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ELECTROHETRIC TITRATION

CF

OXIDIZING AND REDUCING SUBSTANCES

 

A Thesis Submitted to the Faculty of the
Michigan State College of Agriculture and
Applied Science in Partial Fulfillment of
the Requirements for the Degree of Kaster

of Science

 

BY
Charles Henry Spurway

1926

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-A_—————-—__4
_,

Dwight T. Ewing, Ph.D.
Associate Professor of Chemistry

Major Professor

II

II.
III.

IV.

V.
VI.

VII.

III

TABLE OF COHTENTS

Introduction

General Theory

Methods and Apparatus

Special Points Taken from the Literature of
the Subject Concerning the Electrometric
Titration of Several Chemical Substances
Electrometric Titration Curves

Curve Data and Discussion

Bibliography

I

Introduction

Electrometric titration methods of chemical analysis,
particularly for the determination of small quantities of
several of the elements combined in some oxidisable or
reduceable form, are coming into general use in laboratories
devoted to teaching, research. and the process control of
many industries. In addition to the work already done on
these methods, there are still many possibilities and some
Opportunities for applying them to the qualitative and also
to the quantitative determination of other elements, and to
the control of other industrial processes. These methods are
valuable to chemists because of their case of Operation,
accuracy, and time saving qualities. Not only are electrometric
titration methods serviceable for analytical procedures, but
because they deal with quantity and movement of electrons, or
charges of negative electricity, they are applicable to
problems involving chemical hypotheses, and theories relating

to the state or preperties of matter in solution.

As is-the case with practically any division of a general
science. the literature dealing with the theory and the problems
of electrometric titration is scattered through a large number
of scientific and industrial Journals; and hence. is buried
from the sight of the average student and investigator in

chemical science largely because of the time and labor found

I‘D

necessary to assemble, arrange, and mentally digest the
contents of the many papers extant on the subject. Collecting
and cataloging the knowledge on any subject, interesting in
itself,and also applicable to the study of various scientific
problems is considered advisable. because by this means such
knowledge is placed within easy reach of students and other
interested persons; and in addition. complete bibliOgraphies

of the subject may be maintained.

This thesis has for its object, then. four main divisions
of effort: (1) An exposition of the general theoretical and
practical principles upon which electrometric titration
methods are based, more particularly with respect to oxidation
and reduction effects; (2) A description of the several
analytical methods now in use; (3) A search for possible new
correlations that might be exposed by a study of the titration
curves found in the literature; and, (4) A complete, up-to-date

bibliography of the subject.

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II
General Theory

The theoretical basis for electrometric methods dealing
with oxidation and reduction phenomena lies chiefly in the
work of the Swedish physicist Arrhenius tn establishing the
electrolytic dissociation theory of substances in solution (177;);
and lies also in the second law of Faraday which states that
the valence of an ion is a function of the number of electrical
charges it carries C127).Acids, bases. and salts when in water
solution are dissociated into ions (electrically charged
particles) and the number of these electrical charges are
preportional to the valence of the elements involved. This

phenomenon may be graphically illustrated as follows:

NaCl -—+ Na+4k Cl- Monovalent dissociation.
IgSO4 —+ lg”+ SOZ- Divalent dissociation.
A1013 -+ A1”; 013‘" Trivalent dissociation.
SnCl4 ~€P antF 012'"! Tetravalent dissociation.

Some of the chemical elements may exist in combination
in more than one state of valence, and dissociate electrolytically
in accord with this valence state, thus, -

Ferrous Sulfate - Pe804 «eh-Pektk»80;"’

++F

Ferric Sulfate - Fe2{304)3 —e— 3’9 4‘ 530;“-

We can change the valence of iron and some of the other
chemical elements by adding or subtracting electrical charges -

electrons. The valence of iron from two in the ferrous state

' V
e

‘ Q

.

4—..—
o
c. -.

can be raised to three in the ferric state by subtracting one
electron(adding a positive charge), or from the ferric to the
ferrous state by adding one electron(subtracting a positive
charge).

Pe'H' -- Q = Fe+*+

Fem-q- @ :: Fe++

'hen an electron, or electrons, are taken from an element
in solution the.element is said to be oxidized; and when the
Opposite change takes place, the element is said to be reduced.
These changes are accompanied by electrical effects which can
be observed and measured. The common method of measuring these
electrical effects is by comparing them with a normal or a
standard electrode which gives a definite electromotive force.
The calomel standard cell is the one most commonly used.

If an electrode of a noble metal such as a piece of platinum
foil or wire is placed separately into solutions of ferric and
ferrous ions, a definite potential difference will be set up
between the electrode and the ions in each solution, and the
value of this potential difference will depend upon the
concentration of the iron ions of these solutions. There will
be a tendency for the ferrous ions to give up electrons to the
platinum wire, charging it negatively, and for the ferric ions
to take up electrons from the platinum wire charging it positively.
But, if the platinum wire dips into one solution containing a

mixture of ferric and ferrous ions, and is connected externally

with a normal calomel electrode which in turn makes contact

with the ferrous-ferric solution by means of a salt arm making

a liquid-liquid Junction, a potential difference is established

which is determined by the ratio of the ferrous to the ferric

ions in the solution. In the measuring process this potential

difference is compared to that of the normal calomel cell. The

relationship is expressed by Peter's equation.(6).

RT (reduction product)
a? (oxidation product)

observed potential difference

a constant applying to this particular case
(reduction product)

and is equal to B when 4 ------------------ =-l
(oxidation product)

the gas constant =volts X coulombs

absolute temperature

valence of the metal

Faraday =-. 96,540 coulombs

When applied to the case under consideration, the formula

becomes,-

Subtracting numerical values and changing to common

logarithims, we have,-

E

-
-

0.46 4- 0.058 leg I .....

E, or the potential difference, then becomes a measure of

the relative concentration of the ferric and ferrous ions.

As the ratio of these ionic concentrations gradually changes,
due to the addition of oxidizing or reducing reagents, B will
also show a gradual and corresponding change in value until
one of the ionic concentrations becomes vanishingly small as
compared with the other, when E will show an abrupt change of
comparatively great magnitude which may be taken as the end-
point of the reaction. The theory may-be applied to all
oxidation and reduction reactions, but the constant, C, is
different in value for the different chemical elements that
may be oxidized or reduced. Some of these constants are given

by LeBlanc (17?).

The potential at the beginning and end of the reaction
depends upon the nature and concentrations of other ions which
may be present as well as the concentrations of the ferrous or
ferric ions, but the changes in potential due to the addition
of reagents are always slight in magnitude until the end-point

of the reaction is reached when a great change takes place.

It is possible to determine quantitatively two or more
substances present in the one solution providing they can be
oxidized or reduced by the same reagent, and that their oxidation
or reduction potentials are far enough apart in values to give

separate end-points within the voltage range obtainable.

The oxidation or reduction reaction must progress quite

rapidly for the electrometric method to be of value.

0‘ ‘9.

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III
METHODS

A classification of the different systems used in the
methods of electrometric titration has been made by Kolthoff
(Chem. Weekblad 50,559(1920). Three systems are given by
Kelthoff, to which a fourth may be added. 1. The common
system, consisting of a stable electrode, unattacked in an
acid solution, used with a constant half-cell such as the
calomel cell or a similar electrode. 2. The Pinkhof system
(Chem. Weekblad 16,1163-7,1919) in which a compensating
electrode replaces the calomel half-cell, the potential of
the compensating electrode being equal to the stable electrode
at the end of the titration. A reversal of polarity marks
the end-point in this case. The common disadvantages of this
system is that each titration requires a different compensating
electrode, and the concentration of substances foreign to the
titration may interfere with the compensating effects. 3. The
Treadwell system (Helv. Chim. Acta 2,680,1919) is composed of
a half-cell replacing the standard electrode which contains an
electrolyte having the same composition at the titrated
solution at the end-point of the reaction. This system has the
same disadvantages as the Pinkhof system being specific for
each titration. 4. The bimetallic system of Willard and
Fenwick (J.A.C.S. 44,2504,l922) may be mentioned as a fourth
system of electrometric titration. In this system the

standard half-cell is replaced with an inactive metal or a

metal-alloy electrode which shows no great change in potential
at the end-point of the reaction. Some of the metals used are
tungsten, and alloys of the platinum group of metals. Another
form of this system consists of two inactive metallic electrodes,
one of platinum, upon which in imposed an e.m.f. equal to the

hydrogen overvoltage of the platinum electrode.

Some data regarding the specific use of these systems are
discussed in connection with the methods used for the several
elements that may be determined quantitatively by means of

electrometric titration.

ESSENTIAL APPARATUS

 

I —,II) .

(a)

 

 

 

 

 

 

‘3? I—

C.

 

 

 

Figure 1. Diagram of the apparatus.

P - Potentiometer. A Leeds and Northrup Students'
potentiometer is suitable for a great variety of
determinations.

Ba - Battery. Two dry cells, or a four volt accumulator
may be used.

R - Resistance for adjusting the electric current in

in the (a) circuit. A plug resistance of 11, 110
ohms capacity works well with two dry cells.

3 - A current reversing, double pole, double throw
switch. The reversing feature of this switch aids
in making voltage readings if the voltage should become
negative during a titration. By reversing the (a)
circuit current, the potentiometer may be read
directly.

G - Galvanometer. A sensitive, dead beat, galvanometer
is very good for the purpose, although a too sensitive
galvanometer is not to be desired for some work.

3.0. - A standard cell, used for calibrating the potentio-
meter as a diredt reading voltmeter.

r - A high resistance used to protect the standard cell.
This high resistance may also be used in the titrating
cell circuit to prevent drawing a current from it.

C - Titrating cell. These titrating cells may be of a
variety of forms depending upon the nature of the work
to be done.

K1 and K2 - are key-switches used to momentarily close
their respective circuits when calibrating or reading

the potentiometer.

The apparatus illustrated and described belongs to the
commonly usedwcompensating system devised by Poggendorf for
the measurement of voltages. It consists, essentially, of

two electrical circuits, one of which (a) may be set at a

 

10

known value through the potentiometer, and the value of the
other unknown circuit (b) measured by comparison with the

circuit (a).

Titrating Cells.- Various forms of titrating cells are
in general use. If in any particular case it is unnecessary
to protect the solutions being titrated from oxidation by the
oxygen of the air, or from possible contamination by carbon
dioxide gas, or any other laboratory gas or fume; an epen
beaker, or other cpen dish of suitable size and shape is all
that will be required. On the other hand, a closed container,
or cell, is necessary for the titrating of solutions that
oxidize readily in the air, or that require elimination of
the adsorbed oxygen by eleutration with an inert gas passed
through the solution preliminary to the titration, or that
need to be protected from gas injury by maintaining constantly
an atmosphere of an inert gas over them during the titration
procedure. Two common forms of closed containers for electro-
titration work are; 1. the titrating head described by
Hostetter and Roberts (45), and 2. the closed container used
by Ewing and Eldridge (75). Burette tips, electrodes, stirrers,
or the delivery points of other pieces of apparatus, or gas
inlets and outlets are sealed into the cells in some suitable

manner a

In the earlier electrometric titration work, a capillary

electrometer was used in place of the galvonometer; and for

11

end-point determination only, a common sliding resistance,
of the proper value, took the place of the potentiometer. In
this case, the end~point readings were taken by constantly
adjusting the resistance to hold the electrometer at zero as
the titration progressed, and the end-point was shown by a
marked deflection of the electrometer. This arrangement is
now used sometimes in laboratories where rapid work is
required, but the galvanometer has practically replaced the
capillary electrometer as a null-point instrument in the
electrometric titration work. Titration curves may be
plotted from the readings taken in the above manner when the
resistance is first calibrated and provided with a scale for

interpreting the results.

Calibration of the potentiometer.- In calibrating the
potentiometer to read as a voltmeter, the potentiometer is
first set at its voltage reading corresponding to the exact
voltage of the standard cell. The resistance R is then
adjusted until by slight tapping of the key K2 the galvanometer
shows no deflection. Practically no current should be taken
from the standard cell, hence, for accurate work, a rough
standard cell should first be used while the preliminary
resistance adjustments are being made, and the final adjust-
ment then made with the accurate standard cell. For accurate
work this calibrating of the potentiometer should precede each
titration, especially if titration curves are to be plotted
from the voltage readings.

12

Special kinds of electrometric titrating apparatus made
in compact, portable forms and adapted to different types

of analytical work are new on the market.

The vacuum tube generally used in radio sets has also
been used in place of the potentiometer for making e1ectro~
metric titrations. For references see Geode(93), Calhane a

Cushing(134b), Treadwe11(l68), and Gardner(176).

Method of making electrometric titrations.- The actual
method of Operation used will depend somewhat upon the nature
of the work; certain precautions being necessary in most cases.
It is the plan of the writer to give only a general outline
of the methods used at the present time.

The common technis adapted by investigators for making
electrometric titrations is to place a known quantity of a
certain oxidisable or reducible solution into the titrating
vessel, usually in an acid solution, although some titrations
require an alkaline solution, and then titrating this solution
with a reducing or oxidising reagent, as the case may require,
and the voltage readings taken in a manner that allows for the
plotting of a curve with the voltage readings plotted against
cubic centimeters of the reagent used, This procedure is known
as direct titration. An indirect method of titration is used
in cases where a sharp end-point can not be obtained by means
of the direct method. The indirect method consists of adding

an excess of reagent to the solution to be titrated and then

13

titrating this excess with a reagent that does give a sharp
end-point with the first solution used in excess. In making
a first study of a reaction, known solutions are used in
order to prove the accuracy or advantage of the method before
the method is applied to unknown solutions. As stated before,
if the end-point alone is desired in a titration it is only
necessary to determine the point in the titration where the
greatest deflection of the galvanometer takes place. It is
advisable to know just about the quantity of titrating reagent
that will be required for a determination, and as the end-
point is approached, tO add the reagent, drOp by drOp, until
the end-point is exactly reached.

It must be remembered that in making titrations by the
electrometric method the voltage changes are caused by the
changing concentrations Of the ions in the solutions, and
that any ionic concentration change will cause a corresponding
change in voltage. This being the case, the Operator must
know, in order to do accurate work, that the only concentra-
tion changes taking place are due to the additions of the
reagent used, and not due to changes at the electrodes, or
to oxidation or reduction effects not due to the reagent added,
or to side reactions that may take place under the conditions
of the titration. This means that the electrodes should not be
attacked by the original solutions, or by the products formed
during the reaction; and that the oxygen of the air must be

excluded when easily oxidiaable solutions are being used; and

14

also, that possible side reactions do not interfere with
the main determination. It is also necesarry to know that
the reaction is completed when the end-point deflection

occurs .

The common method of protecting a titration against air
oxidation is to bubble an inert gas through the solution
preliminary to the titration Operation in order to remove
absorbed oxygen, and then during the titration to maintain
an atmosphere of the gas over the solution in the titrating
cell. Carbon dioxide, hydrOgen and nitrogen are the inert
gases ordinarly used. However, bubbles of the gas should
not be allowed to adhere to the metallic electrode, as this
tends to form a gas electrode and causes a change in e.m.f.
of the cell; hence, the adsorbed oxygen, or other gas, should
be removed before the metallic electrode is inserted into the

solution.

The metallic electrode generaly used is platinum. Some
workers prefer platinum foil of a certain size for the
metallic electrode, while others have found platinum wire
best, at least for certain determinations. The shape and
size of the platinum electrode is probably specific for some
titrations, but this point seems to have been little investi-
gated. These platinum electrodes must be kept clean and bright
by washing in acids, cleaning solutions, ignition, scouring,

or by electrolyzing deposited substances from them; the most

15

suitable method in any case depends upon the nature of the

deposit on the electrode.

Mechanical stirring of the solution during the titration

process is practically always need.

Temperature control is Lunimportant for reduction or
oxidation reactions that take place rapidly at room tempera-
ture. Where heating is required to hasten the reaction, the
usual method is to heat the solution to the required tempera-
ture and then titrate rapidly for the end-point. Some workers
have found, however, that by cooling the solutions to a low
temperature, better results could be obtained. The prOper
temperature for a reaction, at least within a certain range,
is important for reactions that require heating in order to

increase the speed at which they take place.

The advantage of using the electrometric end-point in a
titration method is its high degree of accuracy, and to make
the fullest use of the delicacy of these methods, all standard
solutions used for making electrometric titrations should be
standardized by the most accurate methods possible for each
particular solution. When a high degree of accuracy can be
attained in the standardization of the titrating reagents, the
electrometric titration method is exceedingly accurate, and
very small quantities of a substance may be determined by

means of it if sharp end-points may be obtained.

16

So great is the accuracy attainable as compared with
the common volumetric analysis, by means Of these methods
that many workers in this field recommend that the standard
reagents be made up by weight instead of by volume, and that
the titrating additions of reagents be made by weight also
instead Of being measured from a burst. The degree of accuracy
required or desired will probably be a determining factor in
this connection with many Operators Of the electrometric

titration methods.

Another important factOr in electrometric titration is
the kind of acid used to acidify the solution being titrated,
and the degree or quantity of this acidity necessary for best
results. This factoris specific for each titration, within
certain limits, and when not known, or can not be determined
from the literature on the subject, it must be experimentally

determined.

Breaking the circuit while taking the voltage readings is
sometimes done in order to prevent deposition of metals on

the electrodes.

The magnitude of the end-point voltage deflection depends
upon the intensity Of the reduction or oxidation processes

and may be used as a measure of them.

0'!

