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INTERNATIONAL CHEMICAL SERIES
JAMES F. NORRIS, PH.D., CONSULTING EDITOR

SEMI-MICRO
QUALITATIVE ANALYSIS

A SELECTION OF TITLES FROM THE
INTERNATIONAL CHEMICAL SERIES

JAMES F. NORRIS, PH.D., Consulting Editor

Adkins and McElvain

Elementary Organic Chemistry
Practice of Organic Chemistry

Arthur and Smith

Semi-Micro Qualitative Analysis
Bancroft

Applied Colloid Chemistry
Britcoe

Structure and Properties of Matter
Burrcll

Chemistry for Students of Agriculture
and Home Economics

Cadv

General Chemistry
Inorganic Chemistry

Coghill and Sturtevant
An Introduction to the Preparation and
Identification of Organic Compounds

Daniels

Mathematical Preparation for Physical
Chemistry

Daniels, Mathews and Williams
Experimental Physical Chemistry

Detha

Organic Chemistry

Dole

Experimental and Theoretical Electro-
chemistry

Eucken, Jette and La Mar-
Physical Chemistry

Oillespie

Physical Chemistry

Griffin

Technical Methods of Analysis

Hamilton and Simpson
Calculations of Quantitative Chemical
Analysis

Hammett

Solutions of Electrolytes
HenderMon and Femeliue

Inorganic Preparations
Lnghou

Chemistry of Engineering Materials
Long and Anderson

Chemical Calculations

Mahin

Introduction to Quantitative Analysis
Quantitative Analysis

Mellon

Chemical Publications
Millard

Physical Chemistry for Colleges
Moore

History of Chemistry
Morton

Laboratory Technique in Organic Chem-
istry

Norris

Experimental Organic Chemistry
Inorganic Chemistry for Colleges
The Principles of Organic Chemistry

Parr

Analysis of Fuel, Gas, Water and Lubri-
cants

Reedy

Qualitative Analysis for College Stu-
dents

Rieman and Neuss

Quantitative Analysis
Robinson

Elements of Fractional Distillation
Schmidt and Allen

Fundamentals of Biochemistry, with
Laboratory Experiments

Stock and Stahler (translated by Patnode
and Dennis)

Quantitative Chemical Analysis
Stone and Dunn

Experiments in General Chemistry

Thomas
Colloid Chemistry

Timm

An Introduction to Chemistry

Wilkinson

Calculations in Quantitative Chemical
Analysis

Williams and Homerberg
Principles of Metallography

Woodman
Food Analysis

a
2

SEMI-MICRO
QUALITATIVE ANALYSIS

BY
PAUL ARTHUR, PH.D.

Assistant Professor of Analytical Chemistry
Oklahoma Agricultural and Mechanical College

AND

OTTO M. SMITH, Pn.D.

Professor ^oJ^CKemistry
Oklahoma Agricultural and Mechanical College

FIRST EDITION

McGRAW-HILL BOOK COMPANY, INC.

NEW YORK AND LONDON
1938

BOOK COMPANY, INC.

PRINTED IN THE UNITED STATES OF AMERICA

parts thereof, may not be reproduced

in any form without permission of

the publishers.

THE MAPLE PRESS COMPANY, YORK, PA.

PREFACE

One of the recent advances made in analytical chemistry,
from the standpoints of both education and practice, is
the development of the semi-micro technique. Although
its spread was greatly retarded by scarcity of textbooks
using semi-micro methods, its marked advantages lower
expense to the student, increased accuracy, greater skill
and respect for cleanliness, and the ability to cover more
and varied work in the allotted time have led many
progressive educators to adopt it.

The purpose of this book is to present a system of semi-
micro qualitative analysis in such manner as to make it
clear, easily understood, and, as far as possible, self -adminis-
tering. No teacher need hesitate to adopt the method,
for any student or instructor can readily acquire the skills
necessary in the handling of small amounts and the identifi-
cation of the product of the reactions.

The scheme used is little different from the customary
macro procedure, only the technique and a few of the
reagents being changed. Little actual dependence is
placed on organic reagents, as the authors feel that inor-
ganic reactions, which are more easily understood by the
student, should be emphasized. Some use of the better
organic reagents is made, however, to acquaint the student
with the advantages and disadvantages of the most depend-
able ones. Thus, while the characteristic reactions of
the ions and the laws of chemistry receive chief emphasis,
the practical aspects are not overlooked.

Since qualitative analysis in most schools is taught as
part of general chemistry in the second semester or the
third quarter of the freshman year, this text purposely

vi PREFACE

has been kept small and the treatment of the theoretical
portion very simple and brief. The discussion in the
earlier part is chiefly qualitative, and develops the theory
in terms of the older, standard conceptions. This is
followed by a section dealing with some of the more modern
theories, presented in such manner as to make it easy for
the student to apply the newer conceptions to the theories
studied previously. Care has been used to keep that
balance between theory and technique which is desired
for the average student in qualitative analysis. Extensive
use has been made of supplementary notes to maintain a
close connection between laboratory and theory.

The derivations of various formulas have been largely
omitted in view of the fact that in many schools the
instructors prefer to leave such derivations to be considered
in physical chemistry. In such cases, inclusion of these
derivations serves only to confuse the students.

It has been the experience of the authors that students
often get different types of problems confused when they
are considered separately. For this reason, the methods
of working problems are considered collectively in a section
following the discussion of the theory. This policy has
the advantage of enabling the students to compare the
various types of problems and the added advantage of
enforcing a review of the theories underlying the problems,
at a time when it is most profitable to the student.

The present book evolved from the experiences and
criticisms of both teachers and students during its use in
manuscript form for seven semesters at Oklahoma Agricul-
tural and Mechanical College, where it has been used in the
instruction of students of the freshman and sophomore
levels. In spite of its small size, the book contains suffi-
cient information, technique, and instructions to permit
combining it with any standard reference book for a longer
course when desired. Sufficient source material and
references have been included to maintain the interest

PREFACE vii

of the more zealous students. A careful study of the
references listed would give any investigator an excellent
background for research in this field.

The authors wish to acknowledge the debt they owe to
Dr. Paul Spoerri, who, through his course at Polytechnic
Institute of Brooklyn, demonstrated that semi-micro
technique could be used successfully with the ordinary
analytical scheme by freshmen and sophomore students
and whose notes were the source of many valuable ideas.
Also grateful acknowledgment is accorded Dr. 0. C.
Dermer, Dr. H. M. Trimble, and other colleagues for
helpful material, criticisms, and suggestions.

PAUL ARTHUR,
0. M. SMITH.

STILLWATER, OKLAHOMA,
January, 1938.

CONTENTS

PAGE
PREFACE v

INTRODUCTION

PART I
SEMI-MICRO METHODS

SEMI-MICRO TECHNIQUE . 3

Filtrations Handling of precipitates Evaporations Tests for
gases.

CONSTRUCTION OP APPARATUS 8

Microbeakers Centrifuge tubes Hydrogen sulfide generators
Capillary delivery tubes Reagent droppers Stirring rods and
spatulas Reagent dropper bottles Gas evolution apparatus.

PART II
THEORY OF ANALYSIS

IMPORTANT PRINCIPLES .... 15

Introduction Concentration units Calculations involving nor-
malities General principles Electrolytes and non-electrolytes
Strong and weak electrolytes Hydrolysis Water and pH values
Solubility product; Supersaturation Colloidal solutions Dis-
solving precipitates Complex ions Amphoteric hydroxides
Oxidation and reduction Balancing oxidation-reduction
equations.

MATHEMATICAL RELATIONSHIPS ... .60

Problems involving weak electrolytes Problems involving solu-
bility products.

MODERN THEORIES OF ELECTROLYTES . . . . 59

Hydronium ion Salt effect; the Debye-Hiickel theory Coprecip-
itation phenomena.

ORGANIC COMPOUNDS IN ANALYSIS 67

Organic reagents.

X CONTENTS

PAOB

PART III
ANALYTICAL PROCEDURE (NATIONS

GROUP SEPARATION CATIONS . . 71

Treatment of solid samples Separation of a general unknown into
groups.

GROUP I METALS . .... 75

Chemical characteristics Analytical aspects Preliminary experi-
ments Group I analysis.

GROUP II METALS ... . 82

Chemical characteristics Analytical aspects Preliminary experi-
ments Group II analysis.

GROUP III METALS. . . 97

Chemical characteristics Analytical aspects Preliminary experi-
ments Group III analysis.

GROUP IV METALS . . . .113

Chemical characteristics Analytical aspects Preliminary experi-
ments Group IV analysis

GROUP V METALS . . 119

Chemical characteristics Analytical aspects Preliminary experi-
ments Group V analysis.

AUXILIARY TESTS .... 124

Silver Mercury Lead Copper Bismuth Cadmium Arsenic
Tin Antimony Aluminum Chromium Zinc Cobalt
Nickel Manganese Calcium Sodium Potassium Ammo-
nium Magnesium .

REACTIONS INVOLVED IN SEPARATIONS .... . . 136

QUESTIONS . . 140

PART IV
ANALYTICAL PROCEDURE ANIONS

PRELIMINARY TREATMENT OF SOLIDS 142

PRELIMINARY TREATMENT OF SOLUTIONS 144

GBOUP ANALYSIS 146

CONTENTS xi

PAGE
EQUATIONS FOR REACTIONS OF ANIONS 156

PROBLEMS . . . 159

NOTES ON ANALYTICAL PROCEDURE . 163

REFERENCES ... 176

APPENDIX . .... ... 179

Table I, Oxidation-reduction potentials- Solubility product con-
stants lonization constants List of apparatus List of reagents
Test solutions and unknowns.

INDEX 193

SEMI-MICRO
QUALITATIVE ANALYSIS

INTRODUCTION

.Too often it happens that a student enters a new course
having only a vague idea of what the course is about and
for what he will be held responsible. This is especially
true of the usual course in qualitative analysis. Quali-
tative analysis is far more than a course in analytical
procedure, and, for the student to get the most value from
the course, it is important that he understand, from the
first, what he is expected to learn. The most important
of the things that should be gained during the course are:

1. Experience in the handling of certain types of appa-
ratus rapidly and efficiently.

2. Experience in making observations and drawing
correct conclusions from them.

3. A greater knowledge of the chemistry of metals,
cations and anions, and of the laws and principles of
chemistry.

4. An appreciation of the great necessity for cleanliness
of apparatus and purity of reagents in all chemical reactions.

5. A knowledge of the methods used in carrying out
practical analyses of samples whose chemical composition
is not known.

It will be noticed that the practical knowledge of
analytical methods is perhaps the least important to most
chemists since an analysis can be made such routine
procedure that little knowledge of chemistry is required for
its use. If some unusual problem comes up, however, a

2 INTRODUCTION

knowledge of the laws and of the chemical characteristics
of different ions becomes exceedingly important. In this
connection, it will be found that a review of inorganic
chemistry will enable the student to work more efficiently
and will eliminate to a large extent the unfamiliarity which
accompanies entrance into a new course. Special emphasis
should be placed on the laws of ionization and ionic reac-
tions and on the chemical properties of the metals.

The complete analysis of a substance consists in the
determination of the chemical composition of that sub-
stance. The first step in any analysis is to determine the
nature of the elements or ions present. This process is
known as qualitative analysis. Quantitative analysis con-
sists of determining the weights of the constituent present
in a given sample of the substance.

PART I

SEMI-MICRO METHODS
SEMI-MICRO TECHNIQUE

During the last few years much effort has been expended
in attempts to develop apparatus and special technique for
the analysis of very small quantities of material. Dealing
with samples of only 3 to 5 mg. (compared with the usual
0.5 to 1.0 g. macro sample), the micro method consumes
much less time for reactions to become complete, requires
much smaller quantities of reagents, requires much less of
the sometimes rare and valuable unknown, occupies much
less laboratory space, and in many other ways proves
superior to the macro methods.

Unfortunately, before a student can successfully under-
take the study of the real micro analysis, he must be experi-
enced in the handling of delicate precise instruments and
must have a fair knowledge of chemistry. For this reason,
a technique which approaches micro methods and incor-
porates many of their advantages, yet remains sufficiently
simple to be used by students in the second year of chem-
istry, has been devised. This procedure, known as semi-
micro qualitative analysis, makes use of the simpler micro
methods but works with samples about one-twenty-fifth
as large as the macro samples (or about forty times as large
as a micro sample).

The unknown may be encountered in solution or as a
solid. An analysis consists chiefly of a series of precipi-
tations and re-solutions. Hence, it is necessary that the
unknown be placed in solution (unless already dissolved)
before the analysis can start. These steps require the

SEMI-MICRO METHODS

careful use of special methods, and it is essential that skill
in these techniques be developed early in the course.
Many analyses fail because the student is awkward in
carrying out separations of solids and liquids or because
evaporations are carried too far. Each of the more
important procedures will be discussed under its own
heading.

Filtrations. In semi-microanalysis the usual filter paper
and funnel are not used. Instead, the mixture of liquid

..- - Micro - beaker
and wafer as
counterbalance

Tightly
packed
cotton

1 -Centrifuge.

and solid is placed in a cone-shaped centrifuge tube or a
microbeaker, and the tube is placed in one of the cups of a
centrifuge (Fig. 1). Another tube the same size as the
first, containing a volume of water equal to the volume of
material in the first tube, is placed in the centrifuge exactly
across the head from the first tube. Thus balanced, the
centrifuge will have a minimum of vibration. If a hand
centrifuge is used, it must be turned with sufficient velocity
that the centrifugal force will pack the precipitate in the

SEMI -MICRO TECHNIQUE

<& 3)

tip of the cone-shaped tube leaving the liquid above the
precipitate clear. The centrifuge should always be allowed
to slow down to a stop of its own accord. Any attempt to
quickly stop the instrument may result in damage to the
centrifuge. The electrically driven centrifuges are so
constructed as to provide for the necessary velocity, auto-
matically. Several types of electrically driven centrifuges
are on the market, many of which have variable speed
control and should be operated according to instructions
accompanying the instrument.

The filtrate may be removed by
means of an ordinary medicine
dropper. Into the tip of the
dropper is inserted a tight twist
of absorbent cotton which is cut
off i- in. below the glass tip with
scissors. This serves as a filter.
The bulb of the dropper is com-
pressed, the tip inserted into the
liquid, and the bulb slowly released
so as to draw the liquid into the
dropper tube. The dropper is
then withdrawn, the cotton plug
removed, and the liquid transferred
to a clean micro beaker (Fig. 2).

Occasionally the filtration may be conducted in another
manner. The solution is given a preliminary centrifuging
after which a very loose cotton plug is placed in the centri-
fuge tube and pushed just below the surface of the liquid.
The tube is replaced in the centrifuge and the latter turned
as described earlier. When the tube is removed, the cotton
will be found packed firmly over the surface of the precipi-
tate and the medicine dropper can be used without danger
of drawing up any of the precipitate.

Handling of Precipitates. It is often necessary to trans-
fer a precipitate from one container to another. For this

^--- Cotton N

Fid. 2 Separating filtrate
sirid precipitate after centri-

SEMI-MICRO METHODS

purpose a small glass spatula (made as described later) is
most suitable. If a liquid reagent is to be added, it can
often be used to wash the precipitate from one container to

.
Pyrex beaker

Porce/aiin
cruc/b/e

wire
triangle benf
+o form fnpoo/

FIG. 3.- -Air ha tli

Hole in If of +o
hotel corks.

^ r Corks bored +o hold
micro-bee* ker

Si-earn out/ef
& net hcmof/e

FIG. 4. -Steam hath made from cut-down tin can.

another by agitating the mixture and withdrawing it with
a medicine dropper.

Evaporations. Evaporations are usually carried out by
placing the liquid in a crucible or a hard-glass microbeaker

SEMI-MICRO TECHNIQUE 7

and heating it in an air bath (Fig. 3). No soft-glass appa-
ratus should ever be used for this purpose unless the evaporation
is to be carried out on the steam bath (Fig. 4) as soft glass will
often crack and spill the liquid.

To remove hot containers from the baths, forceps or
small tongs are convenient. However, since these imple-
ments are generally made of metal which may be corroded,
they may contaminate the unknown unless care is used.

Tests for Gases. In testing for ammonium, arsenic, etc.,
the identification depends on converting the ion into a
volatile compound and passing the gas over a bit of filter
paper impregnated with the proper reagent to produce a
color change. In doing this, the specified amount of the
solution to be tested is placed in the gas evolution apparatus
(see description of this apparatus, and Fig. 8) and the
necessary reagent added. The stopper is inserted, and a
bit of filter paper impregnated with the proper reagent is
placed in the attached tube. The mouth of the tube is
closed with a loose plug of cotton to prevent the entrance
of air, and the test tube is warmed. The results can be
seen through the side of the glass tube.

Spot Tests. In many cases the identification of a
metallic ion will depend on the formation of a precipitate
when a drop of the unknown solution is tested by a drop of
each of the necessary reagents. Sometimes the resulting
precipitate is difficult to see on account of the low concen-
tration of the metallic ion. These tests are best carried out
on a watch glass, spot plate, or microscope slide. The
proper background is of great importance in rendering a
faint precipitate visible. If the precipitate is light in color
(white or yellow), a black background (such as a micro-
scope slide painted black on the reverse side, or a piece of
black glass) should be used. If the precipitate is dark
(red, black, gray, etc.), a white background should be used.
A spot plate is excellent for this type of precipitate.

SEMI-MICRO METHODS

CONSTRUCTION OF APPARATUS

All equipment needed in semi-micro work can be pur-
chased ready for use from any of the larger chemical
apparatus supply houses. As yet, however, the small
demand has kept prices at higher levels than usual. Con-
sequently, instructions are given here for the conversion
of standard equipment into microapparatus. The smaller
sizes of macroapparatus may often be used unchanged, and

(I) (2) (3)

Fid. 5. Steps in making microbeakers.

most of the other things needed can be constructed by the
student from standard equipment. Little knowledge of
glass blowing is needed as the student can acquire the
necessary skill with very little practice.

Microbeakers. In place of commercial microbeakers,
light wall Pyrex test tubes, 75 by 10 mm. diameter, are
recommended. Those desiring to use microbeakers of
greater diameter may construct them from test tubes as
follows (Fig. 5): With a sharp file, make a scratch com-
pletely around a 4 by ^ in. test tube about 2^ in. from the
bottom. Heat the mouth of the test tube in the flame till
it is rather hot then touch it with a drop of water. This
will crack the top of the test tube. With a fine wire gauze,

CONSTRUCTION OF APPARATUS

stroke the broken end of the test tube at a 30 angle, and
thus polish the tube down to the scratched line (a hot wire
glass cutter is more convenient if it is available). Heat the
top of the little beaker thus formed, slowly rotating it in
the flame until it is relatively soft, and flare the top of the
beaker into a rim with the tip of a file or a stick of graphite.
About \y% doz. microbeakers are desirable.

Centrifuge Tubes. These are small containers with cone-
shaped bottoms. They may be purchased ready-made or

(I) (2) (3^ (4)

Fio. 6. Steps in making centrifuge tubes. ^

may be made as follows (Fig. 6) : With a file cut off a 12.5-cm.
piece of 12-mm. soft-glass tubing. In the flame of a good
Bunsen or high temperature burner, heat a section 2.5 cm.
in length exactly in the center of the tube, rotating the tube
slowly in the flame to insure uniform heating. When the
tube is extremely soft, remove it from the flame and quickly
but steadily pull the ends of the tube straight out until

10 SEMI-MICRO METHODS

the softened portion is only half its original diameter.
Allow it to cool to rigidity, then exactly in the center of the
constricted portion heat again, rotating as before, pulling
the ends of the tube a little at first to narrow the con-
striction, and finally pulling the tube in two. Seal the two
tips thus formed by holding each in a flame until the
glass melts into a small drop. If this is done carefully, the
inside of the tube will be rounded at the bottom. Pointed
bottoms should be avoided as they are very difficult to
clean. The sharp edges of the top should be fire-polished
and flared as was done in making the microbeakers. About
}<2 doz. tubes will be needed.

Hydrogen Sulfide Generator. To minimize delays and
reduce the quantity of hydrogen sulfide escaping into the
air, it is advisable that each student prepare his own
generator and keep it in his desk. This is easily constructed
from a widemouthed 8-oz. bottle which is about 5 or 5 \^ in.
high, a No. 9 or 10 rubber stopper, a 6 by -''4 in. test tube, a
one-hole No. 2 stopper, and some glass tubing. To one
side of the center of the rubber stopper, drill a hole into
which a 6 by ? in. test tube will fit closely, and on the
other side drill a small hole about 2 mm. in diameter to act
as an air vent. Now heat the extreme bottom of the test
tube in a flame till soft, and, from the inside with a long
file, push the bottom out into a bulge. Heat the bulge
until it is very soft, and, removing it from the flame, quickly
blow into the mouth of the test tube until the bulge forms a
large bubble. Knock this off with a file. There should
then be a hole in the bottom of the tube about 3 mm. in
diameter. It should be neither smaller nor much larger.
If larger, heat the bottom of the tube once more and press
the edges of the hole inward with a file. Let the test tube
cool ; then wet it with water, and push it through the stopper
so that, when the stopper is placed in the bottle, the end of
the tube will be about 2 mm. above the bottom of the bottle.
Next, fit the top of the test tube with the one-hole stopper

CONSTRUCTION OF APPARATUS

and a small right-angle glass tube. Now cut a piece of
thick- walled 3-mm. glass tubing 15 cm. long, seal one end in
a flame, and blow a bulb (about 7 mm. in diameter) in the
center. Cool, then cut the tube off about 1.5 cm. on each
side of the bulb. By means of a 1.5-cm. piece of rubber
tubing, attach this bulb to the glass right-angled tube.
Fill the bulb with a loose plug of cotton; then attach a

Plug of cotton -

2mm.
d/'am.

'\o

Pinch clamp **'

Delivery tube''

- ' 4 oz wide - mou th bottle
5"x 3 /4 "test tube

6NHC1
~ ^~^\~.^~~^s<"4 cm stick feS

: n: -G/ass tubing

v 2mm. hole
FIG. 7. Hydrogen sulhde generator.

6-in. piece of rubber tube (delivery) equipped with a small
screw clamp (see Fig. 7).

To operate the generator, put a couple of small bits of
broken glass tubing into the 6-in. test tube, followed by a
4-cm. stick of ferrous sulfide. Fill the bottle half full of
QN hydrochloric acid and replace the stopper and its
fittings. When hydrogen sulfide is desired release the
clamp on the delivery tube; close the clamp when it is
desired to stop the flow of gas.

An alternative method of generating hydrogen sulfide is
by heating a mixture of 1 part paraffin and 3 parts of

12 SEMI-MICRO METHODS

powdered sulfur by weight with enough medium grade
asbestos to make the mass porous. This requires only a
15-cm. Pyrex test tube (thin wall) fitted with a one-hole
rubber stopper and delivery tube. The gas generation
stops almost at once when heating is discontinued. How-
ever, as the apparatus sucks back on cooling, it is best to
have a T tube in the delivery tube, and the finger tip should
be held over the side outlet 'while using the generator and
removed at once when heating is stopped.

A commercial mixture prepared for this type of generator
and sold under the trade name "Aitchtuess" can be pur-
chased at very moderate prices from any of the standard
chemical supply houses.

Capillary Delivery Tubes. The delivery tube used with
the hydrogen sulfide generator must be of capillary dimen-
sions in order to give the small bubbles desired these being
more completely absorbed than large bubbles. These
delivery tubes are made by heating a 3-mm. glass tube for
3 to 4 cm. of its length until very soft, then removing it from
the flame and quickly, but steadily, drawing it out into a
capillary of about 1 mm. diameter. Finer capillaries are
not convenient because they clog too easily. When cool,
the tube is cut so that about 1 cm. of the original 3-mm.
tube is left attached to about 7.5 cm. of capillary tubing.
The thick end is used for attaching the capillary to the end
of the rubber delivery tube on the generator and the capil-
lary end for inserting into the liquid. At least six of these
tubes should be prepared.

Reagent Droppers. In each desk there should be 1 doz.
ordinary medicine droppers. Take two of these, and, using
the technique described for the capillary delivery tubes,
draw their tips out to capillaries, cutting off the capillary
tip at a point where its external diameter is about % mm.
so as to leave the total length of the glass about 9 cm.
These two droppers are very handy for transferring small
quantities of reagents from one vessel to another.

CONSTRUCTION OF APPARATUS 13

Capillary pipettes are also useful for this same purpose.
These pipettes are capillary tubes about 2 mm. in external
diameter and 10 cm. long and are best prepared from 5 mm.
soft-glass tubing as described above. The ends are
cautiously fire-polished in a small flame. To use them,
place one end of the tube in the liquid, press the tip of a
finger over the upper end of the capillary, and lift the
tube out. To empty the tube, hold it with the open end
slanted downward, cautiously lift the finger tip off the other
end, and allow the liquid to fall, drop by drop, into the con-
tainer. About five of these pipettes should be made.

Stirring Rods and Spatulas. Stirring rods are made by
cutting 5 in. lengths of 2-mm. glass rod, drawing one end out
to 1 mm. diameter, and fire-polishing both ends. If no
rods of this diameter are available, larger rods may be
drawn out or else capillaries, prepared as if for capillary
pipettes, may be sealed off at both ends and used as stirring
rods.

Spatulas may be made by heating the small ends of
stirring rods and pressing them flat between two flat
metallic surfaces. The flattened ends should be small as
they must reach the bottoms of the centrifuge tubes when
used for the transfer of precipitates. About six stirring
rods and four spatulas should be sufficient.

Reagent Containers. The usual liquid reagent con-
tainers have a capacity of, for those most used, about 30 ml.
and, for the others, about 2 to 8 ml. The 30-ml. con-
tainers, about 15 in number, may be the usual 1-oz. drop-
ping bottles equipped with rubber bulbs. These may be
placed in holes drilled in a wooden block to keep them in a
convenient unit. The other reagents, about 65 in number,
are in vials arranged in alphabetical order in a similar
block. Each vial is fitted with a one-hole cork through
which a 3-mm. glass tube extends to within 1 mm. of the
bottom of the vial, sufficient projection of the tube above
the cork being left to allow the tube to be used as a pipette.

SEMI-MICRO METHODS

A small V-shaped groove may be cut vertically in the cork
to act as an air vent.

Gas Evolution Apparatus. Using the narrow width of a
flame obtained from a wing-top burner, strongly heat a

Co / fon

Test paper -i

Cork

Test tube
or micro
beotkvr ->

Reaction
mix ture

FIG. 8. Gas e\olution apparatus.

piece of 10-mm. glass tubing, and draw it down to a diam-
eter of about 2 or 3 mm. Cut the tube so that about 2.5 cm.
of the narrow portion remains attached to 2.5 cm. of the
thicker portion. Put the narrow portion through a one-
hole rubber stopper fitted in a 4-in. test tube (Fig. 8).

PART II
THEORY OF ANALYSIS

IMPORTANT PRINCIPLES

Introduction. It is perfectly possible for a student to
perform the laboratory portion of a course in qualitative
analysis without knowing many of the laws and principles
of chemistry. The usual unknown is made up using con-
centrations that fall between certain definite limits, and the
instructions are so written as to give good results with such
unknowns. In practical work, however, the unknowns
may vary far beyond these limits. Consequently, it
becomes more and more necessary that the analyst know
the theories that were used in making up this procedure so
that he can vary his tactics whenever he finds it necessary.
A knowledge of the laws and principles is what makes the
difference between the good and the poor analyst.

As a rule, the analytical student makes use of these laws
in laboratory work long before they are reached in the
lecture. It is almost impossible to correlate the laboratory
and lecture properly. For this reason a brief, qualitative
review 7 of the more important principles will be given here
with the expectation that they will be expanded in the
lecture portion of the course.

Concentration Units. It has already been stated that
practically all of analysis deals with solutions. Analysis
is not absolute in nature. If a large granule of lead nitrate
is dissolved in 1 ml. of water, the presence of lead in the
solution is easily demonstrated by using the proper reagents.
If, however, 1 drop of this solution is added to 1000 liters
of pure water, the resulting solution will be so dilute that

16 THEORY OF ANALYSIS

we can not detect the presence of the lead although it is
certainly there. The only difference between these two
solutions is one of concentration, and we shall find that the
concentrations of reagents may have just as much to do
with the results of an analysis as do the concentrations of
the ions for which we are testing. In order, therefore, for
us to give and to understand instructions, it is important
that we define two units in terms of which we can con-
veniently express the concentrations of our solutions.

These are based on the number of mols or the number
of equivalent weights of the material in a liter of solution.
A solution that contains 1 gram-molecular weight (i.e.,
1 mol) of solute in 1 liter of solution is a \M (one molar)
solution; one that contains 2 gram-molecular weights of
solute in 1 liter of solution is a 2M (two molar) solution; etc.

A solution that contains 1 gram-equivalent weight of
solute per liter of solution is a one normal (IN) solution;
etc. For example, a solution containing 231.6 g. of lead
nitrate, Pb(NO 3 )2, per liter of solution is a IM (or 2N)
solution. A solution containing 36.5 g. of hydrochloric
acid per liter of solution is both a IM and a IN solution.
A solution of hydrochloric acid containing 18.25 g. per liter
is a half-normal (0.5AT) or a half-molar (0.5Af) solution.

To comprehend this, it is necessary to understand how to
determine the equivalent weight of any compound. In
most cases, it is only necessary to divide the gram-molecular
weight of the compound by the total valence of that radical
in the molecule that is most important to the reaction that
is to be carried out. For example, in acids, it is usually
the hydrogen ion that is the most important ; in bases, it is
the hydroxyl radical; in sodium sulfate, Na 2 SO 4 , it may
be the sodium or it may be the sulfate; in K 2 A1 2 (SO 4 )4, it
may be either the potassium, the aluminum or the sulfate
radicals. It all depends upon whether the substance is to
be used as a source of hydrogen, hydroxyl, sodium, alumi-
num, potassium, sulfate, or other ion. For example, using

IMPORTANT PRINCIPLES 17

the compound K 2 A1 2 (SO 4 ) 4 , if the solution is to be used for
precipitating some salt of potassium, the equivalent weight
is found by dividing the gram-molecular weight by two
since there are two valences represented by the potassium
in each molecule of the compound. If it is to be used as a
source of aluminum (as in the precipitation of aluminum
hydroxide), the gram-molecular weight must be divided by
six. As a source of sulfate radical, the gram-molecular
weight must be divided by eight since that is the total
number of valences represented by the sulfate radicals
present.

Other methods of expressing concentrations are some-
times used, e.g., grams per liter or per 100 ml. (milliliters)
of the solvent, percentage by weight or volume (parts per
hundred), parts per million by weight, etc. These are,
however, rarely used in qualitative analysis though they
are often used in quantitative work.

Calculations Involving Normalities. In the previous
section the method was given for the preparation of solu-
tions of given normalities from the pure substances and
water. Often, however, it happens that a solution of high,
known concentration is available but a solution of lower
concentration is needed. It is possible to dilute such solu-
tions to the lower concentration more rapidly than to
weigh out the required solute.

For example, let it be assumed that 24 drops of a 3N
solution of sulfuric acid is needed. The ordinary concen-
trated sulfuric acid is approximately 36N, and it is desired
to find out how to dilute it in order to get the concentration
and volume of acid needed. There is an important formula
which applies to all such problems, i.e.,

N, X V, = N, X F 2

where N i is the normality and Vi is the volume of the
original solution used; JV 2 is the normality, and F 2 is the

18 THEORY OF ANALYSIS

volume the solution will have after dilution. Using this
formula in the above problem,

N l = 36 N, = 3

V l = ? F 2 = 24 drops

36 x F! = 3 X 24

= 2 drops

Therefore,

This means that 2 drops of concentrated sulfuric acid must
be diluted to 24 drops (by adding water) to make 24 drops
of 3N sulfuric acid.

In the Appendix there is given a list of the normalities
of the common concentrated acids and bases for use in these
calculations. It must be remembered that these figures are
not exact, for the concentrations of the original concen-
trated reagents may vary in strength. However, the
results obtained are sufficiently close for use in qualitative
analysis though far too inaccurate for use in quantitative
work.

The above equation also applies to volumetric reactions
between ions in solution. For example, if it is desired to
calculate the volume of 2N sulfuric acid required to
neutralize 50 drops of 3N ammonium hydroxide, one again
substitutes in the formula,

V l = ? F 2 = 50 drops
Therefore,

2 X F! = 3 X 50
Fi = 75 drops

This gives the answer not only to this problem but also
tells us that 50 drops of any 3N base will require 75 drops
of any 2N acid for neutralization.

IMPORTANT PRINCIPLES 19

General Principles. Although many different laws and
phenomena will be studied during the course, it will be
found that there are two principles that will be applicable
to every case. These are the law of mass action and
Le Chatelier's principle.

The law of mass action states that in any chemical
reaction the mathematical product of the concentrations of
the products of the reaction divided by the product of the
concentrations of the reacting substances will, when the
reaction has attained equilibrium, equal a constant K
regardless of what concentrations were used at the begin-
ning. This constant will never vary unless the temperature
at which the reaction is carried out is changed. To illus-
trate this law let us assume that two substances, A and B
react to form C and D; e.g.,

1A + 2B ^ 2C + ID

If this reaction is allowed to stand until it has attained
equilibrium and an analysis is carried out for each sub-
stance, it will be found that the following relationship will
hold true

(O(C) X (JP) _ (C) 2 X (D) 1 _
(A) X (B)(B) ~ (A)* X (BY ~ A

where (A), (J5), etc., represent molar concentrations. If to
that same mixture some more of either A, B, C, or D is
added, equilibrium is established, and an analysis is again
carried out, it will be found that the same relationship
gives the same constant if the temperature is kept the
same as before.

Le Chatelier's principle states that, if a stress is applied to
a system in equilibrium, the equilibrium will shift in such
direction as to reduce the stress. In chemistry there are
many ways of applying stress to a system. Changes in
concentration, pressure, temperature, intensity of illumi-

20 THEORY OF ANALYSIS

nation, etc., are common examples. As an illustration let
us assume that the following reaction is taking place.

IN, + 3H 2 ^ 2NH 3 + heat energy

1 volume 3 volumes 2 volumes*

According to the law of mass action, the relationship

(NH 3 )(NH 3 ) (NH 3 ) 2

(N 2 ) X (H 2 )(H 2 )(H 2 ) - (N 2 ) X (H 2 ) 3 ~ A

will hold true. According to Le Chatelier's principle, if a
stress is applied to the system, the reaction will shift in such
manner as to remove the stress. If the mixture is heated,
the reaction will shift to the left so as to remove heat; if the
mixture is subjected to increased pressure, the reaction will
shift to the right since that is the direction which will
decrease the volume of the gases as well as the pressure.
Increasing the concentration of either hydrogen or nitrogen
will shift the reaction to the right, using up these gases in
the production of ammonia. On the other hand, increasing
the concentration of ammonia or removing nitrogen will
shift the reaction to the left, etc. Similarly, any chemical
equilibrium can be shifted by applying the proper stress.

Electrolytes and Nonelectrolytes. Soon after the de-
velopment of the voltaic battery, it was discovered that
certain substances in water solution are excellent con-
ductors of electricity. To distinguish this conductor from
metallic conductors, the former are known as electrolytic
conductors or electrolytes. In general, all salts, acids and
bases are electrolytes.

Another class of compounds is known as the nonelec-
trolytes. To this class belong such substances as sugar,

* The volume relationships indicated here follow from Gay-Lussac's
law of combining volumes. This law states that in reactions between two
or more gases, the ratio of the volumes of the gases undergoing reaction is expres-
sible in small, whole numbers. It should be reviewed by the student, using
any good inorganic text.

IMPORTANT PRINCIPLES 21

alcohol, acetone, etc., which are characterized by the fact
that their aqueous solutions are nonconductors of electricity.

The first successful attempt to explain the differences
between electrolytes and nonelectrolytes was made by the
Swedish chemist, Arrhenius (63)* late in the nineteenth
century. His theory postulated that the chief difference
between an electrolyte and a nonelectrolyte is that the
molecules of electrolytes break up, in solution, into two
or more radicals (atoms or groups of atoms which act as a
unit in a chemical reaction) each carrying its own charac-
teristic positive or negative charge. These charged par-
ticles, known as ions, act as the carriers of electricity when
a solution is undergoing electrolysis.

One of the most important characteristics of solutions
of electrolytes is expressed in the additivity principle. The
aqueous solution of any given electrolyte has properties
which are the sum of the properties characteristic of each
kind of ion present in the solution. For example, a
solution of cupric dichromate gives a yellow precipitate with
lead nitrate solution and a red-brown precipitate with
silver nitrate solution. It oxidizes acid solutions of
bromides giving bromine, and oxidizes many organic
compounds to form highly colored substances. These,
and many other properties of cupric dichromate solution,
are characteristic of solutions of all dichromates and may
therefore be considered the properties of dichromate ions.

Similarly, when cupric dichromate is treated with
sodium sulfide, a black precipitate forms; with excess
ammonium hydroxide a deep blue color forms; and with
K 4 Fe(CN) 6 solution a reddish solid precipitates. These
are characteristics of all solutions containing cupric ions.

These two sets of properties together make up the
properties of cupric dichromate. Even the brown color
of cupric dichromate solutions is the result of a combina-

* The numbers in the parenthesis throughout the text are code numbers
under which the references on the subject may be found in the Appendix.

22 THEORY OF ANALYSIS

tion of the blue of the cupric ions and the orange of the
dichromate ions.

Since the chemical properties of a solution of an elec-
trolyte are due chiefly to the individual ions present, it is
evident that chemical reactions between solutions of
electrolytes must be, essentially, reactions between ions.
To be accurate, therefore, reactions between such solutions
should be written in terms of ions instead of molecules.
For example, the reaction between silver nitrate and
sodium chloride solutions to form insoluble silver chloride
is usually written

AgNO 3 + NaCl -* AgClj + NaNO 3
It is more properly written
Ag+ + NO 3 - + Na+ + Cl- -> AgCU + Na+ + NOr

This equation shows that the only reaction occurring is
one between silver and chloride ions to form silver chloride.
The sodium and nitrate ions remain in the solution,
unchanged.

The above equation is called an ionic equation.

Strong and Weak Electrolytes. Experimentally it is
found that electrolytes vary as to the ease with which they
conduct electricity. Hydrochloric acid, for example, is an
excellent conductor of electricity while acetic acid is much
poorer.* This is owing to the fact that while some elec-
trolytes are highly ionized, i.e., most of the molecules
placed in solution are broken up into ions; others, like
acetic acid, ionize to a much less extent in other words,

* Recent theories by Debye, Hiickel, and others postulate that such
strong electrolytes as sodium chloride and potassium chloride are completely
ionized, even in the most concentrated solutions. Although concentrated
solutions of strong electrolytes behave as if the solute were only partly
ionized, modern theories explain this on the basis of probable mutual inter-
ference of ions of opposite charge on each other's movements. This theory
will be discussed more fully later (p. 63).

IMPORTANT PRINCIPLES 23

only a few of the molecules ionize, the rest remaining in
solution as whole molecules.*

It has been found that both the law of mass action and
Le Chatelier's principle apply to solutions of weak elec-
trolytes. Taking the weak electrolyte, acetic acid,
HC 2 H 3 O 2 , for example, we set up the following equation
(letting Ac stand for C 2 H 3 O2 and Ac"" for acetate ion) :

HAc H+ + Ac-
According to the principle of Le Chatelier, if we add Ac"
to the solution, we should cause H+ and Ac"" to combine,
the result being that the hydrogen ion concentration should
diminish and the HAc concentration (the concentration of
acetic acid molecules) should increase. This is exactly
what occurs if we add sodium or ammonium acetate to the
solution. The addition of H+ (which can be accomplished
by adding any strong acid, such as HC1) should cause a
decrease in the concentration of Ac" and again an increase
in the concentration of HAc molecules. These are examples
of the common-ion effect.

This is further brought out by applying the law of mass
action.

In the expression

(H+) X (Ac")

(HAc)

= K a

(K a for HAc at 18 C. is 1.8 X 10" 5 ) it is easily seen that if
K a (the ionization or dissociation constant} is to remain the
same for all such solutions of acetic acid, an increase in the
concentration of Ac" must result in a decrease in the

* This is true only of ordinary solutions, however. As the concentration
of any weak electrolyte is lowered, it is found that the conductivity of the
solution (measured in a cell arranged to keep the whole solution between the
electrodes) is increased, reaching a maximum at infinite dilution. Arrhenius
explained this by assuming that at infinite dilution ionization is always
complete, i.e., in such solutions none of the solute is present in the form of
molecules. Later experiments have verified this conception.

24 THEORY OF ANALYSIS

concentration of H^ and an increase in the concentration
of HAc. If this were not true, the fraction would have a
larger value than K a which is not the case.

The same relationship holds true for weak bases.
NH 4 OH, for example, dissociates as follows:

NH 4 OH =; NH 4 +
Setting this up in the form used for acetic acid,

(NH 4 OH)~ ~ b

The same rule will apply to this equilibrium as applied to
the acetic acid equilibrium. The dissociation (or ioni-
zatiori) constant for NH 4 OH at 18C. is 1.75 X 1(T S .

To increase the concentration of one ion in a solution
without adding the other ions also, it is only necessary to
choose a highly soluble strong electrolyte which contains
the ions desired and add it to the solution. As salts in
general are, with very few exceptions, strong electrolytes,
they are the most commonly used for this purpose.

A solution of an acid containing a high concentration
of a salt of a weak acid or of a base containing a high con-
centration of a salt of a weak base is called a buffer solution.
The value of such solutions lies in the fact that fairly large
quantities of acids or bases can be added to them without
much change in the hydrogen or hydroxyl ion concentration
of the solution. The concentrations of these ions will be
low, but quite constant. This makes such solutions invalu-
able where changes in the acidity of the solution are
undesirable and much use is made of buffering agents
(such as ammonium acetate, ammonium chloride, sodium
acetate, etc.) throughout the analytical procedure.

Strong electrolytes are somewhat different from weak
electrolytes. Without great error, 1 gram-molecular weight
of HC1 dissolved in water may be considered to break up

IMPORTANT PRINCIPLES 25

completely, even in concentrated solutions, to give 1 gram-
molecular weight of H+ (1 gram-ion of hydrogen) and
1 gram-molecular weight of Cl~". On the basis of this
assumption, the reaction goes to completion as indicated
by the equation

HC1 -> H+ + Cl-

Consequently, a \M solution of HC1 contains 1 gram-ion
per liter each of H+ and Cl~. Similarly, a IM solution of
NaOH (a strong electrolyte) contains 1 gram-ion per liter
each of Na+andOH-*

The case of a strong di- or tri-basic acid or base is not
quite so simple, however. These substances ionize in
steps, in the first step acting as strong electrolytes but in
the latter stages acting as increasingly weak electrolytes.
Sulfuric acid, for example, ionizes in the first stage as
follows, the ionization going to completion:

H 2 SO 4 - H + + HSOr ttrisulfate ion)

The bisulfate ion then undergoes ionization, this step,
however, being reversible as in the case of weak electrolytes:

HSOr ^ H+ + S0 4 ~

Consequently, sulfuric acid exhibits properties intermediate
between those of strong and of weak electrolytes.

In general, experiment has shown that salts, with the
exception of compounds of lead, tin, mercury, some zinc
compounds, and one or two less known types, are all strong
electrolytes. The hydroxides of the alkali and the alkaline
earth family, and the three common acids, sulfuric, nitric,
and hydrochloric, are also strong electrolytes.

* Tables of apparent degrees of ionization as calculated from conductivity
measurements list strong electrolytes such as hydrochloric acid and sodium
hydroxide as being only 90 per cent ionized in 0.1 M solutions. The more
modern theories, however, justify treating them as being completely
ionized. (See Debye-Hlickel theory.)

26 THEORY OF ANALYSIS

On the other hand, ammonium hydroxide, acetic acid,
hydrosulfuric acid, benzoic acid, and the salts noted above
as exceptions are all classed as weak electrolytes. It must
be remembered, however, that there is no sharp line of
division between strong and weak electrolytes, as com-
pounds are known ranging through all degrees from the
weak to the strong.

Hydrolysis. One of the important consequences of the
characteristic reactions of weak electrolytes is found in the
way certain acids and bases react with each other as well
as the way certain salts react with water.

Neutralization is commonly defined for beginning students
as being the reaction between any acid and base to form a
salt and water. On writing the ionic equation for a
neutralization, however, it soon becomes evident that the
reaction is more accurately defined as being a reaction
between hydrogen ions and hydroxyl ions to form water,
e.g., the reaction between HC1 and NaOH is

Na+ + OH- + H+ + Cl-= Na+ + Cl~ + H 2 O

slightly ionized

Even in pure water, however, ionization takes place to
some extent, forming H + and OH~~. In pure water the
concentration of H+ is 10~~ 7 g. or 10~ 7 mol per liter; while
that of the OH~ is 17 X 10~ 7 g. or 1(T 7 mol per liter.
Consequently, any solution containing equal mol concen-
trations of H 4 " and OH~ is neutral. On the other hand, if
the mol concentration of H" 1 " is greater than the mol con-
centration of OH~~, the solution is acid; if the mol concen-
tration of OH~~ predominates over that of the H" 1 " ions, the
solution is basic.

Owing to the fact that water is an electrolyte (though a
very weak one) we find that, with certain types of salts,
the reverse of neutralization will occur. For example, if
pure NaAc is dissolved in water a slight reaction will occur.

