to
of tfbe
\Hniv>er0it^ of Toronto
M. JL. Nasmith, Esq.
The D. Van Nostrand Company
intend this book to be sold to the Public
at the advertised price, and supply it to
the Trade on terms which will not allow
of discount.
f f/3.
c
EXPERIMENTAL
ELECTROCHEMISTRY
v \ . Dl
N. MONROE HOPKINS, PH.D.
Assistant Professor of Chemistry in The George Washington University,.
Washington, D. C
TKHttb One 1)un&reO anO Cbirtg Illustrations
NEW YORK
D. VAN NOSTRAND COMPANY
23 MURRAY AND 27 WARREN STREETS
1905
Copyright, 1905,
BY
D. VAN NOSTRAND COMPANY
ROBERT DRUMMOND, PRINTER, NEW YORK
DEDICATED TO
Cbarles J60war& flbunroe
AS A TOKEN OF ESTEEM
PREFACE.
IT has been the aim of the author to produce a book that will
prove useful in both the lecture room and in the laboratory.
Although a work upon the subject of electrochemistry must pre-
suppose a working knowledge of chemistry and electricity, it has been
the object of the writer to introduce the subject as clearly as possible,
dealing with chemistry and electricity without assuming too much on
the part of the student. The book is so written that it may be read
through as a history by the student, presenting theory and practice
together, with the introduction of an ample number of experiments
to supply experimental evidence for the theories advanced. In the
later and more practical part of the book, exercises in preparing
electrolytic compounds and in isolating metals are introduced. It
has also been the aim of the writer to introduce new material with
suggestions for additional experiments with the hope that the work
will not be unwelcome to those already well informed in the subject.
Electrochemistry is at best a subject for the advanced student, and
in order to carry out the experimental work with profit, as given
here, he must have pursued beforehand laboratory courses in both
chemistry and physics. If the more mature and experienced student
profits by a study of the book, or the instructor obtains assistance in
teaching or suggestions for new work, the object in writing the book
will be attained.
N. MONROE HOPKINS.
LABORATORIES OF THE GEORGE WASHINGTON UNIVERSITY,
WASHINGTON, D. C.
SOME IMPORTANT SUGGESTIONS TO STUDENTS AND
RESEARCH WORKERS IN ELECTROCHEMISTRY.
A PERUSAL of the works devoted to electrochemistry, especially
the earlier ones, will show as a rule a sad and almost complete lack
of important data relative to practical manipulations. This absence
of governing facts in the note-books of many students and research
workers also renders it very difficult if not impossible to repeat the
work with any certainty, or to instruct others to do so. The author
wishes to impress upon the student the absolute necessity of observing
and recording numerous electrical factors in all electrochemical
work.
Chemists, until very lately, have shown a pathetic need of elec-
trical knowledge, a failing only rivalled, it may be said, by the lack
of chemical knowledge exhibited by electricians. It should be con-
stantly borne in mind that electrochemistry is primarily the work of
the chemist. The application of the electric current- and its control,
although of vital importance, is subordinate to the purely chemical
side. Electrochemical operations are essentially chemical and based
upon purely chemical changes, and it is only the man with a broad
and keen insight into theoretical chemistry who can ever hope
to make a successful electrochemist or electrochemical engineer.
Nevertheless a thorough working knowledge of electricity is to-day
absolutely necessary for success. Electrochemical science is only
to be mastered by the man with a chemical and physical equipment.
A few words of caution relative to practical work. Do not begin
an experiment or a preparation by roughly mixing up an electrolyte
and subjecting it to the action of an unmeasured electric current
in a haphazard manner. If you enter in your note-books that a
solution of a certain strength was electrolyzed for a certain time by
viii SUGGESTIONS TO STUDENTS AND RESEARCH WORKERS.
a current of so many amperes, flo not feel that you have recorded
all the important data. You have not begun to take into account
the necessary governing conditions. What of the temperature?
What of the electrode area and current density, as well as the mate-
rial of the electrodes ? What was the electrode tension ? The specific
gravity before and after electrolysis? What was the character of
the apparatus and the dimensions of the cell ? Were anode and
cathode temperatures the same or was there a difference ? These
are but a few questions that could be asked after an experimental
run, all of which throw important light upon what was going on.
Let the student have constantly in mind that he is above all things
a chemist and doing a chemist's work, and, secondly, that he is a
physicist or electrician, and that the electric current which he is
using is capable of application in numerous ways. An electric
current is composed of factors, and its deportment as a reagent is
largely dependent upon these factors. A student who mixes up a
solution in a hurry, sticks in electrodes more or less clean, and turns
on the current will never get the most out of a possibility except
by a remote chance; and if he is a careless worker, he is not likely
to recognize the full significance of a success if by chance he should
succeed. Some cases of electrolysis are extremely complicated, and
it will even prove necessary at times to repeat the same experiment
many times in order to obtain and record all the facts. One run
may be made for anode and cathode temperature changes, another
for density changes, and another for analytical purposes, all with a
constant and set current strength, perhaps. Do not feel that one
must use large quantities. The best work is often done by repeated
experiment, using small quantities of material.
TABLE OF CONTENTS,
CHAPTER I
PAGE
HISTORICAL NOTES AND IMPORTANT CLASSIC RESEARCHES i
Early Experiment of Sir Humphry Davy in isolating the metals
Sodium and Potassium. Mechanism of electrolysis of potassium
hydroxide theoretically considered. Experimental electrolysis of a
solution of sodium or potassium hydroxide. The theory of electroly-
tic dissociation. Practical directions for the application and control of
the electric current in experimental work.
CHAPTER II.
THE THEORY OF ELECTROLYTIC DISSOCIATION 17
Electrolytes and non-electrolytes. Simple experiment to show elec-
trolytes and non-electrolytes. Laws of Boyle, Gay-Lussac, and Avoga-
dro relative to the behavior of substances in solution. Osmotic Pressure
and method of measuring it. The Principle of Soret. Application of
data as experimental evidence in favor of the theory of electrolytic
dissociation.
CHAPTER III.
THE THEORY OF ELECTROLYTIC DISSOCIATION (continued) 32
Additional experimental evidence in its support. The lowering of the
freezing-point of solvents. The elevation of the boiling-points of solvents.
The neutralization of acids and bases. The dissociation of ammonium
chloride. Dissociating action of water and other bodies. Practical
experiment to show dissociation by increase in molecular conductivity
and color change. Summary of experimental facts in support of the
theory of electrolytic dissociation.
CHAPTER IV.
NOVEL EXPERIMENTS IN "ELECTROLYTIC INDUCTION." 50
Ostwald's experiment in static induction to show the presence of
"free ions." Ostwald and Nernst's experiment in static induction to
ix
X TABLE OF CONTENTS.
PAGE
show the presence of free ions. Author's experiment in static induction
to show the presence of free ions through the agency of the reflecting
galvanometer as a chemical indicator. Experiments to show effect of
magnetism upon a coil of electrolyte. Experiment to show the effect
of an electric current traversing an electrolyte upon a magnetic needle.
Experiment to show and to measure the effect of an electric current
traversing an electrolyte, upon a mass of iron. Comparisons between
magnetic effects of conductors of first and second class when carrying
current. Experiment to show electrodeless conduction. Experiment
to show and to study the effect of alternating currents upon electrolytes.
The influence of "frequency, " the energy factor remaining constant.
CHAPTER V.
"THE VELOCITY OF ELECTROLYTIC CONDUCTION " 66
Ostwald's experiment to show instantaneous conduction through
electrolytes. Experiments with a high-speed special band chronograph
capable of dividing a second into a million parts. Author's experiment
to show instantaneous conduction through electrolytes, and to compare
the time required for conduction between conductors of the first and second
class. Lodge's experiment for determining the absolute velocities of
the ions. The significance of the data in hand in relation to the theory
of electrolytic dissociation.
CHAPTER VI.
FARADAY'S LAW 80
Electrochemical and chemical equivalents. Experimental demon-
stration of Faraday's Law. Table of a few electrochemical equivalents
arranged as a guide for the student. Validity of Faraday's Law. Law
of Dulong and Petit in relation to the facts brought out by Faraday's
Law. Short table of atomic heats. The construction and use of the
silver voltameter. The copper voltameter. The gas voltameter.
Distinction between current required to electrolyse, and the energy
absorbed. Experiment to show the energy absorbed in different
electrolytes when a common current is passing. Experiments to show
the mechanical transport of ponderable material. The experimental
electrolysis of zinc chloride. The electrolysis of solid glass when sub-
jected to heat. Experiments with frozen electrolytes. Experiment to
show heat convection in electrolysis.
CHAPTER VII.
ENERGY REQUIRED IN ELECTROLYSIS 100
Practical formula for calculating it. Partial table of the heats of
formation of chemical compounds. The experimental liberation of
TABLE OF CONTENTS. XI
PAGE
metallic magnesium from an igneous or fused electrolyte, with current
and energy determinations. The electrolysis of concentrated sulphuric
acid into its elements, with energy, calculation and theoretical mechanism
of electrolysis. The separation of metals through adjustment of
electrode tension. The construction and operation of a rotating anode
equipment for the rapid electro-deposition of metals. Experiment with
copper sulphate solution to show concentration changes. Theory upon
which the concentrations are believed to take place. Experimental
apparatus of Mather and Jones for determining concentration changes
and calculating the relative velocities of the ions.
CHAPTER VIII.
IMPORTANT CONDITIONS TO BE NOTED IN ELECTROCHEMICAL OPERATIONS. 117
A table showing useful tabulation of data. The influence of current
density on oxidation and reduction. Experiment to show the effect of
current density upon oxidation and reduction. Experimental arrange-
ment of Beckmann thermometers for studying electrode temperatures
in electrochemical research. The electrolytic production of caustic
soda and chlorine from salt. The electrolytic production of white lead
from lead electrodes. The electrolytic production of cadmium yellow.
The electrolytic production of mercury vermilion. The electrolytic
production of Scheele's Green. The electrolytic production of Berlin
Blue. Apparatus and equipment.
CHAPTER IX.
THE ELECTROLYTIC PREPARATION OF POTASSIUM CHLORATE FROM POTAS-
SIUM CHLORIDE 133
The electrolytic preparation of sulphur trioxide. Introduction of
special apparatus for the electrolysis of solutions yielding gaseous prod-
ucts and means for condensing and estimating same. The electrical
production of Ozone. Brief history of ozone and its applications.
Ozone generators using sulphuric acid. Ozone generators using
metal-coated tubes. The commercial production of ozone and the
purification of drinking water through its use.
CHAPTER X.
THE PRODUCTION OF NITRIC ACID FROM THE ATMOSPHERE 150
Brief history of nitric acid and the historic chemical and electrical
means of obtaining it from air. Various forms of combustion chambers
for the treatment of the air or gases drawn in. Influence of the size of
TABLE OF CONTENTS.
PAGB
combustion chambers. Influence of amperage and energy in secondary
of transformers, or in the flaming arcs used. Influence of temperature.
The experimental and commercial pieces of apparatus designed and in
use. The construction of a simple rotating electrode equipment for the
production of nitric acid from air in the lecture-room or laboratory.
CHAPTER XI.
THE ISOLATION OF THE METALS SODIUM AND POTASSIUM 165
Brief historic sketch of the early work. Processes of extraction.
Method of Sir Humphry Davy. Apparatus of Charles Watt. Borscher's
sodium cell, and tabulation of workable conditions. Castner cell and
process. The design of a small Castner cell for isolating sodium and
potassium on a small scale. Calculation for minimum voltage necessary
for isolating sodium.
CHAPTER XII.
THE ISOLATION OF THE METAL ALUMINUM 174
Brief history and outline of researches upon the production of
aluminum. Carbon reduction furnaces. Borscher's furnace. De-
ville's apparatus. Hall's furnace. Heroult's furnace. The pro-
duction of aluminum bronze. Method of charging and operating
furnaces for the isolation of the metal and the production of aluminum
bronze.
CHAPTER XIII.
THE ISOLATION OF CALCIUM 186
Outline of the early work in isolating this metal, enumerating the
difficulties encountered. Bunsen's directions and reference to the
importance of current-density conditions. Illustration in section and
detail of an experimental calcium furnace, with instructions for preparing
the metal upon a small scale. Difficulties met and precautions necessary.
Borscher's experimental calcium furnace and method of operating it.
CHAPTER XIV.
THE ELECTRIC FURNACE AND FURNACE PRODUCTS 194
The construction of a wire-resister muffle-furnace for temperatures
under 1600 C. Vertical type upon same principle, with directions for
making. Author's "Series Carbon" furnace for the production of high
temperatures oja the no- volt circuit. Illustration of furnace, diagram
TABLE OF CONTENTS. xlii
PAGB
of connections, and method of using. Directions for the production
of calcium carbide upon an experimental scale. The construction of
an arc or resister furnace for the production of the highest temperatures
when ample current is available.
CHAPTER XV.
PREPARATION OF ORGANIC COMPOUNDS 207
The electrolysis of sodium acetate with the liberation of hydrogen and
ethane. The electrolytic production of iodoform. Detailed directions
for producing small quantities in the laboratory. The synthesis of
acetylene. The production of carbon disulphide and outline of com-
mercial preparation. The production of chloroform. Electrolytic
oxidation. The preparation of kanarin. Apparatus of Hoffman.
Electrolytic reduction.
CHAPTER XVI.
THE PRIMARY CELL 219
Historic outline of early theories to account for the origin of the
current. The production and origin of current in the light of modern
physical chemistry. The phenomenon of solution tension. Experiment
to show "chemical action at a distance." The theory of the cell and the
theory of electrolytic dissociation. Nernst and Helmholtz's "double
layer." The normal electrode and its uses. Neuman's table of potential
differences between metals and their salts. The tension series of the
metals. The chemistry and electrochemistry of the Daniell cell.
Calculation of electromotive force in primary batteries.
CHAPTER XVII.
THE SECONDARY CELL 237
Conditions to be studied in the secondary cell. The construction
and study of a simple experimental cell. Typical discharge curve of
ideal storage-battery. Author's reflecting hydrometer for studying
density changes in batteries. Sellon's hydrometer. The theory and
chemistry of the storage-battery. Difficulties in the way of a clear and
perfect study. Ayrton's theory of the secondary cell. Theory of
Plante". Views of Treadwell. The charging and care of storage-bat-
teries. Rules for the maintenance of cells. Calculation of capacity of
storage-cells. Calculation of electromotive force.
xiv TABLE OF CONTENTS
CHAPTER XVIII.
PAGE
ELECTRICITY FROM CARBON 250
Primary and secondary powers. Zinc as fuel. Thermoelectricity.
Thermoelectric battery and couples. Warning to student in research
work not to confound thermoelectric phenomena. Becquerel's experi-
ment on the oxidation of carbon to produce the electric current direct.
Work of Jablochkoff and Jacques. Ostwald's outline of necessary con-
ditions for success. Solution and ionization of carbon. Experiment
to show solution of carbon. Lack of data in most cases where research
work has been done. Edison's furnace wrong in principle for the
oxidation of carbon to electricity. Outlook for the solution of the
problem.
CHAPTER XIX.
USEFUL PIECES OF APPARATUS 259
Conductivity cells for studying electrolytes with movable and
stationary electrodes. Delicate electrolytic rheostat. Electrolytic test-
tubes. Electrolytic test-tube with removable porous partition. Hoff-
mann's apparatus. Directions for making one upon a large scale. The
Wenhelt interrupter and its application to induction-coils and apparatus
requiring an intermittent current. The aluminum rectifier for convert-
ing alternating currents into direct currents. Principle upon which
the aluminum rectifier is believed to act.
CHAPTER XX.
BIBLIOGRAPHY CHRONOGRAPHICALLY ARRANGED 267
A review of the historic writings upon electrolysis and electrochem
istry, electrometallurgy, etc., to date. The bibliography is not claimed
to be in anywise complete, but it is thought to cover many classic writings
as well as numerous important researches of a more recent and practical
nature. The bibliography as introduced should be of especial value
to students interested in early work.
EXPERIMENTAL ELECTROCHEMISTRY.
CHAPTER I.
HISTORICAL NOTES, AND IMPORTANT CLASSIC RESEARCHES,
WITH SIMPLE DIAGRAMS OF THE USE OF THE ELEC-
TRIC CURRENT TO ELECTROLYSIS.
BELIEVING that the proper introduction to this volume should
give a review of the historical work, together with a notice of the
more basic experimental evidence obtained in support of the theories
and laws advanced, the opening pages are devoted to recording the
more important researches and discoveries.
As this chapter deals jointly with electricity and chemistry, the
best place to commence the recording of events is the time when the
galvanic or voltaic current was introduced or grafted into chemistry.
The history of electrochemistry before the discovery of the gal-
vanic current requires but a brief description. Ages before the dis-
covery of voltaic electricity it had been observed that various metals,
by being simply immersed in metallic solutions, became coated with
the metal previously dissolved in the liquid.
Thousands of years ago Zosimus mentioned the deposition of
bright metallic copper upon iron immersed in a solution of a copper
salt. In the year 1752 Sulzer remarked: "If you join two pieces of
lead and silver, so that they will be in the same plane, and then lay
them upon the tongue, you will notice a certain taste resembling
that of green vitriol, while each piece apart produces no such sensa-
tion." Becaria demonstrated in 1772 that metallic zinc could be
obtained from its oxide by means of a powerful electric spark, as
from a battery of Leyden jars. Paetz and Van Troostvik in 1790
2 EXPERIMENTAL ELECTROCHEMISTRY.
decomposed water by passing electric sparks through it by means
of very fine gold wires.
Up to the close of the eighteenth century, however, a possible
affiliation of electricity with chemistry was not thought of, the second
celebrated experiment of Galvani upon the nerves and limbs of
recently killed frogs, in 1786, marking the dawn of what is now
known as dynamic electricity. ^ As early as 1780 it was observed
by Galvani that the limbs of dead frogs contracted violently when
hung upon a copper hook in the neighborhood of a frictional elec-
trical machine, at each disruptive discharge of the then known and
so-called static electricity. J Six years later Galvani obtained the
same results with the limbs and nerves of frogs without the agency of
an electrical machine, simply by bringing a copper wire joined to a
nerve and one of the limbs in contact with a piece of iron. The
analogy of these results, although six years separated, caused Galvani
to refer the phenomenon to a common agency, namely, electricity.
Galvani describes his discovery of what he called "animal electricity"
in his famous "De Viribus Electricitatis " of 1791 in the following
words: "It is principally found in the nerves and muscles, and its
path seems to be from the muscles to the nerves, or rather from the
nerves to the muscles by the shortest route, as in the Leyden jar.
There is in every part a double electricity, positive and negative, and
disjunctive. One exists internally in the muscles, the other externally;
so that the muscular fiber acts like a little Leyden jar, and the nerves
simply serve the office of conductors." In the year 1792 Alexander
Volta discarded the theory given by Galvani; and from the fact that
convulsions took place more energetically when there were dissimilar
metals in the connecting circuit, instead of only one variety, attributed
the electricity to their being unlike, and laid the basis for the
contact theory of electricity. In 1792 Prof. Fabroni, of Florence,
first suggested chemical action. The following words are from
Prof. Fabroni's report to the Scientific Academy of Florence, con-
cerning experiments which he had made with metals which he had
immersed in water. He said that he was convinced that "a chemical
action had taken place, and that it was unnecessary to seek else-
where the nature of the new stimulus, that it was manifestly owing
to the slow combustion and oxidation of the metal ; which combustion
must have been accompanied by an attraction of oxygen and by a
HISTORICAL NOTES AND CLASSIC RESEARCHES. 3
disengagement of light and caloric." In 1793 Alexander Volta of
Pavia advanced his contact theory of electricity in the Philosophical
and Medical Journal of Leipsic, and later, in his famous memoir
to the French National Institute, he gives an exposition of his
"electromotive apparatus." It is made, he says, writing in 1801
in the above celebrated communication to the National Institute,
"in the form of a pile or of a range of cups, and consists in the
simple metallic pairs of plates, so arranged as to impel the electric
fluid in one particular direction. The zinc is laid upon the silver,
the moist pasteboard over the zinc, and so on consecutively." He
called the different conducting substances the " motors," and their
arrangement a "circle," "in which an electric stream is occasioned,
which ceases only when the circle is broken, and which is renewed
when the circle is again rendered complete." The power of chemical
decomposition of the voltaic "stream" or current was immediately
noticed by numerous workers, Nicholson and Carlisle being the first to
decompose water by means of such a current of electricity on May 2,
1800, and soon afterward Dr. Henry of Manchester decomposed nitric
and sulphuric acid, and also ammonia, by similar means. With the
discovery of the voltaic current scientists became occupied with two
great questions: First, what is the true principle of the voltaic cell
and the source of the electricity ? And second, what is the mechan-
ism of electrolysis, or in other words, how does the electric current
decompose chemical compounds? Let us take up the question of
electro-decomposition first in the present chapter, and discuss the
origin of the electric current when we are in a better position to
appreciate the various factors.
In 1 80 1 Dr. Wollaston discovered that if a piece of silver in
connection with a more positive metal be put into a solution of cop-
per, the silver becomes coated with copper, which coating will stand
the operation of burnishing. During the same year Gerboin first
noticed the movement produced in mercury during the act of electroly-
sis.
In 1803 Hissinger and Berzelius discovered that by means of a
voltaic current the elements of water and of neutral salts were trans-
formed to the respective polar wires immersed in the liquid; and
Cruickshank, about the same time, observed the electro-deposition
of lead, copper, and silver upon one of the polar wires (the one con-
4 EXPERIMENTAL ELECTROCHEMISTRY.
nected with the zinc end of the battery) immersed in solutions of
salts of those metals, and was thus led to suggest the analysis of
minerals by means of the voltaic current.
In 1805 Brugnatelli observed the electro-deposition of gold upon
silver when the former was made the negative pole in a solution of
"ammoniuret of gold"; he also discovered the electro-deposition
of zinc.
The most brilliant and striking proof, however, of the great
breaking-down power of the electric current when applied to chemical
FIG. i. Reproduction of Sir Humphry Davy's Classic Experiment in Isolating the
Metals Sodium and Potassium.
substances was discovered on October 6, 1807, by Sir Humphry
Davy in the electrolytic decomposition of potash and soda, and the
liberation of their respective metals, by a current from a voltaic
battery of 274 cells.
Let us study this classic experiment, and begin our practical
laboratory work by reproducing it, and, under the stimulus of the
famous experiment, undertake to explain the mechanism of electroly-
sis, or in other words, to learn if possible what takes place when an
electric current is made to pass through the substances Davy used.
First let us look into the actual arrangement of the details of the
HISTORICAL NOTES AND CLASSIC RESEARCHES. 5
experiment. For this purpose we will turn to our illustration. In
our electrochemical studies a fair knowledge of chemistry is pre-
supposed, although the author will deal with the subject through-
out as simply and as clearly as possible. A small cavity was made in
a piece of caustic soda, or sodium hydroxide (NaOH), which was
then moistened with water. This was placed upon a piece of sheet
platinum connected with the positive wire of a voltaic battery.
Mercury was poured into the cavity and connected with the negative
wire of the battery, thus closing the circuit through the system.
FlG. 2. S, Block of Moistened Caustic Soda or Potash. M, mercury in cavity of
caustic soda or potash; P, platinum wire dipping into mercury; P', platinum
sheet for positive-wire connections.
Electrolysis began immediately, the metal sodium, from the sodium
hydroxide, being liberated from the hydroxyl and propelled to the
mercury, with which it amalgamated. After about an hour, having
kept the caustic soda moistened by the addition of water from time
to time, the mobile mercury became quite stiff, due to the presence
of the sodium amalgamated with it. This experiment can be most
easily reproduced, and the sodium be freed from the mercury by
distillation of the mercury, leaving the sodium behind, or the amalgam
may be put into water, when the sodium will react with the water
(2Na + 2H 2 O = 2NaOH + H 2 ), setting hydrogen free, which may be
ignited, and forming a solution of sodium hydroxide, which may be
obtained in the solid form by evaporating to dryness on a watch-glass.
In distilling the mercury from the sodium, the reader is referred to
any general work on chemistry, where the proper precautions are
6 EXPERIMENTAL ELECTROCHEMISTRY.
given for this operation. This is a beautiful experiment, and it is
strongly urged that every student in electrochemistry repeat it for
himself. In the place of the historic battery of 274 cells, six or eight
modern cells of battery will suffice, although the direct current from
a lighting system, properly modified by lamps, is to be desired.
The use of lighting circuits and lamps for electrochemical processes
will be fully dealt with later. Let us now look into the theory of
the breaking up of the sodium hydroxide by the electric current.
Davy might have used a strong solution of sodium hydroxide in
water placed in a dish, with a layer of mercury at the bottom to act
as the negative electrode, and to receive the sodium, and the principle
would have been just the same. Let us represent graphically such
a solution, and illustrate by diagram the various steps in the electroly-
sis. As an exposition of all the ancient theories would lead to con-
fusion in our practical work, the latest views only upon this subject
are given, and we will base our work upon the famous theory of
"electrolytic dissociation." This theory explains in a most satis-
factory manner many chemical and electrochemical phenomena,
which without its aid would be hopeless. This celebrated doctrine
was advanced by Svante Arrhenius in 1887, and although there are
many chemists, physicists, and physical chemists who do not accept
it, they have not advanced anything better to account for the numerous
things it explains. There is the most excellent experimental evi-
dence in support of this doctrine, which will be taken up in detail
later on. For the present we will assume it to be true, for besides
being a theory of exceptional beauty, it will be of great assistance
to us in all our work in electrochemistry. The theory simply states
that the molecules of certain chemical substances, when dissolved in
water, break up into ultimate parts, and that these ultimate parts
carry upon them small charges of electricity. Let us look at the
matter from a diagrammatic point of view. The accompanying
illustration (Fig. 3) shows a series of vessels in which we will electro-
lyze a solution of potassium hydroxide. A represents two molecules
of potassium hydroxide about to be plunged into the vessel of water.
Here the familiar molecular chemical formula of the base is given.
B shows what is supposed to take place according to the theory of
electrolytic dissociation. The potassium atom breaks away from the
hydroxyl group, and takes upon itself a charge of positive electricity,
HISTORICAL NOTES AND CLASSIC RESEARCHES.
K OH
K OH
B
f-
K
K
O~H
O"H
JT*
-
-K OH-
-K O~H-*
*_
r^
and the hydroxyl group takes upon itself a charge of negative electric-
ity. All this is believed to happen
simply upon dissolving in the water, with
no electrical influence whatever being
brought to bear. Here we have, accord-
ing to our theory, free potassium de-
tached, and isolated from the hydroxyl
radical, floating around independently
in the water, but covered with a charge
of electricity. At first sight of such a
diagram the majority of chemical readers
would say that they did not believe a
word of it, for in the first place we could
not have free potassium floating about
in water without a violent reaction taking
place between it and the water; and in
the second place, where did the charge
of electricity come from? Let us not
attempt to answer these questions for
the present, but accept the truth of the
theory for the time being, and take up
the next step in the electrolysis. In C
we have introduced into the vessel two
electrodes, one positive and the other
negative, as they are connected to the
positive and negative ends of a voltaic
battery respectively. We know from our
i i i i , , . i inimcianjii in wcitcr; jj, me
elementary physics and electricity that
> same molecules broken down
charges of like signs repel, and those of into "ions" on being dissolved;
unlike signs attract. In this case, if the c the "ions" being attracted
ultimate parts of the molecule of po- to el <* tr des of PP^ V***-
.... . . ity; D, "ions" arrived at the
tassium hydroxide carry positive and electrodes ready to give up their
negative electrical charges respectively, charges; E, the electrical char-
there should be an attraction between &* neutralized, the "ions" be-
. . . come atoms and react to form
the negative Charge Of One electrode potassium hydroxide again, and
and the positive charge of the potassium water, and setting free oxygen
on the one hand, and an attraction and hydrogen gas.
between the positive charge of the other electrode and the negative
L Li I
i< OH
K O~H
i-
f
HOH
FIG. 3. A, two molecules
of potassium hydroxide before
immersion in water; B, the
8 EXPERIMENTAL ELECTROCHEMISTRY.
charge on the hydroxyl group on the other hand. If all this is true,
there will be a movement of the potassium toward the negative
electrode, and a movement in the other direction of the hydroxyl
group toward the positive electrode, as indicated by the small arrows
in the diagram. These ultimate parts of molecules are called "ions,"
whether they consist of a single atom, like our potassium with its
electrical charge, or whether they consist of a group of atoms, like
our hydroxyl, with its electrical charge. Let us then adopt the
technical term, and speak of the potassium hydroxide molecule as
breaking down, in the presence of water, into a positive potassium
jon and a negative hydroxyl ion. The next diagram, D, shows the
potassium ions arrived at the negative electrode and, the hydroxyl
ions arrived at the positive electrode. We may think of the electrical
charges upon these ultimate parts of the molecule as having a pro-
tective action, that is to say, rendering them inert so far as the water
is concerned. We know that we could not put ordinary metallic
potassium into water without a violent reaction taking place, with
the liberation of hydrogen and the formation of potassium hydroxide.
Now let us account for the passive state of the ion potassium in the
water to be due to the protective action of the electrical charge.
What happens when this ion reaches the electrode ? We have plenty
of negative electricity there with which to neutralize the positive
electricity upon the potassium, and neutralization quickly takes
place. The diagram E shows the next step; there the electrical
charges have been neutralized and removed, and instead of ions we
now have ordinary chemical atoms and groups of atoms. In the
lower left-hand corner of this last diagram two molecules of water
have been graphically inserted; for as soon as the potassium ions
become atoms, we know as general chemists that there will be a
reaction to form potassium hydroxide, with the liberation of hydrogen.
The water was not represented in the previous diagrams simply
because it played the part of solvent only, and did not combine
chemically with our ions. The arrows here indicate the setting
free of two atoms of hydrogen at the negative electrode, and the forma-
tion of a molecule of water at the positive electrode, and the setting
free of one atom of oxygen. Now, what are the facts in an actual
experiment? If we electrolyze a solution of potassium or sodium
hydroxide in water, we will have two volumes of hydrogen set free
at the negative electrode, and one volume of oxygen at the positive
HISTORICAL NOTES AND CLASSIC RESEARCHES. 9
electrode. If our negative electrode consists of mercury, as in
Davy's experiment, the sodium or potassium will amalgamate with
the mercury, which prevents it from acting upon the water so long
as the current of electricity continues to pass. In the experiment
with the block of caustic soda or potash there would be only sodium
or potassium set free in the mercury, and oxygen at the moist surface
of contact of the caustic block and the platinum-sheet base.
To electrolyze such a solution experimentally, set up an apparatus
like that shown in the next illustration. Two large test-tubes may be
used, and it will be observed that just twice the volume of hydrogen
n
FIG. 4. Experimental Electrolysis of Sodium or Potassium Hydroxide Solution. O,
oxygen collected in positive tube; H, hydrogen collected in negative tube; CC,
carbons of battery; ZZ, zincs of battery. Arrows indicate the direction of current.
will be set free; in other words, two volumes of hydrogen to one
volume of oxygen will be liberated. Where the wires dip under the
caustic solution it will be necessary to insulate them with a solution
of rubber, or else several coats of gum shellac, to prevent the libera-
tion of gases from the wires themselves. With the ends properly
insulated, the setting free of the oxygen and hydrogen will be confined
to the platinum plates within the tubes. Now, these platinum plates
or electrodes have technical names, and we must become familiar with
them. The positive electrode is called the "anode," and the negative
electrode is called the "cathode." The current in an electrolytic
bath always flows from the anode to the cathode, all electropositive
ions going to the cathode, and all electronegative ions going to the
anode. As a general rule, all the metals and hydrogen go to the cath-
ode in an electrolytic cell, and all other chemical elements go to the
anode. The following table shows the chemical elements arranged
in their electrochemical order, some of the extremely rare ones not
being included. In this table each chemical element is positive to any
10
EXPERIMENTAL ELECTROCHEMISTRY.
element placed above it, and negative to any one given below it.
These distinctions, although of a relative character, are very impor-
tant, since it seems probable that the very nature of chemical attrac-
tion itself rests upon these electrochemical relations.
Anode or Positive Electrode
Negative
Atoms
Oxygen
Sulphur
Nitrogen
Fluorine
Chlorine
Bromine
. Iodine
Selenium
Phosphorus
Arsenic
Chromium
Vanadium
Molybdenum
Tungsten
Boron
Carbon
Antimony
Tellurium
Tantalum
Columbiunl
Titanium
Silicon
Tin
Hydrogen
Gold
Osmium
Indium
Platinum
/ Rhodium
L Ruthenium
\ Palladium
Mercury
Silver
Copper
Uranium
Bismuth
Gallium
Indium
Germanium
Lead
Cadmium
Thallium
Cobalt
Nickel
Iron
Zinc
Manganese
Lanthanum
Didymium
Cerium
Thorium
Zirconium
Aluminum
Scandium
Erbium
Ytterbium
Beryllium
Magnesium
Calcium
Strontium
Barium
Lithium
Sodium
Potassium
Rubidium
Caesium
Cathode or Negative Electrode.
Positive .
Atoms
.
The above column of elements is arranged seriatim as if placed
in an electrolytic cell.
HISTORICAL NOTES AND CLASSIC RESEARCHES. n
This table very forcibly illustrates the preponderance of positive
elements over negative elements, and also the fact that we only
have about seven simple negative ions. By a simple ion, a single
+
charged atom like our K is meant; a complex ion being one like
OH, which is negative. Here we have a negative atom and a positive
atom, composing a negative ion. In this negatively charged hydroxyl
group, or hydroxyl ion, we can think of the hydrogen striving to go
to the cathode, and the oxygen striving to go to the anode, and the
oxygen having the greatest pull and winning, since it is more strongly
electronegative than the hydrogen, is electropositive, as a glance at the
table will show. By means of this table we should be able to pre-
determine the polarity of a complex ion with facility.
For example, let us take the three acids, hydrochloric, sulphuric,
and nitric, and disolves them in water. How do they ionize ? With
the help of the above table and carefully conducted experiment it is
+
an easy matter to determine. The HC1 gives H Cl, the H 2 SO 4 gives
+ + + -
H 2 SO 4 , and the HNO 3 gives H NO 3 . Sulphuric acid has been
+
shown by the present writer to also ionize into the ions H HSC>4.
Here we have a case where hydrogen goes to the positive electrode
or anode, but it is drawn there by being linked to two more power-
fully electro-negative atoms. There are a few cases where metals
go to the anode in electrolysis, but only under such circumstances as
the hydrogen. No metal goes to the anode in an electrolytic cell,
unless it is part of a powerful group of electro-negative atoms. Now
if we accept the theory of electrolytic dissociation, we are led to
believe in a number of things. Perhaps the most important conse-
quence of such a theory is the fact that we have actually moving
masses of matter in a solution when an electric current is made to
flow through it. Such a solution of a chemical substance, capable
of conducting the electric current, is technically known as an elec-
trolyte. In all electrolytes, therefore, the passage of an electric
current through it is accompanied by the movement of ponderable
particles of matter; in other words, the atoms themselves act as
carriers of electricity. There will be an abundance of "experimental
evidence" later to show this, but for the present we must accept
the theory upon faith. As we shall deal a great deal with anodes and
12
EXPERIMENTAL ELECTROCHEMISTRY.
cathodes, and the direction of the electric current, together with
its management and application, the latter part of the present chapter
will be devoted to the more practical side of the question.
SIMPLE DIAGRAMS OF THE USE OF THE ELECTRIC CURRENT FOR
ELECTROLYSIS.
Direction of the Electric Current. As we must always know
the direction of the electric current in all our electrolytic investi-
gations, it does not seem out of place to introduce at this time
a purely electrolytic pole-finder, or current indicator. This con-
sists of a glass tube with the ends bent up as shown in Fig. 5,
and supported horizontally by a couple of laboratory stands.
M
1
s
ee
ES
ee
n
c
^
z
t>
ee
33
ee
iMH
n
s
FIG. 5. A and C, anode and cathode, respectively, in glass tube; M, globule of mer-
cury which travels with the current; B, cell of battery supplying current; C and
Z, copper and zinc electrodes of battery. The arrows indicate the direction of
flow of the electric current, as Well as the movement of the mass of mercury.
Two loose-fitting stoppers carry the platinum-wire electrodes con-
nected to a battery, or modified electric-light current, or small
dynamo. A globule of mercury is placed in the tube as indicated
at M, and the tube filled to near the level of the stoppers with
a dilute solution of sulphuric acid in water. Upon closing the
circuit, the mercury will immediately travel to the negative pole or
cathode. On reversing the direction of the current, the globule of
mercury will be propelled in the reverse direction, serving as a very
pretty illustration of the behavior of positive ions, and answering
all the requirements of a pole-finder or indicator of current direction,
HISTORICAL NOTES AND CLASSIC RESEARCHES.
if the current is sufficiently strong. If the mass of mercury is large
it will require a stronger current to move it, but if quite small it
will be propelled by about T V of an ampere. For very feeble cur-
rents, the direction of flow must be learned by means of a compass-
needle. Perhaps for all ordinary work there is no source of electrical
current so handy and satisfactory as the modified electric-lighting
current when of the direct type, and from no to 220 volts pressure.
The accompanying diagram, Fig. 6, indicates the use of such a cur-
rent in connection with an electrolytic cell and a lamp-bank, which
may be placed in any convenient part of the laboratory or workroom.
FlG. 6. Diagram of Lamp-bank and Electrolytic Cell in Connection with a no- or
a 22o-volt Direct-current Electric-lighting Circuit.
This lamp-bank, which is the special design of the author, has proven
so useful in many electrochemical processes that an enlarged diagram
of it is given in Fig. 7. When connected with the no- volt circuit
a i6-candle-power lamp inserted in any of the single sockets A allows
about \ ampere to pass. With all eight of the single sockets filled,
it allows about 4 amperes to flow; and if these same sockets are
filled with 32-candle-power lamps, a current of about 8 amperes
will be obtained. Now, for a more feeble current, less than J ampere,
two i6-candle-power lamps are placed in the sockets BB, and a
current flow of about T 2 / will be obtained. With three of these
lamps in the sockets CCC, a current of about -^ will be allowed to
pass; 2 20- volt lamps may be used here in series, when the current
will be less than yf-^ ampere. This is only in accordance with the
EXPERIMENTAL ELECTROCHEMISTRY.
well-known law of Ohm: C=R+V, where C is the current, R the
resistance, and V the voltage. A 3 2 -candle-power no- volt lamp has
a resistance of about no ohms, a i6-candle-power lamp about 220
ohms, and a 2 20- volt lamp about 440 ohms. So it will be seen that
with the three kinds of lamps at hand a very flexible lamp-bank
results from the design given.
D
P
p
:D
M
FIG. 7.~-Laboratory Lamp-bank for Electrolytic Work. Range from -j^ ampere to
8 amperes.
However, for certain work the high potential of the electric-
lighting circuit is not desired, and the lamp-bank, no matter how
designed, will not meet requirements. Again, should we require
20 amperes of current for certain work, the lamp-bank would have
to hold twenty 32-candle-power lamps, or forty i6-candle-power
lamps, and would be very wasteful of energy when we consider
that we are working under a difference of potential of at least no
volts; no volts X 20 amperes would represent 2200 watts, which
would be nearly 3 horse-power; 2200 watts divided by 746 (number
of watts to the horse-power) equal 2.94 horse-power. Whereas we
would require the 20 amperes for our electrolysis, we could not only
get along with 4 volts pressure, but would actually prefer it, so we
HISTORICAL NOTES AND CLASSIC RESEARCHES. 5
use a motor-generator, and consume something like 80 watts, instead
of the 2200. 20 amperes X 4 volts = 80 watts. The photograph,
Fig. 8, shows a simple form of motor-generator used by the author
FlG. 8. Photograph of a Motor-generator used to Convert the no-volt Lighting
Current into a i5-ampere Current at 4 Volts Pressure, which are the ideal con-
ditions for many electrochemical processes for experimental purposes.
FlG. 9. Diagram of Motor-generator and Electrolytic Cell in connection with Elec-
trical Measuring-instruments for Observing Electrical Conditions within the
Electrolyte. Here the comparatively high-voltage electric-lighting current is
stepped down to the ideal voltage for electrolysis of chemical compounds.
for the past six or seven years, which gives about 20 amperes at a
pressure of only 4 volts, the driving-motor taking a trifle more than
the corresponding number of watts.
1 6 EXPERIMENTAL ELECTROCHEMISTRY.
The last illustration in our present chapter is a diagram of a similar
motor-generator, but of the belted type, and represents electrical
measuring instruments properly connected for observing the watts
used in the electrolyte for any kind of electrolytic work. Here M
is the driving-motor, G the generator, V and A the voltmeter and
ammeter respectively, and RR rheostats for controlling the speed
of the motor on the one hand and the current supplied to the elec-
trolytic cell on the other. The generator is of the shunt-wound
type, and it is very necessary to have a good variable resistance
in the outside circuit. The ammeter shows the current taken
and the voltmeter the drop of potential across the electrodes. Such
a small rotary converter can be ordered from almost any of the
manufacturers of small dynamos and motors. Having outlined the
simple apparatus necessary for practical work on a small scale, we
will close the present chapter with definitions of the technical terms
introduced and continue the development of the subject in the
next chapter.
Electrolysis. The breaking up of chemical compounds by the
electric current, and the setting free at the electrodes of the con-
stituents.
Electrode. The terminal of the source of electricity which dips
into the electrolyte.
Electrolyte. A chemical compound, capable of conducting the
electric current when in solution or in the fused state.
Anode. The positive electrode in an electrolyte. The electrode
from which the electric current flows.
Cathode. The negative electrode in an electrolyte. The elec-
trode to which the electric current flows.
Ion. A chemical atom or group of atoms possessed of an elec-
trical charge.
Electrolytic conductivity. The passage of the electric current
through an electrolyte accompanied by the movement of ponderable
material. Ion transfer. The carrying of the electric current by
moving ions.
Electrolytic dissociation. The breaking up of certain chemical
molecules when dissolved in water or other suitable solvents, into
ultimate parts charged with electricity.
CHAPTER II.
THE THEORY OF ELECTROLYTIC DISSOCIATION.
IN the last chapter, that beautiful doctrine known generally as
the theory of electrolytic dissociation was touched upon, and the
reader was asked to accept upon faith the truth of its meaning for
the time being. It is the purpose of the present chapter to advance
some of the best experimental evidence in its support, and leave the
student to formulate his own opinions. Probably there is no generali-
sation in the entire domain of physical chemistry quite so unique and
attractive as the theory advanced as recently as 1887 by Svante
Arrhenius, now professor at the University of Stockholm. Few \
theories in either chemical or physical science have been the subject (,
of greater dissertation, dispute, or attack than the dissociation j
theory, and few have served a more useful purpose in accounting'
for certain vital phenomena. The theory of electrolytic dissociation
has the most excellent experimental evidence in its favor, and accounts
perfectly for many heretofore unexplained facts, and those urging
objections to its truth have never been able to propose a better one,
or even one half as good. It will be the effort of the present writer
to advance what he considers to be the best and most forcible evi-
dence for this doctrine, and adopt it throughout in the practical
electrochemical studies which are to follow. Until something better
is brought forward, we will not take our time in making attacks
upon the doctrine. The arguments against the theory will not be
introduced for fear of confusing the student. We know from previous
experience that we have, broadly speaking, two kinds of conductors
of the electric current the metals and alloys on the one hand, and
solutions of certain chemical substances on the other. In the case
of the metals and alloys they are called conductors of the first class,
and in the case of chemical substances in solution or in a state of
fusion they are called conductors of the second class. The passage
i8
EXPERIMENTAL ELECTROCHEMISTRY.
of an electric current through a conductor of the second class is
believed to be accompanied by the actual movement of ponderable
material, or a mechanical transfer of matter. Good evidence in
support of this will be introduced a little later.
We may now take all known chemical compounds and divide them
into two great groups as follows : ist. Those compounds which when
dissolved in water or other suitable solvent conduct the electric
current; ad. Those compounds which when dissolved do not conduct
the electric current. For this purpose we may draw a dividing line
separating these two great classes, and term those which conduct
when in solution, electrolytes, and all those which do not conduct
the electric current when in solution, non-electrolytes. In the follow-
ing table a few chemical compounds of both kinds are given. This
table could, of course, be indefinitely extended, but a sufficient
number of compounds are given to show the character and meaning
of the division. Upon examining the bodies in the left-hand column
it will be observed that all the electrolytes are among, and constitute,
CHEMICAL SUBSTANCES.
Electrolytes.
Non-electrolytes.
Sodium chloride NaCl
Sodium nitrate NaNO 3
Potassium sulphate KzSO 4
Ammonium hydroxide NH 4 OH
Sodium hydroxide NaOH
Potassium hydroxide KOH
Sulphuric acid HjSO,
Nitric acid HNO 3
Hydrochloric acid HC1
Acetic acid CHaCOOH
Oxalic acid CzH 2 O 4
Silver nitrate AgNO 3
Cane-sugar
Ethyl alcohol C^OK
Methyl alcohol CH 3 OH
Benzene C 6 H 6
Chloroform CHCls
Ether (C Z H 5 ) 2 OH
Acetic aldehyde CH 3 CHO
Formic aldehyde HCHO
Acetone CH 3 COCH,
Propyl alcohol C 3 H S OH
Amyl alcohol C 8 H,,OH
Isopropyl alcohol
the "chemically active" bodies, whereas the non-electrolytes con-
stitute the "chemically inactive" bodies. It will be observed that
certain chemical substances are electrolytes only when dissolved
in water or other suitable solvent, or when in the fused condition,
according to the definition of an electrolyte. We may take any of
the chemical compounds in the left-hand column, including even the
acids, and when absolutely water-free they are non-conductors of
the electric current. Water itself, when properly distilled and air-
THE THEORY OF ELECTROLYTIC DISSOCIATION. 19
free, is also a non-electrolyte (except to an infinitesimal extent), and
yet when certain chemical substances are dissolved in water they
become most excellent conductors of electricity. See Fig. 10 for a.
simple experiment for distinguishing electrolytes from non-electrolytes-
The lamp-bank and electric -lighting circuits, or the motor generator
as described in the first chapter, may, of course, be used instead of
the storage-battery as given here. Here we may have the case where
two bodies, when separated, each prove to be non-conductors, and
when brought together, to conduct highly the electric current. What
is it due to? To take a special case, a crystal of rock salt (sodium,
chloride) and carefully distilled water. Neither of these substances
will conduct the electric current to any appreciable extent. Dissolve
B B
FIG. 10. Simple Apparatus for Distinguishing Electrolytes from Non-electrolytes.
A, glass beaker containing distilled water with platinum electrodes into which the
compound to be tested is dropped; C, milli-amperemeter; BB, cells of storage-
battery. There is always a slight indication of conductivity upon a sensi.ive
milli-amperemeter when only the glass beaker and distilled water are present,
due to the dissolving of a minute trace of glass together with impurities in the
distilled water. The amount of deflection can be noted and applied as a cor-
rection.
the salt in the water and the solution has a high conductivity. Some-
thing must have taken place within the water, and yet we know we
have made only a simple solution of salt in water, which when,
evaporated to dryness, gives back our salt unaltered, and if we catch,
and condense the water driven off, we have ordinary distilled water
again. What is the condition of the salt in the water, then, to so greatly
change its physical behavior toward the electric current? In terms
of the theory of electrolytic dissociation, as was pointed out in the
previous chapter, the chemical molecule is broken up into "ions"
or ultimate parts, and these ultimate parts bear electrical charges
upon them. The molecular formula of sodium chloride is represented
simply thus, NaCl, so familiar to all general chemists. Now upon
immersion in water the molecule is believed to be broken up as
20 EXPERIMENTAL ELECTROCHEMISTRY.
+
iollows into two "ions," Na and Cl, the bond or attraction between
the two former atoms being broken, and the sodium ion with its
electrical charge is existing independently of the chlorine ion with
its electrical charge of unlike polarity or sign. The mere act of
passing into solution is believed, in terms of our theory, to separate
.the atoms of certain molecules, the atoms becoming ions at once
by taking upon themselves electrical charges' of opposite signs, the
one becoming a positive ion, the other a negative ion. In the previous
chapter it was pointed out, by means of a diagram, that the positive
ion traveled to the negative electrode, and that the negative ion
traveled toward the positive electrode. It is the purpose of this
chapter to show that we have excellent reasons for believing in the
existence of these electrically charged particles when certain chemical
substances are dissolved in water. We will take up the study of
: the evidence first by comparing the deportment of substances in
aqueous solution with substances in the state of a gas, and for this
purpose we will first set down the three gas laws so well known to ail
students of modern chemistry and physics.
Law of Boyle. The pressure exerted by a gas, the temperature
remaining the same, is proportional to the concentration of the gas.
The concentration of the gas is directly proportional to the number
of molecules or ultimate parts of molecules present.
Law of Gay-Lussac. The pressure of a gas increases a constant
amount for every increase of i in temperature, and the increase in
pressure is equal to ^g of the original pressure of the gas at o C.
Law of Avogadro. Equal volumes of all substances in the
gaseous state, under the same conditions of temperature and pressure,
contain the same number of molecules or ultimate parts of molecules.
Consequently the molecules of all substances, or the ultimate
parts of all molecules when in the gaseous state, under the same
conditions of temperature and pressure, occupy the same space.
Having the three fundamental gas laws before us, we will take
them up separately in the order given, and learn what bearing they
have upon the behavior of substances in solution. What possible
application can these gas laws have to chemical compounds dissolved
in water ? There appears to the general student to be no connection
whatever, and yet there is the most vital application of the gas laws
in support of the theory of cbctrolytic dissociation. Let us firzt
THE THEORY OF ELECTROLYTIC DISSOCIATION. 2r
take up Boyle's law, which has to do with the pressure exerted by
substances in the state of a gas. This tells us facts based upon
experimentally determining the pressures exerted by gases of differ-
ent concentrations. The pressures of gases confined in given volumes
at constant temperature can be readily measured by manometers
or pressure-gauges, as set forth in detail in any good text-book on
physics. This we know; but can we measure the pressures exerted
by substances in solution ? We can convert a given mass of a given
substance into a gas by heating, and measure the pressure at differ-
ent concentrations, or volumes, by means of suitable manometers-
If we dissolve the same quantity of the compound in water, will the
molecules exert a pressure in the dissolved condition, and can we
measure it ? Both these questions can be answered in the affirmative,
and it is the purpose here to show that such pressures exist, and to>
describe how they may be measured. All substances when dissolved
in water exert a pressure, and this pressure has been termed "osmotic
pressure."
