vm<#ma^ -zwmrw
'€&^
•
QUALITATIVE
CHEMICAL ANALYSIS:
A GUIDE IN QUALITATIVE WORK, WITH DATA FOR
ANALYTICAL OPERATIONS AND LABORATORY
METHODS IN INORGANIC CHEMISTRY.
BY
ALBERT B. PRESCOTT,
AND
OTIS C. JOHNSON,
PROFESSORS IN THE UNIVERSITY OF MICHIGAN.
o KG MI GAL
DOliATi
SIXTH REVISED AND ENLARGED EDITION,
ENTIRELY REWRITTEN.
WITH AN APPENDIX BY H. H. WILLARD CONTAINING
A FEW IMPROVED METHODS OF ANALYSIS.
NEW YORK:
D. VAN NOSTRAND COMPANY
23 MURRAY AND 27 WARREN STS.
1904
Copyrighted 1901, by
D. VAN NOSTRAND COMPANY.
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 lonization, 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
] 874. 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 CHEMICAL 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).
Lead 29
Mercury 37
Silver 45
Comparison of Certain Reactions of the Metals of the Silver
Group ... 51
TABLE FOR ANALYSIS OF THE SILVER OR FIRST GROUP 52
Directions for Analysis wit h Notes ' 53
THE TIN AND COPPER GROUP (SECOND GROUP).
THE TIN GROUP, OR SECOND GROUP, DIVISION A.
Arsenic 56
Antimony 72
Tin 82
Comparison of Certain Reactions of Arsenic, Antimony and Tin. 90
Gold 91
Platinum 93
Molybdenum 97
THE COPPER GROUP, OR GROUP II, DIVISION B.
Bismuth 100
Copper 104
Cadmium 110
Comparison of Certain Reactions of Bismuth, Copper and Cad-
mium 112
Vi CONTENTS.
PAGE
THE PRECIPITATION OF 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
TABLE FOR ANALYSIS OF THE COPPER GROUP (SECOND GROUP, DIVISION B). . 124
Directions for Analysis with Notes 126
RARER METALS OF THE TIN AND COPPER GROUP.
Ruthenium 129
Rhodium 180
Palladium 131
Iridium 132
Osmium 133
Tungsten 134
Vanadium 135
German ium 186
Tellurium 1ST
Selenium . 138
THE IRON AND ZINC GROUPS (THIRD AND FOURTH GROUPS) 140
THE IRON GROUP (THIRD GROUP).
A I ii in in ii in 148
Chromium 147
Iron 151
TABLE FOB ANALYSIS OF THE IRON GROUP (THIRD GROUP) 160
DIRECTIONS FOK ANALYSIS WITH NOTES 161
v
THE ZINC GROUP (FOURTH GROUP).
Cobalt 168
Nickel '.168
Manganese : 172
Zinc .178
Comparison of Some Reactions of the Iron and Zinc Group
Bases 182
TABLE FOR THE ANALYSIS OF THE ZINC GROUP (FOURTH GROUP) 183
DIRECTIONS FOR ANALYSIS WITH NOTES 184
ANALYSIS OF IRON AND ZINC GROUPS AFTER PRECIPITATION BY AMMONIUM
SULPHIDE 186
IRON AND ZINC GROUPS IN PRESENCE OF PHOSPHATES 18^
IRON AND ZINC GROUPS IN PRESENCE OF OXALATES 180
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 Phosphoric Acid by Means of Ferric Chloride and Barium
Carbonate 192
THE RARER METALS OF THE IRON AND Zlltfl GROUPS.
Cerium !?. » 193
Coin in bin in (Niobium) 193
Did v in in in 194
Erbium . .195
CONTENTS. )ivil
PAGE
Gallium ••.'•>. 195
<• I uc in ii in (Beryllium) i JJ. 195
Indium '. . JJ . 196
Lanthanum , '. «/;.!. 197
Neodymium '.>. 197
Praseodymium V. 197
Samarium '..i. 197
Scandium .' • . 198
Tantalum 198
Terbium ,'.-, . 198
Thallium :.i. 199
Thorium 199
Titanium 200
Uranium '. j . "201
Ytterbium ). 203
Yttrium 202
Zirconium •« . 202
THE CALCIUM GROUP (FIFTH GROUP). (THE ALKALINE EARTH METALS) ,. 203
Barium 205
Strontium 208
Calcium , , . 210
Magnesium 214
TABLE FOR THE ANALYSIS or THE CALCIUM GKOUP (FIFTH GROUP) 217
DIRECTION FOR ANALYSIS WITH NOTES « . 218
SEPARATION OF BARIUM, STRONTIUM, AND CALCIUM BY THE USE OF ALCOHOL 220
ALKALINE EARTH METALS AS PHOSPHATES 220
ALKALINE EARTH METALS AS OXALATES 220
THE ALKALI GROUP (SIXTH GROUP) 221
Potassium . 222
Sodium ' 226
Ammonium 229
Caesium 233
Rubidium . 234
Lithium 234
DIRECTIONS FOR ANALYSIS WITH NOTES 236
PART III.— THE NON-METALS.
BALANCING OF EQUATIONS '. 238
Hydrogen 243
Boron 245
Boric Acid 245
Carbon 247
Acei ic Acid 249
Citric Acid .251
Tartaric Acid 253
Carbon Monoxide . 254
Oxalic Acid ' 255
Carbon Dioxid e (Carbonates) 259
Viii CONTENTS.
PAOB
Cyanogen 363
Hydrocyanic Acid 268
Hydroferrocyanic Acid 267
Hydroferricyanic Acid 269
Cyanic Acid 271
Thiocyanic Acid 272
Nitrogen 278
Hydronitric Acid 274
Nitrous Oxide 275
Nitric Oxide 275
Nitrous Acid 276
Nitrogen Peroxide 277
Nitric Acid 277
Oxygen 282
Ozone 284
Hydrogen Peroxide 285
Fluorine 288
Hydrofluoric Acid 289
Flnosilicic Acid 289
Silicon 290
Silicic Acid 290
Phosphorus 292
Phosphine 295
Hypophosphorous Acid 295
Phosphorous Acid 297
Hypophosphoric Acid 298
Phosphoric Acid 298
Sulphur 304
Hydrosnlphuric Acid 306
Thiosulphuric Acid 312
Hyposulphnrons Acid 314
Dithionic Acid 314
Trithionic Acid 315
Tetrath ionic Acid 315
Pentathionic Acid 816
TABLE OF THIONIO ACIDS 317
Sulphurous Acid 318
Sulphuric Acid 321
Persulphnric Acid 326
Chlorine 327
Hydrochloric Acid 330
Hypocblorous Acid 337
Chlorous Acid 837
Chlorine Peroxide 338
Chloric Acid 339
Perchloric Acid 841
Bromine 342
Hydrobromic Acid 345
Hypobromons Acid 348
CONTENTS. ix
I
PAGE
Bromic Acid 348
Iodine '. 350
Hydriodie Acid 353
lodic Acid 357
Periodic Acid 360
COMPARATIVE REACTIONS or THE HALOGEN COMPOUNDS 301
PART IV.— SYSTEMATIC EXAMINATIONS.
REMOVAL OF ORGANIC SUBSTANCES 362
PRELIMINARY EXAMINATION OF SOLIDS 363
CONVERSION OF SOLIDS INTO LIQUIDS 366
CONVERSION OF SOLUTIONS INTO SOLIDS 367
TREATMENT OF A METAL. OR AN ALLOT 367
SEPARATION OF ACIDS FROM BASES 368
TABLE FOR PRELIMINARY EXAMINATION OF SOLIDS 370
BEHAVIOR OF SUBSTANCES BEFORE THE BLOW-PIPE 374
TABLE OF THE GROUPING OF THE METALS 375
TABLE FOR THE SEPARATION OF THE METALS 376
ACIDS — FIRST TABLE 378
ACIDS — SECOND TABLE 386
ACIDS — THIRD TABLE 387
ACIDS— FOURTH TABLE 388
NOTES ON THE DETECTION OF ACIDS 389
PRINCIPLES 393
EQUATIONS 396
PROBLEMS IN SYNTHESIS 397
TABLE OF SOLUBILITIES 398
REAGENTS 402
ABBREVIATIONS.
A __
A. Ch.
Am.
Am. S.
Arch. Pharm.
A in. < 'lifin.
B.—
Bl.
B. J.
Comey.
C. N.
Ch. Z.
C. r.
C. C.
Mil. -I.
Present us.
O. O.
Oazzetta.
Gilb.
Gmelin-Kraut.
J.
J. C.
J. pr.
J. Soc. Ind.
J. Anal.
J. A in. SOC.
J. Pharm.
Lad en burg.
M.
Phil. Mag.
Pogg.
Proc. Boy. Soc.
Pharm. J. Trans.
Ph. C.
Tr.
Watt's.
1868*
* Indicates continuance to tho present time.
Liebig's Annalen. 1832*
Annales do Chimie et dc Physique. 1789*
American Chemical Journal. 1879* ...
American Journal of Science. 1818*
Analyst. 1876*
Archives der Pharmacie. 1822*
American Chemist. 1870-77.
Berichte der Deutschen Cbemischen Gesellscbaft.
Bulletin de la Societe Chimique. 1859*
Berzelius Jahresbericht. 1832-51.
Comey's Dictionary of Solubilities. 1896.
Chemical News. 1860*
Chemiker Zeitung. 1877*
Comptes Kendus des Seances de 1' Academic des Sciences,
Cbemiscbes Centralblatt. 1830*
Dingler's Polytecbnische Journal. 1 820"
Dammer's Anorganische Cbcmie. 1892*
Fehling's Handbuch der Chemie. 1871*
Fresenius: Qualitative Chemical Analysis.
Graham-Otto: Lehrbuch der anorganischen Chemie.
Gazzetta chimica italiaua. 1871*
Gilbert's Annalen di-r Physik und Chemie. 1799-1824.
Gmelin-Kraut: Handbuch der anorganischen Chemie. 1877.
Jahresbericht fiber die Fortschritte der Chemie. 1847*
Journal of the Chemical Society. 1849*
Journal fur 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. 1 882-1 S'I5.
Monatshefte fur Chemie. 1880*
Menschutkin. Locke's 7V«//.«/<//;<>«. 1895.
Philosophical Magazine. 1798*
Poggendorff's 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:; 5*
Wells' Trans., 1897.
1885.
ABBREVIATIONS.
W. A. Wiedemanu's Anualen. 1877*
W. A. (Beibl.) Wiedeinann's Annaleu Beiblatter. 1877*
Wormley. Wormley's Microcheinistry of Poisons. 1867.
Wurtz. Dictionnaire de Chimie. 1868.
Z. Zeitschrift fur analytische Chemie. 1862.*
Z. Ch. Zeitschrift fiir Chemie. 1865-1871.
Z. anorg. Zeitschrift fiir anorganische Chemie. 1891*
Z. angew. Zeitschrift fiir angewandte Chemie. 1888*
Z. phys. Ch. Zeitschrift fiir physicalische Chemie. 1887*
PAET I.
THE PRINCIPLES OF ANALYTICAL CHEMISTRY.
11
Name.
Sym-
bol.
V
O=16.
H=l.
Name.
Sym-
bol.
O=16.
H=l.
Aluminum ....
Antimony ....
Al
Sb
27.1
120.2
26.9
119.3
Neodymium. . .
Neon
Nd
Ne
143.6
20.
142.5
19.9
Argon
A
39.9
39.6
Nickel
Ni
58.7
58.3
Arsenic
As
75.0
74.4
Nitrogen
N
14.04
13.93
Barium . .
Ba
137.4
136.4
Osmium
Os
191.
189.6
Bismuth
BI
208.5
206.9
Oxygen
O
16.00
15.88
Boron
B
'11.0
10.9
Palladium
Pd
106.5
105.7
Bromine
Br
79.96
79.36
Phosphorus
P
31.0
30.77
Cadmium
Cd
112.4
111 6
Platinum
Pt
1948
193.3
Caesium
Cs
132.9
131.9
Potassium
K
39.15
38.85
Calcium
Ca
40.1
39.7
Praseodymium.
Pr
140.5
139.4
Carbon. .
c
12.00
11.91
Radium
Ra
225
223.3
Cerium
Ce
140.25
139.2
Rhodium
Rh
103.0
102.2
Chlorine
CI
35.45
35.18
Rubidium
Rb
85.5
84.9
Chromium
Cobalt
Cr
Co
52.1
59.0
51.7
58.55
Ruthenium
Samarium
Ru
Sm
101.7
150.3
100.9
149.2
Columbium
Cb
94
93 3
Scandium
Sc
44 1
43.8
Copper
Cu
63.6
63'.1
Selenium
Se
79.2
78.6
Erbium
Er
166.
164.8
Silicon
Si
28.4
28.2
Fluorine.
F
19
18 9
Silver
Ag
107.93
107.11
Gadolinium
Gd
156
154 8
Sodium
Na
23.05
22.88
Gallium
Ga
70
69.5
Strontium
Sr
87.6
86.94
Germanium .
Ge
72 5
72
Sulphur
S
32.06
31.82
Glucinum
Gl
9 1
9 03
Tantalum
Ta
183.
181.6
Gold
Au
197.2
195.7
Tellurium
Te
127.6
126.6
Helium
He
4.
4.
Terbium
Tb
160.
158.8
Hydrogen
H
1.008
1.000
Thallium
Tl
204.1-
202.6
Indium
In
115.
114.1
Thorium
Th
232.5
230.8
Iodine
1
126.97
126.01
Thulium
Tm
171.
169.7
Iridium
Ir
193 0
191 5
Tin
Sn
119.0
118.1
Iron .
Fe
559
55 5
Titanium
Ti
48.1
47.7
Krypton . .
Kr
81 8
812
Tungsten
W
184.
182.6
Lanthanum .
La
138 9
137.9
Uranium
U
238.5
236.7
Lead
Pb
206 9
205 35
Vanadium
V
51.2
50.8
Lithium
Li
7 O1?
6 98
Xenon
Xe
128.
127.
Magnesium ....
Manganese
Mg
Mn
24.36
55 0
24.18
54 6
Ytterbium
Yttrium
Yb
Yt
173.0
89.0
171.7
88.3
Mercury
Hg
200.0
198 5
Zinc
Zn
65.4
64.9
Molybdenum. . .
Mo
•96.0
95.3
Zirconium
Zr
90.6
89.9
TABLE OF THE PERIODIC SYSTEM OF CHEMICAL ELEMENTS.
§2-
42
ft
1
-3
V
•c
ot
cos
%
II ^
ii
II *
.Li OO
Z"5
Q. rH
°- rH
1
d
DC
as
O"5
II CO
-C O
II®
11 CO
_ 05
rH
II a>-
3 rH
O *"*
os in
rH •*»«
II «
0
O?
c
ri
g'
id
II co
"• II
ua
1-
II
rH)
II
0
c
m
SCO
o
co eg
^
to
-I
to
p
x"d
ococ
rH CO
II II
IO
n
CO
O5
i-t
II
00
rH
00
eo
eg
O eo
II
tj
CO
0
H
II
II
O
2
^
D
— -
o
-,
irt
o ,_;
to
O
00
**tf* CO
t-
-t
o
« »
rH ..
^1
—
CO
I ^5
,|
•
II
^T"
00
.
^^
e«
11
rH
!
05
rH
*^
CC ff
a
lA
U)
1 1
2
II
^
jj
CO
nj
m
>
O
o •*«
U3
r
05
S s
-. i
CO
^H '
,_,
t-
CO
H
. —
II
0
eg
£
id
ococ
u 1!
o *
00
H
II
0
o
O5
II
C
CO
* CO
ii ®
!i
Q.
CO
eg
II
F
N
O
h
0 rH
o
us
rt
rH t""
1-
—
•
rH eg
II
• —
*
o
*"«d
II II
^;
II
o
II
OO
CO
eg
M
an:
II II
0 ^
0
O5
00
c
rH
II
^
II
H
0
re
CO
>•
-1
—
_
o
CO
id
eg
0
CO
CO
—
^.
o
_
•
^^
CO
^_
t—
eg
us
T ^3
rH ]^"
0
1^
CO
eg
>~l
or
05
1 I
00
||
i— i
||
eg
c
II
II
N
1 1
T3
O)
_ o
re
L.
0
re
I
re
0 5
0
co
m
OC
CO
O
eg
- i°.
ococ
ua
~, °
1 a
n
rH
O5
CO
CO
11
Irt
ITS
00
II
OO
O
rH
ii
05
eg'
CO
rH
O5
rH
II
"
||
0
II
CT
||
3
m
II
^
*(
10
^
-I Z
^
BE
0
__; O
05
00
00
•
0
DC
!l »
O5
CO
rH
00
II
rH
||
T 1
II
b
X
qg
** |
H * |
.8 "o
§-W bl
H| a
H | 8
o" '" rt
"" § 3
5 •£ a
JM
'S £ H
•H .£3 eS 05
§ 1 1
• f|!
02 C 4) 3
- 5 -° W
°- E £ »
M -n S« C
f a, C as
r^ *-» ^ fn
2!& °1
^ 'S w "5
• > •" —
>o - ** 2
2 -2 fa 5
His
s 17 c **
a II o o
§3. DISCUSSION OF TIJE PERIODIC SYSTEM.
§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 potassium
X
* Bases are the oxygen compounds of the metals. Acids are compounds of elements for
the most part not metals. In the chemical union of sodium with chlorine, for example,
these two elements differ widely from each other in their various properties. The chlorine
is the opposite of the sodium in that very power 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
power of opposite action exercised by the one element upon the other, in their combination
together, is represented by the opposite polarity of the one in relation to the other during
'electrolysis. Electrolysis is an exercise of the same energy that is otherwise manifested
in chemical union or in a chemical change. Strictly speaking, it may be said that it is only
in electrical results that a positive or a negative polarity appears. But the term positive
polarity, applied to sodium because it goes to the negative pole of a battery, is a term
which, well designates the oppositeness of the chemical action of sodium' in its union with
chhjrine. 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 combination, or in a
chemical change.
Iii chemical combination, the atoms of each element act with a " polarity," the extent
of which may be expressed in terms of hydrogen equivalence or " valence." 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 compound known as hydro-
sulphuric acid, the sulphur is negative, the hydrogen positive, In the relation of one to the
other, as represented by the diagram,
H+-
H+-S
in which the plus and minus signs of mathematics are 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
(HOJ-+ I +-°
v H —
That is, the sulphur acts with six units of valence, all in positive polarity. In respect to
oxidation and reduction, the difference between the action of sulphur in hydrosulphtiric
acid on the one hand, and in sulphuric acid on the other hand, is a difference equivalent to
eight units of valence, the combining extent of eight atoms of hydrogen. This value is in
agreement with the factors of oxidizing agents in volumetric analysis.
In the same sense there is a change of " polarity " equivalent to the extent of eight units
of valence, in reducing periodic acid to hydriodic acid, in reducing arsenic acid to arsine, or
In reducing carbon tetrachloride to methane. That is, in any of the groups from IV. to
VII. there is 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 HH and its positive polarity in its highest combination with oxygen, such as NO2 (OH).
4 DISCUSSIOX OF THE PERIODIC SYSTEM. §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 VII. (not in VIII.), 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 facts 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, 6, 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 VIII. 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 group of elements so probable that it has been placed in the
chart for convenience of study, subject to further conclusions. (W. Ramsay.
Br. Assoc. Adv. ScL, 1897, 598-601 ; B. 1898, 31, 3111. J. L. Howe, C. N.,
1899, 80, 74; 1900, 82, 15, 52. Ostwald, Grundr. Allg. Chem., 3te Auf.,
1899, S. 45.) In comparison with the members of Group Vll. those of
Group VIII. 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 VIII. is the fact that the return to Group I. from
Group VIII. is less complete than the return from Group VII. That is to
say, the character of copper is divided between Group VIII. and Group I..
and the same is true of silver and of gold. This relation to Group VIII.
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 long
period of seventeen members, Series 6 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 VII., falls
in two columns, a column consisting of the alternate elements in the group.
Thus, H, Li, K, Rb 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. DISCUSSION 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 Alkali Metals.— Li 7.03, (Na 23.05), K 39.11, Rb 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. The Alkaline Earth Metals.— (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. The Zinc Family.— Mg 24.3, (Al 27.1), Zn 65.4, Cd 132.4, - — .
Hg 200.0. These metals, save aluminum, belong to the second alternates of
Group II., and, like those of the corresponding half of Group I., 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,
eadmium and mercury insoluble in water, and these three 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. Mercury, 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 D7,S'frx,x70\ (>/•' 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 nnd bromine, are
counted as metals, is not contrary to the periodic order. Group VI1T. binds
Group I. to Group VII. After Co 59.00 follow Cu <>3.(> and then Zn (55.4.
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. Hie halogens. Of the five, chromium and man-
ganese (nearest the halogens) form the best known acids. Nickel and
cobalt, like copper, 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 be written in symbols as follows:
2-3-6 2-.1-4-0-7 2-3-6 2-3 2-3 1-2 2
Cr , Mn , Pe , 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 water. In the conditions of precipitation sulphides are not formed
with the metal in any valence other than two. Chromium acting us a
base with a valence of three, like aluminum whose only valence is three.
refuses to unite with sulphur. Trivalent iron in precipitation by sulphide*
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 Ifeiftl* not Alkalis in Group /., Second !'«>•/, and their Re1ntu-<>*
in Group VIII.— Cu (i.'i.K, Ag 107.92, , Au 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 the 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. arc 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.(5 with Ni 58.70 and Co 59.00, Ag 107.92 with Pd 107.0, and
Au 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.— N 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 Y. 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 - 3 . HHH. In ammonium chloride : N - 4 + x = - 3 . HHHHC1. 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. Relation 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 V. 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 DISCUSSION OF THE PERIODIC SYSTEM. §12.
sire lead and bismuth. In further comparison, arsenic, in its deportment
iis 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 forms 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, Sb,
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.
Xitrogen 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. Th •
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, HN . of decided
acid power, constituting well marked salts, such as Na N3, in which a ring
of nitrogen alone acts as an acid radical. The first four member* of the
nitrogen family agree with each other in forming 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 oxygon
with two units each of negative valence and a unit of negative valence
of hydroxyl: H — 0 — N^Q. The same constitution is found in mctaphos-
phoric acid HO P 02 , meta-arsenic acid HO As 0., , and in antimonic acid
HO Sb 02. The so-called ortho acids, phosphoric and arsenic, have the
constitution (HO)., P 0 and (HO), As 0 , 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,' AsH, 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. DISCUSSION 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 phosphates 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, Cl 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, HC1 , 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 HgI2 and K2 Pt C1G . 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, HO Cl 02 ; bromic, HO Br 02 ; and iodic, HO I 02. Chloric acid
agrees with nitric acid, HO N 02 , in the fact that it forms soluble salts with
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 03 , also (HO)., I 0 , is pre-
served intact without difficulty.
§14. The Relations of Sulphur.— S 32.07. Sulphur is the first member
of a family including selenium and tellurium. It differs from oxygen
almost as much as phosphoms differs from nitrogen, and we may say more
than silicon differs from carbon. The higher valence of Group VI., excr-
10 THE CLASSIFICATION OF THE METALS AS BASES. §15.
cised toward oxygen, cannot be met by oxygen itself. Of the acids of
sulphur, HJ5 , in which sulphur has two electro-negative units of valence,
is quite unstable, while (H0)2 SO.,, in which the sulphur has six electro-
positive units of valence, is the most stable. The sulphides (salts of H..S)
of the heavier metals quite generally are insoluble in water, an important
means of separation in analysis. The sulphates (salts of H2S04) of the
larger number of the metals are soluble in water, the exceptions being
important 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 C2 H., (H C EEC H). 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 sulphur 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 the 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 analytical
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 iL'ith the properties of the leading bases and adds,
THE CLASSIFICATION OF THE METALS AS BASES.
11
in their action upon each other. Without this acquaintance, the analyst is
the servant of routine, and his results liable to fallacy.
The following named are the bases of more common occurrence.
The Alkali Bases.
The sixth group.*
Potassium
Sodium (Natrium), Na1.
Ammonium,
Not precipitated from their salts
by any of the group 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 combination in potassium hy-
droxide, KOH , and in potassium
salts, such as the chloride KC1 , and
the nitrate, KN03 .
In the base, sodium hydroxide and
its salts.
Forms ammonium hydroxide.
NH4OH , representing ammonia,
NIL , and water, and serving as the
base of ammonium salts, such as
(NH4).,S04 , ammonium sulphate.
(Precipitated by carbonates, which
fact alone does not separate them
from the following named groups.)
Separated by precipitation as a
phosphate afler removing all the fol-
lowing named bases. Forms magne-
sium hydroxide. Mg(OH)2 , and mag-
nesium salts, such as MgS04 .
Separated by precipitation with
Ammonium Carbonate, adding
NH4C1 to keep magnesium from pre-
cipitation. Calcium carbonate, a
normal salt, CaCO, .
* The sixth division of the bases, in the order in which they are separated from each other by
precipitation with the group reagents.
t The Roman numerals (as i) express units of valence, each equivalent to an atom of
hydrogen, in the formation of salts and other combinations.
The Alkaline Earth Bases.
The fifth group.
Magnesium, Mgn.
Calcium. Strontium, Barium, Ca11.
Sr11, Ba11.
CLASSIFICATION OF THE METM^K AS 7?.t,S'/:x.
§16.
The Zinc and Iron Groups.
The Zinc Group.
The fourth group.
Znn: zinc salts.
Ifn": manganous salts.
Mnm: manganic salts..
Mnlv: salts unstable.
MnVI: salts of manganic acid.
Mnvn: salts of permanganic acid.
Hi": nickel salts.
Co11: cobaltous salts.
Co111: cobaltic salts.
The Iron Group.
The third group.
Fe11: ferrous salts.
Fem: ferric salts.
Crni: cbromic salts.
CrVI: chromates.
A1HI: aluminum salts.
Metals falling with Copper and Tin.
The second group.
The Copper Group.
Division B, second group.
4
Mercury (Hydrargyrum).
Hgn: mercuric salts.
Hg1: mercurous salts.
Silver (Argentum).
Ag1: silver salts.
Lead (Plumbum).
Pb11: lead salts.
Bim: bismuth salts.
Cu": copper or cupric salts.
Cu1: 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 with
Ammonium Sulphide, after removal
of all the folloiving named bases as
directed below. (The precipitates
are all sulphides.)
Separated by precipitation with
Ammonium Hydroxide, in presence
of NH4C1 , after the removal of the
groups named following. (The pre-
cipitates are all hydroxides.)
Precipitated by H,S in acidulated
solution. (The precipitates are sul-
phides.)
Separated by the insolubility of
the precipitated sulphides in treat-
ment with Ammonium Sulphide.
THE OPERATIONS OF ANALYSIS.
13
The Tin Group.
Division A, second group.
Sn11: stannous salts.
SnIV: stannic salts and stannates.
. Sbm: antimonons compounds.
Sbv: antimonic compounds.
As111: arsenous compoimds.
Asv: arsenic compounds and arsen-
ates.
Metals Precipitated as Chlorides.
The .Silver Group.
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 CEKTAIN COMMONLY OCCUKKING SALTS.
Name of Acid.
Name of Salt.
Formula.
Showing Hydroxyl.
Anhydride.
Carbonic
Carbonate
H2C03
(HO)2CivO
CO2
Oxalic
Oxalate
H2C204
(HO):!C2ivO2
C203
Nitric
Nitrate
HN03
(HO)NVO2
N205
Nitrous
Nitrite
HN02
(HO) NinO
N203
Phosphoric (ortho)
Phosphate
H3P04
(HO)SPVO
P;;05
Metaphosphoric
Metaphosphate
HP03
(HO)Pv02
P,OB
Pyrophosphoric
Pyrophosphate
H4P207
(HO)4PV20S
P.O3
Sulphuric
Sulphate
H2S04
(HO)2Svi"O2
S03
Sulphurous
Sulphite
H2S03
SO.
Hydrosulphuric
Sulphide
H2S
Hydrochloric
Chloride
HC1
Hydrobromic
Bromide
HBr
Hydriodic
Iodide
HI
Chloric
Chlorate
HC1O3
(HO)ClvO2
C12O5
lodic
lodate
HI03
(HO)IVO2
I.O5
THE OPEEATIOXS 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 have formed it. For example, the analyst may
have in hand a mixture containing sodium sulphate, Na.,S04 ; sodium sul-
phite, Na.,80, , and sodium thiosulphate, Na2S203 , hut 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, either 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-
lined 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, 55/255; Welsbach, M.. 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. pJiys. Ch., 1896, 20, 44).
Again, traces of dissolved matters too minute for other means of detection
~ / •
can be revealed by the difference of electric potential between electrode and
solution (Ostwald, Lelirl., 2 AufL, II, 1, 881; Behrend, Z. pliys. Ch., 1893,
11, 466; Hulett, Z. pliys. CJi., 1900, 33, Gil).
§25. 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. When 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 loniza-
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 be
placed among the substances engaged in chemical change.
§27. No other property of substances has so great importance in analysis
and in all chemical operations, as their solubility in water. It must never
be forgotten that there are degrees of solubility, but there is hardly such a
fact as absolute solubility, or insolubility, regardless of the proportion
of the solvent. There are liquids which are miscible with each other
16 THE OPERATIONS OF ANALYHfX. §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
precipitates obtained, requiring about 430,000 parts of water for its solution
at ordinary temperature (Hollemann, Z. phys. Cli., 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, 5 b); 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 i*
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 OPERATIONS OF ANALYSIS. 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
dissolves. 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-
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. Neglect 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 visible 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 experience, that 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 OPERATIONS OF ANALYHfX. $34.
with help of a microscope and with studies of crystalline form. Further
see Behrens, Z. 1891, 30, 125; and Herrnschmidt and Capelle, Z. 1893, 32,
608.
§34. Precipitates are removed — usually by filtration, sometimes by dccan-
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 precipitations, 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 accomplished, while the operator, proceeding
just as though it was accomplished, undertakes to identify the members of
a group by reactions on a mixture of groups. The washing, on the filter,
is best completed by repeated additions of small portions of water — around
the 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.
$35. In dissolving precipitates — 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 be operated upon, it usually has to be neutralized, and
the solution then becomes so greatly encumbered and diluted that reactions
become faint or inappreciable. Precipitates may be dissolved on the filter,
without excess of solvent, by passing 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 Avater. Washed precipitates may also be dissolved in the
test-tube, by rinsing them from the filter, through a puncture 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 sufficient
reagent to cause the desired change should always be governed by testing
a drop of the solution, on a glass rod, with a piece of litmus paper.
£37. When 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 practically complete when one or
§40. THE OPERATIONS 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 -f- Hydrochloric acid = Cal-
cium chloride -f 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 be 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 : BaCl2 -f MgS04 = BaS04 + MgCl2 , 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 = FeCl2 + H2 .
56 + 72.9 = 126.9 -f 2 .
Also in precipitating ferrous chloride by sodium phosphate, we have the
equation:
Fed,, + Na2HP04,12H2O = FeHPO4 + 2NaCl + 12H20 .
126.9 + (142.1 + 216) = 152 + 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 = 100/x = parts of absolute HC1 for 100 parts of iron wire.
" 1 32/100 = x/y = parts of 32 per cent HC1 for 100 parts of iron wire.
! 126.9/72.9 = 100/x
32/100 = x/y = parts of 32 per cent HC1 for 100 parts of FeCl2 , absolute.
126.9/56 = 100/z = parts of iron wire for 100 parts of FeCL.
{152/72.9 = 100/x
32/100 = x/y = parts of 32 per cent HC1 for 100 parts of FeHP04 .
152/56 = 100/z = parts of metallic iron for 100 parts of FeHP04.
152/358.1 = 100/u = parts of Na,HP04,12H2O for 100 parts of FeHP04.
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 formula4 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 AXD IONIZATIOX.
§43. The Theory of Electrolytic Dissociation, proposed by Arrhenius in
1887 (Z. pliys. Ch., 1887, 1, 631), assumes that salts, acids, und 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 hased 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. pliys. Cli., 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 (NH4C1 , PC13) 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 saponifi cation 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 undergo
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
throughout the solvent, just as a gas by virtue of the gas-pressure occupies the entire space at
its disposal.
22 SOLUTION AND ION fZ AT I ON. £44.
into ions groups together and gives system and meaning to these three
classes of facts, experimentally absolutely independent and up to Arrhenius'
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 are 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.,8 does not give the same precipitation
reactions as that in K,S or H,S04 or H,SO, or H,S,0:! . Further, HgCL 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 l>y 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. CdCl., 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 Berthollet 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 CH ,CO,H + C..H.OH =
CH,COoC,H. -f H20 , when equilibrium sets in ab = kcd , 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. Where the
* The unit of quantity is the molecular weight taken in grams (the " mol "). Where there are
18.23 grams HC1 in a liter either in solution or as gas the concentration Is %, where there are
72.92 grams in the same volume the concentration is 2. and so on.
£45. SOLUTIOy AND WNIZATIOX. 23
reaction is a dissociation, as with gaseous NH4C1 , we have ab = k'c , a and b
representing the concentrations of NH3 and HC1 respectively, c that of the
undecomposed NH4C1 , 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 product? of the reaction and thus increase its concentra-
tion, the concentration of the other product must decrease in the saine^
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 NIL
or HC1 diminishes the dissociation of NH4C1 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 H2S, but on addition of sodium acetate the black FeS
is brought down. Similarly a weak base, as NH4OH , is made still less
effective by the presence of its strongly-dissociated neutral salt, as NH4C1 .
Quantitative agreement is obtained between observed effect of NH4C1 on
NH4OH as saponifying agent and that calculated from the equation:
CNH4 ' COH' = kCNH4OH (Arrhenius, Z. phys. Cli., 1887, 1, 110).
§45. The Solubility-Product. — In the saturated solution which always
remains after precipitation we have the usual dissociation equilibrium, as:
CAe ' °Cr ^ffCl ' -^ow *ne 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 (Nernst).
This Ostwald has called the Solubility-Product. Where the saturated solu-
tion is made by bringing the salt into contact with the solvent c^g. • " QV -
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 (Nernst, Z. phys. Ch., 1889, 4, 372;
^oyes, Z. phys. Ch., 1890, 6, 241; 1892, 9, 603).
The Solubility-Product of the alkaline-earth carbonates is
°M " CCO " ~~ IR the solution of a neutral salt, as CaCL , Ca ions are
present in large concentration. When a substance containing C03 ions in
large concentration is added, as Na^CO., , 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 C03 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.
^r
For the metallic sulphides the solubility-product is jj •• g» ~ •
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 H..S 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 HC1 .
containing as it does H ions, one of the dissociation products of H.S , 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 are 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.
fr. 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 reageni. The fact of the reaction should be staled; e. g.,
lead acetate -f- 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(C2H,O2)3 + 4KOH = Pb;!0(OH):,* (white) + 4KC3H,0, + H.O
Pb,O(OH), + 4KOH (excess) = 2K2PbO, + 3H2O
or Pb(C2H,O2)2 -f 4KOH (excess) = K.PbO, + 2KC,H302 + 2H2O .
* It has been found helpful to require students to underscore all precipitates.
£46. ORDER OF LABORATORY STUDY. So
The results should all be tabulated and then summarized in form of a
carefully worded generalization (§205, Ga).
c. Action of Ammonium Hydroxide (volatile alkali) upon solutions of
the salts of the metals, etc., as in (b) above ; e. g., lead nitrate -j- ammonium
hydroxide = a white precipitate not dissolving in excess. Consult text
(§57, 6a) and write the equation :
3Pb(N03)2 + 41TH4OH = 2PbO.Pb(N03)2 + 4NH4N03 + 2H2O .
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 (§205, Ga).
e. A study of the action of Ammonium Carbonate. Summarize the re-
sults (§207, Ga).
f. A study of the solvent action of acids, HC1 , HN03 , and H2S04 , upon
the Hydroxides and Carbonates obtained by precipitation.
g. Action of Hydrosulphuric Acid as a precipitating agent upon salts of
the metals in neutral and acid solutions.
h. The use of Ammonium Sulphide as a reagent.
i. The solvent action of acids, HC1 , HNOr, and HC2H302 , upon the
sulphides obtained by precipitation.
j. 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.
I. 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 -f- calcium chloride — a
white precipitate. Consult the text (§227, 8), and write the equation:
(NH4)2C204 + CaCl2 = CaC204 -f 2NH4C1 .
n. Application of the above reactions to the Grouping of the Metals
for Analysis.
o. 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 effecting their separation and detection. The new work of
each day to be followed by the analysis of " unknown " mixtures prepared
by the teacher to illustrate the 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.
26 ORDER OF LABORATORY STUDY. $46.
q. The study in the class room of Oxidation and Reduction, 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 " Unknown " Mixtures.
t. A special study of the analysis of Phosphates, Oxalates, Borates,
Silicates, etc., and certain of the Rarer Metals.
n. The analysis of mixtures in solution, illustrating Oxidation and
Reduction.
v. A study of Electrolysis as a means of detection in qualitative analysis.
PAKT II.-THE METALS.
THE SILVER AND TIN AND COPPER GROUPS.
(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 upon 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* Pb 206.92 Germanium Ge 72.5
Mercury Hg 200.0 Iridium Ir 193.1
Silver Ag 107.92 Osmium Os 191.0
Arsenic As 75.0 Palladium Pd 107.0
Antimony Sb 120.4 Rhodium Bh 103.0
Tin Sn 119.0 Ruthenium Bu 101.7
Gold Au 197.2 Selenium Se 79.2
Platinum Pt 194.9 Tellurium Te 127.5?
Molybdenum Mo 98.0 Tungsten W 184.
Bismuth Bi 208.1 Vanadium V 51.4
Copper Cu 63.6
Cadmium Cd 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 Silver. Tin and Copper 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 § 1O4.
28 GENERAL 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, 5, 199, and Cohen,
J., Soc. Ind., 1891, 10, 327 (§12).
§50. 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 bismuth
belong to the Nitrogen Series of Elements.
§51. A large proportion of the compounds of these metals are insoluble
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, borat.i >.
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.
$52. 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 ehlorides, 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.
§54. 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. A (tin group) — As , Sb , Sn , Ge , Au , Ir , Mo . Pt , Se ,
Te , W . and V ; sulphides soluble in i/t'Ihur finininninin sulphide; and /?
(copper group) — Hg , Pb , Bi , Cu , Cd , Os , Pd , Rh , and Ru ; sulphides
not soluble in ijt'lloir ammonium yulpltidf.
§55. Mercury, antimony, silver, and gold do not form hydroxides. The
oxides of gold are very unstable.
§56. The metals of these groups are all easily reduced to the metallic
state by ignition on charcoal. Except mercury and arsenic, which vaporize
£57, 4. LI: AD. 29
readily, and certain rarer metals difficultly fusible, the reduced metals melt
to metallic grains on the charcoal.
THE SILVER GROUP (FIRST GROUP).
Lead, Mercury (Mercurosum), Silver.
§57. Lead (Plumbum) Pb = 206.92 . Valence two and four.
1. Properties. — 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 (Schuller, B., 1883,
16, 1312).
i'ure 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 pewter; 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., 1866, (4), 31.
176); its most abundant ore is galena, PbS ; it also occurs as cerussite. PbC03 :
anglesite, PbSO4; pyromorphite. 3Pb,PaO, + PbCL; krokoite. PbCr04; 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 PbSO4 , 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
PbSO4 , SO2 being formed: PbS04 + PbS = 2Pb + 2SO2 . 2PbO + PbS = 3Pb +
SO2 . (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 -f FeS . Frequently these methods are combined or varied according to
the other ingredients of the ore.
4. Oxides.— Lead forms four oxides, Pb2O , PbO . PbO, , and Pb3O4 . Lead
suboxide (Pb2O) is little known: it is the black powder formed when PbC204 is
heated to 300°, air being excluded. Lead oxide (litharge, or massicot) is formed
by intensely igniting in the air Pb , Pb,O , PbO2 , Pb3O4 , Pb(OH), . PbCO:; .
PbC2O4 , or'Pb(NO:1)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), Pb304 , is formed by heating PbO
to a dull-red heat Avith full access of air for several hours. Strong, non-reduc-
ing acids, such as HNO3 , H,S04 , HC1O3 , etc., convert it into a lead salt and
PbO, (a). But concentrated hot H,SO4 converts the whole into PbS04 , oxygen
being evolved (ft). But with the dilute acid and reducing agents, such as
C3Hr,(OH)3 , C,1H1,0,,.H,C;;04 , H,C4H40(i , Zn , Al , Cd , Mg . As, Pb , etc..
it is all reduced to the dyad lead without evolution of oxygen (<•), (</), and (f).
Hydracids usually reduce the lead and are themselves oxidized (/").
(a) PbsO4 + 2H2S04 (dilute) = PbO, + 2PbSO4 + 2H,O
2Pb304 + 6H2S04 (concentrated and hot) = GPbSO4 + 6H2O + O2
H,C204 + 6HN03 = :{Pb(NO,)2 + 4H2O + 2C02
As4 + 30ELS04 = 30PbS04 + 4H3AsO4 + 24H2O
Zn + 4H2S04 = 3PbSO4 + ZnS04 + 4H2O
8HC1 = 3PbCL, + C12 + 4H,,0
The valence of Pb304 is best explained by the theory t-hat it is a union of the
dyad and tetrad (Pb" and Pbiv) , Pb3O, = :2PbO + PbivO2 .
30 LEAD. §57, 5a.
Lead dioxide or peroxide, PbO, , is formed: (1) by fusion of PbO with KC1O,
or KNO3 ; (2) by fusing- Pb3O4 with KOH : (3) by treating- any compound of
Pb" with Cl , Br, K3Fe(CN),, , KMnO , or H,O2 in presence of KOH; (4) by
treating Pb3O4 with non-reducing- acids:
Pba04 + 4HNOS = PbO, + 2Pb(NO,)s + 2H2O.
Ignition forms first Pb304 and above a red heat PbO, oxygen being given off.
It dissolves in acids on same conditions as Pb304 . Very strong solution of
potassium hydroxide, in large excess, dissolves it. with formation of " potassium
plumbate," K,?bO8 . Lead dioxide is a powerful oxidizing agent, one of the
strongest known. Dige&ted 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. — (t. — Met"7 — 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 concentrated 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 (c). 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.
ft. — O-riV/r*.— Lead oxide, titharur. PbO. and the hydroxides, 2PbO.H,O;
.'iPbO.H.O, are readily dissolved or transposed by acids forming the correspond-
ing salts, i. e., PbO + H,SO, = PbSO4 + H,O . The oxide and hydroxide are
soluble in about 7000 parts of water, to which they impart an alkaline reaction.
They are soluble in the fixed alkalis forming plumbites; soluble in certain salts
as NH4C1. CaCl, , and SrCl, (Andre, ('. r., 1883. 96, 435; 1887, 104, :i59); very
soluble in lead acetate, forming basic lead acetate.
Lead dioxide, PbO. , h'<td p?ro.ridc, is insoluble in water or nitric acid; it is
dissolved bv the halogen hvdracids with liberation of the halogen and reduction
of the lead forming a dyad salt: PbO, + 4HC1 = PbCL + Cl, + 2H,O: it is
attacked by hot concentrated sulphuric acid, forming the sulphate and liberat-
ing oxygen; it is soluble in glacial acetic acid forming Pb(C.,H30,)4 . unstable
(Hutchinson and Pollard, J. ('.. 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. They are, however, very unstable, being decomposed to the dyad salt
upon warming (Fischer. ./. r..'lS79, 35, 282; Nickels, A. Cl\., 1867, (4), 10, 328).
The peroxide is slowlv soluble in the fixed alkali hydroxides forming plum-
bates, /. c.. PbO, + 2KOH = K,PbO3 + H,O .
Trilead tetroxide, Pb30, , red lend, minium, is insoluble in water, is at-
tacked by nearly all acids in the cold forming the corresponding dyad lead
salt and 'lead peroxide, PbO, . 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 lessened by the presence of another substance having an ion in
common with it <§45). In some cases, as with Pbl, and KI, this is offset in concentrated solution
by formation of a complex compound.
£57, 5c. LEAD. 31
facilitates the solution of red lead or lead peroxide in acids, i. 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
iiitrate, which greatly facilitates the further analysis.
c. — Salts. — The carbonate, borate, cyanide, ferrocyanide, phosphate, sul-
phide, sulphite, iodate, chroniatc, and tannate are insoluble in water.
The sulphate is soluble in about 21,000 parts of water at 18° (Kohlrausch
and Rose, Z. phyy. Ch., 1893, 12, 241), the presence of HNO;! or HC1 in-
creases its solubility in water; it is insoluble in alcohol even when quite
dilute; sparingly soluble in concentrated H2S04 , from which solution it is
precipitated by the addition of water or alcohol; less soluble in dilute H.,S04
than in water; soluble in 682 parts 10 per cent HC1 , in 35 parts 31.5 per
cent (Rodwell, J. C., 1862, 15, 59); transposed and dissolved by excess of
HC1 , HBr , or HI forming the corresponding haloid salt; insoluble in
HF (Ditte, A. Ch., 1878, (5), 14, 190); soluble in ammonium sulphate,
nitrate, acetate, tartrate and citrate, and from these solutions not readily
precipitated b}r ammonium hydroxide or sulphate (Fleischer, J. C., 1876,
29, 190; Woehler, A., 1840, 34, 235).. The sulphate is almost completely
transposed to the nitrate by standing several days with cold concentrated
nitric acid (Rodwell, I. 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 166
parts water at 10°, in about 45 parts at 80°; the iodide is soluble in 1235
parts water at ordinary temperature, and in 194 parts at 100° (Uenot, J.
pr., 1834, 1, 425). The chloride is less soluble in dilute HC1 or H,S04 than
in water, but is more soluble in the concentrated acids (Ditte, I. 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 NH4C1 —90 grams dissolving in 200
grams NH4C1 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 precipitated
on adding solutions of KNO., to a solution of basic lead acetate^
The relative insolubility of PbCL in cold water or in dilute HC1 makes
it possible to precipitate the most of the lead (by means of HC1) 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
chlorides of that group (§61). The lead is separated and identified in
the second group as the insoluble sulphate. (§95).
32 LEAD. §57, t!.
6. Reactions, a. — Fixed alkali hydroxides precipitate, from solution.- , . I
lead salts, basic lead hydroxide (!}, Pb,0(OH), (Schaffner, A., 1844, 51, L75),
white, soluble * in excess of the reagent as plumbite (^) (distinction from
silver, mercury, bismuth/ copper, and cadmium). The normal lead hy-
droxide, Pb(OH)o, may be formed by adding a solution of a lead salt to
a solution of a fixed alkali hydroxide.
(1) 2Pb(NO3)2 + JKOH = PbJ0(OH), + 4KN03 + H20
(2) Pb,0(OH),,
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
PbCL.PbO.ILO (Wood and Bordeu, C. J\7., 1885, 52, 43); with the nitrate
2PbO.Pb(NO,), (7?., 2, 2, 558). 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 had 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.!(OH),(CO.,)., . Precipitation in the cold approaches more nearly to the
normal carbonate (Lefort, Pharm. J., 1885, (3), 15, 2(5). Solutions of lead
salts when boiled witli freshly precipitated barium carbonate are com-
pletely precipitated. Carbon dioxide precipitates the basic acetate but
not completely.
6. — Oxalic acid and alkali oxalates precipitate lead oxalatf, PbC.,04, white.
from solutions of lead salts, soluble in nitric acid, insoluble in acetic acid.
A solution of lead acetate precipitates a large number — and a solution of
lead subacetate a still larger number — of organic acids, color sulxtancc-.
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 lend ci/nnidr. Pb(CN); , white, sparingly soluble
in a large excess of the reagent and reprecipitated on boiling. Potassium ferm-
cyanide precipitates lead ferrocyanide. Pb^FeCCN),, , white, insoluble in water
or dilute acids. Potassium ferricyanide precipitates from solutions not too
dilute lead ferricyanide, Pb, (Fe(CN),s)_. . white, sparingly soluble in water.
soluble in nitric acid. Solutions of lead salts are precipitated by potassium
sulphocyanate as lead sulphocyanate, Pb(CNS): . 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
* Nearly all the salts of lead are soluble in the fixed alkali hydroxides, P1>S forming almost
the only notable exception.
£57, 6e. LEAD. 33
potassium 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. phys. Ch., 1890, 6, 385).
$ The higher oxides of lead are all reduced by hypophosphoroux acid, lead
phosphate being formed. Lead phosphite, PbHPO3 , white, is formed by
nearly neutralizing phosphorous acid with lead carbonate or precipitating
Na,HPOs with Pb(NO3), (Amat, C. r.. 1890, 110, 901). Sodium phosphate,
Na"HPO4 , precipitates from solutions of lead acetate the tribaslc lead phoxpliate,
Pb3(P04).> , white, insoluble in the acetic acid which is set free (D., 2, 2, 562):
3Pb(C,H,O,), + 2Na,HPO4 = Pb,(PO4)a + 4NaC,H302 + 2HC3H,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 pyrophosphatc, Pb2P2O7 , white, amorphous, is formed by precipitating a-
lead solution with Na4P2O7 , soluble in excess of the precipitant, in nitric acid,
and in potassium hydroxide: insoluble in ammonium hydroxide and in acetic
acid (Gerhardt, A. Ch., 1849, (3), 25, 305). The metaphosphate, Pb(PO3)2 ,
white, crystalline, is obtained by the action of NaP03 upon Pb(NO3)2 in excess.
e. — Hydrosulphuric 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 , MnS , FeS , CoS , and NiS also
jj^ive 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 (!}, some of the sulphur, especially if the nitric acid
be concentrated, is oxidized to sulphuric acid, which precipitates a portion
of the lead (2), 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.
(1) 6PbS + 16HNO3 = GPb(NO3), + 3S2 + 4NO + 8H2O
(2) 3PbS + 8HNO3 = 3PbSO4 + 8NO + 4H2O
In solutions too strongly acidulated, especially with hydrochloric acid,
either no precipitation takes place, or a brick-red double salt, PboSCL ,
* The condition for equilibrium is that a certain r.itio-of concentration exist between the ions,
in the case of PbSO4 between the S ions and the SO4 ions. These concentrations are the same
as those in a solution obtained by digesting the two salts, PbSO4 and PbS, together in water.
PbSO4 dissolves more freely than PbS. and for equilibrium therefore cso /, must be corres-
pondingly greater than Cg,,. But adding H2S or a soluble sulphide to PbSO4 gives just the
opposite of this condition, and transformation accordingly results, increasing the SO4" con-
centration by formation of soluble sulphate and decreasing the S" concentration by precipita-
tion of PbS, until the equilibrium-ratio is produced or, if the quantity of Pi>SO4 present is in-
sufficient for this, until all the PbSO4 has been transformed to sulphide. On th^ other hand,
treatment of PbS with a very large excess of H2SO4 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
always be formed. This principle determines the direction in which a reaction takes place;
AgCl + KI = Agl + KC1 5 CaSO4 + JVa2CO3 = CaCO3 + Na2SO4 (§44).
•34 LEAD. §57, 6/.
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. O., 1889, (5(57).
When galena, PbS , is pulverized with fused KHSO, , H2S is evolved (Jan-
nettaz, J. C., 1874, 27, 188).
Lend tliiottulphate, PbS2O3 , 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, PbS,O,,2Na,,S^O3 (Lenz, A., 1841, 40,
94); on boiling, all the lead is slowly precipitated as sulphide (Vohl, A., 1855,
v'O. '_'...).
Sodium sulphite precipitates lead sulplnie, PbSO., , white, less soluble in
water than the sulphate, slightly soluble in sulphurous acid; decomposed
by sulphuric, nitric, hydrochloric, and hydrosulphuric acids and by alkali
sulphides; not decomposed by cold phosphoric and acetic acids.
Sulphuric acid and soluble sulphates precipitate from neutral or acid
solutions, lead sulphate. PbS04 , white, not readily changed or permanently
dissolved by acids, except hydrosulphuric acid, yet slightly soluble in
strong acids (or). 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 sulphide, insoluble in thiosulphate, being formed (dis-
tinction and separation from barium sulphate, which does not dissolve in
thiosulphates).
The test for lead as a sulphate is from five to ten times less delicate
than that with hydrosulphuric acid: but lend is quantitatively separated
as a sulphate, by precipitation with sulphuric acid in the presence oi'
alcohol, and wa.»hing with alcohol. When heated with potassium eliminate
transposition takes place and yellow lead chromate is formed (//). 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 completely transposes the lead to the chloride (Matthey. J. C'..
1879, 36, 124). See footnote on previous page.
f. — Hydrochloric acid and soluble chlorides precipitate, from solutions
not too dilute, lead chloride, PbCl.. . white. This reaction constitutes lead
a member of the FIRST GROTJP-a- 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 tlir second t/rnn/i.
£57, 7. LEAD. 35
Hydrobromic acid and soluble bromides precipitate lead bromide, PbBr2 ,
white, somewhat less soluble in water than the chloride (5c); soluble in
excess of concentrated potassium bromide, as 2KBr.PbBr2 , which is decom-
posed and PbBr2 precipitated by dilution with water.
Hydriodic acid and soluble iodides precipitate lead iodide, PbI2 , bright
yellow and crystalline, much less soluble in water than the chloride or
bromide (5c); 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, KIPbL with small excess of KI , and 4KI.PbI,
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, Pb02 .
gives 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.
</. — Arsenous acid does not precipitate neutral solutions of lead salts; from
alkaline solutions or with soluble arsenites a bulky white precipitate of lead
arscnlte is formed, insoluble in water, but readily soluble in all acids and in the
fixed alkali hydroxides. Arsenic acid and soluble arsenates precipitate lead
arscnate, white, from neiitral 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 (SnCL 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 K2Cr207 —
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, PhO (a). After fusion and diges-
tion with warm water, the aqueous solution is t'ested 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 (&); and with excess of charcoal the acid radicle may also be
changed (r). 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 .
(a) PbCL + NmCO, = 2NaCl + PbO + CO,
(?>) 2PbSO4 + 2Na2CO3 + C = 2Pb + 2Na,SO4 + 3COt
(c) 2PbS04 + 2Na,C03 + 5C = 2Pb + 2Na2S -f 7C02
8. Detection. — Lead is precipitated, incompletely, from its solutions by
HC1 as PbCL ; separated from AgCl and HgCl by hot water, and confirmed
by H2S , H2S04 , K2Cr04 , and KI . It is separated (in the second group)
from As , Sb , Sn . etc., by non-solubility of the sulphide in (NH4)._.SX ;
from HgS by HNO ; from Bi , Cu , and Cd by precipitation with dilute
sulphuric acid. Insoluble compounds are 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. Estimation. — (a) As an oxide into which it is converted by ignition (if a
carbonate or nitrate), or by precipitation and subsequent ignition, (ft) As a
sulphate. Add to the solution twice its volume of alcohol, precipitate with
H SO, , and after washing with alcohol ignite and weigh, (c) It is converted
into an acetate, or sodium acetate is added to the solution, then precipitated
with KoCTjO, , and after drying at 100°, weighed as PbCrO4 . (d) It is con-
verted into PbS , free sulphur added, and after ignition in hydrogen gas
weighed as PbS . (r) The lead is precipitated with standardized sodium iodate
and the excess of iodate 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 siiblime upon
ignition (Steen, Z. angew., 1895, 530). (g) Qax coluinetric metlmd. 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. angeic., 1891, 329).
10. Oxidation. — Metallic lead precipitates the free metals from solutions
of Hg . Ag , Au , 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., 1893, 646). Pblv is reduced to Pb° in presence of dilute
H.,S04 by nascent hydrogen, and by all metals capable of producing nascent
hydrogen (such as Al , Zn , Sn , Mg , Fe), and to Pb" bv sohible compounds
of Hg', Sn". Sb'", As'", (AsH, gas), Cu', Fe", Cr"', Mn", Mn'", Mniv.
Mnvl. Also by HaC,04 , HNO, 'I H.PO, , H;?PO:! , P , SO, , H,S , HC1 , HBr ,
HI, HCN, HCNS, H,Fe(CN),, , glycerine, tartaric acid, sugar, urea, and
rrri/ 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, NTT, , to nitrite and nitrate in the course of a
few hours (Pollacci, ArcJi. Pharm., 1886, 224, 176).
From lead solutions Zn , Mg , Al , Co , and Cd precipitate metallic lead.
.§58, 5a. MERCURY. 37
§58. Mercury (Hydrargyrum) Hg = 200.0 . Valence one and two.
1. Properties. — Specific gravity, liquid, 13.595:$ (Volkmann, W. A., 1881, 13, 209);
solid, 14.1932 (Mallet, Proc. R, »S'oc., 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
ordinary 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. c., reduced to the finely divided state, by shaking with sugar, grease,
chalk, turpentine, ether, etc. It is slightly volatile even at — 13° (Regnault,
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,
octahedral and needle-shaped crystals, 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 , Ni , Os ,
Pd , Pt , K , Ag . Na , Tl , Sn . and Zn . An amalgam containing about 30
per cent of copper is used for filling teeth (Dudley, Proa. Am. Assc. for Adv. o/
Sci., 1889, 145).
2. Occurrence. — 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 +
O2 = Hg1 + SO., . (b) Lime is added to the ore, which is then distilled;
4HgS + 4CaO = 3CaS + CaSO4 + 4Hg . (c) The ore is heated with iron
(smithy scales): Hg , FeS , and SO2 arc 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 (BL, 1888, (2), 49, 856).
4. Oxides. — Mercury forms two oxides. Hg.O and HgO . Mercurous oxide,
Hg2O , is a black powder formed by the action of fixed alkalis on mercurous
salts. It is converted by gentle heat into Hg and HgO and by a higher (red)
heat, to Hg and 0 . Mercuric o.ride, HgO . is made (/) by keeping Hg at its
boiling point for a month or longer in a flask filled with air; (2) by heating
HgNO3 or Hg(NO3). with aboiit an equal weight of metallic mercury:
Hg(NOn)2 -f 3Hg = 4HgO + 2NO; (3) by precipitating mercuric salts with
KOH or NaOH . Made by (1) and (2) it is red, by (3) 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 . Mercury 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-(NOs), , if there is sufficient HNO3 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
crj-stallizes out, which a fresh supply of nitric acid will dissolve. Sxilphur
attacks mercury even in the barometric vaciuim, forming HgS (Schrotter,
/. C., 1873, 26, 476). H.SO, concentrated at 25° has no action on Hg (Pitman,
38 MERCURY. £58, 56.
J. Am. Soc., 1898, 20, 100). With the hot concentrated acid SO, is evolved and
Hg SO, is formed if Hg be in great excess: HgSO, if the HSSO4 be in exces-s.
Hydrochloric acid gas at 200° is without action (Berthelot, A. T/*., 185(5, (;!). 46.
492); also the acid sp. gr., 1.20. Bailey and Fowler (J. C., 188S, 53, 759) say that
dry hydrochloric acid gas in presence of oxygen and mercury, at ordinary tem-
perature for three weeks, forms Hg,OCl2 without evolution of hydrogen:
-'Eg + 2HC1 + 0, = Hg,OCl,,H,O . Hydrobromic and hydriodic acids, gasi-s.
both attack mercury, evolve H , and form respectively HgBr and Hgl (Ber-
thelot, /. c.). Hydrosulphuric acid, dry gas, at 100° does not attack dry Hg
(Berthelot, I.e.). 1I\ drosulphuric acid, in solution, and alkali sulphides form
HgS . Chlorine, bromine and iodine, dry or moist, attack the metal: mercurous
salts are formed if the mercury be in excess, mercuric salts if the halogens be
in excess.
1>. — O.ridcx. — Mercurous oxide is insoluble in water or alkalis. Hydrochloric
ncid forms HgCl : sulphuric acid forms Hg,SO, , changed by boiling with
excess of acid to HgSO,: nitric acid forms HgNO3 , changed by excess of acid
to Hg(N03)2 . Mercuric oxide is soluble in acids, insoluble in alkalis, soluble
in 20.000 to 30.000 parts water (Bineau. ('. /'., 1855, 41, 509). It is decomposed
by alkali chlorides forming HgCl,* (Mialhe, A. Ch., 1842, (3). 5. 177), soluble in
NH.C1 , from which solution NH.OH precipitates NH.Cl.NHgH.Cl
NH.HgCl (I)itte. C. r.. 1891, 112, 859), soluble in KI, forming 2KI,HgI,
(Jehu, J. C., 1872, 25, <J87).
c. — Salts. — Mercury forms two well marked classes of salts— mercurous,
monovalent, and mercuric, divalent — most mercurous compounds are per-
manent in the air, but are changed by powerful oxidizing agents to
mercuric compounds. The hitter are somewhat more stable, but are
changed by many reducing agents, first to mercurous compounds and then
to metallic mercury (10). Solutions of mercury salts redden litmus.
Many of the salts of mercury are either insoluble in water, or require the
presence of free acid to keep them in solution, being decomposed by water
at a certain degree of dilution, precipitating a basic salt and leaving an
acid salt in solution. Mercurous chloride, bromide, and iodide are insolu-
ble in water: the sulphate is j-oluble in AGO parts cold and 300 parts hot
water, soluble in dilute nitric acid (Wackenroder, A., 1842, 41, 319). The
acetate has about the same solubilities as the sulphate. Mercirrous nitrate
is completely soluble in water. On standing it gradually changes to
mercuric nitrate, prevented by the presence of free mercury, but if free
mercury be present a precipitate of basic mercurous nitrate gradually
forms. Mercuric chloride is soluble in 1(J parts of cold water and 3 parts
* The Law of Mass-Action requires that where the constituents of a slightly-ionized substance
nre presentthat substance shall form at the expense of those more strongly ioiii7.ec!. Such a
elightly-ioni/.t ;V body is HgCl.,. When HgO is brought into contact with KC1 solution Hg and
C I combine to foV«i the non-dissociated HgCl,, leaving K and O, which unite with water, im-
parting to the solut'iiyi a strong alkaline reaction. KBr and KI act even more strongly. HgO,
nlthouKh from the ready decomposition of its salts by water and from its easy reducibility a
weak base, yet will replai.'- the alkali metals where a little-dissociated Hg compound results.
An excess of Hg(NCVa dissolves chloride, bromide, and iodide of Hg and Ag owing to the
Bame cause, the Hg" ions of the strongly dissociated nitrate decreasing the already slight
dissociation of the mercuric haloids <S«). The failure of HgCl, to give many of the pre-
cipitation-reactions obtainable with other soluble mercuric salts is of course due to the same
Tact— the slight concentration of Hg" ions ($ 45).
£58, Get. MERCURY. 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, .1. Ch., 1884 (G), 3, 429),
soluble in Na2S.,0., (Eder and Ulen, J C., 1882, 42, 806), and in many
alkali salts, forming double salts. Normal mercuric sulphate is decom-
posed by water into a soluble acid sulphate and the basic sulphate, HgS04 ,
2HgO , which is practically insoluble (soluble in 43,478 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, 3HgO.N"2On , changed by repeated washing
into the oxide, HgO (Millon, A. Ch., 184G (3), 18, 3G1). 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. Reactions, a.- — Fixed alkali hydroxides precipitate, from solutions of
mercurous salts, mercurous oxide, Hg20 , 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-ydlow 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 (§45). Ammonium hydroxide and car-
bonate precipitate from solutions of mercurous salts mixtures of mercurv
and mercuric ammonium compounds. The same is true of the action of
ammonium hydroxide on insoluble mercurous salts: 2HgCl -j- 2NH4OH
= Hg -f NHoHgCl -f 2H20 -f NH4C1 ; GHgNO., + 6NH4OH = 3Hg +
(NH.HgNO^oHgO + 4NH4NO , + 5H20 ; 4Hg2S04 -f SNH4OH = 4Hg -f
(HgH2N),S04.2HgO -f 3(NH,)2S04 + GILO ; or uniting the salt in dif-
ferent manner, 4HgCl -f 4NH4OH = 2Hg + Hg,NCl.NH4Cl -f 2NH4C1
-f- 4H20 . Examination with a microscope reveals the presence of Hg° .
The mercuric ammonium precipitate dissolves in a saturated solution of
(NH4)2S04 containing ammonium hydroxide and can thus be separated
from the Hg (Francois, J. Pharm., 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, (NH2Hg)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 (&) is formed. Treat-
40 MERCURY. §58, Gft.
ing with fixed alkali hy'droxide until no more ammonia is evolved changes
the former compound to the latter (Pesci, /. 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-
monium nitrate and carbonate (Johnson, C. N.\ 1889, 59, 234). A soluble
Combination of ammonium chloride with mercuric chloride, 2NH4C1.
HgCl2 , or ammonium mercuric chloride, called " sal alembroth," is not
precipitated by ammonium hydroxide, but potassium hydroxide precipi-
tates therefrom the white mercurammonium chloride, (NHa).,HgCl, (c) :
(a) HgCl2 + 2NH4OH = NH.HgCl + NH.C1 + 2H20
(6) 2HgCl2 + 4NH.OH — NHg.Cl + 3NH4C1 + 4H,0
(c) 2NH,Cl.HgCl, + 2KOH = (NH3)sHgCL + 2KC1 + 2H,0
A solution of HgCl.. in KI with an excess of KOH (Xessler's Reagent) is
precipitated by NH4OH (or by ammonium salts), as NHg.,1 (§207, 6k).
Fixed alkali carbonates precipitate from mercurons salts an unstable mcr-
i-uroiix farliuiuitc. Hg.CO., , gray, blackening to basic carbonate and oxide when
heated. Carbonates of barium, strontium, calcium and magnesium precipitate
niercurous carbonate in the cold. -Mercuric salts are precipitated as rrd-liroini
Itaslc 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, HgCL.(HgO), , s , or 4 ; with mercuric nitrate, a basic carbonate,
(HgO)3HgCO:, . Harium carbonate precipitates a basic salt in the cold, from
the nitrate, but not from the chloride.
h.— Oxalic acid and soluble oxalates precipitate from solutions of mercnrous
salts niriTiironx o-nilati'. Hg,C...O4 . white, slightly soluble in nitric acid; from
solutions of mercuric salts, except HgCl, , mercuric uj-uliitc, HgC O, , white,
easily soluble in hydrochloric acid, difficultly soluble in nitric acid. A solution
of HgCl boiled in the sunlight with (NH,)',C,O, gives HgCl and CO, .
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 incrciirottiini a white gela-
tinous precipitate, soon turning bluish green: with incrciiricum a white precipi-
tate, becoming blue on standing. Soluble ferricyanides form with mercurons
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 (Claus, ./. pr.,
1838, 15, 40(>); from solutions of mercuric salts, mercuric thiocyanate,
Hg(CNS), . soluble in hot water (Philipp, Z. CJi., 1867, 553).
f. — 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 Hg-S , which is insoluble in
the hot acid (sp. yr. 1.42) (Howe, Am., 1887, 8, 75). HgCl is slowly dissolved by
nitric acid on boiling. All mercurous salts are oxidi/,ed to mercuric salts by
•excels*; of nitric acid.
(I. — Hypophosphorous acid reduces mercuric salts to Hg°, but the presence of
hydrogen peroxide causes the formation of HgCl from HgCL and is of value
as a quantitative method for estimation of mercury (Vanino and Treubert, B.,
1S<»7. 30, 1!M)'.)).
Phosphoric acid and alkali phosphates precipitate, from mercurous salts.
mfrcimmx plinxphatr. Hg;,PO, , white, if the reagent be in excess; but if HgNO,
be in excess. Hg^PO^HgNO;, , white, with a yellowish tinge. Mercurous phos-
phate is .soluble in dilute HNO, , insoluble in H3P04 . From mercuric nitrate-.
§58, Ge. MERCURY. 41
mercuric phosphate, Hg3(P04), , white, is precipitated, soluble in HN03 , HC1 ,
and ammonium salts, insoluble in H3P04 . Phosphoric acid does not precipitate
HgCL , and NaL,HP04 does not precipitate the white Hg3(P04), from HgCl, ,
but on standing a portion of the mercury separates as a dark brown pre-
cipitate (Haack, J. C., 1591, 60, 400; 1892, 62, 530).
e. — Hydrosulphuric acid and soluble sulphides, precipitate from mer-
curous salts, mercuric sulphide, HgS , black, and mercury, gray. Mercurous
sulphide, Hg-LS , does not exist at ordinary temperatures. According to
Antony and Sestini (Gazzetta, 1891, 24, i, 193), it is formed at — 10° by
the action of HJ3 on HgCl , decomposing at 0° into HgS and Hg . From
mercuric salts there is formed, first, a white precipitate, soluble in acid's
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
sulphide, as HgCl.,. 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 HNO;! (distinction from all other
metallic sulphides); insoluble in HC1 (Field, J. C., 1860, 12, 158); soluble in.
chlorine (nitro-hydrochloric acid); insoluble in (NH4)2S except when KOH
or NaOH be present (Volhard, A., 1891, 255, 252); soluble in K2S (Ditte,
C. r., 1884, 98, 1271), more readily if KOH be present (separation from
Pb, Ag, Bi, and Cu) (Polstorff and Billow, Arcli. Pliarm.., 1891, 229, 292).
It is soluble in K,CS.S (one part S , two parts CS2 , and 23 parts KOH , sp.
gr. 1.13) (separation from Pb , Cu , and Bi); reprecipitated as sulphide by
HC1 (Tlosenhladt, 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,
J. C., 187fi, 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 boiling with hydrochloric acid. Sodium thiosulphate added to
mercuric chloride forms a white precipitate, which blackens on standing or on
adding excess of thiosulphate, but if excess of thiosulphate be rapidly added to
HgCL 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, J. C., 1886, 49, of>7). Mercuric nitrate
\vith sulphurous acid forms slowly a flocculent white precipitate soluble in
nitric acid. The precipitate and solution contain mercurosum as evidenced by
HC1 . Mercuric nitrate with soluble sulphites forms a voluminous white pre-
cipitate, soluble in HNO;, and containing merciirosum. Mercuric chloride 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 MERCI-RY. §58, 6/.
Sulphuric acid and soluble sulphates 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 H,S04 on HgCl., see next paragraph and
(§269, 8, footnote).
/. — Hydrochloric acid and soluble chlorides precipitate from solutions of
mercurous salts, mercurous chloride, HgCl , " Calomel," white, insoluble in'
water, slowly soluble in hot concentrated HC1 . Boiling nitric acid slowly
dissolves it, forming Hg(NO:!), and HgCl2 ; dissolved by chlorine or nitro-
hydrochloric acid to HgCL ; soluble in Hg(NO:i)., (5& 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 mcr-
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(NO;!), gives a precipitate of HgCL on addition of HC1 , soluble on
addition of water. Mercuric chloride is not decomposed by sulphuric
acid. A compound HgCL,.H.,S04 is formed which sublimes undecom-
posed. The same compound is formed when HgSO, is treated with HC1
and distilled (Ditto, .1. f//., 1879, (5), 17, 120).
Hydrobromic acid and soluble bromides precipitate, from solutions of
mercurous salts, merrurous bromide, HgBr . yellowish white, insoluble in
water, alcohol, and dilute nitric acid; from concentrated solutions of
mercuric salts, mercuric bromide, HgBr., , white, decomposed by concen-
trated nitric acid. Mercuric bromide is soluble in excess of mercuric salts
(5b 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 HgCl, . Excess of concentrated H.,S04 gives some Br with HgBr., .
Hydriodic acid and soluble iodides precipitate from solutions of mer-
curous salts, mercurous iodide, Hgl , greenish yellow — "the green iodide
of mercury " — nearl}' 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 HgI2 , the latter being
dissolved as a double salt with the soluble iodide: 2HgI -f 2KI — Hg -f
HgI2.2KI . Mercurous chloride is transposed by HI or KI to form Hgl ,
excess of the reagent reacts according to the above equation (D., 2, 2, 867).
Ammonium hydroxide in the cold decompo?>es Hgl into Hg and Hgl,
(Francois, J. Pltarm., 1897, (6), 5, 388).
Mercuric salts are precipitated as mercuric iodide, HgI2 , first reddish-
§58, 7. UKRCURY. 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 oi' the iodides of all the more positive metals,
i. e. in excess of its precipitant, by formation of soluble double iodides; as
(KI)2HgI_, variable to KIHgl., . A hot concentrated solution of potas-
sium iodide dissolves 3HgL, for every 2KI . The first crystals from this
solution arc KIHgl . 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, K2HgI4 —
(KI)2HgI» (sometimes designated the iodo-hydrargyrate of potassium), is
precipitated by ammonium hydroxide as mercurammonium iodide, NHg2I
(Xessler's test), and by the alkaloids (Mayer's reagent).
Potassium bromate precipitates, from solutions of mercurous nitrate, mer-
furous bromate, HgBr03 , white, soluble in excess of mercurous nitrate and
in nitric acid; from solutions of mercuric nitrate, mercuric bromate, Hgp(BrO3),,
soluble in nitric acid, hydrochloric acid, and in excess of mercuric nitrate,
soluble in 650 parts of cold and 04 parts of hot water (Rammelsberg, Pogg., 1842,
55, 79). No precipitate is formed when potassium bromate is added to mercuric
chloride (oft, footnote). lodic acid and soluble iodates precipitate solutions
of niercurous salts as merciironx itxlatc, HgIO3 , w7hite with yellowish tint, solu-
ble with difficulty in dilute nitric acid, readily soluble in HC1 by oxidation to
mercuric salt. Mercuric nitrate is precipitated as mercuric iodctte, Hg(I03)2 ,
white, soluble in HC1 , insoluble in HN03 and H,S04 . soluble in NH4C1 , trans-
posed and then dissolved by KI . Mercuric chloride is not precipitated by
KIO3 (5fr, footnote) (Cameron, ('. .V., 1876, 33, 253).
fj. — Arsenous acid or arsenitcs form a white precipitate with mercurous
nitrate, soluble, in HN03 (Simon, Pogg., 1837, 40, 442). Mercuric nitrate is
precipitated by a solution of arsenous acid; the precipitate is soluble in HNO3
(D., 2, 2, 020). Arsenic acid or Na.HAsO, precipitates from mercurous nitrate
;;Hg3As04.HgNO:,.II,0 . light yellow if the HgNOs be in excess (D., 2, 2, 921);
dark red HgjAsO, if the arsenate be in excess. Hg.,AsO4 is changed by cold
HC1 to HgCi and H:,AsO4 , by boiling with HC1 to Hg« . HgCL , and H3AsO4 :
and is soluble unchanged in cold HN03 , insoluble in water and acetic acid
(Simon, Pof/y., 1837, 41, 424). Arsenic acid and soluble arsenates precipitate
from mercuric nitrate, Hg.JAsO,)., , white, soluble in HNO3 and HC1 , 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).
7(. — Soluble chromates precipitate from mercurous solutions mrrnirous
<-JtroiH(itc, Hg2Cr04 . brick-red, insoluble in water, readily transposed by HC1 to
HgCl and H2Cr04 , 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 turning dark brown, easily soluble in dilute
acids and in HgCL. Mercuric chloride forms a precipitate with normal chro-
mates, but not with K;.Cr.O7 .
T. 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) undecomposed — 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.
H. Detection. — Mercury in the mercurous condition belongs to the FIRST
CROUP (silver group), and is completely precipitated by HC1 . 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 (NH4).,SX and in dilute
HNO;t . The sulphide is dissolved in nitrohydrochloric acid, and the pres-
ence of mercury confirmed by the precipitation of Hg° on a copper wire, or
by the reduction to HgCl or Hg° by SnCl, .
0. Estimation. — (fl) As metallic mercury. The mercury is reduced by means
of CaO iu a combustion-tube at a red heat in a current of CO... . 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 by SnCi (or H3PO:) at 100°) and dried as
above. (It) As mercurous chloride. It is first reduced to Hg' by H3P03 (l/slar.
Z., 1895, 34, 3JM), which must not be heated above »:0°, otherwise metallic mer-
cury will be formed: and after precipitation by HC1 and drying <m a weiylu-d
filter at 100°, it is weighed as HgCl . Or enough EC1 is added to combine with
the mercury, then the Hg" is reduced to Hg7 l-v FeSO, in presence of NaOH :
2HgO + 2FeO + :!H,O = Hg.O + 2Fe(OH)., . H,SO, is added, which causes the
formation of HgCl , which is dried on a weighed filter at 100°. (<•) As HgS .
It is precipitated by H...S, and weighed in same manner as the chloride. Any
free sulphur mixed with the precipitate shoidd be removed by OS. . (rf) As
HgO , by heating the nitrate in a bulb-tube in a current of dry air not hot
enough to decompose the HgO. (<•) Voluiuetrically, by Na,S,.6:,; from tin-
nitrate the precipitate is yelloic, from the clilnri<l<' it is white'.
r5Hg(NOa)« + 2Na,S,03 + 2H2O = Hg.S,(NO3), + x'Na.SO, + 4HNO,
3HgCl2 + SNa.S.O, + 2H.O = Hg,S,CL + 2Na2S04 + ^HCl .
(/) Volumetric-ally, HgCL is reduced to Hg O by FeSO, in presence of KOH .
and after acidulating with H,SO, the excess of FeSO4 is determined by K.Cr O-
or KMnO, (Jiiptner, C. ('., 1882, 727). (g) By iodine. It is converted into HgCl
and then dissolved in a graduated solution of I dissolved in KI : 2HgCl + I'KI 4-
I, = 2K,HgI, + 2KC1. The excess of iodine is determined by Na,S,O;l . < // )
The measured solution of HgCL is added to a graduated solution of KI:
;!KI + HgCl, = K,HgI, + 2KC1. The instant the amount of HgCl, six -AH
in the equation is exceeded a red precipitate of Hgl, appears, (i) Volumetric,
§59, 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, J. C., 1873, 26, 570; Tuson, J. C., 1877, 32.
679). (;') Electrolytlcattv, by obtaining the mercury as HgNO3 , Hg(NO3)2 ,
or Hg SO, 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.e.).
10. Oxidation. — Free mercury (Hg°) precipitates Ag , An, 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.,0 , when hot to HgO (Kirchmann, /. C., 1873, 26, 476). Mercury*
and mercurous salts are oxidized to mercuric salts by Br , Cl , I , HNO., ,
H,S04 (concentrated and hot), and HC10, .
Reducing agents, as Pb , Sn , Sn", Bi , Cu2, Cu', Cd, Al , Fe, Co, Zn ,
Th1, Mg, H:!P02, H3PO., and H,SO:i , 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:
2HgCl, + SnCL = SHgCl + SnCl,
2HgCl + SnCl, = 2Hg + SnCl4
or HgCl2 -f SnCl2 = Hg + SnCl4
also 2Hg(NOs)2 + SnCl, = 2HgCl + Sn(N03)4
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 ; 2HgNO , -f Cu = 2Hg -4- Cu(NO,)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 grant}/ 10.512 heated in vacuo (Dumas, C. N., 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 gold and softer than copper. Silver is hardened
by copper; United States silver coin contains 90 per cent silver and 10 per cent
copper. Tn malleability and ductility it is inferior only to gold; and as a con-
ductor of heat and electricity it exceeds all other metals.
2. Occurrence. — 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., 1865, IS, 242 ; - Hcumann, J. C., 1875, 28, 132.
46 SILVER. §59, 3.
silver glance, Ag_S , pyrargyrite, Ag . SbS. , and horn silver, AgCl ; it is fre-
quently found in paying quantities in galena, PbS , and copper pyrites, and
in many other ores.
3. Preparation. — (a) It is alloyed with lead by fusion and the lead separated
by oxidation. (6) It is amalgamated with mercury and the mercury separated
by distillation, (e) 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. — Nilrer oxide, A.g,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 oOO° into metallic silver
and oxygen. Concerning the existence of argentous oxide, Ag4O , and silver
peroxide, Ag- O , and their properties, see Muthmann (B., 1887, 20, 983); Pford-
ten (B., 1887, "20, 1458) and Bailey (C. N., 1887, 55, 263).
5. Solubilities. — fl. — Metal. — The fla-ed alkdiis 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. N., 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, Aril., 1897, 18. 587). Silver is not oxidized by
water or air at any temperature: it is attacked by phosphorus or by substances
easily liberating phosphorus: it is tarnished in contact with hydrosnlphuric
acid, soluble sulphides, and many organic compounds containing sulphur:
except that pure rfr// hydrosnlphuric acid is without action upon pure dry silver
(Cabell, C. .Y., 1884, 50. 20S). Dilute sulphuric acid slowly dissolves finely
divided silver (Lea, /. r.), 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, b. — Ojride. — Silver oxide, Ag,O , soluble in 3000 parts of water, com-
bines with nearly all acids, except CO,, , forming the 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 H20 dissolving
227.3 parts Ag'NO at 19.5°, soluble in glycerol, and sparingly soluble in
alcohol and ether. The chlorate dissolves in about ten parts cold water;
the acetate in 100 parts; the sulphate in about 200 parts cold water and
88 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, carbonate, cyanide, ferrocyanide, ferricyanide, phos-
phate, sulphide, sulphite, chloride, bromide, iodide, iodate, arscnite, arse-
nate, and chromatc are insoluble in water.
The chloride is soluble in 24-1 parts HC1 , but its solubility is very much
lessened b}' the presence of mercurous chloride (liuysseii and Varennc, 111.,
1881, 36, o). If a solution of silver nitrate be dropped into concentrated
hydrochloric acid no precipitate appears until one half per cent of the
HC1 becomes AgCl (Pierre, J. C., 1872, 25, 123). Concentrated nitric acid
upon long continued boiling scarcely attacks AgCl (Thorpe, J. C., 1872, 25,
453); sulphuric acid, sp. gr. 1.84, completely transposes even the fused
£59, 6&. SILVER. 47
chloride on long boiling (Sauer, J. C., 1874, 27, 335). Silver chloride is
a 1^<> soluble in ammonium hydroxide and carbonate; in sodium chloride
forming a double salt; in a concentrated solution of mercuric nitrate
(S68, 1: §58, 5b 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 arc 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, bromatc, iodide, iodate, cyanide, and thio-
cyanatc; in a solution of potassium cyanide (and by many other cyanides)
except the sulphide; and in alkali thiosulphates almost without exception.
6. Reactions, a. — The fixed alkali hydroxides precipitate from solu-
tions of silver salts .(in absence of citrates), silver oxide, Ag1 ,0 , 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 '011 boiling with the fixed alkalis, except the iodide,
which is not thus transposed (Togel, J. C., 1871, 24, 313).
Ammonium hydroxide, in neutral solutions of silver salts, forms the
same precipitate, Ag.,0 , very easily dissolving in excess, by formation of
ammonium silver oxide, NH4AgO : AgNO, -f- 2NH4OH = NH4AgO +
NH4NO;S -f H,0 (Prescott, /. Am. Soc., 1880, 2, 32). In 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.
ft.— Oxalic acid and soluble oxalates precipitate silver oxalatc, Ag,C2O4 , white,
slightly soluble in water, soluble with difficulty in dilute nitric or sulphuric
acids, readily soluble in ammonium hydroxide. When heated it decomposes
with detonation, forming' metallic silver.
Potassium cyanide precipitates from neutral or slightly acid solutions
silver cyanide, AgCN , white, quickly soluble in excess of the reagent as
silver potassium cyanide, AgCN.KCN . Hydrocyanic acid precipitates
solutions of silver salts but the precipitate does not dissolve in excess of
the reagent. Silver cyanide is transposed by H2S04 or HC1 and is soluble
in ammonium hydroxide and carbonate (Schneider, J. pr., 18G8, 104, 83).
The ready solubility of nearly all silver compounds in potassium cyanide
<'.V) affords a means of separating silver from many minerals.
Potassium ferrocyanide precipitates silver ferrocyanide, Ag4Fe(CN)0, yellow-
ish white, soluble with difficulty in ammonium hydroxide and carbonate;
48 siLYKit. §59, 6c.
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 xilrcr fcrrii-i/aiiirtc, Ag>3Fe(CN)l! , reddish yellow,
readily soluble in ammonium hydroxide and carbonate. Potassium thiocyanate
gives aUrir thiocyanute, AgCNS , white, soluble in ammonium hydroxide and
carbonate, insoluble in dilute acids. Concentrated sulphuric acid with 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 by hot
concentrated sulphuric acid only on long-continued boiling (5c). 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 .~>00 parts of concentrated nitric acid (Schultz,
7,. Ch., 18(19, .V51), and is precipitated from its concentrated water solutions by
the addition of concentrated nitric acid. </. — Disodium phosphate precipitates
Kilrer plioHplitilc. Ag3P04 , 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 surer pyropftO*ptoto, white, same
solubilities as the orthophosphate.
e. — Hydrosulphuric acid and soluble sulphides precipitate from neutral
acid or alkaline solution^ silver sulphide, Ag.S , 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 form silver
sulphide from solutions of silver nitrate,* c. (]., cupric sulphide gives silver
sulphide, cuprous sulphide gives silver sulphide and metallic silver, in
both cases cupric nitrate resulting (Schneider, J. C., 1875, 28, 133 and
612).
Thiosulphates precipitate silver thiosulphate. Ag S O , white, unstable.
readily soluble in excess of the precipitant, by formation of double thiosul-
phates; with excess of sodium thiosulphate NajAg.fS.O,, )3 is formed (Cohen.
./. C*., 1890, 70, ii, 1(57). Silver thiosulphate turns black on standing or heating:
Ag,S,O., + H,O = Ag,S + H.SO, . Sulphurous acid and soluble sulphites
precipitate nilriim xii/philr, Ag SO . 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 >*iili>litil<-x precipitate
silver sulphate. Ag,SO4 , white, from concentrated solutions of the nitrate or
chlorate; sparingly soluble in water, quite soluble in concentrated sulphuric
acid.
/'. — Hydrochloric acid and soluble chlorides precipitate fiJr-cr cl\lnri<i<\
AgCl, white, curdy: separated on shaking the solution: turning violet to
brown on exposure to the light; fusible without decomposition; very
easily soluble in ammonium hydroxide as ammonio xilrrr cMoride,
(NH3)3(AgCl), (Jarry, C. r., ISiir! 124, 288). If mercurous chloride bo
present with silver chloride the solubility in ammonium hydroxide is
* Ag2S is one of the least soluble of the sulphides. Sec $57, fie, footnote.
£59, 7. SILVER. 49
greatly lessened, in fact a great excess of mercurous 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(N03)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 HC1 and lead
salts only imperfectly, silver is the only metal which belongs exclusively
to the FIRST OR SILVER GROUP OF BASES (§16).
Hydrobromic acid and soluble bromides precipitate silver 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 KIAg'I , which are decomposed by dilution with much water. The precipi-
tate dissolves in 2f>,000 parts of ten per cent ammonium hydroxide; not at all in
a five per cent solution (Longi, Gazzetta, 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 fcromate to silver nitrate is soluble
in about 600 parts water and in ;?20.4 parts nitric acid (sp. (jr., 1.21) at 25°, and
readily soluble in ammonium hydroxide. Silver ioda-te formed in manner simi-
lar to the bromate is soluble in about 28,000 parts water and in 1044.3 parts
nitric acid (sp. (jr., 1.21) at 25°, and readily soluble in ammonium hydroxide
(Longi, I.e.). 0
9- — Soluble arsenites precipitate silver arscnite, Ag3As03 , yellow, very Readily
soluble in dilute acids and in ammonium hydroxide. Soluble arsenates precipi-
tate silver arsenate, Ag3AsO4 , red-brown, soluble in ammonium hydroxide,
nitric acid, arsenic acid, and almost insoluble in acetic acid.
/
A solution of alkali stannite— as K2Sn02 — 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.
h. — Chromates and dichromates, as K2CrO4 and K2Cr2OT , precipitate silver
chromate, Ag2CrO4 , dull-red, sparingly soluble in water and in dilute nitric
acid, soluble in ammonium hydroxide.
7. Ignition. — Silver nitrate melts undecojfiposed at 218°, at a red heat it is
decomposed into Ag° , O, N, and NO (Fischer, Pogg., 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, b.
8. Detection. — Silver is identified by its precipitation with hydrochloric
acid, the insolubility of the precipitate in hot water, and its solubility in
ammonium hydroxide, with reprecipitation on rendering acid with nitric
acid (§61).
y. Estimation. — (a) As metallic silver, into which it is converted by direct
ignition if it is the oxide or carbonate, or by ignition in hydrogen if the
chloride, bromide, iodide or sulphide (Vogel, J. ('., 1871, 24, 1009). (ft) It is
precipitated as AgCl , and after igniting- to incipient fusion, weighed. (<•) It is
converted into Ag,S by H,S , and weighed after drying at 100°; inadmissible
in case of an acid that might liberate free sulphur, (d) Add KCN until a
solution of KAg-(CN). is formed, precipitate with HNO; , and after drying at
100°, weigh as AgCN . (r) Volumetric-ally, by adding a graduated solution of
NaCl until a precipitate is no longer formed. This may be varied by adding
the measured silver solution to the graduated NaCl solution, containing a few
drops of K.CrO, , until the red precipitate begins to form, (f) Volumetric-ally,
add a graduated solution of ammonium thiocyanate, containing ferric sulphate,
until the red color ceases to disappear, (y) Add the measured silver solution
to a standard solution of KCN until a permanent white precipitate is formed.
10. Oxidation. — Metallic silver precipitates gold and platinum from
their solutions, reduces cupric chloride to cuprous chloride,1 mercuric
chloride to mercurous chloride, and permanganates to manganese dioxide-.
Silver is precipitated from its solutions by: Pb , PbS !, Hg , As1, AsH, .
Sb , SbH;1 , Sn , Sn", Bi , Cu , Cu". Cd . Te , I<'e , FeS1, Al , Mn , Zn , Mg ,
P4, PH, , ELPO, , H,SO:! , SiH4\ H.O/, and H (very slowly)7.
In alkaline mixture silver is also reduced by Hg', As"', Sb'", Bi'", and
Mn"., An amalgam of mercury and tin reduces insoluble compounds of
silver in the wet way, the silver amalgamates with the mercury and the
tin becomes SnIV (Laur, C. r., 1882, 95, 38).
Ferrous sulphate in the cold incompletely reduces silver salts; on boiling, the
ferric salt formed is reduced and the silver dissolved (Lea, 7. <•.). In the gradual
reduction of silver by certain organic reagents, the metal is obtained as a bright
.silver coating or mirror upon the inner surface of the test tube or other glass
vessel. Usually a slightly ammoniacal solution of silver nitrate is used and
allowed to stand some time with the reagent: such as alcoholic solution of oil
of cloves or cassia, formic acid, aldehyde, chloral, tartaric acid, etc. Gentle
warming facilitates the result. If a good mirror is desired, great care must be
taken to free the inner surface of the glass from all organic impurities by
carefiil washing with ether, chloroform, etc. In these deoxidations, generally
the nitric acid radical *>f the silver nitrate is not decomposed, but nitric acid is
left: 4AgN03 + 2ELO = 4Ag + ^HNO3 + 0, .
Light acts upon nearly all salts of silver when mixed with gelatine or other
organic substances used in preparing photographic plates, etc. It is quite
probable that the silver is reduced to metallic silver or argentous oxide, Ag-,0 ,
or both: but the action is not well understood. The nitrate in crystal or pure
water solution, the phosphate, bromide, iodide and cyanide are not decomposed
by light alone: but light greatly hastens their decomposition by organic sub-
stances, 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.
1 Lea, Am. «., 1892, 144, 444. *D., 2, 2, 759. » Skey, C. X., 1871, 23, 232. * Senderens, C. r., 1887,
1O4, 175. 8 D.. 2, 1, 456. * Riegler, J. C., 1896, 7O, ii, 471. 7 Pellet, JB., 1874, 7, 656 ; Sehwarzenbach
and Kritscbewsky, Z., 1886, 25, 374 ; Cooke, C. N.. 1888, 58, 103. * Millon, Am. S., 1863, 86, 417.
ijGO. COMPARISON OF REACTIONS OF METALS OF THE SILVER GROUP. 51
§60. Comparison of Certain Reactions of First-Group (Silver Group) Metals.
Taken in Solution of their Nitrates or Acetates.
W
X, rt ^ooCio .2 otT of i>"
? ^Jo £ ° "3 "" w >n in
» u?~ ""^IS^jiX. -> to to
to ED ~ , * S^H S C T ^ M M
eo* g CD J £s "-'T3 rtcD «,0 W W
W fc W W pq W W
w"
m
vr.
to
W
z
03 •
tol
H
to
W
X g 3«!~ s . slJf il i
0 g 3|* 5 C S S g= fc ;
^^ rt ^J j^ ^ ^( tT O ^T O CC ^^ r-1 00 ^ g|QQ 5,^3
^J 1— r^ PA ^ ffi *" ^ *y| ^-^CQ S W5 C^ 1^ **— ^ N«— ^
to to3 to.S 42 to^ fc . B«S *3 » tS^ S iS.2 to to
WWW W W W W W W
Ck
oo"
co:
to
to
— " 1* uu SO "^ """' r~*~ MH!
^ -t:o ro'" "B^^o
o ^ oj -^ .c -»• ^ _ is f-i •*
tn • -— • F— ^^" > W P^H O W
^D '*"*' C rO ^Q ^ ^ •» CC >, ^ O ^^
C M ^ ^j [^ M * .pH -+— ' ^ ^5
^ °- '"Iw" ^ ^^i" tf«^>f<
SjS _ F>-t^K ^oj^- i— i _, •^^ .oow'^^^r >»r /*r/»r
£ CjQ >»<g * C3 C &• — 'Pi''"' ^
to1;^"3 to-S ^ toccS toS'c to'-^ -5 fcil te ^ to to to to
&
.".s^^^-^iij .S iw ^".S S "So! ^
.I^^.S^^M ^"k-T ^-2^ S> 2^
^ » d ** « ^ 53 S -°o^ >s-2 .2 "tS " S S
o "i- S «*» P^^ ^jPi'-i » ••
•^— ' -f-> »— t C} f-H t ^ Q ;/J IfJ fH _, ^ J^ *• Kj ^ if^ ^^ y-j J^ ^ ^^
3 "i Cd — ^J O G p3 ^^ PQ t* O ^o^ HH P^ii-G ^o*3 CQ P eos r^ ^
X •
0)
.S ^
w 1
o S
Jg! S to
S 8 2 5 „
°«S 1 1 I J 1 E
W^O^H d'ri1*^ r
W t< Q ,p4 Q^ O O ^3
o£W s -g 5 .2 | £
M A O W H BJ O P H ^
Stannous Salts
TABLE FOR ANALYSIS OF THE SILVER OR FIRST VROil'. §61.
a
•
•^Si **^ • "*^ fi 5 rt
^ »d --^ — '^ ^^ ^^
1
^^ ^^ 77 .
£
> O t* •w £ ."tJ a;
I
c
_o
S^S^ &°5 -a
3
53
+s
505 • ^"^ ^ dT* y "^
2
ii
**
i^?> T:CS *
o
*^
• O
t: 3 ? «, <u
f
"o
^ ^
a -X ~ _0 rg ^ £
I
»*s
OTO ^
!> JFJ ^ « ^ ^7- »^* 5k
•
^
s— ^ P
'!a'*i-*i °^Q
P
* •**>
c
* *^ _C *••• ^i
•w
S '5
CO ->,
§ 5g>
-* ' f-\ M — >C A
WW°H g o
.Q ,O rZ ,D "^ O fi
C
a
Q
coo £
V
o c
"^ fe
^ I— ' ,^( ^ *?
O £
o c,2
c v. "? «H ^ « ,-;
1
\J2. "»*•
o o —
-•"3* C'O1^ t**
'a.
ss
•Z V£
> -v ce - ^ "
10
'5
s
* ~ .
tt S fc "Z ^*" 'S ^*
CO
£
r* ~* w
OOQ
^^^
0 0 JS
*• *" ^, *^ aoo ^ OO'J
o>
p £->
w "2
o .Q . | 5 K ^^
p
J3
0 §
^3 ^
^CQO S"St''o ^.S*
a
e
O ~
t
-t-i ' J
3J WMM •2-c,i 2 M r
I
1
H JE
CO t!
'? 5
^E;S3 s^ i.2 >.-
"S
•8
£ fc
s
-*" <
s
M
•o
§
T|
E^ '2^3? 3% i ^'.2 ^"3
1
0 "^
'= "S "X — "S "8 >»*»!« a 2 S
2
• O ^™ -^ .•" w ?
O c ^
I 5?
i j?S |if. 1^!^^
0)
5
o
.
~— - . a ••• "~ — >i- •••£•"£ *•" *^ ^'
g
* I a
c
^ I § ^ 1*0* -rii^l'g^
I
M * *"•
*si.
W y *2 2 ^ C O 'JZ ? TJ 'f~ 'C * ^
a
Ell
5
llf l|*< ^*J*J8?1
Si
i
i
S
o
fa fc
O *
co a
35 °
r* ~~~~
• aT
r-l »O
^> "K
W 2
j IF 551 iSisig
—
s
CO
CO
3
•
2
•3
^ 'C
< 1
, •«
3%
%$
w sS^-j ^5§t: ?f2X^-ol
= &1 S sil a, £ 8 § f -s -i
o — cj -^^^^ ^rc-»-t,?e»:
» a H
•
u
«K
•
i
S
/v <*=.
*•»
M
0 v
« 'S
* " ^l 1
a
d o
jj
w 5
2 i
5 J g £ C
t-3 ®
3 i
OC •*
c fe o e
- t- i - . -«>
t3 "o
EH "I
f*1
o » ,-•/.= § »
bo 2J > a _o ~oo
•C o
o o
. ti
5
«!£•*•" ^ «
s- S 60S
2 g
u o
cos '~
S
M • •— ** <- ~ •— oj
O
CO
•O J1
am
_5s
Q
"w »• o< •— £ ;J 4^ aT
•° ^
g
^m .s u*"o (.a
a
S
be Q)
"S
o ^-^ C — S — ^ *2
>2
u>3
.3 -3
5.
1
w 1 j = l If
o> a
H S P %3 1> ....
| s 3 if £-3
-=
a
»o
8
co
precipitati
i yellow act
*M\ *-" ^ Cw
ro
i^
d S
w flj r* *" . ^^
._: S s .r ^ ^
3
^
|
M ^
r- ft*
X
pi
£4
§63, Ge. DIRECTIONS FOR ANALYSIS WITH NOTES. 53
DIRECTIONS FOR THE ANALYSIS OF THE METALS OF THE FIRST GROUP.
§62. Manipulation. — 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(N03)2 + 2HC1 = PbCL + 2HNO., . 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. Notes. — 1. Failure to obtain a precipitate upon the addition of HC1 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 degree of dilution that the lead chloride
formed will be soluble in the dilute acid (§57, 5c).
2. The solution should not be strongly acid with nitric acid, as it forms
nitrohydrochloric acid with the hydrochloric acid, causing oxidation of the
Hg' (§58, 5c). Lead chloride is also more soluble in nitric acid than in dilute
hydrochloric acid (§57, 5o). By a study of the solubilities of the silver group
metals it will be seen that H,SO4 , HC1 , HBr or HI cannot be used in prepar-
ing a solution for analysis when these metals arc 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 HC1 , when no further precipitation should
occur (§32). If a white precipitate is formed by adding a drop of HC1 to
the filtrate it is evident that the precipitation was not complete and more
HC1 should be added and the group separation repeated (pll).
4. The presence of a slight excess of dilute acid does not aid or hinder the
precipitation of the Hg' or Ag , but as PbCl, is less soluble in dilute HC1
than in water, a moderate excess of the acid causes a more complete precipita-
tion of that metal in the first group.
5. Concentrated HC1 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 HC1 , readily soluble in water (§186, 5c).
b. An acid solution of Sb , Bi , or Sn , with some other acid than HC1 .
and saturated with water as far as possible without precipitation, on the
addition of HC1 , precipitates the oxychloride of the corresponding metal
(§76, Gf). These precipitates are readily soluble in an excess of the HC1 . It
must, however, be remembered that a trace of AgCl will also be dissolved by
an excess of HC1 (§59, 5c).
c. Solutions of metallic oxides in the alkali hydroxides are precipitated when
neutralized with acids, e. g., K.ZnO, + 2HC1 = Zn(OH), -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 HC1 (§69, 6e).
e. Soluble polysulphides and thiosulphates give a precipitate of sulnhur,
white, with HCl'(§256, 3a).
54 DIRECTIONS FOR ANALYSIS WITH NOTES. ;<63, (!>f.
f. Certain soluble double cyanides, as Ni(CN) ,.2KCN , are precipitated
as insoluble cyanides, Ni(CN)2 , on the addition of HC1 (§133, c./y).
g. Solutions of silicates (§249, 4), borates, tungstates, niolybdates: also
benzoates, salicylates, unites, and certain other organic salts, are precipitated
by acidulation with HC1 , many of the precipitates being soluble on further
addition of the acid.
h. Acidulation with HC1 may induce changes of oxidation or reduction,
which in certain mixtures may result in precipitation: for example, Cu" salts
with KCNS in ammoniacal solution (§77, (>/>); mixture of solutions of KI and
KIO:, (§280, (>. If, ?), etc.
7. If the precipitate, obtained by the addition of HC1 to the solution, is
colored or does not give further reactions which 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, §3O1).
8. Compounds of the first group metals insoluble in water or acids are trans- .M
posed to sulphides by digestion with an alkali sulphide. The lead and silver
sulphides thus formed are readily soluble in hot dilute nitric acid. The mer-
curous compounds are changed to mercuric sulphide (§58, ."></ and (>e), a second
group mercury compound insoluble in HNO;, .
9. If but one metal of the first group be present, the action of NH,OH
determines which it is; PbCL 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 tho
following reagents, H..SO, , H,S , K,Cr,07 , and KI (§57, <> e, h, and /):
PbCL + H,SO, = PbS04 (white) 4- 2HC1
PbCl, + H,S = PbS (black) + 2HC1
SFbCL + K.Cr.O, + H.O = :>PbCrO4 (yellow) + 2KC1 + 2HC1
PbCl, + 2KI = PbI2 (yellow) + 2XC1
The yellow precipitate with potassium iodide (the KI must not be used
in great excess (§57, 5r)) 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 PbI2 (charac-
teristic of lead).
§65. Notes. — 1. Lead is never completely precipitated in the first group
(§57, Gf). The presence of a moderate excess of dilute HC1 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, <>P, §95 and §98), but that
would necessitate a regrouping of the metals; as, Ba , Sr , and Ca would also
be precipitated (Zettnow. %.. 18(57, 6, 438).
3. In order to precipitate the lead chloride, not removed in the first group, in
the second group witb^H.S , the solutions must not be strongly acid, either
the excess of HC1 shq^^Lbe removed by evaporation or the solution should be
diluted (§57, Cc, and gl, J, :. and '.)).
4. If the lend chloiW^^Blot all washed out with hot water it is changed to
an insoluble basic salHM|riTe) by the NH<OH , 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.
S68, 3. DIRECTIONS FOR ANALYSIS WITH NOTES. 55
is exceedingly delicate, much more so than the formation of the white PbSO,
(§57, 5e). In extremely dilute solutions no precipitate occurs, merely a brown
coloration to the solution. The presence of free acid lessens the delicacy of
tfce test.
6. PbCrO4 is blackened by alkali sulphides and dissolved by the fixed alkalis
(important distinction from BaCrO4); the solubility in the fixed alkalis is also
an 'important distinction from bismuth chromate (§76, &h).
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.
$66. Manipulation. — The white precipitate remaining on the filter after
washing with hot water consists of HgCl and AgCl , with usually some
PbCl2 which was' not removed. To this precipitate NH4OH , 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 -j- 2NH4OH = Hg + NH.,HgCl, 4-
NH4C1 + 2H20 .
The AgCl is dissolved by'' the NH4OH : SAgCl -f 3NH4OH = 3NH, .
2AgCl -f- 3H20 , and is found in the filtrate; its presence being confirmed
by its reprecipitation on rendering the solution acid with HN03 : 3NH3 .
2AgCl -f 3HN03 = 2AgCl
.§67. Notes. — Mercury. — 1. The"black precipitate on the filter, caused by the
addition of NH4OH to the HgClfniay be examined under the microscope for
the detection of globules of Hg°, tor the precipitate may be digested with
concentrated solution of (NH4)..,S04 , which dissolves the NH2HgCl , leaving
t-he Hg° unattacked (§58, 6«).
.£-2. If the original solution contains no interfering metals, the distinctive
i<fcactions of mercurous salts with iodides, chromates and phosphates should be
oiWhined (§58, 6e, ft and d).
3. The precipitation with HC1 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 analysis, oxidizing agents are usually
employed and the mercurvis then found entirely in the second group as a
sulphide (§96 and §97). ^
5. Additional proof may be obtained by mixing a porti^Pof the black residue
with sodium carbonate, drying -find heating in a glass tube (read §58, 7, also
§97, 7).
§68. Silver. — 1. The presence of a large excess of Hg(N03)o prevents the
precipitation of AgCl from solutions of silver salts by HC1 (§59, n<-). IP this
case the metals should be precipitated by H.S and the well-washed preci'
digested with hot dilute HN03 . The silver is dissolved as AgNO;, . while the
mercury is unattacked: 6Ag,S + 16HNO:t = 12AgNO3 + 2S2 + ^NO + 8H.O .
After evaporation of the excess of HNO3 the solution may be treated with
HC1 as an original solution.
2. A small amount of AgCl with a large amount of HgCl is not dissolved by
NH4OH , but is reduced to Ag° by the Hg° formed by the addition of the
NH4OH to the HgCl (§58, fio, §59, 10 and ?60).
3. If Hg' be present and Ag- is not detected, the black precipitate on ;the-
56 ARSENIC. §68, 4.
filter should be digested for some time with (NH4)..,S , washed, and boiled with
hot dilute niti'ic acid. The Ag , if any be present, is dissolved and separated
from the HgS :
NH.HgCl + (NHJ.S + 2H20 = HgS + NH4C1 + 2NH4OH
Hg + (NH4)aSx = HgS + (NH4)2SX.,
4. If only a trace of silver be present, its detection by adding- HNO3 to the
NH4OH solution of the chloride may fail, unless the excess of the NH4OH be
tirst removed by evaporation (because of the solubility of the AgCl in the
ammonium salt, §59, ~tc).
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, (ia and 8); by warming with grape
sugar in^kkaline mixture. In all cases the well-washed grayish black metal
may be dissolved in nitric acid as AgN03 .
f>. To identify the acid of silver salts which are insoluble in HNO.,(AgCl,
AgBr . Agl). (/) Add metallic zinc and a drop of H SO ; when the silver is all
reduced test for the acid in the filtrate. (2) Fuse with Na,CO., . add water,
and test the filtrate for acids. (3) Add H,S , or an alkali sulphide, digest
warm for a few minutes, filter and test filtrate for acids. (-'/) Boil with KOH
or NaOH (free from HC1), 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, QB).
THE TIN AND COPPER GROUP (SECOND GROUP).
Arsenic, Antimony, Tin, Gold, Platinum, Molybdenum, Mercury, Lead,
Bismuth, Copper, Cadmium (Ruthenium, Rhodium, Palladium, Indium,
Osmium, Tungsten, Vanadium, Germanium, Tellurium, Selenium).
THE TIN GROUP (SECOND GROUP, DIVISION A).
Arsenic, Antimony, Tin, Gold, Platinum, Molybdenum (Indium, Tungs-
ten, Vanadium, Germanium, Selenium, Tellurium).
£69. Arsenic. As = 75.0. Valence three and five.
1. Properties. — <S'/«'n'/?<- yrnvity, pure crystalline 5.727 at 14°; amorphous 4.716
(Bettendorff, A., 1867, 144, 110). Aleltiny ixtint, at dull red heal, under pressure
in sealed tube (Landolt, J., 1859, 182); between the melting point of antimony
and silver (Mallet, C. AT., 1872, 26, 97). V<ilutHizcn in an atmosphere of coal gas
without melting at 450° (Conechy, V. 2V'., ISM), 41, ]*'.»)• I >//«//• doixiti/ (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). ^Vrsenic 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.. sec 1)., 2, 1, 1(>1. 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, C. r., 18GO, 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 burns with a bluis'i
flame and becomes the white arsenous anhydride, As.,03 . 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
T.S a base with oxvacids classes it chemically with the non-metallic element ,
(Adie, ./. C., 1889, 55, 157: Stavenhagen, /. nnync., 1893, 283). Its chief use sis ;i
metal is in mixing with lead for making shot.
$69, 5b. ARSENIC. 57
2. Occurrence. — Arsenic is very widely distributed geographically. Found
native; as As2O3; as an alloy with other metals, e. (j., FeAs, , NiAs , CoNiAs :
as realgar, As.S., ; orpiment, As2Ss ; arsenical pyrites, FeAsFeS2 ; as an arsenate
in cobalt bloom, Co3(AsO4)., ; 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.
.!. Preparation. — (1) Reduced from its oxide by ignition with carbon; 2ASOO;,
-f 3C = As4 -}- 3C02 . (2) From arsenical pyrites, FeAsFeS, , by simple igni-
tion, air being excluded; 4(FeAs.FeS.,), = 8FeS + As4 . (3) From orpiment,
As..S3 , by fusion with sodium carbonate and potassium cyanide; 2As»S3 +
6Na>CO, + GKCN = As4 + (5NasS + (iKCNO + 6CO2 .
4. Oxides. — Arsenic forms two oxides: arsenous oxide or anhydride, As,O3
(Biltz, Z. phys. Ch., 1896, 19, 385; C. C., 1896, 793), and arsenic oxide or anhydride,
As2O5 . Arsenous oxide, As,O3 (white arsenic, arsenous anhydride, arsenous ac-id,
arsenic trioride), is usually prepared by burning arsenic; it may also be prepared
by heating arsenic in sulphuric acid till SO, is evolved, or by decomposing
AsCl3 with H2O . 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,
visually 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, As205 (arsenic anhydride, arsenic oxide), is formed by heat-
ing arsenic acid, H3As64 (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 H3AsO4 . The pentoxide, As205 , forms three acids or
hydroxides: meta -arsenic acid, HAs03 = AsO2(OH); ortho-arsenic acid,
H..,AsO4 = AsO(OH)3; and pyro-arsenic acid, H4As2O7 = As,03(OH)4; each
of these forming a distinct class of arsenates with bases. Ortho-arsenic acid is
formed by adding water to arsenic anhydride, As2O5 + 3H2O = 2H3As04 ,
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 =
H4As.,07 -f- H..O . The meta acid is formed by heating the ortho or pyro acid
to 206°: H,As64 = HAsO3 + H..O (D., I. c.).
5. Solubilities. — a. — Metai. — Arsenic is insoluble in pure water. It is readily
attacked 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 As2S5 , depending upon the proportion of
sulphur present (Gelis, A. Ch., 1873, (4), 3O, 114). Chlorine and bromine in
presence of water oxidize it, first to arsenous then to arsenic acid (Millon,
A. Ch,, 1842, (3), 6, 101): As4 + 10CL + 16H20 = 4H3AsO4 + 20KC1 . It is not
attacked by concentrated hydrochloric acid at ordinary temperature and but
slowly by the hot acid in presence of air forming As2O3 , then AsCl3 ; nitric
acid readily oxidizes it first to AsoO3 then to H3As04 ; upon fusion with KNO3
it becomes K,AsO,; readily soluble as H3AsO4 by nitrohydrochloric acid;
sulphuric acid, dilute and cold, is without action; with heat and the more COTI-
centratecl acid As20;! is formed and the sulphuric acid is reduced to S02 .
Ammonium hydroxide is without action (Guenez, C. r., 1892, 114, 1186). Hot
sohition of potassium or sodium hydroxide dissolves it as arsenite: As4 +
4KOH + 4H,O = iKAsO, + GH; .
b. — O.rides. — Arftcnatis o.rldc exists in two forms, crystalline and amorphous, the
solubilities of which differ considerably (§27). At ordinary temperature 100
parts of v.-ater 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. J. pr., 1885, (2). 31. 247). The presence of acids
greatly increases the solubility in water (Schultz-Sellac, B., 1871, 4, 109).
Arsenous oxide is readily soluble in alkali hydroxides or carbonates to arsenites
58 ARSENIC. §69, 5r.
(Clayton, C. N., 1891, 64, 27). Arsenic pcrvtoride, As.,0. , is deliquescent, soluble
in water forming- H AsO, . The meta and pvro acids are easily soluble in
water forming- the ortho acid (Kopp, A. Ch., 1856, (3), 48, 106).
f- — 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.
. l/-.vr////r* of the alkalis are soluble in water, all others are insoluble or only
partially so; all are easily soluble in acids. Alkali nrxeiiutcx, and acid arsenates
of the alkaline earths, are. soluble in wrater; all are soluble in mineral acids,
including- H3As04 (LeFevre, C. r., 1889, 108, 105S). See also under the respec-
tive metals.
Arsenons sulphide, A82S., , 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 As2S:, in water. Boiling water slowly decomposes the
sulphide forming As20, and H2S (Field, C. N., 1861, 3, 115; Wand, Arch.
Phar., 1873, 203, 290). It is completely decomposed by gaseous HC1 form-
ing AsCl3 (Piloty and Stock, B., 1897, 30, 1049), very slightly decomposed
by hot concentrated acid (Field, /. c.). Chlorine water and nitric acid
decompose it readily with formation of H.,As04; with sulphuric acid
As,0, and SO, are formed (Rose, Pogg., 1837, 42, 536). The alkali hy-
droxides or carbonates dissolve it readily with formation of RAsO . and
RAsS, (R = K, Na and NH4) (/>., 2, 1, "l83); soluble in alkali sulphides
and poly-sulphides forming R4As,S. , and RAsS, (Berzelius, Pogg., 1826,
7, 137; NiLsson, J. f., 1872, 25, 599).
Arsenic sulphide, As.,S- , is insoluble in water; soluble in HC1 gas, as
AsCl, : insoluble in dilute HC1 , soluble in HNO, or chlorine water, as
H.AsO, : soluble in alkali hydroxides and carbonates, as R.AsS, and
R,AsO,S : As,S, + 01HI4OH = (NH4),AsS4 + (NH4):!AsO,S -f 3H,0 (Mc-
Cav. Ch. %., 1891, 15, 476); soluble in alkali sulphides, as R,AsS4 (Xilsson.
J. pr., 1876 (2), 14, 171).
Arsenons chloride, bromide and iodide (AsCl. , AsBr , Asl. ) are decomposed
Ijy small amounts of water into the corresponding1 oxyhalogen compounds,
AsOCl , etc. A further addition of water decomposes these compounds into
arsenous oxide and the hn log-en acids.
G. 'Reactions. — «. — The alkali hydroxides and carbonates unite with «r.vr//oH.s-
and tirm-nir ojridex (acids), the latter with evolution of carbon dioxide, forming-
soluble alkali arxeniieg and nrxenates. These alkali salts are chiefly meta arse-
nites and ortho arsenates (Bloxam, J. C., 1802, 15, 2si: (Iraham, I'oyg., 1834, 32,
47).
•Colloids is a iiainc Driven by (iniham to a class of glue-like bodies in distinction to the crystal-
loids, which have a well-defined solid form. The colloids arc indefinitely soluble in water,
giving the little-understood "pseudo-solutions," which stand midway between the mechanical
suspension or emuls'on and the fruo 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 are as a rule broken up by addition of an acid or a neutral
salt.
;<69, Ge. ARSENIC. 59
'>. — Oxalic acid does not reduce arsenic acid* (Naylor and Braithwaite, Pharm.
J. Trans., 1883, (3), 13, 464). Potassium ferricyanide in alkaline solution oxi-
dizes arsenous compounds to arsenic compounds, very rapidly when gently
warmed, c. Nitric acid readily oxidizes all other compounds of arsenic to
arsenic acid. (I. Hypophosphites in presence of concentrated hydrochloric acid
reduces all oxycompounds of arsenic to the metallic state. 0.00001 gram ol
arsenic may be detected by boiling with 10 cc. strong hydrochloric acid and 0.2
gram calcium hypophosphite (Engel and Bernard, C. r., 1896, 122, 390; Thiele
and Loof, C. C., 1890, 1, 877 and 1078; and Hager, J. C., 1874, 27, 868).
<>. — Hydrosulphuric add precipitates the lemon-yellow arsenous sulphide,
As.S^ , 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 sulphur.
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 (or): As2S, -f 3H20 = As20, + 3H2S . This
lias been given as a method of separating arsenous sulphide from all other
heavy metal sulphides (Clermont and Frommel, /. C., 1879, 36, 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) :
As203 + 3(NH4)»S + ::H,0 = As,S, + ONH4OH
As2S3 + 2(NH4)2S = (NH4)4As»S3 or As,S;i + (NH4)2S = 2NH4AsS,
Arsenous sulphide is also soluble in alkali hydroxides and carbonates,
forming arsenites and thioarscnites (5c). The thioarsenites are precipi-
tated by acids forming As,S, : (NH4)4As,S, + 4HC1 = As2S:! + 2H2S +
4NH4C1 or 2NH4AsS, + 2HC1 = As2S:! + H2S + 2NH4C1 .
The solubility 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 (Hager, J. C.,
1SS."), 48, 838). Arsenous sulphide is soluble in solutions of alkali sul-
phites containing free sulphurous acid (separation from antimony and
tin): 4As,S, + 32KHS08 = 8KAsO, + 12KJ320:! + 3S2 + US02 -f 16H20.
It may also be separated from antimony and tin by boiling with strong
hydrochloric acid, the As._,S:i remaining practically insoluble; the sulphides
of antimony and tin being dissolved. It is easily dissolved by strong
*Patrouillard (PJiann. .7. Troiix., IHSj, (3), 13,362) claims the reduction of As" to Aa'"byoxalic
acid ; and Hager ((.'. C., 18S3, 690) reports a microscopic test for arsenic by reduction to metallic
arsenic on boiling with oxalic and sulphuric acids. Experiments in the authors' laboratory fail
to confirm these results.
60 ARSENIC. §69, 6f.
nitric acid, and by free chlorine or nitrohydrochloric acid, as arsenic acid :
6As2S3 -f 20HN03 -f 8H20 = 12H,As04 + 9S2 + 20NO ; 2As2S3 -f 10C1,
+ 16H20 = 4H,As04 + 3S, -f 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: As2S3 + 14C12 + 20H,0 = 2H,As04 -j-
3H,S04 4- 28HC1 .
Arsenic pentasulphide, AsL.S, , is formed by passing H2S 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 H2S and placing, in
stoppered bottle, in boiling water for one hour; or by passing a rapid
stream of H,S into an HC1 solution of H .AsO, (Bunsen, A., 1878, 192, 305 :
Brauner and" Tomicek, J. C., 1888, 53, 146); 2H,As04 + 5H,S + xHCl =
As S -4- 8H20 -f- xHCl . Carbon disulphide extracts no sulphur from the
precipitate, indicating the absence of free sulphur. The presence of
FeCl., or heating the solution does not reduce the As2S5 to As2S3 . If there
be a small amount of HC1 and the H2S be passed in slowly about 15 per
cent of As2S3 is formed: 2H,As04 -f 5H,S -4- xHCl == As,S3 + S2 +
8H20 -f xHCl . If NH4C1 be present more As,S:t is formed. According
to Thiele (C. C., 1890, 1, 877), arsenic acid cold treated with a slow stream
of H2S 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 ga*
arsenous chloride is formed (AsCl- is not known to exist). The solution*
in the alkali hydroxides, carbonates and sulphides form arsenates and
thioarsenates (or). Ammonium sulphide added to a neutral or alkaline
solution of arsenic acid forms arsenic sulphide which remains in solution
as ammonium thioarsenate (5r). The addition of acid at once forms
arsenic sulphide, not arsenous sulphide and sulphur. The reaction i*
much more rapid than with hydrosulphuric acid and is facilitated by
warming.
Arsine, AsH . does not combine with hydrosulphuric acid until heal il
to 230°, while stibine, SbH, , combines at the ordinary temperature (Brunn,
B., 1889, 22, 3202).
Acidulated solutions of arsenic boiled with thiosulphates form awn on*
sitlphide (separation from Sb and Sn) (Lesser, Z., 1888, 27, 218). Arsenic
may be removed from sulphuric acid by boiling with barium thiosulphute
and no foreign material is introduced into the acid : As.,0., -)- 3BaS20:l -
Ag,S:, + 3BaS04 ; 2H:!As04 + 5Na,S,0:i = As.,S, + 5Na2S04 f S, -f 3H.O .
(Thorn, J. C., 187G, 29, 517; Wagner, Din;/!., 1875, 218, 321).
Sulphurous acid readily reduces arsenic acid to arsenous acid: H.AsO,
H,SO, = H,AsO:i + H,S04 (Woehler, A., 1839, 30, 224).
/. — The arsenic from all arsenical compounds treated with concentrated
§69, 6t. ARSENIC. 61
hydrochloric acid and then distilled in a current of hydrochloric acid gas,
passes into the distillate as arsenous chloride, AsCl3 . 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 AsCL
passes over at 132°, condenses with HC1 and may be tested with SnCL
(g), or, after decomposition with water (5c) 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).
Hydrobromic acid in dilute solutions is without action upon the acids
of arsenic. The concentrated acid reduces arsenic acid to arsenous acid:
H,As04 + 2HBr = H,As03 -f Br, -f H20 . Hydriodic acid reduces
arsenic acid to arsenous acid with liberation of iodine. This is a method
of detecting Asv in the presence of As'". 0.0001 gram of H^AsO^ may be
detected in the presence of one gram of As203 : 2H3As04 -f- 4HI = As20.,
+ 2I2 -f 5H20 (Xaylor, J. C., 1880, 38, 421).
Chloric and bromic acids oxidize arsenous compounds to arsenic acid with
formation of the corresponding- hydracid: 3AsaO3 + 2HBrO3 + 9H2O =
C)H3AsO4 + 2HBr . lodic acid oxidizes arsenous compounds to arsenic acid
with liberation of iodine: 5As2O3 + 4HI03 + 13H2O = 10H3As04 + 21, .
g. — Stannous chloride, SnCl2 , reduces all compounds of arsenic from their
hot concentrated hydrochloric acid solutions, as tlocculent, black-brown, metal-
loidal arsenic, containing three or four per cent of tin. The arsenic, in solution
with the concentrated hydrochloric acid, acts as arsenous chloride: 4AsCl3 +
t>SnCL = As, + 6SnCl4 . 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 gram of the (oxidized) solid or solution
to be tested, add about Ingram of sodium chloride, and 2 or 3 cc. of sulphuric
acid, then about 1 gram of crystallized stannous 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, 6</), 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 hydro-
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 distingiiished from mercury
(§58, T). 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 arscnites arid sodium bicarbonate give chromium
arsenate (Tarugi, J. C., 1896, 70, ii, 340 and 390).
i. — Magnesium salts with ammoniiim chloride and ammonium hydroxide
precipitate from solutions of arsenates, magnesium ammonium arxcnatc.
MgNH,AsO4 , 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-
02 . ARSENIC. §69, Gy.
tAte forms slowly but completely. Compare with the corresponding magnefeium
ammonium phosphate (§189, Gd). Mnt/in^hnu arxcnite is insoluble in water, but
is soluble'in ammonium hydroxide and in ammonium chloride (distinction from
arsenates).
y.— Silver nitrate solution precipitates from neutral solutions of arsenites, or
ammonio-silver nitrate * precipitates from a water solution of arsenous oxide,
silver arttenite, Ag;,AsO3 , yellow, readily soluble in dilute acids or in ammonium
hydroxide (§59, (if/). Neutral solutions of arsc-nutcs are precipitated as silver
arsenatc, Ag'.AsO, , reddish brown, hiving the same solubilities as the arsenite.
k. — 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 copper antenite, CuHAsO3 (Scheele's green), soluble in ammo-
nium hydroxide and in dilute acids. Copper acetate, in boiling solution, pre-
cipitates the green copper aceto-arxe-nite (CvLOA.s.,Qa),Cu(C.:'H.3O.;):, (Schweinfurt
green), soluble in ammonium hydroxide and in acids. Both these salts are
often designated as Paris green (§77, 60). Copper sulphate with excess of free,
alkali is reduced to cuprous oxide with formation of alkali arsenate (10).
K3As03 + ^CuSO, + 4KOH =r K8AsO4 + 2K2SO4 + Cu,O + 2H2O . Solutions
of arsenates are precipitated by copper sulphate as copper arsenate. CuHAs04 ,
greenish blue, the solubilities and conditions of precipitation being the same
as for the arsenites.
7. — Ferric salts precipitate from arsenites, and freshly precipitated ferric
hydroxide (used as an antidote, Wormley, 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
reducing 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, , yellowish white, insoluble in acetic acid (compare
§126, Gd).
in. — Ammonium molybdate, (NH4),MoO, , in nitric acid solution, when slightly
\\arnird with a solution of arsenic acid or of arsenates gives a yellow precipi-
tate of CM^HKi'i"/' iirxt'iio-iiioli/lHlatc, of variable composition. No precipitate is
formed wwh As"'. 'J^ps precipitate is very similar.in appearance and proper-
ties '" ";i:i>"r^r'f-^l'r^TT — pV" m-1-Mn*-' except the latter precipitates com-
()'. Special Reactions, a. — Marsh's Test. — Arsenic, from all of its solu-
ble compounds, is reduced by the action of dilute sulphuric or hydrochloric
acid on zinc, forming at first metallic arsenic and then arsenous hydride,
AsH, , gaseous: As20, -f GZn + (>H,S04 = 2AsH, + OZnS04 + 3H,0 :
H:;As04 + 4Zn -f 4HJ304 = AsH, -f 4ZnS04 -f- 4H,0 . The arsenic is
precipitated with the other metals of the second group by hydrogen
sulphide, separated with antimony, tin (gold, platinum and molybdenum)
by yellow ammonium sulphide. This solution is precipitated by dilute
hydrochloric acid and the mixed sulphides, well iraslif-t]^ are dissolved in
hydrochloric acid using as small an amount of potassium chlorate crystals
;is possible. The solution is boiled (till it does not bleach litmus paper)
•Prepared by adding ammonium hydroxide to a solution of silver nitrate till the precipitate
at first produced is nearly all redissolved.
t If the ammonium salts are not thoroughly removed by washing there is danger of the for-
mation of the very explosive chlori le of nitrogen (§268, 1) when the precipitate is treated
with hydrochloric acid and potassium chlorate.
$69, B'a. ARSENIC. 63
to remove excess of chlorine and is then ready for the Marsh 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 chlo^|e 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 outei1 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 flask. Dilute sulphuric acid (one of acid to three of
water) should now be added, very carefully at first,f until a moderate
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 been 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.
fiAgNO:, + AsH, + 3H,0 = fiAg + H3AsO3 + 6HNO3
The hard glass tube should now be heated J to redness by a flame from
* The zinc and all the reagents should be absolutely free from arsenic. If the zinc be strictly
chemically pure it will bo but slowly attacked by the acid. It should be platinized (§219, 4rt) or
should contain traces of iron. Hote <A. fTi., 188t, (6 . 3, 141) removes arsenic from zinc by adding
anhydrous MgCl, to the molten metal. AsCl3 is evolved. The zinc purified in this way is
ivadily attacked by acids.
t The acid first added decomposes the alkali carbonate forming carbon dioxide which rapidly
displaces the air and greatly lessens the danger of explosion when the gas is ignited. If too
much acid be added befora tho carbonate is decomposed violent frothing may take place and
the liquid contents of the flask forced into the calcium chloride tube.
} Before heating the tube or igniting the gas, a towel should be wrapped around the flask to
insure safety in case of an explosion due to the imperfect removal of the air ; or the tube con- ,
necting the hard glass tubi 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 ttie size of the t ube).
64 . ARSENIC. §69, 6'6.
a Bunsen burner provided with a flame spreader. The flame should be
applied to' the tube between the calcium 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 tul^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.
b. Arsenous Hydride (arsine), AsH , 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: 2AsH;l -)- 302 =
As,0 , -f- 3H..O . 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; I AsH. -j-
30., = As4 -f- GILO . A number of spots should be obtained and all ihe
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, magnesium or iron t may be used, instead of /inc. and
hydrochloric acid instead of sulphuric acid. Arsine cannot be formed in the
presence of oxidizing1 agents as the halogens, nitric acid, chlorates, hypo-
chlorites, etc. Arsinuretted hydrogen (arsine) may also be produced from
<ir*cnoiitt compound* by nascent hydrogen generated in alkaline solution. Sodium
amalgam. J 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, C/); hence when
* Arsine is an exceedingly poisonous gas, the inhalation of the unmixed gas being quickly
fatul. Its dissemination in the air of the laboratory, even in the small portions which nre not
appreciably poisonous, should bo avoided. Furthermore, ns it is recognized or determined, in
its various analytical reactions, only by its decomposition, to permit it to escape undecomposed
isso far to fail in the object of its production. The evolved gas should be constantly run inti-
silver nitrate solution, or kept burning.
t According to Thiele (C. C., 1WO, 1, STTi arsenic may be separated from nntimony in the Marsh
test by using electrolyti^ally deposited iron instead of zinc. Stibine is not evolved. According
to Sautermeister (Analyst, 18fl], 218} ursine is not produced when hydrochloric acid acts upon
iron containing arsenic, but if several grams of zinc be added a very small amount of arsenic in
the iron may be detected.
t Sodium amalgam is conveniently prepared by adding (in small pieces at a time) one part of
sodium to eiirht parts (by weight) of dry mercury warmed on the water bath. When cold the
amalgam becomes solid and is easily broken. It should be preserved in well stoppered bottles.
£69, 6'c. ARSENIC. 65
the arsenic is present in the triad condition (Asv may be reduced to As'" by
SO,) the use of one of the above reagents 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 no^. satisfactory for quanti-
tative work (Hager, /. C., 1885, 48, 838; Johnson, C. AT.','i878, 38, 301; and Clark,
./. C., 1893, 63, 884).
If ferrous sulphide contains metallic iron and arsenic, arsine may be gen-
crated with the hydrogen sulphide. It cannot be removed by washing the
gases with hydrochloric acid (Otto, B., 1883, 16, 2947).
Arxine does not combine with hydrogen sulphide until heated to 230°, while
••t Urine, SbHs , combines at ordinary, temperature (method of separation)
(Brunn, 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. AsH3 + 3I2 = AsI3 + 3HI (Jacobson, B., 1887, 20,
1999). Arsenous hydride is decomposed by passing through a tube heated to
redness (mirror in Marsh test) 4AsH3 = As4 + 6H;, . Nitric acid oxidizes it
TO arsenic acid, 3AsH3 + 8HN03 = 3H3As04 + 8NO + 4H20; and may be used
instead of silver nitrate to effect 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 (Ag3AsO4 , reddish brown precipitate, 67).
Dissolve the residue in hydrochloric or nitrohydrochloric acid and test for
antimony with hydrogen sulphide (Ansell, J. C.,\853, 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 exhibit
the properties of elemental arsenic (5a). The reactions of these deposits
having analytical interest are such as distinguish arsenic from antimony.
ARSENIC 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 432°, 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 re-
water and giving the reactions for actions for antimonous oxide,
arsenous oxide.
66
568, G'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.*
Warmed with a drop of ammon-
ium sulphide form yellow spots,
soluble in ammonium carbonate, in-
soluble in hydrochloric acid (6e).
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 aiWnous
iodide, readily volatile when heated.
The heated mirror combines with
hydrogen sulphide forming the
orange antimonous sulphide, whu-h
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.
Be).
With a drop of hot dilute nitric
acid, turn white. The white fleck,
bv action of nitric acid treated with
silver nitrate and vapor of ammo-
niii. gives no color until warmed
witli 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 bv heat.
*Tho hypochlorite reagent, usually NnCIO, decomposes in the air and light on standing.
It should instantly and perfectly b each litmus paper (not redden it). It dissolves arsenic by
oxidation to arsenic acid. A*4 + lONaCIO + 6H7O = 1 11 \ *< >: + lONaCl.
i<69, (\'<l. ARSENIC. G7
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).
6. The crystals of arsenous oxide obtained above are dissolved in water
and treated with ammonio-silver nitrate forming the yellow silver arse-
nite (G/): or with ammonio-copper sulphate forming the green copper
arsenite (6fc) (Wormley, 259). Any other test for arsenous oxide ;may be
applied as desired.
7. Magnesia mixture (6i) is added to the solution of the mirror or spots
in nitric acid. A white crystalline precipitate of magnesium ammonium
arsenate, MgNH,As04 , is formed (Wormley, 316).
d. — Reinsch's Test. — If a solution of arsenic be boiled with hydrochloric acid
and a strip of bright copper foil, the arsenic is deposited on the copper as a
gray film. Hager (C. C., 1886, 680) recommends the use of brass foil instead of
copper foil. When a large amount of arsenic is present the coating of arsenic
separates from the copper in scales. The film does not consist of pure metallic
arsenic, but appears to be an alloy of arsenic and copper. Arsenous compounds
arc 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 present. This serves as one of the most
satisfactory methods of determining the presence or absence of arsenic in
<>8 ARSENIC. $69, G'e.
hydrochloric acid. Dilute the concentrated acid with five parts of water and
boil with a thin strip of bright 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 (O'o, 5). The crystals 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 using.
The copper should not contain arsenic, but if it does contain a small amount
no film will be deposited due to its presence iinless 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
;;nd gold are deposited upon copper when boiled with hydrochloric 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 this is amorphous and
does not show the octahedral crystals when examined under the microscope
(Wormley, 2(59 and ff.; Clark, J. C., 1893, 63, 886).
c. — 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 used for intentional poisoning, 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 sufficient
1o 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 usual 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 effect 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 arserous oxide. If particles be found they can at once In-
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 material.
Several methods hv.ve 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 odcr of chlorine, but concentration should
£69, 7. ABSEXIC. 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 be obtained even if arsenic
be absent. The precipitate shoxild be filtered, washed, and then dissolved in
dilute ammonium hydi oxide, 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.
Redissolve 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 Eeinsch test (6"d).
(3) Method 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
water, add nitric acid, and again evaporate to dryness; and repeat until the
mass is colorless. 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 (5c)) 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
matter. To a portion of this solution the Eeinsch 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
nnist 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 compound
containing arsenic with potassium carbonate and charcoal,* or with potas-
sium cyanide:
2As,03 + fiKCN = As4 + 6KCNO
2As2S3 + fiKCN — As4 + 6KCNS
2As,S3 + GNa,C03 -f fiKCN = As4 + 6Na2S + 6KCNO -f- GCO, .
4H3As04 + 5C = As, + 5C02 -f 6H2O
* A very suitable carbon for the reduction of arsenic is obtained by igniting1 an alkali tartrate
in absence of air to complete carbonization.
70 ARSENIC. §69, 8.
If this ignition be performed in a small reduction-tube * (a hard glass tube
about 7 mm. in diameter, drawn out and sealed at one end), the reduced
arsenic sublimes and condenses as a mirror in the cool part of the tube.
The test may be performed in the presence of mercury compounds, but
more conveniently after their removal; in presence of organic material, it
is altogether unreliable. If much free sulphur be present the arsenic
should be removed by oxidation to arsenic acid by nitric acid or hydro-
chloric acid and potassium chlorate, then precipitation after addition of
ammonium hydroxide by magnesium mixture and thoroughly drying before
mixing with the cyanide or other reducing agent.
8. Detection. — Arsenic is precipitated, from the solution acidulated with
hydrochloric acid, in the second group by hydrosulphuric acid as the
sulphide (6?). By its solution in (yellow) ammonium sulphide it is sepa-
rated from Hg , Pb , Bi , Cu , and Cd . By reduction to arsine in the
Marsh apparatus it is separated with antimony from the remaining second
group metals. The decomposition of the arsine and stibine with silver
nitrate precipitates the antimony, thus effecting a separation from the
arsenic, which passes into solution as arsenous acid. The excess of AgNO ..
is removed by HC1 or CaCL and the presence of arsenic confirmed by its
precipitation witb H.,8 . For other methods of detection consult the text
(6, 6' and 7). For distinction between Asv and As'" see ((> and §88, 4).
9. Estimation. — (1). As lead arsenate, Pb,(As04), . To a weighed por-
tion of the solution containing arsenic acid, a weighed amount of PbO is
added, after evaporation and ignition at a dull red heat is weighed as
Pb:i(As04)., . The weight of the added PbO is subtracted from the residue.
and the difference shows the amount of arsenic present reckoned as As.,0. .
(2). It is precipitated by MgS04 in presence of NH4OH and NH4C1 , and
after drying at 103°, weighed as MgNH4As04.H,0 : antimony is not
precipitated if a tartrate be present (Lesser, Z., 1888, 27, '218). (3). Th<*
MgNH4As04 is converted by ignition into Mg.,As.,0T , and weighed. (Jf).
The solution of arsenous acid containing HC1 is precipitated .by H.S .
* As much of the reduction-glass tubing contains arsenic (?) Fresenius (X., 2O, 531 and 22, ;('.i7i
rci « nnraeiids the following modification of the above method : A piece of reduction tubing about
JO mm. diameter and 15 cm. long is drawn out to a narrow tube at one end. The other end of the
tube is connected with a suitable apparatus for generating and drying carbon dioxide. The
wimple to be tested is thoroughly dried and mixed with the dry cyanide (or charcoal) and car-
bonate, placed in a small p ircelain combustion boat and put in the middle of the reduction
tube. The air is then driven from the tube by the dry carbon dioxide and the whole heated
pently until all moisture is expellrd. The tube is then heated to redness neiir the j>oint of con-
striction and when this is done the boat is heated, gently at first to avoid spattering of the fus-
ing mass, then to a full redness till all the arsenic has been driven out. During the whole of the
experiment a gentle stream of carbon dioxide is passed through the tube. The arsenic collects
as a mirror in the narrow part of the tube just beyond the heated portion. The small end of the
tube may now be sealed, the mirror collected by a gentle flame, driven to any desired portion of
tho tube and tested with the usual tests (6' c5>. Compounds of antimony when treated in thi*
way do not give a mirror. As small an amount as 0.00001 gram of A«<,O, will give a distinct mir-
tor by this method.