IMPORTANT PRINCIPLES 27

Writing the equation ionically,

NaAc > Na+ + Ac~

+
HOH ^ OH- + H+

HAc

one sees that the acetate ions furnished by the sodium
acetate react to some extent with the hydrogen ions from
the water to form acetic acid molecules. On the other
hand, there is no measurable tendency for the sodium ions
to react with the hydroxyl ions because the product of such
reaction would be NaOH which is a strong electrolyte and,
therefore, completely ionized in solution. The result is
that, since the acetate ions remove some of the hydrogen
ions while the hydroxyl ions are not removed, there will
remain an excess of hydroxyl ions in the solution. Conse-
quently, the solution will give a basic test. Reactions of
this type are known as hydrolysis reactions.

Solutions of salts of weak acids and weak bases may be
either acid, neutral, or basic according to whether the
ionization constant for the acid is greater, equal to, or
less than the ionization constant of the base. A salt of a
strong acid and a strong base will not undergo hydrolysis
at all.

The importance of hydrolysis to the analytical chemist
will be seen later, especially during the analysis for the
third group metals. There the hydrolysis of ammonium
benzoate (the salt of the weak base, ammonium hydroxide,
and the weak acid, benzoic acid) is sometimes used to
furnish the hydroxyl ions needed to precipitate the hydrox-
ides of iron, aluminum, and chromium. Since, in general,
hydrolysis occurs to a much greater extent at higher
temperatures, the mixture is usually heated to about 95C.
during the precipitation of these hydroxides.

Water and pH Values. Water, as a weak electrolyte,
has an important influence on the type and concentration of

28 THEORY OF ANALYSIS

ions found in every aqueous solution. Its equilibrium
expression may be set up as follows :

(H+) X (OH-)

(HlO) = K (1)

This can be re- written as follows:

(H+) X (OH-) = K X (H 2 0) (2)

It can be shown by a few simple calculations that the con-
centration of water molecules in ordinary solutions is not
greatly different from that in pure water in other words,
it remains practically constant. If this is assumed to be
true, the product K X (H 2 O) may be set equal to another
constant, and the following expression results :

(H+) X (OH~) = K w ' (3)

In pure water the concentration of H + , as determined
experimentally at 25C., is 0.0000001 g. per liter and of
OH~ is 0.0000017 g. per liter. If these are expressed in
terms of mols, the concentration of each of H 4 " and OH"
is 10~ 7 mol per liter. Substituting in Equation (3),

10~ 7 X 10~ 7 = K w
.'. K u = 10~ 14

Consequently, for aqueous solutions at 25C.,

(H+XOH-) = 10~ 14 (4)

When an acid is added to water the hydrogen ion concen-
tration is increased. However, no matter how high the con-
centration of hydrogen ions may be, there will be some
hydroxyl ions yet remaining in the solution. Conversely,
an alkaline solution always contains some hydrogen ions.
If the concentration of one of the ions is known, the con-
centration of the other may be calculated by making use
of the final equilibrium expression [Equation (4)]. For

IMPORTANT PRINCIPLES

example, the hydrogen ion concentration in a solution
containing 10~~ 2 mol of hydroxyl ions per liter is

= icr 14

mol per liter

1Q-14

(H+) = -TJ =

From this it can be seen that the alkalinity or acidity
of any aqueous solution can be expressed in terms of the

Hydroxyl Ion
Concentration
Cmoles/liter)

1.0

0.1

0.01

0.001

00001

" 7

10"
10

- 12

pH Value

.14
I 3
I 2
1 I
I

9
8

I 7
6
5
4
3
2

Hydrogen Ion
Concentration
Cmoles/liter)

-' 3

Neutr

- 14

io-
,0-10

io- 9

IO" 8

10-"?
io- 6
io- 5

0.0001
0.001
01
O.I
1.0

FIG. 0. Chart of pH values and ion concentrations.

hydrogen ion concentration alone. A solution is said to be
acid, neutral, or basic according to whether the hydrogen
ion concentration is greater than, equal to, or less than
10~ 7 mol per liter.

Often it is more convenient to express the hydrogen ion
concentration of a solution in terms of the log of the
reciprocal of the hydrogen ion concentration, i.e., in terms
of the pH ("potential hydrogen") of the solution. For
example, if the hydrogen ion concentration of a certain

30 THEORY OF ANALYSIS

solution is 10~~ 4 mol per liter, the pH of the solution is given
by the following :

A neutral solution, therefore, has a pH of seven; a solution
having a pH smaller than seven is acid, while one whose
pH is greater than seven is basic (see Fig. 9).

Solubility Product: Supersaturation. Before starting a
discussion of precipitation reactions and the problems
peculiar to them, it is necessary that the student gain some
new conceptions concerning solubilities and saturated solu-
tions. Often, during beginning chemistry, different sub-
stances are characterized as being "soluble" or "insoluble. 7 '
Properly speaking, there is no such thing as a completely
insoluble substance. A substance may be so insoluble that
it is impossible for us to detect its presence in solution by
any chemical means available at present; nevertheless,
some small quantity is probably present in the dissolved
state. This is especially true of the common so-called
insoluble salts such as AgCl, CuS, CaCO 3 , and many
others. It can be definitely shown that these salts are
soluble to a very slight extent, of course, but soluble,
nevertheless.

To understand what is meant by the term solubility,
it is first necessary that a saturated solution be defined.
A saturated solution is a solution in which the dissolved
material is in equilibrium with excess undissolved material.
This is best explained by describing what happens during
the preparation of a saturated solution. If one places a
lump of sugar in water, molecules of sugar immediately
begin to pass into solution. At first, this is the only thing
that occurs, but, as soon as a few molecules are in solution,
some of them will, in moving about, collide with other sugar
molecules and precipitate out again. This precipitating
process is slow at first, but, as the concentration of dissolved

IMPORTANT PRINCIPLES 31

molecules increases, collisions between molecules will
become more and more frequent until finally a point will
be reached when sugar is precipitating at exactly the same
rate as the sugar is dissolving. When this point is reached,
it appears to the observer that no more sugar is dissolving.
Actually, these two opposing actions are continuing all the
time, but the net result is that the solid left in the mixture
will neither increase nor diminish in quantity since material
is precipitated as fast as it is being dissolved. The dis-
solved molecules are therefore in equilibrium with undis-
solv.ed material and the solution is now saturated.

With this picture in mind, we can define the term
solubility. The solubility of a substance is the concen-
tration of its saturated solution at a given temperature.
As a rule, the solubility of a substance increases with
increased temperature. Therefore, solubility tables usually
give the values of determinations made at temperatures
varying from 15 to 25C.

If these concepts are remembered, it is possible to take
up the problem connected with the formation and re-solu-
tion of precipitates. One of the first precipitation reactions
studied in the laboratory is the formation of silver chloride
by the addition of a solution of chloride ions to a solution of
some silver salt. The reaction is essentially a reaction
between silver ions and chloride ions to form silver chloride,
e.g.,

AgN0 3 -> Ag+ + N0 3 ~

+
HC1 -> Cl- + H+

jjr

AgCl (80hd)

Since HNOs is a strong electrolyte, there is little tendency
for combination of its ions. Precipitation of AgCl occurs
leaving a solution saturated with respect to Ag+ ions and
Cl~~ ions and in equilibrium with solid AgCl. Con-

32 THEORY OF ANALYSIS

sequently, the law of mass action applies and the following
expression holds true :

By using a few justifiable assumptions, it is possible to
simplify this expression to the following form :

(Ag+) X (C1-) = K,

* P

K BfD is known as the solubility product constant while the
mathematical product, (Ag 4 ") X (Cl~), is known as the
solubility product. This simply means that, no matter
whence the ions come, if, after equilibrium between solid
and dissolved material is attained, the concentration of the
silver ions remaining in solution, in mols per liter, is multi-
plied by the concentration of the chloride ions, the resulting
product can never be greater than the value of the solu-
bility product constant for silver chloride. In a saturated
solution of silver chloride, this product will exactly equal
the solubility product constant.

It can be seen from this illustration that the so-called
common-ion effect applies to this equilibrium, i.e., the
greater the concentration of Cl~, the smaller will be the
concentration of Ag + left in solution and the more complete
the precipitation.

Each slightly soluble electrolyte has its own solubility
product constant. At 18C., the constant for AgCl is
2 X 10~ 10 , for Agl, 1 X 10~ 16 , etc.

For more complicated molecules the relationship becomes
a little more complicated. For example, the expression for
PbCl 2 is

(Pb++) X (Cl-) 2 = K B . P = 1.0 X 1(T 4 (at 25.2C.)
and for Pb 3 (PO 4 ) 2

(Pb++) 3 X (POD 2 = #, p . = 8 X 1(T 43 (at 25C.)

IMPORTANT PRINCIPLES 33

Therefore, the solubility product principle states that, in
a saturated solution of any slightly soluble electrolyte,
the mathematical product of the molar concentrations of
the ions of the substance, each concentration raised to a
power equal to the number of those ions in the molecule,
will equal a constant, provided the temperature remains
constant. A given constant will, in general, hold true
only for:

1. One particular temperature. Most substances have
higher solubility product constants at higher temperatures.

2. One particular solvent.

3. One particular slightly soluble electrolyte.

4. A saturated solution, i.e., a solution in which undis-
solved material is in equilibrium with its dissolved ions.

5. One particular pressure. This applies to slightly
soluble gases such as hydrogen sulfide. Figures are usually
given for one atmosphere pressure.

6. Solutions whose total electrolyte concentration is low
(see Salt Effect, page 62).

Occasionally one will find a case where ions are mixed
in such concentration as to exceed the solubility product,
yet no precipitation occurs. Such solutions are in a
metastable state and will often remain that way for some
time unless some disturbing factor is introduced. A solu-
tion that contains more solute in solution than its saturated
solution would contain at that temperature is called a
supersaturated solution.

Most of the identification tests used in qualitative
analysis depend upon the formation of precipitates, many
of which readily give supersaturated solutions. Conse-
quently, it is important to know the methods available for
bringing about the precipitation of such substances.
Some of the more common of these methods are :

1. Scratching or rubbing the inside wall of the vessel
with a stirring rod. This is the best method to use in
analysis.

34 THEORY OF ANALYSIS

2. Sudden chilling of a supersaturated solution will often
cause precipitation to occur.

3. Heating the solution slightly will, in some cases, cause
precipitation to occur.

4. Shaking the solution will often disturb the metastable
state existing in the solution and cause precipitation.

Colloidal Solutions. Occasionally, it will be found that,
on addition of a reagent to an unknown, a solid material
forms but refuses to settle out. The solid material is so
finely divided that no individual particle possesses enough
weight to cause it to go to the bottom, even in the centri-
fuge. Since the analysis depends on being able to separate
this solid material from the filtrate, it is important that
some method be devised to precipitate colloids.

Colloidal solutions consist of extremely fine particles
suspended throughout the liquid. Each of the particles
carries a variable number of electrical charges, all the
particles in a given colloidal solution having the same
polarity. The fact that the charges on all the particles
are of the same sign results in their repelling each other
so that they cannot collect into the larger particles necessary
for settling. This suggests that if these charges could
be neutralized precipitation would occur. It is often
possible to neutralize these charges by adding ions carrying
a charge opposite in nature to those carried by the colloidal
particles. Unfortunately, one can never add a negative
ion without adding a positive ion at the same time for
ions do not exist in groups so that we can add one kind at
will. However, it happens that a bivalent ion is much
more effective than two monovalent ions. Consequently,
one can precipitate a positive colloid by adding ammonium
sulfate, the bivalent negatively charged sulfate ion being
much more effective than the positively charged mono-
valent ammonium ions. Similarly, calcium chloride would
be a good precipitant for negative colloids.

IMPORTANT PRINCIPLES 35

Unfortunately, these compounds cause precipitation of
certain metals in the wrong place and add ions for which we
must test later. However, a high concentration of ammo-
nium nitrate will usually precipitate any colloid, though it
may be a slow process. Magnesium sulfate is one of the
best general colloid precipitants but may be used only when
neither of its ions will interfere with the analysis.

Usually, heating a colloidal solution will cause .it to
precipitate. The heat tends to remove the charges on the
particles, and the internal agitation that accompanies the
heating causes the particles to collect in masses suffi-
ciently large to insure settling.

Dissolving Precipitates. It has already been pointed out
that a saturated solution of any substance is simply an
example of an equilibrium between solid material and its
ions in solution. Therefore, the problem of putting a solid
material into solution resolves itself into a question of the
possible ways one can force the reaction.

MX aolld ^ M lon + X lon

to the right. There are several ways this can be done.

1. If one of the ions can be forced to react with some
other ion to form a slightly soluble gas which will escape
and thus remove that ion from the reaction mixture, the
reaction will proceed to the right and the solid will dissolve.
When ZnS is treated with HC1 it will dissolve owing to the
fact that H 2 S, being only slightly soluble in acid, removes
the S = ion as fast as the latter is formed.

ZnS 80lid ^ Zn++ + Sr
+
2HC1 -* 2C1~ + 2H+

jjr

H 2 S

2. Occasionally, one of the ions can be caused to react
with the ions from some reagent to form a weak electrolyte

36 THEORY OF ANALYSIS

(such as water). When the slightly soluble PbSO 4 is
treated with a strong solution of ammonium acetate, the
PbSO4 will dissolve. This is due to the reaction between
the Pb ++ and acetate ions to form the slightly dissociated
PbAc 2 , the latter being one of the few salts that are weak
electrolytes.

S0 4 =

2NH 4 Ac -> 2Ac + 2NH 4 +

jjr

PbAc 2

An example of a reaction in which the formation of water
from one of the ions causes the reaction to go to completion
and the precipitate to dissolve is the case of ferric hydroxide
dissolving in hydrochloric acid.

Fe(OH) 8 Mllil ^ Fe+++ + 3(OH~)

3HC1 ->3C1" + 3H+

\\\
3H 2 O

3. Destruction of one of the kinds of ions produced will
shift the equilibrium and cause the solid to go into solution.
Copper sulfide, for example, is not soluble in hydrochloric
acid but is soluble in warm 3N nitric acid. This is owing to
the fact that the nitric acid oxidizes the sulfide ions to free
sulfur and sulfate ions. The reaction might be written as
follows :

CuS 801id ^ Cu+ + + S-
The sulfide ion is removed by the following reactions:

3S= + 8HNO 3 - 38 + 4H 2 O + 2NOT + 6NO 3 ~ (a)
3S= + 8HN0 3 - 3SO 4 = + 4H 2 O + 8NOT (6)

As fast as sulfide ion goes into solution, it is destroyed in
these two simultaneous reactions.

IMPORTANT PRINCIPLES 37

4. If a reagent is added which causes the formation of a
complex ion, the reaction will proceed towards the right.
A complex ion is an ion formed by the combination of a sim-
ple ion with other ions or with neutral molecules. Exam-
ples of these are the cupric ammonium ion, Cu(NH 3 )4 4 " f ;
the silver ammonium ion, Ag(NH 3 ) 2 + ; the cuprocyanide
ion, Cu(CN) 3 == ; the cobalticyanide ion, Co(CN) 6 s ; the zinc
ammonium ion, Zn(NH3) 4 " H "; and many others of like
nature. An example of forming complex ions in order to
put a substance into solution is the use of ammonium
hydroxide to dissolve AgCl. The reaction that occurs is as
follows :

AgCUa ^Ag+ + Cl

+
2NH 4 OH ^ 2NH 3 + 2H 2 ()

jjr

Ag(NH,),+

The complex ions, which are stable and dissociate only
slightly to give simple ions, have entirely different proper-
ties than the simple ions from which they were formed.
Therefore, the above reaction removes Ag + from the mixture
and the AgCl goes into solution. Acidification of this
mixture will reprecipitate the AgCl, because the NH 3 is
removed, as is shown in the following equation :

Ag(NH 3 ) 2 + ^2NH 3 + Ag+

+
2HNO 3 -> 2H+ + 2NOr

2NH 4 +

Complex Ions. Many metallic ions combine with certain
types of neutral molecules or with anions to form complex
ions, some of which are exceedingly important. A few of
these were mentioned in the preceding paragraphs.

38 THEORY OF ANALYSIS

The cupric ammonium ion, Cu(NH 3 )4" f+ , is formed by the
combination of Cu" f+ with four NH 3 molecules. The
reaction* may be written

Cu+ + + 4NH 3 -H 2 ^ Cu(NH 3 ) 4 ++ + 4H 2

Similarly, the cobalticyanide ion is formed by the combi-
nation of eobaltic ions with cyanide ions:

GO+++ + 6CN- ^ Co(CN) 6 =

Many of these complex ions are used in the analytical
separation of the metals. A few of those not mentioned
previously are the ammonia complexes of cadmium, cobalt,
and nickel, Cd(NH 3 ) 4 + + , Co(NH 3 ) 6 ++ and NiCNH,^; the
chloride complexes of lead, mercury, antimony, and tin,
PbClr, HgClr, SbCl 6 ", and SnCl<r; and the thio com-
plexes of arsenic, antimony, and bismuth, examples of
which are the thioarsenite, AsS-r; thioantimonate, SbS 3 a ;
and thiostannate, SnS 3 = , ions. The ammonia complexes
are formed by treating solutions of the metallic ions with
excess ammonium hydroxide ; the chloride complexes exist
in strong hydrochloric acid solutions of lead, mercury,
antimony, and tin ions ; and the thio complexes form when
the sulfides of arsenic, antimony, or tin are treated with
solutions of sodium or ammonium sulfide.

* Although a solution of ammonia in water is usually considered to be a
solution of NFUOH molecules, many of its reactions, as well as many spectral
and other physical data, indicate that it is chiefly a solution of simple NH 3
molecules. Consequently, its characteristic reactions may be explained by
the existence, in ammonium hydroxide, of two equilibriums:

NH 4 OH =; NH 4 + + OH- (1)

NH 4 OH ^ NH 3 + H,0 (2)

Equation (1) represents its weakly basic reactions, while Equation (2)
shows that a solution of ammonium hydroxide may yield a high concentra-
tion of NHa molecules.

IMPORTANT PRINCIPLES 39

The phenomenon of ions combining with other ions or
neutral molecules to form complex ions is not explained by
ordinary conceptions of valence. A broader theory was
developed, therefore, by A. Werner, postulating the
existence of two types of valence. The first of these, the
primary valence, is the ordinary kind of chemical valence.
The second, which Werner termed the auxiliary or secondary
valence, is a type of bond which an ion can exhibit in addi-
tion to its ordinary valence.

The reason for the existence of secondary valence is not
fully understood as yet. However, it is believed that in
many cases the secondary valence (sometimes called the
coordinate valence) is due to the ability of ions or neutral
molecules to attach themselves to unused electron pairs in
other ions or neutral molecules. For example, the valence-
electronic configuration of the ammonia molecule might be
written,

H
:N:H

where the dots represent electrons, the hydrogen atoms
being held to the nitrogen through shared electron pairs.
The unshared pair of electrons makes this molecule readily
capable of forming complex ions. For example, the con-
figuration of the cupric ammonium ion may be written :

NH 3 -
H 3 N:Cu:NH 3
NH 3

The maximum number of ions or neutral molecules with
which a given ion can combine is called the coordination
number of that ion. By studying these, it is possible to
predict, with fair accuracy, the probable formula of any of
the common complex ions. The coordination numbers of a
few ions are given in the following table :

40 THEORY OF ANALYSIS

Ag+ ........... 2 Hg+

Cd 4 + . 4 Mn

Co 44 6 Ni 4

Co+ +4 (> O

2 and 3 Pt

4
6
6
3
6
6
6

Cu ++ . . 4 Sh

Fe++ 6 Sn

Fe ++ +. 6

The usefulness of complex ions depends upon the fact
that, like weak electrolytes, the complex body is in equilib-
rium with its simple components. For example, the cupro-
cyanide ion in solution is in equilibrium with cuprous
ions and cyanide ions:

Cu(CN)a" ^ Cu + + 3CN-

Similarly, the cadmium cyanide ion sets up the following
equilibrium :

Cd(CN)r ^ Cd++ + 4CN-

For each of these equilibriums, it is possible to write an
equation similar to the expression used for weak electrolytes :

(Cu+XCN-)'

(Cu(CN)D ~ A <* U ' N >'- - x 1U

<Cd^)(CN-r . K _ j 4 x 10 - 17

(Cd(CN)rj "" Ard(rN) * " J>4 x 1U

The dissociation or instability constants for different com-
plex ions have different values. The smaller the constant,
the lower will be the concentration of simple ions existing in
the solution of a given complex ion.*

* Particular attention might he called to the use of CN" ion in the separa-
tion of copper and cadmium. In solutions of the cyanide complexes of
cadmium and copper, the concentration of the simple ions is sufficiently
great for both CdS and CuS to be precipitated by treatment with HaS.
The dissociation constant of the Cu(CN)3~, however, is much smaller than
that of the Cd(CN) 4 ". Consequently, by putting in excess KCN, the
concentration of Cu + can be repressed to the point where it is too low to
give a precipitate with H2S. On the other hand, the more highly dissociated
Cd(CN)4~, though giving less Cd +4 in excess KCN, will give a sufficient
concentration of Cd ++ to allow it to be precipitated with

IMPORTANT PRINCIPLES 41

Occasionally, use is made of the inverted form of these
functions. For example, if the equilibrium for Cu(CN) 3 sa=
is written:

Cu+ + SON" ^ Cu(CN).-

the expression for the corresponding equilibrium constant
would become :

(Cu + )(CN-) 3 ~ /V ^ N >>-

This constant is called the stability constant of the ion.
While the dissociation constant is more a direct measure of
the concentration of the simple ions in solution, the sta-
bility constant is a measure of the concentration of the
undissociated complex ions. It will be noted that the
stability constant is equal to the reciprocal of the instability
constant, ?'..,

K -1

-** nubility ~~ E"

* v mutability

Amphoteric Hydroxides. As a rule, slightly soluble
metallic hydroxides are soluble in strong acids. Some,
however, are soluble in either strong acids or strong bases.
Such hydroxides are said to be amphoteric. Aluminum
hydroxide, zinc hydroxide, and chromic hydroxide are
common examples of amphoteric substances. They are
very weak bases and, as will be shown, are also slightly
acidic in nature. The dual properties of these substances
are expressed by using Zn(OH) 2 as an example. In a
saturated solution of Zn(OH) 2 , a double equilibrium is set
up. The effect of adding hydrogen or hydroxyl ion is shown
by the following equation:*

* To visualize better the mechanics of such reactions, one might look upon
the zinc hydroxide as being written in the customary form for acids, H*ZnO2,
as well as that for a base, Zri(OH) 3 .

42 THEORY OV ANALYSIS

H+ + HZnOr ^ f"Zn(OH) a l ^ Zn++ + 2(OH~)
OH~ < (strong base) (strong acid) 2H 4 "

H 2 O 2H 2 O

Addition of either an acid or a base will produce water and
result in a neutralization. In the case of the addition of
acid, OH~ ions are removed by the formation of water and
the equilibrium is shifted to the right, Zn(OH) 2 being
dissolved. On adding a base, the H + ions are removed by
the formation of water and the equilibrium is shifted to the
left, Zn(OH) 2 being again dissolved. As can be seen by a
little study of the equation, after the Zn(OH) 2 has been dis-
solved by either a base or an acid, it can be reprecipitated
by cautious addition of successive small amounts of an
acid or a base.

The fact that not all elements react in this way makes it
possible to use this reaction to separate zinc, chromium,
and aluminum from iron, manganese, nickel, and cobalt.
The hydroxides of these three metals are dissolved in an
excess of NaOH and filtered off from the insoluble
hydroxides.

It should be apparent from this discussion that there is a
definite hydrogen ion concentration at which the solubility
of a given amphoteric hydroxide is a minimum. The pH
at which an amphoteric substance exhibits its minimum
solubility, is known as the isoelectric point of that substance.

If, therefore, aluminum hydroxide has been dissolved in
excess of either an acid or a base, it may be reprecipitated
by changing the hydrogen ion concentration to the value
corresponding to the isoelectric point for aluminum
hydroxide.

This may be done in either of two ways :

1. An arid or base may be added to increase or decrease
the hydrogen ion concentration of the solution until it
reaches the proper value ; or

IMPORTANT PRINCIPLES 43

2. In case the aluminum hydroxide was dissolved in
excess base, the hydrogen ion concentration can be increased
to the proper value by addition of a buffering agent (such
as ammonium chloride) to reduce the hydroxyl ion concen-
tration. This method is important to the analytical
chemist (see Group III analysis), as the hydrogen ion
concentration can be controlled more accurately by use of
buffers than by any other readily available means.

Lead, arsenic, antimony, and tin are also readily ampho-
teric. Pb(OH) 2 dissolves in strong sodium hydroxide
solution to form the plumbite ion, HPbO 2 ~~; arsenous
oxide forms arsenite ion, AsO 3 ^; arsenic oxide forms
ar senate ion, AsO 4 ss ; stannous hydroxide, Sn(OH) 2 forms
stannite ion, HSnO 2 ~~.*

Oxidation and Reduction. In some reactions, besides
the simple reactions between ions, there are changes in the
valence of two or more atoms or ions. These reactions,
known as oxidation-reduction reactions, always involve a
transfer of electrons. The simple case of the combination
of copper with bromine, when copper is placed in bromine
water, illustrates the changes that occur in this type of
reaction. Ordinarily the reaction is written thus:

Cu + Br 2 -> CuBr 2

From our knowledge of electrolytes and ionic equations, it
should be apparent that a more correct representation of
the facts is found in the following :

Cu -> Cu++
Br 2 -* 2Br~

* Tin is peculiar in one respect. In the stannic state, it forms two
hydroxides. One, the normal hydroxide (sometimes called stannic add)
H 2 Sn(OH) 6f exhibits the normal amphoterism described; the other, meta-
stannic acid, is very difficult to dissolve in either acids or bases. This form
results when metallic tin is treated with nitric acid. It can be dissolved,
however, by converting it to the starmate by an alkaline fusion or by heating
with concentrated IICl followed by addition of water. The difference
between the reactions of the hydroxides is described as being due to differ-
ences in physical state rather than in composition.

44 THEORY OF ANALYSIS

No combination of these ions should be indicated since
CuBr 2 is highly ionized in solution.

To analyze this further it is necessary for us to recall
something of the structure of atoms. It will be remem-
bered (a review of atomic structure as given in any standard
inorganic text will help the student) that atoms consist of a
centrally located, positively charged nucleus surrounded by
orbits in which electrons are traveling. These electrons
are present in the exact number required to neutralize the
positive charges on the nucleus, so the atom is normally
neutral. The outer orbit has the ability to either lose all
its electrons or take on more. If it takes on electrons, it
will take on just enough to make the total number present
in that orbit exactly eight. The question of whether it
takes on electrons or loses them depends upon the number
already present in that orbit in the neutral atom. If
there are less than four, the atom loses electrons; if more
than four, it takes on electrons. When either of these
processes occurs, the product is an ion.

In the case under consideration, copper has two electrons
it can lose. Bromine, having seven electrons in its outer
orbit, needs one electron to make up its octet. When
copper reacts with bromine, the copper atom gives up its
two valence electrons, the copper atom being left with two
excess positive charges. The bromine atoms take up the
electrons, one to each atom, and thus assume one excess
negative charge. Thus, the reaction between copper and
bromine solution results in the formation of cupric and
bromide ions.

With these facts in mind the above reaction can be
correctly written as follows :

Cu 2 electrons (e) > Cu ++
Br 2 + 2 electrons (c) -+ 2Br~

The truth of the above representation can be tested by
placing a rod of copper in a solution of CuSO 4 , and a

IMPORTANT PRINCIPLES

platinum electrode in a solution of KBr containing bromine,
connecting the two solutions with a U tube filled with KC1
solution and connecting the copper and platinum elec-
trodes to a milli voltmeter (Fig. 10). Although the copper
and bromine are not in contact, reaction will occur and a
strong current of electricity will flow through the milli-
voltmeter showing that the reaction is accompanied by a
transfer of electrons.

Voltmeter

Sof/t bridge ,

(Containing KCI solut/on)

Copper 4
electrode-'

<=~P/cftinum
electrode

IN solution
ofCuS0 4

. >zr

~~ IN solution
of KBr and
bromine
wafer

Copper Bromine

Half Cell Half Cell

FIG. 10. Arrangement of i copper-bromine cell.

Elements vary widely as to the readiness with which they
tend to lose or gain electrons. Sodium, for example, has a
very strong tendency to lose its single outer electron while
copper has much less tendency to lose its outer electrons.
Chlorine readily takes on one electron while iodine has
much less tendency to gain its extra electron. By using
these elements in cells like that described above it has been
possible to measure these relative tendencies in terms of
volts. The element is placed in a solution IN with respect
to its ions. This half -cell* is connected to a hydrogen

* An electrode immersed in a solution of its ions is known as a half-cell.
Two half-cells connected by a salt bridge (a U tube filled with a solution of
some salt such as KCI) constitute a cell.

46 THEORY OF ANALYSIS

electrode, by means of the salt bridge (U tube) described,
and the voltage of the resulting combination is measured.
The hydrogen electrode is used as a reference base, and,
consequently, the voltage measured is assigned to the
element being used. Tables of these normal electrode
potentials are very useful in predicting the probable com-
pleteness of given oxidation-reduction reactions. A few
of these potentials are given in Table I (see Appendix).

The neutral atoms of those elements at the top of this
list have the strongest tendency to lose electrons, and their
ions have the least tendency to take on electrons again and
become neutral atoms. Those at the bottom have the
strongest tendency to gain electrons, their ions having the
least tendency to lose electrons. The voltage of any
combination (which is a measure of the tendency for reac-
tion to occur) is found by taking the algebraic difference
of the individual electrode potentials of the components.
For example, it may be necessary to determine the voltage
of a cell made up of a normal zinc electrode and a normal
chlorine electrode.

Zn = +0.76
Cl = -1.36

Therefore the voltage of this cell will be

0.76 - (-1.36) = 0.76 + 1.36 = 2.12 volts

Where the resulting voltage is as high as this, the reaction
will be complete, i.e., zinc and chlorine will react with great
vigor.

Electrode potentials have a very important application
to some phases of analytical chemistry. For example, it is
necessary to reduce tin from the stannic to the stannous
state before testing for this metal. To do this it is neces-
sary that some substance be chosen as reducing agent that

IMPORTANT PRINCIPLES 47

has a more negative oxidation-reduction potential than the
potential for the reaction

Sn++ - Sn++ ++ + 2e

If the latter is looked up in a table of oxidation-reduction
potentials, it is found that the potential for this reaction is
+0.13 volt. This means that any metal having an elec-
trode potential more negative than +0.13 volt will reduce
stannic ion to stannous ion. The greater the difference
between the two (i.e., the greater the potential of the cell
made up of a combination of the two), the more complete
and rapid will be the reduction of the stannic ion. How-
ever, the reaction

Sn -> Sn ++ + 2e

has a potential of 0.13 volt. Any metal more negative
than this will reduce the stannic ion completely to the
metallic state. Consequently, a metal whose electrode
potential is between +0.13 volt and 0.13 volt is necessary
to bring about the desired reduction without at the same
time reducing Sn^" 1 " to Sn. Lead ( 0.12 volt) is one such
metal and is sometimes used for this reduction. However,
it reacts so slowly that more active metals (e.g., zinc,
aluminum, or magnesium) are usually used, any metallic
tin being redissolved in hydrochloric acid.

Balancing Oxidation-reduction Equations. Balancing
oxidation-reduction equations by the trial-and-error method
is usually difficult owing to the complexity of the reaction.
The above discussion, however, opens other methods to us,
the principle of which consists of the balancing of the
number of electrons involved.
These methods are as follows:

I. Valence-electron Method. Five steps are needed in
this method. To illustrate, let us consider the following
equation :

KMnO 4 + HC1 -> KC1 + MnCl 2 + H 2 O + C1 2

48 THEORY OF ANALYSIS

The steps are as follows :

1. Pick out the elements that undergo changes in valence
and write these changes, with their electron changes, as
equations. In the above equation, Mn ++ ~ l ~~ H ~~ H ~ changes to
Mn 4 " 4 " and some of the Cl" changes to C1 2 .

Mn ++f++ ++ + -* Mn +4 -
Cl- - e -> C1 2

2. Balance the atoms and the electrons in each equation.

5e
2C1- - 2e - C1 2

3. Multiply each of these equations throughout by a
number which will make the number of electrons in each
equal to that in the other.

5c
5(2C1 - 2e -> C1 2 )

4. Add all those atoms on the left side of the equations
and place them on the left side of a new equation; do
similarly for those on the right side.

10C1- -> 2Mn+ 4 + 5C1 2

5. Using these numbers in the original equation balance
the atoms and molecules as far as possible.

2KMn0 4 + 10HC1 -> 2K+ + 2Mn+ + + H 2 O + 5C1 2

Ten Cl" call for ten HC1 molecules. However, there are
eight oxygen atoms in the two KMnCh molecules which
form eight molecules of water requiring sixteen hydrogen
ions. Since HC1 is the only source of H" 1 ", at least sixteen
HC1 molecules are required. The chloride ion from the six
extra HC1 molecules react with the K 4 " and Mn 4 " 4 " and the
equation is thus completely balanced.

2KMnO 4 + 16HC1 -> 2KC1 + 2MnCl 2 + 8H 2 + 5C1

IMPORTANT PRINCIPLES 49

II. Ion-electron Method. There are six steps to this
method of balancing oxidation-reduction equations. Using
the same equation as before to make comparison easy:

KMnO 4 + HC1 - KC1 + MnCl 2 + H 2 O + C1 2

the steps are as follows:

1. Write the equation in terms of the ions present and
cancel all those ions that go through unchanged :

MnOr + H+ + Cl- -> Mn++ + H 2 O + C1 2

The H" 1 " must be retained on the left as it goes to form the
slightly dissociated water. The Cl"" are retained on the
left- -though canceled on the right because some of them
go into the form of C1 2 molecules.

2. Write two equations, one for each ion that undergoes
an electronic change and balance the atoms and ions in each.

Mn0 4 ~ + 8H+ -> Mn++ + 4H 2 O
2C1- - C1 2

3. Take the algebraic sum of the charges on the right side
of each equation and subtract algebraically from the alge-
braic sum of the charges on the left. The difference will
give the number of electrons to add (algebraically) to the
left side of the equation to make it balance electrically.

Thus, for the first equation,

(-1 + 8) - ( + 2 + 0) = 5

which means that five electrons must be added to the left
side of this equation. For the second equation,

(-2) - (0) = -2

which means that two electrons must be subtracted from
the left side of this equation. Therefore, to balance these
completely.

50 THEORY OF ANALYSIS

MriOr + 8H+ + 5e - Mn++ + 4H 2 O
2C1~ - 2 -H. Cl.>

4. Multiply each of these equations throughout by a
number which will make the number of electrons in each
equal to those in the other.

2(Mn()r + 8H+ + 5c -> Mn^ + 4H,O)

5(2Cr + 2e -> C1 2 )

5. Add all those ions on the left side of the equations and
place them on the left side of a new equation; do similarly
for those on the right side. This will give a completed
ionic equation for the reaction.

2MnOr + 16H+ + 10C1- -> 2MH++ + 8H 2 O + 5C1 2

6. Complete the equation by writing the proper coeffi-
cients in the original equation using the same method as
described for the valence-electron method in part 5.

In general, either of these two methods is applicable to
any oxidation-reduction equation. Strictly speaking, how-
ever, the latter method is preferable for reactions in solution
where the reaction is actually between ions.

MATHEMATICAL RELATIONSHIPS

In previous paragraphs there has been given a qualitative
discussion of various laws. For practical use it is usually
important to be able to consider the mathematical aspects
of a given analytical problem. lonization constants and
solubility products mean little to one who does not know
how to use them or how they were derived. For this
reason it is important to practice working such problems
until the necessary understanding is reached.

From the standpoint of the student, the more important
types of problems are the following :

MATHEMATICAL RELATIONSHIPS 51

1. Weak electrolytes: These involve the calculation of
the concentration of one ion of a weak electrolyte when

a. The ionization constant and the concentration of the
weak electrolyte are known.

6. The ionization constant, the concentration of the
weak electrolyte, and the concentration of a second ion are
known (the case of buffer solutions).

c. The ionization constant of the weak electrolyte and
the concentration of all electrolytes present are known.
An example of this case is the adding of a strong electrolyte
to a buffer solution.

2. Solubility product: These involve

a. The calculation of the concentration of one ion of a
slightly soluble salt (in a solution saturated with respect
to its ions) when the concentration of the other ion and the
solubility product constant of the salt are known.

6. The calculation of the solubility product constant from
the solubility of the substance.

c. The calculation of the solubility of a slightly soluble
salt from its solubility product constant.

d. The calculation, from solubility product constants, of
the efficiency of separation of two given ions which are
being precipitated by a common reagent.

In the case of each of these problems, it is only necessary
to choose the appropriate formula* and substitute the
given values in it. An example of each type should serve
to illustrate the method to be used.

Problems Involving Weak Electrolytes. Type 1 a.
The first type of problem always makes use of the equation
for the ionization constant of the weak electrolyte. For
example, the hydroxyl ion concentration in a 0.0 1M solu-
tion of NH 4 OH is found by first deriving the equation for
the equilibrium constant of NH 4 OH, then substituting the

* It is best to learn how to derive the various formulas rather than to
depend upon the memory. The derivations are simple and are much less
likely to lead to the use of the wrong formula.

52 THEORY OF ANALYSIS

proper numerical values in it. The ionization that occurs is
NH 4 OH ^ NH 4 f

Therefore, applying the law of mass action, the equation
is found to be [(NH 4 + )(OH-)]/(NH 4 OH) = K b . K b for
NH 4 OH is 1.8 X 10~~ 5 ; and recalling from earlier dis-
cussions that (NH 4 + ), (OH~), and (NH 4 OH) mean the mot
concentrations of NH 4 + , OH~, and NH 4 OH molecules, re-
spectively, substitutions can be made and the hydroxyl ion
concentration calculated. To do this, however, two things
must be remembered. First, OH~ and NH 4 + must be
present in equal concentrations each molecule that has
ionized to give an OH~ has, at the same time, given an
NH 4 + . Therefore, if we let X equal the OH~ concen-
tration (which is what we are trying to calculate), it must
also equal the ammonium ion concentration. Therefore X
can be substituted for each of these in the above equation.

The second point to remember is that, in the case of any
weak electrolyte, only a very small fraction of the molecules
present are in the ionized state at any given time. Conse-
quently, it is possible, without introducing any serious
error for ordinary purposes, to assume that the concen-
tration of NHtOH molecules is equal to the original con-
centration placed in the solution, i.e., (in this case) O.OlAf .

Thus, using these simplifications, it is found that

(*)(*) - ! 8 x 10 -3
0.01 ~ IB * X 1U

Jf 2 = 1.8 X 10~ 5 X 0.01 = 1.8 X 10~ 7
X* = 18_X 10~ 8
X = Vl8 X 10~ 4 mol per liter = OH~~~ concentration*

* To calculate the hydrogen ion concentration of a solution of ammonium
hydroxide, it is first necessary to calculate the hydroxyl ion concentration
as described here. Then the concentration found for the OH~ is substituted
in the equation

and the hydrogen ion concentration is calculated (see Water and pH Values).

MATHEMATICAL RELATIONSHIPS 53

Type 1 b. These problems refer to solutions of a mixture
of a weak electrolyte with a salt having a similar ion. For
example, calculate the hydrogen ion concentration in a
solution 0.01 M with HAc and 0.1M with NaAc. Again
the formula for the ionization constant of the weak elec-
trolyte (HAc) is used.

(H+XAc-)
(HAc) ~ A

The concentration of HAc molecules is assumed to be
that of the original acetic acid. NaAc, however, is a strong
electrolyte (as are practically all salts) and may be assumed
to ionize completely giving a solution 0.1M with respect to
each of the ions, Na + and Ac~. A small amount of Ac~~
comes from the ionization of the HAc, but this is so small
compared to the quantity coming from the NaAc that it can
be ignored. Therefore, if it is assumed that all the Ac~
comes from the NaAc, its concentration may be taken as
0.1M. K a for HAc is 1.86 X 10~ 5 . Using these values as
in the above equation,

(H+X0.1) _ i g 1Q _ 5

0.01 "" ** X 1U
(H+)(0.1) = 1.86 X 10~ 5 X 0.01 = 1.86 X 10~ 7

H + = L86 Q X 1 10 ~ 7 = 1.86 X l<r 6 mol per liter

Type 1 c. In the case of a mixture of a low concentration
of a strong acid and a higher concentration of the salt of a
weak acid, another factor enters. It is found that in such
mixtures the strong acid reacts quantitatively with the salt
forming a mixture of the weak acid and the salt of the
strong acid. For example, if HC1 is added to NaAc the
following reaction occurs:

54 THEORY OF ANALYSIS

Na+ + Ac~
+

or + H+

JJt

HAc

If the salt is in excess, the result will be a mixture of
(from the NaCl and the excess NaAc), Ac~ (chiefly from the
excess NaAc), Cl~" (from the NaCl), H 4 " (from the ionization
of the HAc formed), and HAc molecules. The only weak
electrolyte present is HAc; therefore use is again made of
the equation for its ionization constant, in calculations.
To take a typical example, calculate the hydrogen ion con-
centration in a solution 0.1 M with NaAc and 0.01M with
HC1. When the solution was first prepared, the above
reaction occurred, forming approximately 0.01 mol of
HAc and, consequently, removing 0.01 mol of NaAc. This
leaves 0.1 0.01, or 0.09 mol per liter, of NaAc. ' This
ionizes, giving 0.09 mol of Ac~~. Therefore, the solution is
now 0.0 IM with HAc and 0.09M with Ac~. Therefore
substituting in the equation,

(HAc) ~
it is found that

(H + )(0.09)

o.oi - li8 x 10

(H+)(0.09) = 1.8 X 10~ 7
1 8 V 10~~ 7
H + = 9 x 10-* = " 2 X 1(r * mo1 per liter *

* The solution is identical with a solution prepared by adding sufficient
NaCl, NaAc, and HAc to water to make a solution 0.01 M with NaCl,
0.09M with NaAc, and O.OlAf with HAc. The presence of the NaCl may
be disregarded here, though, as will be learned later, it has a small effect on
the apparent concentration of the other ions (see Salt Effect) . For this

MATHEMATICAL RELATIONSHIPS 55

In a similar manner it is possible to calculate the con-
centration of OH"" ions in solutions of weak bases containing
a salt of a weak base (Type 16) or in solutions containing a
mixture of a strong base and a greater concentration of a
salt of a weak base (Type Ic). In such cases the formula
for the ionization constant of the weak base is used. For
example, if NH 4 OH or its salts are involved in these
problems, the equation is as follows :

(NH 4 OH) ~ *

The reasoning and substitutions are similar to those in the
cases described above.

Problems Involving Solubility Products. These are very
similar in nature to those for weak electrolytes. Again it
is a question of choosing the proper equation (which is, this
time, the equation for the solubility product constant of the

reason, it is legitimate to treat the problem as dealing with a solution of
HAc and NaAo in pure water.

In case the concentration of the strong acid is greater than the concen-
tration of the salt of the weak electrolyte, the problem is resolved into one
of calculating the concentration of H+ from the unreacted strong acid. For
example, if a solution is made Q.IM with HCl and 0.01 M with NaAc, a
reaction again occurs:

This time, the acetate ion concentration is lower than the hydrogen ion
concentration. Consequently, reaction will occur until practically all the
Ac~ is used. This will leave a solution which is equivalent to a mixture of
NaCl, HCl, and HAc. The HCl, being a strong electrolyte, furnishes so
much more H + than does the HAc that the latter can be regarded as yielding
no H+ at all. Hence, it is only necessary to calculate the concentration
of H+ coming from the HCl. Using reasoning similar to that above, it
will be found that the solution is 0.09M with HCl. Therefore the hydrogen
ion concentration, assuming complete ionization of the HCl, is 0.09 mol
per liter.

56 THEORY OF ANALYSIS

slightly soluble substance involved) and substituting in it
the known values.

Type 2 a. Assume that it is desired to calculate the
maximum concentration of Ag" 1 " that can remain in a solu-
tion in equilibrium with 0.001M I~ ion. Obviously, the
slightly soluble substance involved is Agl (whose K av is
1.5 X 10~ 16 ). The equation for the solubility product
constant is

Substituting in this,

(Ag+) (o.ooi) = 1.5 x ur 16

1Q-16 10~ 16

Ag+ = 1.5 X = 1.5 X =3- = 1.5 X 10~ 13 mol per liter

Type 2 6. This type of problem involves the calculation
of the solubility product constant of a slightly soluble salt
from its solubility. For example, given that the solubility
of Ag 2 S is approximately 1.35 X 10~ 17 mol per liter, calcu-
late the K s p for Ag 2 S. To do this it is necessary to know
and remember one other point.

When one molecule of Ag 2 S ionizes, it gives two Ag+ and
one S = .

Ag 2 S -> 2Ag+ + S-

Consequently, since a saturated solution of Ag 2 S contains
1.35 X 10~ 17 mol of the latter, completely ionized, it must
contain two times that much, or 2.7 X 10~ 17 mol per liter of
Ag+, and one times that much, or 1.35 X 10~ 17 mol per liter,
of 8 s . Reviewing the solubility product principle gives

(Ag+)*(S=) = K 8D
as the equation involved. Substituting in this

(2.7 X 10- 17 ) 2 (1.35 X 1(T 17 ) = K..,

(7.29 X 10- 84 )(1.35 X 10~ 17 ) = K. v

K._ v = 9.84 X l(r 81 = 0.984 X lO" 80

MATHEMATICAL RELATIONSHIPS 57

If the concentration had been given in grams per liter, the
first step would have been to convert it to mols per liter.
This is done by dividing the number of grams in 1 liter by
the molecular weight of the salt.