OSMOTIC PRESSURE AND METHOD OF MEASURING IT.
If a gas, oxygen or hydrogen for example, be liberated in a given
space, the gas will expand in all directions and completely fill the
containing vessel. If all parts of this containing vessel are at the
same temperature, the gas will expand and distribute itself uni-
formly throughout the volume. There will be repellent forces
between the molecules of the gas, driving them to the remotest recesses
of the containing vessel, and consequently there will be a pressure
against the walls of the same. The more concentrated the gas, or,,
in other words, the greater the number of molecules or ultimate
parts of molecules present, the greater will be the pressure within the
fixed or given volume. What can be said about substances in solu-
tion? The behavior is the same. Let us take a large vessel of
water, for example a tall glass jar full, and introduce a little sugar in
it. The sugar will immediately fall to the bottom, a small portion
dissolving and passing into solution on the way down. What will
b2 the ultimate result on standing? The sugar at the bottom will
all pass into solution, rise against gravity, and in time distribute
i:self uniformly throughout the solvent. The sugar in the dissolved
state will behave exactly as it would when in the state of a gas, and
22 EXPERIMENTAL ELECTROCHEMISTRY.
will exert a pressure when in solution which may be measured. This
is due to the phenomenon of diffusion, which, not so many years ago,
was wholly unaccounted for. Here we have a heavy substance dis-
solving at the bottom of a tall glass cylinder filled with water, and
rising to the top against the attraction of gravity. There is a pres-
sure, and this pressure has only recently been accounted for. Another
phenomenon which until recently could not be explained was the
bursting of an animal bladder, filled with a mixture of alcohol and
water, when immersed in a vessel containing pure water. The
bladder under these conditions, if it has been closed up properly at
the openings, will be burst by a gradually developed pressure within.
It is easy to show in this way that we have a pressure, and this pres-
sure has been termed osmotic pressure. This is only a very crude
method of showing qualitatively that we have a positive pressure, and
It does not seem to have occurred to the earliest workers that this pres-
sure was a definite thing and could be quantitatively measured.
This osmotic pressure is a very peculiar thing when one considers
the manner in which the pressure is measured. We cannot place a
solution within a closed vessel and get an indication of pressure
upon a gauge-glass or manometer, as we very well know, but must
resort to some kind of a membrane, such as forms the animal bladder.
Strange to say, the pressure developed depends upon a differential,
or selective action, so to speak, of the necessary membrane. It must
allow the solvent to pass through, but not the dissolved substance, a
sort of filter, roughly speaking, and because of this principle the
membrane has been called "semipermeable." Now if we can
really produce a semipermeable membrane or diaphragm, we will
be able to measure the pressure due to substances in solution. Take,
for example, a solution of cane-sugar. If we have at hand a membrane
which will be permeable to water and impermeable to sugar, we
can by its use ascertain the pressure due to the presence of the sugar
molecules, and demonstrate how this pressure varies with concen-
tration of the solution and with changes in temperature. We have
at hand, in other words, means for comparing the behavior of gas
molecules with the behavior of molecules in solution. Let us pre-
pare such a semipermeable membrane in the laboratory and examine
some of the compounds given in the preceding table constituting
electrolytes on the one hand and non-electrolytes on the other.
THE THEORY OF ELECTROLYTIC DISSOCIATION. 23
The accompanying photograph illustrates some simple forms of
porous pots, and Fig. 12 gives a section through such a typical pot,
as well as a completed piece of apparatus for experimentally measuring
osmotic pressure. Moritz Traube, and Pfeffer, the celebrated plant
physiologist, were the first to discover and make use of the properties
of semipermeable membranes. It is to Pfeffer that we owe the
first really serviceable artificial semipermeable diaphragm or partition.
FIG. ii. Some Forms of Porous Pots with Semipermeable Membranes for the
Measurement of "Osmotic Pressure." The broken exhibit shows the semi-
permeable membrane at M.
It was discovered by Pfeffer that plant and animal membranes could
be discarded, and that certain chemical precipitates, when properly
supported, met the requirements almost perfectly. Copper ferro-
cyanide was found to give the most satisfactory results when formed
right in the walls of a very fine-grained unglazed porous pot. In
order to produce such a precipitate within the walls of the porous
pot, it was filled with a solution of potassium ferrocyanide and im-
mersed in a solution of copper sulphate. The two solutions met
within the walls and there formed the semipermeable membrane
with the resistant support of the porous pot. When such a prepared
24 EXPERIMENTAL ELECTROCHEMISTRY.
pot is broken open, the membrane appears in the form of a fine line,
as indicated in Fig. 1 1 at M. There are many necessary precautions
to be taken in the preparation of successful semipermeable mem-
branes, it being an art requiring not a little patience and skill.
The following is taken from one of Pfeffer's writings on the subject :
"The porcelain cells were first completely injected with water under
the air-pump and then placed for at least some hours in a solution
containing at least 3 per cent of copper sulphate, and the interior
was also filled with this solution. The interior only of the porcelain
cell was then rinsed out quickly with water, well dried as rapidly as
possible by introducing strips of filter-paper, and after the outside
had dried off, it was allowed to stand some time in the air until it
just felt moist. Then a 3 per cent solution of potassium ferrocyanide
was poured into the cell and this immediately reintroduced into the
solution of copper sulphate. After the cell had stood undisturbed
for from twenty-four to forty-eight hours, it was completely filled
with the solution of potassium ferrocyanide and closed. ... A
certain excess of pressure of the contents of the cell now gradually
manifested itself, since the solution of potassium ferrocyanide had a
greater osmotic pressure than the solution of copper sulphate. After
another twenty-four to forty-eight hours the apparatus was again
opened and generally a solution introduced which contained 3 per
cent of potassium ferrocyanide and i| per cent of potassium nirate
(by weight), and which showed an excess of osmotic pressure of
somewhat more than three atmospheres."
In all this work as reproduced by the present writer is was found
most essential to obtain a special close-grained grade of porous cup
or pot. A common porous pot, or one the least faulty, such as
containing minute invisible fissures, will defeat the object of the entire
experiment. With faulty pots the writer has frequently had a com-
pleted piece of apparatus assembled, indicating a height of only 2 or
3 feet of the contained solution, when the semipermeable membrane
gave way and oozed through the side of the porous pot. It has been
found that a dilute solution of cane-sugar in water would rise to a
height of 66 feet. Referring once more to Fig. 12, we will note a rise
of about 2 feet, a one-half normal sugar solution being used in this
case. Strange as it may seem, the sugar solution is placed within
the porous pot A and the pot is in turn immersed in the beaker B,
THE THEORY OF ELECTROLYTIC DISSOCIATION. 25
containing distilled water. The pressure is developed within the
porous pot by a very curious action, forcing the liquid up into the
FIG. 12. A, porous pot; B, glass beaker; C, tight-fitting stopper; D, height to
which the contained solution has risen; E, graduated scale; F, transverse sec-
tion of porous pot; S, semi permeable membrane within the wall of the pot;
G, enlarged vertical section through porous pot with semipermeable membrane
showing at S.
manometer-tube. The conditions are as follows: The semiper-
meable membrane allows water to pass through freely, but does
not allow the sugar molecules to pass. Within we have sugar
26 EXPERIMENTAL ELECTROCHEMISTRY.
molecules and water molecules attempting to get out, and out-
side we have all water molecules attempting to get in. Now,
we have, per unit area of the porous pot and semipermeable
membrane, a more effective bombardment from the pure-water
molecules without than from the mixed molecules within. We
may think about the thing also as follows: Every water molecule
striking the diaphragm from the outside gets in, but many of the
water molecules before striking the diaphragm from the inside
collide with sugar molecules, which cannot get through, and thereby
their effectiveness is lost. As a result of such a differential action
we may have a slow ingress of water molecules tending to dilute the
sugar by driving its molecules farther apart and thereby establishing
a pressure. Let us leave the theory of the apparatus now and look
at the facts in some actual and carefully conducted experiments.
OSMOTIC PRESSURE OF NON-ELECTROLYTES AND ELECTROLYTES.
For the sake of simplicity we will record the result of an osmotic-
pressure determination upon a non-electrolyte. For this purpose we
will choose the first non-electrolyte appearing at the top of the column
in the little table already given. This is ordinary cane-sugar, a
solution of which in water does not conduct the electric current.
The following table shows the result of one of Pfeffer's carefully
conducted determinations upon this substance:
CANE-SUGAR, C^H^On.
Concentration in Osmotic Pressure in
Per Cent by Weight. Millimeters of Mercury.
1 535
2 IOl6
4 2082
6 3075
Let us now examine the figures standing for osmotic pressures and
interpret their meaning.
Concentration, Pressure, P.
C. P. C.
1 per cent 535 535
2 " " 1016 508
4 " " 2082 521
6 " " 375 5U
In the above table the pressure in each case has been divided by the
concentration with practically a constant resulting. What little
THE THEORY OF ELECTROLYTIC DISSOCIATION.
discrepancy exists is due to experimental error. Here we have an
analogy with the law of Boyle as applied to gases. We know that the
pressure of a gas increases with its concentration in a direct proportion,
and we see from the above tabulated data that the osmotic pressure
of a solution increases directly with its concentration. In experi-
mental work of this character there are naturally sources of error
which must be expected. For example, when we start with a i per
cent sugar solution and begin to measure its osmotic pressure by
such a piece of apparatus as described, the solution is weakened by
the inflow of water, and unless the manometer-tube is very small, the
volume of sugar solution rising to make the indication will constitute
a high percentage of the entire volume in the porous cup. Pfeffer
also showed, at the instigation of Van't Hoff, that the osmotic pres-
sure of solutions increases slowly with rise in temperature, and that
this pressure is analogous to the increasing pressure of a gas as set
forth in the law of Gay-Lussac. Here a solution of sugar was taken
again, but instead of varying its percentage strength the temperature
of the solution was gradually increased. For this purpose a one-
tenth normal sugar solution was employed (a normal solution of cane-
sugar is made by dissolving a gram-molecular weight of this com-
pound in a liter of water. A one-tenth normal solution is made by
dissolving the gram-molecular weight in 10 liters of water), gradually
increasing its temperature. The accompanying table shows the
data of an actual experiment.
Osmotic Pressure in Cm. of
Mercury.
Temperature.
Calculated
Experimental.
from Gas Law
of Gay-Lussac.
6.8
S-S
5 -5
13-5
5 2 -i
5i-7
14.2
53-i
51.8
22.0
54-8
53-2
While there are slight discrepancies due to experimental error,
the striking application of Gay-Lussac's law to substances in solu-
tion is to be noted. Although Pfeffer was the first to successfully
measure osmotic pressures, it remained for Van't Hoff to bring out
the striking agreements with the gas laws already set forth. Having
28 EXPERIMENTAL ELECTROCHEMISTRY.
observed the strong tendency of solutions to behave like substances
in the state of a gas, by experimenting with semipermeable mem-
branes this great Dutch scientist investigated other possibilities for
showing analogies, among them being what is known in physical
chemistry as the "Principle of Soret."
The Principle of Soret.
If a vertical glass tube is filled with a solution of a chemical
compound, such as copper sulphate in water, for example, and the
two ends of the tube are kept at different temperatures, the copper
sulphate will eventually become more dilute where the temperature
is highest and more dense where the temperature is lowest. This
distribution of the dissolved molecules by diffusion due to differences
in temperature is known as the principle of Soret. The apparatus
shown in Fig. 13 is the design of the present writer for bringing about
such concentration changes. The tubes filled with various solutions
were allowed to stand for a long time with the top and bottom at
different temperatures, when some of the solution was allowed to
run out from the bottom and analyzed for density, and some of the
solution drawn out from the top by means of a pipette and also ana-
lyzed for density. The early experiments showed smaller differences
in concentration than would be called for if Gay-Lussac's law applied
to the temperature coefficient of the osmotic pressure of solutions.
The tubes were then allowed to stand for longer periods of time,
with the result that the figures obtained approached closer and closer
to the value expected from the law pertaining to gases. Diffusion
of molecules takes place very slowly and the tubes had to stand for
many weeks undisturbed before an equilibrium was finally established.
In one experiment where the tubes stood for about twelve weeks a
copper-sulphate solution gave the following results:
Upper end of tube 80 C. ; lower end of tube 20 C. The differ-
ence in density between the respective ends of the tube upon analysis
was found to be 14.03 per cent. The difference in density calculated
from the law of Gay-Lussac is 14.3 per cent. Another experiment
gave 24.87 per cent change in concentration when according to
Gay-Lussac's law the figure should have been 24.8 per cent.
Here the application of Gay-Lussac's law to the behavior of
compounds in solution is very striking. With the apparatus as
THE THEORY OF ELECTROLYTIC DISSOCIATION. 29
FIG. 13. The "Principle of Soret" (Author's Apparatus). A, vertical glass tube con-
taining the substance; B, water in wooden tub; C, outer wood casing to be filled
in between with charcoal or sawdust; D, thermometer; E, heavy layer of hair-
felt between thick boards supporting top of apparatus; F, water in copper heater;
C, ring gas-burner for heating water; H, bulb of air-thermostat for holding
temperature of water constant; 7, mercury of air-thermostat for cutting off gas-
supply, /, if the temperature rises too high; K, long vertical glass condensing-
tube to prevent loss of water, F, by distillation; L, thermometer for observing
temperature.
3 EXPERIMENTAL ELECTROCHEMISTRY.
illustrated, the large mass of water in the wooden tub, together with
a laboratory kept at practically constant temperature, the lower end
of the tube is consequently very uniformly maintained in tempera-
ture. The upper end of the tube is kept at an elevated temperature
by means of the ring gas-burner and the sensitive air-bulb thermostat.
The expansion of the air contained in the air-bulb H forces the
mercury at I up to the tube / which is slotted, and gradually cuts
off the supply of gas which enters as indicated by the small arrow.
Should the temperature of the water F fall below the required tempera-
ture, the air in the bulb H will contract and allow the mercury to fall
away from the tube J, thus uncovering the slot and allowing more
gas to flow to the burner. Two very sensitive thermometers give
the readings for the top and bottom of the tube respectively. The
apparatus is so designed that the top portion may be lifted off, when
the tube containing the solution experimented upon may be readily
removed. It remains now to compare the behavior of substances in
solution with the third and last gas law, namely, that of Avogadro.
This was also done by the chemist Van't Hoff. He worked again
with a solution of cane-sugar, and compared the osmotic pressure of
such a solution with a volume of hydrogen gas of equal concentration.
For this purpose he made a cane-sugar solution having the same
number of sugar molecules in a given volume of solution as there are
hydrogen molecules in the same volume of the gas. The experiment
fully justified the statement that the sugar solution gave an osmotic
pressure equal to the gas pressure. We may then say that equal
volumes of all chemical compounds in solution giving the same osmotic
pressure at the same temperature contain the same number o] molecules
or ultimate parts o] molecules. Now this is only true for di ute solu-
tions. Very concentrated solutions of chemical substances do not
obey the law, and when we look about we are struck by the fact that
very densely compressed gases do not obey the law of Boyle. This
makes our comparisons all the more striking, for where we have
exceptions in the case of gases, we also have exceptions in the case of
solutions. It has now been shown that the three fundamental gas
laws apply to compounds in a state of solution, but what has this,
although striking and of vital interest to the physical chemist, to do
with the theory of electrolytic dissociation? To answer this let us
turn once more to the first table of this chapter where we have
THE THEORY OF ELECTROLYTIC DISSOCIATION. 3 1
electrolytes on the one hand and non-electrolytes on the other. It was
pointed out that all those bodies have been classified according to their
ability to conduct the electric current. All those on the left conduct
when in solution, and all those on the right do not. All those on the
left are called electrolytes, and in terms of the theory of electrolytic
dissociation, their molec les break up into ions, each ion of course
being an ultimate part of a molecule. Now as a matter of fact, only
the non-electrolytes, when dissolved in water, obey the gas laws.
It is only the non-electrolytes which give an osmotic pressure com-
parable with substances in the state of a gas, the electrolytes all giving
an abnormally high osmotic pressure. This is just what we would
expect if one molecule breaks up into two ultimate parts and each
ultimate part occupies the same space as the original molecule.
Our sugar molecule does not conduct the electric current when in
solution, it does not break up into ions, and gives as evidence a normal
osmotic pressure. Our sodium chloride, or common salt, does con-
duct the electric current when dissolved, and it gives an abnormally
high osmotic pressure. Of course in comparing the osmotic pressure
of sugar with sodium chloride, two solutions are made in which the
same number of molecules are dissolved in each case. In order to
accomplish this the gram-molecular weight of each compound is
taken. By gram-molecular weight of a compound we mean the
molecular weight of the substance expressed in grams. For
example, the gram-molecular weight of sodium chloride is 58.5
grams, 58.5 being the molecular weight of sodium chloride. So
much for the theory of electrolytic dissociation and the gas laws and
the evidence that the measurement of osmotic pressure gives us in
favor of ionization. We will take up additional evidence in support
of the theory of electrolytic dissociation in our next chapter.
CHAPTER III.
THE THEORY OF ELECTROLYTIC DISSOCIATION (Continued).
IT is well known that pure water freezes constantly at o C., and
that this fact has been made the basis for the several thermometric
scales for scientific purposes throughout the world. It has also
been well known from very early times that the addition of salts or
other soluble material to water causes it to freeze at a lower tempera-
ture. Every schoolboy knows that common sea-water will not
freeze except at very low temperatures, but few of us who have not
paid attention to physical chemistry have given the fact more than
a passing thought. We know that substances in solution cause
pure water to freeze at a lower temperature than pure water
alone; in other words, that the freezing-point is lowered by the
presence of dissolved substances. This is purely qualitative knowl-
edge, so to speak, and there remains for us to investigate this matter
quantitatively, to see how much solutions of the same concentration
lower the freezing-point, and if all compounds lower it equally.
Raoult, the celebrated French chemist, took up this matter for experi-
mental investigation, and, to make a long story short, found that all
non-electrolytes of equal concentration lowered the freezing-point of
pure water to the same extent. Raoult worked with solutions con-
taining one gram-molecule of the dissolved substance per liter and
found that the lowering of the freezing-point was the same, being
1.85 C. One gram-molecule of a substance per liter is a normal
solution, and we may say therefore that all normal solutions of non-
electrolytes lower the freezing-point of water 1.86 C. This is com-
parable to saying that the lowering of the freezing-point of pure
water is dependent upon the number of molecules or ultimate parts
of molecules present. This is, of course, an interesting fact, but
what has it to do with the theory of electrolytic dissociation ? This
question can be very quickly answered by determining the lowering
32
THE THEORY OF ELECTROLYTIC DISSOCIATION. 33
of the freezing-point by normal solutions of electrolytes. What
would we expect if the theory of electrolytic dissociation be true?
Will a gram-molecule of an electrolyte dissolved in a liter of
water give us the same depression of the freezing-point, namely,
1.86 C. ? This was done by Raoult, and it was found that in every
case of an electrolyte the depression of the freezing-point was greater
than i.86C. It will be remembered that all electrolytes exerted
a greater osmotic pressure than non-electrolytes, and now we see that
all electrolytes lower the freezing-point to a greater extent than non-
electrolytes. We can only account for these striking phenomena by
attributing the abnormal behavior of electrolytes to the breaking up
of the molecules, upon dissolving, into ions. The practical student,
upon reading the work done by Raoult and noting his constant of
1.86 C., will want to know how much greater the depression of the
freezing-point was found to be in the case of electrolytes, and what
kind of a thermometer was employed when dealing with such small
differences in temperature. The average electrolyte, when dissolved
in water, depresses the freezing-point about twice as much as any
non-electrolyte. As for the thermometer, it is far from the ordinary
pattern, and is used in a special piece of apparatus. The best and
most universally used apparatus is that of Beckmann, and is illustrated
in one of its forms in Fig. 14. The thermometer in this particular
case is simply one of great sensitiveness and refinement, reading
direct to hundredths of a degree. Because of an exceptionally large
bulb, the degree divisions are very long, allowing of very fine sub-
division. With such a thermometer one-tenth of i C. is a large
amount. The accompanying illustration should make the scheme of
the apparatus clear, and it will be seen that it is a simple one to get
up in the laboratory for actual work, the thermometer being the only
costly element. For exceedingly accurate research work thermom-
eters may be had reading to thousandths of a degree. There are
also metallic thermometers with which temperatures are measured by
the change in electrical resistance of a little coil of platinum wire,
and the delicacy is almost without limit. For all practical purposes,
however, a mercury thermometer reading to hundredths meets every
requirement. The practical carrying out of an experiment with
such a Beckmann apparatus as shown in Fig. 14 is as follows:
An accurately weighed quantity of pure distilled water is introduced
34
EXPERIMENTAL ELECTROCHEMISTRY.
in the tube A, which in turn is placed in the tube B and packed
around with a mixture of ice and salt. The large tube B provides
an air-space around the tube A, and causes a more uniform freezing
B
FIG. 14.
FIG. 14. Form of Beckmann's Apparatus for the Study of Electrolytes and Non-elec-
trolytes by Depression of Freezing-points. A, large glass test-tube with side neck;
B, larger glass tube with cork to receive test-tube; C, large glass jar to receive
both tubes and freezing-mixture; D, stirrer; EE, wire stirrers within test-tube;
G, side neck into which the substance to be tested is placed; F, delicate "open-
scale" thermometer.
FlG. 15. Apparatus for Experimentally Determining the Elevation of Boiling-points
of Electrolytes and Non-electrolytes. A, flask with double side necks; B, asbes-
tos ring supporting flask on tripod; C, little cylinder of platinum within flask
to prevent cooled condensed water from striking the thermometer-bulb; D, Bun-
sen burner; E, condenser with water-jacket; F, Beckmann thermometer with
mercury-reservoir at top; G, enlarged view of mercury reservoir.
of the water in the inner tube. The air between the two tubes becomes
chilled below the freezing-point of pure water and freezes the water
in the inner tube. The stirrer E is moved up and down in the dis-
THE THEORY OF ELECTROLYTIC DISSOCIATION. 35
tilled water, and the thermometer is carefully watched. The mer-
cury will fall steadily until the sudden formation of flakes of ice
throughout the water occur, when it will quickly rise a little and
remain stationary, and this reading should at once be taken. With
a correct thermometer, the indication should of course be o C.
If the reading is not exactly o it matters not, so long as we are
merely measuring the differences between the freezing-point of pure
water and water containing compounds in solution. At least three
readings should be made with the same water, allowing the ice to
melt and then freezing over again, and taking the average of the
three temperatures for the freezing-point of the pure solvent. The
sudden rise of the thermometer is due to a small supercooling of the
water (in spite of the fact of its being stirred) , below its freezing-point,
and then its warming up again at the instant of the formation of ice
It is well known to those who have studied physics that water throws
of! heat when it freezes, the phenomenon being attributed to latent
heat. Having determined carefully the experimental freezing-point of
the water, a carefully weighed quantity of the substances to be tested is
introduced through the side tube D and allowed to dissolve. The freez-
ing process is then repeated three times, as with pure water, and the
average of the three readings is taken. If the water and compound
have been so weighed as to give a normal solution, and the com-
pound is a non-electrolyte, we will obtain the figure 1.86, working,,
of course, with a centigrade thermometer. The important point to
observe in making all these freezing-point determinations is to read
the thermometer at once after the sudden rise of the thermome-
ter, at the time of the formation of the ice. If we wait, and keep on
with the freezing process, the thermometer will fall again, due to the
fact that the solution has become concentrated by the freezing out
of some of the water. This, of course, concentrates the solution
and gives it a new and lower freezing-point. Electrolytes are treated
in the same manner as non-electrolytes. There is another method
which we should not pass over without notice, and that is the testing;
of electrolytes and non-electrolytes by the elevation of the boiling-
point. It is well known that pure water boils at a constant tempera-
ture under a constant atmospheric pressure, and that the heights
of mountains have been measured by the decrease in boiling-point
of water with a delicate thermometer. It is also very well known
36 EXPERIMENTAL ELECTROCHEMISTRY.
that the presence of dissolved substances increases the boiling-point
of pure water. Raoult also investigated this phenomenon experi-
mentally, and found that normal solutions of non-electrolytes increased
the temperature of the boiling-point to the same extent. He also
showed that all electrolytes of comparable concentration elevated
the boiling-point to a much greater extent. Fig. 1 5 illustrates a piece
of apparatus for experimentally determining the elevation of boiling-
points with great accuracy. This special type of thermometer has an
arbitrary scale, that is, it is not designed to indicate absolute tempera-
tures, but only differences between temperatures. The little reser-
voir at the top contains a supply of mercury, which may be shaken
down to join on to the column within the bore, thus allowing the instru-
ment to be used with liquids of lower boiling-points. There are only
about eight degree divisions upon the entire scale of such a delicate
instrument, and were it not for the flexible character due to the
mercury-reservoir, the use of such a thermometer would be exceed-
ingly limited. With a set of two such instruments, one designed for
low temperatures and the other for high temperatures, in view of the
little reservoirs, we are equipped for experimental work throughout
a very wide range. With either thermometer we may take from the
reservoir, or return to the reservoir, by shaking the instrument,
thereby making it serviceable for use at almost any temperature.
In conducting experiments with this apparatus, a few fragments of
broken glass are introduced in the flask to prevent "bumping"
when the solution boils. We see, therefore, from these two experi-
mental investigations of Raoult, that we have the most excellent
evidence in favor of the theory of electrolytic dissociation.
ADDITIONAL EVIDENCE. THE NEUTRALIZATION OF ACIDS AND BASES.
One of the commonest and most familiar chemical reactions is
the neutralization of an acid by a base with the formation of a salt
and water. The following is a simple example, where hydrochloric
.acid and sodium hydroxide are brought together in solution:
HCl+NaOH =
Here we have sodium chloride (common salt) and water formed
In the reaction. So much for the general chemistry of the reaction.
We also have a physical side to the reaction, and this concerns the
THE THEORY OF ELECTROLYTIC DISSOCIATION.
37
heat produced when the reaction takes place. The general chemist
has to do with the products formed, and the physical chemist has
to do ...with the energy transformations and their measurement.
Now in the above reaction heat is liberated, and it remains for us.
FIG. 1 6. Calorimeter for Measuring the Heat Liberated when Solutions of Electro-
lytes are Mixed together and Allowed to React. A , brass calorimeter-casing con-
taining mass of water, B. There is also an inner calorimeter-casing of polished
metal, C, and the reaction-chamber, D; E, hair-felt covering; F, stirrer; G,
glass reservoir with stopcock; H and /, two similar thermometers of sensitive
type reading to hundredths of a degree; /, thermometer indicating temperature
of water-jacket; K, wooden wedges to insulate calorimeter.
to determine how much, and see if it has anything to do with our
theory of electrolytic dissociation. Let us carry on such a chem-
ical reaction and experimentally measure the amount of heat given
out. For this purpose we shall require a calorimeter like that
represented in Fig. 16. It is easily made of polished brass by any
38 EXPERIMENTAL ELECTROCHEMISTRY.
good sheet-metal worker, and is a valuable piece of apparatus for
the physical-chemical laboratory. The inner reaction-chamber D
should be of thin platinum, however. For our experiment we will
place a normal solution of sodium hydroxide within the platinum
chamber, and a normal solution of hydrocholoric acid within the
glass reservoir with the stopcock turned off. The two thermom-
eters are inserted, and the entire apparatus is allowed to stand for
a sufficiently long time to allow equilibrium to be established. The
thermometers are then read, and the hydrochloric acid from the
reservoir is allowed to run into the calorimeter while the stirrer F
is operated. The thermometer / is carefully watched until the
mercury rises to the highest point. Now we are to measure the
lieat of the reaction in calories, and in order to do this it is only
necessary to know the mass of the liquid raised through the indi-
cated temperature, and to make the usual calorimetric corrections.
We must know and allow for the specific heat of the mixture, the
weight and specific heat of the platinum vessel, the heat exchange
of the calorimeter, etc. The detailed method of working with
calorimeters can be found in any good laboratory treatise on gen-
eral physics. As a result of such an experiment with hydrochloric
acid and sodium hydroxide we get 13,700 calories, in addition to
the formation of the salt and the water. Now to come to the point;
It matters not what acid and what base we use, or what salt is formed,
we always get experimentally in such a calorimeter determination
33,700 calories. The following table indicates the run of things,
and it now remains for us to interpret the meaning and see what
it has to do with electrolytic dissociation. The first table gives a
varying acid and a constant base, and the second table a varying
base and a constant acid.
HCl + NaOH liberates 13,700 calories.
HBr + NaOH liberates 13,700 calories.
HI + NaOH liberates 13,700 calories.
HNO 3 + NaOH liberates 13,700 calories.
HCl+LiOH liberates 13,700 calories.
HC1 + KOH liberates 13,700 calories.
HCl + Ba(OH) 2 liberates 13,800 calories.
HCl + Ca(OH) 2 liberates 13,900 calories.
THE THEORY OF ELECTROLYTIC DISSOCIATION. 39
In the case of the calcium and barium hydroxides, one-half
normal solutions were taken to normal solutions of hydrochloric
acid, for the reason that calcium and barium are bivalent. The
above tables exhibit to us a remarkable performance, and it re-
mains for the physical chemist to explain it. The theory of electro-
lytic dissociation explains it perfectly, and in so doing gains im-
portant experimental evidence in its own support.
HC1 and NaOH react and give NaCl and H 2 O.
In terms of our theory, however, these bodies would be disso-
ciated and represented thus:
+ - + - + -
H Cl and Na OH react and give Na Cl and H 2 O.
Salt and water are the products, of course, but as the salt is
+ -
born in water, so to speak, it is dissociated as represented as Na Cl,
and not as NaCl, as it would be out of solution. The only
thing really formed in the molecular state is water, and the constant
of 13,700 calories is merely the heat of formation of water. AH
bodies have either a positive or a negative heat of formation, and
if the salt is formed in the molecular condition along with the water,
it would, of course, add its own heat of formation to the sum total,
and as different salts have different heats of formation, we would,
of course, not get a constant, but a different number of calories for
each reaction between an acid and a base. The following table
gives the heats of formation of a number of salts produced by the
acids and bases which we have tabulated:
NaCl 9*760 calories.
NaBr 8,580 calories.
Nal 6,910 calories.
NaNOs 115130 calories.
LiCl '9*380 calories.
KC1 10,430 calories.
BaCl 2 x 9>47o calories.
CaCl 2 16,980 calories.
In the case of the calcium and barium chlorides, one-half the
indicated number of calories must be taken, for the reason that cal-
40 EXPERIMENTAL ELECTROCHEMISTRY.
cium and barium are bivalent and require double the quantity of
acid.
It will thus be seen that without the theory of electrolytic dis-
sociation we would be unable to explain the liberation of a con-
stant number of calories, when an acid reacts with a base.
ADDITIONAL EVIDENCE.
Let us take four electrolytes, for example, and make two mix-
tures. For this purpose we will choose:
First mixture: Potassium nitrate,
Sodium iodide, Nal.
Second mixture: Potassium iodide, KI;
Sodium nitrate, NaNOs.
Dilute solutions of both salts in each mixture are represented
as follows, with all the constituents dissociated:
K NO 3 and Na I;
K I and Na NO 3 .
In terms of the theory of electrolytic dissociation we have ex-
actly the same ions present in both cases, and the properties of the
two mixtures should be absolutely the same. The two solutions,
when equivalent quantities of the different substances are taken,
are found to be identical in every respect.
FURTHER EVIDENCE. EXPERIMENTS WITH PERFECTLY DRY
COMPOUNDS.
Let us take any of the "chemically active" bodies, or electrolytes,
and inquire into their behavior when perfectly dry. In terms of
the theory of electrolytic dissociation, dry electrolytes are in the
molecular condition and are also "chemically inactive." In terms
of the theory, the dissociation products, or ions alone, are capable
of entering into chemical combinations. Let us test this matter
carefully by referring to a number of experiments involving the
careful drying of the substances employed. It is well known to
THE THEORY OF ELECTROLYTIC DISSOCIATION. 41
chemistry that ammonia gas, NH 3 , and hydrochloric-acid gas, HC1,
react at once to form ammonium chloride, NH 4 C1:
The white clouds of ammonium chloride are even manifested
when an ammonia bottle is unstoppered in the neighborhood of
hydrochloric acid. It has been shown by the most careful and
patient workers that thoroughly dry ammonia gas and thoroughly
dry hydrochloric-acid gas do not react to form ammonium chloride,
and may be separated after mixing in a thoroughly dry receiver.
The accompanying illustrations show how this may be accom-
plished after the gases have been produced and thoroughly dried.
The absolute drying of these gases is a difficult and tedious process,
FlG. 17. Diagram Representing an Uncombined Mixture of Dry Ammonia Gas and
Dry Hydrochloric Acid Gas. The respective gases are here being separated by
charged electrodes, where they may be drawn off and tested.
for the slightest trace of moisture in either the gases or the glass
globe will defeat the object of the experiment. They may, of course,
be dried by passing through towers of finely broken lime and phos-
phorous pentoxide. The globe must be heated to a high tempera-
ture by means of a Bunsen flame, while thoroughly dried air is
passed through. In every detail the most elaborate precautions
must be taken against having moisture present. In this case we
have two molecules, NH 3 and HC1, behaving like ions, that is, the
one goes to the positive pole and the other to the negative pole.
The student may ask how it is that we have hydrogen going
to the positive pole, as in the HC1 diagrammatically represented
within the globe in Fig. 17. If he will turn to the first chapter and
examine the elements arranged in their "electrochemical order,"
42 EXPERIMENTAL ELECTROCHEMISTRY.
he will note that chlorine is much more strongly electronegative
than hydrogen is electropositive, and being linked to the hydrogen,
draws it to the positive pole. In the case of ammonia, we may
think of the hydrogen winning and dragging the nitrogen to the
negative pole, because there are three hydrogen atoms to the one
FIG. 18. Glass Globe with Electrodes Leading to Static Machine for Separating a
Perfectly Dry Mixture of Ammonia and Hydrochloric Acid Gases. After mix-
ing together in the globe the gases may be separated by static charges upon the
electrodes and be drawn off through the glass tubes.
of nitrogen in the ammonia molecule. Atom for atom nitrogen is
more strongly electronegative than hydrogen is electropositive, as
can be readily seen from the table, but there are three hydrogen
atoms pulling the one nitrogen atom, and we may compare matters
to a game of football where three players for one goal get hold of
a single player for the other goal. The single player is pulling
harder than any one of the others to make his goal because he is
stronger, but he is overpowered in number. Fig. 18 shows a glass
glebe on an electrical machine for carrying out such an experiment.
THE THEORY OF ELECTROLYTIC DISSOCIATION. 43
So much for this experiment. The following list represents work
done by various experimenters in support of the dissociation theory:
Perfectly dry sulphuric acid has been shown not to act on per-
fectly dry metallic sodium!
Dry hydrocholric acid does not act on carbonates.
Dry hydrogen and chlorine may be mixed together and exposed
to the sunlight without an explosion taking place.
Dry hydrochloric-acid gas does not precipitate silver nitrate
from water-free ether or benzene solution.
Dry acids will not act upon litmus paper, and will not form
salts with dry bases.
Absolutely dry oxygen gas will not support combustion in many
moisture-free substances !
Dry chlorine does not combine with metals, not excepting sodium
and potassium.
Absolutely dry gunpowder could not be ignited!!
Allow the slightest trace of water vapor to enter the field in any
of the above cases and we have immediate reactions. What part
does the water play? In terms of our theory it is the dissociant,
or cause for breaking down the molecules. Fig. 19 represents a
molecule consisting of th atoms A and B, with the "chemical
affinity" between them assigned to electrical attraction of unlike
charges. The atoms in the molecule here are believed to be held
together by electrical attraction. Now bring such a molecule into
the presence of water. The negative atom will induce a positive
charge in the water, and the positive atom will induce a negative
charge in the water. Now, according to J. J. Thomson, one of
England's most distinguished and famous physicists, because of
these induced charges the attraction between the atoms A and B
will be weakened, and when immersed in the water will be lost
altogether. The following are Prof. Thomson's words describing
the condition when such a molecule as represented in AB in the
little diagram Fig. 20 is brought near a conducting sphere: "Thus
let AB represent two atoms in a molecule, placed near a conducting
sphere, then the effect of the electricity induced on the sphere by
A will be represented by an opposite charge A', the image of A in
the sphere. If A is very near the surface of the sphere, then the
negative charge at A' will be very nearly equal to that of A Thus
44
EXPERIMENTAL ELECTROCHEMISTRY.
the effect of the sphere will be practically to neutralize the effects of
A ; as one of these effects is to hold the atom B in combination, the
affinity between the atoms A and B will be almost annulled by
the presence of the sphere. Molecules condensed on the surface
of the sphere will thus be practically dissociated. The same effect
would be produced if the molecules were surrounded by a substance
possessing a very large specific inductive capacity. Since water
~ fc^
FIG. 20.
FIG. 19.
FlG. 19. Diagram Illustrating the Part Played by a Dissociant when an Electrolyte
is Immersed. A and B are the atoms of a molecule.
FlG. 20. Diagram Illustrating Prof. J. J. Thomson's Theory of Electrolytic Disso-
ciation, assuming that the atoms in a molecule are held together by electrical
attraction.
is such a substance, it follows, if we accept the view that the forces
between the atoms are electrical in their origin, that when the mole-
cules of a substance are in aqueous solution the forces between
them are very much less than they are when the molecule is free
and in a gaseous state."
Thus far we have considered only solutions of . electrolytes in
water. Water has therefore been the dissociant in all the cases
which we have so far met with. Although water is the strongest
dissociant known, there are other liquids capable of breaking down
molecules when solutions are made in them. By strongest dissociant
known we mean a solvent which breaks the largest number of mole-
cules down into ions per unit of solvent volume. In very concen-
THE THEORY OF ELECTROLYTIC DISSOCIATION. 45
trated solutions of electrolytes we have a mixture of molecules and
ions. As the dilution is increased the number of ions increases,
for upon the addition of more water more of the molecules are broken
down. The strong acids, bases, and salts are completely dissociated
when a molugram equivalent is dissolved in icoo liters of water.
If we dissolve a molugram of a strong acid in 500 liters of water,
we will have in solution molecules and ions. The solution con-
ducts the electric current solely by the transport of the electricity
by the free ions. If we measure the conductivity of such a solu-
tion, we will, of course, obtain a certain conducting value. Now,
what will be the effect of diluting the solution with water? With
further dilution we get increased ionization up to the point where
there are no molecules left, all having broken down into ions We
should expect the molecular conductivity of the solution to increase
upon diluting with water, if the dissociation theory is true. As a
matter of fact, the molecular conductivity does increase up to the
point where we have a gram-molecular equivalent dissolved in 1000
liters of water. Now, water being the strongest dissociant known,
all other solvents must be present in larger quantity to effect an
equal dissociation. We will now give a table with the dissocia'nts
in order of their strength, and follow it by an easily performed prac-
tical experiment to show that dissociation increases upon dilution.
DISSOCIANTS IN ORDER OF POWER.
Water.
Formic acid.
Methyl alcohol.
Ethyl alcohol.
There are other dissociants, but the above are among the most
common and generally employed. J. J. Thomson has also shown
that the dissociating power bears a relation to the dielectric constants.
This is in support of the theory of the electrical attraction between
the atoms in a molecule. Having stated that dissociation and
electrical conductivity increase upon dilution up to a point where
we have the gram-molecule dissolved in 1000 liters of water,
we will now test it by experiment.
46 EXPERIMENTAL ELECTROCHEMISTRY.
PRACTICAL EXPERIMENT TO SHOW DISSOCIATION AND INCREASE IN
ELECTRICAL CONDUCTIVITY UPON DILUTION WITH WATER.
In the following experiment there is developed a double and
simultaneous indication of ionization, the appearance of a deep-
red color on the one hand and the steady increase of electrical
conductivity, upon the addition of water, on the other hand. The
color change is dependent upon the well-known behavior of phenol-
phthalein as a chemical indicator. To the characteristic color
deportment of this interesting compound the conductivity method
is simultaneously applied. The experiment as heretofore exhibite 1
consists in merely noting the color change which is produced as
follows: A small quantity of phenolphthalein is dissolved in ethyl
alcohol and is poured into a tall glass lecture-jar to a height of about
5 centimeters. A few drops of ammonia water are then carefully
added. There will be a slight momentary yellow coloration, which
will immediately disappear upon shaking if too much ammonia
water has not been added; if too much ammonia water has
been added, add more alcohol. Now, chemists know that a
colorless solution of phenolphthalein turns a beautiful red in the
presence of a base. Here we have the phenolphthalein and the
base, ammonium hydroxide, in alcoholic solution together, and no
red color appears. Why? Ammonium hydroxide cannot show
+
its basic properties until dissociated into the ions NH 4 , OH, the
isolated OH or hydroxyl producing such basic manifestations.
Now, if we look at the table of dissociants, we see that ethyl alcohcl
is a very poor dissociant and is unable to break the ammonium-
hydroxide molecules down into ions. Now, what will happen if
we add some water? Water, as will be seen from the table, heads
the list as the strongest dissociant known, and we should expect it
to ionize the ammonium-hydroxide molecules if it be added. If
the ammonium-hydroxide molecules are dissociated or ionized,
we should expect the red color of the phenolphthalein to appear
and become deeper and deeper as the molecules are broken up
into active ions. This is just what happens. Upon the addition
of water the color begins to appear and continues to get deeper
and more decided as dilution continues. This is an odd sight, to
see the addition of pure water to a faintly colored solution produce
THE THEORY OF ELECTROLYTIC DISSOCIATION.
47
a deeper and deeper color as dilution goes on. So much for the
color indication of dissociation on dilution. Now, molecules do
not conduct the electric current, and it occurred to the author to
perform this same experiment over again, but, instead of using the
48 EXPERIMENTAL ELECTROCHEMISTRY
glass jar, to employ a glass tank provided with electrodes and study
the conductivity behavior at the instant the color appears and follow
the conductivity behavior as the phenolphthalein deepens in color.
For this purpose a piece of apparatus was made as illustrated in
Fig. 21. With such a piece of apparatus we should not only ob-
tain the color reaction with an indicator, but an increasing con-
ductivity of the solution. The experiment is best and most forcibly
shown by first filling the tank with pure distilled water to the top,
having washed it out many times previously with distilled water to get
it perfectly clean, when there will be practically no indication upon the
galvanometer. The water is next poured out and the tank carefully
drained and dried as much as possible. It is then filled to the same
level with a solution of phenolphthalein in ethyl alcohol to which
some ammonium hydroxide solution has been added. This should
be colorless, as will be the case if not too much ammonia was added.
There will be practically no indication upon the galvanometer. We
have then separately tested the conductivity of the water and the
solution. Let us now see what the addition of water accomplishes.
For this phase of the experiment the phenolphthalein solution is
poured out, all but a small quantity. The writer usually leaves
solution in the bottom to a depth of about 5 centimeters. Water
is now very slowly added, when the red color begins to appear, and at
the same instant the galvanometer begins to show conductivity. As
the red color increases the electrical conductivity also increases, as
is plainly shown by the galvanometer. The dilution is continued
until the tank is full. The tank is constructed with a distance
between the glass sides of only i centimeter, and therefore requires
but a small volume of solution. The joint between the glass and
the wood is made* in a deep groove by cement. We will now close
the present chapter with definitions of the new terms introduced.
Gram-molecule, or molugram. Molecular weight of a compound
expressed in grams. The molecular weight of sodium chloride is
58.5. In order to use a gram-molecule of sodium chloride we
would weigh out 58.5 grams of the substance, for example.
Latent heat. The amount of heat required to change the phys-
ical state of a body without changing its temperature. The heat
given out or absorbed when certain bodies change their physical
states.
THE THEORY OF ELECTROLYTIC DISSOCIATION. 49
Specific heat. The amount of heat required to raise a gram
of a substance one degree in temperature as compared with the
amount of heat required to raise one gram of water one degree.
Calorie. The unit of heat. The amount of heat required
to raise one gram of water one degree in temperature. There
are large Calories also, being icoo times the small c;.lorie, for con-
venience.
Heat exchange of calorimeter. Error due to loss of heat by the
calorimeter itself, by radiation, etc. To be determined by experi-
ment with individual calorimeters by blank tests.
Elevation of boiling-point. Often expressed in works on phys-
ical chemistry as "the lowering of the vapor-tension of the' solvent."
"Bumping." The liberation of steam with almost explosive
violence from the smooth interiors of glass flasks when liquids are
boiled in them. Prevented by introducing sharp points, as by
the introduction of broken glass.
Dissociant. A solvent that not only dissolves electrolytes, but
breaks them down into ions at the same time. Benzene dissolves
many electrolytes, but does not dissociate them, and therefore such
a solution would be a non-conductor of the electric current.
Chemical affinity. The attraction between the atoms in a mole-
cule, whether due to electrical attraction or other forces. i
Specific inductive capacity. Dielectric constant. We owe to
Cavendish (1771-81) the discovery of the fact that the amount
of inductive effect which takes place through a dielectric is different
for different substances.
Molecular conductivity. Molecular conductivity of an electrolyte
is equal to the specific conductivity of i cubic centimeter of the
solution times the number of cubic centimeters containing a gram-
molecular weight. M = NS, where M is the molecular conductivity,
N the number of cubic centimeters of the solvent containing the
gram-molecular weight or molugram of the electrolyte, and S the
specific conductivity of a cubic centimeter of the solution.
CHAPTER IV.
NOVEL EXPERIMENTS IN "ELECTROLYTIC INDUCTION."
As it is the purpose of the first few chapters of this work to acquaint
the student with the constitution and behavior of electrolytes under
various conditions, such effort would fail should we neglect to touch
upon the electrostatic and electromagnetic deportment of substances
in solution. Having dealt with electrolytes theoretically and experi-
mentally, and learned the fundamental laws upon which their
behavior depends, we will be in a position to take up the practical
work which is to follow in the later chapters, and from the subject
of electrochemistry as a science touch upon electrochemical engineer-
ing as an art. We are, therefore, acquainting ourselves with elec-
trolytes, the theories upon which they are based, and their capacities
as electrical conductors. We shall also study electrolytes as producers
of the electric current but this phase of substances in solution is
best left until a little later. The first experiment illustrating the
effects of electrical induction upon an electrolyte as given in Fig. 22
was designed by Wilhelm Ostwald, Professor of Chemistry in the
University of Leipzig, and one of the most distinguished physical
chemists Germany has ever produced. Prof. Ostwald's experiment
has for its object to prove the existence of "free ions" in an elec-
trolyte, and to show that they actually migrate and carry the elec-
trical charges upon them. The author became much interested in
Ostwald's work, and repeated the experiments for himself, continuing
the research still farther, as will be described in the present chapter,
developing what may be termed "electrolytic induction." Let us
first take up the experiment of Ostwald referring to the illustration.
In the experiment with potassium chloride, Ostwald writes as
follows: "The following consideration may serve to remove the last
doubts as to the validity of the assumption of free electrically charged
NOVEL EXPERIMENTS IN "ELECTROLYTIC INDUCTION." 51
atoms of chlorine and potassium. Imagine two insulated vessels,
B and C, filled with a solution of potassium chloride and electrically
connected by means of the siphon D. Let a negatively charged
body be brought near B, remove the siphon, and lastly the charged
body A . Then, as is well known, B remains positively electrified, and
C negatively electrified. Now, according to Faraday's law, the
electricity in electrolytes can only move simultaneously with the ions.
Consequently, if an excess of positive electricity is present in B,
there must also be an excess of free potassium ions, i.e., of potassium.
B
J
FIG. 22. Prof. Ostwald's Experiment in Static Induction to Show the Presence of
" Free Ions." A, negatively charged body; B and C, beakers filled with a solu-
tion of potassium chloride; D, siphon-tube filled with the same solution and
joining the two beakers.
atoms, by the electricity of which the charge is determined. If the
electricity is conducted away, * the potassium assumes the ordinary
form, and acting on the water of the solution develops hydrogen,
which can be collected in suitable apparatus and tested. Similar
considerations hold good for the chlorine in the vessel C. It is
consequently not only conceivable that the ions in an electrolytic
solution move about with electrical charges, otherwise quite free,
but solutions may be prepared which contain an excess of any ion.
we choose, e.g., an excess of potassium. The assumption that elec-
trolytes contain free ions is not only possible but necessary."
This experiment as originally proposed by Ostwald was not at all
practical, for the quantity of hydrogen gas liberated was so small that
it could not be seen. The liberation of hydrogen is based upon the
following simple equation:
By inserting in the beaker B a platinum wire to earth. N. M. H.
5 2
EXPERIMENTAL ELECTROCHEMISTRY.
The experiment was eventually modified by Profs. Ostwald and
Nernst, the latter being also one of the most brilliant German physical
chemists of the times. This experiment shows to the eye the libera-
tion of hydrogen under similar conditions of static induction, and
is a practical illustration of great beauty. The arrangement of the
apparatus for this experiment is shown in Fig. 23. At A we have
the positive knob of a static electrical machine connected by a tinsel
FIG. 23. Ostwald and Nernst's Experiment in Static Induction to Show the Pres-
ence of " Free Ions." A, positive knob of electrical machine; B, glass flask cov-
ered with tinfoil; C, wet strings connecting the glass flask and the vessel D, both
containing dilute sulphuric acid ; E, burette drawn out into a fine capillary, G,
through the side of which the platinum wire, F, is fused ; H, glass plate on glass
insulators.