17

IV
SPECIAL POIKTS TAKEK F? n EYE TERATUFE OF TPE

SUBJECT COKCERE HG T3? 333 TROLETRIC TITRLTICK OF

SEVERAL C3EMICAL ELEAEKTS

In the folloWing discussion the several elements and
substances that have been investigated by means Of the
electrometric titration methods are listed alphabetically.
The numbers in parentheses given in connection with the
subjects and paragraphs refer to the numbers of the refer-
ences in the bibliography appended to this thesis. The

"*" were not found in the Michigan

references marked with a
State College library (Decehber 51, 1925.) This list of
subjects with the reference numbers will serve as a cross

reference to the bibliography.

Antimony.- (107), (147)

Antimonous sul hate mav be titrated with potassium
p c} A.

bromate in various concentrations of hydrochloric acid, 3-20
percent. (107) (C.A. 17, 1604)

Antimony in the form of antimonous chloride may be tit-
rated with potassium dichronate in the.presence of hydro-
chloric acid, 100 cc. of the solution containing 15 cc. of
hydrochloric acid (d. 1.18), at room temperature, and with
constant stirring. lntimonous chloride and stannous chloride
may be titrated together in one sample, but mercuric chloride

should be present in the solution in excess when antimony is

titrated in the presence of tin. (147) (0.3. 19, 451)

18
Arsenic.- (137), (151)
Arsenic lay also be titrated by means of potassium
bronate in a 5-20 percent hydrochloric acid solution. ’107)
(C.A. 17, 1604)
A sodium arsenite solution, 0.0222d, was prepared as
recommended by the AmeriCan Society of T sting haterials for
steel analysis, and titrated against an iodide solution which
had been standardized against potassium bromate. The sodium
arsenite was titrated electrometrically with a potassium
-permanganate solution, 0.115N, standardized by means Of
sodium oxalate. 25 cc. of the permanganate solution was
titrated in 100 cc. of nitric acid (d. 1.13) with the arsen-
ite solution.
A tungsten electrode was used in place of the calomel
electrode, as in the presence of potassium chloride in the
solution, the theoretical end-point was not obtained. The

potassium chloride causes a more complete reduction of the

This reaction was also carried out in a 2.5K sulfuric
acid solution with 0.0964N sodium arsenite.

in the nitric acid solution the'manganese is reduced
from a valence of 7 to an average valence of 3.5.-’65n041+
lng.»‘-.s0§+l5 H§+2L~flvff4il$11msogf But in the sulfuric acid
solution, the valence is reduced to an average of 2.5;- Mn04?
9 .u. is 054 14 ngzfliv'ii'sffiiemsog.‘

In the electrometric titration of arsenic attention

should be paid to details especially acid concentration,

19
teaperature, presence of halide or mangan(us salt, and to
standardize the solutions used under the same conditions as
they are to be used in the method of analysis. (151)

(C.A. 17, 2687)

A20 dyes.- (94)

The azo dyes were reduced by means of a 0.25N titanous
chloride solution in excess and the excess titrated with a
0.05N ferric alum solution. The titrations were carried
out in an atmosphere of carbon dioxide. Thiocyanate was not

satisfactory as an outside indicator, because of the presence

f carton dioxide, and becaus the solutions the selves were

0

sometimes either pink, blue, or green in color. The reaction
flask consisted of a 250 cc. Tyrex extraction flask with a

a ”1
ne salt-

E-hole stepper to accomodate a platinum electrode, t
arm of the calomel half-cell, and a buret.
Trocedure.- 0.5 to 1.0 gram, f the finely pondered dye,

sufficient in quantity to require 50 to 40 cc. of the titanous

chloride solution for reduction, was placed into the reaction

'

I‘J

flask; and 25 cc. of distilled water was then added, the flas-
covered loosely, and placed on a steam bath for about 10 min-
utes to dissolve or suften the particles. next, 25 cc. of

40 percent sulphuric acid was added, the flask stoppered, and
a current of carbon dioxide was passed through for five min;
utes. Then the required amount of the titanous chloride
solution was added, at least five co. more tha was required
for reduction. The liouid was then boiled for five minutes
and then cooled to 300 C. in a water bath, and titrated. (94)

{0...}... 16, 11:):)‘.

2O

Bromine.-l(ll)*, (61), (105), (116), (120)*

A known weight of potassium.iodide solution was titrated
with a potassium bromate solution in the presence of enoush
su furic acid to make the solution about 2H at the end of the
experiment, and until about 10 percent of the iodide remained
undecomposed. Then the voltage became constant, the titration
was completed with potassium permanganate. Experiments in-
dicated that in this titration the presence of small anounts
of chlorate nad no influence. The iodide solution was also
successfully titrated with the bromate solution. The last
drops of the bromate solution were added very slowly to
avoid running by the end-point, and the whole titration re-
quired at least 15 minutes for end-point, and much lonser to

\-

acquire the data for plotting a curve. Hydrochloric acid
may not be present in Quantities in excess of 0.2N. (61)
(0.3. 15, 5793)
Dutoit and von Weisse used the electrometric method for
the titration and separation of the halides, and recommended
precipitation With a silver salt or some other suitable
precipitant. They call attention to the suitability of the
method for determining the halides in very dilute concen-
trations as 10-20 milligrams of chloride per liter, or
smaller quantities of bromide or iodide. Two end-points are
obtained in a mixture of the three halides, the first one is.

J.

the iodide prec'pitation, and the second one occu s when all

r“

the halides are precipitated. -heoe authors claim.t?at

21

iodide may be determined in this way in the presence of
a great excess of bromide or chloride. (11)
(C.A. 6, 722)

Eronate may also be titrated with titanous chloride
at a concentration of 0.05N x 1.4663, and protecting the
solution by means of carbon dioxide gas. 50 cc. of a sol-
ution of potassium bromate, 0.05N, in 200 cc. of boiled
water, and 10 cc. of ION sulfuric acid were used for the
titration. Uydrochloric acid may be used in place of the
sulphuric acid, but should not exceed a concentration of 0.5N

‘

at the end of the titration. The chanse of bromide in hydro-
cyanic acid solution to bromine cyanide by titrating with
potassium permanganate is too slow for the bimetallic elec-
trode system to be used; but the end-point in this reaction
may be deternined by means of the usual mononetallic system
'if the titration curve is plotted. This oxidation reouires
the theoretiCal equivalent of oxygen, and the oxalate factor
of the permanganate may be used. (105‘, (116),(C.A. 17, 397)
(C.A. 17, sseex

By means of the mercury electrode, chlorides and brom-
ides may be titrated electrometrically with mercurous nit-
rate, but not in the presence of each oth r. (126)

(C.A. 18, 1793)

COPPGT-- (10)*. (ee)*, ( 20)‘. (157)*

W?

Dutiot and von .eisse made a detailed study of the prin-
ciples of concentration cells as applicable to voluuetric

analysis by precipitation. They produced polarization in

22

the cell by means of a feeble, imposed current and used
revolving electrodes to prevent this current from setting
Up concentration differences. The several points studied
are: (1) Speed of stirring. .2) Effect of foreign substances.
(5) Effect of dilution. (4) Influence of the auxillary current.
(5) Duration of the titration. Cepper was titrated with a
solution of potassium ferr<cyanide, and also with a solution
of sulfide. Sodium or potassium sulfide was found to be the
most satisfactory precipitant. Copper could be determined
with an accuracy of Q.l-O.2 milligram per liter. (10)
(0.3. 6, 722)

COpper may be accurately determined by electrometric
titration by means of reduction with titanous chloride in
excess in hydrochloric acid solution, and titrating the ex-
cess of titanous chloride with potassium bromate or potassium
chromate in an atmosphere of carbon dioxide. The reaction

is slow at room temperature, but instantaneous at BMC C.

k

a

4-8 percent of hydrochloric acid should be present in the
solution. If iron is present, it is determined with the

cepper. (88) (C.A. 17, 702)

Copper may be titrated simultaneously with silver. The
silver is first titrated with potassium cyanide, and the
silver cyanide removed from the solution by filtration, and
the copper titrated vith the potassium cyanide solution. This
is a convenient method for the separation of copper and silver.

(120) (C.d. 18, 640)

25

Cupric salts alone, or when mixed sith_ferric salts,
or both in the presence of antimony, may be titrated by

means of titanous chloride. (157), (0.3. 19, 50)

Cyanides.- (12)*, (36), (4e)*, (7e)*, (77)*, (783*,
(116), (126)*, (1451*

Treadwell titrated potassiur cyanide in the presence of
potassium ferrocyanide with silver nitrate by first adding to
the titrated solution 0.1 gram of potassium iodide in a 0.1N
solution. The solution was slightly alkaline. Addition of
ammonia is not necessary. Increasing the amount of potassium
iodide overcame the disturbing influence of sodium hyposulfite.
The influence of an excess of ammonia, sodium hyposulfite, or
potassium ferrocyanide was fould to be due to the formation of
complex ions and the lowering of the concentration of the
silver ion below that required for the precipitation of the
silver yanide. (12) (0.3. 5, 5591)

Muller states that all cations can be electrometrically
determined by deans of a solution having an ion which will
produce a constant insoluble compound. Potassium ferrocyanide
may be used to determine zinc or lead. All ferrocyanide
solutions contain some ferricyanide. The low solubility of
zinc ferrocyanide makes the value ferro ion-ferri ion large
so a small, constant amount of potassium ferricyanide is
added to the solution to make the ferri ion portion independ-
ent of the other. The temperature used was 750. Several

metals in the sane solution could not be determined by the

24

method because of the variation in the composition of most
ferrocyanides. Cadmium did not interfere with this titra-
tion. A calculated table in which a cc. of potassium ferro-
cyanide, b number of ohms required for compensation in the
determination of the potential against a iormal calomel
electrode with a capillary electrometer used as a null-point
instrument; and taking the number of cc. of titrating solu-
tion correSponding to the highest value of da/db, was found
to be more convenient and more accurate than the graphic

method of plotting the titration curve. (40) (C.A. 14, 2805)

Two end-points are obtained when cyanogen ion is titrated
with silver ion; the first corresponds to the addition of
one-half silver ion for each cyanogen ion present, and the
second for equal amounts and when silver cyanide is precip-
itated. It is possible to titrate cyanogen ion in a halide
mixture and determine each constituent. With iodine and
cyanogen three end-points are obtained; the first when one-

half silver has been added for each cranc-en ion the next
9

,7

when -ilver iodide is precipitated, and the third when both
cyanogen a.d iodide are precipitated. Two end-points were'
found in the presence of chlorine or bromine. The first
corresponds to one-half the silver as above, the second to
the precipitation of silver cyanide and silver chloride or
silver bromide as the case may be. With a mixture of iodine,

bromine and cyanoren, three end-points were obtained. The

first one on one-half the cyanogen, second on the precipitation

25

of silver iodide, and the third on the precipitation of all
the ions. With bromine, chlorine, and cyanogen, only two

end-points are obtained so that only the cyanogen and total

halogen content may be determined. (76) (C.A. 16, 2818)

Kulthoff titrated potassium ferrocyanide With potass'um
permanganate in concentrations of more than 0.02 mol in the
presence of sufficient sulfuric or hydrochloric acid to prevent
the precipitation of potassium or manganese ferricyanide, and
by comparing the color of the solution at the end-point with'

a potassium ferricyanide solution. (77) (C.A. 16, 5281)

The titration of ferrocyanide with potassium perman.an-
ate gives better results when the end-point is determined
electrometrically than when it is obtained by the chanre in
color of the solution. (78) (0.3. 16, 5601)

Cyanide can be titrated by means of mercuric chloride.
(126) ' ' (C.A. 18, 1799)

Ferricyanide can be direCtly titrated with potassium
iodide. If the ferricyanide solution is run into a known
quantity of potassium iodide and encugh zinc is added to
form the double compound, zinc-potassium ferrocyanide, the
end-point is obtained more promptly. (143) (C.A. 18, 3020)

About one gram of potassium ferrocyanide in 250 cc. of
water and 2.5 to 5 cc. of sulfuric acid (sp. yr. 1.58), may

be titrated With 0.05N potassium permanganate. The potassium

J

permanganate should be accuratelv standardized with sodium

oxalate. Ferricyanide does not interfere with this titration.

26

The permanganate should be added to the ferrocyanide with
constant stirring. is the end-point is approached, the
permanganate should be added drop-by-drop, and sufficien
time allowed for the completion of the reaction. Chloride
in amount equivalent to one gram of sodium chloride does not
interfere with the reaction. Any salts, "hich under the con-
ditions of the reaction, produce a precipitate with either
ferro- or ferri yanide do interfere with the reaction. (56)
V (C.A. 14, 39)

Hendrixson titrated a solution of pure potassium ferri-
cyanide of a concentration of 0.05N x 0.72, standardized by
means of the iodine method, with titanous chloride, and he
claims a high degree of accuracy for this determination. (116)

(C.A. 17, 5506)

Chromium.- (14), (17), (20), (25), (49)*, (61), (63)*,
(64)*, (663*, (79), (1441*

Dichromate ion was successfully titrated by means of
ferrous ion using a platinum electrode and a calomel half-
cell. The author used a plunger consisting of a glass rod
which passed through a rubber stopper inserted in the top
of the calomel electrode cell to force the potassium chlor-
ide through the salt arm of the cell and flush it out before
each potentiometer reading. A wheatstone bridge potentio;

meter and a capillary electrometer were used to make the

voltage readings. The dichromate solution was titrated with

27

the ferrous solution. in titrating a dichromate solution,

the ferrous iron solution was found to be more accurate than

than a ferricyanide solution. The presence of chlorides was

found to interfere with the proper end-point. (14)(C.A. 7,5938)
Tiegel titrated a potassium dichromate solution with

ferrous sulfate using the 0stwa d-Luther, electrically con—

trolled automatic buret. The results'obtained were not of

a high degree of accuracy. (17) (0.3. 9, 898)

Kelley and Conant used the electrometric method to de-

Ho

termine the amount of chromium n steel._ They used the
method and apparatus of Forbes and Eartlet.. The chromium
was first oxidized to chromate and then titrated. Large
amounts of ferric iron in the solution affect the accuracy
of the endpoint. Temperature had little effect on the re-
action or delicacy of the end-point. The acid concentration
should be quite high, particularly in the presence of chlor-
ides. Cold solutions containing high concentrations of hydro-
chloric acid gave good results. (20) (C.A. 10, 2179)
Kelly, Adams, and Wiley combined the several pieces of
electrometric titration apparatus into a portable unit for
use in industrial laboratories. Using this portable appar-
atus to deternine chromium in steels, they claim an accuracy
of nearly 0.001 percent. The titrating reagent used was
ferrous sulfate which was standardized electrometrically by

means of potassium permanganate; but the permanganate was

standardized with sodium oxalate in the usual way. The

28

chromate may be titrated with ferrous sulfate, or the fer-
rous sulfate added to the chromate in excess, and the tit?
ration completed with permanganate and ferrous sulfate. (25)
(C.A. 11. 2567)
A method for determining chromium in steel in furnace
tests, while the steel is held in a molten condition, has
been devised. The chromium is oxidized by means of the
bismuthate method and then titrated electrometrically with
ferrous ammonium sulfate and potassium dichromate. The
method is not applicable to forged steels or those that have
been heat-treated, or when chromium carbide is set free in
more than Small amounts while obtaining the solution. (49)
(C.A. 3, 1802)
Pendrixson titrated a potassitn iodide solution with
a potassium dichromate solution in a solution of sulfuric
acid of at least 2N, by allowing sufficient tide near the
end-point for the reaction to reach completion. At least
half an hour was required. One gram of the dichromate sol-
ution set free 0.006556 grams of iodine from the iodine
solution, and the proportion of iodide to permanganate was
one to 0.999; one gram of pernanganate reouired 0.005566
grams of sodium oxalate.(61) (C.A. 15, 3799)
In the determination of chromium and vanadium in ferro-
vanadiun, these elements are oxidized by means of ammonium
persulfate, and then titrated with ferrous sulfate in a

hydrochloric acid solution. In a separate portion of the

0-

29

solution, Vanadium is oxidized to the quinquivalent condition

by means of nitric acid which does not affect the chromium,

J

and is then reduced hy titrating with ferrous sulfate. (63)

(C.A. 15, 3J55)

In the determination of chromium in ferrochromium,-the

sample is fused in an alhaline oxidizing flux, the excess of

oxidizer is removed, and the solution acidified and titrated
with ferrous sulfate. (64) . (C.A. 16, 57)
Fwing to the hydrolysis of potassium chromate, it is

not possible to neutralize chromous or chromic acid directly

towards phenolphthalein hence, potassium dichromate is not

to be recoxmcnded as a standard for acidimetry {Conpare with
-hc electroneiric methods). (66) (0.3. 16, 1056)
Bppley and Vasburgh titrated potassiun dichromate with

ferrous sulfate anC studied the effects of concentration of

{D

alts and acids on the reaction and end-point. She reverse

C+

itrstion of ferrous sulfate with potassium dichronate was

{1‘1

also studied by then. inese authors found that electrolyz-

ing the pLatinum electrode a few minutes in sulfuric or
hydrochloric acid, as anode, Was an efficient method for
cleaning it. The direct and reverse titrations, as stated
before, aoroed with each other. Dissolved air was found to
have a negligible effect on this titration. (79)(C.A.16,5851)
Working with pure amnoniuu vanadate and potassium dich-

romate these authors titrated mixtures of solutions of these

salts using standardized ferrous ammonium sulfate, and pro-

50

J

tectins the solutions with an atmosphere of carbon dioxide
gas. In the mixture only one tremk occurs in the curve
which corresponds to the known composition of the mixtrre.
These elements nay be determined in the presence of large
amounts of iron. Vanadium is deternined by a special method
(see vanadium) and the chromium estimated by difference.