Type 2 c. The reverse of the problems of Type 26 would
be to calculate the solubility of a slightly soluble substance
from the solubility product constant. For example, given
that K 8P for Ag 2 CrO 4 is 3.3 X 1(T 12 , calculate the solu-
bility of Ag 2 CrO 4 . Again it must be remembered that
when 1 mol of Ag 2 CrO 4 ionizes it gives 2 mols of Ag +
and 1 mol of CrO^. If S represents the solubility of
Ag 2 CrO 4 in mols per liter, the concentration of Ag + in the
saturated solution will be 2S and that of CrO 4 = will be S.
Again, using the solubility product constant equation,

(Ag+) 2 (Cr()r) = K 8P
We can substitute and calculate as follows:

(2S) 2 (S) = 3.3 X 10-' 2
(4S 2 )OS) = 3.3 X 10~ 12
4S 3 = 3.3 X 10~ 12

S = \S678 X 10~ 12 = 0.94 X 1(T 4 mol per liter
= solubility of Ag 2 CrO 4

To get the solubility of Ag 2 CrO 4 in grams per liter, it is only
necessary to multiply the above answer by the molecular
weight of Ag 2 Cr0 4 .

Type 2 d. This type of problem involves the calculation
of the ratio of the concentrations of two slightly soluble
substances having a common ion, in a solution saturated
with respect to each. For example, assume that a solution
is 0.1 M with respect to each of the ions, Pb"^" and Mn" 4 " 1 ".
If H 2 S is passed into the solution, PbS will precipitate.
This removes Pb +4 ~ from the solution (leaving all the Mn 4 "*
in solution). After a time, however, the concentration of

58 THEORY OF ANALYSIS

the Pb 4 "* will be so reduced that the sulfide ion concentra-
tion, which, in accordance with the solubility product
principle, has been increasing as the lead ion concentration
decreased, reaches a value sufficiently high for the precipi-
tation of MnS. It is desired, therefore, to calculate what
will be the concentration of Pb 4 "*" left in solution at the
moment MnS first begins to precipitate. This is a measure
of the efficiency of the separation of Pb 4 " 4 and Mn 4 " 1 " by
the use of H 2 S.

It is evident that the sulfide ions present in the solution
at that moment are responsible for the precipitation of both
PbS and MnS. Consequently, write the equations for the
solubility relationship of each

then divide the upper equation by the lower equation

and cancel the sulfide ion concentrations. This may be
done, for, since all three ions are present in the same solu-
tion, the sulfide ion concentration is the same for both
reactions.

The resulting equation will be

Now the given values may be substituted in this equation
= 10~ 28 , K, f , MaS 10~ 15 ) and the concentration of
calculated (at the moment MnS first begins to
precipitate, the concentration of Mn ++ will be O.lAf) :

0.1 _ IP" 15 _ 1

(Ph-*-*-) ~ i<r 28 ~ io- 13

(Ph*- 1 ") = (IO- 13 ) X (0.1) = 10- 14 mol per liter)
This shows that the separation is very complete.

MODERN THEORIES OF ELECTROLYTES 59

More complicated problems than these may be met, but
it will be found that they are merely combinations or varia-
tions of those considered above.*

MODERN THEORIES OF ELECTROLYTES

Hydronium Ion. In the treatment of solutions, it has
been the belief, until recent years, that acids give hydrogen
ions when placed in solution. lonization of acids was
presumed to be the simple dissociation of the acid molecule
into positively charged hydrogen ions and negatively
charged ions. A study of the properties of electrolytes in
nonaqueous solvents shows that ionization is not so simple
as this, however, revealing many phenomena that are not
explained by the earlier theory. After many experimental
data had been studied, Bronsted and others finally arrived
at the conclusion that there is a definite reaction between
the acid molecule and the molecule of the solvent, resulting
in the union of the H + ions and the solvent to form solvated
ions.f For example, the ionization of hydrogen chloride
in water would be represented by the equation

H 2 O + HC1 = H 3 O+ + Cl-
the H 3 O 4 being called the hydronium ion.

* In the Appendix will be found a list of problems of the various types
considered in this and earlier sections. Most of these have practical applica-
tions to the analytical procedure. Consequently, the student should work
them and try to see the application at the same time.

f There is evidence for the belief that all ions in solution, both negative
and positive, are solvated. Conductivity measurements on solutions give
values lower than those calculated on the basis of the mass and charge
of the ions involved, the deviation being explained, in part, by the theory
that solvent molecules attach themselves to the ions and act as a "drag"
on the movements of the ions. As a rule, however, the other hydra ted ions
are much less stable than is the hydronium ion, HaO+, formed when acids
dissolve in water.

An excellent discussion of the solvation of ions is given in the section
on complex ions in "Qualitative Analysis and Chemical Equilibrium' 1 by
Hogness and Johnson, Henry Holt & Company, New York, 1937.

60 THEORY OF ANALYSIS

This theory has almost completely supplanted the earlier
conception of ionization. The chief reasons for the wide-
spread acceptance of the newer theory are as follows :

1. Recent investigations of electrolytes in nonaqueous
solvents having different dielectric constants* have shown
that if ionization were due to the dielectric effect of the
solvent alone (as was formerly believed) it would require a
solvent of much higher dielectric constant than any sub-
stance known.

2. It has been found that such substances as hydrogen
chloride (which is nonpolarf in the solid, liquid or gaseous
state) are electrolytes only when dissolved in certain sol-
vents. The molecules of these solvents are all character-
ized by containing an atom or atoms having one or more
unshared electron pairs which are capable of holding
protons (or H + ) and forming fairly stable complex ions in
this manner. For example, in benzene, C 6 H 6 ; ethyl
bromide, C 2 H 6 Br; toluene, C 7 H 8 ; and similar solvents,
hydrogen chloride exhibits none of the properties of an acid
and does not conduct electricity. No evidence of ionization
has been found. On the other hand, solutions of hydrogen

* It lias long been known that if two oppositely charged bodies were
placed near each other they exerted a force of attraction upon each other.
The magnitude of the force, however, depends, among other things, upon
the nature of the medium between the two bodies. Those substances that
diminish the force to the greatest extent, are said to have high dielectric
constants. Thus, the dielectric constant of a given substance is a measure
of its ability to weaken the force of attraction between charged bodies, if
placed between them.

The basis for the theory that solvents of high dielectric constant should
be the best ionizing solvents should be self-evident when it is recalled that
ions are charged bodies.

t There are two meanings that may be applied to the term polar molecule.
In one case, the term applies to molecules in which the atoms are so arranged
that one side of the molecule has a surplus of positive charges while the
other side has a surplus of negative charges. In the second case, the term
applies to molecules having ionic valences.

The term nonpolar as used here refers to molecules in which all valences
are of the shared-electron-pair type.

MODERN THEORIES OF ELECTROLYTES 61

chloride in water, ether, liquid ammonia, methyl alcohol,
etc., are excellent conductors of electricity and give many
other evidences of containing a highly ionized solute.

The Raman spectra and the infrared absorption spectra
of these solutions show no evidence of hydrogen chloride
molecules whereas solutions of hydrogen chloride in ethyl
bromide and other such solvents give distinct spectra of the
molecules of solute (42).

Molecules of water, ammonia, ether, and methyl alcohol
have the following electronic arrangements to which has
been attributed their ionizing effect. The position taken
by the proton is marked by X.

X H X X

H:0:H X:N:H C 2 H B :O:C 2 H 5 CH 3 :O:H
H

water ammonia ether methyl alcohol

3. To ionize gaseous hydrogen chloride in a discharge
tube requires a large amount of energy. Presumably, the
ionization of hydrogen chloride in solution would require
the same quantity of energy, and the only possible source
of so much energy is that set free by some reaction in the
solution such as that postulated by the new theory.

There are many more reasons for accepting the theory of
solvated ions. However, this is sufficient to give some idea
of the reasons for the wide acceptance of this conception of
ionization.

In the light of this discussion, it is necessary that the
mathematical relationships formulated earlier be revised
slightly. Thus, the ionization of the weak electrolyte,
acetic acid, should be represented as

HAc + H 2 O ^ H 3 O+ + Ac~
The equilibrium expression, therefore, should be written

(H 3 Q+) X (Ac-)
(HAc) X (H 2 0)

62 THEORY OF ANALYSIS

However, the concentration of the water molecules may be
considered as varying so little in different solutions as to
remain constant. Consequently, the above expression can
be changed (as was done for the equilibrium expression
treating water as a weak electrolyte) to the following:

(H 3 U+) X (Ac~) _

(HAc) ~ a

It has been shown experimentally that all the hydrogen
ion present is in the form of hydronium ion. Since the
term H 3 + is expressed in mols, the numerical value will be
the same as if expressed as H+ . Consequently, the value
for K a will be the same as that previously given, i.e.,
1.86 X 10-*.

By similar reasoning, the expression for the ion product
of water is shown to be

(H 3 0+) X (OH~) = K w

and the pH of a solution can be calculated by changing the
expression for this factor to

pH - log

Salt Effect; the Debye-Hiickel Theory (44). It has been
pointed out that the solubility product relationship does
not hold strictly true when the total ion concentration of
the solution is high. If, to a solution saturated with a
slightly soluble salt, a highly soluble salt is added, the
solubility of the slightly soluble salt increases as the con-
centration of the more soluble salt is increased. For
example, it was found by Bray and Winninghoff (43) that
thallous chloride, T1C1, has a solubility of about 0.016 mol
per liter in pure water while in 0.1 M KNO 3 solution, which
has no ions in common with those from T1C1, it has a
solubility of 0.019 mol per liter.

MODERN THEORIES OF ELECTROLYTES 63

This phenomenon, known as the salt effect, seems more
pronounced with some salts than with others, the difference
depending little on the ions added, but much more on the
valence types to which they belong. For example, a salt
consisting of two bivalent ions (e.g., MgSO 4 ) has a much
greater salt effect than does an equivalent concentration of a
uni-univalent salt such as KNOs; a salt consisting of two
trivalent ions has more effect than the bi-bivalent type;
and so forth.

A similar phenomenon is noted with weak electrolytes.
It has been found that in strong solutions of weak elec-
trolytes or solutions of weak electrolytes containing a high
concentration of any salt, the ionization constant has a
greater numerical value than normal.

Even strong electrolytes are influenced by the salt effect.
Conductivity measurements on solutions of NaCl and other
strong electrolytes show an increase with dilution; but,
instead of the conductivity gradually increasing with dilu-
tion as in the case of weak electrolytes, with strong elec-
trolytes it quickly reaches a maximum. This was formerly
thought to be due to an increase in the degree of ionization,
as in the case of weak electrolytes. This explanation,
however, left much to be desired, for it supplied no explan-
ation for the apparently great increase in ionization of
strong electrolytes with dilution as compared to the more
gradual increase exhibited by weak electrolytes. Work
on such solutions finally led to the formulation of the
Debye-Huckel theory, the essential points of which are as
follows :

1. Strong electrolytes (such as NaCl and KC1) are com-
pletely ionized even in strong solutions. This idea is borne
out by the fact that X-ray examination of the crystals of
these two salts shows them to be ionized even in the solid
state. There is no justifiable reason, therefore, for believ-
ing that when these substances are placed in solution the

64 THEORY OF ANALYSIS

ions forsake their greater freedom for an existence as
molecules.

2. Each ion is surrounded by a " sphere of influence "
(due to its electrical charge) which will tend to act on any
ion of opposite charge passing through that region.

3. The size and strength of the sphere of influence of
bivalent ions are greater than those of monovalent ions, and
so forth.

4. As a given ion moves through the solution, it passes
through the fields of attraction of many other oppositely
charged ions with the result that there is a mutual " drag-
ging" effect which slows the movement of the ions and
reduces their activity. This effect is greatest in concen-
trated solutions, diminishing to a negligible effect in very
dilute solutions.

According to this theory, therefore, the ions are restrained
in their movements in concentrated solutions. This makes
them act as if their concentration were lower than it really
is.* With these facts in mind, Debye and Hlickel, and
Bronsted and La Mer were able, for many solutions, to
calculate an effective concentration or activity which, substi-
tuted in the ordinary equilibrium expressions, enabled them
to calculate ionization constants that held true over a much
wider range of concentrations than was formerly possible, t

* It should be obvious, now, why the solubility product relationship does
not apply, with any mathematical accuracy, to saturated solutions of more
soluble salts. For example, at 20C. a saturated solution of NaCl contains
approximately 6.2 mols of solute per liter. If complete ionization is assumed,
such solutions contain 6.2 mols per liter of each, Na + and Cl~~. At much
lower ion concentrations than this, the salt effect is very strong. Conse-
quently, the usual calculations will be thrown off in such solutions.

t The original formulas by which activities were calculated are too com-
plicated for presentation here. A simplified formula, however, developed
by Bronsted and La Mer (62), is

In/ = Q.5z\z>2\/n

where/ is the activity coefficient, zi, and zz are the number of charges on the
anion and cation, respectively, and /* is the ionic strength of the solution.
The latter is a measure of the "drag" exerted (Continued on opposite page)

MODERN THEORIES OF ELECTROLYTES 65

Coprecipitation Phenomena. Imperfect reports are by
no means uncommon for students in qualitative analysis,
but knowledge of some of the causes of such mistakes will
help reduce them to a minimum. There are many such
causes (most of which will be found in the notes on the
laboratory procedure) but some of the greatest errors are
due to the phenomena classified collectively under the term
coprecipitation.

1. Inclusion. Certain gelatinous precipitates, such as
aluminum and ferric hydroxides, consist of extremely tiny
crystals intermeshed in such a fashion as to form minute
pockets in which are enclosed small portions of the original
solution. Washing does not remove these protected por-
tions of solution; and when the precipitate is redissolved,
the resulting solution will be contaminated with ions that
should have remained in the original filtrate. In the
analytical procedure used in some texts, ferric, chromic, and
aluminum hydroxides are formed from hydroxyl ions fur-
nished by the hydrolysis of ammonium benzoate. This
method (recommended by Kolthoff) results in a less
gelatinous, more crystalline precipitate and reduces the
error due to inclusion.

2. Adsorption. Certain precipitates (especially the gela-
tinous hydroxides and certain metallic sulfides) have large
exposed surfaces which are capable of adsorbing other ions

by the ions upon each other. For simple salts M can be calculated from
the formula

where C\ and Ci are the concentrations of the an ions and cations respectively,
and the other terms have the meanings given above. The relationship
between the actual concentration of an ion C and its activity a is given by
the expression

a =/X C

By using the activity instead of true concentrations, calculations involving
solutions of weak electrolytes or slightly soluble substances will be accurate
over a much wider range of concentrations.

66 THEORY OF ANALYSIS

in large quantities. This surface phenomenon is not
thoroughly understood, being due, probably, either to a
surface reaction or to a preponderance of one type of
valence over another which enables the ions on the crystal
face to attract ions of opposite charge and hold them to its
surface. Thus, ions which would otherwise remain in solu-
tion are carried down with the precipitate, often in such
concentrations as to give misleading results.

3. Mixed-crystal Formation. Most of the trouble caused
by coprecipitation is due to the formation of mixed crystals
of isomorphous compounds. Isomorphous compounds are
those having the same type of crystalline structure and
atoms of nearly the same size. Cadmium sulfide and
barium sulfide are isomorphous, as are also sodium chloride
and potassium chloride. When cadmium sulfide is pre-
cipitated in Group II, some barium ions occasionally attach
themselves, momentarily, to the sulfide ions on the surface
of the cadmium sulfide. Many of these barium ions escape
into the solution again, but some of them fail to escape
before more cadmium and sulfide ions attach themselves to
the surface of the solid and enclose the barium ions, pre-
venting their going back into the solution. This results
riot only in contaminating the precipitate with barium ion
but also, in extreme cases where the cadmium ion concen-
tration is much higher than that of the barium ion, in
removing so much barium ion that the test for the latter in
Group IV is missed entirely. In case the cadmium ion con-
centration is extremely high, therefore, it is often necessary
to redissolve the Group II precipitate in as small a volume
of acid as possible and reprecipitate with hydrogen sulfide.
In this more concentrated solution, some of the barium ion
will escape coprecipitation and the two filtrates, combined
and concentrated by evaporation, may be used for the
barium test.

Occasionally, however, coprecipitation can be of service
to the analyst. For example, the tests for zinc, copper and

ORGANIC COMPOUNDS IN ANALYSIS 67

cobalt, using ammonium mercuric thiocyanate reagent,
depends, for their greatest sensitivity, on the formation of
mixed crystals. The test is made for cobalt or copper by
adding zinc ion to the solution and then adding the reagent.
In the absence of cobalt or copper the following reaction
occurs ;

ZD++ + Hg(SCN)r - ZnHg(SCN) 4 (white ppt.)

If cobalt is present, the following additional reaction occurs:

Co++ + Hg(SCN)r -> CoHg(SCN) 4 (blue ppt.)

If both ions are present, mixed crystals of these two com-
pounds will form, the precipitate having a blue color, the
depth of which will be greater, the larger the concentration
of cobalt ion. It is found that, carried out in this way, it is
possible to detect cobalt in solutions that would give no
precipitate whatsoever with the reagent in the absence of
zinc.

In the case of the copper and zinc tests, the same type of
reaction is involved. However, if the copper is present in
concentrations greater than the zinc, an apple-green pre-
cipitate of copper mercuric thiocyanate will form.

ORGANIC COMPOUNDS IN ANALYSIS

Organic Reagents.* It has been found that many
organic compounds give tests for metallic ions that are
more specific and sensitive than those given by any known
inorganic reagent. The reactions involved are often quite
complicated and difficult for one not studied in organic
chemistry to comprehend. Occasionally the reaction is
of the ordinary double decomposition type, but more often
it involves coordinate linkages in addition to the ordinary
type.

* Note to teacher: It is not intended that students who have had no organic
chemistry should attempt the study of this section. This material is
included only for those sufficiently advanced in chemistry.

68 THEORY OF ANALYSIS

With few exceptions organic analytical reagents may be
divided into two groups :

1. Those reacting only through coordinate valences.

2. Those having one or more acidic groups, or an acidic
group and an atom having an unshared pair of electrons.
The acidic groups that are the most important are:

a. The oxime, or =NOH group ( --NOH ^ =NO~ +
H+)

6. The hydroxyl, or OH group

c. The mercaptan, or ^SH group

d. The imino, or ^NH group (=NH ^ =N~ + H f )

e. The carboxyl, or COOH group
/. The amido, or =CONH 2 group
g. The sulfonic, or SO 3 H group

These are all important in analytical work, with the
exception of the carboxyl and sulfonic groups.

Of the many organic reagents, only four of the most
common will be discussed here.

Rhodanine, 5-(p-dimethylaminobenzal) rhodanine, is
one of the reagents used for the identification of mercury,
silver, etc. The reaction with silver may be written as
follows (45, 46) :

HN CO

Ag+ + SO C=CHC 6 H 4 N(CH 3 ) 2 ->

\ /

Ag N CO

Si=C C=CHC 6 H 4 N(CH 3 ) 2 + H+

\ /

It will be noted that this is the case of an imido group
reacting in the normal fashion, together with the formation
of a four-membered ring one part of which is a coordinate
linkage between the metal and the sulfur. The formation

ORGANIC COMPOUNDS IN ANALYSIS 69

of coordinate bonds with reagents is much simplified by
considering that the metallic ions, in such cases, are really
forming complex ions of much the same type considered
earlier. It will be recalled that molecules which readily
form complex ions always have one or more atoms upon
which there are unshared electron-pairs.

Organic reagents that depend upon the use of coordinate
linkages always have one or more atoms of this type.
Consequently, the reaction of a reagent such as 5-(p-di-
methylaminobenzal) rhodanine is partly a normal double
.decomposition reaction and partly a type of complex ion
formation.

It will be found that the reaction of Ni+ + with dimethyl-
glyoxime is of this same type. In compounds where such
rings are possible, it will be found that, in accordance with
the Baeyer Strain Theory, if two rings are possible, that
one which comes closer to giving a six-membered ring will
be the most probable and the most stable, if formed.
Five- and six-membered rings are the most stable of all,
while three-membered rings are almost unknown. The
atoms which may make up the ring are carbon, oxygen,
nitrogen, sulfur, and the metallic atom. Coordinate
linkages, of the type shown above, connect the metallic
atom to either a sulfur, a nitrogen, or an oxygen atom
seldom, if ever, to the carbon as these three atoms each
have one or more pairs of unshared electrons.

a-Benzoin-oxime (cupron) is usually used for copper.
The reaction involved is one of simple double decomposition
as is shown in the following equation:

HOCHC 6 H 5 Cu CHC e H 5
GU+++ | -> | | +2H+

HO CC 6 H 5 O CC 6 H 5

\ / \ /

N N

a-Nitroso-jS-naphthol is used for cobalt. It is an
excellent reagent, as nickel, unless in unusually large con-

70 THEORY Or ANALYSIS

centrations, does not interfere. The reaction is (45, 47,
48):

3C 10 H 6 ONOH + Co+++ -* (CioH 6 ONO) 3 Co + 3H+

Thus, this reagent acts like a normal acid reacting with
metallic ions.

The use of dirnethylglyoxime (diacetyl dioxime) as a
reagent for nickel, depends upon the following reaction (45,
48, 49, 50) :

rCH 3 0-NOH]
2 |

LCH 3 C= NOHj

CH 3 CNO ONCCH 3

\ / + 2H+

CH,CN "NCCH,

OH OH

This reagent may be used for copper, bismuth and many
other elements, the reaction being quantitative as well as
qualitative, in many cases.

The use of aluminon and alizarin S (blue) in the test for
aluminum is not well understood as yet. Engelder (11, 48)
suggests that the dyestuff forms chelate rings as in the above
cases, making use of coordinate linkages and forming
definite compounds with the aluminum hydroxide. Others,
however, believe that these dyes are merely adsorbed on the
surface of the aluminum hydroxide and thus render the
precipitate more visible.

PART III

ANALYTICAL PROCEDURE CATIONS
GROUP SEPARATION CATIONS

The usual sample of unknown may contain any number
of metals which will give such varied results with a given
reagent as to make it impractical to identify any one of
them as long as certain others are present. For this reason
it is necessary to treat the mixture in such fashion as to
separate the ions present into groups, each containing only
certain metals, and then further to subdivide each group
until each metal is obtained in a solution free of interfering
ions. This final solution is then tested with the proper
reagent to show the presence or absence of that particular
metallic ion. It is obvious, therefore, that such separations
must be as complete as possible, for if a small amount of
some metallic ion gets into the wrong group, false results
are almost certain to be obtained. For this reason the
following rules must be observed.

1. Every detail of the instructions must be followed in
each step of the procedure. Appended notes must be
carefully read, as they contain important hints that will
save much trouble.

2. Clean apparatus and distilled water must be used at
all times since many of the tests are extremely sensitive and
a trace of foreign matter may ruin the whole analysis.

3. All precipitations must be complete in other words,
care must be taken to add enough reagent to react with all
the ions that are to be precipitated at that point. Failure
to do this means that the analysis of subsequent groups will
be ruined owing to the presence of the ions that have not
been precipitated because of insufficient reagent. On the

72 ANALYTICAL PROCEDURE CATIONS

other hand, excess reagent must be avoided, loo, as in
many oases an excess of the reagent will redissolve all, or
part of, the precipitate, with the result that the metallic ion
appears in the next group and ruins the separation. To
avoid adding either excess or insufficient reagent, it is
desirable to use the following procedure.

First, add a couple of drops of reagent and stir the mixture
thoroughly. If a precipitate forms, centrifuge, then to the
clear supernatant liquid add another drop of reagent. If a
precipitate forms in the liquid, add one more drop of
reagent and stir once more. Centrifuge and repeat the
above procedure until, on addition of a drop of reagent, no
more precipitate forms in the clear layer of liquid. This
marks the point when exactly enough reagent has been
added.

If, on addition of the first two drops of reagent, no
precipitate forms at once, it is well to rub the inside of the
container with a stirring rod for about 1 min. and let the
solution stand for a few minutes as the solution may be
supersaturated .

4. Tongs and other metallic objects must not be touched
to the tops of containers as acid fumes may attack them and
a little of the resulting compound may get into the unknown
solution and contaminate it.

5. Droppers must not be dipped into reagent bottles, nor
must the droppers in the pipette-type reagent bottles ever
be touched to the vessels used for analysis. A single failure
to observe this rule may easily ruin all analyses -your own,
as well as those of the rest of the class for the rest of the
semester. The dropper will almost certainly pick up some
contamination and, when placed in the reagent again, will
contaminate it too.

Before a complete analysis can be made, it is necessary
that the unknown be in solution. In class work, it is
customary to issue the unknown in the dissolved state.
However, in industrial work such is often not the case, and

GROUP SEPARATION CATIONS 73

it is necessary to dissolve the material. A solid sample
must, therefore, be treated first according to the procedure
outlined in Part A. It is important that each step be
carried out in the order given, as otherwise complications
will arise.

Part A. Treatment of Solid Samples. 1. To about
50 mg. of the powdered solid unknown add 10 drops cold
water. Heat almost to the boiling point, stir, and let stand
until no more solid disappears. Filter, and save the
liquid. If much solid seems to dissolve, repeat the treat-
ment with water, combining the filtrates (see Note 1).

2. To any residual solid from (1) add 67V HNO 3 and treat
as before. Again filter and save the liquid (see Note 2).

3. To any residual solid from (2) add concentrated
HNO 3 , and repeat the rest of the procedure in (2).

4. To any residual solid from (3) add aqua regia (made
by mixing 6 drops dilute HC1 with 2 drops concentrated
HNO 3 ) and repeat the procedure outlined before (see Notes
3 and 4).

5. Any dark residual solid from (4) should be carefully
dried in the air bath and cooled. Mix this with three times
its volume of a mixture of 1 part Na 2 O 2 and 3 parts Na 2 CO 3 .
This mixture is then placed in a porcelain crucible (see
Note 5) and heated over a strong Bunsen flame, to redness,
holding it at that temperature for 10 to 15 min. The fused
mixture is then allowed to cool, after which it is pulverized
with the tip of a heavy glass rod and the resulting powder is
put through the first three steps outlined earlier. Any
residue remaining after that may, as a rule, be discarded
(see Note 6).

Part B. Separation of a General Unknown into Groups.
The separation of the metallic ions in a solution into groups
for analysis is based upon five properties of the metallic
compounds.*

* It would be impractical to cover the analysis of all known metals in a
text of this scope. For that reason, only those metals listed here will be

74 ANALYTICAL PROCEDURE CATIONS

1. The chlorides of all metals except silver, mercurous
mercury, Hg 2 " t " f , and lead are fairly soluble.

2. The sulfides of Hg, Cu, Pb, Cd, Sb, Sn, As, and Bi are
insoluble in Q.3N HC1 whereas the sulfides of the other
metals are soluble in this concentration of acid.

3. The sulfides of Fe, Ni, Co, Mn, and Zn are soluble in
0.3N HC1 but are insoluble in very slightly basic solution.
The sulfides of Ba, Ca, Sr, K, Na, and Mg are soluble in
either acid or slightly basic solutions.

4. The hydroxides of Al and Cr are precipitated by the
basic solution required for the precipitation of the sulfides
of Fe, Ni, Co, Mn, and Zn. The hydroxides of Ba, Ca, Sr,
Mg, Na, and K are not precipitated by this weakly basic
solution, the latter two being very soluble even in strongly
basic solutions.

. 5. The carbonates of Ba, Sr, and Ca are only slightly
soluble in either water or basic solutions buffered with
ammonium salts. The carbonate of Mg is insoluble in
water but is soluble in highly buffered basic solutions.

These properties enable us to separate the metals into
five groups. The groups consist of the following metals:

Group I. Ag+ Hg 2 ++ Pb++

Group II. Cu++, Pb++, Hg++, Cd+ + , Sb+++, Sb+++++,
AS+++ AS+++++ Bi+++ Sn++ Sn f4 ++

Group III. Fe+++ A1+++, Cr+++, Ni++, Co++, Mn++ , Zn++

Group IV. Ba++, Sr++, Ca++

Group V. Mg++, Na+, K+, NH 4 +

The method of precipitating these metals into their
respective groups and of further separating and identifying
them will be outlined in succeeding pages. The procedure
will be given in outline form for clarity, and, for economy of
space, abbreviated sentences will be used throughout.

studied, their analysis serving to cover those metals most commonly expected
in commercial work, as well as serving to illustrate the principles involved
in analysis. For a more comprehensive scheme of analysis see "Analytical
Chemistry, Qualitative Analysis" vol. I, by Treadwell and Hall, John
Wiley & Sons, Inc., New York.

GROUP I METALS 75

The outline will describe the method of precipitating each
group and analyzing it. To analyze a general unknown, it
is necessary only that the solution left from the Group I
precipitation be used as the unknown for the Group II
analysis; the solution from the Group II precipitation, for
the Group III unknown, etc.

For a general unknown, no more than 1 ml. of unknown
should be taken. More will make the analysis difficult.
In the case of group unknowns, 8 to 10 drops is sufficient.

GROUP I METALS
CHEMICAL CHARACTERISTICS

Silver. Metallic silver is rather inactive chemically. It
does not react with dilute nonoxidizing acids or with
alkalies. It reacts with aqua regia, but forms a mass of
AgCl that is very difficult to dissolve. The best solvent
for this metal is HNO 3 .

Compounds of Silver. Silver forms a large number of
complex ions and slightly soluble compounds. A few of its
most characteristic reactions are: Soluble chlorides give
AgCl, white, insoluble in water or HNO 3 , but soluble in
NH 4 OH forming Ag(NH 3 ) 2 + , in concentrated HC1, KC1 or
NaCl solution forming AgCl 2 ~, and in alkali cyanides
forming Ag(CN) 2 ~ ions. H 2 S precipitates Ag 2 S, black,
the least soluble of the ordinary silver compounds, being
insoluble in water, HC1 or NH 4 OH, but soluble in dilute
HN0 3 and in concentrated alkali cyanides. The fluoride is
very soluble in water. Bromides precipitate AgBr, pale
yellow, insoluble in water or HNO 3 ; slightly soluble in
NH 4 OH, very soluble in alkali cyanides. Nal or KI pre-
cipitates Agl, yellow, insoluble in water, HNO 3 , or NH 4 OH;
soluble in alkali cyanides. All the halides of silver are
soluble in solutions of Na 2 S 2 3 . Na 2 CO 3 precipitates
Ag 2 CO 3 , white, insoluble in water; soluble in NH 4 OH and
HNO 3 . NaOH, KOH and NH 4 OH give AgOH (2AgOH ^
+ H 2 O) black or brown, insoluble in water but

76 ANALYTICAL PROCEDURECATIONS

soluble in HNO 3 or in excess NH 4 OH. Ag 2 O is a good
oxidizing agent, being itself reduced to metallic silver.

Lead. Metallic lead is attacked by all the common acids.
Owing to the fact that most lead compounds are insoluble,
however, it is best dissolved in dilute HNO 3 . The nitrate
and acetate are its most soluble compounds.

Compounds of Lead. The following reagents, added to
solutions of Pb(NO 3 )2 give compounds that are insoluble
in water: H 2 S gives PbS, black, insoluble in water, NH 4 OH ;
dilute HC1, (NH 4 ) 2 S, and Na 2 S*; soluble in 2N HNO 3 and
in concentrated HC1.

Soluble chlorides give PbCl 2 , white, only slightly soluble
in cold water, but soluble in hot water and in strong HC1.
Iodides give PbI 2 , yellow, and less soluble than PbCl 2 ;
soluble in concentrated alkali iodides, forming PbI 3 ~ ions.

KOH or NaOH precipitates Pb(OH) 2 , white, insoluble
in water or NH 4 OH; soluble in HNO 3 , acetic acid, or excess
alkali (forming HPbO 2 ~); forming insoluble, brown PbO 2
on addition of H 2 O 2 , hypochlorites, or other oxidizing
agents.

H 2 SO 4 or sulfates precipitate PbSO 4 , white, insoluble in
water; soluble in concentrated H 2 SO 4 [forming soluble
Pb(HSO 4 ) 2 ], in strong ammonium or sodium acetate solu-
tion [forming slightly ionized Pb(C 2 H 3 O 2 ) 2 ], and in strong
alkalies.

K 2 CrO 4 and other soluble chromates precipitate PbCrO 4 ,
yellow, insoluble in water or acetic acid; soluble in HNO 3
and alkalies.

Mercury. Metallic mercury is very inactive, being
between copper and silver in the displacement series. As
it is below hydrogen, it does not react with dilute H 2 SO 4 or
HC1. It dissolves readily, however, in hot HNO 3 to form
Hg(NO 3 ) 2 , or in cold HNO 3 to form Hg 2 (NO 3 ) 2 .

HI dissolves mercury readily, forming hydrogen and the
slightly dissociated Hglr ion. Hot, concentrated H 2 S0 4
dissolves mercury forming the sulfate and S0 2 .

GROUP I METALS 77

Compounds of Mercury. Mercury forms compounds of
two types mercurous, Hgo 4 "*, and mercuric, Kg" 1 "*. The
nitrates of both are soluble, but in many ways their
reactions are so different as to make it necessary to consider
them separately.

Mercurous Compounds. Soluble chlorides precipitate
HgoClo, white, insoluble in water or dilute acids; soluble in
hot, strong HNO 3 or H 2 SO 4 . With NH 4 OH, Hg 2 (NO 3 ) 2
gives a mixture of white Hg 2 O(NH 2 )(NO 3 ) and metallic
mercury. Alkali carbonates give a yellow precipitate of
Hg 2 CO 3 , decomposed on boiling or standing, into a gray
mixture of Hg and HgO. H 2 S precipitates a mixture of
HgS and Hg, the whole precipitate being soluble in NaaS*.
KI precipitates green Hg 2 I 2 , which reacts with excess KI
to form HgI 4 = and Hg. Copper displaces Hgo +4 " ions from
their compounds.

Mercuric Compounds. Soluble iodides precipitate HgI 2 ,
red, insoluble in water but soluble in excess alkali iodide to
form HgI 4 = . HgBr 2 is more soluble in water than is HgI 2 .
HgCl 2 is fairly soluble. Hg(NH 2 )Cl, white, insoluble in
water, is formed by treating Hg 2 Cl 2 or HgCl 2 with NH 4 OH.
It is soluble in hot NH 4 C1 [forming Hg(NH 3 ) 2 Cl 2 ] and in
acids.

HgO, red or yellow, is insoluble in water but soluble in
acids. It is precipitated when hot solutions containing
Hg+ + are treated with KOH or NaOH, or with Na 2 CO 3 or
K 2 CO 3 . Unstable above 300C. giving metallic mercury
and oxygen.

HgCrO 4 , orange, is insoluble in water, soluble in acids.
It is precipitated by adding neutral alkali chromates to
solutions containing neutral or acid-buffered Hg 4 ^".

KSCN precipitates Hg(SCN) 2 , white, insoluble in water;
soluble in excess KSCN or NH 4 SCN, forming Hg(SCN)r.
H 2 S precipitates HgS, black, insoluble in water, dilute
acids, concentrated HC1 or (NH 4 ) 2 S. Soluble in hot, con-
centrated HNO 3 (with long boiling), in aqua regia, in yellow

78 ANALYTICAL PROCEDURE CATIONH

(NH 4 ) 2 Ss, and in Na 2 S (forming HgSr from which it is
rcprecipitated by acidification). When it is precipitated
from solutions of Hg ++ , the first substance to form is a
white, insoluble double salt of the type, (HgS) 2 -HgCl 2 .
This is slowly converted, by more H 2 S, into black HgS.
Dissolving HgS in HNO 3 produces (HgS) 2 -Hg(NO 3 ) 2 as an
intermediate product, this slowly dissolving to form the
normal nitrate.

ANALYTICAL ASPECTS

It will be recalled that the chlorides of the three ions
Ag 4 ", Hg 2 ++ , and Pb" 1 "^ are insoluble whereas those of
other metallic ions are soluble. It is possible, therefore,
to separate these three metals from the others in a general
unknown by adding chloride ions to the solution. This
precipitates the chlorides of lead, silver, and mercurous mer-
cury, leaving the other metallic ions in solution. A solu-
tion of HC1 will furnish the necessary chloride ions without
adding anything undesirable to the solution. Too great
excess, however, will redissolve part of the precipitate,
forming PbCl.r~ and AgCl 2 ~.

The further separation and identification of these three
metals depend upon (1) the high solubility of lead chloride
in hot water and the low solubility of lead chromate ; (2) the
high solubility of silver chloride in ammonium hydroxide
and its reprecipitation on the addition of nitric acid; and
(3) the reaction between mercurous chloride and ammonium
hydroxide to form a mixture of free mercury and mercuric
aminochloride (HgNH 2 Cl).

PRELIMINARY EXPERIMENTS GROUP I

1. Silver ion, Ag 4 ". a. To 2 drops of a solution con-
taining Ag + ions, add 1 drop IN HC1 solution. Filter*

* See Semi-micro Technique Filtrations. Use the technique described
there, whenever the instructions say to filter.

GROUP I METALS 79

and try dissolving a small portion of the precipitate from
(a) in hot water.

Try dissolving a second portion of the precipitate in
concentrated HC1.

6. To another portion of the precipitate from (a) add
2 drops QN NH 4 OH and stir. If any solid refuses to dis-
solve, filter. Place 1 drop of the clear filtrate on a small
watch glass, place the latter on a black surface, and add
2 drops QN HN0 3 .

To another drop of the ammoniacal solution add 1 drop
(UNKI solution.

To a third drop of the ammoniacal solution add 1 drop
rhodanine solution and 2 drops 67V HNO 3 . This test is
capable of detecting as little as 0.001 mg. of Ag + (51, 13).

c. To 1 drop Ag + test solution add 1 ml. water and stir
thoroughly. Now place 2 drops of this solution in a micro-
beaker and add 1 drop 3N K 2 CrO 4 solution.

Place another drop of the diluted test solution on a piece
of filter paper. To the same spot add 1 drop 3AT K 2 CrO 4
solution. Compare the two results (Question 1).

2. Lead Ion, Pb ++ . a. To 2 drops Pb ++ test solution
add 1 drop IN HC1 solution. Filter and wash the precipi-
tate in cold water.

To a small portion of the precipitate add 2 to 3 drops cold
HC1.

To the rest of the precipitate add 5 drops hot water.
Heat on the steam bath to make sure the solution is hot
and notice whether lead chloride seems to be soluble in hot
water (Question 2).

b. To 2 drops of the hot solution from (a) on a watch
glass add 1 drop QN H 2 SO 4 . Place against a black back-
ground for observation purposes.

c. Repeat (6), using 3N K 2 CrO 4 solution instead of H 2 SO 4 .

d. To half of the PbCrO 4 from (c) add 2 drops QN HNO 3 .
To the second half of the precipitate add 2 drops QN NaOH
solution.

80 ANALYTICAL PROCEDURE CATIONS

e. To 1 drop Pb ++ test solution on a filter paper add in
the order mentioned 1 drop H 2 O 2 and 1 drop NH 4 OH.
Warm for 1 min. in a jet of steam from the steam bath.
Examine the color of the spot on the paper. The reaction
that has taken place is :

Pb++ + 2(OH~) + H 2 2 -* Pb0 2 + 2H 2

Now, to this spot add 1 drop benzidine reagent. Note
the new color that appears. PbO 2 is an oxidizing agent;
and, when benzidine reacts with an oxidizing agent, it
forms a complex organic dye of the color observed. Con-
sequently, this test is not specific for lead any oxidizing
agent will give a similar test. Taken at this point, how-
ever, with other elements eliminated, it is a good confirm-
atory test for lead, being capable of detecting 0.0015 mg.
lead (14).

/. To 2 drops Pb+ + test solution add 1 drop 0.1 AT KI
solution. Heat to boiling, noting all occurrences, then cool.

3. Mercurous Ion, Hg 2 ++ . a. To 2 drops Hg 2 ++ test
solution add 1 drop IN NH 4 C1 solution. Filter and wash
the precipitate with cold water.

Try dissolving part of the precipitate in 5 drops hot
water.

Try dissolving another portion of the precipitate in
2 drops QN NH 4 OH.

6. To some of the precipitate from (a) add 2 drops
QN NH 4 OH. Note what occurs. Filter and wash pre-
cipitate in cold water.

To this precipitate add 8 drops aqua regia (a mixture of
1 part concentrated HNO 3 with 3 parts QN HC1) and
evaporate to dryness. To the residue add 3 drops of a
saturated solution of sodium acetate, stir, and add 2 drops
diphenylcarbazide reagent. Note the results (see Note 8
and Question 3).

c. To 4 drops Hg^" 1 " test solution add 1 drop 0.1 N KI
solution. Note the color of the precipitate formed. Now

GROUP I METALS

GROUP I ANALYSIS

For the systematic analysis of an unknown, the following procedure may
be followed (see Note 84).

To the group or general unknown, add 2 drops IN HC1 solution, stir
thoroughly and centrifuge. Test for completeness of precipitation and add
more reagent if necessary (see Note 9). Stir once more, then filter (centri-
fuge). Wash precipitate with cold water.

Filtrate. Save
Groups IT,
and IV.

for
III

Precipitate (see Note 10). Consists of PbCl 2 ,
Hg 2 Cl 2 and AgCl. Add 5 drops boiling water and
heat (while stirring) on steam bath until almost-
boiling (see Note 11). Centrifuge quickly and imme-
diately remove filtrate?. Repeat washing until wawh
water gives no test for lead.

Filtrate. Contains PbCl 2 .
While still hot, put 1
drop on a filter paper
and add 1 drop 3%
H 2 2 , 1 drop dil.
NH 4 OH, heat over
steam bath for 1 miii.
t hen add 1 drop ben-
xidine reagent.

BLUE COLOR
indicates Pb ++ present.

Precipitate. Consists of Hg 2 Cl 2 and AgCl. Add
3-4 drops of dil. NH 4 OH and stir. Centrifuge.

Filtrate. Contains Ag-
(NH a )Cl. To 1 drop of
filtrate on a watch glass
over a black background
add 1 drop cone. HNOa.

WHITE CLOUDY PPT.

indicates Ag * present.

Precipitate.
BLACK or GRAY
PPT.
of Hg and HgNH 2 Cl
proves Hg 2 ++ present.

To the precipitate add
5 drops freshly pre-
pared aqua regia and
evaporate to dry ness.
Dissolve residue in 5
drops saturated so-
dium acetate soln.,
add 1 drop dil. acetic
acid and 1 drop of a
freshly prepared alco-
holic soln. of gallic
acid.

YELLOW or
ORANGE PPT.

confirms Hg 2 " f+ pres-
ent (see Note 12).

To 1 drop of filtrate on a
spot plate add 1 drop
rhodanme reagent and 1
drop GAT HNO,.

RED PPT.
confirms Ag 4 " present.

To 1 drop K 2 CrO 4 soln.
on a black back-
ground add 1 drop of
the hot filtrate.

YELLOW PPT.*
proves Pb ++ present.

To 1 drop of filtrate on a
watch glass add 1 drop
0.1 AT Kl soln.

PALE YELLOW PPT.
proves Ag + present.

* In many places it will be necessary to use the abbreviation "PPT " for "precipitate"
in order to conseive space.

82 ANALYTICAL PROCEDURE CATIONS

try to redissolve this precipitate by adding excess KI
(using a saturated solution of KI for this purpose).

d. To 1 drop Hg 2 " f+ test solution add 1 drop IN acetic
acid, 1 drop saturated sodium acetate solution, and 1 drop
of a freshly prepared alcoholic solution of gallic acid
(see Note 12). This test is specific for mercury (Hg 2 4 " f or
Kg** ) in the presence of any of the metals of Groups I and
II except silver, the latter giving a dark coloration.

GROUP II METALS
CHEMICAL CHARACTERISTICS

Copper. Metallic copper is much like mercury in its
reactions. It is insoluble in dilute, nonoxidizing acids in
absence of air, but dissolves slowly in strong HC1, if air is
present, to form the complex cuprochloride ion, Cu^Cl^.
It dissolves readily in HNO 3 and in hot, concentrated
H 2 SO 4 . Copper displaces mercury and silver from their
compounds.

Two types of copper compounds are formed.

Cuprous Compounds. These are similar to silver and
mercurous compounds in many respects. The chloride,
Cu 2 Cl 2 , is white, insoluble in water; soluble in HC1, forming
Cu 2 Cl 4 =:= , and in excess NH 4 OH, forming Cu 2 (NH 3 )4" f+ .
With NaOH, solutions of Cu 2 +4 ~ give a yellow precipitate of
Cu 2 (OH) 2 , decomposed, on boiling, into red Cu 2 O. With
H 2 S, black Cu 2 S is precipitated, insoluble in dilute non-
oxidizing acids but readily soluble in warm HNO 3 to form
Cu(N0 3 ) 2 .

Cupric Compounds. With solutions of Cu^, H 2 S pre-
cipitates black CuS, insoluble in dilute, nonoxidizing acids,
and only slightly soluble in (NH 4 ) 2 S; readily soluble in
warm HNO 3 , or excess KCN [forming Cu(CN) 3 = ions].
With NaOH, solutions of Cu"^" give green Cu(OH) 2 , insol-
uble in water but soluble in excess NH 4 OH [giving deep
blue Cu(NH 3 ) 4 ++], in KCN, and in acids. With Na 2 CO 3 , a

GROUP II METALS 83

basic carbonate is precipitated, soluble in acids and in
NH 4 OH. Its nitrate, sulfate and chloride are very soluble
in water.