-cord or small metal chain to the little hook on the tinfoil covering
of the glass flask B. This flask is filled with dilute sulphuric acid
'.and is thoroughly insulated upon a glass or hard rubber-plate resting
.upon small insulators also of glass. Cords or strings wet with the
t;same dilute sulphuric acid dip into the flask and connect with the
vessel D also containing some of the same sulphuric-acid solution,
and being insulated in a similar manner. The glass burettte E has
been drawn out into a long and fine capillary G through which a
"fme^platinum wire is fused and which /turns" to" earth:. Now what
happens when the electrical machine is put into operation? The
NOVEL EXPERIMENTS IN "ELECTROLYTIC INDUCTION." 53
tinfoil coating being electrically connected with the electrical
machine becomes positively charged, which, acting through the glass
of the flask, attracts and holds a corresponding amount of negative
electricity, while the positive is repelled. The positive electricity,
or, as we believe, the positive ions, which in this case are hydrogen
+ -
(H2SC>4 ionizes into H2 SO4) is repelled through the moist cord which
leads to the vessel D and the capillary of the burette filled with the
acid and water to a height of a few centimeters, when it meets with
a little column of mercury at G connected to earth. This mercury
was drawn up into the capillary by placing it in the bottom of the
vessel D, when some of the dilute sulphuric-acid solution was allowed
to follow. Now the hydrogen ions are repelled: through this system
and are discharged when they reach the grounded mercury. They
then become ordinary atoms of hydrogen, and may readily be seen
in the capillary. On starting the electrical machine the experimenters
observed a rush of tiny bubbles of gas through the mercury at G,
collecting at the top under the glass stop-cock, the SC>4 ion being
held by the positive attraction on the outside of the flask B. Here
we have a very beautiful experiment based upon an induction phe-
nomenon. The experimenters also conducted a most elaborate
quantitative research upon this phenomenon, to ascertain if the
amount of hydrogen set free at G corresponded to that calculated
from Faraday's law, and found within the limits of experimental
error that it did. We shall take up Faraday's law and the
subject of electrochemical equivalents in a later chapter, but at the
present time it is only wise to state that all ions have definite capac-
ities for the electrical charges according to their valencies. Knowing
the electrochemical equivalent of hydrogen, for example, it would
be an easy matter to calculate what mass, or what volume of hydrogen,
would be set free by a given quantity of electricity. The experi-
menters referred to employed such a course in checking the above
experiment quantitatively. So much for the experiment of Ostwald
and Nernst, depending upon the liberation of hydrogen as proof of
the migration of free ions. Fig. 24 illustrates the author's modifi-
cation of this experiment based upon the use of the reflecting
galvanometer as a chemical indicator, to prove that ions had migrated
under the influence of static induction. Gaugin, and later Prof.
54
EXPERIMENTAL ELECTROCHEMISTRY.
Kuester, employed the reflecting galvanometer as a chemical indica-
tor, and the use of such an instrument for detecting acids and bases
be described before we describe the author's experiment. To
quote from Gaugin's work we have the following: "Electromotive
NOVEL EXPERIMENTS IN "ELECTROLYTIC INDUCTION." 55
Force of Platinum in Contact with Acidulated or Alkalinized Water :
Two plates of platinum were immersed in acidulated water for
some time. One being withdrawn, washed in distilled water, and
returned, was found to be negative. Electromotive force = 0.0136
volt. Water alkalinized with KOH was then substituted. The
washed and returned plate was found to be positive. Identical
results were obtained with plates of platinized platinum. It is
possible to recognize by this means whether a liquid is neutral, or
acid or alkaline, even when its reaction is so feeble as not to affect
test-papers."
Now the author's experiment consists in operating the electrical
machine, when the tinfoil coating of the beaker B will be positively
charged and will hold the negative ions of the potassium chloride
+ -
K Cl, which, as may be readily seen, are chlorine ions, and will repel
the positive ions which are potassium through the moist cord into the
beaker C, where they may be discharged, after the removal of the wet
cord, bv the platinum wire shown at the right of the beaker Upon
discharging the potassium ions they become potassium atoms and
react with the water as before 2K + 2H 2 O =2KOH + H 2 , forming
potassium hydroxide and setting hydrogen free. This experiment does
not attempt to show migration by the setting free of the hydrogen,
but by the formation of the alkali, or base, KOH, potassium hydroxide.
To do this the reflecting galvanometer is employed. It would be
expected that the chlorine ions could be discharged in the same
manner and their presence shown by a drop or two of silver-nitrate
solution. Although there is little doubt of their being discharged in
the same manner, the minute quantity of chlorine present would not
suffice to give a chemical precipitation of silver chloride. Perhaps
if the electrical machine was allowed to run for several days, a slight
opalescence might be observed when a drop or two of silver nitrate is
added. When we complete our studies of Faraday's law involving
the electrochemical equivalents we will be in a position to appreciate
how few chlorine ions would migrate under such circumstances as
we have in this experiment. All ions carry very great electrical
charges, and we know as physicists that there is very little quantity
of electricity to be had from a static machine. The electricity from
a frictional machine is almost all potential difference! The amperage
in a current from a static machine is so small as to be detected and
EXPERIMENTAL ELECTROCHEMISTRY.
measured only by very special means. Now a few ions are capable
of carrying many amperes, as we shall see later, and it is not sur-
prising under the circumstances that our static charges have been
carried by very few ions indeed. We will now leave experiments
with static induction and study the effects of magnetic and galvanic
induction upon electrolytes. All the following experiments are
based upon the original researches of the present author, and are
now published for the first time. It occurred to the writer to com-
pare electrolytes with metallic conductors when under the influence
of magnets and electric currents in neighboring conductors, to see if
inductive effects and inductive currents were produced. Will a
magnet induce a current of electricity in an electrolyte as it does in a
metallic conductor ? This question is not touched upon in the treatises
in physics or chemistry, and it was therefore resolved to answer the
question by experiment. Fig. 25 shows the first comparatively
FIG. 25. Experiment to Learn the Effect of a Magnet upon a Coil of Electrolyte.
The central figure represents a coil of wire of equal resistance and dimensions,
which may be substituted for the coil of electrolyte.
rough plan for learning whether a magnet will induce an electric
current in a coil of electrolyte as it does in a coil of wire. We have
here a sensitive reflecting galvanometer at the right to show any
induced current. As a matter of fact a magnet does induce a current
of electricity in the electrolyte and causes the galvanometer to indi-
cate the same. The coil of wire represented in the center was made
of equal dimensions with and substituted for the coil of electrolyte
to ascertain if the effect was quantitatively the same. The coil of
NOVEL EXPERIMENTS IN "ELECTROLYTIC INDUCTION." 57
electrolyte consisted of a glass tube filled with a dilute solution of
sulphuric acid. It was necessary to introduce in series with the
metal coil some additional resistance, which was of a non-inductive
type, in order to obtain comparable conditions, as the coil of elec-
trolyte had a much higher ohmic resistance than the coil of wire.
The deflection of the galvanometer proved to be the same in both
cases. As it was an impossible matter to place the magnet in the
two respective solenoids in exactly the same manner and at exactly
the same time, the experiment as illustrated in Fig. 26 was conducted.
Here we have at the left a soft-iron bar running horizontally through
a coil of insulated wire which is in series with a storage-battery, the
terminals of the wire being free for connection with a contact key
which may be closed uniformly any number of times. Next to the
coil we have a glass coil filled with any good electrolyte in solution,
into which the terminal wires (which must be of platinum) of a
reflecting galvanometer dip. At the extreme right we have a coil of
resistance wire of equal proportions and equal number of turns as
in the glass coil, and in series with it a rheostat of the non-inductive
type, for bringing the wire to the same resistance as the coil of elec-
trolyte. Of course some wire of high resistance must be used, such as
is employed in resistance sets, in order that we will not have to depend
upon much outside resistance, as by the use of the rheostat. The
coil of alloy wire may now be substituted for the coil of electrolyte,
and by means of the key and storage-battery, operating the electro-
magnet, we can produce the same number of magnetic lines of force
in just the same way and in the same time as we did in the case of
the coil of electrolyte. Experiments with such a piece of apparatus
gave the same deflections of the galvanometer with a coil of electrolyte
as they did with a coil of alloy wire. We can then think of the free
ions being migrated by ordinary magnetic induction so common to all
students of physics and electrical engineering. Let us now study the
effect of an electric current upon a magnetic needle while traversing
an electrolyte. For this purpose set up a piece of apparatus like
that represented in Fig. 27. Here we have a glass tube about a meter
in length by about a centimeter in internal diameter, bent up at the
ends as indicated. This is filled with dilute sulphuric acid, and is
provided with platinum electrodes with loose-fitting stoppers. The
tube is supported on two laboratory stands above a delicate compass-
EXPERIMENTAL ELECTROCHEMISTRY.
needle with a graduated arc or scale. There are also two upright
standards provided with insulators between which an alloy wire is
<U C
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g 3
OS rT
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^1
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stretched taut at the same height as the glass tube, so one may be
substituted for the other above the magnetic needle. The magnetic
NOVEL EXPERIMENTS IN "ELECTROLYTIC INDUCTION." 59
needle is represented upon an adjustable stand in order that it may
always be brought to exactly the same distance below the wire and
electrolyte respectively. The measurement must be made from the
center of the wire and electrolytic tube respectively. Below in the
same illustration we have a plan of the apparatus, looking down upon
it. It will be readily seen with such an arrangement how the wire
FIG. 27. Experiment to Show the Effect of Electrolytic Conduction upon a Magnetic
Needle. The experiment is so designed that a wire carrying an equal current of
electricity may be substituted for the electrolyte and the deflection of the mag-
netic needle quantitatively compared. We can throw in series with the electrolyte
or wire at will the lamp, ammeter, and variable rheostat which join to a lighting
circuit or storage-battery. The arrangement of the two switches at either end
will make this clear.
may be quickly substituted for the electrolyte, and how the electric
current may be controlled and made to flow through the electrolyte.
In conducting this experiment the student will be impressed with
the greatly superior conductivity of metals and even high-resistance
alloys over electrolytes. The result of such a carefully conducted
experiment will show that the magnetic effect of electric currents
traversing electrolytes is quantitatively the same as electric currents
of equal strength traversing conductors of the first class. The thought-
60 EXPERIMENTAL ELECTROCHEMISTRY.
ful student will be likely to ask why the effect is just the same when
we have positive ions going to the cathode carrying positive electrical
charges and negative ions going to the anode carrying negative elec-
trical charges. The only answer that can be given is that a negative
ion traveling from right to left tends to turn the magnetic needle in
the same direction as a positive ion does traveling from left to right.
We know that the same current traversing a wire will turn a magnetic
needle to right and left respectively, according to its position above
or below the needle, and that we can greatly increase the magnetic
effect by carrying the wire over and under the needle a number of
times. We may say that a positively charged particle, or ion, pro-
duces -the same effect upon a magnetic needle, traveling from right
to left, as a negatively charged ion does traveling from left to right.
The author has designed an elaborate experiment to show this by
a rapidly running band of pure silk ribbon, upon which are pasted
little tinfoil disks. The band may be run right-handedly under a
suspended magnetic needle, with positive static charges upon the
tinfoil disks and the deflection and direction of the magnetic
needle noted. The direction of the band may then be reversed, and
the disks be charged negatively, when the deflection and direction
of the needle are again noted. The little disks are charged by passing
under and touching a tinsel brush connected with either pole of
an electrical machine of the static type. This is a mechanical repre-
sentation of migrating ions in opposite directions. Owing to the
small quantity charge of electricity upon the disks which were
placed about 3 centimeters apart the ribbon was run at high speed.
The drums over which the ribbon ran were supported upon solid
glass axles to insulate the same. As a matter of fact ions travel very
slowly, but carry very large charges of electricity. In the running
ribbon we have very small charges of electricity, and therefore to
obtain the same magnetic effects we should be obliged to drive the
ribbon and little tinfoil disks at very high speed. The magnetic
needle must therefore be protected from resulting air-currents in
some suitable manner. Having seen the effects of electric currents
passing through electrolytes on magnetic needles, it remains only to
observe the effect of electric currents passing through electrolytes
on masses of ordinary soft iron. Fig. 28 illustrates a simple experi-
ment to measure the "magnetic pull" upon a soft-iron bar, if such
NOVEL EXPERIMENTS IN "ELECTROLYTIC INDUCTION." 6 1
pull exists. At the right-hand side of *he diagram we have an analyt-
ical ba ance with the left pan removed in order that we may suspend
the bar of soft iron to be experimented upon. Directly under the
iron bar is a block of wood or other suitable support for the glass coil
of electrolyte. The platinum electrodes which dip into this electrolyte
connect through an open scale, or delicate ammeter, in series with
the variable rheostat and lamp-bank. The lamp-bank described
in the first chapter of our series may be used. We can now by
this arrangement admit current to the electrolytic coil, read the
current in amperes, and weigh the magnetic pull to a great degree of
accuracy. Substituting the alloy coil, we can, by means of the lamp-
bank and variable rheostat, cause the same current to flow through
FIG. 28. Experiment to Show and Quantitatively Measure the Magnetic Pull of an
Electrolyte Carrying an Electric Current. At the left is an ammeter, a variable
rheostat, and a lamp-bank; above the ammeter is a coil of alloy wire for substi-
tution purposes.
the alloy coil and weigh the pull. As a result of scores of tests of
this character the author found the pull to be just the same with an
electrolyte as it was with a wire carrying the same current. Of
course it goes without saying that the convolutions, and consequently
the ampere-turns, were the same in both cases. For accurate work
it must be impressed, however, that a rather fine wire of high specific
resistance must be employed for the conductor of the first class.
Otherwise the resistance of the wire coil will be so much less than
the electrolyte that we must turn in a great deal of outside resistance
through the agency of the rheostat. Perhaps the most interesting
of all these experiments is that illustrated in Fig. 29, which has been
termed a demonstration of " electrodeless conduction." Here we
.have simply a closed system, an electrolyte without the customary
62
EXPERIMENTAL ELECTROCHEMISTRY.
electrodes for giving and taking the electric current. Let us refer
to the diagram and describe the method of showing this remarkable
phenomenon. The illustration represents an original experiment of
the author performed a number of years ago, but like those preceding,
it has never been published. A represents an alternating-current
generator connected to a coil of insulated wire on the spool which
encloses a soft-iron bar. This soft-iron bar passes into a glass coil
of tubing containing dilute sulphuric acid, and is joined through
two straight glass tubes about a meter long to a second glass coil
o
FIG. 29. Experimental Demonstration of " Electrodeless Conduction." A repre-
sents an alternating-current dynamo connected to a coil of insulated wire; through
this coil a soft-iron bar is passed which enters a glass coil filled with an electro-
lyte; this glass coil is connected by glass tubes with a second glass coil which is
placed within a calorimeter; a cylinder of thin soft Russia iron is placed within
this second coil, which in turn received a very sensitive thermometer.
filled with the same solution. This second coil of electrolyte, however,
is incased within a calorimeter made from a common pasteboard
muff-box, lined within with hair felt, as indicated by the diagonal
lines. Within this coil is placed a small cylinder of thin Russia iron,
which receives in turn and incloses the bulb of a sensitive thermometer,
like those employed in our previous calorimeter work. One of Beck-
mann's thermometers with arbitrary scale and reservoir at the top is
an excellent type. The cover is placed on the calorimeter, and after
equilibrium has been established the thermometer is read and the
NOVEL EXPERIMENTS IN "ELECTROLYTIC INDUCTION." 63
dynamo started. The temperature will slowly rise when within the
little iron cylinder. If the cylinder is removed the mercury in the
thermometer will fall again, and rise once more upon lowering the
cylinder. What part does the little cylinder play ? It is well known
to all physicists and most electricians that iron heats up when
it is magnetized first in one direction and then in the other
by the alternating current. This heating of iron by an alternating
current under such circumstances is called "hysteresis." Here we
have the heating of the little iron cylinder by being rapidly magnetized
first in one direction and then in the other, which gives us proof that
the closed system, without any electrodes whatever, is conducting
the electric current. On breaking the system anywhere, with the
dynamo still in operation, the heating ceases. Here we have un-
doubtedly the ions driven first in one direction and then in the
other, reversing their magnetic effect with their direction. If we
could insert our thermometer in the electrolyte itself, we would
probably get a heating effect due to the "friction" of the ions among
themselves. Fig. 30 illustrates a plan for carefully studying the
effects of alternating currents upon electrolytes of different composi-
tion. The wiring and apparatus is so arranged in this experimental
study as to allow of supplying alternating currents of the same
energy value, but of various frequencies. It has been shown by the
writer with such an experimental apparatus that the frequency of the
alternations, everything else remaining the same, has a decided
effect upon electrolytes. Only a very few years ago little had been
done with the alternating current as applied to electrolytes, and
nothing involving alternating currents with change of frequency.
By "frequency" we mean the number of double reversals of the
current per second. The frequency varies in practice between 25
and 150. The term "period" used in connection with an alternator
denotes the time elapsing between one complete reversal of the cur-
rent. Now if we have free ions in solution which carry the electric
current, they must move back and forth to some extent under the
influence of an alternating current. In other words, they must
oscillate. Now by varying the frequency of our alternations we
vary the rate of oscillation of the ions, and if the heating is due to
friction between the ions, the heating should be greater at higher
frequencies than with low frequencies. Such was found to be the
EXPERIMENTAL ELECTROCHEMISTRY.
NOVEL EXPERIMENTS IN "ELECTROLYTIC INDUCTION." 65
case, the energy value of the alternating current being kept the same.
By a glance at the last illustration we can readily see how the fre-
quency may be changed without altering the energy value of the
current. We can strongly excite the fields of the alternator by admit-
ting a heavy current through the rheostat and driving the alternator
by means of the motor at low speed, when we will obtain an alternating
current of low frequency and of a definite energy value. We can
experiment with this arrangement. We can now turn in our rheostat
and admit a feeble current to the fields of the alternator, and by
driving the armature at a high speed we will be able to obtain the
same energy value for the current, but at high frequency. Experi-
ments were also conducted with electrolytes of various compositions,
that is, with light and heavy ions present respectively. The electrolytes
consisting of light ions invariably heated up quicker than electrolytes
with heavy ions. This can only be explained on the ground of inertia.
The lighter ions travel through the greater distances when oscillating,
and therefore collide a greater number of times. The heavier ions,
because of their greater inertia, do not respond so readily to the
alternations, and therefore move through a lesser distance. They
consequently do not meet with so many collisions, and the friction
is reduced. This of course is theory, but the fact of experimental
investigation remains that the lighter ions cause a more rapid heating
than heavy ions, and that all electrolytes heat up more quickly with
alternating currents of high frequency than they do with alternating
currents of low frequency. It only remains for us to find a theory
to account for the facts.
CHAPTER V.
THE VELOCITY OF ELECTROLYTIC CONDUCTION.
EXPERIMENTS WITH A HIGH-SPEED SPECIAL CHRONOGRAPH CAPABLE
OF DIVIDING A SECOND INTO A MILLION PARTS.
Absolute Velocity of Ions.
IT will be recalled that in the preceding chapter ions were made
to travel by induction. In the experiment with the electrical ma-
chine the two vessels connected by means of the wet string, and
the capillary, the electrostatic charging of the electrolyte took place
at once. In other words, as soon as the electrical machine was
started, bubbles of hydrogen gas made their sudden and immediate
appearance within the capillary. Now, as a matter of fact, the
bubbles of gas would make their appearance at once, whether this
wet-string conductor was long or short. The electrical conduction
would be instantaneous, and yet we will learn a little later in the
present chapter, under the heading "Absolute Velocities of the Ions,"
that the ions themselves move very slowly and have different veloci-
ties. How can we account, therefore, for the instantaneous con-
duction of an electrolyte, when the ions which carry the electricity
upon them move very slowly and have their respective velocities?
We can only account for the facts in such an experiment by attrib-
uting the instantaneous conduction to be due to free ions already
present about the electrodes. Fig. 31 represents an experiment
of Prof. Ostwald to show the instantaneous electrical conduction
through electrolytes. Here we have a glass tube about 50 centi-
meters long and i centimeter in diameter bent at right angles at
the ends and enlarged into cylindrical terminals as shown. At
the left we have a stick of chemically pure zinc supported in posi-
tion by a cork. At the right we have a bent-tube manometer con-
66
THE VELOCITY OF ELECTROLYTIC CONDUCTION.
6 7
taining a little colored water, supported by a good tight cork also.
At the bend on the right a platinum wire is fused in place to act
as the other terminal or electrode. The tube is filled with dilute
sulphuric acid. Upon connecting this piece of apparatus with a
battery, motor-generator, or lamp-bank as described in the first
chapter, making the zinc the anode and the platinum wire the cathode,
bubbles of hydrogen appear instantly upon the platinum wire, and
a pressure is indicated upon the water-gauge. The instantaneous
appearance of bubbles of hydrogen with the closing of the contact
FIG. 31. Prof. Ostwald's Experiment to show Instantaneous Electrical Conduction
through an Electrolyte.
key in series with the source of electricity goes to prove the presence
of free ions already about the electrodes. These free ions merely
give up their charges and escape upon completing the circuit. Now,
if it was necessary for the electric current to first decompose or break
up the molecule of sulphuric acid, then the two atoms of hydrogen,
replaced by the zinc in the SO 4 radical, must have traveled to the
platinum-wire cathode through the tube, which is 50 centimeters
long. Now there are experiments, as we shall see at the close of
this chapter, to determine the absolute velocities of ions, and
measurements upon the velocity of the hydrogen ion show that it
would require a long time for hydrogen ions to travel through a
68 EXPERIMENTAL ELECTROCHEMISTRY.
tube 50 centimeters long. Now, hydrogen appears at once upon
closing the circuit, and we must attribute the immediate response
or conductivity of the solution, to free ions already around the elec-
trodes in readiness to discharge their electricity. Although this
experiment of Prof. Ostwald is one of great interest, it struck the
present writer as being very crude and rough and capable of great
improvement. It does not answer many vital questions. For
example, do all electrolytes conduct with the same velocity? In
other words, will an electrolyte consisting of heavy ions respond
or conduct as quickly as an electrolyte consisting of light ions?
Will all electrolytes conduct as quickly as a metallic conductor?
This appeared to be neglected as a piece of research work; and
with a view of comparing different electrolytes with each other
both in solution and in igneous fusion, and in comparing electrolytes
with metallic conductors, the special high-speed chronograph was
designed and built as illustrated in the following drawings. Through
the agency of this chronograph, a dynamo current was compared
with the current from a set of accumulators, and light thrown upon
such questions as mechanical movement of ions of different weights,
involving the question of inertia. Let us first compare electrolytic
conduction in an electrolyte with metallic conduction, for if the two
act in the same time, the evidence in favor of free ions is strength-
ened. Fig. 32 outlines in diagram the chronograph cylinder and
the electrolyte and wire respectively. Here A represents the elec-
trolyte in the glass tube, and B the parallel metallic conductor.
C is a rheostat in series with the metallic conductor to bring the
same to an equal ohmic resistance with the electrolyte. D repre-
sents a delicate ammeter in series with the electrolyte and the electro-
magnet E of the chronograph. F represents in dotted lines the
same ammeter shifted in series with the metallic conductor and the
electromagnet G of the same chronograph. H illustrates a double
switch for simultaneously closing both circuits after the resistances
of the two have been balanced or made carefully equal to each
other. By revolving the chronograph cylinder and closing the
switch, the two electromagnets will strike the paper band upon
the chronograph cylinder and draw records by means of soft lead-
pencil points. The chronograph and magnets must first be most
carefully calibrated on one and the same circuit by connecting the
THE VELOCITY OF ELECTROLYTIC CONDUCTION.
69
magnets in series with each other, thereby supplying a common,
current of electricity, and adjusting their springs and striking dis-
tances until a current of common value will cause both magnets to
strike upon the rapidly revolving cylinder at the same instant. This.
FIG. 32. Diagram of Author's Method of Studying Time Required for Electric Cur-
rents to Traverse Electrolytes, and to Compare the Time with that Required by
Metallic Conductors. A, electrolyte; B, parallel wire; C, rheostat for balancing
resistance of wire to that of electrolyte; D, mill-ammeter in series with electro-
lyte and magnet E; at F the mill-ammeter is shown shifted in series with wire
and magnet G; H, double switch for closing both circuits simultaneously.
can most easily be seen by the pencil records. When by careful
experiment and adjustment the two electromagnets strike "abreast'*
upon the flying cylinder, which is driven by a high-speed elecrtic
motor, the series connection is changed and each electromagnet
is placed separately in circuit with electrolyte and wire respectively,
previously made of equal ohmic resistance. The lines upon the
7 EXPERIMENTAL ELECTROCHEMISTRY.
cylinder in this drawing illustrate the appearance of the pencil
record when the cylinder is driven at moderately high speeds. A
photograph of such a chronograph is given in Fig. 33, where an
electric motor is directly connected by means of a flexible coupling
to reduce vibration. It was soon found, however, with such a
FIG. 33. Photograph of Simple Drum High-speed Chronograph Direct Connected to
Electric Motor.
simple chronograph cylinder, when driven at very high speeds,
that the pencil records were drawn all the way around, and it was
impossible to see where the contacts were first made. It became
FIG. 34. Rear View of Electrochronograph provided with Electric Motor, Balance-
wheel, and Revolution-counter. This instrument is a modification of that shown
in Fig. 3, as it drives a long band of paper for receiving record.
necessary to expand the chronograph by driving a long band of
paper. Figs. 34 and 35 will make the plan clear. In Fig. 34 the
THE VELOCITY OF ELECTROLYTIC CONDUCTION. 71
end of the chronograph cylinder is again shown, with its electric
driving motor at the right, and with a heavy balance-wheel to steady
its rapid motion at the left. A revolution counter is also depicted
at the extreme left, pressed against the shaft of the chronograph
7^ EXPERIMENTAL ELECTROCHEMISTRY.
cylinder. This revolution counter was afterward moved to the
shaft of the pulley at the far end of the band, as being a fairer place,
for in case there was a slight creeping of the band upon the chrono-
graph cylinder, there would be no error introduced from this cause.
Fig. 35 illustrates a side view of this special form of band chrono-
graph, showing its band and supporting drum-wheel at the far end
of the work-table over which it runs. The arrangement of the
marking-pencils and electromagnets is> made clear in this illus-
tration. Upon the work bench or table are the electrolyte and
wire respectively, together with a cell or storage battery and a
special form of U tube used for various conduction experiments
with electrolytes. This particular chronograph revolves at the
rate of two thousand revolutions per minute, and it will be seen
that the slightest "lag" in conductivity in either circuit, when the
two are closed simultaneously by a proper key, will be shown accu-
rately and quantitatively upon the moving band. With this ar-
rangement, as will be seen from the following mathematical exposi-
tion, a second may be divided into one hundred thousand parts;
and by higher speeds, the second may be laid off and divided into a
million parts, dependent upon the behaviors of the various con-
ductors experimented with. The diameter of the chronograph
cylinder being 15 centimeters, we can take this as a basis upon
which to start the calculation.
3-I4I59
15 cm. diameter
1570795
47-12385 cm. circumference
With 2000 revolutions per minute, we have 47.12385X2000 =
94247.70 centimeters per minute. The space traveled during one
second is therefore
60)94247.70(1504.1283 cm.
In i/ 10 second we have 150.41283 cm.
In i/ioo second " 15.041283 cm.
In i/iooo second " 1.5041283 cm.
In i/ioooo second " .15041283 cm.
THE VELOCITY OF ELECTROLYTIC CONDUCTION. 73
Working with a chronograph of still higher speed, the cylinder
being driven by a two-horse-power motor belted up for speed, the
scale upon the flying band was of course still more open, and allows
of determinations to be made to 1/100,000 and even 1/1,000,000 of
a second. A tabulated length of spaces upon this high-speed band
is as follows up to hundred-thousandths of a second. The figures are
as follows:
Cm. Circumference. Rev. per Min. Cm. Traveled.
47.12385 X loooo = 471238.5 cm. per min.
60)471238.5(7853.97500 cm. per second.
In i/io second we have 785.397500 cm.
In i/ioo second 78.5397500 cm.
In i/iooo second " 7-85397500 cm.
In i/ioooo second " .785397500 cm.
In i/ioooooo second " .0785397500 cm.
For higher speeds still and correspondingly more minute subdivisions
of the second, a chronograph rigged like that shown in Fig. 36 was
experimented with. Here we have an electrolyte 50 feet long in the
glass tube arranged like a steam-radiator, and the chronograph
cylinder driven at enormous speed by the multiplying system of
belting to the countershaft, etc. The telephone-receiver, cell of
battery, induction-coil, and resistance set depicted here were employed
to balance the resistances, instead of the ammeter employed in
the slower-speed design of instrument. The method of measuring
resistances by means of the telephone and induction-coil is known as
Kohlrausch's method, and consists of the simple Wheatstone bridge
arrangement, with a telephone-receiver in the place of a galvanometer,
and the alternating current from the secondary of a small induction-
coil instead of a simple battery of cells. With this arrangement the
alternating current produces a humming sound in the telephone
when the bridge is out of balance. The alternating current in addi-
tion does not decompose the electrolyte, and allows of conductivity
determinations being made with great accuracy. Having described
the apparatus, some of the results will now be given. The first
experiments were made with an electrolyte consisting of dilute sul-
phuric acid in the proportion of 10 cubic centimeters of H 2 SO4,
specific gravity 1.84664, in 40 cubic centimeters of distilled water, and
74
EXPERIMENTAL ELECTROCHEMISTRY.
a wire of German silver made equal in resistance by means of a
rheostat of the non-inductive type. This is an important point to
IH CX
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observe in determining all resistances where only a momentary
current is to be dealt with. For a dissertation upon the subject of
THE VELOCITY OF ELECTROLYTIC CONDUCTION. 75
non-inductive resistances, the student must be referred to any standard
work on physics dealing with electrical measurements. The present
writer may say, however, that with common coil resistances, or
rheostats, there is a choking or damping effect upon electrical impulses
of short duration, due to the phenomenon of self-induction. Having
balanced the respective resistances of the electrolyte and wire with
its non-inductive resistance rheostat in series, the chronograph was
speeded up, and when a rate of 2000 revolutions per minute was
reached, as counted by the revolution-counter upon the drum-wheel
shaft, the key was closed three times in rapid succession and the
chronograph stopped. Three records had been made. At first it
was found that the electromagnet in series with the electrolyte
struck a trifle in advance of the electromagnet in series with the
wire, the marking on the band leading by 0.75 centimeter, indicating
that the conductivity through the electrolyte was ahead by 1/10,000
of a second. What was this due to ? Although the two resistances
were balanced as carefully as possible, the leading of the electrolyte
was undoubtedly due to the fact that its resistance was slightly
lower than that of the wire. The resistances were carefully
rebalanced, using the most refined means, when the two con-
ductors finally "struck abreast," so to speak, upon the flying
band. Electrolytes of various composition were substituted for
the sulphuric acid and carefully compared with the wire, and
in every case where the resistances had been perfectly balanced
the two electromagnets struck abreast. The highest speeds of
bands were of course obtained with the large instrument as shown
in Fig. 36, and with this equipment, electrolytes of various lengths
were experimented with. An electrolyte 50 feet in length conducted
as quickly as an electrolyte only a few centimeters long. It mattered
not whether we used an electrolyte with light or heavy ions, the rate
of conductivity, or response to the electric current, was the same.
Experiments were also conducted with storage-batteries as a source
of electricity, and it was recorded here again that the electricity left
the free ions in the storage-cells as readily as it did a wire charged
by a dynamo current. Experiments of this character were repeated
many times, and the writer believes one is justified in stating the
law that electrolytes of equal resistance conduct the electric current with
a definite velocity regardless of the composition of the electrolytes or the
76 EXPERIMENTAL ELECTROCHEMISTRY.
length of the containing vessel. It may also be stated that an elec-
trolyte conducts the electric current as quickly as a conductor of the first
class, regardless of Us composition, provided we have an equal ohmic
resistance of a non-inductive type. In working with fused electrolytes
the same quantitative behavior was observed, the electric current
flowing as quickly after contact as with all metallic conductors.
Free ions must therefore be around the electrodes and in contact
with them. If molecules had to be first broken down into ions, and
these ions had to travel, there would undoubtedly be a lag in experi-
menting with electrolytes consisting of heavy ions, for the question
of inertia would be involved. The same impulse which would start
up light ions in a given time would fail to produce the same response
where heavier ions were concerned. Having shown the instantaneous
behavior of electrolytes toward the electric current, we are now in a
position to study the experimental methods for measuring the abso-
lute velocity of ions.
Experimental Methods for Showing the Absolute Velocity of Ions.
Lodge's Apparatus.
It has been stated that all ions had their respective velocities, and
that these velocities were exceedingly small. It has been demon-
strated by Bredig, and also by Ostwald, that the velocity or mechanical
motion of the ions is a function of their atomic weights. This rela-
tionship was brought out by series of long and patient research,
but the reason for such behavior is not understood. We have in
chemistry several striking cases of periodic behavior, although we
have so far been unable to account for them. If we arranged the
ions in a table according to their migration rates, we would find that
hydrogen is the swiftest of them all, although its movement through
an electrolyte requires considerable time. Let us look into the
method of Lodge, and learn just what the speed of the hydrogen ion
is. Fig. 37 illustrates the apparatus of this physicist for determining
the speed of the hydrogen ion under a given potential gradient.
Here we have two beakers or glass jars joined by a siphon-tube bent
at right angles at each end. A centimeter-scale is attached to the
under side of this tube as indicated. This glass siphon-tube contains
an aqueous solution of gelatine put in hot, which solidifies when
THE VELOCITY OF ELECTROLYTIC CONDUCTION.
77
cold, forming a jelly. Now this solution of gelatine also contains
some sodium chloride, NaCl, to serve as the electrolyte, and the
entire solution is colored red by the addition of a little phenolphtha-
lein made alkaline with a few drops of sodium-hydroxide solution.
The gelatine is dissolved in hot water in a beaker and some
common salt is added and stirred until a perfectly homogeneous
solution is obtained. A little phenolphthalein is then stirred in
and made red by adding a few drops of the sodium-hydroxide
solution. This mixture is kept near the boiling-point of water for
a few minutes, and is then poured into a number of tubes bent at
r
FIG. 37. Lodge's Apparatus for Experimentally Determining the Absolute Velocity
of the Hydrogen Ion.
right angles to form siphons like that illustrated. Care must be
taken to avoid the inclosing of air-bubbles, and the tubes are put
away to cool and solidify with the bent ends turned up. To measure
the velocity of the hydrogen ion, one of the tubes after cooling is placed
dipping into the two beakers as shown, and the beakers filled with a
dilute solution of sulphuric acid. Two platinum electrodes are put in
place and connected to our motor-generator or lamp-bank, with a
voltmeter joined across the electrodes to show the potential gradient
under which we are working. All ions have a fixed velocity under a
set potential gradient. Now, what takes place when a current of
electricity is made to pass through this system? The hydrogen ion
from the electrolyte of sulphuric acid starts from the anode in the
right-hand beaker and makes its way to the cathode in the left-
hand beaker through the composition in the siphon-tube. What
happens there? The hydrogen simply displaces the sodium from
the sodium chloride present and forms hydrochloric acid, H + NaCl =
HC1, which decolorizes the gelatinous solution of phenolphthalein.
78 EXPERIMENTAL ELECTROCHEMISTRY.
This indicator is red in the presence of a base and colorless in the
presence of an acid. As the hydrogen ion proceeds through the
siphon-tube, it replaces the sodium in the sodium chloride, and
bleaches out the phenolphthalein marking its way through the
composition. The experiment is an interesting one to watch, as the
decoloration proceeds at a slow rate. Lodge worked with a potential
gradient equivalent to a drop of one volt a centimeter. If we have
a tube 50 centimeters long, therefore we must use a difference in
potential of 50 volts, and must employ our lamp-bank for this, unless
we have at hand a dynamo wound for a current output at 50 volts.
For a short tube we can use our motor-generator. Working with
such a piece of apparatus with a drop of one volt per centimeter,
Lodge found the absolute velocity of the hydrogen ion to be about
if centimeters per minute. In three determinations Lodge found the
hydrogen ion to travel :
1.1560 centimeters per minute.
1.1740
1.1440
The average of these three determinations with the above apparatus
being 1.1580 centimeters per minute, or considerably over an hour
for this, the swiftest of all ions, to travel a meter; and yet as shown
by the electrochronograph work, an electric current leaps through
an electrolyte, so to speak, in exactly the same time as it does
through a wire.
Whetham's Method.
Another experimental method for determining the absolute
velocities of ions was devised and used by Whetham, the apparatus
being illustrated in Fig. 38. He describes his method as follows:
"Suppose we have two solutions like copper chloride and ammonium
chloride, containing one ion in common and having nearly equal
conductivities. Let one solution be colored and have a density
different from that of the other. The denser solution is first poured
into the longer arm of a kind of U tube, and then the other is allowed
to flow gently on to its surface from the shorter arm. If a current
is passed across the junction between the two solutions, it carries
the copper and ammonium ions with it and drives the chlorine ions in
THE VELOCITY OF ELECTROLYTIC CONDUCTION.
79
the opposite direction. Since the color depends on the presence of
the copper ions, the boundary will travel with the current, and by
measuring its velocity the speed of the
ions under unit potential gradient can
be calculated." There are several other
methods for determining absolute veloci-
ties of ions, and in the hands of careful
investigators the results agree very strik
ingly. As will be seen later, there are
methods for determining the relative
velocities of the ions, and it will be seen
at once if we have the absolute velocity
of one ion accurately determined, and
we can ascertain in other ways the rela-
tive velocities of the remaining ions, we
can calculate the absolute velocities of
them all. Having learned about electro-
lytes, dissociation, ionic velocity, etc., we
will be in a good position to take up
and appreciate work of a little more
practical character in our next chapter,
and introduce the student to the beauti-
ful work of Faraday in electrochemical
science. Here we will study the quantitative relation of the electric
current to electrolytes, and take up the energy relations between
chemistry and electricity, and lay the foundation for work of a very
practical and useful character.
FIG. 38. Whetham's Appara-
tus for Experimentally Deter-
mining the Absolute Veloci-
ties of Ions.
CHAPTER VI.
FARADAY'S LAW.
DISTINCTION BETWEEN CURRENT REQUIRED AND ENERGY ABSORBED.
EXPERIMENTS TO SHOW MECHANICAL MOVEMENT OF MATERIAL.
THE idea occurred to Faraday to send a current through several
electrolytes connected up in series, and to make weighed compari-
sons of the materials separated by the same current in a given time.
In 1883, the result of Faraday's most important and valuable work
was framed into the following law by H. Von Helmholtz:
" The same quantity of electricity passing through an electrolyte
either sets free or transfers to other combinations always the same
number of valencies. 1 '
Von Helmholtz, in his "Faraday lecture" delivered in London,
on April 5, 1881, laid the foundation of a new electrochemical theory
which explains the facts embraced by Faraday's law. The most
important of these facts may be stated in this sentence:
"Every single valency of an elementary or compound ion is
charged with exactly the same quantity of positive or negative
electricity, which behaves as if it were an electrical atom that can-
not be further divided."
As the work of Faraday is of such great moment in theoretical
and practical electrochemistry, the actual words of Von Helm-
holtz are given, the author believing that the facts as learned by
Faraday are among the first, if not the most important, of all here
recorded. The words used by Von Helmholtz are these, as trans-
lated by M. M. Patterson Muir:
"The same definite quantity of either positive or negative elec-
tricity moves always with each tmivalent ion, or with every unit
of affinity of a multivalent ion, and accompanies it during all its
motions through the interior of the electrolytic fluid.
80
FARADAY'S LAW. 81
" This quantity we may call the electric charge of the atom.
" 1 'I we accept the hypothesis that the elementary substances are
composed of atoms, we cannot avoid concluding that electricity also,
positive as well as negative, is divided into definite elementary por-
tions, which behave like atoms of electricity. As long at it moves
about in the electrolytic fluid, each ion remains united with its
electric equivalent or equivalents.
"At the surface of the electrodes decomposition can take place
if there is a sufficient electromotive force, and then the ions give
off their electric charges and become electrically neutral."
From this work the valuable table of electrochemical equivalents
was compiled, which is of the utmost importance in all practical
electrochemical work. The definition of an electrochemical equiva-
lent being capable of expression in several ways, it should be care-
fully studied and appreciated. If the quantities of all ions which
stand to one another in the relations of their combining weights
carry equal quantities of electricity, it will at once be appreciated
that it is of great scientific importance to know the exact amount
of electricity which a unit quantity of ions will carry. This can be
determined by passing a given quantity of electricity through a
solution of an electrolyte and weighing the amount of metal de-
posited upon the cathode, or measuring the amount of gas liberated
and calculating its weight from its volume. This has been done
very carefully by Lord Rayleigh and Mrs. Sedgewick, who found
that one coulomb of electricity deposits 1.1179 milligrams of
silver. W. and F. Kohlrausch, working with equal care, found
under the same conditions 1.1183 milligrams. The mean of
these figures is 1.1181 milligrams. A more recent determina-
tion of the electrochemical equivalent of silver by Richards, Col-
lins, and Heimrod gives 1.1172 milligrams of silver as equiva-
lent to one coulomb. A still more recent determination by Pat-
terson and Guthe gives the slightly larger value of 1.1192 milli-
grams as equivalent to one coulomb. This agrees with the mean
result obtained by Pellat and Portier, and is very close to the
number obtained by Kahle, 1.1193. The mass of the ions taken is
purely arbitrary. Here, as in so many other cases, it is convenient
to use the gram-molecular weight for univalent and the gram-
equivalent weight for polyvalent ions. For all practical purposes
82 EXPERIMENTAL ELECTROCHEMISTRY.
the electrochemical equivalent of silver, which is usually referred
to as a standard for determining the other values for the other ele-
ments, is set down as the fraction of a gram as equivalent to one
coulomb, thus: 0.0011193 gram. The atomic weight of silver in
terms of oxygen = 16 is 107.93. I n order to separate a gram-
atomic weight of silver it will require, using W. and F. Kohlrausch's
mean of .0011181,
107.93
=96,530 coulombs of electricity.
This number of coulombs, 96,530, as will be seen, will separate
the gram-atomic weight of any univalent body, and is sometimes
called the electrochemical equivalent of electricity. The fact may
be stated thus:
One chemical equivalent of any electrolyte expressed in grams re-
quires the passage of 96,530 coulombs for its liberation or electrolysis;
96,530 coulombs, therefore, are capable of liberating the chemical
equivalent of any electrolyte.
This is an exceedingly important constant for us to remember
in our practical work. This number varies in value a trifle accord-
ing to different investigators, and will be encountered as 96,540,
etc. In the table opposite the chemical equivalents of some of
the most important elements are given. The student must not
confound a chemical equivalent with an electrochemical equivalent.
These terms must be clearly separated in his mind or else he will
be continually getting into confusion. A chemical equivalent is
simply the atomic weight of a substance divided by its valence. The
atomic weight of oxygen being 16, and its valence 2, the chemical
equivalent of oxygen would be 8. Oxygen = 16. Chemical equiva-
lent 16-^2 = 8.
In the table on page 84 the electrochemical equivalents of some
of the most important elements are given. By dividing the atomic
weight in the second column by the valence in the third column,
the chemical equivalent given in the fourth column is obtained,
and this number multiplied by the electrochemical equivalent of
hydrogen in micrograms per coulomb gives the electrochemical
equivalent of the ion in the fifth column, also in micrograms per
coulomb. The numbers in the sixth column are the reciprocals of
FARADAY'S LAW.
CHEMICAL EQUIVALENTS OF CERTAIN ELEMENTS.
Element.
A. W.
C.E.
Element.
A. W.
C. E.
A1"'
27.1
901
Pb"
206.9
IO3.4X
Ba"
3 valence
137-4
67.8
Li
2 valence
7-03
7-O3
Br
2
79.96
70. 06
MR". .
i
24.36
12. l8
Cd"
I
II2.4
z6. 2
Mn"
2
55
27. s
Ca"
2
4O. I
20 ol
Hz. .
2
203.3
203. 3
Cl
2
35-45
3S 4?
He". . .
I
203-3
100. 15
Cr"
i
52-1
26 oc
N'"
2
14.01
4.67
Cu
2
6 3 .6
63.6
Ni"
3
58.7
29.35
Cu"
I
63.6
31.8
O"
2
16
8.00
F
2
19
i o oo
K
2
39 -15
39. ic
Au
I
197.2
6< 73
Ag...
i
107.93
107.03
H
3
j .008
i 008
Na
i
23-05
23. o?
I
i
126.815
126 8?
Sn""
i
119
26. 7?
Fe"
i
55-9
27 OC
Sr"
4
87.6
43. &
Fe"'
2
55-9
l8.63
Zn"
2
65.4
32. 7
3
2
those in the fifth given in grams. The electrochem'cal equiva-
lents of compound ions, such, for example, as the univalent radical
hydroxile OH and the bivalent radical SO 4, are similarly obtained,
the chemical equivalent of such a radical being the sum of its com-
p nent atomic masses div ded by its valence.
This table brings out the beautiful tru h of Faraday's law, and
forcibly indicates the great value of the facts he was able to point
out as a result of his famous investigations. Let us experimentally
test Faraday's law in the laboratory and put down our results. For
this purpose we will set up a piece of apparatus with five different
EXPERIMENTAL ELECTROCHEMISTRY.
electrolytes in series and weigh the cathode products liberated by
the same current-flow.
TABLE SHOWING RELATIVE WEIGHTS OF BODIES LIBERATED BY
A COMMON ELECTRIC CURRENT.
Element.
Atomic
Mass.
Valence.
Chemical
Equivalent.
Electrochemical Equivalents
Micro-
grams per
Coulomb.
Coulombs
per
Gram.
ISt
Hydrogen
2d
I
15.96
35-37
14.01
27.04
206 . 40
64.88
58.60
199 . 80
199.80
63.18
63.18
107. 70
196. 2
3d
I
2
I
3
3
2
2
2
2
I
2
I
I
3
4th
7.98
35-37
4.67
9.01
103.20
3 2 -44
39-3
99.90
199.80
31-59
63.18
107.70
65.40
5th
10.38
82.83
367.10
48.47
93-5
1071 .00
336.7
304.20
1037.00
2074.00
327.90
655.80
i i i 8 . oo
678.90
6th
96,340
12,070
2,724
20,630
10,700
933-7
2,970
3,287
964-3
482.2
3>5
1,525
894.5
i,473
Oxygen
Chlorine
Nitrogen
Aluminium
Lead
Zinc
Nickel
Mercury
Mercury
Copper. . .
Copper. . .
Silver
Gold
EXPERIMENTAL DEMONSTRATION OF FARADAY'S LAW.
Referring to Fig. 39, we have at the left a Hoffmann apparatus,
A, for the electrolysis of solutions yielding gaseous products at the
electrodes. The gases libe ated es ape into the two tubes and
press the solution up into the reservoir by the central tube. By
opening the stopcocks at the tops of the two side tubes containing
the gases, the weight of the solution in the reservoir will force the
gases out, when they may be collected in a most convenient man-
ner. Hydrogen, for example, may be burned as a jet, after elcctro-
lyzing a dilute solution of sulphuric acid, or as we learned in the
first chapter, by electrolyzing a solution of potassium or sodium
hydroxide. Oxygen and hydrogen will be liberated in such a piece
of apparatus in the ratio of two volumes of hydrogen to one volume
of oxygen whether we use dilute ulphuric acid, or a sodium, or
potassium-hydroxide solution. For the present demonstration of
the law of Faraday, we will fill the Hoffmann apparatus with a
dilute solution of sulphuric acid in distilled water in the proper-
FARADAY'S LAW. 85
tion of about i to 10. In the cell B we will place a concentrated
neutral solution of silver nitrate in distilled water. In the cell C
we will place a solution of cuprous chloride, which may be made by
dissolving a few grams of the salt in hydrochloric acid after
having washed it carefully on a filter paper with warm distilled
water. In the cell D we will place an electrolyte consisting of cop-
per sulphate slightly acidulated with nitric acid. In the cell E
we will use a solution of stannic chloride, best prepared in the fol-
:fg=
ffifS
^>=3
fS"i
FlG. 39. Simple Apparatus for Experimentally Demonstrating Faraday's Law.
A, Hoffmann apparatus; B, C, D, and E, cells containing electrolytes and elec-
trodes; F, delicate ammeter.
lowing manner: Take of stannous-chloride crystals 1000 grams;
hydrochloric acid, specific gravity 1.125, II 7 cubic centimeters;
nitric acid, specific gravity i;22o, 435 cubic centimeters; and dis-
tilled water 1000 cubic centimeters. Put the stannous chloride into
a 1 2 -inch evaporating-dish and add the 1170 cubic centimeters
of hydrochloric acid; warm on the steam-bath and stir until the
salt is dissolved; then dilute with one liter of hot water. If the
solution does not remain clear, there is a deficiency of hydrochloric
acid, in which case add very concentrated hydrochloric acid, a
few drops at a time, until the solution becomes clear. Add the
nitric acid, a few cubic centimeters at a time, to the warm solu-
tion, stirring well after each addition. After a considerable part of
the nitric acid has been added, test a few drops of the solution with
a drop of mercuric-chloride solution. If a white precipitate falls,
86 EXPERIMENTAL ELECTROCHEMISTRY.
stannous chloride is present, and more nitric acid is needed. When
no white precipitate falls, the oxidation is complete, and no more
nitric acid should be added. Put the liquid product into a tightly
stoppered bottle. The cathodes of all the cells are to be of
platinum, but we must have a silver anode in the silver-nitrate
solution and copper anodes in both of the copper solutions. The
anode in the tin solution may be of platinum. It is needless
to say that both of the electrodes in the Hoffmann apparatus are
of platinum. These five electrolytes are now all connected in series
with a delicate ammeter, as shown, and the terminal wires run to
a storage-battery or such a motor-generator as described in the
first chapter. The electrolysis may be allowed to proceed for any
length of time within the capacity of the Hoffmann apparatus.
The longer the run the better, the errors in weighing a decided
increase in the respective cathodes being less than in weighing a
slight increase. In this system we will have hydrogen liberated,
the monovalent element silver, the monovalent copper, the divalent
copper, and the tetravalent tin. If the experiment has been
conducted without error and losses, we will have for each gram of
hydrogen liberated 107.93 grams of silver, 63.6 grams of copper
in our monovalent copper electrolyte, 31.8 grams of copper in our
divalent copper solution, and 26.75 gramms of tin from the tin
solution. The following table shows the result of a carefully con-
ducted experiment with the five electrolytes described above :
8 - 3g - AS/ - 6 3'5g-Cu'. 3i.45g.Cu". 2 8. 29 g.Sn"".
Atomic weight ....... 107.93 63.6 63.6 119.00
Here we can see that the monovalent elements separate in pro-
portion to their atomic masses, the divalent elements in propor-
tion to their atomic masses divided by two, the tetravalent element
in proportion to its atomic mass divided by four. The beauty of
this law is very striking, and it may be said that Faraday's law
knows no exceptions. There can be no electrolytic conduction
without the corresponding setting free of substances in the ratios of
their chemical equivalents. As we shall see presently, we have
ample proof of moving particles, or an actual mechanical transfer
of matter when an electric current is passed through an electrolyte.
