. The
U tube was provided with stoppers of fluor-spar, and platinum delivery tubes
for the gases, and was cooled to — 23°. The gaseous fluorine, which was extri-
cated at the positive electrode, was colorless, and possessed the properties of
chlorine, but uuieh more strongly marked. It decomposed water immediately,
seizing upon its hydrogen, and liberating oxj'gen in the ozonized condition: it
exploded with hydrogen, even in the dark, and combined, with combustion,
with most metals and non-metals, even with boron and silicon in their crystal-
lized modifications. As , Sb , S , I , alcohol, ether, benzol and petroleum took
fire in the gas. Carbon was not attacked by it (Moissan, ISSG, C\ r., 103, 20-
and 25r); ,/. C, 50, 1880, 849 and 97«: A. T/?., 1S91, (6), 24, 224).
riuorine, in several characteristics, appears as the first member of the
Chlorine Series of Elements. It cannot be preserved in the elemental state.
as it combines with the materials of vessels (except fluor-spar), and instantly
decomposes water, forming hf/draffiioric acid, HF , an acid prepared by actintr
on calcium fluoride with sulphuric acid (r/). Fluorine also combines with
silicon as SiF^ , siHron fluoride, a gaseous compound, prepared by acting on
calcium fluoride and silicic anhydride with sulphuric acid (b). On passini?
silicon fluoride into water, a part of it is transposed bj' the water, forroiniT
silicic and hydrofluoric acids (r): but this hydrofluoric acid does not at all
remain free, but combines with the other part of the fluoride of silicon, as
fluosilicic acid {JufdrofluosiUcir arid), (HF)2SiF, or H^SIF, (d) (OfTermann.
Z, anfieic, 1S90, ()17). This acid is uned as a reagent; forming" metallic fl"0'
silicates (silicofluorides), soluble and insoluble (§246).
a, CaF, + H,SO. = CaSO, + 2HF
6. 2CaF, -f SiO, + 2H,S0, = 2CaS0, + 2H3O -f SiF«
c. SiF^ -f 2H,0 = SiO., -h 4HF (not remaining fre«)
d. 2HF -f SiF, = H,SiF«
c and d. 3S1F, -f- 2H3O = SiO, + 2H2SiFe
§247. HYDROFLUORIC ACID-FLUOSILICIC ACID, 289
§246. Hydrofluoric acid. HF = 20.058 .
H'F-', H — F.
A colorless, intensely corrosive gas, soluble in water to a liquid that reddens
litmus, rapidly corrodes glass, porcelain, and the metals, except platinum and
gold (lead but slightly). Both the solution and its vapor act on the flesh as
an insidious and virulent caustic, giving little warning, and causing obstinate
ulcers. The anhydrous acid at 25° has a vapor density of 20, indicating that
the molecule at this temperature is HjFa . But at 100° it is only 10, indicating
that at that temperature the molecule is HP . The anhydrous liquid acid
boils at 19.44° and does not solidify at —34.5°.
The fluorides of the alkali metals are freely soluble in water, the solutions
alkaline to litmuc and slightly corrosive to glass: the fluorides of the alkaline
earth metals are insoluble in water; of copper, lead, zinc and ferricum, spar-
ingly soluble; of silver and mercury readily soluble. Fluorides are identified
by the action of the acid upon glass.
Calcium chloride solution forms, in solution of fluorides or of hydrofluoric
acid, a gelatinous and transparent precipitate of calcium fluoride, CaFg , slightly
soluble in cold hydrochloric or nitric acid and in ammonium chloride solution.
Barium chloride precipitates, from free hydrofluoric acid less perfectly than
from fluorides, the voluminous, white, barium fluoride, BaFj . Silver nitrate
gives no precipitate.
Sulphuric acid transposes fluorides, forming hydrofluoric acid, HP (§245, a).
The gas is distinguished from other substances by etching hard ijla^s — previously
prepared by coating imperviously with (melted) wax, and writing through the
coat. The operation may be conducted in a small leaden tray, or cup formed
of sheet lead: the pulverized fluoride being mixed with sulphuric acid to the
consistence of paste.
If the fluoride be mixed with silicic acid, we have, instead of hydrofluoric
acid, niiicon fluoride, SiP* (§245, 6); a gas which does not attack glass, but when
passed into water produces fluosilicic acid, H.SiF^ (§245, c and rf). See below.
Also, heated with acid sulphate of potassium, in the dry way, fluorides dis-
engage hydrofluoric acid. If this operation be performed in a small test-tube,
the surface of the glass above the material is corroded and roughened: CaF^ -f
2KHSO4 = CaS04 -h K2SO4 -f- 2HF. By heating a mixture of borax, acid
sulphate of potassium, and a fluoride, fused to a bead on the loop of platinum
wire, in the clear flame of the Bunsen gas-lamp, an evanescent yellowish-green
color is imparted to the flame.
§247. Fluosilicic acid. H2SiFe= 144.716.
Fluosilicic acid* (hf/dro fluosilicic acid), (HF)2SiF4 , or HsSiF. , is soluble in
water and forms metallic fluosilicates (silicftfluorides), mostly soluble in water;
those of barium (§186, 6t), sodium and potassium, being only slightly soluble
in water, and made quite insoluble by addition of alcohol.
Potassium fluosilicate is precipitated translucent and gelatinous. Ammonium
hydroxide precipitates silicic acid with formation of ammonium fluoride. With
concentrated sulphuric acid, they disengage hydrofluoric acid, HF . By heat,
they are resolved into fluorides and silicon fluoride: BaSiF« = BaFj + SiF* .
• Fluosilicic acid is directed to be prepared by taking one part each of fine sand and finely pow-
dered fluor-spar, with six to el^ht parts of concoatrated sulphuric acid, In a small stoneware
bottle or a glass flask, provided with a wide del Ivory- tube, dipping into a little mercury In a
pmall porcelain capsule, which is set in a largre beaker containing six or eight parts of water.
The stoneware bottle or flask is set In a small sand-bath, with the sand piled about It, as high as
the material, and srentlj heated from a lamp. Each bubble of gas decomposes with deposition
of grelatinous silicic acid. When the water is filled with this deposit, it may be separated by
straining through cloth and again treating with the gas for grreater concentration. The strained
liquid la finally filtered and preserved for use.
290 SIUCOXSILICOy DIOXIDE. §S48.
§248. Silicon. Si = 28.4 . Valence foiir.
There are three modifications of silicon: (a) Amorphous. — A dark brown
powder; specific fjravUy, 2.0; non-volatile: infusible; bums in the air, forming
SiO, , a nd i n chlorine, forming: SiCl4 . It is not attacked by acids except HP:
81 + 6Hr = HjSir, 4- 2Hj . It is dissolved by XOH^with evolution of
hydrogen. (6) Graphitoidai, — May be fused, but is not oxidized upon ignition
in air or in oxygen. It is not attacked by HF , but is dissolved by a mixture
of HF and HNO, , forming H^SlFo • It is attacked slowly by fused XOH.
(c) Adama4itine silicon, crystalline silicon. — Grayish-black, lustrous, octahedral
crystals, formed by fusing the graphitoidnl form. Specijfic gravity, 2.49 at 10*
(Woehler, A., 1856, 97, 261). It scratches glass but not topaz. It melts between
the melting points of pig iron and steel, 1100° to 1300°. In chemical properties
it is very similar to the graphitoidal form, being attacked \%ith even greater
difficulty. Silicon is never found free in nature, but always in combination as
silica, SiOz , or as silicates.
Amorphous silicon is formed by passing vapor of SiCl, over heated potassium:
by heating magnesium in SiF^ vapor; bj' heating a mixture of Mg and 810,: by
electrolysis of a fused silicate. It is readily prepared by heating a mixture of
magnesium, one part, with sand, four parts, in a wide test-tube of hard glass
(Gattermann, B., 1889, 22, 186). The graphitoidal form is crystalline and by
many is said to be the same as the adamantine form. Method of preparation
essentially the same (Warren, C. X., 1891, 63, 46). The crystalline form is made
by fusing a silicate or K38iF« with Al: by passing vapors of 8iCl4 over heated
Na or'Al in a carbon crucible (Deville, .4. Cft., 1857, (3), 49, 62; Deville and
Caron, A. Ch., 1863, (3), 67, 435; Woehler, /. c).
§249. Silicon dioxide. Si02 = 60.4.
(Silicic anhydride; silica.)
Silicic acid. HoSiOj = T8.416 .
II
Si^O-% and H'^Si^O""., , = Si = and H — — Si — — H.
1. Properties.— Silica, silicic anhydride, Si02 , is a white, stable, infusible solid:
insoluble in water or acids; soluble in lixed alkalis with formation of silicates.
specific yruriiy of quartz, 2.647 to 2.652; of amorphous silica, 2.20 at 15.6°.
Silicic acid, siliam hydroj-idc, HzSiO, , is a white, gelatinous solid, generally
insoluble in water, and soluble in mineral acids. A dilute solution in water ii^
obtained by dialysis of the fixed alkali silicate with an excess of HCl until
the chlorides are all removed. It may be boiled for some time before the acid
precipitates out. I'pon standing silitic acid soon separates.
2. Occurrence. — Silicon is never found free in nature; it is always combinetl
with oxygen in the form of silicon dioxide, SiOa , as quartz, opal, flint, sand.
etc.; or the silicon dioxide is in combination with bases as silicates: asbestos,
soapstone, mica, cement, glass, etc. All geological formations except chalk
contain silicon as the dioxide or as a silicate.
3. Formation. — (.'rystalliue silica is formed by ])assing silicon fluoride into
water, forming silicic acid and fluosilieie acid: USi'E\ 4- iJH.O = H^SiOj +
2H3SiF« . The precipitate of silicic acid is dissolved in boiling' NaOH and then
heated in sealed tubes. Below lso° erystals of tridymite are formed. a»<^
above 1S0° crystals of quartz (Maschke, /V/.^/., 1S72, 145, 549).
4. Preparation. — IMire amorphous silica is prepared by fusing finely j^ow-
dered quartz with six parts of sodium carbonate, dissolving the cooled mass in
water, and pouring into fairly concentrated hydrochloric acid. The precipitate
is filtered, well washed and ignited. Or SiF^ vapors are passed into water
(§246) and the gelatinous precipitate washed, dried and ignited. Crystalline
§249,7. iSlLlCOy DIOXIDE, 291
silica is prepared by fusing silicates with microcosmic salt or with borax
(Rose, ./. pr., 1867, 101. 228).
Silicic acid. — The various hydroxides of silica act as weak acids. Metasilicic
acid, HjSiO, , has been isolated; it is formed by decomposing silicon ethoxide,
Si(0C,Hs)4 , with moist air (Ebelmen, J. pr., 1846, 37, 359). Also by dialysis of
a mixture of sodium silicate with an excess of hydrochloric acid until the
chlorides are all removed, concentrating, allowing to gelatinize, and drying
over sulphuric acid. Other hydroxides, acids, have been isolated, but there is
some question as to their exact composition.
5. Solubilities.— Silica, SiOj , is insoluble in water or acids except HF ,
which dissolves it with formation of gaseous silicon fluoride, SiF^ (§246).
Of the silicates only those of the fixed alkalis are soluble in water, water
glass. These silicates in solution are readily decomposed by acids, in-
cluding carbonic acid, forming silicic acid, gelatinous. While anhydrous
silicic anhydride, SiOj , is insoluble in mineral acids, the freshly precipi-
tated hydroxide, silicic acid, is soluble in those acids. Silicic acid is
decomposed by evaporation to dryness in presence* of mineral acids, with
reparation of the anhydrous SiOj ; which is insoluble in more of the same
acids, which previously had effected its solution.
The most of the silicates found in nature are of complex composition.
They are combinations of SiOg with bases. They are, as a rule, insoluble
in water or acids.
6. Beactions. — Solutions of the alkali silicates precipitate solutions of
ill other metallic salts with formation of insoluble silicates; they are
iecomposed by acids with separation of silicic acid, a gelatinous precipi-
tate, soluble in hydrochloric acid. Evaporation decomposes silicic acid
svith separ«^lion of insoluble silicic anhydride, SiOg . Ammonium salts
precipitate gelatinous silicic acid from solutions of potassium or sodium
silicate. Therefore in the process of analysis the silicic acid, not removed
In the first group by hydrochloric acid, will be precipitated in the third
^oup on the addition of ammonium chloride.
Silica, SiOa , is soluble in hot fixed alkalis forming silicates ; it is not
^oluble in ammonium hydroxide, nor are solutions of alkali silicates pre-
?ipitated on addition of ammonium hydroxide as they are on the addition
Df ammonium salts. Boiling SiOj with the fixed alkali carbonates forms
soluble silicates with greater or less readiness. Nearly all silicates are
Iecomposed by heating in sealed tubes to 200° with concentrated HCl or
7. Igpiition. — Silicates fused with the alkalis form soluble alkali sili-
:ates, and oxides of the metal previously in combination. If alkali car-
Donates are employed the same products are formed with evolution of
DOj . Preferably a mixture (in molecular proportions) of potassium and
iodium carbonates, four parts, should be used to one part of the insoluble
iilicate. Silica, SiOg , is also changed to a soluble silicate by fusing with
fixed alkali hydroxides or carbonates.
292 PHOSPHORUS. §248,8.
SiO, does not react with X3SO4 or Na^SOt , even when fused at a very high
temperature (Mills and Meanwell, J. (\, 1881, 39. 533). In the fused bead of
microcosmic salt particles of silica swim unidissolced. If a silicate be taken,
its base will, in most cases, be dissolved out, leavings a ^^ skeleton of silica*^ un-
dissolved in the liquid bead. But with a bead of sodium carbonate, silica (and
most silicates) fuse to a clear glass of silicate.
Silica is separated from the fixed alkalis in natural silicates, by mixing the
latter in fine powder with three parts of precipitated calcium carbonate, and
one-half part of ammonium chloride, and heating in a platinum crucible to
redness for half an hour, avoiding too high a heat. On digesting in hot water,
the solution contains all the alkali metals, as chlorides, with calcium chloride
and hydroxide.
8. Detection. — SiUcafes are detected by conversion into the anhydride.
SiOg . The silicate is fused with about four parts of a mixture of potas-
sium and sodium carbonates, digested with warm water, filtered, and
evaporated to dryness with an excess of hydrochloric acid. The dry resi-
due is moistened with concentrated HCl and thoroughly pulverized; water
is added and the precipitate of SiOo is thoroughly washed. Further con-
firmation may be obtained by warming the precipitate of SiOj vdth
calcium fluoride and sulphuric acid (in lead or platinum dishes), forming
the gaseous silicon fluoride, SiF^ . This is passed into water where it is
decomposed into gelatinous silicic acid and fluosilicic acid: SSiF^ + ^^®
= HoSiOa + 2HoSiF« (§246). Silica, SiOo , is usually treated as directed
for silicates, but may be at once warmed with calcium fluoride and sul-
phuric acid.
9. Estimation. — The compound containing a silicate or silica is fused with
fixed alkali carbonates as directed under detection, and the amount of well-
washed SiOj determined by weighing after ignition.
§250. Phosphorus. P = 31.0. Usual valence three or five.
1. Properties. — Phosphorus is prepared in several allotropic modifications.
Specific yrurUy of the yellow, solid, at 20°, 1.82321; liquid, at 40°, 1.741)24: solid,
at 44°, 1.80681 (Pisati and de Franehis, /i., 1875, 8, 70). At ordinary tempera-
tures it is brittle and easily pulverized. At about 45° it melts, biit may l>e
cooled in some instances (under an alkaline liquid) as low as -f-4° without
solidifying. When it solidifies from these lower temperatures, as it does l>y
stirring with a .solid substance, the temperature immediately rises to about 45".
Boiling pmnt, 287.3° at 762 mm. pressure (Schroetter, A., 184*8, 68, 247; Kopp, •!••
1855, 93, 120). The density of the vapor at 1040° is 4.50 (Deville and Troost,
C, r., 1863, 56, 801). The computed density for the molecule P4 is 4.294. At a
white heat the density, 3.632. indicates dissociation of the molecule to ?:
(Meyer and Biltz, /?., *^1880, 22, 725). Specific gravity of the red amorphous
modification at 10°, 1.064.
Ordinary crystalline yellow stick phosphorus is a nearly colorless, trans-
parent solid: when cooled slowly it is nearly as clear as water. In water con-
taining air it becomes coated with a thin whitish film. If melted in fairl.^'
large quantities and cooled slowly it forms dodecahedral and octahedral crys-
tals (Whewell, C\ xV., 1870, 39, 144). Heated in absence of air above the boiling
point it sublimes as a colorless gas, depositing lustrous transparent crystals
(Blondlot, C. r., 1866, 63, 397). At low temperatures phosphorus oxidizes slo^'X
in the air v.ith a characteristic odor, jjrobably due to the formation of ozone
^^260, 4. PHOSPHORUS, 293
and phosphorous oxicte, PjO, (Thorpe and Tutton, J. C, 1890, 57, 573). It ignites
spontaneously in the air at 60°, burning with a bright yellowish white light
producing much heat. From the finely divided state, as from the evaporation
of its solution in carbon disulphide, it ignites spontaneously at temperatures
at which the compact phosphorus may be kept for days. It must be preserved
under water. Great precaution should be taken in working with the ordinary
or yellow jjhosphorus. Burns caused by it are very painful and heal with
great diificulty. Ordinary phosphorus is luminous in the dark, but it has
been shown that the presence of at least small amounts of oxygen are neces-
sary. The presence of H^S , SO, , CSo , Br , CI , etc., prevent the glowing
(Schroetter, J. pr,, 1853, 58, 158; Thorpe, Nature, 1890, 41, 523). Upon heating
in absence of air, better in sealed tubes, to 300° it is changed to the red modi-
fication (Meyer, B^ 1882, 15, 297).
Red phosphorus is a dull carmine-red tasteless powder. It is not poisonous,
while the ordinary yellow variety is intensely poisonous, 200 to 500 milligrams
being suflicient to cause death. While the yellow modification is so readily
and dangerously combustible when exposed to the air even at ordinary tem-
peratures, the red variety needs no special precautions for its preservation.
It does not melt when heated to redness in sealed tubes, but is partially
changed to the yellow crystalline form (Hittorf,. Po^flf., 1865, 126, 193). If
amorphous phosphorus be distilled in the absence of air, it is changed to the
crystalline form, action beginning at 200°. Heated in the air from 250° to 260°
it takes fire (Schroetter, l.c). A black crystalline metallic variety of phos-
phorus is described by Hittorf (/. c); also Remsen and Kaiser (Aw., 1882, 4, 459)
describe a light plastic modification. Phosphorus is largely used in match-
making. Yellow phosphorus is used in the ordinary match, and the red
.(amorphous) in the safety matches, the phosphorus being on a separate surface.
2. Occurrence. — It is never found free in nature. It is found in the primitive
rocks as calcium phosphate, occasionally as aluminum, iron, or lead phosphate,
-etc. Plants extract it from the soil, and animals from the plants. Hence traces
of it are found in nearly all animal and vegetable tissues; more abundantly
in the seeds of plants and in the bones of animals.
3. Formation. — Ordinary phosphorus is formed by heating calcium or lead
phosphates with charcoal. The yield is increased by mixing the charcoal with
sand or by passing HCl gas over the heated mixture. By igniting an alkali
•or alkaline earth pbosphate with aluminum (Rossel and Frank, /?., 1894, 27, 52).
Red phosphorus is formed by the action of light, heat or electricity on ordinary
phosphorus (Meyer, /?., 1882, 15, 297). By heating ordinary phosphorus with
a small amount of iodine (Brodie, J. pr., 1853, 58, 171).
4. Preparation.— Ordinary phosphorus is prepared from bones. They are
first burned, which leaves a residue, consisting chiefiy of Ca8(P04)2; then
H2SO4 is added, producing soluble calcium tetrahydrogen diphosphate (a).
After filtering from the insoluble calcium sulphate the solution is evaporated
and ignited, leaving calcium metaphosphate (h). Then fused with charcoal,
reducing two-thirds of the phosphorus to the free state (c). The mixture of
sand, SiOj , with the charcoal is preferred, in which case the whole of the
phosphorus is reduced (d). Hydrochloric acid passed over red-hot calcium
phosphate and charcoal reduces the whole of the phosphorus. This process
works well in the laboratory, and has also been successfully employed on a
larger scale. Either of the calcium phosphates may be used (c) and (f).
(a) Ca,(P002 + 2H,S0, = 2CaS0, + CaH,(PO,),
(6) CaH4(P0J, 4- ignition = Ca(P03)2 + 2H,0
(c) 3Ca(PO,)2 + IOC = Ca,(P0,)2 + lOCO -|- P,
(d) 2Csi(V0,), -f IOC 4- 2SiO, = 2CaSiO, -h P* + lOCO
(c) 2Ca3(PO,)2 + 16C -f 12HC1 = eCaCl, + P* + 16C0 + eH,
(0 2Ca(P0»), -f 12c -f 4HC1 = 2CaCL + P4 + 12C0 + 2H2
Red or amorphous phosphorus is prepared by heating ordinary phosphorus
for a long time (40 h-ours) at 240° to 250° in absence of air. At 260° the reverse
change takes place. If the heating is under pressure and at 300°, the change
to the red phosphorus is almost immediate. It is washed with CSj to remove
all traces of yellow phosphorus and is dried at 100°.
294 PB08PB0R18. §850, 5.
5. Solnbilities. — A trace of phosphorus dissolves in water. Alcohol
dissolves 0.4, ether 0.9, olive oil 1.0, and turpentine 2.5 per cent of it,
while carbon disulphide dissolves 10 to 15 times its own weight. Bed
phosphorus is insoluble in water, ether, or carbon disulphide.
6. Beaotions. — When phosphorus is boiled with a fixed alkali or alkaline
earth hydroxide, phosphorus hydride, phosphine (§249), PH3, and a
hypophosphite (§260) are formed. Phosphorus, when warmed in an
atmosphere of N or CO2 , combines directly with many metals to form
phosphides. These phosphides are usually brittle solids decomposing
with water or dilute acids with formation of phosphoretted hydrogen,
PH3 . In nearly all the reactions of phosphorus both varieties react the
same, the red variety with much less intensity, and frequently requiring
the aid of heaj;. It is ignited when brought in contact with PbOg , PbjO^ .
HgO , AgjO , CrOs , KjCr^O^ and when heated with CuO or MnOj . Solu-
tions of platinum, gold, silver, and copper salts are decomposed by phos-
phorus with separation of the corresponding metal (Bocttger, J, C, 1874,
27, 1060).
With HNO3, HsPO^ and NO are formed; when heated with KNO, a
rapid oxidation takes place.
It combines with oxygen, forming PoOg or PjOg . With yellow phos-
phorus the reaction begins at ordinary temperature; with the red variety
not till heated to 250° to 260° (Baker, J. C, 1885, 47, 349).
Water is decomposed at 250°, forming PH3 and H3PO4 (Schroetter, I c).
Combination with red phosphorus and sulphur takes place at ordinary
temperatures, forming P2S3 or PoS,^ , depending upon the proportion of
each employed (Kekule, .4., 1854, 90, 310). With ordinary phosphorus
the action is explosive.
CI or Br react with incandescence at ordinary temperatures, forming
trihalogen or pentahalogen compounds, depending upon the amount of
halogen employed. With iodine, PI3 is formed.
The halogen compounds of phosphorus are decomposed by water with
formation of the corresponding hydraoids and phosphorous or phosphoric
acids, depending upon the degree of oxidation of the phosphorus. In
the presence of water phosphorus is oxidized to H3PO4 by CI, Br, I,
HCIO3, HBr03, or HIO^ with formation of the corresponding hydracid:
P, + lOCl. + 16H,0 = 4H3PO, + 20HC1 .
7. Ignition. — When sodium carbonate is heated to redness with phosphorus,
the carbonic anhydride is reduced and carbon is set free. Phosphorus heated
with magnesium in a vapor of carbon dioxide forms TJtt.gt , which can be
heated to redness in absence of air without decomposition. Heated in the air
it becomes oxidized (Blunt, A. C/i., 1805, (4), 5, 487). Phosphorus also combines
with Cu , Ag , Cd , Zn and Sn when it is heated with these elements in sealed
tubes. It does not combine with Al and but slightly with Fe (Emmerling.
J. C, 1879, 36, 508).
§252,3. PHOiiPHryE-HYPOPHOSPHOROUS ACID. 295
8. Detection. — By its phosphoreBcence; by formation of PH3 when
boiled with KOH (Hofmann, B., 1871, 4, 200); by oxidation to H3PO4 and
detection as such (§76, 6d),
9. Estimation. — Oxidation to HaPO^ , precipitation with magnesia mixture as
MgNH^PO^ , ignition to, and weighing as Hg^T^Oy (§189, 9),
§261. Phosphine. PH3 = 34.024.
P^"H'3,H — P — H.
I
H
Phoapliiney PH, , is a colorless gas having a very disagreeable odor. As
usually prepared, it is spontaneously inflammable, burning in the air with
formation of metaphosphoric acid: 2PHs + 40, = 2HPOs + 2H,0 . It is
liquiiied and frozen at very low temperatures; boUiny point, about — 85°;
melting point, — 132.5" (Olszewski, if., 1886, 7, 371). It is very poisonous, spar-
ingly soluble in water, which solution has the peculiar odor of the gas and has
an exceedingly bitter taste. It is formed by boiling phosphorus with a fixed
alkali or alkaline earth hydroxide (a); by ignition of H,PO, or HsPG. (6); by
ignition of hypophosphites (r); by the decomposition of the alkaline earth
phosphides with water or dilute acids ((/) :
(a) P, + 3K0H -f 3H3O = 3KH,P0, -|- PH.
(6) 2H,P0, = HPOa + PH. + H,0
4H,P0, = 3HP0, + PH, + 3H,0
(c) 4NaH3PO, = Na^PjO, + 2PH, + H,0
(<f) Ca,P, + 6H2O = 3Ca(0H), + 2PH,
CajPa + 6HC1 = 3CaCla + 2PH,
It is a strong reducing agent; transposes many metallic solutions: 3CuS0« -f-
2PH, = Cu.Pa 4- SHjSO^; reduces solutions of silver and gold to the metallic
state: 8AgN0, + PH, -f 4H2O = HsPG^ + 8HN0, -|- 8Ag; is oxidized to H^PO*
by hot HxSG^ , CI , HCIO , HNO, , HNG, , H^AsO^ , etc. A liquid phosphorus
hydride, P^H^ , and a solid. P4HJ . are known (Besson, C. r., 1890, 111, 972;
Gattermann and Hausknecht, B., 1890, 23, 1174).
§262. Hypophosphorous acid. H3PO2 = 66.024 ,
H
I
H'aP'O-^. H — — P=:0.
1. Properties. — Hyx)ophosphorous acid was discovered in 1816 by Dulong (A. Ch.^
1816, 2, 141). It is a colorless syrupy liquid; specific gravity, 1.493 at 18.8". At
17.4® it becomes a white crystalline solid (Thomsen, B., 1874, 7, 994). Although
containing three hydrogen atoms it forms but one series of salts, e. g,, NaHsPO,,
Ba(H,POa), , etc.
2. Occurrence. — Not found in nature.
3. Ponnatioii. — All orJinary metols form hypophosphites except tin, copper
and mercurosum. Silver and ferric hypophosphites are very unstable. (1) A
296 HYP0PH0SPH0R018 ACID. §282,4.
few metals, such as zinc and iron, dissolve in HsPO, , giving off hydrogen and
forming a hypophosphite. (2) The alkali and alkaline earth salts may be
formed by boiling phosphorus with the hydroxides (Mawrow and Muthmann,
Z. angetc,, 1896, ii, 208). (3) As all hypophosphites are soluble, none can be
formed by precipitation. All may be formed from their sulphates by trans-
position with barium hypophosphite. (4) All may be made by adding HtPOt
to the carbonates or hydroxides of the metals.
4. Preparation. — To prepare pure HsPO, , fiaO and P (in small pieces) are
warmed in an open dish with water until PH, ceases to be evolved. The
liquid is filtered and excess of BaO is removed by passing in CO,. » After again
filtering, the liquid is evaporated to crystallization of the barium salt. This
is dissolved in water and decomposed by the calculated quantity of HjSOt .
The solution is filtered and evaporated in an open dish, care being taken not to
heat above 110** . Upon cooling the white crystalline tablets are obtained.
5. Solubilities. — The free acid is readily miscible in water in all proportions.
The salts are all soluble in water, a number of them are soluble in alcohol.
6. Reactions. — A. — With metals and their compounds. Hypophosphorous
acid is a very powerful reducing agent, being oxidized to phosphoric acid or a
phosphate.
i. Pbiv becomes Pb" in acid or alkaline mixture.
2. Ag' becomes Ag** in acid or alkaline mixture.
3. Hg^' becomes Hg^ and then Hg° in acid or alkaline mixture.
4. Asv and As'" become As° in presence of HCl .
5. Bi'" becomes Bi° in presence of alkalis or acetic acid.
6. Cu" becomes CUjH, and on boiling Cu** (separation from Cd)»
7. Fe'" becomes Fe" , no action in alkaline mixture.
8. Crvi becomes Cr'" , no action in alkaline mixture.
5. Co'" becomes Co" , no action in alkaline mixture.
10. Ni'" becomes Ni" , no action in alkaline mixture,
11. Mn"+n becomes Mn" in acid solution.
12. Mniv+n becomes Mniv in alkaline mixture.
B. — With non-metals and their compounds.
1. H,Fe(CN)e becomes H«Fe(CN)e .
2. HNO, and HNO2 become NO .
3. HaPOz on heating becomes HjPO^ and PHj .
4» H^SOa becomes free sulphur with formation of some H^S (Ponndorf, J. C.^
1877,31, 275).
H2SO4 becomes first HjSOs then S . See above.
5. CI becomes HCl in acid mixture, a chloride with alkalis.
HCIO and HClOg form same products as CI .
6. Br becomes HBr in acid mixture, a bromide with alkalis.
HBrO, forms HBr .
7. I forms HI , in alkaline mixtures an iodide.
HI , dry, reacts violently, forming H3PO3 and PH4I (Ponndorf, I. c).
HIO, forms HI .
7. Ignition.— On ignition hypophosphites leave pyrophosphates, evolving P^»-
The acid decomposes on heating to PH» and H.PO^ (or HPO, if at a red heaij-
8. Detection, — Hypopliosphorous acid may be known from phosphorous^
acid by adding cupric sulphate to the free acid and heating the solution
to 55°. With hypophosphorous acid a reddish-black precipitate of eopp^-^r
hydride (CUgHo) is throwTi down, which, when heated in the liquid to lOO •
is decomposed with the deposition of the metal and the evolution of
hydrogen, whilst with phosphorous acid the metal is precipitated a^^
hydrogen evolved, but no CU2H2 is formed. Further, hypophosphorous
acid reduces the permanganates immediately, but phosphorous acid onl}
after some time. Phosphites precipitate barium, strontium, and calciuni
§253,6. PHOSPHOROUS ACID. 297
salts, while hypophosphites do not. When h3rpophosphorou8 acid is
treated with zinc and sulphuric acid it is converted into phosphoretted
hydrogen. On boiling h3'pophosphorous acid with excess of alkali hydrox-
ide, first a phosphite then a phosphate is formed, with evolution of
hydrogen.
9. Estimation. — (1) By oxidation with nitric acid and then proceeding as
with phosphoric acid. (2) By mercuric chloride acidulated with HCl; the
temperature must not rise above 60**, otherwise metallic mercury will be
formed. The precipitated Hg€l , after washing and drying at 100**, is weighed.
NaHjPOa + 4HgCl, + 2H,0 = 4HgCl + H,PO, + NaCl + 3HC1
§263. Phosphorous acid. H3PO3 = 82.024 .
H
I
H'3F"0""3, H — — P — — H.
1. Properties. — Phosphorous anhydride, P^O, , is a snow-white solid, melting
at 22.5°, and boiling at 173.1** (Thorpe and Tutton, J. C, 1890, 57, 545). The
vapor density of the gaseous oxide indicates the molecule to be P40« . Spe^'ific
(f rarity of the liquid at 21°, 1.9431; of the solid at the same temperature, 2.135.
it has an odor resembling that of phosphorus. Heated in a sealed tube at
200° it decomposes into P^O^ and P (T. and T., J, C, 1891, 59, 1019). It reacts
slowly with cold water, forming HsPO,; with hot water the reaction is violent
and PH, is evolved. Upon exposure to the air it oxidizes to PjOa .
The acid, H,POs , is a crystalline solid, very deliquescent, melting at 74°
(Hurtzig and Geuther, A., 1859, 111, 171). It is a dibasic acid, forming no
tribasic salts (Amat, C. r., 1889, 108, 403). One or two of the hydrogen atoms
are replaceable by metals forming acid or normal salts. The third hj'drogen
is never replaced by a metal, but may be replaced by organic radicles (Railton,
J. C, 1855, 7, 216; Michaelis, J, C, 1875, 28, 1160). Neither meta nor pyro-
phosphorous acids are known, but a number of pvrophosphites have been pre-
pared (Amat, C. r., 1888, 106, 1400; 1889, 108, 1056; 1890, 110, 1191 and 901;
A. Ch., 1891, (6), 24, 289).
2. Occurrence. — Does not occur in nature.
3. Formation. — PjOj is formed together with PaO^ when phosphorus is
ignited in the air. HjPO, is formed together with H,P04 when phosphorus
is oxidized with HNO,; by the oxidation of HaPO,: by the action of P upon a
concentrated solution of GuSO^ in absence of air: 3CUSO4 -f P4 + 6H2O =
Cu,P, -h 2HaP0, 4- 3H2SO4 (Schiff, A., 1860, 114, 200).
4. Preparation To prepare phosphorous anhydride, P2O., , phosphorus is
burned in a tube with an insufficient supply of air (Thorpe and Tutton, I.e.).
The acid, HgPO, , is prepared by dissolving the anhydride in cold water; by
decomposing PCI, with water (Hurtzig and Geuther, I.e.),
5. Soliibilities. — The acid is miscible in water in all proportions. Alkali
phosphites are soluble in water, most others are insoluble (distinction from
hypophosphites) .
6. Beactions. — Phosphorous acid is a strong reducing agent, oxidizing to
phosphoric acid when exposed to the air. It reduces salts of silver and gold to
the metallic state and is changed to phosphoric acid by most of the strong
oxidizing acids and by many of the higher metallic oxides. HgCl, becomes
HgCl and then Hg° , CuCls becomes CuCl then Cu° (Rammelsberg, J. C, 1873,
298 HTP0PH08PH0RIC ACID— PHOSPHORIC ACW. §288,7.
526, 13). Concentrated H.80, with heat forms H,PO, and 80, (Adie« J. C 1.S91.
59. 230). H^SO, forms H,8 and H,PO. (Woehler. .4.. 1841, 39, 252). Naaceot
hydrogen (Zn and HjSO,) produce PH, (Dusart, C. r., 1856, 43 , 1126 ).
7. Ignition. — The acid is decomposed bv ignition, forming HFO, and P or
FH, (Vigier. i?/., 1**69, (2), 11. 125; Hurtzig and Geuther, /. r.). Phosphites are
deco mposecl by heat. leaving a pyrophosphate and a phosphide and evolring
FH, or H (Rammelslierg, B., 1876! 9. 1577: and Krant, A,. 1S75, 177, 274).
8. I>et6ction. — By oxidation to H,PO« and detection as such. It is distin-
guished from hypophosphorous acid by reducing GnSO. to Ca^, while the
latter forms CasH,: also by the solubilities of the salts (S2i52, 8). Its reactions
with oxidizing agents distinguish it with hypophosphorous acid from phos-
phoric acid.
9. Estimation. — By oxidation to H,PO« and estimation as such.
§264. Hypophosphoric acid. H^PjO^ = 162.032 .
II II
H',P^20-"e,H — — P — P — — H.
I I
I I
H H
Hypophosphoric acid is formed together with phosphorous and phosphoric
acids by slowly oxidizing phosphorus in moist air (Salzer, A., 1885, 832. 114
and 271): also by oxidizing phosphorus with dilute HNO, in presence of silver
nitrate (Philipp! H., 1SS5, 18, 749). It consists of small colorless hygroscopic
crv'stals which melt at 55°. It decomposes when heated to 70° into JB[xPOs and
HPO, , and at 120° gives H^P^O^ and PH, (Joly. C. r., 1S86. 102, 110 and 7C0).
It is oxidized to HaPO* by warm HNO., , slowly by KKn04 in the cold, rapidly
when heated. It is not oxidized by H..O2 , chlorine water or H^CrO*: HgCl.
becomes Hg€l (Amat, (\ r., 1890, 11*1, 670). It is not reduced by Zn and HjSO»
(distinction from HjPO, and HaPOj). With a solution of silver nitrate it gives
a white precipitate which does not blacken in the light (distinction from H,POi
and HjPOj). It forms four series of salts, all four hydrogen atoms l)einp
replaceable by a metal. The hypophosphates are much more stable towards
oxidizing agents than hypophosphites or phosphites.
§265. Phosphoric acid. H^PO^ = 98.024 .
II
H'^pvo-", ,H — — P — — H.
I
I
H
1. Properties. — Phosphoric anhydride, PjOj *, is a white, flakey, very delique-
scent solid, fusible, subliming undecomposed at a red heat. It is very soluble
in water, forming three varieties of phosphoric acid: ortho, HSPO4; meta, HPO,;
•According to Tllden and Harnett fJ. C, 1896, 69. 154) the molecule is ^^Ox^ not P,0,; P4O1
not PjO, t Thorpe and Tutton, J. (\ 1801, 59, 1023) : and P**. Dot P,8» (Isambert, C.r„ 1898. !•«.
1386).
§265, 3. PHOSPHORIC ACID, 299
and ptfro, "RJ^tOf . Orthophosphoric acid is a translucent, feebly crystallizable
and very deliquescent soft solid, i^pecific gravity, 1.88 (SchiflP, A., 1860, 113, IS.*]);
inelting point, 41.75*' (Berthelot, Bl,, 1878, (2), 29, 3). It is changed by heat,
first to pyrophosphoric acid, then to metaphosphoric acid. Orthophosphoric
acid forms three classes of salts: M'HjPO^ , primary, monobasic or mono-
metallic phosphates; M'2HP04 , secondary, dibasic or dimetallic phosphates;
and M'aPO^ , tertiary, tri basic, trimetallic or normal phosphates. The first
two are acid salts, but Na...HjP04 is alkaline to test paper. Metaphosphoric
acid, HPO3 ,H — O — P = 0,isa white waxy solid, volatile at a red heat
II
O
(ordinary glacial phosphoric acid owes its hardness to the universal presence of
sodium metaphosphate). It is a monobasic acid, but there are various poly-
meric modifications, distinguished from each other chiefly by physical diifeV-
ences of the acids and their salts (Tammann, Z. phys, Ch,, 1890, 6, 122).
O O
II II
Pyrophosphoric acid, H4P,0t , H— O — P — 0~P — O — H,isa glass-like
I I
O O
I I
H H
solid (Peligot, A. Ch., 1840, (2), 73, 28G), very soluble in, but unchanged by,
water at ordinary temperature; changed by boiling water to HgPO^ . Heated
to redness HPO, is formed. It forms two classes of salts: M%H2P20t and
M'.P^O, .
2. Occurrence. — Phosphates of Al , Ca , Mg and Pb are widely distributed in
minerals. Guano consists quite largely of calcium phosphate. Calcium and
magnesium phosphates are found in the bones of animals and in the ashes of
plants. The free acids are not found in nature.
3. Pormation. — Phosphoric anhydride, P^jOj , is formed by burning phosphorus
in great excess of air: also by burning phosphorus in NO , NO, , or CIO2 .
Orthophosphoric aeid, HsP04 , is formed by long exposure of phosphorus to
moist air, or by oxidation with HNO3; by oxidation of HjPOz or HgPO, with
the halogens, HNO, , HCIO, , etc.: by treating P2O5 , HPO, , or H.PaOT with
boiling water: by combustion of PH, in moist air; and by action of water on
PCI5 . It is also formed from metallic phosphates by transposition with acids
in cases where a precipitate' results, as a lead or barium phosphate with sul-
phuric acid, or silver phosphate with hydrochloric acid. But when the pro-
ducts are all soluble, as calcium phosphate with acetic acid or sodium phosphate
with sulphuric acid, the transposition is only partial; so that unmixed phos-
phoric acid is not obtained. A non-volatile acid, like phosphoric, is not sepa-
rated from liquid mixtures, as the volatile acids are, like hydrochloric. The
change represented by equation (a) can be verified, that is, pure phosphoric
acid can be separated; but the changes shown in equations (h) and (c) do not
comprise the whole of the material taken. In the operation (h) some sodium
phosphate and some nitric acid will be left, and in (r) some trihydrogen
phosphate will no doubt be made.
a. CbMJTO,), -f H.aO^ = CaCO^ -f 2H3PO4
6. Na^HPO. -h 2HN0, = 2NaN0, -|- H.PO^
and Na,HP04 -f HNO, = NaNO, -f NaH2P04
c. 2CaSP04 -f 2HC1 = CaCl^ + CaH^CPO^),
Meiaphosphnric acid is formed by treating P2O5 with cold water; by decom-
position of lead metaphosphate with H2S or of the barium salt with H2SO4;
by ignition to dull redness of phosphorus or any of its acids in the presence
of air and moisture.
Pyrophosphoric acid, 'B.^'P..O^ . is formed by the decomposition of lead pyro-
phosphate, PbaPjOy , with H5S or of the corresponding barium salt with
H3SO4: or by heating HgPO* to a little above 200** until no yellow silver
phosphate, Ag^PO, , is obtained on dissolv'ner in water and treatment with
silver nitrate after neutralization with NH4OH .
300 PHOSPHORIC ACID. §248,4.
4. Preparation. — To prepare PgO- , phosphorus is burned in a slow cur-
rent of dry oxygen heating to about 300°, then in a more rapid current
of the gas, and finally the PoOj is distilled in an atmosphere of oxygen
(Shenstone, Watts' Die, 1894, IV, 141). HgPO^ is prepared by warming
phosphorus, one part, with nitric acid, sp, gr. 1.20, ten to twelve parts,
with addition of 300 to 600 milligrams of iodine to 100 grams of phos-
phorus, until the phosphorus is completely dissolved. The excess of
HNO3 is removed by evaporation, water is added and the solution is sat-
urated with HjS to remove any arsenic that may be present. The solution
is then evaporated to a syrupy consistency at temperatures not above
150° (Krauthausen, Arch, Pharm,, 1877, 210, 410; Huskisson, B,, 1884,
17, 161). Many orthophosphates are formed by the action of H3PO4 upon
metallic oxides or carbonates; by the reaction between an alkali phosphate
and a soluble salt of the heavy metal ; by fusion of a metaphosphate with
the corresponding metallic oxide or hydroxide; also by long continued
boiling of meta or pyrophosphates. Metaphosphates are formed by double
decomposition with NaPO» or by fusion of a monobasic phosphate or any
phosphate having but one hydrogen equivalent substituted for a metal,
the oxide of which is non-volatile, e, g., NaNH^HPO^ . Pyrophosphates
are formed by double decomposition with Na^PjOr ; by action of H4P2O;
on certain oxides or hydroxides; also by ignition of dibasic orthophos-
phates, e. g., NaoHPO^ . NaoHoPoO^ may be prepared by titrating a sat-
urated solution of Na^PaOy w^ith HNO3 ^intil the solution gives a red color
with methyl orange. Upon standing the salt separates in large crystals
(Knorre, Z. angew., 1892, ()39).
5. Solnbilities. — All the phosphoric acids are readily soluble in water,
as are all alkali phosphates. Alkali primary orthophosphates have an
acid reaction in their solutions; alkali secondary and tertiary phosphates
are alkaline in their solutions; the latter is easily decomposed, oven bv
CO2, forming the secondary salt. A number of non-alkali primary ortho-
phosphates are soluble in water, e. g., (jdJl^C^O^)^ . All normal and <li-
metallic orthophosphates are insoluble except those of the alkalis. The
normal and dimetallic phosphates of the alkalis precipitate solutions of
all other salts. The precipitate is a normal, dimetallic, or basic phos-
phate, except that with the chlorides of mercury and antimony it is not
a phosphate but an oxide or an oxychloride.
All phosphates are dissolved or transposed by HNO;, , HCl , or H2SO4 »
and all are dissolved by HCgH^Og except those of Pb , Al and Fe'" . All
are soluble in H3PO4 except those of lead, tin, mercury, and bismuth.
The non-alkali meta and pyrophosphates are generally insoluble in
water. The pyrophosphates of the alkaline earth metals are difficultly soln-
ble in acetic acid. The most of the pyrophosphates of the heavy metals,
€l1. phosphoric acid. oOl
except sDTtT. ^a^ ^oiiable in ^solution? of alkali pvrophosphatOsi, a? ij;^?.?.*:'
f^pkcmpkaU;* minAie in vitter (distinction from orthophosphatos). Forr:o
iron as a double fTn>ph*:»sphaie l-.^sos the ohaniotoristii- proportion of iha:
metal (Peisox. J. C^ i>49. 1, 1S3). Phosphates are insoluble in alcohoi.