(see also the abstract for the use of this method in steel

analusis. (144) (C.A. 18, 5573)

Chlorine.- (ll)*, (51), (69)*, (105), (116), (126)*,
{129)*, (132)*, (136)*, (160), (165)*, (164)*

The electrometric nethod may be applied to the deter-
mination of chlorine by precipitation with silver or some
other suitable reasent by means of which an insoluble chloride
may he formed. The method is of importance for determinins
chlorine in low concentrations, or in the presence of a large
excess of bromide or iodide. Chloride may be determined in
concentrations as low as lO-“O milligrans of chloride per
liter. (ll) (0.3. 6, 22,

For the determination of chlorine, a known solution of
silver, and solutions of iodide and potassium permanganate
were used. in carryins out the determination, an excess of
silver solution was added to a known weight of chloride sol-
ution, with stirrino. The precipitate was then filtered
through asbestos on a Cooch crucible into the titrating
vessel. Sulfuric acid was then added to the solution, and
also an excess of iodide solution of known concentration,

and this excess of iodide was titrated with the permanoanate

31

solution. The silver chloride must be filtered off as it
would change largely to iodide upon the addition of the
iodide solution; but the silver iodide does not interfere
with the titration of excess silver solution by the perman-
ganate. The results were satisfactory. (61) (C.A. 15, 5799)
In the determination of the available chlorine in bleach-
ing powder two difficulties may be encountered. ll) Some
chlorate may be reduced and estimated as hypochlorite. (2)
Some chlorate may be forned from hypochlorite during the
analysis. These difficulties may be overcome by using the
electrometric method of analysis. As a comparison electrode
an eight centimeter long glass tube was taken, the end of
which was drawn out into a capillary and turned upward. The
end of the capillary was stoppered with filter paper or With
gelatin containins potassium sulfate. The tube was filled
with potassium sulfate solution to which a little titrated
hypochlorite solution was added or a drop of very dilute
potassium iodate solution. Two platinum wires were used for
he electrodes, one was placed into the hypochlorite solution
to be titrated, and the other in the comparison electrode.
The electrodes were connected through a millivoltmeter with
high resistance, or through a galvanometer with suitable
ballast resistance.’ The end-point was shown by a sudden drop
of the galvanometer to zero. Excellent results were obtained
with tuo methods: (1) The Penot method of titrating hypoch-
lorite with an arsenite solution containing an excess of

sodium bicarbonate, and (2) by the Fontius method of titrating

32

in alkaline solution With potassium iodide. (69)(C.A.15, 5052)
Willard and Fenwick titrated sodium chloride with silver
nitrate by means of the bimetallic electrode system, cleaning
the electrodes between each titration. The silver nitrate
solution was prepared from pure silver and nitric acid. Two
fine platixum thermocouple wires Wound in loose spirals, and
polarized with a 0.5 volt potential through a resistance of
100,000 ohms, comprised the electrode systeh. The theoret-
ical figures of the reaction were obtained regardless of the
direction of approach to the end-point. 1105) (C.l. 17, 1397)
Willard and Fenwick also titrated chlorate ion with tit-
anous sulfate, in a titrating vessel sealed with mercury, and
the solution protected by an atnosphere of carbonic acid gas.
The solutions affected by light were stored in bottles covered
with black enamel paint, and provided with a siphon for remov-
ing the needed portions of the solutions. Either sulfuric or
hydrochloric acid may be used to acidify the titrated solution.
'116) (C.A. 17, 5506)
Chlorides may be titrated with mercuric nitrate usins the
mercury electrode, but not in the presence of bromides. (126)
(c.i. 18, 1799)
In the titration of bleachin” powder solutions, with
arsenious acid, good asreement was found be ween the electro-
metric and the ordinary titration method with iodo-starch
paper as indicator. Evidence that chlorate vapors were not
formed during the titration was found, which is contrary to

the findinss by Clareus. ( £9) (C.i. 17, 2545)

k.

33

Silver chloride is more soluble in the presence of a
protective colloid like gelatin, but this phenomenon has
little effect on the titration of silver witu a sodium chlor-
ide solution. Some medicinal silver preparations cannot be
titrated with sodium chloride until the organic matter is
destroyed by fusing with sodium peroxide; but they may be
titrated with sodium sulfide. (132 (0.3. 17, 5005)

Sodium chloride was titrated electrometrica ly in t
presence of 0.5 percent soluble starch, egg albumin, Witte
yeptone, boiled and raw milk, and ashes from the same for
comparison. Dilute albumin and 5 percent peptone gave curves
with no flat portion. With milk the curve was better, but
showed an error of 1:30 as compared with the ash. (136)

(C.A. 18, 1960)

duller worked out a method for titrating chloride ion,

in sodium chloride, with silver nitrate using metallic elec-
trodes of palladium, platinum, gold, mercury, silver, and
carbon. Carbon and the metals forminr easily soluble chlor-
ides do not serve this purpose well. The mercury electrode
gave the sharpest end-point. in case of the reaction between
sodium chloride and mercurous nitrate, either platinum or
mercury will serve as a metallic electrode. This author also
makes use of the compensating electrode system in these tit-
rations. (160) (C.A. 19, 1256)

For the titration of chlorous ion with sodium arsenite,
W. D. Treadwell used a platinum electrode and a comparison

electrode containing potassium sulfate and a titrated solution

of chlorous ion in a spillary tube. The salts, sodium sul-
fate, sodium bicarbonate, ammonium acetate, sodium chloride,
ammonium chloride, sodium nitrate, or potassium nitrate in
0.2 M solutions can be used in place of the titrated sol-
ution in the comparison electrode. a reducing reagent fives
readings above the normal, and an oxidizing reagent below the
normal; but a 0.2 volt displacement does not seriously affect
the results. The compensating method of A. Fischer may also
be used to titrate chlorous ion. (163 (C.l. 19, 1829)
Iodium hyposulfite may be used to titrate solutions of
soluble chlorites after adding an excess of potass'um iodide
to the sulfuric acid solution. hypochlorous acid may be tit-
rated Witb sodium arsonate solution with and without the add-
ition of sodium bicarbonate. in a mixture, the chlorous ion
may be titrated first with arsenate and the hypochlorous ion

afterwards as indicated above. (164) . (C.A. 19, 1829)

Cobalt.- (127)*, (140)*

This method of determining cobalt by electrometric tit-
ration depends upon addition of an excess of potassium cyanide
and the titration of the excess cyanide with silver nitrate.
With an excess of cyanide cobalt is supposed to form a quadri—
valent anion. The ratio of cobalt to cyanide was found to be
1:5 both in the presence and absence of air. The titrations
were made with a silver electrode. In the absence of ammonia
the results were low, and the sudden change in potential was

at -O.290V.; but in the presence of ammonia the results were

35

accurate and the potential change was greatest at -0.520V.
(127) (C.A. 17, 942)
Cobalt can be titrated successfully with potassium cyan-
ide when silver electrodes are used and the solution protected
from the air. The pota . ium cyanide solution should be
standardized asainst a known cobalt solution, and the unknown

cobalt solution should be added to the s andardized potassium

cyanide solution. (140) (C.A. 18. 3563)

Cadmium.- (114)*

it was found possible to titrate cadmium with sodium
ferrocya nide to form cadmiun ferrocya nide. In titrating
cadmium v'ith potassium ferrocyanide the ratio of cadmiurn to
ferrocyanide was found greater than 1:1. When cadmium and
zinc are both to be determined in the same solution, it is
advisable to titra te the zinc with potassium ferrocyanide,
and then use the empirical factor recommended by Hedrich
(Luller, Die Uectrometrische mas. -analysis) for computing

the cadmium content in the next titration. (114)(C.A.17 2689)

Cerium.- (175)*

To the cerium solution to be titrated, add about three-
quarters of the necessary amount of potassium ferr icyanide,
and eXpell the air from the solution with carbon dioxide gas.
Then add 50 percent of potassium carbonate solution, with
stirring, until the solution contains about 50 percent of
carbonate ion. The titration is conducted in the usual way

with a platinum electrode.

36

The inflection of the titration curve depends upon the

concentration of the otassium carbonate beins shifted to-
! L .

wards the negative side as the concentration increases. Tie
titration was con ucted in an alkaline medium in order to
avoid the precipitation of cerium salts. The method gave
exact results with several specimens of cerium metal and
alloys. (175) ' (C.A. 19, 2922)

1;

.L

ormaldehyde.- (14e)*

‘

ilver nitrate solution reacts with aldecjde to fern

U)

7:

,ilver and formic acid quantitatively, if at least three
times as much sodium carbonate is present than required to
neutralize the three mclecules of acid from one molecule of
aldihyde. An approximate deternination may be made by add-
ins 25 cc. of saturated soda solution to one cc. of aldehyde,
and titrating with 0.1N silver solution, using a platinum
electrode, until the e.m.f. against the calomel electrode is
0.2 plus volts. Then mix five cc. more of the standard silver
nitrate solution tha was required for the preliminary deter-
mination with one cc. of aldehyde and 25 cc. of the saturated
soda solution. After five minutes add dilute sulfuric acid
until the solution is acid to methyl orange, and titrate the
excess silver With 0.1N potassiun chloride using a silver
electrode. Satisfactory results can be obtained. (148)

(0.1. 19, 452)

Gastric Juices.- (24)*

A Special electrode is used consisting of a tube ending

37

in a glass bell. A side tube is fused to the glass bell,
which contains mercury, and has a platinu: wire fused into

q

the lower end which aims under the liquid to be titrated.

Hydrogen is passed into the former tube and a connection is
made between the liquid to be titrated and the other half-
cell through an inverted test tube containing a gel of sat-
urated potassium chloride and 3 percent asar, connected with
a saturated potassium chloride solution, and then with a
calomel half-cell. The electrometric measurement of hydrogen
ion concentration-is recomuended for determinins free acid

in gastric juice. Combined hydrochloric acid is determined

by titratins to pH 6.5. (24) (C.A. 11, 2213)

Indigo.- (99)*

"If a sample of indigo is dissolved in concentrated sul-

g2.

furic aci and the resulting indisosulfonic acid diluted so

that it is 0.05 - 0.01N with respect to indigo and 0.2 - 0.5N
with reapect to sulfuric acid, passase throush a cadmium re-
ductor will result in the formation of indiso white which car
be titrated electrometrically with ferric chloride solution"

(93) (C.A. 17, 39)

Iodine.- (11)*, (44)*, (53)*, (59), (62)“, (67), (105),
(116), (1263*

The electrometric method may be applied to precipitation
analysis and hence to the determination of the halides by pre-
cipitation with silver compounds. Determination may be made

on solutions as dilute as 10 - 2. milligrams of chloride per

38

liter, or smaller quantities of bromide or iodide. iodide
may be titrated in the presence of large quantities of bro-
mide or chloride. The method may be applied to the determin-
ation of iodide in urine. (ll) (0.3. 6, 722)
Twenty-five cc. of 0.1N pota.sium iodide and 5 cc. of
3 percent potassium iodate may be titrated with 1N hydro-
chloric acid added to 0.5 cc. at a time. The theoretical
quantities were obtained. The presence of bromide up to
four times the iodide does not interfere, althouph longer
time is required to make the readings. Chlorides interfere
with this determination. A solution of 0.01N potassium iodide
may be titrated but about 50 seconds time should be allowed
for each readins. (44) (0.3. 15, 2167)
The oxidizing power of the halides stands in the order
iodine, bromine, chlorine, and in the presence of a strong
oxidizing agent they are liberated from their salts in the
order given, and there is a sharp change in potential between
iodine and bromine. Potassium dichromate, potassium bromate,
and potassium iodate may be used. Fenc,, iodide may be det-
ermined in the presence of considerable amounts of bromide,
nd in the presence of large quantities of chloride by the
oxidopotentiometric method in the presence of hydrochloric
acid. (55) (C.A. 14, 5204)
In place of the ordinary potentiometer, Hendrixon used
a rheostat coil of 169.45 ohms resistance having a graduated

contact beam. 229 additional ohms resistance were added to

39

the end of the coil not traversed by the contact slide. This
made the scale one meter long in effect, and each millimeter
equaled a potential of two millivolts. This improvised pot-
1

entiometer was used in connection witn a galvanometer having

a sensitivitv of 106 mesohms. A three neched Voulff bottle
0 a

'0

erved as a titrating flash, the necks holdins the calomel
cell, platinum electrode, and carbon dioxide tube, respect-
ively, with an extra hole bored in the flash, or cork, to
admit the tip of a buret. Weighing burets were used in some
cases. Carefully prepared iodine was titrated with potassium
pernanganate solution, standardized with sodium oxalate, in
about 1N solution of sulfuric acid. Chloride and bromide in
the titrating solution prevents accurate determinations; but
it was found that chloride equivalent to the iodide, and bro-
mide equivalent to 25 percent of the iodide, could be present
in the titrating solution and not interfere with the accuracy
of the results. This reaction was also carried out by adding
an excess of potassium dichromate and titrating the excess with
potassium permanganate. (59) (C.A. 15, 639)
Kolthoff made some observations on the titration of iod-
ides by means of potassium permanganate, and found that if
this reagent is added drop by drop near the point of equival-
ence, and E determined for each addition, this point is made
known by the relative variation of the potential. The maximum
potential corresponds to the point of equivalence. The method
is exact when the determinations are made in a sulfuric acid

medium of at least 0.153. Iodide may be determined in the

4O

presence of twice as much bromide aid five times as much chlor-
ide although the fall of potential at the equivalent point is
lowered by their presence. (62) (C.l. 15, 3799)
Hendrixon titrated potassium biniodate with potassium
iodide in excess and determined the excess of iodide with
potassium permanganate. This investigator also studied the

effects of hvdrochloric -cid and sulfuric acid on the reaction.

{.2

The iodide solution was standardized by means of thiosulfate
solution which had first-been standardized by resublimed iod-
ine. The permanganate solution was standardized by means of
sodium oxalate. The strength of the iodide solutions was
0.05N and also 0.02N.

Iodide, in sulfuric acid, may be determined in the pres-
ence of chloride up to a concentration of 0.1N in the titrated
solution by means of permanganate; but hydrochloric acid can
not be substituted for sulfuric acid, except in low concent-
rations and within narrow limits of concentration.

iodide, in sulfuric acid, may be titrated directly with
iodate. {67) (C.A. 15, 1668}

Willard and Fenwick showed that iodide could be deter-
mined by oxidizing it to iodine cyanide with potassium per-
manganate in hydrocyanic acid solution. The titration was
made in 100 cc. volume of solution which contained 15-20 cc.
of hydrochloric acid and five cc. of a 10 percent potassium
cyanide solution. This method was applicable in various con-
centrations of chloride ion, and in moderate concentrations

of bromide ion. The authors claim, however, as a more accur-

41

ate method, the oxidation of iodide to iodate by excess of
alkaline hypobronite, and titratiné the excess of hypobrom-
ite with arsenite. Eronide or chloride does not interfere
with this last reaction, and either the bimetallic electrode
system or the monometallic electrode system may be used. The
hypobromite solution was made by pouring 40—50 grams of bro-
mine, slowly, into a solution of about 50 grams of potassium
hydroxide in 250 cc. of water at about zero temperature, then
diluting to five liters. This solution was standardized vith
a 0.1H solution of sodium arsenite prepared by weisht from
pure arsenous oxide dissolved in sodium carbonate. The actual
Operation of the method was as follows: A weighed portion of
a standard potassium iodide solution Was made alkaline with
one tram of potassium hydroxide, an excess of the standardized
hypobromite solution was added, the whole diluted to 100 cc.,
allowed to stand for five minutes and the excess of hypobro-
mite titrated with the standard arsenite solution. (105)
(0.3. 17, 1597)
Uendrixson also titrated iodate with titanous sulfate
so ution. The 0.05a iodate solution was standardized with
iodide and thiosulfate, and the titanous sulfate solution
with permanganate and dichrouate. The volume of the titrated
solution was about 800 cc. and the acidity between 1.0N and
2.0N sulfuric acid. Hydrochloric acid was found to be un-
satisfactory. The correct end-point was found by multiply-
ing the volume of titanous solution at the first fall in

voltare by the factor 1.2 v(hich is in accord with the theory.

42

the strensth of the titanous sulfate solution was 0.75N x
1.727. (116) (6.3. 17, $506)
Iodides, at very sreat dilutions, can be titrated accur-
ately With mercuric chloride and the mercury electrode. Earse
amounts of bromide interfere with this titration. (126)

(3.i. l8, 1799)

Iron (Terrous and ferric\.- (6), {14), (17), ($83, ’59)*,

r45), (495*, (s7), (993*, (1193*, (124l‘, (125)* (144\*,

.
(15;), r137)*, (1651*

As shown by the large number of references, considerable
work has been done on the determination of iron by electro-
metric methods. iron salts are also extensively used as tit-
rating reagents.