In view of the mechanical transfer of matter, experiments were con-
FARADAY'S LAW.
ducted upon electrolytes under heavy pressure to learn if Faraday's
law held true under such conditions. At first it was noted that
the electrical conductivity was increased. In other words, more cur-
rent passed through the electrolyte than was accounted for by weighing
the cathodes. As a matter of fact the method of conducting the
experiment was faulty. The pressure was put upon the electrolyte
by air, some of which was of course forced into solution, and ionizing
carried a portion of the electric current. Professors Nernst and Ostwald,
in Germany, tested Faraday's law most critically by electrolyzing
solutions with exceedingly feeble currents to see if any electricity at
all was conducted without corresponding quantitative decomposition
of the electrolyte. In one experiment upon dilute sulphuric acid
they caused an exceedingly small amount of electricity to pass only
0.000005 coulomb. They determined the minute quantity of gas
set free and found that Faraday's law held even for such a small
electric current. In the large commercial electrolytic copper refiner-
ies the law has been tested upon enormous scales by the passage of
millions of coulombs and found to hold absolutely. The law of
Faraday in the light of the many attacks and investigations upon it
seems to be one of the very few in chemical and physical science which
have stood throughout without suffering exception of any kind.
We have now learned that chemical equivalent quantities of all ions
have the same capacity for electricity. It is a striking and interesting
fact to note that this is analogous to the law of Dulong and Petit,
which states that all atoms have the same capacity for heat. If we
multiply the atomic weights of the elements by their specific heats we
obtain almost a constant, which number we term atomic heat. The
following table containing a few elements for the purpose of illustra-
tion is of interest here.
Element.
A.,
Atomic
Weight.
s.,
Specific
Heat.
A.XS. f
Atomic
Heat.
Potassium
7Q
.166
6 e
Calcium ,
4O
. I7O
6 8
Manganese
CC
. 122
6 7
Tin
118
.CK4
6
Gold
107
.032
6 3
Mercury
2OO
.032
6 4
Lead
2O7
.O3T
6 4
Bismuth.
2OQ
.O3O
6 ?
Silver
108
.0^6
6 o
EXPERIMENTAL ELECTROCHEMISTRY.
To this law there are some exceptions, but in the majority of
cases we have practically a constant.
VOLTAMETERS.
METAL AND GAS TYPE. THE SILVER VOLTAMETER.
The most accurate instrument for measuring current-flow is
without question the silver voltameter, and is at the same time the
most easily constructed.
Fig. 40 will make the arrangement clear. Two comparatively
heavy plates of pure silver are joined together to make one electrode,
between which a thin silver sheet forming the other electrode is
suspended. In the silver voltameter the very high equivalent of
FIG. 40. Construction of Silver Voltameter for the Measurement of Current-flow in
Coulombs. A, top view looking down into cell; B, end view; C, side view.
silver, and consequently the great mass isolated upon the cathode by
comparatively feeble currents, reduce the errors in weighing to a
minimum. There is one disadvantage in the use of the silver voltam-
eter, however, and this is due to the fact that silver tends to precipi-
tate out upon the cathode in crystalline form, and if the electrolyzing
current is strong, some of the crystals form so quickly that they will
drop off and introduce errors from this cause. The electrolyte for
this instrument is of a concentrated neutral silver-nitrate solution.
With such an instrument a current of one ampere deposits upon the
thin silver cathode
0.0011181 gram silver per second,
or 0.067086 gram silver per minute,
or 4.025160 grams silver per hour.
FARADAY'S LAW. 89
With such an instrument in series with an electric current, by
determining the weight of the cathode before and after the run we
are in a position to calculate the number of coulombs that have
passed through the system. If we had an absolutely steady current
we could get at the same thing by putting a correct ammeter in series
and multiplying the amperes indicated by the number of seconds
during which the current passed and get the coulombs used. If the
strength of the current varied, however, this plan would be worthless.
With voltameters of either the metal or gas type the current may
vary, flowing at very different rates in a given time, but the increase
in weight of the cathode will give the true number of coulombs
regardless of such fluctuations.
THE COPPER VOLTAMETER.
Here we have a similar construction, only with plates of pure
copper instead of silver ones. The electrolyte consists of a solu-
tion of 30 grams of crystallized chemically pure copper sulphate
dissolved in 200 grams of distilled water, to which 5 grams of
chemically pure concentrated sulphuric acid is added and 5 cubic
centimeters of ethyl alcohol. Such an instrument is inexpensive and
is adapted to the most general requirements. The copper voltameter
can be left in circuit with work through great lengths of time without
fear of losing any copper by falling from the cathode. The copper
will be deposited upon the cathode as a beautiful, salmon-pink
metal. Such an instrument will answer every purpose as an ampere-
hour meter for heavy work if the plates are made generous enough
in area. With this voltameter a current of one ampere deposits
upon the thin copper cathode
0.00033 gram of copper per second,
or 0.01980 gram of copper per minute,
or 1.18800 grams of copper per hour.
This will be found the most satisfactory instrument for general
work and every student in practical electrochemistry should set one
up for his current measurements. With very small currents, how-
ever, there is apt to be a slight error introduced with the use of the
copper voltameter because of some cuprous oxide being deposited
90 EXPERIMENTAL ELECTROCHEMISTRY.
along with the copper. If we have at least o.i of an ampere
following through the instrument, and its cathode area is at least
100 square centimeters, there will be no trouble from this cause,
especially when the electrolyte contains the sulphuric acid and
alcohol as given in the above formula.
THE GAS-VOLTAMETER.
In this instrument either dilute sulphuric acid or a solution of
potassium or sodium hydroxide may be used and the mixed oxygen
and hydrogen gases determined by volume. Dilute sulphuric acid
is perhaps the best for the purpose, and should be mixed in the
proportion of one part of strong sulphuric acid to ten or twelve
parts of distilled water. The acidulated water is decomposed
between two platinum plates and gas collected and reduced to o
and 760 millimeters by the well-known formula for reducing
gases to a standard for comparison. This type of voltameter,
as illustrated in Fig. 41, is very convenient because it does
FIG. 41. Approved Form of Gas Voltameter for the Measurement of Current -flow
in Coulombs. The student must not confound the word voltameter with the
word voltmeter.
away with all weighings. The volume can be read, and by means
of tables, when the temperature has been taken, the gas volume
can be quickly reduced to a standard. This instrument is not
quite as accurate as the copper voltameter, and, moreover, requires
about two volts of electrical pressure to drive a current through
it. If we are using a storage-battery or the motor-generator we
FARADAY'S LAW. 91
must take into consideration the two volts required by such a volt-
ameter in making any calculations. To use the voltameter the water
in the vertical measuring-tube should first be saturated with oxy-
hydrogen gas, by allowing it to fill about half full, and then discon-
necting and shaking, with the thumb or a stopper closing the end of
the tube. With this type of instrument an ampere flowing liberates
0.1740 cc. of oxyhydrogen gas per second,
or 10.4400 cc. of oxyhydrogen gas per minute,
or 626.4000 cc. of oxyhydrogen gas per hour.
The above volumes are at o and 760 millimeters pressure. The
vertical tube for collecting the gas is best graduated direct into cubic
centimeters, reading to tenths. Such a gas-voltameter must be
ordered from one of the chemical supply houses unless the student
is an expert glass-blower, but one answering every requirement can
be quickly made in the laboratory without the art of glass-blowing,
as illustrated at A in Fig. 42. Here we have a glass cylinder fitted
FlG. 42. Gas-voltameter and Large Collecting-jar in Series with Three Electrolytic
Cells equipped with Voltmeters. A, easily constructed gas-voltameter; B, col-
lecting-jar; C, D, and E, electrolytic cells with voltmeters for determining the
electrical energy expended in each electrolyte.
with a tight stopper (the whole success of the apparatus depends
upon its being tight), containing two concentric platinum cylinders,
one forming the anode and the other the cathode. The large bell
jar, B, allows of a long run before it is necessary to stop and measure
the gas volume. The bell jar may be graduated and its readings
taken for rough work, but for close work the gas must be transferred
to a more delicate graduate and reduced to standard conditions.
92 EXPERIMENTAL ELECTROCHEMISTRY.
Now we have just learned that the same current will deposit
upon the cathodes of electrolytic cells in series chemical equivalents
of the elements as well as chemical equivalents of compound ions.
Chemical equivalents of all bodies are therefore liberated or deposited
by the same current. Is the absorption of energy the same in these
different cases? It is not, and the experiment illustrated in Fig. 42
has been designed by the writer to bring out this most important
point. Let us place three different electrolytes in the cells C, D, and
E, and connect across the electrodes in each case a delicate voltmeter
to indicate the fall of potential in each cell. Although the coulombs
passed will of course be equal in each cell, because they are in series,
and the metals will be deposited in the ratio of their chemical equiva-
lents, the number of joules expended in each cell will be different.
Those who have studied electricity will know that the joule is the
unit of electrical energy, and is the product of the ampere by the
second, by the volt. As a coulomb is the product of an ampere
by a second, we may say that the joule is the product of a coulomb
by a volt. It is not the intention of the author to go into the reason
for unequal absorptions of energy in different electrolytes until the
next chapter, when the matter will be fully dealt with. It is, however,
the wish of the writer to impress upon the student that there is a
difference, and that its explanation, as will be learned later, is in
accordance with the doctrine of conservation of energy, and a point
of great beauty in electrochemistry. With an arrangement of appara-
tus as indicated in Fig. 42, therefore, we would be able to calculate
by means of the gas-voltameter in series, the number of coulombs
which have passed through the system, and by the respective readings
of the voltmeters across the electrodes of the respective cells we
would be able to determine the amount of energy absorbed in each.
So much upon this point for the present. It was stated that we
should have introduced in connection with Faraday's law, experi-
mental evidence proving that the passage of an electric current
through an electrolyte was accompanied by an actual mechanical
transfer of matter. A very simple and a very beautiful experiment
may be quickly performed which goes to show the transport of
ponderable material. Arrange a large U tube, as illustrated in
Fig. 43, with stoppers and platinum electrodes. Through one
stopper bore a small hole to receive a bent glass tube as shown.
FARADAY'S LAW.
93
Prepare a concentrated solution of zinc chloride in distilled water to
serve as the electrolyte. Arrange the tube in series with the lamp-
bank and electric-lighting circuit as described in Chapter I, and
throw in one or two lamps. The zinc will immediately begin to
grow in the form of a beautiful metallic tree and chlorine may be de-
tected issuing from the bent glass tube over the opposite pole. There
is optical evidence that we have an accumulation
of metal at the cathode, and evidence of an
equally striking nature that we have chlorine at
the anode. In the drawing, the bleaching action
is shown upon a piece of calico. What is taking
place at the bend in the bottom of the tube,
however ? Have we mechanical movement there ?
We have been led to believe that we have
chlorine ions moving in one direction and zinc
ions moving in the other. Is it true, and is there
any experiment to prove it? The author has
showing the rapid
growth of a "zinc-
tree" and the libera-
tion of chlorine as
indicated by its
bleaching action.
Mechanical Trans-
port of Matter.
FIG. 43. Electrol-
designed three to furnish evidence in support ys i s of a Solution of
of this, and they are here published for the Chloride of Zinc
first time in book form. The first of the three
experiments is illustrated in Fig. 44. Although
glass is always stated to be the best elec-
trical insulator known, it occurred to the
writer that its constituents could carry the
electric current if the ions were only free to migrate. In other
words, glass was regarded by the writer as being a solid elec-
trolyte. Ordinary glass, as is well known to the general chemist*
consists of silicon dioxide fused with calcium and sodium carbonate.
Ordinary glass is therefore a soda lime, silicon-dioxide glass. Bohe-
mian glass is made with potassium carbonate, and flint glass is made
by melting together lead oxide, potassium carbonate, and silicon
dioxide, while strass is a species of glass very rich in lead. We have,
as may be seen, all the ions necessary to carry electric currents if they
were only free to move about. To test this point little rods of solid
glass of the different varieties were softened and platinum wires
run into the ends to a distance of about a centimeter. The rods
were cut a"bout 3 centimeters long, so there was in each case an "insu-
lating gap" ctf about a centimeter between the ends of the platinum
94
EXPERIMENTAL ELECTROCHEMISTRY.
wires. A glance at the figure shows such a glass rod with the plati-
num wires at B. A delicate milliammeter is joined in series with the
platinum wires and our lamp-bank (not shown in the illustration),
and of course no deflection is shown upon the indicating instrument
because glass is an insulator when cold. Now bring a Bunsen
burner, A, under the glass rod and heat it up gradually. As the
glass begins to soften and flow, the milliammeter begins to show con-
ductivity. With the softening of the glass, therefore, the ions are free
to travel. This is a very pretty and convincing experiment when
performed with a large-scale ammeter so that in a lecture-room an
entire class may see the deflection of the needle. It is of interest
FIG. 44. Simple and Easily Performed Experiment to Show Mechanical Transfer
of Matter through Solid Glass. A, Bunsen burner; B, solid rod of glass; C,
milliammeter to show conduction of the electric current.
at this time to point out the fact that all electrolytes when in solution
in water conduct better when heated up. This is just the reverse
with metals and alloys.
EXPERIMENTS WITH FROZEN ELECTROLYTES.
What will be the effect upon the movement of ions on freezing an
electrolyte? This question presented itself to the writer, and not
coming across records of any experiments with frozen electrolytes,
or any theoretical discussion of the same, it was decided to investigate
the matter experimentally. According to our ionic migration, or
convection theory of electricity through substances in solution, the
conductivity of an electrolyte should cease or approach zero value
FARADAY'S LAW.
95
when frozen. When the medium containing the ions is frozen
solid will the conductivity actually cease ? Will the ions move freely
through as before, or can they be forced through at diminished
velocity? The writer has made numerous experiments upon the
physical properties of ice, and has found it to be viscous-like in
behavior. This property has also been demonstrated by Profes-
sors Tyndall and Agassiz upon a tremendous scale in their studies
of the Swiss glaciers. Ice can be bent, twisted, and pressed
into molds, and be made to flow under pressure like semi-molten
glass. The ions in a frozen electrolyte, therefore, should be in a
measure free to move slowly when subjected to an electric current.
Let us look into the facts of actual experiments. Fig, 45 shows a
FIG. 45. Experiment with Frozen Electrolytes. The"V" tube at the left of the
drawing contains the solution to be frozen and is immersed in a beaker contain-
ing the freezing-mixture. Indicating galvanometer in the center, storage-battery
at right. Three cells of battery at least are necessary in this experiment or else
the lamp-bank shown in Chapter I.
simple arrangement for experimenting with frozen electrolytes.
The following experimental work embraced the freezing of a' dozen
or more solutions, including NaCl, KC1, HC1, H 2 SO4, K 2 Cr 2 O7,
KI, etc., and in every instance, with an ordinary ammeter in series,
the current appeared to be completely stopped when the solvent con-
taining the ions was solidified. The temperature in each case was
well below the freezing-point, the solutions frozen being very dilute
not only for the sake of complete dissociation of the salts, but for the
comparative ease with which they could be frozen. The freezing-
mixtures consisted of pounded ice and salt and also the well-known
96 EXPERIMENTAL ELECTROCHEMISTRY.
mixture of sodium sulphate 6 parts, ammonium nitrate 5 parts,
dilute nitric acid 5 parts. The lowest temperature attained by
the use of this mixture was 30 C., which was more than ample to
freeze any of the above dilute solutions. The plan of conducting
an experiment which is suitable for a lecture-room experiment is
shown in the illustration and consists of a V tube of glass fitted
with stoppers and platinum electrodes, which may be immersed in
the beaker containing the freezing-mixture. A large upright galva-
nometer is shown in the center and a cell of storage battery at the
right. One cell of battery is rarely sufficient for electrolysis, as
will be learned in the next chapter when we discuss the energy
required for the electrolysis of various compounds, and at least
three should be used here. Our lamp-bank is well suited for
this experiment. As the electrolyte freezes the pointer of the
ammeter or galvanometer comes to zero and upon thawing again
travels over the scale. Thus we get no deflection with an ordi-
nary ammeter or galvanometer upon freezing a solvent con-
taining ions, even when the lamp-bank is used, and we have a differ-
ence of potential of no volts between the terminals of the system.
One might be misled in stating that we had a non-conductor just as
one is accustomed to speak of glass. On going a step further, however,
and substituting an exceedingly delicate galvanometer or milliammeter
a good deflection through the frozen electrolyte is at once obtained.
Of course it is necessary to predetermine the freezing-points of the
solutions to be experimented upon and to be sure the temperature
is well below that required to insure the absence of any liquid elec-
trolyte. The slight conductivity can only be explained on account
of the viscous-like behavior of ice. If ice became a perfect solid it
would undoubtedly become a perfect non-conductor. The con-
ductivities of crystals of copper sulphate, iron sulphate, etc., were
tried and proved to be non-conductors with the most delicate instru-
ments, unless fused, when they of course conducted as usual. A
crystal is a true solid. It proves to be a true non-conductor. The
passage of an electric current through chemical bodies, therefore,
must be accompanied by the mechanical movement of matter.
The third experiment to furnish evidence in support of ion transport
is dependent upon heat convection as well as electrical convection.
FARADAY'S LAW.
97
HEAT CONVECTION IN ELECTROLYTIC CONDUCTION.
The apparatus illustrated in Fig. 46 was designed for the purpose
of indicating the movement of ponderable material of which ions
are composed, through the agency of the heat which they may be
made to carry. It was reasoned that if we actually had ponderable
material moving through an electrolyte, this material could be made
to convey heat as well as electricity, and by the proper design of an
apparatus the method could be employed for determining the absolute
FIG. 46. Apparatus to Show Heat Convection in Electrolytic Conduction. A and B
are calorimeters made from muff boxes lined with heavy hair-felt. Beckmann
thermometers dip into "T" connections in the glass tube containing the elec-
trolyte.
velocities of ions, although in a rough way, because of sources of
error difficult to overcome. At the left in the illustration is the anode
calorimeter, A, containing a T joint receiving the bulb of a delicate
thermometer. About this T joint is wound a little coil of platinum
wire for the purpose of heating the contained electrolyte. This
coil is attached to stout copper wires terminating in binding-sleeves
outside the calorimeter, which is a pasteboard muff, box lined with
hair-felt. In the cathode calorimeter B is a similar T tube and
thermometer, but without the heating-coil. The electrolyte in one
experiment consisted of dilute sulphuric acid in the proportion of
20 cubic centimeters of concentrated sulphuric acid to 100 cubic
centimeters of distilled water. The thermometers employed read
98 EXPERIMENTAL ELECTROCHEMISTRY.
from o to 100 C. in 0.1 divisions. The platinum coil was con-
nected to the 1 10- volt lighting wires through a finely graduated
variable resistance, allowing of a very close control of current and
heating effect. The following is an account of an experiment con-
ducted with the electrolyte referred to.
The anode calorimeter was placed upon a higher level than the
cathode calorimeter for the purpose of preventing simple convection
currents due to the expansion of the water molecules within the glass
T surrounded by the heating coil traveling by displacement. The
experiment was started by allowing o.i ampere to flow through
the apparatus, with a potential gradient of i volt per centimeter.
This current was allowed to flow for three hours, when the mercury
in the two thermometers appeared to reach a maximum reading.
It should be stated, however, that in a special experiment con-
ducted purely to note a special phenomenon the thermometers
continued to show a rise in temperature, although at an ex-
ceedingly low rate, even after three hours, it being found that
nine and one-half hours were really required before a perfect
balance between the heat-energy supply and the heat-energy radia-
tion loss through the calorimeters was reached. The rate of rise
of temperature as indicated by the thermometers at the end of
three hours, however, was so low as to allow of an experimental
determination of the velocity of ionic travel. To determine the
velocity of hydrogen ions, therefore, it is only necessary to replace
the ordinary thermometer in the cathode calorimeter by an open-scale
Beckmann thermometer of the most sensitive type, carefully adjusted
to the temperature of the electrolyte into which it is to dip, which of
course may be done from the reading on the ordinary thermometer.
The apparatus is then allowed to stand for an additional hour for
the new thermometer to reach a perfect equilibrium, when the cur-
rent is admitted to the platinum heating-coil, the time being noted at
the moment of closing the circuit. In another separate and special
experiment the electrolyte in the T tube in the anode calorimeter
A was kept at 90 C. for three hours without affecting the thermom-
eter in the cathode T tube but 0.2 C., and this was due to conduction
through the glass, and possibly a little convection in spite of the
more elevated anode calorimeter. Water conducts itself, to a slight
extent, so this error must be learned and applied as a correction. It
FARADAY'S LAW. 99
will be seen that this glass tube is inclosed within a second glass tube
with an air-space around it, and the whole thing is in turn inclosed
in a jacket of hair- felt. With the experiment under way as described
it only remained to watch for the first indication of a marked rise
upon the Beckmann thermometer in the cathode calorimeter. The
distance between the two thermometer-bulbs was 80 centimeters.
The hydrogen ions apparently arrived in one experiment twenty
minutes late according to calculations as to when they were due.
This tardiness may be attributed to the lack of sensitiveness of the
thermometer. Undoubtedly the ions arrived on time, if the phe-
nomenon is really due to the heat which they carry, but it remained
for an accumulative action to take place before there was heat
enough to affect the large mass of mercury in the thermometer. A
thermopile and galvanometer would have acted quicker. So much
for evidence of ion migration and mechanical movement. We are
now in a position to take up electrochemical work of a more practical
character and will open the next chapter with such introduction.
CHAPTER VII.
ENERGY REQUIRED IN ELECTROLYSIS. PRACTICAL FORMULA
FOR COMPUTING. ELECTROLYTIC SEPARATION OF MET-
ALS. THE ROTATING ANODE IN ELECTRO-ANALYSIS.
IN the last chapter it was pointed out that whereas a given cur-
rent flowing for a given time would separate chemical equivalents
of electrolytes, the energy absorbed in electrolytes of different con-
stitutions was not the same. Now, why is it that with some elec-
trolytes more energy is required to isolate the constituent parts
than with others? We have learned that 96,540 coulombs will
isolate the chemical equivalent of any electrolyte, and we now learn
that these 96,540 coulombs must be supplied at different electrical
pressures for different electrolytes. The 96,540 coulombs with-
out an electromotive force would not flow, and it is evident that
we must have some electromotive force in order to have electrical
energy, for the joule which is the unit of electrical energy is the
product of the coulombs by the volts in an electrical circuit. We
may theoretically have any amount of electrical energy we may wish
by multiplying our 96,540 coulombs by volts or fractions of volts.
To determine why some electrolytes require more energy (a higher
electrical pressure with the 96,540 coulombs) than others, we need
but to refer to the "heats of formation" of different electrolytes,
and the doctrine of the "conservation and correlation of energy."
By "heat of formation" of a chemical compound, we mean the
number of calories liberated (and sometimes absorbed) when one
gram-molecule of the substance is produced. To find the heat
of formation of a chemical compound, a gram-molecule of the
.substance is taken and its combustion in oxygen determined. Ac-
cording to the principle discovered by Hess, if we know the heat
of combustion of a gram-molecule of a compound, we may de-
termine its heat of formation by subtracting the heat of combus-
100
ENERGY REQUIRED IN ELECTROLYSIS. ioi
tion of the compound from the heats of combustion of the constitu-
ent elements. For example, the heat of formation of methane,
CH 4 , is determined by measuring the heat of combustion of the
compound in oxygen in a suitable calorimeter, and the heats of
combustion of its elements in the same manner, and subtracting
one from the other as follows :
Heat of combustion of methane, CH 4 , =211,930 calories, yielding
CO 2 and 2 H 2 O.
Heat of combustion of carbon, 0=96,960 calories, yielding CO 2 .
Heat of combustion of hydrogen, H 4 = 136,720 calories, yielding
2 H 2 0.
The heat of formation of a gram-molecule of methane is
found by subtracting 211,930 calories (its heat of combustion)
from 96,960 + 136,720 calories (the heats of combustion of its con-
stituent parts) as given:
96,960
136,720
233,680
211,930
21,750 calories = heat of formation of CH 4 .
Therefore when the gram-molecule of CH 4 is produced a
definite amount of energy is liberated ; and according to the doctrine
of the conservation of energy, this same quantity of energy must
be absorbed again before the compound can be broken up into
its constituent parts. CH 4 is not an electrolyte; the compound
was taken to serve as an illustration. Below we have tabulated
the heats of formation of a number of chemical compounds taken
from the general tables in Ostwald's "Outlines of General Chem-
istry." These have been converted into small calories.
There are of course many others given in a comprehensive table,
including both organic and inorganic compounds, electrolytes and
non-electrolytes, but it is believed that the list given includes a
sufficient number of electrolytes to be of service in the laboratory.
With such a table of heats of formation, and the doctrine of the
conservation and correlation of energy, taken in connection with
IO2
EXPERIMENTAL ELECTROCHEMISTRY.
Compounds.
Formulas.
Calories.
Compounds.
Formulas.
Calories.
Hydrochloric acid. . . .
HC1
22,000
Ferric chloride
FeCl 3
96,100
Water
H 2 O
68,400
Ferrous sulphate
FeSO 4
21S. 600
H 2 SO 4
Nickel chloride
NiCl 2
NH 3
Nickel sulphate
NiSO
HNO 3
49,100
Zinc oxide
ZnO
85 8OO
Potassium hydroxide .
KOH
103,200
Zinc chloride
ZnCl 2
Potassium chloride. . .
KC1
104,300
Zinc bromide
ZnBr 2
Potassium bromide . .
KBr
QC.IOO
Zinc iodide
ZnI 2
Potassium iodide. . . .
Potassium nitrate. . .
Sodium hydroxide. . .
KI
KNO 3
KOH
NaCl
80, 100
119,500
101,900
Q7.QOO
Cadmium chloride . .
Cadmium bromide. .
Cadmium iodide. .. .
Cuprous chloride
CdCl 2
CdBr 2
CdI 2
CuCl 2
93,200
75,200
48,800
NaBr
85 800
Cuprous bromide
CuBr 2
Nal
CuI 2
Na2SO 4
328 800
CuSO
Sodium hydrogen sul-
Cupric nitrate
Cu(NO 3 ) 2
82, 300
NaHSO 4
267,800
IVTercurous chloride
Hg 2 Cl 2
62 600
Ammonium chloride.
Ammonium bromide.
NH 4 C1
NH 4 Br
75,800
6^,400
Mercuric chloride. . .
Silver nitrate
HgCl 2
AgNO 3
53.200
28,700
Ammonium iodide . . .
NH 4 I
40, 300
Lead chloride
PbCl 2
82 800
Calcium hvdroxide. . .
Ca(OH) 2
214,900
Lead bromide
PbBr 2
6j. coo
CaO
133 ooo
Lead iodide
PbI 2
20 800
Calcium chloride. . . .
CaCl 2
169,800
Lead sulphate
PbSO 4
216 200
Calcium bromide. . . .
CaBr 2
140,900
Lead nitrate
Pb(NO 3 ) 2
TO? Coo
CaI 2
IO7 1OO
Stannous chloride
SnCl 2
80 800
Magnesium chloride .
Magnesium oxide. . . .
MgCl
MgO
217,300
143,900
Stannic chloride. . . .
Auric chloride
SnCl 4
AuCl 3
127,300
22,800
Magnesium hydroxide
Aluminium hydroxide
Mg(OH) 2
Al(OH),
AlCls
217,300
297,000
I C,I,OOO
Aurous chloride. . . .
Aurous bromide. . . .
Aurous iodide
AuCl
AuBr
Aul
5,8oo
IOO
c coo
Aluminium bromide. .
Aluminium iodide . . .
Ferrous chloride
AlBr 3
All,
FeCl 2
119,700
70,400
82,100
Chloroplatinic acid. .
Bromoplatinic acid. .
H 2 PtCl 6
H 2 PtBr 6
163,200
88,400
our important constant of 96,540 coulombs, we should be able
to calculate the minimum voltage or electromotive force necessary,
and consequently the energy required, to break up any chemical
compound by electrolysis. It is simply converting heat energy
into electrical energy. We must of course know the relation between
the calorie and the joule, and should commit this to memory as a
very important figure :
i joule =0.00024 Calorie
i joule =0.2 4000 calorie
Let us take one or two examples and work them out, develop-
ing a practical working formula for future use. We have already
learned that electrolytes may be either chemical compounds in
solution or in a state of igneous fusion. As we have dealt with
many electrolytes dissolved in water, let us consider an electrolyte
ENERGY REQUIRED IN ELECTROLYSIS. 103
in a state of fusion by heat. Let us first take an electrolyte con-
sisting of monovalent constituents, and consider it theoretically.
Common salt, or sodium chloride, will serve our purpose, with its
monovalent sodium linked to the monovalent chlorine. In cal-
culating the minimum voltage necessary to isolate a chemical equiva-
lent of sodium and chlorine, and the energy necessary to effect the
electrolysis, we must consider the gram-molecule as the basis of
our calculation, which in the case of sodium chloride is 58.5, since
sodium has an atomic weight of 23, and chlorine 35.5. There-
fore 23+35.5=58.5. 58.5 grams of sodium chloride is the
gram-molecule of the salt. Glancing at the table of "heats of
formation" for the value found for sodium chloride, we see it to be
97,900 calories. When 23 grams of sodium combine with 35.5
grams of chlorine, therefore, to form 58.5 grams of sodium chloride,
97,900 units of heat are set free. According to the doctrine of the
conservation of energy, in order to decompose these 58.5 grams
of sodium chloride, an amount of energy equal to that liberated
at the time of formation must be expended upon it to break it
up. How much electrical energy, for instance, is the equivalent
of 97,900 heat units? The heat unit is the calorie, and we have just
learned that 0.24 Calorie is equal to the joule. 97,900 calories
divided by 0.24 therefore gives us the joules necessary to effect the
electrolysis.
97,900
0.24
=407,916 joules.
We have learned that the passage of 96,540 coulombs will set
free the chemical equivalent of any electrolyte, and we see from the
above exposition that 407,916 units of electrical energy are neces-
sary. In order that the passage of 96,540 coulombs should repre-
sent the expenditure of 407,916 joules, they must be supplied at an
electrical voltage or potential of
=4.22 volts.
96,540
To liberate 23 grams of sodium and 35.5 grams of chlorine,
therefore, from 58.5 grams of sodium chloride, we will have to
expend 407,916 joules of electrical energy, and it will be necessary
104 EXPERIMENTAL ELECTROCHEMISTRY.
to have a voltage of at least 4.22 volts before an electrical current
can be made to pass through. Let us take the case of a compound
with a divalent constituent, for example magnesium chloride (MgC^),
from which to theoretically and practically, if we may so speak, sep-
arate the chlorine from the metal. By referring again to our table
of formation heats, we find for the gram-molecule of magnesium
chloride that 217,300 calories are liberated. We have therefore:
217,300
=005,415 joules.
0.2400
Now we are dealing with a divalent electrolyte, and according
to Faraday's law it will require the passage of 96,540 + 96,540 cou-
lombs to separate the constituents. This gives us:
96,540
96,540
193,080 coulombs,
which must be divided into 905,415 joules in order to obtain the
minimum voltage
95>4i5
* =4.68 volts.
193,080
It is evident that these minimum voltages are dependent upon
the degree of accuracy attained in the determination of the heat of
formation of the compound in the calorimeter, which, of course,
is purely, a thermochemical operation. It must be stated, however,
as a matter of fact in an actual experiment with a fused electrolyte,
that the calculated voltages are a little high, which is accounted
for by the high temperatures of the fused compound. At the tem-
peratures of igneous fusion the heats of formation have a lower
value, apart from the fact that compounds are dissociated in the
fused state. When in a state of igneous fusion, therefore, a voltage
of something less than the calculated pressure will drive the elec-
trical current through. Let us perform the experiment of electri-
cally separating the metal magnesium from the chloride, and measure
the electrical energy required. Fig. 47 shows the equipment which
we can use to good advantage, both as a laboratory method of pro-
ducing the metal, and also as a brilliant lecture experiment, where
a large mass of the metal is produced and hammered out and ignited
ENERGY REQUIRED IN ELECTROLYSIS.
105
to produce the dazzling magnesium light. At the left in this illus-
tration we have the motor-generator, with the controlling rheostats
FIG. 47. Experimental Equipment for Lecture-room or Laboratory for Electrically
Isolating Metallic Magnesium from its Chloride. Motor-generator with rheo-
stats at left. Copper voltameter in center, and small fusion-furnace at right.
for supplying a heavy current at low electrical pressure. The
furnace is of the small gas-fusion type, which contains an iron pot for
the electrolyte, and which, as can be seen, is made the cathode.
The anode consists of a large rod of carbon passing through a tight-
fitting cover. A large copper voltameter is depicted in the center
for determining the number of coulombs passed, and electrical in-
struments are connected to the leads to indicate the voltage and
current strength. For our experimental purposes an artificial
carnalite will best serve our purpose for the electrolyte. The com-
position of this mineral may be set down as follows:
KMgCl 3 , 6H 2 O = KCJ + MgCl 2 + 6H 2 O = MgCl 2 34. 2,
KC1 26.9, H 2 O 38.9 = 100.
For our purpose we can prepare an easily-fused artificial carnalite
by evaporating to dryness on a water-bath a solution of
400 grams of crystallized magnesium chloride,
150 grams of potassium chloride,
60 grams of ammonium chloride.
106 EXPERIMENTAL ELECTROCHEMISTRY.
This residue may be placed in a large salt-mouthed bottle for
use as required. For an experimental run the iron vessel in which
the electrolysis is to be carried on is carefully cleaned on the inside
by means of sandpaper, to remove any rust and to produce a bright
metallic surface from which the magnesium is easily separated.
This iron pot is so connected as to form the cathode, and is placed
in the furnace and the gas lighted. A small piece of charcoal is
put in the pot to prevent oxidation as it heats up. When a faint
red glow is seen in the bottom of the iron pot, the artificial carnalite
is slowly added, allowing each addition to fuse and run before the
next portion is added. The pot may in this way be nearly filled,
when the carbon anode is placed in position through the top. Now,
as magnesium is a very light metal and often rises to the top of
the electrolyte after being isolated, when it takes fire and is de-
stroyed, it is necessary to protect it. For this purpose a tight-fitting
top is the best precaution, fitted with asbestos plugs through which
hard-glass tubes pass as indicated in the drawing. By means of
these tubes a gas can be passed which displaces the oxygen present,
and prevents the combustion of the magnesium. The present
writer has connected these glass intake-tubes direct to the city light-
ing gas-supply with much success, allowing the coal-gas and chlorine
to be led away to a hood. In this way we have the magnesium
completely protected from possibilities of combustion, since coal-gas
being a mixture of hydrocarbons, etc., does not support combus-
tion. For an experiment, therefore, the top is placed in position
after the chlorides have fused to a clear, transparent liquid, and
coal-gas is passed through the space over the electrolyte when
the electrical current is started. The magnesium separates
smoothly, and may be ladled out and cast into molds, care being
taken to allow the electrolyte to cool down considerably before
exposing it to the air, as it is very likely at high temperature to take
fire and burn with its characteristic dazzling light. Of course, for
a full efficiency determination of such a process an accurate gas-
meter, with the necessary observations for gas temperature and
atmospheric pressure, should go in with the burner. For a com-
plete physical and electrical study of the isolation of magnesium,
the heat value of the fuel-gas per cubic foot should be determined
by means of a suitable calorimeter. There are several calorimeters
ENERGY REQUIRED IN ELECTROLYSIS. 107
especially designed for determining the heating value of gaseous
fuels. With our copper voltameter and reliable voltmeter we will
be in possession of the essential data for interesting figures taken
in connection with the weight of metal produced. This experi-
ment is capable of going still further, using the liberated chlorine,
for example, to prepare chloride of lime. For this purpose we
should prepare a lead box with lead shelves, upon which we can
lay out a quantity of moist slaked lime. The temperature is not
allowed to rise above 25, which is controlled by diluting the chlorine
passing into the chamber with air. The constitution of chloride
of lime is not known with certainty, but the action between the
chlorine and the moist slaked lime may be represented as follows:
Ca(OH) 2 + Cl 2 = CaCl(OCl) +H 2 O.
Thus a very pretty by-product may be obtained at the time of iso-
lating the metallic magnesium.
Let us take a case of a non-igneous electrolyte, and calculate
the minimum voltage required. We are almost unlimited in a
selection of these cases, and for this reason a novel case of electrol-
ysis as conducted by the author, to show that sulphuric acid con-
sists of hydrogen, oxygen, and sulphur, may be of special interest.
This is an interesting lecture experiment for both beginners in general
chemistry and for those advanced in physical chemistry as well.
We all know that dilute sulphuric acid electrolyzes into hydrogen
and oxygen gases, and that these gases are liberated in the propor-
tion of two volumes of hydrogen to one volume of oxygen. The
writer has conducted numerous experiments with concentrated
sulphuric acid, and by the proper adjustment of concentration,
current density, temperature, etc., has been able to electrolyze the
acid into hydrogen, oxygen, ozone, and free sulphur, and at ele-
vated temperatures into hydrogen, oxygen, sulphur trioxide, and
ozone. Fig. 48 illustrates the apparatus for conducting such elec-
trolysis. At the left we have a bell jar covering the beaker con-
taining the strongest chemically-pure sulphuric acid over a dehy-
drating agent, such as calcium chloride. A thermometer is fitted
through the stopper as shown, and the electrical equipment includes
electrical instruments for observing the energy conditions. The
io8
EXPERIMENTAL ELECTROCHEMISTRY.
motor-generator is shown at the extreme right with the controlling
rheostats, although the lamp-bank serves equally well, if not better,
FIG. 48. Apparatus for the Electrolysis of Concentrated Sulphuric Acid to obtain as
Electrode Products Hydrogen, Oxygen, Ozone, and Free Sulphur which may be
Exhibited upon the Anode and be Burned. Therefore, sulphuric acid may be
directly broken down into its elements.
for this particular experiment. Upon passing a sufficiently heavy
current through the concentrated acid we get free sulphur, together
with ozone, oxygen, and hydrogen, as indicated in the accom-
panying equation:
The sulphur deposits upon the anode and may be burned with its
characteristic blue flame before a class to show its presence. The
ozone may be detected by moistening a piece of filter-paper in starch
and potassium iodide preparation, and holding it near the anode
during electrolysis. The hydrogen may be collected and burned.
The minimum voltage for conducting such an experiment may
be calculated by means of our formula to a fair degree of precision.
As we are now in a position to follow theoretically as well as practi-
cally an electrochemical process, tha writer introduces what he
believes to be the "mechanism" of this particular electrolysis. In
all our practical work we must endeavor to account theoretically
for the phenomena involved in practice, for it is the man with the
insight into both theory and practice who makes the best invest-
ENERGY REQUIRED IN ELECTROLYSIS.
109
tigator. Having calculated the minimum voltage and energy re-
quired, and conducted an actual electrolysis, we should certainly
endeavor to express that which takes place in concise and scientific
manner. In the almost classic case of sulphuric acid and water,
the acid is believed to dissociate into the ions
+ +
H2 and SC>4.
The SC>4, instead of being set free, decomposes the water present,
as indicated as follows, taking up the two atoms of hydrogen present,
to form sulphuric acid and liberating oxygen. As the result of certain
research work upon the concentrated acid, it is believed to dissociate
+
into the ions H and HSC>4.
Cathode
Jbwde
Now, according to recent research upon the conductivity of
pure water, water itself is found to be slightly dissociated (about
one molecule in a million being broken down thus: H OH) and
taking this fact into account, we may have upon this basis the libera-
tion of two volumes of hydrogen and one volume of oxygen equally
well accounted for in diagram.
;*
ESO
~GE
Now let us write the structural formula of sulphuric acid, and
endeavor to represent the "mechanism" of electrolysis when we
obtain hydrogen, oxygen, ozone, and free sulphur.
H-O- =O.
no
EXPERIMENTAL ELECTROCHEMISTRY.
Here we have the sulphur acting with the va'ehce of six, and a
scheme indicating the linking of the acid. How can it electrolyze
and give such products? The following diagram shows the prob-
able manner taking dissociated water into account:
--
S
[Hl-0- =-.!di
II .,11
i I S I '
!-<L-?iC5?:.;>:p:
Here we also have the formation of a molecule of water within the
electrolyte. If, however, the temperature is allowed to rise above
105 C. we obtain hydrogen, oxygen, and sulphur trioxide, with
the formation of a molecule of water within the electrolyte as follows :
|H) - - r - r
i o !
J ]
= O
So much for this side of electrolysis and the part played by
minimum voltage, or electrode tension as a factor in experimental
as well as in commercial work.
ELECTROLYTIC SEPARATION OF METALS.
Under the present heading we will take up the art of electro-
analysis, and separate one metal from another, when both exist
in the same solution, through the proper adjustment of electrode
tension. Fig. 49 shows a plan for accomplishing this. Here we
have a platinum dish resting upon a platinum plate, connected with
the negative wire of a storage-battery and variable rheostat. A
disk of platinum welded to a platinum wire is suspended in the
dish to serve as anode, and is connected in turn with the battery.
A delicate ammeter and voltmeter complete the equipment, and
are joined up as shown. Only instruments of delicacy and pre-
cision are applicable for this class of work, and a rheostat capable
ELECTROLYTIC SEPARATION OF METALS
III
of fine graduations is absolutely essential. The cells of the battery
must be so arranged that one or more may be joined in series in a
FIG. 49. Method of Separating Metals by Electrodeposition through Adjustment
of Electrode Tension.
convenient manner. Having set up the apparatus, we are in a
position to undertake some experimental work. Fig. 50 illustrates
a practical working equipment for separating one metal from an-
other on the basis of electrode tension. A platinum dish and plati-
num strip acting as the cathode and anode respectively are shown
in the center. Readers of electrochemical literature will in this
connection come across the terms "polarization" and "polariza-
tion current," etc., which refer to the back electromotive force
or tension necessary in order to force a current through an electro-
lyte. Therefore, polarization may be under tood to refer to the
minimum voltage necessary to effect an electrolysis. Le Blanc
made many c reful researches upon the electrode tensions neces-
sary to decompose various salts, acids, and bases when in solution,
and it may be easily seen that we may separate one constituent
from another in an electrolyte by carefully adjusting the electrode
tension to fall between the two different values for the different
112
EXPERIMENTAL ELECTROCHEMISTRY.
CdSO 4
CdCl 2 .
CoSO 4 .
CoCl 2
Volts.
1.98
2.03
electrolytes. Le Blanc found the following values for normal solu-
tions :
Volts.
ZnS0 4 2.35
ZnBr 2 i . 80
NiSO 4 2 . 09
NiCl 2 " 1.85
Pb(N0 3 ) 2 1.52
AgNO 3 0.70
HNO 3 1.69
NaOH 1.69
NH 4 OH i . 74
HC1
KOH
HBr. .
0.94
This is of course only a partial list, but will be sufficient to serve
as a useful guide in the laboratory. Any attempt to outline methods
FIG. 50. The Separation of Metals by Adjustment of Electrode Tension. Rheo-
stats, storage-batteries, and electrical instruments are shown here in practical
operation in the laboratory.
of electro-analysis would be incomplete at this time without intro-
ducing the attractive and useful device known as the rotating anode.
Suppose, for example, it is wished to determine copper in the electro-
lytic way, that is by deposition upon a platinum dish. According
to the old scheme, we had the dish and a stationary anode. With
this arrangement it was necessary to work with a feeble current, or
else the deposit of copper would come down dark and non-adherent.
For this reason it took many hours, very often, to produce com-
plete precipitation. With the rotating anode, the current strength
can be enormoulsy increased, and yet obtain a beautiful pink ad-
herent deposit of copper in a correspondingly shorter time. Here
ELECTROLYTIC SEPARATION OF METALS. 113
we have an electric motor properly wound to run on a storage-battery
circuit, or else especially wound to run on the electric-lighting circuit
with lamps in parallel. Do not try to use a very small motor. Pro-
cure one several sizes larger than is really necessary to revolve the
anode, for it can be much more easily controlled through the agency
of our lamp-bank, or a special short bank as illustrated at the extreme
right in Fig. 51. Here the addition of one lamp will give the motor
FIG. 51. Easily-constructed Rotating Anode for Rapidly Conducting Electro-analy-
sis. With this equipment an electro-analysis may be completed in a few min-
utes, which would require several hours to accomplish in the old way.
a certain speed, which may be increased by the addition of others.
From thirty to one hundred and twenty revolutions per minute
have been found to be excellent speeds, although higher speeds
may be used with advantage so long as there is no danger of losing
electrolyte by its spinning out of the dish. A large disk just above
the platinum dish should be included in the equipment, to prevent
anything falling into the dish from the commutator and contact
brushes above. The rest of the make-up is so simple that the
illustration should serve to make it clear. A few words concerning
the principle upon which the rotating anode accomplishes its rapid
precipitation may be given here. The main thing accomplished
by the rotating anode is to keep the solution homogeneous in character.
Let us take the case of a solution of copper sulphate. If a com-
paratively feeble current of electricity be passed through an electro-
114 EXPERIMENTAL ELECTROCHEMISTRY.
lyte consisting of copper sulphate for a long time, there are concen-
tration changes set up. If we use copper electrodes the anode loses
in weight just as much as the cathode gains, and there is at all times
in the solution the same amount of copper, or to put it in other
words, there is always the same number of copper ions in solution.
Now, even with the use of copper electrodes and a constant number
of copper ions present in the solution, concentration changes will
be set up if the current is allowed to flow for any considerable length
of time. In such an experiment it is found with copper sulphate,
for example, that we get an increase of concentration at the anode
and a decrease of concentration at the cathode. Fig. 52 represents
the condition of affairs in a vertical glass tube containing copper
electrodes and an electrolyte of copper sulphate. The shaded
portion represents the concentration of the solution about the anode.
Now, if this took place, as it does with a platinum dish and station-
ary anode, the copper ions become so few, or in other words the
electrolyte becomes so poor in copper about the cathode dish, that
we are held down to the use of a feeble current or there will be
trouble in getting the adherent, pink deposit, so necessary for all
accurate determinations of copper by electro-deposition. How can
these concentration changes be overcome? The rotating anode
accomplishes this perfectly, and enables us to keep a constant supply
of copper ions about the cathode dish, and allows consequently of
a heavy current being employed. How can we explain the con-
centration changes ? Fig. 53 shows by diagram how such accumula-
tion of copper ions about the anode is accounted for according to
Hittorf. The changes in concentration calculated from one of
Hittorf's researches are shown in this diagram. Here the white
circles represent the anions and the black circles the cathions, and
the dotted horizontal line indicates merely the middle of the solu-
tion in the vertical containing vessel. The electrolyte is perfectly
homogeneous before the electric current is passed, as is seen by
an equal number of anions and cathions respectively on each side
of the line. In the figure we have nine upon each side. Now
allow the current to pass for a given time. We know that different
ions have different velocities, and consequently the Cu ions will
move in one direction at a different rate from the SC>4 ions moving
in the other direction. It is very often confusing to the student
ELECTROLYTIC SEPARATION OF METALS.
to understand how we can have such concentration changes due
to different velocities of the ions when we have the same equivalents
of ions liberated at anode and cathode respectively. No cathion
can separate at the cathode until an anion separates at the anode,
and for each and every ion which is liberated at one electrode there
must be a corresponding ion liberated at the other. There may,
FIG. 52. FIG. 53. FIG. 54.
FIG. 52. Experiment with Copper-sulphate Solution to show Concentration Changes
wrought by a Feeble Current Flowing for a Long Time. The electrodes are of
copper.
FIG. 53. Diagram Representing Hittorf's "Transport Numbers." From Hittorfs
researches the relative velocities of ions are determined experimentally, by con-
centration changes.
FIG. 54. Approved Apparatus of Mather and Jones for Experimentally determining
the Relative Velocities of Ions Based upon Concentration Changes.
of course, be the case where we have one divalent ion separating at
one pole and consequently two univalent ions separating at the
other. Faraday's law teaches us this. As a result of the different
velocities, while we have an anion liberated for every cathion, we
may have a banking up of the swifter kind of ions, and so to speak,
ready to discharge. This is illustrated in the little diagram at the
right, where we have thirteen anions across the dotted line and
II 6 EXPERIMENTAL ELECTROCHEMISTRY
only eleven cathions across the dividing line in the opposite direc-
tion. Such concentrations may be practically shown by experiment
with a piece of apparatus as illustrated in Fig. 54, as devised by
Mather, working with Jones. Here a copper-sulphate solution,
for example, may be electrolyzed, and the solution drawn off from
the two tubes respectively and analyzed for concentration. Know-
ing the original concentration of the homogeneous electrolyte, the
number of coulombs passed, and the respective concentrations of
the respective anode and cathode tubes, we have all the data for
determining the relative velocities of the two ions. This is a very
practical piece of apparatus, and with it very interesting concen-
tration changes may be studied.
CHAPTER VIII.
IMPORTANT CONDITIONS TO BE NOTED IN ELECTROCHEM-
ICAL OPERATIONS. CAUSTIC SODA AND CHLORINE FROM
SALT. ELECTROLYTIC PRODUCTION OF WHITE LEAD.
ELECTROLYTIC PRODUCTION OF CADMIUM YELLOW.
ELECTROLYTIC PRODUCTION OF MERCURY VERMILION.
ELECTROLYTIC PRODUCTION OF SCHEELE'S GREEN.
ELECTROLYTIC PRODUCTION OF BERLIN BLUE.
VARIOUS controlling conditions must be observed in all electro-
chemical operations, and be recorded in connection with every piece
of experimental work. There are many governing adjustments or
conditions in electrolysis, without a working knowledge of which the
student will be unable to meet with any notable success in carrying
out a determination, or be able to obtain the same result twice in
any undertaking. One of the most important factors in all electro-
chemical work is that of "current density," and because of its great
moment and importance it will be dealt with at the opening of this
chapter. Current density depends upon the ratio of electrode area
to the current flow in an electrolytic cell. We may have high-current
density at both electrodes, or low-current density at both electrodes,
or else high-current density at one of them and low-current density
at the other. Fig. 55 has been designed to make this clear. At the
left in this diagram the anode is simply a thin platinum wire affording
but small surface from which the electric current can leave to enter
the electrolyte, whereas the cathode is a platinum sheet affording
a large surface for the same current to be conducted from. At the
right in the same diagram the conditions of current density are
just reversed. Now the point of interest lies in the fact that adjust-
ments of current density have an important effect upon almost all
electrochemical operations, determining the character of the electrode
products, as well as the chemical change which may take place
"7
EXPERIMENTAL ELECTROCHEMISTRY.
throughout the entire electrolyte. Oxidation and reduction are two'
of the most important chemical phenomena, and yet the oxidation or
reduction of an electrolyte may be brought about by the same electric
current, the only difference in its application being that of current'
density. The following rule should be learned by all electrochemical
students :
Oxidation is effected by using concentrated electrolytes and by
a low-current density at the anode, as depicted in the right-hand
cell in the illustration. Reduction is effected by using concentrated
- 4-
1
1
-^
FIG. 55. FIG. 56.