6. BcACtiflBs. — ^— With BfftaJs and their compounds. — Plu^phorio aci^i (iis>
eolres sasae izm-taIs. f.f^ F<e . 2n and Kg with evolution of hY(in>^Mi, It xmitos
with the oxid«« and hjdrDxides of the alkalis and alkaline earths and wiih
other fmhJv precipiiated oxides and hydroxidt^ exeept perhaps antinuMun^>
oxide. It aUo decompoises aU oarlxuiates'evolvinir CO^ . Phosphates are forntiv*,
in the abore reaetiona. the composition of whioh deiH'nds u|hmi the iH>nditions
of the experiment.
Free orthopbosphoric acid is not preeipitateil by onlinary salts of thirtl.
fourth and fifth {rronp metals <in instance of ferric chloride, a distinction fi\Mu
pyrophospfacmc acid and metaphosphoric acidK* but is pnvipitated in i^irt by
riher nitrate* and lead nitrate and acetate. Ammoniacal solution of calcium
ehloride or of barium chloride precipitates the normal pht^sphate.
Free metaphospfaoric acid precipitates solutions of silver nitrate, lead nitrate,
tnd lead acetate, the precipitates beinfr insoluble in excess of nietaphosphorio
add, and soluble in moderately dilute nitric acid. I^irium. calcium and ferrtMis
ehlorides. and mairnesium. aluminum, and ferrous sulphates, are no! prtvipi
tated by free metaphosphoric acid. Ferric chloride is pnH*ipitatiHl, a distinc-
tion from orthopbosphoric acid.
Free pyropfaosphoric acid inves precipitates with solutions of silver nitrate.
lead nitrate or acetate, ard ferric chloride: no prtvipitates with l^irinm or
calcium chloride, or with mafrnesium or ferrous sulphate.
Orthopbosphoric acid — or an orthophosphate with acetic acid- W*>«*.t M«»f «nmii»h-
Iflff e^ albtimen or gelatine. This is a distinction of Inith orthopbospborio
add and pyrophosphoric acid ^rom meta phonephoriv and.
With silyer nitrate soluble orthophosphates form silver orthophospliate,
AgjPO^, yellow: with metaphosphates, silver metaphosphate, AgPO, ,
vhite; and with pyrophosphates, silver pyrophosphate, Ag«P,.0; , white,
all Boluble in ammonium hydroxide. Silver metaphosphate is soluMe ii)
excess of an alkali metaphosphate (distinetion from pyroi)hosphatesV
If a disodium or dipotassium orthophosphate is added to solution of Mixer
nitrate, free acid is formed, and an acid reaction to te.st-i>aper is induced {o),
But^^ith a trisodium or tripotassium phosphate, the .solution remains neutral
(ft)-<i means of diiftinftuishit9g the acid phosphates fnn» the fiunmiL
(a) 13ra,HP0« + 3A«rN0, = Ag;PO. + 2NaN0, + HNO,
(b) Na,PO« 4- 3AgN0, = Ag,PO. -|- liNaNO,
free orthopbosphoric acid forms no precipitate with rcajjrcnt silver nitrate.
With lead acetate or nitrate, Na^>HP04 forms Pb.,PO^ , white, insoluhl»»
^11 acetic acid, as are also the phosphates of aluminum and ferrieum. With
*A8olution containing 5 p. c. ferric c/ilorC/f<r, mixcxl with one-ftmrth its vnhiine ^^f u lo \\ *'
■ J^jUonof ort/iop/uwp/ioric acid. re<iuire8 that near half of the lattor !h» nouirall/«Ml utn that
k !**l*ate ia to phosphoric acid as 1.1 U is to 1.000) before preiMpitiit ion ocours. Oi» the othir
^-4oc.of a5 p. c. solution of ferric chloride, mixed with 1 oo. of ii ft p. o. Mtlution of mct,t
"^acid, form a precipitate, to dissolve which, 20 cc. of the sanio m(>tupho8phorlo arid
' ( 00. of a 24 p. c. solution of hydrochloric acid arc roquln»a. Four eo. of a ft p. o.
"■ernttrote with 1 cc. of a 10 p. c. solution of r>rf/iop/i<wp;ior<r acid rIvo a prtvipi-
• which requires 7 cc. of the same orthophosphorio acid solution. [The Author's
Mr. MorgaD, Am, Jour. Phar.. 1876, 4S, 534. Kratschraor and Sztankovanaky,
502 PHOSPHORIC ACID. §258,61
PbCl, the precipitate always contains a chloride. Free phosphoric acid,
H3PO4 , forms an acid phosphate, PbHPO^ (Heintz, Pogg., 1848, 73, 119).
Lead salts also form white precipitates with soluble pyro and metaphos-
phates; the p}TO salt, PbaPjOy , is soluble in an excess of Na^P^O^ . Bis-
muth salts form BiPO^ , insoluble in dilute HNO3 .
Solutions of orthophosphates give, with soluble ferrio, ohromiey and
idnminnm salts, mostly the normal phosphates, FcPO^, etc. The ferric
phosphate is but slightly soluble in acetic acid, and for this reason it is
made the means of separating phosphoric acid from metals of the earths
and alkaline earths (§152). Solution of sodium or potassium acetate \»
added; and if the reaction is not markedly acid, it is made so by additioB
of acetic acid. Ferric chloride (if not present) is now added, drop by
drop, avoiding an excess. The precipitate, ferric phosphate, ]s brownish-
white.
With zinc and manganous salts, the precipitate is dimetallic or nonnal—
ZnHPO^, or Zn3(P04)2 — according to the conditions of precipitation.
When a manganic compound is mixed ^^ath aqueous phosphoric acid, the
solution evaporated to dryness and gently ignited, a violet or deep blue
mass is obtained, from which water dissolves a purple-red manganic
hydrogen phosphate, a distinction from manganous compounds. With salts
of nickel, a light green normal phosphate is formed; with cobalt, a reddish
normal phosphate.
Soluble salts of the alkaline earth metals, with dimetallic alkali phos-
phates, as Na^HPO^ , form white precipitates of phosphates, two-thmh
metalliCy as CaHFO^ ; with trimetallic alkali phosphates, white precipitates
of phosphates, normal or full metallic, as Ca3(P04)2 . The precipitates art'
soluble in acetic acid, and in the stronger acids. Concerning the am-
monium magnesium phosphate, see §189, 6d,
Magnesium salts with ammonium hydroxide give a precipitate of double
pyrophosphate, soluble in alkali pyrophosphate solution.
^fagnesium salts with ammonium hydroxide are not precipitated by
soluble metaphosphates unless very concentrated.
Ammonium molybdate, in its nitric acid solution (§75, 6flf), furnishes an
exceedingly delicate test for phosphoric acid, giving the pale yellow pre-
cipitate, termed ammonium phosphomolyhdate. The molybdate should be
in excess, therefore it is better to add a little of the solution tested (which
must be neutral or acid) to the reagent, taking a half to one cc. of the
latter in a test-tube. For the full delicacy of the test, it should be set
aside, at 30° to 40°, for several hours.
Ammonium molybdate reacts but slowly with meta or pyrophosphate
solutions — and not until orthophosphoric acid is formed by digestion with
the nitric acid of the reagent solution.
^265, 8. PHOSPHORIC ACID. 30a
B. — With non-metals and their compounds. — Phosphoric acid is not
reduced by any of the reducing acids. Phosphates of the first two groups
(ire transposed by HjS , and of the first four groups by alkali sulphides
with formation of a sulphide of the metal, except Al and Cr , which form
a hydroxide; phosphoric acid or an alkali phosphate is also formed.
HCl, HNO39 and HjSO^ transpose all phosphates and all are transposed
by acetic acid except those of Pb , Al and Fe'" phosphates. Sulphurous acid
transposes the phosphates of Ca, Mg, Mn, Ag, Fb, and Ba, also the
arsenite and arsenate of calcium (Gerland, J. C, 1872, 25, 39). Excess of
phosphoric acid completely displaces the acid of all nitrates, chlorides, and
siulphates upon evaporation and long-continued heating on the sand bath.
7. Ignition with metallic mag^neeium (or sodium) reduces phosphorus from
phosphates to magnesium phosphide, PsMgg , recognized by odor of PH, ,
formed on contact of the phosphide with water. A bit of magnesium wire (or
3f sodium) is covered with the previously ignited and powdered substance in
Ek glass tube of the thickness of a straw, and heated. If any combination of
phosphoric acid is present, vivid incandescence will occur, and a black mass
will be left. The latter, crushed and wet ^ith water, gives the odor of phos-
phorus hydride.
Orthophosphoric add heated to 213° forms pyrophosphoric acid; when heated
to dull redness the meta acid is obtained, which sublimes upon further heatings
without change. Phosphoric anhydride, P2O5 , cannot be prepared by ignition
of phosphoric acid. Tribasic orthophosphates, normal pyrophosphates, and
metaphosphates of metals whose oxides are not volatile and not decomposed
by heat alone are unchanged upon ignition. Dimetaliic orthophosphates,
M'aHPO^ , are changed to normal pyrophosphates upon ignition; also tribasic
orthophosphates when one-third of the base is volatile, e, f/., MgNH^PO^ .
Mono-metallic or primary orthophosphates, M^H^POt , become metaphosphates;
also secondary or tertiary orthophosphates when only one atom of hydrogen
is displaced by a metal whose oxide is non-volatile, e. (/., NaNH4HP04 .
Acid pyrophosphates, M^HoPjOr , form metaphosphates. When meta or pyro-
phosphates are fused with an excess of a non-volatile oxide, hydroxide or
carbonate the tertiary orthophosphate is formed (Wntts\ 1894, IV, 106).
Phosphates of Al , Cr , Fe , Cu , Co, Ni , Mn , Gl and TJ when heated to a
white heat with an alkali sulphate form oxides of the metals and an alkali
tribasic orthophoFphate: phosphates of Ba , Sr , Ca , Mg, Zn and Cd form
double phosphates, partial transposition taking place (Derome, C. r., 1879, 89,
952; Grandeau, A. C/t., 1886, (6), 8, 193).
8. Detection. — The presence of orthophosphoric acid in neutral or acid
solutions is detected by the use of an excess of an ammonium molybdate
solution (§76, 6d). With pyro and metaphosphoric acids no reaction is
obtained except as they are changed to the ortho acid by the reagents
used. Disodium phosphate, NaoHF04 , after precipitation with silver
nitrate, reacts acid to test papers. With trisodium phosphate the solu-
tion is neutral (distinction). Orthophosphates are distinguished from
pyro and metaphosphates by the color of the precipitate with silver nitrate:
AggPO^ is yellow, Ag^PoO- and AgPO., are white. Also by the fact that
only the ortho acid is precipitated by ammonium molybdate. Nearly all
|.yrophosphates are soluble in sodium pyrophosphate, Na^PaO^ (distinc-
304 SLLPHUR, §265,9.
lion from orthophosphates). Hager (J. C, 1873, 26, 940) gives a method
for detecting the presence of H3PO3 , H3A8O3 , or HNO3 in H3PO4 . Sodium
metaphosphate does not give a precipitate with ZnSO^ cold and in excess;
with Na^PjO^ and 'S2l.J1.^^0^ a white precipitate of ZiisFoO^ is obtained
(Knorre, Z. angew., 1892, C39).
9. Estimation. — (a) By precipitation as magnesium ammonium phosphate,
MgNH^PO^ , and ignition to the pyrophosphate, (b) By precipitation and
weighing as lead phosphate, Pb3(P04)2 . (c) By precipitation from neutral or
acid solution by ammonium molybdate and after drying at 140® weighing as
ammonium phosphomolybdate. Consult Janovsky (J. C, 1873, 26, 91) for a
review of all the old methods.
§266. Sulphur. S = 32.07 . Usual valence two, four and six.
1. Properties. — Sulphur is a solid, in yellow, brittle, friable masses (from
meltinff): or in yellowish, gritty powder (from Huhlimatiofi) or in nearly white,
slightly' cohering, finely crystalline powder (by precipitation from its com-
pounds). At — 50° it is white (Schoenbein, J. pr., 1852, 55, 161). The specific
yruvity of native sulphur is 2.0:J48 (Pisati, B,, 1874, 7, 301). Melting point. 111*
(Quincke, ./., 1868, 21). Boilimj point, 444.53° (Callendar and Griflfiths, C, N., 1891,
63, 2). Vapor density at 1160° is 34, indicating that the molecule is S3 (Bineau,
C. r., 1859, 49, 799) ; but at lower temperatures the molecule seems to vary from
S3 to S, . Sulphur is polymorphous, existing in various crystalline forms,
rhombic, monoclinic and triclinic systems, and also in amorphous conditions.
It is also classified by the relative solubilities of the various forms in carbon
disulphide. In chemical activity, volatility and other properties it stands as
the second member of the Oxygen Series:' O, 16.000: S, 32.07; Se, 79.2: and Te,
127.5. On being heated it melts at 111° to a pale yellow liquid: as the tempera-
ture rises it grows darker and thicker, until at about 1S0° it is nearly solid.
!?o that the dish may be inverted without spilling. At 200° it again becomes a
liquid as at first; and at 444.53° it boils and is converted into a brownish-red
vapor. If it is slowly cooled, exactly the same physical changes take place in
the reverse order, becoming thick at 180° and thin again at 111°, and at lower
1cmi)cratures solid. If, at a temperature near its boiling point, it is poured
into cold water, it forms a soft, ductile, elastic string, resembling india-rubber.
In a few hours this ductile sulphur changes back to the ordinary form, the
change evolving heat. But if poured into water from the other liquid form-
that is, at 111°— it forms only ordinary, brittle sulphur. In contact with air
sulphur ignites at 248° (Hill, C, .V., 1890, 61, 125); burning in air or oxygen
with a pale blue flame and penetrating odor to SO2 .
The isolated oxides of sulphur are SO^j , SO3 , S..0, and S.Ox . Sulphur and
oxygen combine directly to form SO., and SO,: the former by burning sulphur
in oxygen, the latter by the action of ozone upon SO.; also by burning sulphur
with oxygen under several atmospheres pressure. SjOa is made bj' dissol^^np
sulphur' in sulphur dioxide; S.O^ by the action of the electric discharge upon
a mixture of SO, and O .
2. Occurrence. — (a) Found in a free state, and as SOa in volcanic districts.
(h) As H.S in some mineral springs, (c) As a sulphide: iron pyrites, Fe^-
copper pyrites. CuFeS.: orpiment, As.S,: realgar, As^S^; zinc blende, ZnS:
cinnabar,' HgS; galena, PbS. (//) As a sulphate: gypsum, CaS0,.2H,0: heav.v
spar, BaSO^; kieserite. 'M.gSOtMiO; bitter spar (Epsom salts), MgS04,THiO-
Glauber salt. NBoSO^.IOH.O , etc.
3. Formation. — {a) By decomposing poly sulphides with HCl (Schmidt. Phnjr-
mfurutisehe Chcmie, 1898, 175). (ft) 13y adding an acid to a solution of a thio-
sulphate. {c) By the reaction between SO™ and HjS: 2S0j -|- 4HaS = 38. +
4H.0 . id) By the decomposition of metallic sulphides with nitric acid: 2BisSi
-f ioHNO, =*^4Bi(NO,)8 + 3S, + 4N0 + SH^O .
§257,4. HYDROSLLPHLRia ACID. 307
2. Occurrence. — Found free in volcanic gases and frequently in mineral
springs. While the inhaled gas is poisonous, the mineral waters containing it
are reputed to be a healthful beverage.
3. Formation of Hydrosulphuric Ax^id. — («) By direct union of the elements
when passed over pumice stone heated to 400** (Coren winder, A. Ch,f 1852, (3),
34, 77). {b) Heating paraffin or tallow with sulphur (Fletcher, C. A\, 187U, 40,
154); and by passing illuminating gas through boiling sulphur (Taylor, C. X.,
1883, 47, 145). (c) The sulphur in coal becomes HjS in the process of gas-
making, (d) From steam and sulphur at 440**. (e) Often occurs in nature from
reduction of gypsum by decaying organic matter (Myers, J, pr., 18G9, 108, 123).
(f) Transposition of sulphides by hydracids or by dilute phosphoric or dilute
sulphuric acid, (y) Decomposition of organic compounds containing sulphur.
Formation of Sulphides. — (/) By fusion of the metals with sulphur, see
§256, 7. (2) By action of H,.S upon the free metals, hydrogen being evolved.
With Hg and Ag this occurs at ordinary temperature, but with most metals a
higher temperature is needed. (J) Action of H..S on metallic oxides or
hydroxides. Those sulphides which are decomposed by water (e. r/., Al.S, ,
Gr^S,) are not formed in its presence, but by action of HjS upon the oxide at
a red heat. (.^) By action of soluble sulphides upon metallic solutions. The
ordinary sulphides of the first four groups are formed thus, except ferric salts,
which are precipitated as FeS , and aluminum and chromic salts as hydroxides,
(o) By action of CS2 upon oxides at a red heat. (6) By action of free sulphur
upon oxides at a red heat. (7) By the action of charcoal upon the oxyacids of
sulphur at a red heat in presence of an alkaline carbonate. To prepare a
sulphide absolutely arsenic free, take BaSOi , 100 grams: coal, pulverized, 25
grams; and NaCl, 20 grams, mix, ram into a clay crucible and ignite to a
white heat for several hours (Winkler, Z., 1888, 27,' 2G). (.S) By the action of
zinc amalgam on sulphuric acid'(Walz, C. X., 1871 23, 245). (V) As a reagent
for the formation of metallic sulphides in analvsis it is recommended by
Schiff and Tarugi (B., 1894, 27, 3437), Schiff (/?., 'lS95, 28, 1204), and Tarugi
{Oazzetta, 1895, 25, i, 269), to use ammonium thioacetate, CHsCGSNH.: prepared
by distilling a mixture of phosphorus pentasulphide and glacial acetic acid
(300 grams each) with 150 grams of cracked glass. A large distilling flask is
used and the distillate is collected to 103°. It is then dissolved in a slight
excess of ammonium hydroxide, diluting to three volumes from one volume
of the acid. Salts of the metals of the first two groups in acid solution are
readily precipitated as sulphides upon warming with this reagent.
1, 2Fe -1-82 = 2FeS
2. 2Ag + H,S = Ag,S -f Ha
5. Pb(OH), -f H,S = PbS + 2H,0
4Fe(0H), -h 6H0S = 4FeS + S, + 12HaO
4. 4FeCl, + 6(NHJ,S = 4FeS + S, + 12NH4CI .'^\
5. 2CaO + CS2 = 2CaS + 00, '^
6. 4CaO -f 3S2 = 4CaS + 280,
7. K,80, + 2C = K28 + 200,
4. Preparation. — For laboratory purposes it is nearly always made by
adding HgSO^ or HCl to FeS . The ferrous sulphide is prepared either
by fusion of the iron with the sulphur, or by bringing red hot iron rods
in contact with sticks of sulphur, and is made to drop into tubs of cold
water. Dilute H,SO^ should be used: FeS + HoSO^ = FeSO^ + H,,S .
Concentrated H.^SO^ has no action on FeS , unless heated and then SOo is
evolved : 2FeS + lOHoSO^ = FCoCSO,)., + 9SO2 + lOHoO ; and frequently
free sulphur is formed by the action of the H^S upon the SOo first formed.
The colorless ammonium sulphide, (NHJoS, is prepared by saturating
308 HIDROSULPHCRIC ACID, §257,5.
ammonium hydroxide with HgS until a sample will no longer give a pre-
cipitate with a solution of magnesium sulphate ; showing that ammonimD
hydroxide is no longer present. Upon standing the solution gradually
becomes yellow with fornuition of the polysulphides or yellow ammoidiun
sulphide, (NH4)2S, This may be hastened by the addition of sulphur
(Bloxam, J. C, 1895, 67, 277).
Sodium sulphide, NaoS, is prepared by neutralizing an alcoholic solution
of NaOH with HjS and then adding an equal amount of NaOH and allowing
to crystallize; air being excluded. The various polysulphides, NaoS, to
NagSj; , are prepared by boiling the normal sulphide with the calculated
amounts of sulphur (Boettger, .4., 1884, 223, 335; Geuther, A., 1884, 2SML
201).
5. Solubilities. — At 15° water dissolves 2.6G volumes of the gas HgS.
Sulphides which dissolve in dilute HjSO^ evolve H2S, c. g,, CdS, FcS,
MnS, ZnS, etc. But if a sulphide requires concentrated HjSO^ for its
solution ; S and SO2 are formed or SO^ alone ; e. g., BigS., , CuS , HgS . If
concentrated HgSO^ be used upon a sulphide that might have been dis-
solved in the dilute acid, then no H.S is evolved: ZnS + 4H2SO4 = ZnSO^
+ 4SO2 + 4HoO . Or with a small amount of water present: 2ZnS +
4HoS0^ = 2ZnS0^ + So + 2SO2 + 4H,0 . The sulphur of the zinc sul-
phide is oxidized to free sulphur and that of the sulphuric acid is reduced
to sulphur dioxide. HgS is almost insoluble in HNO3 , dilute or concen-
trated, readily soluble in chlorine, nitrohydrochloric acid, or chloric acid
if hot. ^lost other sulphides are sohi])le in hot HNO^ (^74, Gp). Long
continued boiling with water more or less completely decomposes the sul-
phides of Ag , As , Sb , Sn , Fe , Co , Ni , and Mn ; no effect with sulphides
of Hg, Au, Pt, Mo, Cu, Cd, and Zn (Clermont and Frommel, A. CI.
1879, (5), 18, 203).
As a reagent, hydrosulphuric aqid, gas or solution in water finds ex-
tended application in the analytical laboratory. The grouping of the
bases for analysis depends very largely upon the relative solubilities of the
sulphides. Plydrosulphuric acid in alkaline solution, alkali sulphide or
poly sulphide, is a scarcely less important reagent, being especially valuable
in the subdivision of the metals of the second group.
The sulphides of the first four groups are insoluble. Hydrosulphuric
acid transposes salts of the first two groups in acid, neutral, and alkato^
mixtures, except arsenic, which is generally imperfectly precipitated un-
less some free acid or salt that is not alkaline to litmus be present. The
result is a sulphide, but mercurosum forms mercuric sulphide and mer-
cury, and arsenic acid may form arsenous sulphide and free sulphur.
Ferric solutions are reduced to ferrous with liberation of sulphur. In acid
mixture other third and fourth group salts are not disturbed, but from
§267, ^A8. HYDROSULPHURW ACID, 309
solutions of their normal salts traces of cobalt, nickel, manganese, and
zinc (§136, 6e) are precipitated.
Soluble sulphides transpose salts of the first four * groups. The result
is a sulphide, except that with aluminum and chromium salts it is a
hydroxide, hydrosulphuric acid being evolved. With mercurous salts,
mercuric sulphide and mercury are formed; with ferric salts, ferrous sul-
phide and sulphur.
The precipitates have strongly marked colors — that of zinc being white ;
manganese, flesh colored; those of iron, copper, and lead, hlach; arsenic
stannic and cadmium, yellow; antimony, orange-red; stannous, brown; mer-
cury, successively white, yellow, orange, and black,
6. Beactions. A, — With metals and their compounds. — Some metals
are converted into sulphides on being treated with hydrosulphuric acid;
e, g,, Ag, Cu, Hg, etc. The alkali polysulphides slowly attack many
metals with formation of sulphides: Sn becomes M'sSuS, ; Ag becomes
AgjS, no action with colorlesc (NHJoS ; Ni forms NiS ; Fe, FeS; Cu,
CuS and then CUoS (with colorless ammonium sulphide, (1^4)28, CUjS
is formed with evolution of hydrogen) (Priwozink, A., 1872, 164, 46).
The hydroxides or non-ignited oxides of Pb", Ag , Hg", 8b , Sn , Bi'",
Cu, Cd, Fe", Co", Ni", Mn", Zn, Ba, Sr, Ca, Mg, K, Na, and NH^
unite with moist H^S at ordinary temperature to form sulphides without
change of the valence of the metal. In other cases the valence of the
metal is changed, usually with liberation of sulphur.
1, Pb"+° becomes PbS and S .
2, As^ in acid solution forms some ASgSs and S . See §69, 6e.
3, Hg' becomes HgS and Hg .
Jf, Cr^^ becomes Cr'" and S, if the HgS be in excess: 2K2Cr207 -f SHjS
= 4Cr(0H)3 + 3So -f 2K2S -f 2H2O .
5, Fe'" becomes Fe" and S : 4FeCl3 + 2Ho8 = 4FeCl2 + 4HC1 -f 8, .
If the solution be alkaline FeS is precipitated: 4FeCl3 + 6KjjS = 4FeS -f
12KC1 + S2 .
6, Co"+" becomes Co" and S .
7, Ni"+° becomes Ni" and S .
8, Mn"+° becomes Mn" and S. In alkaline solution with excess of
XMnO^ , an alkali sulphate is formed and MnOj : SEMnO^ + SKjS =
SKjSO^ + 4KoO + 8Mn02 (Schlagdcnhafen, Bl, 1874, (2), 22, IG).
In the above reactions, if an alkaline sulphide be used instead of hydro-
sulphuric acid, the metal will be precipitated as a sulphide with the
* The normal fixed alkali sulphides (Ka,8, K,8), precipitate solutions of calcium and mag-
nesium salts as the hydroxides : Ca(C,H,0,), + 2Na,8 + 2H,0 = Ca(OH), + 23iaC,B,0, +
2MaUS. No reaction with the acid fixed alkali sulphides (NaH8« KH8) or with ammonium,
sulphides (Pelouze, A. Ch„ 1868, (4), 7, 172).
310 HTDROSULPHIRIC ACID. §267, 6R
formation of an alkali hydroxide ; except that the arsenic will remain in
solution (§69, 5c) and the chromium will be precipitated as the hydroxide.
Dry HgS has no action on the dry salts of Pb, Ag, Hg, As, Sb, Sn,
Bi , Cu , Cd , or Co ; nor does it redden dry blue litmus (Hughes, Phil
Mag,, 1892, (5), 33, 471).
Many insoluble sulphides, freshly precipitated, transpose the solutions of
other metallic salts. In some cases the action is quite rapid at ordinary tem-
perature, in others long-continued heating" (several hours) at 100® is necessary.
PdS is formed by action of FdCl, with sulphides of all the metals following in
the series below named, but PdS is not transposed by solutions of the metals
following". Silver salts form AgjS with sulphides of the metals following in the
series but not with sulphides of Pd and Hg , etc.: Pd , Hg , Ag , Cu , Bi , Cd,
Sb , Sn , Pb , Zn , Ni , Co, Fe , As , Tl and Mn (Schiirmann, A., 1888, 240, 326).
B. — With non-metals and their componnds.
1. H3Fe(CN)e becomes H^Fe(CN)« and S. Proof: Boil to expel the
excess of hydrosulphuric acid, then add ferric chloride (§126, 66).
2. HNO3 becomes NO and S . If the HNO3 be hot and concentrated the
sulphur is oxidized to sulphuric acid.
3. HgS has no reducing action on the acids of phosphorus.
4. H2SO3 becomes pentathionic acid, H2S5O0, and sulphur: lOH^SO, +
IOH2S = 2H2S5O0 + 5S2 + I8H0O . With excess of HoS the product is
entirely free sulphur from both compounds: 2H2SO3 + IH.S = ^8^ +
6H2O (Debus, J. C, 1888, 53, 282).
H2SO4 , dilute no action; concentrated and hot, S and SO2 are formed:
2H2SO, + 2H,S = So + 2SO2 + 4H,0 (§256, i5BJf),
5. CI with H^S in excess forms HCl and S ; with CI in excess forms HCl
and H2SO4 .
HC163 with H.S in excess forms HCl and S ; with HCIO3 in excess HCl
and H0SO4 .
6. Br with H^S in excess forms HBr and S ; with Br in excess HBr and
H2SO, .
HBrOg with HoS in excess forms HBr and S ; with HBrOa in excess HBr
and HoSO^ .
7. I becomes HI and S (Filhol and Mellies, ^1. Ch., 1871, (4), 22, 58).
HIO3 becomes HI and S .
7. Ignition. — Dry hydrosulphuric acid pras is not decomposed when heated to
350** to 360**. At this temperature AsHaJn presence of potassium polysulphide,
K^S, , liver of sulphur, is decomposed: 2AsH, -f SK^S, = 2K,AbS3 + SH^S;
thus furnishing a ready means of purifying* HjS for toxicological work (§69.
6'6) (Pfordten. B., 1884, 17, 2897).
If air be excluded some sulphides may be sublimed unchanged; e.g., HgS.
ASjSs , AB2S5 , SbjS, , etc. In some cases part of the sulphur is separated,
leaving a sulphide of a lower metallic valence: 2reS2 = 2reS -f S^ . Some
sulphides remain unchanged upon ignition in absence of air; e. .7., FeS , MnS.
CdS , etc. All sulphides suflPer some change on being ignited in the air: some
slowly, others rapidly; Sb,S, , CuS, Al^S, , Cr,S, , etc., evolve SO, and leave
|257, 9. HYDR08ULPHURIC ACID, 311
the oxide of the metal; Hg^S , Ag,S , etc., evolve SO, and leave the free metal.
All sulphides, as well as all other compounds of sulphur, when fused with KNO.
or KClOa in presence of an alkali carbonate are oxidized to an alkali sulphate;
forming NO or KCl and evolving CO, . The metal is changed to the carbonate,
oxide or the free metal (§228, 7).
When ignited on charcoal with sodixim carbonate — or {diatinciion from
sulphates) if ignited in a porceUiin crucible with sodium carbonate — soluble aodiwn
sulphides are obtained. The production of the sodium sulphide is proved by the
(ilack stain of AgaS , formed on metallic silver by a moistened portion of the
fused mass. (Compounds of selenium and tellurium, §§112 and 113.)
8. Detection. — (a) The odor of the gas constitutes a delicate and char-^
acteristic test when not mixed with other gases having a strong odor.
{b) The gas blackens filter paper moistened with a solution of lead ace-
tate, delicate and characteristic. In the detection of traces of the gas,
a slip of bibulous paper, so moistened, may be inserted into a slit in the
smaller end of a cork, which is fitted to the test-tube, wherein the material
to be tested is treated with sulphuric acid; the tube being set aside in a
warm place for several hours. If any oxidizing agents are present — as
chromates, ferric salts, manganic salts, chlorates, etc. — hydrosulphuric
acid is not generated, but instead sulphur is separated, or sulphates are
formed (6). (r) The gas blackens silver nitrate solution, delicate but
PH3, AsHs, and SbHg also blacken silver nitrate solution, {d) By its
reducing action upon nearly all oxidizing agents with separation of sul-
phur, which is detected according to §266, 8. EHnO^ is perhaps the most
delicate test but the least characteristic, (e) Its oxidation to a. sulphate
is characteristic in absence of other sulphur compounds. This method
is usually employed with sulphides not transposed by dilute H2SO4 ;
chlorine, nitrohydrochloric acid or bromine being the usual oxidizing
agents. Also, these sulphides and certain supersulphides, attacked with
difficulty by acids, as iron pyrites and copper pyrites^ are reduced and
dissolved, with evolution of hydrosulphuric acid, by dilute snlphuric acid
with zinc. The gas, with its excess of hydrogen, may be tested by method
if)' if) Sodium nitroferricyanide gives a very delicate and characteristic
test for HoS as an alkali sulphide. The gas is passed into an excess of
alkali hydroxide; and to this mixture the reagent is added, producing a
transient reddish-purple color. Free HgS, dilute, remains colorless; a
concentrated solution gives a blue color, due to the reducing action of
the HoS on the ferricyanide.
For method of separation of the various sulphur compounds from each
other consult Kynaston (J. C, 1859, 11, 16G) and Bloxam {C, xV., 1895,
72, 63).
9. Estimation. — Sulphides are usually oxidized to HgSO^ (by chlorine,
bromine, or nitrohydrochloric acid, or by fusion with ENOg and NasCO,)
precipitated with BaCls and weighed as BaSO^ .
312 THIOSVLPHURIC ACID. §858|1.
§258. Thiosnlphuric acid. H^SjO, = 114.156 .
Diihion^^ acid.
II
H',(S,)^0-%, H— — 8 — S — H.*
II
1. Properties.— Thiosulphuric acid, HaSjO, (formerly called hyx>osulphiiroTis
acid), has not been isolated; but it almost certainly exists in dilute solutions,
when a dilute weak acid is added to a solution of sodium thiosulphate, KasSjO, ,
soon beginning to decompose into H2SO. and 8 (Landolt, B., 1883, 16, 2985)!
The thiosulphates are not particularly stable compounds, some decomposing
almost immediately upon forming; e. r/., mercury thiosulphates. Alkali thio-
sulphates decompose upon heating into sulphate and polysulphide: 4K&,S30, =
3Ka,S0« -f NajSs . Other salts give also S and H,S . Boiling solution of a
thiosulphate gives a sulphate and H^S or a sulphide of the metal.
2. Occurrence. — Not found in nature.
3. Foniiation. — Thiosulphates are formed by the oxidation of alkali or
alkaline earth polysulphides by exposure to the air or by SO, or KxCr,0,:
2CaS, + 30, = 2CaS,0, -f 38,; 4Na,8. + 680, = 4Na,8,0, -f 98,; 2K,8, +
4K,Cr,0T 4- 13H,0 = 5K,8,0, + 8Cr(0H), + 2K0H (Doepping, A., 1843, 46,
172; Gueront, C. r., 1872, 75, 1276). Also by heating ammonium sulphate with
phosphorus pentasulphide (Spring, B., 1874, 7, 1157).
4. Preparation. — Thiosulphates are prepared by boiling sulphur in a solu-
tion of. normal alkali sulphite: 2Na280, + 8, = 2Na,830, . Fixed alkali or
alkaline earth hvdroxides with sulphur also form thiosulphates: 3Ca(0H), -f
68, = 2Ca85 -i-'^Ca8,0, -f 3H,0 (Filhol and Senderens, C, r., 1883, 96, 839:
Senderens, C. r., 1887, 104. 58). Commercial sodium thiosulphate is prepared
by passing. SO, into ** soda waste" suspended in water, calcium thiosulphate
being formed. This is treated with sodium sulphate, filtered and evaporated
to crystallization.
5. Solubilities. — The larger number of the thiosulphates are soluble in water:
those of barium, lead and silver being only very sparingly soluble. The thio-
sulphates are insoluble in alcohol. They are decomposed, but not fully dis-
solved, by acids, the decomposition leaving a residue of sulphur.
Alkali thiosulphate solutions dissolve the thiosulphates of lead and silver:
also the chloride, bromide and iodide of silver, and mercurous chloride: the
iodide and sulphate of lead; the sulphate of calcium, and some other precipi-
tates — by formation of soluble double thiosulphates:
Ag^SjO, + Na^SjO, = 2NaAgS,0,
AgCl + Na,8,0, = NaAg8,0, + NaCl
Pb80, -f- 3Na,8,0, =Na,Pb(8,00, + Na,SO,
6. Beactions.— A.— With metals and their compounds.— With soluble thio-
sulphates, solutions of lead and silver salts are precipitated as thiosulphates,
white, soluble in excess of alkali thiosulphate. These precipitates decompose
upon standing, rapidly on warming, into sulphides and sulphuric acid: Agfii^*
+ H,0 = Ag.S -f H,SO, . Soluble mercury salts with sodium thiosulphate
form a white precipitate, almost instantly turning black with decomposition to
mercuric sulphide. NajS-^O, blackens HgCl , a portion of the mercury going
into solution, colorless, reprecipitated black upon warming.
Acid solutions of arsenic and antimony are precipitated by hot solution of
*Bunte,B., 1874,7,646.
§258,8. THI08VLPHURIC ACID. 313
Ka,S,0, as sulphides, ASaS, and SbjS, (a separation from tin,* which is not
precipitated) (6e, §§69, 70 and 71). Solutions of copper salts with thiosul-
phates, on long standing, precipitate cuprous salt, changed by boiling to
cuprous sulphide and sulphuric acid (separation from cadmium, §78, 6e).
Solutions of ferric salts are reduced to ferrous salts with formation of sodium
tetrathionate: 2EeCl, -f aNa.S^O, = 2EeCla -f 2NaCl -f NaaS^O,; used as a
quantitative method of estimation, with a few drops of potassium thiocyanate
as an indicator. Chromic acid (chromates in acid solution) are reduced to
chromic salts with oxidation of the thiosulphate.
FermangazLates in neutral solution become manganese dioxide, in acid solu-
tion the reduction is complete to manganous salt, a sulphate and dithionate
being formed (Luckow, Z., 1893, 32, 53).
Barium chloride forms a white precipitate of barium thiosulphate, BaSaO, ,
nearly insoluble in water. Calcium 'chloride forms no precipitate (distinction
from a sulphite).
B. — ^With non-metals and their compounds. — When thiosulphates are decom-
posed by acids, the constituents of thiosulphuric acid are dissociated as sul-
phurous acid and sulphur. Nearly all acids in this way decompose thiosul-
phates: 2NaaSaO, -f 4HC1 = 4NaCl -f 2HaS03 -f S, .
Thiosulphates are reducing agents — even stronger and more active than the
sulphites to which they are so easily converted. This reduction is illustrated
by the action on arsenic compounds, on ferric .salts and on chromates and
permanganates as given above. Also the halogens are reduced to the halide
salts forming a tetrathionate: 2Na3Sa08 -f la ^^ 2NaI -f- Na2S4 0« . If chlorine
or bromine be in excess the tetrathionate is further oxidized to a sulphate:
KaaSaOs -f 4Cla -f SHjO = NaaSO, -f HaSO^ + 8HC1. Chloric, bromic and
iodic acids are first reduced to the corresponding halogens and then with an
excess of the thiosulphate to the halidcs, always accompanied with the separa-
tion of sulphur. Nitric acid is reduced to nitric oxide with the separation of
sulphur.
7. Ignition. — On Ignition, or by heat short of ignition, all thiosulphates are
decomposed. Those of the alkali metals leave sulphates and poly sulphides (a),
others yield sulphurous acid with sulphides, or sulphates, or both. The
capacity of thiosulphates for rapid oxidation, renders their mixture with
chlorates, nitrates, etc., explosive, in the dry way. Chlorates \yith thiosulphates
explode violently in the mortar. Cyanides and ferricyanides, fused with thio-
sulphates, form thiocyanates, which may be dissolved by alcohol from other
products. By fusion on charcoal with Na^COg , thiosulphates form sulphides
(ft) and (c); and by fusion with an alkali carbonate and nitrate or chlorate,
a sulphate is formed (d). By ignition of a metallic salt with NaaSaOa in a
dry test-tube the characteristic colored sulphide of the metal is obtained
(Landauer, B., 1872, 5, 406).
(a) 4NaaSaO, = NaaS, + 3NaaS0«
(6) NaaSaO, -f NaaCO. + 2C = 2NaaS -f 3C0,
(c) 2PbSa03 + 4Kra,C0, -f 5C = 4NaaS + 2Pb + 9C0,
(d) SNaaSaOa + 3NaaC0, + 4KC10, = 6Na,S0, -f 4KC1 + SCO,
8. Detection. — In analysis, thiosulphates are distinguished by giving a pre-
cipitate of sulphur with evolution of sulphurous anhydride when their solu-
tions are treated with hydrochloric acid; by their intense reducing power,
shown in the blackening of the silver precipitate; and by non-precipitation of
calcium salts.
The precipitation of sulphur ttith evolution of sulphurous anhydrid<*, by addition
of dilute acids — as hydrochloric or acetic — is characteristic of thiosulphates.
It will be understood, however, that in presence of oxidizing agents, which can
be brought into action by the acid, sulphides will likewise give a precipitate of
sulphur.
* According to Vortmann (If., 1886, 7, 418) sodium thiosulphate may be used instead of hydro
sulphuric acid in the second erroup of bases. An excess of the reagent is to be avoided and
nitric acid should be absent.
314 UYPOSi'LPlJiROii<l ACW-OITHIOXrC ACID.
in the presence of a sulphate and sulphite the thiosulphate is detected as
follows: Add BaCl, and NH4CI in excess, then HCl to solution of all but the
BaSO^ . Filter and treat the tiltrate with iodine, forming BaSO^ of the sulphite
and BaSfOe of the thiosulphate. Filter and add bromine to the filtrate, which
then forms BaSO« (Smith, C. .V., 1895, 72, 39).
9. Estimation. — By titration with a standard solution of iodine, or by titrat-
ing the iodine liberated by a standard solution of potassium dichromate (§§125,
10, and 279, 61^7).
§259. Hyposnlphnrons acid. H2SO, = 66.086 .
(Hydrosulphurons ar dithionoiis acid,)
H'2S"0-%, H — — 8 — H.
Obtained by Schiitzenberger (C. r., 1869, 69, 196) by the action of zinc on
sulphurous acid: Zn -f 280, -f H.O = ZnSO, + H^SO, . The sodium salt is
formed bv treating a concentrated solution of sodium acid sulphite with zinc
tilings: Zn -f aNaHSO, = ZnSO, -f NaaSO, + NaHSO, + H,0 . In the forma-
tion of the free acid or of the sodium salt no hydrogen is evolved. It is a very
unstable compound, a strong reducing agent, rapidly absorbs oxygen from the
air, becoming sulphurous acid or a sulphite. According to Bernthsen (B., 1881.
14, 4.18) the sodium salt does not contain hydrogen. He gives the formula as
Na^S^O^: Zn + 4NaHS0, = ZnSO, -f Na^SO, + Na^S^O, + 2H,0 . It is used
in the preparing of indigo white for the printing of cotton fabrics. See also
Duprg, J, C, 18G7, 20, 291.
§260. Dithionic acid. HjS^Oo = 1G2.156 .
II II
R\{S,)^0-\, H — — S — S — 0— H.
I! II
Known only in the form of its salts and as a solution of the acid in water.
The free acid or the anhydride has not been prepared. The manganous salt
is prepared by the action of a solution of sulphurous acid upon manganese
dioxide at a low temperature: MnO. -f 2H2SO, = MnSjO, -f 2H2O . Similar
results are obtained with iiickelic or ferric oxides (Spring and Bourgeois, Bl^
1S86, 46, 151). The acid is obtained by treating the manganous salt with
Ba(0H)2 and the filtrate from this with the calculated amount of HaSOi .
It is a colorless solution and may be evaporated in a Vacuum until it has a
specific gravity of 1.347. It decomposes upon further heating: HsSjOg = H.SO4
-h SOo . All other thioni<* comjHmmls dcrompoHC upon heathuj mih separation of
sulphur. By exposure to the air dithionic acid is oxidized to sulphuric acid.
All dithionates are soluble in water and may be purified by evaporation and
crystallization (Gelis, A. Ch., 1862, (3), 65, 230).
Dithionic acid is also prepared by carefully adding a potassium iodide solu-
tion of iodine to sodium acid sulphite (Hoist and Otto, Arr/i. Pharm., 1S91. 229.
171); Spring and Bourgeois {Areh. Pharm,, 1891, 229, 707) contradict the above
statement.
§262. TRITHIONIC ACID-TETRATHIONIC ACID. 315
§261. Trithionic acid. H^SsOe = 194.226 .
II II
IL\{S^y^O-\, H — — 8 — S — S — — H.
The free acid and anhydride are not known. The potassium salt is prepared
l)y boiling potassium acid-sulphite with sulphur (n); by treating potassium
thiosulphate with sulphurous acid (6) (no action with sodium thiosulphate)
(Baker, C. N., 1877, 36, 203; Villiers, C, r., 1889, 108, 402); by the action of
iodine on a mixture of sodium sulphite and thiosulphate (c) (Spring, B., 1874,
T1157):
(a) 12EHS0, + S, = 4X3830, + 2K3SO3 + 6H,0
(ft) 4X28,0. + 68O2 = 4K,S,0, + 8,
(c) Na,80, + Ka,8,0, + I^ = Na,8,0, + 2NaI
The acid is prepared by adding perchloric or fluosilicic acid to the potassium
salt. The acid is quite unstable; at low temperature in a vacuum it decom-
poses into 80a , 8 and H28O4 . The salts are quite stable; they are not oxidized
by chloric or iodic acids, while the free acid is rapidly oxidized by these acids.
Fixed alkalis or sodjum amalgam change the trithionate to sulphite and thio-
sulphate (Spring, Lc).
§262. Tetrathionic acid. HjS^Oo = 226.296 .
11 II
H'a(S,V^O-%, H — — S — S — S — S — — H.
The salts are soluble in water and are comparatively stable. They are best
obtained in crystalline form by adding alcohol to their solutions in water.
The acid has not been isolated but it is much more stable than the tri or
pentathionic acids. In dilute solution it can be boiled without decomposition.
The concentrated solution decomposes into H28O4 , SO^ and 8 .
Tetrathionates are prepared by adding iodine to the thiosulphates: 2Ba820, -j-
I, = Ba8«0« -f- Bal- (Mauraene, (\ r., 1879, 89, 422). The lead salt is obtained
by the oxidation of lead thiosulphate by lead peroxide in presence of sulphuric
acid; 2Pb820, -f PbO, -f 2H2SO4 = Pb8,0, -f 2PbS04 -f 2Ha0 (Chancel and
Diacon, J. pr., 18C3, 90, 55). To obtain the acid the lead should be removed
by the necessary amount of sulphuric acid, and not by hydrosulphuric acid,
which causes the formation of some pentathionic acid. A number of other
oxidizing agents may be u.sed to form the tetrathionate from the thiosulphate
(Fordos and Gelis, (\ r., 1S42, 15, 920). Sodium amalgam reconverts the tetra-
thionate into the thiosulphate: Na,84 0e -f 2Na = 2Na3830, (Lewes, J. C, 1880,
39, 68; 1881, 41, 300). Tetrathionic acid is also formed with pentathionic acid
in the reactions between .solutions of HjS and 80, (Wackenroder's solution,
A., 1846, 60, 189). See also Curtius and Henkel (J. pr,, 1888, (2), 37, 137). The
acid gives no precipitate of sulphur when treated with potassium hydroxide
(distinction from pentathionic acid).