Feters was one of the first investigators to W(rk with

oxidizins and reducing cells from the standpoint of quantit-

C
CD
*1
:3
(D
pa
0
{7‘
H
(D
H)
l

ativo determination of ions; and his wor}: Was con
1y With the ions of ferrous and ferric comp unds. Fe devel-
oped the formula connonly used as the theoretical basis for
oxidizins and reducing reactions, and established the relat-
ionship between ferrous anl ferric ions in solution. (See
the part of this thesis dealins With the theory of electro-
metric titrations}. The electrodes used by ‘eters were the
Weston electrical element and metallic electrodes of plat-
inum, sold plated platinum, and platinized platinum. Teters
also studied the effects of chlorides and fluorides on the
oxidation potential of ferrous and ferric ion, and came to

the conclusion that the chanses in potential observed by him

43

were due to the forsation of comolex ions by the interaction
of the substances present in the titrated solutions. (6)

Forbes and Bartlett titrated dichromate ion With ferrous
ion. (See under chroniuz, 14)

Tiesel titrated potassium dichrolate solution with fer-
rous sulfate usins an automatic, electrically controlled
buret. (17) (C.A. 9, 8J8)

The electrometric method was used by Ferguson for det-
erminins small quantities of iron in Flees sand. The sand
was first decomposed in platinum by means of sulfuric acid
and hydrofluoric acid, and the residue ignited with iron-
free potassium pyrosulfate, and the cake dissolved in dilute
sulfuric acid. The iron was then reduced with StdiflCF?
chloride and the excess of stannous chloride titrated with
dichromate in a platinum-calomel electrode cell. The pur-
ity of the reasents, and of the platinum were of oreat im-
portance in this deteruination. (28)

Ferrous sulfate was titrated with potassium dichromate
in both hydrochloric an sulfuric acid solutions; with pot-
assium bronate in hydrochloric acid solution; and with pot-
assiun permanganate in both hydrochloric and sulfuric acid
solutions. Potassium dichromate and potassium bronate gave
satisfactory resul s. Potassium permanganate was more Satis-
factory in hydrochloric acid solution than in solutions of
sulfuric acid. A platinum electrode and the calomel half-

cell were used in this work. (59) C.A. 14, 3204)

44

The work of Hastetter and Roberts on the electrometric
determination of ferrous and ferric ion has been widely quoted..
They studied the effects of different metallic electrodes used
with the calomel half-cell; also the effect of varyins the
concentrations of the titrating reagents, and of the acid sol-
utions; and they combined the Various necessary pieces of tit-
rating apparatus into one unit which they Called the titrating
head.

The titrating head.- This apparatus consists of a closed
glass tube of suitable size, througn the bottom of which is
passed the salt arm of the calomel cell, the delivery tube
of a burst, and the platinum electrode which passed through
a jacket tube with a side arm to admit carbon dioxide gas to
the flask. This titrating head was made to fit over an Erlen-

meyer flask, the salt arn of the calomel cell and the platinum

J

o"
L~'-L.

C

electrode being lens enoush to reach the bottom of the fl‘
By means of this form of titrating head, the solutions under
investigation could be heated to the desired temperature, and
also protected against oxidation by the oxygen of the air by
an atmosphere of an inert gas placed over the solution in the
flask.

The end-points were checked in two ways, (1) by balanc-
ing known quantities of two reasents, and (2} by checking
against the pink color of a permanganate solution as used in
the common method of volumetric analysis.

Electrodes.- The kind of electrodes used, and the inverse

45

order of their suitability to the purpose was as follows:
Palladium, 60 percent gold-40 percent palladium, gold, plat-
inized platinum, brisht platinum. The palladium electrode
did not show a noticeable change in voltare at the end-point.
A small, bright platinum electrode of small surface area was
found to be the best kind of electrode. The electrodes must
be carefully cleaned, and the inert gas used to protect the
solutions asainst oxidation, must not come in contact with
them. A calomel half-cell was used with these electrodes.

Temperature.- ”Temperature is specific for certain re-
actions, and sudden chanses in temperature should always be
avoided.

Acidity f the titrated solution.- elerp end-point volt-
ase deflections were obtained over a ride range of acidity
with sulfuric acid.(l7 to 67 percent). Nitric acid cannot be
used because of its oxidizins power. hydrofluoric acid, in
the presence of hydrochloric acid and sulfuric acid, was not
detrimental to the titrations. in using 0.001N dichromate,

the acidity and volume of the solution became important fac-

tors. lt was found that small volume of solution with 25
percent concentrated sulfuric acid, and 0.001N or 0.0005N
dichromate also in 25 percent su furic acid gave sharp end-
points.

The titration of ferrous iron in the presence of various
amounts of ferric iron gave definite end-points, but the range

of the difference of potential at the end-points was shortened.

The method can be used to determine ferrous iron in ferric

46

oHpound . Ten prams of the ferric salt was dissolved in
rater in an atmosphere of carbon dioxide, and an equal volume
of concentrated by Wrociloric acid added to the solution; after
which the titratins head was put into place and the solution
titrated with 0.01N potassium dichromate. This method is also
applicable to determinins small quantities of ferrous iron in
other substances.

Ferric iron Was also successfully tit‘ated with stannous
chloride and the method adapted to the determination of small
amounts of ferric iron in ferrous salts.

Since stannous chloride can be titrated with dichronate
solution, a method for determining the total iron in a sol-
ution is as follows: The iron solution Was first completely
reduced with stannous chloride solution in slight excess, as
shown by the voltage; and then titrated With dichromate. The
first end-point shows the end of the oxidation of the stannous
chloride, and the second one, the er d of the oxidation f the
reduced iron. ”‘e volime of the solution used between these
two points was taken as a measure of the iron present. The
method has been used on sla ss sand, sodium and potassium salts,

barium carbonate, zinc oxide, borax and boric acid, and the
oxides f lead; and can be used for determining iron in any
subs ance not reduced by stannous chloride.

This method for iron may also be 1m _ed to determine blanks
to make corrections neceSSary to be applied to these reactions

when carried out with permanganate or color indicators under

different conditions, and for Various purposes. l’"l(v.n.l.,eJlJ)

47

Ammonium ferrous sulfate was used to titrate chromium in
steel analysis. This reagent was added in excess to the pre-
pared sample, and the excess titrated t3 means of potassium
dichromate in a cooled, sulfuric acid solution.(49)(C.A.13,1802§

Custavson and hnudson titrated iron in the presence of
uranium and Vanadium. Ten cc. of a solution containing 41.1
milligrams of iron as ferrous sulfate, 19.7 milligrams of
vanadium as vanadyl sulfate, and 121.0 milligrams of uranium
as uranyl acetate was diluted to 250 cc. and five grams of
zinc added. dulfuric acid was then added and the whole boiled
frr 3“ minutes, filtered through cotton, and titrated while
hot with permansanate solrtion in an atmosphere of carbon
dioxide sas. four cc. of sulfuric acid was added at the start,
and after tke titration has proceeded to the point of the
second voltage deflection, four co. more were added. This
method of adding the acid proved hishly.successful, as it pro-

for a sharper infleciion at the end-points. it Was

’3:

vide
possible ty this nethod to calculate the quantities of iron,
Vanadium and uranium present in a uixture of salts of these
metals. (87) (0,3, 17, £07)
The reducing power of ferrous solutions is greater than
that of vanadyl solutions. A mixture of these solutions may
be analyzed by means of potassium permanganate. The ferrous
iron is oxidized first and the second potential deflection
is due to the oxidation of the vanadium. The temperature
should be raised to '700 for the vanadium end-point. (89)

(c ;, 17, 1401)

48

If potassium permanganate is added to a ferrous iron

solution in the presence of dilute sulfuric acid, the iron

is oxidized to the ferric state and the mansanese reduced.

If potassium fluride is then added to the solution and the

q I

titration continue a second end-point is obtained when the
9

manganese is in the tervalent condition. In case the orig-
inal solution contained no manganese, the volume of the yer-

A.

manganate used in the second step of the titration will be

("+-
(D

L.-.

in‘ ‘ _
' ). ‘1’.C1‘1J

1.2 times as large as that used in the first s

D

125 cc. of solu

J

tion was used, atout 0.53 in sulfuric acid,
and seven grams of potassium fluoride with five cc. of 2H
sulfuric acid were added for the second part of the deter-
mination. Fluoride cannot he present at the start of the
titration. (119) I (c l, is, 3-51

Ferric chloride equivalent to a 5 M solution was tit-
rated with N sodium h"droxide by means of the oxysen elec-
trode and saturated calomel cell. The curve was continuous,
showing that iron oxychloride does not exist. (124)

(0.3. 18, 1039)

This method is an improvement over one previously given
(119). A mixture of ferrous and nansanous solution is first
titrated, in a platinum dish with potassium dichromate until
the potential against a normal electrode is 0.57 volts, which
corrGSponds to the complete oxidation of the iron. Then heat
the solution to 80C, add seven 5 ams of potassium fluoride,
and continue the titration with pernansanate until the man-

ganese is oxidized. A possibility is indicated of titrating

49

iron, vanadium, uranium, and manganese in a single operation.
(125) (C,A, 18, 1628)
Several references to the subject are cited in this

1‘

article. Vanadium can he oxidized by means of permanfanate

can he determined

r.) a
(.2;

to three states of oxidation. Vanadic ac
in the presence of iron but not in the presence of chromic
acid. Ferrous ammonium sulfate is a good reagent for deter-

mining vanadium. (144) (Col, 18

71)

C»)
()1

(—1-
U)

The oxidation potentials of ferrous and ferric sal
were studied by Carter and Cleevs usinp both concentrated
hydrochloric and ph sphoric acid. Platinum foil electrodes

..
L

and electrodes made from g ass coated with p atinum gave con-

‘

cordant resrlts. The logarithnic relationship of ?eters was
substantiated. The oxidation potential decreases 1"ith in-
creasing concentration of hydrochloric acid. This decreas-
is linear with the nornality of the acid between 1N and lON.
The variation in the diffusion potential does not account for
this change in voltage. Results were nearly the same with

the phosphoric acid. (155) (C,A, 19. 945)

Ferric salts may he titrated at room temperature with

erric ions Can he titrated in the presence of cupric
ions and also quinquivalent antimony ions, with titanous chlor—
ide. Then copper is present, meconic acid is used as an ex-

ternal indicator. (157) (0.3. 19, 450)

5O

Ferrous iron may be titrated with potassium dichronate
in a strongly acid hydrochloric solution. Magnetites not
decomposed by sulfuric-hydrofluoric acid treatment are decom~
posed by hydrochloric acid. Adding excess of stannous chlor-
ide to a ferric iron solution and titrating back with dichro—
mate or iodine is better than the direct titration. Titanium
in the solution does not interfere with the reduction of fer-

ric iron b* means of stannous chloride. (165) (C.A. 19 1852)
9

L.

lead.- (128)*

The platinum-calomel cell system was used in this re-
search. Lead and zinc, in the presence of one anoth r, may
be determined electrometrically by two titrations; one for
total lead and zinc, and one for zinc alone. For the sum of
both metals, to every 100 cc. of solution one cc. of potass-
ium ferri yanide was added and the titration made with pet-
assium ferrocyanide at 75°. The reactions are:

3 22:1 *" + 21$e(c:\’)6“‘ = K22115(F€(CN)6)2
:2 F‘b ” + Fe(cm6““ .__ :u:

The zinc is precipitated first. There are two breaks
in the titration curve, the first at 0.54 plus volts, and
the second at 0.19 plus volts. Three to five minutes tide
should be allowed near the first break for the potential to
become constant. 1f the lead is first precipitated with di-
lute sulfuric acid, then the zinc may be determined alone,

and the lead quantitatively estimated by difference in the

titrating solution used for each titration. (128)(C.A.17, 945)

Manganese.- (50), (ee)*, (117)*, (119)*, (125)*

Kelley, Spencer, lllingworth, and Gray studied the elec-
trometric method for manganese in connection with the deter-
mination of this element in steels, and claim that the pres-
ence of chromium and vanadium is not detrimental to the tit-
ration.- Iercurous nitrate was used as the reducing reag-nt
and this solution Was standardized by means of permanganate.
The oxidation of the manganese may be done by either of two
methods; (1) Solution of the steel in nitric acid and oxidiz-
ing with ammonium persulfate and silver nitrate; or (2), the
oxidation may be carried on by means of sodium bismutlate.
200 cc. volumes were used for titrating, containing 50 cc. of

sulfuric acid (Sp. gr. 1.58). The temperature should not be
above 40o. Iitrous acid interferes with the titration, but
nitric acid does not. (50), (0.3. 12, 458)

A reaction which may be taken as the basis for an elec-
trometric oxidation method for manpanesc is as follows:

1.11104“ + 4th H + 83* : sun t” 4- 414.10
In the presence of B - this reaction proceeds with the quant-
itative formation of a complex ion nnF5". The neutral or
slightly acid manganese solution, containing not more than
0.2 grams of manganese, is diluted to 300 cc. and 5-10 cc.
of sulfuric acid (1:7), and 5 grams of ammonium fluoride added.
The titration is carried out With 0.1K potassium permanganate
in the cold. The presence of ferric iron, chlorine, or nit-

rates do not interfere with the quantitative nature of this

reaction. (68) (0.3. 15, 2800)

52

The Volhard titration method for manganese was Carried
out electrometrically, in which the manga;ese is oxidized
by means of permansanate. About seven minutes time was re-
quired to establish the end-point. Test results were obtained
in a volume of 100 cc. 0.43 with sulfuric acid, and contain-
ins eight grams of potassium fluoride. (117) (C.A. 17, 5306)
According to Muller and Wahle the method discribcd in
the preceding reference may be app plied to th as determi; ation
of iron and manganese together. These investigators used
125 cc. of solution 0.5N in sulfuric acid and titrated with

J

permansanate to the iron end-point, then they added seven
grans of potassium fluoride, and five cc. of 2N sulfuric acid
for the second part of the titration. If fluoride is present
at the start, there will be s~me atmospneric oxidation of the

)

This reference is supplementary to the preceeding one.

iron. (119) (C.A. 18, 56

\)3

The Chang so in the method is as follows: Flace the reduced
iron and manganese in a platinum dish, and titrate with pot~
assium dichromate until 0.57 volts are obtained. This pot-
enti a1 correspor ds to the complete oxidation of the iron.
Then heat to 80°, add seven grams of potassium fluoride, and
finish the titration With potassiun permanfanate. (125)

(C.A. 18, 1628)

Molybdenum.- (108)
Willard and Fenwick used their polarized, bimetallic

electrode system composed of platinum electrodes to determine

the end—point in the reduction of molybdie acid to a salt
of the pentavalent molybdenum. The reaction was found to
be rather slow for ideal results. 1 5 - 10 percent solution,
by volume, of hydrochloric acid was found to be the best for
this determination. sulfuric acid was found to be unsatis-
factor;. The reduction was carried out by means of titanous
sulfate standardized with ferric sulfate prepared from elec-
trolytic iron; and the strength of the titanous solution was
0.05N. a solution of sodiuzn melvbdate of approximately 0.017N
was titrated with the titanous sulfate solution. The authors
tate that the titanous solution should be standardized atai
a mol; bdate solution. The method is not high y accurste.
Ammonium phosphomolybdate was dissolved in ammonia, fil-
tered, acid ci ied and titrated with standard titanous sulfate.
Phosphoric acid may be used provided the s lution is titrated

hot. The method is hishly recomnended. Tungsten does not

interfere with this method. (108) (0.1. 17, 1786)

Lercury.- (135 *, (161)*

Mercuric ions in a nitric acid solution react with am-
monium cyanide to form the insoluble ucrcury cyanide, and a
slisht excess of cyanide gives a red color with ferric iron.
This reaction can be used for the ti ”r tion of mercury in the
absence of halides or mercurous salt. The mercury can be
dissolved in nitric acid, treated with an excess of pota ss1 um

permanganate, and ti o excess of pernanranate removed with

ferrous iron, after which the so1. ution may be titrated electro-

St

54

metrically with standard ammonium cyanide. The mercury elec-
trode~calome1 electrode system was used in these experiments.
(led) 0.1. 18, 1959)
This reference contains a review of the literature on

the use of mercuric salts in electrometric titration methods.
The difficulty experienced in the use of mercu y salts with
the mercury electrode is the reduction of metallic mercury_
on the electrode. Mercurous chlorate is considered the best
titratins mercury reagent. The salt is prepared by heating

a weighed amount of oxide of mercury With the equivalent

mount of concentrated perchlerie acid, free from chlorine

if possible, and excess of mercury in a flash under a con-
denser until in a test portion, after the addition of excess
chlorine, no mercury Can be detected in the filtrate. Chlor-
ine and bromine may be determined accurately with this reagent.
In titratins iodides, cyanides and thiocyanates, solutions of
these compounds should be run into the mercury chlorate sol-
ution. Chlorine may also be determined with mercuric chlorate
as well as With mercurous chlorate. The determination of mer-

curic and mercurous ions simultaneouslv 1”ith chloride was not
Q)

(.0

successful. (161) (C.d. 19. 15 7)

Nitrates.- (187)*

If a nitrate, in the presence of hydrochloric acid, is
treated with a ferrous salt solution, ferric iron will be
formed which can be titrated with a standard titanous chloride

solution in an atmosphere of carbon dioxide was. The method

55

is rapid and accurate. The sign of the potential changes

at the end-point. (157) (C.A. 18, 960)

Kitrites.-

Tendrixson determined nitrous acid by addino to it a
5 - 10 percent excess of a standard potassium permanganate
solution, and then an excess of a standard potassium iodide
solution, and titrating the excess of iodide With the per-
mansanate solution. The titrations were made in a 1.5N sul-
furic acid solution. The method was found to be highly ac-

curate. (61) , (C.A. 15, 5799)

Nickel.- (85)*, (140)*

Huller and auterbach determined nickel with potassium
cyanide and silver nitrate. The 0.1N silver nitrate solution
was standardized electrometrically with potassium cyanide.
The standard cyanide solution was added to 10 - 23 cc. of tne
nickel solution in excess, ard this excess titrated with the
silver nitrate solution. (85) (C.A. 17, 249)

Bickel and cobalt may be titrated together with a sol-
ution of potassium cyanide an“ rith the use of silver elec-
trodes. The unknown nickel and cobalt solution should be
added to the cyanide solution after standardizing the cyanide
solution by means of a known nickel-cobalt solution. (140)

(C.A. 18, 2662)

Reducino Sugars.- (113)

This reference describes the use of the potentiometer

56

method for determining reducing sugars ouantitatively by
means of reduction with Fehling's solution. The Fehling's
solution was standardized against pure dextrose. .The "B"
part of the Fehling's solution was made up with only one-
half of the quantity of sodium hydroxide as recommended in
order to get a more distinct end-point. 1n the actual tit-
rating procedure, equal portions of the "A" and "E" solutions
(10 cc. of each) were mixed and 50 cc. of distilled water
sadded. 'The mixture was then brought to the boiling point,
and the sugar solution added until the end-point of the re-
action was reached. The results were considered as being

satisfactory. (113) . (0.4. 17, £545)