Bismuth. Bismuth is a rather inactive metal, insoluble
in HC1 or cold H 2 SO 4 , but soluble in HNO 3 or hot H 2 SO 4 .
The element exhibits two principal valences, +3 and +5.
The trivalent form is the most stable.

Compounds of Bismuth. The compounds of bismuth
hydrolyze readily in dilute solutions giving white precipi-
tates of the type BiOCl or, BiO(NO) 3 . Some characteristic
reactions of bismuth ion solutions are :

NaOH gives white Bi(OH) 3 , insoluble in water or excess
NaOH; soluble in acids.

Alkali carbonates precipitate basic caibonates of the type
Bi(OH)CO 3 . K 2 Cr 2 O 7 gives yellow (BiO) 2 Cr 2 O 7 , soluble
in strong acids but insoluble in alkalies.

H 2 S precipitates Bi 2 S 3 , brown, insoluble in dilute, non-
oxidizing acids or (NH 4 ) 2 S; soluble in hot concentrated
HC1 or in dilute HNO 3 .

Metallic zinc, or stannite ions, HSn0 2 ~, will reduce Bi +++
to metallic bismuth, usually as a black powder. Soluble
iodides precipitate black BiI 3 , soluble in excess to form
yellow BiI 4 ~~.

Strong oxidizing agents such as C1 2 , C1O~, or H 2 O 2 ,
added to alkaline suspensions of Bi(OH) 3 give a brown
precipitate of meta-bismuthic acid, HBiO 3 , insoluble in
water, but a powerful oxidizing agent.

Cadmium. The metal is moderately active, easily
volatilized element (B.P., 767C). On heating, it com-
bines with oxygen to form brown CdO; it dissolves slowly
in HC1 or H 2 SO 4 , and rapidly in HNO 3 .

Compounds of Cadmium. CdS is yellow or orange in
color; insoluble in (NH 4 ) 2 S, Na 2 S*, or KCN; soluble in
3N HC1, HNO 3 , and in hot, dilute H 2 SO 4 .

With solutions of Cd" 4 " 1 ", NaOH precipitates white
Cd(OH) 2 , insoluble in excess alkali; soluble in acids, or in

84 ANALYTICAL PROCEDURECATIONS

NH 4 OH [forming Cd(NH 3 ). t + + ions]. Na a CO, or
(NH 4 ) 2 CO 3 forms white CdCO 3 , insoluble in water; soluble
in acids, in KCN, and in solutions of ammonium salts.
K 4 Fe(CN) 6 precipitates white Cd 2 Fe(CN) 6 . The nitrate,
chloride, and sulfate of cadmium are soluble in water.

Arsenic. Metallic arsenic is a gray very brittle sub-
stance, subliming readily at 615C. giving a characteristic
garlic-like odor. Its vapors and compounds are extremely
poisonous. The element is amphoteric, its oxides dissolv-
ing in acids or in alkalies. The element does not dissolve
in HC1, but dissolves slowly in strong alkalies, and more
rapidly in NaOCl, aqua regia, or concentrated HNO 3 .
At high temperatures it combines readily with sulfur, oxy-
gen, and chlorine and reacts with metals to form arsenides
of the type Zn 3 As 2 .

Compounds of Arsenic. Arsenic exists in both the
trivalent and the pentavalent state. If in alkaline solution,
the arsenic is present in the form of arsenite or arsenate
ions; if in acid solution, as arsenious or arsenic ions. The
two valence types will be discussed separately. Both types
give arsine (AsH 3 , a very poisonous gas) on reduction with
acid and zinc or aluminum, or on reduction with alkalies
and zinc or aluminum. Trivalent arsenic compounds are
the most easily reduced, however.

Arsenious Compounds and Arsenites. As 2 O3, white, is
slightly soluble in water, but very soluble in HC1 foiming
As 4 "" 1 "*, and in alkalies forming AsO 3 % or arsenite ions.
H 2 S causes no precipitation from neutral or alkaline solu-
tions of arsenites, owing to formation of soluble thio-
arsenites. From slightly acid solutions of arsenites, H 2 S
precipitates As 2 S 3 , yellow, insoluble in dilute, nonoxidizing
acids, but soluble (with long boiling) in concent rated HC1,
and readily soluble in concentrated HNO 3 , aqua regia, or
ammoniacal H 2 O 2 , forming AsO 4 ^; in ammonium carbon-
ate, forming a mixture of AsS 3 s and AsO 3 s ; in alkali
sulfides, forming AsSs s ; and in Na 2 S x , forming AsS 4 s .

GROUP II METALS 85

AgNO 3 , with neutral or faintly acid solutions of arsenites
gives Ag 3 AsO 3 , yellow, insoluble in water, but readily
soluble in NH 4 OH or HNO 3 .

AsH 3 (see above) decomposes on heating, giving As and
H 2 . This is the basis of the Marsh test for arsenic.

Arsenic Compounds and Arsenates. As 2 O 6 dissolves
readily in water to form arsenic acid, H 3 AsO 4 , and in
alkalies to form arsenates.

H 2 S does not form any precipitate with neutral solutions
of arsenates. With weakly acid solutions, it will form
As 2 S 6 only after a long period of time. However, if the
solution of arsenate is first made strongly acid with HC1,
yellow As 2 Ss precipitates at once.*

As 2 S6 is insoluble in water or nonoxidizing acids. It is
soluble, however, in the same solvents as those listed for
As 2 S 3 .

AgNO 3 gives, with neutral solutions of arsenates,
Ag 3 AsOi, chocolate colored, insoluble in water but soluble
in NH 4 OH and in acids.

Ammonium molybdate gives (NH 4 ) 3 AsO 4 -12MoO 3 , yel-
low, insoluble in water or dilute HNO 3 , but readily soluble
in NH 4 OH or alkalies. Excess reagent is needed for the
precipitation.

Magnesia mixture (a mixture of MgCl 2 , NH 4 OH, and
NH 4 C1) precipitates white crystalline MgNH 4 AsO 4 , insolu-
ble in water or dilute NH 4 OH. Arsenites give no precipi-
tate with this reagent.

* The failure of H 2 S to precipitate the sulfides from neutral or weakly
acid solutions of arsenates is due partly to the formation of soluble complex
thioarsonate ions (such as AsS4^, AsOaS", etc.), and partly to the fact that
H.<AsO 4 is quite stable and gives a very low concentration of As 44 " 4 H4 ~. A
high concentration of H 4 , however, favors the formation of As 44 " 4 " 4 " 4 " in
accordance with the equation

HaAsO 4 + 5H+ < a AH++ 4 **- + 4H,O

Therefore, since As 2 S6 is insoluble in concentrated HC1, AsaSs precipitates
readily from solutions of strong HC1.

86 ANALYTICAL PROCEDURECATIONS

Slightly acid solutions of arsenates give a yellow precipi-
tate of As 2 S 6 when heated with Na2S2O 3 . Arsenites give a
similar precipitation.

Antimony. Metallic antimony is a brittle, silvery metal,
with a melting point of 630C., a boiling point of 1380C.,
and a density of 6.7. At high temperatures, antimony
combines readily with oxygen to form Sb 2 O 3 and Sb 2 O 4 ,
with sulfur to form Sb 2 S 3 ; and with chlorine to form SbCl 3
and SbCl 6 . It does not dissolve in nonoxidizing acids, but
dissolves readily in aqua regia. HNO 3 attacks it, but
forms an oxide that is soluble, with great difficulty, in
concentrated HNO 3 .

Compounds of Antimony. Though antimony exists in
both the tri- and pentavalent states, the trivalent com-
pounds are the most stable. Both oxides are amphoteric,
but Sb 2 O 3 is more basic and Sb 2 O 5 is more acidic in nature.
Antimonic compounds are oxidizing agents.

All antimony salts hydrolyze in dilute solution, SbCl 3
giving SbOCl and Sb(OH) 3 , both of them white, insoluble
in water, but soluble in acids and alkalies. Tartrates
repress the precipitation, forming soluble (SbO)C 4 H 4 O 6 ~.

NaOH, NH 4 OH and alkali carbonates precipitate anti-
mony hydroxide, Sb(OH) 3 , white, soluble in acids or strong
alkalies. H 2 S precipitates Sb 2 S 3 , orange, insoluble in dilute
nonoxidizing acids, but soluble in strong HC1; in (NH 4 ) 2 S,
forming SbSs^ ; and in alkalies forming mixtures of thio- and
oxythio-antimonite ions.

Zinc and acid reduce antimony compounds to stibine,
SbH 3 , a gas giving reactions very similar to those of
arsine, AsH 3 . Stibine is not formed in alkaline solution,
however, this being used as a method of identifying arsenic
in the presence of antimony. The Marsh test gives a mirror
of metallic Sb, which is distinguished from the As mirror by
the fact that Sb is not soluble in NaOCl solution.

Tin. Metallic tin is a soft, malleable metal, having a
melting point of 232C., a boiling point of 2270C., and a

GROUP II METALS 87

density of 7.3. Being just above hydrogen in the displace-
ment series, it slowly dissolves in dilute nonoxidizing acids,
and, more readily, in strong, hot, concentrated HC1.
The metal reacts with HNO 3 to form meta-stannic acid,
H 2 SnO 3 , a white substance, insoluble in alkalies or acids.
In neutral or slightly acid solutions, zinc displaces tin
from its compounds, forming the metal.

Compounds of Tin. Tin forms compounds in which it
has a valence of two and of four. The latter are the most
stable as air readily oxidizes the former to the tetravalent
state. Both forms are amphoteric, and the compounds of
both hydrolyze readily in dilute aqueous solutions. Most
of the analytical reactions make use of the bivalent form.

Stannous Compounds. With NH 4 OH, carbonates, and
alkalies, white Sn(OH) 2 is precipitated ; only slightly soluble
in excess NH 4 OH but readily soluble in excess alkali
[forming stannite (HSnO 2 ~~) ions], and in mineral acids.
Stannites decompose slowly, on standing, (or more rapidly
if heated) giving brown SnO and metallic tin.

SnCl 2 (the most common stannous compound) hydrolyzes
in water giving Sn(OH)Cl, or in more dilute solutions giving
Sn(OH) 2 . Solutions of SnCl 2 are kept stable by adding
HC1 to prevent hydrolysis, and metallic tin, to reduce any
stannic ions formed by air oxidation.

H 2 S precipitates SnS, brown, insoluble in water or
(NH 4 ) 2 S, but soluble in moderately strong HC1, in mixtures
of oxalic acid and ammonium salts, and in alkali or ammo-
nium polysulfide forming thiostannate (SnSs^) ions. The
latter decompose on acidification giving SnS 2 .

All stannous compounds are good reducing agents.
They reduce HgCl 2 to Hg 2 Cl 2 , then to free Hg. Bismuth
compounds are reduced by alkaline stannous solutions to
black metallic Bi. Alkaline solutions of lead are reduced, in
a similar manner, to metallic Pb, the reaction being much
slower than with bismuth compounds. Stannous com-
pounds will also reduce ferric ions to the ferrous state.

88 ANALYTICAL PROCEDURE CATIONS

Stannic Compounds. The simple stannic ion seldom
exists in solution as it has strong tendencies to form com-
plex negative ions. In neutral aqueous solutions, stannic
compounds hydrolyze, forming Sn(OH) 4 . To prevent this
the solution must be strongly acid, the tin being in the form
of ions of the type, SnCl 6 = . Alkalies, NH 4 OH or soluble
carbonates precipitate Sn(OH) 4 from solutions of stannic
compounds. The Sn(OH) 4 thus formed is white, insoluble
in NH 4 OH or Na 2 CO 3 , but soluble in alkalies or K 2 CO 3
(forming Sn(OH) 6 == ) and in mineral acids. If the stannic
hydroxide is allowed to stand, or is heated for a while, it is
converted to meta-stannic acid (H 2 SnO 3 ) which is insoluble
in either acids or alkalies. *

H 2 S reacts with solutions of stannic compounds, precipi-
tating SnS 2 , yellow, insoluble in water or ammonium car-
bonate; soluble in (NH 4 ) 2 S or alkali polysulfides (forming
SnSs^) and in acids. Considerable amounts of oxalic acid
prevent this precipitation.

Antimony, lead, or iron reduces acid solutions of stannic
compounds to the stannous state but not to the metal.
Zinc reduces stannic solutions to metallic tin; but if the
solution is strongly acid, the tin redissolves, forming
stannous ion.

Stannic compounds are not reducing agents.

ANALYTICAL ASPECTS

Group II consists of those metals whose sulfides can be
precipitated by H 2 S in a solution whose acidity is equal to
that of 0.3AT HCL The metals of this group are divided

* For one method of dissolving meta-stannic acid (sometimes called
j8-stannic acid), see the section on Amphoterism (also see Note 4).

Besides the method given in these sections a very effective method of
putting meta-stannic acid into solution is to fuse it with Na 2 CO 3 and sulfur,
treating the pulverized residue with water. The reaction is

2H 2 SnO 3 + 2Na 2 CO 3 + 9S -* 2Na 2 SnS 3 + 3SO 2 f + 2CO 2 f + 2H 2 O
On weakly acidifying this solution, SnSa will precipitate.

GROUP II METALS 89

into two subgroups on the basis of the solubility of their
sulfides in (NH^S reagent. Subgroup A (sometimes
called the " copper group") consists of those metals whose
sulfides are not soluble in (NH 4 ) 2 S, i.e., copper, lead,
mercuric mercury, cadmium and bismuth ; subgroup B (the
"tin group") consists of those metals whose sulfides are
soluble in (NH^S, i.e., tin, arsenic, and antimony.

Lead is considered in this group as well as in Group I
because its chloride is not extremely insoluble (it is about
a thousand times as soluble as silver chloride) and some
lead remains in solution and is precipitated as PbS in
Group II.

PRELIMINARY EXPERIMENTS GROUP II*

1. Mercuric Ion, Hg ++ . a. Pass H 2 S through 2 drops
Hg" 1 " 1 " test solution for about 30 sec., noting the color
changes. Filter and wash the precipitate. Divide the
precipitate into two parts.

To one portion of the precipitate add 5 drops 3N HNO 3
and heat on the steam bath for about 3 min. (Question 4).

To the other portion of the precipitate add 5 drops freshly
prepared aqua regia (Question 5) and evaporate to dryness.
Cool, add 2 drops water, stir, then filter. To 1 drop filtrate
on a watch glass add 1 drop SnCl 2 solution (see Note 13).

b. To 3 drops Hg^ test solution add 1 drop 0.1 AT KI
solution. Now add an excess of saturated KI solution
(Question 6).

c. To 2 drops Hg ++ test solution add 1 drop 3N K 2 Cr0 4
solution.

d. To 2 drops Hg ++ test solution add 1 drop IN NH 4 C1
solution and 1 drop concentrated NH 4 OH (Question 7).

e. To 1 drop Hg"^" test solution add 1 drop freshly
prepared solution of gallic acid in alcohol.

* For additional chemical properties that are occasionally used for tests
for these metals, see Auxiliary Tests.

90 ANALYTICAL PROCEDURE CATIONS

2. Copper Ion, Cu*" 4 " . a. Pass H 2 S through 2 drops Cu 4 " 1 "
test solution for about 30 sec. Filter and wash the precipi-
tate. Divide into two parts.

To one portion of the precipitate add 5 drops 3 AT HNO 3
and heat. Filter off any sulfur (recognized by the fact
that it usually floats whereas CuS sinks). To the clear
solution add IN NH 4 OH, drop by drop, noting that a
precipitate first forms, then dissolves in the excess reagent
(see page 38). The reaction is*

CU++ + 2(OH-) -*Cu(OH) 2
Cu(OH) 2 + 4NH 3

To the second portion of precipitate add 5 drops freshly
prepared (NH 4 ) 2 S solution. Warm on the steam bath
(Question 8).

6. To 2 drops Cu++ test solution add 2 drops 6N NH 4 OH.
Then add, drop by drop until the deep blue color just dis-
appears, a 3 per cent solution of KCN (Caution: KCN is an
extremely dangerous poison. If added to acid solutions,
it gives off very poisonous fumes of HCN. For this
reason it is important that every detail in the procedure be
checked before it is used.) Now bubble H 2 S through the
solution for a minute (see Note 14).

c. To 1 drop Cu 4 " 4 " test solution on a watch glass add 1
drop K 4 Fe(CN)e solution.

d. Put 1 drop GU++ test solution on a bit of filter paper.
To the same spot add 1 drop a-benzoinoxime solution.
Hold the spot over a watch glass containing a few drops
concentrated NH 4 OH and observe any color change (52,
53).

3. Bismuth Ion, Bi +++ . Bismuth forms a sulfide under
the same conditions as does copper. Like CuS, Bi 2 S 3 is
black in color and is soluble in 3N HNO 3 . To save time,

* An explanation of the reasons for NH4OH acting as a source of both
OH " ions and NH 3 molecules, will be found on p. 38.

GROUP II METALS 91

only a few of the other characteristic reactions of its ions
will be tried experimentally.

a. To 1 drop Si 4 " 1 " 1 " test solution add 10 drops water and
stir. Observe very closely as the resulting precipitate is
almost transparent. If no precipitate appears, add 1 drop
QN NH 4 OH (see Note 7 and Question 9).

b. To 1 drop Si 4 " 4 " 4 " test solution on a filter paper, add 1
drop IN NaOH solution and 1 drop freshly prepared sodium
stannite solution (see Note 15).

c. To 1 drop Bi + " H " test solution add 1 drop Cu 4 " 4 " test
solution and 10 drops water. Add QN NH 4 OH, drop by
drop, until the whole solution is deep blue. Can you see
any precipitate in the solution? Try to dissolve it by
adding more NH 4 OH (see Note 16).

d. To 1 drop Bi +++ test solution on a piece of filter paper
add 1 drop cinchonine reagent. This test is capable -of
detecting as little as 0.00014 mg. of bismdth. Try the
reagent on a drop of a mixture of Cu 4 " 1 ", Pb^ 4 " , Si 4 " 4 " 4 ", and
Hg 4 " 1 ", noting the concentric rings (6, 14).

4. Cadmium Ion. a. Bubble H 2 S through 2 drops Cd" 1 " 4 "
test solution until precipitation is complete. Filter and
wash the precipitate. Divide the precipitate into two
parts.

To one portion of the precipitate add 5 drops 3N HNO 3
and warm on the steam bath.

To the second portion of precipitate add 5 drops of
1:4 H 2 SO4 (a mixture of 1 drop concentrated H 2 SO 4 and 4
drops water). Warm on the steam bath.

b. To 1 drop Cd^ 4 " test solution on a filter paper add
1 drop IN NaOH solution. In the center of the spot, place
a tiny crystal of thiosinamine (allyl thiourea), and heat in a
jet of steam from the steam bath. A bright yellow spot
should quickly result (16) (see Note 17).

c. To 1 drop Cd +4 ~ test solution add 1 drop Cu 4 " 4 " test
solution and 4 drops water. Now add 2 drops concen-
trated NH 4 OH; then add 3 per cent KCN solution, drop by

92 ANALYTICAL PROCEDURE CATIONS

drop, until the blue color of the copper is discharged.
Bubble H 2 S through this solution for a few seconds.

6. Arsenic Ion, As 4 " f ~ H ~ i ~. a. Pass H 2 S through 2 drops
Aaf*~ H ~ H " test solution until a precipitate forms. Note the
length of time it takes.

To 2 drops As 4 " 1 1 " 1 " 1 " test solution add a small crystal of
NH 4 I, heat to boiling and pass H 2 S through the solution.
Note the length of time required before a precipitate forms
this time (see Note 18). Continue passing in H 2 S until
precipitation is complete. Filter and wash the precipitate.
Divide the precipitate into three portions.

To one portion of precipitate add 5 drops (NH 4 ) 2 S solu-
tion and warm. When solution is complete add, drop by
drop, IN HC1 until the solution is faintly acid (see Note 19).

To another portion of the precipitate add 5 drops con-
centrated HC1 and heat on the steam bath. Does the
precipitate dissolve?

To the third portion of precipitate add 5 drops aqua regia
and evaporate just to dry ness (see Note 20). Dissolve the
residue in 2 drops QN HC1. Add 1 drop 0.1 N sodium
bisulfite and warm 1 min. Then add 5 drops 6N NaOH
solution and stir. Transfer to the gas evolution apparatus.
In the bulb of the latter place a bit of filter paper moistened
with a drop of AgNOa solution and cover with a loose plug
of cotton. Then, to the solution in the main tube, add
four pieces arsenic-free granulated aluminum and quickly
place the top of the apparatus in position. Warm if
necessary and set aside for about 5 min. Then examine
the silver nitrate spot (see Note 21).

6. Antimony Ion, Sb 44 " 1 ". a. Pass H 2 S into 2 drops
gb+++ test solution until precipitation is complete. Filter
and wash precipitate. Divide the precipitate into two
parts.

To one portion of the precipitate add 5 drops (NH 4 ) 2 S
solution and warm on the steam bath. To the solution so

GROUP II METALS

obtained add IN HCl, drop by drop, until the solution is
faintly acid.

To the second portion add 5 drops concentrated HC1 and
heat until the precipitate dissolves. Now add 5 drops
cold water and pass in H 2 S for 1 min.

b. To 1 drop Sb 4 "^ test solution add 5 drops water.
Now try to redissolve the precipitate formed, by adding
concentrated HC1 (see Note 22).

c. To 1 drop Sb+++ test solution add IN NaHCO 3 solu-
tion, drop by drop, until a precipitate barely begins to form.
Add just enough &N HC1 to redissolve the precipitate, heat
almost to boiling, and add a crystal of sodium thiosulfate.
Note the orange precipitate of antimony oxysulfide,
SbOS 2 (54, 55).

GROUP II ANALYSIS

The unknown may be the filtrate from the precipitation of the chlorides
of Group I, or a solution provided by the instructor.

Evaporate barely to dry ness then cool (see Note 25). Add 1 drop of
cone. HCl, let stand 1 min. then add 1 nil. water (see Note 26) and pass
in H 2 S until precipitation is complete (see Note 34). Add a small crystal
of NH 4 I, heat to boiling, and again pass in H 2 S. Cool once more and com-
plete the precipitation with H 2 S (sec Question 10). Filter and wash pre-
cipitate twice with NH 4 NO r H 2 S solution (see Notes 27 and 28).

Filtrate. Make slrongly
acid with HCl and boil
to half its volume. Save
for Groups III, IV, and
V.

Precipitate. Contains CuS, 1M 2 S.<, (MS, HgS,
PbS, SnS 2 , Sb 2 S 3 , and As 2 S 3 .

Add 10 drops (NH 4 ) 2 S soln. (see Notes 19
and 29) and heat 011 steam bath for 1 ruin.,
with stirring. Centrifuge and save nitrate.
Repeat treatment of precipitate with hot
(NH4) 2 S, centrifuging and combining the ni-
trate with that from the first treatment.
Wash precipitate with NH 4 NO 3 -H 2 S soln.
and discard the wash liquid.

Filtrate. Contains members of Group
115 as (NH 4 ) 8 AsS3, (NH 4 ) a -
SbS 3 and (NH 4 ) 4 SnS 4 . Treat as
described under the heading
Group II B.

Precipitate. Contains sulfides of Group
II A metals. Analyze as described
under the heading Group II A.

ANALYTICAL PROCEDURE CATIONS

GROUP II A ANALYSIS

To the precipitate from the subgroup separation, add 10 drops 3N HNOa
and heat for 3 to 4 min. on steam bath. Filter and wash precipitate with
cold water.

Filtrate. Contains any Pb, Cu, Bi and Cd as
nitrates. Add 3 drops 6N H 2 SO4 and evaporate to
, ! i original volume (no smaller). Cool, then add

1 1 nil. of cold water, and stir. Let stand. Filter
and wash ppt. with 50 % alcohol. (See Note 30.)

Precipitate. Con-
tains S and/or HgS.

Divide ppt. into two
parts. (See Note
31.) Test as follows:

Filtrate. Cu++, Bi" + and

Precipitate. PbS0 4

To one portion add

Cd" 1 " 4 . Add an excess of

Heat almost to boil-

aqua regia and

cone. NH 4 OH. (See Note

ing with 10 drops

evaporate to dry-

32. ) Filter even if no ppt. is

saturated NH 4 Ac

ness. Add 2 drops

visible. Wash any ppt. with

soln. (See Ques-

water, stir, and

dilute NH 4 OH.

tion 11 and page

filter. To one drop

36.) Filter. To

solution on a watch

Filtrate. Cu Precipitate.

filtrate on black

glass add 1 drop

and Cd. Bi(OH) 3 . Dis-

background add 2

stannous chloride.

(See next solve in 5

drops K 2 CrO 4 solu-

page.) drops 3 N

tion.

WHITE, GRAY, or

HP1 Tn 1

BLACK PPT.

drop of this solution on

YELLOW PPT.

indicates Hg ++ .

watch glass add 1 drop of

confirms Pb ++ pres-

cinchonine reagent.

ent.

To a second portion

of the ppt. add 10

ORANGE PPT.

drops bromine

indicates Hi*"*"* present.

water, boil one rnin.,

filter TY +Vm fi]

Confirm again by adding to

niircr. A u tiie ni-
trate add 2 drops of

another portion of the soln.

6ATH 2 SO 4 and evap-

on filter paper 2 drops of

orate until SO 3

fresh sodium stannite soln.

fumes appear. Cool,

(See Note 15.)

add 3 drops of water

and 2 drops of sat.

BLACK COLORATION

NaAc solution.

confirms Bi"*" 1 "* present.

Then add 2 drops of

diphenylcarbazide

solution.

DEEP BLUE

indicates Hg++ pres-

ent.

GROUP II METALS

Filtrate from the separation of Cu ++ and Cd++ from Bi +f ^ This filtrate
contains Cu+ + and Cd ++ . Divide it into two equal parts. Use one portion
for the Cu ++ tests and the other for the Cd++ tests as follows :

For Cu + +:

If the solution has ti BLUE COLOR Cu + + is present.

Confirm by testing 1 drop of the solution on a watch glass with 1 drop
QN acetic acid and 1 drop K4Fe(CN) 6 solution.

RED PPT.

indicates Cu present.

To another drop of the solution on filter paper add I drop a-benzoinox-
ime.

GREEN SPOT
confirms Cu + + present.

For

To test for cadmium, treat the portion of filtrate saved for that purpose
in one of the following wjiys:

// copper is present:

Add KCN solution, (Caution: be
sure the solution is strongly ammo-
iiiacal before adding KCN), drop
by drop until the blue of the copper
has disappeared. Add 3 more
drops of KCN; then pass in H 2 S
for 2 to 3 miii.

YELLOW PPT.
indicates Cd ++ present.
(See Note 33.)

Confirm by test described below.

// copper is absent:

See Note 33 and use procedure de-
scribed there. Or proceed as fol-
lows:

Pass HjjS into the solution for 2 to
3 in in.

YELLOW PPT.
indicates Cd+ + present.

Confirm by test described below.

To confirm the presence of Cd f f , dissolve the yellow precipitate obtained
in the procedure used above, in warm 1 : 4 H 2 SO 4 . Filter. To 1 drop of
the filtrate on filter paper add 2 drops QN NaOH and a small crystal of
thiosinamine and warm in a steam jet. YELLOW COLORATION or PPT.
confirms Cd ++ present.

ANALYTICAL PROCEDURE CATIONS

GROUP II B ANALYSIS

To the filtrate from the sub-group separation add dil. HC1 until faintly
acid. (See Note 35.) Filter. (See Note 30.) Wash precipitate with
NH 4 N0 3 -H 2 S soln.

Filtrate.
Discard.

Precipitate. Sulfides of As, Sb and Sn. Add 10 drops com*.
HCl, warm on steam bath and stir for 1 min. Cool and add
2 drops water. Filter.

Filtrate. Containing SbCl 3 and
SnCU. Divide into four equal
parts and test as follows:

Place the first portion in a micro-
beaker. Add an equal volume of
water (exactly) and pass in H 2 S for
1 min.

ORANGE PPT.

indicates Sb +44 present. (See
Note 38.)

Take a second portion of the filtrate
and add IN NaHCO 3 soln. to it
until it is just weakly acid or until
a precipitate starts to form. Re-
dissolve any precipitate by adding
1 drop GN HCl. Heat the soln. to
boiling, remove from flame and
quickly add a crystal of sodium
thiosulfate about the size of a pea.
Let stand about 1 min.

ORANGE PPT. or COLORATION

of SbOS,
Confirms Sb f++ present.

To the third portion of filtrate in a
microbeaker add one small alumi-
num filing and warm slightly (see
Note 24). If any black precipitate
is left after gas evolution ceases,
add 2 drops cone. HCl and warm
again until all traces of gas evolution
cease. Filter quickly (using cot-
ton plug if necessary). Cool, and
add 1 drop cacotheline soln.

VIOLET SOLN.

indicates Sn present. (See Note
39.)

Precipitate. As ? Si (and As 2 S 6 ).
Place in a crucible, add 10 drops of
aqua regia and evaporate barely to
dryncss. Dissolve the residue in
5 drops water and 1 drop of 6JV
acetic acid and stir. Use as follows :

To 1 drop of the soln. on a spot plate
add 2 drops saturated ammonium
acetate and I drop AgNO 3 soln.

CHOCOLATE PPT.
indicates arsenic present. (Sec
Note 40.)

To the remainder of the solution add
1 drop NallSO.i soln. and warm for
I mm. on steam bath. Then add
5 drops (>N NaOH and transfer to a
gas evolution tube. Place a drop of
AgNO< solution on a bit of filter
paper and put the latter in the bulb
of the gas evolution apparatus.
Drop four granules of arsenic-free
aluminum filing into the NaOH solu-
tion and quickly put top of appara-
tus in place. Set apparatus aside
for about 10 min. (See Note 37.)
YELLOW, BROWN or BLACK

STAIN
confirms As present.

Treat the fourth portion of filtrate with an aluminum filing as in previous
paragraph. Filter, and add to filtrate 1 drop HgCl 2 soln.

WHITE, GRAY or BLACJK PPT.
indicates Sn present.
(very specific if obtained, but less sensitive) .

GROUP III METALS 97

7. Stannous and Stannic Ion, Sn" 14 " and Sn 4 " 4 " 4 " 4 ". a.
Pass H 2 S into 2 drops Sn 4 " 4 " test solution until precipitation
is complete. Filter and wash precipitate. Divide into
two portions.

To one portion of the precipitate add 5 drops (NH 4 ) 2 S
solution and warm. Acidify the resulting solution with
IN HC1.

To the second portion add 5 drops concentrated HC1 and
warm until all precipitate dissolves. To 2 drops of this
solution on a watch glass add 1 drop HgCl 2 solution. To
the remainder of the solution add 1 drop cacotheline solu-
tion and note the color change (see Note 23) (6, 56).

6. To 2 drops Sn 44 ' test solution add 1 drop 0.1 AT NaOH
solution. Note the precipitate of Sn(OH) 2 that forms.
Now add 1JV NaOH, drop by drop, until the precipitate
just redissolves. The resulting solution is a solution of
sodium stannite.

c. Pass H 2 S into 2 drops Sn 4 " 4 " 4 " 4 " test solution. Filter
and wash the precipitate. Divide into two parts.

Test the solubility of one part in hot (NH 4 ) 2 S. Reprecip-
itate by making the solution faintly acid.

Dissolve the remainder in 10 drops concentrated HC1.
To this solution add the smallest possible granule of zinc.
Heat until all the zinc is dissolved. To the solution add a
drop of HgCl 2 solution (see Note 24).

GROUP III METALS

CHEMICAL CHARACTERISTICS*

Iron. Iron is a gray metal, having a melting point of
1535C., a boiling point of 3000C., and a density of 7.86.
It is a moderately active metal, reacting slowly at ordinary
temperatures with oxygen, chlorine, and many other non-
metals. It reacts rapidly, at high temperatures, with
oxygen to form Fe 3 O 4 ; with sulfur, to form FeS; and with

* For other typical reactions that have application to analysis, see
Auxiliary Tests.

98 ANALYTICAL PROCEDURE CATIONS

steam, to form Fe 3 O 4 and hydrogen. It dissolves readily
in HC1, H 2 SO 4 , or dilute HNO 3 . Cold, concentrated (or
fuming) HNOs, however, changes it to "passive" iron a
form which will riot show any of the usual reactions of iron. *

Compounds of Iron. Iron forms ferrous, Fe"^, and
ferric, Fe+ ++ , ions. The former readily oxidize to the
ferric state in the presence of oxygen, bromine, chlorine,
and other oxidizing agents. As many of the ferric com-
pounds important in analysis are less soluble than the
corresponding ferrous compounds, any ferrous ion is usually
oxidized to the ferric state before the analysis is attempted.

Ferrous Compounds. From solutions containing ferrous
ions, NH 4 OH and alkalies precipitate Fe(OH) 2 , white, f
insoluble in water or excess alkali; but soluble in high
concentrations of ammonium salts and in acids. The
Fe(OH) 2 quickly oxidizes to Fe(OH) 3 in the presence of
air or other oxidizing agents (such as Na 2 O 2 ).

H 2 S or (NH 4 ) 2 S precipitates FeS from neutral solutions,
or more quickly, from alkaline solutions of ferrous ions.
FeS is black, insoluble in water or alkalies but soluble in
quite dilute acids.

KCN precipitates yellowish Fe(CN) 2 , soluble in excess
KCN [to form Fe(CN) 6 ] and in acids.

K 3 Fe(CN) 6 precipitates blue Fe 3 (Fe(CN) 6 ) 2 (TurnbulPs
blue).

K 4 Fe(CN) 6 precipitates white Fe 2 Fe(CN) 6 [usually tinged
blue, owing to air oxidation forming small amounts of
Prussian blue, Fe 4 (Fe(CN) 6 ) 3 ].

* It is believed that the strong HNOa forms a thin, unstable film of oxide
on the surface of the iron, and that this protects the iron from any reagents
that may be applied. This film is destroyed by heating, scratching, or
striking the iron sharply.

t As usually precipitated, Fe(OH) 2 has a greenish color. This is due to
the fact that it is very difficult to prepare or keep ferrous salts free of ferric
compounds. The trace of Fe(OH) 3 that precipitates colors the Fe(OH) 2 .
It will be found that other ferrous compounds, as ordinarily precipitated,
exhibit similar anomalies.

GROUP III METALS 99

Dimethylglyoxime gives a red coloration which slowly
fades, as the Fe 4 " 1 " ions are oxidized by air to the ferric state.

Soluble carbonates precipitate FeCO 3 , white at first but
turning green on standing; soluble in acids.

Ferric Compounds. NH 4 OH and alkalies precipitate
Fe(OH) 3 , red-brown, less soluble in water than Fe(OH) 2
and insoluble in alkalies or in solutions of ammonium salts ;
soluble in acids.

H 2 S or (NH 4 ) 2 S precipitates a mixture of sulfur and FeS,
the sulfide first reducing the Fe +4 " f to Fe~ H ~.

K 3 Fe(CN) 6 gives a brown coloration with solutions con-
taining Fe H " H ~.

K 4 Fe(CN) 6 gives a dark blue precipitate of Fe 4 (Fe(CN) 6 ) 3
(Prussian blue).

KSCN gives a deep red color with solutions containing
Fe~M~ + , owing to the formation of the slightly ionized
Fe(SCN) 3 .*

Soluble carbonates give a precipitate of basic ferric
carbonate decomposed, by boiling, into Fe(OH) 3 .

Soluble phosphates precipitate brown FePO 4 , soluble in
excess reagent (forming complex ions) or in mineral acids.

Acetates precipitate red-brown Fe(OH) 2 (C 2 H 3 O 2 ) from
hot solutions of Fe ++ " f ions.

Aluminum. Aluminum is a silvery rather soft metal,
having a very low density (2.7), a melting point of 648C.,
and a boiling point of 1800C. It is very active, but due to
the formation of a protective film of oxide that quickly
forms on its surface, it is not readily attacked by oxygen or
moisture. It dissolves readily in HC1, dilute H 2 SO 4 , or in
solutions of alkalies (forming A1O 2 ~) but is unaffected by
HN0 3 .f

* The color does not form if fluorides, phosphates, oxalates or tartrates
are present, as these form very stable complex ions with the iron. If the
cold solution is made acid with nitric acid, however, the test may be used
successfully.

t This is generally attributed to the formation of a protective film of
oxide, as in the case of iron. However, HC1 readily destroys this film.

100 ANALYTICAL PROCEDURE CATIONS

Compounds of Aluminum. All the salts of aluminum
hydrolyze almost completely in neutral solution, precipitat-
ing A1(OH) 3 .

NH 4 OH, (NH 4 ) 2 S (or H 2 S in NH 4 OH solution), alkalies,
and carbonates precipitate A1(OH) 3 , a white gelatinous
precipitate, insoluble in acetic acid but soluble in alkalies
(forming A1O 2 ~) * or in mineral acids.

Phosphates precipitate A1PO 4 , white, soluble in alkalies
or in mineral acids.

Acetates produce no precipitate with cold solutions of
AI 4 "" 1 "*, but on boiling, a white basic acetate, A1(OH) 2 -
(C 2 H 3 O 2 ), is precipitated. This is readily soluble in acids
and in alkalies.

Certain dyes, such as aluminon (the ammonium salt of
aurin tricarboxylic acid) form characteristic colored
"lakes" with Al(OH),.

For other characteristic reactions of aluminum, see
Auxiliary Tests.

Chromium. Chromium is a silvery, rather brittle metal,
having a melting point of 1615C. and a density of 6.92.
It dissolves in HC1, dilute H 2 SO 4 , or alkalies but, like
aluminum, is not dissolved by nitric acid. It is very
similar to aluminum in many respects, but differs in that
chromium forms complex ions and forms compounds in
which it exhibits many different valences.

Compounds of Chromium. Although chromium forms
compounds corresponding to the oxides, CrO, Cr 2 O3, CrO 2 ,
and CrO 3 (as well as several less known, higher oxides) the
most common are those formed from Cr 2 O 3 and CrO 3 . Of
these, CrO and Cr 2 O 3 are basic anhydrides, though Cr 2 O 3 is
also somewhat amphoteric. CrO 3 is an acidic oxide,
dissolving in water to form chromic acid, H 2 CrO 4 . Its
normal salts are the chromates; its acid salts are the
dichromates.

* For a discussion of the reactions of A1(OH) 3 , Cr(OH) 3 , Zn(OH) 2 and
other amphoteric hydroxides see the section on Amphoteric Hydroxides.

GROUP III METALS 101

The chromous salts (salts of CrO) are strong reducing
agents, reacting readily with oxygen and other oxidizing
agents to give chromic, Cr~ H " t ~, ions.

With NH 4 OH, (NH 4 ) 2 S (or H 2 S and NH 4 OH), alkalies,
and carbonates, chromic salts give a precipitate of Cr(OH) 3 ,
green, slightly soluble in excess NH 4 OH (especially if
ammonium salts are present) and readily soluble in excess
alkali (forming chromite, CrO 2 ~, ions) and in acids.

Phosphates give a greenish precipitate of CrPO 4 , soluble
in acetic acid and mineral acids.

Acetates produce no precipitate except when considerable
quantities of aluminum or iron are present. In this case,
chromium coprecipitates with these two metals as the basic
acetate.

Alkaline solutions of chromitos are oxidized to chromates,
CrO 4 = , by addition of Na 2 O 2 , NaOCl, Br 2 , or C1 2 . Fusion
of chromic compounds with Na 2 CO 3 and oxidizing agents
(such as Na 2 O 2 , KC1O 3 , or KNO 3 ) converts them to
Na 2 CrO 4 .

Acid solutions of chromic ion (Cr 4 " H ~) may be oxidized to
dichromate ion (Cr 2 O 7 = ) by use of strong oxidizing agents
such as hot HNO 3 or KC1O 3 .

The chromates are strong oxidizing agents, especially in
acid solutions. Acids convert them to dichromates.*

Cold concentrated H 2 SO 4 gives a red precipitate of
CrO 3 with chromates. On heating, the mixture gives ofif
oxygen and the CrO 3 is reduced to a green solution of

If an acid solution of a chromate is treated with H 2 O 2 , a
green or blue color appears due to the formation of blue
perchromic acid, H 3 CrO 8 (see Note 46). This rapidly dis-

* The dichromates are acid chromates. The relationship may be shown
by the equations for the reaction between H + and CrO 4".

CrOr + H+ ^ HCrO 4 -
2HCrOr ^ Cr 2 O 7 - + H 2 O

102 ANALYTICAL PROCEDURE CATIONS

appears as the compound breaks down giving oxygen arid
0+++. It is more stable in ether solutions, however.

If carried out in acetic acid buffered with ammonium
acetate, a brown solution is formed, remaining for several
minutes.

Many metals form characteristic insoluble chromates.
For these, see the discussions on silver, mercury, lead,
barium and strontium.

Manganese. Manganese is a gray or reddish-white, very
hard, brittle metal having a density of 7.12 and a melting
point of 1260C. It is very similar to iron in activity,
combining readily with oxygen, sulfur, chlorine, and many
other nonmetals, at high temperatures. Manganese dis-
solves readily in dilute acids.

Compounds of Manganese. The most important classes
of compounds formed by manganese are the pink, mariga-
nous, Mn 4 ^", compounds; the green manganates, MnO 4 === ;
and the violet or purple permanganates, MnO 4 ~.

Manganous Compounds. Alkalies and NH 4 OH react
with solutions of manganous ions to form Mn(OH)^, white,
insoluble in excess alkali but soluble in solutions of ammo-
nium salts and in acids. Air or oxidizing agents quickly
convert Mn(OH) 2 to manganous acid, H 2 MnO 3 or
MnO 2 -H 2 0.

Soluble cyanides give Mn(CN) 2 , brown, soluble in excess
of the reagent to form Mn(CN) 6 ~.

Soluble carbonates precipitate MnCO 3 , white, insoluble
in solutions of ammonium salts but readily soluble in dilute
acids.

Soluble phosphates precipitate Mn 3 (PO 4 )2, white, soluble
in acetic and mineral acids.

K 4 Fe(CN) 6 precipitates Mn 2 Fe(CN) 6> white, insoluble
in water but slightly soluble in HC1.

K 3 Fe(CN) 6 precipitates Mn 3 (Fe(CN) 6 ) 2 , brown, insoluble
in water and most acids, but somewhat soluble in HC1.

GROUP III METALS 103

KClOs (solid), boiled with concentrated HNO 3 solutions
of Mn 4 " 4 ", gives H 2 MnO 3 , brown, insoluble in HNO 3 .

When manganous compounds are fused with Na 2 CO 3 and
an oxidizing agent such as air, KC1O 3 , KNO 3 , or Na 2 2 ,
green Na 2 Mn04 (sodium manganate) is formed. A solu-
tion of a manganate will, on acidification, change from
green to violet or pink, and a brown precipitate of E^MnOa
will settle out. The reaction is

MnOr + 4H+ -* MnOr + H 2 Mn0 3 + H 2 O

manganate permanganate

Mn ++ is oxidized in nitric acid solutions to permanganate,
MnO 4 ~~, ions by heating with ammonium persulfate,
(NH 4 ) 2 S 2 O 8 , and AgNO 3 ; with PbO 2 , or with KBrO 3 .
Chlorates or iodates do not act in this way. Sodium bis-
muthate gives the reaction in the cold solution, but gives
H 2 Mn0 3 , if heated.

The metallic permanganates are all quite soluble. Per-
manganate ions are oxidizing agents, being reduced in acid
solution to Mn ++ , and in alkaline solution to MnO 3 = .

Zinc. Metallic zinc is a bluish-gray, rather soft metal,
having a density of 7.14, a melting point of 419C., and a
boiling point of 907C. It is quite active, burning readily
in air to form white ZnO, and combining with many non-
metals (such as chlorine) even at ordinary temperatures.
It dissolves readily in dilute acids, or alkalies giving
hydrogen.

Compounds of Zinc. Alkalies and NH 4 OH react with
solutions of Zn 4 ^ ions precipitating Zn(OH) 2 , white and
gelatinous, soluble in excess alkali, giving zincate, HZnO 2 ~,
ions; in excess NH 4 OH or solutions of ammonium salts,
giving Zn(NH 3 ) 4 " f " f ; and in acids.

H 2 S or (NH 4 ) 2 S precipitates ZnS from neutral, basic,
ammoniacal or acetic acid-sodium acetate solutions, but
not from acid solutions (unless the concentration of zinc is

1 04 ANAL YTICAL P ROC ED URECA TIONS

unusually high). ZnS is white, insoluble in water or
NH 4 OH, but soluble in very dilute mineral acids.

Soluble carbonates precipitate mixed basic carbonates of
the type Zn 2 (OH) 2 CO 3 , white, soluble in NH 4 OH or strong
solutions of ammonium salts, in alkali carbonates, and in
acids. Phosphates precipitate Zn 3 (P0 4 )2, white, soluble in
NH 4 OH or in acids.

Cyanides precipitate Zn(CN) 2 , white, soluble in excess of
the reagent to form Zn(CN) 4 = . The latter are decomposed
by (NH 4 ) 2 S to form ZnS.

K 4 Fe(CN) 6 precipitates white Zn 2 Fe(CN) 6 .

