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A System of Chemical Analysis

(QUALITATIVE AND SEMI-QUANTITATIVE)
for the

Common Elements

by
ERNEST H. SWIFT, PH.D.

PROFESSOR OF ANALYTICAL CHEMISTRY
CALIFORNIA INSTITUTE OF TECHNOLOGY

NEW YORK

PRENTICE-HALL, INC.

PRENTICE-HALL CHEMISTRY SERIES
WENDELL M. LATIMBR, PH.D., Editor

70 FIFTH AVENUE, NEW YORK

REPRODUCED IN ANY FORM, BY MIMEOGRAPH OR ANY

OTHER MEANS, WITHOUT PERMISSION IN WRITING

FROM THE PUBLISHERS.

First Printing. . January 1939
Second Printing. . .October 1940
Third Printing. . .January 1946
Fourth Brinting ... March 1949
Fifth Dinting... August_19 50

PRINTED IN THE UNITED STATES OF AMERICA

ARTHUR AMOS NOYES

Preface

The system of analysis presented in this book is the result of an
experimental investigation by the author and numerous associates
which has been in progress for a period of over ten years, and which
has had for its purpose the development of an analytical system
which would provide not only reliable and sensitive qualitative
information in regard to the constituents present in the material
being analyzed, but also, without an undue expenditure of time or
labor, sufficient quantitative information to eliminate in many
cases the need for further quantitative determinations.

The practical value of the conventional system of qualitative
analysis has been materially reduced by the advances which have
been made in the technique of spectrographic methods, by the
developments in the field of microchemistry, and by the increase
in the number and use of specific organic reagents for the purely
qualitative detection of the elements; especially is this true of those
qualitative systems which provide relatively little quantitative
information. And, although there has also been a corresponding
activity in the development and study of highly precise methods
for the quantitative determination of single constituents, practical
experience has indicated that there is very frequently a demand for
both qualitative and quantitative information regarding a given
material or substance, but that the purpose for which the quantita-
tive information is needed requires only approximate values and does
not seem to justify the time and labor involved in carrying out the
usual series of highly precise quantitative determinations.

This system of analysis is designed to meet this demand and
consists, first, of a system of qualitative analysis, largely conventional
in its methods, in which the' separations of both groups and single
elements have been studied and, where necessary, modified or
changed until each constituent is detected and isolated without
undue loss, but in which an effort has been made not to increase
very materially the time required for the analysis over that necessary
for properly carrying out a conventional qualitative system. In
certain cases, where the procedures have of necessity been time-
consuming or laborious in order that more quantitative separations
might be obtained, simpler and more rapid, but less exact, optional
procedures have been provided which can be used at the judgment

vii

viii PREFACE

of the analyst. Distribution and distillation methods have been
used for certain of the separations, not only because they have
proved effective, but also because it seems that the traditional
adherence to precipitation methods has prevented methods based
upon other principles from being developed to their fullest extent.
Sulfide separations are used frequently because in a surprising
number of cases they were found to be much superior to alternative
methods; the objectionable features of hydrogen sulfide as a labora-
tory reagent did not seem adequate justification for the use of less-
effective analytical procedures.

A system for the separation of the various constituents having
been developed, methods for the estimation of the amount of each
of the constituents thus separated are provided. These methods,
largely volumetric, have been selected primarily on the basis of
rapidity and simplicity of technique in order that the estimation
shall require a minimum of time. However, many of these methods
are conventional quantitative determinations for the particular
constituents and can be used for more accurate analyses if desired.

Considerable experimental data are presented in order to show
both the effectiveness and the limitations of the various separations
and estimations which have been studied; such data also serve to
give the student a better perspective of analytical methods, to
develop a more critical attitude on his part, and to stimulate his
interest in carrying out similar studies. It is realized that the
practical value of such a system as is described above is materially
decreased by being restricted to the so-called "common elements";
however, adequate experimental data have not been collected at the
present time to enable even the more important of the "rarer ele-
ments" to be included.

The material in this book has been used for several years in
mimeographed form as the text for the course in analytical chemistry
given in the sophomore year at the California Institute of Tech-
nology. In addition, as various portions of the system have been
developed, they have been extensively tested by selected groups of
students in this course. From the instructional point of view, the
author has found that the correlation of a more exact system of
qualitative analysis with the quantitative course presents certain
advantages. The qualitative analysis which can be given in the
freshman course in general chemistry is seldom adequate to provide
the student with either a sufficient mastery of the technique of the
subject to give him confidence in his experimental ability, or with

PREFACE ix

that valuable foundation of systematized inorganic chemistry which
is so uniquely given by an intensive study of a qualitative system.
This lack of a background of inorganic chemistry is likely to be
emphasized at the present time because the trend, however desirable
it may be, toward an extensive presentation of physio-chemical
principles in the quantitative course necessarily limits the range of
factual chemistry with which the student is made familiar. On
the other hand, courses in qualitative analysis in the upper years
are likely to be regarded as a continuation of the freshman subject
and may fail to sustain the interest of the abler students.

In the system presented here, training in the precise measurements
of volumetric and gravimetric analysis is provided in the preliminary
preparation and standardization of the solutions required for the
volumetric determinations. In addition, many of the single deter-
minations which are commonly included in courses in quantitative
analysis can be carried out by slight modifications of the systematic
procedures; these modifications will be obvious or are suggested in
notes to the procedures. Even when such single determinations
are being carried out, it has been found of value for the student to
have constantly before him the qualitative separations of that con-
stituent and thus to acquire more familiarity with the treatment
of complex substances and with the steps necessary for the elimina-
tion of interfering constituents.

Many of the standardization procedures and certain of the pro-
cedures for the estimation of a constituent have been discussed in
considerable historical and experimental detail. This has been done
in order that the student may appreciate that such methods are
often the result of constant study over considerable periods of time,
and that he may obtain a better appreciation of the numerous factors
which may have to be considered in the experimental study of a
method. In general, it has been found that the intensive study of a
few well-chosen methods is of more value than a more diffused review
of a wider field. An outline of a suggested course for one year is
given in the Appendix.

It is believed that the advances which have been made in the
content and presentation of the first-year college course in general
chemistry is justification for the assumption that the student has
acquired some familiarity with the laws of definite proportions and
combining weights and even with the simple principles of the mass-
action law. The author feels that the teachers of general-chemistry
courses have a just cause for resentment in the assumption implicit

x PREFACE

in many- of the texts on analytical chemistry that so little of this
material has been absorbed from the course in general chemistry.
However, the author is not in sympathy with the concept that the
courses in analytical chemistry, especially those in qualitative anal-
ysis, should serve primarily as an excuse for an exposition of the
entire field of modern chemical theory. The physical principles
underlying the procedures have been presented in this book only as
and where they have been needed to explain the experimental ob-
servations and facts at hand. It is believed that there is an ad-
vantage in this close correlation between experiment and theory,
even though the presentation of the theoretical material cannot
be made as systematically as when it is collected in a separate section
of the book.

Although a large amount of experimental work has been expended
in the development and testing of this system of analysis, the author
is aware that imperfections and limitations from both the technical
and instructional points of view still exist. Because of this, any
suggestions for improvement or notice of errors of procedure or text
will be received with genuine appreciation.

ERNEST H. SWIFT

Acknowledgments

The author wishes, first, to acknowledge his great indebtedness to
Arthur A. Noyes, late Professor of Chemistry at the California
Institute of Technology, to whom this book is dedicated. He not
only gave freely of advice and assistance during a large part of the
experimental development but also was a constant source of inspira-
tion and encouragement. With his consent and cooperation, many
of the methods used by Professor Noyes in his Qualitative Chemical
Analysis were made the basis for the study of the more quantitative
separations required in this system of analysis and frequently have
been adopted with little modification. In addition, the author
was so fortunate as to have been among those associated with
Professor Noyes and with Professor W. C. Bray in the experimental
development of their System of Qualitative Analysis for the Rare
Elements; any merit which this book may possess can reflect in only
a small measure the value of that association and experience.

It would have been impossible to have carried out the experi-
mental work involved in the development of this system of analysis
without the aid which has been given by a large number of assistants
and students. An attempt has been made to acknowledge this
assistance specifically in the discussion of the various procedures
and methods. The author is also grateful to those of his classes
who have used this material in preliminary mimeographed form for
their spirit of cooperation and for many valuable suggestions. In
addition, certain sections of this book have been so influenced by the
advice, experience, or experimental work of certain associates that
it is desired to mention these more general contributions here.
These sections and those contributing to them are as follows : Prepara-
tion of the sample and solution, Mr. J. B. Hatcher; precipitation
and separation of the Copper and Tin Groups, Mr. R. C. Aussieker;
analysis of the Tin Group, Dr. Chester Wilson and Dr. R. C. Barton;
analysis of the Ammonium Sulfide Group, Dr. R. C. Barton; analysis
of the Alkaline Earth and Alkali Groups, Dr. Carter Gregory; and
the analysis for the acidic constituents, Mr. Theodore Vermeulen,
Mr. R. W. Dodson, Mr. D. K. Beavon, and Mr. R. N. Wimpress.

Dr. Charles D. Coryell, Dr. Clifford S. Garner, Mr. R. W. Dodson,
and Mr. J. B. Hatcher have each read a large portion or all of the
manuscript and have made many valuable contributions.

xii ACKNOWLEDGMENTS

The index has been prepared by Mr. J. B. Hatcher, who has also
rendered very valuable assistance in preparing the manuscript for
the printer and in reading the proof.

The author wishes to express his appreciation to Professor T. H.
Morgan for his kindness in having made available laboratory space
in the Kerckhoff Marine Biological Laboratory of the California
Institute of Technology at Corona del Mar, California. Much
of the experimental work has been carried on there in the summer
periods under exceptionally favorable surroundings; the friendly
cooperation of Professor G. E. MacGinitie and the members of
that laboratory has contributed to the pleasant associations of those
periods.

ERNEST H. SWIFT

Contents

PART I
THE PREPARATION OF STANDARD SOLUTIONS

PROCEDURE PAGE

THE ANALYTICAL BALANCE 3

Principles and Construction of the Analytical Balance 3

The Use of the Analytical Balance 8

The Determination of the Point of Rest 8

I. Method of Long Swings 10

II. Method of Short Swings 12

III. The Determination of the Sensitivity of a Balance. . 13

Methods of Weighing 14

The Correction for Buoyancy 16

IV. The Calibration of a Set of Weights 16

General Rules for tl\e Use of the Balance 21

VOLUMETRIC METHODS OF ANALYSIS 22

General Discussion 22

Precipitation Methods of Volumetric Analysis 26

V. The Preparation of a Standard Silver Nitrate

Solution 26

VI. The Preparation and Standardization of a Thio-

cyanate Solution 32

The Applications of Precipitation Methods 46

Oxidation and Reduction Methods of Volumetric

Analysis 47

Discussion 47

Permanganate Methods of Volumetric Analysis ... 53

VII. The Preparation of a Permanganate Solution. . 59

VIII. The Standardization of Permanganate Solutions 60
IX. The Preparation and Standardization of Ferrous

Sulfate Solutions 66

The Applications of Permanganate Methods. . . 67

lodometric Methods of Volumetric Analysis 68

Discussion 68

xiii

xiv CONTENTS

PBOCIDURE PAOS

X. The Preparation of Iodine Solutions 69

XL The Preparation of Starch Solutions 70

XII. The Standardization of Iodine Solutions 71

XIII. The Preparation of Thiosulfate Solutions 77

XIV. The Standardization of Thiosulfate Solutions .... 80
The Applications df lodometric Methods 84

The Use of Standard Solutions of Other Oxidizing

Agents in Volumetric Analysis 84

The Use of Oxidation-Reduction (Potential) Indi-
cators in Volumetric Analysis . 85

Neutralization (and Displacement) Methods of Vol-
umetric Analysis 89

General Discussion 89

The Preparation of Standard Solutions of Acids

and Bases 97

XV. The Preparation of Carbonate-Free Solutions of

Sodium Hydroxide 99,

XVI. The Standardization of Sodium Hydroxide Solu-
tions 100

XVII. The Preparation and Standardization of Hydro-
chloric Acid Solutions 103

The Applications of Neutralization and Displace-
ment Methods 109

GRAVIMETRIC METHODS OP ANALYSIS 110

General Principles of Gravimetric Methods 110

I. Factors Affecting the Solubility of the Pre-

cipitate 110

II. (1) Factors Affecting the Physical Character-

istics of the Precipitate 116

(2) Factors Affecting the Composition and
Purity of the Precipitate 119

III. Factors Affecting the Composition and Sta-
bility of the Weighed Precipitate 128

The Operations of Gravimetric Analysis 132

The Filtering of Precipitates 132

The Washing of Precipitates 137

XVIII. The Gravimetric Standardization of a Hydrochloric

Acid Solution. 139

CONTENTS xv

PART II

THE SYSTEM OF ANALYSIS FOR THE BASIC
CONSTITUENTS

PROCEDURE . PAGE

THE PREPARATION OF THE SAMPLE, PRELIMINARY OBSER-
VATIONS, AND PREPARATION OP THE SOLUTION FOR

THE ANALYSIS 153

1. Preparation of the Sample 153

2. Treatment of Solutions or Suspensions 157

3. Preliminary Observations and Tests 158

4. The Elimination of Organic Material . . 166

Preparation of the Solution for the Analysis 169

General Discussion of the Solution of Solid Substances 169

5. Treatment of the Sample with Nitric Acid 171

6. Treatment of the Sample with (1) Hydrochloric and (2)

Hydrochloric and Nitric Acids 178

7. (1) Treatment of the Sample with Perchloric Acid; (2)

Detection and Estimation of Silica by the Use of

Hydrofluoric Acid . 182

8. Fusion of the Sample with Sodium Carbonate 189

SEPARATION OF. THE BASIC CONSTITUENTS INTO GROUPS . 103

General Discussion of the Group Separations . . . 193

PRECIPITATION OF THE HYDROGEN SULFIDE GROUP AND

SEPARATION OF THE COPPER AND TIN GROUPS 197

11. Precipitation of the Hydrogen Sulfide Group 197

12. Separation of the Copper Group from the Tin Gr.oup. . 211

13. Precipitation of the Tin Group 216

THE ANALYSIS OF THE COPPER GROUP 221

21. Solution of the Copper Group Sulfides in Nitric Acid. . 221

22. Detection and Precipitation of Silver 223

23. Estimation of Silver 226

24. Precipitation of Lead 229

25. Estimation of Lead 232

26. Precipitation of Bismuth and Detection of Copper. . . . 236

27. Estimation of Bismuth 238

28. Precipitation and Estimation of Copper 242

29. Detection of Cadmium 247

30. Estimation of Cadmium 249

xvi CONTENTS

PAOB

ANALYSIS OF THE TIN GROUP ........................ 251

General Discussion of Methods for the Analysis of the

Tin Group .................................... 251

41 A. Optional Method for the Analysis of the Tin Group. . 254

41. Solution of the Tin Group Sulfides ................... 256

42. Separation and Detection of Arsenic ................. 258

43. Estimation of Arsenic .............................. 263

44. Precipitation of Mercury ........................... 264

45. Estimation of Mercury ............................. 267

46. Precipitation of Antimony .......................... 269

47. Estimation of Antimony ............................ 270

48. Detection and Precipitation of Tin .................. 274

49. Estimation of Tin ................................. 275

PRECIPITATION OP THE AMMONIUM SULPIDE GROUP; SEPA-
RATION OF IRON; REMOVAL OF PHOSPHATE; AND
SEPARATION OF THE ALUMINUM AND ZINC GROUPS 281

51. Precipitation of the Ammonium Sulfide Group ....... 281

52. Detection and Separation of Iron .................... 291

53. lodometric Estimation of Iron ...................... 298

53.4. Optional Method for the Estimation of Iron by Titra-

tion with Permanganate ........................ 300

54. Removal of Phosphate ............................ 305

55. Separation of the Zinc and Aluminum Groups ........ 309

THE ANALYSIS OF THE ZINC GROUP .................... 321

61. The Precipitation and Separation of Zinc as Sulfide. . . . 321

62. Estimation of Zinc ................................. 331

63. Precipitation of Nickel and Cobalt (and Iron), and the

Separation of Iron from Nickel and Cobalt ...... 332

64. Separation of Nickel and Cobalt ..................... 336

644. Optional Method of Detecting Nickel and Cobalt ...... 340

65. Estimation of Nickel ............................... 342

66. Estimation of Cobalt .............................. 344

THE ANALYSIS OF THE ALUMINUM GROUP ............. 347

71. Destruction of Oxalate, Precipitation of Manganese,

and Detection of Chromium .................... 347

72 Estimation of Manganese ........................... 351

73. Precipitation of Aluminum ......................... 354

74. Estimation of Aluminum ........................... 356

75. Estimation of Chromium ........................... 358

CONTENTS xvii

PROCEDURE PAGE

THE ANALYSIS OF THE ALKALINE EARTH GROUP 361

81. Precipitation of the Alkaline Earth Group 361

The Separation of the Alkaline Earth Group Elements. 366

82. Precipitation of Barium 368

83. Estimation of Barium 370

84. Precipitation of Strontium 372

85. Estimation of Strontium 373

86. Precipitation of Calcium 375

87. Estimation of Calcium 378

88. Precipitation of Magnesium 380

89. Estimation of Magnesium 381

THE ANALYSIS OF THE ALKALI GROUP 385

91. The Detection of the Alkali Group and . Estimation of

the Total Amount of Sodium and Potassium
Present 385

92. Precipitation of Potassium 389

93. Estimation of Potassium 393

94. Precipitation of Sodium 394

95. Estimation of Sodium 395

96. Detection and Estimation of Ammonia 401

PART III

THE SYSTEM OF ANALYSIS FOR THE ACIDIC
CONSTITUENTS

THE ANALYSIS OF THE ACIDIC CONSTITUENTS 409

General Discussion 409

111. Preparation of a Solution for the Analysis of the Acidic

Constituents 413

Preliminary Tests for Certain Groups of the Acidic Con-
stituents 415

112. Detection of Oxidizing Anions 415

113. Detection of Reducing Anions 417

114. Detection of Constituents of the Cyanide, Halide, and

Oxy-Halogen Groups 418

115. Detection of Sulfate, Sulfite, Oxalate, Fluoride, or

Chromate 420

THE ANALYSIS FOR SULFIDE, CYANIDE, AND SULFITE .... 422

121. Volatilization and Detection of Sulfide and Cyanide. . 422

122. Separation and Estimation of Sulfide 426

123. Detection and Estimation of Cyanide 427

124. Separation and Estimation of Sulfite 429

xviii CONTENTS

VBOCKDURK PAGE

. THE ANALYSIS FOR FERROCYANIDE AND FERRIC Y AN IDE . . 433

131. Precipitation of Ferrocyanide and Ferricyanide 433

132. Separation and Detection of Ferrocyanide and Ferri-

cyartide 435

133. Estimation of Ferrocyanide 438

134. Estimation of Ferricyanide 440

THE ANALYSIS OF THE HALIDE GROUP 443

141. Precipitation of the Halide Group 443

142. Metathesis of the Silver Halide Group Precipitate by

. Sulfide 446

143. Detection and Separation of Iodide and Detection of

Thiocyanate ' 448

144. Estimation of Iodide 452

145. Elimination of Thiocyanate (and Iodide) 452

146. Detection and Separation of Bromide 455

147. Estimation of Bromide 457

148. Detection and Estimation of Chloride 457

THE ANALYSIS OF THE OXY-HALOGEN GROUP 459

151. Precipitation of the Oxy-Halogen Group. . . 459

152. Separation and Estimation of Hypochlorite 461

153. Detection and Estimation of Porchlorate 464

154. Detection of Periodate 466

THE ANALYSIS OF THE PHOSPHATE GROUP 469

161. Precipitation of the Phosphate Group 469

162. Reparation of Phosphate and Arsenate from Arsenite and

Oxalate 471

163. Detection and Estimation of Arsenate 473

164. Detection of Phosphate 474

165. Detection of Arsenite 475

1 66. Precipitation of Oxalate 476

THE DETECTION AND ESTIMATION OF SULFATE AND OF

FLUORIDE 478

171. Precipitation of Sulfate 478

172. Precipitation of Fluoride 480

THE ANALYSIS OF THE SODIUM CARBONATE SOLUTION

FOR NITRATE, NITRITE, BORATE, AND ACETATE 483

181. Detection and Estimation of Nitrate and Nitrite 483

182. Detection and Estimation of Nitrite 486

183. Detection of Borate 488

184. Detection of Acetate 490

CONTENTS xix

PROCEDURE PAGE

THE ANALYSIS OF THE RESIDUE FROM THE SODIUM CARBON-
ATE TREATMENT 493

General Discussion 493

191. Detection of Sulfide and Cyanide in the Sodium Carbon-

ate Residue 493

192. Detection of the Halides in the Sodium Carbonate

Residue 495

193. Detection. of Borate in the Sodium Carbonate Residue. 497

THE ANALYSIS OF THE ORIGINAL MATERIAL FOR

CARBONATE 499

201. Detection and Estimation of Carbonate 499

APPENDIX

A SUGGESTED COURSE OF INSTRUCTION 505

QUESTIONS AND PROBLEMS 509

The Analytical Balance 509

Accuracy, Precision, Errors, and Significant Figures. . . . 512

Volumetric Methods of Analysis 515

Gravimetric Methods 530

Sulfide Precipitations and Separations 533

Hydroxide Precipitations and Separations 535

Miscellaneous Problems 536

TABLES 539

I. The Solubility-Product Values of Certain Slightly

Soluble Salts 539

II. Molal and Formal Reduction Potentials 540

III. The Approximate Precipitation pR and Solubility

Products of the Oxides and Hydroxides of
Certain Elements 544

IV. lonization Constants of Acids and Bases 545

V. Dissociation Constants of Certain Complex Ions. 546

VI. International Atomic Weights 547

REAGENTS AND CHEMICALS 548

Solutions 548

Solids 552

Test Solutions 552

EQUIPMENT 559

INDEX 563

Tabular Outlines

OUTLINE PAOK

I. Preparation of the Sample, Preliminary Observations,

and Preparation of the Solution 152

II. Separation of the Basic Constituents into Groups. . . . 194

III. Precipitation of the Hydrogen Sulfide Group and

Separation of the Copper and Tin Groups 196

IV. The Analysis of the Copper Group 220

V. The Analysis of the Tin Group 252

V-A. Optional Method for the Analysis of the Tin Group . . 254
VI. Precipitation of the Ammonium Sulfide Group; Separa-
tion of Iron; Removal of Phosphate; and Separation

of the Aluminum and Zinc Groups 280

VII. The Analysis of the Zinc Group 320

VIII. The Analysis of the Aluminum Group 346

IX. The Analysis of the Alkaline Earth Group 360

X. The Analysis of the Alkali Group 386

XL Separation of the Acidic Constituents into Groups. . . . 410
XII. Preliminary Tests for Certain Groups of the Acidic

Constituents 415

XIII. The Analysis for Sulfide, Cyanide, and Sulfite 421

XIV. The Analysis for Ferrocyanide and Ferricyanide 432

XV. The Analysis of the Halide Group 442

XVI. The Analysis of the Phosphate Group 468

XVII. The Analysis of the Sodium Carbonate Solution for

Nitrate, Nitrite, Borate, and Acetate 484

XVIII. The Analysis of the Residue from the Sodium Carbonate

Treatment 492

XXI

PART I
THE PREPARATION OF STANDARD SOLUTIONS

The Analytical Balance

PRINCIPLES AND CONSTRUCTION OF THE ANALYTICAL

BALANCE

Discussion. The analytical balance is an instrument for com-
paring masses, and "weighing" is essentially the operation of making
a comparison of the mass of an unknown object with the known mass
of certain "weights." 1

The balance consists in principle of a horizontal lever, known as
the beam, supported at its center on an agate knife edge (the central
knife edge) and carrying at each end agate knife edges (the terminal
knife edges) which support the pans on which the objects and weights
are placed. As the pans which carry the object and weights are
flexibly supported on the terminal knife edges, it follows that apply-
ing a load to the pan is the same as concentrating this load at the
terminal knife edge. Therefore, when an object is placed on a
pan, a rotational moment, OL, is produced (where L is the distance
from the point of support, or central knife edge, to the terminal
knife edge supporting the applied load). When the object is bal-
anced by applying weights to the opposite pan, an equal and opposing
moment, TFL 2 , will have been produced, and

Ola = WL*. (1)

1 To be exact, distinction should be made between the mass and the weight
of a body. The mass M of an object is the quantity of matter of which it is
composed, whereas the weight W is the force resulting from the attraction
of gravity between this mass and that of the earth; the relation between them
is expresse'd by the equation

W - Mg, (1)

where g is the gravitational constant.

As g may vary at different locations on the earth's surface, it is seen that
W can vary, and that it is the mass which it is desired to determine. In using
the balance, the weight of an object, W\, is usually compared with the weight,
Wi, of certain standard masses or "weights." However, at a given location

(2)
and

W* - M*g. (3)

It therefore follows that the weight is always directly proportional to the
mass and that when W\ = W*, then MI = M* and a comparison of masses is
made. It has become so customary to designate the process of making this
comparison as "weighing" and to speak of the mass as the weight that here-
after the term weight will be used to designate mass.

THE ANALYTICAL BALANCE

If LI and L 2 are made equal, it follows that = W. In order to
protect the knife edges from damage when the instrument is not in
use, means are provided for raising the beam and pans from the knife
edges ; this arresting mechanism is controlled by means of the milled
head labeled "beam-arrest control" shown in Fig. 1. Means are also
provided for arresting the motion of the pans, and this mechanism is

Fig. 1. The Essentials of the Analytical Balance. BA -beam-arrest
control; C column or post; B beam; P pointer; G movable weight for
adjusting sensitivity; R rider; TKE terminal knife edge, CKE central
knife edge; stirrup; AP agate plate; BS beam support; PS pan sup-
port; AS adjusting screw.

controlled by the button labeled "pan-arrest control." In some
balances these controls are combined.

It is obvious that the essential parts of the balance are the beam and
knife edges, and the design and construction of these largely deter-
mine the qualities of the balance. The principal requirements of a
good beam are rigidity and strength with minimum weight, and these

PRINCIPLES AND CONSTRUCTION

conflicting qualifications have led to much research both as to the
design and the material used in its construction. Various designs
can be noted by reference to the catalogues of balance manufacturers.
The materials most commonly used at present are aluminum alloys
and brass or phosphor-bronze alloys.

The material used for the knife edges is a special grade of agate,
used because of its hardness, durability, and resistance to corrosion.
The central and terminal knife edges bear against agate plates
set in the "post" and in the pan "stirrups," respectively.

The relative positions of the knife edges are of fundamental
importance in determining the performance of the balance. First,
and obviously, the central knife edge or point of support should be
above the center of gravity of the beam system in order for it to be in
stable equilibrium; also, as will be shown later, the distance between
these two points is an important factor in determining the "sensi-
tivity," that is, the response of the balance to a given difference in the
load on the pans. The sensitivity is conventionally measured in
terms of the deflection of the pointer produced by an excess in weight
of 1 milligram on either
pan. Referring to Fig. 2,
it is seen that the sensi-
tivity, S, can be expressed
in more general terms as

(2)

Fig. 2. Principles of the Balance.

where Ar is the change in

the point of rest upon

adding an excess of weight,

w, to either pan. Second,

the knife edges should be

parallel to each other and

perpendicular to the plane

of rotation of the beam in order to avoid excessive frictional effects

and possible variation in arm length as the beam oscillates. Third,

the knife edges should be in the same horizontal plane in order

that the sensitivity of the balance remain constant with varying

loads. That this is true can be seen from the following analysis of

the factors determining the sensitivity of the balance.

In Fig. 2 are shown the fundamentals of a balance system. Let Q
represent the weight of the beam and unloaded pan system with the

6 THE ANALYTICAL BALANCE

center of gravity in the position shown; W is the weight applied to
each pan and w is the excess of weight applied to one pan; a is the
angle of deflection made by the rotating beam; L is the arm length;
D is the distance from the point of support to the center of gravity
of the beam system, while D' is the distance from the point of support
to the center of gravity of the applied weights. This latter center of
gravity is obviously on a line joining the two terminal knife edges,
since, owing to the fact that the pans are flexibly suspended from the
terminal knife edges, the pan load can be considered as centered at the
suspending knife edge. It is apparent that as w is applied, a moment
equal to w times its horizontal distance from the point of support
will be set up which will cause displacement of the beam system until
the moment due to the weight Q times its horizontal distance from
the point of support plus that due to 2W times its horizontal distance
from the point of support becomes equal to the displacing moment.
Referring to Fig. 2, the displacing moment due to w is seen to be

wl, (3)

and the restoring moments will be

QD sin a (4)

and

2WD' sin a. (5)

At equilibrium the displacing moments must equal the restoring
moments, and therefore

wl = QD sin a + 2WD' sin a. (6)

For very small angles, which can be justifiably assumed in this case,
as the actual deflections of a balance are relatively small, I approaches
L and the value of sin a approaches that of tan a. Therefore Equa-
tion 5 can be rewritten as follows :

wL = tan a(QD + 2WD'). (7)

As previously shown, the sensitivity, S, of the balance can be ex-
pressed in terms of the applied excess weight and the deflection, Ar,
of the pointer as follows:

= '
and from Fig. 2 it is seen that

< = 4 r . w

PRINCIPLES AND CONSTRUCTION 7

where P is the length of the pointer, and therefore it follows that

tan a = -^ . (9)

Substituting this expression for tan a in Equation 7, the following
general formula for the sensitivity of the balance is obtained :

An inspection of this equation shows that if the terminal knife
edges are in the same plane with the central knife edge, then D'
becomes zero, and the equation is simplified to the form

r p

S - . (11)

This shows that the sensitivity is directly proportional to the arm
length, inversely proportional to the distance from the point of
support to the center of gravity and to the weight of the beam system,
but independent of the applied load. If the terminal knife edges are
below the central knife edge, then, as shown in Equation 10, the
sensitivity becomes less as the load (2W) is increased. If the terminal
knife edges are above the central knife edge or point of support,
Equation 10 has the form

o _.

(QD-2WD')'

and it is seen that as 2W is increased, the sensitivity increases, until
2WD' becomes equal to QD, when it becomes infinite and the balance
becomes unstable.

From these considerations it would appear that if the balance is
constructed with the terminal knife edges in a plane with the central
knife edge, the sensitivity would remain constant; however, owing to
bending of the beam with the applied load and to frictional forces,
the sensitivity would be likely to decrease. Because of this, in
manufacturing practice the terminal edges are often placed slightly
above the central one, and occasionally balances will be found in
which the sensitivity may actually increase with applied load. A
means of changing the sensitivity is provided by the adjustable
weight on the pointer, which by being shifted changes the value of D.
Increasing the sensitivity of balances by increasing the arm length,
L, is limited practically by the mechanical difficulty in building a

8 THE ANALYTICAL BALANCE [P. I

rigid beam of great length and light weight, and by the fact that the
period of oscillation of the balance increases as the arm length is
increased; this last increases the time required for making a weighing.

THE USE OF THE ANALYTICAL BALANCE
P. I. The Determination of the Point of Rest

Discussion. The first operation in the use of the balance is the
determination of the "point of rest," that is, the reading of the
pointer when the beam assembly is released and has assumed its
equilibrium position. 2 The obvious method of doing this would be
to set the beam in motion, allow it to come to rest, and then care-
fully observe the position of the pointer on its scale. Such a static
method is obviously too time-consuming for practical use, and in
addition small frictional effects might cause the beam to stop at
other than its exact equilibrium position. Because of this, dynamic
methods of determining the point of rest are used. The method
most commonly described, which may be termed the "long swing
method," consists in setting the beam in motion so that the pointer
swing covers from 5 to 10 divisions on the scale, and then making a
series of readings of the extreme points of the swings. An odd
number of readings, from 3 to 5, are taken on one side, and an even
number, from 2 to 4, are taken on the other side. The average of
each of these two sets is taken, and the point of rest is the mean
of these two values. For example, if the successive readings are a, 6,
c, d, e t f t and g (see Fig. 3), the point of rest would be

a+c+e+g b +d+f

By taking an odd number of readings on one side, the error due to the
decrease in the amplitude of successive swings (called the decrement)
caused by friction and air resistance is eliminated. This is seen
upon considering Fig. 3, where A, B, C represent the swings of
an ideal balance with no decrement, and a, 6, c the swings of a
balance with a decrement K. The value of K can be assumed to be

1 This is sometimes called the "zero point," because the pointer scales of
balances were often calibrated from zero in the center to 10 divisions on
each side; the term "zero point" is also sometimes used to designate the
equilibrium position of the balance with zero load, and the "rest point 11 the
equilibrium position of the loaded balance.

P.I]

POINT OF REST

constant for the limited number
of swings considered here, as its
value is small in comparison to the
amplitude of the swing. With
no decrement it is obvious that
the true point of rest would be the
mean of the first two observations,
thus:

P.R. =

3 '

Pointer Reading
.

The real observation, 6, is equal to

D Zr xi f /A , D TT\ /O ' ' -

B - A; therefore (A + B - K)/2 ment on the Oscillation of a Balance.
differs from the true point of rest

by K/2. However, if two swings to the left and one to the
right are taken, we have, for the ideal balance,

P.R. =

and for the real observations,

A + C + 2K
P.R. = ?

- K

in which K is eliminated and the true point of rest obtained. The
same treatment can be extended to a larger number of readings.

The method of long swings requires considerable time and com-
putation, and in order to avoid this the so-called "short swing
method" of determining the point of rest may be used with a precision
usually well within the experimental limits desired. In this method
the swing is limited to 1 or, at the most, 2 divisions, and the point
of rest is determined by visually estimating the center of the pointer
swing; for very precise work the extremes of one of the short swings
may be recorded and the mean taken. As the pointer is moving
much more slowly, the period of oscillation being constant, and as the
decrement with such a short swing is hardly appreciable, the center
can be very precisely placed. 3 - 4

8 For a discussion of the relative merits of the long and short swing methods
see Wells, /. Am. Chem. Soc., 42, 411 (1920).

4 Balances are now available which are provided with damping devices.
These serve to bring the beam system to rest within 10 to 20 seconds and thus

10 THE ANALYTICAL BALANCE [P. I

Procedure I: DETERMINATION OF THE POINT OF REST.
A. Method of Long Swings. Examine the balance to ascer-
tain that it is in proper operating condition (Notes 1, 2).
Release the pan rests by pushing in the button controlling
them; this usually may be locked in by a slight turn. If
the pans begin to swing, bring them to a stationary position
by means of the pan rests. Now lower the beam by very
slowly turning the beam-rest control in a counterclockwise
direction as far as it will go (Note 3) . By means of the rider
control momentarily lower the rider to the beam near the
end so that the pointer is caused to swing from 5 to 6
divisions to one side of the center. After one or two swings
(Note 4) begin recording the successive extreme positions
of the pointer, taking 5 such readings, three on one side and
two on the other (Notes 5, 6, 7). Arrest the beam by means
of the beam rest, taking especial care that the beam is in a
horizontal position when it is lifted; then release the pan-rest
control. From the readings so obtained calculate the point
of rest of the empty balance (Note 8). Repeat the deter-
mination (Note 9).

Notes:

1. The balance should be situated in a position which is as free as possible
from the corrosive fumes common to a cheipical laboratory, from air move-
ment, and from sudden changes in temperature, such as would be caused by
an oven in the vicinity or by direct sunlight. Adequate illumination should
be provided. The balance should be solidly mounted so as to avoid vi-
brations.

2. Before the balance is used, it should be examined to see that it is level
(a leveling instrument is usually provided in the case), that the beam and pan
arrests are operating smoothly, and that the "rider" is properly controlled
by its mechanism throughout the entire range of the beam calibrations.

effect a considerable saving of time in determining the point of rest. This
damping is accomplished by air displacement or magnetically. In the first
case cylinders which are closed at the top are suspended from the beam by the
stirrups and move vertically inside cylinders closed at the bottom and sup-
ported in a fixed position by the central post. The two cylinders do not touch
but fit sufficiently closely for the friction of the displaced air to give the
desired damping. An adjustable vent allows the desired amount of damping
to be obtained.

The magnetic damping is obtained by suspending an aluminum plate from
the stirrup between the poles of a magnet which is supported by the balance
post. The current induced by the vertical movement of this plate causes the
damping.

P. I] POINT OF REST 11

Students should never attempt adjustments or repairs, but should report the
condition to ihe instructor. The balance should be kept scrupulously clean
at all times; any evidence of any material spilled on the pans or in the case,
or of corrosion on the pans should be reported immediately. Even if appar-
ently clean, the pans should be gently brushed off with a small soft camers-
hair brush (provided for that purpose only) before each series of weighings.

3. It cannot be too strongly emphasized that this release of the beam
must be done with the utmost care, as by it the knife edges with their loads
are brought into contact with the bearing plates. The operator of a balance
can tell by the feel as the beam control is turned at what point the knife
edges make contact and must use especial care to make this contact as gently
as possible.

4. The first swings of the balance may be somewhat erratic; therefore it is
preferable not to begin immediately recording the swings.

5. In order to avoid errors due to parallax the observer should seat him-
self squarely in front of the case and attempt to have the line of sight per-
pendicular to the scale as the reading is made.

6. See the discussion above for the reason for taking an odd number of
readings, or "an even number on one side and an odd on the other"; seven
or even nine readings may be taken if it is desired.

7. NOTEBOOKS. Because the recording of experimental data is of such
fundamental importance, a few brief rules will be indicated: (1) A permanent
bound notebook should be used and all data immediately recorded in it in ink.
Pencil marks are not permanent, can be erased or altered, and may inspire
doubt as to the integrity of the work. The use of loose scraps of paper is not
to be tolerated. Original data are of more value than those which have been
transcribed for the sake of neatness. (2) Date each entry. This is often of
value in determining the time required for certain effects, in checking stand-
ard solutions, and in determining the sequence of experiments and may be of
great value in research work leading to patent proceedings. (3) Make a com-
plete record of the experiment. Before beginning, note the title and object
of the experiment. Then record not only all the measurements made, but all
experimental conditions, such as temperature, and so forth, however insig-
nificant, that may in any way affect the results or conclusions. Many experi-
ments have had to be repeated because some condition not considered worth
noting was later found to influence the results. The method of procedure
should be so clearly stated that another person could duplicate the results.
Figures and diagrams of apparatus are extremely valuable and space-saving.
Where a mass of data is to be taken, it is more easily comprehended and space
is conserved if a tabular form is provided and the measurements entered
directly therein. Clearly indicate the significance of all data; do not rely
on the memory to interpret them at some later time. (4) Record all experiments ;
if they are later known to be faulty, they can be so labeled with an explana-
tion of the causes or mistakes such a recordJends confidence in the note-
book. (5) After finishing an experiment, record at once a summary of the
results and any conclusions which can be drawn. This may save much time
if it is necessary to refer to, or to interpret, the data at a later time.

8. In the recording of observations and measurements and in calculations
involving these, attention should be paid to the proper use of "significant

12 THE ANALYTICAL BALANCE [P. II

figures" Significant figures are the numbers expressing the value of a meas-
ured quantity, and by convention the digits comprising the number are ex-
tended to, but not beyond, one digit whose value is doubtful. Thus, as the
observation of the extreme of a swing is usually uncertain to approximately
0.2 of a division, it would be recorded to one decimal place and not beyond;
an observation thought to be exactly ten would be recorded as 10.0, not
merely 10, and not 10.00.

In calculations involving the multiplication or division of significant
figures the principle to be kept in mind is that the percentage precision of the
final resulting figure cannot be greater than that of the percentage precision
of the least precise of the measurements entering into the calculation. As an
example, in this procedure the calculated value of the point of rest cannot be
extended to any such number as 10.152, as it would imply a precision obvi-
ously greater than that of the measurements on which it is based; the number
should be rounded to 10.2. In discarding superfluous figures increase the
last figure by one if the discarded figure is 5 or greater.

In adding or subtracting significant figures discard all figures occurring
to the right of the term extending the least number of figures to the right
relative to the decimal point. Thus, in adding the terms 24.6, 8.72, and
0.2564, each term should be rounded off to only one figure to the right of the
decimal point; the sum would be 33.6.

9. The rest points obtained by a series of three or more such determina-
tions should not show a maximum deviation of more than 0.2 of a division.
A constant shift in the point of rest usually indicates uneven temperature
conditions.

Procedure II: DETERMINATION OF THE POINT OF REST.
B. Method of Short Swings. Release the pan and beam
arrests and, if necessary, set the beam in motion by momen-
tarily touching it near the center with the rider so that a
swing of not over 1 to 1.5 divisions is obtained (Note 1).
Determine the point of rest by visually estimating the cen-
tral point of these swings. Record this value, and without
arresting the beam, repeat the determination by recording
the extremes of a swing and taking the mean. Raise the
beam and release the pan rests. Repeat the determination
(Note 2).

Notes:

1. Usually, unless the point of rest is exactly at the center of the scale,
a swing of from 0.5 to 1 division can be obtained without the use of the rider
by carefully releasing the beam.

2. The values obtained by this method should be compared with those
obtained by Method I. It is recommended that the discussion of the relative
merits of the two methods by Wells, J. Am. Chem. Soc., 42, 411 (1920), be
read. From the data obtained the analyst can draw his own conclusions
as to the relative precision of the two methods and decide which method he is
justified in using.

P. HI] SENSITIVITY 13

P. III. The Determination of the Sensitivity of a Balance

Discussion. The structural factors influencing the sensitivity
of the balance have been discussed above. It is advantageous to
know the sensitivity of the balance at various loads, because this
knowledge enables one to estimate the weight to add or subtract in
order to restore the original point of rest when making weighings,
because it is useful in making weighings by the sensitivity method
(see p. 14), and because it indicates the maximum load which the
balance should carry. It is generally stated that a balance should
not be made to carry an excess of weight over that which causes the
sensitivity to decrease to 40 per cent of its maximum value ; analytical
balances are usually designed to carry a maximum of 100 or, for
better instruments, 200 g on each pan. Obviously this criterion
cannot be applied to the somewhat rare case in which a balance shows
a continued increase in sensitivity with loading. A convenient sen-
sitivity for making weighings to 0.1 mg is from 2 to 5 divisions per
milligram.

Procedure III : DETERMINATION OF THE SENSITIVITY OF A
BALANCE. Determine the point of rest of the empty balance,
place upon one of the pans (Note 1) a 1-mg weight (or if a
1-mg weight is not provided in the set of weights being used,
place the "rider" upon the division on the beam indicating
1 mg, Note 2), and again determine the point of rest. The
number of scale divisions through which the pointer has been
deflected is the sensitivity of the balance with zero load.

Repeat the determination with 10-, 20-, and 50-g loads on
each pan (Note 3), and construct a curve by plotting the
sensitivity as ordinate against the load as abscissa.

Notes:

1. In making weighings the object is usually placed on the left pan, since
the manipulation of weights is more conveniently made upon the right pan.

2. Since balance beams are calibrated differently, care should be taken
that the weight of the rider used should correspond to the value of the cali-
bration directly above the terminal knife edge; the exact weight of the rider
should also be checked.

3. The point of rest with a load may not agree with that of the empty
balance, owing to inequality of length of the arms of the balance or small
errors in the weights; if the pointer is deflected from the central portion of the
scale with the heavier loads, it should be brought back by adjusting the rider.

14 THE ANALYTICAL BALANCE

Methods of Weighing

Discussion. The simplest method of weighing an object would
be to determine the point of rest of the empty balance, then to place
the object on one pan and, by means of adding weights and adjusting
the rider, to restore the original point of rest. This is the method
commonly used by chemists for routine work. In order to avoid
having to adjust the weights to the exact mass of the unknown object,
the method of "weighing by sensitivity " is often resorted to, espe-
cially when making precise weighings. In this method the weights
are applied until the total is within 1 or 2 mg of the required amount
and the point of rest determined. Then, if the original point of rest
and the sensitivity of the balance are known, the weight to be added
or subtracted in order to restore the original point of rest can be
calculated. It is convenient in making weighings by this method
to have the sensitivity expressed in milligrams per division of the
pointer scale. Because the sensitivity of the balance may vary
with the load, it is better practice for precise work to determine the
sensitivity under the actual load conditions by again taking a point
of rest after adding or subtracting a milligram from the load on
the pan. 6 - 6

Methods of Weighing Which Eliminate Errors Due to Unequal Arm
Lengths. In the previous methods it has been assumed that when the
original point of rest has been restored, the applied weights are equal
to the weight of the object; this may or may not be a justifiable

6 The so-called "single deflection" method of weighing is frequently used
for routine work where & saving of time is desired. In this method a weight is
first applied to one arm of the balance (by means of the adjusting nut at the
end of the beam or by adding a milligram weight to one pan) . The beam is then
lowered, the pan rests are released, and the position of the pointer at the limit
of the first deflection is noted. An object is weighed by placing it on one pan
and adding weights to the other until, upon releasing the beam as before, the
pointer again reaches the same position at the extreme of its first swing.

This method requires that a balance with separate beam and pan controls
be used; also that the pan controls be so adjusted that they impart no impulse
to the pans when they are released. For a discussion of this method see
Brinton, /. Am. Chem. Soc., 41, 1151 (1919).

8 In order to expedite the weighing process the "chainomatic balance" has
been developed. This balance eliminates the rider and uses in its place a
fine gold chain, one end of which is supported by the beam and the other end
by a support which can be raised or lowered by a control outside of the balance
case, thus varying the fraction of the weight of the chain carried by the beam.
In this way the rider and fractional weights up to 100 mg are replaced by the
chain. The effective weight applied to the beam is read by means of a vernier
scale.

METHODS OF WEIGHING 15

assumption. When the balance reaches an equilibrium, the moments
arising from the object and the applied weights are equal; that is,

Q X L = R X W,

where Q is the weight of the object, L is the length of the left arm,
W is the applied weights, and R is the length of the right arm. Obvi-
ously, only when L = R is the above assumption justified, and this
may not be true within the desired experimental limits. Accordingly,
for precise work, it is necessary either to determine the effect of this
inequality and then correct for it, or to carry out the weighing in
such manner that the effect is eliminated. The latter procedure is
more commonly used and may be accomplished by either of two
methods.

1. The Method of Substitution. This method, commonly called
Borda's method, consists in placing the object on one pan and bal-
ancing it with any suitable material T (commonly called a tore),
preferably weights from an old set, and then removing the object and
balancing the tare with precise weights W. From the principle of
moments, the following relation applies to the first weighing:

QL = RT,
and for the second,

WL - RT,

from which it is seen that L and R can be eliminated from the equa-
tions and that Q = W.

2. The Method of Transposition or Double Weighing. This
method, commonly called Gauss' method, consists in placing the
object Q on one pan and balancing it with weights W\ 9 and then
transferring the object to the other pan and again balancing with
weights TFV For the first weighing,

QL = RWi,
and for the second,

WJU = RQ,
from which it is seen that

Q =

16 THE ANALYTICAL BALANCE [P. IV

and the true weight is the geometric mean of the observed weights.
As the difference between Wi and W% is relatively small, the arith-
metical mean,

will give the true weight well within the desired experimental limits
and is generally used. This latter method is usually more con-
venient, since it does not require an additional set of tare weights
and is more generally used for precise weighings,

The Correction for Buoyancy

Discussion. A further assumption which has been made in the
above methods was that the only force acting on the object and
weights was that of gravity. Actually objects and weights are
buoyed up by the weight of the air which they displace, and unless
they are of the same density, this force will not be the same. Al-
though the effect is so small that it can be neglected for most work,
when calibrating volumetric vessels or carrying out weighings where
unusual accuracy is desired, the correction for buoyancy must be
made.

If W is the observed weight, then the weight in vacua, W v , will be
W 9 = W + (V 1 - V")d, where V and V" are the volumes of the
object and weights respectively, and d is the density of the air at the
time of the weighing. However, V = W 9 /d', and V" = W /d",
where d r is the density of the object and d" the density of the weights.
However, as W v differs from W by such a small amount, as an ap-
proximation V = Wo/d', and therefore

fw, w\ ,

I _ I (l t

\ d d /

> /

W 9 = W

Although the density of the air varies with the barometric pressure,
the temperature, the humidity, and the carbon dioxide content,
it is usually sufficiently accurate to assume that 1 ml of air weighs
0.0012 g.

P. IV. The Calibration of a Set of Weights

Discussion. Unless they have been specifically certified by the
makers or by the Bureau of Standards, or if they have been in previ-
ous use, weights cannot be relied upon for precise work until they
have been calibrated. The method used for carrying out this calibra-
tion consists in principle in assuming a value for one of the smaller

P. IV] CALIBRATION OF WEIGHTS 17

weights, or the rider, and obtaining the value of all the other pieces
in the set in terms of this arbitrary standard. These relative values
may then be used where the absolute weight of the object is not
necessary; however, in order to obtain absolute weighings (as, for
instance, in calibrating volumetric vessels) it is necessary to compare
one of the larger pieces with a piece of known value, usually one
certified by the Bureau of Standards. An absolute value having
been established for this piece, the values of the remaining ones
can be calculated by means of the ratio of the absolute and relative
values. In order to eliminate an accumulation of errors due to
inequality of arm lengths the weighings should be made by substitu-
tion or by double weighing. The latter method avoids the use of a
second set of weights, and the modification suggested by Weatherill 7
is particularly adapted for the calibration process. This method
can be illustrated as follows : Suppose that the two pieces Wi and W*
are to be compared; Wi is placed on the left pan and W 2 on the right
and the point of rest determined. The pieces are exchanged, and the
difference in weight, Aw, necessary to restore the first point of rest is
determined. For the first point of rest

and for the restored point of rest

W*L = R(Wi + Aw).
From this it is seen that

Awfl

W* - W l =

L + R'

and as Aw is small, it can be safely assumed that L and R are equal;
therefore

If an arbitrary value has been assigned to W\, then TF 2 = W\ +
(Aw/2). If weight has to be added to W\ to restore the first point of
rest, Aw is positive; if weight has to be removed, Aw is negative.

This method is illustrated by the actual data taken in calibrating a
set of weights and shown in Table I. The 0.005-g weight was taken
as the arbitrary basis on which to obtain the relative values. The

T Weatherill, J. Am. Chem. Soc,, 52, 1938 (1930).

TABLE I
THE CALIBRATION OF A SET OP WEIGHTS

Face

Value

Weights on Left
Pan

Weights on Right
Pan

Point
of Rest

Au>/2
(mg)

Rel. Value
(tasedon
5-mg wt.)

Absolute
Value

Corr.
(mg)
rounded)

0.005
rider

rider
0.005

10.2
9.1

-0.11

0.00500
0.00489

0.00501
0.00490

-0.1

0.01

0.005, rider
0.01

0.01
0.005, rider

8.7
10.5

+0.18

0.01007

0.01008

+0.1

0.01*

0.005, rider
0.01*

0.01*
0.005, rider

9.1
10.2

+0.11

0.01000

0.01002

0.02

0.01, 0.01*
0.02

0.02
0.01, 0.01*

9.5
10.0

+0 05

0.02012

0.02015

+0.2

0.05

0.02 + 2
0.05

0.05
0.02 + 2

10.4
8.8

-0.16

0.04992

0.04998

0.1

0.05 + 2
0.1

0.1
0.05 + 2

10.5

8.9

-0.16

0.09984

0.09997

0.1*

0.05 + 2
0.1*

0.1*
0.05 + 2

10.4
9.0

-0.14

0.09986

0.09999

0.2

0.1, 0.1*
0.2

0.2
0.1, 0.1*

9.7
9.7

0.19970

0.19995

0.5

0.2 + 2
0.5

0.5
0.2 + 2

9.7
9.4

-0.03

0.49937

0.50000

0.5 + 2
1

0.5 + 2

9.9
9.1

-0.08

0.99869

0.99996

1
1*

1*
1

9.7
9.7

0.99869

0.99996

1**

1
1**

1**
1

9.7
9.7

0.99869

0.99996

1, 1*
2

2
1, 1*

6.7
12.4

+0.57

1.99795

2.00049

+0.5

2, 1, 1*, 1**
5

5
2, 1, 1*, 1**

10.8
7.4

-0.34

4.99368

5.00003

5, 2, 1, 1*, 1**
10

10
5,2,1,1*,!**

10.1
6.3

-0.38

9.98732

10.00002

10*

10
10*

10*
10

7.7
8.7

+0.10

9.98742

10.00012

+0.1

10, 10*
20

20
10, 10*

7.0
6.3

-0.07

19.97467

20.00007

+0.1

20 + 2
50

50
20 + 2

4.7
0.5

-0.42

49.93669

50.00019

+0.2

1 standard
1

1
1 standard

9.8
9.4

-0.04

0.99996

2 indicates the sum of the smaller weights of the series, in this case the
0.01-, 0.01*-, 0.005-g weights and the rider.

P. IV] CALIBRATION OF WEIGHTS 19

sensitivity of the balance did not vary appreciably from 0.20 mg per
division throughout the load applied. It should be noted that the
arms of the balance used for obtaining this data were not exactly
equal, the effect caused by this becoming quite appreciable at the
heavier loads. Column 1 shows the face value of the weight being
compared, columns 2 and 3 the weights on each pan, and column 4 the
point of rest with the weights in the original and transposed posi-
tions. Column 5 shows Aw/2 (in milligrams), and in column 6 the
values of the pieces on the basis of the 0.005-g weight are summed up.
After comparing the 1-g piece with a Bureau of Standards 1-g weight,
the absolute values of the other pieces are calculated by correcting
each of the pieces by the ratio of the absolute weight to the relative
weight, that is, 0.99996/0.99869, and the values shown in column 7
obtained. 8 For convenience in making weighings the value of the
correction in milligrams to be applied to the face value of each weight
to give its absolute weight has been rounded off and is shown in
column 8.

Procedure IV : CALIBRATION OF A SET OF WEIGHTS. Ex-
amine the balance, brushing off the pans, and see that it is in
proper condition. Examine the weights to be calibrated,
noting if they have any dust or dirt sticking to them or if
any pieces are corroded or tarnished (Note 1). Place the
rider on the right arm exactly at the 5-mg calibration and
the 0.005-g weight on the left pan and determine the point
of rest (Note 2); place the rider on the left arm (Note 2)
and the 5-mg weight on the right pan and again determine
the point of rest (Note 3). Similarly compare the 10-mg
weight with the 5-mg weight plus the rider, and continue

8 The labor involved in the calculations required for the calibration can be
greatly reduced as follows: In the data above we find that the relative value
for the weight of face value 1 g is 0.99869, whereas the absolute value is 0.99996;
therefore for this 1-g weight a correction of + 0.00127 g has to be applied to
the relative value in order to obtain the absolute value. As the relative values
are consistent among themselves, the correction to be added to any other
weight can be assumed to be +0.00127 g multiplied by the ratio of the face
value of that weight to that of the 1-g weight; in the case above, the correction
to be added to the relative value of the 0.5-g piece will be +0.00127 X 0.5/1,
which is 0.00063. The advantage of applying a ratio to the correction rather
than to the relative value is that only the same number of significant figures
have to be carried in the ratio as are necessary in the correction. For a dis-
cussion of the assumptions involved in the calibration process and a justifica-
tion of the above method of procedure see the article by Frank H. Hurley, Jr.,
/. Ind. Eng. Chern., Anal. Ed. t 9, 238 (1937).

20 THE ANALYTICAL BALANCE [P. IV

until all the pieces of the set have been compared as indi-
cated in Table I. Next obtain a standard weight (Note 4)
and compare it with a piece of the same denomination
from the set.

From the two points of rest obtained in comparing each
piece calculate the Aw/2, referring to the curve obtained in
P. Ill for the sensitivity with varying loads. From these
values compute the relative values for the pieces on the basis
of the 5-mg piece (or the 10-mg piece, if it is used). From
the ratio of the absolute value (obtained by the comparison
of a piece with the standard weight), and the relative value
obtained for this same piece, calculate the absolute value
for each piece of the set (Note 5). From these values calcu-
late to 0.1 mg the correction to be applied to the face value
of each piece in order to obtain its true value.

Notes :

1. Dust sticking to weights should be brushed off with a camePs-hair
brush; if the weights appear dirty or corroded, the instructor should be
notified. Dirt can be removed by rubbing the weights with a soft silk cloth
or fine chamois skin. Weights that are corroded should be returned to the
manufacturer for repair and adjustment. Weights should never be touched
with the hands. Use the ivory- or bone-tipped forceps provided with the set;
these forceps should not be used for any other purpose.

2. If the balance is equipped with a 10-mg rider, it can be compared with
the 10-mg weight. If the left arm of the balance has no calibrations, the
rider can be placed directly on the pan.

3. It is recommended that a tabular form similar to that shown in Table I
be drawn in the notebook and the data taken directly in it. Only the weights
and points of rest should be recorded while in the balance room ; the remain-
ing calculations can be completed subsequently.

4. Weights accompanied by a certificate as to their mass value can be
obtained from the Bureau of Standards, or individual weights can be sub-
mitted to the Bureau for standardization. The fees are nominal.

The calculations involved in the calibration are somewhat shortened, the
possibility of accidental error is reduced, and more accurate values are ob-
tained for the heavier weights if two standard weights are available, a 1-g
piece being used as a reference standard for the fractional weights and a 50-g
piece for the heavier pieces. 9

5. See footnote 8, p. 19, in regard to these calculations.

If only a 1-g standard weight is available, it is also to be noted that, having
obtained the value of the 1-g weight by comparison with this standard 1-g
piece, the absolute values for the larger pieces can be obtained by building up

'See Hopkins, Ziun, and Rogers, /. Am. Chem. Soc., 42, 2528 (1920), for
suggestions in this regard.

P. IV] GENERAL RULES 21

directly from it; this method should be used to check those obtained from the
relative values by the ratio calculated above.

General Rules for the Use of the Balance

1. Keep the balance case and pans clean. Brush off pans with
camel's-hair brush before making a weighing.

2. Sit directly in front of the balance to avoid error from parallax.

3. Release the beam with a gradual steady motion.

4. Arrest the swing of the beam only when the pointer is at the
center of the scale. Always raise the beam before leaving the bal-
ance. Do not leave weights or objects on pans after weighing.
Keep the case closed.

5. Use only your own weights. Handle the weights w r ith the for-
ceps provided for that purpose; never use your hands.

6. Place weights and objects as near as possible to the center of the
pan. Damp oscillation of the pans before determining the point
of rest.

7. As a general rule, substances should not be placed direr tly upon
the pans. No liquids, unless in stoppered bottles, are to be brought
into the balance case.

8. Allow heated objects to cool to balance-room temperature before
weighing.

9. Do not vigorously rub glass objects (possibly producing an
electrostatic charge) just before weighing.

10. When making successive weighings use the same larger weights
whenever possible.

11. Ascertain rated capacity of the balance, or determine it from
the change in sensitivity, and do not exceed it. . ^

12. Triple count the weights to avoid error; (I) count the Weights
on the pan; (2) count the spaces in the box; and (3) cheek the weights
as they are removed from the pan.

13. Record weighings at once in your notebook. Do not carry
scraps or loose sheets of paper into the balance room.

14. Report any injury to the balance to the instructor at once.
Do not attempt to make adjustments or repairs yourself.

Volumetric Methods of Analysis

GENERAL DISCUSSION

Volumetric methods are used extensively in this system of analysis
for estimating the amount of the constituents which have been de-
tected. This is done because, if standard solutions are available,
volumetric methods, in general, can be carried out much more
rapidly and usually require less experience and special technique
than gravimetric methods. For these reasons the first section of this
book is devoted to a general discussion of volumetric methods and
to the preparation and standardization of the solutions which are
later required.

Volumetric methods differ from gravimetric methods in that the
final weighing of a compound of definite percentage composition,
which is the characteristic feature of gravimetric methods, is re-
placed by the measurement of the volume of a standard solution which
reacts directly or indirectly with the substance being estimated.
Thus, the amount of chloride in a sample of drinking water can be
determined by titrating the sample with a standard silver nitrate
solution by means of the precipitation reaction

OP + Ag + = AgCl ( .) ,

or the amount of iron in an ore can be determined by a titration with
a standard permanganate solution, the iron being first dissolved and
reduced to the ferrous condition by suitable means. In this case the
reaction is represented by the equation

+ Mn07 + 8H+ = 5Fe+++ + Mn^ + 4H,O.

In order for a given reaction to be developed into a volumetric
method, certain requirements must be met, namely:

1 . The fundamental reaction involved, that is, the reaction between
the standard solution and the titrated substance, must, when it
reaches an equilibrium, be complete to within the desired precision.
The equilibrium conditions can usually be predicted from available
data such as solubility measurements, ionization constants, or
potential values.

2. The reaction must proceed with practical rapidity. This "re-
action rate" cannot be predicted from the equilibrium calculations
and usually must be measured experimentally.

GENERAL DISCUSSION 23

3. The reaction must be stoichiometric. That is, the reaction
must proceed according to some definite equation, or equations, so
that from the volume and concentration of the standard solution
used the exact amount of the titrated substance present can be
calculated. These criteria are not met by titrations involving "side
reactions" or "induced reactions." An example of each of these
follows: (a) The reaction between iodine and thiosulfate is assumed
to proceed as follows :

I* + 2SOr = 2I~ + S 4 0r;

however, if the titration is made in a solution which is slightly alka-
line, part of the thiosulfate is oxidized to sulfate. (b) The reaction
between stannous tin and permanganate is assumed to have the
following stoichiometric equation:

2MnO7 + 5Sn + + + 16H+ = 2Mn f + + 5Sn ++ ^ + 8H*O;

however, if the titration is made in the presence of air, the oxidation
of stannous tin by oxygen is induced, and such a large error intro-
duced that this particular determination is not practical. Often the
existence and magnitude of these effects caYi be determined only by
experiment.

4. Some means must be available for determining the completion
of the titration. Various methods are available: (a) The titrating
solution may have such a distinctive and intense color that a very
slight excess becomes apparent (permanganate solutions); (b) indi-
cators (substances which undergo a color change at the desired point)
may be used; or (c) physical methods, such as observing the change
in potential or conductance of the solution, are available. By these
means the "end-point" of the titration is taken. This is to be dis-
tinguished from the "equivalence-point," which is the point at
which an exactly equivalent amount of the titrating substance has
been added. The error in the titration is measured by the difference
between the equivalence-point and the end-point.

The imposition of the above requirements greatly limits the number
of reactions upon which precise volumetric methods can be based.

Types of Volumetric Methods

The methods used in volumetric analyses may be divided into three
general types, as follows :

I. Precipitation methods.
II. Oxidation and reduction methods.

III. lonization methods (including neutralization, complex-ion,
and un-ionized compound formation). This classification is based

24 VOLUMETRIC METHODS

upon the type of reaction taking place between the substance being
titrated and the standard solution. With methods of the first type
the completeness of the reaction is dependent upon the solubility
of the precipitate formed, and the conditions obtaining at the equiva-
lence-point usually can be calculated if solubility-product data are
available. With methods of the second type the completeness of the
reaction is dependent upon the potentials involved. Those of the
third type are dependent upon the formation of un-ionized com-
pounds, and, especially in neutralization reactions, upon the degree of
ionization of the acids and bases involved. The ionization con-
stants of these substances being available, calculation of the equi-
librium conditions can be made. Specific examples of these general
types will be discussed in detail in the subsequent procedures.

Concentration Units

Concentrations may be expressed in many different units. There-
fore it is necessary to state definitely the conventions to be used
hereafter in order to avoid confusion. Two general classifications
are in use. First, volume concentration, which states the amount of
any substance per unit of volume; and second, mass or weight con-
centration, which states the amount of any substance per unit of
mass of the solvent or the solution. In volumetric analysis, interest
is centered on volume measurements and volume concentrations,
and the unit of volume is the liter (defined in Note 8 of P. V).
In this book, volume concentrations will be expressed as formal,
molal, or normal.

A formal solution contains 1 gram-formula weight of the stated
compound per liter and states only the total amount of the substance
present, not the specific molal or ionic species in which it may exist
in the solution.

A molal solution contains 1 gram-molecular weight of the particular
molecule or ion stated per liter. 1 As an example, a solution containing
12.006 g of sodium hydrosulfate (NaHSO 4 ) per liter is 0.1 formal in
NaHSO 4 , and, as, according to modern theories, the sodium hydro-
sulfate is completely ionized into sodium and hydrosulfate ions, the

1 It should be pointed out that there is extensively used another convention
in which the term molar is used to designate the moles of solute per liter of
solution and molal the moles of solute per 1000 g of solvent; when this con-
vention is used, the above distinction is not made between formal and molal
concentrations. It is found that the use of molar to denote volume con-
centration and molal to denote weight concentration is confusing when used
with normal to denote volume concentration. Also, it appears more definite
to specify weight concentrations (which are not extensively used in volumetric
work) by the terms " weight-formal' 1 or "weight-molal, 11 as the case may be.

CONCENTRATION UNITS 25

sodium ion (Na + ) concentration is 0.1 molal. However, the hydro-
sulfate ion (HSO7) is partly ionized into hydrogen (H+) and sulfate
(SO7) ions, and according to measurements which have been made
the solution is 0.034 molal in SOT and H + and 0.066 molal in HSO7.
A normal solution contains 1 gram-equivalent weight per liter.
A gram equivalent of a substance is that weight of it which, directly
or indirectly, reacts with 1 gram-atomic weight (1.0081 g) of hydro-
gen. The same substance may have several equivalent weights,
depending upon the reaction in which it is involved. Thus it may
take part in precipitation or neutralization reactions (broadly termed
metathetical reactions), or this same substance may enter into one
or more oxidation and reduction reactions. As an illustration, iodic
acid in either of the following reactions would have 1 equivalent per
formula weight, and a solution which was 1 formal would likewise
be 1 normal.

HIO 3 + OH" = H 2 O + IOF (Neutralization) (1)

HIO 3 + Ag+ = AgIO 3( .) + H+ (Precipitation) 2 (2)
As an oxidizing agent in the reaction

lOa" + 3H 2 S = I" + 3S () + 3H 2 0, (3)

iodic acid would have 6 equivalents per formula weight and a normal
solution would contain one sixth of a formula weight per liter.
However, in the reaction

+ 6H+ = JI, + 5Fe +++ + 3H 2 O, (4)

it has only 5 equivalents per formula weight, and a normal solution
would contain one fifth of a formula weight per liter, thus being 0.2
formal. It is therefore seen that the concentration of a normal solu-
tion is not uniquely specified, and that confusion is likely to result,
unless the reaction for which the solution is to be used is also stated.
Because of this fact objections have been raised to the use of this
unit; however, in stoichiometric calculations the use of equivalents
and of normal concentrations is of great convenience. Therefore
throughout the procedures of this system reagents will be specified
by their normal concentration except where confusion is likely to

2 In writing equations, solids will be indicated by the subscript (s) and gases
by (g) I* 1 general, equations will be written in their ionic form (as above);
this form is simpler to construct, and shows the fundamental changes taking
place in the reaction and the effect of each ion or molecule upon the equilibrium
of the reaction. If the principles involved in writing the ionic equations are
understood, the completed equation can be readily constructed when this is
desired for stoichiometric calculations.

26 VOLUMETRIC METHODS [P. V

result, in which case their formal concentration will be given. Molal
concentrations will be used in mass-action considerations where it is
desired to express the concentration of some particular species of ion
or molecule as it exists in the solution.

Volume concentration is sometimes expressed as the weight of a
given substance per unit of volume; this unit will not be used here.

Weight concentrations will be expressed in terms of grams of
solute per 100 g of solution, or the weight per cent.

Formal and molal weight concentrations, although extensively
used in thennodynamic calculations, will not be used here.

The Preparation of Standard Solutions

Discussion. Standard solutions for volumetric work can be pre-
pared by two methods: First, directly, by weighing precisely the
required amount of a pure substance, dissolving it, and diluting this
solution at a definite temperature to an exactly known volume in a
calibrated flask; or, second, by making the solution of approximately
the concentration desired and then measuring the volume of it
which reacts with a precisely weighed amount of some pure substance
(called a primary standard), thus standardizing it. The first process
is used when the substance from which the standard solution is made
can be obtained in a high degree of purity, can be precisely weighed,
and does not change upon being dissolved in water or when the solu-
tion is diluted (for example, react with the gases or traces of organic
matter usually present even in distilled water). The second process
is more frequently employed because many of the substances used
for standard solutions are not readily obtained in a sufficiently pure
form, or are not stable when first dissolved in distilled water.

PRECIPITATION METHODS OF VOLUMETRIC ANALYSIS
P. V. The Preparation of a Standard Silver Nitrate Solution

Discussion. As very pure silver nitrate can be prepared by re-
crystallization (or purchased), and as it can be dried and weighed
(even fused at 220 to 225) without change, provided organic matter,
reducing gases, or hydrogen sulfide are avoided, standard solutions
of silver nitrate are usually prepared by precisely weighing out the
desired amount of the salt and then dissolving and diluting this to the

P. V] SILVER NITRATE SOLUTIONS 27

desired volume in a volumetric flask at a specified temperature.
Such solutions are then stable, provided they are protected from
organic matter, reducing gases, and also from light.

When the silver nitrate solution with its concentration precisely
known is available, it can be used for standardizing other solutions
which are not readily directly prepared.

Procedure V: PREPARATION OF A STANDARD SILVER
NITRATE SOLUTION. Clean a weighing bottle, wipe it with a
clean, dry, lintless towel (Note 1), dry it in an electric oven or
by gently heating it with a burner, allow it to cool in a desic-
cator (Note 2), and weigh it to 0.01 g (Notes 3, 4). Weigh
into the bottle approximately 17.0 g of silver nitrate (Note
5) of the highest purity obtainable, place the open bottle in
a small beaker loosely covered with a watch glass, and heat
for at least 2 hours in an electric oven at 105 to 110C.
(Note 6). Again cool in the desiccator and weigh. With
the aid of a large-stem funnel carefully transfer the silver
nitrate to a clean, calibrated 1-liter flask (Notes 7, 8), then
rinse the weighing bottle out repeatedly with 5-ml portions
of water (Notes 9, 10). Use a stirring rod to guide the solu-
tion from the bottle to the beaker and take care to avoid loss
by spattering. Dissolve any silver nitrate crystals remain-
ing on the funnel and fill the flask almost to the mark with
water at room temperature (Note 11); so use the funnel as
to avoid wetting the flask above the calibration mark.
Finally, remove the funnel, being sure to rinse off any solu-
tion adhering to it; then with a dropper or pipet add water
until the lower meniscus just coincides with the calibration
mark. Any droplets of water spilled against the side of the
flask above the calibration mark should be removed with
filter paper before making this final adjustment. Stopper
the flask and shake it vigorously until the contents are per-
fectly mixed (Note 12).

Transfer the solution to a clean ground-glass-stoppered
bottle (Note 13), first rinsing it out with several small por-
tions of the silver nitrate solution. Calculate the normality
of the solution at the standard laboratory temperature
(20C.) and label the bottle (Note 14).

Notes :

1. After being cleaned and dried, a weighing bottle, or any other piece of
apparatus which is to be weighed precisely (to 0.1 to 0.2 mg) should be ban-

VOLUMETRIC METHODS

[P. V

died by means of a piece of paper, lintless cloth, or chamois skin folded around
it. It is preferable not to handle such apparatus with the fingers, since
they may be greasy or moist, and prolonged handling may raise the tem-
perature of the apparatus sufficiently to cause an apparent decrease in
weight. These effects would not be significant in the above case, where the
object is to be weighed to only 0.01 g.

2. A desiccator is a glass container in which objects, especially crucibles,
may be placed and allowed to cool in a dry atmosphere protected from dust,
carbon dioxide, and the fumes common to a laboratory. A common type of
desiccator is shown in Fig. 4. The desiccating agent most frequently used is

anhydrous calcium chloride; for special pur-
poses sulfuric acid or phosphorus pentoxide
may be used. The desiccator should be
opened only when necessary, closed as soon
as possible, and k.ept tightly closed at all
other times; 3 the desiccant should be re-
placed as soon as it begins to lose its effi-
ciency. After being heated, an object should
be allowed to cool from 20 to 30 minutes be-
fore it is weighed.

3. Methods of weighing out samples.
Material prepared for analysis, primary
standards, and similar substances are usually
dried and kept in glass-stoppered weighing
bottles. Samples may be weighed out from
these by either of two general methods,
here termed the "direct" and "difference"
methods. By the direct method (used in this
procedure) an empty weighing bottle, a

watch glass, a piece of glazed paper, or any other suitable container is
weighed, the desired amount of material from the weighing bottle trans-
ferred to it, and the weight of the container plus the material found. The
material is usually transferred by holding the sample bottle over the weighed
receiver, removing the stopper, and then tipping the bottle and slowly rotat-
ing it, holding the top close to or against the receiving vessel, until approxi-
mately the desired amount of material has been transferred. If an exact
weight of sample is desired, material can be added or taken by means of a
pointed steel spatula until this is obtained. In the above procedure time
may be saved by drying the silver nitrate in a clean weighing bottle and then
weighing it on a clean, dry, previously weighed watch glass. Care must be
taken not to get any of the silver nitrate on the balance pans and not to lose
any in the subsequent transfer.

In using the "difference" method, the weighing bottle containing dry
material is first weighed, the desired amount of material transferred from it
directly to the beaker or flask in which it is to be treated, and the weighing
bottle closed and again weighed; the loss or "difference" in weight represents

Fig. 4. Desiccator.

1 Booth and Mclntyre, /. Ind. Eng. Chem., Anal. Ed., 8, 148 (1936), give
data showing the length of time required to dry the air in a desiccator.

P. V] SILVER NITRATE SOLUTIONS 29

the sample taken. The transfer of the material is accomplished by essen-
tially the process described above, care being taken that all of the material
leaving the bottle is caught in the beaker or flask. Since it is not good
practice to return material from the receiving vessel (which may not be
entirely dry or clean) to the sample vessel, it is not convenient to take exact
sample weights by this method. Approximate sample weights can be taken
by removing slightly less than the weight of the desired sample from the
right pan and then intermittently removing material from the bottle until an
approximate balance is again obtained; however, frequent removal and
insertion of the stopper may cause loss of material by "dusting," especially
with finely powdered dry material.

The direct method is usually more convenient where a large amount of
material is to be weighed, presents less chance of loss of finely powdered
material, and pertnits taking an exact sample weight; the "difference"
method is more rapid where several successive weighings are to be made or
where the sample is likely to absorb moisture during weighing, and permits
weighing out the sample directly into the beaker in which it is to be treated.

4. The analyst should always keep in mind the precision to which measure-
ments should be carried out. One is not justified in carrying most chemical
analyses to a precision of greater than one part in a thousand or one tenth
of one per cent; this limit is set by instrumental errors, by errors inherent in
the methods involved, and in some cases by uncertainty in the atomic
weights. This being true, it is obviously a waste of time to weigh out 17 g of
silver nitrate to 0.0001 g, which would represent a precision of one part in
170,000, or approximately six ten-thousandths of one per cent. If the
weighing is made precise to 0.01 g, it will be within the desired limits; ob-
viously a sample of 0.17 g should be weighed precisely to 0.0001 g.

5. Directions are given for preparing a liter of the solution; if less than
this is required, a smaller volume should be prepared. By approximately
weighing out the amount of silver nitrate taken, no excess is dried, and the
solution will be close to the desired normality. If it is desired to make the
solution exactly 0.1 normal, small crystals of the dried salt can be removed or
added by means of a clean, dry, pointed steel spatula until exactly 16.99 g of .
the nitrate are taken.

6. As silver compounds are easily decomposed by organic matter and
reducing agents, the oven should be protected from dust and from reducing
gases. Paper labels should not be used on articles placed in the drying oven,
because the adhesive material used chars, giving off volatile reducing sub-
stances (tarry materials) which cause the reduction of silver compounds.
Weighing bottles should be identified by marking with a pencil on the ground-
glass seal around the stopper. Most commercial silver nitrate will darken
slightly on being heated, but with the better grades this decomposition is
not significant.

7. The cleaning of volumetric apparatus. The vessels used in volumetric
measurements must be scrupulously clean. A vessel may be considered
suitable for use when it is visually clean and when a solution drains from it
without leaving perceptible streaks or dropkts on the surface. The most com-
mon contaminant is a grease film. Glass apparatus for volumetric work
often can be effectively cleaned by scouring it with soap powder and then
rinsing it with distilled water; to economize with distilled water, beakers,

30 VOLUMETRIC METHODS [P. V

flasks, and so forth may be liberally rinsed with tap water and finally with a
small portion of distilled water. Except when a resistant film of grease is
encountered which cannot be removed by treatment with the soap solution,
or when the construction of the vessel makes the use of soap powder ineffec-
tive, the use of the so-called "cleaning solution/' 4 while effective, is to be
discouraged because of its corrosive nature and the difficulty with which it is
completely removed from glass surfaces. When cleaning solution is used,
it may be necessary to allow it to stand in the vessel for an hour or even
overnight. More rapid action is frequently obtained by treating the vessel
with 6 n. NaOH for 5 or 10 minutes and then following this with the cleaning
solution. Calibrated vessels should not be heated if it can be avoided.

8. The calibration and use of volumetric apparatus. Volumetric vessels
are usually calibrated either "to contain" (flasks) or "to deliver" (burets and
pipets) a definite volume of solution at a stated temperature, and because of
drainage an appreciable error will be made if they are used for other than the
indicated purpose.

When using volumetric apparatus for precise work, care should be taken
that the solution is at the temperature for which the vessel is calibrated or
that the proper correction is made if working at a different temperature.
Since there has been some confusion in regard to the units of volume and
since glassware as often purchased is not precisely calibrated, all such
apparatus should be checked before being used. This is usually done by
finding the weight of water which such apparatus will contain or deliver,
and converting this weight into the metric unit of capacity. 5

As the actual calibration of apparatus is carried out at room temperatures
and under atmospheric pressure, it is necessary to calculate the proper
weight of water to be taken. As an example,* if it is desired to calibrate a
liter flask for use at 20, it will be found by reference to a table showing the
relative densities 6 of water at various temperatures that this volume of water
would weigh 998.23 g in vacuo. The apparent weight of this mass of water
in air of mean humidity and 760 mm pressure is 997.19 g. The calculation
for temperatures other than the standard temperature of 20 can be similarly
made, remembering that the average cubical coefficient of expansion (the
increase in volume per unit of volume per degree centigrade) of glass is
0.000025. Flasks which are calibrated "to contain" are checked by weighing
the clean dry flask on a balance of suitable capacity, carefully filling the
flask with water of the proper temperature, taking care that there is no
water above the calibration (as indicated in the procedure above), and again

4 Cleaning solution is made by adding 50 g of commercial sodium dichromate
to 500 ml of commercial concentrated sulf uric acid.

* The liter is the metric unit of capacity and is defined as the volume oc-
cupied by 1000 g of pure water at the temperature of its maximum density and
under a pressure of one atmosphere. One milliliter is not exactly the same as
one cubic centimeter; the latter is derived from the unit of length and is the
volume of a cube with an edge one centimeter in length; 1 ml is 1.000027 cc.

6 The absolute density of water (or any other substance) in C. G. S. units is
the mass per cubic centimeter. The relative density (or specific gravity) is
the ratio of the mass to the mass of an equal volume of water at 4C. There-
fore, in order to obtain the mass of a given volume of water the absolute den-
sity must be used if the volume is in cubic centimeters, the relative density if
the volume is in milliliters.

P. V]

SILVER NITRATE SOLUTIONS

weighing. 7 Pipets and burets are calibrated "to deliver" and are checked by
weighing the water obtained from them (Notes 3-6, P. VI). Reference
books on quantitative analysis should be .consul ted for the details of these
operations.

9. Two methods are used for transferring weighed material to a flask.
If the material is finely ground, crystalline, and readily soluble, it may be
transferred directly by means of a funnel as directed above. Any solid
remaining in the funnel is readily dissolved by the water added. If the
material is of an amorphous nature, tending to "lump" together, or is not
*eadily soluble, it is advisable to transfer it first to a beaker where it can be
more readily brought into solution (heated if necessary) and this solution
then poured through the funnel into the flask. When a solution is being
poured, it should always be guided by means of a stirring rod. This pre-
vents the solution from spattering in the receiving vessel and from running
down the outside of the vessel from which it is poured.

10. Stirring rods, droppers, and wash bottles are constantly used by the
analyst, and an adequate supply should be available. It is recommended
that these be made while the silver nitrate is being dried.

Make the stirring rods by cutting glass rod (3 to 4 mm in diameter) into
lengths varying from 10 to 25 cm, and then heating both ends in a flame until
they are well rounded and smooth (fire polished). If this is not done, they
scratch the inside of vessels in which they are used; precipitates then tend to
crystallize in these cracks, and are difficult to
remove.

Droppers are made by heating and drawing
glass tubing (6 to 8 mm inside diameter, 10 to
15 cm in length) to a capillary at one end and
heating and enlarging the other end with the
end of a file so that it holds a rubber nipple
tightly. It should then be marked at the
position from which it delivers 1 ml, and the
number of drops per ml noted. Droppers are
extremely useful for adding small volumes of
reagents and for washing precipitates with
small volumes of water.

Wash bottles are constructed from Florence
flasks as indicated in Fig. 5. The flask should
be of a resistance glass and the stopper and
rubber tubing should be of the best grade
rubber. The tip of the outlet tube should be
gradually constricted until the desired stream
of water is delivered. When designed for hot
water the neck may be wrapped with asbestos
cord or cork. For greater flexibility, the
mouthpiece can be provided with a piece of
rubber tubing 5 to 8 cm in length. A liter Fig. 5. Wash Bottle.

7 When a number of flasks (or burets) are to be checked, calibrating bulbs
(of the Morse-Blalock or Ostwald type) can be used to advantage. Directions
for their use can be found in reference books on quantitative analysis.

32 VOLUMETRIC METHODS [P. VI

flask is convenient for wash water, smaller flasks for special wash solu-
tions.

11. The solution in the flask should be gently swirled as the water is
added so that it is well mixed before the final adjustment of the meniscus
is made. The solution of a solid or the mixing of solutions may result in an
appreciable volume change.

12. One of the most common sources of error in volumetric analysis is the
incomplete mixing of solutions following their preparation in volumetric
flasks. Such solutions can be effectively mixed by swirling the solution,
inverting, and again swirling, repeating this sequence at least ten times.

13. Solutions which are sensitive to light should be kept in bottles of
brown glass or, if these are not available, in bottles which are covered with
paper or painted with black lacquer.

14. All standard solutions, reagents, and solutions for analysis which are
to be reserved for future use should be carefully labeled, dated, and, if for
general use, initialed by the person preparing them.

P. VI. The Preparation and Standardization of a Thiocyanate Solution

Discussion. As both ammonium and potassium thiocyanatc are
somewhat hygroscopic, are too soluble to be readily purified by re-
crystallization, and cannot be easily dried, standard solutions of these
substances are not usually prepared by direct weighing as was done
with silver nitrate. 8 Solutions of approximately the desired concen-
tration are therefore prepared and then standardized, usually against
pure metallic silver or silver nitrate.

This standardization is accomplished by making use of the reac-
tions involved in the so-called Volhard method for determining
silver; 9 this titration is an example of a volumetric precipitation
method. In carrying out this titration, a definite volume of the
already prepared standard silver nitrate solution is pipeted out,
some ferric nitrate is added as indicator, and the solution is titrated
with the thiocyanate solution. As long as there is an appreciable
amount of silver in the solution being titrated, the concentration
of the thiocyanate remains so small, owing to the precipitation of
silver thiocyanate, that no perceptible amount of the red compounds
which thiocyanate forms with ferric iron (probably Fc(SCN)J" and

1 Kolthoff and Lingane, /. Am. Chem. Soc., 67, 2126 (1935), have prepared
KSCN suitable for use as a primary standard by recrystallization from water,
drying over PxO 5> and finally melting for a short time at 200C. The dried
material is not hygroscopic at a relative humidity of less than 45 per cent but
deliquesces rapidly at relative humidities greater than 50 per cent.

f Charpentier, Bull. Soc. Ing. Civ. France, 135, 325 (1870). Volhard, /. Prakt.
Chem. (2), 9, 217 (1874).

P. VI] THIOCYANATE SOLUTIONS 33

Fe-Fe(SCN)e) are formed. 10 As the amount of thiocyanate added
becomes approximately equivalent to the silver present, further
addition causes its concentration to rise very rapidly; this results
in the formation of perceptible amounts of the red-colored com-
pounds, giving the end-point. 11

This change in the concentration of the silver and of the thio-
cyanate ions during the course of the titration can be predicted from
the mass-action and solubility-product laws. It will be recalled
that the mass-action law states that in the reaction between a mole-
cules of substance A, b molecules of substance B - , to form g
molecules of G and h molecules of H , according to the equation

aA + bB + = gO + hH + . ,

when equilibrium is attained, the following relation will exist between
the molal concentrations of the various molecules, [A], [B] ,
[G], [#], namely:

* [G][H] h - -
-r-v = K (at a given temperature).

LAJ a J#J *

This equation is known as a mass-action or equilibrium expression,
and K is known as the equilibrium constant. 12 The above typo of

10 Schlessinger and Van Falkenburgh, /. Am. Chem. Soc., 63, 1212 (1931),
cite experiments to show that with an excess of thiocyanate Fe(SCN) 6 ta is
formed; with an excess of ferric iron the molecule Fe-Fe(SCN) predominates.

11 Attention is again called to the distinction that should be made between
the equivalence-point, which is the point at which an equivalent amount of the
standard solution has been added, and the end-point, at which the titration is
concluded. The latter is based upon some perceptible evidence, usually
furnished by an indicator, that the desired reaction is completed. The agree-
ment between these two points measures the accuracy of the titration.

12 As an example of the conventions used in formulating mass-action ex-
pressions, consider the reaction previously mentioned (Equation 3, p 25) :

107 + 3H 2 S = I- + 3S ( 8 ) + 3H 2 O.

The complete equilibrium expression for this reaction would be written as
follows:

However, the concentration of the dissolved sulfur would always have a
constant value (if the solution were saturated that is, an equilibrium had
been reached with the solid sulfur) and, therefore, this term may be omitted
from the expression and the value of K changed correspondingly; this omission
is justified in writing mass-action expressions in all cases where the solution is
kept saturated at a definite temperature with a pure solid substance. Next,
the concentration of water in dilute aqueous solutions (0.1 formal or less) does

34 VOLUMETRIC METHODS [P. VI

expression is applicable to equilibria between gases or substances in
solutions, but is subject to increasing deviations as the reactants
become more concentrated, especially in solutions of highly ionized
substances. These deviations will be discussed later (p. 40).
In considering the titration of thiocyanate with silver nitrate solu-
tion, it is seen that as soon as a precipitate of silver thiocyanate is
produced, the equilibrium in the solution can be represented by the
equation

AgSCN ( .) = Ag+ + SON-

(it is assumed that silver thiocyanate, like most other salts, is com-
pletely ionized in dilute solutions), and therefore when equilibrium is
attained (and the solution saturated^ but not supersaturated, with
solid AgSCN), the mass-action expression can be formulated:

[AglfSCN-] - K.

This expression states that in dilute solutions which are in equilibrium
with solid AgSCN the product of the concentration of thfe two ions
has a constant value at any given temperature ; this is a statement of
a general principle applicable to all slightly soluble salts, and the
value of the ion-confeentration product at saturation is commonly
termed the solubility-product constant, and denoted as K QtP% . n
Using this solubility-product principle, the theoretical change in the
thiocyanate and in the silver ion concentrations can be calculated
for the course of a titration, and in Fig. 6 is shown the curve (labeled
Ag" 1 ") obtained by plotting the silver ion concentration, and the curve
(labeled SCN~) obtained by plotting the thiocyanate ion concentra-
tion, as the ratio of the equivalents of thiocyanate added to the
equivalents of silver initially present is gradually increased. (For
convenience it has been assumed that a 0.1 n. solution is titrated
without volume change.) The changes in ionic concentrations

not vary greatly (pure water being about 55.4 f. in HiO at 20); therefore for
dilute solutions and for all except the most precise calculations, its concentra-
tion can be assumed to be constant, and it may also be omitted from the mass-
action expression, the value of K again changing. The equilibrium expression
for the above reaction therefore becomes

- K.

11 It is seen that the form of the solubility-product expression will differ
in considering other types of salts; thus for lead chloride the equilibrium
would be represented by the equation PbCli (8) - Pb++ + 2C1~; and the solu-
bility product would have the form [Pb+^HCl""]* = K*. f .

P. VI]

THIOCYANATE SOLUTIONS

-1
-2
-3

-4

c
o

1-6

3- 7

-9
-10

-11

Ksp=lX10-

0.5 0.6 0.7 0.8 0.9 1.0 1.1 1.2 1.3 1.4
Ratio of Equivalents of Thiocyanate Added
to Silver Present

1.5

Fig. 6. Changes in Silver and Thiocyanate Ion Concentrations When
Titrating Silver Nitrate with Potassium Thiocyanate.

occurring near the equivalence-point during the titration of silver
ion with other ions forming slightly soluble salts are shown in Fig. 7. 14

14 As an example of the method of calculating these values, consider the
case when the ratio of equivalents of thiocyanate to silver has the value 0.90.
For simplicity assume that the volume remains constant throughout the
titration. First, it is necessary to obtain the solubility product of silver
thiocyanate. Upon looking up solubility data it will be found that the solu-
bility of silver thiocyanate at 20C. has been experimentally determined to be
10~ 6 formal ; therefore, assuming complete ionization and neglecting any hydrol-
ysis, the solubility product will be

[Ag+][SCN-]

B.P. = 10- 12 .

If a tenth normal solution of silver nitrate is titrated with thiocyanate and
the thiocyanate added until the ratio of the equivalents of thiocyanate (the
negative ion) to the equivalents of silver nitrate initially present has the value
0.90, the concentration of the excess silver nitrate will be 0.01 (assuming that
there has been no volume change), and the total concentration of the ion will
be 0.01 plus an amount equal to the concentration of the thiocyanate ion re-
maining in solution or to 0.01 plus the molal solubility of silver thiocyanate in a
solution in which the excess of silver ion is 0.01 molal. Substituting the
unknown concentration of the thiocyanate ion in the solubility-product
equation, there is obtained the expression

[Ag+][SCN-] - [0.01 + x][x] = 10- 12
By inspection it is seen that x must be small in comparison to 0.01 so that

VOLUMETRIC METHODS

[P. VI

-2 1

-4

Ratio of Equivalents of Silver Ion to Negative Ion

(During Titration of Negative Ions).
.05 1.04 1.03 1.02 1.01 1.00 0.99 0.98 0.97 0.96 0.95

-5

c
o

o
u

-8
-9

-10

Agl0 3 Ksp=2X10

-AgCI Ksp=lX10" 10

-AgSCN

Ksp=lX10~ 12

Aglp 3

Ksp=2X10- 8

AgCI Ksp~lX10" 10
[Agi

AgSCN Ksp=lX10~ 12

Agl Ksp=lX10' 16

-11
-12
-13

-14

0.95 0.96 0.97 0.98 0.99 1.00 1.01 1.02 1.03 1.04 1.05

Ratio of Equivalents of Negative Ion (Thiocyanate, etc.)

to Silver (During Titration of Silver).

Fig. 7. Concentration Changes of Silver Ion and Various Negative
Ions During Precipitation Titrations.

From the curve for the thiocyanate titration (which expresses the
requirements of the solubility-product equation), it is seen that as
the ratio of the equivalents of thiocyanate to silver is increased (as
the titration with thiocyanate proceeds), the silver ion concentration
decreases and the thiocyanate ion concentration increases until at the
exact equivalence-point they are equal. It is to be noted that the
rate of change of these concentrations with a given change in the

the solution of the equation can be simplified by neglecting the additive term.
It is desirable in solving mass-action equations that the exact expression be
formulated and that methods of simplification then be considered. As will be
shown later, deviations from the mass-action laws of from 1 to 5 per cent are
to be expected even in. dilute solutions; therefore greater exactness than this
in solving such equations is not justified. Upon solving the above equation,
there is obtained a value for x (or for the thiocyanate ion concentration) of
10~ 10 and for the total silver ion concentration of (0.01 + z), or substantially
0.01. Values for other ratios of thiocyanate to silver can be similarly cal-
culated, and the curves illustrated in Figs. 6 and 7 were constructed from such
data.

P. VI] THIOCYANATE SOLUTIONS 37

ratio of equivalents of the ions present increases tremendously when
this ratio approaches unity, or as the equivalence-point is approached.
This rapid rate of change of concentration near the equivalence-point
is of fundamental importance, as by the addition of a small amount
of the standard solution (in this case the thiocyanate) there is pro-
duced such a large change in the ionic concentrations of the solution
that some effect producing an end-point will take place (in this case
the formation of a visible amount of the red compound). As will be
observed later, this rapid change near the equivalence-point is an
essential characteristic of all "titration curves" and one of the fea-
tures which determine the precision of the method.

Consider now in this regard the titration of silver ion with thio-
oyanate ion in the presence of ferric ion. Experiments have shown
that with the ferric nitrate 0.01 formal the color of the red compound
can be detected when the concentration of the thiocyanate ion is
between 1 X 10~ 5 and 2 X 10~ 5 formal. It is seen from Fig. 7 that
at this point the ratio of thiocyanate to silver is between 1.001 and
1.002 and, therefore, results accurate to within 0.2 per cent should be
obtained. 15

From the curves shown in Fig. 7 for the titration of other silver
salts it is seen that the position of the curves and the rate of change
of the concentration of the ions involved near the equivalence-point
is a characteristic of the solubility product of the salt. Thus if a
titration were made of a salt with a solubility product of 2 X 10~ 8
(corresponding to that for silver iodate), the concentration of the
anion would reach 1 X 10~ 6 when the ratio of the equivalents was
only 0.98, and thus an error of as much as 2 per cent might result.
In this regard it is to be remembered that one of the requirements
which was set up for the use of a reaction as the basis of a precise
volumetric method was that it should be quantitatively complete.
Obviously, the more soluble the salt, the less complete its precipita-
tion reaction. As will be observed later, it is a general feature of
titration curves that the less complete the fundamental reaction,
the less pronounced the inflection of the curve near the equivalence-
point. It is also to be pointed out that these curves are characteristic
of other di-ionic salts; for example, the curves for barium sulfate,
Ifgp = 10~ 10 , would coincide with those for AgCl. 16

15 This calculation neglects the amount of the Fe(SCN)? needed to give a
visible color, but experiments have shown this to be of the order of 10~ 6 formal.

16 BaSO 4 is not readily adapted to a volumetric precipitation process because
its solutions tend to supersaturate, and thus the rate at which it attains an
equilibrium is slow.

38 VOLUMETRIC METHODS [P. VI

Summarizing these considerations with respect to the thiocyanate
titration of silver, it is seen that the point at which the color appears,
and the end-point is taken, depends upon the relative solubility of
silver thiocyanate and the stability (or degree of dissociation) of the
colored compound. If silver thiocyanate were much more soluble
or the ferric thiocyanate compound much less ionized, the end-point
would be taken before the equivalence-point; if the ferric thiocyanate
compound were highly dissociated, it would require a larger concen-
tration of thiocyanate to force the equilibrium

+ 6SCN" = Fe(SCN) 6 "

to the right and thus produce a perceptible color; the end-point
would then occur after the equivalence-point. It is also to be seen
from the foregoing discussion that the end-point can be shifted some-
what by varying the concentration of the ferric salt added as indi-
cator; experimentally the relations are such that highly precise
titrations can be made.

Experimental studies of this method have shown that there is a
tendency for the silver thiocyanate precipitate which is first formed
to collect, or, as such a process is termed, to adsorb 17 silver ions on its
surface; therefore, unless the mixture is vigorously stirred, the first
end-point obtained may be premature. In addition, the solution
should be acid in order to repress the partial hydrolysis of the ferric
ion, as the products of this hydrolysis impart a yellow color to the
solution, making the end-point less easily detected. It should be
cold, since ferric ion is more hydrolyzed in hot solutions and the
Fe(SCN)r molecule is more dissociated; in addition, the thiocy-
anate ion is oxidized by nitric acid 'and by ferric ion at elevated
temperatures. 18

Other Indicators for Volumetric Precipitation'Titrations. It has been
seen above that there is an extremely rapid change in the concentra-
tion of both cation and anion near the equivalence-point in a pre-
cipitation reaction involving a slightly soluble precipitate, and, in the
titration just discussed, the sudden increase in the anion (thiocyanate)
concentration caused the appearance of the red color which was taken
as the end-point. Obviously other effects which depend upon the
cation or anion concentration could be used for the same purpose.

17 See "Gravimetric Methods, 11 p. 119, for a discussion of the phenomena of
adsorption.

18 For an experimental study of this method, the article by Kolthoff and
Lingane, /. Am. Chem. Soc. t 67, 2126 (1935), should be consulted.

P. VI] ADSORPTION INDICATORS 39

Thus in P. 95 the end-point of the titration of chloride with silver
ion is obtained by adding a soluble chromate to the solution and
noting the first appearance of a precipitate of silver chromate. As is
seen in Fig. 7, when chloride ion is titrated with silver ion, the silver
ion concentration remains very small until quite close to the equiva-
lence-point, and then it rapidly rises. Because silver chromate is
somewhat more soluble than silver chloride, no precipitation of the
reddish silver chromate occurs until very close to the equivalence-
point, and therefore a precise end-point determination can be made.
A detailed discussion of this end-point is given in P. 95.

Adsorption indicators. It has been found that certain colored
organic compounds tend to be adsorbed (see "Gravimetric Methods,"
p. 119) on the surface of precipitates and that the extent to which this
adsorption takes place is greatly influenced by whether there is an
excess of the precipitating anion or cation in the solution. As an
example, fluorescein, which is a colored organic acid (K A = 10~ 8 ),
is scarcely at all adsorbed on a silver halide or thiocyanate pre-
cipitate so long as there is an excess of halide or thiocyanate ion
present (preferentially adsorbed). However, as soon as an excess
of silver ion is added and the halide or thiocyanate concentration
reduced, the fluorescein is thereupon strongly adsorbed. According
to Kolthoff 19 and his associates, this may be due to an "exchange
adsorption" which may be represented as follows:

Ag+SCN" + Fluorescein~ = Ag+Fluorescein~ + SCN~

(Surface) (Surface)

In addition, upon being adsorbed, the fluorescein undergoes a pro-
nounced change in color, so that this effect can be used to obtain an
end-point. 20

As fluorescein is a very weak acid, it is not useful in other than
practically neutral solutions. Dichlorofluorescein and eosin .are
stronger acids and can be used in hydrogen ion concentrations as large
as 1(T 3 molal.

19 Kolthoff and Larson, J. Am. Chem. Soc., 66, 1881 (1934); Kolthoff and
Rosenblum, J. Am. Chem. Soc., 56, 1264, 1658 (1934).

10 The detailed mechanism of the adsorption process and accompanying
color change has been extensively investigated, especially by Fajans and his
coworkers, who originated and developed the use of these indicators.

Fajans and Hassel, Z. Electrochem., 29, 495 (1923); Fajans and Steiner, Z.
physik. Chem., 126, 309 (1927); Fajans and Wolff, Z. anorg. allgem. Chem., 137,
221 (1924); Hassel, Kolloid Z., 34, 304 (1924); Kolthoff, Kolloid Z., 68, 190
(1934).

For a review of the theories regarding the mechanism of the process and the
factors affecting their application, see Kolthoff, Chem. Rev., 16, 87 (1935).

40 VOLUMETRIC METHODS [P. VI

The end-point with such indicators is likely to be unsatisfactory
with large concentrations of other ions present, since they may be
preferentially adsorbed or cause coagulation of the precipitate before
the end-point.

Because the ferric thiocyanate end-point for the silver-thiocyanate
titration is so precise and can be used in acid solutions as well as in
the presence of even high concentrations of many other anions and
cations, it is a much more generally useful indicator for the titration
of silver ion with thiocyanate than are adsorption indicators. Ad-
sorption indicators are more useful in the titration of chloride with
silver ion and are employed in an optional procedure in P. 95.

For a general discussion of the use of these indicators in other
titrations see Kolthoff and Furman, Volumetric Analysis, Vol. II,
p. 214, and Fajans, Neuere Massanalytische Methoden, Enke, 1935,
pp. 161-207, 194-196.

Limitations to the Use of Concentrations in Mass- Action Expressions.
In the preceding discussion it was stated that there were deviations
from the mass-action law as the reacting substances became more
concentrated, and application of the solubility-product law was
restricted to relatively dilute solutions in predicting the accuracy
of the thiocyanate-silver titration. It has been found that the
mass-action law can be applied to solutions of un-ionized substances
as concentrated as 1 formal with deviations not exceeding a few per
cent, but that with solutions of highly ionized substances the devia-
tions may become greater than this at concentrations as low as 0.01
formal. It is found that the "effective concentration" of a substance,
that is, the concentration which should be substituted in a mass-
action expression in order to express the effect which that substance
exerts on the equilibrium in which it is involved, becomes pro-
gressively less the higher the total ion concentration of the solution.
This effect is now explained on the basis of what is commonly called
the "inter-ionic attraction theory." The principle postulates of this
theory are (1) that most salts and the so-called strong acids and bases
are practically completely ionized and (2) that the "effective concen-
tration," or "activity," of these ions is decreased in solutions because
of the mutual attraction of the opposite charges. The magnitude
of this effect is dependent upon the total ionic concentration of the
solution, the size and number of unit charges on the individual ions,
the dielectric constant of the medium, and the temperature. In order
to obtain the "effective concentration," or, as it is termed, "activity,"
a, to be used in mass-action expressions it is necessary to correct the

P. VI]

INTER-IONIC ATTRACTION EFFECTS

formal concentration, c, by an appropriate factor termed the activity
coefficient , 7; thus a = cy.

The activity coefficients of many substances have been experi-
mentally derived (from vapor pressure, freezing point, and electro-
motive-force measurements), and in order to illustrate their magni-
tude the values for solutions of a few common electrolytes are shown
in Table II. It is 10 be emphasized that these are the activity co-

TABLE II
ACTIVITY COEFFICIENTS OF VARIOUS ELECTROLYTES

Solution

Coefficient

0.001 f.

0.01 f .

0.05 f.

0.1 f.

If.

3f.

HC1

0.98
0.98
0.98
0.97

0.92
0.92
0.92
0.90

0.86
0.84
0.84
0.78

0.81
0.80
0.80
0.72

0.82
0.65
0.75
0.40

1.4
0.70

KOH
KC1

AgNO 3

efficients of the electrolytes in pure solutions of the concentrations
given, and that in the presence of other electrolytes the activity co-
efficients will have other values dependent upon the total ionic con-
centration of the solution and the valence type of the individual ions.
Because of the inter-ionic attraction effects discussed above, it is
unsafe to make quantitative use of the mass-action law unless the
total ionic concentrations are below approximately 0.01 formal or
unless the activities, and not concentrations, are used. In Fig. 8 this
effect is shown by plotting the solubility of thallous chloride against
the equivalent concentration of various added salts. In this same
figure there is shown by means of the dotted line the calculated solu-
bility (assuming complete ionization) in the presence of salts having
a common ion. The increased solubility in the presence of the salts
with no common ion is to be noted, and especially the effect of the
divalent sulfate ion, which, having a molal concentration only half
that of the nitrate, causes a greater solubility increase. The more
pronounced effect on the solubility of salts having divalent ions is
shown in Fig. 9, where the solubility of silver sulfate in various salt
solutions is similarly plotted. It is to be pointed out that thallous
chloride and silver sulfate are moderately soluble salts and that the
deviations due to inter-ionic attraction forces are relatively much
more pronounced than those to be expected in dealing with dilute
solutions containing slightly soluble salts such as silver chloride or
silver thiocyanate.

VOLUMETRIC METHODS

[P. VI

0.030

Calculated Solubility,
Univalent Salts

0.05 0.10 0.15 0.20 0.25 0.30
Equivalent Concentration of Added Electrolyte

Fig. 8. The Solubility of Thallous Chloride in Various Salt Solutions.
(Data from article by A. A. Noyes, /. Am. Chem. Soc., 46, 1107 (1924). Broken
lines are extrapolated values.)

Summarizing, it is seen that the deviations to be expected when
substituting concentrations in mass-action or solubility-product
expressions may be relatively large unless the total ionic concentration
of the solution is small (0.01 formal or less) and unless only univalent
ions are present. Because the calculation of the activity coefficients
of individual ions in complex solutions is too involved to be treated
here, concentrations will be used in mass-action expressions, but it
will be understood that except in extremely dilute solutions of simple
electrolytes the results so obtained are of value only as an approxi-
mation of the effect to be expected.

Procedure VI: PREPARATION AND STANDARDIZATION OF A
THIOCYANATE SOLUTION. Weigh out 8.0 g of ammonium
thiocyanate (or 10.0 g of potassium thiocyanate) on a rough
balance, dissolve it in water, dilute the solution to a liter
(Note 1), transfer it to a clean ground-glass-stoppered
bottle, and thoroughly mix the solution. Label the bottle.

Pipet three 25-ml portions of the silver nitrate solution
into separate 200-ml flasks (Notes 2, 3), add to each of
these 10 ml of chloride-free 6 n. HNO 8 , 3 ml of chloride-free
1 n. Fe(NO) 8 solution, and 25 ml of water. Fill aburet
(Notes 4, 5) with the thiocyanate solution and draw the

P. VI]

THIOCYANATE SOLUTIONS

0.025 0.050 0.075 0.100 0.125 0.150 0.175 0.200
Equivalent Concentration of Added Electrolyte

Fig. 9. The Solubility of Silver Sulfate in Solutions of Various Salts.
(Data from article by A. A. Noyes, /. Am. Chem. Soc., 46, 1107 (1924). Broken
lines are extrapolated values.)

meniscus down to the zero mark (Note 6) ; allow 30 seconds
for drainage, again adjust the meniscus to the mark, and
remove any hanging drop from the tip. Swirl the silver ni-
trate solution and add the thiocyanate to it until the first
perceptible pink color is produced which remains permanent
upon vigorously shaking the mixture for 30 seconds (Note 7).
After allowing time for drainage, read the buret, and imme-
diately record this data in ink in a permanent notebook. Simi-
larly titrate the other portions of silver nitrate. From the
normality of the silver nitrate solution and the volumes of
silver nitrate and thiocyanate used, calculate the normality
of the thiocyanate solution (Note 8).

Notes :

1. As the solution is to be subsequently standardized, this dilution does not
have to be precisely made, and may be done in a 1-liter graduated cylinder.

2. Standard solutions containing nonvolatile compounds should not be
poured from storage bottles, since the solution adhering to the neck of the
bottle evaporates and the residue may fall back into the bottle when the
stopper is withdrawn. If such solutions are to be used frequently, some type
of delivery tube should be arranged. For occasional use the solution may
be removed by means of a clean dry pi pet. When not in use, the tops of

44 VOLUMETRIC METHODS [R VI

the bottles should be protected from dust, and so forth, by inverting over
them small beakers.

3. When measuring a solution with a pipet, the sequence of operations is
as follows: (1) Draw a small portion of the solution into the cleaned pipet
(the pipet should be dry on the outside, should have been carefully drained,
and suction should be applied as it is first immersed in a solution in order to
avoid possible contamination or dilution of standard solutions) and wet the
entire surface of the pipet; discard this solution, and repeat the operation
(the pipet should be again cleaned if streaks or droplets are left on draining).

(2) Draw the liquid into the pipet slightly above the calibration mark.

(3) Wipe off any liquid adhering to the outside of the lower stem. (4)
Holding the pipet vertically, allow the solution to escape very slowly (by
regulating the pressure of the finger on the top of the pipet) until the lower
meniscus of the solution just reaches the calibration mark. (5) While
maintaining the column of liquid in this position carefully touch the tip of
the pipet against a glass surface. (6) Allow the solution to run into the
desired vessel, holding the pipet vertically with the tip sufficiently close to
the receiving surface to prevent spattering. (7) Fifteen seconds after
continuous flow has ceased, lightly touch the tip of the pipet against the wet
side of the receiving vessel. It is not precise technique to blow out the
portion remaining in the tip.

The tip of the pipet should be so constricted by the manufacturer tlrat
the time of outflow is not less than the minimum specified by the United
States Bureau of Standards, 21 or serious errors from drainage effects may
result. The minimum outflow time for 10- and 50-ml pipets is 20 and 30
seconds, respectively.

Before being used for precise work, the calibration of a pipet should be
checked by delivering water of a known temperature from it into a pre-
viously weighed weighing bottle or light-weight conical flask closed by means
of a watch glass. The volume delivered by the pipet can be calculated
from the vacuum weight of water thus obtained and from the relative
density.

4. Burets are of two types (see Fig. 10), those with glass stopcocks
(Geissler), and those fitted with a connection of rubber tubing between the
buret proper and the glass tip (Mohr). The flow is controlled in the latter
type by means of pinch or screw clamps or by glass beads in the tubing. The
Mohr type is not to be used for highly precise work. Such burets are less
expensive and are convenient for use with alkaline solutions (which cause
glass stopcocks to stick if left in contact with them) and for less precise work
with other solutions which do not attack rubber (iodine or permanganate
solutions should not be so used). Glass stopcocks should be lubricated by
means of vaseline, or if there is a tendency to leakage, with stopcock grease,
which can be prepared or purchased. Before applying the grease, the stop-
cock should be dried and only a minimum amount of lubricant applied, or
it is likely to collect in the tip and cause stoppage during a titration.

81 The specifications as to outflow times, tolerances, and so forth, for volu-
metric glassware can be found in United States Bureau of Standards Circular
No. 9, 1916.

P. VI]

THIOCYANATE SOLUTIONS

5. Before being used, the buret
should be cleaned with a buret brush
and soap powder or cleaning solution
and then thoroughly rinsed with
water; it should then drain without
leaving any streaks or droplets.
Before filling it with the thiocyanate
solution it should be rinsed several
times with small portions of the
thiocyanate, care being taken that
the tip is flushed out and that no
air bubbles remain around the stop-
cock.

6. When reading a buret, one
should take care that the line of
sight is perpendicular to the buret.
This can be done by lining up the
front and back calibration mark
(these marks should extend com-
pletely around the buret at each
milliliter) or by encircling the
buret just below the meniscus with
a piece of paper and sighting across
the top of this. With other than
opaque solutions the lowest menis-
cus should be read. As the posi-
tion of this meniscus may vary with
the position of the liquid and the
lighting, it is an aid to darken it by
encircling the buret just below the
meniscus with a strip of darkened
paper or piece of black rubber tub-
ing, while a white card or piece of

s 1

= S

=49

|5Q

Mohr

Geissler

Fig. 10. Types of Burets.

filter paper held back of this makes a contrasting background on which to
view the darkened meniscus. A piece of white cardboard with the lower half
blackened, preferably with india ink, is also quite satisfactory.

If solutions are withdrawn from burets too rapidly, errors m$y result,
since it has been shown 22 that drainage is considerable and continues for a
long period of time. The tip of the buret should be so constructed that a
50-ml buret does not empty in less than 120 seconds; if it is found that the
rate of outflow is greater than this, it should be restricted by the stopcock.
With this rate of delivery the buret can be read immediately with very little
drainage error; however, as an additional safeguard, it is recommended that a
fixed interval of 30 seconds be adopted. Do not allow the buret to stand for
any considerably longer period of time before making the reading, since
errors from leakage or drainage may then become significant.

tf Stott, Volumetric Glassware, Witherby (London), 1928, pp. 109-125.

46 VOLUMETRIC METHODS [P. VI

Before being used for precise measurements, the calibrations of a buret
should be checked as follows: The clean buret is filled with water at the
proper temperature, the water drawn down to within 0.02 to 0.03 ml of the
zero mark, 30 seconds allowed for drainage, and the meniscus then carefully
adjusted to the zero mark. The tip of the buret is then touched with a glass
surface to remove the hanging drop. The water is then slowly drawn out
into a weighed weighing bottle (or light-weight conical flask covered with a
small watch glass) until it is just above the 5-ml (or 10-ml) mark (20 to 30
seconds should be required for the delivery), and after 30 seconds the menis-
cus again exactly adjusted and the tip of the buret touched against the
inside surface of the weighing bottle. The container and water are then
weighed. It is convenient, but not necessary, that the meniscus be stopped
exactly on the unit mark; the buret should be read to 0.01 ml and the water
weighed to 0.01 g. The interval from 5 to 10 ml (or from 10 to 20 ml) is
similarly checked, and so on, until the length of the buret is covered. From
the weights and temperature of the water the corresponding volumes are
calculated. Should any portion of the buret show an abnormal error, it can
be calibrated for each milliliter interval. It is convenient in using the buret
to have a plot of the corrections as abscissae against the volumes as ordinates.

7. When unknown amounts of silver are titrated, the approach to the
end-point is shown by the slower disappearance of the local red color caused by
each drop of thiocyanate. As mentioned in the discussion, this is partly
due to adsorption of silver on the precipitate, and the mixture should be
vigorously shaken so that this may be removed. At this stage of the titra-
tion the thiocyanate should be added by allowing only a fraction of a drop to
form on the tip of the buret, removing this, and adding it to the solution by
means of a stirring rod. Also, when near the en<J of a titration, the inside
walls of the flask, or other titration vessel, should always be washed down by
means of a jet of water from the wash bottle; any of the standard solution,
or of the solution being titrated, adhering thereto is thus recovered. When
very precise measurements are desired, an "end-point correction" should be
made. This is done by titrating the thiocyanate solution into a solution of
the same volume as that at the end-point of the titration, and also containing
the same amount of acid and ferric nitrate, until a color matching that used
for the end-point is obtained. This volume is subtracted from the volume
used in the titration.