Silver.e (10)*, (67), (120)*, (132)*

Dutoit and VonWeisse used a silver electrode in their
Studies on the practical application of the properties of
concentration cells to analeis by precipitation. They also
called attention to the value of silver salts as precipitan s
in volumetric analysis. (lO)| (0.3. 6, 722)

Silver nitrate solution is comnonly used as a reagent
to determine the halides quantitatively by the patentiometric
method. The process can be reversed and silver ion may be
determined by means of potassium iodide solution; or in some
cases, potassium cyanide is preferred as a rearent for silver.
Hendrixson determined silver by means of an excess of potassium
iodide and titrating this excess with permanganate. A solution

of silver sulfate containing 0.0021573 grams of silver per gram

57

of solution was used. The permanganate solution vas standard-
ized with sodium oxalate and the potm sium iodide solution
with the permanganate solution. Nitric acid may be used on
concentrations up to SR. A solution of silver nitrate pre-
pared from pure silver and nitric acid contained o.oo4714
grams of silver per pram of solution, was also used, and tit-
rated by the method without filtering off the precipitate.
The res alts were satisfactory. The author also calls attention
to the value of pure metallic silv or _for use as a standard
to make up kno*n solutions of iodide and permanoa nate. (67)
(C.;. 15, 1668)
Silver can be titrated accurately with potassium cyan-
ide. dilver and copper can be determined simultaneously with
potas siun cyanide provided that th silver cyanide first form-
ed is filtereC off when its end-point has Mela determ Mr ed;
then by adding sodium bisulfate the copper can be determined
at 700 with the same reagent. (120) (C.£. 18, 640)
The presence of gelatin as a protective colloid has
lit le effect upon the electrometric titration of silver with
sodium chloride solution, although the solubility of he sil-
ver is increased. Some medicinal silver preparations cannot
be titrated until the organic matter has been destroyed by
fusins with sodium peroxide, or by titrating directly with

sodium sulfide. (152 . (0.3. 17, 5005)

Sulfur.- (84), (953*, (106), (126)*, (135

Hendrixson and Verbeck attempted to titrate sulfurous

58

acid with permanganate and found that only about 90 percent
of the sulfurous acid was oxidized to sulfuric acid. A
weighed quantity of standard permanganate solution was tit-
rated with a sodium sulfite solution in an acid concentration
'of about 1N sulfuric acid at the end of the titration. The
solutiox was protected from oxidation by means of carbon di-
oxide sas. Fermansanate solution was used with large excess,
slight excess, and to the end-poin ; and the excess of per-
manganate titrated with potassium iodide. The large excess
of permanganate was th most efficient, oxidiziné 93 percent
of the sulfite to sulfate. {84) (C.A. 17, 33)
Free sodium sul”ite in a sulfur black dye bath may be
determined to an accuracy of 0.5 percent by titrating with

iodine and sodium thiosulfate. By a pipet method the iodine

c+
F4

is allowed to react with the dye solu

I,

'on for one minute, and
then the thiosulfate is added in excess and back-titrated
with more iodine. (95) . (0.3. 16, 5208)
Willard and Fenwick titrated sodium sulfide solution
with an ammonical silver solution. The change of e.m.f. at

the end-point v'as great. The sodiu; sulfide solution was

made by dissolvins 9 to 10 grams of hydrosen sulfide in

m

solution of 96 grams of sodium hydroxide and diluting to six
liters. A 0.053 silver nitrate solution containins an excess
of about 30 cc. of 28 percent ammonium hydroxide per liter,
and standardized gravimetrically by precipitation of chloride,
was used as the titrating solution. The addition of more

sodium or amnonium hydroxide was without effect in the end-point.

59

presence of sulfite, thiosulfate, and sulfate was pra ctic-'
ally ithout e feet, but slishtly decreased the quantity
of silver required for the titration. The presence of chlor-
ide reduces the quantity of silver required somewhat more
than does the presence of sulfite or sulfate. The method is
applicable to the determination of sulfur in steel by first
converting it to hydrosen sulfide and then titrati ng with
ammonical silver solution. (106) - (C.A. 17, 1596)
Sulfides can be determined in dilute solution with

d

mercuric chlor by means of the mercury electrode, and in
the presence of sodium hydroxide. The effectiveness of this
method is diminished slightly in the presence of salts of

.1

divalent ions. 3i ver nitrate is a better reagent for det-

ermining sulfides in concentrated solutions, but in dilute
solutions the results are too low. The complex ion mercury
thi osulfate is formed at the potential break in the titration

of thiosulfate olt.tions . Gulfide and thios ulfa- te C; n be

:39

titrated in the same solution with mercuric chloride, and at
th e neutral point. (126‘ (C.A. 18, 1799)

A solution of a soluble sulfate a7 be treated by deans
of a solution of a lead salt in exooss and the excess deter-

mined by electrometric titration with 0.1N pota ss -ium ferro-

cyanide. (155a) (C.A. 18, 162 )

Selenium.- (109)
Willard and Fenwick determined selenium quantatively in
selenious acid b;7 ti trating with titanous sulfate in a cold

hydrochloric acid solution saturated with sodium chloride,

60

and using the bimetallic electrode system. moderate amounts

of sulfuric acid had no harmful effects, and the reduction is
independent of the presence of iron. The reduction of copper
and selenium by titanous sulfate is selective and both may be

determined in the same solution. (109) (0.x. 17, 1766)

Tin.- (147)

Fleysher titrated stannous chloride alone, and in mix-
tures with potassium dichromate in hydrochloric acid solution;
and he also titrated antimonous and stannous chlorides toget-
her with dichromate solution, after first adding an excess of

mercuric chloride. A platinum foi1~calomel electrode cell

vas used. The titrating solutions were protected by means of

.4

an atmosphere of carbon dioxide gas. The reducins'solutions
were prepared by dissolving pure tin and antimonous oxide in
dilute hydrochloric acid. nohr's salt was used to standardize
the dichromate solution. The values for the solutions are

as follows: stannous chloride, 0.2020N; dichromate, c.1u;
antimonous enloride, 0.098353. The titrations were made at

room temperature and with constant stirring. (147)(C.l.19,451))

Titanium.- (72)*, (85), t99)*, (14e)*

By means of cadmium powder used in a Jones' reductor,

c+

itanium was satisfactorily reduced in titanium sulfate sol-
utions. This reduced titanium was titrated with potassium
permanganate with good results in sulfuric acid solutions

containing 0.9 to 18.8 grams of sulfuric acid to 10 cc. of

61

a l to 10 titanium sulfate solution. There was less neut-
ralization of the acid than when zinc was used in the Jones‘
reductor. Iron and titanium was reduced tofiether in the Cai-

miun reductor and titrated electronetrically with permanganate

(“2

solution. (72) ( .A. 15, E041)

Uendrixson and Verbeck used the electrometric method to
investigate methods of accurately standardizing titanous sol-
utions. The titrations were carried on With titanous sulfate
in sulfuric acid solution protected by an atmosphere of hydro-
gen. 50 cc. of about 0.1K ferric alum made up to 150 cc. with
dilute sulfuric acid ”as ”asked first with carbon dioxide sas;
the electrodes (platinum calomel? were then inserted and the
titanous sulfate solution added to a liberal excess of the
ferric alum solution, and this excess of ferric alum titrated
with 0.051653 pernansanate solution.' The pernanranate sol-
ution was also replaced with a 0.05N dichronate solution.
:itanous sulfate solrtions may be accurately standardized by
titratins them into standard pernanfanate or dichronate sol-
utions. ?oth methods were found to be accurate. (85‘

(0.9.. 17, 38)

Titanium may be titrated in the presence of iron by
first redu'ing with cadmium powder in a Jones' redretor tube,
and then titrating with permansanate or ferric chloride sol-
ution. It is a good plan to add some sodium fluoride to the
solutions being titrated. Then tbe titanium content is low
as compared to the iron, reduce the iron with sulfurous acid,

add ammonia, and filter off the precipitate of titanic acid

62
which will be contaminated With some ferric hydroxide, reduce
the titanium with cadmium and titrate with ferric chloride.
(99) (0.x. 17, so!

Zotthoff made quite a complete study of the reducing
action of titanous solutions, and found bat the equation
oH ='o.os + 0.058 log. (tiIV)(H*)/(TiIII) holds at 180. In
this paper are given methods for preparing pure titanium
chloride crystals; for detecting iron as an impurity in tit-
anium solutions; for standardizins the solutions; and for
the estimation of dichromate, vanadate, molybdate and iodate.
Directions are also given for estimating some oxidizing agents
when present tosether. Details _f hese methods are not given

in the abstract. (146) (C.A. 19, 449)

"ranium.- (75), (87), (9a)*, (122)*
Uwins and Tldridge successfully titrated trivalent uran-
ium to the tetravalent form and then to the hexaValent con-
dition; both with potassium permansanate in sulfuric acid
solution, and with potassium dichromate in hydrochloric acid
solution. The method was also used to titrate uranium in the
presence of iron where three end-points were obtained. The
uranium solution was standardized by eVaporation and igniting
to uranium oxide (U505), one cc. containins 0.0151 gram. The
permanoanate solution was standardized With sodium oxalate

I'Ofl OTC

Ho

and was 0.1011N; and the dichromate solution with
and it was 0.09064N. A closed platinum-calomel electrode

cell was used which admitted the electrodes buret, stirrer,

and a Jones' redrctor tube used for reducins the uranium

63

solutions by means of zinc directly before titrating them.

a new form of platinum electrode was used by the authors
which was made by sealing a laroe platinum wire into the end
of a glass capillary tube and grinding the exposed end flush
with the end of the glass tube. This electrode Was }:<pt
bright by rubbing it across corbonundum paper before each
determination.

The solution proportions used for makins the titrations
were 10 cc. of the ur aniuni solution and 40 cc. of a sulfuric
acid solution containing two cc. of concentrated sulfuric acid
per 90 cc. of solution. After Warminp to 80-900, the air was
removed by carbon dioxide gas, and the solution poured slowly
through the reductor tube which was then rinsed with 50 cc.
of the sulfuric acid solution, and the reduced uranium sol-
ution was titrated with permanganate. The iron-oxidation
end-point is last in the series when iron is present. The
titration may be made with dichrouate in hydrochloric acid
solution. The sulfuric acid concentration should be less
than two cc. per 100 cc. (75) (C.A. 16, 2459)

Gustavson and Knudsen used the electrometric method to
determine uranium in the presence of vanadium and iron. @he
common form of apparatus was used, and the soiution were
protected by means of carbon dioxide gas. The; comrared With
platinum foil, a platinum wire electrode was found more suit-
able for this titration. The elements were reduced with zinc
and titrated in a su furic acid solution with potassium per-

manganate. For titrating uranium in the presence of iron,

64

the optimum concentration of sulfuric acid was found to be
five cc. of concentrated acid in 250 cc. of solution. Vhen
uranium, vanadium, and iron were determined together, it was
found best to add the acid in two portions of four cc. each
to the 250 cc. of solution. (87) (C. A. 17, 507)
Uranium may be reduced in a cadmium reductor and titrated
with potassium permanganate 0n the basis of the equation:
5U‘"’ + 2Mn04‘ + 2H20:: 5U02" + 2Mn"'+ 43*. The results
are better than when the uranium is reduced with zinc, and
also when the reduced solution is caught in an Open beaker.
(99) (0.1. 17, 59)
Uranium and iron can be titrated electrometrically in
the same solution, the first break in the titration curve
shows the oxidation of uranium from a valence of four to one
of six, and the second break shows the oxidation of the iron.
These titrations are all carried out with ootassiun perman-
ganate. ”hen uranium and Vanadium are together reduced with
zinc, three end-points are obtained with permanganate, and
best at 800. The first end-point represents in this series
the oxidation of the vanadiim from two to three valence;
the second, oxidation of vanadium from three to four valence
and of uranium from four to six; the third end-point repre-
sents the oxidation of vanadium from four to five valence.
iron, vanadium and uranium may be determined all together
by reduction with zinc and titrating while hot in an acid

solution With permanganate. (122) (C.l. 18, 643)

65

Vanadium.- (201, (21), (4:), (ea), (87), (89)*. (99)*.
(103), (1221*, (1443*

lelly and Conant were probably the first investigators
to determine vanadium by means of electrometric titration
methods. ihey used the method particularly to determine
vanadium in the presence of chromium in steel. The steel
was dissolved in sulfuric acid and heated to 8G0 C. After
adding five grams of sodium phOSphate to the solution, 0.1N

f“

" color. The

potassirm permanganate was added to the

H
C‘r“ ‘7
L. CLO}

. . . 0
solution was then cooled in ice to 10 to 20

C., and more
sulfuric acid added. This solution was then titrated with
ferrous sulfate for vanadium. A portion of the solution
representing one sram of steel was then taken for the chrom-
ium determination. (See under chromium). (20?, (21)

(C.i. lO,’2179)
(0.3. 10, 729)

Kelley, Viiey, thn and 7right develOped a method for
the determination of vanadium in steel in the presence (f
chromium based on the selective oxidation of vanadyl salts
by nitric acid. infter this oxidation, the vanadium was tit-
rated with ferrous sulfate. The authors tried several oxid-
izing agents but nitric acid proved to be t‘e most efficient.

‘

Tne method was carried out by dissolving a E-gram sample

of steel in 100 cc. of sulfuric acid (9p. ”r. 1.20. Two cc.
of nitric acid was then added, drop by drop, and the solutirn

boiled to free from oxides of nitrogen. The solution was

then diluted to 100 to 120 cc. with hot water, 40 cc. of

66

FVJ

concentrated nitric acid added, and boiled for one iorr
keeping the solution volrme above 100 cc. After cooling
to 20° C, or lower, the solution was titrated with ferrous
sulfate and potass iun dicnronate, 0.36fl9 prams per liter.

‘

These same authors further improved the method to include

the determination of vanadium in ferrovanadiun. (45), (63)
(0.33. 15,1985)
(0.1. 15, 5955)
A method for the determination of Vanadium in th 1e pres—

ence of iron and uranium was worked out by Gus tavson and
Knudsen. ”hev determined vanadium in a vanadyl su fate sol-
ution. 20 cc. of solution containing 18.7 milligrams of
vanadium, tooether 1"ith ferrous sulfate and uranyl acetate
was diluted to 250 cc. with water and five grams of zinc
added, and two pertions of four cc. each of sulfuric acid
were added durins the titration, one at th e start of the
titration, and the other at he end of the first end-point.
the solution was titrated directlv with oeriai anate. the
oxidation is selective, and the three elements may be deter-
mined in the same solution. (87) (0.3. 17, 507)
Willard and Fenwich used the polarized bimetallic elec-
trode system for the determination of vanadium in tne preser ice
of iron and chromium, and claim for the method a more sensit-
ive means of determinins the end-point when ferrous sulfate
is the titrating agent. The chromium was reduced, select-
ively, vith‘hydrogenperoxide in the presence of acetic acid

in 20-30 percent solution. 25 to 50 cc. of concentrated

67

hydrochloric acid was then added and the titration made with
ferrous sulfate. The authors have adapted the method for
determining vanadium in special alloy steels. (103)(C.d.l7, 507)
Ferrous solutions are more vigorous in their reducing
power than are vanadyl solutions. Then these solutions are
titrated together with permanganate two sudden chanpes in
e.m.f. are observed. The first change representins the oxi-
dation of the iron, and the second one to the oxidation of
the vanadium. The last part of the titration should be made
at 70°. (89) (C.A. 17, 1401)
Vanadic acid is reduced by cadmium in a Jones' reductor.
The reduced solution can be titrated with permanganate if
protected by carbon dioxide gas. The equation 5V*'+ 5Mn04-
. 5920 .: saves + sun"'+ H’. (99) (0.1. 17, 59)
Vanadium can be reduced to the bivalent condition in a
hinc reductor, and can be oxidized back to the quinquivalent
condition by means of permanganate. The titration is best
conducted at 800 temperature. Vanadium and iron can be tit-
rated together; first hot until two stages are conpleted --
oxidation of the vanadium; then cold for oxidation of the
iron, and hot again until Vanadium oxidation is complete.
(122) (C.l. 18, 645)
This reference contains a review of the literature on
electrometric titration of vanadium. Vanadium and chromium
Lay be determined together by titrating with ferrous amnonium
sulfate, protecting with carbon dioxide . Only one end-point

is obtained which corresponds to the known composition of the

68

mixture. After this determination has been made, the van-
adium is determined b;,7 ti tratins with perman sanate until

the iron is oxidized, heatins to 70-800, and finish'ng the
titration with permanganate. The reasent used between the
first and second break corresponds to the vanadium. Mangan-
ese does not interfere in this titration. A complete method
for the determination of vanadium and chromium in steel is

given. (144) ' (0,1. is, $571)

Zinc.- :27), (72)*, (77)*, (112)*, (114)*, (128)*, {154c)*

(1:8)*

Bichowsky used the platinum— calomel electrode system for
the determination of zinc, by titrating with ferrocvanide.
Zinc ore wa s put into solution with a minimum quantity of
strong hydrochloric acid, and a small quantity of potassium
chlorate added. 'ihe solution was then made a laline with
ammonia, an-aliquoit portion filtered off and acidified with
10 percent excess of hydrochloric acid. Sulfur dioxide {as
was then added to saturation, and the excess of this {as dis-
pelled by boiling. Test lead may be employed to free the
solution of cepper. The solution Was then heated to 700, and
titrated with ferrocyanide solution. if cadmium is present in
considera 1e quantities, it should be separated by some suit~'
able method. 10 to 2. percent of hydrochloric acid should be
present in the solution for best results. (27) C.A. ll, 2313