§69,10. ARSENIC. 71
The precipitate is separated from free sulphur by solution in NH4OH and
reprecipitated with HC1 . It is then dried and weighed as As2S3 . (5). By
precipitation as in (4) and removal of sulphur by washing the precipitate
with CS2 . Dry and weigh as As2S;, . (C>). ITranyl acetate, in presence of
ammonium salts, precipitates NH4TJ02As04 ; by ignition this is converted
into uranyl pyroarsenate (UO.,)2As207 , and weighed as such. (7). Small
amounts may be converted into the metallic arsenic mirror by the Marsh
apparatus and weighed or compared with standard mirrors (Grooch and
Moseley, C. N., 1894, 70, 207). (8). As'" is converted into Asv by A
graduated solution of iodine in presence of NaHC03 . The end of the
reaction is shown by the blue color imparted to starch. (.9). As'" is oxi-
dized to Asv by a graduated solution of KoCr.,0. , and the excess of
K2Cr,07 determined by a graduated solution of FeS04 . (10). As"' is con-
verted to Asv by a weighed quantity of K.,Cr.,07 with HC1 , and the excess
of chlorine is 'determined by Kl^and Na.,S203 . (11} • As'" is oxidized to
Asv by a graduated solution of KMn04 . The end of the reaction is indi-
cated by the color of the KMn04 . (12). Asv is reduced to As'" by a grad-
uated solution of HI . The action takes place in acid solutions. (13). In
neutral solution, as arsenate, add an excess of standard AgN03 , and in an
aliquot part estimate the excess of AgNO., with standard NaCl . (14). Dis-
tillation as AsCl3 (Piloty and Stock, B., 1897, 30, 1649; see also G'e 4).
(15). The arsenic compound is converted into AsH , and this passed into a
solution of standard silver nitrate, the excess of which is estimated with
standard Nad or the excess of AgNO., is removed and the arsenous acid
titrated as in methods (9) or (77). Many other methods have been
recommended.
10. Oxidation.— As-'"H, is oxidized to As'" by AgN03 , H.SO, , H2S04 ,
and HI03 ; and to Asv by KMn04 (Tivoli, Gazzetta, 1889, 19, 630), HNO, .
HNO, , Cl and Br (Parsons, C. N., 1877, 35, 235). As0 is oxidized to As'"
by H202 (Clark, J. C., 1893, 63, 886), HNO:! , H,S04 hot, Cl , HC10 , HC10, ,
Br, HBrO., , HI03 , Ag' (Senderens, C. r., 1887, 104, 175), and to Asv by
the same reagents in excess except H.,S04 and Ag', which oxidize to As"'
only. As"' is also oxidized to Asv in presence of acid by Pb02 , Crvl: by
compounds of Co, Ni , and Mn , with more than two bonds; and in
alkaline mixture by PbO, , Hg20 , HgO , CuO . K,CrOt , K,¥e(CN),; . pi .
(Mayer, J. pr., 1880 (2), 22, 103). Arsine is oxidized to metallic arsenic by
HgCL (Magencon and Bergeret, /. C., 1874, 27, 1008), and by As'", the As'"
also becoming As0 (Tivoli, C. C'., 1887, 1097). Asv and As'" are reduced to
metallic arsenic by fusion with CO , with free carbon, or with carbon com-
bined, as H2C204 /KCN , etc. (7). By SnCL (Gg) and H,PO, (6'0 in strong
HC1 solution; also with greater or less completeness by some free metals,
such as Cu , Cd , Zn , Mg, etc. Rideal (C. N., 1885, 51, 292) recommends
72 A\TIMO.\ 1 . $70, 1.
the use of the copper-iron wire couple for the detection of small quantities
of arsenic by reduction to the elemental state. 0.0000075 grams may be
detected. In solution Asv is reduced to As'" by H3PO, , H..S , H..SO, ,
Na2S20:! (6e), HC1 , HBr , HI (Gf), HCNS , etc. Asv and As'" are reduced
to As~'"H3 by nascent hydrogen generated by the action of Zn and dilute
H.,S04 , or, in general, by any metal and acid which will give a ready
generation of hydrogen, as Zn , Sn , Fe , Mg , etc., and H,S04 and HC1
(Draper, Dingl, 1872, 204, 320). As'" is reduced to As-'"H3 by nascent
hydrogen generated in alkaline solution as, Al and KOH , Zn and KOH ,
sodium amalgam, etc. (separation from antimony) (Davy, Ph. C., 1876,
17, 275; Johnson, C. N., 1878, 38, 301).
§70. Antimony (Stibium) Sb = 120.4. Valence three and five (§11).
1. Properties.— Specific gravity, 6.097 (Schroeder, «/., 1859, 12). Mel tiny point,
432° (Ledebiir, Wird. liribl., 1881, 650). Jioilitin point, between 1090° and 1450°
(Carnelley and Williams, J '. ('., 1879, 35, 566). Its molecular weight is unknown,
as its vapor density has not been taken. 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 ami
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. \., 1888, 57, 64); as stibnite, Sb S : as valentinite, Sb.O :
in very many minerals usually combined with other metals as a double sulphide
(Campbell, Phil. May., I860, (4), 20, :!04; 21, 318).
:!. Preparation.— ( a ) The sulphide is converted into the oxide by roasting in
the air, and then reduced by fusion with coal or charcoal. (It) The sulphide is
fused with charcoal and sodium carbonate: ::Sb S + 6Na2CO, + 3C = 4Sb +
« Na2S + 9CO2 . (r) It is reduced by metallic iroii: Sb2S3 -f 3Fe = 2Sb + 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; 1'feifer, A., 1881, 209, 161).
J. Oxides.— Antimony forms three oxides, Sb O ; , Sb,0, , and Sb.O-, . (a)
Antimonous oxide, Sb 0 , is formed (/) by the action of dilute nitric acid upon
Sb°; (^) by precipitating SbCl3 with Na,CO:, or NH4OH; (5) by dissolving Sb°
In concentrated H.SO, and precipitating with Na^COj: ('/) by burning antimony
nt a red heat in air or oxygen; (.'>) by heating Sb.O, or Sb,O-, to S()0° (Baiibigny,
<'. r., IS'.tT, 124, -:99, 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 (">'<). Anti-
inonons oxide sometimes acts as an acid, Sb 0 + 2NaOH — 2NaSbO, + H.O;
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 (f)., 2, 1. 198).
{&) Diantimony tetroxide. Sb,O4 , is formed by heating Sb° , Sb,Ss , Sb,O3 ,
•or Sb,,O. in the air at a dull red heat for a long time. The antimony in this
compound is probably not a tetrad, b\it a chemical union of the triad and
pentad: 2Sb,04 = 2Sb'"SbvO4 = Sb203.Sb,Or, . It is found native as antimony
ochre, (o) Antimonic oxide, Sb2Os , is formed by treating Sb° , Sb,O8 or
Sb,O, with concentrated nitric acid. When heated to 300° it loses oxygen,
iorming Sb2O4 (Geuther, J. pr., 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, 803' have prepared a number of antimonates and conclude
that the acid is always the met:i. II SbO, .
£70, 5b. ANTIMONY. 73
i. e., ortho, meta and pyro (Geuther, I. c., and Conrad, C. N., 1879, 40, 198). The
ortho acid, H3SbO4 is formed by the decomposition of the pentachloride with
water and washing- until the chloride is all removed (Conrad, 7. c., and Dau-
brawa, A., 1S77, 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;,O3 with potassium nitrate and extraction with
water gives the potassium metantimonate, KSb03 , 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: 2H3SbO4 = H4Sb207 + H,O (Conrad, 1. c.), which upon •
further heating to 200° gives the meta acid. The pyroantimonic acid forms
two series of salts, M.Sb.O, and M,H.,Sb2O7 . The 'sodium salt Na,H,Sb,O7
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, (>#). For the latter
the soluble potassium salt K2H.,S'b,OT 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 tetrapotassium pyroantimonate, K4Sb..07 , which does not
precipitate sodium. The reagent is, of course, not applicable in acid solutions.
The reaction is as follows: K,H,Sb,,O7 + 2Na,Cl = Na,H,Sb,07 + 2KC1 (§11).
The ortho acid, H3Sb04 , is sparingly soluble in water, easily soluble in KOH,
but insoluble in NaOH. The meta acid, HSbO3 , is sparingly soluble in water,
easily soluble in both the fixed alkalis; the pyro acid, H4Sb2O7 , is sparingly
(more easily than the meta) soluble in water; the normal fixed alkali salts,
R4Sb.,O7 , are soluble in water, also the acid potassium salt, K2H2Sb2O7 , but
not the corresponding sodium salt, Na2H.,Sb.,07 .
5. Solubilities. — a. — Metal. — Antimony is attacked but not dissolved by nitric
acid, forming Sb2O3 (a) or Sb.jO.-, (1>), depending upon the amount and degree
of concentration of the acid; it is slowly dissolved by hot concentrated sulphuric
acid, evolving SO2 and forming SbJ(S04)3 (c); it is insoluble in HC1 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,O3 , which is dissolved in the acid without evolution
of hydrogen (Ditte and Me'tzner, A. Ch., 1896, (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-
monous chloride, SbCl3 , is at first formed (d), but if sufficient nitric acid be
present this is rapidly changed to antimonic chloride, SbCln (c). If, however,
too much nitric acid be present, the corresponding oxides (not readily soluble
in nitric acid) are precipitated (tic). The halogens readily attack the metal
forming at first the corresponding trihalogen compounds (rf). Chlorine and
bromine (gas) unite with the production of light, and if the halogen be in
excess, the pentad chloride (f) or bromide is formed (Berthelot and Petit, A. Ch., '
1891, (6), 18, 65). The pentiodide, SbI5 , does not appear to exist (Mac Ivor,
J. C., 1876, 29, 328).
2HN03 = Sb203 + 2NO + H2O
10HNO3 = 3Sb205 + 10NO + 5H2O
( H2S04 = Sb2(SO4)3 + 3SO,, + 6H20
3C1, — 2SbCls
+ CL = SbCl5
l>. — O.ritlcs. — Antimonous oxide, Sb2O3 , is soluble in 53,000 parts of water at
15° and in 10,000 parts at 100° (Schiilze, J. Pr., 1883, (2), 27. :!20) ; insoluble in
alcohol: soluble in hydrochloric (rt), sulphuric and tartaric (ft) 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 dilute 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 AXTUIOXY. §70, .V.
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. /)/«»//.. issr>, 258, 520).
(a) Sb203 + 6HC1 — 2SbCl3 + 3H20
(6) Sb20s + H.C.H.O,, = (SbO)aC4H400 + H2O
(r) Sb203 + 2KOH = 2KSbO2 + H20
(d) Sb20s + Na.,00, = 2NaSb02 + CO,
Antimony tetroxide, Sb2O4 , is insoluble in water, slowly dissolved by hot
concentrated hydrochloric acid. Antimonic oxide, Sb,O,, , is insoluble in water;
soluble in hydrochloric and tartarie acids without reduction; hydriodic acid
dissolves it as antimonous iodide with liberation of iodine (Gf); slowly soluble
in concentrated fixed alkalis; soluble in alkaline solution of gtycerine (Kohler,
J. C., 1886, 50, 428). The hydrated oxides of antimony (acids) have essentially
the same solubilities as the oxides (4).
c. — RaltJt. — Antimonous chloride, SbCl , is very ricliqHCNCfHt, decomposed by
pure water, forming1 a basic salt; soluble in water strongly acidulated with an
inorganic acid, or tartaric, citric, or oxalic acids (fift), 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 deliqueactnt and require moderately concentrated acid to
keep them in solution. The sulphate. Sb_.(SO4)s . 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 CS2; the chloride and bromide are soluble in alcohol
without decomposition, but the iodide is partially decomposed by alcohol or
ether (Mac Ivor, J. C., 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 HC1 have been added first, any desired amount of
water (if added at one time) may be ndded without causing a precipitation of
the basic salt. If after acidulation water be added slowly, the basic salt will
soon be precipitated.
Antimonous sulphide, SbJS... , is readily soluble in K..S . and on evapora-
tion large yellow transparent crystals of K4Sb._,S- are obtained (a) (Ditte,
C. /'., 188(5, 102, 1(58 and 212). It is soluble in moderately concentrated
HC1 with evolution of H..S (/>) ; slowly decomposed by boiling with water
into Sb,0, and H,S (r); and on boiling with NH4C1 into SbCl, and (NH4),S
(de Clermont, C. r., 1879, 88, 972). Dilute H,S04 is almost without action,
dilute HNO, gives Sb,0.t (rf). Sparingly soluble in hot NH4OH solution.
soluble in the fixed alkalis (on fusion or boiling) (r)\ insoluble in (NH.,)2CO:.
(distinction from arsenic); insoluble in the fixed alkali carbonates in the
cold but on warming they effect complete solution (f) (distinction from
tin): very sparingly soluble in ncrmal ammonium sulphide; readily soluble
in yellow ammonium sulphide with oxidation (//} ((>f). The penlasulphide,
Sb..S- . is insoluble in water: soluble in the alkali sulphides (h), and in the
fixrd alkali carbonates and hydroxides; insoluble in ammonium carbonate
and sparingly soluble in ammonium hydroxide, more readily when warmed
(D.} 2, 1, 217). On boiling with water it slowly decomposes into Sb,0.. ,
$70, 5d. ANTIMOXY. 75
HoS and S (Mitscherlich, J. pr., 1810, 19, 455). Hydrochloric acid on
warming dissolves it as SbCl . (•»') :
(a) Sb2S3 + 2K2S = K4Sb2S5
(6) Sb,Ss + GEC1 = 2SbCl3 + 3H2S
(c) Sb2S3 + .1H,O = Sb2O3 + 3H2S
(d) 2Sb,S3 + 4HN03 = 2Sb203 + 2S, + 4NO + 2H2O
(e) 2Sb2S3 + 4KOH = riKSbS, + KSbO, + 2H2O
(f) 2Sb2S3 + 2Na,CO3 —3NaSbS, + NaSbO, + 2C02
(0) 2Sb,S3 + «(NH4)2S, = 4(NH4)3SbS4 + S2
(ft) Sb.S., + :!(NH4).S — 2(NH,)3SbS4
(t) Sb.S; + GHCl = 2SbCl3 + .°>H2S + S2
d. — 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 antimonous chloride the basic salt precipitated is
white antimonous oxychloride, Sb4Cl,0- , " Powder of Algaroth," soluble
in tartaric acid (distinction from bismuth, §76, 5r7) (Mac Ivor, C. N., 1875,
32, 229), 4SbCl, + 5HaO =? Sb4CL.03 + 10HC1 . The basic salt repeatedly
washed witli water is slowly (rapidly if alkali carbonate be used) changed
to the oxide, Sb,0, '(Malaguti, J. pr., 1835, 6, 253), Sb.,Cl20, -f H20 ' =
2Sb,0:, -f 2HC1 . With antiinonic chloride, SbCl, , the basic salt is
SbOCl, ; SbCL -f- H,0 = 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 decomposition (hydrolysis) of the salt by the
ion-- of the water, H« and OH', forming1 ag-ain the original acid an;] base. N!i2CO3, for instance,
is split up into the weak non-dissociated H2CO3 and the strongly-dissociated NaOH, whose
OH ions give the "alkaline reaction." PeCI3 in water forms soluble colloid Pe(OH)3, which
may be separated by dialysis from the free HC1 resulting or precipitated by addition of a
neutral salt-, as NaCl, to the dilute solution; KC1V gives alkaline KOII and non-dissociated
nciV, readily detected by its odor. Tn other cases precipitation is caused, as in the treatment
of bismuth or antimony solutions with water or on heating Xa-jZiiO., solution, hydrolysis in
general being increased by raising tho temrerature. The action of water on soap belongs to
this class.
§70, G«.
dissolved in excess of tartaric or citric acid solution is not precipitated on
addition of water.
6. Reactions. — a. — The alkali hydroxides and Carbonates precipitate from
acidulated solutions of inorganic antimonous salts, antimonous oxide,* Sb.Oj
(a) (Rose, Pogg., 1825, 3, 441), white, bulky, readily becoming crystalline on
boiling; sparingly soluble in water (5ft), readily soluble in excess 'of the fixed
alkalis, forming a metantimonite (ft) (Terreil, 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 antimouitr.
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 (e).
(a) 2SbCl8 + 'KOH = Sb,O3 + <>KC1 + :iH,0
2SbCl3 + :«Na2CO, = Sb,O.., + ONaCl + .tCO,
(6) Sb2O8 + 2KOH = SKSbO, + H2O
or SbCl, + 4KOH = KSbO, + :!KC1 + 2H.O
(c) Sb2O8 + Na.CO, = ^NaSbO, + CO,
(d) 2KSbO;: + 2HC1 = Sb,03 + 2KC1 + H,O
(e) 2NaSb01, + 2NaHCO, = Sb,Os + 2Na2CO3 + H,O
Antimonic saltx are precipitated tinder the same conditions as the antimonous
salts. The freshly formed precipitate is the orthoantimonic acid, H SbO, =
SbO(OH), = 8b,O(,3H,0 (</) (Conrad, C. .V., 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 (ft).
(a) SbCl., + r.KOH = SbO(OH):, + r.KCl + H.O
(ft) SbOtOH ) + KOH = KSbO, + .'H.O
ft. — 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 docs 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, soluble in hydrochloric acid (distinction from tin), or fixed
alkali hydroxides (Warren, (''. A7., 1888, 57, 124). Potassium ferricyanide is
reduced' to ferrocyanide by antimonous salts in alkaline solution (Haumann.
7.. aii'H-ir., islix;, 117).
r.— From the solutions of the fixed alkali antimonites or nntimonaten tin-
oxides or hydrated oxides (acids) are precipitated upon neutralization with
nitric acid (or other inorganic acids); the freshly formed precipitates readily
* Men«chutkin (pa#e 185) says the precipitate formed by the action of alkalis upon antimonmis
salts is the meta acid, HSbOs.
§70, 6e. .4.Y77.1/OAT. 77
dissolving in an excess of the acid. Antimonons nitrate is very 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 antimonv with the acids of phosphorus are not known,
(Na,,HP04 does not precipitate antimony salts, separation from tin, §71, 6d).
c. Hydrogen sulphide precipitates, from acid * solutions of antimo:ious
salts, antimonous sulphide (a), Sb2S3 , orange-red; in neutral solutions
(tartrates) the precipitation is incomplete. In strong fixed alkali solu-
tions (Ga) the precipitation is prevented, or rather the sulphide first
formed (?>) is at once dissolved in the excess of the fixed alkali (c), sparingly
in NH4OH . 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 co"n-
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 (?).
(«) 2SbCl3 + ,?H2S = Sb.S, + CHC1
(6) 2KSb02 + 3H2S = Sb2S3 + 2KOH + 2H2O
(c) 2Sb2S3 + 4KOH — rKSbS2 + KSbO2 + 2H,O
(d) Sb2S3 + K.S = 2KSbS2
(e) 2Sb2S3 + 6(1TH4)S0, = 4(NH4)3SbS4 + S2
(f) Sb2Ss -f m20 = Sb,03 + 3H2S
(</) 2Sb2S3 + 2K2CO3 = CKSbS2 + KSbO2 + 2C02
(70 Sb2S3 + OHC1 = 2SbCl3 + 3H2S
(/) :;KSbS2 + KSb02 + 4HC1 = 2Sb,S3 + 4KC1 + 2H2O
or ::(^H4)3SbS4 + GHC1 = Sb2S5 + GNH4C1 + 3H2S
Hydrosulphuric acid f and alkali sulphides precipitate (under like condi-
tions as for antimonous salts), from solutions of antimonic salts, antimonic
sulphide, Sb2S5 , 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. Sb2S- -)- 3K,S
= 2K3SbS4 ; 4Sb,S3 + 18KOH = 5K,SbS4 + 3KSb03 + 9H2~0 . When
dissolved in HC1 the penta-sulphide is reduced to SbCL with liberation
of sulphur, Sb2S, + GHC1 rzr 2SbCl3 + 3H2S -f S2 .
* According to Loviton (J. C., 1888, 54, 993) 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 liydrogeu 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, 6f.
All salts of antimony when warmed with sodium thiosulphate, Na.,S,O., ,
are precipitated as the sulphide (separation of arsenic and antimony). :.'SbCl
+ 3Na2S2O, + 3H2O = Sb2S3 + 3Na,SO4 + GHC1 . Sulphurous acid reduces
antimonic salts to antimonous salts (Knorre, Z. anyeic., 1888, 155). Sulphates of
antiirony are not prepared by precipitation, but bj- 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, 6. c, d.
Hydrochloric 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.
Potassium 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 antimonous snU with liberation of iodine (distinc-
tion from SnIV: SbCl, -f 2HI = SbCl, -f 2HC1 + I2 . 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., -j-
SKOH -f I, — KjSbO^ .z^Jg? -+ 3KC1 + ^H,0 (Weller, A., 1882, 213,
304). Also the absence of other oxidizing agents which liberate iodine
from hydriodic acid must be assured.
(i. — 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. insoluble (Mensehut-
kin). Stannous chloride reduces antimonic compounds to the antimonous
condition, but in no case causes a precipitation of the metal (distinction from
arsenic).
7*. — Antimonous salts in acid, neutral or alkaline solution, rapidly reduce
solutions of chroinates to chromic compounds. Acid solutions of antimonous
salts reduce solutions of manganates and permanganates to manganous salts;
with alkaline solutions to manganese dioxide. Those reactions are capable of
quantitative application in absence of other reducing agents. The antimony is
oxidized to the (iiitiniviiic condition (!) and 10).
i. — An antimonous compound when evaporated on a water bath with an
ammoniacal solution of silver nitrate gives a black precipitate (Hunsen, A.,
1^ .•>.*,, 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 precipitate of silver antimonate, soluble in ammonium
hydroxide.
§70, 6;. ANTIMONY. 79
;'. — 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; 2SbCl3 + 3Zn = 2Sb -f 3ZnCL . 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'b).
If hydrogen be generated more abundantly than in the operation above
mentioned, by zinc and dilute sulphuric or hydrochloric acid, the gaseous
antimony hydride, stibine, SbH3 , is obtained for examination. For com-
parison with arsine and details of manipulation see " Marsh's Test " under
arsenic (§69, 6'a):
Sb203 -f GZn + 6H2S04 = <>ZnS04 + 3H2O + 2SbH8
SbCl3 + 3Zn + 3HC1 = SZnCl, + SbH3
Stibine is a colorless, odorless gas, not nearly so poisonous as arsine. It
burns with a luminous and faintly bluish-green flame, dissipating vapors
of antimonous oxide and of water (n) ; 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).
(«) 2SbH3 + nO, = Sb.03 + 3H2O
(fc) 4SbHs + 302 = 4Sb + 6H2O
(c) 2SbH3 = 2Sb + 3HS
When the antimony hydride (stibine) is passed into a solution of silver
nitrate, the silver is reduced, leaving the antimony with the silver, as
antimonous argentide, SbAg^ , a black precipitate, distinction from arsenic,
which enters into solution (§69, ()'a and />); SbH, -f 3AgNO;! = SbAg:j +
3HN03 . The precipitate should be filtered and washed free from unde-
composed silver salt (and arsenous acid, if that be present), and dissolved
with dilute hydrochloric acid (HC1 does not dissolve uncombined anti-
80 ANTIMOXY. §70, 7.
mony, 5a) : SbAg3 -f 6HC1 = SbCl:i -f- 3AgCl + 3H2 . 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
silver chloride may be dissolved (§59, 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). .
i
Stibine 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 Sb,S3 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,S3 + 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, Sb.jO^.SboO.., or SbJO4
(Sb"'Sbv'O4), a compound unchanged at a dull red heat, but when heated to
«00° 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, by fusion in the crucible
with sodium carbonate and oxidizing agents, i. €., with sodium nitrate and car-
bonate, the compounds of antimony are converted into non-volatile sodium
pyroantimonate. Na.Sb O; , and arsenic compounds if present are at the same
time changed to sodium orthoarsenate. Na;,AsO4 . If now the fused mass be
digested and disintegrated in cold water and filtered, the antimonate is sepa-
rated as a residue, Na,H...Sb,O7 (4c), while the arsenate remains in solution
with the excess of alkali. The operation is much more satisfactory when the
arsenic and antimony are previously 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 may 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,
SnOj , in the residue with the Na,H2Sb,O7 . 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 sodium carbonate, more rapidly with potassium cyanide; 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 (Johnston*-.
C. N., 1883, 58, 296). The same white oxide is formed on heating antimony or
its sulphides in a glass tube, through which air is allowed to pa.^s.
8. Detection. — Antimony is precipitated, from the solution acidulated
with hydrochloric acid, in the second group by hydrosulphuric acid as the
sulphide (Ge). 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, xfilrine and arsine, into AgNO.
solution, the antimony is precipitated as SbAg., , (infimony argentide, sepa-
* Antimony as sulphide solution in potassium sulphide may be detected electrolytically, being
deposited as Sb°. Delicate to one part in 1,500,000 < Kohn, J. Soc. Intl., 1891, 1O, 327).
; 70, 10. ANTIMONY. 81
rating it from the arsenic which is oxidized and passes into solution as
arsenous acid. The SbAg3 is dissolved in HC1 and the presence of the
antimony is confirmed by the precipitation of the orange colored sulphide
with H2S . Study text at 6 and §84 to §89. For distinction between Sbv
and Sb'" see §89/7.
9. Estimation. — (/) Tartaric acid and water are added to SbCl3 , which is
then precipitated by H2S as Sb,S3 , 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 C02 . (2) Antimonous oxide, sulphide,
or any oxysalt of antimony is first boiled with fuming- nitric acid, which con-
verts it into Sb2O5 , and then by ignition it is reduced to Sb,0t , 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 KC1O3 (the tin must be present as Sniv) and
distil in a current of HC1 . 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 weigh as Na2H2Sb2O7 (7) (Tookey,
J. €., 1862, 15, 462; and Thiele, A., 1894, 263, 361). (5) For estimation of anti-
mony and separation from arsenic and tin by the use of oxalic acid, see Lessen
(Z., 1888, 27, 218) and Clarke (C. N., 1870, 21, 124). (6) Volumetrically. The
antimony compound is converted ir.to 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 (W)), and the true amount of antimony
present computed frqm the two determinations (Houzeau, J. C., 1873, 26, 407).
(7) Sb'" is oxidized to Sbv in presence of NaHCO3 by a standard solution of
iodine. The end of the reaction is shown by the blue color given to starch.
(8) Sb'" is oxidized to Sbv in presence of H2C,H4O6 by KMn04 . (.9) Sb'" is
oxidized to Sbv by K2Cr.,07 , and the excess of K2Cr2O7 used is determined by
a standard solution of FeSO4 , K3Fe(CN),., being used to show the end of the
reaction. (10} The antimony as the triad salt is treated with an excess of
standard K3Fe(CN)0; the excess of which is estimated in a gas apparatus with
HjO, (Banmann, 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 Sb,0;.
and H20 . Passed into a solution of silver nitrate, SbAg3 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 Sbv; but if the SbH3 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 Sbv, depending upon the amount of these reagents and the
temperature. Iodine oxidizes the metal to Sb'" only, except in alkaline
mixtures when Sbv is formed.
Antimonous compounds are oxidized to antimonic compounds by Cl ,
Br , HNO.J , K2Cr.,07 , and KMnO , ; by silver oxide in presence of the fixed
alkalis (6i); by gold chloride in hydrochloric acid solution, gold being
82 Try. §71, 1.
deposited as a yellow precipitate (§73, 10). The antimony is precipitated
as Sb205 unless sufficient acid be present to dissolve the oxide : 4AuCl3 -f-
3Sb,03 + 6H,0 =r 4Au + 3Sb20, + 12HC1 .
Antimonic compounds are reduced to antimonous compounds by HI (6/)
and by SnCl., ($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 antir
mony is still further reduced to stibinc. 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, 51, 292).
Sodium amalgam with dilute sulphuric acid evolves stibine from all
antimony solutions (Van Bylert, 7^., 1890, 23, 2968) but the generation
of hydrogen in alkaKne solution, i. e., Zn -j- KOH , 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. — Specific yrtirily, 7.29:i (Ilammelsberg, B., 1870, 3, 724); melting
point, 231.68° (Callendar and (Jriffiths, C. A:M 1891, 63, 2). Boils between 1450°
and 1600° (Carnelley and Williams, ./. C., 1879, 35, 566). Does not distill in u
vacuum at a red heat (Schuller, </., 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 SnO . 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. Cli., 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, — : 9°) 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. ./. />r., 1879, 2, 19, 322). Tin forms
alloys with ir.any metils. Bronze consists of copper and tin, brass frequently
contains from two to five per cent of tin, solder consists 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 Teninsula, United
States, etc. (D., 2, 1, 643). Tin pyrites, impure SnS, , is found in small quanti-
ties in various tin veins.
3. Preparation. — The reducing agent employed is carbon. The impure ore.
SnO , is first roasted, which removes some of the arsenic as As.Oa , and some
of the sulphur as SO. . 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 by treating the refined lin 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; stannons oxide, SnO . black or blue black, and stannic oxide,
SnO, , white; the latter -acts both as a base, in stannic salts, and as an anhy-
dride, in stannates. fttannoiis njridc is formed (/) by precipitating SnCl, with
K2CO3 , 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, 56. TIN. 83
(Longe, C. C'., 188G, 34); (2) by melting- a mixture of SnCl2 and NaaCO3 with
stirring until it becomes black, and removing- the NaCl by washing (Sandal,
Phil. 1/aj/., 1838, (3), 12, 216; Bottger, /!., 1839, 29, 87). Stannous hydroj-idi;
Sn(OH)2*, white to yellowish white, is formed by adding alkalis or alkali
carbonates" to stannous chloride, washing and drying at a low temperature
(Ditte, .-I. Ch., 1882, (5), 27, 1:5). (§12.)
Sldiinic oxide exists in two forms, crystalline and amorphous. The native
tinstone is nearly pure crystalline Sn02 . 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; (.{) tin filings are ignited in a retort with HgO
(I)., 2, 1, 047). Stannic liydroxide or stannic acid exists in two forms: (1) Nor-
mal stannic acid, SnO(OH)2 = H.SnO., , is formed when a solution of stannic
chloride is precipitated by barium or calcium carbonate (Freing, Pogg., 1842, 55,
519); if an alkali carbonate be used some alkali stannate is also formed. (2)
Metastannic acid, H10Snr,Oir>. , is formed by decomposition of tin with nitric
acid (Hay, C. N., 1870, 22, 298; Scott, C. 2V., 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: 5SnCl4 + 20Na2SO4 + 15H2O =
H10Sn5O15 + 20NaCl + 20NaHSO4 , or according to Fresenius (16th edition),
271: SnCl4 + 4Na2S04 + 4H2O = Sn(OH)4 + 4NaCl + 4NaHSO4 .
5. Solubilities.— a. — Metal. — Tin dissolves in hydrochloric acid slowly when the
acid is dilute and cold, but rapidly when hot and concentrated, stannous
chloride and hydrogen being produced (fl); in dilute sulphuric acid, slowly, with
separation of hydrogen (I), (not at all even in hot acid if more dihite than
H2SO4.6H:,O (Ditte, A. Ch., (5), 27, 145); in hot concentrated sulphuric acid,
rapidly, with separation of sulphurous anhydride and sulphur (c) ; nitric acid
rapidly converts it into metastannic acid, insoluble in acids (d); very dihite
nitric acid dissolves it without evolution of gas as stannous nitrate and am-
monium nitrate (e) (Maiimene, Bl.. (2), 35, 598); nitro-hydrochloric acid dis-
solves tin easily as stannic chloride (f), potassium hydroxide solution dissolves
it very slowly, and by atmospheric oxidation (#) ; or, at high temperatures,
with evolution of hydrogen (ft). Bromine vapors readily attack melted tin
with formation of SnBr4 , colorless crystals, melting point 30° (Carnelley and
O'Shea, J. C., 1878, 33, 55).
Sn + 2HC1 — SnCl2 + H4
Sn + H2S04 = SnS04 + H2
Sn + 2H2S04 = SnSO, + 2H.O + SO2
and then -!SnS04 -f 2SO, + 4H2SO4 = 4Sn(SO4), + S2 + 4H2O
(d) 15Sn + 20HNO3 + 5H20 = 3H10Sn5O13 + 20NO
(e) -iSn + 10HN03 = 4Sn(NO3)2 + 3H2O + NH4NOS
(f) Sn -f 2C12 = SnCl4
(<j) 2Sn ; 4KOH + O2 = 2K2Sn02 + 2H2O
(7i) Sn + 2KOH = K2SnO2 + H2
It. — OxidfS. — Stannous ox'idc is insoluble in water, soluble in acids (Ditte, A. C'h.,
1882, (5), 27. 145; Weber, J. C., 1882, 42, 1266), oxidized by nitric acid when
heated, forming the insoluble metastannia acid. JStannous hydroxide is readily
soluble in all the solvents of the oxide, and is also readily soluble in fixed
alkali hydroxides. Stannic o.ridc, SnO2 , is insoluble in water; soluble with
difficulty in alkalis; insoluble in acids except in concentrated H2S04 (D., 2, 1,
(•>••! s). Sulphur forms SnS., and SO2; chlorine forms SnCl4 (Weber, PofKj., 1861,
112. 619). Normal fttnnnic add, H2Sn03 , freshly precipitated, is soluble in
fixed alkali hydroxides and in acids (Ditte, C. r., 1887, 1O4, 172); insoluble in
water and changed by hot nitric acid to the insoluble metastannic acid.
Al<'l<i*tannic acid, H10Sn0O15 , is insoluble in water and acids, HC1 changes it to
* According to other authorities Sn(OH)2 dors not exist, but a hydratcd oxide is formed,
SnO.Sn(OH)2 (Graham-Otto, '4,2,1267; D., 2, 1,657; Gmelin-Kraut, 3, 107).
84 TIN. §71, 5c.
rnetastannic chloride insoluble in the acid, but soluble in water after removal
of the acid; soluble in the fixed alkalis as metastannates, which are soluble in
water and precipitated by acids. Metastannic acid in contact with HC1 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
I'hloride and bromide dissolve in pure water with little or no decomposition
(Fersonne, C. r., 18(12, 54, 21f>; and Carnelley and O'Shea, J. C., 1878, 33, 55).
Stannous chloride is soluble in less than two parts of water (Kngel, A. Ch., 1891,
((>), 17, 347); but more water decomposes it, unless a strong excess of acid be
present: 2SnCl2 + H,O = SnO.SnCL + 2HC1 . Pure stannic chloride is a
liquid; sp. <jr., 2.2; boiling point, 144°; solidifies at — :',3° (Besson, C. r., 1889, 109,
940). A small amount of water ,added to the liquid combines with heat to form
crystals of SnCl,.:.H O . which are readily soluble in excess of water (/)., 2, 1,
<>f>2). 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, Ga~zctta,
1S92, 22. ::.S4). Stannic iodide is readily soluble in water (Schneider, 1'om/., 18f>fi,
127, 024). 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 deliquescent. Stannous sulphide is insoluble in water, soluble
in HC1 with formation of H S: decomposed by HNO with oxidation to meta-
stannic acid; insoluble in solution of the normal alkali sulphides, but soluble
in the polysulphides with oxidation to a stannic compound (<>f). Stannic sul-
phide is soluble in HC1 , with evolution of H,S: and in solutions of the alkali
sulphides.
G. Reactions. — a. Alkali hydroxides and carbonates precipitate from
solutions of stannous salts, stannous hydroxide, Sn(OH)._, (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 (Schaffner, A., 1844,
51, 174).
SnCL + 2KOH = Sn(OH)2 + 2KC1
Sn(OH)2 + 2KOH = K2SnO2 + 2H2O
SnCL + 4KOH = K2SnO, + 2KC1 + 2H2O
SnCL + Na2C03 + H,O = Sn(OH), + 2NaCl + CO,
By gently heating the solution of potassium stannite, K,SnO, , 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 ; 2K»SnO., -f- H,0 - Sn -f-
K,SnO:! -|- 2KOH . The reaction proceeds more rapidly upon the addition
of a little tartaric acid. Sf/im/ir .suits are precipitated by alkali hydroxides
and carbonates as stannic aril. H.,SnO:, soluble in excess of the fixed alkali
hydroxides, insoluble in ammonium hydroxide and the alkali carbonates
(Ditte, A. Ch., 1897 (G), 30, 282).
SnCl, + 4KOH = H2Sn03 + 4KC1 + H20
H2SnOa + 2KOH = K2Sn08 + 2H.O
SnCl, + ('KOH = K2SnO3 + 4KC1 + 3H20
SnCl4 + 2Na2CO, + H2O = H.SnO, -f 4NaCl + 2CO2
.
* Stannous sulphate is decomposed by an excess of cold water forming 2SiiSO4.4SiiO.3HaO^
and by a small amount of hot water forming: SnSO4.28nO (Ditte, A. Ch., 1882, (5), 27, 161).
§71, Be. TY.V. sr>
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.
ft.- — 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 (TTausmanu
and Loewenthal, A., 1854, 89, 104).
Potassium cyanide precipitates both stannous and stannic salts, white, in-
soluble in excess of the cyanides. Potassium ferrocyanide precipitates from
stannous chloride solution stannous ferrocj/anide, Sn2Fe(CN)6 , white, insoluble
in water, soluble in hot concentrated hydrochloric acid. Stannic chloride is
precipitated as a greenish white gelatinous precipitate, soluble iii hot hydro-
chloric acid, but reprecipitated upon cooling (distinction from antimony)
(Wyrouboff, A. Oh., 1876, (5), 8, 458). Potassium ferricyanide precipitates from
solutions of stannous chloride, stannous ferricyanide, Sn.8(Fe(CN)8), , 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.
c. — 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 SnO2; 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), 58, 471).
d. — Hypophosphorous acid does not form a precipitate with stannous or
stannic chlorides, nor are these salts reduced when boiled with the acid. Sodium
hypophosphite 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., I860, 114,
113). With stannic chloride a white gelatinous precipitate is formed, soluble
in HC1 and KOH . insoluble in HN03 and HC2H3O, . If the stannic chloride be
dissolved in excess of NaOH before the addition of Na2HPO4 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 (Bornemann, Z. angew., 1899, 635).
e. Hydrosulphuric acid and soluble sulphides precipitate from solutions
of stannous salts dark brown hyd rated stannous sulphide, SnS (a), insol-
uble in dilute, soluble in moderately concentrated HC1 (&). It is readily
dissolved with oxidation by alkali supersulphides; the yellow sulphides,
forming thiostannates (r); from which acids precipitate the yellow stannic
sulphide (d). The normal, colorless alkali sulphides scarcely discolve any
t stannous sulphide at ordinary temperature, compare (§69, Qe and §70, Qe),
but hot concentrated K2S dissolves SnS forming K2SnS., and Sn (e) (Ditte,
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 brown stannous sulphide (</). Am-
monium hydroxide and the alkali carbonates do not dissolve it (distinction
from arsenic. §69, fo). The insolubility in fixed alkali carbonates is a
'86 77.V. £71, 67.
distinction from antimony (§70, Ge). Nitrohydrochloric acid (free chlorine)
dissolves it as stannic chloride, with residual sulphur (h). Nitric acid
oxidizes it to metastannic acid without solution (i) (separation from
arsenic. $69, Ge).
(a) SnCl, 4 H2S = SnS -f 2HC1
(ft) SnS + 2HC1 = SnCl, + H,S
(c) SnS 4 (NH4),S, = <NH,),SnS,
(d) (NH,),SnS, 4 2HC1 = SnS, + :.'NH,C1 4 H2S
(e) 2SnS + K,S = K,.SnS, 4 Sn
(/) 2SnS + IKOH = K,SnO, + X2SnS2 4 2H.O
(g) (K2SnO, + K,SnS,) + 4HC1 = 2SnS 4- 4KC1 + 2H.O
(ft) 2SnS 4- -ICL = 2SnCl, + S,
(0 :!OSnS + 40HNO, 4 10H,O = iiH11,Sn5O1!i 4 40NO 4 158,
Solutions of stannic salts are precipitated as stannic sulphide, SnS,, .
hydrated, yellow, having much the same solubilities as those given for
stannous sulphide, with this difference, that stannic sulphide is moderately
soluble in normal, colorless, alkali sulphides. The following equations
illustrate the most important reactions:
SnCl, 4- 2H,S = SnS, + 4HC1
SnS, + 4HC1 = SnCl, + 2H,S
SnS. 4 (NHi),S= (HH,)sSnSs
2SnS2 4 2(NH,),S, = 2(NH4)2SnSs + S2
:iSnS, 4 r.KOH = X,SnO, 4 2K2SnS3 4 3H20
(K,SnO, 4 2K,SnS:>) 4 <>HC1 = .iSnS, 4 «XC1 4 3H20
SnS2 4 2C12 = SnCl, 4 S,
l.r>SnS2 4 20HNO, 4 .*>H2O = :!H10Sn,O,, 4 15S2 4 20NO
Sodium thiosulphate does not form a precipitate with the chlorides of tin
(separation from As and Sb) (Lesser, /., 1888, 27, 218). Sulphurous acid ami
sodium sulphite precipitate from stannous chloride solution not too strongly
acid, xtdiinous KHlphiti; SnSO , white, readily soluble in HC1 . When warmed in
the presence of hydrochloric aekl, sulphur dioxide acts ;is an oxidizing agent
upon the stannous salt. A precipitate of S^O,,^ or SnS2 is formed, or H,S
is evolved and SnCl, formed, depending upon the amount of HC1 present.
dSnCl, 4 2SO2 4 «HSO = Sn,0,,,S, 4 12HC1
GSnCl, 4 2SO2 4 sHCl = SnS, 4 :>SnCl, 4 tH:O
3SnCl2 4 SO, 4 GHC1 = 3SnCl, 4 H2S 4 2H2O
Stannic chloride does not give a pn-cipitatc with sulphurous acid or sodiufn
sulphite.
'I'lif sulphates of tin are formed by dissolving the freshly precipitated
hydroxides in sulphuric acid and evaporating at a gentle heat. They cannot be
formed l>\ precipitation and are decomposed l>\ water (Ditle, .4. <'h., is.^ii, (5),
27, 145).
f. — Potassium iodide added to a concentrated water solution of stannous chlo-
ride forms tirst a yellow precipitate soluble in excess of the SnCl;, . Further
addition of KI gives a yellow precipitate rapidly turning to dark orange needle-
like crystals, often forming in rosette-like clusters. It a drop of the stanndus
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 HC1 . KOH . and C,H-OH , soluble in
large excess of KI and sparingly soluble in H,0 with some decomposition.
*§71, :. 77. \. s;
The yellow precipitate is probably a double salt of stannons iodide and potas-
Kium iodide, and has about the same solubilities as the orange precipitate
(Personne, J., 1862, 171; Boullay, .1. Cli., 1S27, (2), 34, :!72). 1 otassium iodide in
concentrated solution precipitates xtunii-ir indiili', yellow, from very concentrated
water solutions of stannic chloride. The precipitate is readily soluble in water
to a colorless solution (Schneider, ./., 1866, :>29). llydriodic acid does not give
free I with Sniv , distinction from Sbv and Asv (llarroun, /. C'., 1882, 42, 661).
The chlorates, broiiiat.es and iodates of tin have not been thoroughly studied
(Watt*, 1, 5:59, III., 22; D., 2, 1, 675). Stannous chlorate appears to be formed >
when potassium chlorate is added to. a concentrated water solution of stannons
chloride; it dissolves on addition of HC1, and nearly all dissolves in excess of
water. With KBrO3 , bromine is liberated, and with KIO:1 iodine is liberated.
Potassium chlorate, bromate and iodate all form precipitates with stannic
chloride, soluble in HC1 without liberation of the halogen.
g. — Stannous arsenate, 2SnO.As.,O5 , a voluminous flocculent precipitate is
formed by adding a solution of SnCL to a concentrated acetic acid solution of
K.AsO, , decomposed by heating to As , As.,O3 and Sn02 (Lenssen, A., 1860, 114.
115). Stannic arsenate, 28nOa.AS|Og , a white gelatinous precipitate is formed
by adding HUTtX, to a mixture of Na,SnO:, and Na.,AsO4 (Haeffely, J., 1855, :;o:>).
With antimony, tin acts as a base, forming stannous and stannic antimonites
and antimonates (Lenssen, I. c.).
h. — If 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 SnCl4 , insoluble in H,O , soluble with difficulty in HC1 . K2Cr2O7
also gives a precipitate with SnCL and SnCl, fLeykatif, J. pr., 1S-10, 19, 127).
i. An ammoniacal solution of silver nitrate is reduced to metallic silver
by a solution of potasxiwn * fa unite. 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 ,
Sn . Al . Cr , 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, HgCL , reacts with stannous
chloride solution, forming SnCl4 and a precipitate of HgCl (white) or Hg°,
gray, depending upon the relative amounts present (§58, 6y).
k. Stannous salts react with (NH4)2Mo04 , giving a blue-colored
solution of the lower oxides of molybdenum, constituting a delicate test
for Sn" (§75, Qg).
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 strongty, becomes of a blue-green color. SnOx
fused with KCN gives metallic tin (Bloxam, ,/. C., 1865, 18, 97).
88 77 .v. £71, H.
8. Detection. — Tin is precipitated, from the solution acidulated with
hydrochloric acid, in the second group hy hydrOBulphuric acid, as the sul-
phide (tie). By its solution in yellow ammonium sulphide il is separated
from the Copper Group (Hg . Pb , Bi , Cu . and Cd). By the reaction in
ihe Marsh apparatus the tin is reduced to the metal and is not dissolved
;is long as xine is still present. The residue Sn (Zn , Sb . Au , and Pt) in
the Marsh apparatus is wanned with hydrochloric acid, which dissolves
the Sn as SnCl, . This is detected by its reducing action on HgCl., , giving
a white precipitate of HgCl or a gray one of Hg° (O/).
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 AgNO., to sixteen
parts NH4OH). A brown-black precipitate of metallic silver indicates
(hat tin was present in the stannous condition ((>i). Consult also §90
and §92.
9. Estimation. — (/) (Jravimetrically. It is converted into SnO , and after
ignition weighed. (2) Volumetrically. To SnCL add KNaC4H4On and NaHCO, ,
then some starch solution and a graduated solution of iodine, until a perma-
nent blue coloration appears. (3) To SnCl, add slight excess of 3?eCl;, , and
determine the amount of FeCL formed, by a graduated solution of KMnO, .
(-J) Ry 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 , Cu , Pt ,
and Au , to the metallic state. Sn" is oxidized to SnIV by free HNO, .
HNO,1, H,Fe(CN)(, . H.SO, and H,S04 (if hot), Cl , HC10 , HC10, , HC10;l ,
Br , HBrO, , I », and HIO . Also by Pb" (in alkaline solution only). Pblv ,
Ag^1. Hg'. Hg". Asv. As'" (in presence of HC1), Sbv, MoVI. Bi'", Cu'.
PdtNO,), , PtIV •», Fe'", FeVI, CrVI, Co'", Ni'", and Mn2+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 fonns, 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-
i Kestner, A. Ch., 1860, (3), B«. 471. • Ditte, .4. Ch.. 1882. (5). 27, 145. ' Thomas, C. r., 189*i, 122,
1539. «Ditte,C.r., 1882, 04,1114.
§71, 10. TIN. 89
pounds of arsenic in strong HC1 (detection of arsenic) ; bismuth salts, to
metallic bismuth (in alkaline mixture §76, 6</); and ferric salts, to
ferrous salts, left in solution, much used in volumetric analysis of iron
(9, and §126, 6g 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
Muller, J. pr., 1884, 30, 252).
Solutions of SnIV and Sn" are reduced to the metallic state by Cd , Al .
Zn , and Mg . According to Rideal (C. N., 1885, 51, 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, 25, 274),
90
REACTIONS OF MtsH\H'. A.V77A/0AT AND TIN.
372.
I
? I"
I ^
o
en • ~
S I
&, c
<N
I> -<
•//. O
s-
&
I1§li t
•-H /-^ o 0)
O c
.t: at- ~
•f5 =
?l
— HH ^ — ^
;•*>.= •£ 5!5 o"'
g — s x 5 -2 "*?.
co M
n**- *
-> ~Z ~ - -C +-1
* = = -=• 2
C "S "rt P S • !f
O
W
c|
»* X
E - c -r
= :" » "
fl
CO
g
•
03
2§g
; . 4S -iS"*
' 's'-S ^ o
I «D >»•— X
• «=:
•- ^'J^ " ^:
JS t"
j, rrjtos ; OJ
" ?-i^x<3'
AX wO^
O 4>
=!,3
S — fl v:
CO
w
= s 0'
i. T. ^ ^
..= 0
O ^ CQ
CO
CO
«O ha'^
n
OlMgl
*ir?
^a-g si
CO ? C
W. -C 05 —
' Ml2
• H
•W* c-^COo ? «-
CO CO CO
do
«
B
«
'S
v
V
m
E
C.
C
1
55
cc
V
0) ^
*•
/•-O •
i it w C"
1
&fe
It
hJ^'S -
o
^
^§
± c
a o
111
-
d 01
S • 9
"o 0 6 "g
t- -^ 42
«w « es
P4
\0 ® CC
^1 -
5^i
>lu
p *S o **•
~ .5 -= s
B y-
: W
T3^-3 ^ K
O j3 * S,'3
8 « «B . 6?
«?.1.2.E-S
M
o
W
X 4)
: i
13
S 1!
w a\
c"^"§ g£
y *•?* i
= = 1 c
pgsg^i
.•3 £ G fc 4J K
c W
$73. 5fl. GOLD. 91
£73. Gold (Aurum) Au — 197.2. Valence one and three.
1. Properties.— Specific ffravity, 19.30 to 19.34 (Rose, Pogg., 1848, 75, 403), Melt-
ing point, 10(>1.7° (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 metals also affects 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 O.OOQ1 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,
equal to copper, not so good as silver. It has a high coefficient 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 A.g , Cu . etc.
2. Occurrence.- — Gold is usually found native, but never perfectly pure, being
always alloyed with silver, and occasionally also with other metals. It is found
as gold-dust in alluvial sand, sometimes in nuggets, and sometimes disseminated
"in veins of quartz.
3. Preparation. — (1) 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. (3) 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 lend by cupellation. The silver is separated from
the gold by dissolving it in nitric 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. (4) By treatment with a
solution of KCN . ( •>) By amalgamation with mercury and electrolysis at the
same time.
. v.4. Oxides and Hydroxides. — Aiiroiis o.ri(h', Au2O , 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 SO2 in ice-cold solution; heating to
2006 changes it to the oxide (Kriiss, .4.. 1886, 237, 274). Auric hydroxide,
Au(OH)3 , is prepared by precipitation from the chloride solution with MgO
(Kriiss, I. c.). It is a yellow to brown powder, changing to the oxide upon dry-
ing at 100°. Healing to 250° gives the metal and oxygen (^10)-
r,. Solubilities.— a. — Metal. — Gold is not at all tarmsneci or in any way acted
upon by water at any temperature, or by hydrosnlphuric 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 broni-idc, AuBr3: and by iodine: dissolving when finely divided in hydri-
odic acid by aid of the air and potassiinn iodide, as potassium auric iodide,
KIAul,: !Au + Ir.'HI + »KI + 3Oa — 4KIAuI., + 6H,0 . Potassium cyanide
solution, with aid of the air, dissolves precipitated gold as potassium anro-
wa-inilc. KAu(CN).,: ^Au + SKCN + O, + 2H,O = 4KAu(CN), + 4KOH .
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; when 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 f>0 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 GOLD. §73, olf.
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.
c. — 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
Au° and Au"' . Auric chloride is deliquescent; 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.
(i. Reactions, a. The fixed alkali hydroxides and carbonates in excess
do not precipitate AuCl . solutions, as a soluble aurate, KAu02 , readily
forms; but upon boiling and neutralizing the excess of alkali, Au(OH),
is precipitated. Ammonium hydroxide precipitates from concentrated
solutions a reddish-yellow ammonium aurate, (NH.,).,Au20;t , "fulminating
gold." I. 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 golfl separates in metallic flakes or forms
a mirror on the side of the test-tube. 2AuCl3 + 3H,C,04 = 2Au + GOO*
-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
group metals as sulphides. Potassium gold cyanide, KCN.Au(CN)3 , is
formed when a neutral solution of AuCl. is added to a hot saturated
solution of KCN. It is verv soluble in water and by heating above 200°
it is decomposed into CN and KCN.AuCN , which latter product is formed
when gold is dissolved in KCN in the presence of air (5a). c. A solution
of AuCl., is precipitated as Au° by a solution of KNO, . d. Sodium
pyrophosphate forms with AuCl , a double salt which has found application
in gold plating, e. Hydrosulphuric acid precipitates from gold chloride
solution, hot or cold, tjold sulphide, variable from Au.,8 to AuJSL , 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, i. c.
Au,(SO,),.(NH4),SO.,.(iNH., -f ;5H,0 . 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 duronx loiJi/le, 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
colored solution: AuCl3 -f- 4KI = Aul + 3KC1 -f I, with KI . But, on
gradually adding auric chloride to solution of potassium iodide, so that th«;
§74,1. /'/,.r/7.vr.i/. 9:*
latter is in excess at the point of chemical change, there is first a dark-
green solution of potassio-auric iodide, KIAuI3 ; then a dark-green precipi-
tate of auric iodide, AuI3 , very unstable, decomposed in pure water, more
quickly by boiling; changed in the air to the yellow aurous iodide.
(j. Stannous chloride gives a purple precipitate containing the oxides of
tin with the gold, " purple of Cassius " insoluble in acids.
h. Ferrous sulphate is the most common reagent for the detection of
gold, reducing all gold .salts to the metallic state; AuCl3 -4- 3FeS04 =
An + Fe2(S04)3
7. Ignition. — Gold is reduced from many of its compounds by light, and from
all of them by heat — its separation in the dry way being- readily effected by
fusion 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 off 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 H,S , passes into Division A (tin group) by (NH4)2S , and may
be detected in the flask of the Marsh apparatus by the usual methods.
9. Estimation. — Gold is; always weighed in the metallic state, to which form
it is reduced: (1) By ignition alone if it is a salt containing no fixed acid; if m
an ore, by mixing4 with 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 FeS04 , H,G.,Ot , chloral hydrate, or some
easily oxidized metal, such as Zn , Cd , or Mg . (.?) Gold is also estimated volu-
metrically by H2C,O4 and the excess of H2C2O, used, determined by KMnO4 .
10. Oxidation. — Gold is reduced to the metallic state by very many
reducing agents, among which may be mentioned the following: Pb , Ag ,
Hg, Hg', Sn, Sn", As, As"', AsH3 , Sb , Sb'", SbH, , Bi , Cu , Cu',
Pd, Pt, Te, Fe, Fe", Al , Co, Ni, Cr'", Zn, Mg, H,C204 , HNO, , P,
H3P02 , H3PO., , PH., , H2SO., , and a great number of organic substances.
§74. Platinum. Pt = 194.9 . Valence two and four.
1. Properties. — Specific i/rariti/ at 17.0°, 21.48 (Deville and Bebray, ('. r., 18fiO.
50, 1038). Meltiny point, 1775° (Yiolle, (".'. r., 1879. 89, 702). 1'ure platinum is .1
tin-white metal, softer than silver, hardened by the presence of other metals,
especially iridium, which it frequently contains. It is surpassed iu ductility
and malleability only by Au and Ag . Platimnn I/lack i.s the-tuiely divide!
metal, a black powder, obtained by retrmving an -a44t«li««-j solution of tire platinous
•:alt with alcohol (Low, H., 1390, 23, 2S9) ; platinum xponye, a gray spongy mass.
M PLATJM M. §7*»$
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 i'orms 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): c. (/., a current of
hydrogen mixed with air ignites when passed over platinum black, also1
hydrogen and chlorine unite. S02 unites with O to form SO3: alcohol is oxi-
dized to acetic acid, formic and oxalic acids to CO- , As"' to AsV , etc.
2. Occurrence.— Found in nature only in the metallic state, gem-rally alloyed
with palladium, iridium, osmium, rhodium, ruthenium, etc. The Ural Moun-
tains furnish the largest supply of platinum.
?>. Preparation. -Usually by the wet method. The finely divided ore is treated
with nitrohydroehloric acid until the platinum is all dissolved. The filtrate is
then treated with lime water to u slightly acid reaction: this removes the
greater part of the Fe , Cu . Ir , Rh , and a portion of the Pd . The filtrate is
now evaporated to dry ness, 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 HNO3 , which
dissolves out the Pb . Cu , Pd , and Bh . The black powder which remains Is
dissolved in nitrohydroehloric acid, the Pb remaining, removed with H..SO, ,
and the Pt precipitated with NH Cl . The precipitate contains a little rhodium,
which is removed by gently igniting the mass with potassium and ammonium
di-sulphate, and exhausting with ^ater, which dissolves out the rhodium
sulphate (§105, 7). In the laboratory the platinum residues are boiled with
KOH or K=CO, and reduced with alcohol. The fine black powder is filtered,
washed with water and hydrochloric acid and ignited.
4. Oxides and Hydroxides.— Platinum forms two oxides. PtO and PtO(,..
I'latiiiviiK hydroxide is formed by treating a dilute solution of platinous potas-
vium chloride with NaOH and boiling (.lorgensen, J. pr., 1877. (2), 16, :544).
A black powder easily soluble in HC1 or HBr . reduced by formic acid to Pt° ,
gentle heating changes it to the oxide PtO . I'laJhiie hydroxide, Pt(OH)4 , is
formed by treating a solution c.f H_.PtCl,. with Na*CO3 in excess, evaporating
\o dryness. washing with water and then with acetic acid. It is a red-brown
powder, soluble in NaOH. HC1 . HNO3 . and H,S04: insoluble in HC2H:,O, .
<Jentle heating changes it to the oxide PtO, tT°pso«N #-, 1870, 3, 4C2).
.1. Solubilities.- -n — 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 nitrohydroehloric acid (to platinic chloride)
less readily than gold. l>. — Oxide* and hydroxide*. — See 4. <•.--&<///*. — 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 rhloride, PtCl, , and bromide, all the
platinicyanides (as PbPt(CN),). and the platinocyauides of the metals of the
:ilkalis a'nd alkaline earths (as K,Pt(CN),). are soluble in water. The platinous
.-ind platinic nitrates are soluble in water, but easily decomposed by it, with the
precipitation of basic salts. The larger number of the inetnUo-phithilc chloride*
or " ehloroplatiii.-ites " are soluble in water, including those with sodium
jNa,PtCl<: or (NaCl),PtCl,]. barium, strontium, magnesium, /inc. aluminum,
copper: and those with potassium, and ammonium, are sparingly soluble in
water, and owe their analytical importance :is eomplete precipitate* to their
insolubility in alcohol. Of the mctallo-platinoii* chloride* (the "chloroplatinites")
— those with sodium [Na,PtCl,]. and barium, are soluble: zinc, potassium and
ammonium, sparingly soluble: lead and silver, insoluble in water. Platinic
iitlphnle. Pt(S04); . is soluble in water (glO).
f>. Heactions. n. Platinous chloride, PtCL . is precipitated by KOH as
Pt(OH). . soluble in excess of the reagent to K.PtO= , pOfcUliMN ^Infinite, which
solution is reduced by alcohol to "platinum black" (1). Plat hi ir chlnride.
PtCl, . a brown-red solid, soluble in alcohol and water, forms with KOH or
NH4OH , not too dilute, a yellow crystalline precipitate of an alkali (K or NH4)
rltlnridc. e. </.. K PtCl, . sparingly soluble in water, soluble in excess
$74,7,1. PLATINUM. 95
of the alkalis and repi'ecipitated by hydrochloric acid. K2CO3 and (NH4)jCO,
give the same precipitate, insoluble in excess of the reagent. A more complete
precipitation of the K or NH4 is obtained by the use of the chlorides. The
sodium platinum chloride. Na,PtCl,, , is very soluble in water and is not formed
by precipitation with sodium salts, ft. — Oxalic acid does not reduce platinum
salts (distinction from gold). A solution of chloral hydrate precipitates pla-
tinum from its solutions. Plarlnous and platinic salts form with cyanides a
great number of double salts, rr.— See :V. <L— Hypophosphorous acid reduces
platinum salts to metallic platinum. Phosphates do not precipitate platinum
salts.
e. Hydrosulphuric 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,
PtS2 , black; slowly soluble in alkali sulphides (Kibau, C. r., 1877, 85, 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 HC1 for
the removal of the SO, (Birnbaum, A., W71, 159, 116).
f. The chlorides of potassium and ammonium are estimated quantita-
tively by precipitation from their concentrated solutions with a solution
of platinic chloride. Potassium iodide colors a solution of platinum
chloride brown-red and precipitates the black platinic iodide, PtI4 , excess
of the KI forming K.2PtIc , 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
rnetal. 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, IN THE USE or PLATINUM APPARATUS WITHOUT
UNNECESSARY 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 PLAT tX I'M. §74, 7, "..
(2} The alkali hydroxides (not their carbonates) and the alkaline earths,
especially baryta and lithia, with ignited platinum in the air, gradually
corrode platinum (by formation of platinites: 2Pt -f- 2BaO -j- 0., :
2BaPt02 . Silver crucibles are recommended for fusion with alkali
hydroxides.
(3) All metals which may be reduced in Ilie fusion — especially compound*
< f 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, xinc, etc., are liable to be rapidly reduced, and immediately to
melt away platium in contact with them.
(4) 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 crnciblo 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 borax
upon them, and by digestion with nitric acid.
8. Detection. — Platinum is identified by the appearance of the reduced
metal; by its insolubility in HC1 or HNO, and solubility in HNO;l -j- HC1 ;
and by its formation of precipitates with ammonium and potassium
ehlorides. Tt 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 evaporated1-, ignited, and 'trie
residue examined and after proving insolubility in HC1 or HNO, , dissolved
in nitrohydrochloric acid and the presence of platinum confirmed with
NH4C1 . Tf the gold and platinum have been precipitated in the second
group with H..S and dissolved with (NH,)..SX they may be separated from
As , Sb , and Sn by dissolving the reprccipitated sulphides in HC1 -f KC10, ,
evaporating to remove the chlorine and boiling after adding KOH in ex-
cess, with chloral hydrate, which precipitates the Au and Pt . leaving the
As , Sb , and Sn in solution. The Au and Pt may then be dissolved in
HNO.. -)- HC1 and separated as directed above. FeS04 may be use to prc,-
eipitate 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: (1) By simple ignition; (2) by precipitation :is
<NH,)2PtCln , K,,PtCl,, . or PtS, and ignition; (.?) by reduction, using Zn . Mg' .
or FeSO, .
10. Oxidation.— Solutions of platinum are reduced to the metallic state by the
• 75. ('«: MOLVRltKM'M. <>7
following metals:' Pb . Ag , Hg . Sn (Sn" to Pt" only), Bi , Cu , Cd , Zn ,
Fe . Fe" . Co, and Ni . Very many organic substances reduce platinum
<-ompounds to the metallic state.
§75. Molybdenum. Mo = 96.0 . Valence two, three, four and six.
1. Properties. — Specific gravity, 8.5(5 (Loughlien, Am. 8., 1868, (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, i4.
forms first a blue oxide, then MoO ; .
2. Occurrence. — Not found native, but occurs chiefly as molybdenite, MoS, ;
as an oxide in molybdenum ochre, Mo03: and as wulf'enite, PbMo04 .
3. Preparation.— (1) By heating the oxide, sulphide or chloride in a current
<»f oxygen free hydrogen (von der Pfordten, B., 1884, 17, 732; Rogers and
Mitchell, J. Am. Soc., 1900, 22, 3CO); (2) by heating with C and Na2C03; (3) by
heating Mo03 with KCN (Loughlien, I.e.).
4. Oxides and Hydroxides. — Moli/bdous hydrox-ide, MoO.xH.O , is formed when
molybdous chloride or nitrate is precipitated with alkali hj'droxides or carbon-
ates, dark brown becoming blue in the air by oxidation. Mo(OH)3 , black,
turning red-brown by oxidation in the air, is formed by treating MoCl; with
KOH; also by electrolysis of ammonium molybdate (Smith, B., 1880, 13, 751).
By heating the hydroxide in a vacuum Mo2O3 is obtained as a black mass,
insoluble in acids. MoO, , a dark bluish mass, insoluble in KOH or HC1 , 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), MoO3 , 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 MoO3 , evolving NO; in hot concentrated sulphuric acid, evolving SO, . The
Tarious lower oxides of molybdenum 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, Mo03 , 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, B., 1871, 4, 14);
the nitrates in dilute nitric acid. The anhydride MoO3 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, Mo03 , which dissolves in excess of the acids.
0. Reactions. — 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. b. — A solution of a molybdate
acidulated with hydrochloric acid gives no red color with KCNS (distinction
from Fe'"); 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 MOLvnni'MM. §75, 6tf.
i
d. — Tribasic phosphoric acid ami its -alt- precipitate, from strong nitric
acid solutions of ammonium molybdate,* somewhat slowly, more rapidly
on warming, annii'inrum phospko-motybdate, 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 fhoupko-molybdate 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 aruno-molybdatf 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 brown or
red-brown precipitate, MoSL . in<>li/l>tl<'iiirin trisulphidt, is obtained after
some time. The precipitate is soluble in ammonium sulphide, better when
hot and not too concentrated, as ammonium thiomolybdate, (NH4)JBtoS, .
from which acids precipitate the trisulphide (Berzelius, Poyg., 1826, 7,
429), soluble in nitric acid, insoluble in boiling solution of oxalic acid
(separation from stannic sulphide).
If Na2S2O, be added to a solution of ammonium molybdatc, slightly acid.
a blue precipitate and blue-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 hypophosphorons 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.
ff. — Arsenic acid and arsenates form, with a nitric acid solution of ammonium
molybdate, a yellow precipitate of iiiiiiinniiinn nrxriin-nniljilHlnlf, in appearance
and reactions not to be distinguished from the ammonium phospho-molybdate;
except the precipitation docs not take place until the solutions arc slightly
warmed, while with phosphates the precipitation begins even in the cold.
Stamioiis salts give with <NH,)..Mo04 a blue solution of the lower oxides of
molybdenum (a delicate test for Sn") (Longstaff, C. N., 1899, 70, C2S2).
Ji. — Solutions of the alkali molybdates are soluble in water and precipilat*-
solutions of nearly all other metallic salts, forming molybdates of the corre-
sponding metals, insoluble in water. <•.</.. K.MoO, + Fb(JTO»), = PbMoO, +
2KNO, .
•The reagent ammonium molybduto, X1I< , MoO,. is prepared by dissolving molybdic acid.
MoO, (100 grams), in ammonium hydroxide (250 re. sp. gr. O.CO with 250 cc. water) cooling, and
slowly pouring this solution into well cooled fairly concentrated nitric: ncid (750 cc. sp. gr. 1.42
with 750 cc. water; with constant stirring.
t Sodium rhospJitt-mnljilnlati- -Soimcnschcin's reagent for ncid solutions of alkaloids— is pre-
par> d as follows : Tho yellow precipitate formed on mixing acid solutions of ammonium molyb-
dato and sodium phosphate— t ho ammonium phospho-molybdate- is well washed, suspended in
vater, and h--ated with sodium carbonate until completely dissolved. Tho solution is evapor
ated to dry ness, and the residue gently ignited till all amnionin i:\ expelled, sodium being sub-
stituted for ammonium. I f blackening: occurs, from reduction of molybdenum, the residue is
moistened with nitric acid, and heated again. It is then dissolved with water :md nitric ncid
tostrongacidulation; the solution bring made ten parts to ono part of residue. It must be
kept from contact with vapor of ammoni:t, both durin-r fie preparation and when preserved
lor use.
575, 10. MOLYBDENUM. !»'.•
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 bead, 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 . An , and
Pt . The solution remaining in the Marsh apparatus is decanted from
the residue (Sn , Sb , An , Pt and excess of Zn) and heated with concen-
trated HNOa , the molybdenum is oxidized to molybdic acid. This solution,
evaporated to dryness, dissolved in ammonium hydroxide and poured into
moderately concentrated HC1 forms a solution of ammonium molybdate
which may be identified by the many precipitation and reduction tests
(6 &, c, d, e, i, etc., 7, and 9). If the molybdenum be present as a molybdate
it may be precipitated from its nitric acid solution by Na2HP04 , washed,
dissolved in ammonium hydroxide, the phosphate removed by magnesia
mixture (§189, 60), and the filtrate evaporated to crystallization (Maschke,
Z., 1873, 12, 380). The crystals may be tested by the various reduction
tests for molybdenum.
9. Estimation. — (/) Molybdic anhydride and ammonium molybdate may be
reduced to the dioxide by heating in a current of hydrogen gas. 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
weighed. (2) Lead acetate is added to the alkali molybdate, the precipitate
washed in hot water, and after ignition weighed as PbMo04 . (3) Volumet-
rically. The molj'bdic acid is treated with zinc and HC1 , which converts it into
MoCls . This is converted into molybdic acid again by standard solution of
potassium permanganate.
10. Oxidation.- "Reducing agents convert molybdic acid either into the l>!n>-
intermediate oxides, or, by further deoxidation, into the black molybdous oxide,
MoO . In the (hydrochloric) acid solutions of molybdic acid, the blue or black
oxide formed by reduction, will be held in solution with a blue or brown color.
Nitric acidulation is, of course, incompatible with the reduction. Certain
reducing agents act as follows:
Cane sugar in the f.-cbly acid boiling solution, forms the blue color — seen
better after dilution; a delicate test. Stannous chloride forms first the blw,
then the ftrotcn, or the (/rcenish brown to black-brown, solution of both the
intermediate oxide and the molybdous oxide. Zinc, with HC1 or H2SO4 , gives
the blue, then green, then brown 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 1810 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.
]00 BISMUTH. §76, 1.
THE COPPER GROUP (SECOND GROUP, DIVISION B).
Mercury (Mercuricum), Lead, Bismuth, Copper, Cadmium (Ruthenium,
Rhodium, Palladium, Osmium).
I
§76. Bismuth, Bi = 208.1 . Valence three and five.
1. Properties. — Spcciflr //rarity, 9.7474 (Classen, B., 1890, 23, 938); meltimj point,
269.22 (Callendar and Griffiths, C. N., 1891, 63, 2): it vaporizes at 1700° and tin-
density of the vapor shows that the molecule Bi has begun to dissociate (Blitz
and V. Meyer, B., 1889, 22, 723). It is a hard, brittle, reddish-white, lustrous
metal; forming beautiful rhombohaedrul 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, c. y., an alloy
of: Bi two. Sn one, and Pb one part by weight melts at 93.7°; and an alloy of:
Bi fifteen, Pb eight, Sn four, and Cd three parts by weight melts at f>8°
" 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 O ). bismnthite (4BijO3.3COo.4H~O). bismuth
glance (Bi2S,), etc.
3. Preparation. — The rock containing bismuth, usually with large amounts
of cobalt, etc., ns 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 KNO or KCN .
4. Oxides. — Bixmuth trioxitic, Bi O , is formed by heating the metal in the
presence of air, or by igniting the hydroxide: it is a pale cjtron-yellow powder.
The hydroxide, Bii OH i , 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. BiOCl (Strohmeyer, Pong., 1832, 26,
549). The meta hydroxide, BiO(OH) , is formed upon drving the orthohydroxide
at 100° (Arppe, Pogff.. 1845, 64, 237). Ilixmiitli pentoaMe, Bi,O5 , 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,
/. P.); or by treating Bi(OH), with three per cent H,On in strong alkaline solu-
tion (Hasebrock. /?., 1887, 20, 213). It is a heavy dark brown powder. At ir>00
•it gives off O, and at the temperature of boiling mercury becomes Bi 0- . It
is decomposed in the cold by HC1 with evolution of chlorine. Bixiinitliir anil.
HBiO, , or more probably Bi._.O:..H,O , is formed upon conducting a rapid
current of chlorine into BilOH), suspended in concentrated KOH solution.
It is a beautiful scarlet red powder which at 120° gives off its water, becoming
BLO, (Mnir, J. C., 1S7<>. 29, 144; Muir and Carnegie, J. C.. 1887, 51, 8(5). It is
doubtful if any alkali salt of bismuthic acid exists, although mixtures of KBiO:1
and HBiO3 are claimed by Hoffmann (A., 1884. 223. 110), and Andre (C. r., 1S'.)1,
113. 860). The so-called bismiitli tctrn.riilr. Bi,O, . is probably a mixture of the
trioxide and peutoxide i£12).
.">. Solubilities.— -n.— Mctnl. — 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. Tt burns
in chlorine with production of light: it combines with 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 bo found in solution when the metal is boiler! with hydro-
chloric acid, "but no more than whon the metal has been boiled with pure water (Ditto and
Metzner, A. Ch., 1896, fG), 29. 389),
i;76, Ga. BIKML'TH. 101
mercial alloys (1). The halogen derivatives of pentad bismuth are not known
(Muir, J. C., 1876, 29, 144). ft — O.m/r.s- and Inidro.rides. — Bismuth oxide, Bi2O3 ,
and the hydroxides, Bi(OH), 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, BLO; , is solu-
ble in HC1 , 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, 5d footnote). '
The chloride, bromide and sulphate are deliquescent.
d. — Water. — 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., BiCl3 -f- H20 = BiOCl -f 2HC1 ,
insoluble in tartaric acid (distinction from antimony, §70, f>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(NO3)3 + H,0 =BiON03 + 2HNO3
^ 12BiON03 + H2O — fiBi,Os,5iN5O5 -f 2HNO3
This is prevented by the presence of one part ammonium nitrate to five
hundred parts water (Lowe, J. pr., 1858, 74, 341).
Bismuth nitrate crystallizes with ten molecules of water, Bi(NOs);!.
10H20 . It is decomposed by a small amount of water forming the basic
nitrate, BiON03 ; 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 oxy chloride 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 BiOI (Schneider, A. Ch., 1857 (3), 50, 488); the
normal sulphate very readily absorbs water to form Bi,(S04).,.3H,0 , which
is decomposed by more water to Bi,0.t.SO.t .
6. Reactions, a. — The alkali hydroxides precipitate from solution? of
bismuth salts bismuth hydroxide, Bi(OH)., , white; insoluble in excess of
the fixed alkalis (distinction from Sb and Sn), insoluble in ammonium
* Lowe (C. 2V., 1883, 45, 296) dissolves the hydroxides of copper and bismuth in glycerol, adds
glucose and gently warms. The copper is completely precipitated and separated from the bis-
muth. Upon boiling1 the filtrate for some time the bismuth is completely precipitated as the
metal.
102 BISMUTH. §76, fib.
hydroxide (distinction from Cu and Cd). The hydroxide is converted by
boiling into the oxide, Bi203 , yellowish white. The precipitation is pre-
vented by the presence of tartaric acid, citric acid, glycerol, and certain
other organic substances (Kohler, J. C., 1886, 50, 428).
The alkali carbonates precipitate basic bismuth carbonate, Bi203.CO3 , white,
insoluble in excess of the reagent. Freshly precipitated barium carbonate
forms the same precipitate without heating.
I). — Oxalic acid and soluble oxalates precipitate bittmiith o.ru-ldt<; Bi2(C2O4), ,
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.
c. — The action of nitric acid upon bismuth and its salts is fully explained
under (r>X </. — Metallic bismuth is precipitated when bismuth salts are warmed
with hypophosphorous acid (separation from Zn .and Cd) (Muthmann and
Mawron. %., 1874, 13, 209). From solutions of bismuth nitrate (;VZ) phosphoric
acid and soluble phosphates precipitate bismuth phosphate, BiPO4 , white,
readily soluble in HC1: 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 HC1) is obtained with soluble phosphates.
e. — Hydrosulphuric acid and sulphides precipitate bismuth sulphide,
Bi2S3 , 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
bixnnith xnlpliiilc. 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, Gfr). The precipitate is readily
dissolved on addition of more hydrochloric or nitric acid (distinction from
ihc 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-
bromide, 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
sails, unless strongly acid, bismuth iodide, black or brownish gray crystals,
quite readily soluble in excess of the reagent * or in strong HC1 without warm-
*Hismutb iodide dissolves in solution of jmtassiura iodide with an intense yellow color, deli-
cate to one-millionth (Stone J. Soc. Chem. Ind., 1887, 0. 416). The potassium iodide solution of
bismuth iodide is used as Dragendorff's reagent to detect tho presence of an alkaloid. Leger
(B/,, 1888, 50, 91) uror Hnchonine and potassium iodide to prove 1 he presence of bismuth. Dtl-
icate to one-five bum: re 1 thousandth. Other metals must be removed.
§76, 9. 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 HC1 , 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 HNO3 .
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 -(-
3K2Sn02 -f- GKOH = 2Bi -f 6KC1 + 3K2Sn03 + 3H20 (Vanino and Treu-
bert, B., 1898, 31, 1113).
A. — Solutions of bismuth salts, nearly neutral, poured into a hot solution of
potassium bichromate precipitates the orange^red chromate, (BiO)2Cr2O7; but
if poured into a cold solution of the neutral chromate a citron-yellow precipi-
tate, .'{Bi2O:1.2CrO;, , is formed. These precipitates are soluble in moderately
concentrated acids, insoluble in fixed alkalis (distinction from Pb). The pre-
cipitate with X;,Cr2O7 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 flame, form-
ing an incrustation (BLOj), 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 use of the blow-pipe.
The reagent mixed with the \mknown is fused on charcoal or on a piece of
aluminum sheet. A red sublimate indicates bismuth. Mercury gives a mix-
ture of red and yellow sublimates (Hutchings, C. N., 1877, 36, 249).
Bismuth chloride may be, sublimed at the temperature of boiling sulphur;
recommended as a separation from lead (Remmler, B., 1891, 24, 3554).
8. Detection. — Bismuth is precipitated from its solutions by H2S form-
ing Bi2S3 . By its insolubility in (NH4)2Sxand solubility in hot dilute
HN03 it is separated with Pb , Cu , and Cd from the remaining metals
of the tin and copper group. Dilute H2S04 removes the lead and
ITH4OH precipitates the bismuth as Bi(OH), , leaving the Cu and Cd in
solution. The presence of the bismuth is confirmed by the action of a
hot solution of K2Sn02 on the white precipitate of Bi(OH)3 , giving metallic
bismuth (6g) or by dissolving the Bi(OH)3 in HC1 and its precipitation as
BiOCl upon dilution with water (5d).
9. Estimation. — (1) As metallic bismuth formed by fusion with potassium
cyanide. (2) As Bi.2O3 formed by ignition of bisimith salts of organic acids, or
of the salts of volatile inorganic oxyacids. (3) By precipitation by H2S , and
* For a modification of this test sec Muir (J. C., 1877, 32, 45).
104 COPPER. £76, 10.
after drying at 100°, weighing as Bi S . (,J) By precipitation by K,Cr.,O7 , and
after drying at 120°, weighing as (BiO),Cr..,O7 . (.;) Volunu'tricall-i/. By precipi-
tation with K2Cr.,OT . Dissolve the chromate in dilute acid, transfer to an
azotometer and reduce the chromate with hydrogen peroxide (Hauinann, Z.
<tu<i<'w., 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, G74).
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 , Fe , Al , Zn , Mg , and HH,PO, (6d). All
salts of bismuth are oxidized to Bi,0- by Cl or H.,0., in strong alkaline
mixture (Hasebrock, /*., 1887, 20, 213; Schiff, A. Ch., 1861 (3), 63, 474).
All compounds of bismuth arc reduced to the metal by potassium stannite
X.,SnO., (G<7). 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. — Specific gravity, electrolytic, 8.914; melted, 8.921; natural crya-
tals, 8.94; rolled and hammered sheet, 8.952 to 8.958 (Marchand and Scheerer,
J. pr., 1866, 97, 193). Melting point, 1080.5 (Hey cock and Neville, J. C., 1895, 67,
190). A red metal, but thin sheets transmit a greenish-blue light, and it also
shows the same greenish-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 CuFeS : copper glance, Cu.S; green malachite, CUo(OH)oCOs;
blue malachite, Cu3(OH)J(CO3).!; red copper ore. Cu.O: 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; (3) reduction of the oxide by ignition with carbon, carbon
monoxide, illuminating gas, or other forms of carbon; (.'/) reduction of the
•oxide by K or Na at a temperature a little above the melting point of these
metals; (-5) reduction by fusion with potassium cyanide: CuO + KCN = Cu +
KCNO . For its reduction in the wet way, see 10.
4. Oxides and Hydroxides. — Cuprous oxide (Cu,O), red, is found native; it is
prepared: (1) by reducing CuO by means of grape-sugar in alkaline mixture;
(2) by igniting CuO with metallic copper; (3) by treating an ammoniacal cupric
solution with metallic copper; then adding KOH and drying. Cuprous hydrox-
ide, 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 cupric salts in the air; or by
heating- the metal in a current of air. Cupric hydroxide, Cu(OH)2 , is formed
by precipitating cupric salts with KOH or NaOH . It is stated by Rose (l'o<i<}.,
1863, 12O, 1) that tftf(iciii>ric mono-ridc, (Cu,O , is formed by treating a cupric
salt with KOH and a quantity of K,SnO, insufficient to reduce it to the metallic
state. A jxroj'ide of copper, CuO. , is supposed to be formed by treating
Cu(OH), with H,O., at 0° (Kriiss, Ii., 1884, 17, 2593). (§10.)
5. Solubilities. — a.— Metal. — Copper docs not readily dissolve in acids with
evolution of hydrogen; it dissolves most readily in nitric acid chiefly with
evolution of nitric oxide' 3Cu + SHNO3 = £Cu(lTO,), + -!H2O + 2ND (Freer
and Higley, Am., 1899, 21, 377); also in hot concentrated sulphuric acid, with
evolution of sulphurous anhydride: Cu + 2H,S04 = CuS04 + 2H,O + 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 luELO to one HC1 pre-
vents the action (TCngel, C. r., 1895, 121, 528). Hydrobromic 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, (NH4)2S , colorless, acts upon
copper turnings with evolution of hydrogen, forming Cu2S (Heumann, J. C.,
1873, 26, 1105).
b. — 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. Cupric 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 tartratcs
and fixed alkalis (but precipitated as Cu20 by heating with glucose) (sepa-
ration from Cd and Zn) (Warren, C. N., 1891, 63, 193); insoluble in
ammonium hydroxide in absence of ammonium salts (Maumene, /. f .,
1882, 42, 1266):
c.— Salts. — All salts of copper, except the sulphides, arc soluble in am-
monium hydroxide. All cuproiis 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
blue 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 deliquescent; the sulphate, permanent; the acetate, efflorescent.
Cupric basic carbonate, oxalatc, 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 axe-insoluble; the chloride and nitrate, soluble. Ether dissolves
the chloride.
106 COPPER.
(>. Reactions. — a. — Fixed clkcii hydroxides precipitate acid solutions of
cuprous chloride, first as the white cuprous chloride, cl>an£iiig- with more of
the alkali to the yellow cuprous Ji)/dro.ri(lc, insoluble in excess. Ammonium
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 carbo-natc, Cu2CO3 .
Fixed alkalis — KOH — added to saturation in solutions of cupric salts,
precipitate cupric hydroxide, Cu(OH)., , 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, Cu.,0,(OH).,; by boiling, to CuO .
If tartaric acid, citric acid, grape-sugar, milk-sugar, or certain other
organic substances are present, the precipitate cither 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 fi.rcd alkali x): 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 CuS04.(NH,)4 (§77, 56). The deep
blue solution probably consists of this compound in a hyd rated condition,
i. f. Cu(OH),.-,>NH4OH.(NH4),S04 : or (NH4)4Cu(OH)4S04 . Other salts
than the sulphate form the corresponding compounds: CuCl. -j- 4NH(OH
= Cm(OH)2.2HH4OH.2HHiCl . 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, •'>."> 1). Ammonium carbonate, like ammonium hydroxide, precipitates
and redissolves to a blue solution. Carbonates of fixed alkali metals — as
K ,CO, — precipitate the greenish-blue, basic carbonate, Cu.,(OH)2CO.j , 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 concent rated solution of a fixed
alkali precipitates the blue hydroxide, changed on boiling to the black
oxide, CuO .
It. — Oxalates, cyanides, ferrocyanides, ferricyanides and thiocyanates pre-
cipitate their respective cuprous salts from cuprous solutions not too strongly
acid. The ferricyanide is brownish-red, the others are white. The thiocj-anate
is used to separate copper from palladium (Wohler, A. Ch., 1867, (4), 10, 510);
and also from cadmium. In solutions of cupric salts, oxalates precipitate cupric
§77, 6c. COPPER. 107
oxalate, CuC204 , bluish-white, insoluble in acetic acid, and formed from mineral
r,cid salts of copper by oxalic acid added with alkali acetates.
Potassium cyanide forms the yellowish-green cupric cyanide, Cu(CN)2 ,
soluble in excess of the reagent with formation of the double cyanide,
2KCN>.Cu(CN)._) , 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,
€u2Fe(CN)u , 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 ferri-
cyanide precipitates cupric ferricyanide, Cu3(Fe(CN)0)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. (7.,
1880, 38, 748). Ammonium benzoate (10 per cent solution) precipitates copper
salts completely from solutions slightly acidified (separation from cadmium)
(Gucci. B., 1884, 17, 2(>59).
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) (Kiiorre, B., 1887, 20, 28:?).
Potassium xanthate gives with very dilute solutions of copper salt a yellow
coloration; according to Wagner (I.e.) one part copper in 900.000 parts water
may be detected.
c. — Nitric acid rapidly oxidizes cuprous salts to cupric salts, d. — A solution
of cupric sulphate slightly acidulated with hydrochloric acid is precipitated as
cuproTis chloride by sodium hypophosphite (Cavazzi, Gazzetta, 1886, 16, 167); if
the slightly acidulated copper salt solution be boiled with an excess of the
hypophosphite the copper is completely precipitated as the metal. Sodium
phosphate, Na.HPO4 , gives a bluish-white precipitate of copper phosphate,
CuHP04 . if Ihe reagent be in excess and Cus(P04)2 if the copper salt be in
excess. Sodium pyrophosphate precipitates cupric salts, but not if tartrates
or thiosulphates be present (separation from cadmium) (Vortmann, B., 1888,
21, HO.1?).
«. — Cuprous salts (obtained by treating cupric salts with SnCL) when boiled
with precipitated sulphur deposit The copper as CucS (separation from cad-
mium) (Orlowski, ,/. f1., 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 H,S gives CuS , black (with some Cu,S), produced
alike in acirl solutions (distinction from iron, manganese, cobalt, nickel)
* Freshly precipitated cuprous sulphide transposes silver nitrate forming silver sulphide,
metallic silver and cupric nitrato : with cupric sulphide, silver sulphide and cupric nitrate are
formed (Schneider, Poijg., 1874, 152, 471). FiosLly precipitated sulphides of Fe, Co, Zii, Cd,
Pb, Ui, Sii", and Snlv, when boiled with CuCl in presence of NaCl give Cu.,S and chloride of
the mct:il : with CnCl, . CuS and a chloride of the metal are formed, except that SiiS gives
CiijS, CuCl and Si»IT (Kaschig, «., 1884, 17, 697).
108 COPPER. §77, 6f.
and in alkaline solutions (distinction from arsenic, antimony, tin). — Solu-
tions containing' only the one-hundred-thousandth of copper salt are
colored brownish by the reagent. The precipitate, CuS , is easily soluble
by nitric acid (distinction from mercuric sulphide); with difficulty -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).
Concerning1 the formation of a colloidal cupric xitlpliidc, see Spring (B., 1883,
16, 1142). According to Brauner (C. A'., 1896, 74, 99) cuprie salts with excess
of hydrogen sulphide always yield a very appreciable amount of cuprous
sulphide. See also Ditte (C. r., 1884, 98, 1492). Solutions of cuprie salts are
reduced to cuprous salts by boiling with sulphurous acid (Kohner, C. C., 1886,
813). Sodium thiosulphate added to hot solutions of copper salts gives a black
precipitate of cuprous sulphide. In solutions acidulated with hydrochloric
ucid, this is a separation from cadmium (Vortmann, .I/., 1888, 9, 165).
/. — Hydrobromic acid added to cuprie 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 dissolves 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. , H..SO, ,
FeS04' (b).
(a) 2CuSO4 + 4KI = 2CuI + I2 + 2K,S04
(6) :.>CuS04 + 2KI + 2FeS04 = 2CuI + K.S04 + Iea(S04)3
2CuS04 + IKI + H2S03 + H20 = 2CuI + 2K2SO4 + H2S04 + 2HI
//. — Arsenites. ns KAsO , or arsenous acid with just sufficient alkali hydrox-
ide to neutralize it, precipitate from solutions of cupric salts (not the acetate)
the green < •«/>/*•*• ursenitt, chiefly CuHAsO., (Scheele's green. " I'aris green"),
readily soluble in acids and in ammonium hydroxide, decomposed by strong
potassium hydroxide solution. From cupric acetate, arsenites precipitate, on
boiling, copper iHrlo-iirxi-nitc. (CuOAs,0:,),,Cu(C,HsO,), • Scliveinftirt green or
Imperial green, " Paris green." dissolved by ammonium hydroxide and by
acids, decomposed by fixed alkalis.
Soluble arsenates precipitate from solutions of cuprie salts '•///*/•/<• anflMte,
bluish-green, readily soluble in acids and in ammonium hydroxide. .
h. — Potassium bichromate does not precipitate solutions of cupric sult-
*The precipitation is incomplete unless the free iodine.one of the productsof the reaction, is
removed by means of a reducing: agent (|44.i.
$77, 10. COPPER. 109
normal potassium chromate forms a brownish -red precipitate, soluble in am-
monium hvclroxide to a green solution, soluble in dilute acids.
7. Ignition. — Ignition with sodium carbonate on charcoal leaves metallic
copper in finely divided grains. The particles are gathered by triturating the
charcoal mass in a small mortar, with the repeated addition and decantation
of water until the copper subsides clean. It is recognized by its color, and
its softness under the knife. Copper readily dissolves, from its compou-drs in
beads of borax and of microcosmic salt, in the outer flame of the blow^ !pe.
The beads are green while hot, and line* when cold. In the inner flame the
borax bead becomes colorless when hot; the microcosmic salt turns dark green
when hot, both having a reddish-brown tint when cold (Cu.O) (helped by add-
ing tin). Compounds of copper, heated in the inner flame, color the outer flame
green. Addition of hydrochloric acid increases the delicacy of the reaction,
giving a greenish-blue color to the flame.
8. Detection. — Copper is precipitated from its solutions by H2S , form-
ing CuS . By its insolubility in (NH4)J3xand solubility in hot dilute HNO,
it is separated with Pb , Bi , and Cd from the remaining metals of the tin
and copper group. Dilute H.,S04 Avith C..H-OH removes the lead and
ammonium hydroxide precipitates the bismuth as Bi(OH).j , leaving the
Cu and Cd in solution. The presence of the Cu is indicated by the blue
color of the ammoniacal solution, by its precipitation as the brown ferro-
cyanide after acidulation Avith HC1 (6&); 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. Estimation. — (1) It is precipitated on platinum by the electric current or
by means of zinc, the excess of zinc may be dissolved by dilute hydrochloric
acid. (2) It is converted into CuO and weighed after ignition, or the oxide is
reduced to the metal in an atmosphere of hydrogen and weighed as such.
(3) It may be precipitated either by HoS or NaoS.O, , and, after adding free
sulphur and igniting in hydrogen gas. weighed as cuprous sulphide, or it may
be precipitated by KCNS in presence of H..,SO3 or H3PO2 , and, after adding S ,
ignited in H and weighed as Cu,S . Cu,0 , CuO. Cu(NO,)2,CuC03 , CuSO4 ,
and many other cupric salts, are converted into Cu2S by adding S and igniting
in hydrogen gas. (.)) By adding KI to the cupric salt and titrating the liber-
ated I by NaoS.O3: not permissible with acid radicals which oxidize HI.
(5) By precipitation as Cul and weighing after drying at 150° (Browning.
Am. *SY., 18!):? [.'{], 46, 280). (6) By titrating in concentrated HBr , using a
solution of SnCL in concentrated HC1: the end reaction is sharper than with
SnCL alone (Etard and Lebeau, C. r., 1890, 110, 408). (7) By titration with
Na2S. Zinc does not interfere (Borntriiger, Z. angcw., 1893, 517). (8) By-
reduction with SO, and precipitation with excess of standard NH4CNS; dilu-
tion to definite volume and titration of the excess of NH.,CNS in an aliquot
part, with AgN03 (Volhard, A., 1878, 190, 51). (9) Small amounts are treated
with an excess of NH4OH and estimated colorimetrically by comparing with
standard tubes.
10. Oxidation.— Solutions of Cu" and Cu' are reduced to the metallic
state by Zn , Cd , Sn , Al , Pb , Fe , Co , Ni , Bi , Mg *, P , and in presence
of KOH by K2Sn02 . A bright strip of iron in solution of cupric salts
acidulated with hydrochloric acid, receives a bright copper coating, recog-
nizable from solutions in 120,000 parts of water. With a zinc-platinum
* Warren, C. IV., 1895, 71, 93.
110 CADMIUM. §78, 1.
couple the copper is precipitated on the platinum and its presence can be
confirmed by the use of H.,S04 , concentrated, and KBr , an intense violet
color is obtained (Creste, /. C., 1877, 31, 80,3). Cu" is reduced to Cn' bj
Cu° (Boettger, •/. C., 1878, 34, 113), by SnCl, in presence of HC1 , in
presence of KOH by As20, and grape sugar, by HI , and by S02 . Metallic
copper is oxidized to Cu" by solutions of Hg", Hg', Ag', PtIV, and An."',
these salts being reduced to the metallic state. Ferric iron is reduced to
the ferrous condition (Hunt, Am. 8., 1870, 99, 153). Copper is also oxi-
dized by many acids.
§78. Cadmium. Cd = 112.4 . Valence two.
1. Properties. — Specific gravity, liquid, 7.989: cooled, 8.G7; hammered, 8.6944.
Melting point, 320.68° (Callendar and (irittiths, C. A7., 1891, 63, 2). Roiling point,
763° to 772° (Carnelley and Williams, ./. C., 1878, 33, 284). Specific heat is 0.0567.
Vapor density (H = 1), 55.8 (Deville and Troost, .4. Cli., 1860, (3), 58, 257). From
these data the gaseous molecule of cadmium is seen to consist of one atom
(Richter, Anory. Chan., 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 line wire, but at 80° it is brittle.
Upon bending1 it gives the same creaking sound as tin. It may be completely
distilled in a current of hydrogen above 800°, forming silver white crystals
(Kammerer, B., 1874, 7, 1724). Only slightly tarnished by air and water at
ordinary temperatures. When ignited burns to CdO . When heated it com-
bines directly with Cl , Br , I , P , S , Se , and Te . It forms many useful alloys
having low melting-points.
2. Occurrence. — Found as 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 «./•»//•, 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 tytfnKrfcfe-Cd(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. r., 1SS5, 100, 1497). fc.— The oxide and In/droaiidc
are insoluble in water and the fixed alkalis, soluble in ammonium hydroxide,
readily soluble in acids forming salts: soluble in a '-old mixture of fixed alkali
and alkali tartrate. rep'-ecipitated upon boiling (distinction from copper)
(Behal, •/. I'ltann.. is,s:,. (:,). 11, 553). r.— Salts.— The sulphide, carbonate,
oxalate, phosphate, cyanide, ferroeyanide and ferricyanide are insoluble ($27)
in water, soluble in hydrochloric and nitric ::eids, and soluble in NH4OH ,
except CdS. The chloride and bromide are dclitiiiwi'tit, the iodide is perma-
nent; they are soluble in water and alcohol.
6. Reactions, 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(OH)., , white, insoluble
£78, Gi. 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., CdCl2(NH3)3 ,
CdS04(NH.,)4 , Cd(NOn),(NH,),; (Andre, C. r., 1887, 104, 908 and 987;
Kwasnik, Ardi. Pliarm., 1891, 229, 569). The fixed alkali carbonates pre-
cipitate cadmium carbonate, CdCO., , 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.
ft. — Oxalic acid and oxalates precipitate cadmium oxalate, white, soluble in
mineral acids and ammonium hydroxide. Potassium cyanide precipitates
cadmium cyanide, white, soluble in excess of the reagent as Cd(CN)....2KCN;
ferrocyanides form a white precipitate: ferricyanides a yellow precipitate,
both soluble in hydrochloric acid, and in ammonium hydroxide. Potassium
sulphocyanat© 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 boiling, cadmium oxide is precipitated (separation
from copper and zinc) (Aubel and Ramdohr, A. Ch., 1858, (3), 52, 109).
«.— 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 im
hot dilute sulphuric acid and in a saturated solution of sodium chloride *
(distinction from copper) (Cu»hman, Am., 1896, 17, 379).
Sodium thiosulphate, Na,,S203 , does not precipitate solutions of cadmium
salts (Follenius, Z., 1874, 13, 438), but in excess of this reagent, ammonium
salts be'ing absent, sodium carbonate completely precipitates the cadmium as
carbonate (distinction from copper) (Wells, C. A7., 1891, 64, 294). Cadmium
salts with excess of sodium thiosulphate are not precipitated upon boiling1
with hydrochloric acid (distinction from copper) (Orlowski, J. C., 1882, 42, 1232).
/. — The non-precipitation by iodides is a distinction from copper, g. — Soluble
arsenites and arsenates precipitate the corresponding cadmium salts, readilj
soluble in acids and in ammonium hydroxide, h. — Alkali chromates precipitate
yellow cadmium chromate from concentrated solutions only, and sohible OB.
addition of water.
i. — A solution of copper and cadmium salts, very dilute, when allowed to
spread npon a filter paper or porous porcelain plr.te. gives a ring of the cad-
mium salt beyond that of the copper salt, easily detected by hydrogen sulphide
(Ragley, ,/. •#., 1878, 33, 304).
•Owing to the formation of incompletely-dissociated CdCl.,. CdI2 is still less dissociated
and accordingly CdS dissolves more readily in HI than in HC1 and much more readily than in
H NO,, of the same concentration. On tho other hand, of course, precipitation of the sulphidu
takes place with more difficulty from the iodide than from the other salts.
112
REACTIONS OF BIUMl'TH, COl'l'KK AM) CM) MUM.
, 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 K.S, a yellow precipitate
of CdS is obtained (distinction from zinc) (Chapman, J. C.. 1877. 31, 400).
8. Detection. — Cadmium is precipitated from its solutions by H2S form-
ing CdS. By its insolubility in (NH4).,SX and solubility in hot dilute HNO.
it is separated with Pb , Bi , and Cu from the remaining metals of the
second group. Dilute H,S04 with C..H-OH removes the lead and NH4OH
precipitates the bismuth as Bi(OH) , 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 HJ5 .
]f Cu be absent the yellow CdS is obtained at once from the ammoniacal
solution with HJ5 . See also Oi.
{». Estimation. — (1) It is converted into, and after ignition weighed as an
oxide. (^) Converted into, and after drying- at 100°, weighed as CdS. (.'*) Pre-
cipitated as CdC..,O4 and titrated by KMnO4. (.}) Klectrolytically from a slightly
ammoniacal solution of the sulphate or from the oxalate rendered acid with
oxalic acid. (J) Separated from copper by XI: the I removed by heating; the
excess of XI removed by XNO., and H SO : the cadmium precipitated by
Na,CO, and ignited to CdO (Browning, Alii. X., 1893, 146, 280). (6) By adding
a slight excess of H.SO, to the oxide or salt, and evaporation 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.
Tali-en in Solutions of their Chlorides. -Nitrate*.
or Acetates.
Bi
Cu
Cd
KOH i.r NaOH. in
C "PC" '
Bi(OH)», white.
Cu(OH),., dark
bine.
Cd(OH)2, white.
NH4OH, in excess
Dilution of sntu-
Bi( OH),, white.
BiOCl, white (§76,
-,ij)
Mine solution.
Colorless solution.
Iodides
Partial precipita-
Precipitation of
Cul with libera-
Sulphides
n6t very strongly
ai-id (§76, Of).
Bi S . black, in-
tion of iodine
($77, en.
Cu.S and CuS.
CdS. yellow, insol-
Iron or zinc
soluble in KGN.
Bi spongy precipi-
black, soluble in
KCN.
Cu. bright coating
uble in KCN.
Cd. gray sponge
Glucose, KOH, and
tate.
Bi. black.
(§77. 10).
Cu,O. yellow (§77.
- 1,\
with /inc, no ac-
tion with iron.
neat
K.SnO, + KOH..
Bi. black.
.>»).
Cu, precipitated
metal.
§81,4. PRECIPITATION OF METALK OF SECOND GROUP. 113
SYSTEMATIC ANALYSIS OF THE METALS OF THE Tix AND COPPER GROUP.
The precipitation of the metals of the second( group (Tin and Copper
Group) bv hvdrosulphuric acid, and their separation into Division A (Tin
Group) and Division B (Copper Group). See §312.
§80. Manipulation. — The filtrate f-rom Group I. (§62), or the original
solution, if the metals of the silver group be absent, is rendered acid with
a few drops of HC1 . warmed and saturated with hydrosulphuric acid gas.
2H3As04 + xHCl + 5H2S = As2S5 + xHCl + 8H2O
or 2H3As04 + xHCl + 5H2S =As,S3 + xHCl + S3 + 8H2O
SnCl4 + 2H2S = SnS, +• 4HC1
SnCL + H2S = SnS + 2HC1
2Bi(N03)3 + 3H2S = Bi2S3 + 6HN03
CdSO4 + H2S = CdS + H2S04
The precipitate, after being allowed to settle a few minutes, is filtered and
thoroughly washed with hot water containing a little HC1 . 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 ajcid 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 nay be used instead of sulphuric. The gas
should be passed through ?. 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 unknown solution with H2S 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.
/?. H2S is decomposed by HN03 or HNO3 + HC1 (nitrohydrochloric acid)
(§257, 6#), hence these acids must not be present in excess. If these acid;;
were used in preparing the solutions for analysis, they must be removed by
evaporation. Sulphuric acidulation is not objectionable to precipitation with
H,S , 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 HC1
and prepares the solution for precipitation with H2S , if other acids are not
present in excess and if too much HC1 was not employed. The presence of a
great excess of HC1 does not prevent the precipitation of arsenic (§69, <>e), but,
does hinder or entirely prevent the precipitation of the other metals of this
group, especially tin, lead (§57, fie), cadmium and bismuth. The solution must
be acid or traces of Co , Ni and Zn (§135, Ge) will be precipitated. No instruc-
tions can be given as to the exact amount of HC1 to be employed. About one
part of HC1 to 25 of the solution should be present to prevent the precipitation
114 PRECIPITATION OF METALS OF 8ECOXI) GROUP. §81,5.
of Zn , and it is seldom advisable to use more than one part of HC1 to ten of
the solution * (this refers to the reagent HC1 , §324).
5. 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
H2S , and before filtering1 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, tic).
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 , SnSo , As.,8,, and As..,Sr> are lemon-yellow; Sb2S:>
and Sb.S, are orange: SnS . HgS , PbS , Bi^Sj , Cu.S and CuS are black to
brownish-black. If too much HC1 be present, lead salts frequently precipitate
a red double salt of lead chloride and lead sulphide (§57, fie^ Mercuric chloride
at first forms a white precipitate of HgCL.'JHgS , changing from yellow to red,
and finally to black with more H..S , due to the gradual conversion to HgS
(§58, 6e).
7. Addition of water to the solution before passing in HoS may cause the
precipitation of the oxychlorides of Sb . Sn or Bi (5rf; §70, §71 and §76). These
should not be redissolved by the addition of more HC1 , as they are readily
transposed to the corresponding sulphides by H S. 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.
iS. 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, when a further treatment with H...S causes a precipitation of the
arsenic as As S. 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,
t>'e,l).
9. The presence of a strong oxidizing agent as HNO , K.Cr.O, . Fed, , etc.,
causes with H2S the formation of a white precipitate of sulphur (§125, 6e),
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, tif and 10) to effect reduction of a possible higher
oxidized fonn of Cr or Mn before the precipitation with H,.S . thus avoiding
the unnecessary precipitation of sulphur. -
10. Complete precipitation of the metals of the second group with H...S may
fail: (/) from incomplete saturation with the gas (§81, 2); (,?) from the pres-
ence of too much HC1 (§81, -I): (.?) from the presence of much pentad arsenic
(§69. <:<•). The first cause of error may be avoided by card til observance of the
directions in note (2). To prevent the second cause of < rror a portion of the
filtrate, after the removal of the precipitate by filtration, should be largely
diluted with water (10 volumes) and H.S (gas or saturated water solution)
again added. In case a further precipitate is obtained, the whole of the filtrate
should be diluted and again precipitated with H.S. 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, H,S should be passed into the
* Addition of a strong acid, containing » ions in large quantity, diminishes the already slight
dissociation of the H,S (§44>, thus decreasing in number ihe S ions, whose concentration multi-
plied by that of the metal ions must equal the solubility-prod-.ict 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 H2S. This appears strange in view
of the fact that the acid which prevented precipitation and the acid which finally 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-dissociated body as H2S
than where 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 8
tons and so the solubility-product is reached and precipitation results.
$83, 4. PRECIPITATION OF METALS OF SECOND GROU'. 115
hot solution for fully r.O minutes (Note o) or the solution should be treated
with SO2 or some other agent for the reduction of Asv to As'" (§69, 10).
£82. Manipulation.— 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
:^ possible. Yellow ammonium sulphide (NH4)2SX (§83, 2) is then added
and the precipitate digested for several minutes with warming:
As2S3 + 2(NH4),S2 == (NH4)4As2S5 + S2
^SnS + (NH4)3SS = (NH4)2SnS3
2SnS2 + 2(NH4),S2 = 2(NH4)aSnS, + S2
2Sb2S3 + 6(NH4)2S, = 4(NH4)3SbS4 + S2
2MoS3 + 2(NH4)2S2 = 2(NH4)2MoS4 + S2
The precipitate is then filtered and washed once or twice with a small
amount of (NH4)2S* , and then with hot water. The filtrate consisting of
solutions of the sulphides of As , Sb , Sn , An , Pt , Mo (Gr, Ir, Se, Te,
W, V), constitutes the Tin Group (Division A of the second group). The
precipitate remaining upon the filter, consisting of the sulphides of Hg ,
Pb, Bi, Cu, Cd (Os, Pd, /?//, and Ru), constitutes the Copper Group
(Division B of the second group, §95).
§83. Notes.— I. The precipitate of the sulphides of the tin and copper group
must be thoroughly washed with hot water (preferably containing H2S and
about one per cent of reagent HC1 to prevent the formation of soluble colloidal
sulphides (§69, 5c), to insure the removal of the metals of the iron and zinc-
groups, which would be precipitated on the addition of the ammonium sulphide
(§144).
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 HC1 .
The sulphides of the tin group are soluble in the fixed alkali sulphides, K,S
and Na2S; cupric sulphide is insoluble in these sulphides. Mercuric sulphide,
however, is much more soluble in fixed alkali sulphides than cupric sulphide is
in the (NH4)2SX. Tf 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).
Hut if Hg" be present, the (NH4),,SX should be used, and th« 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. ATI excess of the reagent is to be avoided, as the acidulation of the
solution causes the precipitation of sulphur (§256, 3), which may obscure the
precipitates of the sulphides present.
116
TABLE FOR THE ANALYSIS OF THE TIN GROUP.
§84.
C4
CO
coo
w
m
CH
P
O
«
o
a
x
o
o
a
05
«
O
O,
P
O
K
O
W
CL,
Oi
O
O
<n
H
.J
•<
!?5
<1
«
O
h
H
•J
-
* ^3 co
ai r ** «
2 * o
s i •
s£*
O ^5 '
BsM
§ .t; CQ
P*
|| a
KJI
c ® u
_
III
p o 3
•f W •
*• Ja i
-
" ^
. j
c •* ^ a
c eo ^ 2
>.OD T3 ^
SS g «
5 S
TABLE FOR THE ANALYSIS OF THE TIN GROUP.
117
0 o
~
«
-g
e
o " re <* P«.g AS
4) .2
0-= ^
^H W .P^ ^
j^«(-, +j re
re cc -2 ;>>;£
^ c _oo
§ •-'i.oo805
sen , to-, r^
««O ••• G
=fc:-:
-^ --00 00
X u ae»cos
"^'Q «Tto
x .
<u K5 T;
-^ I-
5
-t3 SO5
X
4) T3
-^ c
37+- "o «» 5 P o o «
MOM1^ 2 ^ ^ -^
c
C P
re re
coco
t-05
-8
IF* £
c ^ ^
'— l> 05
'J eosoos
1*
-o
o
-u
CS
8
.1
CO 5)
w*
on zj
S «s
o °^
» fl>
OD 2
« j-
w p,
bl p,
J o
is
a i:
is
=* §
ag
-8s
_2 5,5
^^
S c »
bsen, it is bette
u n any case.
,6).
t,
in
8
0 5
.. o «
S — 1-
^2 — . +J
b'^ «
S d
o 3 c
II a
S£o
•o g £
* o-S
•** ,r -»
0 * S
S 5 8
o i; o>
£ a i-
ft-g &
01 O
& «-
ft *
5 ® X
p, -u o
&3 5
8 * »
t < ®
* +- "S
v n a
fl) f-3 03
ti oo -3
13 ^ S
•g- 2
•So®
log
*1a
^ a
-j H tw
s § S
"^•2
111
s s S
ginal
dicai
r cen
be
OJ ^ W
1 1 §
en .g- o3
* g'S
s! £ O
P ft 9-
In m
A bl
no
fi fi ^ fi
.2 .2 ° '"
5 5 * *
II ^
»
+3 p "5 *"
H "£ JS •*
jf 1 1
f la's
'ag-o*
8 « S fi
fSff
o a B c
32*£
£j ® <w **
o^o-
•5 <e a <
"S 2 $•«
ils'S
?§^'
&2"£
P « ^ 5
o fe <H -S
<H P< O 00
to * ® -S
-u -rt _O >
ic "3 <D a
ail."
^ -ffi S
ss
io
n
•^ co
l!
« «
s *
V fi
1 *
o ^
« ?
fi 5
- *5
r-i >» O
r S
2 £?
^ cs
. ' " »
•*> r3 —
O ®
o 5
s ^
e«
CS
go
« s
•§£
2 5
3 ID
O J3
O -t-
« _
S «
^ .s
O rt
11
.li
X O 03
III
Pi
s
e produ
ucod tin
as above.
e
r
o
d
p
e
ft 0 &
•9 4$ ft
if n
pre
HC
DIRECTfOys l-'Dlf IN l/.r.v/x 117 777 .YO'/7,'X. ^85.
Jj85. Manipulation.— The solution of tin- sulphides in (NH4).,SX is care-
fully acidulated with hydrochloric acid:
2(NH,),S, + -1HC1= 1NH.C1 + S, + 2H,S
(NH4),As2S., + -4HC1 = As.S, + 4NH,C1 + 2H,S
2(NH,)aSbS4 + <>HC1 = Sb,S, 4- c,NH4Cl + :;H,S
(NHJ.SnS, + 2HC1 = SnS, + 2NH4C1 + H,S
The precipitate obtained when the metals of the tin group are present
is usually ycllcr.' or orange-yellow and is easily distinguished from a pre-
cipitate of sulphur alone (SnS and MoS, are brownish-black). It should
be well washed v.'itli hot water and then dissolved in hot HC1 using small
fragments o ! -CIO, ($69, tic] to aid in the solution:
2As,B3 -:- 10C1, + 10H,O = -iH.AsO, + :.'()HC1 + :;S,
CnS, + :HC1 = SnCl, + 2H,S
rts, :- :.-ia = ptci4 4- s.,
The solution is boiled (to insure removal of the chlorine (i<69. 10) until it
no longer bleaches litmus paper.
§86. \nt<'N. — /. If the i>rei'ipitatr obtained is white, it probably consists of
sulphur alone and indicates absence of more than traces of the nietals belong-
ing to this group (GeS, is \\liite. §111, (>).
2. Care should be taken not to use too much HC1 in precipitating the sul-
phides from the (NH,),SX solution. ;is some of the sulphides (especially SnS,)
are quite soluble in concentrated HC1 .
.{. It will lie noticed (585) that the lower sulphides of Sb and Sn arc o\idi/i-d
by the (NH,),SX. and are pi\ c-ipitated by the HC1 as the higher sulphides
Sb^S.-, and SnS. respect ivel\p. This fact may be most readily observed bv the
precipitation of a solution o, SnCL with H,S . frivinjr a brown precipitate of
SnS. then dissolving this precipitate in (NH,)..SX and reprecipitatinir with HC1
as the orange-colored SnS. .
}. Hot reagent HC1 (§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 :it all. or at most only tract's. The
sulphides of Au and Pt are not soluble in HC1 . MoS is soluble in hot con-
centrated HC1 . The relative solubility of these sulphides in HC1 is used by
some chemists as the basis of a separation of As from Sb and Sn ( §69. iw, also
bottom of next note. .7).
•7. The sulphides of arsenic arc readily soluble in ammonium carbonate (§69.
.V) and are thus separate! from the sulphides of Sb and Sn. which are prac-
tically insoluble. The following' table suggests a method of analysis b
upon this property of these sulphides.
§86, 6. DIRECTIONS FOR ANALYSIS WITH NOTES.
Digest with solution of ammonium carbonate and filter.
119
Residue: SnS,, , Sb2S5 , (S) .
Dissolve in hot hydrochloric acid (5c, §70
and §71).
Solution: SnCl4 , SbCl3 .
Treat with zinc and hydrochloric acid in
Marsh's apparatus (§69, (>'«).
Deposit: Sn , (Sb) .
Dissolve by hydro-
chloric acid.
Solution: SnCL .
(Residue, Sb .)
Test by ammoniacal
silver nitrate and
by mercuric chlo-
ride (§71, 61 and ;).
Gas: SbH3 .
(Test the spots,
§69, O'c, 1.)
Receive the gas in
solution of silver
nitrate. Dissolve the
precipitate(SbAg;i )
(§70, 6/), and test
by H,S (§87 and
Solution:
(NH4)3AsS4 + (NH4)3AsO,
and
(NH4)4As2S; + (NH4)4As20B .
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: H3As04 .
Apply Marsh's Test, as directed :n
§69. f/cr, testing the spots (§69.
(i'r) ; receiving the gas in solu-
tion of silver nitrate, and test-
ing the resulting solution (§87).
Examine the original solution, as
indicated in §88, 1.
The plan above given may be varied by separating antimony and tin by ammo-
nium' carbonate in fully oxidized solution, as follows: The Sb2Sr, and SnS, are
dissolved by nitrohydrochloric acid, to obtain the antimony as pyroantimonic
acid. The solution is then treated with excess of ammonium carbonate, in a
vessel wide enough to allow the carbonic acid to escape without waste of the
solution.
The soluble diammoniuir dihydrogen pyroantimonate, (NH4)2H2Sb2O7 , is
formed. Meanwhile the SnCl, is fully precipitated as H..SnO:, (§71, Ca), 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 aiiti-
monous chloride, SbCl3 . Excess of ammonium carbonate does not redissolve
the Sb203 which it precipitates from SbCl3 , as stated in §70, <>//.
The above plan may also be varied as follows: After removal of the arsenic
sulphide with (NH4)2CO3 , the residue is dissolved in strong HC1 , not using
KC1O3 or HNO3 . The solution consists of SnCl4 and SbCl:i . 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
Hg€l, (Pieszczek, Arch. Pharm., 1891, 229, GG7).
6. The sulphides of As, Sb and Sn are all decomposed by concentrated nitric
acid, which furnishes a basis of an excellent separation of the arsenic from the
antimony and tin (Vaughan, American Chemist, 1875, 6, 41). The sulphides
reprecipitated from the (NH4)2SX solution by HC1 are well washed, transferred
to an evaporating dish, heated with concentrated HNO3 until brown fumes are
no longer evolved, and then evaporated to dryness, using sufficient heat to
expel the HNOS and the H2SO4 formed by the action of the HN03 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 H3As04 .
and the antimony and tin remaining as residue of Sb2O5 and SnO2 . The pres-
ence of arsenic may be confirmed by the reactions with AgN03 (§69, 67), CuSO<
(§69, GA") by the Marsh test (§69, 6'er), or by precipitation with magnesia mix-
120 DIRECTIONS FOR ANALYSIS WITH NOTES. §86, 7.
ture (§69, 61). A portion of the residue may be tested in the Marsh apparatus
for the Sb (§70, (\j), another portion may be reduced and dissolved in an open
flish with Zn and HC1 (not allowable if As be present, §71, 10), and the result-
ing- SnCl. identified by the reaction with Hg€L (§71, tii).
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 HC1 with the aid of KC1O;, .
8. Instead of chlorine (HC1 + KC1O3), nitrohydrochloric acid may be em
ployed, but it is liable to caut>e the formation of a white precipitate of Sb,Or,
and SnO, .
0. The chlorine should all be removed, as the metals cannot be reduced by
the Zn and H..SO4 in the Marsh apparatus in the presence of powerful oxidizing
agents as Cl . This would also require evaporation to expel the HNO3 , if
nitrohydrochloric acid were used to effect solution.
10. Hydrogen peroxide. HO , 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. Pltarni., 1886, 224, 772).
§87. Manipulation.— The solution of tho metals of the tiif group is
then ready to he transferred to the Marsh apparatus (the directions for
the use of the Marsh apparatus are given under arsenic (§69, (i'rj), 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>AgN03 + 3H20 = H3AsO, + f,Ag + r,HNOs
SbH3 + :AgNO3 = SbAg3 + .sHNO.,
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, (i'r) : 2SbH, == ^Sb + 3H, . The
ignited gas is brought in contact with a cold porcelain surface for the
production of the arsenic and antimony spots (§69, (57;). Failure to obtain
mirror, spots, or a black precipitate in the AgNO.j is proof of the absence
of both arsenic and antimony. The black precipitate obtained in tho
silver nitrate solution is separated by filtration, washed and reserved to be
tested for antimony. The filtrate is treated with HC1 , or a metallic
chloride, as CaCL, or NaCl , to remove the excess of silver and, after evapor-
ation to a small volume, is precipitated with H2S . A lemon-yellow pre-
cipitate indicates arsenic. The black precipitate from the silver nitrate
solution is dissolved in hot reagent HC1 : SbAg., -f 6HC1 = - SbCl, -+-
3AgCl . The excess of acid is removed by evaporation, a little water is
added (§70, 5d and §59, 5c) and the AgCl removed by filtration. The
filtrate, is divided into two portions. To one portion H2S is added; an
orange precipitate indicates antimony. The H2S may give a black precipi-
tate of Ag,S from the AgCl held in solution by the HC1 . If this be the
§89,.j. DIRECTIONS FOP ANALYSIS WITH NOTES. 121
case, to the other portion one or two drops of KI arc added and the
solution filtered. This nitrate 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, 6'c 5). The spots should be tested with NaCIO 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 by
the Zn and H,SO, 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 Asa03 , 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 containing1 soda lime or through a solution of
KOH (distinction and separation from antimony).
4^ Arsenites and arsenafes are distinguished from each other by the following
reactions: (a) Arsenous acid solution acidulated with HC1 is precipitated in the
cold instantly by H S: arsenic acid under similar conditions is precipitated
exceedingly slowly (§69, 6e). (1>) Neutral sohitions of arsenites give a yellow
precipitate with AgN03; neutral solutions of arsenates give a brick-red pre-
cipitate. Both precipitates are soluble in acids or in ammonium hydroxide
(§59, 6$). (c) Magnesia mixture precipitates arsenic acid as white magnesium
ammonium arsenate, MgNH4AsO4; no precipitate with arsenous acid (§189, Gy).
(d) HI gives free iodine with arsenic acid; not with arsenous acid (§69, 6f).
(e) 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. — Antimony. — 1. If antimony be present in considerable amount,
it (in the form of the sulphide) is most readily separated from arsenic by
boiling with strong HC1 (solution of the antimony sulphide, (§70. <>r)); or by
digesting with (NH4)2CO3 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 AgNOa; this is a good separation from arsenic and
tin, and after filtration it remains to dissolve ths SbAg3 in concentrated HC1
and identifj' the Sb as the orange precipitate of Sb.S,, . The formation of the
black precipitate in the AgN03 solution must not be taken as evidence of the
presence of antimony, as arsine gives a black precipitate of metallic silver with
AgN03 . 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. Sb2S3 is precipitated from solutions quite strongly acid with HC1 , i.e., in
the presence of equal parts of the concentrated acid (sp. 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, Ge). This is a convenient method
of separation. The addition of one volume of concentrated HC1 to two volumes
of the solution under examination before passing in the H,S will prevent 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 HNO3; the residue strongly fused with Na,C03 and NaOH: and the
cooled mass disintegrated with cold water, the filtrate will contain the arsenic
as sodhim arsenate, Na3AsO4 , and the tin as sodium stannate, Na_,SnO3; while
the antimony remains as n residue of sodium pvroantimonate, NaoH.,Sb,07
(§70, 7).
122 DIRECTIONS FOR J.V.-IM'X/N WITH NOTES. §89,5.
5. Stibine is evolved much more slowly than arsine in the Marsh apparatus,
and some metallic antimonv will nearly alwavs 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 K..CO3 as preliminary to the preparation of the sub-
stance for analysis, since they hinder the complete precipitation of the anti-
mony with H2S (§70, Or).
7. Antimonic compounds are reduced to the anlimonous condition by HI with
liberation of iodine (§70, Gf and 10). Chromates oxidize antimonous salts to
antimonic salts with formation of green chromic salts (§70, (>//). KMnO4 also
oxidizes antimonous salts to antimonic salts, a manganous salt being formed
in acid solution (§70, ('>//). 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 bo
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 HNO, ,
which oxidix.es the molybdenum to molybdic acid, MoO, . The residue
is dissolved in NH4OH (the zinc salt present does not interfere) and poured
into moderately concentrated nitric or hydrochloric acid (§75, Gd footnote).
This solution is tested for molybdenum by Na.,HP04 . The original solu-
tion should also be examined for the presence of molybdenum as molybdic
acid or molybdate ($75, He/).
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 HC1 . Sb . Au . and Pt are insoluble (§70, 5a). The
Sn passes into solution as SnCL and gives a gray or white precipitate with
HgCL , depending on amount of the latter present (§71, Gj):
SnCL + HgCL = SnCl, + Hg
SnCL + 2HgCL = 2HgCl + SnCl,
The presence of Sn" should always be confirmed by its action in fixed
alkali solution upon an ammoniacal solution of AgNOo , 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 ((>/<, §§73 and 74). Both metals are precipitated. They are then
dissolved in nitre-hydrochloric acid and evaporated to dryness 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 FeS04 (§73,0/0- by boiling with
oxalic acid (§73, Gb), or by treating with a mixture of SnCl, and SnCl4
(Cassius' purple) (§73, 6*7).
If a portion of the original solution, free from HNO;. , be boiled with
§94,5. DIRECTIONS FOR ANALYSIS WITH NOTES. 123
oxalic acid the gold is completely precipitated as the metal, separation
from the platinum which is not precipitated (§74, 66).
§91. — Notes. — Molybdenum. — 1. In the regular course of analysis, molyb-
'lenum remains in the flask of the Marsh apparatus as a dark colored solution,
the Zn and HoS04 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 phosphomolybdate, 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 HC1 to prevent its pre-
cipitation by H.jS than arsenic or antimony (§89, 3).
2. The yellow ammonium sulphide (NH4)^SX must be used to effect solution
if tin (Sn") be present, SnS being practically insoluble in the normal am-
monium sulphide (§71, 5c).
3. Tin in the stannotis condition, dissolved in the fixed alkalis (stannites),
readily precipitates metallic silver black from solutions of silver salts. An
arsenite (hot) or an antimonite 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 (§7O, (if). 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 HC1 . forms SnCL , which may be detected by the reduction
of HgCl, to HgCl or Hg° (§71, (>/), and by the action in fixed alkali solution
upon the strong ammoniacal solution of silver oxide (§71, 6i).
& 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 as the metal as long as zinc is present (§135, 10).
(>. 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 HC1 .
§93. Gold. — 1. Gold will usually be met with in combination with other metals
as alloys, and is separated from most other metals by its insohibility in all
acids except nitrohydrochloric acid.
2. If more than 23 per cent of gold be present in an alloy, as with silver,
the other metal is not removed by nitric acid (§73, :V/). 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 FeSO4 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 Marsh apparatus, insoluble in HC1 and may be identi-
fied by the reactions for Au° .
•7. The reactions of gold chloride with the chlorides of tin forming Cassius'
purple (§73, (if/) is one of the most characteristic tests for gold.
$94. Platinum. — /. Notes / to } under gold apply equally well for platinum,
except that it is necessary to boil with FeSO, to insure complete precipitation
of the platinum.
2. Oxalic acid is the best reagent for the separation of gold from platinum
(§73. 6fr).
.<?. The most important problems in the analysis of platinum consist in its
separation from the other metals of the platinum ores (§74, 3).
124
TABLE FOR ANALYSIS OF THE COI'I'KK
$95.
to
03
OQ
*
w
£
€
Q
M
09
fc
,
^
o
09
w
03
6
fc
P
n
09
5
. O
09
1 «
O
fc
•o £
Q
„
0
09
be
"» TJ
W
W
^ s
^
1
-•g
t)
0
-
O
id
^
9
03
T3 iJ
3 «
_o «
x D
w
CH
5
d
5 -;5
CS .,
CH
0
O
a
S
Ammonj
• pVecipi
loderatel
H
^ c
fa
^
*- _(j-
O
-^ C
' ft)
5 ^
0 5
CO
-
.2
^ ^
S s
60S ~
"O a: x
l< i/
•5 T3 bt
*>S
_ ,s
ttj W fcn
S— i ID
ft
03 JS
• . .
fe S -r
c^x *-
lii
>• = =
O 55»i
^ I u'Z
W ^-3-
- u &o
t'_
a*i'£
c o «s
*~*
o
d
iSfe
•5 - "»
^
S tc^
0
u o)'o
-
•5^2
• /— ,
2 o^
o"
»
I-0?.^
jvW c
m
;- -^- .—
«- Q *-
*
•s . *
f*
r x ti
•^
8-iU
§
v ^ .^
C S y
•031;
- v- i.
^3
*"S ^
ft
' 'Z. ~~
d
o
o" z t
09 -"H.-+*
M * ?T **
.d
"•£ p ^
§
. "-3
; o
o05
fi
.
HI i
'o S — S-?-2 ^T;
? 2 6 S w * .2 " = •= - c
<1> fl^ ., !""? ™ ^ _ ''^'v r*. .— ^J i
S-d1°
TABLE FOR ANALYSIS OF THE COPPER GROUP.
•e pt»
o
oo o-
"E | _, "o » •* ^
O ^ ^ ^ *H
cc *y -H cd c*>
be g ° i- .-S §
P iv -H c H" PJ 5
>FH +^ o ^~^"in •/>
^+2
•p .
|S 2
o
000
oo o'
. w*-. >r
OOO'CDO
~l-05iH-
Q «y5 w. v/.
r-
0
O
,^o
03 iH
IS"
s
o ""
00 "*=
0) «» ^
^ lO 05 05
coo'co
r; «/-. sew
DIRECTIONS FOR ANALYSIS WITH NOTES 8»fl
yvv,
§96. Manipulation. — The well washed residue after digesting the pre-
cipitated sulphides of the second group (the Tin and Copper Group) in
(NH4)2S, may contain any of the metals of the Copper Group, and in
addition frequently contains sulphur, formed by the action of the H2S
upon oxidizing agents: 4FeCl3 -f 2H2S =: 4FeCL, -f- 4HC1 -f S2 . 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,S3 + 1GHN03 = 4Bi(NOs)3 + 4NO + 8H,O + 3Sa
GCdS + 1GHNO, = (iCd(NO,)a + 4NO + 8H2O + 3Sa
Mercuric sulphide is unattacked (§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: 2 HgS -f 2C1., — SHgCL -J- 3., .
This solution is boiled to expel all chlorine and the presence of mercury
determined by reduction to HgCl or Hg° by means of SnCl2 (§58,0.7):
HgCl. + SnCL, == HE + SnCl4 , 2HgCl, -f SnCl, = 2HgCl + SnCl4 ; or
by the deposition* of a mercury film on a strip of bright copper wire
($50, 10): HgCl, + Cn — Hg -f CuCL . Confirm further by bringing in
contact with iodine in a covered dish: Hg + I2 r= Hgl, (.Jannaosdi, Z.
"tiorf/., ISJMi, 12, 143). The mercury may also be detected by using
NH4OH and KI as the reverse of the Nessler's test (§207, (!/,-) (delicate
1 to 31,000) (Klein, Arch. Pharm., 1889, 227, T3).
§97. Sntrx. — /. The concentration of HNO (1-2) is necessary for the solution
of the sulphides of Pb , Bi , Cu and Cd , and ina.v also dissolve 1 races of HgS .
However, the concentrated HNO (*/j. (jr., 1.42) dissolves scarcely more than
traces of HgS (858. (if). ^ Loiifj-continued boiling of HgS with concentrated
HNO:, changes a portion'of the HgS to Hg(NO )..HgS. a white precipitate,
insoluble in HNO .
£. In the use of nitrohydrochloric^'id to dissolve the HgS . Hie HC1 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 HC1 jj-ives about sufficient HC1 to decompose all the HNO, ,
hence in this reaction a little more than that proportion of HC1 should be
used.
.?. A small amount of black residue left after boiling1 the sulphides with
HNO, may consist entirely of sulphur, which can best be determined by
burning the residue on a platinum foil and noting1 the appearance of t In-
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 oxidi/e
* If preferred the precipitate on the filter may be washed with the boiling hot nitric acid of
the above mentioned strength, pouring the same acid back upon the precipitate, reheating each
time, until no further action takes place.
§99,5. DIRECTIONS FOR ANALYSIS WITH \OTE8. 127
a trace at least of sulphur to H.SO, (§256, GB, 2), which will form PbSO4 if any
lead be present:
S2 + 4HN03 = 2H.SO, + 4NO
3PbS + 8HN03 = 3PbS04 + 4H,0
If the boiling be not continued too persistently, the amount of PbSO4 formed
is soluble in the HNO3 present (§57, 5c), and does not at all remain behind
with the HgS .
5. If the Sb and Sn are not removed, through an insufficiency of (NH4)2SX
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 Sb.05 and SnCK:
<>Sb,S3 + 20HNO-3 = <;Sb,0, + OS, + 20NO + 10H2O '
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 restdts (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,CO3 + K,CO:!) (§58, 7). and then
adding a few drops of HNO3 (concentrated) and a small crystal of KI . Upon
warming the iodine sublimes and combines with the sublimate of Hg , forming
the scarlet red Hgl, . As and Sb both give colored compounds with iodine, de-
composed by HN03 (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 H.,S04 and the mixture evaporated on a sand bath or over
the naked flame in a casserole or evaporating dish until the fumes of
HoS04 are given off:
Pb(N03), + H,SO, — PbSO4 + 2HN03
Cu(NOt)., + H,S04 = CuS04 + 2HN03
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 precipitate of PbSO, will collect
in the center of the beaker, and its presence even in very miall amounts
may be observed. The filtrate from the PbS04 should be decanted througli
a wet filter, and the PbS04 in the beaker may be further identified by its
transference into the yellow chromate with K2Cr04 or into the yellow iodide
with KI (57, G/ and fc).
§99. 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 H.SO, for
lead, and if that metal be not present, the above step may be omitted with
the remainder of the solution and the student may proceed at once to look
for Bi , Cu and Cd . If, however, lead is present, the whole of the solution
must be treated with H..SO, .
2. The nitric acid should be removed by the evaporation, as PbSO4 is quite
appreciably soluble in HNO3 (§57, 5c).
3. The H2SO4 should be present in some excess, as PbSO4 is less soluble in
dilute H2SO4 than in pure water (§57, 5c).
4. Alcohol should be present, as it greatly decreases the solubility of PbSO,
in water or in dilute H2S04 (§57, 5c, 6e).
5. Too much alcohol must not be added, as sulphates of the other metals
present 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.
128 DIRECTIONS FOR AX.lLYxlH WITH \OTE8. §99,6'.
6. If the (NH4)2SX had not been well removed by washing, ammonium sul-
phate would be present at this point, greatly increasing the solubility of
PbS04 (§57, 5c).
§100. Manipulation.— The filtrate from the PbS04 should be boiled Lo
expel the alcohol (or if Pb be absent evaporate the nitric acid solution of
division B) and then carefully neutralized with NH4OH . An excess of
NH4OH should be added to dissolve the precipitates of Cu(OH)0 and
Cd(OH)2 , leaving the Bi(OH), as a white precipitate. The solution should
be filtered, the precipitate thoroughly washed, and then treated upon the
filter with a hot solution of potassium stannite, K ,SnO, . A black piv-
eipitate is evidence of the presence of Bi (§76, Gg).
§101. Notes. — 1. If the precipitate of the sulphides of the second group was
not well washed, the hydroxides of the metals of the iron group (Al , Cr and
Fe) may be present at this point. The precipitate of A1(OH)3 would be white.
but would not give a black precipitate with K SnO .
2. If an insufficient quantity of (NH4),SX was used, Sb and Sn would be
present and give a white precipitate with the NH.OH .
3. If the lead had not been removed it would appear as a white precipitate
with the NH4OH , and would give a brownish-black precipitate with the hot
K SnO (§57, fir;). The presence of a permanent white precipitate with NH4OH
must never be taken as final evidence of the presence of Bi .
}. As a confirmatory test for the presence of Bi , a portion of the white
precipitate with NH.OH should be dissolved in HC1 and the solution evapo-
rated nearly to dry ness to remove the excess of HC1 . Now upon adding
water, a white precipitate of BiOCl , bismuth oxychloride, will be obtained if
Bi is present (§76, 3d).
§102. Manipulation. — If the annnoniacal filtrate from the Bi(OH).t is of
a blue color, that is sufficient evidence of the presence of Cu unless nickel
was precipitated in the second group. In absence of a blue color a portion
of the solution should be acidulated with acetic acid and then to this solu-
tion a few drops of potassium ferroeyanide, K4Fe(CN)0 , should be added.
A brick-red precipitate is evidence of copper. Or to the acidulated solu-
tion a bright nail or piece of iron wire may be added, obtaining a film of
metallic copper. If sufficient copper be present to give a blue color to
the solution, before testing for cadmium a solution of KCN should !>»•
added until the blue color disappears. Then the addition of H2S will
give a yellow precipitate for cadmium.
§103. Notes. — 1. The precipitate of the brick-red Cu,Fe(CN)0 is a much more
delicate test for copper than the blue color to the ammoniacal solution (§77.
Kb). Cd gives a white precipitate, insoluble in the acid.
2. The student should not forget that in the regular course of analysis a
trace of copper may be lost by the solubility of the sulphide in (NH4)._,SX . If
mercury has been shown to be absent, the sulphides of the tin group (second
group, division A) should be dissolved by the addition of a fixed alkali sulphide
(§71, 6e), K2S or Na.S . which does not dissolve CuS . In case mercury br
present, the presence or absence of small amounts of copper must be deter-
mined by the usual reactions for copper upon the original solution, having
due regard for the possible interference of metals which the analysis has
shown to be present.
3. Potassium cyanide, KCN , in excess changes cupric salts to the soluble
;;104, 5. RUTHEXIUM. 120
double salt of cuprous cyanide and potassium cyanide, K:[Cu(CN)4 , which is
colorless and not precipitated by sulphides. With cadmium salts the soluble
double cyanide, K2Cd(CN)4 , is formed, which is decomposed by sulphides
forming- CdS , yellow.
-J. If preferred, the sulphides of Cu and Cd may be precipitated from the
ammoniacal solution by H2S and then the black CuS dissolved with KCN ,
leaving1 a yellow precipitate of CdS .
•j. Copper and cadmium may be separated from each other by reduction of
the copper (from the ammoniacal solution acidulated with HC1) with SnCl,
(§77, 10): 2CuCL + SnCL = 2CuCl + SnCl4 , and its precipitation with milk
of sulphur (§77, 6e), forming- Cu,S , removal of the tin with NH4OH and the
precipitation of the cadmium with H,S .
6. From the solutions of copper and cadmium acichilated with HC1 , a hot
solution of Na2S2O3 precipitates the copper as Cu2S (§77, 6e), while the cadmium
remains in solution. From this solution the cadmium is detected as the sul-
phide by neutralization with NH4OH and precipitation with H^S or (NH4)..S .
7. The ammoniacal solution of Cu and Cd may be precipitated with H2S ,
and the resulting- sulphides, after filtering- and washing-, boiled with, hot dilute
H2SO4 (one of acid to five of water). In this solution the CuS (§77, 5c) is
unattacked while the CdS is dissolved. The filtrate upon dilution with water
gives the yellow CdS with H,S or (NH4)2S (§78, 6e).
RARER METALS OF THE TIN AND COPPER GROUP.
(Second Group.)
Ruthenium. Rhodium, Palladium, Iridium, Osmium, Tungsten, Vanadium,
Germanium, Tellurium, Selenium.
§104. Ruthenium. Ru = 101.7 . Valence two to eight.
1. Properties.— Specific gravity, 11.0 to 11.4 (Deviile and Debray, C. r., 1876, 83,
•926). Next to osmium it is the most difficultly fusible of all the platinum
metals. A black powder or a grayish-white crystalline brittle metal.
2. Occurrence. — In small quantities in platinum ores.
3. Preparation. — Ignite the Pt residues in a st-.am of chlorine in presence of
NaCl . Dissolve the fused mass in H,O , add KNO2 , neutralize with Na3CO3 ,
evaporate to dryness and extract the double nitrites with absolute alcohol
(separation from rhodium). Add water to the solution, distill of the alcohol,
add HC1 and obtain a red solution of potassium ruthenium chloride. This is
changed to the double ammonium salt and then precipitated with HgCL ,
which upon recrystallization and ignition gives pure Ru (Gibbs, Am. &., 1802,
(2), 34, 349 and 355).
4. Oxides and Hydroxides.— The hydroxides, Ru(OH)2 , Ru(OH)3 , and
Ru(OH)4 , are precipitated from the respective chlorides by KOH . They are
dark brown to black. Perruthenic anhydride or acid, Raid), , is a golden
yellow crystalline powder, volatile even at ordinary temperatures. It has a
peculiar odor, somewhat like ozone, is sparingly soluble in water, melts at 50°
and boils at a little over 100° (Deviile and Debray, B., 1875, 8, 339). It is pre-
pared by heating- K2BuCl5 with KOH into which a current of chlorine is
passed or by distillation of a Ru salt with KC1O3 and HC1 . The vapor is
yellow and is strongly irritating to the membrane of the throat.