8. Properly carried out, these titrations should give results which agree to
within 2 parts in a thousand.

The Applications of Precipitation Methods

Volumetric precipitation methods are used most extensively for
determining the anions which form insoluble silver salts (see P. 114
and P. 141 for a list of these salts). These may be determined
either (1) by direct titration, using suitable indicators, or (2) by
adding an excess of a standard silver solution and titrating this excess
in the presence of the precipitate ; sometimes a preliminary filtration
may be necessary. Other metallic ions that form insoluble salts
with these anions may be determined indirectly. Discussions of

OXIDATION AND REDUCTION 47

these methods will be found in P. 22, "The Estimation of Silver"; P.
95, "The Indirect Estimation of Sodium"; P. 27, "The Indirect
Estimation of Bismuth"; P. 122, "The Estimation of Sulfide"; P.
123, "The Estimation of Cyanide"; and P. 148, "The Estimation
of Chloride." The estimation of silver discussed in P. 23 and
of mercury in P. 45 are examples of methods which should not be
classified as precipitation methods (the principle reactions being car-
ried to completion because of ionization effects) but^vhich because
of their similarity to certain precipitation methods ^re frequently
grouped with them.

OXIDATION AND REDUCTION METHODS OF VOLUMETRIC

ANALYSIS

General Principles of Oxidation and Reduction (Electronic) Reac-

tions

Discussion. Oxidation and reduction reactions may be defined
as those in which there is a transfer of electrons from some ion or
compound (called the reducing agent) to some other ion or compound
(called the oxidizing agent), and thus may be more concisely termed
electron-transfer, or for brevity, electronic reactions. Thus, elec-
trons are readily removed from metallic zinc or from bipositive tin,

Zn = Zn ++ + 2E~

Sn ++ = Sn + + + + + 2E~

and, therefore, metallic zinc and solutions of stannous tin are classed
as reducing agents. On the other hand, chlorine, tripositive cobalt,
and permanganate ion in an acid solution very readily take up
electrons,

2CF = C1 2 + 2E~

+ 4H 2 = MnO7 + 8H+ + 5E~

and are classed as oxidizing agents. Also, there are other reactions of
intermediate nature, such as

Fe + + = Fe 4 " 4 " 4 " + E~~
Ag =Ag+ + E~
H 3 As0 8 + H 2 = H 3 As0 4 + 2H+ + 2E~

48 VOLUMETRIC METHODS

which can be made to proceed in either direction with comparative
ease, depending upon whether these substances are brought into
reaction with more reducing or more oxidizing compounds. As all
of these reactions involve electrons, 28 it would seem that this ten-
dency to give off or take up electrons would give rise to an electrical
potential, and such is the case. Thus if a zinc rod is immersed in a
beaker of water, there is a tendency for the zinc atoms to ionize and
pass into the solution as zinc ions. This process cannot proceed to an
appreciable extent, since the electrons remaining in the zinc develop
an opposing potential on the rod which checks the process, the rod
becoming thereby an electrode. Similarly, if we pass chlorine gas
into a separate beaker of water, there is a tendency for the elementary
chlorine to take up electrons and form chloride ions, but if there are
no substances present with available electrons, the reaction cannot
proceed to an appreciable extent. If an inert conducting metal,
such as platinum, is immersed in the solution, there is a tendency
for the chlorine to draw electrons from it; and this results, as was
the case with the zinc rod, in the production of a potential on the
platinum, which likewise becomes an electrode. This potential will
be opposite in sign to that on the zinc, where there was a tendency for
electrons to accumulate. Thus we have two electrodes with poten-
tials of opposite sign, and it would seem that if they were suitably
connected in an electrical circuit, a flow of electrons, or an electrical
current, would take place. There would be thus produced an elec-
trical cell, and each electrode and the solution of the substances in-
volved in the production of its potential would constitute a half -cell.
If only the two electrodes were connected, by means of any
suitable conductor (usually a copper wire), a flow of electrons would
take place momentarily; however, at the zinc electrode positive ions,
Zn* 4 ", would accumulate; at the platinum electrode negative ions,
Cl", would accumulate, and these accumulated charges would check
the further flow of electrons. If, however, some means of transfer
of these charges between the two solutions is provided, a continuous
current can flow. This is accomplished experimentally by connect-
ing the two solutions by means of a "salt bridge/ 7 which is usually
an inverted U-tube dipping into each solution and filled with a solu-
tion of an electrolyte. This permits the migration of positive ions

11 The definition that oxidation-reduction reactions involve a transfer of
electrons does not exclude non-ionic reactions such as S + Oi SO*, for it
seems reasonable to assume that the valence electrons of the sulfur are more
closely associated with the oxygen atoms in the sulfur dioxide molecule.

OXIDATION AND REDUCTION 49

to the chloride-chlorine half-cell and of negative ions to the zinc-zinc
ion half-cell, If the two electrodes are now connected, there will be
obtained a finite and continuous current flow, electrons passing
from the zinc electrode to the platinum electrode, positive ions being
produced in the zinc-zinc ion half-celf, negative ions in the chloride-
chlorine half-cell, and a transfer of ions taking place through the salt
bridge. The electromotive force, or potential of the complete cell,
will be determined primarily by the potentials arising at the two
electrodes (small potentials are also developed at the junctions of
unlike solutions). These potentials will be determined by the in-
herent tendency of the substances involved to take up or give off
electrons, in this case, of the reactions Zn = Zn* + + 2E~" and 2CF =
C1 2 + 2E~ to take place, and by the concentrations of each of the
substances involved in these half-cell reactions. As any two of the
illustrative electron reactions first mentioned above can be similarly
used to construct a complete cell, it is seen that all of these equations
represent essentially the reactions taking place at suitable electrodes
or are what are termed half-cell reactions. If some one of these half-
cells is assumed to have a definite potential value when the ions in-
volved are at definite concentrations, the value of any other can be
measured by combining it with this reference half-cell and then
experimentally measuring the electromotive force of this completed
cell. The reference half-cell thus used is the so-called hydrogen
electrode,

H 2 (g, 1 atm.) = 2H + (unit activity) + 2E~,

which is arbitrarily given the value zero. The more important of
these electronic reactions with the values of these potentials in volts
when all of the reactants are present in unit activity, that is, when
the gases are at 1 atmosphere and the ions or soluble compounds at
1 molal activity, have been collected in Table II of the Appendix;
and these values are referred to as their molal reduction potentials.

These molal potentials are often of little practical value, and may
be misleading, when one attempts to predict from them the behavior
of oxidizing and reducing agents in the relatively concentrated salt
and acid solutions encountered in analytical chemistry. As an
example, Huey and Tartar 24 pointed out that the value 0.154v,
which they calculated for the molal potential of the reaction

Sn + + = Sn++++ + 2E~,
represents the potential of an electrode in a neutral solution where

" Huey and Tartar, /. Am. Chem. Soc., W, 2586 (1934).

50 VOLUMETRIC METHODS

these ions were at unit activity and where no hydrolysis or complex-
ion formation was involved. Because of the pronounced acidic
character of stannic tin and the impossibility of Obtaining stable
solutions except where complex ions exist, such a molal potential
has very little physical significance. Because of such cases, there
also have been collected with this table of molal potentials the poten-
tial values for certain of these reactions when the concentrations of
the substances involved are 1 formal ; these values are designated as
formal potentials. These- formal potentials could be calculated from
the molal potentials (or the reverse) if there were available adequate
data for the hydrolysis constants, for the dissociation constants of
any complex ions which may be formed, and, finally, for the activity
coefficients of the reactants in these relatively concentrated solu-
tions. Since such data are usually inadequate, it is advantageous to
have the experimentally measured formal potentials.

According to the convention adopted here, a positive sign to the
potential indicates that with respect to the reference hydrogen half-
cell the tendency is for the half-cell reaction to proceed to the right,
that is, to give up electrons; and a negative sign indicates a tendency
for the reaction to proceed to the left.

The Effect of Concentrations on Potentials. The potential at other
than unit pressures or concentrations (activities) can be calculated
by the Nernst equation, which is most useful in the form

E = Eo - lo glo Q at 25C.,

where E is the calculated or observed potential, E Q is the molal
(or formal) potential, and Q is the product of the activities of the
substances on the right-hand side of the half-cell equation divided by
the product of the activities of the substances on the. left-hand side
of that equation, each activity having as an exponent the coefficient
of the substances in the equation; n is the number of electrons
involved in the equation (the number of faradays of electricity
produced by the reaction) ; and 0.059 is a numerical constant which
varies with the absolute temperature, being (X058 at 18C. Quali-
tatively, it is thus seen that the potential of a half-cell is changed
approximately 0.06/n by a tenfold change in the value of Q.

This effect of the concentration on the potential set up by the ions
involved in a half-cell reaction and the use of the above equation is
shown by considering the case of the silver-silver ion half-cell,

Ag = Ag + + E~; E Q = -0.799 v.

OXIDATION AND REDUCTION 51

This moderately large negative value indicates that silver ion should
show oxidizing tendencies and, in fact, should act as a better oxidizing
agent than ferric ion, as the ferrous-ferric potential has the more
positive value, namely,

Fe" 1 "* = Fe""" 1 " + E~; E = -0.782 v.

As a matter of experiment, if silver ion is added in excess to a ferrous
sulfate solution, there will be a partial reduction of the silver ion to
metallic silver, and an equivalent amount of ferric sulfate will be
produced. However, if to a solution of silver ion there is added
thiocyanate until the thiocyanate ion activity is 1 molal, the con-
centration of the silver ion is reduced to 2 X 10~ 12 (Ks.p. AgSCN
being 2 X 10~ 12 ) and the silver-silver ion potential becomes

E = -0.799 - Jog 2 X 10" 12 = -0.109 v.

It is worthy of note that this value should represent the molal
potential for the reaction 25

Ag +.SCN" = AgSCN (i) + E~; E Q = -0.109 v.

It is now seen that metallic silver in the presence of thiocyanate ion
should reduce ferric iron, and a quantitative method has been
proposed for so reducing ferric solutions preliminary to titrating
them with standard permanganate solutions. 26 In this method the
thiocyanate ion concentration is made approximately 0.01 molal,
so that the silver-silver ion potential becomes 0.227 v. As, under
equilibrium conditions, there can be only one potential value obtain-
ing in a given solution, other oxidizing or reducing agents present in
this solution must have their concentrations so changed as to have
this same potential value. Therefore, the ratio of ferric to ferrous
iron can be obtained by substitution in the Nernst equation as
follows:

26 As the result of recent measurements, Pearce and Smith, J. Am. Chem.
Soc., 59, 2063 (1937), have calculated for this potential the value -0.095 v;
the discrepancy probably arises from the value of the solubility product used
above.

26 Edgar and Kemp, /. Am. Chem. Soc., 40, 777 (1918).

52 VOLUMETRIC METHODS

which gives the ratio [Fe^J/IFe"^] the value 3.9 X 1(T 10 . It is
thus seen that only 3.9 X ICT* per cent of the ferric iron should
remain unreduced.

As an example of the use~of this same equation in calculating the
conditions existing at the end-point of an electronic reaction, consider
the calculation of the potential of a solution 10~~ 8 molal in Mn 4 " 4 ",
1 X 10~ 6 molal in Mn07, and 1 molal in H* ; these concentrations
approximate those which would exist at the end of a permanganate
titration. Substituting in the Nernst equation, we have

J, = -,.45 - *! log fi .-,.426.

One can now calculate what the ratio of ferric iron to ferrous iron
would be under these conditions, and thereby predict how completely
ferrous iron would be oxidized when titrated by a permanganate
solution. Again simply substituting

we obtain

m 10

This indicates, as is found experimentally, that the reaction would be
complete well within the usual analytical limits. The possibility
that the rate at which the reaction proceeds is too slow for it to be of
practical use is not precluded.

The Calculation of Equilibrium Constants from Potential Values.
The potential of a complete reaction, that is, of a cell, can be ob-
tained by subtracting two half-cells, 27 thus:

Mn"* + 4H 2 O = MnO7 + 8H+ + 5E~ -1.45

= SFe 1 "*" 1 " + 5E~ -0.782

+ 4H 2 O + 5Fe++ + = MnOJ + 8H+ + 5Fe++ -0.67

The relatively large negative value for the molal cell potential,
EQ = 0.67, indicates that the reaction as written should proceed

17 It should be noted that in adding or subtracting half-cell reactions it is
permissible to multiply the constituents on each side of a half-cell equation
without changing the value of the potential; the potential is independent of
the amount of electricity which is involved.

PERMANGANATE SOLUTIONS 53

quantitatively from right to left. Similarly, by subtracting the
Nernst equations for two half-cells, the potential of the cell reaction
can be obtained in terms of the equilibrium constant for the reaction;
thus in the above case

0.059 .
log

and

Subtracting,

0.059 .

[Mn++][Fe+++) B '

When the reaction reaches equilibrium, EI must equal EZ, and
therefore EI E% = 0, and the potential of the cell (E) is zero.
Therefore the following simple expression is obtained for calculating
the equilibrium constant from the molal potentials:

0-059 [Mn07][H + ] 8 [Fe

0.059 , v
= - log K.
n

In the above case K has the value 10~ 57 . It is thus seen that if
molal potential values are available, predictions can be made as to
the completeness of any reaction it is desired to use. It is to be
emphasized again that such predictions refer to equilibrium conditions
only and that no predictions can be made therefrom as to the rate at
which these equilibrium conditions will be attained.

Permanganate Methods of Volumetric Analysis

Discussion. Standard solutions of permanganate are extensively
used in volumetric oxidation and reduction reactions. This is due,
first, to the large negative value, 1.45 v, of the manganous ion-
permanganate potential, and, as a result, the pronounced oxidizing
tendency of that reaction; second, to the intense color of the com-
pound which enables it to serve as its own indicator; and third, to

VOLUMETRIC METHODS

the fact that, properly prepared and kept, permanganate solutions
are stable over long periods of time.

In the discussion to P. VI, there was given a series of curves
showing the calculated changes in the concentrations of the ions
involved during the titrations of a soluble silver salt with various
anions which form insoluble silver salts. In these curves the pre-
dominant feature was the very rapid change in the concentrations
of the silver ion and of the anion near the equivalence-point, and it
was pointed out that these large concentration changes could be

--10

I -30
j

'-50

-60

0.9800.985 0.990 0.995 1.000 1.005 1.010 1.015 1.020

Ratio of Equivalents of MnO; Added to Fe ++

(titrating with 0.1 n. Solution)

Fig. 11. Changes in Permanganate Ion Concentration During the
Titration of a Ferro*us Salt Solution.

Log of Concentration
(molal) of Fe++

I i i

V> 1 H- I

D en o cn c

II.

1 \

0.980 0.985 0.990 0.995 1.000 1.005 1.010 1.015 1.020
Ratio of Equivalents of MnO; Added to Fe++
(titrating with 0.1 n. Sol.)

Fig. 12. Changes in Ferrous Ion Concentration During
Titration with Permanganate.

PERMANGANATE SOLUTIONS

used to cause some effect which would indicate the end-point of the
reaction. This same principle is illustrated in Fig. 11, where the
changes in the permanganate concentration during the titration of
ferrous iron with permanganate have been calculated and are shown
as a function of the ratio of the two substances present. It is noted
that as long as there is an appreciable amount of ferrous iron present,
the concentration of the permanganate is practically negligible, but
that near the equivalence-point there is a very rapid rise in the
permanganate concentration which causes its color to become visible ;
therefore, the appearance of this color can be taken as the end-point
of the titration. It would also be possible to determine the end-
point by testing for the ferrous ion; for, as is seen by Fig. 12, its
concentration very rapidly decreases to a negligible quantity near
the equivalence-point. 28

1. J\J
1 AH

1.30

&
o
> i 9fi

"c i in

i no

HQO

i -

- -"

}

"'"6.990 0.984 0.988 0.992 0.996 1.000 1.004 1.008 1.012 1.016 1.020
Ratio of the Equivalents of Permanganate Added to the
Equivalents of Ferrous Iron Present

Fig. 13. The Potential Change During the Titration of
Ferrous Ion with Permanganate.

Change of Potential During a Titration and Potentiometric Titrations.
In Fig. 13 there has been plotted for the same reaction, not the

28 This method is used in the so-called Penny method for titrating ferrous
iron with standard dichromate solution (which is not so intensely colored as
to be used as its own indicator) by removing a drop of the titrated solution and
adding it to a drop of potassium ferricyanide solution. As long as there is an
appreciable concentration of ferrous iron present, a blue coloration is obtained.
For the details of the titration see Treadwell-Hall, Analytical Chemistry,
Vol. II, Quantitative, 7th Ed., p. 549 A.

VOLUMETRIC METHODS

Fig. 14. Principles of Potentiometric Titrations. B battery of constant
e.m.f.; LAf uniform slide curve resistance; N sliding contact; R variable
resistance; G galvanometer; X tapping key switch; C calomel electrode;
E platinum-wire indicator electrode.

Explanation of method: Within the broken lines are shown the essential
features of the potentiometer method of measuring an unknown electromotive
force. The battery B is connected through a variable resistance to the ter-
minals of the uniform slide wire resistance LM, and should maintain a constant
potential drop through this resistance; this potential drop can be properly
adjusted for the electromotive force to be measured by the variable resistance
R 9 . The cell whose unknown electromotive force is to be measured is con-
nected as shown. In this case this cell is composed of the reference calomel
electrode and the platinum electrode, which assumes the potential set up
by the constituents in the titrated solution. As thus connected, the unknown
electromotive force tends to send a current through the slide wire in a direction
opposite to that of the battery, and, in making a measurement, the sliding
contact N is adjusted until a position is found at which no current exists in
the galvanometer circuit as evidenced by no deflection of the galvanometer,
0, being obtained upon closing the tapping key, K. At this point the opposing
electromotive forces through the resistance are equal. As the ratio of the
lengths, LN/LMj is equal to the ratio of the resistances of these two lengths
and to the potential drop across them, it is seen t^at if the potential drop

PERMANGANATE SOLUTIONS 57

across LM has been fixed at a definite value, E,, the unknown potential drop,
Ex, across LN can be found; that is,

By this means no current need be drawn from the unknown cell as the measure-
ment is made; this avoids polarization effects, which would result if the cell
were directly connected to a voltmeter.

As the electromotive force of the reference calomel half-cell is constant,
the change during the course of a titration of the measured electromotive
force, E x , represents the change in the potential of the solution in which the
platinum indicator electrode is immersed. By the substitution of a reference
standard cell of known electromotive force in place of the measured cell the
electromotive, E, t can be accurately determined, and thus a measurement of
the actual value of E x can be obtained; this is the basis of the electrical method
of determining the molal and formal potential values.

concentration of the ions, but the existing potential as a function of
the ratio of the equivalents of the oxidizing and reducing substances
present. As would be predicted, since the potential is a function of
the ion concentrations, it is again seen that the rate of change of the
potential is much greater near the equivalence-point. This fact is
the basis of the electrometric method of determining the end-point,
in which the potential set up by the ions or compounds present in the
solution is experimentally measured by suitable means at frequent
intervals during the course of the titration, and a curve similar to
those of Fig. 7 obtained, from which the end-point can be deter-
mined. The principle of the method used in making these potential
measurements is shown in Fig. 14 and the accompanying explanatory
material. A more complete treatment of this subject will be found in
textbooks of quantitative analysis or in special reference books deal-
ing with this subject, such as Kolthoff and Furman, Potentiometric
Titrations, 2nd Ed., Wiley and Sons (1931), and Kolthoff, pH and
Electrotitrations, Wiley and Sons (1931).

The Effect of Reaction Rate on Titrations. It has been mentioned
above that it is not proper to attempt to predict how rapidly a
reaction will proceed from calculations of the equilibrium conditions.
In general it is necessary to determine by experiment the effect which
rate phenomena may have upon a given reaction or method. This
factor of rate is of peculiar importance in permanganate titrations.
Referring to the table of molal potentials, it is seen that there are
three reactions, which might be involved in such titrations, namely,

+ 2H 2 O = Mn07 + *H+ + 3E~ -1.59 (1)
+ 4H 2 = Mn07 + 8H+ + 5E" -1.45 (2)
+ 2H 2 = Mn0 2( .) + 4H+ + 2E~ -1.24 (3)

58 VOLUMETRIC METHODS

It is desired in titrating a reducing agent with a standard permanga-
nate solution that the reaction proceed from right to left according to
Equation 2. and not according to Equation 1, with the resultant
formation of manganese dioxide. As the hydrogen ion concentration
enters to the eighth power in the equilibrium expression for Equation
2, it would be expected that as long as an excess of the reducing
agent was present and an acid solution was maintained, no appre-
ciable amount of manganese dioxide would be formed. However,
as the end-point is approached and a perceptible excess of permanga-
nate ion is added to the manganous ion in the solution, it might be
expected that Reaction 1 would proceed to the left, forcing Reaction 3
to the right according to the reaction

3Mn++ + 2Mn07 + 2H 2 = 5MnO 2(8) + 4H+, (4)

and, by the methods illustrated above, it is calculated that the
equilibrium expression for this reaction has an exceedingly small
constant, thus,

lMn ++ ] 8 [Mn07f = 1Q -3,

By substituting reasonable values for the manganous ion concentra-
tion at the end of a titration and for the permanganate ion con-
centration necessary to give a perceptible end-point, it can be cal-
culated that the hydrogen ion concentration would have to be raised
to an impracticable value in order to prevent the precipitation of
manganese dioxide. Therefore if it were not for the fortunate fact
that the rate of the reaction shown by Equation 4 is very slow, the
permanganate end-point could not be used without the formation
of a manganese dioxide precipitate. Experimentally it is found that
unless the hydrogen ion concentration is kept above approximately
0.5 molal, such a precipitate is likely to form as the end-point is
approached. There are also certain methods which are carried out
in neutral solutions in which the permanganate is reduced only to
hydrated manganese dioxide. Such titrations are to be avoided, if
possible, because of the difficulty of detecting the end-point color in
the presence of the precipitate. Rate effects are also an important
factor in permanganate titrations which are carried out in hydro-
chloric acid solutions; again it would be predicted that chloride
would be oxidized by permanganate, but the rate at which this takes
place is so slow that, by properly controlling the conditions, such

P. VII] PERMANGANATE SOLUTIONS 59

titrations are experimentally possible. An extended discussion of
these effects is given in P. 53-4.

P. VII. The Preparation of a Permanganate Solution

Discussion. It has already been stated that, if properly prepared
and stored, permanganate solutions are stable over long periods of
time; however, it is not practical to prepare standard solutions of
permanganate by direct weighing. Even the best grades com-
mercially obtainable almost invariably contain some manganese
dioxide which has been formed on the surface, and, although it is
possible by elaborate means to prepare a very pure product, the
distilled water used for the solution usually has sufficient reducing
gases or organic material in it to cause the production of some
manganese dioxide upon standing. This manganese dioxide then
serves as a catalytic agent for the further decomposition of the
permanganate, and such solutions rapidly decrease in strength,
especially if they are exposed to light. Therefore, it is customary to
prepare the solution, to heat it, or allow it to stand until all reducing
substances in the water are oxidized, and then to filter the solution
through asbestos or sintered-glass filters into the storage bottle.
Thereafter the solution must be protected from light and from
contact with dust, organic matter, or reducing gases.

Procedure VII : PREPARATION OF A PERMANGANATE SOLU-
TION. Weigh out 3.2 g of the best grade of KMn0 4 obtain-
able (Note 1), transfer it to a large beaker, add 1 liter of
water, and heat the solution to boiling, stirring the mixture
until the crystals have dissolved. Cover the solution with a
clock glass and keep it just boiling for 15 minutes. Allow
the solution to cool (Note 2) and filter it through a layer of
asbestos supported on a wad of glass wool in a glass funnel
(Note 3). Collect the solution in a ground-glass-stoppered
bottle (Note 13, P. V) which has just been cleaned
with cleaning solution, rinsed with distilled water, and
drained. Swirl the solution, without wetting the neck of
the bottle, until it is thoroughly mixed (Note 4).

Notes:

1. Even though the solution is later to be standardized, it is an advantage
to use a good grade of KMn0 4 , as it is less likely to contain substances
which will later cause a slow reduction of the permanganate and the forma-
tion of MnCV

2. If time is available, the solution should be kept hot, preferably on a

60 VOLUMETRIC METHODS [P. VIII

water bath for an hour, or left to stand overnight. In this way slowly
oxidized organic substances are more effectively removed.

3. A sintered-glass filter should be used if available. In any case the
filter should be cleaned with cleaning solution and thoroughly rinsed first
with distilled water and then with a small portion of the permanganate
solution.

4. It is desirable that the neck and stopper of the bottle be not wet with
the permanganate, since this then evaporates, and a deposit of KMnO 4 and
Mn02 results, which on again opening the bottle may fall into the solution.
For this reason, when filling a buret, the solution should not be poured, but a
clean dry pi pet used (see Note 2, P. VI).

P. VIII. The Standardization of Permanganate Solutions

Discussion. Primary standards. A substance which is directly
weighed and used for the standardization of a solution is called a
"primary standard"; for very precise work it is preferable that a
solution be thus directly standardized rather than compared with
another standard solution.

A primary standard should meet certain qualifications, namely:

1. It should be a pure substance. It may be an elementary
substance such as iodine, copper, or iron; or it may be a compound
such as sodium oxalate^ arsenious oxide, or potassium iodate; but
it must be capable of purification, usually by recrystallization or
sublimation, to a definitely known composition.

2. It should be stable. Hydrated substances, while frequently
used, are not preferred, since they are difficult to free of extraneous
moisture and are prone to change their moisture content upon storage.
Compounds subject to surface oxidation (FeSO^NEU^SCVCH^O,
Fe wire) or reduction (KMnO-j) are to be avoided if possible.

3. It should not be hygroscopic. Difficulty in storage and in
weighing is likely with such materials.

4. It should react stoichiometrically. Complex apparatus or
special technique should not be required in order to make the reaction
stoichiometric. A particular standard may be useful for only one
solution of a general class. Thus arsenious oxide is an excellent
reducing agent for the standardization of iodine solutions, but
presents specific difficulties if used for standardizing permanganate.

5. It should have a large equivalent weight. It is desirable that
the weight required for a titration should be so large that the weigh-
ing errors will be small relative to the other errors involved in the
standardization.

The standardization of permanganate solutions against sodium
oxalate. The primary standards which have been most commonly
used for permanganate solutions are ferrous ammonium sulfate

P. VIII] PERMANGANATE SOLUTIONS 61

(FeS04.(NH 4 ) 2 SO 4 -6H 2 0), oxalic acid (H 2 C 2 O 4 .2H,O), iron wire,
and sodium oxalate (Na2C204). Of these, the latter most nearly
fulfills the qualifications listed above, 29 and therefore its use as
a standard for permanganate solutions will be considered in detail.

It was pointed out in the general discussion of volumetric methods
that in order for a given reaction to be used as the basis for a precise
method it must (1) be quantitatively complete, (2) proceed with
reasonable rapidity, (3) be stoichiometric, and (4) there must be
some practical means of determining the end-point.

The first of these requirements is satisfactorily met by sodium
oxalate, since, from an equilibrium point of view, the reaction between
permanganate and oxalate is quantitatively complete. This re-
action, when taking place in an acid solution, is represented by the
equation 30

2MnO7 + 5H 2 C 2 O 4 + 6H + = 2Mn+ + + 10CO 2 + 8H 2 O,

2 * Sodium oxalate for standardization purposes can be obtained from the
United States Bureau of Standards.

30 The principles involved in constructing equations for electronic reactions
are aptly illustrated by the permanganate reactions, which have been con-
sidered. These principles are as follows: First, there must be known the
initial and final oxidation states of the reducing and of the oxidizing agents
involved. From this the unit change of each can be noted; and, as the total
change for the two substances must be equal, the proper coefficient for each
substance can be determined. Following this, if any of the substances in-
volved are oxygen compounds, it is usually necessary to include the proper
amount of hydrogen ion, hydroxyl ion, .or water, as the case may be. As an
illustration, consider the equation

MnO 4 - + 5Fe++ + 8H+ = Mn++ + 5Fe++ f + 4H 2 O.

To construct this it must first be known: (1) that the reaction is to take
place in an acid* solution, (2) that permanganate under such conditions is
reduced to bipositive manganese, and (3) that ferrous ion is oxidised only to
the tripositive ion; the oxidizing and reducing substances can then be properly
balanced. The hydrogen ion needed to combine with the oxygen of the
permanganate is then readily seen.

The reaction between permanganate and oxalate in acid solutions given
above, although showing the initial and final products, contains an example
of a compound, namely, the oxalic acid molecule, where there may be doubt as
to the oxidation state of one or more of the atoms involved. In such a case
recourse is had to the principle that in a compound the sum of the positive
oxidation numbers must be equal to the sum of the negative oxidation numbers
of all the atoms present. Therefore, since the oxidation states of the hydrogen
and oxygen atoms are known, it is seen that the average oxidation state of the
carbon atoms is +3. (There are compounds in which atoms of the same ele-
ment may exist in different oxidation states; in such cases the above rule gives
only the average oxidation state.) As the final product of the oxidation of the
oxalate is carbon dioxide, in which the carbon is quadripositive, two oxidation
equivalents are required for each mole of oxalate. Knowing this fact, the
equation can be balanced.

62 VOLUMETRIC METHODS [P. VIII

and is, moreover, au example of what is called an irreversible re-
action, since it has not been found possible experimentally to cause
the reduction of carbon dioxide to oxalate. The second requirement
is not so satisfactorily fulfilled, since the rate of the reaction between
oxalate ion or oxalic acid and permanganate is somewhat slow at
room temperature, and it is necessary to heat the solution in order to
cause the reaction to proceed to completion at a practical rate. It
will be found experimentally that the first portion of the permanga-
nate added will be decolorized very slowly even at 90C., but that
thereafter it disappears more rapidly. This is due to the fact that
manganous ion catalyzes the reaction; and since it is one of the
products, the reaction proceeds more rapidly as soon as an appre-
ciable amount of the permanganate has been reduced. This is an
example of an auto-catalytic reaction, that is, one which is catalyzed
by one of the products of the reaction.

A small amount of manganous ion could be added before beginning
the titration ; however, this is not advised, as the initial delay is slight,
and after the first one or two drops the reaction proceeds quite
rapidly. The mechanism of this reaction and its catalysis is very
complicated and, although extensively investigated, is not yet entirely
clear. 31

In very general terms it seems that, as stated, the reaction between
oxalate or oxalic acid and permanganate is quite slow, but that if
manganous ion is present, it can be oxidized by permanganate to the
tri- or quadripositive state and that this in turn can rapidly oxidize
the oxalate. This simple picture is complicated by the formation of
complex compounds between the various manganese ions and oxalate
ion, and the possible formation from the oxalate of unstable oxidation
products, such as the ion COjT. These considerations 'are extensively
treated in the references cited.

The question of the stoichiometric nature of the reaction has been
exhaustively studied by McBride 32 and more recently by Fowler and
Bright. 83 In considering the titration, it is evident that errors can
arise from the following sources: (1) As oxalic acid is readily sub-
limed, it is possible that it might be lost from a hot acid solution

11 Skrabal, Z. anorg. allgem. Chem., 42, 1 (1904); Launer, J. Am. Chem. Soc.,
54, 2597 (1932) ; Launer and Yost, /. Am. Chem. Soc., 56, 2571 (1934) ; Fessenden
and Redman, J. Am. Chem. Soc., 57, 2246 (1935); Lidwell and Bell, /. Chem.
Soc., 1935, 1303; Polissar, J. Am. Chem. Soc., 58, 1372 (1936).

" McBride, /. Am. Chem. Soc., 34, 393 (1912).

M Fowler and Bright, J. Research Natl. Bur. of Stnds., 15, 493 (1935).

P. VIII] PERMANGANATE SOLUTIONS 63

during the titration. (2) It is known that oxalate solutions are not
stable when kept in glass for considerable periods of time and that
oxalic acid is decomposed by sulfuric acid in concentrated solutions,
so that an appreciable decomposition of the oxalate might occur.
(3) It has been shown that oxalate solutions in the presence of
manganous ions are slowly oxidized by atmospheric oxygen, and it is
entirely possible that, as is known to happen in other cases, this
reaction might be further induced during the permanganate-oxalate
reaction. (4) As permanganate is known to be unstable in hot acid
solutions, decomposition may occur during the course of the titra-
tion. (5) The oxalate might not be completely oxidized by the
permanganate at the end-point, or products other than carbon
dioxide, such as hydrogen peroxide, might result from the oxidation
reaction. (6) The permanganate might not be completely reduced
to manganous ion at the end-point, compounds of tri- or quadri-
positive manganese remaining in the solution.

In order to determine the existence and magnitude of these possible
effects experiments were made by McBride in which one factor was
varied and, insofar as possible, all the other variables were kept
constant. In this manner the effect of (1) temperature, (2) acidity,
(3) volume of titrated solution, (4) rate of addition of the per-
manganate, (5) rate of stirring of the solution, (6) allowing the
solution to stand at an elevated temperature before titrating, (7) the
absence or presence of oxygen during the titration, and (8) the con-
centration of manganous ion were studied. For the details of these
experiments one should consult the original article, but the results
can be summarized as follows: There was no appreciable volatiliza-
tion, decomposition, or air oxidation of the oxalate in solutions
approximately 1 normal in sulfuric acid even at elevated tempera-
tures. There was no evidence of an appreciable error due to in-
complete oxidation of the oxalate. There was some evidence of error
due to decomposition of the permanganate, with apparent loss of
oxygen, but the magnitude of this effect was small, and, by proper
control of the conditions of the titration, results which were repro-
ducible, that is, precise, to 0.2 per cent were obtained.

As the result of his experiments, McBride recommended that the
titration be carried out slowly, at relatively high temperatures
(60 to 90) and with continuous and vigorous stirring. These
conditions, if closely regulated, gave very reproducible results, but
it was later found that there was a tendency for these values to give

64 VOLUMETRIC METHODS [P. VIII

permanganate liters 34 from 0.1 to even 0.4 per cent too high. Fowler
and Bright (loc. cit.) subsequently investigated this effect and are
inclined to attribute it in part (but not wholly, as it could not be
entirely eliminated by working under oxygen-free conditions) to
peroxide formation, which may be represented as follows:

2 + H 2 C 2 4 = H 2 2 + 2C0 2 . (1)

If the peroxide then decomposes,

2H 2 2 = 2H 2 + 2 , (2)

before it is oxidized by permanganate, too little permanganate is
consumed, and a high normality for the permanganate is obtained;
the slower the rate of titration the greater the opportunity for the
peroxide to decompose. Several investigators have found evidence
of peroxide formation during the permanganate-oxalate reaction,
and Kolthoff 35 has shown that this increases with increase in tem-
perature and manganous ion concentration.

As a result of their investigation of this positive error Fowler and
Bright have found that it can be almost entirely eliminated if most
of the permanganate is added rapidly at room temperature and the
solution then heated to 55 to 60 before obtaining the end-point. In
order to increase the rate of the reaction at room temperature, the
initial concentration of the sulfuric acid present was increased from
about 0.7 n. (as recommended by McBride) to about 1.8 n. An
end-point correction is made by matching the permanganate color in
another beaker containing the same amount of hot water and of acid.
Under these conditions (which are provided for in the following
procedure) the stoichiometric error should not exceed 0.1 per cent,
and if weight burets are used, the instrumental errors should also be
less than that value.

The determination of the end-point of the titration is simplified
because the color of the permanganate ion is so intense that no
additional indicator is required. The end-point is consistent, and,
if the proper end-point correction is made, it is highly precise.

84 The liter of a solution may be defined as the concentration expressed in
grams per milliliter. The titer is usually determined by titration and may
refer to the weight of a given substance contained in, or that will react with, a
milliliter of the solution. This convention may be used to advantage where a
standard solution is to be used primarily for the determination of a single
substance; the titer is then given as the grams of that substance reacting with
1 milliliter of the solution.

" Kolthoff, Z. anal. Chern., 64, 185 (1925).

P. VIII] PERMANGANATE SOLUTIONS 65

Bray 36 found that if the titration is carried out under certain condi-
tions, the end-point correction could not be precisely made by match-
ing the color of the titrated solution, because of the presence of some
incompletely reduced manganese compounds. However, McBride
showed that at temperatures above 60C. and by slow addition of the
permanganate near the end-point this effect could be reduced to
within the usual experimental limits.

Procedure VIII : STANDARDIZATION OF A PERMANGANATE
SOLUTION AGAINST SODIUM OXALATE. Transfer about 1 g
of Na2C2O 4 (Note 1) to a weighing bottle and dry the
material at 110 to 200C. for 1 hour (Note 2). Weigh out
0.25- to 0.30-g portions (Note 3) of the oxalate into each of
three 400-ml beakers. Fill a clean buret with permanga-
nate solution (Note 5, P. VI). Add to the oxalate in the
first beaker 250 ml of 1.8 n. H 2 SO 4 (12.5 ml of 36 n. diluted
to 250 ml) which has been previously boiled for 10 to 15
minutes and then cooled to 27 it 3C. Gently stir until
the oxalate has dissolved; a thermometer may be used to
advantage. Add 90 to 95 per cent of the calculated volume
of the permanganate at a rate of 25 to 35 ml per minute
while stirring slowly (Note 4) and let the solution stand
until the pink color disappears (usually about 45 seconds).
Heat to 55 to 60C. (Note 5), and then continue adding
the permanganate at such a rate that no permanganate
color persists in any part of the solution. As the end-point
is approached, each drop must be allowed to decolorize
before the next is added, and finally part drops should be
transferred from the tip of the buret to the solution with the
stirring rod (or thermometer). Keep the solution between
55 and 60C. The end-point should be taken when the
faintest perceptible pink color which persists is obtained,
viewing the solution against a white background. Imme-
diately after the end-point is obtained, add to 300 ml of
1.8 n. H2S04 (which should have been previously boiled
and then cooled to 55 to 60) sufficient permanganate
to match the color in the titrated solution; subtract this
volume from that used in the titration (Note 6). Similarly
titrate the other two samples and calculate the normality
of the permanganate (Note 7).

W. C. Bray, J. Am. Chem. Soc., 32, 1204 (1910).

66 VOLUMETRIC METHODS [P. IX

Notes:

1. Sodium oxalate obtained from the Bureau of Standards or of a special
grade for standardization purposes should be used.

2. Sodium oxalate can be heated to 240C. without decomposition. It is
not hygroscopic, however, and 1 hour at 1QOC. will usually suffice to dry the
material. It may be weighed without danger of absorption of moisture from
the air.

3. The size of the sample should be such that at least 35 ml of the per-
manganate are used in the titration, thus reducing the percentage error
caused by the uncertainty in reading the buret. The size is limited by the
fact that it is undesirable to have to refill the buret, thus introducing the
uncertainty of two additional buret readings. The approximate value of the
permanganate should be known so that most of the required volume can be
added before heating the solution ; if desired, a preliminary titration can be
run in a hot solution.

J 4. The permanganate should be added directly to the solution and not
allowed to flow down the side of the beaker, since this is likely to result in the
formation of a film of Mn02 on the glass. Any permanganate spattering
onto the side of the vessel should be washed down immediately.

5. This part of the titration is advantageously carried out with the beaker
on an electric hot plate or over a low burner. The buret can be protected
from heat and condensation of moisture by inserting the tip tightly through a
piece of cardboard; the latter then acts as a shield.

6. This correction usually amounts to 0.03 to 0.05 ml of 0.1 n.KMnG 4 .

7. By this method even inexperienced analysts should obtain results that
check to at least 2 parts in 1000.

P. IX. The Preparation and Standardization of Ferrous Sulfate

Solutions

Discussion. The application of permanganate methods can be
extended to the determination of certain oxidizing agents if a standard
solution of a reducing agent which reacts stoichiometrically with
permanganate is available. The process consists in adding an excess
of the reducing agent to the oxidizing agent to be determined and then
titrating the excess with permanganate; the reducing agent most
frequently used for this purpose is ferrous sulfate. As has been
shown in the general discussion of electronic reactions, ferrous iror
and permanganate react quantitatively, and, furthermore, the
reaction is both rapid and stoichiometric.

Neutral ferrous solutions are rapidly oxidized by the oxygen of th(
air, and the ferric ion then hydrolyzes, with the resultant precipita-
tion of ferric hydroxide. Hydrochloric acid solutions of ferrous
salts are so rapidly oxidized as to be unsatisfactory standard solu-
tions, and, in addition, the presence of chloride may be objectionable;
sulfuric acid solutions are so slowly oxidized that they can be satis-

P. IX] FERROUS SULFATE SOLUTIONS 67

factorily used. However, for precise work, ferrous sulfate solutions
should be standardized daily unless they are kept under an inert
atmosphere, such as carbon dioxide or nitrogen.

Procedure IX : PREPARATION OP A FERROUS SULFATE SOLU-
TION. Add 28 g of finely crushed FeS0 4 -7H 2 (or 40 g of
FeS0 4 -(NH 4 ) 2 SO 4 .6H 2 0) to 100 ml of 6 n. H 2 S0 4 and 400
ml of water and stir the mixture until the crystals dissolve.
Dilute the solution to a liter, mix it thoroughly and transfer
it to a ground-glass-stoppered bottle.

STANDARDIZATION OF A FERROUS SULFATE SOLUTION. Pipet
25 ml of the ferrous sulfate solution into a 200-ml flask con-
taining 10 ml of 6 n. H 2 S0 4 and 75 ml of water. Titrate
this solution with the permanganate solution until the
first perceptible pink color is obtained when viewing the
solution against a white background (Note 1).

Notes :

1. The appearance of the pink color is most sensitively detected by com-
paring the titrated solution with an equal volume of acid and water in a
similar flask. After obtaining the end-point, permanganate should be
added to the comparison solution until its color matches that of the titrated
solution. The volume of permanganate thus used should be subtracted from
that used in the titration.

The Applications of Permanganate Methods

Permanganate methods are used in this system of analysis for the
estimation of cadmium (P. 30), zinc (P. 62), strontium (P. 85),
calcium (P. 87), for the optional estimation of iron (P. 53 A),
and for the optional estimation of manganese (P. 72); they are
also used in the analysis of the acidic constituents for the
estimation of ferrocyanide (P. 133) and oxalate (P. 166 and P.
87). Other constituents that are frequently determined by titra-
tions involving permanganate or ferrous sulfate solutions are anti-
mony, manganese, and peroxide.

68 VOLUMETRIC METHODS

IODOMETRIC METHODS OF VOLUMETRIC ANALYSIS

Discussion. General principles. The molal potential of the
reaction 21"" = I 2 + 2E~~ is only 0.53 v, and therefore it would
appear that iodine might be used as an oxidizing agent for a con-
siderable number of compounds which are positive to it in the Table
of Reduction Potentials, and that, similarly, iodide ion might be
used to reduce a considerable number of compounds which are
negative to it in that series. It is because both of these conditions
can be quantitatively realized that iodometric methods are so ex-
tensively used in volumetric analysis. For the first-mentioned type
of reaction a standard solution of iodine is used. The second type
usually involves the addition of an excess of a soluble iodide to the
oxidizing substance to be determined, whereby an equivalent amount
of iodine is set free; this iodine is then titrated with a standard
solution of thiosulfate. 37 This latter reaction may be represented as

and is almost unique, since most other oxidizing agents convert
thiosulfate at least partially to sulfate, and, in fact, some sulfate will
be formed by iodine if the solution is at all alkaline. This is ap-
parently due to the fact that iodine reacts with hydroxyl ion,

I 2 + OH" - HIO + r,

and that the hypoiodous acid thus formed tends to oxidize thiosulfate
to sulfate. Because of this same reaction of iodine with hydroxyl
ion and the subsequent decomposition of the hypoiodite to iodate,

3io~ = 2r + ior,

iodine cannot be used as an oxidizing agent in solutions in which the
hydrogen ion concentration is less than approximately 10~ 9 .

Determination of the end-point of iodometric titrations. Several
methods are available for determining the end-point of iodometric
reactions. First, in the absence of other colored ions, the color of the
iodine itself (or more properly of the tri-iodide ion, 17, which iodine
forms with iodide) can be detected in daylight when its concentration
is as small as 3 X 10~ 6 normal, and therefore it can be titrated to
within one drop of one-tenth normal thiosulfate. Secondly, the
purple to violet color which iodine gives in organic solvents (such as

17 Methods of the first type, involving the use of a standard solution of
iodine, are sometimes designated as "iodimetric" methods, and those of the
second type where iodine is liberated and titrated with thiosulfate as
"iodometric" methods; this convention will not be used here.

P. X] IODINE SOLUTIONS 69

carbon tetrachloride, chloroform, or benzene) is so intense that, by
shaking a small amount of one of these with the solution being titrated
and noting when this color disappears, a very satisfactory, although
not so convenient, end-point can be obtained. In most cases,
another almost unique reaction of iodine is made use of, namely, the
fact that tri-iodide forms with starch an intensely colored blue
compound, the exact nature of which is not certain, although it is
generally thought to be an adsorption compound. This color is so
intense that solutions 1 to 2 X 10~ 6 n. in iodine and at least 4 X 10~ 6 n.
in iodide give an easily visible blue color. The intensity of the color
decreases with less than this concentration of iodide, and is more
pronounced in slightly acid than in neutral solutions; above 25C.
it is markedly decreased. The color is also somewhat dependent
upon the presence of other salts, changing to purple or brownish in
concentrated solutions.

P. X. The Preparation of Iodine Solutions

Discussion. Iodine is so slightly soluble in water, 1.34 X 10" 3
moles per liter at 25C., and has such an appreciable vapor pressure,
that advantage has to be taken of the increased solubility caused by
the presence of iodide in preparing the standard solutions. This
effect is due to the reaction

i, + r = IF,

which tends to proceed to the right, as is shown by the value for the
equilibrium expression, thus:

-^-L = 7 j x 10 2 at 25c.
UaJll J

Experiments have shown that in solutions containing approximately
2 to 3 per cent by weight of potassium iodide, tenth normal solutions
of iodine do not rapidly lose iodine if the proper precautions are
observed. The container should be kept stoppered except when
withdrawing solution, and this should be done as rapidly as possible.
The solution should be protected from dust and from reducing gases.
Also, since the reaction

41- + O 2 + 4H+ = 2I 2 + 2H 2 O

is induced by light, the storage bottle should be covered or kept in a
dark place.

Procedure X: PREPARATION OF AN IODINE SOLUTION.
Weigh out into a liter flask or graduated cylinder 25 g of

70 VOLUMETRIC METHODS [P. XI

potassium iodide (Note 1) and dissolve it in 25 to 30 ml (not
more) of water (Note 2). Weigh into a weighing bottle 12.7
g of iodine (Note 3), add this to the potassium iodide solu-
tion, and shake the mixture until the iodine is completely
dissolved (see Note 2). Slowly add sufficient distilled
water to the solution to make the volume 1 liter, swirling
the solution as the water is added. Transfer the solution
to a ground-glass-stoppered bottle (Note 13, P. V), taking
care that no iodine has remained undissolved.

Notes :

1. The KI should be free from iodate and carbonate. The first would
react with the iodide when the solution was acidified, and the latter would by
hydrolysis cause the solution. to become alkaline, which would tend to con-
vert the iodine into iodate and iodide. Iodate is tested for by dissolving
1 g of the KI in 10 ml of water, and adding 1 ml of 6 n. HjSCV, no iodine
color should develop in 30 seconds. A similar solution, 1 g KI in 10 ml of
water, should not give an alkaline reaction with phenolphthalein.

2. Iodine dissolves very slowly in dilute iodide solutions, and for this
reason the iodine should be first completely dissolved in the concentrated KI
solution. The solution of the iodine in the concentrated KI should not be
rapidly diluted, or iodine will be reprecipitated; this falls to the bottom of
the bottle and may not dissolve for days.

3. The weight of iodine taken need not be exact. Iodine, should be
brought into balance cases only in stoppered vessels; its vapors are corrosive.

P. XI. The Preparation of Starch Solutions

Discussion. The nature of the starch-iodine blue color and the
sensitivity with which it can be detected have been mentioned in the
general discussion of iodometric processes.

The indicator solution can be prepared from potato, arrowroot, or
rice starch. These starches are composed of two main products,
/3-amylose, which is the so-called soluble starch, and a-amylose,
which is quite insoluble. When the relatively insoluble starch
grains are subjected to boiling water, they burst, and upon standing
(or centrifuging) the a-amylose and other insoluble material settles
out and can be eliminated by decanting off the clearer portion of the
solution. This is desirable, since iodine forms with a-amylose red-
dish-colored compounds which are not as readily decolorized as is
the blue compound; cornstarch contains a much higher percentage
of a-amylose and should not be used.

If not kept sterile, starch solutions decompose within a few days,
because of bacterial action, and the decomposition products are

P. XII] IODINE SOLUTIONS 71

likely to consume an appreciable amount of iodine as well as form a
reddish color. Similar products are formed by the hydrolysis of
starch in acid solutions, but this reaction is not sufficiently rapid to
cause trouble in titration unless the acid concentration is greater than
3 to 4 normal or the titration is unduly prolonged. The above
effects are accelerated and the starch is likely to be coagulated if it
is introduced into a concentrated iodine solution ; consequently, the
indicator is always withheld from the titrated solution until the
iodine color becomes uncertain.

Procedure XI: PREPARATION OF A STARCH SOLUTION.
Place 2 g of powdered starch (potato, rice, or arrowroot)
and 20 to 30 ml of cold water in a mortar, rub them into a
thin paste, and pour this slowly into 500 ml of boiling
water, stirring constantly. Heat the mixture just to
boiling for 15 to 20 minutes, allow to stand overnight, and
then decant the clear liquid into 100-ml bottles. Place
these in boiling water, insert soft rubber stoppers loosely,
and heat for 2 hours; finally insert the stoppers firmly and
allow the solution to cool (Notes 1,2).

Notes:

1. The preparation keeps indefinitely in the sterile stoppered bottles, but
decomposes rapidly when opened; this is indicated in use by the appearance
of a reddish or brown color which persists after the disappearance of the blue
color.

2. So-called "soluble starch" may be purchased and is more convenient to
prepare. As specimens vary in sensitivity they should be tested before use.
Solutions are prepared as follows: Mix 1 g of the starch to a thin paste with
cold water and pour into 200 ml of boiling water. Prepare the solution only
as it is needed.

P. XII. The Standardization of Iodine Solutions

Discussion. The reaction between arsenious acid and iodine. The
primary standard most suited for the standardization of an iodine
solution is arsenious oxide. This substance can be purified by
recrystallization from hot 6 n. hydrochloric acid and then by subli-
mation; however, it is more practical to purchase it from the Bureau
of Standards. The method used in the standardization depends
upon the oxidation in a neutral solution of the arsenious acid, formed
in dissolving the arsenious oxide, to arsenic acid by means of a
standard iodine solution. The reaction is of sufficient theoretical
interest to be considered in detail.

72 VOLUMETRIC METHODS [P. XII

The equilibrium constant for the reaction

H,As0 3 + 17 + H*0 = H 8 As0 4 + 3I~ + 2H+ (1)

has the value 7 X 10~ 2 , and, consequently, the reaction can be forced
in either direction by controlling the hydrogen ion concentration.
When an aqueous arsenious acid solution is titrated with iodine, the
reaction is incomplete if the acid formed by the reaction is allowed to
accumulate in the solution; in an alkaline solution iodine reacts with
hydroxyl ion,

17 + OH" = 21" + HIO, (2)

causing an excess of iodine to be consumed before the starch-iodine
end-point is obtained. Therefore, the conditions under which the
reaction can be made the basis of a precise iodometric method are
limited but can be approximately calculated as follows: 38 In order to
calculate the upper limit for the hydrogen ion concentration let it
be assumed (1) that the final volume will be 200 ml, (2) that 5 milli-
equivalents of arsenious acid are to be titrated, (3) that the standard
iodine solution is 0.05 formal in iodine (12), and 0.15 formal in potas-
sium iodide, and (4) that an accuracy of 0.001 per cent is desired;
this small value is arbitrarily taken to insure a reasonable factor of
safety.

Since 50 ml of the iodine solution will be used, the final iodide
concentration will be 0.0625 molal; to give the accuracy desired the
ratio of the concentration of H 3 AsOs to H 3 As0 4 39 will have to be
10~ 5 , and, if the minimum detectable starch-iodine color is used as
the end-point, the concentration of the tri-iodide ion may be as low as
2 X 10~~ 7 molal. (This value was experimentally found by Wash-
burn.) Substituting these values in the equilibrium expression for
Reaction 1, it is found that the maximum hydrogen ion concentration
is 2.4 X 10~ 5 .

The factor limiting the lowest hydrogen ion concentration per-
missible is the reaction

I~ + OH" = HIO + 21".

Since the equilibrium constant of this reaction has been calculated to
be 1.3 X 10 2 , and since the final iodide and tri-iodide ion concentra-
tions are essentially the same as in the above calculation, and since,
in order to obtain the assumed accuracy, not over 0.001 per cent of the

Ji These conditions have been calculated and experimentally studied by
Washburn, J. Am. Chem. Soc., 30, 31 (1908); it is recommended that reference
be made to this article.

lf More exactly, the ratio of HjAsOi to H 8 AsO 4 + H 3 AsO 3 .

P. XII] IODINE SOLUTIONS 73

iodine used can be converted into HIO, the final concentration of this
substance should not exceed 1.25 X 10~ 7 molal. From these values
the hydroxyl ion concentration is calculated to be 1.9 X 10~ 5 , cor-
responding to a hydrogen ion concentration of 5.3 X 10~ 10 molal.
Therefore, from an equilibrium point of view, if the hydrogen ion
concentration is controlled between the limits 2.4 X 10~ 6 and 5.3 X
10~ 10 , it would appear that the reaction could be made the basis of an
accurate volumetric method. The rate of the reaction has not been
considered in these calculations, and experimentally it has been
found that it is better to adjust the hydrogen ion concentration
nearer to the lower limit, since in more acid solutions the rate at
which the iodine is removed by the arsenious acid becomes so slow
as to make the titration quite tedious.

Buffer Solutions. Some means now has to be provided by which
the hydrogen iori concentration of the solution can be automatically
controlled within this favorable range. In any solution containing
an acid and its salt the hydrogen ion concentration is determined by
the ionization constant of the acid and the ratio of the molal con-
centration of the un-ionized acid to that of the salt, thus:

[HA]

= K

[A-]'

Therefore, if, first, an acid with a constant having a value close to
that of the desired hydrogen ion concentration is selected and, sec-
ond, relatively large amounts of both the acid and the salt are pro-
vided, the hydrogen ion concentration of the solution will be fixed
and, furthermore, will be little changed by addition or formation of
either acid or base. Such a solution is said to be "buffered." It is
desirable that the constant of the acid be approximately equal to the
desired hydrogen ion concentration; this permits a ratio of acid to
salt of near unity and provides the most effective buffering action
against either acid or base for a given amount of the buffering mate-
rial. For the present case, acids which will not be oxidized by iodine
or reduced by iodide and having constants near the value desired
(approximately 10~ 7 ) are dihydrophosphoric acid (H 2 PO7, K A =
2 X 10~ 7 ) and carbonic acid (H 2 C0 3 , K A = 3 X 10~ 7 ); these
acids and their salts were experimentally found to be satisfactory by
Washburn (loc. cit.). By adding sodium hydrocarbonate (NaHCO 3 )
to an acid solution, carbonic acid is set free, and since the solubility of
carbon dioxide in water has a fixed value at a given temperature, the
concentration of the carbonic acid (H 2 CO 3 ) in a solution saturated
with carbon dioxide is fixed. Thus by neutralizing an acid solution

74 VOLUMETRIC METHODS [P. XII

with hydrocarbonate, and then adding a definite amount in excess,
the solution is buffered at a value determined by the excess of hydro-
carbonate added. The molal concentration of carbonic acid [H 2 CO 3 ]
in a solution saturated at atmospheric pressure at 20C. with carbon
dioxide is 3.4 X 10~ 2 ; therefore, if a sufficient excess of hydrocafbonate
is added to make its molal concentration about 0.34, the hydrogen
ion concentration would be 3 X 10~ 8 molal. Since during the titra-
tion of 5 milli-equivalents of arsenious acid 5 milli-equivalents of
acid are set free, this would cause the hydrocarbonate concentration
to decrease to 0.315 molal, and the hydrogen ion concentration
would be increased proportionally to 3.2 X 10~ 8 molal, approxi-
mately a 7J per cent change. If the same amount of acid is added to
200 ml of pure water, the hydrogen ion concentration would change
from 10~ 7 to 2.5 X 10~ 2 , a 200,000-fold change. If carbon dioxide
were continuously passed through the hydrocarbonate solution, it
would also be "buffered" against a decrease in the hydrogen ion
concentration, for the addition of 5 milli-equivalents of hydroxyl ion
would result merely in the hydrocarbonate ion concentration increas-
ing to 0.365 molal, and proportionally decreasing the hydrogen ion
concentration to 2.8* X 10~ 8 again only a 7| per cent change as
compared with one of 200,000-fold if the hydroxyl ion were added to a
pure water solution.

It is obvious that in any solution containing a base and its salt
similar considerations apply, thus:

rnir-i v [BOHl
[OH] _K B -_.

Therefore, when it is desired to "buffer" a solution to a basic range, a
base with the proper ionization constant and its salt can be used
effectively.

Buffer solutions are of such great importance in analytical, com-
mercial, and biological processes that they have been discussed at some
length here; the principles relating to them may be summarized
from the practical side by recalling that in order to "buffer" a solu-
tion effectively two important considerations should be observed:
First, an acid (or base) should be selected whose constant is of the
same order of magnitude as that of the hydrogen (or hydroxyl)
ion concentration at which it is desired to maintain the solution,
since for a given total amount of acid and salt the buffering is most
effective when the ratio of acid to salt is unity. Second, the amount
of the buffer material provided must be large in comparison to the
amount of hydrogen or hydroxyl ion which may be formed.

P. XII] IODINE SOLUTIONS 75

In this standardization sodium hydrocarbonate is used as the
buffering agent, and, as has been shown by Washburn, a high order of
precision is attainable.

Procedure XII : THE STANDARDIZATION OF AN IODINE SOLU-
TION AGAINST ARSENIOUS OXIDE. Dry approximately 0.6 g
of As 2 O 3 (Note 1) at 100 to 105C. for 1 hour and weigh
out 0.2-g portions into each of three 500-ml flasks (Note 2).
Dissolve the solid in 2 ml of 6 n. NaOH and 10 ml of water,
warming if necessary (Note 3). Dilute with 2 ml of
water, add 2 drops of phenolphthalein indicator solution,
carefully add 6 n. HC1 until the red color disappears, and
then add 1 ml in excess (Note 4). Cool the solution
(Note 5) and, while inclining the flask, add (in small por-
tions as long as vigorous effervescence occurs) 4 g of solid
NaHCOaj rinse the side of the flask. Add 5 ml of starch
indicator and titrate with the iodine solution (Note 6),
swirling the solution gently but not vigorously shaking it
(Note 7), until the first purple or blue color is obtained
which persists for 30 seconds (Notes 8, 9). Similarly
titrate the remaining portions of As 2 O 3 and calculate the
normality of the iodine solution (Note 10).

Notes:

1. As2O 3 for standardization purposes can be purchased from the Bureau
of Standards, and it is recommended that this be used for highly accurate
work; it is not hygroscopic and need not be dried unless extreme accuracy is
desired or it has been exposed to considerable moisture. Since As2O 8 sub-
limes at somewhat higher temperatures, it can be dried by being left over
sulfuric acid in a desiccator for 12 hours.

Commercial material often contains chloride, sulfide, and water. The
product may be purified by dissolving 40 g in 50 ml of hot 12 n. HC1, adding
50 ml of hot water, filtering out any residue, and then adding 50 ml more
water and cooling in ice. The crystals are filtered out and washed with ice
water until free of chloride. Sublime this product from a casserole onto a
round-bottom flask. 40

2. Weighing out three or more separate samples of arsenious oxide can be
avoided by preparing a stock solution in a volumetric flask and taking aliquot
portions of this with a pipet. A 250-ml flask and a 50-ml pipet are convenient
for this purpose. If they have not been calibrated, a relative calibration is
entirely satisfactory, as in this case it is desired only that an exactly
known fraction of the total be taken. Such a calibration is made as follows:

Clean and dry a 250-ml volumetric flask. (To dry the flask, and
similar apparatus, rinse it with several small portions of alcohol,

40 See Chapin, /. Ind. Eng. Chem., 10, 622 (1918), for a discussion of the
preparation of pure

76 'VOLUMETRIC METHODS [P. XII

invert it, insert a glass tube almost to the bottom, attach the tube
to a vacuum line and draw air through it. Compressed air is
not as desirable, since it is likely to contain oily material picked
up from the pumps.) Carefully deliver 5 pipets of water (Note
3, P. VI) into the flask and mark the position of the meniscus
(a small gummed label may be used).

To prepare the arsenious acid solution precisely weigh out
about 1 g of AsaOa into a small beaker, dissolve it in 5 to 10 ml of
6 n. NaOH, and transfer the solution to the 250-ml flask with
the aid of a funnel. Wash out the beaker with repeated por-
tions of water and dilute the solution to the calibration mark
(see Note 9, P. V). Mix the solution (see Notes 11, 12, P. V),
pipet 50 ml into a 500-ml flask, and treat as directed in the fore-
going procedure. The solution should be used at once, since al-
kaline arsenite solutions are slowly oxidized ; neutral solutions are
quite stable.

3. Arsenious oxide is only moderately soluble in water, and the dry powder
may dissolve slowly even in the sodium hydroxide solution.

4. The additional acid is added in order to saturate the solution with
CO 2 when the hydrocarbonate is added. A solution saturated with this
gas is desired for the purpose mentioned in the discussion.

5. The solution should be cooled to avoid loss of CC>2, to decrease the
tendency for the reaction of the iodine to form hypoiodite, and to avoid the
decrease in the sensitivity of the starch color which takes place above 25C.

6. Iodine solutions should not be used in Mohr-type burets.

7. It is desirable that the solution remain saturated with CC>2 until the
completion of the titration (see the discussion). Because of this, the titra-
tion is carried out in conical flasks, not open beakers, and the solution should
not be shaken more than is necessary to cause mixing. If a tank or generator
of C02 is available, it is more effective to keep a very slow stream of the gas
passing through the solution, or a few drops of hydrochloric acid can be
added just before taking the end-point.

8. Most starch preparations will show a transitory purplish tinge, which
can be taken as the end-point. This color will be caused by a small fraction
of a drop of iodine and should be followed immediately by a clear blue. It
should not be confused with the more stable brown or reddish color given by
decomposed starch solutions.

9. Overrunning the end-point should be avoided by noting the rate of dis-
appearance of the local color caused by each drop of iodine. The inside of
the flask should be rinsed down just before taking the final end-point.
Should the end-point be slightly overrun, the solution can be back-titrated
to the absence of color by means of a standard thiosulfate solution.

A common procedure which allows rapid titration of a solution without
overruning the end-point is to reserve a small fraction of the solution and
add it after a rapid titration to a preliminary end-point. This can be done
by drawing up a small portion of the solution into a long "dropper" and
leaving this in the solution until the preliminary end-point is obtained. The
reserved portion is then expelled, the dropper flushed out with the titrated
solution, the end-point obtained, and the dropper again flushed out..

10. Results obtained by this method should agree to within 0.2 per cent.

P. XIII] THIOSULFATE SOLUTIONS 77

P. XHL The Preparation of Thiosulfate Solutions

Discussion. As has been stated, iodometric methods can be
divided into two general classes: (1) those in which a reducing agent
is oxidized by a standard iodine solution, and (2) those in which an
oxidizing agent is estimated by allowing it to react with an excess of
iodide and then titrating the iodine set free with a standard reducing
agent almost invariably sodium thiosulfate. This substance is used
because it reacts stoichiometrically with iodine and because its
solutions are relatively stable with respect to the oxygen of the air.
This is not true of most other standard reducing solutions, such as
stannous chloride and titanous sulfate, which can be used to reduce
iodine quantitatively.

Anhydrous sodium thiosulfate is too hygroscopic for use as a
primary standard and although the pentahydrate (Na^Oa 5H 2 0)
can be prepared in a high degree of purity by repeated recrystalliza-
tions from water and drying under carefully controlled conditions,
this preparation is not justified because of some uncertainty as to the
stability of the solution, especially when first prepared. Numerous
explanations have been advanced as to the causes of the instability of
thiosulfate solutions. It was originally suggested that, since thio-
sulfate is decomposed by acid, the effect was due to carbon dioxide
present in the water, the reaction being

+ H 2 co 3 = HSO^ + Hcor + s,

and it was assumed that the solution would become stable after the
carbon dioxide had been exhausted. However, since sulfite in
neutral solution will reduce two equivalents of iodine, such solutions
should increase in normality, which is not generally observed.
Furthermore, as hydrosulfurous acid is stronger than carbonic acid,
it would seem that the carbonic acid would be regenerated and
serve merely as a catalyst for the decomposition of the thiosulfate;
as a matter of experiment, thiosulfate solutions have been found to be
relatively stable in the presence of carbon dioxide.

It has also been assumed that traces of metallic ions in the water
catalyzed the air oxidation of thiosulfate. Thus, although the
potential of the reaction

2H 2 O - O 2 + 4H+ + 4E~

would indicate that it should oxidize thiosulfate quite completely,
actually the rate of the reaction is such that it has been found that

78 VOLUMETRIC METHODS [P. XIII

solutions saturated with oxygen are very little affected. 41 However,
if a trace of copper salt is present, it may react with thiosulfate as
follows:

The cuprous ion is then oxidized by oxygen,

+ s0 2 + H 2 = 2Cu + + + 20H

and thus the oxidation of the thiosulfate continues as long as oxygen
is present. 42

More recent investigations 43 have shown that the major cause of
the instability of thiosulfate solutions is sulfur-consuming bacteria,
and that sterile solutions are quite stable; slightly alkaline solutions
(H* concentrations of from 10~ 8 to 10~ 9 ) seem to inhibit this bac-
terial action. Numerous experiments have confirmed the fact that
0.1 normal solutions which are prepared from freshly boiled water
to which has been added 100 mg of sodium carbonate per liter and
which are protected from bacterial infection are subject to very little
change and may be used as soon as prepared.

It is assumed in iodometric methods that the equation

i 2 + 2s 2 or - 2r + s 4 or

represents quantitatively the reaction between thiosulfate and
iodine. However, under certain conditions sulfate may be formed,

4I 2 + SsOr + 100H" = 2SO: + SI" + 5H 2 0,

and, as is indicated, its formation is favored by hydroxyl ion. It is
generally assumed that the mechanism of this reaction involves the
formation of hypoiodite, according to the reaction

I 2 + QH~ = HIO + r,

since hypoiodite is known to oxidize thiosulfate to sulfate, as follows:
4HIO + S*0r + 60H~ = 2S07 + 4F + 5H 2 0.

41 Kilpatrick and Kilpatrick, J. Am. Chem. Soc., 46, 2132 (1923).

42 This is an example of u homogeneous catalysis/' inasmuch as only one
phase is present. Heterogeneous catalysis is usually assumed to take place
because of an increase in the activity of the reactants at some surface.

4J Kilpatrick and Kilpatrick, loc. cit.} Mayr, Z. anal. Chem., 68, 274 (1926).

P. XIII] THIOSULFATE SOLUTIONS 79

Experimentally, it is found that acid or even neutral solutions, 0.1 n.
in iodine, may be titrated with thiosulfate or into thiosulfate pre-
cisely, the order of the titration not affecting the results. With more
dilute iodine solutions, where the hydrolysis into hypoiodite is more
pronounced, and with solutions which may be slightly basic, con-
taining, for example, an excess of bicarbonate and not being saturated
with C02, the iodine should be titrated into the thiosulfate* When
titrated in this order, the rate of reduction of the iodine is so rapid
that hydrolysis and subsequent sulfate formation does not take place
to an appreciable extent. Under no conditions should the hydrogen
ion concentration be less than 10~ 9 during an iodometric titration.

It is an easily verified fact that the addition of a thiosulfate solution
to even a dilute acid solution will cause decomposition of the thio-
sulfate with precipitation of sulfur,

SsOr + 2H+ = H 2 S0 3 + S (8) ,

the precipitate of sulfur appearing after an interval of time which
depends upon the concentration of the acid. Since this H 2 S0 3 may
react with iodine or may escape from the solution as S02, in one
case increasing and in the other case decreasing the iodine equiva-
lence of the solution, it is proper to question the use of thiosulfate
in acid solutions. As a matter of experiment, thiosulfate solutions
may be titrated into iodine solutions which are as concentrated in
acid as 3 to 4 n. if the solution is effectively stirred during the addi-
tion of the thiosulfate; the rate of oxidation of the thiosulfate by
iodine is so rapid that no appreciable decomposition of the thio-
sulfate takes place. Thiosulfate solutions should not be acidified
before being titrated. Practically, in the vast majority of cases,
the thiosulfate will be titrated into 'acid solutions of iodine. Since
sodium carbonate (or sodium hydroxide) is often added to thiosulfato
solutions, it is preferable to acidify neutral iodine solutions slightly
before titrating them.

Procedure XIII: PREPARATION OF A THIOSULFATE SOLU-
TION. Boil somewhat more than 1 liter of distilled water in
a flask for 5 minutes, cover the flask, and allow the water
to cool. Weigh out 25 g of Na2S 2 03-5H 2 and 0.1 g of
Na^COs, dissolve them in the freshly boiled water, and
dilute the solution in a graduated cylinder to 1 liter. Trans-
fer the solution to a clean ground-glass-stoppered bottle
(Note 1).

VOLUMETRIC METHODS

[P. XIV

Notes:

1. If a large volume of solu-
tion is prepared, the storage
bottle should be fitted with a
siphon tube and the inlet tube
provided with a soda-lime tube
packed with sterile cotton. Since
most of the solution may be lost
should the rubber or glass stop-
cock of the siphon tube begin to
leak, a convenient modification
of this arrangement is shown in
Fig. 15.

A is an atomizer bulb; B is a
soda-lime bulb packed with sterile
cotton; C is a relief tube fitted
with a pinch clamp; D is a guard
tube fitted over the outlet and
containing a small volume of the
standard solution in order to pre-
vent evaporation. By removing
D and pumping with A, solu-

Fig. 15. Storage Bottle for Standard km * s forced out . ^ needed '
Solutions. The now may be quickly stopped

by releasing the pinch clamp at C.

The relief tube may be eliminated by providing a stopcock on the outlet
tube as shown.

P. XIV. The Standardization of Thiosulfate Solutions

Discussion. Various methods are available for the standardiza-
tion of thiosulfate solutions. Since those solutions are always ti-
trated against iodine solutions, the fundamental primary standard is
iodine. While pure iodine can be readily prepared, it has such a
relatively high vapor pressure that the weighing and transferring of
it to the solution for titration requires a special technique ; therefore
this method is not generally recommended. Recourse is made to
other primary standards, oxidizing agents, which when treated
with an excess of iodide are capable of producing an equivalent
amount of iodine in solution. Among these should be mentioned
potassium iodate, potassium bromate, and potassium dichromate.
While the first two are capable of very exact results, potassium di-
chromate is more commonly available, has a higher equivalent
weight, and, as its solutions are useful for other purposes, is given
preference here.

P. XIV] THIOSULFATE SOLUTIONS 81

The reaction between iodide and dichromate, to give iodine and
tripositive chromium, illustrates several of the factors which have
to be considered in developing iodometric methods. In an alkaline
solution iodine will oxidize tripositive chromium to chromate, while
in an acid solution (10~ 4 molal or greater in hydrogen ion) iodide is
quantitatively oxidized by chromate. However, at this low acid
concentration, the rate of the oxidation is too slow for it to be of
practical use. Experimentally, it has been found 44 that if the hydro-
gen ion concentration is greater than 0.2 molal and the iodide con-
centration greater than 0.05 molal, the oxidation is quantitatively
complete in 5 minutes.

If the hydrogen ion concentration is much greater than 0.4 molal,
another effect one which has to be considered in all iodometric
titrations made in acid solutions becomes appreciable, namely, the
so-called "oxygen error/' As was true with thiosulfate, the poten-
tial of the oxygen half-cell is such that oxygen would be expected to
oxidize iodide quite completely according to the reaction

2 + 4F + 4H+ = 2I 2 + 2H 2 0.

However, the rate of this reaction is so slow that iodide solutions
with the hydrogen ion concentration as high as 0.4 molal can stand
in contact with the air for periods up to 10 minutes without appreci-
able oxidation. Unfortunately, the reaction is induced by the pres-
ence of light and catalyzed by various ions, such as those of copper.
Furthermore, the reaction may be caused to proceed more rapidly,
or, as it is commonly called, "induced/' during the course of another
oxidation-reduction reaction. In the case of the dichromate-iodide
reaction, it is found that if the hydrogen ion concentration is greater
than about 0.4 molal, the above oxygen-iodide-hydrogen ion reac-
tion is caused to proceed to an appreciable extent, is "induced,"
during the reaction between the iodide and chromate. It was
experimentally found by Vosburgh and by Bray and Miller (loc. cit.)
that this effect varied with the iodide concentration and even with
the order of mixing of the iodide and dichromato, being greater
when the dichromate was added to the iodide and acid than when
the iodide was added to the dichromate and acid; these articles
should be consulted for further details.

It is obvious that thiosulfate solutions can also be standardized
against previously standardized iodine solutions, or, by a method

" Vosburgh, /. Am. Chem. Soc., 44, 2120 (1922); Bray and Miller, ibid., 46,
2204 (1924).

82 VOLUMETRIC METHODS [P. XIV

similar to the procedure below, against standard solutions of potas-
sium permanganate, these latter serving as "secondary standards."
In general, secondary standards are not recommended, since they
involve the possibility of error in an additional set of volumetric
measurements and in changes in the standard solutions. However,
the accuracy with which the standardization of permanganate
against oxalate and of thiosulfate against permanganate can be
carried out makes this an exceptional case, and if a recently stand-
ardized permanganate solution is available, the saving in time will
justify its use for all except the most exact work. Because of these
facts, a brief procedure for the method is included in the notes.
Bray and Miller have studied the method and shown that results
obtained by it agree with other methods of standardization to within
0.1 per cent.

Procedure XIV: THE STANDARDIZATION OF A THIOSUL-
FATE SOLUTION AGAINST POTASSIUM DICHROMATE. Precisely

weigh out about 2 g of dry K 2 Cr 2 O 7 (Note 1), transfer
to a 500-ml volumetric flask which has been calibrated
against a 50-ml pipet (Note 2, P. XII), dissolve the crys-
tals, and dilute with water to the mark. Mix the solution.
Pipet 50 ml of the K 2 Cr 2 O7 solution into a 500-ml flask.
Dissolve 3 g of KI in 50 ml of water, add to it 5 ml of 6 n.
HC1 (Note 2), and immediately add this to the dichromate
solution. Gently swirl the solution once (Note 3), then
close the flask with a watch glass, and allow it to stand in a
dark place for 5 minutes. Dilute the solution with 300 ml
of water, washing down the sides of the flask, and, while
swirling the solution slowly, titrate with the thiosulfate solu-
tion. When the yellow color of the iodine becomes indis-
tinct (Note 4), add 5 ml of starch solution and slowly titrate
to the disappearance of the blue color of the starch. Re-
peat the titration with two other portions of the E^CraO?
solution and calculate the normality of the thiosulfate solu-
tion (Note 5).

Notes:

1. As the best grades of K 2 Cr 2 0? often contain Cr0 3 , the salt should be
purified by recrystallization from distilled water, the first portion of the
crystals to appear being discarded; contamination from dust should be
carefully avoided. The crystals should be dried to constant weight at
180 to 200C.

P. XIV] THIOSULFATE SOLUTIONS 83

2. On adding the HC1 to the iodide, no iodine color should develop;
otherwise there is indicated the presence of iodate or, rarely, of oxidizing
agents, usually chlorine, in the acid. If this cannot be avoided, a correction
should be made by dissolving another portion of KI in the same volume of
water and acid and titrating it with thiosulfate. This is preferable to reduc-
ing the iodine with thiosulfate before adding the solution to the dichromate,
as a slight excess of thiosulfate may be added, and as this, as well as the
tetrathionate, may be oxidized to sulfate by the dichromate. The acidified
iodide solution should not be allowed to stand for any prolonged length of
time, or oxidation by the air will take place.

3. Iodine solutions should not be violently shaken in conical flasks or
allowed to stand in open vessels, or loss of iodine vapor will occur. Thus,
it has been found 45 that upon swirling gently for 1 minute approximately
30 ml of 0.1 n. iodine in 4 per cent KI solution in a 250-ml flask that about
0.2 per cent of the iodine was lost. Stoppering the flask did not ap-
preciably reduce the loss. Upon leaving 50 ml of the same solution in an
open beaker for 15 minutes, a loss of 0.9 per cent was found.

4. The titration should be continued without starch until the yellowish
tinge of the iodine can no longer be detected with certainty. This is some-
what difficult in this case because of the green color due to the chromic ion
present, but this color is so reduced by the dilution of the solution that there
is no excuse for adding the starch until within 0.2 to 0.3 ml of the end-point.
It is to be noted also that the final end-point will be a change from the starch
blue to a green; however, the change caused by even half a drop of 0.1 n.
thiosulfate is so distinct that it can be easily detected.

5. The thiosulfate solution can be standardized against a previously
standardized iodine solution as follows :

Pipet 25 ml of the iodine solution into a large conical flask con-
taining 1 g of KI dissolved in 100 ml of water and then add to the
solution 1 ml of 6 n. HC 2 H 3 02. Titrate with the thiosulfate solution
until the iodine color becomes indistinct and then add 5 ml of starch
and titrate slowly, with part drops, to disappearance of the color.
The iodide is added to prevent possible loss of iodine from the dilute
solution ; the acetic acid is added because of the sodium carbonate present in
the standard sodium thiosulfate solution.

The thiosulfate solution can also be standardized against the standard
permanganate solution as follows:

Pipet 25 ml of the permanganate solution into 400 ml of water
containing 3 g of KI and 2 ml of 6 n. H2S0 4 . After 1 minute
titrate the solution with thiosulfate as directed in the above titra-
tion of the iodine solution.

The equilibrium and rate of the above reaction are so favorable that precise
results are obtained as long as sufficient acid is added to provide for that used
in the reaction; as much as 60 milli-equivalents may be present without
harmful results. 46 It is obvious that the standardization of the thiosulfate
by the three methods given should give consistent results, and if the solutions
are available it is recommended that this check be made.

48 Rice, Kilpatrick, and Lemkin, J. Am. Chem. Soc., 45, 1362 (1923)
46 Bray and Miller, loc. cit.

84 VOLUMETRIC METHODS

The Applications of lodometric Methods

lodometric methods are used in this system of analysis for the
estimation of lead (P. 25), copper (P. 28), arsenic (P. 43), antimony
(P. 47), tin (P. 49), iron (P. 53), cobalt (P. 66), manganese (P. 72),
chromium (P. 75), barium (P. 83), magnesium (P. 89), ferricyanide
(P. 134), iodide (P. 144), bromide (P. 147), and arsenate (P. 163).
Other constituents that are often determined by iodometric methods
are the elementary halogens and their oxygen acids (except per-
chlorate), peroxide, sulfide, and sulfite.

The Use of Standard Solutions of Other Oxidizing Agents in Vol-
umetric Analysis

Potassium Bichromate. Standard solutions of potassium dichrom-
ate offer the advantages that they can be directly prepared by weigh-
ing the salt and diluting to volume, and that they are extremely
stable. Bichromate solutions formerly were extensively used for
the titration of ferrous salts in hydrochloric acid solutions, where
the titration with permanganate is troublesome. Bichromate solu-
tions have the disadvantage that the color of the chromate is not
sufficiently intense for it to be used to obtain the end-point of the
titration (especially in the presence of the green chromic ion which
is the reduction product), and, therefore, some form of an indicator
has to be used, or the end-point obtained potentiometrically.

Ceric Sulfate. The use of eerie sulfate, Ce (804)2, solutions has
been developed somewhat recently by Willard and Young 47 and by
Furman and various associates. 48 The solutions are readily pre-
pared by dissolving eerie ammonium sulfate, which can be obtained
commercially, in dilute sulfuric acid solutions; however, the solid is
not sufficiently pure for the direct preparation of standard solutions.
The Ce 111 = Ce IV + E~ potential has been determined by Kunz 49

* 7 Willard and Young, /. Am. Chem. Soc., 60, 1322, 1334, 1368, 1376, 1379
(1928); 61, 139, 149 (1929); 62, 36, 132, 553, 557 (1930).

48 Furman et al, J. Am. Chem. Soc., 60, 755, 1675, (1928); 61, 1128, 1449
(1929); 62, 1443, 2347 (1930); 63, 1283, 2561 (1931).

49 Kunz, /. Am. Chem. Soc., 63, 98 (1931).

OXIDATION-REDUCTION INDICATORS 85

to be 1.44 v in sulfuric acid solutions, which places it between per-
manganate and dichromate as an oxidizing agent. Ceric sulfate so-
lutions offer the advantages that they can be used for the titration of
ferrous salts in hydrochloric acid solutions, that they are relatively
stable, that there are no intermediate oxidation stages (as with per-
manganate), and that the cerous salts give colorless solutions. Ceric
salt solutions are somewhat more intensely orange-yellow colored than
are dichromate solutions of the same normal concentration, but the
color is not sufficiently intense for it to be used for determining the
end-point in highly precise ti* r ations. Details of the preparation
and use of standard solutions ui eerie sulfate can be found in the
newer text and reference books on quantitative analysis. 60

Potassium lodate. Standard iodate solutions are used almost
exclusively with the iodine monochloride end-point, which is dis-
cussed below in connection with the subject of potential indicators.

The Use of Oxidation-Reduction (Potential) Indicators in Volumetric
Analysis

Discussion. As has been stated, solutions of dichromate and
eerie salts are not sufficiently colored to serve as their own indicators,
and this fact formerly required the use of outside indicators or of
potentiometric titrations when these solutions were employed.
To eliminate this disadvantage, the use of oxidation-reduction, or
potential, indicators has been developed. A potential indicator
may be defined as any substance which changes from a colorless to a
colored form, or from one color to another, when it is oxidized or
reduced.

The theory of these indicators can be illustrated by considering
first the use of an inorganic compound in such a capacity. If iodine
in hydrochloric acid solutions is treated with an oxidizing agent, a
reaction which can be represented as follows takes place:

I 2 + 2CP = 2IC1 + 2E~. (1)

It has been shown that, in hydrochloric acid solutions, iodine mono-
chloride is largely combined with chloride ion as follows: 51

_ IC1 + Cl" = ICir. (2)

80 The G. Frederick Smith Chemical Company, 867 McKinley Avenue,
Columbus, Ohio, has compiled a booklet on the applications of eerie sulfate
solutions to volumetric analysis which also contains a complete bibliography
of the literature on this subject.

" Philbrick, /. Am. Chem. Soc., 56, 1257 (1934).

86 VOLUMETRIC METHODS

Therefore the equation for the reaction is more adequately ex-
pressed as

I 2 + 4C1~~ = 2IC12- + 2E, (3)

and the value of the formal potential in 4 f. hydrochloric acid has
been found to have the value 1.05 v. 52 As has been mentioned
previously, very low concentrations of iodine (in the absence of
iodide, less than 10~ 5 f.) can be detected in aqueous solutions by the
use of organic solvents (such as carbon tetrachloride) ; also, iodine
monochloride in hydrochloric acid solutions has only a faint yellow-
ish color and is not extracted by these solvents from hydrochloric
acid solutions. Therefore, when the reaction represented by Equa-
tion 3 takes place, the iodine color disappears from the solution
(or the carbon tetrachloride) and only a very pale yellowish color
remains. Now suppose that one desired to titrate a solution of
ferrous iron in 4 f. hydrochloric acid under which conditions it
would be difficult to obtain a permanent permanganate end-point
because of the reduction of permanganate by this high concentration
of hydrochloric acid. The iodine-iodine monochloride potential
given above indicates that if a small amount of an iodine mono-
chloride solution were added to the ferrous solution, it would be
largely reduced to iodine, which would impart its characteristic color
to the solution. The difference between the formal iodine-iodine
monochloride potential ( 1.05 v in 4 f. HC1) and the formal ferrous-
ferric potential ( 0.70 v in 1 f . HC1 and probably more positive in
4 f. HC1) indicates that if the solution were now titrated with per-
manganate (or, in fact, with any other powerful oxidizing agent),
substantially all of the ferrous iron would be oxidized to ferric before
a large fraction of the iodine was re-oxidized to iodine monochloride;
also that the disappearance of the iodine color could be taken to indi-
cate the completion of the titration.

The reactions taking place if permanganate is used are as follows :

(1) Addition of the iodine monochloride to the ferrous solution:

2Fe ++ + 2IC1F = 2Fe+ + + + I 2 + 4CF.

(2) Oxidation of the remaining ferrous iron by the permanganate :

MnO7 + 5Fe++ + 8H+ = Mn++ + 5Fe w + 4H 2 O.

(3) Oxidation of the iodine to iodine monochloride by the perman-
ganate :

+ 5I 2 + 16H+ + 20CP = 2Mn++ + 1QICI* + 8H 2 O.

11 Unpublished experiments by J. S. Edwards.

OXIDATION-REDUCTION INDICATORS 87

It is to be noted that, since the indicator is added as iodine mono-
chloride and exists in that form at the end-point, the titration is
independent of the amount of indicator added; if the indicator were
added as iodine, the amount added would have to be so small as to
be negligible in comparison with the permanganate used in the
titration, or a correction for the amount added would be necessary.

The agreement of the end-point with the equivalence-point ob-
tained by using this indicator can be predicted by means of the
following calculations: When within 0.2 per cent of the equivalence-
point, the ratio of the equivalents of permanganate added to the
ferrous iron present has the value 0.998, and the ratio of ferric iron
to ferrous iron will be 0.998/0.002. By use of the Nernst equation
the potential existing at that point in the titration can be calculated
as follows: Since the titration is made in 4 f. hydrochloric acid, the
formal ferrous-ferric potential in 4 f. hydrochloric acid should be
used for the EQ value; since this is not available, the formal potential
in 1 f. HC1 ( 0.700 v) will give more nearly correct values than the
molal potential ( 0.782). Therefore the desired potential is

[Fpl

= E - 0.059 log [ -~ = -0.700 - 0.059 log = -0.859.

If the indicator is added in such an amount as to make the solution
initially 10~~ 4 f. in iodine (which is more than the amount needed to
detect the iodine color easily), there can be calculated from the
equation

" 2 ~ 6 [I 2 ][Cl-t

that the molal ratio of the concentration of iodine to iodine mono-
chloride will be 0.87 (assuming the solution to be 4 f. in HC1). There-
fore the iodine color will still be plainly visible in the solution. When
the ratio of permanganate to ferrous iron is 1.002, the molal ratio
of permanganate to manganous ion will be 0.002/1.000, and if the
hydrogen ion concentration is assumed to be 4 molal, calculations
similar to those above will show that the potential existing is ap-
proximately 1.50 v, and that then the ratio of iodine to iodine
monochloride will be 4.8 X 10~ 22 . It is thus seen that 0.2 per cent
before the equivalence-point of the titration the indicator exists
largely as iodine, but that before 0.2 per cent excess of permanganate
has been added, it has been substantially all converted to the iodine
monochloride. Therefore a sharp and precise end-point should be
obtained.

88 VOLUMETRIC METHODS

This so-called iodine monochloride end-point is the basis of an
extensive series of methods in which the titrations are carried out in
2 to 6 f . hydrochloric acid, and various reducing substances are ti-
trated with standard oxidizing solutions, such as potassium iodate,
permanganate, or eerie sulfate. 63

It should also be noted that iodine could be used as a potential
indicator for the titration of a strong reducing agent, such as stan-
nous tin, with standard oxidizing solutions. In this case, the small
amount of iodine added as the indicator would be reduced to iodide
by the stannous tin, and the end-point would be obtained when the
remainder of the stannous tin has been oxidized by the standard
oxidizing solution and the appearance of iodine color is observed.
Thus iodine (or its compounds) can be used as a potential indicator
for two different potential ranges, one corresponding to the oxidation
of iodide ion to iodine and the other to the oxidation of iodine to
iodine monochloride.

With the exception of the iodine monochloride end-point, inorganic
substances have not been extensively used as potential indicators.
However, there are many organic compounds which are capable of
being oxidized and reduced and with which processes there is asso-
ciated an appearance, disappearance, or change of color; such com-
pounds are therefore capable of being used as potential indicators.
Many such compounds are not of practical value because (1) the
color "change is not sufficiently pronounced; (2) they are often so
easily oxidized as to be of no value for most titrations; and (3) the
oxidation-reduction reaction may not be readily reversible, one of the
oxidation stages reacting only slowly or being further changed by
excess of the titrating reagent.

Although these organic compounds are usually of rather com-
plicated structure, the reaction taking place can be represented as
In Red = Incfc + nE~,

and for a large majority of the organic compounds so used n is 2.
Therefore if the value of the molal potential of the indicator is
known, the value of the ratio of the oxidized to the reduced form
can be calculated for any point in the course of an oxidation-reduc-
tion titration, thus:

0.059

,
log

" See Jamieson, Volumetric Iodate Methods, Chem. Cat. Co., 1926, for a
compilation of the methods using standard iodate solutions. For a study of
the conditions under which the iodine monochloride end-point can be applied
to titrations with other oxidizing agents see Swift, /. Am. Chem. Soc. t 62,
894 (1930).

NEUTRALIZATION METHODS 89

It is seen that the same principles which were discussed in con-
nection with the use of the iodine-iodine monochloride end-point
can be applied to these substances. Potential indicators which
have been extensively used are the following: (1) Diphenylamine,
which is colorless in the reduced form and violet in the oxidized
form, the transition range occurring close to 0.76 v. This indi-
cator has been extensively used for the titration of ferrous salts with
dichromate. 64 It is necessary that phosphoric acid be present
(which, by forming an un-ionized compound with the ferric iron,
gives a more positive and therefore less oxidizing potential at the
equivalence-point) or the transition is somewhat sluggish and pre-
mature. It is not a strictly reversible reaction. (2) Diphenylamine
sulphonic acid 86 turns a reddish- violet color at about 0.83 v.
It therefore gives a sharper end-point in the titration of ferrous salts
in the absence of phosphoric acid and in addition is more soluble in
aqueous solutions than is diphenylamine. (3) Orthophenanthroline
ferrous complex 66 has been extensively used with eerie sulfate solu-
tions. It is red in the reduced form and very pale blue when oxi-
dized. The transition occurs at a potential of about 1.2 v
too high to be used very satisfactorily with dichromate solutions.

NEUTRALIZATION (AND DISPLACEMENT) METHODS OF
VOLUMETRIC ANALYSIS

General Discussion* Neutralization Reactions. For the purposes
of this discussion only neutralization in aqueous solutions will be
considered, and the classical definition will be used, namely, that
the characteristic feature of neutralization methods is the reaction
of hydrogen ions with hydroxyl ions to form un-ionized water,* 7

H+ + OHT = H,0.

" Knop, J. Am. Chem. Soc., 48, 263 (1924).

M Server and Kolthoff, /. Am. Chem. Soc., 63, 2902, 2906 (1931); Willard and
Young, Ind. Eng. Chem., Anal. Ed., 5, 154 (1933).

* Walden, Hammett, and Chapman, /. Am. Chem. Soc., 66, 2649 (1933).

67 For simplicity of presentation the hydrogen ion will be written as above
even though it exists as H|O+ in aqueous solutions. It is also felt that neutral-
ization reactions in nonaqueous solutions are not as yet of sufficient analyt-
ical importance to warrant a more generalized treatment of the subject at
this place. The article "Modern Conception of Acids and Bases'' by N. F.
Hall, /. Chem. Ed., 7, 782 (1930), is recommended to those wishing a prelimi-
nary treatment of thip topic; references to more detailed treatments are given
in the article.

90 VOLUMETRIC METHODS

The equilibrium expression [H + ][OHT]/[H 2 O] = K w , which can
be written [H + ][OET] = K w (since in dilute solutions the concen-
tration of the un-ionized water is practically constant), has for K w ,
called the ionization constant for water, the value at 20C. of 1.0
X 10~~ 14 . If the equation for the reaction of a "strong" acid, that is,
one which is highly ionized (such as hydrochloric, nitric, and the first
hydrogen of sulfuric acid), with a strong base (such as sodium or
potassium hydroxide) is written ionically

H+A~ + B+OHT = HOH + B+A~

it is seen that the equilibrium expression for the reaction reduces to
the form of the reciprocal of that for the ionization constant of water.
Therefore the constant for such a neutralization reaction, (X n ),
has the value 10 14 , which shows that the reaction is complete well
within the usual measurements of quantitative analysis.

When a weak acid reacts with a strong base, the reaction can be
represented as follows:

HA + B+OH~ = H 2 O + B + A~

If ionization of the salt is assumed to be complete, the equilibrium
expression for this reaction has the form

[A"]
[HA][OH-]

Here it is obvious that not only the ionization constant for water
[H+][OH~] = K w but also that of the acid [H + ][A~]/[HA] - K A has
a determining effect on the completeness of the reaction and on the
value of K n . By combining these Jast two expressions, thus elimi-
nating [H" 1 "], there is obtained the neutralization expression above,

~ _ K A _

*irr~~ J\. n .

[HA][OH~] K w

This shows that K n for the reaction of a weak acid with a strong base
is equal to the constant for the acid divided by the constant for water.
It is therefore obvious that such reactions will not be as quantitative
as those between strong acids and bases, and the completeness of
such reactions will vary with the strength of the acid. Similar
considerations will show that K n for the reaction between a weak
base and a strong acid will be given by K B /K W .

For the reaction of a weak acid and a weak base, HA + BOH =
H2O + B*A~, the equilibrium expression is

[HA][BOH]

= K n .

NEUTRALIZATION METHODS 91

Here the constant for water, K WJ for the acid, KA, and for the base,
KB, are all involved, and by properly combining the corresponding
expressions it will be seen that the neutralization expression is

[B*][A~1 K A K B
[HA][BOH] K w An *

From this it is seen that the reaction of even moderately weak acids
and bases will not proceed to completion.

The Hydrolysis of Salts in Aqueous Solutions. In the foregoing
considerations the reactions have been considered as taking place
from left to right as written and have been termed neutralizations.
It is obvious that, especially for those cases where the reaction is
incomplete in this direction, if we start with the products on the
right side of the equation, namely, a salt and water, then the reac-
tion will proceed to the left until an equilibrium is set up which
satisfies the equilibrium constant. Such reactions are termed
hydrolysis reactions, and from the considerations given above it is
clear why the hydrolysis of salts occurs, and how, when the con-
stants for the acids and bases involved are known, the extent of such
reactions can be predicted, and the equilibrium conditions calculated.

Since the value of the ionization constant for water (K w ) changes
much more rapidly with temperature than docs that of most acids
and bases, K w being 0.11 X 10~ 14 at 0C. and 48 X 10~ 14 at 100C.,
it is seen that at higher temperatures neutralization will be less
complete and hydrolysis correspondingly greater.

The Changes in Hydrogen and Hydroxyl Ion Concentrations during
Neutralization Titrations. Analytically, the factor of most importance
during the course of a neutralization titration is the change in the
hydrogen ion concentration of the solution, because it is this change
which enables the end-point to be determined. This hydrogen ion
concentration can be calculated for any point during the titration
involving a strong acid and base from the excess of either the acid or
base present (if it is assumed that they are completely ionized and
that the activity is equal to the concentration). At the equivalence-
point

[H+] = [OH~1 = 1(T 7 molal (at 20C.),

and the solution is said to be neutral. In these strong acid-strong
base titrations the hydrogen ion concentration changes very rapidly
with the addition of acid or base, especially near the equivalence-
point. This is shown in Fig. 16 by the curves labeled "Strong
Acids" and "Strong Bases." In this figure there are shown the

VOLUMETRIC METHODS

NEUTRALIZATION METHODS 93

changes in the [H + ] and [OH""] plotted logarithmically as a function
of the ratio of acid to base during the course of titrations with acids
and bases of various strengths. 68 The similarity of these curves
to the titration curves for precipitation and for oxidation and reduc-
tion reactions is to be noted.

The calculation of the hydrogen ion concentrations at various
points during the titration of a weak acid with a strong base is not so
simple; as an example consider the titration of acetic acid (K A =
1.8 X 10~ 5 at 25C.) with sodium hydroxide. The equation for the
reaction is

HC 2 H 3 2 + OH" = CsHaO? + H 2 0,
and the equilibrium expression is

1-8 X 10~ 5

* *

If it is assumed that 100 ml of 0.2 n. HC2H 3 02 were taken and that
to it has been added 99 ml of 0.2 n. NaOH, and if it is further as-
sumed that the reaction has proceeded to completion, it is seen
that the excess of HCtHs0 8 would be 0.2 X 1/199 molal, and that
the concentration of the acetate formed would be 0.2 X 99/199
molal. However, since the reaction between the acid and hydroxide
is incomplete, there will be a certain amount of hydroxyl ion re-
maining in the solution. If the molal concentration of this hydroxyl
ion is represented by x, the molal concentration of HC2H 8 02 will be
0.2 X 1/199 + x and that of the CjHaOj" will be 0.2 X 99/199 - x.
Assuming, as an approximation, that x is negligible in comparison
with these values and solving, it is found that the [OH~] is equal to
5.5 X 10~ 8 and the [H + ] = 1.8 X 10~ 7 . 69

68 In plotting these curves it has been more practical to use, rather than the
concentrations, the logarithms of these concentrations. The same is also true
in many of the calculations involved in the determination and use of hydrogen
ion concentrations. This has led to the convention of designating the hydro-
gen ion concentration of a solution by a quantity called the pH and defined
by the equation pH = log 1/[H+] = log[H+]. (Rigorously, the hydrogen
ion activity rather than the concentration should be specified.) Thus it is
seen that for solutions in which the hydrogen ion concentrations are 10~ 7 and
5 X 10~ 7 respectively, the pH values are 7 and 6.3. As the use of pH values has
become quite widespread, their significance should be clearly comprehended.

t A much simpler approximate solution is obtained by noting that as long
as an appreciable excess of HCtHj0 2 is present (as long as HCjHjOj in the

94 VOLUMETRIC METHODS

At the equivalence-point it is more convenient to consider the
problem as one of hydrolysis, as the solution is the same as one
containing an equivalent amount of salt dissolved in water. In this
case the formal concentration of the C 2 H 3 0^ is 0.2 X 100/200, and
the molal concentration of the HC 2 H 3 2 is equal to that of the [OH""].
If this is represented by x, the molal concentration of C 2 H 3 07 is
O.lo?, and the approximations used above give the [OH~] as
7.5 X 10~ 6 and the [H + ] as 1.3 X 10~ 9 . This value shows that at the
equivalence-point of this titration the solution has already become
distinctly alkaline.

It is thus seen that the hydrogen ion concentration can bo calcu-
lated for frequent intervals during the course of a titration and that
curves similar to the titration curves for other types of reactions
can be constructed from these data. This has been done and is
shown in Fig. 16, where the [H + ] at any point during the titration
of acids with various ionization constants with strong bases, and of
bases with various ionization constants with strong acids, can be
approximated by means of the curves which are shown. These
curves are calculated for 0.2 n. solutions; in general the error will
be less when using more concentrated and greater with more dilute
standard solutions. The case for the titration of weak acids with
weak bases will not be considered as it is relatively unimportant in
analytical work, due to the incompleteness of these reactions and
to the consequent relatively slow rate of change of the [H + ] during
the titration.

The hydrogen electrode. So far we have considered only the course
of the reaction during the neutralization titration without reference
to possible means of determining the end-point. It was seen in the
discussion of oxidation-reduction titrations that the end-point could
be determined potentiometrically by inserting in the solution an
inert conducting electrode and measuring the change in potential
during the course of the titration. In the absence of other strongly
oxidizing or reducing systems it is possible to do the same thing in a

above equation is large in comparison to x), the [H + ] can be calculated from
the familiar relation

[HA]

IH+) = K A

IA-1

Also, as there is a ratio of the concentrations, the [H+] is approximately
independent of the volume or dilution.

NEUTRALIZATION INDICATORS 95

neutralization titration by saturating the solution with hydrogen
gas at constant pressure and measuring the potential of the reaction

H 2 = 2H+ + 2E~

during the course of the titration. Also, as the molal potential of
this reaction is known, it is possible to determine the hydrogen ion
concentration in the solution if the pressure of the hydrogen gas is
maintained at the same known value this is the principle and proce-
dure used in the determination of hydrogen ion concentrations or
pH by means of the so-called "hydrogen electrode." 60

Neutralization indicators. Potentiometric titrations of neutraliza-
tion reactions are to be avoided if possible, since the electrode is
usually slow to reach an equilibrium and considerable time is re-
quired. The end-point is most frequently obtained by the use of
organic compounds called neutralization indicators, which for the
present can be considered as weak acids, possessing the distinguish-
ing characteristic that the free acid exhibits a different color from
that of the basic ion; thus they show a color change upon being neu-
tralized. Assuming them to be simple acids, the ionization of these
indicators can be represented as

HIn = H+ + In', (1)

and the conventional mass- action expression formulated:

[H + ][IiT] _
[HIn] -* 1 " (2)

Here KI R is the ionization constant of the indicator acid, or, as it is
more commonly termed, the indicator constant. It is seen that the
color which the indicator exhibits will be determined by the fraction
of the indicator which has been converted to the salt or basic (In~)
form (red with phenolphthalein) and the fraction remaining in the
acidic (HIn) form (colorless with phenolphthalein). Thus if we
have a solution in which the indicator is c formal and we let x be

60 The electrical equipment used is essentially the same as that outlined
in the method for making potentiometric titrations. Instead of the straight
piece of platinum wire used there as the inert electrode, it is necessfc*y||j$
use a metallic electrode, usually platinum which has been coated wittr&QI
amorphous deposit of that metal; such a surface more quickly reaches a#
equilibrium with the hydrogen gas. Some device must also be provided for
keeping the electrode and the solution immediately surrounding it saturated
with the hydrogen gas at a constant pressure.

96 VOLUMETRIC METHODS

the fraction converted into the basic form, then ex is the molal con-
centration of In~ and c[l x] that of the HIn, and we obtain

1 - x

In;

which is the conventional form of the indicator equation. From
this it is apparent that the indicator is 50 per cent transformed
when the hydrogen ion concentration hereafter indicated for
brevity as the [H*] is equal to the value of Ki n ; also that if the
[H*] of the solution is known and the fraction of the indicator trans-
formed determined (for example, by colorimetric measurements),
then Ki n can be calculated; or, that knowing lC In and determining x,
the [H*] of a solution can be determined. 61

From the foregoing it is seen that an indicator changes in color
when the [H*] of the solution changes through such a range that a
sufficient fraction of the indicator is converted from one form to the
other to result in a color change visible to the eye. The percentage

61 The mechanism of the color change of indicators is not as simple as has
been assumed above, but usually involves equilibria which can be represented
as follows:

HIn' - HIn' and HIn" H+ + In",

where HIn' (with phenolphthalein) would represent the colorless acid, and
this would be in equilibrium in an acid solution with HIn", the colored form.
This latter is present in only small amount but being more highly ionized is
converted into its salt (In"~, colored) upon addition of base. It is seen that
combining the mass-action expressions

[HIn'J "
gives an expression, namely,

-
l * In

[Hln'l

in which, as long as HIn" and In 7 are small in comparison to HIn' and In",
the total indicator in the acid form can be substituted for the HIn' and that
in the basic form for In"~, and that thus it is justifiable to treat indicators
as though they were simple acids. It is also apparent that even though the
indicator be a base, InOH * In"*" + OH", the ionization expression

IIn+HOH-1 _
[InOHJ B

can be combined with that for water [H+][OH~] K* to give

IInOHHH+1

which again gives the relation between the basic and acidic forms in terms of
the (H+l.

SOLUTIONS OF ACIDS AND BASES 97

conversion necessary to cause a visual color change depends upon
the specific nature of the indicator, and whether the change is from
colorless to a colored form, or from one color to another. In general
the human eye cannot detect the presence of less than 10 per cent
of one form of an indicator in the presence of the other, and, there-
fore, the transition range of most indicators extends over approxi-
mately two pH units. However, by titrating to either the acid or
basic color of an indicator, this pH interval can be restricted to
approximately either half of the transition range; by using a refer-
ence solution containing the same volume of solution and indicator
(and approximately the same amount of salts or of other compounds
as are present in the titrated solution at the equivalence-point), the
pH of the titrated solution at the end-point can be adjusted to within
a few tenths of a pH unit. In Fig. 16 is shown not only the change
in (H*) during the titration of various acids and bases but also the
so-called "transition ranges" of several of the more common indi-
cators. Therefore, by the use of this table it is possible to select for
use in titrations those indicators which will give color changes near
the equivalence-point of the titration. Thus it is seen that any
of the indicators listed can be so used as to give a color change within
0.2 per cent of the equivalence-point in titrating strong acids with
strong bases, but that such is not the case when titrating weak acids
with strong bases or weak bases with strong acids. The transition
ranges of indicators are shifted somewhat by changes in (1) tem-
perature, (2) the ionic concentration of the solution (this is usually
termed the "salt effect"), and (3) the solvent (for example, by the
addition of alcohol to the solution). Usually the magnitudes of
these effects are not such as to cause serious errors in titrations,
but they are of importance in the use of indicators to determine the
hydrogen ion concentration of a solution. 62 The selection of the
proper indicator for specific titrations will be taken up in more de-
tail in the following procedures.

The Preparation of Standard Solutions of Acids and Bases

General discussion. It is desirable that, unless intended for some
specific purpose, the acids and bases to be used as standard solutions
in neutralization methods should possess certain general qualifica-

61 For a discussion of these effects and of the indicator method of deter-
mining hydrogen ion concentrations see Kolthoff, "The Colorimetric and
Potentiometric Determination of pH," John Wiley and Sons, 1931; Clark,
'The Determination of Hydrogen Ions," 3rd ed., Williams and Wilkins, 1928.

98 VOLUMETRIC METHODS

tions, as follows: (1) They should be highly ionized, for, as has been
mentioned, if the titration of a weak base with a weak acid is at-
tempted, the rate of change of the hydrogen ion concentration near
the equivalence-point is likely to be so slow as to prevent a precise
determination of the end-point. (2) They should be so soluble that
solutions as concentrated as half normal and in some cases even 1
normal can be prepared. (3) They should form soluble salts, since
the formation of a precipitate during the titration may obscure the
end-point, or the precipitate may adsorb the indicators. (4) They
should be stable compounds. Oxidizing or reducing agents, besides
being affected by the possible presence of extraneous material (such
as dust, organic matter, or even the oxygen of the air) are likely to
react with the indicators, many of which are unstable with respect
to oxidation or reduction, being decomposed or converted into color-
less compounds. Volatile compounds, such as ammonia, are diffi-
cult to preserve in storage without elaborate precautions. Sub-
stances which rapidly attack glass containers also require special
apparatus in storage,

No one acid or base meets all of the above qualifications. Sulfuric
acid, which is nonvolatile, forms insoluble salts with alkaline-earth
hydroxides; nitric acid, although relatively stable when cold and
dilute, may act as an oxidizing agent in hot solutions or may contain
traces of nitrous acid, which reacts with certain indicators, notably
methyl orange and methyl red. Perchloric acid is a strong acid,
is nonvolatile, is stable toward reduction in dilute solutions, but is
expensive, and the potassium and ammonium salts are only moder-
ately soluble. Although hydrogen chloride is a gas, it is so highly
ionized in aqueous solutions that its partial pressure from even 0.5
normal solutions is so small that they can be boiled for considerable
periods of time without appreciable loss if the solution is not allowed
to concentrate by evaporation. Since most chlorides are soluble
. and since hydrochloric acid is relatively inert toward oxidation or
reduction, it is very extensively used as a standard acid,

Standard bases present more difficulties in storage, because they
tend to attack glass containers and to absorb various gases, espe-
cially carbon dioxide. Barium hydroxide is sometimes used as a
standard base but is sparingly soluble and forms insoluble salts;
ammonium hydroxide solutions lose ammonia; potassium hydroxide
has no distinct advantage over sodium hydroxide and is more ex-
pensive; therefore the latter is the more generally used reagent.

P. XV] SODIUM HYDROXIDE SOLUTIONS 99

P. XV. The Preparation of Carbonate-Free Solutions of Sodium

Hydroxide

Discussion. Sodium hydroxide cannot be used for the direct
preparation of a standard solution because it cannot be obtained
sufficiently pure and because it is too hygroscopic and too reactive
with the carbon dioxide of the air. Carbonate is undesirable in
standard alkali solutions because when titrated with acid using a
strongly acidic indicator (for example, methyl orange) it is converted
into carbonic acid, thus using two equivalents of the acid for each
mole of carbonate present. When it is titrated with acid using a
strongly basic indicator (for example, phenolphthalein), it is con-
verted to hydrocarbonate (if the solution is cold) and thus uses
only one equivalent of the acid; furthermore, this latter end-point
is not obtained precisely except under very strictly controlled condi-
tions. For these reasons it is an advantage to have standard alkali
solutions which are relatively free of carbonate. It is very difficult
to obtain sodium or potassium hydroxide which is free from car-
bonate, but, fortunately, sodium carbonate is relatively insoluble in
50 per cent sodium hydroxide solution, and by preparing such a solu-
tion and allowing it to stand until the sodium carbonate settles out
(or filtering), a carbonate-free solution can then be prepared by dilu-
tion with water that is free of carbon dioxide.

Procedure XV: PREPARATION OF A CARBONATE-FREE SO-
LUTION OF SODIUM HYDROXIDE. Weigh out 16 g of sodium
hydroxide sticks, dissolve them in 15 ml of water, transfer
the solution to a 50-ml test tube (Note 1), stopper the mix-
ture with a rubber stopper, and then allow it to stand in a
vertical position until the precipitate settles (Notes 2, 3).
Quickly pipet out 10 ml of the clear solution without stir-
ring up the residue, and transfer it to a bottle which has been
fitted with a two-hole rubber stopper carrying a siphon
tube with a short rubber tube and pinch clamp at the bottom
in one hole and a short inlet tube to which a soda-lime tube
is attached in the other hole (Note 4). Immediately dilute
the solution to a liter with water which has been just boiled
and cooled (Note 5), and fit the stopper into place.

Notes:

1. This tube should be of resistance glass because of the rapidity with
which ordinary "soft" glass would be attacked by the concentrated alkali.
If resistance glass tubes are not available or if it is desired to keep this
mixture for preparing future solutions, the inside of the test tube should be

100 VOLUMETRIC METHODS [P. XVI

first coated with paraffin; the alkaline solution should be cooled before being
transferred to a paraffined container.

2. The settling of the sodium carbonate may require considerable time,
usually several days. Therefore, if a centrifuge is available, the process can be
expedited by centrifuging the mixture until the precipitate is thrown out.
Otherwise the mixture can be filtered with the aid of suction through an
asbestos or sintered-glass filter and collected directly in the container in
which it is to be stored.

3. Since it requires some time for the sodium carbonate precipitate to
settle and since there is considerable loss of sodium hydroxide in preparing
the 50 per cent solution for only 1 liter of the diluted hydroxide, it is suggested
that a stock supply of the concentrated solution be prepared in a paraffined
container and kept available in the laboratory.

4. If the solution is to be kept for a considerable period, it is desirable
that the inside surface of the bottle be protected by a layer of paraffin. To
do this, clean and dry the bottle, warm it in an oven or in hot water, and
pour into it sufficient melted paraffin to cover the inner surface. Roll the
bottle until the sides are completely covered, tilt it until the bottom is
uniformly covered, and allow it to cool in an upright position.

5. According to Kolthoff, Biochem. A., 176, 101 (1926), water in equilib-
rium with air of the usual carbon dioxide content is only 1.5 X 10~ 5 f. in
carbon dioxide. This amount is negligible in alkali solution of the con-
centrations ordinarily used. However, distilled water is often found to be
supersaturated to an amount 10 or 20 times greater than this equilibrium
value; this excess is removed by boiling for a few minutes.

P. XVI. The Standardization of Sodium Hydroxide Solutions

Discussion. Alkali solutions can be precisely standardized against
constant-boiling hydrochloric acid, 63 benzoic acid (CeHaCOOH), or
potassium hydrophthalate (KHCgHUO^; commonly called potassium
acid phthalate. Oxalic acid is sometimes used t but the hydrated
compound (H 2 C20 4 2H 2 0) is somewhat difficult to prepare with an
exactly known water content, and the anhydrous acid is too hygro-
scopic for practical use. Constant-boiling hydrochloric acid is an
excellent standard if it is available; however, for a single standardiza-
tion its preparation is a lengthy process. Benzoic acid can be ob-
tained from the Bureau of Standards, but the resublimed acid is so
voluminous as to be somewhat inconvenient to weigh unless it is
melted; also, the acid is so slightly soluble in water that it must be
dissolved in alcohol for the titration. Potassium hydrophthalate,
which is also supplied by the Bureau of Standards, has neither of the

11 Hulett and Bonner, /. Am. Chem. Soc., 31, 390 (1909); Morey, ibid., 34,
1032 (1912); Foulk and Hollingsworth, ibid., 45, 1220 (1923); Shaw, J. Ind.
Eng. Chem., 18, 1065. (1926).