(C.A. 11, lane)

9

69
Zinc can be titrated electrometrically by means of pot-
assium ferrocyanide, the z'nc being precipitated as the double
ferrocyanide of potassium and zinc. A 1/40 molar solution of
potassium ferrocyanide, and a titrating temperature of 700
degrees is recommended. In a faintly acid solution two milli-
grams of zinc can be determined with an accuracy of one per-
cent. Sulfates in larse amounts interfere with the titration.
This titration can be reversed with accurate results. (77)
6.3. 16, 5281)
muller and Adam determine zinc by adding to a zinc sol-
ution an excess of cyanide with silver nitrate using a silver
electrode. Cadmium, lead, and copper can not be determined
by this method. (112) (0.1. 17, 2095)
Zinc can be determined electrometrically With potassiun
ferrocganide even in the presence of cadmium. (114)
(0.1. 17, 2689)
Zinc and lead may be determined electrometrically with
potassium ferrocyanide. to every 100 cc. of s lution one cc.
of M potassium ferricyanide is added and the titration made
with 0.1 M potassium ferrocyanide at 750. Near the first
break in potential, a wait of three to five minutes is nec-
essary in order for the potential to become constant. A first
treak occurs at - 0.54 volts and a second one at - 0.19 volts.
The results are exact for the sum of both metals, but not for

1

zinc alone, and attempts made at higher temperatures failed

m1

to produce good results. 1ne lead, however, Can be precipit-

ated with a slight excess of dilute sulfuric acid and the zinc

70

determined with ease. Therefore, two titrations are necessary
in order to successfully determine zinc and lead in solution
together. (128) (0.1. 17, 945)

In titrating zinc sulfate solutions with potassium ferro-
cyanide solutions, Kolthoff and Verzyl showed that 15° and
65° about one percent too little of the reagent is used when
the sudden potential change is observed. The addition of
three cc. of 0.25N sulfuric acid to 100 cc. of a solution
containing 1.08 grams of zinc sulfate changed the error to
0.72 percent, and the further addition of 3 grams of potass-
ium sulfate reduced it to 0.48 percent. (154)(C.A. 18, 1627)

Good results are obtained with a few grams of ammonium
sulfate in titrating zinc with ferrocyanide. (see preceding
abstract). The presence of large amounts of sodium, alum-
inum, magnesium, and calcium salts lowers the results. Mag-
nesium, cadmium, and cepper interfere with the method and
should be removed from the solution. Ferric compounds are
made harmless by adding ammonium fluoride and sulfuric acid.

(138) (C.A. 18, 2482)

M’CHIGAN AGRICULTURAL COLLEGE

 

 

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7 ‘

DEPAR'I ML». i' of MAI HLMAI IL:

78
VI

Curve Data and Conclusions

Note:- whe marginal numbers correspond to the curves on the

graphs.

1,2. 25 cc. of ferrous sulfate, 50 cc. of 4N sulfuric acid,

25 cc. of water titrated with 0.1N potassium dichromate.

3. 25 cc. of ferrous sulfate, 20 cc. of 4N sulfuric acid, and

30 cc. of water titrated with O.lN potassium bromate.

4. 25 cc. of ferrous sulfate, 20 cc. of 4N sulfuric acid, and

20 cc. of water titrated with 0.1N potassium permanganate.

Titration of potassium iodide -- 0.25584 grams in 100 cc.
of 10 percent sulfuric acid solution. --

5. lith potassium permanganate —- 3.161 grams per liter.

6. With potassium dichromate -- 4.903 " H n
7. With potassium brcmate —- 2.785 " H
8. With potassium iodate -- 5.567 " " "

Titration of potassium ferrocyanide -- 0.8448 grams in
100 cc. of 10 percent sulfuric acid solution. --

9. With potassium permanganate -- 5.161 grams per liter.

10. With potassium bromate -- 2.783 " n n
11. With potassium dichromate -- 4.903 " " "
Titration with potassium permanganate -- 5.161 grams per
liter.--
12. Of potassium ferrocyanide -- 0.8448 grams in 100 cc. of

10 percent sulfuric acid solution.

13.

14.

15.

16.

17a,
18a,

19.
20.

21a,

22a,

25a,

24.

25.

26.

79

0f ferrous sulfate -- 0.556 grams in 100 cc. of 10 percent
sulfuric acid solution.
0f potassium iodide -- 0.25384 grams in 100 cc. of 10 per-

cent sulfuric acid solution.

Titration with potassium bromate -- 2.785 grams per liter.
0f potassium ferrocyanide -- 0.8448 grams in 100 cc. of 10
percent sulfuric acid solution.
0f potassium iodide -- 0.25384 grams in 100 cc. of 10 percent

sulfuric acid solution.

Titration with potassium permanganate. --
17b. 0f vanadium and iron.
18b. 0f vanadium and uranium.
0f iron.
Of iron, vanadium, and uranium in 700. sulfuric acid.
21b,210. 0f iron, vanadium, and uranium in two additions

of sulfuric acid of 4 cc. each during the titration.

Titration with potassium dichromate. --
22b. 0f iodate in sulfuric acid.
23b. 0f bromate in sulfuric acid.

0f chlorate in hydrochloric acid.

Titration with potassium permanganate. --
0f iodide and bromide in the presence of chloride.

0f iodide in the presence of chloride.

27.

80

0.021 grams of chromium titrated with ferrous sulfate.

28a,28b. Ferrous and titancus ion titrated with potassium

permanganate.

29a,29b. Ferrous and titanous ion titrated with potassium

dichromate.

50a,50b. Reverse of curves 28a and 28b. Ferric ion and perman-

ganate titrated with titanous ion.

51a,5lh. Reverse of curves 29a and 29b. Ferric ion and dich-

32.

34.

35.

romate titrated with titanous ion.

75 cc. of 0.055 ferrous sulfate in a high concentration
of hydrochloric acid titrated with 0.1N potassium dichromate.
Electrodes -- platinum; 0.1N potassium chloride calomel

electrode.

25 cc. of 0.1N ferrous sulfate, 20 percent sulfuric acid,
titrated with 0.1M potassium permanganate.

Electrodes -- platinum; tungsten.

ferrous sulfate in 55 percent hydrochloric acid solution
titrated With 0.1K potassium dichromate.

Electrodes -- platinum; standard silver chloride half-cell.

Ferrous sulfate in 55 percent hydrochloric acid solution
titrated with 0.1N potassium dichromate.

Electrodes -- platinum; silver chloride platinum.

81

Terrous sulfate in 20 percent sulfuric acid titrated
with 0.1K potassium permansanate.

Electrodes -- platinum; silver chloride platinum.

250 cc. of solution containing 75 cc. of a 0.05m pot-
assium dichromate solution titrated with 0.1K ferrous
sulfate in different concentrations of hydrochloric acid.
0.4M hydrochloric acid.
0.8M hydrochloric acid.

2.0M hydrochloric acid.

Titration with potassium permanganate --

(a) Uranium solution -- 1 cc. equals 0.0151 grams.

(b) Permansanate solution -- 0.1011N.

(c) Ferrous solution -~ 0.1001N.

10 cc. of uranium solution, 40 cc. of 2 cc. concentrated

sulfuric acid in so of solution, 80-900 -- reduced.

0a,40b. Uranium titrated with permanganate.

41a,4l‘b. H H H H

42a,42b,42c. Uranium plus 1 cc. of ferrous sulfate solution

titrated with permanganate.

Uranium titrated with permanganate in various concentrat-

ions of sulfuric acid. --

45a,45b. 20 cc. sulfuric acid in 100 cc. total volume.

44.
45.

10 H " H H H '1 H H

5 7' H I? 'Y H H H H

468. , 46b 0 2 H " " N H n n n

47.
48.

49.

50.

51.
52.
55.
54.

55.
56.
57.
58.
59.
60.
61.

82

Potassium biniodate treated with excess iodide and

the excess iodide titrated with potassium permanganate -
0.02N iodate in 150 cc. with sulfuric acid.

0.05N " " 250 " " ” "

iodate titrated with permanganate in sulfuric acid.

4.1 mg. ferrous iron in 70 percent hydrochloric acid
titrated with 0.01 potassium dichromate.
4.24 mg. ferrous iron titrated with 0.01m potassium
dichromate.
In 67 percent sulfuric acid.

H 33 n n u

n 17 n n u

' 50 " hydrochloric acid.

Effect of adding ferric iron --
No ferric iron.
500 ppm. of ferric iron.
1500 ppm. of ferric iron.
2500 ppm. of ferric iron.

0.558 mg. ferrous iron and 1660 mg. ferric iron.
Ferric iron titrated with 0.005N stannous chloride.
Stannous chloride titrated with-0.01N potassium
dichromate.
Ferric iron reduced with stannous chloride in excess

and both titrated with 0.01 potassium dichromate.--

85

62a,62b. In 40 percent hydrochloric acid.

5a,65b. " 50 " n n
64a,64b. " 60 " " n
65a,65b. " 70 " vv n
66a,66b. " 90 " n n

Ferrous iron titrated with 0.01N potassium permanganate -

67. 2.87 mg. ferrous iron in 17 percent sulfuric acid.

68. 2.87 mg. ferrous iron in 10 percent sulfuric acid with
28.7 mg. of ferric sulfate.

69. 2.87 mg. ferrous iron in 2.5 percent of 85 percent phos-
phoric acid with 57.5 mg. of ferric sulfate.

70. 2.87 mg. ferrous iron in 2.5 percent of 85 percent phos-
phoric acid with 57.5 mg. of ferric sulfate and 5 percent
sulfuric acid.

71. 2.87 mg. ferrous iron and 6 percent of the phosphoric-
sulfuric acid mixture with 250 mg. of iron as ferric
chloride.

72. 2.87 mg. ferrous iron and 6 percent of the phosphoric-
sulfuric acid mixture with 250 mg. of iron as ferric

chloride, and 5 percent of hydrochloric acid.

Ferrous iron titrated with potassium dichromate in
various concentrations of hydrochloric acid. --
75. 50 cc. concentrated hydrochloric acid, 100 cc. water,
1.4 gm. ferrous ammonium sulfate.
74. 50 cc. concentrated hydrochloric acid, 100 cc. water,

1.05 gm. ferrous ammonium sulfate.

84

75. 100 cc. concentrated hydrochloric acid, 50 cc. Water,

1.05 gm. ferrous ammonium sulfate.

76. 10 cc. 0.1N potassium bromide titrated with approximately
O.lN mercurous nitrate.
77. 10 cc. 0.1N potassium chloride titrated with approximately

0.1N mercurous nitrate.

A general consideration of these electrcmetric titration
curves shows that it is quite impossible to obtain from them
distinct theoretical correlations of first importance. Nearly
all of the investigational work pertaining to them has been
done from the standpoint of analytical chemistry alone; and
the chief aims of the investigators seems to have been, to
discover suitable reagents for the titrations, to obtain
well-marked end-points, and to determine small quantities
of unknowns. In general, it may be stated that the end
sought was rapidity and refinement of analytical methods.

To this end, various kinds of electrodes, various concent-
rations of reactants and of acid solutions, and different
mechanical operations have been employed; all of which is
contrary in principle to the systematic research necessary
in this kind of work in order to establish fundamental chem-

ical concepts.

Some of the points taken under consideration are as

follows:

85

l. The end-point placement on the graph with reference to
the zero e.m.f. and the magnitude of the deflection of
the curves. .

2. is there a "neutral point"?

5. Correlation with the Periodic Law.

4. Slope of curve deflections or Speed of the reactions with
reference to kind and strength of acid used.

5. Energy relationships.

6. Chemical activities.

7. Summary of rules for electrometric titration.

8. Scope and Opportunities in the field of electrometric

titration.

A cursory examination of the curves will suffice to
show that no distinct relationship exists between the mag-
nitude of their end-point deflections and their position on
the graphs with respect to the zero e.m.f. point. However,
from the standpoint of their position on the graphs alone,
the curves can be classified, roughly, into three groups:
1. Those ranging from 400 millivolts up. 2. Those occupying
‘the intermediate position of O to 400 millivolts. 5. Those
having the e.m.f. range of 0 down to the lowest negative
voltage obtained.

A list of the reactions of class 1. --

Curves l to 4 -- oxidation of ferrous iron.

curves 5 to 8 -- oxidation of potassium iodide.

UUI‘VGS

Curves 12 to 16 -- oxidation of ferrocyanide,

Curve 210

9 to 11 -- oxidation of

86

ferrocyanide.

ferrous iron,

and potassium iodide.

-- oxidation of uranium.

Curves 22b to 25b -- reduction of iodate and bromate.

Curves 25 to 26 -- oxidation of iodide and bromide.
Curve 27 -- reduction of chromium.
Curves 28b and 29b -- oxidation of ferrous iron.
Curves 50b and 51b -- reduction of ferric iron.
Curves 52 to 56 -- oxidation of ferrous iron.
Curves 40b, 41b, and 420 -- oxidation of uranium and ferrous
iron.

Curves 47 to 49 -- oxidation of iodide.
Curves 50 to 59 -- oxidation of ferrous iron.
Curves 62b to 66b -- oxidation of ferrous iron.
Curves 67 to 75 -- oxidation of ferrous iron.

A list of the reactions of class 2. --
Curves 20 to 21b -- oxidation of vanadiun.
Curves 17b, 18b -- oxidation cf vanadium.
Curve 19 -- oxidation of ferrous iron.
Curves 28a, 29a -- oxidation of titanium.
Curves 50a, 51a -- reduction of titanium.
Curve 55 -- oxidation of ferrous iron.
Curves 57 to 59 -- reduction of dichromate.

Curve 60

reduction of

ferric iron.

87

oxidation of tin.

Curve 61 -

Curves 62a, 66a reduction of ferric iron.

Curve 76 bromide and mercurous chloride.

Curve 77 - chloride and mercurous chloride.

A list of the reactions of class 5. --
Curves 17a, 18a, 21a -- oxidation of vanadium.
Curves 40a, 41a, 42a -- oxidation of uranium.

Curves 44, 45, 46a -- oxidation of uranium.

The basis for the above classification is apparent, and
can be referred to the initial and final electrode potentials
of the several substances used in the titrations. The high-
est voltage is obtained with permanganate, and the lowest
with uranium in a low state of oxidation. From these data
we can conclude that permanganate is a strong oxidizer and
uranium a strong reducer. metals with two or more states of
oxidation, or valences, will give end-points for these differ~
ent states, and the end—point curves will occupy different
positions on the graph, the lowest oxidation state correspond-
ing to the lowest position on the graph. The electrode pot-
ential may be negative for a low state of oxidation and pos-

itive for a higher state of oxidation.

The electrode potential of a solution measures 'ts oxid-
izing or reducing power. Therefore, in general, the greater
the difference in electrode potential between an oxidizing

solution and a solution to be oxidized, the greater the act-

88

ivity of the oxidizing reaction and hence the greater the
'speed of reaction and the sharper the end-point deflection

of the measuring potentiometer. The same conditions apply

to reducing reactions. in both cases, however, the strength
of the acid titrating solution is an important factor, as is
also the kind of acid used to acidify the titrating solution.
The disturbing influences are the electrode potentials of the
products of the reaction, and the kind of electrodes used in

the determinations.

The titration of uranium with permanpanate may be taken
as illustration of the consideration. Curves 43a
and 43b represent this titration in 20 percent sulfuric acid
solution. In this case the change in potential is great
being from about -530 to +1150 millivolts for both stages of
oxidation of the uranium compound; but the endpoint deflect-
ions are not sharp. in a 2 percent solution of sulfuric acid,
however, as illustrated by curves 46a and 46b, about the '
same change in voltage is observed while the two end-points
are very distinct.

Another illustration of this same phenomenon may be
'found in the titration of stannous and ferrous ion with dich-'
romate - curves 61 to 66. The influence of kind of acid is
illustrated in the titration of ferrous iron with dichromate -
in both sulfuric and hydrochloric acid solution - curves 51
to 54 - in which ease the hydrochloric acid solution gives

the best end-point.

89

‘In case certain reactions do not give suitable end-points,
these end-points may be improved by reversing the reaction.
Curves 28 to 31 illustrate a case of this kind. Ferrous and
titanous ion titrated with permanganate or dichromate do not
give suitable curves; but when ferric ion and dichromate or
permanganate is titrated with titanous ion, sharp end-points
are obtained in case of the ferric ion and permanganate, and
for the ferric ion alone in case of the dichromate. This
eXperiment shows, also, that titanous ion will reduce ferric
ion, or that one metal ion will satisfactorily reduce or
oxidize another metal ion provided that the values of elec-

trode potentials are far enough apart.

For slow oxidation or reduction reactions, or in cases
where sharp end-points may not be obtained, it is often ex-
pedient to add a known excess of the oxidizing or reducing
reagent, and after sufficient time has elapsed for the re-
action to be completed, this excess of reagent is titrated
by means of another quickly acting reagent that does give
a sharp end-point, and the unknown ion calculated from the
difference between the ouantity of reagent first added and
the excess found by titration. Again, the oxidation process
may be easier applied to a certain metal than the reduction
process in titration. in this case, the metal is first
reduced by some knoun efficient proéess after which it is
titrated with an oxidizing reagent, or this process may be

reversed.

90

in only one case does the end-point e.m.f. approximately
coincide with O e.m.f.; that is in the reduction of perman-

ganate with titanous ion in the presence of ferrous ion.

The number of elements investigated is too small to
provide a means of correlation with the periodic law. More-
over, oxidizine or reducing power may be only relative; for
instance, ferrous iron may reduce permanganate and oxidize
a stannous salt. A metal in a low state of oxidation may
be caused to reduce the same metal in a higher state of

oxidation.