5. Solubilities. — Ru is soluble with difficiilty in nitrohydrochloric acid, in-
soluble by fusion with KHSO , , but is soluble by fusion with KOH , especially
in presence of KNO;, . Soluble in chlorine, forming- a mixture of RuCl2 ,
RuCl3 , and RuCl4 . The double nitrites are soluble in water and alcohol (sepa-
ration from rhodium).
130 RHODIUM. §104, 6.
f>. Beactions. — The alkalis precipitate from ruthenic chloride the dark yellow
hydroxide, soluble in acids, insoluble in the fixed alkalis, soluble in NH.OH
with a greenish-brown color. H,S precipitates slowly tin- black sulphide
(formed at once by (NHJ.S), the solution becoming blue. The sulphide is
insoluble in alkali sulphides. KI gives with hot solutions a black precipitate
of ruthenic iodide. KCNS forms, after some time in the cold, a red coloration,
which upon heating assumes a beautiful violet color (characteristic). The
double nitrites are soluble, and if to the solution (NH4)...S be added. ;i char-
actetristic crimson red liquid is obtained. I'pon standing the solution becomes
brown, or a brown precipitate is caused by excess of the (NH4)1.S.
7. Ignition. — If RuO be heated to a dull-red heat the violet-blue dioxide is
formed (I)ebray and .loly, ('. r., 1SSS, 106, 32S).
8. Detection. -By oxidation and distillation as RuO .
9. Estimation. Reduced to the metal and weighed as such.
10. Oxidation. RuO, heated with HC1 forms RuCl , evolving chlorine.
RU8 — x is changed to RuO, by distilling with KC1O and HC1 . Zn reduces
Bu solutions to the metal, with an indigo-blue color during transition from
v to Bun .
$105. Rhodium. Rh = 103.0. Valence two, three and four.
1. Properties. -N/«r///r »/r«n'///, 1:2.1 (Devi lie and Debray, C. r., 1874, 78, 1782).
Melts more difficultly than Pt . It is a white metal, nearly as ductile and
malleable as Ag . The metal precipitated by alcohol or formic acid appears as
a black spongy mass (\Vilm, It., 1HS1, 14, <>2'.)).
2. Occurrence. — Found in platinum ores.
3. Preparation.— Fusion of the Pt residues with Pb , digestion with HNO:,
and then Cl , converting the Bh. into the chloride, from which solution it is
precipitated as the double ammonium chloride by fractional precipitation.
See Gibbs (./. /»/•., lS(i:,. 94. 10) and \Vilin (/*.. 1SS3, 16, 3033).
4. Oxides and Hydroxides. Rh( OH i is precipitated from a solution of
sodium rhodium chloride by an excrss of KOH . It is a black gelatinous pre-
cipitate, forming the oxide upon ignition. Uhodium fused with KOH and KNO,
gives RhO , a brown powder, insoluble in acids or alkalis.
r>. Solubilities. — The pure iretal or the alloy with Au or Ag is almost in-
soluble in acids; alloyed with Bi , Pb . Cu or Pt , it is soluble in HNO3 (Devillc
and Debray. /.<•.). Attacked by chlorine the most easily of all the Pt metals.
The precipitated metal, a gray powder, is soluble in HC1 in presence of air to
a cherry-red color.
G. Beactions. Alkali hydroxides and carbonates precipitate solutions of Bh
salts as RhiOHi . yellow, insoluble in acids, soluble in excess of NH.OH .
forming a rhodium ammonium base, precipitated by HC1 as a bright yellow
crystalline salt. chloro-p:irpnreo-rhodmm chloride, Bh(NH3);,Cl3 . Alkali
nitrites precipitate alcoholic solutions of rhodium chloride as alkali-rhodium
nitrite ((iibbs. Am. X.. 1HG2, (2), 34, 341) (separation from ruthenium). From
a hot solution of Bh salt, H...S precipitates the sulphide, insoluble in the alkali
sulphides; the sulphide precipitated from the cold solution is soluble in alkali
sulphides. KI precipitates from hot solutions a black-brown rhodium iodide.
7. Ignition. — When the metal or its compounds are repeatedly fused with
HPO; or KHSO4 , the corresponding Bh salts are formed. The mass fused
with KHSO, is soluble in water to a yellow color, turning red with HC1 .
8. Detection. — Hy ignition as given above. Also to the concentrated neutral
solution add fresh NaCIO solution. To the yellow precipitate add a small
amount of HC H O and shake till an orange-yellow solution is obtained.
After a short time the solution becomes colorless, then a gray precipitate
separates out and the solution ass\imes a skv-blue color (Demarcay, C. r., 1885,
101, 9.r>l).
9. Estimation.- It is reduced to the metal and weighed as such.
10. Oxidation. — Solutions of rhodium stilts are reduced to the metal by Zn .
All Bh compounds are reduced to the metal by heating in a current of
hydrogen.
$106, (). PALLADIUM. 131
§106. Palladium. Pd = 107.0 . Valence two and four.
1. Properties. — Specific- gravity, 11.4 (Deville and Debray, C. r., 1857, 44, 1101),
Meltiny point, about 1500° (Violle, C. r., 1878, 87, 981). It conducts electricity
about one-eighth as well as silver (Matthiessen, Poyg., 1858, 103, 428). Palla-
dium has about the color and lustre of silver. The metal when only slightly
heated assumes a rainbow tint from green to violet. Because of its general
properties, it is to be classed with the platinum metals, yet in its reaction
with acids it is markedly different. In the air at ordinary temperature it is
but slightly tarnished, but at a red heat it becomes covered with a coating of
the oxide. The finely divided metal, palladium sponge, absorbs many times its
volume of hydrogen, retaining the most of the hydrogen even at 100° . At a
high heat the hydrogen is all driven off. It is much used in gas analysis for
the separation of hydrogen from other gases (Hempel, B., 1879, 12, 636, 1006).
Also used for scale graduations of the best scientific instruments.
2. Occurrence. — It is a never-failing element in the platinum ores, native or
alloyed with Pt , Au or Ag .
3. Preparation. — The obtaining of pure palladium involves its separation from
the other platinum metals, i. e., platinum, iridium, osmium, rhodium and
ruthenium. The student is referred to the various works on metallurgy; also
to the following: Bunsen, A., 1868, 146, 265; Wilm, B., 1885, 18, 2536; and Mylius
and Forster, B., 1892, 25, 665.
4. Oxides and Hydroxides. — Palladium monoxide, PdO , is the most stable of
the oxides of Pd . It is formed by the gentle ignition of Pd(NO3)2 or the
precipitation of PdCL with Na2CO3 , forming Pd(OH)2 , and then igniting.
Palladic oxide, PdO;, , when gently ignited loses half its oxygen, becoming PdO .
5. Solubilities. — a. — Metal. — It is slowly dissolved by boiling with HC1 or
H2SO4; HNO:, dissolves it, even in the cold, forming Pd(NO3)2 . It is more
readily soluble in nitrohydrochloric acid, forming PdCl4 . It is not at all
attacked by H2S . An alcoholic solution of iodine blackens it, and when fused
with KHS04 it becomes the sulphate (distinction from platinum). &. — Oxides. —
PdO, is soluble in HC1 with evolution of Cl , forming PdCl2 . Pd(OH)2 is
readily soluble in acids forming palladous salts, c. — Salts. — Palladic chloride.
PdCl4 , the most stable of the palladia salts is decomposed by boiling with
water or by much dilution with cold water, forming PdCl2 . It forms double
chlorides with other metals, as calcium palladic chloride, CaPdCl0 , which for
the most part are stable, and soluble in water and alcohol. Potassium palladic
chloride, K.PdCl,; , is but sparingly soluble in water, insoluble in alcohol; par-
tially decomposed by both solvents.
Palladous chloride is readily soluble in Avater with a brownish-red color; with
metallic chlorides, it forms double chlorides, as potassium-palladous chloride,
K,PdCl4 , all of which are soluble in water. — Palladous iodide is insoluble in
water, alcohol or ether; insoluble in dilute hydrochloric acid or hydriorlic acid;
slightly soluble by iodides and bjr chlorides. — Palladous nitrate, Pd(NO3)2 , is
soluble in water with free nitric acid; the solution being decomposed by dilu-
tion, evaporation, or by standing, with precipitation of variable basic nitrates,. — -
Palladous sulphate, PdSO4 , dissolves in water, but decomposes in solution on
standing.
0. Reactions. — Palladous chloride is precipitated by potassium hydroxide or
sodium hydroxide; as brown basic salt or as brown palladous Jiudroa-idc,
Pd(OH), , soluble in excess of the hot reagents. Ammonium hydroxide gives
a flesh-red precipitate of palladio-diammoninm chloride, (NH3)2PdCL . The flesh-
red precipitate is soluble in excess of the ammonia, and from this solution
reprecipitated by hydrochloric acid, with a yellow color. The fixed alkali
carbonates precipitate the hydroxide; ammonhim carbonate acts like the
hydroxide. — Potassium cyanide precipitates palladous ci/anide, Pd(CN)2 , white,
soluble in excess of the reagent. Phosphates give a brown precipitate. —
Hydrosulphuric acid and sulphides precipitate the dark-brown palladons
sulphide, PdS , insoluble in the ammonium sulphides, soluble in nitrohydro-
chloric acid. Potassium iodide precipitates palladons iodide, Pdla , black, visible
in 500,000 parts of the solution, with the slight solubilities stated in 5c, an
important separation of iodine from bromine. In very dilute solutions, only a
132 I RID IV M. §106, 7.
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; 110 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 Ihe
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 H,S.
not dissolved by (NH4):SX (separation from the tin group). It is distinguished
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, Pdl, , to a dark brown
solution. KCNS does not precipitate palladium salts, not even after the addi-
tion of SO. (separation from Cu). The addition of H...SO, 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 K^.PdCl,, . Evaporate the solution of palladic
chloride with potassium chloride and nitric acid to dry ness, 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, Git).
§107. Iridium. Ir — 193.1 . Usual valence three and four.
1. Properties.— N/>m/?r (jraritu, 22.421 (Deville and Debray, ('. r., 1875, 81. 839).
Melting point, 1950° (Violle, C. 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 taking 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 by 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 osmiuin-iridiiiin 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 double 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
H;S just as the brown iridium sulphide begins to form, a complete separation
can be made by filtration. By r^-crystalli/ation th:- pure sodium double salt.
(>NaC1.2lrCl3 -f- 24H...O , is obtained, which is changed to the tetrad ammonium
double salt, (NH4)2IrCl, , by the addition of NH4C1 and oxidation with chlorine
(Wohler, /*«<//;., 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. IrO, is formed by
§108,5. ofuin'U. 133
igniting the metal in the air at a bright reel heat, hence the scaling of platinum
dishes which contain indium. The hydroxide, Ir(OH)4 , is formed by boiling a
solution of the trichloride, IrCl:, , in a fixed alkali hydroxide or carbonate.
Careful addition of KOH to IrCL, in a vessel full of liquid and closed to exclude
air gives Ir(OH), . easily oxidi/ed to Ir(OH), (Clans, ,/. pr., 184(5, 39, 104).
5. Solubilities. — Freshly precipitated iridium may be d;~«3olved in nitrohydro-
chloric acid. The ignited metal is insoluble in all acids. Its proper solvent is
chlorine. Iridium trichloride. IrClr, , is soluble in water and forms with the
alkali chlorides doiible chlorides, soluble in water, insoluble in alcohol. The
tetrachloride with sodium chloride, Na2IrCl,; , is formed when the platinum
residues mixed with HaCl 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 WH4C1 , a reddish-
brown precipitate, soluble in 20 parts of water (Vauquelin, A. Ch., 1806, 59, 150
and 225). Tlie potassium double salt is sparingly soluble in water.
6. Reactions. — Fixed alkali hydroxides or carbonates precipitate from boil-
ing solutions of iridium chloride, IrCl3 or IrCl4 , iridium lii/droxide, Ir(OH)4 ,
dark blue, insoluble in all acids except HC1 . 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 IrCl3 ,
and then precipitates the trisulphide, Ir2S3 , 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 6.
9. Estimation. — It is converted into the oxide by igniting with KNO3 and
then reduced by ignition in an atmosphere of hydrogen.
10. Oxidation. — Formic acid (from hot solution), zinc and H.SO, or HC1
reduce iridiiim compounds to the metal. SnCL , FeSO4 and H,C2O4 reduce
tetrad iridium to triad, but do not further reduce (separation from gold, §73,
figr, h and b).
§108. Osmium. Os = 191.0 . Valence two to eight.
1. Properties. — Specific gravity, 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 Pt without melting or oxidizing. In presence of air,
when heated a little above the melting point of Zn . it burns 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 resembling 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 . By repeated
additions of HNO:! and further distillation, the osmium may all be driven into
the receiver. The distillate is treated with HC1 and Hg and the amalgam
ignited in a current of hydrogen (Berzelius, Pogy., 1829, 15, 208).
4. Oxides. — Osmium, forms five different oxides, OsO , Os203 , Os02 , OsQ;1 .
OsOt . The first three are bases, the salts of which have been but little
studied; OsO3 forms salts with bases, and OsO4 acts rather as an indifferent
peroxide. Perosmic acid, OsO4 , 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. H2S 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 TUNGSTEN. £108,6.
ing nitric acid. By heating- the metal in a current of chlorine a mixture of
OsCl, and OsCl4 is formed. They are both unstable.
(j. Reactions. — Perosmic acid, 6sO4 , when boiled with alkalis, is red need to
osmates, as KJOsO4 . A solution of perosmic aeid decolors indigo, oxidizes
aleohol to aldehyde, and liberates iodine from potassium iodide. In the pres-
ence of a strong mineral acid, BLS precipitates osmium sulphide. OsS, . brown-
ish black (Clans, ./. in:, ISliO, 79, 28); insoluble in alkali hydroxides, carbonates'
or sulphides.
7. Ignition.- Osmium when heated on a piece of platinum i'oil gives an in-
tensely luminous flame of short duration. By holding the foil in the reducing-
flame and then again in the oxidi/ing flame, the luminosity may lie 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. Aleohol 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 perusmic aeid 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. — Specific yrarity, 10.129 (RoHcoe, 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 most common tungsten minerals are scheelite,
CaWO4 , and wolframite, FeWO, and MnWO4 , in variable proportions. It never
occurs native.
3. Preparation. — Hy reduction of WO, in H at a red heat (Zettnow, Poy.i/.,
I860, 111, 1(>); by ignition of WO, and Na under NaCl . Tungstic acid of
commerce is prepared by igniting for several hours: 100 parts Na.CO. , ignited;
150 parts finely ground wolframite; and 13 parts NaNO . The cooled mass is
exhausted with water and the filtrate poured into hot, moderately concentrated
HC1 (Franz, J. /)/-., 1871, (2), 4, 238).
4. Oxides. — WO., is obtained as a brown powder by decomposing WC14 with
water (Koscoe, ./. f.). 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 WO. and
W03.
5. Solubilities. — The metal is scarcely at all attacked by HC1 or H,,S04 , slowly
by HNO ; or nitrohydrochloric, slowly soluble in alkalis. The halogens com-
bine directly upon heating. WO2 is readily soluble on healing with HC1 and
H SO, to a red color. It is also soluble in KOH with red color, evolving
hydrogen. Both the aeid and alkaline solutions deposit the blue oxide on
standing (von der 1'fordten, A., 1SS4, 222, 158). W03 is' insoluble in water or
acids, not even soluble in hot concentrated H.SO, . Soluble in KOH, K.CO,
and NH,OH . In an atmosphere of CO.. it reacts with the chlorides of Ca ,
Mg, Co, Ni and Fe (not with those of Pb , Ag , K and Na), <?.//., MC12 -f-
L'W03 = MWO4 + WO,C1, . Heated with chlorine, WO,C1, is formed, and also
WC1, . decomposed by water. S , H,,S or HgS form WS3 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
tinetioii from MoO3), soluble in NH,OH . 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 aeid causes the
precipitation of tungstic acid. Fusion of WO3 with KHSO, gives a compound
of potassium tungstate and tungstic acid, not readily soluble in water but very
readily soluble in (NH4)SOO, (distinction from silica, §249, 5).
(). Reactions. — Solutions of salts of Ba , Ca , Pb , Ag and Hg produce white
precipitates with solutions of alkali tungstates. H2S precipitates WS;
from acid solutions, the sulphide dissolving readily in (NHJ.S , forming a
thiotungstate (NHJ.WS, . The tungstates, like the molybdates, form complex
compounds with phosphoric acid, /. r., phosphomolybdates and phosphotung-
states, which react very similarly with ammonium salts and with organic bases
(§75, ti</). K4Fe(CN),; 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 FeSO, . Heated on charcoal in presence of Na..CO3 with
the blow-pipe, using the reducing flame, the metal is obtained.
8. Detection. — If a tungstate be fused with Na_,CO, . the mass warmed writh
water and the water then absorbed with strips of filter paper, the tungsten
may be detected by moistening the strip with HC1 and warming, obtaining the
yellow color of WO3; and the blue color of a lower oxide by moistening with
SnCL and warming. (NH4)jS does not color the paper, even after adding HC1 ,
but on warming a blue or green color is obtained.
9. Estimation. — It is converted into WO3 and weighed as such after ignition.
10. Oxidation. — WO3 gives with SnCl2 , or Zn in presence of HC1 or H2SO, ,
a beautiful blue color, due to the formation of oxides between W0; and WO3 ,
blue oxides of tungsten (delicate and characteristic).
§110. Vanadium. V — 51.4 . Valence two to five.
1. Properties. — Specific yravitj/, 5.5. A grayish non-magnetic powder; slowly
oxidized in the air, rapidly on ignition with formation of V2O5 . 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, <-.</., vanadinite, :;Pb3V2Or, + PbCL: volborthite, (Cu.Ca)3V2O8;
mottramite, (Cu.Pb)5V2Olo.2H,0; 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 H...SO, . The vanadic acid is neutralized with NH4OH and
precipitated with NHjCl , in which it is insoluble. This upon ignition gives
V,05 pure (Wohler, A., 1851, 78, 125). The metal is prepared from the dichlo-
ride, VC12 , by long-continued ignition in a current of hydrogen.
4. Oxides. — Vanadium forms four oxides: VO , gray; V.,O3 , black; V02 , dark
blue; and V..O.-, , dark red to orange red.
5. Solubilities. — Vanadium is not attacked by dilute HC1 or H2SO4; concen-
trated H.,SO4 gives a greenishTyellow solution; HN03 a blue solution. VO dis-
solves in acids to a blue solution with evolution of hydrogen. V2O3 dissolves,
in dilute HC1 to a dark greenish-black solution. Chlorine forms with V«OS »
VOC13 and V2O5 . VO. dissolves in acids to a blue solution, from which solu-
tions Na2CO3 gives a precipitate of4V1!O2(OH)4 + 5H2O , grayish-white mass,
losing 4H2O at 100° and turning black, soluble in acids and alkalis. V.O-,
exists in several modifications with different solubilities in water, the red
modification being soluble in 125 parts of water at 20° (Ditte, f. 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 NH4C1 .
136 GEKH \\riM. £110, i».
6. Reactions. — Solutions of vnnaclic acid produce brown precipitates with
alkalis, soluble in excess to a yellowish-brown color. Potassium ferrocyanide
gives a green precipitate, insoluble in acids. Tannic acid givts a blue-black
solution, which is said to make a desirable ink. Ammonium sulphide precipi-
tates V2S5 . 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,S; .
If to a solution of a vanadate. neutral or alkaline, solid NH4C1 be added, the
vanadium is completely precipitated as NH,VO3 , ammonium mctavaiiadate,
crystalline, colorless, insoluble in NH,C1 solution; upon ignition in air or oxy-
gen, ]>ure vanadic oxide, V.jO.-, , is obtained.
7. Ignition. — Borax {fives 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 vanadinm ignited in air or
oxygen give V.O5 .
8. Detection. — Vanadium will almost always be found as a vanadate (2) and
is detected by the reactions used in its purification ('!)• also by the reactions
with reducing agents, forming the colored lower oxidized compounds (10).
9. Estimation. — (/) It is precipitated as basic lead vanadate and dried at
100°. (£) It is precipitated as ammonium vanadate, NH,VO3 , in strong
NH4C1 solution, ignited to the oxide V205 , and weighed.
10. Oxidation. — Zn , in solutions of vanadates with dilute H SO, , reduces the
vanadium to the tetrad, a green to bine solution, then greenish-blue to green,
the triad, and finally to lavender blue, the dyad. H2S 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. — Specific (/rarity-, 5.469 at 20.4°; meltiny point, 900° (Winkler,
./. ;>r., 18S6. (2). 34, 177). A gray-white crystalline metal. Fused under borax
it gives a grayish-white regulns with a metallic lustre. It is stable in the air,
volatilized at a high heat (Meyer, H., 1887. 20. 497), and is easily pulverized.
It burns in oxygen to form germanic oxide, GeO;. .
2. Occurrence. — It is found in small quantity in argyrodite, a sulphide of
silver and germanium. .'!Ag,S + GeS , a silver ore from Freiburg, Saxony.
It is also found in euxenite from Sweden (Kriiss, C. C., 1888, 75).
3. Preparation. — It is formed by reduction of the oxide, GeO , with H , C
or Mg (Winkler, II., 1891, 24, 891); also by reduction of the sulphide in H.
4. Oxides. — It forms two oxides, GeO and GeO, . To prepare pure GeO2 , the
mineral argyrodite is pulverized and intimately mixed with equal weights of
Na,CO, and S and heated to a good full ignition. The mass must be added
carefully to prevent foaming. The fused mass is exhausted with H.O , the
germanium going into solution as a thiosalt. \Yith a decided excess of HLSO4 ,
the sulphide is completely precipitated. The precipitate is now dissolved in
KOH , the sulphides of Ag , Cu and Pb remaining undissolved. lly adding to
the KOH solution H.SO, not quite to neutrali/.ation. the As and Sb sulphides
are precipitated on boiling, while the GeS remains in solution with some
As,S3; H2S is carefully added to the solution until the As_,S3 is all precipitated,
then the nitrate is made strongly acid with H,SO, , and the solution evaporated
till SO, fumes escape. The mass is dissolved in hot water, and upon cooling
GeO2 crystallizes out (Winkler, I.e.).
5. Solubilities.— Germanium is insoluble in HC1 , soluble in nitrohydrochloric
acid as GeCl4 , and oxidized with HN03 to Ge02 . Hot concentrated H2SO4
evolves SO, and forms Ge(SO4), . Insoluble in KOH solution but dissolves
with incandescence in fused KOH. It unites directly with Cl , Br and I
(Winkler, /. e.). Germanic oxide, GeO2 , 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°;
$112, 5. TELLURIUM. 137
it is decomposed by water. If a solution of the oxide in excess of HC1 be
evaporated to dryness the Ge is all volatilized. GeS2 is soluble in 222 parts
water, in alkali sulphides and hydroxides; insoluble in HC1 or H.SO4 , which
precipitate it from its solutions: soluble in nitrohydrochloric acid with separa-
1inn of sulphur. Nitric oxide changes it to GeO2 with separation of sulphur.
(5. Reactions. — Germanium salts give almost no characteristic reactions witli
the various reagents. H2S precipitates germanic sulphide, GeS, , white, from
solutions of the salts quite strongly acid. The sulphide is soluble in ammonium
sulphide, forming a thio salt, thus placing Ge in division A of the second group.
7. Ignition. — Heated before the blow-pipe in the reducing flame without an
alkaline flux the metal is formed, and at the same time a white coating of
the oxide. It forms a colorless bead with borax.
8. Detection. — In the mineral, argyrodite, by heating in an atmosphere of
H2S or illuminating1 gas, an orange-yellow sublimate is obtained, which may be
examined under the microscope and in the wet way (Haushofer, f. C., 1888,
8(57).
9. Estimation. — It is converted into the sulphide, GeS2 , and then heated
with HNO3 and weighed as GeO, .
10. Oxidation. — Zn in acid solutions of Ge salts precipitates the metal as a
dark brown slime. If GeS2 is heated in a current of H , GeS is at first formed
with H.S, finally Ge°.
§112. Tellurium. Te = 127.5 ? Valence two, four and possibly six.
1. Properties.— Specific gravity, 6.2445 (Berzelius, Fogg., 1834, 32, 1 and 577).
Melting point, 452° (Carnelley and Williams, ./. C., 1880, 37, 125). Te is crystal-
line, silver white, brittle, stable in the air and in boiling water; heated in the
air, it burns with a greenish flame. In its general properties and reactions it
stands closely related to S and Se (2).
2. Occurrence. — In few places and in small quantities in Germany, Mexico,
Bolivia, United States and .Japan. Some of the minerals are: tellurite, TeO.,;
tetradymite, 2Bl,Te3.Bi,S3; ferrotellurite, FeTeO4 , etc. It also occurs native.
:?. Preparation. — (/) Fusion with alkali carbonate and C , which converts it
into a telluride, as Na.Te; then solution in (air free) water, the air being
excluded as much as possible, and the filtrate precipitated by passing air
through the solution. The Te is precipitated as a gray metallic powder, con-
taining what Se may have been present. (2) Conversion into TeCl, by distilla-
tion in a current of chlorine, decomposition of the chloride with water to
H.,TeO3 and precipitation of the Te with KHSO., . (3) From lead chamber
scale by digestion with Na,C03 and KCN , forming KCNTe . The decanted
solution is acidified with HNO3 and the Te precipitated with H,S (Schimose,
f. Ni, 1884, 49, 157). (.}) For purification of the commercial Te , see Brauner
(M., 1889, 10, 411) and Schimose (C. N., 1884, 49, 20, and 1885, 51, 199).
4. Oxides and Hydroxides. — TeO is said to be formed by heating TeSO3 in a
vacuum above 180°: TeSO, — TeO -f SO. (Divers and Schimose. C.N., 188:!, 47,
221). TeO., forms when Te is burned in the air, and when TeCl4 is decomposed
by boiling water. It is a white crystalline solid, sparingly soluble in H,O .
more Soluble in acids from which solutions water causes a white precipitate of
TeO, or H,TeO:i . H.TeO, is formed when a HNO, solution of Te is immediately
poured into cold water, warming to -10° changes it to TeO. . H.,TeO+ is made,
by fusing TeO, with KNO., , treating the K,TeO4 so obtained with soluble lead
or barium salt and decomposing this salt with H,SO4 or H.,S , colorless crystals,
insoluble in alcohol or ether-alcohol (separation from H.,SO4). Tt can be
recrystallized from water and upon heating forms TeO3 (Clarke, Am. »s'., 1S77,
114, 281; 1S78, 116, -'01).
5. Solubilities. — Te is insoluble in HC1; HNO:! and nitrohydrochloric acids
oxidize it to H.TeO,: in H,SO4 it becomes ELTeO:t with evolution of SO, (Hilger,
A.. 1*74, 171< 211): soluble in warm concentrated solution of KCN. from which
solution HC1 precipitates all the Te . H,TeO, is fairly soluble in water, red-
dens moist litmus paper and easily decomposes into TeO, and H,0 . Acid solu-
138 t*ELK\ir.v. §112, (J.
•
tions of TeO. are precipitated upon addition of water or upon standing. TeO,
and H.TeO.., form soluble alkali ^;ms \\itn ,1. alkalis from which solutions of
the other metallic salts precipitate the respective itllvrites. H2TeO4 is soluble
in water, acids ai.d alkalis: alkali carbonates form acid teluu«.;\\, less soluble
than the corresponding normal salts. Solutions of the alkali tellurates iui-o;
insoluble tellurates with soluble salts of the other metals, i: </. K,TeO 4-
BaCl, = BaTeO, + 2KC1 . . "
»i. Reactions.— Tellurium is classed with second group metals because of its
precipitation from solutions of tellurites and tellurales by H,S . The precipi*
tate is not a sulphide, but is Te mixed with varying proportions of S , for CS.
removes nearly all the sulphur (Meeker, A., 187<>, 180. :>.~>7). In appearance the
precipitate of Te with H.S very much resembles SnS . and is very soluble in
(NH.hS.
At a high temperature Te and H unite directly, forming H...Te (IVauner, M..
1S89, 10, 44<>). H,Te is best prepared by heating together Te and Te. or Zn. ami
decomposing these tellurides with HC1 (analogous to the corresponding reac-
tions with sulphur, §257, 4). A colorless gas, odor similar to H.S , burns with
a blue flame, fairly soluble in water and is precipitated as Te° from its solution
by the oxygen of the air. H,Te precipitates solutions of metallic salts very
similarly to H.S and H.Se .
7. Ignition.- Te combines on ignition with most metals to form tellurides.
TeO;i ignited, decomposes into TeO, and O . All lower Te compounds ignited
with KNO give K.TeO, . All Te compounds give on charcoal with the blow-
pipe a while powder, which colors the reduction flame green and disappears.
Heated in an open glass tube. Te compounds give a sublimate of TeO.. , which
melts upon heating. Te compounds fused with KCN in a current of hydrogen
form potassium tellurocyanate. KCNTe: soluble in water but precipitated by a
current of air as Te° (distinction and separation from Se). Heated with Na.CO,
on charcoal Te compounds give Na,Te . which blackens silver with formation
of Ag.Te .
s. Detection. — My reduction to Te° and solution in cold concentrated H,SO,
to a purplish-red solution (characteristic). Separated from Se by fusion with
KCN in a current of hydrogen and precipitation from the solution by a current,
of air.
!». Estimation. — The Te compound is heated in a current of Cl . TeCl, being
sublimed. This is decomposed by water to TeO, . 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
H,Te . H,S reduces Te compounds to Te° mixed with S . Fusion with KNO,
oxidi/es ail Te compounds to K.TeO, . SO, reduces Te compounds to Te° .
SnCl 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 HC1 evolve chlorine and are reduced to H._.TeO, ,
which precipitates as TeO, on r.dding water if too much HC1 be not present
(distinction from Se).
£113. Selenium. Se — 79.? . Valence two and four, possibly six.
1. Properties. — A»>rr//?r ///•// rih/. of the red variety. 4.2.">0; of the black variety,
4.790 (SchatVgotseh. ./. /</-.. 1*4*. 43. :U)S). It begins to soften between 40° and
.".O0: it is half fluid at about 100°. but is not completely molten until 250°
(Draper and Moss. ('. A.. 1*7(1. 33. 1). The molten Se does not become com-
pletely solid until cooled to .">0°. Selenium with tellurium is closely related t..
Milphur, 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-
i;;ij- from the molten state is insoluble in CS. . HoiHiif/ jiainl. Cull0 to (>S::°
(Camelley and Williams. C. \.. 1S70. 39. 2Sf>).
:.'. Occurrence. — Tn no place abundantly: never native. It is found in com-
bination with minerals in the Hart/ Mountains. Sweden. Argentine Republic ami
Mexico (Millandot. ('. A"., iss-j. 46. GO). It occurs in very small quantities with
some sulphides of Fe . Cu ami Zn .
ill 13, 10. SELENIUM. 139
3. Preparation. — In the lead chambers of the H,SO, works it is found as a
red deposit with some S , As,O;, , Sb,03 , PbSO4 , etc. The scale is washed with
water and digested with KCN solution at 80° to 100°, until the red color entirely
disappears. The filtrate is then treated with HC1 , which precipitates the Se .
It is further purified by oxidation to SeO, , sublimed and then reduced with
SO3 (Nilson, B., 1874, 7, 1719).
4. Oxides- and Hydroxides. — H.SeOj is prepared by oxidizing Se with HNO ,
or nitrohydrochloric acid. H.jSeO3 evaporated to dryness gives H2O and SeO2 ,
crystalline. SeO... is also formed by burning Se in air or oxygen; it has an
odo~ similar to decaying radish. It sublimes at about 200° as a yellow vapor,
condensing to white needles on cooling. SeO:, is not known. H,SeO4 , pure,
is a white crystalline mass, melting at 5S°. H,Se04.H,O is crystalline at — 38°,
and if recrystallized melts at 2.")°. The selenic acid usually obtained is a thick
oily liquid, resembling HL,SO, and containing about 95 per cent H2SeO4 . It is
obtained by fusing Se or SeO, with KNO:1 and precipitation of the K2SeO4 with
soluble salts oi' Ba . Pb . Ca or Cu and decomposing the washed precipitates,
suspended in water, with H..SO, or H,S .
r>. Solubilities. — Se dissolves in cold concentrated EL.SO, to a green colored
solution without oxidation (dilution with water precipitates the Se) ; if the
solution be warmed SO, is evolved and the green color disappears (dilution
with water gives precipitate), the Se being oxidized to SeO., . HN03 and nitro-
hydroehloric acid oxidize it to SeO . Selenous oxide, ScO, , is soluble in water
in all proportions, forming H,Se03 . 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, r. (/.. Na^SeO^ + BaCl2 = BaSeO + 2NaCl . Many
Of the selenites are soluble in excess of H,SeO, . Selenates are less stable
than selenites. BaSeO, is soluble in HC1 (distinction and separation from
BaS04) and upon long-continued boiling is reduced to BaSeO .
f>. Reactions. — Selenous acid precipitates with H,S a mixture of Se and S ,
lemon yellow, bright red upon heating (Divers and Shimose, C. N., 1885, 51,
199). This mixture is soluble in (NH4).,S , hence in qualitative analysis Se is
classed among the metals of division A, second group, while because of its
geiieral properties it belongs with sulphur. When Se and H are heated to-
gether they begin to combine directly at 2.50°, forming H,Se (Ditte, C. r., 1872,
74, uso) : which in practically all its reactions is similar to H,S . H,Se is also
formed by treating K.Se . FeSe . etc., with dilute HC1 or H,SO,: HN03 gives
H;SeO:( with selenides. HLSe is a colorless gas. odor similar to H,S but more
penetrating. It is more poisonous than H^S , burns when ignited, combines
.slowly but completely with Hg° . c-volving hydrogen. It dissolves in water to a
greater extent than H...S , 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 TCeeb, J. PJiarni., 1869. (4), 9,
]?.'{). With soluble sulphites H,Se gives a precipitate of a mixture of Se and S .
7. Ignition. — When Se or compounds of Se are fused with KCN in a current
of hydrogen, potassium selenocyanate, KCNSe , is formed. Long boiling with
HC1 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 Na,Se , which yields a black stain with Ag°
ami H,Se with dilute acids.
s. Detection. — If in solution as selenites it is precipitated with H2S (soluble
in (NH4),S): 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 solenides. decomposed by HC1 or H2SO4 , forming H,Se . which is conducted
into water and the Se° precipitated by passing air or oxygen through the solu-
tion.
.'. Estimation. — Oxidized to selenic acid and precipitated as BaSeO, and
weighed as such. If BaSO, be present the precipitate is reduced in H, and
the resulting BaSeO separated by solution in HC1 . 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 H,SeO, .
10. Oxidation. — Se° is oxidized to SeO, by HN03 , nitrohydrochloric acid,
140 THE IRON AND ZINC GKOl'PS. §114.
H2S04 hot concentrated, by heating- in air or oxygen, etc. H,SeO3 is oxidi/ed
to H SeO, by continued chlorination. and by fusion with KNO . H SeO, is
reduced to H,SeO by boiling' \vith HC1 . SO, reduces selenous compounds to
the red Se° , even in H SO, solutions (distinction from tellurium) (Keller,
J. Am. /S'or., 1900, 22, 241). H,,S forms a precipitate of Se mixed with S . SnCl,
precipitates Se° from HC1 or H SO solutions of selenous compounds.
THE IRON AXD ZINC GROUPS (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:!8.ti
Chromium Or = 52.1 Neodymium Nd = 143.6
Iron Fe = 55.fi Praseodymium Pr = He.fi
Cobalt Co = 59.00 Samarium Sm = 150..'!
Nickel Ni = 58.70 Scandium Sc = 44.1
Manganese Mn = 55.0 Tantalum Ta = 182.K
Zinc Zn = 65.4 Terbium Tr = IfiO.
Cerium Ce = 139.0 Thallium Tl =204.15
Columbium Cb = 93.7 Thorium Th =232.6
Erbium E =10(5.0 Titanium Ti = 48.15
Gallium Ga = 70.0 Uranium U =239.6
Glucinum Ol = 9.1 Ytterbium Yb = 173.!i
Indium In =114.0 Yttrium Y = 89.0
Zirconium Zr = 90.4
§115. 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 with 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
ugeiits (§8, g9).
§116. The oxides and Jn/i1ro.ri<1f* of these metals arc insoluble in water
nnd they are precipitated from all their salts by alkalis. In tin- case of
zinc, the precipitate redissolves in all the alkalis: the ahnitiintni. 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 rnbtill 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 tar/uric acitl, citric ndil, siu/ar, find sotne other organic
substances, prevents the precipitation <>f hasp* of these i/nui/i* Inj nil-all^-.
.§117. Ammonium salts, as NH4C1 . dissolve moderate nuantities of
£120. THE IRON AND ZINC GROUPS. 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 ammonium hydroxide, with ammonium
chloride, the latter necessary on account of magnesium (§189, 6*7), nvm-
ganese (§134, 6a), and aluminum, will fully precipitate only aluminum,
chromium, and ferricum 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, ;ind ferricum are
precipitated by ammonium hydroxide in the presence of ammonium
chloride (Fe" by its previous oxidation with HNO:! 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 ,
Ni , 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. Hydrosulphuric 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 — FeCl._, -f~ H2S = FeS -j-
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(OH), , 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 ALUMIXl'M. §121.
by hydrosulphuric acid, in solution, with a precipitation of sulphur, a.n<l
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 (§205, 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 (NH4)2CO:. . In the case of ferrous and manganous
*alts, 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 Fe'", 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 , Ni ,
and Zn are redissolved by excess of NH4OH . 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 les<
soluble than the phosphate: FeHP04 + (NH4),S = FeS + (NH4),HPO, .
Hsnee, 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 cliarconl, 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 IROV ORorv (Tinitn GROUP).
Aluminum, Chromium, Iron.
124, Aluminum. Al = 27.1 . Valence three.
1. Properties. N/'m'/'V i/nirHi/. 2.58:! (Mallet. ('. .V.. 1SS2, 46. 178). Melting
point, R54.50 (Heyeock and Neville. ./. ('.. iv.».">. 67. 1ST). It is a tin-white metal
(the powder is gray). odorless and tasteless, very ductile and malleable, about
us hard as silver. It has not been vaporized, impurities increase the melting-
point. when molten it possesses great fluidity. As a conductor of heal il is ;i
little better than tin and about two-thirds as {rood as silver. It conducts
electricity about one-half as well as copper (Poggendorf, P<><>!/., isr>(>, 97, 643),
about one-third as well as silver (Matthiessen, Pogg., 1858, 103, 428), and about
times better than iron. Commercial aluminum is never pure, containing
U24, da. ALUMINUM. 143
small amounts of silicon and iron, and sometimes Cu and Pb , with 96 to 99
per cent aluminum (Hampe, A.. 187(5, 183, 78). It is used for cooking utensils,
canteens and other military equipments, boats, small weights, measures,
articles of ornament and scientific instruments; as an alloy with copper
i aluminum bronze) it finds extensive application.
2. Occurrence. — Not found free iu nature. Is found in coruudrum, ruby and
*apphiro, as nearly pun- A1..O : in diaspore (A1OOH): in bauxite (A1,O(OH)4);
in felspar (K,Al.,SiO6): in cryo'i e (Na.AlF9i. As a silicate in all clays and iu
very many minerals. It is widely distributed, constituting about one-twelfth of tho
earth's crust.
3. Preparation. — (./) By electrolysis of the fused NaAlCl4 . (2) By fusion of
cryolite or the chloride with Na or K . (3) By heating NaAlCl4 with zinc, with
which it forms an alloy from which the zinc is driven off by a white heat.
(4) By fusion of the chloride with potassium cyanide. (5) By fusing A12SS
with iron. A great many i<ew methods have been patented. See Dammer, 3,
79.
4. Oxide and Hydroxides. — ALO;, 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 KHS04 or Na2CO:i . A1(OH)3 is
formed when aluminum salts are precipitated with cold ammonium hydroxide.
A1,O(OH)4 is formed if the precipitation is made at 100°.
5. Solubilities. — a. — Metal. — Pure aluminum scarcely oxidizes at all in dry or
racist 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, BL, 1881, (2)^ 36, 550). Dilute sulphuric acid attacks it
slowly, evolving hydrogen (Ditte, (J. r.,'l890, 110, 57:;); 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, (:J>), 43, 14; Montemartini,
tfazsctta, 1892, 22, :'.97; Ditte, I.e., 782). Hydrochloric acid, dilute or concen-
trated, dissolves it readily with evolution of hydrogen: also attacked readily
by fixed alkalis, sparingly by NH4OH (Gottig, B., 1896, 29, 1671), evolving
hydrogen with formation of an aluminate: 2A1 + 2KOH + 2H2O = 2KA1O2 +
.'fHji . It is attacked by fixed alkali carbonates (D., 3, S7). When ignited with
sodium carbonate, aluminum oxide is formed, sodium is vaporized and a small
•jmount of aluminum nitride produced (Mallet, J. C., 1876, 30, 349). Fused
KOH is decomposed by aluminum at very high temperature, the potassium
being vaporized (Deville, •/.. 1S57. 152). 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 part in 526).
6. — Oxide and hydroxide. — The oxdde is insoluble in water, and when not
too strongly ignited dissolves readily in dilute acids and in fixed alkalis.
Corundum, crystallized A1,.O:1 , is insoluble in acids, but is rendered soluble
fiy 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, c. — Xdlts. — 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 Cr 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. Cfi., 1859, (ri), 56, 102).
Solutions of normal salts of aluminum have an acid reaction.
f>. Reactions, a.— The alkali hydroxides and carbonates* precipitate
aluminum hydroxide (1), A1(OH)S (4), grayish-white, gelatinous insoluble
* According to Langlois (A. Ch., 1856, (3), 48. 502) the precipitate with alkali carbonates always
fontains COS. He assigns the formula 3(A13O3 CO2) + 5( Al,oa.8H,O).
144 ALU.MISU.V. S124, fib.
in water, soluble in excess of the fixed alkali hydroxides* (2) (Preseott,
'•/. Am. 8or., 1880, 2, 27; Ditte. .-I. Ch., 1897 (6), 30, '200), 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 alu initiate,
is precipitated as aluminum hydroxide by careful neutralization of the
alkali with acids including hydrosulphuric (3}, and carbonic, as basic
hydroxide, by adding excess of ammonium chloride (4) (distinction from
zinc which is precipitated by a small amount of NH4C1 , but redissolves on
adding an excess) (Lowe, /?., 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. (St-e
£126, 6fl.) The presence of citric, oxalic, or tartaric acid greatly hinders
the precipitation of aluminum hydroxide, and an excess may entirely pro-
vent its precipitation by the formation of a soluble double salt, e. g,,
KA1(C4H40,:).; . Other organic substances, as sugar, pieces of filter paper,
etc., hinder the precipitation. To obtain complete precipitation all or-
ganic substances should be decomposed.
(/) A1C1, + HXOH = A1(OH)3 + liKCl
:>A1C18 + 3K.COS + .5H2O = 2A1(OH), 4- f.KCl + !CO,
(2) A1(OH):, 4- KOH = KA10, + 2H,0
or A1C1, + 4 KOH = KA10, + .iKCl + 2H.O
(3) 2KAHX + H2S + 2H20 = 2A1(OH)S + K,S
(.J) 2KA1O, + 2NH4C1 + H,0 = A1.O(OH)4 + 2KC1 4- 2NH3
(.7) 2A1C1, 4- P,BaCO3 + 3H=O = 2A1(OH), + :iBaCl. + SCO,
1). — Oxalates do not precipitate aluminum salts. The <i<rl«lt' of alum-
inum is decomposed upon boiling, forming the insoluble basic acetate
(separation of iron and aluminum from the fourth group): A1(C.,H..O.,), 4-
H20 — A1(C,H,0,),OH + HC.,H,0, . The basic acetate is best formed w
follows: To the solution of aluminum salt add a little sodium or am-
monium carbonate, as much as can be added without leaving a preeipitate
on stirring, then add excess of sodium or ammonium acetate, and boil for
some time, when the precipitation at length becomes very nearly com]'
Phenyl hydrazine, C..H-NHNH., . completely precipitate^ aluminum ».s
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.. 180!). 21; 776).
*A solution of barium hydroxide may he used to dissolve the A1(OH)3 in separating frum
Fe(OH), and Cr(OHi3; especially valuable in detecting the presence of small amounts of
aluminum when the reagents N«OH and KOH contain aluminum •'Neumann, IT.. 1S1W, 15. W
§124, «Sr. 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
A14OS(NO3), (Ditte, C. r., 1890, 110, 782).
d. — Alkali phosphates precipitate aluminum phosphate, A1P04 , white,
insoluble in water and acetic acid, soluble in mineral acids, and in the
fixed alkalis (separation from FeP04) (Grueber, Z. angew., 1896, 741).
A separation of Al and P04 may be effected by dissolving in hydrochloric
acid adding- tartaric acid and then ammonium hydroxide, and digesting
Home 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 nitrate
contains nearly all .of the aluminum. The same method may be employed
with Fe'" and P04 . 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. N., 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 H2S is evolved: 2A1CL + 3(NH4)2S -f 6H,0 = 2A1(OH)3 + 6NH4C1
-f- 3H..S , from solutions in the fixed alkalis ammonia is evolved, fixefj
alkali sulphide being formed: 2KA102 -f (NH4)2S -f 2H20 = 2A1(OH)3 -f
K2S + 2NH3 .
Sodium thiosulphate precipitates, from aluminum salts, in neutral solutions,
aluminum hydroxide with free sulphur and liberation of sulphurous anhydride:
2A12(S04)3 + 6Na,S203 + 6H2O = 4A1(OH)3 + 3S2 + 6Na,S04 + 6S03 . A
small 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: 2A1C13 + 3Na2SO3 -i
iH.O = 2A1(OH)3 -f- GNaCl + 3SQ, . 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 tho
potassium aluminum compound is perhaps best known, KA1(S04)2.12H.,0 ,
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 wate1*. 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.
14G ALUMINUM. §124, i>f.
f.— Aluminum chloride is a very powerful dehydrating agent and is much
vsed in organic chemistry as a halogen carrier. An impure aluminum chlorate,
mixture of KC1O3 and A12(SO4)3 , is much used in calico printing (Schlum-
berger, Diny!., 187.'!, 207, 6i!). y. — Aluminum salts are precipitated by solu-
tions of alkali arseiiites and arseirites. but not by arsenous or arsenic acids.
A. — J'otassium chromatc forms a yellow gelatinous precipitate, potassium
bichromate gives no precipitate with aluminum salts. /. — Solution of horav
precipitales an ackl aluminum borate, quickly changed to aluminum hydroxide.
7. Ignition. — Compounds of aluminum are not reduced to the metal, but
most of them are changt-d 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 lilac 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 ALS... (liuch-
erer, Z. <iii</cir., 1892, 48.'>).
To scpnrnfr Al from PO, . fuse the precipitate* or powdered substance with
\y» parts finely divided silica and <1 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 by dissolving it in hydrochloric acid,
evaporating to dryness to render the silica insoluble. Treat with hydrochloric
acid and tilter: the filtrate containing aluminum chloride.
8. Detection. — After the removal of tlu> first two groups it is precipi-
tated with Cr and Fe'" as the hydroxide, Al(OH) , , by NH.OH in. the pres-
ence of NH.C1 . It is separated from Fe(OH)., and Cr(OH), by boiling
with KOH . From the filtrate acidulated with HC1 it is precipitated as
hydroxide with (NH4)..CO, ; or it is precipitated from the KOH solution
by an excess of NH4C1 ((>«).
9. Estimation. — Aluminum is usually weighed as the oxide, after ignition.
It is separated from /ine as a basic acetate; from chromium by oxidi/ing tin-
latter to chromic acid, by boiling with potassium chlorate and nitric acid, or
by fusing with KNO, and Na C0:, . or by action of Cl or Br in presence of
KOH. and after acidulating with HC1 precipitating the aluminum with am-
monium hydroxide. It may be separated from iron by boiling with KOH (<>a).
by Na...S..O., (lie), or by phenylhydra/.ine (lib). It is separated from iron by
conversion into the oleale and dissolving the oleate of iron (Fe'" or Fe") in
petroleum ( Horntraeirer. 7,.. 1S9:>. 32. 1S7). It is sometimes precipitated and
weighed as the phosphate.
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 stale: ferric salts to ferrous salts;
As and Sb with HC1 become respectively AsH, and SbH:! with alkalis As'"
is reduced to AsHa . Asv is unchanged (§69, (!'/; and 10). and Sb'" and
Sbv become Sb°. Ahiminum salts are not reduced to the metallic state
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. G., 1893. 301 ; Wislicenus, B. !•««>. 2S, 1333. t Tommasi, Bl., 18»2. <»', 37. ' Ci
J Flavitsky, B., 1873, 6. 195 ; Zimmermen, Z., 1888, 27, 61.
£125, 5c. CHROMIUM. 14?
§125. Chromium. Cr — 52.1 . Valence two, three and six.
1. Properties. — Upccifie orarity, t>.81 (Woehler, A., 1859, 111, 2:;i). Melt* with
greater difficulty than platinum ((Hat/el, II., 1890, 23, .'5127). 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.e.). It
burns to the oxide Cr.Oa when heated to 200° to M)0° in the air (Moissan, ('. r..
1879, 88, 180).
2. Occurrence. — Not found native. It is found in several minerals. Chrome-
ironstone or chromite (FeOCr.,0:i) is the chief ore of chromium, and is usual Iv
employed in the manufacture of chromium compounds. Chromite and also a
double sulphide of iron and chromium, FeCr.,S4 , are found in many meteors.
3. Preparation. — (1) By electrolysis of the chloride. (2) By fusing the
chloride with potassium or sodium. (3) By ignition of the oxide with carbon.
(//) Hy fusing CrCi, with Zn . Cd or Mg , using KC1 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, A., 1898, 301, 19).
4. Oxides and Hydroxides. — ('liroiiunix o.ride, CrO , has not been isolated. The
corresponding hydroxide, Cr(OH)^ , is made by treating CrCL with KOH .
Chromic o.i-idf, CrL,O3 , 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 diehromates of the alkalis:
4Hg2Cr04 = 20^03 + SHg + 5O2
(NH4)8CrsOT = Cr203 + N2 + 4H20
4K2Cr2O7 = 2Cr203 + 4K2CrO4 + 3O,
In the last the K2CrO4 may be separated by water. After heating to redness,
Cr.O3 is insoluble in acids. Chromic hydroxide, Cr(OH)3 , 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 trioj-idc or chromic anhydride, CrO3 , is formed as brown-red
needles upon addition of concentrated sulphuric acid to a concentrated solution
of K2Cr2O7; to be freed from sulphuric acid it must be recrj'stallized from
water, in which it is readily soluble, or treated with the necessary amount of
BaCrO4 (Moissan, A. Ch., 1885, (6), 5, 568). It is also prepared by transposi-
tion of BaCrO4 with HN03 or H,S04; PbCrO4 with H,SO4: and Ag,CrO4 with
HC1: etc. It melts at about 170° (Moissan, I.e.), decomposing at higher tem-
perature into Cr,O3 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, H-Cr04 , is disputed (Moissan, I.e.;
Field, C. N., 1892, 65, 153; and Ostwald, 'Zeit. plii/s. Ch., 1888, 2, 78). Two
series of salts are formed as if derived from chromic acid, H,CrO4 , and
dichromic acid, H,Cr2O7 . The salts are quite stable and find an extended
application in analytical chemistry (67<, §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 Cr.,O:, . rapidly in pres-
ence of KOH . It is soluble in HC1 or dilute H,S04; insoluble MI concentrated
H.jSO, or in HNO:; , dilute or concentrated. Chlorine or bromine attack it
with formation of the corresponding halides (Woehler, I.e.; Ufer. A., 1859, 112,
l'02). li. — Oxides and Hydroxides. — Chromie <i.ride, Cr.,O:; . is insoluble in water,
slowly soluble in acids, but not at all if previously ignited (Tranbe, .4.. 184*,
66. 88); the Iti/drii.ritle is insoluble in water, soluble in acids, sparingly soluble
in ammonium hydroxide, soluble in fixed alkalis to '.-hromites, reprecipitated
again upon boiling. The presence of other metallic hydroxides, as iron, etc.,
hinders the solution in fixed alkalis. Chromic anhydride, CrOn , is very soluble
in water, soluble in reducing acids to chromic salts.
c. — Salts. — Chromic sulphide is not formed in the wet way, being-
decomposed by water; the phosphate is insoluble in water. The chloride
148 CHROMIUM. §125, 6a.
t
exists in two modifications; a deU</u<>xcoil soluble chloride, which also
forms a soluble basic chloride (Ordway, Am. >'., lsr>S (•>). 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. Cli., 1846 (3), 16, M!>S); ',!)<• 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
compounds. 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.
(>. Reactions.* 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(OH), -f KOH = KCrO, + 2H,0 : 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 cold
ammonium hydroxide to a violet solution, completely reprecipitated on
boiling. The precipitate formed with the alkali carbonates is almost
entirely free from carbonate: ?CrCl, + 3Na,CO.. + 3H,0 = '2Cr(OH)n +
<>NaCl + 3CO.. . 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 eold (separation
from the fourth group). For removal of excess of reagent, add H..SO,
and the filtrate will contain the chromium as a sulphate.
Alkali dichromates are changed to normal eliminates by alkali hydrox-
ides or carbonates.
h. — Chromium forms no basic acetate and remains in solution when the
basic acetates of alnmir.nm and ferric iron are formed (<>&, §124 and §126).
Potassium cyanide precipitates chromium hydroxide. Oxalates and ferro-
cyanides cause no precipitate. H^CrO, is reduced to chromic compounds
«
•Chromous salts are very unstable, they are great reducing agents, oxidizing rapidly when
exposed to the air. They are almost never met with in analysis. Chromous chloride, CrCla, is
formed when the metal is heated in contact with hydrochloric acid gas (Ufer, I. c ); also by re
duction of CrCI3 with hydrogen in a heated tube (Moberg. J. pr . 1K48. 44, 323). Precipitates are
formed in its solutions by the alkali hydroxides, carbonates, sulphides, etc. (Moissan. 75/.. ISN*
(2), 37. 296).
£125, G/t. CHROMIUM. 149
by K3Fe(CN)6 and KCNS. <: — 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 ohromate is formed (separation from Fe and
Al). d. — Hypophosphorous acid reduces chromates to chromic salts. Soluble
phosphates, as Na,HPO4 , precipitate chromic phosphate, CrP04 , insoluble in
acetic acid, decomposed by boiling- with KOH , 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 KCr02 are precipitated as chromium
hydroxide; SKCrO, + HJ3 + 2H,0 == 2Cr(OH)3 + KJ3 . The hexad
chromium of chromates is reduced to the triad condition with liberation
of sulphur, in neutral or alkaline solutions, chromium hydroxide being
formed: 2K2Cr,07 + 8H,S = 4Cr(OH):, -f 2K2S + 3S2 + 2H20 ; in acid
solutions a chromium salt is formed (10). Alkali sulphides precipitate
chromium, salts as the hydroxide liberating H2S :
2CrCl3 + 3(NH4)2S + GH2O = 2Cr(OH)3 + 6NH4C1 + 3H,S
Chromates are reduced and precipitated as chromium hydroxide with sepa-
ration of sulphur: 4K2Cr04 + 6(NH4)2S + 4H20 -I 4Cr(OH), + 8KOH
-f- 3S2 -f- 12iyH3 . Soluble sulphites and thiosulphates reduce chromates
in acid solution (Donath, J. C., 1879, 36, 401; Longi, Gazzetta, 1896, 26,
ii, 119).
/. — Hydrochloric acid reduces chromates to chromic chloride on boiling,
with evolution of chlorine: 2K2Cr04 + 1GHC1 = 2CrCl3 -f 4KC1 -f 3C12 +
8H20 ; more readily without evolution of chlorine in presence of other
easily oxidized agents, as alcohol, oxalic acid, etc.: K2Cr,07 -f- 8HC1 -f-
3C2H5OH = 2KC1 + 2CrCl3 -f 3C2H40 (acetaldehyde) + 7H20 . If the
dry chromate be heated with sulphuric acid and a chloride (transposable
by sulphuric acid) (§269, 5), brown fumes of chromium dioxydichloride
are evolved: K2Cr207 -f 4NaCl + 3H,S04 = 2CrO,Cl2 + K2S04 + 2Na2S04
-f 3H20 (§269, Sd) (Moissan, BL, 1885 (2), 43, 6). To obtain a quantity of
CrOoCl, , Thorpe (J. C., 1868, 21, 514) recommends 10 parts of NaCl and
12 parts K2Cr207 fused together and distilled with 30 parts of H2S04 .
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. K2Cr207 + 6HI + 4H2S04 = K2S04 -f Cr,(S04)3 +
3I2 -f" 7H20 . 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
gives CrAs04 (Neville, J. (7., 1877, 31, 283).
ft. — Potassium chromate colors an acid solution of chromic salt brown-yellow;
150 CHROMIUM. £125, 7.
on addition of ammonium hydroxide, a precipitate of the same color is obtained.
chromic chromate (Mans, Pogtf., 1827, 9. 127). The alkali metals form two
classes of chromates: yellow normal chromates and reddish dichromates
(Schulernd. •/. ('., 1879, 36, 2!)8). The chromates of the alkalis, and those of
magnesium, calcium, xinc and copper are soluble: those of strontium, mercury
(Hg") are sparingly soluble: and those of barium, manganese, bismuth, n.-r-
cury (Hg/), silver and lead are insoluble in water. Alkali chromales <>r
dichromates are precipitated as normal chromates (in some cases as dichro-
mates) (Preis and Hayman, Jf., 1880, 13, :$4()) by solutions of si her. lead, mer-
cury (Hg^) and barium salts. Silver eliminate is dark red, soluble in nitric
acid and ammonium hydroxide (§59. fill): lead chromate is yellow, transposed
with difficulty by nitric acid (Duvillier, A. Cli.. 1S7I!, (4), 30, 212), insoluble in
acetic acid (§57, <>7;); barium chromate, yellow, is soluble in hydrochloric and
nitric acids, sparingly soluble in chromic acid (§186. (57;).
7. Ignition. — Chromic oxide, chromic salts and ehromates 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. Cr.,0, . is left. When
chromium compounds are fused with an alkali carbonate, and a nitrite, nitrate,
chlorate, bromatc or iodate, an alkali chromate is formed, soluble in water
(distinction from Al and Fe).
8. Detection. — It" present as chromate (solution red or yellow), it is
reduced by HC1 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 NH4C1 . Boiling with KOH separates the Al and leaves the Cr
with the Fe , as hydroxides. The precipitate is fused on a platinum foil
with Na..CO:1 and KNOa which oxidizes the Cr to an alkali chromate, soluble
in water (separation from the Fe). The Cr is identified after acidulation
with HCoH.,0., by the formation of the yellow lead chromate, using
Pb(C,Hs02), . "
'.). Estimation. — Chromium is usually estimated gravimetrically (/) as the
oxide. It is brought into this form either by precipitation as a hydroxide (fin)
and ignition or, in many cases, by simple ignition (4). (,?) 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(C,H:,0,), . dried and
weighed. 'Volumetrically, as a chromate (if present as chromic salt it may be-
oxidi/ed to a chromate). (.?) By titration with a standard solution of ferrous
sulphate. (.}) By liberation of iodine from hydriodic acid (<>'/) and measuring
the amount of iodine liberated with standard sodium thiosulphate solution.
10. Oxidation.— fhromous compounds arc very strong reducing agents,
changing HgCl2 to HgCl , CuS04 to Cu°, SnCl, to Sn°, etc. Chromic com-
pounds are oxidized to chromates by chlorates ((iiacomelli, L'Oroxi, 1895.
18, 48; Storer, Am. 8., 1869,98,190) (<;/'), Na,0,, MnO, (Marchal and Wier-
nick, Z. ant/en'., 1891, 511), and PbO., in acid solution; in alkaline mixture,
by reducing PbO, to PbO, Ag.O to Ag°, Hg,0 and HgO to Hg°, CuO to
Cu.O , KMn04 and K,Mn04 to MnO, (Donath and .Idler, ('. (\, 1887, 151);
by Cl , Br, and I, forming the corresponding halide; and by H202*
* The use of H,O, in alkaline solution is proposed by Uiggs (Am. S., 1894, 148, 409) in the sepa-
ration of Al, F"e and Cr. lOOcc. water, 10 cc. H...O,, and one grram of NaOH are added to tho
freshly precipitated hydroxides and digested until effervescence ceases. Filter off the precipi
tate of ferric hydro.xido, 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
(Baumann, Z. aiic/ew.. 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 Cr02Cl2 (Moissan, BL, 1880 (2), 34, 70).
Chromic acid and chromatcs are reduced to chromic compounds by
H2C204 (Werner, J. C., 1888, 53, 602), K4Fe(CN)6 , KCNS, H,S , (NHJ2S,
Na2S20.. , S02 , H202 , 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. N., 1878, 38, 228).
2K..Cr,O^ + 16HC1 + <>H,S = 4CrCl3 + 4KC1 + 3S2 + 14H2O
12H2Cr04 + 3S2 = 4Cr,,0,Cr04 + fiSO, + 12H.O
2H2CrO, + ?,BO, =Cr,(SOt), + 2H20
While H202 in alkaline solution oxidizes Or'" to CrVI, in acid solution the
reverse * action takes place : 2H2Cr04 -f 3H2S04 + 3H202 = Cr2(S04)3 -4-
302 + 8H20 (Baumann, 1. c.).
§126. Iron (Ferrum). Fe •= 55.9 . Usual valence two and three.
1. Properties. — Specific g rarity, 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 1000°: 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
ignited; 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 J.7 to 4. (5 per cent, and wroiight 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 (Fe2O3), brown hematite
(2Fe2O3.3H2O), magnetic iron ore (Fe3O4), iron pyrites (FeS2), spathic iroa
ore (FeCO3), clay iron-stone (FeC03 with clay), black band (FeCO3 mixed with
bituminous matter).
3. Preparation. — Pure iron is not usually found in the market. It is made:
(1) by electrolysis; (2) by heating its purified salts with hydrogen; (3) by
heating the purified salts with some form of carbon; (//) in metallurgy iron is
made from the ores, and the reducing agents are coal, coke, charcoal and
natural gas.
* With a chromate in acid solution H2O2 at first gives a deep blue solution (probably of per-
chromic acid, HCrO4) a very delicate reaction, followed by the reduction to a chromic salt.
152 IRON. $126, 4.
4. Oxides and Hydroxides.— Fcrrtmx o.riilc. FeO . is made from Fe,O3 by heat-
ing it to :$0()° in an atmosphere of hydrogen: also by heating1 Fe.C.O, to 160°,
air being excluded. It takes tire spontaneously in the air, oxidizing to Fe.jO3 .
Ferrous hudrojri<li; Fe(OH), . is formed by precipitating ferrous salts with KOH
or NaOH , perfectly white when pure, but usually green from partial oxidation.
Ferrw oj-iilc, Fe^O3 , is formed by heating FeO , Fe(OH),. . or any ferrous salt
consisting of a volatile or organic acid in the air: more rapidly by heating
Fe(OH), . Fe(NO3)3 . or Fe.,(SO4)3 . 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 washing than when
NH4OH is employed. Ily increasing the temperature and concentration of the
solutions, the following definite compounds may be formed: FeO(OH) ,
Fe,O(OH)4 . Fe4O,(OH), . Fe4O,(OH),, . Fe:,O,(OH),'. Fe;)04 is slowly formed
by heating FeO or Fe.O, to a white heat. Its corresponding hydroxide may be
made by precipitation: FeCL + 'JFeCL, + SNH4OH = Fe:i(OH)s + sNH4Cl .
Fe3(OH), when heated to (.K)° forms FesO4 . The black color and magnetic
properties show that it is a chemical xtttt and not a mechanical mixture of FeO
and Fe,Oa . Fe'" acts as an acid towards the Fe"; this oxide, Fe.,O4 , or
FeFe_.O4 , may be called ferrous ferrite. Other ferrites have been formed, r. g..
calcium ferrite, CaFe.O, ; MgFe,O4 and BaFe,O, (List, II., 1878, 11, 151.2): zinc
ferrite, ZnFe^.0, . Compare potassium aluminate, KA1O (§124, (>a). and potas-
sium chromite. KCrO, (§125, <ia). Ferric aciil. H,FeO4 , and its anhydride,
FeO, , have not been isolated, rutaxxiuin frrnitt-, K_.FeO4 , is made (/) by elec-
trolysis; (2) by heating iron-tilings, FeO or Fe.O, , to a red heat with KNO, :
(3) by heating Fe(OH)a with potassium peroxide K^O, ; ( J) by passing Cl 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: 4K..,FeO4 +
10H.O -- sKOH + 4Fe(OH)3 + :iO, . \Vith barium salts it precipitates a
stable barium ferrate, BaFeO, .
5. Solubilities. — a.—Mcttil. — Iron dissolves, in hydrochloric acid and in dilute
sulphuric acid, to ferrous salts, with liberation of hydrogen (n)', concentrated
cold H SO, has no action, but if hot. SO... is evolved and a ferric salt formed (/>) :
in moderately dilute nitric acid, with heat, to ferric nitrate, liberating chiefly
nitric oxide (r): in cold dilute nitric acid, forming ferrous nitrate with pro-
duction of ammonium nitrate (f/). of nitrous oxide (e), or of hydrogen f)
(Lauglois, .4. Ch., 1856, [3], 48, 502).
(«) Fe + H,S04 = FeSO4 + H2
(h) 2Fe + OH.SO; = Fe,(S04)3 + 3SO2 + GH20
(c) Fe + 4HN03 = Fe(NO,)3 + NO + 2H,0
(d) 4Fe + 10HNO, = 4Fe(NO8), + NH4N03 + :5H20
(e) 4Fe + 10HNO3 = 4Fe(NO3), + NaO + 5H2O
(0 Fe + 2HN03 = Fe(NO,), + 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. Am., 1896, 18. s::r>), and so far as uncombined
will remain undissolved, as graphitic carbon. The metal is attacked by moist
air, forming chiefly 2Fe,O3.:tH,0 , iron rust. \Vhen hot iron is hammered, scale
oxide, Fe.Oj.liFeO , is formed. Cold concentrated 'HNO, forms passive iron.
6. — Oxides and hydroxides.— Ft mm* n.ndc and liiidn>.r'nlv 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 FeS . Fe.,O4 ,
treated with an insufficient amount ofHCl. forms FeCL and Fe,0,.: treated with
HC1 sufficient for complete solution, a mixture of FeCL and FeCL is obtained.
which, when treated with excess of ammonium hydroxide and dried at 100°
again exhibits the magnetic properties of the original. Ferric ii.rhlf, Fe.O., , dis-
solves in acids, quite slowly if the temperature of preparation of the oxide has
been high. Mitscherlieh (•/. //>'.. IMIO. 81. 110) recommends warm digest ion with
ten parts of a mixture of sulphuric acid and water (8-3). If the oxide be
£126, 6a. IRON. 153
heated with alkalis or alkali carbonates, it then dissolves much more readily in
acids. Ferric liydroaridc, Fe(OH)3 , is insohible 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 Fe.,03 are present to one of the FeClr. (Be-
champ, A. Ch., 1859, (3), 56, 306)
c.— Salts. — Ferrous salts, in crystals and in solution, have a light green
color. Solutions of the salts have a slight acid reaction toward litmus.
The sulphate FeS04.7H.,0 , 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, ferro-
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 sohible in alcohol
.(separation from ferrous sulphate). Ferric chloride is soluble in ether
saturated with hydrochloric acid, separation from aluminum (Gooch and
Havens, Am. 8., 1896, 152, 416). Solutions of ferric salts, when boiled,
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, U footnote).
6. Reactions, a. — The alkali hydroxides precipitate ferrous hydroxide,
Fe(OH)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 arid
117). The soluble carbonates precipitate, from purely ferrous solutions,
ferrous carbonate, FeC03 , 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, Fe(OH)3 , variable to FeoO^HoO — FeO(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
barium 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 -(- 3BaC03 -|-
3H20 = 2Fe(OH)8 + 3BaCl2 + 3C02). 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 -f- BaCO., = NiC03 + BaS04 . 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 , Mg , Mn . Zn , and Cu react similar to the
barium carbonate.
b. — Oxalic acid and soluble oxalates precipitate from solutions of ferrous
salts, ferrous oxalate, FeC_.O4 , yellowish-white, crystalline, sparingly soluble in
hot water, soluble in HC1 . HNO and H SO,: ferric salts are not precipitated
by oxalates except as reduction to ferrous oxalate takes place.
The acetates, as NaC.,H.,0., , form in solutions of ferric salts a dull red *
solution of ferric acetate, Fe(C.,H:iO.,)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).,). The basic
precipitates are soluble in HC1 , HNO., and H.,S04 and are transposed by
alkali hydroxides.
Tannic acid precipitates concentrated solutions of ferrous salts: ferric salts
are precipitated as blue-black ferric tannate (the Ixixix of I-HIHIIUHI ink), insoluble
in water or acetic acid, very soluble in excess of tannie acid. Ferric salts arc
completely precipitated by ammonium succinate from hot solutions (Younjf.
J. ('., 1880, 37, 074). Moth ferrous and ferric salts (not nitrates) slightly acid
are completely precipitated by a solution of nitroso B. naphthol (separation
from aluminum and chromium) (Knorre, /*., lss7. 20. 'Js:;: Meiiickc, /. tni<i<'\r.,
1SS8, r>). If the Fe"' be in excess of the PO, the phosphate will all be pr« -
eipitated. Hydrochloric acid should be absent. /'. *'.. excess of NaCIT.O; should
be added (Knorre. Z. anyetc., 1893, 267).
Potassium cyanide g-ives with solutions of ferrous salts a yellowish-red pre-
cipitate, which dissolves in excess of the reagent to potassium ferrocyanide,
K.FefCN),,: with solutions of ferric salts, ferric hydroxide is precipitated with
evolution of hydrocyanic acid (equation («), page l.Vi).
Potassium ferrocyanide precipitates ferrous salts as potassium fcrrnu<
frrrocyaiiiile (ft), K,FeFe(CN)0 , (Everitt's salt), bluish-white, insoluble in
*Meconic acid and formic acid form red solutions with ferric salts : benzoic acid gives a tlcsli
colored precipitate; phenol, creosote, saligonin, and other hyrtroxy aromatic derivatives give
a blue to violet color. Morphine g ives 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 CuSO4 in five
cc. of water, add five cc. of a ten per cent. « therial solution of salicylic acid. If the layer of
contact assumes a violet color iron is present (Grigge, Z., 1895, 34, 450).
£126, M. y/fo.v. 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)< Fe4(Fe(CN)6)3 , 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 Fe.,0;, (Koningh, Z. angeir., 1898, 463). Potassium ferri-
cyanide precipitates from dilute solutions of ferrous salts ferrous ferri-
cyanide (V/), Fe,(Fe(CN),.)., (TurnbulFs blue), dark blue, insoluble in acids,
transposed by alkali hydroxides (/?): with ferric salts no precipitate is
obtained, but the solution is colored brown or green (i). 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 Wood red ferric Hiiocyaiiatc, Fe(CNS), (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
(7.. 1881, 20, 350) one part of iron", as ferric salt, may be detected in
1,KOO,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, .1., 18G4, 130, 24<>). The red color of the liquid is destroyed by
mercuric chloride (I-}, also by phosphates, berates, acetates, oxalates, tar-
trates, racemates, malates, citrates, succinates, and the acids of these salts.
Nitric and chloric acids give red color with potassium thiocyanate, re-
moved by heat.
* The quantity of non-dissociated Fe(CNS;3 , 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 breaking up of the
Fe'CNS)3 to form Hg(CNS), which is («ven less dissociated in water solution than HgCl3.
156 IRON. §126, (XT.
o
(a) FeCl3 + 3KCN + 3H2O = Fe(OH)s + 3KC1 + 3HCN
(6) FeS04 + K,Fe(CN),!=:K,FeFe(CN)0 + K2SO4
(c) K,FeFe(CN)0 + 2KOH = Fe(OH)2 + X,Fe(CN).
(d) 4K,FeFe(CN)0 + O, + 4HC1 = Fe4(Fe(CN)0)3 + K.FeCCN), + 4KC1 + 2H2O
(e) 4FeCl3 + 3K4Fe(CN)tt = Fe4(Fe(CN).)a + 12KC1
(0 Fe4(Fe(CN)9)s + 12KOH = 4Fe(OH)., + 3K4Fe(CN)(J
(0) 3FeSO4 + 2K3Fe(CN),, = Fe..(Fe(CN),.), + 3K,SO4
(h) Fe3(Fe(CN)6)2 + fiKOH = 3Fe(OH), + SKJXCN),
(1) FeCl3 + K3Fe(CN)0 = FeFe(CN)0 + :;KC1
(/) FeCl3 + 3KCNS = Fe(CNS)3 + 3KC1
(fc) 2Fe(CNB)s + 3HgCL = :iHg(CNS)2 + 2FeCl3
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. — Hypophosphorous acid reduces ferric salts to ferrous salts. From
solutions of ferrous salts, alkali phosphates, as Na.,HP04 , precipitate
secondary ferrous phosphate, FeHP04 , mixed with the tertiary salt,
Fe.,(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: 3FeS04 + 2Na,HP04 + 2NaC,H,0, == Fe:1(P04), + ':JNa2S04 +
2HC,H,0., . Ferric salts are precipitated as ferric phosphate, FeP04 ,
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:
Fed, + H,P04 + 3HaC,H8Oa = FeP04 + 3NaCl + 3HC,H,0, . It phos-
phates of the fourth group 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 //.): phosphates of Co. Ni. and Zn being
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 trc;itt'<! with an
* Equilibrium requires that a weak ucirt, as phosphoric, be present for the most part as the
non-dissociated molecule. Hut FePO4 , ns any neutral salt, is dissociated, so far as it dissoh s
in water, into its ions, as is also the strong hydrochloric acid. Hringrinjir these together will r»--
Bult in the union of the H ion of the acid and the PO4 ion to non dissociated H3P«»4, thus
maintaining the equilibrium for H3PO4 , but disturbing that between solid and dissolved
PePO4, which requires a certain concentration of PO4 ions. To restore the latter more FePO4
dissolves, only to react with the H ions as before, and this process roiitinucs 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 HC1 is in excess, until the FePO4 is entirely dissolved. This process takes
place whenever a strong acid dissolves the >a't of a weak one. It is analogous to the solutin
of a base in an acid, forming non-dissociated water.
§126, 6«. IRON. 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: 4FeP04 -f- 6(NH4)oS
— 4FeS -f 4(NH4).,P04 + S2 .
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 H2S in alkaline mixture, form ferrous
sulphide, FeS, 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, Fe,>0(S04)o . Ferric salts are reduced to ferrous salts with
liberation of sulphur by H2S (1), or soluble sulphides, the latter at once
reacting to precipitate ferrous sulphide (2):
(1) 4FeCl3 + 2H,S = 4FeCl2 + 4HC1 + S2
(2) 4FeCl3 + 6(NH4)J5 = 4FeS + 12NH4C1 + S2
After the removal of the metals of the second group by H2S, 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 H2S 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 (Go).
Ferrous sulphite is but little soluble in pure water, easily soluble in excess of
sulphurous acid, to a colorless solution. The moist salt oxidizes rapidly on
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):t , rapidly reduced to the ferrous condition accord-
ing- to the following- equation: Fe,.(SO3)3 = FeSO,, + FeS20() (Gelis. (7. C., 1862,
896). Ferrous thiosulphate, FeS.O.,, is formed, together with some FeS and FeSO ,
by the action of SO2 upon Fe° (Fordos and Gelis, I. c.). Ferric snlts are reduced
by sodium thiosulphate to ferrous salts in neutral solutions with formation of
sodium tetrathionate: 2FeC31n + 2Na,.S,0,. = 2FeCL + 2NaCl + Na,S4O,; (Fordos
and Gelis. (\ r.. 1842. 15, 920): in acid solutions sulphuric acid and sulphur are
'formed: 4FeCl, + 2Na2S2O3 -f 2H2O = 4FeCl, + 4NaCl + ^H,SO4 + S2 (Men-
158 IRON. §126, bf.
schutkin, 78). Ferric iron is precipitated as basic nitrate by the addition of a
solution of ammonium sulphate to a solution of iron in HNO:: 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 salts
by hydriodic acid with liberation of iodine.
tj. — Soluble arsenites and arsenates precipitate solutions of ferrous and ferric
salts, forming1 the corresponding- arsenites and arsenates. Basic ferric arsenite,
4Fe.JO.1.Asi03 + ">H...O , 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 poisoning1 (§69, f>/
and 6V; 'D., 3, 352).
//. Ferrous salts are rapidly oxidized to ferric salts by solutions of chro-
mates, the chromium being reduced to the triad condition (9 and 10).
AVith ferric salts potassium chromate forms a reddish-brown precipitate.
{. — Zinc oxide precipitates solutions of Fe"' , Al , Cr'" and Cu completely and
Pb partially, effecting1 a separation of these metals from Mn , Co and Ni
(Meineke, Z. angcu:., 1888, 258).
7. Ignition. — The larg-er number of iron salts are decomposed, as solids, by
heat; FeCl3 vaporizes partly decomposed, at a very little above 100°. Igni-
tion in the air changes ferrous compounds, and ignition on charcoal or by
reducing1 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 FesO4 from 330° to 440°, FeO
from 500° to (500°. and Fe° above <)00° (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 ?/r//oir when cold, or colorless, if feebly saturated. The same bead, held
persistently in the iTiliirini/ flame, becomes colorless unless strongly saturated,
when it shows the pale grcrn color of ferrous compounds. The reactions with
microcosmic 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 NH4C1 with Al and Cr'" by an excess of NH.OH . The Al is re-
moved by boiling with excess of KOH . Tf more than traces of Fe be
present it is detected in presence of the Cr(OH), , by dissolving in HC1
and obtaining the blood-red solution with KCNS . Tn case Cr be present
in great excess the Cr(OH), and Fe(OH) ; are fused on a platinum foil with
Na,CO, and KNO.t . oxidizing the Cr to a chromate soluble in water. After
filtering, the precipitate of Fe203 is dissolved in HC1 and tested with KCNS.
The original solution must be tested to determine whether the iron was
present in the ferrous or ferric condition. A portion of the original
solution acidified with HC1 gives blood red color with KCNS if Fe'" is
£126, 10. ntox. 159
present, no color for the Fe". Another portion gives a blue precipitate
with KjFe(CN)6 if Fe" is present, only a brown or green color for the
Fe'" (6&).
0. Estimation. — (1) After oxidation to Fe'" , if necessary, it is precipitated
with irH,OH , dried, ignited to a dull-red heat and weighed as Fe20;, . (2) By
precipitation with nitroso-/?-naphthol in slightly acid solution (Knorre, B., 1887,
20, 283). Volumetrically: (3) As ferrous iron, by titration with a standard
solution of KMnO4: 10FeS04 + 2KMn04 + 8H..SO, = 5Fe,(S04)3 + K,SO4 +
2MnS04 + 8H..O . (4) By titration with a standard solution of K,Cr,07 , using
a solution of K,Fe(CN)0 as an external indicator: 6FeSO4 + K,Cr,67 + 7H™SO4 =
:iFe2(SO4)3 + K2S04 + Cr,(S04)3 + 7H.O . (J) As ferric iron, by titration with
a standard solution of Na.SoO., , using KCNS as an indicator: 2FeCl, + 2Na2S,O3
= 2FeCL + NaJS4O6 + 2NaCl . A few drops of a solution of CuSO4 are added,
which seems to hasten the reaction and gives more accurate results; or use
excess of the Na2S,O;< and titrate back with standard iodine (Crafts, J. €., 1873,
26, 1162). (6) The iron as ferric salt is treated with an excess of a standard
SnCl2 solution, the excess of the SnCL being determined by a standard solution
of iodine in potassium iodide: 2FeCl.," + SnCL = 2FeCL + 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 Na^O,
(very accurate for small amounts of iron). (8) When present in traces it is
determined colorimetrically as Fe(CNS), in etherial solution (Lunge, Z. angew.,
, 669).
10. Oxidation. — 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 An, Ag, CrVI, Mnvn, Mnvl, and H,0., . In presence of some dilute
acid, such as H.,S04 or H3P04 by PbO, , Pb304 , Mn304 , Mn02 , Mn20 . ,
Co.O., , Ni,0, . The following acids also oxidize Fe" to Fe'", HNO, , HIM), ,
HC10 , HC10, , HC10, , HLS04 (if concentrated and hot), HBrO~ , HBrO,
HIO:! , also Br , Cl . Br and Cl in presence of KOH changes Fe" and Fe'"
To K2Fe04 . Barium ferrate is the most stable of the ferrates; they are
strong oxidizers, acting upon nitrites, tart-rates, glycerol, alcohol, ether,
ammonia, etc. (Rose^p. Am. Sor., 1895, 17, 760).
Fe'" is reduced to Fe" by solutions of Sn", Cu',, H,PO, , H3PO, , H,S ,
H.,S03 , Na.,S20, , and HI . Also by nascent hydrogen, or by any of the
?netals which produce hydrogen when treated with acids, including Pb ,
As . Sb , Sn , Bi , Cu *, Cd 1 , Fe , Al , Co , Ni . Zn , and Mg f.
* Carnegie, J. C., 1888, 53, 468. t Warren, C. N., 1889, 60, 187.
160
AXALYXIS OF THE IRON GROUP.
5127.
§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 H...S will cause no pre-
cipitate (§80), and freed from ELS 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, Co) of
Ammonium Hydroxide (§116). If there is a precipitate, filter and wash.
Precipitate: A1(OH)5 , Cr(OH)3 , Fe(OH)3 .
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 />«// for several minutes. If a residue remains, filter
and wash.
Residue: Cr(OH)3, Fe(OH)3 .
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).
Residue: Fe,O3 .
Dissolve the residue in
HC1 and test for iron
with potassium thio-
cyanate (§126, <>b).
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 HC1 and test-
ed for the blood-red
color with KCNS .
Iron being found, to de-
termine whether it is
ferric or ferrous, or
bothf, in the original
solution, test the latter,
after acidulating with
hydrochloric acid, with
KCNS for ferricum,
and with K,Fe(CN)n for
ferrosum (§126. 6ft).
Solution: Na,CrO4 ,
K,Cr04 (Na,CO;i) .
Acidify with HC,.H,O and
precipitate the chro-
mium as lead eliminate
(yellow) with a solu-
tion of lead acetate
(§57. c./i).
If the original solution
contains a chromate it
will be yellow (normal
chromate), or red (acid
chromate), and will
^ give the reactions for
eliminates with
Pb(CH,0) . Bad, .
etc. (§125. oh). If the
chromium is present as
a chromic salt, Cr,(SO,)n,
the solution will have
a green or bluish-green
color and will give the
general reactions as de-
scribed at §125. fi.
Chroma tes should be re-
duced bv boiling with
HC1 and C.H.OH be-
fore proceeding with
the regular course of
analysis (§125. fif).
Study §136, §128, §129,
§130 and §131.
Solution : KA1O , .
Make the solution slight-
ly acid with hydro-
chloric acid, and then
add ammonium car-
bonate. A precipitate
is A1(OH)3.
The same result is ob-
tained with nearly
equal certainty by add-
ing an excess of NH,C1
to the alkaline solution
(§124, <>a: §130).
Lead and antimony give
similar results if
(through carelessness)
they have not been
removed (§131, 6).
Study §136. §128. §129. Study §136. §128, §129.
§130,' §131.' §131, 6, and §124. <,.
* In the filtrate from the Second Group iron is necessarily in th" ferrous condition (1
+ Ferrous salts, which have been kept in tho air.aro novcr wholly freo from ferric compound*.
. DIRECTIONS FOR ANALYSIS WITH NOTES.
DIRECTIONS FOR THE ANALYSIS OF THE METALS OF THE THIRD (TROUP.
§128. Manipulation. — Boil the filtrate from the second group (§80) to
expel the H2S and then oxidize any ferrous iron that may be present by
the addition of a few drops of HNO , continuing the boiling to a clear
straw-colored solution (§126, Gc):
3FeSO, + 4HN03 = Fe,(S04)3 + Fe(N03)3 + NO + 2H,O
Add to the solution about one-half its volume of NH4C1 (56, §§134 and-
189) and warm and then add NH4OH in a decided excess (§135, 6a):
MgCl, + NH4C1 + NH4OH =^H4MgCl, + NH4OH
Fe2(S04)3 + ONH4OH = 2Fe(OH)7Sr B(NH4)2S04
ZnS04 + 4NH4OH = (NH4),ZnO, + (NH4)2S04 + 2HX>
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 H2S is not all expelled, it becomes oxidized by the
HNO3 with deposition of a milky precipitate of sulphur (§257, Ml), which
tends to obscure the reactions following: (iH2S + 4HN03 = 3S"2 + 4NO + 8H2O.
Also any H2S not decomposed by the HN03 would cause a precipitate of the
sulphides of the fourth group upon the addition of the NH4OH: H»S + NiCl, -f-
3NH4OH = NiS + 2NH4C1 + 2H,O .
(2) Any iron that may have been present in the original solution in the
ferric condition is reduced to the ferrous condition by the H,S (§126, 6e) :
4FeCla + 2H2S = 4FeCl2 + S, + 4HC1 . The ferrous hydroxide is not com-
pletely insoluble in the ammonium salts present (§117), and hence unless the
oxidation with the HNO3 be complete, some of* the iron will be found in the
next group.
(3) If considerable iron be present thet solution .becomes nearly black upon
addition of nitric acid, due to the combination of the nitric oxide with the
ferrous iron (§241, 80). Therefore the boiling, and addition of HNO3 , a drop
or two at a time, must be continued until the solution assumes a bright straw
color.
(4) 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, 6«). The careful addition of the nitric acid (avoiding an
m excess) prevents this oxidation of the manganese.
(5) Ammoniiim hydroxide precipitates a portion of Mn (§134, 6a) and Mg
(§189, 6«), but these hydroxides are soluble in NH4C1 (5c, §§134 and 189);
hence if that reagent be added in excess the Mn (§134, 6a) and Mg are not at
all precipitated by the NH4OH:
2MnCl2 + 2NH4OH = Mn(OH), + (NH4)2MnCl4
Mn(OH), + 4NH4C1 = (NH4)2MnCl4 + 2NH4OH
2MgCL + 2NH4OH = Mg(OH)2 + NH4MgCl3 + NH4C1
Mg(OH). + 3NH4C1 = NH4MgCl3 + 2NH4OH
(6) Ammonium chloride lessens the solubility of A1(OH)3 in the NH4OH
solution and effects an almost qiiantitative 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 NH4OH the odor should be noted after shaking the test tube
and after the solution has been heated.
(8) 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 Go).
162 DIRECTIONS FOR ANALYSIS WITHMWKX. >j!29, 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 NH4OH (§125, »;</). Boiling the solution to return »• excess of ammonia
prevents this.
(10) A small portion of the nitrate of the second gro\tp after the removal of
the H.S by boiling should be tested for the presence of phosphates by am-
monium tuolybdate (§75, <></). 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 :
A1(OH):, + NaOH = NaAlO, + 2H2O
Cr(OH)3 + NaOH = NaCrO- + 2H2O (in the cold)
NaCrO, + 2H.O = Cr(OH)3 + 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 KNO, and Na,CO.t: 2Cr(OH)::
+ 2KNO, + eNa,CO, := K2Cr04 + Na,Cr04 -f 2NO + C02 + 3H,0
($125, ?). The fused mass is then dissolved in water, filtered, rendered
acid with acetic acid and tested for chromium with Pb(CJEI.,02)., , a yellow
precipitate at thi« point being sufficient evidence of the presence of
chromium : Na..Cr04 + K.,Cr04 -f 3Pb(C,H:I02). = 2PbCr04 + 2NaC,H ,0,
+ 2KC2H,0, (§57, Gfc).
The residue of the fused mass not soluble in water should be washed
with hot water and then dissolved in HC1 : Fe,0:! -f OHC1 == 2FeCl, -f
3H20, and tested for iron with KCNS : Fed, + 3KCNS == Fe(CNS), +
SXC1 .
If iron has been found to be present, the original solution acidulated
with HC1 (or a few drops of the filtrate from the first group) should 1><
tested with KCNS for the presence of ferric iron (§126,06) and with
K,Fe(CN)0 for the dark blue precipitate of Fe:,(Fe(CN)0), indicating tin-
presence of ferrous iron (§126, 06): 3FeS04 + 2K.Fe(CN)0 = Fe.,(Fe(CN)li) .
-f 3K2S04 .
The alkaline filtrate obtained after boiling the precipitated hydrox-
ides with NaOH , is slightly acidulated with HC1 : KA10, -f- 1HC1 --
A1C13 + KC1 + 2H20, and then precipitated with (NH4),CO, , a white
gelatinous precipitate being evidence of the presence of aluminum:
2A1C1, + 3(NH4),CO, + 3H,0 ~ 2A1(OH)3 + 6NH4C1 + 3CO, . Or an
excess of NH4C1 may be added. directly to the alkaline filtrate, giving the
white gelatinous precipitate of aluminum oxide-hydroxide: 2EA102 -\-
2NH4C1 -f H20 = A1,0(OH)4 + 2KC1 + 2NH, (§124, Ga).
§132, 1. COBALT. KM
§131. Notes. — (/) Chromium hydroxide when precipitated from solutions of
pure chromic salts by NaOH is readily soluble in an excess of the cold reagem
(§125. (''/)• 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 th;- 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 HC1 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, (ifl), alu-
minum hydroxide being a white gelatinous precipitate.
(Ji) If the fused mass has a green color, manganese (§134, 7) is evidently
present 'in large quantities and was not completely separated by the NH4C1
and NH,OH (§134, 6fl). ~By dissolving the fused mass in water and carefully
warming with HC1 . the manganate, KL.MnO4 , may be reduced (a) (§134, .~>r)
without effecting a, rediiction of the chromate, which may be precipitated as
BaCrO4 by BaCL 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 KNO3 and K2CO3 a green mass is
obtained, the operation should be repeated:
K2Mn04 + 8HC1 = MnCL + 2KC1 + 201, + 4H2O
2K,Cr04 + 10HC1 + :iC,H,;O = 2Cr013 + 4KC1 + 3C2H40 + 8H2O
(5) The presence of chromium as chromic salts is usually indicated by the
green or bluish-green color of the original solution. Chromium as chromates
(red or yellow) should be reduced to chromic salts by boiling with HC1 and
C.jHnO before proceeding with the regular group separations (§125, 6e and f).
H..S will effect this reduction but gives also a precipitate of sulphur which
should be avoided when convenient to do so: 2K.,Cr.,O7 + 16HC1 + (>H2S =
4CrCl3 + 4KC1 + 3S3 + 14H.O .
(6) Too much stress cannot be 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 throtigh lack of sufficient HjS or too much HC1 , lead or
antimony are not completely removed in the second group, they will give all
the reactions for aluminum* (§57, fifl, and §70, 6«); 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, 6«, 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 Zinc.
§132. Cobalt. Co = r>9.00 . Usual valence two and three.
1. Properties. — Specific t/mriti/, powder from the oxide reduced by hydrogen,
mean of five samples, 8.957 (TCammelsberg, I'offg., 1849, 78, 93); melting point,
3500° (Pictet, C. r., 1879, 88, 1317). Cobalt is similar to iron in appearance, is
164 COBALT. §132, 2.
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. G'ft., isr>i>,
(3), 4(5, 202). The fine powder oxidi/es in the air quite rapidly and may evert
take fire spontaneously ; in a compact mass it is but little tarnished in moist air.
At a white heat it burns rapidly to Co O . It is attracted by the magnet and
can be made magnetic, retaining (unlike steel) its magnetism at a white heat.
2. Occurrence. — Cobalt does not occur in a free state, except in meteoric
iron. It is found in linnaeite (Co3S4): skutterudite (CoAs3); speiss cobalt
(CoNiFeAs,); glance cobalt (CoFeAsS,): wad (Co.MnO,2MnO, + 4H,O); etc.
3. Preparation. — (/) By electrolysis of the chloride. (2) By heating with
potassium or sodium. (3) By heating any of the oxides, hydroxides or the
Chloride in hydrogen gas. (}) By fusion of the oxalate under powdered glass.
(5) Also reduced by carbon in various ways.
4. Oxides and Hydroxides. — ('oltaltoug oxide, CoO , is made (/) by heating
any of its oxides or hydroxides in hydrogen to (not above) 350°; (^) by ignition
of Co(OH) = or CoC03 . air being excluded: (3) by heating Co3O4 to redness in
a, stream of CO, (Hussell, J. C., 1863, 16, r>l): (.'/) by heating any of the higher
oxides to a white heat (Moissan, A. Ch., 1880, (5), 21, 2^2). Col'xiltoux Iti/dntJ'Mc
is made from cobaltous salts by precipilation with fixed alkalis; oxidizes if
exposed to the air (6«). The most stable oxide is the calHiltono-cobaltir (Co,O4)
tricolxilt tetrofidc', 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, e.g., Co30,(OH),. Co:,O(OH),,, Co3O:i(OH).,.
Cobaltic oj-idi; Co,©, , is made by heating the nitrate just hot enough for de-
composition, but not hot enough to form Co30, . Cohnltic 1ii/dni.ridr, Co(OH), ,
is made by treating any cobaltous salt with Cl , HC10 , Br or I in presence of
a fixed alkali or alkali carbonate. It dissolves in HC1 with evolution of chlo-
rine, in H SO, with evolution of oxygen, forming a cobaltons salt. CoO, has
not yet been isolated, but MeConnell and Banes (./. f., 1897, 71, 584) have
shown that it exists as H CoO and in certain cobaltites.
5. Solubilities.—*/. — .Vr/«/.--Slowly soluble on warming in dilute HC1 or
IL.SO, , more rapidly in HNOn . not oxidi/ed on exposure to the air or when
heated in contact with alkalis. Like iron, it may exist in a passive form
(Nickles, J. ;>r., 1854, 61, 168; St. Edme, T. r., 1889, 109, "04). With the halogens
it forms cobaltous compounds (Hartley, J. C., 1874, 27, 501). ft. — O.ridex 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, A., 1886, 232. :!24) ;
the various higher oxides and hydroxides are insoluble in ammonium hydroxide
or chloride (separation from niekelous hydroxide after treating with iodine
in alkaline mixture) (Donath, Z., 1881, 20, 38(>), and are decomposed by acids,
evolving oxygen with non-reducing acids, or a halogen from the halogen acids,
and forming cobaltous salts. CO;,O4 is said to be soluble in acids with great diffi-
culty (Ciibbs and Geuth, Am. 8., 1857, (2), 23. 257). c.—Kalt*.— Cobalt forms two
classes of salts: roltaltmift, derived from CoO , and <-nlmlti<; from Co,O, . 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 crystallization
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 ilclii/iir^rnt: chloride, hygroscopic; sulphate,
efflorescent. The chloride vaporizes, undecomposed, 'at a high temperature.
The carbonate, sulphide, phosphate, borate, oxalate. cyanide, ferroeyanidc
and ferricyanide are iiixohililc in water. The potassimn-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 nitrate;
ether dissolves the chloride, sparingly, more so if the ether be saturated with
HC1 gas (separation from Ni) (Pinerua, C. r.. 1897. 124, 802). Most of the
salts insoluble in water form soluble compounds with ammonium hydroxide.
§132, 66. COBALT. 165
G. Reactions, 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(OH)2 .
The cobaltous hydroxide is not soluble in excess of the reagent, but is
somewhat soluble in a hot concentrated solution of KOH (distinction from
Ni) (Eeichel, Z., 1880, 19, 468). Freshly precipitated Pb(OH)2 , Zn(OH).,,
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,
different soluble compounds noted for their bright colors, as (NH3)4CoCl2 ,
(NH3)6CoCl2,(NH3)4CoCl3,etc.
Alkali carbonates precipitate cobaltous basic-carbonate, Co8Or,(C03)3 ,
peach-red, which when boiled loses carbonic 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. Carbonates of Ba , Sr , Ca , or Mg do not precipitate cobaltous
chloride or nitrate in the cold (separation from Fe"', Al, and Cr'"), 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 BaCO,
(distinction from Ni).
6. — Oxalic acid and oxalates precipitate reddish-white cobaltous oxaldte,
CoC204 , soluble in mineral acids and in ammonium hydroxide.
Alkali cyanides — as KCN — precipitate the brownish-white cobaltous
cyanide, Co(CN)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 — (KCN).,Co(CN)2 — potassium cobaltous cyanide, the solu-
tion having a brown color: CoCl2 + 2KCN = Co(CN)2 -f- 2KC1 . Then
Co(CN)2 + 2KCN = (KCN)2Co(CN)2 . Dilute acids, without digestion,
reprecipitate cobaltous cyanide from this solution (the same as with Ni) :
(KCN),Co(CN)2 -f 2HC1 = Co(CN)2 -f- 2HCN + 2KC1 . But if the solu-
tion, with excess of the alkali cyanide and with a drop or two of hydro-
chloric acfd,* insuring free HCN , be now digested hot for some time, the
* Moore (C. JV., 1887, 56, 3) adds glacial phosphoric acid to the neutral solutions of cobalt and
nickel, until the precipitate fi^st formed begins to redissolve ; then he adds KCN and boils,
continuing the boiling and addition of KCN 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 leaving the cobalt in solution. See al-o Hambly (C. N., 1892, 63, 289).
166 COBALT. S132, tjr.
cobaltous cyanide is oxidized and converted into alkali cobalticyanide -us
K3Co(CN)0 — corresponding to ferricyanides, but haruuf no corresponding
nickel compound :
4Co(CN)2 + 4HCN + 0, = 4Co(CN), (eobaltic cyanide) + 2H.O
Co(CN)s + 3KCN = K3Co(CN)0 (potassium cobalticyanide).
In the latter solution acids cause no precipitate (important distinction from
nickel, whose solution remains (KCN).,Ni(CN).> , 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.t
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:
<5Co(CN), + 24KCN + 2Cr03 + 3H.O = GK3Co(CH),. + Cr,0:, -4- 6KOH
(McCulloch, C. N., 1889, 59, 51).
Ferrocyanides, as K,Fe(CN)B , precipitate titbaltou* fcrrotwtnidr-, Co=Fe(CN). ,
gray-green, insoluble in acids. Ferricyanides, as KsFe(CN)a , precipitate cobalt-
uux ferriciianidr, Co?(Fe(CN)6). , brownish-red, insoluble in acids. But a more
distinctive test is made by adding ammonium chlwiilc 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, ('•!>).
Nitroso-^-naphthol completely precipitates solutions of Cu , Fe , and Co ;
Ag . Sn . and Bi salts are partially precipitated; and Pb , Hg , As , Sb , Cd .
Al , Cr . Mn . Ni . Zn . Ca , Mg , and Gl remain in solution (Burgass, Z.
anyett'., 189(1, 590). In analysis for the separation of cobalt and nickel it is
recommended to proceed as follows: The rnetals 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 HC1, then with hot 12 per cent HC1 , 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, B., 1887, 20, 283 and Z. angew., 1893, 264).
c. — Potassium nitrite forms with both cobaltous and nickelous salts the
double nitritt-s, Co (N02)2.2KN02 and Ni(N02)2.2KNO, , 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 eobaltic nitrite, Co(N02)3.3KN02 (separation from
Ni): CoCl. + GKNO, + HC2H,02 + HN02 = Co(N02)3.3KN02 + 2KC1 -f
KC^O, -f H,0 + NO .
§132, 7. COBALT. 167
</. — Phosphates, as Na^HPC^ , precipitate cobaltous salts as Iho reddish
eobaltoitK phosphate, CoHPO4 , soluble in acids and in ammonium hydroxide.
Sodium pyrophosphate forms a gelatinous precipitate with solutions of cobalt
salts, soluble in excess of the reagent. The addition of acetic acid causes u
precipitation of the cobalt even in the presence of tartrates (separation from
Ni, 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 phosphate
and hydrochloric acid, and when hot treated with an excess of ammonium
hydroxide, a bluish precipitate of CoNH4PO4 will appear on stirring (separa-
tion from nickel *) (Clark, C. N., 1883, 48, 262; Hope, J. Soc. Ind., 1890, 9, 375).
e. — Hydrosulphuric 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 thiosulphate 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 (HC1 does not reduce
the cobalt in K3Co(CN)6).
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. — KMnO4 added to an ammoniacal solution of cobaltous salts
oxidizes the cobalt and prevents its precipitation by KOH (separation from
Ni) (Delvaux, C. r., 1881, 92, 723).
/. — Cobaltous salts in ammoniacal solution, warmed with H2O2 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 Uue 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, nickels or iron interfere somewhat. By
ignition, with sodium carbonate on chaiWal or with the reducing flame,
"V\!
* Krauss (Z., 1891, SO, 327) gives a grood review of the most important methods for the separa-
tion of cobalt and nickel.
168 NICKEL, j §132, 8.
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 off the light of yellow flames (§205, 7). The
black cobaltoso-cobaltic oxide, Co304 , 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 H2S in ammoniacal solution with Ni , Mn and Zn .
The sulphides are digested with cold dilute HC1 which dissolves the Mn
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. HC1 , using a
few drops of HNO., . 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 fib, and further identified by the bead test.
9. Estimation. — (1) As metallic cobalt, nil compounds that may he reduced
by ignition in hydrogen gas, r. </., CoCl , Co(NO3), . CoCO , 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 CoiOHi by precipitation with a fixed alkali, ahd 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 oxidi/ed, and
then again ignited in CO. . (3) After converting into a sxilphate it is ignited
at a dull-red heat and weighed as a sulphate. (}) After converting into the
oxalate, titrated with KMnO, . (.>) In presence of nickel, it is oxidized in
alkaline solution by H.O2 , KI and HC1 are added, and the liberated iodine
titrated with sodium tHosulphate (Fischer, C. C., 1889, 116). (6) Electroly-
tically. (7) Separated frcra nickel by nitroso- i-naph.tb.ol, and after ignition
in hydrogen weighed as the metal (6ft).
10. Oxidation. — Co" is oxidized to Co'" in presence of a fixed alkali by
Pb02 , Cl , KC10 , Br , KBrO , I and H20,* ; in presence of acetic acid by
KN02 (Or). Co'" is reduced to Co" by H~C204 , H.,P02 , H,.S , H2SO:! , HC1 ,
HBr , and HI . Metallic cobalt is precipitated from solution of CoCl2 by
Zn, Cd. and Mg.*
§133. Nickel. Ni = 58.70 . Usual valence two and three.
1. Properties.— RpeHfic urfiriti/, 8.9 (Schroeder, Po{ig., 1859, 106, 226). Melting
paint, H50° (I'ictet, 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 incandescence when heated in O , Cl , Br or S . It is much
*Durrant, C. .V., 1897, 75, 43.
§133, Qa. NICKEL. 169
used in plating other metals, in making coins of small denominations, in
hardening armor plate, projectiles, etc. The presence 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, NiS2; as nickel blende, NiS; as iron nickel blende,
NiFeS; as cobalt nickel pyrites, (NiCoFe)3S4 , etc.
3. Preparation. — (Jf) By electrolysis. (2) By heating in a stream of hydrogen.
The oxide is reduced in this manner at 270° (W. Miiller, Pogg., 1869, 13Q, 51).
(3) By fusing the oxalate under powdered glass (CO2 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(CO)4 to 200°.
4. Oxides and Hydroxides. — Xld'Clmis oxide is formed when the carbonate,
nitrate, or any of its oxides or hydroxides are strongly ig-nited. Nickeloux
hydroxide is formed by precipitation of nickelous salts with fixed alkalis.
Nickclic oxide, Ni2O3 , is made from NiC03 , l\Ti(NO.,)2 or NiO by heating in the
air not quite to redness, with constant stirring. It is changed to NiO at a red
heat. Nickelic hydroxide, Ni(OH)3 , is formed by treating nickelous salts
first with a fixed alkali hydroxide or carbonate and then with Cl , NaCIO , 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,
Ni3O4 , magnetic (corresponding to Co3Ot , Fe3O4 , Mn304 and Pb3O4), is formed,
according to Baubigny (C. r., 1S7S, 87, 1082), by heating NiCL, in oxygen gas
at from 350° to 440°; and by heating Ni.O3 in hydrogen at 190° (Moissan, A. (Th.,
1880, (5), 21, 199).
5. Solubilities.— a. — Metal. — Hydrochloric or sulphuric acid, dilute or con-
centrated, attacks nickel but slowly (Tissier, C. r., I860, 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. It. — Oxides and
hydroxides. — Nickelous oxide and hydroxide, are insoluble in water or fixed
alkalis, soluble in ammonium hydroxide and in acids. Nickelic oxides and
hydroxides are dissolved by acids with reduction to nickelous salts, with halogen
acids the corresponding halogens are liberated. The moist nickelic hydroxide,
formed by the action of Cl , Br , etc., in alkaline solution, after washing with
hot water liberates free iodine from potassium iodide (distinction from cobalt).
Nickelic hydroxide when treated with dilute sulphuric acid forms NiSO4 ,
oxygen being evolved. With nitric acid the action is similar, distinction from
cobaltic hydroxide, which requires a more concentrated acid to effect a similar
reduction, c. — Raits. — 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 alcoJjL and the nitrate in
dilute alcohol. Most salts of nickel form soluble co^Bunds by action of
ammonium hydroxide.
6. Reactions, a. — Alkali hydroxides precipitate solutions of nickel
salts as nickel hydroxide, Ni(OH)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, forming 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 liquid, sp.
gr., 1.3185, boiling at 43° and freezing at -25°. When heitcd t > 200° it is decomposed into >"1 and
CO (Berthelot, C. r., 1891, 112, 1343; 113, 679; Mond, J. Soc. Iiid., 1892, 11, 750).
170 NICKEL. §133, M.
will slowly precipitate nickel hydroxide from the ammoniacal solutions
(distinction from cobalt). Alkali carbonates precipitate green basic
nickelous carbonate, Ni-(OH)0(C03)2 (composition not constant), soluble in
ammonium hydroxide or ammonium salts, with blue or greenish-blue color.
Carbonates of Ba , Sr , Ca , and Mg 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.
b. — Oxalic acid and ozalates precipitate, very slowly but almost completely,
after twenty-four hours, nickel oxalate, green. Alkali cyanides, as KCN , pre-
cipitate -nickfl cyanide, Ni(CN), , 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 correspond* exactly to that for
cobalt (§132, fib); and the solution of double cyanide is reprecipitated as
Ni(CN)2 by a careful addition of acids (like coba'lt); 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), . Ferrocyanides, ;is K4Fe(CN)« , precipitate a
greenish-white nickel ferroci/diiide, NLFe(CN),., , insoluble in acids, soluble in
ammonium hydroxide, decomposed by fixed alkalis. Ferricyanides precipitate
greenish-yellow nickel fcrricjurnidc, insoluble in acids, soluble in ammonium
hydroxide to a green solution (§132, Ob). A solution of nitroferricyanide
precipitates solutions of cobalt and nickel salts, the latter being soluble in
dilute ammonium hydroxide (Cavalli, Gazzetta, 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 (NH4),S (Phipson, C. N., 1877, 36, 150).
The xanthate also precipitates nickel in alkaline solution in presence of
Na4P.,O7 (a separation from Fe'") (Campbell and Andrews, J. Am. Soc., 1895,
17, 125).
Nickel xaltx are not precipitated by ati acetic acid solution of nitroso-/?-
naphthol (separation from cobalt) (Knorre, /?., 1885, 18, 702).
c. — Potassium nitrite in presence of acetic acid does not oxidize nickelous
compounds (distinction from cobalt). (I.- — Sodium phosphate, Na..HPO4 , pre-
cipitates nickel phosphate, Ni3(PO4), , greenish-white.
f. — Hydrosulphuric acid precipitates from neutral solutions of nickel
salts a portion of the nickel as nickel siilpliide, black (Baubigny, C. r.. 1882,
94, 1 183: 95, 34). The precipitation takes place slowly, and from nickel-
ous acetate is coii^fete. In the presence of mineral acids no precipita-
tion takes place. ^Pkali 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 sulphide. 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 solutions of nickel and cobalt, and an
excess of sodium hydroxide. He then passes in 11 3S. The cohult is completely precipitated
while the nickel remains in solution, and can be precipitated upon acidulating1 the nitrate.
£133, 10. NICKEL. 171
sulphur) on addition of acetic acid (distinction from cobalt). Freshly pre-
cipitated nickel sulphide is soluble in KCN and reprecipitated as Ni(CN)2 on
adding- HC1 or H2SO4 (separation from cobalt) (Guyard, BL, 1876, (2), 25, 509).
When nickel salts are boiled with a solution of Na2S2O3 , a portion of the nickel
is precipitated as the black sulphide.
f. — The halogen acids reduce the higher oxides of nickel to nickelous
salts with liberation of the corresponding halogen. Potassium iodide
added to freshly precipitated nickelic hydroxide gives free iodine (distinc-
tion from cobalt).
0. — Nickel salts are precipitated by arsenites and arsenates, white or green-
ish-white, soluble in acids, including1 arsenic acid. ft. — Potassium chromate
precipitates basic nickel chromate, yellow, soluble in acids, including chromic
acid (Schmidt, A., 1870, 156, 19). K,Cr2O7 forms no precipitate.
7. Ignition. — Nickel compounds dissolve clear in the borax bead, giving with
the oxidizing flame 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 finally 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 microcosmic 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 and 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 H..S (after neutralizing with NH4OH), and its presence con-
firmed by the usual tests. Or dissolve the sulphides of Ni and Co in
HNO.j , evaporate nearly to dryness, add an excess of KOH or Na2C03 ,
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 CS2) is evidence of
the presence of nickel.
0. Estimation.* — (/) Nickel hydroxide, oxide, carbonate or nitrate is ignited
at a white heat and weighed as NiO . (2) It is converted into the sulphate and
deposited on platinum as the free metal by the electiic* current. (3) Volu>
metrically, fty titration in a slightly alkaline solution with KCN , using a
small' amount of freshly precipitated Agl as an indicator (Campbell and,
Andrews, J. Am. Koc., 1805, 17, 127).
10. Oxidation. — Ni" is changed to Ni'" in presence of fixed alkalis by
Cl, NaCIO , Br, and NaBrO (not by I, distinction from cobalt, Donath,
B., 1879, 12, 1868). Ni'" is reduced to Ni" by all non-reducing acids with
evolution of oxygen; by reducing acids, HoC204 is oxidized to C02 , HNOa
•Goulal (Z. aiujcic., 1W»8, 177) gives a summary of the methods proposed for the volumetric
estimation of nickel.
172 MANGANESE. §13*, 1.
to HN03 , H3P02 to H3P04 , H2S to S , H,S03 to H2S04 , HC1 to Cl , HBr to
Br , HI to I , HCNS to HCN and H2S04 , H4Fe(CN)(i to H3Fe(CN)8 . Ni"
is reduced to the metal by finely divided Zn . Cd . and Sn .
§134. Manganese. Mn •=. 55.0 . Valence two, three, four, six and
seven.
1. Properties.— Specific gravity, 7.138 to 7.206 (Brunner, Pogg., 1857, 101, 264);
melting point, at a high white heat (blue heat) (Deville, A. Ch., 1856, (3), 46,
199); volatilizes at the highest heat of the blast furnace (Jordan, C. r., 1878,
86, 1374). 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
ordinary temperature (Deville, /. <•.). 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 manganates 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, Mn,_.O3 ; hausman-
nite, Mn.,04: manganite, MnO(OH): manganese spar, MnCO3; mangam.e
blende, MnS: and as a constituent of many other minerals.
3. Preparation, — (/) By electrolysis of the chloride. (2) By reduction with
metallic sodium or magnesium (Cilatzel, /*.. 1S89, 22, 2857). (3) By reduction
with some form of carbon. It has not been reduced by hydrogen. (4) By
ignition with aluminum ((ioldschmidt, .4., ISDN, 301, 19).
4. Oxides and Hydroxides.- («) Manganous oxide, MnO . represents the only
base capable, of forming stable manganese salts. It is formed (/) by simple
ignition of Mn(OH), , MnCO or MnC..O, . air being excluded; (2) by ignition
of any of the higher oxides of mangane.se with hydrogen in a closed tube
(Moissan, .1. Ch., 1880, (5), 21, 199). If prepared at as low a temperature as
practicable, it is a dark gray c;r greenish-gray powder, and oxidizes quickly
in the air to Mn:,0, . If prepaied at a higher heat it is more stable. Man-
ganous hydroxide, Mn(OH), , is formed from manganous salts by precipita-
tion with alkalis. It quickly oxidizes in the air. forming MnO(OH), thus
changing from white to brown. (I) Manganic oxide, Mn,03 , is formed by
heating any of the oxides or hydroxides to a nd heat in oxygen ij;is or in air
(Schnieder,' POM., 18")9, 107, f>6:>). Manganic oxide-hydroxide. MnO(OH) , is
formed (/) by oxidation of Mn(OH), in the air; (^) by treating MnO., with
concentrated H SO, at a temperature of about l."0°, forming Mn.. (S0,)3 and
then adding water: Mn,(SO4), + 4H.O == 2MnO(OH) -f :iH,SO4 (Carius, A..
1856, 98, (53). (P) 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 v\onld be Mn:,(OH)s: this has not been isolated.
A corresponding OXide-hydroxide is formed bv adding freshlv formed and
moist MnO, to an exeew of MnCl, contain ^ NH.C1 (Otto. .1..' is:,.-), 93, 372).
(d) Manganese peroxide. MnO, , is formed (/) by heating Mn(NO^)., to 200°
(Gorgeu, C. r., 1879, 88, 79(5): (>) by heating MnCO:, with KC1O:> to ::00°; (.?) bv
boiling any manganous salt with concentrated HNOS and KC1O3 . A correspond-
ing hydroxide, Mn(OH)4 , has not been isolated. Several other hydroxides.
e.g., MnO(OH),. Mn.OJOH), . Mn,O,(OH)4 etc., have been produced. The
chief use of manganese dioxide is in the preparation of chlorine or bromine.
(e) Manganates. — .l/ffj?.w»i> acid. H2Mn04 , is not known in a free state. The
corresponding salt. K..,MnO4 , is formed when any form of manganese is fused
with KOH or K.CO, (/) in the air, oxygen being absorbed: or (:i) with KNO:;
or KC10, . NO or KC1 being formed.' A-manganate of the alkali metals is
soluble in water. jr//7i f/radniil (IcroniiH'xition into manganese dioxide and per-
manganates: aK.MnO, + 2H,0 = 2KMnO, + MnO, + 4KOH . Free alkali
§134, 5c. MANGANESE. 173
retards, and free acids and boiling- promote, this change. Manganates have
a green color, which turns to the red' of permanganates during the decomposi-
tion inevitable in solution. This is the usual method of manufacturing KMnO4 .
(f) Permanganic acid is not- in use as an acid, but is represented by the per-
manganates, as KMnO4 . The permanganic acid radical is at once decomposed
by addition of hot HuS04 to a solid permanganate (/), but in water solution
this decomposition does not at once take place, except by contact with oxidiz-
able substances. The oxidizing power of permanganates extends to a yreat
number of substances, possesses different characteristics in acid and in alka-
line solutions, and acts in many cases so rapidly as» to be violently explosive.
The reactions with ferrous salts (.2) and with oxalic acid (3) are much used in
volumetric analysis.
(1) 4KMnO4 + 2HoS04 := 2K2SO4 + 4MnO, + 3O2 + 2H2O
and 2MnO2 + 2H2SO4 = 2MnS04 + 2H20 -f 0,
or 4KMn04 + CH2S04 = ;MnS04 + 2K2SO4 + 50, + 6H2O
(2) KMn04 + SFeCL + SEttl = MnCL + KC1 + sPeCl, + 4H2O
(3) 2KMn04 + 5H2C2O4 + 6HC1 = 2MnCL + 2KC1 + 8H2O + 10CO3
5. Solubilities. — n. — Mvtal. — Manganese dissolves readily in dilute acids to
form manganous salts. Concentrated H2SO4 dissolves it only on warming, S02
being evolved. It combines readily with chlorine arid bromine. 1>. — Oxides
and liydrofid'.'x. — All oxides and hydroxides of manganese are insoluble in
water. They are soluble, upon warming, in hydrochloric acid, forming- man-
ganous chloride; the higher oxides and hydroxides being reduced with evohi-
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 HC1). In the cold, hydrochloric
acid dissolves MnO2 to a greenish-brown solution, containing, probably. MnCl:i
or MnCl4 , unstable, giving chlorine when warmed and forming MnO2 when
strongly diluted with water (Pickering, J. C., 1879, 35, 654; Nickles, A. Ch.,
1865, (4), 5, 161). Nitric and sulphuric acids dissolve manganous oxide and
hydroxide to manganous salts. Manganese dioxide (or hydrated oxide) is
insoluble in nitric acid, dilute or concentrated: concentrated sulphuric acid
with heat decomposes it, evolving oxygen and forming manganous sulphate:
2MnO2 + 2H2SO4 — 2MnS04 + 2H,O + O2 . Manganous hydroxide is insoluble
in the alkalis but soluble in solutions of ammonium snlts.
c. — Salts. — Manganous sulphide, carbonate, phosphate, oxalate, borate,
and sulphite are insoluble in water, readily soluble in dilute acids. Man-
ganic salts are somewhat unstable compounds, of a reddish-brown or
purple-red color, becoming paler and of lighter tint in reduction to the
manganous combination. MnCL and MnS04 are deliquescent. Man-
ganic chloride, MnCl3 , exists only in solution, which is reduced to
MnCL, by boiling, also by evaporation to a solid. Manganic sulphalf,
Mn2(S04)3 , is soluble in dilute sulphuric acid, but is reduced to MnS04 by
the attempt to dissolve it in water alone; potassium manganic sulphate
and other manganic alums are also decomposed by water. Alkali rnangan-
ates and permanganates are soluble in water, the former rapidly changing
to manganese dioxide and permanganate, which is much more stable in
solution. In presence of reducing agents both manganates and perman-
ganates are reduced to lower forms.
K2MnO4 + 8HC1 =: MnCl2 + 2KC1 + 2C12 + 4H2O
2KMn04 + 3MnSO4 + 2H2O = 5Mn02 -f K2SO4 + 2H2SO4
174 MANGANESE. §134, Ga.
Concentrated H2S04 in the cold dissolves KMnO , , forming (Mn03)2S04
(a sulphate of the heptad manganese : 2KMn04 -f 3H,S04 = (Mn03)2S04 -f-
3KHS04 -f 2H20 (Franke, J. pr., 1887, 36, 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(OH)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 * (1). Ammonium 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 (3) (separation of manganese from the metals of the third
group) (Pickering, J. C., 1879, 35, 672; Langbein, Z., 1887, 26, 731).
Manganic 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 manganouft 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 (106).
(/) MnCOH), + ^NH.Cl = MnCl,.2NH4Cl -f 2NH4OH
(2) 2MnSO4 + 2NH4OH = MnS04.(NH4),S04 + Mn(OH)z
(3) MnClz + 2NH4C1 = MnCl^NH.Cl or (NH4),MnCl«
&. — Oxalic acid and alkaline oxalates precipitate manganous 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 o.ralnli'x in presence of some mineral acid:
+ r,H,C,04 .+ 3HL,S04 = K,S04 + 2MnS04 + 10CO, + SH20 .
Soluble cyanides, ;is KCN , .precipitate iiKini/aiioHx <•//»/«/</<-, Mn(CN), , white,
but darkening- in the air; soluble in excess of the precipitant by formation of
double cyanides, as Mn(CN),.2KCN . This solution, exposed to the air. pro-
duces inunudiiici/inndt'N (analogous to ferricyanides), with oxidation of the
*It has IwH'ii ijurstioncd whether the solubility of Mn(OH)2 in ammonium salts is rtuo to com-
bination between the two. As has boon already stated, the Law of Mass- Action causes that
reaction to take place which leads to the formation of a slightly dissociated substance. Thus
Fe(OII):, dissolves in HC1 and A»...O3 in NnOII because in each caso water, a non-dissociated
substance, results; and FeS and A*»S., dissolve in HCI and NaHS respectively because the
little-dissociated II2S is a product. Similarly, Bf H4C1 with Mn(OH!2 gives opportunity for the
formation of NH4OH, a compound of small dissociation-constant. Solution due to this cauco
can take place only with hydroxides having a comparatively large solubility-product i§45)_
Sce Ostwald on the solubility of Mg(OH)», '• Wissenschaftliche Grundlagen der analytisclum
Chemie," 2d ed.. i>. 133.
§134, 6e. MANGANESE. 175
manganese: 12(Mn(CN)2.2KCN) + 3O2 + 2H.O = 8K3Mn(CN)6 + 4MnO(OH).
Fe'" 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 MnO. and the iron remains in solution (Beilstein and Jawein, B., 1879, 12,
1528). Ferrocyanides piecipitate white manganous lerrocyanlde, Mn2Fe(CN), ,
soluble in hydrochloric acid. Ferricyanides precipitate brown manganous fcrrl-
vyaniilc, Mn3(Fe(CN)0)2 , insoluble in acids (separation, with Co and Ni , from
Zn) (Tarugi, Gazzctta, 1895, 25, ii, 478). If an alkali or alkali carbonate be
present, potassium ferricyanide oxidi/es 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 Pb02 ,
KC103 , etc., to oxidize manganous compounds to manganese dioxide or to
permanganic acid. Reducing agents as HC1 , etc., should be absent.
Sulphuric acid may be used instead of nitric acid.
2Mn(NO3);, + 5PbO, + GHN03 = 2HMn04 -f 5Pb(N03)2 + 2H2O
5MnSO4 -f 2KC103 + H2S04 + 4H.O = 5TO.nO, + K2SO4 + 01, + 5H2SO4
In using Pb02 and HN03 to detect manganese, the compound should first
be reduced with hydrochloric acid, precipitated with potassium hydroxide
and this precipitate dissolved in nitric acid, as Mn02 is not all oxidized
by Pb02 and HNO, (Koninck, Z. angeu\, 1889, 4).
<£. — Hypophosphorous acid reduces all higher forms of manganese to the
manganous condition. Alkali phosphates, as Na2HP04 , precipitate, from
neutral solutions of manganous salts, normal manganous phosphate, Mn;;(PO4)2 ,
white, slightly soluble in water, and soluble in dilute acids. It turns brown in
the air. The manganous hydrogen phosphate, MnHPO4 , is more soluble in
water, and is obtained by crystallization from a mixture of manganous sul-
phate acidulated with acetic acid and disodium phosphate, Na,HPO4 , added
till a precipitate begins to form. From the ammonium-manganese solution,
freshly formed (6a), phosphates precipitate all the manganese as ina<nfiannu»
ammonium phosphate.
e. — Hydrosulphuric 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'.
Ammonium sulphide precipitates from neutral solutions, anfl forms from
the recent hydroxide of mixtures made alkaline, the flesh-colored man-
ganous sulphide, MnS . Acetic acid, acting on the precipitated sulphides,
separates manganese from cobalt and nickel, and from the greater part of
zinc. All the higher oxidized forms of manganese (in solution or freshly
precipitated) are reduced to the manganous condition, with separation of
sulphur (10), by hydrosulphuric acid or soluble sulphides: 4KMn04 -f-
14(NH4),S + 16H26 == 4MnS -f 4KOH -f S8NH4OH + 5S, . The green
manganous sulphide, MnS , crystalline, anhydrous, is formed by the action
of H2S on a hot ammoniacal manganous solution not containing an excels
of ammonium salts (Meineke, Z. antjeiv., 1888, 3).
Soluble sulphites precipitate from solutions of manganous salts, manganous
sulphite, MnSO , white, insoluble in water, soluble in acids (Gorgeu, C. r.,
1883, 96. 341). Solutions of manganates or permanganates are immediately
176 .u. LYfM.v /•>•/•;. §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.
f. — HC1 , 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 HC1 without heat,
the dark brownish colored solution is said to consist of manganese tetra-
chloride, MnCl, , 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 iodato
(distinction from chloride and bromide). Potassium chlorate or bromate
when boiled with concentrated nitric or sulphuric acids and manganous
compounds forms manganese dioxide (c).
!l. — Soluble arsenites precipitate inannnnoiia arscnitc, and arsenates precipitate
miinwnioitx nrnriintr, insoluble in water, soluble in acids. Arsenous acid and
arsenites reduce solutions of manganates or permanganates, forming a brown
flocculent precipitate: or a colorless solution if warmed in presence of a
mineral acid. /». — Normal potassium chromate precipitates manganous salts,
brown, soluble in acids and in ammonium hydroxide; no precipitate is formed
with potassium dichromate. i. — Soluble manganates and permanganates pre-
cipitate mangnnous salts as manganese dioxide, being themselves reduced to
the same form: : MnSO, + :2KMnO, + rJH.O = r>MnO, -f K,SO, + 2H,SO4 .
7. Ignition with alkali and oxidizing agents, forming a briyhi green mass
of alkaline manyanate, 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 test 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 (fc). 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) ^,Mn(OH)a + 4KNO3 + Na.CO, =
2K,MnO4 -f- Na,Mn04 + 4NO + CO. + 3H,O
(6) Mn(OH)2 + 2KOH + O2 = K,Mn04 + 2H,O
(c) Mn(OH)2 + Na.CO, + O, = Na,Mn04 + H2O + CO,
With beads of borax and microcosmic salt, before the outer blow-pipe flame,
manganese colors the bead violet while hot, and (inicthi/xt-n'd when cold. The
color is due to the formation of iii(iii(/iinic u.riili', the coloring material of the
amethyst and other minerals, and is slowly destroyed by application of the
inner flame, which reduces the manganic to manganous oxide.
8. Detection.— After the removal of the metals of the first three group-
(the third group in the presence of NH4C1 in excess, ">fr and 6«), the Mn
£134, 10ft. MANGANESE. 177
with Co , Ni and Zn is precipitated in the ammoniacal solution by H2S .
By digestion in cold dilute HC1 the sulphides of Mn and Zn are dissolved,
and after boiling to remove the H2S , Mn is precipitated as the hydroxide
by excess of KOH , which dissolves the Zn . The precipitate of the man-
ganese is dissolved in HN03 and boiled with more HN03 and an excess of
Pb02 . A violet-colored solution is evidence of the presence of manganese.
9. Estimation.— (1) By converting- into Mn: 0, (4c), and weighing as such.
(2) By precipitating as MnNH4P04 , and after ignition weighing as Mn2P2O7 .
(3) By treating the neutral manganous salt with a solution of KMnO4 of
known strength (6t). If some ZnSO, is added the action is more satisfactory
(Wright and Menke, /. C., 1880, 37, 42). (//) By boiling the manganous com-
pound with Pb02 and HNO3 , 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 KC1O3 and HNO3 . This is
then reduced by an excess of standard H202 , H2C2O4 or FeSO4 , and the excess
of the reagent estimated by the usual methods. (6) MnO.. , obtained as in (5),
is treated with H2C2O4 and the evolved CO, measured or weighed. (7) MnO2 ,
obtained as in (5), is boiled with HC1 and the evolved Cl estimated.
10. Oxidation.— (a) Mn" is oxidized to Mn'" in alkaline mixture on
exposure to the air; to Mn™ in neutral solution by K2Mn04 and KMn04 ,
in alkaline mixture by Cl , Br , I , K3Fe(CN)0 , KC10 , KBrO , H202S etc. ;
in acid solution by boiling with concentrated HN03 or H2S04 , and KC103
or KBrO., . MnVI~n is oxidized to MnVI by fusion with an alkali and an
oxidizing agent, or by fusion with KC103 alone (Boettger, Z., 1872, 11,
433). Mnvn-n is oxidized to Mnvu by warming with PbO, or Pb,,04 and
HN03 or H2S04 . 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 HN03 or H2S04 , 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. (&) All compounds of man-
ganese having a higher degree of oxidation than the dyad, (Mn"+n) are
reduced to the dyad (Mn") by H,C204 , HH2P02 , H2S4, K2S , H2SO:! , H2022
(in neutral or alkaline solution to MnIV), HC1 , HBr , HI , HCNS , Hg', Sn",
As"', Sb'", Cn', Fe", Cr", Cr'", etc.; the reducing agents becoming respec-
tively C02 , Pv, S° to SVI (depending upon the temperature, concentration,
and the agent used in excess), Cl , Br , I , HCN and SVI, Hg", SnIV, Asv, Sbv,
Cu", Fe'", and CrVI. MnIV+n is reduced to MnIV (or Mn'") by H 3, AsH,3,
SbH33, PH33, Na,S034, Na2S2034, NH4OH3 (slowly), Mn", etc. KMn04 is
reduced to K2Mn04 on boiling with concentrated KOH : 4KMn04 -f 4KOH
= 4K2Mn04 -f 2H20 + 0, (Rammelsberg, B., 1875, 8, 232).
'Klein, Arch. Pharm., 1889, 227, 77; Jarmaesch and von Cloodt, Z. anwrg., 1895, 10, 398 and 410;
Carnot, C. r., 1888, 1O7, 997 and 1150.
2Carnot, HI., 1889, (3), 1,277; Gorgeu,C. r., 1890, 110, 958. 3 Jones, J. C1., 1878, 33, 96. 4 Hoenig
and Zatzck, M., 1883, 4, 738 ; Glaeser, M., 1835, 6, 329.
178 ZINC. §135, 1.
§135. Zinc. Zn = 65.4 . Valence two.
1. Properties.— Specific gravity, 7.142 (Spring, B., 1883, 16, 2723). Mcltin;/
point, 418.5° to 419.35° (Heycock and Neville, J. f., 1895, 67, 185). Boilinn point,
940° (Violette, C. r., 1H82, 94, 720). It is a bluish-white metal, retaining its
histre 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 by percussion
(Schwarz, B., 1882, 15, 2505). At ordinary temperature it breaks with a course
crystalline fracture. It is more malleable at 100° to l.">0° 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: PtCl4 + 2Zn —
Pt + 2ZnCl2 .
2. Occurrence. — It is found as ealamir.e (ZnCO..,), as zinc-blende (ZnS); also
associated with other metals in numerous ores.
3. Preparation.— The process usually employed consists of two operations:
(/) Roasting: in case of the carbonate the action is: ZnCO = ZnO + CO..,; if it
is a sulphide, 2ZnS -f :;O, = 2ZnO + 2SOX . (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
puritied 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, ZniOHt , is made from solutions of zinc,
salts by precipitation with fixed alkalis (6«).
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,
«7. C., 1885, 47, 3-49). 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 hydrogen;
in very dilute nitric acid, without evolution of gas (3); in moderately dilute
cold nitric acid, mostly with evolution of nitrous oxide (.)); and, in somewhat
less dilute nitric acid, chiefly with evolution of nitric oxide (5). Concentrated
nitric acid dissolves zinc but slightlv. the nitrate being very sparingly soluble
in nitric acid (Montemartini, QaaxtUt, 1S92, 22, 277). Hot concentrated sul-
phuric acid dissolves it with evolution of sulphur dioxide (6).
(/) Zn + H2SO, = ZnS04 + H3
(2) Zn + 2KOH = K2ZnO, + H2
(5) 4Zn + 10HNO, = 4Zn(NO3)2 + NH4NO3 + 3H2O
(4) 4Zn + 10HNO3 = 4Zn(NO3)2 + N2O + -r>H2O
(5) 3Zn + 8HN08 = nZn(N03)2 + 2NO + 4H3O
(6) Zn + 2H,S04 = ZnSO, + SO2 + 2H2O
(ft) Oxide and Hydro-ridc— All the agents which dissolve the metal, dissolve also
its oxide and hydroxide.
*Mulr ami Kobbs, C. A'., 1882, 45, 69.
§135, tie. ZINC. 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 KOH and NaOH except
zinc sulphide, and all in NB.'4OH except ZnS and Zn2Fe(CN)8 .
6. Reactions, a. — The alkali hydroxides precipitate zinc hydroxide,
Zn(OH)2 , white, soluble in excess of the precipitant forming an alkali
zincate :
ZnCl2 + 2KOH — Zn(OH) , + 2KC1
Zn(OH)2 + 2KOH = K2ZnO2 + 2H2O
ZnCL + 4NH4OH = (NH4)2ZnO2 + 2NH.C1 + 2H2O *
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 zinc 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(OH)6(C03)2 , white,
soluble in ammonium carbonate, readily in alkali hydroxides (Kraut, Z.
anorg., 1896, 13, 1). Carbonates of Ba , Sr, Ca , and Mg 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, Zn(CN)2 , white,
soluble in excess of the precipitant. Alkali ferrocyanides, as K4Fe(CN)n ,
precipitate zinc ferrocyanide, Zn2Fe(CN),, , white (5c). Alkali ferricyanides,
as K3Fe(CN)(i, precipitate zinc ferricyanide, Zn3(Fe(CN)0)2 , yellowish, c. —
See 5c. d. — Sodium phosphate, Na2HPO4 , precipitates zinc phosphate, soluble
in alkali hydroxides and in nearly all acids.
e.— Hydrosulphuric acid precipitates a part of the zinc from neutral
solutions 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 inclines to the view that the solubility in NTH4OH Is due to the formation of a
complex ammonium-zinc ion (Scientific Foundations, p. 151 ; see also second German edition, p.
147). The fact that NH4C1 precipitates Zn(OH)a from its solution in fixed alkali, and on further
addition redissolves it and also that NH4CI hinders precipitation by heat from the ammoniacal
so.ution of the hydroxide speaks against the assumption that solution in the latter case arises
from the formation of a zincate.
180 ZINC. §135, 6f.
Z., 1886, 25, 512): ZnCl2 + 2KC2H,02 -f H2S = ZnS + 2KC1 +
2HC2H30, .* 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 precipitate
of the basic sulphite (Seubert, Arch. Phtmii., 1891, 229, :31(i)'. f.— If a hot con-
centrated zinc chloride solution be treated with ammonium hydroxide until
u precipitate begins to form, a basic chloride. 2ZnC1..9ZnO , will separate out
upon cooling- as a white precipitate (llabermann, M.. 1884, 5, 4:i2).
y. — Zinc salts are precipitated by solutions of alkali arsenites and arsenates,
forming1 respectively zinc arsenite or arsenate. white, gelatinous, readily solu-
ble in alkalis and acids, including1 arsenic acids, h. — Normal potassium chro-
mate forms, with solutions of zinc salts, a yellow precipitate readily soluble
in alkalis and acids, including1 chromic acid. Xo precipitate is formed with
X2Cr,O7 .
7. Ignition. — 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 coating1, yellow when
hot and white when cold. Tf this coating1, or zinc oxide otherwise prepared,
be moistened with solution of cobalt nitrate and again ignited, it assumes a
green color (Bloxam, ./. ('., ISO.1), 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 is
precipitated with Co , Ni and Mn from the ammoniacal solutions by HJ5 .
Digestion of the precipitated sulphides with cold dilute HC1 dissolves the
Mn and Zn as chlorides. The solution is thoroughly boiled to expel the
H..S and the zinc changed to Na,ZnO, by an excess of NaOH , which precipi-
tates the manganese as the hydroxide. From the alkaline filtrate H2S gives
a white or grayish-white precipitate — evidence of the presence of Zn .
9. Estimation. — (1) 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 (Ku'nzel,
Z., 1863, 2, 373). (3) It may be converted into ZnNH4PO, , and, after drying
at 100°, weighed. Ignition converts it into Zn,P,07 , with slight loss of /inc.
(4) Volumetrically, by converting into ZnJFe(CN)« and titrating with potas-
sium permanganate or by using Fed, acidulated with HC,H3O2 as external
*In the equation for acetic acid, ab = lcc, a and b, the concentrations of the H and C2II3O,
ions respectively, are small, c is large, and It, the so-called " dissociation-constant," to which
the strength of the acid is proportional, is very small. But addition of the fully-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 HC1 unite with the acetic ions of the sodium acetate,
leaving Na and Cl ions in the solution. The displacement of a weak acid from its salt by a
strong one lies then not so much in an attraction of the strong acid by the base as in the ten-
dency of the weak acid to form the non-ionized molecule.
£135, 10. ZINC. 18>
indicator (Voigt, Z. angcic, ISSU, I107). (J) Hy precipitation as Zn3(Fe(CN)8)± ,
treating the precipitate with potassium iodide and titrating the liberated iodine
(Mohr. Di>i<>l., 1858, 48, 115). (H) l!y litratiou in hydrochloric acid solution
with K4Fe(CN)8 , using a uranium salt as an indicator (Fahlberg, Z., 1874, 13,
379; Koninck and Prost, Z. (in(/(nc., 18<H>, .~>(>8). (?) By titration in alkaline
solution with Na3S , using a copper salt as an indicator. (N) The zinc is pre-
cipitated as ZnNH,AsO4 , the precipitate decomposed with HI and the liber-
ated iodine titrated with standard Na.,S,0, (Meade, J. Am. Soc., 1900, 22, 353).
10. Oxidation. — Metallic zinc precipitates the free metal from solutions
of Cd , Sn , Pb , Cu , Bi , Hg , Ag , Pt , An , As , Sb , Te , In , Fe \ Co ,
Ni, Pd, Rh, Ir, and Os (Gmclin-Kraut, Handbucli, 1875, 3, 6). Zinc
with copper (zinc-copper couple, used in water analysis) reduces nitrates
and nitrites to ammonia, chlorates to chlorides, iodates to iodides, ferri-
cyanides to ferrocyanides, etc. (Thorpe, <7. C., 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, (7. N., 1895, 71, 92).
The oxide is reduced to the metal by heating in a current of hydrogen
(Deville, A. Ch., 1855 (3), 43, 477).
i Daviea, .T. C., 1875, 88, ill.
REACTIONS OF IRON AND. ZINC GROUP
CO
U
-t—
W
g
?T
1
w
u
CQ
O
• fc
: W
1
£
o
•
a
d
0
N
fc
«
tq
M
•^
*.
W
d
E
u
d
P4
y
| Z
o
= s.
09
: W
P,
1
&
fl
&
£
J<
1
'o S
W
o
: I
!