316 PENTATHTOXW ACID. §ttt
§263. Pentatliionic acid. K^Sfi^ = 258.366 .
II II
IL\(S,y^O-\, H — — 8 — 8 — 8 — 8 — 8 — — H.
II II
Only known in the salts and in the solution of the acid in water. It is formed
by the action of H,S upon SO, in the presence of water (a) ; by the action of
water on sulphur chloride (b) ; by the decomposition of lead thiofiulphate with
HaS (Persoz, Pogg,, 1865, 124, 257):
a. 10H,SO, + 10H,S = 2H,Ss0« + 58, + 18H,0
6. 10S,C1, + 12H,0 = 2H,S,0« + 5S, + 20HC1
The filtrate from the decomposition of SO, by HjS is known as Waekenroder's
solution (Arch, Phami,, 1826, 48, 140). It has been shown to contain the tri
and tetrathionic acids in addition to the pentathionic acid (Debus, C. .V., 1888.
57, 87). Pentathionic acid may be concentrated in a vacuum until it has a
specific gravity of 1.6; farther concentration or boiling heat alone decomposes
it into HaSO* , SO, and S . The solution of the acid does not bleach indigo.
^Vhen treated with a fixed alkali hydroxide an immediate precipitate of sulphur
is obtained (distinction from H^SfOe): 4H3SBO, + 20NaOH = 6Ha,S0, +
4Ka,S,0, + 38, + 14H,0 (Takamatsu and Smith, J. C, 1880, 37, 592); or if the
NaOH be added short of neutralization: 10Ha850e -f 20NaOH = 10Na,S«0, 4-
5Sa + 20H2O . Neutralization of pentathionic acid with barium carbonate gives
barium tetrathionate and sulphur (Takamatsu and Smith, /. C, 1882, 41, 162;
Lewes, J, C, 1881, 89, 68). See also Spring, A., 1879, 199, 97.
TABLE OF TBIOVIC ACIDS.
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9
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2
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il8 SULPHUROUS ACID. §28S,L
§265. Snlphurons anhydride. SO^ = C4.07 .
Snlphnrous acid. H2SO3 = 82.086 .
II
8^0-% and H'^S^O^,, = 8 = and H — — S — — H.
1. Properties. — Sulphurous anhydride, SO, , sulphur dioxide, is a colorless gas
of a strong suflPocatingr odor of burning sulphur. Specific gravity of the liquid
at 0**, 1.4338 (Cailletet and Matthias, C. r., 1887, 104, 1563); of the gas at 0* and
760 mm. pressure, 2.2369 (Leduc, C. r., 1893, 117, 219). It is liquefied at atmos-
pheric pressure upon cooling to —10** (Pierre, C r., 1873, 76, 214). In an opes
dish it evaporates rapidly, the temperature of the remaining liquid dropping
to — 75**; or by evaporating rapidly under diminished pressure it becomes s
white wooly solid. Cooled to — 70.1** it becomes a snow-white solid (Faraday.
C. r., 1861, 53, 846). The dry gas is not combustible in the air, does not react
acid to litmus, but in i^resence of water it has a marked acid reaction. The gas and
the free acid, not the salts, are quite poisonous, due to the absorption of the
SO, by the blood and oxidation to H2SO4 . The gas is soluble in water, form-
ing probably sulphurous acid, H2SO, . The pure acid has not been isolated,
but forms salts mono and dibasic as if derived from such an acid (Michaclis
and Wagner, B„ 1874, 7, 1073). It has a strong odor from vaporization ol
sulphurous anhydride, which is soon completely expelled upon boiling. The
acid oxidizes slowly in the air, forming H3SO4 , hence sulphurous acid usually
gives reactions for sulphuric acid. Light seems to play an important part in
this oxidation (Loew, Am, »Sf., 1870, 99, 368). The moist gas or a solution of the
acid is a strong bleaching agent, however not acting alike in all cases. Wool,
silk, feathers, sponge, etc., are permanently bleached; also many vegetable sub-
stances, straw, wood, etc.; yellow colors and chlorophyll are not bleached; red
roses are temporarily bleached, immersion in dilute H2SO4 restoring the color.
2. Occurrence. — Found free in volcanic gases (Ricciardi. B., 1887, 20, 464).
3. Formation. — (a) By burning sulphur in air. (h) By heating sulphur with
various metallic oxides, (c) By decomposition of thiosulphates with HCl. (d)
By burning HjS or OS, in air. (e) By the action of hot concentrated sulphuric
acid on metals, carbon, sulphur, etc. (f) By heating sulphur with sulphates.
{g) By decomposition of sulphites with acids;
(a) S, -f- 2O2 = 2S0,
(&) MnOj + So = MnS -f SO,
2Pb5 04 + 5S, = 6PbS -f 4SO2
(c) 2Na2S20, -f 4HC1 = 4NaCl + 2S0, + S, + 2H,0
((f) 2H,S + 30, = 2SO, -f- 2H,0
CS, + 30, = 2S0, -f- CO,
{€) Cu -f 2H,S04 = CUSO4 + SO, + 2H,0
S, + 4H,S04 = 6S0, + 4H,0
C + 2H,S04 = 2S0, 4- CO, + 2H,0
(f) FeSO^ 4- S, = FeS -f 2S0,
{fj) Na,SO, + 2H,S0, = 2NaHS04 -f SO, + H,0
4 Preparation.— (a) By heating moderately concentrated sulphuric acid \\ith
copper turnings; Cu + 2H,S04 = CuSO* + SO, -f 2H,0 . The gas is dried bv
passing through concentrated sulphuric acid. (&) By heating a mixture of
sulphur and cupric oxide in a hard glass tube, (c) In a Kipp's generator by
decomposing cubes composed of three parts calcium sulphite and one part of
calcium sulphate, with dilute sulphuric acid (Neumann, B„ 1887, 20, 1584).
Preparation of sulphites.— The sulphites of the ordinary metals are usually
made by action of sulphurous acid upon the oxides or hydroxides of the metals.
They are normal, except mercurous, which is acid, and chromium, aluminum
§266, G^l. iSlLPHl ROUS ACID. 319
and copper, which are basic. Sulphurous acid precipitates solutions of metals
of the first and second groups, except copper and cadmium.
The sulphites of the alkalis precipitate solutions of the other metals except
chromium salts; and some normal sulphites may be made in this manner.
The sulphites of silver, mercury, copper and ferricum (known only in solution)
are unstable, the sulphurous acid becoming sulphuric at the expense of the
base, which is reduced to a form having a less number of bonds. With the
unstable stannous sulphite the action is the reverse. (See 0.4.) All sulphites
by exposure to the air slowly absorb oxygen, and are partially converted into
sulphates.
5. Solubilities. — One volume of water at 0° dissolves OS.SCl volumes of sul-
phurous anhydride; at 20**, 3G.200 volumes (Carius, A., 1855, 94, 148); or at 20°,
0.104 part by weight (Sims, J, C, 1802, 14, 1). Charcoal absorbs 105 volumes,
camphor 308 volumes, glacial acetic acid 318 volumes of the gas. Liqufd sul-
phurous anhydride dissolves P , S , I , Br and many gases.
The sulphites of the metals of the alkalis are freely soluble in water; the
normal sulphites of all other metals are insoluble, or but very .slightly soluble
in water. The sulphites of the metals of the alkaline earths, and some others,
are soluble in solution of sulphurous acid, the solution being precipitated on
boiling. The alkali bases form acid sulphites (bisulphites), which can be
obtained in the solid state, but evolve sulphurous anhydride. The sulphites
are insoluble in alcohol. They are decomposed by all acids except carbonic
and boric, and in some instances, hydrosulphuric.
6. Beactions. A. — With metals and their compounds. — Sulphurous acid
reacts with Zn , Fe , Sn , and Cu to form hyposulphuroiis acid, H.SO.
(Schutzenberger, C. r., 1869, 69, 19G). With Zn in the presence of HCl
it is reduced to hydrosulphuric acid: 3Zn + 6HC1 + HoSO^ = 3ZnClo +
HjS + 3HoO . Free sulphurous acid precipitates solutions of first and
second group metals except those of copper and cadmium; solutions of
other metallic salts are not precipitated owing to the solubility of the
sulphites in acids.
Alkali sulphites precipitate solutions of all other metallic salts. The
precipitates, mostly white, are soluble in acetic acid. The precipitates
of Pb ^ Hg , Ba , Sr , and Ca are usually accompaijied by sulphates, due to
the fact that soluble sulphites nearly always contain sulphates (4).
Solution of lead acetate precipitates, from solutions of sulphites, lead
sulphiiCy PbSO., , white, easily soluble in dilute nitric acid; and not blacken-
ing when boiled (distinction from thiosulphate). Solution of silver nitrate
gives a white precipitate of silver sulphite, AgoSO^ , easily soluble in very
dilute nitric acid or in excess of alkaline sulphite, and turning dark-
brown when boiled, by formation of metallic silver and s\di)huric acid.
Solution of mercurous nitrate with sodium sulphite gives a gray precipi-
tate of metallic mercury. Solution of mercuric chloride produces no
change in the cold ; but on boiling, the white mercurous chloride is precipi-
tated, with formation of sulphuric acid. Still further digestion, with
sufficient sulphite, reduces the white mercurous chloride to gray metallic
mercury (§68, Cye).
Solution of ferric chloride gives a red solution of ferric sulphite,
Fe2(S03)3 ; or, in more concentrated solutions, a yellowish precipitate of
820 SULPHUROUS ACID, §265, 6B.
basic ferric sulphite, also formed by addition of alcohol to the red solu-
tion. The red solution is decolored on boiling; the acid radical reducing
the basic radical, and forming ferrous sulphate.
Solution of barium chloride gives a white precipitate of barium sul-
phitey BaSOs , easily soluble in dilute hydrochloric acid — distinction from
sulphate, which is imdissolved, and should be filtered out. Now, on adding
to the filtrate nitrohydrochloric acid, a precipitate of barium sulphate
is obtained — evidence that sulphite has been dissolved by the hydrochloric
acid:
BaSO, + 2HC1 = BaCl, -f- H^SO,
BaCl, + H,SO. + CI, + H,0 = BaSO* + 4HC1
Calcinm chloride reacts similar to barium chloride, the precipitate of
calcium sulphite being less soluble in water than the corresponding sul-
phate.
Sulphurous acid and sulphites are active reducing agents by virtue of
their capacity for oxidation to sulphuric acid and sulphates.
The reactions with silver, mercury and ferricum given above illustrate
the reducing action, and the following should also be noted:
PbOo becomes lead sulphate.
As^ forms arsenous and sulphuric acids.
Sbv forms Sb'".
Cu" becomes cuprous sulphate.
Cr^^ forms chromic sulphate.
Co'" forms cobaltous sulphate.
Ni'" forms nickel sulphate.
Mn"+" forms manganous sulphate.
With MnOo in the cold, manganous dithionate, HnSsOe, is formed
(Gmelin's Ha'nd-hool% 2, 174).
With sfannovs chloride sulphurous acid acts as an oxidizing agent, form-
ing stannic sulphide and stannic chloride or stannic chloride and hydro-
sulphuric acid, according to the amount of hydrochloric acid present
(§71, 6e).
B, — With non-metals and their compounds. — Upon other acids sul-
phurous acid acts as a reducing agent, except with hypophosphorous, phos-
phorous, and hydrosulpluiric acids.
i. H3Fe(CN),, forms H.FeCCN)^ and H^SO^ .
2. HNO, and HNO^ form NO and HoSO^ .
3. PH3 + 2H0SO3 = H,PO^ + S2 + 2H2O (Carvazzi, Oazzetta, 1886, 16,
169). HoPO. becomes H3PO4 and the SOj is reduced to S , and with excess
of H3PO. to H,S . H3PO, forms H3PO, and H^S (§263, 6).
4. H,S forms S from both compounds: 4H28 + 2SO2 = SS^ + 4S^0 .
See also 8263 .
§286, 1. SULPHURIC ACID. 321
5. CI , HCIO , and HCIO3 form hydrochloric and sulphuric acids.
6. Br forms hydrobromic and sulphuric acids. HBrOg forms first
bromine then hydrobromic acid, sulphuric acid in both cases.
7. I forms hydriodic and sulphuric acids. In presence of hydrochloric
acid and a barium salt it serves as a means of detecting a sulphite
mixed with a sulphate and a thiosulphate (Smith, C, N., 1895, 72, 39).
HIO3 forms first iodine then hydriodic acid, sulphuric acid in both cases.
7. I^ition. — Acid sulphites heated in sealed tube to 150° are decomposed
into sulphates and sulphur (Barbaglia and Gucci, B., 1880, 13, 2325; Bert helot,
A. (:h„ 1864, (4), 1, 392). Dry SO, at high heat with many metals is decom-
posed, forming a sulphide and sulphate or sulphite (Uhl, R, 1890, 23, 2151).
Sulphites are decomposed by heat into oxides and sulphurous anhydride:
CaSO, = CaO -f SO^; or into sulphates and sulphides: 4Na3SO, = 3NaaS04 -h
8. Detection. — Free sulphurous acid is detected by its odor and by its
decolorizing action upon a solution of EHnO^ or I (Hilger, J. C, 1876,
29, 443). The reaction with iodic acid is also employed as a test for
sulphurous acid (as well as for iodic). A mixture of iodic acid and starch
is turned violet to blue by traces of sulphurous acid or sulphites in vapor
or in solution, the color being destroyed by excess of the sulphurous acid
or the sulphite. Sulphites are distinguished from sulphates by failure to
precipitate with BaCls in presence of HCl . After removal of the BaSOf
by filtration the sulphite is oxidized to sulphate by chlorine water and
precipitated by the excess of BaCL present.
Normal potassium sulphite, EoSO., , is alkaline to litmus but when
treated with BaCL, gives a neutral solution. The acid sulphite, EHSO3 ,
is neutral to litmus but with BaCls gives an acid solution: 2EHSO3 +
BaClo = BaSO., + 2KC1 + SO^ + HoO (Villiers, C, r., 1887, 104, 1177).
9. Estimation. — (a) After converting into H2S04 by HNO, or CI it is precipi-
tated bj' BaClz and weighed as BaS04 . (6) The oxidation is effected by fusing
with NajCO, and KNO, (equal parts), (c) A standard solution of iodine is
added, and the excess of iodine determined by a standard solution of Na^SsO, .
§286. Sulphuric acid. H^SO^ = 98.086 ,
H'2SviO-%,H — — S — — H.
1. Properties. — ^Absolute sulphuric acid, H^SO^ , is a colorless oily liquid
(oil of vitriol) > specific gravity, 1.8371 at 15** (Mendelejeff. B., 1884, 17, 2541).
According to Marignac (A. Ch., 1853, (3), 39, 184), it begins to boil at about
290*', ascending to 338** with partial decomposition. At temperatures much
below the boiling point (160**) it vaporizes from open vessels, giving off heavy,
white, suffocating vapors, exciting coughing without giving premonition by
322 SULPHURIC ACID. §266,2.
odor. At ordinary temperature it is non-volatile and inodorous. At low tem-
peratures it solidifies to a crystalline mass. The freezing point is greatly
influenced by the amount of water present. When the acid contains one mole-
cule of water, H2SO«.HaO , the melting point is highest, -f-7.5* (Pierre and
Puchot, A, Ch., 1874, (5), 1G4).
H2SO4 is a very strong acid and, because of its high boiling point,
displaces all the volatile inorganic acids; on the other hand it is displaced,
when heated above its boiling point, by phosphoric, boric, and silicic acids.
It is a dibasic acid, forming two series of salts, M'HSO^ and H'oSO^ . It is
miscible with water in all proportions with production of heat; it abstracts
water from the air (use in desiccators), and quickly abstracts the elements
of water from many organic compounds, and leaves their carbon, a char-
acteristic charring effect. It dissolves in alcohol, without decomposing it
— but if in sufficient proportion producing ethylsulphuric acid, HCoH^^SO^ .
Sulphuric anhydride, SO, , is a colorless, fibrous or waxy solid, melting at
14.8** (Rebs, A., 188S, 246, 379), boiling at 46'' (Schulz-Sellak, B.. 1870, 3, 215).
and vaporizinpT with heavy white fumes in the air at ordinary temperatures.
It is very deliquescent, and on contact with water combines rapidly, forming
sulphuric acid with generation of much heat.
2. Occurrence. — P'ound free in the spring w^ater of volcanic districts. Found
combined in gypsum. CaSOf -f 2H2O: in heavy, spar, BaSO^: in celestine, SrSO*:
in Epsom salts, MgSOf -\- 7H2O: in Glauber salt, Na^SO^ -f 10H,O , etc.
3. Formation. — (a) Hy electrolyzing HjO , using Ft electrodes with pieces of
S attached (Beoquerel, C. r., 1863^56, 237). (/>) Hv oxidizing S or SO, in presence
of water by 01 . Br , HNO, , etc. (r) By heating S and H^O to 200°. (d) By
adding HjO to SO,, (r) By passing a mixture of SO, and O over platinum
sponge and then adding water.
4. Preparation. — Industrially, sulphuric acid is made by utilizing the
SO2 evolved as a by-])roduct in roasting various sulphides — e, g., iron and
copper pyrites, l)lendcs etc. (a) and (h); or by burning sulphur in the air
to form the SOo . The SOo is passed into a large leaden chamber and
brought into contact with HNO;, , steam, and air. The HNO3 first oxidize?
a portion of the SOo (r); the steam then reacts upon the NOo , forming
HNO3 and NO (d). This NO is at once oxidized again by the air to NO. ,
so that iheoreticalhj no nitric acid is lost, but all is used over again.
Practically, traces of it are constantly escaping with the nitrogen intro-
duced as air, so that a fresh supply of nitric acid is needed to make up for
this loss. The absolute H^SO^ cannot be made by evaporation or distilla-
tion; it still contains about two per cent, of water. It may be made by
adding to water, or to the H.SO^ containing the two per cent of water,
a little more SO^ or H^JQ, than would be needed to make HjSOf ; then
passing perfectly dry air through it until the excess of SO3 is removed,
leaving absolute H0SO4 . Pyrosulphuric, or Xordhausen sulphuric acid.
HoS.Oy, is made by solution of sulphuric acid in sulphuric anhydride (e)\
by drying FeSO^ + '^HjO until it becomes FeSO^ + HjO , and then di-^-
tilling (f). Sulphuric anhydride is made by the action of heat on sodium
§266, 6A. SULPHURIC ACID, 323
pyrosulphate, Na^SzO, (g), prepared by heating NaHSO^ to dull redness; by
distilling pyrosulphurie acid, the anhydride is collected in an ice-cooled
receiver; by heating HjSO^ with PjOg (h):
(a) 2ZnS H- 30, = 2ZiiO + 2S0,
(6) 4reS, + 110, = 2Fe,0, + SSO,
(c) SO, + 2HN0, = H,SO, H- 2N0,
(d) 3N0, H- H,0 = 2HN0. + NO
(e) HaSO, H- SO, = H,S,0,
(f) 4FeS0« H- HaO = 2Fe,0, + H,S,Of + 2S0,
(g) Na,3sOf = Na,SO« + SO,
(h) H,SO« + P,0. = 2HP0, + SO,
Sulphates are made: (a) by dissolving the metals in sulphuric acid;
(i) by dissolving the oxides or hydroxides; (c) by displacement. All salts
containing volatile acids are displaced by sulphuric acid and a sulphate
formed (except the chlorides of mercury). The excess of acid may gener-
ally be expelled by evaporation, or the crystals washed with cold water or
alcohol. The insoluble sulphates are best made by precipitation.
5. Solubilities. — Sulphuric acid is miscible with water in all proportions ;
the concentrated acid with generation of much heat. Sulphuric, acid
transposes the salts of nearly all other acids, forming sulphates, and either
acids (as hydrochloric acid, §269, 4) or the products of their decomposi-
tion ( as with chloric acid, §273, 6). Chlorides of silver, tin, and antimony
are with difficulty transposed by sulphuric acid, and chlorides of mercury
not at all. Also, at temperatures above about 300° phosphoric and silicic
acids (and other acids not volatile at this temperature) transpose sulphates,
with vaporization of sulphuric acid.
The sulphates of Pb , Hg', Ba , Sr , and Ca are insoluble, those of Hg^
and Ca sparingly soluble. Sulphuric acid and soluble sulphates precipi-
tate solutions of the salts of Pb , Hg', Ba , Sr , and Ca ; Hg' and Ca salts
incompletely. The metallic sulphates are insoluble in alcohol which pre-
cipitates them from their moderately concentrated aqueous solutions.
Alcohol added to solutions of the acid sulphates precipitates the normal
sulphates, sulphuric acid remaining in solution: 2EHSO4 = E2SO4 +
H2SO4 . PbSO^ is soluble in a saturated solution of NaCI in the cold,
depositing after some time crystals of PbCIj , complete transposition being
effected. A solution of PbCIj in NaCl is not precipitated on addition of
H2SO4 (Field, J, C, 1872, 25, 575).
6. Beactions. A. — With metals and their compounds. — Sulphuric acid,
dilute, has no action on Pb , Hg , Ag , Cu *, and Bi . Au , Pt , Ir , and Bh
are not attacked by the acid, dilute or concentrated; other metals are
attacked by the hot concentrated acid with evolution of SO, . The f ol-
•Andrews, J. Am. Soe., 1806, 1S« 251.
324 SULPHURIC ACID.
lowing metals: Sn , Th , Cd , Al , Fe , Co , Ni , Hn , Zn , Mg , E , and la
are attacked by the acid of all degrees of concentration ; the dilute and
the cold concentrated, with evolution of hydrogen; the hot concentrated
with evolution of SOo . The degree of concentration and the tempera-
ture may be regulated so that the two gases may be evolved in almost
any desired proportions. A secondary reaction frequently takes place,
the metal decomposing the SOj forming HoS or a sulphide; and the HjS
decomposing the SOg with separation of sulphur (Ditte, A. Ch,, 1896, (6),
19, 68; Muir and Adie, J. C, 1888, 63, 47).
Sulphuric acid or soluble sulphates react with soluble barinm salts to
give barium sulphate, white, insoluble in hydrochloric or nitric acids. This
insolubility is a distinction from all other acids except selenic and fluo-
silicic. The precipitate formed in the cold is ver}' fine and difficult to
separate by filtration; if formed in hot acid solution and then boiled it is
retained by a good filter. In dilute solution for complete precipitation
the mixture should stand for some time. Solutions of lead salts give a
white precipitate of lead sulphate not transposed by acids except H^S (5).
soluble in the fixed alkalis. The presence of alcohol makes the precipi-
tation quantitative (§57, 9). Solution of calcitim salts not too dilute form
a white precipitate of calcium sulphate (§188, 5c).
Dilute sulphuric acid docs not oxidize any of the lower metallic oxide?.
The concentrated acid with the aid of heat effects the following changes:
HgoO forms mercuric sulphate, and sulphurous anhydride is evolved.
SnClg forms, first, sulphurous anhydride, then hydrosulphuric aeiil.
stannic chloride at the same time being produced.
Fe" is changed to Fe2(S04)^ by hot concentrated sulphuric acid.
Mn"+° forms MnS04 and 0. That is, all compounds of mangane^('
having a degree of oxidation above the dyad are reduced to the dyad with
evolution of oxygen.
Potassium permanganate dissolves in cold concentrated stdphuric acid
with formation of a green solution of a sulphate of the heptad manganese.
(Mn03)2S0, (§134, 5c).
Similarly the hot concentrated acid also reduces Pb^ to Pb", Co'" to
Co", Ni'" to Ni", Fe^^ to Fe'", and Cr^i to Cr"', oxygen being liberated
(oxidized) and the metal reduced while the bonds of the SO^ radical are
not changed ; a sulphate of the metal being produced.
B, — With non-metals and their compounds. — When dilute sulphuric acid
transposes the salts of other acids, no other change occurs if the acid set
free be stable under the conditions of its liberation. In ordinary reactions
sulphuric acid never acts as a reducing agent.
L Many organic acids and other organic compounds are decomposed by
the hot concentrated acid, the elements of water being abstracted and
§288, 7. SULPHURIC ACID. 325
carbon set free. Continued heating of the carbon with the hot concen-
trated acid oxidizes it to COg with liberation of SOg .
H2C2O4 becomes CO2 , CO , and H2O . The bonds of the HjSO^ remain
unchanged.
K^Fe(CN)e with dilute HoSO, forms HON : 2K^'Fe{C'S)^ + SHjSO^ =
6HClf + K2FeFe(Clf)e + skjSO, .
Cyanates are decomposed into CO2 and NH3: 2KClfO + 2H2SO4 + 2HoO.
= K2SO, + (NHJsSO, + 2CO2 . . -
Tliiocyanates are also decomposed by concentrated sulphuric acid.
2. Nitrites are decomposed with formation of nitric acid and NO :
6KNO2 + 3H2SO, = SKoSO, + 2HNO3 + 4N0 + 2H2O .
3. H3PO2 or hypophosphites are oxidized to phosphoric acid with re-
duction of the sulphuric acid to sulphurous acid and then to sulphur.
^. Sulphur is slowly changed by hot concentrated sulphuric acid to
sulphurous acid with reduction of the sulphuric acid to the same com-
pound. Hydrosnlphurio acid with hot concentrated sulphuric acid is
oxidized to sulphur with reduction of the sulphuric acid to sulphurous
acid. Further oxidation may take place as indicated above.
5. Chlorates are transposed and then decomposed when treated with
concentrated sulphuric acid : 3ECIO3 + 2H2SO4 = 2KHSO4 + KCIO^ +
2CIO2 + HoO .
6. HBr forms Br and SO2 . No action except in concentrated solution.
7. HI forms I and SO2 .
7. Ignition. — All sulphates fused with a fixed alkali carbonate are
transposed to carbonates (oxide or metal if the carbonate is decomposed
by the heat used, §228, 7) with formation of a fixed alkali sulphate
(method of analysis of insoluble sulphates). If the sulphate, or any other
compound containing sulphur, is fused in the presence of carbon, as
fusion with a fixed alkali carbonate on a piece of charcoal, the resulting
mass contains an alkali sulphide, which, when moistened, blackens metallic
silver.
The sulphates of Cu , Sb , Fe , Hg , Nl and Sn are completely decomposed at
a red heat: 2reS04 = Fe^O, + SO, + SO,; 2CUSO4 = 2CuO -f 2S0, + O, . A
white heat decomposes the sulphates of Al , Cd . Ag , Pb , Mn and Zn . An
ordinary white heat has no action on the sulphates of the alkalis and alkaline
earths; but at the most intense heat procurable the sulphates of Ba, Ca and
Sr are changed to oxides; and at the same temperature K2SO4 and Na2S04 are
completely volatilized, preceded by partial decomposition.
Lead sulphate heated in a current of hydrogen is reduced according to the
following equation: 2PbS04 -f 6H, = Pb -f PbS -f SO, + CH^O . After a
distinct interval the remainder of the sulphur is removed as HjS: PbS -f H, =
Pb -f HjS (Rodwell, J, C, 1863, 16, 42). Potassium sulphate heated in a
current of hydrogen is reduced to potassium acid-sulphide: K3SO4 -f 4H2 =
XOH + KHS -f 3H5O (Berthelot, A. Ch., 1890, (6), 21, 400). Potassium acid-
sulphate, EHSO4 , heated to 200** evolves H2SO4 . The sodium acid-sulphate
^lecomposes mor€ readily.
326 PERSVLPHIKJC ACID. §266,8.
8. Detection. — Free sulphuric acid or the soluble sulphates are detected
by precipitation in hot hydrochloric acid solution with barium chloride^
forming the white, granular, insoluble barium sulphate.
The sulphates insoluble in water are decomposed for analysis — (Ist) by
long boiling with solution of alkali carbonate; and more readily (2d) by
fusion with an alkali carbonate. In both cases there are produced— alkaU
sulphates soluble in water, and carbonates soluble by hydrochloric or nitric
acid, after renu)ving the sulphate (a). If the fusion be done on charcoal,
more or less deoxidation will occur, reducing a part or the whole of the
sulphate to sulphide (7), and the carbonate to metal (as with lead, §67, %
or leaving the metal as a carbonate or oxide (7, §§222 and 228).
a. BaSOf + NajCOs = KajSOf (soluble in water) -f BaCO, (soluble in acid).
A mixture of H3SO4 and a sulphate may be separated by strong alcohoU
which precipitates the latter. A test for free sulpMric acid, in diMinction from
sulphates, may be made by the use of cane sug^r, as follows: A little of the
liquid to be tested is concentrated on the water-bath; then from two to four
drops of it are taken on a piece of porcelain, with a small fragment of white
sugar, and evaporated to dryness by the water-bath. A greenish -W«rik residue
indicates sulphuric acid. (With the same treatment, hydrochloric acid gives a
brownish-black, and nitric acid a yellow-brown residue.) A strip of white
glazed paper, wet with the liquid tested, by immersing it several times at short
intervals, then dried in the oven at 100°, will be colored black, brown or reddish,
if the liquid contains as much as 0.2 per cent of sulphuric acid.
9. Estimation. — («) By precipitation as barium sulphate and weighing as
such. The s«olution should be hot and acidified with hydrochloric acid, and
the mixture should be boiled a few minutes after the addition of the barium
chloride, (b) By precipitotion as barium sulphate with an excess of an hydro-
chloric acid solution of barium chromate (three per cent hydrochloric acid).
Add NH4OH . fill to a definite volume, and filter through a dry filter-paper.
Transfer an aliquot portion to an azotometer with H^O. , and after acidifying,
determine the oxygen evolved (Baumann, Z. nmjeir., 1891, 140) (§244, GAT 12).
(r) When present in small amounts in drinking water by a photometric method
(Hinds, C. i^., 1896, 73, 285 and 299).
§267. Persnlphuric acid. HSO4 = 97.078 .
The anhydride, SoO^ , was discovered by Berthelot (C. r., 1878, 86, 20 and 71).
It is obtained by the action of the silent electric discharge upon a mixture of
equal volumes of dry SOj and O . In solution, the acid is obtained by the
electrolysis of concentrated HjSO^; also by the action of HjO, on concentrated
H,SO« .
At 0° persnlphuric anhydride, ^^0^ , consists of flexible crystalline needles,
remaining stable for several days. The solution in water decomposes rapidly:
more stable when dissolved in concentrated H3SO4 . When heated it decom-
poses into SO3 and O . With SO^ it combines to form SO,: SjOt + SO2 = 3S0, .
Although in its reactions it acts as a strong oxidizing agent, it is weaker than
chlorine or ozone; oxalic acid and chromium salts are not oxidized (Traube, 5.,
1889, 22, 1518, 1528; 1892, 25, 95). Marshall (J. C, 1891, 59, 771) has prepared a
number of salts of persnlphuric acid. The potassium salt, KSO4 , is prepared
by electrolysis of a saturated solution of KHSO4 with a current of 3 to 3.5
amperes. It is a white crystalline powder, which may be recrystallized from
hot water with almost no decomposition. Continued heating of the solution
effects decomposition. The ammonium salt is prepared by electrolysis of a
saturated solution of ammonium sulphate. It is soluble in two parts of water
^268, 3a. CHLORIXE, 327
and can be purified by recrystallization if not heated above 60°. The dry salt
is stable at 100°. With a solution of K2GO, it gives an abundant crystalline
precipitate of KSO4. It is used in the cyanide process for the recovery of gold
(Klbs, Z, anoeic, 1897, 195). The potassium salt, soluble in 50 parts of water at
0°, appears to be the least soluble salt; it gives no precipitate with other metal-
lic salts. Salts of Mn" , Co" and Fe" are oxidized; KI is rapidly decomposed
upon warming; organic dyes are slowly bleached; K4re(CN)« becomes
KiFe(CN),; alcohol is slowly oxidized to "aldehyde, rapidly upon warming.
The barium and lead salts are readily soluble in water (distinction from
H,S0J.
§268. Cliloriiie. CI = 35.45 . Valence one, three, four, five, and seven.
1. Properties.— if oiecK/ar weight, 70.9. Vapor density, 35.8. The molecule con-
tains two atoms, Q, . Under ordinary air pressure it liquifies at — 33.6° and
solidifies at — 102° (Olszewski, if., 1884, 5, 127). Under pressure of six atmos-
pheres it liquefies at 0°. It is a greenish-yellow, suffocating gas, not com-
bustible in oxygen, burns in hydrogen (in sunlight combines explosively),
forming HCl . On cooling an aqueous solution of the gas to 0°, crystals of
Clj.lOHjG separate out (Faraday, Quart. Jour, of Sci., 1823, 15, 71). Chlorine
when passed into a solution of KOH produces, if cold, KCl and KCIO , if hot,
XCl and KCIO,; 2K0H + 01, = KCl + KCIO + H^O; 6K0H + 301, = 5KC1 +
KOlO, 4- 3Hj0 . Passed into an excess of NH4OH , NH4CI and N are formed;
8NH4OH -f 301, = 6NH4CI -f N, -f 8H2O; if chlorine be in excess chloride of
nitrogen is formed; NH4OH + 301, = NCI, -f 3H01 -f H3O . The NCI, is one
of the most dangerous explosives known; hence chlorine should never be passed
into NH4OH or into a solution of ammonium salts without extreme caution.
Chlorine bleaches litmus, indigo and most other organic coloring matter.
The three elements, chlorine, bromine and iodine, resemble each other in
almost all their properties, reactions and combinations, differing (as do their
atomic weights, 35.45, 79.95, 126.85) with a regular progressive variation; so
that their compounds present themselves to us as members of progressive
series. In several particulars fluorine (atomic weight, 19.05) corresponds to the
first member of this series.
Two oxides of chlorine have been isolated; 01,0 , hypochlorous anhydride
(§270), and 010, , chlorine dioxide. The latter is made by the addition of
H3SO4 to KCIO, at 0°. It is a yellowish-green gas, condensing at 0° to a red-
brown liquid. At — 59® it becomes a crystalline solid, resembling KaCr,OT . It
may be preserved in the dark, but becomes explosive in the sunlight.
The most important acids containing chlorine are discussed under the
sections following. They are:
Hydrochloric acid, HCl .
Hypochlorous acid, HCIO .
Chlorous acid, HClOo .
Chloric acid, HCIO., .
Perchloric acid, HCIO4 .
2. Occurrence. — It does not occur free in nature, but its salts are numorui;.-,
the most abundant being NaCl .
3. Formation. — (a) By the action of HCl upon higher oxides as indi-
cated in §289, 64. The usual class-room or laboratory method is illus-
trated by the following equations :
MnO, + 4HC1 = MnCl, -f- CI, + 2H,0
MnO, + 2KaCl + 3H2SO4 = MnS04 -f 2NaHS0, + CI, -f 2H,0
328 CHLORIDE. §868t 36.
(6) By fusing together NH4NOJ and NH4CI : 4NH«K0, -f 2IIH4CI = 5H, +
CI, H- 12H,0 . ((•) By ignition of dry MgCl, in the air: 2MgGl, -f O, = 2]i[gO
4- 2C1, (Dewar, J. Soc, Ind,, 1887, 6, 775). (d) Some chlorides are dissociated
by heat alone: 2AuCl, = 2Au + 3C1, .
4. Preparation.— (a) WeldoiCs pi-ocens'. UnO, is treated with HCl , and the
MnClx formed is precipitated as UnCOH), by adding Ca (OH), . The ]Kn(OH),
is warmed by steam, and air is blown into it, oxidizing it again to KnOs , and
by repeating this process the same manganese is used over again. See Lun ge
and Prett (Z. angeic, 1893, 99) for modification of this method, using HHO, .
(6) Deacon's procrHn: HCl , mixed with air, is passed over fire-bricks moistened
with CuCl, and heated to about 440°. The heat first changes the CuCl, to
CuCl , evolving chlorine; then the oxygen of the air, aided by the HCl, 03d-
dizes the CuCl to CuCl, . It is not certain that the explanation is correct.
It is only known that the hydrochloric acid which is passed into the apparatus
comes out as free chlorine, and that the copper chloride (small in amount)
does not need renewing, (c) Electrolysis now seems likely to supersede other
methods where large amounts are needed.
5. Solubilities. — The maximum solubility of chlorine in water is at 10**.
At 0** one volume of water dissolves 1.5 volumes of chlorine; at 10** three
volumes; at 30° 1.8 volumes (Riegel and Walz, J., 1846; 72). Boiling
completely removes the chlorine from water.
6. Reactions. A. — With metals and their compoondB. — Chlorine is one
of the most powerful oxidizing agents known, becoming always a chloride
or hydrochloric acid. All metals are attacked by moist chlorine, forming
chlorides, many of them combining with vivid incandescence. With per-
fectly dry chlorine many of the metals are not at all attacked. Sn ,
Sb , and As are rapidly attacked, forming liquid chlorides (Cowper, J. C,
1883, 43, 153; Vcley, J. C, 1894, 65, 1). In the presence of acids the
oxidation of the metal takes ])laoe to the same degree as when that metallic
compound is acted upon by HCl (§269, 6A); a chloride is formed having
the same metallic valence that would have resulted from treating the
oxide or hydroxide with hydrochloric acid, e, g., adding HCl to COoO, makes
C0CI2 not C0CI3 , hence adding chlorine to metallic cobalt makes CoCL and
not C0CI3 . In alkaline mixture usually the highest degree of oxidation
possible is attained, as indicated by the following:
1. Pb" becomes PbOj and a chloride in alkaline mixture. With PbClj , it
is claimed that the unstable PbCl^ is formed (Sobrero and Selmi, .4. Ch.^
1850, (3), 29, 1G2; Ditto, A. Ch., 1881, (5), 22, 566).
2. Hg^ becomes Hg" in acid and in alkaline mixture; also HCl or a
chloride.
3. As'" becomes As'^ in acid and in alkaline mixture. Some water must
be present or the reverse action takes place, forming AsClg (§269, 6*42).
Jf, Sb'" becomes Sb^ and a chloride with acids and alkalis.
5. Sn" becomes Sn^'^ and a chloride with acids and alkalis.
6. Mo'^^"" becomes Mo^^ and a chloride with acids and alkalis.
7. Bi'" becomes Bi'^ and a chloride with alkalis only.
8. Cu' becomes Cu" and a chloride with alkalis and with acids.
§268, 6^7. CHLORIXE, 329
9. Cr'" becomes Cr^^ and a chloride in alkaline mixture only.
10. Fe" becomes Fe'" and a chloride with acids and alkalis, but with
alkalis it is also further oxidized to a ferrate.
11. Co" becomes Co(0H)3 and a chloride with alkalis only.
12. Ni" becomes Ni(0H)3 and a chloride with alkalis only.
13. Mn" becomes HnOg and a chloride with alkalis only. See Ditte, I. c,
for formation of HnCl4 .
B. — With non-metals and their compoTmdB.
1. H2C2O4 in acid mixture: H^CjO^ + CI2 = 2CO2 + 2H(a, the HoCjO^
must be in excess and hot (Guyard, Ely 1879, (2), 31, 299); in alkaline
mixture: K^Cfi^ + 4K0H + Clj = 2K2CO3 + 2KC1 + 2H2O .
HCN becomes CNCl and HCl (Bischoff, J9., 1872, 6, 80).
HCNS forms NH3 , H0SO4 , CO2 , and other variable products, and HCl
(Liebig, A., 1844, 50, 337). •
H4Fe(CN)e becomes H3Fe(CW)6 and HCl ; an excess of CI finally decom-
poses the H3Fe(CBr)« .
2. Chlorine does not appear to have any oxidizing action upon the
oxides or acids of nitrpgen.
3. Phosphoms and all lower oxidized forms become H0PO4 with forma-
tion of HCl .
4. Sulphur, and all its lower oxidized forms are oxidized to H2SO4 with
formation of HCl . In an alkaline solution a sulphate and a chloride are
formed. With H2S , S is first deposited, which an excess of CI oxidizes to
H2SO4 . A sulphide in an alkaline mixture is at once oxidized to a sul-
phate without apparent intermediate liberation of sulphur.
5. In alkaline mixture chlorine oxidizes chlorites, and hypochlorites to
chlorates with formation of a chloride: KCIO2 + 2K0H + CL = KCIO3
+ 2KC1 + H2O . With NaOH a hypochlorite is formed if cold, if hot a
chlorate :
2NaOH -f CI, = NaClO + NaCl + H,0
6NaOH + 3C1, = NaClO, -f 5NaCl + 3H,0
6. Chlorine does not oxidize bromine in acid mixture, in alkaline mix-
ture a bromate and a bromide are formed. HBr in acid solution becomes
free hromine, in alkaline mixture a bromate ; hydrochloric acid or a chloride
being formed.
7. Iodine is oxidized to HIO3 in acid mixture, forming HCl ; in an
alkaline mixture a periodate and a chloride are formed. From hydriodic
acid or iodides, iodine is first liberated, followed by further oxidation as
indicated above: 2HI + CI2 = 2HC1 + I2 ; lo + SClg + GH.O = 2HIO3 +
lOHCl ; HI + 8K0H + 4CI2 = KIO, + 8KC1 + 4H2O .
By comparing the oxidizing action of CI with that of Br and I, the
following facts will be observed, and should be carefully considered. The
330 HYDuOf liLORlC ACIP. §988,7.
elements chlorine, bromine, and iodine have an oxidizing power in reverse
order of their atomic weights, chlorine being the strongest. That is, if all
three have the same oxidizing effect, the chlorine acts with the greatest
rapidity; and in some cases, as with cuprous salts, the chlorine oxidizes
while the iodine does not. Their hydracids are reducing agents graded
in the reverse order. If any increase of bonds takes place in presence of
an acid, by chlorine, bromine or iodine, the same increase always occurs in
presence of a fixed alkali. But the oxidation frequently goes further in
presence of a fixed alkali. Thus, with chlorine and potassium hydroxide
we form PbOj, NiCOH), , Bi^O. , Co(0H)3, KjFeO^, and MnO/, which
cannot be formed in presence of an acid.
It is Tery important to remember that those oxides which are formed hy
chlorine, in presence of a fixed alkali, but not in presence of an acid, are the
only ones which can he reduced hy hydrochloric acid. And further^ that this
reduction proceeds not always to the original form, never proceeding beyond
that number of bonds capable of being formed in presence of an acid. Thus,
any lead salt, with potassium hydroxide and chlorine, forms PbOo , and
this treated with hydrochloric acid again forms the lead salt, PbCln . And
ferrous chloride with potassium h^^droxide and chlorine forms K^PeO^ , in
which iron is a true hexad, and KoFeO^ with hydrochloric acid forms, not
the ferrous chloride with which we began, but ferric chloride, for it could
only be oxidized to that point in presence of an acid.
The above is true for bromine and iodine, as well as for chlorine,
7. Ignition. — See 1.
8. Detection. — Free chlorine is recognized by its odor, by its liberation
of iodine from potassium iodide, by its bleaching action upon litmus,
indigo, etc., and by its action as a powerful oxidizing agent (see above).
9. Estimation. — (a) It is added to a solution of potassium iodide and the
liberated iodine determined by standard sodium thiosulphate. (6) It is eon-
verted into a chloride by reducing" agents, and estimated by the usual methods
(S269, 8).
§269. Hydrochloric Acid. HCl = 36.458 .
H'Cl-', H— CI.
1. Properties. — Vapor density, 18.22. At ordinary pressure it liquifies at
—102°, and solidifies at —112.5° (Olszewski, i/., 1884, 5, 127). At 10° under
pressure of 40 atmospheres it condenses to a colorless liquid (Faraday, Tr.,
1845. 155). Critical temperature, 52.3°; nitical pressure, 86 atmospheres ( Dewar.
C A'., 1S85, 51, 27). Dissociated into H and CI at about 1500°, but combines
again upon cooling (Deville, (\ r., 1865, 60, 317). It is a colorless gas, having
an acrid, irritating odor. Headily absorbed by water. The chemically pure
concentrated acid has usually a specific gravity of 1.20, and contains 39.11 per
cent HCl (Lunge and March'lewski, Z. angew,, 1891, 4, 133). The IT. S. P. acid
has a fipcoific fjrariiu of 1.163 at 15° and contains 31.9 per cent HCl . A concen-
trated solution of HCl gives oflf gaseous HCl faster than HjO: a dilute solution
|269, 5. HYDROCHLORIC ACID. 331
^ves oft H2O faster than HCl , as a final result in both cases an acid sp. gr. 1.1
distils unchanged at 110° and contains 20.18 per cent HCl (Bineau, A. C*., 1843,
(3), 7, 257).
2. Occurrence. — Found native only in the vicinity of volcanoes. Found as a
chloride in many minerals, sodium chloride being the most abundant.
3. Formation. — (a) All chlorides except those of mercury are trans-
posed by H2SO4 ; sHver chloride must be heated nearly to the boiling point
of the H2SO4 before the action begins. Lead, antimony and tin chlorides
are slowly transposed.
(6) By the action of sunlight on a mixture of H and CI , or by heating the
mixture to 150**. (c) Platinum black, palladium, charcoal, and some other sub-
stances which rapidly absorb gases will cause the union of the hydrogen and
the chlorine, (cf) When h3'drogen is passed over the heated chlorides of the
most of the metals of the first four groups, the metals are set free and hydro-
chloric acid is formed, (f) Slowly formed by the action of chlorine upon
water in the sunlicrht; rapidlv by its action upon reducing acids such as
BT.CaO, , HH2PO, , H,S, H.SOa .etc.: HH.PO, + 2C\^ -f 2H,0 = H3PO4+ 4HC1 .