P. XVI] SODIUM HYDROXIDE SOLUTIONS 101

difficulties mentioned in connection with benzoic acid and in addition
has a high equivalent weight; it is therefore used here.

The ionization constant for the second hydrogen of phthalic acid
is 3.1 X 10~ 6 , and it is calculated (assuming that the titration is made
with 0.2 n. solutions and that therefore the final concentration of the
salt is 0.1 f.) that the [H*] at the equivalence-point will be 5.6 X
10~ 10 . This value is obtained approximately by referring to Fig. 16
and tracing to the equivalence-point an imaginary curve for an acid
whose constant is 3 X 10~ 6 . It is seen that this value lies within
the transition range for phenolphthalein and that, furthermore,
when the ratio of the equivalents of base to acid is 0.998, the indi-
cator is largely in its acid form (the indicator constant, # In , of
phenolphthalein has the value 2 X 10~ 10 ) and that when this ratio
is 1.002, it is largely in its basic form; thus it is indicated that, using
phenolphthalein as the indicator, phthalic acid can be titrated with
sodium hydroxide (or any strong base) to an accuracy within 0.2
per cent.

It is obvious that methyl orange could not be used as the indicator
for the above titration. However, for the titration of ammonium
hydroxide (K B = 1.8 X 10~ 5 ) with a strong acid it is seen that the
titration curve would indicate that the [H 4 "] at the end-point would
be approximately 7 X 10\/ close to the transition range of methyl
orange, and that when the ratio of acid to base was 0.998, the indi-
cator would be very largely in the basic form, and that when the
ratio was 1.002, it would be sufficiently converted to the acid form for
a color change to be observed; therefore, an accuracy within 0.2 per
cent could be attained. From these examples it is seen that with
the aid of Fig. 16 the accuracy with which any acid can be titrated
with a strong base, or with which any base can be titrated with a
strong acid, can be predicted if the transition range of the indicator is
known.

As either methyl orange or phenolphthalein can be used for the
titration of strong acids with strong bases, and as weak acids can be
titrated by the use of phenolphthalein and weak bases by the use of
methyl orange, these two indicators can be used for nearly all of the
neutralization titrations usually made. However, the indicators
shown in Fig. 16 have certain specific advantages: Bromphenol blue
is useful for the titration of carbonate or hydrocarbonate (HC07)
to carbon dioxide; methyl orange for the titration of phosphoric
acid to dihydrogen phosphate (H 2 P07) ; methyl red for the titration
of dilute solutions of strong acids and bases where methyl orange

102 VOLUMETRIC METHODS [P. XVI

would give an appreciable error; thymol blue for the titration of
carbonate to hydrocarbonate ; and thymolphthalein for titration of
phosphoric acid or dihydrogen phosphate to monohydrogen phos-
phate (HPOr).

Since the transition ranges of indicators are somewhat broad and,
in many cases, not as sharply defined as would be desired for a pre-
cise titration, reference solutions are frequently necessary for ob-
taining the proper end-point. Such a reference solution should
contain the same volume and the same amount of indicator as the
titrated solution; in addition, it is desirable* that it contain approxi-
mately the same amount of the salts or other products as are formed
by the titration reaction. The titration is then carried out until the
color in the titrated solution matches the color of the reference solu-
tion. Such a procedure is especially useful in the titration of poly-
basic acids to intermediate neutralization stages, since the pH change
near the equivalence-point in such titrations is usually not suffi-
ciently rapid to give very sharp end-points. As an example, if a titra-
tion of phosphoric acid to monohydrogen phosphate is to be made,
an amount of monohydrogen phosphate (conveniently added as
Na2HP04-2H 2 0) equivalent to that to be formed in the titration
should be dissolved in the same volume of water as would be present
at the end-point, the same amount of indicator added, and this solu-
tion used as the indicator reference standard.

Procedure XVI : STANDARDIZATION OF A SODIUM HYDROX-
IDE SOLUTION AGAINST POTASSIUM HYDROPHTHALATE. Dry

about 6 g of KHC 8 H 4 4 at 120C. for 1 hour and allow
the crystals to cool in a desiccator (Note 1). Precisely
weigh out about 1.5g of the phthalate and dissolve it in
50 ml of water in a 300-ml flask (Note 2). Add 3 drops
of phenolphthalein indicator to the solution and titrate it
with the 0.2 n. sodium hydroxide solution to be standardized
(Note 3) until the first perceptible pink color is obtained
(Note 4). Stopper the flask (Note 5). Add to a similar
flask a volume of water equal to that of the titrated solu-
tion and the same amount of indicator solution, and care-
fully add the sodium hydroxide solution to it until a color
matching that in the first flask is obtained (Note 6). Sub-
tract the amount required for the blank from the first titra-
tion. Weigh and similarly titrate two more samples of the
phthalate, and from the data thus obtained calculate the
normality of the sodium hydroxide solution.

P. XVII] HYDROCHLORIC ACID SOLUTIONS 103

Notes:

1. KHC 8 H 4 04 is not hygroscopic. If properly prepared and dried, and
thereafter protected from excessive moisture, the material should not con-
tain over 0.05 per cent water; because of this, the preliminary drying may
often be dispensed with. It has also been shown 64 that standard solutions
of the hydrophthalate are stable over long periods of time; it is therefore
possible to prepare such a solution and use it from time to time.

2. Carbon dioxide, if present in considerable amounts in this water, would
cause an error in this standardization, because it would be converted to
hydrocarbonate (HCOa ) at the phenolphthalein end-point, thus using an
equivalent amount of the sodium hydroxide. The extent of this error is
usually small; however, if desired, this possibility can be eliminated by first
boiling the water in the flask, and then inserting a stopper carrying a soda-
lime tube and cooling the water. (See Note 5, P. XV)

3. Alkaline solutions should not be allowed to stand in burets with glass
stopcocks for any longer time than is absolutely necessary, and the buret
should be rinsed as soon as the alkali is removed. Burets with tips connected
by short pieces of rubber tubing (Mohr type) can be used for all work in
which a high degree of precision is not required.

4. Care should be taken to use the same amount of indicator for each
titration and for the blanks, since, especially with a one-color indicator, the
pH range at which the color change is observed will be dependent upon the
concentration of the indicator.

5. A solution which is alkaline to phenolphthalein will slowly absorb
carbon dioxide from the air and become colorless.

6. In order to make a theoretically precise blank, the solution should
contain an amount of sodium or potassium phthalate, I^CgH^, equivalent
to the hydrophthalate taken for the titration. Experiments have shown
that the correction as made will give results within the precision of the other
measurements of the standardization.

P. XVII. The Preparation and Standardization of Hydrochloric

Acid Solutions

Discussion. Standard solutions of hydrochloric acid can be
obtained by preparing constant-boiling acid, 65 which is approxi-
mately 6 normal, and then diluting weighed amounts of this to the
desired volume. More frequently, such solutions are prepared by
diluting the concentrated acid of commerce to approximately the
desired normality and then standardizing this solution against an
appropriate primary standard, such as sodium carbonate or borax. 66
The former substance is more commonly available and easily pre-
pared, and therefore it is used here. An optional procedure for the

M Hoffman, J. Research Natl. Bur. Standards, 15, 583 (1935).

w See references, p. 100.

" Kolthoff, J. Am. Chem. Soc., 48, 1447 (1926).

104 VOLUMETRIC METHODS [P. XVII

preparation of borax, NatB 4 07-10HjO, and its use as a standard is
given in Note 8 below.

Since sodium hydrocarbonate is more readily obtained pure than
is the normal carbonate, it is prepared (or purchased) and converted
into the latter. There have been many investigations and some
contrary evidence as to the possibility of converting sodium hydro-
carbonate to the anhydrous carbonate without obtaining some
sodium oxide (or hydroxide) in the process. 67 However, there ap-
pears to be conclusive experimental proof that by heating pure
sodium hydrocarbonate at 270 to 300C. complete conversion
to the anhydrous carbonate is effected and a product suitable for use
as an acidimetric standard obtained.

Displacement Reactions. The titration of a carbonate with a strong
acid is not, strictly speaking, a neutralization reaction, since the
formation of water is not primarily involved. The reaction

+ 2HC1 = H 2 CO 8 + 2NaCl
is written ionically

C07 + 2H+ = H 2 C0 3 ,
with the equilibrium expression being

[HCO]

where K D is obviously the reciprocal of the total ionizatiou constant
of carbonic acid. 88 Such reactions are termed displacement reae-

7 Lunge, Z. angew. Chem., 10, 552 (1897); 17, 231 (1904); 18, 1520 (1905);
Schmidt, Z. anal. Chem., 70, 321 (1927); Smith and Hardy, J. Chem. Ed., 10,
507 (1933); Waldbauer, McCann, and Tuleen, /. Ind. Eng. Chem., Anal. Ed.,
8, 336 (1934); Smith and Croad, ibid., 9, 141 (1937).

18 It should be pointed out that the above reaction proceeds in two steps as
follows:

(1) CO," + H+ - HCO,~
and

(2) HCO~ + H+ - H,CO,,

The end-point of the first reaction can be obtained by properly using phenol-
phthalein (or other indicators having approximately the same pH range).
However, the pH change near the equivalence-point is so much less pronounced
than it is for the second reaction that it should not be used for highly precise
titrations. Simpson, J. Ind. Eng. Chem., 16, 709 (1924), has recommended a
mixed indicator of 6 parts thymol blue and 1 part cresol red when it is desired
to use this end-point for the analysis of carbonate-hydrocarbonate mixtures.
The indicator changes from a violet-purple color in carbonate solution to a
rose at the end-point and becomes orange-yellow with excess acid.

P. XVII] HYDROCHLORIC ACID SOLUTIONS 105

tions, since it was formerly considered that carbonic acid was dis-
placed from its salt. It is obvious that the weaker the acid the
greater the value of K D , the constant for the displacement reaction,
and the more complete the reaction. In considering the change in
the [H*] of the solution with regard to determining the end-point
of such titrations, it is to be observed (1) that the salt of a strong
base and a weak acid will be alkaline by hydrolysis; (2) that during
the titration the [H + ] added will be used up as long as an appreciable
concentration of the salt remains; (3) that at the equivalence-point
one has essentially a solution of the weak acid; and (4) that as an
excess of strong acid is added, the [H + ] will rise very rapidly. There-
fore, if one selects an indicator which has a transition range close to
or slightly above that given by a dilute solution of the weak acid,
a satisfactory titration can be made. For example, in considering
the titration of sodium carbonate, it is to be observed that a solution
saturated with carbon dioxide ^is 3.4 X 10~ 2 molal in H 2 C0 3 ; and,
as the ionization constant for the first hydrogen of carbonic acid is
3 X 10~ 7 , it can be calculated that the [H 4 ] of the solution at the
equivalence-point is approximately 10~ 4 molal. The transition
range for methyl orange extends from about 4.5 X 10~ 5 to 10~ 8 ;
therefore methyl orange is perceptibly, though not completely,
transformed to the acid form in such a solution, and though the
further addition of acid causes a pronounced additional change
toward the pink color, the end-point is not as sharp as would be
desired. Because of this it is advisable to titrate with the acid to the
first perceptible change, to heat the solution sufficiently to expel the
carbon dioxide, then to cool it (since the methyl orange transition
range is shifted in hot solutions), and to finish the titration with the
cold solution. By this means a very satisfactory end-point is
obtained.

The change in the [H" 1 "] during the titration of a salt can be pre-
dicted from the curves of Fig. 16 if one considers that upon hydrolysis
the salt of a weak acid and a strong base produces hydroxyl ions and
can therefore be considered as a weak base, thus:

A~ + HOH = HA + OH"

The equilibrium expression for this hydrolytic reaction is

[HAKOH-] _

106 VOLUMETRIC METHODS [P. XVII

where KH, called the hydrolysis constant, is obviously the reciprocal
of the neutralization constant, thus:

It is to be noted that the above equilibrium expression for a hydrol-
ysis reaction has the same form as the ionization expression for a
base in which the concentration of the un-ionized base is replaced
by [A~] and that of the basic ion by [HA]. It follows that upon
adding acid to such a solution the [OH""] and the [H + ] will change,
as does a solution of a base with an ionization constant of the same
values as the hydrolysis constant of the salt. Therefore, to use
Fig. 16 for displacement titrations, the hydrolysis constant is calcu-
lated, and the curve for a base with an ionization constant of that
value is traced, or if titrating the salt of a weak base and a strong
acid, the curve for the corresponding acid is followed. The use of
Fig. 16 is thereby extended to displacement titrations.

Consider the application of this principle to the titration of borax,
which is not complicated, as is the titration of carbonates, by the
escape of carbon dioxide from the solution. Boric acid can be
treated as a monobasic acid having a constant equal to 6.4 X 10~ 10 ;
therefore the hydrolysis constant is calculated as follows:

*" = E - 6-OTr. - L6 x 1(r> -

Tracing a curve for a base whose constant is 1.6 X 10~ 5 , we see that
at the equivalence-point the [H + ] is approximately 8 X 10~ 6 and
that methyl red (or methyl orange) would be a very satisfactory indi-
citor for the titration. Such is the case, and if pure borax is avail-
able, it can be used accurately and conveniently as a standard for the
acid. (See Note 8 below.)

The Use of Mixed Indicators. As will be seen by an inspection of
the titration curves for the above displacement reaction, the change
in the pH near the equivalence-point is much less pronounced than
that obtained in titrating strong acids with strong bases; the same
is also true of the titration of weak acids or bases. It is therefore
desirable to employ indicators which have very narrow transition
bands at some definite pH value. In order to obtain this, use has
been made of a mixture of two indicators or of an indicator and an
inert dye stuff. The use of such an indicator is of advantage in this
procedure. (See Note 7 below.) For a discussion and extensive
list of such indicators see Kolthoff and Furman, Volumetric Analysis,
Vol. 2, pp. 62-66.

P. XVII] HYDROCHLORIC ACID SOLUTIONS 107

Procedure XVII : STANDARDIZATION OF AN HYDROCHLORIC

ACID SOLUTION AGAINST ANHYDROUS SODIUM CARBONATE. Cal-

culate the volume of the concentrated hydrochloric acid of
commerce required to prepare 1 liter of 0.2 n. solution and
measure it into a liter volumetric flask or cylinder. Dilute
it to volume, transfer it to a ground-glass-stoppered bottle,
mix it thoroughly, and allow it to cool to room tempera-
ture. Weigh approximately 5 to 6 g of NaHCOs (Note 1)
into a clean platinum or nickel crucible (Note 2), support
the crucible within a larger iron crucible by means of a
triangle (Note 3), hang a thermometer so that the bulb is
close to the crucible containing the NaHCOs, and then heat
the outer crucible so that a temperature of 270 to 300C.
is maintained for 30 minutes (Note 4). Stir the material
occasionally with a clean platinum wire. Remove the
crucible to a desiccator, allow it to cool, and weigh it.
Repeat the heating taking care that none of the material
adheres to the wire (Note 5) and weighing until the weight
remains constant to within 0.5 mg. Transfer the sodium
carbonate to a weighing bottle and keep it in a desiccator
(Note 6).

Weigh out precisely 0.4 to 0.5 g of the prepared Na 2 C03
(Note 6) into a 200-ml flask, dissolve it in 50 ml of water,
add 2 drops of methyl orange indicator solution (Note 7),
and titrate with the HC1 solution, adding the acid slowly
while inclining the flask, until the first perceptible change
in color of the methyl orange (from a clear yellow toward
pink) is obtained. Use the same volume of water and
2 drops of methyl orange in a similar flask as a reference
solution. Heat the titrated solution to boiling for 2 to 3
minutes, frequently swirling it, cool it to room temperature
by running tap water over the outside of the flask while
swirling the contents, and again titrate to the first per-
ceptible change of the indicator. Repeat this process
with two more samples of the NaaCOa and calculate the
normality of the HC1.

Notes:

1. A special brand of NaHCOa for analytical purposes can be purchased
and is usually sufficiently pure to be used directly. If this grade is not
available, the NaHCOs can be prepared by saturating a solution with Na^COa,
filtering, and then saturating this solution with CO*; pour off the solution,

108 VOLUMETRIC METHODS [P. XVII

wash the precipitate with a small amount of water, collect the crystals on a
sintered-glass filter, and'dry at 120C.

2. A porcelain crucible or even a weighing bottle of resistance glass can be
used as a container, but these are less effective, due to their smaller heat
conductivity.

3. A sand bath may be used if the sand is clean and dry and the outside of
the crucible carefully wiped clean before weighing.

4. Care should be taken not to exceed 300C., since the decomposition of
the NaaCOa into Na^O may become appreciable at higher temperatures.

5. After the carbonate has been once weighed, care should be taken to
prevent any loss. Therefore, after stirring the material, the stirring wire
should be carefully brushed with a small camel's-hair brush while held above
the crucible. A small glass stirring rod may be used if care is taken to brush
all particles from it.

6. Na^COs is noticeably hygroscopic, and therefore the bottle should be
opened only when necessary and then closed as soon as possible.

7. The titration can be made more rapidly by adding 1 drop of phenol-
phthalein and rapidly titrating until the red color disappears; this is the
end-point of the reaction

CO,- + H+ - HCO,-.

The methyl orange can then be added and slightly less than the same amount
of acid added without overrunning the end-point.

A mixed indicator, made by dissolving 0.1 g of methyl orange and 0.25 g
of indigo carmine in 100 ml of water, gives a much sharper end-point, es-
pecially if working by artificial light. The indicator is greenish in alkaline
solutions, and turns very sharply to gray at the end-point (pH = 4) and to
violet with an excess of acid. If it is available, the use of this indicator is
advised; 2 to 3 drops per 100 ml of solution are most satisfactory.

8. Borax, Na^Oy-lOB^O, has the advantages over sodium carbonate
of a higher equivalent weight and of not being hygroscopic, but has not been
extensively used because of uncertainty in drying the hydrate without
decomposition. It has been shown that pure borax can be obtained by
Simple recrystallization, that the product can be dried by washing with
alcohol and ether, and that it can be stored in ground-glass-stoppered
containers for a year without a change in composition of greater than* 0.1
per cent. 69 The material can be prepared and the standardization carried
out as follows:

Procedure: Preparation of the Borax. Dissolve 15 g of borax in
35 ml of water, filter out any residue, and cool to room temperature
or below. Collect the crystals on a sintered-glass filter and wash
with two 10-ml portions of cold water. Mix all the wash solutions
thoroughly with the crystals with the aid of a stirring rod and then
use suction to draw off the liquids. Wash with two 10-ml portions
of 95 per cent alcohol and three 10-ml portions of ethyl ether.
Spread the crystals on a watch glass and (1) allow them to stand
in air (protected from dust and so forth) until they are free of ether

Hurley, /. Ind. Eng. Chem., Anal. Ed., 8, 220 (1936); 9, 237 (1937).

P. XVII] NEUTRALIZATION APPLICATIONS 109

and have reached a constant weight (usually from 5 to 8 hours), 'or
(2) place them in a vacuum desiccator and evacuate with the full
force of an efficient water aspirator for 25 minutes, weigh, then
again evacuate for 5-minute intervals until constant weight is ob-
tained. Store the crystals in a ground-glass-stoppered container.

Standardization of the Acid Against the Borax. Weigh out
precisely 1.3 to 1.5 g of the prepared borax, dissolve it in 50 ml of
water, add 3 drops of methyl red indicator solution, and titrate with
HC1 until the color matches an equal volume of water to which has
been added 3 drops of methyl red and 1 g of H 3 B03 and 0.5 g of NaCl.
Repeat the process with two more samples of the borax and calculate
the normality of the HC1.

A considerable saving of time can be effected by the use of the vacuum
desiccator. Experiments have shown that standardization values obtained
from material prepared by its use agree within 0.1 per cent with those obtained
by other methods. It is quite desirable that the comparison flask be used,
or a slightly premature end-point may be taken.

The Applications of Neutralization and Displacement Methods

These methods, often collectively designated as acidimetry and
alkalimetry, can be used for the estimation of acids and bases, the
estimation of salts of weak acids or of weak bases, and the indirect
determination of other substances which can be converted into, or
stoichiometrically precipitated by, compounds of this type. Thus
in this system of analysis, ammonium is estimated by being dis-
tilled as ammonia from an alkaline solution, collected in a known
amount of acid, and the excess acid titrated (P. 96). Nitrate and
nitrite are reduced to ammonia and similarly determined (P. 181).
This same principle is employed in the well-known Kjeldahl method
for the determination of the nitrogen in organic substances; in this
method the organic nitrogen is converted into ammonium salt by
being heated with concentrated sulfuric acid containing various added
substances to aid in the decomposition of the organic material.
The resulting solution is then made alkaline with sodium hydroxide,
and the distillation and titration of the resulting ammonia is carried
out as described in P. 96. Reference books on quantitative analysis
should be consulted for the various modifications of the original
Kjeldahl method, as well as for a more complete treatment of the
determinations which can be made with the use of standard solutions
of acids and bases.

Gravimetric Methods of Analysis

The Gravimetric Standardization of a Hydrochloric Acid Solution

Discussion. Although the chloride present in a hydrochloric acid
solution can be very accurately determined gravimetrically by pre-
cipitating it as silver chloride and weighing the precipitate, in general
it is preferable to standardize acid solutions volumetric ally, rather
than gravimetrically, if they are intended for use in acidimetric
analyses. This is because the volumetric standardization (1) can be
made quite accurately and by methods similar to those for which the
solutions are being prepared, (2) can be carried out more quickly
and easily, and (3) is based upon the acid concentration, whereas the
gravimetric standardization is based upon the anion (in this case
chloride) concentration.

The procedure for the gravimetric standardization is included here
because (1) it can be carried out with exceptional accuracy; (2) it
illustrates exceptionally well many of the principles and much of the
technique involved in gravimetric analysis; and (3) the same
principles and technique form the basis for a majority of the methods
for the qualitative and quantitative separations of various elements
or groups of elements. Because of this, the discussion of these
principles and the description of the technique can be collected here
and reference made to them when particular applications occur in the
later procedures.

GENERAL PRINCIPLES OF GRAVIMETRIC METHODS

Discussion. It was pointed out in the discussion of volumetric
methods that a given reaction had to meet certain requirements
before it could be made the basis for a precise volumetric analysis.
Similarly, certain factors have to be considered in the development
of a gravimetric method. These may be classified into groups as
they affect the following:

I. The solubility of the precipitate.

II. The physical characteristics and purity of the filtered pre-
cipitate.

III. The stability and composition of the weighted precipitate.

I. Factors Affecting the Solubility of the Precipitate

The factors affecting the amount of a given constituent which will
remain in the filtrate when the precipitate is filtered may be classified
as follows:

110

SOLUBILITY EFFECTS 111

1. The common ion and activity effects (the solubility-product
principle).

2. The formation of complex ions.

3 r Hydrogen (and hydroxyl) ion effects.

4. Solvent effects.

5. Temperature effects.

6. Effect of time on the completeness of precipitation (super-
saturation effects). These factors are discussed briefly below.

1. The common ion and activity effects (the solubility-product
principle). It is a fundamental necessity for a gravimetric method
that the precipitate be so insoluble (under the conditions of the
precipitation) that the amount left in the saturated filtrate and in the
wash solution be negligible in comparison with the other errors of the
method (or, as is possible in certain cases, that a reliable correction
can be made for the amount so lost). In some cases calculations
can be made from solubility data as to whether the amount of the
compound lost in the filtrate and wash solution will be within the
accuracy that is desired, or can be reduced to the desired limits by
the addition of a common ion; for making such calculations the
solubility-product principle is employed. This principle has been
discussed in P. VI, where it was applied in predicting the course of
certain precipitation reactions, and reference should be made to that
discussion. There it was pointed out that the solubility-product
principle (and the mass-action law in general) is obeyed experi-
mentally only when the activities of the substances are used instead
of their concentrations, or if the latter are used, only under the
following limitations: (1) With quite insoluble precipitates, (2)
when the total ion concentration of the solution is low, and (3) with

TABLE III
THE SOLUBILITY OP SILVER CHLORIDE IN POTASSIUM CHLORIDE SOLUTIONS

KC1

moles/liter

Ag
moles/liter

KB.P.

0.00670
0.00833
0.01114
0.01669
0.03349

1.7 X lO- 8
1.39 X 10~ 8
1.07 X 10~ 8
0.738 X 10-'
0.388 X 10-

1.14 X 10- 10
1.16 X 10- 10
1.19 X 10~ 10 .
1.23 X 10~ 10
1.30 X lO' 10

salts of simple valence type. In Table III the experimental data of
Jahn 1 have been used to show (1) the solubility of silver chloride in

Jahn, Z. phyaik. Chem., 33, 454 (1900).

112

GRAVIMETRIC METHODS

-5.0

? -6.0

ttt
C

?
5

-9.0

-10.0

-7.0

-8.0

0.01 0.02
Cone, of Cl in g

0.03
Equiv.

0.04

Fig. 17. The Solubility of Silver
Chloride in Dilute Chloride Solutions.

dilute solutions of potassium
chloride and (2) the constancy
of the calculated solubility-prod-
uct value when making the
assumption that the potassium
chloride is completely ionized
and that the. activities of the ions
are equal to their formal concen-
trations. It is to be noted that
the values for the solubility prod-
uct agree reasonably closely in
solutions which are below 0.01
formal in KC1. This agreement
is to be compared with the values
(shown in Fig. 9) for the solubil-
ity of silver sulfate, a more solu-
ble salt and one containing a
binegative ion, where especially
in the presence of an excess of
sulfate ion the solubility varies considerably more from the pre-
dicted value ; in fact, the common ion causes very little decrease in the
solubility of the salt. As previ-
ously stated, these deviations
are due to the attractive forces
existing between the ions in the
solution, and these are responsi-
ble for the marked increase in the
solubility of silver sulfate in solu-
tions of salts not having a com-
mon ion.

2. The formation of complex
ions. A still different effect from
those mentioned above is shown
in Figs. 17 and 18. In Fig. 17
is shown the decrease in the solu-
bility of silver chloride caused by
the common ion effect in low con-
centrations of chloride, and this
should be contrasted with Fig. 18,
where the data of Forbes 2 for the

1.0 2.0 3.0 4.0
Cone, of Cl in g Equiv.

of Silver
Chloride

Fig. 18. The Solubility
Chloride in Concentrated
Solutions.

> Forbes, J. Am. Chem. Soc., 33, 1939 (1911).

SOLUBILITY EFFECTS 113

solubility of silver chloride in concentrated solutions of potassium
chloride are plotted. It can be shown that the remarkably rapid
increase in the solubility of the silver chloride in the more concen-
trated solutions can be explained on the basis of the formation
of complex ions (of the type AgQI, AgCl^, and AgClD, in which
the positive silver ion is surrounded by the negative chloride ions. 3
It is thus seen that if these effects are ignored, calculations on the
basis of the solubility-product principle would lead to erroneous and
misleading predictions. Furthermore, the haphazard addition of a
large excess of a precipitation reagent may defeat the purpose
sought, since it is entirely possible that a large excess of even a
common ion may cause an increase in the solubility of the precip-
itate. It is not to be inferred that the solubility-product principle
does not apply to cases where complex formation exists, but merely
that the additional equilibria involved in their formation must also
be considered. It should also be noted that, as in the large majority
of cases complex ions are formed as the result of adding an excess of
a negative ion to a positive one, less pronounced effects will follow the
addition of an excess of a positive ion as precipitant to a negative
ion than the reverse. Thus, the solubility of silver chloride in 2
formal silver nitrate is only 5 X 10~ 5 formal, less than one tenth
that in 2 formal chloride solution, although greater than it is in water.
Other obvious applications of the tendency of certain elements to
form complex ions are demonstrated in the solubility of silver chloride
and of copper oxide in the presence of ammonia (because of the for-
mation of the di-ammino silver ion, Ag(NH 3 )J, and the tetrammino
copper ion, Cu(NH 3 )t*). 4 Where the values of the dissociation
constants of these compounds have been determined and are avail-
able, calculations can be made as to the effect of such compound
formation on the solubility of a precipitate. Thus, as the value
for the equilibrium-dissociation constant for the expression

-68X10-"
~ b ' b x 1U '

* On the basis of his calculations, Forbes did not find evidence of the ion
in these concentrated solutions, but concluded that AgClF and AgClT
were the species present; later work with less concentrated solutions, Forbes
and Cole, /. Am. Chem. Soc., 43, 2497 (1929), and evidence from other sources
have indicated that the simpler ion is also formed.

4 These and other complex ions are discussed further in P. 11, in connection
with their significance in the precipitation of the.sulfides of the various ele-
ments from acid solutions, and in P. 64, where the complex compounds of
cobalt are treated.

114 GRAVIMETRIC METHODS

it can be predicted that silver chloride would be very soluble in a
solution 0.5 f. in ammonia, but that silver iodide would be relatively
insoluble arid that a separation of iodide from chloride could be thus
effected; this separation is an experimental fact.

3. Hydrogen (and hydroxyl) ion effects. The effects of hydrogen
and hydroxyl ions on the solubility of a substance are treated in the
general discussion of the preparation of the solution for the analysis
and in P. 5. Examples of the various solubilities of the salts of a
weak acid in various hydrogen ion concentrations are also given in
P. 11, Table XI, where the solubilities of the sulfides in different
solutions of acids and bases are shown. The adjustment of the
hydrogen ion concentration is also of importance in gravimetric
precipitations because it often effects the physical character of the
precipitate (this effect is discussed later) and because it can be used
to prevent the precipitation of undesired compounds. Thus if an
attempt were made to precipitate silver chloride from a neutral
solution, there would also precipitate any carbonate, phosphate,
arsenate, or other anions of weak acids which form insoluble silver
salts; the same would be true of the attempt to precipitate barium
sulfate from a neutral solution.

When precipitating the hydroxides by addition of hydroxyl ion,
each specific case has to be considered separately; thus a high con-
centration of hydroxyl ion is required for the complete precipitation
of magnesium hydroxide, whereas with aluminum hydroxide, because
of its amphoteric nature, an excess of hydroxyl ion has to be carefully
avoided.

4. The effect of the solvent. In many cases the solubility of a
salt can be reduced by altering the properties of the solvent or
changing to a different medium. In general, inorganic salts, espe-
cially if they are highly ionized, are less soluble in organic solvents,
such as alcohol and ether, than they are in water. Because of this
effect, strontium chromate, which is appreciably soluble in water
solutions, can be quantitatively precipitated from a solution con-
taining a considerable proportion of alcohol. Likewise, the quanti-
tative precipitation of potassium perchlorate has to be made from an
alcoholic or other nonaqueous medium, and nickel chloride, which is
very soluble in water, can be precipitated by passing hydrogen
chloride gas into an ether solution.

5. Temperature effects. Where the precipitate is sufficiently
insoluble and stable, and where other undesirable effects (such as
the hydrolysis of certain salts or the oxidation of some of the con-

SOLUBILITY EFFECTS 115

stituents present) are not encountered, there are definite advantages
in carrying out the precipitation, filtering, and washing of precipitates
at elevated temperatures, since, in general, the precipitates are more
readily coagulated and brought into a filterable form. As an ex-
ample, barium sulfate which has separated in such a finely dispersed
condition as to pass through even a dense filter can be converted
into a much coarser and more readily filtered form by keeping the
solution near the boiling temperature for a short period. Solutions
are more rapidly filtered when hot, mainly because their viscosity is
much less; the specific viscosity of water is 0.637 at 15C. and only
0.158 at 100C. Experimentally it has been found that hot solutions
will pass through a paper filter from five to ten times as fast as those
at room temperature. Another advantage arises from the fact that
the substances contaminating the precipitate are usually more soluble
in hot solutions and are thus more readily extracted.

6. The effect of time on the completeness of precipitation. It is
not generally realized that in many precipitation reactions there is a
considerable lapse of time before equilibrium is reached. There may
be two general causes for this. The first is the familiar phenomenon
of supersaturation, which is most obvious with moderately soluble
compounds and occurs even when the constituent ions of a salt are
present in high concentrations. In the second case the substance
may be extremely insoluble, and yet the rate of precipitation may be
very slow because the constituent ions are present in such very minute
concentrations. The slow precipitation of arsenic sulfide, As 2 S 6 , from
acid solutions is a striking example of the latter case. This is
undoubtedly due to the very low concentration of simple penta-
positive arsenic ions in the solution and to a slow rate of formation of
these ions from the other molecular species in which the arsenic
exists. 6 Another example of this type of rate effect is involved in the
precipitation of cobalt as potassium cobaltinitrite, K 3 Co(N02)e,
where the precipitation has to be preceded by the oxidation of the
cobaltous ion to the cobaltic state and by the formation of the
complex cobaltinitrite ion.

Although no general rule can be given, it is quite important that a
sufficient time elapse between adding the precipitant and beginning
the filtration, not only for the precipitate to develop into a filterable
form, but also for it to reach an equilibrium. Precipitation can be
induced in supersaturated solutions by vigorous agitation, such as

6 This precipitation is treated from the experimental point of view in the
discussion of P. 11.

116 GRAVIMETRIC METHODS

stirring or shaking; the expedient (often used in preparative work) of
adding a crystal of the precipitate is usually not practical in analytical
procedures. In cases of the second effect, precipitation can be
hastened by providing a high concentration of the reactants, by
increasing the temperature, and, in certain cases, by mechanical
agitation.

II. (1) Factors Affecting the Physical Characteristics of the Pre-
cipitate

It would seem that the physical characteristics of a precipitate
that is, whether it is crystalline or amorphous, granular or gelatinous,
whether it separates in a highly hydrous condition, and even its color
would be determined by its chemical composition. Thus it would
seem that sulfides are of an amorphous nature, hydroxides gelatinous,
and sulfates crystalline because of their chemical composition.
However, owing in a large measure to the work of Von Weimarn, 6
it has been shown that these characteristics are influenced more by
the conditions under which the precipitate forms than by its chemical
composition.

Considering the mechanism of precipitation, it seems reasonable to
expect that if a precipitate forms in a solution in which its normal
solubility is only slightly exceeded, initially only a relatively few
crystal nuclei will be formed, and that after these are present, the
subsequent precipitation will consist mainly of an enlargement, or
"growth," of these crystals. This subsequent growth of crystals is
in accord with the experimental fact that the solubility of extremely
small particles is appreciably greater than that of larger ones. This
behavior is predicted from theoretical considerations and has been
studied experimentally by various workers, 7 who have found solu-
bility increases ranging from 15 to 80 per cent when studying small
particles (ranging from 0.0001 to 0.0002 mm in diameter) of barium
sulfate and of calcium sulfate monohydrate, CaS0 4 -H 2 0. Therefore
a solution which is saturated with respect to the very small particles is
obviously supersaturated with respect to larger ones, and as a result
precipitation takes place on the larger ones, and the smaller ones
tend to pass into solution.

If the initial precipitation takes place in a solution which is greatly
supersaturated, then there will be formed a large number of crystal

' Von Weimarn, Zur Lehre von den Zustanden der Materie (1914).

7 Ostwald, Z. physik. Chem., 34, 495 (1900); Hulett, ibid., 37,385(1901);
42, 581 (1903); 47, 357 (1904). Hulett and Allen, /. Am. Chem. Soc., 24, 667
(1902); Dundon and Mack, ibid., 45, 2479 (1923); Dundon, ibid., 46, 2658 (1923).

CHARACTERISTICS OF PRECIPITATES 117

nuclei, and the precipitate will appear in a highly dispersed condition.
With very insoluble precipitates ordinary methods of mixing the
solutions will always result in a relatively high degree of super-
saturation, and therefore the precipitate will appear as a large number
of very small particles; in fact, they may be so small that they will
remain colloidally dispersed. Furthermore, if a precipitate is
extremely insoluble, the concentration of the saturated solution will
'be so small that the growth of large crystals at the expense of the
more soluble smaller ones will be slow. According to Von Weimarn,
this degree of supersaturation, which can be expressed as the ratio
of the initial supersaturation of the substance (before precipitation
begins) to its equilibrium solubility (or [AS s]/[s], where S represents
the initial concentration of the substance before precipitation begins
and 5 the equilibrium solubility, each expressed in equivalents per
liter), is a major factor in determining the physical characteristics of
precipitates. From this principle the generalization is made that the
physical characteristics of two precipitates will be largely the same,
irrespective of their chemical nature, if they are produced under
corresponding conditions, and the most important factor determining
these conditions is the value of [S s]/[s]. It follows that if the
same substance is precipitated under conditions where [S s]/[s]
has widely varying values, its physical characteristics will be quite
different. As an example, barium sulfate, which normally appears
as a finely divided but distinctly crystalline precipitate, can be made
to separate in a colloidal or gelatinous form by causing it to precipi-
tate from a highly supersaturated solution, or one where the value of
[S s]/[s] is large. This can be accomplished in two ways: (1)
by decreasing its solubility (for example, by the addition of alcohol
to the aqueous solution), thus decreasing s, or (2) by mixing highly
concentrated solutions, thus increasing S. As examples, it is found
that the precipitate resulting from mixing equal volumes of 0.025 n.
cobalt sulfate and 0.025 n. barium thiocyanate in 50 per cent alcohol
has at first a distinctly colloidal appearance, and on coagulating
might easily be mistaken for aluminum hydroxide, while the precipi-
tate resulting from rapidly mixing equal volumes of 7 n. manganous
sulfate and 7 n. barium thiocyanate is a viscous gel. Since colloidal
solutions consist of finely dispersed particles (usually with diameters
ranging from 10~ 4 to 10~ 6 mm), it would be expected that a high
value of the ratio of [S s]/[s] would be favorable to their formation;
this is true, as is shown by the fact that barium sulfate (a typically
crystalline precipitate) can be produced as either a suspensoid or

118 GRAVIMETRIC METHODS

even as a gel. On the other hand, by precipitating barium sulfate
from a 4 n. hydrochloric acid solution, where s is considerably greater
than it is in a neutral solution, the precipitation takes place quite
slowly, and relatively large crystals can be obtained. 8

Apparent exceptions to this rule can be noted in the fact that
substances of approximately the same solubility precipitate in
distinctly different physical form; barium sulfate and silver chloride
furnish an example of this anomaly. These two substances have
about the same solubility, yet normally appear as precipitates of a
very different type silver chloride as a curdy flocculated colloid
and barium sulfate in a definitely crystalline form. The difference
is explained by the fact that the solubility of barium sulfate increases
much more rapidly as the particle size is decreased than does silver
chloride, and therefore the initial supersaturation is much less with
barium sulfate than with silver chloride. 9

It is thus seen that larger and more readily filtrable particles can
usually be obtained (1) if the initial supersaturation of the solution
during the precipitation process is kept as small as possible and (2)
if time is given for the extremely small particles first formed to
increase in size. In addition to this "growth/' which can be
attributed to the increased solubility of the small particles, if a
crystalline precipitate is caused to form very rapidly from relatively
concentrated solutions, it is probable that many imperfections and
strains will result in the process, and that if such a precipitate is
allowed to age, a recrystallization and adjustment will take place.
This process will usually result in a decrease in the total surface and
gives a more filtrable precipitate. This process frequently is re-
sponsible for the changes taking place upon "digesting" (that is,
maintaining the solution for a period of time at or near its boiling
point) a very fine or colloidal precipitate in order to convert it into a
form that can be filtered. 10

Summing up the above with reference to most analytical precipita-
tions where, in order to facilitate filtering and washing and to avoid
colloidal solutions, it is desirable to obtain a coarsely granular or

1 A more detailed discussion of these effects is given by Washburn, Princi-
ples of Physical Chemistry, McGraw-Hill (1915), and by Smith, Analytical
Processes , Longmans, Green (1929).

Kolthoff, J. Phys. Chem., 36, 860 (1932). An extensive discussion of pre-
cipitation and coprecipitation phenomena is given in this article.

10 For a discussion of this subject of the "ageing" of precipitates, see
Kolthoff and San dell, Quantitative Analysis, Macmillan, 1936, pp. 98-102.

COPRECIPITATION 119

crystalline precipitate, it is obviously advantageous (1) to mix the
reagents very slowly and with effective stirring, (2) to use dilute
solutions, (3) in many cases to increase the solubility of the precipi-
tate, usually by the addition of an acid, or by working in hot solu-
tions, and (4) to allow the precipitate to stand until the particle size
is such that it will be retained by the filter.

II. (2) Factors Affecting the Composition and Purity of the Pre-
cipitate

Coprecipitation. It is a universally observed phenomenon that
when a precipitate separates from a solution, it will carry out with
it in varying amounts the soluble constituents of the solution; here-
after the term coprecipitation will be used to designate this phe-
nomenon without implying any specific cause or mechanism. Co-
precipitation effects are one of the most important factors involved
in precipitation methods and have to be considered regardless of
whether a separation or a gravimetric determination is desired. In
general, coprecipitation is not the result of any one effect but may
result from the operation of any one or more of the four following
causes:

1. Adsorption.

2. Compound formation.

3. Solid-solution formation.

4. Mechanical inclusion.

Adsorption as a Cause of Coprecipitation. Adsorption may be
broadly defined as the process which causes an increase in the con-
centration of a gas, liquid, or dissolved substance at an interface
that is, at the surface between two phases. In analytical chemistry
there is special interest in the process whereby constituents of a
solution are concentrated on the surfaces of precipitates. Charcoal
is probably the best-known example of an effective adsorbing agent,
and consequently is extensively used scientifically and industrially
for collecting gases and for removing many substances from dilute
solutions; an example of the latter is the commercial use of special
charcoals for the recovery of iodine from oil-well brines.

The adsorption of acetic acid from aqueous solutions by charcoal
was investigated quite early in the study of adsorption phenomena,
and data showing the relation between the amount of adsorption
taking place on a given amount of charcoal and the concentration of
the acetic acid are shown in Table IV. If the adsorption process

120

GRAVIMETRIC METHODS

TABLE IV
THE ADSORPTION, OP ACETIC ACID FROM AQUEOUS SOLUTIONS BY CHARCOAL

Equilibrium concentration (c)
of HC 2 HiOj moles/liter

Millimoles HCjHsC^ per gram
of charcoal (x/m)

0.0181
0.0309
0.0616
0.1259
0.2677
0.4711
0.8817
2.785

0.467
0..624
0.801
1.11
1.55
2.04
2.48
3.76

Freundlich, Kapillarchemie, 1909.

were of a chemical nature due conceivably to the formation of
insoluble compounds it would be expected that definite saturation
values would be found limiting the amount of material taken up by a
given amount of adsorbing agent. If the process were of a physical
nature, it would be expected to be more analogous to the distribution
of a substance between two phases, and such a process should obey
the distribution law. This law is formulated as follows :

C 2

7T - K >
Ci

where C\ and 2 represent the concentrations of a common substance
in two phases which are in equilibrium and where the substance
exists in the same molecular form in each phase. Should the sub-
stance exist in different molecular aggregations in the two phases
for example, be polymerized in one phase the equation would have
the more general form

(CO" _ K

(CO A)

where n is the number of molecules associated in the polymer, in
agreement with the general mass-action law.

It has been experimentally found that adsorption is an equilibrium
process and that it follows an empirical equation, called the adsorp-
tion isotherm, which has the following form :

where x represents the weight of material adsorbed by the weight m of
adsorbing material, c represents the equilibrium concentration of the

COPRECIPITATION

121

adsorbed material, and K and n
are constants. The constant K
is highly dependent upon the
specific conditions to which the
equation is being applied, and n
is always greater than 1, usually
less than 5, and frequently ap-
proximately 2. It is seen that
this equation bears a formal
resemblance to the distribution
law. In Fig. 19 is shown the
curve obtained by plotting the
data of Table IV. It is also
evident that the equation may
be written

. x log c
log _ = _*_

m n

, r .
log K,

'0 1.0 2.0 30

Equil. Cone, of HC 2 H 3 2 m ~

Fig. 19. The Adsorption of Acetic
Acid from Aqueous Solutions by
Charcoal.

and that a straight line should result from plotting the data loga-
rithmically; this is also shown in Fig. 20. This graphical method

is frequently used to determine
whether a given process is due to
adsorption, or to the presence of
other effects which would cause a
departure from the straight line.

Not only are molecular com-
pounds adsorbed from solutions by
solids, but it is also found that ions
are similarly concentrated at the
surface of precipitates; the charge
associated with colloidal particles
is in general due to this effect, and
the presence of this charge con-
tributes to the stability of the sus-
pension. 11

It is found that, as a general
rule, precipitates tend to adsorb the
ions of which they are composed.

-0.5

-1.0

-4 -2

Log of Equil. Cone.

Fig. 20. The Adsorption of Ace-
tic Acid from Aqueous Solutions
by Charcoal (Logarithmic Plot).

11 It should be pointed out that the solution does not acquire an electro-
static charge because of the adsorption of ions of a particular sign. This is
prevented by the formation of a second loosely attached and more mobile
layer of ions of opposite sign (thus producing the so-called electrical double
layer). Even upon coagulation or filtration the double layer is carried with
the precipitate, and the solution remains electrically neutral.

122

GRAVIMETRIC METHODS

A colloidal suspension of silver chloride, bromide, or iodide which
has been formed in the presence of an excess of the halide ion will be
found to be negatively charged because of the adsorbed negative ions;
in the presence of an excess of silver ions such a precipitate will be
positively charged. The adsorption of silver nitrate by silver iodide
has been extensively studied, and the results of one such study are
shown in Fig. 21.

,-0.006

i^o

! 0.004

2*0.002

0.004 0.008 0.012 0.016

Equilibrium Concentration in Millimoles
per Milliliter of Silver Nitrate

-2.2

-3.5 -3.0 -2.5 -2.0 -1.5

Log of Equilibrium Concentration

Fig. 21. The Adsorption of Silver Nitrate by Silver Iodide.

The silver halide first formed by the addition of silver nitrate to a
solution containing halide ions is negatively charged and usually
remains in suspension until quite near the equivalence-point; then
as the concentration of the halide ion in the solution is reduced to a
very small value, the adsorbed halide ions are withdrawn from the
precipitate and coagulation takes place. 12 By adding a large excess
of silver nitrate, the precipitate can be peptized, that is, again
colloidally dispersed, and the particles will now have a positive charge.

12 In many precipitation reactions this effect can be used to indicate when
an approximately equivalent amount of the precipitant has been added and an
unnecessary excess thus avoided. In certain cases the effect is so pronounced
that the volumetric end-point can be taken with considerable precision by
noting when this sudden coagulation or clearing of the solution takes place.
Such a process is termed titrating to the "clear point." This effect has been
studied for the titration of silver iodide by Lattermoser, Seifert, and Forst-
mann, Kolloid Z., 36, 230 (1925); also a volumetric determination of lead by
titration with a molybdate solution to the clear point has been proposed by
Sacher, Kolloid Z., 276 (1926).

COPRECIPITATION 123

Such a suspension may be coagulated by the addition of negative
ions to the solution; however, the coagulated precipitate will be found
to have carried down with it some of the negative ions, and these
may be an undesirable contamination for quantitative purposes.
Since it is probable that all precipitates pass through the colloidal
state to some extent in the mechanism of their formation, it is not
surprising to find that marked adsorption may occur even with
typically crystalline precipitates. Thus it has been found that the
coprecipitation of potassium nitrate and of sodium chloride by a
barium sulfate precipitate follows the adsorption equation.

The magnitude of the adsorption taking place in precipitation
processes will vary from an amount not exceeding the usual quanti-
tative errors to cases where even qualitative separations cannot be
carried out. Adsorption phenomena are probably at least partly
responsible for the coprecipitation effects shown in Tables XIX and
XXI, where the limitations of the "ammonia precipitation" and the
sodium hydroxide-peroxide process for separating the elements of
the Ammonium Sulfide Group are shown . In many cases it is difficult
to distinguish between adsorption and compound formation. Thus
in the hydroxide separations just mentioned it is probable that the
precipitates adsorb hydroxyl ion on their surfaces and that this
results in the formation of the hydroxides of the soluble elements.
A somewhat similar effect has been noted 18 in a study of the co-
precipitation of zinc sulfide by copper sulfide, and has been attributed
to the induced precipitation of zinc sulfide caused by the adsorption
of hydrogen sulfide on the copper sulfide precipitate.

Space does not permit of an adequate treatment of even the most
important of those adsorption and colloidal phenomena which are
encountered in analytical precipitations and separations. For a
more detailed treatment reference should be made to Analytical
Processes, by T. B. Smith; Elementary Quantitative Analysis, by
Willard and Furman; Quantitative Inorganic Analysis, by Kolthoff
and Sandell; or to reference books and textbooks of colloidal and
surface chemistry.

Compound Formation as a Cause of Coprecipitation. Compound
formation was once thought to be a much more important cause of
coprecipitation than is believed to be the case at present. As
examples, the coprecipitation of zinc by chromium hydroxide in the
ammonia precipitation was attributed to the formation of a zinc

"Kolthoff and Pearson, J. Phys. Chem., SB, 649 (1932); also see the dis-
cussion of P. 61.

124 GRAVIMETRIC METHODS

chromite. This is unlikely because of the fact that in the hydroxide
concentration provided in $n ammonia precipitation there would
hardly be a significant concentration of chromite formed; in fact,
present evidence indicates that the solution of chromic hydroxide in
more concentrated hydroxide solutions is due not to chromite forma-
tion but to the formation of a stable colloidal suspension in which
there is a strong adsorption of hydroxyl ion. Such a suspension
would tend to be coagulated by,. and to carry down with it, a positive
ion, such as zinc; in general, the greater the charge the more effective
the coagulating action. Many cases formerly attributed to "double-,
salt" formation are now believed to be due to adsorption; the adsorp-
tion of cadmium salts by cadmium sulfide precipitate is such a case. 14

Compound formation is a more probable cause of coprecipitation in
cases where partially hydrolyzed products are formed, where complex
compounds may exist, or where a preliminary adsorption process
may be followed by compound formation. The attempted precipita-
tion of bismuth hydroxide by ammonium hydroxide is likely to
result in a mixture of hydroxide and basic salts in varying proportions
depending upon the anions present, the hydroxyl ion concentration
of the solution, the temperature, &nd other factors such as the method
and rate of mixing the solutions. The coprecipitation of sulfate
with a ferric hydroxide precipitate has been attributed to the forma-
tion of a basic ferric sulfate.

Solid Solutions (Mixed Crystals) as a Cause of Coprecipitation. It
is possible to have a constituent uniformly distributed throughout a
solid, and thus form what is known as a solid solution. For example,
it was observed quite early 16 that if a solution of iodine in benzene was
cooled until solid benzene separated, iodine would be uniformly
distributed throughout the solid phase and the concentration of the
iodine in the solid phase was proportional to the concentration of the
iodine in the liquid phase.

The occurrence of isomorphism was also observed quite early.
Isomorphous substances are those having so nearly the same crystal
configuration that the constituents of one can be substituted in the
crystal lattice of the other without causing an appreciable change in
the crystal.. Thus on evaporating a solution containing two iso-
morphous substances, homogeneous crystals of each substance will be
obtained ; however, each will be found to contain a certain proportion
of the other substance. Since it seemed reasonable to assume that

14 Weiser and Durham, /. Phy9. Chem., 32, 1061 (1928).
i* Van't Hoff, Z. physik. Chem., 5, 322 (1890).

COPRECIPITATION 125

isomorphous substances should have analogous molecular formulas,
this rule was used by Mitscherlich as early as 1820 in deciding which
multiple of the combining weight of an element should be its atomic
weight. An experiment which strikingly shows mixed crystal
formation can be performed by crystallizing potassium perchlorate
from a solution containing even a very low concentration of per-
manganate ions. The crystals will remain pink even after washing
(preferably with a saturated solution of potassium perchlorate) until
the wash solution is colorless.

It has been found, however, that isomorphism is not necessary for
solid-solution formation. The striking coprecipitation of nitrate by
a barium sulfate precipitate, which previously had been attributed to
compound formation (compounds of the type Ba(NOaVBaSO4
having been postulated), has been recently studied 16 by means of
X-ray diffraction experiments with results that indicate that solid
solutions are formed.

Mechanical Inclusion (Occlusion) o$ a Cause of Coprecipitation. In
many cases it has been made evident, especially from microscopic
studies, that portions of the solution have been either entrapped by
crystals or entrained in a mass of crystal capillaries and in this manner
carried out with the precipitate. Such a process is properly termed
occlusion, even though this term is sometimes used to designate
coprecipitation in general. This effect can often be minimized by
proper adjustment of conditions and method of precipitation and is
probably the least serious of the causes of coprecipitation discussed.
An example of it is found in the precipitation of potassium per-
chlorate, where it has been observed 17 that sodium perchlorate and
perchloric acid are entrained by the precipitate.

General observations on coprecipitation effects. 1. Coprecipita-
tion effects often appear to be specific. At the present time the funda-
mental principles of coprecipitation are too little understood and the
conditions affecting it are too varied for general predictions to be
made in regard to the occurrence or magnitude of the phenomenon;
in many cases the effects appear to be controlled entirely by specific
conditions. However, certain general tendencies have been noted
and are discussed below. It should be mentioned that a considerable
amount of intensive work is being carried out at the present time by
various workers in regard to the mechanism of the formation of

" Walden and Cohen, /. Am. Chem. Soc., 67, 2591 (1936).
17 Smith and Ross, /. Am. Chem. Soc., 47, 774 (1926).

126 GRAVIMETRIC METHODS

precipitates, their changes on ageing, the role of adsorption in co-
precipitation, and other related topics; a much better general under-
standing of the subject should result from such investigations.

2. The effect of the concentration of the copredpitated substance.
Where coprecipitation results from adsorption alone, the effect of
the concentration of the adsorbed substance can be predicted from
the adsorption equation; regardless of the cause it is qualitatively
apparent that the greater this concentration the greater the co-
precipitation, other factors being constant.

3. The effect of temperature on coprecipitation. Since adsorption
processes are exothermic, it is to be predicted that the higher the
temperature the less the amount of adsorption taking place. On the
other hand, where hydrolysis is involved in the coprecipitation
process, raising the temperature is likely to increase coprecipitation.
In certain cases where separations are being made which depend upon
maintaining some substance in a supersaturated condition, raising
the temperature appears to break up this supersaturated state and
result in extensive coprecipitation. This effect is noted in the
precipitation of zinc sulfide in the presence of cobalt (see discussion of
P. 61).

4. The effect of the solubility of the copredpitated substance. It has
been observed that, where other conditions are constant, the amount
of coprecipitation occurring increases as the conditions of the pre-
cipitation more nearly approach the saturation value of the co-
precipitated substance. Thus in a determination of the sulfate in a
solution which was approximately 0.04 n. in calcium chloride, a
barium sulfate precipitate weighing approximately 1 g showed a loss
in weight of 16 mg (replacement of barium by calcium), whereas
under the same conditions with magnesium chloride present the error
was 3 mg positive (coprecipitation of barium chloride). 18 In the
separation of the bipositive elements manganese, cobalt, nickel, and
zinc from the tripositive elements aluminum, iron, and chromium
by a carefully controlled ammonia precipitation it was found that
coprecipitation increased in the order named; zinc, the least soluble
of the bipositive hydroxides, was the most extensively carried out
with the tripositive hydroxides (see discussion of P. 51).

5. Coprecipitation and complex-compound formation. It is fre-
quently observed that where a constituent of the precipitate tends
to form a complex compound with some constituent of the solution,
there will be coprecipitation of that compound. When sulfate is

11 Blasdale, Quantitative Analysis, Van N oat rand, 1928, pp. 140 and 188.

COPRECIPITATION 127

precipitated as barium sulfate, ferric iron, which forms a complex
compound of uncertain type with sulfate (or hydrosulfate), is much
more extensively coprecipitated than aluminum (where any such
complex is much less stable), or than ferrous iron.

6. Coprecipitaiion varies with the methods and conditions of pre-
cipitation. Since the physical characteristics of a precipitate,
especially its surface, can be varied so extensively by the conditions of
precipitation, it would be expected that such effects as adsorption
would be similarly influenced. Also, as the concentration of the
adsorbed ion is a factor in coprecipitation effects, it is possible to
vary the effective concentration of such an ion during the formation
of the precipitate (when coprecipitation is usually most pronounced)
by the order of mixing the solutions. Thus in the precipitation of
sulfate by addition of an excess of barium chloride, the coprecipita-
tion of barium chloride will be less if it is added to the sulfate solution
than if the mixing is made in the reverse order; however, if alkali-
metal ions are present in the solution, more alkali sulfate will be
coprecipitated when the barium chloride is added to the sulfate solu-
tion; because of this Popoff and Neumann 19 have recommended the
reverse order of precipitation. The possible effects of the tem-
perature at which the precipitation is made have been already
mentioned above.

General methods for minimizing coprecipitation effects. Before
leaving this topic, certain general methods for avoiding or minimizing
coprecipitation effects are mentioned, even though some repetition is
involved.

1. Avoid introducing or remove substances known to cause co-
precipitation. As an example, in the determination of sulfur in a
pyrite (FeS 2 ) material it is preferable to avoid a fusion process in the
solution of the material because of the introduction of large amounts
of sodium or potassium salts; if nitric acid is used in the solution
process, it should be removed by repeated evaporation with hydro-
chloric acid; and the ferric iron should be either precipitated as
ferric hydroxide or reduced to the ferrous form, which, as mentioned
above, is relatively little coprecipitated.

2. Keep the concentration of coprecipitated substances low. This
can be accomplished by avoiding large excesses of precipitants, or of
acids and bases which have to be subsequently neutralized, and by
making the precipitation from dilute solutions.

3. Properly control the method and conditions of precipitation.

11 Popoff and Neumann, /. Ind. Eng. Chem., Anal. Ed. t 2, 45 (1930).

128 GRAVIMETRIC METHODS

Factors involved would be (1) the adjustment of the volume of the
solution, (2) the order of mixing the reagents, (3) the temperature at
which the precipitation is carried out, (4) the rate of addition of the
precipitant, (5) the stirring of the solution, (6) the time allowed before
filtering the precipitate, and (7) the selection of a suitable medium
for the precipitation.

4. Reprecipitation. As a last resort, it is often necessary to dissolve
the precipitate and to repeat the precipitation. Since in the second
solution the concentration of the coprecipitated substance is relatively
small, the coprecipitation is correspondingly decreased. It is obvious
that a single reprecipitation is effective only when a small proportion
of the absorbed substance is carried out in the first precipitation.

III. Factors Affecting the Composition and Stability of the Weighed
Precipitate

In order to effect a gravimetric determination, it is necessary, after
filtering and washing the precipitate, that it be (1) completely dried,
(2) of definite composition, and (3) sufficiently stable to be precisely
weighed.

Assuming that volatile contaminants are not present (usually as
coprecipitated material), precipitates can be divided into two types,
as follows:

(1) Those which are weighed in the same form as they were filtered.

(2) Those which have to be converted into a more stable and uni-
form compound before they can be weighed.

With precipitates of the first type it is necessary only to remove the
superficial water (adsorbed or entrained) or other solvent medium.
Examples of such precipitates are silver chloride, bromide, iodide, or
thiocyanate, lead sulfate, barium sulfate, bismuth oxychloride, and
nickel dimethylglyoxime. With such precipitates the drying can
usually be accomplished by heating at a relatively low temperature,
in some cases as low as 110C. In certain cases where the water
may be more firmly held, or where a high order of precision is desired
as in atomic-weight determinations higher temperatures must
be used; for example, in order to remove the last traces of moisture
(usually less than 0.01 per cent) from a silver chloride precipitate, it
must be heated until fused. If the precipitate has been collected on a
paper filter, the necessity of completely burning the paper will often
require a much higher temperature than is necessary to dry the
precipitate within the usual analytical limits; for this reason, and
because of the danger of reduction reactions occurring in the burn-

CHARACTERISTICS OF PRECIPITATES 129

ing process, filtering crucibles are advantageously used for such
precipitates.

The possibility of drying precipitates of this first type by treating
them with some volatile solvent which has a great tendency to take
up water, such as ethyl alcohol or a combination of alcohol and ether,
and then removing this solvent by placing the crucible and pre-
cipitate in a desiccator and evacuating for a short time offers certain
possibilities. A considerable saving of time is effected, since the
whole process is carried out at room temperature and the crucible
can be weighed without having to wait the 20 to 30 minutes necessary
for it to cool to room temperature. Also, certain precipitates can
be dried by this process which are unstable if dried by heat. Thus it
has been found that magnesium ammonium phosphate hexahydrate
(MgNH 4 P04 6H 2 0) can be dried in this manner. 20 More recently,
the use of alcohol and ether followed by the aspiration of air through
the precipitate or a brief treatment in a vacuum desiccator has been
proposed by Dick 21 for the rapid drying of quite a number of pre-
cipitates. The accuracy of the method has been questioned 22 and
confirmed by others. 23 Experiments 24 have indicated that when
using sintered-glass filtering crucibles certain precipitates can be
effectively dried by the process, but that the use of ether following
the alcohol is not only unnecessary but in certain cases objectionable,
since its rapid evaporation tends to cool the crucible until moisture
condenses on it. Confirmatory determinations, 25 the results of
which are shown in Table V, show that by sucking the wash water
from the precipitate, washing it with first 95 per cent and then 99.5
per cent ethyl alcohol, aspirating the alcohol, and finally placing the
crucible in a vacuum desiccator and evacuating it with an efficient
water aspirator for 5 minutes (essentially the method used in Optional
Procedure D given below) silver chloride, lead sulfate, and biSmuthyl
chloride precipitates can be so effectively dried that results agreeing
surprisingly well with those resulting from conventional methods of
drying can be obtained.

Use of the alcohol method with students for the gravimetric
standardization of hydrochloric acid solutions has shown that, unless

20 Fales, Inorganic Quantitative Analysis, p. 222, Century, 1925; Wor-
sham, Thesis, Columbia University, 1923.

" Dick, Z. anal. Chem., 77, 352-363 (1929); 78, 414 (1929); 83, 105 (1931).

11 Moser and Von Zombory, Z. anal. Chem., 81, 95 (1930).

" Wassiljew and Sinkowskaja, ibid.. 89, 262 (1932).

24 Unpublished experiments by K. Watanabe and D, DeVault.

M Unpublished experiments by R. C. Custer.

TABLE V

A COMPARISON OF THE USE OF ALCOHOL AND OF HEAT FOR THE DRYING OF

CERTAIN PRECIPITATES

In carrying out these experiments, a definite volume of a solution of the
element being determined was pipeted and the precipitation and washing of
the precipitate carried out in the conventional manner. The crucibles, either
sintered glass or Gooch-type asbestos, were prepared and dried by the alcohol
method outlined above and then weighed. The crucible was then heated (at
150 to 200 when used for silver chloride and bismuthyl chloride precipitates
and at 400 to 500 when used for the lead sulfate) and again weighed. The
precipitate was then collected in the crucible, washed, treated by the alcohol
method, and weighed. It was then heated to the temperature stated and again
weighed. The weights obtained by the "alcohol method" for the empty
crucible, for the crucible and precipitate, and forthe precipitate are shown in
Columns V, VI, and VII following the designation "alcohol" in Column IV;
the weights obtained after drying by heat for the same crucible, for the crucible
and precipitate, and for the precipitate are shown directly below and following
the designation "heat" in Column IV.

Exp.
No.

II
Type
Crucible

III
Pre-
cipitate

IV
Drying
Agent

V VI VII

Weight Found

Crucible

Crucible
and Pre-
cipitate

Pre-
cipitate

Glass

AgCl

Alcohol

26.9606

27.2925

0.3319

Heat

26.9606

27.2925

0.3319

Glass

AgCl

Alcohol

24.5138

24.8440

0.3302

Heat

24.5138

24.8440

0.3302

Asbestos

AgCl

Alcohol

16.2366

16.5678

0.3312

Heat

16.2366

16.5677

0.3311

Asbestos

AgCl

Alcohol

15.1706

15.5017

0.3311

Heat

15.1706

15.5017

0.3311

Glass

PbSO*

Alcohol

26.9602

27.3263

0.3661

Heat

26.9600

27.3257

0.3657

Glass

PbS0 4

Alcohol

24.5091

24.8752

0.3661

Heat

24.5089

24.8747

0.3658

Asbestos

PbS0 4

Alcohol

16.3679

16.7329

0.3650

Heat

16.3671

16.7324

0.3653

Asbestos

PbSO 4

Alcohol

15.2093

15.5749

0.3656

Heat

15.2091

15.5746

0.3655

Glass

BiOCl

Alcohol

26.9605

27.2493

0.2888

Heat

26.9599

27.2487

0.2888

Glass

BiOCl

Alcohol

24.5085

24.7958

0.2873

Heat

24.5082

24.7951

0.2869

Asbestos

BiOCl

Alcohol

15.2893

15.5765

0.2872

Heat

15.2889

15.5758

0.2869

Asbestos

BiOCl

Alcohol

14.7413

15.0290

0.2877

Heat

14.7411

15.0287

0.2876

a On cleaning and weighing the crucible for subsequent determinations, a
loss in weight of 0.0047 g was noted. It has been noted that an occasional glass
crucible will show an erratic loss in weight from time to time, suggesting that
the mat had not been perfectly sintered together.

130

CHARACTERISTICS OF PRECIPITATES 131

a very thin mat of asbestos is used, complete drying is frequently not
obtained unless the precipitate and crucible are repeatedly evacu-
ated; therefore, the method may lose much of its rapidity unless
used with a fabricated filtering crucible. Because of the quickness
with which this method of drying certain precipitates can be carried
out when used with sintered-glass crucibles, an optional procedure
providing for its use is given below.

Precipitates of the second type can be divided into two different
sub-groups. In drying those of the first sub-group a relatively high
temperature is necessary because there has to be a complete elimina-
tion of what may be called water of constitution. This group in-
cludes the hydrous oxides (ferric and aluminum hydroxides, and
silicic and "metastannic" acid are examples), which are of very
uncertain composition and which have to be converted to the oxides
before a stable compound of definite composition is obtained. The
second sub-group of precipitates is composed of those which may be
partly decomposed upon drying by heat, and therefore it is necessary
to convert them into a compound of more uniform and stable com-
position before they can be weighed. Examples are magnesium
ammonium phosphate (MgNHiPCVBH^O), which is converted by
heat to magnesium pyrophosphate (Mg 2 P207), and calcium oxalate
(CaC 2 O4-H 2 0), which is weighed as either the carbonate or the
oxide.

The treatment of these precipitates will vary greatly; thus in order
to convert calcium oxalate to calcium carbonate the temperature
must bo controlled between 475 and 525C., 26 or if the ignition is
carried out in an atmosphere of carbon dioxide, the temperature
can be raised to 700C. 27 The complete dehydration of aluminum
hydroxide or of silicic acid requires heating above 1000C. for some
time.

It should be emphasized that certain precipitates which are
weighed in the same form as they separate are heated at higher
temperatures because of the presence of coprecipitated compounds;
thus barium sulfate, which can be dried at a relatively low tem-
perature, is usually heated to approximately 900 in order to reduce
the error caused by coprecipitated material. In Table VI are
shown some of the precipitates most frequently used in gravimetric
analysis, the form in which they are weighed, and the temperatures
to which they are usually heated.

28 Willard and Boldyreff, /. Am. Chem. Soc., 52, 1888 (1930).
87 Foote and Bradley, /. Am. Chem. Soc., 48, 676 (1926).

132 GRAVIMETRIC METHODS

TABLE VI
. PRECIPITATES FREQUENTLY USED IN GRAVIMETRIC ANALYSIS

Compound
Precipitated

Compound
Weighed

Temperature
Used CC.)

Remarks

AgCl

130 to 150

Estimation of Ag or Cl.

PbSO 4

500 to 600

Precipitate washed with

dilute H 2 SO 4 .

BiP0 4

200 to. 300

Decomposes above 650.

BiOCl

105 to 110

CdS

CdSO 4

450 to 500

CdS treated with H 2 S0 4

HgS

105 to 110

HR.C1,

Hg.Cl,

105 to 110

As 2 S,

As 2 S 8

105 to 110

MgNH 4 AsCV6H 2 6

Mg 2 As 2 O7

850 to 950

Estimation of As or Mg.

Sb 2 S 3

Sb 2 Ss

280 to 300

SnO 2 (H 2 0)x

Sn0 2

1100 to 1200

Fe 2 8 (H 2 0)x

Fe 2 Os

1000 to 1100

In oxidizing atmos-

phere.

ZnNH 4 P0 4 -6H 2

/ZnNH 4 PO 4
\Zn 2 P 2 O 7

110 to 135
850 to 950

/T^Q

/ZnSO 4

450 to 500

ZnS treated with H 2 SO 4|

IZnO

850 to 950

decomposes above 700.

NiC 8 H 14 N 4 4

110 to 120

CoS

CoS0 4

450 to 500

CoS treated with H 2 S0 4 .

MnS

MnSO 4

450 to 500

May decompose above

550.

MnNH 4 P0 4 -H 2

Mn 2 P 2 7

850 to 950

A) 2 0,(H 2 0),

A1 2 0,

1050 to 1200

BaSO 4

350 to 900

For Ba or S.

SrSO 4

SrSO

450 to 500

CaC 2 4 -H a O

(CaCO,
ICaO

475 to 525
1200

MgNH 4 PO 4 -6H 2

Mg 2 P 2 7

1000 to 1100

For Mg or P.

KC1O 4

KC10 4

350

NaCl

600

SiO.(H,0).

SiO,

1200

THE OPERATIONS OF GRAVIMETRIC ANALYSIS
The Filtering of Precipitates

The operation of filtering has for its object the separation of the
precipitate from the solution in which it has been formed. This is
accomplished by passing the solution through some porous medium
which will be capable of passing the solvent and the constituents
which are in true solution, but which will retain even very finely
divided precipitates down to those having diameters of approxi-
mately 1 to 2fjL (lju, called one micron, is 0.001 millimeter). With
particles of less than 0.2 to 0.5^ colloidal properties become evident,

FILTERING OF PRECIPITATES 133

and although it is possible to devise filtering media which will retain
even colloidal suspensions, special apparatus is required, and the
rate of filtration is so slow that it is not practical for analytical pur-
poses to try to retain particles of less than approximately 1/z. Cer-
tain restrictions are imposed upon the material used for a filtering
medium. (1) It must be relatively inert to the various solvents and
solutions to be filtered. This implies that it must not be weakened
or disintegrated during the filtering process, and that it must not
introduce any contamination into the solutions passed through it.
(2) It must not retain, either by adsorption or by absorption the
soluble constituents of the solution. (3) For quantitative purposes
(where the precipitate is to be subsequently weighed) it must either
remain constant in its weight or be capable of complete removal
for example, by being burnt or volatilized. The various media
most commonly used in analytical processes are discussed below.

Filtering media. The filtering media which are most extensively
used in analytical work are paper, asbestos, glass, quartz, porcelain,
and platinum. The relative advantages of these different agents
are discussed below.

Paper filters. The medium used almost exclusively for qualitative
work and still used most frequently for quantitative purposes is pure
cellulose paper. Paper filters have several disadvantages: (1) They
are not inert, being attacked by concentrated solutions of both alka-
lies and acids, and by powerful oxidizing agents; (2) they are lacking
in mechanical strength, are not readily adapted to vacuum filtration,
and often disintegrate, introducing fibers into otherwise clear solu-
tions; (3) they have the property of adsorbing constituents from the
solutions passed through them; and (4) they cannot be dried to a
constant weight for precise quantitative work. Their chief advan-
tages lie in their cheapness, availability, and superior filtering effi-
ciency, especially in the filtration of gelatinous precipitates. This
advantage over other filtering media is in part due to the larger
surface exposed and a larger ratio of pore space to total surface.
Paper filters are obtainable in various degrees of porosity, those of
more open texture being suited for the rapid filtration of easily re-
tained precipitates and the more dense ones being adapted for finer
precipitates; the finer filters will retain particles of approximately
2ju in diameter, and the coarser ones will retain particles of about
6/u in diameter.

When used for gravimetric work, paper filters have a serious dis-
advantage in that they cannot be dried to constant weight and

134 GRAVIMETRIC METHODS

therefore have to be ignited or burnt before the precipitate can be
weighed. This process is often unsatisfactory for the following
reasons: (1) It may require a higher temperature than is necessary
to dry the precipitate and one at which the precipitate may be un-
stable. (2) Many precipitates (for example, the silver halides and
barium sulfate) are appreciably reduced during the burning of the
paper and have to be subsequently treated to correct this reduction.
Where this reduction is serious, it is necessary to remove most of the
precipitate from the paper and burn it separately, a process which
invites mechanical loss. (3) There always remains after burning
the paper a nonvolatile residue, or ash.

Quantitative (ashkss) paper filters. In order to reduce this error,
paper filters which are intended for quantitative work are washed
in the process of manufacture with hydrochloric and hydrofluoric
acids to remove as completely as possible the nonvolatile inorganic
salts and the silicon compounds. This process increases the cost
of these filters but reduces the weight of the ash to a very low value.
The average weight of this ash in various types of filters is shown in
Table VII.

TABLE VII
THE WEIGHT OP THE ASH OP VARIOUS TYPES OP FILTERS

Diameter of Filter
(cm)

Qualitative
(mg)

Quantitative
(mg)

7
9
11

0.2 to 0.9
0.4 to 1.3
0.6 to 3. 3

0,02 to 0.05
0.03 to 0.08
0.05 to 0.11

It is seen that the ash of the quantitative filters has been reduced
to such a small value that it can be precisely corrected for or even
neglected in most routine work. Because of the development of
several types of satisfactory filtering crucibles, the use of filter
papers is no longer recommended for precipitates which can be dried
at low temperatures (100 to 500C.) or which have to be separated
from the filter before the paper is burned.

Asbestos filters. For quantitative work asbestos is almost uni-
versally used in porcelain crucibles with perforated bottoms (so-
called Gooch crucibles), although the form originated by Gooch 28
was of platinum (see Fig. 22), When used for qualitative work,

Gooch, Proc. Amer. Acad., 13, 342 (1878).

FILTERING OF PRECIPITATES

135

asbestos filters can be made upon cir-
cular bevel-edged perforated porcelain
plates or upon a wad of glass wool sup-
ported in an ordinary funnel (see Fig. 23).
Asbestos is a calcium-magnesium sili-
cate, and that used for quantitative pur-
poses should be of a white, silky, long-
fibered variety, free from magnetite and
iron. It has been found that even the
asbestos purchased and designated specif-
ically "for quantitative analysis" should
be acid-treated and washed before being
used. Asbestos can be safely heated to
moderately high temperatures and, when
used for gravimetric work, Gooch
crucibles (and other filtering crucibles)
have the great advantage over paper that
they can be dried to constant weight,
and the disadvantages connected with
the burning of a paper filter are elimi-
nated. Asbestos filters are intended for
use with suction. With coarse crystalline
precipitates and with a thin properly
made mat, they afford a more rapid fil-
tration than paper, although asbestos
(and other filtering media such as glass
and porcelain) is not as satisfactory as
paper for gelatinous precipitates, since
these tend to clog the relatively small
filtering surface. Asbestos is not so satisfactory for precipitates
which have to be heated to very high temperatures, since some
specimens tend to lose weight, and the asbestos insulates the pre-
cipitate from the heat of the burner; the perforated porcelain cruci-
bles are ateo subject to cracking at high temperatures unless used
with extreme care. Asbestos filters (and filtering crucibles in gen-
eral) should be heated to a constant weight at the temperature subse-
quently to be used for drying the precipitate. They possess the
disadvantage that they require more time in their preparation than a
paper filter; also, some asbestos tends to absorb water so rapidly as to
cause difficulty in weighing.

Filtering crucibles. Filtering crucibles of various materials are

Fig. 22. Asbestos Filter
in Gooch Crucible with
Holder for Quantitative
Work.

136

GRAVIMETRIC METHODS

Glass Woo!

Fig. 23. Asbestos Filters for Qualitative Work. Suction Flasks.

now offered which have a filtering element of the same material
fabricated in place as an integral part of the crucible; among the
materials so used are glass, quartz, porcelain, and platinum. Such
crucibles possess the advantage over paper (as does the Gooch-typc
asbestos filter) that they can be dried to constant weight. They
possess the advantages over the asbestos filter that they are ready
for immediate use, that the filtering elements can be obtained with
various pore sizes, that they are more quickly brought to a constant
weight, and that they are less hygroscopic thereafter. They possess
the disadvantage that special methods are necessary in order to clean
the precipitate from them. The characteristics of these crucibles
are discussed briefly below.

Sintered-glass crucibles. 29 The filtering element in the bottom of
these crucibles is composed of glass particles which have been pow-
dered and graded as to size, and then sintered together (in place)
to form a filtering crucible of the desired porosity (see Fig. 24).
They possess the advantage of being transparent, of quickly coming
to constant weight, and of being easy to clean. With care in heating

29 At the present time these crucibles seem to be made only of Jena glass by
Schott and Gen. of Jena.

WASHING OF PRECIPITATES 137

and cooling, they can be used up to 600C. They
can be used with strong acids (except hydro-
fluoric) and with moderately concentrated bases.
A tendency of some of these crucibles to lose
weight after being heated much above 150C. has
been noted; otherwise they are excellent.

Sintered-quartz crucibles. These crucibles are
similar to the glass crucibles but are constructed
of clear quartz and possess the advantage that Fig. 24. Sin-
they can be used for temperatures up to 1200C. tered-Glass Fil-
They are expensive. tering Crucible -

Porous porcelain crucibles. These crucibles are similar to sin-
tered-glass crucibles, but are constructed of porcelain and have a
porous porcelain filtering element in the bottom. They can be ob-
tained in any desired porosity, and can be heated up to 1000 to
1200C.

Platinum filtering crucibles (Munroe type). These crucibles, pro-
posed by both Munroe 30 and Neubauer, 31 are similar in design to a
Gooch crucible except that they are constructed of platinum and the
filtering element is a porous layer of platinum; this is prepared by
building up a layer of ammonium chloroplatinate of the desired thick-
ness and then igniting. These crucibles have many advantages.
They can be made to retain fine precipitates, yet filter rapidly; they
quickly attain constant weight and can be heated to high tem-
peratures, and because of their superior heat conductivity, precipi-
tates in them are not insulated as they are by asbestos mats. They
are usually formed with a cap which fits over the bottom and protects
the precipitate from possible action of gases from burners. They
arc expensive, however, and the mats have to be frequently re-
newed, especially when used at high temperatures. They are recom-
mended for use in highly precise work. The following references are
valuable for details of the preparation, use, and cleaning of these
crucibles:

Snelling, Chem. News, 99, 229 (1909); J. Am. Chem. Soc., 31, 456
(1909); Swett, ibid., 31, 928 (1909).

The Washing of Precipitates

In washing a precipitate, the first consideration is the selection
of the most effective washing medium. It is not often that this will
be pure water. In making this selection, the following factors have

Munroe, /. Anal. App. Chem., 2, 241 (1888); Chem. News, 38, 101 (1888).
11 Neubauer, Z. anal. Chem., 39, 601 (1900).

138 GRAVIMETRIC METHODS

to be considered: (1) the solubility of the precipitate, (2) the stability
of the precipitate, (3) the nature of the material to be washed out,
(4) the tendency of the precipitate to become colloidal, and (5) the
effect of the wash solution or its constituents on the drying and heat-
ing of the precipitate.

If the precipitate is appreciably soluble, it will be necessary to re-
duce its solubility in the wash medium. This is often accomplished
by (1) the addition of a common ion; (2) the use of some special solu-
tion, an example being the washing of an magnesium ammonium
arsenate precipitate (P. 88) with dilute ammonium hydroxide; (3)
the addition of organic liquids (frequently alcohol) ; and (4) the use
of cold wash solutions (especially where the solubility of the salt has
a high temperature coefficient).

Certain precipitates may undergo changes if washed with water;
thus bismuth phosphate may be partly converted into a basic salt.
Other precipitates, certain sulfides for example, may be partly oxi-
dized unless the wash solution is freed of oxygen, or an inhibiting
substance (such as hydrogen sulfide) is present.

It is essential that the nature of the ions and compounds in the
filtrate be considered in selecting the wash medium. This is espe-
cially important in separations, for, if the precipitation has taken
place in a strongly acid solution and the precipitate were washed
with water, some of the constituents of the filtrate may be partly
precipitated. Thus, arsenic can be separated from antimony by
precipitation as sulfide from a 9 f . hydrochloric acid solution, but if
the solution adhering to the precipitate were diluted by addition of
wash water, precipitation of antimony sulfide would be likely to
result.

The method of washing will be determined by the nature of the
precipitate and the method of filtering. The advantages of decan-
tation apply to the washing as well as the filtering process; settling
the precipitate by centrifuging will often facilitate a decantation
process. The type of filter used and the nature of the precipitate
will indicate whether suction can be us'ed to advantage.

In washing a precipitate, it is usually desirable, because of solu-
bility effects and to prevent unduly diluting the filtrate, to use the
smallest possible volume of wash solution. For these reasons it is
usually more effective to wash with repeated small portions of wash
solution than with a smaller number of larger portions. This can

be shown by an application of the expression

/nr / ' r_ \ s^i

~ \y + v r j

P. XVIII] HYDROCHLORIC ACID SOLUTIONS 139

where Co is the original concentration of any material in the solution
left with a precipitate, C n the concentration after n washings, V the
volume of wash solution added, and V r the volume of solution left
after each washing. Thus if a precipitate weighing 0.2 g has left
with it and the filter 0.5 ml of solution and 0.1 g of contaminating
material, and is washed with 20 ml of water, added in one case in
10-ml portions and the other in 5-ml portions, it is seen that in the
first case there will be left 2.3 X 10~ 4 g of material and in the second
only 6.8 X 10~ 6 g, these corresponding to an error of 0.1 per cent
and 0.004 per cent respectively.

These calculations indicate that precipitates can be washed with
relatively small volumes of water; however, the assumption has been
made that the foreign material has been merely held in the solution re-
maining on the precipitate and filter and that it has been uniformly
distributed through each portion of wash water. In many cases, espe-
cially when dealing with bulky precipitates, it may be difficult to
get the wash water into effective contact with all portions of the
mass, and, more important yet, the foreign material is likely to be
adsorbed on the precipitate. Because of this latter effect, calculations
such as the above are likely to be extremely misleading, and the
only safe way to determine when a precipitate has been washed is by
making suitable tests on the wash solution

Procedure XVIII: THE GRAVIMETRIC STANDARDIZATION
OF A HYDROCHLORIC ACID SOLUTION. In the procedures below
the hydrochloric acid solution is standardized by pre-
cipitating and weighing the chloride as silver chloride.
The discussion of the general principles and operations of
gravimetric analysis ju3t preceding should be read and
understood before carrying out these procedures.

The procedures be ? ow provide for the use of various fil-
tering media and foif the optional use of alcohol as a drying
agent when using a sintered-glass or other fabricated filter-
ing crucible.

Precipitation of .he silver chloride. Pipet a 25-ml portion of
the solution to be standardized into a 400-ml beaker (Notes
1, 2) and add approximately 150 ml of water. Add drop-
wise the calculated volume of a 0.2 n. AgNOa solution (Note
3), stirring the rjiixture vigorously (do not spatter or splash
it) during the addition, and then add 5 ml in excess (Notes
4, 5, 6). Do not carry out this precipitation or the subse-
quent operations in sunlight or very intense artificial illumi-
nation. Heat :he mixture to approximately 60C. and stir

140 GRAVIMETRIC METHODS [P. XVIII

it frequently until the precipitate settles rapidly, leaving
a clear solution (Note 7). Keep the beaker covered when
not stirring the solution.

Filtering and washing the precipitate. Optional method A .
Using a Gooch-type filtering crucible (a perforated porcelain
crucible with asbestos filtering mat). Preparation of the
crucible. Select a crucible of the proper size (for this pre-
cipitate one of about 20 to 30 ml capacity), clean it thor-
oughly, and fit it into a filter crucible holder (be sure that
the rubber does not extend beyond the bottom of the cru-
cible so that the filtrate comes in contact with it). Fit the
stem of the crucible holder or of the funnel in a one-hole rub-
ber stopper and support it in a suction filter flask of the type
shown in Fig,22 (Note 8). Shake up a water suspension of
acid-washed chloride-free asbestos (Note 9) and pour it
through the crucible (with the suction turned off) until a
uniform mat of asbestos of about 1 to 2 mm thickness is ob-
tained(Note 10.) Place a perforated porcelain plate on the
mat and, applying very gentle suction, pour into the cru-
cible just enough of the asbestos suspension to fill around the
edges and the perforations of this plate (Note 11), Wash the
mat until the water passing through is free of asbestos fibers.
(Carefully guide the first portion of water onto the plate with
a stirring rod so that the mat is not disturbed, and thereafter
do not allow one portion to run out entirely, thus exposing
the mat, before adding the next.) Place the crucible in a
small beaker, cover it with a glass and place it in an oven
at 110 to 120C. for 1 hour (Note 12). Place the crucible
in a desiccator (see Note 2, P. V), allow it to cool for 20 to
30 minutes, and weigh it. Again heat the crucible for 20
minutes, cool, and weigh. Repeat this process until the
weight remains constant within 0.2 mg (Note 13).

Filtering and washing the precipitate. Insert the crucible
in the holder and, applying a very slight suction (Note 11),
decant the clear solution through the filter (Note 14).
Carefully guide the first portion of the solution onto the
plate with a stirring rod, and thereafter do not allow the
crucible to become empty (Note 15). Note carefully if
any asbestos fibers appear in the first portions of the filtrate ;
in this case it must be refiltered (Notes 16, 17). Prepare a
wash solution by adding 5 ml of 6 n. HNOa to 500 ml of

P. XVIII] HYDROCHLORIC ACID SOLUTIONS 141

water and put this in a small wash bottle. (Label this
"special wash bottle. ") Wash the precipitate by decanta-
tion with three 25-ml portions of this wash solution, trans-
. fer it to the crucible (Note 18), and wash with the dilute
acid until 2 to 3 ml of the washings give no precipitate
upon the addition of a single drop of HC1 (Note 19). Fi-
nally, wash with 5 ml of water, added dropwise.

Drying the precipitate. Place the crucible in a covered
beaker and heat it at 120C. for 1 hour (Note 20). Cool
the crucible for 30 minutes in a desiccator and weigh. Re-
peat the heating and weighing until the weight is constant
to 0.2 mg. Calculate the normality of the hydrochloric
acid solution.

Notes:

1. The concentration of the HC1 will determine the volume to be taken.
If less than 25 ml are pipeted, the precision of that measurement is not
sufficient for a standardization; it is also desirable that the weight of the
silver chloride precipitate be not less than 0.3 g (to reduce the weighing
error) and not larger than 1 g (to avoid difficulty in the filtering and wash-
ing of the precipitate). Duplicate (for very precise work, triplicate) deter-
minations should be made.

2. The above procedure can be applied to the determination of the chloride
in soluble salts. In that case weigh out an amount of the sample which
contains from 0.15 to 0.25 g of chloride, and dissolve this in 150 ml of water
to which has been added 0.25 ml of 6 n. HNOa. If bases or the salts of weak
acids are present, the solution should be first neutralized and the excess of
acid indicated above added.

3. The calculated volume of the AgNOa may be measured in a graduate
and added from a dropper, or it may be added from a graduated pipet or
buret. In order to avoid splashing, the solution should be delivered very
close to the surface of the solution or flowed down the side of the beaker.

4. Stirring rods used for this purpose should have their ends fire-polished
and should not be rubbed against the walls of the vessel; otherwise small
particles of glass may be chipped off or the inside of the vessel scratched;
precipitates form in such cracks and are difficult to remove.

5. Except when standardizing a solution whose concentration is already
closely known, the amount of silver nitrate to be added cannot be previously
calculated. In such cases the equivalent amount can be determined by
noting when no more precipitation takes place upon adding a portion of the
reagent. This is done by adding the AgNOs in successively smaller por-
tions as the rate of precipitation is seen to be decreasing, until finally the
effect of a few drops is noted. As the equivalence-point is approached,
the mixture must be shaken vigorously after each addition of AgNOs and
the precipitate allowed to settle so that the effect of the next portion can be
observed in the supernatant liquid. Usually the precipitate will remain

142 GRAVIMETRIC METHODS [P. XVIII

colloidally dispersed until near the equivalence-point, when a distinct coagu-
lation will be observed. See P. 22 and P. 27 for discussions of |he applica-
tion of this method of determining the "end-point" to the quantitative
estimations of silver and of chloride. After the equivalent amount of silver
nitrate required has been noted, the 5-ml excess is added.

6. Vigorous shaking is very effective in coagulating a silver chloride pre-
cipitate. For this reason a conical flask with a ground-glass stopper can be
used advantageously in this procedure, as the stopper can be inserted after
each addition of reagent and the flask shaken vigorously. If, before re-
moving the stopper, a few drops of water are poured around it, this water
will be drawn into the flask as the stopper is withdrawn and will effectively
wash down any precipitate adhering to the stopper or the neck of the flask.

Do not heat these flasks or allow them to cool with the stopper in place; it is
not advisable, nor should it be necessary, to use vaseline or other lubricants
on the ground-glass joint.

7. It is a general rule that a test should be made to ascertain if sufficient
of the precipitating reagent has been Used by adding a small additional
amount to the clear supernatant liquid, or to the first portion of the filtrate
coming through the filter, and then, after allowing sufficient time, to note if
any additional precipitate forms.

8. Suction flasks are conical flasks made with a side-neck and with thick
walls to withstand the external pressure when a vacuum is produced. It is
unsafe to heat solutions in them.

When a water aspirator is used for suction, the possibility of contamina-
tion of the filtrate by tap water being sucked back into the filter flask should
be prevented. This is done by fitting a conical flask with a two-hole stopper
carrying an inlet and outlet tube and then connecting the inlet tube to the
filter flask and the outlet tube to the aspirator; such an arrangement also
safeguards against solution in the filter flask overflowing into a vacuum
system.

9. It has been found that even the asbestos purchased and designated
specifically for "quantitative analysis" should be further acid-washed and
treated before being used (see the Appendix). When asbestos is used for a
chloride or silver determination, the analyst should assure himself that it is
chloride-free.

10. The thickness of mat to be used will vary with the type of precipitate
to be retained. With a silver chloride precipitate that has been properly
coagulated, a thin, moderately fast filtering mat may be used. When
looking at the light through such a mat, the perforations of the crucible
should be dimly visible; by such an inspection imperfections in the mat may
often be detected. If the holes are not discernible, or if, after applying the
top plate with its film of asbestos, water runs through the filter slowly with
full suction applied, the mat had better be remade. The thinner the mat
(which will retain the precipitate) the more rapid the filtration and washing
of the precipitate, and the more quickly the crucible can be dried and brought
to constant weight.

11. This plate (the so-called Witte plate) and the additional asbestos
serve to prevent the disruption of the mat, and not as an aid to filtering.
If care is taken when pouring solutions upon the mat, the plate may be dis-

P. XVIII] HYDROCHLORIC ACID SOLUTIONS 143

pensed with and more rapid filtering obtained. When adding a solution
to the empty crucible, a slight suction should always be applied, and the
solution should be guided with a stirring rod held against the plate, or, if the
plate is not used, with the stirring rod held just above the mat. Too strong
suction should be avoided, since it may break the mat and in many cases
will so pack the precipitate against the mat as to clog it and slow the rate
of flow.

12. After being cleaned, and especially when hot, crucibles should not be
placed directly on the floor or shelves of ovens, or on desk tops and such
surfaces ; materials thus picked up may contaminate the analysis or change
the weight of the crucible. For the same reason, crucibles should be handled
with a clean pair of crucible tongs and as little as possible with the hands.

The beaker containing the crucible should be covered to prevent particles
of dust, rust, and so forth, from falling into the crucible. More rapid drying
will be obtained if the cover glass is supported by a glass triangle, or on three
small lengths of glass rod which have been bent into the shape, 0, and
hung over the side of the beaker.

To prevent confusion when running duplicate analyses, crucibles should
be marked before being used, either by filing notches on the edges or, prefer-
ably, by means of a "china-marking" pencil. The pigment of these has an
inorganic base, usually iron, and, if the crucible is lightly marked and
then heated, this residue fuses into the surface. The beaker itself should
be marked on the roughened place provided for that purpose. (See Note
6, P. V.)

13. As a general rule when a crucible is being prepared, it should be
treated with the same wash solution and heated to the same temperature
as is to be subsequently used with the precipitate.

14. As a general rule, precipitates should be coagulated and then allowed
to settle before beginning the filtration. Then as much as possible of the
clear solution should be decanted through the filter, care being taken not to
stir up the precipitate. In this manner it is frequently possible not only to
filter a mixture completely, but also to wash the precipitate without remov-
ing it from the precipitating vessel. This is termed filtering by decantation
and should be practiced whenever the nature of the precipitate permits,
since it is a more rapid and efficient process. This is due to the fact that,
since the precipitate usually partly clogs the filter, making the filtration
much slower, it can be much more thoroughly treated with the wash solu-
tion in the beaker than on the filter.

After adding a portion of wash solution and thoroughly mixing it with the
precipitate, it is an advantage to allow the vessel to stand in an inclined
position. The precipitate then tends to settle compactly into the space at
the juncture of the bottom and side of the vessel and is less likely to be
stirred up when the vessel is further tilted to pour out the clear solution.

15. Whenever a solution is poured from one vessel to another or to a
filter, the liquid should be guided by means of a stirring rod placed against
the lip of the vessel. This not only prevents splashing of the solution on the
filter or in the receiving vessel, but prevents the solution from running down
the under side of the lip of the vessel from which it is poured. With some
vessels the lip is so constructed that this may take place even with the use

144 GRAVIMETRIC METHODS (P. XVIII

of a stirring rod; in such cases rub the under side of the lip with the thinnest
possible film of vaseline- or other grease.

16. Having successfully filtered a precipitate from the solution in which
it was formed, it is usually an advisable precaution to remove this solution
from the receiving flask before beginning the washing process. Many pre-
cipitates that filter perfectly from their original solution tend to become
colloidal (or, as it is termed, to "peptize") and pass through the filter when
treated with a wash solution; should this occur, or should a break subse-
quently occur in the filter, much less solution will have to be refiltered.

17. The original filtrate and the first portion of wash water should be
treated with a slight excess of hydrochloric acid, the precipitate allowed to
settle, the solution decanted away, and the residue transferred to the "silver
residue" container. After this procedure is finished, the weighed silver
chloride precipitates should also be added to the same container for later
recovery.

18. When transferring a precipitate from a vessel, extreme care should be
taken that none is lost by spattering. Suspend the precipitate in a small
amount of the wash solution and slowly drain it from the beaker into the
crucible, using the stirring rod to guide it to the bottom of the crucible.
While holding the inclined beaker with the stirring rod held against the lip
in one hand, direct a gentle stream of the wash solution (contained in a-
small wash bottle) against the bottom and sides of the beaker until all of the
precipitate has been washed out. If the precipitate sticks to the sides
of the beaker (and as a general precaution), gently rub the entire inner
surface of the beaker with a "policeman/' a stirring rod capped at one end
with a short piece of soft gum rubber tubing specially made for the purpose
by being sealed at one end. ("Policemen" should not be used as stirring
rods, since they are attacked by some solutions, may disintegrate, or may be
hard to clean if left in the solution when the precipitation takes place.)

If difficulty is experienced in removing the precipitate from the sides of
the vessel (or the "policemen"), it may be dissolved by adding about 10
drops of ammonia to 2 to 3 ml of water and flowing this solution over the
surfaces to be cleaned. This solution should then be made just acid with
HNOs, a drop of the silver nitrate solution added, the solution diluted to
about 10 ml, and the precipitate coagulated by shaking or heating and then
collected on the filter and washed with the remainder of the precipitate.

19. It is impossible to calculate or state how much wash water will be
required for a given precipitate, since the physical nature of the same com-
pound will vary enormously with the conditions of precipitation, as well as
with the nature of the other ions present in the solution. Therefore, it
should be an invariable rule to continue the washing until some substance
known to be present in appreciable amounts in the original filtrate, and for
which a sensitive test can be made, is shown to be absent from the wash
solution.

When washing a precipitate, the original solution and each successive
portion of wash water should be allowed to drain from the filter before the
next portion is added. It has been shown in the discussion that the amount
of wash solution required will be determined largely by the volume of solu-
tion left each time with the precipitate and filter, and also that it is more
effective to wash with several small portions than with one large portion
of the same total volume. However, a precipitate should never be allowed

P. XVIII] HYDROCHLORIC ACID SOLUTIONS 145

to dry or cake and crack before and during the washing; when using an
asbestos filter, it is also desirable that sufficient solution be left in the crucible
to protect the mat when another portion is added.

20. If an oven or electric furnace is available in which the temperature
can be controlled between 150 and 200C., the precipitate can be first dried
for 20 minutes at 110C. and then for the same length of time at the higher
temperature. The precipitate should be gradually heated at first, or it
tends to "cake" and enclose the water, so that a fusion is necessary there-
after to complete its expulsion. If the precipitate is heated to this higher
temperature the crucible should be similarly treated before the initial
weighing.

Optional method B. Using a sintered-glass , porous porce-
lain, or platinum-sponge filtering crucible. Since the filtering
element has been fabricated in place, these crucibles are
ready for use as soon as they are cleaned (Note 1). The
precipitation of the silver chloride, the preliminary heating
of the crucible, the filtration, the washing, and the final
heating of the crucible and precipitate are carried out as
directed above.

Note:

1. Sintered-glass and porcelain crucibles may be cleaned by the use of any
acid cleaning solution except HF; prolonged treatment with concentrated
alkalis is not advisable, although ammonia and cyanide solutions can be
used. The proper solution to use will be determined by the substances it is
desired to remove from the crucible. The solvents mentioned below are
applicable.

Platinum ware should not be treated with an acid solution containing an
oxidizing agent and a chloride, bromide, or iodide; even ferric chloride will
attack platinum. Swett, J. Am. Chem. Soc., 31, 932 (1909), has made an
extended study of the various solvents which may be used with platinum
crucibles and which are most effective in dissolving the precipitates com-
monly used in gravimetric analysis. Those most generally effective include

(1) sulfuric acid with either nitric acid or with ammonium chloride added,

(2) nitric acid, (3) hydrochloric acid with ammonium chloride or oxalic acid
added, (4) sodium hydroxide, (5) sodium sulfide, (6) ammonium hydroxide,
and (7) sodium cyanide. The original article should be consulted for a list
of about 40 solvents with the precipitates for which each is most effective.

Optional method C. Using a paper filter. Carry out the
precipitation as directed in Method A. Select a quantita-
tive filter of the proper size and porosity (Note 1), fit it
to a funnel, and then filter and wash the precipitate by
decantation as directed in Method A above. Cover the
funnel (Note 2), place it in a drying oven (at 100 to 110C.),
and leave until dry.

146 GRAVIMETRIC METHODS [P. XVIII

Clean and mark a porcelain crucible, place it on a clean
triangle (Note 3), warm it gently at first, and then heat it
for 10 minutes (Note 4). Do not cause the crucible to
glow. Allow the crucible to cool somewhat, place it in a
desiccator for 20 minutes, and weigh. Repeat the heating,
cooling, and weighing until the weight is constant to 0.2 mg.

Remove the filter paper from the funnel (Note 5) and
transfer most of the precipitate to a clock glass (or smaller
piece of glazed paper) placed in the center of a piece of
black glazed paper, about 10 inches square (Note 6).
Cover the clock glass containing the precipitate with a
larger clock glass. Fold the filter paper, turning in the
edges, and roll it up so that the portion containing the pre-
cipitate is at one end of the roll. Wind spirally around the
other end a piece of platinum wire about 4 inches long, one
end of which has been sealed into a glass rod (Note 7).
Place the crucible on the glazed paper, hold the roll of
filter paper vertically over the crucible with the end con-
taining the precipitate uppermost, and ignite it at the top
with a very small oxidizing flame. As the roll burns down,
use the flame to burn completely any particles of charred
paper which remain. Collect, the ash in the crucible.
Brush any ash from the wire or from the paper into the
crucible with a small camel's-hair brush or feather. Place
the crucible on a triangle, inclining it with the bottom
against one side of the triangle, and heat it with a small
flame until any particles of carbon are burned. Do not
overheat the crucible (Note 8). Let the crucible cool
and moisten the residue with 2 drops of 16 n. HNO 3 , then
with 1 drop of 6 n. HC1. Evaporate the acids by playing
a small flame around the upper walls of the crucible. Care
must be taken not to hurry the process, or spattering will
result. Cool the crucible sufficiently to place it on the
glazed paper, and very carefully transfer the main portion
of the precipitate to it from the watch glass. Add to the
precipitate 2 drops of 16 n. HNOs and 1 drop of 6 n. HC1
and again evaporate the acids. Heat the crucible, slowly
raising the temperature, until the precipitate just begins
to fuse around the edges. Let the crucible cool somewhat,
place it in a desiccator for 20 minutes, and weigh. Again
heat to incipient fusion, cool, and weigh, repeating the
process until the weight is constant to 0.2 mg.

P. XVIII] HYDROCHLORIC ACID SOLUTIONS

147

Notes :

1. Quantitative filters are made from paper which has been treated with
hydrochloric and hydrofluoric acids until the nonvolatile materials (prin-
cipally compounds of iron, calcium, and silicon) are reduced to a very small
value. The average weight of the nonvolatile residue, or ash, in various
types of filters is shown in Table VII above, and is usually indicated on the
package.

It is also possible to obtain paper filters, both qualitative and quantitative,
of varying degrees of retentiveness and rate of filtration. If the precipitate
is very finely divided or colloidal, the more retentive, although less rapid,
filtering type must be used; with curdy or gelatinous precipitates the more
rapid filtering type may be used.

The size of the filter to be used should be determined largely by the size
of the precipitate (or residue) which is to be collected and not by the volume
of the solution to be filtered. The sizes most commonly used are 7, 9 and
11 cm in diameter. In general, the precipitate should not fill the filter over
one-third full in order that the washing can be carried out effectively and
safely. The funnel used should be of such size that the paper never comes
closer than 0.5 to 1.0 cm to the top; papers that are too large should be cut
to size after being folded. The paper should be folded so as to fit the funnel
closely (it is usually necessary to fold the paper on a slight bias and to use the
larger half), and, before beginning the filtration, it should be tested by filling
it with water. The stem should fill with water and hold this column of water,
which, by the suction it produces, greatly increases the rate of filtration.
Air leakage frequently occurs down the folded edge of the outside half of the
paper; this can be minimized by tearing off the upper corner of this fold.
A filter and funnel thus fitted are shown in Fig. 25. The beveled tip of the
funnel should always be made to
touch the inside of the receiving
vessel.

When the above precautions
are observed, the use of suction
with paper filters is usually un-
necessary, unless the precipitate
is large and of a bulky nature.
When suction is used, the lower
portion of the filter must be sup-
ported by a smaller filter of
hardened paper, a perforated
platinum cone, or a small circular
piece of cloth; the latter aid may
be purchased or cut from surgical
gauze or other similar cloth.
Only a moderate suction can be
used, and care must be taken in
applying it in order to avoid break-

Fig. 25. Paper Filter and Funnel.

ing the filter. Precipitates which are so colloidal in nature as to pass through
the filter can often be retained by adding paper pulp to the mixture before
beginning the filtration. This may be purchased in the form of readily

148 GRAVIMETRIC METHODS [P. XVIII

disintegrating tablets, or may be made by tearing a quantitative paper into
small bits and violently shaking it with water in a test tube. A silver
chloride precipitate which has been properly coagulated should be retained
by a medium or rapid filtering paper and should not require the use of either
paper pulp or suction filtration.

2. The funnel can be covered by a watch glass 'or by moistening the outer
edge of a filter paper and crimping it over the edge of the funnel.

3. Triangles for high-tern perature ignitions are preferably made of fused
silica, of clay tubes on heavy metal wire, or of platinum. Triangles of
nichrome alloy are satisfactory for medium temperatures but are not recom-
mended for high-temperature ignitions, especially with platinum crucibles.
Triangles should be kept scrupulously clean, since any material is likely to
fuse to the crucible.

4. For general purposes, care should be taken that the burner be adjusted
so that an oxidizing (nonluminous) flame is produced. A luminous yellow
flame causes carbon deposition on vessels and crucibles, does not produce as
high temperature, and is very injurious to platinum ware.

A properly adjusted Bunsen burner can produce a temperature of from
900 to 1050C. in a covered platinum crucible; a Tirrill burner, 1050 to
1150C.; a Meker burner, 1150 to 1250C.; and a blast burner, 1100 to
1300C. The temperature attained in a porcelain crucible will be from
200 to 300 less (Hillebrand and Lundell, Quantitative Inorganic Analysis,
p. 99).

In judging temperatures, it can be remembered that at approximately
500C. a perceptible red glow is produced; this becomes bright red at around
1000C. and an intense white glow at 1500C.

5. In removing a filter from a funnel, the paper may be raised by inserting
a clean spatula between it and the funnel and then folded forward so that
only the outside of the filter is handled. An alternative method is to tear
a small piece of paper from a quantitative filter and use this against the
inside of the filter when lifting it; this paper is then burned with the filter.
In case any precipitate has been spattered or has "crept" onto the funnel
above the filter, it can be first wiped up with this paper.

6. Most of the precipitate can be removed by inverting the filter and
gently rubbing the insides together; this should not be done so vigorously as
to cause particles of the paper to be rubbed off. This operation and the
subsequent burning of the filter should be carried out in a place which is abso-
lutely free of air currents.

7. The wire should be wrapped about that part of the filter where there is
no precipitate; otherwise this may fuse or be reduced during the ignition and
adhere to the platinum.

8. Since silver chloride fuses at 455C. and becomes appreciably volatile
at higher temperatures, the crucible should not be heated above the faintest
red glow.

Optional method D. Using a filtering crucible and drying
the precipitate with alcohol. The precipitation of the silver
chloride, the filtration, and the washing of the precipitate
are carried out as above in A. The filtering crucibles are

P. XVIII] HYDROCHLORIC ACID SOLUTIONS 149

cleaned, washed with water, and then treated as directed be-
low (Note 1).

Treat the clean crucible with 10 ml of boiling water and
by means of an efficient aspirator draw a rapid stream of
air through it for 2 to 3 minutes. Remove the crucible
from its support and dry the outside of it (and the part of
the filter support which is in contact with it) with a clean
lintless cloth. Replace it and before applying suction,
wash down the inside, and soak the mat with a 4-ml portion
of 95 per cent C2H 6 OH. Suck off this alcohol with the as-
pirator and then similarly treat the crucible with two 3-ml
portions of 95 per cent C 2 H 5 OH and with two 3-ml portions
of 99.5 per cent C 2 H 6 OH. Again draw air through the cru-
cible for 3 to 5 minutes. Remove the crucible, wipe the out-
side with the lintless cloth, place it in a vacuum desiccator
connected to an efficient water aspirator, evacuate the
desiccator, and maintain the vacuum for 5 minutes. Re-
move the crucible and weigh it. Repeat the evacuation
in the desiccator until the weight is constant to 0.2 mg.

Collect the precipitate on the crucible and wash it as
directed in the procedures above. After washing the pre-
cipitate, treat the crucible and precipitate as directed in
the preceding paragraph (Note 2).

Notes:

1. Experiments have shown that asbestos filters are difficult to dry in a
short time by this method; therefore their use is not recommended.

2. Silver chloride precipitates frequently coagulate, so that water may be
held in the precipitate and not be extracted by the alcohol. Therefore,
such precipitates should be broken up with a stirring rod as the alcohol is
applied; any precipitate adhering to the rod should be washed off with
alcohol.

PART II

THE SYSTEM OF ANALYSIS FOR THE
BASIC CONSTITUENTS

TABULAR OUTLINE I

PREPARATION OP THE SAMPLE, PRELIMINARY OBSERVATIONS, AND
PREPARATION OP THE SOLUTION

P. 1. Obtaining and Preparing the Sample.
P. 2. Treatment of Solutions and Suspensions.
P. 3. Preliminary Observations and Tests.

I. Observation of Physical Characteristics.
II. Detection of Organic Substances and of Water.
III. Solubility Tests (Detection of Oxidizing Agents).
P. 4. Elimination of Organic Substances.

Preparation of the Solution.

P. 5. Weigh a 1 -g sample (0.5 g of a metallic substance).

Treat with water, then 6 n. HNO*. (If completely dissolved, treat

by P. 11.)
// not dissolved:

Treat with 16 n. HNO t . Evaporate. (Heat with HCHO 2 if oxidizing
agents are present.)

Solution:
// sample dis-
solves, treat by
P. 11.

Solution :
Treat by
P. 11.

Solution:
Treat by
P. 11.

Residue: AgCl, AgBr, Agl, BaSO 4 , SrSO 4 , PbSO 4 , HgS,
A1 2 O 3 , Cr 2 Os, Fe 2 O 3 , Cr 2 (SO 4 ) 3 , SnS 2 , Sn 3 (PO 4 ) 2 , SnO 2 ,
Sb 2 O s , SiO 2 , silicates, silicides, SiC, C, Si; also most of
many cyanides. Ag 2 (CN) 2 , Ni(CN) 2 , complex cyanides,
Fe 4 (Fe(CN) fl ) 3 , and of CaSO 4 , CaF 2 , BiPO 4 , BiAsO 4 , and
so forth (see Note 12, P. 5).

P. 6. Treat with 6 n. HCL
If not dissolved:
Treat with IB n. HCl and 16 n. HNO t .

Residue: AgCl, AgBr, Agl, BaSO 4 , SrSO 4 , PbSO 4 , A1 2 O 3 ,
Cr 2 O,, Cr 2 (SO 4 ) 3 , SnS 2 , SnO 2 , SiO 2 , silicates, silicides, SiC,
C, Si, Ag 2 (CN) s ; also most of any CaSO, CaF 2 (see Note 3,
P. 6).

P. 7. Fume with HCIO* (or // 2 S0 4 ), adding UNO*.
Cool, add HF (bubbles; presence of silica).
Add more HF, fume (or evaporate to dryness).

Residue: BaSO 4 , SrSO 4 , PbSO 4 , AljOs, SnOi, a few silicates,
C (graphite), SiC (carborundum) (see Note 9, P. 7).
P, 8. Fuse with Na^CO*. Boil with water.

Solution :
Acidify with HCl
Treat by P. 6.

Residue: Carbonates and hydroxides.
Treat with HCl.
Treat by P. 6.

152

The Preparation of the Sample, Preliminary

Observations, and Preparation of the

Solution for the Analysis

P. 1. Preparation of the Sample

Discussion. Methods of sampling. The importance which at-
taches to the method of securing the sample of the original material
is often not sufficiently realized by an inexperienced analyst. In
order to appreciate the difficulties involved in this process, especially
when the material is a nonhomogeneous solid, the following facts
should be considered: First, the weight of the sample used for an
analysis usually does not exceed 1 g; second, if the analysis is to be of
value, this small amount of material must have a composition which
is the same, within the accuracy of the analysis, as that of the entire
mass of the material being analyzed; third, the mass of the original
material it may be a carload or a shipload or a mineral deposit of
considerable size is usually very large relative to the size of the
sample; and fourth, this material may be quite heterogeneous, as,
for example, an ore deposit composed of widely differing minerals.
It must, therefore, be strongly emphasized that the analyst is not
justified in undertaking an analysis unless he has himself sampled
the material or has adequate information as to the method by which
the sample has been secured.

The methods used in securing samples will vary with the physical
nature of the material. A more detailed discussion of the procedure
used for securing samples of non-metallic heterogeneous solid sub-
stances will be given, as it is probable that material which can be
included under this classification will be analyzed most frequently.
Alloys and metallic substances cannot be assumed to be perfectly
homogeneous and, as they are often non-friable, have to be sampled
by special methods. Solutions, if well shaken, are homogeneous, and
sampling is therefore usually simplified. The analysis of gases
requires special apparatus and technique and will not be treated in
this book.

The principles involved in securing a sample of a non-homogeneous
solid are simple. The operations, which may be laborious and time-
consuming, consist in taking a representative portion of the original
material, reducing it to a smaller particle size, and then taking a
representative portion of this finer material, this process being

153

154 PREPARATION OF THE SOLUTION [P. 1

repeated until a sample of the proper size for laboratory use is
obtained. It is fundamental that each portion be representative
of the entire mass from which it is taken; in other words, its average
composition must be, within the limits of the errors of the proposed
analysis, the average composition of the original material. It
follows that the size of the first portion taken will be determined by
two factors, the nature of the original material and the accuracy of the
analysis; the more heterogeneous the material and the more accurate
the analysis the larger the first portion will have to be. It also
follows that the smaller the particle size to which each portion is
reduced the finer it is crushed or ground the smaller the fraction
of it which then can be taken. In general, each portion should be
of such size that the largest particle of that portion, no matter what
its composition, could be removed or added without changing the
average composition of the portion by more than the errors of the
subsequent analysis.

The directions given in the following procedures must of necessity
be of a very general nature. More detailed treatments of the
sampling of various types of materials will be found in the following
texts and reference works:

The sampling of non-metallic materials:

Fales, Inorganic Quantitative Analysis, Century, 1925.
Mellor, A Treatise on Quantitative Inorganic Analysis, Griffin
and Co., 1913.

The sampling of metallic materials:

Lord and DeMorest, Metallurgical Analysis, 5th Kd., McGraw-
Hill, 1924.