OQ
"*7 d
V
o 'o
^O
*,
6
£
'1
0
O
o
0
£ &
o
Js
_
0 CT
O *~2
t3
i
i
b
w
0
- ^
d
0
09
°o °
£
'i
0
1
3
i
c
a
H
%
o
a
S
a
•-0 fej
o
^
1
1
-<
"e
* £
1 -
CO*
•3
•» •—
§-
u.
Fe(OH);
P.
"c«
; i
d
o
<D
PH
OQ
0)
&-
oj do
Us*
.ol3|
^ S^
a
—
44
1
S
1
•c
OQ
5
.
^
M
-M-
OJ
•
CO
0
E
a
— "•
t u.
-
O
•*
B
&
o QJ
f*
1
§
•
u.
V
a>
0 • fl}
2
0
sf
6
fe
QQ fe
n
H
.0
•^
++
rs
S
'o
•
-§
^3 j
Jo
W
0
- ;
-
.
•rr
i °'
1
3
^
5
!2
«
o
£.4
O h
O
cd
'S
•£
0
X O
PQ
-
E
5
tt
£i
a
c
|
W
o
3 -
.
»
i. g
£
«C
5
>° ^
o
j
ft
0
•j>
;2
W
•- i
7J
o>
"
g
bo
w
"S
:
to
ca
o
j-
1 J
~ ~
c
0
-*j
*i
^ ^
'5
"C
' •„.
T3
; ^C
'C
s
o
W K
1
tc
'§ d
5-1
1
1
w
o
0 0
W W
|
12
•s g
°? C8 .
3.g
c
1
1
E
M
y
"3
W fe >
r
0
C
® ^
- 2 B
§137.
TABLE FOR .IVIMN/.s' OF THE ZIXC (IROL'I'.
183
., §137. TABLE FOR ANALYSIS OF THE ZINC GROUP (FOURTH GROUP)
(Phosphates and Oxalates being absent).
Into the clear ammoniacal filtrate from the Third Group pass HYDROSUL-
PHURIC ACID GAS, and if a precipitate appears, warm, until it subsides.
Filter and wash with a one per cent solution of NH4C1 . (Test filtrate, in
which H2S gives no precipitate for the Fifth Group.)
Precipitate: CoS , NiS , MnS , ZnS .
Treat on the filter with cold dilute Hydrochloric Acid.
Residue: CoS, NiS* (black).
Solution: MnCL , ZnCL(H2S,HCl).
Boil the Noliition tlwrouyhly to remove the
BLS , cool, and add a decided excess
For Cobalt:
For Nickel:
Dissolve in nitro-
Dissolve the sul-
of potassium or sodium Kydroxide and
hydrochloric
phides in nitro-
digest without warming (§135, 6a).
acid, evaporate
hydrochloric
Filter and wash.
and add NaHCO3
acid, evaporate
%
and H2O2; warm
and acid an ex-
1
gently and filter.
cess of uitroso-/?-
Precipitate :
Solution :
' A green color to
naphthol in acet-
the filtrate indi-
ic solution to re-
Mn(OH),*
K2ZnO2.
cates cobalt
move the cobalt
(§140).
§132, 6ft). Filter
Dissolve in nitric
Test for zinc by
and add to fil-
acid and boil
adding H,S. A
Test the black resi-
trate ammonium
with an excess of
white precipitate
due with the
hydroxide till al-
PbO2 and HNO3.
(ZnS) indicates
borax bead (blue
kaline, filter and
Violet solution
zinc.
-;. color characteris-
to the filtrate
(HMnO4) indi-
tic of cobalt,
add H2S. A black
cates manganese
§132,7).
precipitate, NiS,
(characteristic
indicates nickel.
reaction, §134,
If sufficient nickel
Or: Dissolve the
6c).
be present to ob-
CoS and NiS,
.
scure the blue
add excess of
Dark-colored orig-
bead (§133, 7),
hot KOH and
inal solutions in-
dissolve the sul-
Br, boil, filter,
dicating an alka-
phides in nitro-
wash (until fil-
li salt of manga-
hydrochloric acid,
nese should be
evaporate and add
an excess of ni-
troso-/? -naphthol
trate gi^ ^s no
precipitate with
AgN03), add so-
lution of hot KI
reduced by
warming with
HC1 before pro-
in acetic acid so-
*i 11 d test the fil-
ceeding with the
lution (§132, 6ft);
trate with CSa.
analysis (§134,
1 filter, wash and
If free iodine ap-
~iO and 6f).
test the brick-red
pears, nickel is
precipitate with
the borax bead.
present(§ 133,60-
Study the text at
Confirm bv study
§133, f>ff, ft, e and
of the text, §134,
Study the text at
Study §132, 6e,
f; §132. 6ft and c;
7, §136, §138,
§135, 6rt and r,
§136, §138, §139,
§136, §138, §139,
§139, §142. §143,
U36, §138, §139.
§140, §141, §144,
?140, §141, §144,
§144, §145 and
§142, §143, §144,
§145 and ff.
§145 and ff.
ff.
§145 and ff.
*Small portions of cobalt and nickel sulphides may be dissolved by tho cold dilute HCi.and
will be precipitated with the Mn OH)a . These traces will not interfere with the further tests .
for manganese.
is-l- W7W770.YN roit (\ALY8IS WITH NOTES. $138.
DIRECTIONS FOR THE ANALYSIS OF THE METALS OF THE FOURTH GKOIH-.
§138. Manipulation. — Into the warm strongly ammoniacal filtrate from
the third group (§128), H.,S gas is passed until complete precipitation is
obtained :
MnCL.-'NH.Cl + 'JNH.OH + H,S = MnS + ^NH.Cl + 2H,O
(NH,),ZnO, + 2H...S = ZnS + (NH,),S + 2H,0
The solution is wanned 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 ol' NH4C1 ($139, ^). The filtrate
should be again tested with ELS and if complete precipitation has been
obtained it is set aside to be tested for the metals of the succeeding groups
ijjl91). The well wasbed precipitate of the sulphides of Co , Ni , Mn , and
Zn is digested on the filter or in a test-tube with cold dilute HC1 (one part
of reagent HC1 to four of water): MnS + 2HC1 = MnCl, 4- H,S . The
black precipitate remaining undissolved contains the sulphides of Co and
Ni , the filtrate contains Mn and Zn as chlorides with an excess of HC1
and the HJ5 which has not escaped as the gas.
§139. A"o/n».— (/) Instead of passing the H.S into the ammoniacal solution, a
freshly prepared solution of ammonium sulphide may be used. The yellow
ammonium sulphide, (NHJ-.-Sx, should not be employed to precipitate the
metals of the fourth group, as nickel sulphide is quite appreciably soluble in
that reagent (§133, fie).
(2) The sulphides of the fourth group, especially MnS and ZnS , should not
be washed with pure water, as thej* may be changed to the colloidal sulphides,
soluble in water. The presence of a small amount of NH.C1 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 HC1,
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.
(.'l) The sulphides of Co and Ni are not entirely insoluble in the cold dilute
HC1 . 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,
<>e). If desired, dilute acetic may be used, first removing the Mn and then
adding dilute HC1 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, fia), and is detected by the green color of the
fused mass when testing for Cr: :iMn(OH), + -»KNO, + Na,,C03 = 2K,MnO4 +
Na,MnO4 + 4NO + CO2 + :iH2O . 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 black precipitate of cobalt and nickel sul-
phides should first be tested with the borax bead (§141, 3) 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
HC1 , using a few drops of HN03 (§141, 1), and boiled to expel excess of
HNO:! : (iCoS + 12HC1 + 4HN03 = GCoCL, + 3S2 + 4NO + 8H,0 .
Divide the solution into three portions: To one portion of the solution
ijl42. DIRECTIONS FOR ANALYSIS WITH NOTES. 185
add an excess (§142,2) of nitroso- /5-Naphthol, filter, and wash with hot
water and then with hot HC1 (§132, 6&). Test the red precipitate with
the borax bead for cobalt. Render the filtrate ammoniacal, filter again
and test this last filtrate with H..S for the black precipitate of NiS (§133,
(56 and e). To another portion of the solution add NaHCO, in excess,
then add H202 , 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, 6f), catching this last
filtrate in a test-tube containing CSL, (§141, 6). Free iodine is evidence of
the presence of nickel.
§141. Notes. — (1) HN03 interferes with the nitroso- /3-naphthol reaction that
follows the solution of the sulphides of Co and Ni , hence an excess is to be
avoided. A crystal of KC103 may be used instead of HNO3 .
(2) If an insufficient amount of nitroso-/3-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.
-(3) 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.
(-}) 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^CCX, . The black
precipitate so obtained will filter rapidly.
(.7) The test for nickel by adding KI to the mixed higher oxides of cobalt
aifd 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;, + ">KI + <>HC1 =
3I2 + 6KC1 + :!HVO
(7) If the sulphides of Ni and Co be digested with yellow ammonium sul-
phide, a portion of the NiS will be dissolved (§133, fie) 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 liyrfrosulplniric acid (§143, 1], cooled (§135, (>"). a.nd then
treated with an excess of sodium hydroxide. The zinc forms the soluble
zaneate. Na.,ZnO., , while the manganese is precipitated as the hydroxide,
•wfeite. rapidly turning brown by oxidation:
Lj MnCL + 2NaOH == Mn(OH); + 2KC1
ZnCl, + ^NaOH = Na,ZnO, + :>NaCl + 2K.O
180 A\ALY8I8 OF IRO* AM) %l\c <ilt(H'l>S. §143, /.
Filter and test the filtrate with H,S , a white or grayish-white precipitate
indicates zinc (characteristic). Dissolve the well washed precipitate of
Mn(OH)., 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(OH), + oPbO, + 10HNO, = 2HKnO4 + r»Pb(NO:>), + OSLO
§143. A'ofot. /. If the H2S is not completely removed the Zn will be pre-
cipitated as the sulphide upon adding' the NaOH . :md will not be separated
from the manganese: ZnCl, + H2S + :.'NaOH = ZnS + 2NaCl + 2H20 .
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
HC1 and will appear with the precipitate o. MnfOH); . They do not interfere
with the final test for manganese.'
-}. The precipitate of Mn(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.
.?. Instead of PbOs , red lead, Pb304 , is frequently employed with the nitric
acid to oxidize the manganese to permanganic acid:
L'Mn(OH), + :.Pb.,O, + :iOHNO3 = SHMnO, + i:»Pb(N03), + lt>H,O .
. (!. It is very difficult to proem e PbO. or Pb^O, 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 PbO.. 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, Gr).
ANALYSTS OF IROX AXD ZINC GROUPS AFTKH PRECIPITATION BY AMMONIUM
SULPHIDE.
£144. It is preferred by sonic to precipitate tin- metals of the third
and fourth groups together, by means of ammonium sulphide; using
ammonium chloride to prevent the precipitation of magnesium (§189, 5ft
and 6</), and to insure the complete precipitation of the aluminum as the
hydroxide $124, Or;). In the manipulation for tln> method of separation,
the H._,S is not removed from the second group-filtrate, nor is nitric acid
used to oxidi/e any iron that may be present. To the second group filtralr
(^80), wanned, an excess of NH4C1 is added (£189, -V), then NH4OH till
strongly alkaline, and, paying no attention to any precipitate that may be
formed (Cxi. ^124, 125 and 126), normal ammonium sulphide is added (or
what is equivalent H..S is passed into the alkaline mixture). Aluminum
and chromium are precipitated as the lii/dro.i-i'lr*. the remaining metals n.-
ihe sHljiliidrs. The following table illustrate.- a plan of separation of the
ammonium sulphide precipitates of the third and fourth group metals,
phosphates being absent:
ANALYSIS OF IRO\ A\f> ZINC GROUP.
1ST
; 1
K
o
3^ • f^
0
. — ^2 d
u
02
e* ^^ cd
«p*
oa
• f^ Cs
tt
o I
.. o ^ - i*
fl rt ii co
5 N . rt ^
^
1)
ii
In
a
.S
aS C
I? rt
"S ^ uj'ft
•s w w-8 s
• & ^
'x
-' ^
M ^ S .
• a
o
* °
d' cc
l-H
^ %*%>*:
M
<J -a
Co '"'
< m
«
o '°
* S
, ,_ ^
K?
.;— <
•rt «^
^ C fl
o
Uj
^ -
s 3
rr &£
2
IS
^ 6
0 *
o H
• • (/2
<u • 22 H -S • fc"
"S /^ *o S ti 3 °*
5 w >•§.•§ 5 ^
l>
"o
as
t 1
llg
'« X r, c'S^ e
i
o
o
3 |-g
w nrt K
fS
"o
* •itll|fe ,• ". '.'.:'
i
o
_ b
N -Z ^
*3
O «oo
9
a
JH ' -0
"3
^
62 5
— ~
CS
fe ^,c3
1
rt o ^
•S
rn
•9
; ||
I ^
•"-< ---
*^ O>
^ g oi'So"
HJ
•o
o «*-*
£H ^ »«^*
% s
*i C5
g w S ^^W |
8
W -'
HP S £ ."S j^ "^
S ^ H ^ « •"£ t~
§
>^
n ^ rrt
o o
O rf T3 -. ^ ^ r-H
o
*
Q r— S
^W w^§^f*.
2
r '
^ |y
o S
Q « i fO ^2
S o -ts 3 * « *
•«•
" °?.°
M 2
O M* *• IS O
9
••*
3 MS
ij G" «
^ PS
« % '3 fc »
I
'o
is
r CJ , .
•O O r-T
r— COO
i
•il precipitate, 'on
-^ C
§ ll
+j X S
w 3 &
.2 ~~~
•••! ^-^ O •
K d M fc
M W 1>
W 3 l_i
r^
^^ ^3 ••
fl^ > , ^ y
•' ••:"! H«l!!i!?
» F^ 1 '•iC^illm
|5oi^ H c° "w"
te HGl frequently dii
^i , m S IH
1 ^ «S ca ^5
w
•jj
P
x S
* s
^C »£ «H
W "c
OD 03
-^ > S
^ 'So
O 02 X
OJ 4;
R M « '^ H « «Cra£ S,
S c^^ dijv"~'flo
So)'8!!! -dW^«W «*=:
« i; '3 O £^
PH^TS^ ^^°O2 of 12
O3T!li? «5*02l^rO <*'p
g S ^ WpH.!*W« Sc8
ie cold dilu
&
-
S
o p 51
H
-
2 "i " "J-- c
i
2
« ^2 "1 1^
(5 fe U o as ••"
" £ *" a .2
«f c ? •? £ '
O M i^. s: ^- "C
0 .|X
188 IROy AXD ZIXC 11ROUP8. §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 arc 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 HC1 ; 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(OH), and leaving the alkali phosphates in solution.
£147. In analysis a portion of the filtrate from the second group (after
the removal of the HJ3) (§128) should be tested for phosphoric acid with
ammonium molybdate (§75, <><7). Tf 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, H2S , 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-molybdatc has completely formed and settled
($75, (></). 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 H.,8 (§75, tie)-.
From this point proceed by the usual methods of analysis (§§127, 128
and //.).
§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 earth* are soluble,
and the phosphates of Al . Cr"' and Fe'". insoluble in acetic achl.
To the filtrate from the second group, freed from HJ3 by boiling (128),
and nearly neutralized with Na^GO, , an excess of NaC,H.,0., is added and
then FeCl, solution, drop by drop, as long as a precipitate is formed.
Care must be taken to avoid an excess of FeCl, . 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 FeCl, . The mixture should be boilod
IKON AXD ZLNC GROUPS. 18!)
to precipitate the ferric acetate as basic ferric acetate (§126, 6&) and at
once filtered.
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 bo
precipitated and the first drop of FeCL 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 earth 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 (§152) illustrates the separation of the metals in
presence of the phosphates by the use of FeCL, 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 BaCO, (Ga/§§124, 125 and 126): 2A1C1, + 3BaCO, + 3H20 =
2Al(OH)a + 3BaCL + 3C02 . Solutions of the chlorides or nitrates of
the fourth group and of the alkaline earths are not transposed by cold
digestion with BaC03 . Sulphates of the fourth group are transposed by
freshly precipitated BaCO, in the cold: CoS04 -j- BaCO, = BaS04 4-
CoC03 , etc.; and must not be present in this method of separation
(§126, Go).
If an excess of ferric chloride be present the phosphates will all \>\-
precipitated as ferric phosphate and the Al , Cr'" and excess of Fe'" as
the hydroxides upon the digestion with BaCO, . The table (§153) gives
an illustration of the use of the BaC03 in effecting the separation.
It should be observed that presence or absence of Fed, or of BaCO, in
the sample must be fully determined before their addition as reagents.
§151. 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
IRON AND eiKC CROUPS.
*190
organic matter, can readily be removed !>y ignition. Tf the presence <>f
an oxalate has been established (§§188, (>/> and 227, 8). the second group
filtrate should be evaporated to dryness, moistened with concentrated
UNO.; and gently ignited. The residue, dissolved in HC1 , 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, 8 and 221, 8).
§152.
'RON, ZINC AND CALCIUM GROUP METALS.
191
2
*0
p
€
2
13
•S-a
02 O
*
'
S |
•S 5
.2 <1
OS
p<
flj
co
oi
- r- .
__
S «~rC
i "** -^ *S fl
~ t-
'O
~ 5*
5
O •«-> .*•'
r-5' ,H fln ft'™ 2
s 2
fj
03 cj *^ ^^ P ^fi*(T\ i* F-
rH
«2 K lH
fl_) ^
"c
•
H r* -|H .^ 5\ r» *H
0
~ Z VA
'? a; "~"
o
0)
o+j IF? Kr*3>-i®o
Jj
a
g
^
<S
^ 'S ef ' - . M ^ &B 1
«C
s «T
rrj ~ P
6
S* § .^' r4' r^ ^ S S g S
—
0 '« ^
•§
C o
K
S1
o c 6 § ! SP ** o **
E
03
rfl '-^'5
a
'1-^5
£
o
,C
3s oo^"j"C"5ra
1
P" S &
9
3c5x
'i
PH
" rfl ^ • . « t£ ^ • ^ "C "^
1
5
S « g
•o
* OK X
O C/3
1
a
.2"? i? § x 5 -o * ,-,
ro^r P4yOi^r^'£'aj
s
s
^ o "ft
CS «H
rC S°
o"
2
fl
o a>
S 'S
frj «j f^
r~^/
OJ "IT c/
TH • .i-i TT
oi>
•*-• /11
eOQ
2 H
'3 « «
c"
a
o ?-
PI- -5
h ^ e~ «j -^ ;a
s
s ^ s
o o a
n o
ss
CS M
,° i— i a>
C
2 r»»
**-* rH *
"*£
CS
"Si s^
ffe SH"~< ^r* ^ jj
o
h 'g.i
as ^
°^£
g
O
° 50
-£ o d/
df '.:,§« T &
f>
PH
la
9^ E
B*
h
OK ft
* tf >• ^ "?' "« rrt
^
2 ^
•S/3
^ S
£
W
g >>
PH jj X1"-^ cu ^
ft 9
EC O
^3 9 a;
a/
n
bcT:
. pq ^ '5 rS S v
s
,§ a
•|H ""^
r*
CS
B CS 85
n c"* '^ X1 ^^ ^
^
1,
& a
£ ti ^
B
m
r- CV 05
pj *t^ rM ^ P~l fli ^
•PH
W
O> O (H
rH
i-« 2
o S ^ 25 g
1
§ rA
^
"1-2
I
1
w
0.0 0
CS °5 rQ
^tg •« ss^J>o &
r-? " 0 ^2 * 0
23 3|| ;^
1
QJ
1
&
• —
O
5 w
bine with
;tpitated
C c£ 05
-r->
I
o^l
TS i»»il
•
0)
A
h?tS ^
M Ll «
a g
83.
S X ^
C
' ^
§1.0
G SH -^ S pH
2
O J^
O 3
o al "|J
fll ^^
O Q) CS O n
^
rH
•S S
rQ 4)
rC ifc~
+J rQ fl)
£ PH ."pH ^ aJ o
o
< *
'•*. =
r^'fl ^
•^
tq
'S S B
& Ti ^ O "cS *
2
_0
i -o
rQ CS g
"S C
^
g Qj ,-i
Q rQ O «H -M W
w
"o
w ^
W CO P
facM
02
o
« .t:
w •<
' 1
IJ
CO -3-
i the Second Group, freed from H2
precipitate by ammonium chlorid
lium sulphide (§137). Filter and
both precipitates separately, and di
wash.f (For rare metals, see §154
0
W
0)
PH
W
r-l
O
o"
rl
u
W
™
at dilute hydrochloric acid, with th
e, and filter out any free sulphur,
sodium acetate strongly acidified w
r while hot.
CS
eg
• ¥
P4 CS
0) ^,
O -i- eos
•- o '
o 53
Solution:
KAIO, . or NaA102 with an
alkali phosphate.
Acidulate with HNO3 and add
an excess of NH,OH .
Precipitate: A1PO4 ,
insoluble in acetic acid.
Or, separate the Al and P04
as directed in §124, P>d and
7.
Confirm the presence of a
phosphate by (NH4)2M:o04
in the solution acidulated
with HN03, and also test for
a phosphate in the filtrate
from the NH4OH precipi-
tate of A1PO4 .
ti this filtrate are in excess of the amount n
, shows that some phosphates, (Fe'" or M«
sZ.
G rj o
Tl
.5 G «H a;
s-l 'C
" Qt
" 2
® rH ri
pG, ^
n
"C ®S
** - r>j
O rH
- ®
<& o a
0) -M (j
| ffl
o
^ 3
PH c
11,
+J CS
cy
r«t
r^ rt ^
p( £L,
».
cc 3 *o
E fl ^*
i;
<J
'O 4) r2 CS
•rH *7J
fw
— "— ' —
*J 4;
^S
<u
'« «S * r"
2 S
w |
B-d*
'If
t- .."S
10 <u
T)
^ B gj
* *
2- •-
•S .2 a
cS +j ,^j
CS [S
'in
^ ft'O ^
rG
4>V
o
"3 cs •-
-C
rS
X B r—
-t SS
Z- J3 fl
rH ^
W
S 'oc
•f-t
0)
* ft §
III
S.1
r^ - — 1
83 *'
X
O
PC
•d 1
'S H
>> O
G si a
•<! P-i o*
o
H
S
H
* ""p
192
7/J02V, Z/A"C' .1A7> ( ALCll M (lliOl 1' !//."/ l/.N.
i<155. CERIUM— COLUMBll'U. 193
THE RARER METALS OF THE IRON AND Zixc GROUPS.
Cerium, Columbium (Niobium), Didymium, Erbium, Gallium, Glucinum
(Beryllium), Indium, lanthanum, Neodymium, Praseodymium, Sama-
rium, Scandium, Tantalum, Terbium, Thallium, Thorium,
Titanium, (Jranium, Ytterbium, Yttrium, Zirconium.
§154. Cerium. Ce = 139.0 . Valence three and four.
Specific gravity, 0.628. Melts higher than Sb :ind lower than Ag (Hillebrandt
and Norton, 1'oyy., 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) (Poyy., 1842, 56, 503). The metal is obtained
from the chloride, CeCl3 , 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 Cl ,
Br and in vapor of I , S and P . Soluble in acids. Two oxides are known,
Ce2O3 and CeO2 , forming two classes of salts, cerous and eerie, the latter being
less stable. Ignition in air or oxygen changes Ce2Os to CeO. . Ce2Os is white
or grayish-white, soluble in acids and formed by igniting Ce2(C03).j , Ce2(C2O4)s
or CeO. in an atmosphere of hydrogen. Cerous salts are white and form color-
less solutions in water. Ceric oxide, CeO3 , is yellowish-white, orange-yellow
when hot, soluble in acids with difficulty; the hydroxide dissolves readily.
Ceric salts are yellow or red, fdrming yellow solutions. Ceric hydroxide,
Ce(OH)1 , dissolves in HC1 with evolution of chlorine, forming colorless cerous
chloride. Sulphurous acid decolorizes solutions of eerie 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 oxj'gen, with formation of eerie hydroxide. The precipitate is in-
soluble in excess of the fixed alkalis (distinction from Al and Gl). The pre-
cipitation 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 dihition with cold water. A con-
centrated solution of K,SO4 forms the douUe sulphate, K.,Ce(SO4)., , white,
sparingly soluble in water, insoluble in K2SO4 solution (distinction from Gl).
Ua,S2O3 does not precipitate cerium salts. BaC03 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 eerie hydroxide. If cerous nitrate be boiled with
PbO, and HNO3 , eerie nitrate, a deep yellow solution is formed (delicate test
for cerium). Cerium givss 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-
MUM. £i5<i
chloric acids, but is readily soluble in hot concentrated sulphuric acid, forming
a colorless solution (Koscoe, ('. N., 1878, 37, 25). It forms several oxides, CbO
CbOa and Cb,O.- . Columbic acid (anhydride) Cb_.O, , is a white powder, yellow
:\vhen hot (distinction from tantalum); it is obtained by ignition of the lower
oxides, or by decomposition of solutions of the salts by water or alkalis and
igniting. CbO, , black, is prepared by strongly igniting Cb.O, in a current of
hydrogen. Cb,O, , not too strongly ignited, is soluble in acids, from which
solutions NH.OH and (NH,),S precipitate iitliimliic acid containing some am-
monia. By mixing Cb_.O, with charcoal and heating in a current of chlorine, a
mixture of CbOCL and CbCl, is obtained. CbCl., is a yellow crystalline solid
(needles), melting at 1<)4° and distilling at 2-10.5° (Deville and Trobst, ('. r., 18fi7,
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 H SO, (distinction from tantalum). Cb.,O... not previously ignited dissolves
in HF; which solution when mixed with KF , the HF being in excess, gives
a double fluoride. ^KF.CbF,; if the HF be not in excess, a double oxy-fluoride
is obtained, :.'KF.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 givew
the columbates. the imttixxiiim ttalt being quite soluble in water and in pot as
sium hydroxide: the xiiiliinn xiilt 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 cnlumbate. Carbon
dioxide precipitates iitliinihii- acifl from solutions of columbates. Soluble salta
of Ba . Ca and Mg form white bulky precipitates with a solution of potassium
'columbate. Ag-NO; gives a yellowish-white precipitate, CuSO, a green pre-
cipitate. Cb,O., in presence of HC1 or H SO, gives a blue to broim color with
Sn or Zn, due to partial reduction of the Cb (distinction from tantalum).
1 Fused with sodium meta-phosphate. columbic acid gives in the inner flame a
violet to blue bead; a red bead by addition of FeSO4 .
8156. Didymium = J Neodymium- N* = 143.6 . Valence thrr, -.
i Praseodymium. Pr — 140.5 . Valence three.
Specific iiraritu, <'>.."»44. Mrltx with greater difticulty than Ce or La . Present
in cerite in Sweden and in mona/ite sand from Brazil. IMdymium was reported
about 1840 by Mosander, having been separated from cerium and lanthanum
In 1885 Welsimch (.!/., 1885, 6, 477) separated didymium salts into two distinct
salts, neodymium and praseodymium. Hy the absorption sped rum bands
other chemists are of the opinion that the so-called didymium consists of a
group of elements, nine or more (Kruess and Nilson, It., 1887, 20, 'Jl ('•('.; Kreuss,
A., 1892, 265. 1). Concerning the separation of didymium compounds, see
Dennis and Chamot (•/. Am. »S'w., 1897, 19, 799). By repeated fractionation of
the nitrate (several thousand times) Wclsbach 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 CO2 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 tin III ln'iiJfil tit Itftituitr, while Iiintlmninii mtlphate
l>rtcii>ittttcx from the saturated solution at 20°. Fixed alkalis precipitate the
hydroxide; NH.OH, a basic salt: insoluble in ex-cess of the reagents. Alkali
carbonates form a bulky precipitate, insoluble in execs.-, of the reagent, barium
carbonate precipitates slowly but completely. Precipitation by alkalis is pre-
vented by tartarie acid. Oxalic acid precipitates didymium salts completely,
soluble with difficulty in HC1 . The double potassium sulphate forms much
more slowly and less completely than with cerium. The salts give a distinct
and characteristic absorption spectrum. Consult Jones (Am., 1898, 20, 345),
Schele (Z. it miry.. 1898, 17. 319). Boudard (C. r., 1898, 126, 000), Demarcay
.(C. r., 1898, 126, 10:59), and Brauner (C. N., 1898, 77, 161).
£159. ERHfr.V—GAUJrM—GLUCINUAf. ]!»:.
. • $157. Erbium. Er=16fi.O. Valem-o three.
Erbium metal has not been prepared. As oxide or earth it is described by
Cleve (('. r., 1880, 91. 381) as that yttrium earth the most beautiful rose
colored. It forms a characteristic 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, 10Q.1;'; Soret, C. r., 1880, 91, 378; Crookes, C. N., 1886, 54,
18). The oxide ''gives upon ignition an intense green light; it is not fusible or
volatile.
£158. Gallium. Ga =: 70.0 . Valence three.
>s'/;<r///c gruritu, the solid, at 23° to 24.5°, 5.935 to 5.956; the melted, at 24.7°,
6.069. Melting point, ::0.15°; frequently may be cooled to 0° 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
gives a very we,ak, and fugitive flame spectrum; the spark spectrum shows two
.beautiful violet;lines. When heated in the air or in oxygen it is but slightly
onsijdized,' 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 Kkaalumiuum of Mendelejeff, who described the general prop-
,'ertjes 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
iPy.renees; and in some American zinc blendes (Cornwall, Ch. Z., 1880, 4, 443).
Hit 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, Ga O:, , 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 KHS04 . The alkalis and the alkali carbonates precipitate the salts
as the hydroxide, perceptibly soluble in fixed alkali carbonates, more easily
in ammonium hydroxide and in ammonium carbonate, and very readily in
t£e fixed alkalis. Tartrates hinder the precipitation of the hj^droxide. 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 HC1 , noticeable in very dilute solutions
(1-175,000). H.f} does not precipitate gallium salts from solutions acid with
mineral acids; from the acetate or in presence of ammonium acetate the wTiitr.
sulphide, Ga,S3 , 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 Ga2Cl8 , which decomposes
to GaCl3 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,SO, 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. Glucinum (Beryllium). Gl = 9.1 . Ypfenoc 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. H., 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
Heareely attacked by oxygen or sulphur. It is a strongly positive element,
.1% I \IHUM. §160.
in general properties between aluminum and tlie 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 chrysoberyl. G1(A1O2)3 . in phcnakite,
Gl_SiO, , and in some other silicates. It is prepared by heating the chloride,
G1C1, , with Na in a closed iron crucible (>'ilson and Pettersson. I.e.); or by
heating the oxide. G1O . with Mg (Winkler, H.. 1S90, 23, 120). The oxide, G1O ',
is obtained by igniting1 the hydroxide. It is a white infusible powder, soluble
in acids and in fixed alkalis. The hydroxide is prepared by precipitating- the
salts with NH4OH , soluble in the fifed nlknlix aiiid in a in in nn in in ntr'.oiMte,
concentrated: precipitated on dilution and boiling- (distinction and separation
from Al). The metal is soluble in acids except that when in the compact
form it is scarcely attacked by UNO, . The hydroxide is soluble on continued
boiling1 with NH.C1 . forming1 G1C1 . The more common salts of g-lucinum 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 NH,OH
(Joy, Am. ft.. 18(53, (2). 36, H.'i). The salts are not precipitated by H,S , but are
precipitated by (NH,)2S as the hydroxide. BnCO does not precipitate Gl saHs
in the cold, but precipitates them upon boiling. G1CL melts at about 600°
nnd sublimes at a white 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 HC1 and the Al and Gl
chlorides filtered from the SiO, . An excess of ammonium carbonate precipi-
tates both metals, but redissolves the Gl . After repeating this separation
several times pure glucinum hydroxide, Gli OH i . 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 yrarity, 7.11 to 7.28 at 20.4°. Mel thin point . 176°. Indium was discov-
ered in Freiberg zinc blende by Reich and Richter (,/. pr., 18G3, 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 (Hoettger, J. pr., 186(5, 98,
2(5). In the preparation of indium the Freiberg /inc is dissolved in HC1 or
IT.SO, , leaving an excess of the /.inc. 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 sulphuric1 acid; then
taken up with water separating from lead. The solution is precipitated with
NH.OH . which precipitates the In and Fe: this precipitate is dissolved in
HC1 and boiled for some tiire with NaHSO:: . Tin- indium sulphite is obtained
as a fine crystalline pcwder, which is treated with HNO:1 and HLSO4 , forming
indium sulphate, from which the metal is precipitated l>v /inc (Bayer, .-!., Ls.l,
158, :!72; Boettger, J. pr., 1869, 1O7, :i<»; 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 0^ . Indium
does not decompose water at 100°. At a red heat it combines with sulphur
and the halogens. Ry ignition with charcoal or in a current of hydrogen it is
reduced to the metal from its compounds. It is soluble in HC1 and H.SO4 ,
evolving H; in HNO, . evolving NO. In the reactions of its salts indium
deports itself quite similar to Fe'" and Al . Its most rharartrrifttic property is
its spectrum; two lines, an indium a , intense blue, and an indium ft , less
intense violet (Schroetter, J. pr., 1865, 95, 441). In203 is browu when hot.
£164. LANTHA\UM-SAMARIUM. 197
light yellow when cold, slowly soluble in cold acids, rapidly when heated.
Indium salts are precipitated by the alkalis as In(OH)3 , soluble in excess of
the fixed alkalis, reprecipitated by boiling or treating with NH4C1 . Tartrates
prevent the precipitation by alkalis. Alkali carbonates precipitate the indium
carbonate, soluble in ammonium carbonate, but reprecipitated on boiling.
BaCO., carbonate precipitates the indium completely as a basic salt (separation
from Co, Ni , Mn , Zn and Fe"). Phosphates form white precipitates from
neutral solutions. H2S precipitates from neutral solutions, or solutions acid
with acetic acid, yellow indium sulphide. In alkaline solutions HJ3 , or in
neutral solutions (NH,)..S , forms a white precipitate containing In.S:; . Yellow
In,S3 boiled with (NH4).,Sx becomes white and is partly dissolved. Upon cool-
ing the solution a bulky white precipitate separates out. K4Fe(CN),, gives a
white precipitate; KoCr64 gives a yellow precipitate; K2Cr2OT , K3Fe(CN)8 and
KCNS do not form precipitates.
§161. Lanthanum. La = 138.6 . Valence three.
Specific gravity, G.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 NHjOH is also used to separate La from Ne and Pr , the latter precipitat-
ing first (Cleve, Rl., 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 La is higher than that of Ce; it is also not so readily attacked
by HN03 . In cold water La is slowly attacked, but in hot water the action
is violent (Winkler, B., 1890, 23, 787). The oxide, La2O3 , is a white powder,
readily soluble in acids; with water it forms the hydroxide, La(OH)3, which
reacts alkaline tpwards litmus and absorbs CO. from the air. La(OH)3 is
soluble in a solution of NH.C1 (similar to Mg(OH)2). The salts are colorless.
K.SO4 and H2C.,04 form precipitates with lanthanum salts as with cerium salts.
Fixed alkalis precipitate lanthanum salts as La(OH)3 , white, insoluble in
excess of the reagent and not changing color on exposure to the air (distinc-
tion from Ce). Alkali carbonates precipitate La^GO^s , insoluble in excess.
BaCO.., precipitates the salts completely in the cold. NH4OH precipitates basic
salts. ITS forms no precipitate; (NH4),S precipitates the hydroxide. Lantha-
num gives a number of characteristic lines in the spark spectrum (Bettendorf,
A., 1889, 256, 159).
§162. Neodymium. 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,
B., 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 SCANDIUM— TAXT ALUM— TKRBIUM. $165.
§165. Scandium. 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, Sc.0;;,/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 white bulky precipitate, insoluble in excess of the precipitant. Tartrates
hinder the precipitation in the cold, but not upon heating. Na^COj gives a
bulky white precipitate, soluble in excess of the reagent. H.S is without
action, but (NH4)2S precipitates the 1indro.ride. K SO; precipitates the double
scandium sulphate. .'!KS04.ScL.(SOj3 . soluble in water but not in a saturated
K.SO. 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, Puffy., 1850, 99, 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..,O5 . It is insoluble
in acids except HF , in which it dissolves with evolution of H . Upon ignition
in a current of chlorine, TaCl3 , volatile, is formed. Solution of alkalis has
no actJon, upon fusion with the fixed alkalis an alkali tantalate is formed.
Ta^.0.1 is a white infusible powder, specific gravity, 8.01 (Marignac, A. Ch., I860,
(4), 9, 254). The oxide fused with fixed alkalis gives also an alkali tantalate,
M'TaO.., . When KOH is \ised, the fused mass is soluble in water. When NaOH
is used, water removes the excess of alkali, leaving the NaTaO3 as a white
residue, which dissolves in pure water, but not in NaOH solution. Tantalum
chloride is a yellow solid, melting at 211..'}° 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, 2HTa03.H,0 =
H^a.jOj . The freshly precipitated acid is soluble in acids and reprecipitated
"by NH.OH . The acid is readily soluble in HF , which solution with KF forms
a characteristic double salt, 2KF.TaP- , crystallizing in fine needles, insoluble in
•water slightly acidulated with HF (distinction and separation from colum-
bium). A solution of alkali tantalate gives with HC1 a precipitate of tantalic
acid, soluble in excess of the HC1 . From this solution NH4OH or (NHJ.S
precipitates tantalic acid; H SO, precipitates tantalic sulphate. Tartaric ;ici<l
prevents the precipitation with NH.OH and (NHJJ5 . A solution of tantalie
acid gives no coloration with zinc (distinction from Cb). Solutions of alkali
tantalates form tantalic acid with C02 . The acid fused with sodium meta-
phosphate gives a colorless bead (distinction from SiO ). 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, Tr,03 , is the darkest colored of the yttrium earths, soluble
§169.
THALLIUM— THORIUM. 190
in acids and sets NU3 free from ammonium salts. The hydroxide is a
gelatinous precipitate which absorbs CO2 from the air. It is quite probable
that terbia is a mixture of rare earths (Boisbaudran, C. r., 188fi, 102, 153, 'M?>..
483 and 899).
§168. Thallium. Tl = 204.15 . Valence one and three.
Thallium was discovered by Crookes by means of the spectroscope in 1801.
in selenium residues of the H,SO, factory at Tilkerode in the Hartz Mountains.
Germany (C. N., 1861, 3, 193, 303; 1863, 7, 290; 1863, 8, 159, 195, 219, 231, 24:;.
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,
-A., 1867, 144, 127). It is prepared by reduction from its solutions with Zn or
A'l; by electrolysis; by precipitation with KI , and then reduction by Zn or Al
or by electrolysis. Specific gravity, 11.777 to 11.9 (Werther, J. pr., 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 H2S04 and HNO3 , in HC1 with great difficulty; combines directly with
Cl , Br , I , P , S , Se , and precipitates from their solutions Cu , Ag , Hg .
An 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.
Thallious oxide, TLO , is black; on contact with water it forms an hydroxide,
T10H , 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.
Hydrochloric acid precipitates, from solutions not very dilute, thallioux
chloride, T1C1 , 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, Til , on adding a drop of potassium iodide solution, the
precipitate being slightly soluble in excess of the reagent. H.S pi'ecipitates
the acetate, but not the acidified solutions of its other salts. (NH4)^S pre-
cipitates TLS , which, 011 exposing to the air, soon oxidizes to sulphate.
Ferrocyanides give a yellow precipitate, Tl4Fe(CN)0; phosphomolybdic acid a
yellow precipitate; and potassium permanganate a red-brown precipitate, con-
sisting in part of TLO3 . Chroinates precipitate yellow normal chromate: and
platinic chloride, pale orange, tlmllious platinic chloride, TLPtCl0 . Thallium
compounds readily impart an intense green color to the flame, and 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, T1LO:1 , 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 bromides do not precipitate thallic solutions; iodides
precipitate Til with I. Sulphides and H._S precipitate tluilllnus sulphide, with
sulphur. Thallic oxide, suspended in solution of potassium 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 = 232.0 . Valence four.
Thorium is a rare element found 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
with potassium. The metal is a gray powder; specific (jrarit;/, 11.000; stable in
air at ordinary temperature, but igniting when heated; attacked by vapors of
Cl , 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 TITANIUM. ^170.
and N'ilson, /?., 1887, 20, 1665). Thorium forms one oxide, ThO2 , upon ignition
of the oxalate. It is u snow-white powder, not easilj- soluble in acids if highly
ignited (Cleve, ./., 1874, 261). The hudruj-id*-, Th(OH)4 , 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, ThCl4 , and the nitrate, Th.(NO.,)4 ,
"re deliquescent. The chloride is a white body melting at a white heat and then
subliming in beautiful white needles (Kruess and Xilson, /. c.). The sulphate
is soluble in five parts of eold water. The curlon'ttc, n.ralatc and pJio^}>Jiatc are
insoluble in water: the turnlate is scarcely soluble in dilute mineral acids.
Alkali hydroxides or sulphides precipitate thorium hydroxide, Th(OH)4 ,
insoluble in excess of the reagent. Tartaric and citric acids hinder the pre-
< ipitation. Alkali carbonates precipitate the basic carbonate, soluble in ex-
cess, if the reagent be concentrated. The solution in (NH4),COn readily repre-
cipitates upon warming. BaCO., precipitates thorium salts completely. Oxalic
ncid and oxalates form a white precipitate (distinction from Al and Gl), not
soluble in oxalic acid or in dilute mineral acids: soluble in hot concentrated
(NH,),CO, and not reprecipitated on cooling and diluting (distinction from
Ce and La). A saturated solution of K,SO, slowly but completely precipitates
a solution of Th(SO,).. , forming potassium thorium sulphate: innolulla in a
saturated K,SO, solution, sparingly soluble, in cold water, readily soluble in
hot water. HF precipitates Th.F4 , insoluble in excess, gelatinous, becoming
crystalline on standing. Roiling freshly precipitated Th(OH)4 with KF in
presence of HF forms K,ThF(1.4H.,O , a heavy fine white precipitate almost
Insoluble in water. The dixtiniiuixniny reactions of thorium are the precipitation
with oxalates and with KS04 , and failure to form a soluble compound on
fusion with Na.,CO3 (distinction from Si02 and TiO:).
§170. Titanium. Ti = 48.15 . Valence three and four.
Titanium is found quite widely distributed as rutile, brookite, anatase,
titanite, titaniferous iron, FeTiO, , and in many soils and clays. Never found
native. It is prepared by 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
flame 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 on/// metal that combines directly with nitrogen when heated in the
air (Woehler and Deville, A., 1857, 103, 230; Merz, J. pr., I860, 99, i:>7). The
most common oxide of titanium is the dioxide, TiO2 , analogous to CO, 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 TiO, in dry hydrogen gives Ti,0s , an
amorphous black powder, dissolving in H.SO, 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, Ti02 , is a white powder, melts somewhat easier
than SiO, , soluble in the alkalis unless previously strongly ignited. Mixed
with charcoal and heated in a current of chlorine TiCl4 is formed. The
bromide is formed in a similar manner. Ti02 acts as a bane, forming a series
of stable salts; also as an acid, forming titanates. TiCl4 is a colorless liquid,
fuming in the air; it boils at 136.41° (Thorpe, J. C., 1880, 37, 329); it is de-
composed by irv/fT. forming titanic acid, which remains in solution in the HC1
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 HJSO4 no TiF4 is volatilized (distinction
from EiF4). When evaporated with HF alone, TiF4 is volatilized. The double
pittasxiitin titanium flintridc. K.TiF8 , formed by fusing TiO, with acid KF , is
sparingly soluble in water (96 parts), readily soluble in HC1 . Solutions of
titanic salts in water or acid solutions of titanic acid are precipitated by
§171. URANIUM. 201
alkali hydroxides, carbonates and sulphides as the hydrated titanic acid, insolu-
ble in excess of the precipitants and in ammonium salts. BaC03 gives the same
precipitate. K4Fe(CN),, gives a reddish-yellow precipitate; K3Fe(CN)6 a yellow
precipitate. Na.,HPO4 precipitates the titanium almost completely, even in the
presence of strong HC1 . An acid solution of TiCX, when treated with Sn or
Zn gives a pale blue to violet coloration to the solution, due to a partial reduction
of the titanium to the triad condition. These colored solutions are precipitated
by alkali hydroxides, carbonates and sulphides. H.S is without action. The
solution reduces Fe'" to Fe" , Cu" to Cu' , and salts of Hg , Ag and Au to the
metallic state; the titanium becoming again the tetrad. The reduction by Sn
or Zn takes place in presence of HF (distinction from columbic acid). Titanium
compounds fused in the flame with microcosmic salt give in the redticing flame
a yellow bead when hot, cooling to reddish and violet (reduction of the tita-
nium). With FeSO, in the reducing flame a Mood-red bead is obtained.
§171. Uranium. TJ — 239.6 . Valence four and six.
Specific gravity, 18.685 (Zimmermann, A., 1882, 213, 285). Melts at a bright
red heat (Peligot, A. Ch., 1869, (4), 17, 368). Found in various minerals; its
chief ore is pitch-blende, which contains from 40 to 90 per cent of T730R .
Prepared by fusing UC14 with K or Na (Zimmermann, A., 1883, 216, 1; 1886,
232, 273). It has the color of nickel, hard, but softer than steel, malleable,
permanent in the air and water at ordinary temperatures; when ignited burns
with incandescence to TTSOS; unites directly with Cl , Er , I and S when heated:
soluble in HC1 , HJ5O4 and slowly in HN03 . Uranous ox-ide, TJO., , formed by
igniting the higher oxides in carbon or hydrogen, is a brown powder, soon
turning yellow by absorption of oxygen from the air. Uranous hydroxide is
formed by precipitating uranous salts with alkalis. Uranic oxidf, TJO3 , is
formed by heating uranic nitrate cautiously to 25°, and upon, ignition in the
air both this and other uranium oxides, hydroxides and uranium oxysalts with
volatile acids are converted into TJ30, = TJO;>2TT03 . Uranium acts as a base in
two classes of salts, uranous and uranijl salts. Uranous salts are green and give
green solutions, from which alkalis precipitate uranous hydroxide, insoluble in
excess of the alkali; alkali carbonates precipitate U(OH)4 , soluble in
(NH4)2C03; with BaC03 the precipitation is complete even in the cold. H^S is
without action; (NHJJS gives a dark-brown precipitate; K,Fe(CN)0 gives a
reddish-brown precipitate. In their action toward oxidizing and reducing
agents uranous and uranyl (uranic) salts resemble closely ferrous and ferric
salts; uranous salts are even more easily oxidized than ferrous salts, e.g., by
exposure to the air, by HNO3 , Cl , HCl63 , Br , KMnO4 , etc. Cold, silver and
platinum salts are reduced to the free metal. The hexad uranium (TJ^'1) acts
as a base, but usually forms basic salts, never normal: we have T70^(NO:;)j .
not TI(NO3)C; TJOJ304 , not TT(SO4)3 . These basic salts were formerly called
nranic salts, but at present (TJO;,)" is regarded as a basic radical and called
uru-ni/l, and its salts are called uranyl salts, e. g., U02CL uranyl chloride,
(ITO..):.(PO.l)2 urany] orthophosphate. Solutions of uranyl salts rre yellow:
KOH and NaOH give a yellow precipitate, uranates, KJI.O; and Na,,U,0T .
insoluble in excess. Alkali carbonates give a yellow precipitate, soluble in
excess; BaCO3 and CaCO3 give UO3 . HS does not precipitate tho nrMiiium,
but slowly reduces uranyl salts to uranous salts (Formanek, A., 1S<!0, 257, 115).
(NH4),S gives a dark-brown precipitate. K4Fe(CN)0 gives a reddish-brown
precipitate. Used in the analysis and separation of uranium compounds
(Fresenius and Hintz, Z. ungew., 1895, 502). Sodium phosphate gives a yellow
precipitate. The hexad uranium acts as an aeid toward some stronger bases.
Thus we have K,TJ,07 and Na,TT,O7 , formed by precipitating uranyl salts with
KOH and NaOH; compare the similar salts of the hexad chromium, K,Cr,O7
and Na..Cr,07 . Other oxides of uranium are described, but are doubtless com-
binations rf TJO. and U0:; . Zn . Cd , Sn , Pb , Co. Cu . Fo , and ferrous sails
reduce urnnyl salts to uranous salts. Solutions of Sn . Pt . Au . Cu , Hg and
Ag are reduced to the metal by metallic uranium (Zimmermann. 1. c.). For
method of recovery cf waste uranium compounds, see Lnube (Z. <nir/cir., 1SC9,
575).
202 YTTERBIUM— YTTRIUM— ZIRCONIUM. $172*
§172. Ytterbium. Yb = 173.2. Valence three.
Obtained as an earth by Marignac (C. r., ISTs, 87, 578) from a gadolinite
earth; by Delafontaine (C. r.. 1878, 87, 933) from sipylite found at Amherst, Va.
Nilson (/?., 1879. 12, 550; 1880, 13, 14:53) 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 give no
absorption spectrum. For the spark spectrum see Welsbach (M., 1884, 5, 1).
The o.ridc, Yb.03 , is a white powder, slowly soluble in cold acids, readily upon
warming. The hydroxide forms a gelatinous precipitate, insoluble in NH.OH
but soluble in KOH . It absorbs CO... 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. »S'., 1889, 38. 474). The metal has been prepared by electro-
lysis of the chloride; also by heating the oxide, Y=03 , with Mg (Winkler, /?..
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'. N., 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 been
shown to be mixtures of the oxides of several c'losely related elements (Wels-
bach, M., 1883, 4, fi41; Dennis and Chainot. -7. .-1m. flw., 1897, 19, 799). Yttrium
salts are precipitated by the alkalis and by the alkali sulphides as the
hydroxide, Y(OH), , a white bulky precipitate, insoluble in the excess of the
7-eagents (distinction from Gl). The ovide and hydroxide are readily soluble
in acids; boiling with NH.C1 causes solution of the hydroxide, as the chloride.
The alkali carbonates precipitate the carbonate Y... (CO3)3 , soluble in a large
excess of the reagents. If the solution in ammonium carbonate be boiled, the
lil/dn:.ridr is precipitated. Soluble oxalp.tes precipitate yttrium salts as the
white oxalate (distinction from Al and Gl): soluble with some difficulty in
HC1 . 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 s;ilts is not
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 . Oe and La). Fuse with alkali car-
bonate, decompose with HC1 , ami filter from the SiO, . Neutralize the filtrate
and precipitate the Y, La and Ce as oxalates with (NH,),C,0< . Ignite the
precipitate and dissolve in HC1 . Precipitate the La and Ce as the double
potassium sulphates, and from the filtrate precipitate the yttrium as the
hydroxide with NH.OH . 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 ;i raiv 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 (I'wifj., 1825,
4, 117). Also prepared by electrolysis of the chloride (Becquerel. A. C!h., 1831.
48. 337). The metal exists in three modifications: crystalline, graphitoidal and
amorphous. The amorphous zirconium is a velvet-black powder, burning when
£175. THE CALCIUM GROUP. 203
heated in the air. Acids attack it slowly even when hot, except HF , which
dissolves it in the cold. It forms but one oxide, Zr02 , analogous to Si02 and
TiOj . ZrO2 is prepared from the mineral zircon by fusion with a fixed
alkali. Digestion in water removes the most of the silicate, leaving the
alkali zircnnntc as a sandy powder. Digestion with HC1 precipitates the last of
the Si02 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 K2SO4 , and after resolution in acids precipitated by NH4OH and ignited
to ZrO2 (Berlin, J. pr., 1853, 58, 145; Roerdam., C. C., 1889, 533). Zr02 is a white
infusible powder, giving out an intense white light \vhen heated; it shows no
lines in the spectrum. It is much used with other rare earths, La2O3 , Y2O8 ,
etc., to form the mantles used in the Wclsbach gas-burners (Drossbach, C. C.,
1891, 772; Welsbach, ,/.. 3887, 2C>70; C. A7., 1887, 55, 192). The oxide (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).., = H.ZrO., , forms zirconates, decomposed by
acids. As a base 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 zirconi\im salts by Na,HPO4 or
H3P04 . The silicate, ZrO._..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. Solutions 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
NH4C1 solution (difference from Gl). Tartaric acid prevents the precipitation.
Alkali carbonates precipitate basic zirconium carbonate, white, soluble in
excess of KHCO3 or (NH^-CO..,; boiling precipitates a gelatinous hydroxide
from the latter solution. BaCO 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 HC1 . A saturated solution of K2S04
precipitates the double potassium zirconium sulphate, white, insoluble in excess
of the reagent if precipitated cold, soluble in excess of HC1; if precipitated
hot, almost absolutely insoluble in water or HC1 (distinction from Th and Ce).
Zirconium salts are precipitated on warming with Na._,SL,O., (separation from
Y, Nd and Pr). Solution of H2O, completely precipitates zirconium salts,
Tumerte paper moistened with n solution of zirconium salt and HC1 is colored
f>range upon drying (boric acid gives the same reaction) (Brush, J. 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 GROUP (FIFTH GROUP).
(THE ALKALIXE 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 GROUP. §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: (§7) :
§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. Cli., 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 they
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: BaCO. is soluble in 45,566 parts at 24.2° (Hollemann,
Zeit. phys. Cli., 1893, 12, 125); SrC03 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 (Chcvalet, Z., 1869, 8, 91). The
presence of NH4OH and (NH4),CO, lessens the solubility of the carbonates
of Ba . Sr , and Ca , while their solubility is increased by the presence of
NH4C1 . MgCO:! is soluble in ammonium carbonate and in ammonium
chloride, so much so that in presence of an abundance of the latter it is
not at all precipitated by the former, i. e. (NHJ..CO., does not precipitate a
solution of MgCl, as the NH4C1 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 difficulty (§145 and //.).
§180. The oxalates of Ba , Sr, and Mg are 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
precipitation with carbonate in presence of ammonium chloride, after
whii-li magnesium is precipitated from the filtrate, as phosphate.
§182. The hydroxides of Ba , Sr . and Ca , in their saturated solution.-,
necessarily dilute, precipitate solutions of s.alts 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 205
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 Mg(OH)o 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-
tated by 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 + H8PO4 + 2NH4OH = CaHPO, + 2NH4C1 + 2H20
CaH4(P04)2 + 2NH4OH = CaHP04 + (NH4)2HPO4 + 2H20 .
If excess of the ammonium hydroxide be added the precipitate is Ca3(P04)2.
In the case of a magnesium salt the precipitate is LIgNH4P04 .
§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 HC1) impart char-
acteristic colors to the non-luminous flame, and readily present well-defined
spectra.
§186. Barium. Ba = 137.40 , Valence two.
1. Properties. — Specific gravity, 3.75 (Kern, C. N., 1875, 31, 243); melting point,
above that of cast iron (Frey, A., 1876, 183, 368). It is a white metal, stable in
dry nir, but readily oxidized in moist air or in water at ordinary temperature,
hydrogen being- evolved and barium hydroxide formed. It is malleable and
ductile (Kern, I.e.).
. 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
oxide and hydroxide are so basic, absorbing acids readily from the air. Its
most common forms of occurrence are heavy spar, BaSO, , and witherite,
BaC03 .
3. Preparation. — (1) By electrolysis of the chloride fused or moistened with
strong HC1 . (2) By electrolysis of the carbonate, sulphate, etc., mixed with
Hg and HgO , 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, Ba(OH)2 , 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. Po.w/., 1832. 26, 172) or cupric oxide (Wanldyn, 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, HoO2 ,
which is formed by treating it with dilute acids: BaO., + 2HC1 = BaCL +
* The presence of an excess of fixed alkali renders these hydroxides much less seluble, th»>
high concentration of the hydroxyl ions, one of the factors of the solubility product, diminish-
ing' the other factor. (§45).
£06 BARIUM. §186, 5«.
5. Solubilities. — a. — Metal. — Metallic barium is readily soluble in acids with
evolution of hydrogen, ft. — Oxides and hydroxides. — Barium oxide is acted iip6n
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
(KoPtMistheil and Ruehlmann, J., 1870, 314). Barium peroxide, Ba(X , is very
sparingly soluble in water (Schone, A., 1877, 192, 257); soluble in acids with
formation of H.O, .
c. — Salts. — Most of the soluble salts of barium are permanent; the
acetate is efflorescent. The chloride, bromide, bromate, iodide, sulphide,
ferrocyanide, nitrate, l^pophosphite, chlorate, acetate, and phenylsul-
phate, are freely soluble in water; the carbonate, sulphate, sulphite,
chromate* phosphite, phosphate, oxalate, iodate, and silico- fluoride, are
insoluble in water. The sulphate is perceptibly soluble in strong HC1 .
The chloride is almost insoluble in strong hydrochloric acid (separation
from Ca and Mg) (Mar, Am. 8., 1892, 143, 521); likewise the nitrate in
strong hydrochloric and nitric acids. The chloride and nitrate are insolu-
ble in alcohol.
G. Reactions, 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, BaCO , white. The precipitation is promoted by heat and
by ammonium hydroxide, but is made slightly incomplete by the presence
of ammonium salts (Vogel, J. pr., 183G, 7, 455).
Barium Carbonate — BaCO, — 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 twice 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): Na2S04
-f BaCO., == BaS04 + Na,,CO, . When salts of non-alkali metals are so
decomposed, of course, they are left insoluble, as carbonates or hydroxides,
nothing remaining in solution:
FeSO4 4- BaCO,, = BaSO, + FeCOa
Fe.(S04)5 + ;BaC03 + :!H,O = :;BaSO, + 2Fe(OH), + 3CO2
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
* Kohlrauscli and Rose, Z. phys. Ch., 1893, 12, 341 ; Schweitzer, Z., 1890, 28, 414.
§186, 7. BARIUM. 207
acid, citric acid, sugar, and other organic substances, hinder or prevent
the decomposition by barium carbonate.
6. — Ammonium oxalate precipitates barium oxalate, BaC2O4 , 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, fulj metallic, or two-thirds metallic,
as Na2HPO4 , precipitate barium phosphate, white, consisting of BaHPO4
when the reagent is two-thirds metallic, and Ba3(PO4)2 when the reagent is
full metallic. Soluble phosphites precipitate barium salts, hypophosphites dp
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, H2S04 , and all soluble sulphates, precipitate barium
sulphate (BaS04), 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, 5c) solution of strontium sulphate
is a more certain test between Ba and Sr . BaS04 is not transposed by
solutions of alkali carbonates (distinction from Sr and Ca , §188, 6a foot-
note).
f. — Solutions of iodates, as NaIO3 , precipitate, from barium solutions not
very dilute, barium iodate, Ba(IO3)., , white, soluble in 600 parts of hot or
1746 parts of cold water (distinction from the other alkaline earth metals).
ff. — Neutral or ammoniacal 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 cliromate, BaCr04 . 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 K.,Cr207 ,
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).
i. — Fluosilicic acid, H2SiF8 , precipitates white, crystalline barium fluo-
silicate, BaSiF0 , 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 nitrate (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 g, green glass. The spectrum of barium is readily distinguished
from the spectra of other metals by the green bands Ba«, fi and } . Barium
carbonate is very stable when heated, requiring a very high heat to decompose
it into BaO and CO2 .
208 STROM' I I'M §188, S.
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 K,Cr,0- as BaCr04 which separates it
from strontium and calcium. The barium is further identified by the
non-solubility 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 sxilphate).
9. Estimation. — Barium is weighed as a sulphate (Fresenius and Hurtz, Z.
angnf., 1890, 253), carbonate or fluosilicate (BaSiF, ). It is separated from
strontium and calcium: (/) By digesting1 the mixed sulphates at ordinary tem-
peratures for 12 hours with ammonium carbonate. The calcium and strontium
are thus converted into carbonates, which are separated from the barium
sulphate by dissolving in hydrochloric acid. (2) By hydrofluosilicic acid.
(3) By repeated precipitation as the chromate in an acetate solution.
Tt 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. Volumetric-ally it is precipitated as the chromate, thoroughly
washed, dissolved in dilute HOI and the Crvi determined by H...O, (Baumann,
/. iitiiicir.. 1891, 331).
10. Oxidation. — Barium compounds are reduced to the metal when heated
with Na or K (:i). BaO.. oxidizes MnCl, to Mn..O3 (Spring and Lucion, Bl.,
1890, (3), 3, 4).
£187. Strontium. Sr = 87. (50 . Valence two.
1. Properties.— Spf-riflr ffrnritii, 2.4 (Franz. ./. pr., 1869, 107, 234). 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 wheti exposed to
the air. and when heated in the air burns, as does barium, with intense
illumination (Franz, /. <:).
2. Occurrence. — Strontium occurs chiefly in strontianite. SrCO:, . and in
celestine. SrSO, .
3. Preparation.- First isolated in 1HOS by Davy by electrolysis of the hydrox-
ide (Twin*. Koi/iil .S'w., 345). It is made by electrolysis of the chloride (Frey.
A., 187(5, 183, 307): by heating a saturated solution of SrCl, with sodium
amalgam and distilling off the mercury (Franz, /. r.): by heating the oxide with
powdered magnesium the metal is obtained mixed with MgO (Winkler, /?., 1890,
23, 125).
•>. Oxides and Hydroxides. — Strontium oxide. SrO . is formed by igniting the
hydroxide, carbonate (greater hcai required than with <alcium carbonate),
nitrate and all organic strontium salts. The hylioxide, Sr(OH).. , is formed
by the action of water on the oxide. The peroxide, Sr0....sH...O , is made by pre-
cipitating the hydroxide with H.O..: at 100° this loses water and becomes SrO, .
a white powder, melting at a red heat, used in bleaching works (Conroy,
J. Roc. Ind., 1892, 11. si 2).
r>. Solubilities. — «. — Mrtnl. — Strontium decomposes water at ordinary tem-
perature (Winkler. /. c.). it is soluble in acids with evolution of hydrogen.
h. — O.rirfcx and liiidrn.r'ulrx. — The oxide. SrO , is soluble in about 100 parts \vater
at ordinary temperature, and in about five parts of boiling water forming the
hydroxide' (Scheibler. Xrnr /.ritxrlirift fur Rncliniziirkn: 1881, 49, 257). The
peroxide is scarcely soluble in water or in ammonium hydroxide, soluble in
acids r.nd in ammonium chloride.
£187, Mi. STRONTIUM. 209
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,
5c). The sulpliats is very sparingly soluble in water (1-10,090 at 20.1°)
(Hollemann, Z. pliys. Cli., 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 HC1 or HN03 ; 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. Reactions, a. — The fixed alkalis precipitate strontium salts when
not too dilute, as the hydroxide, Sr(OH)2 , less soluble than the barium
hydroxide. No precipitate with ammonium 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).
c. — 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 6e, §§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
(5c). 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 (5c). It is almost insoluble in a solution of
ammonium sulphate (separation from calcium).
/. — The halides 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, (j. — Neutral solutions of arsenites 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, 60).
h. — Normal chromates 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 solutions
a solution of K2Cr04 plus one-third volume of alcohol. The calcium
CALCIUM. $187, ti».
chromate is about 100 times as soluble as the strontium chroinate (Fre-
senius and Kubbert. Z., 1891, 30, 672). No precipitate is formed with
potassium bichromate (separation from barium).
i. — Fluosiiicic 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. Ignition. — 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 «, at the red end of the spectrum; the two red lines, Sr /3 and Sr y,
and the blue line, Sr rf , 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 K,Cr20- . The strontium and calcium are separated from the
excess of chromate by reprecipitation with (NHJ..CO., . The precipitate is
again dissolved in HC,H:i02 and from a portion of the solution the stron-
tium is detected by a solution of CaS04 (6e). The flame test (7) is of value
in the identification of strontium.
M. Estimation. — Strontium is \\eighed :is ;i sulphate <>r a carbonate. The
§:lm\ide .-ind earbonate may In- determined by alkalimetry. It is separated
in e:ilfiiiin: (I) l!v the itisol uhilit v (if its -iilpliate in ammonium sulphate.
By tailing the nitrates with nmyl alcohol (§188, !»)• (-r) I'-.v treatii
rut'eM \\ith eqiml vol. line of absolute nlctfhol and ether (§189, 9). For
separation frdtrt barium flee §180, 9.
SI 88. Calcium. Ca = 40.1 . Valence two.
1. Properties.— ffpcHfic //rarity, 1.6 to 1.8 (Caron, C. r., I860, 50, 547). Melting
point, 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. — (1) 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
CaCL , 400 parts 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, /. r.). (.'?) By electrolysis of the chloride (Frey, I. c.).
(4) By 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 corresponding1 hydroxide, Ca(OH)2 (slaked lime), is made by treating the
oxide with water. Its usefulness when combined with sand, making mortar,
is too well known to need any description here. The peroxide, CaO.j.SHoO , is
made by adding hydrogen peroxide or sodium peroxide to tho hydroxide.:
Ca(OH), + H.Oo = Ca02 + 2H2O (Conroy, J. Soc. Ind., 1892, 11, 80S).' Drying
at 130° removes all the water, leaving a white powder, CaO2 , which at a red
heat loses half its oxygen (Schoene, A., 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)2 .
6. — Oxide and hydroxide. — CaO combines with dilute acids forming cor-
responding salts, it absorbs C02 from the air becoming CaCO., .* In moist
air it becomes Ca(OH)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)2 , is soluble in acids, being capable of
.jitration 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 Eotondi, B., 1874, 7, 817); and in 1712 parts at 100°
(Lamy, 1. 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, I. c.). A clear solution of the hydroxide in water is lime
water (absorbs C02 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 chloral ,<• are
deliquescent; the acetate is efflorescent.
The carbonate, oxalate, and phosphate an1 insoluble in water. Thi>
chloride, iodide, and nitrate are soluble in alcohol. The nitrate is soluble
in 1.87 parts of equal volumes of other and alcohol (Freseniua, Z., 1893,
32, 191); readily soluble in 'boiling amyl alcohol (Browning, Am. 8., 1892,
143, 53 and 314) (separation from barium and strontium). The carbonate
is -viable in water saturated with carbonic acid (as also are barium, stron-
tium., and magnesium carbonates), giving hardness to water. The oxalato
is insoluble in acetic acid, soluble in hydrochloric and nitric acids. The
sulphate is soluble in about 500 parts of water f at ordinary temporaturc,
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. Ammonium sulphate (1-4) requires 287 parts for the
solution of one part of CaS04 (Fresenius, Z., 1891, 30, 593) (separation
from Ba and Sr). Eeadily soluble in a solution of Na2S.,0;. (separation
from barium sulphate) (Diehl, J. pr., 1860, 79, 430). It is soluble in 60
parts hydrochloric acid, 6.12 per cent at 25°, and in 21 parts of the same
* Dry CaO does not absorb dry CO2 or SO2 below 350<\ < Veley, J. C., 1893, 63, 821).
t Goldhammer, C. C., 1888, 708; Droeze, B., 1877, 1O. 330; Hoisbaudran, A. Ch., 1874. <5), 3, 477
Kohlrausch and Rose, Z. phys. Ch., 1893, 12, 241 ; Raupenstrauch, if., 1885, 6, 563).
CALCIUM. §188, 6a.
acid at 103° (Lunge, J. Soc. Ind., 1895 14, 31). The chromate is soluble
in 214.3 parts water at 14° (Siewert, J., 1862, 149); in dilute alcohol it is
rather more soluble (Fresenius, I. c., page 672); very readily soluhle in
acids including chromic acid.
6. Reactions, 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), i. e. potassium hydroxide will form a
precipitate with calcium sulphate since the sulphate requires less water
for its solution than the hydroxide (56 and c) ; also the calcium hydroxide
is less soluble in the alkaline solution than in pure water. Ammonium
hydroxide does not precipitate calcium salts. The alkali carbonates pre-
cipitate calcium carbonate, CaC03 , -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(OH)., , is used as a reagent for the detection of
carbon dioxide (51 and §228, 8).
b. — Alkali oxalates, as (NH4),C.,04 , precipitate calcium oxalate, CaC204 ,
from even dilute solutions of calcium salts. The precipitate is scarcely at
all soluble in acetic or oxalic acids (separation of oxalic from phosphoric
acid (§315), but is soluble in hydrochloric 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 digesting a few minutes,
if a precipitate appears, SrS04 , 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 neutral or
alkaline solution be treated with a mixture of (NH4).,S04 and (NH.,).,C204 ,
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, B., 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) (Baubignj', BL, 1895, (3), 13, 326).
* Here experiment shows that for equilibrium the 8O4 ions must be present in solution in large
excess of CO3 ions. With strontium also an excess of SOt ions 1=1 required, nlthough nut so
great as in the case of calcium. For barium, however, equilibrium demands that the concen-
tration of CO3 ions exceed that of SO4. This condition is alrendy fulfilled when an alkali CT-
bonate is added to BaSO4 and therefore no change takes place in this case, while in tho o'hors
the sulpha'e is transformed into carbonate. It is important to notice that the relative or ab-
solute quantities of solid carbonate and sulphate pr -sent do not affect the equilibrium, which
is determined solely by the substances in solution (§57, 6e, footnote).
$188,9. c \LCIUM. 21.T
.e. — See .V. 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 CaS04 . The precipitate is Ca(OH),: CaCL + 2N&.S + 2H2O =
Ca(OH), + 2NaCl + 2NaHS . The acid sulphide', HaHS , does not precipitate
calcium salts (Pelouze, A. CJi., I860, (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 CaS04 ,
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)2 (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 CaNH4As04 (distinction from strontium after the addition of sulphuric
acid) (Bloxam, C. N., 1886, 54, 16).
/<.— Normal chromates, as K^CrO, , precipitate solutions of calcium salts as
calcium chromate, CaCrO4 , yellow, provided the solution be not too dilute (5c),.
The precipitate is readily soluble in acids and is not formed with acid chro-
mates as K.jCr.,O7 (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, CaSO,.2HoO , gypsum , loses its water of
crystallization at 80° and becomes the anhydrous sulphate, CaS04 , plaster of
Paris; which on being moistened forms the crystalline CaS04.2H2O , expands
and " sets." Calcium carbonate, limestone, when heated (burned) loses carbon
dioxide and becomes lime, CaO .
• Compounds of calcium, preferably the chloride, render the flame yellowish
r.ed. 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 tho
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 MAGNESIUM. ^189, •.
alcohol (">r). The best method of separation from strontium is 10 treat the
nitrates with a mixture of equal volumes of alcohol and ether. The calcium
nitrate dissolves, but not more than one part in G0,000 of the strouliiim -
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. X. (irtycir., 1898, 770). See also 9. §186 and §187.
$ 189. Magnesium. Mg = ->4.H . Valence two.
1. Properties.— ,s>r///r j/rari///, 1.75 (Deville and Caron, .1. ('It., isr,::. (:t), 67,
:»4(>); melting point, a little below 800°, does not appear to be volatile (Meyer,
It., 1887, 20, -197). A white, hard, malleable and ductile metal: not acted upon
by water or alkalis at ordinary temperature and only slightly at 100° (Hallo.
li., lS8:i, 16, (>94). When heated in air or in oxygen it burns with mcandeseemv
to MgO . It combines directly when heated in contact with N . P . As , S
and Cl . It forms alloys with Hg and Sn . forming compounds which decom-
pose water.
2. Occurrence. — Magnesite. MgCO-: dolomite, CaMg(CO:i) ,: brueite. Mg(OH);:
epsom salts, MgS04.7H..O: and combined with other metals in a great variety
of minerals.
H. Preparation. — (/) By electrolysis of the chloride or sulphate (Hunsen, .1.;
1852. 82, 137). (£) By ignition of the chloride with sodium or potassium
(Wohler. .1., 1857, 101, 502). (.?) Mg,Fe(CN)a is ignited with Na CO , and
this product ignited with zinc (Lanterbronn, German Patent No. :'.9.915).
4. Oxide and Hydroxide. — Only one oxide of magnesium. MgO , 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.
,"). Solubilities. — a. — Metal. — Magnesium is soluble in acids including
carbonic acid, evolving hydrogen: Mg -f COL, -4- H.,0 == MgCO, -4- H,
(Ballo, #., 1882, 15, 3003): it is also attacked by the arid alkali carbonat.--.
as NaHC03 , to form MgCO, , Na.CO, and H (Ballo, /. r.). Soluble iri
ammonium salts: Mg + 3NH4C1 == NH4MgCL 4- '.'NIL + H, . With
the halogens it acts tardily (Wanklyn and Chapman, •/. ('.. ISCC, 19, 141).
ft. — Oxide and hydroxide. — Insoluble in water, soluble in acids. Mg(OH),
is soluble in 111,111 parts of water at 18° (Kohlrausch and Rose, Znt.
phys. CJi., 1893, 12, 241). In contact with water the nxidc is slowly
changed to the hydroxide, Mg(OH), , and ahsorlts CO, from the air. Sol-
uble in ammonium salts:* Mg(OH), -f :5NH4C1 NH4MgCl, •£
2NH4OH . r.— &nltis. — The chloride, bromide, iodide, chlorate, nitrat.-'.
and acetate (4 aq) are <l<'li</ii<'x<rnf \ the sulphate (7 :u\) slightly dlloi-cs-ent .
The carbonate, phosphate, borate, arsenite, and arsenate arc 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( OH» ,
* The conditions here are the same as in the case of Mn(OH),, $134, 6a, footnote.
ij 1 89, 7 . MAGXESIUU. 215
I
white, gelatinous, from solutions of magnesium salts; insoluble in excess
of the reagent but readily soluble in ammonium salts: Mg(OH)2 -f- :>NH4C1
= MgCl,.NH4Cl 4- 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: 2MgS04 -4- 2NH4OH = Mg(OH),
+ (NH4)2Mg(S04)2 (Kheineck, DingL, 1871, 202, 268). The fixed
alkali carbonates precipitate basic magnesium carbonate, Mg4(OH).J-
(C0,):, , variable to Mg.,(OH)2(CO,)4 : 4MgS04 -f 4Na2CO, + H20 \
Mg4(OH)2(CO;s):! + 4Na,S04 -f C02 . If the above reaction takes place in
the cold the carbon dioxide combines with a portion of the magnesium
carbonate to form a soluble acid magnesium carbonate: 5MgS04 -(-
5Na,CO, + 2H,0 = Mg4(OH)2(CO,), -f MgH2(CO,)2 -f :>Na,S04 . On
boiling, the acid carbonate is decomposed with escape of C02 . Ammonium
carbonate does not precipitate magnesium salts, as a soluble double salt is
at once formed. Acid fixed alkali carbonates, as NaHCO , , do not precipi-
tate magnesium salts in the cold; but upon boiling, C02 is evolved and the
carbonate is precipitated (Engel, A. Ch., 1886, (6), 7, 260).
ft. — Soluble oxalates do not precipitate solutions of magnesium salts, as they
form soluble double oxalates. Tf to the solution of double oxalates, preferably
magnesium 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 Na2HP04 — precipitate magnesium phosphate,
MgHP04 , if the solution be not very dilute. But even in very dilute
solutions, by the further addition of ammonium hydroxide (and NH4C1).
a crystalline precipitate is slowly formed, magnesium ammonium phosphate
— MgNH4P04 . 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 (56).
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., 1860, 8, 125).
e. — Magnesium xulphidc is decomposed by water, and magnesium salts are
not precipitated by hydrosulphuric acid or ammonium sulphide; but MgO -(-
H2O (1-10) absorbs H2S , forming in solution MgH^S;, , which readily gives
off US upon boiling (a very satisfactory method of preparing H;,S absolutely
arsenic free) (Divers and Shmidzu, J. C., 1884, 45, (599). Normal sodium or
potassium sulphide precipitates solutions of magnesium salts as the hydroxide
with formation of an acid alkali sulphide: MgSO4 -f- 2Na,S + :2H..O = Mg(OH),
+ Na,SO4 + 2NaHS (Felouze, A. Ch., 1806, (4), 7, 172). Sulphuric acid and
soluble sulphates do not precipitate solutions of magnesium salts (distinction
from Ba , Sr and Ca).
f. — Magnesium chloride, in solution, evaporated on the water bath evolves
hydrochloric acid (7). .</. — Soluble arsenates precipitate magnesium salts in
deportment similar to the corresponding phosphates.
7. Ignition. — Magnesium ammonium phosphate when ignited loses ammonia
216 MAGNESIUM. ^189, 8.
and water, and becomes the pyrophosphate: 2MgNH^O« — MgjP20, + HaO +
2NH.: . The carbonate loses CO., and becomes MgO . In dry air magnesium
chloride may be ignited without decomposition, but in the presence of steam
MgO and HC1 are formed: MgCl, + H,O = MgO + 2HC1: a technical method
for preparing- HC1 (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 in
precipitated as the white magnesium ammonium-phosphate, MgNH4P04 ,
by Na,HPO, .
!). Estimation. — After removal of other non-alkali metals, magnesium is pre-
cipitated as MgNH,PO4 , then changed by ignition to Mg2P.,07 (magnesiunt
pyrophosphate) and weighed as such. Separated as MgCL from KC1 and NaCl
by solution in uinyl alcohol, evaporated with H...SO, and weighed as MgSO,
I Higgs, AHI. X.. 1S<J2. 44. ]().'{). It is estimated volumetrically by precipitation
as MgNH4PO, . drying at about SO0 until all free NH.OH 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, ./. Am. .Sfoe., 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, CJ, Si, Pb. P|, As, Sb , Bi , Cr, Mo. Mn , Fe , Co, Ni , Cu,
Cd , Zn , Gl , Ba . Sr , Ca , Rb , K , Na , and Li . In some eases the reaction
takes place with 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. r., 1893, 116, 1524).
• Winkler, B., 1890, 23, 44, 120 and 772 ; 1891, 24, 892.
t Kern, C. N, 1876, 83, 112 and 236.
t Seubert and Schmidt, A., 1892, 267, 218.
ANALYSIS OF THE CALCIUM GROUP.
2ir
PH «M C«
t> o O5
O ,^ T~l
p^ ."^ coo
^5 a> a>
<S> <n w
B ;.§.••«
C!
f5
H 0> .S
w •» .2
H ® ^
^ TH CS
I-H COS X>
ANALYS
ABLE FOR
ithod of S
ation given
ss of trace
X cS
g
g B13.A g
£ «3 -53 "S
W^ .. «.S
Jill*
^j eL, .. ' *
x -W 8
Btf|Sf
.H «^ °
n-^£l
"Bd
^^££ £^> ;fc g
«~-g,.H?*!j o| «§« £
*rffp|S3f **$Ǥ,
precipitate appe
produced (§186,
§
&
C
. »
2H3
; if
3*
d -Sa
O 02 O
•rt ^
a ^
03 .,
S D
X >;
o o
'O 05
ff8
11
.5 *
go:
tj
£3 '*"'
i:&
a^
fl
^"N ^
>o^3
'S "**
5 tn
-. J3
^' !S
« ,
05 f
iH £
MCS
^^«C
-5?
•a§S
'3 S«
aj -^ +i
•L f-> ^
a, cs o
57 o cs
2H302
b
I
M g
• • O
e) r
O M
28.
SI
§s
T3 cs
-
05
b-OJ
^ _j
0 §
-t- c.S . K:
3 as
0)
lH
«»
S5
t- 05
r-l iH
vr. vx-
f-> f
05
t-05
tHl-l
tar.ten
•218 niKEUTIOXS FO/f A\M.YNIK 1I777/ NOTES. §191.
DIRECTIONS FOR ANALYSIS OK TIII: MK.TAI.S OK THE CALCIUM GROUP.
(Tin: ALKALINE KAKTIIS.)
§191. Manipulation. — To the filtrate from the fourth group in which
H2S (£192, 1) gi'ves no precipitate (£138) add NH4OH and ammonium
carbonate as long as a precipitate is formed: Bad . -|- (NH4)..CO;! = BaQO,
-f- 2NH4C1 . 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
jind 211).
The well washed white precipitate is dissolved in acetic acid, using as
little as possible: SrCO, + 2HC,H.(02 = Sr(C,H,02), + CO, + H,0 .
To a small portion of the acetic acid solution add a drop of KL,Cr,07 ;
if a precipitate— BaCr04 — is obtained, the K..Cr.,07 must be added to the
whole solution: 2Ba(C:,H,0,) + K,Cr,07 -f H,0 = 2BaCr04 -f 2KC2H30.,
-4- 2HC.,H..O., . Filter, wash the precipitate, dissolve it in HC1 and pre-
cipitate the barium as barium sulphate, with a drop of sulphuric acid.
To the filtrate from the barium chromate add NH4OH and (NH4),C03 ,
warm, filter, and wash. Dissolve the white precipitates of SrC03 and
CaCO., 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 SrS04 may be
moistened with HC1 and the crimsoii 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
HC1 . A white precipitate — CaC,0, — insoluble in acetic acid by its forma-
tion in that solution, and soluble in HC1 is proof of the presence of calcium.
§192. \otcy.-l. The failure of (NH,),S (or H,S in presence of NH4OH) 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 a Tier the addition of ammonium carbonate, as this will drive,
off ammonium hydroxide and carbonate, increasing the solubility of the CaCO,
(note :; and §178).
.?. 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 analyses. The carbonates of barium, strontium and
calcium are all slightly soluble in ammonium chloride solution; and while 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 qmintiliitiiT analyses, if
il94, 4- l)lltK('TIO\X FOIt .LV-LLTN/N \YITH NOTES. 219
barium and calcium art- precipitated as carbonates, it must be done in the
absence of ammonium chloride or sulphate, and the precipitate washed with
water containing ammonium hydroxide.
.'/. If barium be absent, as evidenced by the failure to obtain a precipitate
with K.jCroOT , 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,,Cr,.O7 , 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 (NH,).CO- , an excess of ammonium hydroxide
should lie added to prevent the liberation of CO, when the ammonium car-
bonate is added.
7. Strontium sulphate is so sparingly soluble in water (§187, 5c) that its
precipitation by CaSO, (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 SrSO, will not precipitate sohitions of calcium salts.
Solutions of strontium and barium salts (except SrSO4) are all precipitated
by CaSO,. The presence of excess of calcium salts lessens the delicacy of the
precipitation of strontium salts by calcium sulphate.
X. — In very dilute solutions the sulphates of the alkaline earths are not
precipitated rapidly. Time should be allowed for the complete precipitation.
Boiling and evaporation facilitates the reaction.
!>. It should be noticed that the test for calcium as an oxalate is made upon
that portion of the calcium not removed by K,S04; or in other words upon a
'.solution of CaSO, (1-.">()0). A solution of SrSO, (1-10.000) may be present but
is not precipitated by (NH4);,C,.O4 . The presence of a great excess of
lNH,).SO, prevents the precipitation of traces of calcium salts by (NH4)2C2O4 .
£193. Manipulation. — To a portion of the filtrate from the carbonates
«f Ba . Sr . nnd Ca add a drop or two of (NH4).2S04 and then a few drops of
<NH4),C,04 : filter if a precipitate is obtained and test the filtrate for Mg
with Na.,HP04 . A white precipitate — MgNH4P04— 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
§194. Notes. — /. By some, magnesium is classed in the last or alkali group
instead of in the alkaline earth group. Tt is not precipitated by the (NH4)2C03,
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 #'.).
2. Traces of Ba . Sr and Ca may remain in solution after adding (NH4),CO:i
and warming; due to the solvent action of the ammonium salts present. To
..prevent these traces giving a test for magnesium with Na,HP04 , a drop or
two of (NH^.SO, is added to remove barium or strontium and a few drops of
<NH,),C,O, to remove calcium. The precipitate (if any forms) is removed by
filtration, before the Na.HPO, is added.
./. The precipitate of MgNH,PO, does not always form rapidly if only small
;. mounts of Mg are present, and the solution should be allowed to stand.
Rubbing the side.s of the test tube with a glass stirring rod promotes the pre-
• pitation.
}. The precipitation of Mg as MgNH4PO, is fairly delicate (1-71,492) (Kissel,
Z.. 18(59, 8, 173); but not at all characteristic, 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, 8TROM']IM AXD CALCIUM. $194, /;'.
magnesium depends upon the riyid exclusion of the other metals (not alkalis)
by the previous processes of analysis.
J. Lithium phosphate is not readily soluble in water or ammonium salts and
may give a test for magnesium. See §210, •*/.
§195. The unlike solubilities in alcohol, of the chlorides a, ml 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 HC1 , 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
(be, §§186, 187 and 188).
Residue: BaCl .
Dissolve in water, test
with CnSO, , SrSO. ,
K2Cr:O, , etc.
Filtrate: SrCl, and CaCl .
Evaporate to dryness, dissolve in water, change to
nitrates by precipitating with (NH4)oCO, , 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).
Residue: Sr(NOs), .
Precipitation by CaSO
in water solution:
flame test, etc.
Filtrate: Ca(NOs), .
Precipitation by H.SO,
in alcohol solution, by
<NH4),CS04 , etc.
Or, -the alcoholic filtrate of SrCl2 and CaCl2 may be precipitated with (a
drop of) sulphuric acid, the precipitate filtered out and digested with
solution of (NH4),S04 and a little NH.OH . Residue, SrS04 . SoluHf/n
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 (§145 and ff.) 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 umyl alcohol. Calcium
{titrate is dissolved making a complete separation from the strontium nitrate i$18«, 5r).
$200. THE ALKALI GROUP. 221
THE ALKALI GROUP (SIXTH GROUP).
Potassium. K = 39.11. Caesium. Cs = 132.9.
Sodium. Na = 23.05. Rubidium. Rb = 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 , Rb , K ,
Na , 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 (§6).
Ammonium is the basal radical of ammonium salts, and as such ha8
many of the characteristics of an alkali metal. The water solution of the
gas ammonia, NIL (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 salt*
of the alkali metals (including NH4) are soluble in water, not excepting their
carbonates, phosphates (except lithium), and silicates; while all otlwr
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 metal*,
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 the precipitates they at first pro-
'(luce with salts of certain metals, viz. : the hydroxides of Pb , Sn , Sb (oxide),
* Except those of potassium O.S75 :in:l sodium (0.9735).
'? -2 '.' I>0 TA SSI I'M. § 20 1 .
Zn, Al, and Cr dissolve in the fixed alkalis; and oxide of Ag and hy-
droxides of Cu , Cd , Zn , Co , and Hi dissolve in the volatile alkali.
§201. Solutions of the alkalis are caustic to the taste and touch, and
turn red litmus 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. Sodium nitroferricyanide, with
hydrogen sulphide, gives a delicate reaction for the alkali hydroxides
(§807,66).
§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.
£205. Potassium. K = 39.11. Valence one.
1. Properties.— .S/wr/Y/'r yniritu. 0.875 at i:i° ( Haunihaner, It., 1873, 6, 055).
Multinu ]»ihi4, (i2.1° (Hagen, ('. (.'., 1883, 129). Hoilintt JHt'int, 719° to 731° (Car-
uclley and Williams, /*., 187!), 12, 13(iO); (>«7° (1'ernian, ./. ('., 1889, 55, 328).
Silver-white metal with a bluish tinge. At ordinary temperature of a wax-like
'•onsisteney. duetile and malleable: at 0° it is brittle. It is harder than Na
and is seratehed by Li , Pb . Ca and Sr . The glowing- vapor is a very beautiful
intense violet (Dudley. Am., 1892, 14, 185). It is next to caesium and rubidium,
The most electro-positive of all metals, remains unchanged in dry air, oxidi/rs
rapidly in moist air, anil decomposes water with great violence, evolving
hydrogen, burning with a violet flame. At a red heat CO and CO. are
decomposed, at a white heat the reverse action takes place. Liquid chlorine
does not attack dry potassium ((iautier and ('harpy, ('. •/•., Is'.il, 113, 597). Acids
attack it violently, evolving hydrogen.
'2. Occurrence.— Very widely distributed as a portion of many silicates. In.
M'a water in small amount as KC1 . In numerous combinations in the large
-alt deposits, especially ;:1 Stassfurt: <• »/.. carnal lite. KCl.MgCl, -f (iH^O;
kainite, K.SO,.MgSO,.MgCl + <iH,0 . etc. As an important constituent of
nany i.iants .-.-rape, potato, sugar-beet, tobacco, linuaria. rumcx. oxalis. etc.
.'!. Preparation.- (/) I5\ reduction of the carbonate with carbon. (.!) lly
T-lectroh sis of the hydroxide (Horning and Kasemeyer. It.. 1889, 22, 277c:
Castner. //.. 1M»:.'. 25. i79c). (.?) By reduction of K.CO3' or KOH with iron car-
bide: f.KOH + 2FeC, = GK + 2Fe + 2CO + 2CO; + :$H, (Castner, C. Ar., 1880,
54, 218). (.}) Ry reduction of the carbonate or hydroxide with Fe or Mg
(Winkler. It., 1S90, 23, 44).
4. Oxides and Hydroxide. — Potassium oxide* K,O , is prepared by carefully
* The existence of the oxides M'2O of K, Nn and Rb Is disputed (Erdmann and Koethner, A.,
1896, 2»4 , .V, .
;205, 6&. POTASSIUM. 223
beating potassium with the necessary amount of oxygen (air) (Kuhnemann,
'.<'.. 186.3, 491); also by heating K2O4 with a mixture of K and Ag (Beketoff,
C. C., 1881, 643). It is a hard, gray mass, melting above a red heat. Water
• •hanges it to KOH with generation of much heat. PotdMiinii li ml nu-idf , KOH,
is formed by treating K or K.O with water: by boiling a solution of K,CO,
with Ba , Sr or Ca oxides; by heating KoCO:i with FejO., to a red heat and
decomposing the potassium ferrate with water (Kllershausen, C. C., 1891, (1),
3047; (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. Potaxxium
ftuperoj-idc, K.,O4 , 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. Ar., 1886, 53, 280). It is an amor-
phous powder of the color of lead chromate. Upon ignition in a silver dish
oxygen is evolved and K..O and Ag,O formed (Harcourt, I.e.). 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, 7. c.; Brodie, Proc. Roy. Soc., 1863, 12, 209).
:>. 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 tartrato, silico-fluoride, picrate, phos-
phomolybdate, perchlorate, and chlorate are only sparingly soluble in
fold water, and nearly insoluble in alcohol. The carbonate and sulphate
are insoluble in alcohol.
6. Reactions, a. — Potassium and sodium hydroxides are very strong
bases, fixed alkalis, and precipitate solutions of the salts of all the other
rnetals (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, KoPbO., , potas-
sium plumbite, etc. Potassium carbonate is deliquescent, strongly alkaline,
and precipitates solutions of the salts of the metals (except Cs , Rb , 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.
6.- -The potassium salts of HCN , H4Fe(CN), , H3Fe(CN)(i , and HCNS
find extended application in the detection and estimation of many of the
''•envy metals.
Tartaric acid, H2C4H40C , or more readily sodium hydrogen tartrate,
NaHC4H400 , precipitates, from solutions sufficiently concentrated, potas-
sium hydrogen tartrate, KHC4H40,; , granular-crystalline. If the solution
]>e alkaline, tartaric acid should be added to strong acid reaction. The
must be made in absence of non-alkali bases. The precipitate is in-
• •reased by agitation, and by addition of alcohol. It is dissolved by fifteen
Darts of boiling water or eighty-nine parts water at 25°, b}r mineral acids,
by solution of borax, and by alkalis, which form the more soluble normal
•tartrate, K2C4H400 , but not by acetic acid, or at all by alcohol of fifty
; er cent
-224 POTAssii'M. §205, 6c.
Picric acid, C,.H.,(NO.,)3OH , precipitates, from solutions not very dilute,
the yellow, crystalline potassium jiirnitr. C,VH.,(NO.,);,OK , insoluble in alco-
hol, by help of which it is formed in dilute solutions. The dried precipi-
tate detonates strongly when heated.
r. — If a neutral solution of a potassium salt be added to a solution of cobaltic
nitrite,* a precipitate of the double salt potassium cobaltic nitrite, K,Co(NO,). ,
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, (k-).
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
KC1 (/>., 2, 2, 72). It finds extended application in the manufacture of gun-
l>owder. d. — See §206, (if/.
p.— Potassium sulphide may be taken as a type of the soluble sulphides
which precipitates solutions of the metals of the first four groups as
sulphides except: Hg' becomes HgS ;ind Hg°, Fe'" becomes FeS and S,
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 H..S or soluble
sulphides as a reducing agent see the respective metals. Potassium sul-
phate is used to precipitate barium, strontium, and lead. It almost always
occurs in nature as double salt with magnesium, K.,SO(.MgS01.MgCl2 -f-
<iH,0 . kainite, and is used in the manufacture of KA1(S04)., , K,CO., and
KOH . 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 K..CO, , 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 iodides;
it is also extensively used in medicine. As a type of ;i soluble iodide it
precipitates solutions of the salts of Pb , Ag , Hg . and Cu'. Cu" salts
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 arsenous acid with liberation of iodine.
* One cc. of cobaltous nitrate solution and three cc. of acetic acid are added to five cc. of a ton
per cent solution of sodium nitrite. This gives a yellowish solution having an odor of nitrous
acid.
§205,1. J'OTASSH'M. 225
Potassium chlorate is used as a source of oxygen and as an oxidizing agent
:n acid solutions. Sodium perchlorate, NaClO^ , precipitates from solutions of
potassium salts potassium perch liiratc, KC1O4 , sparingly soluble in water and
ulrnost insoluble in strong alcohol (Kreider. Z. anon)., 189.), 9-. 342). Potassium
iodate is used as a reagent in the detection of barium as Ba(IO:,)... . <j. — 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. — Fluosilicic acid, H2SiF0 , precipitates from a neutral or slightly
acid solution of potassium salts, potassium fluosilicate (silico-fluoricje),
XJ3iFc, 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 HC1 at 14° (Stolba, J. pr., 1868, 103, 396V
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)2PtCl4 , 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.
ft. — An alcoholic solution of BiCl3 in excess of Na.SoO;, 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;, + 2H.O . 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, C. 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 K2CO:! and Na2CO;! 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 + K2C03 =
BaCO., + K2S04 .
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 HC1, 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
J?fcG SODIUM. £205, *!
prism filled with indigo solution, sufficiently thick, entirely cuts oil' tip-
yellow light of sodium, and enables the potassium flame to he seen. The
red rays of the lithium flame are also intercepted by the blue gla —
indigo prism, a thicker stratum being required than for sodium. I"
organic substances are present, giving luminosity to the flame, they mu-t
be removed by ignition. Certain non-alkali bases interfere with the
examination. Silicates may he fused with pure gypsum, giving vapor of
potassium sulphate. Bloxam (J. C., 1865, 18, 329) 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
NH4OH and (NHj,CO, 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 »•, 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). Sodium 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, hem •••
the test should be made after the removal of the heavy metals (;<§2li
and 212).
Potassium may be precipitated as the platinichloride (G/); as the per-
chlorate ((>/); as the silico-fluoridc ((>/); as the acid tartrate ((>&); etc.
Certain of these reactions are much used for the quantitative estimation
(9) of potassium but are seldom used for its detection qualitatively.
(.». Estimation. — (/) Potassium is converted into the sulphate or phosphate
and vfeighed as such. (2) It is precipitated and weighed as the double chloride
with platinum. (3) If present as KOH or K CO it is titrated with standard
acid (Kippenberrer, K. n-mjcic.. 18'.)4, 493). (.}) It is precipitated with H,SiF,,
and strong alcohol. (•!) 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. (6) It is pre-
cipitated as the bitartrate in presence of alcohol and, after filtration and
solution in hot water, titrated with dcci-normal KOH. (7) By precipitation as
the perchlorate. KC1O, (Wense. Z. dinjcir., 1S92, 233; Caspar!. Z. (nujeir., 1893, 68).
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 K204 see 4.
S206. Sodium. Na = 23.05 . Valence one.
1. Properties.— Specific //rrm'///, 0.973:> at i:;.r>° (Baumhauer, B., 1873, 6, 665);
0.7414 at the boiling point (Ramsay, li., 18SO. 13, :»145). Melting point, 97.6°
(Hagen, B., 1883, 16, 1668). BoiUng point, ~42° (Perman, C. N., 1889, 59, 237).
£206, c»l. SODIUM. •>•>'
A silver-white' metal with a strong1 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 iu 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 + 2H.O = :>NaOH + Ha . It burns, when
heated to a red heat, with a yellow flame. Vure dry Na is scarcely at a P.
attacked by dry HC1 (Cohen, O. 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 sonic
mineral waters and in the sea water. It is found in enormous deposits as rock
salt. NaCl; as Chili saltpeter, NaNO3: in lesser quantities as carbonate, borate.
sulphate, etc.
o. Preparation. — (/) By igniting the carbonate or hydroxide with carbon;
(.2). by igniting the hydroxide with metallic iron; (3) by electrolysis of the
hydroxide; (4) by gently heating the carbonate with f/Lg .
4. Oxides and Hydroxides. — Sodium oxide, NaoO , is formed by burning
sodium in oxygen or in air and heating again with Na to decompose the Na,0s
(§205, 4, footnote). Sodium hydroxide, NaOH , is formed by dissolving the
metal or the oxide in water (Rosenfeld, J. pr., 1893, (2), 48, 599); by treating
a solution of sodium carbonate with lime; by fusion of NaNO;, with CaCO. ,
CaO and Na,,CO3 are formed and the mass is then exhausted with water; by
igniting Na.CO., with Fe.O,, , forming sodium ferrate, which is then decom-
posed with hot water into NaOH and Fe(OH)3 (Solvay, C. C., 1887, 829). It is
a white, opaque, brittle crystalline body, melting xinder a red heat. The
fused mass has a sp. gr. of 2.13 (Filhol. A. Ch., 1847, (3), 21, 415). It has a very
powerful affinity for water, gradually absorbing water from CaCL (Muller-
Erzbach, #., 1878, 11, 409). It is soluble in about 0.47 part of water according
to Bintau (C. r., 1855, 41, 509).
Sodium peroxide, Na,O, , is formed by heating sodium in CO, free air or
oxygen (Prud'homme, C. C., 1893, (1), 199). It reacts as H2O2 , partly reducing
and partly oxidizing. It may be fused without decomposition. Water decom-
poses it partially into NaOH and H2OS .
5. Solubilities. — 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 fluosilicato
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.
O. Reactions, a. — As reagents sodium hydroxide and carbonates act in
all respects like the corresponding potassium compounds, which see.
/;. — By the greater solubility of the picratc and acid lartrntc- of sodium, that
metal is separated from potassium (§205, G?>). c. — Sodium nitrate occurs in
nature in large quantities as Chili saltpeter, used as a fertilizer, for the manu-
facture of nitric acid, with KC1 for making KNO3 , etc.
- (1. — Sodium phosphate, Na,HP04 , 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, 5c), soluble in acids). Solu-
228 SODIUM. §206, fa:
tions of alkali phosphates precipitate solutions of all other metallic salts
as phosphates (secondar}r, tertiary or basic) except : HgCL precipitates as
j\ basic chloride (§58, (id), and antimony as oxide or oxychloride (^70, M).
e, f, ff, h. — As reagents the sodium salts react similar to the corresponding
potassium salts, which see. i. — Sodium fluosilicate is soluble in 15:t.:5 parts
H2O at 17.5° and in -40.06 parts at 100° (Stolba. Z., 1872, 11, 199); hence is not
precipitated by fltiosilicic acid except from very concentrated solutions
(separation from K). /. — Sodium platinic chloride, (NaCl)..,PtCl, , crystallines
from its concentrated solutions in red 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 glass 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.
A-. — Solution of potassium pyroantimonate, K.,H.,Sb.,07 , produces in
neutral or alkaline solutions of sodium salts a slow-forming, white, crystal-
line precipitate, XaL.H.,Sb.,0T , almost insoluble in cold water. The reagent
must be carefully prepared and dissolved when required, as it is not per-
manent in solution (§70, 4c).
7. Ignition. — Sodium bicarbonate, NaHCO, , loses H,0 and CO., at 125°
becoming Na.,CO:t , no further decomposition till 400° when a very small
amount of NaOH is formed (Kirsling, 7,. angeir., 1889, 332).
Sodium compounds color the flame intensely yellow, the color being
scarcely affected by potassium (at full heat), but modified to orange-red
by much lithium, and readily intercepted by blue glass. Infusible com-
pounds may be ignited with calcium sulphate. The test is interfered with
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 in
the yellow of the solar spectrum separable into two bands, D, and D//? by
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 metal-. In the usual process of analysis
the presence or absence of sodium is determined in the presence of
magnesium (as Na.,HP04 is the usual reagent for the detection of mag-
nesium, it is evident that the presence or absence of the sodium must be
determined before the addition of that reagent); and as that metal gives
a yellowish color to the flame it must be removed if small quantities of
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.
filtration the barium is removed by (NHJ..CO., or H2S04 and the filtrate
tested for sodium by the flame or by the pyroantimonate test
9. Estimation. — (1) If present as hydroxide or carbonate, by titration with
standard acid (Lunge, Z. angew., 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 AgNO3 and computing the amounts of K and Na . (4 It is precipitated
by K.H2Sb,O7 and dried and weighed as Na2H,Sb207 .
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
H202 in alkaline solution, which see (§244, 6).
§207. Ammonium. (NH4)'. Valence oiie.
1. Properties. — Specific gravity of NH3 gas, 0.589 (Fehling, 1, 384); of the
liquid, 0.6234 at 0° (Jolly, A., 1861, 117, 181). The liquid boils at — 33.7°, at
0° 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-
bines energetically with acids to form salts, the radical NH4 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, 99,
144). One gram 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, e. g., 8A1 + 5KOH + 3KNO3 + 2H,O =
8KA102 + 3NH3; by the reduction with the hydrogen of the zinc-copper couple;
by boiling organic compounds containing nitrogen with KMnO4 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 salfwith 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° contams 28.33 per cent of NH3 (Lunge and Wiernik,
Z. angew., 1889, 183).
5. Solubilities.— Ammonia, NH3 , 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.
280 AMMONIUM. $207, 6a.
6. Reactions, a. — The fixed alkali hydroxides and carbonates liberate
ammonia, NH3 , 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 NH4OH . 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, Jc and I).
With salts of Ag , Cu , Cd . Co , Ni , and Zn the precipitate redissolves in
excess of the reagent. Ammonium carbonate, (NN4)2CO:1 , 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.
b. — 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 K,CO;. upon ignition. Sodium nitroferri-
cyanide, Na,Fe(NO)(CN)5 , added to a mixture of NH4OH and H2S
[ (NH4).,S ] gives a very intense purple color, characteristic of alkali
sulphides and the manipulation may be 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 H2S 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 H2S ;
and this solution added to the Na,Fe(NO)(CN). .
c. — Ammonium nitrite, NH4NO, , is used in the preparation of nitrogen
(§235, 3); ammonium nitrate in the preparation of nitrous oxide, N3O ,
" laughing: gas " (§237). d. — Ammonium phosphate, as a reagent, acts
similarly to sodium phosphate. When sodium phosphate, Na2HPO4 , is used to
precipitate metals in the presence of ammonium hydroxide, a double phosphate
of the metal and ammonium is frequently formed as MnNH4P04, MgNH4PO4,
etc. By some chemists microcosmic salt, NaNH4HPO4 , is preferred to sodium
phosphate. Na;,HP04 , as a reagent.
e. — When ammonium hydroxide is saturated with H2S , ammonium sul-
phide, (NH4)2S , is formed. Complete saturation is indicated by the failure
§207, 6k. AMMONIUM. 281
to precipitate magnesium salts, that is, NH4OH precipitates magnesium
salts while (NH4)2S does not. Freshly prepared ammonium sulphide is
colorless, but upon standing becomes yellow with loss of ammonia and
formation of the poly-sulphides, (NH4).,SX . 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 Sn" are with great difficulty
soluble in the normal ammonium sulphide, but readily soluble in the
poly-sulphide. ISnckel sulphide, NiS, is insoluble in normal ammonium
sulphide but is sparingly soluble in the yellow poly-sulphide (distinction
from cobalt). (NH4)2S gives a rich purple color with sodium nitroferri-
cyanide (&). Ammonium sulphate as a precipitating reagent acts similar
to all soluble sulphates (§205, 6e). A 25 per cent solution of (NH4)2S04
is used to dissolve CaS04 (§188, 5c) (distinction from Ba and Sr).
f.— Ammonium chloride is much used as a reagent. It prevents pre-
cipitation of the salts of Mn by the NH4OH , and is of special valuejn 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(OH)3 is diminished
by the presence of NH4C1 (§124, Ga, and §117). *
g, h. — Similar as reagents to the corresponding1 potassium salts, i. — Fluo-
silicic acid, H2SiFa , does not precipitate ammonium salts, the ammonium
fluosilicate being- very soluble in water (distinction from potassium). ;'. — Plat-
inum chloride, PtCl4 , forms with ammonium salts the yellow double ammonium
platinum chloride, (NH4)2PtClti , very closely 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 KC1 .
k. — A solution of potassium mercuric iodide, K2HgI4 , containing also
potassium hydroxide— Nessler's test *— produces a brown precipitate of
nitrogen dimercuric iodide, NHg2I , dimercur-ammonium iodide (§58, 6a),
soluble by excess of KI and by HC1 ; not soluble by KBr (distinction from
HgO):
NH3 + 2HgI2 = NHgvI + SHI
NH4OH + 2K2HgI4 + 3KOH = NHg.,I + 7KI + 4H20
* This reagent may be prepared as follows : To a solution of m rcuric chloride add solution
of potassium i dido till the precipitate is nearly all redissolved ; then add solution of potassium
hydroxide sufficient to liberate ammonia from ammonium salts ; leave until tho liquid becomes
clear, and decant from any remaining sediment.
232 AMJUOMIM. §207,6?,
This very delicate test is applicable to ammonium hydroxide or salts;
traces forming only a yellow to brown coloration. The potassium mercuric
iodide, " Meyers lleagent," 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 Nessler's reagent, but from this solution a precipi-
tate is formed with HgCl, (De Koninck, X., 1893, 32, 188).
/.-Mercuric chloride, HgCL , forms, in solutions of ammonium hy-
droxide or ammonium carbonate, the "white precipitate" of nitrogen
dihydrogen mercuric chloride, NH.,HgCl , 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 substance?
(Wittstein, Arch. Pharm., 1873, 203, 327).
in. — Add a small quantity of recently precipitated and well-washed silver
chloride, and, if it does not dixmlre after agitation, then add a little potassium
hydroxide solution. The solution of the AgCl , before the addition of the fixed
alkali, indicates free ammonia; niter the addition of the fixed alkali, ammonium
salt. (Applicable in absence of thiosulphates, iodides, bromides and sulpho-
eyanates.)
H. — Sodium phosphomolybdate (§75, Crf) precipitates ammonium from neutral
or acid solutions; also precipitates the alkaloids, even from very dilute solu-
tions, and, from concentrated solutions, likewise precipitates K, Eb and Cs
(all the fixed alkalis except Na and Li).
7. Ignition. — Heat vaporizes the carbonate, and the haloid salts of am-
monium, undecomposed ('dissociated but reuniting1 upon cooling); decomposes
the nitrate with formation of nitrous oxide and water, and the phosphate and
borate with evolution of ammonia. NH, heated to 780° or hig-her is dissociated
into N ai d H (Hamsay and Young, J. C., 1884, 45, 88).
8. Detection. — As ammonium hydroxide and chloride are used in the
regular process of analysis, the original solution must be tested for the
presence or absence of ammonium compounds. The hydroxide or the
carbonate may be detected by the odor (1); the action on red litmus paper
suspended in the test-tube above the heated solution; the blue color im-
parted to paper wet with copper sulphate; the blackening of mercurous
nitrate paper; and if in considerable quantity, the white vapors when
brought into contact with the vapors of volatile acids. In combination
as salts the gas is liberated by the fixed alkali hydroxides or carbonates
(oxides or hydroxides of Ba, £r , or Ca may be used) and distilled into
Nessler's reagent, or collected in water and the test with HgCl2 (Gl) applied
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
§208, :>. CAESIUM. 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" (NCl3?) (a).
The same product is liable to arise from solid ammonium salts treated with
chlorine. Gaseous ammonia, and ammonium hydroxide, with chlorine gas,
generate free nitrogen (6), a little ammonium chlorate being formed if the
ammonia is in excess. Hypocliloritcs or hypobromites (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 ft) ; 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 (d). 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 (#). 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 (ft); 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.
(a) NH4C1 + 3C12 = NC13 + 4HC1
(&) 8NH3 + 301, = 6NH4C1 + N2
2NH4C1 + 3C1, = 8HC1 + N3
(c) 2NH, + I2 = NH4I + NH2I
(d) (>NH4OH + 31, — 5NHJ + NH4I03 + 3H2O
(e) <>NH4OH + 8HMn04 = :JNH4NO3 + SMnO(OH)2 + 5H2O
(f) 12Cu + 2NH3 + 9O2 = 12CuO + 2HNO2 + 2H2O
(fir) 3HN03 -f 8A1 + 8KOH = 8KA1O2 + 3NH, + H20
(ft) HCN + KOH -f H,0 = NH:, + KCH02 (formate).
§208. Caesium. Cs = 132.9 . Valence one.
1. Properties.— Specific 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 imder 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., 18G4, 58, 715); in
many mineral springs (Miller, C, N., 1864, 10, 181); in the ash of certain plants,
tobacco, tea, etc.
?,. Preparation. — By electrolysis of a mixture of CsCN with Ba(CN)2; by
ignition of CsOH with Al in a nickel retort (BeketofE, C. C., 1891, (2), 4oO).
4. Oxide and Hydroxide. — An oxide has not yet been prepared. The
hydroxide, CsOH , is a grayish-white solid, very (leliiiiicscoit, absorbs CO2 from
the air; dissolves in water with generation of much heat, forming a strongly
caustic solution.
f>. Solubilities. — Caesium dissolves with great energy in water, acids or
alcohol, liberating hydrogen and forming the hydroxide, salts or alcoholate
respectively. The hydroxide is soluble in water and alcohol. The salts are
all quite readily soluble. The double platinum chloride, Cs2PtCl4 , and the
acid tnrtrate, CsHC,H40,, , being least soluble and used in preparation of the
salts free from the other alkali metals.
234: RUBIDIUM— LITHIUM. §208, G.
6. Reactoiis. — In all its reactions similar to the other fixed alkalis.
7. Ignition. — Caesium salts color the non-luminous flame violet. The spec-
trum gives two sharply defined lines, Cs u and Cs ,i, in the blue and a third
faint line in the orange-red Cs , , also several taint lines in the yellow and
green. \Vith the spectroscope three parts of CsCl may be detected in presence
of 300,000 to 400,000 parts KC1 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. — (1) As the double platinum chloride; (2) as the chloride with
RbCl , estimation of the amount of Cl and calculation of the relative amounts
of the metals; (.<) as the sulphate obtained from ignition of the acid tartrate
and treatment with H2SO4 (Bunsen, Pogg., 18G3, 119, 1).
§209. Rubidium. Rb = 85.4 . Valence one.
1. Properties.— Specific yniritu, 1.52 (Bunsen, A., 1863, 125, 367). Melting
point, o8.5°; 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.
:;. Frepaiation.- From the mother liquor obtained in the preparation of Li
salts (Heint/., J. pr., 1862, 87, 310): (/) By ignition of the acid tartrate with
charcoal; (^) electrolysis of the chloride; (J) by ignition with Mg or Al
(Winkler, li., 1SOO, 23. 51: Beketoff, II., 1SSS, 21, c, -424).
4. Oxide and Hydroxide. — The oxide Rb.O has not been with certainty pre-
pared. The hydroxide. RbOH . is formed when the metal is decomposed by
water; also through the action of Ba(OH), upon Rb..S04 . 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
water 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 perehlorate. the fluosili-
cate, the double platinum chloride, the silicotungstate, the picrate, and the
phosphomolybdate. The alum is less soluble than the corresponding potassium
alum.
:''. Reactions. — Similar to the other fixed alkalis.
7. Ignition.- The salts give a violet color to the flame. The spectrum gives
two characteristic lines in the violet, Rb a and Rb ,1; two less intensive in the
outer red, Rb 7 and Rb rf; a fifth Rbf in the orange; and many faint lines in the
orange, yellow and green. As small a quantity as 0.0000002 gram of RbCl can
be detected (Bunsen, 7. f.).
8. Detection. — By the spectroscope (7 and §210, 7).
9. Estimation.— '(/) By weighing with CsCl as the chlorides, determining the
amount of Cl and calculating the proportion of the metals; (2) as the double
platinum chloride.
£210. Lithium. Li = 7.03 . Valence one.
1. Properties.— Rpirifir (jnirUij. 0. .-,<):;<•,. the lightest of all known solid bodies
(Bunsen and Matthiessen, '.-!., 1855, 94, 107). MrUiny point, 180° ; does not vaporize
at a red heat. It is a silver-white metal with a grayish tinge: harder than
X 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 chemictil 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 Cl , Br , I , O , i and dry CO2 . 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;
€. 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 NH4C1 or KC1
(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, LLO , 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 sohible in water with evolution of
hydrogen, forming the hydroxide; soluble in acids with formation of salts.
The oxide, Li2O , dissolves in water, forming the hydroxide. The most of the
lithium salts are soluble in water. A mimber 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. Soc. Ind.,
1888, 7, 730); Li, CO., in 75 parts at 20°; Li3PO4 in 2539 parts pure water and
3920 parts ammoniacal water, more soluble in a solution of NH4C1 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 NH4C1 solution than in pure water (distinction from mag-
nesium). In dilute solutions the phosphate is not precipitated iintil 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, .1. Cli., 1818, (2), 7,
157). The phosphate dissolved in HC1 is not at once precipitated by neutraliz-
ing with NH.OH (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-mi 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 a, and a faint orange line, Li /3. The color
tests have an intensity intermediate between those of sodium and potassium.
8. Detection. — BII the spectroscope. — To the dry chlorides of the alkali metals
a few drops of HC1 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 B/b 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 Bib and Cs appear.
Evaporate to dryness the filtrate from the precipitate of the platinum double
salts, add oxalic acid and ignite, moisten with HC1, 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 DIRECTIONS FOR ANALYSIS WITH NOTES. §210, 9.
9. Estimation. — After separation from other elements it may be weighed as
a sulphate, carbonate or phosphate, Li3PO4 . It may also be estimated by the
comparative intensity of the lines in the spectroscope (Bell. Am., 1886, 7, 35).
DIRECTIONS FOR THE ANALYSIS OF THE METALS OF THE ALKALI GROUP
(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 (NH4)2S04 and (NH4)2C204 (§193), may at once be tested for the pres-
ence of potassium and sodium. 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, i. e. Na.,HP04 .
It is recommended by many to use ammonium phosphate. (NH4)2HP04 .
This reagent removes the magnesium, and permits the application of the
flame test for the fixed alkalis; but the presence of tho phosphate obstructs
the gravimetric determination of the alkalis. The phosphate may be
removed by lead acetate and the excess of the lead by hydrogen sulphide.
§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 Ba(OH), to
precipitate the magnesium as Mg(OH)2 (§§177 and 182). After filtration,
the excess of barium in the filtrate is removed by H2S04 , 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 EOH or NaOH to the solution and warm gently. Notice
the odor (§207, 1). Suspend a piece of moistened rod 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
Nessler's Reagent (§207, Gk) or by the precipitation with HgCl2 (§207, 61).
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).
§215. Lithium, because of the insolubility of its phosphate (§210, 5c),
interferes with the detection of magnesium. If the nitrate 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 sulphuric
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 Minus Numbers,* 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+^l"1.
That of water, (H+I).,0~11. (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.f
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 the sum of the bonds of two elements
which act upon each other.
The use of thjs notation is illustrated in the following equations :
::SnCL + fi,S03 + fiHCl = ,!SnCl4 + H2S + 3H20
In this equation the three atoms of tin gain six bonds; the bonds of the
sulphur in the H,SO:; have then been diminished by six; that is, it has
given up six bonds to the tin,' and having only four in the first place must
now have minus two (4 -6 = -2).
* O C. Johnson, C. JV., 1880, 42, 51. See also Ostwald, Orundr. allg. Chem., 3te Aufl., 1899, 8. 439.
tlf there is polarity in the union of like atoms with each other in forming an elemental
molecule, the sum must bo 7.ero, as in the formation of the molecules of compounds.
§217, f. BALANCING OF EQUATIONS. 239
.3SnCl, + HI03 + GHC1 — :iSnC!4 + HI + ;H,0
Here also the three atoms of tin gain six bonds, and these are furnished
by the iodine of the HI03 . It has five in the first place, and being
diminished by six, has one negative bond remaining (5-6 = -]). [In
other words, unless we deny that iodine has five bonds in HIO, , we must
admit that it has one negative bonfl in HI (written H'I~').]
8HMn04 + 5AsH3 + 8H,S04 = 5H3As04 + 8MnSO4 + 12H.O
In this equation eight atoms of manganese in the first member have 56
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 oT arsenic in arsenous hydride has
-3 bonds (As-"'Hs).
SnCl2 + HgCl2 — Hg + SnCl4
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 HgCl2 .
§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 Pb is -zero, but in PbCl2 the lead has two posi-
tive bonds, and each atom of chlorine has one ne.gati.ve bond.
In Bi2S3 , each atom of Bi has three positive bonds (e)r and each atom of
S has two negative bonds (f).
In ammonium nitrite, NH4N02 , or H4= N — 0 — N = 0 , the nitrogen
of the NH4 has four negative bonds and one positive bond. The other
nitrogen, that of the acid radical N02 , has three positive bonds. Each
atom of hydrogen has one positive bond and each atom of oxygen two
negative bonds, the sum being zero : -4-4 — 4-4-1 + 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 fol-
240 BALANCING OF EQUATIONS. §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 seen to be aero.
Hg"(NvO-"a)2.2 + 10 — 12 = 0
Bi'"2 (SviO-"4)3.6 + 18 _ 24 = 0
Ba"(MnviiO-"4)2.2 + 14 — 16 = 0
Fe'"(NvO-"3)3.^ + 15 — 18 = 0
As'"2S— "3.6 — 6 = 0
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.
(1) 3As4 + 20HNO, + 8H,O = 12H3AsO4 + 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 As4
are taken.
(2) (>Sb + 10HNO, = .sSbjO; -f 10NO + :>H3O
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 Sb.O. with an odd number of atoms of
antimony, we double the ratio and take six and ten.
(3) :5H2S + SHNO3 = liH.SO, + sNO + 4H,0
The S in the first member has 2 negative bonds (« and </); in the second
member it has fi positive, gaining s bonds: hence 8 molecules of HNO:1 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 HaS
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 com-
pound; in the second member it has six bonds (valence of six), but positive
because combined with oxygen (S03 or -pro S ~r\)- The valence of the
hydrogen 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;Or. 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 (,S) must equal the bonds lost by the HNO3 (ft).
The least common multiple, twenty-four, indicates the least possible total
change of valence for each compound: this requires that three molecules of
H,S and eight of HNO3 be taken, giving for the products three molecules of
H.SO4 and eight of NO with four of water to complete the equation.
(4) 3Sb38a + 28HNO:: = riSb.O, + 9H,S04 + 28NO + 5H,O
In this case, both the Sb and the S in the molecule gain bonds, and must be
§218, 12. BALANCING OF EQUATIONS. 241
considered. It is plain (from d and c) that each atom of Sb gains 2 bonds, and
the two in the molecule will gain 4.
The S in SbjS:, has 2 negative bonds, and in the second member (in H SO,)
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 Sb..,S3 ; hence 28 molecules of HNO,
must be taken. We take 3 of Sb2S3 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 a
gain for one molecule of Sb.Sa 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 Sb2S3 three mole-
cules, and for the HN03 twenty-eight molecules. The products are as indicated
in the equation.
(5) 2Ag3AsO4 + HZn + 11H,S04 = 2AsH3 + 6Ag -f HZnS04 + SH3O
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 + 5Pb304 + :;OHNO3 = 2HMn04 + loPbCNO,). + 14H20
The manganese gains five bonds, therefore five molecules of Pb3O4 are taken.
1 The three atoms of lead in one molecule of Pb3O4 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:! + TPbO, + 1-4HNO, = 2HMn04 + 2Br2 + 7Pb(NO3), + GH2O
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 Pb02 are taken. The lead loses two,
therefore two of MnBr2 are taken.
(S) MnS + 4KNO3 + K2C03 , fusion = K,Mn04 + K,S04 + 4NO + K.,CO»
The manganese gains four bonds and the sulphur eight, making twelve;
therefore twelve of KN03 would be taken, and since the nitrogen loses three
bonds, three of MnS would be taken, but since three is to twelve as one iis to
four, the latter amounts are taken.
(9) 2Cr(OH)3 + 3Mn(N03)2 + 5K2C03 , fusion = 2K2CrO4 +
2K2MnO4 + GNO + 5CO2 + 3H2O
.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(OH)3 are taken. The chromium gains three, therefore
three of Mn(NOs)2 are taken.
(10) 3Ag + 4HNO3 = 3AgN03 + NO + 2H2O
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) 2FeI2 + 6H2SO4 , cone., hot = Fe2(S04)3 + 3S02 + 2I2 + 6H20
The rule here calls for two of FeI2 and three of H2S04 , but three more of
H SO, that are not reduced are needed to combine with the iron, making six
in all.
(12) 3HNOS + 8A1 + 8KOH = 3NH3 + 8KA1O2 + H20
The nitrogen has five bonds in HN03 , and in NHS it has minus three,
losing eight, therefore eight of aluminum are taken. The aluminum gains
three, therefore three of HNO, are taken.
242 BALANCING OF EQUATIOXs. §218. 13.
(13) 3BiONOs + HAl + 11KOH — :iBi + 3NH, + UKAIO, + H,O
The bismuth loses three bonds and the nitrogen loses eight, therefore eleven
of aluminum are taken; the aluminum , gains three, therefore three of the
BiONOg are taken.
(14) MnO. + 4HC1 = MnCL + CL -f 2H2O
The manganese loses two bonds and the chlorine gains one, but two more of
unoxidized HC1 are needed to combine with the manganese, hence four are
taken.
(15) 2CrI3 + C.4KOH + 27C1, = 2K,CrO, + (iKIO, + 54KC1 + 32H3O
The chromium gains three bonds and the iodine (in the molecule) gains
twenty-four, therefore twenty-seven of C12 are taken and the Cl, loses two,
therefore two of CrL, 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-4H'«. —4+4 = 0
CHSC1 C-3 + /H'8Cl-'. -3+1 + 3-1 = 0
CH*C1. C-' + sH'aCl-',. -2+2+2 — 2 = 0
CHOI, C-I + 1H/C1-',. —1 + 3+1 — 3 = 0
CCL C4Cl-'<. 4 — 4 = 0
HCjH.O, H'(C,) + 3— ''H'sO-V 1 + 3 — 3+3 — 4 = 0
CaHaO (ai'-'H'eO-*. 1 — 5+6-2 = 0
C3HeO, (C3)-' »-'H'BO-2,. -6 + 3 + 8 — 6 = 0
C6HiaO, (C,.)-717H',,O-S,. -7+7+12—12 = 0
(/) CH4 + 4CL = CC1, +
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) ^C2H0O + 2X,Cr,O7 + SH1,SOi = 3HC*H,O2 + 2K,S04 +
2Cr2(S04)8 + 11H.O
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 four
molecules of the potassium dichromate. Thi two atoms of the chromium lost-
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) 3C3HbO3 + HHNO3 = 9CO, + 14NO + 19H,O
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.
(//) C0H12O0 + 12H,S04 = (CO, + 12SO2
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 the non-metallic elements will be de-
scribed in the order of their atomic weights; and the acids in the order
of the degree of oxidation of the characteristic element, e. g., N before S ,
HC1 before HC10 , HC10, before HC10, , etc.
§219, 6. HYDROGEN. 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 (jracitu, 0.06949 (Crafts, C. r., 1888, 106, 1062). 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 (Olszewski, C. r., 1884, 99, 133; 1885, 101, 238; Wroblewski, C. r.,
1885, 100, 979). Critical temperature , — 234.5°; critical pressure, 20 atmospheres;
boiling point, —213.5° (Olszewski, PMl. May., 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 manj^ metals, especially
palladium; which, heated to 100° in an atmosphere of Irydrogen and then
cooled in that atmosphere, absorbs at ordinary temperatures 982.14 volumes of
hydrogen (Graham, J. C., 1869, 22, 419). This occluded hydrogen acts as a
strong reducing agent, reducing FeCl3 to FeCL , 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 burns with a non-luminous flame and with generation
of much heat (more than an equal weight of any other substance or mixture
of substances). Hydrogen forms two oxides: water, H2O , and hydrogen
peroxide, H.O. (§244).
2. Occurrence. — In volcanic gases (Bunsen, Poyg., 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 inmimerable minerals (H^O and OH) and in organic compounds.
3. Formation. — (a) By the reaction of alkali metals with water. (6) 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
HN03). By heating potassium formate or oxalate with KOH : K2C204 -}-
2KOH = 2K2CO:, + H, (Pictet, A. Cli., 1878, (5), 13, 21G).
4. Preparation. — (a) By the action of dilute sulphuric acid (one to
eight) on commercial or platinized zinc * (§135, 5a). The solution must
he kept cold or traces of S02 and H2S will be evolved. (&) By the elec-
trolysis of acidulated water.
5. Solubilities. — Water at ordinary temperature dissolves nearly two per
cent (volume) of hydrogen. Charcoal dissolves or absorbs fully ten times its
volume of the gas (1).
6. Reactions. — 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 hydrogen 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 platinum chloride ; the zinc is made
readily soluble in dilute sulphuric acid and a uniform and rapid generation of hydrogen can be
obtained.
244 HYDROGEN. §219, 7.
conditions combines with O , S , Se , Te , Cl , Br , I , N , P , As , Sb and Si
with comparative readiness. The reduction of salts by nascent hydrogen in
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 KOBE and
Al does not reduce AsV; that formed by zinc and acids gives AsH3 . Sbv
with sodium amalgam and acids gives Sb°; with zinc and acids, SbH:, (Chu-
brier, C. r., 1872, 75, 484; Tommasi. BL. 1882, (2), 38, 148).
Hydrogen occluded in metals as Pd , Ft, etc., is even more active than
" nascent hydrogen "; often causing combination with explosive violence
(Berthelot, A. Ch., 1883, (5), 30, 711); 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
Or'"; KC103 to KC10; CH3CO2H to CH.CHO and C2H-,OH; and CaH,lTO, to
C8H5NH2 . The reactions are quantitative. Salts of Pb . Bi . Cd , As , Sb , W.
Mo , Zn , Co , Ni , Al , Ce , U , Rb , Cs , K , Na , Ba , Sr and Ca are not reduced
(Schwarzenbach and Kritschewsky, /?., 1S8C>, 25, 374). In the presence of
platinum black hydrogen reduces very much as described above; also K3Fe(CN),,
becomes K4Fe(CN)n; dilute HNO., becomes NH4NO2 . concentrated HNO, be-
comes HNO;.; Cl , Br and I combine with the hydrogen in the dark; KC1O,
and KC10 are reduced to chlorides. KC10, is not reduced; H,SO< , concen-
trated, is reduced to H SO (Cooke, C. AT., 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
Ag-NCX; hindered by HNO, or KNOn . Solutions of Au , Pt and Cu are also
acted upon (Russell, J. C., 1874, 27, 3; Leeds, B.. 187<>, 9, 1450; Reichardt, Ar<-li.
Pharm., 1883, 221, 585; Poleck and Thuemmel, B., 1883, 16, 2435; Senderens, Bl.,
1897, (3), 15, !)!)!). KMnO, in acid, neutral, or alkaline solution slowly
oxidizes hydrogen. It is not at all oxidized by nitrohydrochloric acid, in
diffused daylight, CrO3 , at ordinary temperature'. Fed, /KnFe(CN)0 , HNO., ,
sp. gr. 1.42, or H SO, , xp. fir. 1.84 (Wanklyn and Cooper, Phil. .!/«//., 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; €. ff.,
hydrogen at 100 atmospheres pressure precipitates Hg° from HgCl2 (Loewcn-
thal, J. pr., I860, 79, 480).
7. Ignition. — Chlorine and bromine combine with bydrogen directly in
the sunlight, but heat is required to effect its combination with iodine,
fluorine, and oxygen.
All oxides, hydroxides, nitrates, carbonates, oxalates, and organic salts
of the following elements are reduced to the metallic or elemental state by
ignition in hydrogen gas : Pb , Ag . Hg , Sn . Sb , As . Bi , Cu , Cd , Pd ,
Mo , Ru , Os , Eh , Ir , Te , Se , W , Fe , Or , Co , Ni , Zn , Tl , Nb , In , V
Compounds of aluminum, manganese, and of the fifth and sixth group
metals have not been reduced by hydrogen.
8. Detection. — (a) Method of formation it' known. (6) Its explosive
union with oxygen when the mixture with air is ignited, (c) Absorption
by palladium sponge, (d) Explosive union with chlorine in the sunlight
to form HC1 . (e) Separated from most other gases by its non-absorption
by the chemical reagents used in gas analysis.
9. Estimation. — By volume niensurement, almost never by weight, except
when determined in its compounds by combustion to H.O .
§221, 4. BORON— BORIC ACID. 245
§220. Boron. B = 11.0 . Valence three (§2).
Boron does not occur free in nature. It is found chiefly as borax, Na2B4Or ,
and as boric acid, H:BO , 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, A., 1867, 141, 268). Elemental boron is prepared (o) by electrolysis;
(ft) by fusing B2O., with Al , Na or Jff.g; (c) by igniting BC13 with hydrogen;
(d) by fusing borax with red phosphorus. Specific orarity of the crystalline,
2.53 to 2.68 (Hampe, A., 1876, 183, 75); of the amorphous, 2.45. Amorphous
boron is a greenish-brown, opaque powder, odorless, tasteless, insoluble in
water, alcohol or ether. It is a non-conductor of electricitj'. Heated in air or
oxygen it burns with incandescence. In air it forms B2O3 and BN . It is
oxidized by molten KOH or PbCr04 , with incandescence. It is dissolved by
concentrated HN03 or H2S04 , forming boric acid. At a red heat it decom-
poses steam. When. heated it combines directly with S , Cl , Br , N and many
metals. It forms BC13 with chlorine, not BC15 . Fused with P2O5 it forms
B,03 and P; with KOH, K;(BOS and H; with K2CO3 , K;BO and C. Boron
forms but one oxide, B203 , boric anhydride. Three hydroxides are known:
2H:tBO, = B2O:i.3H2O , orthoboric acid; 2HB02 = B2O3.H20 , metaboric acid; and
H2B4O7 =: 2B2O3.H,,O , pyroboric acid.
§221. Boric acid. H3B03 = 62.024 .
H'3B'"0-"3, H-0-B~o~£
1. Properties. — Boron trioxide, B2O3 , fton'.c anhydride, is a brittle vitreous
mass; sp. gr. at 12°, 1.8476 (Ditte, A. Ch., 1878, (5), 13, 67). Melting point, 577°
(Carnelley, J. C., 1878, 33, 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,
85, 1069). In some respects boron trioxide deports itself as a weak base. It
forms a sulphide, B2S3 , decomposed by water (Woehler and Deville, A. Ch.,
1858, (3), 52, 90); a sulphate, B(HS04)3 (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, 1. c.~). Soluble in 25 parts
water at 20°, and in 3.4 parts at 102° (Ditte, 7. c.)." 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, Mg7B10030Cl2 (62.5
per cent B203). In the volcanic regions of Tuscany and the Liparic Islands as
steam saturated with boric acid.
3. Formation. — The anhydride is formed by burning the metal in air
or oxygen, or by heating the acids. Orthoboric acid, Hr>BO, , is formed
by dissolving the oxide in water; the meta acid, HB02 , H — 0 — B = 0 ,
by heating the ortho acid a little above 100° (Bloxam, J. C., 1860, 13,
177); the pyroboric acid, tetraboric acid, H2B407 , by heating the ortho
or meta acid for some time at 160° in a current of dry air (Merz, J. pr.,
1866, 99, 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. £221, 5.
from water. (&) In Nevada the boronatrocalcite, Ca2B0011.Na.,E407 -f-
18H20 (45.6 per cent B203), is evaporated in lead pans with HJ504 to a
stiff 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, Na.,B407.10H.,0 , is dissolved in hot water, twelve parts, and acidi-
fied with hydrochloric acid. Upon cooling, the boric acid, H(BO.t , 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 burns with a beautiful green
flame. 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. Reactions. — Silver nitrate forms, in solutions of acid borates, a white
precipitate of silver borate, AgBO., , but normal borates form in part silver
oxide, brown. Lead acetate gives a white precipitate of lead borate,
Pb(BO.,).> ; 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(BO.,)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 yreen 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 CO., ; but being non-volatile except at a very high heat, it dis-
places most other acids upon ignition.
By 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. CARBOX.. 247
If a powdered borate (previously calcined), is moistened with sulphuric-
acid and heated on platinum wire to expel the acid, then moistened witli
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 3 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 colored glass,
used in the detection of certain metals (§132, 7). The fused borax form.-;
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 H2S and barium salts
should be removed or converted into the sulphate). Also by the red color
imparted to a strip of turmeric paper.
9. Estimation. — 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 Na,,CO3 , fusing
and weighing; then after determining the CO.., subtracting its weight and
that of the Na2O present (calculated from Na2CO3 first added). The differ-
ence is the weight of B2O3 present. See also Will (Arch. Pliartn., 1887, 225, 1101).
In the presence of glycerine, boric acid may be accurately titrated with sodium
hydroxide, using phenolphthalein as an indicator: B.Oj + 2NaOH = 2NaBO2 +
H2O . Sodium carbonate must be absent or we get: 2B2O3 -j- Na2CO3 —
Na,B4O, + CO2 (Honig and Spitz, Z. anyeic., 1896, 549; Joergensen, Z. anyew.,
1897, 5).
§222. Carbon. C = 12.0 . Usual valence four (§15).
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, J., 1873, 237); graphite, Ceylon, 2.25 to 2.26
(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. CJi., 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 CO.., .
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, ?.
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, etc.
Amorphous carbon is black, lighter than diamond or graphite. 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 have
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, ii, 20); in California (C. N., 1868, 17, 209); in Canada (Dawson,
Am. 8., 1870, (2), 50, 130); in New Zealand (Mac Ivor, C. N., 1887, 55, 125);
in Russia, Germany, Greenland, etc. 1'ure 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 Fe304 at 400°. Amor-
phous carbon is formed by passing CC14 over Na in a tube heated to red-
ness (Porchcr, C. N., 1881, 44, 203).
4. Preparation. — Pure graphite is prepared by heating the commercial
graphite on a water bath with KC10t and H,S04 and repeatedly washing.
If it contains Si02 it should also be treated with NaF and H.,S04 . 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.* Amorphous carbon is prepared approximately pure by
heating pure cane sugar in a closed platinum crucible; then boiling in
succession with HC1 , KOH , and H.,0 ; 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. Reactions. — N"ot attacked by acids or alkalis. It slowly oxidizes to
C02 when heated with concentrated H..SO^ and K,Cr,07 . Upon gently
warming graphite with KC10. and HNO-, , graphitic acid, CUH400, is
said to be formed (Stingl, B., 1873, 6, 301). 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 C02 . If the carbon and oxygen
have been previously very thoroughly dried the action is very slow,
especially with graphite. By fusion with KNO., or KC10.( carbon is oxid-
ized to C02 . With vapors of sulphur, carbon disulphide is formed; i. e.,
§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, C2H2 , is formed. By
igniting in an atmosphere of carbon dioxide, C02 , the whole of the carbon
becomes carbon monoxide : C -(- C02 = SCO .
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 , Sb , Sn , Pd , Mo , Ru , Os , Rh , Ir , Te , Se , W , K ,
Wa , Kb , Cr , Fe , Mn , Co , Ni , Zn , Ti , Tl .
(a) Pb304 + 20 = 3Pb + 2CO2
(6) 2PbCL + 2Na,C03 + C = 2Pb + 4NaCl + SCO,
(c) CuO + C (excess) = Cu + CO
(d) C + 2CuO (excess) = 2Cu + CO2
With excess of carbon CO is formed (c). With excess of the oxide C02 is
formed (d). In the reduction of iron ore, the process is conducted so as
to give some CO and some C02 . 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 -f- 2C = BaS + 2C02 .
8. Detection. — By its appearance ; failure to react with general reagents ;
and by its combustion to C02 with oxygen (air), or with K2Cr207 and con-
centrated H0SOj (Fritsche, A., 1896, 294, 79), .then by identification with
Ca(OH)2 (§228, 0). %
9. Estimation. — By combustion to CO2 and weighing- after absorption in KOH
solution. See works on ultimate organic analysis.
§223. Acetic acid. HC,H302 = 60.032 .
H 0
H'4(C2)+'"-'"0-".) , H — C — C — 0 — H = CH3C02H.
• I
H
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, ami lias a specific gravity of 1.0481 at 1">°. 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. Formation. - (a) During the decay of many organic compounds, (b)
By gently heating sodium methylate, NaOCH, , in a current of carbon
monoxide: NaOCH3 + CO = GH.,GO»Na (I^aC,H30,). (c) By boiling
methyl cyanide with acids or alkalis: CH.,CN -f HC1 -f 2H20 = HC,H302
-f NH4C1. (d) By the oxidation of alcohol: 30^.0 + 2K,Cr,67 -+-
8H.SO, = 2K,S04 -f 2Cr,(S04), + 3HC,HaO, •+ IIH.,0.
4. Preparation. — (a) By the dry 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 g1 1: ial
acetic acid.
5. Solubilities. — Miscible 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 Fe'", Al , etc., are
insoluble in water. Very many of the acetates are soluble in alcohol.
(>. Reactions. — 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, CO.. and SO, are evolved.
Solution of ferric chloride forms, with solutions of acetates, a red solu-
tion containing ferric acetate, Fe(C2H.,0.,).t . which on boiling precipitates
brownish-red, basic ferric acetate. The red solution is not decolored by
solution of mere-uric chloride (distinction from thiocyanatc); but is de-
colored by strong acidulation with sulphuric acid or hydrochloric acid (dis-
tinction from thiocyanatc and from meconato). The ferric acetate is pre-
cipitated by alkali hydroxides.