Chlorides may be made: {a) By direct union of the elements, mostly
without heat. Whether an ous or ic salt is formed depends upon the
amount of chlorine used, {h) By the action of hydrochloric acid upon the
<?orresponding oxides, hydroxides, carbonates, or sulphites. The solutions
formed may be evaporated to expel excess of acid. If the chlorides thus
formed contain water of crystallization it cannot be removed by heat alone,
for part of the acid is by this means driven off, and a basic salt remains.
If the anhydrous chloride is desired, it may always be made by (a), and
when thus fonned may be sublimed without decomposition, (r) Chlorides
of the first group are best made by precipitation, {d) Metals soluble in
hydrochloric acid evolve hydrogen and form chlorides. In these cases
ouSy and not tV, salts are formed, {e) Many chlorides may be formed by
bringing HgClj in contact with the hot metal.
4. Preparation. — For commercial purposes, made by treating NaCl with
H2SO4 and distilling.
5. Solubilities. — Hydrochloric acid (gas) is very soluble in water as
stated in (1); forming in its solutions of various strengths the hydro-
chloric acid of commerce. Its combinations with metals, forming chlor-
ides, are for the most part soluble in water. AgCl and HgCI are insoluble
in water. PbCIg is only slightly soluble in cold water (§57, be). These
three chlorides constitute the first or silver group of metals, and are pre-
cipitated from their solutions by hydrochloric acid or soluble chlorides
(§61). Solutions of lead salts are not precipitated by mercuric chloride;
green chromic chloride is incompletely precipitated and a sulphuric acid
solution of molybdenum oxychloride not at all by silver nitrate. The chlo-
rides of Sb'", Sn", and Bi require the presence of some free acid to keep them
in solution. AsCl.j , PCI3 , SbCI., , and SnCI^ are liquids at ordinary tem-
332 HYDROCHLORIC ACID. §269, 6A.
perature. The first two are decomposed by water liberating HCl : AiCl^
+ 3H2O = H3A8O3 + 3HC1 . A saturated solution of bismuth nitrate
is precipitated by HCl as the oxychloride (§76, 6/). Hydrochloric acid
increases the solubility of the chlorides of Pb , Hg , Ag , Sb , An , Pt .
Bi and Cn'; it decreases the solubility of Cd , Cu", Co , Ni , Hn , Th , Ba .
Sr , Ca , Hg , An , K and NH^ . Chlorides of Th , Ba , Na , K and NH«
are nearly insoluble in strong HCl (Ditte, C. r., 1881, 92, 242; A. Ch,
1881, (5), 22, 551 ; Berthelot, A. Ch., 1881, (5), 23, 86).
Silver chloride is readily soluble in ammonium hydroxide (separation
from lead and mercurous chlorides) (§69, ^a)'y lead chloride is soluble in
fixed alkali hydroxides (§67, 6a).
HCl dissolves or transposes all insoluble oxalates, carbonates, hypophos-
phites, phosphates, and sulphites. Sulphides of Fe", Mn, and Zn are
dissolved readily; those of Pb , Ag, Sb, Sn, Bi , Cu, Cd, Co, and Hi if
the acid be concentrated; AB2S3 and ASoSg are insoluble in the cold con-
centrated acid, very slowly soluble in the hot concentrated acid; HgS .
red, is insoluble; black, very slowly soluble in the hot concentrated acid.
HgSO^ is only partially transposed by HCl (§68, 6/"), BaSO^ not at all.
The insoluble sulphates of Pb , Hg', Sr , and Ca are slowly but completely
dissolved by the hot concentrated acid. Many of the metallic chlorides
are soluble in alcohol, a few are soluble in ether.
6. Reactions. — A. — With metals and their compounds. — Hydrochloric
acid acts upon the following metals, forming chlorides with evolution of
hydrogen: Pb (slowly but completely), Sn , Cu (very slowly), Cd , Fe , Cr,
Al, Co, Ni, Mn, Zn, and the metals of the fifth and sixth groups:
Ag , Hg , As , Sb , Au , Pt , and Bi are insoluble in HCl (Ditte and Metzner,
A. Ch., 1893, (6), 29, 389).
The following metallic oxides and hydroxides are acted upon by hydro-
chloric acid, forming chlorides of the metal without reduction, water be-
ing the only by-product : Pb" , Ag , Hg , As'" (only with very concentrated
acid), Sb , ^Sn , An'", Pt , Mo^^ Bi'", Cu , Cd , Fe , Al , Cr'", Co", Ni",
Mn", Zn , Ba , Sr , Ca , Mg , K , and Na . The ignited oxides unite with
HCl more slowly than when freshly precipitated or when dried at 100°.
Ignited Ct^O^ is insoluble in HCl ; other ignited oxides, as FCoO., , AloO. ,
etc., require very long continued boiling with the HCl to effect solution.
The following metallic compounds are attacked by hydrochloric acid
with reduction of the metal and evolution of chlorine:
1. Pb"+° becomes PbCL ; no action with a chloride in presence of a
three per cent solution of acetic acid, while bromine is completely set
free from a bromide by PbOg in presence of three per cent of acetic acid
(detection of a chloride in presence of a bromide) (Vortmann, 3f., 1882, 3,
510; 5., 1887, 16, 1106).
§269, 6B6. HYDROCHLORIC ACID, 333
2. As^ becomes AsCl, . (The presence of very concentrated HCl is
required; Fresenius, Z,, 1862, 1, 448; Smith, J. Am, Soc, 1895, 17, 682
and 735.)
: S, Br becomes BiClg .
; 4. Cr^^ becomes CrClg . With KaCroO^ , bromine is completely liberated
from a bromide in presence of 4 cc. of HgSO^ to 100 cc. of water. The
chlorine of a chloride is not liberated, and the bromine may be removed
by boiling. Test the solution for 9, chloride (Dechan, J, C, 1886, 49,
682). Dry HCl does not reduce Cr^^ but combines with it to form the
volatile CrOgClj , chlorochromic anhydride (method of detecting a chloride
in the presence of a bromide).
5' With the exception of ferrates the salts of iron are not reduced by
hydrochloric acid.
6. Co"+° becomes CoCU .
7. lfi"+^ becomes NiCJlj .
8. Mn"+° becomes MnClj . MnOg with small amounts of dilute H2SO4
(1-10) may be used to detect a chloride in presence of an iodide or bromide.
Boiling the mixture removes the iodine first, then the bromine; while the
chlorine is not set free until considerable H2SO4 has been added (Jones,
C, N,, 1883, 48, 296). A mixture of KHSO4 and KMnO^ completely liber-
ates the bromine from a bromide in the cold. A chloride remains unde-
composed until warmed. Aspirate off the bromine, warm and collect the
chlorine (Berglund, Z., 1885, 24, 184).
B, — With non-metals and their componnds.
1. No reducing action with Hfi^O^ , HjCOg , HCN , HCNS , I[^'Ee{CII\ ,.
and H3Fe(CN)e .
2. HNO2 forms chiefly NO and CI. HNO3 forms NOjCl and CI, or
NOCI and CI , or merely NO2 and CI . In case excess of HCl is used the
reaction is: 2HNO3 + 6HC1 = 2N0 + 3CI2 + 4H2O (Koninck and Nihoul,
Z. anorg,y 1890, 477). Dry HCl gas, passed into a cold mixture of con-
centrated H2SO4 and HNO3 , reacts according to the following equations :
2HC1 4- 2HNO3 = 2H2O + 2NO2 -f CI2 (Lunge, Z. angew., 1895, 4, 8,.
and 11).
3. No reducing action with HoS , H2SO3 , or H2SO4 . With thiosulphates
the unstable H2S2O3 is liberated which decomposes as follows : 2Na2S20s +
4HC1 = 4NaCl + S2 + 2SO2 + 2H2O . Sulphates of Ag and Eg" are
completely transposed by HCl , those of Ba , Sr , and Ca not at all, all
others partially (Prescott, C, N,, 1877, 36, 179).
-4. With an excess of HCl , hypophosphites, phosphites, and phosphatea
are dissolved or transposed without reduction.
5, Hypochlorons acid forms chlorine and water: HCIO -}- HCl = H2O +
CI2. Chlorio acid fonns CIO,, Cl^O, and CI in varying proportions^
334 HYDROCHLORIC ACID, §288, 6£^.
but with HCl in excess the following reaction takes place : KCIO, + 6HC1
= KCl + 3CI2 + 3H2O (Koninck and Nihoul, Z, anorg., 1890, 481).
6. EBrOa is decomposed by boiling with HCl, the bromine being set
free: 2KBr03 + 12HC1 = 2KC1 + Br^ + SCl^ + 6H2O (Kaemmerer,
J. pr., 1862, 86, 452).
7. With HIO3, ICI3 ^^^ CI arc formed, no action in dilute solutions:
HlOa + 5HC1 = ICI3 + CI2 + SHjO (Ditte, A.y 1870, 156, 336). According
to Bugarsky (Z. anorg,, 1895, 10, 38'i) YSLJO^ with dilute HjSO^ does not
liberate chlorine from a chloride even on boiling (separation from a
bromide).
7. Ignition. — The chlorides of metals are, generally, more volatile than the
other compounds of the same metals: example, ferric chloride.
Insoluble chlorides are readily transposed by fusion with sodium carbonate:
PbCla + Na^COs = PbO -f 2NaCl -h CO, . If the carbonate be mixed with
charcoal, or if the fusion is done on a piece of charcoal, the metal is also
reduced: 2PbCl, + 2Na.C0, + C = 2Pb + 4NaCl -f 300, .
Heated in a bead of microcosmic salt, previously saturated with copper
oxide in the inner blow-pipe flame, chlorides impart a blue co/or to the outer
flame, due to copper chloride.
Dry sodium sulphate at 150° is transposed by dry HCl (Colson, C. r., 1897.
124, 81). Gaseous HCl transposes potassium and sodium sulphates completely
at a dull-red heat. With the sulphates of the alkaline earths the transposition
is nearly complete (Hensgen, B., 1876, 9, 1671). The silver halides heated with
bismuth sulphide on charcoal before the blow-pipe give distinguishing colored
incrustations: Agl , bright red; AgBr , deep vellow; AgCl , white (Goldschmidt,
C, C, 1870, 297).
8. Detection. — {a) In its soluble compounds, when not in mixtures
with bromides and iodides, hydrochloric acid is readily detected by pre-
cipitation with solution of silver nitrate, as a white curdy precipitate,
opalescence if only a trace be present, turning gray on exposure to the
light
The properties of the precipitate of silver chloride are given in §59, ^r
and Gf. It is of analytical interest in that it is freely soluble in ammonium
hydroxide (considerably more freely than the bromide, and far more freely
than the iodide of silver); soluble in hot, concentrated solution of am-
monium carbonate (which dissolves traces of bromide, and no iodide of
silver); insoluble in nitric acid, temporarily soluble in strong hydrochloric
acid, precipitating again on dilution. It should be observed, that it is
appreciably soluble in solutions of chlorides.
(J) A test for traces of free hydrochloric acid, in distinction from metallic
chlorides, is made by heating the solution with HnOs , without adding an
acid, and distilling into a solution of potassium iodide and starch. Larger
proportions of HCl are more frequently separated by distilling it intact.
(c) Gaseous hydrochloric acid (formed by adding sulphuric acid to dry
chlorides, 3a) is readily detected by the white fumes formed w^hen brought
in contact with ammonia vapor. Also by bringing a stirring rod moist-
§269, 8/. HYDROCELORIC ACID. 336
ened with silver nitrate in contact with the hydrochloric acid gas. Con-
firm by proving the solubility of the white precipitate in ammonium
hydroxide.
(d) The reaction with chromic anhydride is in use as a test for hydro-
chloric acid, more especially in presence of bromides :
{a) 2HC1 -f CrOg = CrO^Cl, (chlorochromic anhydride) -f H,0
(6) 4NaCl + K^CraO, + SHaSO, =
2CtO^CU + 2NaaS04 + K,S04 + 3HaO
To obtain a rapid production of the gas, so that it may be recognized
by its color, the operation may be made as follows: Boil a mixture of
solid potassium dichromate and sulphuric acid, in an evaporating-dish
until bright red, and then add the substance * to be tested, in powder —
obtained, if necessary, by evaporation of the solution. If chlorides are
present, the chromium dioxydichloride rises instantly as a bright brownish-
red gas. The distinction from bromine requires, however, that the mate-
rial, which may be in solution, should be distilledy by means of a tubulated
flask or small retort, the vapors being condensed in a receiver, and neutral-
ized with an alkali (c and d). The chromate formed makes ^ yellow solu-
tion (bromine, a colorless solution). As conclusive evidence of chlorine,
the chromate (acidified with acetic acid), with lead acetate, forms a yellow
precipitate (bromide, a white precipitate, if any):
(c) CrOXla -h 2H,0 =: H^CrO, -f 2HC1
(d) CrO^Cl^ + 4(NH4)OH = (NHJ^CrO, -f 2NH4CI + 2H3O
(e) To detect a chloride in the presence of a cyanide or thiocyanate,
add an excess of silver nitrate, filter and wash. To the moist precipitate
add a few drops of silver nitrate (§318, 2^) and then several cubic centi-
meters of concentrated sulphuric acid and boil for two or three minutes.
The silver cyanide and thiocyanate are completely dissolved with decom-
position, while the silver chloride is not changed except on long continued
boiling. The student should confirm by tests on known material.
According to Borchers (C. N., 1883, 47, 218), to detect a chloride in
the presence of a cyanide or a thiocyanate add silver nitrate, filter, wash,
and boil the precipitate with concentrated nitric acid to complete oxida-
tion of the cyanogen compound. See Mann (Z., 1889, 28, 668) for detec-
tion of a chloride in presence of an alkali thiocyanate by use of CUSO4
and HoS .
(/) If a solution containing iodides, bromides, and chlorides be boiled
with Feo(SO Ja , all the iodine is liberated and may be collected in a
solution of KI and estimated with standard NaoSsOs . The solution should
* With tho chlorides of merctry no brown fumes are obtained as these chlorides are not
transposed by the sulphuric acid ; and the chlorides of lead, silver, antimony, and tin are so
slowly transposed that the formation of the chromium dioxydichloride may escape observation^
Refore relying upon this test the absence of the above named metals should be assured.
386 HYDROCHLORIC ACID. §269,8^.
be cooled to about 60° and a slight excess of E[Mn04 added. The bromine
is all liberated and may be collected in NH^OH and estimated as a bromide
after reduction with SOj . The chloride may now be detected in the
-filtrate and may be estimated by one of the usual methods. Aspiration
aids the removal of the iodine and bromine (Weiss, C. C, 1885, 634 and
712; Hart, C, N., 1884, 60, 268).
{g) Villiers and Fayotte (C. r., 1894, 118, 1152, 1204 and 1413) detect
a chloride in presence of an iodide and bromide by passing the liberated
halogens into a solution of aniline in acetic acid (400 cc. of a saturated
water solution of aniline to 100 cc. of glacial acetic acid) use 3 to 5 cc.
of this solution for each test. Iodine gives no precipitate; bromine gives
a white precipitate; and chlorine a black precipitate. If the bromide be
present in large excess, add silver nitrate, digest the precipitate with
ammonium hydroxide, add hydrogen sulphide and test the filtrate as the
original solution. Liberate the halogen with EMnO^ and H^SO^ .
(70 Deniges (J5Z., 1890, (3), 4, 481; 1891, (3), 6, Q(S) uses HoSO, and
Te'" to liberate the iodine, and KjCrO^ to liberate the bromine; then
after boiling off the I and Br he adds KHnO^ to liberate the chlorine.
The iodine he detects with starch paper, the bromine fumes are absorbed
on a rod moistened with EOH , which then gives an orange-yellow color
with aniline. The chlorine he collects as the bromine and obtains a violet
color with aniline.
({) Dochan (J. C, 1886, 60, 682; 1887, 61, 690) removes iodine of
iodides by distilling with a concentrated solution of KoCrgO^; then the
bromine of bromides by adding dilute HoSO^ and again distilling. The
chloride is precipitated by AgNO., after dilution and addition of HNO^ .
{']) Yortman (.¥., 1882, 3, 510; Z., 1886, 25, 172) detects chlorine in
presence of bromine and iodine as follows: The solution containing the
halogens combined with the alkali or alkaline earth metals is heated with
acetic acid and peroxide of lead until the supernatant liquid is colorless
and has no longer the slightest odor of iodine or bromine; in this way the
whole of the bromine and part of the iodine are driven off, the remainder
of the latter remaining as iodate of lead along with the excess of lead
peroxide. This is filtered off, the precipitate washed with boiling water,
and the chlorine precipitated from the filtrate by addition of silver nitrate.
9. Estimation. — (a) — It is precipitated by AgNO, , washed, and, after igni-
tion, weighed as AgCl. (ft) By a standard solution of AgNO, . A little
NajHPO^ , or, better, KjCraO, , is added to the ehloride to show the end of the
reaction. When enough AgNO, has been added to combine with the chlorine
the next addition gives a yellow precipitate with the phosphate, or a red with
the chromate.
§271, 2. HTP0CHL0R0U8 ACID-CHLOROUS ACID, 337
§270. Hypochlorous acid. HCIO = 52.458 .
H'CrO-", H — — CI.
1. Properties. — Hypochlorous anhydride, Cl^O , is a reddish-yellow gas, con-
densing at about — 20° to a blood-red liquid, which boils at about — 17° (Pelouze,
A, Ch., 1843, (3), 7, 176). Rise of temperature causes decomposition, explo-
sively, into chlorine and oxygen (Balard, A. Ch., 1834, 57, 225). Molecular weight,
86.9. Vapor density, 43.5 at 10°. The acid, HCIO , has not been isolated. Its
aqueous solution smells like CljO , decomposing rapidly, especially in the sun-
light, into CI and HCIO, .
2. Occurrence. — Not found in nature.
3. Fonnation. — (a) By adding chlorine to HgO in the presence of water:
2HgO -f 2CI2 -f HaO = Hg^OCl, -f 2HC10 (Carius, A., 1863, 126, 196). {h) By
adding five per cent nitric acid to calcium hypochlorite and distilling at li
low temperature (Koffer, A., 1875, 177, 314). (c) By passing chlorine into the
sulphates of Mg , Zn , Al , Cu , Ca or Na: Na3S04 + CI, -h H^O = NaHS04 +
NaCl -h HCIO . (d) By heating a mixture of KCIO, and H^CjO^ to 70° (Calvert
and Davies, A, C/*., 1859, (3), 55, 485).
I. Preparation. — For commercial purposes, as a bleaching agent and as a
disinfectant; used as calcium hypochlorite with calcium chloride, chlorinated
lime, made by bringing chlorine in contact with calcium hydroxide, without
heating. Lunge and Schoch (^., 1887, 20, 1474) give the formula Ca~cf^
to chlorinated lime. See also Kraut (i.., 1882, 214, 244), Also as sodium
hypochlorite, made by treating sodium hydroxide with chlorine short of satu-
ration in the cold: 2NaOH + CI, = NaClO -f NaCl -f H^O . The sodium
hypochlorite-and-chloride — mixed as formed by chlorine in solution of sodium
hydroxide or sodium carbonate, or by double decomposition between solution
of the calcium hypochlorite-and-chloride and solution of sodium carbonate — is
pharniacopoeial, under the name of solution of chlorinated soda (NaCl.NaClO).
5. Solubilities. — Hypochlorites are all soluble in water and are decomposed
by heating.
6. Beactions. — The hypochlorites are all unstable. They are decomposed by
nearly all acids, including CO^: 2Ca(C10)2 + 2C0j = 2CaCOs -f 2C1, + Ol\
4NaC10 -f 4HC1 = 4NaCl + 2H,0 -h 201^ -f O^ . They are very powerful
oxidizing agents, acting in acid solution as free chlorine, as the above equa-
tions indicate. Hypochlorites act as chlorine in alkaline mixture (§268, 6)
(Fresenius, Z. anyeic, 1895, 501).
7. Ignition.— All hypochlorites are decomposed by heat: 2KC10 = 2KC1 -f O2 .
8. Detection. — Although silver hypochlorite is soluble in water, it decora-
poses very quickly, so that on adding silver nitrate to sodium hypochlorite
the final reaction is as follows: 3NaC10 -h 3AgN0, = 2AgCl -f- AgClO, -+-
?.NaNO, . When KCIO is shaken with Hg° , yellowish-red Hg^OCl, is formed;
the other potassium salts of chlorine, /. r., KCl , KClOj , KCIO, and KCIO^ ,
have no action upon Hg° . An indigo solution is decolored by hypochlorites,
"While KMnOf is not decolored. If arsenous acid be present, the indigo solution
is not decolored until the arsenous acid is all oxidized to arsenic acid.
9. Estimation.— It is estimated as AgCl after reduction with Zn and H,SO« .
Rosenbaum (Z. angctc.j 1893, 80) gives a method for estimating the various
chlorine compounds in chlorinated lime.
§271. CUorous acid. HClOg = 68.458 .
H'Cl'"0-"2 , H — — CI = .
1. Properties. — The anhydride, CI3O3 , has not been isolated and the free acid
is known only in solution, and this generally contains some HCIO, . It has an
intense yellow color and is very unstable.
2. Occurrence. — Neither the acid nor its salts are found in nature.
338 CHLORINE PEROXIDE. §871,3.
3. Formation. — An impure chlorous acid is said to be formed when XCIO, is
treated with HNO, and A8,0, , Cj,H„0,, or C.H, (Millon, A, CK 1843, (3), 7.
298; Schiel; A., 1859, 109, 318; Carius, A,, 1866, 140, 317). Chlorites of a number
of metals have been made by adding the bases to a water solution of the acid;
also from KClOa by transposition.
4. Preparation. — EClO, is prepared by adding an aqueous solution of CIO5 of
known strength to the proper quantity of KOH , and evaporating in a vacuum.
The crystals of EClO, which are formed in the reaction are removed and the
mother liquor is crystallized from alcohol.
5. Solubilities. — All chlorites which have been prepared are soluble in water,
lead and silver chlorites sparingly soluble.
6. Beactions. — Chlorouc acid or potassium chlorite in dilute acid solution is
a powerful oxidizing agent, acting similar to chlorine.
7. Ignition. — Chlorites when heated evolve oxygen and leave a chloride, or
first a chloride and a chlorate (Brandau, A., 1869, 151, 340).
8. Detection. — A concentrated solution of a chlorite gives a white precipitate
with silver nitrate, fairly readily soluble in more water. XlCn04 is decolored,
a brown precipitate being formed. A solution of indigo is decolored even in
presence of arsenous acid (distinction from hypochlorous acid). Chlorites
when slightly acidulated give a transient ameth^'st tint to a solution of ferrous
sulphate.
9. Sitimation. — By reduction to chloride and estimation as such. By meas-
uring the amount of ferrous iron oxidized to the ferric condition: 4FeS0« -h
HCIO, -f 2H,S04 = 2re,(S0*), + HCl + 2H,0 .
§272. Chlorine Peroxide. ClOj = 67.45 .
W^O-'o, ^Z. CI — — C1 = or Oi=:Cl = 0*.
Chlorine peroxide, CIO2 , at ordinary temperature, is a dark greenish-yellow
gas. In concentrated solution it has very much the odor of nitrous" acid.
Cooled in a mixture of ice and salt it condenses to a bromine-red liquid; and
in a mixture of solid CO3 and ether it forms a mass of orange-yellow, brittle
crystals. When warmed to about 60° it explodes with violence. In direct
sunlight at ordinary temperature it decomposes slowly into chlorine and
oxygen, while in the dark it is quite stable. In contact with many substances,
as phosphorus, sulphur, sugar, ether, turpentine, etc., it explodes at ordinary
temperature. In moist condition it bleaches blue litmus-paper without pre-
viously reddening it.
One volume of water absorbs about 20 volumes of the gas at 4° (Millon.
A. Ch,, 1843, (3), 7, 29H). The solution in water contains HCIO. and HCIO, .
It is prepared by carefully adding ECIO3 to cold concentrated H2SO4: the
mixture is then carefully warmed to 20°, later somewhat higher. The gas is con-
densed in a tube cooled by a mixture of ice and salt: 3EC10, -f- 2H.S0, =
2KHS0« -f ECIO4 + H-O -f 2CIO2 (Millon, I.e.). It is also made by warming
a mixture of oxalic acid and potassium chlorate. When prepared in this man-
ner it is mixed with CO^: 2KC10, -f 21B^,C^0, = K,C,0, + 2H3O -f 2C10, +
2CO3 (Calvert and Davies, A., IsriO, 110, 344). It is also formed, mixed with
chlorine, when ECIO, is warmed with HCl . HI is oxidized to I; SO, to H.SO. .
Indigo is bleached even in presence of AsjO, .
•Pebal, A., 1875, 177,1.
§273, 6A. CHLORIC ACID. 339
§273. Chloric acid. HCIO3 = 84.458 .
H'CrO-% . H — — CI ^ J
1. Properties. — A solution of chloric acid may be evaporated in a vacuum
until its specific gravity is 1.282 at 14°. The composition is then HClOg.THjO ,
containing 40.1 per cent HCIO, (Kaemmerer, Fogg,, 18G9, 138, 390). Farther
attempts at concentration result in evolution of chlorine and oxygen, forming
HC10«: 8HC10. = 4HCIO4 + 2HaO + 30, + 2C1, (SeruUas, A, Ch., 1830, 45, 270).
Its solution in the cold is odorless (ind colorless; first reddening and then
bleaching litmus. It is a strong oxidizing agent, paper soaked with the acid
takes fire on drying. The anhydride, ClaOa , has not been isolated.
2. Occurrence. — Does not occur in nature.
3. Formation. — The free acid may be formed by adding an excess of HaSiF^
to a hot solution of EClOg; the filtrate is evaporated in vacuo, the excess of
HaSiF« volatilizes, leaving the HClOg . Many chlorates are formed by treating
the metallic hydroxides with the free acid. Also by the action of Ba(C10s)i
upon the sulphate of the metal whose chlorate is required; or by the action
of the chloride of the chlorate needed, upon a solution of AgClO, .
4. Preparation. — By adding H2S04 in molecular proportions to a solution of
BaCClOj), . Chlorates of the fifth and sixth group metals are prepared by
passing chlorine into the respective hydroxides dissolved or suspended in water.
By repeated crystallization the chlorate is separated from the chloride which
is also formed: 6K0H -f 3CI2 = 5KC1 -f EClO, -f 3HaO .
5. Solubilities. — All chlorates are soluble in water, the chlorates of
Hg , Sn , and Bi require a little free acid. Mercurous and ferrous chlorates
are very unstable. Potassium chlorate is the least soluble of the stable
metallic chlorates; soluble in about 21 parts water at 10° (Blarez, C. r.,
1891, 112, 1213).
G. Reactions. .-1. — With metals and their compounds. — Chloric acid
attacks Mg evolving hydrogen and forming a chlorate only. With Zn ,
Fe , Sn , and Cu some chloride is also formed. With Zn and HgSO^ the
reduction to chloride is complete, and with sodhim amalgam no reduction
whatever (Thorpe, J. C, 1873, 26, 541). With the zinc-copper couple *
the reduction to a chloride is rapid and complete. The hot concentrated
acid attacks all metals. With oxides or hydroxides the acid forms chlor-
ates provided a chlorate of that metal can by any means be formed. Free
chloric acid is a strong oxidizing agent, and if an excess of the reducing
agent is used, it is converted into hydrochloric acid, or a chloride. With
the aid of heat the chloric acid splits up, forming some chlorine and
oxides of chlorine.
Hg' forms Hg".
As'" forms As^.
Sb'" forms Sb^.
Sn" forms Sn^.
Cu' forms Cu". ,
• Gladstone and Tribe's copper-zino couple is prepared by treating thin zinc foil with a 1 per
cent solution of copper sulphate until the zinc is covered with a black deposit of reduced cop-
per. When washed and dried it is ready for use.
340 CHLORIC ACID, §278,68.
Cr'" forms Cr^*^, chromic salts are readily oxidized to chromic acid on
boiling with KCIO., and HNO3 .
Fe" forms Fe'" (a distinction from perchloric acid) (Carnot, C. r., 189C,
122, 45-2).
Mn" fonns Mn'^\ man,<ranous salts arc rapidly oxidized to MnOj on warm-
ing with KCIO3 and HNO,. .
Salts of lead, cobalt, and nickel do not appear to be oxidized on boiling
with KCIO3 and HNO3 .
B. — ^With non-metals and their compounds.
1, H2C2O4 forms COo and varying proportions of CI and HCl . Heat
and excess of oxalic acid favors the production of HCl (Guyard, B/., 18T9,
(2), 31, 299). All oxalates are decomposed, CO2 and a chlorate or chloride
of the metal being formed. Carbonates are all transposed.
HCNS forms H^SO^ , HCN , and HCl .
H4Fe(CN)« first forms B.^Yt{C'S)^ and HCl ; a great excess of HCIO,
decomposes the HgFeCCN),, .
2. HNOo forms HNO3 and CI . Nitrites are transposed and oxidized,
forming chlorates or nitrates of the metal.
S, PH3 , HH.POo , and H.,P03 form H3PO4 and HCl . Hvpophosphite>
and phosphites are transposed and then oxidized, H3PO4 and a chlorate or
a chloride of the metal being produced.
Jf. S^'^~° forms S^'^ and HCl ; that is, the sulphur of all compound^
becomes HoSO^ with formation of HCl. All sulphides, sulphites, thin-
sulphates, etc., are transposed, forming a chlorate, chloride, or sulphate
of the metal.
5. HCl in excess forms only CI and HoO (§269, (jBo). NaCl warmed with
HCIO3 evolves CI , leaving only NaClOo .
0. HBr forms Br and HCl . KBr warmed with HCIO3 evolves Br , leav-
ing only KCIO3 .
7. I and HI form HIO3 and HCl . Soluble iodides form iodic acid or
an iodate.
7. Ignition. — All chlorates are resolved by heat into chlorides and
oxygen : 2KCIO3 = 2KC1 + ^Oj . Some perchlorate is usually formed as
an intermediate product: 2KCIO3 = KCIO^ + KCl + O2 (Serullas, .4. C/i..
1830, (2), 45, 270). In presence of various metallic oxides, etc., tho
oxygen is separated more easily, the metallic oxides remaining unchanged.
With manganese dioxide, the oxygen of potassium chlorate is obtained at
about 200°; ferric oxide, platinum black, copper oxide, and lead dioxide
may be used (§242, 3). If chlorates are rapidly ignited some chlorine i>
given off (Spring and Prost, /?/., 1889, (3), 1, 340). When friturafed or
heated with combustible substances, charcoal, organic substances, sulphur,
sulphites, cyanides, thiosulphates, hypophosphites, reduced iron, etc. —
§274, 1. PERCHLORIC ACID. 3-1:1
chlorates violently explode, owing to their sudden decomposition, and the
simultaneous oxidation of the combustible material. This explosion is
more violent than with corresponding mixtures of nitrates.
Alkali chlorates when fused with an alkali, or an alkali carbonate, and
a free metal or a lower oxide, or salt of the metal, generally oxidizes it to
a higher oxide, or to a salt having an increased number of bonds; and
the chlorate is reduced to a chloride — e. g,, Mn^^""° becomes Mn^ . That
is, any compound of manganese having less than six bonds is oxidized to
the hexad (a), Cr'" becomes Qt^ (h), As^-" becomes As^ (c), Pb^-°
becomes Pb^ (d). Co'"-° becomes Co'" (e). C^-° becomes C^^ (yr^^ pv-n
becomes P^ (g). F-° becomes F (h). S^-"" becomes S^^ (t).
(a) 3Mn,0, + 18K0H + 5KC10. = OK^MnO^ + 5KC1 + OH-O
(6) 2CrCl, + lONaOH + NaClO, = 2Na,Cr04 + 7NaCl + 5H,0
(r) 3AS4 + 36EOH + lOKClO, = 12KsA80« + lOKCl + ISHjO
id) 3Pb,04 + NaaCO, + 2NaC10, = 9PbOj + 2NaCl + Na,CO,
(e) 6C0CI, + 12K0H + KCIO3 = 300,0, + 13K01 + 6H,0
(f) 3K,04H40, 4- 5K010. = 5K01 + 3K,00, + 9C0, + 6H,0
(g) 3Pb(H,PO,)2 + 18K0H + 5X010^ = r^PbO^ + 6K,P0, -f 5K01 + 15H,0
(h) Znl, + KaOO, + 2K010, = ZnO +. 2KI0, + 2K01 + 00,
(i) BK.SsO, 4- 12K,00. + lOKOlO, = 15E,S0« + lOKOl + 1200,
8. Detection, yu) Dry chlorates when warmed with concentrated sul-
phuric acid, detonate evolving yellow fumes : 3ECIO3 + 2H2SO4 = 2KHSO4
H- KCIO4 + 2CIO2 + H2O . This action is modified by reducing agents;
some acting rapidly, increase the detonation; others acting slowly, lessen
it. (h) HCIO3 , like HNO3 , decolors indigo solution and gives colors with
brucine, diphenylamine, paratoluidine, and phenol similar to those formed
by HNO3. (c) By ignition a chloride is left: 2EC108 = 2KC1 + 30^ .
(d) It is changed to a chloride by nascent hydrogen: 2ECIO3 + 6Zn +
THjSO^ = eZnSO^ + K2SO4 + 2HC1 + 6H2O; or by reducing acids or
bases: 2KCIO3 + H2S0/+ 6H2SO3 = KgSO^ + 6H2S0^ + 2HC1 . The
resulting HCl is then identified in the usual manner. Chlorides, if origin-
ally present, should first be removed by silver nitrate.
9. Estimation. — (a) Reduction to a chloride and estimation as such, (h) Addi-
tion of HOI and KL and estimation of the liberated iodine with standard
17a,S,0, .
§274. Perchloric acid. HCIO4 = 100.458 .
=
H'Cl^^O-", , H — — CI =
1. Properties.— fifpec//?c gravity, 1.782 at 15**. The anhydrous H010« is a color-
less oily liquid, volatile but cannot be distilled without partial decomposition,
often with explosive violence. Only its solution in water can be safely handled.
Paper, charcoal, ether, phosphorus, and many other substances when brought
342 BROMINE, §274^2.
in contact with the anhydrous acid take fire. The dilute acid is very stable, not
being easily reduced (Berthelot, A, Ch,, 1882, (5), 27, 214). It does not bleach,
but merely reddens blue litmus paper.
2. Occurrence. — Not found in nature.
3. Formation. — (a) By electrolysis of a solution of CI or HCl in water
(Riche, C. r., 1858, 46, 348). {b) XCIO4 is formed by electrolysis of XGIO. ,
using platinum electrodes (Lidoff and Tichomiroff, J, C, 1883, 44, 149). {e)
KCIO, is heated with an excess of HsSiF« , after cooling and filtering, the
filtrate is carefully distilled (Roscoe, J. C, 1863, 16, 82; A., 1862, 121, 346).
((/) By treating the sulphate of the metal, the perchlorate of which is desired,
with Ba(C10«)a in molecular proportions, (e) By treating the chloride of the
metal, the perchlorate of which is desired, with AgC10« in molecular propor-
tions.
4. Preparation. — ECIO4 is made by carefully heating KCIO, until no more
oxygen is evolved: 2EC10, = KCl + KCIO4 + O, (7). The residue is dissolved
in water and upon cooling crj'^stals of KCIO4 separate. The free acid, nearly
pure, is obtained by cautiously distilling ECIO4 with concentrated H2SO4 .
5. SolubilitieB. — All of the perchlorates of the ordinary metals are soluble
in water, and all are deliquescent except NH.CIO4 ECIO4 , PbCClO*), and
HgC104 (Serullas, A. Ch,, 1831, 46, 362). Potassium perchlorate is soluble ia
142.9 parts of water at 0**, in 52.5 parts at 25*'. and in 5 parts at 100** (Muir,
C, N., 1876, 33, 15). KCIO4 is insoluble in alcohol (distinction from NaC104)
(Schloessing, A. Ch„ 1877, (5), 11, 561).
6. Beactions. — Iron and zinc evolve hydrogen when treated with perchloric
acid. The acid reacts with the hydroxides of many metals to form per-
chlorates. It is not reduced by HCl , HNO. , HjS or SO, . Iodine is oxidized
to HIO4 with liberation of chlorine: Ij -f 2HCIO4 = 2HIO4 -H CI, . A solution
of indigo is not decolored by HCIO4 even after the addition of HCl (distinction
from all other oxj'acids of chlorine). It is not reduced by the zinc-copper
couple (distinction from chlorate). Sodium perchlorate, NaClOf , is used as a
reagent to precipitate potassiunr salts.
7. Ignition. — Perchlorates strongly ignited evolve oxygen and leave a chloride
(§242, 3).
8. Detection. — In presence of a hypochlorite, chlorite, chlorate and chloride
boil thoroughly with HCl; the first three are decomposed, leaving chloride and
perchlorate. Remove the chloride with AgNO, and fuse the evaporated filtrate
with NaoCO, . Dissolve the fused mass in water and test for a chloride; its
presence indicates the previous presence of a perchlorate.
9. Estimation. — (a) After being changed to a chloride as indicated above, it
is estimated in the usual manner, {h) It is fused with zinc chloride and the
amount of chlorine liberated measured by the amount of iodine set free from a
solution of potassium iodide (separation from chlorate, chlorides and nitrates).
(c) ECIO4 is heated to 200** with HPO, and KI; the iodine liberated showing
the amount of perchlorate present (Gooch and Kreider, Am. 8.<, 1894, 48, 33; and
1895, 49, 287).
§275. Bromine. Br = 79.95 . Valence one and five.
1. Properties— Mo7€Pt/7ar weight, 159.90; vapor density, 80; specific gravity, 3.18S2S
at 0°; hoiling point, 59.27*^ (Thorpe, J. C, 1880, 37, 172). At —7.2' it becomes a
brown solid (Philipps, B., 1879, 12, 1421). At ordinary temperatures bromine
is a brown-red, intensely caustic liquid, freely evolving brown vapors, corro-
sive vapors of a suffocating chlorine-like odor. As a solid it is still darker in
color. It reacts with KOH in all respects similar to chlorine (§268, 1). Indigo,
litmus and most other organic coloring matters are bleached. A solution of
starch is colored slightly yellow.
Bromine decomposes hydrosulphuric acid with separation of sulphur, and
subsequent production of sulphuric acid; changes ferrous to ferric salts, and
(in presence of water) acts as a strong oxidizing agent. It displaces iodine
from iodides, and is displaced from bromides by chlorine; its character being
intermediate between that of chlorine and that of iodine.
§275, 6.4, 11. BROMIXE, 343
No oxides of bromine . have, with certainty, been iso lated. The well-estab-
lished acids are: Hydrobromic, HBr; hypobromous, HBrO; bromic, HBrO, .
2. Occurrence. — Not found free in nature. As a bromide in sea water, mother
liquor from salt wells, mineral springs, and in a few minerals.
3. Formation. — (a) Hydrobromic acid or any soluble bromide is warmed with
MnOa and H3SO4 . (b) Any soluble bromide is treated with chlorine water
and the solution warmtd.
4. Preparation. — The bromine of commerce is obtained chiefly from the
mother liquor of the salt works: (a) By treating with MnO, and H3SO4: Mg^Br,
4- MnO, -h 2H,S04 = MgSO« -f MnSO^ -f Br, -f 2H2O . (b) By leading a
current of steam and chlorine into the bottom of a vessel filled with coke,
into which a stream of the mother liquor flows from above: MgBrj -f CI3 =
MgCla + Bra . (c) By adding to the mother liquor a mixture of Mg(OH)a ,
suspended in water and saturated with chlorine, rendering acid and distilling
in a current of steam: MgCClO.), -f 6MgBr, -f 12HC1 = TMgCU -f GHjO +
6Br2 . (d) By electrolysis of the mother liquor at a low temperature and then
distilling in a current of steam.
Commercial bromine is freed from chlorine by adding KBr and distilling. If
iodine be present it is first removed as Cul .
5. Solubilities. — Bromine dissolves in 30 parts of water at 15**, forming an
orange-j'ellow solution (Dancer J, C, 1862, 15, 477). Its water solution is
permanent, but slowly decomposes: 2Brj 4- 2H3O = 4HBr -f O. . Much more
soluble in HCl , HBr , KBr , BaCL , SrCl, , and in many other salts than in
water. Soluble in carbon disulphide, chloroform, ether and alcohol. Readily
removed from its solution in water by shaking with carbon disulphide or
chloroform, imparting a brown color to the solvent.
6. Beactions. A. — With metals and their compounds. — Bromine unites
directly with gold, platinnm, and all ordinary metals to form bromides.
Silver salts are precipitated, yellow- white, as bromide and bromate:
6AgN0s + 3Br2 + SH^O = 5AgBr + AgBrOg + 6HN0^ . In the follow-
ing metallic compounds the valence of the metal is changed; the bromine
being reduced to HBr or, if in alkaline mixture, to a bromide. The reac-
tion is less violent than with chlorine.
I. Pb" becomes PbOj in alkaline mixture only.
£. Hg' becomes Hg" in acid and in alkaline mixture.
S, As'" becomes As^ in acid and in alkaline mixture. With AsH^ and
a solution of bromine in water H3A8O3 is first formed, and if the bromine
be in excess the final products are H3A8O4 and HBr .
4. Sb'" becomes Sb'^ in acid and in alkaline mixture.
5. Sn" becomes Sn^'^ in acid and in alkaline mixture.
6. Bi'" becomes Bi205 in alkaline mixture only.
7. Cu' becomes Cu" in acid and alkaline mixture.
8. Cr"' becomes Cr^ in alkaline mixture only.
9. Fe" becomes Fe'" in acid mixture; in alkaline mixture the iron is
further oxidized to a ferrate, HBr or a bromide being formed.
10. Co" becomes Co'" in alkaline mixture only.
II. Ni" becomes Ni'" in alkaline mixture only (Bolpius, J. C, 1876,
29, 742).
344 BROMINE. §275, 6^, U.
12. Mn'^~° becomes Mn^^ in alkaline mixture only.
B. — With non-metals and their compounds.
1. H2C2O4 becomes a carbonate and a bromide in alkaline mixture. An
excess of hot saturated oxalic solution changes Br to HBr .
HCNS fonns, among other products, H.^S04 ^^^ a bromide in acid mix-
ture, and a sulphate and a bromide in alkaline mixture.
H4Fe(CN)g in acid mixture forms lL^t{GS)^ and HBr , in alkaline mix-
ture a ferricyanide and a bromide (Wagner, J. C, 1876, 29, 741).
2. HNO2 becomes HNO3 and HBr if dilute and cold.
3. PH3 , HH2PO2 and H3PO8 become H3PO4 and HBr with acids, and a
phosphate and a bromide in alkaline mixture. P and Br unite to form
PBrg or PBrj , depending upon relative amounts of the elements present.
The phosphorus bromides are decomposed by water, forming HBr and
the corresponding acids of phosphorus.
4. S°, HjS , H2SO3 , H0S2O3 , S^^-" becomes H2SO4 and HBr with acids,
a sulphate and a bromide in alkaline mixture.
5. Br does not act as an oxidizing agent upon the compounds of chlorine,
but may, at low temperatures, combine with chlorine to form a chlorine
bromide, BrCl (Bornemann, A., 1877, 189, 183).
6. In alkaline mixture hypobromites by boiling are oxidized to bromates
with formation of a bromide.
7. Iodine becomes an iodate and a bromide in alkaline mixture; the
elements may combine to form the unstable bromiodide, IBr (Bornemann,
/. c). HI and iodides form I and HBr , but in alkaline mixture an iodate
and a bromide are produced.
7. Ignition. — WarminpT drives off all the bromine from its solutions in water
or other solvents. Heat favors all reactions with bromine.
8. Detection. — Bromine is usually detected by shaking its solution in
water with CSo , which dissolves it with a reddish-yellow color; if present
in large quantities the color is brown to brownish black. In this case
a large excess of CSo must be used or a very small portion of the unknown
taken, in order that the solution be dilute enough for the reddish-yellow
bromine color to be distinguished from the violet color of iodine.
Ether or chloroform may be used instead of carbon disulphide, but the
solution is of a paler yellow. Starch solution gives a yellow color with
bromine, but the reaction is less delicate than vnih CSo .
9. Estimation. — (a) The bromine is made to act npon KI , and the iodine
which is liberated is estimated by standard solution of Na-SjO, . (b) It is
estimated by the amount of As.O, which it oxidizes in alkaline solution, (r) It
is converted into HBr by HjS or H.SO^ , and then precipitated by AgNO, .
and weighed as AgBr .
§276, 6A. HYDROBROMIC ACID. 345
§276. Hydrobromic acid. HBr= 80.958.
H'Bi^', H — Br.
1. Properties.— if o/ccw/ar weight, 149.9. Vapor density, 39.1. A colorless gas,
condenses to a liquid at — 69" and solidities at — 73° (Faraday, A,, 1845, 56, 155),
Its aqueous solution is colorless and is not decomposed by exposure to the
air. The specific gravity of the saturated solution at 0** is 1.78; containing 82.02
per cent HBr, or very nearly HBr.HxO . If a saturated solution is boiled,
chiefly HBr is given off, and if a dilute solution is boiled, chiefly HjO is given
off, until in both cases the remaining liquid contains 47.38 to 47.86 per cent
of HBr , its sp. gr. 1.485, its boiling point constant at 126**, and its composition
almost exactly HBr.5H20 , which distils over unchanged. Its vapor density
of 14.1 agrees with the calculated vapor density pf HBr.5HsO .