FIG. 55. Diagram to Show Two Different Conditions of Current Density. At the
left there exists high-current density at the anode and low-current density at the
cathode. In the cell depicted at the right we have low-current density at the
anode and high-current density at the cathode.
FIG. 56. Four Cells in Series receiving a Common Current, but because of Dissimilar
Current Density Adjustments, Electrolytes in the Several Cells will Yield Differ-
ent Decomposition Products. The electrolysis in the two cells at the left will
be the same, but will differ from the products in the two cells at the right.
electrolytes and a low-current density at the cathode, as depicted
in the left-hand cell in the illustration.
It will, therefore, be fully appreciated how very important it is to
note and take fully into account the conditions of current density in
any piece of experimental work. Fig. 56 illustrates four cells in
series, the whole system, therefore, receiving a common electric
current. The conditions of current density, however, are not the
same, and we will obtain different results in the two cells at the right
from those in the two cells at the left. To test this we can perform
the following simple experiment, using a solution of oxalic acid for
the electrolytes, to which has been added a quantity of sulphuric acid.
Take 60 grams of oxalic acid to the liter of water, and add 50
grams of sulphuric acid, and place an equal portion of this solution
in each of the four cells. A current of about half an ampere is
allowed to flow for an hour, when the oxalic solution in each cell
is determined by means of permanganate of potassium. Oxidation
ELECTROCHEMICAL OPERATIONS.
119
will be found to have taken place in the two right-hand cells if the
current adjustment is as shown in the diagram, and is equal for
each cell. At the left there will be no notable increase, although
we do not get a correspondingly great reduction. It should be stated
that these current density conditions exert a strong tendency to
oxidize and reduce respectively, but of course all electrolytes are
not oxidizable or reducible any more than many compounds are
which go to make them up. We know that oxidation is usually
accompanied by liberation of heat, and it is therefore of great moment
to know both the anode and cathode temperatures in an electro-
chemical research. Fig. 57 illustrates the plan of the author for
FIG. 57. Author's Arrangement of Two Sensitive Beckmann Thermometers to
Study Anode and Cathode Temperatures, when Making a Research upon an
Electrolyte.
investigating such differential heat liberation. What are some of
.the other important conditions to be observed? They are many
and vital, and it is deemed that a concise tabulation of them, as
arranged by the author for use in the laboratory, will be perhaps a
120 EXPERIMENTAL ELECTROCHEMISTRY.
r good way of presenting them. In making any kind of a research
x upon a solution when subjected to the action of an electric current,
T the conditions tabulated here should be taken account of. If a piece
of experimental work is to be undertaken, a neatly-kept notebook
should of course be opened, and a careful record kept of each thing
observed, together with all the existing conditions. It will be neces-
sary to make a number of repeated special runs to secure all the data
as advised in the accompanying table, as there are too many conditions
to be usefully observed and recorded during any one run. For
example, a special run may have to be made for differential tem-
peratures, another for specific gravity determinations, etc. In
several runs the following table may be compiled for reference.
This particular table was the result of the author's work upon
sulphuric acid.
Duration of run One hour
Compound electrolyzed H 2 SO. ;
Character of solution No solution
Sp. gr. before electrolysis i .84664
Sp. gr. after electrolysis i .84001
Quantity of compound taken 100 c.c.
Character of apparatus See illustration
Dimensions of cell 7X8 cm.
Source of electricity Motor-generator
Temperature of electrolyte 21 . 5 C.
Temperature at anode 21 . 5 C.
Temperature at cathode 2 1 . 5 C.
Amperes flowing ... 4 . 250
Volts indicated 16.00
Area of anode immersed 4 sq. cm.
Area of cathode immersed 4 sq. cm.
Current density at anode ^ioo = 106.2
Current density at cathode C 100 = 106.2
Distance between 3 cm.
Phenomenon at anode SO 3 and O
Phenomenon at cathode Hydrogen
Phenomenon between Floating S in 3 mins.
Secondary action at anode None at once
Secondary action at cathode None at once
Secondary action between None
Later phenomenon at anode SO 3 at 103.5 C.
Later phenomenon at cathode None
Later phenomenon between Increased S
Material of anode Platinum
Material of cathode Platinum
Material of containing cell Glass
Special peculiarities
ELECTROCHEMICAL OPERATIONS. 1 21
Many operations, of course, will not require the setting down of
so many data, but for all research purposes the student will do
well to tabulate his facts as completely as possible. We are now
in a position to produce electrolytic preparations, and a few interest-
ing examples for laboratory practice are given here.
CAUSTIC SODA AND CHLORINE FROM COMMON SALT.
This is one of the first laboratory exercises the student in experi-
mental electrochemistry should take up in the way of preparations.
The experiment is a very practical and easily carried out introduction
to electrochemical manufacture. The apparatus as illustrated in
the photograph in Fig. 58 is easily and quickly put together in any
laboratory, and serves a most useful purpose in many cases of elec-
trolysis where the anode gas is to be collected. The apparatus
simply consists of a large beaker glass containing a good-sized porous
pot, about which a cylinder of nickel- wire gauze is placed to form
the cathode of the cell. A cylindrical lamp-chimney is next procured,
and fitted with a heavy rubber stopper, through which passes a rod
of carbon to serve as the anode. There is a second hole in this
stopper, to receive a small glass tube, through which the chlorine
escapes from the glass lamp-chimney chamber. The rubber stopper
should be given several coats of paraffin wax inside and out with a
good brush dipped into a melted mass. The electrolyte consists
simply of a saturated solution of common salt in water, and our
lamp-bank with two or three lamps in parallel in connection with
an electric-lighting system completes the equipment. Chlorine gas
escapes copiously from the anode chamber, and a corresponding
quantity of sodium hydroxide is formed in the cathode chamber.
Hydrogen of course escapes from the nickel-wire gauze when the
salt breaks up in accordance with the following equation:
2 NaCl + 2 H 2 O = 2 NaOH + H 2 + C1 2 .
The chlorine should be led into a lead-lined box containing lead
shelves holding moist calcium oxide, as outlined in the electrolysis
of magnesium chloride, for the simultaneous production of chloride
of lime, or the chlorine may be led into water to saturate it for
122 EXPERIMENTAL ELECTROCHEMISTRY.
bleaching purposes. The liquid from the cathode chamber is
poured off after the run, and evaporated to dryness in a porcelain
dish to obtain the solid caustic soda.
This experiment can and should be run
quantitatively by including a copper
voltameter in series with it, and noting
the fall in voltage between the electrodes,
in order that we may state the number of
joules absorbed per gram of sodium
hydroxide produced, and per gram of
bleaching-powder, etc. The porous-pot
partition in this apparatus plays a most
important part, as it keeps the chlorine
set free from acting upon the caustic
soda formed to produce another com-
pound, namely, sodium hypochlorite. As
the use of porous pots is of great impor-
FIG. 58. Large Beaker Ar- A . . .
ranged with Porous Pot and tance m a g reat manv electrolytic opera-
Glass Anode Chamber for Pro- tions as well as in research and investi-
ducing Caustic Soda and Chlo- gations, a group of the various desirable
rine from Common Salt. t i i i i , j
shapes and sizes has been photographed,
which constitutes Fig. 59. A good supply of these should be at
hand in every electrochemical laboratory.
ELECTROLYTIC PRODUCTION OF WHITE LEAD FROM METALLIC LEAD
ELECTRODES.
A very beautiful electrolytic preparation is that of white lead
from the metallic lead electrodes in an electrolytic cell. White lead,
or technically .the basic lead carbonate, has the following formula,
which is sometimes called hydrate-carbonate of lead:
2 PbCO 3 .Pb(OH) 2 .
For our purpose we will require either a rectangular glass jar or
cell, or else a large beaker glass, and heavy sheet-lead electrodes.
The adjustment for current density in this preparation is the same
for both electrodes, that is, they are of the same immersed area in
the electrolyte. A good working current density for this experiment
ELECTROCHEMICAL OPERATIONS. 123
is 0.5 ampere for every 100 square centimeters of anode and
cathode surface immersed. The expression for current density is
frequently met with in abbreviated ways, which the reader of electro-
FIG. 59. Porous Pots of Various Shapes and Sizes which should be Supplied to
every Electrochemical Laboratory.
chemical processes will come across, and they are therefore given
here in their usual manner, in order that he may become familiar
with them. For example, the following expression,
Da =Dc =0.5 ampere per 100 square centimeters,
means that the anode current density is the same as the cathode
current density, and that they each equal 0.5 ampere per 100 square
centimeters of immersed surface, measuring usually both sides of
the two electrodes in making the calculation of area. The expression
N.D.ioo=o.5 ampere
will also be met with, expressing the current flow from an electrode
for every 100 square centimeters of electrode area.
Let us now prepare a few grams of electrolytic white lead, and
tabulate the data in such concise form that we should at any time be
able to repeat the experiment with certainty, or be able to direct
others to do so. The electrolyte in this case will be 12 grams of
sodium chlorate and 3 grams of sodium carbonate dissolved in i liter
of water. A rectangular glass cell, two sheet-lead electrodes, and our
lamp-bank equipment in connection with a lighting circuit will
meet the electrical requirements, and it only remains to fit up a
generator for the production of carbon-dioxide gas from dilute
124
EXPERIMENTAL ELECTROCHEMISTRY.
acid and fragments of marble. The electrolysis is conducted at
about 20 C., and a slow current of carbon-dioxide gas is led into
the electrolyte in contact with the cathode, the electrolyte being
kept in motion by a stirrer. The white lead flows down in thick
clouds from the anode to the bottom of the cell, and may be col-
lected in a little bag of tow attached to the electrode, when it may
be removed and ground with oil to make the well-known basis for
oil colors. The yield of white lead in this experiment is excellent,
and the operation forms a very pleasing lecture preparation, for
the formation and falling down of the white lead from the solution
is very beautiful and pleasing, especially when one is familiar with
the unattractive old Dutch method, dependent upon the chemical
action of the vapors of acetic acid, carbonic acid, and oxygen upon
masses of lead in pots, which must be buried for long periods of
tmc in horse-manure, in order that fermentation may assist chemi-
cal action by an increase in temperature. Our electrolytic process
may be made continuous, and has attained commercial importance
in recent years. In this experiment the electrolyte contains two
salts in very dilute solution. The sodium chlorate, which is present
in four-fifths of the total amount, has as an anion C1O 3 which forms
a soluble salt with the anode lead, producing lead chlorate, which
passes into solution. The sodium carbonate, whose anion CO 3
FIG. 60. FIG. 61.
FIG. 60. Experimental Apparatus for the Electrolytic Production of White Lead
from Lead Electrodes.
FIG. 61. Experimental Apparatus for the Electrolytic Preparation of Cadmium Yel-
low from a Stick Cadmium Anode.
forms an insoluble salt, lead carbonate, produces the precipitation.
As a result of using such proportions, the insoluble salt does not
ELECTROCHEMICAL OPERATIONS. 125
deposit at once upon the anode, but is precipitated some distance
from it, and does not give trouble by forming an insoluble crust on
the electrode. The caustic soda produced at the cathode co-
bines with the carbon-dioxide gas which is bubbling through the
solution, and regenerates sodium carbonate. Fig. 60 will make
the apparatus clear. Here in the center of the cell is shown a stirrer
to be operated by a small electric motor when it is desired to show
the experiment in the lecture-room. Below will be found tabulated
the principal data in such an experimental run:
Duration of run i hour
Electrolyte 12 grams NaClO 3 ,
3 grams Na 2 CO 3
in i liter water
Character of apparatus See figure
Dimensions of cell 8 cm.Xso cm.X2o cm.
Source of electricity Lighting circuit and
lamp-bank
Temperature of electrolyte 20 C.
Amperes flowing 0.5 ampere
Volts indicated 60
Area of anode immersed 100 sq. cm.
Area of cathode immersed 100 sq. cm.
Current density at anode N.D. 100 = o. 5 ampere
Current density at cathode N.D. JOO = o.5 ampere
Distance between Approximately 18 cm.
Material of anode Soft sheet lead
Material of cathode Soft sheet lead
Phenomenon at anode White lead flowing
down in streams
Phenomenon at cathode Bubbling of carbon di-
oxide gas and the
formation of Na2CO 3
Another plan for producing a carbonate of lead can be shown
with this same piece of apparatus, but with a different electrolyte.
A solution of sodium nitrate is used, which when electrolyzed forms
nitric acid, which attacks the lead electrode and puts it into solution
as lead nitrate. The following equations show how this scheme
probably works :
1. NaNO 3 + H 2 O=NaOH + HNO 3 .
2. 2 HNO 3 + Pb = Pb(NO 3 ) 2 +H 2 .
3. Pb(NO 3 ) 2 + 2 NaOH=Pb(OH) 2 + 2NaNO 3 .
4. Pb(OH) 2 + HNaCO 3 = PbCO 3 +NaOH+H 2 O.
126 EXPERIMENTAL ELECTROCHEMISTRY.
Instead of reactions i and 2 taking place as shown, the following
may be the true state of affairs, since hydrogen is liberated at the
cathode :
2NaNO 3 + 2H 2 O + Pb = ( 2 NaOH + H 2 ) + PbO 2 + (NO 3 ) 2 .
Equation 4 results from the addition to the lead hydroxide of
a solution of sodium bicarbonate. Other modifications of this
very beautiful method will doubtless suggest themselves to the in-
genious student, apart from the interesting quantitative figures
he is in a position to obtain by working with such an equipment
and electrical measuring instruments. Of course, for economy in
operation the motor-generator should be used, as we do not require
the electric current at anything like no volts pressure.
Having produced the white lead, which is the basis for most
oil-color paint, we can next try our hand at the electrolytic pro-
duction of pigments. Perhaps the easiest and most satisfactory
pigment to take up first is that of cadmium yellow.
THE ELECTROLYTIC PRODUCTION OF CADMIUM YELLOW.
This very brilliant and beautiful pigment may be easily pro-
duced electrolytically in a cell similar to that employed in the prepa-
ration of white lead. This cell is shown in Fig. 61, a cylindrical
stick or rod of cadmium acting as the anode, and a strip of platinum
acting as the cathode. In the place of the CO 2 generator as used
in the previous preparation, a hydrogen-sulphide generator is em-
ployed. For this purpose, as is well known by every chemist, we
require some fragments of iron sulphide and a little dilute hydro-
chloric or sulphuric acid. This generator is depicted in its simplest
form, and may be replaced to advantage by one of the approved
"automatic" types, whereby the supply of gas controls the action
of the acid upon the iron sulphide. The electrolyte in this experi-
ment consists of a saturated solution of common salt in water, and
when electrolyzed under these conditions forms cadmium chloride
at the anode, and sodium hydroxide at the platinum cathode. For
the production of any quantity of this pigment, both the anode
and cathode should be placed in porous pots to prevent the mixing
ELECTROCHEMICAL OPERATIONS. 127
together of the respective electrode products. The cadmium chloride
produced is immediately precipitated as the brilliant yellow cad-
mium sulphide by the stream of hydrogen-sulphide gas. The
following simple equations indicate the steps in the production of
the pigment:
2 NaCl + 2H 2 O - 2 NaOH + C1 2 ,
2 Cd + C1 2 + 2 H 2 S = 2 CdS + 2 HC1.
If the electrolyte is kept stirred by a mechanical device, the
effect is very beautiful indeed. The tabulation of the data in the
electrolytic preparation of cadmium yellow is given below:
Duration of run i hour
Electrolyte Saturated solution of NaCl
in water
Character of apparatus See figure
Dimensions of cell 8 cm.X 20 cm.Xzo cm.
Source of electricity Lighting circuit and lamp-
bank, except for econ-
omy and quantitative
work. Then use motor-
generator.
Temperature of electrolyte 30 C.
Amperes flowing i . oo ampere
Volts indicated 4.5
Area of anode immersed 50 sq. cm.
Area of cathode immersed 100 sq. cm.
Current density at anode N.D. 100 = 2 amperes
Current density at cathode N.D. 100 = i ampere
Distance between Approximately 18 cm.
Material of anode Rod or stick of cadmium.
Material of cathode Strip of platinum
Phenomenon at anode Solution of CdCb
Phenomenon at cathode Liberation of hydrogen
and the formation of
NaOH
Phenomenon between The precipitation of CdS
Special peculiarities^ Use of porous pots for the
production of the pure
pigment in quantity to
prevent mixing of elec-
trode products.
Note here the double current density at the anode as expressed
in the abbreviated manner, as the result of using a stick of cad-
128
EXPERIMENTAL ELECTROCHEMISTRY.
mium having one-half the area of the platinum strip. The next
pigment is that of mercury vermilion described as follows:
THE ELECTROLYTIC PRODUCTION OF MERCURY VERMILION.
The electrolytic production of this brilliant sulphide is a little
more difficult to accomplish, as the conditions must be exactly
right or the scheme does not work out as smoothly as that for the
preparation of the cadmium sulphide. Fig. 62 shows the arrange-
FIG. 62. Experimental Apparatus for the Electrolytic Preparation of Mercury Ver-
milion from a Mercury Anode.
ment of the apparatus, where a mass of mercury acting as anode
is shown within a small porcelain dish, with a strip of platinum
as cathode. This little dish may rest upon a block of glass, as for
example a rectangular glass paper-weight, for effect, if shown to
a number of persons as a lecture experiment. A platinum wire
runs 4own into the mercury, and is protected by a covering of thin
rubber tubing to prevent its acting as an electrode. A hydrogen-
sulphide generator similar to that used in the preparation of the
cadmium sulphide is employed, as shown at the right. The electro-
lyte consists of a solution of 8 per cent each of ammonium and sodium
nitrates,
NH 4 NO 3 and NaNO 3 ,
which electrolyze into NH 4 OH and NaOH at the cathode, and the
setting free of the two corresponding NO 3 groups at the anode,
which is of mercury, and consequently the soluble mercury nitrate
is- formed, Hg(NOs)2. The hydrogen sulphide acting upon the
nitrate of mercury produces the sulphide
ELECTROCHEMICAL OPERATIONS. 129
with the formation of two molecules of nitric acid, which, acting
upon the ammonium and sodium hydroxides, reforms ammonium
and sodium nitrate.
NaOH + HNO 3 = NaNO 3 + H 2 O,
NH 4 OH + HNO 3 - NH 4 NO 3 + H 2 O.
The following tabulation shows the important points to observe
in the successful preparation of this vermilion electrolytically :
Duration of run i hour
Electrolyte 8 grams NH 4 NO 3 , 8
grams NaNO 3 in I
liter of water
Character of apparatus See figure.
Dimensions of cell 8 cm.X 20 cm.X 20 cm.
Source of electricity Motor-generator
Temperature of electrolyte 50 C.
Amperes flowing 5 amperes
Volts indicated 5.5 volts
Area of anode immersed 200 sq. cm. approx.
Area of cathode immersed 100 sq. cm.
Current density at anode N.D. 100 = 2 .5 amperes
Current density at cathode N.D. 100 =5 amperes
. Distance between 12 cm. approximately
Material of anode Metallic mercury (see il-
lustration)
Material of cathode Strip platinum
Phenomenon at anode Formation of Hg (NO 3 ) 2
Phenomenon at cathode Production of NH 4 OH
and NaOH
Special peculiarities Arrangement of mercury
in small porcelain dish
For an effective lecture exhibit the vermilion sulphide should be
agitated by means of a mechanical stirrer. Fig. 63 illustrates four
electrolytic cells in series with electric motors attached to stirrers,
producing a very striking apparatus for the simultaneous production
of white lead and pigments. It is absolutely necessary to insulate
the stirrer rods from the shafts of the motors if a common electric-
lighting circuit and lamp-bank is employed for power for the
motors and electrolyzing current. If this is not carefully done,
there will, with most of the small motors on the market, be trouble-
some short circuits, and a failure to accomplish the electrolysis from
this cause. These motors may be joined in series, and the cells
130 EXPERIMENTAL ELECTROCHEMISTRY.
should be mounted upon a board, which may be lowered by remov-
ing the support in order to withdraw the electrodes and stirrers.
FIG. 63. Arrangement of Electrolytic Cells with Electromechanical Stirrers for the
Experimental Preparation of White Lead and Colored Pigments Simultaneously.
This apparatus is designed as a striking lecture-room experiment. It is necessary
to insulate the stirrer rods with hard-rubber connections as indicated in black if
a common electric -lighting circuit is employed for motors and electrolysis.
It is almost needless to say that this piece of apparatus will be also
useful for any other operations where mechanical agitation is re-
quired for prolonged periods of time. There are other pigments
which may be easily produced in the same general manner, the
details of which will be left to the ingenuity of the student. For
example, a beautiful green may be produced electrolytically as
follows :
THE ELECTROLYTIC PRODUCTION OF SCHEELE'S GREEN.
For this preparation dissolve 10 grams of sodium sulphate in
i liter of distilled water and place in the electrolytic cell as employed
for the previous compounds. The electrodes are cut from pure
soft sheet copper about 5 centimeters by 25 centimeters for the size
ELECTROCHEMICAL OPERATIONS. 131
of cell we are using, and about No. 18 gauge. The cell and electro-
lyte must be heated to a temperature of about 100 C. by means of
a water-bath. A little bag of tow is made and filled with white
arsenious oxide, which is suspended in the electrolyte. A current
of about 3 amperes is necessary for a cell of this size, and it is better
to employ the lighting current and the lamp-bank for the electrolyzing
current. Copper sulphate and sodium hydroxide are formed, the
sodium hydroxide dissolving the arsenious oxide and forming sodium
arsenite. The sodium arsenite immediately reacts with the copper
sulphate which separates as a beautiful green precipitate to form the
copper arsenite, regenerating sodium sulphate. The operation may be
conducted until the copper electrodes are consumed, and the arsen-
ious oxide completely converted to the arsenite of copper.
THE ELECTROLYTIC PRODUCTION OF PRUSSIAN BLUE.
In the same general manner Prussian blue may be prepared in-
the electrolytic cell. A potassium ferrocyanide solution of 10
grams to the liter is precipitated by means of a normal ferrous
salt solution. This precipitate is stirred in water by means of our
mechanical stirrer. This must be placed in a large porous pot of
sufficient diameter to admit a suitable stirrer in which a platinum-
anode is placed About 25 cubic centimeters of nitric acid is added
to the solution in the pot and a platinum cathode is placed outside.
About 5 amperes must be passed through the system for several
hours, when we will obtain for our product a dark Berlin blue.
A few words concerning experimental apparatus may be of ser-
vice to the electrochemist. Fig. 64 illustrates a most convenient
and satisfactory electrolytic stand for a great variety of purposes.
The column is solid glass, which serves to mount the two electrode
arms and effectually insulate them from each other. The electrode
arms not only slide up and down the column and are set by means
of a thumb-screw, but the electrodes may be slid in and out from
the center of the column and set at any required distance. For
rapid adjustment and flexibility of use these stands are unsurpassed.
A half dozen or more of these stands should be a part of every electro-
lytic equipment. They are always ready lo receive electrodes of
various kinds and materials, and are quickly connected to the source
EXPERIMENTAL ELECTROCHEMISTRY.
of electricity by means of binding screws. Many electrochemical
investigations may be begun by electrolyzing small volumes of
PIG. 64. Convenient Electrolytic Stand for Conducting Experimental Work. The
column is of solid glass, and the electrodes are easily adjusted for cells and beakers
of various kinds.
electrolytes in beakers with these stands until data warranting the
use of larger cells with separately fixed electrodes are secured.
CHAPTER IX.
ELECTROLYTIC PREPARATION OF POTASSIUM CHLORATE
FROM POTASSIUM CHLORIDE. PREPARATION OF SOLII>
TRIOXIDE OF SULPHUR. PRODUCTION OF OZONE FROM
THE ATMOSPHERE. EXPERIMENTS WITH OZONE AND ITS
PRACTICAL APPLICATIONS.
ELECTROLYTIC PREPARATION OF POTASSIUM CHLORATE.
AN outline of the purely chemical method for preparing this
important compound should be given before undertaking the elec-
trolytic plan. Potassium chlorate, or chlorate of potash, is prepared
in the chemical way by passing an excess of chlorine gas rapidly
into a strong solution of potassium hydroxide. When the liquid
becomes hot enough to decompose the hypochlorite first formed into
potassium chloride, which remains in solution, the potassium chlorate
deposits in tabular crystals. The ultimate result of such a chemical
method for this compound may be expressed by the following
equation :
6KOH + C1 6 = KC1O 3 + 5KC1 + 3 H 2 O.
If potassium carbonate or a weak solution of potassium hydroxide
be employed, the liquid will require boiling after saturation with
chlorine, in order to convert the hypochlorite into chlorate. This
chemical process for the preparation of potassium chlorate is far
from being economical, since five-sixths of the potash is converted
into chloride, being employed merely to furnish oxygen to convert
the chlorine into chloric acid. In manufacturing potassium chlorate
upon a large scale, a much cheaper material is used, namely lime, to
furnish the oxygen. The lime is mixed with water, and saturated
with chlorine gas in closed leaden tanks, when we obtain the following
reaction :
133
134 EXPERIMENTAL ELECTROCHEMISTRY.
The liquid is boiled down, when the calcium hypochlorite is decom-
posed into calcium chlorate and chloride,
The calcium chlorate is now decomposed by boiling with potassium
chloride, when it yields calcium chloride which remains in solution,
.and potassium chlorate which separates in crystals as the solution
cools. The following equation expresses this step in the process:
Ca(ClO 3 ) 2 + 2KC1 = CaCl 2 + 2KC1O 3 .
In the preparation of potassium chlora e by electrolytic means,
certain difficulties are met with. The simple plan of subjecting a
hot solution of potassium chloride to electrolysis, and allowing the
free mixing of the chlorine and caustic soda produced, is workable
only with low concentrations. To obtain the potassium chlorate
f.om an electrolyte containing chloride is a rather difficult matter
by means of any process of crystallizing out. Without a porous
partition or means of keeping the electrode products separated, the
anode or oxidized product will reach the cathode and will there be
reduced. At the same time the potassium hydroxide formed at the
cathode can serve independently as an electrolyte, yielding at the
electrodes oxygen and hydrogen gases, with the useless expenditure
of electrical energy. The student in electrochemical processes
should be on the lookout for every such possibility, and should have
such losses in mind in the design and construction of his apparatus.
Let us take up the experimental preparation of this most important
compound electrolytically in the simplest manner, using first an open
beaker glass without porous partitions The student can make his
determinations, and then modify the apparatus with a porous pot
with a view to increasing the working efficiency. For the simplest
workable plan we will employ an electrolyte of the following composi-
tion :
100 grams of potassi m chloride,
i gram of potassium carbonate,
i gram of potassium dichromate,
250 grams of hot distilled water.
Tig. 65 shows the arrangement of the necessary apparatus in its
POTASSIUM CHLORATE.
135
simplest form. In making records of experiments in note books, it
cannot be too strongly impressed upon the student to make concise
and neat little sketches of the apparatus and the manner in which it
was assembled, to be accompanied of course by a full tabulated list of
conditions as set forth in a previous chapter. In our first illustration
we have a large beaker resting upon a square of asbestos, supported
by a low iron ring tripod over a special low type of Bunsen burner
for maintaining the electrolyte at such an elevated temperature, as
is necessary. The anode and cathode are both of sheet platinum,
FIG. 65. Experimental Apparatus for the Electrolytic Preparation of Potassium
Chlorate from Potassium Chloride.
and are most conveniently supported by a couple of electrolytic
stands with insulating glass columns, as described and recommended
in the last chapter. A thermometer is immersed in the electrolyte
together with a glass tube as indicated, through which a current of
carbon-dioxide gas is passed. The current density at the anode
should be at least 20 amperes per 100 square centimeters of immersed
electrode, and the current density at the cathode should be about
double this for best results. The electrolyte is kept at a temperature
of 60 C. and a feeble acid reaction is maintained by a current of car-
bon-dioxide gas, which also keeps the solution agitated and of uniform
composition. According to Dr. Karl Elbs,'of the University of Giessen,
at least 60 ampere hours are necessary for this quantity of electro-
lyte, since one ampere hour yields 0.75 gram of potassium chlorate.
According to Dr. Elbs, if a separation of potassium chlorate has
136
EXPERIMENTAL ELECTROCHEMISTRY.
commenced during the electrolysis, after allowing the electrolyte to
cool down, a considerable quantity of potassium chlorate crys-
tallizes out, which is purified by a single recrystallization. By
evaporation of the mother liquor to about one-half its original volume
and then allowing it to cool down, a second lot of crystals is obtained.
The current efficiency amounts to about 70 per cent of the theoretical,
and only begins to fall off to a serious extent if over 50 per cent of the
potassium chloride has been converted into potassium chlorate. It
is very evident from the foregoing that if the current efficiency be
calculated from the amount of potassium
chlorate in the crystalline solid form, the
figure obtained will be too low because
of the quantity remaining behind in the
mother liquor. Sodium chlorate may be
prepared with the same apparatus by using
as an electrolyte,
80 grams of sodium chloride,
2 grams of sodium carbonate,
i gram of sodium dichromate,
250 grams of hot distilled water.
The electrical conditions are the same as
in the preparation of potassium chlorate
and the current efficiency is also about
70 per cent. The electrolyte in this case
is evaporated to a small volume, when
the sodium chloride will separate out of
FIG. 66.-Useful Design of U the solution and is filtered off. The
Tube for Conducting Elec- .
trolytic Preparations and ^ product IS not pure, being COn-
Maintaining an Elevated taminated with sodium chloride and
Temperature by the Current so dium chromate. Sodium chroma te is
Itself. The center tube ,. , , , . . ,
receives the thermometer. exceedingly soluble, and a complete pun-
fication is therefore a difficult matter.
Potassium and sodium chlorates may be quickly prepared in small
quantities without the aid of external heating by employing a special
U tube with a center limb like that illustrated in Fig. 66. Because
of the comparatively small cross-section of the electrolyte, the tempera-
ture can easily be kept at the required point by controlling the electric
PREPARATION OF SOLID TRIOXIDE OF SULPHUR.
137
current. The central tube serves for the reception of the thermometer.
Such a piece of apparatus not only serves for this particular line of
electrolytic work, but is of a generally useful and flexible character.
Fig. 67 illustrates such a tube in connection with an ammeter and
FIG. 67. Combination of Special U Tube with Ammeter and Lamp-bank, where-
by Electrolytes can be Maintained at their Boiling-points by the Electrolyzing
Current.' By means of the ammeter and the thermometer conditions can be
easily controlled.
lamp-bank, by means of which electrolytes can easily be maintained
at their boiling-points by the electrolyzing current if necessary.
PREPARATION OF SOLID TRIOXIDE OF SULPHUR.
It will be remembered that in a previous chapter the electrolysis
of concentrated sulphuric acid was conducted with the liberation of
hydrogen, oxygen, ozone, and free sulphur, and at elevated tempera-
tures the electrode products become hydrogen, oxygen, and sulphur
trioxide. If now we select a thin glass Woulf bottle intended for heat-
ing, and subject concentrated sulphuric acid to a process of electrolysis
with a heavy current, we will elevate the temperature sufficiently to
produce sulphur dioxide in quantity, which may be condensed to a
white solid with ease by passing the gas into a suitable condenser.
This experiment is not intended to show a method for preparing
sulphur trioxide, to compete with any existing chemical schemes,
138 EXPERIMENTAL ELECTROCHEMISTRY.
but merely as an exercise of interest in conducting practical work.
Fig. 68 shows the arrangement of the apparatus, which consists of a
thin Woulf bottle with platinum electrodes and a thermometer for
noting the temperature. There is a glass tube leading to a condenser
fitted with stop-cocks immersed in a freezing- mixture, and a lamp-
bank for modifying the lighting current. Electrical measuring instru-
ments are shown in this illustration, although the production of this
interesting compound upon this plan hardly warrants the setting
down of efficiency data. Like the experiment with the cold acid,
FIG. 68. Experimental Preparation of Solid Sulphur Trioxide by the Electrolysis of
Concentrated Sulphuric Acid.
the experiment is useful in demonstrating the composition of sul-
phuric acid by electrolytic means. Such a piece of apparatus is
very useful in conducting research work upon compounds when the
"mechanism of electrolysis is doubtful. For example, in electrolysis
we may be obtaining a condensible gas from one of the electrodes,
and it would be highly desirable to obtain this gas in a liquid or a
solid form for identification. In a certain piece of research work
undertaken by the author, oxides of nitrogen were given off on elec-
trolys's, and it was convenient to pass the electrode gas through such
a condenser immersed in liquid air, when the oxides of nitrogen con-
THE ELECTRICAL PRODUCTION OF OZONE.
densed to a light-blue solid, and were estimated in this form. Sul-
phur dioxide, as is well known by all general chemists, can be lique-
fied by passing it through such a condenser immersed in a simple
freezing- mixture of ice and salt. It is believed that, apart from
introducing an experiment, or a method of preparing a compound,
the description of special and useful apparatus will prove as valuable
to the experimenter and investigator as many theories and their
applications. Having produced a number of typical electrolytic
preparations in inorganic chemistry, we would neglect some very
important applications of electricity to chemistry, unless we prepared
certain important gases of commercial value.
Perhaps the first gaseous preparation should be T^T /
that of ozone, and the following pages will
therefore be devoted to this valuable product.
THE ELECTRICAL PRODUCTION OF OZONE.
Ozone is the chemical name applied to a
peculiar form of oxygen the exact nature of
which is open to some little discussion, as it
has been impossible to obtain absolutely pure
ozone. It is always accompanied with ordinary
oxygen, but there are good reasons for believ-
ing that ozone consists of three atoms of oxygen.
Three atoms of oxygen, occupying three vol-
umes, therefore, combine to one molecule of
OZOne which Occupies two volumes. Ozone, FlG - 6 9- ~ Experimental
, . , , . e f , . , , Apparatus for the Pro-
according to this scheme of formation, would duction of Ozone from
be one-half as heavy again as ordinary oxygen,
and experiments upon its rate of diffusion go
to support this theory. In the year 1785,
Von Marum noticed that oxygen upon
being subjected to electrical discharges, acquired an odor like
that experienced after the atmosphere had been subjected to an
active electrical storm. In 1840, Schoenbein called attention to
the similarity between the odors produced when air was treated
to electrical discharges, and the odors noticed when water was
electrolyzed between platinum electrodes, or gold electrodes, for with
the baser meta]s the production of ozone was not so marked. The
the Atmosphere. Dilute
sulphuric acid fills both
outside cylinder and in-
ner chamber of the cen-
tral tube.
140 EXPERIMENTAL ELECTROCHEMISTRY.
production of ozone was probably just as great, only it was largely
expended in oxidizing the electrodes when they were not of platinum
or of gold. The same investigator also observed that a like odor
accompanied the slow oxidation of phosphorus, and also of sulphur,
and that in each case a piece of filter-paper moistened with a solution
of starch and potassium iodide was turned blue. About the same
time, two other investigators, Marignac and De la Rive, showed
that ozone was only a changed condition of oxygen. In 1852,
Becquerel and Fremy demonstrated that pure oxygen could be
converted into ozone. In 1860, Andrews and Tait called attention
to the fact that a marked contraction in volume accompanied the
formation of ozone from oxygen, and in the same year Soret showed
that oil of turpentine absorbed ozone completely, and in this way
determined its relative density, confirming his results in 1867 by the
method based on the rate of the diffusion of gases as already men-
tioned. Andrews suggested at this time that the substance present
in the atmosphere which affected starch and potassium iodide paper
was this modified form of oxygen. We have already seen that ozone
may be obtained in the electrolysis of concentrated sulphuric acid
between platinum electrodes. Is this the most efficient and econom-
ical way of producing it? It is not, for there have been designed
and put into practice many forms of special ozone-generators of
greater efficiency. For laboratory purposes the generator as
illustrated in Fig. 69 has proven very useful. Here the apparatus
is shown in section in order that its several parts and their relations
may be clearly seen and understood. The tall outside cylinder glass
is filled with dilute sulphuric acid, one to ten, as well as the inner
chamber of the tube which is immersed therein. Two platinum
wires are dipped respectively into the acid in the outer glass cylinder,
and in the acid in the inner tube. When these platinum wires are
connected to the secondary terminals of a good induction-coil, the
two portions of the sulphuric acid become electrified, and there is
believed to be a stress set up which resolves th$ oxygen molecules
into its atoms, with a recombination to form molecules of ozone.
Dry air or dry oxygen is passed through the space between the two
portions of acid, which space for most economical action should be
as narrow as possible. Another important condition is the tem-
perature of the air or oxygen used. It should be cooled down to a
THE ELECTRICAL PRODUCTION OF OZONE. 141
low temperature before sending through, and should not be supplied
too fast. Fig. 70 illustrates a special horizontal form of ozone ap-
FIG. 70. Another Form of Ozone Apparatus Based upon the Same Principle. The
outer and inner tubes contain dilute sulphuric acid. The inner tube is sealed
up, only a small air-bubble remaining.
paratus based upon the same general principle. The longer the tube
and the greater the area exposed to the oxygen or air the greater the
ozonizing power of the apparatus for a stream of air or oxygen of
a given velocity and volume. The ordinary chemical test for ozone,
as has been intimated, is a damp mixture of starch and potassium
iodide, preferably made into test-paper. To make a good test-paper
for our work with ozone preparation, take 100 grams of starch
and grind thoroughly in a mortar with 50 cubic centimeters of distilled
water. The thin paste is then poured into 250 cubic centimeters
of boiling distilled water in a beaker, and one-half a gram of
potassium iodide crystals are added and made to dissolve by stirring.
Narrow pieces of filter-paper are drawn through this solution with
glass rods and hung up to dry. When these pieces of paper are
moistened, and exposed to ozone, they turn blue, but are not affected
by ordinary oxygen. The ozone abstracts the potassium from the
potassium iodide, and sets free the iodine, which has the specific
property of imparting a deep-blue color to starch. Papers impreg-
nated with manganese sulphate, lead acetate, or thallous oxide,
become brown, in the first two cases by the formation of the peroxide
of the metal, and in the last case from the formation of thallic oxide,
under the influence of ozone. Thus it will be seen that ozone is
an excellent and energetic oxidizing agent. If ozone is passed into
a solution of indigo, the blue color will soon disappear, since the
ozone oxidizes the indigo, and gives rise to products which, in a
diluted state, are nearly colorless. Ordinary oxygen is not capable
of bleaching indigo in this manner. If ozone is passed through a
tube of vulcanized caoutchouc, this will*-soon -be perforated by the
142 EXPERIMENTAL ELECTROCHEMISTRY.
corrosive effect of the gas, while ordinary oxygen would be without
effect upon it. If ozonized air be passed into a flask with a little
mercury in the bottom, the surface of the mercury will soon be
tarnished as the result of the formation of oxide, and when the mercury
is shaken around in the flask, it will adhere to the sides, which will
not be the case with pure mercury, as is well known. Let us look
into other forms of ozone apparatus and the application of the
electricity. Fig. 71 illustrates a simple Siemens tube, which is
FIG. 71. Cross-section and Elevation of an Original Siemens Apparatus for the
Production of Ozone from Atmospheric Oxygen. The sulphuric acid is dis-
placed in this tube by some metal not readily oxidized, such as tinfoil.
perhaps the simplest to construct in the laboratory. Here we have
merely an inner and an outer tube as shown in the sectional view.
The inner tube is held in position by a good cork and is lined on
the inside with tinfoil. This tube is closed at one end by sealing
off, and at the other end by a cork stopper through which a platinum
wire passes and makes contact with the tinfoil lining. The outer
tube has a side and end tubulure through which the air or oxygen
gas is passed. This outer tube is coated with tinfoil on the outside
with which electrical connection is also made. A good induction-
coil with a condenser of large capacity is employed with this apparatus
as with the acid tubes. In a lecture by Froelich he has given an
account of the experiments which have been made in the laboratory
of Messrs. Siemens & Halske in Berlin for the commercial production
of ozone. The original Siemens ozonizing tube was like that
illustrated, the two coatings being supplied with an alternating
current of high potential, while oxygen was made to traverse the
annular space between the two tubes. It appears that only one
dielectric is necessary, mica, celluloid, porcelain, and the like being
THE ELECTRICAL PRODUCTION OF OZONE. 143
available, as well as glass, and the ozone tube having either a metal
tube within and a metal-coated non-conducting tube without, or a
metal tube without, while the inner tube is made of the non-con-
ducting material and lined with metal. The metals to be used are
of course those which are not attacked by ozone, such as platinum,
tin, or aluminium. Cold water flows through the inner tube, and
through the annular space pure, dry air. Several such tubes may
of course be combined into one system and be worked equally well
with the alternating currents, although for single tubes it is stated
that direct currents operated by interrupters may be used to advan-
tage. The apparatus as described by Froelich in the German
laboratory yields 2.4 milligrams of ozone per second. Experiments
have been made with a view of producing compressed ozone for
technical uses, this having already been done on a small scale under
a pressure of about 10 atmospheres. One use of ozone especially
dwelt upon by Froelich is the oxidation of organic impurities in
drinking-water. Experiments with ozone upon water-supplies have
been made in this country, nevertheless, for a long time. In all
productions of ozone it is very important to keep the temperature
low, and the following experiment will serve to impress this fact
upon the student: If ozone be made to pass slowly through a glass
tube heated in the center by a Bunsen burner, it will be found to
lose its power of affecting starch and potassium iodide paper, the
ozone having been reconverted into ordinary oxygen under the
influence of heat. The formation of ozone may be compared with
the production of hydrogen peroxide, and we may express this break-
ing up of ozone symbolically. Just as hydrogen dioxide, H 2 O 2 ,
may be regarded as formed by the combination of a molecule of
water H 2 O, with an atom of oxygen, so ozone may be viewed as a
combination of a molecule of oxygen O 2 , with an atom of oxygen.
The breaking up by heat of ozone may, therefore, be expressed,
2(00 2 )= 3 (0 2 ).
A temperature of 250 C. is sufficient to completely bring about
this breaking up of the ozone. As we have learned, a given volume
of oxygen diminishes when a portion of it is converted into ozone by
the silent electrical discharge, and it regains its original volume
144 EXPERIMENTAL ELECTROCHEMISTRY.
when the ozone is reconverted by heat. It is of interest to note that
the conversion of oxygen into ozone is attended by the absorption
of heat; in other words, it is endothermic. The value of this heat
absurption may be expressed thus :
302 = 2Os 59,200 calories.
Ozone is, therefore, a very unstable body theoretically, and practically
we find such a state to be the case. To work any one of the ozonizers
economically in the laboratory the author has found it very convenient
and satisfactory to first dry the oxygen or air by passing it through
concentrated sulphuric acid, and then through a long tin or lead
worm immersed in iced water Chilled, dry oxygen of course gives the
richest yield of ozone. . A very efficient design of ozone-generator
and one of easy manufacture in the laboratory is illustrated in con-
nection with an induction-coil in Fig. 72. This simply consists of
a large outer glass tube of any length, containing a number of small
thin-walled glass tubes closed by sealing off at one end. Narrow
strips of tinfoil are pushed into each tube, as indicated, and joined
together in the manner of connecting up an electrical condenser.
A stream of chilled, dry air or oxygen is passed through the large
tube. Of course for a large outer tube of great length a very
powerful induction-coil must be employed. The adjustment of the
size and len'gth of tube to the induction-coil, the temperature, dry-
ness and rate of supply of air or oxygen, etc., make a very valuable
set of conditions for the student to experiment with with a view to
obtaining a combination for maximum efficiency. The electrical
energy supplied to the coil should of course be measured, as well as
every other chargeable item in the production of the body in question.
There are other ways of producing ozone electrically, although in
point of efficiency they caanot be compared with those described.
Among the experiments in electrolysis yielding as one of the products
ozone, may be mentioned the electrolysis of aqueous solutions of
nitric acid, hydrofluoric acid, sulphuric acid, and phosphoric acid,
as well as solutions of potassium nitrate, potassium phosphate, and
sodium sulphate. Hydrochloric, hydrobromic, or strong nitric
acid do not produce ozone when electrolyzed. Neither do aqueous
solutions of metallic chlorides, bromides, iodides, or ferrous sulphate.
THE ELECTRICAL PRODUCTION OF OZONE. 145
According to Houzean, the electrolysis of water furnishes only 3
to 5 milligrams of ozone per liter. When a given quantity of
oxygen is subjected to the action of charged surfaces, as presented
in any of the foregoing descriptions of ozonizers, only one-fifth, at
most, according to Bloxam, is converted into ozone; but if the ozone
be now removed by some substance which absorbs it, a fresh quantity
of the oxygen may be ozonized. The proportion of ozone formed
depends upon many conditions, the intensity and frequency of the
FIG. 72. Efficient Type of Ozone Apparatus and Induction-coil Assembled for
Production.
electrical discharge, the pressure, quantity, temperature, etc. The
question of temperature in all ozone processes appears to be by far
the most important. According to the older researches, about the
time 1880, 20 per cent of the oxygen becomes ozone at 25 C.,
and only 12 per cent at 20 C., the ordinary working temperature
of a chemical laboratory. At the temperature of boiling water, the
production is but 2 per cent. In more recently made determinations,
1893, it is stated that the production is 5.2 per cent at 20 C. The
matter of ozone production has occupied the minds of many inventors
during the last few years, and many forms of commercial appardtus
have been designed and constructed to produce this gas upon a
very large scale. There are drum-shaped chambers containing
many sets of stationary tinsel brushes arranged around the cylindrical
interior, with corresponding revolving sets of brushes to constitute
the opposite pole, for example. High potential discharges of vary-
ing current strength have been experimented with in such pieces of
146 EXPERIMENTAL ELECTROCHEMISTRY.
apparatus, and it has been found, to produce the best effects, that
high potential difference with small current and energy value must
be used. Probably this is due to the absence of heating effect with
electrical discharges of small quantity. A thick, heavy discharge
appears to break up the ozone formed almost as rapidly as it is
produced.
COMMERCIAL PRODUCTION OF OZONE AND ITS APPLICATIONS.
Among the early plants for the commercial production of ozone
for its practical application may be mentioned two equipments in
Germany, where ozone is used as a sterilizing and oxidizing agent
in connection with commercial water-works and supplies. These
plants are situated in Schierstein, near Wiesbaden, and in Paderborn
respectively. Ozone is well adapted for water sterilization and
purification, being a powerful destructive agent to all organic matter
both animal and vegetable. Of its efficiency there is no longer any
doubt, as experiments conducted over a long space of time all go to
prove its great efficiency. The question lies in the cost entirely,
and in order to secure figures of value upon this question we must
study a commercial working equipment. The two German plants
referred to were installed by the Siemens & Halske Company, of
Berlin. The plant at Schierstein is designed to sterilize 66,000 gallons
of water per hour as a maximum. This maximum output is called
for only on occasions when exceptionally large quantities of water
are required, while the average demands are about one-half of this,
or 33,000 gallons per hour. This plant is divided into two inde-
pendent units, like lighting and power units, the one being in opera-
tion while the other is at rest. The electrical generators installed
furnish electricity for operating pumps for the air and water, and
also for working the step-up transformers. These transformers
receive the current at a pressure of 120 volts and supply a current
from the secondary winding, at a pressure of 8.000 volts, the second-
aries of course being connected with the ozone-generators. The
ozone produced by the action of this discharge is driven by fans into
the sterilizing-towers, in which the ozone rises and comes in contact
with the water flowing down. The water is thus intimately mixed
up with the gas, and is thereby purified and flows off to a reservoir.
PRODUCTION OF OZONE AND ITS APPLICATIONS. 14?
In the ozonizer-room there are 48 ozone-generators mounted on
stout shelves on both walls opposite an aisle, the battery being
divided into two halves to be connected with the two generating
units respectively. Each half, therefore, contains 24 ozone pro-
ducers or generators, and these in turn are divided into three series,
of 8 ozonizers each, and each series of eight is connected to the 8000-
volt secondary winding independently. We have, therefore, three
sets of 8 ozonizers in series, connected in parallel to the secondary
of the transformer. One series of 8 ozone-generators produces a
sufficient quantity of ozone for one sterilization-tower. The ozon-
izers are of the Siemens-tube type, constructed of glass and metal,
one pole being grounded on account of the method employed of
cooling the inner tube with water. The connections to the positive
pole, the pole which is not grounded, are well protected, so that
it is not possible for the attendants to meet with accidents from the
high potential. Eight ozonizing tubes or generators are contained
in an iron box, the upper, lower, and front sides of which are provided
with heavy glass windows. The ozonizer-room is usually kept
darkened, and the attendant whose duty it is to enter it sees at once
from the blue light passing through the glass windows of the ap-
paratus whether it is working satisfactorily or not. As all the metallic
parts connected to the 8ooo-volt poles are carefully protected on the
one hand, and grounded on the other, there is little danger of a fatal
accident. The sterilizing-room in this plant contains two series,
each of four towers, of brick. Each tower is divided into four
sections by two partition walls perpendicular to one another. These
towers are filled with a coarse gravel through which the water trickles
on its downward flow, presenting a great surface to the action of
the ascending ozone. Through the combined four sections of each
tower, 11,100 gallons of water are flowing downward per hour,
while in the same time 2800 cubic feet of ozonized air pass upward
through the gravel. For the operation of each half of the plant,
sterilizing 33,000 gallons of water per hour, 50 horse-power are
required, of which 27 horse-power are used for the ozonizers, 22
horse-power for the pumps, and i horse-power for various minor
purposes. The cost of the process at Schierstein is 0.35 cent per
cubic meter of water sterlized, o.i cent being the cost of the coal
required for the operation of the ozonizing apparatus. To this
148 EXPERIMENTAL ELECTROCHEMISTRY.
must of course be added the usual interest on the money invested
in the plant, and the cost of depreciation and repairs. This par-
ticular plant has to operate pumps in addition to its electrical appara-
tus, which is very unusual, and can be deducted from the cost in
almost any other water-works. Tests of the sterilized water, made
by leading German chemists and bacteriologists, proved that the
sterilization process by means of this apparatus is in every way
successful so far as the annihilation of germs and bacteria is concerned.
The process has therefore been shown to be well adapted to the purifi-
cation of drinking-water. The installation at Paderborn is similar
to that at Schierstein, with the exception of the sterilizing-towers.