Lundell, Hoffman, and Bright, Chemical Analysis of Iron and Steel,
Wiley, 1932.

The treatment of samples in the laboratory, with especial reference to
minerals:

Hillebrand, Analysis of Silicate and Carbonate Rocks, Bulletin 700,
U. S. Geological Survey, 1924.

Hillebrand and Lundell, Applied Inorganic Analysis, Wiley, 1929.

Procedure 1: PREPARATION OF THE SAMPLE. In the
following procedure the method is varied according to
whether the material to be sampled is (I) a non-metallic
solid, (II) a metallic solid, or (III) a solution or suspension.

(I) The material is a non-metallic solid: Take uniformly
throughout the body of the material from 20 to 50 portions

P. 1] SAMPLING 155

of such size that from 7 V to -^ of the total mass is
withdrawn (Notes 1,2).

Reduce the material to a finer state of aggregation (Note
3), pile it into a cone-shaped mound, flatten out this mound
into a disk, divide the disk into quarters, and reject the two
opposite quarters (Note 4). Repeat this quartering process
until the proper amount of material (determined by the
particle size) remains (Note 2). Again crush, grind, or
powder this portion and quarter it, repeating this process
until a bulk sample of from 1 to 5 Ibs. is obtained.

Again reduce the material to a smaller particle size, spread
it uniformly over a clean piece of glazed paper, oilcloth, or
rubber sheeting (about 20 in. square), pull one corner over
until it touches the opposite corner, and then repeat this
with each successive corner until the material is thoroughly
mixed. Pull each corner over in such a manner that the
material rolls over and over and does not merely slide
together. The material may now be "coned" or mounded
and quartered by means of a spatula (Note 11). These
operations are repeated until a portion of 50 to 100 g is
obtained.

Grind this portion in a hardened steel or agate mortar
and again quarter; repeat this process until a portion of
5 to 10 g is reduced to a very fine powder (Note 5). Trans-
fer this laboratory sample to a clean, dry, ground-glass-stop-
pered bottle. Treat this material as directed in P. 3.

(II) The material is a metallic solid: Clean the surface
(Note 6), and, by means of drilling, chipping, or cutting,
take portions uniformly from the material (Note 7) . Crush,
grind, or cut these portions into a finely divided condition
(Note 8), spread the material out evenly on an oilcloth or
glazed paper, divide it into a suitable number of small
squares, and take sufficient material from each of these to
give a laboratory sample of 5 to 10 g. Treat this material
as directed in Section II of P. 5.

(III) The material is a solution or suspension: Take, in a
clean glass bottle, that volume of the solution or suspension
which is thought to contain from 5 to 10 g of solid material
(Notes 9, 10). Close the bottle with a ground-glass (or
if it is not attacked by the solution, a new cork) stopper.
Shake the bottle vigorously before withdrawing a portion
for an analysis. Treat this material as directed in P. 2.

156 PREPARATION OF THE SOLUTION [P. 1

Notes:

1. The method of taking these portions will vary with the substance.
With gross materials, such as cement clinker or ore bodies, shovelfuls of the
material may be withdrawn from uniformly spaced locations. An endeavor
should be made to secure coarse and fine particles in the same proportion as
they occur in the material. Sampling may be done advantageously when
material is being moved for instance, unloaded from a car or being carried
on a conveyor as portions can then be taken at definite intervals.

2. The ratio of sample withdrawn to the total mass of material is deter-
mined by the accuracy desired in the analysis and by the homogeneity and
physical state of the substance. Thus this ratio would be much larger in
sampling a material such as coal or uncrushed ore than it would be in sampling
a well-ground and mixed material such as cement. It is necessary to make
the portion of such size that the loss of the largest particle, whatever its
composition, will not change the average composition by more than the
accuracy desired in the analysis.

3. This can be done with a hammer or other suitable instrument, or, if
available, a power crusher or pulverizer can be used. It is of course neces-
sary that no instrument or containers be used in crushing, grinding, or
powdering the substance which will introduce an appreciable amount of
foreign material into the sample. If iron or steel apparatus (crushers,
grinders, mortars, and so forth) have been used on non-magnetic materials,
it is common practice to remove particles of Iron thus introduced into the
sample by spreading the material on a clean piece of paper in a very thin
layer and moving a strong magnet slowly over the entire sample. The
magnet must not touch the sample, or particles of iron will be brushed off,
but it must be held as closely as possible. The sample should be then rolled
and the process repeated.

The size to which the particles must be crushed is determined by the
weight of the portion; thus with a portion of 1000 Ibs. the particle size
should not exceed 1 in. in diameter, with 50 Ibs. the particles should not
exceed about 0.4 in., and with J Ib. th6 particles should pass a No. 10
sieve (the openings of this size sieve are squares with sides approximately
0.079 in.).

It is often desirable to reserve a portion of the original untreated material
for an examination of its physical characteristics by Part I of P. 3.

4. 'This process, known as "quartering," is very generally used in the
taking of samples.

5. An agate mortar and pestle is commonly used for the final grinding.
Care should be taken not to pound the material in such a mortar but to use a
rotary grinding motion of the pestle. Mortars made of hardened steel are
also satisfactory for this operation. Power grinders must be used with
caution, as they introduce extraneous material.

6. The surface of the material should be cleaned of all rust, dirt, sulfide,
or any other acquired coating. This surface layer may be removed with a
wire brush, an emery wheel, or a file. Grease and organic matter usually
may be removed by washing with ether.

7. Alloys, castings, and other fabricated materials are often not homo-
geneous, and therefore samples must be taken from all sections. This can

P. 2] SOLUTIONS OR SUSPENSIONS 157

be done by drilling through or cutting sections from the object at uniform
distances. A lathe often can be used to advantage in taking cuttings.
Molten metals can be most effectively sampled when they are being poured,
by withdrawing small portions at regular intervals. Sections can be taken
from sheet metals by means of shears.

8. Brittle metals may be crushed in a hardened steel mortar; softer
material may be milled, cut, or filed (with a clean file) into finer particles.
This is desirable in order to obtain a representative sample and because the
more finely divided the material the greater the surface and the more rapid
the rate at which it will be attacked by solvents.

9. The volume of solution taken for an analysis will be determined by
the amount of inorganic solid material present. This may amount to as
much as several gallons in the case of potable waters or only a few milliliters
with concentrated commercial preparations. In most cases the approximate
concentration of the solution can be estimated from preliminary information,
or by evaporating a small portion and noting the amount of residue.

10. Most solutions, unless they contain suspended matter, can be assumed
to be homogeneous. Samples of drinking water from wells or reservoirs can
be taken by submerging the bottle (after washing it with the water) and
allowing it* to fill under the surface; if taken from a stream, portions should
be taken at various places across the line of flow; if taken from a pipe line, the
water should be allowed to run for some time in order to clear the line of
accumulated material. Commercial preparations vary so greatly that no
general procedure can be used, and the analyst must suit the method to the
specific case, being sure to take representative samples in the case of suspen-
sions and emulsions, to avoid loss in handling volatile solutions, and to
guard against change or contamination after taking the sample.

11. This "rolling" process secures thorough mixing of the sample and
prevents the loss or segregation of particles of a given size. This is important
where it is desired to take a small sample from a large amount of material
by the "spotting" process, that is, by removing numerous small uniformly
spaced portions from the mass of the material.

P. 2. Treatment of Solutions or Suspensions

Discussion. If the material is a solution (with no suspended
material), it is necessary to determine the total concentration of
solid matter present so "that such a volume may later be taken for the
analysis as will give about 1 g of solid inorganic constituents. This
determination is made by evaporating a small volume of the solution
on a water bath and weighing the residue. The residue is also used
to determine if any non-volatile organic material is present, and
for carrying out the preliminary test of P. 3. Before the solution is
evaporated, it is neutralized in order to avoid possible loss of volatile
constituents such as certain acids, or AsCl 8 from a hydrochloric acid
solution.

158 PREPARATION OF THE SOLUTION [P. 3

If the material contains suspended matter in significant amount,
it may be desirable to filter out the suspended material and to
analyze it and the solution separately; by this means information
both as to the composition of the solution and of the suspended
material is obtained. Otherwise, depending upon the information
desired from the analysis, the suspended material may be filtered
out and rejected, or it may be disregarded and the suspension treated
as though it were a clear solution.

Procedure 2. TREATMENT OF SOLUTIONS OR SUSPENSIONS.
If the material is a clear aqueous solution (Note 1), pipet 10
ml of it into a weighed porcelain crucible, and test it with
litmus paper. If it is acid, make it just alkaline with
NH 4 OH, evaporate it to dryness on a steam bath, cool the
crucible in a desiccator, and weigh it. Treat the residue
(Note 2) by P. 3.

If the material is a suspension, either (a) treat the suspen-
sion by the paragraph above as though it were a clear
solution, or (b) filter out the suspended material and
treat the residue by P. 3 and the solution by the paragraph
above.

Notes :

1. Inorganic material may sometimes be dissolved in a volatile organic
solvent such as ether, alcohol, or benzene. Such solutions may be treated
in a similar manner in order to determine the total solids present. Such a
solvent may often be identified by distilling it off and noting its boiling
point, odor, and other physical characteristics.

2. If an insufficient amount of solid material is obtained, a larger volume
of the solution will have to be evaporated.

The residue is treated by P. 3 in order to detect organic material and to
note if any residue remains after the heating, since the solution may have
contained only an acid or an ammonium salt.

P. 3. Preliminary Observations and Tests

Discussion. Valuable information as to the composition and
physical characteristics of an unknown substance often can be
obtained by means of an examination made with the aid of a magni-
fying lens or, preferably, a low-power microscope. It is usually
possible to determine whether the material is a homogeneous sub-
stance or a mixture and whether it is crystalline or amorphous or
composed of particles of each type, and the color and crystalline

P. 3] PRELIMINARY OBSERVATIONS 159

nature of individual crystals can be noted, so that often an identifica-
tion can be made. 1

Organic compounds may often interfere with the course of an
analysis by separating in a voluminous form, thus making filtrations
and such operations difficult; by causing the precipitation of other-
wise soluble cations; or by forming un-ionized or complex compounds
with certain cations and preventing their precipitation. Thus
oxalate may cause the precipitation of certain alkaline earth elements
and at the same time prevent the precipitation of aluminum or
chromium with the Ammonium Sulfide Group. For these reasons
organic material must be tested for before beginning an analysis
and, if present, must be removed or destroyed. Organic matter is
tested for here by heating a small portion of the solid material in a
test tube made of resistance glass, or, if a more certain indication is
desired, by heating with concentrated sulfuric acid and noting if
charring occurs. Indications of other constituents may be also
obtained in both of these processes.

When the material is heated in the tube, any water which may be
present is driven off and may be detected by causing it to condense
in the upper, and cooler, portions of the tube. By noting the
amount of condensate, a rough guess as to the amount of water
present can be made; if an approximate estimation is desired, a
supplementary procedure for determining the "loss on ignition' ' is
provided in Note 5. This information is often of value, as it enables
one to sum up all the constituents present in the material. Further
information can be obtained by determining the water lost by drying
at 110C.

It is often advantageous to obtain some information as to the
solubility of the material in different solvents before beginning the
systematic preparation of the solution. Therefore, preliminary
tests on small samples are made in order to ascertain the solvent
effect of water, of nitric acid, and of hydrochloric acid. Information
as to the presence of silver and of large amounts of sulfides may also
be obtained; this information is useful in reaching a decision as to
whether to treat the sample first with nitric or hydrochloric acid.
The relative advantages of these acids for this purpose are mentioned
in the discussion of P. 5, last paragraph.

1 A microchemical technique has been developed whereby, even when the
amount of the material is very small, quite complete analyses can be per-
formed. References: Chamot and Mason, Handbook of Chemical Microscopy,
Wiley, 1930 % - Emich, Microchemical Laboratory Manual, Wiley, 1932.

160 PREPARATION OF THE SOLUTION [P. 3

Because the presence of considerable quantities of oxidizing agents
would cause the formation of sulfur and sulfate in the hydrogen
sulfide precipitation, it is desirable that these be reduced prior to
that treatment. Accordingly, an incidental test for the presence of
oxidizing agents is made, and, if they are found present, formic acid
is used to insure their reduction during the preparation of the solu-
tion. This test depends upon the oxidation of chloride in a hydro-
chloric acid solution to chlorine, which can be detected by its color,
smell, or the use of starch-iodide paper. Unfortunately, the precipi-
tation of sulfur in the hydrogen sulfide precipitation is not entirely
eliminated. Certain less powerful oxidizing agents) ferric salts, for
example, do not give the chlorine test and may not be completely
reduced by the formic acid, and yet oxidize hydrogen sulfide to sulfur.

Procedure 3:1. PRELIMINARY OBSERVATIONS. Note closely
the physical characteristics of the material, such as color,
approximate density, odor, and feeling when rubbed be-
tween the fingers. Examine a small portion of it with the
aid of a magnifying lens or low-power microscope. Observe
whether it is homogeneous, and whether it is crystalline
or amorphous or a mixture of both types of substances. If
a mixture, try to identify the individual constituents by
their color and form (Note 1).

II. SOLUBILITY TESTS (DETECTION OF SILVER AND OF
CERTAIN OXIDIZING AGENTS). Transfer approximately 0.1 g
of the finely divided material to a test tube (25 ml or larger),
add to it 10 ml of water, shake the mixture vigorously, and
note the result. Heat the mixture to boiling and note any
changes. Test the solution with litmus (Note 6) . If a residue
remains, add 5 ml of HNOa and note if any reaction occurs
(Note 7) ; again heat the solution, keeping it almost boiling
as long as the material appears to be dissolving (Note 8).
If the material is completely dissolved, treat the solution by
the second sentence of the next paragraph.

If complete solution was not obtained, treat another 0.1-g
portion with 5 ml of 16 n. HNOa, warming the mixture as
long as a reaction appears to be taking place (Notes 9, 10).
Add 10 ml of water (Note 11) and just 0.1 ml of 3 n. NH 4 C1,
and then heat the mixture almost to boiling. (White pre-
cipitate, presence of silver. Note 12.)

Treat another 0.1-g portion with 5 ml of 12 n. HC1; gently

P. 3] PRELIMINARY OBSERVATIONS 161

warm the mixture as long as the material appears to be dis-
solving. (Chlorine gas in the tube, presence of oxidizing
constituents. Notes 13, 14.) If complete solution is not
obtained, add to the mixture 1 ml of 16 n. HN0 3 and again
warm it (Note 15).

III. DETECTION OF ORGANIC SUBSTANCES AND WATER.
Transfer 0.1 to 0.2 g of the material to the closed end of a test
tube, 10 to 15 cm long and about 1 cm in diameter, made of
resistance glass (Pyrex or Jena) . Clartip the tube in a nearly
horizontal position, wrap a moist strip of paper or cloth
near the upper, open end, and heat the lower, closed end of
the tube, very gently at first and then finally with the full
heat of the burner (Note 2). Note any changes taking
place in the material, any deposit forming on the inside of
the tube, or any gas issuing from the tube. (Charring of
the material, accompanied by a tarry deposit, or smoking,
accompanied by a "burnt" odor, presence of organic sub-
stances. Notes 3, 4. Aqueous deposit, presence of water,
Note 5.)

If it is desired to estimate the water present, treat a por-
tion of the material as directed in Note 5 below.

If organic substances have been found present, proceed
as directed in P. 4.

If the material does not contain organic substances, pro-
ceed as directed in P. 5.

Notes:

1. Although in the large majority of cases the Analyst can, and should,
obtain some previous information as to the nature of the material, occasion-
ally he has to deal with a completely "unknown" sample. In such cases the
preliminary examination may furnish information which will materially assist
in the subsequent analysis. Often by the color of individual crystals or by
the smell of the substance, especially if it is warmed, predictions can be made
as to the constituents present. It may be advantageous to make these
observations on representative specimens of the original material which have
not been crushed or ground. It is to be emphasized, however, that these
are only predictions, and cannot be relied upon until fully substantiated by
the analysis.

2. When an unknown substance is heated, the tube should be kept in
such a position that personal injury will not result should an explosion occur.
Caution should be taken in smelling any gases given off.

3. A black residue resulting from heating the material does not con-
clusively prove the presence of organic substances, as certain salts, such as
the nitrates and carbonates of copper, cobalt, and nickel, may be decom-

162 PREPARATION OF THE SOLUTION [P. 3

posed to give black oxides. The characteristic charring and the formation
of a tarry deposit which accompanies the heating of organic material are
much more conclusive.

From the amount of material taken and the residue remaining, an estimate
of the amount of organic matter present can be made. If desired, this
can be more accurately determined by transferring a larger portion of the
material to a weighed porcelain crucible, again weighing the crucible, then
gently heating it over a burner until the organic matter is completely charred
without flaming, and finally heating with the full heat of the burner until all
the carbon is burned off. Upon allowing the crucible to cool and weighing
it, the weight of the inorganic constituents is obtained. It is to be realized,
however, that water and certain other inorganic constituents, such as CC>2
and mercury compounds, may be also driven off by this treatment.

A more distinctive and sensitive test for organic substances can be made
by carefully adding 0.1 to 0.2 g of the substance to 5 ml of 36 n. H2S04 and
then heating the mixture. Under these conditions, most organic substances,
even in small quantities, darken the solution; thus 1 mg of tobacco or starch
and 2 mg of tartaric acid were easily detected by this test.

Considerable information regarding the acidic constituents present may
also be obtained by heating the substance with sulfuric acid, especially if
the material is treated with 2 to 4 ml of 6 n. H2S04, warmed, any reaction
noted, 5 ml of 36 n. I^SCU slowly added, and the mixture heated. Upon
being warmed with the dilute acid, many sulfides will evolve H2S; carbonates
will give C02j cyanides will give HCN, detected by its odor; sulfites will
give S02; thiosulfates will give S02 and free sulfur; acetates will smell of
acetic acid; nitrites will give a brown gas, N02j and chlorine, bromine, or
iodine will result from the presence of these halogens together with oxidizing
agents. Upon heating the concentrated acid, S02 will be given off if reduc-
ing agents such as organic matter, iodides, metals, or sulfur are present;
volatile acids such as HC1 and HNOs will be expelled. Oxalates will not
be decomposed by the dilute acid, but will give CO and C02 upon being
heated in the concentrated acid.

4. Although this so-called "closed-tube" test is used primarily to detect
organic substances and water, much supplementary information may also
be obtained. Thus, if the material is completely volatilized, non-volatile
substances cannot be present, and the material is composed wholly of such
substances as ammonium salts, mercury compounds, or organic substances.
If a deposit other than water or the characteristic tarry, gummy material
caused by organic substances is obtained on the sides of the test tube, it
should be carefully examined. A white, more or less crystalline deposit
indicates the possible presence of the following: ammonium salts; the chlo-
rides of mercury; when anhydrous, the chlorides of tin, arsenic, aluminum,
iron, and antimony; arsenic trioxide; phosphorus oxides; and many organic
compounds which may sublime without charring. A gray deposit may be
caused by mercury compounds, and examination of this deposit with a lens
may show the presence of droplets of free mercury; metallic arsenic gives a
black deposit, which in a mixture will appear gray. A yellow deposit indi-
cates arsenic sulfide, sulfur due to sulfide, or elementary sulfur.

P. 3] PRELIMINARY OBSERVATIONS 163

Gases that may also be evolved may aid in identifying the presence of
certain acidic constituents. Carbon dioxide, from a carbonate, oxalate, or
other organic substance, is indicated if a turbidity is produced upon sus-
pending a drop of Ba(OH)2 solution in a small loop of platinum wire and
inserting this in the tube. Sulfur dioxide, detected by its smell, indicates
sulfites, thiosulfates, or sulfides. Oxygen, detected by inserting a glowing
splinter in the tube and noting if it flames, indicates a nitrate, chlorate,
peroxide, or higher oxide such as Pb02. Hydrogen sulfide may sometimes
be detected when a moist sulfide is heated. A brown gas indicates N02
from a nitrate or nitrite. A violet vapor indicates iodine, which can also be
recognized by a mirror-like or crystalline sublimate and by its smell. Bro-
mine or chlorine may be expelled from certain salts.

5. Adsorbed water, or the so-called hygroscopic moisture, will almost
always be found present in small amounts in materials which are ex-
posed to the air; this water will usually be less than 1 per cent of the total
and can be fairly completely expelled by drying the material at 110C. for
an hour. Water may also be held mechanically enclosed within crystals,
causing them to decrepitate when heated; in certain cases water is held by
capillary forces in minerals having, as do certain of the zeolites, a porous
structure.

Water present owing to the causes just discussed is sometimes classified
as non-essential water to distinguish it from water which is assumed to be
present in some systematic molecular or crystal structure and which bears
some definite stoichiometric relation to the other constituents of the mole-
cule in which it is present. This essential water may be also divided into
two classes: water of crystallization and water of constitution. Water of
crystallization can usually be expelled below 200C. Water of constitution
is usually expelled only at higher temperatures and is accompanied by a
distinct change in the chemical nature of the molecule; acidic substances
such as Ba(H2P04)2 or KH^A^SiO-Oa and basic compounds such as Mg(OH)2
or Fe(OH) 2 C2H 3 02 contain water of constitution.

A determination of the "loss on ignition" and an approximate estimation
of the hygroscopic and total water present can be made as follows:

Heat a porcelain crucible over a Meker-type burner, cool, and
weigh. Weigh 0.5 to 1 g of the material into the crucible and heat
in an oven at 110C. for 1 hr. Cool and weigh. Again heat the
crucible over the burner, raising the temperature slowly, for 20 to
40 min. Cool and weigh. Repeat the heating and weighing until
constant weight is obtained.

The hygroscopic water will be obtained from the loss in weight on heating
at 110C.; however, hydrated salts may lose much of their water at this
temperature. As the material may contain other volatile constituents
(such as organic matter, carbonates, and mercury compounds) or constitu-
ents (such as sulfides) which may be changed by being heated in the air to
the temperature given by a Meker burner, the results of this procedure should
be interpreted only after the complete analysis has been made. The above
procedure is sometimes used for making an approximate determination of
the carbonate in limestones.

164 PREPARATION OF THE SOLUTION [P. 3

6. If the material appears to have partly dissolved, another 10-ml por-
tion of water should be added and the effect noted: a moderately soluble
substance may be thus completely dissolved. An examination of the
water-insoluble residue will sometimes show that one of the constituents of
the original material has been extracted or all but one insoluble constituent
have been dissolved. The test with litmus will show whether the solution
is alkaline indicating soluble hydroxides (NaOH, Ca(OH)2, and so forth)
or salts formed from slightly ionized (weak) acids and highly ionized (strong)
bases (NajCOa, NaaPC^, and so forth) or acidic indicating acids (PgOe,
HaCaC^), acid salts (NaHSC>4), or salts formed from weak bases and strong
acids (FeCl 3 , Bi(N0 3 )s, and so forth).

Indicator test papers can now be purchased 2 which give more quantitative
information than does litmus in that they show a continuous change in color
as the hydrogen ion concentration of the solution is varied over a wide
range (from pH 2 to pH 10), and by their use the pH of a solution may be
determined to within 1 to 2 pH units. If such papers are available, they
can be used to advantage for determining the approximate pH of this solu-
tion. These papers should be dipped in the solution being tested and im-
mediately withdrawn, as their coloring material readily passes into the
solution, and future color indications of the constituents of the solution may
be thereby obscured.

7. Carbonates or sulfides will be indicated by effervescence, the latter
being detected by the smell of HaS; sulfites will give S02, which can also be
detected by smell; nitrites will be decomposed, and brown fumes of NOa
will form in the test tube; and cyanides will produce the characteristic odor
of HCN (DANGER).

8. A residue that is insoluble in water but soluble in dilute HNOa indi-
cates the presence of water-insoluble salts of weak acids, the more common
ones being sulfides, carbonates, sulfites, phosphates, arsenates, borates, and
chromates (see Table Vllf and Note 12, P. 5). A precipitate forming on
adding the acid suggests the presence of a soluble silicate, or of antimony or
tin dissolved in an alkaline solution.

An insoluble residue indicates the presence of insoluble salts of strong
acids, of certain very insoluble salts of weak acids, such as mercuric sulfide,
or of native or high-temperature products. Substances of the latter classes,
even if soluble when newly precipitated, are often converted by heat or
time into very resistant forms which dissolve with extreme slowness. Ex-
amples of these substances are the silicates, aluminum and iron oxides, and
oxides and sulfides of stannic tin.

9. Upon treating the material with concentrated HNOa, the appearance
of brown fumes will indicate the presence of reducing material. Thus, most
metallic substances present will be dissolved and most metallic ions in a
lower oxidation state will be oxidized, mercurous, arsenious, antimonous,
stannous, cuprous, and ferrous salta all being converted into a higher oxida-
tion state. Many acidic ions will also be changed. Thus sulfide will form

1 "Alkacid," Fisher Scientific Co., 711-723 Forbes St., Pittsburgh, and
"Hydrion," R. P. Cargille, 118 Liberty St., New York. "Nitrazine" is use-
ful for a more restricted range (see Note 1, P. 161).

P. 3] PRELIMINARY OBSERVATIONS 165

sulfur and sulfate, sulfite will give sulfate, nitrite will give N(>2, and chloride,
bromide, and iodide will all be oxidized to the free halogen.

10. If the material has been insoluble in dilute HNOa but has dissolved
in the concentrated acid, the presence of difficultly soluble sulfides is strongly
suggested; thus silver sulfide would be rapidly dissolved by this treatment.
A white residue, apparently precipitated by the HNOa, indicates the pres-
ence of silicates or of antimony or tin. The nitrates of lead, barium, and
strontium are only moderately soluble in concentrated HNOsj however,
these precipitates are crystalline and dissolve when the solution is diluted.
If a large amount of free sulfur is formed, it is desirable that the material be
first treated with HC1 when a solution is being prepared for the analysis.
This treatment will in most cases cause the sulfide to be expelled as HaS,
and avoid the formation of the large amounts of sulfate which result from
treating a sulfide with concentrated HNOa.

11. If there is a residue which remains suspended in the solution and
which would prevent the analyst from noting the formation of a precipitate,
the mixture should be filtered before adding the NHiCl.

12. The formation of a precipitate upon adding only 0.1 ml of NEUC1
indicates the presence of silver or of a large quantity of lead. If the precipi-
tate remains upon heating, the solution, the presence of silver is proved; the
small amount of lead chloride which would be formed by only 0.1 ml of the
NHiCl would dissolve. Mercurous mercury, if originally present, would
have been oxidized by the treatment with hot HNOs. If silver is present,
the solution resulting from treating the original material with HNOs (P. 5)
should be filtered and treated separately by P. 11. This avoids precipitating
all of the silver as chloride upon adding HC1 in P. 6 and the difficult treat-
ment necessary to redissolve this precipitate. Only 0.1 ml of NH4C1 is
added, as this will cause 0.1 mg of silver to givs a detectable precipitate in
the hot solution even in the presence of 500 mg of lead.

13. The more powerful oxidizing agents, such as permanganates, chlo-
rates, bromates, iodates, of chromates, which are also usually soluble, will
give chlorine without the solution being warmed. Nitrates and insoluble
oxidizing compounds, such as lead dioxide or lead chromate, may require
warming. Nitrites would cause a brown gas, NOa, in the tube, but this is
darker colored and will be more quickly expelled than the chlorine. The
presence of chlorine in the tube may be shown by its yellow color and charac-
teristic odor. Chlorine in small amounts may be detected by inserting in
the tube a strip of filter paper which has been moistened with a solution con-
taining potassium iodide and starch; the immediate formation of a brown or
dark color shows the presence of chlorine. However, any other oxidizing
gas in the tube will cause the same reaction, and the smell of C1 2 is so sensi-
tive a test that it is usually sufficient.

14. The dissolving in the HC1, without evolution of CU, of material in-
soluble in nitric acid indicates the presence of slowly dissolving oxides,
notably those of iron, chromium, and aluminum. Concentrated HC1 is a
more effective solvent than HNOa for such substances, owing to the forma-
tion of stable complex ions.

15. Certain substances, such as HgS, will dissolve in the mixed acids, ow-
ing to the combination of a powerful oxidizing action and the above-men-

166 PREPARATION OF THE SOLUTION [P. 4

tioned tendency toward complex formation. Thus the sulfide ion will be
removed by oxidation to sulfur or sulfate and the mercuric ion will combine
with the chloride ion to form the very slightly ionized complex ion HgCl 4 "
(see the discussion of P. 6).

P. 4. The Elimination of Organic Material

Discussion. Organic material can be destroyed in a "dry way"
by heating to a high temperature in air and thus volatilizing or
burning it off. Such a procedure is not used except in special cases,
as it involves the possible loss by volatilization of certain of the
inorganic constituents. Thus all mercury compounds would be lost;
the chlorides of many of the other metals are volatile at higher tem-
peratures; and upon being heated with carbon many of the basic
constituents are reduced to the metal, which may be volatile.
Furthermore, upon being heated to a high temperature, many
metallic compounds are decomposed into oxides which are very
slowly dissolved upon treatment with acid; thus the ignited oxides
of aluminum, chromium, iron, and tin are not appreciably at-
tacked by even prolonged treatment with concentrated acids.
Organic material is more commonly destroyed in a "wet way"
by heating with concentrated sulfuric and nitric acids. Although
an effective treatment, this introduces sulfate into the analy-
sis and causes the precipitation of lead, barium, strontium, and
calcium as the slightly soluble sulfates; also, after fuming with sul-
furic acid, chromium is converted into the compound CrJH^SOOr,
which dissolves with extreme slowness. Because of these facts,
perchloric acid is used in place of sulfuric in this procedure.

The Properties of Fuming Perchloric Acid. Heretofore, perchloric
acid has not been used extensively in analytical processes for two
reasons. First, the cost has been prohibitive, and, second, there
apparently has been a general apprehension as to the possibility of
explosions. The cost of the acid is now more reasonable, and ex-
tensive experience has shown that, if properly handled, perchloric
acid can be used relatively safely. Aqueous perchloric acid is an
extremely stable reagent when dilute, reduction being impossible
except by a few very powerful reducing agents such as chromous
and titanous salts, and even these react slowly. When the acid
is heated to fuming, it becomes a powerful oxidizing agent, con-
verting chromic and manganous ions to chromate and manganese
dioxide, respectively. If the acid is heated to fuming with organic
substances, especially alcohols, or with certain reducing agents,

P. 4] INTERFERING ACIDIC CONSTITUENTS 167

such as antimonous salts, violent and dangerous explosions will
occur. However, it has been shown by Noyes and Bray 8 that, if
nitric acid is added to such mixtures, they can be fumed, or even
evaporated to dryness, and in no case was it possible to cause the
mixture to explode. 4 Furthermore, it has been shown by Noyes
and Bray that the mixture of nitric and perchloric acid oxidizes
organic substances more rapidly and with less charring than
does sulfuric acid. The only insoluble perchlorate which may pre-
cipitate is the potassium salt; this, however, is readily soluble in
hot water.

Many types of organic material, such as fatty acids, oils, greases,
and so forth, can be eliminated by extracting them with ether (in
which most inorganic substances are insoluble) and then collecting
the residue on a filter paper or separating the ether by means of a
separating funnel. Information as to the nature of the material
being analyzed, or a trial with a small portion, will indicate the most
effective treatment.

In addition to organic material there are certain acidic constituents
which may interfere with the course of the systematic basic analysis.
The effect of phosphate in precipitating the alkaline earth elements
with the Ammonium Sulfide Group is discussed in P. 54, and a
method for eliminating phosphate is provided there (arsenate would
behave similarly but is removed by the hydrogen sulfide precipita-
tion). Although oxalate is an organic anion, it may decompose
without charring sufficiently to give a positive test for organic
material, and it would tend to prevent the precipitation of aluminum
and chromium hydroxides (because of the formation of complex
ions) with the Ammonium Sulfide Group and to cause the precipita-
tion of alkaline earth oxalates with that group. Fluoride may cause
the partial or complete precipitation of alkaline earth fluorides with
the Ammonium Sulfide Group, although these fluorides are usually
so resistant to solvents that in preparing the solution the sample
would have to be treated with fuming perchloric acid (P. 7"!, which
would volatilize the hydrofluoric acid. Ferrocyanide and ferri-
cyanide are troublesome, as they form slightly soluble precipitates
with many of the cations and as they are partially decomposed in
hot acid solutions with the formation of blue precipitates. Because

3 Noyes and Bray, Qualitative Analysis for the Rare Elements, Macmillan,
1927.

4 Recent reports of explosions with nitric and perchloric acid mixtures,
/. Ind. Eng. Chem., News Ed., 15, 214, 332 (1937), indicate that care should
be taken to carry out the operations in the procedure below so that, should an
explosion occur, personal injury will not result.

168 PREPARATION OF THE SOLUTION [P. 4

of these effects it is desirable that the analysis for the acidic constit-
uents be made before carrying out the basic system of analysis,
If these interfering anions are found present, they can be removed
or decomposed by the treatment with nitric and perchloric acid of
the procedure below. In addition, certain anions (such as nitrite,
sulfite, thiosulfate, and certain of the oxy-halogens) which would
cause the formation of sulfur in the hydrogen sulfide precipitation
(P. 11) are decomposed by fuming perchloric acid.

Procedure 4: THE ELIMINATION OF ORGANIC MATERIAL.
If organic matter has been found present (P. 3), and if the
original material is a non-metallic solid, weigh and transfer
to a casserole that quantity of the material which, from the
information obtained in P. 3, is estimated to contain about
1 g of inorganic residue (Note 3, P. 5). Treat as directed in
the second paragraph below.

If organic matter has been found present and if the origi-
nal material is a solution or suspension, measure out into a
casserole that volume of the solution which is estimated to
contain 1 g of solid inorganic material (Note 3, P. 5). If the
solution is acid, make it neutral with NH4OH, and evaporate
it just to dryness.

Add slowly 10 to 20 ml of 16 n. HN0 3 and cautiously
warm the mixture; cool it (by immersing the bottom of the
casserole in cold water) if the reaction becomes too vigorous
(Note 1). Heat the mixture until no more action seems to
be taking place, add slowly 5 ml of 9 n. HC10 4 , place the
casserole on a steam bath, and heat it, adding more HNOs
as this evaporates, as long as brown fumes are given off.
Finally, evaporate the mixture until dense white fumes are
evolved (Caution: Note 2). If the solution is dark colored,
due to charred material, add 2 ml more HNOa and again
fume, repeating this process until the solution is clear
(Note 3).

If there is no residue (Note 4), dilute the solution to 50
ml and treat it by P. 11 (Note 5).

If there is an insoluble residue, treat the mixture by P. 6.

Notes:

1. When organic substances are treated with concentrated nitric acid,
the reaction, although it starts slowly, may become so rapid as to cause loss
by fepattering. For this reason the acid should be added first in small por-

P. 4] SOLUTION OF SOLID SUBSTANCES 169

tions and the mixture heated slowly, so that, if the reaction becomes too
vigorous, it can be controlled by cooling the casserole.

2. The possibility of obtaining an explosion when the procedure is carried
out as directed is extremely small; however, since perchloric acid explosions
are so violent as to be quite dangerous, it is advised that the casserole be
placed on a ring stand and allowed to evaporate behind a screen in a hood or
other place such that, should an explosion occur, no one will be injured.

If perchloric acid is not available, or if it is desired to avoid its use, an
equal volume of 18 n. sulfuric acid may be substituted. The disadvantages
of this acid are pointed out in the discussion above.

3. If much manganese is present, a dark precipitate of Mn02 may form
on vigorously fuming the acid. This is readily recognized and can be dis-
solved by adding a few drops of formic acid and again heating the mixture.

4. Salts, especially KC104, may crystallize in the fuming acid. There-
fore, if a crystalline residue is obtained, it is admissible to dilute the mixture
before deciding whether or not there is an insoluble residue.

5. If the mixture has been fumed for a considerable period, so much of the
acid may have been lost that proper conditions for the HzS precipitation in
P. 11 may not be obtained. In this case, the solution should be made just
neutral to litmus with NFUOH, and 5 ml of 6 n. HNOa should then be added.

THE PREPARATION OF THE SOLUTION FOR THE ANALYSIS
General Discussion of the Solution of Solid Substances

Two general methods are available by which solid substances
may be brought into a soluble form: first, by treatment in the "wet
way" with solutions of various solvents such as nitric or hydro-
chloric acid alone, mixed acids such as nitric and hydrochloric or
sulfuric and hydrofluoric together, or an alkaline solvent like sodium
hydroxide; and second, by treatment in the "dry way," which usually
consists in fusing the material with some flux such as sodium car-
bonate or potassium hydrosulfate. Each of these methods has its
merits. The treatment with acids is favored in qualitative work
because no basic ions are introduced into the analysis, and also
because fusion processes may cause the volatilization of certain
constituents, may require more specialized technique and more
elaborate apparatus, and are usually more time-consuming. How-
ever, there are a large number of insoluble salts of strong acids and
many native and high-temperature products which are completely
decomposed only by a fusion process. When adequate information
as to the qualitative composition of the material is available, it is
usually possible to select at once the proper method of attack; how-
ever, for qualitative work it seems desirable to use a systematic

TABLE VIII

THE PREPARATION OP THE SOLUTION
(Solvents, Types of Solvent Action, and Substances Dissolved)

Proce-
dure

Solvents Used

Type of Solvent Action

Type of Substances
Dissolved (Examples)

P.5

(a) H,0

Solution (ionization,

Water-soluble com-

solvation)

pounds

(b) HNO, (dilute)

I. Hydrogen ion ef-

fects

1. Hydroxides, basic

1. Neutralization

oxides, basic

2. Displacement

salts

3. Oxidation (2H +

2. Salts of weak acids

+ 2E- - Hi)

3. Certain metals

(Zn, Al)

II. Oxidation (NOr

II. Certain vigorous re-

+ 4H+ + 3E- -

ducing agents (fer-

NO + 2H 2 O)

rous and stannous

salts)

I, II. Oxidation

Reducing compounds

hot)

(NO,- + 2H+ +

(sulfides, alloys, .met-

E- = NO 2 + H 2 O)

als, etc.)

P.6

(a) HC1 (cone.)

I, III. Reduction

Oxidizing compounds

(2C1- - 01, + 2E-)

higher oxides (MnO2),

oxidizing salts

(PbCrO 4 )

IV. Complex ion for-

Compounds of cations

mation

forming complex ions
(HgClr, SnClr)

(b) HC1 (excess)

I, II, III, IV

Those above; also noble

and HNO,

metals (Pt, Au) re-

(cone.)

quiring both oxida-

tion and complex for-

mation

P.7

(a) HC1O 4 , fuming

I, 11 (only when

Salts of volatile acids

(or H 2 SO 4 )

fuming)

(sulfides, halides,

V. Displacement by

fluorides, etc.)

volatilization (of

lower boiling acids)

(b) HClO 4 andHF

Primarily silicates

(excess)

(formation of H 2 SiF 6

andSiF 4 )

P.8

(a) Na 2 CO 3 (solu-

VI. Metathesis (car-

Compounds of cations

tion)

bonate and hydrox-

forming insoluble car-

ide)

bonates and hydrox-

VII. Hydroxyl ion

ides

effect

1. Acidic oxides

1. Neutralization

2. Salts of weak

2. Displacement

bases

3. Oxidation

3. Certain reactive

metals (Al, Zn)

(b) Na 2 CO 3 (fu-

VI, VII

As above, more rapid

sion)

VIII. High tempera-

and extensive

ture

VI, VII, VIII, and

As above; also com-

NaNO 3 (fu-

pounds of elements

sion)

forming acids in their

higher oxidation

states (Cr, Mn)

170

P. 5] TREATMENT WITH NITRIC ACID 171

treatment of the material, first with single acids, then with mixtures,
and finally, only when it is absolutely necessary, with a fusion process.
The successive reagents used in bringing materials into a soluble
form are outlined in Table VIII. There is also indicated the type of
solvent action exerted by each reagent or combination of reagents
and, furthermore, the type of compounds usually brought into solu-
tion by these reagents. A more detailed discussion of the reactions
of the solvents and the specific compounds dissolved by each will
be found in the subsequent procedures. 6

P. 5. Treatment of the Sample with Nitric Acid

Discussion. The first step in the systematic treatment to pre-
pare a solution for the analysis of a material is to treat it with water
alone. A compound may dissolve in a pure solvent, such as alcohol
or water, as undissociated molecules, as does sugar or glycerin, or
it may pass into the solvent partly or completely as ions, as does
acetic acid or sodium chloride. In general, the greater the ionizing
tendency of the liquid (measured largely by its dielectric constant)
the more effective it is as a solvent for most inorganic compounds.
The solution process may also be accompanied by solvation (in the
case of water as the solvent, hydration) ; this involves the formation
of some type of a compound between the molecules or ions of the
solute and the molecules of the solvent. Thus it is known that
in many cases the positive metallic ions form stable compounds
with water molecules, an example being the hydrated copper ion,
Cu(H 2 O)4 ++ . It is now generally believed that hydrogen ions exist
largely in solution as H*(H 2 O), usually written HsO 4 " and called the
"hydronium ion." lonization and solvation are probably the most
important factors determining the solvent properties of a pure liquid.

The amount of a compound which will pass into a given solvent
is increased by the addition to the solvent of any substance which
will react with the molecules or ions of the solute already present.
As the concentration of these ions or molecules is decreased by this
reaction, more of the solute must pass into the solvent in order to
maintain the saturation equilibrium. Therefore, after the treatment
with water the preparation of the solution consists in the addition
of reagents which, because of various types of reactions with the
solute ions or molecules, increase the amount of the solid passing
into the solvent.

5 For a general discussion of methods of preparing a solution of various
materials, see Falee, Inorganic Quantitative Analysis; for methods especially
adapted to silicate and carbonate rocks, see Hillebrand, Analysis of Silicate
and Carbonate Rocks , or Hillebrand and Lundell, Applied Inorganic Analysis.

172 PREPARATION OF THE SOLUTION [P. 5

The Solvent Action of Nitric Acid. The first of the reagents so
added is nitric acid. As was shown in Table VIII, two general
types of solvent action are obtained by the use of this acid. The
first of these includes what may be termed the hydrogen ion effects
and would be obtained by the use df any strong acid. These effects
are the result of the tendency of hydrogen ions to enter into three
different types of reactions, namely: (1) neutralization, (2) dis-
placement, and (3) oxidation. The first of these hydrogen ion
effects results from the fundamental neutralization reaction when
water is the solvent medium:

H+ + OH" = H 2 O.

Because of this, those substances, such as the hydroxides, basic
oxides, and basic salts, which ionize in aqueous solution to produce
hydroxyl ions are dissolved by strong acids. In certain cases these
substances may have been converted usually by high tempera-
tures into forms which dissolve so slowly that they are considered
as being practically insoluble; examples are ignited aluminum and
stannic oxides.

The second hydrogen ion effect, displacement, is a result of the
type reaction

H+ + A" = HA,

where A~ represents the anion of a weak acid. This accounts for
the general increase in the solubility of the salts of weak acids in the
presence of other acids. The term displacement arises from con-
sidering that the weaker acid is displaced from its salt. The action
of hydrogen ion as an oxidizing agent, in consequence of the elec-
tronic reaction

2H+ + 2E~ = H 2 ,

is of most importance as a solvent effect in the solution of certain
reducing metals, such as zinc or aluminum. Neutralization, dis-
placement, and electronic reactions have been discussed from the
mass-action point of view in the section dealing with the preparation
and standardization of solutions.

The second general solvent effect obtained from the use of nitric
acid is due to its strong oxidizing tendency. This tendency makes it
especially effective in dissolving metals, alloys, and sulfides. Sul-
fide ion is oxidized to elementary sulfur or to sulfate, depending upon
conditions, such as the concentration and temperature of the acid.
The course of the reaction and the products resulting from the

P. 5] TREATMENT WITH NITRIC ACID 173

reduction of nitric acid are determined largely by the concentration
of the solution and the potential of the reducing agent. The two
principal reduction reactions with their potentials are as follows:

1. H 2 + N0 2 = N<V + 2H+ + E~; -0.79 v.

2. 2H 2 O + NO = N0 3 ~ + 4H+ + 3E~; -0.94 v.

In dilute solutions, except with very powerful reducing agents,
nitric oxide is the product; in concentrated solutions nitrogen dioxide
is the invariable product. The latter is true because, even though
the direct product of the reaction is nitric oxide, it is oxidized by
concentrated nitric acid, the equilibrium constant for the reaction

NO + 2NO 3 ~ + 2H + =, 3NO, + H 2

being 5 X 10~ 9 . With dilute solutions of nitric acid and very
powerful reducing agents, such as zinc, the reduction product may
be ammonia, as follows:

3H 2 O + NH 4 + = NOr + 10H+ + 8E~; -0.77 v.

It is to be emphasized that the rates with which the abova reactions
take place are often so extremely slow that predictions made from
the potential values would be misleading. Thus cold solutions 1 n.
in nitric acid can be saturated with hydrogen sulfide without appre-
ciable reduction taking place; the reaction becomes rapid only when
the acid concentration is greater than 2 n. and the solution is hot.
The specific behavior of many of the more common substances upon
treatment with nitric acid is mentioned in Note 12 of this procedure.
Either nitric or hydrochloric acid could have been used first in
this systematic treatment of an unknown solid, and each has its
merits and disadvantages. Thus hydrochloric acid furnishes a
non-oxidizing solvent which is more generally effective in bringing
metallic oxides into solution than is nitric acid. Owing to its reduc-
ing character, it dissolves insoluble oxidizing substances, such as
manganese dioxide and lead chromate, and it reduces oxidizing
agent3 such as permanganates and chlorates, thus minimizing the
formation of sulfur during the precipitation of the Hydrogen Sulfide
Group- Against these desirable features of the use of hydrochloric
acid are to be considered the danger of loss of certain elements,
especially arsenic in the tripositive state, and the formation of in-
soluble silver chloride, which makes necessary, whenever silver is
present, a long and troublesome treatment before complete solution
of the substance is obtained. The initial use of nitric acid avoids
the possibility of loss of arsenic by volatilization of the trichloride,

174 PREPARATION OF THE SOLUTION IP. 5

causes solution of most silver compounds, and is more generally
effective in attacking alloys. Its objectionable oxidizing action,
especially that of converting sulfides to sulfate, is largely avoided by
using first a cold dilute solution so that most of the hydrogen sulfide
formed may be expelled without oxidation. Also, it has been found
possible to provide for the reduction of most oxidizing substances
by using a simple treatment with formic acid. For these reasons,
a treatment with nitric acid is provided for first in this system,
although its use is governed to a large extent by the information
gained in the preliminary tests in P. 3.

Procedure 5: TREATMENT OF THE SAMPLE WITH NITRIC
ACID. (I) Non-Metallic Solids. If the original material
is a non-metallic solid, weigh out a sample of approximately
1 g and transfer it to a 200-ml flask (Notes 1, 2, 3).

Add to the sample 20 ml of water and, if the substance does
not dissolve completely, heat the mixture almost to boiling
(Note 4). Test the solution with litmus paper, make it
just acid with HNOs (or first just basic with NH 4 OH, if
it is already acid) and then add just 5 ml more HNOs (Note
5). (If a clear solution is obtained, treat it by P. 11.)

If the substance has not dissolved, cover the flask with a
watch glass and slowly heat the mixture almost to boiling;
hold the temperature at that point, without evaporating
any of the acid, as long as the substance seems to be dis-
solving (Notes 6, 7). (If the substance has dissolved, treat
the solution by P. 11.)

If the substance has not dissolved, remove the watch
glass, evaporate the mixture to 2 to 3 ml, add to it 5 ml of
16 n. HNO 3 , and then evaporate it almost to dryness (Note
8), using a steam bath to remove the last 1 to 2 ml of the
liquid.

If, in P. 3, the presence of oxidizing substances was not
detected, omit the remainder of this paragraph. If oxidiz-
ing agents were found present, add to the residue 5 ml of
water 1 ml at a time, warming after each addition, and 1 to
15 ml of a 90 per cent solution of HCHOj, formic acid (Note
9). Heat the mixture just to boiling as long as any reaction
seems to be taking place, and then evaporate the mixture
almost to dryness as directed above. Add 5 ml of HN0 3
and again evaporate almost to dryness.

P. 5] TREATMENT WITH NITRIC ACID 175

If the original substance has not been converted into a
soluble form (Notes 11, 12), treat the residue by P. 6 (Note
10).

If the original substance has apparently dissolved, add to
the residue just 5 ml of HN0 3 and then 10 to 30 ml of water
(Note 13). If the residue dissolves slowly, cover the flask
and heat the mixture to 80 to 90C. If a clear solution is
obtained, treat it by P. 11 (Note 14); if not, treat the mix-
ture by P. 6

(II) Metallic Solids, If the original material is a metallic
solid, weigh out a sample of approximately 0.5 g and trans-
fer it to a 200-ml flask (Notes 1, 2, 3)

Add to the substance 10 ml of HN0 3 , cover the flask
with a watch glass, slowly heat the mixture just to boiling,
and gently boil it as long as the substance seems to be dis-
solving. If there seems to be a slowly dissolving residue,
add 5 ml of 16 n. HNO 3 to the mixture and again boil it.

If the substance has not been converted into a soluble
form (Notes 11, 12, 14), treat the mixture by P. 6 (Note 10).

If the substance has apparently dissolved, evaporate the
solution just to dryness, using a steam bath to remove the
last 1 to 2 ml of liquid.

Treat the residue by the last paragraph of Part I of this
procedure.

(III) Solutions. If the original material is a solution or
suspension, measure out into a flask that volume of the
solution which is estimated to contain 1 g of solid material
(Note 15), make it neutral with NH 4 OH or HNO 3 , evaporate
(or dilute) it to 50 ml, add just 5 ml of HN0 3 , and treat it by
P. 11 (Note 16).

Notes:

1. The sample should have been subjected to the preliminary observa-
tions and tests of P. 3, and the decision as to whether it should be treated
first with HN0 3 , by this procedure, or directly with HC1, by P. 6, should be
based upon any previous knowledge of the substance and upon the informa-
tion gained in the tests of P. 3 (see the last paragraph of the discussion
above and also Note 10, P. 3).

If, in P. 3, the sample has appeared to be insoluble in the HC1, HN0 3 ,
and the mixed acids, time will be saved by omitting both P. 5 and P. 6 and
proceeding directly to P. 7.

2. The sample should have been prepared by P. 1 and should have been
reduced to a very fine powder. It is recommended that the sample be kept

176 PREPARATION OF THE SOLUTION [P. 5

in a closed glass "weighing bottle" and that the sample be weighed by
difference. That is, the weight of the bottle and material is found, a portion
of the substance thought to be slightly less than 1 g is tapped directly into
the flask, and the bottle is then closed and approximately weighed. From
the weight thus found, the additional amount to be taken to give a sample
of approximately 1 g can be estimated, this amount added to that in the
flask, and the weighing bottle then weighed to the precision desired.

3. The precision with which the sample should be weighed or measured
is determined by the accuracy desired in the analysis. Thus, a 1-g sample
weighed to within 1 mg will justify results to db 0.1 per cent, which, except
in special cases, is greater than the precision which can be attained by the
separations and estimations of this system of analysis.

4. It is sometimes desired to determine separately the portion of the
substance which is "water soluble. " In this case, after treating the sub-
stance a sufficient length of time to be sure that the soluble constituents are
extracted, the mixture should be filtered, decanting as much as possible, the
residue washed, the acidity of the filtrate adjusted as directed in the next
sentence of the procedure, and this solution then analyzed separately.

5. A careful adjustment of the acidity is necessary at this point, as the
subsequent separation of the Hydrogen Sulfide Group from the other ele-
ments is highly dependent upon attaining a certain acid concentration
namely, 0.3 n. when the solution is later treated with H^S.

If the test papers mentioned in Note 6, P. 3, are available, they can be
used to advantage in this procedure, as, by neutralizing this solution to a
pH of approximately 3, the extensive precipitation of easily hydrolyzable
constituents, such as bismuth and ferric salts, may be prevented and a
more precise adjustment of the acidity for the hydrogen sulfide precipitation
may be obtained.

6. The mixture is warmed slowly in order that, insofar as possible, any
sulfides present may dissolve and the hydrogen sulfide formed may be
expelled without oxidation. Oxidation of the hydrogen sulfide would result
in the formation of sulfur and sulfate, the latter causing precipitation of any
barium or strontium, or, if present in considerable amounts, of lead and
calcium.

7. If there is a crystalline residue which seems more soluble in the hot
solution, or which seems to require only a larger volume of solution in order
to dissolve completely, the mixture should be diluted to 50 ml and, if a clear
solution is obtained, treated by P. 11.

8. When evaporating a solution almost to dryness, only sufficient liquid
should be left to keep the residue wet. This operation is used when drying
the residue might cause it to be converted into an insoluble or slowly dis-
solving form, or cause loss due to spattering or volatility of some of the con-
stituents. The use of a steam bath to remove the last few milliliters of
solution eliminates spattering and the danger of overheating portions of
the residue. This operation can be carried out over a flame if the vessel is
kept in constant motion and if all parts of the residue are continuously
kept wet; however, the danger of loss by spattering has to be carefully
guarded against.

9. The volume of HCHO^ to be added can be judged by the amount of

P. 5] TREATMENT WITH NITRIC ACID 177

oxidizing substances found in P. 3. Reduction of a permanganate is indi-
cated by the disappearance of pink color; reduction of a chromate is indi-
cated by the formation of a clear green solution with disappearance of the
yellowish color. The HCH02 is oxidized to carbon dioxide and water.

10. If an unchanged residue of the original sample is apparent in con-
siderable amount, this residue should be treated directly by P. 6, unless the
presence of silver has been indicated in P, 3. In that case the residue should
be treated by the next paragraph'of this procedure. The formic acid may
produce a dark precipitate, owing to the precipitation of metallic mercury
or silver; this residue should dissolve when it is treated with the 5 ml of 16 n.
HN0 3 and the mixture is again evaporated almost to dryness.

11. If there is doubt as to whether or not the substance has dissolved, or
if there is a considerable crystalline residue which may dissolve when treated
with acid or when the solution is diluted, the residue should be treated by
the next paragraph of the procedure.

12. Of the more common substances which would remain undissolved
by treatment with concentrated HN0 3 , there may be listed the following
groups: (1) AgCl, AgBr, Agl, BaS0 4 , SrS0 4 , PbS0 4 ; (2) HgS; (3) A1 2 3 ,
Fe 2 3 , Cr0 3 , Cr 2 (S0 4 ) 3 , SnS, Sn 3 (P0 4 ) 4 ; (4) Mn0 2 , Pb0 2 , Sn0 2 , Sb 2 3 ,
Si0 2 ; (5) many silicates; (6) silicides of such metals as iron, manganese,
and chromium, SiC (carborundum), C (graphite), Si, and many resistant
alloys. Substances dissolving to only a slight extent would include CaSC^,
CaF 2 , PbCr0 4 , BaCr0 4 , BiP0 4 , BiAs0 4 , and many cyanides and complex
cyanides. It is seen that these compounds include several classes of sub-
stances, and their insolubility is explained by various causes. Thus the
first group is composed of insoluble salts of strong acids, and it would not
be expected that they would greatly increase in solubility in acid solutions.
It would be expected that HgS would dissolve in an acid solution, because
HgS is a salt of a weak acid and, moreover, sulfide is oxidized by concen-
trated HN0 3 ; however, HgS is so extremely insoluble that it is only very
slowly attacked. The group including A1 2 3 , Cr 2 (S0 4 ) 3 , SnS, and Sn 3 (P04)4
exhibits a rather common phenomenon, being composed of substances which,
when first precipitated, are quite soluble but which when heated or found
native are not appreciably attacked by acids even on prolonged treatment.
Chromic sulfate behaves in this manner, even a short fuming with sulfuric
acid converting it into a very resistant compound having the unusual formula
Cr 4 H 2 (S0 4 ) 7 . 6 Mn0 2 , Pb0 2 , and, more typically, Sn0 2 , Sb 2/ 3 , and Si0 2
are acidic oxides which are insoluble in non-reducing or non-complex-forming
acids. Many silicates are not appreciably attacked by acids other than HF,
and those which are decomposed give a precipitate of hydrous silicic acid.
Finally, there is a large group of substances, such as certain silicides, car-
bides, elementary substances (graphite, silicon), and alloys, which are very
resistant to nitric acid. lit some cases this may be due partly to the forma-
tion of protective coatings.

By the treatment with formic acid most of any insoluble oxidizing com-

* Caley, J. Am. Chem. Soc., 55, 3947 (1933), has shown that this and many
other resistant compounds such as SnCh, PbSCh, and the alkaline earth ele-
ment sulfates are readily decomposed by hot concentrated hydriodic acid.

178 PREPARATION OF THE SOLUTION [P. 6

pounds, such as Mn02, PbOa, or PbCrO 4 , are reduced and consequently
dissolve in the HNOa next added.

13. The larger volume of water should be added if there is considerable
crystalline residue which seems to be dissolving with the addition of more
water. If, however, the residue dissolves completely in the acid and a
precipitate forms upon adding water, the mixture should be treated, without
filtering, by P. 11. Such a precipitate is caused by the hydrolysis of com-
pounds of antimony, tin, or bismuth; as these precipitates are metathesized
into sulfides in P. 11, it is not necessary to redissolve them.

14. If only a small residue is left, it is usually advisable to filter the solution
through an ashless filter paper, decanting as much as is possible in order to
retain the residue in the flask, to wash the residue and paper by decantation
with 5 ml of hot water, and to treat the filtrate and washings by P. 11 and
the residue by P. 6. Also, if silver has been detected in P. 3 or is thought
to be present, it is better to treat the solution obtained at this place (which
will usually contain all of the silver) by P. 11, thus avoiding the formation
of difficultly soluble AgCl in the next procedure.

If the substance is an alloy containing antimony or tin, these two ele-
ments will be rendered insoluble by the evaporation with HNOa and will
remain as a white residue. Often, especially in the case of alloys, it is
advisable to filter out this precipitate and thus separate these elements at
this point. Furthermore, as antimony is often known to be absent, the
residue will consist only of hydrated SnO2 and can be either ignited and
weighed or dissolved in HC1 and treated directly by P. 49 to determine the
amount present.

15. If the material is a suspension, either (a) treat the suspension as
though it were a clear solution, or (b) filter out the suspended material and
treat this residue by Part I and the solution by Part III of this procedure.
If it is desired to know the weight of suspended matter in a given volume of
solution, a known volume of the original solution should be filtered, and the
precipitate dried and weighed.

16. If a precipitate forms on neutralizing the solution and does not redis-
solve on adding HN0 3 , it should be filtered out and treated by Part I of P. 5.
From an alkaline solution such a precipitate would result from sulfides held
in solution by an excess of an alkaline sulfide, from silver halides dissolved
in a cyanide or ammonia solution, from amphoteric hydroxides (antimony
or tin) dissolved in an excess of a hydroxide, or from soluble silicates which
on acidification precipitate silicic acid. If it is known that the precipitate
is caused by elements of the Hydrogen Sulfide Group for example, by the
hydrolysis of antimony or bismuth salts, or by silver halides held in an alka-
line solvent the precipitate need not be filtered out, as it will be meta-
thesized to sulfide by the H2S treatment in P. 11.

P. 6. Treatment of the Sample with (1) Hydrochloric Acid, and
(2) Hydrochloric and Nitric Acids

Discussion. In this procedure the systematic treatment of the
sample is continued by the use of hydrochloric acid and, where it is
necessary, a mixture of nitric and hydrochloric acids.

P. 6] HYDROCHLORIC AND NITRIC ACIDS 179

The Solvent Action of Hydrochloric Acid. By the use of a concen-
trated hydrochloric acid solution three general solvent effects are
obtained. The first of these is the hydrogen ion effect, which is
obtained by the use of any strong acid. The second is obtained be-
cause of the tendency of chloride ion to reduce certain higher oxides
(for example, MnO2 and PbO2) and also the anions of certain oxygen
acids, especially the insoluble chromates. The third general effect
is due to the tendency of chloride ion to form soluble complex com-
pounds with most of the elements of the Hydrogen Sulfide and
Ammonium Sulfide Groups. These complex compounds and their
influence on the solubility of the sulfides of the Hydrogen Sulfide
Group elements are discussed in P. 11.

The Solvent Action of a Mixture of Nitric and Hydrochloric Acids.
A mixture of nitric and hydrochloric acids, the so-called "aqua
regia" solution, is a tremendously effective solvent, as it combines
in one solution the solvent effects of both nitric and hydrochloric
acids. Thus there is obtained the oxidizing effect of nitric acid with
the complex-forming tendency of hydrochloric acid. The particular
effectiveness of this mixture in dissolving the so-called noble metals
(gold and the platinum group elements) is due to the strong tendency
of these metals to form very stable chloride complexes (AuCU" and
PtCl~ are examples), and not to the existence of an exceptionally
high oxidizing potential frequently attributed to the presence of
nitrosyl chloride (NOC1) and other products of the reaction between
the nitrate and chloride ions.

After this treatment, the mixture is evaporated almost to dryness
in order to render more insoluble any silicic acid which may have
been formed. Care has to be taken not to boil the hydrochloric acid
solution before addition of the nitric acid (or other oxidizing agent),
as arsenious chloride is quite volatile; on evaporating 25 ml of 6 n.
HC1 containing 250 mg of arsenic (as arsenious chloride) to 5 ml,
97 per cent of the arsenic was lost.

Procedure 6: TREATMENT OF THE SAMPLE WITH (1)

HYDROCHLORIC ACID, AND (2) HYDROCHLORIC AND NITRIC

ACIDS. If the sample has been treated by P. 5, treat the
mixture or the residue obtained in that procedure, in the
same 200-ml flask (Note 1), as directed in the third para-
graph of this procedure.

If the sample has not been treated by P. 5, weigh out a
sample of approximately 1 g (Notes 2, 3, P. 5) if it is a non-
metallic solid or 0.5 g if it is a metallic solid, transfer it to a
200-ml flask, and treat it as directed in the next paragraph.

180 PREPARATION OF THE SOLUTION [P. 6

Add to the residue (or sample) 10 ml of 6 n. HC1, cover
the flask with a watch glass, and heat the mixture to 80 to
90C. (do not boil it!) as long as the substance seems to be
dissolving. If the substance has dissolved, add to the solu-
tion 10 to 30 ml of water (Note 2; also Notes 11, 13, P. 5)
and just 4 ml of NH 4 OH and treat it by P. 11.

If the substance does not seem to be dissolving, or is only
slowly dissolving, evaporate the mixture to 3 to 5 ml (Note
3), add to it 10 to 20 ml of 12 n. HCi and 5 ml of 16 n. HN0 3 ,
cover the flask with a watch glass, and again heat the mix-
ture just to boiling as long as the substance seems to be
dissolving (Note 4). Finally evaporate the mixture almost
to dryness (Note 3), add to the residue just 5 ml of 6 n. HCI,
warm the mixture with the flask covered until all the soluble
salts seem dissolved, and then add to the mixture 10 to 20 ml
of water (Note 13, P. 5). Filter the solution through a
quantitative (ashless) paper filter (see p. 134) and wash
the residue with two 2-ml portions of 1 n. HCI and then
with five 2-ml portions of hot water, collecting these wash-
ings with the filtrate. Treat the filtrate by P. 11. Treat
the residue by P. 7.

Notes:

1. If the insoluble residue obtained in P. 5 has been filtered, it should
have been retained in the original flask insofar as was possible by filtering
and washing by decantation. The residue carried onto the filter can be
combined with that in the flask in the following manner: Open the filter and
carefully tear away that portion of it to which no residue has adhered. Lay
the filter against the side of a funnel (the funnel just used, or, for this and
similar operations, a larger one with the stem, cut off short is more con-
venient; a funnel with a capillary stem cannot be used). Wash the residue
from the paper into the flask, using the 10 ml of HCI next to be added and
applying it in a fine stream by means of a dropper. Finally, the paper can
be similarly washed with hot water. Usually such a small fraction of the
residue will adhere to the paper that it can be neglected and the paper can '
be discarded. If an appreciable amount of the residue remains on the
paper, or if a very accurate analysis is desired, the paper should be folded
into a compact roll, and a stout platinum wire, which has one end sealed
into a glass rod, wrapped in a spiral around it. The paper is then held above
a crucible (which is preferably placed in the center of a square of black
glazed paper), and the filter is gradually heated with a small oxidizing flame
from a gas burner until it has completely burned. Any ash remaining on the
wire can then be brushed into the flask with a small camel's-hair brush, and
any ash which has dropped into the crucible or onto the paper can be simi-
larly added to the flask. If unburned portions of the filter paper drop into

P. 6]

RECOVERY OF VOLATILE CHLORIDES

181

the crucible, they can be ignited by heating the crucible. The black oxides
of certain metals should not be mistaken for charred paper. After cooling,
the residue is transferred to the flask.

2. Certain moderately soluble chlorides, such as PbCh and HgCl?, may
precipitate and not be dissolved completely with this volume of water; in
such cases, the mixture, or a portion of it, may be further diluted and the
effect noted.

3. The hydrochloric acid solution has not been boiled before this in order
to avoid possible loss of arsenic, or, to a much smaller extent, mercury,
antimony, or tin. Especially if arsenic is thought to be present, the evapora-

Fig. 26. Apparatus to Prevent the Loss of Volatile Chlorides.

tion should be carried out in such a way that the volatile chlorides of these
elements can be condensed and recovered. This can be accomplished by
transferring the mixture to a distilling flask and distilling it to the volume
indicated, at the same time collecting the distillate in 50 ml of cool water.
The operation can be more rapidly and effectively performed as follows:
Place in the flask a two-hole rubber stopper carrying in one hole a glass inlet
tube which extends about two thirds of the distance to the solution (see Fig.
26). In the other hole insert a tube which is bent in the form of a U so that
it can be extended through a hole in a similar stopper in another 200-ml
flask. The second flask should contain 50 ml of water, and the tube should
be made to extend just under the surface of this water. In the second hole
of the stopper of this flask insert a short glass tube and connect this by
means of rubber tubing to a water aspirator, which is adjusted so that a

182 PREPARATION OF THE SOLUTION [P. 7

rapid stream of air is drawn through the two flasks. By boiling the mixture
in the first flask and cooling the second flask with a bath of cold water,
rapid evaporation is obtained and any volatile chlorides are quite completely
condensed in the second flask. The condensed solution, after being used
also for the second evaporation in this procedure, is made neutral to litmus
with NEUOH, 5 ml of 6 n. HC1 are added to it, and, disregarding any precipi-
tate, it is treated by P. 11. Any precipitate obtained there is combined with
any precipitate obtained by P. 13 or is treated separately by P. 41 to P. 49.
The filtrate is discarded.

4. If there is a residue which seems to be slowly dissolving, more of the
two acids should be added and the boiling continued. Of the insoluble sub-
stances mentioned in Note 12, P. 5, those which would probably dissolve
would include HgS, Fe20a, SnaCPO^, Sb20a, most cyanides and complex
cyanides, BiP04, and BiAs04. The solution of HgS is due to the stability
of the HgCU" ion which is formed, combined with the oxidation of the
sulfide by the mixture. Oxides and salts of the metals of the Hydrogen
Sulfide Group and the Ammonium Sulfide Group are in general more soluble
in HC1 than in HNOa, because of the tendency of these elements to form
complex compounds in HC1 solutions. Any oxidizing substances, such as
MnOa or PbCrOi, not reduced by the formic acid would be reduced and dis-
solved by the concentrated HC1. The slowly dissolving substances, such
as ignited or native A^Os or SnS2, would be only slightly attacked by the
mixed acids.

P. 7. (1) Treatment of the Sample with Perchloric Acid; (2) Detec-
tion and Estimation of Silica by the Use of Hydrofluoric Acid

Discussion. By heating the residue from the "aqua regia" treat-
ment with nitric and perchloric acids until the latter fumes, several
effects are obtained. First a powerful oxidizing action is obtained.
This is due at first entirely to the oxidizing effect of the nitric acid
in the dehydrating medium furnished by the mixed acids; perchloric
acid does not exhibit its powerful oxidizing tendency until it be-
comes quite concentrated and hot. Then it may oxidize manganous
ion to manganese dioxide or chromium to chromic acid (see discus-
sion of P. 4).

The Solvent Action of Fuming Perchloric (or Sulfuric) Acid. Next,
upon fuming the perchloric acid, many insoluble salts of volatile
acids are dissolved, the acids being volatilized from the solution at
the high temperature of the fuming acid (280C.). Thus, such
salts as calcium fluoride, silver bromide, and silver iodide are dis-
solved. The solvent action on the last two is made more effective
because of oxidation of iodide and bromide by the fuming acid;
some of the iodide is even converted to iodate. Silver chloride is
more slowly dissolved, however, as suggested in Note 2, if the chlo-

P. 7] PRECIPITATION OF SILICIC ACID 183

ride is metathesized into bromide by addition of hydrobromic acid;
this can then be dissolved by further fuming with the perchloric
acid. Sulfuric acid is even more, effective, because of its higher
boiling point, in dissolving the insoluble salts of volatile acids; thus
silver chloride is readily dissolved by fuming sulfuric acid.

The Precipitation of Silicic Acid by Fuming Perchloric Add
Finally, the fuming perchloric acid has the effect of partially de-
hydrating and quantitatively precipitating silicic acid. This effect
is obtained because (1) silicic acid is a very slightly ionized acid,
and therefore most silicates are decomposed by treatment with a
strong acid: (2) silicic acid is an insoluble acid, although when first
precipitated it tends to form colloidal solutions, which are not
coagulated except by partially dehydrating the silicic acid; (3)
fuming perchloric acid is a good dehydrating medium and is thus
effective in causing the precipitation of colloidal silicic acid; and (4)
perchloric acid does not form any soluble complex acids with silicic
acid, as does hydrofluoric acid.

Silicic acid is more commonly dehydrated and made insoluble by
evaporating a hydrochloric acid solution to dryness, heating the
residue for some time, digesting with hot dilute hydrochloric acid
until the soluble constituents are dissolved, and then filtering out
the silicic acid. Disadvantages of this process are (1) the time re-
quired for the evaporation and drying, (2) the fact that the pre-
cipitate usually retains an appreciable amount of the soluble basic
constituents of the solution, and (3) the fact that the process has to
be repeated with the filtrate in order to precipitate the silica quan-
titatively. The amount of silica remaining in the first filtrate de-
pends upon the quantity and nature of the other constituents pres-
ent, and various other factors such as the temperature and the time
for which the residue is dried; often as much as several per cent of
the total will pass into the first filtrate.

The dehydration and precipitation of silicic acid by fuming per-
chloric acid has been the subject of several investigations; 7 these
have shown that practically complete precipitation can be obtained
and that the precipitate is cleaner than that obtained by the evapo-
ration process. Because of these facts and the quickness of the
method, it is used here.

Fuming sulfuric acid will also dehydrate and precipitate silicic

f Willard and Cake, /. Am. C/iem. Soc., 42, 2208 (1920); Gibson, Rock Prod-
ucts, 35, 70 (1930); Meier and Fleishman, Z. anorg. Chem., 88, 84 (1932); Fish
and Taylor, /. Chem. Ed., 10, 246 (1933).

184 PREPARATION OF THE SOLUTION [P. 7

acid; this process, while frequently used in the analysis of ferrous
alloys, is not practical for qualitative systems because of the pre-
cipitation of insoluble sulfates.

In the large majority of cases, the residue remaining after this
treatment will consist entirely of silicic acid, which is estimated by
heating it to a temperature sufficiently high (1000 to 1200C.) to
convert it to anhydrous silicon dioxide and then weighing it. Most
other basic constituents remaining in the residue are converted into
oxides at this temperature and are weighed as such. In order to
eliminate the silica, the ignited residue is treated with hydrofluoric
acid in the presence of concentrated perchloric acid. The hydro-
fluoric acid reacts with silica (or silicates) as follows :

SiOi + 4HF = SiF 4 + 2H,0.

If an excess of the acid is added, the Complex fluosilicic acid is formed,
thus:

SiF 4 + 2HF = H*SiF fl .

However, upon fuming the solution, this reaction is reversed as the
volatile silicon 'tetrafluoride and hydrogen fluoride are expelled.
Therefore, by this operation the complete elimination of silicon is
accomplished. Often, after this process, no significant residue
remains and further treatment is obviated. If any residue remains,
the perchloric acid is evaporated, and the residue is heated as before
and again weighed; the weight of this residue is subtracted from the
first weight to give the true weight of the silica present. The
residue, -if appreciable, is treated with hydrochloric (or nitric) acid,
and the solution so obtained is analyzed for the basic constituents.
Only rarely will there be any residue after the treatment with these
acids; in case there is, it should be fused with Na2COs as directed in
the next procedure. Reference should be made to Hillebrand,
Analysis of Silicate and Carbonate Rocks, Bulletin 700, U. S. Geo-
logical Survey, 1924, or to Hillebrand and Lundell, Applied Inor-
ganic Analysis, Wiley, 1929, for an extensive discussion of the methods
for determining silica.

If it is not desired to estimate the amount of silica present, the
dehydration and precipitation of the silica and the ignition and
weighing of the silica and the residue remaining after the hydro-
fluoric acid treatment can be dispensed with. The silica is volatilized
by adding the hydrofluoric acid to the fumed perchloric acid solu-
tion, and the fluorides and excess hydrofluoric acid are volatilized
by again fuming the perchloric acid solution.

P. 7] ESTIMATION OF SILICA 185

Modified procedures are given below for the case (A) where it is
desired only to detect and remove the silica and (JB) where an estima-
tion of the amount present is wanted. In Note 11 below is given a
procedure whereby the estimation of the silica in silicates which are
decomposed by acids can be inade directly without going through the
previous procedures of this system.

Procedure 7: (1) TREATMENT OF THE SAMPLE WITH PER-
CHLORIC ACID; (2) DETECTION AND ESTIMATION OF SILICA.
Transfer the residue from P. 6 to a 200-ml round-bottom
flask (Note 1, P. 6), add to it 5 ml of 9 n. HC10 4 (Note 1)
and 1 ml of 16 n. HNO 3 , and heat the mixture until dense
white fumes are evolved. If there is a residue which seems
to be slowly dissolving, add cautiously 1 ml more of 16 n.
HNO 3 , insert a test-tube condenser in the neck of the flask,
and boil the mixture as long as the residue appears to be dis-
solving (Note 2); then again evaporate the mixture to
fuming (Note 3). If the substance has dissolved, add to it
10 to 20 ml of water and treat it by P. 11.

(A) Procedure for Use when Only the Detection and Elim-
ination of Silica Is Desired.

Transfer the mixture to a platinum crucible, using as little
water as possible, again evaporate it just to fuming, and
allow it to cool (Note 4). Place the crucible under a hood,
add to it 5 or 6 drops of 48 per cent HF (Caution: Notes 5, 6),
and then warm it to 70 to 80C. (Formation of bubbles,
presence of silica.) If silica is present, add 5 ml more HF
and keep the mixture just boiling (or evaporate it on a steam
bath) as long as the residue seems to be dissolving. If the
residue seems to be slowly dissolving, add repeatedly 1-ml
portions of HF and of 16 n. HNOa after the previous portion
has almost entirely evaporated or the HC10 4 begins to fume.
Finally, evaporate the mixture almost to dryness (Note
7) and treat it by the last paragraph of this procedure.

(B) Procedure for Use when an Estimation of the Silica
Is Desired.

Cover the flask with a watch glass and keep the mixture
just fuming for 10 min., adding more HC104 if the mixture
tends to solidify. Cautiously add 50 ml of water, heat the
mixture almost to boiling for 2 to 3 min., and filter the hot
solution through a quantitative paper. Wash the residue

186 PREPARATION OF THE SOLUTION [P. 7

with 10 to 20 ml of hot 0.12 n. HNO 3 , collecting the washings
with the filtrate (treat the filtrate by P. 11). Wash the
precipitate and the entire filter with hot water until it is
free of acid (Note 10).

Remove the filter from the funnel (see Note 5, P. XVIII
C), fold it, and transfer it to a platinum crucible which has
been previously heated and weighed. Heat with the low
flame of an ordinary burner until the paper is dry, and then
slowly increase the heat until the paper has completely
charred and the carbon has burned without any flaming.
Cover the crucible and slowly raise it to the highest tem-
perature obtainable with a Meker or similar burner for 10
min., taking care that the platinum is exposed to an oxidiz-
ing flame only (Note 1, P. 8). Cool the crucible in a desic-
cator and weigh it. Repeat the heating and weighing of
the crucible until its weight is constant to 0.2 mg. Add to
the residue 2 ml of 9 n. HC10 4 . Place the crucible under a
hood, slowly add to it 5 or 6 drops of 48 per cent HF (Cau-
tion: Notes 5, 6), and warm it to 70 to 80C. (Formation
of bubbles, presence of silica.) If silica is present, add 3 ml
more HF and evaporate the mixture almost to dryness, add
0.3 g of solid oxalic acid (H2C 2 4 -2H20), evaporate to
dryness, and heat the crucible with a burner as before for
5 min. (Note 12). Allow it to cool and again weigh it.
From the loss in weight calculate the amount of silica
which was present. If there is any residue, place the
crucible under a hood, add to it 1 to 2 ml of 9 n. HC1O 4 ,
evaporate the mixture almost to dryness, and treat it by
the next paragraph.

Transfer the residue to a 200-ml flask with the aid of as
little water as possible, add to it 5 ml of 6 n. HC1 (Note 8),
cover the flask with a watch glass, and heat the mixture
just to boiling as long as the residue seems to be dissolving.
Add to the mixture 10 to 20 ml of water, filter the hot solu-
tion through an ashless filter, and wash it with five 2-ml
portions of hot water. Treat the filtrate by P. 11. Treat
the residue by P. 8 (Note 9).

Notes:

1. If HC10 4 is not available, H 2 SO 4 may be substituted for it. In that
case, 10 ml of 6 n. H 2 S0 4 should be used. No danger is involved in using
perchloric acid at this point, as all organic material has been removed and

P. 7] PERCHLORIC AND HYDROFLUORIC ACIDS 187

any reducing material has been oxidized by the treatments with concen-
trated nitric acid.

2. If much silver is present (as AgCl), it will dissolve only incompletely
even on prolonged treatment. If the preliminary treatment in P. 3 has
indicated the presence of silver, or if a characteristic residue of AgCl is seen
to be present, it can be decomposed as follows:

Add to the flask 1 to 3 ml of 9 n. HBr, heat the mixture, add
1 ml of 16 n. HNOa, and again fume the HC104 (using a test-
tube condenser to prevent undue evaporation) until the residue
is dissolved. Replenish the HN0 3 as it is distilled. By this
treatment, even 500 mg of silver can be brought into solution (see
Note 9).
A test-tube condenser (see Fig.

27) is made by selecting a test tube

which fits loosely in the neck of the

flask and closing it with a two-hole

rubber stopper. One hole of the

stopper carries a glass tube which

extends to the bottom of the test

tube and is connected by rubber tub-
ing to a water supply; the other hole

carries a shorter glass tube with a

piece of rubber tubing leading to a

drain. Cold water is run in through

the longer tube and out the shorter,

the test tube being inserted in the

flask so that it extends slightly

below the neck.

3. If the residue has not been
appreciably attacked by the HC104
and HN0 3 and is thought to consist
only of silica, the mixture can be
evaporated to dryness in a platinum
crucible, after which the residue can
be heated and weighed as directed
in the second paragraph of B of the

procedure above without filtering out the precipitate. If the residue dis-
solves completely in the perchloric acid, the absence of silica is shown.

If manganese is present, it may be oxidized to the dioxide by the fuming
HC10 4 . If a dark precipitate forms on fuming the HC10 4 , add HCHO 2
dropwise to the warm (but not fuming) mixture as long as this precipitate is
being dissolved. Chromium would also be oxidized to chromate by the
HC10 4 and would be reduced by the HCH0 2 .

4. The HC104 can be cooled rapidly by repeatedly immersing the lower
portion of the crucible momentarily in a dish of cold water.

5. HF is an extremely dangerous chemical. It produces painful burns if
spilled on the hands, especially if it gets under the fingernails; in such
cases a paste of borax and dilute acetic acid should be applied immediately.
The fumes are irritating to the nose and lungs, and all operations with HF
solutions should be carried out under an efficient hood.

Fig. 27. A Test-Tube Condenser.

188 PREPARATION OF THE SOLUTION [P. 7

6. Solutions containing HF cannot be handled in glass vessels. The
most convenient apparatus funnels, beakers, graduates, and so forth
for handling this reagent are those made from "transparent bakelite." If
these are not available, apparatus made of hard rubber, or glass vessels
which have been thoroughly coated with paraffin, should be employed.

7. If no residue is apparent, the solution should be completely evaporated
in order to see if silica has been the only constituent of the residue brought
to this procedure.

8. Hydrochloric acid is preferable for general use since it causes the reduc-
tion of any Mn(>2 or chromate formed during the fuming with HC1O4; but
as these would liberate chlorine the residue is transferred from the platinum
vessel before adding the HC1. However, if the preliminary tests of P. 3
have indicated the presence of silver, HNOs should be substituted for the
HC1, and in this case the acid may be added directly to the residue in the
crucible.

9. By the above procedure most of the salts of volatile acids are taken
into solution, and C^Oa and Cr 2 (S04)s are slowly attacked (the chromium
being oxidized to chromate by the hot HCICM. Upon adding the HF, most
silicates are decomposed and silicon and most silicides are also dissolved by
the HF and HNOs, or hot HCICV The residue remaining will be mainly
sulfates of barium, strontium, and lead, a few silicates Such as cyanite
(AUSiOs) and beryl (BesAUSieOis), and also graphite, carborundum (SiC),
and part of some slowly dissolving substances such as A^Oa or SnC>2.

10. If the paper were ignited without washing out all of the HC1O 4 , it
would burn with explosions that might cause loss of material -from the
crucible.

11. If only a determination of the silica in a cement, limestone, or similar
material is desired, proceed as follows:

Weigh out samples of about 1 g into 250-ml beakers. Pour 5 ml
of water over each and then slowly add 5 ml of 12 n. HC1 and 1 ml
of 16 n. HNOs. Let the mixture stand until vigorous action ceases,
add 10 ml of 9 n. HClOi, cover the beaker with a clock glass sup-
ported on glass hooks, and evaporate until the perchloric acid
fumes copiously. Lower the clock glass to the rim of the beaker
and keep the mixture fuming for at least 10 min., adding more
HC1O 4 if the mixture tends to solidify. Dilute the mixture with
50 ml of water, heat almost to boiling, and filter through a quantita-
tive paper filter. Wash the precipitate and the entire filter with
25 ml of 1.2 n. HC1 and then with hot water until it is substantially
free from chlorides (Note 10).

Treat the precipitate as directed in the second paragraph of jB of
the procedure above. Discard the filtrate. Correct the weight
of silica first found by the weight of "non-volatile residue" re-
maining.

Water is added to the sample, and the acids are first added slowly, to pre-
vent loss by spattering when treating a carbonate; hydrochloric acid is used
because of its solvent action on metallic oxides, and nitric acid is used because
of its oxidizing action. Larger volumes of the perchloric acid and of solution
are used here than in the procedure above to minimize coprecipitation. If

P. 8] FUSION WITH SODIUM CARBONATE 19

it is not desired to correct for the "non-volatile residue," a porcelain crucible
may be used instead of a platinum one.

12. If the perchloric acid mixture were evaporated to dryness and heated,
all of the silica would be eliminated as SiF4 and the excess hydrofluoric acid
would be volatilized, but the perchlorates remaining after volatilizing the
excess perchloric acid would be converted to chlorides upon being heated,
whereas it is necessary to weigh the residue as oxides; also, the decomposition
of these perchlorates in contact with the platinum at high temperature might
lead to damage to the crucible. By discontinuing the evaporation when a
small amount of perchloric acid is still present, the decomposition of the
perchlorates into chlorine or other oxidizing constituents which might at-
tack the platinum is avoided. By then adding a large excess of oxalic acid
and again heating the crucible, the small amount of perchloric acid remaining
is eliminated, the excess oxalic acid is sublimed, and the perchlorates are
apparently converted to oxalates, which, upon further raising the tempera-
ture, are rapidly converted to oxides. Sulfuric acid is generally used for
this process of volatilizing the silica and the excess hydrofluoric acid, but it
has the disadvantage that it leaves sulfates, which have to be heated to a
very high temperature in order to decompose them to oxides.

An experimental study of this oxalic acid method 8 has shown that silica
can be eliminated from residues containing iron, aluminum, and calcium
and that these latter elements can be quantitatively converted to their
oxides by this treatment. In most cases the loss in weight of the platinum
crucible during the process did not exceed 0.2 mg.

P. 8. Fusion of the Sample with Sodium Carbonate

Discussion. In most cases, after the treatment with perchloric
and hydrofluoric acids, there will be no residue; however, in order
to insure complete decomposition of a few insoluble sulfates and of
very resistant materials, a fusion with sodium carbonate is provided.
This treatment will usually decompose these more resistant sub-
stances, so that upon treating the melt with water the acidic con-
stituents are extracted and the carbonate-oxide residue of the more
basic elements can be separately dissolved in an acid. For a discus-
sion of the principles involved in the metathesizing action of sodium
carbonate, see the discussion of P. 81. The water extract is acidi-
fied and the two solutions are separately evaporated in order to
dehydrate silica; they are not united, since in some cases this would
cause the reprecipitation of insoluble salts. The group precipitates
later obtained from these solutions may be combined.

It has been shown by Caley and Burford 9 that many of the insol-
uble compounds which would have resisted the previous treatments

1 Unpublished experiments by J. Billheimer.

Caley and Burford, /. Ind. Eng. Chem., Anal. Ed. t 8, 63 (1936).

190 PREPARATION OF THE SOLUTION [P. 8

and which would usually have to be fused with sodium carbonate
are dissolved by concentrated hydriodic acid. Thus, either ignited
or native tin dioxide, which is only slowly, and usually incompletely,
brought into solution by fusion with sodium carbonate, is rapidly
dissolved by being heated to 90 to 95C. with concentrated hydri-
odic acid; the reaction taking place is as follows:

Sn0 2 + 4HI = SnI 4 + 2H 2 O.

The Snl4 is insoluble in the concentrated acid, and the formation
of the characteristic reddish precipitate affords a sensitive indication
of the presence of tin. As this treatment is much shorter and more
easily carried out than the fusion with sodium carbonate, its op-
tional use is provided for in Note 5 below; it is recommended that
it be employed when the residue is thought to consist wholly or
partly of tin dioxide.

The concentrated hydriodic acid is also effective in dissolving
insoluble sulfates. Thus lead and strontium sulfates are readily
dissolved, and barium sulfate and chromium dihydroheptasulfate
are attacked more slowly. This solvent action is due to the reduc-
tion of sulfate to hydrogen sulfide, the type reaction being as follows:

MS0 4 + 10HI = MI 2 + 4I 2 + H 2 S + 4H 2 O.

The evolution of H 2 S affords a sensitive indication of the presence of
a sulfate. The article by Caley and Burford should be consulted
for further details of the action of hydriodic acid with various in-
soluble compounds.

Procedure 8: FUSION OF THE SAMPLE WITH SODIUM CAR-
BONATE. Transfer the dry residue (Note 1, P. 6) from P. 7
to a clean agate mortar and grind it to a powder (Note 5).
Mix it thoroughly with 1 to 3 g of the purest Na 2 COa avail-
able, and transfer the mixture to a platinum (or nickel)
crucible (Note 1) which has 1 g of the Na 2 C0 3 spread over
the bottom. Spread 1 g of the Na 2 COa over the top of the
mixture and gradually heat the crucible with the oxidizing
flame of a blast burner (or a large burner of the Meker
type) until the entire mass is a viscous liquid. If particles
of undecomposed substance remain, or if dark particles
seem to be formed, cool the crucible somewhat and add to
it in small portions 0.2 to 0.5 g of a mixture of equal weights
of Na^COa and NaNOa, and then again heat the crucible as
directed above.

P. 8] FUSION WITH SODIUM CARBONATE 191

When decomposition of the substance seems complete,
allow the crucible to cool (Note 2). Place the crucible and
melt in a 200-ml beaker, add 30 ml of water, and boil the
mixture until the melt is completely disintegrated. Filter
the hot mixture through an ashless paper filter and wash
the residue with 5 to 10 ml of ? n. Na 2 COa, collecting the
washings with the filtrate. Treat the filtrate as directed
in the last paragraph of this procedure.

Transfer the residue to a casserole (igniting the paper),
moisten it with 1 n. HC1 (Note 3), and evaporate it just to
dryness without overheating any part of the residue. Treat
the residue by the third paragraph of P. 6 (Note 4).

Make the filtrate just acid with HC1 (Note 3), taking care
to avoid spattering, and evaporate it just to dryness without
overheating any part of the residue. Treat the residue by
the third paragraph of P. 6 (Note 4).

Notes:

1. A platinum crucible is to be preferred, as, when properly used, it is
not appreciably attacked by the flux and no extraneous constituents are
introduced into the analysis.

There are certain precautions to be taken in order to avoid damage to
platinum ware: Fusions of the above type should not be made in platinum
vessels if elements of the Hydrogen Sulfide Group are present; these elements
are easily reduced and then alloy with the vessel. Compounds containing
sulfur, phosphorus, or arsenic should not be heated with reducing agents,
especially organic matter, as when reduced they react with the platinum.
When platinum vessels are heated with a gas flame, care should be taken to
secure an oxidizing flame, as, in a reducing flame, carbon is deposited on the
vessel, forming carbides of platinum, and hydrogen gas from the flame
penetrates the platinum; either of these results in causing the metal to become
brittle and more rapidly attacked. Platinum vessels are very little affected
by fused sodium carbonate, but are rapidly attacked by more alkaline fluxes
for example, the hydroxides of the alkali and alkaline earth metals and even
sodium or potassium cyanide. Strongly oxidizing fluxes, such as sodium
peroxide or nitrate, should not be used. Moderate amounts of NaNOa
added to a Na2C03 fusion do not rapidly attack the crucible, but only a
minimum amount of the nitrate should be used. Solutions which evolve
the free halogens should not be put into platinum vessels. Sharp implements
should not be used to remove "cakes" from platinum vessels, as the surface
is easily scratched and then more rapidly attacked.

If a platinum crucible is not available, one of nickel may be substituted.
This metal is not so satisfactory, because its greater thickness makes com-
plete fusion more difficult to obtain and because it is appreciably attacked
by the flux, several milligrams of nickel being dissolved with each fusion.

192 PREPARATION OF THE SOLUTION [P. 8

For this reason, nickel cannot be tested for after such a fusion. However,
if much NaNOs is to be used, or if a stronger oxidizing flux, such as Na202,
is to be employed, a nickel crucible should be used.

2. Several methods of facilitating the removal of the melt from the
crucible may be employed. While still quite hot, the lower part of the
crucible may be immersed in a vessel of cold water until the melt is cooled.
After this treatment the melt will usually separate from the crucible quite
readily when treated with a solvent. Another effective method is to insert
a platinum wire into the molten mass and then to let it cool until it is solid.
Upon reheating the crucible, the outer edges of the melt soften and the mass
of the melt can be removed.

3. If arsenic, antimony, or tin compounds are thought to be present,
HN0 8 should be substituted for HC1.

4. The only residue remaining after these residues are treated by P. 6
should be silica, resulting from the decomposition of a silicate which had not
been entirely dissolved by the acid-HF treatment, but which had been
decomposed by the Na2COa fusion. If it is desired to estimate the amount
present, this can be done as directed in P. 7. Owing to the large amount of
Na2COs used, and to the fact that in nearly all cases they would have been
previously brought into solution, the analysis for the Alkali Group metals
should not be made on the solutions obtained from the above fusions.

5. If the residue is thought to consist wholly or partly of tin dioxide,
treat it as follows:

Transfer the mixture to the bottom of a large test tube and add
2 to 5 ml of concentrated HI (sp. gr., 1.70). Sweep the air from
the tube with a stream of C02, and then heat the mixture almost
to boiling. (Red precipitate, usually accompanied by a yellowish
sublimate on the tube just above the solution, presence of tin. Evo-
lution of hydrogen sulfide, presence of insoluble sulfates.) Keep
the mixture almost boiling as long as a reaction appears to be
taking place.

Add to the mixture 25 ml of water, filter out any precipitate,
wash it with 20 ml of water, and treat the solution by P. 1 1. Treat
the residue as directed in the procedure above.

See the discussion above in regard to the reactions taking place. In
Jbhe concentrated acid, less than 0.2 mg of tin will give an easily recog-
nizable red precipitate, which dissolves upon dilution of the solution.

Separation of the Basic Constituents
into Groups

General discussion of the group separations. In devising a scheme
for the systematic detection and separation of a large number of
constituents, either cations or anions, the first step is the considera-
tion of means of separating these constituents into a smaller number
of groups. These groups are composed of constituents having cer-
tain common characteristics, so that a single reagent effects the
detection and separation of the entire group. It is obviously more
efficient to detect and separate these groups first than to attempt to
detect and isolate each constituent separately. Thus entire groups
of constituents are frequently absent and their individual analysis
is avoided. Also, methods of singly separating each element or
constituent from all of those remaining in the solution are often not
available; furthermore, the accumulation in the solution of the
many reagents that would have to be added would finally seriously
interfere with further operations.

Tabular Outline II shows the methods which are used in this
system of analysis for the separation of the basic elements into
groups. It will be seen that there are first precipitated those ele-
ments whose sulfides are insoluble in a solution approximately 0.3
n. in hydrogen ion and from 0.3 to 0.6 n. in chloride ion; these ele-
ments constitute the Hydrogen Sulfide Group. This group is still
too large for individual separations to be made effectively, so it is
further divided by a treatment with sodium disulfide reagent into
two smaller groups, consisting of, first, those elements of the Hydro-
gen Sulfide Group whose sulfides are insoluble in the sodium sulfide
reagent, called the Copper Group, and, second, those elements
whose sulfides dissolve in the sulfide reagent; these constitute the
Tin Group.

By making the filtrate from the Hydrogen Sulfide Group slightly
alkaline with ammonium hydroxide and providing for the presence
of an excess of ammonium sulfide, there is precipitated a second
large group, composed of those elements which form sulfides or
hydroxides which are insoluble in a slightly alkaline solution; this
group is termed the Ammonium Sulfide Group.

This large group is subdivided by dissolving the precipitate in
hydrochloric acid, extracting the iron with ether, and then dividing

193

g
o

S5
O

5; o-a-S

&i3 g *

.tr^4 P r

O ~ _fe * m

2
O

-S

IS"
i ">

2^2,
S

6 a *

CX -

g -3

*d O << ^
C3 fcj.

-+J b/D CQ

*cL *
5 S c o

o IS3 CSJ O
t-.

P-.

^3 o

C *

S.S

W * x s,

5 ^a

S
CO

a-S
o> -^

oL-o
'S >>

"3
S S

^5 . 2

CS ^s

1 CQ qa

-"i

|^S

.S O* bCl3

194

GROUP SEPARATIONS 195

the remaining elements into two groups by a treatment with hydro-
gen sulfide in an almost neutral buffered solution containing a rela-
tively high concentration of oxalate ion. Those elements which are
precipitated as their sulfides from this solution constitute the Zinc
Group; those remaining in the solution constitute the Aluminum
Group.

The filtrate from the Ammonium Sulfide Group precipitation is
treated (after evaporation) with ammonium carbonate, ammonium
hydroxide, and ethyl alcohol. The elements thus precipitated as
their carbonates constitute the Alkaline Earth Group.

Of the basic elements provided for in this system there now remain
in the original solution only sodium and potassium, and these con-
stitute the Alkali Group. Ammonia is tested for in a separate por-
tion of the sample, as it has been introduced with the reagents in
the course of the analysis.

The relative advantages of these separations, as compared with
other methods which are at times used for effecting the same or
similar separations, are discussed later in connection with these
procedures.

Tabular Outline II shows the reagents used in these group sepa-
rations and the elements which constitute each of the groups.

TABULAR OUTLINE III

PRECIPITATION OF THE HYDROGEN SULFIDE GROUP AND SEPARATION OF THE

COPPER AND TIN GROUPS

Solution of the Original Material (30 milli-equivalents of acid; volume 50 ml) :
Add SO mini-equivalents of NH*Cl (no precipitate, absence of Ag).
Heat, saturate with H^S.
Dilute to 100 ml, again saturate with HvS. (P. 11)

Precipitate:

Hydrogen Sulfide Group
Ag 2 S, PbS, Bi,S,, CuS, CdS
Hg, HgS, As,S,, As 2 S 6 , Sb 2 S,, Sb 2 S,, SnS,
SnSi
S

Treat with Na^S-Na^S* reagent. (P. 12)

Filtrate:
Elements of the
Ammonium Sulfide,
Alkaline Earth, and
Alkali Groups (to P. 51)

Residue :
Copper Group sulfides

Ag 2 S, PbS, Bi 2 S,

CuS, CdS
Analyze by P. 21-SO.

Solution:
Tin Group as sulf o-salts

HgS 2 , AsS", SbS*, SnSr
Add H^SOt. (P. 13)

Precipitate:
Tin Group

HgS, As 2 S B ,

Sb 2 S 4 -S, SnS 2 , S.
Analyze by P. 41-49.

Filtrate:
Discard.

196

Precipitation of the Hydrogen Sulfide Group

and
Separation of the Copper and Tin Groups

P. 11. Precipitation of the Hydrogen Sulfide Group

Discussion. A marked departure from conventional systems of
qualitative analysis is found in the elimination of the so-called Silver
Group, usually separated by filtering out the precipitate produced
by the addition of chloride ion to the cold nitric or sulfuric acid solu-
tion, of the original substance. This precipitate would contain,
as chlorides, all of the silver and mercurous mercury, most of the
lead when a large quantity is present, and, less frequently, part of
any bismuth or antimony as the oxy chloride. The instructional
value of this treatment is apparent. The analytical value, which is
in the quick detection and removal of silver, the information gained
as to the state of oxidation of mercury, and some information as to
the amount of lead present, is more than offset by the following facts:
The information regarding the state of oxidation of mercury is ob-
tained only rarely, as nearly all mercurous compounds are so in-
soluble as to require the action of oxidizing agents in their solution;
lead has to be detected and removed from two places in the system
of analysis; and any precipitate produced by the addition of the
chloride causes the complete analysis of the Silver Group to be
carried out, involving the detection and separation of lead, mercury,
and possibly bismuth and antimony, all of which are then again
tested for and separated in the analysis of the Copper and Tin
Groups. Furthermore, the detection and estimation of silver and
mercurous mercury, when they are precipitated together as chlorides,
offer some difficulty. Again, as silver nearly always occurs, both
native and in alloys, together with elements of either the Tin or
the Copper Group, its previous separation usually does not obviate
an analysis of these groups. Finally, an incidental, though usually
conclusive, test for silver is made in this first procedure; if its pres-
ence is indicated, it is easily detected and removed immediately after
dissolving the Copper Group sulfides, at which place it can ad-
vantageously be precipitated and coagulated in a hot solution; if
it is shown to be absent, that procedure can be omitted entirely.

Sulfide Separations. Separations which are dependent upon the
solubility of the sulfides of the elements are more commonly used

197

198

HYDROGEN SULFIDE GROUP

[P. 11

in qualitative analysis than any other method. This is true because
so many of the elements form insoluble sulfides, because these sulfides
vary through such a wide range of solubility, and because the sulfide
ion concentration can be varied and controlled over such a wide
range by controlling the hydrogen ion concentration of the solution.
This range is shown in Table IX, where the hydrogen and sulfide
ion concentrations, and also the concentration of the undissociated
hydrogen sulfide, are shown for various solutions. The data for
this table are taken mostly from the calculations of Knox 1 and are
for solutions at 25C.

TABLE IX
ION CONCENTRATIONS IN SULFIDE SOLUTIONS

Solution

H+

H 2 S

Solutions saturated with H 2 S:
HC1, HN0 8 , 1 f.

2 X 10~ M

HCiHaOa, If.

4 2 X 10~ 8

6 X 10" 1 '

Water

9 X 10-*

1 3 X 10~ 16

Solutions of the following:
NH 4 HS, 1 f .

7 X 10~ 9

1 6 X 10~ 7

(NH<) 2 S, If

5 X 10~ 10

5.5 X 10~*

2.4 X 10" 1

NaHS, If

3.3 X 10- u

3.6 X 10~ 4

3.6 X 10" 6

Na 2 S, If.

1.3 X 10~ 14

1.3 X 10~ 7

The activity coefficient of HC1 is 0.83, and that of HN0 8 is 0.73 in 1 f .
solutions.

The reason for this wide range in the sulfide ion concentrations
is evident when the ionization of hydrogen sulfide as a dibasic acid
is considered, thus:

[Hi [HS

and

[H 2 S]

fH + ] [SI
[HS-]

= 9.1 X 10"

= 1.2 X 10

-15

If these two are combined, there is obtained the expression

[H + ] 2 [si

[H 2 S]

1.1 x 10

-22

(1)
(2)

(3)

which more clearly shows the effect of the hydrogen ion concentra-
tion on the sulfide ion concentration. A solution saturated with
hydrogen sulfide at atmospheric pressure and room temperature is

1 Knox, Trans. Faraday Soc., 4, 47 (1908).

p. 11]

SULFIDE SEPARATIONS

199

approximately 0.1 m. in HjS, so that for a saturated solution there
can be used the simplified expression

[HTIS-] = 1.1 x 10

,-23

(4)

The extreme difference which exists in the calculated solubility
products of sulfides is shown in Table X, where the values of this
constant for a few typical sulfides are given.

Using the value 1 X 10~ 24 for the solubility product of zinc sulfide,
it is calculated that, in order to hold 100 mg of zinc in 100 ml of
solution, the hydrogen ion concentration must be greater than
approximately 0.4 m. Experimentally, more than 100 mg of zinc
can be held for a reasonable length of time in 100 ml of a solution

TABLE X

THE SOLUBILITY PRODUCTS OP CERTAIN SULFIDES
(Arranged in the Order of Their Precipitation as the Hydrogen Ion Is

Decreased)

Sulfide

Solu-
bility
Product

Sulfide

Solu-
bility
Product

HgS
CuS
Ag,S
CdS
PbS
ZnS (a form)

10-"

10-60
1Q-28

10-*

ZnS (/9 form)
CoS
NiS
FeS
MnS (flesh color)
MnS (green)

10-*
10-"

10-*

a The data for this table were taken from the calculations of Kolthoff, /.
Phys. Chem., 35, 2711 (1931). Reference should be made to this article for a
general treatment of the problem of the solubility products of the metallic
sulfides.

0.3 m. in hydrogen ion, and this apparent discrepancy illustrates a
fact which should be emphasized at this point, namely, that it is
extremely hazardous to use these values indiscriminately in solu-
bility calculations and to attempt to predict therefrom the condi-
tions under which separations may be made. This can be done in
certain cases, but often other factors, such as complex ion formation,
rate of precipitation, and adsorption effects, may make such calcu-
lations of little value or even misleading. It is also impossible to
derive reliable solubility products from the solubilities of the sulfides
of arsenic, antimony, and tin, as, owing to their acidic nature and
to their tendency to form complex ions, especially in chloride solu-
tions, even an approximate estimate of the concentration of their

W
ft

O
W

05
OH
W

dfl
to*- 1

rfw

55X

H+, 10-' n
(COsandH(
oxalate

nOO

.T3

r S , S r

^Or?

c ^
oo

MB 2^ m g.

.sa S-^-og

^so

** o

"*^ .2

-o

O2 O2 CO O2
G fl O M-*

a S

CS5 *

O p,

;a*oo

d g
o o

d S

002

oSoi
200

i be present in 50 to 100 ml of the solution with-
illy precipitated. Formulas in parentheses show

,n by the charge on the ions for example, Hg++,
mple ionic species predominate in the solution;
r ten not the case.

.s that of cupric sulfide.

a
.2

5 5

3 o

O3 GQ

T3 13

's s

*
g CQ

g T3

o a

*S o3

r3
a

o
o

aj t^

-S g

** .^i

^*CQ
"

s-i

^1.
si

a> o

Ss~

^ 3

CO S

a> OQ

-ss

tern, and may be reprecipitate4 by boiling. It

i-i ,-r 00 ^.T
03-2 rt OQ
rf T3 <U QQ

^S S. S 2 ."22

09
0>

-) >

13
"3

*-! CO
O QTJ

O c?

*l ^^

QQ
~t

-*3

1-3"

.1 ^

'oS

<u <D ^3 -C3 ja

fl 2 - "
II JS J.I

H3
a?

-u

bO
03

a>
a

"-3

S s

<L>

"S >^

'S S
*R

o fc;

'H
'o

'o
o

o S -2

S^ 1.1 1

^ 1 8 S

oS
--?

"S,

4~
O

*T3

^ I

v S

^C3 oS

^ bO

"GO

t-l ^

^ 1

bO g

a s

ecipitate is indicated, it has been found that 1 mg of
recipitate under the conditions indicated; larger am
ns formed in solution,
ion number of the elements is indicated by Roman r
ten done. The latter is likely to convey the impre
se of hydrolysis to an acidic form, or to complex ion :

03
hO
W
T3

1/3

,a
-*-

*T3
<L>
-*^>
03
O>
I-.
-*

.22

-+^>

ity of cuprous sulfide is approximately of the same o

a>
fe

'-+*

O
en

12
"c^
o5

^
w

.&
S

<f^>

icentration is also frequently obtained by the uset)f
upon the temperature and other conditions, varyin

-4^

525.

*OQ

>->

*-.
OJ
>

c
^o

!s<
*s

upon conditions, these elements may be coprecipit

a
O
O

-*J

Q.
<-i i

o
-i-i

*o
a?
>-.

"T3

a>
N
^>>

QQ -
4->

es to a slight extent. See discussion of P. 12.

^O
QQ

-o

o>
.2

.2
-u

W
x^

o
^Q

-4-3

^5
bO

*o3

> T H 4 )2S 2 and a higher concentration of NH 4 OH, the
ntly used for separating arsenic, antimony, and tin f

ssolves in dilute alkali, owing to the formation of a
in concentrated hydroxide solutions.

.S"-2;*<o s

g > 8 H -2 8

S -| -a 3 ^

jus,'!

^ - OJ + **

o>
oS

a
v

-C3

'Thesolubil

-o

> (

o bo

HH "^

W 55

.2 8.
jq ^
H Q

s
s

-3
a

*a3
H

' CuS dissolv

QQ
CO

g'g-

-5 S
^

> 03

PQ g

3 S

-^*

S-S

O W

SS -3-g

*-+3

c g

201

202 HYDROGEN SULFIDE GROUP [P. 11

various ionic species cannot be made. Rate effects may also be of
great importance; thus there is considerable evidence that the pre-
cipitation of zinc sulfide is an extremely slow process, and that, if
an equilibrium were attained, zinc sulfide would precipitate from
solutions 0.2 to 0.3 m. in hydrogen ion. Probably because of similar
effects, nickel and cobalt sulfides cannot be precipitated from solu-
tions in which the hydrogen ion is 0.1 m.; yet they are dissolved so
slightly, or slowly, in 1 m. hydrochloric acid that this treatment is
often used to separate them from the other elements of the Am-
monium Sulfide Group it is to be noted that the values given for
their solubility products would lead one to predict that they would
precipitate from more strongly acid solutions than does zinc. It is
also to be seen that, by working in alkaline solutions and increasing
the sulfide ion concentration, the insoluble sulfides of certain acidic
and amphoteric elements are dissolved as a result of the formation of
soluble sulfo-salts; this behavior is the basis for the very important
separation of the so-called Copper and Tin Groups.

Because of the importance of the sulfide separations, and because
of the difficulty in treating them from theoretical considerations,
the behavior of the elements of the Hydrogen Sulfide and Ammonium
Sulfide Groups under the conditions attained in certain of the various
sulfide separations is presented in Table XL In this table the
precipitates or distinctive ions which are formed are shown; where
no, or in some cases incomplete, precipitation is obtained under
the conditions indicated, the space is left vacant. The elements
have been arranged in the order of their precipitation as the hydrogen
ion concentration of the solution is decreased.

This table emphasizes a very important principle, namely, that
not only does the hydrogen ion concentration affect the precipitation
of the sulfides, but that the anions present also exert a very decided
effect. Thus it is seen that in 9 n. hydrochloric acid only the sulfides
of arsenic are precipitated, but that from 9 n. sulfuric acid there are
precipitated with these the sulfides of mercury, copper, silver, anti-
mony, and bismuth. This difference could be attributed (1) to
the abnormal activity of hydrochloric acid in solutions of this
concentration (the activity coefficient having been found to be
much greater than unity), (2) to the incomplete ionization of the
hydrosulfate ion, which would correspondingly decrease the hydro-
gen ion activity of the sulfuric acid solution, or (3) to the formation
of what are termed complex or coordination compounds. These
compounds are in general formed because of the tendency of positive
ions to attract negative ions or neutral groups around them. The

P, 11] SULFIDE SEPARATIONS 203

greater the positive charge on the ion, the more stable such com-
pounds are likely to be, and the greater the number (termed the
coordination number) of negative or neutral groups which can be
held. As will be observed later, this tendency is of significance,
especially in concentrated solutions, in determining the reactions
of all of the metallic elements except those of the Alkali Group and
the Alkaline Earth Group. The following are examples of some of
the chloride complex ions of the Hydrogen Sulfide Group elements:
HgCir, CuCir, AgCl 2 ~, BiCir, SnCl 6 ~, SnCir, CdCir, and PbCV".
The other halogen acids tend to form similar compounds; the tend-
ency is much less pronounced with sulfuric acid and practically
inappreciable with nitric or perchloric acid.

In order to decide which of the three effects mentioned above is
the most important, a study of the precipitation of the Hydrogen
Sulfide Group elements from 9 n. perchloric acid was made. This
acid was used because the hydrogen ion activity in the 9 n. solution
has been found to be much greater than that in 9 n. hydrochloric
acid, because its solutions do not oxidize hydrogen sulfide, even at
this concentration (as would nitric acid), and because (as mentioned
above) there is practically no tendency toward complex ion forma-
tion. It was found from these experiments that, in addition to the
sulfides precipitated from 9 n. sulfuric acid, stannic sulfide also
formed in the 9 n. perchloric acid. This shows that the predominant
factor preventing the precipitation of these sulfides from 9 n. hydro-
chloric acid is the formation of the chloride complex ions; this
conclusion is emphasized by the fact that bismuth and stannic
sulfides are not precipitated from even 3 n. hydrochloric acid.

It would seem possible to effect the separation of a group of ele-
ments from the main solution by precipitation with hydrogen sulfide
in a solution made 9 n. in sulfuric or perchloric acid, or even 3 n.
in hydrochloric acid. However, the addition of sulfuric acid would
cause the precipitation of lead and alkaline earth metals as sulfates;
perchloric acid is somethat expensive and might cause partial pre-
cipitation of potassium, and, furthermore, these separations are so
critical upon the adjustment of the acid concentration as to be un-
satisfactory for general use. For these reasons the first group sepa-
ration in this system of analysis is made by the precipitation of
those elements which form sulfides insoluble in a solution approxi-
mately 0.3 n. in hydrogen ion and from 0.3 to 0.6 n. in chloride ion;
these elements will be hereafter designated as the Hydrogen Sulfide
Group.

In order to begin the precipitation of these elements under condi-

204 HYDROGEN SULFIDE GROUP [P. 11

tions which lessen the tendency of certain other elements (especially
zinc, nickel, and cobalt) to be precipitated with them, ammonium
chloride is added to the solution of the original substance, which
should contain 30 milli-equivalents of acid, and it is first heated and
saturated with HaS in a volume of 50 ml, the hydrogen ion concen-
tration being 0.6 n. This, incidentally, causes certain of the sulfide
precipitates to coagulate into a more easily filtered form, and to
precipitate more completely small quantities of arsenic when it is
present in the quinque-positive state. The solution is finally diluted
to 100 ml, cooled, and again saturated with H2S in order to precipi-
tate more completely the sulfides of lead and stannous tin.

The experiments of Noyes and Bray* show that under conditions
similar to those of this procedure 1 mg of any of the metals of the
Hydrogen Sulfide Group will give a perceptible precipitate, and that
even 500 mg of zinc, nickel, and cobalt, the more readily precipi-
tated of the other elements, remain in solution. Noyes and Bray
(loc. cit.) also show that with 500 mg of certain elements of the
Copper Group even 1 mg of zinc, nickel, and cobalt pass into the
filtrate in sufficient quantity to be detected.

The work of Feigl 3 indicates that with large quantities of the
Ammonium Sulfide Group elements coprecipitation may be expected
to occur to an appreciable extent. Feigl found that, when mercuric
sulfide was precipitated from an acid solution containing zinc chlo-
ride, it carried with it 2 to 3 per cent of the zinc, and that, when
stannic sulfide was precipitated from a hydrochloric acid solution
with cobalt present, it contained considerable cobalt sulfide. In
experiments made to study this effect it was found that, upon taking
250 mg of mercury (as mercuric chloride) and 250 mg of zinc (as
zinc nitrate) in 100 ml of a solution containing 30 milli-equivalents
of nitric acid and saturating the cold solution with hydrogen sulfide,
about 50 mg of the zinc were coprecipitated with the mercuric sulfide.
However, when the precipitation was made under the conditions of
the procedure below less than a milligram of zinc was coprecipitated ;
the coprecipitation of cobalt with stannic tin was also shown to be
negligible. Kolthoff and Pearson 4 have shown that under certain
conditions zinc is carried down by copper sulfide; this is apparently
due to the adsorption of zinc on 'the surface of the copper sulfide,

2 Noyes and Bray, J. Am. Chem. Soc., 29, 188 (1907); Noyea and Bray,
Qualitative Analysis for the Rare Elements, Macmillan, 1927, p. 374.