The formula eXpressing change in free energy for an
electrometric reaction change is ~45F — E.N F. While
the change in-energy accompanying a change in valence for
1 gram mole is l Farady, it is now known if this value Can
be eXperimental y realized. The difficulties enco‘ntored
in electrometric titration processes are the unknown liquid-
liquid potential values, and the unknown values of chemical
activities of electrolytes in the presence of other electro-
lytes - acids or salts. The data at hand give no values
that may be used to calculate energy changes from a theor-
etical standpoint. The magnitude of the curve deflections

is, however, an effect of this free energy.change.

From an examination of the data on electrometric tit-
ration methods, with special reference to oxidation or re~

duction reactions, some rules and precautions may be evolved

91

as follows: 1. Choose active oxidizinr and reducing reagents.
2. Suitable metallic electrodes are of importance. 5. Select
chemical reactions that are complete. 4. Slow reactions may
be used if an excess of reasent is added and this excess
back-titrated or titrated With a third substance. 5. A pro-
per temperature is important for some reactions. 6. The
interaction of components should not disturb the titration
process. 7. A certain strength and kind of acid will be
found necessary for sharp end-points. 8. Avoid substances
that interfere with the reaction -- external oxidizing or
reducing substances. 9. Weighing the solutions is more
accurate than measurino them, especially for the determin-
ation of small quantities of unknowns. 10. All reasents
should be accurately standardized by the best known methods,
and by applying the electrometric titration method, in the
same manner as the resular titration is to be conducted,

when possible.

Only a limited number of elements and chemical com-
pounds have been studied by means of the electrometric
titration methods, and an opportunity exists to study those
elements that have not already been employed in these methods.
very little work has been done on phoSphates, carbonates,
sulfides, chromates, and oxalates, and the formation of
colloids by means of oxidation or reduction methods. Another

possible field for the use of these methods lies in a study

92

of the effects of catalyzere on chemical reactions and the
energy changes involved. Almost no work has been done with
these methods in the field of organic analyses, except on

aldehydes and reducing sugars.

...]

VII
BIETIOJRAPUY

1892
Oxidniion cells.

13, EE7-408.

1835
leotronetric

Chef]. , 11. 4.36.910

9

The potential of hydrofo
neuman. Z. physig. Chem.,
1937

ihe use of the eloctronoter as

of 7.

“016° -‘- 1r“ hunec‘
u \L- 90-“. kafi‘ k. .

53- :01.

r
Lad

A gaannoneter titrition

Oxidation nnn reduction cells,

plex ions

Z. physik. Chen., 25, 195- 2 6.
190?

Stvdies on oxidation potentials.

Chem., 24, 2? 5- -62.

W. D. Eencroft. Z.

Robert

and several

Bottfer.

notk(d.

nnd

on *Veir e1 eotroitlotive

phfi’si};

.3
"J

Tehrei . Z

P

14,1JoJ-2500

indicator in

"o . 1'
I

(‘1' \Jllx.‘

I

Z.

pk

Salomon. 3

force. Ru

metals.

th e inflrunune

95

rw.
. b:](:"—A1. ,

Bernhard

tho titration

-\

f‘ t

blflleijo , L1.

. tilelctrcwflfcx1.,

of com-

CO

dolph Feter

9b.
10.

11.

‘4.

1'7
1‘.

14.

15.

94

1902

The theory of oxidation and reduction cells. Carl

0

Fredenhngen. Z. anorg. Chem., 49, 538-463-

1911

m 1

ine employment of the electrometric method for the

~

estimation of the acidity of tan liquors. H. c. Qand
Lind U0 ('0 T131137. . :0 SOC. Chcji'lo lnlo , 3"), 30
III. Frecipitation followed ty change of potential.

1?. Determination of copper and silver.

7. Titration and separati(n of the halides. Dntoit and

”9 q
C
I

n , 7 ‘ fir:-
sse. .. Chen. phgs. 9, 516

all

H.

C ‘64'... (COA. 6, 722‘
The titration of potass'um cyanide in the presence of
potassium ferrocynnide. 7. D. Treadwell. u. anorf. Chem..

71, 219-225. ’C.A. 5, 55911

1913
Gene applications of the hydrogen electrode in analysis,

\
7v 3

research and teaching. Joel i. Hildebrand. J.n.C.S.,
:5, £347-71. ‘ (0.31. 7' £307"),
The increase in the oxidizing potential of the dichromate

ion on platinum Caused ty certain reducing aren s. an

improved method for the electrometric determination of

ti]

ferrrvs salts. 1. 3. Forbes and . P. ?artlett. J.A.C.S..

I .

35 1527-58. (C.A. 7 5v

()1

a)

.
Separation of chlorides and hronides by silver nitrate
accordine to the conductivity method._P. Dutoit and Reeb.
463. (C.A. 8, 472i

9

17.

18.

19.

F;
‘9

0

‘5

‘u ‘J.

(‘7)
()1

(“D

95

1914

 

Automatic titrations. 1. Automatic electrometric tit-

ration of dichromate and ferrous iron. E. Tierel. Z.
anal. Chem., 53, 755-62. (C.A. 9, 838)

The electro-titrametric method and its application to
general analytical chemistry. T. P. Vesselink van
Vuchtelen & Arno ltano. J.A.C.W., 36, 1795-1805.

(Co-(A10 8, 341““)

1916

 

conductivity cell for electro~titration. H.E.Rohhins.
An. Cour. 501., 41, 249-50. {0.1. 10, 987)
The determination of chromium and vanaditm in steel by
electrometric titration. u. L. Kelley & Joe. E. Conant.
ind. Eng. Chem. 8, 719-23. ‘ (0.3. 10, 2179)

-

Electronrtric titration of vanadium. G. n. Kelley &

.09. L. Conant. 3.3.0.3., 38, 341-51. (C.A. 10, 729)
The volumetric determinaticn of alkaloids ty physio-

chemical methods. P. Dutoit & leyer-Uev*. Jour. Chem.

ehvs., 14, 553-60. (0.1. 11. 1eoo)

o)

1 17

KO

 

new form of conductivity cell for electro-titration.

F. E. Robbins. 3.3.0.3., 59, 646-8. 'C.A. ll, 1062)
Electrometric methcd of titration and its application
to the examination of gastric juices. Leonor Lichaelio.

tiochem. z., 79, 1-54. (0.1. 11, so 3

(‘3

‘4

['21

'7

CC]

C)
O

07.

l.

96

1917L7cont.

 

Convenient a paratuo for electrometric titration depend-
ins on cha nse of oxidation potential. Determination of

fl -

chronium in steel. ~. u. Lelley, J. R.

(i

small amounts oi
Adams a C. A. Wiley. 1nd. Ens. ihen., 9, 730-2.

10.1. 11, £567)
Conductivity measurements upon oxidation-reduction re-
actions. ”rahen Edgar. 3.1.0.3., 59, 914-28.(C.L.11,15J0.

11 1'"

'.1:e elee.ro1etric titration of zinc. 1. us sell V. Bich-
OWO1::’7. ind. Eng. Uhefll., 9, 668-71. (00“. 1.1, 2:: f3)
:. “fad-9h. .‘lxcad. r:Cj.. , 7, 141.3. (8.1L. 1]., lEOC’)

‘

Tne determination of iron in glass

('3

and. {ohn B. Ferguson.

'1 '3' '7' T; (‘1
J. , (_1 L-" . C

()1

r1 A '1
. J-

8'
p

Ind. Eng. 0hem., 9, 941-
The meas Ireu1ent of oxidation potentials at mercuzy cloc-
trodes. G. S. Forber & H. 7. Richter. 5.1.0.3., SJ, 1140.

(C.A. 11, 2365)

1J18

 

The determination of manganese in steel in the preser cc
of chromium and vanadium by electrometric titration. G.

t. Kelley. I-I. G312 ncer, C. 3. 11111151101111, and '3. Gay.

F-‘I

nd. Eng. Chen., 10,19J-2.. (0.A. 12, 4581

a method for making electronotric titrations of solutions

conta inino proteins. iohn C. Baler & Lucius 1. Van Dyke.
(our. Eiol. 0hen., 35 157-45. (0.4. 12, 1383)

New reduction methods in volvnetric a11alys's. inc 1t &

Hibhert. p. 65. Zonsmans, Green a Company.

35.

56.

58.

59.

40.

41.

42.

43.

97

1919

 

Electrometric titrations. I. D. Treadwell & L. Weiss.
Helv. Chim. Acta., 2, 680-97. (C.A. 15, 1262)
The theory of electrometric titrations. 7. D. Treadwell.
Helv. Chim. Acta., 2, 672-80. (C.A. 15, 1265)
End point in oxidation titrations determined by means

of the potentiometer. I. M. Kotthoff. Chem. Weekblad,

16, 408-16. (C.A. 15, 1263)
Electrometric method for the determination of ferrocyan-
ides depending on a change in oxidation potential. 'G. L.
Kelley & R. T. Bohn. J.A.C.S., 41, 1776-83.(C.A. 14, 30)
Electrometric titration of mixtures of acids. P. A. Meer-
burg. Chem. Weekblad, 16, 1358-47. (C.A. 14, 257)
The application of electrometric titrations. J. Pinkhoff.
Chem. Weekblad, 16, 1165-7. (C.A. 14, 257)
The estimation of ferrous iron by electrometric titration.
I. M. Kolthoff. Chem. Weekblad, 16, 450-61.(C.A. 14, 3204)
Electrometric analysis with potassium ferrocyanide. E.
Mflller. angemo. Chem., 32, I, 351-2. (C.A. 14, 5205)
The electrometric titration of plant juices. A.R.C. Uass.
Soil Sci., 7, 487-91. (C.A. 14, 420)
Titration by means of changes in potential. J. Pinkhoff,
Pharm. Weekblad, 56, 794. (C.A. 13, 1793)
The determination of vanadium in steels by electrometric
titration. The selective oxidation of vanadium salts by

nitric acid in the presence of chromic salts. 0. L. Kelley,

98
1919, cont.
J.A. Wiley, R.T. Bohn, & W. C. Wright. Ind. Eng. Chem.
11, 632-34. (0.4. 13, 1985)
44. Electrometric titration of iodides. i. M. Kolthoff. Chem.
Weekblad, 16, 926-9. ’C.A. 13, 2167)
45. Electrometric titrations, with Special reference to the
determination of ferrous and ferric iron. J. C. Hastetter
and H. L. Roberts. J.A.C.S., 41, 1337-57. (C.A. 13, 2319)
46. The electrometer as an indicator for titrations. J. Pink-
hoft. Pharm. Weekblad, 56, 1218-34. (C.A. 13, 3102)
47. Electrical apparatus for use in electrometric titrations.
R. S. Roberts. J.A.C.9., 41, 1358-62. (C.A. 13, 2302)
48. The effect of dilution in electro-titrimetric analyses.
G. A. Freak. Jour. Chem. 800., 115, 55-61.(C.A. 13, 2645)
49. The rapid electrometric determination of iron in sone
Optical glasses. J. B. Ferguson and J. C. Hostetter.
Jour. Am. Ceram. Soc., 2, 608-21. (C.A. 13, 2580)
49a. Rapid determination of chromium in steel by electrometric
titration. G. L. Kelley and W. C. Wright. iron Age, 103,
1507. (C.A. 13. 1802)
50. A simple method for the analysis of bearing metal and
similar alloys. 0. Oesterheld and P. Houegger. helv. Chim.
Acta. 2, 398-416. (0.3. 13, 3106)
51. G. Redrick. Diss. Dresden.

52. J. Pinkhoff. Diss. p. 44. Amsterdam.

99

1920

 

53. The oxidopotentiometric titration of iodides in the
presence of chlorides and bromides. I. M. Kolthoff.
Rec. trav. Chim., 39, 208-4. (C.A. 14, 3204)
54. The use of electrometric titration methods. 0. 0. Bann-
ister. Electrician, 85, 535-7. (C.A. 15, 352)
55. Review of the application of potentiometer titrations.
I. M. holthoff. Chem. Weekblad, 17, 659-64.
(C.A. 15, 996)
56. The electrometric titration of phenols. i. M. Kolthoff.
z. anorg. allegem. Chem., 112, 187-95. (0.3. 15, 1669)
57. The electrometric titration of alkaloids and their salts.
I. M. Kolthoff. Z. anorg. allegem. Chem., 112, 196-208.
(0.4. 15, 1670)
57a. The determination of hydrogen ions. W. M. Clark. Chapt.
XIV. The relation of hydrogen electrode potentials to
reduction potentials.
58. The acidimetric estimation of heavy metals in their salts.
1. M. holthoff. z. anorg. allegem. Chem., 112, 172-86.
(0.2. 15, 1670)
1921

 

59. The electrotitration of hydriodic acid and its use as
a standard in oxidimetry. W. 9. Hendrixson. J.A.C.S.,
43, 14-23. (C.A. 15, 639)
60. The theory of electrometric titrations. W. D. Treadwell.

Relv. Chim. Acta., 2, 672-80. (0.1. 15, 1263)

61.

62.

63.

64.

67.

68.

100

1921, cont.

 

Electrometric determination of bromate, dichromate,
nitrite, and chloride ions. 7. S. Fendrixson. J.A.C.S.,
43, 1509—17. (C.A. 15, 3793)
The electrometric titration of iodides by means of per-
manganate. 1. M. Kolthoff. Rec. trav. Chim., 40, 532-8.
(0.4. 15, 3799)
The determination of vanadium and chromium in ferrovan-
adium by electrometric titration. C. L. Kelley, J. A.
Wiley, R. T. thn, and 7. C. Wright. Ind. Eng. Chem.,
13, 939-41. (C.A. 15, 3955)
The determination of chromiun in ferrochromium by elec-
trometric titration. G. L. Kelley and J. A. Wiley.
1nd. Eng. Chem., 13, 1053-4. (0.4. 16, 37)
Simple and rapid electrometric method for the deter-
mination of cobalt in an ammonical electrolyte and its
recovery from cobalt nitro-B-naphthal. K. Wagonmann.
Metall u. Erz, 18, 447-9. (C.A. 16, 695)
The acidimetric titration of dichromate. I. m. Kolthoff
and E. H. Vogelenzanz. Rec. trav. Chim., 40, 681-5.
(C.A.—l6, 1056)
The determination of iodic acid and silver by electro-
metric titration. W. S. uendrixson. J.A.C.%., 43, 858-66.
(0.4. 15, 1668)
Oxidimetric determination of manganese in hydrofluoric
acid solution. I. Josef. Halluta and Josef. Christ.

nonateh., 41, 555.71. (0.2. 15, 2800)

69.

70.

71.

72.

73.

74.

75.

76.

77.

101

1921LAcont.

 

The electrometric titration of hypochlorous acid. W. D.
Treadwell. Helv. Chim. Acta, 4, 396-405. (0.4. 15, 3052)
Electrometric control in chemical industry. Erik. K.
Rideal. Chem. Age, 5, 232-3. (0.4. 15, 3797)
Modified laboratory apparatus. 0. 0. Kiplinger. Ind. Eng.
Chem., 15, 715. , (0.4. 15, 3569)
Reductions with zinc and cadmium in volumetric analysis.
7. D. Treaduell, n. Luthy, and A. R. Rheiner. Helv. Chim.

Acta, 4, 551-65. (0.4. 15, 5041)

1922

 

The application of amalgams in volumetric analysis. III.
Estimation of iodic, bromic, and chloric acids. Suetaro
Kikuchi. J. Chem. Soc. Japan, 45, 175. (0.4. 16, 1716)
The use of the iodine (I2) electrode in potentiometric
titrations. l. M. Kolthoff. Rec. trav. Chim., 41, 172-91.
(0.4. 16, 2277)
Electrometric titration of uranium with potassium per-
manganate and potassium dichromate. D. T. Ewing and E.
F. Eldridge. 3.4.0.3., 44, 1484-9. (0.4. 16, 2459)
The electrometric determination of cyanogen in the pres-
ence of halogens. Erich. muller and Hans Lauterbach.
Z. anorg. allpun. Chem., 121, 178-92. (C.A. 16, 2818)
The use of potassium ferrocyanide in potentiometric

titrations. I. The titration of potassium ferrocyanide

78.

79.

£1
80.

81.

82.

83.

84.

102

1922, cont.
by means of potassium permanganate. I. M. Kolthoff.
Rec. trav. Chim., 41, 343-52. (0.4. 16, 3281)
II. The potentiometric titration of zinc. Ibid. 41,
425-37. (0.4. 16, 3281)
Electrometric titration of ferrocyanic acid. Erich
Muller and Pans Lauterbach. Z. anal. Chem., 61, 398-403.
(0.4. 16, 3601)
The electrometric titration of dichromate with ferrous
sulfate. Jarion Eppley and 7. 0. Vasburgh. J.4.0.S., 44,
2148-56. (0.4..16, 3931)
lectrometric titrations. 4. Lassieur. Bull. soc. Chim.,
31, 817-31. (CJA. 17, 37)
Bimetallic electrode systems in electrometric analysis.
I. Systems comprising two dissimilar metals. H. R.
Willard and F. Fenwick. J.A.C.S., 44, 2504-15; 2516-19.
(0.4. 17, 37;38)
The electrometric determinations of end points. E. Muller.
3. ans. Chem., 35, 563-6. (0.4. 17, 38)
The electrometric standardization of titanous solutions.
W. S. Rendrixscn and L. M. Verbeck. J.4.C.S., 44, 2382-7.
(0.4. 17, 38)
Electrometric titration of sulfurous acid with perman-
ganate. W. S. Hendrixson and L. M. Verbeck. Ind. Eng.