The purified water in this plant is allowed to flow off in the form
of a cascade, in order that the contained ozone dissolved may be
liberated and given off. The cost of purifying a cubic meter of
water in this equipment is placed at 0.4 cent, including interest,
depreciation, wages, etc.
Apparatus has been designed, and constructed, and tested at
Niagara Falls in this country, using an electrical discharge under
a potential difference of 50,000 volts against the 8000 employed at
the German works. Ozone produced with this equipment was
passed through two pipes to a water-sterilizing-tower partly filled
with broken rock, one supply-pipe entering the bottom and the other
the top of the tower. The water flowing through the tower down-
ward meets with the gas from all sides, above and below, and is
purified by its contact with the several ozone streams. It is claimed
that 30 to 40 horse-power are sufficient to sterilize 6000 tons of water
daily. Some interesting experiments have been conducted by
Warburg, who investigated the production of ozone by discharges
in a closed volume of oxygen. In this case the formation of ozone
has a limit which varies widely with the conditions of the experiment.
Besides the ozonizing effect of the electrical discharges, there exists
also a contrary effect which counterbalances the production after a
time. Since for the limited duration of the ozonizing process the
spontaneous decomposition of the ozone is negligible, it follows that
the electric discharge itself must produce the contrary effect. In
endeavoring to obtain a measure of both effects, Warburg starts
from the assumption that the ozonizing effect is proportional to the
number of molecules of oxygen present, while the decomposition
PRODUCTION OF OZONE AND ITS APPLICATIONS. 149
is proportional to the number of ozone molecules present. There
is a decided difference between the behavior of a positive discharge
and that of a negative discharge; the maximum percentage of
ozone produced, according to Warburg, is about three times as high
for negative as it is for positive discharges. If after the negative
maximum is reached, the mixture is subjected to the positive dis-
charge the percentage falls to the positive maximum. The de-
composing activity is the same for positive and negative electricity,
but the ozonizing action itself is three times higher for negative
electricity. The effect of temperature upon both discharges is the
same. The author has not verified these data, but introduces them
as the work of a careful and reliable experimenter. If partly ozonized
oxygen is left to itself, the proportion of ozone decreases, according
to Warburg, in an interesting manner for different conditions. This
decrease is known under the name of spontaneous deozonization, and
may be due to a great variety of causes. The experimenter referred
to points out that since the amount of ozone at 200 C. in a state
of equilibrium is inappreciable, such equilibrium can be attained at
ordinary temperatures only by complete dissociation of the ozone.
The factors bringing about this dissociation may be external, as
for instance, a contact with foreign bodies or with the walls of the
vessel; or they may be internal and due to the collision of two
molecules of ozone, or of a molecule of ozone with a molecule of
oxygen. Warburg formulates a mathematical theory of the process,
and finds by suitable experiments that the internal effects are almost
entirely due to collisions between the ozone molecules themselves.
Heat increases this internal disintegration much more rapidly than
it does the external process, and it becomes the chief agent at high
temperatures. The experimenter in question finds that moisture
makes no difference in the stability of ozone at 100. The various
forms of ozone- generators suggested and in use to-day would fill a
volume in themselves, and it is believed that the fundamental prin-
ciples and typical forms introduced here will be a sufficient guide to
the student who wishes to experiment with this interesting and
valuable modified condition of oxygen. In the next chapter another
gaseous product will be dealt with of no less scientific interest or
technical value. This will pertain to the electrical production of
nitric acid from the atmosphere.
CHAPTER X.
THE PRODUCTION OF NITRIC ACID FROM THE ATMOSPHERE
WITH HISTORICAL NOTES.
WE know, as general chemists, that nitric acid is one of the most
important of all chemical compounds, and that it is usually obtained
from potassium nitrate, which is obtained as an incrustation upon
the surface of the soil in hot and dry climates, as in certain parts of
India and Peru. The salt imported for the chemical production of
nitric acid from Bengal and Oude consists of potassium nitrate,
while the Peruvian or Chilian saltpeter is sodium nitrate. Either
of these nitrates serve for the production of this important acid.
On a small scale in the laboratory, nitric acid is prepared by the
distillation of sodium or potassium nitrate with an equal quantity
by weight of concentrated sulphuric acid. The decomposition of
potassium nitrate by an equal weight of sulphuric acid is represented
by the following chemical equation:
KNO 3 + H 2 SO 4 = HNO 3 + KHSO 4 .
It would appear from a study of this equation that one-half of the
sulphuric acid might be saved, inasmuch as one molecule could be
made to decompose two molecules of potassium nitrate as follows:
2 KNO 3 + H 2 SO 4 = 2HNO 3 + K 2 SO 4 ,
but it is found that when a smaller quantity of the sulphuric acid is
used a very much higher temperature is necessary to bring about the
complete breaking up of the saltpeter, and that much of the nitric
acid is decomposed, and the normal potassium sulphate, K 2 SO 4 ,
which is the final result, is much more troublesome to get into solu-
tion in order to remove it from the retort or still. For the prepara-
tion of large quantities of nitric acid sodium nitrate is used instead
150
NITRIC ACID FROM THE ATMOSPHERE. 151
of the more costly potassium nitrate. The chemical production of
nitric acid dates back to very early times. Geber, the old Arabian
alchemist, produced this acid by distilling saltpeter, copper vitriol,
and alum. As early as the thirteenth century, Raymund Lullius
described its preparation by distilling saltpeter with iron vitriol.
Glauber, in the seventeenth century, produced it by the same method
as is employed to-day, that is by the distillation of potassium or
sodium nitrate with sulphuric acid. Cavendish, in his studies of
the atmosphere, showed that it consisted of a mixture of oxygen and
nitrogen. Although these elements in their pure condition show
no attraction for each other, five distinct compounds of oxygen and
nitrogen are prepared in various indirect ways. These compounds,
which should be familiar to all of us who have given attention to
general inorganic chemistry, are as follows:
N2O, nitrous oxide;
NO, nitric oxide;
N 2 Os, nitrogen trioxide ;
NO 2 , nitrogen peroxide;
N 2 O 5 , nitrogen pentoxide.
When a succession of powerful electric sparks were made to pass
through dry air in a flask, a red gas, NO 2 , nitrogen peroxide, was
formed, and when these discharges were allowed to take place in the
presence of moisture, this gas was absorbed and converted into a
mixture of nitrous and nitric acids,
2 NO 2 + H 2 O = HNO 2 + HNO 3 .
If instead of water we place in the flask a weak solution of potassium
hydroxide, instead of obtaining the mixed acids, we obtain the two
salts, potassium nitrite, and potassium nitrate:
HNO 2 + KOH = KNO 2 + H 2 O,
HNO 3 + KOH = KNO 3 + H 2 O.
By evaporating to dryness such a solution, we obtain a saltpeter
consisting of potassium nitrite and potassium nitrate, and if we
distill this product with strong sulphuric acid we will obtain nitric
acid. Cavendish went through this process and prepared the above
EXPERIMENTAL ELECTROCHEMISTRY.
compounds without difficulty. If the experiment of passing sparks
through air (for best effects the air should be mixed with some pure
oxygen) be repeated in a U tube having one of the vertical limbs
of the tube provided with a stoppered globe into which the two
sparking wires are sealed, and we fill this system, all but the globe,
with water colored blue with litmus, the solution will be reddened
by the acids formed, and the air in the globe will be found to diminish
in volume. Fig. 73 shows this form of "U " tube and the arrange-
FIG. 73. Special Form of U Tube for Showing Absorption of the Oxides of Nitrogen.
by Water,
ment of the experiment. The blue litmus solution is placed in the
tube with the stopper in place so that we will have a difference in levels
as indicated in the illustration. A little scale, as shown at the left,
assists us in determining the diminution in volume when the oxides
of nitrogen combine with the contained water. This formation of
oxides of nitrogen may be regarded as the combustion of nitrogen in
oxygen due to the kindling effect of the electric spark. A simple non-
electrical experiment to illustrate the probable combustion of nitrogen
in oxygen may be performed by igniting a piece of magnesium ribbon
in a tall glass jar of carefully dried air. If such an experiment be
performed and one looks down into the jar red fumes may be seen
and the presence of oxides of nitrogen may be shown by drawing the
contents of the jar through a mixture of potassium iodide with starch
and acetic acid, when the iodine will be set free, which blues the
NITRIC ACID FROM THE ATMOSPHERE. 153
starch. This experiment makes it appear likely that the electric
spark causes the combustion of nitrogen in oxygen on account of a
kindling effect. When ozonized air is passed into water nitric acid
is found in solution. It has also been shown that rain-water contains
about one part per million of nitric acid. When hydrogen gas,
mixed with a small quantity of nitrogen, is burned, the water collected
as the result of the combustion is found to be slightly acid, due to the
presence of some nitric acid, due to the combination of the nitrogen
with the oxygen of the air under the high temperature of the burning
hydrogen. With the recording and appreciation of such experiments
a new method of preparing nitric acid began to dawn upon chemists
in general. In the year 1859, Newton produced nitric acid from
the atmosphere by constructing a modified form of U tube and
bulb, differing from that already illustrated only in the shape of the
chamber. Newton employed a more spacious design of chamber
for the air and water into which he inserted his sparking wires.
Means were provided in this chamber for sending in a fresh supply
of air and for allowing the acidulated water produced to run off.
Prim, in 1882, modified this arrangement and used both a silent
electrical discharge and a series of sparks to bring about the combina-
tion of the nitrogen with the oxygen present in a chamber. In the
year 1392 Sir William Crooks produced what he termed an electric
flame, through the agency of a high voltage alternating electric cur-
rent discharge between pieces of platinum, and by its means brought
about the combustion of nitrogen in oxygen. Lord Rayleigh and
Professor Ramsay employed such a flaming arc in a number of their
researches upon the atmosphere in connection with the oxidation of
nitrogen. They employed an alternating current stepped up to a
difference of potential of 8000 volts. Lord Rayleigh in his work upon
the production of nitric acid from the atmosphere employed a spherical
glass chamber with a capacity of about 50 liters in which could be
maintained a fountain of sodium or potassium hydroxide solution.
Into this chamber air and oxygen were slowly pumped containing
a flaming discharge between platinum terminals. It is stated that
Lord Rayleigh produced with this equipment about 40 grams of
nitric acid per hour with the expenditure of about one electrical
horse-power in energy. McDougall and Howies describe an equip-
ment for producing nitric acid from the atmosphere, and its per-
154
EXPERIMENTAL ELECTROCHEMISTRY.
formance, which is of special interest to the student in this line of
work. They outline an alternating current-generator with a fre-
quency of 60, equipped with two separate armature circuits. One
of these windings delivers a current of 24 amperes at 200 volts
pressure, and the other 12 amperes at 400 volts pressure. The
current of this alternator, from either winding at will, can be led to
1
c
^
FIG. 74. Stone-ware Pipe used as Combustion Chamber for Producing Nitric Acid
from Atmosphere.
a step-up transformer with a ratio of i to 40, so it will be seen that
a voltage at the secondary winding of the transformer of either
8000 or 16,000 volts may be obtained. The current in the primary
of the transformer may be very closely measured by inserting an
ammeter, and through the agency of a voltmeter at the terminal of
the alternator the initial electrical pressure can be recorded. A
NITRIC ACID FROM THE ATMOSPHERE.
'55
wattmeter was also included in the primary circuit by means of
which the readings of the voltmeter and the ammeter could be checked
and the energy consumed in the flaming discharge could be ascer-
tained. An early form of combustion chamber equipped with
platinum-iridium electrodes for the production of nitric acid from
the atmosphere with this electrical equipment, consisted of a stone-
ware pipe of the shape depicted in Fig. 74. Air is drawn through
this chamber during the time of an electrical discharge, and the
oxidized nitrogen is drawn through a series of Woulf bottles con-
taining either water, or a solution of sodium hydroxide. Instead
of employing the Woulf bottles a series of towers may be employed
to great advantage, each tower containing broken glass for the
FIG. 75. Modified Form of Combustion Chamber for the Production of Nitric Acid.
water to run over and cause it to present a great surface for absorp-
tion to the oxides of nitrogen which are made to pass through. We
can produce a solution of sodium nitrate, or free acid, at will with
this equipment. The form of the combustion chamber itself under-
went numerous evolutions, one or two of the shapes being given here
as a matter of interest and guide to the student. Fig. 75 shows one
156
EXPERIMENTAL ELECTROCHEMISTRY.
of these modifications. This chamber consists of a large stoneware
bottle with vertical supply and outlet tubes for the gases. Lord
Rayleigh showed that these combustion chambers should be quite
spacious; in other words, that there should be a considerable amount
of room about the flaming discharge. If the air is not passed through
at a certain maximum rate, the oxides of nitrogen formed will be
broken up again by the discharge, very much like the behavior of
ozone when subjected to heat after its formation. The oxides of
nitrogen should be removed from the field, or influence of the arc as
soon as possible for high efficiency in yield. A too rapid supply of
air causes the electrical discharge to become unsteady, and a loss in
efficiency results. . Some interesting and valuable data were obtained
on nitric acid from air by McDougall and Howies using various
forms of combustion chambers, and supplying the electricity under
different conditions. Fig. 76 is another form of combustion chamber
FIG. 76. Another Modification of Combustion Chamber.
as described in the work of these experimenters. They varied the
forms and sizes of the chambers, keeping the electrical conditions
constant, and kept the forms of combustion chambers constant,
NITRIC ACID FROM THE ATMOSPHERE. 157
and varied the character of the electrical discharges, making quanti-
tative determinations of the nitric acid obtained in each case. By
varying the current value in the secondary circuit of the transformer,
keeping the voltage constant, these experimenters were able to vary
the temperature of the flaming discharge, and to study its effect upon
the yield of nitric acid produced. They formulated the following
table which brings out the fact that a high temperature discharge is
unfavorable to high efficiency, and consequently we can produce
more acid with a fewer number of watts, or, in other words, with less
electrical horse-power.
Watts Used Current Used in Yield of Acid per H.P.
in Flame. Flame. per 12 Hours.
302 .3 to .38 ampere 180 grams
225 .2 " .25 " 270 "
172 .15 " .2 " 300 "
Too great a decrease in the current value of the flame caused the
flaming arc to become unsteady and liable to extinction. The
production of nitric acid by these experimenters with the foregoing
electrical equipment and design of chambers, at the rate of 300
grams of nitric acid per horse-power for twelve hours, represents
51.5 per cent of the amount theoretically obtainable from the amount
of air supplied. This result compares very favorably with the
figures obtained by Lord Rayleigh, working with a mixture of oxy-
gen and nitrogen in the proportion of two volumes of oxygen to one
volume of nitrogen, when he obtained 440 grams of nitric acid
in the same time and with the same consumption of electrical energy.
When ordinary air is used in these chambers, the theoretical pro-
portions of oxygen and nitrogen for the production of the oxide
are diluted with an excess of nitrogen, which of course is detrimental
to the best effects. An experiment was conducted in one of these
chambers with a mixture of pure oxygen and nitrogen gases in the
proportion of one volume of oxygen to two volumes of nitrogen, and
the yield was 590 grams of acid. To study the effects of tem-
perature upon the rate of oxidation, the air supplied to the com-
bustion chamber was raised in temperature by passing it through
a porcelain tube packed with asbestos, around which a heating
coil of platinum wire was wound. On passing a strong current
of electricity through this wire in the tube, the entire system was
158 EXPERIMENTAL ELECTROCHEMISTRY.
elevated in temperature to incandescence, and the air after passing
through was immediately treated in the flaming discharge. This
porcelain tube was cemented directly into the stoneware combus-
tion chamber. A marked decrease in the yield of nitric acid was
noted with the same supply of energy. This behavior is entirely
in accord with the behavior of air when treated to form ozone.
Kowalski describes his apparatus and method for producing nitric
acid from the atmosphere, and states that the yield of product is
largely dependent upon the frequency of the alternating current
used for the flaming discharge. At high frequencies the best re-
sults are attained. Kowalski and Moscicki, working with an
alternating current with a frequency between 5000 and 6000 cycles
per second, obtained 43.5 grams of nitric acid per kilowatt hour
with an amperage in the secondary of .2. They have also con-
firmed the work of McDougall and Howies and others, that the
amperage in the secondary of the transformer has a direct bearing
upon the yield, and that for high current strength the yields of acid
are not so great. The present writer has also fully confirmed this
statement. Kowalski and Moscicki also found that the influence
of the length of the flaming discharge is decided. With a current
in the secondary of only .05 ampere and a difference of potential of
50,000 volts, and a frequency of 6000 to 10,000 cycles per second,
they obtained a maximum yield of nitrous vapors for the energy
employed. They have obtained from 52 to 55 grams of nitric acid
per kilowatt hour, which yield could be nearly doubled by adding
about 50 per cent of pure oxygen gas to the air employed in the
combustion chamber. The work of Bradley and Lovejoy for the
production of nitric acid from the atmosphere upon a large scale is
especially noteworthy. They point out that whereas the silent
electrical discharge, and the spark or disruptive discharge can
cause the combination of oxygen and nitrogen gases, they have but
feeble capacity in point of efficiency, which is also the case with the
ordinary arc. To obtain the best results they point out, as a result
of lengthy researches, that it is necessary to employ an arc divided
into numerous thin and flat subdivisions in order to present a large
surface for a small amount of electrical energy. This work is only
in keeping with that of previous experimenters. They point out that
the thinner the arc the greater the efficiency of the process up to
NITRIC ACID FROM THE ATMOSPHERE. 1 59
the point where the arc breaks. It is necessary for best effects to
greatly subdivide the current by arranging the arc circuits in par-
allel. In the experimental apparatus of Bradley and Lovejoy, they
employ a direct current of .75 ampere at a pressure of 8000 volts,
which can be increased to several ampers at 15,000 volts. There
are 138 arcs between which this current is divided, each of which
is made and broken by a revolving mechanism 50 times per second.
Each arc has a current value of only .005 ampere. The arcs are
all produced successively, and not at one time, by a special arrange-
ment of the wire electrodes on the revolving drum which carries them.
In their recent apparatus there are 6900 arcs formed and extinguished
per second, each arc lasting only for the brief period of 1/20,000 of
a second. As each little arc tends to increase in volume, due to
increased conductivity as soon as it is formed, the tendency to short
circuit the others is avoided by placing inductance-coils in series
with the arcs. These small inductance-coils are so designed and
calculated that during about 1/40,000 of a second they delay or im-
pede the flow of the current, thereby preventing a rapid growth
of the arc, and during the succeeding 1/40,000 of a second, while the
arc is being drawn out by the revolving-drum carrier, it sends an im-
pulse which increases the current flow and so prolongs the arc.
With a difference of potential of 8000, and about i /2oo of an ampere
of current, the arcs are drawn out 4 to 6 inches, and the oxygen and
nitrogen treated to this discharge. As in the production of ozone,
the molecules of nitrogen oxide must be removed from the field before
the atoms of oxygen and nitrogen dissociate, and this point is
especially remarked upon by these later investigators. The appara-
tus installed at Niagara Falls by Bradley and Lovejoy is about
5 feet high by 4 feet in diameter, built of iron of cylindrical form.
Six rows of inlet wires, well insulated by porcelain sleeves, enter the
sides of this chamber. The terminals of the electrodes are of plati-
num wire, turned downwards to spread out the arc in a thin flat
discharge. In the center of the iron cylinder is a shaft vertically
arranged, carrying a series of 23 radial arms, corresponding to the
23 rows of points which enter the sides, there being six radial arms
in each plane. These radial arms are each tipped with platinum
wire, which come within 1/25 of an inch of the platinum wires of the
opposite stationary poles. As the two platinum wire points approach
160 EXPERIMENTAL ELECTROCHEMISTRY.
upon the turning of the shaft, a spark jumps a gap of about ^ of
an inch to meet the approaching platinum point, and then the arc is
drawn out from 4 to 6 inches by the retreating point until it breaks.
If it were not for the little inductance-coils in series with these arcs,
it is evident that we would have a comparatively heavy arc at the
time of formation, increasing as the points come nearer together.
Each of these coils is immersed in oil, and is 5 inches in diameter
by 12 inches long, and contains several thousand turns of fine in-
sulated wire. The central shaft carrying the moving electrode
points, is turned at the rate of 500 revolutions per minute, and
takes a little over i horse-power. The efficiency of this piece
of apparatus is stated to be i pound of nitric acid, per 7 elec-
trical horse-power per hour. Carefully dried air is used in this
combustion chamber, which is protected by a coating on the inside
of asphalt varnish. If moisture was present in the air, nitric acid
would be formed within the chamber, and would in time lead to
serious corrosions. The oxides of nitrogen are led to a tower down
and through which water trickles, for the production of nitric acid,
or a solution of sodium hydroxide, for the production of sodium
nitrate. It has been suggested that milk of lime be made to flow
through one of these towers for the production of calcium nitrate
for fertilizing purposes. This is reported to be a cheaper method
of getting nitrogen into the soil than by using sodium nitrate, with
lime, at $1.50 per ton. The present writer has experimented with
both ozone production and nitric-acid production, employing volt-
ages as high as 250,000, and is able to confirm the results of these
experimenters from his own note-book. This field is one of great
charm to the student of electrochemistry, and a few concise details
as to methods of producing nitric acid from the air may prove
welcome to him. Although the foregoing work by others in this
line will give the student the fundamental principles involved, and
the basic information upon which to experiment, a few specific
directions may not be out of order. Let us construct ourselves
a simple piece of apparatus of our own design for producing nitric
acid by electrical means for use in the laboratory or lecture-room.
Fig. 77 shows a simple and easily constructed design of chamber
for rotating electrodes, and which has proven most satisfactory in
the laboratories of The George Washington University, where it
NITRIC ACID FROM THE ATMOSPHERE.
161
was built and operated. This chamber consists of a casing of
pine wood with an outside diameter of 14 inches, and an internal
diameter of n inches. The internal width of the chamber is 3
inches. The shell for this chamber was cut out on a band-saw at
162 EXPERIMENTAL ELECTROCHEMISTRY.
a lumber mill, and two wooden side plates securely screwed against
the sides. The chamber was carefully protected on the inside by
several coats of acid-proof paint. The steel spindle which carries
the hub into which the four electrodes are screwed, is supported, and
turns into holes drilled in two disks of hard rubber, screwed against
the sides of the casing. The wood of the casing is cut away at the
center so the steel spindle constituting one pole of the high-tension
alternating current is in contact with hard rubber only. These
circular openings in the wood are 6 inches in diameter, and the
diameter of the hard-rubber plates screwed against the sides is
8| inches. The stationary electrode which protrudes through
the top in a vertical position is connected to the other terminal of
the high potential transformer, and is carefully insulated from the
wooden casing by passing through a hard-rubber tube. There is
a brass commutator wheel on the outside, as shown, upon which a
brass brush rests, which is in turn supported by a hard-rubber block,
as shown. A hard-rubber grooved pulley is on the end of the shaft
for the belt of the driving motor. On no account construct this
apparatus without the generous use of hard rubber, for in our ex-
perimental work it may be desirable to increase the voltage to 16,000
or to even 32,000, and ordinary wooden insulation would be value-
less as an effective insulator. An inlet and outlet tube must be
provided, as shown, for the air-supply, and a little glass window
should be provided to enable us to see the condition of the electrical
flame within. Fig. 78 shows this piece of apparatus assembled in
connection with a transformer for high potential electrical discharges,
a driving motor for the electrodes, a foot-bellows for air-supply, and
a Woulf bottle for the absorption of the oxides of nitrogen in
either water or caustic soda. The mechanical and electrical con-
ditions set down in one experiment with the present equipment
are as follows :
Revolutions per minute of electrodes 150
Alternating current cycles per second 60
Complete reversals, therefore, per second 120
Voltage at secondary of transformer 10,000
Amperage in naming arc 1/25
Jump gap when electrodes were opposite each
other in fraction of an inch 1/16
Length of naming arc at time of breaking by
being drawn out, in inches 3^
Appearance of naming arc Pale yellow
Type of transformer Oil immersion
NITRIC ACID FROM THE ATMOSPHERE.
163
With these conditions the yield of nitric acid was excellent, taking
in the air at the temperature of the laboratory. This apparatus
was operated under various conditions and the oxides of nitrogen
estimated in several ways. In the place of the Woulf bottle, U
1 64 EXPERIMENTAL ELECTROCHEMISTRY.
tubes were substituted in several runs, which were immersed in large
beakers of liquid air, when the oxides of nitrogen were condensed
to a light-blue solid and estimated in this form. It is believed that
the construction and operation of this piece of apparatus will fully
repay the student having at hand the necessary high potential
electrical discharge for the production of nitric acid.
CHAPTER XI.
THE ISOLATION OF THE METALS SODIUM AND POTASSIUM.
BY giving specific attention to the metal sodium in this chapter
we will also be covering in a general manner the means employed
for the isolation of potassium. We will, therefore, refer to sodium
in this chapter, and it will be understood that potassium may also be
isolated by similar apparatus introduced and used under like con-
ditions. Sodium is much cheaper than potassium, as it is well
known by all chemists, because of the greater abundance of cheap
sodium salts. As we learned in the opening chapter Sir Humphry
Davy was the first to obtain metallic sodium through the agency of
the electric current and a mercury cathode. We know that sodium
is one of the most abundant of all chemical elements, and that it occurs
in immense quantities in combination as rock-salt deposits in saline
springs and in sea- water. Sodium also occurs in the form of
nitrates, borates, carbonates, etc., etc.
The cheapest source of sodium is, of course, from sodium chloride
or common salt. Rock salt forms very considerable deposits in
many regions. Among the most important are those at Northwich,
in Cheshire, England, where very large quantities are extracted by
mining processes.
Sodium has been prepared by an ordinary chemical process by
reducing its oxide by carbon at a white heat. The following chemical
equation indicates the character of the reaction:
This old process was worked by taking 30 kilograms of dry-
sodium carbonate, 13 kilograms of charcoal, and 3 kilograms of
chalk. These were thoroughly mixed together, calcined, and intro-
duced into iron cylinders heated in reverberatory furnaces. At
165
l66 EXPERIMENTAL ELECTROCHEMISTRY.
a bright red heat the sodium distills over and is collected in suitable
receivers. It is purified by redistillation, and then melted under
petroleum into ingots which are preserved under naphtha or other
suitable hydrocarbon.
Another chemical process devised by Castner, consisted in reducing
sodium hydroxide by heating it to a temperature of 850 C. with an
intimate mixture of finely divided iron and carbon prepared by
mixing the iron with molten pitch. These old chemical methods
have been entirely replaced by electrolytic processes, and it is the
purpose of this chapter to outline the principle upon which the
electrochemical method is dependent. As we have learned, the
electrolytic decomposition of sodium and potassium hydroxides led
to the discovery of these metals. Sir Humphry Davy, writing in the
Philosophical Transaction in 1810, describes his research as follows:
"By means of a stream of oxygen gas from a gasometer applied
to the flame of a spirit-lamp, which was thrown on a platina spoon
containing potash, this alkali was kept for some minutes in a strong
red heat, and in a state of perfect fluidity. The spoon was preserved
in communication with the positive of the battery of the power of
100 of 6 inches, highly charged, and the connection from the negative
side was made by a platina wire."
This method of Sir Humphry Davy, although theoretically
attractive, does not work very smoothly in practice and we will see
that very special precautions must be taken in order to get a satis-
factory yield of either metallic sodium or potassium. Many modifi-
cations of this classic experiment have been made, among which
may be mentioned the use of a platinum dish containing a strong
solution of potassium hydroxide and metallic mercury in the bottom,
which is connected to the negative electrode of a suitable battery.
We have here the dawn of the practical processes which followed
where the containing vessel is made the cathode in electrolysis.
One of the earliest designs of commercial apparatus is that of Charles
Watt, which is described in his specifications in 1851. The following
account is from his own specifications: "The second part of my
invention consists of a mode of preparing or obtaining the metals of
the alkalies and alkaline earths by the united action of electricity
and heat. For performing this part of my invention by the united
action of electricity and heat, I employ a vessel [of the form shown
THE ISOLATION OF THE METALS SODIUM AND POTASSIUM. 167
in Fig. 79], which is made of iron or other suitable material capable
of bearing a full red heat. In this figure A is the vessel, which
should be at least one-half an inch thick, and, if made of iron, previ-
ously to its being used should be coated over its exterior with clay
or other substance to preserve it from the action of the fire; B,
movable head for the collection of the metals; C, electrodes, with
B
FIG. 79. Watt's Electrolytic Cell for the Production of Sodium and Potassium.
their attachments E; D, flanges to support the vessel upon the furnace.
The covered compartment F, being that in which it is intended to
eliminate the metals, is supplied with a carbon electrode and the
uncovered compartment is supplied with a gold electrode; but I
wish it to be understood that I do not restrict myself to the particular
form of apparatus, or to the material to be used for electrodes. The
vessel is filled with dry saline matter, so that when it is in a state of
fusion it shall reach the dotted lines [the author has shown a full
black line drawn across the interior of the cell]; the partition keeps
the eliminated substances from reacting on each other, and also
1 68 EXPERIMENTAL ELECTROCHEMISTRY.
excludes air from the compartment in which the metal is eliminated,
the access of which would cause the metal to be oxidized. The
vessel is placed in a furnace where it can be subjected to the action
of a full red heat, and when the saline matter is in a state of fusion
contact is made between the decomposing vessel and the apparatus
supplying the electric current or currents, the intensity of which
should, at least, be equal to that which would be supplied by 10 cells
of DanielPs battery arranged for intensity, but, of course, this depends
upon the nature of the salt which is being decomposed. The fused
salt is maintained at that temperature which will ensure the instantane-
ous volatilization of the metal as it is eliminated, and a proper receiver
(such a one as is usually employed for the preparation of such metals
will answer) is connected air-tight with the narrow tube projecting
from the head. The metal is received and preserved in any con-
venient fluid hydrocarbon. The salts which I usually employ are the
chlorides, iodides, or bromides of the metals of the alkalies or alka-
line earths."
This historic piece of apparatus proved to be absolutely worthless
in commercial practice, for it is impossible to successfully distill such
metals in a retort chamber of this peculiar design. There are
many weak features about this apparatus which condemn it for all
serious uses. The gold anode, apart from its prohibitive cost, would
have but a short existence in a fused electrolyte such as he describes
where it would be subject to the action in addition of nascent chlorine
gas. Numerous forms of sodium and potassium cells had their rise
and fall, the majority of them being designed without a suitable
knowledge of the severe conditions and requirements for the success-
ful preparation of these metals. One of the early workable designs
for a practical sodium cell was that of Borcher, which is illustrated
in Fig. 80. The melting-vessel A had an opening surrounded by the
socket tube B, and two other openings with tubular necks C. The
double socket, which consists of a porcelain tube E, fits into the socket
as shown, and this receives and supports the upper chamber with its
electrode and side tube. This electrode is an iron rod, which is made
the cathode of the cell, and is immersed to a proper depth in the
fused electrolyte. The anode F is of carbon, and is supported in a
porcelain chamber provided with a side tube, as shown at G. The
sodium separates upon the lower part of the cathode and floats
THE ISOLATION OF THE METALS SODIUM AND POTASSIUM. 169
upward, where it is allowed to overflow through the side tube and
is collected in a suitable vessel containing a hydrocarbon. An
equivalent of chlorine is set free at the anode and escapes by the side
tube G, where it is either allowed to escape into the atmosphere or
be utilized for the production of a by-product. This particular piece
of apparatus was designed to take a current varying between 30 and
50 amperes, and returns a yield of about 65 per cent of the weight of
FIG. 80. Borcher's Design of Cell for the Production of Sodium and Potassium.
sodium theoretically obtainable. The principal objection to be
urged against this apparatus of Borcher is its lack of durability and
costliness of its parts. It is well known that cast iron has but a
limited existence when subjected to the action of alkaline chlorides at
a red heat. The porcelain tube B, which insulates the cathode com-
partment from the electrolytic cell is rather intricate in its design.
Although it would be out of place here to give such a minute and
detailed account of the evolution of the sodium cell as will be found
in special works treating the subject of electrometallurgy from the
commercial point of view, a brief rev ew of the governing requirements
will be given. In many of these pieces of apparatus, metallic sodium
170 EXPERIMENTAL ELECTROCHEMISTRY.
at high temperature and porcelain are brought into direct contact,
and there must necessarily be a loss of sodium resulting from the
action of the hot metal upon the aluminum silicates of the porcelain.
According to Borcher, the following conditions must be observed in the
successful design of the electrolytic cells for sodium:
1. "A refractory metal only may be used as a material for a
cathode, preferably the better sorts of iron.
2. "The alkaline metal must be collected in, and conveyed
from, the cathode cell without coming into contact with any reducible
substance.
3. "The walls of the cathode chamber may be made to serve also
as cathodes, but in that case they must not be in contact with the
electrolyte on the outer surfaces.
4. "The anode must be made of carbon.
5. "The anode compartment must allow of an easy escape for
the halogen, and its walls must be made of some material that will
withstand the action of the halogens and haloid salts.
6. "The walls of the anode compartment must not be in contact
with the separated metal.
7. "No metallic object must be immersed in the electrolyte in
any position between the poles or in the path of the current.
8. "The whole apparatus must be of a fire-resisting material."
Without dwelling upon the very numerous forms and patterns of
sodium and potassium cells, which have met with more or less success,
we will describe the sodium cell as designed and operated by Castner,
which fulfills the conditions as tabulated above and has proved
itself to be commercially successful. Let us look into the design of
a workable Castner cell on a small scale, and describe its mode of
operation.
By referring to Fig. 81, the student may become familiar with the
design of this cell, which is here illustrated in elevation and section.
This particular experimental cell consists of a large inverted iron
bottle with a rather large elongated neck. An insulating stopper
carrying an iron cathode is passed up through the neck of the bottle,
which rests upon a suitable support in order that the bottle portion
may be heated by a ring-burner. A metal water-jacket is slipped
over the lower part of this neck in order that the insulating stopper
and lower portion of the neck may be kept cold. Caustic soda or
THE ISOLATION OF THE METALS SODIUM AND POTASSIUM. 1 71
potash, as the requirements may dictate, is put into the iron bottle
in a molten condition. That portion of the fluid electrolyte which
runs down into the lower extremity of the neck, kept cold by the
water-jacket, solidifies and forms a seal for the fluid portion of the
electrolyte which is kept at the necessary high temperature by the
FIG. 81. Elevation and Section Through Castner's Cell. Experimental Design.
ring-burner. The anodes which, with the present electrolyte, may
also be of iron, are suspended from the cover, or they may be cast in
one piece of hollow cylindrical form as illustrated in the engraving.
Immediately over the end of the cathode is suspended a little cylin-
drical chamber or receiver for the isolated sodium or potassium.
At the lower extremity of this chamber we have a cylindrical wire-
gauze guard which, because of the high surface tension of melted
sodium, prevents this metal from flowing through, and thereby wander-
ing away from the mouth of the inverted receiver.
For the removal of the fluid metallic sodium, Castner uses a per-
forated ladle, which retains the metal because of its high solution
172
EXPERIMENTAL ELECTROCHEMISTRY.
tension, while the caustic soda drains away through the perforations.
The various parts of the apparatus are insulated by asbestos. With
this apparatus hydrogen gas is evolved at the cathode with the sodium,
accompanied by the expenditure of a certain amount of electrical
energy. We have escaping oxygen gas from the anode compartment.
Very large quantities of metallic sodium are produced upon this
principle. Fig. 82 shows the various parts of such a Castner cell in
FIG. 82, Essential Parts of an Experimental Cell of Castner's Type.
detail, and it will be seen that they may be easily made and put
together by the student in the laboratory. As will be seen there are
only about six or seven important pieces which enter into the con-
struction of this laboratory or experimental cell.
In order to operate this furnace we should have, at least, 50
amperes available and a pressure of about 6 volts. As will be seen
from the following simple calculation, an electromotive force of 4.4
volts is just sufficient to drive a current through a fused sodium-
hydroxide electrolyte. The heat of combination of NaOH is 102
Calories. The minimum pressure necessary therefor is obtained by
dividing the number of Joules represented by 102 Calories, by the con-
stant 96,540. We will remember that the Joule is equivalent to
0.00024 Calorie. 102 divided by .00024 gives us the figure 425,000,
42 5000
from which we obtain the following: =4-4 volts.
96540
HE ISOLATION OF THE METALS SODIUM AND POTASSIUM. 173
It is a matter of interest to note that hundreds of horse-powers
are regularly expended in such sodium cells for the supply of the
market of the world. In our next chapter another process, using a
fused electrolyte and of still greater commercial importance, will be
taken up. This will pertain to the electrolytic manufacture of
aluminum.
CHAPTER XII.
THE ISOLATION OF THE METAL ALUMINUM.
ALUMINUM is distinguished among metals as silicon is among
non-metals for its immense abundance in the solid mineral portions
of the earth, to which indeed it is almost entirely confined, for it is
present in vegetables and animals in so small a quantity that it
can be scarcely regarded as forming one of their necessary compo-
nents. Aluminum, as we know, is an extremely important element,
both in nature and in the arts. It occurs very widely distributed
and very abundantly in many different forms of combination; among
them are feldspar, mica, cryolite, and bauxite.
Feldspar is a silicate of aluminum and potassium, of the formula
AlKSIsOg. Mica is a general name applied to a large number of
minerals which are silicates of aluminum and some other metal, as
potassium, lithium, magnesium, etc. The simplest form of mica is
that represented by the general formula KAlSiO 4 , according to
which the mineral is a salt of orthosilicic acid, Si(OH) 4 . Cryolite
is a double fluoride of aluminum and sodium, or the sodium salt
of fluoaluminic acid, NaaAIFe. Bauxite is a hydroxide of aluminum
in combination with a hydroxide of iron. Besides, in the above
forms, aluminum occurs in the products of decomposition of minerals.
One of the most important of these is clay, which is found in all
conditions of purity from the white kaolin to ordinary dark-colored
clay. Kaolin is the aluminum salt of orthosilicic acid of the formula
Al4(SIO4)3+4H 2 O. Aluminum silicate is found in all soils, but is
not taken up by plants, and does not find entrance into the animal
body. The name aluminum has its origin, in the fact that the salt
alum was known at an early date, and the metal was afterwards
ioslated from it.
All the compounds of aluminum may be derived from the oxide
174
THE ISOLATION OF THE METAL ALUMINUM. 1 75
A1 2 O3 and the hydroxide A1 2 (OH) 6 . From the oxide the sulphide
Al 2 Ss and the salts which contain alumina as the base are derived,
and from the hydroxide the aluminates which are salts containing the
aluminum in the acid radical. The preparation of aluminum on
a large scale has involved a problem of the highest importance to
modern technology. A rough outline of the chemical means for
isolating this important metal may not be out of place. As early
as 1842 Oersted attempted the decomposition of aluminum chloride
through the agency of a potassium amalgam, but the success of his
work is open to considerable doubt, for subsequent workers endeav-
oring to follow his directions were unsuccessful in obtaining any
metallic aluminum. Three years later, however, Wohler successfully
reduced the chloride by using potassium. Deville working at a later
date produced this important metal through aluminum chloride,
by resorting to the use of the double chloride of aluminum and
sodium. Instead of the costly potassium, the far cheaper metal,
sodium, was used. For twenty-five or thirty years this process was
carried on in France, and for a time it was also used in England.
Rose in 1853 proposed the substitution of cryolite for the chloride
and used magnesium in the place of sodium. Grabau's process, which
is of much later date, is of special interest, for it is of unusual merit,
as may be seen from the following equation, where solutions of sul-
phate of alumina are first treated with cryolite to obtain aluminum
entirely as fluoride A1 2 (SC>4) + A1 2 F 6 = 2A1 2 F 6 + 3Na 2 SO 4 . The alu-
minum fluoride being insoluble in water, is filtered off, washed and
dried, and heated to a low red heat, when it is at once charged into a
cold vessel lined with pure cryolite. The required quantity of metallic
sodium is now placed in upon the hot material and the vessel covered.
Accompanied by a great liberation of heat energy the reaction
takes place, which may be represented by the following equation:
2Al 2 F 6 + 3Na 2 =Al 2 +Al 2 F 6 .6NaF. The aluminum, after the reac-
tion has taken place, is recovered melted into a metallic mass at
the bottom of the chamber, but covered over with a slag of cryolite,
which itself has been completely fused through the high temperature
of the reaction. This is a workable scheme for the production of
aluminum, but it is evident at once that the economy of such a process
depends upon the cost of sodium.
The electrical production of metallic aluminum may be divided
176
EXPERIMENTAL ELECTROCHEMISTRY.
into two different schemes, one where the metal is obtained by the
reduction of its oxide and the other by a typical case of electrolysis.
Let us first look into the processes of reduction. For a long time
alumina, which is the oxide of aluminum, was held to be unreducible.
In the electric furnace, however, with sufficient current density, the
oxide may be reduced in the presence of carbon. This electro-
reduction is non-electrolytic, being simply brought about by the
FIG. 83. Laboratory Furnace for the Reduction of Alumina. Experimental Design.
intense heat of the electric furnace. With a sufficient current den-
sity, it appears to-day that no oxide can withstand the high tempera-
ture of the electric furnace. Let us impress upon a small quantity
of alumina this powerful reducing action of carbon at the tempera-
tures accompanying a high current density electric arc. Fig. 83
shows a handy laboratory furnace for accomplishing this. We have
firmly clamped in an iron ring a graphite crucible which also rests
upon the iron base of the ring-stand. A large carbon rod fed through
an opening in a fire-clay cover serves as the other electrode. The
charge of alumina and pulverized carbon is placed in the crucible
THE ISOLATION OF THE METAL ALUMINUM.
177
around a slender conducting pencil of carbon to start the electric
current. A very heavy current is essential to bring about the reduc-
tion. According to Borcher, a current density of about 3500 amperes
per square inch will bring about this reduction. A current of 6500
amperes per square inch is sufficient to reduce any metallic oxide
known. The electromotive force need not be high, 15 volts being
ample for a small furnace. About 53 electrical, horse-power are
therefore necessary. Fig. 84 shows a horizontal furnace for the
FlG. 84. Furnace for the Reduction of Alumina. Practical Design.
reduction of alumina on a somewhat larger scale. In our small
crucible furnace, the carbon pencil can be, of course, less than an
inch in sectional area, permitting the use of a smaller amount of
current. With one-tenth of an inch in sectional area for our carbon
pencil, we can reduce a small quantity of alumina with one-tenth the
current strength. These furnaces are known as resistor furnaces
and are easily constructed and operated. Carbon pencils of
inch sectional area effect a very complete reduction with a current
of 40 amperes. Let us look into the electrolytic methods of isolating
aluminum, and to this end we will first refer to Deville's apparatus.
The following is a translation from Deville's original paper: "Up
to the present time it has appeared to me impossible to obtain alu-
minum from aqueous solution by means of a galvanic battery; and
I should even now believe in the absolute impossibility of doing so
if the brilliant experiments of Bunsen in the production of barium,
chromium, and manganese had not shaken my convictions. How-
ever, I am compelled to say that all the processes of this kind which
have been published recently in reference to the preparation of
178 EXPERIMENTAL ELECTROCHEMISTRY.
aluminum have given me only negative results. Every one knows
the beautiful process by means of which Bunsen has produced
magnesium by decomposing magnesium chloride with the aid of a
galvanic battery. The illustrious Professor at Heidelberg has
opened a way which may lead to results that will be interesting from
many points of view. However, there can be no hope of applying
the battery to the direct decomposition of aluminum chloride,
which is a substance which does not fuse, but that volatilizes at a
low temperature; it is nececsary, therefore, to find a composition
for the metallic bath that shall involve the use of a fusible material,
from which aluminum alone can be deposited by the electric current.
I have found such a substance in the double chloride of aluminum
and sodium, the production of which is a necessary feature of the
extraction of aluminum by sodium. This chloride, which is fusible
at about 185 C., and remains fixed at a sufficiently high tempera-
ture, although it is volatile at a temperature above the fusing point
of aluminum, fulfills all the required conditions. I introduced this
substance into a porcelain crucible, which was imperfectly separated
into two compartments by a plate of biscuit porcelain, decomposed
it by means of a battery of five elements, using carbon electrodes,
the crucible being heated and the temperature being increased con-
tinually in order that the charge might be maintained in a fluid con-
dition as it became gradually less and less fusible; but the fusing
temperature of aluminum was not exceeded. Arrived at this point,
I stopped the experiment, and, after lifting out the diaphragm and
the electrodes, I heated the apparatus to a bright red heat, and found
at the bottom of the crucible a regulus of aluminum, which was
rolled and was exhibited to the Academy at its meeting on March
20th, 1854. It was accompanied by a considerable quantity of
carbon, which had prevented a notable portion of the metal from
uniting into a single mass. This carbon resulted from the disinte-
gration of the very dense sample of retort carbon that served as
electrode; and as a result of this action the positive electrode was
entirely eaten away in spite of its thickness, which was very con-
siderable. This disposition of apparatus (as used by Bunsen for
magnesium) was not convenient in the case of aluminum; and the
process to which I have been led, after many experiments, is as
follows: The aluminum bath is prepared by weighing 2 parts of
THE ISOLATION OF THE METAL ALUMINUM.
179
aluminum chloride and adding to it i part of marine salt in the
state of dry powder. The whole is mixed in a porcelain crucible
heated to about 200 C. Combination shortly sets in with evolution
of heat, and there results a very fluid mixture, which is a bath used
for the decomposition.
"The apparatus [as shown in Fig. 85] consists of a glazed porcelain,
crucible A, which, as a measure of precaution, is placed within the
somewhat larger fire-clay crucible B] the whole is surmounted by a
FlG. 85. Deville's Cell for Experimental Work.
crucible cover C pierced with a slot D, through which is placed a,
wide and stout sheet of platinum E to serve as negative electrode,
and with an aperture in which is tightly fixed a well-dried porous
cell F. Within the latter is placed a rod of retort carbon G as positive
electrode. The bottom of the porous cell should be kept at the
distance of some centimeters from that of the porcelain crucible.
The porcelain crucible and the porous cell are filled to the same
level with the fused aluminum-sodium chloride, and the apparatus
is heated after the manner described. The electrodes are then
introduced and the current is passed through the apparatus. Alumi-
num is deposited with some sodium chloride upon the platinum plate,
and chlorine together with some aluminum chloride is disengaged
l8o EXPERIMENTAL ELECTROCHEMISTRY.
in the porous cell; fumes are thus produced which are destroyed by
introducing dry and powdered marine salt at intervals into the
porous cell. This salt is transported to the negative pole during the
operation along with the aluminum. A small number of elements
(two are actually sufficient) are required to decompose the chloride,
"which presents only a feeble resistance to the electric current.
"The platinum plate is raised from time to time as it becomes
sufficiently charged with metallic and saline deposit. It is allowed
to cool, the mass of salt is rapidly broken, and the plate is replaced in
the circuit. The crude material detached from the electrode is fused
in a porcelain crucible enclosed within a fire-clay crucible. After
cooling, the mass is treated with water, which dissolves a large quantity
of sodium chloride; and a gray metallic powder is left, which is
reunited into a regulus by several successive fusions. Addition of
double chloride of aluminum and sodium is necessary during each
fusion."
This process of Deville may be said to have marked the dawn
of the successful processes dependent upon the electrolysis of fused
aluminum compounds. As it would require a great deal more
space than we are able to give to the subject in a general work, to
enumerate and describe the various steps in the evolution of alumi-
num reduction cells, we must content ourselves with an outline of the
more important processes. Let us look into the method of Hall,
and for this purpose we may best turn to the specifications and draw-
ings of one of his patents. The following is from the patent of Hall
of 1886, and reads as follows: "The invention described herein
relates to the reduction of aluminum from its oxide by dissolving
such oxide in a bath containing a fused fluoride salt of aluminum,
and then reducing the aluminum by passing an electric current
through the bath, substantially as hereinafter more fully described
and claimed. In the accompanying drawings, [Fig. 86] represents a
sectional elevation of a form of apparatus applicable in the practice
of my invention, and [Fig. 87] is a view partly in elevation and partly
in section of a modified form of apparatus.
" In the practice of my invention I prepare a bath for the solution
of the aluminum by fusing together in a suitable crucible, A, the
fluoride of aluminum and the fluoride of a metal more electro-
positive than aluminum, as, for example, the fluoride of sodium,
THE ISOLATION OF THE METAL ALUMINUM.
lol
potassium, etc., these salts being preferably mingled together in
the proportions of 84 parts of sodium fluoride and 169 parts
fluoride, represented by the formula Na 2 F 8 . A convenient method
of forming the bath consists in adding to the mineral cryolite - of
its weight of aluminum fluoride. The object of thus adding alu-
FIG. 86. Hall's Electrolytic Cell and Furnace for the Production of Aluminum.
minum fluoride is to secure in the bath the proper relative propor-
tions of the fluorides of aluminum and sodium. To the fused bath
is added alumina, or the oxide of aluminum, in sufficient quantities,
and the alumina being dissolved by the fused bath an electric current
is passed through the solution by means of suitable electrodes, C and
Z>, connected with a dynamo-electric machine or other suitable
source of electricity, and immersed in the solution. By the action
of the electric current, which preferably has an electromotive force,
of about 4 to 6 volts, oxygen is released at the positive electrode, C,.
and aluminum is released at the negative electrode, D, which, on
account of the affinity of aluminum for other metals, is formed of
182
EXPERIMENTAL ELECTROCHEMISTRY.
carbon when it is desired to produce pure aluminum. The positive
electrode may be formed of carbon, copper, platinum, or other
suitable material. When formed of carbon the electrode, C, is
gradually consumed, and must therefore be renewed from time to
time; but when formed of copper an oxide coating is formed over
FIG. 87. Experimental Cell of Hall.
the surface of the electrode. This coating serves to protect the
electrode from further destruction by the action of the oxygen, but
does not interfere materially with the conducting qualities of the
electrode.
"On account of the affinity of the aluminum for other metals,
and also the corrosive action of the materials, I prefer to form the
crucible, or melting-pot A, of metal as iron or steel and protect
the same from the action of the aluminum by a carbon lining, A.
This crucible is placed in a suitable furnace, B, and subjected to a
sufficient heat to fuse the materials placed therein, such materials
fusing at approximately the same temperature as common salt.
THE ISOLATION OF THE METAL ALUMINUM. 183
"In lieu of the electrode D [Fig. 86], the carbon lining, A, may be
employed as the negative electrode, as shown in [Fig. 87], the con-
ductor from the negative pole of the electric generator being suitably
connected, as shown at N, to such lining.