If acetic acid or an acetate be warmed with sulphuric acid and a little
alcohol, the characteristic pungent and fragrant odor of ethyl acetate or
acetic ether is obtained:
HC H:,O, + C.HaOH = H2O + C,HAC:H,OS
Acetic acid does not act as a Reducing Agent a< readily as do most of
the organic carbon compounds. It docs iio/Qlitce permanganates even in
boiling solution; reduces auric chloride only in alkaline solution, and docs
not reduce alkaline copper solution. Tt 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 <>f in chnii'. C..H..O . inflammable and of an agreeable odor. By pro-
longed ignition of alkali acetate< in the air. carbonates are obtained free
from charconl. By ignition with alkali hydroxides in dry mixtures,
methane. ni<trdi-<i(ix, CH4 . is evolved. By ignition with alkalis and
arsenous anhydride, the poisonous and offensive vapor of cacodyl oxide
.^224. 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:
4KC2H302 -f As203 = As2(CH3)40 + 2K2C03 -f SCO, .
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 §152).
(d) By ignition of the dry acetate alone to acetone, CH3COCH3 ; with
NaOH to methane, CH4 ; or with As203 to cacodyl oxide, (e) As a delicate
test for formates or acetates it is directed to warm a solution of CuCl2 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. Estimation. — 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. H3CBH307 = 192.064 .
H2C — COoH
|
H'3(C6)+lfr-4H'nO-"7 , H — 0 — C — C02H
H2C — CO,,H
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
citrates 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 dropping1 the solution tested
into the reagent) gives no precipitate with citric acid or citrates in the cold
(distinction from tartaric aAty, but on heating, the white calcium citrate,
Ca3(C0Hr,O7), , is precipitated (not soluble in cold potassium hydroxide solu-
tion). By filtering before boiling, the tart rate 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, Pb3 (C,iH5O7)2 , soluble
in ammonia. Silver nitrate gives a white precipitate of silver citrate,
Ag:jC0Hr,OT , which does not blacken on boiling (distinction from tartrate).
For action of citric acid or citrates in hindering many of the \is\ial analytical
reactions, see Spiller, J. C., 1858, 10, 110.
One part of citric acid dissolved in two pavts 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 isomcrs}.
252 TARTARIC ACID. ^225, 1.
Citric acid does not act very readily as a reducing agent; does not reduce
alkaline copper solution, or silver solution; reduces permanganate very slowly
Concentrated nitric acid produces from it, acetic and oxalic acids; and diges-
tion with manganese dioxide decomposes it, with formation of acetone, acrylic
and acetic acids. Citrates carbonize on ignition, with various empy*eumatic
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,C4H4Ofl = 150.048 .
H 0
I II
H — 0 — C — C - 0 — H CH(OH)C02H
or
H — 0 — C — C — 0 — H CH(OH)C02H
1. Properties. — 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, racemic acid, and mesotartaric acid.
They differ from each other in crystalline form, in solubility, and especially
in the deportment of their solutions towards polari/.ed light. Racemic 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 (Kiliaili, .1., 1880, 205, 175). By action of sodium amalgam on
oxalic ether in alcoholic- solution (Debus, /!., 1873, 166, 124). By synthesis
from succinic acid by formation first of the dibromsuccinic acid, HJC4Br;H3O4;
then substitution of the OH group for the bromine by means of water and
silver oxide.
4. Preparation. — The crude argol deposited during the fermentation of grape
juice is recrystallized, giving the commercial cream of tartar, KHC,H(O .
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 crysl;illi/ation and
purified by recrystallization (Ficinius, Arch. Pharm., 1879, 215,' 14 and 310).
5. Solubilities. — Th(\ Tartrates of the alkali bases arc soluble in water;
the normal tartrates being freely soluble, 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, KSbOC4H400 .
Hydrochloric, nitric, and sulphuric acids transpose the tartrates
(whether forming solutions 01 not). Most of the tartrates are also dis-
§225, 8d. TARTARIC' ACID. 253
solved (and, if abeady 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,
Rochelk salt: Sb , SnIV, Bi , Cu, Fe , Al, Cr, Co, Ni, Mil, and Zn ; Ba,
Sr, Ca, and Mg to quite an extent. CdCO;i is not dissolved (Warron,
. C. N., 1888, 57, 223).
6. Keactions. — Solution of calcium hydroxide, added to alkaline reac-
tion, precipitates from cold solution of tartaric acid, or of soluble tartrates,
calcium tarlrate, white, CaC4H400 . 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 &alts of the following metals : Al , Bi , Co , Ni , Cr , Cu , Fe , Pb , Pt ,
and Zn (Grothe, J. pr., 1864, 92, 175).
Silver nitrate precipitates, from solutions of normal tartrates, silver
tartrate, Ag2C4H4Oa , 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, lOb),
distinction from citric acid. Free tartaric acid does, not reduce silver
salts. Permanganate is reduced quickly by alkaline solution of tartrate-;
(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
burnt sugar, separating carbon, and becoming finally converted to carbon-
ates.— Strong sulphuric acid also blackens tartrates, on warming. Melted
potassium 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. (&) By
the deportment of the calcium salt with cold and hot KOH (6). (c) By the
formation of the silver mirror (§59, 10b). (/T) By its action as an alkali
354 CARBON MONOXIDE. §225, 1).
tartrate in preventing precipitation of the solutions of the heavy metals
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).
§226. Carbon monoxide. CO = 28.0 .
C"0-", C = 0 .
1. Properties. — Carbon monoxide, carbonic oxide, formic anhydride, CO , is a
colorless, tasteless gas. Specific yrarity, 0.9678. By maintaining1 a pressure of
200 to 300 atmospheres at — 136° and then reducing1 the pressure to CO atmos-
pheres the g'as becomes a colorless transparent liquid (Wroblevvski and Ols-
zewski, A. Oh., 1884 (6), 1, 128). It is, when inhaled, a virulent poison, abstract-
ing oxygen from the blood and combining with the haemoglobin. It burns in
the air with a pale blue flame to CO. , 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. Occurrence. — In combination as formic acid in ants and in nettles.
3. Formation. — (a) By the incomplete combustion of coal, charcoal or
organic material. (6) From the reduction of metallic oxides in the blast
furnace with excess of charcoal: Fe20:i -f 3C — 2Fe -j- 3CO . (c) By
heating sodium sulphate with excess of charcoal (LeBlanc's soda process):
Na2S04 -f 4C = Na2S -f 4CO . See also Grimm and Kamdohr (A., 185(5,
98, 127).
4. Preparation. — (a) By passing steam over charcoal at a white heat
(water gas): H,0 + C == CO -f H2 (Naumann and Pistor, R., 1885, 18,
164). (b) By passing C02 over red hot charcoal, (c) By heating
K4Fe(CN)(. with concentrated H,S04: K4Fe(CN)(i + 6H2S04 -f GH20 =
2KJ304 -f 3(NH4)2S04 + FeS04 -f GCO . With dilute acid HCN is formed.
(d) By heating a formate with concentrated sulphuric acid: 2KCH02 -j-
H^SO, == K,S04 + 2CO -f 2H,0 . («) By heating an oxalate with con-
centrated sulphuric acid: K.,C.,04 -f 2H0S04 = K,S04 + H,S04.H,,0 +
CO + C02
5. Solubilities. — It is not absorbed by KOH or Ca(OH), (distinction
from CO.,). It is absorbed bv charcoal, cuprous chloride, and by several
metals, e. (j., K , Ag , and An .
6. Reactions. — It is an energetic reducing agent. Combines with moist
fixed alkalis to form a formate (Froelich and Geuther, A., 1880, 202, 317).
In the sunlight it combines directly with chlorine or bromine. It is
oxidized to C02 by warming with K,Cr,07 and concentrated HoS04 ; also
§227,46. OXALIC ACID. 255
by palladium sponge saturated with hydrogen, and in presence of oxygen
and water (Eemsen and Keiser, E., 1884, 17, 83). A solution of PdCl2 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 C02 by its failure to be absorbed by
KOH or Ca(OH)2 . By its combustion to C02 and detection as such. By
its combination with hot concentrated KOH to form a formate. It is
detected in the blood by the absorption spectrum (Vogel, B., 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, C. N., 1878, 37, 6).
§227. Oxalic acid. H2C204 = 90.016.
0 0
I! II C02H
H'2(C2)+60^'4 ,H — 0 — C — C — 0 — H or |
C02H
1. Properties. — Absolute oxalic acid, HoC2O4 , is a white, amorphous solid,
\vhich may be sublimed at 150° with only partial decomposition: H^C..^ ==
CO2 + CO + H.,O . Crystallized oxalic acid, H,C.,O4,2H2O , effloresces very
slowly in warm, dry air, and melts in its water of crystallization at 98°; at
which temperature the liquid soon evaporates to the absolute acid. Oxalic
anhydride is not formed.
2. Occurrence. — Found in many plants in a free state or as an oxalate. In.
sorrel it is found as KHC2O4; in rhubarb as CaC,O4 , 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, (ft) By the oxidation of
glycol with nitric acid, (c) By heating potassium formate above 400°
(Merz and Weith, B., 1882, 15, 1507). (d) By passing C02 over a mixture
of sodium and sand at 360° (Drechsel, Bl, 1868, 10, 121).
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 C02 . (b) By heating sawdust with KOH or NaOH . Hydrogen is
evolved, the cellulose, CGH100. , ^eing converted into oxalic acid. Under
certain conditions, additional products are formed. It is also formed in
the oxidation of a great many organic compounds.
256 OXALIC AVID. §227,4a,
C^H^Ou + 12HNO, = 6HAO, + 12NO + HH2O
3H.C,O4 + 2HNOS = 6CO2 + 2NO + 4H2O
C0H100, + fiKOH + H,O = 3K2C2O4 + 9H2
Oxalates are formed: a. — By treating the oxide, hydroxide, or car-
bonate with oxalic acid. In this manner may be made the oxalates of
Pb, Ag, Hg', Hg", Sn", Bi, Cu", Cd, Zn , Al , Co , Ni, Mn , Fe", Fe'",
Cr"', Ba , Sr , Ca , Mg , Na , and K . And some others.
b. — By adding oxalic acid to some soluble salt of the metal. In this
manner the above oxalates may be made, except alkali, magnesium,
chromic, ferric, aluminum and stannic oxalates, which arc not precipitated.
Antimonous salts are precipitated, but the precipitate is basic.
c, — Alkali oxalates will precipitate the same solutions as oxalic acid,
but many of the precipitates are soluble in excess of the alkali oxalate,
and, as a rule, the salt formed is a double one, e. y., AgNH4C.,04 . Ba , Ca
and Sr are well-defined exceptions to this rule — their precipitates, formed
by this method, being normal oxalates.
d.— Some of the metals when finally divided are attacked by oxalic acid,
hydrogen being evolved.
5. Solubilities. — Oxalic acid is very soluble in water and in alcohol.
Alkali oxalates are freely soluble in water, as is also chromic oxalate.
Nearly all other metallic oxalates are insoluble in water or only sparingly
soluble (Luckow, J. C., 1887, 52, 529).
The metallic oxalates, soluble and insoluble, are transposed by dilute
sulphuric, hydrochloric, and nitric acids, with formation of oxalic acid:
CaC.,04 + 2HC1 = CaCl, + H,C204 . That is : the precipitated oxalates
of those metals, which form soluble chlorides, dissolve in dilute hydro-
chloric acid; of those metals which form soluble sulphates, in dilute sul-
phuric acid; and all precipitated oxalates dissolve in dilute nitric acid
Acetic acid does not dissolve precipitated oxalates, or but slightly.
Certain of the oxalates dissolve, to some extent, in oxalic acid (as acid
oxalates).
G. Reactions. — A. — With metals and their compounds. — Oxalic acid and
soluble oxalates precipitate solutions of many of the metallic salt*. With
excess of the alkali oxalates soluble double oxalates of the heavy metals
are frequently formed (4). An excess of alkali oxalate transposes par-
tially the alkaline earth carbonates. On the other hand, the alkali car-
bonates in excess partially transpose the alkaline earth oxalate:* (Smith,
J. C., 1877, 32, 245). See also under 6& of the respective metals.
Oxalic acid is a decided reducing agent, being converted to irater and
•carbonic anhydride (a), and the metallic oxalates to Carbon&tM and carbonic
anhydride (b), by all strong oxidizing agents.
£227, 115. OXALIC ACID. 257
(a) 2H,C1;O1 + O, = 2H,O + 4C02
(&) 2K2C2O4 + 0, = 2K,C03 + 2CO,
1. — Pb02 with oxalic acid forms lead oxalate and C02 . Oxalic acid has
no action upon Pb304 , but reduces it quickly in presence of any acid
capable of changing the Pb304 to PbO., .
2. Oxalic acid or ammonium oxalate boiled in the sunlight with HgCl2
gives HgCl and C02 [Gmelin's Hand-look, 9, 118].
5. — HjtAs04 becomes H;!As08 , and C02 is evolved. To prove that Asv
becomes As'", add excess of potassium hydroxide, and then potassium per-
manganate. The latter will be quickly decolored.
4. — Bi205 becomes bismuth oxalate and C02 .
5. — Mn"+n becomes Mn". (That i«, all compounds of manganese having
more than two bonds are reduced to the dyad.) In absence of other free
acid, MnC204 is formed, and C02 is given off. If some non-reducing acid
be present, such as H2S04 , it unites with the manganese, and all of the
oxalic acid is converted into C02 .
6. — Co203 and Co(OH)3 form cobaltous oxalate, and C02 is evolved.
7. — Ni203 and Ni(OH)3 become nickelous oxalate, and C02 is evolved.
8. — H2Cr04 is reduced to chromic oxalate, and C02 is evolved.
As a rule, reducing agents have no action on oxalic acid at ordinary
temperatures. By fusion, however, a few metals, K , Na , Mg , etc., reduce
it to free carbon.
B. — With non-metals and their compounds.
1.— HCN , HCNS , H4Fe(CN)G , and H;!Fe(CN)6 seem to be without action
upon oxalic acid.
2. — HN02 seems to have no action upon H2C204 . With HN03, C02 ,
NO , and H20 are formed. The nitric acid should be concentrated. Test
for the C02 by passing the gases into a solution of BaCL containing KOH .
' 3. — H3P02 , HSF03 , and H3P04 do not act upon oxalic acid.
4. — Concentrated sulphuric acid, with a gentle heat, decomposes oxalic
acid, by removing the elements of water from it, with effervescence of
carbon dioxide and carbon monoxide: H2C204 -f- H2S04 = H.,S04.H20 -|-
C02 -(- CO . With oxalates, the decomposition generates the same gases.
Other strong dehydrating agents produce the same result.
The effervescing gases, C02 and CO , give the reactions for carbonic anhy-
dride; also, if in a sufficient quantity, the CO will burn with a blue flame,
when ignited.
5. — With chlorine, hydrochloric acid is formed and the oxalic acid
becomes C02 (Gmelin's Hand-book, 9, 116). This reaction takes place
more readily in the presence of KOH , forming KC1 and K2CO:! . HC10
forms C02 and Cl . If the oxalic be in excess HC1 is formed. The action
258 OXALH- ACID. ^227, lln.
is more rapid in the presence of a fixed alkali, an alkali chloride and
carbonate heing formed. HC10, forms C02 and varying proportions of
Cl and HC1 . A high degree of heat and excess of oxalic acid favoring
the production of HC1 (Calvert and Davies, J. C., 1850, 2, 193).
6. — Bromine decomposes oxalic acid in alkaline mixture, forming 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 CO.. are formed ; with excess of oxalic acid and heat hydro-
bromic acid is formed.
7. — HIO forms C02 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 (6). Th<;
oxalates of metals, whose oxides are decomposed by heat, leave the metal,
and give off carbonic anhydride (r). 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, = CaCO:, + CO
(h) ZnC20, = ZnO + CO, + CO
(<•) Agr.ao, = 2Ag + -'CO,
8, Detection. — (a) By warming with concentrated sulphuric acid after
decomposition of carbonates with dilute sulphuric acid; showing the pres-
ence of CO., by absorption in Ca(OH), or in a solution of BaCL alkaline
with KOH ; and showing the presence of CO by its combustibility. (&) 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 HC1 solution containing BaCL and CaCL he adds
carefully a solution of NaC.,H.,0.. and watches the zone of contact.
0. Estimation. — (a) It is precipitated as CaC,04; after washing, the Ca is
determined by §188, 9. from which the oxalic acid is calculated. (/>) By the
amount of KMnO, which it will reduce, (c) 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 AuCl, .
£228, 4. CARBON DIOXIDE. 2«VJ
§228. Carbon dioxide. C02 = 44.0 .
, (Carbonic anhydride.)
Carbonic acid (hypothetical). H2C03 = 62.016 .
0
CIV0-"2 and H'2CIV0-"3 ,0 = C = OandH — 0 — C — 0 — H.
1. Properties. — The specific y rarity of the gas CO2 is 1.52897 (Crafts, C. r., 1888,
106, 162); of the liquid at —34°, 1.057 (Cailletet and Mathias, C. r., 188G, 102,
1202); of the solid (hammered), slightly under 1.2 (Landolt, B., 1884, 17, :^09).
Critical temperature, 30.92° (Andrews, Trans. Roy. 8oc., 1809, 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 CO. is more compres-
sible than other liquids (Natterer, ./., 1851, 59). It diffuses through porous
plates more rapidly than oxygen (Graham, C. N., 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 CO;, 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. pr., 1834, 3, 109). Solid CO2 at 767.3
mm. barometric pressure melts at — 77.94° (Regnault, A. Cft., 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
great 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. (&) By
burning CO . (c) By the reduction of many metallic oxides upon ignition
with charcoal, (d) During fermentation or decay of organic material.
(e) By the reaction between acids and carbonates.
Liquid C02 is made by compressing the gas with pumps at a reduced
temperature.
Solid C02 is made by allowing the liquid to escape freely into woolen
bags and then compressing in wooden moulds (Landolt, I. c.).
4. Preparation. — CaC03 (chalk or marble) in small lumps is treated with
hydrochloric acid in a Kipp's gas generating apparatus. The gas is passed
through a solution of NaHC03 to remove any HC1 that may be carried
over, and then dried by passing through a tube filled Vith fused CaCl2 .
It .is also prepared on a large scale for making the liquid C02 , and for
use. in sugar factories by the ignition of limestone: CaCO, = CaO -4- C02 .
Preparation of Carbonates. — Na2C03 is made by converting NaCl into
Na2S04 , by treating it with H2S04 ; then by long ignition with coal and
calcium carbonate, impure sodium carbonate is formed (Leblanc's process).
Na2SO< + 4C.+ CaCO; = CaS + 4CO + Na2CO3
260 CARBON DIOXIDE. §228, 5.
It is separated by lixiviation with water, and farther purified. The
other method, known as the ammonia, or Solvay's process, consists in pass-
ing NH3 and C02 into a concentrated solution of NaCl (rf). The NaHCO.,
is converted into Na2C03 by heat, and the evolved C02 used over again (6).
The NH4C1 is warmed with MgO (c), and the NH:! which is given off is
used over again. The IlgCL, is strongly heated (d) and the MgO is used
over again, and the evolved gas sold as hydrochloric acid. This continu-
ous process has nearly superseded the Leblanc process.
(a) NaCl + NH3 + H,O + CO2 = NaHCO, + NH4C1
(6) 2NaHCO, + heat = Na,COa + CO, + H2O
(c) 2NH4C1 + MgO = MgCL + 2NH, + H2O
(d) MgCl2 -|- H2O + heat = MgO + 2HC1
The other carbonates are mostly made from the sodium salt (6).
5. Solubilities. — C02 is soluble in water, forming the hypothetical
HoG03 , 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 -f CO, = KHC03 or 2KOH + C02 =
K2C03 + H,0 . The carbonates of the alkalis are soluble in water (acid
alkali carbonates are less soluble than the normal carbonates), other
carbonates are insoluble in water or only sparingly soluble. The prest-iu r
of some other salts, especially ammonium salts, increases the solubility oi'
carbonates, notably magnesium carbonate (§189, 5c). Many of the car-
bonates are soluble in water saturated with C02 ; forming acid carbonates
of variable composition. Boiling removes the excess of C02 , causing pre-
cipitation of the carbonate.
G. Reactions. — Dry carbon dioxide does not unite with dry calcium
oxide at ordinary temperature (Birnbaum and Maher, B., 1879, 12, 1547;
Scheibler, B., 188G, 19, 1973). Also at 0° no reaction takes place between
dry C02 and dry Na,0 , but at 400° combination takes place with incan-
descence (Beketoff, BL, 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-
mal 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 HC1 , HC10;,HBr , H3r03 ,
HI , HI03 , H2C204 , HN03 , H3P04 , H2S04 , and even by H2S , completely
from carbonates of the first four groups, incompletely from those of the
fifth and sixth groups (Nandin and Montholon, C. r., 1876, 33, 58).
Ammonium carbonate precipitates solutions of all the non-alkali metals,
$228, li. CARliOX DIOXIDE. 2G1
chiefly as carbonates; except magnesium salts which are not at all pre-
cipitated, a soluble double salt being at once formed (separation of barium,
Btrontium, 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 C02 which reddens moist litmus paper: Na.,C03 -{-
H2S04 = Na.SO, + H20 -f 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 : 2Na2CO:! + H,S04 = = Na2S04 -f 2NaHCO:! ; 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-
bonate solution is added drop by drop to the acid, so that the latter is con-
stantly in excess, even slight traces of carbonate give notable effervescence.
The effervescence of carbonic acid gas, C02 , is distinguished from that of
H2S or S02 by the gas being odorless, from that of N^O, by its being color-
less and odorless : from all others by the effervescence being proportionally
more forcible. It should be remembered, however, that C02 is evolved
(with CO) on adding strong sulphuric acid to oxalates or to cyanates.
On passing the gas, C02 , 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 b}r 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(OH) , = CaC03 + H,,O
(b) CO, + Ba(OH), = BaCO3 + H.O
(c) CO, + CaCL + 2NH4OH = CaCO3 + 2NH4C1 + H2O
(d) C02 + Pb,O(C2H30,)2 = PbC03 + Pb(C,H30,),
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
ino.\nn:.
it cannot be preserved in bottles partly full and frequently opened, and
cannot be diluted clear, unless with recently boiled water.
Solutions of tlie acid carbonates effervesce, with escape of CO., , on boiling
or heating, thus :
I2KHC03 = K,COS + H,O + CO, . (Gradually, at 100°.)
2NaHC03 = Na,CO., + H,0 -f CO, . (Gnulually, at 70°; rapidly at 90° to 100°.)
2NH4HCO3 = (NHJ.CO, + H,0 + CO, . ( llegins to evolve CO2 at :^G°.)
(NH4)4H2(CO,)a = :.>(NH,),CO, + H,0 + CO, . (Begins at 49°.)
7. Ignition. On ignition, the normal carbonates of the metals 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 CO., . At a lower temperature, ignition changes
all other carbonates to the oxide and CO., , except that the carbonates of
silver at 250°, mercury, and some of the rarer metals are reduced to the
metallic state, C02 and oxygen being evolved. Stannous and ferrous
oxides ignited in an atmosphere of CO., are changed to SnO.., and Fe20,. ,
respectively, with evolution of CO (Wagner, Z., 1879, 18, 559).
8. Detection. — Carbonates are detected: (a) By the sudden effervescence
when treated with dilute acids, (b) By the precipitate which this ga«
forms with solutions of Ca(OH), , Ba(OH), , or Pb,0(C,H.O,), . If but a
small amount of carbonate be present, the mixture must be warmed to
drive the CO., over into the reagent (G). A non-volatile acid as H.,S04 or
H3P04 should be used, as a volatile acid might pass over with the C02 and
prevent the formation of a precipitate, (c) Phenolphthalein 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. Pliarm., 1892, (5), 25, 448).
To detect free carbonic acid in presence 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 cc. of this rosolic acid solution be added to about 50 cc.
of the water to be tested — spring water, for instance — the liquid will bo
colorless, or at most faintly yellowish if it contains free carbonic acid,
whereas, if there be no free carbonic acid, but only double salts, it will
be red (Pettenkofer, Dingl, 1875, 217, 1 58).
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 (£207,6^) docs not take place in presence of free car-
bonic 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 brown precipitate is obtained (Dobbin, J. Soc. Ind.t
1888, 7, 829)!
§230, 1. CYANOGEN— HVDKOCYAMC ACID. 263
9. Estimation. — («) By decomposition of a weighed sample with acids and
determining the CO. by loss of weight, after taking into consideration the
gain in weight due to the acid used, (h) I5y decomposition of the weighed
sample and collection of the CO. in a weighed KOH solution. (e) By decom-
position with an excess of a standard acid, boiling to expel the CO2 and
titrating the excess of acid, (d) Sodium bicarbonate may be estimated by
titration with sodium hydroxide: NaHC03 + NaOH = Na,CO3 + H2O . The
first excess of sodium hydroxide beyond the reaction gives a brown precipi-
tate with silver nitrate (Lunge, Z. anyew., 1897, 109; Bohlig, Arch. Pharm., 1888,
326, 541).
§229. Cyanogen. CN = 26.04 .
N=C — C =N.
A colorless, intensely poisonous gas; tipmifie {/rarity, 1.8064 (Gay-Lussac, Gilb.,
1816, 53, 145). The molecular weight shows the molecule to be C2N2 . At
•rdinary atmospheric pressure it liquifies at — 22° (Drion, J., 1860, 41); at 20°
under four atmospheres pressure (Hofmann, li., 1870, 3, 658). The gas has
an odor of bitter almonds and burns with a red color to the flame forming
CO2 and N . When cooled to about the free/ing point of mercury it solidifies
to a crystalline ice-like mass (Hofmann, I.e.). drlt'wal temperature, 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,
J. pr., 1847, 42, 145). It is prepared: («) By heating the cyanides of mercury,
silver or gold: Hg(CN), = Hg + C,N, . ' (b) By the dry distillation of am-
monium oxalate: (NH4).,C..04 = 4H.,0 + C..N.,'. (c) By fusing KCN with
HgCl2: 2KCN + HgCL — Hg + 2KC1 + C2N2 . (rf) By heating a solution of
CuSO4 with KCN . Half of the CN is evolved and CuCN is formed. If the
CuCN be heated with FeCl3 or MnO, and HC2H3O,, , 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. CJi., 1886, (6), 6,
140). It combines with Cl , Br , I, S, P, and with many of the metals,
reacting very much like the halogens. It dissolves in water, alcohol and
ether; but gradually decomposes with formation of ammonium oxalate and
carbonate (Vauquel'in, A. Ch., 1823, 22, 132; Buff and Hofmann, A., 1860, 1-I3,
129). At 500° it combines with hydrogen to form HCN (Berthelot, Bl., 1880,
(2), 33, 2). With Zn it forms Zn('CN)J , rapidly at 100°. With HC1 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 KCNO are formed: C.N, + 2KOH = KCN + KCNO -J- H2O . 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. »S"/,m'/7r (jrai-itu at 19°, 0.697
(Bleekrode, Proc. Roy. Soc., 1884, 37, 339). It burns with a bluish-red flame,
forming H2O , CO. 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
sohition 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 CO2 . The free acid or soluble salts when warmed
HTDKOCYAMC A('ll). ij230, 2.
with dilute alkalis or acids (with strong acids in the cold) becomes formic
acid and ammonia: HCN -f 2S..O = HCO,H + NHS .
2. Occurrence. — The free acid dots 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 routs of the
last-named tree, when soaked in water for a time and then distilled, yield
hydrocyanic acid, together with bitter-almond oil. Potassium cyanide appears
in the deposits of blast furnaces for the smelting of iron ores.
;f. Formation. — (a) Decomposition of amygdaline by emulsine and distilla-
tion, (b) By the action of the electric spark on a mixture of acetylene :m<l
nitrogen (Berthelot, •/., 1ST4, li:i). (c) T?y heating a mixture of cyanogen and
hydrogen (§229). («/) By the dry distillation of ammonium formate: NH4CHO,
= HCN + 2H;.O . (<) By boiling or fusing many organic compounds contain-
ing nitrogen with KOH , forming KCN (Post and Hnebner, B., 1872, 5, 408).
(f) By decomposition of metallic cyanides with mineral acids. (0) By heating
chloroform with a mixture of ammonium and potassium hydroxides (llof-
mann. .1.. 1S(17, 144, 11(1).
4. Preparation. — (a) By the action of dilute sulphuric acid on potassium
ferrocyanide: -jK.FelCN),, + ,iH,S04 = (iHCN + KsFe.(CN)d -f :!K,SO, .
(b) By action of acids upon metallic cyanides. (<•) By the action of sulphuric
:u.-id upon mercuric cyanide in the presence of metallic iron: Hg(CN) . + Fe -4-
H2S04 = 2HCN + FeS04 + Hg .
Metallic cyanides are prepared: («) By the action of HCN on metallic
hydroxides. (/>) By the action of soluble cyanides on metallic salts, (c) By
igniting potassium ferrocyanide: K,Fe(CN),i = 4KCN + FeC, + N2 . (d) By
heating potassium ferrocyanide with potassium carbonate. If prepared in
this manner it contains some cyanate: K.Fe(CN),, -f K CO, = 5KCN + KCNO
+ Fe + CO, .
r>. 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 22.5° ; it also increases slightly in volume (Bussy and
Huignet, A. f'A., 18f>5, (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 tvetalUe ci/anidex.
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 -f 2KCN + 2H,O = I-'CuCN + :.'KOH -f- H. . Iron or steel
wire are not attacked ((Joyder, (7. N., 1K94, 69, 2(52, 2«8 and 280).
G. 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 exeess of the precipitant: HgCL + 2KCN = Hg(CN)., + 2KC1; and with
excess of KCN: Hg(CN), -f 2KCN = (KCN),Hg(CN). .
Class I. Double cyanides irhich arc not a ft eel rd. li)i dlkiili lii/<Iro.n<1ex, Init xnffrr
dissociation ichen treated iritli dilute acid*: (KCN),Hg(CN);, + 2HC1 = Hg(CN)2
+ 2KC1 + 2HCN . These closely resemble the double iodides (potassium
mercuric), and the double sulphides or thiosalts (§69, r>r and f>e). The most
frequently occurring of the double cyanides of this class, which dissolve in
water, are given below:
Potassium (or sodium) zinc cyanide, K,Zn(CN)4 or (KCN)2Zn(CN)2 .
Potassium (or sodium) nickel cyanide, K;Ni(CN)4 or (KCN),Ni(CN)2 .
Potassium (or sodium) copper cyanide, KJCu(CN)4 or (KCN);,Cu(CN)2 .
Potassium cadmium cyanide, K=Cd(CN)4 or (KCN),Cd(CN);, .
Potassium (sodium or ammonium) silver cyanide, KCNAgCN or KAg(CN), .
Potassium (or sodium) mercuric cyanide, K,Hg(CN)4 or (KCN),.Hg(CN), .
Potassium (or sodium) auric cyanide. KAu(CN)4 or KCNAu(CN), .
$230, 6. HYDROCYANIC ACID. 2(j~j
Class II. Double cyanides irliich, <ix precipitates, are transposed by alkali Jiydro.r-
idcs, in dilute solution (a), and arc transposed, icithout dissociation, by dilute acids
(b). In these double cyanides, as potassium ferrous cyanide, K4Fe(CN),, , 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)« as " ferrocyanogen,"
giving K,Fe(CN)G 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 cyanides of this class, which are soluble in
water, are given below.
CuJFe(CN), + 4KOH = 2Cu(OH)2 + K4Fe(CN)6
K4Fe(CN)0 + 2H,S04 = 2K,S04 + H4Fe(CN)(i
2K3Fe(CN)0 + ::H,S04 = :;K,S04 + 2H,Fe(CN).
Alkali ferrocyanides, as K4Fe"(CN)0 , potassium ferrous cyanide.
Ferricyanides, as K3Fe"'(CN),i , potassium ferric cyanide.
Cobalticyanides, as K3Co'"(CN)0 , potassium cobaltic cyanide.
Manganicyanides, as K3Mn/"(CN)0 , potassium manganic cyanide.
Chromicyanides, as K3(Cr'") (CN)c; , 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 HJ3 , 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
cyanide, AgCN , white, insoluble in dilute nitric acid, soluble in ammonium
hydroxide, in hot ammonium carbonate, in potassium cyanide, and in
thiosulphates — uniform with silver chloride. Cold strong hydrochloric
acid 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-
cipitable as chloride (distinction and means of separation from chloride).
Solution of mercurous nitrate, with cyanides or hydrocyanic acid, is
resolved into metallic mercury, as a gray precipitate, and mercuric cyanide
and nitrate, in solution. Salts of copper react, as stated in §77, 6&; salts
of lead, as stated in £57, Oft'.
Ferrous salts, added to saturation, precipitate from solutions of cyan-
ides, not from hydrocyanic acid, ferrous cyanide, Fe(CN)L> , while, if free
from the ferric hydroxide formed by admixture of ferric salt, and, with
the same condition, soluble in excess of the cyanide, as (with potassium
cyanide), (KCN)4Fe(CN),, = K4Fe(CN)0 , potassium fenwyanith (a). On
?66 /n'l)]tO('Y.\\H' .ir//>. ^230, 7.
acidulating this solution, it gives the blue precipitates with iron salts,
more marked with ferric salts (?;):
(a) 2KCN + PeSO, = Fe(CN), + K,SO4
Fe(CN), + 4KCN = K,Fe(CN),,
(ft) liK.FetCN),, + 4FeCl3 = Fe,(Fe(CN),;):i -f- 12KC1
This production of the blue ferric ferrocyanide in made a delicate test foi
hydrocyanic acid, 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 appeal-, if hydrocyanic acid was present (Link and Moeckel, Z., 1878,
17, 456).
Solution of nitrophenic acid, picric acid, C,jH,(N02);,OH , 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 (aw
KCSH4N-06). This test is very delicate, but not very distinctive, as var-
ious reducing agents give red products with nitrophenic acid (Vogel,
T. N., 1884, 50, 270).
The fixed alkali hydroxides, in boiling solution, strongly alkaline, gradu-
ally decompose the cyanides with production of ammonia and formate:
HCN -f KOH 4 H,0 = KCHO, + NH, . Ferrocyanides and fcrricyanides
finally yield the same products.- Dilute alkalis, not heated, transpose, as
by equation a, class Tl 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 oxvgen with formation of
a cyanate: 2KCN -f- 0.. = VlKCNO . 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 (8).
Hydrocyanic acid reduces PbO., , forming Pb(CN)., and CN : PbO, -f 4HCN
= Pb(CN), + C,N, + 2H,o"(Liebig, .-1., 1838," 25, 3). With HOT and
H,0, oxamide is formed (Altfield, J. C., 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, /?., 1887, 20, TIT, 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 cya nates or thiocyanates formed in
the reaction are not readily decomposed by heat alone.
£231. iiYi)ifori-:i(ii<>('YAMC ACID. 267
8. Detection. — Cyanides may be deteeted: (a) By the odor of the free
acid upon decomposition of the cyanide with acids. This test must be
applied with extreme caution as the evolved HCN or CN is very poisonous.
(b) By formation 'of a ferrocyanide and its reaction with ferric salts, as
described in (5. (r) The production of the red ferric thiocyanate is a test
for hydrocyanic acid, more delicate than formation of ferrocyanide. By
warm digestion this reaction occurs: 2KCN -(- S., = 2KCNS ; or:
• 2(NH,),S4 + 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 H2S), 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, /. 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 (>0°. 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.KCN is
formed. As soon as the alkali cyanide is all used in the formation of the
double cyanide, the next molecule of silver nitrate decomposes a molecule of
tJie 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). (It) By titration with a standard solution of HgCL , applicable
in presence of cy a nates and thiocyanates (Ilannay, J. C., 187S. 33, 245).
§231. Hydroferrocyanic acid. H4Fe(CN),; — 216.172 .
H',Fe"(CN)-'(! .
Absolute hydroferrocyanic acid (§230. C>, Class IT.), 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: 7H4Fe(CN)(. + O.. = Fe4 (FeCCN),,);,
4- 2H,0 -f 24 HCN .
Potassium ferrocyanide' is the usual starting point in the preparation of the
free acid or any of the salts. Tt is prepared by fusing together in an iron
268 HYDROFEKROCYAXIC ACID. 5<231.
kettle nitrogenous animal matter (blood, hair, horn, hoof, etc.), commercial
potash (KOH). and scrap iron. The ferrocyanide is formed when this mass is
digested with water. The filtrate is evaporated to crystallization (lemon-yellow
prism), soluble in four parts of water.
Hydro ferrocyanic acid is formed by transposition of metallic ferrocyanides
in solution, with strong acids («). When the solution is heated, hydrocyanic
acid is evolved; in the case of an alkali ferrocyanide, without absorption of
oxygen (6). Potassium ferrocyanide and sulpluiric acid are usually employed
for preparation of hydrocyanic acid (c):
(a) K.FeCCN), + 2H2SO4 = 2K,SO4 + H4Fe(CN)0
(6) 3H4Fe(CN)0 + K4Fe(CN)0 = 2K,FeFe(CN)0 + 12HCN •
(c) 2K4Fe(CN)a + 3H,SO4 = 3K,SO4 + K,FeFe(CN)» + GHCN
The ferrocyanides 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 ferntcunnitlcs; soluble and insoluble; that
of barium and potassium is soluble, but potHUxiinii calcium ferrocyanide is in-
soluble. The most of the ferrocyanides of a heavy metal and an alkali metal
are insoluble. Potassium and sodium ferrocj-anides 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 ferrocyanides have marked and very diverse
colors, as seen below.
Solutions of alkali ferrocyanides, as K4Fe(CN),: , give, with soluble salts of:
Aluminum, a white precipitate, A1(OH)3 and Fe(CN);, (formed slowly).
Antimony a white " Sb4[Fe(CN),;]3.25H,,O .
Bismuth, a white " Bi4(Fe(CN),,)s .
Cadmium, a white " Cd2Fe(CN),, (soluble in HC1).
Calcium, a white " K,CaFe(CN),; .
Chromium, no "
Cobalt, a green, then gray " Co,Fe(CN)a .
Copper, a red-brown " Cu..Fe(CN)a •
Gold, no "
Iron (Fe"), white, then blue " K,FeFe(CN)0 .
Iron (Fe'"), a deep blue " Fe,(Fe(CN)(1)a .
Lead, a white " Pb,Fe(CN)a .
Magnesium, a white " (NH,),MgFe(CN)0 (in presence of NH4OH)
a yellow-white " K.MgFe(CN), (only in concentrated solu-
tion).
Manganese, a white " Mn,Fe(CN),l (soluble in HC1).
Mercury (Hg7), a white " Hg4Fe(CN),, (gelatinous).
Mercury (Hg"). a white " Hg,Fe(CN)0 , turning to Hg(CN), and
Fe3(Fe(CN)J,. blue.
Molybdenum, a brown "
Nickel, a greenish-white " Ni,Fe(CN)0 .
Silver, a white " Ag4Fe(CN)0 , (slowly turning blue).
Tin (Sn" and Sniv), white " (gelatinous).
TTranium (uranous), brown " UFe(CN),., .
TJranium (uranyl), red-brown " (UO,);,Fe(CN)rt .
Zinc, a white, gelatinous " Zn,Fe(CN),, .
See Wyrouboff (A. CJi., 1876 (5), 8, 444; and 1877, (5). 10, 409).
Insoluble ferrocyanides are transposed by alkalis (§230, 0. Class II.)
It will be observed (§230, 6) that fn-roci/iiiihli * arc frrroitx combinations, whl
fcrr'u-iintiiili'M nrr {''>'>'"• combinations. And. although ferrocyanides arc far l^ss
easily oxidized than simple ferroiis salts, being stable in the air. they art
§232. HYltJtOFKRRICYANIC ACID. 369
nevertheless reducing agents, of moderate power: 2K4Fe(CN)0 + Cl., =
2K8Fe(CN)ci + 2KC1 .
PbO2 with sulphuric- acid forms Pb" and H3Fe(CN)0 .
Ag' with fixed alkali forms an alkali ferricyanide and metallic silver.
Crvi with phosphoric acid, gives Cr'" and H.,Fe(CN)(1 (Schonbein, J. pr., 1840.
20, 145).
Co'" with phosphoric acid forms Co" and H3Fe(CN)(i .
Ni'" with acetic acid gives Ni" and H3Fe(CN),, .
MnO with phosphoric acid gives Mn" and H3Fe(CN)ti .
Mnvri forms with potassium hydroxide Mn<X and potassium ferricyanide.
With sulpluiric acid, manganous sulphate and hydroferricyanic acid.
Ferricyanides when boiled with NH4OH give ferrocyanitles (Playfair, J. C'..
1857,9, 128).
HNO., forms first hydroferricyanic acid, then hydronitroferricyanic acid and
NO .
HNO3 forms hydroferricyanic acid, and then hydronitroferricyanic acid, NO
being evolved.
Cl forms first hydroferricyanic and hydrochloric acids. Excess of chlorine to
be avoided in preparation of ferricyanides.
HC1O3 forms hydroferricyanic and hydrochloric acids.
Br forms hydroferricyanic and hydrobromic acids.
HBrO 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 (Gmelin'a 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 6ft, §126 and §77). Separatee!
from ferricyanide, by insolubility of alkali salt in alcohol.
Ferrocyanides are estimated in solution with sulphuric acid by titrating with
standard KMnO4 . Also by precipitation with CuSO4 either for gravimetric de-
termination or volumetrically, using a ferric salt as an external indicator.
§232. Hydroferricyanic acid. H,Fe(CN)t; = 215.164 .
H'3Fe"'(CN)-'t; .
Absolute hydroferricyanic acid. H:iFe(CN),, . is a non-volatile, crystallizablt-
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 hydru-
ferricyanic acid radical is broken up.
Potassium ferricyanide is the usual starting point in the preparatiou of most
ferricyanides. It is prepared by passing chlorine into a cold solution of
K4Fe(CN)0 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 K.FelCN), ).
The free acid is made by adding to a cold saturated solution of K:..Fe(CN)
three volumes of concentrated HC1 and drying the precipitate which forms,
in a vacuum (Joannis, f. r., 1SS2. 94, 449. 541 and 725). Lustrous, brownish-
green needles, very soluble in v\a1er and alcohol, insoluble in ether.
The ferricyanides of the metals of the alkalis and alkaline earths arc 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, <>).
270 H7DROFERRICYAXIC ACID. JJ232.
Solutions of metallic ferricyanides give, with soluble salts of:
Aluminum, no precipitate.
Antimony, no precipitate.
Bismuth, light-brown precipitate, BiFe(CN)« , insoluble in HC1 .
Cadmium, yellow precipitate. Cd3[Fe(CN)(1]..j , soluble in acids and in a mum-
mum hydroxide.
Chromium, no precipitate.
Cobalt, brown-red precipitate, Co3[Fe(CN)0]2 , 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.
Copper, a yellow-green precipitate, Cua[Fe(CN),,]« , insoluble in HC1 .
Gold, no precipitate.
Iron (ferrous), dark blue itrwiitifutc, Fe3[Fe(CN),,]2 , insoluble in acids.
Iron (ferric), no prrciiritute, a darkening of the liquid.
Lead, no precipitate, except in concentrated solutions (dark brown).
Manganese, brown precipitate, Mn., [Fe(CN),,], , insoluble in acids.
Mercury (mercurous), red-brown precipitate, turning white on standing.
Mercury (mercuric), no precipitate.
Nickel, yellow-green precipitate. Ni:,[Fe(CN)H]- . insoluble in hydrochloric acid.
With ammonium chloride and hydroxide, excess of ferricyanide gives a
copper-red precipitate.
Silver, a red-brown precipitate, Ag3Fe(CN),, , soluble in NH^OH .
Tin (stannous), white precipitate, Sn.1[Fe(CN),1], . soluble in hydrochloric acid.
Tin (stannic), no precipitate.
Uranium (uranoiis). no precipitate.
Zinc, orange precipitate. Zn,|Fe(CN).,]» . soluble in HC1 and in NH,OH .
Ferricyanides. ferric combinations, are capable of acting as oxidizing agents,
becoming ferrocyanides, ferrous combinations.
4K;,Fe(CN),; + 2H.S = :!K,Fe(CN), + H<Fe(CN)fl + 8,
L'K3Fe(CN)., + -'KI = 2K,Fe(CN),. + 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 nltroferrtcyanidcs or M//ro/>nm/</r« (1'layfair, I'hil. May., 1845,
(3), 26, 197, 271 and :i4K). These salts are generally held to have the composi-
tion represented by the acid H Fe(NO) (CN):. . Sixliinn nitroprUBflde is used as
a reagent for soluble sulphides — -that is, in presence of alkali hydroxides, a
test for hydrosulphuric acid: in presence of hydrosulphuric acid, a test for
alkali hydroxides (§207. Ofc).
K;,Fe(CN), is reduced to K,Fe(CN),. by Pd . Th . Mg and As. but not. i.y
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. ./. r., is7::, 26, 47::).
Pb" with potassium hydroxide forms PbO. and potassium ft-rrocyanide (Vfnttx'
Dirtioiiarn, issii, 2, :t-IO).
Sn" with potassium hydroxide forms potassium stannate, K_.SnO... and potas-
sium ferrocyanide (Wttttx' IHctioiHtri/. I.e.).
Cr'" forms in alkaline mixture a chroma te and a ferroeyanide (I'loxam, ''. \ '.,
1885, 52, 10!)).
Mn" with ]X)tassium hydroxide forms MnO. and potassium ferrocyanide
(Boudault. ./. //;•., 1845, 36. :.'.:).
Co" and Ni" are not oxidi/ed.
In alkaline solutions K:iFe(CN),; oxidizes sugar, starch, alcohol, oxalic acid
and indigo (Wallace. ./. C'., 185o, 7, 77; Mercer, /'////. M<i</., 1S47. (:;), 31, 12t>).
HNO, and HNOr. both form hydronitroferricyanic acid. H,Fe(NO)(CN), .
NO in alkaline solution becomes a nitrate (Wallace, 7. r.).
P in alkaline solution becomes a phosphate (Wallace, J. c.).
$233. CYANIC ACID. 271
HH..PO, forms H4Fe(CN),i and H.,PO4 .
HSS forms S, then H2SO, and H4Fe(CN),, (Wallace, I.e.).
SO; forms H2SO4 and H4Fe(CN)u .
Cl decomposes ferricyanides.
HC1OS acts upon K3Fe(CN),, , forming ])otassinin superferricyanide, K,Fe(CN),
(Skraup, A., 1877, 189, IHiS).
HI forms H4Fe(CN)0 and I .
Ferricyanides in solution are detected by the reactions with ferrous and
ferric salts (§126, 6b). 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 rediiction to ferrocyanide with KI in presence of con-
centrated HC1; the liberated iodine being titrated with standard Na,S,O5 .
Or it is reduced to ferrocyanide by boiling with KOH and FeSO4 , filtering,
acidulating with H.SO, and titrating with KMnO, .
§233. Cyanic acid. HCNO = 43.048 .
H — 0 — 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.
C.N, + 2KOH = KCNO + KCN + HO
KCN + PbO = KCNO + Pb
4KCN + Pb304 = 4KCNO + :;Pb
The free acid may be obtained by heating cyanuric acid, H3C3N303 , to
redness, better in an atmosphere of CO., . Cyanic acid is found in the dis-
tillate. H3C.N3O., = 3HCNO .
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 -+- H.O = NH, + CO, . The cyanates, therefore, when treated with
hydrochloric or sulphuric acid, effervesce with the escape of carbonic (inltjidndc
(distinction from cyanides), the pungent odor of cjiiniic ucid being perceptible:
2KCNO + 2H,SO,* + 2H.O = K,SO, + (NHJ.SO, + 2CO, . The iniimnnia
remains ?n 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 evolution of ammonia, tf-ilrcr C'jHiitntc is sparingly
soluble in hot water, readily soluble in ammonia: soluble, with decomposition,
in dilute nitric acid (distinction from cyanide). Copper ci/andfc is precipitated
greenish-yellow.
Ammonium rifimate in solution changes gradually, or immediately when boiled,
to iircfi, or carbamide, with which it is isomeric: NH.CNO = CO(NH,)o . 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 nitric
acid. Also, solution of urea with solution of mercuric nitrate, forms a white
precipitate, CH4NJ0(Hg'0)i; , not turned yellow (decomposed) by solution of
sodium carbonate (no excess of mercxiric nitrate being taken). Solution of
urea, on boiling, is resolved into ammonium carbonate, which slowly vaporizes:
CH4N,O + 2H,O = (NH4),CO:i . 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.
272 THIOCYANIC ACID. §234.
§234. Thiocyanic acid. HCNS = 59.118.
H — S — C = N.
An aqueous solution of HCNS may be obtained by treat ing^lead thioeyanate
suspended in water with H.S , also by treating1 barium thiocyanate with H»SO4
in molecular proportions. The anhydrous acid is obtained by treating dry
Hg(CNS)2 with H.S . Potassium thiocyanate is formed by fusing KCN with
S. Or two parts of K4Fe(CN),, with one part of sulphur. Also by fusing1 the
cyanide or ferrocyanide of potassium with potassium thiosulphate, K.S.Oj
2KCN + S2 = ^
K4Fe(CN),; + :'.S, = IKCNS + FelCNS),
4KCN + 4K,S,0, = 4KCNS + .iK.SO, + K,S
1-JK,S,0, = i -KCNS + OK,SO,
Thiocyanic acid is quite as frequently called sulphocyanie acid, and its salt.K
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 evolution of carbonic anhydride,
carbon disulphide, hydrosulphuric acid, hydrocyanic acid, ammonia, and othe,r
products.
The same products result, in greater or less degree, from transposing soluble
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, hixolublv
in water, are not all readily transposed. Thiocyanates of metals, whose sul-
phides are insoluble in certain at ids, 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 — arc soluble in
water. Mercuric thiocyanate, sparingly soluble: potassium mercuric thiocyanate,
more soluble. Nilrcr tliiori/oinitf 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), . crystalli/able in blue
needles, soluble in alcohol, not in carbon disulphide. The coloration is
promoted by warming, and the test is best made in an evaporating dish.
In strictly neutrnl solutions, iron, nickel, zinc and manganese, do n<»t,
interfere.
Copper, if concentrated, a black crystalline precipitate, Cu(CNS), , soluble in
thiocyanate. With sulphurous acid, a white precipitate. CuCNS: also with
hydrosulphuric acid (used to separate a thiocyanate from a chloride)
(Mann, Z., 1HH9, 28, 068).
Iron (ferrous), no precipitate or color.
Iron (ferric), an intensely blood-red solution of Fe(CNS), . decolored by solu-
tion of mercuric chloride (§126, tifr, <lixti>i<-H»ii from nt-ctic <tritl)\ decolored
by phosphoric, arsenic, oxalic and iodic acids, etc., unless with excess ,f
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.
Xiead. gradually, a yellowish crystalline precipitate. Fb(CNS). . changed by
boiling to whit" basic salt.
Mercury (mercurous), a white precipitate. HgCNS , resolved by boiling into
Hg and Hg(CNS), . The mercurous thiocyanate. HgCNS. swells greatly
on iiniiJian (being used in " Pharaoh's serpents "). with evolution of mer-
cury. nitrogen. thioc\ anogen. cyanogen and sulphur dioxide.
£235, 1. yiTROGEX. 273
Mercury (mercuric), in solutions not very dilute, a white precipitate,
Hg(CNS).. , somewhat soluble in excess of the thiocyanates, sparingly
soluble in water, moderately soluble in alcohol. On ignition, it swells like
the mercurous precipitate.
Platinum. 1'lj^kk- chloride, gradually added to a hot, concentrated solution
of potassiuraRiiocyanatc, forms a deep-red solution of double thiocyanate of
potassium and platinum (KClJS),Pt(CNS)4 , or more properly. K,Pt(CNS)(;,
potassium thiocyanoplattoate. 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
H...SO4 (separation from AgCl) (Volhard, A., 1877, 190, 1).
Certain active oxidizing agents, viz., nascent chlorine, and nitric acid contain-
ing nitrogen oxides, acting in hot, concentrated solution of thiocyanates, pre-
cipitate pcrthioci/anoycii, H(CNS)., , of a yellow-red to rose-red color, even blue
sometimes. 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,
perthiocj/anic acid, (HCN)2S3 , slightly soluble in hot water, from which it
crystallizes in yellow needles. It dissolves in alcohol and in ether.
Potassium 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 final 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 tame reaction is assumed as probable.
PbO.. and Pb3O4 form Pb" and sulphuric acid, in acid mixture only (Hardow,
J. C., 1859, 11, 174).
H. AsO; forms H AsO; , hydrocyanic and sulphuric acids.
Co'" forms Co" , hydrocyanic and sulphuric acids.
Ml"' forms. Ml" , hydrocyanic and sulphuric acids.
CrVi forms Cr"' , hydrocyanic and sulphuric acids.
Mn"+n forms Mn" . hydrocyanic and sulphuric acids. In alkaline mixture, a
cyanate and sulphate are formed (Wurtz's Diet. CJiim., 3, 95). '
HNO2 forms sulphuric acid and nitric oxide.
HN03 forms sulphuric acid and nitric oxide.
Cl forms at first a red compound of unknown composition, then HC1 , H.SO,
and HCM are produced. In alkaline mixture a chloride and sulphate are
formed.
HC10 same as with Cl .
HC1O3 forms sulphuric, hydrochloric and hydrocyanic acids.
Br forms HBr and H2SO4; but with alkalis, a bromide and sulphate.
HBr03 forms HBr and H2SO4 .
HIO3 forms H2S04 and free iodine.
§235. Nitrogen. N = 14.04 . Valence one to five
1. Properties.— Weight of molecule, M", , 28.08. Vapor density, 14 (Jolly, W.
A., 1879, 6, 536). At — 123.8°, under pressxire of 42.1 atmospheres, it condenses
to a liquid (Sarrau, C. r., 1882, 94, 718). Boiling point, —194.4° (Olszewski, W. A.,
1897, 31, 58). Liquid nitrogen is colorless and transparent. The gas is taste-
less, odorless and colorless. Not poisonous, but kills by excluding air from the
lungs. Does not burn or support combusion. It is very inert, not attacking
other free element?. Its simplest combinations are the following: N— '"H'8 ,
M2O , NO , N20,, , NO2 and N2O, . The number of organic compounds contain-
ing nitrogen is very large. The nitrogen in all compounds that are the
274 HTDKOMTItlV ACID. Jj235, t.
immediate products of vegetable growth has a valence of minus three anil
may without change of bonds be converted into N— '"H'a . This statement is
made with a limited knowledge of the facts and without, at present, having
conclusive proof; and merely predicting that future researc^^'ill verify it.
2. Occurrence. — It constitutes about four-fifths of the vol^le 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.
.'{. Formation. — («) From the air, the oxygen being removed by red-hot
copper, the CO2 by potassium hydroxide, the ammonia and water by passing
through H SO, . (l>) Ignition of ammonium dichromate, (NHJ.Cr.OJ = N^ -}
Cr,O3 -f 4H;,0 . (c) Uy heating ammonium nitrate and peroxide of manganese
to about 200° ((Jatehoiise, C. A., 1*77, 35, I IX). (d) Ignition of NH.C1 and
KjCr.O;: ^'NH4C1 + K,Cr,0T = 2KC1 + N, -\- Cr.O, + ^H,O . Unless the
temperature be carefully guarded traces of NO are formed, which may be
removed by passing the gases through FeSO4 . (r) Action of chlorine upon
NH,: sNH, + :!CL = t.NH^Cl + N2 . The NH3 must be kept in excess to
avoid the formation of the dangerously explosive chloride of nitrogen, NCI.! .
(f) Removing the oxygen from the air by shaking with NH,OH and copper
turnings. (//) Burning phosphorus in air over water. (/*) Uy passing air
through a mixture of FeS and sawdust; then through a pyrogallate solution,
and finally through concentrated H,SO, . (/) Hy shaking air with Fe(OH)z
and Mm OH).. ( /) J5y ]>assing air through an alkaline pyrogallate. (k) By
passing air, from which CO... has been removed, mixed with hydrogen over
heated plat i mini black, the hydrogen having been added in just sufficient
quantity to form water with all the oxygen ( Damon I'm. ./., 1851, :»21). (I) By
warming a concentrated solution of NH.NO, or a mixture of KNO., and NH<C1:
NH4NO. = NJ. + -H...O . Potassium dichromate is added to oxidize to nitric
acid any of the oxides of nitrogen that may be formed (Gibbs, tt., 1H77, i:'»87).
(m) Bv action of potassium or sodium hvpobromite upon ammonium chloride:
.(NaBrO + :»NH,C1 = N, -f :;NaBr + 2HC1 + 3H,O .
4. Preparation.— Nitrogen has been economically produced by most of the
above methods.
">. Solubilities. — Nitrogen is nearly insoluble in all known liquids.
(>. Reactions. — At ordinary temperatures nitrogen is not acted upon by other
compounds. Nodules growing on the roots of leguminous plants absorb nitro-
gen and build up nitrogenous compounds therewith.
7. Ignition. — Under electric influence it combines slowly with hydrogen;
also with B . Cr . Mg . Si and V .
s. Detection.— .Nitrogen is more easily detected by the nature of its com-
ponnds than by the properties of the liberated clement.
y. Estimation.- — («) As free nitrogen by measuring the volume of the gas.
(fo) 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. (<•) 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 CuO and Cu° , absorbing
the COj by KOH and determining the nitrogen by volume.
(For Hydroxylamine, see foot-note, page 278.)
§236. Hydronitric acid A/oimide). 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 :t7°. 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
$238, tf. NITROUS OXIDE— NITKIC OXIDE. 275
and the alkaline earths, are soluble in water and erystallizable (Dennis, •/. Am.
J$OC., 1S9S, 20, 225). Potassium trinitride precipitates from thorium salts, the
hydroxide of this metal in quantitative separation from cerium, lanthanum,
neodymium and praseodymium (Dennis, ./. Am. Roc-., 1H96, 18, 947). Hydro-
nitric acid is Ityrmed by treating- ammonia with sodium, and the resulting
sedamide, NaNH, , with nitrous oxide: 2NaNH., + N2O = NaN3 + NaOH -f-
JTH, (Wislicenus, B., 1892, 25, 2084).
§237. Nitrous oxide. N,0 = 44.08 .
N',0-", N — 0 — 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, Pi/roflH'miscli. I'litcrsuch., 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: («) By heating ammonium nitrate in a
retort from 170° to 260°: NH4NO, = N2O + 2H2O . (ft) By passing NO through
solution of SO2 . (c) By action of HNO:i; */>. <>r., 1.42, diluted with an equal
volume of water, upon metallic zinc, (d) A mixture of five parts of SnCl2 , ten
parts of HC1 , sp. (jr., 1.21, and nine parts of HNO3 , ftp. gr., 1.3, is heated to
boiling: 2HN03 + 4SnCL + 8HC1 = 4SnCl4 + N,O + 5H-O (Campari, J. ('..
J889, 55, 569).
§238. Nitric oxide. NO = 30.04 .
N"0-", N = 0 .
1. Properties. — The vapor density (15) shows the molecule to be NO (Daccomo
and Meyer, B., 1887, 20, 1832). Under pressure of one atmosphere it is
Mquified 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. — (a) Reduction of nitric acid by means of ferrous sulphate
previously acidulated with H SO, . (ft) Action of cold nitric acid, ,s'/j. (jr., 1.2,
upon metallic copper; unless great care be used other oxides of nitrogen are
produced, (c) SO, is passed into slightly warmed HN03 , sp. </•»•., 1.1.">, and
excess of SO2 removed by passing through water, (c/) According to Emich
(M., 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 HoS04 . sp. (jr., 1.84, and
1.50, dissolve 3.5 and 1.7 volumes respectively (Lunge, /?., 1885, 18, !;!)!). 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. Reactions. — 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 Cu,O: the higher oxides of
these metals not being thus produced (Sabaticr and Senderens, C. r.. 1892, 114,
1429). Oxidized to KNO., by KMnO,: KMnO, + NO = MnO, + KNO3 (Wank-
lyn and Cooper, Phil. M<i</., 'l878, (5), 6, 288).
276 \mtorx .!<•//>. §239,1.
§239. Nitrous acid. HNO, = 47.048 .
H'N"0-", , H — 0 — N = 0 .
1. Properties. — Nitrous acid is known only in solution. Almle liy adding
NS0S to water. Jt has a blue color and, owing to its tendency to dissociation
(tiHNO., = 2HNO + :NO + 2H.O), is very unstable (Fremy, ('*. i:, 1S7(). 70, (.1).
Nitrous anhydride is obtained when a mixture of one volume of oxygen and
four volumes of nitric oxide are passed through a hot tube, 4NO + 0. = -N...O., .
It is a deep red gas, eonder.sing to a blue liquid at 14.4° under 7."> mm. pressure
(Cuins, r. IV., 18S-J, 48, <>7).
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.
::. Formation. — liy action of nitric acid, x/i. <n\, 1.33, upon starch or arsenous
oxide. At 70° nearly pure N^O;. is obtained, which passed into cold water
forms HNO_. . 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 silver nitrite may be
made from these by precipitation: and from this salt m:iny nitrites may be
made nearly pure bv transaction: c. »/., Bad, -f 2AgNO. = Ba(NO,).J +
-'AgCl and then Ba(NO,)2 + ZnSO, = Zn(NO,), -f BaSO, .
4. Preparation.- Same as above.
."). 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 tin alkali reduces nitrates to nitrites; e.g.,
sodium amalgam, aluminum wire in hot KOH . etc. Used in excess the nascent
hydrogen reduces Ilu» nitrogen still further, forming NH:; .
(i. 'Reactions.- I. With metals and their compounds. — Nitrons acid acts
sometimes as an oxidi/er. sometimes as a reducer: in the former case NO is
uswtUu 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.
/. PbO2 becomes Pb" and nitric acid.
1. Hg' becomes Hg° and nitric acid.
.?. Crvi becomes Cr'" and nitric acid.
.'i. Co" becomes Co"' and nitric oxide. Excess of KNO, with acetic acid is
used to separate cobalt from nickel (§132. Or).
.7. Ni'" becomes Ni" and nitric acid.
<>'. Mn" -+- n becomes Mn" and nitric acid.
/*.— With non-metals and their compounds.
/. H,Fe(CN),, becomes first II1Fe(CN)« and then hydronitroferricyanic acid.
Solution of indigo in sulphuric acid is bleached by nitrites.
2. Nitrites are decomposed by nitric acid.
,t. HH,.PO, becomes H3FO4 and NO.
4. H..S does not displace or transpose alkali nitrites, but if acetic- acid In-
added to liberate the nitrous acid, then S° and NO are produced. H...SO, be-
comes H.SO4 and chiefly NO. With excess of H.SO:, . N,O or NH., is formed.
See Weber, Pogti.. ISfiC,, 127, 34:'. and isc,7. 130. 277: Fremy, C. >:. 1^70, 70. ''.I.
.7. HC1O becomes Cl° and HNO, .
6. HBrO becomes Br° and HNO3
7. HI becomes 1° and NO .
HI03 becomes 1° and HNOS .
7. Ignition.— In general nitrites are changed to oxides, but with potassium
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, (i) \
§241,4. NITROGEN PEROXIDE— NITRIC ACID. 27?
mixture of a nitrite and ZI liberates iodine on addition of acetic acid (nitrates
requiring a stronger acid for transposition). (3) Nitrous acid with iodic acid
liberates iodine, and nitric acid is produced. (J/) Solution of potassium per-
manganate acidified with sulphuric acid is reduced by nitrites (distinction from
nitrates).
9. Estimation. — Acidify with acetic acid, distil and titrate the distillate with
standard solution of permanganate.
§240. Nitrogen peroxide (dioxide). N02 = 46.04 .
Vapor density, 2?, (Rumsay, J. C., 1890, 57, 590). Melting point, —10°
(Deville and Troost, C. r., 1867, 64, 257). Boils at 21.64° (Thorpe, J. C., 1880,
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: 2NO2 +
2KOH = KNOa + KNO, + H20 .
§241. Nitric acid. HN03 = 63.048 .
0
II
H'NvO-";, , H — 0 — N = 0 .
1. Properties. — Nitric anhydride, N,O5 , is a colorless solid, melting at 30°
with partial decomposition to NO, and O, and if exposed to direct sunlight
decomposition begins at lower temperatures.
Nitric acid, HNO3 , has not been perfectly isolated; that containing 99.8 per
cent of HN03 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°: 4HN03 = 4NO2 + 2H,O + O2 (Carius,
B., 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 HNO3 (Kolbe, A. Ch., 1867 (4), 10, 136). This is the acid usually
placed, on the market. The reagent usually employed has a sp. 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 , NH4 , Ca , Mg , and
of a few other metals, the most abundant supply coming from Chili and
Bolivia as sodium nitrate, " Chili saltpeter."
3. Formation. — (a) Oxidation of nitrogenous matter in presence of air,
moisture and an oxide or alkali; (ft) by oxidation of NO , N,O3 or NO, by
oxygen (or air) in presence of moisture; (r) from NH3 , by passing a mixture
of NH3 and oxygen through red-hot tubes.
4. Preparation. — By treating nitrates with sulphuric acid and distilling.
Mtrates may be made: (a) By dissolving the metal in nitric acid, except
those whose metals are not attacked by that acid, e. g., An , Pt , Al and Cr ;
and also, antimony forms Sb.,05 , arsenic, H3As04 and \\rith excess of hot
acid tin forms metastannic acid H10Sn5013 . (ft) By adding HN03 to the
oxides, hydroxides or carbonates. All the known nitrates can be made
278 X1TRIC ACID. i<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 , Au 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. Sci., 1894, 163.)
The anhydride is made: (a) By passing chlorine over silver nitrate:
4AgNO., -f 2C1, == 4AgCl -f 2N,0, -f 02 . (6) By adding anhydrous P20,
to HN03 : ?NHO:1 + P.O. = 8HPO, + N20. .
5. Solubilities. — All normal nitrates are soluble. A few are decom-
posed by water, e. g., Bi(N03)3 + H,0 = = BiONO., -f 2HNO, . Most
nitrates are less soluble in nitric acid than in water, e. (j., Cd , Pb , Ba , etc.;
the barium nitrate being completely insoluble in HN03 , sp. gr., 1.42.
Nitric acid decomposes the sulphides of all ordinary metals, except
mercuric sulphide which by long continued boiling with the concentrated
acid becomes 2HgS.Hg(N03)2 , insoluble in the acid.
G. Reactions. A. — 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 ,
NH, , H.NOH *, N , N,0 , NO , HNO, , NO, . If the acid is concentrated,
in excess and hot, the product is usually entirely nitric oxide, colorless,
but changing to the red colored N02 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 NH3 , 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 HinSn.Oir, , Sb.,0, , and HnAs04 .
With the respective metals it forms Hg' or Hg", Sn" or Sn"", As'" or Asv,
Sb'" or Sbv, 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.
* Hydroxylamine, NH,OB. is formed by the reducing action of Sn and HC1 upon NO. N2Oa,
HNO,. etc. (Lessen, A., 18S8, 252,170); also by the action of II 2S, SO,. K. Ka, MS Z<i, and A 1 upon
HXO3. or by the action of II 28 upon certain nitrates (Divcrsand Haga, C. N., 1C80, 5 1, 271 . By
action of sodium amalgam upon s diumni rite solution, WII2OH Is pr duct d along wih nitrous
oxide, free nitrogen, ammoni.i, sodium hyponitritc, and sodium hydroxide the highest yield of
the hydroxylamino beins obtained when the nitrite solution is as dilute aa one in fifty the mix-
ture kept cold (Divers, J. G., 1890, 75, 87 and £9 . It is a bis •< with an alkaline reaction and a
strong reducing agent. When i tiro it is a crystalline solid, odorless, mclti: g at 33.C50, boilingat
68° at 22 mm. pressure; oxidised by oxygen t • UNO, (Lobry do nruyn.C., IS'JS, 2.). 3. i:0ai.d634>.
It is a good antiseptic a :d pr scrv.itivo. It combines with acids tj form salts: NH2OII + IK 1=
N H ..Oil . HCl. Ilydr.).\ylamino hydrochlorido is dccomprscd 1 y alkalis forming the free base,
\vhichisdecomposedbytbchalogens, KM aO^. K.,Ci-,o,, BaO, and ri>O . Its solution in eth«r
reacts with sodium forming: a white precipitate of 3iH,ONa.
§241, 7. NITRIC ACID. 279
1. Pb02 is not changed. Pb304 is changed thus: Pb.,04 4- tHNOa =
Pb02 + 2Pb(NO:!)2 + 2H20 .
2. Hg' becomes Hg".
3. Sn" becomes Sniv. Stannous chloride and hydrochloric acid, heated
with a nitrate, form stannic chloride, and convert nitric acid to ammonia
(which remains as ammonium salt). See §71, 6c.
4. Sb'" becomes Sbv, forming Sb20r, , insoluble.
5. As'" becomes Asv, forming H3As04 .
6. Qu' becomes Cu".
7. Fe" becomes Fe'".
B. — With non-metals and their compounds.
1. Carbon (ordinar}r, not graphite) becomes C02 if the nitric acid be
hot and concentrated.
H.,Co04 becomes C02 , in hot concentrated acid.
H4Fe(C]Sr)6 becomes first H,Fe(CN)6 and then hydronitrof erricyanic acid.
HCNS is oxidized, the sulphur becoming H2S04 .
2. Nitrites are all decomposed, nitrates being formed, the nitric acid
not being reduced. The nitrous acid liberated immediately dissociates:
3HNO, = 2NO -f HNO, + H20 .
8. P°, PH3 , HH.,P02 and H3P03 become H3P04 . That is Pv~n becomes
Pv.
4. S becomes H2S04 .
H2S becomes first S° and then H2S04 .
H2S03 becomes H2S04 ; and in general SVI~n becomes SVI.
5. HC1, nitrohydrochloric acid: 2HN03 + 6HC1 = 2NO + 4H20 -f 3C12
(Koninck and Nihoul, Z. anorg., 1890, 477). See §269, &B2.
HC103 is not reduced. Chlorates are all transposed but not decom-
posed until the temperature and degree of concentration is reached that
would dissociate the HC10a if the nitric acid were absent.
6. Br° is not oxidized. HBr becomes Br° and is not further oxidized.
All bromates are transposed but the HBrO , is not decomposed until a tem-
perature and degree of concentration is reached that would cause the
dissociation of the HBrO., if the nitric acid were absent.
7. 1° becomes HI03. 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.,
are inflamed by the strong acid (Kraut, B., 1881, 14, 301). For action
on starch, see Lunge, B., 1878, 11, 1229, 16-11. With many organic bodies
substiti'.tion products are formed, the oxides of nitrogen taking the place
of the hydrogen.
7. Ignition.— Nitric acid is dissociated by heat: 4HNO3 = 4NO, + 2H,O
complete if at 256° (Carius, B., 1871, 4, 828). No nitrates are volatile as such;.
280 NiTKir ir//>. §241,8.
ammonium nitrate is dissociated: NH4NO: = N,0 + ~'H20. 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 N2O*
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 oxidi/er: c. //., changing Sn" to Sn'v , As'" to Asv , Sb'" to Sbv ,
Te" 1o Fe"' , Cr'" to Crvi , Mnvi-n to Mnvi . Svi-n to Svi , etc.
Heated on charcoal, or with potassium cyanide, or sugar, sulphur or other
easily oxidi/able substance (as in gunpowder), nitrates are reduced with
dfflayrutton or cxploHion, more or less violent. With potassium cyanide, on
platinum foil, the deflagration is especially vivid. In this reaction free nitrogen
is evolved.
Strongly heated with excess of potassium hydroxide and sugar or other
carbonaceous compound, in a dry mixture, nitrates are reduced to (immonin,
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 of! abundant brown vapors; the colorless nitric
oxide, NO , which is set free from the mixture, oxidizing immediately in
the air to nitrogen peroxide, NO., :
2KNO, + IH.SO, + nCu = K2SO, + :!CuSO4 + 4H,O + 2NO
4H:S04 + GFeSO, = K2SO4 + :iFe2(S04), + 4H20 + 2ND
The three atoms of oxygen furnished by two molecules of nitrate suffice to
oxidize three atoms of copper; so that 3CuO with 3HJ504 , may form
:iCuS04 and 3H..O . The same three atoms of oxygen (having six bonds)
suffice to oxidize six molecules of ferrous salt into three molecules of
ferric salt; so that GFeS04 with 3H2S04 , can form 3Fe,(S04), and 3H.O .
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
black-brown liquid, (FeS04).,NO , decomposed by beat and otherwise un-
stable: 2KN03 + 4H2S04 + 10FeS04 = K,S04 + 3Fe,(S04), + 4H20 +
2(FeS04),NO .
a. — This exceedingly delicate " Brown ring " test for nitric acid or
nitrates in solution may be conducted as follows: If the solution of a
nitrate is mixed with an equal volume of concentrated IL,S04 , the mixture
allowed to cool and a concentrated solution of FeS04 then cautiously jidded
to it, so that the fluids do not mix, the junction shows at first a purple,
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, 8ft. NITRIC 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 H2S04 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.
b. — Indigo solution. — In presence of HC1 heat moderately and blue
color is destroyed. Interfering substances, HC10.( , HIO., , HBr03 , Fe'",
Crvl, Mnvn, and all that convert HC1 into Cl .
c. — Sodium salicylate is added to the solution, H2S04 is slowly added,
test-tube being inclined. Avoid shaking, keep cool for five minutes. A
yellow ring indicates HNO:j . To increase the brilliancy of the color,
shake, cool and add to HN4OH .
d. — Ammonium test. — Treat the solution with KOH and Al wire, warm
until gas is evolved. Pass the gas into water containing a few drops of
Nessler's reagent. A yellowish-brown precipitate indicates HNO.. :
3HN03 + 8A1 + 8KOH = 3NHy + 8KA102 + H20 . Nothing interferes
with this test, but action is delayed by Clv , Iv and many other oxidisers.
e. — Nitrite test. — Keduce 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 HC2H30, and test for I with CS2 . This test should
always be made in connection with (d). Other oxidisers including Clv,
Brv, 1^, and Asv are reduced before the reduction of the HNO., begins :
3HNO3 + 2A1 + 5KOH = ::KN02 + 2KA102 + 4H20
2KN02 + 2KI + 4HC2H:,0, = I. + 4KC2H:,O2 + 2H20 + 2ND
Other means of making the nascent hydrogen are sometimes preferred;
e. g., 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 H3P04 .
/. — Add three drops of the solution to be tested to two drops of
diphenylamine, (C(.H, )2NH , dissolved in H2S04 . A blue color indicates
a nitrate, Cl°, Clv, Brv, Iv, Mnvn, CrVI, SeIV", and Fe'" interfere with this
test.
g. — Brucine, 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.)
ft- — Phenol, C8Hr,OH , gives a deep red-brown color with nitric acid, by for-
mation of nitrophenol (mono, di or tri), C6H4(NO.,)OH to CBH2(N02)3OH ,
"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 OXYGEX. §241, 8r.
; •
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 (,/. (,'., 18&3, 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 (('. A"., 1SSS, 58, 1H>), 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
HC1 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 pyrogallol is dissolved in the liquid to be tested (less than one
ing. to one cc.) and ten drops of concentrated H..SO4 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 appears if nitric acid is present. One nag. of
nitric acid in one litre of potable water can thus be detected (Curtman, Arch,
/'harm., 1880, 223, 711).
9. Estimation. — («) If the base is one capable of readily forming a silicate,
the nitrate is fused with SiO, and estimated by the difference in weight. (I)) By
treating with hot sulphuric acid, passing the distillate into BnCO and esti-
mating the nitric acid by the amount of barium dissolved, (r) Treating with
Al and KOH and estimating the distillate as NH:i . (<l) Neutrali/.ing the free
acid with ammonium hydroxide, and after evaporation and drying at 115°,
weighing as ammonium nitrate, (e) In presence of free H SO, 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 NH3 that has been formed.
£242. Oxygen. 0=16.000. Usual valence two.
1. Properties. — A colorless, odorless gas; xprcifir <;rari4i/. 1.] 05(52 (Crafts, G. r.,
1888, 106, 1002). When heated it diffuses through silver tubing quite rapidly
(Troost, C. r., 1K84, 98, 1427). It liquifies by cooling the gas under great pres-
sure and then suddenly allowing it to expand under reduced pressure. It Itoils
jit — 11.'!° under 50 atmospheres pressure: and at — 184° under one atmosphere
pressure (Wroblewski, C. r., 1884, 98, .'!04 and 982). Its critical tcinpcrntiirc is
about — 118°, and the criliciil prrmturc 50 atmospheres. Specific (/niritii of the,
liquid at — 181.4°, 1.124 (Ols/.ewski, M., 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 (Carius, .4.., 1855, 94, 134). Molten silver absorbs
about ten volumes of oxygen, giving it up upon c-ooling (blossoming of silver
beads) (Levol, C. r., 1852. 35, ('>:;). It transmits soui-.d better than air (Bender,
B., 1873, 6, f>(55). It is not combustible, but supports combustion much better
than air. In an atmospln re of oxygen, a glowing splinter bursts into a flame;
phosphorus burns with vivid incandescence: also 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 of
oxygen with elements or compounds is termed combustion or oxidation. The
temperature at which the combination takes place varies greatly: Phosphorus
at <)0°: hydrogen in air at 5."2°: in pure oxvgen at 530° (Mallard and Le Cliate-
lier, /.'/., 1*v;. (2), 39. 2): carbon disnlpliide at 149°; carbon at a red heat;
while the halogens do not combine by heat alone.
2. Occurrence. — The rocks, clay and sand constituting the main part of the
«-a nh's crust contain from 44 to -4S per cent of oxygen; and as water contains
§242, 4/\ OXYGEN. 283
88.81 per cent, it has been estimated that one-half of the crust is oxygen.
Except in atmospheric air, which contains about 23 per cent of uncombiued
oxygen, it is always found combined.
3. Formation.— (a) By igniting HgO . (&) By heating KC103 to 350°, KC10<
is produced and oxygen is evolved; at a higher temperature the KC104 becomes
KCi . In the presence of Mn02 the KC103 is completely changed to KC1 at
200°, without forming KC104 , the Mn02 not being changed. Spongy platinum,
CuO , Fe2O3 , PbO2 , etc., may be substituted for Mn02 (Mills and Donald, J. C.,
1882, 41, IS; Baudrimoiit, Am. S., 1872, 103, 370). Spongy platinum, ruthenium,
rhodium and indium with chlorine water or with hydrogen peroxide evolve
oxygen. The spongy ruthenium acts most energetically (Schoenbein, A. Ch.,
1866, (4), 7, 10o). (c) Action of heat on similar salts furnishes oxygen; e.g.,
KC1O and KC102 form KCI , KBrO:t forms KBr , KIO, and KIO4 form KI ,
and KNO3 forms KNO2 (at a white heat K,O , NO and O are formed), (d) By
the action of heat on metallic oxides as shown in the equations below. ((?) By
heating higher oxides or their salts with sulphuric acid. CrVi is changed to
Cr'" , Co"' to Co" , Ni'" to Ni" , Biv to Bi"' , Fevi to Fe"' , Pbiv to Pb" , and
Mn"+n to Mn"; in each case a sulphate is formed and oxygen given off:
a. 2HgO (at 500°) = 2Hg + O2
ft. 10KC103 (at 3^0°) = GKC104 + 4KC1 + 3O2 (Teed, J. C., 1887, 51, 283)
2KC103 (at red heat) = 2KC1 + : O,
2KC1O, + nMnO, (at 200°) = nMnO2 + 2KC1 + 3O2
c. KC1O2 = KCI + O2
2KBr03 = 2KBr + 3O2
2KIOS — 2KI + 302
KIO4 = KI + 2O2
2KNO3 = 2KNO2 + O,
4KNO2 (white heat) = 2K2O -f 4NO + O2
d. 2Pb3O4 (white hrat) = CPbO + O2
2Sb2O5 (red heat) = 2Sb20, + O,
Bi2O5 (red heat) == Bi2O3 + O,
' 4CrOs (about 200°) = 2Cr,O, + 30,
4K2Cr2Or (red heat) = SCr.,0, + 4K2CrO4 + 3O,
6Fe2O3 (white heat) = 4Fe30, + O,
•5MnO, (white heat) = lCr.,0, + O,
f)Co2O3 (dull-red heat) = CoO, -f 0:
2Ni2O3 (dull-red heat) = 4KiO + O,
2Ag,O (300°) = 4Ag- + O,
2Ba02 (800°) = 2BaO + 0,
e. 2K,Cr2O7 + 8H2SO4 = 4KCr(S04), + 3O2 + 8H20
4KMnO4 + 6H..SO, — 2KJSO4 + 4MnSO4 + 5O, + 6H2O
2Pb3O4 + 6H2S04 — fiPbSO, + (iH.O + O,
4. Preparation. — (a) By heating KC1O3 to 200° in closed retorts in the pres-
ence of Mn02 or Fe203'. If KC1O3 be heated alone, higher heat (350°) is
required, and the gas is given off with explosive violence. About equal parts
of the metallic oxide and KC103 should be taken. (6) BaO heated in the air
to 550° becomes Ba02 , and at 800° is decomposed into BaO and O , making
theoretically a cheap process, (c) By heating calcium plumbate. The calcium
'plumbate is regenerated by heating in the air (Kassner, J. C., 1894, 66, ii, 89).
(d) By passing sulphuric acid over red-hot bricks: 2H2S04 = 2SO2 + 2H,O + O2;
the SO, is separated by water, and after conversion into H2S04 (§266, 4) is
used over again, (c) By warming a saturated solution of chloride of lime with
a small amount of cobaltic oxide, freshly prepared and moist. The cobaltic
oxide seems to play the same role as NO in making H,SO4 (Fleitmann, A. Cli.,
1865, (4), 5, .'JO"). '(/") The following cheap process is now employed on a large
scale. Steam-is passed ov< r sodium manganate at a dull-red lirnt: Mn...O3 and
284 OZONE. §242,5.
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 Mn20,
is thus again oxidized to Na,MnO, , and free nitrogen is liberated:
4Na;iMnO4 + 4H2O (dull-red heat) = SNaOH + 2Mn2O3 + CO2
8NaOH -f 2Mn203 + air, 3(Oa + 4NZ) = 4NaaMnO4 + 4H2O + 12N2
5. Solubilities.— See 1.
6. Reactions. — Pure oxygen 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, \\hen 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 except fluorine; the combination
occurring directly, at high temperatures, except with Cl , 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 NO2; by
its combination with phosphorus, etc. It is separated from other gases by
its absorption by a solution of ehromous chloride, pyrogallol or by phosphorus.
In combination in certain compounds it is liberated in whole or in part by
simple ignition; as with KC1O., , KMnO, . HgO , Au,O3 , PtO, , Ag,O , Sb.O,; ,
etc. In other combinations by ignkion with hydrogen, forming water.
9. Estimation. — Free oxygen is usually estimated by bringing the gases in
contact with phosphorus or with an alkaline solution of pyrogallol (CO., having
been previously removed), and noting the divmmition in volume. Oxygen in
combination is usually estimated by difference.
£243. Ozone. 0., = 48.000.
0 — 0
Ozone was first noticed by Van Marum in 1785 as a peculiar smelling ga.s
formed during the electric discharge; and which destroyed the lustre of
mercury. Schoenbein (Poyg., 1840, 50, 616) named the gas ozone and noticed
its powerful oxidizing properties. 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
Quntz, C. r., 1888, 107, 344; Wills, B., 1873, 6, 769). By the oxidation of moist
phosphorus at ordinary temperature (Leeds, A., 1879, '198, 30; Marignac, C. r.,
1845, 20, 808). By electrolysis of dilute 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, /. 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 ^..iall
animals which have been made to breathe it. For further concerning the
physiological action, see Binz, C. C., 1873, 72. Its specific gravity is 1.658 (Soret,
A., 1866, 138, 4). It has been liquified to a deep-blue liquid,, boiling at — 106°
(Olszewski, M., 1887, 8, 230). The gas is sparingly soluble in water (Carius, B.,
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. ff., 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 ozone oxidizes all metals except gold and platinum to the highest pos-
sible oxides.
Pb" becomes PbOo
Sn" becomes Sn02
Hg' becomes Hg"
Bi'" becomes BLO.,
Pd" becomes PdO2
Cr'" becomes Crvi
Fe" becomes Fe,O3 ; in presence of KOH , K,Fe04
Mn" becomes MnO2; in presence of H2SO4 or HNO3 , HMnO4 is formed.
Co" becomes Co'"
Ni" becomes Ni'" . With the salts of nickel and cobalt the action is slow,
rapid with the moist hydroxides.
K4Fe(CN)8 becomes K3Fe(CN)8
N2O3 becomes HN03 , in absence of water NO2 is formed
SO, becomes H2SO4
H2S becomes S and H,0 , the sulphur is then oxidized to H2SO4 (Pollacci,
G. C., 1884, 484)
P and PH3 become H3PO4
HC1 becomes Cl and H2O
HBr becomes Br and H..O
I becomes HIO3 and HE04 (Ogier, C. r., 1878, 86, 722)
HI and KI become I and H=O , then IV
Most organic substances are decomposed; indigo is bleached much moro
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 HN02 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 i's
much less delicate than with potassium iodide starch paper. It is estimated
quantitatively by passing the gas through a solution of KI rendered acid with
H2SO, , and titration of the liberated iodine: O3 + SHI = O., + I., + H..O .
§244. Hydrogen peroxide. H20, = 34.016 .
H — 0 — 0 — H.
1. Properties.— Pure hydrogen peroxide (99.1 per cent) is a colorless syrupy
hquid, boiling at 84° to 85° at 68 mm. pressure. It .does JUik, ^readily moisteii
the containing vessel. It is volatile in the air, irritating to the skin, and
286 HYDROGEX 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 bO per cent HO , losing
about one-half by volatilization. The presence of impurities causes its decom-
position with 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 preserved 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 100°. It contains the most oxygen of any known compound;
one-half of th« oxygen being available, the other half combining with the
hydrogen to form water.
2. Occurrence. — In rain water and in snow (Houzeau, C. r., 1870, 70, 510).
It is also said to occur in the juices of certain plants. '
3. Formation. — (a) By the electrolysis of 70 per cer.t H..SO, (Richarz, W. A.,
1887, 31, 912). (/>) 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, (c) By the action of oxygen and water on palladium sponge saturated
with hydrogen (Traube, /?., 1883, 16. 1201). (f) By the action of moist air On
phosphorus partly immersed in water (Kingzett, J. ('., 1SSO, 38. 3).
4. Preparation. BaO., is decomposed by dilute H,SO, . the BaSO, being
removed by filtration. The BnO is obtained by heating BaO in air or oxygen
to low redness. At a higher heat the BaO. is decomposed into BaO and O
(Thomsen, B., 1874, 7, 73). Sodium peroxide. Na,,O2 , is formed by heating
sodium in air or oxygen (Harcourt, J. C., 1802, 14, 2f>7); by adding H2O2 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,SiFn with formation of H,O, . NsuO, is soluble in
water with generation of much heat. It is a powerful oxidizing agent.
6. Reactions. .!.— With metals and their compounds. — Hydrogen
peroxide usually acts as a powerful oxidizing agent to the extent of one-
half its oxygen. Under certain conditions, however, it acts as a strong
reducing agent. Some substances decompose it into H.,0 and 0 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 highest oxides, e. //., Al , Fe , Mg , Tl , As , etc. Gold and
platinum are not attacked.
1. Pb" becomes Pb02 (Schoenbein, /. pr., 1862, 86, 129; Jannasch and
Lesinsky, /?., 1S93, 26, 2334).
2. Ag .0 becomes Ag and 0 .
8. HgO becomes Hg and 0 .
4. Au.,0, becomes Au and 0 .
5. As'" becomes Asv.
6. Sn" becomes Sniv.
7. Bi"' becomes Biv.
8. Cu" in alkaline solution (Fehling's solution) becomes Cu20 (Hanriott,
BL, 1886, (2), 46, 468).
9. Fe" becomes Fe'" (Traube, B., 1884, 17, 1062).
10. Tl' becomes Tl,03 (Schoene, A., 1879, 196, 98).
§244, 9&. HYDROGEN PEROXIDE. 287
11. Cr'" becomes CrVI in alkaline mixture (Lenssen, J. pr., 1860, 81,
278).
12. Cr" with H2S04 gives a blue color, HCr04 , percliromic 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 (Knglor and Wild, B., 1890, 29, 1940).
13. Mn" in alkaline mixture becomes MnO., . In presence of KCN a
separation from Zn (Jannasch and Mederhofheim, B., 1891, 24, 3945;
Jannasch, Z. anorg., 1896, 12, 124 and 134).
Mn"+n with H2S04 forms MnS04 , oxygen being evolved both from the
H202 and from the Mn compound (Brodie, J. C., 1855, 7, 304; Lunge,
Z. angew., 1890, 6).
14. BaO , SrO , and CaO become the peroxides.
15. NaOH becomes Na202.8H20 .
16. NH4OH becomes NH4N02 (Weith and Webber, B., 1874, 7, 17 and
45).
B. — With non-metals and their compounds.
1. K4Fe(CN)6 becomes K3Fe(CN)(. (Weltzien, A., 1866, 138, 129); in
alkaline solution the reverse action takes place : 2K3Fe(CN')(, -f- 2KOH +
H202 = 2K4Fe(CN)6 + 2H,0 -j- 02 (Baumann, Z. angew., 1892, 113).
2. 03 becomes 02 (Schoene, I. c., page 239.).
3. H3P02 becomes H3P04 .
4'. H2S and sulphides, and S02 and sulphites, become H2S04 or sulphates
(Classen and Bauer, B., 1883, 16, 1061).
5. Cl becomes HC1 (Schoene, /. c., page 254). It is a valuable reagent
for the estimation of chloride of lime : CaOCL -|- H202 = CaCl., -j- H00 -|r
02 (Lunge, Z. angew., 1890, 6).
6. I becomes HI (Baumann, Z. angew., 1891, 203 and 328). KC1 , KBr ,
and KI liberate oxygen f rom H202 but no halogen is set free ; except that
with commercial H202 free iodine may always be obtained from KI
(Schoene, A., 1879, 195, 288; Kingzett, J. C., 1880, 37, 805).
7. Ingition. — The peroxide of barium is formed by igniting- BaO to dull red-
ness; strong ignition causes decomposition of the BaO into BaO and O . The
peroxide of calcium cannot be formed by ignition of lime in air or oxygen.
8. Detection. — In a dilute solution of tincture of guaiac mixed with malt
infusion, a blue color is obtained when H^O, is added. To the solution sup-
posed to contain H.;O2 add a few drops of lead acetate; then KI , starch, and a
little acetic acid; with H.,O2 a blue color is produced (Schoenbein, 7. c.; Struve,
Z., 1869, 8, 274). As confirmatory, its action on KMnO4 and on K2Cr,O7 should
be observed. A ten per cent solution of ammonium molybdate with equal
parts of concentrated sulphuric acid gives a characteristic deep yellow color
with H2O, (Deniges, C. r., 1890, 110, 1007; Crismer, BL, 1891, (.1), 6. 22). H,O,
gives some extremely delicate color tests with the aniline bases (Tlosvay, /?.,
1895, 28, 2029; Deniges, J. PMrm., 1892, (5), 25, 591).
9. Estimation. — (a) By measuring the amount of oxygen liberated with MnO,
(Hanriott, BL, 1885, (2), 43, 468). (&) By the amount of standard KMnO,
288 FLUORIM:. $2*5.
reduced, or by measuring the volume of oxygen set i'ree: SKDCnO, -j- : H.SO, -!-
5H202 = K,S04 + :.'MnSO, + 8H,O -f 5O2 . (c) By decomposition of" KI in
presence of an excess of dilute H,SO. : and titration of the liberated iodine with
standard Na S 0 . (d) Dissolve a weighed sample of BaO in dilute HC1 , add
K,Fe(CN)9; transfer to an azotometer and add KOH . The volume of oxygen
is a measure of the amount of H202 (Baumann, I. c.).
§245. 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 limbs of which the
electrodes were inserted; 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 — 2iJ°^ The gaseous fluorine, which was extri-
cated at the positive electrode, was colorless, and possessed the properties of
chlorine, but much more strongly marked. It decomposed water immediately,
seizing upon its hydrogen, and liberating oxygen 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, 1886. C. r., 103. 202
and 256; J. C., 50, 1886, 8-49 and 976; .1. CJi., 1891, (6), 24, 224).
Fluorine, in several characteristics, appears as the first member of th<:
Chlorine Series of Elements. It cannot be preserved in the elemental stat.-,
as it combines with the materials of vessels (except fluor-spar), and instantly
decomposes water, forming Jij/dro fluoric arid. HF , an acid prepared by acting
on calcium fluoride with sulphuric acid («). Fluorine also combines with
silicon as SiF, , xUicoii fluoride, a gaseous compound, prepared by acting on
calcium fluoride and silicic anhydride with sulphuric acid (b). On passing
silicon fluoride into water, a part of it is transposed by the water, forming
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 (Jiitfrofttioxilicii- acid). (HF).,SiF4 or H.SiF,, (d) (Offermann.
Z. angew., 1890, 617). This acid is ur-ed as a reagent; forming metallic fluo-
silicates (silicofluorides), soluble and insoluble (§246).
a. CaF, + H,S04 = CaS04 + 2HF
6. 2CaF, + SiO, + 2H2SO. = 2CaSO4 + 2H20 + SiF4
c. SiF4 + 2H,O = SiO, + ^HF (not remaining free)
d. 2HF + SiF4 = H,SiFa
c and d. 3SiF4 + 2H,O = SiO, + 2H2SiF,
$247. HYDROFLl ORIC 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 platimim 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 H2F., . But at 100° it is only 10, indicating
that at that temperature the molecule is HF . 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 litmus 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, Cal\ , 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, BaF., . Silver nitrate
gives no precipitate.
Sulphuric acid transposes fluorides, forming hydrofluoric odd, HF (§245, a).
The gas is distinguished from other substances by etching hard (jlass — 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
ncid, silicon fluoride, SiF4 (§245, 6); a gas which does not attack glass, but when
passed into water produces fluosilicic acid, H2SiFfi (§245, c and d). 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 siirface of the glass above the material is corroded and roughened: CaF., +
2KHSO4 = CaSO4 -f- K2SO4 + 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 yellomsh-green
color is imparted to the flame.
§247. Fluosilicic acid. H2SiF0= 144.716.
Fluosilicic acid* (Ju/drofluosUicic and), (HF),£iF4 , or H2SiF6 , is soluble in
water and forms metallic fluosilicates (silicmfliHtridwf), mostly soluble in water;
those of barium (§186, 6i), 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, B"F . By heat,
they are resolved into fluorides and silicon fluoride: BaSiFu = BaF2 + SiF4 .
* Fluosilicic acid is directed to be prepared 1>v ta'dng o-ie "art each of fine sand and flne'y pow-
dered fluor-spar, with six to right parts of concentrated sulphuric acid, in a small stoneware
bottle or a glass flask, provided wi-h a wide de'iv ry-tube, dipping into a little mercury in a
small porcelain capsule, \vhich is set in a large beaker containing six < r 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 gentlj heated from a lamp. E .ch bubble of gas decomposes with deposition
of gelatinous silicic acid. When the water is filed with this deposit, it may be separated by
straining through cloth and again treating with the gas for greater concentration. The strained
liquid is finally filtered and preserved for use.
290 SILICON— SILICON DIOXIDE. $248.
§248. Silicon. Si = 28.4. Vuleaee lour (§15j.
There are three modifications of silicon: (a) Amorphous. — A dark brown
powder; specific yrarity, 2.0; non-volatile; infusible; burns in the air, forming
SiO, , and in chlorine, forming- SiCl4 . It is not attacked by acids except HlT:
Si + GHF = H,SiFn + 2H, . It is dissolved by KOH'with evolution of
hydrogen. (6) Graphitoidiil. — 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,SiF,, . It is attacked slowly by fused KOH .
(c) Adamantine silicon, crystalline silicon. — Grayish-black, lustrous, octahedral
crystals, formed by fusing the/graphitoidal form. Specific gravity, 2.49 at 10°
(Woehler, A., 1856, 97, 2(>1). lyscratches 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 graf! hitoidal form, being attacked with even greater
difficulty. Silicon is never found free in nature, but always in combination as
silica, SiO;. , or as silicates.
Amorphous silicon is formed by passing vapor of SiCl, over heated potassium;
by heating magnesium in SiF, vapor; by heating a mixture of Mg and SiO_.; 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, R., 1889, 22, 18(5). 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. N., 1891, 63, 46). The crystalline form is made
by fusing a silicate or K SiF with Al; by passing vapors of SiCl, over heated
Na or Al in a carbon crucible (Deville, *.l. Ch., 1857, (3), 49, 62; Deville and
Caron, A. Ch., 1863, (3), 67, 435; Woehler, /. c.).
$249. Silicon dioxide. SiO, = 60.4.
(Silicic anhydride; silica.)
Silicic acid. H,Si03 = 78.416 .
Si^O-2, and H'2SiIV0-", ,0 = Si = OandH — 0 — Si — 0 — H.
1. Properties. — Silica, xilicic anhydride, SiO2 , is a white, stable, infusible solid;
insoluble in water or acids; soluble in tixed alkalis with formation of silicates.
Specific y rarity of quartz, 2.647 to 2.652; of amorphous silica, 2.20 at 15.6°.
Silicic acid, xilicim hydnifidr, H SiO , is a white, gelatinous solid, generally
insoluble in water, and soluble in mineral acids. A dilute solution in water is
obtained by dialysis of the fixed alkali silicate with an excess of HC1 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 combined
with oxygen in the form of silicon dioxide, SiO,, as quartz, opal, flint, sand,
etc.; or the silicon dioxide is in combination with bases as silicates: asbestos.
Boapstone, mica, cement, glass, etc. All geological formations except chalk
contain silicon as the d'oxide or as a silicate.
3. Formation. — Crystalline silica is formed by passing silicon fluoride into
water, forming silicic acid and llimsilicie acid: :SiF, + :iH,0 = H.SiO;, +
•JH SiF,, . The precipitate of silicic aeid is dissolved in boiling NaOH and then
heated in sealed tubes.' Below ISO0 crystals of tridymite arc formed, and
above ISO0 crystals of quartz (Maschke, /V///., is: 2, 145, r>-i9).
4. Preparation. — lure amorphous silica is prepared by fusing finely pow-
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 p-elnti-iors rrceipitr.t ;• \\aslied, dried and ignited. Crystalline
$249, 7. SILICON DIOXIDE. 291
silica is prepared by fusing silicates with microcosmic salt or with borax
(Rose, ./. /»•., 1867, 101, 228).
Silicic acid. — The various hydroxides of silica act as weak acids. Metasilicic
acid, H SiO, , has been isolated; it is formed by decomposing1 silicon ethoxide,
Si(OC2H5)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 gelatinise, and drying
over sulphuric acid. Other hydroxides, acids, have been isolated, but there is
some question as to their exact composition.
5. Solubilities.— Silica, SiO,, is insoluble in water or acids except HF ,
which dissolves it with formation of gaseous silicon fluoride, SiF4 (§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, Si02 , 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
separation of the anhydrous Si02 ; which is insoluble in more of the samo
acids, which previously had effected its solution.
' The most of the silicates found in nature are of complex composition.
They are combinations of Si02 with bases. They are, as a rule, insoluble
in water or acids.
6. Reactions. — Solutions of the alkali silicates precipitate solutions of
all other metallic salts with formation of insoluble silicates; they are
decomposed by acids with separation of silicic acid, a gelatinous precipi-
tate, soluble in hydrochloric acid. Evaporation decomposes silicic acid
with separation of insoluble silicic anhydride, Si02 . 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
group on the addition of ammonium chloride.
Silica, Si02 , is soluble in hot fixed alkalis forming silicates; it is not
soluble in ammonium hydroxide, nor are solutions of alkali silicates pre-
cipitated on addition of ammonium hydroxide as they are on the addition
of ammonium salts. Boiling Si02 with the fixed alkali carbonates forms
soluble silicates with greater or less readiness. Nearly all si'.icates aro
decomposed by heating in sealed tubes to 200° with concentrated HC1 or
H2S04 .
7. Ignition. — Silicates fused with the alkalis form soluble alkali sili-
cates, and oxides of the metal previously in combination. If alkali car-
bonates are employed the same products are formed with evolution of
C02 . Preferably a mixture (in molecular proportions) of potassium and
sodium carbonates, four parts, should be used to one part of the insoluble
silicate. Silica, Si02 , is also changed to a soluble silicate by fusing with
fixed alkali hydroxides or carbonates.
292 PHOSPHORUS. §249, 8.
SiOj does not react with K SO or Nn_SO . even when fused at a very high
temperature (Mills and Meanwell. J. C., 1881. 39, 533). In the fused bead of
microcosmic salt particles of silica swim iintiixxnlrrti. If a silicate be taken,
its base will, in most cases, be dissolved out, leaving 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 w:itcr,
the solution contains all the alkali metals, as chlorides, with calcium chloride
and hydroxide.
8. Detection. — Silicates are detected by conversion into the anhydride,
SiO., . 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 HC1 and thoroughly pulverized; water
is added and the precipitate of Si02 is thoroughly washed. Further con-
firmation may be obtained by warming the precipitate of SiOJ \vit.li
calcium fluoride and sulphuric acid (in lead or platinum dishes), forming
the gaseous silicon fluoride, SiF4 . This is passed into water wlinv it i<
decomposed into gelatinous silicic acid and fluosilicic acid: 33iF4 + 3H20
= H2SiOa -f 2H,SiF(. (§246). Silica, SiO, , 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 SiO, determined by weighing after ignition.
§250. Phosphorus. P = 31.0. Usual valence three or five (§11).
1. Properties. — Phosphorus is prepared in several allotropic modifications.
Specific ijrtirUy of the yellow, solid, at 20°, 1.82:121: liquid, at 40°, 1.74024; solid,
at 44°, 1.80681 (1'isati and cle Franchis, B., 1S7J, 8, 70). At ordinary tempera-
tures it is brittle and easily pulverized. At about 4.">° it melts, but may be
cooled in some instances (under an alkaline liquid) as low as +4° without
solidifying. When it solidifies from these lower tempi ranin -s. as it docs l>\
stirring with a solid substance, the temperature immediately rises to about 45°.
Boilinu point, 287.3° at 7(52 mm. pressure (Schroetter, A., 1848, 68, 247; Kopp, A.,
1855, 03, 129). The density of the vapor at 1040° is 4.50 (Deville and Troost,
C. r., 1863, 56, 891). The computed density for the molecule P4 is 4.294. At a
white heat the density. ::.«,::.'. indicates dissociation of the molecule to P3
(.Meyer and Biltz, B., 1889, 22, 72.1). N/«r///V <iniritn of the red amorphour.
modification at 10°, 1.964.
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 fairly
large quantities and cooled slowly it forms dodecahedral and octahedral crys-
tals (Whewell. C. A".. 1ST!), 39. 144). Heated in absence of air above the boiling
point it sublimes as a colorless gas. depositing lustrous transparent crystals
(Blondlot, C. r.t 186(i, 63. ::(.»7). At low temperatures phosphorus oxidizes slowly
in the air with a characteristic cdor, probably due to the formation of ozone
$250,4. PHOSPHORUS. 293
and phosphorous oxide, P2O3 (Thorpe and Tutton, «7. C., 1890, 57, 573). It ignites
spontaneously in the air at 00°, 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 phosphorus. Burns caused by it are very painful and heal with
great difficulty. 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, , CS., , Br , Cl , etc., prevent the glowing
.(Schroetter, ,/. 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 sufficient 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, Poffff., 1865, 126, 193). If
amorphous phosphorus be distilled in the absence of air, it is changed to the
crystalline form, action beginning at 260°. Heated in the air from 250° to 260°
it takes fire (Schroetter, I. c.). A black crystalline metallic variety of phos-
phorus is described by Hittorf (I.e.); also Remsen and Kaiser (Am., 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 istiever 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 phosphoriis is formed by heating calcium or lead
•phosphates with charcoal. The yield is increased by mixing the charcoal with
sand or by passing HC1 gas over the heated mixture. By igniting an alkali
or alkaline earth phosphate with aluminum (Rossel and Frank, B., 189-4, 27, 52).
Red phosphorus is formed by the action of light, heat or electricity on ordinary
.phosphorus (Meyer, B., 1882, J5, 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 chiefly of Ca3(PO4);,; then
H.,S04 is added, producing soluble calcium tetrahydrogen diphosphate (a).
After filtering from the insoluble calcium siilphate the solution is evaporated
and ignited, leaving calcium metaphosphnte (1>). Then fused with charcoal,
reducing two-thirds of the phosphorus to the free state (c). The mixture of
'sand, SiO., , 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 (e) and (f).
(a) Ca3(P04)2 + 2H,S04 = 2CaSO, -f CaH4(PO4).
(6) CaH4(P04), + ignition = Ca(PO,)2 + 2H2O
(c) 3Ca(F03), + IOC = Ca3(P04)2 + 10CO + P4
(d) 2Ca(PO,)2 + IOC + ^£10, = SCaSiO, + P4 + 10CO
(c) 2Ca3(P04);: + ICC + 12HC1 = CCaCL + P4 + 1«CO + GH2
(f) 2Ca(PO3), + 120 + 4HC1 = 2CaCl, + ?4 + 12CO + 2H3
Red or amorphous phosphorus is prepared by heating ordinary phosphorus
for a long time (40 hours) at 240° to 250° in absence of air. At 2fiO° 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 CS2 to remove
all traces of yellow phosphorus and is dried at 100°.
294 PHOSPHORUS. S250, r>.
'5. Solubilities. — 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, lied
phosphorus is insoluble in water, ether, or earbon disulphide.
6. Reactions. — When phosphorus is boiled with a fixed alkali or alkaline
earth hydroxide, phosphorus hydride, phosphine (§249), PH3 , and a
hypophosphite (£250) are formed. Phosphorus, when warmed in an
atmosphere of N or CO., , 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,
PH.J . 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 heat. It is ignited when brought in contact with PbO.> , Pb304
HgO , Ag.O , CrO., , K,Cr,07 and when heated with CuO or MnO, . Solu-
tions of platinum, gold, silver, and copper salts are decomposed by phos-
phorus with separation of the corresponding metal (Boettger, J. C., 1874,
27, 10GO).
With HN03 , H,P04 and NO are formed ; when heated with KN03 a
rapid oxidation takes place.
It combines with oxygen, forming P203 or P205 . 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 H3P04 (Schroetter, /. c.).
Combination with red phosphorus and sulphur takes place at ordinary
temperatures, forming P2S3 or. P2SB , depending upon the proportion of
each employed (Kekule, A., 1854, 90, 310). With ordinary phosphorus
the action is explosive.
Cl 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 hydracids and phosphorous or phosphoric-
acids, depending upon the degree of oxidation of the phosphorus. In
the presence of water phosphorus is oxidized to H3P04 by Cl , Br , I ,
HC103 , HBrO:! , or HIO., with formation of the corresponding hydracid:
P4 + 10C12 + 16H,0 = 4H3P04 -f 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 P2Mga , which can be
heated to redness in absence of air without decomposition. Heated in the air
it becomes oxidized (Blunt, A. Ch., 1865, (4), 5, 487). I'hosphorus also combines
with Cu , Ag , Cd , Zn and Sn when it is heated with these elements in sealnJ
tabes. It does not combine with Al and but slightly with Fe (Emmerlintf,
J. C., 1879, 36, 508).
§252,3. PBOKl'HIMl—lJYI'Ol'llOHl'HOHOrN ACID. 295
: 8. Detection. — By its phosphorescence; by formation of PH3 when
boiled with KOH (IJ of maim, /?., 1871, 4, 200); by oxidation to H,P04 and
detection as such (§75, 6d).
'J. Estimation. — Oxidation to H3PO4 , precipitation with magnesia mixture as
MgNH4PO4 , ignition to, and weighing as Mg,P,O7 (§189, 9).
§251. Phosphine. PH3=: 34.024.
P-'"H'S , H — P — H .
Phosphine, PH3 , 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: 2PH3 + 40. = 2HPO3 + 2H,O . It is
fiquified and frozen at very low temperatures; lioiUiiff point, about — 85°;
melting point, — 132.5° (Olszewski, M., 1886, 7, :>71). 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 H3PO2 or H3PO3 (6) ; by
ignition of hypophosphites (c); by the decomposition of the alkaline earth
phosphides with water or dilute acids (d) :
(a) P4 + 3KOH + 3H2O = 3KH2P02 + PH3
(6) 2H3P02 = HP03 + PH, + H20
4H3POS = 3HPO3 + PH3 + 3H2O
(c) 4NaH2PO, — Na4P207 + 2PH3 + H2O
(d) Ca.P, + 6H,O = 3Ca(OH)a + 2PH,
Ca,P, + 6HC1 = 3CaCl2 + 2PH3
It is a strong reducing agent; transposes many metallic solutions: 3CuSO4 -f-
2PH, = Cu3P2 + 3H2SO4; reduces solutions of silver and gold to the metallic
state: 8AgNO3 + PH3 + 4H=O = H3P04 + 8HN03 + 8Ag; is oxidized to H3PO4
by hot H2S04 , Cl , HC1O , HNO., , HNO:! , H:!AsO4 , etc. A liquid phosphorus
hydride, P2H4 , and a solid, P4BL , are known (Besson, C. r., 1890, 111, 972;
Gattermann and Hausknecht, B., 1890, 23, 1174).
§252. Hypophosphorous acid. H3P02 = 66.024 .
H
I
H'3P'0-"2 . H — 0 — P = 0 .
H
1. Properties. — Hypophosphorous acid was discovered in 1816 by Dulong (A.Ch.,
1816, 2, 141). It is a colorless syrupy liquid; specific yravity, 1.493 at 18.8°. At
17.4° it becomes a white crystalline solid (Thomsen, H., 1874, 7, 994). Although
containing three hydrogen atoms it forms but OHO series of salts, e. g., NaH2POn,
BaCK^PCX)., , etc.
2. Occurrence. — Not found in natuie.
3. Formation. — All ordinary iiu-t;ils form liypophosphites except tin, copper
and mercurosum. Silver and ferric hypophosphites are very unstable. (1> A
arPOPBOaPBOMOUS ACID. $252, I.
few metals, such as /inc and iron, dissolve in H:.PO, , giving- off bydrog^O and
forming a bypophoaphite. (2) The alkali and alkaline earth salts may In-
formed by boiling- phosphorus with the hydroxides (Mawrow and Miithmann,
X. anyew., 185)0, ii, 2(58). (3) As all hypophosphites are soluble, none ean be
formed by precipitation. All may be formed from their sulphates by trans-
position with barium hypophosphite. (4) All may be made by adding- H3PO,
to the earbonates or hydroxides of the metals.
4. Preparation. — To prepare pure H^PO., , BaO 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 H,SO, .
The solution is filtered and evaporated in an open disli, care being taken not to
heat above 110° . t'pon 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. — .1. — With metals and their compounds. Hypophosphorous
aeid is a very powerful reducing agenl. being1 oxidized to phosphoric acid or a
phosphate.
/. Pbiv becomes Pb" in aeid or alkaline mixture.
2. Ag' becomes Ag° in acid or alkaline mixture.
3. Hg" becomes Hg' and then Hg° in acid or alkaline mixture.
-}. Asv and As'" become As0 in presence of HC1 .
5. Bi'" becomes Bi° in presence of alkalis or acetic acid.
6. Cu" becomes Cu H 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.
9. Co'" becomes Co" , no action in alkaline mixture.
10. Ni"' becomes Ni" , no action in alkaline mixture.
//. Mn"+n becomes Mn" in acid solution.
12. Mniv+n becomes Mniv in alkaline mixture.
B. — "With non-metals and their compounds.
/. H3Fe(CN)0 becomes H4Fe(CN), .
2. HNO., and HNO, become NO .
3. H3PO. on heating becomes H.,P04 and PH3 .
4. HjSOj becomes free sulphur with formation of some HaS (Ponndorf, J. O.,
1877, 31, 27.'.).
H SO, becomes first H SO then S . See above.
5. Cl becomes HC1 in acid mixture, a chloride with alkalis.
HC10 and HC1O:1 form same products as Cl .
6. Br becomes HBr in acid mixture, a bromide with alkalis.
HBrO3 forms HBr .
7. I forms HI, in alkaline mixtures an iodide.
HI. dry, reacts violently, forming H,PO:I and PH4I (Ponndorf, I.e.).
HIO3 forms HI .
7. Ignition.- -On ignition hypophosphites leave pyrophosphates, evolving PE,.
The acid decomposes on heating to PH, and H3P04 (or HPOS if at a red heat).
8. Detection.— II vpophosphorous acid may bo 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 copper
hydride (Cn2H2) is thrown down, which, when heated in the liquid to 100°,
is decomposed with the deposition of the metal and the evolution of
hydrogen, whilst with phosphorous acid the metal is precipitated and
hydrogen evolved, but no Cu.,H2 is formed. Further, hypophosphorous
acid reduces the permanganates immediate!}, but phosphorous acid only
after some time. Phosphites precipitate barium, strontium, and calcium
§253, fi. PHOSPHOROUS ACID. 297
salts, while hypophosphitcs do not. When hypophosphorous acid is
treated with zinc and sulphuric acid it is converted into phosphoretted
hydrogen. On boiling hypophosphorous 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 HC1; the
temperature must not rise above 60°, otherwise metallic mercury will be
formed. The precipitated HgCl , after washing and drying at 100°, is weighed.
NaH2P02 + 4HgCl, + 2H20 = 4HgCl + H3PO4 + NaCl + 3HC1
§253. Phosphorous acid. H.5PO, = 82.024 .
H
H'8F"0-", , H — 0 — P — 0— H.
I
0
1. Properties. — Phosphorous anhydride, P208 , is a snow-white solid, melting
a* 22.5°, and boiling at 173.1° (Thorpe and Tutton, J. C., 1890, 57, 545). The
Vftpor density of the gaseous oxide indicates the molecule to be P4O6 . Specific
gravity of the liquid at 21°, 1.9481; of the solid at the same temperature, 2.135.
Tt has an odor resembling that of phosphorus. Heated in a sealed tube at
260° it decomposes into P2O4 and P (T. and T., ./. (7., 1891, 59, 1019). It reacts
slowly with cold water, forming H3PO:!; with hot water the reaction is violent
and PH3 is evolved. Upon exposure to the air it oxidizes to P205 .
The acid, H:tPO:i , is a crystalline solid, very deliquescent, melting at 74°
(Hurtzig and Geuther, .4., 1859, 111, 171). It is a dibasic acid, forming no
tribasic salts (Amat, C. />, 1889, 108, 403). One or two of the hydrogen atoms
are replaceable by metals forming acid or normal salts. The third hydrogen
ie never replaced ity 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 pyrophosphites have been pre-
pared (Amat, C. r., 1888, 106, 1400; 1889, 108, 1050; 1890, 110, 1191 and 901;
A. Ch., 1891, ((>), 24, 289).
2. Occurrence. — Does not occur in nature.
3. Formation. — P20:, is formed together with P205 when phosphorus is
ignited in the air. HaPO- is formed together with H3P04 when phosphorus
is oxidized with HNO3; by the oxidation of H3P02; by the action of P upon a.
concentrated solution of CuS04 in absence of air: :>CuSO4 + P4 + <>HoO =
Cu3P, + 2H:1PO:! + I]H,SO4 (Schiff, ,1., 1SCO, 114, 200).
4. Preparation To prepare phosphorous anhydride, P20:; , phosphorus is
burned in a tube with an insufficient supply of air (Thorpe and Tutton, I.e.).
The acid. H:,PO:, . is prepared by dissolving the anhydride in cold water; by
decomposing PCI, with water (Hurtzig and Geuther, 1. <•.).
5. Solubilities. — The acid is miscible in water in all proportions. Alkali
phosphites are soluble in water, most others are insoluble (distinction from
hypophosphites).
<>. Reactions. Phosphorous acid is a strong reducing agent, oxidi/ing 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
oxidi/ing acids and by many of the higher metallic oxides. HgCL becomes
and then Hg° . CuCl, becomes CuCl then Cu° (Uammelsberg, •/. C., 1873,
298 ffYPOPHOSPHORIC ACin—PHObPUOKir A<'ll>. ^253,7.
26, 13). Concentrated H.SO, with heat forms H,PO, and SO. (Adie, J. C., 1SD1,
59, 2:JO). H.SO, forms ITS and H:,P04 (Woehler, .1., 1841, 39, 252). Nascent
hydrogen (Zn and H.SO,) "produce PH3 (I)nsart, C. r., IS.™, 43, 1126).
7. Ignition. — The acid is decomposed by ignition, forming HPO3 and P or
PH:; (Vigier, Rl., 1809, (2), 11. 125; Hurtzig and Geuther, I.e.). Phosphites are
decomposed by heat, leaving a pvrophosphate and a phosphide and evolving
PH3 or H (Rammelsberg, /?., 187<i! 9. 1577; and Kraut, A., 1S75, 177, 274).
8. Detection. — By oxidation to H,PO4 and detection as such. It is distin-
guished from hypophosphorous acid by reducing CuSO to Cu°, while the
latter forms Cu.H;.; also by the solubilities of the salts (§252, 8). Its reactions
with oxidizing agents distinguish it. with hypophosphorous acid from phos-
phoric acid.
9. Estimation. — By oxidation to H3PO4 and estimation as such.
$254. Hypophosphoric acid. H4P,0(; = 1(52.032 .
0 0
II II
H'4PIV,0-"0, H — 0 — P — P — 0 — H.
I I
0 0
I I
H H
Hypophosphoric acid is formed together with phosphorous and phosphoric
acids by slowly oxidizing phosphorus in moist air (Salzer, A., 1885, 232, 114
and 271); also by oxidizing phosphorus with dilute HNO, in presence of silver
nitrate (Philipp, /*., 1SH5. 18, 74'.)). It consists of small colorless hygroscopic
crystals which melt at 55°. It decomposes when heated to 70° into H3PO, and
HPO3 , and at 120° gives H,P,O? and PHa (.loly, C. r., 18805 102, 110 and 7GO).
It is oxidized to H;PO, by warm HNO . slowly by KMnO, in the cold, rapidly
when heated. It is not oxidized by H,O: , chlorine water or H,CrO4; HgCI,
becomes HgCI (Amat, C. /'.. IS'.K). Ill, <>7<>). It is not reduced by Zn and H,SO,
(distinction from H,PO, and H.PO,). With a solution of silver nitrate it gives
a white precipitate which does not blacken in the light (distinction from HjPO.
.and H,PO3). It forms four series of salts, all four hydrogen atoms being
replaceable by a metal. The hypophosphates are much more stable towards
oxidizing agents than hypophosphites or phosphites.
£255. Phosphoric acid. H,PO, = 98.024 .
0
II
H',PvO-"4 ,H — 0 — P — 0 — H.
I
0
I
H
1. Properties.— Phosphoric anhydride, P20., *, 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: orthn. H3PO4: mcta. HPO, ;
•According to Tildeii and Hanu-tt J. C., 1896, 69, 154) tne moleciiU- is P4O10 not P,Ot ; P4O.
not P,03 < Thorpe and Tutton, J. ('., 1W»1, 59. 1032) ; un<, P^S, not P,..S3 (Tsamlxjrt. C.r., 1886. 1 oa.
1386.
§255, 3. PHOSPHORIC ACID. 299
and pyro, H4P2O7 . Orthophosphoric acid is a translucent, feebly crystallizable
and very deliquescent soft solid. Specific gravity, 1.88 (Schiff, A., 1860, 113, 183);
melting 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'H2P04 , primary, monobasic or mono-
metallic phosphates; M'2HPO4 , secondary, dibasic or climetallic phosphates;
and M'3PO4 , tertiary, tribasic, trimetallic or normal phosphates. The first
two are acid salts, but Na2HPO4 is alkaline to test paper. Metaphosphoric
acid, HPO3 , H — O — P = O , is a 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 differ-
ences of the acids and their salts (Tammann, Z. phi/s. Ch., 1890, 6, 122).
0 O
II II
Pyrophosphoric acid. H4P2O7 . H— O — P — O — P — O — H,isa glass-like
O O
I I
H H
solid (Peligot, A. Ch., 1840, (2), 73, 286), very soluble in, but unchanged by,
water at ordinary temperature; changed by boiling water to H3PO4 . Heated
to redness HPO- is formed. It forms two classes of salts: M'2H2P2O7 and
M'4P207 .
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., Formation. — Phosphoric aiilii/drifle, P2O5 , is formed by biirning phosphorus
in great excess of air: also by burning phosphorus in NO , NO2 , or CIO, .
Orthophasphnric acid, H3PO4 , is formed by long exposure of phosphorus to
moist air, or by oxidation with HNO3; by oxidation of H:,PO2 or H3PO3 with
the halogens, HN03 , HC1O3 , etc.: by treating P,,O., , HPO3 , or H4P2O7 with
boiling water; by combustion of PH3 in moist air: and by action of water on
PC15 . 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 («) can be verified, that is, pure phosphoric
acid can be separated; but the changes shown in equations (?J) and (e) do not
comprise the whole of the material taken. In the operation (6) some sodium
phosphate and some nitric acid will be left, and in (r) some trihydrogen
phosphate will no doubt be made.
a. CaH4(P01).J + H2C,04 = CaC.O4 + 2H3PO4
ft. Na.HPO, + 2HNO- = 2NaN03 + H3PO4
and Na2HPO, + HN03 = NaNO, + TSTaH2P04
c. 2CaHP04 + 2HC1 = CaCL + CaH4(P04)s
Metaphosphoric acid is formed by treating P20.-. 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.
Pyrophosphorie arid. H,P.O7 . is formed by the decomposition of lead pyro-
phosphate, Pb2P2O7 . with H,S or of the corresponding barium salt with
H2SO4; or by heating H3PO, to a little above 200° until no yellow silver
phosphate. Ag-PO, . is obtained on dissolving in water and treatment with
silver nitrate after neutralization with NH,OH .
300 PHOSPHORIC ACID. §255,4..
4. Preparation. — To prepare P205 , 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 P205 is distilled in an atmosphere of oxygen
(Shenstone, Walts' Die., 1894, IV, 141). H3P04 is prepared by warming
phosphorus, one part, with nitric acid, sp. (jr. 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
HN03 is removed by evaporation, water is added and the solution is sat-
urated with H2S to remove any arsenic that may be present. The solution
is then evaporated to a syrupy consistency at temperatures not above
150° (Krauthausen, Arch. Pliarm., 1877, 210, 410; Huskisson, B., 1884,
17, 161). Many or tho phosphates are formed by the action of H3P04 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 met a or pyrophosphates. Metaphosphaies 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., NaNH4HP04 . Pyro phosphates
are formed by double decomposition with Na4P207 ; by action of H4P207
on certain oxides or hydroxides; also by ignition of dibasic orthophos-
phates, e. g., Na2HP04 . Na2H2P207 may be prepared by titrating a sat-
urated solution of Na4P,,07 with HNO., until the solution gives a red color
with methyl orange. Upon standing the salt separates in large crystals
(Knorre, Z. angew., 1892, 639).
5. Solubilities. — 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, even by
C02 , forming the secondary salt. A number of non-alkali primary ortho-
phosphates are soluble in water, e. //., CaH4(P04)., . All normal and di-
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 riot
a phosphate but an oxide or an oxychloride.
All phosphates are dissolved or transposed by HNO,, HC1 , or H._,S04 ,
and all are dissolved by HC.-H.O., except those of Pb , Al and Fe'" . All
are soluble in H3P04 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 metal* are diflicultly solu-
ble in acetic acid. The most of the pyrophosphates of the heavy metals,
i}255, dA. PHOSPHORIC ACID. 301
except silver, are soluble in .solutions of alkali pyrophosphates, as double
pyrophosphates soluble in water (distinction from orthophosphates). Ferric
iron as a double pyrophosphate loses the characteristic properties of that
metal (Persoz, J. C., 1849, 1, 183). Phosphates are insoluble in alcohol.
<>. Reactions. — A. — With metals and their compounds. — Phosphoric acid dis-
solves some metals, c. y., Fe , Zn and Mg with evolution of hydrogen. It unites
with the oxides and hydroxides of the alkalis and alkaline earths and with
other freshly precipitated oxides and hydroxides except perhaps antimonous
oxide. It also decomposes all carbonates evolving- CO. . Phosphates are formed
in the above reactions, the composition of which depends upon the conditions
of the experiment.
Free orthophosphoric acid is not precipitated by ordinary salts of third,
fourth and fifth group metals (in instance of ferric chloride, a distinction from
pyrophosphoric acid and metaphosphoric acid),* but is precipitated in part by
silver nitrate, and lead nitrate and acetate. Ammoniacal solution of calcium
chloride or of barium chloride precipitates the normal phosphate.
Free metaphosphoric acid precipitates solutions of silver nitrate, lead nitrate,
and lead acetate, the precipitates being- insoluble in excess of metaphosphoric
acid, and soluble in moderately dilute nitric acid. Barium, calcium and ferrous
chlorides, and magnesium, aluminum, and ferrous sulphates, are not precipi-
tated by free metaphosphoric acid. Ferric chloride is precipitated, a distinc-
tion from orthophosphoric acid.
Free pyrophosphoric acid gives precipitates with solutions of silver nitrate,
lead nitrate or acetate, and ferric chloride; no precipitates with barium or
calcium chloride, or with magnesium or ferrous sulphate.
Orthophosphoric acid — or an orthophosphate with acetic acid — docs not coagu-
late egg albumen or gelatine. This is a distinction of both orthophosphoric
acid and pyrophosphoric acid from meta phosphoric acid.
With silver nitrate soluble orthophosphates form silver orthophosphate,
Ag3P04 , yellow; with rnetaphosphates, silver metaphosphate, AgPO;, ,
white; and with pyrophosphates, silver pyrophosphate, Ag4P207 , white,
all soluble in ammonium hydroxide. Silver metaphosphate is soluble in
excess of an alkali metaphosphate (distinction from pyrophosphates).
If a disodium or dipotassium orthophosphate is added to solution of silver
nitrate, free acid is formed, and an acid reaction to test-paper is induced (a).
Rut with a trisodium or tripotassium phosphate, the solution remains, neuiral
fft) — a means of diatinffviahinfl the acid phosphates from the normal.
(«) Na,HPO4 + :iAgNO, = Ag,P04 + 2NaNO3 + HNO,
(6) Na,P04 + 3AgNO, = Ag3POt + sNaNO,
Free orthophosphoric acid forms no precipitate with reagent silver nitrate.
With lead acetate or nitrate, Na2HP04 forms Pb3P04 , white, insoluble
in acetic acid, as are also the phosphates of aluminum and ferricum. With
* A solution containing 5 p. c. ferric chloride, mixed with one-fourth its volume of a 10 p. c
sohiti"n of orthophosphoric acid, requires that near half of the latter be neutralized (so that
phosphate is to phosphoric acid as 1.114 is to l.OCO) before precipitation occurs. On the other
hand, 4 cc. of a 5 p. c. solution of ferric chloride, mixed with 1 cc. of a 6 p. c. solution of meta-
phoaphoric acid, form a precipitate, to dissolve which, 20 cc. of the same metaphosphoric aci'i
solution ( r 5 cc. of a 24 p. c. solution of hydrochloric acid are required. Four cc. of a 5 p. c.
solution of stiver nitrate with I cc. of a 10 p. c. solution of orthophosphoric acid give a precipi-
tate, to dissolve which rpquiri s 7 cc. of the^sam^jartiinphosphoric acid solution. [The Author'*
report of work by Mr. Morur.m, A>n. Jour. P/in1-., 1876, !^, .r3!.. Kratschmor and S/.tankOvansky,
Z., 1883, 21, 520.)
302 PHOSPHORIC ACID. §255, fi/1 .
PbCl2 the precipitate always contains a chloride. Free phosphoric acid,
H3P04, forms an acid phosphate, PbHP04 (Heintz, Fogg., 1848, 73, 119).
Lead salts also form white precipitates with soluble pyro and metaphos-
phates ; the pyro salt, Pb.,P.,07 , is soluble in an excess of Na4P207 . Bis-
muth salts form BiP04 , insoluble in dilute HNO;, .
Solutions of orthophosphatos give, with soluble ferric, chromic, and
'aluminum salts, mostly the normal phosphate*, FeP04 , 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 is
added; and if the reaction is not markedly acid, it is made so by addition
of acetic acid. Ferric chloride (if not present) is now added, drop by
drop, avoiding an excess. The precipitate, ferric phosphate, is brownish-
white.
With zinc and manganous salts, the precipitate is dimetallic or normalr-
ZnHP04 , or Zn,(P04)., — according to the conditions of precipitation.
When a manganic compound is mixed with 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 redd-i.sh
normal phosphate.
Soluble salts of the alkaline earth metals, with dimetallic alkali phos-
phates, as Na,HP04 , form white precipitates of phosphates, two-third*
metallic, us CaHP04 : with trimetallic alkali phosphates, white precipitates
of phosphates, normal or full metallic, as Ca.,(P04)., . The precipitates an-
soluble in acetic acid, and in the stronger acids. Concerning the am-
monium magnesium phosphate, see £189, 6(1.
Magnesium salts with ammonium hydroxide give a precipitate of double
pi/rophosphatr, soluble in alkali pyrophosphate solution.
Magnesium salts with ammonium hydroxide are not precipitated by
soluble meta phosphates unless very concentrated.
Ammonium molybdate, in its nitric acid solution (§75, ti<l), furnishes an
exceedingly delicate test for phosphoric acid, giving the pale yellow pre-
cipitate, termed ammonium phosphoinolyldate. 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 pyrophosphato
solutions— and not until orthophosphoric acid is formed by digestion with
the nitric acid of the reagent solution.
§255, 8. PHOSPHORIC ACID.
B. — With non-metals and their compounds. — Phosphoric acid is not
reduced by any of the reducing acids. Phosphates of the first two groups
are transposed by H2S , 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.
HC1 , HN03 , and H.,S04 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, Pb , 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
sulphates upon evaporation and long-continued heating on the sand bath.
7. Ignition with metallic magnesium (or sodium) reduces phosphorus from
phosphates to magnesium plwsphidc, P2Mg3 , recognized by odor of PHS ,
formed on contact of the phosphide with water. A bit of magnesium wire (or
<of sodium) is covered with the previously ignited and powdered substance in
a 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 with water, gives the odor of phos-
•phorus hydride.
Orthophosphoric acid, heated to 213° forms pyrophosphoric acid; when heated
to dull redness the meta acid is obtained, which sublimes upon further heating1
•without change. Phosphoric anhydride, PL,Or) , 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. Bimetallic orthophosphates,
M'.,HPO4 , are changed to normal pyrophosphates upon ignition; also tribasic
orthophosphates when one-third of the base is volatile, e. (/.. MgNH4PO4. .
Mono-metallic or primary orthophosphates, M'H2P04 , become metaphosphates;
also secondary or tertiary orthophosphates when onty one atom of hydrogen
is displaced by a metal whose oxide is non-volatile, e. (/., NaWTH4HP04 .
Acid pyrophosphates, M'2H2PoO7 , form metaphosphates. When meta or pyro-
phosphates are fused with an excess of a non-volatile oxide, hydroxide or
carbonate the tertiary orthophosphnte is formed (Watts'. 1S94. IV. 106).
Phosphates of Al , Cr , Fe , Cu , Co . Ni . Mn , Gl and "0" when heated to a
white heat with an alkali sulphate form oxides of the metals and an alkali
tribasic orthophosphate; phosphates of Ba , Sr , Ca . Mg- , Zn and Cd form
double phosphates, partial transposition taking place (Derome. C. r.. 1879, 89,
952: Grandeau, A. Ch., 1886, (6), 8, 19::).
8. Detection. — The presence of orthophosphorie acid in neutral or acid
solutions- is detected by the use of an excess of an a mm on him molybdate
solution (§75, 6d). With pyro and mctaphosphoric acids no reaction is
obtained except as they are changed to the ortho acid by the reagents
used. Disodium phosphate, Na.,HP04 , after precipitation with silver
nitrate, reacts acid to test papers. With trisodium phosphate the polu-
tion is neutral (distinction). Ortho-phosphates are distinguished from
pyro and metaphosphates by the color of the precipitate with silver nitrate:
Ag,P04 is yellow, Ag4P,07 and AgPO, ;ire white. Also by the ('act that
only the ortho acid is precipitated by ammonium molybdate. Nearly all
pyrophosphates are soluble in sodium pyrophosphate, Na4P,,07 (distinc-
304 nrL/'iirK. $255, !).
lion from orthophosphates). Ilager (J . C., ]Sl:>, 26, 940) gives a method
for detecting the presence of H..PO., , H AsO , or HNO in H,P04 . Sodium
metaphosphate docs not give a precipitate with ZnS04 cold and in excess;
with Na4P,07 and Na,H,P,07 a white precipitate of Zn.^P.^0. is obtained
(Knorre, Z. angew., 1892, r,;]9).
9. Estimation. — (a) By precipitation as magnesium ammuniuiu pliospbate,
MgNH,PO, , and ignition to the pyrophosphate. (6) By precipitation and
weighing as lead phosphate, Pb., (PO4), . (r) By precipitation from neutral or
acid solution by ammonium molybdate and after drying at HO0 weighing as '
ammonium phosphomolybdate. Consult .lanovsky (J. C., 1873, 26, 91) for a
review of all the old methods.
§256. Sulphur. S = 32.07. Usual valence two, four and six (§14).
1. Properties. — Sulphur is a solid, in yellow, brittle, friable masses (from
melting); or in yellowish, gritty powder (from sublimation) or in nearly whitei,
slightly cohering, finely crystalline powder (by invcinituti<ni from its corn-
pounds). At — 50° it is white (Schoenbein, •/. /)/'., 1852, 55, Mil). The xiwvific
yrnntu of native sulphur is 2.o:!4.s (Pisati, R., 1874, 7, 'Ml). .Melting point, 111°
(Quincke, J., 18(58, 21). Itiiiling point, 444..V10 (Callendar and Criflitlis, C. \., 18tfl,
63, 2). Vapor dcnxitii at ll(iO° is :M, indicating that the molecule is S, (liineau,
C. r., 1859, 49, 799); but at lower temperatures the molecule seems to vary from
S2 to S, . Sulphur is polymorphous, existing in various vrystalline forms,
rhombic, monoclinie and trielinie 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, 1(5.000; S, "2.07: 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 180° it is nearly solid,
so that the dish may be inverted without spilling. At 2<>0° it again becomes a
liquid as at first; and at 444.5,'t0 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 ISO0 and thin again at 111°, and al lower
temperatures solid. If, at a temperature mar 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. Hut 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 24S° (Hill, C. N., 1890, 61. 125): burning in air or oxygen
with a pale blue flame and pent trating odor to SO. .
The isolated oxides of sulphur are SO. . SO:1 , S,O:, and S,O- . Sulphur and
oxygen combine directly to form SO. and SO:1: the former by burning sulphur
in oxygen, the latter by the action of o/one upon SO ; also by burning sulphur
with oxygen under several atmospheres pressure. S..03 is made by dissolving
sulphur' in sulphur dioxide: S,O7 by the action of the electric discharge upon
a. mixture of SO., and O .
2. Occurrence.— (</) Found in a free state, and as SO, in volcanic districts.
(b) As H.S in some mineral springs, (r) As a sulphide: iron pyrites, FeS, :
copper pyrites, CuFeS. ; orpiment. As.S;: realgar, As,S,: /ine blende. ZnS;
cinnabar,' HgS; galena," PbS. ('/) As a sulphate: gypsum. CaSO,.2H,O; heavy
spar, BaSO,: kieserite. MgSO,,H"..O: bitter spar (Epsom sails). MgSO^TH.O;
(Jlauber salt. Na,S04.1<)H,O . etc.
3. Formation. — (a) By decomposing polysulphides with HC1 (Schmidt, I'hitr-
maceiitlscJic Clicmir. IS'.IN. 175). (ft) Bv adding an acid to a solution of a thio-
Kiilphate. (c) By the reaction between SO, and H,S: 2S02 + 4H.S = :!S, -i
4H..O . (d) By the decomposition of metallic sulphides with nitric acid: 'JBijS,
+ 16HNO., = JBi(NO3)3 + :iS2 -f 4NO + SH.O .
£256, (W8. SULPHUR. ;505
4. Preparation. — (a) The native sulphur is separated from the clay and rock
IB which it is embedded, partly by melting- and partly by distillation, (ft)
From FeS2 by heating- in close cylinders 3FeS2 = Fe3S4 + S2; or at a higher
temperature: 2FeS2 = 2FeS + S. . Much of the sulphur contained in pyrites
is converted into and utilized as sulphuric acid.
5. Solubilities. — Ordinary (not precipitated) sulphur is soluble in carbon di-
sulphide; the ductile variety is insoluble. There are several allotropic forms
of sulphur. Samples of commercial sulphur are almost never found which are
entirely soluble or insoluble in carbon disulphide. Forms of sulphur insoluble
in CS2 are changed to soluble forms upon heating- to the melting point;
also amorphous sulphur insoluble .in CS~ (formed by adding acids to thiosul-
phates or SO, to H,S) is changed to the soluble form by mixing with a solution
of H2S in water. It dissolves readily in hot solutions of the hydroxides of
potassium, sodium, calcium or barium, forming polysulphides and thiosul-
phates: 3Ca(OH)2 + CS2 — 2CaS5 + CaS2O3 + .°,H2O . These can be separated
by alcohol, in which the sulphides dissolve. These products are also readily
decomposed by acids with separation of sulphur (method of preparation of
precipitated sulphur).
Precipitated sulphur (in analysis, HC1 upon (NH4)2SX) is soluble in benzol or
low boiling petroleum ether; of value in analysis for the removal of the sulphur
to detect the presence of traces of As or Sb sulphides (Fresenius, Z., 1894, 33,
573).
6. Reactions. A. — With metals and their compounds.. — Sulphur does
not combine with metals without the aid of heat (see 7), except that under
very great pressure (6500 atmospheres) it combines with Pb , Sn , Sb , Bt,
Cu , Cd , ye , Zn , and Mg (Spring, B., 1883, 16, 999).