Cobalt and Nickel.* These elements are both gray
metals, nickel being more silvery, in appearance, than is
cobalt. Both have a density of 8.9 and a boiling point of
2900C. The melting point of cobalt is 1480C. while that
of nickel is 1452C.

Both dissolve readily in HNO 3 ; but while cobalt dissolves
readily in dilute HC1 and H 2 SO 4 , nickel dissolves in them
with difficulty.

Both metals readily form a large number of complex ions,
and both are capable of existing in either the bivalent state
or the trivalent state. In both cases, the most common
ions are the bivalent form. However, cobalt is more easily
oxidized to form the trivalent ion than is nickel.

Compounds of Nickel. Alkalies precipitate, from solu-
tions of nickelous, Ni ++ , ions, Ni(OH) 2 , green, insoluble in
excess alkali but soluble in mixtures of ammonium salts and
NH 4 OH [forming Ni(NH 3 ) 4 + + ] and in acids.

Oxidizing agents such as Br 2 , C1 2> or KOC1, oxidize
alkaline suspensions of Ni(OH) 2 to black nickelic hydroxide,
Ni(OH) 3 , insoluble in NH 4 OH or KCN.

NH 4 OH precipitates a basic salt of the type Ni(OH)Cl,
soluble in the same reagents as is Ni(OH) 2 .

* These metals are so similar in nature that it is logical to consider them
together in order to bring out their similarities and differences. They are
also quite similar to iron.

GROUP III METALS 105

Carbonates precipitate green NiCOa, soluble in mixtures
of ammonium salts and NH 4 OH, in ammonium carbonate,
and in acids.

H 2 S or (NH 4 ) 2 S precipitate NiS from basic, neutral, or
weakly acid (acetic acid-acetate) solutions. NiS is insolu-
ble in (NH 4 ) 2 S, very slowly soluble in cold IN HC1, but
readily soluble in aqua regia or concentrated HNO 3 .

Phosphates precipitate green Ni 3 (PO 4 ) 2 , soluble in dilute
acids.

KNO 2 produces no precipitate in dilute solutions of Ni^ 4 "
ions. KCN precipitates green Ni(CN) 2 , soluble in excess
reagent to form the stable Ni(CN) 4 = . This is decomposed
by acids. Solutions of Ni(CN) 4 = will give NiS with H 2 S,
but with strong oxidizing agents they decompose giving
Ni(OH) 3 and cyanate, CNO~, ions.

Chromates precipitate a basic chromate from hot, nickel-
ous ion solutions, but react only slowly in the cold.

K 4 Fe(CN) 6 precipitates Ni 2 Fe(CN) 6 , green, almost insol-
uble in dilute HC1.

KSCN gives no visible reaction with nickel solutions.

Compounds of Cobalt. Alkalies or NH 4 OH precipitate a
blue basic salt of the type Co(OH)NO3, insoluble in excess
reagent; but soluble in acids, in an excess of KCN, or in
strong solutions of ammonium salts. On heating, the basic
salt hydrolyzes, forming pink Co(OH) 2 . The latter
changes to brown Co(OH) 3 , slowly, in air, but rapidly if
oxidizing agents such as Br 2 , C1 2 , H 2 O 2 , or NaOCl are
added.

H 2 S or (NH 4 ) 2 S precipitates black CoS from neutral or
basic solutions. CoS is insoluble in (NH 4 ) 2 S, and dissolves
but slowly in cold dilute (IN) HC1. It dissolves readily,
however, in hot, concentrated HNO 3 or in aqua regia.

Phosphates precipitate blue-violet Co 3 (PO 4 ) 2 , soluble in
NH 4 OH.

Carbonates precipitate basic cobalt carbonates, red,
soluble in NH 4 OH or in ammonium carbonate solution.

106 ANALYTICAL PROCEDURE CATIONS

KCN precipitates Co(CN) 2 , brown, soluble in excess
reagent to form very stable cobaltocyanide, Co(CN) 6 B , ions.
The latter oxidize readily to cobalticyanide, Co(CN) 6 ", ions.

Chromates precipitate a dark brown basic chromate*
from cold solutions of Co^ 4 ". The precipitate is readily
soluble in NH 4 OH.

K 4 Fe(CN) 6 precipitates Co 2 Fe(CN), green, almost insol-
uble in HC1 but soluble in NH 4 OH.

ANALYTICAL ASPECTS

Group III consists of those metals whose sulfides or
hydroxides are precipitated in mildly alkaline solutions of
H 2 S. To prevent the precipitation of the hydroxides of the
fourth group metals and magnesium, it is necessary that the
hydroxyl ion concentration, during precipitation of Group
III, be kept low, yet constant. For this reason an excess of
NH 4 OH buffered with ammonium chloride is used to furnish
the necessary alkalinity.

On making the solution alkaline, the hydroxides of all
the metals of Group III precipitate, with the exception of
those of zinc and nickel which remain in solution as ammo-
nia complexes. When H 2 S is passed into this solution, the
zinc and nickel complexes and the hydroxides of iron, man-
ganese and cobalt are converted into sulfides. The sulfides
of aluminum and chromium are peculiar in that, even if
they were formed momentarily, they would at once hydro-
lyze to form the hydroxides again. This is easily under-
stood if one recalls the fact that the sulfides of these metals
belong to that class of salts where hydrolysis is most com-
plete, i.e., salts of very weak acids and very weak bases.

The metals of this group are divided into two groups.
This division is based on the fact that treating a solution
containing these ions with excess alkali will cause the

* One of the separations of cobalt and nickel is based on the fact that
cobalt precipitates the basic chromate from cold solutions whereas nickel
precipitates very slowly (64).

GROUP III METALS 107

amphoteric hydroxides of aluminum, chromium and zinc to
dissolve while the hydroxides of iron, nickel and cobalt, and
the hydrated oxide of manganese, being insoluble in excess
alkali, may be filtered off (see page 42) .

PRELIMINARY EXPERIMENTS GROUP III

1. Ferric Ion, Fe +++ . On account of the fact that the
hydroxides and sulfides of ferric iron are less soluble than
those of ferrous iron, it is necessary that all ferrous ions be
converted to the ferric state before the analysis of this
group begins. This is easily done by adding a few drops of
bromine water to the unknown and boiling for a minute.
Since this is done first, the reactions that must be con-
sidered, are restricted to those of the ferric ion.

a. Pass H 2 S into Fe" 1 " 1 " 4 " test solution for a few seconds
(Question 12).

b. To 2 drops Fe~ HH ~ test solution add 1 drop of a satu-
rated solution of NH 4 C1 and 2 drops 6N NH 4 OH. Saturate
with H 2 S. Filter and wash the precipitate (see Note 41).

Test the solubility of the FeS precipitate in &N HC1.

c. To 2 drops Fe +++ test solution add IN NaOH solution
until a copious precipitate forms. Now try to redissolve
the precipitate in QN NaOH.

d. To 2 drops Fe^ 44 " test solution add, drop by drop,
IN NH 4 OH until a faint precipitate just appears. Then
add 1 drop IN acetic acid and 5 drops 5 per cent ammonium
benzoate solution. Dilute to 3 ml. with water; stir, and
heat on the steam bath for 5 min. Compare the appear-
ance of the Fe(OH) 3 precipitate with that obtained in part
(c) (Note 42). Filter and wash the precipitate with cold
water. Dissolve the precipitate in 2 drops QN HC1. Add
2 drops water and test the solution as follows :

To 1 drop of the solution on a watch glass or spot plate
add 1 drop K 4 Fe(CN) 6 solution.

To another drop of the solution on a watch glass or spot
plate add 1 drop KSCN solution (Question 13).

108 ANALYTICAL PROCEDURE CATIONS

2. Aluminum Ion, Al + ~ f ~ + . a. Pass H 2 S through 2 drops
A1 + " H ~ test solution for a few seconds.

6. To 2 drops A1+++ test solution add 2 drops 6N NH 4 OH.
Now add excess concentrated NH 4 OH to the precipitate
formed.

c. To 2 drops of Al +++ test solution add IN NaOH solu-
tion, drop by drop, until a precipitate forms. Then add
6A r NaOH in excess (Question 14).

To this solution add an equal volume of saturated NH 4 C1
solution. Compare the appearance of the precipitate with
that of the A1(OH) 3 precipitated at first (see Note 43).

d. To 1 drop Al 4 ""*" 4 " test solution add 5 drops water and
3 drops aluminon reagent. Now add 1 drop 6A^ NH 4 OH,
centrifuge, and observe the appearance of the precipitate
at the bottom (see Note 44).

e. Repeat (d) using alizarin S (blue) instead of aluminon
(57).

3. Chromic Ion, Cr ++ +. a. Repeat with Cr + ++ test
solution the tests outlined under the headings Aluminum
(a), (6), and (c).

6. To 2 drops 0+++ test solution add QN NaOH until
the precipitate that first forms redissolves. Now add some
Na 2 O2 (about half the size of a grain of rice). (Caution:
Handle only with a dry spatula as this substance is very
corrosive and decomposes when moistened.) Warm on
the steam bath for 1 min. then divide the solution into two
equal portions (see Note 45).

Acidify one portion with QN HNO 3 . Then make the
solution strongly basic with ammonium hydroxide (see Note
46).

To the other portion add 2 drops saturated NH 4 C1 and 2
drops BaCl 2 solution, stir and filter. Wash the precipitate
with cold water twice, then dissolve it in the smallest possi-
ble volume of &N HC1. Dilute with 2 drops water and
then add a volume of saturated ammonium acetate equal
to the total volume of the solution. If no precipitate

GROUP III METALS 109

appears, scratch the side of the container with a stirring
rod for a few seconds (see Note 47).

4. Manganese Ion, Mn" 1 " 1 ". a. Repeat parts (a) and (6)
under Ferric Ion using Mn" 1 " 1 " test solution in place of
Fe" 1 " 1 " 1 " test solution.

b. To 4 drops Mn +H ~ test solution add 3 drops IN NaOH
solution. Filter and divide the precipitate into four parts.

To one portion of the precipitate add an excess of 67V
NaOH solution.

To the second portion add 1 drop water and a little Na 2 O 2 .

Allow the third portion to stand exposed to air for about
15 min., then compare the results with those obtained with
the Na 2 O 2 .

Dissolve the fourth portion in 3 drops QN HNOs, add 3
drops AgNO 3 solution and a little ammonium persulfate,
(NH 4 ) 2 S 2 O 8 , about the size of a grain of rice. Heat on the
steam bath for 2 to 3 min. Note color of the solution (see
Note 48). This test is both sensitive and specific.

6. Zinc Ion, Zn"^. a. Pass H 2 S through 2 drops Zn ++
test solution for a few seconds.

6. To 2 drops Zn ++ test solution add 1 drop saturated
ammonium acetate solution and saturate with H 2 S. Filter
and wash the precipitate. Dissolve the precipitate in
1:9 HC1 (1 volume concentrated HC1 mixed with 9 volumes
water). Evaporate the solution to dryness, being careful
not to overheat the residue, cool and redissolve in 4 drops
water. Divide this solution into two equal parts.

To one part add 1 drop 1:9 HC1, 1 drop 0.1 per cent
CuSO 4 solution, and 1 drop HgCl 2 -4NH 4 SCN solution.
Scratch and let stand if no precipitate appears at first
(8, 9, 35). This test is sensitive to 0.05 mg. zinc (see page
67). To the second portion of solution add 1 drop NH 4 -
SCN solution and 1 drop pyridine (26).

c. To 2 drops Zn*" 1 " test solution add cautiously, drop by
drop, IN NaOH solution until a copious precipitate forms.
Filter and divide the precipitate into two parts.

110 ANALYTICAL PROCEDURE CATIONS

To one portion of the precipitate add 1 drop saturated
ammonium chloride and 2 drops QN NH 4 OH.

To the second portion add 2 drops QN NaOH (Question
16).

6. Cobalt Ion, Co++. a. Pass H 2 S through 2 drops 00++
test solution for a few seconds (Question 17).

b. To 2 drops Co" 1 " 4 " test solution add 1 drop saturated
NH 4 C1 solution, 2 drops 6N NH 4 OH, and saturate with
HzS. Filter and wash the precipitate with (NH 4 )2S0 4
solution (see Note 49). Divide the precipitate into two
equal parts.

Try dissolving one portion in 1:9 HC1 (Question 18).

To the second portion add 4 drops aqua regia. Evapo-
rate to dryness, then let the residue cool.

Bend the tip of a platinum wire (a 2-in. wire, sealed into
the end of a glass tube which serves as a handle) into a loop

GROUP III ANALYSIS

The unknown may be the filtrate from the precipitation of the sulfides
of Group IT, or an unknown furnished by the instructor. Add 5 drops
bromine water and boil I min. Keeping the soln. hot, add 15 drops saturated
NH 4 C1 soln. and 5 to 6 drops 6AT NH 4 OH. (See Note 52.) Dilute to 4 ml.
with water; heat, and saturate with H^S. Test for completeness of pre-
cipitation, adding more II 2 S if necessary. (See Note 53.) Filter, and wash
precipitate with NlI 4 NOrH 2 S soln. (See Note 54.)

Filtrate. Acidify at
once with HC1 and boil
to half its volume.
Save for Groups IV
and V.

Precipitate. Consists of FeS, ZnS, MnS, CoS,
NiS, A1(OH),, and Cr(OH)..

Add 10 drops aqua regia and warm until pre-
cipitate dissolves. Neutralize with QN NaOH
soln., heat to boiling, and add a few granules
Na2O-2 while stirring. Dilute with 10 drops
water, stir, and filter. Wash precipitate with
cold water.

Filtrate. Contains A1O 2 ~,
CrO 4 ", and HZnOr.

Analyze as directed under
the heading, Group III A.

Precipitate. Consists of Fe(OH) s , MnO 2 ,
Ni(OH)i, and Co(OH) 3 . Analyze as directed
under the heading, Group III B.

GROUP III METALS

111

about the diameter of a match stick. Heat this loop to
redness in a flame and quickly dip it into a heap of pure-
borax. Heat the mass of borax that clings to the loop
until it fuses into a transparent bead ; then, with the bead
still hot, dip it into the residue from the aqua regia evapo-

GROUP III A ANALYSIS

To the filtrate from the subgroup separation add 10 drops saturated
ammonium acetate soln., stir and heat for 2 min. on the steam bath. Filter
and wash the precipitate with cold water. (See Note 55 and p. 42.)

Filtrate. Contains CrO 4 - and HZnO-r
ions. Use for the following tests:

To one-third of the filtrate add 6N
acetic acid with stirring, until the
soln. is faintly acid to litmus. To
1 drop of this soln. on a white spot
plate add 2 drops 50 % AgNO 3 soln.

RED-BROWN PPT.
indicates Cr present.

(See Note 58.)

To the other two-thirds of the filtrate
add 2 to 3 drops of BaCl 2 . Stir, let
stand a minute, then filter.

Filtrate. Contains

HZnOr. Acidify with
6N acetic acid and pass
H 2 S into the soln.
WHITE PPT.
indicates Zn ++ present.
(See Note 59.)

Precipitate.

A
YELLOW

PPT.

confirms Cr
present.

Precipitate. A1(OH) 3 . Dissolve
in 6 drops 2N HC1, and divide
into two portions.

To one portion in a centrifuge tube
add 3 to 5 drops of freshly filtered
aluminon reagent, and make
strongly basic with NH 4 OH. (See
Note 56.)

RED PPT.

indicates Al +++ present.

To the second portion in a centri-
fuge tube add 3 to 4 drops freshly
filtered alizarin S (blue) reagent,
and make strongly alkaline with
NH 4 OH.

BLUE PPT.
confirms Al ff4 present.
(See Note 57.)

Filter off the precipitate and dissolve it in 6N HC1. Evaporate to dryness,
cool and redissolve residue in 5 drops (1:9) HC1. Add to the solution
1 drop 0.1 % CuSO 4 soln., 1 drop HgCl 2 4NH 4 SCN reagent, scratch walls
of tube with stirring rod and let stand.

LAVENDER or BLACK PPT.
confirms presence of Zn++.

(See Note 60.)

112

ANAL YTICAL PROCEDURE C A T1ONS

ration. Heat the bead again until it has become homo-
geneous, let it cool and look through the bead towards a
window or towards a piece of white paper held in a strong
white light. The blue color serves as a test for cobalt (see
Note 50, and Cobalt under Auxiliary Tests) (33).

c. To 1 drop Co"*" 4 " test solution add 5 drops water, 1 drop
Zn" 4 " 4 test solution, 1 drop 1:9 HC1 solution, and 2 drops

GROUP III B ANALYSIS

Dissolve the precipitate from the subgroup in 6 to 10 drops cone. HNOs,
warming, if necessary. (See Note 61.) Add 3 to 4 small crystals of KClOs
and heat for 2 to 3 min. on the air bath, keeping it just below the boiling
point. Cool, add an equal volume of water and filter. Wash precipitate
with cold water.

Filtrate. Contains Ni++, Fe +f+ , Co f + ions.
Make the soln. strongly basic with NH 4 OH, stir,
and filter. Wash precipitate with cold water.

Precipitate. MnO 2 -H 2 O.
DARK BROWN PPT.
indicates Mn present.

Filtrate. Contains Ni 4 + and
Co ++ . Use as follows:

Precipitate :
Fe(OH).,. Dis-
solve in 6
drops 3JVIIC1.
Divide into
two parts.

To one part of
the soln. on a
spot plate add
1 drop KSCN
soln.
RED COLOR
indicates Fe f+4
present.

To half of the precipitate
add 5 drops 6N HNO 3 ,
1 drop 50% AgNO 3
soln. and a few crystals
of ammonium persul-
fate. Heat for 3 to 4
min.
PINK or VIOLET
COLOR
confirms Mn present.

To 1 drop of the filtrate on filter
paper add 1 drop dimethyl-
glyoxime soln.
CRIMSON SPOT
shows Ni ++ present.

To 1 drop of the filtrate on filter
paper add 1 drop a-nitroso-0-
naphthol, and 1 drop 6JV
H 2 S0 4 .
REDDISH SPOT
indicates Co+ + present.

To the other half of the
precipitate add 2 drops
67V HC1 and warm until
it dissolves. Put 1
drop of this soln. on a
filter paper, add 1 drop
6N NaOH, and let
stand about 2 min.
Then add 1 drop benzi-
dine reagent.
BLUE SPOT
confirms Mn present.

To 3 drops of filtrate in a micro-
beaker add 2 drops 6N HC1, 1
drop Zn ++ test soln. and 1 drop
of HgCl 2 -4NH 4 SCN soln.
Scratch walls of beaker arid let
stand.
BLUE PPT.
confirms Co+ + present.
(See Note 62.)

To the second
part of the
soln. add 1
drop K 4 Fe-
(CN) fl soln.
DEEP BLUE
confirms Fe ++ +
present.

GROUP IV METALS 113

HgCl 2 *4NH 4 SCN. If no precipitate forms, scratch the
side of the beaker for a few seconds* (see Coprecipitation).

d. Place 1 drop Co 4 "*" test solution on a piece of filter
paper and hold the moist spot over a diop of concentrated
NH 4 OH on a watch glass for a few seconds. To the same
spot on the filter paper add 1 drop a-nitroso-jS-naphthol
solution and 1 drop QN H 2 SO4.

7. Nickel Ion, Ni 4 "*". a. Repeat tests (a), (6), (c), and
(d) as described under the heading Cobalt Ion, using Ni 4 " 4 "
test solution instead of Co" 1 " 1 " test solution.
. b. Place 1 drop Ni"^ test solution on a piece of filter
paper and hold the moist spot over 1 drop concentrated
NH 4 OH for a few seconds. Then add 1 drop dimethyl-
glyoxime solution (see Note 51 and Question 19).

GROUP IV METALS

CHEMICAL CHARACTERISTICS

Barium. Barium is a silvery metal, having a density of
3.5, and a melting point of 850C. It is extremely active,
tarnishing quickly in air and reacting readily with most
nonmetals. It reacts with cold water giving hydrogen and
a precipitate of Ba(OH) 2 and reacts vigorously with dilute
acids. As its nitrate and chloride are readily soluble, the
best solvent for the metal is very dilute HC1 or HNO 3 .

Compounds of Barium. Neither H 2 S, (NH 4 ) 2 S, NH 4 OH,
nor chlorides cause any precipitation with solutions con-
taining Ba ++ ions.f

* The authors have found this test to be specific in Group III and more
sensitive than the a-nitroso-0-naphthol test. Ferric ion, if present, is
removed by addition of microcosmic salt or weakened, after precipitation,
by dilution of the red solution.

t BaS, does not precipitate in aqueous solutions by the addition of H2S
owing to the fact that it hydrolyzes and forms the acid salt, Ba(HS)2, a
fairly soluble substance.

As Ba(OH)2 is moderately soluble, NH4OH does not furnish a concen-
tration of OH~ sufficiently large to cause its precipitation from any but the
most concentrated solutions of Ba"*"*".

114 ANALYTICAL PROCEDURE CATIONS

Soluble carbonates precipitate BaCO 3 , white, insoluble
in NH 4 OH or alkalies, but soluble in acetic acid and dilute
HC1 or HNO 3 .

Alkalies precipitate some Ba(OH) 2 from strong solutions
of Ba" 1 " 1 ". The precipitate is white, insoluble in excess
alkali, but soluble in acetic acid, dilute HC1, or HNO 3 .

Phosphates precipitate Ba 3 (PO 4 )2, white, soluble in
mineral acids.

Chromates precipitate BaCrO 4 , yellow, only slightly
soluble in strong acetic acid, but soluble in mineral acids.*

Sulfates precipitate BaSO 4 , white, extremely insoluble in
water, alkalies, or acids. Slightly soluble in hot concen-
trated H 2 SO 4 .f

Oxalates give a white precipitate of BaC 2 O 4 , only
slightly soluble in water but soluble in hot dilute acetic acid.

K 4 Fe(CN) 6 gives a white precipitate of Ba 2 Fe(CN) 6 or
BaK 2 Fe(CN) 6 , depending on the relative amounts of Ba ++
and reagent present.

* Solutions of chromates contain chromate, CrC>4~, ions in equilibrium
with dichromate, CraO?"", ions, according to the equation

2Cr0 4 - + 2H + ^ Cr 2 0r + H 2 O

The addition of hydrogen ions favors the formation of dichromate ions, and
lowers the concentration of the chromate ions.

Acetic acid, being slightly ionized, does not furnish hydrogen ions in
sufficient concentration to repress the chromate ions to the point where
BaCrO4 will dissolve. However, HC1 and HNOa will do so.

t In order to get the barium in BaSC>4 into solution either of the following
procedures may be used:

1. By boiling BaS(>4 with successive portions of strong Na2CO3, a
precipitate of BaCO 3 will be left behind. The reaction is:

BaSO 4 + Na 2 CO 3 ^ BaCO 3 + Na 2 SO 4

The BaCOa can then be dissolved in acetic acid, HC1, or HNOa.

2. Fusion of BaSO4 with Na 2 CO3 causes the conversion of BaSO 4 to
BaCOa. The Na 2 S(>4 formed is dissolved in water, leaving the insoluble
BaCOa behind. The BaCO 3 is then dissolved in acetic acid, HC1, or HNO 8 .

3. If a mixture of BaSO 4 , NaaCOg, and powdered charcoal is heated on
a charcoal block in the reducing flame of a blowpipe, the BaSC>4 is reduced
to BaS. The latter is readily soluble in acids.

GROUP IV METALS 115

BaCl 2 and Ba(NO 3 )2 are soluble in water but insoluble in
absolute alcohol or acetone.

Strontium and Calcium. These metals are very similar
to each other and to barium. They are both very active
metals reacting with the same reagents and in the same
manner as does barium. Calcium is a rather soft metal,
being slightly harder than lead. Calcium has a density of
1.55 and a melting point of 810C. Strontium has a
density of 2.6 and a melting point of 752C. The solvents
for these metals are the same as for barium.

Compounds of Strontium and Calcium. With few
exceptions, solutions of Ca"^ or of Sr 4 " 4 " give the same type
of compounds as would solutions of Ba 44 ". The most
notable exceptions are as follows:

Chromates do not give precipitates with neutral or acid
solutions of Ca" f+ or Sr +H ". If the solution is slightly
alkaline, however, Sr ++ gives SrCr0 4 , yellow, readily
soluble in acetic acid, HC1, or HNO 3 , but insoluble in
alcohol or alcohol-water mixture.

Sulfates precipitate SrSO 4 from solutions containing
Sr ++ , but give no precipitate with Ca ++ unless the latter are
present in very high concentrations.

Oxalates precipitate SrC 2 O 4 and CaC 2 O 4 , neither of which
dissolves readily in acetic acid, but both of which are
soluble in mineral acids.

K 4 Fe(CN) 6 gives no precipitate with either Ca 44 " or
Sr 4 " 4 . The anhydrous chloride and nitrate of strontium are
insoluble in either absolute alcohol or anhydrous acetone.
The corresponding salts of calcium, however, are soluble in
both of these solvents.

ANALYTICAL ASPECTS

Owing to the facts that the chlorides of barium, stron-
tium, and calcium are very soluble and that their sulfides
cannot be precipitated from aqueous solutions, they are not

116 A NALYTICAL PROCEDURECA TIONS

precipitated with either Group I, II, or III. However,
their carbonates are precipitated by addition of ammonium
carbonate to an alkaline solution of their ions. Magnesium
carbonate is also precipitated under these conditions, but,
if the alkalinity of the solution is lowered by addition of a
large concentration of an ammonium salt followed by the
addition of the ammonium carbonate, magnesium car-
bonate will not precipitate.

The identification of these elements is based upon (1) the
low solubility of their carbonates, (2) the difference in
solubility of their chromates (see Note 63), and (3) the low
solubility of calcium oxalate, CadO 4 .

PRELIMINARY EXPERIMENTS GROUP IV

1. Barium Ion, Ba^" 1 ". a. To 2 drops Ba f j " test solution
add 1 drop 3N K 2 CrO 4 solution. Filter and wash the
precipitate twice with cold water. Now add 5 drops IN
acetic acid and 5 drops 2.5N ammonium acetate solution
and stir well. Centrifuge (Question 23).

To some of the filtrate add dilute NH 4 OH until faintly
basic. Add 10 drops alcohol and stir (Question 24).

b. To 2 drops of the hot test solution add 1 drop saturated
ammonium chloride solution and 1 drop ammonium car-
bonate reagent (see Note 64). Filter and wash precipitate
with cold water.

Dissolve the precipitate in 2 drops dilute acetic acid,
warming if necessary. Now add 1 drop 3N K 2 CrO 4 solu-
tion, filter, and wash the precipitate once. Make a small
loop in the end of a clean platinum wire (see Note 65) ; dip
the loop into the precipitate until a little adheres to the wire.
Hold for a few seconds in the fumes from a little concen-
trated HC1 in a microbeaker, then hold in the edge of the
hot nonluminous flame of your burner. Note the faint
greenish tinge imparted to the flame. This is an excellent
confirmatory test for Ba"^ (see Note 66).

GROUP IV METALS 117

2. Strontium Ion, Si* 4 " . a. To 2 drops Sr+ 4 test solution
add 1 drop saturated NH 4 C1 solution, heat, and add 1 drop
ammonium carbonate reagent (Question 25). Filter, wash,
and add 2 drops dilute acetic acid to the precipitate.

6. To the solution thus obtained add 1 drop 3N K 2 CrO 4
soln. (Question 26 and Note 67.) Now add 2 drops
dilute NHiOH and 5 drops alcohol. Filter and wash the
precipitate once.

c. To the precipitate from (a) add 5 drops IN acetic acid
and 5 drops 2.57V ammonium acetate solution. Stir (see
Question 27 and Note 68). To the solution thus obtained
add NH 4 OH until it is faintly basic and then add an equal
volume of alcohol. Filter and wash precipitate.

d. Using the precipitate from (6), try the flame test as
described under Barium Ion 6. The crimson flame is a
confirmation of the presence of strontium.

3. Calcium Ion, Ca+ 4 . a. Repeat (2a) using Ca 4 " 4 " test
solution instead of Sr~ f f test solution. Save the solution
for (b).

b. To the solution from (a) add 1 drop 3N K 2 CrO 4 solu-
tion and 2 drops dilute NH 4 OH. Then add an equal
volume of alcohol (Questions 28 and 29). Save the solution
for (c).

c. To the clear solution from (6) add 1 drop 3N potassium
oxalate (KoC^O^ solution. Heat and let stand for 15 min.
or more (see Note 69). Filter. Wash precipitate and save
it for part (d).

d. With the precipitate from part (c) try the flame test
as described under Barium Ion 6. The faint, momentary
brick-red tinge imparted to the flame is an excellent con-
firmation of the presence of Ca +4 \

e. To 1 drop Ca 4 "* test solution add 2 drops water and 1
drop IN K 2 SO 4 solution. Scratch the side of the micro-
beaker and let stand for 5 min. Filter. To the clear*
filtrate add 1 drop 3N potassium oxalate (K 2 C 2 O 4 ), scratch
the walls of the container again, and let stand.

118

ANALYTICAL PROCEDURE CATIONS

GROUP IV ANALYSIS

Unknown (or nitrate from Group III). Acidify the nitrate from Group
III, or the group unknown with 6N HC1 and evaporate to half its volume.
Filter and discard any precipitate. To the nitrate add 6N NEUOH until
the soln. is faintly alkaline, heat on the steam bath, and add ammonium
carbonate reagent, drop by drop, until precipitation is complete. (See
Note 70.) Filter and wash the precipitate twice with cold water.

Filtrate. Save for
Group V.

Precipitate. Contains BaCO 3 , SrCO 3 , and CaCO 3
Dissolve in smallest possible volume of 6N acetic acid
(see Note 71), warming if necessary. Evaporate just
to dryness and cool the residue. Dissolve the residue
in 10 drops of \N acetic acid and add 10 drops 2.5N
ammonium acetate solii. Then add 3N K 2 CrO4
soln., drop by drop, until the clear liquid, after stirring,
and centrifugmg, has a distinct yellow color. Add
3 drops 6W NH 4 OH and more K 2 CrC>4 soln. until the
nitrate again has a persistent yellow color. Add half
its volume of alcohol and stir. Filter, and wash the
precipitate with a mixture of 5 drops alcohol, 5 drops
water, and 1 drop 6N NH 4 OH.

Filtrate. Contains CaCrO 4
and possible small amounts
of SrCrO 4 . Add an equal
volume of alcohol arid 1
drop IN K 2 SO 4 soln. and
let stand. Filter off any
precipitate that forms,
and discard it.

To the clear nitrate add 2
drops 3N potassium oxa-
latcj K 2 C 2 O 4 , soln., heat to
boiling (see Note 72), and
let stand for about 15 min.
Centrifuge.

WHITE PPT.

indicates Ca+ + present.

Confirm by flame test, using
the precipitate above as
described in Preliminary
Experiment 16.

FAINT BRICK-RED
flame confirms Ca ++ present.

Precipitate. Contains BaCrO 4 and SrCrO 4 .
Add 5 drops IN acetic acid and 5 drops 2.5N
ammonium acetate soln., stir, and let stand
for about 3 min. Filter and repeat the
treatment of the precipitate. Combine the
two filtrates. Wash any precipitate with
cold water.

Filtrates, (combined):
contains SrCrO 4 .
Add 3 drops 6N
NH 4 OH and a vol-
ume of alcohol equal
to that of the total
filtrate present. Stir
and centrifuge.
YELLOW PPT.

or cloudiness indicates
Sr++ present.

Confirm by flame test
using the precipitate
above.

CRIMSON FLAME
confirms Sr++ present.

Precipitate. BaCrO 4

The presence of

YELLOW PPT.

at this point indicates
Ba f + present.

Confirm by flame test
using precipitate
above.

GREENISH FLAME
confirms Ba ++ .

GROUP V METALS 119

GROUP V METALS

CHEMICAL CHARACTERISTICS

Magnesium. Metallic magnesium is a silvery metal,
having a density of 1.74, a melting point of 651C., and a
boiling point of 1110C. It is quite active, reacting slowly
with boiling (but not cold) water to give hydrogen and the
rather insoluble magnesium hydroxide Mg(OH) 2 . It com-
bines readily with oxygen and, at high temperatures, reacts
readily with such nonmetals as the halogens, sulfur, and
.even nitrogen (the last reaction resulting in the formation
of magnesium nitride, Mg 3 N 2 ). The metal dissolves readily
in dilute acids forming hydrogen gas and magnesium ions.

Compounds of Magnesium. The compounds of mag-
nesium have properties very similar to those of calcium, for
which it is often mistaken in analysis. The chromate,
MgCrO 4 ; the ferrocyanide, Mg 2 Fe(CN) 6 ; the chloride,
MgCl 2 ; and the nitrate, Mg(NO 3 ) 2 , are very soluble in
water. The sulfides cannot be precipitated from aqueous
solutions as hydrolysis of the sulfide readily occurs.

Phosphates precipitate Mg 3 (PO 4 ) 2 from neutral or
slightly basic solutions of Mg"^ 4 ". From ammoniacal solu-
tions of Mg 4 " 4 ", containing a large concentration of ammon-
ium salts, white, crystalline MgNH 4 P0 4 precipitates. The
latter is soluble in acids.

Arsenates give white Mg 3 (AsO 4 ) 2 with neutral solutions
of Mg 4 " 4 " or white, crystalline MgNH 4 AsO 4 with ammon-
iacal solutions of Mg 4 " 4 ", containing ammonium salts.

Oxalates precipitate MgC 2 O 4 , white, soluble in acids and
in excess alkali oxalates.

Sulfates produce no precipitate, as MgSO 4 is very soluble
in water.

Soluble carbonates produce a precipitate of MgCO 3 ,
white, insoluble in NH 4 OH, but soluble in solutions of
ammonium salts (even in the presence of NH 4 OH) and in
acids.

120 ANALYTICAL PROCEDURE CATIONS

NH 4 OH or alkali hydroxides produce a precipitate of
Mg(OH) 2 , white, gelatinous, insoluble in excess alkali, but
soluble in solutions of ammonium salts and in acids.

Potassium, Sodium and Ammonium.* Sodium and
potassium are very soft metals, having, when freshly
cut, a silvery luster. The former has a density of 0.97, a
melting point of 97.5C., and a boiling point of 880C.
Potassium has a density of 0.86, a melting point of 62.3C.,
and a boiling point of 760C. Both are extremely active
metals, potassium being slightly the more active. Both
tarnish almost instantly in air forming oxides, Na 2 O 2 ,
K 2 O4, and K 2 O; both react vigorously with water giving
the hydroxide and hydrogen; both combine readily with
nonmetals, many of them at ordinary temperature. With
acids, their reaction is so rapid as to be dangerous. The
safest solvent is ethyl alcohol.

Potassium, sodium, and ammonium are monovalent in
all their compounds.

Compounds of Potassium, Sodium, and Ammonium.
Practically all the compounds of these three ions are soluble
to a greater or less extent. The hydroxides, sulfates,
chlorides, nitrates, carbonates, phosphates, arsenates,
chromates, sulfides, and ferrocyanides are all very soluble.
However, NH 4 OH is very unstable towards heat, and
breaks up giving NH 3 and water. Sodium oxalate is less
soluble than the oxalates of the other two ions.

A few of the compounds of sodium, potassium, and
ammonium are sufficiently insoluble to allow their use in

* Although ammonium is not an element, the radical NH4 exhibits certain
peculiarities that justify its treatment as a metal. Its base- and salt-forming
properties are already familiar to the student. In addition, however, it has
the ability to form an amalgam with mercury. This is easily demonstrated
by either electrolysing a strong solution of NH 4 C1, using a pool of mercury
as the cathode, or by adding some sodium amalgam (a solution of metallic
sodium in mercury) to a saturated solution of NH^Cl. If either of these
operations is carried out at a low temperature, an amalgam of ammonium
is formed, the truth of this being demonstrated by allowing the amalgam to
warm, and noting the escape of hydrogen and NH 3 gases from the mass.

GROUP V METALS 121

analysis. As these will be considered under Preliminary
Tests and under Auxiliary Tests, no further discussion is
necessary here.

ANALYTICAL ASPECTS

The members of this group remain in solution throughout
all the previous precipitations. Magnesium readily forms
a number of slightly soluble compounds but sodium,
potassium, and ammonium ions form very few slightly
soluble substances. In view of the fact that ammonium
ion is added many times during the analysis, it is necessary
to test the original solution for that ion.

Most of the reagents that precipitate potassium, also
precipitate ammonium ions. It is therefore necessary to
destroy the ammonium ion before testing for potassium.
This is done by oxidizing the ammonium ion with aqua
regia and igniting the residue to drive off the last traces of
ammonium compounds. Since none of the other ions
interfere with each other in these tests, it is not necessary
to do any separation other than this.

PRELIMINARY EXPERIMENTS GROUP V

1. Magnesium Ion, Mg +t ~. a. To 1 drop Mg"^ test
solution add 1 drop ammonium carbonate solution (without
NH 4 C1).

Repeat this test using 1 drop Mg"*" 4 " test solution, 1 drop
saturated NH^Cl solution (see Question 30), and 1 drop
ammonium carbonate reagent.

b. To 1 drop Mg++ test solution add 1 drop QN NH 4 OH.
Now try redissolving the precipitate in saturated NH 4 C1
solution (see Question 31).

To this solution add 1 drop NaH 2 PO 4 solution. Scratch
inside wall of microbeaker and let stand a few minutes.

c. To 1 drop Mg ++ test solution add 1 drop p-nitro-
benzeneazoresorcinol reagent. Stir and centrifuge.
Divide the "lake" or precipitate into two portions (58).

1 22 A NALYTICAL PROCEDURE C A TIONS

To one portion of the "lake" add QN HC1.
To the second portion of the "lake' 1 add QN NaOH (see
Note 73).

2. Potassium Ion, K+. a. To 1 drop K+ test solution
add 1 drop tartaric acid solution (see Note 74).

6. To 1 drop K+ test solution add 1 drop freshly prepared
and filtered sodium cobaltinitrite, Na 3 Co(NO 2 )6, solution
(see Note 75).

Repeat using instead of the pure K" 1 " ion solution a mix-
ture of K+ and Ag+ (59) (see Note 76).

c. To one drop K + test solution add 1 drop picric acid
solution and let stand.

d. Evaporate 1 drop K + test solution to dryness. With
a platinum wire try a flame test on the residue, looking at
the flame through a piece of blue (cobalt) glass. The flame
test is capable of detecting as little as 6 X 10~ 4 mg. potas-
sium (see Note 77).

Repeat this test once more using a mixture of K + test
solution and an equal volume of Na 4 " test solution.

Repeat this test once more using Na+ test solution alone
(see Note 78 and Question 32). It is possible to detect as
little as 3 X 10~ 6 mg. sodium by this method.

3. Ammonium Ion, NH 4 +. a. Repeat the tests under
the heading Potassium Ion, parts a, 6, and c, using NH 4 +
test solution instead of K+ test solution (Question 33).

6. Precipitate some (NH 4 ) 2 NaCo(NO 2 )6 by adding 1 drop
saturated sodium cobaltinitrite solution to 1 drop NH 4 +
test solution. After precipitation occurs, heat the mixture
until the precipitate dissolves. Cool, then add more
sodium cobaltinitrite solution.

Repeat the above test, using K+ in place of NH 4 + (see
Note 79 and Question 34).

c. In the test-tube portion of the gas evolution apparatus
put 3 drops NH 4 + test solution and evaporate just to dry-
ness (see Question 35). Cool, add 2 drops QN NaOH, and
put the top of the apparatus in place. In the top tube

GROUP V METALS 123

place a strip of moist red litmus paper, close the top with a
loose plug of cotton, and warm the test tube for a few
minutes, observing any change in the color of the litmus.

4. Sodium Ion, Na + . a. Repeat all tests described
under the title Potassium Ion using Na+ test solution
instead of K+ test solution (see Question 36).

b. To 1 drop Na+ test solution on a watch glass (using
a. black background) add 8 drops zinc uranyl acetate
solution. Rub the glass in contact with this mixture with a
stirring rod a few seconds and let stand. This test will
detect as little as 0.00001 g. sodium. Neither K+, NH^,
nor Mg"*" 1 " interferes (60, 61).

GROUP V ANALYSIS

Group unknown (or filtrate from Group IV). Divide into six equal parts
and perform the following tests.

To one portion add 1 drop saturated NH 4 C1, 1 drop 3N NH 4 OH, and 1 drop
NaH 2 PO 4 . Rub inside wall of beaker with a stirring rod and let stand
10 minutes.

WHITE PPT.

indicates Mg + + present.

To another portion of the unknown made faintly acid (see Note 73) add
an equal volume of p-nitrobenzeneazoresorcinol reagent. Centrifuge.

BLUE LAKE or PPT.

indicates Mg++ is present.

Combine two portions of the unknown and evaporate to dryness. Add
10' drops aqua regia and evaporate again to dryness, using a crucible
as the container. Ignite the residue (heat the crucible to redness) until
no more white fumes escape (Question 37). Cool, then dissolve the residue
in 3 drops water. Use this solution for the following tests.

To 1 drop of this solution on a watch glass add 1 drop of sodium cobalti-
nitrite solution (see Notes 74 and 75). Let stand.

PALE YELLOW PPT.
indicates K + present.

Confirm by heating in a microbeaker until precipitate redissolves, cooling
and adding more reagent. Reappearance of the precipitate is excellent
confirmation for K 4 " ion.

124 ANALYTICAL PROCEDURECATIONS

To another drop of the prepared solution apply the flame test using the
platinum wire and blue glass.

LAVENDER FLAME visible through BLUE GLASS

con fir ins presence of K+.

(See Note 77.)

Combine two portions of the unknown and evaporate to dryiiess Cool
and dissolve residue in 3 drops water. Tost as follows.

To 1 drop of this solution on a watch glass (black background) add 8
drops zinc uranyl acetate solution. Rub inside of container with a
stirring rod arid let stand 5 min.

PALE YELLOW PPT.

indicates Na + present.

Try the flame test with the remainder of the concentrated solution (see
Note 80).

BRIGHT YELLOW LASTING FLAME

confirms presence of Na + .

The test for NH 4 + ion must be made on the original general or gioup
unknown before any reagents are added to it. The test is carried out as
follows :

Evaporate 3 drops of original unknown just to dryness in the gas evolution
tube (see Note 81). Cool, then add 2 drops 6N NaOH solution to the
residue. Quickly place a strip of moist red litmus paper inside of the bulb.
Warm gently on the steam bath.

Uniform BLUE COLOR to litmus
indicates NHU+ present.

(See Note 82.)
(For additional tests, see Ammonium under Auxiliary Tests.)

AUXILIARY TESTS

Sometimes it happens that, owing to the presence of
unusual elements in unknowns or to slight errors in pro-
cedure, the final identification tests do not give clear-cut
results. In such cases it is often necessary to carry out
additional tests in order to confirm or disprove the con-

AUXILIARY TESTS 125

elusions drawn from the usual tests. A few such tests are
given here under the headings of the ions for which they are
used. For the sake of brevity, only the essentials of the
tests are included ; and if any extensive use is to be made of
them it is suggested that a study be made of the literature
cited here. If the original publications are not available,
the abstracts should be used, as the information included in
them is sufficient for many purposes.

Silver. 1. Zinc purpurate gives a violet precipitate with
solutions containing silver ions (see also Mercury).
Instructions for preparing the reagent are given by
G. Deniges (1).

2. Diphenylthiocarbazone (dithizon) reagent with silver
ion in slightly alkaline solutions gives a fine violet precipi-
tate (see also lead, copper, zinc, cadmium, and cobalt)
(2, 16).

Mercury. 1. A drop of solution containing mercurous
ion placed on a piece of filter paper and treated with a drop
of NaNO 2 solution, gives a dark stain of metallic mercury.
Silver interferes, but the colored products of copper, iron,
nickel, cobalt, and chromium can be washed off (3).

2. If a drop of a solution containing mercurous ion and a
drop of aniline are added to a drop of SnCl 2 solution on
filter paper, a black stain of metallic mercury will result (4).

3. A small strip of copper placed in a solution containing
either mercurous or mercuric ion will become coated with
metallic mercury. To distinguish this from the silvery
coating given by other metals below copper in the displace-
ment series, the mercury may be rolled into tiny balls by
rubbing with the edge of the thumb nail. If the quantity
is too small for this, the copper strip may be heated in a
tube made by sealing one end of a 5-in. piece of 3-mm.
tubing, and the condensation of mercury in the cooler
portions of the tube observed.

4. A solution of zinc purpurate gives (in absence of
silver ion) a peach-colored precipitate, with mercuric ion.

126 ANALYTICAL PROCEDURE CATIONS

The reagent is added to the sodium acetate solution of
mercuric ion formed after evaporation of the Hg, HgNH 2 Cl,
or HgS precipitates with aqua regia (1).

5. Diphenylthiocarbazone has also been suggested as a
reagent for mercury (16).

Lead. 1. A drop of a solution containing lead ions, on
filter paper, converted to PbO 2 by treatment with H 2 O 2 and
NH 4 OH and heated in a jet of steam to decompose the
unused H 2 O 2 , will give a blue spot with a 0.1 per cent solu-
tion of tetramethyl-diaminodiphenylmethane in 10 per cent
acetic acid. Bismuth does not interfere (5, 6).

2. In an alkaline solution containing KCN, lead ions
will give a red-violet coloration with diphenylthiocarbazone
reagent. Under these conditions, silver, copper, cadmium,
nickel, zinc, and antimony do not interfere (2, 6).