"In order to render the bath or solvent more fusible, fluoride
of lithium may be substituted for a portion of the fluoride of sodium ;
as, for example, for one-fourth the fluoride of sodium an equivalent
amount of lithium fluoride by molecular weights may be substituted.
Thus 26 parts of lithium fluoride displacing 42 parts of sodium
fluoride, the resulting combination contains 26 parts of lithium
fluoride for every 126 parts of sodium fluoride, and 338 parts of
aluminum fluoride.
"While I consider the proportions of fluorides of sodium and
aluminum, and of the fluorides of sodium, lithium, and aluminum
hereinbefore stated, are best adapted for the purpose, such propor-
tions may be varied within certain limits without materially affecting
the operation or function of the bath, as in fact, any proportions
which may be found suitable may be employed. The aluminum,
as it is reduced at the negative electrode, is melted and collects
thereon in globules, and then drops down to the bottom of the bath,
which is of lower specific gravity than the molten aluminum, and
can be removed by suitable means; or the bath may be poured out,
and after being cooled the aluminum can be picked out."
Fig. 88 shows the scheme finally adopted by Hall for the pro-,
duction of aluminum by the electrolysis of fused salts. According
to this later patent the following bath is employed: "Fluoride of
calcium, 234 parts; cryolite, the double fluoride (Na 6 Al 2 Fi 2 ), 421
parts; the fluoride of aluminum, 845 parts, by weight, and about
3 to 4 per cent of a suitable chloride, e.g., calcium chloride. Alumina
is then added to this bath, preferably in sufficient quantities to form
a saturated solution. Electrodes are then inserted in the bath, the
negative electrode being formed of carbon when pure aluminum is
desired. The .positive electrode may be formed of carbon or other
suitable material. This piece of apparatus, as the drawing indicates,
consists of an iron trough lined with plates of carbon. A heavy
copper connector is riveted to the outside of this trough and connects
with the negative lead of the dynamo. The iron trough, therefore,
with its carbon lining serves as a cathode. The anodes consist
1 84
EXPERIMENTAL ELECTROCHEMISTRY.
of carbon rods suspended from a heavy copper bar in such a way
that they may be lowered into the electrolyte. As these carbons
are slowly consumed by the oxygen liberated in contact with them,
it is necessary to have such an adjustment for feeding them into the
bath. It is interesting to note that with such a system the weight of
carbon burned by the nascent oxygen is about equal to the metallic
aluminum produced. The carbon lining is but very slowly altered
and lasts for a long time. The electrolysis consists in the decom-
position of alumina dissolved in the fused bath. Several such
troughs in practice are connected up in series for the commercial
production of this important metal. In this plan of operation the
FIG. 88. Hall's Approved Aluminum Cell.
heat required to keep the electrolyte fluid is derived from the elec-
trolyzing current. Immense quantities of aluminum are produced
upon this general principle. The next device which we will consider
is that of Heroult, as depicted in Fig. 89. This process may produce
either metallic aluminum or aluminum bronze, the latter being
an- alloy of aluminum and copper. In this piece of apparatus the
electrolyte consists of alumina dissolved in fused cryolite, or in an
artificial mixture of aluminum fluoride with sodium fluoride. The
electrolyte is maintained in the molten state by the heat generated
by the passage of the electrolyzing current. In the illustration we
have a heavy iron vessel lined with carbon plates, with tap-hole
for allowing the molten alloy to be drawn off. A heavy carbon cover
with a suitable opening for the anodes is put in place to prevent
undue loss of heat from radiation. The anodes are connected as
shown to a common holder-bar, and are of carbon hung in such a
THE ISOLATION OF THE METAL ALUMINUM.
185
way that they may be lowered into the electrolyte as they are con-
sumed. The operation of the process is started by placing some
pure copper in the bottom of the furnace and lowering the anodes
until they come in contact with the metal. The intense heat of an
arc so established fuses the copper, when the electrolyte is added,
which, of course, rapidly assumes a fluid state. Cryolite is added
until a sufficient depth has been attained and the electrolysis of the
FlG. 89. Heroult's Aluminum Cell.
alumina which is dissolved therein takes place between the molten
copper as cathode and the immersed carbon rods as anodes. Alu-
minum separates at the molten copper and alloys with it, the product
being allowed to run off at intervals. Additional quantities of copper
and alumina are fed in in the proper proportions from time to time to
form a suitable alloy. It may be seen that the furnace and its opera-
tion are equally well adapted for the production of pure aluminum,
if some of this metal is placed in the bottom of the furnace instead
of copper, when the process is first started, which then approaches
very closely the device of Hall.
CHAPTER XIII.
THE ISOLATION OF CALCIUM.
CALCIUM is more generally met with in a state of chemical com-
bination than any other metal, for it occurs in enormous quantities
in limestones and chalks, and in the minerals gypsum, fluorspar,
apatite, etc., etc. Yet, notwithstanding its great abundance in com-
bination, because of the great difficulties in isolating it, the metal is
exceedingly scarce and seldom met with in the laboratory. Let
us look a little into the properties of this interesting metal, and the
chemical methods of preparing it, before undertaking to obtain it
by electrolysis, which is by no means an easy task. Calcium is
usually described in the text-books as a brass-yellow metal of lustrous
appearance, which in moist air soon becomes covered with a coating
of calcium hydroxide and calcium carbonate. Calcium decom-
poses water just as potassium and sodium do, but the heat of the re-
action is not sufficient to set fire to the hydrogen which it evolves.
Calcium, because of the expense and difficulty in obtaining it, has
never had any useful application except in the most special cases
on very small scales in some research work. Before the electrolytic
method of isolating calcium, it was obtained at a great cost by purely
chemical means. The following method will be of interest to the
student in general chemistry and will serve as an introduction to
the electrical means. For example, calcium may be obtained as
a finely divided substance by heating powdered lime with powdered
metallic magnesium, as shown by the following simple equation:
CaO+Mg = MgO + Ca.
After the reaction has taken place, the presence of calcium may be
shown by adding a little distilled water to the mass in a tube, when
the free calcium reacting with the water will set hydrogen free, and
186
THE ISOLATION OF CALCIUM. 187
in addition, the presence of lime-water in the tube may be shown.
Another chemical method consists in making a zinc calcium and
distilling off the zinc by heating to a high temperature in a gas-retort
carbon crucible. The zinc calcium is made by melting together
a mixture of calcium chloride, zinc, and metallic sodium. The
sodium decomposes the chloride, and the reduced calcium dissolves
in the zinc as soon as it is liberated. Metallic calcium when heated
to redness burns with a very brilliant white light and is converted
into its oxide. Modern text-books on general chemistry state that
it is obtained to-day by electrolyzing its fused chloride, but they
do not, of course, undertake to dwell upon the details of the opera-
tion or to even state the great difficulties involved. The electrolytic
isolation of metallic calcium is far from easy, and its electrolytic
preparation is introduced here to show the student in electrochemistry
that his skill and resources will often be taxed, and at times very
severely.
As intimated, therefore, while theoretically simple, the extraction
of calcium from its chloride by electrolysis is attended by extreme
difficulty as compared with certain other metals. It may be stated
in general that the three metals of the alkaline earths are exceedingly
difficult to isolate. Bunsen and Matthiessen were the first to isolate
these metals from their chlorides in their pure condition. Failure
after failure attended their efforts to extract these metals in Bunsen's
laboratory by electrolyzing their chlorides, using apparatus such as
we described in a previous chapter for the production of metallic
magnesium from its chloride. In the year of 1854, however, Bunsen,
as the result of his carefully conducted research, was able to point
out the reason for the failures. The following interesting and note-
worthy observation was made by Bunsen, which should be impressed
upon the students of electrochemistry, as Bunsen's discovery has
great bearing upon many cases in this field. We will quote his
own words, which are as follows: "The density of the current used
for electrolysis that is, the ratio of current volume to electrode
area exerts a most important influence on its chemical effects. The
power of the current to overcome affinities increases with this density.
Of no less importance is the relative mass of the constituents of the
electrolyte through which the current passes." To put Bunsen's
suggestion into practical operation we must have, therefore, means
1 88 EXPERIMENTAL ELECTROCHEMISTRY.
for electrolyzing calcium chloride and a suitable container for the
same with adjustment for high current density. It has been found
in experiment and practice, that exceedingly high current density
at the cathode is absolutely essential for the isolation of calcium.
In order to secure conditions of exceedingly high cathode cur-
rent density special designs of apparatus, are necessary. High
cathode current density involves small cathode area, and a current
which gives rise to very high temperature, even beyond the melting-
point of iron and steel. It is, therefore, necessary to provide some
means for keeping the cathode cold, or comparatively cold. Means
must also be provided for preventing, as far as possible, the recom-
bination of the liberated chlorine with the freed calcium. The
device as illustrated in Fig. 90 embodies these set requirements
upon an experimental scale. Here we have a small calcium reduc-
tion-furnace illustrated in both elevation and section. The chamber,
for the calcium chloride to be electrolyzed, consists of a large graphite
crucible, not less than 5 or 6 inches in diameter, with the bottom
sawed off, giving us in reality a large graphite collar. This graphite
crucible has clamped to its exterior a heavy iron band and serves
as an anode in the operation. This bottomless crucible, or collar,
rests upon and may be cemented to a disk of mica, which in turn
rests upon a cylindrical water-bath, as shown. This water-bath has
a tube soldered within its center, which receives with a tight fit
the turned rod or bar of iron not less than an inch in diameter.
This rod is long enough to go up through the bottom of the bath and
attached mica covering, which forms the bottom of the graphite
crucible, extending a couple of inches below and having a stout
clamp of iron connected with the lead of the dynamo. The cathode
proper consists of a piece of steel wire about 1/16 of an inch in
diameter and about 3 inches in length, which is securely screwed
into the upper end of the iron bar. When the water-bath is filled
with cold water and means provided for a continuous circulation
through it from a spigot, the cathode wire may give up its heat by
conductance down into the iron bar, which in turn gives up its
acquired heat to the circulating water. In this way it is prevented
from getting too hot. A cylinder of platinum- wire gauze about i
inches in diameter goes over the cathode wire, as shown, to prevent
the liberated calcium from wandering about in the electrolyte. To
THE ISOLATION OF CALCIUM.
189
put such a furnace in operation it is assembled, with the exception
of the platinum-gauze cylinder and cover, when the cathode bar is
put in proper electrical connection with a suitable dynamo. The
dynamo should be capable of giving at least 100 amperes at a pres-
1 90 EXPERIMENTAL ELECTROCHEMISTRY.
sure of about 60 volts. The calcium chloride in lumps is slowly
added, a small quantity at a time, and by means of an iron rod a
small arc is established between the cathode wire and the side of
the crucible until a small quantity of the calcium chloride has been
fused down to a fluid state, which will then conduct the electric
current. More chloride is added until the crucible is about three
fourths filled. If means are at hand for melting a sufficient quantity
of calcium chloride separately, and pouring into the crucible, the
process may be more quickly put in operation. When the crucible
is filled with fluid electrolyte, the temperature is easily maintained at
the melting-point of the substance by the current, and then the plat-
inum-wire cylinder should be put in position. The calcium will sepa-
rate in little globules from the cathode and be retained within the plat-
inum-gauze cylinder, while the chloride will escape from walls of the
crucible, which act as anode, and pass out through the covering.
The author is describing a successful run upon this design of furnace.
It is very easy to have an unsuccessful run, and if the conditions are
not just right the attempt will result in failure. For instance, it
is not really possible, so far as the experiments of the present writer
go to show, to isolate calcium with less than 60 amperes, with a
design and dimensions similar to those given. If, on the other
hand, too heavy a current is used, the steel- wire cathode will, in spite of
its connection with the mass of iron in the water-bath, rise so rapidly
in temperature that it will melt off. On the other hand, if the cur-
rent is not strong enough to keep the entire mass in fusion, a solid
crust of calcium chloride will form on top of the molten chloride
and offer a resisting seal to the chlorine gas, which is being given
off from the lower portion of the crucible walls. Should such a
crust form it should immediately be punctured to allow the chlorine
to escape, or it will lift the crucible with almost explosive violence
from the mica disk, even if it has been securely cemented down,
and a stream of fluid electrolyte will be forced out. With everything
working smoothly, the calcium may be ladled out from the interior
of the platinum-wire cylinder by means of a small iron spoon per-
forated to allow the fused electrolyte to run through, the calcium
remaining within the ladle, because of its high surface tension.
This calcium is liable to take fire in the air and burn with a fierce
white light, and a suitable hydrocarbon in a wide-mouth vessel
THE ISOLATION OF CALCIUM.
IQI
should be ready in which to immediately plunge the liberated metal.
By referring to Fig. 91 the separate essential parts of such a laboratory
furnace may be seen, which are, namely: a water-bath, insulating
mica disk with a small hole through its center, just the size of the
cathode wire, and a platinum-wire-gauze cylinder together with
a heavy iron bar with its cathode screwed in, and a bottomless
crucible. As the calcium industry is comparatively unimportant,
FIG. 91. Essential Parts of Experimental Furnace for the Isolation of Calcium.
we will content ourselves with only one more design of calcium
furnace, and for this purpose we will turn to Fig. 92. This design
we owe to Borscher, as given here in elevation and section. The
outer casing is in the form of a long thimble and may be of almost
any convenient size. This thimble serves as anode, which may
be of iron, brass, or nickel. In the bottom of this elongated thimble
tube is placed a small porcelain crucible of such a size that it will
just slip within the tube. The cathode is a piece of steel wire between
1/16 and 1/8 of an inch in diameter and about i inch in length,
screwed into a concave end of a similar tube which is supported
by an insulating collar as shown. Within the center of this tube,
which may be supported by a middle collar brazed or soldered
in position, is the cooling water, which falls directly upon the end
I9 2 EXPERIMENTAL ELECTROCHEMISTRY.
of the tube carrying the cathode wire, and discharging from the outlet
tube at the right, as shown. With such a device the inner tube
is kept cold, giving up the heat generated at the cathode wire
which it supports. The insulating collar or support is provided
with a small side tube for the escape of chlorine gas. To put the
furnace in operation the little porcelain crucible is dropped within
the elongated thimble, and the whole tube is filled about two
FIG. 92. Experimental Calcium Furnace. Borscher's Design.
thirds full of calcium chloride fragments, which may be melted down
by holding the tube in the flame of a Bunsen burner. When in the
fluid state, the electrolyte receives the water-jacket tube and cathode
wire, which, being at a low temperature, immediately chills the cal-
cium chloride to the point of solidification. This has been indicated
in the drawing by the white mass surrounding the water-chamber.
The cell is immediately placed in circuit with the electrical supply,
when the Bunsen burner may be removed and the temperature of
electrolysis maintained by the passage of the current. Calcium iso-
lates from the steel- wire cathode in small globules, and if it rises
it is caught in the concave end of the water-jacket. The little
THE ISOLATION OF CALCIUM. 193
porcelain crucible serves a double purpose, namely, in catching any
metallic calcium which may fall if specific gravity conditions of the
electrolyte so induce, but more especially to prevent any chlorine
gas from rising and reuniting with the isokted calcium. It will
be at once appreciated that no chlorine will be liberated from the
interior of this porcelain crucible, for being of a non-conductible
material it does not act as an anode. There is, in consequence, no
chlorine given off which may reach the cathode at a point lower
than the upper edge of this crucible. This device of Borscher is
one intended for producing small quantities of calcium and must
be directly taken apart in order to secure such fragments of metal
isolated. It is, nevertheless, a furnace of neat design, and very
useful for experimental work upon a small scale.
CHAPTER XIV.
THE ELECTRIC FURNACE AND FURNACE PRODUCTS.
ELECTRIC furnaces may be roughly classified into two general
kinds, those for the attainment of moderate temperatures and those
for the attainment of the highest temperatures within the reach
of man. By moderate temperature, we may consider furnaces
capable of running up to about 1500 or 1600 C. It is this first
type of furnace which will now occupy our attention. Both general
types of electric furnaces are converters of electrical energy into heat
energy and both types depend upon resistors. The resistors in the
type of furnace which we will consider first consist of platinum
wire, and because of the facility and ease with which the temperature
may be regulated, find a most useful place in all chemical and elec-
trochemical laboratories. Having had considerable success with
the furnace here depicted it is deemed of value to describe more or
less in detail the method of assembling and constructing such small
furnaces of a great range. of general utility. Fig. 93 illustrates a,
side view of a small muffle furnace which may be very easily put
in operation and regulated. It consists, as shown, of an iron retort
stand with a clamp holding in a horizontal position a fire-clay tube
which is wound with platinum wire, having slipped over the platinum
wire winding a second somewhat larger fire-clay tube. The terminals
of the platinum wire are connected with a lamp-bank, together
with a 1 10- volt or 220- volt electric lighting system. In order to
make the construction of this furnace clear, we will turn to Fig. 94..
where the smaller tube is shown turned down and spirally threaded
to receive the platinum-wire winding. It is over this portion of the
fire-clay tube that the outer jacket is placed. There are numerous-
kinds of fire-clay, which, before baking may be turned on the lathe
with facility and alter but little in shape and dimension after the
194
THE ELECTRIC FURNACE AND FURNACE PRODUCTS.
firing process. In this last figure referred to, an end view of the
furnace is given, together with a side elevation of the muffle tube.
A good size for. such a furnace is to have the inner tube, upon which
the platinum wire is wound, about 5 inches long with an internal
diameter of about i| inches, and an external diameter of about 2
inches. Where the tube has been turned down and the spiral cut
FIG. 9^. Elevation of Muffle Furnace and Lamp-bank.
on, the thickness should not be over 3/16 of an inch. Of course,,
these dimensions may be varied to meet different requirements,,
but for studying the behavior of certain bodies at different tempera-
tures only a small quantity of these bodies is necessary for the
examination. At least two meters of platinum wire should be
wound on a single furnace, the wire to be about number 22 gauge.
It will require some preliminary experiment with each furnace in
connection with the lamp-bank and suitable ammeter in series to
196
EXPERIMENTAL ELECTROCHEMISTRY
ascertain how much current the furnace will stand with the out
side muffle in place. It is, of course, an easy matter to burn out
3
the platinum wire, but such a mishap is not usually very serious,
for the wire is apt to fuse at a certain point, when it may, of course,
THE ELECTRIC FURNACE AND FURNACE PRODUCTS.
197
be quickly welded together again. After a mishap of this kind
the experimenter will be in possession of valuable data in connection
with his furnace if he has slowly admitted the current to it through
a suitable ammeter. By referring to Fig. 95 a sectional view through
3Q8
EXPERIMENTAL ELECTROCHEMISTRY.
a. completely assembled furnace may be seen. The furnace is here
indicated with a fire-clay plug in one end and a similar plug of
light design at the other end. The furnace may be brought to
bright incandescence within a very few moments after the current
is turned on, and the temperature may be held between that of the
laboratory and the melting-point of platinum with great precision.
The temperature within the muffle tube can be ascertained, of course,
by the method depending upon the melting-point of pure metals, or
by exploring the interior of the furnace with a platinum loop carry-
ing an electrical current, in connection with the proper electrical
FIG. 96. Top View of a Vertical Type of Wire Resistor Furnace.
instruments. A hint of practical value may not be out of place
at this time. With certain kinds of fire-clay material there is
shrinkage upon the first firing, and the platinum wire, because of
its own expansion by heat, is apt to come out of the threaded groove,
causing the separate convolutions of the adjacent coils to come
into contact and cause serious trouble by short circuiting. This
may be avoided by rewinding the furnace after the shrinkage has
taken place, winding on the platinum wire quite tight, when there
"will be no more difficulty from this source.
Another design of furnace of the platinum-wire resistor type,
is depicted in Figs. 96 and 97. The former being a top view looking
down into the furnace which is of a vertical type. Here we have
THE ELECTRIC FURNACE AND FURNACE PRODUCTS.
199
a number of small tubes or pipe-stems arranged around the interior
of a thick outer casing of fire-clay and held in position by a fire-
clay plug fitting snugly in between them and by the lacing back
and forth of the platinum-wire resistor. The second illustration
here shows a section through the vertical type of furnace. With this
FIG. 97. Section Through a Vertical Type of Wire Resistor Furnace.
design the platinum wire may expand without the slightest danger
of short circuits being formed, because it is entirely enveloped within
these small vertically and cylindrically arranged pipe-stems or
tubes. In experimenting with the platinum wire winding of an
electric furnace to ascertain its maximum current-carrying capacity,
EXPERIMENTAL ELECTROCHEMISTRY.
it must be borne in mind that the maximum current-carrying ca-
pacity of a platinum wire is very dependent upon its surrounding
conditions. If we ascertain the current-carrying capacity of the
platinum winding without the muffle tube in place and then give the
coil the same current after covering with the muffle tube, we will,
without question, burn out the furnace, for the reason that the heat
can no longer dissipate so freely. It has been found in practice
with a furnace of the general character of either of the foregoing
designs that a platinum-wire coil will stand only about one half as
much current when the muffle is in place as it did when freely
exposed to the air. One must regard these resistor furnaces purely
as converters, as stated at the opening of the chapter, and we
must maintain our furnace in operation at such a point, where
the energy supply as electricity is carried off as heat energy, the
balance between the supply on the one hand of electrical energy and
the liberation of heat energy on the other hand, taking place within
the limit of the melting-point of platinum. Because of the melting-
point of platinum, this design of furnace is limited for work under
about 1600 C. But this vertical type may be so modified as to
allow of the temperature being carried up to the very melting-point
of the fire-clay itself, enabling the experimenter to melt down platinum,
gold, iron, and steel. For this purpose this vertical pipe system
of tubes must be packed with finely granulated carbon, and instead
of being connected in series must be joined in multiple-arc by con-
necting all the lower ends together by means of a carbon disk, as
well as the upper ends by means of a similar carbon disk. Such a
modified furnace will, of course, require a very much heavier cur-
rent to operate it, but the suggestion is made here for the benefit of
those who may wish to experiment with small muffle furnaces at
exceedingly high temperatures. For the production of electrical
products requiring extremely high temperatures a furnace of very
different design must be employed. Although at the opening of
this chapter the author classified all electrical furnaces on the prin-
cipal of suitable resistors, the type of furnace we are about to describe
is sometimes considered to be of the arc type. Although we may
have an arc it may still be maintained that this type of furnace is
on the resistor principle, for in the present case 'the resistor con-
sists of a stream of incandescent gaseous carbon. A very convenient
THE ELECTRIC FURNACE AND FURNACE PRODUCTS 20 1
laboratory or lecture-room furnace for the production of such bodies
as calcium carbide, is illustrated in elevation by the photograph
constituting Fig. 98. With this practical design of furnace, which
was gotten up several years ago by the author, calcium carbide
202
EXPERIMENTAL ELECTROCHEMISTRY.
may be produced on the lecture table in a very few minutes by
drawing only about 20 amperes from a no- volt lighting system.
Fig. 99 represents a sectional view through this furnace, but
it is only by referring to Figs. 100 and 101 that the peculiar design
THE ELECTRIC FURNACE AND FURNACE PRODUCTS. 203
of the furnace is understood. Here we have 6 carbon electrodes
so connected that we have three electric arcs in series, allowing of
the use of this furnace on incandescent lighting systems, protected
only by a fuse of moderate capacity. It occurred to the writer a
number of years ago in wiring a couple of arc lamps across the
feeders of an incandescent lighting system, that a small experimental
electric furnace could be gotten up on this plan. Only four carbons
FlG. 101. Diagram Showing Connections and Mode of Controlling Series
Carbon Furnace.
were employed at first on the principle of the two arc lamps in
series, but it was found upon the addition of such a charge of lime
and coke as is utilized in the production of calcium carbide that
it had a decided short circuiting effect and allowed too heavy
a current to flow through the furnace. An additional pair of carbon
electrodes were then added, giving three arcs in series instead of two.
With such a furnace a suitable charge for the production of cal-
cium carbide may be employed, but the resistance to the passage
of electric current is not that offered by the three small arcs, but
considerably less, as will be appreciated by any one familiar with
electricity. The resistance is sufficiently high, nevertheless, to
204 EXPERIMENTAL ELECTROCHEMISTRY.
enable one to produce calcium carbide in considerable quantities
on 20 amperes of current.
Fig. 101 illustrates the connections of these series carbon furnaces
FlG. 102. Vertical Type of Furnace with a single Arc.
coupled with a variable rheostat for controlling the intake of cur-
rent by the furnace. If, however, ample current is at hand,
say 50 or 60 amperes, a double crucible furnace, like that illustrated
in Fig. 102, is found very convenient. Here we simply have two cru-
THE ELECTRIC FURNACE AND FURNACE PRODUCTS. 205
cibles, one within the other, separated by some good non-com-
bustible heat insulating material. With this equipment, calcium
carbide, carborundum, etc., may be prepared on a small scale.
For the reduction of metallic oxides in the presence of carbon this
type of furnace is most convenient. For the production of quanti-
ties of calcium carbide on a small scale on the lecture table the
following directions should be carefully followed: Good unslacked
lime and hard carbon are weighed out in the requisite combining
proportions. The following equation indicates theoretically the pro-
duction of calcium carbide, and from the same, the amounts to be
weighed out may be learned'
CaO + C 3 =CaC 2 + CO.
By hard carbon, it is meant that charcoal be not employed, for be-
cause of its lightness it is apt to bum away without combining
with the calcium of the lime. A convenient and most satisfactory
carbon is obtained by crushing up in a large mortar fragments of
old electric light carbons. Both the carbon and the lime should
be ground to a fine granulation and intimately mixed together, and
for this purpose the writer has found an old iron coffee-mill to meet the
requirements in a most satisfactory manner. After a run of half an
hour at a full incandescent temperature, fragments of calcium car-
bide will be obtained as large as an English walnut, which yield a
large supply of gas, and may be burned in a large jet if thrown
into a cylinder jar of water equipped with a small glass outlet tube
and tight fitting stopper. As we experiment with hydrogen the
jet should not be lighted until one is sure that all of the air has been
driven from the cylinder. It must be remembered that the
preparation of calcium carbide requires extreme temperature. The
furnace must be allowed to be well under way in temperature before
the timing of the run is begun. An electric furnace of this type is
nothing more or less than a box of poor heat conducting material,
in which electrical energy is poured, so to spe,ak, until the entire
interior assumes a temperature of the electric arc. The temperature
of the arc has been carefully computed by many experimenters
and is found to be in the neighborhood of 3500 C., which is 6332 F.
We will not go into the commercial question of electrical iurnaces
206 EXPERIMENTAL ELECTROCHEMISTRY.
here as it would constitute a treatise in itself. We are, nevertheless,
fully enabled at this time and at this stage of our work to conduct
an efficiency research upon electrochemical processes of this kind,
if we have profited by the previous chapters on the theoretical side
of our subject. Believing that we have here outlined the simplest
types of electric furnace available for experimental work, we will
take up another subject in the next chapter.
CHAPTER XV.
PREPARATION OF ORGANIC COMPOUNDS.
THE ELECTROLYSIS OF SODIUM ACETATE.
IN the present chapter we will produce electrolytically certain
organic chemical compounds, and we will start our work by taking
a typical case of organic electrolysis and one of peculiar beauty
for demonstration purposes, as a combustible gas is set free at each
electrode. To accomplish this we will electrolyze a strong solution
of sodium acetate: CHs.COONa in the assembled apparatus, as
illustrated in Fig. 103. Here we have at A the electrolytic cell,
which consists of a wide mouth glass cylinder, carrying a large
rubber stopper. Through the center of this rubber stopper is a
glass tube of large diameter terminating in a bell mouth, as shown
at B. This glass tube, which is somewhat the shape of a lamp-
chimney, is provided with a tight fitting stopper at its top through
which passes the wire attached to the anode and the glass tube C,
leading to the wash-bottle F and gas-collecting tube D at the extreme
left. Immediately under the bell-mouth opening of the tube B is
a cylindrical porous, pot E containing the anode G, which is of
platinum. The cathode H is a large cylinder of pure sheet copper
surrounding the porous pot E and leaving considerable space for
electrolyte between it, the walls of the outside containing vessel
and the porous pot. The tube / passes through the stopper of
the outside container, the electrolytic cell A and runs to the bottom
of the wash-bottle /, thence to the vertical gas collecter K. The
wash-bottle J contains plain water, whereas wash-bottle F contains
lime-water for the absorption of carbon dioxide. The terminals
of the electrolytic cell are connected to the lighting system through
our lamp-bank and two or three i6-c.p. lamps will suffice for the
207
208
EXPERIMENTAL ELECTROCHEMISTRY.
current. At the cathode we will have two atoms of hydrogen
isolated, as a result of the setting free of sodium, as shown by the
secondary reaction in the following equation :
2Na 4- 2 H 2 O = 2 NaOH + 2 H.
PREPARATION OF ORGANIC COMPOUNDS. 209
This hydrogen will, of course, be collected in the tube K after pass-
ing through the wash-bottle /. At the anode the following processes
take place :
2 CH 3 .COO +H 2 O= 2 CH 3 .COOH + O,
2 CH 3 .COOH + O = C 2 H 6 + 2CO 2 +H 2 O.
The gas ethane C 2 H 6 , and carbon dioxide CC>2 escape through
the tube C. The carbon dioxide is absorbed by the lime-water
with the formation of calcium carbonate, whereas the ethane col-
lects in the vertical tube D. It is interesting to note that the volume
of ethane is about the same as the volume of hydrogen produced.
According to Jahn, we would obtain more ethane if it were not
for the oxidation of some of the acetic acid at the anode by the
oxygen, as suggested by the following equation:
CH 3 .COOH -f 4 O = 2 CO 2 + 2 H 2 O.
The hydrogen and ethane produced may be ignited to show
that both gases are of a combustible nature.
THE ELECTROLYTIC PRODUCTION OF IODOFORM.
For the production of iodoform we will require a beaker of
about 500 cubic centimeters capacity, a cylinder of nickel wire gauze
to serve as cathode, a porous pot and a suitable platinum anode.
The beaker is to be mounted upon a tripod in order that the process
may be conducted at an elevated temperature. Free iodine when
allowed to react with a heated aqueous alkaline solution of ethyl
alcohol produces iodoform, (CHI 3 .) Fig. 104 shows the' assembled
apparatus with a thermometer for observing the temperature of
the reaction. The nickel gauze cylinder, porous pot, and anode
are also separately shown at the right in this illustration. Without
taking into account the intermediate products formed, the reaction
may be expressed very simply by the following equation:
CH 3 CH 2 OH + 10! + H 2 O - CHI 3 + CO 2 + yHL
One will observe that we have hydriodic acid formed, which, of
course, will combine with the sodium hydroxide present to produce
sodium iodate and carbonic acid. The following directions for
the actual carrying out of an experiment may be followed to advan-
tage: The cathode of nickel-wire gauze is placed in the beaker,
2IO
EXPERIMENTAL ELECTROCHEMISTRY.
together with the cathode liquid, which consists of a strong solution
of sodium hydroxide. Within the porous pot, which is next placed
in position is a solution consisting of 15 grams of sodium hydroxide,
10 grams of potassium iodide, 10 cubic centimeters of ethyl alcohol,
and 100 cubic centimeters of distilled water. The thermometer is
FIG. 104. Apparatus for the Electrolytic Production of lodoform.
placed within the porous pot and the temperature elevated to about
70 C. The best working current density at the anode for this
preparation is about one ampere per square decimeter. It is well
to allow the current to run for 4 hours, when the process may be
interrupted. The liquor from the interior of the porous pot is
poured out into an evaporating dish, when after standing for some
PREPARATION OF ORGANIC COMPOUNDS. 211
time, from i to 2 hours, a beautiful crystalline deposit of iodoform
is filtered off and allowed to dry at the temperature of the laboratory.
There will be formed as a secondary product in the mother liquor,
sodium iodate. The yield of iodoform is about 70 per cent. This
is a very satisfactory organic preparation, and lends itself to some
interesting efficiency determinations when conducted in connection
with suitable electrical measuring instruments.
THE ELECTROLYTIC PRODUCTION OF CHLOROFORM.
Chloroform may also be prepared electrolytically. For this
purpose a suitable still, which may be heated by a steam jacket
and containing a set of revolving paddles, is employed. These
paddles consist of carbon plates and are made the anode in the
electrolytic process. The interior of the still, which must be of
lead, serves as cathode. A 20 per cent solution of common salt
is placed in the still to which acetone is admitted from the bottom >
as shown, by means of the tube which leads to the elevated reservoir.
The acetone is converted into chloroform by the combined action.
of chlorine and sodium hydroxide. The reaction may be theoretically
illusttated in two stages according to the following equations:
(i)
Chloracetone.
(2) CH 3 COCCI 3 +NaOH = CH 3 COONa + CHC1 3 .
Sodium Acetate. Chloroform.
The chloroform produced distils off because of the elevated tempera-
ture maintained by the live steam and is collected in a suitable
receiver. It is claimed that from 100 parts by weight of acetone
1 80 parts by weight of chloroform are produced. The theoretical
yield figures out 206 parts by weight of chloroform, so it will be
seen that the process is quite economical. It is interesting to note
from a study of these equations that only one of the two available
methyl groups in the acetone is utilized for the production of chloro-
form. In the drawing, Fig. 105, it will be observed that the current.
is sent into the revolving anodes by means of a brush and the com-
mutator.
THE PRODUCTION OF ACETYLENE.
There is to-day a big field for the organic chemist with refer-
ence to electricity. Seme of the reactions brought about by the
212
EXPERIMENTAL ELECTROCHEMISTRY.
aid of electricity possess the greatest field for synthetic organic
chemistry. Berthelot showed that carbon and hydrogen combined
to form acetylene on causing the electric arc to pass between carbon
electrodes in an atmosphere of hydrogen.
For this purpose a glass globe was employed with two openings
opposite each other in the form of tubulures, into which were fitted
FIG. 105. Apparatus for the Production of Chloroform Electrolytically.
large stoppers carrying electrodes and entrance and exist tubes
for the gas. The globe was first carefully swept free of air by a
current of hydrogen when an electric arc was established through
the carbon pencils within. A good yield of acetylene results from
such a combination. This is, of course, simply of scientific interest
and has no practical application. The following organic synthesis,
however, is not only of scientific interest in the experimental labora-
PREPARATION OF ORGANIC COMPOUNDS.
213
tory, but has found a commercial application. It pertains to
the production of carbon disulphide in the electric arc.
THE PRODUCTION OF CARBON DISULPHIDE.
If, instead of supplying hydrogen to the enclosed electric arc
between carbon electrodes, we supply sulphur or roll brimstone
we get quite another product, namely: The mobile and volatile
liquid, known as carbon -disulphide. This interesting compound
may be prepared upon an experimental scale by assembling and
operating such a piece of apparatus as illustrated in Fig. 106. Here
FIG. 106. Glass Globe, Carbon Electrodes and Sulphur for the Experimental
Production of Carbon Disulphide.
we have a large glass globe equipped with four tubulures or necks
to receive stout stoppers. A metal rod with insulating handle passes,
as indicated, through two of these tubulures in a horizontal position,
terminating in holders for supporting carbon pencils. Before these
stoppers are put in place a narrow strip of thin asbestos is drawn
through and held in position as a bridge by the stoppers which
carry the electrodes. Upon this piece of asbestos, between the ends
214
EXPERIMENTAL ELECTROCHEMISTRY.
of the carbon electrodes, the roll brimstone is placed. The globe
is swept out by carbon dioxide gas and the arc is started between
the carbon pencils. Carbon disulphide in the state of a gas is
formed within the arc, condensing in minute drops over the interior
surface of the glass globe. After a suitable run, enough carbon
disulphide will condense to run to the bottom of the globe, where
FlG. 107. Experimental Equipment Complete for the Preparation of
Carbon Disulphide.
it flows out through the bottom opening and may be collected in a
test-tube. So much for the experimental side of this work. It will
be of interest to know that this process is being carried on upon a
commercial scale in New York State where specially designed
furnaces are in operation. These furnaces, erected at Penn Yan,
are 16 feet in diameter and 41 feet high. The process as carried
on by Taylor originally consisted in building a furnace with an
experimental shell of iron, in which the brimstone is placed and
melted down by the heat radiated from the inner metal shell of the
furnace where the electric arc had been maintained. Brick walls
were subsequently substituted for those of metal and the sulphur
in the cold state fed directly to the furnace surrounds the interior
PREPARATION OF ORGANIC COMPOUNDS. 215
so completely as to practically make a blanket, which, in melting,
carries back into the furnace the heat absorbed. This regeneration
proved to be efficient in the extensive production of carbon disul-
phide. It is said that in a building containing the furnace there
are no unpleasant gases that are in the least in evidence, the
entire building being at times as comfortable as any ordinary;
manufacturing plant; in fact other operations could be conducted
within the same building without inconvenience. Arrangements
were made in this process to keep the electrodes constantly and
automatically supplied with broken carbon which provides the
electrodes, themselves of carbon, with large contact surface,
from which the broken carbon tapers off to the interior of the fur-
nace where the current resistance converts the electrical energy
into heat just where it is required for effective work. The sulphur
rises in the bottom of the furnace, and its heat is regulated by feed-
ing cold sulphur into the surrounding chamber to meet the require-
ments. The sulphur, being a non-conductor of electricity, itself
plays an important part in regulating the amount of current which
flows through the furnace. The alternating current is used in Mr.
Taylor's furnace, which has practically revolutionized the manufac-
ture of carbon disulphide in America.
ELECTROLYTIC OXIDATION.
THE PREPARATION OF KANARIN.
As we learned in a previous chapter, oxidation and reduction may
be effected by the proper adjustment of current density, etc., and
we will now apply this important electrolytic oxidation method for
the practical production of certain organic oxidation products.
Organic electrolysis may, in the majority of cases, be divided into
two general processes entirely independent of current density con-
ditions. On the one hand we will have oxidation, and on the other
we will have reduction. Therefore, most cases of organic electrolysis
are either oxidation or reduction processes. The electrolysis of
organic compounds is of comparatively recent development, and a
sound knowledge of organic chemistry is essential for work in the
new field of organic electrochemistry. It would be impossible in a
216 EXPERIMENTAL ELECTROCHEMISTRY.
general experimental work on this subject to go deeply into the
chemistry and electrochemistry of organic oxidation or reduction
products, and we will content ourselves by introducing one or two
experiments which illustrate the effect of the electric current upon
organic compounds when applied under proper conditions. Per-
haps the production of the organic yellow dye, known as kanarin,
will serve best to illustrate a typical case of organic oxidation. For
carrying out an actual experiment we will employ a Hoffman appara-
tus of a general type, which is shown by the photograph in Fig. 108.
For this experiment the Hoffman apparatus is provided with plati-
num electrodes, and the electrolyte consists of a solution of potassium
sulphocyanide in the proportion of one part by weight of the sul-
phocyanide, and five parts by weight of distilled water. This solution
is placed within the Hoffman apparatus, the terminals of which
connect through two or three i6-c.-p. lamps to the lamp-bank and
ico-volt lighting system. Hydrogen is given off at the cathode
and streams up through the solution in the cathode tube of the
Hoffman apparatus, where it collects, as in the ordinary case of
electrolysis of dilute sulphuric acid. The interesting optical feature
of this experiment is the non-appearance of the corresponding
oxygen at the anode. Here we have, therefore, hydrogen streaming
off the cathode and collecting in the cathode tube, with no gas
liberated at the anode, for the oxygen set free immediately oxidizes
and combines with the sulphocyanide acid to produce kanarin,
which appears as a yellow mass extending over the surface of the
anode. The theory of this oxidation may be represented by the
following equation :
6HCNS + 1 iQ + H 2 O = C 6 H4O 2 N4S 5 + H 2 SO 4 + 2HNO 3 .
The kanarin which, after fifteen or twenty minutes, will have
collected in sufficient quantities for an experiment in dyeing, for
which purpose the kanarin is dissolved in a basic solution.
CASES OF ELECTROLYTIC REDUCTION.
Let us now look into a case of reduction, and for this purpose
we will choose nitrobenzene. The nitrobenzene is first dissolved
in strong sulphuric acid and placed in a porous pot, into which a
PREPARATION OF ORGANIC COMPOUNDS.
217
platinum cathode is immersed. The anode, of platinum, is placed
outside the porous pot in an 80 per cent solution of sulphuric
acid in distilled water. The porous pot with its cathode and
FIG. 108. Hoffman's Apparatus which Deserves a Special Place of Honor in
Electrochemical Work.
the anode are placed within a larger beaker glass and the electric
current turned on. The product is para-amido-phenol sulphonic
acid of the following chemical composition: C 6 H 4 (NH 2 )(OH).
The reaction within the cell is believed to take place in two stages,
21 8 EXPERIMENTAL ELECTROCHEMISTRY.
as shown in the following equations with the intermediate production
of phenylhydroxylamine :
C 6 H 5 (N0 2 ) + 2 H 2 = C 6 H 5 (NH)(OH)H 2 0;
Phenylhydroxylamine.
C 6 H 5 (NH) OH = C 6 H 4 (NH 2 ) (OH) .
Amido-phenol.
The ultimate product, which is para-amido-phenol sulphuric acid,
separates from the electrolyte in the form of fine crystals, which
are filtered off through asbestos. To illustrate the working range
of organic electrolytic reduction it may be well to refer to two other
preparations by the reduction of nitrobenzene. Nitrobenzene, when
in a diluted sulphuric acid solution, yields under similar conditions
of electrolysis, the aniline direct as shown by the following equation:
C 6 H 5 (N0 2 ) =C 6 H 5 (HNH 2 ) + 2 H 2 0.
When nitrobenzene is in an alkaline solution azobenzene is produced,
as shown by the following equation:
2 C 6 H 5 (N0 2 )H 8 = C 6 H 5 NNC 6 H54H 2 0.
Such reductions may, of course, be brdught about by ordinary
chemical means, but with the electrolytic process, there are many
outside points in its favor, among which may be mentioned an
ease of control of the course of the reaction. There are, however,
cases where electrolytic reduction brings about results different
from those obtainable by ordinary chemical means. It will now
be seen from the fundamental and typical cases given here, that
oxidation and reduction in cases of organic electrolysis do not
depend so much upon conditions of current density at anode or
cathode, but more especially upon the composition of the electrolyte
itself. The cases that might be here cited in the field of organic
chemistry are almost without limit, and we must content ourselves
in the present work with the typical cases presented.
CHAPTER XVI.
THE PRIMARY CELL.
IN the present chapter we will deal with the production of the
electric current through chemical action, and this phenomenon may
be regarded as the converse of what we have hitherto considered.
We have, up to the present time, applied the electric current to com-
pounds in a state of solution, and we will, in the present chapter, look
into the production of the electric current by the chemical action of
substances in solution within a suitably arranged and assembled cell.
Two great problems in electrochemistry presented themselves to the
early workers in this field, the first being: How does the electric
current decompose electrolytes and what is the mechanism of such
decomposition, which we- term electrolysis ?
The second problem relates to the origin of the electric current.
What produces it, and how can we satisfactorily explain the phenom-
ena observed when, for example, dissimilar metals are immersed in
an acid? We have already considered, at considerable length, the
decomposition side, or electrolysis, when compounds are subjected
to the influence of the electric current.
In our very first chapter we touched upon the origin of the
electric current in a general way and referred to the work of Galvani
with his frogs, and also to Professor Fabroni, of Florence, as well
as to Alexander Volta, of Pavia. For the sake of historic interest,
and also to illustrate to the student the hopeless condition of
affairs pertaining to the origin of the electric current, we
will turn for an instant to the time when Galvani and his famous
experiments upon frogs attracted the attention of the scientific
world. It will be remembered that in the famous experiment of
Galvani, conducted in the year 1786, the dawn of dynamic electricity
was marked. It is true that six years prior to this, Galvani observed
219
220 EXPERIMENTAL ELECTROCHEMISTRY.
that the limbs of dead frogs when hung upon a copper hook in the
neighborhood of a frictional electric machine contracted violently
at each spark or discharge of the then known static electricity.
It is the later experiment, however, conducted in 1786, which we
may refer to as the first to attract universal attention. It will be
remembered that Galvani obtained the twitching of the limbs of
dead frogs without the agency of any electrical machine whatever,
by bringing a copper wire joined to a nerve in contact with a piece
of iron wire, which was attached to one of the limbs of the frog.
The analogy of these results, although six years apart, nevertheless
impressed upon Galvani the belief that the two distinct phenomena
were due to one and the same cause, namely : that of electricity, and
he described his discovery of what he styled "animal electricity"
in his celebrated memoir, "De Viribus Electricitatis," in 1791.
A complete history of the voltaic cell and its origin would debar
us, for lack of space available here, from considering, as we must,
the later theories and the more experimental and practical side of
the electric battery. To give an idea, however, of the hopeless
state of affairs existing until the theory of electrolytic dissociation
threw some light upon the subject, we will refer to the views of Stur-
geon who, writing on the subject of Voltaism in 1842, expressed himself
as follows: " Voltaism is the production of electricity by the associa-
tion of metals and other inorganic bodies by the simple contact of
inorganic bodies, whether solid or fluid." Galvanism is the produc-
tion of electrical currents, " either by a natural or artificial associa-
tion of animal matter, whether alive or dead." The " animal
matter " element of this definition was eventually abandoned by most
of the investigators, when two distinct theories were formulated,
namely: "The contact theory of the cell" and the "chemical theory
of the cell." De la Rive, writing in 1853, defined the voltaic cell in
the following words : "An apparatus in which electricity is developed,
according to some, by the contact of two metals of a different nature,
and according to others, by the chemical action of the liquids with
which it is charged upon one of the two metals which enter into its
formation." Faraday expressed himself upon these two theories of
the cell as follows: "The contact theory assumes that when two
different bodies being conductors of electricity are in contact, there
is a force at the point of contact by which one of the bodies gives a
THE PRIMARY CELL. 221
part of its natural portion of electricity to the other body, which the
latter takes in addition to its own natural portion; that, though
the touching points have thus respectively given and taken electricity
they cannot retain the charge which their contact has caused, but
discharge their electricities to the masses respectively behind them;
that the force which at the point of contact induces the particles
to assume a new state cannot enable them to keep that state; that
all this happens without any permanent alteration of the parts that
are in contact, and has no reference to their chemical forces."
"The chemical theory assumes that at the place of action, the
particles which are in contact, act chemically upon each other and
are able, under the circumstances, to throw more or less of the acting
force; that, in the most favorable circumstances, the whole is con-
verted into dynamic force; that then the amount of current force
produced is an exact equivalent of the original force employed
and that in no case can any electric current be produced without
the active exertion and consumption of an equal amount of chem-
ical force ending in a given amount of chemical change."
Gore writes upon the theory of the cell as follows: "The essential
cause is the stored-up and ceaseless molecular energy of the corroded
metal and of the corroding element of liquid with which it unites,
whilst contact is only a static condition, and chemical action is the
process or mode by which the molecular motions of those substances
are more or less transformed into heat and current."
So much for these old theories. What have they taught us?
Can we, in the light of them, satisfactorily explain the cause of the
electric current, its origin, or birth? What must we do in order
to produce the electrical current by chemical action? In the light
of our chemical knowledge, let us see what takes place in the pro-
duction of electricity in the following simple experiment: A little
dilute sulphuric acid is placed in a beaker glass and two strips of metal
of dissimilar character, for example, platinum and zinc, are partly im-
mersed in the liquid so that they do not touch each other. If now the
uppermost ends of these metal strips be joined by a suitable wire, an
electric current in the direction from the platinum to the zinc will
be produced, as may be proven by a magnetic needle or galvanometer.
Bubbles of hydrogen may be seen to make their appearance on the
surface of the immersed platinum strip. So much for the physical
222 EXPERIMENTAL ELECTROCHEMISTRY.
manifestation. What can we say of the chemical? As general
chemists, we can merely analyze the solution and weigh the metal
strips for information relative to their loss or increase in weight. If we
analyze the diluted sulphuric acid solution we will find that it is no
longer merely a sulphuric acid solution, but that we also have zinc
sulphate present, and that the strip of zinc has lost in weight, whereas
the platinum is unaltered. An electric current has been produced.
What was its origin and how can we, as physical chemists, explain its
production? Before the advancement of the theory of electrolytic
dissociation this was a matter veiled in great obscurity. We could, of
course, say that the electric current was the result of chemical action,
or that it accompanied the formation of sulphate of zinc, or that it
was produced when zinc was dissolved in sulphuric acid, but does this
really take us as deeply into the inquiry as we wish to go ? In order
to comprehend the modern theory of the cell, as based on the theory
of electrolytic dissociation, we must know something concerning the
solution tension of metals in addition to facts in connection with the
theory of electrolytic dissociation. The solution tension of metals
when immersed in liquids may be compared with the vapor tension
of liquids. When a bar of metal is immersed in a liquid it tends
to dissolve, and does dissolve to a greater or less extent. When, on
the other hand, for example, an open vessel containing a liquid is
placed in the laboratory, the liquid tends to evaporate and does
evaporate to a greater or less extent. A bar of common zinc will
dissolve in dilute sulphuric acid much more rapidly than a bar of
iron, and, on the other hand, an open vessel of ether will evaporate
much more rapidly than a similar vessel of water. We may term
the tendency of the zinc to go into solution, the solution tension
of zinc, and the tendency of iron to go into solution, the solution
tension of iron, and from the fact that they go into solution at
different rates, we may logically deduce that they have different ten-
dencies to go into solution, or, in other words, we may say that they
have different solution tensions. We will, a little later in this
chapter, take up the experimental determination of the solution
tension of metals. We will see, when we come to this work, that
we have means for quantitively measuring and comparing what we
now term the solution tensions of the metals. We can at a fixed
temperature establish a solution tension value for any metal in
THE PRIMARY CELL.
223
a given liquid, as we can measure the vapor pressure or tendency of
a liquid to volatilize at a given temperature. A most important
fact for us to take into account at this time is the fact that the solu-
tion tension of any metal immersed in any given liquid is indepen-
FlG. tog. Experiment to show that Solution Tension is Independent of Surface Area.
dent of the size of surface of the metal immersed, which, we will see
is the case as indicated by E.M.F. phenomena, if we perform an
experiment as indicated in Fig. 109. We know, and can experi-
mentally demonstrate that the vapor prezsure of ether, for example,
is independent of the surface exposed under the same conditions
of temperature and barometric pressure. For experiment with
vapor tensions the reader is referred to any good text-book on
224 EXPERIMENTAL ELECTROCHEMISTRY.
experimental physics. As we will presently see, we may arrange
the known metals in a series, in the order of their solution tensions,
and such a table of the metals may be referred to in reference to
a liquid or electrolyte as the tension series of the metals. This
preliminary outline has been necessary, in the opinion of the writer,
before presenting to the student the following experiment, which,
upon the basis of the theory of electrolytic dissociation, explains
in a beautiful and satisfactory manner the origin of the electric
current in the voltaic cell.