2. Occurrence. — Not found free in nature, in combination as bromides in sea
water and in some minerals.
3. Forxiiation. — («) By action of bromine upon phosphorus immersed in
water, the amorphous phosphorus is preferred: P* + lOBr, -|- 16HjO = 4H,P04
-f 20HBr . (6) By action of H3PO4 or H2S0« on KBr (Bertrand, J, C, 1876, 29,
877). (c) By transposition of BaBrj by cold dilute HjSO^ added in molecular
proportions, (d) By passing a mixture of Br and H over platinum sponge.
{€) By action of Br on HsPO, . (f) By adding Br to Na^SOs .
Metallic bromides are formed: (i) By direct union of the elements, but in a
few cases heat is required to effect the combination. (2) By action of ECBr
upon the metallic oxides, hydroxides and carbonates. (J) Many bromides are
formed by action of HBr on the free metal, ous salts and not \c being formed.
(^) Bromides of the first group are best made by precipitation, (o) Bromides
of E , Na , Ba , Sr and Ca are made by the action of bromine on their hydrox*
ides and subsequent fusion:
6E0H + 3Br, = EBrOs + 5KBr -f 3H,0
2KBrO, (ignited) = 2KBr -f 30,
4. Preparation. — (a) HaS is added to a solution of bromine in water until
the yellow color disappears; the solution is then distilled. The first portion
of the distillate is rejected if it contains HjS, and the latter portion if it con-
tains H2SO4 (Recoura, C. r., 1890, 110, 784). (6) HsS04 is added to a concen-
trated solution of KBr; after twenty-four hours the greater portion of the
KHSO4 has crystallized out. The remaining liquor is then distilled. The
product usually contains traces of HsSOa. (c) By passing bromine into hot
paraffine (Crismer, B., 1884, 17, 649).
5. Solubilities. — Silver and mercurous bromide are insoluble in watcr^
lead bromide is sparingly soluble; all other bromides are soluble. Hydro-
bromic acid and soluble bromides precipitate solutions of the metals of
the first group, lead salts incompletely. Lead bromide is less soluble than
the corresponding chloride. The presence of soluble bromides increases
the solubility of lead bromide. A small amount of hydrobromic acid
decreases its solubility, but a larger excess increases it (Ditte, C. r.. 1881,
92, 718).
In akoholy the alkali bromides are sparingly or slightly soluble : calcium
bromide, soluble; mercuric bromide, soluble; mercurous bromide, insolu-
ble. Silver bromide is soluble in NH^OH .
6. Reactions. — A, — With metals and their compoTinds. — Hydrobromic
acid dissolves many metals with the formation of bromides and evolution
of hydrogen, e. ^., Pb , Sn , Fe , Al , Co , Ni , Zn , and the metals of the
84:6 HYDROBROMIC ACID. §276, 64, L
calcium and the alkali groups. It unites with salt forming oxides and
hydroxides to produce bromides without change of valence: PbO + 2EBr
= FbBrj + HjO. But if the valence of the metal in the oxide or
hydroxide is such that no corresponding bromide can be formed, then
reduction takes place as follows :
i. Pb"+^ becomes PbBrg and Br .
2, As^ becomes As'" and Br . The HBr must be concentrated and in
excess, and the As^' compound merely moistened with water: HjAsO^ +
2HBr = H3A8O3 + Brj + HjO . In presence of much water the reverse
action takes place: H3A8O3 4- Brj + HjO = HsAsO^ + 2HBr.
3. Sb^ becomes Sb'" and Br .
4. Bi^ becomes BiBr3 and Br .
5, Fe^^ becomes Fe'" and not Fe" , and Br .
6, Cr^ becomes CrBr3 and Br (a separation from a chloride if the solu-
tion be dilute) (Friedheim and Meyer, Z, anorg,, 1891, 1, 407). KBr is not
decomposed by a boiling concentrated solution of 'Kfitjd^ (separation
from KI) (Dechan, J. C, 1887, 51, 690).
7. Co"+° becomes CoBrj and Br .
8, Ni"+° becomes NiBrj and Br .
9. Mn"+° becomes MnBr. and Br (§269, 8; Jannasch and Aschoff, Z.
anorg., 1891, 1, 144 and 245). KMnO^ liberates all the bromine from KBr
in presence of CuSO^ (a separation of bromide from chloride (Baubigny
and Rivals, C. r., 1807, 124, 859 and 954).
Silver nitrate solution precipitates, from solutions of bromides, silver
hromide, AgBr, yellowish-whito in the light, slowly becoming gray to
black. The precipitate is insoluble in, and not decomposed by, nitric acid,
soluble in concentrated aqueous ammonia, nearly insoluble in concentrated
solution of ammonium carbonate, slightly soluble in excess of alkali
bromides, soluble in solutions of alkali cyanides and thiosulphates. It is
slowly decomposed by chlorine.
Solution of mercurous nitrate precipitates mercurous bromide, HgBr,
yellowish-white, soluble in excess of alkali bromides.
Solutions of lead salts precipitate, from solutions not very dilute, lead
hromidey PhBr^ , white.
B. — With non-metals and their compounds.
i. H3Fe(CN)o becomes H^Fe(CN)6 and Br . The HBr must be in excess
and concentrated, also the ferricyanide should be merely moistened with
water, as in the presence of much water the reverse action takes place:
2K,Fe(CN)e + Br^ = 2K,'Fe{CJl), + 2KBr .
3. HNO2 , in dilute solutions, no action (distinction from BEI) (Gooch and
Ensign, Am. S., 1890, 140, 145 and 283).
HNO3 becomes NO and Br .
§276, 8. HYDROBROMIC ACID, 347
S, Phosphorus compounds are not reduced.
Jf, H2SO4 becomes SOj and Br . Both acids must be concentrated and
hot, otherwise the reverse action takes place: SO2 + Brj + 2H2O = HoSO^
+ 2HBr . With H2SO4 , sp, gr. 1.41, no bromine is set free even when
solution is boiled (Feit and Kubierschky, J. Pharm.y 1891, (5), 24, 159).
The bromine of bromides is all liberated when warmed to 70° or 80° with
ammonium persulphate (separation from a chloride) (Engel, C. r., 1894,
118, 1263).
5. Chlorine liberates bromine from all bromides, even from fused silver
bromide (Nihoul, Z, amjew,, 1891, 441).
HCIO3 becomes HCl and Br . If the HClOg be concentrated other pro-
ducts may appear.
6. HBrO liberates Br from both acids ; the same with HBrO, .
7. HIO3 becomes I and Br .
8. Hydrogen peroxide liberates the bromine from hydrobromic acid at
100° (a distinction and separation from chloride). The bromine can best
be removed by aspiration (Cavazzi, Gazzetta, 1883, 13, 174).
7. Ignition. — Some bromides can be sublimed undecomposed in presence of
air; e. (/., AbBt, , SbBr, , HgBr and HgBr, . Some can be sublimed only by
exclusion of air and moisture; e. f/., AlBr, and NiBr, . Bromides of sodium and
potassium are not chang-ed bj' heat. Silver bromide melts undecomposed.
Many bromides, however, are more or less decomposed when ignited in pres-
ence of air and moisture: CuBr. becomes CuBr and Br .
8. Detection. — Bromides are usually oxidized to free bromine, which is
detected by its physical properties and by its color when dissolved in
CS2 (§275, 5). The oxidizing agent used to liberate the bromine varies
according to the conditions. Chlorine is more commonly employed and
acts when cold {QB5). A large excess of chlorine is to be avoided, as it
decolorizes bromine solutions with formation of a chlorbromide. Mtric
^cid when dilute acts slowly unless hot. H0SO4 , dilute, fails to oxidize
the HBr even when hot: but when concentrated and hot is sometimes
preferred. If chlorine be used, the mixture if alkaline must first be
acidified; otherwise a colorless bromato will be formed, free bromine not
beinor a visible intermediate step in the oxidation : KBr + ^>KOH + ^Clj
= KBrOs + 6KC1 + SH.O . If an iodide be present: (a) In absence of a
chloride precipitate with silver nitrate, and digest the precipitate with
NH^OH , which will dissolve the AgBr and none of the Agl . The filtrate
may be treated with H2S , which precipitates the silver as Ag^S , leaving
the bromine in the filtrate as NH^Br , which may be detected in the usual
way. (h) To the acid mixture add chlorine water and carbon disulphide,
shake and continue the addition of the chlorine water until the violet
color of the iodine solution disappears, when the brown color due to the
Iromine may be observed: 2BI -f 2KBr + 7CI2 + GHjO = 2HI0., + Bt^
348 EYPOBROMOIS ACID—BROMIC ACID. §876,9.
-|- 4KC1 + lOHCl . (c) To the solution from which the bases have been
removed add a cold saturated solution of potassium chlorate and dilute
sulphuric acid (one of acid to four of water); wann until the solution is
of a pale straw color, or colorless if only iodides are present. It may be
necessary to add more of the solution of potassium chlorate to complete
the oxidation of the iodine. Dilute the solution with water, cool, and
shake with carbon disulphide. See also §269, 8.
6KI + 6KBr + 2KC10. + 7H,S0« = 3l, + 3Br, + TK^SO* + 2HC1 -h 6H,0
61, -h xBr, +. lOKClO, + 5H,S0« + GTL^O = 12HI0, + xBr, + SKjSO* -h lOHCl
9. Estimation. — (a) It is converted into AgBr, and after gentle ignition
weighed as such. (6) The bromide is oxidized to free bromine, which is
passed into a solution of KI and the liberated iodine titrated with standard
KajSjO, . (c) The bromide is oxidized to bromine, which is passed into an
alkaline solution of arsenous acid. The excess of the arsenous acid is titrated
with a standard solution of KHX1O4 .
§277. Hypobromous acid. HBrO = 96.958 .
H'Br'O-", H — — Br.
The anhydride, BrjO , has not been isolated. The acid, HBrO , is a very
unstable yellow liquid, a strong oxidizing and bleaching agent. The hypo-
bromites are less stable than the corresponding hypochlorites. The calcium
and the alkali group hyjKjbromites may be prepared by adding bromine to the
respective hydroxides in the cold. The free acid is obtained by the action of
bromine upon mercuric oxide: 2HgO + 2Brj + HjO = HgjOBr, -|- 2HBrO;
also by the action of bromine upon silver nitrate: AgNO, + Br, -f- H^O =
AgBr 4- HBrO -f- HNO, (Dancer and Spiller, C. A'., IhGO, 1, 38; 1862, 6, 249).
The free acid as an oxidizing agent reacts in many cases similar to free
bromine. With HBr free Br is obtained from both acids (Schoenbein, J. pr.,
1863, 88, 475).
§278. Bromic acid. HBrOj = 128.958 .
H'Sr^O-^'a . H — — Br ~ J
1. FropertieB. — The anhydride, Br20B , has not been isolated: and the acid,
HBrO, , is known only in solution. It is a colorless liquid, smelling like bro-
mine. It is a strong oxidizing agent. The solution of HBrO, is decomposed
upon boiling, but by evaporating in a vacuum a solution containing about
50 per cent of the acid may be obtained.
2. Occurrence. — Neither the acid nor its salts are found in nature.
3. Pormation.— (a) Bv the electrolysis of HBr (Riche, C, r., isr>s, 46, 34S).
(ft) By the decomposition of AgBrO, by Br: SAgBrO, -|- aBr, + 3H2O = oAgBr
-h 6HBrO, . (c) An alkali bromate is made by adding bromine to a solution
of chlorine in sodium ca.'bonate (Kaommerer, J. pr., 1862, 85, 452).
4. Preparation.— Broniates of Ba , Sr , Ca , K and Na are made by the action
of bromine upon the respective hydroxides at 100'': CEOH -f ^^Br- = 5KBr +
EBrO, 4- 3H.0 . The free acid is prepared by adding dilute H,SO< in slight
excess to BaCBrO,)^; the slight excess of H,SO, being removed by the cautious
addition of Ba(OH)a .
§278, 8. BROMIC ACID, 34^
5. Solubilities. — AgBrOg is soluble in 123 parts of water at 24.5**
(Noyes, Z. pkys. Ch,, 1890, 6, 246). Ba(Br03)2 is soluble in 124 parts of
water at ordinary temperature and in 24 parts at 100° (Rammelsberg,
Pogg., 1841, 52, 81 and 86). With the exception of some basic bromates,
all other bromates are soluble in water.
6. Reactions. — A. — With metals and their compounds. — Bromic acid is
a powerful oxidizing agent, acting in most respects like free bromine.
It is usually reduced to hydrobromic acid, sometimes only to free bromine :
1. Hg' becomes Hg" and a bromide.
2. As'" becomes As^ and a bromide.
3. Sb'" becomes Sb^ and a bromide.
-4. Sn" becomes Sn^ and a bromide.
5. Cu' becomes Cu" and a bromide.
6. Fe" becomes Fe'" and a bromide.
7. Mn" becomes VLnO^ and bromine.
8. Cr"' becomes HjCrO^ and bromine.
Silver nitrate precipitates in solutions not very dilute, silver bromate,
AgBrOs , white, sparingly soluble in water, soluble in ammonium hydroxide,
easily soluble by nitric acid, its color and solubility in ammonium hydroxide
differing a little from the bromide (§276, 5). It is decomposed by hydro-
chloric acid with evolution of bromine — a distinction from bromides and
from other argentic precipitates.
B. — With non-metals and their compounds.
1. H2C2O4 becomes COo and Br. An excess of hot H2C2O4 changes the
Br to HBr (Guyard, Bl, 1879, (2), 31, 299).
HCNS becomes H2SO4 , HBr and other products.
H^FeCCN)^ becomes H3Fe(CN)e and HBr . An excess of HBrOs carries
the oxidation farther.
2. HNO2 reduces HBrOs , forming HNOa and Br .
3. PH3, HH2PO2 and HaPOj become HsPO^ and HBr.
^. S and SO2 become H2SO4 and HBr .
HjS forms first S then ttjSO^ .
5. HCl becomes CI and Br .
6. HBr forms Br from both acids.
7. HI becomes I and Br . With an excess of HBrO, the products are
HIO3 and Br (Kaemmerer, I c, Wittstein, Z., 1876, 15, 61).
7. Ignition. — All bromates are decomposed upon heating. EBrO.
NaBrOs and Ca(Br03)2 evolve oxygen and leave the bromides. Co(Br03)
Zn(Br0j)2 and other bromates evolve oxygen and bromine, leaving an oxide.
8. Detection. — The bromine is first liberated by some reducing agent
that docs not carry the reduction to the formation of HBr. H2C2O4 is a
3 y
2 >
350 IODINE. §278, 9.
very suitable agent for this purpose, since it does not change Br to HBr
except when hot and concentrated. The Br is detected by CS^ (§275, 8).
Sulphuric and nitric acids liberate bromic acid from metallic bromates,
the HBrOs remaining for some time intact, and the solution colorless. The
gradual decomposition of the HBrOs is first a resolution into HBr and 0,
and as fast as HBr is formed it acts with HBrOg , so as to liberate the
bromine of both acids. Now, if the solution contained bromide as well as
bromate, an abundance of free bromine is obtained immediately upon the
addition of dilute sulphuric acid in the cold. Hence, if dilute sulphuric
acid in the dilute cold solution does not color the carbon disulphide, and
if the addition of solution of pure potassium bromide immediately develops
the yellow color, while it is found that no other oxidizing agent is present,
we have corroborative evidence of the presence of a bromate. And, if we
treat a solution known to contain bromide with dilute sulphuric acid and
<jarbon disulphide, and obtain no color, we have conclusive evidence of the
absence of bromates. Hydrochloric acid transposes bromates and quickly
decomposes the bromic acid, liberating both bromine and chlorine.
A mixture of bromate and iodate, treated with hydrochloric acid, fur-
nishes bromine without iodine, coloring carbon disulphide yellow.
The ignited residue of bromates, in all cases if the ignition be done with
sodium carbonate, will give the tests for bromides.
9. Estimation. — The bromate is reduced to free bromine or to a bromide and
determined as such.
§279. Iodine. I = 120.85 . Usual valence one, five and seven.
1. Properties.— ASpfd/yc gravity, 4.948 at 17** (Gay-Lussac). Melting poinU
114.2**. Boiling point, 184.35** at 760 mm. pressure (Ramsay and Young-, J. C,
1886. 49, 453). At ordinary temperature iodine is a soft g^ray-black crystalline
solid with a metallic lustre. The thin crystals have a brownish-red" appear-
ance. Precipitated iodine Is a brownish-black powder. It vaporizes very
appreciably at ordinary room temperature with a characteristic odor, and may
be distilled with steam. The molecule of iodine vapor under about 800** is I;;
above that temperature dissociation takes place, until at 1700** it is complete
and the molecule consists of single atoms (Biltz and Meyer, B., 1889, 22, 725).
The vapor of iodine unmixed with other gases is deep blue, mixed with air
or other gases it is a beautiful violet. It is sparingly soluble in water to a
brown or yellowish-brown solution, which slowly bleaches litmus paper. It
stains the skin yellow-brown. The solution gradually decomposes in the sun-
light with formation of HI. It reacts similarly to bromine and chlorine, but
with much less intensity. The free element combines with starch,* forming
a compound of an intense blue color. This colored body is quite stable in the
cold; decolors upon warming, the color returning upon cooling. The reaction
of iodine with starch constitutes a very delicate reaction for the detection of
the presence of iodine. It also serves as an indicator in the volumetric estima-
tion of iodine, as all reducing acrents destroy the color by taking the iodine
into combination. Combined iodine does not react with starch.
♦The compound formed when Iodine unites with 8t«rch is re^rded by Bondonneau (Bl., 1877,
(2), 28, 452) as an addition compound of the cjmposltion (CaHioOft)sI .
§279,64,5. IODINE, 351
Colorless solutions are formed by all the alkali hydroxides with iodine; the
fixed alkali hydroxides forming iodides and iodates. With ammonia in water
solution it dissolves moi^e slowly, becoming colorless; the solution contains the
most of the iodine as ammonium iodide, and deposits a dark-brown powder,
termed " iodide of nitroyeVf'^ very easily and violently explosive when dry.
According to Chattaway (Am., 1900, 24, 138) this compound has the composi-
tion N,Hal3 .
The anhydride of iodic acid, IjOa , is the only stable compound of iodine and
oxygen. The chief acids of iodine are: Hydriodic acid, HI; iodic acid, HIO.;
periodic acid, HIO4 .
Hypoiodous acid is said to be formed by the action of alcoholic iodine upon
freshly precipitated mercuric oxide (Lippmann, C. r., 1866, 63, 968). Lunge and
Schoche (B., 1882, 15, 1883) prepared iodide of lime which seemed to contain
calcium hypoiodite, Ca(I0)2 .
2. Occurrence Found free in some mineral waters (Wanklyn, C, N., 1886, 54,
300). As iodides and iodates in sea water (Sonstadt, C. N., 1872, 25, 196, 231
and 241). In the ashes of sea plants. In small quantities in several minerals,
especially in Chili saltpeter as sodium iodate.
3. Formation. — From iodides by nearly all oxidizing agents: 2KI + Br, =
2KBr 4- !•: and from iodates by nearly all reducing agents: 2HIOs + 5H,C,04
= la + lo'CO, + 6HjO.
4. Preparation. — (a) The ashes of the sea plants are digested in hot water
and from the filtrate most of the salts removed by evaporation and crystalliza-
tion. The iodides remain in the mother liquor and from this the iodine is
obtained by treatment with IfnO, and H2SO4 . (ft) The sodium iodate in the
mother liquor of the Chili saltpeter is reduced with SO2 . the iodine precipitated
as Cul with CUSO4 . From the precipitate the iodine is recovered by distilla-
tion with MnOj and H,S04 . By far_the greatest portion of the iodine and
iodides of commerce is obtained from the Chili saltpeter deposits.
5. Solubilities. — It is soluble in about 5300 parts water at 10° to 12**
(Wittstein, J,, 1857, 123), differing from CI or Br in that it forms no
hydrate. It is much more soluble in water containing hydriodic acid or
soluble iodides. From a concentrated solution in KI the compound KI3
has been obtained. Iodine dissolves in very many organic solvents as
alcohol, ether, chloroform, glycerol, benzol, carbon disulphide, etc. Car-
bon disulphide readily removes the iodine from its solution or suspension
ill water; with small amounts of iodine imparting to the carbon disulphide
a beautiful violet color, with large amounts the CSg solution is almost
black.
6. Beactions. — A, — With metals and their compounds. — It unites slowly
by the aid of heat with Pb and Ag; more rapidly with Hg, As, Sb, Sn,
Bi, Cu, Cd, Al, Cr, Fe, Co, Ni, Mn, Zn, Ba, Sr, Ca, Mg, K and Na.
In oxidizing metallic compounds the iodine invariably becomes HI or
an iodide, depending upon whether the mixture be acid or alkaline. It
may, however, with certain substances act as a reducing agent, becoming
oxidized to iodate or periodate.
i. Hg' becomes Hg" in acid and in alkaline mixture.
2. As'" becomes As^ in presence of alkalis only.
3. Sb'" becomes Sb^ in presence of alkalis only.
i. Sn" becomes Sn^ in acid or in alkaline mixture.
5. Cr"' becomes Cr^ in presence of alkalis only.
^52 IODISE. §279, 64, e.
6, Fc" becomes Fc'" in presence of alkalis only.
7. Co" becomes Co"' in presence of alkalis only.
8, Ni" is not oxidized.
9. Mn" becomes Mn^^ in presence of alkalis only.
B. — With non-metals and their componnds.
i. K4Fe(CN)e is oxidized, forming KaFeCCN)^ and KI , action slow and
incomplete.
^. HNO3 forms HIO3 and NO . Strong HNO3 must be used (at least
sp. gr. 1.42). Action is slow. A very good method of making HIO3 .
J. HH2PO2 becomes H3PO4 with acids and with alkalis.
Jf, H2S becomes S and HI; no action if both substances be perfectly dry
(Skraup, C. C, 1896, i, 469) (separation of HjS from A8H3). According
to Saint-Gilles (.4. C/t., 1859, (3), 57, 221), in alkaline mixture from sLx
to seven per cent of the sulphur is oxidized to a sulphate.
H2SO3 becomes 92SO4 and HI. With a thiosulphate a tetrathionate is
fonned: 2Na2S203 + Ij = NajS^Oe + 2NaI (Pickering, J. C, 1880, 37,
128).
5. CI becomes ICl or ICI3 , depending upon the amount of chlorine
present, water should be absent. In the presence of water HCl and HIO,
are formed; in alkaline mixture a chloride and a periodate: L + TClo +
IGNaOH = 14NaCl + 2NaI0, + 8H0O . HCIO3 forms HIO3 and HCl:
oHClOa + 3I2 + SH^d = 6HIO3 + 5HC1 .
6. Br becomes IBr, decomposed by water (Bornemann, A,, 1877, 189,
183). In alkaline mixture with an excess of Br a bromide and an iodate:
I. + 5Bro + 12K0H = 2KIO3 + lOKBr + 6H2O . HBr03 becomes Br
and HIO3 .
7. Iodine combines with KI in concentrated solution to form Kl3(Kn2) .
7. Ignition. — See I.
8. Detection. — Iodine is recognized by the yellow to black color when
mixed with water; the violet color when dissolved in carbon disulphide;
the reddish color when dissolved in chloroform or ether; the blue color
when added to a cold solution of starch ; the violet color of the vapors, etc.
The presence of tannin interferes with the usual tests for iodine unless a
drop or two of ferric chloride solution be added (Tessier, Z., 1874, 11, 313).
9. Estimation. — (0) It is reduced to an iodide, precipitated with AgNO, , and
after drying- at 150°. weighed as Agl . It is estimated vohimetrically with a
standard solution of NajSjO, , using- starch as an indicator, (h) The iodine
dissolved in potassium iodide is treated with an alkaline solution of hydrogen
peroxide in an azotometer, the oxygen liberated being a measure of the amount
of iodine present (Baumann, Z. angew,, 189 1, 204).
§280, 5. • HYDRIODIC ACID, 353
§280. Hydriodic acid. HI = 127.858 .
HI-', H — I.
1. Properties.— lfo?erii?ar weight, 127.858. Vapor density, 63.927. A colorless
incombustible gas. At atmospheric pressure it solidifies at — 51°. At 0** it
liquefies under a pressure of 3.97 atmospheres (Faraday, A, Ch,, 1845, (3), 15,
260). The constant boiling point of the aqueous solution of the g^s is 127**,
M'hich solution contains 57 per cent of HI and has a specific gravity of 1.694
(Roscoe, J, C, 1861, 13, 160). Gaseous HI is dissociated by heat, slowly at 260";
rapidly at 240° (Lemoine, A. Ch., 1877, (5), 12, 145). Iodine separates from the
water solution of the acid when exposed to, the air.
2. Occurrence. — Not found free in nature, but in combination as iodide or
iodate.
3. Formation. — (a) By direct union of the elements at a full red heat (Merz
and Holzmann, B,, 1889, 22, 869). (b) By direct union of the elements in pres-
ence of platinum black at 300° to 400° (Lemoine, C. r., 1877, 85, 34). (c) From
Bal, by adding H3SO4 in molecular proportions, (d) By the action of iodine
upon NajSO, or Na^SaO, (Mene, C. r., 1849, 28, 478). (e) By the action of iodine
upon moist calcium hypophosphite : Ca(H,P0,)2 + 41, 4- 4HaO = CaH4(P04)a
-f 8HI (Mcne, /. c).
Iodides are formed by the direct action of iodine upon the metals; or better,
by* the action of HI upon the oxides, hydroxides or carbonates of those metals
Avhose iodides are soluble in water. Iodides of lead, silver and mercury are
formed by precipitation.
4. Preparation. — (a) By passing HjS into a mixture of finely divided iodine
suspended in water, adding more iodine as fast as the color disappears: 21, +
2H,S = 4HI + Sa (Pellagri, Gazzetta, 1875, 5, 423). (6) By bringing moist red
phosphorus in contact with iodine: P^ -|- lOl, + 16HaO = 4H,P04 + 20HI
(Meyer,!?., 1887, 20, 3381). (c) By passing vapors of iodine into hot liquid
paratline (Crismer, B., 1884, 17, 649). (d) By heating iodine with copaiba oil
(Bruylants, B., 1879, 12, 2059). It cannot be prepared by adding H^SO^ to an
iodide and distilling (5).
5. Solubilities. — Iodides of lead, silver, mercury and cuprosum are in-
soluble. Iodides of other ordinary * metals are soluble, those of bismuth,
tin and antimony requiring a little free acid to liold them in solution.
Lead iodide is sparingly soluble in water (§57, 5c). Mercuric iodide is
readily soluble in excess of potassium iodide, forming a double iodide,
KjHgl^; most other iodides are more soluble in a solution of potassium
iodide than in pure water. The iodides of the alkalis, Ba, Ca and Hg"
are soluble in alcohol; Hgl and Agl are insoluble. All iodides in solution
are transposed by HCl or by dilute H2SO4 . Hot concentrated HsSO^
decomposes all iodides, those of Pb , Ag and Hg slowly but completely,
SO. and I being produced: 2KI + 2IL^S0^ = KjSO^ + I2 + SOj + 2H2O .
HNO3 in excess first transposes then decomposes soluble iodides: 6KI +
8HNO3 = 6KNO3 + 3I2 + 2N0 + 4HoO . If the HNO3 be concentrated
the iodine is further oxidized: SI. + IOHKO3 = 6HIO3 + lONO + 2H2O .
Long-continued boiling with HNO3 , sp. gr. 1.42, decomposes the insoluble
iodides. Chlorine in the cold decomposes all soluble iodides, by heating
with chlorine the insoluble iodides are also decomposed: 2KI + CI2 =
• Thallium iodide, Tl I, is perfectly insoluble In cold water, a distinction and separation from
bromides and chlorides (Huebner. Z., 1872, 11, 897). Palladous iodide is insoluble in water.
354 BYDRIODIC ACID, §280, 6i.
2KC1 + I2. With an excess of chlorine the iodine is further oxidized:
I2 + 5CI2 + 6H2O = 2HIO3 + lOHCl . Silver iodide is ahnost insoluble
in ammonium hydroxide or ammonium carbonate (distinction from silver
chloride). It is soluble in KCN . Agl and Pblj are soluble by decomposi-
tion in solution of alkali thiosulphates : Agl -|- NagSjOj ^ Nal +
NaAgSjOs • Lead iodide is soluble in a solution of the fixed alkalis.
G. Reactions. — .4. — With metals and their compounds. — Silver nitrate
solution in excess precipitates, frgm solutions of iodides, silver iodide^ Agl ,
yellow-white, blackening in the light without appreciable separation of
iodine. For solubilities see paragraph above.
Solution of mercoric chloride precipitates the bright, yellowish-red to
red, mercuric iodide, HgL . The precipitate redissolves on stirring, after
slight additions of the mercuric salt, until equivalent proportions are
reached, when its color deepens. For the solubilities of the precipitate
see §58, 6/. Solution of mercorous nitrate precipitates mercurous iodidty
Hgl , yellow to green (§68, G/).
Solution of lead nitrate or acetate precipitates, from solutions of iodides
not very dilute, lead iodide, Pbia , bright-yellow — soluble, as stated in full
in §67, 5c.
Palladous chloride, PdClj , precipitates, from solutions of iodides, pal-
ladous iodide^ PdL , black, insoluble in water, alcohol or dilute acids, and
visible in 500,000 parts of solution. The reagent does not precipitate
hromine at all in moderately dilute solutions, slightly acidulated with HCl .
Palladous iodide is slightly soluble in excess of the alkali iodides, and is
soluble in ammonium hydroxide (§106).
Copper salts precipitate from solutions of iodides cuprous iodide (white)
mixed with iodine (black) : 2CuS0^ + 4KI = 2CuI + SKjSO^ + I, . If
sufficient reducing agents (as sulphurous acid) are present to reduce the
liberated iodine to HI, only the white 'cuprous iodide will be precipitated
(a distinction from bromides and chlorides).
When metals are attacked by HI an iodide is formed and hydrogen is
evolved. Ifydriodic acid unites with all metallic oxides and hydroxides
(expect ignited CroOg) to form iodides; frequently, however, iodine is
liberated arid an iodide of lower metallic valence is formed:
L Pb"+° becomes Pb" .
2. As^ becomes As'" ; KI has no action upon normal EjAsO^ (Friedheim
and Meyer, Z. anorg., 1891, 1, 409).
3, Sb'v becomes sV" .
i. Biv becomes Bi'" .
5. Cn" becomes Cu' . Soluble iodides reduce normal cupric salts, but
have no reducing action in alkaline mixture or upon cupric hydroxide.
With phenylhydrazine sulphate and cupric sulphate the iodine of iodides is
§280, 6B, 6, HTDRIODIC ACID. 355
completely precipitated (separation from chlorides) (Eaikow, Ch, Z., 1894,
18, 1661).
6. Fc"' becomes Fc" (§269, 8).
7. Cr^ becomes Cr"' . K^CtO^ is not reduced by KI even upon boiling
the concentrated solutions. K^Ct^Oj with KI slowly gives I and Cr"' in
the cold. When KI is boiled with a concentrated solution of KjCrjOj the
iodine is completely liberated (separation from bromides and chlorides
which are unchanged): 6KI + SKjCraO^ = SKjCrO^ + CijOg + SIj
(Dechan, J. C, 1886, 60, 682; 1887, 61, 690). When Agl is boiled with
K2Cr207 and H2SO4 no iodine is evolved, chromium is reduced and the
iodide becomes silver iodate: K^Ct^Oj + Agl + SHjSO^ = 2KHSO4 +
Cr2(S0j3 + Agio, + 4H2O (Macnair, J. C, 1893, 63, 1051).
S. Co"+° becomes Co"; KI has no reducing action upon cobaltic hy-
droxide.
9. Hi"+° becomes Hi''; KI reduces Ni'" , liberating iodine.
10. Mn"+° becomes Mn" . When KI is boiled with KMnO^ the manga-
nese becomes HnOj and the iodide is oxidized to an iodate: 6K][n04 +
SKI + 3H2O = 3KIO3 + 6Mn02 + 6K0H (Groeger, Z. angew., 1894, 13
and 52) (distinction from bromides, which do not decolor permanganates).
B. — With non-metals and their compounds.
1. H3Fe(CH)e forms H4Fe(CN)e and I; the reaction also takes place in
neutral mixture.
2. HNO2 forms NO and I (separation of iodide from bromide and
chloride) (Jannasch and Aschoff, Z. anorg., 1891, 1, 144 and 245).
HNO3 forms HO and I, with further oxidations to HIO3 with concen-
trated HHO3 . The HHO2 acts much more rapidly than the HHOg .
3. No reduction with phosphorous compounds.
-4. H2SO4 dilute no action; with the concentrated acid in excess, SO2 and
I are formed: 2KI + 3H2S0^ = I2 + SO2 + 2KHSO4 + 2H2O ; if KI be
added in excess to boiling H2SO4 , HjS and I are formed: SKI + 9H2SO4 =
4I2 + H28 + 8KHSO4 + 4H2O (Jackson, J. C, 1883, 43, 339). Ammo-
nium persulphate liberates iodine from iodides at ordinary temperature
(Engel, C. r., 1894, 118, 1263).
5. Cl in excess forms HCl and HIO3 ; with excess of HI , HCl and I are
formed. In the presence of a fixed alkali a periodate and a chloride are
formed: KI + 8K0H + 4CI2 = 8KC1 + KtO^ + 4H2O . HypocUorous
acid oxidizes to iodine, then to iodic in acid solution; in alkaline solution
to periodate.
HCIO3 with excess of HI forms HCl and I; with excess of HCIO3 HCl
and HIO3 .
6. Br forms I and HBr or a bromide.
366 HTDRIODIC ACID. §280, 6B, 7.
HBrO^ with excess of HI forms HBr and I ; with excess of HBrO^^ Br
and HIO3 .
7. HIO3, iodine is liberated from both acids: HIO, + 5HI = 3Ij +
3H2O . HIO^ gives iodine.
8, H2O2 becomes H^O , and I (§244, 6B6) (Cook, J. C, 1885, 47, 471).
P. Ozone promptly liberates iodine from soluble iodides. Atmospheric
oxygen decomposes HI and ferrous and calcium iodides slowly, the alkali
iodides not at all.
7. Ignition. — As a general rule iodides strongly ignited in presence of air
and moisture evolve iodine, leaving the oxide of the metal. Ignited in absence
of air or moisture the following iodides are not decomposed: KI , Nal , Bal. ,
Cal, , Sri, , Mnia , AH, , SnI* , Pbl, , Agl and Hgl, . See Mitscherlich (Pogg.,
1833, 29, 193), Personne (C. r., 1862, 54, 216) and Gustavson (A., 1873, 172, 173).
8. Detection. — The iodide is oxidized to free iodine by one of the re-
agents mentioned in (6) above. With a dr}' powder hot concentrated
H2SO4 is usually employed when the iodine is detected by the violet fumes
evolved, condensing in the cooler portion of the test tube. With solu-
tions the usual reagent is chlorine water. The iodine is recognized by
the violet color when shaken with CSj , or the bright-red color with CHCl, .
In case a large amount of iodine be present the CSg solution may be almost
black. In this case large dilution with CSo is necessary to detect the violet
color. If but a small amount of iodine be present the chlorine must be
added very cautiously or the iodide will all be oxidized to the colorless
iodic acid.* With small amounts of iodide, nitric acid is less liable to
cause error as relatively much more nitric acid is required to oxidize the
iodine to iodic acid. For the detection of small amounts of iodide a
cupric salt strongly acidulated with HCl is an excellent reagent for the
oxidation : 2CuCl, + SKI = 2CnCl + 2KC1 + I, .
If insoluble iodides are present they should be transposed by H2S.
the insoluble sulphide removed by filtration, the excess of HjS removed
by boiling, and the solution then tested for hydriodic acid. Or the
insoluble iodide should be reduced by Zn and H0SO4: 2AgI + Zn + H.>SO^
= 2Ag -f ZnSO^ + 2HI . The filtrate may then be tested for hydriodic
acid. The insoluble iodide may also be fused with NajCOg , and after
digestion with water the filtrate acidulated and tested for hydriodic aci<l
That is, the solution must be acidulated before chlorine water is added,
else the iodine will be oxidized to an iodate or periodate.
9. Estimation. — Gravimetrically by precipitation as Agl and weigrhingr a^
such after gentle ignition. Volumetrically by oxidation to iodine and titration
with standard N&.S.O, (Grower, Z. angctc,\ 1894, 52).
• To test potassium bromide for traces of an iodide It is recommended to add CS, and cupric
sulphate or a small amount of ferric alum. Or add chlorioe water and then a few crystals of
ferrous sulphate ; then shake with CS, (Brito« C. N„ 1884, 50, 210\
§281, BA. IODIC ACID, 357
§281. Iodic acid. HIO3 = 175.858 ,
H'F0-"8, H — — I
1. Properties. — Iodic acid is a white crystalline solid; its solution saturated
at 14° contains G8.5 per cent HIO, , and has a specific gravity of 2.1629 (Kaem-
merer, Pogg., 1869, 138, 390). At 170° it loses water, forming iodic anhydride,
IjOa , a white crystalline solid, which, at 300°, dissociates into iodine and
oxygen. See Ditte, A, Ch., 1870, (4), 21, 5. It is readily soluble in water and
in alcohol; the solutions redden litmus and afterwards bleach it.
2. Occurrence.— The free acid is not found in nature. It is found as Ca(I0,)2
in sea water, and as sodium iodate in Chili saltpeter (Sonstadt, C, N., 1872, 25,
196, 231 and 241; Guyard, BL, 1874, (2), 22, 60).
3. Formation. — (a) By elect roly zing a solution of I or HI (Riche, C. r., 1858,
46, 348). (h) By the action of chlorine on iodine in the presence of much
water. The HCl formed cannot be expelled by boiling without decomposing
the HIO, . It must be removed by the careful addition of Ag,0 . (c) By
adding water to ICl, and washing with alcohol: 2101, + 3H,0 = HIO, -f
5HC1 -h ICl . (d) KIO, is made by treating iodine with KOH: 3l, + 6K0H =
:KI0, -I- 5KI -h 3H2O . And then washing with alcohol to remove the KI . (c)
By heating potassium chlorate and iodine: lOKClO, -h 61, -h 6H,0 = 6KHI,0«
-f 4KC1 + 6HC1 (Bassett, J. C, 1890, 57, 760). (f) By boiling iodine with barium
hydroxide until neutral, filtering and decomposing with sulphuric acid (Steven-
son, C. N., 1877, 36, 201). (g) By the action of I upon AgKO,: 5AgN0, + 3l, +
3H,0 = 5AgI + 5HN0, + HIO, .
lodates of the alkalis and alkaline earths are easily made by the action of
iodine on the hydroxides, and separation by alcohol or by crystallization from
the iodides which are formed in the reaction. All iodates may be made by
action of the acid on the hydroxides or carbonates.
4. Preparation. — (a) Iodine is oxidized by boiling with nitric acid sp, gr,
1.52, and removing the excess of the nitric acid by evaporation. (6) By adding
a slight excess of H3SO4 to Ba(I0,)2 and removal of the excess of H2SO4 by
the careful addition of BaCIO,), . (c) By boiling a solution of potassium
iodide with an excess of potassium permanganate in neutral or alkaline solu-
tion: KI -h 2KMn04 + H^O = KIO, -h 2K0H + 2Mn Oa (G roger, Z. angew.,
1894, 13 and 52). (d) The very stable potassium biiodate, KHIaO, , is formed by
recrvstallizing a water solution of equal portions of KIO, and HIO, . It is
soluble in 18.66 parts water at 17° (Meineke, A., 1891, 261, 359).
5. Solubilities. — Ba(I03)2 is soluble in about 3000 parts water at ordi-
nary temperature; and in about 600 parts at 100° (Kremers, Pogg,, 1851,
«4, 27; Spica, Oazzetta, 1894, 24, i, 91). AglOg is soluble in 27,700 parts
of water at 25° ; in 2.1 parts NH^OH (10 per cent) at 25° (separation from
silver iodide); in 1044.3 parts HNO3, sp. gr, 1.21 at 25° (Longi, Oazzettay
1883, 13, 87). The iodates of Ag , Ba , Pb , Hg , Sn , Bi , Cd , Fe and Cr
require at 15° more than 500 parts of water for their solution and the
following require less : Cn , Al , Co , Ni , Mn , Zn , Sr , Ca , Mg , K and Na .
They are all transposed by concentrated HNO3 or HgSO^; and are decom-
posed by concentrated HCl . They are soluble in the alkalis in so far as
the corresponding metallic oxides are soluble in those reagents. Most
of the iodates are insoluble in alcohol (with E , Na , Ba and Ca iodates a
separation from iodides).
6. Beactions. — A. — ^With the metals and their compoiinds. — A few metals
358 IODIC ACID, §281, 64, 1.
are attacked evolving hydrogen, forming iodates, sometimes traces of
iodides. With the following metallic compoimds the valence of the metal
is changed:
i. As'" becomes As^ with liberation of iodine. AsH, in excess forms
A8° , with the HIO, in excess As^ (Ditte, /!., 1870, 166, 336).
2. Sb'" becomes Sb^ with liberation of iodine. SbH, forms Sb** .
3. 8n" becomes Sn^^ ^nd HI .
Jf. Cu' becomes Cn" with liberation of iodine.
5. Fc" becomes Fc'" with liberation of iodine.
Solution of silver nitrate precipitates, from even very dilute solutions of
iodates and from solutions of iodic acid if not very dilute, silver iodak,
AglOj , white, crystalline, soluble in ammonium hydroxide, soluble in an
excess of hot HNO3 . In the ammoniacal solution, hydrosulpliario acid
forms silver sulphide, sulphur and ammonium iodide.
Barium chloride precipitates barium iodate, Ba(I03)2, slightly soluble
in cold, more soluble in hot water, insoluble in alcohol, soluble in
hot dilute nitric acid, readily soluble in cold dilute hydrochloric acid.
Hence, dilute solutions of free iodic acid should either be neutralized or
tested with barium nitrate. This precipitate, by addition of alcohol, is a
complete separation from iodides, and, when well washed, decomposed with
a very little sulphurous acid (8), and found to color carbon disulphide
violet, its evidence for iodic acid is conclusive. Barium iodate is trans-
posed by ammonium carbonate.
Salts of lead give a white precipitate of had iodate, Pb(I08)2 . Ferric
chloride gives, in solutions not dilute, a yellowish-white precipitate of
ferric iodate, Fe(I03)3, sparingly soluble in water, and freely soluble in
excess of the reagent. Boiling decomposes it.
Alcohol precipitates potassium iodate from water solution, an approxi-
mate separation from iodide.
B. — With non-metals and their compounds.
1. H2C2O4 becomes COg and I . Action is slow unless solutions are hot.
Carbon (except diamond) heated in sealed tubes becomes COj with sepa-
ration of I (Ditte, l. c),
H,Fe(CN)e becomes H3Fe(CN)e and I .
HCNS forms H0SO4 , 1 and some other products.
3, HNO2 becomes HNO3 and I .
3. PH3 becomes H3PO4 and I . With an excess of PH3 , HI is formed.
Water in which phosphorus has stood reduces iodic acid to iodine (Corne,
J. Pharm,, 1878, (4), 28, 386).
HH0PO2 becomes H3PO4 and I .
4. H28 becomes S and I . Thiosulphates form first iodine then an iodide.
§281, 9c. IODIC ACID. * 359
H2SO3 , with excess of HIO, , becomes HjSO^ and I; with excess of H2SO3 ,
HjSO^ and HI .
5, HCl , if concentrated, forms ICI3 and CI , iodine not being liberated.
G. HBr forms Br and I .
7. HI forms I from both acids. The addition of tartaric acid to a mix-
ture of KI and KIO3 is sufficient to give an immediate test for free iodine
with CSg . It must be remembered that an iodide alone rendered acid will
^ive a test for free iodine after a short time.
8, Morphine reduces iodic acid with separation of iodine.
7. Ignition. — Potassium . and sodium iodates on ignition form iodides
and evolve oxygen (Cook, J. C, 1894, 65, 802). Many other iodates evolve
oxygen but the iodide formed is further decomposed as stated in §276, 7.
Iodates in dry mixture with combustible bodies are reduced, on heating
or concussion, with detonation, but much less violently than chlorates or
nitrates.
8. Detection. — It is usually detected, after acidulation, by treatment
with some reducing agent for the formation of free iodine. H2SO, is
often employed because it acts rapidly and in the cold ; but traces of HIO3
frequently escape detection for the least excess of H2SO3 at once reduces
the iodine to colorless hydriodic acid. A desirable reagent for this reduc-
tion is one that will act rapidly in the cold, and in no case cause the
further reduction to hydriodic acid. The following reducing agents have
been used: K^FeCCN)^ acidulated with dilute HgSO^ , H3ASO3 , CnCl , FeSO^ ,
morphine sulphate and uric acid. To detect KIO3 in KI it is recom-
mended by Schering (J. C, 1873, 26, 191) to add a crystal of tartaric
acid to the solution. The formation of a yellow zone is indicative of an
iodate. Hydrochloric acid may be used, but if it contains a trace of
chlorine it will give the test for an iodate. Iodine frequently occurs in
nitric acid as iodic acid. Hilzer (J. C, 1876, 29, 442) directs to add equal
volumes of water, carbon disulphide, and then coarse zinc filings. It may
he necessary to warm the solution slightly. Biltz (C. C, 1877, 86) dilutes
the HITOa with water, adds starch solution and then HjS solution drop
hy drop. A blue zone is formed if HIO^ be present.