Flowers of sulphur boiled with SnCl2 gives SnS and SnCl4 ; with HgNO;t
almost exactly one-half of the mercury is precipitated as HgS . No action
with sulphates of Cd , Fe", Mn", Ni and Zn ; with acid solutions of SbCL,
and BiCl3 ; or with solutions of Asv and As'" (Vortmann and Padberg,
B., 1889, 22, 2642). Sulphur boiled with hydroxides of K , Na , NH4 , Ba ,
Ca, Sr, Mg, Co, Ni, Mn, Hg", Bi, Cu', Cu", Cd, Pb, Ag, and Hg'
forms sulphides and thiosulphates; also some sulphates are formed. No
action with hydroxides of Fe, Zn and Sn (Senderens, BL, 1891, (3), 6,
800).
B. — With non-metals and their compounds.
1. HCN warmed with sulphur or a polysulphide becomes a thiocyanate:
2KCN + S2 = 2KCNS or 4HCN + 2(NH4)2S4 == 4NH4CNS + 2H2S -f S2 .
2. HN03 becomes NO' and H2S04 . Strong acid and long continued
boiling are necessary to the complete oxidation of the sulphur. The
crystallized variety is attacked with much greater difficulty than the
amorphous or flowers (Saint-Gilles, A. Ch., 1858, (3), 54, 49).
8. Red phosphorus combines readily at ordinary temperature, forming
P2S3 or P2S5 , depending upon the relative amounts of the elements used.
Ordinary phosphorus combines explosively. See §252, 6. Tribasic sodium
or potassium phosphate when boiled with sulphur forms alkali polysul-
phide and thiosulphate, changing the phosphate to dibasic phosphate
•(Filhol and Senderens, C. r., 1883, 96, 1051).
306 HYDROXULI'HURIC ACID. ><256, (\Ftl+.
' 4- H2S04 , concentrated and hot, becomes S02 from both the S and the
H2S04 : 4H,S04 + S2 = = GSO, -f- 4H20 . S03 when added to S at 12°
forms the blue, hyposulphurous anhydride, S20;< (not the anhydride of
thiosulphuric acid, S.,0.,). S02 reacts with .S even at ordinary tempera-
tures, forming thiosulphuric acid and tri or tetrathionie acid (Colcfax\
J. (7., 1892, 61, 109).
5. Cl in presence of water forms HC1 and ILS04 . HC10n becomes HC1
and H,S04 .
6. Br in presence of water becomes HBr and H2S04 . HBr03 become*
HBr and H,S04 .
7. Sulphur does not appear to have any action upon iodine or upon
iodine compounds.
7. Ignition.— In the air, at ordinary temperatures, finely divided sulphur is
very slightly oxidized, by 0/0110, to xitlphuric acid; at 248° it begins to oxidi/e
rapidly to tftilptninnix iinlii/dridc, burning with a blue flame.
Sulphur, when fused with the following elements, combines with them to
form sulphides: Pb . Ag , Hg . Sn . As. Sb , Bi , Cu , Cd , Zn , Co, Ni , Fe .
Sr , Ca , Mg , K . Na . In . Tl , Pt . Pd , Bh , Ir , Li , Ce , La , Ne , Pr .
Svi— n becomes Svi when fused with alkaline carbonate and nitrate or chlorate.
That is, free sulphur, S° , or any compound containing sulphur with valence
less than six, is oxidi/ed to a sulphate if fused with an alkaline nitrate or
chlorate, nitric oxide or a chloride being formed and carbon dioxide escaping.
8. Detection. — (a) By burning in the air to a gas having the odor of
burning matches, (ft) By its solubility in CS2 . (c) By formation of
H.,S04 with oxidizing agents, (rf) By the formation of sulphides upon
fusion with metals, (c) By the blackening of silver coin after boiling
with alkali hydroxide, (f) Formation of reddish-purple with sodium
nitroferri cyanide after boiling with alkali hydroxide, (g) In organic
compounds by heating with Na and testing the Na,S with sodium nitro-
ferricyanide (Vohl, K., 1870. 9, 875).
9. Estimation.— Sulphur is usually estimated by oxidation to a sul-
phate and weighing as BaSO, .
§257. Hydrosulphuric acid. H2S = 34.086.
H',8-", H — S — H.
1. Properties. — Molecular iceiyht, 34.0SG. Vapor density^ 17. JMliny point,
— 61.8°. Freezing point, — 85.56°. Under a pressure of 14.6 atmospheres it be-
comes a liquid at 11.11° (Faraday, A., 1845, 56, 156). It is a colorless poisonous
gas. It burns readily, forming sulphur dioxide and water: 2H2S + :iO2 = 2SO,
+ 2H2O . The aqueous solution slowly decomposes upon exposure to the air
with separation of sulphur. The gas is readily expelled from its anurous
solution by boiling; slowly when exposed at ordinary temperature. I'.olli tin-
gas and the water solutions have a feebly acid reaction towards moist litmus
paper. They also possess a strong characteristic odor, resembling that of
rotten eggs. In acid or in alkaline solutions it is a strong reducing agent.
See 6.
§257,4. HYD&OSitLPBURlO .4.CID. 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 Acid. — («) By direct union of the elements
when passed over pumice stone heated to 400° (Corenwinder, A. Cli., 1852, (3),
33:, 77). (6) Heating paraffin or tallow with sulphur (Fletcher, C. N., 187'J, 40,
154); and by passing illuminating gas through boiling sulphur (Taylor, C. N.,
1883, 47, 145). (c) The sulphur in coal becomes H2S in the process of gas-
making, (d) From steam and sulphur at 440°. (c) Often occurs in nature from
reduction of gypsum by decaying organic matter (Myers, J. pr., 1869, 108, 123}.
(f) Transposition of sulphides by hydracids or by dilute phosphoric or dilute
sulphuric acid. (.(/) Decomposition of organic compounds containing sulphur.
Formation of Sulphides. — (1) 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. (3) Action of H S on metallic oxides or
hydroxides. Those sulphides which are decomposed by water (c. </., A1...S,, ,
Cr2S3) are not formed in its presence, but by action of IL>S upon the oxide at
a red heat. (4) 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.
|5) 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 BaSO4 , 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, 26). (8) By the action of
zinc amalgam on sulphuric acid (Walz, C. N., 1871 23, 245). (9) As a reagent
for the formation of metallic sulphides in analysis it is recommended by
Schiff and Tarugi (#., 1894, 27, 3437), Schiff (#., 'l895, 28, 1204), and Tarugi
tGaszetta, 1895, 25, i, 269), to use ammonium thioacetate, CH3COSNH4; 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, dihiting 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 + S2 — 2FeS
2. 2Ag + H2S = Ag2S + H2
3. Pb(OH)2 + H,S — PbS + 2BLO
4Fe(OH)3 + 6H2S = 4FeS + S2 4- 12H2O
!t. 4FeCls + 6(NH4)2S = 4FeS + S2 + 12NH4C1
5. 2CaO + CS2 — 2CaS 4- CO2
6. 4CaO -f 3S2 = 4CaS + 2SO2
7. K2SO4 + 2C = K2S 4- 2C02
4. Preparation. — For laboratory purposes it is nearly always made by
adding H2S04 or HC1 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 H2S04 should be used: FeS + H2S04 == FeS04 + H2S .
Concentrated H2S04 has no action on FeS , unless heated and then SO, is
evolved: 2FeS + 10H,S04 = Fe,(S04)3 -f OS02 + 10H,0 ; and frequently
free sulphur is formed by the action of the H2S upon the S02 first formed.
The colorless ammonium sulphide, (NH4)2S , is prepared by saturating
H7DROSULPHURIC ACID. §257,5.
ammonium hydroxide with H,S until a sample will no longer give a pre-
cipitate with a solution of magnesium sulphate; showing that ammonium
hydroxide is no longer present. Upon standing the solution gradually
becomes yellow with formation of the poll/sulphides or yellow ammonium
sulphide, (NH4),SX This may he hastened by the addition of sulphur
(Bloxam, J. C., 1895, 67, 277).
Sodium sulphide, Na.,S, is prepared by neutralizing an alcoholic solution
• »f NaOH with H,S and thcii adding an equal amount of NaOH and allowing
to crystallize; air being excluded. The various polysulphides, Na,S, to
Na,S5 , are prepared by boiling the normal sulphide with the calculated
amounts of sulphur (Boettger, .1., 1884, 223, 335; Geuther, A., 1884, 224,
201).
5. Solubilities. — At 15° water dissolves 2.GG volumes of the gas H2S\
Sulphides which dissolve in dilute H,S04 evolve H2S , e. y., CdS , FeS ,
MnS , ZnS , etc. But if a sulphide requires concentrated H2S04 for its
solution; S and SO, are formed or SO, alone; e. y., Bi2S.t , CuS , HgS . If
concentrated H,S04 be used upon a sulphide that might have been dis-
solved in the dilute acid, then no H,S is evolved: ZnS -f 4H,S04 = ZnS04
-|- 4SO, -+- 4H,0 . Or with a small amount of water present: 2ZnS -j-
4H,S04 == 2ZnS04 + S, + 2S02 + 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 HNO , dilute or concen-
trated, readily soluble in chlorine, nitrohydrochloric acid, or chloric acid
if hot. Most other sulphides are soluble in hot HNO, (§74, 6e). Long
rontinued 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 . Ft . Mo . Cu . Cd , and Zn (Clermont and Frommel, A. Ch.,
1879, (5), 18, 203).
As a reagent, hydrosulphuric acid, gas or solution in water finds ex-
tended application in the analytical laboratory. The grouping of the
bases for analysis depends verj' largely upon the relative solubilities of the
sulphides. Hydrosulphuric acid in alkaline solution, alkali sulphide or
jiolysulphide, 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 alkaline
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 aeid may form arscnous sulphide and free sulphur.
Ferric solutions are reduced to ferrous with liberation of sulphur. In acid
mixture other third and fourth group sails are not disturbed, but from
§257, 6.15. HYDRORl'LPHURW ACID. 309
solutions of their normal salts traces of cobalt, nickel, manganese, and
zinc (§135, 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, flesli colored; those of iron, copper, and lead, black; arsenic
stannic and cadmium, yellow; antimony, orange-red; stannous, brown; mer-
cury, successively white, yellow, orange, and black.
6. Reactions. A. — With metals and their compounds. —Some metals
are converted into sulphides on being treated with hydrosulphuric acid;
0." g., Ag , Cu , Hg , etc. The alkali polysulphides slowly attack many
metals with formation of sulphides: Sn becomes M'2SnS3 ; Ag becomes
Ag2S, no action with colorless (NH4)2S ; Hi forms NiS ; Fe , FeS ; Cu,
CuS and then Cu2S (with colorless ammonium sulphide, (NH4)2S, Cu2S
is formed with evolution of hydrogen) (Priwozink, A., 1872, 164, 46).
The hydroxides or non-ignited oxides of Pb", Ag , Hg", Sb , Sn , Bi"',
Cu, Cd, Fe", Co", Ni", Mn", Zn , Ba, Sr , Ca, Mg, K, Na, and 1TH,
unite with moist H2S 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"+n becomes PbS and S .
2. Asv in acid solution forms some As2S3 and S . See §69, 6e.
3. Hg' becomes HgS and Hg .
4. CrVI becomes Cr'" and S , if the H2S be in excess : 2K2Cr20T -f 8H2S
= 4Cr(OH)3 + 3S2 + 2K2S + 2H20 .
5. Fe'" becomes Fe" and S : 4FeCl3 + 2H2S = 4FeCl2 -f 4HC1 + S2 .
If the solution be alkaline FeS is precipitated : 4FeCl3 + 6K..S = 4FeS -f
12KC1 + S2 .
6. Co"+n becomes Co" and S .
7. Ni"+n becomes Ni" and S .
8. Mn"+n becomes Mn" and S. In alkaline solution with excess of
KMn04 , an alkali sulphate is formed and Mn02 : 8KMn04 -f- 3KJ3 =
3K2S04 + 4K20 -f 8Mn02 (Schlagdenhafen, Bl, 1874, (2), 22, 16).
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 (Bfn2S, KaS), precipitate solutions of calcium and mag-
nesium salts as the hydroxides : CatCgHjO,', + 2\n,S + 2H,O = Ca(OH)., + 2SaC2HsO., +
2NatlS. No reaction with the acid fixed alkali sulphides (NaHS, KHS) or with ammonium
sulphides (Pelou/.e, A. Ch., 1866, (4), 7, 172).
310 &7DROSULPHUR1C ACID. J<257, ()R.
formation of an alkali hydroxide; except that the arsenic will remain in
solution (§69, or) and the chromium will be precipitated as the hydroxide.
Dry H2S 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 PdCL 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 Ag,S 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, 249, 326).
B. — With non-metals and their compounds.
/. H:!Fe(CN)(! becomes H4Fe(CN)t: and S. Proof: Boil to expel the
excess of hydrosulphiiric acid, then add ferric chloride (§126, 6&).
2. HNO;i becomes NO and S . If the HNO:t be hot and concentrated the
sulphur is oxidized to sulphuric acid.
S. H.S has no reducing action on the acids of phosphorus.
Jt. H..SO., becomes pentathionic acid, H2S-0U , and sulphur: 10H2SO, -4-
10H,S =: ?H,S,0, + 5S, + 1SH...O. With excess of H,S the product is
entirely free sulphur from both compounds: 2H,S03 + 4H.S = 3S2 -f-
()H,0 (Debus. J. C., 1888, 53, 282).
H,S04 . dilute no action; concentrated and hot, S and S02 are formed:
2H,S04 + ->H,S = S, + 2SO, + 4H,0 (§256, (\B4).
5. Cl with ELS in excess forms HC1 and S ; with Cl in excess forms HC1
and H ,S04 .
HC10, with H,S in excess forms HC1 and S ; with HC103 in excess HC1
and H,SO, .
6. Br with ELS in excess forms HBr and S ; with Br in excess HBr and
H,S04
HBrO., with H2S in excess forms HBr and S ; with HBrO:! in excess HBr
and H,S04 .
7. I becomes HI and S (Filhol and Mellies, A. Ch., 1871, (4), 22, 58).
HIO, becomes HI and S .
7. Ignition. — Dry hydrosnlphnric acid gas is not decomposed when heated to
350° to 360°. At this temperature AsH in presence of potassium polysulphide,
K«S3 , liw of .sM//>7<«/\ is decomposed: 2AsH, + 3K2S, = '_>K3AsS3 + 3H2S;
thus furnishing a ready means of purifying H2S for toxicological work (§69,
G'ft) (Pfordten, B., 1884, 17, 2897).
If air be excluded some sulphides may be sublimed unchanged; e. g., HgS ,
As,S3 , As,S5 , Sb2S, , etc. In some cases part of the sulphur is separated,
leaving a sulphide of a lower metallic valence: 2FeS2 = 2FeS + S2 . Some
sulphides remain unchanged upon ignition in absence of air; e. g., FeS , MnS .
CdS , etc. All sulphides suffer some change on being ignited in the air; some
Slowly, others rapidly; Sb2S, , CuS , A12SS , Cr,Sa , etc., evolve SO2 and leave
£257, 9. HYDROSULPHURIC ACID. 311
the oxide of the metal; HgS , A.g,S , etc., evolve SO, and leave the free metal.
All sulphides, as well as all other compounds of sulphur, when fused with KNO:j
or KC1O3 in presence of an alkali carbonate are oxidized to an alkali sulphate;
forming NO or KC1 and evolving- CO2 . The metal is changed to the carbonate,
oxide or the free metal (§228, 7).
, When ignited on charcoal with sodium carbonate — or (distinction from
sulphates) if ignited in a porcelain crucible with sodium carbonate — soluble sodium
sulphides are obtained. The production of the sodium sulphide is proved by the
black stain of Ag2S , 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.
(6) 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 (G-). (c) The gas blackens silver nitrate solution, delicate but
PH., , AsH3 , and SbH3 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 §256, 8. KMn04 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 H2S04 ;
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 sulphuric acid
with zinc. The gas, with its excess of hydrogen, may be tested by method
(/). (/) Sodium nitroferricyanide gives a very delicate and characteristic
test for H2S 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 H2S , dilute, remains colorless ; a
concentrated solution gives a blue color, due to the reducing action of
the H2S on the ferricyanidc.
For method of separation of the various sulphur compounds from each
other consult Kynaston (J. C., 1859, 11, 166) and Bloxam (C. N., 1895,
72, 63).
9. Estimation. — Sulphides are usually oxidized to H.,S04 (by chlorine,
bromine, or nitrohydrochloric acid, or by fusion with KN03 and Na2C03)
precipitated with BaCl2 and weighed as BaSO , .
312 THIOSULPHVRIC ACID. §258, 1.
§258. Thiosulphuric acid. H,S,0, = 114.156 .
Di tli lotions acid.
0
II
H'2(S2)IV0-"3 ,H — 0 — S — S — H.*
II
0
1. Properties. — Thiosulphuric acid, H S O (formerly called hyposulphurous
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, Na.S.O, ]
soon beginning to decompose into H2SO3 and S (Landolt, B., 1883, 16, 2985)!
The thiosulphates are not particularly stable compounds, some decomposing
almost immediately upon forming; c. </., mercury thiosulphates. Alkali thio-
sulphates decompose upon heating into sulphate and polysulphide: iNa.S.O. =
3Na2SO4 -f Na-S., . Other salts give also S and H2S . Boiling solution of a
thiosulphate gives a sulphate and H2S or a sulphide of the metal.
2. Occurrence. — Not found in nature.
3. Formation. — Thiosulphates are formed by the oxidation of alkali or
alkaline earth polysulphides by exposure to the air or by SO2 or K,Cr,OT-
2C&S, + 30, = 2CaS,03 + :iS,; 4Na,S., + <iSO2 = 4Na2S26s + yS2: 2k,Ss +
4K2Cr2O7 + i:;H,O == :,K,S,Oa + sCr(OH), + :»KOH (l)oepping, A., 1843, 46,
172; Gueront, C. r., 1HT2, 75, 1276). Also by heating ammonium sulphate with
phosphorus pentasulphide (Spring, K., 1874, 7, ll.">7).
4. Preparation. — Thiosulphates are prepared by boiling sulphur in a solu-
tion of normal alkali sulphite: L'Na.SO, + S2 = L'Na..,S;.Os . Fixed alkali or
alkaline earth hydroxides with sulphur also form thiosulphates: 3Ca(OH), +
(iS2 = -'CaS., +'CaS.,O, + ::H2O (Filhol and Senderens, C. r., 1883, 96, 839;
Senderens, f. r., 1887, 104, 58). Cnmrnrrcml sodium thiosulphate is prepared
by passing SO, into " sod.a 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 aleoliol. They are decomposed, but not fully dis-
solved, by acids, the decomposition leaving a residue of sulphur.
Alkali thiosulphate solutions dissolve the thiosulphatee 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,S,O3 + Na S O, = 2NaAg-S 0
AgCl + Na2SzOs = NaAgS,03 + NaCl
PbSO, + :5Na,S20, = Na4Pb(S,03)3 + Na,S04
6. Keactions. — 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: Ag,SjO,
+ H2O = Ag.S -f H,,SO, . Soluble mercury salts with sodium thiosulphate.
form a white precipitate, almost instantly turning black with decomposition to
mercuric sulphide. Na,S,03 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
* Dunte, B., 1874, 7. 646.
§258,8. THIOSULPHURIC ACID. 313
Na-.S.O;, as sulphides, As,S:. and Sb2S3 (a separation from tin,* which is not
precipitated) (&e, §§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: 2FeCl3 -f- 2N"a,S2O3 = 2FeCl2 -f 2NaCl -f- Na2S4O6; used as n
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.
Permanganates in neutral solutioa 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, BaS203 ,
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: 2Na2S2O3 + 4HC1 = 4NaCl + 2H2S03 + S2 .
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: 2Na2S203 + I, = 2NaI + Na,S400 . If chlorine
or bromine be in excess the tetrathionate is further oxidized to a sulphate:
Na2S2O3 + 4CL + 5H,O = Na2SO4 + H2SO4 -f 8HC1 . Chloric, bromic and
iodic acids are first reduced to the corresponding halogens and then with an
excess of the thiosulphate to the haliclcs, 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 polysulphides (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 with thiosulphates
explode violently in the mortar. Cyanides and ferricyanides, fused with thio-
sulphates, form thiocyauates, which may be dissolved by alcohol from other
products. By fusion on charcoal with Na.2CO3 , thiosulphates form sulphides
(6) and (c) ; and by fusion with an alkali carbonate and nitrate or chlorate,
a sulphate is formed (rf). By ignition of a metallic salt with Na2S2O3 in a
dry test-tube the characteristic colored sulphide of the metal is obtained
(Landauer, B., 1872, 5, 406).
(c) 4Na2S2O3 = Na2S5 + 3Na2SO4
(6) Na2S203 -j- Na.,CO3 + 2C = 2Na2S + 3CO2
(c) 2PbS2O3 + 4Na2C03 -f- 5C = 4Na2S + 2Pb + 9CO2
(d) 3Na2S2O, + 3Na2C03 + 4KC103 = 6Na2S04 + 4KC1 + 3CO2
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 with evolution of sulphurous anhydride, 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
rulphur.
•According to Vortmann (M., 1886, 7, 418) sodium thiosulphate may be used instead of hydro
sulphuric acid in the second group of bases. An excess of the reagent is to be avoided and
nitric acid should be absent.
314 HTPOSULPHUROU8 ACID—DITIUONIC ACID. §258,9.
In the presence of a sulphate and sulphite the thiosulphate is detected as
follows: Add Bad, and NH4C1 in excess, then HC1 to solution of all but the
BaSO, . Filter and treat the filtrate with iodine, forming BaSO, of the sulphite
and BaS,O,, of the thiosulphate. Filter and add bromine to the filtrate, which
then forms BaSO, (Smith, ('. A., 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,
5259. Hyposulphurous acid. H,SO, = 6G.086 .
(Hydrosulphurous or difhionotu acid.)
H',S"0-"2 , H — 0 — S — H .
Obtained by Schiitzeuberger (C. r., 1809, 69, 19(5) by the action of zinc on
sulphurous arid: Zn + 2SO, + H=0 — ZnSO + H.SO, . The sodium salt is
formed bv treating a concentrated solution of sodium arid sulphite with zinc
tilings: Zn -f :;NaHSO:, = ZnSO, + Na.,SO3 -f NaHSO, + H.O . In the forma-
tion of the free aeid or of the sodium salt no hydrogen is evolved. It is a very
unstable compound, a strong1 reducing1 agent, rapidly absorbs oxygen from the
air, becoming sulphurous acid or a sulphite. According to Bernthsen (B., 1881,
14. 4.'!S) the sodium salt does not contain hydrogen. He gives the formula as
Na,S204: Zn + 4NaHSO:, = ZnSOs -f Na,SOu + Na,S,04 + 2H2O . It is used
in the preparing1 of indigo white for the printing- of cotton fabrics. See also
Dupre, J. C\, 1867, 20, 291.
§260. Dithionic acid. H,S20,, = 162.156 .
0 0
II II
H'0(S.,)X0-"0 , H — 0 — S — S — 0 — H.
II II
0 0
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 2H2SO3 = MnS,On + 2H2O . Similar
results are obtained with nickelic or ferric oxides (Spring and Bourgeois, III..
I^M'I. 46, l.")l). The acid is obtained by treating- the manganous salt with
Ba(OH): and the filtrate from this with the calculated amount of HjSO, .
li is a colorless solution and may be evaporated in a vacuum until it lias ;\
specific gravity of 1.347. It decomposes upon further heating: H...S2O0 = H,SO,
-f S02 . AH other tJiionic cnniitoitntlx ilivoiiiixixr ni>on ln-iitin<i irilh xciHtrnttnn of
.s»/////»r. By exposure to the air dithionie acid is oxidi/.ed to sulphuric acid.
All dithionatcs arc soluble in water and may be purified by evaporation and
crystallization (Gelis. .1. C)i.. lsc,2, (.'!), 65, 230).
Dithionic acid is also prepared bv carefully adding a potassium iodide solu-
tion of iodine to sodium acid sulphite (Hoist and Otto, Arch. Pharm.. isoi, 229,
171); Spring and Bourgeois (Arcli. Pharm., 1891, 229, 707) contradict the above
statement.
§262. TRirnroxic ACT n— TETRATHIONIC ACID. 315
§261. Trithionic acid. H,S,06 = 194.226 .
0 0
II II
H'2(S,)100-26 , H — 0 — S — S — S — 0 — H.
The free acid and anhydride are not known. The potassium salt is prepared
by boiling potassium acid-sulphite with sulphur (a); 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,
7, 1157):
(a) 12KHSO3 + S2 = 4K2S,00 + 2K2SO:! + 6H2O
(ft) 4K2S2O3 + GSO2 = 4K,S3Oti + S,
(c) Na.,S03 + Nax83O, + I, = Na2S300 + 2NaI
The acid is prepared by adding perchloric or fluosilicic acid to the potassium
salt. The a.cid is quite unstable; at low temperature in a vacuum it decom-
poses into SO2 , S and H2SO4 . 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 sodium amalgam change the trithionate to sulphite and thio-
sulphate (Spring, /. c.).
§262. Tetrathionic acid. H2S4Ofi = 226.296 .
0 0
' II II
H'2(S4)100-2f) , H — 0 — S — S — S — S — 0 — 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 solttion it can be boiled without decomposition.
The concentrated .solution decomposes into H_,S04 , SOU and S .
Tetrathionates art- prepared by adding iodine to the thiosulphates: :.'BaS,0 +
I2 = BaS4O(i + Bal, (Maumeno, C. r., 1879, 89, 422). The lead salt is obtained
by the oxidation of lead thiosulphate by lead peroxide in presence of sulphuric
acid: 2PbS2O, f PbO, + 2H,S04 — PbS4On + 2PbSO4 + 2H20 (Chancel and
Diacon. ./. j>r.. isiiii. 90. .">.">). 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 used to form the tetrathionate from the thiosulphate
(Fordos and Gelis, C. r., 1842, 15, 920). Sodium amalgam reconverts the tetra-
thionate into the thiosulphate: Na.S.O,, + 2Na = 2Na2S2O3 (Lewes, J. C., 1880,
39, 68; 1881, 41, 300). Tetrathionic acid is also formed with pentathionic acid
in the reactions between solutions of H.S and SCX (Wackenroder's solution,
A., 1846, 60, 189). See also Curtius and Henkel (./. pr., 1888, (2), 37, 137). The
acid gives no precipitate of sulphur when treated with potassium hydroxide
(distinction from pentathionic acid).
816 PENTATH IONIC ACID. $263.
§263. Pentathionic acid. H2S506 = 258.366 .
0 0
II II
H'2(S5)100-% , H — 0 — S — S — S — S — S — 0 — H.
Only known in the salts and in the solution of the acid in water. It is formed
by the action of H2S upon SO2 in the presence of water (a); by the action of
water on sulphur chloride (6); by the decomposition of lead thiosulphate with
HjS (Persoz, Pogg., 18(55, 124, 257):
or. lOHoSO, + 10H.S = 2H2S5O0 + 5S2 + 18H2O
b. 10S2C12 + 12H2O = 2H2S50S + 5S2 + 20HC1
The filtrate from the decomposition of SO2 by H,S is known as Wackenroder's
solution (Arch. I'harm., 1826, 48, 140). It has been shown to contain the tri
and tetrathionic acids in addition to the pentathionic acid (Debus, C. N., 1888,
57, 87). Pentathionic acid may be concentrated in a vacuum until it has a
specific gravity of !.<>; farther concentration or boiling1 heat alone decomposes
it into H SO. , S02 and S . The solution of the acid does not bleach indigo.
When treated with a fixed alkali hydroxide an immediate precipitate of sulphur
is obtained (distinction from H]S406): 4»H,S5O, + 20NaOH = <>NajSO, -+-
4Na2S2O., + :!S, + 14H2O (Takamatsu and Smith, J. C., 1880, 37, 592); or if the
NaOH be added short of neutralization: lOH^O, + 20NaOH = 10Na2S4O? -f
&8, + 20H20 . Neutralization of pentathionic acid with barium carbonate gives
barium tetrathionate and sulphur (Takamatsu and Smith, •/. C., 1882, 41, 162;
Lewes, «/. C., 1881, 39, 68). See also Spring, A., 1879, 199, 97.
TABLE OF THIONIC ACIDS.
31T
S
•2
^S
"S
oo "3 3
g£
= d
"o 3
o
2'3^ |
li
_o
q
o?
3d
*j«B
0 g d .
i.
I1.
d O
.-Sj^ d "33
3
ft
1
Pentathionic acid, 11,
nmediate precipitate
phur, redissolving g:
on standing, if not i
excess and coagulated
o action.
t first yellow precipita
ing white with exces
agent on standing.
ellow precipitate, g
darkening; black on
ammonium hydroxide
Imost immediate bro
oration, becoming b
wa ming.
t first yellow pre<
gradually turning b
heating, with evolu
liCN.
n warming, whitish
precipitate.
rhite precipitate of sa
ecolori/ed. without
of dilute of H.,SO4.
- M
J?;
<$
p*1
<<
-<) O
?
a
^
4. to
'•r.
"•a m
Q
"3
Ej
.C J
C +i
-*J -^ ;3
'
•2
O
3
•*"d
d *
gs
ft &C ^*
a
•S
•d
2
O
3
CO
B"
bo
o d oj
Si
11 . 1
3
3
2
2'
*'J2 o
— '5
*g D
"o
O
'?
1
ft
,5ro >,
>. 2
IJ
l§= i
2
•4 i
ItS
fl
2
T OS r-
>^
•**
5
£ 5*
Vl -
xo° t«
ft
•d'T"
2
•3.
d
a.S
ft°~
° °x;
3 ° 'd
i
*2
cS
*s
0
d
J« S
-^ v
"CD ^
B
'r^ d
c
Q^
•-5
£ ^
^ ^ d
erf ^ ^
^•23 ^
*"*
O ^H
2
H
ft
0
0
=3 c«
fl>-d 3j
o £ §
3 > "
4^4^ V C
2
3«4
QJ O
55
h
>*
SE
< o
£
Q
O
£
B"
•d"
i
0
3
'ft
'3
|
•d
0
black precipitate,
vhite on stanuing.
pitate, soon becom-
oloratiou, even on
On warming, Ag2S
jipitate, becoming
excess of reagent.
mediate brown pre-
ren in presence of
04.
Trithion
1
'ft
ft
0
Evolution ol
tate of S.
Immediate
becoming i
Yellow preci
ing black.
No brown c
standing,
formed.
£5
: - ftg
: ^2
; af
QoM
PI
d'SS
0
Q
(D
49
q
1
B"
sjj
2*
2
1
®
• 3
5
i
1
3 |
'g
2
8
3
o
3
'o
£
'ft
3 3
|c
|g
E
ej
a
A
c, c.
2
0
0
o
O
o o
c ^
fe
K
55 55
o
2
•3
2
03
»
1
2
"©2
1
1
£
•d
2
2
1
1
'53
1 ^
2S
<H d
d fj
S
Ja
g
'd
2
•3
>. a
Sid
O gg
"3 a
1
a
P
•d
S
3
a
•S
U 0
JH tH
Bg
3.0
3 C
1ft
o
6
1
S
t.
11
8 1
Q) O
So
II
S to
£
S
i
Cfl
^
<5 ~-
(2"
O
.1
318 SULPHUROUS ACID. £285, 1.
§265. Sulphurous anhydride. SO, = 64.07 .
Sulphurous acid. H,SO, = 82.080 .
0
II
SIV0-"2 and H'2SIV0-"n , 0 = S = 0 and H — 0 — S — 0 — H.
1. Properties. — Sulphurous anhydride, SO. , sulphur dioxide, is a colorless gas
of a strong- suffocating- 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 (Lednc, C. r., 1893, 117, 219). It is liquefied at atmos-
pheric pressure upon cooling to — 10° (Pierre, C. r., 1873, 76, 214). In an open
dish it evaporates rapidly, the temperature of the remaining liquid dropping
to — 75°; or by evaporating rapidly under diminished pressure it becomes a
white wooly solid. Cooled to — 76.1° it becomes a snow-white solid (Faraday,
C.r., 1861, 53, S4(i). The dry gas is not combustible in the air, does not react
acid to litmus, but in presence 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 H.SO, . The gas is soluble in water, form-
ing probably sulphurous ac-id, H SO . The pure acid has not been isolated,
but forms salts mono and dibasic as if derived from such an acid (Michnelis
and Wagner, It., 1S74, 7, 1073). It has a strong odor from vaporization of
sulphurous anhydride, which is soon completely expelled upon boiling. The
acid oxidizes slowly in the air, forming H SO, . hence sulphurous acid usually
gives reactions for sulphuric acid. Light seems to play an important part in
this oxidation (Loew, Am. K., 1S70, 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 H.SO, restoring the color.
2. Occurrence. — Found free in volcanic gases (Uicciardi, #., 1887, 20, 464).
3. Formation. — (n) By burning- sulphur in air. (l>) Ry heating sulphur with
various metallic oxides, (r) Ry decomposition of thiosulphates with HC1. (d)
Ry burning H,S or CS, in air. (r) Ry the action of hot concentrated sulphuric
acid on metals, carbon, sulphur, etc. (f) Ry heating sulphur with sulphates.
(0) By decomposition of sulphites with acids:
(ft) S, + 2O, = 2SO:
(6) MnO, + S, = MnS + SO,
2Pb3O4 + :>S, — GPbS + 4SO,
(r) 2Na,S,03 + IHC1 = 4NaCl + 2SO2 + S, + 2H,.O
(d) 2H,S + 30, = 2SO, 4- 2H,0
CS, + 3O, = 2SO, + CO,
(p) Cu + 2H2SO4 = CuSO, + SO: + 2H2O
S, + 4H,SO4 = f.SO, -f 4H,O
C + 2H,SO, = 2SO, + CO, 4- 2H,O
(/) FeSO4 + 8, = FeS + 2SO,
(0) Na,SOa + 2H,SO, = 2NaHS04 + SO, + H,O
4 Preparation. — (fl) Ry heating moderately concentrated sulphuric acid with
copper turnings: Cu -f 2H,SO, = CuSO, + SO, + 2H2O . The gas is dried by
passing through concentrated sulphuric acid, (b) Ry heating a mixture of
sulphur and cupric oxide in a hard glass tube, (c) In a Kipp's generator by
<1. •composing cubes composed of three parts calcium sulphite and one part of
. -alcium sulphate, with dilute sulphuric acid (Neumann, /*., 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
§265, 6A. SULPHUROUS 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 6^1.) 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 68.861 volumes of sul-
phurous anhydride; at 20°, 36.206 volumes (Carius, A., 1855, 94, 148); or at 20°,
0.104 part by weight (Sims, J. C.. 1862. 14, 1). Charcoal absorbs 165 volumes,
camphor 308 volumes, glacial acetic acid 318 volumes of the gas. Liquid 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. Reactions. A. — With metals and their compounds. — Sulphurous acid
reacts with Zn . Fe , Sn , and Cu to form hyposulphurous acid, H,SO,
(Schiitzenberger, C. r., 1869, 69, 196). With Zn in the presence of HC1
it is reduced to hydrosulphuric acid: 3Zn -j- 6HC1 -j- H.SO, = 3ZnCl, -f
H2S -f 3ILO . 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 accompanied "by sulphates, due to
the fact that soluble sulphites nearly always contain sulphates (4).
Solution of lead acetate precipitates, from solutions of sulphites, lead
sulphite, 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, Ag2SO.< , easity soluble in very
^dilute nitric acid or in excess of alkaline sulphite, and turning dark-
brown when boiled, by formation of metallic silver and sulphuric acid.
Solution of mercurous nitrate with socliiim 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 (§58, Ge).
Solution of ferric chloride gives a red solution of ferric sulphite,
Fe2(S03)., ; or, in more concentrated solutions, a yellowish precipitate of
320 SULPHUROUS ACID. $265, G£.
basic ferric sulphite, also formed by addition of alcohol to tin- rod 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-
phite, BaSO . , easily soluble in dilute hydrochloric acid — disliiirli.oji from
sulphate, which is undissolved, and should be filtered out. Xow, on adding
to the filtrate nitrohydrochloric acid, a precipitate of barium sulphate
is obtained — evidence that sulphite has been dissolved by the hydrochloric
acid:
BaSO3 + 2HC1 = BaCL + H.SO-
Bad. + H,S03 + C12 + HaO = BaSO4 + 4HC1
Calcium 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:
Pb02 becomes lead sulphate.
Asv forms arsenous and sulphuric acids.
Sbv forms Sb'".
Cn" becomes cuprous sulphate.
CrVI forms chromic sulphate.
Co'" forms cobaltous sulphate.
Ni'" forms nickel sulphate.
Mn"+n forms manganous sulphate.
With MnO, in the cold, manganous dithionats, HnS.,0,, , is formed
(Omelin's Hand-book, 2, 174).
With stannoiifi 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, 6.).
B. — With non-metals and their compounds. — Upon other ;ici(N sul-
phurous acid acts as a reducing agent, except with hypophosphorous, phos-
phorous, and hydroeulphuric acids.
1. H3Fe(CN),. forms H4Fe(CN), and H,SO, .
2. HNO, and HNO, form NO and H,S04 .
3. PH3 + 2H,S03 = H,P04 -f Sa + 2H,0 (Carvazzi, Gazzelta, 1886, 16",
K59). ILPO... becomes H,P04 and the SO, is reduced to S , and with excosg
of H3PO, to H,S . H,PO; forms H,P04 and H,S (^253. (i).
4. H,S forms S from both compounds: 4H2S -f 2SO, == :5S, + 4H30 .
See also §263 .
§266, 1. SULPHURIC ACID. 321
5. Cl , HC10 , and HC103 form hydrochloric and sulphuric acids.
6. Br forms hydrobromic and sulphuric acids. HBrO;! 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).
forms first iodine then hydriodic acid, sulphuric acid in both cases
7. Ignition.— Acid sulphites heated in sealed tube to 150° are decomposed
into sulphates and sulphur (Barbaglia and Gucci, B., 1880, 13, 2325; Berthelot,
A. Ch., 1864, (4), 1, 392). Dry S02 at high heat with many metals is decom-
posed, forming a sulphide and sulphate or sulphite (Uhl, B., 1890, 23, 2151).
Sulphites are decomposed by heat into oxides and sulphurous anhydride:
CaSO3 = CaO -f SO-; or into sulphates and sulphides: 4Na.,SO3 =: 3Na2SO4 -j-
8. Detection.: — Free sulphurous acid is detected by its odor and by its
decolorizing action upon a solution of KMn04 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 Bad, in presence of HC1 . After removal of the BaS04
by filtration the sulphite is oxidized to sulphate by chlorine water and
precipitated by the excess of BaCL present.
Normal potassium sulphite, K2S03 , is alkaline to litmus but when
treated with Bad , gives- a neutral solution. The acid sulphite, KHS03 ,
is neutral to litmus but with BaCL gives an acid solution: 2KHSO-, -f-
BaCl2 = BaSO, + 2KC1 -f S02 -f H20 (Villiers, C. r., 1887, 104, 1177).
9. Estimation.— (a) After converting into H2S04 by HNO3 or Cl it is precipi-
tated by BaCL and weighed as BaSO, . (6) The oxidation is effected by fusing
with Na2CO3 and KNO3 (equal parts), (c) A standard solution of iodine is
added, and the excess of iodine determined by a standard solution of Na2S2O3 .
$66. Sulphuric acid. H2S04 = 98.086 .
0
II
H'2SVI0~"4 ,H — 0 — S — 0 — H.
1. Properties. — Absolute sulphuric acid, H2S04 , 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, •>.
odor. At ordinary temperature it is non-volatile and inodorous. A1 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 out- mole-
cule of water, H SO..H O , the melting point is highest, +7.5° (Pierre and
1'uchot, .4. Ch., 1S74, (5), KM).
H2S04 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'HS04 and M'.,S04 . 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, HC2H,S04 .
Sulphuric anhydride, SO, , is a colorless, fibrous or waxy solid, melting at
14.8° (Hebs, A., 188S, 246. 379). boiling at 4f>° (Schulz-Sellak, H., 1870, 3, 215),
and vapori/.ing 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. — Kound free in the spring water of volcanic districts. Found
combined in gypsum. CnSO, + 2H.O: in heavy spar. BaSO, : in celestine, SrSO,;
in Epsom salts. MgSO, 4- 7H,O: in (ilauber salt. Na,S04 + 1()H,0 , etc.
:!. Formation.— («) Hy electrolysing H20 . using Pt electrodes with pieces of
S attached (Hecqnerel. ('. r., 1863,56. 2:!7). (ft) Hv oxidizing 8 OF SO, in presence
of water by Cl . Br , HNO., , etc. (r) Hy heating S and HaO to 200°. ((/) Hy
adding H.6 to S03. (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
S02 evolved as a by-product in roasting various sulphides — e. (]., iron and
copper pyrites, blende, etc. (a) and (fc); or by burning sulphur in the air
to form the SO, . The SO., is passed into a large leaden chamber and
brought into contact with HNO., , steam, and air. The HNO., first oxidizes
a portion of the SO, (/•): the steam then reacts upon the NO,, form ing
HNO., and NO (</). This NO is at once oxidized again by the air to NO, ,
so that theoretically 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,S04 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.,S04 containing the two per cent of water,
a little more SO, or H,,S,07 than would be needed to make H2S04 ; then
passing perfectly dry air through it until the excess of S03 is removed,
leaving absolute H,S04 . Pyrosulphuric, or Xordhausen sulphuric acid,
H2So07 , is made by solution of sulphuric acid in sulphuric anhydride (e)\
by drying FeS04 + ?H20 until it becomes FeS04 + H.,0 , and then dis-
tilling (/). Sulphuric anhydride is made by the action of heat on sodium
$266, GA. SULPHURIC ACID. 323
pyrosulphate, Na2S207 (#), prepared by heating NaHS04 to dull redness ; by
distilling pyrosulphuric acid, the anhydride is collected in an ice-cooled
receiver; by heating H2S04 with P20, (/t):
(a) 2ZnS + CO2 = 2ZnO + 2S02
(6) 4FeS2 + 11O2 = 2Fe2O3 + SSO,
(c) SO2 + 2HNO3 = H2SO, + 2NO2
(d) 3NO2 + H2O = 2HN03 + NO
(e) H2SO, + S03 = H2S207
(0 4FeSO< + H2O = 2Fe203 + H2S2O7 + 2SO,
(g) Na2S207 = Na2S04 + S03
(ft) H2S04 + P205 = 2HP03 + S03
Sulphates are made: (a) by dissolving the metals in sulphuric acid;
(5) 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: 2KHS04 = K2S04 -f-
H2S04 . PbS04 is soluble in a saturated solution of NaCl in the cold,
depositing after some time crystals of PbCL , complete transposition being
effected. A solution of PbCL in NaCl is not precipitated on addition of
H2S04 (Field, J. C., 1872, 25/575).
6. Reactions. A. — With metals and their compounds. — Sulphuric acid,
dilute, has no action on Pb , Hg , Ag , Cu *, and Bi . An . Pt , Ir , and Rh
are not attacked by the acid, dilute or concentrated; other metals are
attacked by the hot concentrated acid with evolution of S02 . The fol-
*Andrews, J. Am. Soc., 1896, 18, 251.
324 SULPHURIC ACID. §266, 6B.
lowing metals : Sn , Th , Cd , Al . Fe , Co , Ni , Mn , Zn , Mg , K , and Na
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 S02 . 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 S02 forming HJ5 or a sulphide; and the H2S
decomposing the S02 with separation of sulphur (Ditte, A. Ch., 1890, (6),
19, 68; Muir and Adie, J. C., 1888, 53, 47).
Sulphuric acid or soluble sulphates react with soluble barium 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 very 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 calcium salts not too dilute form
a white precipitate of calcium sulphate (§188, 5c).
Dilute sulphuric acid docs not oxidize any of the lower metallic oxides.
The concentrated acid with the aid of heat effects the following changes:
Hg.,0 forms mercuric sulphate, and sulphurous anhydride is evolved.
SnCl2 forms, first, sulphurous anhydride, then hydrosulphuric acid,
stannic chloride at the same time being produced.
Fe" is changed to Fe.,(S04);; by hot concentrated sulphuric acid.
Mn"+n forms MnSO , and 0 . That is, all compounds of manganese
having a degree of oxidation above the dyad are reduced to the dyad with
evolution of oxygen.
Potassium permanganate dissolves in cold concentrated sulphuric acid
with formation of a green solution of a sulphate of the heptad manganese,
(Mn03)2S04 (§134, 5c).
Similarly the hot concentrated acid also reduces PbIV to Pb", Co'" to
Co", Ni'" to Ni", FeVI to Fe'", and CrVI to Cr"', oxygen being liberated
(oxidized) and the metal reduced while the bonds of the S04 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.
1. Many organic acids and other organic compounds are decomposed by
the hot concentrated acid, the elements of water being abstracted and
§266, 7. SULPHURIC ACID. 325
carbon set free. Continued heating of the carbon with the hot concen-
trated acid oxidizes it to CO, with liberation of S02 .
H2C204 becomes CO, , CO , and H,0 . The bonds of the H2S04 remain
unchanged.
K4Fe(CN)6 with dilute H2S04 forms HCN : 2K4Fe(CN)(i + 3H2S04 =
6HCN -f K,FeFe(CN)ti + 3K2S04 .
Cyanates are decomposed into CO, and NH3 : 2KCNO + 2H,S04 + 2H20
= K,S04 + (NH4)2S04 -f- 2C02 .
Thiocyanates are also decomposed by concentrated sulphuric acid.
2. Nitrites are decomposed with formation of nitric acid and NO :
6KNO, -f 3H,S04 = 3K,S04 + 2HNO, -f 4NO + 2H20 .
3. H3F02 or hypophosphites are oxidized to phosphoric acid with re-
duction of the sulphuric acid to sulphurous acid and then to sulphur.
4- Sulphur is . slowly changed by hot concentrated sulphuric acid to
sulphurous acid with reduction of the sulphuric acid to the same com-
pound. Hydrosulphuric 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: 3KC103 + 2H,S04 = 2KHS04 + KC104 +
2C102 + H,0 .
6. HBr forms Br and SO, . No action except in concentrated solution.
7. HI forms I and SO, .
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- , Ni and Sn are completely decomposed at
a red heat: 2FeSO4 = Fe2O3 + S03 + SO.,; 2CuS04 = 2CuO + 2SO2 + 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 Na2SO4 are
completely volatilized, preceded by partial decomposition.
Lead sulphate heated in a current of hydrogen is reduced according to the
following equation: 2PbS04 + 6H2 = Pb + PbS + SO2 + 6H,O . After a
distinct interval the remainder of the sulphur is removed as H.S: PbS + Ho =
Pb + H..S (Rodwell, J. C., 1863, 16, 42). Potassium sulphate heated in a
current of hydrogen is reduced to potassium acid-sulphide: K,SO4 + 4H, =
KOH -|- KHS + 3H,O (Berthelot, A. C7f., 1890, (6), 21, 400). Potassium acid-
sulphate, KHS04 , heated to 200° evolves H2S04 . The sodium acid-sulphate
decomposes more readily.
326 PER8ULP11UR1C ACID. $236, S.
8. Detection. — Free sulphuric acid or the soluble sulphates are detected
by precipitation in hot hydrochloric acid solution with barium chloridq,
forming the white, granular, insoluble barium sulphate.
The sulphates insoluble in water are decomposed for analysis — (1st) by
long boiling with solution of alkali carbonate; and more readily (2d) by
fusion with an alkali carbonate. In both cases there are produced — alkali
sulphates soluble in water, and carbonates soluble by hydrochloric or nitric
acid, after removing the sulphate (a). If the fusion be done on charcoal,
more or less dcoxidation will occur, reducing a part or the whole of the
sulphate to sulphide (7), and the carbonate to metal (as with lead, §57, 7),
or leaving the metal as a carbonate or oxide (7, §§222 and 228).
a. BaSOt + Na,C03 — Na,SO4 (soluble in water) + BaCO (soluble in acid).
A mixture of H,SO, and a sulphate may be separated by strong alcohol,
which precipitates the latter. A test for free Httlphuric acid, in distinction from
sulphates, may be made by the use of cane sugur, 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-black 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. — (a) By precipitation as barium sulphate and weighing as
such. The solution should be hot and acidified with hydrochloric acid, ami
the mixture should be boiled a few minutes after the addition of the barium
chloride. (/<) By precipitation as barium sulphate with an excess of an hydro-
chloric acid solution of barium chromate (three per cent hydrochloric acid).
Add NH.OH . fill to a definite volume, and lilter through a dry filter-paper.
Transfer an aliquot portion to an azotometer with HO , and after acidifying,
determine the oxygen evolved (Baumann, Z. angcic., 1891, 140) (§244, 6A, J2).
(c) When present in small amounts in drtnkmg water by a photometric methotl
(Hinds, C. N., 1896, 73, 285 and 299).
§267. Persulphuric acid. HS04 = 97.078.
The anhydride, S,O7 , 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 SO2 and O . In solution, the acid is obtained by the
electrolysis of concentrated H.SO,: also by the action of H2O2 on concentrated
H SO, .
At 0° persulphuric anhydride, S2O7 , consists of flexible crystalline needles,
remaining stable for several days. The solution in water decomposes rapidly;
more stable when dissolved in concentrated H SO, . When heated it decom-
poses into S03 and O . With SO2 it combines to form SO,: S207 + SO, = 3SOS .
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, #.,
1889, 22, 1518, 1528; 1892, 25, 95). Marshall (J. C., 1891, 59, 771) has prepared a
number of salts of persulphuric acid. The potassium salt, KSO, , is prepared
by electrolysis of a saturated solution of KHSO« 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.
and can be purified by recrystallization if not heated above (50°. The dry salt
is stable at 100°. With a solution of K2C03 it gives an abundant crystalline
.precipitate of KSO,. It is used in the cyanide process for the recovery of gold
(Elbs, Z. angew., 1897, 195). The potassium salt, soluble in 5G 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 oxidi/ed: KI is rapidly decomposed
upon warming; organic dyes are slowly bleached: K,Fe(CN)« becomes
K3Fe(CN)6; alcohol is slowly oxidized to aldehyde, rapidly upon warming.
The barium and lead salts are readily soluble in water (distinction from
H2S04).
§268. Chlorine. 01 = 35.45. Valence one, three, four, five, and seven.
1. Properties. — Molecular weight, 70.9. Vapor density, 35.8. The molecule con-
tains two atoms, C12 . Under ordinary air pressure it liquifies at — 33.6° and
solidifies at — 102° (Olszewski, M., 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 HC1 . On. cooling an aqueous solution of the gas to 0°, crystals of
C1:.10H2O separate out (Faraday, Quart. Jour, of /Set., 182:!, 15, 71). 'Chlorine
when passed into a solution of KOH produces, if cold, KC1 and KC1O , if hot,
KC1 and KC103: 2KOH + C12 = KC1 + KC10 + H,O; liKOH + ;5C12 = 5KC1 +
KC1O3 + 3H2O . Passed into an excess of NH4OH , NH4C1 and N are formed :
8NH4OH + 3C1, = 6NH4C1 + N2 + 8H,0; if chlorine be in excess chloride of
nitrogen is formed: NH4OH + 3C12 = NC13 -+- 3HC1 + H2O . The NC13 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 (§13).
Two oxides of chlorine have been isolated: C12O , hypochlorous anhydride
(§270), and C1O2 , chlorine dioxide. The latter is made by the addition of
EL SO., to KClOg at 0°. It is a yellowish-green gas, condensing at 0° to a red-
brown liquid. At — 59° it becomes a crystalline solid, resembling K2Cr207 . It
may be preserved in the dark, but becomes explosive in the sunlight.
The most important acids containing chlorine are discussed under tlio
sections following. They are:
Hydrochloric acid, HC1 .
Hypochlorous acid, HC10 .
Chlorous acid, HC10, .
Chloric acid, HC10.. .
Perchloric acid. HC104 .
2. Occurrence.— It does not occur free in nature, but its salts are numerous,
the most abundant being ITaCl .
3. Formation. — (a) By the action of HC1 upon higher oxides as indi-
cated in §269, 6A. The usual class-room or laboratory method is illus-
trated by the following equations :
MnO2 + 4HC1 = MnCL + Cl, + 2H..O
MnO, + 2NaCl + 3H.SO, = MnSO; -f- 2NaHSO4 + Cl. + 2H2O
328 CHLORINE. §268, 36.
(b) By fusing together NH4N03 and NH4C1 : 4NH4NOS + 2NH4C1 = 5N, -f
CL + 12H.,O . (c) By ignition of dry MgCL in the air: 2MgCl2 + O2 = 2MgO
+ 2C12 (Dewar, J. Soc. Ind., 1887, 6, 775). (d) Some chlorides are dissociated
by heat alone: 2AuCl3 = 2Au + ::CL .
4. Preparation. — (a) Wcld^tn'n ;>wvxx: MnO, is treated with HC1 , and th;e
MnCL formed is precipitated as Mn(OH), by adding CaiOH) . The Mn(OH),
is warmed by steam, and air is blown into it, oxidizing it again to MnO. , and
by repeating this process the same manganese is used over again. See Lunge
and Prett (Z. nnyric., 18!Kt, 09) for modification of this method, using HNO^ .
(b) Dearun'x //•wrtfx: HC1 , mixed with air, is passed over fire-bricks moistened
with CuCL and heated to about 440°. The heat first changes the CuCl2 to
CuCl , evolving chlorine: then the oxygen of the air, aided by the HC1 , oxi-
dizes the CuCl to CuCL . Tt 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 when- large amounts are needed.
5. Solubilities. — The maximum solubility of chlorine in water is at 106.
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 compounds. — 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; Veley, J. C., 1894, 65, 1). In the presence of acids the
oxidation of the metal takes place to the same degree as when that metallic
compound is acted upon by HC1 (£269, 0/1); a chloride is formed having
the same metallic valence that would have resulted from treating the
oxide or hydroxide with hydrochloric acid, e. f/., adding HC1 to 00,0., makes
CoCl2 not CoCl.., hence adding chlorine to metallic cobalt makes CoCl, and
not CoCl., . In alkaline mixture usually the highest degree of oxidation
possible is attained, as indicated by the following:
1. Pb" becomes Pb02 and a chloride in alkaline mixture. With PbCl,, it
is claimed that the unstable PbCl4 is formed (Sobrero and Selmi, A. Ch.,
1850, (3), 29, 102; Ditte, A. Ch., 1881, (5), 22, 500).
2. Hg^ becomes Hg" in acid and in alkaline mixture; also HC1 or a
chloride.
3. As"' becomes Asv in acid and in alkaline mixture. Some water must
be present or the reverse action takes place, forming AsCl:! (§269, GA2).
4- Sb"' becomes Sbv and a chloride with acids and alkalis.
5. Sn" becomes Snlv and a chloride with acids and alkalis.
6. JS.ovl~0 becomes MoVI and a chloride with acids and alkalis.
7. Hi'" becomes Biv and a chloride with alkalis only.
8. On' becomes Cu" and a chloride with alkalis and with acids.
§268, ()B1. CHLORINE. 329
9. Cr'" becomes Crvl and a chloride in alkaline mixture only.
i 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(OH)y and a chloride with alkalis only.
12. Hi" becomes Ni(OH)3 and a chloride with alkalis only.
13. Mn" becomes Mn02 and a chloride with alkalis only. See Ditte, I. c.t
for formation of MnCl4 .
B. — With non-metals and their compounds.
1. H2C204 in acid mixture: H2C204 + C12 = 2C02 + 2HC1 , the H2C204
must be in excess and hot (Guyard, Bl, 1879, (2), 31, 299) : in alkaline
mixture: K2C204 -4- 4KOH + C12 = 2K2C03 -f 2KC1 + 2HL>0 .
HCN becomes CNC1 and HC1 (Bischoff, B., 1872, 5, 80).
HCNS forms NHr! , H2S04 , C02 , and other variable products, and HC1
(Liebig, A., 1844, 50, 337).
H4Fe(CN)(1 becomes H,Fe(CN)(i and HC1 ; an excess of Cl finally decom-
poses the H,Fe(CN)0 .
2. Chlorine does not appear to have any oxidizing action upon the
oxides or acids of nitrogen.
3. Phosphorus and all lower oxidized forms become HrP04 with forma-
tion of HC1 .
4- Sulphur and all its lower oxidized forms are oxidized to H2S04 with
formation of HC1 . In an alkaline solution a sulphate and a chloride are
formed. With H2S , S is first deposited, which an excess of Cl oxidiz.es to
H2S04 . 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: KC10, + 2KOH + Cl, = KC10S
-4- 2KC1 '~j- H20 . With NaOH a hypochlorite is formed if cold, if hot a
chlorate :
2NaOH + Cl, = NaCIO + NaCl + H,O
<>NaOH + :;CL — NaCIO,
6. Chlorine does not oxidize bromine in acid mixture, in alkaline mix-
ture a bromate and a chloride are formed. HBr in acid solution becomes
free bromine, in alkaline mixture a bromate; hydrochloric acid or a chloride
being formed.
7. Iodine is oxidized to HIO., in acid mixture, forming HC1 ; 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 + C12 = 2HC1 + I, ; I., + r,CL, + OH ,0 - ?HIO, -f
10HC1 ; KI + 8KOH + 4C12 = KI04 + 8KC1 + 4H20 .
": By cornparing the oxidizing action of Cl with that of Br and I, the
following facts will be observed, and should be carefully considered. The
330 HYiniociiiAHtw ACID. §268,7
elements chlorine, bromine, and iodine have an oxidizing power in reverse
order of their atomic weights, chlorine b,eing 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 i^ot. 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 ^ways 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 Pb02, Ni(OH)., , Bi.,03 , Co(OH), . K,Fe04 , and Mn02", which
cannot be formed in presence of an acid.
It is very important to remember that those oxides which* are formed by
chlorine, in presence of a fixed alkali, but not in presence of an acid, are the
only ones ufiich can be reduced by hydrochloric acid. And further, tluit this
reduction proceeds not alicays to the original form, never proceeding Wfjond
fhat number of bonds capable of being formed in f&sence of an acid. Thus,
any lead salt/ with potassium hydroxide and chlorine, forms PbO., , and
this treated with hydrochloric acid again forum the lead salt, PbCl, . And
ferrous chloride with potassium hydroxide and chlorine forms K.,Fe04 , in
which iron is a true hexad, and K,Fe04 with hydrochloric acid forms, not
the ferrous chloride with which we began, but ferriq chloride, for it could"
only be oxidized to th.it point in presence of an acid. ^1* <
The above is true for bromine and iodine, as well as for chlorine.
7. Ignition. — See L
8. -Detection.— Fret* chlorine is recognized by its odor, by its liberation
of iodine from potassium iodide, by its bleaching action upon litmus,
indi^-i.. dr.. and by its net ion :is ;i powerful oxidizing agent (see above).
'.». Estimation. — (a) It is added to a solution of potassium iodide and the
liberated iodine determined by standard sodium thiosulphate. (ft) It is con-
verted into n chloride by reducing" agents, and estimated bv the usual methods
(§269, 8).
.
S269. Hydrochloric Acid. HC1 = 3(!.458.
H'Cl-'. H — Cl .
1. Properties. — IVipor </<«*////, ls.:^. At ordinary pressure it liquifies al
—102°, and solidifies at —112.5° (Ols/ewski, *f.. 1HS4, 5, 127). At 10° under
pressure of -40 atmospheres it condenses to a colorless liquid (Karaday, TV.,
1845, 155)u. Critical 1cni]icralnrc. ~>2.'.',°: critical ///> X.S///Y. s<; atmospheres ( Dewar,
(,'. X., 1SS5, 51. 27). Dissociated into H and Cl at about l.'OO0. but rombines
a .train ivpcrti cooling1 ^Devillr. r. ;•„ isi;."), 60, ::17). It is :i colorless jr. as, having-
an adfel, 'irritating- odor, lieadily absorbed by water. The chemically pure,
concemrated acid Las usually :i specific 'gravity of 1.20, and contains .19.11 per
cent HC1 (Lung-e and March'lewski. '/.. nntjnr.. 1891, 4, 133). The U. S. P. acid
lias a .s-/<rr///'- iiriirUj/ of 1.1(>:; at 15° and contains ::!.!) per cent HC1 . A eoncen:-
tiated solution of HC1 yives off yaseous HC1 faster than H,O: a dilute solution
'JJ269, .r>. HYDROCHLORIC ACID. 331
gives off H2O faster than HC1 , as a final result in both cases an acid sp. gr. 1.1
distils unchanged at 110° and contains 20.18 per cent HC1 (Bineau, A. Cft., 1843,
,(*). 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 H2S04 ; silver chloride must be heated nearly to the boiling point
of the H2S04 before the action begins. Lead, antimony and tin chlorides
.are slowly transposed.
(6) By the action of sunlight on a mixture of H and Cl , or by heating the
mixture to 150°. (c) Platinum black, palladium, charcoal, and some other sub-
stances which rapidJy absorb gases will cause the union of the hydrogen and
the chlorine, (d) When hydrogen 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, (c) Slowly formed by the action of chlorine upon
water in the sunlight: rapidly by its action upon reducing acids such as
H«C,04 , HH2PO2 . H2S, H-SO, . etc.: HH2PO2 + 2C12 + 2H2O = 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. (&) By the action of hydrochloric acid upon the
corresponding 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,
fpr 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 formed may be sublimed without decomposition, (c) Chlorides
of the first group are best made by precipitation, (d) Metals soluble in
hydrochloric acid evolve hydrogen and form chlorides. En these cases
ous, and not ic, salts are formed, (e) Many chlorides may be formed by
bringing HgCl, in contact with the hot metal.
4. Preparation. — For commercial purposes, made by treating NaCl with
H2S04 and distilling.
r>. 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 HgCl are insoluble
in water. PbCL is only slightly soluble in cold water (§57, 5c). These
three cmorides constitute the first or silver group of metals, and are pre-
cipitated frQjn their solutions by hydrochloric acid or soluble chlorides
(§61). Solutions of load salts arc 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 . PCI. , SbCL . and SnCl, arc liquids at ordinary tern-
332 HYDROCHLORIC ACID. §269,64-
perature. The first two are decomposed by water liberating HC1 : AsCl3
-f- 3H20 = H AsO., + 3HC1 . A saturated solution of bismuth nitrate
is precipitated by HC1 as the oxychloride (§76, Gf). Hydrochloric acid
increases the solubility of the chlorides of Pb , Hg , Ag , Sb , Au , Ft ,
Bi and Cu'; it decreases the solubility of Cd , Cu", Co , Ni , Mn , Th , Ba ;
Sr , Ca , Mg , Au , K and NH4 . Chlorides of Th , Ba , Na , K and NH4
are nearly insoluble in strong HC1 (Ditte, C. r., 1881, 92, 242; A. Oh.,
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) (§59, 6a); lead chloride is soluble in
fixed alkali hydroxides (§57, 6a).
HC1 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 Ni if
the acid be concentrated; As.,S3 and As..S- are insoluble in the cold con-
centrated acid, very slowly soluble in the hot concentrated acid; Hg3 ,
red, is insoluble; black, very slowly soluble in the hot concentrated acid.
HgS04 is only partially transposed by HC1 (£58, (>/), BaS04 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 HC1 (Ditte and Metther,
A. Ch., 1893, (6), 29, 389).
The following metallic oxides and hydroxides aro 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, Au'", Pt , MoVI. Bi'", Cu , Cd , Fe , Al . Cr'", Co", Ni",
Mn", Zn , Ba . Sr , Ca , Mg , K , and Na . The ignited oxides unite with
HC1 more slowly than when freshly precipitated or when dried at 100°.
Ignited CraO:, is insoluble in HC1 : other ignited oxides, as Fe,0, , ALO, ,
eto., require very long continued boiling with the HC1 to effect solution.
The following metallic compounds are attacked by hvdrochl^ic acid
with reduction of the metal and evolution of chlorine:
1. Pb"+n 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 Pb02 in presence of three per cent of acetic acid
(detection of a chloride in presence of a bromide) (Vortmann, 3/., 1882, 3,
510: 7?., 1887, 15, 1106).
§269, 655. HYDROCHLORIC ACID.
2. Asv becomes AsCl3 . (The presence of very concentrated HC1 Is
required; Fresenius, Z., 1862, 1, 448; Smith, J. Am. Soc., 1895, 17, 682
and 735.)
3. Biv becomes Bid, .
4. CrVI becomes CrCl3 . With K2Cr207 , bromine is completely libcrnted
from a bromide in presence of 4 cc. of H2S04 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 a chloride (Dechan, J. C., 1886, 49,
682). Dry HC1 does not reduce CrVI but combines with it to form the
volatile Cr02Cl2 , 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"+n becomes CoCl2 .
7. Ni"+n becomes NiCl2 .
8. Mji"+n becomes MnCL . Mn02 with small amounts of dilute H2S04
(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 tho
chlorine is not set free until considerable H2S04 has been added (Jones,
O. N., 1883, 48, 296). A mixture of KHS04 and KMn04 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 compounds.
• 1. No reducing action with H ,C.,04 , H2CO, , HCN , HCNS , H,Fe(CN), ,
and H3Fe(CN)0 .
2. HN02 forms chiefly NO and Cl . HN03 forms NO,C1 and Cl , or
NOC1 and Cl , or merely N02 and Cl . In case excess of HC1 is used tho
reaction is: 2HN03 + 6HC1 = 2ND + 3C12 + 4H20 (Koninck and KihouL
Z. anorg., 1890, 477). Dry HC1 gas, passed into a cold mixture of con-
centrated H2S04 and HNO., , reacts according to the following equations :
2HC1 -f 2HN03 = 2H20 -f 2N02 + C12 (Lunge, Z. angew., 1895, 4, 8,
and 11).
3. No' reducing action with H,S , H2S03 , or H,S04 . With thiosuluhates
the unstable H2S203 is liberated which decomposes as follow?: 2N'a,S.,0.. 4-
4HC1 = 4NaCl + S, + 2SO, + 2H20 . Sulphates of Ag ami Hg' ar^
completely transposed by HC1 , those of Ba , Sr , and Ca not ai all, all
others partially (Prescott, C. N., 1877, 36, 179).
4. With an excess of HC1 . hypophosphites, phosphites, and phosphates
are dissolved or transposed without reduction.
5. Hypochlorous acid forms chlorine and water: HC10 -4- HC1 = H«0 -j-
•C|12 . Chloric acid forms C102 , C120 , and Cl in varying proportions.
334 H7DROCHJ4>RIC ACID. §269,
but with HC1 in excess the following reaction takes place: KC10 -4- GllCl
— KC1 -f 3C12 -f 3H20 (Koninck and Nihoul, Z. an on/., 1890, 481). |r''
6. KBrO. is decomposed by boiling with HC1, the bromine bemgM&<it
free: 2KBr03 -f 12HC1 == 2KC1 -f Br, -f 5C1, + 6H20 (Kaemmerer,
J. pr., 18G2, 85, 452).
7. With HI03 , IC13 and Cl are formed, no action in dilute solutions:
HI03 + 5HC1 = IC13 -+- C12 -f 3H20 (Ditte, A., 1870, 156, 336). According
to Bugarsky (Z. anorg., 1895, 10, 387) KHI,00 with dilute H2S04 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:
PbCl, + Na,CO;, = PbO + 2NaCl + CO2 . If the carbonate be mixed with
charcoal, or if t'he fusion is done on a piece of charcoal, the metal is also
reduced: 2PbCl3 + ^Na2CO, + C = 2Pb + 4NaCl + :.CO2 .
Heated in a bead of microcosmic salt, previously saturated with copper
oxide in the inner blow-pipe flame, chlorides impart a blue color to the outer
•flame, due to copper chloride.
Dry sodium sulphate at 1">0° is transposed by dry HC1 (Colson, C. r., 1897,
124, SI). Gaseous HC1 transposes potassium and sodium sulphates completely
'at a dull-red heat. With the sulphates of .the alkaline earths the transposition
is nearly complete (Hensgen, /*., 187<>, 9, 1(>71). The silver halides heated with
bismuth sulphide on charcoal before the blow-pipe give distinguishing colored
incrustations: AgT . bright red; Ag-Br , deep yellow: AgCl , white (Goldschmidt,
C. C., 1870. 297). ..,,»;
8. Detection. — (a) In its soluble compounds, when not in mixture
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 .sf/ivr cliloride arc given in §59, 5r
and Gf. It is of analytical interest in that it is freely soluble in ammonium
hydroxide (considerably more freely than the bromide, ;ind far more freely
than the iodide of silver); soluble in hot, concent rated solution of am-
monium carbonate (which dissolves traces of bnmiide. 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.
(6) A test for traces of free hydrochloric acid, in distinction from metallic
chlorides, is made by heating the solution with Mn02 , without adding an
acid, and distilling into a solution of potassium iodide and starch. Larger
proportions of HC1 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 when brought
in contact with ammonia vapor. Also by bringing a stirring rod moist-
§269, Sf. HYDROCHLORIC ACID. 335
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 + Cr03 =: CrO.CL (chlorochromic anhydride) + H2O
(6) 4NaCl + K2Cr2O7 + 3H2SO4 =
2CrO2Cl2 + 2Na2S04 + K2SO4 + 3H2O
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 should be dry, should be distilled, 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 a 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) CrO2Cl3 + 2H2O = H2Cr04 + 2HC1
(d) CrO,Cl, + 4(NH4)OH= (NH4)2Cr04 + 2NH4C1 + 2H20
(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, 21) 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 Borchcrs (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 CuS04
and H2S .
(/) If a solution containing iodides, bromides, and chlorides be boiled
with Fe2(S04)3 , all the iodine is liberated and may be collected in a
solution of KI and estimated with standard Na2S20:i . The solution should
* With the chlorides of mercftry no brown fumps 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.
Before relying upon this test the absence of the above named metals should be assured.
336 HYDROCULOKH' ACID. £269, V
be cooled to about 60° arid a slight excess of KMnO, added. Tho bromine
js all liberated and may be collected in NH,OH and estimated as a bromide
after reduction with S02 . 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, (534 and
712; Hart, C. N., 188-1, 50, 2B8).
(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 ^.
of this solution for each tost. Iodine gives no precipitate; bromine gives
a white precipitate; and chlorine a black precipitate. If the bromide be
present in large excess, add silver nitrato, digest the precipitate with
ammonium hydroxide, add hydrogen sulphide and test the filtrate as the;
original solution. Liberate the halogen with KMnO, and H.SO, .
: (h) Deniges (BL, 1890, (3), 4, 481 ; 1891, (3), 5, 66) uses H.SO, and
Fe'" to liberate the iodine, and K.,Cr04 to liberate the bromine; then
after boiling off the I and Br he adds KMn04 to liberate the chlorine.
The iodine he detects with starch paper, the bromine fumes are absorbed
on a rod moistened with KOH , which then gives an orange-yellow color
with aniline. The chlorine ho collects as the bromine and obtains a violet
color with aniline.
(t) Dechan (J. f., 188(J, 50, 082; 1887, 51, 690) removes iodine of
iodides by distilling with a concentrated solution of K2Cr,07; thenf the
bromine of bromides by adding dilute ILS04 and again distilling. The
chloride is precipitated by AgNO., after dilution and addition of HNO ,
(;) Vortman (M., 1882, 3, 510; Z., 1886, 25, 172) detects chlorine w
presence of bromine and iodine as fallows: The flotation containing thi»
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. :md, after igni-
tion, weighed as AgCl. (ft) By a standard solution of AgNO3 . A little,
NaaHPO4 , or, better, KjCr20T , is added to the chloride to show the end of the.
reaction. When enough AgNO., lias been added to combine with the chlorine
the next addition gives a yellow precipitate with the phosphate, or a red witb
the chromate.
§271,2. HYPOCHLOROUS ACID— CHLOROUS ACID. 33?
§270. Hypochlorous acid. HC10 = 52.458 .
H'Cl'O-", H — 0 — Cl .
1. Properties. — Hypochlorous anhydride, CLO , is a reddish-yellow gas, con-
densing at about — 20° to a blood-rtd 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, HC10 , has not been isolated. Its
aqueous solution smells like ClaO , decomposing rapidly, especially in the sun-
light, into Cl and HC103 .
2. Occurrence. — Not found in nature.
3. Formation. — (a) By adding chlorine to HgO in the presence of water:
2HgO -f 2CL + H,O = Hg2OCl2 + 2HC1O (Carius, A., 1863, 126, 196). (b) By
adding five per cent nitric acid to calcium hypochlorite and distilling at ii
low temperature (Koffer, A., 1875, 177, 314). (c) By passing chlorine into the
sulphates of Mg ,-Zn , Al , Cu , Ca or Na: NanS04 + C12 + H.O = NaHSO4 +
NaCl -f- HC1O . (d) By heating a mixture of KC1O3 and H2C264 to 70° (Calvert
and Davies, A. Ch., 1859, (3), 55, 485).
4. 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 (B., 1887, 20, 1474) give the formula Ca"0,^1
to chlorinated lime. See also Kraut (A., 1882, 214, 244). Also as sodium
hypochlorite, made by treating- sodium hydroxide with chlorine short of satu-
ration in the cold: 2NaOH + CL = NaCIO + NaCl + H2O . 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
pharmacopoeial, under the name of solution of chlorinated soda (NaCl.NaCIO).
5. Solubilities. — Hypochlorites are all soluble in water and are decomposed
by heating.
6.. Reactions. — The hypochlorites are all unstable. They are decomposed by
nearly all acids, including CO,: 2Ca(ClO), + 2CO2 = 2CaCO3 + 2C1, + O2;
4NaClO + 4HC1 = 4NaCl + 2H2O + 2CL + O2 . 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. anyew., 1895, 501).
7. Ignition. — All hypochlorites are decomposed by heat: 2KC1O = 2KC1 + O2 .
8. Detection. — Although silver hypochlorite is soluble in water, it decom-
poses very quickly, so that on adding silver nitrate to sodium hypochlorite
the final reaction is as follows: SNaCIO + 3AgN03 = 2AgCl + AgClO3 -}-
3NaNO3 . When KC1O is shaken with Hg° , yellowish-red Hg,OCL is formed;
the other potassium salts of chlorine, i. r., KC1 , KC1O, , KC103 and KClOt ,
have no action upon Hg° . An indigo solution is decolored by hypochlorites,
while KMnO4 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 H2SO4 .
Rosenbaum (Z. angew., 1893, 80) gives a method for estimating the various
chlorine compounds in chlorinated lime.
§271. Chlorous acid. HC102 = 68.458 .
H'Cl'"0-"2 , H — 0 — Cl = 0 .
1. Properties. — The anhydride, CLO3 , has not been isolated and the free acid
is known only in solution, and this generally contains some HC1OS . It has an
intense yellow color and is very unstable.
2. Occurrence. — Neither the acid nor its salts are found in nature.
CHLORIXE PEROxini:. §271, ;J.
3. Formation — An impure chlorous acid is said to be formed when KC1O, is
treated with HNO, and As.O, , C.jHojO,, or C,H. (Millon, A. Ch., 1843, (3), 7,
298; Scbiel. A., 1859, 109, 318; Carius, .4., 18(56, 140, 317). Chlorites of a number
of metals have been made by adding1 the bases to a water solution of the acid;
also from KC1O by transposition.
4. Preparation. — KC1O, i:s prepared by adding- an aqueous solution of CIO, of
known strength to the proper quantity of KOH . and evaporating1 in a vacuum.
The crystals of KC1O., 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. Reactions. — Chlorous- 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., 18f>9, 151, 340).
8. Detection. — A concentrated solution of a chlorite gives a white precipitate
with silver nitrate, fairly readily soluble in more water. KMnO, is decolored,
a brown precipitate being formed. A solution of indigo is decolored even ip
presence of nrsenoiis acid (distinction from hypoehlorous acid). Chlorites
when slightly acidulated give a transient amethyst tint to a solution of ferrous
sulphate.
9. Estimation. — By reduction to chloride and estimation as such. By meas-
uring1 the amount of ferrous iron oxidized to the ferric condition: iFeSO, -+-
HC10, + 2H,SO4 = 2Fe,(S04), + HC1 + 2H,O .
$272. Chlorine Peroxide. C102 = 67.45.
C11V0-"... , ° '~=_ Cl — 0 — Cl =- 0 or 0 = Cl == 0 *.
Chlorine peroxide, CIO, , 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 CO, and ether it forms a mass of orange-yellow, brittle
• crystals. When warmed to about 00° 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, 298). The solution in water contains HC10 and HC10 .
It is prepared by carefully adding KC10 to cold concentrated H SO,; 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: 3KC1O, + 2H.SO, =
2KHSO4 + KC1O, + H.O + 2C1O, (Millon, /. c.). 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,: 2KC1O, + 2H,C,O4 = K2C=O« + 2H2O + 2C1O? H
2CO2 (Calvert and navies. A., 1859, 110. 344). It is also formed, mixed with
chlorine, when KC1O, is warmed with HC1 . HI is oxidized to I; SO, to H2SO4
Indigo is bleached even in presence of As.O .
•Pebal,4.,1875, 177, 1.
§273, 6A. CHLORIC ACID. 339
§273. Chloric acid. HC10:! — 84.458 .
H'ClvO-"3 , H — 0 — Cl =
! == ®
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 HC1O8.7H2O ,
containing 40.1 per cent HC1O3 (Kaemmerer, Fogg., 1869, 138, 390). Farther
attempts at concentration result in evolution of chlorine and oxygen, forming
HC104: 8HC1O3 = 4HC1O4 + 2H2O + 302 + 2C12 (Serullas, A. Ch., 1830, 45, 270).
Its solution in the cold is odorless and colorless; first reddening and then
bleaching litmus. It is a strong oxidizing agent, paper soaked with the acid
takes fire on drying. The anhydride, CLO5 , has not been isolated.
2. Occurrence Does not occur in nature.
3. Formation. — The free acid may be formed by adding an excess of H.,SiF.
to a hot solution of KC103; the filtrate is evaporated in vacuo, the excess of
H2SiF6 volatilizes, 'leaving the HC103 . Many chlorates are formed by treating
the metallic hydroxides with the free acid. Also by the action of Ba(C103),
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 AgC103 .
4. Preparation. — By adding H2SO4 in molecular proportions to a solution of
Ba(ClO3)2 . 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: 6KOH + 3C12 = 5KC1 + KC1O3 + 3H2O .
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 stablo
metallic chlorates; soluble in about 21 parts water at 10° (Blarez, C. r.,
1891, 112, 1213).
6. Reactions. A.— 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 H2S04 tho
reduction to chloride is complete, and with sodium 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 Asv.
Sb'" forms Sbv.
Sn" forms Sniv.
Cu' forms Cu".
;* Gladstone and Tribe's copper-zinc couple is prepared by treating thin zinc foil with a 1 per
cent solution of copper sulphate uutil the zinc is covered with a black deposit of reduced cop-
per. When washed and dried it is ready for use.
340 CHLORIC ACID. §273, 6B;
Cr'" forms Crvr, chromic salts are readily oxidized to chromic sic id on
boiling with KC103 and HN03 .
Fe" forms Fe'" (a distinction from perchloric acid) (Carnot, C. r., 189G,
122, 452).
Mn" forms Mnlv, manganous salts are rapidly oxidized to MnO, on warm-
ing with KC103 and HNO , .
Salts of lead, cobalt, and nickel do not appear to be oxidized on boiling
with KC103 and HNO;! .
B. — With non-metals and their compounds.
1. HoC204 forms CO., and varying proportions of Cl and HC1 . Heat
and excess of oxalic acid favors the production of HC1 (Guyard, BL, 1879,
(2), 31, 299). All oxalates are decomposed, C02 and a chlorate or chloride
of the metal being formed. Carbonates are all transposed.
HCNS forms H,S04 , HCN , and HC1 .
H4Fe(CN)0 first forms H,Fe(CN),. and HC1 ; a great excess of HC10,
decomposes the H:1Fe(CN)(1 .
2. HN02 forms HNO, and Cl . Nitrites are transposed and oxidized,
forming chlorates or nitrates of the metal.
3. PH., , HH..PO, , and H,P03 form H3P04 and HC1 . Hypophosphites
and phosphites are transposed and then oxidized, H3P04 and a chlorate or
a chloride of the metal being produced.
4. SVI~n forms SVI and HC1 ; that is, the sulphur of all compounds
becomes H.,S04 with formation of HC1 . All sulphides, sulphites, thio-
sulphates, etc., are transposed, forming a chlorate, chloride, or sulphate
of the metal.
5. HC1 in excess forms only Cl and H.O ($269, f>/tf). NaCl warmed with
HC10., evolves Cl , leaving only NaCIO, .
6. HBr forms Br and HC1 . Or warmed with HC10 , evolves Br , leav-
ing only KC10., .
7. I and HI form HI03 and HC1 . Soluble iodides form iodic acid or
an iodate.
7. Ignition. — All chlorates are resolved by heat into chlorides and
oxygen: 2KC10;! = 2KC1 -|- 302 . Some perch lo rale is usually formed as
an intermediate product: 2KC103 = KC104 + KC1 + 02 (Serullas, A. Ch.,
1830, (2), 45, 270). In presence of various metallic oxides, etc., the
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 is
given off (Spring and Prost, BL, 1889, (3), 1, 340). When triturated or
heated with combustible substances, charcoal, organic substances, sulphur,
sulphites, cyanides, thiosulphates, hypophosphites, reduced iron, etc.—
§274, 1. PERCHLORIC ACID. 341
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., MnVI~n becomes MnVI . That
is!, any compound of manganese having less than six bonds is oxidized to
the hexad (a). Cr'" becomes CrVI (6). Asv-n becomes Asv (c). Pblv~n
becomes PbIV (d). 'Co'"-" becomes Co'" (e). Clv~n becomes C1V (jf). Pv~n
becomes Pv (g). Iv~n becomes Iv (h). SVI-n becomes SVI (t).
(a) 3Mn,O4 + 18KOH + 5KC103 = 9K2MnO4 + 5KC1 + 9H,O
(6) 2CrCl3 + lONaOH + NaC103 = 2Na2CrO4 + 7NaCl + 5H2O
(c) 3As4 + 36KOH + 10KC1O3 = 12K3AsO4 + 10X01 + 18H2O
(d) 3Pb3O4 + Na.CO, + 2NaC103 = 9Pb02 + 2NaCl + Na2C03
(e) GCoCl, + 12KOH H- KC103 = 3Co2O3 + 13X01. + 6HSO
(f) 3K,C4H4O8 + 5X010, = 5KC1 + 3K2C03 + 9CO2 + 6H2O
(fir) 3Pb(H2P02)2 + 18KOH + 5KC103 = 3PbO2 + GX,P04 -f 5KC1 + 15H2O
(h) ZnI2 + K2CO3 + 2KC103 = ZnO +, 2KIO3 + 2KC1 + CO,
(0 SKoS.O,. + 12K,CO3 + lOKClOa = 15K2SO4 + 10X01 + 12CO2
8. Detection. va) Dry chlorates when warmed with concentrated sul-
phuric acid, detonate evolving yellow fumes: 3KC103 -f- 2H2S04 — 2KHS04
-f- KC104 -j~ SClOo -f- H.,0 . This action is modified by reducing agents;
some acting rapidly, increase the detonation; others acting slowly, lessen
it. (6) HC10.5 , like HNO., •, decolors indigo solution and gives colors with
brucine. diphenylamine, paratoluidine, and phenol similar to those formed
by HNOa. (c) By ignition a chloride is left: 2KC103 = 2KC1 + 302 .
(d) It is changed to a chloride by nascent hydrogen: 2KC10H -+- 6Zn -f-
7H2S04 == 6ZnS04 + K2S04 + 2HC1 + 6H20; or by reducing acids or
bases : 2KC103 + H2S04 4> 6H,SO, = K,S04 -f- 6H2S04 + 2HC1 . The
resulting HC1 is then identified in the usual manner. Chlorides, if origin-
ally present, should first be removed bv silver nitrate.
0
9. Estimation. — (c) Reduction to a chloride and estimation as such. (6) Addi-
tion of HG1 and KI and estimation of the liberated iodine with standard
Na2S2O3 .
§274. Perchloric acid. HC104 = 100.458 .
= 0
H'ClVII0-"4 , H — 0 — Cl = 0
= 0
1. Properties.— Specific gravity, 1.782 at 15°. The anhydrous HC1O4 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, '?.
in contact with the anhydrous acid take fire. The dilute aeid is very stable, not.
being easily reduced (Berthelot, A. Ch., 1882, (5), 27, 214). It does not 1)1 each,
but merely reddens blue litmus paper.
2. Occurrence. — Not found in nature.
3. Formation. — («) By electrolysis of a solution of Cl or HC1 in water
(Riche, ('. r., 1858, 46, 34«). (^)'KCIO, is formed by electrolysis of KC10, ,
using platinum electrodes (Lidoff and Tichomiroff. •/. C., 188:?. 44. H'.t). (c)
KC103 is heated with an excess of H SiF0 , after cooling and filtering, the
filtrate is carefully distilled (Koseoe, J. C., 1863, 16, 82; A., 1862. 121, :!4<i).
((/) By treating the sulpliate of the metal, the perch lornte of which is desired.
with Ba(ClOJ2 in molecular proportions. (<•) By treating the chloride of the
metal, the perchlorate of which is desired, with AgCIO, in molecular propor-
tions.
4. Preparation.— KC1O, is made by carefully heating KC1O until no more
oxygen is evolved: 2KC10, = KC1 + KC1O. + O2 (7). The residue is dissolved
in water and upon cooling crystals of KC1O, separate. The free acid, nearly
pure, is obtained by cautiously distilling KC10, with concentrated H SO, .
5. Solubilities.— All of the perehlorates of the ordinary metals are soluble
in water, and all are deliquescent except NH,C1O, KC1O, , Pb(ClO4), and
HgClO4 (Serullas, A. Ch., 1831, 46, 362). Potassium perehlorate is soluble in
142.9 parts of water at 0°, in 52.5 parts at 25°, and in 5 parts at 100° (Muir,
C. A'., 1876, 33, i:>). KC1O4 is insoluble in alcohol (distinction from NaC104)
(Schloessing, .1. Ch., 1877, (5), 11, 561).
«>. Reactions — Iron and /.ine evolve hydrogen when treated with perchloric
acid. The acid reacts with the hydroxides ot many metals to form per-
ehlorates. It is not reduced by HCl , HNO . H:S or SO. . Iodine is oxidized
to HIO, with liberation of chlorine: I, -f 2HC1O4 = 2HIO, + Cl, . A solution
of indigo is not decolored by HC1O, even after the addition of HCl (distinction
from all other oxyacids of chlorine). It is not reduced by the zinc-copper
couple (distinction from chlorate). Sodium perchlorate, NaCIO, . is used as a
reagent to precipitate potassiun salts.
7. Ignition. — I'erchlorates strongly ignited evolve oxygen and leave a chlorid.-
(§242, 3).
8. Detection. — In presence of a hypochlorite, chlorite, chlorate and chlorui.-
boil thoroughly with HCl; the first three are decomposed, leaving chloride and
perchlorate. Remove the chloride with AgNO, and fuse the evaporated tilt rat..-
with Na,CO, . 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. (It) It is fused with y.inc 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),
(r) KC1O, is heated to 200° with HPO, and KI; the iodine liberated showing
the amount of perchlorate present (Ciooch and Kreider, Am. S., 1894, 48, 315; and
1895, 40, 287).
§275. Bromine. Br = 79.95 . Valence one and five.
1. Properties — Molecular weight, 159.90; vapor tenuity, SO; specific gravity, 3.18828
at 0°; boiling point, 59.27° (Thorpe, J. C., 1880, 37, 172). At —7.2° it becomes a
brown solid (Philipps, B., 187(.», 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 ibdjne
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. BROMINE. 343
No oxides of bromine have, with certainty, been isolated. The well-estab-
lished acids are: Hydrobromie, HBr; hypobrbmous, HBrO; bromic, HBrOa .
2. Occurrence. — Not found free in nature. As a bromide in sea water, mother
liquor from salt wells, mUgMul springs, and in a few minerals.
3. Formation. — (a) Hydrobromie acid or any soluble bromide is warmed with
MiiO and H,SO4 . (ft) Any soluble bromide is treated with chlorine water
and the solution warnitd.
4. Preparation. — The bromine of commerce is obtained chiefly from the
mother liquor of the salt works: («) By treating with MnO and H.,SO4: MgBr,
-t- MnO2 + 2H2SO4 = MgSO4 + MnS04 + Br, + 2H2O " (b) By leading a
current of steam and chlorine into the bottom of a vessel tilled with coke,
into which a stream -of the mother liquor flows from above: MgBr._, -4- CL =
MgCL + Br, . (c) By adding to the mother liquor a mixture of Mg(OH), ,
suspended in water and saturated with chlorine, rendering acid and distilling
in a current of steam: Mg(C10.,), + OMgBr, + 12HC1 = 7Mg€L + GH2O +
6Br2 . (rf) 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-yellow solution (Dancer J. C., 1862, 15, 477). Its water solution is
permanent, but slowly decomposes: 2Br2 -+- 2H...O = 4HBr + OL • Much more
soluble in HC1 , 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. Reactions. A. — With metals and their compounds. — Bromine unites
directly with gold, platinum, and all ordinary metals to form bromides.
Silver salts are precipitated, yellow-white, as bromide and bromate:
6AgN03 + 3Br2 -f 3H,0 = SAgBr -f AgBr03 -f 6HNO, . 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.
1. Pb" becomes Pb02 in alkaline mixture only.
2. Hg' becomes Hg" in acid and in alkaline mixture.
3. As'" becomes Asv in acid and in alkaline mixture. With AsH3 and
a solution of bromine in water ILAs03 is first formed, and if the bromine
be in excess the final products are H3As04 and HBr .
.aad in alkaline mixture.
5. Sn" becomes Snlv iffl&id *&d in alkaline mixture.
6. Bi'" becomes Bi205 in\|k.rf!i|ifc mixture only.
7. Cu' becomes Cu" in ac* and^flk^line mixture.
8. Cr'" becomes CrVI in alkaline^ mixture only.
9. Fe" becomes Fe'" in acid mixture^jji alkaline mixture the iron is
further oxidized to a ferrate, HBr or a%romide being formed.
10. Co" becomes Co'" li^lkaline mixture only*.
11. Ni" becomes Ni"' in 'alkaline mixture only (Kilpius,- J. C., 187G,
29, 742).
344 BROMINE. §275, 6.4, 12.
1.2. MnIV~n becomes MnIV in alkaline mixture only.
B. — With non-metals and their compounds.
1. H2C204 becomes a carbonate and a bromide in alkaline mixture. An
excess of hot saturated oxalic solution changes Br to HBr .
HCNS forms, among other products, H.,S04 and a bromide in acid mix-
ture, and a sulphate and a bromide in alkaline mixture.
H4Fe(CN)(! in acid mixture forms H;!Fe(CN)0 and HBr , in alkaline mix-
ture a ferricyanide and a bromide (Wagner, J. C., 1876, 29, 741).
2. HN02 becomes HN03 and HBr if dilute and cold.
3. PH.,, HH,PO, and H,PO:! become H.,P04 and HBr with acids, and a
phosphate and a bromide in alkaline mixture. P and Br unite to form
PBr., or PBr- , 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°, H,S , H,SO, . H,S,0:t , SVI~n becomes H,S04 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. Tn 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
dements 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. — Warming1 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 CS., , 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 CS., 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 he distinguished from the violet color 6f 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 with CS, .
9. Estimation. — (a) The bromine is made to act upon KI . and the iodine
which is liberated is estimated by standard solution of Na.jS20, . (b) It is
estimated by the amount of As O which it oxidizes in alkaline solution, (r) It
is converted into HBr by H,S or H2SO, , and then precipitated by AgNO, .
and weighed as AgBr .
§276, 6A. HYDROBROMIC ACID. 345
§276. Hydrobromic acid. HBr = 80.958 .
H'Br-' , H — Br .
1. Properties. — Molecular weight, 149.9. Vapor density, 39.1. A colorless gas,
condenses to a liquid at — 159° and solidifies 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 b2.02
per cent HBr, or very nearly HBr.H2O . If a saturated solution is boiled,
chiefly HBr is given oft', and if a dilute solution is boiled, chiefly H2O 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 of HBr.5H,O .
2. Occurrence. — Not found free in nature, in combination as bromides in sea
water and in some minerals.
3. Formation. — («) By action of bromine upon phosphorus immersed in
water, the amorphous phosphorus is preferred: P4 + lOBr^ + 16H2O = 4H3PO4
+ 20HBr . (6) By action of H3PO4 or H2S04 on KBr (Bertrand, /. C'., 1876, 29,
877). (<?)• By transposition of BaBra by cold dilute H,S04 added in molecular
proportions, (d) By passing a mixture of Br and H over platinum sponge.
(?) By action of Br on H3PO, . (/) By adding Br to Na2S03 .
Metallic bromides are formed: (Jf) By direct union of the elements, but in a
few cases heat is required to effect the combination. (2) By action of HBr
upon the metallic oxides, hydroxides and carbonates. (3) Many bromides are
formed by action of HBr on the free metal, ous salts and not ic being formed.
(4) Bromides of the first group are best made by precipitation. (5) Bromides
of K , Ha , Ba , Sr and Ca are made by the action of bromine on their hydrox-
ides and subsequent fusion:
GKOH + :;Br3 = KBrO3 + 5KBr + :;H2O
2KBr03 (ignited) = 2KBr + 302
4. Preparation -(a) H2S 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 H2S, and the latter portion if it con-
tains H2SO4 (Kecoura, C. r., 1890, 110, 784). (b) H,S04 is added to a concen-
trated solution of KBr; after twenty-four hours the greater portion of the
KHSO, has crystallized out. The remaining liquor is then distilled. The
product usually contains traces of H.SO4. (c) By passing bromine into hot
paraffine (Crismer, B,, 1884, 17, 649).
5. Solubilities. — Silver and mercurous bromide are insoluble in water,
lead bromide i& 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 alcohol, the alkali bromides are sparingly or slightl}' soluble: calcium
bromide, soluble; mercuric bromide, soluble; mercurous bromide, insolu-
ble. Silver bromide is soluble in NH4OH .
6. Reactions. — .1. — With metals and their compounds. — Hydrobromic
acid dissolves many metals with the formation of bromides and evolution
of hydrogen, e. g., Pb , Sn , Fe , Al , Co , Ni , Zn , and the metals of the
346 HYDROBROM1U ACID. §276, (L4, 1.
calcium and the alkali groups. It unites with salt forming oxides and
hydroxides to produce bromides without change of valence: PbO -f- ;JHBr
= PbBr2 -f- H20 . 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:
1. Pb"+n becomes PbBr, and Br .
2. Asv becomes As'" and Br . The HBr must be concentrated and in
excess, and the Asv compound merely moistened with water: H,As04 -f-
2 HBr =- H.(AsO;t -f- Br2 -j- H20 . In presence of much water the reverse
action trfkes place : H,AsO, + Br2 -f H20 = H3As04 -4- 2HBr .
3. Sbv becomes Sb'" and Br .
.'/. Biv becomes BiBr and Br .
5. FeVI becomes Fe'" and not Fe" , and Br .
6. CrVI becomes CrBr, and Br (a separation from a chloride if the solu-
tion be dilute) (Friedheim and Meyer, Z. anon/., 1891, 1, 407).. KBr is not
decomposed by a boiling concentrated solution of K2Cr207 (separation
from KI) (Dechan, J. ('., 1887, 51, 690).
7. Co"+n becomes CoBr, and Br .
<<?. Ni"+n becomes NiBr, and Br .
0. Mn"+n becomes MnBr, and Br (£269, 8: Jannasch and Aschoff, Z.
aiiory.. 1891, 1, 144 and 245). KMn04 liberates all the bromine from KBr
in presence of CuS04 (a separation of bromide from chloride (Baubigny
and Rivals, (\ /•., 1897, 124, 859 and 954).
Silver nitrate solution precipitates, from solutions of bromides, silver
bromide, AgBr . yellowish-white 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 nieirnnnift bromide, HgBr,
yellowish-white, soluble in excess of alkali bromides.
Solutions of lead salts precipitate, from solutions not very dilute, lead
bromide, PbBr., , white.
B. — With non-metals and their compounds.
1. H,Fe(CN),. becomes H4Fe(CN)t; 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:
2K4Fe(CN),; + Br. = 2K,Fe(CN)(i + 2KBr .
2. HN02 , in dilute solutions, no action (distinction from HI) (Gooch and
Ensign, Am. S., 1890, 140, 145 and 283).
HNO:1 becomes NO and Br .
£276, 8. UYIUIOBUOMIC ACID. 347
•3. Phosphorus compounds are not reduced.
J/. H2S04 becomes S02 and Br . Both acids must be concentrated and
hot, otherwise the reverse action takes place: SO,, -j- Br2 -f- 2H20 = H2S04
-f 2HBr . With H.,S04 , sp. gr. 1.41, no bromine is set free even when
solution is boiled (Keit and Kubierschky, J. Pharm., 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 (Xihoul, Z. anyew., 1891, 441).
HC103 becomes HC1 and Br . If the HC10, be concentrated other pro-
ducts may appear.
G. HBrO liberates Br from both acids ; the same with HBr03 .
7. HIO;; becomes I and Br .
S. 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 vmdecomposed in presence of
air; e. </., AsBr. , SbBr, , HgBr and HgBr., . Some can be sublimed only by
exclusion of air and moisture; c. //., AlBr3 and NiBr, . Bromides of sodium and
potassium are not changed by 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
CS., (£275, 0). The oxidizing agent used to liberate the bromine varies
according to the conditions. Chlorine is more commonly employed and
acts when cold (67/5). A large excess of chlorine is to be avoided, as it
decolorizes bromine solutions with formation of a chlorbromide. Nitric
acid Avhen dilute acts slowly unless hot. H2S04 , 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 bromate will be formed, free bromine not
being a visible intermediate step in the oxidation: KBr -4- 6KOH -4- 3C12
= KBrO, -f 6KC1 -f 3HL,0 . If an iodide be present: (a) In absence of a
chloride precipitate with silver nitrate, and digest the precipitate with
NH4OH , which will dissolve the AgBr and none of the Agl . The filtrate
may he treated with H2S , which precipitates the silver as Ag.,S , leaving
the bromine in the filtrate as NH4Br , which may be detected in the usual
way. (ft) 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
"bromine may be observed: SKI -f- 2KBr -f 7CU -4- 6H,0 = ?HIO.. -f- Bra
348 HYPOBROMOUS ACID—BROMIC ACID. £276, !).
-f- 1KC1 -|- 10HG1 . (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); warm 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.
OKI + GKBr + 2KC103 + 7H.SO4 = 3l2 + 3Br2 + 7K,SO4 + 2HC1 + f,H,0
61, + xBr, + 10KC10, + r,H2SO4 + f,H2O = 12HIO, -f xBr2 + r>K,SO« + 10HC1
9. Estimation. — (a) It is converted into AgBr , and after gentle ignition
weighed as such, (b) The bromide is oxidixed to free bromine, which is
passed into a solution of KI and the liberated iodine titrated with standard
Na.S,O . (p) 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 KMnO. .
§277. Hypobromous acid. HBrO — 96.958 .
H'Br'O-" , H — 0 — Br .
The anhydride, Br,O , 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 hypobromites 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: 2figO -f- 2Br2 + HjO =: Hg,OBr, + 2HBrO;
also by the action of bromine upon silver nitrate: AgNO^ + Br.. + H..O =
AgBr + HBrO + HNO, (I)anctr and Spiller, C. .Y., I860, 1, 38; 1SG2, 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. HBrO, = 128.958 .
H'BrvO-" ., , H — 0 — Br =
= O
1. Properties. — The anhydride, Br.Oj , 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 abort
80 per cent of the acid may be obtained.
2. Occurrence. — Neither the acid nor its salts arc found in nature.
3. Formation. — («) Bv the electrolysis of HBr (Kiche. r. /-., is;,s. 46. 348
(ft) By the decomposition of AgBrO, by Br: .-)AgBrO, + 3Br, + :!H,O = :>AgBr
+ 6HBrO3 . (<•) An alkali bromate is made by adding bromine to a solution
of chlorine in sodium carbonate (Kaemmerer, ./. />/•., 18(i2, 85, 4.52).
4. Preparation. — Bromates of Ba . Sr , Ca , K and Na are made by tin- act in::
of bromine upon the respective hydroxides at 100°: OKOH + iiBr, = :>KBr -f-
KBrOj + 3H..O . The free acid is prepared by adding dilute HSO, in slight
excess to Ba(BrO,),; the slight excess of H2SO4 being removed by the cautious
addition of Ba(OH). .
§278,8. BROMIC ACID.
5. Solubilities.— AgBr03 is soluble in 123 parts of water at 2-4.5°
(Noyes, Z. phys, Ch., 1890, 6, 246). Ba(Br03)2 is soluble in 124 parts of
water at ordinary temperature and in 24 parts at 100° (Rammelsberg,
Fogg., 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 Asv and a bromide.
3. Sb'" becomes Sbv and a bromide.
4. Sn" becomes SnIV and a bromide.
5. Cu' becomes Cu" and a bromide.
6. Fe" becomes Fe'" and a bromide.
7. Mn" becomes Mn02 and bromine.
8. Cr'" becomes H2Cr04 and bromine.
Silver nitrate precipitates in solutions not very dilute, silver bromate,
AgBrO:! , 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 an-1
from other argentic precipitates.
B. — With non-metals and their compounds.
1. H2C204 becomes CO., and Br. An excess of hot H2C204 changes the
Br to HBr (Guyard, El., 1879, (2), 31, 299).
HCNS becomes H2S04 , HBr and other products.
H4Fe(CN)(i becomes H,Fe(CN)6 and HBr . An excess of HBr03 carries
the oxidation farther.
2. HN02 reduces HBr03 , forming HN03 and Br .
3. PH3 , HH2P02 and H3P03 become H3P04 and HBr .
4. S and S02 become H2S04 and HBr .
H2S forms first S then H2S04 .
5. HC1 becomes Cl and Br .
6. HBr forms Br from both acids.
7. HI becomes I and Br . With an excess of HBr03 the products are
HIO:, and Br (Kaemmerer, I c., Wittstein, Z., 1876, 15, 61).
7. Ignition. — All bromates are decomposed upon heating. KBrO:. ,
NaBrO:! and Ca(BrO,)o evolve oxygen and leave the bromides. Co(Br03)., ,
Zn(BrOt!)., and other bromates evolve oxygen and bromine, leaving an oxide.
8. Detection. — The bromine is first liberated by some reducinir agent
that does not carry the reduction to the formation of HBr. H.,C,04 -is a
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 IIBrO:l remaining for some time intact, and the solution colorless. The
gradual decomposition of the HBrO. is first a resolution into HBr and 0,
and as fast as HBr is formed it acts with HBrO. . 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
carbon disulphide, and obtain no color, we have conclusive evidence of the
•licence of bromates. Hydrochloric acid transposes bromates and quickly
•lecomposes the bromic acid, liberating both bromine and chlorine.
A mixture of Itrnmnte and iodaie, 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
podium carbonate, will give the tests for bromides.
0. Estimation. — The bromate is reduced to free bromine or to a bromide and
determined as such.
§279. Iodine. I = 126.85. Usual valence one, five and seven (§12).
1. Properties.— Specific fjrnrity, 4.948 at 17° (Gay-Lussac). Mcltino point,
114.2°. Boilinii point, 184.3.r>° at 7f>0 mm. pressure (Ramsay and Young1, ./. ('.,
1886. 49. 4">.'{). At ordinary temperature iodine is a soft gray-black crystalline
solid with a metallic lustre. The thin crystals have a brownish-red appear-
ance. Precipitated iodine is a brownish-black powder. It vapori/es 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 I2:
above that temperature dissociation takes place, until at 1700° it is complete
rmd the molecule consists of sinple atoms (Biltz and Meyer, B., 1889. 22, 725).
The vapor of iodine unmixed with other pases is deep blue, mixed with air
or other pases 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,* forming1
a compound of an intense blue color. This colored body is quite stable in the
oold: decolors upon warmintr. 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 agents destroy the color by taking the iodine
into combination. Combined iodine does not react with starch.
*The compound formed when iodine unites with starch is regarded by Bondonneau (Bf., 1877,
(2). 28, 452) as an addition compound of the c imposition (C,H10O»)»I .
i 279, 6.4,5. IODINE. 351
Colorless solutions are formed by all the alkali hydroxides with iodine; the
rixed alkali hydroxides forming iodides and iodates. With ammonia in water
solution it dissolves more slowly, becoming colorless; the solution contains the
most of the iodine as ammonium iodide, and deposits a dark-brown powder,
Termed " iodide of nitrogen," very easily and violently explosive when dry.
According to Chattaway (Am., 1900, 24, 138) this compound has the composi-
tion N,H3I3 .
The anhydride of iodic acid, I:O5 , 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(IO), .
2. Occurrence Found free in some mineral waters (Wanklyn, C. N., 1886, 54,
r;00). 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 + Br2 =
2KBr + !•; and from iodates by nearly all reducing agents: 2HI03 + 5H2C2O4
= I. + 10OO, + 6H20 .
4. Preparation. — (a) The ashes of the sea plants are digested in hot water
and from the nitrate 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 MnO2 and H,S04 . (1>) The sodium iodate in the
mother liquor of the Chili saltpeter is reduced with SO, . the iodine precipitated
as Cul with CuSO4 . From the precipitate the iodine is recovered by distilla-
tion with MnO2 and H2SO4 . 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 5500 parts water at 10° to 12°
(Wittstein, J., 185*, 123), differing from Cl 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, glycorol, benzol, carbon disulphide, etc. Car-
bon disulphide readily removes the iodine from its solution or suspension
in water; with small amounts of iodine imparting to the carbon disulphide
a beautiful violet color, with large amounts the CS2 solution is almost
black.
6. Reactions. — .4. — 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.
1. Hg' becomes Hg" in acid and in alkaline mixture.
2. As'" becomes Asv in presence of alkalis only.
3. Sb'" becomes Sbv in presence of alkalis only.
4. Sn" becomes Snlv in acid or in alkaline mixture.
5. Cr'" becomes CrVI in presence of alkalis only.
352 IODISE. §279, 6.1 , 6.
6. Fe" becomes Fe'" in presence of alkalis only.
7. Co" becomes Co'" in presence of alkalis only.
8. Ni" is not oxidized.
9. Mn" becomes MnIV in presence of alkalis only.
B. — With non-metals and their compounds.
1. K4Fe(CN)6 is oxidized, forming K.,Fe(CN)tt and KI, action slow and
incomplete.
2. HN03 forms HIO;i and NO . Strong HN03 must be used (at least
sp. gr. 1.42). Action is slow. A very good method of making HI03 .
J. HH,P02 becomes H.P04 with acids and with alkalis.
4. H>S becomes S and HI; no action if both substances be perfectly dry
(Skraup, C. C., 1890, i, 4G9) (separation of H2S from AsH;!). According
to Saint-Gilles (.4. Ch., 1859, (3), 57, 221), in alkaline mixture from six
to seven per cent of the sulphur is oxidized to a sulphate.
H»S03 becomes H,S04 and HI. With a thiosulphate a tetrathionate is
formed: 2Na,S,0, + I, == Na,S40,., + ->NaI (Pickering, J. C., 1880, 37,
128).
5. Cl becomes IC1 or IC13 , depending upon the amount of chlorine
present, water should be absent. In the presence of water HC1 and HIO:t
are formed; in alkaline mixture a chloride and a periodate: I2 -(- 7C12 -f-
KiNaOH =: MNaCl + ?NaI04 + 8H..O . HC10, forms HIO, and HC1:
5HC10., + 3I2 + 3H.O = (5HIO, + :>HC1 .
(I. Br becomes IBr, decomposed by water (Bornemann, ^1., 1877, 189,
183). In alkaline mixture with an excess of Br a bromide and an iodate:
I2 + 5Br, -4- 12KOH = 2KIO, -4- lOKBr -f GH,0 . HBr03 becomes Br
and HIO, .
7. Iodine combines with KI in concentrated solution to form KI.;(KII,) .
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. <-t. .
The presence of tannin interferes with the usual tests for iodine unless ,->,
drop or two of ferric chloride solution be added (Tessier, Z., 1874, 11, 31 :r.
9. Estimation. — (a) It is reduced 1o an iodide, precipitated with AgNOn . aiut
after drying at 150°, weighed as Ag-I . It is estimated volumetric-ally with a
standard solution of Na.S.CK . using starch as an indicator, (b) 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, anyeic., 1891, 204).
£280, 5. HYDRIODIC ACID. 353
§280. Hydriodic acid. HI = 127.858 .
H'l-' , H — I .
1. Properties. — Molecular 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,
266). The constant boiling- point of the aqueous solution of the gas is 127°,
which solution contains 57 per cent of HI and has a specific gravity of 1.694
(Itoscoe, •/. (-'.. 1861, 13, 160). Gaseous HI is dissociated by heat, slowly at 260°;
rapidly at 240° (Lemoine, A. Cti., 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 liolzmaim, B,, 1889, 22, 869). (b) By direct union of the elements in pres-
ence, of platinum black at ;>00° to -400° (Lemoine, C. r., 1877, 85, 34). (c) From
Bnl, by adding H2SO4 in molecular proportions, (d) By the action of iodine
upon Na2SO3 or Na2S203 (Mene, C. r., 1849, 28, 478). (e) By the action of iodine
vjpon moist calcium hypophosphite: Ca(H2P02)2 + 4I2 + 4H2O = CaH4(PO4),
-h 8HI (Mene, /. 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
whose iodides are soluble in water. Iodides of lead, silver and mercury are
formed by precipitation.
4. Preparation. — (a) By passing H2S into a mixture of finely divided iodine
.suspended in water, adding more iodine as fast as the color disappears: 2I2 +
L'H2S = 4HI + S2 (Pellagri, Guszetta, 1875, 5, 423). (b) By bringing moist red
phosphorus in contact with iodine: P4 + 10I2 + 16H2O = 4H3P04 + 20HI
,Meyer,#., 1887, 20, 3381). (c) By passing vapors of iodine into hot liquid
juiraffine (Crismer, B., 1884, 17, 649). (d) By heating iodine with copaiba oil
(Bruylants, B., 1879, 12, 2059). It cannot be prepared by adding H2S04 to an
iodide and distilling (5).
a. Solubilities. — Iodides of lead, silver, mercury and cuprosum are in-
boluble. Iodides of other ordinary * metals are soluble, those of bismuth,
tin and antimony requiring a little free acid to hold 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,
X2HgI4; 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 HC1 or by dilute H,S04 . Hot concentrated H2S04
decomposes all iodides, those of Pb , Ag and Hg slowly but completely,
SO, and I being produced: 2KI -f 2H,S04 = K2S04 + I2 + SO, + 2H20 .
HNO, in excess first transposes then decomposes soluble iodides: 6KI -|-
HHNO, = GKN03 + 31, -f 2NO -f 4H,0 . If the HN03 be concentrated
the iodine is further oxidized : 31, + 10HNO, = 6HI03 + 10NO -f 2H20 .
Long-continued boiling with HN03 , sp. gr. 1.42, decomposes the insoluble
iodide?. Chlorine in the cold decomposes all soluble iodides, by heating
with chlorine the insoluble iodides are also decomposed: SKI -+- C12 =
* Thallium iodide, Tl I, is perfectly insoluble in cold water, a distinction and separation from
bromides and chlorides (Huebner. Z., 1872, 11, 397). Palladous iodide is insoluble in water.
354 HYDRIODIC ACID. £280, <U.
2KC1 -j- I2 . With an excess of chlorine the iodine is further oxidixed:
I2 -f 5C12 + (5H20 = 2HH>3 -f- 10HC1 . Silver iodide is almost insoluble
in ammonium hydroxide or ammonium carbonate (distinction from silver
chloride). It is soluble in KCN . Agl and PbI2 are soluble by decomposi-
tion in solution of alkali thiosulphates: Agl -f Na.&^O., = Nal -j-
NaAgS,03 . Lead iodide is soluble in a solution of the fixed alkalis.
(i. Reactions. — .1. — With metals and their compounds. — Silver nitrate
solution in excess precipitates, from solutions of iodides, silver iodide, Agl .
yellow-white, blackening in the light without appreciable separation of
iodine. For solubilities see paragraph above.
Solution of mercuric 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, (>/. Solution of mercurous nitrate precipitates mercurous iodide,
Hgl , yellow to green (£58, (i/).
Solution of lead nitrate or acetate precipitates, from solutions of iodides
not very dilute, lead iodide, PbI2 , bright-yellow — soluble, as stated in full
in £57, or.
Fallacious chloride, PdCL , precipitates, from solutions of iodides, pal-
ladous iodide, PdI2 , black, insoluble in water, alcohol or dilute acids, and
visible in 500,000 parts of solution. The reagent does not precipitate
bromine at all in moderately dilute solutions, slightly acidulated with HC1 .
Fallacious 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): 2CuS04 -f 4KI -- 2CuI + 2K2S04 + 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. Hydriodic acid unites with all metallic oxides and hydroxides
(expect ignited CrL,03) to form iodides; frequently, however, iodine is
liberated and an iodide of lower metallic valence is formed:
1. Pb"+n becomes Pb" .
2. Asv becomes As'" ; KI has no action upon normal K .AsO, (Friedheim
and Meyer, Z. anorg., 1891, 1, 409).
8. Sb'v becomes Sb'" . *
4. Biv becomes Bi'" .
5. Cu" 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, (>/*, G. HTDRIODIC ACID. 355
completely precipitated (separation from chlorides) (Baikow, Ch. Z., 1894,
18> 1661).
6. Fe'" becomes Fe" (§269, 8).
7. CrVI becomes Cr'" . K2Cr04 is not reduced by KI even upon boiling
the 'concentrated solutions. K2Cr207 with KI slowly gives I and Cr"' in
the cold. When KI is boiled with a concentrated solution of K2Cr207 the
iodine is completely liberated (separation from bromides and chloride.s
•whicli arc unchanged): GKI -f- 5K2Cr207 = 8K2Cr04 + Cr203 -f 31,
(Dechan, J. f'., 188(5, 50, 682; 1887, 51, 690). When Agl is boiled with
K2Cr.,07 and H2S04 no iodine is evolved, chromium is reduced and the
iodide becomes silver iodate: K,Cr,07 -f Agl + 5H2S04 = 2KHS04 -f
Cr2(S04), + AglO, -f- 4H20 (Macnair, J. C., 1893, 63, 1051).
S. Co"+n becomes Co"; KI has no reducing action upon cobaltic hy-
droxide.
9. Ni"+n becomes Ni"; KI reduces Ni'" , liberating iodine.
10. Mn"+n becomes Mn" . When KI is boiled with KMn04 the manga-
nese becomes Mn02 and the iodide is oxidized to an iodate: 6KMn04 -\-
3KI + 3H20 = 3KI03 + 6Mn02 + 6KOH (Groeger, Z. angew., 1894, 13
and 52) (distinction from bromides, which do not decolor permanganates).
B. — With non-metals and their compounds.
1. H3Fe(CN)0 forms H4Fe(CN)6 and I; the reaction also takes place in
neutral mixture. >
2. HN02 forms NO and I (separation of iodide from bromide and
chloride) (Jannasch and Aschoff, Z. anorg., 1891, 1, 144 and 245).
HN03 forms NO and I , with further oxidations to HIO.{ with concen-
trated HN03 . The HN02 acts much more rapidly than the HN03 .
3. Xo reduction with phosphorous compounds.
4. H2S04 dilute no action; with the concentrated acid in excess, S02 and
I are formed: 2KI + 3H2S04 = I8 + S02 + 2KHS04 + 2H20 ; if KI be
added in excess to boiling H2S04 , H2S and I are formed: SKI -|- 9H2S04 =
4I2 -f H2S -f 8KHS04 + 4H,0 (Jackson, J. C., 1883, 43, 339). Ammo-
nium persulphate liberates iodine from iodides at ordinary temperature
(Engel, C. r., 1894, 118, 1263).
it. Cl in excess forms HC1 and HI03; with excess of HI , HC1 and I are
formed. In the presence of a fixed alkali a periodate and a chloride are
formed: KI -f 8KOH + 4C12 = 8KC1 + KI04 + 4H20 . Hypochlorous
acid oxidizes to iodine, then to iodic in acid solution; in alkaline solution
to periodate.
HC103 with excess of HI forms HC1 and I; with excess of HC10, HC1
and HIO:; .
6. Br forms I and HBr or a bromide.
356 HYDRIODIC ACID. §280, 6B, 7.
HBrO, with excess of HI forms HBr and I ; with excess of HBrO , Br
and HI03 .
7. HI03, iodine is liberated from both acids: HIO. -(- 5HI = 3I2 -j-
3H20 . HI04 gives iodine.
8. H202 becomes H20 , 0 and I (§244, GB6) (Cook, J. C., 1885, 47, 471).
9. 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, , SrI2 , MnI2 , AIL, , SnI4 , PbI2 . Agl and Hgl, . See Mitscherlich (POff.,
1833, 29, 193), Personne (C. r., 1802, 54, 21(5) and (Justavson (A., 1H73, 172, 173).
8. Detection. — The iodide is oxidized to free iodine by one of the re-
agents mentioned in (6) above. With a dry powder hot concentrated
H2S04 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 CS2 , or the bright-red color with CHC1 , .
In case a large amount of iodine be present the CS2 solution may be almost
black. In this case large dilution with CS2 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
oupric salt strongly acidulated with HC1 is an excellent reagent for the
oxidation : 2CuCl, + 2KI = 2CuCl + 2KC1 + I, .
If insoluble iodides are present they should be transposed by H,S ,
the insoluble sulphide removed by filtration, the excess of H,S removed
by boiling, and the solution then tested for hydriodic acid. Or the
insoluble iodide should be reduced by Zn and H,S04: 2AgI -f Zn + H,SO,
= 2Ag -f ZnS04 + SHI. The filtrate may then be tested for hydriodic
acid. The insoluble iodide may also be fused with Na.CO., , and after
digestion with water the filtrate acidulated and tested for hydriodic acid.
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 weighing as
such after gentle ignition. Volumetrically by oxidation to iodine and titratiou
with standard Na S.O. (Groger, Z. anf/cir.', 1894, 52).
*To test potassium bromide for traces of an iodide it is recommended to add C8, and cupric
sulphate or a small amount of ferric alum. Or add chlorine water and then a few crystals rf
ferrous sulphate ; then shake with CS, (Brito, C. If., 1884, 50, 210 .
$281, 6A. WDIC ACID. 357
§281. lodic acid. HIO.; = 175.858 .
H'IvO-"3 , H — 0 — I ^ £
1. Properties. — lodic acid is a white crystalline solid; its solution saturated
at 14° contains 68. 5 per cent HIO3 , and has a specific gravity of 2.1629 (Kaem-
merer, Poyy., 1869, 138, 390). At 170° it loses water, forming iodic anhydride,
I2O5 , 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(IO3)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.— («) By electrolyzing a solution of I or HI (Riche, C. r., 1858,
46, 348). (?>) 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 Ag2O . (e) By
adding water to IC13 and washing with alcohol: 2IC13 + 3H2O = HIO3 -f
r>HCl + IC1 . (d) KI03 is made by treating iodine with KOH: 31, + 6KOH =
KIO -j- 5KI + 3ELO . And then washing with alcohol to remove the KI . (e)
By heating potassium chlorate and iodine: 10KC103 + 6I2 + 6H2O = 6KHI2O,
-j- 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 AgN03: 5AgNO3 + 31, +
3H.O — 5AgI + 5HN03 + HIO3 .
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, (ft) By adding
a slight excess of H2SO4 to Ba(IO3)2 and removal of the excess of H2SO4 by
the careful addition of Ba(IO-)2 . (c) By boiling a solution of potassium
iodide with an excess of potassium permanganate in neutral or alkaline solu-
tion: KI + 2KMnO4 + H2O = KIO3 + 2KOH + 2MnO2 (Groger, Z. angew.,
1894, 13 and 52). (d) The very stable potassium biiodate, KHI,O,, , is formed by
recrystallizing a water solution of equal portions of KIO3 and HIO3 . 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,
84, 27; Spica, Gazzetta, 1894, 24, i, 91). AgI03 is soluble in 27,700 parts
of water at 25°; in 2.1 parts NH4OH (10 per cent) at 25° (separation from
silver iodide); in 1044.3 parts HN03 , sp. gr. 1.21 at 25° (Longi, Gazzetta,
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 : Cu , Al , Co , Ni , Mn , Zn , Sr , Ca , Mg , K and Ha .
They are all transposed by concentrated HNO., or H2S04; 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 K , Na , Ba and Ca iodates a
separation from iodides).
6. Reactions. — A. — With the metals and their compounds. — A few metals
358 IODIC ACID. $2&l,iiA,l.
are attacked evolving hydrogen, forming' iodates, sometimes traces of
iodides. With the following metallic compounds the valence of the metal
is changed:
J. As'" becomes Asv with liberation of iodine. AsH, in excess forms
As0 , with the HIO., in excess Asv (Ditte, .4., 1870, 156, 33G).
^. Sb'" becomes Sbv with liberation of iodine. SbH. forms Sb° .
3. Sn" becomes Snlv and HI .
4- Cu' becomes Cu" with liberation of iodine.
5. Fe" becomes Fe'" 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 iodate,
AglO, , white, crystalline, soluble in ammonium hydroxide, soluble in an
excess of hot HNO, . In the ammoniacal solution, hydrosulphuric acid
forms silver sulphide, sulphur and ammonium iodide.
Barium chloride precipitates barium iodate, Ba(IO )_, , 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 lead iodate, Pb(I03)2 . Ferric
chloride gives, in solutions not dilute, a yellowish-White precipitate of
ferric iodale, Fe(I03):! , sparingly soluble in water, and freely soluble in
excess of the reayent. Boiling decomposes it.
Alcohol precipitates potassium iodate from water solution, an approxi-
mate separation from iodide.
B. — With non-metals and their compounds.
/. H,C,04 becomes C0a and I . Action is slow unless solutions are hot.
Carbon (except diamond) heated in sealed tubes becomes C02 with sepa-
ration of I (Ditte, /. c.}.
H,Fe(CN), becomes H;iFe(CN)0 and I .
HCNS forms H.,S04 , I and some other products.
2. HN02 becomes HN03 and I .
3. PH3 becomes H,P04 and I . With an excess of PH8 , HI is formed.
Water in which phosphorus has stood reduces iodic acid to iodine (Corne,
J. Pharm., 1878, (4), 28, 386).
HH,P02 becomes H,P04 and I .
4. H2S becomes S and I . Thiosulphates form first iodine then an iodide.
§281, 9f. IODIC ACID. 359
H2S03 , with excess of HI03 , becomes H2S04 and I; with excess of H2S03 ,
H2S04 and HI .
5. HC1 ," if concentrated, forms IC13 and Cl , iodine not being liberated.
0. HBr forms Br and I .
?. HI forms I from both acids. The addition of tartaric acid to a mix-
ture of KI and KI03 is sufficient to give an immediate test for free iodine
with CS2 . It must be remembered that an iodide alone rendered acid will
give 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 §275, 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. H2S03 is
often employed because it acts rapidly and in the cold; but traces of HIO.,
frequently escape detection for the least excess of H.,S03 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 hydriodie acid. The following reducing agents have
been used : K4Fe(CN)fi acidulated with dilute H2S04 , H:!AsO., , CuCl ; FeS04 ,
morphine sulphate and uric acirl. To detect KIO, in KI it is recom-
mended by Schering (J. f., 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. (7., 1870, 29, 442) directs to add equal
volumes of water, carbon disulphide, and then coarse zinc filings. It may
be necessary to warm the solution slightly. Biltz (C. C., 1877, 86) dilutes
the HNO., with water, adds starch solution and then H2S solution drop
"by drop. A blue zone is -formed if HIO, "be present.
0. Estimation. — (a) Bv precipitation with A#N(X, , and after drying at 100°
weighing as AgIO3. (1>) By reducing to an iodide and estimating as such,
(c) By treating with KI noidnlated with H2SO4 , and titrating the iodine lib-
erated with standard Na,S2O3 .
360 PERIODIC ACID. £282.
§282. Periodic acid. HI04 = 191.858 .
H H H
N I /
0 000
ii si/
H'IVII0-"4 or -R'rr"0-"r , H — 0 — 1 = 0 orH — 0 — I — 0 — H.
The anhydride, I..O- , has not been isolated, and but one acid is known in the
free condition. HIO..2H.O or H^IO,, . This acid exists in colorless monoclinic
crystals, which do not lose water at 100°. It melts at 133°, and at n
higher temperature it decomposes into iodic anhydride, water and oxygen
(Kimmins, ./. ('.. 1SS7, 51. ::."><>; and 1889, 55. 148). Numerous periodates have
lieen prepared as if derived from one or the other following named acids:
IH04 , H,IO, . HIOa, H4I,0,. , H.r.0,, , H.,1,0,,, H10I4OI9 , H10I.O,e
(Rammelsberg, P<>!H>., 1805, 134, 3f>8, 499).
The free periodic acid, Hr,IOu . is prepared. («) Ry oxidising iodine with per-
chloric acid: L'HCIO, + I, + 4H,O = 2H-,IO« + Cl. (Kaemmerer, Pogff., 18G9,
138. 106). (l>) }\y heating iodine or barium iodide with a mixture of barium
oxide and barium peroxide, digesting with water, and transposing the
Ba iIO, ) thus obtained with the calculated amount of sulphuric acid (Ram-
melsberg, /'«//.'/., 1S<>9, 137, .105). (r) By conducting chlorine into sodium iodate
in presence of sodium hydroxide: NalO, + :!NaOH + C12 = Na2H,IO( -f-
2NaCl . This acid pcriodate dissolved in water with a little nitric acid and
then precipitated with silver nitrate, forms the silver salt, Ag,H;IO,. . This
precipitate is dissolved in nitric acid and evaporated on the water-bath, when
orange-colored crystals of silver meta periodate are formed according to the
following: 2Ag,H,XO, + 2HNO, = 2AgIO4 + 2AgNOs + 4H.O . Water decom-
poses this precipitate: 2AgI04 + 4ELO = Hr,IOB + Ag.H^IOg . Or the silver
periodate, AglO, . is decomposed by Cl or Br (Kaemmerer, /. c., p. 390).
The silver salts vary in color: AglO, is orange: Ag HIO , dark brown;
Ag,r.O,, , chocolate colored: while silver iodate is white (a distinction). In the
general reactions periodic acid and periodates resemble iodic acid and iodatea.
H.C.O, becomes CO, and I.
HtPO, becomes HaPO, and HI .
H.S becomes S and HI .
H.SO.T becomes H..SO, and HIO, without separation of iodine when the two
acids are present in molecular proportions. The presence of a greater pro-
portion of H SO causes, first, separation of iodine with final complete reduc-
tion to HI (Selmoua, R., 1888, 21, 230):
HIO, + H,SOS = HI03 + H,S04
:',HIO, + SH..SO, = HIO, + I* + 8H,S04 + H2O
2HI04 + 7HSSO, = I, + 7H2SO4 + H20
HI04 + 1H.SO, = HI + 4H2SO,
HC1 becomes Cl and IC1,
HI forms I from both acids.
According to Lautsch (</. pr., 1867, 1OO, 86), its behaviop' with mercurotia
nitrate is characteristic. The pentasodic periodate, NasIO, , gives a light-
yellow precipitate, Hg-.IO,. .
!i283.
COMPARATIVE REACTIONS OF HALOGEN COMPOUNDS.
3C1
eo
r-t
coo
U
*o
OB
T3
•i-t
O
•§ -
CB ^
«
* I
8 *
eo
oo
2
S
3
,
«r
a
CD
a
e
5
3
•§
DO
0)
03
jj:
[3.
a
a
o«3
•O i
y
'3
Q
'3
oS ••
0
2
2
2
•*5
*5B
a
ft
ft
a
C'
K
0
o
0
_* //-
^
S ua"
B
Sr
5r
0
5
B B
1
V
M
3
•S
5
•S
W A
4
^
5
®
^
3
Jr
• -0
e
M
O
M
sc
1
S
O
B
ft
'3
2
£
a
•— i
O
2
£
o"
5i w
08 0
PQ y
o
B
0
fl
S c
•
r a s
5
§ §
.
j
S
3
3
«
3
1
00
M
Broma
o precipi
ft
'3
2
a
o
nstable.
a
3
0
pa
rand Cl.
e
rt * v3
M fa i.O
faB pa pa8?.
K
P
B
PS
PB 8 8B
a
CD
•
-S
S
S
•o
M a
U *"
S
a
:§
5
3
S
B «
"a
p,
a
^
*O
z
Xo preci
No preci
No preci
«5
F«
B
N
9)
Ml
0
« 1|
3 S us
8 B 8
1
y>
•d
"^
3 •
2
CD
o
3
•d
i
fi
3
A
*
""
S
^
1
a
3
0
H
•8
CD
^
•7^
^
'a
D
i
M
e
.9
T3
«
V
2
SB
|
8
3
a
Mt
a
CD
1"?
' 1
43
O
to
o
te
M
s
H
J -a
M +J
+3
<
co
B
fe
»P4
8
B -
M - M"
5
<D
3
1
8
3
1
3
J3
<3
a:
g
2
£
1
•5?
a
'3
A
s
.5
E
of
1
91
e
fa
pa
03
y
_ft
'3
2
0.
Dissoh
2
ft
o
CD
ft
O
S5
aa
Ml
B
i«
8
2 *•'
B pa
b
pa
CD
3
S
£
3
fl
v
6
d
"o
«
2
*'
S.
e
*
1
•o
^
00
j
O
Mi
•r-t
O
2
&
S
5
Ml
B
8
0
o
3
2
3
S~"
3
S.
— *
00
O
h
8
€ 2 •
00
_>
3
<H
M
0
09
2- 1 ®
B s yi,
AgXOa, in excel
A«NO3t excess o
lion tested.
NH4OH. to the
cipitate.
00
®
5
H
0)
c
a.
HjjClj, excess o
tion tested.
UgNO3, excess.
B"
J3
5* S*
a s
S U
•3
O
B"
u •
B S
3 2 B
? ! i
1 o' '^
p 5 M !«
fa "^ S f^ s* "
• SB 0 M IK •
PART IV.-SYSTEMATIC EXAMINATIONS.
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. (>'r/); :td, by oxidation with nitric acid in
presence of sulphuric acid, at a final temperature of the boiling point of the
latter (§79. (iV.J): -1th. by solvents of certain classes of organic substances;
"ith, by dialysis. These operations are conducted as follows:
§285. ('ombiixtioH at <i red or trhite 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, '.'>). If chlorides are present
some iron will be lost at temperatures above 1()0°. and potassium and sodium
waste notably at n white heat, and slightly at a full red heat. Certain acids
will be expelled, and oxidi/.ing 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 .ill sub-
stances. Complete combustion, without use of oxidi/ing 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 nf 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 mny be added, as a solid, in the same
way.
§286. O.ridalion irifli i>otanxinni chlorate and liudmcliloric acid on tlic irntrr-bath
does not wholly remove organic matter, but so far disintegrates and changes
it that the nitrate will give the group precipitates, pure enough for most tests.
It does not vaporize any bases but ammonium. l>ut of course oxidizes or
chlorinates all constituents. It is especially applicable to viscid liquids: it may
be followed by evaporation to dry ness 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. 11 is then evaporated to remove chlorine,
diluted and filtered. If potassium and chlorine arc to be tested for. another
portion may be treated with 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. PItEIJM I \.\lfY K\ A FIXATIONS OF SOLIDS. 363
5:287. The action of sulphuric irith nitric acid at a ynidufiUjf 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 iii 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 vapori/.e chlorides. The material is now transferred to a platinum dish
and heated until the 8>ilphiiric (/<•!<! bei/ins to vaporize. Then add small portions
of nitric acid, at intervals, until the liquid ceases to darken by digestion, after
a portion of nitric acid is expelled. Finally, evaporate off the sulphuric acid,
using the lowest possible heat at the close.
§288. The Koli-cnts 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 sohitions 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. Jin Diali/tsix, 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 iipon the kind of organic material. Thus albuminoid compounds are
almost fully rejected: but saccharine compounds pass through the membrane
<juite 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 flume of the candle, or of the nas-ict, burning under ordinary circum-
stances, consists' of three distinct parts: a dark nucleus or zone in the centre,
surrounding the wick, consisting of un burnt 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 unconstimed hydrocarbons, is called
the reducing- flame.
§292. The flame produced lit/ the liloir-]>ipe (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.
Jn order to produce a good oxidizing flame, the jet of the blow-pipe is placed
ji7Pt 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 PRELIMINARY EXAUIXATIOXS OF SOLIDS. Jj293.
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 flame 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.-
\vhat 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 Glass 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 paper,
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.
Ry 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 us
to their alkalinity or acidity, by small strips of moist red and blue litmus
paper inserted in the neck of the tube.
/
Heat in Glass 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 arc very characteristic;
these metals, if present, are generally easilv detected in this way (§69, 7).
Heat in Blow-pipe Flame on Charcoal.
§296. For this test, a well-burned piece of charcoal is select ed. 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 I.,- :i
powder, it may be moistened with a drop of water. The coal is held horizont-
ally, and the flame made to play upon the JISSMV at an angle of about twenty-
five degrees. The substance is brought to a moderate heat, and finally t.o
intense ignition. Any escaping gases are to be tested for their odor: the
£300. PRELIMINARY EXAMINATIONS OF SOLIDS. 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,
dome 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
eobalt 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, as
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 Bimsen burner alone will suffice; the substance
being heated in the loop of platinum wire, which, in all cases, should be first
dipped in hydrochloric acid and ignited, in order to secure against the presemv
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 mor»-
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
exces"s, 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 Microcosmic Salt.
§300. This is best effected in the loop of platinum wire. This is heated and
dipped into the borax or mieroeosmic salt and heated to a colorless bead; n.
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 tlu-
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 CONVERSION Of SOLIDS fXTO LIQUIDS. §301.
-mall quantitj- will, in most cases, suffice. Certain bodies, as the alkaline
earths, dissolve in borax, forming1 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, they become opaque. This operation is called flaming.
As reducing agents, certain metals are employed in the bead of borax or
microcosmic salt. For this purpose tin is generally chosen, letnl ;md silver
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 KO
obtained. A number of these are now taken on charcoal and fused into a large
)>ead. which is charged with the substance to be tested, and then with the tin
<>r other metal. For this purpose tin foil (or lead foil) is previously cut in
•-trips half an inch wide, and the strips rolled into rods. The end of the rod
s touched to the hot bead to obtain as much of the metal as required. Lead
may be added as precipitated lead (" proof-lead "), and silver as precipitated
silver. My aid of tin in the bead, 'cuprous oxide, ferrous oxide and metallic
.-mtiinom are obtained and other reductions effected, as directed in §77, 7,
;ind elsewhere.
COX VERSION OK SOLIDS INTO LIQUIDS.
§301. Me fore the fluid reagents can be applied, solids must be reduced to
liquids. To obtain a complete solution, the following steps must be observed:
Firxt. 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 clear
liquid poured oil' or separated by filtration. A «/»•»>/* or ticv evaporated on glass,
• «r 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
washed with hot water.
Si'fiiiiil. The residue, insoluble in water, is digested some time with hot
hydrochloric acid. (Observe §305.) The solid, if any remain, is separated by
*i It ration and washed, first with a little of this acid, then with water. The
solution, with the washings, is reserved.
Ttiiril. The well-washed residue is next digested with hot nitric acid.
Observe it' then1 are vapors of nitrogen oxides, indicating that a metal or other
body is being oxidi/.ed. Observe if sulphur separates. 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. T?nt if a first-group
base be present. HNO, should be added first, for HC1 would form an insoluble
chloride. If the substance contain tin (especially an alloy of tin) HNO,
would form insoluble metastannic acid. H,0Sn,0,- , in which case HC1 should
lie 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.
Fiftli. 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 muxl In' fnxetl irith n /i.red alkali eiirtmmite. when the con-
stituents are transposed in such manner as to render them soluble. Tin-
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, mid
analyzed separately. The separate remit.*, in nnnii/ ease*. Indicate fix- orhjinal
• •• >itfiiii(tti<»i nf ffich metal.
§303. TREATMENT OF A METAL OR A2V" ALLOY. 367
CONVERSION OF SOLUTIONS INTO SOLIDS.
§302. 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 clone 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 in their detection. Thus:
Gold 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.
GSb 4- 10HNO, = ;:8bsO, + 10NO + 5H20
15Sn + 20HNO. + 5H20 = aH10Sn5O13 + 20NO
All the other metals are oxidized and converted into compounds that dissolve
either in water or an excess of the acid; c. y.:
3Pb -f 8HNO:1 = ,lPb(NO3)2 + 2ND + 4H2O
Bi + 4HNO, = Bi(NO:!) , + NO + 2H20
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. J
» This section is furnished by Dr. J. W. Baird, Dean of the Massachusetts College of
Pharmacy.
1 When gold or platinum constitutes more than one-quarter of the alloy, nitric acid fails to
extract the whole of the base metals that otherwise are readily soluble. In such a case the
amount of gold or platinum must be reduced to at least 25 per cent, by fusing the alloy with
the requisite amount of that base metal whose absence is surely known.
J Arsenic, if present in the original alloy, now exists in the form of arsenic acid, the precipi-
tation of which requires heat and long-continued passage of H2S (§69, 6' e 7).
3C8 SEPARATION OF A(.'II>K FROM BASES. £303, B.
B. — The Residue Insoluble in Xilric Acid.
This may contain gold and platinum in their metallic forms, and tin * and
antimony * in the form of metastannic and antimonic acids. 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 SnClt
and SbCl .
Digest, therefore, the well-washed residue in concentrated hydrochloric acid
at a boiling temperature for from ."> 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 tint, again add hydro-
chloric acid, heat, and again ilccnnt.
The Hydrochloric Acid Solution,
This solution may have a turbid appearance, especially when cold, due to the
action of the water upon the SbCl : but without filtering proceed with the
separation and detection of the tin and antimony by the usual process. t
The Dark-colored Hettidiie.
Add, after washing, two volumes of hydrochloric and one of nitric acid:
evaporate almost or quite to dry ness, 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 nqua-regia formed, and
BOW exist as auric and platinic chlorides.
Firxl Portion — Test for Gold.
Dilute with at least ten times its bulk of water; add a drop or two of a mix-
ture of staunous and stannic chlorides; a purple or brownish-red precipitate
(or coloration), purple of Cassius, constitutes the test for gold.