3 FeiRl, Z. anal Chem., 66, 25 (1924).

4 KolthofT and Pearson, Z. phys. Chem., 36, 540 (1932).

P. 11] 8ULFIDE SEPARATIONS 205

with the subsequent formation of zinc sulfide. Since they found
that other active surfaces would cause the same effect, it was assumed
that the solution was supersaturated with respect to zinc sulfide.
Experiments have shown that under the conditions of this procedure
less than a milligram of zinc is carried out when 250 mg each of
copper and of zinc are present. (For a detailed discussion of the
precipitation of the sulfides of zinc, nickel, and cobalt, see the dis-
cussion of P. 61.)

The precipitation of quinque-positive arsenic by hydrogen sulfide
from acid solutions has been extensively studied, 6 and the experi-
mental facts are as follows : A cold solution of an arsenate' in dilute
acid will absorb considerable hydrogen sulfide gas before any precipi-
tation takes place; after some time, a precipitate of arsenious sulfide
is formed. From cold concentrated hydrochloric add arsenic pentasulfide
is precipitated ; if the solution is hot, a mixture of trisulfide and penta-
sulfide is obtained. From the investigations which have been made
it appears that the first reaction taking place in a cold dilute acid
solution is the formation of sulfo-arsenic acids (of the type HsAsOsS),
which then slowly decompose to give sulfur and arsenious acid.
Arsenious acid then rapidly reacts with hydrogen sulfide to give
arsenious sulfide. In a cold concentrated acid solution the basic
ionization of the arsenic into ions such as AsC^ 4 " takes place to a
much larger extent, and these react directly with hydrogen sulfide
to give the pentasulfide. When the solution is heated, there appears
to be a much more rapid decomposition of the sulfo-arsenic acids,
producing sulfur and arsenious sulfide, and there may also be a direct
oxidation of hydrogen sulfide by the arsenic acid, represented by
the equation

H 3 AsO< + H 2 S + 3H+ - As+ + + + S (l) + 4H 2 0.

Thus a mixture of the two sulfides and sulfur is often obtained. It
is seen that the rate of precipitation is favored by increasing the
concentration of the hydrogen ion and of the hydrogen sulfide, and

6 McCay, Am. Chem. /., 9, 174 (1887), 10, 459 (1888); Z. anorg. allgem. Chem.,
29, 36 (1910); /. Am. Chem. Soc., 24, 661 (1902); McCay and Foster, Z. anorg.
allgem. Chem., 41, 452 (1904). Bosek, J. Chem. Soc., 67, 515 (1895), and Brau-
ner, ibid., 87, 527 (1895), have shown that, when hydrogen sulfide is passed
into a solution of quinque-positive antimony, rapid precipitation occurs but
the ratio of SbaSs to Sb2S$ is greater the more rapid the treatment with HjS,
the colder the solution, and, up to 3 to 4 n. HC1, the greater the acidity; only
under very carefully controlled conditions can pure SbsSs be precipitated;
above 95C., largely Sb 2 S 3 is obtained.

206 HYDROGEN SULFIDE GROUP [P. 11

by raising the temperature. Under the conditions of this pro-
cedure, precipitation would be incomplete even after saturating the
solution at 90 to 100C. with H 2 S at atmospheric pressure for an
hour. Rapid precipitation can be obtained by increasing the
hydrochloric acid concentration to 6 n. or greater, but this is not
desired, as the solution is to be next neutralized. Precipitation
can also be obtained by evaporating the solution to dryness, dis-
solving the residue in 6 n. hydrochloric acid, and precipitating the
arsenic from this solution. However, the evaporation of such a
large volume of solution without danger of loss is time-consuming,
and, in addition, considerable amounts of sulfate are formed, caus-
ing possible loss of barium and strontium. The addition of hydri-
odic acid to the solution, which catalyzes the reduction of arsenic
acid by hydrogen sulfide, was tried under various conditions and
did increase the rate of precipitation of the arsenic, but not suffi-
ciently to justify its use. Finally, the effect of heating the solution
with hydrogen sulfide under pressure was investigated, and experi-
ments have shown that all but 1 or 2 of 500 mg of arsenic, when
heated with H 2 S under pressure as directed in the procedure below,
will be precipitated in 15 min. and that precipitation is practically
complete in 30 min.

Procedure 11 : PRECIPITATION OF THE HYDROGEN SULFIDE
GROUP. The solution of the sample, prepared by P. 5 to
P. 8, should be in a 200-ml flask and should contain 30
milli-equivalents of acid in a volume of 50 ml.

Heat the solution of the sample almost to boiling (Note
1), and, if HC1 has not been used in the preparation of the
solution, add to it drop by drop, 1 ml of 3 n. NH 4 C1. (No
precipitate, absence of silver. Note 2.) Add to the solu-
tion 10 ml of 3 n. NH 4 C1, heat it almost to boiling, and
saturate it with H 2 S (Note 3) ; add slowly 50 ml of water,
cool, and again saturate with H 2 S. (Black or colored pre-
cipitate, presence of the Hydrogen Sulfide Group. Note 4.)

If there is no precipitate, immediately pour the solution
into a 500-ml flask (Note 5), boil it until the H 2 S is com-
pletely expelled (Note 1, P. 51), and treat it by P. 51
(Note 6).

If there is a precipitate, let it settle and filter the mixture
with the aid of suction through an asbestos filter (Note 7),
decanting as much as possible of the clear solution (Note 8).

P. 11] PRECIPITATION 207

Wash the precipitate (Notes 9, 11) with three to six 5-ml
portions of 0.12 n. HNOa (Note 12), add these washings
to the filtrate, and treat it as directed in the next para-
graph (Note 13). Wash the precipitate thoroughly with
hot water (or, if it begins to run through the filter, with
0.12 n. HNOa saturated with H 2 S), discarding these wash-
ings (Note 14). Treat the precipitate by P. 12, after
combining with it any further precipitate obtained in the
next part of this procedure.

Heat the filtrate almost to boiling and again saturate
it with H 2 S (Note 15). If no further precipitate forms,
immediately pour the solution into a 500-ml flask (Note 5),
boil it until the H 2 S is completely expelled (Note 1, P. 51),
and treat it by P. 51 (Note 6).

If a precipitate forms, transfer the solution, with the aid
of as little 0.3 n. HN0 3 as possible, to a stout 200-ml large-
mouth pyrex bottle or short-neck, round-bottom flask.
Cool the mixture, saturate it with H 2 S, insert a clean rubber
stopper into the mouth of the flask, and tie it into place
with wire or stout cord. Immerse the mixture in a beaker
of boiling water for 15 min. Remove the flask and allow
the contents to cool before removing the stopper. If a large
precipitate has formed, again saturate the cool mixture
with H 2 S, stopper it, and heat it as before for 10 min. Filter
and wash the precipitate as directed in the third paragraph
of this procedure. Immediately pour the filtrate into a
400-ml flask (Note 5), boil it until the H 2 S is completely
expelled (Note 1, P, 51), and treat it by P. 51.

Notes:

1. When it is directed that a solution be heated "almost to boiling," it
should be brought to 80 to 90C. only. In the above case, if the solution
is heated to boiling, it frequently superheats, so that, when it is shaken (or
when H 2 S is passed into it), such a rapid evolution of steam takes place as to
cause loss by spattering.

When it is directed that a solution be heated "just to boiling," it should
be brought to the boiling point but not heated so as to cause rapid loss of
volume by evaporation.

2. If no precipitate is produced here, the absence of silver is shown, and,
if no oxidizing agents have been used in the preparation of the solution, the
absence of mercurous mercury. A local precipitate which redissolves does
not prove the presence of these elements, as it may be caused by a large
quantity of lead, bismuth, or antimony; the precipitate produced by the
latter substances is likely to be more crystalline or granular, and differ in

208

HYDROGEN 8ULFIDE GROUP

[P. 11

appearance from the more "curdy" silver chloride; even a large quantity
of lead will remain in the hot solution. The formation of a permanent pre-
cipitate upon later adding the additional NHUCl indicates the presence of
these elements; such a precipitate should not be filtered out, as it will be
metathesized by the [28. Only 1 ml of NH4C1 is added at this point in
order to minimize the possibility of precipitating lead chloride; also the
solubility of AgCl is increased by a large excess of chloride ion.

If desired, an approximate estimate of the amount of silver present can
be made by adding the ammonium chloride from a 1-ml measuring pipet
and noting the volume added when the next drop no longer produces a per-
ceptible precipitate. The solution should be vigorously shaken after each
drop, and the precipitate should be allowed to settle somewhat so that the
effect of the next portion can be observed.

If mercurous mercury is known to be absent, the silver may be precipi-
tated and removed at this point by carefully
avoiding an excess of ammonium chloride,
filtering the solution while keeping it hot, and
washing the precipitate with hot 0.3 n. HNOs.
3. When a solution is to be saturated with
H2S, the flask should be fitted with a tight-
fitting two-hole rubber stopper with an inlet
tube in one hole extending about one-half the
distance to the solution and a shorter tube in
the second hole (see Fig. 28). Then pass the
gas into the flask at the pressure of the gen-
erator until the air has been swept out, and
close the second tube by holding a finger over
it or by providing it with a short piece of
rubber tubing carrying a pinch clamp. The
solution should be carefully swirled until it is
saturated; this is indicated when the gas no
longer bubbles through the wash bottle. If
the H^S is generated by the use of iron sulfide
and acid, it may also contain some hydrogen
and volatile hydrocarbons formed from the
iron and carbon usually present in commercial

ferrous sulfide. Because of this, it is advisable occasionally to let the gas
escape for a few seconds so that these gases do not accumulate in the flask.
It is not recommended that the gas be bubbled directly into the solution
and allowed to escaptf; this practice is wasteful of the gas, causes spattering
of the solution, and often causes difficulty in removing the precipitate which
forms on and inside the inlet tube.

4. If only a fine, entirely white precipitate is obtained by the treatment
with H2S, it is due to the oxidation of H2S to S by some oxidizing substance,
such as ferric iron or permanganate, present in the solution. Such a pre-
cipitate should be ignored, and the mixture should be treated as if no pre-
cipitate were obtained.

The color of the various sulfides is as follows: Ag 2 S, black; PbS, black
(the first precipitate produced in hydrochloric acid solutions may be orange-

Fig. 28. Precipitation with
Hydrogen Sulfide.

P. 11] FILTRATION 209

red; this changes to dark brown, and then to black; the first precipitate is
mainly Pb2SCl2, lead sulfodichloride; the subsequent precipitates are mix-
tures of this and lead sulfide in varying proportions); 1^283, brown to black;
Cu2S, black; CuS, black; CdS, yellow; HgS, black (the precipitate first pro-
duced in hydrochloric acid solutions is white mercury sulfodichloride,
HgjSCU; this changes to yellow, brown, and then black; in solutions above
3 to 4 n. in hydrochloric acid, the ictermediate colors, which are mixed
compounds similar to those formed by lead, may persist for a considerable
period of time); As2Ss, yellow; A82Ss, yellow; Sb2Sa, orange-red, changing to
black on long standing or on heating; Sb2Sb, orange-red; SnS, brown; and
SnS2, yellow.

5. When a solution is poured from one vessel to another, or through a
filter, it should be guided with a stirring rod in order to prevent it from
splashing or running down the underside of the lip of the vessel from which
it is poured. The solution adhering to the sides of the original vessel should
be carefully drained off, and then the sides should be washed down with
several successive small portions of water or other suitable wash solution
(in this case two to four portions of 2 or 3 ml of water).

6. The KfoS should be expelled from the solution as soon as possible, as
it tends to become oxidized on standing in air. This increases the amount
of sulfate which may be subsequently formed and increases the possibility
of the loss of barium or strontium in later procedures.

7. Asbestos filters are frequently used in these procedures where it has
been found that the precipitate can be more readily dissolved from the
asbestos than from paper and where successive operations would lead to an
accumulation of pulp, which makes more difficult the subsequent operations.
Asbestos filters are also more readily used with suction. The general discus-
sion of the filtering of precipitates accompanying P. XVIII, "The Gravi-
metric Standardization of a Hydrochloric Acid Solution/' should be con-
sulted.

An asbestos filter for use here is made by placing in a medium-size glass
funnel (or a funnel such as is used in quantitative determinations for holding
perforated porcelain crucibles) a perforated porcelain plate (the so-called
Witte plate) about 2 cm in diameter, and preferably having a beveled edge,
and then pouring through this a water suspension of acid-washed asbestos
fibers until a layer 1 to 2 mm thick is formed. The layer should be built up
slightly thicker around the edges of the plate. The mat should then be
washed until the wash water is free from fibers in order that a clear filtrate
may be obtained. The asbestos should be prepared for use according to the
directions in the Appendix.

Suction flasks (heavy conical flasks with thick walls and. a side neck)
of the conventional type may be used with the asbestos filters. As these
are often available only in large sizes and as solutions cannot be safely heated
in them, it is at times convenient to filter directly into the conical flask in
which the filtrate is next to be treated. This can be done by selecting a
stout flask and fitting it with a two-hole rubber stopper, carrying the funnel
in one hold and an outlet tube connected to the suction in the other. Flex-
ible rubber tubing and not the more rigid vacuum or pressure tubing should
be usfed, as the latter will often tend to overturn the lighter flask; also, the

210 HYDROGEN SULFIDE GROUP [P. 11

lighter tubing, even if collapsed, will carry as large a volume of air as can be
safely used. As an additional precaution the funnel should be supported in
a funnel stand. When a water aspirator is used for the suction, a water trap
should be provided consisting of a flask with a two-hole stopper fitted with
an outlet tube to the aspirator and an inlet tube from the filter flask outlet;
the possibility of tap water being sucked into the filter flask is thus eliminated.

The first portion of any solution poured upon an asbestos filter should
be added with great care in order not to stir up the mat, and a slight suction
is generally advisable at this time as it helps to keep the filter in place. When
flocculent or gelatinous precipitates are being filtered, a very thin layer of
asbestos should be used and a second plate placed on top of this mat to keep
it in place. Suction should be applied only as it is needed to keep the solu-
tion moving through the filter, for otherwise the first part of the precipitate
will be tightly packed against the filter and may completely clog it.

It is possible to use a paper filter for this filtration; however, owing to the
accumulation of paper pulp which occurs in the operations of P. 12 and P.
21, filtrations and washings are made difficult. Moderate suction may be
used with a paper filter if the lower portion of the filter is supported in the
funnel by a smaller (5-cm) filter of hardened paper, a perforated platinum
cone, or a small cloth support. If a paper filter is employed, see Note 1,
P. XVIII C, and the discussion of filtering media, p. 133, for suggestions in
regard to its use.

8. In general, precipitates should be allowed to settle before being
filtered, and then as much as possible of the clear solution should be de-
canted through the filter. In this manner, it is often possible to filter a
mixture completely and then to wash the precipitate without removing it
from the flask. This is termed filtering by decantation and should be prac-
ticed whenever the nature of the precipitate permits. In this case, if the
precipitate settles rapidly, it need not be brought upon the filter but can be
washed in the flask, the wash solution can be decanted through the filter,
and the precipitate can then be treated directly in the flask with the sodium
sulfide reagent with which it is to be treated in P. 12. Any precipitate
carried onto the filter can be treated there or transferred to the flask with
the aid of the sodium sulfide reagent.

9. Many precipitates have a tendency to become colloidal and to run
through the filter when they are being washed, so that it is advisable to
remove the filtrate before beginning to wash the precipitate, especially if the
volume is large, and to collect the washings separately. Then, if a break
occurs in the filter or the precipitate becomes colloidal, a much smaller
volume of solution will have to be refiltered.

10. When washing a precipitate, the original solution and then each suc-
cessive portion of wash water should be allowed to drain from it before the
next portion is added. It can be shown that the amount of wash solution
required will be determined largely by the volume of solution left with the
precipitate and filter each time. However, a precipitate should never be
allowed to dry or cake before or during the washing. It can also be shown
that it is more effective to wash with several small portions of wash water
than with one portion of the same total volume. See p. 137 for a general
discussion of the washing of precipitates.

P. 12] COPPER AND TIN GROUP SEPARATION 211

11. When limiting amounts of wash solution are specified, the volume
used should be adjusted to the size of the precipitate to be washed. As a
1-g sample (0.5 g for alloys) is used, it is assumed that not more than 500 mg
of any one constituent will be found.

12. It is not essential that this wash solution be accurately 0.12 n. in
HNOa; 0.1 n. would be satisfactory. However, as the more common dilute
acids and bases are usually prepared to be 6 n., it is seen that 0.12 n. acid
can be conveniently made by diluting 2 ml to 100 ml. Likewise, 0.6 n.
wash solutions are often specified where 0.5 n. would be equally satisfactory
but could not be so conveniently made from a 6 n. solution.

13. Experiments have shown that six washings with 5 ml of wash solution
should leave less than 0.1 mg of elements of other groups in even a large
sulfide precipitate. Thus, with 500 mg of iron present, less than 0.1 mg of
it was left in a precipitate containing 500 mg of arsenic as As2Ss. Accord-
ingly, in order not to dilute the filtrate unduly, only the volume of the solu-
tion used to wash the precipitate is added to it. Any additional wash solu-
tion, which serves to remove traces of elements which might interfere in the
group analysis, is accordingly discarded.

The precipitate is washed with dilute acid to prevent possible precipita-
tion of any Ammonium Sulfide Group elements and to diminish the tendency
of the precipitate to become colloidal ; it is also advantageous to saturate the
wash solution with H^S.

14. The washing of a precipitate should in most cases be continued until
portions of the wash solution give no test for the most easily detected of the
substances originally present. In this case, if the washing is done with
water, the washings as they come from the funnel can be tested for hydrogen
ion with litmus paper.

15. The second treatment with H^S is to provide for detecting any
quinque-positive arsenic not already precipitated. See the discussion.

P. 12. Separation of the Copper Group from the Tin Group

Discussion. The Hydrogen Sulfide Group precipitate contains
such a large number of elements that it is advisable to separate it
into two smaller groups, rather than to make a large number of
individual separations. In considering the possibilities of making
such a separation, an inspection of Table XII shows that certain
of the elements comprising this group are acidic, or at least distinctly
amphoteric in nature and therefore tend to dissolve in alkaline
solutions with the formation of acidic ions. Thus arsenic and
antimony, especially when in their higher oxidation states, are pre-
dominantly acidic in character, and tin and lead are amphoteric,
tending to form soluble acidic ions in alkaline solutions. The other
elements of the group namely, copper, mercury, cadmium, bis-
muth, and silver form ions which are successively more basic in
their reactions and precipitate relatively insoluble hydroxides in

TABLE XII

THE BEHAVIOR OF THE HYDROGEN SULFIDE GROUP ELEMENTS IN SOLUTIONS
OF VARIOUS HYDROXYL (AND HYDROGEN) ION CONCENTRATIONS

Element
and Oxida-
tion State

(OH-),2to4m.

(H+), (OH-), approx.
10-' m.

(H + ),10- i tolO-m.

Agi

Ag,O

(Ag+)<

(Ag + )

(Pb(OH)D

PbCOHV

(Pb+ + )

Hi"'

Bi(OH),

Bi(OH),'

BiO(OH)'

Cul

Cu(OH), CujO

(Cu 1 )"*

(Cu')*

Cu(OH) 2 , CuO"

(Cu ++ )"'

(Cu++)

Cd(OH) 2

(Cd++y

(Cd++)

Hgl

Hg,0

(H gl ++)/

(Hg,++)

Hg"

HgO

( Hg ++)

(Hg ++ )

As" 1

(HAsO,-)

(HaAsOa)*

(H 3 A 8 0,)

AsV

(AsOD

(HAsOr)

(HjAaOr)

Sb" 1

(HSbOr)

Sb s O,

SbV

(SbO<-)'

Sb 2 Os- (H 2 O),

Sb,0,-(H 2 0),

Sn"

(Sn(OH)r)

Sn(OH) 2

Sn'V

(Sn(OH)e-)

SnOr(HjO).,

SnO 2 -(H 2 O) x

The composition of the so-called hydroxides varies with the conditions of
formation, such as the concentration of the salt and of the hydroxyl ion, the
temperature of the solution, the anions present (basic salts being formed),
and the time for which they are allowed to stand in solution or in the air.
They often separate from the solution with considerable amounts of adsorbed
water, and can be dried without their temperature-vapor pressure curves
showing the abrupt breaks characteristic of definite hydrate formation. Be-
cause of this uncertainty, the term hydrous oxides, or, where definite hydrates
obviously exist, hydrous hydrated oxides, has been used (Weiser, The Hydrous
Oxides, McGraw-Hill, 1926); however, the term hydroxide will be retained in
this book, but with the limitation that an exact formula is not implied.

The ions which are formed when the amphoteric elements are dissolved by
an excess of hydroxyl ion are usually assigned dehydrated structures (such as
PbOa"" and SnO 3 ~), not because it was thought that these were necessarily the
predominating species, but because the correct formula was not known, and
frequently an equilibrium between two or more forms may exist. However,
there now seems to be justification for assigning to these ions more specific
formulas, not only because of evidence from work on the crystal structures of
their compounds and from the principles applying to coordination-compound
formation,* but also because the analogy with other coordination compounds
is more clearly demonstrated and, in addition, in many cases a more adequate
explanation of their chemical behavior is thereby obtained. f

a Bi(OH) 3 dissolves to a very slight extent in concentrated (6 f. or more)
hydroxide solution.

5 Cu(OH) 2 dissolves slightly in concentrated hydroxide solutions, forming
the blue CuO 2 " ion.

c Complex ammino ions are formed when ammonia is present.

d Pb(OH) 2 dissolves to an appreciable extent.

* For a discussion of the formulas of the antimonates and the anions of a
number of other elements see Pauling, J. Am. Chem. Soc., 56, 1895 (1933).

t Hammett, Solutions of Electrolytes, McGraw-Hill, 1936, pp. 105-113, dis-
cusses the reactions of certain amphoteric hydroxides.

212

P. 12] ALKALINE SULFIDE SEPARATIONS 213

* Cu(OH)j may be incompletely precipitated, depending upon temperature,
concentration, and anions present.

* Ammonia causes a mixed precipitate composed of finely divided metallic
mercury and HgNHjCJ or HgO-HgNH^'NOt, depending upon the anion
present.

9 Ammonia in chloride solutions forms HgNHjCl ; in nitrate solutions it
forms HgO-HgNHjNOs.

* AsjOf is only moderately soluble in water.

* This is usually a mixture of basic salts.

' In less alkaline solutions the, ion Sb(OH) 6 ~ probably predominates.

* Cuprous ion is unstable in aqueous solution, tending to give copper and
cupric ion, unless stabilized by the formation of a complex ion.

such solutions. Therefore it is seen to be theoretically possible to
treat a solution containing all of these elements with sodium or
potassium hydroxide in large excess, and thus to precipitate the last
five named and to dissolve the first four. Practically, owing to the
solubility of certain of the precipitates in strongly alkaline solutions
and to the tendency of certain of the soluble acidic elements to be
carried out with those precipitated, such a separation is not gen-
erally applicable.

Separations by Alkaline Sulfide Solutions. If, instead of the
tendency of these acidic elements to form the oxygen acids, advan-
tage is taken of their tendency to form the analogous sulfur acids,
a very similar effect is obtained; thus an inspection of Table XI
shows that, when their sulfides are treated with ammonium sulfide,
arsenic, antimony, and tin form sulfo-salts analogous to their oxy-
salts, and that with a higher concentration of sulfide (provided by
using a sodium sulfide solution containing sodium hydroxide to
repress tKe hydrolysis of the sulfide) mercury also forms a sulfo-
salt. 6 It has been found experimentally that the use of these sulfide
reagents gives group separations of these elements which are more
quantitative than those obtained by the use of hydroxide solutions.
The reagent most commonly used for this purpose is one containing
ammonium sulfide, ammonium disulfide, and ammonium hydroxide.
The disulfide is provided in order to oxidize stannous tin, as stan-
nous sulfide is but slightly soluble in ammonium monosulfide, and
as, in general, the acidic nature of the element becomes more pro-
nounced the higher the oxidation state. Although arsenious sul-
fide, As2S 3 , and antimonous sulfide, Sb 2 S 3 , are soluble in a mono-

8 Knox, Trans. Far. Soc. t 4, 29 (1908), from a study of the solubility of HgS
in sodium sulfide and disulfide solutions, gives for the dissociation constant.

[HgS,-]
the value 1.96 X 10-.

214

HYDROGEN SULFIDE GROUP

[P. 12

sulfide reagent, they are oxidized by disulfide ion and are present
in such solutions in their higher oxidation states as the ions AsS* m
and SbS^, respectively. In this system a sodium sulfide reagent,
which is approximately 3 n. in sodium sulfide, 1 n. in sodium disulfide,
and 1 n. in sodium hydroxide, is used in place of the more commonly
used ammonium sulfide reagent 7 for separating the Tin Group from
the Copper Group sulfides.

An experimental study of the behavior of the elements of the two
groups when their sulfide precipitates are treated with the sodium
sulfide reagent as directed in the procedure below has been made, and
the results of these experiments are shown in Tables XIII and XIV. 8

TABLE XIII

EXTRACTION OF SMALL AMOUNTS OF THE TIN GROUP SULFIDES FROM THE
COPPER GROUP SULFIDES BY THE SODIUM SULFIDE REAGENT

Copper

Group

Element

Tin Group Elements
Amount Dissolved by Na2S Reagent

500 mg

1 mg Taken

2mg
Taken

Taken

Hg"

As 111

Sb m

Sn 11

Sn lv

Hg"

Ag 1

0.5

1.5

Pb 11

tr.

Bi m

0.5

Cu"

tr.

1.5

Cd 11

0.5

It is seen that, when the sodium sulfide reagent is used as directed,
1 mg of any element of the Tin Group, except mercury, will be
extracted in sufficient quantity to give a satisfactory test even when
present with 500 mg of any element of the Copper Group. When
1 mg of mercury is present with large amounts of lead or copper,
its detection is uncertain; with a large amount of cadmium it fails.
When 2 mg of mercury is present, enough is always extracted to
insure its detection. When larger amounts of the Tin Group ele-
ments (250 mg) are extracted, considerable mercury and antimony
are left if a large amount of cadmium is present. If desired, these

7 Both the ammonium and sodium sulfide reagents are more complex in
their composition than is indicated. On dissolving sulfur in alkaline sulfide
solutions, thiosulfate, S2O 8 -, trisulfide, 83", and even higher sulfides are formed.
These latter compounds appear to be present in appreciable amounts even if
the amount of sulfur added is inadequate to convert the sulfide to disulfide;
for this reason these reagents are frequently termed "polysulfide" solutions.

8 Unpublished experiments by R. C. Aussieker.

P. 121 COPPER AND TIN GROUP SEPARATION

215

elements can be recovered from the residue resulting when the
Copper Group sulfides are dissolved in nitric acid and do not inter-
fere with the analysis of that group. Of the elements of the Copper
Group 20 ml of the sodium sulfide reagent dissolves about 1 mg of
bismuth and only a trace of copper. Occasionally, small amounts
of the Copper Group sulfidos may pass into the sodium sulfide
solution in colloidal form.

In comparison, the ammonium sulfide reagent, which is commonly
used for this separation, leaves mercuric sulfide with the Copper
Group, where its detection and separation are much more difficult,
especially when silver is included in the Copper Group. Further-
more, the data of Noyes and Bray 9 on the behavior of the Copper

TABLE XIV

EXTRACTION OF LARGE AMOUNTS OF THE TIN GROUP SULFIDES FROM THE
COPPER GROUP SULFIDES BY THE SODIUM SULFIDE REAGENT

Copper Group
Element
250 mg Taken

Tin Group Elements
250 mg Taken
Amount Found with the Residue

Hg 11

As 111

Sb m

Sn 11

Ag'
Pb 11

0.5-1
1

tr.
2-3

1-3
5

tr.
1

Bi IU

1-2

tr.

Cu 11

tr.

Cd n

tr.

and Tin Group sulfides with ammonium mono- and polysulfide,
show that, when tin is present with large amounts of any of the
Copper Group elements, it may remain undetected in amounts as
large as 3 to 5 mg, and even 15 ing of stannous tin may be wholly
left with 500 mg of cadmium. Finally, 10 ml of the ammonium
sulfide reagent dissolves 5 to 10 mg of copper, 1 mg of mercury,
and 0.1 to 0.3 mg of cadmium.

Procedure 12: SEPARATION OF THE COPPER GROUP FROM
THE TIN GROUP. Combine the H 2 S precipitate on the filter
(Note 1 ) with that remaining in the flask and treat it there
with 5 to 20 ml of sodium sulfide reagent (Note 4). Heat
the mixture to 60 to 80C. (Note 2) for 5 min., stirring it
and disintegrating any residue (Note 3). Add to the mix-
ture 5 to 20 ml of hot 0.6 n. NaOH and filter it at once
through an asbestos filter, catching the filtrate in a 200-ml
flask. Wash the precipitate with two to four 10-ml portions

Noyes and Bray, /. Am. Chem. Soc., 29, 192 (1907).

216 HYDROGEN SULFIDE GROUP [P. 12

of a solution made by diluting 1 volume of the sodium sulfide
reagent with 3 volumes of water and heating it to 60 to
80C. Add these washings to the original filtrate (Notes
9, 10, P. 1 1). Treat the filtrate by P. 13. Wash the residue
thoroughly with hot water (or, if it begins to run through
the filter, with a hot solution made by diluting 1 volume of
the sodium sulfide reagent with 9 volumes of water) and
treat it by P. 21.

Notes:

1. A precipitate on an asbestos filter can be transferred by tilting the
funnel containing it and then scraping away the precipitate with any part of
the filter to which it adheres, or by inverting the funnel in a casserole and
pushing out the entire filter and precipitate with a thin glass rod inserted
through the stem of the funnel. Any precipitate adhering to the funnel can
then be loosened with a stirring rod or policeman and washed with the solu-
tion with which the precipitate is next to be treated in this case sodium
sulfide reagent.

If only a small precipitate has been obtained, or if only a small portion of
a larger precipitate has been carried onto the filter, it may be treated directly
by pouring the sodium sulfide reagent repeatedly through it. After treating
the precipitate in the flask with the reagent, the same filter may be used
for filtering the mixture.

2. A flask or casserole at 60 to 80C. can be held in the hand for only a
second or two without discomfort. The mixture should not be boiled.

3. When a precipitate is treated with a solution which is intended to dis-
solve all or part of it, any residue or solid material left after adding the splu-
tion should be carefully broken up and disintegrated with a stirring rod;
one which has been heated and flattened at one end is advantageous for this
purpose.

4. When limiting volumes of a reagent are specified, the amount used
should be adjusted to the size of the precipitate or residue to be treated. It
is advantageous to use as little of the sulfide reagent as possible because of
the large amount of sulfur which is precipitated when the solution is made
acid in P. 13.

P. 13. Precipitation of the Tin Group

Discussion. Upon acidifying the sodium sulfide solution, which
contains the Tin Group elements as sulfo-salts, the sulfide present is
converted into US and the polysulfides are converted into H 2 S and
S, and at the same time the Tin Group elements are reprecipitated
as sulfides. The reactions occurring may be represented as follows:

S + 2H = H2S( g )
ST + 2H + = HA,, + S (i)
2AsS 4 " + 6H + . = AsjSsf,) + 3H z S (g) .

P. 13] PRECIPITATION OF TIN GROUP 217

Sulfuric, rather than hydrochloric acid is used, because, first, these
sulfides are more soluble in high concentrations of chloride ion (see
the discussion of P. 11) and, second, by adding an excess of the sul-
furic acid to the sulfate already formed by the neutralization, a
"buffered" sulfate-hydrosulfate solution (see the discussion of P.
XII and P. 61) is obtained in which the hydrogen ion concentration
is sufficiently high to cause complete decomposition of the sulfo-
salts without danger of dissolving the sulfides. Arsenic, tin, and
mercury are precipitated as the higher sulfides, As2Ss, SnS2, and
HgS; with antimony there is evidence that a precipitate of the tetra-
sulfide, Sb 2 S4, and sulfur results. 10

It has been found that heating the mixture while keeping it satu-
rated with H 2 S insures the complete and more rapid precipitation of
the sulfides and their coagulation into a more rapidly filtered form.

In order to test the detection of the Tin Group sulfides when the
sodium sulfide solution is acidified, 1 mg of each of the Tin Group
elements was dissolved in 20 ml of sodium sulfide reagent, 20 ml of
0.6 n. sodium hydroxide and 20 ml of a solution made by diluting 1
volume of sodium sulfide reagent with 3 volumes of water were added,
the solution was then slowly acidified with hydrochloric acid, and
2 ml more were added. The precipitate was then examined and
compared with that produced by acidifying a similar sodium sulfide
solution containing no Tin Group element. The mixture was then
heated to 80 to 90C. The blank solution, when first acidified,
produced a finely divided pure white precipitate, which upon coagu-
lation by heating acquired a very slight yellowish color; 1 mg of
mercury produced a distinct gray precipitate, which coagulated into
dark flocculent particles; 1 mg of arsenic first produced a precipitate
only a little different from that of the blank solution, which upon
coagulation became deeper yellow and more flocculent in nature;
2 mg of arsenic produced an easily recognized yellow flocculent pre-
cipitate; 1 mg of antimony produced an orange precipitate; 1 mg
of tin produced a buff precipitate. A comparison with a blank
would be necessary only in determining the presence of as little as
1 mg of arsenic.

Procedure 13: PRECIPITATION OF THE TIN GROUP.
Slowly acidify the filtrate from P. 12 with 18 n. H 2 S0 4
(Note 1), note the volume of the acid used, and then add
one-half that volume more. Shake the mixture vigorously

10 Currie, /. Phys. Chem., 30, 205 (1926).

218 HYDROGEN SULFIDE GROUP [P. 13

and heat it to 80 to 90C. If the precipitate is white
(Note 2), discard the mixture.

If the precipitate is not entirely white, heat the mixture
nearly to boiling, saturate it with H 2 S, and allow the precipi-
tate to settle. If the precipitate does not settle rapidly,
add to the mixture 5 g of solid (NH^SO^ shake it vigor-
ously, and again heat it (Note 3). Filter the mixture
through an asbestos filter (Note 4). It is not necessary to
wash the precipitate; drain it as dry as possible (Note 5).
Treat the precipitate by P. 41 (Note 6).

Notes:

1. Unless otherwise stated, when it is directed to acidify, neutralize, or
make a solution alkaline, it is intended that litmus, in the form of test
papers, be used as the indicator. Usually the necessary amount of acid or
alkali can be estimated and most of this volume can be added at once, care
being taken not to cause any loss through spattering of the solution in cases
such as the above, where BUS or C02 are liberated; then the remainder is
added in gradually decreasing portions, the solution being shaken after each
addition, and a drop of it held on a stirring rod being tested with a corner
of a strip of litmus paper. The amount of solution used in making these
tests is generally insignificant; however, if the volume of the solution being
used is small, the litmus paper can be finally washed with a few drops of
water. It is bad practice to drop strips of litmus paper into the solution being
tested. The paper usually absorbs the strongly alkaline or acid solution and
then responds slowly to changes in the solution when the neutral point is
approached, the litmus is leached from the paper and may impart an unde-
sirable color to the solution, and the paper may so disintegrate as to make it
difficult to detect small precipitates which may be present.

Aid in making the above neutralization can be gained by observing the
first appearance of the white, very finely divided, precipitate of sulfur
the precipitate first appearing being yellow; complete neutralization of the
base present is also indicated when, after vigorous shaking of the mixture,
bubbles of H^S are no longer formed with further addition of acid. These
indications are of value in this case, as the H2S set free tends to cause a slight
acid reaction of the litmus paper before all of the Na2 has been decomposed.

2. A finely divided, uniformly white precipitate is always produced on
acidifying this solution, and is caused by the liberation of sulfur from the
disulfide present in the sodium sulfide reagent. The absence of any appre-
ciable dark, yellow, orange, or brown precipitate with this sulfur indicates
the absence of as much as 1 mg of any of the Tin Group elements. One
not familiar with the type of precipitate produced by the sodium sulfide
reagent should obtain a basis of comparison by acidifying the same volume
of sodium sulfide reagent and NaOH as was used in the procedure, and
comparing the color and nature of this precipitate with that obtained on
acidifying the filtrate.

P. 13] PRECIPITATION OF TIN GROUP 219

The precipitate produced with even 1 mg of mercury is unmistakably
darkened, coagulating into dark grayish particles, that with antimony is
orange, and that with tin is buff-colored. The precipitate produced by 1
mg of arsenic, while not as definite as those produced by the other three
elements, has a more decided yellow tinge and coagulates into more floccu-
lent particles than does the sulfur precipitate alone. Two mg of arsenic
produces an easily recognized, floccuient, yellowish precipitate. See the
discussion.

The precipitate may be darkened by the presence of Copper Group ele-
ments (see the discussion of P. 12). If this seems probable, or if there is
doubt as to the presence of Tin Group elements, or if only a small precipitate
is obtained and a quick identification of the elements is all that is necessary,
the precipitate, after being filtered, should be treated as directed in Note 6.

3. If tin is present, the precipitate may not settle and the filtration may
be very slow, as stannic sulfide tends to be colloidal. The addition of a
large amount of an electrolyte, such as ammonium sulfate, usually assists
in coagulating the precipitate.

4. If the precipitate is colloidal in nature, indicating the presence of tin,
a paper filter may be used to advantage. See Note 1, P. XVIII C, and pp.
133-134, as to the characteristics, selection, and use of paper filters.

In general, the filtering and washing of precipitates will proceed more
rapidly and efficiently with hot solutions; and advantage should be taken
of this fact unless the precipitate is appreciably soluble or some undesirable
reaction is caused by the higher temperature.

5. As the solution in this case contains nothing which will interfere with
the subsequent treatment of the precipitate, the usual washing is not neces-
sary. As the precipitate can be more easily transferred to the beaker to be
used in P. 41, from the flask than from the filter, it should be filtered by
decantation as much as possible.

6. If there is doubt as to the presence of any Tin Group elements, the
precipitate can be rapidly analyzed by the optional method given in P. 41.

3 O

-J O

3 n

D H

< S

S
2

TJ
U?

O
I

Sodium Sulfide Trea
u**/i //ATO,. (P. 21)

+"*
n

|-S
at
t"^

..oo

-
w-g

*"?
8l

V o

CQ v
o v

HH .

z^ S

-s
3 1

fl - C3

^ s

li-1
lll!

'Suo o is Jb

55 | I
jj(^ ^ S

.So

-S

+ -*.

'S.O

034* -

'42 CO

220

The Analysis of the Copper Group

P. 21. Solution of the Copper Group Sulfides in Nitric Acid

Discussion. The first operation in the analysis of the Copper
Group is to bring the sulfides into solution. Nitric acid is employed,
as both hydrochloric and sulfuric acids cause the precipitation of
certain elements of the group. Nitric acid is also a more effective
solvent for sulfides; when hot and more than about 2 n., it acts as
an oxidizing agent on sulfides and hydrogen sulfide, forming sulfur
and smaller amounts of sulfate. In addition, when nitric acid is
used, any mercuric sulfide not extracted by the sodium sulfide
reagent will be left after the Copper Group sulfides are dissolved,
and any antimony sulfide will be converted into antimonic oxide.

When ammonium sulfide is used for separating the Copper and
Tin Groups, mercury remains with the Copper Group and such a
treatment with nitric acid is used to separate it from the other ele-
ments; it is also sometimes used for the quantitative separation of
mercury from these elements. 1 As is seen in Table XIV, mercury
is likely to be found here in small amounts when large amounts of
cadmium are present.

Dilute nitric acid is first used and the solution is gradually heated
in order to expel as much as possible of the hydrogen sulfide without
oxidation. If the sulfides were treated directly with concentrated
nitric acid, or quickly heated, the hydrogen sulfide would be largely
oxidized to sulfur, which would tend to enclose the remaining pre-
cipitate and prevent it from being dissolved. Solution could be
obtained then only by using concentrated nitric acid, which would
oxidize the sulfur to sulfate with possible precipitation of lead sul-
fate, even though, owing to the solubility of this salt in hot nitric
acid, considerable amounts remain in solution. Upon treating 500
mg of silver (as freshly precipitated silver sulfide) by the procedure
below, less than 2 mg of silver were found in the dark residue which
remained. The nitric acid residue, which is very largely sulfur, is
washed with dilute nitric acid in order to avoid the tendency of
bismuth salts to hydrolyze and precipitate as oxy-salts when diluted
with water. It should be noted that even if the residue is white it

1 See Hillebrand and Lundell, Inorganic Analysis, p. 171, for a discussion
of this method.

221

222 ANALYSIS OF THE COPPER GROUP [P. 21

does not prove the absence of mercury; mercuric sulfide, on treat-
ment with nitric acid, is often transformed into white compounds,
such as 2HgS*Hg(N03)2, which are difficultly soluble.

Procedure 21: SOLUTION OF THE COPPER GROUP SUL-
FIDES. Treat the residue from the sodium sulfide treat-
ment (P. 12) with 10 to 20 ml of chloride-free 3 n. HN0 3
(Note 1), and stir the mixture in such a way as to break up
any lumps (Note 2). Slowly heat the mixture and boil it
gently for 5 min. If there is a dark residue which seems
to be slowly dissolving, add to the mixture 2 to 5 ml of
chloride-free 16 n. HN0 3 and again boil it until no more of
the residue appears to dissolve (Note 3). Prepare a paper
filter (Note 4, P. 13) and filter out the residue while keeping
the solution hot. Collect the filtrate in a 200-ml flask.
Wash the residue with 5 to 10 (Notes 10, 11, P. 11) 5-ml
portions of hot 1.2 n. HN0 3 (Note 4), adding these washings
to the filtrate. Treat the filtrate by P. 22. Treat the
residue as directed in Note 5 or discard it.

Notes:

1. If, in P. 12, the mixture has been filtered by decantation, the residue
should be retained in the same flask and treated there with the acid, combin-
ing with it any precipitate on the filter.

The HNOs used should be tested in order to be sure that it is free from
chloride; this would cause the partial or complete loss of any silver present
by precipitating it as silver chloride, which would be filtered out with the
residue and asbestos.

2. After adding the HN0 3 , the mixture should be stirred and the residue
disintegrated so that any solvent action taking place in the cold can proceed.
The mixture should then be heated only as much as is necessary to keep the
solution of the precipitate in progress. If there is a large sulfide precipitate,
and especially if silver is present, it may dissolve only slowly in the dilute
HN0 3 ; in this case the additional HN0 3 is added, and the solution is again
heated as long as the precipitate seems to be dissolving. If the sulfide pre-
cipitate is large or if it is heated too rapidly, a spongy residue of sulfur,
darkened by enclosed traces of the precipitate, will frequently form. This
residue can be readily distinguished from the undissolved sulfides, and, as
the amount of them which it contains is usually negligible, it should be
filtered out, pressed free of solution with a stirring rod, and discarded.

3. During the boiling of the solution the flask should be covered with a
watch glass to prevent loss of acid by evaporation. If, in spite of this
precaution, considerable volume is lost, it should be replaced with 3 n.
HN0 3 ; the volume of the mixture when filtered should be 10 to 15 ml.

4. The precipitate should be gently stirred up with each portion of wash

P. 22] DETECTION OF SILVER 223

water, added with either a dropper or a wash bottle, depending upon the size
04 the precipitate and paper, and then the remainder of that portion should
be applied dropwise or in a very fine stream around the upper edge of the
filter, more being used on the triple-folded section of the paper (Note 10,
P. 11).

5. This residue will usually consist of only sulfur and asbestos. How-
ever, under certain conditions some mercury or antimony (see the discus-
sion) or some silver (Note 1) may be present. The silver may be detected
by pouring 10 ml of warm NH 4 OH repeatedly through the residue, adding
to the clear solution 1 ml of 3 n. NH 4 C1, and acidifying it with HN0 3 ; a
white precipitate, dissolving when the solution is again made alkaline,
proves the presence of silver.

If desired, the mercury and antimony can be recovered as directed in the
procedure for the rapid analysis of the Tin Group (P. 41A), or the mer-
cury may be detected by treating the residue as directed in Note 1, P. 45.

P. 22. Detection and Precipitation of Silver

Discussion. Silver is separated from lead and the other Copper
Group elements at this point by precipitation as the chloride. This
separation depends upon the fact that silver chloride can be quanti-
tatively precipitated from a boiling nitric acid solution; lead chloride
is not precipitated; because an excess of chloride can be avoided,
ard because lead chloride is much more soluble in hot than in cold
water. Owing to the moderately high concentration of nitric acid
and to the small excess of chloride present, even 500 mg of bismuth
do not cause a precipitate (of BiOCl), although hydrolysis, and,
therefore, the tendency of bismuth to precipitate, is more pro-
nounced in the hot solution. The precipitation of silver chloride
and the factors affecting the solubility of the precipitate have been
extensively discussed in P. XVIII, "The Gravimetric Standardiza-
tion of a Hydrochloric Acid Solution/'

Experiments have shown that the amount of silver present can be
quite precisely estimated, even in the presence of large amounts of
bismuth and lead, by titrating with a dilute standard solution of
sodium or ammonium chloride until the formation of any precipi-
tate ceases. This principle is the basis of a very precise volumetric
method for the quantitative determination of silver.' 2 While this
method is precise, it is time-consuming and somewhat uncertain
when carried out for the first time. In addition, if a large quantity
of silver were titrated, a considerable volume of solution would be
added, requiring the addition of more acid to prevent the precipita-

2 See Treadwell-Hall, Analytical Chemistry, Vol. II, Quantitative, 7th Ed.,
p. 600, for a discussion of this method, known as the Gay-Lussac titration.

224

ANALYSIS OF THE COPPER GROUP

[P. 22

tion of oxy-salts of bismuth and increasing the volume of solution to
be evaporated in P. 24. In order to avoid this dilution, an appro^i-
mate estimate may be made using a 1 n. chloride solution. The
results of tests of this method are shown in Table XV. 3 The sodium
chloride solution used was prepared to be exactly In.; the silver and
other elements listed were added as solutions of their nitrates to 30
ml of 3 n. HNOa; the time required for the titrations varied from 10
to 20 min. This method, while not adopted for general use in this
procedure, may be employed at the judgment of the individual ana-
lyst, thus eliminating P. 23.

The precipitate is washed with HN0 3 , and not pure water, to
avoid precipitating oxy-salts of bismuth. It is washed first with
sufficient solution to remove practically all of any other Copper

TABLE XV
THE TITRATION OF SILVER "WITH 1 N. CHLORIDE SOLUTIONS

Experiment

Elements Taken
(mg)

NaCl Used

(ml)

Ag Found

(mg)

Ag 50

0.45

AglOO

0.92

Ag250

2.38

257

Ag500

4.72

509

/Ag 250\
\Bi 250/

2.40

259

/Ag 250\
\Pb 250J

2.35

254

Group elements present, and these washings are combined with the
filtrate. It is finally thoroughly washed in order to remove any
traces of bismuth or lead which might interfere with the subsequent
determination of silver by forming precipitates when the silver
chloride is dissolved in ammonia.

If a standard solution of potassium cyanide is not available, the
estimation of silver can be made by one of the methods used in the
gravimetric standardization of a hydrochloric acid solution that
is, by weighing the precipitated silver chloride. Because of the
rapidity with which it can be carried out, the methpd in which the
precipitate is dried by means of alcohol (P. XVIII D) is recom-
mended. See pp. 129-131 of the discussion accompanying that
procedure.

* Unpublished experiments by Everett G. Trostel.

P. 22] DETECTION OF SILVER 225

Procedure 22: DETECTION AND PRECIPITATION OP SILVER.
Evaporate the filtrate from P. 21 to a volume of 20 to 30 ml
and add to the hot solution just 0.1 ml of 1 n. NH 4 C1 solu-
tion. (White precipitate, presence of silver. Notes 1, 2;
Note 6, P. XVIII.)

If there is no precipitate, treat the solution by P. 24.

If there is a precipitate, add 1 n. NE^Cl, 0.5 ml at a time,
until no more precipitate forms. Avoid adding more than
0.5 ml excess of the NEUCl by keeping the mixture hot,
shaking it, and letting the precipitate settle after adding
each portion of the reagent so that the effect of the next por-
tion can be observed in the clear solution. (It should not be
necessary to add more than 6 ml of the 1 n. NHiCl.) When
no more precipitate forms, heat the mixture almost to boil-
ing and shake it vigorously until the precipitate settles
rapidly.

If a gravimetric estimation of the amount of silver present
is desired, this can be carried out by any one of the optional
methods given in P. XVIII, "The Gravimetric Standardiza-
tion of an Hydrochloric Acid Solution." Method D,
involving the use of a sintered-glass filter and the use of
alcohol for drying the crucible and precipitate is recom-
mended. The precipitate is washed as directed in the
paragraph below.

Prepare a small paper filter, not larger than 7 cm, and,
with the aid of a stirring rod, decant the hot solution
through the paper, receiving the filtrate in a 200-ml flask
(Note 3) . Carefully drain the solution from the precipitate,
carrying as little of the precipitate as possible onto the
paper. Wash the precipitate remaining in the flask with
three 5-ml portions of hot 1.2 n. HNOs, heating each portion
just to boiling after adding it to the precipitate, and then
again decanting (Note 4). When decanting the solution,
direct it to the lower part of the filter paper with a stirring
rod. Add these washings to the filtrate and treat it by
P. 24. Wash the precipitate on the filter paper and that in
the flask with 0.6 n. HNO and treat it at once by P. 23
(Note 5).

Notes:

1. If silver is found present, direct sunlight should be avoided in carrying
out the remainder of P. 22 and P. 23. The silver halides are decomposed

226 ANALYSIS OF THE COPPER GROUP [P. 23

by sunlight, and the metallic silver formed would dissolve in either the
NEUOH or KCN used in P. 23 only as it was oxidized by any oxygen present.

2. It is recommended that an approximate estimation of the silver present
be made at this point by slowly adding a standard 1 n. NH 4 C1 or NaCl
solution from a graduated pipet or tfuret and noting when no more precipi-
tate forms. The NH 4 C1 should be added in successively smaller portions
as the rate of precipitation is seen to be decreasing, until finally the effect
of each drop is noted. After the end-point is obtained by this method, an
excess of 0.1 to 0.2 ml of NH 4 C1 should be added in order to insure complete
precipitation. The necessity of carrying out P. 23 can often be eliminated
by this process (see the discussion).

If a precipitate which appears to contain less than 10 mg of silver is ob-
tained, it should be estimated by comparing it with what is thought to be
an equal amount of silver which has been precipitated under similar condi-
tions. With such small quantities of an element, this method is in general
more trustworthy than a titration.

3. It should be an invariable practice to test either the clear solution above
the precipitate or the first portion of the filtrate to see that an excess of the
precipitant has been added. In this case a few drops of ammonium or
sodium chloride should be added to the clear filtrate and the solution exam-
ined for any evidence of further precipitation.

4. When it is noted that the next operation to be carried out with the
precipitate can be performed in the vessel in which it is already contained,
it is better not to transfer the precipitate to the filter but to wash by decan-
tation, as was done in this case. The small amount of precipitate un-
avoidably carried onto the filter can then be much more easily dissolved
or transferred, as the case may be. In this case, if a large silver chloride
precipitate were transferred to the filter, it would be difficult to dissolve it
in the ammonia used in P. 23.

5. If the AgCl precipitate is allowed to stand for a considerable length
of time, it dissolves slowly when treated by P. 23.

P. 23. Estimation of Silver

Discussion. The method used here for the estimation of silver
is based upon the classical Liebig 4 method for determining cyanide.
The principles involved are as follows : A silver chloride precipitate,
although very slightly soluble in water, dissolves in an ammonia
solution, owing to the formation of the complex ion, Ag(NH3) 2 + ,
called the diammino silver ion. Silver iodide is even less soluble
in water than the chloride and dissolves scarcely at all in ammonia
solutions; therefore, upon the addition of an iodide to the ammonia
solution, silver iodide precipitates. Still less dissociated than the
diammino silver ion is the complex ion which silver forms with
cyanide, having the composition Ag(CN) 2 ~. Therefore, when KCN
is added to an alkaline solution containing both the diammino ion

Liebig, Liebig's Ann. d. Chem., 77, 102 (1861).

P. 23]

ESTIMATION OF SILVER

227

and a silver iodide precipitate, there is a quantitative conversion of
first the ammino ion and then the iodide precipitate into the soluble
cyanide ion. If a standard cyanide solution is used and the point
at which the silver iodide precipitate disappears is noted, the amount
of silver present can be calculated. Because the silver iodide acts
only as an indicator for the titration, it is not necessary to measure
exactly the amount of iodide that is added ; for a very precise deter-
mination, this quantity, as well as the volume of the solution and
the amount of ammonia present, should be closely controlled, as
they each affect the end-point. 5 As, under the conditions of this
procedure, the silver chloride precipitate may so coagulate as to
dissolve slowly in an equivalent amount of cyanide, a sharper end-
point is obtained by adding a slight excess of cyanide and then back-
titrating with standard silver nitrate solution to the appearance of
the characteristic yellowish turbidity of the silver iodide precipitate.

In an experimental study 6 of the procedure used here, it was found
that, if (in P. 22) the silver chloride precipitate was thrown on the
filter, it collected into large particles which dissolved very slowly in
even concentrated ammonia, thus causing results which were from
1 to 2 per cent low. This loss, as
well as the time required for the
process, was reduced by carefully
filtering by decantation. The table
at the right shows the results ob-
tained by treating solutions con-
taining silver and lead by P. 22 and
P. 23.

This method of estimating the
silver present has been adopted
because the titration can be made
directly in the ammonia solution
of the silver chloride precipitate.
Other more commonly used methods
for estimating silver, such as the
titration with standard thiocyanate

in nitric acid solution with ferric salt as indicator (the Volhard
method), would require a rather difficult preliminary separation of
the silver from the chloride.

TABLE XVI

THE TITRATION OF SILVER WITH
STANDARD CYANIDE SOLUTION

Experi-
ment

Elements
Taken

(mg)

Silver
Found

(mg)

AgSOO

499

Ag250

249

fAg 54.81
\Pb 400 J

54.6

(Agl64.2\
\Pb300 J

164.1

[Ag 331.41
1 Pb 100 |

331.0

* For a detailed discussion of the conditions affecting the Liebig method for
determining cyanide, see Kolthoff and Furman, Volumetric Analysis, Vol. 1,
pp. 40-64 and Vol. II, pp. 238-240.

' Unpublished experiments by Robert J. Coulter.

228 ANALYSIS OF THE COPPER GROUP [P. 23

Procedure 23: ESTIMATION OF SILVER. Place the flask,
which should still contain most of the precipitate, under the
filter through which the filtrate has been passed. Pour,
drop by drop, 5 ml of warm (40 to 60C.) 6 n. NH 4 OH
through the filter (moistening thoroughly every portion of
. it) into the flask (Note 1). Wash the paper in the same
manner with 10 ml of water, collecting these washings in
the flask. (Discard the paper.) Shake the mixture in the
flask; crush any undissolved particles of precipitate with a
stirring rod and add to it 1 ml of 1 n. KI solution. Take in
a clean 50-ml buret (Note 7, P. V; Notes 4, 5, 6, P. VI) what
is thought to be a sufficient volume of standard 0.2 f . KCN
solution for the titration (Notes 2-5) and note the reading
of the meniscus. Add this solution to the mixture in the
flask in 0.5-ml portions until the precipitate dissolves
(Notes 3, 4), shaking the mixture vigorously and breaking
up any lumps of precipitate present after adding each
portion of the KCN solution. Then titrate the solution
with standard 0.1 n. AgNOs, added drop by drop with con-
stant shaking, until a permanent yellowish turbidity is pro-
duced. From the volume of the KCN and AgNOa solutions
used calculate the amount of silver present (Note 5).

Notes:

1. The small quantity of finely divided precipitate which is carried onto
the filter is readily dissolved by the ammonia. It is essential that most of
the silver chloride be kept in the flask, since it is difficult to dissolve a large
precipitate from the filter even with concentrated ammonia. If much pre-
cipitate has been carried to the filter, it should be returned, by means of a
stirring rod, to the flask; the ammonia should be poured repeatedly through
the filter and then collected in the flask.- Tl^e precipitate in the flask, if
large, may not completely dissolve in the ammonia, and, although this is
desirable, it is not necessary, as it will dissolve in the KCN. The titration
with the KCN can be carried out more rapidly, however, if a clear ammonia
solution is obtained before adding the iodide.

2. It is an extravagant waste to fill a buret for every titration. Estimate
the required amount of the standard solution from the size of the precipitate
and take only slightly more than that amount.

3. Under certain conditions, and especially if it has been allowed to
stand, particles of the silver chloride precipitate may dissolve slowly upon
addition of the KCN. Accordingly, they should be carefuHy broken up,
preferably with a stirring rod which has been heated and flattened out at
one end. A small, flogculent, white precipitate, caused at times by traces
of bismuth or lead which have been incompletely washed from the AgCl

P. 24] PRECIPITATION OF LEAD 229

precipitate or by small amounts of aluminum in the reagents, should not
be mistaken for the more granular silver chloride particles or for the very
characteristic, usually colloidal, yellowish silver iodide.

4. If the silver chloride precipitate completely dissolved in the NH40H,
it is often possible to avoid the back-titration with AgNOa by adding the
KCN in successively smaller portions as the precipitate of Agl begins to
diminish, and then drop by drop until the precipitate dissolves. Even if the
end-point is slightly overrun, the few drops of AgNOa needed for the back-
titration then can be added with sufficient accuracy from a 1-ml measuring
pipet or even with a dropper from a small graduate. The appearance (or
disappearance) of the precipitate is best noted by using a black background.

5. The potassium cyanide solution should be standardized each day it is
used, as its concentration gradually decreases. Proceed as follows:

Pipet 25 ml of the cyanide solution into a 200-ml flask, add 5 ml
of 6 n. NH 4 OH and 1 ml of 1 n. KI, and dilute to 50 ml. Titrate
the solution as directed above with standardized AgNOa (prepared
in P. V) until the first yellowish turbidity is obtained.
The 0.2 f. KCN solution is prepared by dissolving approximately 13 g of
KCN and 5 g of KOH in 1 1 of solution. The KOH represses the hydrolysis
of the KCN and,makes the solution more stable.

P. 24. Precipitation of Lead

Discussion. After silver .has been detected and removed, the
next separation in the analysis of the Copper Group is made by pre-
cipitating the lead as sulfate. This is accomplished by adding sul-
f uric acid and evaporating the mixture to fuming in order to remove
the nitric acid and any chloride present. As is seen from the solu-
bility-product principle, any factor which decreases either the sulfate
or the lead ion concentration will increase the solubility of lead
sulfate. Nitric acid, or any strong acid, has this effect because it
tends to convert SO^ into HS04~~; hydrochloric acid has a still
greater effect, as, in addition to the hydrogen ion effect, the chloride
ions tend to combine with the lead to form complex ions of the type
PbCl 4 ". The solubility of lead sulfate in increasing concentrations
of sulfuric acid at first decreases, owing to the expected common
ion effect, but the decrease is less than would be calculated, because
of the decrease in the activity of the sulfate ion as the result of
inter-ionic attraction forces. Then there is a concentration range
(from somewhat less than 1 f. to approximately 11 f. acid at 25C.)
throughout which there is very little change in the solubility and in
which the minimum solubility is found. Finally, as the concentra-
tion of the sulfuric acid is increased (and the concentration of the
water becomes relatively small), the fraction of the acid ionized into

230 ANALYSIS OF THE COPPER GROUP [P. 24

sulfate decreases to such a small value that in very concentrated
acid there is a marked increase in the solubility of the lead sulfate. 7
This separation of lead, as the sulfate, from copper and cadmium
is very satisfactory. The separation from bismuth is somewhat
less so, for, although the sulfates of copper, cadmium, and bismuth
are readily soluble, bismuth ion tends to hydrolyze, as shown by the
equation

Bi f ++ + H 2 O = BiO+ + 2H+.

On standing, the bismuthyl ion 8 may precipitate from sulfuric acid
solutions as bismuthyl sulfate, (BiO)2S0 4 , or under certain conditions
as a compound having the formula (810)2820 7 -3H 2 O, hydrated
bismuthyl pyrosulfate. These compounds ordinarily remain in the
solution largely because of the formation of supersaturated solutions ;
however, with a large amount of bismuth, or when lead sulfate forms,
they may precipitate in considerable amounts. As the result of an
experimental study of this difficulty, 9 the procedure below was devel-
oped, whereby, if a large precipitate is obtained, the solution is
decanted from it and the precipitate dissolved in hydrochloric acid
and again fumed with a smaller portion of sulfuric acid. It was
found that, under the conditions of the procedure below, if 250 mg
each of lead and bismuth were present, from 30 to 50 mg of bismuth
would precipitate with the lead; that, after decanting, adding 3 ml
of 36 n. sulfuric acid, and again fuming, 7 to 10 mg of bismuth still
remained; but that, after adding 5 ml of 12 n. hydrochloric acid and
then fuming with the sulfuric acid, only 1 to 1.5 mg of bismuth
remained. The concentrated hydrochloric acid almost completely
dissolves or metathesizes the precipitates, thus facilitating the
extraction of the bismuth.

Other details of the separation are as follows: Experiments have
shown that the concentrated sulfuric acid should not be diluted by
pouring it into water, as is often done in qualitative systems for
reasons of safety, since the first portions of acid are thereby so
diluted as to favor the precipitation of the bismuth oxy-salt. The
solution should be kept cold when the sulfuric acid is being diluted, in
order to minimize the hydrolysis of the bismuth and to decrease the
solubility of the lead sulfate, which increases rapidly as the tempera-

7 For the solubility of lead sulfate in sulfuric acid of various concentra-
tions, see Creckford and Brawley, J. Am. Chem. Soc., 56, 2600 (1934).

Such hydrolyzed or "oxy" ions are frequently designated by the suffix
yl. For example, the antimonyl ion is SbO+; the vanadyl ion is VO" 1 " 1 ".

' Unpublished experiments by Randal Maass.

P. 24] PRECIPITATION OF LEAD 231

turc is raised. As lead sulfate also forms supersaturated solutions,
the mixture is allowed to stand for 5 min. after it is diluted; upon
standing much longer, bismuthyl compounds tend to precipitate.
The precipitate is washed first with 1.2 n. H 2 S04, and not with
water, to remove any bismuth sulfate without causing it to hydrolyze.
The H 2 SO4 is then washed out in order to make the precipitate more
soluble in the ammonium acetate later used.

In spite of these precautions, if lead and bismuth ar6 both present
in considerable amounts, a small amount of bismuth may precipi-
tate with the lead ; however, it does not interfere with the determina-
tion of lead in P. 25.

A volumetric method for estimating the lead in the precipitate is
provided in P. 25. However, if standard solutions are not available,
a rapid gravimetric estimation can be made by collecting the lead
sulfate on a previously weighed sintered-glass filter, drying it by
means of alcohol, and then weighing the filter and precipitate.
See pp. 129-131 for a discussion of this method and for data con-
cerning it.

Procedure 24: PRECIPITATION OF LEAD. Cool the solu-
tion from the silver chloride precipitation (P. 22), slowly
add to it just 3 ml of 36 n. H 2 S0 4 (Note 1), and evaporate it
on a sand bath or hot plate (or directly over a burner if it is
continuously kept in swirling motion) until the H2SO4 gives
off copious dense white fumes (Note 2). Allow the mix-
ture to stand until it cools somewhat (Note 3), cool the
flask to room temperature with running water, and, while
continuously cooling the flask, pour slowly into it, 1 ml a* a
time, 20 ml of water. Allow the solution to stand for not
less than 5 min. and not longer than 10 min. (White pre-
cipitate, presence of lead.)

If a gravimetric estimation of the amount of lead present
is desired, the precipitate should be collected on a sintered-
glass filter, treated and washed as directed in the following
paragraphs, and dried with alcohol as directed in Optional
Method D of P. XVIII, "The Gravimetric Standardization
of an Hydrochloric Acid Solution. "

If only a small precipitate forms, decant the solution
through a paper filter (Note 4), collecting the filtrate in a
200-ml flask. Wash the precipitate by decantation with
three 5-ml portions of 1.2 n. H 2 SO4, taking care to retain as
much as possible of the precipitate in the flask, and add

232 ANALYSIS OF THE COPPER GROUP [P. 25

these washings to the original filtrate. Treat the filtrate by
P. 26. Wash the precipitate with two 5-ml portions of cold
water and treat it by P. 25.

If considerable precipitate forms (Note 5), decant the
solution through a paper filter, retaining as much as possible
of the precipitate in the flask, and collect the filtrate in a 200-
ml flask. Add to the precipitate in the flask 5 ml of 12 n.
HC1 and 2 ml of 36 n. H 2 S04, evaporate the solution as
directed above until the H 2 S04 fumes, allow the mixture
to cool, and add to it in the manner directed above 15 ml
of water. Treat the mixture as directed in the preceding
paragraph, using the same filter, and combine the filtrates.

Notes:

1. The acid should be added slowly and the solution kept cool; otherwise
spattering may result.

In general, concentrated sulfuric acid should be added to an aqueous
solution, not in the reverse order. Elsewhere in this procedure this order is
reversed, because a better separation results and because transferring of a
precipitate from one vessel to another is thereby avoided. This mixing
should be carried out very slowly and the solution constantly cooled.

2. It is essential that the E^SO* be made to fume in order to expel all of
the HNOs and HC1. The more transparent HNOa fumes produced as the
solution becomes concentrated should not be mistaken for the dense white
fumes of H2S04. The latter can be recognized by the choking sensation
which even a very small amount causes. It is to be noted also that the
HjSOi will not begin to fume until the solution is reduced to about the
volume of 36 n. E^SCU added.

3. The flask should be allowed to cool somewhat before it is cooled fur-
ther with tap water, or it is likely to crack.

4. If a precipitate corresponding to only a few milligrams is obtained, the
presence of lead should be confirmed and a visual estimation should be made
as directed in Note 4 of P. 25.

5. The precipitate need not be treated by this last paragraph unless it
ifil thought to contain a considerable amount (50 mg or more) of bismuth,
and unless the analyst desires to make a precise estimation of the lead or
bismuth present.

P. 25. Estimation of Lead

Discussion. Lead is estimated in this procedure by dissolving
the lead sulfate precipitate in an ammonium acetate solution,
reprecipitating the lead as chromate from an acetic acid-acetate
solution by adding a measured volume of standard chromate solu-
tion, and then determining the excess of chromate iodometrically.

P. 25] ESTIMATION OF LEAD 233

Lead sulfate dissolves in an ammonium acetate solution because
of the formation of the slightly ionized lead acetate molecule. 10
If a highly ionized acid, such as nitric, is added to this solution, the
acetate concentration would be so greatly decreased (the ionization
constant for acetic acid being 1.8 X 10~ 6 ) that the lead sulfate
might be at least partially reprecipitated; for this reason, the solu-
tion is acidified with acetic acid before precipitating the chromate.
Although chromic acid is not a highly ionized acid and is converted
into dichromate by hydrogen ion according to the reaction

2CrOr + 2H+ = Cr 2 7 " + H 2

(see the discussion of P. 82), extensive investigations 11 have shown
that lead chromate is so very insoluble that it can be almost quanti-
tatively precipitated from acetate and acetic acid solutions. Con-
firmatory experiments 12 have shown that for the estimation of lead
the precipitate should be formed in an acid solution, or low results
are obtained, owing probably to the formation of basic lead chro-
mates. The precipitation of the yellow lead chromate serves as a
sensitive confirmatory test for the presence of lead as well as pro-
viding a means for its indirect determination. Bismuth, even if
dissolved by the ammonium acetate, would not give a precipitate
because of the acidity of the solution.

The lead chromate precipitate could be dissolved in hydrochloric
acid and directly treated with potassium iodide, and the resulting
iodine could be titrated. Even large amounts of lead chromate are
readily dissolved by hydrochloric acid and by concentrated chloride
solutions, owing to the formation of complex ions of the type PbCU".
This procedure is recommended for use when a small amount of lead
is present. With considerable amounts of lead it was found that
the yellowish lead iodide precipitate which usually formed interfered
with the determination of the end-point and also that it required
considerable time to wash the lead chromate precipitate free of the
excess chromate.

10 Noyes and Whitcomb, /. Am. Chem. Soc., 27, 747 (1905); Fox, /. Chem.
Soc., 95, 878 (1909).

"Patten, /. Assn. Off. Agr. Chem., 4, 217 (1920) ; Goode and Summers,
Soc. Chem. Ind., Victoria, Proc., 32, 689 (1932); Fairhall and Akatsuka, /.
Am. Chem. Soc., 56, 14 (1934); Huybrechts and Degard, Bull. Soc. Chim. Belg.,
42, 331, (1933); Guzelj, Z. anal. Chem., 104, 107 (1936).

12 Unpublished experiments by R. W. Hoeppel.

234 ANALYSIS OF THE COPPER GROUP [P. 25

The chromate-iodide reaction and the subsequent titration are
carried out under conditions similar to those which have been used
in standardizing the thiosulfate solution against potassium dichro-
mate (P. XIV); that procedure should be consulted for the details
of the method.

Procedure 25: ESTIMATION OF LEAD. Place the flask-
containing most of the PbSOi precipitate (from P. 24) under
the funnel and with a stirring rod carefully tear a small hole
in the bottom of the filter. Wash through as much of the
precipitate as possible with 5 ml of water, and then pour
dropwise over every portion of the filter 5 to 20 ml of 3 n.
NH4C 2 H 3 02 (Note 1). Discard the filter.

Dilute the filtrate to 50 ml, heat it to boiling, and add to
it 2 to 10 ml of HC2H 3 2 (about one-half the volume of
NH 4 C2H 3 02 used above). Add to the hot solution from
a buret or pipet (Note 2) standard 0.1 n. (oxidizing equiva-
lents) K2Cr 2 O 7 solution, 5 ml at a time, until no more precipi-
tate appears to form. The mixture should be heated almost
to boiling, swirled vigorously, and then allowed to settle
before each addition so that it can be observed if more pre-
cipitate is formed (Note 3). When no more precipitate
forms, heat the mixture and frequently swirl it until the
precipitate settles rapidly (Note 4) ; if the solution does not
show a distinct orange-red color, more K 2 Cr 2 0? should be
added.

If less than 40 to 50 mg of lead are thought to be present,
cool the mixture to room temperature and filter the solution
through a small paper filter. Wash the precipitate and
filter with cold water, added dropwise, until the wash solu-
tion runs through colorless. Discard the filtrate. Dis-
solve the precipitate by pouring dropwise through it a solu-
tion made by adding 10 ml of HC1 to 50 ml of water and
saturating the solution with NaCl. Collect the solution in
a 400-ml flask and treat it as directed in the second para-
graph below.

If more than 40 to 50 mg of lead are thought to be present,
cool the mixture to room temperature, transfer it with the
aid of cold water to a 250-ml volumetric flask (Note 5), and
dilute it exactly to the mark. Mix the solution thoroughly
and then, while the precipitate settles, prepare a small

P. 26] ESTIMATION OF LEAD 235

paper filter (Note 6). Decant 10 to 15 ml of the super-
natant solution through the filter into a 100-ml flask. Use
this portion to rinse the filter and flask and then discard it
(Note 7). Decant 60 to 70 ml of the solution through the
filter into the flask, pipet 25 ml of it into a 400-ml flask, add
to it 10 ml of HC1 and 25 ml of water, and treat it as directed
in the next paragraph.

Add to the solution 2 to 3 g of KI dissolved in 40 ml of
water. Close the flask and allow it to stand for 3 min. in a
dark place. Dilute the solution to 300 to 400 ml and
titrate it at once with standard 0.1 n. Na2$203 solution
until the iodine color becomes indistinct (Note 8), and
then add 5 ml of starch indicator solution and again titrate
until the blue color just disappears. From the volume of
standard dichromate solution used to precipitate the lead,
and from the volume of standard thiosulfate solution used
in this tit rat ion, calculate the amount of lead present.

Notes:

1. The volume of NH^HaC^ used should be adjusted to the size of the
precipitate, noting that PbS0 4 is a compact, heavy substance. As an aid
to judging the volume of I^C^Oy to be later used, a visual estimate of the
amount of lead .present should be made. If experience with this type of
precipitate is lacking, a known amount of lead, obtained from the test solu-
tion provided in the laboratory, should be fumed with H^SO^ the mixture
then diluted as in P. 24, and the precipitate carefully compared with that
already obtained.

The minimum amount of NH^HsC^ required should be added; the
solubility of PbCr04 becomes appreciable in very concentrated acetate
solutions.

2. The approximate volume of 0.1 n. K^Cr^Oj required for the amount
of lead judged to be present should be calculated, as, when this is known, a
pipet can be used more quickly than a buret. An excess of at least 3 ml of
the K2Cr 2 O 7 should be added.

3. The PbCr04 precipitate tends to remain suspended until an excess of
the dichromate has been added, and then it settles much more rapidly.
This behavior is so characteristic that one familiar with the precipitation
can judge very closely when the equivalent amount of dichromate has been
added.

4. If less than about 5 mg of lead are thought to be present, it is usually
more satisfactory to make the estimation by comparing the size of the pre-
cipitate with known amounts of lead which have been precipitated under
similar conditions; the latter is essential, as the conditions of the precipita-
tion may markedly affect the character and apparent size of the precipitate.

5. Hot solutions should not be poured into volumetric measuring flasks.

236 ANALYSIS OF THE COPPER GROUP [P. 26

These flasks are usually calibrated for use at room temperature (20C.) and
return slowly to their original size after being heated.

6. If the mixture can be conveniently left until the precipitate completely
settles, the clear supernatant solution may then be pipeted directly from
the flask and a filtration can be avoided. Care should be taken not to stir
up the precipitate. The error introduced by the volume of the precipitate
is so small that it can be neglected.

7. It is more convenient to rinse the paper and the collecting flask and
then discard the first portion of the solution than to use a dry filter and
flask. Also, in more precise quantitative work this procedure would be
advisable, as the filter paper may adsorb an appreciable quantity of the con-
stituents from the first portion of the solution passing through it.

8. One accustomed to this method can carry the titration to within a few
drops of the end before adding the starch. This is quite desirable, for starch,
when used as an indicator, should not be added to a solution which contains
a considerable quantity of iodine, as it then tends to coagulate, making the
end-point much less definite.

P. 26. Precipitation of Bismuth and Detection of Copper

Discussion. After silver and lead have been removed from the
solution of the Copper Group, the separation of bismuth from
copper and cadmium is next made by adding an excess of ammonium
hydroxide to the sulfuric acid solution. Bismuth hydroxide is
relatively insoluble, exhibits no appreciable amphoteric tendency
in the low hydroxyl ion concentration provided by an excess of
ammonia in the presence of ammonium salts, and does not form an
ammino complex ion; therefore, it is quantitatively precipitated.
As the solution is neutralized, the bismuth ion hydrolyzes and is
precipitated as a basic salt, probably (BiO) 2 S04. This may not be
completely converted into the hydroxide by an excess of ammonia,
and therefore the precipitate varies in its composition, so that bis-
muth cannot be precisely determined by attempting to ignite this
precipitate to the oxide (Bi 2 03) and weighing it as such.

In the presence of an excess of ammonia, copper and cadmium
form the tetrammino ions CuCNHs)/* and CdCNHa)/*. These
ions remain in solution, as their hydroxides and salts are soluble.
The tetrammino copper ion has an intense blue color, which serves
as a distinctive means for detecting copper. In this procedure the
solution is evaporated to a small volume in order to make this detec-
tion more sensitive; the blue color given by mg of copper is readily
perceived under these conditions.

Experiments carried out under the conditions of the procedure
below have shown this separation to be very satisfactory; with 250

P. 26] BISMUTH AND COPPER 237

mg of bifemuth and 250 mg of copper or cadmium, not more than 1
mg of copper or cadmium was found in the precipitate. For a
more detailed discussion of the separation of various other elements
by the use of an ammonia precipitation, see the discussion of P. 55.
The precipitate is washed with only an amount of 1.2 n NEUOH
which will remove practically all of any copper and cadmium present,
and these washings are added to the filtrate. It is not further
washed, as none of the constituents of the solution interfere with the
subsequent estimation of bismuth.

Procedure 26: PRECIPITATION OF BISMUTH AND DETEC-
TION OF COPPER. Evaporate the filtrate from P. 24 to about
15 ml (Note 1) and cool it with running water. Add slowly
with a dropper 15 n. NH 4 OH until it is alkaline (Note 2)
and then add 2 ml more. (White precipitate, presence of
bismuth; blue solution, presence of copper. Note 3.)

Heat the solution almost to boiling and let the precipitate
settle. Filter the solution through a paper filter, decanting
as much as is possible, and wash it with two or three 5-ml
portions of hot 1.2 n. NH 4 OH, adding these washings to the
filtrate. Treat the filtrate by P. 28 if copper is present, or
by P. 29 if copper is absent. Treat the precipitate by P. 27.

Notes:

1. Solutions can be safely and rapidly evaporated from conical flasks by
giving the contents a continuous rapid swirling motion while the flask is
heated directly over the flame until the solution boils rapidly. Vigorous
boiling without keeping the flask and contents in motion is dangerous with
most solutions, because of the danger of "bumping."

2. The solution is cooled and the concentrated ammonia is added slowly
to the acid solution in order to avoid danger of spattering.

When a solution is made alkaline with ammonia, the approximate volume
to be added can, in general, be estimated, and the solution can then be
added in small portions until the smell of ammonia can be detected. Avoid
splashing the ammonia on the sides of the flask and mix the contents thor-
oughly by swirling.

3. When the solution is intensely colored, it should be examined in a
bright light so that a small bismuth precipitate will not be overlooked, If
the presence of copper is doubtful, owing to the presence of a precipitate or
the indefiniteness of the blue color, the solution, after being filtered, should
be compared against a white background with the same volume of water in
a similar flask. Under these conditions the blue color caused by 5 mg of
copper can be readily detected.

238 ANALYSIS OF THE COPPER GROUP [P. 27

P. 27. Estimation of Bismuth

Discussion. The volumetric method used for determining bis-
muth depends upon precipitating the oxychloride (bismuthyl chlo-
ride, BiOCl), dissolving this in nitric acid, precipitating the chloride
ion with an excess of a standard silver nitrate solution, and then
titrating the excess of silver with a standard thiocyanate solution,
using ferric nitrate as the indicator. This titration has been dis-
cussed in P. VI.

The volumetric determination of bismuth is usually not attempted,
the BiOCl precipitate being filtered, washed, dried at 100 to 110C.,
and weighed; if desired, this procedure can be followed here. How-
ever, a study of the advantages of the two methods, especially as
regards rapidity and accuracy, has indicated that the volumetric
procedure is the more satisfactory, as it can be carried out in about
one-half the time and with comparable accuracy. 13 If standard
solutions of silver nitrate and potassium thiocyanate are not at hand,
the gravimetric method is more expedient for a single determina-
tion; a procedure for its optional use is provided.

The bismuth hydroxide precipitate is dissolved in hydrochloric
acid, as the precipitate dissolves much more readily in it than in
nitric or sulfuric acids, it having been shown that the ions BiCU"
and Bids'" are formed; 14 also, the chloride for the subsequent pre-
cipitation is thereby provided.

When the hydrochloric acid solution of the bismuth hydroxide
is sufficiently diluted, bismuthyl chloride precipitates; that this
should occur is shown by the following equations:

BiCir + HOH = BiOCl ( .) + 2H+ + 3CF (1)

BiCir + HOH = BiOCl ( ., + 2H+ + 4CF. (2)
The precipitation reaction is often represented as follows:

Bi++ + + Cl~ + HOH = BiOCl (8 ) + 2H+. (3)

However, an inspection of the equilibrium expression for this reac-
tion shows that, except for activity effects, it should be independent
of dilution, as the concentrations of the rcactants occur to the same
power in both the numerator and the denominator; it is also seen
that dilution should markedly favor the formation of the precipitate

18 Unpublished experiments by Roland C. Hawes (1928). Migray, Chem.
Ztg. 9 57, 744 (1933), has recently investigated the same method.
14 Noyes, Hall, and Beattie, /. Am. Chem. Soc., 34, 2626 (1917).

P. 27] ESTIMATION OF BISMUTH 239

by either of the first two reactions. Since an excess of acid has been
used in dissolving the ammonia precipitate, and as this (with the
acid formed from the precipitation of a large quantity of bismuth)
would tend to prevent complete precipitation of the oxychloride,
the acidity of the solution is finally adjusted with the aid of methyl
orange. Methyl orange is used because it is desirable that the
solution not become actually alkaline during this adjustment, for
BiO(OH) might then precipitate and not be metathesized to oxy-
chloride during the subsequent boiling. Heating the solution
favors the hydrolysis reaction. Experiments have shown that,
under the conditions of this precipitation, 1 mg of bismuth can be
detected; that with 500 mg not over 0.1 mg remained in the filtrate;
and that with 50 mg of lead no precipitate was produced. The
precipitation of the bismuthyl chloride also serves as a distinctive
confirmatory test; any lead not precipitated as sulfate (in P. 25)
would be precipitated with the bismuth by the ammonia but would
not precipitate with the bismuthyl chloride.

The precipitate is dissolved in nitric acid and the chloride is pre-
cipitated with standard silver nitrate, a large excess being avoided
by noting when precipitation ceases (see the discussion of P. 22).
For determining this excess, the titration with thiocyanate is used
(see P. VI). As the method is used here, the excess of silver nitrate
is titrated in the presence of the silver chloride precipitate. This
may lead to some error, for silver thiocyanate is less soluble than
silver chloride, and therefore the thiocyanate may be removed from
the solution by the metathetical reaction

AgCl (8) + SON- = AgSCNu, + CP.

This causes the end-point to fade and an excess of thiocyanate to be
added. While this titration of silver with thiocyanate in the pres-
ence of a precipitate of AgCl is theoretically inaccurate, it has been
shown 16 that, if the AgCl precipitate is thoroughly coagulated, the
rate at which it reacts with the excess of thiocyanate necessary for
the end-point is so slow that quite accurate results can be obtained. 16
If standard solutions of thiocyanate and silver are not available,
it is recommended that the bismuth be estimated gravimetrically

11 Rothmund and Burgstaller, Z. anorg. Chem. y 63, 330 (1909).

16 Caldwell and Moyer, /. Ind. Eng. Chem., Anal. Ed., 7, 38 (1935), have
suggested that 1 ml of nitrobenzene be added to the solution before making
the back-titration. This is adsorbed on the precipitate, causes it to coagu-
late, and further decreases the rate of the metathesis; this method has been
found to be very effective and its use is recommended.

240 ANALYSIS OF THE COPPER GROUP [P. 27

by collecting the bismuthyl chloride precipitate on a weighed* sin-
tered-glass filter, drying it with alcohol, and again weighing. See
pp. 129-131 for a discussion and for data relative to this method.

Procedure 27: ESTIMATION OP BISMUTH. Dissolve the
BiO(OH) precipitate (P. 26) remaining in the flask with
5 to 15 ml of warm HC1 (Note 1), and then pour this solu-
tion dropwise through the precipitate on the filter, collecting
the solution in a 400-ml flask. Wash the flask and filter
with 10 ml of hot 1.2 n. HC1. Cool the solution and add
NEUOH until a slight permanent precipitate is produced
(Note 2); then add HC1, 0.1 ml at a time and shaking
vigorously after each addition, until the precipitate just dis-
solves.

Heat the solution almost to boiling, slowly add to it 100
ml of boiling water (Note 3), and boil the mixture until the
precipitate settles rapidly. Add to the boiling solution
2 or 3 drops of methyl orange indicator solution and then
NH 4 OH, drop by drop, until the pink color first begins to
change to yellow (Note 4), and then just 3 drops more than
enough HC1 to restore the distinct pink color. Add to the
mixture 100 ml of boiling water and again heat it just to
boiling until the precipitate settles rapidly (Note 5).

If a gravimetric estimation of the bismuth is to be made,
dry and weigh a sintered-glass filtering crucible as directed
in Optional Method D of P. XVIII. Collect the precipi-
tate on the filter, and wash as directed in the next para-
graph. Dry and weigh the precipitate and crucible by the
same procedure used with the eftipty crucible.

Filter the hot mixture through a paper filter and wash
the precipitate, by decantation, with 10-ml portions of hot
water, retaining as much of the precipitate as possible in
the beaker until the wash water no longer gives a pro-
nounced test for chloride (Note 6). Discard the filtrate
and washings.

Tear a small hole in the bottom of the filter with a stirring
rod and wash most of the precipitate through the funnel
into the original flask with 10 ml of water. Dissolve the
precipitate remaining on the filter by pouring dropwise
through it 5 to 25 ml of HNOj (chloride-free), and then dis-

P. 27] ESTIMATION OF BISMUTH 241

solve the precipitate remaining in the flask with this solu-
tion (Note 7).

Add standard 0.1 n. AgN0 3 from a buret to the HN0 3
solution. 1 ml at a time, until no more precipitate forms.
After adding each portion of the AgNOs, shake the mixture
vigorously and let the precipitate settle so that the effect
of the next portion can be observed in the supernatant
liquid (Note 8; Notes 2, 3, P. 22). Carefully note and
record the amount of AgNOa used.

When no more precipitate forms, heat the mixture just to
boiling and shake it vigorously until the precipitate coagu-
lates into lumps and settles rapidly. Cool the mixture to
room temperature, add to it 5 ml of 1 n. Fe(NOa)3, and
titrate it with 0.1 n. KSCN until the first pink color, per-
manent for 15 to 20 sec., can be observed in the solution.
Calculate the amount of bismuth present from the volumes
of standard AgN0 3 and KSCN used.

Notes:

1. A small precipitate may be due to aluminum, silica, iron, or elements
of the Tin Group. These constituents may have been introduced in the
reagents or may be present owing to incomplete washing of the sulfide pre-
cipitates in P. 11 and P. 12. Such a precipitate should be compared with
known amounts of bismuth, similarly precipitated. The following con-
firmatory test should then be made on both precipitates:

To 0.5 ml of SnCl2 reagent add 2 ml of water and (cooling the
solution) NaOH until the precipitate first produced is dissolved.
By means of a dropper, treat the precipitate on the filter with this
solution. (Immediate blackening of the precipitate, presence of
bismuth.)

The bismuth hydroxide is immediately converted into intensely black
metallic bismuth by the alkaline sodium stannite solution. A slow darken-
ing of the residue is of no significance, as the stannite slowly decomposes
into stannate and metallic tin. Because of this, one unfamiliar with the
test should also apply it to the known bismuth hydroxide precipitate.

2. A volume of ammonia nearly equivalent to the acid present can be
rapidly added, and smaller portions can be added until a transitory pre-
cipitate appears; it should then be added dropwise until the last drop forms
a permanent precipitate.

3. If the solution is rapidly diluted, the resulting precipitate may be so
finely divided as to be difficult to filter or wash.

4. It is essential that the addition of the NH 4 OH be stopped when the
pink color first begins to change. Enough HC1 should then be immediately
added to restore an unmistakably pink color to the solution. With 3 drops

242 ANALYSIS OF THE COPPER GROUP [P. 28

of methyl orange indicator solution these changes can easily be seen even in
the presence of a large precipitate.

5. The BiOCl at times separates in such a finely divided state that it will
pass through the filter unless it is coagulated by boiling, as is directed.

6. These tests should be made by adding a few drops of HNOs and then
a few drops of 0.1 n. AgNOa to the 10-ml filtrate to be tested. It is suggested
that a standard be prepared containing the amount of chloride which would
correspond to not more than 0.5 mg of bismuth, and that the washings be
continued until not more than this amount is found in a 10-ml portion of
wash water. The test for chloride with silver ion is so sensitive that it is
unnecessary and time-consuming to wash until a perfectly clear solution is
obtained.

7. If the precipitate on the filter or in the beaker dissolves slowly, the
HNOs can be heated to 40 to 50C., thereby increasing the rate of solution.
It should not be heated more than this or oxidation of the chloride by the
nitric acid may occur.

8. If it is desired, the estimation of bismuth can be made by this titration,
as suggested in Note 3, P. 22. In general, if a standard thiocyanate solution
is available, the end-point can be more quickly obtained as directed in the
next paragraph of the procedure.

P. 28. Precipitation and Estimation of Copper

Discussion. If copper has not been found present in P. 26, this
procedure for its estimation should be omitted.

Upon adding an excess of a soluble iodide to a solution containing
cupric copper, a precipitate of cuprous iodide is formed and an
equivalent amount of iodine is set free. By titrating this iodine
with standard thiosulfate, an estimation of the copper present is
obtained.

From an inspection of the potentials (see Appendix) of the elec-
tronic reactions

Cu + = Cu++ + E~
and

3F = I 3 ~ + 2E~,

it is obvious that the equilibrium of the reaction

+ 3F =

should greatly favor the formation of the products on the left side
of the equation. If, however, the effect of the solubility of cuprous
iodide is considered, it is seen that it may have a large effect in
shifting the equilibrium toward the right by keeping the concentra-
tion of the cuprous ion at a very low value. In fact, from the
potential for the reaction

Cul (i) = Cii"" + r + E"

P. 28] ESTIMATION OF COPPER 243

it is seen that the equilibrium for the reaction

2Cu++ + 5I~ = 2CuI (l) + I 3 ~

should favor the formation of the products on the right; and a calcu-
lation of the equilibrium constant enables the prediction to be made
that the reaction should proceed to an extent well within the usual
quantitative limits, especially if an excess of iodide is present in the
.solution.

This method for the determination of copper is extensively used
commercially and has been the subject of numerous investigations. 17
These have shown that the solution should contain approximately
4 per cent by weight of potassium iodide, that the cuprous iodide
may absorb iodine, and that in an acid solution there may be some
iodine liberated by the oxygen of the air; these latter effects are not
large and are compensating, thereby permitting results precise to
0.1 to 0.2 per cent to be readily attained.

Cuprous iodide is very slightly soluble in water, so that all but a
few milligrams of the copper is precipitated; this amount remains
in solution, owing mainly to the formation of the CuI 2 " ion. The
precipitate is filtered out and this small amount of copper is left in
the solution, as it does not interfere with the test for cadmium under
the conditions of P. 29. Experiments have shown that no cadmium
is carried out by the precipitate.

The copper can be precipitated as cuprous thiocyanate 18 by the
addition of an excess of a soluble thiocyanate before the titration
of the liberated iodine. An extensive series of experiments 19 have
shown that under very closely regulated conditions the precipitation
of the copper can be made so complete that no precipitate is obtained
upon passing hydrogen sulfide into the filtrate, and, consequently,
the later use of cyanide to prevent the precipitation of cuprous sul-
fide in the precipitation of cadmium as sulfide is avoided. The
accuracy of the method was proved by confirmatory experiments
made with various amounts of copper, both alone and in the presence
of cadmium. However, experience with the procedure in the hands

17 Gooch and Heath, Am. J. Sci., 24, 65 (1907); Bray and McCay, J. Am.
Chem. Soc., 32, 1199 (1910); Peters, ibid., 34, 422 (1912); Popoff, ibid., 51, 1299
(1929); Whitehead and Miller, J. Ind. Eng. Chem., Anal. Ed., 5, 15 (1933).

18 Bruhns, Chem. Zeit., 42, 301 (1918); Kolthoff, ibid. t 42, 609 (1918); Kruger
and Tschirch, Z. anal. Chem., 97, 161 (1934).

19 Unpublished experiments by Karl A. Ganslee (1927).

244 ANALYSIS OF THE COPPER GROUP [P. 28

of persons unfamiliar witn the method showed that often the proper
conditions were not obtained, and the test for cadmium was ob-
scured by a black copper sulfide precipitate. This fact made the
use of an additional confirmatory test necessary, and, as the precipi-
tate of cuprous thiocyanate was no more readily filtered and washed
than that of cuprous iodide, the use of the thiocyanate was aban-
doned. However, if it is desired to avoid the use of potassium
cyanide in P. 29, this method is available and instructions for its use
are provided in Note 4 of this procedure and Note 1 of P. 30. Foote
and Vance 20 have shown that the iodine adsorbed on the cuprous
iodide precipitate is removed by the addition of thiocyanate just
before the completion of the titration with thiosulfate. While pos-
sible error caused by the reaction between the iodine and thiocyanate
is thus eliminated and the end-point is made much more distinct,
especially for one not accustomed to the titration, it was found in the
series of experiments mentioned above that the precipitation of the
copper was not so complete as when the thiocyanate was added
before the titration was begun. The effect may be caused by a slow
precipitation of the CulV ion. When the complete precipitation
of copper is not necessary, it is recommended that thiocyanate be
added just before taking the end-point.

A rapid determination of the copper could be made by titrating the
ammoniar*! solution with a standard cyanide solution until the dis-
appearance of the blue color. This method is capable of giving
fairly consistent results when the amount of copper being deter-
mined does not vary widely, when the conditions as to temperature,
volume, and concentration of ammonia and ammonium hydroxide
are closely regulated, and when the cyanide solution is standardized
with copper under these same conditions. However, as the reaction
taking place, which is essentially represented by the equation

2Cu(NH,) 4 " H " + 7CN~ + H 2

= 2Cu(CN) 8 - + CNO~ + 2NH 4 + + 6NH 8 ,

depends appreciably upon the conditions existing in the solution
and is, moreover, influenced by the presence of cadmium, which
likewise forms a complex cyanide ion, this method is not as reliable
as the iodo metric titration, and, therefore, it is not generally advised.

10 Foote and Vance, J. Am. Chem. Soc., 57, 845 (1935).

P. 28] ESTIMATION OF COPPER 245

In case only a quick approximate estimation is desired, it may be
used. 21 As is mentioned above, the iodometric method is exten-
sively used commercially for determining the copper in alloys and
ores. The only elements commonly occurring in copper ores which
would oxidize iodide are arsenic, antimony, and ferric iron. It has
been shown that the effect of the iron can be eliminated by adding
fluoride (which forms the complex ion FeFe*), 22 and oxidation by
the arsenic and antimony is prevented by buffering the solution to
a pH of approximately 3.3; this buffering effect can be obtained by
providing the proper ratio of fluoride and hydrofluoric acid in the
solution. 23 A procedure for the direct determination of copper in
ores is given in Note 8 below. The copper in alloys may usually
be determined by dissolving the metal in nitric acid and then fol-
lowing P. 24, P. 26, and P. 28.

Procedure 28: PRECIPITATION AND ESTIMATION OF COP-
PER. If in P. 26 copper was found to be present (Note 1),
carefully neutralize (Note 2) the NH 4 OH solution with
H 2 S04 and then add just 2 ml more. Have available a buret
containing standard 0.1 n. Na2S20 3 solution. Add to the
acidified solution 2 to 3 g of solid KI (Note 3), swirl the
mixture until the KI dissolves, and titrate it immediately
with the Na2S20 3 solution (Note 4). Add the thiosulfate
rapidly until the color of the iodine becomes indistinct,
add 5 ml of "starch solution/' and again titrate just to the
disappearance of the starch color (Note 5) . Note the buret
reading and from the volume of standard thiosulfate used
calculate the amount of copper present.

Shake the mixture vigorously for several minutes, let
the precipitate settle, and filter it out on a paper filter (Note
6). Wash the precipitate with two or three 5-ml portions of
cold water, adding these washings to the filtrate. Treat
the filtrate at once by P. 29 (Note 7). Discard the pre-
cipitate.

21 See Treadwell-Hall, Analytical Chemistry, Vol. II, Quantitative, 7th
Ed., p. 616.

Mott, Chem. Analyst, 5, 7 (1912).

11 Park, /. Ind. Eng. Chem., Anal. Ed., 3, 77 (1931); Crowell, Hillis, Ritten-
berg, and Evenaon, ibid., 8, 9 (1936); Foote and Vance, ibid., 8, 119 (1936);
Crowell, Silver, and Spiher, ibid., 10, 80 (1938).

246 ANALYSIS OF THE COPPER GROUP IP. 28

Notes:

1. If the amount of copper present is thought to be less than 10 to 15 mg,
it can be estimated by adding to the same volume of solution, in a similar
flask containing the same amount of ammonium sulfate and of ammonia, a
standard solution of copper, until, when compared as directed in Note 3,
P. 26, the intensity of the blue color in the two solutions is the same. If it
is estimated that only this amount of copper is present, the remainder of this
procedure can be omitted, as it will not interfere with the detection of cad-
mium.

2. The disappearance of the intense blue color due to the tetrammino
copper ion may be used to indicate the neutralization of the solution.

3. The solution should be titrated as soon as the KI is dissolved in order
to avoid possible loss of iodine vapors from the flask.

4. If it is desired to avoid the use of KCN in P. 29, after dissolving the KI
add to the mixture 2.5 g of KSCN. Care should be taken to obtain an abso-
lutely clear filtrate when the mixture is later filtered.

If only the titration of copper is involved, it is recommended that the
thiocyanate be added just before the end-point is taken (see the discussion).

5. One accustomed to this titration can carry it to within 1 to 2 ml of the
end-point before adding the starch; however, the presence of the buff-
colored precipitate usually confuses one unfamiliar with the process. The
addition of the starch can be deferred and the final end-point approached
more closely by letting the mixture in the flask come to rest, and then holding
the tip of the buret close to the surface of the solution and noting if the next
drop of thiosulfate bleaches the solution locally as it is added. If it is
doubtful whether or not the end-point has been passed, it should be noted if
an additional millimeter of starch darkens the solution.

6. Unless Note 4 has been followed, it is not essential that a perfectly clear
filtrate be obtained, as a small amount of Cul will be held in solution in
P. 29 by the KCN added there. Should the iodine color reappear, it should
be removed by additional ^28263 before proceeding to P. 29; an excess of
the thiosulfate should not be added.

7. Upon standing in an acid solution for a considerable length of time,
the tetrathionate may decompose into, among other products, sulfurous
acid. This will cause a precipitate of sulfur when the solution is treated
with H 2 S in P. 29.

8. If only the determination of the copper in an ore is desired, proceed
as follows:

Weigh out a 0.3 to 0.4 g sample into a 200-ml flask, add 10 ml of
18 n. H 2 S0 4 and, slowly, 10 ml of 16 n. HNO 3 and 5 ml of 12 n. HCL
Cover the flask with a watch glass and heat the mixture so that the
reaction continues until decomposition of the ore appears complete
and sulfur is no longer present. Remove the watch glass and
evaporate just to fuming. Add 10 ml of 16 n. HNOa and 5 ml of
12 n. HC1 and evaporate to fuming. Cool the solution (Note 3,
P. 24), add 20 ml of water (Note 1, P. 24, second paragraph), and
then add NHiOH until its odor is distinct above the solution.
Add 1.5 g of ammonium acid fluoride, (NH4)HF2, and titrate as

P. 29] DETECTION OF CADMIUM 247

directed in the procedure above, beginning with the second sentence.
Add 2 g of KSCN just before the end-point.

P. 29. Detection of Cadmium

Discussion. Cadmium is detected and precipitated as the sulfide.
The formation of tho yellow precipitate is a very distinctive test
for this element, especially when the precipitation is made from an
ammoniacal solution, in which the sulfides of arsenic, antimony, or
tin are soluble. As copper sulfide is less soluble than cadmium
sulfide, copper has to be absent or its precipitation prevented. The
latter is accomplished in this procedure by carrying out the precipi-
tation in an alkaline solution to which cyanide has been added; the
complex ion Cu(CN) 3 " is so stable that the concentration of the
copper ions left in the solution is not sufficient to exceed the solubility
product of the sulfide. 24 Experiments have shown that 0.5 mg of
cadmium will form an easily visible precipitate, while a large quan-
tity of copper will remain in the solution to which NH 4 OH and
KCN have been added. It is an advantage to have removed most
of the copper in P. 28, for, although even 500 mg of this element
would not cause an immediate precipitate of cuprous sulfide upon
passing H 2 S into the solutiop, a red precipitate of hydrorubianic
acid, (CSNH 2 )2, is sometimes formed. Since ammonium sulfide
solutions usually have a yellowish color (because of the presence of
polysulfides), the test for cadmium is made by passing hydrogen
sulfide into the ammoniacal solution for a short time only. The
solution should not be saturated with hydrogen sulfide, as the large
amount of ammonium sulfide formed would be partially oxidized
to polysulfides and these would impart a yellow color to the solution,
and as the excess of hydrogen sulfide would convert the cyanide
into the slightly ionized hydrocyanic acid, thus permitting the pre-
cipitation of any copper present.

Cadmium sulfide, when precipitated from neutral or slightly acid
solutions, usually tends to be colloidal and is filtered with difficulty,
but, under the conditions of this procedure, separates in a form that
is readily coagulated.

Procedure 29: DETECTION OF CADMIUM. Make the
filtrate from P. 28 alkaline with NH 4 OH, and then add 5 ml

24 Bonner and Kaura, Chem. Met. Eng., 34, 84 (1927), and Glasstone, J.
Chem. Soc., 1929, 702, have shown that Cu(CN)j" is the predominant ion ob-
tained when cuprous cyanide is dissolved in an excess of cyanide.

248 ANALYSIS OF THE COPPER GROUP [P. 29

more (if P. 28 has been omitted, treat the ammoniacal fil-
trate from P. 26 as next directed). Add 2 ml of 1 n. KCN
(Caution: Note 1); filter, wash, and discard any precipitate
which forms, collecting the washings with the filtrate (Note
2). Pass a moderate current of H 2 S into the filtrate for 20
sec. (Yellow precipitate, presence of cadmium. Notes
3, 4.) If a yellow precipitate is obtained, add to the mix-
ture 5 ml of NH 4 OH and again pass a slow current of H 2 S
into it for 5 min. (do not saturate the solution). Shake the
mixture vigorously for 2 or 3 min., heat it to boiling (Note
5), and allow the precipitate to settle. Treat the mixture
by P. 30.

Notes:

1. KCN is a deadly poison. Solutions containing it should not be acidi-
fied unless under a well-ventilated hood. Care should be taken in the
disposal of such solutions that they are not inadvertently mixed with acids.

If, in order to avoid the use of cyanide, thiocyanate has been used in P.
28, an additional 5 ml of NH 4 OH should be added here in place of the KCN.

2. A precipitate forming upon addition of the ammonia is probably due
to lead or mercury, or to iron or aluminum introduced by reagents or appa-
ratus. As any of these, except aluminum, would cause a dark sulfide pre-
cipitate with H2S, they should be carefully filtered out, washed with a mini-
mum of water, and discarded.

3. Even 0.5 mg of cadmium will produce a distinct yellowish, opalescent
precipitate of CdS, so that, if no precipitate forms after passing in the H^S,
the solution may be discarded. The H^S should be passed into the solution
for only 20 sec., and under no conditions should the solution be saturated
with the gas. Occasionally a moderate amount of sulfur will be liberated
when H2S is passed into the alkaline solution; this may be due to incomplete
reduction of iodine by the ^28263 in P. 28 or to the presence of sulfite
resulting from the decomposition of thiosulfate or tetrathionate. This
sulfur redissolves, forming disulfide and imparting a yellowish color to the
solution which should not be mistaken for the yellowish turbidity caused
by small amounts of cadmium sulfide.

4. If a black precipitate forms, so that a small yellow precipitate also
present might be obscured, collect it on a small filter and transfer it to a
casserole with the aid of 25 ml of 1.2 n. H2S04. Boil the mixture gently for
10 min. Filter out any residue through a small paper filter, wash it, and dis-
card it. Collect the filtrate in a 200-ml flask, dilute it to 100 ml, and satu-
rate it with H2S. If cadmium is present, a yellow precipitate should be
obtained.

By this procedure, lead, mercury, or copper are discarded in the residue.
Iron sulfide is not precipitated from sulfuric acid solutions.

5. Except when only a small amount of cadmium is present, the precipi-
tate is effectively coagulated by boiling the mixture.

P. 30] ESTIMATION OF CADMIUM 249

P. 30. Estimation of Cadmium

Discussion. The amount of cadmium present is estimated by
treating the cadmium sulfide precipitate under suitable conditions
with an excess of ferric sulfate, with the result that an equivalent
amount of ferrous salt is produced, according to the reaction

2Fe + ++ + CdS (8) = 2Fe++ + Cd++ + S (i) .

The ferrous iron is titrated with permanganate, since the precipi-
tated sulfur does not react sufficiently rapidly with either the excess
of ferric sulfate or with permanganate to cause serious error. Al-
though cadmium can be detected and precipitated as the sulfide, the
precipitate is usually of uncertain composition. This is usually
stated to be due to the fact that it contains a considerable propor-
tion of double salts of the type CdCU-CdS; however, studies have
indicated that the effect is probably caused by the adsorption of
cadmium salts by the precipitate. 25 Because of this, it is not gen-
erally recommended that the precipitate be weighed or used as the
basis for volumetric methods. An experimental study 26 of the
method used here has indicated that, if the precipitation is made
from an alkaline solution in the presence of cyanide, the precipitate
is more uniform in composition; for this reason it is recommended
that cyanide be added in P. 29, even if copper is absent.

Before being treated with the ferric salt, the cadmium sulfide pre-
cipitate is heated and washed with a hydrosulfate-sulfate solution
in order to remove any adhering ammonium sulfide without at the
same time dissolving any of the cadmium. The "buffering" of a
solution by means of hydrosulfate and sulfate is discussed in P. 61.
Before titrating with the permanganate, phosphoric acid is added
to diminish the color caused by the large amount of ferric salt pres-
ent, thus making the end-point more distinct. Experiments have
shown that the results obtained are usually high, the error ranging
from 0.3 to 2 per cent; that complete washing of the precipitate is
necessary even with the hydrosulfate-sulfate treatment; and that
the greater the amount of cadmium present and the longer the
calcium sulfide is heated with the ferric sulfate the greater the error.

The precise determination of cadmium is usually effected either
(1) electrolytically, or (2) by precipitating the element from an acid
solution as the sulfide (after having separated it from other elements

11 Weiser and Durham, J. Phya. Chem., 32, 1061 (1928).
28 Unpublished experiments by F. N. Laird.

250 ANALYSIS OF THE COPPER GROUP [P. 30

that would be similarly precipitated), dissolving the sulfide in hydro-
chloric acid, evaporating this solution to dryness with an excess of
sulfuric acid, and heating to 500C. until the anhydrous cadmium
sulfate thus obtained is constant in weight.

Procedure 30: ESTIMATION OF CADMIUM. Filter the mix-
ture from P. 29 by decantation through an asbestos filter.
Sprinkle about 1 ml of solid Na2S04 on the precipitate and
wash it by decantation with a 50-ml portion of boiling water
(Note 1). Add to the precipitate in the flask 50 ml of a solu-
tion made by adding 3 g of solid Na2SO 4 and 1 ml of 6 n.
H 2 SO 4 to 100 ml of water, and heat the mixture almost to
boiling (Note 2). Again decant the solution and wash the
precipitate with the remainder of the sulfate solution. If
the precipitate does not coagulate readily, add solid NH 4 C1.

Transfer the precipitate and filter to a 400-ml flask and
add to it a solution made by dissolving 4 g of solid Fe 2 (S0 4 ) 3
in 50 ml of boiled water. Slowly heat the mixture to boil-
ing, add to it 3 ml of 6 n. H 2 S0 4 , and heat it on a water bath
until the CdS has dissolved and only a coagulated sulfur
precipitate remains (usually 5 to 10 min., depending upon
the size of the precipitate).

Immediately cool the solution, add to it 3 ml of 85 per cent
H 3 PO 4 and 10 ml of 6 n. H 2 SO 4 , and dilute it with 150 ml of
water. Titrate the solution with standard KMn0 4 solu-
tion, stirring continuously, until a drop of the perman-
ganate causes a pink color to persist throughout the solution
for 15 to 20 sec. (Note 3). From the volume of KMn0 4
used, calculate the amount of cadmium present (Note 4).

Notes :

1. Adding the Na 2 SO 4 and heating the mixture for 1 to 2 min. usually
causes the precipitate to coagulate, so that after 2 to 5 min. most of the solu-
tion can be readily decanted. If the precipitate is small, most of it may have
to be caught on the filter.

2. The precipitate is treated with H 2 S0 4 -Na 2 S0 4 solution in order that
any adsorbed sulfide may be extracted from it and expelled as H 2 S. The
hydrogen ion concentration of the hydrosulfate-sulfate solution is so low
that there is no appreciable solution of the CdS upon expelling the H 2 S.

3. Although the end-point is not permanent, it is easily noted; the per-
manganate is rapidly decolorized, and the entire solution is not colored until

the end-point is reached, especially if the solution is continuously stirred.

4. The Fe 2 (SO 4 ) 3 should be tested for ferrous iron by titrating 4 g of it
with KMn0 4 under the conditions of this procedure.

The Analysis of the Tin Group

General discussion of methods for the analysis of the Tin Group.

The satisfactory separation of the elements of the Tin Group when
these are present in moderately large amounts was found to offer
considerable difficulty. Some systems of qualitative analysis at-
tempt the separation of antimony and tin from arsenic and mercury
by treating the precipitate of the Tin Group sulfides with 12 n.
hydrochloric acid. By such treatment antimony and tin sulfides,
when present alone, can be dissolved, while, if the solution is kept
saturated with hydrogen sulfide, arsenic sulfide is insoluble. How-
ever, it has also been found that considerable amounts of mercury
are dissolved by the concentrated hydrochloric acid, and that, if a
mixture containing considerable amounts of these sulfides is treated,
the extraction of antimony from the arsenic sulfide is incomplete.
Attempts were made to separate mercury from the other elements
of this group by adding an excess of an ammonium salt to the sodium
sulfide solution of the Tin Group sulfides. This causes the precipi-
tation of mercuric sulfide, which, as was mentioned in the discussion
of P. 12, is insoluble in an ammonium sulfide reagent. It was found
that by this means even small amounts of mercury could be detected
and that its precipitation was complete when it was present alone
or even when it occurred with arsenic or antimony. When tin was
present in even moderate amounts, the precipitation of mercury
was incomplete; also, the precipitate contained a considerable amount
of tin and was so colloidal that filtration was often impossible. Be-
cause of these facts the above methods are not used in this system of
analysis.

As a result of an extensive investigation by Dr. Chester E. Wilson,
a process has been developed in which the sulfide precipitate is
dissolved in a hydrochloric acid solution with the aid of bromate as
an oxidizing agent; the arsenic is then reduced and distilled as the
relatively volatile trichloride, the mercury is reduced to the metallic
form by means of phosphorous acid, and the antimony is then pre-
cipitated as the sulfide from a solution in which the acid and chloride
concentrations are carefully controlled.

The tin is precipitated as sulfide by making the solution alkaline,
passing in hydrogen sulfide, and then making the solution again

251

TABULAR OUTLINE V
THE ANALYSIS OF THE TIN GROUP

(Method for the More Quantitative Separation and Estimation of the
Elements of This Group)

Precipitate of the Tin Group Sulfides (from P. 13): As 2 S 6 , HgS, Sb 2 S 4 , SnS 2 , S
Treat with HCl and KBrOi. (P. 41)

Residue:

Solution: H 3 AsO 4 , H 2 HgCl 4 , HSbCU, H 2 SnCl 6 , Br 2

Add HBr and H,PO 3 . Heat.

Solution: AsCl 3 , HSbCl 4 , H 2 HgCl 4 , H 2 SnCl 6 , Br~, H 3 PO 3 , H*PO 4

Distill. Collect distillate in water.

Distillate:

AsCl 3 , HCl
Neutralize with
NaliCOz

Titrate with KI Z
solution, starch
indicator.

(H 3 As0 4 , I")
(P. 43)

Solution: H 2 HgCl 4 , HSbCl 4 , H 2 SnCl 6 , BIT, H 3 PO 3 , H 3 PO 4
Add HzSO* and dilute. Heat. (P. 44)

Precipitate:
Hg 2 Cl 2 , Hg
Treat with
Na^HPOa and
NaOH.
(Hg,0, Hg)

Dissolve in
UNO,.

Add Fe(NOi)*,
titrate with
KSCN.
Solution:
Hg(SCN),,
un-ionized;
Fe(SCN) 6 -
(red)
(P. 45)

Solution: HSbCl 4 , H 2 SnCl 6 , H 3 PO 4 , H 3 PO 8 ,
HSOr, Br-, H+CI-.
Dilute to 75 ml. Heat. Saturate with //j*S r .
(P. 46)

Precipitate:
Sb 2 S 3
Dissolve in HCl.
Add tartaric
acid.
Neutralize.
(HSbOC 4 H 4 6 )

Add NaHCO*,
titrate with K/ 3
solution, starch
indicator
Solution:
HSbO 2 C 4 H 4 O 6 ,
I- (P. 47)

Solution: H 2 SnCl 6 , etc.
Neutralize, pass in 7/ 2 >S',
make just acid.

Precipitate: Filtrate:
SnS 2 Discard.

Dissolve in HCl.
Heat with Ni foil.
(H 2 SnCl 4 , NiCU)

Cool, dilute, add starch,
titrate with Kit solution.
Solution: H 2 SnCl 6 , I~
(P. 49)

252

GENERAL DISCUSSION

253

acid. It was found that, because of the phosphoric and phosphorous
acids present, the precipitation of the tin by saturating even a very
slightly acid solution with hydrogen sulfide was very slow; by form-
ing the sulfo-salt in an alkaline solution and then acidifying, rapid
precipitation is obtained. The details of these procedures are shown
in Tabular Outline V; the results obtained from a series of test
analyses 1 are shown in Table XVII.