Chem., 14, 1152-5. (0.4. 17, 59)

85.

86.

87.

88.

89.

90.

92.

94.

105

1922, cont.

Electrometric determination of nickel with silver nitrate.
Erich. muller a Bans Lauterbach. Z. anal. Chem., 61,
457-64. (0.4. 17, 249)
Some applications of the oxygen electrode, air electrode,
and oxidation potential measurements to acidimetry and
alkalimetry. N. H. Furman. J.4.0.S., 44, 2685-97.

(0.4. 17, 505)
The successive electrometric titration of iron, vanadium,
and uranium. R. T. Gustavson and 0. M. knudson. J.4.0.S.,
44, 2756-61. (0.4. 17, 507)
Potentiometric titration of copper. E. Lintl and E.
Wattenberg. Ber. 558, 3366-70. (0.4. 17, 702)
The electrometric determination of iron and vanadium
when present together. Erich Muller and H. Just. Z. anorg.
allegem. Chem., 125, 155-66. (0.4. 17, 1401)
Acidimetric titration of magnesium in its salts. I. M.
Kolthoff. Rec. trav. Chim., 41, 787-94. (0.4. 17, 2092)
4 simple method of electrcmetric titration in acidimetry
and alkalimetry. P.F.Sharp and E.B. McDougall. J.4.C.S.,
44, 1193-6. (0.4. 16, 2277)

4 continuous reading electrotitration apparatus. K. R.

‘GCOde. JOAQCOS.’ 44, 26-9. (GOA. 16, 665)

Electrometric titration of azo dyes. D.0. Jones and H.R.

189. Ind. Eng. Chemo, l4, 46*8. (00A. 16. 1155)

95.

96.

97.

98.

99.

100.

101.

104

1922, cont.

 

Electrometric titration as a means of determining the

free sodium sulfide in a sulphur black dye bath. W. W.

Russell and T.T. Arnold. 4m. Dyestuff Rep., 10, 346;

375-6; 451-2. (0.4. 16, 3208)

Construction of platinum coated glass electrodes and

methods of measuring with them. 4. Eilert. Z. angew. Chem.,

35, 445-6; 452-5. (0.4. 16, 4095)

Titrations in ethyl alcohol as solvent. Edna R. Bishop,

Esther B. Kittredge, and Joel E. hildebrand. J.4.0.S.,

44, 155-40. (0.4. 16, 885)

The effect of alkali on the titration of certain metals

with ferrocyanide. W. D. Treadwell and D. Cheruet. nelv.

chim. 40ta, 5, 633-9. (0.4. 17, 38)

Reductions with cadmium in volumetric analysis. W. D.

Treadwell. nelv. Chim. Acta, 5, 732-43. (0.4. 17, 39)

4 simple apparatus for electrometric titration. 7. E.

Garner and C. 4. Waters. Jour. Soc. Chem., 41, 337-8.
(0.4. 17, 229)

4 new vessel for electrometric titration. W. T. Eovie.

J.4.C.3., 44, 2892-5. (C.A. 17, 654)

1925

 

Electrometric acidimetry and alkalimetry without the use
of hydrogen. P.A. Van der meulen and F. Wilcoxson. Ind.

Eng. Chem., 15, 62-3. (0.4. 17, 505)

105.

104.

105.

106.

107.

108.

109.

110.

105

1923L cont.

 

4 new method for the electrometric titration of van-
adium in the presence of iron and chromium. H. H.
Willard and F. Fenwick. 3.4.0.3., 45, 84-92.

(0.4. 17, 507)
The bimetallic electrode applied to neutralization
reactions. H. R. Willard and F. Fenwick. J.4.0.S., 45,
715-16. (0.4. 17, 1396)
The electrometric titration of the halides in the pres-
ence of one another. R. H. Willard and F. Fenwick. J.4.
0.8., 45, 623-33. (0.4. 17, 1397)
Electrometric determination of sulfur in soluble sulfides.
8. R. w'illard and F. Fenwick. J.4.C.S., 45, 645-9.

(0.4. 17, 1398)
Potentiometric titration of arsenic and antimony. E.
Lintil and H. Wattenberg, Ber. 56, 472-80.

(0.4. 17, 1604)
Electrometric titration of molybdenum with a titanous
salt. H. R. Willard and F. Fenuick. J.4.C.S., 45, 928-55.

(0.4. 17, 1766)
The electrometric titration of selenium in the presence
of tellurium, iron, and 00pper. H. H. Willard and F.
Fenwick. J.4.0.5., 45, 933-9. (0.4. 17, 1766)
An electrometric study of the neutralization of phos-
phoric acid by calcium hydroxide. G. R. Tendt and 4. R.

Clarke. 3.4.0.8., 45, 881-7. (C.4. 17, 1767)

111.

112.

113.

114.

115.

116.

117.

118.

119.

106

1923, cont.

 

Electrotitration with the aid of the air electrode. N.
H. Furman. Trans. Amer. Electrochem. Soc., 43, 1923.

(0.4. 17, 1930)
Electrometric determination of zinc with silver nitrate.
Erich. Muller and 4. Adam. Z. Electrochem., 29, 49-53.

(0.4. 17, 2093)
Electrometric titration of reducing sugars. Wanda E.
Daygett and 4. W. Campbell with J. L. Whitman. 3.4.0.8.,
45, 1043-5. - (0.4. 17, 2545)
The electrometric titration of zinc and cadmium. T.
Muller. z. anorg. allgem. Chem., 128, 125-30.

(0.4. 17, 2689)
Electrometric titration of acids and bases with an anti-
mony indicator electrode. Alfred Uhl and Wilhelm Kes-
tranck. Monatsh. 44, 29-34. (0.4. 17, 3303)
Electrometric titration of iodate, bromate, chlorate,
and ferrocyanide with titanous sulfate. 7. S. Fendrixson.
3.4.0.8., 45, 2013-7. (0.4. 17, 3306)
The electrometric determination of manganese. Eric. muller
and 0. Wahle. Z. anorg. allgem. Chem. 129, 33-40.

(0.4. 17, 3306)
Recent applications of electrotitrimetry to chemical
analysis. Jean Barbandy. Technique Moderne, 15, 545-53.

(0.4. 17, 3846)
Simultaneous electrometric determination of iron and
manganese, Erick. Muller and 0. Wahle. Z. anorg. allgem.

Chem., 130, 63-8. (0.4. 18, 363)

122 O

125.

124.

125.

126.

107

1923, cont.
Simultaneous electrometric determination of copper and
silver with potassium thiocyanate. Erick. Muller and

4. Rudolph. 3. Anal. Chem., 63, 102-11. (C.4. 18, 640)

4 titrometric micro-method for the determination of
sodium. E. muller. Helv. Chim. Acta 6, 1152-61.

’ «0.1. 18, 641)
Electrometric determination of vanadium and uranium

alone, together, and in the presence of iron. Eric.

Muller and 4. Flath. 3. Electrochem. 29, 500-8.
(00:10 18, 643)

Electrometric method of measuring acidity and alkalin-

ity. H. T. Bovie. J. Optical Soc. 4mer., 8, 149-68.
(0.4. 18, 797)
The titration of ferric chloride with sodium hydroxide,
usins the oxygen electrode, a proof of the non-existence
of iron oxychloride. R. B. Smith and P. M. Giesy. J. 4m.

Pharm. Assoc., 12, 855-6. (0.4. 18, 1099)

Simultaneous electrometric determination of iron and

manganese. Eric. Muller and O. Wahle. Z. anorg. allgem.

Chem., 132, 260-4. (0.4. 18, 1628)
The use of the mercury electrode in potentiometric

titrations. Determination of halides, cyanides, sulfides,
and thiosulfates. I. M. Lolthoff and E. J. A. H. Verzyl.

Rec. tran. chim., 42, 1055-64. (0.4. 18, 1799)

127.

128.

129.

130.

132.

133.

134.

108

1923, cont.

Electrometric determination of cobalt with silver nitrate.

Eric. Muller and Kust Gabler.:3. anal. Chem.. 62, 23-8.
(0.4. 17, 942)

The electrometric titration of zinc and lead in the

presence of one another with potassium ferrocyanide. E.

Muller and Kust Gablerx Z. anal. Chem., 62, 29-34.

(0.4. 17, 943)

The titration of hypochlorous acid. 4. Schleichter.

z. anal. Chem., 62, 329-35. (0.4. 17, 2543)
Standardization of solutions used in iodimetry. 3. Pop-
off and F. L. Chamhers. 3.4.0.5., 45, 1358-60.

1 (0.4. 17, 2247)

‘The titration of solutions of permanganate and sodium
arsenite. 7. T. Hall and C. E. Carlson. 3.4.0.9., 45,
1615-20. (0.4. 17, 2687)
The titration of silver ions and chloride ions in the

presence of protective colloids. W. D. Treadwell, S.
Janett, and M. Blumenthal. Helv. Chim. 4cta, 6, 13-8.
{0.4. 17, 3003)
The influence of alkali on the titration of certain
metals with ferrocyanide. W. D. Treadwell and D. Chervet.
Helv. Chim. Acta, 6, 550-9. (0.4. 17, 3005)
An electrometric method of following certain inorganic

hydrolytic reactions. G. S. Tilley and O. C. Ralston.

Trans. Am. Electrochen. Soc. 44 (preprint)(0.4.17, 2836)

109

 

1923LAcont.
134b. An application of the vacuum tube to chemistry. D. F.
Calhane and R. E. Cashing. 1nd. Eng. Chem., 15, 1118-20.
$0.4. 17, 3846)

1924.

 

1340. The accuracy of the potentiometric titration of zinc
with ferrocyanide. I. M. Zolthoff and E. 3. 4. H. Verzyl.
Z. anorg. allgem. Chem., 132, 318-20. (0.4. 18, 1627)
135. The electrometric determination of soluble sulfates.
E. muller and R. Wertheim. z. anorg. allgem. Chem., 133,
411-6. (0.4. 18, 1626)
135b. Electrometric titration of mercury with ammonium thic-
cyanate. R. Muller and O. Eeuda. z. anorg. allgem. Chem.,
134, 102-4. (0.4. 18, 1959)
136. The potentiometric determination of chlorides in the
presence of colloids. P. Siebert. Chem. Weekblad. 21,
167-9. (0.4. 18, 1960)
137. The estimation of nitrates by electrometric titration.
3. B. Robertson and 4. 3. Belling. 3. S. African Chem.
Inst., 7, 9-13. (0.4. 18, 1960)
138. The practical utility of the potentiometric titration
of zinc with ferrocyanide. E. 3. 4. H. Verzil and I. M.
Kolthoff. Rec. tran. chim., 43, 380-8. (0.4. 18, 2482)
139. The electrometric titration of boric acid in the pres-
ence of polyphenols and organic acids. M. G. Mellon
and V. N. Xorris. P. Indiana 4cad. Sci., 35, 85-91.

(0.4. 18, 2484)

140.

141.

142.

143.

144.

145.

146.

110

1924, cont.

The electrometric titration of nickel and cobalt with
potassium cyanide. E. nuller and W. Schluttip. Z. anorg.
allgem. Chem., 134, 327-43. (0.4. 18, 2662)
Electrometric titration - investigation of its methods

and application to certain metallurgical analyses. G. A.

to

Shires. 3. Chem. net. Mining 800. S. Africa. 24. 9-45.
(0.4. 18, 2851)
The electrometric determination of barium alone, and in
the presence of calcium. E. Muller and R. Wertheian z.
anorg. allgem. Chem., 135, 269-72. (0.4. 18, 3017)
Electrometric titration of ferrocyanic acid with potass-
ium iodide. E. Muller. Z. anorg. allgem. Chem., 135,
265-8. (0.4. 18, 3020)
The potentiometric determination of vanadium, chromium,
and iron in the presence of each other and the use of
the method in steel analysis. I. M. Kolthoff and 0.
Tomicek. Rec. trav. chim., 43, 447-56. (0.4. 18, 3571)
The theory and practice of electrometric titration. Fr.
Auerbach and E. Smolczyk. Z. physik. Chem., 110, 65-141.
,(0.4. 19, 222)
The application of titanous chloride to potentiometric
titrations. I. General considerations - the reducing
action of titanous solutions. I. n. Kolthoff. Rec. tran.
chim., 43, 768-74. II. Purity, preparation, and standard-

ization of the titration liquid. I. m. Iolthoff and 0.

111

1924, cont.

 

Tomicek. Ibid. 775-83. III. Estimation of oxidizing
anions. O. Tomicek. ibid. 784-97. IV. Estimation of
oxidizing cations. ibid. 798-807. (0.4. 19, 449)
147. Electrometric titration of antimony and tin by potass-
ium dichromate. n. H. Fleysher. 3.4.0.3., 46, 2725-7.
(0.4. 19, 451)'
148. Electrometric determination of formaldehyde. E. muller
and W. Low. 3. anal. Chem., 64, 297-302.(0.4. 19, 452)
149. The electrometric titration of hydrazine and its salts.
E. 0. Gilbert. 3.4.0.9., 46, 2648-55. ‘ (0.4. 19, 452)
150. Electrometric titrations. E. Muller. Metall 1L Erz, 21,
265-9. (0.4. 19, 620)
151. The electrometric determination of tin from hydrochloric-
oxalic acid solution. E. Eckert. Metall 11 Erz, 21,
202-4. (0.4. 19, 621)
152. The behavior of electrodes of platinum and platinum
alloys in electrometric analysis. R. G. Van Naue and
F. Fenwick. 3.4.0.3., 47, 9-29 (1925) (0.4. 19, 795)
153. Potentiometric determination of bismuth alone and in
the presence of lead. E. Lintl and 4. Rauch. £1.anorg.
allgem. Chem., 139, 397-410. (0.4. 19, 798)
154. Theory of certain electrometric and conductometric tit-

rations. E. D. Eastman. 3.4.0.8., 47, 332-7.

155.

156.

157.

158.

159.

160.

161.

112

1924, cont.

 

Oxidation potentials of ferrous and ferric salts in
concentrated hydrochloric and phosphoric acids. S. B.
Carter and F. B. Class. J. Chem. Soc. 125, 1880-8.

(C.A. 19, 208)
A convenient power-line circuit for the potentiometer.
A. Eeighton. Ind. Eng. Chem. 16, 1189-90.(0.A. 19, 422)
The titration of ferric and cupric salts separately and
in the presence of one another, also in the presence of
antimony, by means of titanous chloride. 1. M. Kolthoff.
Rec. trav. Chim. 43, 816-22. (0.A. 19, 450)
A new dial potentiometer. A. berthelot. Bull. soc. chim.
biol. 6, 683-6. (0.A. 19, 1211)
Quantitative electro-analyses. J. H. Frydelender. Rev.
prod. chim. 27, 757-62. (C.A. 19, 1233)
Electrometric titration of chlorides. E. Muller. Z.
Electrochem. 30, 420-3. (0.2. 19, 1236)
Potentiometric determinations with mercuric salts. E.
Muller and H. Aarflat. Rec. trav. chim. 44, 874-8.

(0.1. 19, 1387)

1925

 

Electrometric titrations with special reference to the
use of titanous chloride for one analysis. A. Mcmillan
and W. C. Ferguson. J. Soc. Chem. Ind. 44, 141-2 T.

(C.A. 19, 1828)

164.

165.

166.

167.

168.

169.

170.

171.

1925,_cont.

The electrometric titration of hypochlorous acid. A.

Schleicher and L. Toussant. Z.ana1. Chem. 65, 399-405.
(0.4. 19, 1829)

The electrometric titration of chlorous acid and its

determination in the presence of hypochlorous acid. A.

’Schleicher and W. Wesly. Z. anal. Chem. 65, 406-11.

(C.A. 19, 1829)
Electrometric titration in the estimation of ferrous
and ferric iron in magnetites. H. 8. Adam. J. S. African
chem. Inst. 8, mo. 1, 7-10. (C.A. 19, 1832)
An improved calomel electrode. 0. J. Schollenberger.
Ind. Eng. Chem. 17, 649. (0.A. 19, 1969)
Electrometric titration. J. C. brumich. ind. Eng. Chem.
17, 631-2. (0.A. 19, 1999)
The use of radio receiver tubes for electrometric tit-
rations. W. D. Treadwell. Helv. Chim. Acta. 8, 89-96.

(C.A. 19, 1999)
New calomel half cells for industrial hydrogen-ion
measurements. H. 0. Parker and C. A. Darmerth. Ind. Eng.
chem. 17, 637-9. (C.A. 19, 2147)
The potentiometric standardization of potassium perman-
ganate solutions with sodium oxalate. 0. del Fresno.
Z. Elektrochem. 31, 199-200. (C.A. 19, 2314)
Electrometric titrations using quinhydrone. H. niklas

and A. Hock. z.angswu. chem. 38, 407-9. (0.A. 19, 2314)

 

114

1925, cont.
172. Differential electro-titration. D. C. Cox. J.A.U.S. 47,
2138-43. . (0.A. 19, 2610)
173. Electrode vessel for liquids heavier and lighter than
the liquid junction potential eliminator. L. E. Dawson.
J.A.C.S. 47, 2172-3. (0.A. 19, 2761)
174. Standardization of solution used in iodometry. II. S.
Popoff and J. L. Whitman. J.A.U.S. 47, 2259-75.
(0.1. 19, 2921)
175. The potentiometer determination of cerium. O. Tomicek.
Rec. trav. Chim. 44, 410-5. (C.A. 19, 2922)
176. continuous measurement of the e. m. f. of titration
cells. Arthur W. Gardner, Thesis for degree of B. 3.,
M. 3. 0., 1925.
177. A Text-Book of Electra-Chemistry. Max Le Blane. Trans.
by Whitney and Brown. The Macmillan 00., N.Y., 1910.

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