A convenient wav of preparing this mixture of stannous and stannic chlorides
is to
(a) Add a few drops of chlorine-water to a solution of staunous chloride; or
(It) Add to a small quantity of stannous chloride enough ferric chloric!*' to
produce a faint coloration.
Second I'orlhn — Text for Platinum.
Add, without dilution, au equal volume of a strong solution of ammonium
chloride. The formation, either at tirst or on standing, of a lemon-yellow
crystalline precipitate, consisting of the double chloride of platinum :I-M|
ammonium, (NH4Cl)3PtCl, . 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 chloridf
remains behind as a yellow crystalline powder.
SEPARATION 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 Hliojr
contains arsenic, part of it will combine with the antimony and tin, and be held in the residue.
£309. SEPARATION OF ACIDS FROM BASES. 369
decrepitation caused by water in crystals) indicates chlorates, nitrates, bro-
mates, iodates. Explosion or deflagration will occur if these, or other oxygen-
furnishing' salts — as permanganates, chroinates — are in mixture with easily
combustible matter (§273, 7). Hypophosphites, heated alone, deflagrate in-
tensely. A brotniish-ijelloic capo? indicates nitrates or nitrites (§241, 7); a
green flame, borates (§221, 7). The odur of burn-ing sulphur: sulphides, sulphites,
thiosulphates, or free sulphur. The separation of carbon Mack: 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 Avet way, indications
of the more volatile acids will be obtained. Sudden effervescence: a carbonate
(oxalate or cyanate, .§228, (i). Greenish-uello-ic vapors: a chlorate (§272).
Brownish-yellOU), 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? Thus, if
barium be among the bases, we need not look for sulphuric acid, 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 b]/ transposition with 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 antimon}/ is among the bases, they may all be
removed by boiling with slight excess of sodium or potassium carbonate, and
filtering. Arsenic 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 add 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 work
of the third group of metals, and is explained in §§152 and 153. For removal
•f boric acid, sec §221; oxalic acid. §151: and silicic acid. §249, G 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
PRELIMINARY EXAMINATION OF SOLIDS.
^310.
SOLIDS.
^
•o
5
-3
S
•2
3
o
So
0
u
•f.
tl
"u
£
a
t
_£
•« "5 2
M|
? | 5
h
!
h
1
.5
S?
1
•8
•
1
t*>
e
—
S
again bj- vaporization.
_o
.u
sublimate.
'loic ichile hot, ichite irhen culd.
Hon.
§
•0*
•H
a
•
o
-ye/low «f he/* co/d
blimates.
** —
=
- -
TB ii *
•
IS
a
f
41
a-. £
-3
« —
QQ M>
XAMJXATIOX 0
s Sfructure, Grarili
if/'1.-
ies (including combi
be
.5
C3
Oi
^3
0
u
^
*c
5
il
S
*i
5 ^=
? J=
C .be
C .
t s
^ ~
t •£
b -i
•?• ^
^ « p *
* * "* ~
n S — -s
>.5 si
numerous other sali
heated, becoming sol
ially or wholly:
combination, or abso
4-
o/
0
s
A
I
o
z
rubbed to globules,
forms »r//i/e crystalli
. forms a stiblimate,
. turning red on tritu
: an r lie odor.
' octahedral crystals
2 '•«
1 |
•« 0
- 1
— aT
5 "^
•* tn
IX,
« i 5
J ~
•J. ^
= ^ -o -
•0
c
"C .:
be
^' c
tc a,
c *•
~ .c:
£' ^
r"
"S t —
3: ^
"" •;<
» j
^ "v
H -5 ° 1
c
*3
at
£S,
&|
s
'* Ji
a —
._ c-
s s
S •*
P=
<
V. ^ ~
ar *"
c ^
~ X
a 5
« 5
Z x S3 -~
5
&
» C8
V N
"uo
S
0) •*-•
5* ,*
1 1
3 -S
O •-"
i §
K
t— i
S
t— t
*-3
W
«
PU
0
1— 1
"c
lie xiibMtam-c miffcn
Absenco of vela
c -5
S *
°s
~ —
0 0
0 9
fl f5
N -
S ~ S *
T X t3 5
> 55 r* *"
o o" o 0-
.4' «' -a a
W PM O OT
/i« substance fuses:
Most alkali sal
crystallizatiol
/«e substance xubliiii
H,O of crystalli
Sublimate cond
oos ;,'
bo bo
W W
w w
As . steel-gray si
As..0a sublimes
cd —
§ 2i.
1 i
^ £
> «C
09 00
"3 rf?
««5 CQ
CO
tit
g«
«^
h
E*
etn
ft.
VD
**
N
CO
-r
•«x
o
>> >>
rt
fe;
|B
TJ
oS
1 j
£*
«a
O QJ
^
P< N
H ^
C
S
4J
>
I
jj *-«
cd v
2"~ 3
1-1 f)
P.
C) o
tt
H
55310. J'KELIMIXARY EXAMINATION OF SOLIDS. 371
>7, 7).
0-
i drops, yellow when solidified.
;e, iodate or peroxide. A small
heated,
blackens lead-paper. Kecognized
etc. Recognized by its odor and
S
•2
1
0)
-
"3
5
3
-3
O
«0
Intensely poisonous.
•oicn, acrid vapor.
on I. are modified by oxidation.
&.U and Pt .
1 action on litmus-paper.
CD
•fH
m
2
'o
a
£
fi
m
T-l
T-l
••"••
In
'O
a
i
in
P
O
A
t)
"5
tion is avoided.)
U
"li
Ol
a
at
be
^^ *• W •
to t* aj
"S be
•>
v
• ~^
5* 3
*
**
w
3
NH4 salts, those not decomposing, white sublimate (§!
FeCl3 slowly sublimes as a reddish-yellow stain (§126,
S , free or by reduction of sulphide, gives reddish-brov
H2C2O4 , a heavy white vapor and crystalline sublimati
I , a violet vapor and blue-black sublimate.
he substance evolves a gas or rapor:
O indicates the presence of a nitrate, chlorate, brom:
piece of coal placed upon the assay glows upon bein
H2S , from hydrated sulphides and some sulphites,
by its odor.
S02 , from sulphites, thiosulphates, certain sulphates
bleaching effect.
NH3 , from its compounds which decompose, characte
litmus.
CN , recognized by characteristic odor and violet flame
Oxides of Nitrogen, from nitrates or nitrites, reddish-
Acetone, from acetates, characteristic fragrant odor.
:ain of the changes stated above as occurring in opera
Oxides are obtained from metals, except from Ag ,
S and sulphides yield SO, . Eecognized by its odor ai
As yields a sublimate of As203 .
Sb yields a sublimate (white), of Sb,03 and Sb203 .
Bi , a sublimate, dark-brown while hot, lemon-yellow
Te, gray sublimate of tellurous anhydride (Te02).
Se and selenidcs evolve Se02 , odor resembling that of
Hg , sublimate of metallic mercury.
'he substance decrepitates:
Crystals as NaCl. (If finely pulverized, the decrepit
'he substance deflagrates:
Nitrates, Chlorates, lodates, Hypophosphites, Pevma
e«<
t.
t-i
i^
0)
10
w
—
N
<O 10
, ,
,0 rd
js ^
3 fl
r2 A
•*
pq o
W £
^
0 0
$ *
5 ®
s
CD"
0)
_» ^
C3
<& •
.2 S
an
,
.3 *>
4)
rt 03
"ri ^
"3
-u « CM
o
« *P4
o
W
B
372
PRELIMINARY EXAMINATION OF SOLIDS.
§310.
the charcoal:
alkaline earths.
cen/:
Ikaline to test-paper,
hot), not alkaline to test-paper.
charcoal :
0)
-M
2
1
f
G
-
C
i— (
"o
u
c
V
if
(L
emon-yellow when cold,
old, greenish-white flame.
\v flame.
9
E
jt
0)
P
-^
=
"3
•—
1
1
V
1
e
-
"o
u
c
II
j»> Q
*} ..i
<j B ti »
c.
°
-E
^
1
'^stance fuses, and is absorbed 1
ts of alkalis and some salts
'tstancc is infusible and incand
, Sr , Ca , Mg, the residue is
03 , MgO , ZnO (yellow whili
tstancr forms an incrustation o
' 0)
£ B
"- es
w ^
< X
t. •-
VI
0 V
~ 5
o
- ;. -
*= .C >>
^* fe £C
s « 5
t .t; rs
|1 «
^ « "p
5 1^ 6
w is i
III
1
ce
:
r.
V
—
-'
a
V
-' ^
6 ^'
a a
«c «
c £
6, -~
& t
f 8
"n £
&^
C —
t< •£-
«s or incrustation is colored:
) , yelloicixh-grccn.
) , bluish-green.
O, , dirty dark-green.
O, , SiO, , phosphates, &!u«.
it«
_c
*Q.
O
•- e
S -e |
o j, S
15Q
^13
c grains are obtained:
Sb , brittle.
, Cu , Sn , Au , malleable.
X
0
b
L
5 0
!?•
!^
= Oj r-T 3
fi"
8 N 0
<
-a fl
W 03
1 §
to
\^
B^
to° 0
|Sfi
•C!
« -=
iS
^
w
t. &,
K
i
^
CO
•* 0
H
H
M
C 1C
1
O ff
fl *
° fl
M ^~*
3
-^
^
tfj
fl V
o •»->
^
g ® I
o> °
o
*3 °
'3
" 5 i5
§ G
(fl
o is
j?
^5 fl *
w .2
•rt
to «
>,
1H Q)
M i
^
5 2
0 0)
ill
^ g ^
00
a
|S
"
13
10 ^
lls
M
&
>
$310.
PRELIMINARY EXAMINATION OF SOLIDS.
373
pG
O
B
7J
u
a>
a
C3
O
o
P-H
^3
13
a;
O>
to
-<->
.S
c
O
c
X
II
o
a
to
CO
a*
'3
03
O
*C
and silicates, give bluish-
ities of sodium compounds,
the reaction. In presence
;=
3
cs
Tr,
c
-S
0
S
00
00
much Ba (§187, 7).
Sodium interferes with the
and its compounds.
o
-M
CS
w
'3
_0
"3
CO
13
a
CS
V
w
C
'3
a
Hi
cs
p5
3
SI a «
CU M V
-1-1 cS '•%
">
too
^
0
o
a
i
i
9
.5
f
co~
£
s
13
2 Z«
PH ^ B
CO S
0 K 0
A S ^2
* s *
"o
2
1c
a)
|
cS
13
u
13 P
|o"
03 *-•
PI cos
o"
CO
addition
pd
-*->
13"
V
"cB
CO
la.
o
O
ju
"3
CD
.. a
s v
§- £
. C IH
"» -U
0 0
^> M
,2 ^
v ^
~ «rH
00 1-^
~ ^~
t-H
a)
'3
i
O C
O) .«
a g
"fl 0?
a) ^
to c3
PH CO
M -2
.S 13
PH C
S es
0) OJ
£ s
p2 5
~
"H
c
ef
0
CO
^.'
ag
a
a>
a>
01
<y
O
a green glass, appears orange-yellow; moistene
delicate (§206, 7).
Violet: TS. and most of its salts, except borates,
violet flame, distinguished in presence of very
Excess of the lattsr prevents the reaction; Li
of sodium, the potassium flame appears reddi
O
11
>i
cs
13
m
13
i
s
o
PH
a
O
^ O
I- 03
^^ ro
« R
co o
CO W
03 ..
'&'§
Sr and many of its salts yield a crimson flame,
Li and its salts produce a carmine-red flame i
reaction; potassium does not.
Green: Yellowish-green, Ba and most of its salts
Emerald-green, Cu and most of its compounds.
Bluish-green, B203 .
Yellowish-green, B,O3 , best obtained by the
A
M
V
CO
2
'3
CS
^o
"PH
P
,c
A
3
CO
a;
r=3
-»->
"-*-!
fl
P ^
0) t-
PH
'S *-•
£ Oi
a c°n
p ^
•S.-?
rt '5
-a-sp
Whitish-green, Zn .
Blue: Light blue, As and many arsenic compound
Azure-blue, Pb , Se . Also, CuCl2 .
Greenish-blue, Sb , CuBr2 .
K. ®
PI «
d
<C -H
|
i
^
S t
^
a) o
43 S
"P<
.
-u oT
pu
•
,s|
_ 'S
0
«
H O
wn
*o
O o
p.
^U
"S .2*
3 A
o
.rt
CU 'Si
M 0
,2
fj
M
371
SVBSTAM'EN BKFOKK 'HIE III.OW-PIPE.
§311.
J 41 4
.
ja
0
g C § 49
5
• "C
t/
if .^ -~
_B
v R
S Hi
1
^ ~ ^
*»
g^-w
<— s
E-1 * -2 ^~
.
te
? a
i
£
""^ w ^ ®
'M'
t
3C —
§_
*
^ jfe ""* js
t
•J!
x' T «>
0
i/i «i
zz • .p*
r 3 V tk
HH ° ft
o
6
•
to
6
M
i
E
3
a>
h
K Ts
<u w
B«
C3 c
(A C
<-J ' "55 v. cS
§
.2
^1 L:0
H
6
H
c
L *•
'/
Sg
o %^ i 5
"S
5
S3: • "•
!«.•*•*
«
is
0
*
og
S
C£ 2 r «>
~- *» (^
"w
c
6
^,
tu £ *» 5*
0
.2
U) M
a
••* *
5.
2- . H
•
C5 <*•• •
2 ~~ «B
o
°n
i « 4>
bM
6
0
g 3 J ?
|
•5
'i
M
•"^ ••-<
3
£
'^
w .g.'S 2
^
u
^,
e *:
3
.*>
6
f
&
^H ^^ ^o O
QJ ^_ 00 ' ^r
• *&
-,
•I
O •
m
•7 s t
"s
^ 1 i°
n §
o> -^
^:S
^
53
o •* % *f
*
« < &
i r •
*:c
«^
•i to
L:«:
_3 ~ ••> 0
^. -
c
.i
g
Sg 111
ionium
f
17
E
o
4
.1^
m
S ^ "5
B
1
^=M
^
•o
o ^ ^ "T
|
1
L^.
•I
°^
w ^ ^ 'S
ll
^
~
1 tfl 0
S
H
6
6
x>"
"§ §*
•5
i*^*
(SS
5S
H
"&
u P
p<
fT^ Q "*^ CJ C
C/3 2
* 5! *
Q
bo
8 H £ £ ^ *
-M'S
*
W<1 0
•e to
o«:
to
*s C
to:«
^~\ -^^ I ^ <
H 2 e
S =S '§ «•• o ,
"a a
S|
to
9« ^3
6
0
o
po -^ ^-_ ^5 — -
5 fci.
,*" "^ ^J
fH
C/3
o
•^
1« .a
0)
Q
6
T*
i
Ji, **— -< 0 -^
o
^
'r ^' 'CxT
P
m
V
O ^ "^3 —
i
?
y rf 1-0 ^
0)
i
o
PH g " >Si *f
5
?
8 O o
••
d
^
2 1 "2 ' 1 S
•*H
•5
'? g1 x c "5
OD .
n
M
6
a" . ^
K^ O S ^ ^^
Q
-i, ac S
ac £5
oo
"• O
u ^ u
"*! "t" 'g &; *
e
>— ' ..
^^
c
•:O •:
W '^ c ^ >>
^
00 ^1 H
i«: to
*:'
•e:
« to
i«J to
0>
pq § ^ S c
£
j.
^=
•*-•'
-B
^ « ^S i to
co -^i
009
o
s<
d
a
%
j.
o
"o
Yellow
or
rovvnis
« >j
%$*
r <u
^ S
aJ
a
S
d
o>
I
b
O
ll
§
o
«
<;
°°
URQUMNQ OF THE MKTA.L8.
375
eo
CO
o
bfl
2
CJ
04
i— l
CO
S
§
a
£ s
>> < „•
.c t N ^*
s
C3
C8 S
5 rf «
0
£
c
3
«
2
.5
^? e
5 "SS1 tt5g
g
S
t
I
M
2
««
4) T3
bC "C
ft" ts^'Oi® ?5 *
18 f **' * Sl*i«s
£
9
03
o ^* "Jn P c*3 h ^ W fl Cl C^B C*l
•o
c
S
^
p ^
go at' 5 8^3"^ g ^
CS
V
2
R
te
1
8
s
?|
c*5 ft -2 t; ^ .D '-''S a 3 a
o w * a 5 « « o § g §
e
J
c
+-
•pa
g
"S
e
gj" O
^ PH t» O Q M 2
OK)
A
IQ
S
& <W
Q1 _«««•«
1
s
H
•js
i
os
S
(B
c o
f fl
b O
O jj p.'^S^"*^'"^^
s
s
p
8
"3
8t
^ »- A ^ * A '
«• flT •• ILL •* •* ZT •
<a
"JJ
•»•*
_n
c3
Q
QJ vt .,H »^ ^^ GB-
a
M
a
3
0
0
0
<a
g JJ1
ife § «* | » ^0;
a
c
±|
Q
i.
S
•S o
»
S
bo
m
•iH
n
^
«
8 H
tf'S^jj^; g IS SB*.
•s
e
•c
•§
text a;
«
if
5
•d"
D
S
O
ammo
filter.
l£«" tf JSl«f§l*
S -^ d *• 3-rj fl j .2» a^
i.£| E |I.Ula|i
1
'o
£
"S
ri
H
PH P9 W 0 S
3
-w
cS
03
C
3
"3
"o
<H
ihiiric ;•
.tate is fi
j a
c*
M
^ ••e »
e * S ess
s-'fl1
of
ft
3
g
Q
33
'I
S
M
"s
& I g S
3 •§ °
1 21 *
-
O
03
41
M
$
7j
£
•
0
g 'ft ^3
ft '3 ^
M
es ^ . "
1
§
C
O
A
"5
IH £
B A
H —
e«=
%tn 9 '• **
K
5
+2
fl
m
a
1
0
«'
rtl c
m 3
-1 «
0
i
^
€»
|« 6? *»
15 §g Is
•c
c
ce
3
fe"
£
Cu
0
fc5 S N
•o
-2
<H
Op
_x
. *
i
S
2
1
c3
i
g|
e
w .
^ O
H
S
!«'
I
0)
*-S
*
w
o
43
5
s5:
C
-^ p,
n w-s «s
— _ w_s «-g
<H
O
—
3
2 <§
C
I s
w
s i^
« "I
a
«
fl
•
I
foil the fill
hydroxi
f phospha
Prec:
Third G
luminum,
\l (Ml
rS ~ Q '^y.O
& P it G **
"3
1
1
§§154 to 1
«
g
S
O
H
«
•
s
«
1
Mi
ft.-s E
s
1
*aS§§'«.i .oco
SB S .g.2 .2 ,* w o • <0^ ^ ^
0 |
n
p»
Pass hydrosulph
Precipitate :
Second Group.
I Separate the precipit
sulphides into two
tions by warm digei
with yellow am
nium sulphide.
•
c*
9
%
ID
> -a
d
"5 fe Os •* s s *ji »-o . a
* ^o • flS ^ *8 3
1* A * M J* ^ af cB* w
jS* « S! L» a tt « l»
,^ Q(t » ^^ 5 c^ * C^B
^9 * " d^ O "•• t ""
C»p« '""nE ** >>"'j3
^95 .S'a5» 'SO.Sfc'oS «B«P<.2a
OH OP^S gijco
Copper, §§77, 1O2.
CuS b
Cadmium, §§78, 102,
CdS j/e
Rare metals, §§1O4 to
2
£
*"* ^
S i»
,—
1 E
® ft . s *
-
00
•8 S W ad
g . ^
^
sow »
• c
"3
ft ** C^
n
B
rH r(j l> ^^
O
i & "b
^ *
M
—
vy" PH M
03
£ 2 "?U 3
»
LJ" CJ !C
3
±-0,80
S yjMr
O
^^
S?
in ce
S
376
TABLE FOR THE SEPARATION OF Till; METALS.
55313.
§313. TABLE FOR HE VIEW 01
\ / \
f H,SO« = PbS04 White.
Pb
a.
3
PbCI, £ i
«
Pbci. ]HaS =PbS Black'
| K2Cr04=PbCr04 Yellow.
0
I K i = Phi., Yellow.
Hg
1
HgCI |
HgCI 1 q NH HgCI Hg Hlack.
•3
r Z {
Ag
AgCI
AgCI j | |(NH,i3(AflCI).,|-Add HN03]AgCI White.
As
As,S3
i
2
As,S,
\ *^ '
H3As04
\ ' "
"3
—
Ar
As,S.
NH,. AsS,
i
As S
OS
•
2
- —
Sbv
Sb.S
(NH«),SbS4
1 I
Sb,S,
SbCi
12.
*^ a
Sb'
Sb.S
0
•3
: o
P ~
9
c-a
p
Sniv
a.
SnS,
I
NH, SnS
E i-
•O
SnS,
5
SnC'4
|J
Sn'
O
SnS
"3
z
fl *
t, §
x
1*
Au'"
•
1
u,S,
Solution.
3 5
AuaS3
9
AuCI,
S"S(
3
K
O
PtS,
0
Solution.
•3 5
PtS,
§~
PtCI4
— - <
00 ^
0
PJiv
O
B
•3
*O
— O
S 3
C
: Mo"
c
MoS
S
•z
.NH, ..MoS,
MOS:,
-^ MoCI«
H|"
S
_c
* I \ / ~\ i_™
1
8
, Hg"
S ^
HgS
2. i i
•5 . / HgS j HgS Dissolve in nitrohydrochlori
! po
_o ~"
PbS
2 |
PbS
^o
Pb(N03), ^ -j PbSO, Confirm 1».
3
n Bi
_3
B
us,
< "
Bi,S,
= x
*i O
Bi'NO, . :7.o BiNO.i,
X
O .
2 Cu
•3
1
a
CM S
||
Cu,S
tn •-
S^
Cu NO t 2x Cu NO i
*^^ ^J
f an
xf
Z i
* ^
| Cu
2
•3 *
CuS
I
CuS
S"3
2||
— x
•3 ®
•< "
I Cd
1
,& 0
1 5
C
/
\ >
Cd'NO,),
\
<5S
Cd(N03),
i
£
3. Q /
. / \
-. ..
•3
•-T +•» /
a 5 T
) 2 ^
AKOHi,
J NaAlO, Add NH4CI or ivn.l.-r urid
; Al
j-
~
1 * *
fl
I Cr-
•3
o
\
II
Cr(OH),
!1
Cr(OH),
Fuso on platinum )
K
i
N-i
Crvl
'•
~ -3
i
'Sz
foil with KNO,
ni it
5 Fe"
§ !
j U si
FelOHi,
••
Fe OH',
and Na.CO J
Ft
: Fe"'
3 'C
i su \ / \ /
i
3 3
:
/
5 Co
1 1
3
—
s
» o -a
' , ,
° - ? a.
CoS
"O
1
CoS
fl!
(
Ni
i
•a
LI
^S S
* . 23
"i 5 . | a
NiS
NiS
nl
= =f
r
&
« C - =
z "3 - B °
> X
Mn"
3
/ z
^ ~. i £ JQ
MnS
*»«
MnCI.
•3'!^
V
Mnv11
i
Is * 1
H o "5. c
JS
§2§
/
Zn
s | S."
1=1 |
ZnS
S13
ZnCI
»" «
/ N
—
J •o
0 | | ^ / ri. / / \ -
Ba
I
c_
= " •%
a
BtCO,
—
B
c "* •—
* "3 5
•3 • ,1
^
•=< ,.
/
K n
. .. = c
2 0 =
- 0
Sr
Z J
S ? * "3
a « 1
s
SrCO,
ic>'c5
S
i
»•* S
11= §
I x 1
•;
•-z,^
Ca
E ^ =
5- c ? ^
- * "
o
CaCO,
x
c.
J
B
- *" ^ 3
- -3 £\
Mg
B
51
z ^
5-f 1 •§,
4* o 8 ~
4J« j QO-100 tit- «j.'
£ z 1 n • ' - =
£
H
5 « 3 «
2£ ^» «h rr -^
•• i ' S'S ^"^ - - —
^
K
/
1*
z <5
c x
r o^ 5^ 22 "5 •
S
-
i :, •— "3 £
.2 z
^ C -*-( , — — x - '&•£!
Na
e
s
.2 5 = '23.
_3 Ig
I-'P1 z.l -FT -tU'2-
1
NH.
3
•*
"n - = "3 _
-/. ^
-2 Sl||?a|o
5
"
1 *1I ^
— — ~ ? £.-—
§313. TABLE FOR THE SEPARATION OF THE METALS.
E SEPARATIONS OF THE METALS.
377
AsH3
HsandSb
Sn
&Q 55 o
*-z~ .-
c "*• s< ~ "^
\ bta5"c w
H3As03 |- Remove AgN03 with CaCI2 and add H2S -J As2S3 Lemon yellow.
SbAg:1 ^ Dissolve in hot HCI, dilute, filter and add HsS-j Sb2S3 Orange.
SnCI2 | Test with HgCI2. -j HgCI, White ; or Hg Gray.
Sb lsi-r'f SbCU 1 ,-r | SbCI5 reject or test in Marsh apparatus.
Au LllU AuCI ' 'i 30 I All 1 DiMolveinnltroliydro-j AuCI3.NH4CI El
f 1 1 Z I 3 ,r = | 5 j AU f chloric acirt, evaporate to } and ignite to Au°,
\ Pt
ptcu
pt
Iryiitus with exce
NH,CI and digest with;
Evaporate
Yellow.
(NH4)2PtCI6 Ignite to Pt°,
. Gray,
to green- ) Evaporate to dryness with excess of HN03. Dissolve residue in NH4OH and add to an excess
lutionatk i of HCL Tost this solution with Na2HP04 ^'Ammouium phosphomolybdate, Yellow.
d and test with SnCI2 and Cu wire.
•mation of Pbl2 or PbCrO...
H)3 Add hot K,SnO, £ Bl° Black.
H)2.2NH4OH.2NH1N03 Deep blue solution evidence of copper.
For traces add HC.,H30, and test with K4Fe(CN)6-j Cu2Fe(CN)6 Red-brown.
H)22NH4OH.2NH4NO
Add KCN till blue color disappears, then H2S-| CdS Lemon-yellow.
b. HCI and precipitate with (NH4)2C03-j AI(OH)3 White, gelatinous.
V Acidify with HC2H30... and add Pb(C2H30..,)2-j PbCr04 Lemon-yellow.
4 I
I Dissolve in HCI and add KCNS { Fe(CNS>3 lilood red.
(Test original solution (acid) with KCNS for Fe' " and with K3Fe(CN)6 for Fe'' ] Fe3LFe CN)6]2 Blue.
r a. Test with borax bead. Blue bead. ~|
Lb. Add NaHCO , and H202, Green solution. J
r a. Tost with borax bead. Brown bead. -,
b. Heat with I N:,nux ' add Kl.
L Br and NaOH i N|(OH>3 (- Free | in C$2 J
H)J Boil with Pb02and HN03[-HMn04 Purple.
Add H2Sj-ZnS White.
Co— Red
precipitate.
• Test with borax bead.
Ni - )AddNH4OH
Solution. ( filter and add H2S < NiS, Black.
j [ Dissolve in HCI and add H2S04 1 BaS04 White.
\i .
f \
a
\
\ -g i .. 1. Add CaSO , . set aside ten minutes -[ SrS04, White.
I30o)2
a-'n-jO
SrC03
- C
Sr(C2H302*2
'^ a? Moisten SrSO , with HCI and apply flame test.
Z rH ~~
-^x
•O c c
•0*1
So'
Zf*-1 2. Add K..SO,. boil, set aside ten minutes.
H30,,
•o z
CaCO,
^ /
31
Ca(C,,H302),
1 11NH^2iC204add ' CaC^°*' White, soluble in HCI.
«H,P04, White.
—Apply flame test using cobalt glass. Violet,
a— After removal of Mg apply flame test, yellow.
H,t— To the original solution add KOH in strong excess, warm (note odor) and test with moist litmus
paper; pass jras into Nessler's reagent -j NHg.,1, Brown.
378
ACIDS. F1KXT TABLE.
S314.
to *
1/1
V
'-5 pq
icati
Remo
9) -
5 8
o SS,
rH
CO 0)
V
r-J «3
C -
C3 "
cu
SC
'5
eo F
tan J"
eos 41
_T C8
£> ^r
* f- rr?
.500
.s tc
s is
.'£ °
^ cs ^ 3
^ ^
O ~C3 -y 5C
— i j_, Cv
•/. G •-*-( -*i
c "
o ° 2 -2
.
g eft bo
r— ; O2 cc O)
"3 S -2 -
D
fcX)
£ > .
^ J o o
° ° .~s ~
£ a '& >*
T3 -2 ^ -°
01^0^3
S ''/3 -^ O
•-H Q CC Cft
-M 2 o c c
.* £r.bc A §
-Q O
&?
= 'I"
~ >H
rr g
•3 &
a -3
•8 'i
2 -3
•&!
3 2
CO -^
. C
>i o>
-O CJ
^ o
s s
o o
l«
3
*
oa
**-~ cs -e*.
odo
A
s, bu g
evolve
0.
1
acid, HC,H:
ether (§223
,
a
p
senous ox
pre
no
a
e
^ «o
— : <N
- a
^3 CS
^j jj
^ ^
s- x^
-t-> cc
CS r^2
i §
o o
O f*1 )
5 £
§ §
"fr •-
•r '5
•* cs
i_i' . >- 3 -S •- 5
^:W ^ es i.«n,S
•^* o bo a * '3 cs
'>^ ^fc S^4)^
o o
* S
- ^
„ "o
I S
T c
1 Hs
-5 a .g
id'
/. u
5^
113
fi* g
j *5 —
§314.
ACIDS. FIRST TABLE.
379
t-
s
_2
CD
-c
CD
-f
+
;J
CO
0
CO
CD
M
C
i>r
<M
(N
'o
r— 1
a
0
"o
^
cT
rt
0
0
'c
r>D
J^
o
PH
c
£
CD
r5
•vy.'
=+-!
• r— 1
p*j
o
o
-M
CD
a
Ki
C
CD
fc9
Cfl
— •
o"
03
M
II
r—1
0
M
a
^:
-M
S
f-
c
^_-
rr
CD
<T3
_'§H
M
^
W
_g
r£5
+3
0
CD
j3
'o
CD
—
M
W
o^
2
e?
c
co"
CD
^^
+^
'-P
CO
3
PH
,
1—1
0
^
^2
pC
CD
03
^
_5?
O
, ^
3
^
13
PH
cij
_s
3
O
'1
+
o
CO
X
O
CO
o
,'-.
03
^5
'3
cT
CO
in
?f-
Q
d'
o
r— i
0
"3
oT
CJ
M
-IJ
•s
-r-*
rO
O
o
bD
•— (
t-,
CJ
/— N
03
0
CO
CD
CO
O
c
-<— >
G
58
f~t
"3
^
PH
UJ
'o
«\
-C^
CO
-z
•*•"*
' '
CD
0
p,
"3
g
^H
C
_C
]bC
C
CJ
tJ
d
^
'C
o
0
^
1,
CO
S
IW
o
rt
;-
— •
*J
53 <X>
•2 +
CD -f
+
o o
«
* * >
fl » ° g
•s s |
t3 53 ^ 'Trt
^ bJO 00 *
0 rrt ~ .23
> o oo
® S g *
. r^ OJ O
|S S 8
o
CD
P
.S «3
CD CD
-u r&
CD 'P
> O
Tfl ^^
S tJ1
C W
O CO
co
CX70 CD
«5-£3 g
w
. CD
CD -rJ
S f=
J s §
% co s
o oo «
'
O -rH
ca r^
^
o
0
-I
si
rH O
G -*-•
Is
Kn ^>
be -<o
c e
•g K
O3 i-O
M C
0
o
0
« C '-?
o
^5
CD
be X
S *
'3 CD
CD r*
-
00
C g
"-1 0)
^_,
2 PI
CD -.
ri CD
— ^Q
^ C
' — CO
P C
o 55
•^ * "r* .° §
0 5J5 ^ fcf .S
* PH^.
0) P
r/3 r
- o
.
»s
5 fe ^
r/2 CO
l'/J <M
_2 <V>D
CO
CS -i->
S ^ « 0
8 'S °X)
O O
~ ^ "*-" o3
Iplll
WH 03 .rH CD
CD
co
CJ
's
CD O C3
C3 ^3 0
rrj
5
o rS
0 PH
'
03 CO ^_
CD Pi -—
3^5"
'S "% T3
22 'i
fH O rC
s fee
-S ^
=4-1
CD 03
CD CJ
Us
03 O
CD ""O "
rl CD ^-.
r"~( _LJ 1 - 1
- g-r
Jz; 5 *
0 g T3
W | £
CD O O
cS
«
O
o
-s
03
to ^
111
r1 r1 .2
<U <U rCj
-380
ACIDS. FIRST TABLE.
and
bona
• S _«
r2 ^ CB
±i -»-> cu
c3 {)•*-»
- £• 2
be K, *5
r- -C &i
£ .s 3
^ O co
!3 " 0>
T3 S 3
'0^3
rt 2 -o
£» 5 »
|1 |
—
V
•S « 'S
2 &i c
J3 ° S
CS ~ ^
04 N '^
s.|2
CO ?' -t->
— ^ O
CO <U Pi
O ^^
•P ^ -*
>•> S __
o ~ 'o
o> « co
bya u «S
C g
•3 £ g
*v|
s ic -
bO.S ^2
§55
E -^ 2
08 CO W
c?^3 cj
^ I
EH H
eo
CO
VS3
0
u ^^
O
p<
CO
S
o
s
I
5314. ACIDS. FIRST TABLE.
CD r o IM ' j3 ' ^. ^ ^ „; O> <D CO CD G CD
|i:| || -is g; iJlfBtf
.siH ^ 1-s-s is olif^-s
o3 o * tn --^ c ° S^SS'-' j u G <+-^
s sJfgJ:;- |£p ®!l «i lloiii
t=^ «,_, £ * t< 2 „ 0 ^ ^ CD ° cS bCrJ « S
W Ofn^-^S * ^ S o «> M r/) « ScD'o-2^^
a>£S;^ *^8 p-co^ ^ ijM3>9
B 1*5*1 Itf? sle -3 '8 * illl
If 1 1,1 •!§•! |t* l.s li'sifl
^ S | &~ g 3^1^ S^g^Sg ^&1..-
d. PH^^^^ ^§1 •S'-S^ff O 8 -g'S-ls
••§ *jT § 2 i * ^ 3 I? ^ I I -a <*£ * § I * **. |^
1 52 5 2 r^ 2* ^ S -S S S 'C ^ « ^ | 'I -^ § , '£ J
?^ if 1*1 III ^tl^gS 3^1 li«
S« J^ill 21| ^l-g^l-? ll^Vif
^ CT ^'^S^g -2^S -C^^^gg r^O^^O^
_rrt etJ'SricoS .^CD ^PHdpnOS*1 ^^OiE-lrfCS
1 *c ,t?o 2 *c w'S^^cj ^StJ
_i ''O i u O r? "S O P Q; c 03 03 3 rC
3 1 Il1^| SH °M^Sso| °^ls|l
•|rt f^»l^ IF--? S*|lli .^litjg
& o qo)>jp!. '>^>aj +» r3 C s t-> d H r-H >^ J^
9'« ^S^0r2 3^=5 S I* J'^ J&'fl §1 'o^S
rH 9 ^ .-£ — ! -M ^^i^'-'Q^
O ^ G 'O PH
-^ _°J O O ^->
O a) Tl P
CD
03
1
f*
ad
0
^
^_J
a
03
'-*^
.
D
0
'5
G 5
a ^
-t-3
2
CO
CD
CC
^
9
09
<*T
?H rrt
CD .rt CO
-t^> O CD
•^^
o
o
^S o
-(j
rt
CD
O
fl
O
0
g
n3
•fl
0
13
«
o
1
• rH
T3
• rH
0
03
>-j
^2
T5
CD
03
O
PH
O3
). Insoluble
0
PH
C5
«
4)2
OB
0
9
os"
COO
CO
03
soluble in wa
hosphorous a
or phosphati
Changed by
d transposes
08
:i
CJ
O
^~ .
CO
bo
H>-i
considerable <
by the reacti
»— i
CO
GO
»cT
bo
_fl
O
03
_o
_S
s
rr
G
_0
Hp
o3
PH ~
01 o -3
S &-
°;1
00
^
~S ""§
R -tJ
0 0
OJ
"o3
CD
^
0
CD
"o
1
O3
CD
-4-J
'o
ej
o
•S
.„ hH 0
' _^ . — -
0> H ?H
•S °
rfj
"rt
[p-
PH <~O
uf
^>
^3
03
I
d
_0
'£
_c
03
'o
^ ^ft
Sf G co
T^ ^
"o 3
38
CD
02 o.
O) •-
COS
ooo
02
o
a,
CD
fH
03
CD
O
'co
CD
3
«
X
15
-^
?a
2
CJ
0
—
T3
>.
_c
>•
^
^3
CD
««
0 rG 0
* ^-5
PH^ ^
0 o5 ^
PH bJO
^>3.S
O CD
'o ^3
'ftQ
CO
"5
g
p
c
>•
2
03 [^
PH S
co
o o>
9
CO
CD
PH
o
^
^^
O
o
CD
Q
O
g
PH
EH
—
03 S --i .^!
d
s a
c ac
« e -gs * -a s
co ^_ ' « ^5
CJ Si -M '-J CD O .
^3 « « O ^ ^
;s.s * &*a s
L, § t. S 9 S
i-T = CD o *" ^
? § PH § ^^
50 53 03 cy ^ ^>
, fc5 PH^T -G rG
I ,11 I
03 O ft &* ~3
ft « ft
.'M ^* ^ ^ "3
o) P4 ft PH co
382
ACIDS. FIRST TABLE.
.^314.
o;
_- -
GO oo
»0 0
02
1 1
cS
"° •£
o> -"g
42 9
1 ft
o .2
CO 00
»M O
0
2 ~
a
42
— <**
C a>
.2 -5
be F
C 0
'
Scd
o bo
CO
*
eu p *^
~ oo
EH a. „
£ -~||*.
'o H ^"°°
ii 1 1
0
CM
cor
O
a
~*~ 2 -2 <"
2 <jj O '3
M° « -T3 o
cu . —
SB K C
~r -2 c p
°* J2 5 "§*
C
0*1
- 3
cc PQ 73
I
CQ r"!
"E
£
'^ 13
_^2 ~^ o f^4
M
W
g
t
1 «
<» -
1
o
c
^
bt
_c
-2 £
"3 -S
t< «t
j § 1 S-
(2
d
coo
_L
J
S'i
s ~
-*S *
•§• *o
S
3
c
CS i-Q
cc ^
S oc ^-:-
fr-
.2 o I-H
42 - O
.^^ § 0
S 2 £ c
CS
.1
T3
5
i
i
^
'S "
03 1
1
Odor of burning s
es precipitate salts
2
3
42
j=U
"5
o
c
_o
o c
C-«+-
s bt
cc T
C ^X
c -S
i/ ~
fc. -^ CO
^_ bt 12
0 C CO
CS
OB 2 ••>
C 'o t
'•§ £ W-
- S CO
-r co
S £ «
42 "cs 5-
lit effervescence,
lution of silver nit
1 but soluble in am
; ammonia. Obtai
chlorine; bleach in
chlorine; bleach i
M
CO
"3
'o
n
*3
CO
«b
3
?
d by hydrochloric ;
Reduce to chlorid
|
u
a
2
o
.~ti
o
i) -^
•g *S
.3> g •« 6
•H
•£
/
-
9
5
• ^2
o
"r"
' — CD'
^n ^- ~ t-
O
O
«c
C '
VI
N ^"
— *^™
— oo
^
c
-4— t
f^ — ^
O
si
•5
"H- w
t! ''- S
3 o «*
'B 'E I"
O
ec
5
2 to .
a
o
oT
|j
a
.£ ^
cc
r' -S <y>'-
o ^ '5 ^
W « o •£
"p
o>
42
3
'o
<u
concen
~7
1 +
1 1
n applicat
r* CO
.^
•**"• (J^
"iL^ o
'o , 'S 'O
-^^
-^^
_ f*
9
*fl
j- W
03 y'-
4=-'
»s >
•^ 4=
3 . *-
« O cS
etf =i . 3
o S -2 ,2
'5
'^
r^H
'S "
*
a
11
p -^
r^: e&
omposed
—
S
. CS
^ cc
-4-> O
• o es -^-
•2 'ftg
Hi * "^
'C •- o °
^ .— 3 o
3 H*
g H. S -s
T3
m
CO
O
1
to
o
.2
5
es «
-1
^ §
~
^T*
o
^3
O
o
i >
s
_
o
^~^
•^
V
0
O)
•*
60
c
<
w
C
P
S
s
o3
H
•
2
•0
o
cn
OJ
g
^j
^5*
2
4>
2
2
• i-)
^
•
'^
."yT
o
CCj
4
43
en
IH
*#.
^.
o
«
2
0
Pi
O
o
PH
o
1
4
CO
H
S
§
0
w
3
0
§314.
ACIDS. FIRST TABLE
383
0
-f-
o>
03
-i->
o>
,C
-M
•—
O
O)
-M
03
co
00
la
B
'|
O O
G CJ
.-- f^
i— i ^t—
O -M
N 'g
4*
a>
'o
«f-l 1 O)
0 'En ^
fl 'a
'"*
cT
f-i
tJL
^
o
M
v/.
CD
j=i
M
GC
77^
' -^
o a>
-(J
O C1J
-t-3 ni
•rt fH
-1
+
'S
O>
1
-M
1
S,
1
H
OJ
QJ
^J
3
cc
0
f-i ^2
0 «+H
fl °
co
t-
OJ
co:
^ s
O 03
P> 3s
i •»
o
1 ^8
cu i~c OJ
^ o. <D CO3
• I—I SH
S
'C
o
3
\~
o
0
rf
to
o
W
CO
o"
M
CO
+
~i
+
o'"
Ig
'S
03
g
'I
C
.2
03
!§"
'S
OJ
rd
o>
o
p
S CO
CL.CO
02
* COO
O) o
r^ S-t
0 6
CD
'o
T
-i
CO
n3
O
f-,
O
O)
"S
IS
bio
>»
C
03
o5
o
-»J
5
s
^a
to
BO
o
_c
"p
$
4-;
'o
d
HH
ofl
SM
o
'S
CO
£
^o
CD'
i>
0*
*/r.
romates, but the heat of combination, aided by t
5
S
~S
1
c
0)
fcx;
0
•V
0)
c
's
0
E
O
'G
'£
o
sulphuric acid bromates are merely transposed,
eration of bromine is due to the decomposition
Dinates with an alkali carbonate forms a bromide I
3 sulphuric acid, hydriodic being liberated but no1
With the concentrated acid decomposition takes
3
•9
'3
3.
3
o
=H
— '
X
'i
"* o 4J
X 'o -°
_-M o 53
: > >^
W '5b co
rrt . fl
1.1 -J
ell
- s "
M 0}
i « i
W M &
w S
>»*: o
X5 ^ O
^ .2
OJ >
CC ^ '~^
0 ^
^^ »
0 £ ©~
0 PnCO
OJ 75 0«
^ S S2?
2 ^ a>
03 9 3
§!/.;§
3 S S
HD w PL,
-ed by AgNO, and any of the reduction tests applied to the flltra
M*
^
-fro"
O
^3
O
CO
OS
r£2
02
•*
o ,_c
•r~
+->
-3 <-•
J1
-2i!
0
+
a;
'"§ §
_o
0
£
'~
%
A
Q
ob
O
+J
r^ 'a
3
^ 1
•+•>
<U O3 bX)
CO
j3
rrt 1
-tJ
^
+3
^
CO
CO
_ ^
«J— i
^•33
2
3 .J3 aa
^~
o
f^j
•2 '
'oj
c
—
vS
i.
f^
cq
V
^ r«
b
^ 0
; —
3 2 ^rj
"*
02
C
0) ^
73 bO
0 <J
« 11
O
O
|
|
si)
45
CO
cxT
03
_0
'S
GQ
o
p«
s
o
CJ
V
O
-t->
'3
bt
n3 &i
CU ?-•
S 9
8-*
c ^
<U
•1-1
j- 0) «O
. £ *£ CO
^ "cc °O ^
il solution,
."^
03 '
g
.5
^r
^
eo
J3
—
^^
03 ^-i
— j
0
'Q bJO 03 °S
c
W
"S
"1
o
O
cc
T:
61
_c
3
_v-^
t-
<M
<X7.
«X)
O
__Pn
3
O2
nU
OJ
irected.
•_, O3
-(-. 0)
C ^2
C ^
'o
03
j^ ^^
O M
03 a,
t/3 -i_j
O> T-
~o
rt
o
aS •? 13 °o"
o w "
«\ fe coo
nt In orig
d
"5
'o
cT +
I
^O
C
c
|
O
o>
O)
central
—
-i-»
03
O>
X
r^ O
o ^
CS t-£2
1
O
•^ o
^3 X!
o X
^ *
^
O>
E g „ «»
o 9 s -,
o "*! <V O
2^ T3 -*J 02
TO rs rrt i-^
> .S -M S3
i be prese
0
H
fl
o
CO
O
|
PQ
O
S
«
*
M
09
tJ
"S
n
g
g
7s
O
o
rti
f-l
E
s
pq
pq
M
384
ACIDS. FIRST TABLE.
0) O) o
O ^ :
w - •>
1S -2 g
* v a t-,
° +* cS 03
O —
O "^ o S
* ^
•3 -2
o
CO C
V 0
s -S
o o
s Q
CJ O3
03 C
J§ be
o S
• S 03
O -t->
03 O
CD ^^
o *o
ll
es "3
O
S X <U
H •« •£
O >~> 03
£ 'o 3
T3 S JS»
o c o
a § 6
S .S O
x c x
.ir o cu
O «; oo
£ £ °
~ cs P.
o ^ c
11
£ s
bOT3
C C
'S *
03 >
S3o;3
b^2 oo
o 13 ^
s; s
?« 2
-o .t
"8 'S 3 fl
s
cu
T3
CO 0
IP §
* s &
"^ O r-
a '-5 S
>> 3 o
11 I
00 — -i
T3 O " " -J.
a} fl *M O QO
c i» ^;
O M CD
•— ' ;j t- r-
GO • «3
co . -3 rr o
^ i v. -5 ^
-4- .S r- 53 -tf
^ -5 I .2 §
S — P S je
N
•*•* O oo
C '*"*'
<U ^> r^
o 9> B
« ~ f3
o LT -a
W '•v
O co
0) w
^c .>
T3 - bfi
i -1? 2
* ss •£
P-I «t_i
-5 C ~
C **' V
o> rs xf
0 ""C .2
C "Tl K»*
'
OO
•^3
=:
c?
S
i
^ . o
03
•8
-2 » ?a
'S c s
« c S
oo '** '~5
3 O •— I
C «
I -U-s
^N »^H j-
^ ^^
3 ^ be
7, ^ «
— -
cu o
>> A
- » —
<-s —- *- ' —
Si -2 a
<J ? o c
co O O w
»Hi|l.§
| ^| "2 '^?i
0 o ^§ 3 -2 c^-
fl a ^
I
g
S
3 J3 **
JS '"S r-J
OH ~- O
•g 'C be
% y> G
& « i "o
-*->-: co
^^ cs r^
i* _e t> ~
s ^ ^|
co S C2 ^-~
V
-»->
CB
S
o
o
P
§314.
ACIDS. FIRST TABLE.
385
03
O
T3
0)
be
d
03
^5
O
'-3
T3
O)
o3 as
CO TO
^ S
11 • s
i—1 «••» . .— Q.)
f*i t - <—* -+—
C % S S
•s'S^ S
•§1
d H
g g
§ «
- -
S O
X OJ
ated
g
r^ =0
."S -c
^ ^
r^
bo
d
KH
fcc'o o
CO
^ I
1 8^
^ es
_0. cc
cc
n3 pi
O) O
££ «
C 03
03 be
-§ §
ci.
03 [,.
p^'o
S $
o
o
o
119
-
s;
* JT;
bc^
d KJ
03 ^
C "
H /^^
ed
anganous s
lead peroxid
ess
and
a
g
e
o
o :s
^ «
r- *| i O3 GO
s o H be
o
OJ
p
386
ACIDS. SECOND TABLE.
$315.
^
•
.-
^~ .2
^
K
0
P
•— > a
0
^
-_r
s
js SI
5
1
S
1
N
1 »
S
a
Jj
—
aT
^a
S
.a
V
1
C« rt
•8
D
£3
•_"
2 §
0
-
q
1
"1
S
'1*
E
2
•5
-?'
COO
P.
l!
a
4
b
e
•
•a
3
H
«-H
o
o
1
S
'i ^
O
~,
O
&
u
£•
1
'p .S
of
—
4
§
03
0
h.
Vlil
1
6
S
- 1? "^
^
8
B
«
.^
—
^i
**- ^j
a
40
q
coo
R
.1 <
•^
s«
r-4
o
T3 '&, OJ
S
•
C
•3
0
Q
e
ۥ
u
t~,
»
S
S '9 s
a
*
f
co
09
O
1 °
0
&
oS
O
a
.S S-eT
C3 m O
w
3
S
<
q
M
Q
a
q
<
H
4
H
a
-
H
S
•o
•it
3
h
&
3
~
CO
5
o
i
^
9
08
«
I
•w
Q
0
n
•<
'$
o
-*
"i 5
«T*
5 ?
^ S
||
i
~
>.
-=
tf
"•-
'S
•-"
1
•_
a
1
s
B
§314, or during the examination
Solution: H,PO,.
Add excess of ammoniui
§255, 64.
With ainmonium molybdate, arsei
exccjit that the phosphoric acid ii
acid not until warmed to about 3
Study §§146 to 153.
i
1
a
i
as
a
O
a
x)unds are oxidized to ll . so,. Const
sulphates see $266, 8.
1 acids are also precipitated.
i).
H
•°
g
£
^ «
to
•o
a-s
H
d
g
*0
"s
=2
J3 """
1
1
CO "S
1
•0
c
§•§
38
E-fi
0 «
OS
Q
O 1*
_J
"^
^** i^»
A a
5 55
~.
^
^ .1,
-
g
> c
$
o» ***
*> .„
\
Q
C? C3
g
^^ t<
O $
1
g
?
4
4
^
h
(D
to
-
•s
s
1
«
O c«
5-s --
E
C .a
~
a
—
I
J ,
•= a
f -2
^1
£
13
d ~
O 1
fl S
'b
•S
1
e
i
s
"S
0
wer oxidized su
ulphateor othe
nd some of the
:. both cannot 1
5
s
~ 5
<S
5 *n
p
3
O oo
» a
2
i
4
4
i
Concern
•2
1
-« a
A, 09
S 93
0
^ 1
4^" ^S5
0
0
-
o
I'S
11
o
BL
t, sulphur
led and all '.
> of barium
ite3($113',
y In' i'1-t.^i'
4
4
1
|
4
H
i
*
1
ipitate, barium #
precipitate to settl
0
0
r:
0
0
"S
~
^D V5
'o »«
09
o e
^ ^
'"" -2
^ «
3 "
o ^
Residue
W «
ll
1--
"8 *
00 85
S< (N
S IM
cos
5 rf
o O 10
a
a
8
I
a
A
id may be decompos
cerning the analysis
urates($112), seleni
CrO4 or (..<<» ma
3
/;
S «
'On
i s
o o
'3 Jjl Q
a>
a Q
* o
T a
2 5
•i-:
j- o
— « o
S S ^>
<M
hH
W O
H ,
a"
I
s
'
§ -1
o M |
r
*+ *
-«j
^
^
Q
Z
0
§316.
ACIDS. THIRD TABLE.
^387
§
CO
o"
S
D
li
O
M
bO
01
M
O
ja
fffffjlfill
1
2
0
CO
bo
<l
"1
H
W
w-aǤsl ^S^I'gS
•
•
§
ooo
M
o
100
A
bO
*
03*
M P
MM~i ^ "fCd 0,^-^2
s
•
0)
^^
^1
hv
£-4 C
fcJD w . S^ ^^* *^ -*^ o (— ' **<3
B
<D
b
"^
O
O
Q
0 r°^al
S
<o
r.
PI
W
_bc'§ 'C °°"
•a
H
bo
be
te
S
1
<j
^
4^
o
+3
e *" •>
|
M
H
£
^;
W
d
^ . 0 ^ (M $J
CO
i— i
£
O
m
bo
pq
||
il^I'f
1
W
M
GQ
6
1
bo
5
bO
W
o
K
6°
H
bO
<J
•|H
A
EH
lull
as an indie
ssolved.
O- O .^— *-2 ^ *S -*J
*~3 £j ^ g^ ^ t^ c3
'O ^ *" ^ ^~ <& &
c3 *n ^ ^ift*-i *M ^ o
O ^ fl ^\i *^ _o *^
CIPITATBD B^S
1
1
oo
rt
0
bO
0
d
bO
d
09
bo
d
03
bo
H-"
d
Is, representing
ilation.
CN, AgCNS,
phenolphthalein
ration,
be completely di
H
03
WJ
Q
S
W
W
H
O
s
1
i
03
1
O
d
W
d
do
^C
d
e>
bO
<5
03
^
03
bO
d
PH
*-i
0
*^
tj
?
s
8
HH
bo
d.
h
q
be
^B*
v\
03
O
be
1
-B1
"s
%>
«=s
e
S
'e
s
e
mmonitim silver compounc
e HNOs to a slight acid>.
be
<J
«
bO
<1
O
bO
<J.
O
-M
C3
ith nitric acid (use litmus or
tate form, remove it by fllti
jr c iromate and iodate may
For Bromides.
Test the original so-
^ £"C
crs
"3
S ;1
-5. 2O -«'2Qi
o
2
^
hO
^3
CD
< 8
g
|O t-
" oi m ~ <j fe
£ 'p, >
OQ
O
<o
a
O
h
o
bb
O
03
1
fe
O
bO
droxid
• • 'S
ll
£ ^
ft
•rH
O
CD
b
Pk
1 ^
°f£gS|l|4
|1|
3&-S
2 =j **
*" T3
3 T3 ••=«
d
fi
d
CO
be
;d
'o
c6
0
3
bn
P>5
A
1
+~
r-H
S
r£ "?
H
a-j^i |^§s
23 §
5l|
3
J- 1
50
W
•rH
• 1-1
r^ c
CO rw •
q
be
03
'*
•i
a
0
"i g£
« S co
O <S 00
1 a s
M
d
bO
'S
pq
bO
amm
•-H 2 ••"
tp'-g £
00 U3 N
^ CO 5
•-§1 §
s "° a
OQ
•^ tn oo
t— i
^
*X3
|H
%^
. _ •*
^> *•» c£
h- 1
bC
bO
1
bO
V
i
fl rt T
.S a o
03
£0 c« .£
IS"!
i
O
03" M
S
<D
i
<&
«•
)"•( . — s—
fcO 0 -=
<1 S
~ o
Ill
aj" a> aj
5 5 '«
"3 — i o.
H
S
W^ a
-o
s
.. V
1 .3 E*
^^ 2
'ft
•^ -,
be bn
8
•l-l
m
o
•w
§^ ^
'•5 f>8
c3 ^
<4H +3 ^
CO
o
<J <j
(^
Pn
•2 "S i
^5 'H H-
3
i-H
o
co cu T:
—* p! "^
J/r -t-
CO
PH
CD E-1 ^
— v ""O
•^H
e«o
P3 .2
c
r^ CS
'5
388
ACIDS. FOURTH TABLE.
§317, 1.
§317. TABLE FOR IDENTIFICATION AND SEPARATION OF THE COMMONLY
OCCURRING AN IONS ( ACIDS).*
1. Boil the material with dilute HNO,
CO.
S02
N203
H2S
HCN
C2H402
There results:
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,CO:! 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
C03" . Acidulation with* HN03 sets free C02, and Si02 is precipitated;
identified in the microcosmic salt bead. The filtrate is made ammoniacal.
3. Ca(NO.,)2 solution precipitates:
in
ilute
HC1
/heated
F'
as CaF,
CN'
as CaC,04
as Ca(CN),
insoluble
in acetic
acid;
v
C4H406" as CaC4H40(1
HAs03" as CaHAs03
HAs04" as CaHAsO,
C insoluble; H2S04 liberates HF .
1 soluble, reappearing with NH ;
| decolors KMnO, solution.
with Fe" H--FC"" + OH' gives
Prussian blue on acidifying.
with K' ions in concentrated solution po-
tassium bitartrate precipitated.
In the filtrate from the above,
H..S precipitates As2S3 at once in the cold.
In the filtrate from the above,
H2S slowly precipitates from hot solution
~S2 + As,S3 .
In the filtrate from the above,
ammonium molybdate gives yellow pre-
cipitate; or Mg" + NH4* + OH' gives
MgNH4P04 .
4. In the filtrate from 3. Ba(NO:i)o precipitates:
Cr04"(Cr,07") as BaCr04 , yellow, soluble in HC1; the yellow color of the
solution becoming green on boiling with alcohol.
HP04" as CaHPO,
J
•From Chem. Prakt. Abegg and Herz (1900), Breslau, Page 113; reviewed by Fresenius. Z~,
1900, 39, 566.
•8.
by OH' / coloration.
gives with \ Prussian
§318, 2. NOTES ON Tin; DETECTION or AC//JN. 389
SO " as BaS04 \ / unchanged^, remains insoluble in
HC1.
SiF6" as BaSiF6 \ insoluble on J gives off SiF4 , which deposits
in HC1 ; ignition \ Si02 in a drop of water ; the
residue, BaF2 , is soluble in
HC1.
5. The filtrate from 4. is exactly neutralized with HNO,*; Zn(N03)2
then precipitates :
Fe(CN)/" as Zn3[Fe(CN)6]2 brownish-yellow dissolved ( brown
Fe(CN)/'" as Zn2Fe(CN)6 white
Fe'" and H' ( blue.
G. A few drops of the filtrate from 5. are treated with as little Fe'" as
possible :
Eed | Fe(CNS), ) on j permanent red color,
coloration 1 Fe(C2H302)3 j heating 1 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, one part of the solution is
treated with CS2 and a little Cl-water:
I' violet coloration, disappears with ) , _.
.... .,, V much Cl-water.
Br brown coloration, does not disappear with j
The second portion is evaporated to dryness with K2Cr207 , fused, and
the mass after cooling distilled with concentrated H2S04; appearance of
oily brown drops of Cr02Cl2 , forming Cr04" with water: Cl' .
7. A concentrated water-extract of the original substance is treated
with concentrated H2S04 and solid FeS04 or Fe" solution, prepared cold;
a brown coloration shows the presence of NO/.
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., SO/ and S"; S03" and NO/; NO/ and CN';
NO/ and S"; NO/ and I'; NO/ and HAsO/'; S" and HAs04" , etc. It is
to be noticed that this always simplifies the analytical procedure.
§318. NOTES ON THE DETECTION OF ACIDS.
1. 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 boiling; again soluble on cooling (separation from citrate). Calcium tartrate
is soluble in acetic acid (separation from oxalate).
2. A number of basic carbonates give almost no effervescence when treated
*In the original German text it is directed to use HC1 at this point.
390 &NOTES ON THE DETECTION OF ACIDS. §318, 3.
with acids. To detect tlic presence of small amounts of mrlm-nah; 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 Ca(OH)2 is redissolved (solution of lime in spring water). Boiling
drives off the excess of CO2 and causes the reprecipitation of the CaCO; .
Barium hydroxide may be used instead of calcium hydroxide.
If. If compounds have been strongly ignited previous to solution for analysis,
oxalates cannot be present.
5. In Table II (§315), if strong oxidizing agents are present, as KC1O .
X2Cr,O7 , KMn04 , 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 hydrochloric acid, heated to boiling and
tested for sulphate by the addition of a few drops of barium chloride (§317).
6. In Table n, if sulphites or thiosulphates are present, the solution in
hydrochloric acid must be heated sufficiently 1o drive off all the sulphurous
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 fcrricyanides 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. Failuw
to lilcocli hidif/o xohitiim in the presence of an ej-eess of liiidrorhloric acid may be
taken as conclusive evidence of the absence of nitrates.
9. 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, SiO2 .
10. Meta- or pyrophosphates do not react promptly with ammonium molyb-
date. In the usual course of analysis they are changed to the orthophosphatr
(§255, 6A).
11. Phosphoric acid may be detected in the presence of arsenic acid by
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, fi;). See also note 26.
12. Sulphides which are transposed ft;/ hiidrncltlnric arid are best detected by
the odor of the evolved gas, and by passing the evolved gas into ammonium
hydroxide and testing with sodium nitroferricyanide. Other sulphides are
decomposed by nitric acid or by nitrohydrochloric acid with separation of
sulphur as a leathery mass or as a yellow precipitate. Persistent heating of
$318, 19. NOTES ON THI-; DETKCTIOX OF ACIDS. 391
Ilic 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.
IS. A sulphite (or other lower oxidized compound of sulphur) is readily
detected by adding barivfm chloride in excess to a portion of the solution,
dissolving in HC1 , 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, thjo
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 sxibstance. The student will be aided by studying §§257, 8; 258, 8; and
265, 8.
15. The recognition of chlorides, bromides and iodides — by evolving their
chlorine, bromine arid 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 thiocyanutes, see §234.) The chlorine is vaporized (from
another portion) as chlorochromic nnhi/dnde, and the latter identified by its
color and its various products, as described in §269, 8<7. Before the bromine
is identified the iodine is to be either rcmored as free iodine, or oxidized to iodate
(§276, 8ft). The oxidation to iodic acid is effected 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 11).
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 , A.gBr , 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. 8rt).
17. The direct removal of iodides ft// precipitation, learini) bromides and chlorides
in solution, can be effected (approximately) by copper sulphate with sulphurous
acid (§77, 6/% or quite completely by palladous chloride (§106, 6).
18. Chloric acid is separated from hydrochloric and all other acids of chlorine,
bromine and iodine (except from hypochlorous acid, and from traces of bromic
acid), by remaining in solution during the precipitation by silver nitrate
(§273, 5).
W. Chloric acid is separated from nitric acid — after finding that silver nitrate
!?:rex no precipitate in another portion of the solution, acidulated — by evaporat-
ing and igniting the residiie, 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 nitroiis
acid.
*In consequence of tho relative commercial values of bromine and iodine, and the medicinal
relations of bromides and iodides, it is of great importance to search commercial iodides for
intentional and considerable mixtures of bromides— an impurity likclv 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 occurring: from the
difficulty of exact separation in the manufacture of iodine from kelp.
392 NOTES OJV THE DETECTION OF ACIDS. ij318, 20
20. If we have to separate- chloric acid loth from nitric and hydrochloric acids,
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, ignited,
dissolved and tested for silver chloride, indicating chlorate in the original
solution, and another portion is tested for nitric acid. Also, chlorates arc
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 ferric acetate (§§126, (>l>; 152; 223, fi). 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 iri/7i 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 40 grammes of K,Cr,07 , dissolved in 100 cc. of water, which lib-
erates and expels all of the iodine without disturbing the bromine and chlorine.
oKjCrjOr + OKI = Cr2O3 + 8K,Cr04 + 31,
(2) 8 cc. of a dilute solution of sulphuric acid (consisting of equal volumes of
ITSO4 sp. ar. 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 ferricyanides,
see C. TV., 1883, 47, 209.
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, 8f), a drop or two of silver nitrate should be added to the precipitate
before heating with the acid or a black precipitate will be obtained, apparently
carbon.
25. 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
ncid as described in §276, Rr. 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 heating
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 sulphxirous acid it will not interfere
with the ammonium molybdate 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 jn 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.
2$. Manganates are readily decomposed by water with formation of KMnO4
and MnO, . 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.
PRINCIPLES. 393
§319. PRINCIPLES.
lii the following statements, the term salt includes only cases where
the metal acts as a base, e. g., chromium salts include CrCl3 , not K2Cr04 .
Only salts of ordinary metals arc 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 A12S3 and Cr2S3 are not formed in presence of
water. (Some oxides after ignition fail to unite with all acids, e. g., Sn02 ,
Fe203 , A1203 , but by long boiling unite with a few acids ; while ignited
Cr203 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 Snlv
it is SnO(OH)2 , with Sbv, SbO(OH),, , in all other cases a normal hydroxide.
[At boiling heat instead of normal hydroxides other hydroxides are some-
times formed, e. //., Fe4On(OH)0 , and Cu302(OH)2]. This precipitate re-
dissolves in eight cases, forming, if potassium hydroxide be used . . .
K,Pb02 , K.SnO, . K2SnO;! , KSb02 , KSb03 , K2Zn02 , KA102 , KCr02 . The
last precipitates on boiling.
4. Ammonium hydroxide precipitates solutions of the first four groups,
manganese and magnesium imperfectly and not at all if ammonium
chloride be present. The precipitate i? a normal hydroxide, except that
with SnIV it is SnO(OH)2 , with Sbv, SbO(OH), , with Ag and Sb'" the
oxide, with Pb a basic salt, and with Hg a substituted mercuric ammonium
compound, Hg' in addition forms Hg°. The precipitate redissolves in six
cases, viz.. silver, copper, cadmium, cobalt, nickel and zinc. With silver,
NH4AgO 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, fi/t).
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/t).
7. The iodides of lead, silver, mercury and ouprosum are insoluble.
Hydriodic acid and soluble iodides precipitate solutions of lead, silver,
394 ;•/,•/, \r//'/,/-;x. £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 red need to arscn-
ous acid with liberation of iodine.
(Bismuth, stnnnotis and antimonous iodides are really insoluble in water, and
are readily formed by the aetion of hydriodie aeid or soluble iodides on the dry
or merely moistened salts. Hut the dissolved salts of these three melals fre-
quently contain so much free aeid that it prevents their precipitation by
hydriodie aeid or by soluble iodides. Arsenous iodide is deeomposed by water.
It may be formed from the chloride, best by adding hydriodie aeid or a soluble
iodide to a solution of arsenous aeid in strong hydroehlorie acid. Bismuth
iodide is black: stannotis, antimonous and arsenous iodides are yellowish red.)
8. The sulphates of lead, mercurosum, barium, strontium and calcium ar >
insoluble, those of calcium and mercurosum slightly soluble. Sulphur:)
acid and soluble sulphates precipitate solutions of lead, mereuro.-
barium, strontium and calcium; calcium and mercurosum incompletely.
!). (a) The sulphides of the first four groups are insoluble. Hydro-
sulphuric 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 oxide>.
$257, fi.l).
(b) 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
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.
ft. A normal hydroxide with Sn". Al . Cr'" and Fe'": with Sniv. 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.
'/. A basic carbonate in other cases, except mercuric chloride, which
forms an oxychloridc. Carbonic is completely displaced by strong acids,
for example, from all carbonates, by HC1 , HC10,HBr , HBrO, , HI , HIO ,
H,C.,04, HNO, , H,P04, H,SO, . and even by H,S . completely from
carbonates of first four groups, incompletely from those of the fifth and
sixth groups (Nandin and Montholon, C. r.. 1876, 83, 58).
§319, 13e. PRINCIPLES. 395
11. All normal and di-metallic orthophosphates 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, biit an oxide, or an oxy-
chloride.
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. All 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 Pb304 and Fe,0;! , but ^at a white heat form PbO and Fe304 . Other
oxides, if ignited in the air to a white heat, when changed, either take up
or lose oxygen and leave ultimately the following: Ag, Hg , Au . Pt ,
SnO, , Sb26, , As,0, , BLO, , CuO , ' CdO , Fe:,04 , Cr,0:; . Al,0, , Co,04 ,
KiO , Mn,04 , ZnO , BaO . SrO , CaO , MgO , K,0 , Na,0 . In a few cases
ignition at a lower temperature gives other results, e. g., Co20., , Ba02 ,
Na202 , Sb204 , etc.
1). 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 CO., 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 are 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 burns to C02 . •
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 present.
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.
1). 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.
31)t>
EQUATIONS..
,<320.
§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 the
rule as illustrated there. When the time permits, the operations represented
by the equations studied in the class-room should be performed by each
student at his laboratory work-table. At first the teacher should select simpler
equations illustrating' analytical operations and the principles (§319). Later.
the more difficult equations involving* oxidation and reduction should be studied.
The student should give the authority for every reaction. The accompanying-
list of equations is merelj* 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 in the
first substance. For description and methods of making the basic salts used
in this list, see Hammer's AnoryiniixJir
8,
9,
10,
11.
12,
]::.
14.
15.
if,.
17.
is.
19.
20.
21.
22.
24.
25.'
2G.
27.
28.
20.
30.
31.
32.
33.
34.
35.
36.
37.
38.
39.
40.
41.
Sb + HNO.
As4 + HNO,
As..O, + HNO,
Mn(OH) . + PbO, 4- HN03
, MnSO4 + Pb.,04 4- H S04 . dilute
MnO, + KNO, + K.CO, , fusion
S., 4- KNO, 4- K..CO, , fusion
MnS + KNO3 + K.CO, , fusion
Mn.,0,, + Pb,O4 + HNO,
Cr(OH):1 + KNO3 + KC03
fusion.
Pb3(As04), 4- Zn + H,SO4 , dilute
Cu.AsjO, 4- Zn 4- H.,S04 . dilute
Pb(NO,). + Al 4- KOH
Cu(N03). + Al 4- KOH
Bi(NO,).. 4- Al + KOH
Hg,,,0,(NOr,),, + Al 4- KOH
MnS 4- Mn(NO.,), 4- K..CO3 , fus.
Mn:iO, 4- Pb;0, 4- HNO:;
Fe + H.SO, . con., hot.
KI 4- KIO, + H S04 , dilute
MnSO, 4- KMnO, + H..SO, , dilute
(NaCl + K Cr04 + H,S04), dry,
hot
KNO, 4- FeS04 + H,S04 , con.,
cold
K,Cr,O(CrO4), + HC1 , hot
Hg*,O(NOs)8 4- Al + KOH
Ag3AsO4 + SnCL 4- HC1 , up. gr.
1.18
PbO, + K..C..O, 4- H2S04 , dilute
Pb,,O4 , white heat
NaH.PO, . ignition
Fe.O^AsO,) + FeS 4- HC1
FeBr. + HNO
Sn + HN03 , hot
KOH + Br, , hot
Fel, 4- H,SO, , .<?/>. fir. 1.84, hot
KBr + KBrO., + H,SO4 , dilute
FeS04 4- KMnO4 + H S04 , dilute
K,Cr2O(CrO,), + KOH + Br.
4Hg,0.(N,0,)3 + Al 4- KOH"
Ag3AsO3 4- SnCL. 4- HC1 , t>p. gr.
1.18
C0o03 , ignition, white he.it
H,S 4- HNO3 , sp. (jr. 1.42, not
42. Hg.,(AsO,).. + FeS + HC1
43. FehO,(As03):: + KOH + Cl,
44. Fel. + HNO3 , -v/(. fir. 1.48, hot
45. CrJSO.h + Cr(NO3)3 + K,C03 ,
fusion
40. Pb^AsO,), + Zn + H.,S04 , dilute
47. KOH -f CL. , cold
48. KBr + KIO, -f H.SO4 , dilute
49. (Cr.OHCl, + K.CrO, + H2SO4),
dry, hot
;,0. Zn4O,(NO,), + FeSO4 + H2SO« ,
concentrated, cold
51. Hg3(As04), + SnCL + HC1, sp.ffr.
1.18
.")*.'. Mn,O-, , ignition
53. Fe,O..SO, + Zn + H.,SO4 , dilute
54. Bi.S, + HNO, . dilute, hot
.15. Hg.AsO, + FeS + HC1
56. Cr,(OH),SO4 + KOH + Cl,
57. Fe(H.PO.), + HNO,
58. Cr.O., 4- KC1O, + K.CO., , fusion
59. Cu,0,(AsO4), + Zn + H,S04 , dil.
(10. KOH + Cl, . hot
til. Mn,.O,, + KC1O., + K.CO, , fusion
62. HIO, + SnCl, + HC1
(..!. Bi^O.^NO,),,, + FeSO. + H,SO4 ,
con., cold
04. CrO3 , ignition
05. KMn04 + H..C..O, + H..SO,, dilute
06. FeAsO, + SnCl, 4- HC1 , ftp. (jr. 1.18
07. Fe3CL + FeS + HC1
OS. 5CuO.As.O, + Fe + HC1
09. HIO, + H.C,O, , hot
70. t'Cr,(OH),Cl + K,Cr,07 + H,SO4),
drv, hot
71. Fe(NO,)- + FeSO4 4- H,SO4 , con.,
cold
72. Ag.,SO4 + Zn
73. H..SO, + HNO, . .s'/>. fir. 1.42
74. FeAs64 4- FeS + HC1
75. Pb(AsO) , + KOH 4- Cl.
70. Fe(NO,), + HNO,
77. Mn,O, + Mn(N03)s + K,CO, ,
fusion
78. FesO.(AsO,), + KOH 4- Br2
7!). Fb1.0,(OH).(irOs)( + Al + 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 mereuroiis 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.
2f>. 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 iotlate 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. Manganoiis 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 siibnitrate from metallic bismuth and nitric acid.
49. Barium perchlorate from sodium chloride and barium hydroxide.
50. Lead iodate from metallic lead and iodine.
398 TADLE OF tiOLL'HIUriKti. £322.
§322. TABLE OF SOLI'BILITIES.*
Showing the classes to which the compounds uf the commonly occurring elements
belong in respect to their solubility in water, hydrochloric acid,
nitric acid, or a^ua retjia.
PRELIMINARY K KM ARKS.
For the sake of brevity, the classes to which the compounds belong are
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 acid,
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 :
\V — A or w — a. ditlicultly 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 normal, 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 dichromate. W. (2) Potassium borotartrate. W. (H) Hydro-
jren potassium oxnlate. W. (4) Hydrogen potassium carbonate. W. (5) Hydro-
gen potassium tartrate, W. (6) Ammonium potassium tartrate, W. (7)
Sodium potassium tartrate. W. (S) Ammonium sodium phosphate, W. (9) Acid
sodium hornte . W. (10) Hydrogen sodium carbonate. W. (11) Tricalcium
phosphate. A. (12) Ammonium magnesium phosphate. A. (Kl) Potassium
aluminum sulphate. W. (1-4) Ammonium aluminum sulphate. W. (15) Potas-
sium chromium sulphate. W. (10) Zinc sulphide, as sphalerite, soluble in
nitric acid with separation of sulphur: in hydrochloric acid, only upon heating-.
(17) Manganese dioxide, easily soluble in hydrochloric acid: insoluble in nitric
acid. (IS) Nickel sulphide is rather easily decomposed by nitric acid: very
difficulty by hydrochloric acid. (19) Cobalt sulphide, like nickel sulphide.
(20) Ammonium ferrous sulphate. W. (21) Ammonium ferric chloride, W.
The following table of solubilities, is taken from Freseniua Qualitative Analysis, Well's
translation of 16th German edition.
,;322. TABLE OF SOLUBILITIES. 395)
(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. (2.">) Tribasic lead acetate, W. (26) Mercurius solubilis Hahne-
manni, 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.
by hydrochloric acid, nor by nitric acid, but it is dissolved by hot aqua regia.
(37) Potassium plantinic chloride, W — 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 b\- 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 SOLUBILITIES.
SOLUBILITY
Potassium.
Sodium.
Ammonium.
Barium.
Strontium.
Calcium.
a
1
a
Aluminum.
Chromium.
K
7
=
|
=
Nickel.
Cobalt.
Oxide
W W
W W
w
W-A
A
A
A&I
A
a,,
A
A
Cbromate.
W, w
w
a
w-a
w-a
u
a
w
w
a
a
Sulphate..
W,..,.
W
w,4.30.JO
I
I
W-I
W
W,,.I4
W&I,,
W
W
W
W
Phosphate
w
w.
w..,,
a
a
A,,
>«,,
a
a
a
a
a
a
Borate —
Wo
w,
w
a
a
a
w-a
a
a
a
a
a
a
Oxalate . . .
w,
W
W
a
a
A
a
a
w-a
a
w-a
a
a
Fluoride.
w
w
w
w-a
w-a
A-I
a i
w
w
w-u
a
w-a
w-a
Carbonate
w,
w,.
w
A
A
A
A
A A A A
Silicate . . .
W
W
a
a
a
a
a-i
a a a
n
a
Chloride . .
w§1
w,,
w
"SI-SB
W
W
W
W
w
W \ I \V \V
\v
W
Bromide . .
w
w
w
w
w
w
w
w
w & i w
w
w
w
Iodide
w
w
W
w
w w
w
w
w
w
w
w
w
Cyanide. . .
w .
w
w
W-ll
w ; w
w
a
A
a
a-i
a-i
FcrrocyVle
w
w
w
w-a
w
w
w
A-I
a
i
i
Ferricy'de
w
w
w
w
w
a
i
i
i
S'lphocy'de
w
w
W
w
w
w
w
w
w
w
w
w
Sulphide. .
Nitrate.. J
Chlorate . .
w
» w*
w
W
W
w
W
w
w
w
w
w
w
\V
W
w
w
a
w
w
a
w
w
a-i
\\
XV
A,,
w
\v
A
w
w
w
w
a,,
W
w
Tartrate . .
W§...t.M.lt
W7
w.
a
a
A
« a
w
w
a
w a
a
w
Citrate....
w
w
w
a
a
w-a
W
w
w
w-a
a
w
w
Miilat <•....
w
w
w
w&a
w
w-a<7
W
w
w
w
Succinate.
^w
w
w
w-a
\v a
w-a
W
w-a
w-a
w
w
w-a
Hcn/.oalc..
4-
vr
w
w
w
w
w
Salicylate.
^
w
\\
\v
w-a
u a
w-a
XX-
Acetate...
w
W
W
W
W
W
W
\v
w
w
w w
w
Formate..
w
w
w
w
•w
w
w
w
W
\\
XV w
w
Arsenite..
w
w
wfc
a
a
a
a
a
.'
a
Arsenate..
W
W
w
a
a
a
a
a
a
a
a
a J.
.
TABLE OF SOLUBILITIES.
401
TABLE.
en
a
.
3
o
j
S
B
00
g
0
a
3
'E
f*
+j
3
3
o
o
2
QJ
h
8
B
i
B
!ercu
3
'C
0.
3
3
E
00
a
•«
C3
2
"o
a
fl
=
CB
'3
1
J
a
a
to
to
33
ri
•
8
0
3
O
CS
E
OQ
02
*
a
A
a
A24
A
A
A
a
a
a
a
a&i
A«
Oxide
'.V
a
A-I
a
w-a
w
a
a
a
a
Chromate
W21,
W
W -A
A-I
w-a
W27
W30
w
W
w
w
a
Sulphate
a
A a
a
a
a
a
a
a
a
a •
w-a
Phosphate
a
a a
a
a
a
w-a
a
Borate
a
a
a
a
a
a
a
a
a
w
a
w
a
Oxalate
w-a
w
w
a
w-a
a
w
w-a
w
w
w
Fluoride
A
a
A
a
a
A
a
a
Carbonate
a
a
a
a
a
Silicate
W
Wal
I
W-I
A-I
W28
W
W-AS,
W
^36
^37-38
W
W40
W-A42
Chloride
w
w
i
w-i
a-i
w
w
w-a
W
W
tr
w-a
Bromide
W
w
i
W-A
A
A
w
a
W
a
i
w
w
w-a
Iodide
a-i
I
a
\V
a
a
W
w
Cyanide
i
I
i
a
i
i
i
Ferrocy'cle
I
w
i
w-a
i
Ferricy'de
w
w
i
a
A
w
a
w-a
a
w
Sulpbocy'de
A
a
a23
A
A
A29
a,,
a
A
a3«
a3»
OH
a41
A44.4B
Sulphide
w
w
W
W
*„
W
W
W34
w
w
Nitrate
w
w
w
w
w
w
w
w
w
w
Chlorate
w-a
w,.
a
a
w-a
a
w
a
w-a
a
84.
Tartrate
w
w
a
a
a
w-a
w
a
Citrate
\v
w-a
w-a
a
w-a
w
w
w
Malate
w-a
a
a
a
a
w-a
w
w
a
i
Buccinato
w
a
w-a
a
a
w-a
a
w
Ben/.oatc
w-a
w-a
w
Salicylate
w
W
w
W2B
w-a
w
W32
w
w
w
w
Acetate
w
w
w
w-a
w
w
w
w
w
w
Formate
a
a
a
a
a
a
A
a
Arsenite
a
a
a
a
a
a
a
a
a
a
Arsenaif
402 REAGENTS. §323.
§323. Reagents.*
During the past two years the reagents for use in qualitative chemical
analysis at the University of Michigan have been made up on the basis
of the normal solution: /. e., the quantity capable of combining with one
gram of hydrogen or with its equivalent is taken in a litre for the normal
solution. For example: Normal potassium hydroxide, KOH . requires 5G.1
grains per litre of solution (not •">(>.! grams to a litre of water), but the usual
pure product contains about ten per cent of moisture, so it is directed to
use 62..'} grams or 312 grams for a solution five times the normal strength,
•~>y. Barium chloride, BaCl._,.2H,0 , has a molecular weight of 244.2, but
the hydrogen equivalent is (244.2 -f- 2) 122.1, so for a litre of half-normal
solution. N/2, take (il 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. HC,H3O, (: 00. 5N). up. f/r. 1.04, 30 per cent acid.
Arsenic, H.AsO, .' . H.O (i:>. '/> H AsO, ~ 5).
Fluosilicic, H,.SiF.,' . §247.
Hydrobromic, HBr HO. N/2).
Hydriodic, HI (r,4, N/2).
Hydrochloric, HC1 (1*2. 5X, ftp. f/r. 1.0S4), ftp. f/r. 1.12, 24 p. c. acid.
Hydrosulphuric, H.S. saturated aqueous solution, §257, 4.
lodic, HIO, (15. %, HXO.+ 6).
Nitric, HNO, (:U5. 5N. up. gr. 1.165), *p. f/r. 1.2, 32 p. c. acid.
Nitrohydrochloric, about one part of concentrated HNO, to three parts
HC1 .
Nitrophenic, C,.H (NO,).OH (picric acid).
Oxalic, H.C O..2H.O . crystals dissolved in 10 parts water.
Phosphoric. H.FOi (10, N/2).
Sulphuric, H SO, , concentrated. «/>. f/r. 1.84.
Sulphuric, dilute (245, 5N. ftp. f/r. 1.153), one part acid to five parts water.
Sulphurous, H.SO^ . saturated aqueous solution.
Tartaric, H C.H.O., , crystals dissolved in three parts water.
Alcohol, C.H.,0 . up. f/r. O.S15, about 95 p. o.
Aluminum Chloride, A1C1. (22. N/2).
Nitrate, A1(HO.),.7%H-0 (58. N/2).
Sulphate, Al,(SO,):,.1sH,O (.">. N/2).
Ammonium Carbonate, (NH,)..CO:1 (2-'0, 5N), one part crystalli/ed salt in four
pnrts water, with one part ammonium hydroxide.
Ammonium Chloride, NH.C1 <cl^~. 5N). one part salt in eifjht parts water.
Hydroxide, NH.OH (^.">NH, . 5N. *t>. ///'. O.'.MU). .s-/>. <ir. 0.90, 10 p. c.
.,,
Ammoniiun Molybdete. (NH.) .MoO, (::r>MoO, . N/2. §75. Or/), 150 £. salt in one
litre of NH.OH . pour this into one litre of HNO^ . x/>. f/r. 1.2.
Ammonium Oxalate, (NH,).C,O,. H2O. (40, N/2), one part crystallized salt in
2 I parts water.
Ammonium Sulphate, (NH,) SO, (::3. N/2).
Sulphide, (NH,) S. colorless, three parts NH.OH . saturate with
H S and add two parts of NH.OH .
*In the greater number of cases, reagents should be "chemically pure." Different uses
require different degrees of purity. An article of sodium hydroxide contaminated with
chloride may be uiod in aim" operations; not in others. Those who have had training? in
analysis can do without specific directions, which cannot be made to cover all circumstances;
and the beginner must depend on others for the selection of reagents.
§323. REAGENTS. 403
Ammonium Sulphide, (NH4)2SX , yellow, allow the colorless to stand for some
time or add sulphur.
Antimonic Chloride, SbCl5 (30, N/2).
Antimonous Chloride, SbCl8 (38, N/2),
Arsenous Oxide, As2O3 (8, N/4), saturated aqueous solution.
Barium Carbonate, BaC03 , freshly precipitated.
Chloride, BaCl,.2H2O (61, N/2), one part salt to 10 parts water.
Hydroxide, Ba(OH)2.8H2O (32, N/5), saturated aqueous solution.
Nitrate, Ba(NO3)2 (65, N/2), one part to 15 of water.
Bismuth, Chloride, BiCl, (52, N/2, use HC1).
Nitrate, Bi(N63)3.5H,O (40, N/4, use HN03).
Cadmium Chloride, CdCL, (46, N/2).
Nitrate, Cd|N63).,.4H..O (77, N/2).
Sulphate, CdSO4.4H2O (70, N/2).
Calcium Chloride, CaCl2.6H2O (55, N/2), dissolve in 5 parts water.
Hydroxidez Ca(OH)., , a saturated solution in water.
Nitrate, Ca(NO3),.4HaO (59, N/2).
Sulphate, CaSO4.2H2O , a saturated solution in water.
Carbon Bisulphide, CS2 , colorless.
Chromic Chloride, CrCl3 (26, N/2).
Nitrate, Cr(NO3)3 (40, N/2).
Sulphate, Cr2(SO4)3.18H2O (60, N/2).
Cobaltous Nitrate, Co(NO3)2.6H2O (7?., N/2), in 8 parts of water.
Sulphate, CoSO4.7H,O (70, N/2).
Copper Chloride, CuCl2.2H,O (43, N/2).
Nitrate, Cu(NO3)2.6H2O (74, N/2).
Sulphate, CuSO4.5H2O (62, N/2). in 10 parts water.
Cuprous Chloride, CuCl (50, N/2, use HC1).
Perric Chloride, FeCL, (27, N/2), 20 parts water to one part metal.
Nitrate, Fe(NO3)3.9H2O (67, N/2).
Ferrous Sulphate, FeSO4.7H2O (80, N/2"), use a feAv drops of H2SO4 .
Gold Chloride, HAuCl4.3H2O , 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.
Lead Acetate, Pb(C2H3O2)2.3H2O (95, N/2), dissolve in 10 parts of water.
Chloride, PbCL , saturated solution, N/7.
Nitrate, Pb(NO3)2 (83, N/2).
Magnesia Mixture: MgSO, , 100 g\; NH4C1 , 200 g-.; NH4OH , 400 cc.; H2O , 800
cc. One cc. = 0.01 g-. P.
Magnesium Chloride, MgCL.6H,O (51, N/2).
Nitrate, Mg(NO3)2.6H2O (64. N/2).
Sulphate, MgSO4.7H,O (62, N/2), in 10 parts of water.
Manganous Chloride, MnCl.,.4H.,6 (50, N/2).
Nitrate, Mn(NO,) ,.6H,O (72, N/2).
Sulphate, MnSO4.7H2O (69, N/2).
Mercuric Chloride, HgCL. (08, N/2), in 16 parts of water.
Nitrate, Hg(NO3)n (81, N/2).
Sulphate, HgS04 (74, N/2).
Mercurous Nitrate, HgN03 (131, N/2), one part salt, 20 parts water and one
part HNO3 .
Nickel Chloride, NiCL.6H.,O (59, N/2).
Nitrate, Ni(N63),.6H,O (73, N/2).
Sulphate, NiSO4.6H,O (66, N/2).
Palladous Sodium Chloride, Na:.PdCl4 , in 12 parts water.
Potassium Arsenate, K3As04 (26, */, K3As04 -4- 5).
Arsenite, KAsO, (24, %"KAsO2 -^ 3).
Bromate, KBrO, (14, i//KBrO( -4- 6).
Bromide, KBr (60, N/2).
Carbonate, K,CO., C207, 3N).
Chlorate, KC10. , the dry salt.
Chloride, KC1 (37, N/2).
404 REAGENTS.
Potassium Chromate, K.CrO4 (49, X/2), in 10 parts water.
Cyanide, KCN (33, X/2), in four parts water.
Bichromate, K,Cr,O, (:5S, y,, K,,Cr,07 -r- 4), in 10 parts water.
Ferrocyanid'e, k4Fe(CN),,.:,H,O (.VI, X/2), 12 parts water.
Ferricyanide, K:,Fe(CN), (:>:>, N/2), in 10 parts water.
Hydroxide, KOH (:H2 [!)0 p. c. KOH |. 5N).
lodate, KIO3 (is, ys KIOa -r- (>).
Iodide, KI (S3, X/2), dissolve in 20 parts water.
Mercuric Iodide, K.Hgl, , Nessler's solution, §207, Gfc.
Nitrate, KNO, (:>0, X/2). the crystallized salt.
Nitrite, KNO, . the dry salt.
Pyroantimonate, KjH^St^Of.GHtO , see §70, 4c.
Permanganate, KMn04 (16, y, KMnO, -i- 5).
Thiocyanate, KCNS (4!i. N/2), in 10 parts water.
Hydrogen Sulphate, KHSO, , fused salt.
Sulphate, K,SO, (44, N/2), in 12 parts of water.
Platinic Chloride, H,PtCl«.GH,O , in 10 parts of water.
Silver Nitrate, AgNO3 (4iJ, N/4), in 'JO parts of water.
Sulphate, Ag.SO, . saturated solution, N/13.
Sodium Acetate, NaC,H,0...;H,O , in 10 parts of water.
Carbonate, Na,CO, (159, 3N), one part anhydrous salt or 2.7 parts of
the crystals. Ha,CO,.10H,O . in 5 parts of water.
Chloride, NaCl (2<». N/2).
Tetraborate, Na,B(07.l()H,O. tumif, the crystallized salt.
Hydroxide, NaOH (220 [1)0 p. c. NaOHl. 5X), dissolve in 7 parts of
water.
Hypochlorite NaClO. §270, 4.
Nitrate, NaNO, (4:1,, N/2).
Phosphate, Na,HPO,.12H,O (00, X/2), dissolve in 10 parts of water.
Phosphomolybdate, §75, (if/.
Sulphate, (:\:>. N/2).
Sulphide, Na,S , one part NaOH saturated with H2S to one part of
NaOH . unchanged.
Acid Sulphite, the dry salt.
Sulphite, Na^SO^.TH...© (63, N/2), in 5 parts of water.
Acid Tartrate, NaHC,H,On , in 10 parts of water
Thiosulphate. Na.S,O...'>H O , in 40 parts of water.
Stannic Chloride, SnCl, (:::!, N/2).
Stannous Chloride, SnCl,.2H..O (5<>, X/2), in 5 parts water strongly acid with
HC1 .
Strontium Chloride. SrCl...(iH,O (f»7. X/2).
Nitrate, Sr(NO.,)2 (">3. X/2).
Sulphate, SrSO, , a saturated aqueous solution.
Zinc Chloride, ZnCl, (:t4. X/2).
Nitrate, Zn(NOn) ,.r,H..O (74. X/2).
Sulphate, ZnSO4.7H26 (72, N/2).
INDEX.
Acetates, detection of 251
ignition of 259
with ferric salts 154
Acetic acid 249-251
estimation of 251
formation of 250
glacial 250
occurrence of 249
preparation of 250
properties of 249
'•factions of 250
solubilities of 250
Acids, detection of, notes on .'!89
displacement of weak by strong-. 180
i-ffect of concentrated sulphuric
upon 378
list of 1".
precipitated by barium and cal-
cium chlorides 388
preparation of 395
separation from bases 368
table of. precipitated by silver
nitrate 387
table of separation of 38S
Alkali carbonates, with third and
fourth group salts 142
group 221
hydroxides, action on double
cyan ides 265
hydroxides, detection of in pres-
ence of carbonates 2f>2
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... 201
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
Aluminum 142-146
acetate . v 144
compounds, ignition of 146
PAGH
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 14:»
oxidation of 140
oxide and hydroxides 143
phosphate, separation of ...145, 146
preparation of 143
properties of 142
reduction of 146
salts, reactions of 143
salts, with hydrosulphuric acid. 145
salts, with phenylhydrazin 144
separation of, from iron by
NaoS.O;; and Na,SO 145
separation of, from Cr and 4th
group by basic acetates 143
separation of, from glucinum... 196
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 23f
carbonate, in separation of As ,
Sb and Sn 119
chloride, as a reagent 231
chloride, in the third group 161
chloride, with PtCl4 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
ISDEX.
Ammonium hydroxide, detection
by Jnercuric chloride 232
hydroxide, preparation and prop-
erties of 229
molybdate. preparation of 98
molybdate, test for phosphates. 302
molybdate, with arsenic acid.... (57
oxidation of 23:5
phosphomolybdate OS
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 2:t(i
sulphate, in separation of stron-
tium and calcium 220
sulphide, as a reagent 231
sulphide, formation of 2150
sulphide, preparation of 307
sulphide, on iron and y.inc
groups 186
sulphide, yellow, formation of.. 115
sulphide, yellow, in separation of
cobalt and nickel 185
sulphide, yellow, in eupric salts. 113
test for nitric acid 281
thioncetate as a substitute for
hydrosulphuric acid 307
Analysis of alkali group 23(5
proximate 14
operations of 13, 20
ultimate 14
Anions, table of separations of... 388
An.timon.ic acid 76
distinction from ant imonous. . . . 122
reduction to antimonous by stan-
nous chloride 78
salts, action of hydriodic acid
on 78
sulphide, precipitation of 77
Antimonites 74
Antimonous argentide 79
compounds with silver nitrate. .. 78
iodide, formation of 78
<>\ide, formation of 76
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 alloys 367
Antimony, detection of >o
detection of traces of 121
distinction from arsenic TS
estimation of ^1
in the test for aluminum Ii53
metal with hydrosulphuric acid. Oil
mirror ii.">
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 79
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 6»5
Argol, purification of 2.VJ
Arsenates, distinction from arsen-
ites 70. 71
separation from phosphates. . . . 290
Arsenic r>i>-72
acid, precipitation by hydrosul-
phuric acid 114
acid, reduction by hydrosul-
phuric acid and hydriodic acid. Cl
acid, reduction with sulphurous
acid (>o
acid, with ammonium molybdate 67
acid, with molybdatcs (52
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
detection of 70
-detection of, in poisoning f>*
distinction from antimony 78
INDEX.
40"
Arsenic, 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
propert ies 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
Sb2S:1 121
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
Arsenites, distinction from arsen-
ates 121
Arsenous hydride 64
oxide, crystals, identification of. 67
sulphide, solubilities of 58
sulphide, with HC1 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
Barium carbonate, as a reagent to
precipitate chromium . .A . . . . 148
carbonate, and ferric salt . . . 15
carbonate, to separate
phates from third, four mr and
fifth groups JT. . 189
chloride, separation omi from
SrCl, and CaCL by Hcl 206
detection of /. 208
estimation of A 208
hydroxide, formation fif 20,")
iodide, properties 358
occurrence of 205
oxide, preparation of 20.~>
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
strontiiim 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 gro^p of ^, , 12
silver group of £*
sixth group of :-ty 1
third group of 1
tin group £>f 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
IXDEX.
PAGE
Bismuth hydroxide, solubility in
glycerol 101
iodid^Htability toward water... 103
nitracBprecipitation with HC1. 101
nitratSreactions 101
notes 017 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.. 101
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 102
Blowpipe, examination of solids.. 374
Blue vitriol 105
Bonds, plus and minus 238
Borates, green flame by ignition of 246
in analysis -r^. 54
reactions of ." . i. 246
Borax, bead, formation of 247
bead, test for Mn 184
bead, use of 365
Boric acid 243-247
estimation of 247
formation of 245
occurrence of 245
preparation of 245
properties of 245
solubility of 24f>
Boron 245
Bromates, detection of ::4'.)
estimation of 350
ignition of IM'.i
preparation of :; is
solubilities of .'540
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 . . 34. "i
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
Cacodyl oxide, test for acetates... 250
Cadmium 1 10-112
detection of 112
estimation of 112
hydroxide 110
notes on analysis of 129
occurrence of 110
oxide 110
properties of 110
reactions of, comparison with
Bi and Cu 112
salts, absorption by gaseous sub-
stances, separation from Cu.. Ill
salts, fused with K2S 112
salts, with alkaline tartrates,
separation from Cu Ill
salts, with alkalis 110
salts, with ammonia Ill
salts, with barium carbonate... Ill
salts, with pyrophosphates, sepa-
ration from Cu Ill
salts, reactions with Na2S..,O., ,
separation from Cu Ill
salts, reduction of by inetals.... 112
salts, reduction of by ignition.. 112
separation of 110
separation from Cu by KCNS. .. Ill
separation from Cu by glycerol. 1<).r>
separation from Cu by Na^-SzO,
and Na2SO, * Ill
solubilities of 110
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 2i:*>
group 203
group, directions for analysis of. 218
INDEX.
409
PAGE
Calcium hydroxide, formation and
properties 211
hydroxide, formation by Na,,S.. 213
hydroxide, to detect C02 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
phosphoric acid by 212
separation from Ba and Sr by
(NH4)2S04 211
separation from Ba and Sr by
amyl alcohol 211
solubilities of 211
sulphate, separation from stron-
tium sulphate 209
sulphate, solubility in ammo-
nium sulphate 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
relations of 10
solubilities of 248
dioxide 258-263
dioxide, absorption by Ca(OH)2. 261
dioxide, detection in sodium car-
bonate 262
dioxide, detection by calcium hy-
droxide 212
dioxide, distinction from H2S ,
SO2 , N2O3 , 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
Cassius' purple 93
Cerium 193
Chili saltpeter, occurrence of 277
PAGE
Chloric acid 339-341
formation of 339
preparation of 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
ignition of 340
oxidation by ignition of 341
preparation of 339
reactions with 339
solubilities of 339
Chlorides, detection of 149
detection of, in presence of bro-
mides 335, 336, 391
detection of, in presence of cy-
anides or thiocyanates. . . .335, 392
formation of 331
ignition of 334
Chloride of lime, formation of 337
estimation of, by H202 287
Chlorine 327-330
action on metals 328
as an oxidizer 328
detection of .330
estimation of 330
formation of 327
occurrence of 327
peroxide, formation and proper-
ties 338
properties of 327
solubilities of 328
Chlorochromic test for chlorides. . 335
anhydride 149
Chlorous acid,- formation and de-
tection 338
properties of 337
Chromates 150
in test for HC1 149
reduction of, by hydrochloric
acid 149
reduction of by H2S. 149
use in separation of barium 207
with antimonous salts 78
with As'" 149
with ferrous salts 158
Chromic acid, detection of 150
formation of 149
identification of 392
Chromium 147-151
distinction from aluminum 148
estimation of . . 150
410
INDEX.
Chromium hydroxide, solubility in
ammonium hydroxide :
and manganese in third group
separation 103
metal, solubility of 147
occurrence of 147
oxidation of 150
oxides and hydroxides 147
oxide, solubilities of 147
preparation of 147
properties of 147
reduction of 150
salts, solubilities of 147
salts, reaction of 148
separation from Al and Fe by
H2O2 150
separation from fourth group. . 148
separation from Fe by Na S..O
and Nn SO. 145
Chromous salts 148
Cinchonine as a test for bismuth. 102
Citric acid 251-252
detection of oxalic acid in 251
distinction from tartarie 2."il
properties and reactions 231
Colloidal sulphides of the fourth
group 184
Color, flame tests 365
Columbium, distinction from Ti.. 201
properties and reactions 193
separation from tantalum 193
Cobalt 163-168
bead test 167
cobalticyanide separation from
nickel 166
detection of 168
detection of in presence of Ni
by HaO2 185
estimation of 168
hydroxide 1 <> 1
metal, solubilities of 164
nitrate, effect of ignition with.. 365
occurrence of 164
oxidation of 168
oxides and hydroxides 1 >4
phosphate, a distinction from Ni 167
preparation of 164
proper! ies of 163 •
reduction of 168
salts, solubilities of 164
salts, with alkalis 165
salts, with barium carbonate... 16")
separation from nickel by ether 164
separation from nickel by KN02 16G
Cobalt, separation from nickel by
KMnO , 1G7
separation from nickel by ni-
troso-/?-naphthol 166, is.-)
Colloidal sulphides of fourth group 184
Color, flame tests 355
distinction from Ti 201
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 103
compounds with cyanogen 107
detect ion of 100
detection of, in alloys 367
detection of traces of, with H2S 10s
detection of, with HBr 108
electrical conductors 104
estimation of 109
ferroeyanide, formation of 107
group, metals of 56, 100
hydroxide of 104
occurrence of 104
oxides of 104
precipitation of, by iron wire... 109
preparat ion of 104
properties of 104
reactions of, comparison with
Bi and Cd 112
reduction by ignition 109
reduction of, by KCNS 107
salts, detection by potassium
xanthate 107
salts, reaction with zinc-plati-
num couple 109
salts, reduction of, with H:1PO2. 107
salts, separation of, from Cd by
Na.P.O, 107
salts, solubilities of lO.j
separation of, from Bi by gly-
cerol 101
separation of, from Cd by gly-
cerol 105
separation of, from Cd by
NajS.O, and Na,C03 Ill
separation from Cd by nitroso-
0-naphthol 107
separation from Cd by ammo-
nium benzoate 107
separation from Pd 106
traces, loss of 115
traces of, with KiFe(ClT) 107
INDEX.
411
PAGE
Cream of tartar, formation of 252
Cupr ammonium, salts 106
Oupric hydroxide in NH4OH 105
hydroxide, effect of boiling 10(5
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 Na2S2O3 .... 108
salts, reduced by SO, 108
sulphide colloidal 10S
sulphide, formation of 107
sulphide, separation from Cd by
H2SO4 108
sulphide, solubility in (NH4)2SX. 108
sulphide, solubility in KCN 108
sulphide, with K2S ; . . 115
sulphide, with (NH4),Si 115
Cuprous iodide 108
oxide, formation of, by glucose. 105
salts, oxidation of, by As.,0....... 110
salts, separation, from Cd by S. 107
salts, with metallic sulphides... 107
sulphide, formation by Na2S20:1. 108
thiocyanate, formation of 107
Cyanates, detection of, in presence
of cyanides 271
Cyanic acid 271
Cyanide of silver, distinction from
chloride 2<>.>
Cyanides, detection as thiocyan-
ate 267
double, dissociated by acids 261
double, not dissociated by acids. 265
estimation of 267
guaiacum test 267
ignition of 266
preparation of 264
reactions with 264
simple, with mineral acids 265
solubility of 264
transposition by acids 267
Cyanogen properties and reac-
tions 263
Danger and Flandin, detection of
arsenic 69
Decomposition of organic mate-
rial 362, 363
Dialysis, separation from organic
material by 363
Didvmium 194
DnJhenylamine test for nitric acid 281
Dissociation, electrolytic 20
Dithionic acid, formation and
properties 314
Dragendorff' 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 153
and ferrous compounds, distinc-
tion 162
hydroxide, antidote for arsenic. 62
phosphate, formation of 156
salts, detection of traces 155
salts, with acetates 154
salts, with BaCO:i 153
salts, with HI and iodides 158
salts, with H2S 157
salts, with H3PO2 156
salts, with K;Fe(CN)6 155
salts, with K4Fe(CN)8 155
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 Ni T66
reactions of 270
Ferrocyanides, detection of 269
detection and estimation 271
reactions of 268
Ferrous iron, detection of, in
ferric salts 153
in the third group 161
in the third group with phos-
phates 189
salts, traces in ferric salts 155
salts, with chromates . , 158
salts, with HN03 156
salts, with KCN «, ». 154
412
INDEX.
PAGE
Ferrous salts, with KnFe(CN), 155
-salts, with K4Fe(CN),, 154
sulphate, with gold salts 93
First group metals, table of 52
Fixed alkalis 221
alkali hydroxides on stibine 79
alkalis with salts of tin 84
Flame, blowpipe, production of... 364
or color tests 373
oxidizing and reducing1 363
reactions with copper salts 109
Fluorides, solubilities of 289
Fluorine 288
Fluosilicates, formation of 289
Fluosilicic acid 247-248
in detection of potassium 225
in separation of Ba , Sr and Ca. 207
Formates, formation from cyan-
ides 266
Fourth group, directions for anal-
ysis 184
reagents 141
sulphides colloidal 184
table of 183
Fresenius and Babo, detection of
arsenic 68
Froehde's reagent 99
Fulminating gold 92
Gallium (oka-aluminium) 195
Gases, absorption of by palladium 131
Germanium, properties and reac-
tions 1 36
sulphide 118
Glass, etching1 by hydrofluoric
acid 289
Glauber's salts 304
Glucinum (Beryllium) 195
distinction from yttrium 202
separation from aluminum 190
separation from cerium 193
Glucose, in formation of cuprous
oxide 105
Gold 91-93
Detection in alloys 367, 368
!et$ction of 93
distinction from Pd 132
estimation of 93
fulminating1 92
notes on analysis 123
occurrence, properties, etc 91
reduction by ferrous sulphate... 93
reduction with oxalic acid 92
salts with alkalis 92
salts with stannous chloride.. 89
Gold, separation from Ir 133
Greenockite 101, 110
Gypsum 213
Halogens y
as oxidizers 330
compounds, comparative table of 361
hydracids as reducers 330
Heat, upon substances in closed
tubes :564. 370
upon substances in open tubes.
364, 371
Hydriodic acid 353-356
action on antimonic salts 78
action on arsenic salts 61
action on ferric salts 158
as a reducer 354, 355
formation of 353
Hydrobromic acid 344-348
detection of Cu with 108
formation of 345
occurrence of 345
preparation of 345
properties of 345
reactions of 345
Hydrochloric acid 330-330
action on Sb S 77
action on bismuth nitrate 102
effect of excess in second group. 113
formation of 331
formation from MgCl, 216
gas on arsenic sulphide 07
occurrence of 331
preparation of 331
properties of 330
reactions with 332
solubilities of ."-31
Hydrocyanic acid 263-267
formation of 204
occurrence of 264
on PbO, 204
preparation of 264
properties of 203
solubilities of 264
Hydroferricyanic acid 269-271
Hydroferrocyanic acid 267-269
separation from hydroferri-
cyanic acid 269
Hydrofluoric acid 289
Hydrofluosilicic acid (fluosilieic
acid) 289
Hydrogen 243-244
absorption by Pd sponge 131
detect ion of 244
cst inisil ion 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 of 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 Ni 185
peroxide, separation from ozone 237
peroxide, separation of Al , 3Te
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 antimony 67
gas on arsenic 67 1
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 30(5
uses as a reagent 308
with arsenic acid 114
with oxidizing agents 114
Hydrosulphurous acid 314
Hydroxylamine, formation and
properties 278
Hypobromous acid, formation and
properties 348
Hypochlorites, detection of 392
formation of 337
Hypochlorites, . formation from
chlorine 32!)
on arsenic 66
Hypochlorous acid 337
Hypoiodous acid, existence of.... 351
Hyposulphites, detection of 296
ignition of 290
Hypophosphites in formation of
PH3 290
Hypophosphoric acid 29S
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
Hyposulphurous 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
lodic acid 357-359
formation of 357
preparation of 357
properties of 357
reactions of 35S
Iodide of nitrogen 351
Iodides, decomposition by HNO:! . 281
detection as Pdl, 131
detection of 356
estimation of 356
formation of 353
ignition of 356
occurrence of 353
reactions of 354
separation of, from bromides
and chlorides by KMnO4 176
solubilities of 353
Iodine 350-352
detection of 352
estimation of 352
formation of 351
liberation by copper salts 108
occvirrence of 351
on antimonous salts 78
on antimony 66
414
INDEX.
PAGE
Iodine, on arsenic 66
preparation of 352
properties of 350
reactions of 351
separation from Br by Pd 131
solubilities of 351
lens 21
lonization and solution 20, 24
Iridium 132-133
Iron 151-159
and zinc groups 140
detection of 162, 163
detection of traces in copper... 154
detection of traces 154, 155
estimation of 159
group 142
group, separation from Co, Ni ,
and Mil by ZnO 158
hydroxides i52
in relation to metals 6
occurrence of 151
oxidation of 159
oxides 152
preparation of 151
properties of 151
reduction 159
salts, ignition of . ... 158
salts, solubilities of 153
salts, with alkalis 153
salts, with nitroso-/?-naphthol. . 154
salts, separation from Al as
basic nitrate 158
separation from Al and Cr by
nitroso- /?-naphthol 154
separation from Cr and Al 154
separation from Ni by xanthate 170
solubilities of 152
Lanthanum 197
Lead 29-36
acetate, properties of 32
chloride 34
chloride, precipitation of 53
chromate, formation of 3.5
compounds, ignition of 35
detection in alloys 367
detection of 36
estimation of 36
in the test for Al 163
iodide, formation and proper-
ties 35
notes on analysis of 127
occurrence of 29
oxidation and reduction 36
oxides of 29
Lead oxides', solubilities of 30
preparation of 29
properties of 29
red 29
relation to nitrogen family 7
suits, 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 54
Leblanc-soda process 259
Lithium 234-23fi
Lime, slacked 211
stone ( CaCO ) 213
Light, action on silver salts 50
Magnesia mixture 14.r>
Magnesium 214-210
as a reducing agent 21(5
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
propert ies of 214
removal for detection of sodium. 236
salts, with ammonium salts 215
salts, with arsenic acid 01
salts, with Na S 215
salts, solubilities of 214
Malachite 104
Manganates, identification 392
Manganese 172-177
detection of, 176, 18(5
est imation of 177
hydroxides of 172
hydroxides, solubilities of 17'.',
ignition of 176
in third group 161, 163, 184
occurrence of 172
oxidation of 177
oxidation to permanganic acid.. 175
oxides 172
oxides, solubilities of 17.'!
preparation and properties 172
reduction of 177
reduction by sulphites 175
salts, reactions with oxalic acid. 174
salts, solubilities of . . 173
INDEX.
415
Manganese salts, with alkalis.... 174
salts, with sulphides 175
separation from zinc with acetic
acid 184
solubilities of 173
with KI 176
Manganic acid 172
Marsh's test (-2
Mass action, law of 22, 38
Mayer's reagent 41!, 232
Mercurammonium compounds ... 39
Mercuric chloride with stamious
chloride 88
sulphide, formation and proper-
ties 41
sulphide, with K2S 115
Mercury 37-45
chlorides 42
compounds, ignition of 43
detection and estimation of 44
iodides 42
metallic, analysis of 367
occurrence of 37
oxidation of 45
oxides 37
preparation and properties of... 37
salts, reactions 39, 43
salts, solubilities of 38
solubilities of 37
sulphide, analysis of 126
Metals, classification 10
grouping 375
table of separation 376
Metaphosphoric acid 299
Metastannic acid 83
Kicrocosmic salt 230
use in ignition 365
Kilk of lime 211
Molybdates in analysis 54
with phosphates 98
Molybdenite 97
Molybdenum 97-99
deportment in second group 99
detection of 99, 122
estimation of 99
ignition tests 99
notes on analysis of 123
occurrence of 97
oxides and hydroxides 97
preparation and properties 97
reduction tests 99
solubilities of 97
Molybdic acid 97
Nascent hydrogen on nilric acid.. 278
Neodymium 194, 197
Nessler's reagent 43, 231
Nickel 168-172
detection of 171
detection of, in presence of Co
by KI 185
distinction from cobalt 170
estimation of 171
hydroxides 169
hydroxides with KI 171
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-
/3-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 NiS in ammonium
sulphide 170
xanthate, separation from Fe. . 170
Niobium (Colunibium) 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 HC1 279
detection of 280
detection by diphenylamine 281
detection by reduction to NH .
278, 281
detection by reduction to nitrite. 281
dissociation, by heat 279
estimation of 282
formation of 277
indigo test 281
in separation of Sn , Sb and As. 119
sodium salicylate test 281
with phenol 281
with pyrogallol 282
wit h brucine 281
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
Nitrof erricyanides 270
Nitrogen 273-274
chloride 62, 120, 327
combination with elements 274
detection and estimation 274
family 7
formation, occurrence 274
peroxide 277
properties 273
Nitroso- i-naphthol, separation of
Co and Ni 166, 185
separation of Cu from Cd 107
with iron salts 154
Nitroprussides 270
Nitrous acid .276-277
as an oxidizer 276
as a reducer 276
formation of 276
occurrence of 276
proper! ies 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
on 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 •> 255-258
as a reducer 2.">ti
decomposition of by H.SO, 2">7
formation of 255
in separation of gold 92
occurrence of 255
preparation and properties of... 255
solubility of 256
Oxidation, balancing equations in. 238
Oxidizing flame 363
Oxygen 282-284
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 HO 287
Palladium 131-132
distinction from gold and plati-
num 131, 132
separation from copper 106
sponge 131
Palladous iodide in analysis 131
Paris green 62, 10S
Pentathionic acid, formation and
properties ::it>
Perchlorates, preparation and
properties 341, 342
Perchromic acid l.r>l
Periodic acid ::r>i)
system, table of 2
Permanganates, identification ::i»i
action on ant imonous salts 7*
Permanganic acid 1715
Persulphuric acid 326
Phenol reaction for nitric acid....
Phenylhydrazine, on aluminum
salts m
Phosgene, formation •-'"•!
Phosphates, changes by ignition
detection 162, 30.:. :.'.•»
distinction between primary.
secondary and tertiary n|
estimation of '•(>!
in presence of third and fourth
group metals. 142, 188, 189. KM
occurrence of '-"•'•>
reaction with ammonium molyl>-
date L88,
separation as ferric phosphate..
INDEX.
41?
Phosphates, solubilities of 300
Phosphides, formation of 303
Phosphine 295
Phosphoric acid 298-304
preparation of 300
properties of 298
Phosphoric anhydride, formation
of 299
Phosphorous acid 297-298
detection of 298
preparation and properties of... 297
Phosphorus 293-295
detection and estimation of 295
•" in combination with the halo-
gens 294
occurrence and preparation of.. 291)
properties of 292, 294
use in match-making 293
Phosphotungstates 135
Picric acid, in detection of potas-
sium 224
Plaster of Paris (calcium sul-
phate) 213
Platinized asbestos 94
Platinum 93-97
apparatus, care of 95
black 93
chloride, as a reagent 95
detection of 96, 122, 367
distinction from palladium. .131, 132
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 226
ierricyanide, formation of 269
ferrocyanide, formation of ..265, 267
hydroxide, as a reagent 223
iodate, in separation of alkaline
earths 207
PAGB
Potassium iodide, as a reagent 224
iodide, in separation of AgCl
from SbCl3 120
iodide, in the test for nickel.... 185
iodide, on nickelic hydroxide.... 171
iodide, on permanganates. ...... 176
nitrite in separation of cobalt
from nickel 166
occurrence, preparation and
properties of 222
picrate 224
pyroantimonate 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 403
Reducing flame, description of... 363
Reduction, balancing equations in 238
with charcoal 364, 365, 371
Reinsch's test for arsenic 07
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 197
Scandium, properties and reac-
tions . . 196
418
INDEX.
Scheele's green and Schweinfurt's
green 62, 108
Selenic acid, separation from sul-
phuric acid 139
Selenium, 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
cyanate in distinction from chlo-
rides 271
detection of 50, 367
estimation of 50
in presence of mercury salts.... 55
iodate, properties of 358
mirror, formation by tartrates. 25.'5
nitrate, action on stibine 79
nitrate with stannous and anti-
monous salts 78, 79, j88
occurrence and properties of.... 45
salts, action of light upon 50
solubilities of 46
thiocyanate, separation from
silver chloride 272
Soda lime on stibine 79
process. Le Blanc's 259
process, Solvay's 260
Sodium 226-229
amalgam, action with arsenic... 64
as a reducing agent 229
detection of 73, 228
estimation of 229
flame test 228
hydroxide, formation of 227
nitroferricyanide as reagent. 230, 311
occurrence of 227
phosphate as reagent 227
Sodium phosphomolybdate as re-
agent 98, 232
preparation and properties of..
226, 227
pyroantimonate 73, 80
pyrophosphate with copper and
cadmium 107
salicylate test for nitric acid... 281
sulphide, preparation of 308
thiosulphate on cupric salts.... 108
thiosulphate with antimony salts 78
thiosulphate with third group
. metals 145
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 98
Stannic salts, solubilities 84
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 reducing1 agent.... 88
chloride with gold salts 93
chloride with molybdic acid 99
salts, distinction from stannic
salts 12.'>
salts, solubilities 64
salts with silver nitrate 87
salts with sulphurous acid 86
sulphide, formation and proper-
ties 85
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, 21.'i
estimation of 210
hydroxide, formation 20*
occurrence of 208
INDEX.
419
PAGE
Strontium, preparation and prop-
erties of 208
sulphate, distinction from CaSO , 209
. sulphate, separation from BaSO, 209
Sulphates, detection and estima-
tion of 32f>
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 cadmium 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
occurrence of 318
preparation and properties of.. 318
formation of 318
Sulphurous acid, reduction of
cupric salts 108
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.
Tellurium, distinction from sele-
nium 138, 140
properties and reactions of 137
separation from sulphuric acid.. 137
Tenorite 104
Terbium 198-190
Tetrathionic acid, formation and
properties 315
Thallious iodide 199
Thallium, properties and reac-
tions 199
Thioacetate in formation of sul-
phides 307
Thiocyaiiates, reactions with 272
Thiocyanic acid as a reducer 273
properties of 272
Thionic acids, table of compari-
sons 317
Thiosulphates. detection of 313
distinction from sulphates and
sulphites 314
estimation of 314
ignition of 313
formation and properties of. . . . 312
Thiosulphuric acid 312-314
Third group reagents 141
Thorium 199-200
Tin 82-89
creaking of 82
detection of 88, 122, 3G7
estimation of 88
Group, metals of .">(>
Group, separation from Copper
Group 1 l.i
Group, sulphides with (NHJ.Sx 11. "»
occurrence of 82
oxidation of 88
oxides and hydroxides H2
preparation and properties of... s-j
notes on the analysis of I-':;
relation to Nitrogen Family 7
reduction by ignition 87
salts with the alkalis 84
salts with hydrosulphuric acid.. 85
separation from antimony ..... 81
separation from antimony sul-
phides 121
separation from arsenic US
solubilities of 83
sulphides, colloidal 1 1 r,
Tin with antimony and with arsenic 87
Titanium 200-201
distinction from columbium 201
distinction from tantalum 198
properties and reactions of 200
separation from thorium 200
Trithionic acid, formation and
properties 315
Tungsten, properties and reac-
tions 134
TurnbulFs blue 155
Unit of quantity 22
Uranium, properties and reactions 201
Urea, from ammonium cyanate... 271
Valence, negative 3
Vanadium 135-136
Volatile alkali (ammonium hy-
droxide) 221
Water, action on bismuth salts... 101
action on antimonous salts 75
Welsbach burners 203
Wolframium (tungsten) 134
Wulfenite 97
Ytterbium properties and reac.
tions 202
Yttrium 202
Zincates, formation of 179
Zinc 17!j-181
detection and estimation of 180
Family 5
granulated 63, 178
Group, table for analysis 183
Group, comparative reactione. . . 182
hydroxide and oxide 178
ignition of 180
occurrence of 178
oxidation of 181
platinized 178, 243
preparation and properties 178
reduction of 181
salts, solubilities and reactions
of 179
sulphide, formation in presence
of acetic acid 179
Zirconium . 202-203
The methods described in this Appendix have been in use in the Uni-
versity of Michigan for the past two years, and have proved to be
thoroughly reliable. They are based fundamentally on older methods,
which have, however, been so modified as to make them more satisfactory
for qualitative tests.
The modifications are due to Mr. Hobart H. Willard, Instructor in
Qualitative Analysis in this University.
APPENDIX.
SEPARATION OF ARSENIC, ANTIMONY, AND TIN.
These metals are precipitated from a N/5 HCI solution by H..S and therefore
belong to the 2nd group, but their sulphides, unlike those of Pb Cu , Cd , Bi ,
are soluble in yellow ammonium sulphide, (NH4)2SX, and are thus separated
from them before treatment with 2NHNO3. This operation is necessary only
when both divisions of the group are present, and is to be avoided when
unnecessary. Hence a little of the 2nd group precipitate is tested by warm-
ing with 1 or 2 cc. (NH4),,SX. If it all dissolves, only As, Sb , Sn can be
present; if nothing dissolves, none of these can be present; if part dissolves,
then the whole 2nd group precipitate must be so treated. To see if anything
has dissolved in the (NH4),,SX it is acidified slightly with HCI (test with litmus) ;
a milky, white precipitate of S will always be formed, but if any sulphides are
present they will appear as a flocculent, colored precipitate. If the whole 2nd
group precipitate is treated with (NH4)2SX, the solution is filtered and acidified
just as the test portion was, the precipitated sulphides of As , Sb , Sn well
washed with hot water and removed from the filter to a casserole by a spatula,
or, if the amount is small, treated with the filter; a convenient amount of
concentrated HCI (sp. gr. 1.2) is added and boiled a minute or two to expel
H,S . The sulphides of Sb and Sn are dissolved to form the chlorides SbCI3 and
SnCI, while the As,S3 is hardly attacked. Since the strong acid attacks the
filter, the solution is diluted a little, which should cause no reprecipitation if all
H2S was expelled, filtered, and the residue well washed. It may be either As^Sj
and S, or S alone. A few cc. of warm NH.OH are poured over it, the solution
being passed through again if necessary. The As2S, dissolves and the S remains.
To the solution, which must be clear, add 1 or 2 cc. H2O2 , 2 to 3 cc. NH4CI , and
2 to 3 cc. " magnesia mixture," which is MgCI, -j- NH4CI -|~ NH.OH . Cool, and let
stand for a time. The As ' ' ' ' ' is precipitated as NH4MgAsO4 , a white, crystalline
precipitate exactly like NH,MgPO4 in appearance. This "magnesia mixture" is
not intended to take the place of the " Marsh apparatus " but to confirm its
results.
2H3AsO3 + 3H,S -f [ HCI1 — As.S, -f 6H,O .
2H3AsO4 -f 5H,S + [ HCI] = As2S3 -f S, + 8H2O .
Sx = 4(NH4)sAsS4-j-(3x — 5)S2.
As2S0+ HCK12N, hot) =no action.
As,S5 + 16NH4OH -f- 20HA, = 2(NH4)3AsO4 + 5(NH4),SO4 + 28H2O.
(NH4)3As04 -j- MgCI2 -f [NH4OH + NH4CI] = MglMH4AsO, + 2NH4CI .
The filtrate from As2S, is to be tested for Sb and Sn . For the Sb , place a few
drops on a clean silver coin; it should produce no discoloration. A piece
of tin, bent into the shape of a broad U, is now placed on the coin so that
one end is in the center of the drop and the other in contact with the silver
outside. Allow to stand about 5 minutes. If Sb is present it will be deposited
as a brown spot on the silver covered by the drop, the Sn and Ag acting as a
424 APPENDIX.
galvanic couple to reduce the Sb ' ' ' to metal. Another test consists in treating
the solution with pure, fine Fe wire, the Sb being precipitated in black metallic
form, while the Sn ' " ' is merely reduced to Sn ' ' but not precipitated.
Test the rest of the solution for Sn by heating with fine Fe wire until the
solution is colorless or greenish, with no trace of yellow, to make sure
that all the Sn ' ' " is reduced to Sn " . Ten minutes or more may be required.
Filter and add the filtrate slowly (a few drops at a time) to a few cc. of ammo-
nium molybdate, (NH4),MoO4 solution. A deep blue color or precipitate will
appear if Sn " is present, due to the reduction of the MoO;; to a lower oxide. Or,
instead of adding this filtrate to molybdate solution, it may be treated with
HgCL, a white precipitate of HgCI being formed if Sn " is present. Note that
this is reversing the test for Hg " with SnCI, . Remember that it is always
necessary to reduce the Sn ' ' " to Sn " with Fe , since the former cannot reduce
the HgCI;.. or (NH4),MoO4. The HgC!2 test is most characteristic. The precipi-
tation of As._.S;f unlike that of the other sulphides, is not prevented by the pres-
ence of any amount of HCI , however large, but, on the contrary, is aided. It
may, therefore, be necessary, after removing all other sulphides in the N/5 HCI
solution, to add several cc. of concentrated HCI , heat to boiling, and pass in
H,S for some time to precipitate the rest of the As. In the cold, H:tAsO4 is very
slowly precipitated by H,S , but strong HCI and heat accelerate the reaction
very much.
It is essential that the sulphides be thoroughly washed before treatment
with HCI .
CuS is slightly soluble in (NH.K.S, and may give a coloration when the solu-
tion is acidified.
(NH,),S, which is colorless, gives no precipitate of S upon addition of
excess of acid; (NH,),SX . yellow, always gives more or less S , white and difficult
to filter.
2(NH,)I8. -f- 4HCI = 4NH4CI + 2H:S-}-(x — 1)S=.
(NH4),S -f 2HCI = 2NH4CI-f H,S .
Make a blank test on the Fe wire used, to see that its solution in HCI gives no
test for Sn with molybdate.
Additional Tests for Bismuth. — First. — Dissolve the separated hydroxide in
hydrochloric acid, then one drop of iodide of potassium will produce a black color
and a larger quantity forms a yellow solution not decolorized by stannous
chloride. A very delicate test. The yellow color of ferric salts and the green of
copper salts are removed by stannous chloride.
Second. — To the separated bismuth hydroxide add sodium hydroxide and for-
maldehyde. Metallic bismuth is produced. 2Bi(OH):, -f 3HCHO -f SNaOH =
2Bi -f SHCOONa 4- 6H2O. Since sodium stannite reduces lead hydroxide to
metallic lead and formaldehyde does not, it is preferred as a reducer. Owing to
imperfect separation traces of lead might be present at this point.
Separation of Third Group Cations, Al ' ", Fe ' ", Cr " '.
Boil the filtrate from the second group to expel H2S ; any Fe present is in the
ferrous condition. Since the precipitation of Fe " by NH4OH is not quite com-
plete, it must be oxidized to Fe ' ' '. To effect this, add 5 to 10 drops HNO3 and
boil till there is no further change of color, even upon adding another drop of
acid. A slight, white precipitate here is sulphur from oxidation of H2S not
APPENDIX. 425
expelled by boiling. Add to the solution half its volume of NH4CI , heat to boil-
ing, and add NH4OH , two or three drops at a time, till there is a slight excess.
Boil a minute or two, and at once filter and wash. If precipitated in a boiling
solution the hydroxides filter more rapidly. The precipitate consists of Fe(OH)3,
AI(OH)3, Cr(OH)s, and more or less Mn(OH)3 ; sometimes all the Mn is precipi-
tated here. The more rapidly the precipitation and filtration are performed, the
less Mn(OH)3 will be included; hence never delay at this point. Wash the
precipitate with hot water, stirring it up with the jet from the wash-bottle. If
the filtrate is amethyst colored and no Co " is found, too much NH4OH was added,
and a trace of Cr(OH)3 has dissolved. Further boiling may precipitate this, but
often the color remains. If the third group precipitate is greenish or black and no
Cr " ' is present, it shows that not all the Fe ' ' was oxidized, and the precipitate
must be dissolved in HGI and again treated with HNO3.
The precipitate is transferred to a casserole or beaker, either by means of a
spatula or with a fine jet of water, but any considerable dilution is to be avoided.
Add 5 to 10 cc. NaOH , 10 to 15 cc. H2O2 and boil for several minutes — until
effervescence ceases. The H2O2 oxidizes the Cr " ' to CrO4", while the Al ' '
goes into solution as the ion AIO/. The precipitate then consists of Fe(OH)3
and some Mn(OH)3. Filter and wash with hot water.
The filtrate, yellow if Cr is present, and containing NaAIO2 and Na2CrO4, is
acidified with dilute H2SO4 . (Test with litmus.) This converts the AIO/ ion into
Al' ' ' and CrO4" into Cr2O7", the latter having an orange color. Divide the
solution into two parts; cool one under running water and add a little H2O2 ;
a dark and beautiful blue color, due to perchromic acid, HCrO4 , is formed if Cr
is present; the color soon disappears, becoming green, and the more rapidly the
warmer the solution. It is not formed at all in a warm solution. To the other
portion add excess of (NH4)2CO3, warm, and if no precipitate appears at once
set aside a few minutes. A white, gelatinous, flocculent precipitate of AI(OH)3
will be formed if Al ' " is present. This precipitate is sometimes almost trans-
parent; it is nearly always very light and liable to escape notice.
The precipitate of Fe(OH)3 and Mn(OH)3 is dissolved in dilute H2SO4 with the
addition of a few drops of H2O2 if solution does not readily occur, and a few
drops of this solution are tested for Fe with KCNS or K4Fe(CN)8. Test the
solution for Mn " " by adding Pb3O4 and HNO3, boiling several minutes and allow-
ing to settle without filtering. If it shows a reddish purple color, this is due to
the presence of HMnO4 formed by the oxidation of the Mn". Test the Pb3O4
for Mn by boiling it with HNO3 ; some samples give a slight color.
If, in the original solution, Cr,O7" or CrO4" is present, it must be reduced to
Cr " ' before precipitating the second group. Add considerable HCI and several
cc. of alcohol, and boil till the color changes to green. Be sure that reduction is
complete. Boil down to expel all the alcohol, then neutralize and proceed in the
regular way for the second group. If not reduced by alcohol and HCI , H2S will
at once give a green color, a large amount of sulphur being liberated, difficult to
filter.
Do not fail to test the reagent NaOH and H2O2 for A I .
The fact that a solution is green does not exclude the possibility of some Cr.O7"
being present; but a very small amount of Cr " ' will give a distinct color to the
orange Cr2O7" solution. If further proof of the presence of Cr " ' is desired, the
Cr,O7" may be precipitated by making alkaline with Na2CO3, adding excess of
acetic acid (this is to avoid the presence of free mineral acid), and then BaCI2
426 APPENDIX.
or Pb(C2H3O2)2. Cr"" remains in the filtrate, the anion being precipitated as
BaCrO, or PbCrO4 .
The reagent NaOH may contain Al or silica, which might give the test if no
Al " - was present in the solution being analyzed. Test, therefore, about 10 cc.
of it by acidifying with HCI and adding excess of (NH4).,CO3. Note amount of
precipitate formed, and do not report AT * ' unless a heavier precipitate than this
is obtained in the analysis. This is known as a " blank test"; it is often neces-
sary to apply such checks in analytical work, e. g., testing the Pb3O4 for Mn .
Strong solutions of NaOH or KOH attack and soften filter paper, interfering
with filtration. If this difficulty is experienced, dilute with water.
Remember that it is impossible to keep a solution of Fe ' ' free from Fe " "
unless all oxygen is carefully excluded, so that the latter is always associated
with the former.
NEVER FAIL to test the precipitate remaining after treatment with NaOH
and H2O2 for Mn ; it may all appear here and none in the fourth group.
ANALYTICAL EQUATIONS.
Cr'"
Reduce Cr,O," or CrO," to Cr * * ' thus:
K2Cr207 + 3C2H.O + 8HCI = 2CrCI3 + 3C2H4O -f 2KCI + 7H2O .
CrCl, -f 3NH4OH =Cr(OH)3 -f 3NH4CI .
2Cr(OH)3 + 4NaOH -f H2O2 = 2Na:CrO4 -j- 8H,O .
2Na2CrO4-f-4H:SO, -f H:O2 = 2HCrO4 + 4NaHSO4 -f 2H.O . (Cold solution.)
2Na2CrO4 + 2HNO3 = Na-CrO, + 2NaNO3 -f H,O .
Na2Cr2O7 -f 2 NaOH = 2Na,CrO4 + H,O .
2Na,CrO4 -f 5H-SO, + 3H2O2 (warm) = 2Na2SO4 + Cr:(SO4), -f- 3O2 f 8H2O .
Separation of Nickel and Cobalt.
Dissolve the CoS and NiS in HCI with a small crystal of KCIO,, boll, filter,
add two cc. of NH.CI, and NH4OH in large excess, at least ten cc. more than
enough to make the solution alkaline. Dilute to at least 25 cc. in a beaker or
casserole, add .2 to .3 gram potassium persulphate, K2S..OS , and warm, with fre-
quent stirring, until crystals dissolve. Boil for some time to expel most of the
ammonia. The blue color of the (NH3),NiCI2 does not change, but the cobaltous
compound is oxidized to a complex cobaltic ammonium salt, the color changing
to red. There should be no precipitate in the solution now (if there is, either
the solution was not dilute enough or too little NH4OH was added. Dissolve in
HCI and repeat process without adding more NH4CI).
Add a few more crystals K2S2OS and boil again to make sure that oxidation is
complete. When no further change of color occurs, cool under the tap to room
temperature, add 10 to 15 cc. NaOH and shake. If Ni is present it will be
precipitated as Ni(OH)3, dark brown, turning black. The precipitate may form
slowly, and requires some time for complete separation. Let stand at least
15 minutes; filter, and if Co is present, the filtrate will be pink or red. The
amount of NaOH necessary to precipitate the Ni depends on the amount of
NH4CI and NH4OH present: if a very large excess is present, more than 10 cc.
NaOH may be required. No Co(OH)3 is precipitated unless the solution is
warm.
APPENDIX. 437
Separation of 01', Br', and I', by the Persulphate Method.
To ten cc. of the original solution, add slight excess of Na2CO., , free from
chlorine, and boil, to precipitate the heavy metals. The solution must react
alkaline. Filter and add to the nitrate acetic acid, several cc. more than enough
to neutralize it, dilute to 50-60 cc., add about one-half gram of K2S2O8 , and heat.
If an iodide is present, free iodine will be liberated, and may be identified by
shaking a few drops of the solution with CS2 . Boil in a casserole until all iodine
is expelled, which should require three to four minutes. If action is slow, more
persulphate should be added. When the solution is colorless, add a few more
crystals of persulphate and boil again, to make sure that no iodine remains. As
the solution evaporates add distilled water to maintain the original volume. To
remove Br' add two cc. of H.,SO4 , previously diluted with water, a little more
K^S2OS , and heat to boiling point, but do not boil. A yellow or red coloration, if
the separation of I has been properly conducted, indicates Br . Pour a little of
the solution into a test tube, cool, and shake with CS2 , which should be colored
yellow or red but not violet, which would indicate that the I had not been com-
pletely removed. If bromine is present, add one-half gram of K.SoOj, to the main
part of the solution, and boil until it is all expelled and the solution is colorless;
then test with a little more K2S2OS and boil five minutes longer to make sure of
the complete expulsion of the bromine. BE SURE THAT THE VOLUME OF
THE SOLUTION DOES NOT FALL BELOW 50 TO 60 cc. Add distilled water
from time to time to replace that lost by evaporation. When all bromine is
removed, cool and add a few drops of silver nitrate; a white, curdy precipitate
of silver chloride indicates the presence of Cl . If too much silver nitrate is
added, a white crystalline precipitate may be formed, but will dissolve upon
dilution and warming. '-'
If ClO'j is present, the above procedure cannot be followed, for the I' would
be oxidized to IO'3 . In this case it is necessary to precipitate the Cl', Br', and I'
by adding to the original solution excess of silver nitrate and then nitric acid;
this effects a separation, silver chlorate being soluble. Wash the precipitate of
AgCI , AgBr, Agl , transfer to a test tube, add a piece of zinc, a little water,
and a drop of sulphuric acid. Let it stand until it is perfectly black all the way
through, showing complete reduction to metallic silver. Filter and treat the
filtrate containing ZnCL, ZnBr2, ZnL, according to the above method, starting
at the beginning. Even if no heavy metals are present, Na.CO, should be added
to neutralize any mineral acid that may be present and to form some sodium
acetate when acetic acid is added.
The persulphate method should be used only when the presence of I' or Br'
has been proved by some short test (H2SO4, Cl , HNO2, HNO3, or other oxidizer).
In presence of a great excess of Br', CuSO4 , KNO,, or HgCI is an excellent test
for I'.
EQUATIONS.
2KI 4 K2S2O8 = 2K2SO4 4. I2 .
2KBr+ K2S208+H2S04 = 2K2S04-f Br2 4 H2SO4 .
KCI + K2S2O8+ H2SO4 (1.5— 2N) = No action.
2Agl -|- 2AgBr -f 2AgCI -j- 3Zn = 6Ag 4. ZnL 4 ZnBr2 4 ZnCla .
428
APPENDIX.
CORRECTIONS IN ATOMIC WEIGHTS, MADE BY THE
INTERNATIONAL COMMITTEE, JANUARY, 1905.
Journal of the American Chemical Society, Vol. XXVII.
PAGE.
72— Antimony 120.2
100— Bismuth ' 208.5
342 — Bromine 79.96
193— Cerium 140.25
193 — Columbium 94.
288 — Fluorine 19.
196— Indium 115.
350— Iodine 126.97
132— Iridium 193.
197— Lanthanum 138.9
29— Lead 206.9
214— Magnesium 24.36
131— Palladium 106.5
93— Platinum . . 194.8
222— Potassium 39.15
234— Rubidium 85.5
45— Silver 107.93
304 — Sulphur 32.06
198— Tantalum 183.
137— Tellurium 127.6
199— Thallium 204.1
199— Thorium 232.5
200— Titanium 48.1
201 — Uranium 238.5
135 — Vanadium 51.2
202— Ytterbium 173.
202 — Zirconium . 90.6
Catalogue of Scientific Publications
and Importations of the D. Van Nostr and Company,
23 Murray Street and 2j Warren Street, New York.
ABC CODE. (See Clausen-Thue.)
ABBOT, H. L., Gen'l. The Defence of the Seacoast of
the United States. Lectures delivered before the U. S. Naval
War College. 8vo, red cloth $2.00
ABBOTT, A. V. The Electrical Transmission of Energy.
A Manual for the Design of Electrical Circuits. New edition, re-
vised, and entirely rewritten. Fully illustrated. 8vo, cloth. $4.50
ADAM, P. Practical Bookbinding. With illustrations
and figures. Translated from the German by Thomas E. Maw.