TABLE XVII

TEST ANALYSES OF THE METHOD FOR THE ANALYSIS OF THE
TIN GROUP ELEMENTS

Amount of Each Element Taken and Found (mg) ; S Denotes a Satisfactory

Detection

Arsenic

Mercury

Antimony

Tin

Experi-

ment

Taken

Found

Taken

Found

Taken

Found

Taken

Found

500

490

500

481

500

499

505

504

250

245

250

232

253

251

255

220

* Twenty mg of tin were found in the sulfur residue (P. 41); probably as a
result of too rapid addition of KBr0 3 in that procedure.

However, this method for the more quantitative separation of
these elements was found to require such an expenditure of time
and the use of such special technique that it was decided to offer
a simpler and more rapid optional method for the analysis of this
group. This method can be used when there is question as to
whether or not any Tin Group elements are present in the precipi-
tate obtained on acidifying the sodium sulfide solution, or when
only small amounts (50 mg or less) of these elements are present and
only a rapid qualitative analysis is desired. When larger amounts
are present and when a more quantitative separation and estimation
of these elements is required, the use of the longer but more exact
method is recommended.

1 These test analyses were part of the investigation carried out by Dr.
Chester E. Wilson.

254 ANALYSIS OF THE TIN GROUP [P. 4L4

TABULAR OUTLINE V-A

OPTIONAL METHOD FOR THE ANALYSIS OP THE TIN GROUP
(For Use When Only a Rapid Qualitative Analysis Is Desired)

Tin Group Precipitate (from P. 13): HgS, As 2 Si, SbaS*, SnS 2 , S
Treat with NH 4 OH and

Residue: HgS, S I Solution: AsS 4 ", SbS 4 a , SnSr, HS~ HS 2 - NH 4 OH
-- _ -- 1 Heat with 6 n. HCl.

Residue: As2S&, S

Solution: HSbCU, H 2 SnCl 6
Dilute, treat with H^S.

Precipitate: Sb 2 S Solution: H 2 SnCl 6
Partly neutralize, treat with

Precipitate: SnS 2 Solution: Discard.

P. 414. Optional Method for the Analysis of the Tin Group (For
Use When Small Amounts of the Tin Group Elements Are
Present and Only a Rapid Qualitative Analysis
Is Desired)

Discussion. In this method, which is shown in detail in Tabular
Outline V-A, the sulfide precipitate (from P. 13) is treated with
ammonium hydroxide and a small amount of an ammonium sulfide
reagent. These dissolve the sulfides of arsenic, antimony, and tin
but leave mercuric sulfide (and any Copper Group sulfides) in the
residue, where a confirmatory test for the presence of mercury can
be made. Only a small amount of the ammonium sulfide is used
in order not to dissolve much of the sulfur; thus, upon acidifying
this ammonium sulfide solution the detection of the presence of any
Tin Group sulfides is made more certain. By then adding an equal
volume of 12 n. hydrochloric acid to this acid mixture and heating,
most of any antimony and tin will be dissolved, leaving arsenic sul-
fide in the residue. The acid concentration of the filtrate is then
suitably adjusted, and the filtrate is treated with hydrogen sulfide,
which precipitates the antimony as sulfide. Thus, only tin is left
in the filtrate, from which it is precipitated as sulfide by partial
neutralization of the acid.

Although this method gives very imperfect separations when con-
siderable amounts (50 mg or more) of these elements are present, it
is recommended for use with smaller amounts for the following
reasons: It is very much more rapidly and easily carried out than

P. 414] OPTIONAL METHOD 255

are the more exact procedures; it provides an even more sensitive
detection of traces of these elements; and, by comparison of the
precipitates obtained with known amounts of the element precipi-
tated under similar conditions, an approximate estimate of the
amount of the elements present can be made.

Procedure 414: OPTIONAL METHOD FOR THE ANALYSIS
OF THE TIN GROUP. Transfer the sulfide precipitate (from
P. 13) to a casserole, add to it 10 ml of NH 4 OH and 2 ml of
ammonium monosulfide reagent, and heat the mixture
to 60 to 70C.; disintegrate any residue. (Dark resi-
due, presence of mercury or of Copper Group sulfides.)
Filter out the residue, and, if it is darkened, treat it as
directed in Note 1, P. 45, to confirm the presence of mer-
cury.

Collect the filtrate in a 100-ml flask (do not add any wash
water to it) and boil it vigorously for 5 min. (Note 1).
Acidify the solution with 12 n. HC1, adding 1 ml in excess.
(Colored precipitate, presence of arsenic, antimony, or tin.)

If a colored precipitate is obtained, transfer the mixture
to a graduated cylinder, add an equal volume of 12 n. HC1,
return the mixture to the flask, and heat it just to boiling for
10 min. (Note 2). Pass a few bubbles of H 2 S into the mix-
ture, again heat it to boiling, and filter the hot solution
through a small paper filter which has been moistened with
6 n. HC1. (Yellow residue, presence of arsenic.) Do not
add any wash water to the filtrate. Treat the residue by
Note 9, P. 42, to confirm the presence of arsenic.

Dilute the filtrate to just three times its volume with
water, heat it almost to boiling, and again saturate it with
H 2 S. (Orange precipitate, presence of antimony.) Filter
out the precipitate (Note 1, P. 47).

Add to the filtrate a volume of 15 n. NH 4 OH just two-
thirds that of 12 n. HC1 originally added (the solution
should be acid), again saturate the hot solution with H 2 S,
close the flask, and allow it to stand for 15 min. (Yellow
precipitate, presence of tin. Note 2, P. 48.)

Notes:

1. The solution is boiled in order to expel most of the ammonia and thus
reduce the volume of acid required to make it acid.

256 ANALYSIS OF THE TIN GROUP [P. 41

2. By adding an equal volume of 12 n. HC1, the solution is made approxi-
mately 6 n. in hydrochloric acid. This is the composition of the constant-
boiling ^solution, so that its concentration is not changed by boiling. This
is important as it permits a close adjustment of the acid for the subsequent
separation of antimony from tin.

P. 41. Solution of the Tin Group Sulfides

Discussion. In the more exact method for the analysis of the
Tin Group elements, the first step involves the complete solution of
the sulfide precipitates. In the procedure below, the antimony and
tin sulfides are dissolved by concentrated hydrochloric acid alone,
and then arsenic and mercury sulfides are brought into solution by
the addition of an oxidizing agent, such as potassium bromate (potas-
sium chlorate or nitric acid may also be used). Bromate is used in
preference to chlorate as it was found to be much more effective in
dissolving the sulfur which is always precipitated upon acidifying
the disulfide reagent. The use of liquid bromine was found to be
equally effective in this respect, but it formed large amounts of
sulfur monobromide, S2Br 2 , a red liquid which retained considerable
amounts of arsenic. This was probably due to the formation of a
compound containing arsenic, sulfur, and bromine, analogous to the
chlorine compound As4S5Cl2, which is produced under similar condi-
tions. When bromate is used and when a large amount of arsenic
is present, a reddish residue containing a few milligrams of arsenic
will frequently be left. Usually, after the treatment with bromate,
there remains only a small residue, which is composed mostly of
elementary sulfur with small amounts of undissolved sulfides and of
sulfur monobromide enclosed within it. Experiments have shown
that when 500 mg of each element of the Tin Group as the sulfide
is treated separately by this procedure, there remain in these resi-
dues less than 2 mg of mercury, 3 mg of arsenic, not over ^ mg of
antimony, and an insignificant amount of tin.

Although arsenic trichloride is readily volatilized from hot hydro-
chloric acid solutions, the use of the bromate, which is required to
dissolve arsenic sulfide, insures the complete oxidation of this ele-
ment to the quinque-positive state. The solution cannot be boiled
or unduly heated, as quinque-positive arsenic is reduced by the
bromide resulting from the reduction of the bromate and would then
volatilize.

Procedure 41: SOLUTION OF THE TIN GROUP SULFIDES.
Transfer the sulfide precipitate (from P. 13) to a 100-ml

P. 41] SOLUTION OF TIN GROUP SULFIDES 257

beaker with the aid of 20 ml of 12 n. HC1 (Note 1). Cover
the beaker with a glass and warm the mixture (do not boil)
for 2 to 3 min. (or longer if the precipitate seems to be dis-
solving). Disintegrate the residue frequently and thor-
oughly with a stirring rod. If the precipitate is not all
dissolved, cool the beaker and add to the mixture, in 0.1-ml
portions, solid KBrOa (Note 2) until the sulfides appear to
be dissolved and a permanent orange color is obtained (Note
3). Carefully disintegrate the residue before each addition
of the K3r0 3 (Note 4), and add a fresh portion only after
the previous one stops reacting.

Cool the mixture and filter it through an asbestos filter
which has been moistened with 12 n. HC1 and drained, col-
lecting the filtrate in a 200-ml distilling flask which has also
been washed with 12 n. HC1 (Notes 5, 6). Wash the beaker
and residue with 10 ml of 12 n. HC1, added in small por-
tions, and collect these washings with the filtrate. Treat
the filtrate by P. 42. Discard the residue.

Notes:

1. Only 12 n. HC1 (and no water) should be used to aid this transfer, as
the later distillation of the arsenic depends upon having concentrated acid
present. The precipitate is most easily transferred by inserting a thin glass
rod through the stem of the funnel and pushing the perforated plate and
asbestos into the beaker; any precipitate adhering to the sides can be washed
through the funnel with the acid. A beaker is used in preference to a cas-
serole to avoid the danger of loss by spattering during the addition of the
bromate.

If a paper filter has been used and only a small amount of the precipitate
has been carried onto the filter, this precipitate may be transferred by push-
ing a hole through the paper with a stirring rod and then washing it off the
paper and through the funnel with the HC1. It is also possible to open the
filter against the side of a larger funnel and wash the precipitate from it by
squirting the HC1 repeatedly against it with a dropper. The filter and fun-
nel can then be washed with the HC1. If a large amount of precipitate has
been carried onto the filter, it is usually better to transfer as much of the
precipitate to the beaker as is possible with a stirring rod or porcelain spat-
ula; the precipitate adhering to the paper can then be treated as suggested
above.

2. As long as the precipitate appears to be dissolving in the HC1 alone,
the KBrOa should not be added, as it causes the liberation of considerable
sulfur; this mechanically encloses undissolved particles of the sulfides, which
then dissolve more slowly. Of the sulfides of the Tin Group, antimony and
tin dissolve readily in concentrated HC1 alone, mercury dissolves to a slight
extent, and arsenic dissolves scarcely at all. When treated with HC1 under

258 ANALYSIS OF THE TIN GROUP [P. 42

certain conditions, mercuric sulfide changes into a white residue which still
contains mercury; a dark precipitate changing in this way should be further
treated with KBr0 3 .

3. When properly used, 2 g of the bromate are adequate to dissolve 500
mg of arsenic sulfide, even when the latter is present with the sulfur resulting
from the using of the maximum amount of the disulfide reagent. When the
sulfides have been dissolved, the addition of bromate will give a permanent
orange color to the solution; a large excess of the bromate should be avoided,
as it oxidizes the phosphorous acid later to be added.

4. When a large sulfide precipitate is treated, there usually remains a
small dark or colored residue of sulfur enclosing small amounts of undis-
solved precipitate; after being treated as directed, these residues generally
contain insignificant amounts of antimony or tin and not over 2 mg of mer-
cury or 3 mg of arsenic, even when these elements are present in very large
quantities. If the KBrO 3 is added rapidly, if the mixture is boiled, or if the
residue is not continually kept broken up, larger amounts of the sulfides
may remain uridissolved. Therefore, if it is desired to recover this amount
of these elements, the solution should be completely decanted, the residue
should be heated with 1 to 2 ml of 6 n. NaOH and evaporated almost to
dryness, and 5 ml of 12 n. HC1 and then sufficient KClOa to dissolve the
sulfide present should be added. After this solution has been boiled to
expel any C\2 left, it can be filtered and collected with the main filtrate or
can be treated separately. In this treatment the sulfur is dissolved by the
hydroxide, forming sulfide and thiosulfate, and any Tin Group elements
again dissolve as the sulfo-salts. Upon acidifying the solution, the sulfides
are reprecipitated and it can be noted if an appreciable amount of arsenic
or mercury is present. If there is, the precipitate is then in a form which
readily dissolves on treatment with KC10 3 .

5. The filter and flask are moistened with HC1 to prevent the precipita-
tion of antimonyl chloride when large quantities of antimony are present,
and to prevent dilution of the solution before the distillation of arsenic.

6. The distilling flask should be of resistance glass and should preferably
have the side arm bent as shown in Fig. 29.

P. 42. Separation and Detection of Arsenic

Discussion. After a solution of the Tin Group sulfides has been
obtained by means of the treatment with hydrochloric acid and
bromate, methods for separating the individual elements are to be
considered. Arsenic can be precipitated as its sulfides from a solu-
tion 9 to 12 n. in hydrochloric acid, and a quantitative separation
from antimony and tin can thus be obtained. It was found, how-
ever, that mercury, when present in considerable quantities, would
divide between the precipitate and the solution; if the acid concen-
tration was reduced sufficiently to precipitate the mercury completely
(to less than 6 n.), antimony would be coprecipitated by the arsenic
sulfide precipitate.

P. 42] DETECTION OF ARSENIC 259

A method often used for the quantitative separation of arsenic
from all the other common basic elements consists in the distillation
of the trichloride from a hydrochloric or hydrobromic acid solution.
This method can be advantageously applied at this point, as the dis-
tillation is most effectively made from a concentrated acid solution;
however, as all of the Tin Group elements are present in their higher
oxidation states, a means of reducing the arsenic has to be provided.
This could be done by the addition of considerable bromide to the
solution, but the rate of the reduction is somewhat slow, and the
bromine which is evolved reoxidizes the arsenic in the distillate. By
the addition of phosphorous acid, which is to be used later to reduce
the mercury, the liberation of the bromine is prevented, because it
is reduced by the phosphorous acid as fast as it is formed. Phos-
phorous acid alone would not serve the purpose, as the reduction of
arsenic acid by phosphorous acid proceeds at an extremely slow rate.

In the procedure which has been adopted, 7 ml of 6 f . phosphorous
acid and 1 ml of 9 f. hydrobromic acid are added to the 30 ml of
12 f. hydrochloric acid, and the resulting solution is distilled to a
volume of 15 ml. In the development of this process it was found
that, with 500 mg of arsenic present, about 90 per cent of it passes
over with the first 10 ml of distillate, and not over 1 mg remains
after the distillation. When two 5-ml portions of hydrochloric acid
are subsequently added and distilled, the amount of arsenic remaining
(approximately 0.1 mg) is so small that it does not interfere with the
detection of antimony. Under these same conditions, with 500 mg
of each of the other elements, 10 to 12 mg of mercury, 3 to 5mgof
antimony, and only a trace of tin are distilled. When the distillation
is carried to 7 ml, an additional 4 to 6 mg of mercury are distilled,
but only a trace of either antimony or tin is distilled; when it is con-
tinued until only 5 ml remain, 6 to 8 mg more of mercury, but no
appreciable amount of antimony or tin, come over.

From the fact that the boiling point of arsenious chloride is 130C.,
while that of stannic chloride is 114C., it is obvious that the be-
havior of these elements in this distillation is not due entirely to a
difference in the volatility of the various chlorides. However,
arsenious chloride exists as such to a large extent in hydrochloric
acid solutions, while stannic chloride is converted into the very
stable chlorostannate ion, SnCl 6 ". Quinque-positive arsenic does
not volatilize, because it is largely hydrolyzed to arsenic acid; the
antimony chlorides are likewise to a large extent converted into
complex ions or hydrolyzed. Mercuric chloride is not appreciably

260

ANALYSIS OF THE TIN GROUP

IP. 42

ionized in aqueous solutions but forms a complex ion, HgCU", with
an excess of chloride ion; the mercury which distills over does so
as metallic mercury because of a slight reduction by the phosphorous
acid.

Procedure 42: SEPARATION AND DETECTION OF ARSENIC.

Add to the solution in the distilling flask 7 ml of 6 f . H 3 P0 3

% and 1 ml of 9 f. HBr (Note 1). Close the end of the side-arm

outlet tube with
a short piece
of rubber tubing
fitted with a
clamp or glass
bead, and in-
sert in the neck
of the flask a
closely fitting
test tube filled
with cold water
and wet with
12n.HClpnthe
outside. Heat
the solution to
80 to 90C.
until any bro-
mine color dis-
appears (Note
2), and then
cool the solu-
tion. Remove
the rubber tub-
ing and the test
tube, rinsing
HC1, and clamp

0.5 mm

Fig. 29.

Apparatus for the Distillation of
Arsenic Trichloride.

the latter with 1 to 2 ml of 12 n.
the flask in position (see Fig. 29) with the side arm extend-
ing to the bottom of a 100-ml volumetric flask containing
40 ml of cold water and immersed in a large beaker of ice
water. Insert an "ebullition tube" (Note 3) and close the
distilling flask with a cork stopper carrying a glass tube 7
mm in diameter and about 70 cm long which extends to the
bottom of the flask. Heat the solution and distill until
just 15 ml of solution remain in the flask (Notes 4, 5).

P. 42] DETECTION OF ARSENIC 261

Add 10 ml of 12 f . HC1 to the distilling flask through the
safety tube by means of a dropper and again distill until
only 15 ml of solution remain (Note 4).

Remove the flask with the distillate, rinsing the delivery
tube, cool to room temperature, dilute with water to the
mark on the flask, and thoroughly mix the contents (Note 6).

Pipet 25 ml of the distillate into a 200-ml flask, add 10 ml
of 12 n. HC1, saturate the solution with H 2 S, and then heat
it almost to boiling while continuing the flow of H 2 S. (Yel-
low precipitate, presence of arsenic, Notes 7, 8, 9, 10.)

If arsenic is present, treat the remainder of the distillate
by P. 43.

Treat the residual solution in the distilling flask by P. 44.

Notes:

1. Hydrobromic acid can be obtained as the 48 per cent solution, which
is 9 f., or as the 42 per cent solution, which is 7.2 f. If only the less concen-
trated solution is available, an equivalent amount should be added.

2. After adding the hydrobromic acid, there will probably be a yellowish
color due to bromine in the solution, because in the concentrated acid both
arsenic and antimonic acids are reduced by bromide. If the distillation
were started immediately, this bromine would pass into the distillate and
would reoxidize the arsenic, introducing an error in the subsequent estima-
tion of arsenic; therefore the solution is heated until the bromine is reduced
by the phosphorous acid. To prevent loss of arsenic, a test tube of cold
water is inserted in the neck of the distilling flask. When antimony is pres-
ent, a slight yellowish color may remain after the bromine has been reduced;
otherwise the solution becomes colorless. It has been found that, when the
solution is heated to between 80 and 85C., the arsenic and bromine are
completely reduced in less than 5 min. Care should be taken not to boil
the solution, or arsenious chloride will be lost.

3. "Bumping" and irregular boiling of the solution during the distilla-
tion can be minimized by the use of an "ebullition tube" in the flask. Such
tubes are made as follows : Heat a piece of soft glass tubing of 8 to 10 mm
in diameter and draw from it a long capillary of approximately 0.5 mm in
diameter. Cut this capillary into lengths somewhat shorter than the height
of the distilling flask. Heat one end of each of these lengths until it seals,
and then heat at a point about 1 cm from the other end until the capillary
constricts and closes, thus leaving a capillary space about 1 cm long in that
end.

"Bumping" is caused by the lower portion of the solution superheating
and not maintaining an equilibrium with its vapor phase, until, when finally
gas formation does start, there is a sudden formation of a large volume of
vapor. The ebullition tube is placed in an upright position in the flask
with the open end down and just above the heating flame. When the solu-
tion is heated, the air is expelled from the capillary, the tube fills with the
vapors of the liquid, and thus, as long as the solution is boiled, a vapor phase

262 ANALYSIS OF THE TIN GROUP [P. 42

is kept in contact with the hottest part of the solution; superheating is thus
minimized. It is desirable that the solution should not completely stop
boiling at any time during the distillation, because, if the vapor condenses
and the tube completely fills with liquid, it ceases to function.

A supply of these tubes can be readily made and kept available in the
laboratory.

4. In order to avoid distilling over considerable amounts of mercury,
the solution should not be distilled below 10 ml. Furthermore, in order
that an exact adjustment of the acid present in the subsequent procedures
may be made, the amount of solution remaining in the flask should be just
15 ml. This volume is best estimated by pouring 15 ml of water into a
similar flask and making a direct comparison.

5. To avoid possibility of loss through the safety tube, it should be occa-
sionally cooled during the distillation by adding 0.1 to 0.2 ml of 12 n. HC1
to it with a dropper.

6. A small gray or black precipitate floating on the surface of the dis-
tillate or collected on the bottom of the flask is probably mercury. In this
case, before the solution is diluted, it should be filtered through a tightly
packed asbestos filter which has been washed with 6 n. HC1 and the precipi-
tate and filter should then be washed with 3 to 5 ml of 6 n. HC1 added drop-
wise. This filter should be reserved and the mercury obtained in P. 44
should be collected on it. The HC1 is used on the filter and in washing to
prevent the precipitation of antimony.

7. As the arsenic may be present in the quinque-positive state, the mix-
ture is heated to boiling in order to hasten the formation of the precipitate
and to coagulate the precipitate more effectively.

8. If no precipitate (or only a very small one) is obtained, the remainder
of the distillate should be added to the flask, a corresponding volume of 12 n.
HC1 should be added, and the solution should be treated with H2S. The
amount of arsenic in any precipitate thus obtained can be estimated by a
comparison with standard precipitates.

9. If desired, the presence of arsenic in the precipitate may be further
confirmed as follows:

Filter and wash the yellow precipitate. Transfer it to a small
casserole with the aid of 10 ml of NH 4 OH. Evaporate the ammonia
solution almost to dryness, add to it 5 ml of 16 n. HNOs, and again
evaporate just to dryness. Treat with 5 ml of 3 n. HNOa, and then
filter out and discard any insoluble residue. Make the solution
alkaline with NH^OH, add 1 ml excess, and then add 5 ml of 1 n.
Mg(NOs)2 reagent. (White crystalline precipitate, presence of
arsenic.)

The precipitate so obtained is MgNHUAsC^. For a discussion of its
properties, see P. 88.

The precipitate may be further confirmed by collecting it on a small
filter, washing it with 1 to 2 ml of cold water added dropwise, and then
pouring dropwise over it a solution made by adding 5 drops of HC2Hs02
to 15 drops of AgNOs. (Red precipitate, presence of arsenic.) By this
treatment the MgNH 4 As04 is metathesized to the less soluble Ag 3 As04,
which is brick red in color.

P. 43] ESTIMATION OF ARSENIC 263

10. Any antimony which may have distilled over may be recovered and
estimated approximately by filtering out the arsenic precipitate, diluting
the filtrate to 100 ml, adding 10 ml of 15 n. NEUOH, and again saturating
it with H2S. An orange precipitate shows the presence of antimony; the
amount present can be estimated by comparison with a standard precipitate.

P. 43. Estimation of Arsenic

Discussion. The method used here for the estimation of arsenic
depends upon the oxidation of arsenious acid to arsenic acid in a
neutral solution by means of a standard iodine solution. This
method was discussed in detail in P. XII, p. 71, and was used there
for standardizing an iodine solution. It was mentioned there that
this reaction can be caused to proceed quantitatively in either direc-
tion by the proper adjustment of the hydrogen ion concentration,
and the conditions for using the reverse reaction are treated in detail
in the discussion of P. 89. As, in spite of the precautions taken in
P. 41, traces of bromine may pass over into the distillate and oxidize
an equivalent amount of arsenic, the reverse reaction is utilized here.
Any quinque-positive arsenic is reduced before the titration with
iodine by adding a small amount of iodide to the strongly acid
solution; the iodine thus set free is reduced with thiosulfate, an
excess being avoided, and the solution is then neutralized and buf-
fered for the titration with standard iodine solution.

Procedure 43: ESTIMATION OF ARSENIC. Pipet 25 ml of
the distillate (from P. 42) into a 200-ml flask. Add to it 10
ml of 6 n. HC1 and, in 0.1 g portions, 1 g of solid NaHCOs.
Add just 0.5 g of solid KI, close the flask with a clean rubber
stopper, swirl the mixture until the KI dissolves, and allow it
to stand for 5 min. (Note 1). If there is a perceptible yel-
low color, carefully add 0.1 n. Na2S20a, a drop or half-drop
at a time, until the solution becomes colorless (Note 2).

Dilute the solution at once to 50 ml, add to it 3 drops of
phenolphthalein indicator solution and then sodium hy-
droxide until a pink color is obtained (Note 3). Add HC1
until the pink color disappears, and then add 1 ml in excess.
Cool the solution, add solid NaHCO 3 , 0.2 g at a time, until
rapid effervescence no longer occurs (about 0.5 g), and then
add 3 g more (Note 4). Add 5 ml of starch indicator and
titrate with standard iodine solution, swirling the mixture
gently but not vigorously shaking it (Note 7, P. XII), until
the first permanent blue (or purplish, Note 5) color is ob-

264 ANALYSIS OF THE TIN GROUP [P. 44

tained which persists for 30 sec. (Notes 8, 9, P. XII). From
the volume of standard iodine used calculate the amount
of arsenic present.

Notes:

1. NaHCOs is added to provide an atmosphere of CO* in the flask and
thus prevent possible oxidation of iodide by the oxygen of the air. If too
much KI is added, or if the HC1 is too concentrated, a precipitate of yellowish
arsenious iodide, Asia, may separate. As this will obscure the end-point,
it should be dissolved by adding water, 5 ml at a time, until a clear colorless
or yellow solution is obtained.

2. It is not necessary to note the amount of thiosulfate which is added,
but great care must be taken not to add an excess, as it will cause a corre-
sponding error in the subsequent titration. Not more than 5 to 10 drops
should ever be required to decolorize the solution. Starch is not necessary
in the small volume of solution used here and is not satisfactory as an indi-
cator in strongly acid solutions, 4 n. or greater.

If large amounts of antimony are present, traces may be found in the
distillate, imparting a slight yellowish color to the solution and causing the
disappearance of the iodine color to be difficult to detect. In this case
add 5 ml of CCU to the mixture and titrate until the CCU layer loses its
iodine color.

3. The solution is neutralized with NaOH in order to avoid using a large
amount of NaHCOs, with the attendant danger of loss by spattering.

4. The following titration should be carried out in a flask (not in an open
beaker) and the solution should be quite cold in order to prevent the loss of
COj. If this occurs, the solution becomes more alkaline and, due to the
reaction between iodine and hydroxyl ion, the end-point will fade. For
very precise titrations it is advisable to pass a slow stream of COj continu-
ously through the solution.

5. Certain starch solutions, especially in hydrocarbonate solutions, give
an intermediate purplish color with iodine.

P. 44. Precipitation of Mercury

Discussion. After arsenic has been distilled from the solution,
the mercury is separated from antimony and tin by reduction with
phosphorous acid; mercurous chloride and, finally, metallic mer-
cury are formed.

In a cold dilute hydrochloric acid solution mercuric chloride is
reduced only to a mercurous chloride by phosphorous acid, and the
complete precipitation requires at least several hours. The reac-
tion can be represented as follows:

2HgCir + H 3 P0 3 + H 2 = Hg 2 Cl 2(8 ) + H 3 P0 4 + 2H+ + 6C1~.
A gravimetric determination of mercury is often made by collecting

P. 44] PRECIPITATION OF MERCURY 265

the mercurous chloride precipitate, drying it at 100 to 110C., and
weighing it.

However, if a large excess of phosphorous acid is provided and
the solution is heated as directed in this procedure, 0.5 mg of mer-
cury will cause a white precipitate of mercurous chloride within 5
min. ; the reduction then proceeds to the formation of metallic mer-
cury, and the precipitation of even 500 mg is complete within 10
min. As would be predicted from the above equation, the precipi-
tation of mercury by phosphorous acid is markedly influenced by
the concentration of the hydrochloric acid present; thus no appre-
ciable precipitation takes place during the distillation of the arsenic,
but, by diluting the residual solution, complete precipitation can be
obtained. If the concentration of the hydrochloric acid is much
greater than that provided for, approximately 1.5 n., the time re-
quired for the precipitation of the mercury is longer.

In order to prevent the precipitation of antimony as the oxychlo-
ride, the residual solution from the distillation is diluted with a sul-
furic acid solution, and some additional hydrochloric acid is added
to prevent the precipitation of tin when it is present in large amounts.
The solution is heated on a water bath during the precipitation so
as to avoid excessive loss of the acid by evaporation and loss of mer-
cury by volatilization.

Procedure 44: PRECIPITATION OF MERCURY. Cool the
residual solution (from P. 42), and transfer it from the dis-
tilling flask to a 200-ml flask with the aid of a solution made
by adding just 20 ml of 6 n. H 2 S0 4 and 5 ml of 6 n. HC1 to 25
ml of water (Notes 1,2). Add 1 ml of 6 f . H 3 P0 3 , cov.er the
flask with a watch glass, heat the solution just to boiling,
and immerse it for 5 min. in a beaker of boiling water.
(White precipitate, turning to black, presence of mercury.
Note 3.)

If no precipitate forms, add to the solution just 5 ml of 6
n. HC1 and 5 ml of water and treat it by P. 46.

If a precipitate forms, heat the mixture in the boiling
water for 5 min. more, remove it, heat it just to boiling for
15 to 30 sec., and allow the precipitate to settle (Note 4).
Decant the solution through a compact asbestos filter' (Note
5) which has been moistened with 6 n. HC1 and drained.
Wash the precipitate and filter by decantation with just 10
ml of 3 n. HC1 aded in 2-ml portions (Note 6). Collect

266 ANALYSIS OF THE TIN GROUP IP. 44

these washings with the filtrate in a 200-ml flask. Treat
the precipitate by P. 45 (Note 7). Treat the filtrate by
P. 46.

Notes:

1. The solution is cooled to avoid the precipitation of mercury in the
distilling flask when it is diluted. Even so, with large amounts of mercury
a small precipitate may be formed, and extreme care must be taken that it
is all transferred to the conical flask. The safety tube should be carefully
washed.

2. Sulfuric acid has to be added with the water in order to avoid the pre-
cipitation of antimonyl chloride, SbOCl. A larger amount of HC1 would
cause the mercury to precipitate slowly. An additional amount of HsPOa
is next added in case most of that added in P. 42 has been oxidized by a
large excess of bromate.

3. In order not to overlook a small precipitate of mercury, one should
examine the solution in a bright light for finely divided suspended particles.
Then, after being heated, the solution should be swirled, and the bottom of
the flask should be carefully searched for small particles which should be
collected near the center.

4. The heating of the solution and especially the final boiling cause the
crystalline white precipitate of Hg2Cl2 which first forms to be changed into
finely divided metallic mercury; this coagulates into compact black particles
which quickly settle out or even coalesce into droplets of the metal. Ac-
cordingly, the solution can be quite completely decanted, and this precipi-
tate can be washed with a small volume of solution. Care should be taken
not to overlook or lose any of the precipitate because of its compactness
and density.

5. Unless a compact filter is used, there is danger of loss of the finely
divided mercury.

6. If a precipitate, usually gray, continues to form as the solution is being
filtered, it indicates incomplete precipitation of the mercury, and the mixture
should be again heated for 5 min. and refiltered.

7. The precipitation of mercury in this procedure gives no indication
as to the state of oxidation of the mercury in the material being analyzed.
If it is desired to obtain information in this regard, proceed as follows:

Introduce 0.5 g of the sample into a flask, add 50 ml of water,
and heat to 40 to 50C. for 3 min. Cool the solution, allow any
residue to settle, and add to it first 3 drops (noting the result) and
then 5 ml of HC1. (White crystalline precipitate, presence of mer-
curous compounds or of silver.) Filter the mixture through a small
paper filter and wash the residue with 10 ml of hot water. Treat
the filtrate as directed in the next paragraph. Pour 2 to 5 ml of
NHUOH through the residue on the filter. (Blackening of the
white precipitate, presence of mercurous compounds.)

Add to the filtrate 0.1 ml of SnCU reagent, allow it to stand for
1 to 2 min., and then add 5 ml more of the SnCU. (White crystal-
line precipitate, darkening on standing, presence of mercuric
compounds.)

P. 45] ESTIMATION OF MERCURY 267

As mercurous compounds are unstable, tending to decompose into metallic
mercury and mercuric compounds, it is not advisable to attempt to dissolve
the material by boiling it with an acid; also, by so doing, mercurous com-
pounds might be oxidized if certain oxidizing agents, such as nitrates, were
also present. Because of this, the material is first treated with only warm
water, and it should be noted, if the first few drops of HC1 cause a precipi-
tate; this would indicate *he presence of a soluble mercurous compound.
As many mercurous compounds, notably the halides, are insoluble, the
residue remaining from the water treatment is not filtered out but is also
treated with the ammonia. If it is desired to confirm further the presence
of mercury in this residue, it should be treated as directed in Note 1 of P. 45.

Most mercurous compounds are darkened upon application of ammonia,
because of the formation of finely divided metallic mercury. The reaction
occurring with the chloride is as follows:

Hg 2 Cl 2( .) + NH 3 - Hg w + HgNHiClw + NH 4 + + CK

Mercuric sulfide would not be affected by either the dilute HC1 or the
ammonia.

P. 46. Estimation of Mercury

Discussion. This estimation of mercury depends upon the
fact that mercuric thiocyanate is an extremely un-ionized sub-
stance. Therefore, when a standard thiocyanate solution is added
to a cold solution containing both mercuric and ferric ions, un-
ionized mercuric thiocyanate is first formed; a sufficient amount of
the ferric thiocyanate compound (see P. VI, p. 37) to give a detect-
able pink color will not form until an amount of thiocyanate closely
equivalent to the mercury present has been added.

As mercuric chloride is also only slightly ionized, the presence of
even a small amount of chloride in the solution makes the end-point
very uncertain; and for that reason the precipitate from P. 42, which
may consist partly of mercurous chloride, is treated with an alkaline
phosphite solution, by which it is more completely reduced to metal-
lic mercury and by which any mercurous chloride remaining is
metathesized into mercurous oxide. After this treatment the pre-
cipitate is dissolved in concentrated nitric acid and the solution is
heated in order to expel any nitrous acid present, as this substance
rapidly reacts with thiocyanate, forming a transitory red compound
and finally oxidizing it to hydrocyanic acid and sulfuric acid. In
order to be sure that the mercury is all in the dipositive state and
that the last traces of nitrous acid are removed, permanganate is
added to the solution until a permanent color is produced. As per-
manganate also oxidizes thiocyanate, the excess is reduced by ferrous
sulfate. The solution must be cold during the titration or the end-

268 ANALYSIS OF THE TIN GROUP [P. 45

point occurs prematurely; also, when considerable HNOs is present,
oxidation of the thiocyanate may proceed at an appreciable rate.

The titration of mercuric ion with thiocyanate is a very precise
method for the determination of mercury; 2 however, because of the
loss of mercury by volatilization in P. 42, and possible mechanical
loss in this procedure on account of the compact nature of the pre-
cipitate, the estimation of the amount of mercury present is usually
lower than the correct value by from 1 to 5 per cent. 3

Procedure 45: ESTIMATION OF MERCURY. Pour drop-
wise through the precipitate (from P. 44) on the filter a solu-
tion made by adding 1 ml of 6 f. H 3 PO 3 to 5 ml of NaOH
(Note 1). Collect the filtrate in the flask containing the
remainder of the precipitate. Heat the mixture just to
boiling for 2 to 3 min. and then slowly decant the solution
through the same filter. Discard the filtrate (Note 2).
Wash the precipitate by decantation with three 5-ml
portions of hot water. Discard the- wash water.

Dissolve the precipitate on the filter by pouring dropwise
(and repeatedly if necessary) through every portion of it 10
ml of warm 16 n. HNOs. Collect the acid with the precipi-
tate in the flask and heat it just to boiling until the precipi-
tate is completely dissolved. Wash the filter with two
10-ml portions of hot water, collecting the washings with
the solution. Boil the solution vigorously for 2 to 3 min.
and cool it to room temperature or below.

Add 0.2 f . KMnO 4 dropwise until the first perceptible per-
manent pink color (Note 3) is obtained, and then add 0.1 f.
FeSO4 dropwise until the solution is colorless, avoiding an
excess. Add 2 ml of ln.( f.) Fe(NOa)3 and titrate the cold
solution with 0.1 n. KSCN until the first permanent pink
color is produced (Note 4).

Notes:

1. If only a very small precipitate (less than 5 mg) is obtained, it is
preferable to omit this treatment and the volumetric estimation. In that
case it is advisable to confirm the presence of mercury and visually estimate
the amount present as follows:

Pour repeatedly through the filter 5 ml of HC1 to which 1 or 2
drops of liquid bromine (hood) has been added; add more Br2 if its

1 Kolthoff and Lingane, J. Am. Chem. Soc., 57, 2377 (1935).

* For a discussion of the errors involved in the gravimetric estimation of
mercury by reduction to the metal with various agents, see Willard and Boldy-
reff, /. Am. Chem. Soc., 62, 569 (1930).

P. 46J PRECIPITATION OF ANTIMONY 269

color disappears. Add 5 ml of water and boil the solution until
the bromine is expelled. Cool, add 2 ml of SnCU solution, and let
the solution stand for 5 min. (White precipitate, turning gray to
black, presence of mercury.) Compare the precipitate with
known amounts of mercury precipitated under similar conditions.

2. If this filtrate is not perfectly clear, it should be carefully refiltered.

3. As the permanganate may react somewhat slowly, the solution should
be swirled for a few seconds after obtaining a perceptible pink color in
order to be sure that it is permanent. An excess of permanganate should
be avoided, or a precipitate of Mn02 may result; should this occur, ferrous
sulfate should then be added until it is completely dissolved.

4. When a large quantity of mercury is present, a precipitate of white
crystalline Hg(SCN)2 may appear during the titration, but this does not
interfere with the detection of the pink color.

P. 46. Precipitation of Antimony

Discussion. The separation of antimony from tin by precipita-
tion of the antimony as sulfide from a hydrochloric acid solution of
the proper concentration has been shown by Noyes and Bray 4 to be
satisfactory for the qualitative detection of the two elements, but
experiments have shown that when both elements are present in
considerable quantity and the antimony is completely precipitated,
considerable amounts of tin are carried down with the precipitate.
Thus, from a solution containing 250 mg of each element, about
25 mg of tin were found to be carried down with the antimony sul-
fide precipitate.

It has been found by Vortmann and Metzel 6 that a quantitative
separation can be obtained under very similar conditions if phos-
phoric acid is present in the solution, and confirmatory experiments
have shown that a similar effect is obtained with phosphorous acid.
Thus under the conditions of this procedure, with 250 'mg of each
element present, less than 2 mg of tin were found with the antimony
sulfide precipitate, and with 500 mg of tin and 1 mg of antimony the
antimony was detected without precipitation of the tin. Therefore,
regardless of whether or not any of the phosphorous acid added in
P. 42 has been oxidized, satisfactory conditions are obtained for the
quantitative separation of antimony from tin. The precipitate ob-
tained under these conditions of acidity and temperature coagulates
readily and often is darkened by the conversion of the orange modifi-
cation of antimony sulfide to the black form.

4 Noyes and Bray, J. Am. Chem. Soc., 30, 481 (1908).
1 Vortmann and Metzel, Z. anal. Chem., 44, 533 (1905).

270 ANALYSIS OF THE TIN GROUP IP. 47

Procedure 46: PRECIPITATION OF ANTIMONY. Heat the
filtrate from P. 44, which should have a volume of 75 ml and
contain approximately 120 milli-equivalents of HC1 (Note
1), almost to boiling, saturate it with H2S, and immerse the
flask containing it in a large beaker of water at 75C. Pass
a slow current of H 2 S (3 to 5 bubbles a second) through the
solution for just 10 min., keeping the water in the beaker at
75C. (Orange to black precipitate, presence of antimony.)
Filter the hot solution immediately through an asbestos
filter which has been washed with 3 n. HC1 and drained,
and wash the precipitate with three 5-ml portions of hot
1.2 n. HC1, adding the washings to the filtrate in a 200-ml
flask (Note 2). Treat the filtrate by P. 48. Wash the
precipitate thoroughly with 1.2 n. HC1 and treat it by P. 47.

Notes :

1. It has been found that after the distillation the hydrochloric acid in
the flask is not of constant-boiling composition (approximately 6 n.) but is
about 4 n. This difference is due mainly to the presence of the phosphorous
and phosphoric acids.

2. The solution is kept hot during the precipitation and filtration and is
filtered immediately through an acid-washed filter in order to avoid the
possible precipitation of tin should it be present in large quantities.

P. 47. Estimation of Antimony

Discussion. This method for the estimation of antimony is based
upon the oxidation of antimony in a neutral solution from the tri-
to the quinque-positive state by means of a standard iodine solution.
The principles involved are very similar to those applying to the
standardization of an iodine solution by means of arsenious oxide
(P. XII) and to the estimation of arsenic (P. 43); in general, the
discussions given in those procedures apply to the method for anti-
mony. More time should be taken near the end of this titration, as
the reaction between the iodine and the antimony is somewhat slow. 6

The antimony, being in the form of sulfide, is dissolved in 6 n.
hydrochloric acid, whereby the hydrogen sulfide set free insures
the reduction of all of the antimony to the tripositive state; the
hydrogen sulfide is then expelled by boiling. If the hydrochloric
acid solution were diluted largely, or neutralized, precipitation of
antimonyl chloride, SbOCl, or the hydrous oxide, would result and

5 For a study and very complete bibliography of various methods for the
determination of antimony, see McNabb and Wagner, /. Ind. Eng. Chem.,
Anal. Ed., 2, 251 (1930).

P. 47] ESTIMATION OF ANTIMONY 271

the iodine would react only slowly with the precipitate. This pre-
cipitation is prevented by the addition of tartaric acid, which forms
a soluble ion, SbOC 4 H4Oe~. This allows the solution to be neu-
tralized and the excess of hydrocarbonate required for the titra-
tion to be added without the formation of a precipitate. Anti-
monic oxide also dissolves in tartrate solutions, forming the ion

The antimony in stibnite (native SbaSa) materials may be deter-
mined by the procedure given below, provided no interfering ele-
ments, such as iron or arsenic, are present. As the native material
is more resistant to solution, it may be necessary to use an additional
amount of concentrated hydrochloric acid; in that case the solution
process should be carried out in a flask fitted with a test-tube con-
denser to avoid possible volatilization of antimony as the trichlo-
ride; instructions for carrying out this determination are outlined in
Note 6 below.

Hillebrand and Lundell 7 show that the loss of antimony is not
significant when dilute hydrochloric acid solutions in covered
vessels are boiled but not concentrated, but that upon evapo-
ration to a small volume the loss is large. Thus over half of
approximately 250 mg of antimony was lost when a solution of the
trichloride in hydrochloric acid was twice evaporated to a syrup.

If a standard solution of potassium iodate is available, the anti-
mony can be estimated by making use of the iodine monochloride
end-point. 8 In this method the tripositive antimony is titrated
in a hydrochloric acid solution with standard iodate until the iodine,
which is the first reduction product of the iodate, is oxidized to
colorless iodine monochloride and can no longer be detected in a
small amount of carbon tetrachloride or similar solvent. The suc-
cessive reactions taking place can be represented as follows :

Reduction of iodate to iodine and oxidation of the antimony :

5HSbCl 4 + 2IO 8 ~ + 12H+ + 10CF = 5HSbCl 6 + I 2 + 6H 2 O. (1)

Oxidation of the iodine to iodine monochloride by further addition
of iodate:

2I 2 + I0 3 ~ + 6H+ + 5CF = 5IC1 + 3H 2 0. (2)

Experiments have shown that, if the hydrochloric acid is much

1 Hillebrand and Lundell, Applied Inorganic Analysis, Wiley, 1929, pp.
221-222.

8 See pp. 85-88; also Andrews, J. Am. Chem. Soc., 25, 756 (1903); Jamieson,
/. Ind. Eng. Chem., 3, 250 (1911); Jamieson, Volumetric Iodate Methods, Chem.
Cat. Co., 1926, pp. 7-17.

272 ANALYSIS OF THE TIN GROUP [P. 47

more concentrated than 3 n., the reaction apparently proceeds
mainly as follows:

2HSbCl 4 + IO 8 ~ + 6H+ + 5CF - 2HSbCl 6 + IC1 + 3H 2 0, (3)

and so little iodine is formed that the end-point is not distinct. If
the hydrochloric, acid concentration is much less than 2 n., precipi-
tation of antimony 1 chloride may result and the rate at which the
reaction represented by Equation 2 takes place is so slow that an
excess of iodate may be added before the end-point is obtained.

As the method can be readily applied to the hydrochloric acid
solution of the antimony sulfide and as, when the hydrochloric acid
concentration is properly regulated, it is so precisely and rapidly
carried out, an optional procedure for its use is provided in Note 2
of the procedure below.

Procedure 47: ESTIMATION OF ANTIMONY. Transfer the
precipitate from P. 46 (Notes 1, 6) to a 400-ml beaker, add to
it 10 to 15 ml of HC1, cover the beaker with a clock glass,
heat the mixture slowly just to boiling, and boil it gently
until the H2S is completely expelled (as shown by the fail-
ure of a strip of filter paper moistened with Pb(C 2 H 3 O2)2
solution to darken when held in the vapors). Cool the solu-
tion and add to it 2 g of finely powdered tartaric acid
(Note 2).

If the amount of antimony present is estimated to be less
than 100 'to 150 mg (Note 3), dilute the solution slowly
(Notes 4, 5) to 50 ml and treat the diluted solution as
directed in the last paragraph of P. 43, "Estimation of
Arsenic. "

If the amount of antimony present is estimated to be more
than 100 to 150 mg, transfer the solution to a 100-ml volu-
metric flask, dilute it slowly to the mark, and mix the con-
tents thoroughly (Notes 4, 5). Pipet 25 ml of this solution
into a 200-ml flask, add to it 2 g of powdered tartaric acid,
dilute it to 50 ml, and treat it as directed in the last para-
graph of P. 43, "Estimation of Arsenic. " From the volume
of standard iodine used calculate the amount of antimony
present.

Notes:

1. In case only a small precipitate without the characteristic orange
color of Sb2Sa was obtained, it should be treated as follows:

Transfer the precipitate to a 200-ml beaker, add 13 ml of HC1,

P. 47] ESTIMATION OF ANTIMONY 273

boil until slightly less than 10 ml are present, cool, transfer to a
10-ml graduate, and dilute with 6 n. HC1 to just 10 ml. Transfer
the mixture to a 100-ml flask, taking care not to dilute it, heat it
just to boiling, and saturate it with H 2 S. Filter out and discard
any precipitate (AsaSa). Dilute the solution to just 25 ml, heat it
almost to boiling, and again saturate it with H 2 S. (Orange-red
precipitate, presence of antimony.)
Arsenic and mercury, but not antimony, are precipitated as sulfides from

a solution 6 f. in HC1; antimony, but not tin, is precipitated from a solution

2.4 f. in HC1.

2. In case it is desired to titrate the antimony with a standard iodate
solution, omit the addition of the tartaric acid and proceed as follows:

If less than 100 mg of antimony are thought to be present, dilute
the solution with exactly an equal volume of water and transfer it
to a 200-ml ground-glass-stoppered flask with the aid of 10 to 20 ml
of 3 n. HC1. Treat it as directed in the second paragraph below.
If more than 100 mg of antimony are thought to be present,
dilute the solution with exactly an equal volume of water, transfer
it to a 100-ml volumetric flask, and dilute it to the mark with 3 n.
HC1. Pipet out a 25-ml portion into a 200-ml ground-glass-
stoppered flask and treat as directed below.

Add to the solution 5 ml of CC1 4 and titrate with a 0.1 n. (0.025 f.)
iodate solution until the iodine color which first appears can no
longer be observed in the CCU. Add 10 ml of 6 n. HC1 with each
10 ml of the iodate solution. As the end-point is approached,
stopper the flask and shake the mixture vigorously before observing
the CC1 4 .

The iodine color can be most effectively observed by stoppering the flask
tightly, inverting it, and looking through the CCU after it has collected in
the neck of the flask. Loss of solution from the flask can be prevented by
cooling the flask with tap water and then pouring 0.5 to 1 ml of 3 n. HC1
around the stopper before it is withdrawn.

3. This estimation should be based upon the size of the Sb2Sa precipitate
obtained in P. 46.

4. The solution should be diluted slowly to avoid precipitation of SbOCl.
If this substance is precipitated locally (by too rapid dilution of the solu-
tion), it redissolves very slowly.

5. If all the HaS has not been expelled by the previous boiling, an orange-
colored precipitate will form when the solution is diluted. In this case the
mixture should be immediately boiled, more acid being added if necessary,
until this precipitate is dissolved, and the dilution should then be continued.

6. In case it is desired to determine the antimony in a native stibnite
(SbjSs), directly proceed as follows:

Weigh out a 0.5 to 1-g sample of the ore, transfer it to a 300-ml
flask, and add 20 ml of 12 n. HC1. Provide the flask with a test-
tube condenser (Note 2, P. 7), and treat the mixture as directed in
the procedure above.

Arsenic and iron should not be present in significant amounts, as they
would be reduced to their lower oxidation states by the above treatment
and would be oxidized in the subsequent titration.

274 ANALYSIS OF THE TIN GROUP [P. 48

P. 48. Detection and Precipitation of Tin

Discussion. The conventional method of precipitating the tin as
sulfide by partly neutralizing the filtrate from the antimony sulfide
precipitate and saturating it with H 2 S is modified here, because the
presence of the phosphorous and phosphoric acids cause the precipi-
tation of the tin to take place very slowly, probably because it
exists largely as a complex molecule. However, by first making
the solution alkaline and then treating it with H 2 S, sufficient sulfide
is formed in the solution to convert the tin into a soluble sulfo-salt,
which, when the solution is then acidified, rapidly decomposes,
with the precipitation of the tin as sulfide (see the discussion of
P. 13).

Procedure 48: DETECTION AND PRECIPITATION OF TIN.
Cool the filtrate from P. 46 and add to it 15 n. NH 4 OH until
it is alkaline, avoiding a large excess (Note 1). Saturate the
solution with H 2 S and then pass a slow current of the gas
through it for 2 to 5 min. Heat the solution almost to boil-
ing, slowly add H 2 S04 until it is just acid, and then add 2
ml more. (Yellowish precipitate, presence of tin. Note 2.)

If a precipitate is produced, immerse the flask containing
the mixture in a beaker of boiling water and pass a slow
current of H 2 S through it for 5 min. Let the precipitate
settle and filter the solution by decantation, using suction
if the precipitate is large. Add to the precipitate 0.5 g of
solid NH 4 C1 and wash it by decantation with a 15-ml portion
of hot water. Discard the filtrate and washings (Note 3).
Treat the precipitate by P. 49 (Note 4).

Notes:

1. A volume of 15 n. NH40H equivalent to the HC1 present can be calcu-
lated and added at once. The additional ammonia required to complete
the neutralization of the solution can then be added 1 ml at a time until the
solution turns red litmus paper blue.

Phosphorous acid of commerce frequently contains iron, and this is
precipitated from the alkaline solution as ferrous sulfide, causing the tin
sulfide precipitate to be dark in color even after the solution is acidified.
As iron does not interfere with the estimation of tin (in P. 49), it can be neg-
lected in case a precipitate of considerable size (more than 10 to 20 mg) is
obtained.

2. If a small dark precipitate is obtained which makes doubtful the de-
cision as to the presence of tin, proceed as follows:

Filter out the precipitate on a paper filter and pour dropwise
through it 2 nil of warm sodium sulfide reagent. Collect the solu-

P. 49] ESTIMATION OF TIN 275

tion in a flask, make the solution just acid with HC1, and then add
just one-half its volume of 6 n. HC1 in excess. Heat the solution
to boiling and saturate it with H^S. Filter out and discard any
precipitate (Sb2Sa). Collect the filtrate in a flask, dilute it with
twice its volume of water, saturate it with H 2 S, and allow it to stand
for 10 min. (Yellow precipitate, presence of tin.)

3. The complete and rapid precipitation of the tin sulfide depends upon
the formation of the sulfo-salt, and therefore, if a considerable precipitate
has been obtained, it is advisable to repeat the first paragraph of this pro-
cedure with the filtrate before it is discarded.

4. The precipitation of the tin at this point gives no information as to
the state of oxidation of the tin in the original material. If it is desired to
know if stannous tin is a constituent of the sample, proceed as follows:

Boil 10 ml of HC1 in a small flask (covered with a watch glass)
for 3 min. Quickly introduce 0.5 g of the sample into the flask,
cover it, and again boil for 3 min. (if H2S is evolved, continue the
boiling until it is expelled). Add in small portions 0.2 g of NaHC0 3
and cool the mixture to room temperature with running tap water,
holding the watch glass tightly in place. Pour the cold solution
into 10 ml of a saturated solution of HgCl 2 (if a residue is present,
pass the solution through a rapid-filtering paper filter and collect
the filtrate in the HgCl 2 solution). (White precipitate, presence
of stannous tin.)

The HC1 solution is first boiled and then treated with NaHC0 3 to prevent
the oxidation of stannous chloride by the oxygen of the air. Upon adding
a solution of stannous tin in hydrochloric acid (in which the tin is present
largely as the ion SnCU") to a mercuric chloride solution, a precipitate of
mercurous chloride is formed; the reaction taking place can be represented
as follows:

SnCl 4 - + 2HgCl 2 - SnCU" + Hg 2 Cl 2 (.>.

The stannous chloride solution is cooled and added to a large excess of
the mercuric chloride in order that the finely divided, crystalline precipitate
characteristic of mercurous chloride will be produced; if the solutions were
hot or an excess of stannous tin were present, the reduction would proceed
to the formation of a dark precipitate of finely divided mercury.

p. 49. Estimation of Tin

Discussion. This method for the estimation of tin, which has
been studied by Hallett, 9 is based upon obtaining complete reduc-
tion of the stannic salt to the stannous state by means of metallic
nickel, and then oxidizing the reduced solution with a standard
solution of iodine. Experiments have shown that the time neces-
sary for complete reduction of the tin is reduced from about 30 min.
to about 10 min. by the addition of an antimony salt to the solution.

Hallett, J. Soc. Chem. Ind., 35, 1087 (1916). This reference should be
consulted for a discussion of this and various other methods for the volumetric
estimation of tin.

276 ANALYSIS OF THE TIN GROUP IP. 49

This is probably because the metallic antimony produced reacts
more rapidly with the stannic tin than does the nickel foil. This
metallic antimony usually adheres to the nickel foil and is removed
from the solution. It may, especially if the mixture is vigorously
boiled, become detached and remain in the solution. This does
not introduce an error if the solution is kept cold during the ti-
tration, as neither the iodine nor the oxidized tin is reduced at an
appreciable rate. An experimental study 10 has shown that, unless
the hydrochloric acid is at least 3 n., a longer time is required for
the reduction, and that the acid concentration can be raised to 6 n.
without any harmful results being noticed. By diluting the solution
before the titration, the acidity is reduced so that starch can be more
suitably used as the indicator; also, the color of the dissolved nickel
is decreased and the 'end-point is made more distinct.

Complete reduction having been obtained, the precision of this
estimation is determined largely by the effectiveness of the pre-
cautions taken to prevent the oxidation of the stannous chloride by
the oxygen of the air.

Other methods for the reduction of the tin have been investigated.
Thus Ibbotson and Brearly 11 have proposed the use of finely divided
metallic antimony in a boiling hydrochloric acid solution. After
the mixture is cooled, the stannous tin is titrated with iodine without
removal of the antimony, the assumption being that in the cold
solution the rate of reduction of the stannic tin by the antimony
becomes so slow as to be negligible, so that a permanent starch-
iodine end-point can be obtained. A careful study of this method 12
has shown that the reduction of even large amounts of tin can be
made complete in 2 to 3 min. and that very consistent results can be
obtained by one familiar with the method. It was found, however,
that the method was quite dependent upon the fineness of the anti-
mony, material passing a 150- to 200-mesh screen being necessary.
If the antimony was finer than this, it remained suspended when
the mixture was cooled and apparently reacted with the iodine so
that permanent end-points were not obtained; coarser material
caused slower reduction. It was also necessary to heat the mixture
in such a way as to cause the antimony to remain uniformly sus-
pended, or reduction was incomplete. Finally, practical use showed

10 Unpublished experiments by G. S^ Schull.

11 Ibbotson and Brearly, Chem. News, 84, 167 (1901).
11 Unpublished experiments by Robley D. Evans.

P. 19] ESTIMATION OF TIN 277

that persons unfamiliar with the method found it difficult to detect
the starch-iodine end-point in the presence of antimony, and for
these reasons the method was not adopted, even though under
strictly controlled conditions it is much shorter and as precise as
the method given below.

The use of metallic zinc, in a modified method developed by Milon
and Fouret, 13 was also investigated. In this method the stannic tin
is reduced more or less completely to the metal by the zinc in a
hydrochloric acid solution. When the solution is then heated, the
excess of zinc is rapidly dissolved and the precipitated tin dissolves
as stannous chloride. It was found, however, that under certain
conditions the tin dissolved slowly and particles remained unde-
tected in the solution. Other reducing agents, such as "test-lead"
pellets, 14 metallic iron (usually with an antimony salt), and alumi-
num, have been suggested but were found to be not so rapid or so
generally satisfactory as the method of this procedure.

Procedure 49: ESTIMATION OF TIN. Add to the precipi-
tate in the flask 10 ml of HC1 and heat the mixture until
the precipitate dissolves. Pour the hot solution slowly
through the precipitate on the filter, catching the filtrate
in a 400-ml flask (Note 1). Wash the filter with 10 ml of
hot HC1, collecting the washings with the filtrate. Discard
any residue.

If the amount of tin present is estimated to be less than
150 to 200 mg, treat the solution by the last paragraph of
this procedure.

If the amount of tin present is estimated to be more than
150 to 200 mg, cool the solution to room temperature, pour
it into a 100-ml volumetric flask, dilute it to the mark, and
thoroughly mix the contents. Pipet 25 ml of this solution
into a 400-ml flask and treat it by the next paragraph.

Add to the solution 5 ml of 0.1 f. SbCU solution and 30 ml
of 12 n. Hpl (Note 2), and dilute the solution to a volume of
100 ml. Cut a thin sheet of nickel foil about 3 cm by 30 cm,
having a tab about 12 cm long extending perpendicularly
frqm one end (Note 3). Roll the nickel into a loose spiral

18 Milon and Fouret, Kept. 8th Int. Cong. App. Chem., 1, 373 (1912).

" Powell, /. Soc. Chem. Ind., 37, 287 T (1918). Lundell and Scherrer,
J. Ind. Eng. Chem., 14, 426 (1922). Patrick and Wilsnack, J. Am. Chem. Soc.,
4, 597 (1882).

278 ANALYSIS OF THE TIN GROUP [P. 49

and introduce it into the solution by means of the tab.
Heat the solution to boiling and keep it boiling for at least
10 min. (Note 4). Remove the flame and just as the solu-
tion stops boiling drop into the flask a marble chip about
1 cc in size (Note 5). The marble should be free from iron
and should be boiled vigorously in distilled water for 2 to 3
min. just before being used. Cool the flask at once with
tap water. If the marble chip is dissolved during this cool-
ing, another should be added. Cool the flask at such a rate
that the C0 2 evolved prevents air from being sucked into it.
Remove the nickel, washing it with cold, recently boiled
water as it is withdrawn. Add rapidly to the cool solution
100 ml of cold, freshly boiled water containing 5 ml of starch
indicator solution and titrate it immediately and rapidly
with 0.1 n. iodine solution until the appearance of the starch
blue color (Note 6). The mouth of the flask should be cov-
ered with a card, and the tip of the buret should be intro-
duced into it through a small hole. From the volume of
iodine used, calculate the amount of tin present.

Notes:

1. If a large amount of the precipitate has been carried onto the filter, it
may not dissolve readily. In this case transfer the precipitate and filter
to a small casserole and heat with the HC1 until the precipitate dis-
solves. Filter out the disintegrated paper upon another small paper fil-
ter, wash the residue with 10 to 20 ml of HC1, discard it, and treat the
solution as directed in the next paragraph of this procedure.

2. The concentration of the HC1 should be at least 3 n. during this reduc-
tion.

3. The nickel spiral should be of at least the size recommended; the more
surface exposed the more rapid the reduction. If used previously, the spiral
should be cleaned, by immersion in warm 6 n. HC1 to which a little 16 n.
HNOa has been added, until any adhering antimony has been removed and
a bright surface has been obtained.

4. It is necessary that the solution be boiled steadily so that it is con-
tinuously stirred and all parts are brought into contact with the nickel.
Very vigorous boiling may detach the precipitated antimony from the nickel,
but, if the solution is cooled to room temperature before making the titra-
tion, this antimony then reacts so slowly with both the stannic tin and the
iodine that it does not introduce an appreciable error. By applying the heat
directly under the nickel spiral, a continuous flow of solution is maintained;
if the solution does not cover the spiral, clamp the flask in an inclined posi-
tion. If the solution does not cover the nickel foil, if the nickel foil has not
the surface recommended, or if the solution is not effectively stirred, it is
recommended that the boiling be continued for 20 min.

P. 49] ESTIMATION OF TIN 279

5. If a carbon dioxide generator (or tank) is available, it is preferable
to keep a flow of the gas passing continuously through the solution during
the boiling as well as during the subsequent cooling and titration.

6. It is essential that the sequence of operations following the stopping
of the boiling of the solution be carried out as rapidly as possible. Accord-
ingly, the freshly boiled marble and the freshly boiled water with the starch
added to it should be prepared previously. The solution should then be
titrated rapidly with a minimum amount of shaking. If a slight excess of
iodine is added, it can be back-titrated with standard thiosulfate solution.

TABULAR OUTLINE VI

PRECIPITATION OF THE AMMONIUM SULFIDE GROUP; SEPARATION OF IRON;
REMOVAL OF PHOSPHATE; AND SEPARATION OF THE ALUMINUM AND

ZINC GROUPS

Filtrate from the Precipitation of the Hydrogen Sulfide Group:
Elements of the Ammonium Sulfide, Alkaline Earth, and Alkali Groups;
H+, NH 4 +, C1-, NOr, H 2 S

Boil out the H^S. Neutralize with NH>H.

Precipitate: Al(OH),, Cr(OH),, Fe(OH),, Alkaline Earth phosphates

Treat with H^S, add excess of NH<OH. (P. 51)

Precipitate: Al(OH),, Cr(OH),, FeS, Fe 2 S 3 , ZnS, NiS,
CoS, MnS, Alkaline Earth Phosphates

Dissolve in 6 n. HCl, adding Br*. Boil out excess Br t .

(Test small portion for Fe with KSCN.)

Shake with ether. Separate ether from aqueous layer.
(P. 52)

Filtrate:

Alkaline

Earth

and

Alkali

Group Elements

Ether solution:
Fed,

Evaporate the

ether.

Add HCl, KI.

(Fe++, Ir)

Titrate with

Aqueous solution: A1+++, Cr+++, Zn++, Ni++, Co++,

Mn++, (Ba++, Sr++, Ca++, Mg++), H,PO 4 , H+C1~
Phosphate Absent: Treat by P. 55.

Phosphate present:

Evaporate, add 16 n. HNO*. Evaporate.
Make 0.5 n. in HNO t .
Add Bi(NO t )t reagent. (P. 54)

JvaaSjfi. 1*6^,
I-, S 4 0,-)

(P. 53)

Precipitate:
BiP0 4
Heat.
Weigh.

Solution: Positive ions listed above and
31+++, H+NOr.
Saturate with H*S.

Precipitate :
Bi 2 S,
Discard.

Solution: Positive ions
listed above, H+NOr
Treat by P. 51.

P. 55. Neutralize with NH 4 OH, make just acid with HCl.
Add (AT// 4 ) 2 C 2 4 , NaHCd.
Treat with H^S.

Precipitate: ZnS, NiS, CoS, (FeS)
The Zinc Group

(P. 61-P. 66)

Filtrate: Al(C 2 O 4 ) l -! ", Cr(C 2 O 4 )i",
Mn(C,0 4 )r

The Aluminum Group
(P. 71-P. 75)

280

Precipitation of the Ammonium Sulfide Group;

Separation of Iron; Removal of Phosphate;

and Separation of the Aluminum

and Zinc Groups

P. 51. Precipitation of the Ammonium Sulfide Group

Discussion. Those elements whose sulfides are insoluble in 0.3 n.
acid having been precipitated, it is now desirable to find some means
of separating another group of elements from the solution. As
several methods are available, a study of their relative merits should
be made.

Hydroxide Separations. Methods for separating the elements
which depend upon their behavior in solutions of various hydrogen
ion concentrations are, after the sulfide methods, most extensively
used in qualitative and quantitative analysis; because of this there
have been collected in Tables XII, p. 212, and XVIII data showing
the behavior of the common elements in solutions of various hydro-
gen ion concentrations.

Surveying the elements remaining after the Hydrogen Sulfide
Group is separated, it is seen that, if the solution is made approxi-
mately neutral, only the hydroxides of the elements present in the
tripositive state- aluminum, chromium, and iron are precipitated.
The isolation of these tripositive elements as their hydroxides would
make an ideal group separation, both pedagogically and analytically,
and is sometimes attempted (any ferrous iron being first oxidized)
by neutralizing the solution and adding a slight excess of ammonium
hydroxide, which results in making what is known as an "ammonia
precipitation/' As this precipitation with ammonia is so extensively
used for various separations in qualitative systems and in quantita-
tive work, an experimental study of it was made ; a discussion of the
results obtained is given below. 1

The Precipitation by Ammonium Hydroxide. The precipitation by
ammonium hydroxide can well be termed one of the classical ana-
lytical separations and is, as stated by Hillebrand and Lundell, 2

1 The following discussion is taken in part from a study of this method by
Swift and Barton, /. Am. Chem. Soc., 64, 2219 (1932).

2 Hillebrand and Lundell, Applied Inorganic Analysis, Wiley, 1929, p. 69.

281

282

AMMONIUM SULFIDE GROUP

[P. 51

TABLE XVIII

THE BEHAVIOR OF THE AMMONIUM SULFIDE AND ALKALINE EARTH GROUP
ELEMENTS IN SOLUTIONS OF VARIOUS HYDROXYL (AND HYDROGEN)

ION CONCENTRATIONS
(See Table XII, p.212, for the behavior of the Hydrogen Sulfide Group elements.)

Precipitates, Compounds, or Ions Formed

Element and Oxida-

tion State

[OH-]

[H+]

EH*]

2 to 4 m.

10-' to KT 7 m.

10- to 10-" m.

Ali"

Al(OH)r

Al(OH),

Cr'"

Cr(OH)r a d

Cr(OH) 8

Cr(OH),'

Fe" 1

Fe(OH),

Fe"

Fe(OH) 2 -

(Fe++) '

(Fe**)

Mn(OH) 2 -

(Mn++)

(Mn**)

Ni"

Ni(OH) 2

(Ni++)'

(Mi- 1 ""-)

Co"

Co(OH) 2

(Co++)'

(Co+ + )

Zn"

Zn(OH)r

(Zn++)>

(Zn**)

Ba"

Ba(OH) 2 fc

(Ba++)

Sr"

Sr(OH) 2 "

(Sr++)

(Sr ++ )

Ca"

Ca(OH) 2 >

(Ca++)

Mg(OH) 2

(Mg++)

Separations which are dependent upon having the hydroxyl ion concentra-
tion from 2 to 4 m. (column I) are usually made by the use of solutions of
sodium or potassium hydroxide; by the addition of an oxidizing agent (usually
Na 2 O 2 ), chromium is oxidized to chromate, ferrous hydroxide to ferric, man-
ganese to manganese dioxide, and cobalt to cobaltic oxide. Separations
depending upon maintaining the solution approximately neutral (column II)
are obtained (1) by the use of ammonia in the presence of a large excess of
ammonium salts, whereby the specific effect of the formation of complex
ammino compounds is also obtained; (2) by using an excess of solid barium
carbonate in a cold solution; (3) by using a solution containing an excess of
iodate and iodide; or (4) by boiling the solution with sodium thiosulfate.
The conditions of column III are most commonly obtained by the use of a
boiling solution containing acetic acid and a soluble acetate the so-called
"basic acetate method" (see P. 63); other organic acids and their salts which
have been similarly used are formic, succinic,* and benzoic.f

See General Note to Table XII in regard to the composition of the hydrox-
ides and the formulas of the ions formed by the amphoteric elements.

Cr(OH)s dissolves in dilute hydroxide solutions, owing to formation of a
stable colloidal system; with more concentrated solutions Cr(OH) 4 ~ is formed.
b Moderately soluble.
Incompletely precipitated.
d Oxidized to CrOr by Na 2 O 2 , NaCIO, and so forth.

* Oxidized upon exposure to air.

* Forms complex ammino ions with ammonia.

8 With presence of ammonia, is oxidized in air and forms complex ammino
ions.

* Treadwell-Hall, Analytical Chemistry, Vol. II, Quantitative, 8th Ed.,
p. 160.

t Kolthoff, Stenger, and Moskovitz, J. Am. Chem. Soc., 66, 812 (1934).

P. 51] HYDROXIDE SEPARATIONS 283

"one of the commonest operations the analyst has to perform . . . ,
with the object either of weighing the precipitated compound or of
effecting a joint separation of two or more metals from others."
That it may be inadequate even as a qualitative separation in certain
cases is shown by the experiments of Noyes and Bray 8 in which,
with large amounts of aluminum or ferric iron (100 to 200 mg) and
with amounts of cobalt, zinc, or nickel up to 20 mg, from 75 to 99
per cent of the latter elements were found to be carried down in the
precipitate. Noyes, Bray, and Spear 4 also state that "a large quan-
tity of zinc may be quantitatively precipitated by ammonium
hydroxide when a larger proportion of chromium is present; and
manganese will in any case be partially precipitated by that reagent
owing to its oxidation by the air." However, the somewhat con-
tradictory results which have been obtained in various investigations 6
of the quantitative use of the method have indicated that the effec-
tiveness of the separations which can be obtained are highly de-
pendent upon the exact conditions under which the method is carried
out, especially with respect to the excess of ammonia added in mak-
ing the precipitation. The effects caused by an excess of ammonia
depend upon whether the bipositive elements remain in solution
because of the solubility of their hydroxides (as is most probable
with manganese) or because of the formation of the soluble ammonia
complexes, and also upon the relative tendency of these two molecu-
lar types to be carried down with the precipitate. The effect of the
hydrogen ion concentration of the solution upon the physical nature,
and, therefore, the adsorbing tendency of the precipitate, must also
be considered. An experimental study of the method, with especial
reference to these points, was made, 6 and the data obtained are
tabulated in Table XIX and are discussed below.

An inspection of the data in Table XIX seems to lead to two
general conclusions: First, in about half of the separations studied
the separation is quite unsatisfactory even under the most favorable
conditions. Second, in every case studied the separation is more
effective when an excess of ammonia is avoided and, in most cases,
unless this excess is avoided, the separation is so imperfect that little
is gained by reprecipitations; thus, under the conditions studied,

8 Noyes and Bray, A System of Qualitative Analysis for the Rare Elements ,
Macmillan, 1927, pp. 153-155.

4 Noyes, Bray, and Spear, /. Am. Chem. Soc., 30, 482 (1908).

5 Blum, J. Am. Chem. Soc., 38, 1291 (1916); Lundell and Knowles, ibid.,
45, 676 (1923).

6 Swift and Barton, loc. cit.

SEPARATI
In the
the elem
proeedur

CSI < k

S^H t^ 1-1 O>
^-.<N<N-

o
(3

8
g

> 1 HH > 1
H-H

^SSS

"S.

rH rH rH

CO CO CO CO

s^ss

.s a

S co S a>

O CO QO cO CO
CO 1-1 CO W

CO -* 00

1 HH > i

t-H HH

> 1 (

> <

C<l CO ^ *O

r- H rH f r-l

8SSSS

O* CO -f

.2 ?

"^ CO O 00 ^ *O

rH CO 35 <N 1 CO
r- 4 H i ( JO

co ^ co co

"ciok""

1 >-H I-H

< > 1

H-t 1 4 > 1

"S.

CO l>- OO O5 O -i

i-H <

SS8

OS O -H QO Ci

CO "'T ^ *^ Tt^

QO Tt* O

o c^ co o

O 5 QO t^

d .

*H +*

SO O O 00

^f <N

W I-H O <N

^ a

^ '

PL,

Q >A -O

^^ ^^

-NCO^US

2SSSJ

CO CO CO CO

+=
c
a.

1
5

J3
O

9 S-g-

Sb 2

'S3 -C^

13 .'13 C
J 3

o fce

.5 oT-S^ ^ 53
TJ S g J X a

S Jj ^ 8 TJ
rt o a) g os

O y *** . ^H O

T3- w) gig s
2 "S .9 2 g

2- rs

"** "

o SR > *" **
2 v *M 2 08

2ll||^

S?5|1^

"^ >H

#> i *H *

ced

im* in his study of the precipitation of aluminum

>ride and 6 to 12 milli-equivalents of HCI in a
ihange was noted, with methyl red, or in a few
ented the use of the internal indicator, litmus
matched that obtained from a similar solution
ad been added until the methyl red (or rosolic
ions; if the color transition was overrun, the
d for 1 or 2 min. and filtered; it was kept hot
irate filters were used in most cases. The pre-
the washings gave no appreciable precipitate

r volume, a larger quantity of ammonium chloride
In order to compare more exactly the erTect of an

ows :
t, after the solution was carefully neutralized

n the specific notes to the table. The precipitate

100 ml of wash solution; the precipitate in Expen-
n of ammonia and hydrogen peroxide.
&ras added until the methyl red was distinctly pink,

cperiment 6.
,he hydrochloric acid,
ion with an internal indicator. Probably slightly

made, and an excess of ammonia was avoided, as

se Experiments 48^ind 49 and Experiments 50 and
>nia.

i-*

s
1

onium chic
til a color c
ution previ

the litmus ]
ydroxide h
neutralizat

tt> w ^^ *z
P^-S So
8 % S

- o *

8*3

ho o *"

c "**
+* <D
3

contained i
rried down

lie *$%

%%* -5!

- t c 1 'i s

O ^ CD -G w cj

2 -a o ^.E2 S

zation was

;d; otherwi
ss of amm<

S So

<2 3 CO

**-i fl)
d QQ

.32 'g

2j3 S

-C2

]2>-r

3 0) - "I h

& ^>S

VU Sf

"O C3

-5 x

'he second procedure conformed more closely to the procedure 01

was as follows:

Procedure II : To a boiling solution which contained 10 g of ar
volume of approximately 250 ml, 6 n. NH 4 OH was added dropwise
cases with rosolic acid, as the indicator. Where the color of the

test papers were used and the ammonia was added until the color
of ammonium chloride and hydrochloric acid, to which ammoniui
acid) color transition occurred. Extreme care was taken in the
mixture was made acid and the process was repeated. The mixt
until the filtration was completed. To expedite filtering and was
cipitate was washed with a hot 2 per cent solution of ammonium
when tested with ammonium sulfide solution,
rocedure II differs from Procedure I in that the solution was dilu

present, and extreme care was taken to avoid an excess of ammonium
iss of ammonium hydroxide, .the experiments labeled III were carr
Procedure III : This was performed exactly as was Procedure 1
with ammonia, an excess of 2 ml of 6 n. NH 4 OH was added,
observations and variations from the procedures outlined above a
analyzed in order to determine the amount of the soluble element

The chromium hydroxide precipitate obtained in Experiment 3 re
t 5 required about 500 ml. The washings were tested for mangam
In this neutralization a slight excess of ammonia was added; there
the neutralization was repeated.
A slightly larger excess of ammonia was added in Experiments 7 an
Five grams of NH 4 C1 were used in addition to that formed by neut
The solution was made just neutral to litmus without using a re
e ammonia was added than in Experiment 13.
The volume and other conditions were as in Procedure I. Neuti

'rocedure II.
Only 160 mg of zinc were taken.
In Experiments 49 and 51 an excess of 5 ml of 15 n. NH 4 OH was a<
ere carried out in exact duplicate to note the effect of the larger e

o 3

o w io n3 *w

t-t ^

P-i *

.2 S

285

286 AMMONIUM SULFIDE GROUP [P. 51

when an excess of ammonia is added, more than 50 per cent of the
nickel, cobalt, or zinc is carried out by either chromium or aluminum.
This would seem to indicate that the ammonia separation is more
effectively carried out, at least in dilute solutions, as a process of
selective hydrolysis and not as one depending on the formation of the
soluble complex ammino ions. In support of this, it is to be noted
that manganese, where the tendency toward this complex formation
is least, is much less coprecipitated than nickel, cobalt, or zinc.

Also, the pH. values at which these bipositive elements are pre-
cipitated from solution are given by Britton 7 as follows: zinc, 5.2;
nickel, 6.7; cobalt, 6.8; and manganese, 8.5 to 8.8. It is seen that
the coprecipitation in the experiments carried out by Procedure II
in every case varies in amount in the sartie order zinc, the least
soluble hydroxide, showing the greatest tendency to be carried with
the precipitate. The same order, in general, holds for the experi-
ments by Procedures I and III, which is somewhat surprising, as,
with an excess of ammonia present, it would be expected that the
formation of the soluble ammonia complexes would be a more de-
ciding factor; for the same reason it would have been predicted that
the large excess of ammonia added in Experiments 49 and 51 would
have decreased markedly the amount of coprecipitation; however,
the difference is within the experimental variations. That the effect
is due to an adsorption process and not to mechanical inclusion or
local precipitation is indicated by experiments which have been made
by Ibbotson and Brearly 8 and Noyes and Bray. 9 These show that,
when an ammoniacal solution of the bipositive element is added to
a suspension of the freshly precipitated hydroxide, the coprecipita-
tion approaches in amount that obtained by precipitation in the
presence of the bipositive element. That the complex ammonia
compounds are not extensively carried down was shown by the fact
that relatively little ammonia was found upon analysis of an alumi-
num precipitate, produced by Procedure III, which had coprecipi-
tated with it about 200 mg of nickel.

Specifically, in addition to confirming the conclusions of Lundell
and Knowles 10 that, by proper methods of neutralization, a satisfac-
tory separation of manganese from aluminum and from iron is

7 Britton, Hydrogen Ions, Van Nostrand, 1929, pp. 254, 278.
1 Ibbotson and Brearly, Chem. News, 81, 193 (1900).

9 Noyes and Bray, A System of Qualitative Analysis for the Rare Elements,
Macmillan, 1927, p. 154.

* Lundell and Knowles, /. Am. Chem. Soc., 45, 676 (1923).

P. 51] OTHER GROUP SEPARATIONS 287

obtained, it is shown that under these same conditions manganese
can be separated from chromium. However, it is to be noted that
the coprecipitation of manganese is increased much more by an
excess of ammonia in the separation from chromium than it is in
the separation from either aluminum or iron. The separations of
nickel from aluminum and chromium show from 50 to 90 per cent of
this element brought down when an excess of ammonia is added,
demonstrating the futility of reprecipitations; the same general
behavior is obtained with these elements and cobalt. Even under
the most favorable conditions, these separations are hardly adequate
for quantitative work. The separations of nickel and cobalt from
iron, made from carefully neutralized solutions, show about 2 per
cent coprecipitation, so that a reprecipitation would probably reduce
this to satisfactory limits. The separation of zinc under these con-
ditions is unsatisfactory regardless of the method of neutralization.
It is to be noted that the coprecipitation in Procedure I is usually
greater than that in Procedure III, showing the favorable effect of
an increased volume and the thereby decreased concentration of the
coprecipitated substance. It is perhaps worthy of note that, al-
though it is almost universal in textbooks of qualitative analysis,
the procedure which directs that a slight excess of ammonia be added
to a relatively small volume of solution produces conditions which
are apparently the least favorable of those studied for the separations
desired.

Because the separation of the tripositive elements, as a group,
from the solution by an ammonia precipitation is so unsatisfactory,
under even the most closely regulated conditions, it is not used in
this system of analysis.

Other Possible Group Separations. If one considers the further
use of sulfide separations, an inspection of Table XI, p. 200, suggests
several possibilities. First, if the hydrogen ion concentration is
reduced from 0.3 m. to 0.01 m., of the elements remaining in the
solution zinc alone is precipitated. This method is not used, how-
ever, because it is not expedient to employ a procedure which may
take considerable time in order to detect and separate a single ele-
ment from the original solution. This element will often be absent,
and any system which generally used single separations would require
a large expenditure of time even when very few elements were present.

If the hydrogen ion concentration is reduced to approximately
10~ 6 m., nickel and cobalt sulfides are included with zinc; and by the
use of acetic acid and an acetate this hydrogen ion value can be very

288 AMMONIUM SULFIDE GROUP [P. 51

effectively maintained. The. formation of this group, including zinc,
nickel, and cobalt, would offer several advantages if it could be
quickly done and if clean separations could be obtained. However,
the complete precipitation of nickel and cobalt under these condi-
tions is slow, and, as is indicated in Table XI, p. 200, iron is likely to
be coprecipitated, while aluminum and chromium tend to hydrolyze
and precipitate under these same conditions.

The Precipitation by Ammonium Sulfide and Ammonium Hydroxide.
Finally, if the hydrogen ion concentration is reduced until the solu-
tion is alkaline, the hydroxyl ion concentration being approximately
10" 6 , practically complete precipitation of zinc, cobalt, nickel, iron,
and manganese as their sulfides, and of aluminum and chromium
as their hydroxides is obtained. The alkalinity of the solution is
usually adjusted by the use of ammonium hydroxide and ammonium
salts; hydrogen sulfide is then passed into the solution in a limited
amount, forming ammonium sulfide, or a solution of the latter may
be added. This method effects a satisfactory separation of the
above-mentioned elements from those of the Alkaline Earth and
Alkali Groups, except when phosphate is present; then the alkaline
earth elements may be precipitated as their phosphates and their
presence has to be provided for in the subsequent group analysis
(the absence of carbonate from the reagents has to be assured, as it
also causes precipitation of the Alkaline Earth elements). This
separation is very commonly used and has been adopted in this
system. Other sulfide separations in more alkaline solutions are
not practical, as the solubility of aluminum hydroxide is increased
and as magnesium hydroxide, which is not amphoteric, will be pre-
cipitated.

In the procedure which has been adopted, the filtrate from the
hydrogen sulfide precipitation is boiled until the hydrogen sulfide
is expelled, in order that the effect of adding ammonium hydroxide
alone may be noted. For, upon just neutralizing the solution, which
contains a considerable amount of ammonium salts, the formation
of a precipitate will indicate the presence of aluminum, chromium,
and iron (the latter upon heating the solution so that any ferrous
iron is partially oxidized to ferric). If phosphate and alkaline earth
elements are present, a precipitate of an alkaline earth phosphate
may also be obtained; manganese and cobalt may also be precipitated
as phosphates. Thus if no precipitate is obtained, the absence of
aluminum, chromium, and iron can be assumed and the analysis of
the Ammonium Sulfide Group is simplified.

P. 51] GROUP PRECIPITATION 289

The neutralized solution is then treated with hydrogen sulfide,
as zinc, nickel, and cobalt will be precipitated almost completely
from the neutral solution, and the precipitate will be much more
readily coagulated and filtered than if it were produced in an alkaline
solution. Possible information as the presence of these elements is
also obtained. The solution is finally made distinctly alkaline in
order to cause the complete precipitation of the sulfides of man-
ganese and iron.

Procedure 51: PRECIPITATION OF THE AMMONIUM SUL-
FIDE GROUP. Transfer the filtrate from the H 2 S precipita-
tion (P. 11) to a 500-ml conical flask and boil it until the
H 2 S is completely expelled (Note 1). While keeping the
solution just boiling, slowly add NH 4 OH until it is neutral
to litmus (Note 2), and then boil the mixture for 2 to 3 min.
(No precipitate, absence of aluminum, chromium, and of
alkaline earth elements as phosphates, probable absence of
iron. Note 3.)

Saturate the neutral solution with H 2 S. (If considerable
precipitate forms, add 1 ml of NH 4 OH to the solution and
resaturate it with H 2 S. Note 4.) Finally make the solu-
tion just alkaline with NH 4 OH, add 2 ml in excess, and pass
in a moderate stream of H 2 S (2 or 3 bubbles a second) for 15
sec. at a time until, after the contents of the flask are thor-
oughly mixed, the gas above the solution rapidly darkens
a strip of filter paper which has been moistened with
Pb(C 2 H 3 2 ) 2 solution. Cover the flask with a watch glass
(or loosely stopper it, Note 5), heat the mixture nearly to
boiling, shake it vigorously as long as the precipitate seems
to be coagulating, and allow the precipitate to settle. If
the solution is not alkaline to litmus, add NH 4 OH (0.5 ml
at a time) until it is, and again test for the presence of
H 2 S above the solution (Note 6). (No precipitate, absence
of Ammonium Sulfide Group elements.)

Filter the hot solution through a paper filter, using suction
if the precipitate is large, and wash the precipitate with two
to five 5-ml portions of a solution made by passing HjS for
10 sec. through a solution containing 0.5 ml of NH 4 OH and
0.5 g of NH 4 C1 in 25 ml of water. Add these washings to
the filtrate. Keep the solution hot during the filtration;
cover the funnel with a watch glass and collect the filtrate

290 AMMONIUM SULFIDE GROUP [P. 51

in a 400-ml flask provided with a two-hole rubber stopper,
inserting the stem of the funnel through one hole. Wash
the precipitate with three 5-ml portions of hot water, dis-
carding the washings (Notes 7, 8). Treat the precipitate
by P. 52 (Note 9). Immediately acidify the filtrate with
HC1, boil out the H 2 S (Note 10), and treat it by P. 81.

Notes:

1. To determine when the H2S has been completely expelled, smell
the escaping vapors or test them with strips of paper moistened with
Pb(C2H 3 02)2i as long as H^S is present, PbS will be formed and the strips
will darken.

2. The solution is kept boiling and an excess of NH 4 OH is avoided in
order to prevent the precipitation of manganese. In an alkaline solution
manganous salts are oxidized by air, and a brown precipitate (of hydrous
Mn20 8 ) is obtained. The ammonia should be added until litmus paper is
turned to an intermediate purplish color not to a distinct blue. If any of
the test papers mentioned in P. 3, Note 6, are available, they can be ad-
vantageously used for this neutralization and the solution can be adjusted
to a pH of 6.

3. Unless there is considerable iron present, ferrous hydroxide is not
precipitated under these conditions. However, there is such a pronounced
tendency for ferrous iron to be oxidized (partly because of the very slight
solubility of Fe(OH)3) that upon boiling the solution, even with small
amounts of iron present, a precipitate of ferric hydroxide will be obtained;
a black precipitate, consisting of a mixture of ferric and ferrous hydroxides,
may also result.

If only a small precipitate is obtained, and if phosphate is absent (see
Note 1, P. 54), considerable time may be saved by separating the tripositive
elements at this point. This can be done as follows:

Add HC1, 1 ml at a time and heating after each addition, until the
precipitate is dissolved. To the hot solution add saturated bromine
water, 1 ml at a time, until an excess is present (indicated by its
odor above the solution or color in the solution). Boil the solu-
tion vigorously (under a hood) until the excess of bromine is ex-
pelled and then for 2 min. longer. Keep the mixture just boiling
and neutralize with NHiOH as directed in the procedure above.
Filter while keeping the mixture hot and wash the precipitate with
hot water.

Treat the precipitate by the procedure given in paragraphs 2 to 8
of Note 3, P. 55, for the separation and detection of small amounts
of iron, aluminum, and chromium. Treat the filtrate as directed in
the second and third paragraphs of the procedure above in order
to detect and precipitate the bipositive elements.

4. If, when the neutral solution is saturated with H2S, a white precipitate
is obtained, the presence of zinc is indicated; a black precipitate will indicate
nickel or cobalt. If the color of the solution has indicated the presence of

P. 52] SEPARATION OF IRON 291

nickel, it is advisable to avoid admitting air to the flask (by keeping a con-
stant flow of HaS) and separately to filter and wash the precipitate obtained
from the neutral solution. This is done because nickel sulfide, when pre-
cipitated from an alkaline solution containing disulfide (usually formed by
oxidation of the sulfide by the air), tends to form a dark brownish colloidal
solution which is exceedingly difficult to coagulate. The precipitation of
the remaining elements is then made in an alkaline solution as directed, and
the two precipitates are united for treatment by P. 52.

5. The solution and precipitate are protected from exposure to air during
the precipitation and filtration in order to avoid oxidation of the sulfide with
formation of disulfide. This not only causes nickel sulfide to become col-
loidal but increases the amount of sulfate which is later formed; the latter
may cause the premature precipitation of barium and strontium.

6. The solution should be distinctly alkaline, but a large excess of am-
monia should be avoided; the solubility of aluminum hydroxide becomes
appreciable under such conditions, and, unless the ammonium hydroxide
solution has recently been distilled, it is likely to contain carbonate (by
absorption of carbon dioxide), which may cause the partial precipitation of
alkaline earth elements as carbonates.

7. The precipitate is finally washed with hot water in order to remove
ammonium salts, which are likely to precipitate as NH 4 C1 in the HCl-ether
treatment of P. 52.

8. If nickel is present and a brownish colloidal filtrate is obtained, this
can usually be coagulated by making the filtrate just acid with HC1, heating
it almost to boiling, again making it just neutral with NHUOH, and saturating
it with H2S. This precipitate is filtered separately and combined with the
first one.

9. If phosphate has not been present and the tripositive elements have
been precipitated and removed by ammonium hydroxide as suggested in
Note 3, this precipitate (which contains the bipositive elements of this
group) should be dissolved as directed in the first paragraph of P. 52, and
the solution should then be treated as directed in P. 55.

10. The filtrate is at once made acid and the H^S is expelled in order to
avoid oxidation of the sulfide by air, with final formation of sulfate. This
takes place to a considerable extent if the alkaline solution is allowed to
stand for any length of time. If a white precipitate of sulfur forms, it should
be coagulated, filtered out, and discarded.

P. 52. Detection and Separation of Iron

Discussion. The Ammonium Sulfide Group precipitate having
been obtained, it would be logical next to separate this large group
of elements into two smaller groups. However, iron is detected
and separated singly, as this can be done with so little loss of time
and so effectively from the small volume of solution which is here
obtained. When the ammonium sulfide precipitate is dissolved in
hydrochloric acid, the presence of iron is indicated by the yellow

292

AMMONIUM SULFIDE GROUP

IP. 52

color which it imparts to the solution. If the color of other elements
masks this effect, a simple test with thiocyanate is made on a portion
of the solution and this portion is then combined with the main
solution.

The separation of iron which is used here is based upon experi-
ments 11 which have shown that when a solution which contains 125
to 250 mg of iron and which is 5.5 to 9.0 n. in hydrochloric acid is
shaken with an equal volume of isopropyl ether, over 90 per cent
(99 per cent at 6.5 to 8.5 n. hydrochloric acid) of the ferric iron
passes into the ether. From the data collected in Table XX it is
seen that of the elements of the Ammonium Sulfide Group none
passes into the ether in appreciable quantity (more than 0.01 per
cent).

TABLE XX

BEHAVIOR OF THE AMMONIUM SULFIDE GROUP ELEMENTS ON SHAKING A 7.75 N.

HYDROCHLORIC ACID SOLUTION WITH AN EQUAL VOLUME OF

ISOPROPYL ETHER

Element

Amount Taken
(mg)

Amount Extracted
(mg)

Per Cent
Extracted

Co"

500

0.0

500

0.00

500

0.00

500

0.00

Cr in

500

<0.03

<.01

Zn"

500

0.0

P (as H,P0 4 )

125

0.1

/P (as H,P0 4 )

62.5

\Fe nl

250

Fe"

250

0.0

Fein

248.8

248.6

99.9

It has long been known that ferric iron could be extracted with
ethyl ether from solutions approximately 6 n. in hydrochloric acid,
and thus could be separated from all the elements of the Ammonium
Sulfide, Alkaline Earth, and Alkali Groups, 12 Although such an

11 Dodson, Forney, and Swift, J. Am. Chem. Soc., 58, 2573 (1936).

12 For discussions of this separation with ethyl ether, see Rothe, Stahl u.
Eisen, 12, 1052 (1892); Ledebur, ibid., 13, 333. (1893); Langrauir, /. Am.
Chem. Soc., 22, 102 (1900); Kern, ibid., 23, 689 (1901); Sprier, Chem. News,
83, 124 (1901). For a table showing the behavior of the chlorides of the less
common metallic elements in this treatment, see Swift, /. Am. Chem. Soc.,
46, 2378 (1924).

P. 52]

EXTRACTION OF IRON BY ETHER

293

extraction process offers obvious advantages over precipitation
methods, this separation has not been used as widely in general
qualitative and quantitative procedures as would seem justified.
This has been due to several factors: The efficiency of the extraction
process is critically dependent on the hydrochloric acid concentra-
tion; the solubility of ether in the acid is so large that the volume of
the aqueous phase is increased by as much as 25 per cent, making

2468

Initial HCI Concentration,
Moles per Liter

Fig. 30. Extraction of Ferric Chloride from Aqueous Hydrochloric Acid
Solutions by an Equal Volume of Ethyl and of Isopropyl Ether. --, ethyl
ether; -o~, isopropyl ether.

the concentration of the acid after an extraction uncertain or requir-
ing the use of acid previously saturated with the ether; it has been
feared that peroxide or alcohol, which are frequently present in
ether, would reduce the iron to the ferrous state, in which form it is
not extracted; and, finally, the volatility of ethyl ether makes the
quantitative technique of the extraction somewhat difficult at room
temperatures and involves a serious fire hazard, especially for stu-
dent use.
Because of these facts the behavior of isopropyl ether was investi-

294 AMMONIUM SULFIDE GROUP [P. 52

gated, 18 and it was found that more complete extractions could be
made over a wider range of acid concentrations than could be made
with ethyl ether. Curves showing the percentage of iron extracted
by these two ethers at various acid concentrations are shown in
Fig. 30. It is seen that an equal volume of isopropyl ether extracts
more than 99 per cent of the ferric iron from solutions which are
from 6.5 to 8.5 n. in hydrochloric acid. With ethyl ether 99 per
cent is extracted only when the acid is approximately 6.2 n., and the
efficiency of extraction diminishes rapidly at higher concentrations,
owing in part to the more rapid increase in the solubility of the latter
ether in the more concentrated acid.

Because of the above facts and also because isopropyl ether is less
volatile and therefore can be used with less danger of mechanical
loss in separating funnels and with less fire hazard, it is recommended
for use in the procedure below. 14 Should it not be available, pro-
vision is made in Note 8 of this procedure for the substitution of
ethyl ether. This extraction separation is most advantageously
used when a large amount of iron is present, as it entirely avoids the
coprecipitation errors inherent in a precipitation method and, if
properly carried out, is quite precise. However, the operations in-
volved have to be carefully performed, or loss of the solutions may
result ; it is also necessary that the concentration of the hydrochloric
acid be closely controlled. The time required is usually less than
that required to filter and wash an iron hydroxide precipitate. If
only a small amount of iron is present, it is recommended that the
ether extraction be omitted and the iron be separated in the optional
procedure provided in P. 55, Note 3, or allowed to precipitate with
the other Zinc Group sulfides. It will remain in the solution when
the zinc is precipitated as sulfide in P. 62 and can be precipitated as
directed in the notes to P. 63.

Procedure 52: DETECTION AND SEPARATION OF IRON.
Transfer, by means of a stirring rod, as much as possible of
the ammonium sulfide precipitate (P. 51) from the filter to
a 100-ml flask. Transfer any precipitate remaining in the
original flask to this flask with the aid of 10 to 25 ml of warm
HC1, heat the mixture as long as the residue seems to be

11 Dodson, Forney, and Swift, loc. cit.

14 The technical grade of isopropyl ether has been found to be satisfac-
tory for this separation; it is less expensive than a comparable grade of ethyl
ether.

P. 52] SEPARATION OF IRON 295

dissolving, and pour this solution through any precipitate
left on the filter. Finally, add liquid Br 2 to the solution, a
drop at a time, until the residue becomes light colored and
an excess is present (Note 1) ; also pour this solution through
any residue left on the filter. Wash the filter and residue
with 2 to 5 ml of HC1. Filter out any residue on a small
filter and wash it with HC1 (Note 2). Add to the filtrate
2 ml of 12 n. HC1, evaporate the solution to a volume of
approximately 10 rnl (or until the excess Br 2 is expelled),
and cool it (Note 3).

If the solution is colorless, treat it by P. 54 (Note 4).

If the solution is colored, remove with a dropper 0.5 ml
to a small test tube and add 0.1 ml (3 drops) of 1 n. KSCN.
(Pink or red color, presence of iron Note 5.) Pour this
solution back into the main filtrate and wash the dropper
and test tube with 2 ml of 6 n. HC1.

If iron is not present, treat the solution by P. 54.

If iron is present, treat the solution by the next para-
graph (Note 6).

Pour the cold HC1 solution into a 50-ml separating funnel
(Note 7), wash out the flask with two 2-ml portions of 12 n.
HC1, and add to the solution an equal volume of isopropyl
ether (Note 8). Use the ether to wash out the flask further.
(Caution: Ether is very inflammable. Do not carry out these
operations near aflame.) Cool the mixture (Note 9), shake
it vigorously for 5 to 10 sec., and then allow the layers to
separate completely (Note 13). Again cool the mixture
and draw off all the water layer into a second separating
funnel (Note 10). Add 1 ml of 7.5 n. HC1 (made by mixing
equal volumes of 6 n. and 9 n. acids) to the ether in the first
funnel, rinsing the stopper as it is added, again shake the
mixture, let the layers separate, and add this water layer to
the first. Draw off the ether layer into a flask, washing out
the funnel with two to five 1-ml portions of 1 n. HC1. Re-
peat these operations with the water layer in the second
funnel, using only 5 ml of ether each time, until the ether
layer remains colorless (Note 11). Draw off the final water
layer into a 200-ml flask and treat it by P. 54 if phosphate is
present, or by P. 55 if phosphate is absent. Return the
previous ether extracts to the funnel containing the final
one and treat these by P. 53 (Note 12).

296 AMMONIUM SULFIDE GROUP [P. 52

Notes:

1. When the precipitate from P. 51 is treated with HC1, the sulfides of
manganese, iron, and zinc are readily dissolved with evolution of H2S; how-
ever, the sulfides of nickel and cobalt dissolve so slowly that an oxidizing
agent is required to bring them into solution. The Br2 is not added at first
because its action would cause the formation of considerable sulfate; this
is undesirable, as barium or strontium might be present. It is also necessary
for the iron to be in the ferric form for it to be extracted by the ether, and,
as it would have been reduced by the H2S first set free, Br2 is added until
an excess is present, thus insuring that the iron is oxidized.

2. If there are any sulfides present, there is usually a small residue of
sulfur left after this treatment. If the mixture is heated until this is light
colored, it will contain only insignificant amounts of the metallic elements
and can be filtered out and discarded.

3. It is necessary that the Br 2 be completely expelled, as it would oxidize
the thiocyanate which is later added according to the reaction:

SON- + 3Br 2 + 4H 2 - HCN + HS0 4 - + 6H+ + 6Br-.

It is necessary that the volume be approximately 10 ml, as the adjustment
of the acid concentration for the ether extraction is based upon this volume.
Upon evaporation of the solution, the acid approaches 6 n. the constant-
boiling solution. There are added 2 ml of 6 n. and 6 ml of 12 n. acid, giving
a solution ,of approximately 8 n. HC1. Ferric bromide solutions are orange-
yellow, and therefore the complete expulsion of the Br2 can best be deter-
mined by smelling the vapors above the boiling solution.

4. Even 1 mg of iron will impart a distinct yellow color to this volume of
6 n. HC1; if this distinctive yellow color is obtained, it is not necessary to
make the thiocyanate test. If any chromium is present, it may make the
detection of iron uncertain because of the green color which it causes. One
mg of chromium will also cause an easily perceptible color, so that, if the
solution is colorless, chromium need not be further tested for.

5. The presence of 0.5 mg of iron in the original solution will cause a
distinct pink color in the solution being tested. The color change is notice-
able even when a large quantity of chromium is present.

6. As mentioned in the discussion, if it is not desired to use the ether
separation, especially if only a small amount of iron (less than 10 mg) is
thought to be present, this solution can be carried directly to P. 54 (or to
P. 55 if phosphate is absent) and the iron present can be separated and
detected as indicated in Note 3 of P. 55, or it may be later separated as
Fe(OH) 3 and Fe(OH) 2 C 2 H 3 O 2 in P. 63.

7. A separating funnel with the stem cut off as close to the stopcock as
possible is desirable because of the difficulty in removing the solution from
a long stem.

8. As mentioned in the discussion, it is possible to substitute ethyl ether
for the isopropyl ether when the latter is not available. In this case only
1 ml of 12 n. and 4 ml of 6 n. HC1 should be used for washing out the flask:
otherwise the procedure remains the same.

When working with ethyl ether, extreme care should be taken to extinguish

P. 52] SEPARATION OF IRON 297

all flames in the vicinity; also, the operations should be carried out near a
hood so that the ether vapor does not accumulate where it may be subse-
quently ignited.

9. Unless the mixture is kept cold, the ether develops a pressure in the
funnel and tends to force the solution out around the stopper and stopcock.
For this reason, the funnel should be cooled under running tap water or,
preferably, with ice water. Since the volume of solution is small, great care
should be taken to avoid loss during this operation.

10. The successive shaking with ether can be more easily carried out by
drawing the water layer directly into a second separating funnel, thus avoid-
ing additional transfers of the solution.

11. The color of the iron in the ether layer is so distinctive that when this
layer is colorless there remains so little iron (less than 0.5 mg) in the water
layer that it will not interfere with any of the subsequent separations; the
aqueous layer may be colored by other constituents, thus, titanium (often
present in ores) forms a yellow compound with peroxide (present in the
ether) very similar in color to that of ferric chloride solutions.

12. No information as to the state of oxidation of the iron in the original
material is obtained from its detection in the above procedure. In case this
information is desired, proceed as follows:

Boil 10 ml of 6 n. H2&O4 for 2 to 3 min. in a flask which is covered
with a watch glass; then quickly introduce 0.1 g of the original
sample and again boil for 5 min. Add, in small portions, 0.2 g of
NaHC0 3 and cool the flask, with running tap water, to room tem-
perature. Filter the mixture through a rapid-filtering paper filter
and collect the solution in a flask to which has been added 0.1 g
of NaHC0 3 .

Pour half of the solution into a small flask (or test tube) and add
to it, 0.1 ml at a time and as long as a precipitate forms, K 3 Fe(CN) 6
solution. (Blue precipitate, presence of. ferrous iron.)

Pour the other half of the solution into a test tube containing
5 ml of KSCN solution. (Red color, presence of ferric iron.)
The H2S04 solution is first boiled in a covered flask, and the NaHCOg
is added to insure that any ferrous iron present is not oxidized by oxygen in
the solution or in the air above it. For a discussion of the precipitates
formed by ferrous and ferric iron with ferro- and ferricyanide, see P. 133.
If silver or copper were present, they would give orange and green precipi-
tates, respectively; the blue compound formed with ferrous iron is so very
insoluble and intensely colored that even small amounts can be recognized
in the presence of these other precipitates, especially if the ferricyanide is
added gradually so that the blue precipitate is formed first. The red-
colored compound formed by thiocyanate with ferric iron is a very distinctive
test and has been used in the above procedure; a large excess is added to
increase the sensitivity of the test and to provide an excess in case any
oxidizing agents (which might oxidize thiocyanate to sulfate and cyanide)
are present.

13. When isopropyl ether is used and large amounts of iron are present,
two ether layers may form; both of these layers should be retained in the
funnel.

298 AMMONIUM SULFIDE GROUP [P. 63

P. 53. lodometric Estimation of Iron

Discussion. The ferric chloride is recovered from the ether solu-
tion by shaking it with water. As is seen from the curves shown in
Fig. 30, under these conditions the iron will pass quantitatively into
the aqueous layer, and the ether can be recovered and reserved for
future use. The method for estimating iron which is recom-
mended for use here depends upon the fact that ferric chloride in a
hydrochloric acid solution oxidizes iodide to free iodine and is
itself reduced to the ferrous state; the iodine can then be titrated
with a standard thiosulfate solution. The principal reaction can
be represented by the equation

and from the potentials involved it can be calculated that the reac-
tion would reach an equilibrium with a considerable amount of ferric
iron present if iodide were added in only an equivalent amount.
However, it can be shown that the equilibrium would be so shifted
by the addition of an excess of iodide that less than 0.2 per cent of
the iron would remain in the ferric form; this amount should be
greatly reduced as the iodine is removed during the titration with
the thiosulfate. Several factors not shown by the above equation
have to be considered: First, the rate of the reaction is somewhat
slow, and therefore it is necessary to allow the iodide and ferric iron
to react in a small volume for an appreciable length of time before
beginning the titration. Second, although it is not evident from the
equation given above, an experimental study of the reaction has
shown that the accuracy of the method is dependent upon the
hydrochloric acid concentration. With insufficient acid the reaction
becomes slow and incomplete, because of the hydrolysis of the ferric
iron; with the hydrochloric acid concentration too high the reaction
is also incomplete, probably because of the formation of complex
molecules of the type HFeCU (such salts as KFeCU having been
prepared). In addition to these effects, oxidation of the iodide by
the oxygen of the air becomes appreciable as the acid concentration
is increased. It is thus indicated that the accuracy of the process
is very dependent upon the conditions obtaining and to a certain
extent upon a compensation of errors. It has been found experi-
mentally that, if the reaction is allowed to proceed under the condi-
tions given in the procedure below, results accurate to 0.2 to 0.3

P. 53] IODOMETRIC ESTIMATION OF IRON 299

per cent can be obtained. 16 If sulfates are present, higher concen-
trations of acid and of iodide are required; the reaction can be re-
versed by the addition of phosphoric acid. The reaction of phos-
phoric acid with ferric iron is discussed in P. 53 A.

Iron is more commonly determined by titration with an oxidizing
agent, most frequently standard permanganate solution. This re-
quires that the iron be in the ferrous state, and various reagents and
methods are used for this preliminary reduction (see the discussion
of P. 53 A). These methods are often time-consuming, and special
precautions have to be taken to prevent oxidation of the ferrous salt
after the reducing agent has been removed. Also, the titration of
ferrous iron with permanganate in the presence of chloride has to be
carried out under very closely controlled conditions, or appreciable
reduction of permanganate by the chloride may occur.

As in this procedure the iron is already in the ferric form and in a
hydrochloric acid solution, and as the iodometric method is so very
rapidly carried out, its use is recommended. However, as the titra-
tion with permanganate in a hydrochloric acid solution is so fre-
quently used for the analysis of iron ores, it is possible that the
special reagents required may be available and that this method may
be preferred; therefore a procedure for this method is given in P. 53 A.

Procedure 53: IODOMETRIC ESTIMATION OF IRON. Add
an equal volume of water to the ether extracts, cool, shake,
allow the layers to separate, and draw the water layer into
a 200-ml flask. Repeat this process with two 5-ml portions
of water. Reserve the ether solution for future extractions.

Evaporate the aqueous extracts on a steam bath (Caution:
Note 1) until all dissolved ether is expelled, and then evap-
orate over a flame until just 15 ml of the solution remain.
Estimate this volume by comparing the solution left in the
flask with a measured 15-ml volume of water in a similar
flask. Cool the flask with tap water and add to it 4 ml
of HC1.

If less than 75 to 100 mg of iron are thought to be present,
treat the solution as directed in the last paragraph of this
procedure.

If more than 75 to 100 mg of iron are thought to be pres-
ent, transfer the cold solution to a 100-ml volumetric flask,

11 Swift, J. Am. Chem. Soc., 51, 2682 (1929).

300 AMMONIUM SULFIDE GROUP [P.

dilute it to the mark, and thoroughly mix the solutions.
Pipet 25 ml of this solution into a 200-ml flask, preferably
one with a ground-glass stopper, add to it 4 ml of HC1,
and treat it as directed in the next paragraph.

Add to the solution 3 g of solid KI, close the flask, gently
swirl the mixture until the KI is dissolved, and allow it to
stand for 5 min. (Note 2). Dilute the solution to 100 to
125 ml with cold water, and rapidly titrate with 0.1 n.
Na2S20a solution as long as the iodine color is apparent.
When the iodine color becomes indistinct, add 5 ml of starch
indicator solution and carefully titrate until the blue color
just disappears (Note 3). From the volume of thiosulfate
solution used, calculate the .amount of iron present.

Notes :

1. As some ether may remain in these solutions and as its vapors are very
inflammable, the first part of this evaporation should be performed under a
well-ventilated hood or on a steam bath heated with an electric heater (it
is usually sufficient to heat a large beaker of water to boiling and then ex-
tinguish the flame before beginning to evaporate the ether). See Note 8,
P. 52.

2. The results obtained by this method are dependent upon the amount
of HC1 in the solution, upon the volume during the reaction between the
ferric iron and the iodide, upon the amount of potassium iodide added, and
upon the length of time for which the solution is allowed to stand. The
effect of the acid is mentioned in the discussion. Increasing the volume of
the solution decreases the concentration of acid, ferric iron, and iodide, and
slows the rate of reaction. With insufficient iodide the reaction is again
slow; with too much the oxidation of the iodide by air is increased. If
insufficient time is allowed, the reaction may not be complete when the end-
point is taken; if the solution is allowed to stand too long, oxidation of the
iodide by the air takes place.

3. Both ferrous chloride and ferrous iodide are oxidized by air in a HC1
solution, and the rate is dependent upon the concentration of the HC1; for
this reason the solution is diluted before making the titration. The titration
should be carried out as rapidly as possible, and after reaching an end-point
it is advisable to stopper the flask and allow it to stand for 2 to 3 rnin. If a
blue color reappears, it can usually be permanently removed with less than
a full drop of the thiosulfate.

P. 53A. Optional Method for the Estimation of Iron by Titration

with Permanganate

Discussion. This method is based upon the oxidation of ferrous
iron, in a hydrochloric acid solution, with a standard permanganate
solution.

P. 53A] PERMANGANATE ESTIMATION OF IRON 301

The Reduction of Ferric Salt Solutions. In the procedure below,
the iron is obtained in solution in the ferric state (as is usually the
case in the analysis of iron ores) and must be reduced before the
titration. This reduction can be accomplished by (1) metals (such
as zinc, cadmium, or aluminum), (2) reducing gases (such as hydrogen
sulfide or sulfur dioxide), or (3) solutions of a reducing agent (usually
stannous chloride). After such a reduction the excess of the reduc-
ing agent must be removed. In the case of metals this may involve
a filtration or the use of a special reduction apparatus, 16 and the
gases have to be expelled by long boiling or swept out with an inert
gas. If the iron- is in a hydrochloric acid solution (as is frequently
the case because of the effectiveness of this acid in dissolving iron
ores), the reduction can be very quickly carried out with a stannous
chloride solution and the excess stannous tin can be removed by
oxidation with mercuric chloride. This reagent is uniquely fitted
for the purpose, as (1) mercuric chloride does not oxidize ferrous
iron, and as (2) the reduction product (mercurous chloride) is so
insoluble that, especially if it is caused to precipitate in a crystalline
form, it is not appreciably oxidized by either ferric iron or permanga-
nate during the course of the titration.

The Titration of Ferrous Salts in Hydrochloric Acid Solutions. As
was mentioned in the general discussion of permanganate methods,
it would be calculated from the potentials involved that, in a strongly
acid solution, permanganate, even in the small concentration neces-
sary to give an end-point, would be reduced by chloride. Fortu-
nately, the rate at which this reaction takes place is so slow under
most conditions that such titrations can be carried out; thus oxalate
(P. [85 and P. 87) and ferrocyanide (P. 133) are titrated in the
presence of relatively high concentrations of hydrochloric acid.
However, it was realized quite early in the history of this titration 17
that, when ferrous salts in hydrochloric solutions are titrated with
permanganate, there is oxidation of the chloride.

Investigations have shown that this effect is due to the reaction
between permanganate and ferrous iron inducing the oxidation of
chloride; ferric ions do not catalyze the reaction between permanga-

l< This is usually the " Jones Reductor," which is a vertical glass tube
filled with amalgamated zinc. The solution to be reduced is passed through
this tube, the rate of flow being controlled by a stopcock at the bottom end.

17 The ferrous iron-permanganate titration appears to be the first per-
manganate method to have been used (Marguerite, Ann. Chem. P/iys., 18,
244 (1846)). The first mention of the effect of chloride was made bv Lowenthal
and Lenssen, Z. anal Chem., 1, 329 (1862).

302 AMMONIUM SULFIDE GROUP [P. 63A

nate and chloride. Experimental evidence has been obtained 18 that
this induction is caused by the permanganate first oxidizing the
ferrous iron to an unstable higher oxidation state, probably ferrate
(Fe0 4 "), as follows:

MnOr + Fe+ + - Fe0 4 "" + Mn+++.