3. A trace of bismuth, followed by a drop of sodium
stannite solution, added to an alkaline solution containing
lead ions will produce a dark stain of metallic lead in a
short time. The bismuth alone gives the test if added in
sufficient quantities, so a blank should be run at the same
time (7).

Copper. 1. If, to a faintly acid solution containing
copper ions, an excess of zinc ions and a drop of HgCl*-
4NH 4 SCN solution are added, a colored precipitate will be
obtained. The color will range from lavender, for minute
traces of copper, to a purple-black, for greater concentrations
of copper. If copper is present in greater concentration
than zinc, an apple-green precipitate of CuHg(SCN) 4 will
form (8, 9).

2. Diphenylthiocarbazone reagent gives a yellow-brown
coloration with copper in neutral, or slightly ammoniacal,
solutions. Lead, cadmium, zinc, and nickel do not
interfere (2, 11).

3. If a solution of copper ions is sufficiently concen-
trated, the addition of KI solution will produce free iodine
and a white precipitate of cuprous iodide. Many metallic

AUXILIARY TESTS 127

ions interfere with this test, the reaction serving better as a
test for moderate concentrations of iodide ion.

Bismuth. 1. If a solution of BiCl 3 or Bi(NO 3 ) 3 is
treated with a 1 per cent solution of dimethylglyoxime and
made strongly ammoniacal, a voluminous yellow precipi-
tate will form (10). If Bi 2 (SO 4 )3 is used in place of the
chloride or nitrate, a white precipitate will result (11).

2. If a drop of a solution containing lead ions on filter
paper is treated with a solution containing bismuth ions,
a drop of QN NaOH, and a drop or two of sodium stannite
solution, a black deposit of metallic lead will form (7).
The bismuth acts as a catalyst for the reduction of the lead
and this reaction serves as a very sensitive test for bismuth.
A blank should be run at the same time, however, as the
lead will be reduced, in time, whether bismuth is present or
not. Easily reduced ions, such as silver and mercury,
interfere.

3. A drop of bismuth ion in 10 per cent HNO 3 on filter
paper, treated with a dilute solution of KI, will give a black
stain of BiI 3 . The stain should be washed with distilled
water to remove iodine set free by ferric or cupric ions.
Excess KI converts it to a solution of yellow BiI 4 ~ ion
which soon hydrolyzes to orange BiOI. Mercury inter-
feres (12).

Cadmium. 1. If an ammoniacal solution of cadmium ion
containing excess KCN is treated with formaldehyde and
boiled, a precipitate of Cd(OH) 2 will form. In the presence
of dinitrodiphenylcarbazide, the precipitate will be blue.
The precipitate of Cd(OH) 2 may be dissolved in dilute
HC1 and other confirmatory tests applied (13).

2. If a drop of a solution containing cadmium ion is
added to a drop of diphenylcarbazide reagent on filter
paper, a reddish-violet coloration will be produced. In the
presence of both lead and cadmium, a yellow ring of PbI 2
will be surrounded by the reddish-violet ring due to the
cadmium. Copper, bismuth, and many other ions interfere.

1 28 A NALYTJCAL PROCED9 WE- CM T/ONS

The reagent is a saturated solution of diphenylcarbazide
in 90 per cent alcohol to which has been added a little
KCNS and KI (14).

3. A solution containing cadmium ions, treated with a
mixture of solid Na 2 CO 3 and charcoal in a test tube, evapor-
ated to dryness, then heated strongly, will give a mirror of
cadmium metal edged with brown, in the cooler portions of
the test tube. If sulfur is added to the test tube and
distilled over the mirror, a layer of CdS will form which will
be orange while hot, but yellow when cool (15).

4. Diphenylthiocarbazone (2, 6, 16), a 0.5 per cent
solution of l-(2-quinolyl)-4-allyl thiosemicarbazide in 50 per
cent alcohol (17), and HgCl 2 -4NH 4 SCN solution (8) have
also been suggested as reagents for cadmium. The latter
is very slow to produce a precipitate.

Arsenic. 1. CuSO 4 and KOH solutions added to a solu-
tion of As*" 1 " 1 " will, on heating, give a yellow precipitate
which, on longer heating, turns orange or red. The
precipitate is Cu 2 O, formed by the reducing action of the
arsenious ion.

2. The solution of As"*" 1 "*" 1 " 1 " obtained by evaporating the
arsenic sulfide precipitate in Group II with aqua regia and
dissolving the residue in 3N HNO 3 will give a white
crystalline precipitate of MgNH 4 AsO 4 if treated with
magnesia mixture [a solution of MgCl 2 and NH 4 OH to
which has been added enough NH 4 C1 to prevent precipi-
tation of Mg(OH) 2 ] and allowed to stand. The usual
precautions against supersaturation must be taken,
however.

3. A neutral solution of AsO 4 EH will give a chocolate-
colored precipitate with AgNO 3 . AsO 3 s gives a white
precipitate with the same reagent. Chlorides, iodides,
bromides, and many other anions interfere.

The test may be carried out by bringing a drop of a
solution of As 4 "^"^ and AgNO 3 in 6N HNO 3 into contact
with the edge of a drop 6N NH 4 OH on a glass slide. At

AUXILIARY TESTS 129

the interface between the liquids a chocolate-colored
precipitate will form.

4. A hot solution of ammonium molybdate (3.6 g.
MoO 3 dissolved in 7 ml. concentrated NH 4 OH and 15 ml.
water, followed by addition of a mixture of 25 ml. concen-
trated HNO.J and 53 ml. water with stirring) will give with
AsO/^ a yellow precipitate of ammonium arsenomolybdate,
(NH 4 ) 3 AsO 4 -12MoO 3 .

Tin. 1. A separation of tin and antimony may be made
by adding to the slightly acid solution of the mixture a
solution of oxalic acid, passing in H 2 S, precipitating and
filtering off the antimony as Sb 2 S 3 . The filtrate is treated
with NH 4 OH and calcium chloride solution to precipitate
the oxalate ion. The solution is made very weakly acid
with HC1, arid H 2 S is passed into the clear filtrate to give a
yellow precipitate of SnS 2 . This test is based on the fact
that H 2 S will not precipitate SnS 2 in the presence of
oxalate ion.

2. A slightly acid solution containing tin and antimony,
treated with excess of chlorine water and boiled, will pre-
cipitate the antimony as kSb 2 O 5 . The clear filtrate will then
give a yellow precipitate of SnS 2 on adding H 2 S (25).

3. If zinc is added to a strongly acid (HC1) solution of tin
ions, a small test tube filled with water held in the reacting
mixture for about 5 min. and the bottom of the test tube
then held in the nonluminous flame of a burner, a pale blue
coloration will be imparted to the flame. Many other
metals (e.g., copper and lead) interfere, giving their own
characteristic colors. The test is significant only if
positive (18, 19).

4. If a drop of a solution containing tin (Sn + " f or Sn ' ' ' ' )
is mixed with a drop of Q.Q5N KI solution and a drop of this
mixture is brought into contact with a drop of concentrated
H 2 SO 4 on a glass slide, a fine yellow precipitate of SnI 2 or
SnI 4 will form at the interface between the two liquids, grad-
ually spreading into a broader band. Antimony and arsenic

130 ANALYTICAL PROCEDURE CATIONS

interfere somewhat, giving iodine and a yellow orange
precipitate. This, however, usually appears only as a nar-
row line at the interface without spreading. Controls
should be run in case of doubt (20, 21).

5. A borax bead tinted blue by dipping in strong
Cu(NOs)2 solution and heating will give with tin (even in
traces, and in any type of compound) a red or reddish-
violet bead. The color is best developed by heating in a
reducing flame. The bead is colorless while hot, the color
appearing rather suddenly while cooling (22).

6. Nitrophenylarsonic acid has also been suggested as a
precipitation reagent for stannic ions in dilute acid solu-
tion (23). Diphenylthiocarbazone gives reddish color-
ations with stannous ions (16). Stannous ions added to a
dried precipitate of Hg 2 Cl 2 on filter paper followed by the
addition of a drop of aniline gives a dark spot of metallic
mercury (very sensitive) (4).

Antimony. 1. A strongly acid (HC1) solution containing
antimony ion treated with a drop of 0.05N KNO 2 and
heated to decompose excess HN0 2 will give a violet color
on addition of a drop of a 0.1 per cent solution of rhodamine
B (tetraethylrhodamine) in water. Tungsten and oxidizing
agents interfere. As little as 0.0005 mg. antimony can be
detected in the presence of 12,400 times as much tin by
this test (24).

2. By boiling a slightly acid solution with excess chlorine
water, antimony can be precipitated as Sb 2 O 6 (separation
from tin). The resulting precipitate dissolved in HC1 may
be tested by any of the usual methods (25).

Iron. 1. Ferrous ion gives a dark blue precipitate of
ferrous ferricyanide with freshly prepared K 3 Fe(CN) 6
solution. Ferric ion gives only a light brownish solution
with this reagent.

2. A solution containing ferrous ions gives a light red
precipitate with dimethylglyoxime. Nickel interferes.

AUXILIARY TESTS 131

3. Solutions containing ferric ions give a red-brown pre-
cipitate of basic ferric acetate on boiling with acetate acid
mixture.

Aluminum. 1. A drop of a solution of aluminum ion
containing a trace of cobalt on a filter paper which was
previously treated with saturated KC1O 3 solution and dried
will, on treatment with NH 3 fumes and burning, leave an
ash that varies in color from an olive green to a light blue.
In the absence of aluminum ion, the ash is black.

2. If solutions of ammonium chloride and KCNO are
added to a solution containing aluminum ions and the mix-
ture is heated, a granular precipitate of A1(OH) 3 will form.
Prolonged heating reduces the effectiveness of this pre-
cipitation (26).

Chromium. 1. A solution containing CrC^" gives a blue
coloration with benzidine reagent. The test may be used
on a HC1 extract from the precipitate of BaCr0 4 (27).
The test is not very satisfactory, however.

2. Orcin (s-dihydroxytoluerie) gives a brown stain with
CrOr 1 and strong HC1 (6). A 1 per cent solution of
strychnine in concentrated H 2 SO 4 added to a solution of
CrO 4 == gives a blue- violet color, slowly turning red. Manga-
nese, cobalt, Fe(CN) 6 Ss , and oxidizing agents interfere (28).

3. Tetramethyldiamino-diphenylmethane gives a bluish-
purple coloration if added to a buffered acetic acid solution
of chromate ion. To test for chromic ions, it is best to add
the reagent (a 0.1 per cent solution in 10 per cent acetic
acid) to the solution, make the latter alkaline with NaOH,
and add a few grains of Na 2 2 . Warm on the steam bath
until bubbling has practically ceased, cool, and make acid
with acetic acid. A bluish-violet color appears becoming
more intense and more reddish on standing. Carried out
in this manner on a portion of the filtrate from the alumi-
num test (see Group III Procedure), this serves as an
unusually certain, moderately sensitive test for chromium.

132 ANALYTICAL PROCEDURE CATIONS

In very dilute solutions the test is faint and fleeting, espe-
cially if too great excess of Na 2 O 2 is used.

Zinc. 1. Diphenylthiocarbazone forms with zinc ions a
red-purple precipitate soluble in the chloroform of the
reagent solution. The change is startling, the color going
from the rich green of the reagent to a pink or red. Moder-
ate quantities of manganese, cobalt and nickel will not
interfere if the solution is made strongly ammoniacal and
allowed to stand for a minute or two before adding the
reagent. Interference by silver, copper, mercury, gold,
bismuth, cadmium, and lead is masked by a preliminary
treatment of the solution with Na 2 S 2 O 3 . With Na 2 S 2 O3
and KCN (in faintly acid solution to mask nickel, cobalt,
and palladium) the reagent is specific for zinc (2, 29).

2. A test developed by Benedetti-Pichler makes use of
ashless filter paper soaked in a solution of 4 g. K 3 Co(CN) 6
and 1 g. KC1O 3 per 100 ml. water and carefully dried.
A drop of zinc solution in the center of a 1-in. square of such
paper, dried to a brown spot and burned, will leave a green
ash if zinc is present to a concentration of 1 mg. per milli-
liter. Special methods of using this test will detect as little
as 0.0006 mg. of zinc. The ash is best observed if caught
on white porcelain (30).

3. A solution containing zinc ions will give a white
precipitate if treated with pyridine and solutions of KCNO.
The test is less sensitive, however, than the test with
pyridine and NH 4 CNS (see Group III) (26).

4. Jeffreys and Swift (31) have shown that zinc may be
precipitated as the sulfide by passing H 2 S into a sulfate-
hydrosulfate buffer solution, the hydrogen ion concen-
tration of which may be as high as 2.5 X 10~ 2 mol per liter.
In this manner zinc may be separated from nickel, iron,
manganese, chromium, and aluminum. Cobalt, however,
comes down with the zinc.

Cobalt. 1. A cold, strongly ammoniacal solution of
cobalt will give, with solid Na 2 S 2 O 4 (sodium hyposulfite), a

AUXILIARY TESTS 133

yellow solution which gradually changes to orange, red,
dark brown, and finally forms a black precipitate. Other
metals form a brownish-red tint with the solid Na 2 S2O 4 .
On filtering the mixture, however, an orange filtrate is
obtained whose color deepens on addition of more reagent
(32).

2. If a solution containing cobalt ions is saturated with
KC1, treated with a drop of KNO 2 solution, acidified with
acetic acid and warmed, a fine yellow precipitate of K 3 Co-
(NO 2 )e will form. The solutions must be stirred with
scratching and allowed to stand awhile, as the salt super-
saturates readily.

3. Diphenylthiocarbazone may be used in Group III as a
specific test for cobalt if the solution is first made alkaline
with 2 per cent NaOH. A bluish-violet color first appears,
soon fading to a colorless gray (2).

4. The borax bead test described in the preliminary tests
on Group III serves as an excellent test for cobalt, giving,
according to Curtman and Rothberg (33), a blue bead in a
mixture of as high as 95 per cent nickel with 5 per cent
cobalt.

Nickel. 1. If, to a neutral or slightly alkaline solution
containing nickel ions, a little ethylenediamine and Na 2 S 2 O 3
are added, a violet crystalline precipitate will form.
Although this test is not so sensitive as the test using
dimethylglyoxime, it is specific in the presence of large
concentrations of iron, cobalt, copper, and chromium
(34).

2. In the absence of cobalt, the borax bead test (see
Preliminary Experiments) gives a reddish-brown color if
used on the nickel filtrate in Group III.

Manganese. 1. A solution of manganese ions in con-
centrated nitric acid gives, on boiling with solid Pb0 2 , a
pink or violet solution of HMnO 4 (very specific).

2. A drop of a solution of manganese ions made alkaline
with NaOH and evaporated to dryness will, on treatment

134 A NALYTICAL P ROC ED URECA TIONS

with a 1 per cent solution of strychnine in concentrated
H 2 SO 4 , give a blue-violet color, slowly turning red (28).
Cobalt interferes, but its interference is removed by pre-
liminary treatment with KCN. Oxidizing agents interfere.

3. A sodium carbonate bead treated with manganese
ions, touched while hot to a small crystal of KC1O 3 and
fused, will give a green bead owing to the formation of
sodium manganate (Na 2 MnO 4 ).

Calcium. 1. Calcium may be separated from barium
and strontium by converting them to the nitrates, heating
strongly to 'form the anhydrous salt, and treating with
anhydrous acetone or absolute alcohol. Calcium nitrate is
soluble in these liquids, whereas the nitrates of barium and
strontium are insoluble.

Sodium. 1. A neutral solution containing sodium ions
will give a white precipitate of Na 2 H 2 Sb 2 O7 if treated with a
saturated alkaline solution of potassium pyroantimoniate
(K 2 H 2 Sb 2 O 7 ) and allowed to stand. Neither potassium nor
ammonium ion will give this test.

2. A solution containing sodium ions will give a yellow
precipitate of NaCo(UO 2 )3 < (C 2 H 3 O 2 )9 if treated with a solu-
tion made by mixing equal volumes of a 4 per cent solution
of uranium acetate in 3 per cent acetic acid and a 20 per cent
solution of cobalt acetate in 3 per cent acetic acid. The
reagent should be allowed to stand over night and should
be filtered before using (36).

Potassium. 1. A saturated solution of picric acid in
water gives a yellow precipitate with solutions containing
potassium ions. Ammonium ion interferes and must be
destroyed before the test is applied.

2. Strong perchloric acid added to solutions containing
potassium ions produces a white precipitate of KC1O 4 .
Ammonium ions interfere. Chloroplatinic acid will pre-
cipitate potassium as yellow potassium chloroplatinate,
K 2 PtCl 6 . Ammonium ions give a precipitate of the same
appearance, but on igniting, then dissolving the residue in

AUXILIARY TESTS 135

water and adding more reagent, the ammonium gives no
precipitate, while the potassium does.

3. A solution of tartaric acid reacts with neutral solutions
of potassium ion to produce a precipitate of KHC 4 H 4 06.
Ammonium ion produces a similar precipitate of
NH 4 HC 4 H 4 O 6 .

4. A 0.1 N solution of 5-nitrobarbituric acid gives a
precipitate with as little as 0.09 mg. of potassium per milli-
liter. All other ions of Groups IV and V interfere. How-
ever, to give the same amount of precipitate requires one

4 hundred times as much sodium as potassium (40).

Ammonium. 1. The tests described under the title
Potassium may be used for ammonium ions in the absence
of the former. The best way to use them is to make the
solution alkaline with NaOH and distill, immersing the
delivery tube in a drop or two of distilled water containing
a little HC1. As NH 3 is readily volatile, it will distill over
and the tests can be applied to the resulting solution. An
alternative is to hold a drop of the reagent in the tip of a
glass tube over the solution made alkaline with NaOH and
to warm the latter, observing the appearance of cloudiness
in the reagent.

2. Nesslcr's solution, an alkaline solution of KaHg]^,
turns brownish if added to solutions containing ammonium
ions. The depth of color is proportional to the concen-
tration of ammonium ions.

Magnesium. 1. An ammoniacal solution of magnesium
ions buffered with NH 4 C1 will give a greenish-yellow pre-
cipitate on addition of 8-hydroxyquinoline (Oxine) (37, 38).
No members of Group V interfere, but several other metals
give precipitates of various kinds.

2. An ammoniacal solution of magnesium ions free of
ammonium salts, placed on a piece of filter paper and
warmed to drive off all ammonia, will, on treatment with
phenolphthalein, give a red spot. This will disappear on
drying and reappear on moistening (39).

136 ANALYTICAL PROCEDURE CATIONS

3. Sodium or ammonium carbonate will, in absence of
ammonium salts, precipitate MgCO ;t , soluble in excess of
saturated ammonium chloride solution.

REACTIONS INVOLVED IN SEPARATIONS

[The student should balance each of these for practice,
indicating precipitates by ( j. ) and gases by ( | ).]

GROUP I

Ag + + Cl~ -> AgCl (white)
AgCl + NH 3 -> Ag(NH 3 ) 2 + + Cl-
Ag(NH 3 ) 2 + + H+ + Cl--> AgCl + NH 4 +
Ag(NH 3 ) 2 + + I- -* Agl (yellow) + NH 3
Pb++ + Cl- - PbCl 2 (white)
PbCl 2 + CrO 4 = -> PbCrO 4 (yellow) + Cl~
Hg 2 ++ + Cl- - Hg 2 Cl 2 (white)

Hg 2 Cl 2 + NH 3 - Hg (black) + HgNH 2 Cl (white) +
NH 4 + + Cl-

GROUP II

Hg++ + 8- -* HgS (black)

Bi+++ + S -> Bi 2 S 3 (black)

Bi 2 S 3 + HNO 3 - Bi +++ + NO + S + H 2

Bi(OH) 3 + SnOr -* Bi + H 2 O + SnOr
Cu++ + S" - CuS (black)
CuS + HNO 3 -> (see page 36)
Cu++ + OH- -> Cu(OH) 2 (soluble in NH 4 OH)
Cu++ + NH 3 -> Cu(NH 3 ) 4 + -*- (deep blue color)
Cu(NH 3 ) 4 ++ + H+ - Cu ++ + NH 4 +
Cu++ + Fe(CN) 6 B - Cu 2 Fe(CN) 6
Cu(NH 3 ) 4 ++ + CN- - Cu(CN)r + NH 3 + C 2 N 2
Cu(CN) 3 ~ + H 2 S no reaction.

Cd" 1 "*" all reactions are similar to those of copper except the
following:

REACTIONS INVOLVED IN SEPARATIONS 137

Cd(NH 3 ) 4 ++ + ON- - Cd(CN)r + NH,
Cd(CN),- + H 2 S -> CdS (yellow) + CN~ + H +

+ J- _ AS+++ 4. l t

AS+++ + 8- - As 2 S 3 (yellow)

As 2 S 3 + S- (from (NH 4 )tS) - AsS,"

AsS 3 = + H+ - As 2 S + H 2 S

As 2 S s + HC1 > no reaction

As 2 S s + HC1 + HNO, -* AS+++++ + H 2 O + Cl~ + NO +

S
AsO 4 s + Ag+ -^ Ag 3 AsO 4 (chocolate colored)

+ H 2 O + SO,- - AS+++ + SOr + H+
+ OH" + Zn - AsH s + HZnOr

+ g= _ Sbag3 ( orange )

Sb 2 S s + H + -> Sb^+ + H 2 S

Sb+++ + S 2 O 3 = - Sb 2 OS 2 (orange) + SO

+ H+ + Zn - SbH 3 + ZH++
+ S~ - SnS (brown)
SnS + H+ -> Sn++ + H 2 S
Sn 4-f^+ + g- _^ Sn s 2 ( ye n ow )

SnS + (NH 4 ) 2 S no reaction

SnS 2 + (NH 4 ) 2 S - SnS 3 = + NH 4 +

SnS 3 = + H + -* SnS, + H 2 S

SnS 2 + H+ -* Sn ++++ + H 2 S

(+H+) + Zn -* Sn++ + ZH++
+ Hg++ + Cl- -* Sn++^ + Hg 2 Cl 2 (white)
+ H g2 Cl 2 -* Sn++ ++ + Hg (black) + Cl~

GROUP III

+ Br 2 -* Fe- f++ 4- Br~

+ OH- - Fe(OH) 3 (red-brown)
Fe(OH) 3 + S- - FeS + S + OH~
FeS + H+ -> Fe++ + H 2 S
Fe++ 4- O 2 = + H 2 O - Fe(OH) 3 + OH~
Fe(OH) $ + H + -* F6+++ + H 2 O

+ Fe(CN) 6 - Fe 3 (Fe(CN) 6 ) 2 (dark blue)

138 ANALYTICAL PROCEDURE CATIONS

Fe+++ + SON- - Fe(SCN) 3 (deep red)

Mn++ + OH- - Mn(OH) 2 (white)

Mn(OH) 2 + O 2 -> MnO(OH) 2 (brown)

Mn(OH) 2 + S = - MnS (flesh-colored) + OH~

MnS + H+ - Mn + + + H 2 S

Mn++ + H 2 O + Or -* MnO 2 (brown) + OH~

Mn-"- + CIO,- -> Mn0 2 + Cl~

MnO 2 + H+ + S 2 O 8 = (+Ag+) - MnOr + H 2 O + SO 4 "

Cr-*-"- + OH" - Cr(OH) 3 (purple or green)

Cr(OH) 3 + S = > no reaction

Cr(OH) 3 + H + -* Cr +++ + H 2 O

Cr(OH) 3 + OH- (strong NaOH) - CrOr + H 2 O

CrO 2 ~ + O 2 = + H,O - CrO 4 = + OH~

CrO 4 = + Ag + -> Ag 2 CrO 4 (red-brown)

CrO 4 ~ + Ba ++ -* BaCrO 4 (yellow)

BaCrO 4 + H+ - Ba++ + Cr 2 O 7 = + H 2 O

A1+++ + OH- -* Al(OH),

A1(OH) 3 + S = -> no reaction

A1(OH) 3 + OH- -> A1O 2 - + H 2 O

Al(OH), + H+ - A1+++ + H 2 O

Zn ++ + OH" - Zn(OH) 2

Zn(OH) 2 + OH- - HZnOr + H 2 O

Zn(OH) 2 + ST - ZnS + OH~

ZnS + H + - Zn ++ + H 2 S

+ (NH 4 ) 2 Hg(SCN) 4 - (see page 67)
+ C 6 H 5 N (pyridine) + SON" - Zn(CH.N) 4 fSCN) s
(white)

Zn(OH), + NH 3 - Zn(NH 3 ) 4 ++ + OH~
Co++ + OH" -* Co(OH) 2
Co(OH) 2 + NH 3 - CoCNH,),^ + OH~
Co(NH 3 ) 6 ++ + S -> CoS (black) + NH 3
CoS + HC1 (1:9) - Very slow reaction
CoS + HNO 3 + HC1 - Co++ + Cl- + NO + H 2 O + S
Co++ + (NH 4 ) 2 Hg(SCN) 4 - (see page 67)
NI+++ OH- - Ni(OH),.
Ni(OH) 2 + NH S -> Ni(NH 3 ) 4 ++ + OH~

RKACTIONS INVOLVED IN SEPARATIONS 139

Ni(NH.,), ++ + S- - NiS (black) + NH 3
NiS + HC1 (1:9) - (Very slow reaction)
NiS + HC1 + HNO 3 - Ni++ + S + H 2 + NO + Cl~

GROUP IV

Ba++ + COr - BaGOs (white)
BaCO, + H+ -> Ba ++ + H 2 O + C0 3
Ba++ + CrO<- - BaCrO 4 (yellow)
BaCrO 4 + H+ - Ba++ + Cr 2 O 7 = + H 2 O
Sr 4 ^ (same reactions as Ba 4 " 1 ")
Sr++ + SOr - SrSO 4 (white)
Ca++ + CO 3 = - CaCO 3 (white)
CaC0 3 + H+ - Ca ++ + H 2 + CO 2
Ca ++ + CrO 4 = no reaction

+ SOr -* CaSO 4 (more soluble than SrSO 4 )

+ C,Or - CaC 2 4 (white)

GROUP V

Mg" 1 " 1 " + CO 3 = > MgCO 3 (prevented by large excess of

NH 4 + ion)
Mg++ + OH" - Mg(OH) 2 (soluble in excess NH 4 + ion

solution)

Mg ++ + NH 4 + + HPO 4 = - MgNH 4 PO 4 (white) + H+
K+ + H 2 C 4 H 4 6 - KHC 4 H 4 O 6 (white) + H+
K+ + Na 3 Co(N0 2 ) 6 - K 2 NaCo(NO 2 ) 6 + Na +
K + + Ag+ + Na 3 Co(NO 2 ) 6 - K 2 AgCo(N0 2 ) 6 + Na +
K + + H 2 PtCl 6 - K 2 PtCl 6 (yellow) + H+
NH 4 + + OH- -> NH 3 + H 2 O
NH 3 + H + -* NH 4 +
(Other reactions similar to those of potassium.)

QUESTIONS

1. Does a color reaction seem to show up best as a spot test or as an
ordinary beaker precipitation reaction?

2. How could you separate lead chloride from silver chloride?

3. In view of the fact that mercuric acetate, Hg ^2^0)2)2, is a very weak
electrolyte show by ionic equations why HgCl2 should dissolve readily in
concentrated solutions of sodium acetate.

4. Is 3N HNO 3 a practical solvent for HgS?

5. Why is it best to prepare fresh aqua regia each time when it is needed
as a solvent?

6. Comparing the results obtained from the addition of KI solution to
solutions of Hg 2 ++ and Hg++, how could you distinguish between these
two ions?

7. How do these results compare with similar tests on Hg 2 ~ 1 " 1 "?

8. Is CuS soluble in (NH 4 )S? HgS, CdS and Bi 2 S 3 act similarly.

9. Why is NH 4 OH used so much more than, for example, NaOH (two
reasons) ?

10. Why is it necessary to add the NHJ?

11. By means of ionic equations, show why lead sulfate dissolves in
saturated ammonium acetate solution. In terms of solubility products
explain why lead chromate precipitates from this solution though lead
sulfate does not.

12. Will FeS precipitate when H2S is added to a neutral solution of Fe + "*"?
Explain in terms of solubility product theory.

13. What background is best for a spot test for the following ions:

a. Pb+ + using K 2 CrO4 solution as the reagent?

b. Fe+ + + using KSCN as the reagent?

c. Fe"*" 4 ' 4 ' using K 4 Fe(CN)6 solution as the reagent?

14. Does this reaction suggest a way to separate Fe" 1 " 4 " 4 " from A1+++?
Explain.

15. For what two purposes is ammonium acetate usually used? For
what purpose is it used here?

16. Does this suggest a method of separating Zn ++ from Mn++? Explain.

17. Does it seem possible to precipitate any of the members of Group III
by passing H 2 S into a neutral solution of the metallic ion?

18. CoS is soluble in 1 :9 HC1. How do you explain its apparent lack of
solubility here? How could this be used to separate ZnS from CoS?

19. Does this test for nickel seem to be very sensitive, i.e., capable of
detecting very low concentrations of nickel?

20. By ionic equations and a few words explain the action of ammonium
benzoate here. Why is the solution diluted? Why is it heated (see page
27).

140

QUESTIONS 141

21. What is the action of the NH4C1 here? Explain by ionic equations
why it causes the aluminum hydroxide to precipitate. (HINT: See Note 42.)

22. What is the function of the ammonium sulfate in the wash water?
(HINT: See Note 49.)

23. Does any of the precipitate seem to have gone into solution?

24. From this (arid from your answer to Question 23) would you say that
BaCrO 4 is soluble in the buffered solution of acetic acid used?

25. How do these results compare with those from similar treatment of a
solution of Ba++? Of Ca++ (see Calcium ion Tests)?

26. Does a precipitate form readily? How does this compare with the
precipitation of BaCrC>4 from a similar acid solution?

27. How do your results compare with those obtained for barium?
How could you separate BaCrC>4 and SrCrO 4 ?

28. What is the function of the alcohol in these tests?

29. What can you say, from your results, as to the relative solubility of
these three chromates? Is this in accordance with what you would expect
from a comparison of their solubility products?

30. Explain the results obtained in this experiment.

31. What is the function of the ammonium chloride in this experiment?
Explain in terms of the solubility product principle and by means of ionic
equations.

32. Do these three tests suggest to you any means of detecting potassium
in the presence of sodium, by flame test?

33. Could you distinguish between potassium ions and ammonium ions by
means of any of these tests? Give your reasons.

34. How could you distinguish between potassium ions and ammonium
ions?

35. Can you suggest any reason why overheating at this point might ruin
the test for ammonium ions?

36. Will sodium interfere with the tests for potassium ions? Give your
reasons. Could you use the flame test to identify sodium in the presence
of potassium ions?

37. What is the purpose of this treatment with aqua regia and ignition?
How does the aqua regia act here?

PART IV
ANALYTICAL PROCEDURE ANIONS

In the analysis of solutions for their metallic constituents,
the separations and identifications have been based on the
use of solutions of certain anions as reagents. Conversely,
the analysis of anions is carried out by making use of their
characteristic reactions with cations, such as H + ions and
certain metallic ions. The laws and rules that apply to
cation analysis apply similarly to ariion analysis; and it-
should be easy for the student, with the knowledge gained
up to this time, to apply the technique learned earlier, to
anion analysis.

As in cation analysis, the anions are divided into groups.
Group I consists of those anions which form volatile com-
pounds when treated with HC1; Group II, those anions
which form compounds with AgNO 3 which are insoluble in
dilute nitric acid; Group III, those anions the calcium or
barium salts of which are insoluble ; and Group IV, those ani-
ons which do not form either volatile compounds with dilute
HC1 or insoluble compounds with barium, calcium or silver.

PRELIMINARY TREATMENT OF SOLIDS

As in the case of the cation analysis, most of the tests are
performed on solutions of the unknown. For this reason,
it is necessary to know how to dissolve any solid unknowns
that may be issued.

A study of the procedure used in anion analysis makes it
obvious that the procedure used in dissolving samples, in
preparation for the cation analysis, cannot be used here for
any but a few of the anions. For example, the treatment
with acids would result in the loss of the members of Anion
Group I, as volatile substances. Also, the use of HNO 3 ,
HC1, and Na 2 CO 3 would add NO 3 ~, Cl~, and CQ = to the
solution, and make it difficult to tell whether or not these

142

PRELIMINARY TREATMENT OF SOLIDS 143

ions were present in the original unknown. Addition of
some anions cannot be avoided, but it is possible to use the
different solvents in places where this causes no trouble.
For this reason, it is best to prepare the solid unknown,
using a special procedure for each group.
The best and simplest procedure is as follows :

1. Anion Group I. The tests for the members of Group I
may be performed, in most cases, on the solid unknown.
Thus, where the procedure calls for 1 to 2 drops of a solution
of the unknown, a portion of the pulverized solid about the
size of a grain of rice may be used.

This procedure is adequate for the detection of the
carbonates, sulfites, thiosulfates, nitrites, cyanides, and some
sulfides. However, the sulfides of a few metals are insoluble
in the HC1 used (see Group II Metals), and the sulfide test
may be missed if the sulfur is combined with these metals.
In case negative results are obtained when testing for sul-
fides using the solid, heat a sample of the solid unknown
(about the size of a match head) for 10 to 15 min. with a
strong solution of Na 2 CO 3 , and filter. The sulfide test
should then be repeated on both the solid residue (if any)
and on the filtrate.*

2. Anion Group II. For this, use a sample about the
size of an ordinary match head. To prepare the sample for
the Group II analysis, first dissolve as much of it in hot
water as is possible. Filter, and save the filtrate.

Heat the residue, on the steam bath, with saturated
Na 2 CO 3 solution for 10 to 15 min. Filter, discarding the
precipitate.

Combine the aqueous and Na 2 CO 3 filtrates and acidify
with dilute H 2 SO 4 .t Filter, if any precipitate forms, and
use the filtrate for the Group II tests.

* Many of the insoluble sulfides will react, to some extent, with strong,
hot solutions of NaaCOs, forming the metallic carbonate and some Na 2 S.
The latter is soluble in water and reacts readily with dilute HC1.

f Add the acid cautiously as otherwise the CO2 given off on acidifying the
Na 2 CO 8 will cause the solution to overflow the sides of the container.

144 ANALYTICAL PROCEDURE ANIONS

3. Anion Group III. Treat a small sample of the solid
unknown with hot water, then with saturated Na 2 CO 3
solution, as described in the preparation for Anion Group
II. * To the precipitate from the Na 2 C0 3 treatment add
6JV HNO 3 . Heat for 2 to 3 min. on the steam bath, then
filter, saving the filtrate. Repeat this treatment, using
concentrated HNO 3 this time. Filter, and discard any
residue. Finally, combine the filtrates from the treatments
with water, Na 2 CO 3 , 6N HNO 3 , and concentrated HN0 3 .
Be sure the solution is acid, then add AgNO 3 until precipi-
tation of Group II anions is complete. Filter, discarding
precipitate. Evaporate the filtrate to half its volume.
Cool, and neutralize with IN NaOH until the solution is
just faintly acid to litmus. Use this solution in the tests for
CrO 4 = , F~, C 4 H 4 6 == , AsO<r, AsOr, POr, and B 4 O 7 = .

4. Anion Group IV. Since the anions in this group form
soluble salts with all the common metals, those compounds
in the unknown that contain the members of this group
should be soluble in water. Therefore, treat the solid
sample with water, heat for 2 to 3 min., and filter. Treat
the filtrate in the manner described under the heading
Preliminary Treatment of Solutions, Anion Group IV.

PRELIMINARY TREATMENT OF SOLUTIONS

As in cation analysis, the members of one group of anions
may interfere with the tests for those of another group.
For this reason it is necessary to remove interfering ions
before starting the analysis. Since the different groups
need slightly different treatment, the procedure used in each

* The combined water and Na 2 CO 3 filtrates should be acidified with
dilute HNO 3 , filtered (the precipitate being discarded), and a portion of the
filtrate saved to use in testing for 804". Otherwise, if the unknown contains
Ba++, Sr ++ , or Pb++, a precipitate of the sulfates will form when the solutions
are combined, later. This might result in missing the sulfate test entirely.
Lead, barium, and strontium, if present, remain in the residue from the
Na 2 CO 3 treatment as carbonates. Later, however, they will be dissolved in
acid; and on combining filtrates, the sulfates would precipitate.

GROUP ANALYSIS 145

case will be discussed under the name of that group for
which the final solution is to be used (see Note 84).

1. Anion Group I. For the members of this group, no
preliminary treatment is necessary, as the tests are carried
out on the vapors escaping from the hot, acid solution.

2. Anion Group II. As the members of Groups III and
IV do not interfere with the tests for this group, only Group
I need be removed. To do this acidify 1 ml. of the solution
with 6N H 2 SO 4 and heat on the steam bath for about 5 min.
Cool, filter, and use the filtrate to test for the anions in
Group II.

3. Anion Group III. Acidify 1 ml. of the unknown solu-
tion with QN HNO 3 , heat just to boiling, and keep the solu-
tion hot on the steam bath for 4 to 5 min. Then add
AgNO 3 solution, drop by drop, until precipitation is com-
plete.* Filter, discarding the precipitate, and use the
filtrate in testing for the members of Group III.

4. Anion Group IV. Acidify 1 ml. of the unknown with
dilute H 2 SO 4 and evaporate to two-thirds its original
volume. Make the filtrate faintly basic with NaOH, and
add BaCl 2 *CaCl 2 reagent until precipitation is complete.!
Acidify, add AgNO 3 until precipitation is complete, and
filter, discarding the precipitate. Filter, and again discard
the precipitate. Use the filtrate in testing for C1O 3 ~ and
C 2 H 3 O 2 ~~". For the nitrate test, prepare as before, only use
a saturated solution of Ag 2 SO 4 or a little solid Ag 2 CO 3
instead of AgNO 3 in removing the anions of Group II.

GROUP ANALYSIS

The analysis of these solutions will be much simplified if
the student keeps in mind the characteristics of the various

* The treatment with AgNOa is unnecessary if the analysis has shown that
the anions in Group II are absent.

t The acidification and evaporation remove members of Group I; the
AgNOs precipitates the members of Group II; and the BaCU-CaCU reagent
in basic solution removes the members of Group III. If any one of these
groups has been shown by analysis to be absent, the procedure for removing
members of that particular group may be omitted.

146 ANALYTICAL PROCEDURE AN IONS

ions that may be in the solution. For example, strong
oxidizing ions, such as MnO 4 ~, would never be found in
solution with strong reducing ions, such as SO.-r, S", NO 2 ~,
or I~~, especially if the solution were acid. Similarly, acid
solutions will never contain both S = and SO 3 = , for these
react, in acid, to form free sulfur and water.

GROUP I ANIONS

This group may contain NO 2 ~, S = , S0 3 = , S 2 0-r, CO 3 = , or
CN~. Using the solution prepared as described for this
group under Preliminary Treatment, perform the following
tests :

1. Nitrite Ions, NO 2 ~. a. In the gas evolution apparatus
place 3 drops of the prepared solution. Moisten with
dilute H 2 SO 4 , and put the top of the apparatus in place.
In the testing chamber place a strip of moist starch-iodide
paper. Warm the solution and note the color of both
vapors and the starch-iodide paper.

Yellow or Brown Gas Giving a Blue Coloration

on the starch-iodide paper indicates that NO 2 ~~ are present.
b. To 2 drops of the solution add 1 drop 0.17V ferrous
ammonium sulfate and acidify with 6N H 2 SO 4 . Warm and
test vapors as before. If NO 2 ~" are present, the results will
be the same as in (a).

2. Sulfide Ions, S = . a. In the gas evolution apparatus,
place 2 drops of the prepared solution and acidify with dilute
HCL Test the gases in the usual manner with a piece of
filter paper moistened with lead acetate solution, warming,
if the test is not obtained at first.

Brown or Black Stain

on the filter paper shows 8 s " are present.

3. Sulfite, SOs" and Thiosulfate, S 2 OT Ions (see Note
85). a. Place 2 drops of the prepared solution in the gas
evolution apparatus and acidify with dilute HCL Use a

GROUP I AN IONS 147

piece of filter paper moistened with a mixture of 1 drop
0.05 per cent KMnO 4 and 1 drop QN H 2 SO 4 , as test paper.
If nothing occurs, warm the solution.

Bleaching of Spot

indicates either SOi* or S 2 O-r present (see Note 86).
If, on warming the solution and letting it stand, a yellow
precipitate of sulfur forms, S 2 O 3 = may be considered
present (see Note 87).

b. To one drop of the prepared solution add 1 drop of
50% AgNO 3 solution and warm almost to boiling.
Yellow or Orange PPT.

turning very Dark Brown on standing or heating, indicates
S 2 O-r present (see Notes 88 and 89).

4. Carbonate Ions, CO.r (see Note 90). a. In a micro-
beaker place 2 drops of the prepared solution, acidify with
dilute H 2 SO 4 and at once hold a drop of Ba(OH) 2 solution
(in the tip of a glass tube) inside the beaker. Warm the
solution on the steam bath, and observe the drop of
Ba(OH) 2 solution.

White Precipitate or Cloudiness

in the Ba(OH) 2 solution indicates COs"* present (see Note
91).

5. Cyanide Ions, CN . a. Place 2 drops of the prepared
solution in a microbeaker and acidify with dilute H 2 SO 4 .
Hold a very small drop of (NH 4 ) 2 S X solution in the end of a
glass tube and lower it inside the beaker to within a short
distance of the solution. Warm the solution gently for
30 to 60 sec. (do not boil). Withdraw the tube, then
holding the drop high above a small flame, evaporate it
barely to dryness. Dip the end of the tube in IN HC1 and
stir to dissolve the residue off the tube. To the solution so
obtained add 1 drop FeCl 3 solution.

Red Coloration

indicates that CN"" are present.

148 ANALYTICAL PROCEDURE AN IONS

b. If FeCCNV and Fe(CN) 6 " are absent, the following
test may be worked successfully :

To 1 drop of the solution add 1 drop FeCl 3 solution, make
alkaline with NaOH, and heat for 3 to 5 min. Acidify with
dilute HC1.

Dark Blue Color or Precipitate

indicates that CN~~ ions are present (see Note 92).

GROUP II ANIONS

This group may contain SCN~ Fe(CN)<r, Fe(CN) 6 ^,
Cl~~, Br~, and I~~. Using the solution prepared for this
group, test 1 drop with 1 drop AgNO 3 and 1 drop cone.
HN0 3 . A white precipitate indicates Group II present.
In such case, perform the following tests (see Note 93).

1. Thiocyanate Ions, SCN~. a. To 1 drop of the pre-
pared solution add 1 drop FeCl 3 solution.

Red Coloration

indicates that SCN~ ions are present (see Note 94).

6. To 1 drop of the prepared solution, add 1 drop Co ++
test solution and 3 drops amyl alcohol. Shake vigorously,
then examine the alcohol layer.

Blue Layer

indicates SON"" ions are present.

c. To 3 drops of the prepared solution, add a mixture of
1 drop Zn" 4 " 1 " test solution, 1 drop HgCl 2 solution, and 1 drop
of Co 4 " 1 " test solution. Rub the inside of the container with
the tip of a stirring rod and set aside.

Blue Precipitate

indicates SCN~ are present (see Note 95).

2. Ferricyanide Ions, Fe(CN) 6 ". a. To 1 drop of the
prepared solution add 1 drop ferrous ammonium sulfate
solution.

GROUP II ANIONS 149

Blue Color or Precipitate

indicates probable presence of Fe(CN) 6 s ions. A pale blue
precipitate slowly turning darker on exposure to air indi-
cates Fe(CN) 6 B instead of Fe(CN)<r.

b. To 1 drop of the prepared solution add 1 drop AgNOs
solution.

Orange Precipitate

indicates that Fe(CN) 6 = are present.

To the precipitate add 2 to 3 drops QN NH 4 OH. If the
'precipitate dissolves readily, it may be considered con-
firmation for the above, in the absence of CrO 4 = and
AsOr.

3. Ferrocyanide Ions, Fe(CN) 6 ~. a. To 1 drop of the
prepared solution add 1 drop FeCl 3 solution.

Deep Blue Color or Precipitate

indicates that Fe(CN) 6 " are present.

4. Iodide Ions, I~~. a. To 2 drops of the prepared solu-
tion add 1 drop bromine water. Add 4 to 5 drops of CC1 4
and shake vigorously.

Violet or Pink

layer of CC1 4 indicates I~ are present.

5. Bromide Ions, Br~ (see Note 96). a. To 2 drops of
the prepared solution add 2 drops H 2 O 2 and 5 drops CC1 4 ,
and shake the mixture vigorously.

Yellow or Brown Layer

of CC1 4 indicates that Br~ ions are present.

6. Chloride Ions, Cl"~. a. To 1 drop of the prepared
solution add 2 drops AgNO 3 solution. Filter, discarding
the filtrate, and wash the precipitate.

To the precipitate add 2 drops IN NH 4 OH. Stir once
and filter at once. To 1 drop of filtrate on a watch glass
with a black background add 1 drop 6N HN0 3 .

150 ANALYTICAL PROCEDURE AN IONS

White Precipitate
indicates that Cl~~ are present (see Note 97).