0. Estimation. — (a) By precipitation with AgNO, , and after drying at 100®
weighing as AglO.. (h) By reducing to an iodide and estimating as such.
(e) By treating with KI acidulated with HaSO^ , and titrating the iodine lib-
erated with standard NajSjO, .
360 , PERIODIC ACID.
§282. Periodic acid, HIO^ = 191.858 .
H H H
\ I /
II \l/
HT^O-^, or TL\T^Qr\, H — — 1 = or H — — I — — H.
The anhydride, IjOt , has not been isolated, and but one acid is known in the
free condition, HI04,2H30 or HsIO, . This acid exists in colorless monoclinic
crystals, which do not lose water at 100**. It melts at 133", and at a
higher temperature it decomposes into iodic anhydride, water and oxygen
(Kimmins, J. C, 1887, 51, 356; and 1889, 55, 148). Numerous periodates have
been prepared as if derived from one or the other following named acids:
HIO4 , H3IO5 , H,IO, , HJaO„ , HJjOii , Hi,I,Oi, , HioI^Oj, , H^oI^O.^
(Rammelsberg, Pogg,, 1865, 134, 368, 499).
The free periodic acid, HglOs , is prepared: (a) By oxidizing iodine with per-
chloric acid: 2HCIO4 + I, -f 4HaO = 2H5IO, -}- CI, (Kaemmerer, Pogg.^ 1869,
138, 406). (6) By heating iodine or barium iodide with a mixture of barium
oxide and barium peroxide, digesting with water, and transposing the
BasCIO.), thus obtained with the calculated amount of sulphuric acid (Ram-
melsberg, Pogg,, 1869, 137, 305). (r) By conducting chlorine into sodium iodate
in presence of sodium hydroxide: KalO, + 3NaOH + Cl^ = KajHalOe +
2KaCl . This ^cid periodate dissolved in water with a little nitric acid and
then precipitated with silver nitrate, forms the silver salt, A^zHsIOq . This
precipitate is dissolved in nitric acid and evaporated on the water-bath, when
orange-colored crvstals of silver meta periodate are formed according to the
following: 2Ag,H3lOe -f 2HN0, = 2AgI04 + 2AgN0, -f 4H3O . Water decom-
poses this precipitate: 2AgI04 + 4HsO = HsIO, -f AgsH.IO, . Or the silver
periodate, AglO^ , is decomposed by CI or Br (Kaemmerer, I. c, p. 390).
The silver salts vary in color: AgI04 is orange; AgjHIOs , dark brown;
Ag^IjOg , chocolate colored; while silver iodate is white (a distinction). In the
general reactions periodic acid and periodates resemble iodic acid and iodates.
H^CnOt becomes CO. and I .
HaPOa becomes H3PO4 and HI .
HjS becomes S and HI .
H3SO, becomes H3SO4 and HIO, without separation of iodine when the two
acids are present in molecular proportions. The presence of a greater pro-
portion of H.SOs causes, first, separation of iodine with final complete reduc-
tion to HI (Selmous, 7?., 1888, 21, 230):
HIO, + H,SO, = HIO, + H3SO,
3HI0, + SH2SO, = HIO, -f I2 + RH,SO, + H,0
2HI0, + 7H,S03 = I3 + 7H3SO, -h H,0
HIO, + 4H,S03 = HI + 4H,S0,
HCl becomes CI and ICl,
HI forms I from both acids.
According to Lautsch (J. pr., 1867, 100, 86), its behavior with mercurous
' nitrate is characteristic. The pentasodic periodate, Na^IO, , gives a light-
yellow precipitate, HgalO, .
§283.
COMPARATIVE REACTIONS OF HALOGEN COMPOUNDS.
361
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PAET IV -SYSTEMATIC EXAI^ONATIONS.
REMOVAL OF ORGANIC SUBSTANCES.
§284. The methods of inorganic analysis do not provide against interference
by organic compounds; and, in general, it is impossible to conduct inorganic
analysis in material containing organic bodies. The removal of the latter can
be effected, 1st, by combustion at a red or white heat, with or without oxidiz-
ing reagents: 2d (in part), by oxidation with potassium chlorate and hydro-
chloric acid on the water-bath (§69, 6W); 3d, by oxidation with nitric acid in
presence of sulphuric acid, at a final temperature of the boiling point of the
latter (§79, 6'eS): 4th, by solvents of certain classes of organic substances:
5th, by dialysis. These operations are conducted as follows:
§285. Combustion at a red or white heat, of course, excludes analysis for mer-
cury, arsenous and antimonous bodies (except as provided in §70, 7), and
ammonium. The last-named constituent can be identified from a portion of the
material in presence of the organic matter (§207, 3). If chlorides are present
some iron will be lost at temperatures above 100°, and potassium and sodium
waste notably at a white heat, and slightly at a full red heat. Certain acids
will be expelled, and oxidizing agents reduced.
The material is thoroughly dried and then heated in a porcelain or platinum
crucible, at first gently. It will blacken, by separation of the carbon of the
organic compounds. The ignition is continued until the black color of the
carbon has disappeared. In special cases of analysis, it is only necessary to
char the material: then pulverize it, digest with the suitable solvents, and
filter: but this method does not give assurance of full separation of all sub-
stances. Complete combustion, without use of oxidizing agents, is the way
most secure against loss, and entailing least change of the material: it is. how-
ever, sometimes very slow. The operation may be hastened, with oxidation of
all materials, by addition of nitric acid, or of ammonium nitrate. The material is
first fully charred: then allowed to cool till the finger can be held on the
crucible: enough nitric acid to moisten the mass is dropped from a glass rod
upon it, and the heat of the water-bath continued until the mass is dry, when
it may be very gradually raised to full heat. This addition may be repeated
as necessary. The ammonium nitrate may be added, as a solid, in the same
way.
§286. Oxidation with potassium chlorate and liydrochloric acid on the tcater-bath
does not wholly remove organic matter, but so far disintegrates and changes
it that the filtrate will give the group precipitates, pure enough for most tests.
It does not vaporize any bases but ammonium, but of course oxidizes or
chlorinates all constituents. It is especially applicable to viscid liquids: it may
be followed by evaporation to dryness and ignition, according to the paragraph
above.
The material with about an equal portion of hydrochloric acid is warmed on
the water-bath, and a minute portion of potassium chlorate is added at short
intervals, stirring with a glass rod. This is continued until the mixture is
wholly decolored and dissolved. It is then evaporated to remove chlorine,
diluted and filtered. If potassium and chlorine are to be tested for, another
portion may be treated wnth nitric acid, on the water-bath. The organic
matter left from the action of the chlorine or the nitric acid may be sufficient
to prevent the precipitation of aluminum and chromium in the third group of
bases: so that a portion must be ignited. As to arsenic and antimony, see
§70, 7.
§292. PRELIMINARY EXAMINATIONS OF SOLIDS. 363
§287. The action of sulphuric with nitric acid at a gradually increasing heat
leaves behind all the metals (not ammonium), with some loss of mercury and
arsenic (and iron?) if chlorides are present in considerable quantity. In this,
as in the operations before mentioned, volatile acids are lost — sulphides partly
oxidized to sulphates, etc.
The substance is placed in a tubulated retort, with about four parts of con-
centrated sulphuric acid, and gently heated until dissolved or mixed. A funnel
is now placed in the tubulure, and nitric acid added in small portions, gradu-
ally raising the heat, for about half an hour — so as to expel the chlorine, and
not vaporize chlorides. The material is now transferred to a platinum dish
and heated until the sulphuric arid begins to vaporize. Then add small portions
of nitric acid, at intervals, until the liquid ceases to darken by digestion, after
n portion of nitric acid is expelled. Finallj', evaporate off the sulphuric acid,
using the lowest possible 'heat at the close.
§288. The solvents used are chiefly ether for fatty matter^ and alcohol or ether,
or both successively, for resins. Instead of either of these, benzol may be
used; and many fats and some resins may be dissolved in petroleum ether.
It will be observed that ether dissolves some metallic chlorides, and that
alcohol dissolves various metallic salts. Before the use of either of these sol-
vents upon solid material, it should be thoroughly dried and pulverized. Fatty
matter suspended in water solutions may be approximately removed by filter-
ing through wet, close filters; also by shaking with ether or benzol, and decant-
ing the solvent after its separation.
§289. By Dialysis, the larger part of any ordinary inorganic substance can
be extracted in approximate purity from the greater number of organic sub-
stances in water solution. The degree of purity of the separated substance
■depends upon the kind of organic material. Thus albuminoid compounds are
almost fully rejected; but saccharine compounds pass through the membrane
quite as freely as some metallic salts. (Consult Watts' Dictionary, 1894, IV, 172).
PRELIMINARY EXAMINATION OF SOLIDS.
§290. Before proceeding to the analysis of a substance in the wet way, a
careful study should usually be made of the reactions which the substance
undergoes in the solid state, when subjected to a high heat, either alone or in
the presence of certain reagents, before the blow-pipe, or in the flame of the
Bunsen burner. This examination in the dry way precedes that in the wet,
and should be carried on systematically, following the plan laid down in the
tables, and noting carefully every change which the substance under investiga-
tion undergoes, and if necessary making reference to some of the standard
works on blow-pipe analysis. In order to understand fully the nature of these
reactions, the student should first acquaint himself with the character of the
different parts of the flame, and the use of the blow-pipe in producing the
reducing and oxidizing flames.
§291. The flame of the candle, or of the gas-jet, burning under ordinary circum-
stances, consists of three distinct parts; a dark nucleus or zone in the centre,
surrounding the wick, consisting of unburnt gas — a luminous cone surrounding
this nucleus, consisting of the gases in a state of incomplete combustion. Ex-
terior to this is a thin, non-luminous envelope, where, with a full supply of
oxygen, complete combustion is taking place: here we find the hottest part of
the flame. The non-luminous or outer part is called the oxidizing flame; the
luminous part, consisting of carbon and unconsumed hydrocarbons, is called
the reducing flame.
§292. The flame produced by the lloic-pipr (or Bunsen burner) is divided into
two parts: the oxidizing flame, where there is an excess of oxygen, correspond-
ing to the outer zone of the candle-flame; and the reducing flame, where there
is an excess of carbon, corresponding to the inner zone of the candle-flame.
Upon the student's skill in producing these flames depend very largely the
results in the use of the blow-pipe.
In order to produce a good oxidizing flame, the jet of the blow-pipe is placed
just within the flame, and a moderate blast applied — the air being thoroughly
mixed with the gas, the inner blue flame, corresponding to the exterior part
364 PRELIMiyARY EXAUIXATIOyS OF SOLIDS, §298-
of the candle-flame, is produced: the hottest and most effective part is just
before the apex of the blue cone, where combustion is most complete.
The reducing^ fl&me is produced by placing the blow-pipe just at the edge of
the flame, a little above the slit, and directing the blast of air a little higher
than for the oxidizing flame. The flame assumes the shape of a luminous cone,
surrounded by a pale-blue mantle; the most active part of the flame is some-
what beyond the apex of the luminous cone.
§293. The blast with the blow-pipe is not produced by the lungs, but by the
action of the muscles of the cheek alone. In order to obtain a better knowledge
of the management of the flame, and to practise in producing a good reducing
flame, it is well to fuse a small grain of metallic tin upon charcoal, and raising
to'a high heat endeavor to prevent its oxidation, and keep its surface bright:
or better, perhaps, to dissolve a speck of manganese dioxide in the borax bead
on platinum wire — the bead becoming amethyst-red in the outer flame and
colorless in the reducing flame. The beginner should work only with sub-
stances of a known composition, and not attempt the analysis of unknown
complex substances, until he has made himself perfectly familiar with the
reactions of at least the more frequently occurring elements.
The amount of substance taken for analysis should not be too large; a
quantity of about the bulk of a mustard-seed being, in most cases, quite
sufficient.
The physical properties of the substance under examination are to be first
noted; such as color, structure, odor, lustre, density, etc.
Heat in Olass Tube Closed at One End.
§294. The substance, in fragments or in the form of a powder, is introduced
into a small glass tube, sealed at one end, or into a small matrass, and heat
applied gently, gradually raising it to redness, if necessary with the aid of the
blow-pipe. When the substance is in the form of a powder it is more easily
introduced into the tube by placing the powder in a narrow strip of papier,
folded lengthwise in the shape of a trough; the paper is now inserted into the
tube held horizontally, the whole brought to a vertical position, and the paper
withdrawn; in this way the powder is all deposited at the bottom of the tube.
By this treatment in the glass tube we are first to notice whether the sub-
stance undergoes a change, and whether this change occurs with or without
decomposition. The sublimates, which may be formed in the upper part of the
tube, are especially to be noted. Escaping gases or vapors should be tested as
to their alkalinity or acidity, by small strips of moist red and blue litmus
paper inserted in the neck of the tube.
Heat in Olass Tube Open at Both Ends.
§295. The substance is inserted into a glass tube from two to three inches
long, about one inch from the end, at which point a bend is sometimes made:
heat is applied gently at first, the force of the air-current passing through the
tube being regulated by inclining the tube at different angles. Many sub-
stances undergoing no change in the closed tube absorb oxygen and yield
volatile acids or metallic oxides. As in the previous case, the nature of the
sublimate and the odor of the escaping gas are particularly to be noted. ^ The
reactions of sulphur, arsenic, antimony and selenium are very characteristic;
these metals, if present, are generally easily detected in this way (§69, 7).
Heat in Blow-pipe Flame on Charcoal.
§296. For this test, a well-burned piece of charcoal is selected, and a small
cavity made in that side of the coal showing the annular rings; a small frag-
ment of the substance is placed in the cavity, and, if the substance be a
powder, it may be moistened with a drop of water. The coal is held horizont-
ally, and the fiame made to play upon the assay at an angle of about twenty-
five degrees. The substance is brought to a moderate heat, and finally to
intense ignition. Any escaping gases are to be tested for their odor; the
§300. PRELIMiyARY EXAMINATIONS OF 80LID8. 365^
change of color which the substance undergoes, and the nature and color of
the coating which may form near the assay, are also to be carefully noted,
sjome substances, as lead, may be detected at once by the nature of the coating.
Ignition of the Substance previously Moistened with a Drop of Cobalt
Nitrate.
§297. This test may be effected either by heating on charcoal, in the loop of
platinum wire, or in the platinum-pointed forceps. A portion of the substance
is moistened with a drop of the reagent, and exposed to the action of the outer
flame. When the substance is in fragments, and porous enough to absorb the
cobalt solution, it may be held in the platinum-pointed forceps and ignited.
The color is to be noted after fusion. This test is rather limited; aluminum,
zinc and magnesium giving the most characteristic reactions.
Fusion with Sodium Carbonate on Charcoal.
§298. The powdered substance to be tested is mixed with sodium carbonate,,
moistened and placed in the cavity of the coal. Some substances form, with
sodium carbonate at a high heat, fusible compounds; others infusible. Many
bodies, as silicates, require fusion with alkali carbonate before they can be
tested in the wet way. Many metallic oxides are reduced to metal, forming
globules, which may be easily detected.
When this test is applied for the detection of sulphates and sulphides, the
flame of the alcohol lamp is to be substituted for that of the gas-flame, aa
the latter generally contains sulphur compounds.
Examination of the Color which may be imparted to the Outer Flame.
§299. In this way many substances may be definitely detected. The test may
be applied either on charcoal or on the loop of platinum wire, preferably in the
latter way. When the substance will admit a small fragment is placed in the
loop of the platinum wire, or held in the platinum-pointed forceps, and the
point of the blue flame directed upon it. If the substance is in a powder it may
be made into a paste with a drop of water, and placed in the cavity of the
charcoal, the flame being directed horizontally across the coal. The color
which the substance imparts to the outer flame in either case is noted. In
most cases the flame of the Bunsen burner alone will suffice; the substance
being heated in the loop of platinum wire, which, in all cases, should be first
dipped in hj'drochloric acid and ignited, in order to secure against the presence
of foreign substances. Those salts which are more volatile at the temperature
of the flame, as a rule give the most intense coloration. When two or more
substances are found together it is sometimes the case that one of them masks
the color of all the others: the bright yellow flame of sodium, when present in
excess, generally veiling the flame of the other elements. In order to obviate
this, colored media, as cobalt-blue glass, indigo solution, etc., are interposed
between the flame and the eye of the observer. The appearance of the flame
of various bodies, when viewed through these media, enables us often to detect
very small quantities of them in the presence of large quantities of other
substances.
Treatment of the Substance with Borax and Hicrocosmic Salt.
§300. This is best effected in the loop of platinum wire. This is heated and
dipped into the borax or microcosmic salt and heated to a colorless bead; a
small quantity of the substance under examination is now brought in contact
with the hot bead, and heated, in both the oxidizing and reducing flames. Any
reaction which takes place during the heating must be noticed; most of the
metallic oxides are dissolved in the bead, and form a colored glass, the color
of which is to be observed, both while hot and cold. The color of the bead
varies in intensity, according to the amount of the substance used; a very
366 CONVERSIOy OF SOLIDS INTO LIQUIDS, §801.
small quantity will, in most cases, suffice. Certain bodies, as the alkaline
earths, dissolve in borax, forming" beads which, up to a certain degree of satura-
tion, are clear. When these beads are brought into the reducing flame, and an
intermittent blast used, the\' become opaque. This operation is called flamiiig.
As reducing agents, certain metals are employed in the bead of borax or
microcosmic salt. For this purpose tin is generally chosen, lead and silttr
being taken in some cases. These metals cannot be used in the loop of plat-
inum wire, as they will alloy the platinum. The beads are first formed in the
loop of wire; then, while hot, shaken off into a porcelain dish, several being so
obtained. A number of these are now taken on charcoal and fused into a large
bead, which is charged with the substance to be tested, and then with the tin
or other metal. For this purpose tin foil (or lead foil) is previously cut in
strips half an inch wide, and the strips rolled into rods. The end of the rod
is touched to the hot bend to obtain as much of the metal as required. Lead
may be added as precipitated lead (*' proof-lead "), and silver as precipitated
silver. By aid of tin in the bead, cuprous oxide, ferrous oxide and metallic
antimony are obtained and other reductions eflPected. as directed in §77, 7,
and elsewhere.
CONVERSION OF SOLIDS INTO LIQUIDS.
§301. Before the fluid reagents can be applied, solids must be reduced to
liquids. To obtain a complete solution, the following steps must be obser\ed:
Firftt. The solid, reduced to a fine powder, is boiled in ten times its quantity
of water. Should a residue remain, it is allowed to subside, and the cleaV
liquid poured off or separated by filtration. A drop or tiro evaporated on glass,
or clean and bright platinum foil, will give a residue, if any portion has dis-
solved. If a solution is obtained, the residue, if any, is exhausted, and well
w^ashed with hot water.
F^rrnnd. The residue, insoluble in water, is digested some time with hot
hydrochloric acid. (Observe §305.) The solid, if any remain, is separated In-
filtration and washed, first with a little of this acid, then with water. The
solution, with the washings, is reserved.
Third. The well-washed residue is next digested with hot nitric acid.
Observe if there are vapors of nitrogen oxides, indicating that a metal or other
body is being oxidized. Observe if sulphur se])arates. If any residue remains
it is separated by filtration and washing, first with a little acid, then with
water, and the solution reserved.
Sometimes it does not matter which acid is used first. But if a first-group
base be present, HNO^ should be added first, for HCl would form an insoluble
chloride. If the substance contain tin (especially an alloy of tin) HNO,
would form insoluble metastannic acid, HioSUjOto . in which case HCl should
"be used first.
Fourth. Should a residue remain it is to be digested with nitrohydrochloric
acid, as directed for the other solvents.
The acid solutions are to be evaporated nearly to dryness, and then redis-
solved in water, acidulating, if necessary, to keep the substance in solution.
Fifth, Should the substance under examination prove insoluble in acids, it
is likely to be either a sulphate (of barium, strontium or lead); a chloride, or
bromide, of silver or lead: a silicate or fluoride— perhaps decomposed by sul-
phuric acid — and it vniM he fvf<ed irith a fixed alkali carbonate, when the con-
stituents are transposed in such manner as to render them soluble. The
water solution of the fused mass will be found to contain the acid: the residue,
insoluble in water, the metal, now soluble in hydrochloric or nitric acids
(compare §266, 7).
If more than one solution is obtained, by the several trials with solvents,
the material contains more than one compound, and the solutions, as sepa-
rated by filtration, should be preserved separately, as above directed, and
analyzed separately. The separate results, in many cases, indicate the oiifjinal
combination of each metal.
§303. TREATMENT OP A METAL OR AN ALLOY. 367
CONVERSION OF SOLUTIONS INTO SOLIDS.
§802. Before solids in solution can be subjected to preliminary examination,
either for metals or for acids, they must be obtained in the solid state. This
is done by evaporation.
TREATMENT OF A METAL OR AN ALLOY.*
§303. On account of the different effect that nitric acid has upon the un-
combined metals, it is used as a solvent jn their detection. Thus:
Oold and platinum are not attacked by nitric acid.
Tin and antimony are oxidized and converted into compounds that are insolu-
ble both in water and an excess of the acid.
6Sb + lOHNO, = SSb.O, -f lONO + 5H,0
ISSn -h 20HNO3 + 5H,0 = 3HioSnaOift + 20NO
All the other metals are oxidized and converted into compounds that dissolve
either in water or an excess of the acid; e. g,:
3Pb 4- 8HNO3 = r>Pb(NO,), + 2N0 -f 4H,0
Bi -h 4HN0, = Bi(NO,), + NO + 2H,0
Method of Procedure.^
Place a small quantity of the metal or alloy, about equal in bulk to a pea,
having previously obtained it in as finely divided a state as possible, in an
evaporating-dish, or any suitable vessel, cover well with nitric acid, sp. (jr, 1.20,
and apply heat. Continue the application of heat, replacing from time to time
the acid lost by evaporation, until the metal or alloy is dissolved or wholly
disintegrated.
If complete solution takes place immediately, pass on to A.
If a residue remains, decant the liquid portion upon a filter; again add nitric
acid to the residue, heat, and again decant upon the same filter. Then thor-
oughly wash with hot water, either by boiling with water and decanting, or
by transferring the whole to and pouring hot water through the filter. Add
the first portions of the hot-water filtrate to the nitric acid filtrate already
obtained, and treat the mixture as directed in A, after having first evaporated
a drop or two on platinum foil, to ascertain whether anything has really
been dissolved.
Treat the residue as directed in B,
A. — The Nitric Acid Solution.
This solution may contain any of the metals, except those mentioned under B.
If the nitric acid has effected a whole or partial solution of the original
metal or alloy, evaporate almost to dryness to remove excess of acid, add about
ten times its bulk of water, and proceed with the separation and detection of
the metals in the regular way.
Should the concentrated liquid become turbid when diluted with water, the
presence of bismuth is indicated. In this case enough acid must be added to
clear up the solution.^
• This Beotion is furnished by Dr. J. W. Baird, Dean of the Massachusetts Collet of
Pharmacy.
t When gold or platinum constitutes more than one-quarter of the alloy, nltrlo add failo tn
extract the whole of the base metals that otherwise are readily soluble. In such a case the
amount of erold or platinum must be reduced to at least 26 per cent, by fusing the alloy with
the requisite amount of that base metal whose absence Is surely known.
t Arsenic, if present in the origrinal alloy, now exists in the form of arsenic acid, the precipi-
tation of which requires heat and long-continued passage of H,S ($69, 6' e 7)*
3G8 SEPARATIOX OF ACIDS FROM BASES. §308, J?.
B. — The Residue Insoluble in Nitric Acid.
' This may contain gold and platinum in their metallic forms, and tin * and
antimony * in the form of metastannic and antimonic acid&. The separation
of the two former from the two latter depends upon the fact that the meta-
stannic and antimonic acids are soluble in hydrochloric acid, forming^ SnCl^
and SbClft .
Digest, therefore, the well-washed residue in concentrated hydrochloric acid
at a boiling temperature for from 5 to 10 minutes; then add at once an equal
volume of water (to dissolve the stannic chloride), and bring to the boiling
point.
If gold or platinum existed in the original metal or alloy it will now be
found in the form of a dark-brown or black powder or mass, insoluble in the
hydrochloric acid. If such a residue exists, decant trhile hot, again add hydro-
chloric acid, heat, and again decant.
The Hydrochloric Acid Solution.
This solution may have a turbid appearance, especially when cold, due to the
action of the water upon the SbClj: but without filtering proceed with the
separation and detection of the tin and antimony by the usual process.f
The Dark-colored Residue .
Add, after washing, two volumes of hydrochloric and one of nitric acid:
evaporate almost or quite to dryness, dissolve in a small quantity of water
(to obtain a concentrated solution), and divide into two portions.
The gold and platinum have been dissolved by the aqua-regia formed, and
now exist as auric and platinic chlorides.
First Portion— Test for Gold.
Dilute with at least ten times its bulk of water; add a drop or two of a mix-
ture of stannous and stannic chlorides; a purple or brownish-red precipitate
(or coloration), purple of Cassius, constitutes the test for gold.
A convenient way of preparing this mixture of stannous and stannic chlorides
is to
{a) Add a few drops of chlorine-water to a solution of stannous chloride; or
{h) Add to a small quantity of stannous chloride enough ferric chloride to
produce a faint coloration.
Second Portion — Test for Platinum.
Add, w^ithout dilution, an equal volume of a strong solution of ammonium
chloride. The formation, either at first or on standing, of a lemon-yellow
crystalline precipitate, consisting of the double chloride of platinum and
ammonium, (NH4Cl)2PtCl< , constitutes the test for platinum.
Addition of alcohol favors the precipitation.
If the proportion of platinum is very small, the mixture, after ammonium
chloride has been added, should be evaporated to dryness on a water-bath and
the residue treated with dilute alcohol. The ammonium platinic chloride
remains behind as a yellow crystalline powder.
SEPARATIOX OF THE ACIDS FROM THE BASES.
§304. The preliminary examination of the solid material in the dry way will
give indications drawing attention to certain acids. Solutions can be evapo-
rated to obtain a residue for this examination. Thus, detonation (not the
• Traces may sometimes be dissolved.
t Arsenic must be looked for in this as well as in the nitric acid solution. For when the alloy
contains arsenic, part of it will combine with the antimony and tin, and be held in the residue.
§309. SEPARATION QF ACIDS FROM BASES. 369
<lecrepitation caused by water in crystals) indicates chlorates, nitrates, bro-
mates, iodates. Explosion or deftayration will occur if these, or other oxygen-
furnishing' salts — as permanganates, chromates — are in mixture with easily
combustible matter (§273, 7). Hypophosphites, heated alone, deflagrate in-
tensely. A broiaiish-yellow rapar indicates nitrates or nitrites (§241, 7); a
tfrefH flame, borates (§221, 7). The odor of burning sulphur: sulphides, sulphites,
thiosulphates, or free sulphur. The separation of carbon black: an organic acid.
The formation of a silver stain: a sulphur compound (§266, 7).
§305. When dissolving a solid by acids for work in the wet way, indications
of the more volatile acids will be obtained. Sudden effervescence: a carbonate
(oxalate or cyanate, §228, 6). Green ish-yellmc vapors: a chlorate (§272).
Broirnish-ycllow, chlornitrous vapors on addition of hydrochloric acid: a nitrate.
The characteristic odors: salts of hydrosulphuric acid, sulphurous acid, hydro-
bromic acid, hydriodic acid, hydrocyanic acid, acetic acid. The separation of
sulphur: a higher sulphide, etc. It will be remembered that chlorine results
from action of manganese dioxide, and numerous oxidizing agents, upon
hydrochloric acid.
*^§306. If the material is in solution, the bases will be first determined.
(Certain volatile acids will be detected in the first-group acidulation — by indica-
tions mentioned in the preceding paragraph.) Now, it should first be con-
sidered, what acids can be present in solution with the bases found f Thus, if
barium be among the bases, we need not look for sulphuric add, nor, in a
solution not acid, for phosphoric acid.
§307. As a general rule, the non-alkali metals must be removed from a
solution before testing it for acids, unless it can be clearly seen that they will
not interfere with the tests to be made.
Metals need to be removed: because, firstly, in the testing for acids by precipi-
tation^ a precipitate may be obtained from the action of the reagent on the
base of the solution tested, thus: if the solution contain silver, we cannot test
it for sulphuric acid by use of barium chloride (and we are restricted to use
of barium nitrate). And, secondly, in testing for acids by transposition icith a
stronger acid — the preliminary examination for acids — certain bases do not
permit transposition. Thus, chlorides, etc., of lead, silver, mercury, tin and
antimony, and sulphide of arsenic, are not transposed by sulphuric acid, or
not promptly.
§308. If neither arsenic nor antimony is among the bases, they may all be
removed by boiling with slight excess of sodium or potassium carbonate, and
filtering. Arsetiic and antimony, and all other bases of the second group, may
be removed by warming with hydrosulphuric acid, and filtering. When the
bases are removed by sodium or potassium carbonate, the filtrate must be
exactly neutralized by nitric acid, with the expulsion of all carbonic acid by boil-
ing. Then, for nitric acid, the original substance may be tested.
§309. The separation of phosphoric acid from bases is a part of the w^ork
of the third group of metals, and is explained in §§152 and 153. For removal
of boric acid, see §221; oxalic acid, §151; and silicic acid, §249, 6 and 8.
The non-volatile cyanogen acids can be separated from bases by digesting
with potassium or sodium hydroxide (not too strong, §§231 and 232), adding
potassium or sodium carbonate and digesting, and then filtering. The residue
is examined for bases, by the usual systematic process. The solution will
contain the alkali salts of the cyanogen acids, and may contain metals whose
hydroxides or carbonates are soluble in fixed alkali hydroxides.
370
PREUMISARY EXAMINATION OP SOLIDS.
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PRELIMINART EXAMIXATION OF SOLIDS.
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PRELIMINARY EXAMINATION OF 80LID8.
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SVBSTAyCES BEFORE THE BLOW-PIPE.
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OROUPliXa OF THE METALS,
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378
AVWS. FIRST TABLE.
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ACIDS. FIRST TABLE.
379
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380
ACIDS. FIRST TABLE.
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ACIDS, FIRST TABLE.
381
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382
ACJDS. FIRST TABLE.
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ACIDS. FIRST TABLE. 383
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384 ACIDS, FIRST TABLE. §814.
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ACIDH. FIRST TABLE.
385
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iClDS. SECOKD TABLE.
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ACIDS. THIRD TABLE.
387
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388
ACID8. FOURTH TABLE.
§317, 1.
§317. Table for Identification and Separation of the Commonly
Occurring Anions (Acids).*
1. Boil the material with dilute HNO, . There results:
CO
Effervescence; turbidity in a drop of lime-water.
Effervescence, penetrating odor.
Effervescence, red-brown fumes, odor.
Odor, blackening of paper moistened with lead acetate, separa-
tion of sulphur in the solution.
Odor ) Often masked by the others; see special tests
Vinegar odor j below.
2. Boil with concentrated Na^COg solution; all cathions (bases) except
the alkalis are precipitated as carbonates or hydroxides and removed by
filtration. The filtrate contains all the anions (acids) and the excess of
COa" . Acidulation with HNO3 sets free CO2, and SiOj is precipitated;
SO2
HCN
identified in the microcosmic salt bead.
3. Ca(N03)2 solution precipitates:
F' as CaPg 1 insoluble in
V in acetic
as CslC^O^ j acid;
as Ca(CN)2 \
as CaC^H^Oo
The filtrate is made ammoniacal.
CN'
dilute
HCl
{
CAOe
HAsO o" as CaHAsO,
HAsO/ as CaHAsO,
HPO/ as CaHPO^
o
C3
insoluble; H2SO4 liberates HF .
soluble, reappearing with NH3;
decolors EMnO^ solution,
/heated with Fe" + Fe"* + OH' gives
Prussian blue on acidifying,
with K' ions in concentrated solution po-
tassium bitartrate precipitated.
In the filtrate from the above,
HjS precipitates ASgSa at once in the cold.
In the filtrate from the above,
HoS slowly precipitates from hot solution
S2 + A8,S3 .
In the filtrate from the above,
ammonium molybdate gives yellow pre-
/ cipitate; or Mg*' + NH/ -f OH' gives
/ \ MgNH.FO,.
4. In the filtrate from 3. Ba(N03)2 precipitates:
CrO/'(Cr20/') as BaCrO^, yellow, soluble in HCl; the yellow color of the
solution becoming green on boiling with alcohol.
• Prom Chem, Prakt. Abegg and Here (1900), Breslau, Page 118 ; reviewed by Freaenlufi. Z^
1900, 89, 566.
§318, 2. NOTES ON THE DETECTION OF ACIDS. 389
SO J' as BaSO^ \ / unchanged, remains insoluble in
HCl.
SiPe'' as BaSiFe \y insoluble on / gives oflE SiP^ , which deposits
in HCl ; ignition V SiOj in a drop of water; the
I residue, BaFj, is soluble in
\^ HCl.
5. The filtrate from 4. is exactly neutralized with HNO3*; Zn(N03)2
then precipitates:
Pe(CN)/" as Zn3[Fe(CN)e]2 brownish-yellow \ dissolved f brown
/ by OH' / coloration.
Fe(CN)o"" as Zn2Fc(CN)6 white / gives with \ Prussian
] Fe*" and H* ( blue.
G. A few drops of the filtrate from 5. are treated with as little Fc"* as
possible:
Red f Fe(CNS)3 ) on j permanent red color,
coloration | Fe(C2H302)8 ) heating | precipitate and colorless solution.
In the absence of CNS' another drop is tested with Ag' for the halogens;
if a precipitate results or if CNS' is present, 07ie part of the solution is
treated with CS2 and a little Cl-water:
r violet coloration, disappears with ) , p. ,
Br' brown coloration, does not disappear with J
The second portion is evaporated to dryness with KoCroO^ , fused, and
the mass after cooling distilled with- concentrated HgSO^; appearance of
oily brown drops of CrOgClg , forming Cr04" with water: CI' .
7. A concentrated water-extract of the original substance is treated
with concentrated H^SO^ and solid FeSO^ or Fe" solution, prepared cold;
a l)rown coloration shows the presence of NO3'.
The anions mentioned above to some extent exclude one another, being
unstable when together in solution owing to their power of mutual oxida-
tion and reduction, e. g,, SO3" and' S"; SO3" and NO./; NO2' and CN';
NO2' and S"; NO2' and I'; NO.' and HA8O3"; S" and HAsO/' , etc. It is
to be noticed that this always simplifies the analytical procedure.
§318. Notes ox the Detection of Acids.
7. The precipitation of tartrates by calcium salts is incomplete; from
calcium sulphate solution a precipitate forms slowly or not at all. Calcium
tartrate is soluble in the cold in a solution of KOH , precipitating" gelatinous
on hoiling; again soluble on cooling (separation from citrate). Calcium tartrate
is soluble in acetic acid (separation from oxalate).
2, A number of basic carboxiates give almost no efPervescence when treated
•In the orlfirlnal German t«xt it Is directed to use HCl at this point.
390 NOTES ON THE DETECTION OF ACIDS. §318, S.
with acids. To detect the presence of small amounts of carbonate, it is recommended
to place the dry powder' in a test-tube and fill about three-fourths full of
distilled water. Close the test-tube with a two-holed rubber stopper contain-
ing a thistle tube reaching nearly to the bottom of the test tube, and a
delivery tube reaching just through the stopper. Add dilute sulphuric acid
and warm gently. The carbonate is decomposed, driven from the solution,
and, owing to the limited air space, readily passes through the delivery tube
into the solution of calcium hydroxide.
3. With the generation of an abundance of CO2 , the precipitate first formed
in the CaCOH), is redissolved (solution of lime in spring water). Boiling
drives off the excess of CO, and causes the reprecipitation of the CaOO, .
Barium hj'droxide may be used instead of calcium hydroxide.
4. If compounds have been strongly ignited previous to solution for analysis,
oxalates cannot be present.
5. In Table H (§315), if strong oxidizing agents are present, as XCIO, .
KsCrsOr , KM11O4 , etc., the oxalic acid will be decomposed on warming with
hydrochloric acid. This may be avoided by adding calcium chloride to the
solution, neutral or alkaline with ammonium hydroxide. The oxalate will be
precipitated and thus separated from the oxidizing agents. After filtering,
the precipitate is digested with dilute acetic acid, filtered and the filtrate
tested for phosphate with ammonium molybdate. The residue is dissolved in
hydrochloric acid, filtered if necessary (calcium sulphate does not dissolve
readily), and the filtrate made alkaline with ammonium hydroxide. The pre-
cipitate thus obtained is washed, dissolved in nitric acid and tested with
potassium permanganate. The filtrate from the solution after the addition
of calcium chloride is acidified with hj'drochloric acid, heated to boiling and
tested for sulphate by the addition of a few drops of barium chloride (§317).
6. In Table H, if sulphites or thiosulphates are present, the solution in
hydrochloric acid must be heated sufficiently to drive off all the eulpharoiis
anhydride, or reactions for oxalates will be obtained, due to the sulphurous
acid alone. If there be any doubt as to the complete removal of the sulphur-
ous anhydride, the gas evolved by the reaction of the potassium perman-
ganate should be passed into a solution of calcium hydroxide. A precipitate of
calcium carbonate at this point is positive evidence of the previous presence
of oxalic acid or oxalates.
7. Alkali ferro- and ferricyanides are separated from each other by the
solubility of the latter in alcohol.
8. In testing for nitric acid the student must not be content with good
results from one test. At least four tests should be made, and all of them
should give positive results before final affirmative judgment is passed. Failun
to bleach indigo solution in the presence of an cj-crss of hydrochloric acid may be
taken as conclusive evidence of the absence of nitrates,
0. In the analysis of minerals, silica or silicates will usually be present.
The silica should be removed before proceeding with the analysis. Fuse the
finely divided material with an excess of sodium carbonate, digest the cooled
mass thoroughly in hot water, filter and evaporate the filtrate to dryness.
Moisten the residue with concentrated hydrochloric acid, and again evaporate
to dryness. Pulverize thoroughly, digest in water acidulated with hydro-
chloric acid and filter. The residue, white, consists of the silica, SiOj .
10. Meta- or pyrophosphates do not react promptly with ammonium molyb-
date. In the usual course of analysis thev are changed to the orthophosphate
(§255, GA).
11. Phosphoric acid may be detected in the presence of arsenic acid bj
ammonium molybdate if the solution be kept cold: it is preferable to remove
the arsenic before testing. In absence of interfering substances the color of
the silver nitrate precipitate will indicate the presence or absence of arsenic
acid (§69. r>/). See also note 20.
12. Sulphides which are transposed hi/ hydrochloric acid are best detected by
the odor of the evolved gas, and by passing the evolved gas into ammonium
hydroxide and testing with sodium nitroferri cyanide. Other sulphides are
decomposed by nitric acid or by nitrohydrochloric acid with separation of
sulphur as a leather^' mass or as a yellow precipitate. Persistent heating of
§318, 19, NOTES OA' TEE DETECTION OF ACIDS, 391
the sulphur with the reagent decomposing the sulphide will cause the oxida-
tion of a portion of the sulphur to a sulphate which may be detected in the
usual manner. A portion of the precipitated sulphur should be burned on a
platinum foil with careful observance of the odor of the evolved gas.
13, A sulphite (or other lower oxidized compound of sulphur) is readily
detected by adding barium chloride in excess to a portion of the solution,
dissolving in HCl , filtering if residue remains, and adding bromine or chlorine
to the clear filtrate. A precipitation of barium sulphate indicates the oxidation
of a lower compound of sulphur to a sulphate.
14- It frequently becomes necessary to detect and estimate sulphides, thio-
sulphates, sulphites and sulphates in mixtures containing two or more of the
compounds. The method of procedure will vary according to the nature of
the substance. The student will be aided by studying §§257, 8; 258, 8; and
265, 8.
15. The recognitioxi of chlorides, bromides and iodides — by evolving their
chlorine, bromine and iodine, in presence of each other — can be accomplished as
follows — for the iodine the test being very easy; for chlorine, indirect but
unmistakable: for bromine, dependent upon much care and discretion.*
The iodine is liberated with dilute chlorine-water, added drop by drop, and
is readily detected by starch, or carbon disulphide, according to §280, 8. (As
to interference of thiocyanates, see §234.) The chlorine is vaporized (from
another portion) as chlorochromic m\h}fdrule, and the latter identified by its
color and its various products, as described in §269, M, Before the bromine
is identified the iodine is to be either removed as free iodine, or oxidized to iodnte
(§276, 86). The oxidation to iodic acid is efTected as follows: Treat with
chlorine-water till free iodine no longer shows its color; add a drop or two
more of the chlorine-water, and dilute with water, keeping cool: then add the
carbon disulphide, agitate and leave the solvent to settle, for the yellow color
of bromine. The removal of free iodine may be accomplished as follows: Add
chlorine-water, drop by drop, as long as the iodine tint seems to deepen by
the addition: add the carbon disulphide, agitate, allow to subside, and remove
the lower layer, either by taking it out with a pipette or by filtration through
a wet filter. Repeat, if need be, till iodine color is no longer obtained; then
continue, with dilute chlorine-water, in test for bromine.
If iodide in large proportion is to be removed, it is well, first, to precipitate
it out, as far as possible, by copper sulphate and a reducing agent (Note 17).
The filtrate is then to be treated by either method above given.
16. The separation by ammonium hydroxide, as a solvent of the silver pre-
cipitates — ^AgCl , AgBr, Agl — when conducted with dilute ammonium hydrox-
ide, may be made complete between the chloride and the iodide, also between
the bromide and the iodide, but it is very imperfect between the bromide and
the chloride. The hot and strong solution of ammonium acid carbonate
separates the chloride from the bromide (§269. 8(7).
17. The direct removal of iodides hy precipitation, leaving bromides and chlorides
in solution, can be effected (approximately) by copper sulphate with sulphurous
acid (§77, 6f), or quite completely by palladous chloride (§106, 6).
IS, Chloric acid is separated from hydrochloric and all other acids of chlorine,
bromine and iodine (except from hj-pochlorous acid, and from traces of bromic
acid), by remaining in solution during the precipitation by silver nitrate
(§273, 5).
19, Chloric acid is separated from nitric acid — after finding that silver nitrate
(fives no precipitate in another portion of the solution, acidulated — by evaporat-
ing and igniting the residue, then dissolving and testing one portion of the
solution by silver nitrate for the chloride formed from chlorate during igni-
tion (§273, 7). The other portion of the solution is tested for nitric or nitrous
acid.
* Tn consequence of the relative commercial values of bromine and iodine, nnd the medioinal
relatlODS of bromides and iodides, it is of great importance to search commeroial iodides for
intentional and considerable mixtures of bromides— an impurity likely to escape cursory
chemical examination. There are, however, very slight and usually unobjectionable propor-
tions of bromides frequently to be found in the iodides of commerce, and occurringr from the
difficulty of exact separation in the manufacture of iodine from kelp.
392 yOTES Oy THE DETECTION OF ACID^. §818, 20
20, If we have to separate chloric acid boih from nitric and hydrochloric acidt^
a solution of silver sulphate must be used instead of the nitrate, to precipitate
out all the hydrochloric acid. The filtrate from this is evaporated, igpiiitad,
dissolved and tested for silver chloride, indicating chlorate in the original
solution, and another portion is tested for nitric acid. Also, chlorates are
distinguished (not separated) from nitrates, by oxidation of ferrous sulphate
in solution with acetic acid on heating, and the consequent formation of the
red solution of feiric acetate (§§126, (W>; 162; 223, 6). The solution tested must
contain no free acids, and no nitrites or other oxidizing agents beside the two
in question, but may contain chlorides; and*, of course, the ferrous sulphate
must be pure enough not to color when heated alone with the acetic acid.
Mix the ferrous sulphate solution with the acetic acid, boil, then add the solu-
tion to be tested, and heat nearly to boiling, for some minutes. If no red
color appears, chlorates are absent, and nitrates may be present.
21. Hypochlorites are separated with chlorates from' chlorides (bromides), etc.
by silver nitrate; and distinguished from chlorates (in the filtrate from AgCl,
etc.) by bleaching litmus, and by their much more rapid decomposition and
consequent precipitation of any silver in solution. They are also more active
than chlorates, as oxidizing agents.
22, M. Dechan's method (§269, Si) consists (/) in boiling the mixture with a
solution of ^0 grammes of KzCt^O-i , dissolved in 100 cc. of water, which lib-
erates and expels all of the iodine without disturbing the bromine and chlorine,
oKjCFaOx + OKI = Cr^O, -f SKjCrO* + 31,
(2) 8 cc. of a dilute solution of sulphuric acid (consisting of equal volumes of
HiSO^ sp, (jr. 1.84, and water) are added to 100 cc. of the dichromate solution,
and on boiling, the bromine is distilled off without disturbing the chlorine;
after which the chlorine is detected in the usual manner. For other methods
of detecting chlorides in presence of bromides and iodides, see §269, 8.