The ferrate then oxidizes chloride ion to hypochlorous acid, it having
been shown that the oxidized chlorine is largely present as this
compound. 19 It was also observed quite early 20 that this effect
was decreased by the presence of a relatively high concentration of
manganous salts. The mechanism of the inhibiting effect of the
manganous ion has not been clearly established. In partial expla-
nation the suggestion has been advanced that raising the concen-
tration of the manganous ion decreases the oxidizing potential of the
permanganate manganous ion half-cell. However, the change in
potential which can be thus produced is so limited that the validity
of this explanation is doubtful. It is more probable that the effect
is one in which the mechanism of the reaction is changed by the
presence of the manganous ion. Thus the manganous ion, by react-
ing rapidly with the permanganate to give an intermediate oxidation
state or states (as is the case in the permanganate-oxalate catalysis),
may eliminate the formation of the higher oxidation states of iron;
or, alternatively, the ferrate (or the hypochlorous acid which it
forms) may react very rapidly with manganous ion to give tri- or
quadripositive manganese, which then rapidly oxidizes ferrous iron
to the ferric state. Zimmermann 21 appears to have been the first
to have made practical use of this phenomenon by the addition of
large amounts of manganous salts to the solution to be titrated.

Ferric iron in hydrochloric acid solutions causes an intense yellow-
ish color (due to complex compounds of the type HFeCU) which
obscures the permanganate color and necessitates the addition to
such solutions of a higher concentration of permanganate in order
to obtain an end-point. This yellow color can be bleached by the
addition of phosphoric acid, which forms more stable (but colorless)
complex compounds with ferric iron; 22 sulfuric acid forms similar

u Manchot, Ann. Chem. Pharm., 325, 105 (1902); Bohnson and Robertson,
/. Am. Chem.oc., 45, 2493 (1923); Hale, /. Phys. Chem., 33, 1633 (1929).

19 Baxter and Frevert, Am. Chem. /., 34, 109 (1905).

"Kessler, Ann. J. Chem., 118, 17 (1863).

11 Zimmermann, Ann. d. Chem., 213, 305 (1882).

"These were thought by Weinland and Ensgraber, Z. anorg. Chem., 84,
340 (1913), to be complex acids of the type HsFetPO^. However, the experi-

P. 534] PERMANGANATE ESTIMATION OF IRON 303

but less stable compounds. Thus, by the addition of these acids a
colorless solution and a much sharper end-point are obtained. The
addition of phosphoric acid may decrease the tendency toward
chloride oxidation, as, by decreasing the ferric ion concentration, it
facilitates the oxidation of ferrous iron and, by bleaching the solu-
tion, decreases the concentration of the permanganate required for
recognition of the end-point. It has been found experimentally
possible, under rather closely controlled conditions, practically to
eliminate the chloride oxidation by the use of large amounts of
phosphoric acid or of phosphates. 23 Reinhardt 24 suggested that
manganous sulfate, phosphoric acid, and sulfuric acid be combined
in one reagent, and this solution is still quite universally used and
known as the Zimmermann-Reinhardt "preventative" solution.

By the use of this solution, by avoidance of an excess of stannous
chloride and therefore of mercurous chloride precipitate, and by slow
titration in a cold dilute solution, the method can be made to give
quite precise results, and, because of the rapidity with w r hich it can
be carried out, it is very extensively employed, especially for the
analysis of iron ores.

Procedure 53A : ESTIMATION OF IRON BY TITRATION WITH
PERMANGANATE. Treat the ether extracts as directed in the
first paragraph of P. 53.

Evaporate the aqueous extracts on a steam bath (Caution:
Note 1, P. 53) until any ether is expelled, add 0.2 f. KMn04
until a distinct reddish-pink color is obtained, and then add
2 or 3 drops in excess (Notes 1,2).

Evaporate the solution to a volume of about 15 ml and,
while keeping it hot, add 1 n. SnCU drop by drop, swirling
the solution, until the yellow color disappears and the solu-
tion is colorless (or only a very slight greenish tinge remains),
and then add 1 drop of SnCU in excess.

Cool the solution to at least room temperature and add
rapidly 10 ml of a saturated mercuric chloride solution
(Note 3). Allow the mixture to stand 2 min. (Note 4),
transfer it to a 600-ml beaker containing 400 ml of water
and 25 ml of Zimmermann-Reinhardt solution (Note 5),

ments of Bonner and Romeyn, /. Ind. Eng. Chem., Anal. Ed., 3, 85 (1931),
do not confirm this but indicate that an un-ionized molecule of the type
Fe(H a PO 4 )i is formed.

" Hough, /. Am. Chem. Soc., 32, 539 (1910); Barneby, ibid., 36, 1429 (1914).

" Reinhardt, Chem. Zeit., 13, 323 (1889).

304 AMMONIUM SULFIDE GROUP [P.

and titrate immediately with standard KMn0 4 . Do not
add the KMn0 4 rapidly at any time. Stir the solution
continuously and approach the end-point drop by drop.
Take the end-point when the first perceptible pink tinge
spreads uniformly throughout the solution and persists for
at least 15 sec. (Note 6). Use a white background for
viewing the solution and avoid overrunning the end-point.
Make an end-point correction (Note 7). From the cor-
rected volume of standard permanganate used, calculate the
amount of iron present.

Notes:

1. The KMn0 4 is added to oxidize any reducing materials (organic com-
pounds or peroxide) which may have been introduced into the solution by
the ether separation. The excess permanganate is reduced by the chloride,
and the chlorine thus formed is expelled during the subsequent boiling or
is reduced by the SnCl2 next added.

2. Many iron ores can be brought into solution by treatment with HC1
and the iron can be determined by directly treating the resulting solution
by the procedure given above; elements which would cause errors in the
titration (vanadium, molybdenum, tungsten, and platinum) are not com-
monly present in the iron ores. Proceed as follows:

Weigh out an amount of the ore which will require from 25 to
35 ml of the standard permanganate, add to it in a 150-ml beaker
20 ml of 6 n. HC1 and 1 ml of SnCl2, cover the beaker, and heat
the mixture almost to boiling until only a white residue remains
(if the solution becomes yellow, add SnCU in 0.1-ml portions until
it is decolorized). When solution is complete, add 0.2 f. KMn0 4
until the first yellow color is obtained and treat the solution by
the second and subsequent paragraphs of the procedure above.
The presence of an excess of SnCU greatly increases the rate at which

iron ores are dissolved by HC1, probably by causing a surface reduction of

the iron.

3. The solution is cooled and the HgCU solution is added all at once in
order to prevent the formation of metallic mercury. If a large excess of
SnCl2 is added, or if the solution is warm, this may occur and will cause a
black precipitate. In this case the analysis must be discarded, as metallic
mercury rapidly reduces both permanganate and ferric iron.

4. The mixture is allowed to stand for 2 min. in order that complete pre-
cipitation of the mercurous chloride may occur and so that it may change
from the finely divided, often colloidal, form in which it first separates to
a more compact crystalline form. In this state it is oxidized much more
slowly by ferric iron or permanganate.

The solution should not stand much longer than this time, or oxidation
of the ferrous iron by the oxygen of the air will become appreciable. For
the same reason the solution should be titrated immediately after adding
the Zimmermann-Reinhardt solution.

P. 54] REMOVAL OF PHOSPHATE 305

5. The Zimmermann-Reinhardt solution is approximately 0.3 f . in MnS0 4 ,
3 f . in H3P04, and 4.6 f . in H2S04. See the Appendix for directions for its
preparation.

6. The end-point will slowly fade because of the oxidation of the mer-
curous chloride, but, if the above conditions have been adhered to, it will
persist for at least 15 sec. and usually considerably longer.

7. The end-point correction should be made by taking the volume of
hydrochloric acid and of stannous chloride originally used, oxidizing the
stannous chloride with the 0.2 f. KMn(>4, then decolorizing this solution
with stannous chloride, and carrying out the remainder of the operations as
in the procedure above. This correction will usually vary between 0.03
and 0.05 ml of 0.1 n. KMn0 4 .

P. 64. Removal of Phosphate

Discussion. If phosphate has been found (in P. 164, or by the
test given in Note 1 below) to be one of the constituents of the
material being analyzed, and if a precipitate was obtained on adding
ammonium hydroxide to the filtrate from the hydrogen sulfide
precipitate, this may have been due to the precipitation of one or
more of the alkaline earth elements as phosphate. In order to
recover the alkaline earth elements, and because phosphate might
cause the precipitation of manganese with the Zinc Group, a sup-
plementary procedure is necessary to remove phosphate before
separating the Zinc and Aluminum Groups.

The method used here depends upon the fact that bismuth phos-
phate is the only phosphate of the common elements which is only
slightly soluble in 0.5 n. nitric acid. Therefore, upon the addition
of bismuth nitrate to such a solution, the phosphate is precipitated
and the other basic elements remain in solution. The solubility
of bismuth phosphate is much greater in hydrochloric acid because
of the formation of compounds such as HBiCU; therefore, because
of this fact and because of the slight solubility of BiOCl, chloride
must be removed before the precipitation is made; this is done by
evaporating the hydrochloric acid solution with nitric acid. If it is
desired to estimate the amount of phosphate present, the precipitate
of bismuth phosphate can be collected on a previously heated and
weighed Gooch-type crucible, and the precipitate can be dried at
400 to 500C. and weighed as BiPCX. Confirmatory experiments 25
have shown that, of the basic elements which may be present in this
solution, less than 1 mg of iron, aluminum, cobalt, nickel, or zinc
is carried out with the precipitate when 500 mg of phosphate are

" Unpublished experiments by Dr. R. C. Barton.

306 AMMONIUM SULFIDE GROUP IP. M

present with 250 mg of any one of these elements. However, with
the same quantity of chromium, the normally white BiPO* precipi-
tate is greenish in color and may contain as much as 12 to 15 mg
of chromium.

The excess of bismuth is removed by precipitating it as sulfide
from the solution, which is 0.5 n. in nitric acid. After filtering out
the sulfide precipitate, the Ammonium Sulfide Group elements can
be precipitated (by P. 51) and the Alkaline Earth Group elements
are left in the filtrate.

Procedure 54: REMOVAL OF PHOSPHATE. If either (a)
phosphate has not been found present (Note 1) or (b) if no
precipitate was obtained on adding NH 4 OH to the filtrate
from the H 2 S precipitation (P. 51), treat the HC1 solution
from P. 52 by P. 55. If phosphate has been found present
and if a precipitate was also obtained on adding NH^OH to
the filtrate from the H^S precipitation, treat the HC1 solu-
tion as directed in the next paragraph (Note 2).

Evaporate the HC1 solution from P. 52 (Caution: It con-
tains ether. See Note 1, P. 53) to 1 to 2 ml (colorless solu-
tion, absence of chromium, cobalt, and nickel. Note 3.)
Add 5 ml of 16 n. HN0 3 and evaporate the solution to 1 to
2 ml. Add 5 ml of 16 n. HNO 3 and again evaporate the
solution to 1 to 2 ml (Note 4). Dilute the solution to 50
ml and add 3 n. Na 2 CO 3 solution, 0.5 ml at a time, until the
solution is neutral (Note 5) or slightly alkaline to litmus.
(White precipitate, presence of aluminum, manganese,
zinc, or alkaline earth elements. Colored precipitate,
presence of chromium, cobalt, or nickel. Note 6.) Add
4 ml of 6 n. HN0 3 , heat the solution nearly to boiling, and,
while keeping it hot, add dropwise to it 0.5 ml portions of
0.1 f. bismuth nitrate reagent until a precipitate no longer
forms (Note 7). Boil the mixture for a minute or two, keep-
ing the flask in motion to prevent bumping, and then let it
stand until the precipitate has settled. Filter the mixture
by decantation through an asbestos filter (Note 8) and wash
the precipitate with three 5-ml portions of hot 0.3 n. HNO 3 ,
collecting the washings with the filtrate. Saturate the
warm filtrate with H 2 S and filter the mixture. Wash the
precipitate with three 5-ml portions of hot water, collecting
the washings with the filtrate. Discard the precipitate
(Note 9). Treat the solution by P. 51 (Note 10).

P. 54] REMOVAL OF PHOSPHATE 307

Notes:

1. If the analysis for the acidic constituents has not been made and a
quick qualitative test for phosphate is desired, it can be made on a portion
of the original material as follows:

Take approximately 0.1 g of the original sample in a casserole
(if arsenic has been found present, add 5 ml of 12 n. HC1, 2 ml of
9 n. HBr, and 2 drops of liquid bromine and evaporate the mixture
just to dryness under a hood), add to it 5 ml of 16 n. HNOa, and
evaporate the mixture almost to dryness. Add to the residue 5
ml of 6 n. HN0 3 and filter out the residue on a small paper filter.
Collect the filtrate in a small flask in which has been mixed 5 ml of
HNOa and 5 ml of (NEUhMoCU reagent. Warm the mixture to
40 to 60C. and let it stand for 10 min. (Canary yellow precipi-
tate, presence of phosphate.)

This method for the precipitation of phosphate which, because it can
be made from an acid solution, is extensively used for the separation of
phosphate from the basic constituents with which it is associated depends
upon the formation of the tri-ammonium salt of a complex molybdophos-
phoric acid. The reaction can be represented as follows:

3NH 4 + + 12 MoOr 4- H 3 PO 4 + 21H+ = (NH 4 )sPO 4 -12Mo0 8 (,) + 12H 2 0.

As is indicated by the equation, the sensitivity of this test is dependent
upon (1) a large excess of molybdic acid, required to insure the formation
of the complex molybdophosphoric acid; (2) a high concentration of am-
monium ion, to reduce the solubility of the tri-ammonium salt of this acid;
and (3) an acid solution, to convert the rnolybdate ion to the acidic form.

If the solution is made too acid or is heated much above 70C., white or
pale yellow molybdic acid, MoCVH^O, may precipitate. In this case the
precipitate should be dissolved in NHUOH and the solution should be reacidi-
fied with HNO 3 . The yellow precipitate is of too uncertain composition
and is too unstable for it to be used as the basis of a precise gravimetric or
volumetric determination. Accordingly, when a gravimetric determination
is desired, the yellow precipitate is dissolved in ammonium hydroxide
and the phosphate is reprecipitated as magnesium ammonium phosphate
(MgN^PCVG^O) from the ammoniacal solution.

If arsenic is present, the sample is evaporated with HC1 and HBr in
order to volatilize it as arsenic trichloride. Bromine is added to insure
the solution of arsenic sulfides, and bromide to insure that all the arsenic is
reduced to the tripositive state. The removal of arsenic is necessary be-
cause arsenic acid, as would be predicted from the periodic relationships,
forms a yellow precipitate, (NH4) 3 AsO4*12Mo0 3 , analogous in its properties
to the phosphorous compound.

2. It is possible that the precipitate obtained on adding ammonia to
the H2S filtrate (in P. 51) may have been due to iron phosphate
alone, and in that case the removal of phosphate is not necessary.
Also, if alkaline earth elements have not been brought into this group and
if manganese is not present in considerable amount, the presence of phos-
phate is permissible. Therefore, one may add a slight excess of NHiOH
to the HC1 solution above, and, if no precipitate is obtained, no alkaline

308 AMMONIUM SULFIDE GROUP [P. 54

earth phosphates are present; also, if no brownish precipitate results upon
adding to this slightly ammoniacal solution a few drops of 3 per cent H202,
the absence of manganese is shown and the removal of phosphate is unneces-
sary. If a brownish precipitate is obtained, the mixture should be made
acid with HC1, boiled until the precipitate dissolves, and treated as directed.

3. This volume of HC1 solution is colored a distinct blue by 0.5 mg of
cobalt and green by the same quantity of chromium, while 1 mg of nickel
imparts a pale yellowish-green color. However, it should be remembered
that the presence of as little as 0.1 mg of ferric iron would give a distinct
yellow color to this small volume of HC1, and this may vitiate the color
detection of traces of cobalt, chromium, and nickel.

4. The solution is evaporated repeatedly with 16 n. HNOs to remove
chloride ion, which interferes with the precipitation of BiP(>4 because of
the tendency of bismuth to form such complex compounds as HBiCU. The
formation of these complexes reduces the concentration of bismuth ion and
requires the addition of a large excess of bismuth nitrate reagent to precipi-
tate completely the phosphate present.

5. This neutralization is conveniently followed by the effervescence
produced upon the addition of Na2COa. The neutralization of the acid is
complete when the addition of 0.5 ml of the Na2COa solution gives no evo-
lution of CO2. Avoid an excess of Na2C0 3 , because the acidity of the
solution is to be adjusted for the removal of phosphate.

6. A white precipitate is produced upon neutralizing with Na2COs an
acid solution containing phosphate and aluminum, manganese, zinc, or
alkaline earth elements. Under the same conditions cobalt gives a purplish-
blue, chromium a bluish-green, and nickel a light green precipitate.

7. The reagent is added dropwise because a more easily filtered precipitate
is obtained if the precipitation takes place slowly. The precipitate should
be allowed to settle before each addition of the bismuth nitrate reagent.
This enables the completeness of the precipitation to be followed closely
and prevents the use of too large an excess of the reagent.

The bismuth nitrate reagent is 0.1 f. in bismuth nitrate and 0.5 f. in
HNOa. The reagent is made acid with HN0 3 to prevent the precipitation
of oxy-salts of bismuth. About 65 ml of the reagent are required to precipi-
tate 500 mg of phosphate.

8. If it is desired to estimate the amount of phosphate present, a filter
of the Gooch type should be used. In this case, after the precipitate has
been washed thoroughly, it should be heated to constant weight at 300 to
500C. The crucible should have been heated to constant weight at this
temperature before being used.

9. If the BiP0 4 precipitate is greenish, it indicates that chromium has
been coprecipitated (see the discussion). If desired, this chromium can
be determined as follows :

Transfer as much as possible of the precipitate from the filter to
a 300-ml flask. Dissolve the precipitate remaining on the filter
with 1 to 5 ml of HC1 and wash the filter with two 5-ml portions
of 2 n. HC1, collecting the solution and washings in the flask.
Shake the flask until the precipitate dissolves and dilute the solu-
tion until the acid concentration is about 0.3 n. Saturate the
solution with BUS and filter it, collecting the filtrate in a 300-ml

P. 66] ZINC AND ALUMINUM GROUPS 309

flask. Wash the precipitate with three 5-ml portions of hot water
and combine the washings with the solution. Add 6 n. NaOH to
the solution until it is definitely alkaline, cool it, and treat the cold
mixture with 1 to 2 g of Na2Oz. Boil the solution until the excess
peroxide has decomposed and the volume of the solution has been
reduced to about 80 ml. Cool the flask and treat the solution by the
last paragraph of P. 75.

10. The solution, which may contain alkaline earth elements, is treated
by P. 51 in order to again precipitate any Ammonium Sulfide Group ele-
ments. This precipitate should then be dissolved in HC1 as directed in
the first paragraph of P. 52, and the solution should be treated by P. 55.
The filtrate, containing any alkaline earth elements, can be combined with
the original filtrate from the ammonium sulfide precipitation if the alkali
elements are not to be tested for or can be treated separately by P. 81.

P. 55. Separation of the Zinc and Aluminum Groups

Discussion of general methods for separating the Ammonium
Sulfide Group elements into two groups. After the iron has been
removed from the group of elements obtained by the ammonium
sulfide precipitation, it would be advantageous to separate the
remaining elements into two smaller groups. An obvious method
would be the precipitation of the tripositive hydroxides of aluminum
and chromium by suitably adjusting the hydrogen ion concentration
of the solution. The unsatisfactory results obtained when this is
done by the "ammonia precipitation" have been shown in the dis-
cussion of P. 51. Similar conditions of acidity are obtained by
treating an acid solution of these elements with an excess of solid
barium carbonate; however, this introduces barium into both the
precipitate and the solution, and the presence of the excess of barium
carbonate makes it difficult to detect small precipitates of aluminum
or chromium hydroxides. Various other methods of obtaining
this desired hydrogen ion concentration are used in quantitative
separations, but these are in general not suitable for use in qualita-
tive systems, owing to the wide variations in the quantities and
combinations of elements which may be encountered. 28 Other
methods of separating these elements into two groups are discussed
below.

The Separation by Sodium Hydroxide and Peroxide. As is shown
in Table XVIII, p. 282, if a high concentration of hydroxyl ion is used,
the amphoteric nature of aluminum, chromium, and zinc is taken

86 Descriptions and discussion of these methods will be found in Hillebrand
and Lundell, Applied Inorganic Analysis, pp. 68-701, and in Treadwell-Hall,
Analytical Chemistry, Vol. II, Quantitative, 8th Ed., pp. 155-162.

310 AMMONIUM SULFIDE GROUP [P. 55

advantage of and these elements are dissolved (chromium hydroxide
dissolves in dilute hydroxide solutions largely because of its tendency
to form a stable colloid ; in concentrated hydroxide solutions chromite
is formed), while manganese, nickel, and cobalt are precipitated as
hydroxides or hydrated oxides. Owing to the tendency of chromium
to be carried down with the precipitate and to the appreciable solu-
bility of cobalt in high concentrations of alkali, this method is nearly
always modified by adding an oxidizing agent (usually sodium per-
oxide), which oxidizes the chromium to chromate, the manganese
to manganese dioxide, and the cobalt to cobaltic oxide. The process
is used extensively in qualitative systems, and in quantitative
methods for the separation of the individual elements. 27

Statements in the literature as to the effectiveness of these indi-
vidual separations are somewhat incomplete and often not definite.
Noyes, Bray, and Spear 28 state: "The separation of the two groups
by this process is entirely satisfactory, at any rate, from the stand-
point of qualitative analysis, with the single exception that when
5 or 10 mg. of zinc are present this may be carried down completely
when elements of the Iron Group (especially manganese) are present
in large quantity/ 7 However, there appears to be a certain amount
of distrust of the method; this is perhaps explained by the statement
of Hillebrand and Lundell: 29 "The use of sodium hydroxide for pre-
cipitations has not been in good repute among analysts because of
the uncertain quality of the reagent and the slimy character of the
solution and precipitate." It is explained, however, that the first
objection is not as well founded today as in the past, that the diffi-
culty in filtering about 5 per cent solutions is not serious, and that
"in the hands of one of us (L) the method has proved very satis-
factory. "

The confirmatory experiments and test analyses of Noyes, Bray,
and Spear give quite complete information as to the qualitative
value of the separation, as they perform it, in the detection of small
amounts of an element in the presence of even large amounts of
another; additional data showing the completeness of the individual
separations, when considerable amounts of each element are present,
have been collected 30 and are presented in Table XXI.

27 Frequently in quantitative methods other oxidizing agents, such aa
bromine, chlorine, or peroxysulfate, are substituted for the peroxide, but these
are not as convenient as the peroxide, and, in qualitative systems, may intro-
duce objectionable ions into the analysis.

28 Noyes, Bray, and Spear, J. Am. Chem. Soc., 30, 484 (1908).

29 Hillebrand and Lundell, Applied Inorganic Analysis, Wiley, 1929, p. 76.

30 Swift and Barton, /. Am. Chem. Soc., 64, 4155 (1932).

P. 55] ZINC AND ALUMINUM GROUPS 311

In Procedure II of this table (see p. 312) the slightly acid solution con-
taining the hydrogen peroxide was poured into the hot hydroxide
solution, as there is agreement that a better separation results from
this order of mixing. Qualitative experiments also indicated that
a more rapid and complete oxidation of chromium was obtained
when hydrogen peroxide was present in the acid solution; when
sodium peroxide was added to an alkaline solution, as in Procedure
I, or an acid solution was added without peroxide to a hot sodium
hydroxide solution, a precipitate which oxidized slowly was obtained.
The acid solution could be added to a cold sodium peroxide solution ;
however, the preparation of a solution of sodium peroxide, even when
it is kept cold, results in excessive loss of peroxide by decomposition.
Furthermore, the precipitate produced upon addition of the acid
solution to the hot hydroxide seemed more readily handled than that
formed by precipitating in the cold and then heating.

The data in the table show that the separation of aluminum from
manganese and iron can be made sufficiently complete for qualita-
tive and for most quantitative work; with 250 mg of each element,
from 4 to 5 per cent of the aluminum is left with the cobalt precipi-
tate and from 12 to 14 per cent with the nickel precipitate.

The separation of chromium, beginning with chromic ion and with
sodium peroxide as the oxidizing agent, is unsatisfactory in every
case regardless of method. That this is due to incomplete oxidation
of the chromium by the peroxide was shown by analysis of the pre-
cipitates; 81 only tripositive chromium was found to be present.
This is also shown by Experiments 20 to 23, where, beginning with
the chromium as chromate, satisfactory separations were obtained,
except in the case of manganese, in which the coprecipitation is
reduced from 20 per cent to about 3 per cent.

The separation of zinc from the elements precipitated by sodium
hydroxide is unsatisfactory, from 10 to 50 per cent of the zinc remain-
ing in the precipitate. Because of these difficulties, this method has
not been used in this system of analysis.

Separations Obtained by Controlling the Sulfide Concentration and
also Adding a Complex-Forming Anion. Aluminum (and chromium)
are often quantitatively separated from iron by treating the solu-
tion with tartaric acid or a tartrate -which in alkaline solutions
forms soluble complex ions with both aluminum and chromium
and then adding ammonium sulfide, which precipitates the iron as
sulfide. Experiments have shown that this method provides ex-

11 Experiments by Harrison Backus.

TABLE XXI

THE SEPARATION OF MANGANESE, IRON, COBALT, AND NICKEL FROM

ALUMINUM, CHROMIUM, AND ZINC BY PRECIPITATION WITH

SODIUM HYDROXIDE AND PEROXIDE

In these experiments 250 mg of one of the elements in the first column were
precipitated from a solution containing 250 mg of one of the elements listed
at the top of the three major columns. In each of the three major columns
are shown the experiment number, the procedure used, and the amount
in milligrams of the soluble element found in the precipitate.

Element
Precipi-
tated

Aluminum

Chromium

Zinc

Expt.

Pro-
ce-
dure

Alin
Ppt.
(nig)

Expt.

Pro-
ce-
dure

Cr in
Ppt.

Expt.

Pro-
ce-
dure

Zn in
Ppt.

Manganese

12 to 15

50 to 70

2 to 3

!!*

ll d

Iron

1 to 2

103

(ferric)

1 to 2

II'

0.2

Cobalt

20 to 25

65 to 80

!<*

12 to 15

II*

II-

Nickel

30 to 35

30 to 40

0.5 to 1

The separations were carried out according to two general procedures.
The first of these, called Procedure I, conforms closely to that given by A. A.
Noyes* and by Noyes and Bray,t and more nearly represents the usual qualita-
tive technique. It was as follows :

Procedure I: To a solution of the elements to be separated, con-
taining 1 to 2 ml of 6 n. hydrochloric acid and having a volume of 30
to 40 ml was added 6 n. sodium hydroxide until the solution was alka-
line to litmus, and then 5 g of sodium peroxide. The sodium per-
oxide was sprinkled in very gradually, and the mixture was kept cold
during the addition. Then the mixture was boiled until the sodium
peroxide was decomposed, diluted to 60 ml, and filtered through
hardened filters; in order to facilitate washing, two separate filters

* Noyes, Qualitative Chemical Analysis, Macmillan, 1922, 9th Ed.,
p. 95.

t Noyes and Bray, A System of Qualitative Analysis for the Rare Elements,
Macmillan, 1927, pp. 164, 168.

312

P. 56] ZINC AND ALUMINUM GROUPS 313

were used if the precipitate was bulky. The precipitate was washed
with hot water until the washings no longer turned red litmus blue.
In Procedure II certain modifications were used which are recommended
as giving more quantitative separations. It was as follows:

Procedure II: To a solution of the elements to be separated, con-
taining 1 to 2 ml of 6 n. hydrochloric acid and having a volume of 15
to 20 ml, 6 n. sodium hydroxide was added dropwise until the first
permanent turbidity was produced. To this was added 20 ml of 3 per
cent hydrogen peroxide, and the resulting solution was poured drop-
wise into 100 ml of 5 per cent sodium hydroxide which was kept just
boiling during this addition. The mixture was then cooled and kept
cool during the slow addition of 5 g of sodium peroxide. Finally,
the mixture was boiled until the peroxide was decomposed, and it was
then filtered, two separate papers being used in most cases. Hardened
filters were not necessary in filtering these solutions. The precipi-
tate was washed with 50 ml of hot 5 per cent sodium hydroxide and
then with a hot 1 per cent solution until no test for the soluble element
being separated was obtained. Observations and variations from
the procedures outlined are contained in the notes to the table.

Only 125 mg of aluminum were taken.

b The color of the solution indicated that the chromium was completely
oxidized when the acid solution containing hydrogen peroxide was poured
into the sodium hydroxide.

Difficulty was experienced in washing the precipitate free of chromate.
d Only 160 mg of zinc were taken.

The chromium was taken as potassium chromate; and no hydrogen per-
oxide was added to the acid solution.

f Duplicates of Experiments 32 and 33, except that the sodium hydroxide
was added to the acid solution instead of the reverse order.

ceptionally fine separations of iron and zinc from aluminum and
chromium, and that the separation of cobalt from either aluminum
or chromium was equally satisfactory, although these separations
are surprisingly imperfect by either the ammonium hydroxide or
sodium hydroxide-peroxide methods. 32 However, there was some
uncertainty as to the complete precipitation of manganese and
nickel; and, in addition, the removal of the tartrate from the filtrate
by fuming with sulfuric and nitric acids (and the re-solution of the
chromic sulfate, apparently Cr 4 H 2 (S04)7, which formed) proved so
tedious and time-consuming that the method was considered as
undesirable.

It is known that oxalate tends to form complex compounds with
all of these elements, and many of these compounds have been pre-
pared. Very stable ions of the type M (CaOi)/" are formed with
aluminum and chromium (as well as with ferric iron), and ions,
apparently somewhat less stable, of the type M (0264)2" are formed
with manganese, zinc, nickel, and cobalt. As oxalate can be much

11 Unpublished experiments by F. N. Laird.

314 AMMONIUM SULFIDE GROUP [P. 55

more easily oxidized and removed from a solution, its use as a sub
stitute for tart rate was investigated.

It was found that moderate amounts of aluminum or chromium
could be held in an ammoniacal solution if sufficient oxalate was
present, but that the precipitation of manganese as sulfide from such
solutions was quite incomplete, as much as 50 mg of manganese
being present before a precipitate was obtained, and that in slightly
acid solutions a precipitate of manganese oxalate was obtained.
However, upon neutralization of the solution with sodium hydro-
carbonate, no precipitate of manganous oxalate formed when an
excess of oxalate was present, and it was found possible to have even
large amounts of manganese present without formation of a precipi-
tate upon saturation of the solution with hydrogen sulfide.

A detailed study of the behavior of manganese and of the other
elements of the Ammonium Sulfide Group in solutions containing
sodium hydrocarbonate and an excess of oxalate was therefore made.* 3
These experiments showed that the manganese remains in solution
largely because of the formation of a supersaturated solution, but
that this state of supersaturation can be maintained for a consider-
able length of time even when another precipitate is present. They
also showed that, by properly adjusting the amount of sodium hydro-
carbonate added, the volume of the solution, and the time for w r hich
the hydrogen sulfide is passed, even 500 mg of aluminum, chromium,
and manganese (hereafter designated as the Aluminum Group) can
be held in solution, while 1 mg of zinc, nickel, and cobalt (hereafter
designated as the Zinc Group) are precipitated. Iron, if present, , is
likewise precipitated. Furthermore, it was shown that 1 mg of any
one element of the group can be detected in the presence of 500 mg
of any element of the other group, and that with 250 mg of any
element of the Zinc Group and 250 mg of any element of the Alumi-
num Group not over 1 mg of the Aluminum Group element will be
carried into the precipitate. These results were so satisfactory that
this method has been adopted for use in this procedure when a pre-
cise separation of considerable amounts of these elements is desired.

Optional Method for Separating the Aluminum Group. As the
sulfide-oxalate procedure requires a fairly close adjustment of condi-
tions and is somewhat time-consuming, and as the destruction of
the oxalate in P. 71 is also time-consuming, an operation is provided
whereby an indication of the amount of the Aluminum Group ele-

" Swift, Barton, and Backus, J, Am. Chem. Soc., 54, 4161 (1932).

P. 55] ZINC AND ALUMINUM GROUPS 315

ments present is obtained. For the case that these are present in
small amounts, an optional and much shorter, though less exact,
method is provided in Note 3 of this procedure for their separation
and detection. By this method aluminum and chromium (and
iron, if it is present) are precipitated by ammonium hydroxide and
separated from manganese, zinc, nickel, and cobalt. Manganese is
then separated from these elements by the addition of bromine,
which oxidizes and precipitates it as hydrated manganese dioxide.
The filtrate is then carried to the Zinc Group for analysis. The
ammonia precipitate is subjected to a "sodium hydroxide-peroxide"
treatment, which precipitates any iron, leaving aluminate and
chromate in the solution. The aluminum is precipitated as hy-
droxide by acidifying the solution and then neutralizing with am-
monium hydroxide; chromate remains in the filtrate.

Procedure 56: SEPARATION OF THE ZINC AND ALUMINUM
GROUPS. Transfer the HC1 solution (from P. 52) to a 200-
ml flask and evaporate it to 4 to 5 ml (Note 1). Dilute the
solution to 50 ml, heat it to boiling, and Carefully make it
just alkaline with NH 4 OH, avoiding an excess (Note 2).
(White or greenish precipitate, presence of aluminum or
chromium; reddish precipitate, presence of iron. See Note
3 in regard to optional method of analysis.)

Slowly add HC1 to the hot mixture until any precipitate,
is dissolved (avoiding an excess), dilute the mixture to 100
ml, add 2 to 10 g of solid (NH 4 )2C 2 O4, and heat the mixture
until this dissolves (Notes 4, 5). Cool the solution and
carefully add 1 f. NaHC0 3 , in 1-ml portions, until it is
neutral to litmus that is, until it turns both blue and red
litmus to an intermediate purplish color (Note 6). Satu-
rate the cold solution with H 2 S under pressure, first sweeping
the air from above the solution with the gas, and then
bubble a rapid stream through the solution for 10 to 15 sec.
and a slow stream (2 or 3 bubbles a second) for 5 min. (Note
7) ; avoid exposing the solution to the air insofar as is pos-
sible (Note 8). Test the solution (which should be slightly
alkaline) with litmus, and, if it is neutral or at all acid, add
10 ml of the NaHCOs (taking care to avoid loss due to
effervescence), and again saturate it; repeat this process
until the solution remains slightly alkaline (Note 9). Fi-
nally, while continuing the slow stream of HaS, heat the

316 AMMONIUM SULFIDE GROUP [P. 56

mixture just to boiling and keep it at that temperature
until the precipitate settles rapidly. (White precipitate,
presence of zinc; black precipitate, presence of cobalt or
nickel. Note 10.) If the precipitate does not settle
rapidly, tightly close the flask with a clean rubber stopper,
cool the mixture to 60 to 70C., and, holding the stopper
firmly in place, vigorously shake the mixture for a few min-
utes. Immediately filter the hot mixture through a paper
filter, decanting as much as possible, and wash the precipi-
tate with a hot wash solution which has been prepared by
adding 1 ml of NaHCO 3 and 1 g of (NH 4 )2C 2 O4 to 100 ml of
hot water and passing a stream of H 2 S through it for 30 sec.
Add the first 10 to 15 ml of the wash solution to the filtrate
and treat it by P. 71 (Note 11). Treat the precipitate by
P. 61.

Notes:

1. If iron has been present, this solution will contain a considerable
amount of ether dissolved in it. In this case the first part of the evapora-
tion, which should be carried out under a hood, can be effectively done by
heating a beaker of water to boiling, extinguishing the flame, and immersing
the flask in the water. After the ether is expelled, the flask may be heated
directly with a flame.

The solution is evaporated to a small volume in order that less NF^OH
be required for the neutralization. A high concentration of ammonium
ion might later cause ammonium oxalate to crystallize from the solution.
See Note 3, P. 54, as to the information concerning the elements present
which may be obtained from the color of the HC1 solution.

2. The neutralization should be carefully made and an excess of am-
monia should be avoided; if an excess is added, the solution should be acidified
and the neutralization should be repeated. If the solution is colorless,
methyl red may be used as an indicator and ammonia may be added until
the transition from pink to yellow is obtained; if the solution is colored,
litmus paper can be used and an intermediate purplish color can be taken
as the neutralization point, or, if they are available, use one of the indicator
test papers mentioned in P. 3, Note 6, and adjust the solution to a pll of (>.
The first perceptible smell of ammonia from the boiling solution is also a
sensitive indication of when the proper amount has been added; the NH 4 OH
should be added with a dropper and care should be taken not to spill it on the
sides of the flask.

If the solution is made too alkaline, manganese may be oxidized and may
appear as a brown precipitate (hydrous M^Oa). If present in large amounts
(300 mg or more), zinc may precipitate as the white hydroxide; therefore,
if a white precipitate is obtained (and zinc may be present in large amounts),
add HC1 until the precipitate is dissolved and then 1 ml in excess. Repeat
the neutralization with ammonia, taking extreme care to avoid an excess.

P. 55] ZINC AND ALUMINUM GROUPS 317

By providing a higher concentration of ammonium salts and of chloride,
the precipitation of zinc is prevented.

3. Under these conditions less than a milligram of aluminum or chro-
mium will cause a precipitate; therefore, if no precipitate is obtained, the
absence of these elements is proved. If no precipitate is obtained here and
no color was present in the HC1 solution when it was evaporated to a small
volume, the absence of cobalt and nickel is also proved; and only mag-
ganese and zinc can be present. In this case treat the solution at once by
P. 61 (second paragraph), beginning with the addition of 5 ml of 36 n. HjSC>4.
By this procedure (P. 61) the zinc is precipitated as sulfide, the manganese
remaining in solution; by then expelling the E^S and making the solution
alkaline with sodium peroxide, the manganese will be precipitated as hydrous
manganese dioxide. Each of these precipitates can then be treated directly
by the procedures provided for their estimation.

If only a small precipitate of aluminum, chromium, or iron is obtained
(one corresponding to 25 to 50 mg, depending upon the exactness which is
desired in the separation and estimation of these elements), it may be filtered
out and analyzed as follows:

While keeping the solution hot, filter out the precipitate on a
small paper filter. Wash the precipitate with hot water, collecting
only the first 5 ml of wash water with the filtrate. Stopper and
reserve the filtrate.

Dissolve the precipitate by pouring 5 to 10 ml of hot HC1 drop-
wise through the filter; wash the filter with 2 to 3 ml of water.
Evaporate the solution to 3 to 5 ml, cool, and dilute to 10 ml.
Make the solution alkaline by sprinkling in small portions of Na202,
add about 0.5 ml of the solid in excess, and then boil until small
bubbles are no longer evolved. (Brownish-red precipitate, presence
of iron; yellow solution, presence of chromium.) Filter out the
precipitate and treat it by Note 6, P. 63.

(The solution should be cold before adding the Na2C>2, as otherwise this
substance may decompose explosively. A porcelain spatula or spoon
should be used not paper. The solution is boiled until no more oxygen
bubbles are evolved, as the excess peroxide must be decomposed or it would
reduce any chromate present when the solution is subsequently acidified.)

Make the filtrate from the Na2C>2 treatment just acid with HC1,
then just alkaline with NH 4 OH, avoiding an excess, and heat it to
boiling. (White precipitate, presence of aluminum.) Estimate the
amount by comparison with known amounts of aluminum precipi-
tated under similar conditions. Filter out the precipitate and
treat it as directed in Notes 1 and 2, P. 73.

If chromium is present, estimate the amount by adding a stand-
ard chromate solution to the same volume of a solution made just
alkaline with NH40H. To confirm the presence of chromium,
make the solution just acid with HC2H302 and add 1 ml of 1 n.
Pb(N0 3 )a. (Yellow precipitate, presence of chromium.) Discard
the mixture. ,

As has been pointed out, the "ammonia precipitation" and the "sodium
hydroxide-peroxide" separation cannot be used when large amounts of the

318 AMMONIUM SULFIDE GROUP [P. 55

above elements are to be separated or when quantitative separations are
desired; however, their use as indicated above effects a considerable saving
of time for separating and detecting smaller amounts. This is especially
advantageous in detecting and removing small amounts of aluminum, as
this element is often introduced into the analysis in small amounts as im-
purities in the reagents, especially by alkaline solutions which have stood in
glass containers. In such cases it is an advantage to remove it as expedi-
tiously as possible.

Aluminum and chromium (and iron) having been removed from the
solution, the reserved filtrate can now also be tested for manganese, which,
if present in small amounts, can be separated at this point. This is desir-
able, as the necessity for carrying out P. 55 and P. 71, both long and trouble-
some procedures, is thus eliminated. Proceed as follows:

Add to the filtrate (from the ammonia precipitation) 2 ml of
NEUOH and 1 ml of bromine water. (Brownish precipitate, pres-
ence of manganese,) If a large precipitate is produced and a pre-
cise separation is desired, make the mixture acid with HC1 and,
without dissolving the precipitate (Note 4), treat it by the second
paragraph of P. 55 above. If only a small precipitate is produced,
add 2 ml more of bromine water and of NH 4 OH, warm the mixture,
filter out the precipitate, and treat it as directed in Note 1, P. 72.
Treat the filtrate by P. 61, beginning with the addition of 5 ml of
36 n. H 2 S0 4 .

By the addition of Br2 to an ammoniacal solution, manganese is precipi-
tated as hydra ted manganese dioxide. Zinc, nickel, and cobalt remain
in solution as complex ammino ions. Complete precipitation of the man-
ganese requires that the solution be alkaline whan filtered and that an
excess of bromine be added. As bromine and hydroxyl ion are used by the
precipitation reaction

+ Br 2 + 40H~ - MnO(OH) 2 + 2Br~ + H 2 0,
and as bromine and ammonia react and are removed by the reaction
3Br 2 + 2NH 3 + 60H~ - 6Br~ + N 2 + 6H0,

more of each is added in case manganese is found present.

4. If the solution has been made too alkaline or has been allowed to
stand exposed to the air for a considerable time, the manganese may be
partially oxidized and a brownish precipitate may appear. This precipitate
may not rapidly dissolve in the dilute HC1; and, because the addition of a
large excess of acid is undesirable, it is left in the solution, as it will be re-
duced and dissolved by the oxalate next added.

5. Ammonium oxalate dissolves slowly in a cold solution; for this reason
the solution is heated until it is dissolved. If too much acid is present, and
if large quantities of cobalt, nickel, or zinc are present, the oxalates of these
elements may precipitate; it is not necessary to dissolve them, as they will
be completely metathesized by the subsequent H*S treatment.

6. If methyl red has been used in the first neutralization, it may also
be used here, a distinct yellow color being obtained.

P. 56] ZINC AND ALUMINUM GROUPS 319

7. This flow is preferably regulated by passing the gas into the solution
at the pressure of the generator or tank, and then connecting a short piece
of rubber tubing to the outlet tube with a screw clamp. By adjusting the
screw the flow of gas is evenly controlled. This arrangement causes the
H 2 S to bubble through the solution at the generator pressure, thus expedit-
ing the complete precipitation of nickel; unless the directions of this pro-
cedure are followed exactly, large amounts of nickel may not be completely
precipitated.

8. Repeated experiments have shown that, when nickel is present,
exposure to the air causes the precipitate to become colloidal, apparently
because of the formation of disulfide; for this reason the air is swept from
the flask. 34 When the neutrality of the solution is subsequently being
tested, the flow of gas should be continued, the stopper being raised just
sufficiently to permit removing a drop of the solution on a stirring rod.

9. In order to secure rapid and complete precipitation of nickel (or of
iron, if it is present), the solution should be distinctly alkaline to litmus
after precipitation is complete. For this reason it is advisable to test the
solution again with litmus just before stopping the flow of H2S and begin-
ning the filtration.

10. If a white precipitate is obtained, the presence of zinc alone is indi-
cated; and, as even 1 mg of nickel or cobalt will cause a large precipitate of
zinc to be discolored, these elements are absent. In this case the precipitate
should be treated directly by P. 62. A black precipitate indicates the
presence of cobalt or nickel, but a smaller amount of zinc may also be present.

11. If the filtrate is dark colored, this is probably due to colloidal nickel
sulfide, which often partly coagulates after passing through the filter. In
this case test the neutrality of the solution, making sure that it is slightly
alkaline, again saturate it with H2&, slowly add to it 5 ml of 6 n. HNOa
(which should make the solution distinctly acid), heat it almost to boiling
for 2 to 3 min., cool it to 60 to 70C., and again vigorously shake it. Filter
out the precipitate through a separate smaller paper, wash it, and combirie
it with the first one. Treat the filtrate by P. 71. Although it causes an
intensely dark solution and is voluminous in appearance, this second pre-
cipitate usually contains not over 1 to 2 mg of nickel.

34 It should be mentioned that a black colloidal solution has been obtained
by various investigators in ammonium sulfide solutions supposedly free of
disulfide; for a discussion of the phenomenon, see Weiser, The Colloidal Salts,
McGraw-Hill, 1928, p. 103.

TABULAR OUTLINE VII
THE ANALYSIS OF THE ZINC GROUP

H,S precipitate from P. 55: ZnS, NiS, CoS, [FeS] a
Dissolve the precipitate in HCl and HNOt.
Add 6 ml of S6 n. H^SO*. Fume. Cool. Dilute.
Neutralize with NaOH. Add S ml of 6 n. H t SO*.
Treat with H t S. (P. 61).

Precipitate: ZnS
Add to excess of

Fe 2 (SO 4 ),, heat,

add

Precipitate: S
Solution:

Zn++, Fe++, ex-

cess Fe"^ 1 " 1 "

Titrate with

KMnOi.

(Fe+++, Mn++)
(P. 62)

Filtrate: Ni++, Co++, [Fe++], HSOr, SOr

Boil out H^S.

Add excess NaOH.
Precipitate: Ni(OH) 2 , Co(OH) 2 , [Fe(OH) 2 ]

[If iron is present:

Dissolve with HCl, add Br*, boil out excess Br^.
Solution: Ni++, Co++, Fe+++

Neutralize, add HC^H^O* and NaC 2 Hi0 2 . Boil.
Precipitate: Fe(OH) 8 , FeOC 2 H,O 2
Filter and treat by P. 5$.]

Add excess Na z 2 . (P. 63)

Precipitate: Ni(OH) 2 , Co 2 O 3
Dissolve in HCl.
Evaporate. Add ethyl ether.
Saturate with HCl gas. (P. 64)

Filtrate:
Discard.

Precipitate: NiCl 2
Dissolve in H^O.
Add NH 4 Cl, NHtOH.
Add AgNOi, KI.
(Agl; Ni(NH,) 4 ++)
Titrate with KCN.
(Ni(CN)r, Ag(CN)r)
(P. 65)

Filtrate: H 2 CoCl 4
Neutralize , add NaBO 3 .
Precipitate: Co 2 O 8
Boil to decompose NaBOi.
Add HCl, KI.
(Co++, Ir)
Titrate with Na 2 S 2 O^
(Co^, I-, S 4 0r)
(P. 66)

a If iron has not been separated by the ether treatment (P. 52), it will be
precipitated with this group. Its subsequent behavior and separation is
indicated in brackets.

320

The Analysis of the Zinc Group

P. 61. The Precipitation and Separation of Zinc as Sulfide

Discussion of Methods for Separating the Elements of the Zinc
Group. After the sulfides of zinc, cobalt, and nickel have been
obtained, several methods might be used for the separation of these
elements. A method commonly used in qualitative systems con-
sists of treating the mixed sulfides with cold 1 n. hydrochloric acid
in an effort to dissolve the zinc sulfide and leave undissolved the
nickel and cobalt sulfides. This method appears inconsistent with
the fact that nickel and cobalt do not precipitate with the Hydrogen
Sulfide Group elements and with the fact that zinc can be precipi-
tated as sulfide from a solution with a higher hydrogen ion concen-
tration than can either nickel or cobalt. The phenomenon is at-
tributed to the existence of two or more allotropic forms of nickel
and cobalt sulfides of differing solubilities; only the more soluble
modification can be formed in solution, but, once formed, it is
changed into a less soluble form or forms (which may be polymers
of the soluble form). This transformation seems to be favored by
heat and by acid; in general, the longer the sulfides have stood, the
less readily soluble they become. It is also possible that the phe-
nomenon may be at least partly explained by a slow rate of solution ;
many compounds which dissolve readily when recently precipitated
dissolve only slowly after they have been allowed to stand for a
considerable time. This method was not used here because it has
been found that the zinc sulfide also dissolves slowly when it is
present with large amounts of nickel or cobalt, and that appreciable
amounts of nickel and cobalt may be dissolved if the treatment is
prolonged or if considerable amounts of other sulfides are present.

It would be expected that nickel and cobalt could be separated
from zinc by precipitating them from a strongly alkaline solution,
the zinc remaining in solution as the zincate ion; however, as is
shown in Table XXI, large amounts of zinc are carried down with
the nickel and cobalt and small amounts were found to be so com-
pletely coprecipitated that zinc could not be detected in the filtrate.
This method was tried under many different conditions, such as
varying the order of mixing the reagents, using powerful oxidizing
agents (for example, peroxysulfate, which would oxidize both cobalt
and nickel), and adding an excess of sodium hydroxide to an am-

321

322 ANALYSIS OF THE ZINC GROUP [P. 61

inoniacal solution and then boiling off the ammonia; but in no case
were the results satisfactory.

As is shown in Table XI, zinc is precipitated as sulfide from a
solution in which the hydrogen ion concentration is approximately
10~~ 2 , while nickel and cobalt are not. The precipitation of zinc
sulfide has been the subject of many investigations, and it has been
realized for over half a century that there are advantages in precipi-
tating zinc sulfide from an acid solution rather than from one which
is neutral or alkaline. Specifically, certain separations can be made
by proper adjustment of the hydrogen ion concentration, and the
precipitate usually separates in a form which is more readily handled.
Because of these effects, many early workers 1 recommend the use
of various organic acids for acidifying the solution and thereby ob-
tained, somewhat uncertainly, the effect of a buffered solution prior
to the development of the modern theory and use of such solutions.

Somewhat more recently the use of sulfuric acid solutions was
studied by Weiss, 2 who recommended that the precipitation be made
from a solution 0.01 n. in sulfuric acid by means of a very rapid
stream of hydrogen sulfide, but apparently provided no means for
controlling the increase in hydrogen ion concentration taking place
during the precipitation; by Waring, 3 who once more recommended
the use of formic acid; by Funk, 4 who studied the effect of various
organic acids on the precipitation and on the separation of zinc from
various other elements; and by Glixelli, 5 who made a comprehensive
study of the effect of the nature of the solution on the forms and
apparent solubility of the precipitate, and also on the time factor
which enters into the establishment of the precipitation equilibrium.

In the excellent study of Fales and Ware, 6 the limits of the hydro-
gen ion concentration between which quantitative precipitation
could be effectively made were first precisely determined, and means
were provided for controlling the acidity within these limits. They
used formic acid-formate buffered solutions and found that for
such solutions the hydrogen ion concentration most favorable for

1 Deeffs, Chem. News, 41, 279 (1880); Bragard, Dissertation, Berlin, 1887;
Z. anal. Chem., 27, 209 (1888); Beilstein, Ber., 11, 1715 (1885); Miihlhauser,
Z. angew. Chem., 15, 731 (1902); Berg, Z. anal. Chem., 25, 512 (1886); Dohler,
Chem. Ztg., 23, 399 (1899); Neumann, Z. anal. Chem., 28, 57 (1889).

2 Weiss, Dissertation, Munchen, 1906.

1 Waring, J. Am. Chem. Soc., 26, 26 (1904).

4 Funk, Z. anal. Chem., 46, 93 (1907).

Glixelli, Z. anorg. Chem., 55, 306 (1907).

6 Fales and Ware, J. Am. Chem. Soc., 41, 487 (1919).

P. 61] PRECIPITATION OF ZINC SULFIDE 323

quantitative precipitation of zinc sulfide was between 1(T 2 to 10~ 8 .
Lundell, Hoffman, and Bright 7 recommended the precipitation of
zinc sulfide from a 0.01 n. sulfuric acid solution and stated that the
best precipitation is obtained from a sulfuric acid-sulfate solution.
They did not present any discussion or experimental data in support
of this recommendation. As there are certain advantages in using
sulfate-hydrosulfate buffered solutions for this precipitation, a study
of this method was made, 8 and a discussion of the results obtained
is given below.

As the value of the second ionization constant of sulfuric acid has
been determined 8 to be 1.15 X 10~ 2 , and as the optimum hydrogen
ion concentration for the precipitation of zinc sulfide from sulfate
solutions has been found to be about 10~ 2 , it would seem that the
required ratio of sulfate to hydrosulfate should be about one the
most favorable ratio for effective buffer action. As compared with
this, Fales and Ware had to use an acid-to-salt ratio of 84 to 1, with
the formic acid concentration 4.7 f., in order to obtain an initial pH
of 1.86, a value at which, in sulfate solutions, quantitative precipi-
tation was obtained. This pH was experimentally obtained in the
sulfate-hydrosulfate solution by using an acid-to-salt ratio of 20.7
to 66, with the total concentration of sulfate and hydrosulfate 0.347 f .
In addition, Fales and Ware recommended the use of considerable
ammonium citrate and 6.25 g of ammonium sulfate in 200 ml of
solution. They state that "ammonium citrate is used for the pur-
pose of forming complexes with interfering metals/' thus assisting
in holding in solution such elements as iron and manganese. It was
found that from a sulfate-hydrosulfate solution of the proper pH
such an agent was not needed. The ammonium sulfate was added
as a "salting-out agent/' In the sulfate solutions the need of adding
an additional salting-out agent was not apparent; quantitative pre-
cipitation was obtained, and the form of the precipitate left nothing
to be desired as to ease of filtration and washing; furthermore, it
seemed that the hydrogen ion concentration was the predominant
factor governing the form as well as the solubility of the precipitate.
It should also be mentioned that, when the precipitation of zinc

7 Lundell, Hoffman, and Bright, Chemical Analysis of Iron and Steel, Wiley,
1931, p. 388.

Jeffreys and Swift, J. Am. Chem. Soc., 54, 3220 (1932).

Sherrill and Noyes, /. Am. Chem. Soc., 48, 1873 (1926); as the result of
more recent measurements, Hamer, J. Am. Chem. Soc., 68, 860 (1934), gives
the value 1. 2 X 10-*at25C.

324

ANALYSIS OF THE ZINC GROUP

[P. 61

sulfide is used as a means of effecting a qualitative separation, it is
a decided advantage not to have to add organic matter to the solu-
tion, for, especially in a group analysis, material such as citrates
may cause difficulties in subsequent operations. The solubility of
zinc sulfide in solutions with varying ratios of hydrosulfate to sulfate,
and therefore of hydrogen ion concentration, is shown in Table XXII.

TABLE XXII

THE SOLUBILITY OP ZINC SULFIDE IN SULFATE-HYDROSULFATE SOLUTIONS

OF VARIOUS HYDROGEN ION CONCENTRATIONS

(Volume, 250 ml; zinc taken, 257 mg.)

Expt.

Initial Ratio
NaHSCV
Na,S0 4
(millimoles)

Initial pH

Final pH

Zinc Found in
Filtrate
(mg)

8.3/66

2.18

1.83

0.0

12.4/66

2.08

1.78

Trace.

20.7/66

1.86

1.66

0.20

22.8/66

1.78

1.62

Trace'

24.9/66

1.77

1.62

0.1 to 0.2

27.0/66

1.72

1.57

0.3 to 0.5

27.0/66

1.72

1.56

0.3 to 0.5

27.0/66

1.72

1.58

0.3 to 0.5

27.0/66

1.72

1.58

0.3 to 0.5*

31.0/66

1.67

1.47

0.3 to 0.5*

37.3/66

1.57

1.46

0.5

45.0/66

1.49

1.39

0.8 to 1.0

58.0/66

1.38

1.29

1.5

70.5/66

1.31

1.22

4.0

87.2/66

1.19

1.14

6.0

a Filtered after standing for 3 hr.
* Filtered after standing overnight.

From the above data it may be concluded that from these solu-
tions precipitation will be complete to less than 0.25 mg at a pH of
1.6 or at a hydrogen ion concentration as high as 2.5 X 10~ 2 . Less
than 1 mg remains in solution at a pH of 1.54 or a hydrogen ion con-
centration of 2.9 X 10~ 2 . It would seem that the conditions of
Experiment 4, in which the ratio of NaHS0 4 to Na 2 S0 4 is about 1 to
3 and the initial pH is 1.78, are best adapted for securing the desired
completeness of precipitation of zinc and for holding other elements
in solution. If desired, the buffer action could be made considerably
more effective and the change in hydrogen ion concentration during
precipitation could be made much smaller by increasing the concen-
trations of sulfate and hydrosulfate. It was considered desirable, in

P. 61)

PRECIPITATION OF ZINC SULFIDE

325

order to minimize the difficulty in washing the precipitate and avoid
the use of excessive quantities of buffer material, to keep the buffer
concentrations at the lowest values which would adequately control
the acidity. It should be pointed out that, because of the high con-
centration of electrolyte in these solutions, the evaluation of the
activities of the various ions is subject to considerable uncertainty,
and a simple calculation of the hydrogen ion concentration from the
formal salt-acid ratios may lead to results considerably divergent
from the experimental values. This is evident from a series of
experiments made to show the effect of dilution on the pH of such a
solution. The results are given in Table XXIII.

The decrease in acidity

TABLE XXIII

THE EFFECT OF DILUTION ON THE pH OP
SULFATE-HYDROSULFATE SOLUTIONS

Expt.

NaHS0 4

(moles/
liter)

Na 2 SO 4
(moles/
liter)

0.1

0.5

1.99

0.05

0.25

1.48

0.025

0.125

2.17

0.01

OD5

2.43

0.0025

0.0125

2.82

in the more dilute solutions
is due to the fact that there
is not adequate buffering
material to keep the ratio
of hydrosulfate to sulfate
constant, the fraction of the
hydrosulfate ionized into
hydrogen and sulfate ions
becoming large at these
concentrations; the increase
in the more concentrated

solutions is due to the effect of the high salt concentrations on
the activities of the various ions. It is therefore seen that, if, in
order to control more closely the acidity, a higher concentration of
the buffer is used, the initial ratio of hydrosulfate to sulfate should
be adjusted accordingly. Thus, with the sodium sulfate concen-
tration 0.25 f., a ratio of NaHS0 4 to Na 2 S0 4 of 1 to 3 is required
to produce an initial pH of 1.78 (Table XXII, Experiment 4);
with the sodium sulfate concentration 0.5 f., this ratio has to be
reduced to 1 to 5 (Table XXIII, Experiment 1). As it was desired
to keep the volume small in the procedure below, the sulfate con-
centration is approximately 1 f. and the ratio is 1 to 7.

It was found that the presence of a high concentration of chloride
ion lengthens the induction period in the beginning of precipitation,
apparently slows down the precipitation, and, with the same pH,
increases the solubility of zinc sulfide by three or four times. This
is shown by comparing the results in Table XXIV with those ob-
tained at corresponding pH values in Table XXII.

326

ANALYSIS OF THE ZINC GROUP

[P. 61

TABLE XXIV

THE EFFECT OF CHLORIDE ION ON THE SOLUBILITY OF ZINC SULFIDE
(Potassium chloride added, 132 milli-equivalents; zinc taken, 257 mg;

volume, 250 ml.)

Expt.

Initial Ratio
NaHS0 4 /
Na 2 S0 4
(millimoles)

Initial pH

Final pH

Zinc Found
in Filtrate
(mg)

16
17
18

20.7/66
27.0/66
37.3/66

1.73
1.62
1.43

1.51
1.46
1.37

1.5
2.0
2.0

The Precipitation of Zinc Sulfide in the Presence of Other Elements.
The separation of zinc from cobalt and nickel and also from the
other elements of the Ammonium Sulfide Group was also studied,
and Table XXV shows the results obtained.

TABLE XXV
THE SEPARATION OF ZINC FROM OTHER METALS BY PRECIPITATION

AS SULFIDE

(Volume, 250 ml; initial ratio NaHSO 4 /Na 2 SO 4 in millimoles, 20.7/66, except

in Experiment 20; 257 mg of zinc and 250 mg of other metal taken

tinless otherwise stated under remarks.)

Expt.

Other
Metal

Initial
pH

Final
P H

Other Metal
in ZnS Ppt.
(nig)

Remarks

1.82

1.64

200 mg Co.

1.35

Trace

200 mg Co; initial ratio

NaHSO 4 /Na,SO 4 = 58/66.

1.82

Large quan-

lOOmgZn, 400mgCo.

tity

1.82

No ppt.

No Zn, 600 mg Co; solution

stood 3 days.

1.82

1.66

0.3 to 0.4

1.81

1.65

0.1

1.82

1.65

0.5 to 0.6

<0.25

From the results shown in Table XXV it is seen that the separa-
tions are satisfactory in all cases except with cobalt. Even in a
solution whose initial pH was 1.35, too high in acid for complete
precipitation of zinc, cobalt is still coprecipitated. In Experiment

P. 61] PRECIPITATION OF ZINC SULFIDE 327

21, where 100 mg of zinc were precipitated from a solution contain-
ing 400 mg of cobalt, the zinc sulfide precipitate came down quite
white for several minutes, then darkened rapidly, and finally was
almost black in color.

This effect appears to be similar to the induced precipitation of
zinc sulfide in more concentrated acid solutions, which has been
studied by Kolthoff and Pearson. 10 This effect they designate
"post-precipitation" and attribute to an induced precipitation of the
zinc sulfide from a supersaturated solution, the induction being
caused in their experiments by adsorption of hydrogen sulfide on
the copper sulfide. Under the conditions of certain of their experi-
ments, this adsorption became effective in promoting the precipita-
tion of the zinc sulfide only after quantitative precipitation of the
copper sulfide had taken place. That the effect obtained here is a
case of either promoted precipitation or coprecipitation is shown by
Experiment 22. In this experiment the solution, which was identical
with that in Experiment 21, but contained 500 mg of cobalt alone,
was saturated with hydrogen sulfide in the usual way, except that
the gas was bubbled through for more than an hour; then the flask
was closed and allowed to stand for 3 days. No cobalt sulfide pre-
cipitated. Fales and Ware also report that cobalt cannot be sepa-
rated from zinc by a hydrogen sulfide precipitation from formic
acid-formate solutions. It is of interest to note that when a few
milligrams of cobalt are coprecipitated with zinc sulfide, as in Experi-
ment 19, the resulting precipitate is green. This phenomenon was
noted in a large number of experiments. 11

Experiments have shown that the amount of cobalt which is
carried down with the zinc can be very much reduced if the solution
is not heated to boiling; thus a zinc sulfide precipitate which had
been formed at 60C. in the presence of cobalt became only slightly
greenish in color even after standing 24 hr., but turned quite dark
when the mixture was heated to boiling. It is not practicable to
make the precipitation at room temperature, as under these condi-
tions zinc sulfide forms very slowly and complete precipitation
would require many hours.

10 Kolthoff and Pearson, J. Phys. Chem., 36, 549 (1932).

11 Caldwell and Mayer, /. Am. Chem. Soc., 67, 2375 (1935), have found that
the addition of a small amount of acrolein to the solution greatly reduced
the amount of cobalt precipitated with the zinc sulfide, and, ibid., 57, 2372
(1935), that a very small amount of gelatin or agar-agar causes much more
rapid coagulation of the zinc sulfide precipitate. The original articles should
be consulted for a discussion of the phenomena.

328

ANALYSIS OF THE ZINC GROUP

[P. 61

The hydrogen ion concentrations at which cobalt and nickel can
be quantitatively precipitated are shown in Table XXVI. The
data for this table are taken from studies 12 which were made in solu-
tions in which the hydrogen ion concentration was controlled by
means of acetic acid and ammonium acetate.

TABLE XXVI
THE PRECIPITATION OF COBALT AND NICKEL AS SULPIDES IN SOLUTIONS OP

VARIOUS HYDROGEN ION CONCENTRATIONS

Known amounts of cobalt and nickel were taken. The cobalt in the pre-
cipitate was weighed after reduction to the metal; the nickel sulfide precipi-
tate was converted to the oxide and weighed. More complete and exact
. data are given in the original articles.

(Volume, 100 ml; Ni or Co taken, about 0.2 g.)

_TJ

Error in the Determination

Cobalt

Nickel

(nig)

(mg)

4.9

+0.2

4.7

-0.1

4.4

-0.2

4.2

-0.3

4.0

-0.1

-0.4

3.9

-0.1

-0.8

3.7

-0.2

-0.6

3.4

-1.4

-1.6

3.3

-1.3

-9.9

3.1

-13.5

Procedure 61: PRECIPITATION OF ZINC. If the hydro-
gen sulfide precipitate from P. 55 is white, treat it by
P. 62 (Note 1).

If the precipitate is discolored (Note 2), transfer as much
as possible of it from the filter to the flask. Dissolve that
remaining on the filter by pouring repeatedly through it 5
ml of hot HC1 to which has been added 1 ml of 16 n. HN0 3
(Note 3). Add the solution to the precipitate in the flask,
together with 1 to 5 ml of 12 n. HC1 and 2 to 10 ml of 16 n.
HNOa. Evaporate the mixture slowly until the sulfides
are dissolved (Note 4). Slowly (Caution) add just 5 ml
of 36 n. H 2 S0 4 and heat the solution until it begins to fume
(Note 5). Cool the solution to room temperature (P. 24,

12 Haring and Heatherman, /. Am. Chem. Soc., 62, 5135 (1930); Haring and
Westfall, ibid., 52, 5141 (1930).

P. 61] PRECIPITATION OF ZINC 329

Note 3) and, while continuously cooling the flask with run-
ning water, pour slowly into it, in 1-ml portions, 20 ml of
water, and then dilute the solution to about 70 ml. Add
carefully 6 n. NaOH until the solution is just neutral to
litmus or until a faint permanent turbidity is obtained, and
then add exactly 2 ml of 6 n. H*SO 4 (Notes 6, 7). Heat the
solution to about 60C. and pass in a moderately rapid
stream of H 2 S for about 10 min. (White precipitate, pres-
ence of zinc. Notes 8, 9.)

Filter the mixture, retaining as much as possible of the
precipitate in the flask. Wash the precipitate by decanta-
tion with three 5-ml portions of a hot wash solution which
has been prepared by adding 2.5 g of NajSOi and 15 drops
of H 2 SO 4 to 100 ml of water (Note 10). Collect these wash-
ings with the filtrate (after combining with it any cobalt
or nickel recovered in Note 12). Again wash the precipi-
tate with the wash solution until it is free of sulfide (Note
11). Treat the precipitate by P. 62 (Note 12). Treat the
filtrate by P. 63.

Notes:

1. Since even 1 mg of nickel or cobalt will distinctly discolor the precipi-
tate, it should be treated directly by P. 62 if it is quite white.

2. Small amounts of cobalt or nickel may not make the precipitate dark
or merely gray, as would be expected. Thus a precipitate containing 500
mg of zinc and 5 mg each of cobalt and nickel was greenish-gray in color.

3. Cobalt and nickel sulfides do not dissolve readily in even hot 6 n.
HC1; therefore an oxidizing agent, such as HNOs or KClOa, must be used
to facilitate the solution.

4. A small dark residue consisting of sulfur and some enclosed sulfide
may remain when this mixture is evaporated. When this becomes light in
color, the residue may be removed with a stirring rod and discarded. If the
analyst wishes to be sure that no nickel or cobalt remains in this residue,
it should be placed in a small porcelain crucible and gently heated with a
burner until the sulfur is burned. If a black residue remains, it should be
dissolved in HC1 and added to the main solution.

5. The fuming must be stopped at the first escape of dense white fumes
from the mouth of the flask, in order to prevent excessive loss of sulfuric
acid. The quantity of acid present must be known closely in order that
the solution may be buffered effectively for the separation of zinc.

6. This neutralization of the acid must be done precisely, as the ratio
of HS04~ to S04~, and therefore of the H+ concentration, will depend upon
this and upon the amount of H2S04 added. If a large precipitate is pro-
duced, it should be dissolved in the smallest possible volume of HjS(>4 and
the neutralization should be repeated.

330 ANALYSIS OF THE ZINC GROUP [P. 61

This neutralization can be carried out more exactly and the formation
of a precipitate can be avoided by the use of "Nitrazine" indicator test
papers. The sodium hydroxide should be added until the color given by
the paper corresponds to a pH of 5 to 5.5 on the chart provided with the
indicator papers. (See Note 1, P. 161; see also Note 6, P. 3, in regard to
the use of other indicator test papers.)

7. If iron has not been separated by the ether treatment (or as suggested
in P. 55, Note 3), it will be in the ferric form here, having been oxidized by
the HNOj, and will cause a precipitate upon treatment of the solution with
H 2 S, because of the reaction

+ + H 2 S - 2Fe++ + S ( ., + 2H+.

As this precipitate may be mistaken for ZnS, the iron should be reduced by
passing an excess of S0 2 into the solution. As S0 2 and H 2 S also react,

S0 2 + 2H 2 S . 3S (i) + 2H 2 0,

the solution should be boiled until the S0 2 is completely expelled. The
water lost by evaporation should be replaced.

8. When a small quantity of zinc, 1 mg or less is present, only a faint
turbidity may be observed after the solution has been treated with H 2 S for
10 min. at 60C. In this case, heat the solution slowly to 80C. while pass-
ing in a slow stream of H 2 S, then stopper the flask, and let it stand 5 min.
before filtering. The coagulation of the precipitate is hastened by this
treatment.

If a large precipitate is obtained after 10 min., it is advisable to continue
the passage of H 2 S for 10 min. more at 60C. Do not heat the solution to
boiling, because a large amount of cobalt may precipitate at the boiling
temperature when large quantities of both cobalt and zinc are present.

9. When much cobalt is present with zinc, the ZnS which first forms is
pure white, but on prolonged treatment with H 2 S it may turn a greenish
color, owing to the coprecipitation of cobalt. With 250 mg each of cobalt
and zinc present, 4 to 8 mg of cobalt may coprecipitate.

10. As the precipitation of zinc may be incomplete because of slow pre-
cipitation or improper adjustment of the acid, the filtrate should be tested
by being heated to 60C. and again saturated with H 2 S while the precipitate
is being washed.

11. The washing should be continued until the wash solution gives no
test for sulfide; this is necessary, as any sulfide left in the precipitate will
cause an error in the estimation of zinc by P. 62. The test may be made
conveniently by adding a drop of copper nitrate to the solution.

12. If the precipitate is dark colored, indicating that cobalt (or nickel, if
the acidity of the solution has not been properly adjusted) has been carried
out with it, it may be treated as follows:

Dissolve the precipitate in 20 ml of hot 6 n. H 2 S(>4, adding Br 2
water to the solution if it is needed. Boil out the excess of Br 2 , neu-
tralize the solution with NaOH as directed in the above procedure,
dilute it to 50 ml, add 1 ml of 6 n. H 2 S0 4 , saturate it with H 2 S, and
heat the mixture to 60 to 80C., keeping the gas flowing. Finally,
close the flask and allow it to stand for 10 min., filter out the precipi-

P. 62] ESTIMATION OF ZINC 331

tate, wash it, and treat it as directed in the last paragraph of the
procedure above. Combine the filtrate with the original filtrate.

P. 62. Estimation of Zinc

Discussion. This estimation is based upon the same principle
as was involved in the estimation of cadmium. As there, the method
consists in treating the sulfide precipitate with an excess of an acidi-
fied ferric sulfate solution and thereby causing the formation of
an equivalent amount of ferrous iron as a result of its reduction by
the sulfide; the ferrous iron is then titrated with permanganate.
Sulfur is precipitated but does not react sufficiently rapidly with
either the ferric iron or the permanganate to introduce serious
error. Experiments 18 have shown that this method is more rapid
and requires less experienced technique than the titration of a
zinc solution with standard ferrocyanide solution, 14 but that the
results are usually from 1 to 2 per cent high.

Procedure 62: ESTIMATION OF ZINC. Prepare a ferric
sulfate solution by dissolving 3 to 4 g of solid Fe 2 (S04)s (Note
4, P. 30) in 50 ml of boiling water in a 400-ml beaker.
Transfer the washed ZnS precipitate to the solution (Note
1), and heat the mixture (almost to boiling) until only a
coagulated residue of sulfur remains (usually 2 to 3 min.) ;
cover the beaker with a clock glass during this heating.
Add to the mixture 2 ml of H2S04 and again heat it for 1
min. Quickly cool the mixture with tap water, dilute it to
150 to 200 ml, and add to it 10 ml of H 2 S0 4 and 5 ml of 15
f. HsP04. Titrate the solution at once with standard
KMnO4 until a faint pink color persists throughout the solu-
tion for 10 sec. Stir the solution continuously during the
titration. From the volume of standard permanganate
solution used, calculate the amount of zinc present.

Notes:

1. If a very large precipitate is obtained, it is suggested that it be dis-
solved and an aliquot portion be used for the following analysis. By this
means, check titrations can be made and a better estimate can be obtained.
This can be done as follows:

Dissolve the precipitate as directed in Note 12, P. 61, boiling
out the Brj and cooling the solution. Transfer the cold solution to

18 Unpublished experiments by F. N. Laird.

14 Treadwell-Hall, Analytical Chemistry, Vol. II, Quantitative, 8th Ed.,
p. 672; Hillebrand and Lundell, Applied Inorganic Analysis, p, 335.

332 ANALYSIS OF THE ZINC GROUP [P. 63

a 100-ml volumetric flask, dilute it to the mark, mix the solution
thoroughly, and pipet 25 ml of it into a 200-ml conical flask. Neu-
tralize this solution and treat it further as directed in Note 12, P. 61.
Treat the precipitate obtained there by the above procedure; dis-
card the filtrate.

P. 63. Precipitation of Nickel and Cobalt (and Iron), and the
Separation of Iron from Nickel and Cobalt

Discussion. The filtrate from the zinc sulfide precipitation is
boiled to expel the hydrogen sulfide, and the nickel and cobalt are
then detected by the addition of an excess of sodium peroxide.
The nickel is precipitated as pale greenish nickelous hydroxide,
while cobalt is oxidized to the tripositive form and appears as the
black hydrous oxide. If the precipitate is not darkened, the absence
of cobalt is shown. The elements could be precipitated as their
sulfides, but this method is more rapid and furnishes more informa-
tion, and the precipitate is more easily handled and redissolved.

The Basic Acetate Separation. If iron has not been separated by
the ether treatment or removed as suggested in Note 3, P. 55, it
would precipitate with the Zinc Group sulfides (in P. 55) and appear
as ferric hydroxide when the solution is neutralized with sodium
peroxide. The separation of iron from the cobalt and nickel can be
effectively made at this place by redissolving the precipitate and
carrying out what is known as a "basic acetate" precipitation. In
this method the solution is so buffered by adding acetic acid and an
acetate in the proper ratio that, upon heating the solution, the ferric
salt hydrolyzes and precipitates. The precipitate is a mixture of
ferric hydroxide and basic ferric acetate (FeOC 2 H 3 O2) in which the
molal ratio of iron to acetate is usually approximately 2 to 1.

This process is extensively used in quantitative work for sepa-
rating ferric iron (and aluminum) from the bipositive elements of the
Ammonium Sulfide Group, and, by the addition of an excess of
iron, it affords a means of separating phosphate from the alkaline
earth elements. As this method is extensively used in both quanti-
tative and qualitative work, the results of a series of experiments 10
made to investigate the effects of various factors on the precipitation
and separation of certain elements are shown in Table XXVII.

It is seen from the data of this table that the proper ratio of acetic
acid to acetate is dependent upon the separations desired. Iron
can be precipitated from solutions in which this ratio is greater than

11 Experiments by the author. For a much more complete study of the
basic acetate method, especially in regard to its application to the rarer ele-
ments, see Noyes and Bray, Qualitative Analysis for the Rare Elements, Mac-
mill an, 1927, pp. 396-410.

TABLE XXVII

THE BEHAVIOR OF VARIOUS ELEMENTS IN THE BASIC ACETATE SEPARATION
In these experiments, 100 ml of a solution containing the elements listed,
NH 4 C1, NH 4 C2H 8 O2, and HC 2 H,O2, in the amounts shown, were heated just
to boiling in a covered flask for the time indicated, and filtered. The com-
position of the precipitate and filtrate was determined by analysis.

Milli-enuiva-
lents Present

Amount Found (mg) In

Expt.

Elements
Taken (mg)

Time
Soiled

'm in ^

O
W

L llllll . )

Precipitate

Filtrate

Fe 1

Fe 10

2 to 3

7 to 8

Fe 10

1 to 2

8 to 9

Fe 10

No ppt.

Fe 10

2 to 3

7 to 8

Fe 10

/Fe 101
lAl 10)

No ppt.

/10 Fe
\10 Al

/Fe 10\

IA1 10)

Ppt.

/Trace Fe
\3 to 5 Al

Fe300

300

Fe 300

Large ppt.

1 to 2

Fe 300

Large ppt.

15 to 20

/Fe 200 \
\Ce m 100)

No ppt. fr

Al 1

Turbidity

Al 2

Small ppt.

Al 10

Al 300

15 to 30

Cr 300

2 to 3

/Fe 300 \
ICr 300)

Large ppt.

/3 to 5 Fe
115 to 20 Cr

Zn 500

No ppt.

/Fe 300\
IZn 300)

/300 Fe \
\15 to 25 Zn)

Mn 200

No ppt.

/Fe 300 \
\Mn 200)

/Fe 300\
IMn 1 )

Large ppt.

/O Fe
10.6 Mn

/Fe 300 \
INi 1 )

Large ppt.

/O Fe
\0.5Ni

/Fe 300 \
\Ni 200)

/Fe 300\
\Co 1 )

Large ppt.

0.8 to 1.0 Co

/Fe 300 \
\Co 200)

The precipitate was colloidal and difficult to filter.

b No precipitation was observed until the solution was neutralized with
NH 4 OH.

333

334 ANALYSIS OF THE ZINC GROUP IP. 63

10 to 1, while with aluminum a much lower ratio is required; also,
if aluminum is present, a lower ratio is required for the precipitation
of the iron. Chromium is very incompletely precipitated and so
inhibits the precipitation of iron and aluminum that its presence
precludes the use of the method. The inhibiting effect of certain
elements is strikingly shown in Experiment 14, where the precipita-
tion of iron is completely prevented by the presence of a cerous salt;
this is probably due to stabilization of the precipitate in a colloidal
system by adsorption of the tripositive cerium. By comparison
with Table XIX, p. 284, it is seen that the separation of iron from
nickel and cobalt by this method is not so satisfactory as that which
could be obtained by an ammonia precipitation. However, only a
small amount of iron should be present here, and the presence of
ammonia in the solution is undesirable, since the nickel and cobalt
would have to be precipitated as their sulfides rather than as their
hydroxides, and, as stated above, the sulfide precipitates are much
more difficult to handle.

Other organic acids and their salts have been used for buffering
systems to approximately the pH used in the basic acetate method.
Included in this group are formic acid (K A 2 X 10~ 4 ), succinic
acid (KA, first hydrogen = 6.6 X 10~ 5 ), and benzoic acid (K A =
6.6 X 10 5 ). The benzoic acid system has been shown to give
unusually complete separations. 16

Procedure 63: PRECIPITATION OF NICKEL AND COBALT

AND THEIR SEPARATION FROM IRON. Boil the filtrate

from the zinc sulfide precipitation until the H 2 S is com-
pletely expelled (Note 1), and then add NaOH to the hot
solution until it is alkaline to litmus. (Light green pre-
cipitate, presence of nickel; blue precipitate, turning pink,
presence of cobalt; greenish-gray precipitate, presence of
iron. Note 2.)

If iron has been previously separated, treat the mixture
as directed in the last paragraph of this procedure.

If iron has not been previously separated, proceed as
follows: Make the mixture acid with 12 n. HC1, add 1 ml
in excess, and boil the mixture until the precipitate dissolves.
Cool the solution somewhat and add liquid bromine, 1 drop
at a time until an excess is present, and then boil it again
until the excess of bromine is expelled (Note 3). Cool the

18 Kolthoff, Stenger, and Moskovitz, J. Am. Chem. Soc., 66, 812 (1934).

P. 63] NICKEL AND COBALT 335

solution and add NaOH until the first permanent precipitate
is obtained. Dissolve this in HC1 (avoiding an excess),
add 1 ml of HC 2 HaO 2 and 10 ml of 3 n. NaC 2 H 3 O 2 (Note 4),
and boil the solution for 2 min. (Reddish precipitate,
presence of iron.) If a large precipitate forms, add 5 ml
more of the NaC 2 H 3 O 2 and again boil for 1 min. Filter
the mixture through a paper filter and wash the precipitate
with hot water (Note 5). Treat the precipitate as directed
in Note 6. Treat the filtrate as directed in the next para-
graph.

Cool the solution and add in 0.1-g portions, shaking after
each addition, Na 2 O 2 until the solution is alkaline, then
1 to 2 g in excess. (Light green precipitate, presence of
nickel; black precipitate, presence of cobalt. Note 7.)
Filter the precipitate on a paper filter, wash it with hot
water, and treat it by P. 64 (Note 8). Discard the filtrate.

Notes:

1. The H 2 S must be completely expelled, or the black sulfides will be
precipitated when the solution is made alkaline.

2. Information as to the elements present may be obtained from the
color of the precipitate. Nickelous hydroxide is pale green. Cobaltous
hydroxide, when first precipitated, is blue but turns pink on standing. This
effect is usually attributed to the formation of a blue basic salt which is con-
verted by excess of alkali to the normal hydroxide, 17 while according to
Hantzsch 18 the precipitate which first forms has the formula CoO-H 2
and the pink form has the formula Co(OH) 2 . In contradiction to these
theories, X-ray studies show no difference in the crystal structure of the
two forms; and from these and other studies Stillwell 19 has concluded that
the difference in color is due to the difference in degree of dispersion of the
precipitate. If iron is present, it will be mostly in the ferrous form. Pure
ferrous hydroxide is white, but it nearly always contains some ferric hy-
droxide and appears gray-green; on standing exposed to the air, it is rapidly
oxidized and becomes reddish brown. If a solution containing ferrous and
ferric iron is made alkaline, the precipitate may be black, corresponding to
the oxide FesC^.

3. Ferrous salts are not precipitated by the basic acetate process. Fer-
rous hydroxide is a stronger base than ferric hydroxide, and therefore ferrous
salts are not so easily hydrolyzed; also, ferrous hydroxide is considerably
more soluble.

The excess of bromine must be expelled, as, in a solution containing

17 Treadwell-Hall, Analytical Chemistry, Vol. I, Qualitative, 9th Ed.,
1937, p. 239.

18 Hantzsch, Z. anorg. C/iem., 73, 305 (1912).
i Stillwell, J. Phys. Chem., 33, 1256 (1929).

336 ANALYSIS OF THE ZINC GROUP [P. 64

HCaHaC^ and NaC2Ha02, cobalt may be oxidized by it and precipitated

4. NaC2H 3 O 2 , rather than NH^HaO*, is used, since the presence of
ammonium salts would prevent the subsequent precipitation of nickel and
cobalt by NaOH and Na 2 2 .

5. If the precipitate tends to become colloidal and pass through the filter,
use a wash solution prepared by adding 1 ml of 3 n.-NaC2H 3 02 and 0.1 ml
of HC 2 H 3 02 to 100 ml of hot water.

6. If only a small precipitate is obtained, the presence of iron should be
confirmed by dissolving it in 5 ml of HC1, adding 1 ml of 1 n. KSCN, and
noting if a red color is produced. The amount present can be estimated by
adding the ferric iron test solution to 5 ml of 6 n. HC1 and 1 ml of KSCN
until a color match is obtained, or by precipitating with NEUOH and com-
paring the precipitate with known amounts of iron precipitated under similar
conditions.

If iron has not been separated by the ether treatment and a considerable
precipitate is obtained, proceed as follows:

Pour 10 ml of warm 6 n. NH 4 OH repeatedly through the precipi-
tate and discard the solution. Dissolve the precipitate with 10 to
20 ml of HC1, evaporate the solution to 4 to 5 ml, dilute to 25 ml
and treat as directed in P. 53, beginning with the third paragraph.
The precipitate is treated with NKUOH in order to metathesize most of
the basic acetate to hydroxide and thus remove the acetate. The presence
of acetic acid would tend to inhibit the reaction between ferric iron and
iodide in P. 53.

7. By the addition of Na 2 02 the nickel hydroxide precipitate is not ap-
preciably changed in appearance, although it may be partly converted into
a nickelous peroxide; when the mixture is heated, the peroxide is decom-
posed. More powerful oxidizing agents, such as hypochlorite or peroxy-
sulfate, convert the nickel to a black oxide of approximately the compo-
sition NiaOa. The cobaltous hydroxide rapidly darkens because of the
formation of Co20 5 , which appears as a brown or a black precipitate, de-
pending upon the alkalinity and temperature of the solution. When cobalt
is being precipitated, the cobaltous solutions should not be treated with a
high concentration of alkali in the absence of an oxidizing agent, as co-
baltous hydroxide dissolves slightly under these conditions, forming a blue
solution which supposedly contains the ion Co02 =s .

8. If the precipitate is light green in color and not at all darkened, the
absence of cobalt can be assumed. In that case the precipitate can be
dissolved in 3 to 10 ml of HC1, and the solution can be evaporated to 2 to
3 ml and treated directly by P. 65; or, if the precipitate is small, it can be
treated as suggested in Note 2 of P. 64.

P. 64. Separation of Nickel and Cobalt

Discussion. The separation of nickel and cobalt which is used
here when large amounts of the two elements are present and a
quantitative separation is desired depends upon the difference in
the solubilities of the chlorides of the two elements in an ether

P. 64] SEPARATION OF NICKEL AND COBALT 337

solution saturated with hydrochloric acid gas. Under these con-
ditions, nickel chloride is precipitated, and cobalt remains in solu-
tion because of the formation of a complex ion. This compound is
so intensely colored that even 1 mg of cobalt will impart to the
ether solution a distinct blue color, which serves as the detection of
this element. This method for separating nickel and cobalt, de-
veloped by Havens, 20 is usually considered inadequate for quanti-
tative work because of the inclusion of cobalt in the precipitate.
Experiments 21 have shown that, when, under the conditions of this
procedure, 250 mg of nickel are precipitated in the presence of 250
mg of cobalt, not over 2 mg of cobalt are coprecipitated ; it was also
shown that 1 mg of nickel causes a precipitate and 1 mg of cobalt
causes a blue color, each of which is easily perceptible.

Coordination Compounds of CoboM. The color changes which
cobalt solutions undergo upon treatment with various reagents, such
as an alkaline cyanide, ammonia, or ammonia plus an oxidizing
agent, or upon evaporation of its solutions with concentrated solu-
tions of the halogen acids are due to the formation of complex, or
coordination, compounds of the type mentioned in the discussion of
P. 11. Thus, in dilute aqueous solutions there is evidence that the
cobaltous ion surrounds itself with six water molecules, forming the
ion Co(H20)e^ which is pink colored. This behavior is charac-
teristic of many other ions: thus, the blue color of aqueous cupric
solutions is due to the hydrated ion Cu(H 2 0)4 +f ; on dehydration
by heating with concentrated sulfuric or perchloric acid, the solution
becomes colorless. If an excess of a soluble cyanide is added to a
cobaltous solution, there is first formed the complex cobalt ocyanide
ion, Co(CN) 6 s ; if an oxidizing agent is present (even the oxygen
of the air), the much more stable cobalticyanide ion, Co(CN) 8 ,
is formed. These ions are analogous to the ferro- and forricyanide
ions, Fe(CN) 6 s and Fe(CN) 6 ~, respectively.

Cobaltous ion does not readily form ammonia complexes, but, as
was mentioned in P. 11, the tendency toward complex formation
increases as the positive charge on a given ion increases, so that, if
an oxidizing agent is added to an ammoniacal cobaltous solution,
an intense dark red to brown solution results. This is due to the
formation of various ions and compounds which vary according to
the concentration of the ammonia and of the other anions in the

10 Havens, Am. J. Sci. (4), VI, 396.

21 Unpublished experiments by Edwin McMillan and Carter Gregory.

338 ANALYSIS OF THE ZINC GROUP IP. 64

solution. Examples of such compounds are as follows:
Co(NH,)Cl" H ' f _ Co(NH 8 ) 4 Cl 2 ~, Co(NH 3 )3(N0 2 ) 3 (un-ionized),
Co(NH 8 )2(NO2)4~, and so forth. It is to be noted that in all cases
a total of six negative or neutral groups are attached, cobalt having
a coordination number of six in these compounds.

The cause of the change in color from pink to blue as a cobaltous
chloride solution is evaporated or treated with concentrated hydro-
chloric acid has been the object of much study, with two general
theories being advanced: first, that the color change was due to a
change in the degree of hydration of the ion, this being represented
by an equation of the type

Co(H,0). + + + 2C1" = CoCl 2 -2H 2 O + 4H 2 0,

and, second, that the change was due to the transformation from a
hydrated complex compound to one containing halogen ion, probably
as follows:

Co(H 2 0) 6 ++ + 4CF = CoCir + 6H 2 0.

However, it has been shown by Donnan and Bassett 22 that in the
blue solution the cobalt migrates to the anode, while in the pink
solutions it migrates toward the cathode. Further work by Bassett
and Croucher 23 has demonstrated that the blue hydrochloric acid
solutions contain the negative ions CoCir* and Co(Cl 3 H 2 0)~. In
the ether solution saturated with hydrochloric acid, the cobalt
probably exists as H 2 CoCl 4 .

Procedure 64: SEPARATION OF NICKEL AND COBALT.
Method for the More Exact Separation of Large Amounts of
Nickel and Cobalt. Dissolve the hydroxide precipitate
(from P. 63) on the filter by pouring repeatedly over it 10 to
20 ml of hot HC1 (Note 1). If it dissolves slowly, pour 1 ml
of 3 per cent H 2 2 over the precipitate, again treat it with the
HC1, and then evaporate the solution to a volume of 4 to
5 ml (Note 2).

Transfer the HC1 solution to a large test tube (15 mm by
200 mm; Note 3), using 3 to 5 ml of HC1 to wash out the
flask, and again evaporate the solution to 4 to 8 ml (Notes 4,
5). Cool the solution, immerse the test tube in a beaker

" Donnan and Bassett, J. Chem. Soc., 81, 939 (1902).
13 Bassett and Croucher, J. Chem. Soc.. 1930, 1784.

P. 64] SEPARATION OF NICKEL AND COBALT 339

containing a mixture of ice and water, add 15 to 25 ml of
ethyl ether, and (working under a hood and with no flames
in the vicinity) pass a fairly rapid stream, 2 or 3 bubbles a
second, of dry HC1 gas (Note 1, P. 94) through the solution
until it is saturated, as shown by the gas being no longer
absorbed (Note 6). (Yellow precipitate, presence of nickel;
blue solution, presence of cobalt.)

If nickel is present, cork the tube and let it stand; in a
similar test tube saturate another 15- to 30-ml portion of
ether with HC1 gas. Filter the mixture by decantation
through an asbestos or, preferably, a sintered-glass filter
(which has been sucked dry of water and then washed with
2 to 3 ml of the ether-HCl solution), and wash the precipi-
tate by decantation with 2- to 5-rnl portions of the ether-
HCl solution. Collect the washings with the filtrate in a
400-ml flask. Treat the precipitate by P. 65.

If the filtrate has a detectable blue color, treat it by P. 66;
if it is colorless, discard it.

Notes:

1. If the precipitate is large, it can be more advantageously dissolved
by transferring most of it, using a stirring rod, to the vessel in which it was
precipitated, and heating it there with the HC1.

2. If the precipitate is small, or if it is of moderate size (containing not
over 200 mg of the two elements) and an exact separation and estimation is
not required, the two elements may be detected and the amount of each
more rapidly estimated by treating the HC1 solution by Optional Pro-
cedure 64 A.

3. If a test tube of this size is not available, a small flask may be used,
although it is not as efficient a vessel in which to saturate the solution with
HC1 gas.

4. The volume of solution remaining should be adjusted to the amount
of the two elements thought to be present. The precipitation of nickel is
more complete, and the color of the cobalt is more pronounced, when only
the minimum volume of water is left and the minimum amount of ether is
added later. When large amounts of the elements are present, the larger
volumes should be used to prevent coprecipitation of cobalt with the nickel
chloride.

5. The solution tends to bump during this evaporation so that it has to
be kept in continual motion. It is advisable to keep an ebullition tube
(Note 3, P. 42) in the solution, or, if one is not available, a light-weight
stirring rod (one cut from capillary tubing without rounding the edges is
very effective) may be used.

6. Saturation of the solution is also indicated by the aqueous and ether
layers becoming miscible.

340 ANALYSIS OF THE ZINC GROUP [P. 64,4

P. 644. Optional Method of Detecting Nickel and Cobalt

Discussion. An optional method of detecting and estimating
nickel and cobalt is provided for use when the amount of the two
elements is small (as judged by the precipitates obtained in P. 63),
or where exactness is not required and only a rapid method of detect-
ing and estimating the two is desired. In this method the solution
is divided, and the nickel is detected in one portion by precipitating
it with dimethylglyoxime (diacetyl dioxime). Dimethylglyoxime,
(CH 3 ) 2 C2(NOH) 2 , is an organic compound which acts in the nature
of a weak acid and forms a salt with nickel, ((CH 8 ) 2 C 2 (NOH)NO) 2 Ni,
in which only one of the hydrogen atoms is displaced. This com-
pound is so insoluble in alkaline or weakly acid solutions and is so
voluminous and intensely red colored that it affords an extremely
sensitive test for nickel; thus, in 50 ml of a solution made slightly
alkaline with ammonia, 0.05 mg of nickel gave the characteristic
precipitate within 1 min., while 0.02 mg could be detected when the
solution was allowed to stand for 15 min. When large amounts of
cobalt are present, an intense brown color is formed, the precipita-
tion of nickel is inhibited, and an excess of the reagent must be added.
These effects are due to the formation by the cobalt of a stable
soluble compound with the dimethylglyoxime. They can be partly
overcome by making the precipitation from a solution containing
acetic acid and acetate; thus, under the conditions of this procedure,
0.05 mg of nickel can be readily detected in the presence of 50 mg
of cobalt.

Cobalt is detected in another portion of the solution by precipi-
tating it as potassium cobaltinitrite, K3Co(NO 2 ) 6 . This compound
is formed by the addition of potassium nitrite to the cobaltous salt
in a slightly acid solution. Under these conditions the dipositive
cobalt is oxidized to the tripositive state and forms the nitrite com-
plex, Co(N0 2 )e", the potassium salt of which is insoluble. The
oxidation of cobalt to the tripositive state by nitrite, an apparent
anomaly, is due to the stability of the complex ion, which reduces
the concentration of the simple tripositive ion to an exceedingly
small value, and to the concentration of the complex ion being
maintained at a small value by its precipitation as the potassium
salt. The dipositive nickel is not appreciably oxidized by the nitrite,
and, although nickelous nitrite ion, Ni(NO 2 ) 6 % is formed, the salt
K 4 Ni(NO2) is relatively soluble. If alkaline earth elements are
present, the separation is unsatisfactory, as yellow precipitates of
the type BaK 2 Ni(N02)e are then formed. The proper conditions

P. 644] DETECTION OF NICKEL AND COBALT 341

for the quantitative precipitation of cobalt require a high concen-
tration of nitrite ion in order to cause the formation of the complex
ion at an appreciable rate and a high concentration of potassium ion
in order to reduce the solubility of the precipitate. The solution
should be only slightly acid, as the precipitate is soluble in strong
acids; it should not be basic, as the complex ion is decomposed by
hydroxyl ion with precipitation of cobaltic oxide.

Procedure 64LA: DETECTION OF NICKEL AND COBALT.
Optional Method for Use with Small Amounts or When Only
Qualitative Information Is Desired. Dissolve the hydroxide
precipitate (from P. 63) as directed in the first paragraph
of the procedure above.

Cool the solution, transfer it to a 50-ml volumetric flask,
dilute it to the mark, mix it thoroughly, and pipet 5 to 25 ml
(Note 1) into a 200-ml flask. Add a slight excess of
NH 4 OH (filter out any precipitate), dilute the solution to
50 ml, and add 5 to 20 ml of a 0.1 n. solution of dimethyl-
glyoxime in ethyl alcohol. (Red precipitate, presence of
nickel. Note 2.) Compare the precipitate with that pro-
duced under similar conditions by known amounts of nickel,
or collect it on a weighed Gooch crucible or, preferably, one
of sintered-glass (Note 3), wash it with hot w r ater, dry it at
110C. for 1 hr., and cool and weigh it.

Transfer 10 to 25 ml of the remainder of the solution
to a 100-ml flask and evaporate it to 3 to 5 ml. Neutralize
the solution with KOH, make it acid with HC 2 H 3 O 2 , and
then add 5 ml in excess. Add 5 ml of 6 n. KN02, cool the
mixture, shake it vigorously, and let it stand, shaking it
frequently, for at least 10 min. (Yellow precipitate, pres-
ence of cobalt.) If a large precipitate is obtained, add
2 ml more of HC 2 H 3 O 2 and 5 ml of 6 n. KNO 2 (Note 4).
Compare the precipitate with that produced under similar
conditions by known amounts of cobalt, or, if a more precise
estimation is desired, treat another portion of the solution
in the volumetric flask by the last paragraph of P. 66
(Note 5).

Notes:

1. Too large a portion of the solution should not be taken, as an excessive
volume of the reagent may be required; also, over 10 to 20 mg of nickel
produce so voluminous a precipitate that it is difficult to filter and wash.

342 ANALYSIS OF THE ZINC GROUP [P. 65

2. If considerable cobalt is present, the solution may be so dark colored
as to prevent determining if a red precipitate is present. In this case make
the solution just acid with HCaHsOj and let it stand for 10 min.

3. If a quantitative precipitation is desired, the solution should be allowed
to stand for 30 min. to 1 hr. before being filtered. The filtrate should be
tested by adding more of the dimethylglyoxime reagent and allowing it to
stand.

4. The quantitative precipitation of considerable amounts of cobalt
requires from 12 to 24 hr.; smaller amounts are precipitated in a much
shorter time. Since 0.1 mg of cobalt can be readily detected, small portions
of the original solution should be used if considerable cobalt is present.
Completeness of precipitation should be tested for by adding more nitrite to
the filtrate and allowing it to stand.

5. If no more of the original solution is available, add cautiously 5 ml
of 36 n. HjSO* to the mixture, evaporate it until the acid fumes, cool the
solution, dilute it with 25 ml of water, and treat it by the last paragraph of
P. 66.

P. 65. Estimation of Nickel

Discussion. In the method for estimating nickel which is used
here, the nickel chloride solution is titrated with a standard cyanide
solution, the principal reaction being the formation of the nickelous
cyanide complex ion, Ni(CN) 4 ". As this compound is unstable in
an acid solution, the nickel chloride precipitate (from P. 64) is
dissolved in an ammoniacal solution and is thus largely converted
into the tetrammino nickel complex ion, Ni(NH 3 )4 + *. However,
upon addition of cyanide, the cyanide complex is rapidly formed,
as it is much less dissociated. The tetrammino nickel ion is blue,
but this color is not sufficiently intense for its disappearance to be
used as the end-point. The end-point is therefore determined by
first adding a known volume of standard silver nitrate and then an
excess of potassium iodide, which produces a colloidal precipitate of
yellowish silver iodide. When this mixture is titrated with the
cyanide, the complex nickelous cyanide anion is formed first; after
an amount of cyanide equivalent to the nickel has been added,
further addition of cyanide forms the Ag(CN) 2 ~ anion, which causes
the precipitate of silver iodide to dissolve. The end-point is taken
when the last yellowish opalescence due to the silver iodide precipi-
tate disappears. For a discussion of this end-point, see P. 23.

It is, of course, necessary to correct for the amount of silver which
has been added. As cobalt also forms a cyanide complex ion, it
should not be present in appreciable amounts. The solution should
be cold during the titration. Only a slight excess of ammonia should

P. 65] ESTIMATION OF NICKEL 343

be present, as otherwise the conversion of the ammino complex to
the cyanide complex takes place more slowly.

Procedure 66: ESTIMATION OF NICKEL. Dissolve the
NiCU precipitate (from P. 64) from the filter by pouring
through it, in small portions, 10 to 20 ml of hot water. Col-
lect the solution in the test tube and dissolve the precipitate
remaining there. Transfer the solution to a 200-ml flask,
and add 5 ml of 3 n. NEUC1 and then NH^OH until a clear
blue color is obtained and the solution is alkaline, taking
care to avoid an excess of the NH 4 OH (Note 1). Cool the
solution.

If less than 100 mg of nickel are thought to be present,
treat the solution as directed in the last paragraph of this
procedure.

If more than 100 mg of nickel are thought to be present,
transfer the cold solution to a 100-ml volumetric flask, dilute
it to the mark, mix the contents thoroughly, and pipet from
it into a 200-ml flask a volume of solution which is thought to
contain 50 to 100 mg of nickel.

Dilute the solution to 100 ml, and add to it 1 ml of 1 n.
KI and then (from a buret or graduated pipet) 0.10 ml of
standard 0.1 n. AgN0 3 solution. Add the AgN0 3 dropwise
while swirling the solution. Titrate the solution with a
standard 0.2 f . KCN solution, adding the KCN rapidly as
long as the turbidity does not decrease and then slowly,
swirling the mixture constantly, until the yellowish pre-
cipitate dissolves. If it is thought that the end-point has
been overrun, titrate the clear solution with 0.1 n. AgNOs
solution until a permanent precipitate is again produced.
From the volumes of standard AgN0 3 and KCN used,
calculate the amount of nickel present (Note 2).

Notes:

1. If a flocculent precipitate, usually due to aluminum from reagents or
vessels, persists after the solution is made alkaline, it can be neglected if it
is not sufficient to interfere with the detection of the end-point; however,
it is preferable to filter it out.

2. The standard KCN solution is not stable and may be standardized at
the completion of the titration by pipeting a known volume of the standard
AgN0 3 into the solution and again titrating with KCN to the end-point.

344 ANALYSIS OF THE ZINC GROUP [P. 66

P. 66. Estimation of Cobalt

Discussion. This method for the estimation of cobalt, developed
by Erigle and Gustavson, 24 and studied further by Willard and Hall, 25
is based upon the oxidation of cobalt in an alkaline solution to
cobaltic oxide by sodium peroxyborate, NaB0 3 . Nickel, if present,
is not oxidized to the higher oxidation state; it may form a peroxide
compound, but this is decomposed by heating. The excess peroxy-
borate is destroyed by merely boiling the alkaline mixture, with
the evolution of oxygen. The precipitate is then reduced by an
excess of iodide in an acid solution, an equivalent amount of iodine
being liberated; this is titrated with a standard thiosulfate solution.
Experiments 26 have shown that the cobaltic oxide dissolves more
rapidly if hydrochloric acid is used for acidifying the mixture instead
of sulfuric acid, as originally proposed, and also that low results are
obtained when large amounts of cobalt (over 50 mg) are present.
The article by Willard and Hall contains a discussion of other
methods for the volumetric estimation of cobalt.

Procedure 66: ESTIMATION OP COBALT. In case a blue
color was obtained in P. 64, evaporate the ether solu-
tion (under a hood by immersing the flask in a beaker of
water which has been heated to boiling and from which the
flame has then been removed) until the ether is expelled;
then evaporate the aqueous solution almost to dryness.
Dissolve the residue in 25 ml of water.

If less than 50 mg of cobalt are thought to be present,
treat the solution as directed in the last paragraph of this
procedure.

If more than 50 mg of cobalt are thought to be present,
cool the solution and transfer it to a 100-ml volumetric
flask, dilute it to the mark, mix the contents, and pipet
from it into a 400-ml flask (preferably one with a ground-
glass stopper) that volume of the solution which is thought
to contain 40 to 60 mg of cobalt.

Make the solution just alkaline with NaOH, gradually
sprinkle into it 1 to 2 g of sodium peroxyborate (NaBOs),
and then add 10 to 15 ml of NaOH (Note 1). Swirl the
mixture until the contents are thoroughly mixed, and then

14 Engle and Gustavson, Ind. Eng. Chem., 8, 90 (1916).
Willard and Hall, J. Am. Chem. Soc. t 44, 2237 (1922).
16 Unpublished experiments by Dr. Lucas Alden.

P, 66] ESTIMATION OF COBALT 345

boil it vigorously for 10 min. Cool the mixture (Note 2),
add to it 2 g of KI dissolved in 10 ml of water, washing
down the sides of the flask and mixing the solutions, and
then add 50 ml of HC1. Again cool the mixture, close the
flask, and let it stand for 5 min., or until the precipitate
dissolves. Dilute the solution with 100 ml of water which
has been recently boiled and cooled, and then titrate it with
0.1 n. Na 2 S 2 3 until the brownish-yellow iodine color be-
comes indistinct. Add 5 ml of starch indicator solution
and again titrate until the starch blue color disappears.
From the volume of standard Na 2 S 2 8 used, calculate the
amount of. cobalt present.

Notes:

1. If sodium peroxyborate is not available, substitute for it 2 to 3 ml
of solid Na 2 2 . Cool the mixture during the addition of the Na 2 2 .

2. If a carbon dioxide generator (or tank) is available, the flask should be
swept with the gas sufficiently to remove all the air from it before adding the
KI and acidifying. This minimizes the probability of error due to oxidation
of the iodide in the hydrochloric acid solution by the oxygen of the air.

TABULAR OUTLINE VIII
THE ANALYSIS OF THE ALUMINUM GROUP

Filtrate from P. 55: A1(C,O 4 ) 3 -, Cr(C a O 4 ),-, Mn(C 2 O 4 )r, C,Or, HCOr, HS-,

Na+, NH 4 +, H,S

Make acid with HNOi, heat. (H 2 S and CO 2 expelled.)
Evaporate. Add 16 n. HNOt, evaporate almost to dryness.
Add HNO Z and KCIO>, heat. (CO,, C1O 2 ) (P. 71)

Precipitate: MnO 2

Dissolve in HCl and

Solution: Mn++, I 3 ~
Titrate with

starch indicator.
Solution: Mn**,

S 4 6 -

(P. 72)

Filtrate: AI+++, Cr 2 O 7 -,
Neutralize with NH 4 OH. (P. 73)

Precipitate:

Al(OH),
Heat and weigh.
(P. 74)

Filtrate: CrOr, NH 4 +, NO,-,
NH 4 OH

Make acid with HiSO*.
Add KI.
Solution :Cr+++, Ir

Titrate with NatStOi, starch indi-
cator.

Solution: Cr+++, I-, S 4 Or
(P. 75)

The Analysis of the Aluminum Group

P. 71. Destruction of Oxalate, Precipitation of Manganese, and
Detection of Chromium

Discussion. The filtrate from the H 2 S precipitation of the Zinc
Group contains aluminum, chromium, and manganese in the form
of complex oxalates. It is necessary that the oxalate be removed
from the solution, as neither aluminum nor chromium can be pre-
cipitated in its presence. The procedure adopted for destroying
the oxalate consists in evaporating with concentrated nitric acid,
which causes partial oxidation of the oxalate, and then completing
this process by adding chlorate to the nitric acid solution. Two
other desirable results are attained: The manganese is oxidized by
the chlorate under these conditions and precipitates as the hydrated
dioxide, thus being separated from aluminum and chromium, and
the chromium is oxidized to chromate, in which form it can be sepa-
rated from the aluminum by precipitating the latter with ammonium
hydroxide. The oxalate could have been destroyed by fuming the
filtrate with sulfuric acid, but this process causes the formation of
the difficultly soluble chromic dihydroheptasulfate, C^HaCSOO?,
which is so slowly dissolved by ordinary reagents that the process
was not adopted; furthermore, manganese dioxide is not readily
precipitated by chlorate from solutions containing sulfuric acid,
as soluble complex compounds of tripositive manganese are formed.

It was found that the presence of manganese very markedly cat-
alyzed the oxidation of oxalate by nitric acid; for this reason, a
simple preliminary test for manganese is made on a portion of the
solution, and, if it is absent, a moderate amount is added. A further
advantage of having manganese present is that the precipitation of
the dioxide is proof that all oxalate has been destroyed and that
further addition of chlorate for that purpose is unnecessary; the
chromium is also oxidized to chromate before precipitation of the
manganese dioxide occurs.

The Precipitation and Separation of Manganese as Dioxide by Chlo-
rate in a Nitric Acid Solution. The precipitation of manganese from
a nitric acid solution by chlorate is a very distinctive test for this
element and also can be used to separate it from all the other ele-
ments of the Ammonium Sulfide Group. The extent to which other
elements are carried down with the manganese was determined by

347

348 ANALYSIS OF ALUMINUM GROUP [P. 71

treating 250 mg of manganese and 250 mg of each of the other ele-
ments of the Ammonium Sulfide Group in separate solutions by the
last paragraph of this procedure. The precipitates were then
analyzed to determine the amount of the other element coprecipi-
tated. The amounts found (mg) were as follows: aluminum, 3;
chromium, 10; iron, 25; zinc, 3 to 5; nickel, 4; and cobalt, 10. As
there has been some question as to how quantitative the precipitation
of manganese can be made by this method, the filtrates were also
examined; in none of the above experiments was more than 0.1 to
0.2 mg of manganese found in the filtrate, and in repeated experi-
ments 0.3 mg of manganese was detected without difficulty, even in
the presence of large amounts of aluminum. When large amounts
of chromium were present, the detection of 0.3 mg was uncertain but
0.7 mg gave a readily detected precipitate. Small amounts of man-
ganese such as these often gave first a pink color, which changed into
a brownish suspension and then, on heating, into a dark precipitate.

Procedure 71: PRECIPITATION OF MANGANESE AND DE-
TECTION OF CHROMIUM. Make the filtrate from P. 55 (con-
tained in a 500-ml flask) just acid to litmus with HNO 3 and
add 5 ml of acid in excess. Evaporate the solution until
the H 2 S has been expelled and the volume has been reduced
to slightly less than 100 ml, and then cool and dilute it to 100
ml (Notes 1 and 11). Pipet 10 ml of the solution into a
100-ml flask and add to it 5 ml of 6 n. KOH (Note 2) and
0.5 ml of solid Na2C>2. Boil the solution for 1 or 2 min.
(Brownish-black precipitate, presence of manganese; yel-
low solution, presence of chromium. Note 3.)

If there is no precipitate, return this portion to the main
body of the solution. Add 2.5 ml of 1 n. manganous nitrate
solution (Note 4) and treat the solution according to the
last paragraph of this procedure.

If the precipitate is small, 1 mg or less (Note 5), filter the
mixture through an asbestos filter and wash the precipitate
with two 5-ml portions of hot water. Combine the filtrate
and washings with the main body of the solution. Add 2.5
ml of 1 n. manganous nitrate solution (Note 4) and treat
the solution according to the last paragraph of this proce-
dure. Treat the precipitate by P. 72.

If the precipitate is large, greater than 1 mg (Note 5),
return the mixture to the main body of the solution and
treat it by the next paragraph.

P. 71] MANGANESE AND CHROMIUM 349

Evaporate the solution until salts crystallize or until the
volume has been reduced to 3 to 5 ml. Add cautiously
(Note 6), in 1-ml portions, 10 ml of 16 n. HNO 3 , allow the
flask to stand without heating as long as a vigorous reaction
continues, and then evaporate the solution slowly to 3 to 5
ml. Repeat the treatment with HNOs and again evaporate.
Add 20 ml of 16 n. HNO 3 , heat the solution to boiling, and
add finely powdered KClOs in 0.1-g portions (Care: Note 7)
until a precipitate of manganese dioxide has been obtained
which does not disappear upon boiling the mixture (Note 8)
for 1 min. Boil the solution gently between each addition
of chlorate until any greenish-yellow vapors of ClOa are
expelled from the flask. Continue to add KC10 3 (1 to 5 g)
in small portions until no more manganese is precipitated
and C1C>2 is no longer rapidly evolved. Dilute the mixture
to about 30 ml with water and filter it through an asbestos
(or sintered-glass) filter, retaining as much as possible of the
precipitate in the flask (Notes 9, 10). (Orange-red solution,
presence of chromium.) Wash the precipitate by decanta-
tion with three 5-ml portions of hot water, combining these
washings with the filtrate. Treat the filtrate by P. 73. If
manganese was added in the earlier part of this pro-
cedure, discard the precipitate; if no manganese was added,
wash the precipitate with hot water until it is free of acid
and treat it by P. 72.

Notes:

1. A 100-ml graduate is sufficiently precise for this dilution, for the
aliquot portion to be used for the detection of manganese is analyzed for
that element only when 1 mg or less is present; and the volumetric precision
obtained with the use of a graduate is well within the limits of accuracy of
the colorimetric method used for the estimation of small quantities of
manganese.

2. Because chromium and manganese are present largely as the oxalate
complexes, the quantity of KOH added must exceed the amount of am-
monium salts (about 17 milli-equivalents) present so that the OH~ con-
centration will be sufficiently high for the peroxide to oxidize these elements
completely. KOH rather than NaOH is used to prevent the precipitation
of Na 2 C 2 4 .