GROUP III ANIONS

This group consists of CrO 4 == , Cr 2 O 7 5== , SO 4 == , SiO 3 = , F~,
AsOa 5 *, AsO^, PO 4 S and B 4 O 7 = . To test for presence of
the group, add 2 drops IN NH 4 C1 to 2 drops of the prepared
solution and filter. To the filtrate add 2 drops QN NH 4 OH
and 1 drop BaClo-CaCl 2 reagent. A precipitate indicates
that Group III is present. In such case, using separate
portions of the solution prepared for this group, perform
the following tests :

1. Chromate, CrO 4 = , and Bichromate, Cr 2 O 7 == , Ions (see
Note 98). a. Extinguish any flames nearby. To 2 drops
of the prepared solution in a microbeaker add ether to form
a layer about 4 mm. deep. Then add 1 drop saturated
ammonium acetate solution, 1 drop 67V H 2 SO 4 , and 1 drop
H 2 O 2 . Place the thumb over the top of the microbeaker
and shake. Observe the ether and the aqueous solution.

Blue Ether Layer and /or Green Water Solution

indicates CrO 4 = are present.

6. In the absence of AsO 4 % perform the following test:
To one drop of the prepared solution on a spot plate, add
1 drop QN acetic acid, 2 drops saturated ammonium acetate
solution, and 1 drop 50 per cent AgNO 3 solution.

Red or Red -brown Precipitate

indicates CrOr" are probably present (see Note 99).

2. Sulfate Ions, SOA a. To 2 drops of the prepared
solution on a watch glass with a black background add 1
drop 6N HN0 3 and 1 drop BaCl 2 solution. Rub the glass
lightly with the tip of a stirring rod and let stand.

White Precipitate
indicates that SC^ are present.

GROUP III ANIONS 151

6. To confirm a test obtained in (a), filter off the white
precipitate, discarding the filtrate. Mix the precipitate
with twice its volume of Na 2 CO 3 , add enough water to make
a paste, and transfer to a charcoal block. Heat strongly
with the reducing flame of a blowpipe for 1 to 2 minutes.
Transfer the mixture to a freshly polished silver coin,
moisten with 1 to 2 drops of distilled water, and let stand
for 15 min. Wash the coin with distilled water.

Yellowish or Brown Spot

on the coin proves that SC^ are present (see Note 100).

3. Silicate Ions, SiO 3 = . a. Evaporate 5 drops of the
prepared solution to dryness. Scrape the residue loose
with the tip of a stirring rod, transfer to a small lead dish,
and mix with an equal volume of pure CaF 2 or Na2F 2 .
Place on the steam bath, add 1 drop concentrated H 2 SO 4 ,
and hold a drop of water in a tiny loop of platinum wire
close to the mixture. Warm the mixture for 2 to 3 min.
and observe the drop of water.

Turbidity

in the drop of water indicates that SiOr are present.

4. Fluoride Ions, F~~. a. Place 5 drops of the prepared
solution in a microbeaker, make it slightly alkaline with
dilute NaOH, and evaporate to dryness. Cool, then add
2 drops concentrated H 2 SO 4 to the residue. Warm gently,
holding a small drop of water in the tip of a glass tube
inside the beaker a few millimeters above the mixture.
Observe the drop of water.

Turbidity

in the water indicates F~ are present.

6. Confirm any test obtained in (a) by washing and
drying both the glass tube and the microbeaker and
examining the surfaces of both.

152 ANALYTICAL PROCEDURE A NIG NX

Frosted or Etched

appearance inside the beaker or on the end of the glass tube
confirms the presence of F~.

5. Arsenite, AsO 3 ^ and Arsenate, AsO 4 % Ions. a. In
the absence of other metals of Group IT, perform the
following test :

To 3 drops of the prepared solution add dilute HC1 until
all the Ag+ added in the preliminary treatment are pre-
cipitated as AgCl. Filter, discarding the precipitate, and
evaporate the filtrate barely to dryness. Dissolve the
residue in 0.3 N HC1 and pass in H 2 S for 30 sec.

Yellow Precipitate
forming immediately indicates that AsO : r are present.

If no immediate precipitation takes place, add 2 to 3
crystals of NHJ, heat almost to boiling, and saturate with
H 2 S.

Yellow Precipitate

indicates the presence of AsO 4 ^ (see Note 101).

6. To 2 drops of the prepared solution, add NH 4 C1 solu-
tion until all the Ag + are precipitated as AgCl. Filter,
discarding precipitate. To the filtrate, add 1 drop Cu 4 " 1 "
test solution and an excess of 6N NaOH, and boil.

Blue Solution forming a Yellow or Red Precipitate of Cu 2 O,
on boiling, indicates that AsO 3 ^ are present (see Note 102).

A bluish-green precipitate obtained by the above treat-
ment, insoluble in excess NaOH and giving no yellow or red
precipitate, indicates the probable presence of AsO 4 s and
the absence of AsO 3 =.

c. In the absence of CrO 4 s= or PO 4 S , use the following
tests:

GROUP III AN IONS 153

To 1 drop of the prepared solution add 1 drop AgNOs
and 2 drops saturated ammonium acetate.

Curdy Yellow Precipitate
indicates that AsOs^ are present, and AsO 4 s are absent.

Chocolate-colored Precipitate

indicates that AsO 4 - is present. AsOs 35 may be present or
absent.

6. Phosphate Ions, PO 4 % If arsenites and arsenates are
present, they should be removed before performing the
following tests. The procedure is the same as that used in
precipitating the sulfides of the Group II metals (which see).
The filtrate from the H 2 S precipitation (or the prepared
solution, if AsO : r and AsOi** are absent) is used as follows:

a. Make 2 drops of the solution slightly basic with QN
NHjOH (see Note 103). Add 1 drop magnesia mixture,
rub the inside of the container with a stirring rod, and set
aside for about 15 min.

White Crystalline Precipitate

indicates that PO 4 ^ are present.

6. To 2 drops of the solution add NH 4 C1 and filter to
remove Ag + . To the filtrate add 1 drop QN HNOs and

1 drop ammonium molybdate reagent. Warm on the
steam bath 2 to 3 min. and set aside for about 15 min.

Yellow Precipitate

indicates PO 4 - are present.

c. In the absence of CrO 4 = (and AsO 3 ^ and AsO 4 = ), use
the following test :

To 1 drop of the solution add 1 drop AgNO 3 solution and

2 drops saturated ammonium acetate solution.

Yellow Precipitate
indicates PO 4 ^ are present.

154 ANALYTICAL PROCEDURE AN IONS

7. Borate, BO 3% and Tetraborate, B 4 O7 = , Ions (see

Note 104). a. Place 3 drops of the prepared solution on a
Pyrex watch glass or a crucible lid, add 1 drop QN NaOH,
and evaporate just to dryness. Cool, add 1 drop of con-
centrated H 2 SO 4 and 3 to 4 drops methyl alcohol. Stir,
then ignite the alcohol.

Yellow-green Tinge

to the flame indicates these ions are present (see Note 105).
b. To 1 drop turmeric solution on a crucible lid add 1 drop
IN HC1 and 1 drop of the prepared solution. Evaporate
to dryness, very cautiously.

Red-brown Spot
indicates the presence of these ions.

GROUP IV ANIONS

This group consists of NO 3 ~, C 2 H 3 O 2 ~, C1O", and MnO 4 ~
ions. The tests for the members of this group should be
made on the special solutions prepared for the ions to be
identified.

1. Nitrate Ions, NO 3 "". a. To 5 to 10 drops of the prepared
solution in a microbeaker add 5 drops ferrous ammonium
sulfate and stir. Tilt the tube at a 45-degree angle and by
means of a medicine dropper let 5 to 6 drops of concentrated
H 2 SO 4 run down the side of the beaker to form a layer under
the solution.

Reddish or Brown Ring

at the interface between the two liquids indicates that NO 3 ~~
are present (see Note 106).

b. In the absence of ferric, chlorate, and other oxidizing
ions the following test (due to Lunge) may be used (65) :

Place 1 drop of the prepared solution on a glass slide or
spot plate. Beside it, place 1 drop of diphenylamine
reagent, and, with a stirring rod, draw the solution into
contact with the reagent.

GROUP IV AN IONS 155

Blue Color

at the interface between the liquids indicates that N0 3 ~
are present.

2. Acetate Ions, C 2 H 3 O 2 ~~. a. To 2 drops of the prepared
solution in a 4-in. test tube add 1 drop ethyl alcohol and
1 drop dilute H 2 SO 4 . Plug the top of the test tube with
cotton, place the test tube on the steam bath, and warm
4 to 5 min. Remove the cotton plug and smell the vapors
coming from the mixture.

Fruity Odor

indicates C 2 H 3 O 2 ~ are present. (Caution: It is best to
run a blank reaction at the same time. Otherwise, the
odor of the alcohol may be misleading.)

6. To 1 drop of the prepared solution on a spot plate add
1 drop saturated iodine solution and 1 drop 5 per cent
lanthanum nitrate solution. Let stand for 1 to 2 min.,
then add 1 drop IN NH 4 OH.

Bluish Coloration

indicates C 2 H 3 2 ~~ (see Note 107).

3. Chlorate Ions, C1O 3 ~. a. Place 3 drops of the pre-
pared solution in a microbeaker, make it strongly alkaline
with NaOH solution, and add 3 to 4 granules of aluminum
to the solution. Place the beaker on the steam bath and
heat for 5 to 10 min. Cool, acidify with dilute HNO 3 , and
add 1 drop AgNO 3 solution.

White Precipitate

indicates that C10 3 ~~ is present (see Note 108).

4. Permanganate Ions (MnO 4 ~). a. If manganese was
found during the cation analysis of this unknown, a violet
color in the solution indicates that MnO 4 ~ are present. A
colorless solution indicates that MnO 4 ~ are absent.

156 ANALYTICAL PROCEDURE ANIONS

b. To 3 drops of the prepared solution (if it is violet in
color) add 1 drop dilute H 2 SO 4 . Then add, drop by drop,
some 3 per cent H 2 O 2 .

Fading of Violet Color

with evolution of oxygen, confirms the presence of MnO 4 ~
(Note 109).

EQUATIONS FOR REACTIONS OF ANIONS

Anion Group I :

la. NOr + H + * HNO 3 + NOT (colorless)

NO + O 2 ^ NO 2 | (brown)

N0 2 + I" + H + - I, (blue, with starch) + H,O +

NOT
6. Fe ++ + SO- + NO 2 - + H + -> Fe(NO)S0 4 (brown) +

H 2 O

Fe(NO)SO 4 + SO 4 = ^ Fe 2 (S0 4 ) 3 + NOT
2a. S- + H+ ^ H,S

H 2 S + Pb(C,H,O,), -* PbS (bla(;k)I + H-C,H 3 O,
3o. SO 3 = + H + ^ H 2 S0 3 ? H 2 O + SO 2 |

S 2 O 3 = + H + ? H s SiO. ^ H 2 + Si + SO,T

MnOr (violet) + SO 3 = + H+ - Mn++ (colorless) +

H 2 o + sor

6. Ag+ + S 2 O 3 Ag 2 S 2 3 |

Ag 2 S 2 3 - A g2 S( I ) +S0,(|)
4o. CO 3 = + H + ? H.CO, ^ H 2 O + C0 2 T
COr + Ba++ -* BaCOsi

60. CN- + H + -* HCNt

CN- + 8,- - SCN- + S"
6. CN- + Fe ++ - Fe(CN) 6 =

Fe(CN) 4 B + Fe +++ - Fe 4 (Fe(CN) 6 ) 3 i (dark blue)
Anion Group II :

lo. SCN- + Fe +++ - Fe(SCN) 3 (deep red)
6. SCN- + Co ++ -> Co(SCN) 2 (blue)
c. (See equations for Co++ ion reactions)
2o. Fe(CN),- + Fe 4 ^ - Fe 3 (Fe(CN) 6 ) 2 i (blue)

EQUATIONS FOR REACTIONS OF ANIONS 157

6. Fe(CN),- + Ag+ - Ag 3 Fe(CN) 6 J (orange)

Ha. Fe(CN) 6 - + FC+++ -* Fe 4 (Fe(CN),U (dark blue)

4a. T + Br, -> I 2 + Br~

b. Cu 4+ + I- -* Cul,

CuI 2 - Cu 2 I 2 | (white) + I 2 (violet in CC1 4 )
5. Br- + H 2 2 + H+ - Br 2 (yellow in CC1 4 ) + H 2 O
60. Cl- + Ag+ - AgCU (white, soluble in NH 4 OH)

Br~ + Ag + -> AgBrl (yellowish, insoluble in NH 4 OH)

I- + Ag+ -> Agll (yellowish, insoluble in NH 4 OH)
Anion Group III :

1ft. CrOr + Ag + -^ Ag 2 CrO 4 | (red-brown)
2a. SO 4 = + Ba ++ -> BaSO 4 i (white)
ft. BaSO 4 + NasCOj - BaCO 3 l + Na 2 SO 4

Na,SO + C - Na 2 S + COT

Na,S + Ag + H+ + O, - Na+ + H 2 O + Ag 2 S| (brown

or black)

3. Na 2 F 2 + H,SO 4 - H 2 F, + Na 2 SO 4
H,F, + SiOr -* SiF 4 f + H 2 O + F~
SiF 4 + H 2 - H,SiF + H 4 SiO 4 | (white)

4. [Same as (3).]

5a. AsO 3 3 + H 2 S - As 2 S 3 l (yellow) + H 2 O
AsO 4 = + H 2 S - As 2 S 6 | (yellow) + H,O
As0 4 - + I" + H+ - AS+++ + H 2 O + I,
6. Cu(OH), + AsO 3 - - Cu 2 Ol (red) + AsO 4 = + H 2 0.
Cu(OH) 2 + AsO 4 - -*Cu 3 (AsO 4 ) 2 i (bluish gre,en) +

c. AsO 3 D + Ag + - Ag 3 AsO 3 J (yellow)

AsO 4 = + Ag + > Ag 3 AsO 4 | (chocolate brown)
60. P0 4 = + Mg++ + NH 4 + - MgNH 4 PO 4 | (white)
b. POr + NH 4 + + MoO 4 = + H+ - (NH 4 ) 3 PO 4 -
12MoO,| (yellow) + H 2

PROBLEMS

1. What weight of each of the following substances would be required
to prepare 100 ml. of a QAAf solution of each? What weight would be
required to prepare 100 nil. QAN solution?

(a) NaCl (d) Fe 2 (SO 4 ),i

(6) KOH (e) HNO 3

2. If 25 ml. of Q.2N NaOH solution is required to neutralize 40 ml. of a
solution of H 2 SC>4, what is the normality of the acid? What is its molarity?

3. A solution is prepared by diluting 25 rnl. of an acid to a volume of
100 ml. Thirty milliliters of this diluted acid will neutralize 75 ml. of Q.SN
KOH solution. What is the normality of the original acid?

4. A solution is prepared by adding 10 ml. concentrated H 2 SC>4 to 50 ml.
water and diluting the resulting solution to a total volume of 180 ml. On
trial, it is found that 50 ml. of this solution will neutralize 200 ml. of Q.5N
NaOH solution, (a). What was the normality of the concentrated H 2 SO4?
(6). What was its molar concentration? (c). What weight of H 2 SO4 is
there in 1 liter of the concentrated H 2 SO4?

6. A certain solution of H 2 SO4 has a specific gravity of 1.5263 and is
62.18 per cent H 2 SO4 by weight. What is the (a) molar and (6) normal
concentration of this acid?

6. A solution of CuSOi, having a specific gravity of 1.206 is 20 per
cent CuSOi by weight. What is its (a) molar and (6) normal concentra-
tion?

7. What is the maximum equilibrium concentration of Ag+ ions that can
exist in a solution containing 0.01 mol per liter of I~ ions?

8. What is the maximum concentration of CO 3 "" that can be added to
0.1M AgNO 3 solution without causing Ag 2 CO 3 to precipitate?

9. Silver nitrate solution is added slowly, drop by drop, to a solution
which is QAM with K 2 CrO 4 and 0.01M with NaCl. What will be the
concentration of the first ion, at the moment the second ion first begins to
precipitate?

10. A solution is Q.1M with Pb(NO 3 ) 2 , 0.01A7 with AgNO 3 , 0.1M with
Hg(NO 3 ) 2 , 0.1M with Fe(NO 3 ) 2 , and 0.1M with Mri(NO 3 ) 2 . If H 2 S is
passed into the solution until precipitation is complete and until the sulfide
ion concentration in the solution at the end is 10~ 22 mol per liter, which of
the metallic sulfides will have precipitated? What will be the concentration
of each metallic ion in the final solution?

11. From calculations using the solubility product constants of Agl,
AgBr, Ag 2 <X>3, Ag 2 CrO4, and AgCN, list these compounds in the order of

159

160 PROBLEMS

decreasing concentrations of Ag+ needed to precipitate these compounds
from 0.1M solutions of I~, Br~, CO 3 ", CrO 4 ~, and CN~.

12. The solubility of CaCrO 4 is 14.1 g. per liter. Calculate its solubility
product constant.

13. The solubility of SrC 2 O 4 is 0.05 g. per liter. Calculate its solubility
product constant.

14. From the solubility product constants of the following substances,
calculate their solubilities in (a) mols per liter; (b) grams per liter; (c) grams
per 100 cc.

(1) BaCCX (3) PbS0 4

(2) CaF 2 (4) Ag a S

15. The solubility of Ca 3 (PO 4 )2 is 0.02 g. per liter. What is its solubility
product constant?

16. The solubility of CaHPO 4 is 0.2 g. per liter. What is its solubility
product constant?

17. From the solubility product constants of CMS and ZnS calculate tho
efficiency with which Cd ++ and Zn + ^ may be separated using II. S sis (he
reagent.

18. In the 0.37V HC1 solution used during the precipitation of the Group II
metals it is possible for the H 2 S to furnish S~ in concentrations as high as
10~ 21 molar. The usual unknown in Group TT contains each ion at a con-
centration of 0.02 mol per liter. With these points in mind, what v\ould
be the concentration of the following ions left in the filtrate at the end of the
Group II precipitation?

(a) Cu ++ (b) CdS (/C. p , or GIB = 10" 2 ")

19. Show by calculations that it is possible to precipitate ZnS in the
second group if the concentration of Zn M in the solution is greater than
1.2 X 10~ 2 mol per liter.

20. What is the hydrogen ion concentration in a 0.1J7 solution of HCN?
Of O.OlAf HCN?

21. What is the hydrogen ion concentration in a 0.1 M solution of
HC 2 H 3 O 2 ? Of O.OU7 H-C 2 H 8 O 2 ? Of O.OOOOOU/ H-C%H 3 O>? (HINT:
Remember Arrhenius' theory.)

22. What is the hydroxyl ion concentration in QAAf NH 4 OH? In
O.OOlAf NH 4 OH?

23. What is the hydrogen ion concentration in each of the following
solutions :

(a) 0.1 A/ HC1? (d) O.IM KOH?

(6) O.OlAf HNO 3 ? (f) 0.1A/ NH 4 OH?

24. What is the hydrogen ion concentration in a solution that is 0.1 AT
with NaC 2 H 3 O 2 and 0.01A/ with H C 2 H 3 O 2 ? One that is 0.1M with
NaC 2 HO 2 and 0.1M with HC 2 H 3 O2?

PROBLEMS 161

25. What is the hydroxyl ion concentration in each of the following
solutions :

(a) A solution 0.1 A/ with NH 4 C1 and 0.01 A/ with NH 4 OH?
(ft) A solution O.lAf with NH 4 C1 and 0.1A/ with NH 4 OH?

26. What would be the hydroxyl ion concentration in a solution made by
mixing equal volumes of the solution in Problem 25 (a) and distilled water?
Work the same problem for Problem 25 (6).

27. At the beginning of the Group III precipitation, the unknown is
made strongly basic with NH 4 OH and then NH 4 C1 is added. Assuming
that the resulting solution is 2A7 with NH 4 C1 and I At with NH 4 OH what
is the hydrogen ion concentration of the solution? If the K 8P . n,s = 10" M ,
what is the maximum S" concentration, in this solution, obtainable by
adding H a S to the solution?

t 28. Using the answer obtained in Problem 27, calculate the concentration
of Fe ++ ion, of Mn M and of Zn^ left in the filtrate after the Group III
precipitation has been completed.

29. Show by calculation that H 2 S will precipitate Cu" 4 " 1 " from a solution
0.1 A/ with Ou(NO) and \M with HOI.

30. What is the hydrogen ion concentration in a 0.2/17 solution of NH 4 Ac
to which has been added an equal volume of

(a) 0.02717 HC1? (c) 0.02A/ NaOH?

(6) 18Af HOI? (d) 0.18A7 NaOH?

From those results could you call NH 4 Ao a universal buffering agent?

31. A solution is made by mixing equal volumes of 0.22A7 acetic acid and
0.20^7 NH 4 OH What is the hydrogen ion concentration of the resulting
solution? What is its hydroxyl ion concentration?

32. What is the pH of each of the following solutions:

(a) 0.1 A7 HOI? (c) 0.003A7 NH 4 OH?

(b) 0.5A/ NaOH? (d) A solution that is made 0.1A/

with NH 4 Cl and 0.01M with
KOH?

33. What is the electrode potential of a combination of a normal man-
ganese half-cell and a normal antimony half -cell?

34. If a cell is made up of two half-cells which contain the same elements,
the concentration of the ions in the two being different, the electromotive
force of the combination is given by the expression

,-, , 0.059, Ci
Krn.f. = ---log^

where n is the valence of the ion involved, and C\ and Ct are the molar
concentrations of the ions in the respective half-cells.

Using this expression calculate the electromotive force of the following
combinations:

162 PROBLEMS

(a) 0.1M Cu ++ against 0.00001M Cu + +.

(b) O.U/ Ag< against 0.00001M Ag+

35. Using the simplified expression by Bronsted and La Mer, calculate
the ionic strength /u of

(a) 0.01A/ NaCJl. (6) 0.001A/ CuCl 2 .

NOTES ON ANALYTICAL PROCEDURE

1. If the unknown is an alloy the water treatment may be omitted here.
However, if it is necessary to carry the alloy through a fusion, the fused
substance should be tested with water as instructed.

2. Nitric acid is used before hydrochloric acid, owing to the fact that, if
any of the metals of Group I are present, they will cover the sample with an
insoluble coating of the metallic chloride and render its solution difficult.

3. Aqua regia is sometimes prepared by mixing one volume of concen-
trated nitric acid with three volumes of concentrated HC1. It must always
be prepared just before using as it decomposes very quickly on standing.

4. If the unknown is an alloy containing tin, a white residue of meta-stan-
nic acid is often left at this point. This should be treated with hot concen-
trated HC1 for a few minutes, then filtered and the residue treated with
water, the filtrates being combined with the other filtrates (see Amphoteric
Hydroxides).

A mineral should receive the above treatment at this point, in any case,
as any white residue may appear colored owing to the presence of dark-
colored substances in the mixture.

6. If the analysis is to be made for silicon and sodium, it is necessary
to use a crucible made of some other material, as part of the crucible itself
reacts with the fusion mixture, and the crucible contains these materials.
Iron or, even better, nickel crucibles can be used in such cases. If a com-
plete analysis is desired it is best to conduct two fusions, using one kind of
crucible for the first and another for the second. Those elements, found in
both cases, may be assumed to be present in the unknown sample.

6. The filtrate from the aqueous extraction of the fusion mixture should
be acidified with dilute HNO 3 , heated 2 to 3 min. and filtered. The filtrate
should be combined with the other filtrates as before. The precipitate is
silicic acid, its source being either the unknown, or, if a porcelain crucible
was used for the fusion, the glazed surface of the crucible. In the latter
case, the precipitate should be discarded.

7. When the filtrates are combined, a white precipitate often forms
especially when silver, lead, bismuth, mercury, or antimony is present.
This precipitate may be due to either of two things: (a) If bismuth or anti-
mony are present, their salts may have undergone hydrolysis to form
insoluble basic salts of the type Bi(OH) 2 (NO 8 ); (&) if the other three metals
are present, their chlorides will precipitate when the chlorides from the
aqua regia or hydrochloric acid treatments are mixed with the nitric acid
solution containing these metallic ions.

It is necessary, in such cases, to make the solution strongly acid with
concentrated hydrochloric acid. This will dissolve the hydrolytic products

163

164 NOTES ON ANALYTICAL PROCEDURE

of bismuth and antimony and part or all of the lead chloride (owing to the
formation of complex ions of the type PbCh" and PbCU""). The precipitate
left should be used as the precipitate for the Group I analysis. Any lead or
mercury that redissolves will be found in Group II.

8. When nitric and hydrochloric acids are mixed, as in aqua regia, some
reaction occurs forming a mixture of chlorine, and nitrosyl chloride (NOC1),
which is unstable. In dissolving substances such as HgS, therefore, two
reactions are responsible for the great ability of aqua regia to put insoluble
substances into solution : (a) The nitrate ions present, as well as the nitrosyl
chloride, act as a powerful oxidizing agent, oxidizing sulfide ion to free
sulfur and to sulfate ion; (6) chloride ions may combine with many metallic
ions to form complex ions of the type HgCU".

9. After precipitation seems complete scratch the inside of the beaker
with a stirring rod and let stand 3 to 5 min. PbCl 2 is very slow to precipitate
and may be overlooked in the group if not given plenty of time.

10. The absence of a precipitate at this point means that the metals of
Group I, with the possible exception of lead, which can be found in Group II,
are absent.

11. Lead chloride is quite soluble in hot water but reprecipitates at once
if the solution is allowed to cool. The small volume of liquid used here
cools very rapidly, and the test for lead is often lost through failure to keep
the solution hot during the whole procedure. It is best to preheat the
medicine dropper, before using it for the filtration or for transferring the
solution, by placing its tip in a microbeaker full of hot water on the steam
bath for a few seconds. Also the filtrate should be kept hot on the steam
bath during the running of the tests for lead.

12. The gallic acid test works only with absolutely fresh gallic acid solu-
tion, 5 minutes standing being sufficient to ruin the reagent. Acid solutions
render the test less sensitive, while alkaline solutions oxidize in air imme-
diately, to give a color similar to that produced by mercury. It is necessary,
therefore, to have the solution faintly acid a condition best reached by
buffering an acid solution with sodium acetate.

13. Stannous chloride, SnClz, is a strong reducing agent. It reduces the
mercuric chloride (from the aqua regia treatment) in two steps, forming
mercurous chloride, a white insoluble substance, and hydrated stannic
chloride. More stannous chloride reduces the mercurous chloride to black
metallic mercury:

2HgCl 2 + SiiCl* -> Hg 2 Cl 2 + SnCU
Hg 2 Cl 2 + SnCl 2 - 2Hg + SnCU

14. When cyanide ion is added to a solution containing cupric ion, two
successive reactions occur. First, the ions combine to form cupric cyanide
which is unstable and immediately decomposes to form insoluble cuprous
cyanide and cyanogen, a very poisonous gas. The cuprous cyanide then
reacts with more cyanide ions to form the highly soluble, extremely stable
cuprocyanide ion. The reactions are

NOTES ON ANALYTICAL PROCEDURE 165

Cu + + + 2CN- -* Cu(CN).
2Cu(CN) -* Cu 2 (CN) 2 -h C.N.

Cu 2 (CN) 2 + 4CN- -* 2Cu(CN),-

As cuprocyanide ions are very stable, their solutions contain few copper ions.
Consequently, hydrogen sulfide will not precipitate the sulfides of copper
from solutions of cuprocyanide ions in the presence of excess cyanide ions.

15. Sodium stannite is another example of a strong reducing agent. In
this case it reduces the bismuth hydroxide to metallic bismuth (black).
The solution is made by adding 6N NaOH, drop by drop, to 2 drops stan-
nous chloride solution (stirring during the addition) until the precipitate
of stannous hydroxide that forms just redissolves. The reactions involved
in this preparation are similar to those given by zinc ions under similar
conditions, stannous ions, also, being amphoteric.

When used as a reducing agent, stannite ions, SnO 2 ", are oxidized to
stannate ions, SnO 3 ".

1 16. In the analysis of Group II, bismuth hydroxide is precipitated in this
same manner in the presence of the blue copper ammonia complex. The
deep color of the latter and the transparent nature of the former make it
difficult to see the bismuth hydroxide in the mixture. Consequently, the
student will be saved much trouble if he notes the appearance of this mixture
very carefully and keeps it in mind when testing for bismuth.

17. If the reagent is allowed to remain in this alkaline solution, it will
slowly deposit a yellow precipitate even in the absence of Cd ++ ions. This
takes a few minutes, however, and as the coloration appears quickly in the
presence of cadmium, this phenomenon need not cause trouble. In case of
doubt, however, it is best to run a blank test for comparison.

18. The precipitation of As 2 Ss from solutions of arsenic ions, As +++++ t
by means of hydrogen sulfide, is very slow except in hot, strongly acid
solutions. High acidity is prohibited in this separation, as the sulfides of
some of the other Group II metals would not precipitate in highly acid
solutions. Therefore, use is made of the fact that arsenious sulfide is
easily precipitated from solutions of arsenious ion, arsenic ions being reduced
by NHJ to the arsenious state and then the H 2 S being added.

The reduction is preceded by a preliminary precipitation with hydrogen
sulfide, however, as it is essential that the tin present be precipitated before
reduction. As stannous sulfide is not soluble in ammonium monosulfide
(NH4) 2 8, any tin that is present as stannous ions will be overlooked.

19. Insoluble As 2 S 3 and SnS 2 react with the sulfide ions from ammonium
monosulfide to form the complex ions AsSs w , SbSa^ and SnSa". These
ions decompose if weakly acidified, to form the original sulfides and hydrogen
sulfide.

Ammonium monosulfide, (NH 4 ) 2 S, should be prepared fresh each labora-
tory period, as it quickly oxidizes to ammonium polysulfide, (NH 4 ) 2 S,, a
reagent which dissolves not only the above sulfides but also, to some extent,
the sulfides of mercury, copper and bismuth. The reagent is prepared by

166 NOTES ON ANALYTICAL PROCEDURE

passing a rapid stream of hydrogen sulfide through 3 ml. 6N NH 4 OH for
about 3 min.

20. The product of the evaporation of As 2 S 3 with aqua regia will be arsenic
oxide which is produced by the hydrolysis of the AsCU first formed and the
loss of hydrogen chloride during evaporation. As As 2 O 6 is somewhat
volatile, prolonged or intensive heating at this point should be avoided.

21. The experiment described here is a modification of the Gutzeit test
for arsenic. It depends upon the reduction of the arsenic, by means of the
sodium hydroxide and aluminum, to arsine, AsH 8 , a gas which reacts with
the silver nitrate giving a brownish stain of silver, arsenic, and complexes
of these elements. The test works best if, before the sodium hydroxide is
added, the residue of arsenic pentoxide is reduced to the arsenious state by
the addition of HC1 and sodium bisulfite, NaHSOs, solution.

22. This is another example of a hydrolysis reaction (see page 26 and
Note 7). It is reversed if a strong acid is added to the mixture.

23. The change of color of the cacotheline is due to its reduction by the
stannous ion to a new dye. The success of this test depends upon the
successful reduction of the stannic ion to the stannous state by the treatment
with acid and aluminum or with zinc.

24. An unusually successful way to carry out the reduction of the Sn+ +++
ions is as follows:

In a 4 by ^-in. test tube put the solution to be tested. Make it strongly
acid with HC1 and add 2 to 3 granules of zinc or aluminum. In the top
of the tube, suspend a 3 by 3 -in. test tube, two-thirds full of cold water;
then heat the solution in the larger test tube very gently, keeping it as hot as
is possible without causing the mixture to boil over. When the metal is all
dissolved, use the solution for the usual tests.

The advantage of this method is that it prevents, by condensation, the
escape of HC1 and water, and it prevents the entrance of air into the reaction
mixture. The reduction proves to be more dependable using this method.

26. It is essential that the residue from this evaporation be not overheated.
Certain metallic salts, such as those of iron, hydrolyze readily and, in
evaporations such as this, are converted to the oxides. Many of these
oxides are changed, by heating, into forms which are only slowly soluble in
Q.3N HC1 and thus may be found mixed with the Group II precipitate.
This results in the tests for members of Group II being spoiled and in failure
to find certain metals in later groups. Certain ions in the second group,
such as those of bismuth, antimony and tin, remain undissolved in the
0.3W HC1; but if the evaporation is carefully accomplished, this residue
need not worry the student. It is only necessary to continue the treatment
with hydrogen sulfide until all the original residue has changed color
becoming either black, orange, or yellow.

26. This process will give a solution whose acidity is approximately that
of 0.3JV HC1, all excess acid being lost during the evaporation.

27. The sulfides of certain metallic ions (arsenic, cobalt and nickel are
notable examples) tend to peptize readily, i.e., go from a precipitate to a
colloidal form. Others, of only moderately low solubility (such as cadmium),

NOTES ON ANALYTICAL PROCEDURE 167

partly redissolve in water. This wash solution opposes both these tenden-
cies, the ammonium nitrate neutralizing the charges on the colloidal particles
(see page 34) and the sulfide ions from the hydrogen sulfide preventing
much of the metallic sulfide going into solution (see page 32). The
NH 4 NO 3 -H 2 S solution should be prepared fresh by passing H 2 S through
1 ml. NH4NOg solution for 1 min.

28. Much information can be gleaned by observing the appearance of the
mixture during various stages of the precipitation of the Group II sulfidea.
No precipitate, of course, means that these metals are all absent. A yellow
precipitate means that lead, copper, bismuth, and mercury are absent, as
their sulfides are black. A heavy yellow precipitate, forming immediately
upon adding H 2 S to the solution after treatment with NH J, indicates the
probable presence of As + +~ H+ ~ h .

These indications cannot be taken as tests for these ions; but their appear-
ance often serves as a warning of something having gone wrong, if the test
for the element does not appear later.

29. In case a large quantity of precipitate is found in Subgroup II B t the
treatment of the precipitate with ammonium mo no sulfide should be repeated
until all members of the tin group are dissolved out.

30. In general, the solubilities of salts in alcohol or mixtures of alcohol
and water are lower than in water itself. Thus, by adding alcohol, redis-
solving of the PbSO 4 may be prevented. Sometimes a hydrolysis product
of bismuth sulfate precipitates out at this point. However, if the solution
is allowed to stand only 2 min., the bismuth salt will not come down as it
supersaturates readily.

31. If the precipitate is light in color or remains floating on top the
liquid, it is most likely sulfur, and the tests for mercury may be omitted.

32. It is essential that a moderate excess of NH4OH he added here, and
that the solution be stirred thoroughly as Cu(OH) 2 may otherwise not
redissolve, and both the copper and bismuth may be overlooked.

33. Whether copper is present or not, the CdS precipitate may be colored
black due to traces of lead or bismuth which have escaped removal. Conse-
quently, the following procedure is usually more successful in detecting
cadmium. Proceed as follows:

Take the remainder of the ammoniacal solution from the bismuth and
copper tests and evaporate barely to dryness. Dissolve the residue in
4 drops 3N HC1 and pass in H^S until precipitation of any dark colored
sulfides is complete. Centrifuge, using a cotton plug, and discard the
precipitate. To the filtrate add 2 drops 6./V NHiOH and 1 drop ammonium
acetate and pass in more H 2 S. A YELLOW PRECIPITATE indicates
that Cd + + are present. Confirm as directed in the original procedure using
thiosinamine.

34. If the FeS type of generator is used, the outlet tube should be thor-
oughly cleaned and fresh cotton should be placed in the filter bulb of the
delivery tube each laboratory period. Otherwise, spray may carry traces
of iron into the unknown, and, since the tests for iron are very sensitive, it
may be found, even though the unknown originally contained none. It

168 NOTES ON ANALYTICAL PROCEDURE

may he necessary to use a wash bottle with this apparatus if contamination
appears.

35. If, on acidification, traces of precipitate appear at first hut dissolve
later, it is prohahle that too much acid has been added. In that case, the
solution should he tested with litmus and, by addition of 17V NH 4 OH,
brought to a weakly alkaline condition. One normal HC1 should again be
added (with stirring) until the solution is very faintly acidic to litmus.

36. If the precipitate refuses to settle or is very pale yellow (almost white),
it is probably sulfur and the members of the tin group may be considered
absent.

37. In carrying out this modified Gutzeit test for arsenic, it is necessary
to avoid direct sunlight, H 2 S in the air, or NH 8 fumes. Any of these will
darken the AgNOs spot giving it a color similar to that obtained in the
presence of arsenic.

38. A yellow precipitate forming after some time, here, may be dis-
regarded. It is probably stannic sulfide, which will precipitate under these
conditions if present in large concentration.

39. It is essential that all the aluminum be removed from the solution,
either by letting the excess metal dissolve in acid or by filtration. If this
is not done, the metal and acid will reduce the cacotheline and give the test
for tin whether the latter is present or not.

40. In both the precipitation, and the Gutzeit test for arsenic, it is best
to run a blank test at the same time as the regular test is being run. To do
this, all of the reagents are added to a drop of distilled water as if the latter
contained arsenic. By comparing the results, errors due to contamination
from the air or reagents are easily avoided.

41. During this precipitation, it is necessary that a moderate excess of
base be present. The reaction of a salt with hydrogen sulfide results in
excess hydrogen ion being formed as illustrated here:

Ni(NO 3 ) 2 -> Ni+ f + 2NO 8 -
H 2 S -> S- + 2H+

This acid must be neutralized by the base in order to prevent reversal of the
reaction. At the same time, the hydroxyl ion concentration must be kept
low to prevent precipitation of the hydroxides of barium, calcium, strontium,
and magnesium, widi the metals of Group III. These apparently opposing
conditions are brought about by the use of a strong solution of NH^OH
buffered with NH 4 C1. This gives a large reserve of base, but a low immedi-
ate hydroxyl ion concentration.

42. The ferric hydroxide precipitated by the hydrolysis of ammonium
benzoate is more crystalline in nature and has less tendency to include
portions of the solution containing the other ions.

43. A large excess of hydroxyl ion is necessary if the aluminate ions are
to remain stable in solution. If the hydroxyl ion concentration is lowered

NOTES ON ANALYTICAL PROCEDURE 169

by adding a buffering agent, such as ammonium chloride, the sodium alum-
inate will decompose and aluminum hydroxide will precipitate. This
precipitates the aluminum hydroxide without the necessity of first acidifying
the solution, and avoids certain difficulties otherwise found in the analysis
of chromium (see Note 46).

44. Both aluminon reagent and alizarin 3 (blue) reagent tend to precipi-
tate on standing in the reagent bottle. It is therefore necessary that these
reagents be freshly filtered by the student just before using. The best
procedure is to take about eight drops of the reagent and centrifuge it,
drawing off enough of the clear solution for the test, using a medicine
dropper.

46. When chromic ion is treated with excess sodium hydroxide, a solution
of sodium chromite (NaCrO-2) is formed. On addition of sodium peroxide,
which is a strong oxidizing agent, the chromite ion is oxidized to chromate
ion.

46. The addition of sodium peroxide to water always results in the forma-
tion of small amounts of hydrogen peroxide. Hence, when sodium chromite
solution is treated with sodium peroxide, the result is a solution of chromate
ion and hydrogen peroxide. If this is acidified, perchromic acid (a highly
unstable, blue substance of uncertain formula) is formed, this rapidly
decomposing to give oxygen and the original chromic ion, Cr ++ +. For
this reason, in the separation of chromate ion from the aluminum, a buffering
agent is used to precipitate the latter instead of, as some procedures direct,
first acidifying and then making the solution alkaline with ammonium
hydroxide.

47. Barium chromate is soluble in strong acids. However, if such solution
is treated with a buffering agent to reduce the hydrogen ion concentration
of the solution, the barium chromate will reprecipitate.

48. Ammonium persulfate in the presence of silver ion is a very powerful
oxidizing agent. When the HNOs solution of a compound of manganese
is treated with this mixture, the manganese is oxidized to permanganic
acid, HMnO 4 , and the persulfate ions, 8208", are reduced to sulfate ions.

49. As mentioned earlier (see Note 27), the sulfides of cobalt and nickel
have a strong tendency to peptize. The ammonium sulfate prevents the
formation of these colloids by neutralizing any charges on the colloidal
particles, as fast as they form.

50. Many metals give characteristic bead tests with borax, these tests
often being used for confirmation tests. The test for cobalt is extremely
sensitive and quite specific. r

61. Cobalt gives a very light, brownish color with dimethylglyoxime if
the metallic ions are present in high concentration. This coloration, how-
ever, is so weak that it will not interfere with the test for nickel in any
ordinary analysis. The test for nickel as described here is sensitive to one
part in a million. Dimethylglyoxime is a reagent for several other metallic
ions but the interfering ions are removed before this point in the procedure.

62. Some helpful information can be gained by carefully watching the
changes that occur in the unknown on the addition of the NH 4 OH. A

170 NOTES ON ANALYTICAL PROCEDURE

gelatinous, reddish-brown precipitate, insoluble in excess NH 4 OH, may
indicate that iron is present. A colorless, gelatinous precipitate that fails
to dissolve in excess NH 4 OH indicates the probable presence of aluminum.
No precipitate at all, in excess NH 4 OH, means that both iron and aluminum
are absent. If, on. adding H 2 S, a white precipitate appears with no other
colored precipitate forming, zinc is probably present, and iron, nickel, and
cobalt are absent.

53. After precipitation is apparently complete, test the solution with
litmus. It should be strongly basic. If the solution is weakly basic or
acidic, add 2 to 3 drops more QN NH 4 OH solution, stir and add more H2S
(see Note 41).

64. The ammonium nitrate-hydrogen sulfide solution is used here for the
same purpose as in Group II (see Note 27). Ammonium sulfate cannot be
used at this point as it would precipitate barium, strontium and, perhaps,
calcium from the filtrate that still wets the precipitate.

55. The action of the ammonium acetate here is as described in Note 46.

66. If the alurninon reagent is weak, it may be necessary to add 8 to 10
drops instead of 3 to 5 drops.

57. If a precipitate is obtained that is not of the proper color, it is best to
filter it, redissolve in dilute hydrochloric acid, add 8 to 10 drops of alizarin,
and reprecipitate with ammonium hydroxide.

68. The high concentration of acetate ions, added during the precipitation
of A1(OH) 8 , causes the formation of slightly soluble silver acetate and pre-
vents the formation of Ag 2 CrO 4 unless excess AgNO 3 is added. If the
solution is yellow before the AgNO is added and no Ag 2 CrO 4 appears, it is
best to add another drop of the reagent.

A white precipitate should be disregarded, as silver acetate will precipitate
whether chromium is present or not. A basic solution should be avoided,
also, as it would then precipitate dark-colored Ag 2 O.

69. If the solution becomes somewhat milky in appearance, evaporate to
half its volume and add more H 2 S. ZnS readily forms colloidal solutions, so
the usual precautions must be taken against this. Plenty of time should
be taken for the precipitation.

60. Another portion of the solution, obtained by dissolving the ZnS, may
be tested by adding 1 drop NH 4 SCN solution and 1 drop pyridine to the
solution. A white precipitate of Zn(C6H 6 N) 4 (SCN)2 confirms the presence
of zinc. Other tests are given under Auxiliary Tests.

61. A dark brown precipitate that fails to dissolve readily in the HNOa
is usually MnO 2 , and suggests the probable presence of Mn++. This may
be dissolved by adding 1 drop QN HC1 to the mixture and warming. If this
procedure is followed, however, it it necessary to evaporate the resulting
solution to half its volume, then add more concentrated HNOa before
adding the crystals of KC1O 3 . Otherwise, the HC1 delays, and may even
prevent, the precipitation of the manganese with KClOs.

62. The formation of a red or pink solution on the addition of the
HgCl 2 -4NH 4 SCN reagent indicates that some iron is still present. The

NOTES ON ANALYTICAL PROCEDURE 171

red color may be removed by the addition of a solution of microcosm io
salt or of a soluble fluoride (see page 99).

63. Barium chromate is the least soluble of the three chromates. Stron-
tium chromate is about 500 times as soluble as barium chromate but is still
quite insoluble, while calcium chromate is moderately soluble in water.
Therefore, it is easy to precipitate the chromates of barium and strontium
without that of calcium unless calcium is present in great concentration.
The strontium chromate may then be dissolved out of the mixed precipitate
in a buffered acid solution, leaving the barium salt behind. It is necessary
however, that the proper buffer mixture be used, since more strongly acid
solutions will dissolve barium chromate. The proper mixture is obtained
by mixing equal volumes of IN acetic acid and 2.5N ammonium acetate.

64. Heating during the addition of reagents will cause the formation of
larger crystals and make filtration easier.

v 65. The platinum wire must be cleaned thoroughly each time before it is
used. This is done by heating the wire red hot and at once plunging it into
clean, concentrated HC1. This process is repeated until the wire no longer
causes any color in the flame.

66. The chromates, sulf ates, and oxalates of the metals are not sufficiently
volatile to color the flame much. The chlorides, on the other hand, give
excellent results. If the precipitate were dipped into HC1, however, it
would dissolve off. The vapor treatment described will convert some of the
salt to chloride and still avoid loss of precipitate.

67. Notice that whereas barium chromate was easily precipitated, stron-
tium chromate does not precipitate until the acid is neutralized.

68. If not all the precipitate dissolves, filter and treat the residue with a
fresh mixture of the acetic acid-ammonium acetate mixture.

69. Calcium oxalate, though very insoluble, has a strong tendency to
supersaturate. Therefore, it is necessary to heat the solution, scratch the
side of the beaker with a stirring rod, and let it stand.

70. The acidification with HC1 followed by addition of ammonium hydrox-
ide furnishes some ammonium chloride. However, caution should be used
to add just enough reagent to complete the precipitation for, if too much is
used, any magnesium present may precipitate and spoil the analysis of this
group.