23. For A. Longi's process for the analysis of a mixture of chlorides, bro-
mides, iodides, chlorates, bromates, iodates, ferrocyanides and ferricvanides,
see C. y.j 18s:^ 47, 200.
2'i. In the detection of chlorides in presence of cyanides and thiocyanates
by the decomposition of the silver salts with concentrated sulphuric acid
(§269, Sr), a drop or two of silver nitrate should be added to the precipitatf
before heating with the acid or a black precipitate will be obtained, apparently
carbon.
2'). For the detection of a bromide in the presence of an iodide, the most
satisfactory method is by the use of potassium chlorate and dilute sulphuric
acid as described in §276, 8r. The student should carefully note the change
in color of the solution. The first very dark color is due to the liberation of
iodine. There is usually a sudden change of color on the complete oxidation
of the iodine, but if much bromine be present the solution will be quite dark
straw color. This should be tested with carbon disulphide and the heatinp
continued if free iodine is still present.
26. Arsenic acid should not be present when testing for a phosphate. If the
arsenic acid be reduced to arsenous acid by sulphurous acid it will not interfere
with the ammonium niolybdate test for a phosphate. The excess of sulphur-
ous acid must be removed by boiling before testing for the phosphate. Arsenic
is best removed by precipitation as sulphide in the second group.
27. Chromic acid is always identified by reduction and precipitation as
chromic hydroxide, green, in the third group. The red or yellow^ color to the
original substance usually gives evidence of the probable presence of the
hexad chromium. The reduction is effected in the course of analysis by hydro-
sulphuric acid with precipitation of sulphur. It is advisable to reduce all
chromates by warming with hydrochloric acid and alcohol before proceeding
with the analysis. Another portion of the substance may be reduced with
sulphuric acid and alcohol and tested for chlorides.
2S. Manganates are readily decomposed by water with formation of XKnO,
and MnOa . In the presence of a fixed alkali the manganate solution is green
and does not rapidly change to permanganate. The manganates and perman-
ganates in solution are all dark colored (green, purple-red) and should be
reduced by warming with hydrochloric acid before proceeding with the
analysis.
§319, 7. PHINCIPLES. 393
§319. PRINCIPLES.
In the following statements, the term salt includes only eases where
the metal acts as a base, e. g., chromium salts include CrClg , not KjCrO^ .
Only salts of ordinary metals are included.
1. Hydroxides when brought in contact with acids form salts, provided
they can be formed by any means in the presence of water. The same
is true of oxides. But AlgSj and CrgSj are not formed in presence of
water. (Some oxides after ignition fail to unite with all acids, e. g., SnOj ,
FCaOa, AI2O3, but by long boiling unite with a few acids; while ignited
CTjOs is insoluble in all acids).
2. All nitrates, chlorates and acetates are soluble, but salts of cuprosum,
bismuth, tin, antimony and the oxysalts of mercury require some free acid
to hold them in solution.
3. All oxides and hydroxides are insoluble, except those of the alkalis^
those of Ba, Sr and Ca slightly soluble. The fixed alkalis precipitate
solutions of all other metallic salts, Ba, Sr and Ca incompletely. The
precipitate with silver, antimonosum and mercury is an oxide, with Sn'^
it is Sn0(0H)2 , with Sb^ Sb0(0H)3 , in all other cases a normal hydroxide.
[At boiling heat instead of normal hydroxides other hydroxides are some-
times formed, e, g., 'Etfi^{0'R)Q, and Cu302(0H)ol. This precipitate re-
dissolves in eight cases, forming, if potassium hydroxide be used . , .
KjPbOo , KgSnOg , K2Sn03 , KSbOo , KSbOg , EoZnO. , E:A102 , KCr02 . The
last precipitates on boiling.
4. Ammoninm hydroxide precipitates solutions of the first four groups,
manganese and magnesium imperfectly and not at all if ammonium
chloride be present. The precipitate is a normal hydroxide, except that
with Sn^v it is Sn0(0H)2, with Sb^ SbO(OH),, with Ag and SV" the
oxide, with Pb a basic salt, and* with Hg a substituted mercvrtc ammonium
compound, Hg' in addition forms Hg°. The precipitate redissolves in six
cases, viz., silver, copper, cadmium, cobalt, nickel and zinc. With silver,
inS^AgO is formed, with zinc (NH4)2Zn02 .
5. The chlorides of the first group are insoluble, lead chloride slightly
soluble. Hydrochloric acid and soluble chlorides precipitate solutions of
salts of the first group, lead salts incompletely, and normal lead salts are
not precipitated by mercuric chloride. (For action on higher oxides, etc.,
see §269, 6A).
6. The bromides of the first group are insoluble, lead bromide slightly
soluble (less than the chloride). Hydrobromic acid and soluble bromides
precipitate solutions of the salts of the first group, lead salts incompletely.
(For action on higher oxides, etc., see §276, 6^4).
7. The iodides of lead, silver, mercun^ and ouprosum are insoluble.
Hydriodic acid and soluble iodides precipitate solutions of lead, silver.
394 PRINCIPLES, §319, 8.
mercury and cuprosum. Cupric salts are precipitated as cuprous iodide
with liberation of iodine. Ferric salts are merely reduced to ferrous
salts with liberation of iodine. Arsenic acid is merely reduced to arsen-
ous acid with liberation of iodine.
(Bismuth, stannous and untimonous iodides are reany insoluble in water, and
are readily formed by the action of hydriodic acid or soluble iodides on the dry
or merely moistened salts. But the dissolved salts of these three metals fre-
quently contain so much free acid that it prevents their precipitation by
hydriodic acid or by soluble iodides. Arsenous iodide is decomposed by water.
It may be formed from the chloride, best by adding hydriodic acid or a soluble
iodide to a solution of arsenous acid in strong hydrochloric acid. Bismuth
iodide is black; stannous, antimonous and arsenous iodides are yellowish red.)
8. The sulphates of lead, mercurosum, barium, strontium and calcium are
insoluble, those of calcium and mercurosum slightly soluble. Sulplmri?
acid and soluble sulphates precipitate solutions of lead, mercurosuir.,
barium, strontium and calcium; calcium and mercurosum incompletely.
9. (a) The sulphides of the first four groups are insoluble. Hydro-
sulphurio acid transposes salts of the first two groups in acid, neutral,
and alkaline mixtures, except arsenic, which is generally imperfectly
precipitated unless some free acid or salt that is not alkaline to litmus
paper be present. The result is a sulphide, but mercurosum forms mer-
curic sulphide and mercury, and arsenic acid forms arsenous sulphide and
free sulphur. Ferric solutions are reduced to ferrous with liberation of
sulphur. In acid mixture other third and fourth group salts are not
disturbed, but from solutions of their normal salts traces of cobalt, nickel,
manganese, and zinc are precipitated. (For action on higher oxides, seo
§257,6.4)..
(&) Soluble sulphides transpose salts of the first four groups. The
result is a sulphide, except that with aluminum and chromium salts it is
a hydroxide, hydrosulphiiric acid being evolved. With mercurous salt5»
mercuric sulphide and mercury are formed; with ferric salts, ferrous
sulphide and sulphur.
10. The carbonates of the alkalis are soluble. Carbonates of the fixed
alkalis precipitate solutions of all other metallic salts. The precipitate is:
a. An oxide with antimonous salts.
&. A normal hydroxide with Sn", Al , Cr'" and Fe'"; with Sn^^, SnO(OH)-> :
with Sbv SbO(OH), .
c, A normal carbonate with Ba, Sr and Ca salts and, if cold, with silver,
mercurosum, cadmium, ferrosum and manganosum.
d. A basic carbonate in other cases, except mercuric chloride, which
forms an oxychloride. Carbonic is completely displaced by strong acids,
for example, from all carbonates, by HCl , HClO^^HBr , HBrO., , HI , HIO, ,
HjCoO, , HNO, , H^PO^, HoSO^, and oven by HgS, completely from
carbonates of first four groups, incompletely from those of the fifth and
sixth groups (Nandin and :Montholon, C. r., 1870, 83, 58).
§319, 13e. PRINCIPLES. 39o
11. All normal and di-metallic orthopliospliates are insoluble except
those of the alkalis. The normal and di-metallic phosphates of the alkalis
precipitate solutions of all other salts. The precipitate is a normal, di-
metallic, or basic phosphate, except that with mercuric chloride and with
the chlorides of antimony it is not a phosphate, but an oxide, or an oxy-
ehloride.
All phosphates are dissolved, or transposed by nitric, hydrochloric and
sulphuric acids, and all are dissolved by acetic acid except lead, aluminum
and ferric phosphates. AH are soluble in phosphoric acid except those of
lead, tin, mercury and bismuth.
12. Ignition. — a. The oxides of lead and iron heated in the air to a red heat
form PbgO^ and FCgOa , but ^t a white iicat form PbO and FCgO^ . Other
oxides, if ignited in the air to a white heat, when changed, either take i^p
or lose oxygen and leave ultimately the following: Ag, Hg, An, Pt,
SnOj, SboOg, AsoO^, Bi203 , CuO, CdO, Fe304, Cr.Og, ALO., Co.O^,
HiO , MttgO^ , ZnO , BaO , SrO , CaO , MgO , K2O , HagO . In a few cases
ignition at a lower temperature gives other results, e. g,, COjO.t , BaOj ,
HEjOj , SbgO^ , etc.
b. Alkali hydroxides ignited in air at a white heat are not changed.
Other hydroxides give same as in (a).
c. Alkali carbonates ignited in air at a white heat are not changed.
Other carbonates evolve CO2 and leave as in (a).
d. Fixed alkali oxalates ignited at a white heat in absence of air are
changed to carbonates, evolving CO . Ba , Sr and Ca oxalates and a few
others at a red heat, in absence of air, form carbonates evolving CO , at
a white heat these carbonates arc changed to oxides evolving CO. . All
oxalates ignited in presence of air at a white heat are changed as in (a),
except the fixed alkali oxalates which are left as carbonates. In all cases-
when air is present the CO bums to COg .
e. All organic salts ignited at a white heat, in a current of air, leave
residues as in (a), but forming carbonates if fixed alkalis are ])resent.
The products evolved depend upon the composition of the organic por-
tion of the salt.
13. The following acids may be made by adding sulphuric acid in
excess to their respective salts and distilling:
a. Carbonic from all carbonates.
b. Nitric from all nitrates.
d. Sulphurous from all sulphites.
e. Hydrochloric from all chlorides except those of mercury. But sul-
phuric acid transposes the chlorides of Ag, Sn and Sb with extreme
difficulty, so that practically other methods are used to separate hydro-
chloric acid from these metals.
396
EQUATIONS,
§820.
§320. Equations.
It is recommended that in the class-room some attention be paid to the
balancing of equations as representing' the important analytical and synthetic
operations, especially those involving oxidation and reduction. The work will
be simplified by a careful study of §§216, 217 and 218 and application of tbc
rule as illustrated there. When the time permits, the operations represented
by the equations studied in the class-room should be performed by eich
student at his laboratory work-table. At first the teacher should select simpler
equations illustrating analytical operations and the principles (§319). Lcter,
the more difficult equations involving oxidation and reduction should be studied.
The student should give the authority for every reaction. The acconDLi>ai]yiDg
list of equations is merely suggestive and may be expanded by the teacher as
the time permits. In each equation the second substance is to be considered
as in excess; that is, sufficient to produce the greatest possible change ii the
first substance. For description and methods of making the basic salts used
in this list, see Dammer's Anorganishe Chemie.
1. Sb + HNO,
2. As« + HNO,
3. As,0, -f HNO,
4. Mn(OH), -h PbO, + HNO,
5. MnSO« + Pb,04 + H.SO^, dilute
6. MnO, -f KNO, -f- K,CO, , fusion
7. S, 4- KNO, -f KjCOg , fusion
8. MnS -h KNOs + K3CO, , fusion
9. Mn.O,x + Pb.O, + HNO,
10. Cr(OH), -f KNOa -f K,CO, ,
fusion.
11. PbjAsOJs 4- Zn -I- H,SO, , dilute
12. Cu,As20t -f Zn -f H2SO4 , dilute
13. PbCNOJ, -f Al -f KOH
14. CuCNOa), + Al-f KOH
15. Bi(NO,), -f Al -f KOH
16. Hg,oO,(NO,), + Al -f KOH
17. MnS + Mn(N0,)2 -f K,CO, , fus.
18. Mn,0, -h Pb,0, -f HNO3
19. Fe -f HoSO« , con., hot.
20. KI -f KiO, -f H,SO, , dilute
21. MnSO^ -f- KMnO* -f H^SO* , dilute
22. (NaCl -f K^CrO, + H.SO4), dry.
hot
23. KNO, -f- Peso, -f H^SO, , con.,
cold
24. K,Cr,0(CrO,), -f HCl , hot
25. Hg,0(N0,)« -f- Al -h KOH
26. Ag,AsO« -I- SnCl, + HCl , sp. qr.
1.18
27. PbO^ -f KX^O, -f- H,SO, , dilute
28. Pb,04 , white heat
29. NaH.POz , ignition
30. re,Oo(AsO,), -f- Pes + HCl
31. FeBr^ -f HNO,
32. Sn -f HNO, , hot
33. KOH + Br, , hot
34. Pela 4- HjSO^ , sp. gr. 1.84, hot
35. KBr 4- KBrO, 4- H^SO* , dilute
36. PeSO^ 4- KMn04 4- H^SO^ , dilute
37. K.Cr^OfCrOJ, 4- KOH 4- Br,
38. 4Hg20,(N,0„), + Al 4- KOH
39. Ag, AsO, 4- SnCl- 4- HCl , sp, gr,
1.18
40. Co^On , ignition, v«^hite heat
41. H,S 4- HNO, , sp, gr, 1.42, hot
42.
43.
44.
45.
46.
47.
48.
49.
50.
51.
52.
53.
54.
55.
56.
57.
58.
59.
60.
61.
62.
63.
64.
65.
66.
67.
68.
69.
70.
Hg.CAsOJ, 4- I'eS 4- HCl
Pe,0„(A80a)2 4- KOH 4- CI,
Pel, 4- HNO, , sp. gr, 1.48, hot
Cr,(SO,). 4- Cr(NO,). 4- K,CO, ,
fusion
Pb,(As04), 4- Zn 4- H,S04 , dilute
KOH 4- CI, , cold
KBr 4- KIO, 4- HjSO, , diliite
(Cr^OHCl. 4- K,CrO, 4- HaSOJ,
dry, hot
Zn,0,(NO,), 4- PoSO, 4- HJiO, ,
concentrated, cold
Hg,(A80J, 4- SnCl, 4- ^Cl. sp.gr.
1.18
MnaO, , ignition
PejOjSO, 4- Zn 4- H^SO* , dihite
Bi,S, 4- HNO, , dilute, hot
Hg.AsO^ 4- PeS 4- HCl
Cr,(OH),SO, 4- KOH 4- CI,
PeCH.POJ, 4- HNO,
Cr.O, 4- KCIO, 4- K„CO, , fusion *
CusOjCAsOJ, 4- Zn 4- H,SO« , diL
KOH 4- CL , hot
Mn^Oi, 4- kClO, 4- K2CO, , fusion
HIO, 4- SnCl, 4- HCl
Bi,,0,,(N03)xo + Peso, 4- H,SO, ,
con., cold
CrO, , ignition
KMnO, 4- HjC.O, 4- HjSO,, dilute
PeAjsO, 4- SnCl, 4" HCl , sp. gr. I.IS
Pe,Cl, 4- Pes 4- HCl
5CUO.AS3O, 4- Pe 4- HCl
HIO, 4- HX.O, , hot
(Cr2(OH)5Cl 4- K3Cr,0T 4- H,SOJ.
dry, hot
PeCNOa), 4- Peso, 4- H,SO, , con.,
cold
Ag.SO, 4- Zn
H2SO, 4- HNO, . sp. gr, 1.42
PeAsO^ 4- Pes 4- HCl
Pb(AsO0^ 4- KOH 4- 01=
Pe(NO,), 4- HNO,
Mn.O, 4-" Mn(NO,), 4- K,CO, ,
fusion
PenOoCAsO,), 4- KOH 4- Br,
Pb,oO.(OH)e(NO,)e 4- Al 4- KOH
§321. PROBLEMS IN SYNTHESIS. 397
§321. Problems in Synthesis.
For the sake of a more thorough drill in the principles of oxidation and
other reactions, a few problems are here given; a part of them the student
should practically work at his table, but they are chiefly designed for class
exercises. Special care should be taken that a pure product be formed and
that the ingredients be taken from the sources indicated. In each case the
authority for every step in the process should be stated.
1. Silver oxide from metallic silver.
2. Mercuric bromide from mercurous chloride and sodium bromide.
3. Chromic chloride from potassium chromate and hydrochloric acid.
4. Arsenic acid from potassium arsenite.
5. Potassium arsenate from arsenous oxide and potassium hydroxide.
6. Lead nitrate from lead chloride and potassium nitrate.
7. Mercurous nitrate from mercuric chloride and sodium nitrate.
8. Mercurous oxide from mercuric oxide.
9. Mercuric bromide from metallic mercury and potassium bromide.
10. Lead nitrate from lead dioxide and potassium nitrate.
11. Lead chromate from lead hydroxide and chromium hydroxide.
12. Barium chromate from chrome alum and barium carbonate.
13. Mercuric chromate from mercuric sulphide and chromium hydroxide,
14. Chromium sulphate from potassium dichromate and zinc sulphide.
15. Phosphoric acid from sodium phosphate.
16. Phosphorus from calcium phosphate.
17. Lead iodate from sodium iodide and lead sulphide.
18. Silver iodate from silver chloride and iodine.
19. Ferric arsenate from ferrous sulphide and arsenous oxide.
20. Mercuric bromide from mercuric sulphide and sodium bromide.
21. Ammonium sulphate from ammonium chloride and sulphur.
22. Sodium chloride from commercial salt.
23. Phosphorus from sodium phosphate.
24. Lead sulphide from trilead tetroxide and ferrous sulphide.
25. Ferrous sulphate from ferric oxide and sulphuric acid.
26. Ammonium hydroxide from potassium nitrate.
27. Cadmium sulphate from cadmium phosphate and ferrous sulphide.
28. Mercurous nitrate from mercuric sulphide and nitric acid.
29. Barium sulphate from potassium thiocyanate and barium chloride.
30. Mercurous chloride from mercuric oxide and sodium chloride.
31. Sodium iodate from potassium iodate and sodium chloride.
32. Sodium phosphate from calcium phosphate and sodium chloride.
33. Strontium nitrate from sodium nitrate and strontium sulphate.
34. Potassium sulphate from potassium nitrate and sulphur.
35. Barium sulphate from barium chloride and zinc sulphide.
36. Potassium permanganate from manganese dioxide and potassium nitrate.
"37. Arsenous chloride from lead arsenate and sodium chloride.
38. Potassium chromate from potassium nitrate and lead chromate.
39. Potassium iodide from potassium chloride and iodine.
40. Barium chlorate from sodium chloride and barium nitrate.
41. Arsenous sulphide from arsine and ferrous sulphide.
42. Copper sulphate from copper sulphide.
43. Silver nitrite from silver chloride and sodium nitrate.
44. Cuprous chloride from metallic copper and sodium chloride.
45. Manganous carbonate from manganese peroxide and sodium carbonate.
46. Manganous pyrophosphate from manganese peroxide and ammonium phos-
phate.
47. Lead arsenate from lead sulphide and arsenous oxide.
48. Bismuth subnitrate from metallic bismuth and nitric acid.
49. Barium perchlorate from sodium chloride and barium hydroxide.
50. Lead iodate from metallic lead and iodine.
898 TABLE OF SOLUBfLITIES. §88L
§322. Table of Solubilities.*
Showing the classes to which the compounds of the commonly occurring elenufUs
belong in respect to their solubility in water, hydrochloric acid,
nitric acid, or aqua regia.
Preliminary Remarks.
For the sake of brevity, the classes to which the compounds belong aie
expressed in letters. These have the following signification:
W or w, soluble in water.
A or a, insoluble in water, but soluble in hydrochloric acid, nitric add,
or in aqua regia.
I or i, insoluble in water, hydrochloric acid, or nitric acid.
Further, substances standing on the border-lines are indicated as fol-
lows:
W — A or w — a, difficultly soluble in water, but soluble in hydrochloric
acid or nitric acid.
W — I or w — i, difficultly soluble in water, the solubility not being
greatly increased by the addition of acids.
A — I or a — i, insoluble in water, difficultly soluble in acids.
If the behavior of a compound to hydrochloric and nitric acids is essen-
tially different, this is stated in the notes.
Capital letters indicate common substances used in the arts and in
medicine, while the small letters are used for those less commonly occur-
ring.
The salts are generally considered as nomial, but basic and acid salts,
as well as double salts, in case they are important in medicine or in the
arts, are referred to in the notes.
The small numbers in the table refer to the following notes.
Notes to Table of Solubilities.
(1) Potassium dichroraate, W. (2) Potassium borotartrate. W. (3) Hydro-
pen potassium oxalate, W. (4) Hydropren potassium carbonate, W. (5) Hydro-
pren potassium tartrate, W. (6) Ammonium potassium tartrate, W. (7)
Sodium potassium tartrate, W. (8) Ammonium sodium phosphate, W. (9) Acid
sodium born+p W. MO) Hydrogen sodium carbonate, W. (11) Tricalcium
phosphate, A. (12) Ammonium maprnesium phosphate, A. (13) Potassium
aluminum sulphate, W. (14) Ammonium aluminum sulphate, W. (15) Potas-
sium chromium sulphate, W. (16) Zinc sulphide, as sphalerite, soluble in
nitric acid with separation of sulphur: in hydrochloric acid, only upon heatinp.
(17) Mangranese dioxide, easily soluble in hydrochloric acid: insoluble in nitric
acid. (18) Nickel sulphide is rather easily decomposed by nitric acid: very
diflficulty by hydrochloric acid. (10) Cobalt sulphide, like nickel sulphide.
(20) Ammonium ferrous sulphate, W. (21) Ammonium ferric chloride, W.
•Th** followlnir table of Bolubilities, is taken from Fresenius Qualitative Analyais, Well**
trauelation of 16th German edition.
§322. TABLE OF SOLUBILITIES. 399
(22) Potassium ferric tartrate, W. (23) Silver sulphide, only soluble in nitric
acid. (24) Minium is converted by hydrochloric acid into lead chloride; by
nitric acid, into soluble lead nitrate and brown lead peroxide which is insoluble
in nitric acid. (25) Tribasic lead acetate, W. (26) Mercurius solubilis Hahne-
mann!, A. (27) Basic mercuric sulphate, A. (28) Mercuric chloride-amide, A.
(29) Mercuric sulphide, not soluble in hydrochloric acid, nor in nitric acid, but
soluble in aqua regia upon heating. (30) Ammonium cupric sulphate, W.
(31) Copper sulphide is decomposed with difficulty by hydrochloric acid, but
easily by nitric acid. (32) Basic cupric acetate, partially soluble in water, and
completely in acids. (33) Basic bismuth chloride, A. (34) Basic bismuth
nitrate, A. (35) Sodium auric chloride, W. (36) Gold sulphide is not dissolved
by hydrochloric acid, nor by nitric acid, but it is dissolved by hot aqua regia.
(37) Potassium plantinic chloride, \V — I. (38) Ammonium platinic chloride,
W — I. (39) Platinum sulphide is not attacked by hydrochloric acid, is but
slightly attacked by boiling nitric acid (if it has been precipitated hot), but
is dissolved by hot aqua regia. (40) Ammonium stannic chloride, W. (41)
Stannous sulphide and stannic sulphide are decomposed and dissolved by hot
hydrochloric acid, and are converted by nitric acid into oxide which is insoluble
in an excess of nitric acid. Sublimed stannic sulphide is dissolved only by hot
aqua regia. (42) Antimonous oxide, soluble in hydrochloric acid, not in nitric
acid. (43) Basic antimonous chloride, A. (44) Antimony sulphide is com-
pletely dissolved by hydrochloric acid, especially upon heating; it is decom-
posed by nitric acid, but dissolved only to a slight degree. (45) Calcium
antimony sulphide, W— A. (46) Potassium antimony tartrate, W. (47) Hydro-
gen calcium malate, W.
400
TABLE OF 80LUBIUTIE8.
§S2L
SOLUBILITY
1
j
1
a
i
1
1
1
5
3
1
1
^
5
1
n
1
9^
^
Ozldu
W
w
w
w
w
W-A
A
A
AAI
A
»ii
Ctannoate^
Wi
w
w
a
w^
w-a
-w
a
w
w
Sulphate..
W'lt-ii
w
'^M^I'H
£
I
w-i
W
^IVI*
W»I,,
W
W
Phoapbate
w
w.
Wp.,1
a
a
A„
Alt
a
■
a
a
Bomt«....
w,
w.
w
a
a
a
W-a
a
a
a
D
OaalAt^*..
w.
w
W
a
a
A 1
a
a
w~«
a
'w-a
Tliiorlde.
w
w
W
W-M
w-a
A'l
a^i
w
w
W-a
a
W-*
W-«
CftrboQste
w.
w,.
W
A.
Jl
A
A
A
A
Silicate ^^^
w
w
a
a
a
a
a-l
a
a
a
Chloride -^
w,.
w«
^II'IB
W
w
W
W
w
W&I
W
W
W
Bromide . .
w
w
W
w
w
w
w
w
wftl
w
w
w
Iodide
w
w
W
w
w
w
w
w
w '
w
w r
Cj-iDlde. . .
w
w
w
w-a
w
w
w
it
A
a Ia4 -'J
Ferrocy'de
. w
w
w
w-a
w
w
w
A-I
a' iN
Fen-ioy^l^
w
w
w
w
w
a 1 , i 1
BlplitfcrMe
w
w
w
w
w
w
w
w
W|W
w w
^ulphldG..
w
W
w
w
w
WA,,
a
a
a-l
A,. A
ttlJ*'«
Nitrate.-.
w
W
w
W
W
w
w
w
W
w
w
w W
Chlorate , .
w
w
w
w
W
w
w
w
TV
w
w
w w
1
Tartrate--
^8'*»T"ai»*i
w.
w.
a
a
A
WfO
w
W
a
w-a
a
w
Citrate....
w
w
w
a
a
"W-H
w
w
W
w-a
a
w
T
Malate..-.
^■
w
w
W&H
w
w-a^.
w
w
w
w
Buccinate.i
w
w
w
w-a
nr-a
w-a
w
w-a
w-«
w
w
r-i
BeiiM>ftte,.
w
w
MV
w
w-
w
w
1
1
Salit-ylftte.
w
w
^V
w-a
W-fl
w-a
w
1
Auetate...
W
w
w
W
w
W
w
W
W
W
w
w
w
Formate..
w
w
w
W
IV
w
w
w
W
w
w
w
If
Areeulte..
W
w
w
a
a
JL
a
a
a
»
-Arsenate,.
W
W
w
a ,
a
a
a
a
a
a
a
A
i
§322.
TABLE.
TABLE OF aOLUBILITIES.
401
&
A
i
^
1
A
1
A
6
A
s
s
a
Bf
i
^
^
d
^
a
a
A„
a
a
a
a&r
A«
OxW©
w
a
A-I
a
W-a
w
a
a
a
a
Chromaie
w„
W
W-A
A-I
w-a
Wp
w„
w
W
w
w
a
Suipbate
a
A
a
a
a
a
a
a
a
a
a
w^
Pboepliaia
tt
a
a
a
a
a
w-a
a
Borate
a
a
a
a
a
a :
a 1
a
a
TV
A
w
a
Oxalate
Tr-«
w
w^
a
w-a
a
w
w-a
W
V
w
Fluoride
A
&.
A
a
a
A
a
a
Cartw>[jato
&
a
a
a
a
SQlcate
W
W,,
I
W-l
a4
w„
w
W.A„
W
w„
WiT-l»
W
w„
W A„
Chlortde
w
w
i
wA
a-i
w
w
w^
w
w
W
w-a
Hromltle
W
w
i
W-A
A
A
w
a
w
a
I
W j
w
w-a
Jqdld^j
o-l
r
a
W
a
a
W
W
CTanlde
1
I
1
a
1
t
1
I 1
w
i
w-a
i
F^may^de
w
w
i
a
A
w
a
w-a
a
w
Hulpboai'dn
A
a
«ii
A
A
An
Sji
a
A
A|«
dit
f^i
ii4i
All-IB
Sulphide
w
w
W
W
W»i
W
W
W„
w
w
Nitrate
w
w
w
w
w
w
w 1
w
w
w
Chlorate
w-i
w,.
B
a
w-a
a
w
a
W-fl
a
a4i
Tartrate
w
w
a
a
a
w-a
w
a
Citrate
w
w-a
w-a
a
w-a
w
w
W
Malate
w-»
a
a
a
a
w-a
w
w
a
Su&clDate
V
A
a
w-a
a
w-a
a
w
»
BoQZoato
Sallcylati^
w
W
w
w„
w-a
w
w*.
w
w
w
w
Aootate
w
W
w
w-a
w
w
w
w
w
w
Formats
a
a
B
a
a
a
A
.»
Arsenlt^
a
a
a
a
a
a
a
a
A
B
Arfl«DaCjt<
402 REAGEyrs. §sn.
§323. Reagents.^
During the past two years the reagents for use in qualitatiye chemical
analysis at the University of Michigan have been made up on the basis
of the normal solution; i. f., the <juantity capable of combining with one
gram of hydrogen or with its equivalent is taken in a litre for the normal
solution. Kor example: Normal potassium hydroxide, KOH , requires 56.1
grams per litre of solution (not 56. 1 grams to a litre of water), but the usual
])ure ])roduct contains about ten per cent of moisture, so it is directed to
use 62.3 grams or 312 grams for a solution five times the normal strength,
r)y. Barium chloride, BaClo.2H.O , has a molecular weight of 244.2, but
the hydrogen equivalent is (244.2-^-2) 122.1, so for a litre of half-normal
solution, X/2, take 61 grams.
In the following list of reagents, in the parenthesis immediately follow-
ing the formula are given the grams per litre necessary for a solution
of the strength indicated. Fresenius' standard follows the parenthesis.
Acid, Acetic, HCjH.O, (.?00. 5X), sp. gr, 1.04, 30 per cent acid.
Arsenic, H,AsO,.i<, H.O (15, % H^AsO^ ~ 5).
Fluosilicic, H,SiF; , §247.
Hydrobromic, HBr (40, X/2).
Hydriodic, HI (64, X/2).
Hydrochloric, HCl (182, 5X, «/). gr. 1.084), «p. gr, 1.12, 24 p. c. acid.
Hydrosulphuric, H,S . saturated aqueous solution, §257, 4.
Iodic, HIO, (15, J/o, HI0,-7- 6).
Nitric, HNO, (:n5. 5N. sp. gr. l.ir>5), np. gr. 1.2. 32 p. c. acid.
Nitrohydrochloric, about one part of concentrated HNO, to three parts
HCl .
Nitrophenic, C.HJNOj) .OH (picric acid).
Oxalic, H.C.O,.2H.O , crystals dissolved in 10 parts water.
Phosphoric, H,PO. (10, N/2).
Sulphuric, H..SO4 , concentrated. Rp. gr. 1.84.
Sulphuric, dilute (245, 5N, ftp. gr. 1.153), one part acid to five parts water.
Sulphurous, H..SO:, , saturated aqueous solution.
Tartaric, HjC.H^O^ , crystals dissolved in three parts water.
Alcohol, C,H^O , ftp. gr. 0.si5, about 05 p. c.
Aluminum Chloride, AlCl, (22. N/2).
Nitrate, Al(NO,),.7%H.O (58, X72).
Sulphate, Al2(SO,)3.1sH,0 (55. X72).
Amm^onium Carbonate, (NH4)..C0, (240, 5N), one part crystallized salt in four
parts water, with one part ammonium hydroxide.
Ammonium Chloride, NH.Cl (267, 5X), one part salt in eig-ht parts water.
Hydroxide, NH4OH (85NH5 . •'>N, sp. gr. 0.9G4), »p. gr. 0.96, 10 p. c.
NH, .
Ammonium Molybdate, (NHJ^MoO, (36MoOn , X/2, §75, 6(f), 150 g. salt in one
litre of NH.OH , pour this into one litre of HNO, , sp. gr. 1.2.
Ammonium Oxalate, (NH4).C204.2H30 (40, N/2), one part crystallized salt in
24 parts water.
Ammonium Sulphate, (NHJ.SO^ (33, N/2).
Sulphide, (NH«)oS, colorless, three parts NH«OH , saturate with
H,S and add two parts of NH«OH .
•In the (rreater number of cases, reagents should be "chemically pure." DilTerentuses
require dllTerent degrees of purity. An article of sodium hydroxide contaminated with
chloride may be used in some operations ; not in others. Those who have had training io
analysis can do without speciflo directions, which cannot be made to cover all circumstances:
and the beginner must depend on others for the selection of reagents.
REAGENTS, 403
Ammonium Sulphide, (NH«)2Sx , yellow, allow the colorless to stand for some
time or add sulphur.
Antimonic Chloride, SbCl^ (30, N/2).
Antimonous Chloride, SbCl, (38, N/2).
Arsenous Oxide, As.Os (S, N/4), saturated aqueous solution.
Barium Carbonate, BaCO. , freshly precipitated.
Chloride, BaCl2.2H20 (01, N/2), one part salt to 10 parts water.
Hydroxide, Ba(OH)2.vSH2 (32, N/5), saturated aqueous solution.
Nitrate, Ba(N0,)2 (05, N/2), one part to 15 of water.
Bismuth Chloride, BiCl, (52, N/2, use HCl).
Nitrate, Bi(NO.),.5H20 (40, N/4, use HNO,).
Cadmiimi Chloride, CdCl^ (46, N/2).
Kitratc, CdiNO, \ ,- H (77, N/2).
Sulphate, CdS0,.4H,0 (70, N/2).
Calcium Chloride, CaCl,.tiH,0 (r»5, N/2), dissolve in 5 parts water.
Hydroxide, Ca(OH) , , n saturated solution in water.
Nitrate, Cb(NO,),.)H,0 (59, N/2).
Sulphate, CaSOi.i^H^O , a saturated solution in water.
Carbon Disulphide^ CS.. , color k-i^s.
Chromic Chloride, CrCl^ (25, N/2).
Nitrate, CrrNO,), (40, S/2).
Sulphate, CrjSOJs.lsHjO (60, N/2).
Cobaltous Nitrate, Co(NOJ,j'>H.O (7?., N/2), in 8 parts of water.
Sulphate, CoSO,.TH,0 (70, N/2).
Copper Chloride, CuCi,.:^H-.0 (Ki. N/2).
Nitrate, Cu<NO0>.'JH,O |74, N/2).
Sulphate, CUSO4.5H2O (62, N/2), in 10 parts water.
Cuprous Chloride, CuCl (50, N/2, use HCl).
Perric Chloride, FeCl, (27, N/2), 20 parts water to one part metal.
Nitrate, Fe(NO,),.9H20 (67, N/2).
Ferrous Sulphate, FeS04.7H20 (80, N/^, use a few drops of H2S0« ,
Gold Chloride, HAuCl^.nHjO , solution in 10 parts water.
Hydrogen Peroxide, 3 p. c. solution.
Indigo Solution, 6 parts fuming* H2SO4 to one part indigo, pulverize, stir and
cool, allow to stand 48 hours and pour into 20 parts water.
I^ead Acetate, Pb(C2H,02)a.3H30 (95, N/2), dissolve in 10 parts of water.
Chloride, PbCL , saturated solution, N/7.
Nitrate, Pb(N6,)2 (83, N/2).
Magnesia Mixture: MgS04 , 100 g.; NH4CI , 200 g.; NH«OH , 400 cc; H2O , 800
cc. One cc. = 0.01 g. P.
Magnesium Chloride, MgCh.eH,© (51, N/2).
Nitrate, Mg(NO,)2.6E20 (64, N/2).
Sulphate, MgS04.7HoO (62, N/2), in 10 parts of water.
Manganous Chloride, MnCl2.4H26 (50, N/2).
Nitrate, Mn(NO,)2.6H20 (72, N/2).
Sulphate, MnS04.7H20 (69, N/2).
llTercuric Chloride, HgCl. (68, N/2), in 16 parts of water.
Nitrate, Hg(N6,)2 (SI, N/2).
Sulphate, HgS04 (74, N/2).
Mercurous Nitrate, HgNO. (131, N/2), one part salt, 20 parts water and one
part HNO, .
Nickel Chloride, NiCls.eHsO (59, N/2).
Nitrate, Ni(NO,)2.6H,0 (73, N/2).
Sulphate, NiS04.6H20 (66, N/2).
Palladous Sodium Chloride, Na,PdCl4 , in 12 parts water.
Potassium Arsenate, K2AsO« (26, 1/' K3A8O4 -i- 5).
Arsenite, KAsO, (24, % KAsO, -f- 3).
Bromate, KBrO, (14, % KBrOj -^ 6).
Bromide, KBr (60, N/2).
Carbonate, KXO, (207, 3N).
Chlorate, KClOn , the dry salt.
Chloride, KCl (37, N/2).
404 REAGEXTS. §323.
Potassium Chromate, X^CtOa (49, >72), in 10 parts water.
Cyanide^ KCN (^3, N/2), in four parts water.
Bichromate, K^CrsOY (38, Va. K^Ct^O^ -t- *)• in 10 parts water.
Ferrocyanide, K«re(CN),.3H,0 (33, N72), 12 parts water.
Ferricyanide, K,Fe(CN)« (55, N/2), in 10 parts water.
Hydroxide, KOH (312 [90 p. e. KOH], 5N).
lodate, KIO, (18, y^ KIO, -ir 6).
Iodide, KI (83, N/2), diKsolve in 20 parts water.
Mercuric Iodide, K^Hgl* , Nessler^s solution, §207, 6k.
Nitrate, KNO, (50, N/2), the crystallized salt.
Nitrite, KNO, , the dry salt.
Pyroantimonate, K^HsSbjOr.GHsO , see §70, 4c.
Permanganate, KMnO^ (16, 1/2 KMnO« ~ 5).
Thiocyanate, KCNS (49, N/2), in 10 parts water.
Hydrogen Sulphate, KHSO4 , fused salt.
Sulphate, K,SO« (44, N/2), in 12 parts of water.
Platinic Chloride, HsPtCl^.eHjO , in 10 parts of water.
Silver Nitrate, AgNO, (43, N/4), in 20 parts of water.
Sulphate, Ag..S04 , saturated solution, N/13.
Sodium Acetate, NaC.jH,03.3H20 , in 10 parts of water.
Carbonate, NajCO, (159, 3N), one part anhydrous salt or 2.7 parts of
the crystals, NajCGs.lOH.O , in 5 parts of water.
Chloride, NaCl (29, N/2).
Tetraborate, Na3B«O7.10H3O, ftorax, the crystallized salt.
Hydroxide, NaOH (220 [90 p. c. NaOH], 5N), dissolve in 7 parts of
water.
Hypochlorite NaClO, §270, 4.
Nitrate, NaNO, (43, N/2).
Phosphate, NaaHP04.12H20 (60, N/2), dissolve in 10 parts of water.
Phosphomolybdate, §75, (icf.
Sulphate, (35, N/2).
Sulphide, Na^S , one part NaOH saturated with H^S to one part of
NaOH , unchang"ed.
Acid Sulphite, the dry salt.
Sulphite, Na,S0,.7H:,0 (63, N/2), in 5 parts of water.
Acid Tartrate, NaHC4H40« , in 10 i)arts of water
Thiosulphate. Na.S..03.5H.O , in 40 parts of water.
Stannic Chloride, SnCl, (33, N/2).
Stannous Chloride, SnC1..2H.O (50, N/2), in 5 parts water strongly acid with
HCl.
Strontium Chloride, SrCL.OHjO (67, N/2).
Nitrate, Sr(N6,)3 (53, N/2).
Sulphate, SrSO« , a saturated aqueous solution.
Zinc Chloride, ZnCL (34, N/2).
Nitrate, Zn(N6,)-6H20 (74, N/2).
Sulphate, ZnS04.7H30 (72, N/2).
INDEX.
PAGE
Acetates, detection of 251
ignition of 259
with ferric salts 154
Acetic acid 240-251
estimation of 251
glacial 250
formation of 250
occurrence of 249
preparation of 250
properties of 249
reactions of 250
solubilities of 250
Acids, detection of, notes on 389
displacement of weak by strong. 180
effect of concentrated sulphuric
upon 378
list of 13
precipitated by barium and cal-
cium chlorides 386
preparation of 395
separation from bases 368
table of, precipitated by silver
nitrate 387
table of separation of 388
Alkali carbonates, with third and
fourth group salts 142
group 221
hydroxides, action on double
cyanides 265
hydroxides, detection of in pres-
ence of carbonates 262
hydroxides, reactions with 221
Alkalis, on third and fourth group
metals 140
Alkali metals 5
Alkaline earth metals 5
earth metals in presence of phos-
phates 220
earths, relative solubilities of . . . 204
Alkali sulphides, as reagents. 308, 309
action of, on stannic salts 86
action of, on stannous salts.... 85
Alloys, analysis of 367
with copper 104
Alumlntim 142-146
acetate 144
compounds, ignition of 146
PAOB
Aluminum, detection of 146, 162
distinction from chromium 148
estimation of 146
hydroxide, formation and prop-
erties 144
hydroxide, solubility in ammo-
nium chloride 161
occurrence of 143
oxidation of 146
oxide and hydroxides 143
phosphate, separation of 145
preparation of 143
properties of 142
reduction of 146
phosphate, separation of 146
separation of, from iron by
Na^SaO, and NajSO; 145
separation of, from Cr and 4th
group by basic acetates 143
separation of, from glucinum... 196
salts, reactions of 143
salts, with hydrosulphuric acid.. 145
salts, with phenylhydrazin 144
solubilities 143
Alums 145
Ammonia, occurrence 229
formation of, from nitric acid.. 278
. preparation of 229
properties of 229
Ammonium 229-233
arsenomolybdate 62, 98
benzoate, in separation of Cu
from Cd 107
carbonate, as a reagent 230
carbonate, in separation of As ,
Sb and Sn 119
chloride, as a reagent 231
chloride, in the third group.... 161
chloride, with PtCl* 95
compounds, solubilities of 229
cyanate in formation of urea... 271
detection of 232
directions for detection 236
estimation of 232
hydroxide, as a reagent 230
hydroxide, as a distinguishing
reagent for the first group.... 54
406
IXDEX.
PAGE
Ammonittm hydroxide, detection
by mercuric chloride 232
hydroxide, preparation and prop-
erties of 229
molybdate, preparation of 98
molybdate, test for phosphates. 302
molybdate, with arsenic acid 67
oxidation of 23:^
phosphomolybdate 98
picrate, formation of 230
polysulphide, formation of 231
salts, detection by Nessler's re-
agent 231
salts, ignition of 232
solution to be tested for 23G
sulphate, in separation of stron-
tium and calcium 220
sulphide, as a reagent 231
sulphide, formation of 230
sulphide, preparation of 307
sulphide, on iron and zinc
groups 18'*,
sulphide, yellow, formation of.. 115
sulphide, yellow, in separation of
cobalt and nickel 185
sulphide, yellow, in cupric salts. 115
test for nitric acid 281
thioacetate as a substitute for
hydrosulphuric acid 307
Analysis of alkali group 230
proximate 14
operations of 13, 20
ultimate 14
Anions, table of separations of... 38S
Antimonic acid 70
distinction from antimonous 122
reduction to antimonous by stan-
nous chloride 78
salts, action of hydriodic acid
on 78
sulphide, precipitation of 77
antimonites 74
Antimonous argcntide 70
compounds with silver nitrate. . , 78
iodide, formation of 78
oxide, formation of 70
salts with permanganates 78
salts with chromates 78
sulphide 74
sulphide, precipitation of 77
Antimony 72-S2
acids of 72
compounds, reduction with char-
coal 80
detection of, in allovs 367
PA6I
Antimony, detection of 80
detection of traces of ; 121
distinction from arsenic 7s
estimation of 81
in the test for aluminum 16;^
metal with hydrosulphuric acid. 6<»
mirror 65
notes on analysis of 121
occurrence of 72
oxidation of 81
oxides of 72
pentachloride 74
preparation of 72
properties of 72
reduction of 81
reduction to metallic 70
salts 74
separation from arsenic by per-
oxide of hydrogen 120
separation from arsenic 64
separation from tin by sodium
thiosulphate 78
separation from tin 81
solubility of 73
spots 66
sulphide, separation from arsen-
ous sulphide 121
sulphide, separation from stan-
nous sulphide 121
with iodine 6i>
Argt>l, purification of 252
Arsenates, distinction from arsen-
ites 70, 71
separation from phosphates 290
Arsenic 56-72
acid, precipitation by hydrosul-
phuric acid 114
acid, reduction by hydrosul-
phuric acid and hydriodic acid. 61
acid, reduction with sulphurous
acid 60
acid, with ammonium mol^-bdate 67
acid, with molybdates 62
acid, with nitric acid 66
acid with silver nitrate 67
antidote for 62
compounds, ignition of 69
compounds, with concentrated
hydrochloric acid 61
compounds, with magnesium salts 61
compounds, with stannous chlor-
ide 80
detection of 70
detection of. in poisoning 6-^
distinction from antimonv 78
INDEX,
407
PAGE
ArseniCy estimation of 70
in glass tubing 70
metal with hydrosulphuric acid. 66
method of Fresenius and Babo.. 68
mirror 64, 65
notes on analysis of 121
oxidation of 71
oxides of 57
occurrence of 57
pentasulphide, formation and*
properties of 60
preparation of 57
properties of 56
reaction with alkali sulphides... 59
reaction with hydrosulphuric
acid 59
reduction of 71
reduction by stannous chloride. 61
separation from antimony 64
separation from antimony by
peroxide of hydrogen 120
separation from Sb and Sn by
use of thiosulphates 60
spots, formation of 64
spots, properties of 66
sulphide, separation of, from
Sb,S3 121
sulphides with ammonium car-
bonate 118
trichloride, formation in analysis 61
with peroxide of hydrogen 71
with hydrosulphuric acid gas... 67
with iodine 66
with nitric acid 66
ArseniteSy distinction from arsen-
ates 121
Arsenous hydride 64
oxide, crystals, identification of. 67
sulphide, solubilities of 58
sulphide, with HCl gas 67
Arsine 64
from alkaline mixtures 64
reactions with KOH 121
separation from stibine 65
with hydrosulphuric acid 60, 65
Atomic weights, table of 1
Azoimide (hydronitric acid) 274
Barium 205-208
carbonate, action on ferric salts. 154
carbonate, as a reagent 206
carbonate, as a reagent for third
and fourth groups 142
PAQB
Barium carbonate, as a reagent to
precipitate chromium 148
carbonate, and ferric salts 153
carbonate, to separate phos-
phates from third, fourth and
fifth groups 189
chloride, separation of, from
SrCL and CaCl, by HCl 206
detection of 208
estimation of 203
hydroxide, formation of 205
iodide, properties 358
occurrence of 205
oxide, preparation of 205
peroxide, ignition of 287
peroxide, preparation 205
preparation of 205
properties of 205
salts, separation of sulphites
from sulphates 207
salts, spectrum of 207
separation of, from Sr , Ca and
Mg by sulphates 207
solubilities of 206
strontium and calcium, separa-
tion of by alcohol 220
sulphate, separation 209
Bases, alkali 11
alkaline earth 11
copper, group of 12
definition of 3
fifth group of 11
first group of 13
fourth group of 12
iron group of 12
need for separation from acids..