3. Experiments have shown that under these conditions 0.1 mg of
manganese and 0.5 mg of chromium are easily detected.

If iron, nickel, or cobalt were incompletely removed from this solution,
they would behave as follows: Iron would give a grayish-green precipitate
of Fe(OH)2, which upon the addition of Na2C>2 would be oxidized to reddish-
brown gelatinous Fe(OH) 3 ; cobalt would give a purplish-red solution due to

350 ANALYSIS OF ALUMINUM GROUP [P. 71

the formation of a complex ion with the ammonia, and this would decom-
pose and precipitate black Co 2 0| when the alkaline peroxide solution was
heated; nickel would remain in solution because of the formation of the
bluish tetrammino ion. These effects will be more or less masked when
much manganese or chromium is present.

4. The presence of 10 mg or more of manganous ion catalyzes the oxida-
tion of oxalate by concentrated HNOa and reduces appreciably the quantity
of KClOa required to complete the oxidation. Furthermore, since Mn02
is reduced in acid solution by oxalate, no permanent precipitate of Mn02
will be formed until all the oxalate has been destroyed. Hence the forma-
tion of a permanent black precipitate indicates the complete elimination
of oxalate.

If a 1 n. manganous nitrate solution is not available, 5 ml of manganese
test solution containing 10 mg/ml may be substituted.

5. If there is any doubt as to the size of the precipitate, it should be
compared with the precipitate obtained by oxidizing 1 mg of manganese to
the dioxide according to the first paragraph of this procedure.

6. Concentrated nitric acid in the presence of manganous salts reacts
very vigorously with oxalate, and, unless the acid is added slowly, the solu-
tion may foam out of the flask.

7. The action of KClOa in concentrated HNOa is rapid and vigorous,
giving as one of its decomposition products unstable C102, an easily identi-
fied greenish-yellow gas. The chlorate is added in small quantities, and the
solution is boiled between additions to prevent the accumulation of C102J
large quantities of this may explode dangerously when overheated. Under
the conditions of this procedure, when the chlorate has been added in small
quantities, the only evidence of such explosions has been gentle puffs of gas
from the mouth of the flask. However, to eliminate the danger of explo-
sions, the addition of large quantities of KClOa at one time must be avoided.

8. With 10 g of (NH 4 ) 2 C20 4 present, from 5 to 10 g of KC10 3 , depend-
ing upon the quantity of manganous ion present, are required to oxidize
the oxalate completely.

9. A paper filter cannot be used, as it would reduce the Mn02 and be
attacked by the concentrated HNO 3 .

10. The filtrate should be tested for manganese by adding a small amount
of KClOs and boiling the solution for a few minutes. If any precipitate
does form, evaporate the solution to 5 ml, add 10 ml of 16 n. HNOa, and
treat the solution with small portions of KClOa until no additional pre-
cipitate forms.

Very often the oxidation of manganese with KClOa in HNOa solution
produces a visible permanganate color; the quantity is usually so small
that it does not interfere in the subsequent procedures.

11. If an adequate supply of the sample of the original material is avail-
able, a more sensitive test for manganese can be made and the division of
the solution described in the remainder of this first paragraph can be avoided
by proceeding as follows:

Transfer 0.2 g of the original sample to a large test tube, add 5
ml of 16 n. HNOa, heat the mixture almost to boiling until solvent
action seems complete, cool, and add 10 ml of cold water. Add 2 to

P. 72] ESTIMATION OF MANGANESE 351

3 g of solid sodium bismuthate, shake the mixture, and allow any
residue to settle. (Purple color, presence of manganese.)
See the discussion and Note 1 of P. 72 in regard to this method of detect-
ing and estimating manganese. Organic matter would in most cases pre-
vent this test by reducing the permanganate; it should be eliminated by
P. 4,) and the bismuthate should be added to the perchloric acid solution
obtained there.

P. 72. Estimation of Manganese

Discussion. After the manganese has been obtained in the form
of the dioxide (P. 71), it is conveniently estimated by treating it
with an excess of potassium iodide in a hydrochloric acid solution
and titrating the iodine liberated with standard thiosulfate solution.
The reaction by which the precipitate is dissolved may be repre-
sented as follows:

Mn0 2( .) + 3I~ + 4H+ = Mn + + + I 3 ~ + 2H 2 0.

Native manganese dioxide (the mineral pyrolusite) is dissolved only
slowly by iodide in a cold hydrochloric acid solution, but the freshly
precipitated material is much more reactive and usually dissolves
in a few minutes.

An optional method consists in treating the manganese dioxide
with an excess of standard ferrous sulfate, which reduces the man-
ganese dioxide and is itself oxidized to the ferric state; the excess of
the ferrous salt is then titrated with a standard permanganate solu-
tion. This method is commonly used when standard solutions of
ferrous sulfate and permanganate are readily available and when
it is desired to eliminate the use of the relatively expensive potassium
iodide. A procedure for the use of this method is given in Note 3
below; this method can also be utilized for the analysis of pyrolusite
ores. However, it is a different method, is not as rapid as the iodide
method, and (where the amount of manganese is not approximately
known, so that the excess of ferrous sulfate can be kept small) is not
so conveniently performed as is the iodide method.

A very sensitive test for manganese can be made by oxidizing it in
an acid solution to permanganate; the color thus obtained can be
compared with known amounts of manganese and a colorimetric
estimation can be made, or with larger amounts the permanganate
can be reduced with a standard solution of a reducing agent. It is
recommended that the presence of manganese in the sodium per-
oxide precipitate from the preliminary test in P. 71 be confirmed by
this method; if it is desired, a colorimetric estimation can be made.

352 ANALYSIS OF ALUMINUM GROUP [P. 72

If the chlorate precipitate from P. 71 is small (25 mg or less), the
presence of manganese in it can be confirmed and the amount pres-
ent can be estimated by this procedure. A procedure for this
purpose is given in Note 1 below.

In that procedure, sodium bismuthate is used as the oxidizing
agent, the reaction taking place being represented by the following
equation :

5HBiO 3 + 2Mn+ + + 9H + - 2Mn0 4 ~ + 5Bi w + 7H 2 0.

The HBiO 3 is formed by adding commercial sodium bismuthate
(usually a mixture of sodium bismuthate and bismuth tetroxide)
to a nitric acid solution; HBiO 3 is unstable in acid solutions, oxygen
and tripositive bismuth being formed. This oxidation of manganese
to permanganate is frequently used as the basis of both colorimetric
and oxidimetric determinations of manganese. The colorimetric
method is restricted to small amounts (usually less than 2 mg in 100
ml), as the color comparison is unsatisfactory with more concen-
trated solutions; in the oxidimetric method the permanganate formed
is titrated with a standard solution of a reducing agent, usually
ferrous sulfate. In the oxidimetric method difficulty is experienced
in obtaining complete oxidation of the manganese if more than 50
to 100 mg per 100 ml are present. Other oxidizing agents which are
similarly used for oxidizing manganese to permanganate are lead
dioxide, periodate, and peroxysulfate; 1 bismuthate is used here
because it reacts rapidly in a cold solution.

Procedure 72: ESTIMATION OP MANGANESE. Transfer
as much as possible of the MnO2 precipitate (Note 1) on the
filter back to the flask containing the remainder of it by
tipping the perforated porcelain plate on edge and washing
the precipitate and asbestos from the funnel with about 25
ml of water (Note 2) . Dissolve 2 to 5 g of KI (or see Note 3)
in 40 ml of water, add 10 ml of HC1, and then wash the
funnel and plate with this solution, collecting it in the flask.
Close the flask and gently swirl the contents until all of the
MnO2 precipitate has apparently dissolved. Titrate the
solution with standard Na2S 2 Os until the iodine color be-
comes indistinct, add 5 ml of starch indicator solution, and
again titrate until the blue color disappears. From the
volume of Na2S2Oa used, calculate the amount of manganese
present.

1 For discussions of these methods, see Lundell, Hoffman, and Bright,
Chemical Analysis of Iron and Steel, Wiley, 1931, pp. 190-208.

P. 72] ESTIMATION OF MANGANESE 353

Notes:

1. If the precipitate is small (25 mg or less), it can be estimated more
quickly and a confirmatory test can be obtained by using the method given
below. In this case proceed as follows:

Dissolve the precipitate on the filter by pouring dropwise through
it 5 to 10 ml of HNOa to which has been added 2 ml of 3 per cent
H202. Collect the solution in the original flask and dissolve any
precipitate there; boil for 2 to 3 min. (to destroy the peroxide).
Transfer the solution to a large test tube, washing the filter and
flask with 3 n. HNOs, and cool it to room temperature. The
total volume should be from 20 to 50 ml, depending upon the
amount of manganese thought to be present. Add 2 to 3 g of
solid sodium bismuthate and allow the mixture to stand until the
excess settles. (Purple color, presence of manganese.) If all of
the bismuthate dissolves, add more, in 0.2-g portions, until an
excess is present. Add standard permanganate to the same
volume of 3 n. HNO 3 in a similar test tube until the colors match.

If the color is too intense to make a comparison, filter the mix-
ture through a sintered-glass (or asbestos) filter, wash the residue
with 25 ml of cold 0.6 n. HNOs, and collect the filtrate and washings
in a 200-ml flask. Add 2 ml of 85 per cent H 3 P0 4 and titrate
dropwise with standard ferrous sulfate solution until the per-
manganate color disappears. From the volume of ferrous sulfate
used, calculate the amount of manganese present.

2. Possible loss of precipitate or solution is minimized by inserting a
large wide-stem funnel in the flask.

3. As KI is an expensive chemical, the following method may be used
instead of the one described if standard permanganate and ferrous sulfate
solutions are available:

Add, with a 10- or 25-ml pipet, standard ferrous sulfate solution
to the precipitate, running the first portions through the funnel
and shaking the mixture vigorously after each addition. (The
amount of ferrous sulfate required should be approximately calcu-
lated from the size of the precipitate, and a 20 to 50 per cent excess
should be added.) Add 50 ml of 3 n. H2S04, close the flask with a
clean rubber stopper carrying an inlet and outlet tube (similar to
that used for the H 2 S treatment in P. 11), and sweep out the flask
with a stream of carbon dioxide (from a tank or generator). Heat
the mixture almost to boiling until all of the precipitate is dis-
solved (keeping a very slow stream, 1 or 2 bubbles a second, of
C02 passing through the flask), immediately cool the solution to
room temperature, add 3 ml of 85 per cent H 3 PO 4 , and titrate with
standard permanganate solution until the first permanent pink color
is obtained. From the volumes of standard ferrous sulfate and
permanganate used, calculate the amount of manganese present.
If a standard solution of ferrous sulfate is not available, an equivalent
amount of solid FeS04-(NH4)2S04-6H20 can be weighed out and added.
Solid sodium oxalate can be used in place of the ferrous sulfate; it reacts
more slowly, but it has the advantage that it. is not oxidized by oxygen of
the air, and therefore the carbon dioxide can be dispensed with. The excess

354 ANALYSIS OF ALUMINUM GROUP [P. 78

oxalate is titrated with permanganate without the solution being cooled
(see P. 87).

This procedure and the modifications mentioned can be used to determine
the manganese dioxide (or so-called "available oxygen") in pyrolusite,
native manganese dioxide. This is more resistant than the freshly pre-
cipitated material and has to be heated for a longer time. Proceed as
follows:

Weigh out a 0.25- to 0.5-g sample (which should have been
ground to a very fine powder in an agate mortar and dried for at
least 1 hr. at 120C.) into a 300-ml flask and treat it with ferrous
sulfate (or sodium oxalate) as directed in the procedure above.
Continue the heating until no more dark particles of pyrolusite are
present (a white silicious residue may remain).

P. 73. Precipitation of Aluminum

Discussion. After the oxalate has been destroyed, the aluminum
can be precipitated as hydroxide by neutralizing the solution with
ammonium hydroxide; the chromium, having been oxidized to
chromate, remains in solution. The quantitative precipitation of
aluminum has been carefully investigated by Blum, 2 and his results
show that the hydroxide begins to precipitate when the hydrogen
ion concentration of the solution is reduced to between 10~ 3 and
10~ 4 (depending upon the concentration of the aluminum and the
anions present) ; that it is least soluble in an approximately neutral
solution, where the precipitation is complete within usual quantita-
tive limits; and that the precipitate becomes appreciably soluble if
the hydrogen ion concentration is reduced to as low as 10~ 9 or 10~ 10 .
This last effect is due to the amphoteric nature of aluminum hydrox-
ide. By neutralization of the nitric acid with ammonium hydroxide
a high concentration of ammonium salts is provided, and by the
addition of only a small excess of ammonia the hydroxyl ion concen-
tration is controlled at a very favorable value. As the precipitate
tends to become colloidal and run through the filter, it is washed
with an ammonium nitrate solution and not with water alone.

The precipitate produced by ammonia may consist of silicic acid
or chromic hydroxide. Silica is frequently introduced during the
analysis, especially by alkaline solutions which have stood in glass
containers, or by the action of hot alkaline solutions on glass or
porcelain vessels. Any chromium not oxidized by the chlorate or
subsequently reduced would also produce a precipitate which would
make the presence of aluminum doubtful. Because of this, if one

Blum, J. Am. Chem. Soc., 38, 1291 (1916).

P. 73] PRECIPITATION OF ALUMINUM 355

is interested in the detection of small amounts of aluminum, the pre-
cipitate (if small or not of characteristic appearance) should be sub-
jected to a confirmatory test. This is done by making use of the
dyestuff ammonium aurintricarboxylate (commonly known by the
trade name of Aluminon), which forms a compound with aluminum
hydroxide in a very weakly acid solution, giving a bright red precipi-
tate which is not rapidly decomposed when the solution is made
'alkaline with an excess of ammonium carbonate. Such colored
compounds (which are frequently formed by adsorption processes)
of a dyestuff with a hydrous oxide are known as "lakes." This par-
ticular compound does not form in an alkaline medium, but, once
formed, it is stable in ammonium carbonate solution; chromium, if
present, forms a similar lake, but this is decomposed by ammonium
carbonate; alkaline earth elements give red compounds which are
decomposed by carbonate; iron forms a violet precipitate which is
changed to reddish-brown by ammonia and is thus likely to obscure
the aluminum test; phosphates in excess prevent the color with
aluminum, but after an ammonium hydroxide precipitation the
ratio of aluminum to phosphate is usually such that the test is ob-
tained; silicic acid gives a white precipitate. 8

Procedure 73: PRECIPITATION OF ALUMINUM. Dilute
the filtrate from the Mn02 precipitate to 100 ml, heat it
almost to boiling, and, while keeping the solution hot, add
NH^H slowly until the solution just changes red litmus
paper to blue. (White flocculent precipitate, presence of
aluminum. Notes 1, 2.) If an excess of NH 4 OH is added,
make the solution just acid with HNO 3 and repeat the neu-
tralization. Heat the solution for 2 to 3 min. and filter at
once through an ashless paper filter. Wash the precipitate
at once with 2-4 ten-mi portions of a hot wash solution con-
taining 2 g. of NH 4 NO3 in each 100 ml of solution. Collect
this wash water with the filtrate and treat it by P. 75 if
chromium is present. Finish washing the precipitate with
the same wash solution and treat it by P. 74.

1 The following articles give a discussion of this confirmatory test and of
its use for the colorimetric estimation of aluminum: Hammett and Sottery,
/. Am. Chem. Soc., 47, 142 (1925); Lundell and Knowles, /. Ind. Eng. Chem.,
18, 60 (1926); Yoe and Hill, J. Am. Chem. Soc., 49, 2395 (1927); Middleton,
ibid., 48, 2125 (1926); Corey and Rogers, ibid., 49, 216 (1927); Thrun, J. Phys.
Chem., 33, 997 (1929); Winter, Thrun, and Bird, /. Am. Chem. Soc., 61, 2721
(1929); Roller, ibid., 66, 2437 (1933).

356 ANALYSIS OF ALUMINUM GROUP [P. 74

Notes:

1. For the reasons mentioned in the discussion, if only a small precipitate
is obtained, the presence of aluminum should be confirmed. In that case
filter and wash the precipitate and proceed as follows:

Dissolve the precipitate by pouring repeatedly through the
filter 5 ml of warm 1.2 n. HC1. (If the precipitate is reddish
colored, indicating iron, make this solution just alkaline with 6 n.
NaOH, add 2 ml in excess, warm the mixture, filter out the precipi-
tate, make the filtrate just acid with HC1, and add 1 ml in excess.)
Add to the solution 5 ml of 3 n, NH^^HsC^) and 5 ml of a 0.1 per
cent solution of ammonium aurintricarboxylate (Aluminon).
Allow the solution to stand for 5 min., make it just alkaline with
6 n. (NHU^COa reagent, and add 5 ml in excess. (Red precipitate,
presence of aluminum.)

2. If it is important to know whether a small precipitate of aluminum is
due to the presence of this element in the original substance, a blank analysis
should be made by adding together in the same manner the reagents which
have been used to produce the aluminum hydroxide precipitate. Aluminum
is almost always introduced into an analysis in small quantities because of
the action of reagents on vessels and containers and because of impurities
in the chemicals used.

P. 74. Estimation of Aluminum

Discussion. Since no distinctive and rapid volumetric method
for the estimation of aluminum is yet available, especially when the
element may be present in greatly varying amounts, a gravimetric
method is used. This consists in simply heating the precipitate of
aluminum hydroxide (or more properly, of hydrous aluminum oxide),
which is very indefinite in its composition, at a high temperature
until it is converted to the oxide, A^Oa, which is then cooled and
weighed. Blum 4 has shown that 5 to 10 min. at the highest tem-
perature of the blast lamp is sufficient to bring most aluminum pre-
cipitates to constant weight. A somewhat longer time may be
required for large precipitates.

For a general discussion of the heating of precipitates, see p. 128.

Procedure 74: ESTIMATION OF ALUMINUM. Place the
funnel containing the paper with the precipitates in it in a
drying oven at 100C. (Notes 1, 2).

Select a clean porcelain crucible (20- to 30-ml capacity),
cover it, and slowly heat it on a clean clay or nichrome
triangle until the full heat of the blast lamp is used. Allow

4 Blum, J. Am. Chem. Soc., 38, 1282 (1916).

P. 74] ESTIMATION OF ALUMINUM 357

the crucible to cool until it is still noticeably warm when
the hand is held near it. With a pair of crucible tongs
place the crucible in a desiccator and allow it to cool to
room temperature (Note 3). When it is cool, weigh the
crucible to 0.1 mg. Take the funnel with the precipitate
from the drying oven, and carefully remove the paper from
the funnel. Fold the filter paper over the precipitate so
that none can drop out, and place it in the crucible. If any
precipitate adheres to the funnel, wipe it off with a clean
piece of filter paper and put it in the same crucible.

Replace the crucible on the triangle, inclining it, and
then very gradually heat it with the flame from a burner
until the paper is dry and begins to char; then maintain a
temperature such that the paper slowly chars and "smokes
off" without at any time ignitihg (Note 5). After the
paper has completely charred, increase the temperature to
the full heat of the burner and maintain this temperature
until most of the carbon from the filter has been burned.
Turn the tilted crucible frequently. Finally place the
crucible upright, cover it, and heat it at the full temperature
of the blast lamp for 5 to 10 min. Cool the crucible as
before and again weigh it. Repeat the heating with the
blast lamp until a constant weight (Note 4) is obtained.
From the weight of A^Oa obtained, calculate the amount of
aluminum present.

Notes :

1. If the precipitate is small (10 mg or less), an estimate of the amount of
aluminum present can be made by dissolving the precipitate in 5 ml of HC1,
collecting the solution in a test tube, neutralizing with NH 4 OH, and com-
paring the precipitate with varying amounts of aluminum precipitated in a
similar manner. A correction should be made for any aluminum found in
the blank (Note 2, P. 73).

2. Aluminum oxide is not reduced by being heated with moist paper;
therefore if a weighed crucible is ready, the moist precipitate can be trans-
ferred to the crucible and dried by direct heating. Care must be taken not
to raise the temperature so rapidly as to cause spattering of the precipitate.

3. This cooling usually requires from 15 to 30 min. Do not keep the
desiccator open any longer than absolutely necessary.

4. The complete dehydration of Al(OH)s may take considerable time.
It is judged to have been accomplished when no further change (within a
few tenths of a milligram) takes place in the weight of the precipitate.

5. Should the paper ignite, immediately extinguish it by momentarily
placing the cover on the crucible.

358 ANALYSIS OF ALUMINUM GROUP [P. 75

P. 75. Estimation of Chromium

Discussion. After the ammonium hydroxide precipitation of
aluminum, chromium (as chromate) is the only element of the
Aluminum Group remaining in the solution. Therefore, by properly
adjusting the acid concentration and adding an excess of iodide,
the chromate is quantitatively reduced to the tripositive state and
the iodine set free can be titrated with a standard thiosulfate solu-
tion. The procedure is essentially the same as that used for the
standardization of a thiosulfate solution against potassium dichro-
mate; for a discussion of the reaction and of the conditions affecting
it, see the discussion of P. XIV.

If, through faulty procedure, any aluminum or manganese were
present, they would not interfere with this estimation of chromate.
Any chlorate not decomposed by the concentrated nitric acid in
P. 71 reacts so slowly with the iodide under the conditions of the
procedure that it does not cause an appreciable error.

Procedure 75: ESTIMATION OF CHROMIUM. If in P. 71
the presence of chromium has been indicated, evaporate
the filtrate (from P. 73) to 70 to 80 ml and cool it.

If the amount of chromium present is thought to be less
than about 50 mg (Note 1), transfer the cold solution to a
500-ml flask and treat it as directed in the last paragraph of
this procedure.

If the amount of chromium present is thought to be more
than 50 mg, transfer the cold solution to a 100-ml volumetric
flask, dilute it to the mark, and thoroughly mix the solu-
tion. Pipet into a 500-ml flask a volutne of solution
thought to contain about 50 mg of chromium, dilute it to
80 ml, and treat it as directed in the next paragraph.

Add EUSC^ slowly, 1 ml at a time, until the solution
changes from yellow to orange or quickly turns blue litmus
red, and then add 6 ml more of K^SOi and 2 g of KI. Close
the flask with a clean rubber stopper, gently swirl the mix-
ture until the KI is dissolved, and let it stand for 5 min.,
preferably in the dark. Add to the solution 200 ml of
cold water (washing down the sides of the flask) and titrate
it with standard 0.1 n. Na2S2Os until the brownish-yellow
color of the iodine becomes indistinct (Note 2). Swirl the
solution slowly as the thiosulfate is added. Add to the
solution 5 ml of a starch indicator solution, and again titrate

P. 75] ESTIMATION OF CHROMIUM 359

until the blue starch color disappears (Note 2). From the
volume of standard Na 2 S 2 3 used, calculate the amount of
chromium present.

Notes :

1. The amount of chromium present can be judged by adding a standard
chromate solution to a solution of the same volume containing 1 g of NHiNOs
and 1 drop of NEUOH until the intensity of the yellow color in the two solu-
tions is the same.

2. It should be remembered that the solution does not become entirely
colorless, as the tripositive chromium gives it a greenish color.

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360

The Analysis of the Alkaline Earth Group

P. 81. Precipitation of the Alkaline Earth Group

Discussion. Two general methods are used in systems of qualita-
tive analysis for the precipitation of the elements known as the
Alkaline Earth Group. In the most generally used of these, the
filtrate from the Ammonium Sulfide Group precipitation is con-
centrated by evaporation and then ammonium carbonate is added.
The large amount of ammonium chloride present, which has accu-
mulated during the course of the analysis, tends to prevent the
precipitation of magnesium; thus only barium, strontium, and cal-
cium are included in the group. In the alternative method, which
is used here, magnesium is caused to precipitate with the other
elements of the group (1) by the removal of the ammonium chloride
before the precipitation, (2) by using a high concentration of am-
monium carbonate and ammonium hydroxide (the latter being added
to repress the hydrolysis of the carbonate) and (3) by the addition
of alcohol, as most inorganic salts, especially if they are ionized, are
much less soluble in alcohol than in water. This method has been
adopted because the experiments of Bray 1 have indicated that it is
impossible to find conditions under which barium, strontium, and
calcium can be completely precipitated with ammonium carbonate
without causing partial precipitation of magnesium when it is present
in large amounts. It is also true that the precipitation of the first
three elements is much more rapid and complete under the conditions
which cause complete precipitation of magnesium. Bray (loc. cit.)
has also shown that, with conditions similar to those of this procedure,
0.5 mg of any element of the group will give a distinct turbidity in
15 min.; that with 500 mg of any element not more than 0.2 mg
remained in the filtrate when the precipitates were filtered after
30 min.; that the precipitates became crystalline more rapidly and
were more easily filtered if the mixtures were shaken frequently;
and that, if the alcohol was omitted, no precipitates were obtained
when 0.5 to 1 mg of calcium or magnesium was present. Gooch
and Eddy 2 have shown that magnesium can be quantitatively pre-
cipitated by an ammonium carbonate reagent in the presence of

1 Bray, J. Am. Chem. Soc., 31, 611 (1909).

2 Gooch and Eddy, Am. J. Sci. (4), XXV, 444.

361

362 ALKALINE EARTH GROUP [P. 81

alcohol and that neither sodium nor potassium are carried with the
precipitate in significant amounts.

The Removal of Ammonium Salts. The ammonium chloride is
removed from the solution, in the absence of perchlorate, by evapora-
tion with concentrated nitric acid; in this process the ammonia is
oxidized to nitrogen by the products of the reaction between the
nitric acid and the chloride. This method is more convenient for
removing ammonium salts and less liable to loss from spattering than
the more familiar method of heating to a sufficiently high temper-
ature to voltilize them. If perchlorate is present (from the use of
perchloric acid in the preparation of the solution), it must be re-
moved or destroyed in order to prevent the possible precipitation of
the potassium as perchlorate with this group. It has been found
by Noyes and Bray 8 that the perchlorate is destroyed by simple
ignition because of its reaction with the ammonium salts present.
Neither the ammonium chloride nor the perchlorate has to be
quantitatively removed at this point, as a small amount of the first
would not prevent the precipitation of magnesium, nor would a trace
of perchlorate cause the precipitation of potassium; large amounts
would cause these effects.

Either of these methods of removing the ammonium salts produce
a vigorous oxidizing action, which results in the conversion of any
sulfur or sulfur compounds present (various thionic acids, usually
introduced into the analysis during the Hydrogen Sulfide and Am-
monium Sulfide Group precipitations) into sulfate. This sulfate may
cause the partial or, if they are present in small amounts, complete
precipitation of barium and strontium; if it is present in large
amounts, calcium may be partially precipitated. As the residue
from the treatment to destroy ammonium salts is normally treated
with water and any insoluble material is discarded, these elements
may be partially, or even completely, lost.

To avoid this possibility, a treatment for the recovery of these
elements is provided in Note 3 of the procedure. This consists in
treating any insoluble residue with a solution containing sodium
carbonate and a small amount of sodium hydroxide; the latter is
added to repress the hydrolysis of the carbonate. It is intended that
this treatment metathesize the sulfates into carbonates, as repre-
sented by the equation (for BaSO 4 )

BaS0 4( .) + C0 8 "" = BaC0 3( .) + S0 4 ".

1 Noyes and Bray, Qualitative Analysis for the Rare Elements, Macmillan,
1927, p. 236.

P. 81] METATHESIS OF SULFATES 363

The carbonate precipitate is then combined with the subsequent
ammonium carbonate precipitate of this group and dissolved in
nitric acid.

The Metathesis of Slightly Soluble Compounds. Upon reference to
solubility data, it will be found that barium carbonate is considerably
more soluble than barium sulfate; therefore it is appropriate to ques-
tion to what extent the desired metathetical reaction will proceed.
This can be approximately predicted as follows :

As soon as a precipitate of barium carbonate is formed, the solu-
bility-product principle for both barium sulfate and barium carbonate
can be applied (assuming that these precipitates are in equilibrium
with the solution). Therefore the two solubility-product expressions
can be written as follows :

(1)
(2)

As [Ba 4 "*"] is common to both and has the same value (the expres-
sions apply to the same solution), it can be eliminated and the two
expressions can be combined to obtain the following :

[S04 ]

[00,1

This expression shows that, in any solution which is saturated with
both barium sulfate and barium carbonate, the ratio of the con-
centration of sulfate ion to that of carbonate ion will be determined
by the ratio of the solubility product of barium sulfate to that of
barium carbonate. For the case being considered, this ratio has
the value

X 1(T 10 =

U ' U1 '

[con #Baco, 7 x 10-'

Therefore (assuming the conditions obtaining in the proposed treat-
ment), if a solution 1.5 m. in carbonate is added to a BaSOi precipi-
tate, metathesis will take place, the carbonate concentration decreas-
ing and the sulfate concentration increasing until the above ratio is
attained or the supply of BaSO 4 is exhausted. If it is assumed that
an excess of BaS0 4 is present and if x represents the equilibrium
sulfate concentration, then the carbonate concentration will be
1.5-x. 4 Solving for x, we find the concentration of the sulfate to be

4 This assumes (as is justified in this case because of the very slight solu-
bility of these precipitates) that for each mole of SO 4 " appearing in the solu-
tion a mole of COi" is precipitated.

364 ALKALINE EARTH GROUP [P. 81

0.021 m., which indicates that, with 25 ml of 3 n. sodium carbonate,
approximately 0.12 g of barium sulfate should be metathesized. 1

It is seen from these calculations that the amount of a given
precipitate which can be metathesized into some other compound
by such a process is determined by three factors: (1) the relative
solubilities of the two compounds, (2) the concentration of the
metathesizing ion, and (3) the volume of the metathesizing solution.
It is possible, as was calculated above for the case of barium sulfate,
and as is experimentally true, to cause the complete metathesis of a
limited amount of a precipitate into a more soluble compound.
However, as solubility data would indicate, a much larger amount
of strontium or calcium sulfate is metathesized under the same
conditions.

It is to be noted that, in applying the above principles to the
metathesis of a precipitate into another of a different ionic type,
the simple ratio of ion concentrations to solubility products does
not hold; thus, for the metathesis of silver chloride into silver car-
bonate, which can be represented by the equation

2AgCl (8) + C0 3 " = Ag 2 C03 ( .) + 2CP,
the expression

ten 1 _ (K AlCl r

[C0 3 "] ^AtsCOa

is obtained. However, the same principles apply, and calculations
similar to the one above can be made.

It will also be seen, particularly in the analysis of the Alkaline
Earth Group, that the completeness with which two ions may be
separated by a common precipitant can be predicted from the proper
treatment of the solubility-product principle.

Procedure 81: PRECIPITATION OF THE ALKALINE EARTH
GROUP. Make the filtrate from the Ammonium Sulfide
Group precipitation acid with HC1 and evaporate it in a
large flask until salts begin to crystallize or the mixture
begins to bump (Note 1).

If HC104 has not been used in the preparation of the solu-
tion, add to the mixture 5 to 10 ml of 16 n. HN0 3 and warm
it gently as long as a vigorous evolution of gas takes place.

It is to be emphasized that this figure should be considered as only a very
approximate indication of the experimental facts. In these relatively con-
centrated solutions the activities of the ions are undoubtedly quite different
from their concentrations.

P. 81] GROUP PRECIPITATION 365

Finally, evaporate the solution almost to dryness, keeping
the mixture in constant motion or allowing the evaporation
to take place on a steam or sand bath. Add 5 ml more of
16 n. HN0 8 , washing the sides of the vessel, evaporate the
mixture just to dryness, and treat it by the last paragraph
of this procedure.

If HC1O4 has been used in the preparation of the solu-
tion, transfer the mixture to a capacious casserole, add 5 ml
of 16 n. HNO 3 (Note 2), and evaporate it to dryness over a
steam bath. Remove the casserole from the steam bath
and keep a flame from a burner continuously moving over
the surface of the residue until no more burning appears to
take place and no more salts appear to be volatilized. Allow
the dish to cool, moisten all parts of the residue with 1 to
2 ml of 3 n. NKUCl, and again heat as directed above.
Finally, heat the dish, keeping it in constant motion, until
no more fuming of the salts occurs. Do not heat the dish
so intensely that the porcelain at any time shows a reddish
glow.

Add to the residue 1 ml of HNO 3 and 10 ml of water, heat
the solution just to boiling, and transfer it with the aid of
5 ml of water to a 200-ml flask (Note 3). Add slowly 20 ml
of (NH 4 )*C0 3 reagent (Note 4) and 20 ml of 95 per cent
C 2 H 6 OH and shake the solution frequently for 15 min.; or,
preferably, let it stand for several hours. (White precipi-
tate, presence of the Alkaline Earth Group elements.) If
a large precipitate forms, add to the mixture 10 ml more of
each of the last two reagents. Allow the precipitate to
settle, and decant the solution through a paper filter.
Wash the precipitate by decantation with three to five
5-ml portions of a solution containing 2 volumes of the
(NH^COs reagent to 1 volume of alcohol. Collect these
washings with the filtrate and treat it by P. 91. Treat
the precipitate by P. 82.

Notes:

1. If, on evaporation of the solution, there forms a precipitate of sulfur,
or of traces of the sulfides of the previous groups, these should be filtered
out and discarded in order to minimize the formation of sulfate in the subse-
quent treatment with HN0 3 .

2. The HNOs is added to prevent the possibility of an explosion on
evaporating the perchloric acid mixture (see the discussion of P. 4).

3. If there is a white residue after the solution is boiled, it probably con-

366 ALKALINE EARTH GROUP IP. 81

sists of the sulfates of barium, strontium, or calcium. In this case proceed
as follows:

Filter the mixture through a small filter, wash the residue, and
transfer it to a casserole with 10 to 25 ml of 3 n. Na2COa and 1 ml
of 6 n. NaOH. Boil the mixture for 5 min. Reserve the mixture
until the ammonium carbonate precipitate of the Alkaline Earth
Group elements (last paragraph of the procedure above) has been
washed, and then filter it through the same filter which has been
used to collect the ammonium carbonate precipitate. Wash the
precipitate with water and discard the filtrate. (Do not add this
NaaCOs- NaOH filtrate, or the washings, to the filtrate from the
ammonium carbonate precipitate, as the latter is to be treated by
P. 91 and tested for the Alkali Group elements.)

Any sulfur or sulfur compounds introduced during the treatments with
hydrogen or ammonium sulfide will be oxidized to sulfate during the evapo-
ration with nitric acid and will precipitate the sulfates of the elements men-
tioned. By this treatment these small sulfate precipitates are converted
into carbonates (see the discussion above) and can be united with the am-
monium carbonate precipitate.

4. See the Appendix for the preparation of this reagent, which is 3 f . in
(NH 4 )aC0 8 and 6 f . in NH 4 OH.

THE SEPARATION OF THE ALKALINE EARTH GROUP

ELEMENTS

General discussion. The elements included in the Alkaline Earth
Group are all found in the second group of the periodic table; cal-
cium, strontium, and barium appear in the same sub-group, and, as
would be predicted, the compounds of these elements do not vary
widely in their properties. They exhibit only one stable valence
state; they are distinctly basic and show no amphoteric reactions;
there is relatively little tendency toward complex ion formation;
and the solubilities of their salts, on which their analytical detection
and separation depend, vary gradually on passing from one member
to the next of the series.

Although various separations were investigated, the method used
for the separation of these elements is essentially that developed by
Bray 6 for their qualitative separation. Briefly stated, this method
consists in precipitating first the barium as chromate from a weakly
acid solution, then strontium as chromate by making the solution
alkaline with ammonia and adding alcohol, next calcium as the
oxalate, and finally magnesium as magnesium ammonium ar senate.
Because of the relatively small solubility differences existing between
the compounds of the elements of this group, it has been necessary
in precipitating these elements to adopt the expedient of adding at

Bray, /. Am. Chem. Soc., 31, 611 (1909).

P. 82]

PRECIPITATION OF BARIUM

367

first only sufficient precipitant to give a sensitive test for the element
being precipitated. If the amount needed to precipitate completely
a large quantity were added at first, it would tend to precipitate the
next least soluble of the remaining elements. If a large precipitate
is obtained with the first addition, then sufficient precipitant is added
to cause complete precipitation; as a definite amount of sample is
taken for analysis, the amount of any two elements which can be
present is limited.

Ow T ing to the small difference in the solubility of their salts, and
probably to their similar crystal structure, coprecipitation is found
to occur to a marked extent. Various efforts are made to minimize
this, such as adding the precipitant very slowly while vigorously
shaking the solution; these conditions also favor the formation of
precipitates which can be more easily filtered. In spite of these
efforts, experiments have shown that, when large amounts of two
elements adjacent in the periodic table are present, a considerable
fraction of the soluble one may be found in the precipitate. A
second expedient is therefore provided; this consists in filtering the
solution by decantation, dissolving the main portion of the precipi-
tate, and reprecipitating it in a manner involving a relatively small
expenditure of time. When a more exact separation of the elements
of this group is desired, this process can be utilized.

The results of confirmatory analyses of this group are shown in
Table XXVIII. These analyses were begun with the ammonium
carbonate precipitation of P. 81.

TABLE XXVIII
TEST ANALYSES OF THE ALKALINE EARTH GROUP ELEMENTS

Amount of Each Element Taken and Found (mg)

Analysis
No.

Barium

Strontium

Calcium

Magnesium

Taken

Found

Taken

Found

Taken

Found

. Taken

Found

125

127

119

116

124

127

120

121

297

293

240

243

250

251

440

432

240

243

49.6

37.5-

37.6

0.3

37.5

70.6

37.2

26.2

Trace

26 6

94 6

Note: Analyses 4 to 8, inclusive, were carried out by a selected group of
students in a sophomore course in analytical chemistry at the California
Institute of Technology. They had had no previous experience with the
analysis of this group.

368 ALKALINE EARTH GROUP [P. 82

P. 82. Precipitation of Barium

Discussion. The separation of barium is based upon the fact
that, of the elements of this group, barium forms the least soluble
chromate. Experiments have shown that, when a limited amount
of chromate is added to an acid solution in which the hydrogen ion
is closely controlled by means of a carefully regulated ratio of acetic
acid to acetate, 0.5 mg of barium will give a perceptible precipitate,
while 500 mg of strontium, which forms the next least soluble
chromate, remain in solution. If a considerable precipitate is ob-
tained, an additional amount of chromate is then added in order to
precipitate the barium completely. Experiments 7 have shown that,
when 250 mg each of barium and strontium are present, from 5 to
10 mg of the strontium will be found with the precipitate. Because
of this, the precipitant is added, and the mixture is heated, shaken,
and treated, in such a manner that the solution can be filtered by
decantation and the precipitate can be quite completely retained in
the flask. The precipitate is then dissolved in a minimum amount
of hydrochloric acid and reprecipitated, and the filtration is carried
out through the original filter; the entire process takes only a few
minutes.

The confirmatory experiments of Bray 8 have shown that, if the
chromate is added all at once and if 500 mg of barium are present,
as much as 3 mg of strontium may be so completely carried out with
the precipitate as to escape subsequent detection in the filtrate, but
that, if the chromate is added slowly, 1 mg of strontium can be
readily detected.

The effect of acid on the solubility of barium (and other chromates)
is due, first, to the fact that chromic acid is not a highly ionized
acid. The second hydrogen ion is ionized to about the same extent
as is the first hydrogen ion of carbonic acid or the second one of
phosphoric acid; the equilibrium constant for the reaction HCrOi"
= CrOr + H 4 " has been estimated 9 to be 3.2 X 10~ 7 . In addition
to this, there is an equilibrium between hydrochromic acid and
dichromate ion, as follows:

2HCrO 4 - = H 2 O + Cr 2 O 7 ~.
The equilibrium constant for this reaction is 43. Therefore, as

7 The author is indebted to Dr. Carter Gregory for an extensive experi-
mental investigation of tho methods for the analysis of this group.

8 Bray, loc. cit.

9 Neuss and Rieman, J. Am. Ckem. Soc., 66, 2238 (1934).

P. 82] PRECIPITATION OF BARIUM 369

there is usually only a small concentration of hydrochromate ion in
solutions, the predominant equilibrium is between the chromate and
the dichromate ions, as follows:

2Cr0 4 " + 2H+ = Cr 2 7 ~ + H 2 0.

The equilibrium constant for this reaction can be calculated from
those given above to be 4.2 x 10 14 . It is seen that the hydrogen ion
enters into this equilibrium as the square of its concentration and
that, by proper adjustment, the chromate ion concentration can be
controlled quite closely and can be varied through a wide range.
The dichromates are, in general, soluble.

Procedure 82: PRECIPITATION OF BARIUM. Transfer
most of the (NH 4 ) 2 CO 3 precipitate (from P. 81) from the
filter to the flask with the aid of a stirring rod (Note 1).
Dissolve the precipitate left on the filter by pouring drop-
wise through it a 5- to 25-ml portion of cold HN0 3 (collect-
ing the acid in the flask with the precipitate) and wash
the filter with 3 to 5 ml of water added dropwise. Add
15 n. NH 4 OH until the solution is alkaline and then HNOs
(dropwise) until it is again acid, and evaporate or dilute it
until the volume is 35 ml. Add to the solution just 1 ml
of HC 2 H 3 2 and 10 ml of neutral (Note 2) 3 n. NH 4 C 2 H 3 2 ,
heat nearly to boiling, and add, 2 drops at a time and shak-
ing after each addition 3 ml of 3 n. K 2 CrO 4 (Notes 3 and 4).
Heat the solution nearly to boiling for 2 min., swirling it
continuously. (Yellow precipitate, presence of barium.)
If a considerable precipitate forms, add, in the same manner
as above, 2 ml more of K 2 CrO4 and again heat the solution
for 2 min. Allow the precipitate to settle completely and
decant the solution through an asbestos filter.

If there is only a small precipitate (50 mg or less) , wash it
with two to three 5-ml portions of hot water, collecting the
wash water with the filtrate. Treat the filtrate by P. 84.
Wash the precipitate and filter with hot water, added drop-
wise, until the wash water is no longer colored (Note 5).
Discard these washings. Treat the precipitate by P. 83.

If there is a large precipitate and if the presence of stron-
tium is suspected (Note 6), dissolve the precipitate remain-
ing in the flask with 2 to 5 ml of cold HC1 and 5 ml of
water (Note 7). Add 1 ml of K 2 Cr0 4 and then NH 4 OH,
a few drops at a time, until the color changes from orange

370 ALKALINE EARTH GROUP [P. 82

to light yellow, and 2 ml more. Then add HC 2 H 3 02 until
the original color returns; heat the mixture almost to boil-
ing for 2 to 5 min., shaking it frequently; allow the precipi-
tate to settle and wash it as directed in the preceding
paragraph. Treat the precipitate and filtrate as directed
there.

Notes:

1. If the precipitate were treated on the filter with HNOa, considerable
loss by spattering might result from the evolution of C02.

2. The NH 4 C2H 3 02 of commerce frequently contains HC2Ha02 in con-
siderable amounts; for this reason the reagent should be tested and, if acid,
made neutral by addition of NtUOH until litmus is turned neither red nor
blue but remains an indeterminate purplish color. Neutral 3 n. NH^HaC^
can also be made by mixing equal volumes of exactly 6 n. HC2H 3 O2 and
6 n. NH 4 OH.

3. It is essential that the precipitation of the BaCrO 4 be made very
slowly; otherwise it separates in such a finely divided form that it will not
settle, may even pass through the filter, and will tend to carry down
strontium.

4. The solution should be absolutely clear before the chromate is added,
as it is difficult to detect the small yellow turbidity caused by 1 mg or less
of barium. If the detection of these amounts of barium is of importance,
the solution should be filtered after the chromate is added, even when it is
apparently clear, and the filter should be washed and then examined.

5. This washing should be done with the smallest possible volume of
water; however, the estimation of barium in the next procedure requires
that the chromate be removed so completely that its color cannot be seen
when 2 to 3 ml of the wash solution is collected in a test tube.

6. This next operation is carried out for the purpose of removing the
strontium coprecipitated with the BaCr04. This amount will be small
unless considerable amounts of both barium and strontium are present, and
even then it should not exceed 10 mg. Therefore, this step should be
omitted if the size of the (NEU^COa precipitate (P. 81) or a knowledge of
the substance being analyzed indicates that strontium is unlikely to be
present in considerable amount, or if a precise estimate of the two elements
is not desired.

7. If a very small fraction of the precipitate has been carried onto the
filter, only that in the flask need be dissolved. If a considerable portion of
it is on the filter, it should be dissolved by pouring the HC1 dropwise through
it and then washing the filter with the water to be added.

P. 83 . Estimation of Barium

Discussion. In the method used here for the estimation of ba-
rium, the precipitate of barium chromate (obtained and washed free
from excess of chromate in P. 82) is first dissolved by means of cold

P. 83] ESTIMATION OF BARIUM 371

hydrochloric acid. (The effect of hydrogen ion on the solubility of
the chromates has been pointed out in the discussion of P. 82.)
The dichromate in the acid solution is then determined iodomet-
rically, an excess of potassium iodide being added and the liberated
iodine being titrated with standard thiosulfate. In an experimental
study 10 of this process, results exact to 0.3 per cent have been ob-
tained with amounts of barium varying from 50 to 350 mg. Hydro-
chloric acid is used because it very readily dissolves the precipitate
from the filter and because it was found that barium chromate was
not completely metathesized by sulfuric acid, even upon the addi-
tion of the iodide. This incomplete metathesis was probably due to
mechanical inclusion of the chromate by the precipitated barium
sulfate. It was found that, if the barium chromate was precipitated
in P. 82 by rapid addition of the chromate, the results were likely to
be from 1 to 1.5 per cent too high, probably because chromate is
carried down by the precipitate. For a discussion of the conditions
under which the dichromate-iodide reaction should be carried out,
see the references cited in the discussion of P. XIV, "The Standardi-
zation of a Thiosulfate Solution Against Potassium Dichromate. "

Procedure 83: ESTIMATION OF BARIUM. Dissolve the
precipitate on the filter by pouring dropwise through it just
10 ml of HC1, collecting the solution in a 600-ml beaker,
and washing the filter with just 40 ml of water (Note 1).
Transfer the precipitate from the flask to the same beaker
with the aid of just 50 ml of water. Break up the precipi-
tate, stir it until it dissolves, add 2 to 3 g of solid KI, and
gently swirl the mixture until the KI is dissolved. Cover
the solution with a watch glass and allow it to stand in a dark
place for 5 min.

Dilute the solution to 400 ml, titrate it with 0.1 n. Na 2 S2O 3
solution (Note 2) until the color of the iodine becomes indis-
tinct, add 5 ml of a starch indicator solution, and again
titrate the solution until the blue starch color changes to a
light green. From the volume of standard Na2S2O 3 used,
calculate the amount of barium present.

Notes :

1. If a very small precipitate (1 to 2 mg) has been obtained, it should be
dissolved in 2 to 5 ml of HC1, treated by the last paragraph of P. 82, and
then compared with known amounts of barium precipitated under similar

10 Unpublished experiments by K. H. Lau.

372 ALKALINE EARTH GROUP [P. 84

conditions; or, if desired, the presence of barium can be confirmed by evapo-
rating the HC1 solution to 0.1 to 0.2 ml, dipping a platinum wire (one which
has been cleaned and then held in a flame until it no longer causes a color)
into it, and then holding the wire in a small colorless flame. A green color
in the flame indicates the presence of barium.
2. For notes on this titration, see P. XIV.

P. 84. Precipitation of Strontium

Discussion. The precipitation of strontium is the next step in
the analysis of this group. Strontium is also precipitated as the
chromate. This is accomplished by the addition of an excess of
ammonia to the acetic acid filtrate from the barium precipitation
and by the addition of alcohol to the solution. As was pointed out
in the discussion of P. 82, the chromate ion concentration is tre-
mendously increased by decreasing the hydrogen ion concentration
of the solution. Also, by changing from a solvent medium of water
to one containing a considerable proportion of alcohol, the solu-
bility of the strontium chromate is further reduced. As was ex-
plained in P. 82, only such amounts of these reagents are first added
as will give a sensitive test for small amounts of strontium; larger
amounts are added if they are needed to precipitate large quantities
of strontium. Even if the larger amounts of alcohol and chromate
are added, calcium chromate will not precipitate if calcium is present
alone; however, it tends to be coprecipitated with the strontium
chromate. For that reason, if a more complete separation of these
elements is desired, a double precipitation should be made when
considerable strontium is found and calcium is also likely to be
present. Thus, it was found that, with 250 mg of each element,
about 10 mg of calcium were coprecipitated; no calcium could be
detected with the strontium after making a reprecipitation.

Procedure 84: PRECIPITATION OF STRONTIUM. Add 15 n.
NH 4 OH to the filtrate from the BaCr0 4 precipitation
(which should have a volume of not over 60 ml) until the
color changes from orange to yellow, and then add 5 ml
more. Add to the solution 3 ml of 3 n. K 2 Cr0 4 , cool, and
then add, 2 ml at a time and shaking the mixture vigorously
after each addition, 30 ml of 95 per cent C2H50H (ethyl
alcohol). If it is thought that more than 50 to 100 mg of
strontium are present, add 3 ml more of 3 n. K 2 Cr0 4 and, as
directed above, 10 ml more of CiHtOH. Let the mixture
stand at least 10 min (Note 1). (Light yellow precipi-

P. 85] ESTIMATION OF STRONTIUM. 373

tate, presence of strontium.) Let the precipitate settle and
decant the solution.

If there is only a small precipitate (50 mg or less), wash
it with 2 portions of a solution made by adding 1 ml of 15 n.
NH 4 OH to 10 ml of 50 per cent alcohol. Collect the wash-
ings with the filtrate. Treat the filtrate by P. 86. Treat
the precipitate by P. 85.

If there is a large precipitate and calcium may be present
(Note 2), do not wash it but completely drain the solution
from it. Dissolve the precipitate in the flask in 5 'to 10 ml
of warm HC2H 3 02, neutralize the solution with 15 n. NH 4 OH
as directed above, add 2 ml more, and dilute it to 30 ml.
Add 2 ml of 3 n. K 2 Cr0 4 and, as directed above, 20 ml of
95 per cent alcohol, and then let the mixture stand for 10
min. Allow the precipitate to settle and decant the solu-
tion through an asbestos filter, draining all the solution
from the precipitate. Add this solution to the original
filtrate and treat it by P. 86. Treat the precipitate by
P. 85.

Notes:

1. If other operations can be carried out in the meantime, it is advisable
to let this precipitate stand for 30 min. to 1 hr., as it is then much more
readily filtered. If the flask is tilted during this time (by inclining it in a
large casserole or clamping it), the precipitate so settles that it is not stirred
up when the solution is decanted.

2. In the majority of cases, this procedure to remove any coprecipitated
calcium may be omitted; the considerations discussed in Note 6, P. 82, also
apply here.

P. 85. Estimation of Strontium

Discussion. It would seem that the strontium could be estimated
by an iodometric process identical with that used with the barium
chromate precipitate; however, washing the rather soluble stron-
tium chromate precipitate free of chromate without appreciably
dissolving it was found to be extremely difficult. The method
which was adopted depends upon converting the strontium chro-
mate to oxalate, washing this precipitate free from the excess of
oxalate, dissolving it in hydrochloric acid, and then titrating the
oxalate obtained with standard permanganate solution. The
strontium chromate precipitate is readily dissolved in acetic acid,
and, because strontium oxalate is less soluble than the chromate,

374 ALKALINE EARTH GROUP [P. 85

and also because of the presence of an added excess of oxalate,
strontium oxalate precipitates as the solution is neutralized. Hydro-
chloric instead of sulfuric acid is used for dissolving the precipitate
of strontium oxalate, since it avoids the formation of strontium
sulfate, which tends to enclose particles of undissolved oxalate.

The titration of oxalate with permanganate, which has been used
in P. XIV for the standardization of a permanganate solution against
sodium oxalate, is preferably carried out in the presence of sulfuric
acid; when hydrochloric acid is used, there is a pronounced tendency
for chloride to be oxidized by the permanganate. However; studies
by Gooch and Peters, 11 by Baxter and Zanetti, 12 and by Kolthoff 18
have established that, by titrating in the presence of a relatively
high concentration of manganous ion or at an elevated temperature
(70 to 80C.), the oxalate-permanganate reaction can be made to
proceed quantitatively and without significant reduction of the
permanganate by the hydrochloric acid.

Procedure 85: ESTIMATION OF STRONTIUM. Dissolve the
precipitate on the filter (Note 1) by pouring through it 3 to
5 ml of HC 2 H 3 2 and then wash the filter with 10 to 20 ml
of water, collecting the solutions in the flask with the re-
mainder of the precipitate. Warm the mixture until the
precipitate dissolves, dilute it to 50 ml, and heat it to
boiling.

Add to the solution, a few drops at a time and shaking
after each addition, 3 to 5 ml of 3 n. K 2 C 2 04; slowly add
NH 4 OH until the solution is alkaline and then add 5 ml
more (Note 2). Heat the mixture to boiling for 3 to 5 min.,
cool it with tap water, shake it vigorously, and let it stand
until the precipitate settles. Decant the solution through
a paper filter and wash the precipitate by decantation with
5- to 10-ml portions of hot 0.6 n. NH 4 OH until no precipitate
forms when 1 drop of 1 n. Ca(NOa)2 is added to a 5-ml por-
tion of the wash water.

Dissolve the precipitate in the flask with 50 ml of warm
(40 to 50C.) 3 n. HC1 and transfer the solution to a 400-
to 600-ml beaker (Note 3). Remove the paper with the
precipitate from the funnel, open it against the side of the

" Gooch and Peters, Z. anorg. allgem. Chem., 21, 185 (1899).
11 Baxter and Zanetti, Am. Chem. Jour., 33, 500 (1905).
i Kolthoff, Z. anal. Chem. t W, 204 (1924).

P.86] PRECIPITATION OF CALCIUM 375

beaker, and then wash the precipitate from the paper with
a stream of water from a wash bottle until the total volume
of the solution is 100 to 125 ml. Heat the mixture to 70
to 80C., add to it 5 ml of the Zimmermann-Reinhardt
"preventative" solution (see P. 53-4), and titrate it slowly
with 0.1 n. KMnO 4 , stirring the solution constantly, until
the first pink color is obtained. Shove the filter paper
from the side of the beaker into the solution and again
titrate until a pink color is obtained which remains for 10
sec. (Note 2, P. 87). From the volume of standard per-
manganate used, calculate the amount of strontium present.

Notes :

1. If only a small precipitate is obtained, it may be due to a small amount
of barium which had not been precipitated in P. 82. In order to confirm
the presence of strontium, proceed as follows:

Pour repeatedly through the filter a hot solution made by adding
10 drops of HC 2 H 8 02 and 1 ml of 1 n. K 2 Cr0 4 to 5 ml of 3 n.
NH 4 C2H 3 O2. Add 15 n. NH 4 OH until the orange color changes
to yellow and 1 ml more ; then add 5 ml of alcohol. (Yellow precipi-
tate, presence of strontium.)

Barium chromate would not be dissolved by the solution which is poured
through the filter; calcium chromate would be dissolved, but the amount
that might be precipitated in P. 84 would not be likely to cause a precipitate
here. Even 0.3 mg of strontium will give an easily perceptible turbidity
upon addition of the alcohol. The amount present can be estimated by
comparison with precipitates produced by known amounts of strontium.

Strontium gives a red color in a flame. An additional test can be made
in the same manner as was suggested for the barium precipitate.

2. By adding the reagents in this manner a coarsely crystalline precipi-
tate is obtained which is very readily washed by decantation; strontium
oxalate precipitated from an alkaline solution by rapid addition of oxalate
is likely to be difficult to filter.

3. If a large precipitate has been obtained, it is recommended that the
HC1 solution be poured repeatedly through the precipitate, cooled, and
transferred to a 100-ml volumetric flask, the filter be washed with cold 3 n.
HC1, and the solution be collected in the volumetric flask until the calibra-
tion is reached. After the contents are thoroughly mixed, portions may be
pipeted out and duplicate titrations may be made.

P. 86. Precipitation of Calcium

Discussion. This procedure is based upon the conventional
quantitative separation of calcium from magnesium by precipitation
of calcium oxalate monohydrate, CaC 2 O 4 -H 2 0, from an alkaline
solution.

376 ALKALINE EARTH GROUP [P. 86

When precipitated from a cold neutral or alkaline solution, cal-
cium oxalate is so finely divided as to be difficult to filter and wash.
Therefore in quantitative work the oxalate is usually added to a
warm acid solution and ammonia is then slowly added in excess;
by this means coarse, granular, easily filtered crystals can be ob-
tained. This method cannot be used here, as in an acid solution
the chromate previously added would be at least partially reduced
by the oxalate and alcohol present. Therefore in the procedure
below the oxalate is added dropwise to a hot solution; the precipi-
tates thus produced usually can be filtered without undue difficulty.
Potassium oxalate is used as the precipitant, as it is more soluble,
and less coprecipitated, than either the ammonium or the sodium
salt.

Magnesium oxalate is only moderately soluble in water (0.3009 g
of MgC 2 04-2H 2 per liter at 18C.); therefore it would appear to be
impossible to effect a separation of calcium from considerable
amounts of magnesium by this procedure. That such a separation
is possible is due to two specific properties of magnesium oxalate
solutions. First, magnesium oxalate forms unusually stable super-
saturated solutions in which the normal solubility can be greatly
exceeded. W. M. Fischer 14 has shown that under certain conditions
a solution of magnesium oxalate will remain supersaturated to 60
times its normal solubility for \ hr. ; to 40 times for 2 hr. ; to 25 times
for 4 hr.; to 10 times for 40 hr.; and to 4 times for greater than 10
weeks. The time for which the solution will remain supersaturated
is decreased by the presence of other ions, by scratching of the con-
taining vessel, by mechanical agitation, and by increase in tempera-
ture, the latter effect being quite pronounced above 80C. Second,
the solubility of magnesium oxalate is increased, rather than de-
creased, by an excess of oxalate, being about seven times as great
in a solution approximately 0.35 f. in ammonium oxalate as in pure
water. 16

It has long been known that calcium oxalate was more soluble
in solutions of magnesium salts than in water, and the effect has
been attributed to the formation of either un-ionized magnesium
oxalate or to a complex ion; the data cited above, showing a marked

14 Fischer, Z. anorg. allgem. Chem., 163, 62 (1926).

11 Bobtelsky and Malkawa-Janowskaja, Z. angew. Chem., 40, 1436 (1927),
present extensive data showing that the solubility of magnesium oxalate
becomes greater in solutions of ammonium oxalate, oxalic acid, and am*
monium chloride.

P. 86] PRECIPITATION OF CALCIUM 377

increase in solubility with an excess of oxalate, indicate complex ion
formation. Because of this it is necessary to add sufficient oxalate,
not only to precipitate the calcium, but to combine with the mag-
nesium and convert it into the complex ion. Too large an excess
should not be added here, as it has been shown by Hall 18 that the
subsequent precipitation of magnesium may be thereby prevented.
A considerable amount of ammonium chloride has also been found
to aid in preventing the precipitation of magnesium oxalate with
the calcium oxalate. 17

Procedure 86: PRECIPITATION OF CALCIUM. To the
filtrate from P. 84 add 5 g of solid NH 4 C1, heat the solution
to 70 to 80C. ; and add to it, dropwise and swirling the
mixture continuously, 3 to 12 ml of 3 n. K 2 C204 (Notes 1,2).
Let the solution stand for 10 min., shaking it frequently
(Note 3). (White precipitate, presence of calcium.)

If a small precipitate forms, let it settle. Decant the
solution through a paper filter and wash the precipitate by
decantation with four 5-ml portions of warm water, col-
lecting these washings with the filtrate in a 400- to 600-ml
conical flask. Treat the filtrate by P. 88. Wash the
precipitate with hot 0.6 n. NH 4 OH until no precipitate
forms when 1 drop of Ca(NO 3 )2 solution is added to a 5-ml
portion of the wash water. Transfer most of the precipitate
to the filter during the washing. Treat the precipitate by
P. 87.

If there is a large precipitate and magnesium may be pres-
ent (Note 4), decant the solution through a paper filter.
Add 5 to 15 ml of water to the precipitate, heat to boiling,
let the precipitate settle, and decant the solution. Repeat
this washing once and collect the washings with the filtrate
in a 400- to 600-ml conical flask. Dissolve the precipitate
on the filter by pouring dropwise through it 5 to 10 ml of HC1,
collecting the solution with the precipitate in the flask.
Wash the filter with 5 to 10 ml of water added dropwise.

" Hall, J. Am. Chem. Soc., 50, 2707 (1928).

17 A detailed discussion of the quantitative precipitation of calcium as
oxalate and the separation of calcium from magnesium is given by Kolthoff
and Sandeii, Quantitative Inorganic Analysis, Macmillan, 1936, pp. 323-337.

Kolthoff and Sandell, J. fhys. Chem., 37, 443, 459 (1933), have also in-
vestigated the coprecipitation of various substances by calcium oxalate
precipitates.

378 ALKALINE EARTH GROUP [P. 87

Add to the mixture in the flask 3 g of NH 4 C1 and swirl it
until most of the precipitate dissolves (Note 5); then add
to it 40 ml of water and 1 ml of 3 n. K 2 C 2 4 . Add 15 n.
NH 4 OH dropwise to the solution, shaking vigorously, until
it is just alkaline, and then add 2 ml more. Heat the mix-
ture to 70 to 80C. until the precipitate settles rapidly.
Decant the solution through the original filter, wash the
precipitate as directed in the paragraph above, and add the
filtrate and the first two washings to the original filtrate.
Treat the filtrate by P. 88. Treat the precipitate by P. 87.

Notes :

1. The volume of KaCjOi used should be based upon the total amount
of calcium and of magnesium thought to be present. This can be estimated
by the size of the ammonium carbonate precipitate obtained in P. 81 and
by the amount of barium and strontium already found present.

2. The more slowly the KjC204 is added to the solution the more favor-
able are the conditions for obtaining a precipitate that can be readily filtered.
A coarse crystalline precipitate of calcium oxalate can be obtained by begin-
ning the precipitation in an acid solution and then slowly neutralizing. This
cannot be done here, however, as chromate is reduced by both alcohol and
oxalate in hot acid solution.

3. The precipitate should not be allowed to stand for much longer than
10 min. because of the possible precipitation of magnesium oxalate from its
supersaturated solutions.

4. In the majority of cases, this procedure to remove any coprecipitated
magnesium may be omitted; the considerations discussed in Note 6, P. 82,
also apply here.

5. This acid mixture should not be heated, because of the possible reduc-
tion of the chromate by the oxalate present.

P. 87. Estimation of Calcium

Discussion. The principle involved in this estimation of calcium
is the same as that in the procedure for the estimation of strontium ;
that is, the oxalate precipitated with the calcium is titrated with
permanganate after the precipitate has been dissolved in acid. The
procedure is varied somewhat, depending upon whether a large or
small precipitate is obtained. If less than 100 mg of calcium are
thought to be present, the precipitate is dissolved in sulfuric instead
of hydrochloric acid, because the end-point obtained when oxalate
is titrated with permanganate is more stable in this acid (the pos-
sible oxidation of chloride by the permanganate is avoided) and
because any small precipitate of calcium sulfate which may be

P. 871 ESTIMATION OF CALCIUM 379

obtained would cause no trouble. If a large precipitate is obtained,
it is desirable that only a portion of the solution be titrated, because
of the excessive volume of permanganate which would be required.
Therefore, because hydrochloric acid more effectively dissolves
the precipitate of calcium oxalate, and in order to avoid a large pre-
cipitate of calcium sulfate (which may enclose calcium oxalate
particles), this acid is then used; the conditions for the titration of
the hydrochloric acid solution are the same as those in the estimation
of strontium.

The precise titration of oxalate in a sulfuric acid solution has been
discussed in P. XIV. There, in order to obtain very accurate re-
sults, most of the calculated volume of permanganate is added
rapidly to the cold solution of the oxalate. This procedure is not
practical when unknown amounts of oxalate are present and is not
justified in this procedure, as the titration can be more conveniently
carried out in a hot solution and as the titration errors caused by
so doing (see the discussion of P. XIV) are smaller than those likely
to arise from solubility and coprecipitation effects.

Procedure 87: ESTIMATION OF CALCIUM. If the precipi-
tate is thought to contain less than 100 mg of calcium (Note
1), transfer that in the flask to a 600-ml beaker with the aid
of 50 to 100 ml of water and add to it 30 ml of H 2 SO 4 . Re-
move the paper with the remainder of the precipitate from
the funnel, open it against the side of the beaker, and wash
the precipitate from the paper with a stream of water from
a wash bottle, using 50 to 100 ml of water. Dilute the solu-
tion to approximately 200 ml, heat it to 80 to 90C., and
titrate it with standard 0.1 n. KMn04. Stir the solution
constantly and do not add the KMn04 more rapidly than
its color is bleached. When the first permanent pink color
is obtained, push the filter paper into the solution and again
titrate the mixture slowly until a pink color is obtained
which remains for 30 sec. (Note 2). From the volume of
standard KMn0 4 used, calculate the amount of calcium
present.

If the precipitate is thought to contain as much as 100 mg
of calcium, remove the paper with the precipitate from the
funnel, open it against the side of a 400-ml beaker, and wash
the precipitate to the bottom of the beaker with a stream of
water, using not over 25 ml. Dissolve the precipitate in the

380 ALKALINE EARTH GROUP [P. 88

flask in 20 ml of hot 3 n. HC1 and pour this solution into the
beaker. Wash the beaker with 30 ml of hot 3 n. HC1, pour-
ing this solution slowly over the filter paper as well. Wash
the paper with 10 ml of hot water and discard it. Finally
cool the solution to room temperature, transfer it to a 100-ml
volumetric flask, dilute it exactly to the mark, and thor-
oughly mix the contents. Pipet that volume of the solution
which is thought to contain 50 to 75 mg of calcium into a
400-ml flask, dilute it to 100 to 125 ml with 1 n. HC1, add to
it 5 ml of the Zimmermann-Reinhardt preventative solu-
tion (P. 534), and heat it to 70 to 80C. Titrate the solu-
tion slowly with 0.1 n. KMnO 4 solution until the first pink
color is obtained which remains for 10 sec. From the
volume of standard KMnO 4 used, calculate the amount of
calcium present.

Notes:

1. If a very small precipitate is obtained, it should be estimated by com-
parison with known amounts of calcium precipitated under similar condi-
tions. Any strontium not precipitated in P. 84 would appear here. If
desired, the presence of calcium in the precipitate may be confirmed as
follows :

Pour repeatedly through the precipitate 5 ml of HjSC^ containing
just 0.5 ml of 95 per cent C2HsOH. Add to the clear solution 10
ml of C2HBOH. (White precipitate, presence of calcium.)
A small amount of any calcium oxalate present is always dissolved by the
sulfuric acid (oxalic acid being less ionized than sulfuric), and is reprecipi-
tated as CaSO4 2H2<3 by the addition of the large amount of alcohol. Stron-
tium oxalate would be metathesized to SrSO 4 by the sulfuric acid but, in
the presence of the small amount of alcohol first added, does not dissolve in
appreciable amounts. Magnesium sulfate is so soluble that, even if pres-
ent, it does not precipitate on the addition of the alcohol. Calcium salts
give an orange-red color to a flame; however, the intensity of the color is not
adequate for a sensitive confirmatory test.

2. If the paper were introduced into the solution at the beginning of the
titration, it would be disintegrated and make the detection of the end-point
more difficult; also, it would cause an error by reducing an appreciable
amount of the permanganate. 18 When the titration is carried out as di-
rected, this effect is negligible.

P. 88. Precipitation of Magnesium

Discussion. In both qualitative and quantitative methods,
magnesium is usually precipitated as magnesium ammonium phos-
phate. This compound is relatively soluble and tends to form

ll McBride and Scherrer, J. Am. Chem. Roc., 39, 928 (1917); Simpson, .7.
Ind. Eng. Chem., 13, 1152 (1921).

P. 89] ESTIMATION OF MAGNESIUM 381

supersaturated solutions. Therefore the precipitation is most
effectively made from a small volume of cold solution and by means
of a large excess of phosphate, ammonium ion, and ammonium
hydroxide. The ammonium hydroxide is added to repress the
hydrolysis reaction,

MgNH 4 PO 4 (.) + HOH = Mg + + + NH 4 OH + HPO 4 ",

which is a large factor in the solubility of the compound in water.
A more detailed discussion of the solubility of this compound is
given in the discussion of P. 162.

As would be predicted from their periodic relationships, phos-
phoric and arsenic acids are very similar in many of their properties.
This is especially true in regard to the solubilities of their salts, and
magnesium can be effectively precipitated as magnesium ammonium
arsenate under conditions similar to those of the phosphate pre-
cipitation. This substitution is made here because it provides a
means for the rapid volumetric estimation of magnesium in P. 89.
With the large volume of solution which is present here, precipita-
tion would be very slow if the alcohol added in P. 84 were not present.
The alcohol decreases both the solubility of the precipitate and its
tendency to form supersaturated solutions. In spite of the increased
solubility of the magnesium salt in the presence, of the excess 6f
oxalate added in P. 86, experiments have shown that by this pro-
cedure 0.5 mg of magnesium can be easily detected and that larger
amounts are quantitatively precipitated.

Procedure 88: PRECIPITATION OF MAGNESIUM. To the
filtrate from P. 86 add 10 ml of 15 n. NH 4 OH and then, 1
ml at a time and shaking after each addition, 10 to 25 ml of
1 f . Na2HAsO 4 solution. Cool the solution and let it stand
for at least 15 min., shaking it vigorously at frequent inter-
vals. (White precipitate, presence of magnesium.) Allow
the precipitate to settle and decant the solution through a
paper filter, using suction; then wash the precipitate with
10-ml portions of 1.2 n. NH 4 OH until 1 drop of Ca(NO 8 )2
solution gives no precipitate when added to a 5-ml portion
of the washings. Discard the filtrate and washings. Treat
the precipitate by P. 89.

P. 89. Estimation of Magnesium

Discussion. The method used here for the estimation of mag-
nesium is based upon the iodometric determination of the arsenic in
the magnesium ammonium arsenate precipitate.

382 ALKALINE EARTH GROUP [P. 89

In this method the magnesium ammonium arsenate precipitate
is dissolved in hydrochloric acid, and the arsenic acid thus produced
is reduced by the addition of potassium iodide to arsenious acid. 19
The iodine thus liberated is then titrated with standard thiosulfate
solution. The reaction

HsAs0 4 + 2H+ + 31"" = H 8 As0 3 + Is"" + H 2

is the reverse of the one discussed in P. XII and P. 43 and, as was
stated there, can be caused to proceed quantitatively in either
direction by proper control of the hydrogen ion concentration.
However, in solutions in which the hydrogen ion concentration is
such that it would be calculated that the reaction as written above
would proceed quantitatively to the right, the rate is likely to be too
slow for the method to be practical.

Williamson, 20 Gooch and Morris, 21 Rosenthaler, 22 Fleury, 23 Kolth-
off, 24 and others too numerous to tabulate have studied experi-
mentally the effects caused by the equilibrium and the rate of the
reaction between arsenic acid and iodide, and by other factors in-
volved in the determination, such as the oxidation of iodide in the
strongly acid solution by the oxygen of the air. They have also
investigated the influence of various conditions, such as the volume
of the solution, the concentration of the iodide, the amount of acid
present, and the temperature.

To review briefly these experimental studies: It has been found
that, if the concentration of acid is below approximately 4 n., per-
manent end-points are not obtained, and that with this concentra-
tion of acid some time should be allowed for completion of the
reaction. With too high a concentration of iodide and of acid, a
yellow precipitate of arsenious iodide is formed which obscures the
end-point. With the relatively high concentrations of iodide and
acid which are present, the air has to be excluded from the reaction
to avoid "oxygen error." The solution should be stirred vigorously

l Washburn,'/. Am. Chem. Soc., 35, 682 (1913), has shown that in acid solu-
tions as concentrated as the one used here the tripositive arsenic exists to a
considerable extent as AsO*. The volatility of arsenious chloride and bro-
mide and the possible precipitation of arsenious iodide in this procedure show
that in concentrated solutions of these halogen acids there also exists an
appreciable concentration of the corresponding arsenic halide.

10 Williamson, J. Soc. Dyers and Colorist, 1896, 86-89.

21 Gooch and Morris, Z. anorg. allgem. Chem., 25, 227 (1900).

" Rosenthaler, Z. anal Chem., 45, 596 (1906).

" Fleury, /. de Pharmacie et de Chemie, 21, 385 (1920).

" Kolthoff, Z. anal. Chem., 60, 399 (1920).

P. 89] ESTIMATION OF iMAGNESIUM 383

as the thiosulfate is added, to avoid the decomposition of the latter
by the high concentration of acid which is present. Also, the acid
concentration is so high that the use of starch is not advisable; the
end-point can be determined solely by the disappearance of the
iodine from the otherwise clear solution, or some organic solvent
for iodine, such as carbon tetrachloride, may be used.

Experiments have shown that under the conditions Of the pro-
cedure given below, results precise to within 0.3 per cent can be
readily attained in the iodometric determination of arsenic acid.

As the other alkaline earth elements form insoluble phosphates
and arsenates and would be precipitated in P. 88 if they had not
been previously completely removed, a confirmatory test for the
presence of magnesium is desirable if only a small precipitate is
obtained. This can be made by precipitating magnesium hy-
droxide in the presence of the colored organic compound para-
nitrobenzene-azo-resorcinol, whereby the precipitate is colored
blue as a result of the formation of an adsorption compound (a lake)
between the precipitate and the organic dyestuff. 25 The method
can be used for the detection of very small traces of magnesium, and
experiments 28 have shown that, as carried out below, 0.01 mg can
be easily detected. The presence of arsenate (and phosphate) in
the absence of other alkaline earth elements did not interfere with
the test; calcium, strontium, and barium did not give any color;
and when 10 mg of calcium and arsenate were both present, the
colorless precipitate of calcium arsenate was so milky that it was
impossible to detect less than 0.1 mg of magnesium. When am-
monium salts, phosphate, or arsenate are present, sufficient excess
of a strong base must be added to insure the precipitation of the
magnesium hydroxide; a tenfold excess of sodium hydroxide over
that specified below did not diminish the intensity of the color with
given amounts of magnesium.

Procedure 89: ESTIMATION OF MAGNESIUM. Dissolve
the precipitate on the filter and wash the filter by pouring
slowly through it 20 to 50 ml of 6 n. HC1, collecting the solu-
tion with the precipitate in the flask (Note 1). Dissolve
the precipitate in the flask and cool the solution to room
temperature.

26 Suiteu and Okuma, /. Soc. Chem. Ind., Japan, 29, 132 (1926); Ruigh, J.
Am. Chem. Soc., 51, 1456 (1929); Stone, Science, 72, 322 (1930).
10 By Erwin Baumgarten.

384 ALKALINE EARTH GROUP [P. 89

If less than 50 mg of magnesium are thought to be present
(Note 2), add to the cold solution, 0.2 g at a time and
shaking after each addition, 1 g of solid NaHCOa and then
1 g of solid KI. Close the flask, swirl the mixture until the
KI is dissolved, and allow it to stand for 5 min. Slowly
titrate the mixture with 0.1 n. Na 2 S20 3 , swirling the solu-
tion continuously, until the iodine color disappears (Note
3). From the volume of standard Na2S2O 3 used, calculate
the amount of magnesium present.

If more than 50 mg of magnesium are thought to be pres-
ent, transfer the cold solution to a 100-ml volumetric flask,
dilute it to the mark with HCl/and thoroughly mix the con-
tents. Pipet that volume of the solution which is thought
to contain 30 to 40 mg of magnesium into a 200-ml flask
(preferably one fitted with a ground-glass stopper), and
treat it as directed in the paragraph above.

Notes:

1. If only a small precipitate (corresponding to 3 to 5 mg) is obtained or
only a qualitative confirmation of the presence of magnesium is desired,
proceed as follows:

Dissolve the precipitate in 2 to 5 ml of HC1 (if there is a large
precipitate, take only an amount corresponding to 5 to 10 mg of
magnesium). Add 1 drop of the para-nitrobenzene-azo-resorcinol
reagent (a 0.5 per cent solution in 0.3 n. NaOH), make the solution
alkaline with NaOH, and add 1 ml in excess. Shake the mixture
for 10 to 15 sec. and allow it to stand for 2 min. (Blue precipitate,
presence of magnesium.) If uncertain as to a bluish-colored pre-
cipitate, because of the reddish-violet color of the reagent, filter the
mixture through a small paper filter and wash the precipitate and
filter dropwise with 5 to 10 ml of cold water to which has been added
3 drops of NaOH.

2. As the amount of iodide added is based on not more than an amount
of arsenic corresponding to 50 mg of magnesium being present, it is essential
that only a portion of the solution be taken when larger amounts have been
precipitated. If more than 1 g of KI is added, a precipitate of Asia is likely
to separate.

3. The color of iodine in a solution containing iodide is so pronounced
that this end-point can be readily placed within 1 drop (0.02 to 0.04 ml)
of 0.1 n. Naj^Oa. If desired, 3 to 5 ml of carbon tetrachloride can be
added and the solution can be titrated until no more color is observed in
this solvent. If this method is used, it is essential that the solution be
thoroughly shaken as the end-point is approached, in order that equilibrium
between the aqueous and carbon tetrachloride solutions be attained; also,
the thiosulfate should not be added more rapidly than the color is removed
from the aqueous layer, as it is rapidly decomposed by this concentration
of acid.

The Analysis of the Alkali Group

P. 91. The Detection of the Alkali Group and Estimation of the
Total Amount of Sodium and Potassium Present

Discussion. Sodium and potassium, which constitute the Alkali
Group, should have remained in the solution through all the pre-
vious group separations and now have to be detected in the filtrate
from the Alkaline Earth Group precipitation. This is done by first
removing the ammonium salts, then precipitating the potassium as
perchlorate from a mixed alcohol-ether solution, and, finally, pre-
cipitating the sodium as chloride by saturating the alcohol-ether
solution with hydrogen chloride gas. Ammonium, which has many
of the properties of an alkali metal ion, behaves similarly but cannot
be tested for in this solution because it has been introduced as a
reagent in the preceding group separations. However, because
of the striking similarity of ammonia to an alkali metal, a procedure
for its detection and estimation, using another portion of the sample
of the original material, is given immediately after the procedures
for the analysis of this group.

Before beginning the analysis of the Alkali Group, it is necessary
that the ammonium salts introduced in precipitating the Alkaline
Earth Group be completely removed, as potassium is separated
from sodium by precipitation as the perchlorate, and as ammonium
perchlorate is insoluble under the same conditions. The ammonium
salts are removed by treatment with concentrated nitric acid (see
the discussion of P. 81) and by heating at a temperature sufficiently
high to cause their complete volatilization. Also, if, after vola-
tilizing the ammonium salts, there is no residue, further procedures
are unnecessary.

Sulfate has to be tested for and removed, as sodium sulfate is
insoluble in a mixed alcohol-ether solution and would cause the
precipitation of sodium with the potassium perchlorate. Even if
sulfate is not present in the original material, sulfur compounds are
frequently introduced into the analysis during the hydrogen sulfide
and ammonium sulfide precipitations. These compounds are then
oxidized to sulfate by the nitric acid treatment (P. 81) used to
destroy the ammonium salts. The sulfate can be precipitated as
either the barium or the lead salt; the latter is used, as lead can be
removed from the solution by hydrogen sulfide, whereas the barium

385

4 OH, C,HOH

&
3 ^

00 ^2

S~\

(2)

I 5

h?*^

N 'l^

O :

ei 1 I

TABULAR OUTLINE X
THE ANALYSIS or THE ALKALI GROU

e Earth Group Precipitation: Na+, K+, NH 4 +, COr, SOr
NOt, evaporate to dryness, heat just below dull redness,
more precipitate forms. (P. 91)

'"^

Filtrate: Na+, K + , XO 3 -, H,S
Evaporate to dryness. Add HCI, evaporate.
Heat at SOOC.
Weigh residue of NaCl and KCI. (P. 91)

Dissolve residue in H 2 0. Add HClO*;fume;
Add 15 ml 95 per cent C t H>OH and 15 ml (dH
Filter the precipitate on a weighed crucible.

Filtrate : NaClO 4 in C*H 6 OH and
Cool in ice water. Saturate wit

Precipitate: NaCl
Method I:
Heat and weigh.
Method II:
Dissolve in H*0. Add K 2 Cr
Titrate with AgN0 3 .

Precipitate

The Detection and Estimation of Ammonia
a portion of the original material with NaOH.
tection: Test escaping gas (NH t ) with litmus. (Red lit
bimation: Distill into standard HCI solution.
stillate: NH 4 + in excess HCI. Titrate excess HCI unth

'S.

af?i

|lll

"*"* 'W

f"!

5*1

2'-
flSl^
53 S O 4
t ^

was

J3^

a -if

||1

1^1

P. 91] DETECTION 387

has to be removed by precipitation with ammonium carbonate,
which again introduces ammonium salts into the solution. It is
an advantage that the alkali metal salts have been converted to
nitrates, as this prevents the possible precipitation of lead chloride.

Upon evaporation of the filtrate and removal of the hydrogen
sulfide, the elements present are changed from their nitrates to their
chlorides, which are more stable under heat, and the total weight of
sodium and potassium chloride is determined. This often elimi-
nates the subsequent precipitation and estimation of sodium, as,
after estimating the potassium, the amount of sodium present can
be obtained by difference.

Optional Method for the Qualitative Detection of Sodium and Po-
tassium. In case only small amounts of these two elements are
present, or only their qualitative detection is desired, there is pro-
vided in Note 4 of the procedure below an optional method for this
purpose, which is much shorter than the perchlorate separation.
In this procedure the residue obtained after the ammonium salts
are removed is dissolved, and this solution is divided into two por-
tions. One portion is tested for potassium by the addition of a
sodium cobaltinitrite (Na 3 Co(N0 2 )6) reagent. Even 0.2 mg of
potassium causes the formation of a detectable yellow precipitate,
potassium sodium cobaltinitrite (K 2 NaCo(NO 2 )6), similar to that
used for the optional separation of cobalt from nickel (P. 64). The
other portion is tested for sodium by addition of a reagent containing
magnesium acetate, uranyl acetate (U0 2 (C2H 3 02)2), and acetic
acid. 1 As little as 0.3 mg of sodium causes the formation of a volu-
minous, greenish-yellow, crystalline precipitate having the compo-
sition NaMg(UO 2 ) 3 (C2H 3 2 V9H 2 0.

Procedure 91: DETECTION OF THE ALKALI GROUP.
Evaporate the filtrate from the Alkaline Earth Group pre-
cipitation (P. 81) to a volume of 5 to 10 ml, filter out any
precipitate on a small filter, and wash the filter with a few
milliliters of cold water. Collect the filtrate and washings
in a 50-ml casserole. (Discard the precipitate. Note 1.)
Cover the casserole with a clock glass, evaporate the filtrate
almost to dryness, add 2 to 5 ml of 16 n. HN0 3 , and care-

1 This reagent was proposed for use in the quantitative determination of
sodium by Blanchetiere, Bull. Soc. Chim. (4), 33, 807 (1923); it, and a similar
reagent in which the magnesium acetate is replaced by zinc acetate, have
been investigated by Barber and Kolthoff, /. Am. Chem. Soc., 50, 1625 (1928).

388 ANALYSIS OF THE ALKALI GROUP [P. 91

fully evaporate the solution to dryness, keeping the mixture
in motion or allowing the evaporation to take place on a
steam bath. Wash the clock glass and the sides of the
flask with 1 to 2 ml of 16 n. HNO 3 , and again evaporate the
solution to dryness. Remove the clock glass and heat
the casserole over a flame at a temperature just below dull
redness until fumes are no longer given off (Note 2) . (White
residue, presence of the Alkali Group.)

Dissolve the residue in 10 ml of hot water (Notes 3 and 4)
and add to it 0.1 ml of 1 n. Pb(N0 3 ) 2 . If a white precipitate
forms, continue adding the Pb(N0 3 )2 in 0.1-ml portions until
precipitation ceases. Shake the mixture and let the pre-
cipitate settle after adding each portion of the Pb(NO 3 )2i
avoid a large excess. Filter the solution through a small
paper filter, and wash the precipitate with 3 to 5 ml of cold
water added dropwise, collecting the filtrate and washings
in a 100-ml flask. (Discard the precipitate.) Saturate
the solution with H 2 S under pressure (not by passing the gas
through the solution Note 3, P. 11), heat the mixture al-
most to boiling, and at once filter out the precipitate on a
small paper filter. Wash the precipitate with 5 to 10 ml of
hot water, added dropwise, and collect the wash water with
the filtrate in a 100-ml flask (Note 5) which has been previ-
ously heated to 250 to 300C. and then allowed to cool in a
desiccator and weighed. Immediately evaporate the solu-
tion to dryness (Note 6), moisten it with 12 n. HC1, and
again evaporate it, repeating this process once. Finally,
heat the flask to 250 to 300C. for 10 min., allow it to cool in
a desiccator, and weigh it. Note the total weight of alkali
metal chloride present. Treat the residue by P. 92.

Notes:

1. Traces of the alkaline earth elements, of aluminum hydroxide, or of
silica may precipitate as the solution is evaporated. These should be
carefully filtered out through a small filter so that, in case alkali elements
are absent, no residue will remain after the ammonium salts are removed.

2. The sides and entire bowl of the casserole should be heated in order to
remove all traces of ammonium salts. The vessel should not be heated to
a red heat, as the alkali salts would then be appreciably volatile.

3. An insoluble residue which is dark in color is usually due to organic
material introduced into the analysis from filter papers or in the reagents,
especially alcohol. A white residue of silica may also be obtained. These
are usually not sufficiently large to interfere with the Pb(N0 3 )2 treatment
and can be filtered out with the PbSOi.

P. 92) PRECIPITATION OF POTASSIUM 389

4. If it is desired only to detect the presence of the two metals, treat the
solution as follows:

Detection of Potassium: Pour half of the solution into a test tube,
add 5 ml of NasCo(N02)6 reagent, and let the solution stand,
shaking it frequently, for 10 min. (Yellow precipitate, presence
of potassium.) Compare the precipitate with known amounts of
potassium precipitated under similar conditions. (If the mixture
is allowed to stand until the precipitate settles in the bottom of the
test tube, 0.2 mg of potassium can be detected. Ammonium salts
must be removed, as they cause a similar precipitate.)
Phosphate would be precipitated by the uranium in the magnesium uranyl
acetate reagent; therefore, if phosphate has been found present (see P. 54),
it should be removed before testing for sodium. This can be done by the
procedure given above for removing sulfate; therefore, if it is present,
proceed as follows:

Treat the second half of the solution with Pb(NOa)2 and H^S
as directed in the second paragraph of the procedure above, evapo-
rate the H2S filtrate to dryness, dissolve the residue in 5 ml of hot
water, and treat it as directed in the next two paragraphs.

Detection of Sodium: If less than 10 mg of potassium are present
(in the half portion), pour the remaining half of the solution into a
test tube, add 10 ml of magnesium uranyl acetate reagent, and let
the solution stand, shaking it frequently, for 10 min. (Pale greenish-
yellow crystalline precipitate, presence of sodium.)

If more than 10 mg of potassium are present, evaporate the solu-
tion in the flask almost to dryness and add to it 5 ml of 12 n. HC1
and 5 ml of 95 per cent C2H60H. Cool the mixture and decant
the solution through a small filter. Evaporate the solution just to
dryness, dissolve the residue in 5 ml of water, and treat it by the
paragraph above.

If the mixture is allowed to stand for 10 min., 0.3 mg of sodium can be
detected. Not more than 20 mg of potassium can be present without
likewise causing a precipitate; therefore, larger amounts are removed by
treating the residue with a mixture of hydrochloric acid and alcohol. Only
about 20 mg of sodium remain in the alcoholic solution, so that the estima-
tion of larger amounts than this should be made from the original residue
and from the amount of potassium found.

5. If only a qualitative separation of the alkali elements is desired, it is
not necessary to weigh the residue. In this case the solution should be
collected in a 200-ml flask and treated directly by P. 92.

6. The solution should not be allowed to stand for any length of time
before expelling the H^S, as the latter may oxidize on contact with the air
and form a white precipitate of sulfur. In case such a precipitate does
form, it should be carefully filtered out.

P. 92. Precipitation of Potassium

Discussion. This method for the precipitation of potassium and
its separation from sodium is based upon the difference in solu-

390

ANALYSIS OF THE ALKALI GROUP

[P. 92

bilities of the perehlorates of the two elements. As potassium
perchlorate is relatively soluble in water, it has been customary to
effect the separation in nearly anhydrous ethyl alcohol, 97 to 99
per cent, and even in some cases to saturate this alcohol first with
potassium perchlorate. 2 As a substitute for the ethyl alcohol,
various other organic solvents have been investigated, and the use
of a mixture of butyl alcohol and ethyl acetate has been recom-
mended. 8

This reagent, while very effective, may not be available, especially
as the solvents should be anhydrous, and the butyl alcohol often con-
tains impurities which give it an offensive odor. A study by Walter

TABLE XXIX
THE SOLUBILITY OP SODIUM AND POTASSIUM PBRCHLORATBS IN

VARIOUS SOLVENTS

The dry perchlorates were shaken with the reagents indicated at room

temperature (approximately 20C.) for 20 min.; then 50-ml

portions were analyzed.

Expt.

1
2
3
4
5
6
7
8

Volume Taken (ml)

Solubility (gin 50 ml)

Water

HC1O 4

C 2 H 6 OH
(94 per
cent)

(C 2 H,)20

Sodium

Potassium

3.0
3.0
3.0

2.60
2.20

0.4
0.8

31.25
46.88
105.0
105.0
46.88
46.88

0.000
2.039
1.130
0.996
0.336
0.2085

0.603

0.0000
0.0049
0.0020
0.0011
0.0003
0.00001
0.0004
0.0003

60.0
31.25
15.63
20.50
50.0
15.63
15.63

D, Wilkensen, Jr., has shown that the addition of ethyl ether to
ethyl alcohol very markedly decreases the solubility of potassium
perchlorate and that, although a decrease in the solubility of sodium
perchlorate is also caused, the ratio of the solubility of the sodium
salt to that of the potassium is greatly increased. Thus, as can be
calculated from the data shown in Table XXIX (Experiment 2),
with 94 per cent ethyl alcohol this ratio of the formal solubilities

f Davis, J. Agr. Sci., 6, 52, 512 (1912); Gooch and Blake, Am. J. Sci., 44,
381 (1917); Baxter and Kobayashi, /. Am. Chem. Soc., 39, 249 (1917); 42, 735
(1920).

Willard and Smith, /. Am. Chem. Soc., 44, 2816; 46, 293 (1923); G. F.
Smith, tWd., 47, 762 (1925); Smith and Ross, ibid., 47, 774, 1020 (1925).

P. 92]

PRECIPITATION OF POTASSIUM

391

is approximately 700, while with a mixture of 1 volume of 94 per
cent alcohol and 3 volumes of ethyl ether it is about 50 times as
great (Experiment 6).

It is an advantage that the solubilities of the salts are so reduced
that, instead of treating a dry mixture with the solvent and extract-
ing the sodium, it is possible to dissolve the salts in a small amount
of water and to precipitate the potassium perchlorate gradually by
slowly adding first the alcohol and then the ether; by this precipita-
tion process there should be less tendency for sodium to be carried
out with the potassium precipitate. The quantitative value of this

TABLE XXX

THE SEPARATION OF POTASSIUM AND SODIUM

A slight excess of perchloric acid was added to the alkali salts and the mix-
ture was evaporated just to complete dryness. Water was then added, 1 ml
at a time, the mixture being heated after each addition, until the residue was
dissolved; the solution was then evaporated until just 1 ml of water remained.
The mixture was cooled and 15 mi of 95 per cent CjH*OH and 45 ml of (CjHsJiO
were slowly added; then the mixture was allowed to stand for 10 min., with
frequent shakings. The solution was decanted through a Gooch-type fil-
ter, and the precipitate was washed with 50 ml of a solution containing 1
volume of alcohol to 3 volumes of ether. The precipitate was dried at 350C.
and weighed. The filtrate was evaporated to dryness, dried at 350C., and
weighed.

Expt.

1
2
3
4
5
6

Potassium (g)

Sodium (g)

Taken

Found

Taken

Found

0.3642

0.3670

0.2916

0.2899

0.3642

0.3641

0.2916

0.2917

0.0009

0.0011

0.3871

0.3870

1.4449

1.4445

0.0008

0.0011

0.4642

0.4638

0.0015

0.0019

0.3641

0.3636

None

0.0005

method was tested by the procedure given with Table XXX, and
the results obtained are shown there.

For the purpose of this system, where the occlusion of small
amounts of perchloric acid within the potassium perchlorate pre-
cipitate would not introduce serious errors, and where the presence
of phosphate and borate would cause the sodium salts to precipitate
from a neutral solution, it seemed more expedient to use a solvent
containing a small amount of perchloric acid and equal volumes of
alcohol and ether. In testing the procedure given below, it was
found that 0.5 mg of potassium gave an easily detected precipitate

392 ANALYSIS OF THE ALKALI GROUP [P. 92

and also that, when 250 mg each of sodium and potassium were
taken, less than 1 mg of sodium was found in the precipitate and
no detectable amount of potassium remained in the filtrate.

Potassium is often separated from sodium and gravimetrically
determined by precipitation as the chloroplatinate (K 2 PtCle); this
method is not used here because of the cost of the reagent and be-
cause of the necessity of removing the excess of the reagent before
the precipitation of the sodium can be made.

Procedure 92: PRECIPITATION OF POTASSIUM. Transfer
the residue obtained by P. 91 to a 200-ml conical flask with
the aid of as little water as possible. Add to the solution 1
to 3 ml of 9 n. HC104 and evaporate until dense white fumes
are copiously evolved and not over 1 ml of HC1O 4 remains
(Note 1). Keep the contents of the flask in constant mo-
tion during this evaporation. Let the residue cool some-
what, add 1 ml of water, warm the mixture, and shake it
until the residue dissolves or is thoroughly mixed. Cool
the flask to room temperature with tap water, add slowly 15
ml of 95 per cent alcohol, and then add 15 ml of ethyl ether.
Shake the mixture and let it stand for 5 min. in a beaker
of ice water. (White crystalline precipitate, presence of
potassium.) Filter the mixture through a sintered-glass
crucible (Notes 2, 3) which has been previously heated to
300 to 350C., cooled, and weighed (Note 4). Wash the
precipitate with 20 to 30 ml of a cold solution made by
adding 1 volume of alcohol to 2 volumes of ether. Transfer
all of the precipitate from the flask by squirting this wash
solution from a small wash bottle onto the precipitate
while the flask is held in an inclined position. Collect the
washings with the filtrate in a large test tube (25 mm by
200 mm) and treat it by P. 94. Treat the precipitate by
P. 93.

Notes:

1. Enough HC104 should be added to combine with all of the alkali
elements present. If, after adding the HC10 4 and evaporating, no fumes
are evolved before the mixture becomes nearly dry, 1 ml more acid should
be added and the mixture should again be evaporated; an excess of more
than 1 ml of HClOi should be avoided.

2. If available, sintered-glass crucibles should be used here. These cru-
cibles are rapidly prepared for use and brought to constant weight, and are
very satisfactory in this determination. If they are not available,
Gooch-type asbestos filters can be used.

P. 93] ESTIMATION OF POTASSIUM 393

3. Caution: Ether should not be handled near a flame (see the second
paragraph of Note 8, P. 52).

4. The heating prior to weighing may be conveniently carried out in a
manner similar to the final heating and drying of the precipitate in P. 93
(see also Note 2 of that procedure) .

P. 93. Estimation of Potassium

Discussion. A gravimetric method is used for estimating the
potassium, because this element has been detected and separated as
the perchlorate (which can be conveniently heated and then weighed)
and because no rapid volumetric estimation of either the potassium
or the perchlorate is available. Smith and Ross* have shown that
it is necessary to heat the precipitate to approximately 350C. in
order to expel the occluded perchloric acid. If only a very small
precipitate is obtained, the presence of potassium can be confirmed
by dissolving the precipitate and reprecipitating the potassium as
potassium sodium cobaltinitrite, K 2 NaCo(NO 2 )6, as directed in
Note 4, P. 91. This confirmatory test is advisable, as, if either
chloride or sulfate were left with the alkali metals, sodium chloride
or sulfate would separate out from the alcohol-ether solution. The
only other anions likely to be present are phosphate and borate;
however, these are almost completely removed with sulfate by the
treatment with lead in P. 91 and are much less ionized than per-
chloric acid and do not cause the sodium to precipitate.

Procedure 93: ESTIMATION OF POTASSIUM. Place the
crucible with the precipitate from P. 92 (Note 1) in an oven
at 110C. until it is dry (Note 2), suspend it well down in a
large iron or nickel crucible by means of a nichrome or clay-
covered triangle, and heat it to between 300 and 350C. for
10 min. Place the crucible in a desiccator, allow it to cool
to room temperature, and weigh it. Repeat the heating and
weighing until the precipitate no longer changes weight.
From the weight of precipitate obtained, calculate the
amount of potassium present (Note 3).

Notes :

1. If only a small precipitate is obtained, the presence of potassium
should be confirmed by dissolving the precipitate in 5 ml of water and
treating the solution as directed in Note 4, P. 91.

2. The crucible and precipitate may be more quickly dried by placing it
directly in the metal crucible next used and then very gradually raising
the temperature; too abrupt heating will cause spattering. A thermometer

4 fiUvUk a *t,l Dsv T JM rrj.^^, .Q^ A.f T7A

394 ANALYSIS OF THE ALKALI GROUP [P. 94

should be suspended inside the crucible so that the temperature can be
controlled. If an electric furnace with temperature control is available,
it can be used to advantage.

3. If only an approximate estimation of the amount of sodium present
is desired, this may be calculated from the amount of potassium found here
and the total weight of the alkali chlorides found in P. 91.

P. 94. Precipitation of Sodium

Discussion. As the filtrate should now contain only sodium
perchlorate, perchloric acid, alcohol, and ether, the obvious method
of detecting the sodium would be to evaporate the solution to dry-
ness and note if any residue remains. However, if the evaporation
of solutions containing alcohol and perchloric acid is not carried out
under very carefully controlled conditions, a violent and dangerous
explosion will result. Because of this, and because a salt which is
more readily estimated is obtained, the sodium is precipitated as
the chloride by saturating the filtrate with dry hydrochloric acid gas.
By this means, 0.3 mg of sodium can be detected and larger amounts
can be completely precipitated. Any potassium not precipitated
by the perchloric acid will also be precipitated.

The precipitate has to be washed free from perchloric acid with
ether, as any perchloric acid which remains on the precipitate will
displace the more volatile hydrochloric acid from the salt when it
is dried. This would result in a mixture of perchlorate and chloride
and cause an error in the subsequent estimation of the sodium.

Procedure 94: PRECIPITATION OF SODIUM. Immerse the
test tube containing the filtrate from P. 92 in a mixture of
ice and water and pass a fairly rap'd stream of dry HC1 gas
(Note 1) through the solution until it is saturated, as shown
by the gas no longer being absorbed. (White crystalline
precipitate, presence of sodium.) Filter the solution
through a sintered-glass filter (Note 2) which has been
moistened with alcohol and then sucked dry. Wash the
precipitate with three to five 5-ml portions of ether, and dry
it at 110C. for 20 min. (see the second paragraph of Note
8, P. 52). Treat the precipitate by P. 95-. Caution: Discard
the filtrate; do not heat or evaporate it (Note 3).

Notes:

1. The dry gaseous HC1 should be generated as follows: Fit a 500-ml
conical flask with a two-hole rubber stopper carrying in one hole a delivery
tube and in the other a thistle tube. Place in the flask the necessary amount

P. Ml' ESTIMATION OF SODIUM 395

of solid NaCL Support a short test tube upright on the bottom of the
flask by having the end of the thistle tube extend just below the top of the
test tube. Pour 95 per cent HfoSC^ through the thistle tube until the test
tube overflows; thereafter, add the acid as it is needed to cause a vigorous
flow of gas. After a volume of acid equal in milliliters to the grams of salt
taken is added, the mixture should be heated to 60 to 80 C. as the gas is
needed. When the gas is no longer needed, the evolution can be stopped by
cooling the flask. Whatever the type of generator used, it should be pro-
vided with an efficient safety tube, since the HC1 is so vigorously absorbed
at first that it tends to suck the alcohol-ether solution back into the generator.

2. If desired, the sodium may be estimated by drying at 350C. and
weighing the NaCl precipitate. In this case, the precipitate should be
collected on a sintered-glass or Gooch-type asbestos filter which has previ-
ously been dried at 350C., cooled, and then weighed. If a volumetric
estimation is to be made, a simple asbestos filter is satisfactory (see the
discussion of P. 95 in regard to these optional procedures).

3. Solutions containing perchloric acid and alcohol should not be heated,
as a violent explosion may result. Noyes and Bray 5 have found, and
repeated experiments have confirmed, that it is possible to evaporate solu-
tions containing perchloric acid and alcohol by the following procedure :

Transfer the solution to a casserole, dilute it with one-fourth its
volume of water, and evaporate it on a steam bath until the volume
is only slightly greater than that of the 9 n. perchloric acid present
(do not cause the HCIO* to fume). Cool the solution and add 16 n.
HNOa, a drop at a time, as long as any reaction is observed, and then
add 5 ml more. Finally evaporate the solution to dryness, taking
care that, should an explosion occur, a minimum of danger will
result.

It is therefore possible to detect the sodium by evaporating the filtrate
from P. 92 as indicated above and noting if any residue of sodium perchlo-
rate remains. If it is desired to estimate sodium after detecting it by this
process, the perchlorate can be converted to chloride as follows: Add to
the casserole 0.5 to 2 g of solid NH4C1, just moisten the entire mass, and
then heat the casserole as was directed for removing ammonium salts in P. 91.
If a large amount of soflium is present, the process should be repeated to
insure complete decomposition of the perchlorate.

P. 95. Estimation of Sodium

Discussion. Several methods are available for estimating the
amount of sodium present. First, after the total weight of the
chlorides has been obtained in P. 91 and the potassium has been
separately estimated in P. 93, the sodium present can be calculated.
This is a difference method and is subject to error when the amount

1 Noyes and Bray, Qualitative Analysis for Ike Rare Elements, Macmillan,
1927, pp. 259 and 472.

396 ANALYSIS OF THE ALKALI GROUP [P. 95

of sodium is small, the more so since the precision in weighing the
relatively heavy flask with its large surface (in P. 91) would probably
be low.

Second, as the crucible has been weighed before the filtration, the
crucible and precipitate can be heated so as to dry the salt com-
pletely and a gravimetric determination can be made.

Finally, after the precipitate of sodium chloride has been heated
until it is free of hydrogen chloride, the sodium can be determined
indirectly by either of two volumetric methods for determining the
chloride. First, with standard solutions of silver nitrate and thio-
cyanate available, the chloride present can be estimated by adding
an excess of the silver nitrate and then titrating this excess as was
done in estimating bismuth in P. 27. Or, second, the chloride can
be determined by a titration with a standard silver nitrate solution
and the end-point can be determined either (1) by adding a soluble
chromate to the solution and noting the appearance of the first
reddish precipitate of silver chromate this is known as the Mohr
titration or (2) by the use of an adsorption indicator. This direct
titration with a standard silver nitrate solution (using either means
of determining the end-point) is recommended because of the pre-
cision and rapidity with which it can be carried out; procedures are
given below which permit the optional use of either end-point. Be-
cause it has been sio extensively used, especially for the determination
of the chloride in drinking waters, the principles involved in the
Mohr method of obtaining the end-point will be discussed in detail.
A general discussion of adsorption indicators has been given in
P. VI, "The Standardization of a Thiocyanate Solution. "

The Use of Chromate as an Indicator in Titrating Chloride with
Silver Ion. In both methods of determining the end-point, the
sodium chloride is dissolved and a neutral solution is thus obtained.
Then for carrying out the Mohr titration, chromate is added as an
indicator, and the solution is titrated with standard silver nitrate;
silver chloride is formed until the chloride ion concentration is re-
duced to a very low value, when a reddish precipitate of silver
chromate begins to form, and the first appearance of this color is
taken as the end-point. This method, which is one of the older
analytical processes, was first described by Mohr 8 and has been
extensively investigated. The theoretical accuracy of the method
can be calculated as follows: At the end-point it is desired that the
equivalents of silver added be equal to the equivalents of chloride

Mohr, Liebig'* Ann. d. Ctoro., 97, 335 (1866).

P. 95] TITRATION OF CHLORIDE 397

originally present. Therefore, the percentage titration error, T.E.,
will be the total equivalents of silver added, 2 Eq.Ag, minus the
total equivalents of chloride present, 2 Eq.Cl, times 100, divided
by the total equivalents of chloride present, thus;

= 2 Eq.Ag - 2 Eq.Cl
2 Eq.Cl

The silver which has been added is present at the end-point as
the silver chloride precipitate, as the silver chromate precipitate,
and as silver ion; the chloride is present as the silver chloride pre-
cipitate and as chloride ion.

As, at the end-point, the total equivalents of silver, 2 Eq.Ag, are
equal to the equivalents of silver chloride, Eq.AgCl, plus the equiva-
lents of silver chromate, Eq.Ag 2 CrO 4 , plus the equivalents of silver
ion, Eq.Ag + , and as the total equivalents of chloride, 2 Eq.Cl, are
equal to the equivalents of silver chloride, Eq.AgCl, plus the equiva-
lents of chloride ion, Eq.Cl"", Equation 1 can be rewritten

(Eq.AgCl + Eq.Ag 2 Cr0 4 + Eq.Ag+)

T.E. - s " (Eq ' AgC1 + Eq - Cr) X 100. (2)

S Eq.Cl

If these quantities can be evaluated, the titration error can be
calculated.

It is seen that the equivalents of the silver chloride precipitate
drop from the equation; this is obvious, as the equivalents of silver
and of chloride are equal in the precipitate. The equation then
reduces to the form

T = (Eq.Ag 2 Cr0 4 + Eq.Ag+ - (Eq.CD ()

2 Eq.Cl

The equivalents of silver chromate represent the amount of that
precipitate which is required before it is visible. This quantity is
variable, being dependent upon the individual observer, the con-
centration of the chromate, the coagulation of the precipitate, and
the conditions obtaining at the time of the observation; under
favorable conditions it appears to range from between 0.08 and
0.2 mg of silver chromate, but, if care is not taken, it will be much
larger. Thus, in a series of observations made with 100 ml of a
solution which was 1.2 X 10~ 2 n. (6 X 10~ 3 f .) in potassium chromate,
it was found that 0.28, 0.24, and 0.22 ml or an average of 0.25 ml
of 0.01 n. silver nitrate was required to give the first perceptible

398 ANALYSIS OF THE ALKALI GROUP IP. 95

color change. Assuming that an equilibrium was attained in these
experiments, and taking the average value given above, it is seen
that 2.5 X 10~ 6 equivalent of silver has been added. As the silver
added is small in comparison to the chromate, it can be assumed
that the initial and final chromate ion concentrations have approxi-
mately the same value, that is, 6 X 10~ 3 m. 7 The solubility product
of Ag 2 CrO 4 is

] = 2 X 10

Therefore it is calculated that the silver ion concentration in equi-
librium with the above concentration of chromate is 1.8 X 10~ & ,
or that there will be 1.8 X 10~ 6 equivalent of silver as silver ion
remaining in the 100 ml of solution; therefore, 0.7 X 10~ 6 equivalent
of silver chromate has been precipitated. Applying these conditions
to a titration and taking the solubility product of silver chloride as

[A g + ][cr] = i<r 10 ,

it is seen that, in a solution in which the silver ion concentration is
1.8 X 10~ 5 , the chloride ion concentration will be 5.6 X 10~ 6 , or in
100 ml of solution there will be 0.56 X 10~ 6 equivalent of chloride
ion. If we assume some quantity of chloride initially present, say
25 ml of 0.1 n. solution or 2.5 X 10~ 3 equivalent, all of the quantities
are available for substitution in Equation 3 above for the titration
error, thus:

(0.7 X 10~ 6 + 1.8 X 10" 6 ) - (0.56 X 10~ 6 ) v inn
T.E. = ______ X 100

= 0.04 per cent, positive error.

From an inspection of the factors in the expression, it is seen that
the percentage error is determined by the following: (1) The amount
of Ag2CrO4 precipitated. As stated, this quantity is dependent
upon the individual observer and upon the conditions, such as back-
ground and lighting used, and under most conditions will be greater
than the value calculated above, thus leading to a larger positive
error. It was found that the sensitivity of the detection was appar-
ently somewhat increased in the presence of a white precipitate, but
that with too high a chromate concentration the color of the solution

7 If the hydrogen ion concentration of the solution is less than 10~ 7 , the frac-
tion of the chromate existing as OaO?" or HCrOi~ can be neglected for an
approximate calculation.

P. 95] T1TRATION OF CHLORIDE 399

begins to mask the first appearance of the precipitate. (2) The
chromate ion concentration. This is the most important factor
involved, as it can be controlled over a considerable range; and, as
can be seen, increasing its concentration causes an increase in the

chloride ion concentration and a decrease in the silver ion concen-

tration. Practically, the chromate concentration cannot be increased
beyond the point that its color becomes objectionable. (3) The
equivalents of chloride taken. This quantity is limited by the vol-
ume of solution required for the titration. It is apparent that the
error becomes quite serious if 0.01 n. solutions are used, as the term
Eq.Cl (in the denominator of Equation 3) becomes smaller for a
given volume of standard solution.

Practically, the error is most effectively reduced by applying an
"end-point correction.' 7 This consists in subtracting from the silver
nitrate required for the titration a volume equivalent to the end-
point error, that is, referring to Equation 3, (Eq.Ag 2 Cr0 4 + Eq.Ag*)
- (Eq.CF).

In the case considered, the correction would be 1,94 X 10 6
equivalent of silver nitrate, corresponding to 0.02 ml of 0.1 n. AgNO 3
solution. Frequently such an "end-point correction" is approxi-
mated experimentally by adding the silver nitrate solution to a
solution containing the same concentration of chromate as was used
in the titration until a precipitate comparable to that occurring in
the titration is obtained. The volume of the silver nitrate thus
used represents the correction applied.

The accuracy of the method is also dependent upon the hydrogen
ion concentration, which should be between 5 X 10~ 7 and 10" 10 .
The upper limit is fixed by the solubility of silver chromate in acid
solutions and the lower limit by the solubility of silver oxide in
alkaline solutions. As the solubility of silver chromate increases
rapidly with the temperature, the solution should not be much above
20C. during the titration.

A further factor which affects the accuracy of the method and
which is not considered in the theoretical calculations shown above
is the tendency of a silver chloride precipitate to adsorb chloride
ion and thus cause, a premature end-point. This is minimized by
vigorously shaking the solution as the end-point is approached and
noting if the first color obtained is permanent. 8 This Mohr titration

8 Meldrum and Forbes, J. Chem. Ed., 5, 205 (1928), advocate boiling the
solution after obtaining a preliminary end-point, cooling, and then titrating
to a permanent end-point.

400 ANALYSIS OF THE ALKALI GROUP [P. 95

can be used for the determination of bromide as well as chloride,
but, because of adsorption effects, it is not suitable for the deter-
mination of iodide or thiocyanate.

The Use of Adsorption Indicators in Titrating Chloride with Silver
Ion. The conditions are ideal here for the use of an adsorption
indicator, as the chloride is present as a neutral solution of pure
sodium chloride; because of this and because difficulty is sometimes
experienced in detecting the silver chromate end-point, the use of
an adsorption indicator is recommended for this titration. As was
stated in the discussion of these indicators in P. VI, their usefulness
is limited in the presence of a high concentration of neutral salt or
hydrogen ion.

Fluorescein is used here because it gives an excellent end-point and
is commonly available; it can be used satisfactorily only when the
pH of the solution is between 7 and 10. If it is desired to carry out
the titration in an acid solution (up to pH 3), dichlorofluorescein may
be used. Eosin (tetra brom fluorescein) may be used for the titra-
tion of bromide, iodide, or thiocyanate in solutions with a pH up
to 3 but, as it is appreciably adsorbed before the equivalence-point,
it cannot be used for titrating chlorides.

Procedure 95: ESTIMATION OP SODIUM. Gravimetric
Method. Treat the crucible and precipitate as directed in
P. 93. Calculate the amount of sodium present from the
weight of sodium chloride found (Note 1).

Volumetric Method. Dissolve the precipitate by pouring
dropwise through the filter 10 to 35 ml of hot water and
then collecting the solution in a 200-ml flask. Cool the solu-
tion.

If less than 100 mg of sodium are thought to be present,
treat the solution as directed in either of the last two para-
graphs of this procedure.

If more than 100 mg of sodium are thought to be present,
transfer the solution to a 100-ml volumetric flask, dilute it
to the mark, mix it well, and pipet into a 200-ml conical
flask that volume which is thought to contain 50 to 75 mg
of sodium.

Using Chromate as Indicator. Dilute the solution to
50 ml; if it reacts acid to litmus, make it just neutral with
1 f. NaHCOa (Note 2), and add to it 5 drops of 3 n. K 2 Cr0 4 .

P. 96] DETECTION AND ESTIMATION OF AMMONIA 401

Titrate the solution slowly with 0.1 n. AgN0 8 , swirling
vigorously, until the first permanent reddish precipitate is
produced (Note 3). From the volume of standard AgN0 3
used, calculate the amount of sodium present (Note 4).

Using an Adsorption Indicator. Dilute the solution to
50 ml; if it reacts acid to litmus, make it just neutral with
1