71. In place of acetic acid, HNO 8 may be used and the resulting solution
evaporated. Heat strongly to dehydrate, and treat at least twice with
anhydrous acetone. The Ca(NOa)2 dissolves, leaving Ba(NO 8 )2 and
Sr(NO 3 )2 as residue. The ions are then identified by an adaptation of the
procedure described.

72. It is well to recall the preliminary experiment on calcium, using this
reagent, and to reread Notes 64 and 69. If a precipitate forms at once, it
means either that some strontium has escaped precipitation earlier or that
calcium is present in unusually high concentration usually the former.

73. The reagent (p-nitrobenzeneazoresorcinol) is destroyed by either
strongly acid or strongly basic solutions. If this occurs (indicated by the
solution turning yellow on addition of the reagent), it is necessary to test

172 NOTES ON ANALYTICAL PROCEDURE

the solution with litmus and neutralize it until it is very faintly acidic, then
add fresh reagent.

74. All the slightly soluble salts of potassium, ammonium, and sodium
readily form supersaturated solutions. Consequently, when trying to pre-
cipitate these ions, use all the precautions possible against supersaturation.

76. This reagent is prepared by the student just before using. Make a
saturated solution of sodium cobaltmitrite by shaking a couple small crystals
of the solid with 3 drops water. Dilute with an equal volume of water and
filter. Use at once, as the reagent decomposes in a comparatively short
time.

76. When sodium cobaltmitrite solution is added to a solution containing
a, mixture of K+ and Ag 4 , a yellow precipitate of K2AgCo(NO2)e is formed.
This compound is much less soluble than is K-2NaCo(NO 2 )6 and is capable
of detecting as little as 0.0009 mg. K + in a drop of solution.

77. In case sodium is present in moderately large concentrations it may
be necessary to use a double thickness of cobalt glass to cut out the sodium
light, in using the flame test for potassium.

78. The light from the lavender potassium flame will penetrate a piece of
blue glass while the yellow light of the sodium flame will not. This makes
it possible to see the potassium flame even if sodium is present. Occasion-
ally, however, a double thickness of glass is required.

79. On heating, the cobaltmitrite ions decompose. If ammonium ions
are present, the two ions react in the hot solution and the ammonium ion is
destroyed. Any potassium can then be reprecipitated by adding more
sodium cobaltinitrite to the cold solution. This is an excellent confirmatory
test for potassium.

80. It is never to be considered an acceptable flame test if the yellow color
lasts leas than 30 sec., as sodium contaminations are almost invariably
present, either from the glass or from the air or fingers. The platinum wire
must be carefully cleaned before each test by repeated heating to redness and
plunging into some clean concentrated hydrochloric acid.

81. Overheating will cause loss of the ammonium salts present as most
ammonium salts decompose or volatilize on heating.

82. If the paper does not become uniformly blue, it indicates that any
blue spots present are due to spattering of the sodium hydroxide. The test
must then be tried with a fresh sample of unknown.

83. If the unknown is a general, the test for ammonium ion must be tried
before anything else is done. If it is allowed to stand, the unknown may
pick up ammonia fumes from the air and give the test for ammonium ions.

84. The original solution should be tested with litmus before adding any
reagents, and basic solutions should be reported as containing OH~. A
solid material may be treated with water and the aqueous solution used for
this test.

86. In neutral or alkaline solutions it is possible to have a mixture of 8 %
SOs", and 8203". These may be separated and identified in the following
manner:

NOTES ON ANALYTICAL PROCEDURE 173

Add CdCOa to the solution, heat for 30 min. on the steam bath, then filter.
Test the precipitate for sulfide.

To the filtrate from this treatment add an excess of SrCl 2 solution, rub
the walls of the container with a stirring rod, then let stand at least 30 min.
Filter. Test the precipitate for sulfites and the filtrate for thiosulfates in
the usual manner.

86. In the presence of S~ or NO 2 ~ this test is riot specific, as the H 2 S and
the NO 2 given off when sulfides and nitrites are acidified will bleach per-
manganates also. If the original solution is acid, however, there will be
no conflict, for, as pointed out earlier, these ions cannot exist together in
acid solutions.

87. If S~ were found present, the appearance of sulfur at this point does
not necessarily mean S 2 O 3 ~ are present. Sulfide ions react with sulfurous
acid, in acid solution, as follows:

S- + SOr -f 6H*- 3H 2 O + 2S( I )

88. This reaction is specific for S 2 O 3 " in the presence of SOs". It depends
upon the fact that silver ions react with thiosulfates in neutral solutions to
form Ag 2 S 2 Og which breaks down to form black Ag 2 S.

89. A test recommended by C. Spacu and P. Spacu (34) is as follows:
To 1 drop of the prepared solution add 1 drop NV + test solution and 2 to 3

drops ethylenediamine. Stir, and let stand. If S 2 Oj~ are present, a violet,
crystalline precipitate will form. This test is specific as far as S", SOa"", and
SCN~ are concerned.

90. If SOa"* or S 2 O 3 ~ are present, they must be destroyed by acidifying
the solution with dilute H 3 PO 4 , adding 1 to 2 drops 3N K 2 CrO 4 solution and
letting it stand for 2 to 3 min. This solution is then tested for COa~ in
the usual manner.

91. This test depends on the fact that the CO 2 given off reacts with the
Ba(OH) 2 to form insoluble BaCOa. Since the air in most laboratories
contains a fairly large concentration of CO 2 , it is best to hold another drop
of Ba(OH) 2 solution near the apparatus and report COa" only if the drop
above the reaction mixture becomes cloudy more quickly than does the
blank test.

92. If interfering ions, such as SCN~ or Fe(CN)e a f are present, make the
solution very slightly acid, suspend 1 drop 2N NaOII in the tip of a glass
tube a short distance above the mixture and warm gently for 2 to 3 min.

Dip the end of the glass tube containing the NaOH, into 2 to 3 drops
water and stir a moment. Using this solution, perform the test for cyanide
as described.

93. The accurate analysis of solutions containing all six of the members
of Group II is too involved to include here. For a more detailed outline
of the procedure to use for such solutions, reference should be made to a
more comprehensive text such as the treatise by Tread well and Hall (22).

94. This test works only in the absence of I~ and Fe(CN)e~. The former

174 NOTES ON ANALYTICAL PROCEDURE

reduces Fe ++ + to Fe f +, giving at the same time a yellow or brown solution
of 1 2. The reaction is:

2Fe++ + + 21- 2Fe ^ + I 2

The Fe(CN)6 S give a deep blue color or precipitate which covers up the red
color of the Fe(SCN) 3 .

96. If the cation analysis showed copper to be present, it must be removed
before applying this test. This may be done by the usual Group TI metal
sulfide precipitation. The nitrate from this should be boiled to expel H 2 S
and the test then applied.

96. If iodide ions were found present, treat 2 to 3 drops of the prepared
solution with excess (about 5 drops) Cu + + test solution, warm, rub the inside
of the container with a stirring rod, and let stand about 5 min. Filter, dis-
carding the precipitate. Evaporate the filtrate barely to dryness, cool, and
dissolve the residue in water. Use this solution in testing for Br~. The
reactions involved are:

Cu++ + 21- CuI 2

2CuI 2 Cu 2 I 2 ( I ) + I 2

An alternative method is to add 1 drop H 2 O 2 and 5 drops CCU to about
4 drops of the prepared solution, shake well, then withdraw the CC1 4 layer with
a medicine dropper. Repeat this process until the CCU layer is no longer
colored violet or pink. If Br" are present, the CCU layer will then be yellow
or orange.

97. This test does not work if SCN~, Fe(CN) b ", or Fe(CN) 6 s are present
as these ions give a similar reaction. For the treatment of such cases, see a
more comprehensive text.

98. If the cation analysis on this solution showed chromium is absent,
chromates and dichrornates need not be sought. However, if chromium is
present, these test should be carried out.

If the original unknown gives an acid reaction with litmus, any positive
results with these tests should be reported as Cr 2 O7~; if alkaline, CrO4~
should be reported. Similarly, if S~", COa", or SOa" are found in a solution
that was, originally, acid, they should be reported as HS", HCOa", arid
HSOs~, respectively.

99. Unless the precipitate is definitely reddish or brownish in appearance,
do not report CrO4~ on the basis of this test, as I", PC>4~, and several other
ions give yellow precipitates with Ag + .

100. The coin test shows merely that the precipitate contains sulfur. At
this point in the procedure, however, it becomes specific for SO 4 ~.

The use of this test prevents the analyst's being misled by a precipitate
of BaF 2 in case F" are present.

101. This test depends upon the fact that H 2 S precipitates As 2 S 6 very
slowly from weakly acid solutions of arsenates, whereas it precipitates
As 2 Ss readily, from solutions of arsenites (see pages 84 and 85).

102. As many organic compounds reduce alkaline copper solutions in a
similar manner, this test is not always specific for arsenites. However,

NOTES ON ANALYTICAL PROCEDURE 175

with organic compounds eliminated, the test may be considered quite
specific at this point in the procedure.

103. It is best to remove the excess of Ag f , from the AgNOa used in the
preliminary treatment, before adding the NH 4 OH. This is done by adding
NH 4 C1 solution and filtering, using the filtrate for the test described here.

104. In the tests described here, no distinction is made between BOa**
and 640?"". The reason for this is that acid solutions of B 4 O7~ contain
some BOa 8 *, and the tests used here are characteristic of these ions.

106. If the cation analysis showed copper to be present, it should be
removed by the H 2 S treatment described in Note 95. Copper compounds
also give a green color to flames; and if it is not removed, this test may be
misleading.

106. Iodide, thiocyanate, nitrite, chromate, ferrocyanide, and ferri-
cyanide ions interfere with the "ring" test for nitrates. However, these
have been removed by the preliminary treatment.

It is essential that care be taken to use the solution that was prepared
especially for this test, as the other Group IV prepared solution always
contains nitrate ions introduced when AgNOa was added to remove the
Group II anions.

107. Another test for acetate ions is to evaporate 3 drops of the solution
to dryness, add a little dry As 2 O 3 and heat the mixture. A very poisonous,
foul-smelling substance, cacodyl oxide, ((CHs^As^O, is given off. This is a
very sensitive test but is specific only in the absence of other organic acids.

108. When chlorates are treated with alkali and aluminum or zinc, the
chlorate is reduced to chloride. Since any chloride ions present in the original
unknown were removed during the preliminary treatments, the formation
of AgCl at this point is a good test for CIO*".

If excess AgNO 3 were added during the preliminary treatment, a white
precipitate of AgCl may be found even before the AgNOa solution is added
in this test. This, however, merely confirms the presence of ClOa".

109. If the prepared solution, containing MnO 4 ~, is made alkaline
with NaOH before the HgOa is added, the same evolution of oxygen will
occur. In this case, however, a brown precipitate will form instead of the
solution becoming colorless. This is due to the fact that MnO 4 ~ are reduced
to MuOa-HaO in alkaline solution instead of to Mn+ + .

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APPENDIX

TABLE I. OXIDATION- REDUCTION POTENTIALS*

Electrode Reaction E.m.f., Volts

Li -*Li+ + e. . . . -2 959

K-K+ + .. -2 924

Ca-Ca^+4-2e .... . -27

Na - Na+ + e. . . . . . -2 714

Mg - Mg++ 4- 2c . -2 40

A1->A1 +++ 4-3* ... -17

Zn 4- 2(OH~) -> ZiiO 4- H 2 O + 2 e - 1 . 248

Mn -* Mn ++ 4- 2e - 1 1

Tl 4- I- - Til + e -0 77

Zn - Zn+ + + 2. . . -0 76

Pb 4- 2(OH~) -* PhO 4- H 2 + 2e -0 57

Cr -> Cr+ + + 2e -0 55

S--S +2 e . -0 51

Fe -> Fe ++ + 2 . . -0 44

Or-"- - Or-" + + e -0 40

Cd-Cd ++ +2e -0 40

2Cu + 2(OH-) - CujO + HO + 2e. . . . -0 34

Cu 4- H 2 S -* CuS + 2H+ + 2. . . -0 26

Ni - Ni ++ + 2c . -0 23

Sb -1- H 2 O -> SbO + + 2H+- + 3* -0 21

Cu 2 O + 2(OH~) -> CuO -f H 2 O + 2e . -0 15

Sn -> Sn ++ + 2e . -0 13

Pb-Pb ++ 4-2e.. ~0 12

Cu 2 O + 2(OH~) + H,O -* 2Cu(OH), + 2 -0 08

2Hg + 21- -* Hg 2 I 2 + 2e -0 041

2Ag + H 2 S -> Ag 2 S + 2H+ + 2c -0 036

HH 2 -*H + + e . . . 000

Ag + Br~ -> AgBr + . +0 073

Hg + 2(OH-) -* HgO + H,O + 2* . + 097

Cu + Cl- -> CuCl + e +0 128

Sn++ - Sn ++++ + 2e +0 13

2Sb 4- 3H 2 O - Sb 2 O 3 + 6H+ + 6c 4-0 144

H 2 S -* 2H+ 4- S 4- 2 . 4-0.17

Bi - Bi +++ + 3. . . . . . 4-02

* Most of these values were taken from the " International Critical Tables." In the cahes
where the values given in the above source were at wide variance with the values most
widely accepted, these were changed to values reported in the literature, due consideration
being given to the source of such information.

179

180 APPENDIX

TABLE I. OXIDATION- REDUCTION POTENTIAL. (Continued)

Electrode Reaction

E.m.f., Volts

Ag + Cl- - AgCl + e. . .

. ... +0.222

PbO + 2(OH-) -> PbO 2 + H 2 O + 2e

. .. +0.27

Cu -Cu +4 - + 2e. . . .

+0 344

Ti 4 " 4 " ^ r ['j-4-+4- r \ e

+0 37

CuCl 2 - Cu + + + 2C1~ +

. . +0 45

Fe(CN) fl M -> Fe(CN) 6 - +

. . +0 486

2Ag + COs" > Ag 2 COa + 2c .

+0.50

Cu -> Cu + +

+0 51

+0 534

AsOa 31 + H 2 O -> AsO 4 E + 2H - + 2e

. . +0 574

2Hg + SOr - Hg 2 S0 4 + 26

+0 621

MnO 4 - - MnO 4 ~ + e

+0 66

Fe ++ > Fe +f+ + e .

+0.747

Ag - Ag + + e

+0 797

2Hg -* Hg 2 ++ + 2c . .

+0 798

H 2 O 2 -^2H + + 2 +2 .

+0 84

Hg -> Hg ++ + 26 . .

+0 86

JDF ^ /*y> J3r 2 ~T~ c

+ 1 064

1 / T | O TJ /^ T /"\ 1 f* TT f- | }

/ X2i 2 + orljjLf > lv/3 ~r O-H. -p oc

+ 1 19

Tl 4 ^ * Tl 4 ++ + 2c .

+ 1 2

Au - Au +4+ + 3*

+ 1 3

Mn ++ + 2H 2 O -> MriO 2 + 4H + + 2c

+ 1 33

Cl" > J^C1 2 +

. +1 358

Pb+ f + 2H 2 O -> PbO 2 + 4H 4 + 2e

. +1 44

Au Au + + e

+ 1 5

2SO 4 *" > S 2 OH"" + 2e.

+ 1 5

Mn^ + + 4H 2 O -> MnOr + 8H 4 + 5e

. +1.52

MnO 2 + 2H 2 O -* MnO 4 - + 4H 4 + 3

+ 1 58

PbSO 4 + 2H 2 O -> PbO 2 + 4H+ + SO 4 - + 2

+ 1 679

Co 4 " 1 - > Co ++4 " + c ....

...+18

F~ > ^F +

+ 1 96

APPENDIX
SOLUBILITY PRODUCT CONSTANTS*

181

Salt

Temperature,
C.

Constant

AgBr

7.7 X 10-"

AgCl

1.6 X 10~ lfl

AgCN

2.7 X 10~ 12

Ag 2 CO 3

8.2 X 10~ 12

Ag 2 CrO 4

3.3 X 10~ 12

Agl

1 5 X 10- lft

BaCOs

8 X 10~ 9

BaCrO 4

2 2 X 10~ 10

BaS0 4

1.0 X 10' 10

CaCOa

8 X 10 -

CaC 2 O 4

2 6 X 10-"

CaF 2

3.8 X 10- 11

CaSO 4

2 3 X 10-

MgCOa

1 9 X 10-

Mg(OH) 2

2 X 10' i

PbCO 3

3 3 X 10- *

PbCrO 4

1 8 X 10- 4

PbSO 4

1 X 10-

SrCO 3

1.6 X 10~ 9

SrCrO 4

3.6 X 10-*

SrS0 4

2 8 X 10-'

A g2 S

10 -w

CuS

10-42

FeS

10-"

HgS

10-54

MnS

10-W

PbS

10-28

ZnS

1 2 X 10- 23

* Hammett, "Solutions of Electrolytes " (1929).

182

APPENDIX
IONIZATION CONSTANTS*

Reaction

Temperature,
C.

Constant

H 3 BO 3 = H+ + HjBOr

6 5 X 10- 10

H 2 CO 3 = H+ + HCOr

3 3 X 10- 7

HCO 3 " = H+ + COr

3 9 X 10- 11

HCN = H+ + CN-

7 2 X 10-w

H 3 PO 4 = H+ 4- H 2 P0 4 -

8 3 X 10--*

H 2 PO 4 - - H+ + HP0 4 -

2 X 10~ 7

HP0 4 - = H + -j- PO 4 S

4 X 10-"

HNO 2 = H+ -f NO 2 -

4 5 X 10~ 4

HSOr = H+ -f S0 4 -

1 15 X 1C- 2

H 2 S - H+ -f HS-

9 X 1C- 8

HS- H + + S-

1 X 10-

H 2 SO 3 = H+ + HSOs-

1 2 X 10~ 2

HSOr = H+ + S0 3 =*

5 X 10-

HC 2 H 3 2 = H+ -f C 2 H.G 2 -

1 86 X 10~ 6

NH 4 OH = NH 4 + + OH

1 8 X 10-*

* Hammett, "Solutions of Electiolytes" (1929)

LIST OF APPARATUS*

Number (per student) :
2 Beakers, one 250 ml. and one 400 ml.
1 Beaker rack. Made by boring six holes > in. diameter and \fa in. deep

about 2 in. apart (center to center) in a softwood block 8 by 4 by 1 in.
1 Bottle (about 4> in. high and capable of taking a No. 6 or 7 rubber

stopper).

Centrifuge, hand or electrical (1 for each 10 to 12 students).

1 Clamp (screw).
Cotton, absorbent.

2 Crucibles, 10 ml.

1 Cylinder, graduated, 10 ml.

2 Dishes, evaporating, 60 mm.

Filter paper (2 doz. sheets, acid- treated).
1 File, triangular.
1 Flask, Florence 500 ml. (to be made into wash bottle with capillary

tip).

6 Glass rods, 1 to 2 mm. by 15 cm. long.
1 Lead dish (about 2 cm. diameter and 1 cm. deep. May be shaped from

a 1 in. square of He in. sheet lead).
12 Medicine droppers.
12 Microscope slide.
1 Rubber stopper, No. 7.
1 Rubber stopper, No. 2.
1 Spot plate
1 Steam bath (may be made by boring three holes capable of taking

* It is desirable to have individual reagent kits whenever possible, as this
saves much time in the laboratory. These kits may be either purchased or
made. A convenient kit is made by arranging two blocks about 8 by 12
by 1 in. thick, with holes bored in each to accommodate 50 vials each.
About 80 of the vials should be arranged with dropper tops (see Reagent
Containers) for dispensing liquids; about 20 of the vials should have ordinary
cork or screw tops for holding solids.

About 10 of the vials should have a capacity of 30 ml. each. These
hold the reagents most commonly used, such as HC1, HNOz, H-CaHsOc,
H 2 SO 4 , saturated NH^HsOi, saturated NH 4 C1, etc. The other vials
should have a capacity of 8 ml. each.

If individual reagents kits are not available, the reagents may be placed
in 4-oz. dropper-type bottles on the side shelf for use by the whole class.

183

184 LIST OF APPARATUS

a No. 7 cork in the lid of a Ji-pint "Karo" sirup can, fitting them with
corks bored to loosely hold the microbeakers. A hole should be bored
in the side of the can near the top to art as a steam outlet).
20 Test tubes, soft glass 4 by ><j in. (or 18 Pyrex test tubes 75 by 10

mm., and two 4 by J^-in. Pyrex test tubes).
1 Test tube, soft glass, 6 by % in.
1 Test tube, Pyrex, 6 by % in.
1 Tongs, or forceps.
2 ft. Tubing, glass, 6 mm O.D.
2 ft. Tubing, glass, 10 to 12 mm.
1 ft. Tubing, rubber.
1 Wire gauze.
1 Wire triangle.

In addition to the apparatus listed above, blowpipes, platinum wires,
standard small ring stands fitted with rings and small burette clamps, and
Bunsen burners should be available. The latter may be used as micro-
burners by unscrewing the burner tube and lighting the gas at the orifice
in the base. In this way the Bunsen burner serves two purposes.

Experience has shown that the small electrically driven centrifuges with
angle type head, which have been designed for semi-micro work, are more
convenient and economical of time and cost than the less expensive hand-
operated ones.

LIST OF REAGENTS

(All chemicals should be reagent grade)

REAGENTS FOR CATIONS

SOLID REAGENTS

Acid, gallic

Aluminum (granular, As free)

Ammonium chloride

Ammonium iodide

Ammonium persulfate

Ammonium sulfate

Borax (c.p.)

Ferrous sulfide (sticks) (see Hydrogen Sulfide Generator)

Potassium chlorate (c.p. crystals)

Sodium carbonate (anhydrous, c.p.)

Sodium cobaltinitnte

Sodium hydroxide (c.p. pebble form)

Sodium thiosulfate

Sulfur-paraffin mixture (see Hydrogen Sulfide Generator)

Thiosinarnine (allyl thiourea)

LIQUID REAGENTS AND SOLUTIONS

Acid, acetic

Concentrated (approximately 17 N)

Dilute (6AO

Dilute (exactly IN)
Acid, hydrochloric

Concentrated (approximately 13 AT)

Dilute (6#)
Acid, nitric

Concentrated (approximately 16 AT)

Dilute (6AO
Acid, picric

Saturated aqueous solution
Acid, sulfuric

Concentrated (approximately 36JV)

Dilute (6AT)

Dilute (1:4. Made by adding 1 volume of concentrated acid to
4 volumes of water)

185

186 LIST OF REAGENTS

Acid, tartaric

10 g. acid to 100 ml. water
Alcohol, ethyl (95 per cent)
Alizarin S (blue)

Dissolve 0.05 g. in 100 ml. of 50 per cent acetone and add 2 drops

glacial acetic acid
Alurninon, Ci9H 9 O2(COONH 4 ) 3

Dissolve 0.1 g. of aluminon (the ammonium salt of aurin tricarboxyhc

acid) in 100 ml. water
Ammonium acetate

Saturated solution in water

A '2.5N solution (18 g. dissolved in a little water and the solution
diluted to 100 ml. A better way of preparing this is to mix equal
volumes of exactly 5AT acetic acid and 5JV ammonium hydroxide)
Ammonium benzoate (7 per cent)

Dissolve 7 g. in 100 ml. water
Ammonium carbonate reagent

Dissolve 20 g. ammonium carbonate and 30 g. NH 4 C1 in 100 ml.

6JV NH 4 OH
Ammonium carbonate

Dissolve 15 g. ammonium carbonate in a mixture of 8 ml. concen-
trated ammonium hydroxide and 50 nil. water, then dilute to
100 ml.
Ammonium chloride

Saturated solution in water

IN solution (5.35 g. dissolved in water and diluted to 100 ml.)
Ammonium hydroxide

Concentrated (approximately 15./V)

Dilute (6JV)
Ammonium nitrate

Wash solution. Dissolve 8 g. in 100 ml. water
Ammonium sulfate

Wash solution. Dissolve 13 g. in 100 ml. water
Ammonium thiocyanate

Dissolve 8 g. in 100 ml. water
Barium chloride

Dissolve 2 g. of the anhydrous salt (or 2.4 g. of the dihydrate) in

100 ml. water
Benzidine

Dissolve 0.05 g. of the pure substance in 10 ml. glacial acetic acid

and add 90 ml. water
a-Benzoinoxime

Dissolve 5 g. in 100 ml. ethyl alcohol.
Bromine water

Saturated solution in water
Cacotheline

Saturated solution in water

LIST OF REAGENTS 187

Cinchonine

Dissolve 1 g. of reagent in 100 ml. water, add about 10 drops nitric

acid, and warm until dissolved. Cool, and add 2 g. of solid KI
Copper sulfate (0.1 per cent)

Dissolve 0.1 g. of CuSO 4 5H 2 O in 100 ml. water
Dimethylglyoxime

Dissolve 1 g. in 100 ml. ethyl alcohol
Diphenylcarbazide

Prepare 100 ml. of a saturated solution of reagent in 50 per cent

acetone. Filter and saturate with KSCN. Then add 1 g. KI
Ether, diethyl

Hydrogen peroxide (3 per cent)
Mercuric chloride

Dissolve 2.7 g. in 100 ml. water
Mercuric chloride-ammonium thiocyanate reagent (NH 4 )2Hg(SCN)4

Dissolve 30 g. of ammonium thiqcyjuuite in 100 ml. water, then add

27 g. of HgCl-2, and stir until dissolved
p-Nitrobenzeiieazoresorcinol

Dissolve 0.001 g. in 100 ml. 0.5 per cent NaOH and filter if necessary
a-Nitroso-0-naphthol

Dissolve 1 g. in 50 ml. glacial acetic acid arid add 50 ml. water
Potassium chromate (3AT)

Dissolve 58.2 g. of KaCrC^ in enough water to make 100 ml. solution
Potassium cyanide (3 per cent)

Dissolve 3 g. in enough water to make 100 ml.
Potassium ferrocyanide

Dissolve 8.5 g. in 100 ml. water
Potassium iodide

0.1 AT solution. Dissolve 1.7 g. in 100 ml. water

Saturated solution in water
Potassium oxalate (3N)

Dissolve 55 g. in 100 ml. water
Potassium sulfate (0.3W)

Dissolve 5.2 g. in 100 ml. water
Potassium thiocyanate

Dissolve 10 g. in 100 ml. water
Pyridine (pure)
Rhodanine (dimethylaminobenzal rhodanine)

Dissolve 0.03 g. in 100 ml. acetone
Silver nitrate

1 per cent solution. Dissolve 1 g. AgNOa in 100 ml. water

50 per cent solution. Dissolve 50 g. of AgNO 8 in 100 ml. water
Stannous chloride

Dissolve 10 g. in 100 ml. concentrated hydrochloric acid and add two

or three pieces of tin to each bottle
Sodium acetate

Saturated solution in water

188 LIST OF REAGENTS

Sodium bicarbonate

Dissolve 8.4 g. in 100 ml. water
Sodium bisulfite (0.1 N)

Dissolve 1 g. in 100 ml. water
Sodium hydroxide (6AO

Dissolve 24 g. of NaOH in 100 ml. water
Sodium phosphate (Na 2 HPO 4 )

Dissolve 10 g. of Na2HPO 4 -12H 2 O in 100 ml. water
Zinc uranyl acetate

Dissolve 10 g. of uranyl acetate and 30 g. of zinc acetate in 100 ml.
3 per cent acetic acid. Let stand over night and filter

ADDITIONAL REAGENTS FOR ANION ANALYSIS

Alcohol, amyl

Alcohol, methyl

Ammonium molybdate reagent

Dissolve 4 g. MoO 3 in 20 ml. QN NH 4 OH, add 75 ml. GN HNO 3 , and

dilute to a volume of 100 ml.
Ammonium polysulfide

Saturate 100 ml. 6JV NH 4 OH with H 2 S, add 1 g. sulfur, and shake

vigorously.
Barium chloride-calcium chloride reagent

Dissolve 24 g. BaCl 2 -2H 2 O and 22 g. CaCl 2 -6H 2 O in 100 ml. of water
Barium hydroxide

Saturated solution
Carbon tetrachloride
Copper nitrate (0. 1 N)

Dissolve 1.2 g. Cu(NO 3 ) 2 -3H 2 O in 75 ml. water, add 10 drops concen-
trated HNO 3 and dilute to 100 ml.
Diphenylamine reagent

Dissolve 0.4 g. diphenylamine in 100 ml. 80 per cent H 2 SO 4
Ferric chloride (0.05JV)

Dissolve 0.5g. of FeCl.v6H 2 O in 70 ml. water, add 5 ml. concentrated

HC1, and dilute to 100 ml.
Ferrous ammonium sulfate

Dissolve 2 g. FeSO 4 -(NH 4 ) 2 SO 4 -6H 2 O in 80 ml. water, add 3 ml.

6N H 2 SO 4 , and dilute to 100 ml.
Iodine solution

Saturated solution.
Lanthanum nitrate

Dissolve 4 g. La(NO 3 )3 in 100 ml. water
Lead acetate (0.1 N)

Dissolve 1.9 g. Pb(C 2 H 3 O 2 ) 2 -3H 2 O in 80 ml. water, add 1 ml. glacial

acetic acid, and dilute to 100 ml.
Magnesia mixture

Dissolve 10 g. MgCl 2 -6H 2 O and 10 g. NH 4 C1 in 40 ml. water, add 50 ml.
concentrated NH 4 OH, and dilute to 100 rnl.

LIST OF REAGENTS 189

Potassium permanganate

Dissolve 0.5 g. KMnO 4 in 100 ml. water and add 3 ml. concentrated

H 2 S0 4

Silver carbonate (solid)
Starch-potassium iodide paper
Turmeric solution

Saturated solution in 95 per cent alcohol

TEST SOLUTIONS AND UNKNOWNS

In making up unknowns and test solutions, some instructors want the
solution to contain a definite weight of the cation or the anion in a given
volume. Others, wishing to emphasize the laws of chemistry rather than
the practical aspects of analysis, prefer that the solutions have a standard
molarity, as most of the calculations in theory are made using the concen-
trations of solutions in mols per liter.

Semi-micro methods do not require any special concentrations for their
use. The usual concentrations can be used in semi-microanalysis as well
as they could in macroanalysis.

For the instructors convenience, however, the following table is included,
giving the weights of the salts of different cations needed to make a solution
of the cation for test solutions and unknowns.

It is suggested that these cation stock solutions be diluted 1 : 9 in actual use.

Anion test solutions should be 0.1 M, as a rule. Anion unknowns should
be made from stock solutions that are 0.5 A/ with the sodium or potassium
salts of the desired anioris.

On an average it will be found that 50 ml. of each test solution and 50 ml.
of each stock solution for making unknowns will be more than ample for
40 students, during a one-semester course.

190

TEST SOLUTIONS AND UNKNOWNS

191

Weight of salt per

100 ml. to make

Metal

Formula of salt

O.lAf,

100 mg. per

milliliter,

grams

Aluminum

A1(N0 3 ) 3 -9H 2

3 75

Antimony *

SbCla

2 28

18 7

Arsenic t

As 2 O 3

0.98

Barium

Ba(N0 3 ) 2

2 61

Bismuth % . .

Bi(NO 3 ) 3 -5H 2 O

4 85

23 3

Cadmium

Cd(N0 3 ) 2 -4H 2

3 08

27 5

Calcium

Ca(NO 3 ) 2 -4H 2 O

2 36

Chromium

Cr(N0 3 ) 3 9H 2 O

4 00

Cobalt .

Co(NO 3 ) 2 6H 2 O

2 91

49 5

Copperf . .

Cu(NO 3 ) 2 3H 2 O

2 41

38 3

Troii

Fe(NO 3 )j-9H 2 O

4 04

72.3

FeSO 4 (NH 4 ) 2 SO 4 -6H 2 O

3 92

70.2

Lead

Pb(N0 3 ) 2

3 31

Magnesium

Mg(NO 3 ) 2 -6H 2

2 56

106

Manganese

Mn(NO 3 ) 2 -6H 2 O

2.87

52 3

Mercury (Hg++)

Hg(N0 3 ) 2

3 24

15 7

Mercury (Hg 2 ++)||

Hg 2 (N0 3 ) 2 -2H 2

2 80

13 6

Nickel

Ni(NO 3 ) 2 -6H 2 O

2 90

49 6

Potassium

KNO 3

1 01

25 6

Silver

AgN0 3

1 69

15.8

Sodium

NaN0 3

Strontium

Sr(N0 3 ) 2 -4H 2

2 83

32 5

Tin (Sn ++++ )f ....

SnCl 4 3H 2 O

3 14

26 4

Tin (Sn++)H

SnCl 2 -2H 2 O

2 25

Zinc

Zn(NO 3 ) 2 -6H 2 O

2 97

45 4

* Use 6AT HC1 as the solvent,
t Dissolve in hot 6N HC1.
t Dissolve in 3N HNOa.

Dissolve in water, then add 1 ml concentrated HNO to each 100 ml of solution.
|| Dissolve in 100 ml. water and add 3 ml. concentrated HNO 3 and a small globule of
metallic mercury.

If Dissolve in 50 ml. concentrated HC1, dilute to 100 ml , and add a little pure metallic tin.

INDEX

Acetates, tests for, 155

Acidity and pH, 29

Acids, di- and tri-basic, ionization

of, 25
Activity, coefficient of, 64

meaning of, 64
Additivity, principle of, 21
Adsorption, 65
Alizarin S (blue) test for aluminum,

70, 108, 111

Alloys, treatment of, 73
Aluminon test for aluminum, 70,

108, 111
Aluminum, amphoteric nature of, 41,

100, 108

analysis for, 106, 110
properties and compounds of, 99,

100

reactions of, 108, 131
Ammonia complexes, 37, 38, 82, 84,

103, 104, 105

Ammonium, analysis for, 121, 124
properties and compounds of, 120
reactions of, 122, 135
Ammonium hydroxide, dual nature

of, 38
Ammonium mercuric thiocyanate,

67
Ammonium persulfate, oxidation

with, 103, 109, 112
Ammonium sulfide, 165
Amphoteric hydroxides, action of, 41
elements forming, 43
precipitation of, 42
Analysis, anion, discussion of, 142
Group I, 143, 145, 146
Group II, 143, 145, 148

Analysis, anion, Group III, 144, 145,

150

Group IV, 144, 145, 154
groups in, 142
notes on, 172-175
preliminary treatment of solu-
tions for, 144
cation, 71, 74

general discussion of, 73
Group I, 78, 81
Group II, 88, 93
Group III, 106, 110
Group IV, 115, 118
Group V, 121, 123
notes on, 163-172
qualitative, meaning of, 2
quantitative, meaning of, 2
rules for, 71

Antimony, analysis for, 88, 93, 96
Marsh test for, 86
properties and compounds of, 86
reactions of, 92, 130
Apparatus, construction of, 8
conversion of macro into, 8
gas evolution, use of, 7

construction of, 14
list of, 183
Appendix, 179
Aqua regia, preparation of, 73, 163

reactions of, 164
Arrhenius, theory of, 21, 23
Arsenates, action of H2S on, 85
reactions of, 85
tests for, 152

Arsenic, analysis for, 88, 93, 96
Gutzeit test for, 92, 166
Marsh test for, 85
properties and compounds of, 84
reactions of, 92, 128

193

194

INDEX

Arsenites, reactions of, 84

tests for, 152
Arsine, 85
Aurin tricarboxylic acid (alurninon),

70, 108, 111
Auxiliary tests for cations, 124-136

Barium, analysis for, 115, 118

properties and compounds of, 113

reactions of, 116

Barium sulfate, dissolving, 114, 144
Bead tests, borax, 110, 111, 113, 130

sodium carbonate, 134
Benedetti-Pichler test for zinc, 132
Benzidine as a reagent, 80, 112, 131
a-Benzoinoxime, reaction of, 69, 90,

95
Bismuth, analysis for, 88, 93

properties and compounds of, 83

reactions of, 90, 127
Borates, tests for, 154
Bromides, tests for, 149
Bronsted, theories of, 59
Bronsted and La Mer, formula of, 64
Brown ring test, 154
Buffers, 24, 43

Cacotheline test for tin, 96, 97
Cadmium, analysis for, 88, 93

properties and compounds of, 83

reactions of, 91, 127
Calcium, analysis for, 115, 118

properties and compounds of, 115

reactions of, 117, 134
Carbonates, tests for, 147
Cations, analytical aspects of, 71,
78, 88, 106, 115, 121

equations for reactions of, 136-139
Cells, oxidation-reduction, 45
Centrifuge, counterbalancing of, 4

types of, 4, 5, 184

Centrifuge, use of, 4
Centrifuge tubes, construction of, 9
Chelate compounds, 70
Chlorates, tests for, 155
Chloride complexes, 38, 75, 76
Chlorides, tests for, 149

tests with, 31, 74, 78
Chromates, properties of, 101

tests for, 150
Chromium, analysis for, 106, 110

properties and compounds of, 100

reactions of, 108, 131
Cinchonine, test for bismuth with,

91, 94
Cobalt, analysis for, 106, 110

properties and compounds of, 104,
105

reactions of, 110, 132
Cobalt uranyl acetate, sodium test

with, 134
Colloidal solutions, nature of, 34

precipitation of, 34
Common-ion effect, 23, 32
Complex ions, formation of, 37

theories of, 39

as weak electrolytes, 40
Concentration, importance of, 15

units of, 15
molar, 16
normal, 16

Conductivity, 20, 22, 23
Constants, dissociation, 23

ionization, 23, 182

solubility product, 32, 181
Coordination number, 39
Copper, analysis for, 88, 93

properties and compounds of, 82

reactions of, 90, 126
Coprecipitation, 65
Cupric compounds, 82
Cupron, 69

(See also a-benzoinoxime)
Cuprous compounds, 82
Cyanides, complex ions of, 37, 40, 90

tests for, 147

INDEX

195

Debye -Hiickel, theory of, 62
Delivery tubes, capillary, 12
Bichromates, nature of, 114

tests for, 150
Dielectric constant, 60
Dimethylglyoxime, for bismuth, 127
for nickel, 110, 113
reactions of, 70, 112, 113, 127, 130
Diphenylamiiie, test for nitrates

with, 154
Diphenylcarbazide reagent, 80, 94,

127
Diphenylthiocarbazone as a reagent,

125, 126, 128, 130, 132
Dithizon, 125

(See also Diphenylthiocarba-
zone)

Droppers, reagent, construction of,
12

Electrolytes, additivity principle of,
21

characteristics of, 20

strong, 22, 24

weak, 22, 26

Equations, analytical, for anions,
156-157

for cations, 136-139

ionic, 22

oxidation-reduction, 44, 47
Equilibria, ionic, 23

law of mass action and, 23

principle of Le Chatelier and, 23

solubility and, 30
Equilibrium constant, 19
Equivalent weights, gram, 16
Ethyienediamine, 133
Evaporations, 6

Ferric compounds, 99
Ferricyanides, tests for, 148

Ferricyanides, tests with, 130
Ferrocyanides, tests for, 149

tests with, 90, 95, 107, 112
Ferrous compounds, 98
Filtrations, centrifuge, 4

through cotton, 5
Flarne tests, 122, 123, 129
Fluorides, tests for, 151
Fusion, mixture for, 73, 114

Gallic acid, test for mercury with,

82, 89
Gases, tests for, 7

pressure and solubility product

of, 33

Group separation of cations, dis-
cussion of, 71, 73
Gutzeit test for arsenic, 92, 96

Half-cells, 45

Hydrogen electrode, 46

Hydrogen sulfide generator, acid, 10

"Aitchtuess" type of, 11

paraffin-sulfur type of, 11
Hydrolysis, effect of temperature on,

importance in analysis, 27
types of salts undergoing, 27
Hydronium ions, 59
8-hydroxyquinoline, for magnesium,

135
Hyposulfitc, test for cobalt with, 132

Inclusion, 65

Instability constants, 40

Iodides, tests for, 149

Ion-electron method, 49

Ionic strength, 64

lonization, and dielectric constant, 60
theories of, 21, 23, 61, 64
theory of complete, 22, 24, 61

196

INDEX

lonization constants, 23

table of, 182

Ions, reactions between, 22
Iron, analysis for, 106, 110

properties and compounds of, 97,
98

reactions of, 107, 130
Isoelectric point, 42
Isomorphous compounds, applica-
tions of, 67

definition of, 66

Lanthanum nitrate, test with, for

acetates, 155
Lead, analysis for, 78, 81

properties and compounds of, 76,

126

reactions of, 79
Le Chatelier, principle of, 19

application of, to weak electro-
lytes, 23
Literature, references to, 176

Magnesia mixture, 85, 152
Magnesium, analysis for, 121, 123

properties arid compounds of, 119

reactions of, 121, 135
Manganates, 103
Manganese, analysis for, 106, 110

properties and compounds of, 102

reactions of, 109, 133
Marsh test, 85
Mass action, law of, 19

application of, to slightly soluble

electrolytes, 32
to weak electrolytes, 23
Mercury, analysis for, 78, 81, 88, 93

properties and compounds of, 76,
125

reactions of, 80, 89
Metals, analysis of, 71
Microbeakers, 8, 183
Mixed crystals, 66

Mol, meaning of, 16
Molar solutions, 16
Molecular weights, gram, 16
Molybdate test, for arsenic, 129
for phosphate, 153

Nessler's solution, for ammonium,

135

Neutral solutions, H* and OH~ in,
26

pH of, 29

Neutralization, meaning of, 26
Nickel, analysis for, 106, 110

properties and compounds of, 104

reactions of, 113, 133
Nitrates, tests for, 154
Nitrites, tests for, 146
a-Nitroso-0-naphthol, 69, 112, 113
Normal solutions, 16
Normality, calculations involving,
17

dilution and, 17

volumetric reactions and, 18
Notes on analytical procedure, 163

Organic reagents, 67
linkages in, 67

types of compounds used as, 68
Oxidation and reduction, 43
electron changes in, 44, 46
equations, balancing of, 47
potentials, determination of, 45
table of, 179
uses of, 46
Oxine, 135

(tfee also 8-hydroxyquinoline)

Passivity, 98, 99
Permanganates, reactions of, 103

tests for, 155
pH, meaning of, 29

INDEX

197

Phosphates, tests for, 153
Pipettes, capillary, construction of,

Polar, distinction in meaning of, 60
Potassium, analysis for, 121, 123
properties and compounds of, 120
reactions of, 122, 134
Potential, oxidation and reduction,

46, 179

" Potential Hydrogen," 29
Precipitates, dissolving, 35

handling of, 5
Precipitation, test for completeness

of, 72
Preliminary experiments, Group I,

Group II, 89
Group III, 107
Group IV, 116
Group V, 121
Problems, list- of, 15&-162
types of, 50

involving solubility products, 55
involving weak electrolytes, 5J
Pyridine, 109, 132
Pyroantimoniate, for potassium, 134

Questions, cation, 140
R

Reactions, completeness of, 46
Reagent blocks, 13, 183
Reagent containers, 13, 183
Reagents, preparation, of anion, 188

of cation, 185

References to literature, 176
Rhodanine, reaction of, 68, 79, 81
Rhodamine B, test for antimony, 130
Rods, stirring, construction of, 13

Salt effect, 62

Samples, treatment of solid 73 142

Saturated solution, 30
Semi-micro technique, 3

quantities used in, 3

steps in, 3

uses of, 3

Separation of ions, efficiency of, 58
Silicates, tests for, 151
Silver, analysis for, 78, 81

properties and compounds of, 76,
126

reactions of, 78, 125
Sodium, analysis for, 121, 124

properties and compounds of, 120

reactions of, 123, 134
Solubility, 30
Solubility product, limitations on, 33

principle of, 32
Solubility product constants, 32

table of, 181
Solutions, colloidal, 34

molar, 16

normal, 16

saturated, 30

supersaturated, 33

test, 190

unknown, 190-191
Solvation of ions, 59
Solvents, nonaqueous, 60
Spot tests, backgrounds for, 7

methods of making, 7
Stability constants, 41
Stannic acid, meta, 43, 88
Stannic compounds, 88
Stannites, preparation of, 43, 93

reducing properties of, 83, 91, 94
Stanrious compounds, 87
Stibine, 86
Strontium, analysis for, 115, 118

properties and compounds of, 115

reactions of, 117
Sulfates, tests for, 150
Sulfidcs, tests for, 146
Sulfites, separation of, from sul fides
and thiosulfates, 172

tests for, 146
Supersaturated solution, 33

198

INDEX

Tetraborates, tests for, 154
Tetramethyldiamino-diphenyl-

mcthane, 126, 131
Thio-eomplexes, 38, 78, 84, 86-88
Thiocyanates, tests for, 148

tests with, 107, 109, 112, 113, 120
Thiosinamine, test for cadmium

with, 91, 95

Thiosulfate, tests for, 146
antimony with, 93, 96
nickel with, 133
Tin, analysis for, 88, 93, 96

properties and compounds of, 87
reactions of, 97, 129

Unknowns, sample size of, 75
V

Valence, coordinate, 39, 67, 69
primary, 39

Valence, Werner's theory of, 39
Valence-electron method, 47

Water, as an electrolyte, 26

H 4 and OH" concentrations in, 28
and hydrolysis, 26
and pH values, 27

Werner's theories of complex ions,
39

Zinc, amphotcric nature of, 41, 103,

109

analysis for, 106, 110
properties and compounds of, 103
reactions of, 109, 132
Zinc purpurate as a reagent, 125
Zinc uranyl acetate, sodium test
with, 123, 124

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