368, 369
second group of 12
silver group of 13
sixth group of 11
third group of 12
tin group of 13
zinc group of 12
Beryllium 195
Bismuth 100-104
blowpipe, reactions of 103
chloride, sublimation of 103
detection of 103
detection in alloys 367
detection by cinchonine 102
detection as iodide 103
detection by alkaline stannite... 103
detection of traces of 102
dichromate 103
estimation of 103
408
INDEX.
PAGE
Bismttth hydroxide, solubility in
glycerol 101
iodide, stability toward water... 103
nitrate, precipitation with HCl. 101
nitrate, reactions 101
notes on analysis of 128
occurrence of 100
oxidation of 104
oxides and hydroxides of 100
oxychloride, formation of 101
pentoxide, reaction with halogen
acids 101
preparation of 100
properties of 100
reactions of, comparison with
Cu and Cd 112
reduction by grape sugar 104
salts, reaction with the alkalis.. 101
separation from Cu by glycerol. 101
solubility of 100
sulphide, formation of 102
sulphide, separation of, from
CuS 102
. sulphide, separation of, from tin
group 302
Blowpipe, examination of solids. . 374
Blue vitriol 105
Bonds, plus and minus 238
Borates, green flame by ignition of 246
in analysis 54
reactions of 246
Borax, bead, formation of 247
bead, test for Mn 184
bead, use of 365
Boric acid 245-247
estimation of 247
formation of 245
occurrence of 245
• preparation of 245
properties of 245
solubility of 246
Boron 245
Bromates, detection of 349
estimation of 350
ignition of 349
preparation of 348
solubilities of 349
Bromic acid 348-350
properties of 348
reactions of 349
Bromides, detection of 347
detection in presence of iodides.
391, 392
estimation of 348
formation of 345
Bromides, ignition of 347
solubilities of 345
with first group metals 346
Bromine 342-344
detection of 344
estimation of 344
formation of 343
occurrence of 343
preparation of 343
properties of 342
reactions with 343
solubilities of 343
Brown ring, test for nitric acid. . . 280
Brucine, reactions with nitric acid 281
Gacodyl oxide, test for acetates... 250
Cadmium 110-112
detection of 112
estimation of 112
hydroxide 110
notes on analysis of 129
occurrence of 110
oxide 110
separation of 110
properties of 110
reactions of, comparison with
Bi and Cu 112
separation from Cu by KCNS... Ill
separation from Cu by glycerol. . 105
solubilities of 110
separation from Cu by Na^SsO,
and NOsCO, Ill
salts, absorption by gaseous sub-
stances, separation from Cu... Ill
salts, fused with K^S 112
salts, with alkaline tartrates,
separation from Cu Ill
salts, with alkalis 110
salts, with ammonia Ill
salts, with barium carbonate 111
salts, with ])yrophosphates, sepa-
ration from Cu Ill
salts, reactions with Na:2S203 ,
separation from Cu Ill
salts, reduction of by metals.... 112
salts, reduction of by ignition. . . 112
Caesium 233-234
Calcium 210-214
carbonate in spring water 211
carbonate, solubility of 218
detection of 213
detection of by spectrum 213
estimation of 213
group 203
group, directions for analysis of. 218
INDEX.
409
PAOE
Calciiun hydroxide, formation and
properties 211
hydroxide, formation by NajS.. 213
hydroxide, to detect COj 212
oxide, formation and properties. 210
occurrence of 210
peroxide 211
preparation of 210
properties of 210
salts with Na^S 213
salts, separation of oxalic from
phosphorJf (it'id by 212
^eparatioti from Ba and Sr by
(NH,),SO, .., 211
separation from Ba and Sr by
muyl alcohol 211
solubilitiea of 211
Biilphatf', separation from stron-
tium hiilphate 209
siilphnte, solubility in ammo-
iihnu sulphnti^ 220
sulphate, to detect strontium... 213
Carbon 247-249
detection of 249
preparation of 248
properties of 247
reactions of 248
reduction by ignition with 248
rel:iti*pti*s t>f 10
solubilities of 248
dioxide 238-263
dioxide, absorption by CaCOH),. 261
dioxide, detection in sodium car-
bonate 262
dioxide, detection by calcium hy-
droxide 212
dioxide, distinction from H,S ,
SO, W O, etc 261
dioxide, formation of 259
dioxide, occurrence of 259
dioxide, properties of 259
monoxide 254, 255
Carbonates, acid, decomposition of 230
decomposition of, by acids 262
detection of 262
detection of traces 390
estimation of 263
ignition of 262
occurrence of 259
preparation of 259
reactions with 260
Carbonic oxide, formic anhydride. 254
Caasins' purple 93
Cerium 193
Chili saltpeter, occurrence of 277
PAOB
Chloric acid 339-341
formation of 339
preparation of . . i 339
properties of 339
separation of, from nitric acid.. 391
Chlorates, detection of 341
distinction from nitrates 392
estimation of 341
formation from chlorine 329
ifiCiiIliou of 340
oxlrlatiod by ignition of 341
preparation of 339
rt'Qctioui^ with 339
sohiblllties of 339
Chlorides, detection of 149
detection of, in presence of bro-
mides 335, 336, 391
detection of, in presence of cy-
auide.s or thi*»eyaiii\tes, . . .335, 392
formiition of 331
ifriiition of 334
Chloride of lime, formation of.... 337
estimation of, by H^Oa 287
Chlorine 327-330
action on metals 328
detection of 330
estimation of 330
formation of 327
occurrence of 327
peroxide, formation and proper-
ties 338
as an oxidizer 328
properties of 327
solubilities of 328
CMorochromie test for chlorides. 335
an hyd ride 149
Chlorous acid, formation and de-
tection 338
properties of 337
Chromates 150
with As'" 149
with antimonous salts 78
use in separation of barium.... 207
in test for HCl 149
with ferrous salts 158
reduction of by HjS 149
reduction of, by hydrochloric
acid 149
Chromic acid, identification of... 392
formation of 149
detection of 150
Chromium 147-151
distinction from aluminum 148
estimation of 150
410
IXDEX.
Chromium hydroxide, solubility in
ammonium hydroxide
and manganese in third group
separation
metal, solubility of
occurrence of
oxides and hydroxides
oxidation of
properties of
preparation of
reduction of
oxide, solubilities of
salts, solubilities of
salts, reaction of
separation from Fe by NajSgOs
and Na,SOs
separation -from Al and Fe by
H,0.
separation from fourth group . .
Chromous salts
Cinchonine as a test for bismuth.
Citric acid 251
detection of oxalic acid in
distinction from tartaric
properties and reactions
Colloidal sulphides of the fourth
group
Color, flame tests
Columbium, distinction from Ti..
separation from tantalum
properties and reactions
Cobalt 103-
bead test
detection of
detection of in presence of Ni
by H,0,
cobalticyanide separation from
nickel
estimation of
hydroxide
metal, solubilities of
nitrate, effect of ignition with..
occurrence of
oxidation of
oxides and hydroxides
properties of
phosphate, a distinction from Ni
preparation of
reduction of
saits, with alkalis
salts, with barium carbonate...
salts, solubilities of
separation from nickel by ether,
separation from nickel by KNO,
1G2
163
147
147
147
150
147
147
150
147
147
148
145
150
14S
148
102
-252
251
251
251
184
305
201
103
193
-16S
167
168
185
166
16S
161
164
365
164
168
164
16:i
167
164
16S
165
165
164
164
166
PAOt
Cobalt, separation from nickel by
KMnO« 167
separation from nickel by ni-
troso-/?-naphthol 166, 185
Colloidal sulphides of fourth
group 184
Color, flame tests 365
Columbium, properties and reac-
tions of 193
distinction from Ti 201
separation from tantalum 198
Copper 104-110
acetoarsenite 108
analysis of, notes 128
arsenite lOS
compounds with cyanogen 107
detection of 109
detection of traces of, with H^S. 108
detection of, in alloys 367
detection of, with HBr 108
electrical conductors 104
estimation of 109
ferrocyanide, formation of 107
group, metals of 56, 100
hydroxide of 104
occurrence of 104
oxide of 101
preparation of 104
properties of 104
precipitation of, by iron wire... 109
reactions of, comparison with
Bi and Cd 112
reductit)n of, by KCNS 107
reduction by ignition 109
salts, detection by potassium
xanthate 107
salts, reaction with zinc-plati-
num couple 109
salts, reduction of, with HaPO, . 107
salts, separation of, from Cd by
Na,P,07 *. 107
salts, solubilities of 105
traces, loss of 115
traces of, with K^FeCCN), 107
separation of, from Bi by gly-
cerol * \. 101
separation of, from Cd by gly-
cerol 105
se])aration of, from Cd by
NajS.Os and NajCO, ill
separation from Cd by nitroso-
/y-naphthol : 107
separation from Cd by ammo-
nium benzoate 107
INDEX.
4H
PAGE
Copper, separation from Pd 106
Cream of tartar, formatiou of.... 252
Cuprammonium salts 106
Oupric hydroxide in 13^4 OH 105
hydroxide, effect of boiling 106
hydroxide, formation of 106
hydroxide, with glucose 106
hydroxide, with tartrates... 105, 106
salts, reaction with glucose 105
salts, reaction with iodides..... 108
salts, reaction with Na^SsO, .... 108
salts, reduced by SO, 108
sulphide colloidal lOS
sulphide, formation of 107
sulphide, separation from Cd by
HsSO^ 108
sulphide, solubility in (NH4),S,. 108
sulphide, solubility in KCN 108
sulphide, with K^S 115
sulphide, with (KH4)aS, 115
Cuproua iodide 108
oxide, formation of, by glucose. 105
salts, oxidation of, by ASsO,.... 110
salts, with metallic sulphides... 107
salts, separation, from Cd by S. 107
sulphide, formation by NaaS^O,. 108
thiocyanate, formation of 107
Cyanates, detection of, in presence
of cyanides 271
Cyanic acid 271
Cyanide of silver, distinction from
chloride 265
Cyanides, detection as thiocyan-
ate 267
double, dissociated by acids 264
double, not dissociated by acids. 265
estimation of 267
guaiacum test 267
ignition of 266
cyanides preparation of 264
reactions with 2">4
simple, with mineral acids 265
solubility of 264
transposition by acids 267
C3ranog^n properties and reac-
tions 263
Danger and Flandin, detection of
arsenic 69
Pecomposition of organic mate-
rial 362, 363
Didymiunx 194
Dialysis, separation from organic
material by 361
PAOB
Diphenylamine test for nitric
acid .., 281
Dissociation, ele6t)rDlytic 20
Dithionic acid, formation and
properties 314
Dragendorif's reagent 102
Electrolytic dissociation 21
Epsom salts 304
Equations illustrating oxidation
and reduction 396
rule for balancing 239
Erbium. 195
Ethyl acetate, odor of 250
Everett's salt 154
Fatty material, removal of 363
Ferric acetate, formation of 250
acetate, separation of from chro-
mium 154
basic nitrate, separation from
aluminum 158
and ferrous compounds distinc-
tion 16*>
hydroxide, antidote for arsenic. 62
phosphate, formation of 156
salts, detection of traces 15S
salts, with acetates 154
salts, with BaCO, 153
salts, with HI and iodides 158
salts, with H,S 157
salts, with H.iFO, 156
salts, with K,Fe(CN)a 155
salts, with K^FeCCN). 155
salts, with KCNS 155
salts, with stannous chloride. ... 89
salts, separation from ferrous
sulphate 153
Ferric thiocyanate, distinction
from ferric acetate 154
hindrance to reactions of 155
Ferricyanides, in distinction be-
tween Co and Nl 166
reactions of 270
Ferrocyanides, detection of 269
detection and estimation 271
reactions of 268
Ferrous iron in the third group... 161
iron in the third group with
phosphates 189
detection of, in ferric salts 155
salts, traces in ferric salts 155
salts, with chromates 158
salts, with HNO^ 156
salts, with KCN 154
412
INDEX,
PAGE
Ferroua salts, with KaPeCCN)..... 155
salts, with K^FeCCN). 154
sulphate, with f^old salts 93
rirst g^roup metals, table of 52
Fixed alkalis 221
alkali hydroxides on stibine .... 79
alktilis with ^alts of tin 84
Flame, blowpipe* production of... 364
or color tests 373
oxidizing and reducing 363
reactions with copper .saltsj 109
Fluorides, soUibillties of 289
Fluorine 288
FluosUicates, formation of 289
Fluoslliclc acid 247-248
in detection of potassium 225
in separation of Ba , Sr and Ca. 207
FonnateSy formation from cyan-
ides 266
Fourth group, directions for anal-
ysis 184
reagents 141
sulphides colloidal 18-1
table of 183
Fresenius and Babo, detection of
a tf^enic 68
Fioehde^a reagent 99
Fulminiitiiig gold 92
Gallium (eka-aluminium) 195
Oases, absorption of by palladium 131
CJermanium, properties and reac-
tions 136
sulphide 118
OlasSy etching by hydrofluoric
acid 289
Glauber's salts 304
Glucinum (Beryllium) 195
distinction from j^ttrium 202
separation from aluminum 196
separation from cerium 193
Glucose, in formation of cuprous
oxide 105
Gold 91-93
detection in alloys 367, 368
detection of 93
distinction from Pd 132
estimation of 93
fulminating 92
notes on analysis 123
occurrence, properties, etc 91
reduction by ferrous sulphate... 93
reduction with oxalic acid 92
with alkalis 92
salts with stannous chloride.... 89
Gold, separation from Ir 133
Greenocklte 101, 110
Gjrpsum 213
Halogens 9
as oxidizers KSO
compounds, comparative table of 361
hydracids as reducers 330
Heat, upon substances in closed
tubes ^64, 370
upon substances in open tubes.
364, 371
Hydriodic acid 353-356
action on antimonic salts 73
action on arsenic salts 61
on ferric salts 158
as a reducer 354, 355
formation of 353
Hydrobromic acid 344-348
detection of Cu with lOS
formation of 345
occurrence of 345
preparation of 345
properties of 345
reactions of 345
Hydrochloric acid 330-336
action on SbjSj 77
action on bismuth nitrate 102
effect of excess in second group. 113
formation of 331
formation from MgrClz 216
gais on iirf^LMjit^ sulphide 67
occurrence of 331
preparation of 331
properties of 330
reactions with 332
solubilities of 331
Hydrocyanic acid 263-267
formation of 264
occurrence of 264
on PbOa 264
preparation of 264
properties of 263
solubilities of 264
Hydroferricyanic acid 269-271
Hydro ferrocyanic acid 267-269
«e pa ration from hydroferri-
fviHTie acid 2^0
Hydrofluoric ncid 2S9
Hydrofluo silicic acid (fluosilicic
acid) 289
Hydrogen 243-244
absorption by Pd sponge 131
estimation of 244
detection of 244
INDEX.
413
PAGE
Hydrogen, formation of 243
nascent 244
occluded 244
occurrence of 243
preparation of 243
properties of 243
reactions with 243
reducing" action of, with ignition 244
solubilities of 243
peroxide, detection of 287
peroxide, estimation of 287
peroxide, estimation of bismuth
with 104
peroxide, formation Cfi 286
peroxide, occurrence of 286
peroxide, on sulphides of arsenic
and antimony 120
peroxide, preparation of 286
peroxide, properties of 285
peroxide, reactions with 286
peroxide, reagent to separate Co
from Nl 185
j)eroxide, separation from ozone 237
peroxide, separation of Al , Fe
and Cr with 150
peroxide, solubilities of 286
peroxide, with arsenic 71
Hydronitric acid 274-275
'Hydrosulphuric acid 306-311
action on copper salts 107
action on ferric salts 157
aqueous solution 113
dissociation of 114
formation of 307
gas as a reagent 113
gas on arsenic 67
gas on antimony 67
occurrence of 307
on aluminum salts 145
on stannic salts 86
on stannous salts 85
on third and fourth group
salts 141, 161
preparation of 307
properties of 306
uses as a reagent 30S
with arsenic acid 114
with oxidizing agents 114
Hydrosulphurous acid 314
Hydroxylamine, formation and
properties 278
Hypobromotis acid, formation and
properties 348
Hypochlorites, detection of 392
formation of 337
PAOB
Hypochlorites, formation from
chlorine 329
on arsenic 66
Hypochlorous acid 337
Hyi>oiodous acid, existence of. . . . 351
Hyposulphites, detection of 296
ignition of 296
Hypophosphites in formation of
PHs 296
Hypophosphoric acid 298
Hypophosphorous acid 295-297
estimation of 297
formation of 295
preparation of 296
properties of 295
reactions of 296
solubilities of 296
with bismuth salts 102
formation and properties 295
Hjrposulphurous acid 314
Imperial green 108
Indigo test for nitric acid 281
Indium 196
Ink, sympathetic 154
lodates, detection of ....'. 359
estimation of 359
formation of 359
ignition of 357
reactions of 358
Iodic acid 357-359
formation of 357
preparation of 357
properties of 357
reactions of 358
Iodide of nitrogen 351
Iodides, detection of 356
decomposition by HNO, 281
detection as Pdl, 131
estimation of 356
formation of 353
ignition of 356
occurrence of 353
reactions of 354
separation of, from bromides
and chlorides by "KULiiO^ 176
solubilities of 353
Iodine 350-352
detection of 352
estimation of 352
formation of 351
liberation by copper salts 108
occurrence of 351
on antimonous salts 78
on antimony 66
414
INDEX.
PAOB
Iodine, action on arsenic 66
preparation of 352
properties of 350
reactions of 351
separation from Br by Pd 133
solubilities of 351
Ions 21
Ionization and solution 20, 24
Iridium 132-133
Iron ,. 151-159
and zinc groups 140
group 142
group, separation from Co , Ni ,
and JCn by ZnO 158
detection of 162, 103
detection of traces in copper 154
detection of traces 154, 155
estimation of 159
hydroxides 152
solubilities of 152
in relation to metals (>
solubilities of 152
occurrence of 151
oxidation of 150
oxides 152
preparation of 151
properties of 151
reduction 159
salts, ignition of 15S
salts, solubilities of 15 I
salts, with alkalis 15.)
salts, with nitroso-/3-naphthol. . . 154
salts, separation from Al as
basic nitrate 15S
separation from Al and Cr by
nitro80-/?-naphthol 154
separation from Cr and Al 154
separation from Ni by xanthate. 170
Lanthanum 197
Lead 29--M)
acetate, properties of 32
chloride 34
cnni])ounds, ignition of 35
cliromate, formation of 35
chloride, precipitation of 53
detection of 3r>
detection in alloys 3fi7
estimation of 3(>
in the test for Al lO:*.
iodide, formation and properties 35
notes on analysis of 127
occurrence of 20
oxidation and reduction 3r>
oxides of 20
Lead oxides, solubilities of 90
preparation of 29
properties of 29
red 29
relation to nitrogen family 7
salts, reactions 32, 35
salts, solubilities of 31
solubilities of metallic 30
sulphate, formation and proper-
ties of 34
sulphide, formation and proper-
ties of 33
tests for i 54
Leblanc-soda process 259
Lithium 234-236
Lime, slacked 211
stone (CaCOs) 213
Light, action on silver salts 50
Magnesia mixture 145
Magnesium 214-216
as a reducing agent 216
detection of 216
estimation of '. 216
hydroxide, formation 214
occurrence of 214
oxalate, separation of, from K
and Na 215
oxide, formation of 214
preparation of 214
properties of 214
removal for detection of sodium. 236
salts, with ammonium salts 215
salts, with arsenic acid 61
salts, with Na.S 215
salts, solubilities of 214
Malachite 104
ManganateSy identification 392
Zlianganese 172-177
detection of 176, 186
estimation of 177
hydroxides of 172
hydroxides, solubilities of 171
ignition of 176
in third group 161^ 1G3, 184
occurrence of 172
oxidation of 177
oxidation to permanganic acid.. 175
oxides 172
oxides, solubilities of 173
preparation and properties 172
reduction of 177
reduction by sulphites 175
salts, reactions with oxalic acid. 174
salts, solubilities of 173
INDEX,
415
PAGE
salts, with alkalis.... 174
sulphides 175
from zinc with acetic
184
of 173
176
?id 172
i 62
law of 22, 58
)niiun compounds ... 30
loride with stannous
88
ormation and proper-
41
► ith KaS 115
37-15
42
, igfnition of 43
,nd estimation of 44
42
nalysis of 367
of 37
)f 45
• ■.*.... •11
1 and i^roperties of... 37
tions 39, 43
lilities of 38
of 37
nalysis of 126
jification 10
375
paration 376
)ric acid 299
acid 83
gnt 43, 232
salt 230
tion 365
; 211
in analysis 54
hates 98
97
d 97
; 97-90
b in second group 99
)f 99, 122
of 99
its 99
lalysis of 123
of 97
hydroxides 97
1 and properties 97
^ests 99
of 97
rogeu ou nitric acid.. 278
PAOB
N'eodymium 194, 197
Nessler's reagent 43, 231
Nickel 168-172
distinction from cobalt 170
detection of 171
detection of, in presence of Co
by KI 185
estimation of 171
hydroxides 169
ignition of 171
occurrence of 169
oxidation of 171
oxides 169
properties and preparation 168
reduction 172
salts with alkalis 169
separation from Co , cyanide
method 166
separation from Co , by nitroso-
/i-naphthol 168
separation from Co, by KNO,.. 166
separation from Co , by sulphide 170
separation from Co , by xan-
thate 170
solubilities of 169
solubility of WiS in ammonium
sulphide 170
xanthate, separation from Fe.. 170
hydroxide with KI 171
Niobium (Columbium) 193-194
Nitrates, decomposition by igni-
tion 280
distinction from chlorates 392
occurrence of 277
preparation of 277
proof of absence 390
solubilities of 278
Nitric acid 277-282
as an oxidizer 278
brown ring test 281
decomposition of, by HCl 279
detection of 280
detection by diphenylamine 281
detection by reduction to NH,.
278, 281
detection by reduction to nitrite. 2S1
dissociation, by heat 279
estimation of 282
formation of 277
indigo test 281
sodium salicylate test 281
with phenol 281
with pyrogallol 282
with brucine 281
in separation of Sn , Sb and As. 119
416
INDEX.
PAGE
Nitric acid, occurrence of 277
on antimony 66
on arsenic 66
preparation of 277
products of reduction 278
properties of 277
Nitric anhydride, formation of.... 278
oxide 104, 275, 215
Nitrites, decomposition by igni-
tion 276
detection of 276
test for nitric acid 281
Nitroferricyanides 270
Nitrogen 273-274
chloride 62, 120, 327
xjombination with elements 274
detection and estimation 274
family 7
formation, occurrence 274
peroxide 277
properties 273
Nitro80-/i -naphthoic separation of
Co and Ni 166, 185
separation of Cu from Cd 107
with iron salts 154
NitroprUssides 270
Nitrous acid 276-277
as a reducer 276
formation of 276
occurrence of 270
as an oxidiz^T 276
properties of 276
reactions with 276
solubilities of 276
Noble metals, enumeration 7
Nordhausen sulphuric acid 322
Notes on detection of acids 389
on analysis of calcium group. 218, 219
071 analysis of third group 161
Order of laboratory study 24
Organic substances, removal of. .
362, 363
Osmium 133
Osmotic pressure 21
Oxalates, decomposition by igni-
tion of 390
decomposition by oxidation 390
detection of 258
distinction from tartrates. .253, 389
estimation of 258
ignition of 258
in 3d, 4th and 5th groups 189
reactions of 256
solubilities of 256
Oxalic acid
as a reducer 8S6
decomposition of by St80« tSi
formation of 2SS
in separation of gold 92
occurrence of 8S5
preparation and properties of... SS3
solubility of 216
Oxida t Jaa , ^ balancing equations in. S38
Oxidizing flame 363
Oxygen 282-884
as a poison 284
combinations with ignition 284
detection of 284
estimation of 284
formation of 283
occurrence of 282
preparation of 283
reactions with 284
Ozone 284
separation from HjOs 287
Palladium 131-132
distinction from gold and plati-
num 131,132
separation from copper 106
sponge 131
Falladous iodide in analysis 131
Paris green 62. l^
Pentathionic acid, formation and
properties 316
Perchlorates, preparation and
properties 341, 342
Perchromic acid 151
Periodic acid 360
system, table of 2
Permang^anates, identification .... 392
action on antimonous salts "!*
Permanganic acid 1T3
Persulphuric acid 336
Phenol reaction for nitric acid 2S1
Phenylhydrazine, on aluminum
salts 144
Phosgene, formation 254
Phosphates, changes by ignition. 30.1
detection 162, 303, 390
distinction between primary,
secondary and tertiary 301
estimation of 304
in presence of third and fourth
group metals. 142, 188. 189. 191. 192
occurrence of 299
reaction with ammonium molyh-
date 188, 302
separation as ferric phosphate.. 188
INDEX.
417
PAOB
Fhosphatesy solubilities of 300
FhosphideSy formation of 30.i
Phosphine 205
Pliosphoric acid 298-304
preparation of 300
properties of 298
Pliosphorio anhydride, formation
of 299
PhosphorouB acid 297-298
detection of 298
preparation and properties of... 207
Phosphorus 293-295
detection and estimation of 295
in combination with the halo-
gens 294
occurrence and preparation of. . 29:)
properties of 292, 294
use in match-making 293
Phosphotungst&tes 135
Picric acid, in detection of potas-
sium 224
Plaster of Paris (calcium sul-
phate) 213
Platinized asbestos 94
Platinum 93-97
black , 93
apparatus, care of 95
chloride, as a reagent 95
distinction from palladium. .131, 132
detection of 96, 122, 367
estimation of 96
iridium alloys, properties 132
notes on the analysis of 123
occurrence of 94
preparation and properties. . . .93, 94
reduction of 95, 96
sponge 93
Polarity 3
Potassium 222-226
as a reducing agent 226
bichromate, in test for stron-
tium and calcium 219
carbonate, as a reagent 223
chlorate, in preparation of oxy-
gen 283
chloride with platinum chloride. 95
cyanide with copper salts 107
cyanide with ferrous salts 154
detection of 223, 226
estimation of 220
ferricyanide, formation of 269
ferrocyanide, formation of ..265, 267
hydroxide, as a reagent 223
iodate, in separation of alkaline
earths 207
PAOB
Potassium iodide, as a reagent 224
iodide, in separation of Ag^l
from SbCla 120
iodide, in the test for nickel.... 185
iodide, on nickelic hydroxide.... 171
iodide, on permanganates 176
nitrite in separation of cobalt
from nickel 160
occurrence, preparation and
properties of 222
picrate 224
py roantimonate 73, 228
salts, flame test 225
thiocyanate with copper salts... 107
thiocyanate with iron salts 155
xanthate, for detection of copper 107
Powder of algaroth 75
Praseodymium 194, 197
Precipitates, formation and re-
moval of 17, 18
Principles 393
Problems in molecular propor-
tions 19
in synthesis 397
Prussian-blue, formation of.. 155, 266
Purple of Cassius 89, 93
Pyroantimonic acid 73
Pyrogallol, as a test for nitric
acid 282
Pyrophosphoric acid, formation.. 299
Pyrosulphuric acid, formation . . . 322
Reagents, care in the addition of. 17
list of 402
Reducing flame, description of . . . 363
Reduction, balancing equations in 238
with charcoal 364, 365, 371
Reinsch's test for arsenic 67
Rhodium, distinction from ruthe-
nium 133
properties and reactions 130
Rochelle salts, composition of 253
Rosolic acid as a test for carbon
dioxide 262
Rubidium, properties, and reac-
tions 234
Rule for balancing equations 239
Ruthenium, properties and reac-
tions 129
Saltpeter, occurrence 277
Samarium, properties and reac-
tions 107
Scandium, properties and reac-
tions 198
418
INDEX,
PAGE
Scheele's green and Schweinfurt's
green 62, 108
Selenic acid, separation from sul-
phuric acid 139
SeleniunXy properties and reac-
tions 138, 139
Silica (silicon dioxide) 290
detection and estimation of 292
in the borax bead 292
in the third group 163
removal of 390
solubilities of 291
Silicates, decomposition by igni-
tion 291
in analysis 54
Silicic acid 290-292
Silicon 290
distinction from tantalum 198
Silico-fluoride (fluosilicate) 289
Silicon fluoride, formation 288, 289
preparation and properties 290
separation from thorium 200
Silver 45-50
arsenate and arsenite, formation 62
.bromate, properties of 349
chloride, formation and proper-
ties 48
cyannte in distinction from chlo-
rides 271
detection of 50, 367
estimation of 50
in presence of mercury salts.... 55
iodate, properties of 35.S
mirror, formation by tartrates. 25:i
nitrate, action on stibine 79
nitrate with stannous and anti-
monous salts 78, 79, 8S
occurrence and properties of.... 45
salts, action of lipfht upon 50
solubilities of 46
thiocyanate, separation from
silver chloride 272
Soda lime qn stibine 70
process, Le l?lanc's 259
process. Solvay's 2r)0
Sodium 226-221)
amalfifam, action with arsenic... 64
as a reducing,'' apj-ent 229
detection of 73, 22S
estinuition of 229
flame test 22S
hydroxide, formation of 227
nitroferricyanide as reagent. 230, 311
occurrence of 227
phosphate as reagent 227
Pia
Sodium phosphomolybdate as re-
agent 98, 23S
preparation and properties of..
226,22?
pyroantimonate 73,80
pyrophosphate with copper and
cadmium 107
salicylate test for nitric acid... 281
sulphide, preparation of 30$
thiosulphate on cupric salts.... 108
thiosulphate with third group
metals 145
thiosulphate with antimony salts 78
Solids, conversion into liquids.... 366
decomposition upon ignition. 370, 371
effect on ignition with cobalt
nitrate 372
preliminary examination of.... 363
separation of 17
table for preliminary examina-
tion 370
Solubility, degrees of 15, 16
Solubility-product 23
Solutions, conversion into solids.. 367
Solution and ionization 20-24
Solvay soda process 260
Sonnenschein's reagent 9S
Stannic salts, solubilities S4
sulphide, formation and proper-
ties of 86
Stannite, alkali, as a test for bis-
muth 103
Stannous chloride on mercury salts 43
chloride as a reducing agent SS
chloride with gold salts 93
chloride with molybdic acid 99
salts, distinction from stannic
salts 123
solubilities 84
salts with sulphurous acid S6
salts with silver nitrat-e 87
sulphide, formation and proper-
ties S3
Stibine, decomposition by soda
lime \ 79
formation of 79
reaction with fixed alkali hy-
droxides 79
reaction with silver nitrate 79
separation from arsine 65
Strontium 208-210
detection of 210,213
estimation of 210
hydroxide, formation 20^
occurrence of , 208
IXDEX.
419
PAGE
Strontium^ preparation and prop-
erties of 208
sulphate, distinction from CaSOf 209
sulphate; separation from BaSOf 209
Sulphates, detection and estima-
tion of 326
ignition of 325
preparation of 322
reduction by ignition with carbon 249
solubilities of 323
Sulphites, detection of 321
distinction from sulphates 321
estimation of 321
ignition of 321
interference in test for oxalates. 390
preparation of 318
separation from sulphates by Ba
salts 207
solubilities of 319
Sulphides, detection and estima-
tion of 311
formation of 307
ignition of 310
reactions of 309, 310
solubilities of 28, 308
Sulphur 304-306
combinations on ignition of 306
detection and estimation of 306
formation of 304
in the tin group 118
occurrence of 304
oxidation by reagents 305, 306
oxides 304
precipitation of 53, 114, 115
preparation and properties of...
304, 305
reactions in forming sulphides.. 305
relations of 9
separating copper from cad-
mium 107
solubilities of 305
Sulphuric acid 321-326
detection in presence of sul-
phates 326
formation and occurrence of.... 322
properties of 321
reactions with 323, 324, 325
separation from Se 139
separation from Fe 137
anhydride, preparation of 322
Sulphurous acid 318-321
on arsenic acid 60
and sulphites as reducers 320
formation of 318
reduction of cuprie salts 108
PAGE
Sulphurous acid, occurrence of... 318
preparation and properties of... 318
solubilities of 319
on stannous salts 86
Synthesis, problems in 397
Table for acids as precipitated by
barium and calcium chlorides. 386
for acids precipitated by silver
nitrate 387
for acids, preliminary 378
for analysis in presence of phos-
phates by the use of alkali ace-
tates and ferric chloride 191
for analysis in presence of phos-
phates by use of ferric chloride
and barium carbonate 192
for analysis of the Silver Group
(first) 52
for analysis of the Copper Group
(second) 124
for analysis of the Tin Group
(second) 116
for analysis of the Iron Group
(third) 160
for analysis of the Zinc Group
(fourth) 183
for analysis of the Calcium
Group (fifth) 217
of grouping of the metals 375
of separations of the metals.... 376
of separation of the ammonium
sulphide precipitates of the
Iron and Zinc Groups 187
of solubilities 398
Tannic acid with iron salts 154
Tantalum, distinction from silica. 198
distinction from titanium 198
properties and reactions of 198
separation from columbium 198
Tartar emetic, composition of 252
Tartaric acid 252-254
in detection of potassium 223
distinction from citric acid 251
formation and properties 252
Tartrate calcium, deportment with
water 253
detection of • 253
distinction from citrates 253
distinction from oxalates. .. .253, 389
estimation of 254
Tartrates, ignition 253
reactions 253
solubilities 252
Tellurium 137-138
420
INDEX.
PAOE
Tellurium, distinction from sele-
nium 138, 140
properties and reactions of 137
separation from sulphuric acid.. 137
Tenorite 104
Terbium 198-199
Tetrathionic acid, formation and
properties 315
Thallious iodide 199
Thallium, properties and reac-
tions 199
Thioacetate in formation of sul-
phides 307
Thiocyanates, reactions with 272
Thiocyanic acid as a reducer 273
properties of 272
Thionic acids, table of compari-
sons 317
ThioBulphateSy detection of 313
distinction from sulphates and
sulphites 314
estimation of 314
igTiition of 313
formation and properties of 312
Thiosulphuric acid 312-314
Third group reagents 141
Thorium 199-200
Tin 82-S9
creaking of 82
detection of 88, 122, 307
estimation of 88
Group, metals of oii
Group, separation from Cojiper
Group 115
Group, sulphides with (NHJ.Sx ll.">
occurrence of 82
oxidation of 88
oxides and hydroxides 82
preparation and properties of... 82
notes on the analysis of 123
relation to Nitrogen Family 7
reduction by iguition 87
salts with the alkalis 84
salts with hydrosnly)huric acid.. 8r)
separation from antimony 81
separation from arsenic 118
solubilities of 83
sulphides, colloidal 115
separation from antimony sul-
phides 121
Tin with antimony and with arsenic 87
Titanium . .'. 200-201
distinction from columbium 201
distinction from tantalum .. W
•properties and reactions of 200
separation from thorium 200
Trithionic acid, formation and
properties 315
Tungsten, properties and reac-
tions 134
Tumbuirs blue . ; 155
Unit of quantity 22
Uranium, properties and reactions 201
Urea, from ammonium cyanate... 271
Valence, negative 3
Vanadium 135-136
Volatile alkali (ammoninm hy-
droxide) 221
Water, action on bismuth salts... 101
action on antimonous salts 75
Welsbach burners 203
Wolf ramium (tungsten) 134
Widfenite 07
Ytterbium properties and reac-
tions 202
Yttrium 202
Zincates, formation of 179
Zinc nSlJ^l
detection and estimation of ^'^
Family 5
granulated 63, 17S
Group, table for anah'sis 1^^
Group, comparative reactions... 1^-
hydroxide and oxide 1'''
ignition of l^^'*
occurrence of 1'**
oxidat ion of 1^^
platinized 178, 24-^
preparation and iiroperties 1^^
reduction of 1^^
salts, solubilities and reactions
of 1T9
sulphide, formation in presence
of acetic acid 1'^
Zirconium 202-203
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SCIENCE SERIES.
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Ko, 4B, THEORY OF SOLID AND BRACED ARCHES. Bv
William Cain, CE. '
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Thomas Crafg, l^h.D. '
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^^ A, de Varona.
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D. VAX XOSTRAXD COMPANY'S
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Nate. STRENGTH OF WROUGHT-IRON BRIDGE MEM-
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No. 72. TOPOGRAPHICAL SURVEYING. By George J. Specht,
Prof. A. S. Hardy. John H. McMastcr. and H. F. Walling.
No. 73. SYMBOLIC ALGEBRA ; OR, THE ALGEBRA OP
ALdFJlRAIC XUMHFKS. I5y Trof. W.Cain.
No. 74. TESTING MACHINES: THEIR HISTORY, CON-
STRUCTION. AND ISF. I5y Arthur V. Abbon.
No. 75. RECENT PROGRESS IN DYNAMO-BLBCTRIC MA-
Tin NFS. Ucing a Supplcnjcnt to Dynamo-Electric Machinery. By
Pr«»f. .Sylvanus P. Thompson.
No. 76. MODERN REPRODUCTIVE GRAPHIC PROCBSSBa
By Ficut. James S. IVttit, T.S.A.
No. 77. STADIA SURVEYING. The Theory ot Stadia Measurements.
By .Arthur Winsjow.
No. 78. THE STEAM-ENGINE INDICATOR, AND ITS UbE
By W. B. I.c V.m.
No. 79. THE FIGURE OF THE EARTH. By Frank C. RfibertB^CE.
No. 80. HEALTHS FOUNDATIONS FOR HOUSES. Hy CSenB
biu\»n
SCIENCE SERIES.
N0.81. WATER METERS: COMPARATIVE TESTS OF
ACCURACY, DELlVKkY, ETC. Distinctive features of the Worth-
ington, Kennedy, Siemens, and Hesse meters. IJy Ross E. Browne.
No. 8a. THE PRESERVATION OF TIMBER BY THE USE
OF ANTISEPTICS. JJy Samuel Bagster Boulton. C.E.
No. 83. MECHANICAL INTEGRATORS. By Prof. Henry S. H.
Shaw, C.E.
No. 84. FLOW OF WATER IN OPEN CHANNELS, PIPES,
CONDUITS, SEWERS, ETC. With Tables. By P. J. Flynn, C.E.
No. 85 THE LUMINIFEROUS ^THER. By Prof, de Volson Wood
Na86. HAND-BOOK OF MINERALOGY; DETERMINATION
AND DESCRIPTION OF MINERALS FOUND IN THE UNITED
STATES. By Prof. J. C. Foye.
No. 87. TREATISE ON THE THEORY OF THE CON-
STRUCTION OF IIELICOIDAL OBLIQUE ARCHES. By John
L Culley, C.E.
No. 88. BEAMS AND GIRDERS. Practical Formulas for their Re-
sistance. By P. H. Phil brick.
No. 89. MODERN GUN-COTTON: ITS MANUFACTURE,
PROPERTIES, AND ANALYSIS. By Lieut. John P. Wisser, U.S.A.
No. 90. ROTARY MOTION, AS APPLIED TO THE GYRO-
SCOPE. By Gen. J. G. Barnard.
No. 9". LEVELING: BAROMETRIC, TRIGONOMETRIC, AND
SPIRIT. By Prof. I. O. Baker.
No. 92. PETROLEUM : ITS PRODUCTION AND USE. By
lioverton Redwood, F.I.C., F.C.S.
No. 93. RECENT PRACTICE IN THE SANITARY DRAIN-
AGE OF BUILDINGS. With Memoranda on the Cost of Plumbing
Work. Second edition, revised. By William Paul Gerhard, C. E.
No. 94. THE TREATMENT OF SEWAGE. By Dr. C. Meymotl
Tidy.
No. 95. PLATE GIRDER CONSTRUCTION. By Isami Hiroi, C.E.
Second edition, revised and enlarged. Plates and Illustrations.
Na 96. ALTERNATE CURRENT MACHINERY. By Gisbeit
Kapp, Assoc. M. Inst,, C.E.
No. 97. THE DISPOSAL OF HOUSEHOLD WASTE. By W.
Paul Gerhard, Sanitary Engineer.
No. 98. PRACTICAL DYNAMO-BUILDING FOR AMATEURS.
HOW TO WIND FOR ANY OUTPUT. By Frederick Walker.
Fully illustrated.
Tr
RIALS. By Prof. Osborne Reynolds. Edited, with notes, etc., by
F. £. Idell. M. £.
SCIENCE SERIE&
No. xoo. HOW TO BECOME AN ENGINEER ; OK, THE
THEORETICAL AND PRACTICAL TRAINING NECESSARY IN
FITTING FOR THE DUTIES OF THE CIVIL ENGINEER. The
Opinions of Eminent Authorities, and the Course of Study in the
Technical Schools. By Geo. W. Plympton, Am. Soc. C.E.
No. xox. THE SEXTANT AND OTHER REFLECTING
MATHEMATICAL INSTRUMENTS. With PracUcal Suggestions
and Wrinkles on iheir Errors, Adjurtnients, and Use. With thirty-
three illustrations. By F. R. Braiiiard, U.S.N.
No. Z02. THE GALVANIC CIRCUIT INVESTIGATED
MATHEMATICALLY. By Dr. G. S. Ohm, Berlin, 1827. Translated
by William Francis. W..h Preface and Notes by the Editor, Thomas
D. Lockwood, M.I.E.E.
No. X03. THE MICROSCOPICAL EXAMINATION OF POTA-
BLE WATER. With Diagrams. By Geo. W. Rafter.
No. 104. VAN NOSTRAND'S TABLE-BOOK FOR CIVIL AND
MECHANICAL ENGINEERS. Comoiled by Geo. W. Plympton. C.£
No. X05. DETERMINANTS, AN INTRODUCTION TO THE
STUDY OF. With examples. By Prof. G. A. Miller.
No. X06. TRANSMISSION BY AIR-POWER. Illustrated. By
Prof. A. B. W. Kennedy and W. C. Unwin.
No. X07. A GRAPHICAL METHOD FOR SWING-BRIDGES.
A Kational and EUisy Graphical Analysis of the Stresses in Ordinary
Swing- Briilges. With an Introduction on the General Theory of Graphi-
cal Statics. 4 Plates. By Benjamin F. LaRue, C.E.
No. X08. A FRENCH METHOD FOR OBTAINING SLIDE-
VALVE DIAGRAMS. 8 Folding Plates. By Lloyd Bankson. B.S..
Assist. Naval Constructor, U.S.N.
No. X09. THE MEASUREMENT OF ELECTRIC CURRENTS.
Electrical Mkasurinc. Instruments. By Jas. Swinburne. Meters
FOR Elkctricai. Energy. By C. H. Wordingham. Edited by
T. Commerford Martin. Illustrated.
No. XXO. TRANSITION CURVES. A Field Book for Engineers,
containing Rules and Tables for laying out Transition Curves. By
Walter G. Fox.
No. III. GASLIGHTING AND GAS-FITTING, including Specifica-
lions and Rules for Gas Piping, Notes on the Advantages of Gas for
Cooking and Heating, and useful Hints to Gas Consumers. Second
edition, rewritten and enlarged. By Wm. Paul Gerhard.
No. 112. A PRIMER ON THE CALCULUS. By E. Sherman
Gould, C.E.
No. 113. PHYSICAL PROBLEMS AND THEIR SOLUTION.
By A. Huurgougnun, lormcrly Assistant at Bcllevue Hospital.
No. 114. MANUAL OF THE SLIDE RULE. By F. A. Halscy of
tiie American Machinist.
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