8vo, cloth, illustrated net, $2 . 50
ADAMS, J. W. Sewers and Drains for Populous Dis-
tricts. Embracing Rules and Formulas for the dimensions and
construction of works of Sanitary Engineers. 8vo, cloth. ... $2 . 50
ADDYMAN, F. T. Practical X-Ray Work. Part I,
Historical. Part II, Apparatus and its Management. Part III,
Practical X-Ray Work. Illustrated with twelve plates from
photographs. 8vo, cloth, illustrated net, $4.00
A 1 CODE. (See Clausen-Thue.)
AIKMAN, C. M., Prof. Manures and the Principles of
Manuring. 8vo, cloth . , $2 . 50
ALEXANDER, J. H. Universal Dictionary of Weights
and Measures, Ancient and Modern, reduced to the Standards of
the United States of America. New Edition, enlarged. 8vo,
cloth $3 . 50
ALEXANDER, S. A. Broke Down: What Should I Do?
A Ready Reference and Key to Locomotive Engineers and Fire-
men, Round-house Machinists, Conductors, Train Hands and
Inspectors. With 5 folding plates. 12mo, cloth $1 .50
2 D. VAN NOSTRAND COMPANY'S
ANDERSON, G. L., A.M. (Captain of II. S. Artillery).
Handbook for the use of Electricians in the operation and care
of Electrical Machinery and Apparatus of the Lmted States Sea-
coast Defenses. Prepared under the direction of the Lieutenant-
General Commanding the Army. With tables, diagrams and
illustrations. 8vo, cloth, illustrated $3 .00
ANDERSON, J. W. Prospector's Handbook. A Guide
for the Prospector and Traveller in search of Metal-bearing or other
Valuable Minerals. Eighth Edition, revised. 8vo, cloth. . . .$1 .50
ANDERSON, W. On the Conversion of Heat into Work.
A Practical Handbook on Heat-engines. Third Edition. Illus-
trated. 12mo, cloth $2.25
ANDES, L. Vegetable Fats and Oils: Their Practical
Preparation, Purification and Employment for Various Purposes.
Their Properties, Adulteration and Examination. A Handbook
for Oil Manufacturers and Refiners, Candle, Soap and Lubricating-
oil Manufacturers, and the Oil and Fat Industry in general. Trans-
lated from the German. With 94 illus. Svo, cloth. . . .net, $4.00
Animal Fats and Oils. Their Practical Produc-
tion, Purification and Uses for a great variety of purposes; their
Properties, Falsification and Examination. A Handbook for
Manufacturers of Oil and Fat Products, Soap and Candle Makers,
Agriculturists, Tanners, etc. Translated by Charles Salter.
With 62 illustrations. Svo. cloth net, $4 . 00
- Drying Oils, Boiled Oil, and Solid and Liquid Driers.
A practical work for manufacturers of Oils, Varnishes, Printing
Inks, Oilcloth and Linoleum, Oil-cakes, Paints, etc. Svo, cloth,
illustrated net, $5.00
Iron Corrosion, Anti-fouling and Anti-corrosive
Paints. Translated from the German by Charles Salter. Illus-
trated with engravings and half-tone cuts. Svo, cloth, .net, $4.00
Oil Colors, and Printers' Ink. A Practical Hand
book treating of Linseed-oil, Boiled Oil, Paints, Artists' Colors,
Lampblack, and Printers' Inks (black and colored). Translated
from the German by Arthur Morris and Herbert Robson. With
56 figures and diagrams. Svo, cloth, 212 pages net, $2. 50
ARNOLD, E. Armature Windings of Direct-Current
Dynamos. Extension and Application of a General Winding
Rule. Translated from the original German by Francis B.
DeGress, M.E. With numerous illustrations. Svo, cloth. . . $2 . 00
SCIENTIFIC PUBLICATIONS. 3
ARNOLD, R., Dr. Ammonia and Ammonium Com-
pounds. A Practical Manual for Manufacturers, Chemists, Gas
Engineers and Drysalters. Second Edition. 1 2mo , cloth. . . . $2 . 00
Art of Dyeing Wool, Silk and Cotton. Translated from
the French of M. Hellott, M. Macquer and M. Le Pileur D'Apligny.
First published in English in 1789. 8vo, cloth, illustrated, net, $2 . 00
ASHE, S. W., and KEILEY, J. D. Electric Railways,
Theoretically and Practically Treated. With 200 diagrams and
engravings. Ready Sept., 1904.
ATKINSON, A. A., Prof. (Ohio University). Electrical
and Magnetic Calculations, for the use of Electrical Engineers and
Artisans, Teachers, Students and all others interested in the
Theory and Application of Electricity and Magnetism. Second
Edition, revised. 8vo, cloth, illustrated net, $1 . 50
ATKINSON, P. The Elements of Electric Lighting.
including Electric Generation, Measurement, Storage and Dis
tribution. Tenth Edition, fully revised and new matter added-
Illustrated. 12mo, cloth $1 . 50
The Elements of Dynamic Electricity and Mag-
netism. Fourth Edition. 120 illustrations. 12mo, cloth. . $2.00
Power Transmitte " by Electricity and its Appli-
cation by the Electric Motor, including Electric Railway Con-
struction. Fourth Edition, fully revised, new matter added.
12mo, cloth, illustrated $2.00
AUCHINCLOSS, W. S. Link and Valve Motions Sim-
plified. Illustrated with 29 woodcuts and 20 lithographic plates,
together with a Travel Scale, and numerous useful tables. Four-
teenth Edition, revised. 8vo, cloth $2 . 00
AXON, W. E. A. The Mechanic's Friend. A Collection
of Receipts and Practical Suggestions relating to Aquaria, Bronz-
ing, Cements, Drawing, Dyes, Electricity, Gilding, Glass-working,
Glues, Horology, Lacquers, Locomotives, Magnetism, Metal-work-
ing, Modelling, Photography, Pyrotechny, Railways, Solders,
Steam-engine, Telegraphy, Taxidermy, Varnishes, Waterproofing,
and Miscellaneous Tools, Instruments, Machines and Processes
connected with the Chemical and Mechanic Arts. With numerous
diagrams and woodcuts. 12mo, cloth, illustrated $1 .50
AYRTON, H. The Electrical Arc. With numerous fig-
ures, diagrams and plates. 8vo, cloth, illustrated $5.00
4 D. VAN NOSTRAND COMPANY'S
AYRTON, W. E., M.I.C.E. Practical Electricity. A Labo-
ratory and Lecture Course for the first-year students of Electrical
Engineering, based on the International Definitions of the Electri-
cal Units. Vol. I, Current, Pressure, Resistance, Energy, Power,
and Cells. Completely rewritten and containing many figures
• and diagrams. 12mo, cloth ............................ $2 . 00
BACON, F. W. A Treatise on the Richards Steam-engine
Indicator, with directions for its use. By Charles T. Porter.
Revised, with notes and large additions as developed by American
practice; with an appendix containing useful formulae and rules
for engineers. Illustrated. Fourth Edition. 12mo, cloth. . $1.00
BADT, F. B. New Dynamo Tender's Handbook. With
140 illustrations. 16mo, cloth ......................... $1 . 00
- Bell-hangers' Handbook. With 97 illustrations.
Second Edition. 16mo, cloth .......................... $1 . 00
— Incandescent Wiring Handbook. With 35 illustra-
tions and 5 tables. Fifth Edition. IGmo, cloth ........... $1.00
— Electric Transmission Handbook. With 22 illus-
trations and 27 tables. 16mo, cloth ..................... $1 .00
BAKER, Arthur L., Prof. (Univ. of Rochester). Quater-
nions ............................................. In Press.
BAKER, M. N. Potable Water and Methods of Detect-
ing Impurities. 16mo, cloth. (Van Nostrand's Science Series) .60
BALCH, G. T., Col. Methods of Teaching Patriotism
in the Public Schools. 8vo, cloth ....................... $1 .00
BALE, M. P. Pumps and Pumping. A Handbook for
Pump Users. 12mo, cloth ............................. $1 . 50
BARBA, J. The Use of Steel for Constructive Purposes.
Method of Working, Applying and Testing Plates and Bars.
With a Preface by A. L. Holley, C.E. 12mo, cloth ...... $1 .50
BARKER, A. H. Graphic Methods of Engine Design.
Including a Graphical Treatment of the Balancing of Engines.
12mo, cloth .......................................... $1 . 50
BARNARD, F. A. P. Report on Machinery and Pro-
cesses of the Industrial Arts and Apparatus of the Exact Sciences
at the Paris Universal Exposition, 1867. 152 illustrations and
8 folding plates. 8vo, cloth ........................... $5. 00
SCIENTIFIC PUBLICATIONS. 5
BARNARD, J. H. The Naval Militiaman's Guide. Full
leather, pocket size $1 . 25
BARRUS, G. H. Boiler Tests: Embracing the Results
of one hundred and thirty-seven evaporative tests, made on
seventy-one boilers, conducted by the author. 8vo, cloth. $3.00
Engine Tests: Embracing the Results of over one
hundred feed-water tests and other investigations of various
kinds of steam-engines, conducted by the author. With numer-
ous figures, tables, and diagrams. 8vo, cloth, illustrated. . $4.00
The two books sent prepaid for $6 . 00
BARWISE, S., M.D. (London). The Purification of
Sewage. Being a brief account of the Scientific Principles of
Sewage Purification and their Practical Application. 12mo,
cloth, illustrated $2.00
The Bacterial Purification of Sewage. Being a
Practical Account of the Various Modern Biological Methods of
Purifying Sewage. With plates and diagrams. 8vo, cloth,
illustrated net, $2.50
BAUER'S Marine Engines and Boilers In Press.
BAUMEISTER, R. The Cleaning and Sewerage of Cit'es.
Adapted from the German with permission of the author. By
J. M. Goodell, C.E. Second Edition, revised and corrected, to-
gether with an additional appendix. 8 vo, cloth, illustrated. $2.00
BEAUMONT, R. Color in Woven Design. With 32
colored plates and numerous original illustrations. Large,
12mo $7.50
BEECH, F. Dyeing of Cotton Fabrics. A Practical
Handbook for the Dyer and Student. Containing numerous
recipes for the production of Cotton Fabrics of all kinds, of a great
range of colors, thus making it of great service in the dye-house,
while to the student it is of value in that the scientific principles
which underlie the operations of dyeing are clearly laid down.
With 44 illustrations of Bleaching and Dyeing Machinery. 8vo,
cloth, illustrated net, $3 . 00
— —Dyeing of Woolen Fabrics. With diagrams and
figures. 8vo, cloth, illustrated net , $3 . 50
BECKWITH, A. Pottery. Observations on the Materials
and Manufacture of Terra-cotta, Stoneware, Firebrick, Porce-
lain, Earthenware, Brick, Majolica, and Encaustic Tiles. Second
Edition. 8vo, paper 60
6 D. VAN NOSTRAND COMPANY'S
BEGTRUP, J., M.E. The Slide Valve and its Functions
With Special Reference to Modern Practice in the United States.
With numerous diagrams and figures. 8vo, cloth $2.00
BERNTHSEN, A. A Text-book of Organic Chemistry.
Translated by George M'Gowan, Ph.D. Fourth English Edition,
revised and extended by author and translator. Illustrated.
1 2mo, cloth $2 . 50
BERSCH, J., Dr. Manufacture of Mineral and Lake
Pigments. Containing directions for the manufacture of al!
artificial artists' and painters' colors, enamel colors, soot and
metallic pigments. A text-book for Manufacturers, Merchants,
Artists and Painters. Translated from the second revised edition
by Arthur C. Wright, M.A. 8vo, cloth, illustrated net, $5.00
BERTIN, L. E. Marine Boilers: Their Construction and
Working, dealing more especially with Tubulous Boilers. Trans-
lated by Leslie S. Robertson, Assoc. M. Inst. C.E., M. I. Mech. E.,.
M.I.N.A., containing upward of 250 illustrations. Preface by
Sir William White, K.C.B., F.R.S., Director of Naval Construc-
tion to the Admiralty, and Assistant Controller of the Navy.
New and Revised Edition. 8vo, cloth, illustrated $7.50
BIGGS, C. H. W. First Principles of Electricity and
Magnetism. A book for beginners in practical work, containing
a good deal of useful information not usually to be found in
similar books. With numerous tables and 343 diagrams and
figures. 12mo, cloth, illustrated $2.00
BINNS, C. F. Ceramic Technology. Being Some Aspects
of Technical Science as applied to Pottery Manufacture. 8vo,
doth net, $5.00
Manual of Practical Potting. Compiled by Experts.
Third Edition, revised and enlarged. 8vo, cloth net, $7 . 50
BIRCHMORE, W. H., Dr. How to Use a Gas Analysis.
12mo, cloth, illustrated In Press.
BLAKE, W. P. Report upon the Precious Metals. Being
Statistical Notices of the Principal Gold and Silver producing
regions of the world, represented at the Paris Universal Exposi-
tion. 8vo, cloth $2 . 00
BLAKESLEY, T. H. Alternating Cur ents of Electricity.
For the use of Students and Engineers. Third Edition, enlarged.
12mo, cloth $1 . 50
SCIENTIFIC PUBLICATIONS. 7
BLYTH, A. W., M.R.C.S., F.C.S. Foods: Their Com-
position and Analysis. A Manual for the use of Analytical
Chemists, with an Introductory Essay on the History of Adultera-
tions. With numerous tables and illustrations. Fifth Edition,
thoroughly revised, enlarged and rewritten. 8vo, cloth $7.50
Poisons: Their Effects and Detection. A Manual
for the use of Analytical Chemists and Experts, with an Intro-
ductory Essay on the Growth of Modem Toxicology. Third
Edition, revised and enlarged. 8vo, cloth $7.50
BODMER, G. R. Hydraulic Motors and Turbines. For
the use of Engineers, Manufacturers and Students. Third Edi-
tion, revised and enlarged. With 192 illustrations. 12mo,
cloth $5.00
BOILEAU, J. T. A New and Complete Set of Traverse
Tables, showing the Difference of Latitude and Departure of
every minute of the Quadrant and to five places of decimals.
8vo, cloth $5.00
BONNEY, G. E. The Electro-platers' Handbook. A
Manual for Amateurs and Young Students of Electro-metallurgy.
60 illustrations. 12mo, cloth $1 . 20
BOTTONE, S. R. Electrical Instrument Making for
Amateurs. A Practical Handbook. With 48 Illustrations.
Fifth Edition, revised. 12mo, cloth 50
Electric Bells, and All About Them. A Practical
Book for Practical Men. With more than 100 Illustrations.
Fourth Edition, revised and enlarged. 12mo, cloth 50
Electro-motors: How Made and How Used.' A
Handbook for Amateurs and Practical Men. Second Edition.
12mo, cloth 75
BOURRY, E. Treatise on Ceramic Industries. A Complete
Manual for Pottery, Tile and Brick Works. Translated from
the French by Wilton P. Rix. With 323 figures and illustrations.
8vo, cloth, illustrated net, $S . 50
BOW, R. H. A Treatise on Bracing. With its applica-
tion to Bridges and other Structures of Wood or Iron. 156 illus-
trations. 8vo, cloth $1 . 50
BOWIE, AUG. J., Jr., M.E. A Practical Treatise on
Hydraulic Mining in California. With Description of the Use
and Construction of Ditches, Flumes, Wrought-iron Pipes and
Dams; Flow of Water on Heavy Grades, and its Applicability,
under High Pressure, to Mining. Ninth Edition. Small quarto,
cloth. Illustrated $5.00
8 D. VAN NOSTRAND COMPANY'S
BOWKER, Wm. R. Dynamo, Motor and" Switchboard
Circuits. For Electrical Engineers. A practical book, dealing
with the subject of Direct, Alternating, and Polyphase Currents.
With over 100 diagrams and engravings. 8vo, cloth. . net, $2.25
BOWSER, E. A., Prof. An Elementary Treatise on
Analytic Geometry. Embracing Plane Geometry, and an Intro-
duction to Geometry of three Dimensions. Twenty-first Edition.
12mo, cloth $1 .75
— An Elementary Treatise on the Differential and
Integral Calculus. With numerous examples. Twentieth Edition.
12mo, cloth $2.25
— An Elementary Treatise on Analytic Mechanics. With
numerous examples. Sixteenth Edition. 12mo, cloth $3.00
An Elementary Treatise on Hydro-mechanics. With
numerous examples. Fifth Edition. 12mo, cloth $2.50
— A Treatise on Roofs and Bridges. With Numerous
Exercises, especially adapted for school use. 12mo, cloth.
Illustrated net, $2.25
BRASSEY'S Naval Annual for 1904. Edited by T. A.
Brassey. With numerous full-page diagrams, half-tone illustra-
tions and tables. Eighteenth year of publication. £vo, cloth,
illustrated net, $0.00
BRAUN, E. The Baker's Book: A Practical Handbook
of the Baking Industry in all Countries. Profusely illustrated
with diagrams, engravings, and full-page colored plates. Trans-
lated into English and edited by Emil Braun. Vol. I., 8vo,
cloth, illustrated, 308 pages $2 . 50
Vol. II. 363 pages, illustrated $2.50
British Standard Sections. Issued by the Engineering
Standards Committee, Supported by The Institution of Civil
Engineers, The Institution of Mechanical Engineers, The Institu-
tion of Naval Architects, The Iron and Steel Institute, and The
Institution of Electrical Engineers. Comprising 9 platee of
diagrams, with letter-press and tables. Oblong pamphlet,
8JX15 $1 00
BROWN, WM. N. The Art of Enamelling on Metal. With
figures and illustrations. 12mo, cloth, illustrated net, $1 . 00
Handbook on Japanning and Enamelling, for Cycles,
Bedsteads, Tinware, etc. 12mo, cloth, illustrated net, $1 .50
SCIENTIFIC PUBLICATIONS. 9
BROWN, WM. N. House Decorating and Painting. With
Numerous illustrations. 12mo, cloth net, $1 .50
History of Decorative Art. With Designs and Illus-
trations. 12mo, cloth net, $1 . 25
Principle and Practice of Dipping, Burnishing, Lac-
quering and Bronzing Brass. Ware. 12mo, cloth net, $1 . 00
Workshop Wrinkles for Decorators, Painters, Paper-
Hangers and Others. Svo, cloth net, $1 . 00
BUCHANAN, E. E. Tables of Squares. Containing the
square of every foot, inch, and sixteenth of an inch, between one
sixteenth of an inch and fifty feet. For Engineers and Calcu-
lators. 16mo, oblong, cloth $1 . 00
BURGH, N. P. Modern Marine Engineering, Applied to
Paddle and Screw Propulsion. Consisting of 36 colored plates,
259 practical woodcut illustrations and 403 pages of descriptive
matter. The whole being an exposition of the present practice
of James Watt & Co., J. & G. Rennie, R. Napier & Sons, and
other celebrated firms. Thick quarto, half morocco $10 . 00
BURT, W. A. Key to the Solar Compass, and Surveyor's
Companion. Comprising all the rules necessary for use in the
field; also description of the Linear Surveys and Public Land
System of the United States, Notes on the Barometer, Sugges-
tions for an Outfit for a Survey of Four Months, etc. Seventh
Edition. Pocket size, full leather , $2 . 50
CAMPIN, F. On the Construction of Iron Roofs. A
Theoretical and Practical Treatise, with woodcuts and plates of
roofs recently executed. Svo, cloth $2 . 00
CARTER, E. T. Motive Power and Gearing for Elec-
trical Machinery. A treatise on the Theory and Practice of the
Mechanical Equipment of Power Stations for Electrical Supply
and for Electric Traction. Svo, cloth, illustrated $5 . 00
CATHCART, WM. L., Prof. Machine Design. Part I.
Fastenings. Svo, cloth, illustrated net, $3 .00
and CHAFFEE, J. I. Course of Graphic Statics Applied
ta Mechanical Engineering In Press.
CHAMBER'S MATHEMATICAL TABLES, consisting of
Logarithms of Numbers 1 to 108,000, Trigonometrical, Nautical
and other Tables. New Edition. Svo, cloth $1 . 75
10 D. VAN NO3TRAND COMPANY'S
CHARPENTIER, P. Timber. A Comprehensive Study
of Wood in all its Aspects, Commercial and Botanical. Show-
ing the Different Applications and Uses of Timber in Various
Trades, etc. Translated into English. 8vo, cloth, illus. . . net, $6 . 00
CHAUVENET, W., Prof. New Method of Correcting
Lunar Distances, and Improved Method of Finding the Error
and Rate of a Chronometer, by Equal Altitudes. 8vo, cloth. $2 . 00
CHILD, C. T. The How and Why of Electricity. A
Book of Information for non-technical readers, treating of the
Properties of Electricity, and how it is generated, handled, con-
trolled, measured and set to work. Also explaining the opera-
tion of Electrical Apparatus. Svo, cloth, illustrated $1 . 00
CHRISTIE, W. W. Chimney Design and Theory. A
Book for Engineers and Architects, with numerous half-tone
illustrations and plates of famous chimneys. Second Edition,
revised. Svo, cloth $3 . 00
- Furnace Draft : its Production by Mechanical Methods.
A Handy Reference Book, with figures and tables. 16mo, limp
cloth, illustrated 50
- Boiler- waters, Scale, Corrosion, Foaming In Press.
CLAPPERTON, G. Practical Paper-making. A Manual
for Paper-makers and Owners and Managers of Paper Mills, to
which is appended useful tables, calculations, data, etc., with
illustrations reproduced from micro-photographs. Ilium, cloth,
illustrated $2 . 50
CLARK, D. K., C.E. A Manual of Rules, Tables and
Data for Mechanical Engineers. Baaed on the most recent inves-
tigations. Illustrated with numerous diagrams. 1012 pages, bvo,
cloth. Sixth Edition $5.00
Half morocco $7 . 50
Fuel: its Combustion and Economy; consisting of
abridgments of Treatise on the Combustion of Coal. By C. W.
Williams, and the Economy of Fuel, by T. S. Prideaux. With
extensive additions in recent practice in the Combustion and
Economy of Fuel, Coal, Coke, Wood, Peat, Petroleum, etc.
Fourth Edition. 12mo, cloth $1 . 50
The Mechanical Engineer's Pocket-book of Tables,
Formulae, Rules and Data. A Handy Book of Reference for
Daily Use in Engineering Practice. 16mo, morocco. Fifth
Edition, carefully revised throughout $3 . 00
SCIENTIFIC PUBLICATIONS. 11
CLARK, D. K., C.E. Tramways: Their Construction and
Working. Embracing a comprehensive history of the system, with
accounts of the various modes of traction, a description of the
varieties of rolling stock, and ample details of Cost and Working
Expenses. Second Edition, rewritten and greatly enlarged, with
upwards of 400 illustrations. Thick 8vo, cloth $9 . 00
CLARK, J. M. New System of Laying Out Railway Turn-
outs instantly, by inspection from tables. 12mo, cloth. . . $1 .00
CLAUSEN-THUE, W. The A B C Universal Commercial
Electric Telegraphic Code; specially adapted for the use of
Financiers, Merchants, Ship-owners, Brokers, Agents, etc. Fourth
Edition. 8vo, cloth $5 . 00
Fifth Edition of same $7 . 00
• The A 1 Universal Commercial Electric Telegraphic
Code. Over 1240 pages and nearly 90,000 variations. 8vo,
cloth $7 . 50
CLEEMANN, T. M. The Railroad Engineer's Practice.
Being a Short but Complete Description of the Duties of the
Young Engineer in Preliminary and Location Surveys and in
Construction. Fourth Edition, revised and enlarged. Illustrated.
12mo, cloth '. $1 . 50
CLEVENGER, S. R. A Treatise on the Method of Gov-
ernment Surveying as prescribed by the U. S. Congress and Com-
missioner of the General Land Office, with complete Mathemati-
cal, Astronomical, and Practical Instructions for the use of the
United States Surveyors in the field. 16mo, morocco $2.50
CLOUTH, F. Rubber, Gutta-Percha, and Balata. First
English Translation with Additions and Emendations by the
Author. With numerous figures, tables, diagrams, and folding
plates. 8vo, cloth, illustrated net, $5.00
COFFIN, J. H. C., Prof. Navigation and Nautical Astron-
omy. Prepared for the use of the U. S. Naval Academy. New
Edition. Revised by Commander Charles Belknap. 52 woodcut
illustrations. 12mo, cloth net, $3 . 50
COLE, R. S., M.A. A Treatise on Photographic Optics.
Being an account of the Principles of Optics, so far as they apply
to photography. 12mo, cloth, 103 illus. and folding plates. .$2 . 50
COLLINS, J. E. The Private Book of Useful Alloys and
Memoranda for Goldsmiths, Jewelers, etc. 18mo, cloth 50
12 D. VAN NOSTRAND COMPANY'S
COREY. Water-supply Engineering. Fully illustrated.
8vo, floth In Press,
COOPER, W. R., M.A. Primary Batteries: Their Con-
struction and Use. With numerous figures and diagrams. 8vo,
cloth, illustrated net, $4 . 00
CORNWALL, H. B., Prof. Manual of Blow-pipe Analysis,
Qualitative and Quantitative. With a Complete System of
Determinative Mineralogy. 8vo, cloth, with many illustra-
tions $2 . 50
COWELL, W. B. Pure Air, Ozone and Water. A Prac-
tical Treatise of their Utilization and Value in Oil, Grease, Soap.
Paint, Glue and other Industries. With tables and figures.
121110, cloth, illustrated net, $2.00
CRAIG, B. F. Weights and Measures. An Account of
the Decimal System, with Tables of Conversion for Commercial
and Scientific Uses. Square 32mo, limp cloth 50
CROCKER, F. B., Prof. Electric Lighting. A Practical
Exposition of the Art. For use of Engineers, Students, and
others interested in the Installation or Operation of Electrical
Plants. Vol. I. The Generating Plant. New Edition, thoroughly
rerised and rewritten. 8vo, cloth, illustrated In Press.
Vol. II. Distributing Systems and Lamps. Second Edition. 8vo,
cloth, illustrated $3 . 00
-and WHEELER, S. S. The Practical Management
of Dynamos and Motors. Fifth Edition (Sixteenth Thousand), re-
vised and enlarged. With a special chapter by H. A. Foster.
12mo, cloth, illustrated $1 . 00
CROSSKEY, L. R. Elementary Perspective: Arranged to
meet the requirements of Architects and Draughtsmen, and of
Art Students preparing for the elementary examination of the
Science and Art Department, South Kensington. With numer-
ous full-page plates and diagrams. 8vo, cloth, illustrated . . $1 .00
and THAW, J. Advanced Perspective, involving the
Drawing of Objects when placed in Oblique Positions, Shadows
and Reflections. Arranged to meet the requirements of Archi-
tects, Draughtsmen, and Students preparing for the Perspective
Examination of the Education Department. With numerous full-
page plates and diagrams. 8vo. cloth, illustrated $1 . 50
DAVIES, E. H. Machinery for Metalliferous Mines.
A Practical Treatise for Mining Engineers, Metallurgists and
Managers of Mines. With upwards of 400 illustrations. Second
Edition, rewritten and enlarged. 8vo, cloth net, $8 . 00
SCIENTIFIC PUBLICATIONS. 13
DAVIES, D. C. A Treatise on Metalliferous Minerals and
Mining. Sixth Edition, thoroughly revised and much enlarged by his
son. 8vo, cloth net, $5.00
Mining Machinery In Press.
DAVISON, G. C., Lieut. Water-tube Boilers In Press.
DAY, C. The Indicator and its Diagrams. With Chap-
ters on Engine and Boiler Testing; including a Table of Piston
Constants compiled by W. H. Fowler. 12mo, cloth. 125 illus-
trations ' $2.00
DE LA COUX, H. The Industrial Uses of Water. With
numerous tables, figures, and diagrams. Translated from the
French and revised by Arthur Morris. 8vo, cloth net, $4 . 50
DENNY, G. A. Deep-level Mines of the Rand, and their
future development, considered from the commercial point of
view. With folding plates, diagrams, and tables. 4to, cloth,
illustrated net, $10.00
DERR, W. L. Block Signal Operation. A Practical
Manual. Pocket Size. Oblong, cloth. Second Edition. .. .$1 .50
DIBDIN, W. J. Public Lighting by Gas and Electricity.
With tables, diagrams, engravings and full-page plates. 8vo,
cloth, illustrated net, $8.00
Purification of Sewage and Water. With tables,
engravings, and folding plates Third Edition, revised and
enlarged. 8vo, cloth, illus. and numerous folding plates .... $6 . 50
DIETERICH, K. Analysis of Resins, Balsams, and Gum
Resins: their Chemistry and Pharmacognosis. For the use of
the Scientific and Technical Research Chemist. With a Bibliog-
raphy. Translated from the German, by Chas. Salter. 8vo.
cloth net, $3 .00
DIXON, D. B. The Machinist's and Steam Engineer's
Practical Calculator. A Compilation of Useful Rules and Prob-
lems arithmetically solved, together with General Information
applicable to Shop-tools, Mill-gearing, Pulleys and Shafts, Steam-
boilers and Engines. Embracing valuable Tables and Instruc-
tion in Screw-cutting, Valve and Link Motion, etc. Third Edition.
16mo, full morocco, pocket form $1 . 25
DOBLE, W. A. Power Plant Construction on the Pacific
Coast . In Press.
14 D. VAN NOSTRAND COMPANY'S
DODD, GEO. Dictionary of Manufactures, Mining, Ma-
chinery, and the Industrial Arts. 12mo, cloth $1 .50
DORR, B. F. The Surveyor's Guide and Pocket Table-
book. Fifth Edition, thoroughly revised and greatly extended.
With a second appendix up to date. 16mo, morocco flaps. . $2 . 00
DRAPER, C. H. An Elementary Text-book of Light,
Heat and Sound, with Numerous Examples. Fourth Edition.
12mo, cloth, illustrated $1 . 00
-Heat and the Principles of Thermo-dynamics. With
many illustrations and numerical examples. 12mo, cloth. . . $1 . 50
DYSON, S. S. Practical Testing of Raw Materials. A
Concise Handbook for Manufacturers, Merchants, and Users of
Chemicals, Oils, Fuels, Gas Residuals and By-products, and
Paper-making Materials, with Chapters on Water Analysis and
the Testing of Trade Effluents. 8vo, cloth, illustrations, 177
pages net, $5 . 00
EDDY, H. T., Prof. Researches in Graphical Statics.
Embracing New Constructions in Graphical Statics, a New General
Method in Graphical Statics, and the Theory of Internal Stress
in Graphical Statics. 8vo, cloth $1 . 50
Maximum Stresses under Concentrated Loads.
Treated graphically. Illustrated 8vo, cloth $1 . 50
EISSLER, M. The Metallurgy of Gold. A Practical Treatise
on the Metallurgical Treatment of Gold-bearing Ores, including
the Processes of Concentration and Chlorination, and the Assay-
ing, Melting and Refining of Gold. Fifth Edition, revised and
greatly enlarged. Over 300 illustrations and numerous folding
plates. 8vo, cloth $7 . 50
- The Hydro-Metallurgy of Copper. Being ?n Account
of processes adopted in the Hydro-metallurgical Treatment of
Cupriferous Ores, including the Manufacture of Copper Vitriol.
With chapters on the sources of supply of Copper and the Roasting
of Copper Ores. With numerous diagrams and figures. Svo,
cloth, illustrated net, $4 . 50
- The Metallurgy of Silver. A Practical Treatise on the
Amalgamation, Roasting and Lixiviation of Silver Ores, including
the Assaying, Melting and Refining of Silver Bullion. 124
illustrations. Second Edition, enlarged. Svo, cloth $4.00
SCIENTIFIC PUBLICATIONS. 15
EISSLER, M. The Metallurgy of Argentiferous Lead. A
Practical Treatise on the Smelting of Silver-Lead Ores and the Re-
fining of Lead Bullion. Including Reports on Various Smelting
Establishments and Descriptions of Modern Smelting Furnaces and
Plants in Europe and America. With 183 illustrations. 8vo.
cloth $5 . 00
Cyanide Process for the Extraction of Gold and its
Practical Application on the Witwatersrand Gold Fields in South
Africa. Third Edition, revised and enlarged. Illustrations and
folding plates. 8vo, cloth $3 . 00
A Handbook on Modern Explosives. Being a Prac-
tical Treatise on the Manufacture and Use of Dynamite, Gun-
cotton, Nitroglycerine, and other Explosive Compounds, in-
cludjng the manufacture of Collodion-cotton, with chapters on
Explosives in Practical Application. Second Edition, enlarged
with 150 illustrations. 12mo, cloth $5 . 00
ELIOT, C. W., and STORER, F. H. A Compendious
Manual of Qualitative Chemical Analysis. Revised with the co-
operation of the authors, by Prof. William R. Nichols. Illus-
trated. Twentieth Edition, newly revised by Prof. W. B. Lindsay.
12mo, cloth net, $1 . 25
ELLIOT, G. H., Maj. European Light-house Systems.
Being a Report of a Tour of Inspection made in 1873. 51 en-
gravings and 21 woodcuts. 8vo, cloth $5.00
ERFURT, J. Dyeing of Paper Pulp. A Practical Treatise
for the use of paper-makers, paper-stainers, students and others,
With illustrations and 157 patterns of paper dyed in the pulp,
with formulas for each. Translated into English and edited,
with additions, by Julius Hiibner, F.C.S. 8vo, cloth, illus-
trated .' ... net, $7.50
EVERETT, J. D. Elementary Text-book of Physics.
Illustrated. Seventh Edition. 12mo, cloth $1 . 50
EWING, A. J., Prof. The Magnetic Induction in Iron
and other metals. Third Edition, revised. 159 illustrations.
8vo, cloth $4 . 00
FAIRIE, J., F.G.S. Notes on Lead Ores : Their Distribu-
tion and Properties. 12mo, cloth $1 . 00
Notes on Pottery Clays: The Distribution, Properties,
Uses and Analysis of Ball Clays, China Clays and China Stone.
With tables and formulae. 12mo, cloth . $1 . 50
16 D. VAN NOSTRAND COMPANY'S
FANNING, J. T. A Practical Treatise on Hydraulic and
Water-supply Engineering. Relating to the Hydrology, Hydro-
dynamics and Practical Construction of Water-works in North
America. 180 illus. 8vo, cloth. Fifteenth Edition, revised, en-
larged, and new tables and illustrations added. 650 pp $5.00
FAY, I. W. The Coal-tar Colors : Their Origin and Chem-
istry. 8vo, cloth, illustrated In Press.
FISH, J. C. L. Lettering of Working Drawings. Thir-
teen plates, with descriptive text. Oblong, 9X12$, boards. $1.00
FISHER, H. K. C., and DARBY, W. C. Students' Guide
to Submarine Cable Testing. Third (new and enlarged) Edi-
tion. 8vo, cloth, illustrated $3 . 50
FISHER, W. C. The Potentiometer and its Adjuncts.
8vo, cloth $2. 25
FISKE, B. A., Lieut., U.S.N. Electricity in Theory and
Practice ; or, The Elements of Electrical Engineering. Eighth
Edition. 8vo, cloth $2.50
FLEISCHMANN, W. The Book of the Dairy. A Manual
of the Science and Practice of Dairy Work. Translated from
the German, by C. M. Aikman and R. Patrick Wright. 8vo,
cloth $4.00
FLEMING, J. A., Prof. The Alternate-current Trans-
former in Theory and Practice. Vol. I., The Induction of Electric
Currents; 611 pages. New Edition, illustrated. 8vo, cloth, $5.00
Vol. II., The Utilization of Induced Currents. Illustrated. 8vo,
cloth $5 . 00
— Centenary of the Electrical Current, 1799-1899.
8vo, paper, illustrated 50
— Electric Lamps and Electric Lighting. Being a
course of four lectures delivered at the Royal Institution, April-
May, 1894. 8vo, cloth, fully illustrated $3 .00
-Electrical Laboratory Notes and Forms, Elementary
and Advanced. 4to, cloth, illustrated $5.00
— A Handbook for the Electrical Laboratory and Test-
ing Room. 2 volumes. 8vo, cloth each $5.00 j
FLEURY, H. The Calculus Without Limits or Infinitesi-
mals. Translated by C. O. Mailloux In Press.
SCIENTIFIC PUBLICATIONS. 17
FOLEY, N., and PRAY, THOS., Jr. The Mechanical
Engineer^' Reference Book for Machine and Boiler Construction,
in two parts. Part 1— General Engineering Data. Part 2 —
Boiler Construction. With 51 plates and numerous illustrations,
specially drawn for this work. Folio, half morocco $25.00
FORNEY, M. N. Catechism of the Locomotive. Second
Edition, revised and enlarged. 46th thousand. 8vo, cloth. .$3.50
FOSTER, J. G., Gen., U.S.A. Submarine Blasting in
Boston Harbor, Massachusetts. Removal of Tower and Corwin
Rocks. Illustrated with 7 plates. 4to, cloth $3 . 50
FOSTER, H. A. Electrical Engineers' Pocket-book. With
the Collaboration of Eminent Specialists. A handbook of useful
data for Electricians and Electrical Engineers. With innumer-
able tables, diagrams, and figures. Third Edition, revised
Pocket size, full leather, 1000 pp $5.00
FOSTER, J. Treatise on the Evaporation of Saccharine,
Chemical and other Liquids by the Multiple System in Vacuum
and Open Air. Third Edition. Diagrams and large plates.
8vo, cloth $7 . 50
FOX, WM., and THOMAS, C. W., M.E. A Practical
Course in Mechanical Drawing. Second Edition, revised. 12mo,
cloth, with plates $1 . 25
FRANCIS, J. B., C.E. Lowell Hydraulic Experiments.
Being a selection from experiments on Hydraulic Motors on
the Flow of Water over Weirs, in Open Canals of uniform rect-
angular section, and through submerged Orifices and diverging
Tubes. Made at Lowell, Mass. Fourth Edition, revised and
enlarged, with many new experiments, and illustrated with 23
copper-plate engravings. 4to, cloth $15 . 00
FULLER, G. W. Report on the Investigations into the
Purification of the Ohio River Water at Louisville, Kentucky,
made to the President and Directors of the Louisville Water
Company. Published under agreement with the Directors.
3 full-page plates. 4to, cloth net, $10 . 00
FURNELL, J. Students' Manual of Paints, Colors, Oils
and Varnishes. 8vo, cloth, illustrated net, $1 .00
GARCKE, E., and FELLS, J. M. Factory Accounts:
their principles and practice. A handbook for accountants and
manufacturers, with appendices on the nomenclature of machine
details, the rating of factories, fire and boiler insurance, the
factory and workshop acts, etc., including also a large number
of specimen rulings. Fifth Edition, revised and extended. 8vo,
cloth, illustrated $3 .00
18 D. VAN NOSTRAND COMPANY'S
GEIPEL, WM., and KILGOUR, M. H. A Pocket-took
of Electrical Engineering Formulae. Illns. 18mo., mor. . .$3.00
GERBER, N. Chemical and Physical Analysis of Milk,
Condensed Milk, and Infants' Milk-food. 8vo, cloth $1 . 25
GERHARD, WM. P. Sanitary Engineering. i2mo,
cloth $1 . 25
GESCHWIND, L. Manufacture of Alum and Sulphates,
and other Salts of Alumina and Iron; their uses and applications
as mordants in dyeing and calico printing, and their other appli-
cations in the Arts, Manufactures, Sanitary Engineering, Agri-
culture and Horticulture. Translated from the French by
Charles Salter. With tables, figures and diagrams. Svo, cloth,
illustrated net, $5 . 00
GIBBS, W. E. Lighting by Acetylene, Generators, Burners
and Electric Furnaces. With 66 illustrations. Second Edition,
revised. 12mo, cloth $1 . 50
•
GILLMORE, Q. A., Gen. Treatise on Limes, Hydraulic
Cements and Mortars. Papers on Practical Engineering, United
States Engineer Department, No. 9, containing Reports of nu-
merous Experiments conducted in New York City during the
years 1853 to 1861, inclusive. With numerous illustrations.
Svo, cloth $4 00
-Practical Treatise on the Construction of Roads,
Streets and Pavements. Tenth Edition. With 70 illustrations
12mo, «loth $2 . 00
— Report on Strength of the Building Stones in the
United States, etc. Svo, illustrated, cloth $1 .00
GOLDING, H. A. The Theta-Phi Diagram. Practically
Applied to Steam, Gas, Oil and Air Engines. 12mo, cloth,
illustrated net, $1 . 25
GOODEVE, T. M. A Text-book on the Steam-engine.
With a Supplement on Gas-engines. Twelfth Edition, enlarged.
143 illustrations. 12mo, cloth $2.00
GORE, G., F.R.S. The Art of Electrolytic Separation of
Metals, etc. (Theoretical and Practical.) Illustrated Svo,
cloth $3 50
GOULD, E. S. The Arithmetic of the Steam-engine.
8vo, cloth $1 .00
SCIENTIFIC PUBLICATIONS. 19
GOULD, E. S. Practical Hydrostatics and Hydrostatic For-
mulas. With numerous figures and diagrams (Van Nostrand's
Science Series, No. 117.) 16mo, cloth, illustrated, 114 pp... .50
GRAY, J., B.Sc. Electrical Influence Machines: Their
Historical Development, and Modern Forms, with instructions
for making them. With numerous figures and diagrams. Second
Edition, revised and enlarged. 12mo, cloth, illus., 296 pp. . . .$2.00
GRIFFITHS, A. B., Ph.D. A Treati e on Manures, or
the Philosophy of Manuring. A Practical Handbook for the
Agriculturist, Manufacturer, and Student. 12mo, cloth. .. $3.00
Dental Metallurgy. A Manual for Students and
Dentists. 8vo, cloth, illustrated, 20S pp net, $3.50
GROSS, E. Hops, in their Botanical, Agricultural and
Technical Aspect, and as an article of Commerce. Translated
from the German by Charles Salter. With tables, diagrams,
and illustrations. Svo, cloth, illustrated net, $4 . 50
GROVER, F. Practical Treatise on Modern Gas and
Oil Engines. Svo, cloth, illustrated . $2 . 00
GRUNER, A. Power-loom Weaving and Yarn Number-
ing, according to various systems, with conversion tables. An
auxiliary and text-book for pupils of weaving schools, as well
as for self-instruction, and for general use by those engaged in
the weaving industry. Illustrated with colored diagrams. Svo,
cloth net, $3 .00
GURDEN, R. L. Traverse Tables: Computed to Four-
B'ace Decimals for every single minute of angle up to 100 of
istance. For the use of Surveyors and Engineers. New Edition.
Folio, half morocco $7 . 50
GUY, A. E. Experiments on the Flexure of Beams,
resulting in the Discovery of New Laws of Failure by Buckling.
Reprinted from the "American Machinist." With diagrams and
folding plates. Svo, cloth, illustrated, 122 pages net, $1 .25
GUY, A. F. Electric Light and Power: Giving the Result
of Practical Experience in Central-station Work. Svo, cloth,
illustrated $2 . 50
HAEDER, H., C.E. A Handbook on the Steam-engine.
With especial reference to small and medium-sized engines. For
the use of Engine-makers, Mechanical Draughtsmen, Engineer-
ing Students and Users of Steam Power. Translated from the
German, with considerable additions and alterations, by H. H.
P. Powles. Third English Edition, revised. Svo, cloth, illus-
trated, 458 pages $3 . 00
20 D. VAN NOSTRANf) COMPANY'S
HALL, W. S., Prof. Elements of the Differential and
Integral Calculus. Fourth Edition, revised. 8vo, cloth, illus-
trated net, $2.25
- Descriptive Geometry, with Numerous Problems and
Practical Applications. Comprising an 8vo volume of 76 pages
of text and a 4to atlas of 31 plates. 2 vols., cloth. . . . net, $3.50
Postage, 0.32
HALSEY, F. A. Slide-valve Gears. An Explanation of
the Action and Construction of Plain and Cut-off Slide Valves.
Illustrated. Seventh Edition. 12mo, cloth $1 . 50
-The Use of the Slide Rule. With illustrations and
folding plates. Second Edition. 16mo, boards. (Fan Nos-
trand's Science Series, No. 1 14.). 50
- The Locomotive Link Motion, with Diagrams and
Tables. 8vo, cloth, illustrated $1 .00
- Worm and Spiral Gearing. Revised and Enlarged
Edition. 16mo, cloth (Van Nostrand's Science Series, No. 116.)
Illustrated 50
- The Metric Fallacy, and " The Metric Failure in
the Textile Industry," by Samuel S. Dale. 8vo, cloth, illus-
trated $1.00
HAMILTON, W. G. Useful Information for Railway
Men. Tenth Edition, revised and enlarged. 562 pages, pocket
form. Morocco, gilt $2 . 00
HAMMER, W. J. Radium, and Other Radio-active Sub-
stances; Polonium, Actinium and Thorium. With a considera-
tion of Phosphorescent and Fluorescent Substances, the Proper-
ties and Applications of Selenium, and the treatment of disease
by the Ultra-Violet Light. Second Edition. With diagrams,
engravings and photographic plates. 8vo, cloth, illustrated, 72 p.,
$1.00
HANCOCK, H. Text-book of Mechanics and Hydro-
statics, with over 500 diagrams. 8vo, cloth $1 .75
HARRISON, W. B. The Mechanics' Tool-book. With
Practical Rules and Suggestions for use of Machinists, Iron-
workers and others. Wit n 44 engravings. 12mo, cloth. . . .$1 .50
HART, J. W. External Plumbing Work. A Treatise on
Lead Work for Roofs. With numerous figures and diagrams.
8vo cloth, illustrated net, $3 .06
SCIENTIFIC PUBLICATIONS. 21
HART, J. W. Hints to Plumbers on Joint Wiping, Pipe
Bending, and Lead Burning. Containing 184 figures and dia-
grams. 8vo, cloth, illustrated ...................... net, $3 . 00
-- Principles of Hot-water Supply. With numerous
illustrations. 8vo, cloth ........................... net, $3 . 00
- Sanitary Plumbing and Drainage. With numerous
diagrams and figures. 8yo, cloth, illustrated ......... net, $3.00
HASKINS, C. H. The Galvanometer and its Uses. A
Manual for Electricians and Students. Fourth Edition. 12mo,
cloth ................................................ $1 . 50
HAUFF, W. A. American Multiplier: Multiplications and
Divisions of the largest numbers rapidly performed. With index
giving the results instantly of all numbers to 1000x1000 = 1.000,-
000; also tables of circumferences and areas of circles. Cloth,
$5.00
HAUSBRAND, E. Drying by Means of Air and Steam.
With explanations, formulas, and tables, for use in practice.
Translated from the German by A. C. Wright, M.A. 12mo,
cloth, illustrated ...................................... $2.00
-- Evaporating, Condensing and Cooling Apparatus:
Explanations, Formulae, and Tables for Use in Practice. Trans-
lated from the Second Revised German Edition by A. C. Wright,
M.A. With numerous figures, tables and diagrams. 8vo, cloth,
illustrated, 400 pages ............................... net, $5.00
HAUSNER, A. Manufacture of Preserved Foods and
Sweetmeats. A Handbook of all the Processes for the Preserva-
tion of Flesh, Fruit, and Vegetables, and for the Preparation of
Dried Fruit, Dried Vegetables, Marmalades, Fruit-syrups, and
Fermented Beverages, and of all kinds of Candies, Candied Fruit,
Sweetmeats, Rocks, Drops, Dragees, Pralines, etc. Translated
from the Third Enlarged German Edition, by Arthur Morris and
Herbert Robson, B.Sc. 8vo, cloth, iUustrated ........ net, $3 .00
HAWKE, W. H. The Premier Cipher Telegraphic Code,
containing 100,000 Words and Phrases. The most complete and
most useful general code yet published. 4to, cloth ....... $5.00
-- 100,000 Words Supplement to the Premier Code.
All the words are selected from the official vocabulary. Oblong
quarto, cloth ......................................... $5 . 00
HAWKINS, C. C., and WALLIS, F. The Dynamo: its
Theory, Design, and Manufacture. 190 illustrations. 12mo,
cloth ............................................... $3.00
22 D. VAN NOSTRAND COMPANY'S
HAY, A. Principles of Alternate-current Working. I2mo,
cloth, illustrated $2.00
HEAP, D. P., Major, U.S.A. Electric 1 Appliances of
the Present Day. Report of the Paris Electrical Exposition of
1S&1. 250 illustrations. 8vo, cloth $2.00
HEAVISIDE, 0. Electromagnetic Theory. 8vo, cloth,
two volumes each, $5 . 00
HECK, R. C. H., Prof. (Lehigh University). The Steam-
Engine. A text-book for Engineering Colleges In Press.
HEERMANN, P. Dyers' Materials. An Introduction to
the Examination, Valuation, and Application of the most impor-
tant substances used in Dyeing, Printing, Bleaching and Finish-
ing. Translated by Arthur C. Wright, M.A. 12mo, cloth, illus-
trated net, $2.50
HENRICI, 0. Skeleton Structures, Applied to the Build-
ing of Steel and Iron Bridges. 8vo, cloth, illustrated $1 . 50
HERMANN, F. Painting on Glass and Porcelain and
Enamel Painting. On the basis of Personal Practical Experience
of the Condition of the Art up to date. Translated by Charles
Salter. Second greatly enlarged edition. 8vo, cloth, illustrated,
net, $3 . 50
HERRMANN, G. The Graphical Statics of Mechanism.
A Guide for the Use of Machinists, Architects and Engineers; and
also a Text-book for Technical Schools. Translated and anno-
tated by A. P. Smith, M.E. Fourth Edition. 12mo, cloth, 7
folding plates $2.00
HERZFELD, J., Dr. The Technical Testing of Yarns and
Textile Fabrics, with reference to official specifications. Trans-
lated by Chas. Salter. With 69 illustrations. 8vd, cloth net, $3 . 50
HEWSON, W. Principles and Practice of Embanking
Lands from River Floods, as applied to the Levees of the Missis-
sippi. 8vo, cloth $2 . 00
HILL, J. W. The Purification of Public Water Supplies.
Illustrated with valuable tables, diagrams, and cuts. 8vo,
cloth, 304 pages $3 . 00
Interpretation of Water Analysis. i2mo, cloth.
In Press.
SCIENTIFIC PUBLICATIONS. 23
HOBBS, W. R. P. The Arithmetic of Electrical Measure-
ments, with numerous examples, fully worked. Revised by
Richard Wormell, M.A. Ninth Edition. 12mo, cloth 50
HOFF, WM. B., Com., U.S.N. The Avoidance of Collisions
at Sea. 18mo, morocco 75
HOLLEY, A. L. Railway Practice. American and Euro-
pean Railway Practice in the Economical Generation of Steam,
including the Materials and Construction of Coal-burning Boilers,
Combustion, the Variable Blast, Vaporization, Circulation, Super-
heating, Supplying and Heating Feed Water, etc., and the
Adaptation of Wood and Coke-burning Engines to Coal-burning;
and in Permanent Way, including Road-bed, Sleepers, Rails,
Joint Fastenings, Street Railways, etc. With 77 lithographed
plates. Folio, cloth $12.00
HOLMES, A. B. The Electric Light Popularly Explained.
Fifth Edition. Illustrated. 12mo, paper 40
HOPKINS, N. M. Model Engines and Small Boats. New
Methods of Engine and Boiler Making, with a chapter on Ele-
mentary Ship Design and Construction. 12mo, cloth $1 . 25
HOUSTON, E. J., and KENNELLY, A. E. Algebra Made
Easy. Being a clear explanation of the Mathematical Formulae
found in Prof. Thompson's "Dynamo-electric Machinery and
Polyphase Electric Currents. " With figures and examples. Svo,
cloth, illustrated 75
The Interpretation of Mathematical Formulas. With
figures and examples. Svo, cloth, illustrated $1 .25
HOWARD, C. R. Earthwork Mensuration on the Basis
of the Prismoidal Formula^ Containing Simple and Labor-saving
Methods of obtaining Prismoidal Contents directly from End
Areas. Illustrated by Examples and accompanied by Plain
Rules for Practical Use. Illustrated. Svo, cloth $1 .50
HOWORTH, J. Art of Repairing and Riveting Glass,
China and Earthenware. Second Edition. Svo, pamphlet, illus-
trated net, $0. 50
HUBBARD, E. The Utilization of Wood-waste. A Com-
plete Account of the Most Advantageous Methods of Working Up
Wood Refuse, especially Sawdust, Exhausted Dye Woods and
Tan as Fuel, as a Source of Chemical Products for Artificial Wood
Compositions, Explosives, Manures, and many other Technical
Purposes. Translated from the German of the second revised
and enlarged edition. Svo, cloth, illustrated, 192 pages. . net, $2.50
24 D. VAN NOSTRAND COMPANY'S
HUMBER, W., C.E A Handy Book for the Calculation
of Strains in Girders, and Similar Structures, and their Strength;
consisting of Formulae and Corresponding Diagrams, with numer-
ous details for practical application, etc. Fourth Edition. 12mo,
cloth $2 . 50
HURST, G. H., F.C.S. Color. A Handbook of the Theory
of Color. A practical work for the Artist, Art Student, Painter,
Dyer and Calico Printer, and Others. Illustrated with 10 colored
plates and 72 illustrations. Svo, cloth net, $2 . 50
- Dictionary of Chemicals and Raw Products Used
in the Manufacture of Paints, Colors, Varnishes and Allied Prep-
arations. Svo, cloth net, $3 . 00
- Lubricating Oils, Fats and Greases : Their Origin,
Preparation, Properties, Uses and Analysis. 313 pages, with
65 illustrations. Svo, cloth net, $3 . 00
— Soaps. A Practical Manual of the Manufacture of
Domestic, Toilet and other Soaps. Illustrated with 66 engrav-
ings. Svo, cloth net, $5 . 00
- Textile Soaps and Oils : A Handbook on the Prepara-
tion, Properties, and Analysis of the Soaps and Oils Used in
Textile Manufacturing, Dyeing and Printing. With tables and
illustrations. Svo, cloth net, $2.50
HUTCHINSON, W. B. Patents and How to Make Money
out of Them. 12mo, cloth $1 . 25
BUTTON, W. S. Steam-boiler Construction. A Practical
Handbook for Engineers, Boiler-makers and Steam-users. Con-
taining a large collection of rules and data relating to recent
practice in the design, construction and working of all kinds
of stationary, locomotive and marine steam-boilers. With up-
wards of 540 illustrations. Fourth Edition, carefully revised and
much enlarged. Svo, cloth $6 . 00
Practical Engineer's Handbook, comprising a Treatise
on Modern Engines and Boilers, Marine, Locomotive and Station-
ary. Fourth Edition, carefully revised, with additions. With
upwards of 570 illustrations. Svo, cloth $7 . 00
The Works' Manager's Han book of Modern Rules,
Tables and Data for Civil and Mechanical Engineers, Millwrights
and Boiler-makers, etc., etc. With upwards of 150 illustrations.
Fifth Edition, carefully revised, with additions. Svo, cloth. . . $6 . 00
INGLE
clot
, H. Manual of Agricultural Chemistry. Svo,
h, illustrated, 388 pages net, $3 . 00
SCIENTIFIC PUBLICATIONS. 25
INNES, C. H. Problems in Machine Design. For the
use of Students, Draughtsmen and others. Second Edition, 12mo,
cloth net, $2.00
— Centrifugal Pumps, Turbines and Water Motors. In-
cluding the Theory and Practice of Hydraulics. Third and enlarged
edition. 12mo, cloth net, $2 . 00
ISHERWOOD, B. F. Engineering Precedents for Steam
Machinery. Arranged in the most practical and useful manner
for Engineers. With illustrations. Two volumes in one. 8vo,
cloth $2.50
JAMIESON, A., C.E. A Text-book n Steam and Steam-
engines Specially arranged for the use of Science and Art, City
and Guilds of London Institute, and other Engineering Students.
Thirteenth Edition. Illustrated. 12mo, cloth $3 . 00
Elementary Manual on Steam and the Steam-engine.
Specially arranged for the use of First-year Science and Art, City
and Guilds of London Institute, and other Elementary Engineer-
ing Students. Third Edition. 12mo, cloth. $1 . 50
JANNETTAZ, E. A Guide to the Determination of Rocks :
being an Introduction to Lithology. Translated from the French
by G. W. "Plympton, Professor of Physical Science at Brooklyn
Polytechnic Institute. 12mo, cloth $1 . 50
JEHL, F., Mem. A.I.E.E. The Manufacture of Carbons
for Electric Lighting and Other Purposes. A Practical Handbook,
giving a complete description of the art of making carbons, electros,
etc. The various gas generators and furnaces used in carbonizing,
with a plan for a model factory. Illustrated with numerous dia-
grams, tables and folding plates. 8vo, cloth, illustrated. ... $4 . 00
JENNISON, F. H. Th3 Manufacture of Lake Pigments
from Artificial Colors. A useful handbook for color manufac-
turers, dyers, olor chemists, paint manufacturers, drysalters,
wallpaper-makers, enamel and surface-paper makers. With 15
plates illustrating the various methods and errors that arise in
the different processes of production. 8vo, cloth net, $3.00
JOHNSON, W. McA. "The Metallurgy of Nickel." In Press.
JOHNSTON, J. F. W., Prof., and CAMERON, Sir Chas.
Elements of Agricultural Chemistry and Geology. Seventeenth
Edition. 12mo, cloth $2 . 60
JONES, H. C. Outlines of Electrochemistry. With
tables and diagrams. 4to, cloth, illustrated $1 . 50
26 D. VAN NOSTRAND COMPANY'S
JONES, M. W. The Testing and Valuation of Raw Mate-
rials used in Paint and Color Manufacture. 12mo, cloth, . net, $2 . 00
JOYNSON, F. H. The Metals Used in Construction.
Iron, Steel, Bessemer Metal, etc. Illustrated. 12mo, cloth. .. .75
- Designing and Construction of Machine Gearing.
Illustrated. 8vo, cloth $2 . 00
JUPTNER, H. F. V. Siderology: The Science of Iron.
(The Constitution of Iron Alloys and Iron.) Translated from
the German. 8vo, cloth, 345 pages, illustrated net, $5.00
KANSAS CITY BRIDGE, THE. With an Account of
the Regimen of the Missouri River and a Description of the
Methods used for Founding in that River, by O. Chanute, Chief
Engineer, and George Morison, Assistant Engineer. Illustrated
with 5 lithographic views and 12 plates of plans. 4to, cloth. $6 . 00
KAPP, G., C.E. Electric Transmission of Energy and
its Transformation, Subdivision and Distribution. A practical
handbook. Fourth Edition, revised. 12mo, cloth $3.50
Dynamos, Motors, Alternators and Rotary Con-
verters. Translated from the third German edition, by Harold
H. Simmons, A.M.I.E.E. With numerous diagrams and figures.
8vo, cloth, 507 pages $4 . 00
KEIM, A. W. Prevention of Dampness in Buildings.
With Remarks on the Causes, Nature and Effects of Saline Efflo-
rescences and Dry Rot. For Architects, Builders, Overseers,
Plasterers, Painters and House Owners. Translated from the
second revised, German edition. With colored plates and dia-
grams. 8vo, cloth, illustrated, 115 pages net, $2 . 00
KELSEY, W. R. Continuous-current Dynamos and
Motors, and their Control: being a series of articles reprinted
from The Practical Engineer, anal completed by W. R. Kelsey.
With many figures and diagrams. 8vo, cloth, illustrated. . .$2.50
KEMP, J. F., A.B., E.M. A Handbook of Rocks. For Use
without the microscope. With a glossary of the names of rocks
and of other lithological terms. 8vo, cloth, illustrated $1 . 50
Third Edition, revised.
KEMPE, H. R. The Electrical Engineer's Pocket-book
of Modern Rules, Formulae, Tables and Data. Illustrated.
32mo, morocco, gilt $1 . 75
SCIENTIFIC PUBLICATIONS. 27
KENNEDY, R. Modern Engines an* Power Generators.
Vol. I. A Practical Work on Prime Movers and the Transmission
of Power: Steam, Electric, Water, and Hot-air. With tables,
figures, and full-page engravings. 4to, cloth, illustrated. . $3.50
KENNEDY, R. Electrical Installations of Electric Light,
Power, Traction, and Industrial Electrical Machinery. With nu-
merous diagrams and engravings.
Vol. I. The Electrical Circuit, Measurement, Ele-
ments of Motors, Dynamos, Electrolysis. 8vo, cloth, illus. . $3.50
Vol. II. Instruments, Transformers, Installation Wir-
ing, Switches and Switchboards. 8vo, cloth, illustrated. ... $3.50
Vol. III. Production of Electrical Energy, Prime
Movers, Generators and Motors. 8vo, cloth, illustrated. . . $3.50
Vol. IV. Mechanical Gearing; Complete Electric In-
stallations; Electrolytic, Mining and Heating Apparatus; Electric
Traction; Special Applications of Electric Motors. 8vo, cloth,
illustrated $3.50
• Vol. V. Apparatus and Machinery used in Telegraphs,
Telephones, Signals, Wireless Telegraph, X-Rays, and Medical
Science. 8vo, cloth, illustrated $3 . 50
KENNELLY, A. E. Theoretical Elements of Electro-
dynamic Machinery. 8vo, cloth $1 . 50
KILGOUR, M. H., SWAN, H., and BIGGS, C. H. W. Elec-
trical Distribution: its Theory and Practice. 174 illustrations.
12mo, cloth $4.00
KING, W. H. Lessons and Practical Notes on Steam,
the Steam-engine, Propellers, etc., for Young Marine Engineers,
Students and others. Revised by Chief Engineer J. W. King,
United States Navy. 19th Edition, enlarged. 8vo, cloth. . . . $2 . 00
KINGDON, J. A. Applied Magnetism. An Introduction
to the Design of Electromagnetic Apparatus. 8vo, cloth. . $3.00
KIRKALDY, W. G. Illustrations of David Kirkaldy's
System of Mechanical Testing, as Originated and Carried on by
him during a Quarter of a Century. Comprising a Large Selec-
tion of Tabulated Results, showing the Strength and other Proper-
ties of Materials used in Construction, with Explanatory Text
and Historical Sketch. Numerous engravings and 25 lithographed
plates. 4to, cloth $20 . 00
28 D. VAN NOSTRAND COMPANY'S
KIRKBRIDE, J. Engraving for Illustration: Historical
and Practical Notes, with illustrations and 2 plates by ink
photo process. 8vo, cloth, illustrated, 72 pages net, $1 .50
KIRKWOOD, J. P. Report on the Filtration of River
Waters for the Supply of Cities, as practised in Europe, made
to the Board of Water Commissioners of the City of St. Louis.
Illustrated by 30 double-page engravings. 4to, cloth .... $7 . 50
KLEIN, J. F. Design of a High-speed Steam-engine.
With notes, diagrams, formulas and tables. Second Edition,
rcriscd and enlarged. 8vo, cloth, illustrated, 257 pages, .net, $5.00
KNIGHT, A. M., Lieut.-Com. U.S.N. Modem Seaman-
ship. Illustrated with 136 full-page plates and diagrams. 8vo,
cloth, illustrated. Second Edition, revised net, $6.00
Half morocco $7 . 50
KNOTT, C. G., and MACKAY, J. S. Practical Mathematics.
With numerous examples, figures and diagrams. New Edition.
8vo, cloth, illustrated $2 . 00
KOLLER, T. The Utilization of Waste Products. A
Treatise on the Rational Utilization, Recovery and Treatment
of Waste Products of all kinds. Translated from the German
second revised edition. With numerous diagrams. 8vo, cloth,
illustrated, 280 pages net, $3 . 50
- Cosmetics. A Handbook of the Manufacture, Em-
ployment and Testing of all Cosmetic Materials and Cosmetic
Specialties. Translated from the German by Chas. Salter. 8vo,
cloth net, $2.50
KRAUCH, C., Dr. Testing of Chemical Reagents for
Purity. Authorized translation of the Third Edition, by J. A.
Williamson and L. W. Dupre. With additions and emendations
by the author. 8vo, cloth net, $4 . 50
LAMBERT, T. Lead, and its Compounds, With tables,
diagrams and folding plates. 8vo, cloth, 22S pages. . . .net, $3.50
Bone Products and Manures. An Account of
the most recent improvements in the manufacture of Fat, Glue,
Animal Charcoal, Size, Gelatine and Manures. With plans and
diagrams. 8vo, cloth, illustrated net, $3 . 00
LAMBORN, L. L. Cotton-seed Products. With numerous
diagrams and engravings In Press.
SCIENTIFIC PUBLICATIONS. 29
LAMPRECHT, R. Recovery Work after Pit Fires. A
description of the principal methods pursued, especially in fiery
mines, and of the various appliances employed, such as respira-
tory and rescue apparatus, dams, etc. With folding plates and
diagrams. Translated from the German by Charles Salter. 8vo,
cloth, illustrated , net, $4.00
LARRABEE, C. S. Cipher and Secret Letter and Tele-
§raphic Code, with Hog's Improvements. The most perfect
ecret Code ever invented or discovered. Impossible to read
without the key. 18mo, cloth 60
LASSAR-COHN, Dr. An Introduction to Modern Scien-
tific Chemistry, in the form of popular lectures suited to University
Extension Students and general readers. Translated from the
author's corrected proofs for the second German edition, by
M. M. Pattison Muir, M.A. 12mo, cloth, illustrated $2.00
LEASK, A. R. Breakdowns at Sea and How to Repair
Them. With 89 illustrations. Second Edition. 8vo, cloth. $2 . 00
Triple and Quadruple Expansion Engines and Boilers
and their Management. With 59 illustrations. Third Edition,
revised. 12mo, cloth $2 . 00
Refrigerating Machinery: Its Principles and Man-
agement. With 64 illustrations. 12mo, cloth $2.00
LECKY, S. T. S. "Wrinkles" in Practical Navigation.
With 130 illustrations. 8vo, cloth. Fourteenth Edition, revised
and enlarged $9 . 00
LEFEVRE, L. Architectural Pottery: Bricks, Tiles, Pipes,
Enameled Terra-Cottas, Ordinary and Incrusted Quarries, Stone-
ware Mosaics, Faiences and Architectural Stoneware. With
tables, plates and 950 cuts and illustrations. With a preface by
M. J.-C. Formige. Translated from the French, by K. H. Bird,
M.A., and W. Moore Binns. 4to, cloth, illustrated net, $7. 50
LEHNER, S. Ink Manufacture : including Writing, Copy-
ing, Lithographic, Marking, Stamping ,and Laundry Inks. Trans-
lated from the fifth German edition, by Arthur Morris and
Herbert Robson, B.Sc. 8vo, cloth, illustrated net, $2.50
LEMSTROM, Dr. Electricity in Horticulture. Illus-
trated In Press.
LEVY, C. L. Electric-light Primer. A simple and com-
prehensive digest of all the most important facts connected with
the running of the dynamo, and electric lights, with precautions
for safety. For the use of persons whose duty it is to look after
the plant. 8vo, paper 60
30 D. VAN NOSTRAND COMPANY'S
LIVACHE, ACH. (Ingenieur CivU des Mines). The Man-
ufacture of Varnishes, Oil Crushing, Refining and Boiling, and
Kindred Industries. Describing the Manufacture and Chemical
and Physical Properties of Spirit Varnishes and Oil Varnishes;
Raw Materials; Resins; Solvents and Coloring Principles; Dry-
ing Oils, their Extraction, Properties and Applications; Oil
Refining and Boiling; The Manufacture, Employment and Test-
ing of Various Varnishes. Translated from tne French, by John
Geddes Mclntosh. Greatly extended and adapted to English
practice. With numerous original recipes by the translator.
Illustrated with cuts and diagrams. 8vo, cloth net, $5.00
LIVERMORE, V. P., and WILLIAMS, J. How to Become
a Competent Motorman. Being a Practical Treatise on the
Proper Method of Operating a Street Railway Motor Car; also
giving details how to overcome certain defects. 16mo, cloth,
illustrated, 132 pages $1 . 00
LOBBEN, P., M.E. Machinists' and Draftsmen's Hand-
book, containing Tables, Rules, and Formulas, with numerous
examples, explaining the principles of mathematics and mechanics,
as applied to the mechanical trades. Intended as a reference book
for all interested in Mechanical work. Illustrated with many
cuts and diagrams. 8vo, cloth $2 . 50
LOCKE, A. G. and C. G. A Practical Treatise on the
Manufacture of Sulphuric Acid. With 77 constructive plates,
drawn to scale measurements, and other illustrations. Royal
8vo, cloth $10 . 00
LOCKERT, L. Petroleum Motor-cars. i2mo, cloth, $1.50
LOCKWOOD, T. D. Electricity, Magnetism, and Electro-
telegraphy. A Practical Guide for Students, Operators, and
Inspectors. 8vo, cloth. Third Edition $2 . 50
— Electrical Measurement and the Galvanometer: its
Construction and Uses. Second Edition. 32 illustrations. 12mo,
cloth $ 1 . 50
LODGE, 0. J. Elementary Mechanics, including Hydro-
statics and Pneumatics. Revised Edition. 12mo, cloth ... $1 . 50
- Signalling Across Space, Without Wires : being a
description of the work of Hertz and his successors. With numer-
ous diagrams and half-tone cuts, and additional remarks con-
cerning the application to Telegraphy and later developments.
Third Edition. 8vo, cloth, illustrated net, $2 . 00
SCIENTIFIC PUBLICATIONS. 31
LORD, R. T. Decorative and Fancy Fabrics. A Valuable
Book with designs and illustrations for manufacturers and de-
signers of Carpets, Damask, Dress and all Textile Fabrics. 8vo,
cloth, illustrated net, $3 . 50
LORING, A. E. A Handbook of the Electro-magnetic
Telegraph. 16mo, cloth, boards. New and enlarged edition. . .50
LUCE, S. B., Com. Text-book of Seamanship. The
Equipping and Handling Of Vessels under Sail or Steam. For
the use of the U. S. Naval Academy. Revised and enlarged
edition, by Lieut. Wm. S. Benson. 8vo, cloth, illustrated. . $10 . 00
LUCKE, C. E., Prof. Columbia Univ. Gas Engine Design.
LUNGE, G., Ph.D. Coal-tar and Ammonia: being the
third and enlarged edition of "A Treatise on the Distillation of
• Coal-tar and Ammoniacal Liquor," with numerous tables, figures
and diagrams. Thick 8vo, cloth, illustrated net, $15.00
A Theoretical and Practical Treatise on the Man-
ufacture of Sulphuric Acid and Alkali with the Collateral Branches.
— Vol. I. Sulphuric Acid. In two parts, not sold separately.
Second Edition, revised and enlarged. 342illus. Svo, cloth. . $15.00
Vol. II. Salt Cake, Hydrochloric Acid and Leblanc
Soda. Second Edition, revised and enlarged. Svo, cloth. . . $15.00
Vol. III. Ammonia Soda, and various other processes
of Alkali-making, and the preparation of Alkalis, Chlorine and
Chlorates, by Electrolysis. Svo, cloth. New Edition, 1896 . . $1 5 . 00
-and HURTER, F. The Alkali Maker's Handbook.
Tables and Analytical Methods for Manufacturers of Sulphuric
Acid, Nitric Acid, Soda, Potash and Ammonia. Second Edition.
12mo, cloth $3.00
LUPTON, A., PARR, G. D. A., and PERKIN, H. Elec-
tricity as Applied to Mining. With tables, diagrams and folding
plates. Svo, cloth, illustrated, 280 pages net, $3 . 50
LUQUER, L. M., Ph.D. Minerals in Rock Sections.
The Practical Method of Identifying Minerals in Rock Sections
with the Microscope. Especially arranged for Students in Techni-
cal and Scientific Schools. Svo, cloth, illustrated net, $1 . 50
MACCORD, C. W., Prof. A Practical Treatise on the
Slide-valve by Eccentrics, examining by methods the action of
the Eccentric upon the Slide-valve, and explaining the practical
processes of laying out the movements, adapting the Valve for
its various duties in the Steam-engine. Second Edition. Illus-
trated. 4to, cloth $2 . 50
32 D. VAN NOSTRAND COMPANY'S
MACKIE, JOHN. How to Make a Woolen Mill Pay.
8vo, cloth net, $2.00
MACKROW, C. The Naval Architect's and Ship-builder's
Pocket-book of Formulae, Rules, and Tables; and Engineers' and
Surveyors' Handy Book of Reference. Eighth Edition, revised
and enlarged. 16mo, limp leather, illustrated $5.00
MAGUIRE, E., Capt, U.S.A. The Attack and Defence
of Coast Fortifications. With maps and numerous illustrations,
8vo, cloth $2 . 50
MAGUIRE, WM. R. Domestic Sanitary Drainage and
Plumbing Lectures on Practical Sanitation. 332 illustrations.
8vo $4 . 00
MAILLOUX, C. 0. Electro-traction Machinery. 8vo,
cloth, illustrated In Press.
MARKS, E. C. R. Mechanical Engineering Materials:
Their Properties and Treatment in Construction. 12mo, cloth,
illustrated 60.
— Notes on the Construction of Cranes and Lifting
Machinery. With numerous diagrams and figures. New and
enlarged edition. 12mo, cloth. . .« net, $1 . 50
- Notes on the Construction and Working of Pumps.
With figures, diagrams and engravings. 12mo, cloth, illus-
trated net, $1 . 50
MARKS, G. C. Hydraulic Power Engineering. A Prac-
tical Manual on the Concentration and Transmission of Power
by Hydraulic Machinery. With over 200 diagrams and tables
8vo, cloth, illustrated $3 . 50
MAVER, W. American Telegraphy: Systems, Apparatus,
Operation. 450 illustrations. 8vo, cloth $5.00
MAYER, A. M., Prof. Lecture Notes on Physics. 8vo,
cloth $2.00
McCULLOCH, R. S., Prof. Elementary Treatise on the
Mechanical Theory of Heat, and its application to Air and Steam-
engines. 8vo, cloth $3 . 50
McINTOSH, J. G. Technology of Sugar. A Practical Treatise
on the Manufacture of Sugar from the Sugar-cane and Sugar-
beet. With diagrams and tables. 8vo, cloth, illustrated . net, $4 . 50
SCIENTIFIC PUBLICATIONS. 33
McNEILL, B. McNeill's Code. Arranged to meet the
requirements of Mining, Metallurgical and Civil Engineers, Direc-
tors of Mining, Smelting and other Companies, Bankers, Stock
and Share Brokers, Solicitors, Accountants, Financiers and
General Merchants. Safety and Secrecy. 8vo, cloth. ... $6 . 00
McPHERSON, J. A., A. M. Inst. C. E. Waterworks
Distribution. A practical guide to the laying out of systems of
distributing mains for the supply of water to cities and towns.
With tables, folding plates and numerous full-page diagrams.
8vo, cloth, illustrated $2. 50
MERRITT, WM. H. Field Testing for Gold and Silver.
A Practical Manual for Prospectors and Miners. With numerous
half-tone cuts, figures and tables. 16mo, limp leather, illus-
trated $1 . 50
METAL TURNING. By a Foreman Pattern-maker. Illus-
trated with 81 engravings. 12mo, cloth $1 . 50
MICHELL, S. Mine Drainage: being a Complete Prac-
tical Treatise on Direct-acting Underground Steam Pumping
Machinery. Containing many folding plates, diagrams and
tables. Second Edition, rewritten and enlarged. Thick 8vo,
cloth, illustrated $10.00
MIERZINSKI, S., Dr. Waterproofing of Fabrics. Trans-
lated from the German by Arthur Morris and Herbert Robson.
With diagrams and ligures. 8vo, cloth, illustrated... net, $2.50
MILLER, E. H. Quantitative Analysis for Mining Engi-
neers. 8vo, cloth net, $1 . 50
MINIFIE, W. Mechanical Drawing. A Text-book of
Geometrical Drawing for the use of Mechanics and Schools, in
w.hich the Definitions and Rules of Geometry are familiarly ex-
plained; the Practical Problems are arranged from the most
simple to the more complex, and in their description technicalities
are avoided as much as possible. With illustrations for drawing
Plans, Sections, and Elevations of Railways and Machinery; an
Introduction to Isometrical Drawing, and an Essay on Linear
Perspective and Shadows. Illustrated with over 200 diagrams
engraved on steel. Tenth Thousand, revised. With an Appen-
dix on the Theory and Application of Colors. 8vo, cloth. . $4.00
Geometrical Drawing. Abridged from the octavo
edition, for the use of schools. Illustrated with 48 steel plates.
Ninth Edition. 12mo, cloth $2.00
34 D VAN NOSTRAND COMPANY'S
MODERN METEOROLOGY. A Series of Six Lectures,
delivered under the auspices of the Meteorological Society in
1870. Illustrated. 12mo, cloth $1 . 50
MOORE, E. C. S. New Tables for the Complete Solu-
tion of Ganguillet and Kutter's Formula for the flow of liquids in
open channels, pipes, sewers and conduits. In two parts. Part I,
arranged for 1080 inclinations from 1 over 1 to 1 over 21,120 for
fifteen different values of (n). Part II, for use with all other
values of (n). With large folding diagram. 8vo, cloth, illus-
trated net, $5.00
MOREING, C. A., and NEAL, T. New General and Mining
Telegraph Code. 676 pages, alphabetically arranged. For the
use of mining companies, mining engineers, stock brokers, financial
agents, and trust and finance companies. Eighth Edition. 8vo,
cloth $5.00
MOSES, A. J., and PARSONS, C. L. Elements of Miner-
alogy, Crystallography and Blowpipe Analysis from a Practical
Standpoint. Second Thousand. 8vo, cloth, 336illus.. net, $2.00
MOSES, A. J. The Characters of Crystals. An Intro-
duction to Physical Crystallograpky, containing 321 illustrations
and diagrams. 8vo, 21 1 pages net, $2 . 00
MOSS, S. A. The Lay-out of Corliss Valve Gears. (Van
Nostrand's Science Series, No. 119.) 16mo, cloth, illus 50
MULLIN, J. P., M.E. Modern Moulding and Pattern-
making. A Practical Treatise upon Pattern-shop and Foundry
Work: embracing the Moulding of Pulleys, Spur Gears, Worm
Gears, Balance-wheels, Stationary Engine and Locomotive
Cylinders, Globe Valves, Tool Work, Mining Machinery, Screw
Propellers, Pattern-shop Machinery, and the latest improve-
ments in English and American Cupolas; together with a large
collection of original and carefully selected Rules and Tables
for every-day use in the Drawing Office, Pattern-shop and Foundry.
12mo, cloth, illustrated $2 . 50
MUNRO, J., C.E., and JAMIESON, A., C.E. A Pocket-
book of Electrical Rules and Tables for the use of Electricians
and Engineers. Fifteenth Bdition, revised and enlarged. With
numerous diagrams. Pocket size. Leather $2 . 50
MURPHY, J. G., M.E. Practical Mining. A Field Manual
for Mining Engineers. With Hints for Investors in Mining
Properties. IGino, morocco tucks $1 .00
SCIENTIFIC PUBLICATIONS. 35
NAQUET, A. Legal Chemistry. A Guide to the Detection
of Poisons, Falsification of Writings, Adulteration of Alimentary
and Pharmaceutical Substances, Analysis of Ashes, and Exami-
nation of Hair, Coins, Arms and Stains, as applied to Chemical
Jurisprudence, for the use of Chemists, Physicians, Lawyers,
Pharmacists and Experts. Translated, with additions, including
a list of books and memoirs on Toxicology, etc., from the French,
by J. P. Battershall, Ph.D., with a Preface by C. F. Chandler,
Ph.D., M.D., LL.D. 12mo, cloth $2.00
NASMITH, J. The Student's Cotton Spinning. Third
Edition, revised and enlarged. 8vo, cloth, 622 pages, 250 illus-
trations $3 . 00
NEUBURGER, H., and NOALHAT, H. Technology of
Petroleum. The Oil Fields of the World: their History, Geog-
raphy and Geology. Annual Production, Prospection and Develop-
ment; Oil-well Drilling; Transportation of Petroleum by Land
and Sea. Storage of Petroleum. With 153 illustrations and 25
plates. Translated from the French, by John Geddes Mclntosh.
8vo, cloth, illustrated net, $10.00
NEWALL, J. W. Plain Practical Directions for Drawing,
Sizing and Cutting Bevel-gears, showing how the Teeth may
be cut in a Plain Milling Machine or Gear Cutter so as to give
them a correct shap3 from end to end; and showing how to get
out all particulars for the Workshop without making any Draw-
ings. Including, a Full Set of Tables of Reference. Folding
plates. 8vo, cloth $1 . 50
NEWLANDS, J. The Carpenters' and Joiners' Assistant:
being a Comprehensive Treatise on the Selection, Preparation
and Strength of Materials, and the Mechanical Principles of
Framing, with their application in Carpentry, Joinery and
Hand-railing; also, a Complete Treatise on Sines; and an Illus-
trated Glossary of Terms used in Architecture and Building.
Illustrated. Folio, half morocco $15.00
NIPHER, F. E., A.M. Theory of Magnetic Measurements,
with an Appendix on the Method of Least Squares. 12mo,
cloth $1 .00
NUGENT, E. Treatise on Optics; or, Light and Sight
Theoretically and Practically Treated, with the Application to
Fine Art and Industrial Pursuits. With 103 illustrations. 12mo,
cloth $1 . 50
O'CONNOR, H. The Gas Engineer's Pocket-book. Com-
prising Tables, Notes and Memoranda relating to the Manu-
facture, Distribution and Use of Coal-gas and the Construction
of Gas-works. Second Edition, revised. 12mo, full leather, gilt
edges $3 . 50
36 D. VAN NO3TRAND COMPANY'S
OLSEN, J. C., Prof. Text-book of Quantitative Chemica1
Analysi? by Gravimetric, Electrolytic, Volumetric and Gasometric
Methods. With Seventy-two Laboratory Exercises giving the
Analysis of Pure Salts, Alloys, Minerals and Technical Products.
With numerous figures and diagrams. 8vo. cloth., net, $4.00
OSBORN, F. C. Tables of Moments of Inertia, and Squares
of Radii of Gyration; supplemented by others on the Ultimate
and S ife Strength of Wrought-iron Columns, Safe Strength of
Timber Beams, and Constants for readily obtaining the Shearing
Stresses, Reactions and Bending Moments in Swing Bridges.
12mo, leather $3 . 00
03TERBERG, M. Synopsis of Current Electrical Litera-
ture, compiled from Technical Journals and Magazines during
1895. 8vo, cloth $1 . 00
OUDIN, M. A. Standard Polyphase Apparatus and Systems.
With many diagrams and figures. Third Edition, thoroughly
revised. Fully illustrated 13 . 00
PALAZ, A , Sc.D. A Treatise on Industrial Photometry,
with special application to Electric Lighting. Authorized trans-
lation from the French by George W. Patterson, Jr. Second
Edition, revised. 8vo, cloth, illustrated . $4 . 00
PARR, G. D. A. Electrical Engineering Measuring Instru-
ments, for Commercial and Laboratory Purposes. With 370
diagrams and engravings. 8vo, cloth, illustrated net, $3.50
PARRY, E. J., B.Sc. The Chemistry of Essential Oils
and Artificial Perfumes. Being an attempt to group together
the more important of the published facts connected with the
subject; also giving an outline of the principles involved in the
preparation and analysis of Essential Oils. With numerous dia-
grams and tables. 8vo, cloth, illustrated net, $5.00
and COSTE, J. H. Chemistry of Pigments. With
tables and figures. 8vo, cloth net, $4 . 60
PARRY, L. A., M.D. The Risks and Dangers f Various
Occupations and their Prevention. A book that should be in
the hands of manufacturers, the medical profession, sanitary
inspectors, medical officers of health, managers of works, foremen
and workmen. 8vo, cloth net, $3 00
PARSHALL, H. F., and HOBART, H. M. A mature
Winding.? of Electric Machines. With 140 full-page plates, 65
tables and 165 pages of descriptive letter-press. 4to, cloth.. $7 . 50
SCIENTIFIC PUBLICATIONS. 37
PARSHALL, H. F., and PARRY, E. Electrical Equipment
of Tramways In Press
PASSMORE, A. C. Handbook of Technical Terms used
in Architecture and Building, and their Allied Trades and Sub-
jects. 8vo, cloth net, $3 . 50
PATERSON, D., F.C.S. The Colo- Printing of Carpet
Yarns. A useful manual for color chemists and textile printers.
With numerous illustrations. 8vo, cloth, illustrated .... net, $3 . 50
Color Matching on Textiles. A Manual intended for
the use of Dyers, Calico Printers and Textile Color Chemists.
Containing colored frontispiece and 9 illustrations, and 14 dyed
patterns in appendix. 8vo, cloth, illustrated net, $3.00
The Science of Color Mixing. A Manual intended
for the use of Dyers, Calico Printers, and Color Chemists. With
figures, tables, and colored plate. 8vo, cloth, illustrated, net, $3 . 00
PATTEN, J. A Plan for Increasing th? Humidity of
the Arid Region and the Utilization of Some of the Great Rivers
of the United States for Power and other Purposes. A paper
communicated to the National Irrigation Congress, Ogden, Utah,
Sept. 12, 1903. 4to, pamphlet, 20 pages, with 7 maps $1 . 00
PAULDING, C. P. Practical Laws and Data on the Con-
densation of Steam in Covered and Bare Pipes; to which is added
a translation of P6clet's "Theory and Experiments on the Trans-
mission of Heat Through Insulating Materials." 8vo, cloth,
illustrated, 102 pages net, $2.00
PEIRCE, B. System of Analytic Mechanics. 4to,
cloth $10.00
PERRINE, F. A. C., A.M., D.Sc. Conductors for Elec-
trical Distribution: their Manufacture and Materials, the Calcu-
lation of Circuits, Pole Line Construction, Underground Working
and other Uses. With numerous diagrams and engravings. 8vo,
cloth, illustrated, 287 pages net, $3 . 50
Postage . 25
PERRY, J. Applied Mechanics. A Treatise for the use
of students who have time to work experimental, numerical and
graphical exercises illustrating the subject. 8vo, cloth, 650
pages net, $2 . 50
38 D. VAN NOSTRAND COMPANY'S
PHILLIPS, J. Engineering Chemistry. A Practical
Treatise for the use of Analytical Chemists, Engineers, Iron
Masters, Iron Founders, students and others. Comprising methods
of Analysis and Valuation of the principal materials used in
Engineering works, with numerous Analyses, Examples, and
Suggestions. Illustrated. Third Edition, revised and enlarged.
8vo, cloth $4 . 50
Gold Assaying. A Practical Handbook giving the
Modus Operand! for the Accurate Assay of Auriferous Ores and
Bullion, and the Chemical Tests required in the Processes of
Extraction by Amalgamation, Cyanidation and Chlorination.
With an appendix of tables and statistics and numerous diagrams
and engravings. 8vo, cloth, illustrated net, $2.50
PICKWORTH, C. N. The Indicator Handbook. A Prac-
tical Manual for Engineers. Part I. The Indicator: its Con-
struction and Application. 81 illustrations. 12mo, cloth... $1.50
The Indicator Handbook. Part II. The Indicator
Diagram: its Analysis and Calculation. With tables and figures.
12mo, cloth, illustrated $1 . 50
Logarithms for Beginners. 8vo, boards 50
The Slide Rule. A Practical Manual of Instruction for
all Users of the Modern Type of Slide Rule, containing Succinct
Explanation of the Principle of Slide-rule Computation, together
witn Numerous Rules and Practical Illustrations, exhibiting the
Application of the Instrument to the Every-day Work of the
Engineer — Civil, Mechanical and Electrical. Seventh Edition.
12mo, flexible cloth $1 . 00
Plane Table, The. Its Uses in Topographical Survey-
ing. From the Papers of the United States Coast Survey.
Illustrated. 8vo, cloth $2.00
"This work gives a description of the Plane Table employed at
the United States Coast Survey office, and the manner of using it. "
PLANTE, G. The Storage of Electrical Energy, ard
Researches in the Effects created by Currents, combining Quan-
tity with High Tension. Translated from the French by Paul
B. Elwell. 89 illustrations. 8vo, cloth $4 . 00
PLATTNER'S Manual of Qualitative and Quantitative
Analysis with the Blow-pipe. Eighth Edition, revised. Translated
by Henry B. Cornwall, E.M., Ph.D., assisted by John H. Caswell,
A.M. From the sixth German edition, by Prof. Friederich Kol-
beck. With 87 woodcuts. 463 pages. 8vo, cloth . . . net, $4 . 00
SCIENTIFIC PUBLICATIONS. 39
PLYMPTON, GEO. W., Prof. The Aneroid Barometer:
its Construction and Use. Compiled from several sources.
Eighth Edition, revised and enlarged. 16mo, boards, illus-
trated $0 . 50
Morocco 1 . 00
POCKET LOGARITHMS, to Four Places of Decimals,
including Logarithms of Numbers, and Logarithmic Sines and
Tangents to Single Minutes. To which is added a Table of
Natural Sines, Tangents and Co-tangents. 16mo, boards 50
POPE F. L. Modern Practice of the Electric Telegraph.
A Technical Handbook for Electricians, Managers and Operators.
Fifteenth Edition, rewritten and enlarged, and fully illustrated. 8vo,
cloth $1 . 50
POPPLEWELL, W. C. Elementary Treatise on Heat and
Heat Engines. Specially adapted for engineers and students of
engineering. 12mo, cloth, illustrated $3 .00
Prevention of Smoke, combined with the Economical
Combustion of Fuel. With diagrams, figures and tables. 8vo,
cloth, illustrated net, $3 . 50
POWLES, H. H. P. Steam Boilers In Press.
Practical Compounding of Oils, Tallow and Grease, for
Lubrication, etc. By an Expert Oil Refiner. 8vo, cloth, net, $3 . 50
Practical Iron Founding. By the Author of " Pattern
Making," etc. Illustrated with over 100 engravings. Third
Edition. 12mo, cloth $1 . 50
PRAY, T., Jr. Twenty Years with the Indicator: being
a Practical Text-book for the Engineer or the Student, with no
complex Formulae. Illustrated. 8vo, cloth $2 . 50
Steam Tables and Engine Constant. Compiled from
Regnault, Rankine and Dixon directly, making use of the
exact records. 8vo, cloth. '. . $2 . 00
PREECE, W. H. Electric Lamps In Press.
- and STUBBS, A. T. Manual of Telephony. Illus-
trations and plates. 12mo, cloth $4 . 50
PRELINI, C., C.E. Tunneling. A Practical Treatise con-
taining 149 Working Drawings and Figures. With additions
by Charles S. Hill, C.E., Associate Editor "Engineering N«ws."
311 pages. Second Edition, revised. 8vo, cloth, illus. . . . $3.00
40 D. VAN NOSTRAND COMPANY'S
PRELINI, C., C.E. Earth and Rock Excavation. A Manual
for Engineers, Contractors and Engineering Students. 8vo,
cloth, illustrated. About 350 pages In Press.
PREMIER CODE. (See Hawke, Wm. H.)
PRESCOTT, A. B., Prof. Organic Analysis. A Manual
of the Descriptive and Analytical Chemistry of certain Carbon
Compounds in Common Use; a Guide in the Qualitative and
Quantitative Analysis of Organic Materials in Commercial and
Pharmaceutical Assays, in the estimation of Impurities under
Authorized Standards, and in Forensic Examinations for Poisons,
with Directions for Elementary Organic Analysis. Fifth Edi-
tion. 8vo, cloth $5.00
- Outlines of Proximate Organic Analysis, for the
Identification, Separation and Quantitative Determination of
the more commonly occurring Organic Compounds. Fourth
Edition. 12mo, cloth $1 . 75
- and SULLIVAN, E. C. (University of Michigan). First
Book in Qualitative Chemistry. For Studies of Water Solution
and Mass Action. Twelfth Edition, entirely rewritten. I2mo,
cloth net, $1 . 50
- and JOHNSON, 0. C. Qualitative Chemical Analysis.
A Guide in the Practical Study of Chemistry and in the Work of
Analysis. Fifth revised and enlarged edition, entirely rewritten.
With Descriptive Chemistry extended throughout net, $3 . 50
PRITCHARD, 0. G. The Manufacture of Electric-light
Carbons. Illustrated. 8vo, paper 60
PULLEN, W. W. F. Application of Graphic Methods
to the Design of Structures. Specially prepared for the use of
Engineers. A Treatment by Graphic Methods of the Forces
and Principles necessary for consideration in the Design of En-
gineering Structures, Roofs, Bridges, Trusses, Framed Structures,
Wells, Dams, Chimneys, and Masonry Structures. 12mo, cloth,
profusely illustrated net, $2 . 50
PULSIFER, W. H. Notes for a History of Lead. 8vo,
cloth, gilt top $4 .00
PUTSCH, A. Gas and Coal-dust Firing. A Critical Review
of the Various Appliances Patented in Germany for this purpose
since 1885. With diagrams and figures. Translated from the
German by Charles Salter. 8vo, cloth, illustrated. . . . net, $3.00
SCIENTIFIC PUBLICATIONS. 41
PYNCHON, T R., Prof. Introduction to Chemical Physics,
designed for the use of Academies, Colleges and High Schools.
Illustrated with numerous engravings, and containing copious
experiments, with directions for preparing them. New Edition,
revised and enlarged, and illustrated by 269 wood engravings, 8vo,
cloth $3 .00
RADFORD, C. S., Li ut. Handbook en Naval Gunnery.
Prepared by Authority of the Navy Department. For the use
of U. S. Navy, U. S. Marine Corps, and U. S. Naval Reserves.
Revised and enlarged, with the assistance of Stokely Morgan,
Lieut. U. S. N. Third Edition, revised and enlarged. 12mo,
flexible leather $2.00
RAFTER, G. W. Treatment of Septic Sewage (Van
Nostrand's Science Series, No. 118). 16mo, cloth 50
Tables for Sewerage and Hydraulic Engineers. In Press,
and BAKER, M. N. Sewage Disposal in the United
States. Illustrations and folding plates. Third Edition. 8vo,
cloth $6.00
RAM, G. S. The Incandescent Lamp and its Manufac-
ture. 8vo, cloth. $3 .00
RAMP, H. M. Foundry Practice In Press.
RANDALL, J. E. A Practical Treatise on the Incan-
descent Lamp. 16mo, cloth, illustrated 50
RANDALL, P. M. Quartz Operator's Handbook. New
Edition, revised and enlarged, fully illustrated. 12mo, cloth, $2.00
RANDAU, P. Enamels and Enamelling. An introduction
to the preparation and application of all kinds of enamels for
technical and artistic purposes. For enamel-makers, workers
in gold and silver, and manufacturers of objects of art. Third
German Edition. Translated by Charles Salter. With figures,
diagrams and tables. 8vo, cloth, illustrated net, $4.00
RANKINE, W. J. M. Applied Mechanics. Comprising
the Principles of Statics and Cinematics, and Theory of Struc-
tures, Mechanism, and Machines. With numerous diagrams.
Seventeenth Edition, thoroughly revised by W. J. Millar. 8vo,
cloth $5.00
Civil Engineering. Comprising Engineering Sur-
veys, Earthwork, Foundations, Masonry, Carpentry, Metal-
work, Roads, Railways, Canals, Rivers, Water-works, Harbors,
etc. With numerous tables and illustrations. Twenty-first
Edition, thoroughly revised by W. J. Millar. 8vo, cloth. ... $6 . 50
42 D. VAN NOSTRAND COMPANY'S
RANKINE W. J. M. Machinery and Millwork. Com-
prising the Geometry, Motions, Work, Strength, Construction, and
Objects of Machines, etc. Illustrated with nearly 300 woodcuts.
Seventh Edition, thoroughly revised by W. J. Millar. Svo,
cloth $5.00
The Steam-engine and Other Prime Movers. With
diagram of the Mechanical Properties of Steam. Folding plates,
numerous tables and illustrations. Fifteenth Edition, thor-
oughly revised by W. J. Millar. 8vo, cloth. $5.00
Useful Rules and Tables for Engineers and Others.
With Appendix, Tables, Tests and Formula* for the use of Elec-
trical Engineers. Comprising Submarine Electrical Engineering,
Electric Lighting and Transmission of Power. By Andrew
Jamieson, C.E., P.R.S.E. Seventh Edition, thoroughly revised
by W. J. Millar. 8vo, cloth $4 . 00
- and BAMBER, E. F., C.E. A Mechanical Text-book.
With numerous illustrations. Fifth Edition. 8vo, cloth. . $3.50
RAPHAEL, F. C. Localization of Faults in Electric
Light and Power Mains, with chapters on Insulation Testing
With figures and diagrams. Second Edition, revised. 8vo'
cloth, illustrated net, $3 . 00.
RAYMOND, E. B. Alternating-current Engineering Prac-
tically Treated. With numerous diagrams and figures. 12mo,
cloth net, $2.50
RAYNER, H. Silk Throwing and Waste Silk Spinning.
With numerous diagrams and figures. 8vo, cloth, illustrated,
net, $2 . 50
RECIPES for the Color, Paint, Varnish, Oil, Soap and
Drysaltery Trades. Compiled by an Analytical Chemist. 8vo,
cloth $3 . 50
RECIPES FOR FLINT GLASS MAKING. Being Leaves
form the mixing-book of several experts in the Flint Glass Trade.
Containing up-to-date recipes and valuable information as to
Crystal, Demi-crystal, and Colored Glass in its many varieties.
It contains the recipes for cheap metal suited to pressing, blowing,
etc., as well as the most costly Crystal and Ruby. British manu-
facturers hare kept up the quality of this glass from the arrival of
the Venetians to Hungry Hill, Stourbridge, up to the present
time. The book also contains remarks as to the result of the
metal as it left the pots by the respective metal mixers, taken
from their own memoranda upon trie originals. Compiled by
a British Glass Master and Mixer. 12mo, cloth net, $4.. 50
SCIENTIFIC PUBLICATIONS. 43
REED'S ENGINEERS' HANDBOOK, to the Local Marine
Board Examinations for Certificates of Competency as First and
Second Class Engineers. By W. H. Thorn. With the answers
to the Elementary Questions. Illustrated by 358 diagrams
and 37 large plates. Seventeenth Edition, revised and enlarged.
8vo, cloth $5.00
Key to the Seventeenth Edition of Reed's Engineer's
Handbook to the Board of Trade Examination for First and
Second Class Engineers and containing the workings of all the
questions given in the examination papers. By W. H. Thorn.
8vo, cloth $3 .00
Useful Hints to Sea-going Engineers, and How to
Repair and Avoid "Breakdowns"; also Appendices containing
Boiler Explosions, Useful Formulae, etc. With 42 diagrams
and 8 plates. Third Edition, revised and enlarged. 12mo,
cloth $1 .50
Marine Boilers. A Treatise on the Causes and Pre-
vention of their Priming, with Remarks on their General Manage-
ment. 12mo, cloth, illustrated $2 . 00
REINHARDT, C. W. Lettering for Draftsmen, Engineers,
and Students. A Practical System of Free-hand Lettering for
Working Drawings. Revised and enlarged edition. Thirteenth
Thousand. Oblong, boards $1 . 00
The Technic of Mechanical Drafting. A Practical
guide to neat, correct and legible drawing, containing many illus-
trations, diagrams and full-page plates. 4to, cloth, illus. . . $1 .00
REISER, F. Hardening and Tempering of Steel, in Theory
and Practice. Translated from the German of the third and
enlarged edition, by Arthur Morris and Herbert Robson. 8vo,
cloth, 120 pages $2 . 50
REISER, N. Faults in the Manufacture of Woolen Goods,
and their Prevention. Translated from the second German
edition, by Arthur Morris and Herbert Robson. 8vo, cloth,
illustrated net, $2 . 50
Spinning and Weaving Calculations with special
reference to Woolen Fabrics. Translated from the German by
Chas. Salter 8vo, cloth, illustrated net, $5.00
RICE, J. M., and JOHNSON, W. W. On a New Method
of Obtaining the Differential of Functions, with especial refer-
ence to the Newtonian Conception of Rates or Velocities. 12mo,
paper 50
44 D. VAN NOSTRAND COMPANY'S
RIDEAL, S., D.Sc. Glue and Glue Testing, with figures
and tables. 8vo, cloth, illustrated net, $4 . 00
RINGWALT, J. L. Development of Transportation Sys-
tems in the United States, comprising a Comprehensive Descrip-
tion of the leading features of advancement from the colonial
era to the present time, in water channels, roads, turnpikes,
canals, railways, vessels, vehicles, cars, and locomotives; the
cost of transportation at various periods and places by the differ-
ent methods; the financial, engineering, mechanical, govern-
mental and popular questions that have arisen, and notable
incidents in railway history, construction and operation. With
illustrations of hundreds of typical objects. Quarto, half mo-
rocco $7 . 50
RIPPER, W. A Course of Instruction in Machine Drawing
and Design for Technical Schools and Engineer Students. With
52 plates and numerous explanatory engravings. Folio, cloth,
net, $6.00
ROBERTSON, L. S. Water-tube Boilers. Based on a
short course of Lectures delivered at University College, London.
With upward of 170 illustrations and diagrams. 8vo, cloth,
illustrated $3.00
ROEBLING, J. A. Long and Short Span Railway Bridges.
Illustrated with large copper-plate engravings of plans and views.
Imperial folio, cloth $25 00
ROSE, J., M.E Th? Pattern-makers' Assistant. Embrac-
ing Lathe Work, Branch Work, Core Work, Sweep Work and
Practical Gear Constructions, the Preparation and Use of Tools,
together with a large collection of useful and valuable Tables.
Ninth Edition. With 250 engravings. 8vo, cloth $2.50
Key to Engines and Engine-running. A Practical
Treatise upon the Management of Steam-engines and Boilers for
the use of those who desire to pass an examination to take
charge of an engine or boiler. With numerous illustrations,
and Instructions upon Engineers' Calculations, Indicators,
Diagrams, Engine Adjustments and other Valuable Information
necessary for Engineers and Firemen. 12mo, cloth. Illus. . $2.50
ROWAN, F. J. The Practical Physics of the Modern
Steam-boiler. With an Introduction by Prof. R. H. Thurston.
With numerous illustrations and . diagrams. 8vo, cloth, illus-
trated $7 . 50
SABINE, R. History and Progress of the Electric Tele-
f-aph. With descriptions of some of the apparatus. Second
dition, with additions. 12mo, cloth $1 25
SCIENTIFIC PUBLICATIONS. 45
SAELTZER, A. Treatise on Acoustics in Connection
with Ventilation. 12mo, cloth $1 . 00
SALOMONS, Sir D., M.A. Electric-light Installations.
A Practical Handbook. Eighth Edition, revised and enlarged,
with numerous illustrations. Vol. I., The Management of Accu-
mulators. 12mo, cloth. . $1 . 50
Vol. II., Apparatus. 296 illustrations. 12mo, cloth $2.25
Vol. III., Applications. 12mo, cloth $1 . 50
SANFORD, P. G. Nitro-explosives. A Practical Treatise
concerning the Properties, Manufacture and Analysis of Nitrated
Substances, including the Fulminates, Smokeless Powders and
Celluloid. 8vo, cloth, 270 pages $3 . 00
SAUNDERS, C. H. Handbook of Practical Mechanics
for use in the Shop and Draught ing-room; containing Tables,
Rules and Formulae, and Solutions of Practical Problems by
Simple and Quick Methods. 16mo, limp cloth $1 . 00
SAUNNIER, C. Watchmaker's Han; book. A Workshop
Companion for those engaged in Watchmaking and allied Mechan-
ical Arts. Translated by J. Tripplin and E. Rigg. Second Edi-
tion, revised, with appendix. 12mo, cloth $3 . 50
SCHELLEN, H., Dr. Magneto-electric and Dynamo-
electric Machines: their Construction and Practical Application
to Electric Lighting, and the Transmission of Power. Trans-
lated from the third German edition by N. S. Keith and Percy
Neymann, Ph.D. With very large additions and notes relating
to American Machines, by N. S. Keith. Vol. 1, with 353 illus-
trations. Second Edition. 8vo, cloth $5 . 00
SCHUMANN, F. A Manual of Heating and Ventilation
in its Practical Application, for the use of Engineers and Archi-
tects. Embracing a Series of Tables and Formulae for Dimensions
of Heating, Flow and Return Pipes for Steam and Hot-water
Boilers, Flues, etc. 12mo, illustrated, full roan $1 .50
SCHMALL, C. N., and SHACK, S. M. Elements of Plane
Geometry. An Elementary Treatise. With many examples and
diagrams. 12mo, half leather, illustrated net, $1 .25
SCIENCE SERIES, The Van Nostrand. (Follows end of
this list.)
46 D. VAN NOSTRAND COMPANY'S
SCRIBNER, J. M. Engineers' and Mechanics' Com-
panion. Comprising United States Weights and Measures.
Mensuration of Superfices and Solids, Tables of Squares and
Cubes, Square and Cube Roots, Circumference and Areas of
Circles, the Mechanical Powers, Centres of Gravity, Gravitation
of Bodies, Pendulums, Specific Gravity of Bodies, Strength,
Weight and Crush of Materials, Water-wheels. Hydrostatics,
Hydraulics, Statics, Centres of Percussion and Gyration, Friction
Heat, Tables of the Weight of Metals, Scantling, etc., Steam
and Steam-engine. Twenty-first Edition, revised. 16mo, full
morocco $1 . 50
SEATON, A. E. A Manual of Marine Engineering. Com-
prising the Designing, Construction and Working of Marine
Machinery. With numerous tables and illustrations reduced from
Working Drawings. 14th Edition, revised throughout, with an
additional chapter on Water-tube Boilers. 8vo, cloth $6 . 00
and ROUNTHWAITE, H. M. A Pocketbook of
Marine Engineering Rules and Tables. For the use of Marine
Engineers and Naval Architects, Designers, Draughtsmen,
Superintendents and all engaged in the design and construction
of Marine Machinery, Naval and Mercantile. Seventh Edition,
revised and enlarged. Pocket size. Leather, with diagrams. $3.00
SEELIGMANN, T., TORRILHON, G. L., and FALCONNET,
H. India Rubber and Gutta Percha. A complete practical
treatise on India Rubber and Gutta Percha, in their historical,
botanical, arboricultural, mechanical, chemical and electrical
aspects'. Translated from the French, by John Geddes Mclntosh.
8vo, cloth, illustrated, 412 pages net, $7 .50
SEVER, G. F., Prof. Electrical Engineering Experi-
ments and Tests on Direct-current Machinery. With diagrams
and figures. 8vo pamphlet, illustrated •. . . . .net, $1 .00
-and TOWNSEND, F. Laboratory and Factory Tests
in Electrical Engineering. 8vo, cloth, illustrated, about 225
pages . In Press.
SEWALL, C. H. Wireless Telegraphy. With diagrams
and engravings. Second Edition, corrected. 8vo, cloth, illus-
trated net, $2.00
- Lessons in Telegraphy. For use as a text-book
in schools and colleges, or for individual students. Illustrated.
12mo, cloth $1.(;0
SEWELL, T. Elements of Electrical Engineering. A
First Year's Course for Students. Second Edition, revised, with
additional chapters on Alternating-current Working and Ap-
pendix of Questions and Answers. With many diagrams, tables
and examples. 8vo, cloth, illustrated, 432 pages net, $3 . 00
SCIENTIFIC PUBLICATIONS. 47
SEXTON, A. H. Fuel and Refractory Materials. 8vo,
cloth $2.00
Chemistry of the Materials of Engineering. A Hand-
book for Engineering Students. With tables, diagrams and
illustrations. 12mo, cloth, illustrated $2 . 50
SEYMOUR, A. Practical Lithography. With figures and
engravings. 8vo, cloth, illustrated net, $2 . 50
SHAW, S. The History of the Staffordshire Potteries, and
the Rise and Progress of the Manufacture of Pottery and Por-
celain; with references to genuine specimens, and notices of
eminent potters. A re-issue of the original work published in
1829. 8vo, cloth, illustrated net, $3 . 00
Chemistry of the Several Natural and Artificial
Heterogeneous Compounds used in Manufacturing Porcelain,
Glass and Pottery. Re-issued in its original form, published in
1837. 8vo, cloth net, $5.00
SHAW, WM. J., Lieut-Col. Studies in Map Reading
and Field Sketching. An aid to passing outdoor examinations in
these subjects. Illustrated with 15 folding plates. 12mo, cloth,
illustrated, 148 pages net, $2 . 50
Tactical Operations for Field Officers : being Up-to-
date schemes worked out on training grounds at home stations.
With folding platea and maps. 12mo, cloth, illustrated, 321.
pages $3 .00
SHELDON, S., Ph.D., and MASON, H., B.S. Dynamo-
electric Machinery: its Construction, Design and Operation.
Direct-current Machines. Fourth Edition, revised. 8vo, cloth,
illustrated net, $2 . 50
Alternating-current Machines : being the second
volume of the author's "Dynamo-electric Machinery: its Construc-
tion, Design and Operation. " With many diagrams and figures.
(Binding uniform with volume I.) Third Edition. 8vo, cloth,
illustrated net, $2 . 50
SHERRIFF, F. F. Oil Merchants' Manual, and Oil Trade
Ready Reckoner. With two sheets of tables. Second Editiont
revised and enlarged. 8vo, cloth net, $3.50
SHIELDS, J. E. Notes on Engineering Construction.
Embracing Discussions of the Principles involved, and Descrip-
tions of the Material employed in Tunneling, Bridging, Canal and
Road Building, etc. 12mo, cloth $1 . 50
48 D. VAN NOSTRAND COMPANY'S
SHOCK, W. H. Steam Boilers: their Design, Construc-
tion and Management. 4to, half morocco $15.00
SHREVE, S. H. A Treatise on the Strength of Bridges
and Roofs. Comprising the determination of algebraic formulas
for strains in Horizontal, Inclined or Rafter, Triangular, Bow,
string, Lenticular and other Trusses, from fixed and moving loads-
with practical applications and examples, for the use of Students
and Engineers. 87 woodcut illustrations. Fourth Edition. 8vo,
cloth $3.50
SHUNK, W. F. The Field Engineer. A Handy Book
of practice in the Survey, Location and Truck-work of Railroads,
containing a large collection of Rules and Tables, original and
selected, applicable to both the Standard and Narrow Gauge,
and prepared with special reference to the wants of the young
Engineer. Sixteenth Edition, revised and enlarged. With
addenda. 12mo, morocco, tucks $2 . 50
SIMMS, F. W. A Treatise on the Principles and Practice
of Leveling. Showing its application to purposes of Railway
Engineering, and the Construction of Roads, etc. Revised and
corrected, with the addition of Mr. Laws' Practical Examples for
setting out Railway Curves. Illustrated. 8vo, cloth $2 . 50
Practical Tunneling. Fourth Edition, Revised and
greatly extended. With additional chapters illustrating recent
practice by D. Kinnear Clark. With 36 plates and other illustra-
tions. Imperial 8vo, cloth $12 .00
SIMPSON, G. The Naval Constructor. A Vade Mecum
of Ship Design, for Students, Naval Architects, Ship Builders and
Owners, Manne Superintendents, Engineers and Draughtsmen.
12mo, morocco, illustrated, 500 pages $5 . 00
SLATER, J. W. Sewage Treatment, Purification and
Utilization. A Practical Manual for the Use of Corporations,
Local Boards, Medical Officers of Health, Inspectors of Nuisances,
Chemists, Manufacturers, Riparian Owners, Engineers and Rate-
payers. 12mo, cloth $2 .25
SMITH, I. W., C.E The Theory of Deflections and of
Latitudes and Departures. With special applications to Curvi-
linear Surveys, for Alignments of Railway Tracks. Illustrated.
IGmo, morocco, tucks $3 . 00
SMITH, J. C. Manufacture of Paint. A Practical Hand-
book for Paint Manufacturers, Merchants and Painters. With
60 illustrations and one large diagram. 8vo, cloth. . . .net, $3.00
SCIENTIFIC PUBLICATIONS. 49
SNELL, A T. Electric Motive Power: The Transmission
and Distribution of Electric Power by Continuous and Alternate
Currents. With a Section on the Applications of Electricity to
Mining Work. Second Edition. 8vo, cloth, illustrated $4 .00
SOXHLET, D. H Art of Dyeing and Staining Marble,
Artificial Stone, Bone, Horn, Ivory and Wood, and of imitating
all sorts of Wood. A practical Handbook for the use of Joiners,
Turners, Manufacturers of Fancy Goods, Stick and Umbrella
Makers, Comb Makers, etc. Translated from the German by
Arthur Morris and Herbert Robson, B.Sc. 8vo, cloth, 170
pages net, $2 . 50
SPANG, H. W A Practical Treatise on Lightning Pro-
tection. With figures and diagrams. 12mo, cloth $1 .00
SPEYERS, C. L. Text-book of Physical Chemistry.
8vo, cloth $2 .25
STAHL, A. W., and WOODS, A. T. Elementary Mechan-
ism. A Text-book for Students of Mechanical Engineering.
Thirteenth Edition, enlarged. 12mo, cloth $2 . 00
STALEY, C., and PIERSON, G. S. The Separate System
of Sewerage: its Theory and Construction. Third Edition,
revised and enlarged. With chapter on Sewage Disposal. With
maps, plates and illustrations. 8vo, cloth $3 . 00
STANDAGE H. C. Leatherworkers' Manual: being a
Compendium of Practical Recipes and Working Formulae for
Curriers, Boot-makers, Leather Dressers, Blacking Manufac-
turers, Saddlers, Fancy Leather Workers and all persons en-
gaged in the manipulation of leather. 8vo, cloth net, $3 . 50
Sealing Waxes, Wafers and Other Adhesive \ For
the Household, Office, Workshop and Factory. 8vo, cloth, 96
pages net, $2.00
STEWART R. W. A Text-book of Light. Adapted to
the Requirements of the Intermediate Science and Preliminary
Scientific Examinations of the University of London, and also
for General Use. Numerous diagrams and examples. 12mo
cloth $1 .00
Text-book of Heat. Illustrated. 8vo, Cloth, $1.00
Text-book of Magnetism and Electricity. 160 Illus-
trations and numerous examples. 12mo, cloth $1 .00
Elementary Text-book of Magnetism and Electricity.
With numerous figures and diagrams. 12mo, cloth $1 .00
60 D. VAN NOSTRAND COMPANY'S
STILES, A. Tables for Field Engineers. Designed for
Use in the Field. Tables containing all the Functions of a One
Degree Curve, from which a corresponding one can be found for
any required Degree. Also, Tables of Natural Sines and Tangents.
12mo, morocco, tucks $2.00
STILLMAN, P. Steam-engine Indicator and the Improved
Manometer Steam and Vacuum Gauges; their Utility and Appli-
cation. New edition. 12mo, flexible cloth $1 . 00
STODOLA, A. Dr. The Steam Turbine. With 300 dia-
grams and illustrations. Authorized translation by Dr. L. C.
Loewenstein In Press.
STONE, R., Gen'l. New Roads and Road Laws in the
United States. 200 pages, with numerous illustrations. 12mo,
cloth $1 . 00
STONEY, B. D. The Theory of Stresses in Girders and
Similar Structures. With Observations on the Application of
Theory to Practice, and Tables of Strength, and other Properties
of Materials. New revised edition, with numerous additions on
Graphic Statics, Pillars, Steel, Wind Pressure, Oscillating Stresses,
Working Loads, Riveting, Strength and Tests of Materials.
777 pages, 143 illus. and 5 folding-plates. 8vo, cloth $12.50
STUART, C. B., U.S.N. Lives and Works of Civil and
Military Engineers of America. With 10 steel-plate engravings.
Svo, cloth $5.00
- The Naval Dry Docks of the United States. Illus-
trated with 24 fine engravings on steel. Fourth Edition. 4to,
cloth $6.00
SUPPLING, E. R. Treatise on the Art of Glass Painting.
Prefaced with a Review of Ancient Glass. With engravings and
colored plates. Svo, cloth net, $3 . 50
SWEET, S. H. Special Report on Coal, Showing its Dis-
tribution, Classification, and Costs delivered over Different Routes
to Various Points in the State ol New York and the Principal
Cities on the Atlantic Coast. With maps. Svo, cloth $3.00
SWOOPE, C. W. P actical Lessons in Electricity: Prin-
ciples, Experiments and Arithmetical Problems. An Elementary
Text-book. With numerous tables, formulae, and two large in-
struct ion plates. Svo, cloth, illustrated. Fifth Edition . . net , $2 . 00
SCIENTIFIC PUBLICATIONS. 51
TAILFER, L. Practical Treatise on the Bleaching of
Linen and Cotton Yarn and Fabrics. With tables and diagrams.
Translated from the French by John Geddes Mclntosh. 8vo,
cloth, illustrated net, $5 . 00
TEMPLETON, W. The Practical Mechanic's Workshop
Companion. Comprising a great, variety of the most useful
rules and formulae in Mechanical Science, with numerous tables
of practical data and calculated results facilitating mechanical
operations. Revised and enlarged by W. S. Hutton. 12mo,
morocco $2 . 00
TESLA, N. Experiments with Alternate Currents of High
Potential and High Frequency. A Lecture delivered before the
Institution of Electrical Engineers, London. With a portrait
and biographical sketch of the author. With figures and dia-
grams. 12mo, cloth, illustrated. New Edition In Press.
THOM, C., and JONES, W. H. Telegraphic Connections:
embracing Recent Methods in Quadruplex Telegraphy. 20 full-
page plates, some colored. Oblong, 8vo, cloth $1 .50
•
THOMAS, C. W. Paper-makers' Handbook. A Practical
Treatise. Illustrated •. . In Press.
THOMPSON, A. B. Oil Fields of Russia and the Russian
Petroleum Industry. A Practical Handbook on the Explora-
tion, Exploitation, and Management of Russian Oil Properties,
including Notes on the Origin of Petroleum in Russia, a Descrip-
tion of the Theory and Practice of Liquid Fuel, and a Translation
of the Rules and Regulations concerning Russian Oil Properties.
With numerous illustrations and photographic plates and a map
of the Balakhany-Saboontchy-Romany Oil Field. 8vo, cloth,
illustrated net, $15.00
THOMPSON, E. P., M.E. How to Make Inventions;
or, Inventing as a Science* and an Art. A Practical Guide for
Inventors. Second Edition. 8vo, boards $1 .00
Roentgen Rays and Phenomena of the Anode and
Cathode. Principles, Applications, and Theories. For Students,
Teachers, Physicians, Photographers, Electricians and others.
Assisted by Louis M. Pignolet, N. D. C. Hodges and Ludwig
Gutmann, E.E. With a chapter on Generalizations, Arguments,
Theories, Kindred Radiations and Phenomena. By Professor
Wm. Anthony. 50 diagrams, 40 half-tones. 8vo, cloth. . . $1 .50
52 D. VAN NOSTRAND COMPANY'S
THORNLEY, T. Cotton Combing Machines. With Nu-
merous tables, engravings and diagrams. £vo, cloth, illustrated,
343 pages net, $3 .00
Contents. — Preface; List of Illustrations; The Silver Lap Ma-
chine; Ribbon Lap Machine and Draw-frame; General Descrip-
tion of the Heilmann Comber; The Cam Shaft; The Detaching
and Attaching Mechanism of the Comber; The Duplex Comber;
Resetting of Combers; The Erection of a Heilmann Comber;
Stop Motions; Various Calculations; Various Notes and Dis-
cussions; Cotton Combing Machines of Continental Make; Index.
TODD, J., and WHALL, W. B. Practical Seamanship
for Use in the Merchant Service: including all ordinary subjects;
also Steam Seamanship, Wreck Lifting, Avoiding Collision. Wire
Splicing, Displacement and everything necessary to be known
by seamen of the present dav. Fifth Edition, with 247 illus-
trations and diagrams. 8vo, cloth net, $7 . 50
TOOTHED GEARING. A Practical Hand book for Offices
and Workshops. By a Foreman Patternmaker. 184 illustra-
tions. 12mo, cloth $2.25
TRATMAN, E. E. R. Railway Track and Track-work.
With over 200 illustrations. 8vo, cloth $3 . 00
TRAVERSE TABLE, Showing Latitude and Departure
for each Quarter Degree of the Quadrant, and for Distances from 1
to 100, to which is appended a Table of Natural Siw»s and Tan-
gents for each five minutes of the Quadrant. (Reprinted from
Scribner's Pocket Table Book.) Van Nostrand's Science Series.
16mo, cloth $0.50
Morocco $1.00
TREVERT, E. How to Bu.ld Dynamo-electric Machinery,
embracing Theory, Designing, and Construction of Dynamos and
Motors. With appendices on Field Magnet and Armature
Winding, Management of Dynamos and Motors, and Useful
Tables»of Wire Gauges. 8vo, cloth, illustrated $2 50
Electricity and its Recent Applications. A Practical
Treatise for Students and Amateurs, with an Illustrated Dictionary
of Electrical Terms and Phrases. 12mo, cloth. $2 00
TUCKER, J. H., Dr. A Manual of Sugar Analysis, in-
cluding the Applications in General of Analytical Methods to the
Sugar Industry. With an Introduction on the Chemistry of
Cane-sugar, Dextrose, Levulose and Milk-sugar. 8vo, cloth,
illustrated $3.50
SCIENTIFIC PUBLICATIONS. 53
TUMLIRZ, O., Dr. Potential and it; Application to the
Explanation of Electric Phenomena, Popularly Treated. Trans-
• lated from the German by D. Robertson. 12mo, cloth, ill. $1 .25
TUNNER, P. A. Treatis3 on Roll-turning for the Manu-
facture of Iron. Translated and adapted by John B. Pearse, of
the Pennsylvania Steel Works, with numerous engravings, wood-
cuts. 8vo, cloth, with folio atla* of plates $10 . 00
UNDERBILL, C. R. The Electro-magnet. N w and
revised edition. 8vo, cloth, illustrated $1 . 50
URQUHART, J. W. Electric Light Fitting. Embodying
Practical Notes on Installation Management. A Handbook for
Working Electrical Engineers. With numerous illustrations.
12mo, cloth $2.00
Electro-plating. A Practical Handbook on the Depo-
sition of Copper, Silver, Nickel, Gold, Brass, Aluminium, Plati-
num, etc. Fourth Edition. 12mo. $2 . 00
Electrotyping. A Practical Manual Forming a New
and Systematic Guide to the Reproduction and Multiplication of
Printing Surfaces, etc. 12mo $2.00
Dynamo Construction. A Practical Handbook for the
Use of Engineer Constructors and Electricians in Charge, em- .
bracing Frame Work Building, Field Magnet and Armature
Winding and Grouping, Compounding, etc., with Examples of
Leading English, American and Continental Dynamos and Motors,
with numerous illustrations. 12mo, cloth $3 . 00
Electric Ship Lighting. A Handbook on the Practical
Fitting and Running of Ship's Electrical Plant. For the Use of
Ship Owners, and Builders, Marine Electricians and Sea-going
Engineers-in-Charge. Numerous illustrations. 12mo, cloth,
$3.00
UNIVERSAL TELEGRAPH CIPHER CODE. Arranged
for General Correspondence. 12mo, cloth $1 . 00
VAN NOSTRAND'S Engineering Magazine. Complete Sets,
1869 to 1886 inclusive. 35 vols., in cloth $60 . 00
in half morocco $100.00
Year Book of Mechanical Engineering Data. With
many tables and diagrams. (First Year of issue 1905.) In Press.
64 D. VAN NOSTRAND COMPANY'S
VAN WAGENEN, T. F. Manual of Hydraulic Mining.
For the Use of the Practical Miner. Revised and enlarged edition.
18mo, cloth 81 . 00
VILLON, A. M. Practical Treatise on the Leather Industry.
With many tables and illustrations and a copious index. A trans-
lation of Villon's "Traite Pratique de la Fabrication des Cuirs et
du Travail des Peaux," by Frank T. Addyman, B.Sc. 8vo,
cloth, illustrated net, $10.00
VINCENT, C. Ammonia and its Compounds: their
Manufacture and Uses. Translated from the French by M. J.
Salter. 8vo, cloth, illustrated net, $2.00
VOLK, C. Haulage and Winding Appliances Used in
Mines. With plates and engravings. Translated from the Ger-
man. 8vo, cloth, illustrated net, $4 . 00
VON GEORGIEVICS, G. Chemical Technology of Textile
Fibres: their Origin, Structure, Preparation, Washing, Bleaching,
Dyeing, Printing, and Dressing. Translated from the German
by Charles Salter. With many diagrams and figures. 8vo, cloth,
illustrated. 306 pages net, $4 . 50
Contents. — The Textile Fibres; Washing, Bleaching, and Car-
bonizing; Mordants and Mordanting; Dyeing, Printing, Dressing
and Finishing; Index.
— Chemistry of Dyestuffs. Translated from the Second
German edition by Chas. Salter. 8vo, cloth, 412 pages. . . net, $4 . 50
WABNER, R. Ventilation in Mines. Translated from
the. German by Charles Salter. With plates and engravings.
8vo, cloth, illustrated, 240 pages net, $4 . 50
WADE, E. J. Secondary Batteries: their Theory, Con-
struction and Use. With innumerable diagrams and figures.
8vo, cloth, illustrated, 492 pages net, $4 .00
WALKER, F., C.E. Aerial Navigation. A Practical
Handbook on the Construction of Dirigible Balloons, Aerostats,
Aeroplanes and Aeromotore. With diagrams, tables and illus-
trations. 8vo, cloth, illustrated, 151 pages net, $3.00
WALKER, W. H. Screw Propulsion. Notes on Screw
Propulsion ; its Rise and History. 8vo, cloth $0 . 75
WALKER, S. F. Electrical Engineering in Our Homes
and Workshops. A Practical Treatise on Auxiliary Electrical
Apparatus. Third Edition, revised, with numerous illustrations.
8vo, cloth $2.00
SCIENTIFIC PUBLICATIONS. 55
WALKER, S. F. Electric Lighting for Marine Engineers,
or How to Light a Ship by the Electric Light and How to Keep the
Apparatus in Order. 103 illus., 8vo, cloth Second Edition . . $2 . 00
WALLING, B. T., Lieut. Com. U.S.N., and MARTIN, JULIUS.
Electrical Installations of the United States Navy. With many
diagrams and engravings. 8vo, cloth, illustrated In Press.
WALLIS TAYLER, A. J. Modern Cycles, a Practical
Handbook on Their Construction and Repair. With 300 illustra-
tions. 8vo, cloth $4 . 00
Motor Cars, or Power Carriages for Common Roads.
With numerous illustrations. 8vo, cloth SI .80
Bearings and Lubrication. A Handbook for Every
user of Machinery. Fully illustrated. 8vo, cloth $1 . 50
Refrigerating and Ice-making Machinery. A Descrip-
tive Treatise for the use of persons employing refrigerating and
ice-making installations, and others. 8vo, cloth, illustrated. $3 . 00
Refrigeration and Cold Storage: being a Complete
practical treatise on the art and science of refrigeration. 600
pages, 361 diagrams and figures. 8vo, cloth net, $4 . 50
Sugar Machinery. A Descriptive Treatise, devoted
to the Machinery and Apparatus used in the Manufacture of
Cane and Beet Sugars. 12mo, cloth, illustrated $2 . 00
WANKLYN, J. A. A Practical Treatise on the Exam-
ination of Milk and its Derivatives, Cream, Butter and Cheese.
12mo, cloth $1 . 00
Water Analysis. A Practical Treatise on the Exam-
ination of Potable Water. Tenth Edition. 12mo, cloth .... $2 . 00
WANSBROUGH, W. D. The A B C of the Differential
Calculus. 12mo, cloth $1 . 50
WARD, J. H. Steam for the Million. A Popular Treat-
ise on Steam, and its application to the Useful Arts, especially to
Navigation. 8vo, cloth. . . $1 . 00
WARING, G. E., Jr. Sewerage and Land Drainage.
Illustrated with woodcuts in the text, and full-page and folding
plates. Quarto. Cloth. Third Edition $6 . 0,0
66 D. VAN NOSTRAND COMPANY'S
WARING, G. E., Jr. Modern Methods of Sewage Disposals
for Towns, Public Institutions and Isolated Houses. Second
Edition, revised and enlarged. 260 pages. Illus. Cloth. . . $2 . 00
-How to Drain a House. Practical Information for
Householders. Third Edition, enlarged. 12mo, cloth $1 .25
WATSON, E. P. Small Engines and Boilers. A Manua'
of Concise and Specific Directions for the Construction of Small
Steam-engines and Boilers of Modern Types from five Horse-
power down to model sizes. Illustrated with Numerous Dia-
grams and Half-tone Cuts. 12mo, cloth $1 .25
WATT, A. Electro-plating and Electro-refining of Metals :
being a new edition of Alexander Watts' "Electro-Deposition."
Revised and largely rewritten by Arnold Philip, B.Sc. With
numerous figures and engravings. 8vo, cloth, illustrated, 680
pages net, $4 . 50
Electro-metallurgy Practically Treated. Eleventh
Edition, considerably enlarged. 12mo, cloth $ 1 . 00
- The Art of Soap-making. A Practical Handbook of
the Manufacture of Hard and Soft Soaps, Toilet Soaps, etc. In-
cluding many New Processes, and a Chapter on the Recovery of
Glycerine from Waste Lyes. With illustrations. Fifth Edition,
revised and enlarged. 8vo, cloth $3 . 00
The Art of Leather Manufacture : being a Practical
Handbook, in which the Operations of Tanning, Currying and
Leather Dressing are Fully Described, and the Principles of
Tanning Explained, and many Recent Processes Introduced.
With numerous illustrations. Second Edition. 8vo, cloth. . $4 . 00
WEALE, J. A Dictionary of Terms Used in Architecture,
Building, Engineering, Mining, Metullargy, Archeology, the Fine
Arts, etc., with explanatory observations connected with applied
Science and Art. Fifth Edition, revised and corrected. 12mo,
cloth $2.50
WEBB, H. L. A Practical Guide to the Testing of Insu-
lated Wires and Cables. Illustrated. 12mo, cloth $1 . 00
- The Telephone Handbook. 128 Illustrations. 146
pages. 16mo, cloth f 1 . 00
WEEKES, R. W. The Design of Alternate Current Trans-
formers. Illustrated. 12mo, cloth $1 .00
SCIENTIFIC PUBLICATIONS. 57
WEISBACH, J. A Manual of Theoretical Mechanics.
Ninth American edition. Translated from the fourth augmented
and improved German edition, with an Introduction to the Calculus
by Eckley B. Coxe, A.M., Mining Engineer. 1,100 pages and 902
woodcut illustrations. 8vo, cloth $6.00
Sheep $7.50
Mechanics of Air Machinery. Translated with an
appendix on American Practice by Prof. A. Trowbridge, Columbia
University In Press.
WESTON, E. B. Tables Showing Loss of Head Due to
Friction of Water in Pipes. Second Edition. 12mo, cloth. . $1 . 50
WEYMOUTH, F. M. Drum Armatures and Commutators.
(Theory and Practice.) A complete Treatise on the Theory
and Construction of Drum Winding, and of commutators for
closed-coil armatures, together with a full r6sum6 of some of the
principal points involved in their design, and an exposition of
armature reactions and sparking. 8vo, cloth $3 . 00
WHEELER, J. B., Prof. Art of War. A Course of
Instruction in the Elements of the Art and Science of War, for
the Use of the Cadets of the United States Military Academy,
West Point, N. Y. 12mo, cloth $1.75
Field Fortifications. The Elements of Field Forti-
fications, for the Use of the Cadets of the United States Military
Academy, West Point, N. Y. 12mo, cloth $1 . 75
WHIPPLE, S., C.E. An Elementary and Practical Treatise
on Bridge Building. 8vo, cloth $3 .00
WHITE, W. H., K.C.B. A Manual of Naval Architecture,
for use of Officers of the Royal Navy, Officers of the Mercantile
Marine, Yachtsmen, Shipowners and Shipbuilders. Containing
many figures, diagrams and tables. Thick, 8vo, cloth, illus-
trated $9.00
WHITELAW, J., Jr. Surveying, as Practiced by Civil
.Engineers and Surveyors; including the setting-out of work for
construction and surveys abroad, with examples taken from
actual practice. Intended as a handbook for Field and Office
use; also as a text-book for Students. With numerous tables,
full-page plates and diagrams. 8vo, cloth, illustrated, 516
pages net, $4 . 00
WILKINSON, H. D. Submarine Cable-laying, Repairing,
and Testing. 8vo, cloth $5.00
68 D. VAN NOSTRAND COMPANY'S
WILLIAMSON, R. S. On the Use of the Barometer on
Surveys and Reconnoissances. Part I. Meteorology in its Con-
nection with Hypsometry. Part II. Barometric Hypsometry.
With illustrative tables and engravings. 4to, cloth.... $15.00
-Practical Tables in Meteorology and Hypsometry, in
connection with the use of the Barometer. 4to, cloth $2 . 60
WILSON, G. Inorganic Chemistry, with New Notation.
Revised and enlarged by H. G. Madan. New edition. 12mo,
cloth ". $2.00
WILLSON, F. N. Theoretical and Practical Graphics.
An Educational Course on the Theory and Practical Applications
of Descriptive Geometry and Mechanical Drawing. Prepared
for students in General Science, Engraving, or Architecture.
Third Edition, revised. 4to, cloth, illustrated net, $4 . 00
Note-taking, Dimensioning and Lettering. 4to, Cloth,
illustrated net, $1 . 25
- Third Angle Method of Making Working Drawings.
4to, cloth, illustrated net, $1 . 25
Some Mathematical Curves, and Their Graphical
Construction. 4to, cloth, illustrated net, $1 .50
— Practical Engineering Drawing and Third Angle
Projection. 4to, cloth, illustrated net, $2.80
Shades, Shadows, and Linear Perspective. 4to, Cloth.
illustrated net, $1 . 00
Descriptive Geometry — Pure and Applied, with a
chapter on Higher Plane Curves, and the Helix. 4to, cloth
illustrated net, $3 . 00
WTNKLER, C., and LUNGE, G. Handbook of Technical
Gas-Analysis. With figures and diagrams. Second English edi-
tion. Translated from the third greatly enlarged German edition,
with some additions by George Lunge, Ph.D. 8vo, cloth, illus-
trated, 190 pages $4.00
WOODBURY, D. V. Treatise on the Various Elements
of Stability in the Well -proportioned Arch. With numerous
tables of the Ultimate and Actual Thrust. 8vo, half morocco.
Illustrated $4 .00
SCIENTIFIC PUBLICATIONS. 59
WRIGHT, A. C. Analysis of Oils and Allied Substances.
8vo, cloth, illustrated, 241 pages net, $3 . 50
Simple Method for Testing Painters' Materials. 8vo,
cloth, 160 pages : net, $2.50
WRIGHT, T. W., Prof. (Union College.) Elements o'
Mechanics; including Kinematics, Kinetics and Statics. With ap-
plications. Third Edition, revised and enlarged. 8vo, cloth. . $2 . 50
WYNKOOP, R. Vessels and Voyages, as Regulated by
Federal Statutes and Treasury Instructions and Decisions. 8vo,
cloth $2.00
I
YOUNG, J. E. Electrical Testing for Telegraph Engineers.
With Appendices consisting of Tables. 8vo, cloth, illus ... $4 . 00
YOUNG SEAMAN'S MANUAL. Compiled from Various
Authorities, and Illustrated with Numerous Original and Select
Designs, for the Use of the United States Training Ships and the
Marine Schools. 8vo, half roan $3 .00
ZIPSER, J. Textile Raw Materials, and Their Conversion
into Yarns. The study of the Raw Materials and the Technology
of the Spinning Process. A Text-book for Textile, Trade and
higher Technical Schools, as also for self-instruction. Based upon
the ordinary syllabus and curriculum of the Imperial and Royal
Weaving Schools. Translated from the German by Chas. Salter.
8vo, cloth, illustrated net, $5.00
Catalogue of the Van Nostrand
Science Series.
are put up in a uniform, neat, and attractive form. i8mo,
boards. Price 50 cents per volume. The subjects are of an
eminently scientific character and embrace a wide range of topics, ana
are amply illustrated when the subject demands.
No. i. CHIMNEYS FOR FURNACES AND STEAM BOILERS. By
R. Armstrong, C.E. Third American Edition. Revised and
Partly rewritten, with an Appendix on "Theory of Chimney
Jraught," by F. E. Idell. M.E.
No. 2. STEAM-BOILER EXPLOSIONS. By Zerah Colburn. New
Edition, revised by Prof. R. H. Thureton.
No. 3. PRACTICAL DESIGNING OF RETAINING-WALLS. By
Arthur Jacob, B.A. Fourfrh edition, with additions by Prof.
W. Cain.
No. 4. PROPORTIONS OF PINS USED IN BRIDGES. By Charles
E. Bender, C.E. Second edition, with Appendix.
No. 5. VENTILATION OF BUILDINGS. By W. F. Butler. Second
edition, re-edited and enlarged by James A. Greenleaf, C.E.
New edition in press.
No. 6. ON THE DESIGNING AND CONSTRUCTION OF STORAGE
Reservoirs. By Arthur Jacob, B.A. Second American edition,
revised, with additions by E. Sherman Gould.
No. 7. SURCHARGED AND DIFFERENT FORMS OF RETAINING-
walls. By James S. Tate, C.E.
No. 8. A TREATISE ON THE COMPOUND STEAM-ENGINE. By
John Turnbull, Jr. 2nd edition, revised by Prof. S. W. Robinson.
No. 9. A TREATISE ON FUEL. By Arthur V. Abbott, C.E. Founded
on the original treatise of C. William Siemens, D.C.L.
No. 10. COMPOUND ENGINES. Translated from the French of A.
Mallet. Second edition, revised with results of American Prac-
tice, by Richard H. Buel, C.E.
No. it. THEORY OF ARCHES. By Prof. W. Allan.
No. 12. THEORY OF VOUSSOIR ARCHES. By Prof. Wm. Cain.
Second edition, revised and enlarged.
60
BINDING US I JUW 1
>.
CTJ • 0)
LO O JH
IO •
CO O •>
CD CO
fcO •» •
O S
to
a •
C rt
cd C
o3
.Sn
6 .ctj
rf O
•r-j -H
G H
CD 0
OQ r2
O
-P
fn
£
>
^ -H
H ao
<
-P
^ «H
O
o
0)
fn
P-,
g
O
CV2
O5
O
University of Toronto
Library
DO NOT
REMOVE
THE
CARD
FROM
THIS
POCKET
Acme Library Card Pocket
LOWE-MARTIN CO. LIMITED
Live Music Archive
Librivox Free Audio
Metropolitan Museum
Cleveland Museum of Art
Internet Arcade
Console Living Room
Books to Borrow
Open Library
TV News
Understanding 9/11