The experiment to which we now refer has been styled by
Ostwald and others, "Chemical Action at a Distance." Ostwald
presented a paper in 1891 entitled, "Chemische Fernewirkung."
An exceedingly interesting point to which Ostwald draws atten-
tion is the fact that amalgamated zinc is not dissolved by dilute
acids. It has also been pointed out that chemically pure zinc
will not dissolve in dilute acids. This will appear to those
familiar only with general chemistry to be a remarkable statement.
The general chemist, without a knowledge of physical chemistry,
will find it difficult to believe that sulphuric acid, for example, will
not dissolve a stick of zinc if the latter be chemically pure. If, on
the other hand, the zinc is impure, that is to say, contains traces
of other metal as alloy or other metals adhering or cast within or
upon its surface, the zinc will dissolve with rapidity. It is quite
well known in this connection, that in the preparation of hydrogen
by throwing granulated zinc in dilute sulphuric acid that the evo-
lution of the gas is greatly promoted by the addition of some scrap
platinum or a few bright iron nails. Ostwald, in his writing con-
cerning this very point, speaks first of wrapping a platinum wire
around the top of a rod of zinc and immersing its lower end in
dilute acid. If the rod of zinc is chemically pure there will be
no appreciable chemical action if the free end of the platinum
wire does not dip in the acid. If, however, the end of this platinum
wire is immersed in the acid, the zinc will go into solution rapidly,
and hydrogen will be liberated from the immersed end of the plati-
num wire. Ostwald also states that it is not necessary for the zinc
to be surrounded by the platinum wire, for if such a wire touches
the zinc at any point where it is immersed in the acid, solution
will take place. It was also suggested that the zinc and platinum
THE PRIMARY CELL.
225
wire be joined at one place and then the free lower ends of both
zinc and platinum be immersed in a vessel containing, for example,
potassium sulphate. A porous partition is placed between the
immersed zinc and the platinum so that the electrolyte around the
zinc is separated from the electrolyte around the platinum.
Ostwald then asked the following question: "To which metal must
we add sulphuric acid in order that the zinc may be dissolved?"
This question, to the general chemist without a knowledge of
certain fundamental principles of electrochemistry, would appear
to be an absurd one, for it would seem to him to be very evident
that the acid should be poured into the partition containing the
zinc. If such an experiment is carried out we will find that in
order to secure the solution of the zinc, strange as it may appear,
the acid must be added to the compartment containing the platinum.
When the zinc dissolves, a brisk liberation of hydrogen gas is
observed from the platinum. This experiment is illustrated in
Fig. no and consists of two beaker glasses containing a solution of
X
Fie. no. Experiment to Show " Chemical Action at a Distance."
potassium sulphate, joined together by a siphon-tube containing
some of the same solution. In the left of our illustration may be seen
a cell of glass, cylindrical in shape, on the pattern of an ordinary
lamp-chimney, its upper end being fitted with a perforated stopper,
as shown, carrying a rod of pure zinc. The lower end of this
cylinder is tightly closed with vegetable parchment, simply to act
as a porous diaphragm. In the present arrangement of two beakers
226 EXPERIMENTAL ELECTROCHEMISTRY.
separated by a siphon-tube this precaution to prevent diffusion
currents is not absolutely necessary. The glass cylinder carrying
the zinc is, of course, also filled with the solution of potassium sul-
phate. A piece of platinum wire is wrapped around the top of the
zinc bar, and a second piece, terminating in a spiral to give it more
surface, is immersed in the distant beaker. If now a few drops of
sulphuric acid be introduced within the cylindrical cell containing
the zinc, there will be practically no solution of the metal. If, on
the other hand, a few drops of sulphuric acid are poured into the
beaker at a distance, and connected by the siphon containing the
platinum coil, a brisk evolution of hydrogen will take place from this
platinum wire, and after a few minutes the presence of zinc sulphate
within the glass cylinder surrounding the rod of zinc may be proven
by analytical means. In the drawing one will observe an
arrow pointing toward the rod of zinc in the left-hand beaker, and
just above the arrow the platinum wires are shown in contact. If
after adding sulphuric acid to the right-hand beaker, as shown in
the drawing, the liberation of hydrogen from the platinum spiral
and solution of zinc in the left beaker is noted, we separate these
platinum wires from contact, the liberation of hydrogen ceases
together with the solution of zinc. Upon touching the platinum
wires together again, however, the chemical action recommences,
and if we bring a sensitive magnetic needle in the neighborhood
of the platinum wire we will find that an electric current is flow-
ing through it in the direction of the arrow, namely : from the right-
hand beaker containing the platinum to the left-hand beaker con-
taining the zinc.
The explanation of these phenomena is comparatively simple
when based upon the theory of electrolytic dissociation. The
source of the current in a voltaic cell may be, therefore, understood
from the following consideration: When a rod of zinc is immersed
in a solution of a neutral salt, potassium sulphate for example,
zinc ions are sent off in solution because of the solution tension of
the zinc. These zinc ions are driven into solution because of the
solution tension of zinc in an analogous manner to the sending
off of ether molecules into the air when an open vessel of ether is
allowed to stand in the laboratory. In the case of the immersed
zinc it is made negative in electrical sign,, and the solution which
THE PRIMARY CELL. . 227
has received the ions, which are positive in sign, becomes itself posi-
tive. Thus solution continues until a difference of potential in
solution is established. Again comparing the phenomena of vapor
tension; ether in an open vessel would continue to evaporate until
an equilibrium is established. In the case of the zinc rod immersed
in the sodium sulphate solution, the number of zinc ions driven
off, although very small, establish an equilibrium after a while.
Referring once more to the case of the zinc rod in the electrolyte,
the driving off of the zinc ions will cease after a certain point has
been obtained, because of an excess of positive ions in the solution.
In order that more of the zinc ions may be driven out some of these
positive ions must be removed. If the zinc is connected with
another metal, for example our platinum wire, such platinum wire,
of course, takes the same negative charge as the zinc. When the
negative end of this platinum wire, therefore, is coiled and immersed
in the solution it attracts the excess positive zinc ions which exist
in the solution. We might expect, from the description of the
experiment as far as we have gone, that the zinc ions in the solu-
tion would be attracted to the platinum spiral, give up their charges
and deposit thereon, or, in the case of potassium, decompose the
water which is present with the liberation of hydrogen. The be-
havior in such a state of affairs depends upon, not only the nature
of the ion, but of the electrode also. In such an experiment, with
the positive ion, which is the potassium resulting from potassium
sulphate, the difference in potential produced upon the introduc-
tion of the zinc is insufficient to cause the potassium ion to give
up its charge to the platinum. If, however, a little sulphuric acid
is added to the beaker at the right containing the platinum coil,
and the wires are in contact, the difference in potential produced by
introducing the bar of zinc is sufficient to compel the hydrogen
to give up its positive charge to the platinum spiral and appear upon
its surface as minute bubbles. The hydrogen ions in their becom-
ing hydrogen atoms give up positive electricity to the platinum,
neutralizing the negative charge which the latter carries. A current
of electricity will then flow along the wire in the direction of the
arrow, as indicated, to the zinc, which will, of course, become less
negative than before the hydrogen separated at the platinum spiral,
and the difference in potential between the zinc rod and the sur-
228 EXPERIMENTAL ELECTROCHEMISTRY.
rounding solution becomes less. More zinc will, therefore, dissolve
or be driven into solution as zinc ions. Additional hydrogen ions
give up other charges to the platinum spiral and separate as gas,
which, of course, in turn tends to make the zinc still less negative.
Now, as long as we have the circuit closed, we will have an electric
current, one from the platinum to the zinc as a result of the con-
version of positive hydrogen ions to the ordinary hydrogen atoms.
We have already referred to the fact that pure zinc does not dissolve
in acids, while zinc which is impure displaces the hydrogen of an
acid with readiness, or, as we may say in popular language, dissolves.
The tendency of zinc, whether pure or impure, to go into solution
is the same, only in one case it does and the other case it does not.
Pure zinc, however, dissolves readily in an acid, or as we may say,
technically speaking, displaces the hydrogen in an acid with avidity
when in contact with some other metal of lower solution tension,
such as platinum for example, immersed in the acid. The differ-
ence is not in the solution of the zinc, but is due to the ease with
which hydrogen may escape from the solution. The presence of
such a metal as platinum with a very low solution tension, allows
the hydrogen to escape from its surface with ease, and upon this
principle we may see why impure zinc dissolves in acids, when
such impurities of low solution tension exist and act as points, or
surfaces from which the hydrogen ions may discharge their elec-
tricity and escape as hydrogen gas. With a stick of impure zinc
we have numerous impurities in the way of specks upon its surface
of a .lower solution tension than the zinc itself, and we will have
a multiplicity of little galvanic circuits between the zinc impuri-
ties and through the electrolyte from the impurities to the zinc
through the point of metallic contact. The reason why chemically
pure zinc will not dissolve may also be ascribed to the fact that this
metal itself has very high solution tension and sends its own positively
charged ions into solution under a high solution tension opposing
the tension of any other positive ions, like hydrogen for example,
upon it. The rod of pure zinc will not dissolve in acids, therefore,
because the hydrogen ions cannot give up their positive charges to
it to escape as hydrogen molecules. As we saw by referring to
Fig. 109, electromotive force is an exponent of solution tension,
and from the experiment as depicted it may be seen that this solu-
THE PRIMARY CELL. 229
tion tension is independent of the size or area of the metals immersed
in an electrolyte. A difference in potential between the metal and
the electrolyte is therefore established, which proves to be the fun-
damental origin of the E.M.F., produced in any given combination.
We may ascribe the direct cause of this difference of potential
to the solution tension of the metal which tends to drive ions
from the metal into solution, making the metal itself negative in
sign and the solution positive in sign, because of the presence of posi-
tively charged ions. It may be pointed out here that we also have
a pressure in such a cell antagonistic to the solution tension of the
metal, and this is the osmotic pressure of the solution itself, which
tends to cause the ions driven out to separate on the electrode in
the metallic condition. As a consequence of these opposing forces
we have the formation of a double layer so to speak, and the dif-
ference in potential between the metal and the solution. This
double-layer phenomenon is referred to in the electrochemical works
as a Helmholtz double layer. Dr. Nernst very clearly describes
the phenomena of the Helmholtz double layer in such a concise
and lucid manner that it is deemed of value to the student to trans-
late his own words upon the subject here. "Let us now consider
what will take place if we dip a metal whose electrolytic solution
tension is P into a solution of one of its salts; the osmotic pressure
of the metal ions in this solution being p. Let at first P> p, at the
moment of contact a number of positively charged metallic ions,
driven by this large pressure, will pass into solution. Since by the
latter a certain amount of positive electricity is carried from the
metal into the solution, the liquid receives a positive charge, which
arranges itself in the form of the positive ions contained in the
solution on the surface of the metal. At the same time there is, of
course, a corresponding amount of negative electricity set free in
the metal, which also passes to the surface of the metal. We recog-
nize at once that at the surface of contact of metal and electrolyte
the two kinds of electricity must accumulate in the form of a double
layer, whose existence, as is well known, was made probable some
time ago by Von Helmholtz in an entirely different way.
" This double layer furnishes one component of force, which
acts at right angles to the surface of contact of the metal and the
electrolyte, and which tends to drive the metallic ions from the
230 EXPERIMENTAL ELECTROCHEMISTRY.
electrolyte on to the metal, and thus acts in opposition to the solu-
tion tension. Equilibrium will be, of course, established when these
two forces equalize one another. The final result will be the appear-
ance of an electromotive force between the metal and the electro-
lyte, which will give rise to a galvanic current from the metal to the
liquid, if by any device its existence is made possible.
"IiP<p the reverse of course takes place. Metallic ions separate
from the electrolyte and are precipitated on to the metal until the
electrostatic component of force of the positive charge of the metal
and the negative charge of the liquid thus produced are in equilibrium
with the excess of osmotic pressure. An electromotive force again
appears between the metal and the electrolyte, which, under suitable
conditions, gives rise to a galvanic current, but in this case opposite
in direction to the case first considered.
"If, finally, P = p the metal and electrolyte are in equilibrium
at the first moment of contact; therefore, no difference in potential
exists between the two."
In our practical and experimental work it is very important that
we note quantitatively the potential differences existing between
metals and solutions, and in order to practically determine this
we will adopt the method involving a "normal electrode." This
method employs the use of a previously prepared electrode, which
has been termed a normal electrode, because its potential is known.
This normal electrode is connected with a metal whose difference
in potential we wish to ascertain and the E.M.F. of the entire
system determined. As we know the potential of the normal elec-
trode, that of the metal under examination is easily determined,
the E.M.F. of the two when properly combined being the difference
between the potentials on the two sides. The use of the normal
electrode may be understood by referring to Fig. in, where at the
extreme left we have the normal electrode equipment, the beaker
glass in the center containing the metal x, whose tension we are
about to study, while at the right we have a sensitive standard
volt-meter. The normal electrode equipment consists of a wide-
mouth glass bottle containing a layer of purified mercury in the
bottom, as indicated. On top of the mercury a layer of mercurous
chloride is placed, and on top of this a normal solution of potassium
chloride is poured. A platinum wire which passes through a pro-
THE PRIMARY CELL.
231
tecting glass tube is suspended from the stopper, as shown. The
wire is sealed in a glass tube in such a way that its lower extremity,
which is in the form of a tiny coil, comes into contact with the metallic
mercury in the bottom of the bottle. A siphon-tube passes through
232
EXPERIMENTAL ELECTROCHEMISTRY.
a second hole in the stopper of this bottle and dips in the normal
potassium chloride solution, after having itself been filled with the
same electrolyte. Connection is made with the mercury through
the platinum wire, insulated by the glass tube, and this mercury
serves as one -electrode. The siphon is inserted in the liquid, whose
FIG. 112. Special and Convenient Design of Normal Electrode.
potential against the metal under examination we wish to learn.
The metal in question serves as the other electrode and is connected
up with a sensitive volt-meter, as shown, and the E.M.F. of the entire
system determined. The electromotive force between the mercury
and the potassium chloride is .56 volt at the ordinary temperature
of the laboratory. This .56 volt may be used as a constant for our
THE PRIMARY CELL 233
normal electrode, no matter what its size, if put together with chemi-
cally pure materials upon the plan as indicated. If the liquid in
the beaker containing the metal, whose tension is to be measured,
reacts chemically with potassium chloride, the solution of some
indifferent compound may be interposed between the two. Fig. 112
shows a very neat and convenient design for a normal electrode.
Here we have a glass vessel containing a platinum wire fused into
its lower end to make contact with the mercury. The mercurous
chloride is then put in position, as in the previous case, and the
normal potassium chloride solution put on top and made to fill the
siphon, which in this instance is fused into the side of the vessel.
This form of normal electrode is conveniently held in an iron retort
clamp as shown.
The following table is taken from the work of Neumann and
represents the differences of potential between several metals and
normal solutions of their salts.
Metal. Chloride.
Magnesium 1.231 Volts.
Aluminium i .015 ' '
Zinc .' -53 "
Cadmium 0.174 "
Iron o .087 c '
Cobalt -0-015 "
Nickel 0.020 tl
Tin -0.085 "
Lead 0.095 "
Gold -1.356
Platinum 1.066 "
We have learned that the solution tension of a metal is respon
sible for the difference in potential between itself and the electro-
lyte into which it is immersed. If we determine the potential differ-
ence and the value of the osmotic pressure of the positive ions in
solution, we have the data necessary for calculating the solution
tension of the metal.
A few common metals arranged in the order of their solution
tension is given below, and this may be termed a tension series.
234 EXPERIMENTAL ELECTROCHEMISTRY.
Magnesium,
Zinc,
Aluminum,
Cadmium,
Iron,
Cobalt,
Nickel,
Lead, '
Mercury,
Silver,
Copper.
A metal anywhere in the above series will tend to precipitate
from its salt a metal located lower in the series, for example, zinc
will precipitate copper from its salts, etc. A metal at any point
in the series, when made an electrode in a cell of battery against a
metal lower in the series, serving as the other electrode, will throw
off ions in the solution, and thereby become the negative pole. Zinc
is the negative pole in almost all cells of battery. The position
of a metal in the tension series is of extreme interest to us in the
design of primary batteries.
We will close the present chapter by a consideration of the
energy of the primary cell based upon a known chemical reaction.
If we know the chemistry of a cell of battery we can, by a simple
mathematical process, predict what its electromotive force will be.
This calculation may be very simple or very complex, depending
upon which way we attack the problem. As it has been the aim
of the author to avoid the higher mathematics in the present work,
we will proceed with a comparatively simple formula. Let us,
therefore, select a typical simple cell, look into its chemistry,
and predict therefrom the maximum electromotive force which
such a cell is capable of giving. This work will take us back into
certain of our fundamental principles as introduced in the early
portion of the present book and involve Faraday's Law, together
with several fundamental units, and the most important constant
96,540. For this purpose of illustration we may do best by select-
ing the well-known Daniell type of element, or cell. Helmholtz has
pointed out that a relationship may be established between the energy
THE PRIMARY CELL. 235
of the chemical process, or, in other words, the chemical energy of
a cell, and the electrical energy produced in exchange. It will be
seen that under ordinary conditions, most of the energy which could
be obtained as the heat of chemical action can be converted into
electrical energy and be made to do work as an electric current.
If we allow Q to represent the available heat energy, for one gram-
equivalent of a compound that enters into chemical combination
in a cell, it may be assumed in some cases that an amount of elec-
trical energy equivalent to this heat energy can be obtained from
the cell for each gram-equivalent of chemical transfer. In the case
of a Daniell cell, the conditions may be represented by the following
equation :
Cu, CuSC>4 Solution; ZnSO* Solution Zn.
After the cell has been allowed to do work, the condition of affairs
may be represented as follows:
Zn, CuSO 4 = Cu, ZnSO 4 .
Zinc is therefore dissolved at one pole of the battery, and copper
is deposited at the other. The heat value of such a reaction is
the difference between the heat of formation of copper sulphate and
zinc sulphate in aqueous solution. We have, therefore, 106,090
55,960 = 50,130 calories per gram-equivalent. When 32.5 grams
of metallic zinc displace an equivalent of metallic copper from a
solution of sulphate of copper, 25,065 calories are set free. We
will remember from our study of Faraday's Law that about 96,500
coulombs of electricity are obtained for every 32.5 grams of zinc
transported, so if this displacement takes place in a suitable cell
we will have 96,500 coulombs of electricity delivered to us. We
also learned that the Joule, which is the product of i volt by i cou-
lomb, is equal to .00024 large Calories or .24 small calories. The
electrical energy equal to this number of heat units is 25,065 -^-.24 =
104,240 Joules, therefore - = 1.08 volts. This calculation is
96500
quite similar in character to the method introduced when we were
calculating the minimum E.M.F. required to decompose the gram-
molecular weight of any electrolyte when its heat of formation is
known. The Daniell cell, therefore, should give us an E.M.F. of
236 EXPERIMENTAL ELECTROCHEMISTRY.
i. 08 volts; and direct measurement of this cell gives us 1.09 to i.i
volts. It will, therefore, be seen how close we may come to
the prediction of E.M.F. of a cell when we know its general
chemistry.
CHAPTER XVII.
THE SECONDARY CELL.
THE secondary cell in many ways is immensely superior to all forms
of primary battery and is of the utmost interest to electrochemists.
Unlike the primary cell, it is not susceptible to polarization, or
counteraction, resulting from the formation of a film of hydrogen
gas upon the surface of the negative electrode. As hydrogen is
negative to zinc, for example, a counter electromotive force is set
up and the conditions very materially modify the output in the
external circuit. Efforts have been made to diminish this in the
primary cell by adding depolarizers, which substances combine
with the hydrogen liberated at the cathode to form water and are,
therefore, oxidizing agents; bichromate of potash is frequently used
for this purpose. Whereas the chemistry of the primary cell is well
known, that of the secondary cell or storage battery using lead
plates and sulphuric acid is far from being understood by chemists.
Almost any reversible type of primary cell may be termed a storage
battery, because of its reversibility. The original condition of the
electrodes and electrolyte may be re-established after a general altera-
tion has been wrought. To give a simple example, let us immerse
in a beaker containing dilute sulphuric acid, a stick of chemically
pure zinc and a strip of platinum. We have learned from our
previous study of the primary cell that no chemical action will take
place unless suitable wires are joined to these two metals and brought
into contact as in an external circuit. We will then have, as we
know, zinc ions being forced into solution and hydrogen ions discharg-
ing upon the platinum electrode. An electrical current flows through
the wire connector and after a time the electromotive force will be
found to fall. An analysis of the dilute sulphuric acid will reveal
the presence of zinc sulphate, and we have seen that a certain quan-
237
238 EXPERIMENTAL ELECTROCHEMISTRY.
tity of hydrogen has escaped as gas. There will be minute bubbles
of hydrogen clinging to the platinum, however, and it is due to this
fact that we have the phenomenon of polarization. By adding
potassium dichromate to the dilute acid solution we can effectually
prevent the formation of the hydrogen bubbles and maintain a
more constant electromotive force. We can also bring about this
depolarization by substituting a rod of copper oxide, for example,
for the platinum, which will be reduced to metallic copper by the
hydrogen liberated upon its surface. The cell will then furnish a
certain amount of current and do a definite amount of work.
There will be found no sulphuric acid left, but a concentrated solu-
tion of zinc sulphate instead. The zinc electrode has been partly
consumed. The zinc, of course, has displaced in the acid the
hydrogen, which is set free. Upon sending an electrical current
through the cell in the reversed direction, however, the cell
is, incorrectly speaking, recharged. The "charge" in the present
case consists in the deposition of the zinc through the zinc sulphate
upon the zinc electrode and the reformation of sulphuric acid.
Theoretically speaking, therefore, the only thing lost during the
discharge and charge of this particular cell is hydrogen and oxygen
gas in the proportion in which they unite to form water. Such a
storage cell is not of a practical nature, however, and has simply
been introduced to show the general principle upon which reversible
cells depend. Let us, in our study of the lead -lead -sulphuric acid
accumulator, or storage battery, experiment a little, and for this
purpose we may best begin by constructing a simple cell, study its
practical behavior by charging and discharging it, and examine
the plates and electrolyte, by experimental methods, before we
look into the theory. Fortunately, there is no more easily con-
structed type of cell than a simple storage battery. For our experi-
mental purpose let us construct a cell, as illustrated in Fig. 113.
We may use a rectangular glass cell or large beaker. A rectangular
glass cell is preferable, however, and one about 6 inches high and
about 6 inches long by 2 J inches wide will answer our requirements.
A rectangular glass jar is to be preferred, for the reason that the lead
plates required may be cut in a convenient shape to hang over
the mouth of the jar. These lead plates should be cut from sheet
lead, not over of an inch in thickness, and may be provided with
THE SECONDARY CELL.
2 39
lugs and shoulders, as shown by the diagram of the plate in the center
of the illustration. To assemble and charge such a cell we will
put into the jar an electrolyte, consisting of one part of concentrated
sulphuric acid sp. gr. 1.84 to ten parts of distilled water. The
lead electrodes are immersed therein after having been thoroughly
cleaned by dipping them into dilute nitric acid, if the lead was not
perfectly free from impurities beforehand. A ' small cell of this
character may be easily charged in a very few minutes by connecting
it to the motor-generator and supplying about 8 amperes to it under
a potential difference of not less than 4 volts. The cell may also
be charged by including it in the no-volt electric lighting circuit
FIG. 113. Easily Constructed Experimental Secondary Cell.
with a couple of lamps in multiple arc. The primary effect of
the electric current is to decompose the water between the lead
electrodes. The liberated hydrogen escapes from the cathode,
therefore, and the electrolytic oxygen from the anode. The anode
plate, if perfectly bright, may be seen to darken under the oxidizing
action of the electrolytic oxygen, whereas the cathode assumes a
characteristic lead -gray color. For the first few moments, until the
two electrodes are reduced and oxidized respectively, there may be
no oxygen and hydrogen liberated. The oxygen is oxidizing
the metallic lead on the one hand, and the hydrogen is
reducing any oxides which may be upon the opposite electrode,
on the other hand. After a while, however, bubbles of gas will
240
EXPERIMENTAL ELECTROCHEMISTRY.
appear at the cathode and soon afterwards at the anode, when any
further current furnished will not go to charging the cell, but be
expended in electrolizing the acidulated water present without
any further useful end.
The central diagram in Fig. 113 illustrating the form of lead
electrode also indicates in the shaded portion A, the manner in
which the lead oxide leaves the surface of the oxidized plate on
discharge. By referring to Fig. 114, this plate is shown in three
FIG. 114. Diagram Showing the Appearance of the Positive Plate in Three
Stages of Discharge.
stages of discharge When the cell is completely charged the entire
surface of the annode or positive electrode is darkened with the
oxide film. Upon discharge, however, the oxide film begins to
leave or retreat from the bottom and also the sides of the plate,
the diminishing area being depicted by the shaded areas B, C }
and D in the figure. It is of interest to note that so long as we
have a small area of this oxide film, the E.M.F of the cell is prac-
tically constant. This is analogous to the E.M.F. of a primary
cell which, as we saw, is a fixed value whether the electrodes are
immersed to a large or small extent. The E.M.F., therefore, is again
due to the phenomenon of solution tension. If we fully charge
a lead-lead-sulphuric-acid accumulator we will find that for a short
period of time the cell is capable of giving us an electromotive force
of over 2 volts. This E.M.F. has but a very short duration, and
is believed to be due to the occlusion of hydrogen by the cathode
plate, which, after repeated charging and discharging, assumes
a more or less porous and spongy character. We will therefore fre-
quently obtain from storage cells having porous or spongy cathode
THE SECONDARY CELL.
241
plates, an electromotive force as high as 2.4 volts for short
periods of time. After we have charged and discharged the experi-
mental cell, with which we are dealing, we will be able to notice
this interesting phenomenon. If now we allow the cell to do work
by discharging it through a suitable external resistance, interrupting
the current momentarily at regular intervals and measuring its
E.M.F., or electrical pressure, we will find it to be a trifle under 2
volts and constant up to a certain point, when its value falls abruptly.
In other words, a storage battery will yield a good current at a con-
stant voltage for a certain length of time, when the bottom will, so
to speak, fall out all at once. By referring to the diagram in Fig. 115,
Z.3
LI
I.I
/.7
/.s
/.*
s^
^v
s,
^
V
\
\
/ 2 3 t f 6 ye f /<?///* S3 /* /? /t /r /a
/foi/fls
FIG. 115. Diagram Illustrating the Character of Discharge of a Secondary Cell.
the discharge of a typical storage battery is plotted in the form of a
curve for twelve hours. According to this diagram the cell gave a
current under a difference of potential of a trifle less than 2 volts,
when an abrupt falling off of the electromotive force was noted.
For twelve hours, therefore, we may refer to the cell as having a
horizontal line of discharge. This is an idealized curve, for in
practice this discharge -line is never perfectly horizontal, although
very nearly so. We should plot such a curve with our storage cell
by reading a delicate volt -meter across the electrodes at, let us say,
minute intervals without interrupting the flow of the current in
the external circuit. We should now experimentally determine the
changes which take place in the electrolyte, that is to say the density
242
EXPERIMENTAL ELECTROCHEMISTRY.
changes that are wrought in the sulphuric acid solution upon charge
and discharge.
During the discharge of a cell of battery the density or specific
gravity of the electrolyte falls, and increases again on charging. By
knowing the density, therefore, at the point of full charge, and at
full exhaustion, we may learn something of the state of the cell
between those points by specific gravity determinations. For our
experimental cell a sensitive type of hydrometer may be employed,
as depicted in Fig. 116. This peculiar construction is the design
FIG. 116. Author's Sensitive "Reflecting Hydrometer" for Observing Minute
Density Changes.
of the author as applied to storage-battery study in 1895. It
consists of an ordinary glass hydrometer equipped with a metal cap
with a V-shape bearing to support a small knife edge on the end
of an aluminum beam. The other end of this aluminum beam
is equipped with a similar knife edge, which rests in a like V-shape
bearing, mounted upon a standard which is carried by a float. This
float is provided with a central aperture, not unlike that in a cake
dish through which the stem of the hydrometer passes. A tiny
mirror is mounted on one end of this aluminum bar, to receive
a beam of light which it reflects upon a scale analogous to that on
the reflecting galvanometer, only in the present case the scale is ver-
tical in place of being horizontal. It will now be appreciated, from
a glance at the diagram, that the minutest changes in specific gravity
may be noted by the movement of the spot of reflected light upon
the graduated scale. Another form of specific gravity indicator
which will prove useful in our experimental study of the cell is
THE SECONDARY CELL.
243
illustrated in Fig. 117, which is the design of Mr. J. S. Sellon, and
although not so delicate as the one just described, has a useful and
practical application for experimental work. We can now charge
and discharge our cell, measuring the energy supplied and the
energy delivered, examine the electrodes and their behavior, and the
FIG. 117. Sellon's Design of Hydrometer as Applied to Secondary Cells.
physical changes which take place, such as the density alterations
in the electrolyte. What now can be said regarding the theory of this
type of battery ? It would appear at first sight to be an easy matter
when we have such simple elements entering into the construction
as pure lead plates and dilute sulphuric acid. Unfortunately, if we
attempt to follow theoretically the transformations which takes place
in a complete cycle, that is to say the charging and discharging of a
secondary cell, we will find that we become lost a number of times in
the determination.
Regarding the chemistry of the storage battery in view of the
exceedingly difficult chemical problems involved, we will do well
to quote from Professor Ayrton's paper on the ''Chemistry of
Secondary Batteries.'''
244
EXPERIMENTAL ELECTROCHEMISTRY.
"The physical qualities of the cells are capable of very accurate
estimation and investigation. But when you come to attempt to
ascertain the chemical changes that occur in the charging and dis-
charging of a storage cell you encounter formidable difficulties.
The outsider has no idea of these difficulties. Nothing seems more
simple than to determine the chemical changes that take place in
either the positive or the negative plate of a storage battery. It
is not so in reality. The substances used as actve materials are
in the first place mixtures, and the materials obtained at the end
of the reactions are also mixtures, and these mixtures are insoluble
in any reagent which does not decompose them. They cannot be
volatilized ; they cannot be subjected to any process of solution and
crystallization in order to separate and purify their elements."
There are, however, several theories advanced to account for
the charging and discharging of a storage battery, and we would
do well to consider some important ones here. Let us first take
up the theory of Plante.
The first diagram illustrates the effect of the electric current in
the formation of sulphate of lead by the substitution of lead for the
hydrogen in the sulphuric acid.
T
H 2 SO 4 H 2 SO 4 H 2 SO 4 Pb
The second diagram shows the production of persulphuric acid
by the elimination of a further molecule of hydrogen.
T
H 2 SO 4 H 2 SO 3 O
The third diagram shows how the formation of peroxide of lead
may take place by the reaction between persulphuric acid and sul-
phate of lead in the presence of water.
I
H 2 S 2 Og
2H 2 O
+ PbS0 4
THE SECONDARY CELL. 245
The fourth diagram shows the production of a second molecule
of persulphuric acid by the current proceeding from the peroxide-
coated anode, and also the decomposition of the persulphuric acid,
in the presence of water, which results in the formation of peroxide
of hydrogen and the reformation of sulphuric acid.
T
H 2 SO 4
H 2 S0 3 g
H 2 S 2 Og
2 H 2 O
1 j H 2 2
f / 2 H 2 SO 4
Another theory of the storage battery is as follows: During the
discharge of a cell both electrodes are converted into lead sulphate
with the abstraction of SO 4 from the electrolyte, which diminishes
its specific gravity. The change on the anode or positive electrode is
believed to take place in two separate stages: First, the reduc-
tion of the peroxide of lead to the monoxide, and then the conversion
of the monoxide of lead into the sulphate of lead. When the
cell, is charged we have the reverse, the sulphate of lead being
converted into the peroxide on the positive plate and metallic lead
on the negative plate or electrode. According to Treadwell, the
generally accepted theory at present is that of the direct formation
of lead sulphate at both electrodes, "each molecule of the peroxide
is supposed to loose an atom of oxygen, and each atom of spongy
lead to gain an atom of oxygen. Two molecules of hydrogen sul-
phate are thus abstracted from the electrolyte to react with the
peroxide or spongy lead, and their places are taken by 2 molecules
of water.
PbO 2 +H 2 SO 4 = PbSO 4 +H 2 O + O;
and that for the negative plate is
Pb-fO+H 2 SO 4 = PbSO 4 + H 2 O;
or, including both reactions in one equation,
PbO 2 + 2 H 2 SO 4 + Pb = PbSO 4 + 2 H 2 O + PbSO 4 .
"Thus the final result of the complete discharge of a cell is to
form lead sulphate and water by removing sulphuric acid from the
electrolyte and depositing sulphate of lead upon each plate.
246 EXPERIMENTAL ELECTROCHEMISTRY.
" The above is, fortunately, a self -limiting process, since the sul-
phate is a poor conductor. All the peroxide is therefore not acted
upon, and at the end of the discharge we have peroxide of lead
crystals covered with a coating of sulphate."
So much for the study of conditions. The cell illustrated in Fig.
113 is of the simplest possible type, and, because of the limited
electrode area of smooth metallic lead, has but a very low current
capacity. When fully charged it will give a current for but a few
moments, and for this reason is only adapted for the briefer kind of
research and study. This battery can be increased in capacity by
roughening the lead plates through the agency of a knurl or other
suitable tool.
The battery can also be wonderfully increased in efficiency by
not only roughening the anode or positive plate, but by preparing
it by filling it in with active material. The practical cells of battery
on the market to-day consist of such specially prepared plates, which
are termed grids. It would be almost impossible to describe or
illustrate the various designs or types of grids intended to meet the
rather trying conditions of practice. Perforations of every conceivable
shape and size, grooves, mat-work, applications of disks, buttons,
chambers, etc., etc., have been devised in order to hold the active
material, prepared in numerous ways, to form the positive plate. It
would require a special treatise on the storage battery to begin to
illustrate and to do justice to the numerous ingenious designs of
storage -battery electrodes. Apart from designing a positive plate
for maximum efficiency for receiving and holding the active peroxide
of lead, we must design an electrode which will not "buckle" under
working conditions. Buckling, in storage-battery parlance, is the
warping, twisting, and bending of an electrode due to the inequalities
in expansion between the active material and the supporting grid
or lead work during charge or discharge. With certain types of
electrodes this evil is manifested if a battery is charged too quickly
or allowed to give up its available current in too short a period of
time. Twisting or buckling of the electrode of a storage cell, not
only ruins the plates, but interrupts the duty of the cell in a battery
because of the short -circuiting which is brought about. A modern
storage battery of large capacity is a rather costly thing, and in
the hands of incompetent attendants may be easily damaged.
THE SECONDARY CELL. 247
Among the precautions in the care of the storage battery may be
mentioned the desirability of having the electrolyte perfectly homo-
geneous, that is to say, free from strata of different acid concentra-
tions, which will bring about uneven action upon the electrode.
Great care should be taken to prevent foreign bodies from falling
in between the plates, and to supply pure water to make up for loss
due through electrolysis and evaporation. The battery should
be charged at either constant voltage or constant current, the
current being supplied to the battery under a potential difference
only a little higher than that of the battery itself. For example, if
we wish to charge ten cells of storage battery in series, a current
under a pressure of about 30 volts, will be good practice. In the
charging oi a single cell, therefore, we should not seek a high-pres-
sure current, but would prefer one with a voltage of about 2% to 3
volts. In charging the battery, we must know the current density
conditions which the battery will stand, that is to say, we should not
supply too many amperes per square foot of electrode area. The
same question is involved when the battery is allowed to do work,
and may be injured as a dynamo may be injured if allowed to dis-
charge through an external circuit of too low a resistance. In the
case of the battery we are liable to have buckling due to the heavy
current passing, and in the case of the dynamo we are liable to burn
out the armature. Too large a charging current, apart from the
liability of injuring the battery, is very uneconomical, for a good
deal of the electrical energy is transformed into heat. We should,
therefore, keep the charging current comparatively low if we seek
economy, unless the time of charging is to be counted as one of
the elements in cost. Too low a charging current, on the other
hand, is also injurious to the cell, for it is found to produce the white
sulphate of lead upon the positive plate, instead of the active per-
oxide, which we seek. Gladstone and Tribe, in their work entitled,
"Chemistry of Secondary Batteries," bring out this point as follows:
"If we take two plates of lead in dilute sulphuric acid and
pass a current from only one Grove cell, a film of white sulphate,
instead of peroxide, makes its appearance on the positive plate,
and the action practically ceases very soon. If, however, the current
is increased in strength, the sulphate disappears, and peroxide
is found in its place."
^8 EXPERIMENTAL ELECTROCHEMISTRY.
With a good battery a safe rule to follow is to furnish current
to the extent of about 8 amperes per square foot of electrode area.
After a battery has been fully charged, oxygen and hydrogen
gases are, of course, given off at both electrodes, and the energy
which we are furnishing the cell merely goes to the decomposition
of water with the setting free of the component gases. This con-
dition is termed "boiling," which does not, as we have seen, refer
to any phenomenon due to heat, but merely to the setting free of
large volumes of gas. In charging a battery it should never be
allowed to fully discharge, about 30 per cent of the available energy
being left in every case. The voltage of a storage cell under load
should not be allowed to fall below 1.8. When a cell is left to stand
without doing work, it should be fully charged.
. There are several methods for calculating the capacity of a
storage cell in ampere-hours, and as they are purely electrochemical
in character should be touched upon before we close the present
chapter. The number of coulombs maintained by the consump-
tion of a chemically active substance varies with the change of valence
and inversely with the molecular weights of the transforming sub-
stance. According to Treadwell, the combustion or liberation of
one .pound of hydrogen corresponds to 12,160 ampere-hours.
The theoretical current capacity, therefore, in ampere-hours
may be figured from the following rule:
" F = the change of valence of the ions;
JF = the sum of the molecular weights affected, and
1 2, 1 60= the capacity per pound of hydrogen.
i2i6oXF
Then capacity per pound = ^ .
Regarding lead sulphate, which is the ultimate product at both
electrodes, we obtain as the value of the active material by the use
of the above formula 40.24 ampere-hours, or 80.48 watt-hours, per
pound of lead sulphate, with the lead-lead-sulphuric-acid battery.
For the calculation of electromotive force of storage cells we
have the following from "Chemical Theory of Accumulators," by
E. J. Wade.
THE SECONDARY CELL. 249
W = the work in joules.
(?=the coulombs of electricity that are passed through the
electrolyte.
H = ihe number of calories liberated by the recombination of a
unit weight of one of the decomposed ions.
e = its electrochemical equivalent.
c = its chemical equivalent.
& = the electrochemical equivalent of hydrogen
= .00001038.
/ = Joule's coefficient = 4.2
= the E.M.F. required.
W-QE,
W^QJeH-
therefore E = deH
and e = he;
therefore E = JhcH = 4.2 X .ooooio38c.H"
= .0000436^.
heat of formation
Now, cH = -
E =
valency
. 0000436 X heat of formation
valency
Since nearly all the battery equations are expressed in terms of the
transfer of two atoms of hydrogen, or their equivalent (that is,
they are bivalent), and since
.0000436X46000 = 1 volt
heat of formation in calories
we have E = 7 .
46000
CHAPTER XVIII.
ELECTRICITY 'FROM CARBON.
THERE are few problems known to man that promise such a
fruitful reward to the advancement of scientific triumph as the one
pertaining to the direct conversion of the energy of carbon into elec-
tricity. The benefits to be obtained from the solution of this greatest
of electrochemical possibilities are practically without words to
express. v We have on earth but few prime movers when we carefully
look into- and consider the situation. By prime movers we mean a
direct source of energy. Electricity in its dynamic form, and with
its applications for purposes of doing work, is but a secondary power.
When "we consider the question from the basic point of view, we find
as tabulated by Joseph Henry in his Scientific Writings, published
by the Smithsonian Institution, the following interesting and historic
list:
1. Water-power.
2. Wind-power.
3. Tide-power.
4. The power of combustion.
5. The power of vital action.
Continuing, he writes: "To this list may hereafter be added the
power of the volcano and the internal heat of the earth ; and, besides
these, science at the present time gives no indications of any other.
These are denominated primary powers, though in reality, when
critically studied, they may all, except the two last mentioned, be
referred to actions from without the earth, and principally to emana-
tions from the sun.
" Gravitation, electricity, galvanism, magnetism, and chemical
affinity can never be employed as original sources of power. At
the surface of the earth they are forces of quiescence, the normal
250
ELECTRICITY FROM CARBON. 251
condition of which must be disturbed before they can manifest
power, and then the work which they are capable of performing is
only the equivalent of the power which was communicated to them.
" There is no more prevalent and mischievous error than the idea
that there is in what are called the 'imponderables' a principle of
spontaneous activity. Heat is the product of chemical action, and
electricity only manifests power when its equilibrium is disturbed
by an extraneous force, and then the effect is only proportional to
the disturbing cause. It was for this reason that the existence of
electricity remained so long unknown to man. Though electricity is
not in itself a source of power, yet from its extreme mobility and high
elasticity it affords the means of transmitting power with scarcely any
loss and almost inconceivable velocity to the greatest distance. A
wave of disturbance starting from the impulse given at the battery
will traverse the circumference of the earth in less time than I have
been occupied in stating the fact.
" Besides electricity and the principle before mentioned, there are
other agents employed between the primary power and the work,
namely, the elastic force of steam, of air, and of springs ; also vari-
ous instruments called machines. But these must not be confounded,
as they frequently are, with the sources of power. It is not the
engine which is the source of motion of the cars, nor yet the steam,
but the repulsive energy imparted to the expanding water from the
burning fuel."
Through the agency of the steam-engine we obtain mechanical
energy which drives the dynamo for the production of electricity.
In this system, with its several transformations, the electricity
delivered at the terminals of the machine is traced back through
the rotating armature with its necessary losses, through the belt or
shaft to the engine with its friction and radiation, to the boiler with
its many sources of waste, and thence to the grate-bars where heat
results from the oxidation of the carbon supported there. It is
estimated that only 5 or 6 per cent of the available energy of the
carbon is transformed into useful work. If now it were possible
within a suitable cell to obtain as a result of the oxidation of carbon
the electric current direct without the production of heat, what a
majestic discovery would be made!
There is no principle of science standing in our way, and yet
252 EXPERIMENTAL ELECTROCHEMISTRY.
the brightest minds in physical and chemical science have been
unable to solve the problem. Heat energy may be transformed into
electrical energy, mechanical energy into heat energy, mechanical
energy into electrical, etc., etc., etc., without gain or loss, as we
learned from the great doctrine of the conservation and correlation
of energy. Instead of getting heat energy from carbon we can,
as far as existing conditions of science indicate, obtain electricity
direct. From the grate-bars, boiler, engine, and dynamo with their
combined miserable showing of 5 or 6 per cent energy yield, we may
turn for an instant to the primary cell. Here we obtain electricity
direct from what we may term the combustion of zinc, by immers-
ing pure zinc and platinum electrodes in dilute sulphuric acid,
with a current yield as high as 90 per cent of that theoretically
possible. Zinc is too costly a fuel, however, except in special cases
and upon a very small scale, and it has been the goal of chemists
and physicists to discover a method for obtaining electricity as a
direct result of the combustion of carbon or coal. In taking up the
consideration of this problem in the present chapter, the writer
wishes to warn the student not to confuse thermoelectric deport-
ment with the problem we really have in hand. Many investi-
gators have obtained electricity by using carbon rods in connection
with electrodes of different composition immersed in a fused elec-
trolyte, but the source of electricity has been ultimately traced, not
to the primary combustion of carbon, but to thermoelectric action.
Because of the still existing likelihood of research students going
astray in working on this great problem, it may be wise to introduce
the subject of thermoelectric action first, and acquaint him with the
conditions for its existence before the problem of the direct con-
version of the energy of carbon into electricity is dealt with. It
was in 1821 that Prof. Seeback of Berlin, noticed that by heating
a junction of two metals in a circuit an electric current was produced.
The thermoelectric current has been shown by heating the junction
between two dissimilar metals of a circuit which surrounds a mag-
netic needle whereby its direction is noted. If now the source of
heat be removed from this junction, and it be cooled by a little ice
or some absorbent cotton moistened with ether, an electrical current
will be indicated by the magnetic needle but of opposite direction
of flow. A thermoelectric couple for experimental purposes may
ELECTRICITY FROM CARBON.
2 53
be easily constructed by joining together two different metals at
their lower ends, as indicated in Fig. 118, and leading wires from
the upper ends of a galvanometer, or milli-volt-meter. If the lower
junction is now heated there will be a current set up, as indicated
by the little arrow, from the positive to the negative metal. The
electromotive force of this current mayi>e increased by connecting
together a number of bars, as shown at the right in this illustration,
and heating all the lower junctions and keeping the upper ones
FIG. 118. Typical Thermoelectric Couplers.
cool. It is upon this plan that the most delicate thermometers,
or heat detectors have been devised. Numerous designs of thermo-
electric batteries have also been produced, some few of them having
proven quite practical in operation, giving as a direct exchange
for heat energy energy in the form of electricity. Fig. 119 shows
an old form of thermoelectric battery of more scientific interest than
practical value. It has been found that a couple made of bismuth
and antimony heated at the point of union corresponds very closely
in electromotive force to a couple consisting of zinc and copper
immersed in sulphuric acid. Couples have been made of many of
the available metals, as well as from carbon, and electrical currents
obtained under different potentials. At the point of contact between
dissimilar conductors of electricity, whether of the metals or of carbon
and a metal, electricity is set up in the form of a current when such
contact or junction is heated. It is for us to realize this and take
254
EXPERIMENTAL ELECTROCHEMISTRY.
it into account in any work we may do with direct conversion of
carbon into electricity. Of all the attempts to reach a solution
of this great problem by men with every type of equipment, from
the haphazard-try-and-learn character to the research worker with
every theoretical equipment, only failure has resulted. One of the
great difficulties in the way of solving this problem is in the fact
that carbon does not dissolve in suitable electrolytes by the simple
throwing off of the positive ions of the carbon.
FlG. 119. Historic Thermoelectric Battery.
As early as 1855 Bacquerel conducted an experiment as illustrated
in Fig. 1 20. A rod of carbon was immersed in a bath of fused
nitre contained in an iron or platinum spoon in order to bring about
the oxidation of the carbon and the production of the electric current.
A galvanometer included in the circuit gave a marked indication.
In 1877 Jablochkoff repeated the experiment on a large scale and
under certain modified conditions. It will be interesting for us to
ELECTRICITY FROM CARBON.
2 55
note at this time the possible theoretical electromotive force obtain-
able from a carbon cell when the carbon is oxidized completely to
carbon dioxide. From thermochemical data, it has been estimated
that a cell in which carbon is completely oxidized to carbon dioxide,
the electromotive force obtainable is 1.04 volts. In the majority
of experiments which have been made with carbon for the direct
production of electricity, the most incomplete data are given. Apart
from electromotive force indications on open circuit regarding the
deportment of cells, no current data are to be found. One of the
most recent carbon cells is that of W. W. Jacques, in which a carbon
rod is immersed in fused sodium hydroxide, contained within an
iron vessel, heated externally by a furnace; the iron pot serving as a
FIG. 120. Bacquerel's Experiment on the Oxidation of Carbon to Electricity.
positive electrode, while the carbon is the negative. Oxygen is
pumped to the bottom of the fused caustic soda through a tube ending
in a ring around the lower end of the carbon, the ring being provided
with perforations, through which oxygen under pressure escapes
and comes in contact with the heated carbon. There is an opening
in the cover of this cell whereby the gaseous products of oxidation
escape. Of course, air may be pumped into the electrolyte in the
place of oxygen, although, for obvious reasons, it is not claimed to
be so efficient. According to Jacques, the carbon within is oxidized
to carbon dioxide, and the energy produced is given in the form of
an electrical current. It was claimed that with a battery of 100 cells
a current of about 15 amperes was obtained for 18 hours under
256 EXPERIMENTAL ELECTROCHEMISTRY.
a difference of potential of 90 volts, the consumption of carbon
being about 8 pounds. This performance corresponds with an
efficiency of about 80 per cent based on that theoretically possible
from the weight of carbon oxidized. This calculation does not
take into account the energy expended in the form of heat to keep
the electrolyte fluid or the energy necessary to pump the air or
oxygen through the fused salt. The device is not what has been
claimed for it, as it is entirely wrong in principle. Research upon
this cell has proven that the electricity does not come at all from the
oxidation of the carbon, but from purely thermochemical action.
The substitution of other electrodes in the place of the carbon rod
gives us practically the same current. If the carbon was oxidized
to carbon dioxide gas in the molten, caustic soda, we know, as general
chemists, that it would be rapidly converted into sodium carbonate,
which would ruin it for further usefulness.
Probably the majority of research workers on this problem have
immersed a carbon rod against some metal in a strong oxidizing
agent, fused nitre for example, in the hope that the oxidation of the
carbon would take place with the production of the electric current.
In other words, the carbon is brought in direct contact with oxidizing
agents. Such a course is manifestly wrong in principle when we look
a little into the question of primary cells. If, for example, we wish
to obtain electricity by the consumption of zinc, we may do so in
several ways. We would undoubtedly obtain a current of electricity
if we put a rod of zinc and a strip of platinum in strong nitric acid,
but the practice would be a most wasteful one. There would be
some electrical current produced, but most of the energy would
appear as heat, as the result of the energetic local action of the
nitric acid on the zinc. To quote Ostwald upon the subject of
electricity from carbon, he says: "The carbon cell of the future
must have an oxidizing agent in the place where the carbon is not."
To make this statement clear he refers the reader to the experiment
where zinc and platinum are immersed in two separate vessels con-
taining a solution of potassium sulphate and separated by a siphon-
tube, as illustrated in Fig. no, Chapter XVI. In order to bring
about the economical consumption of zinc in this system we learn
that the sulphuric acid must not be added to the beaker containing
the zinc, but to that containing the platinum. It is very evident that
ELECTRICITY FROM CARBON. 257
if we wish to obtain electricity direct from carbon, as we do from
zinc in the primary cell, some electrolyte must be discovered in which
carbon will dissolve with the formation of ions. It is to-day uncer-
tain whether or not carbon ionizes in certain solutions under the
action of electricity, opinion being divided upon this question. Al-
though the majority of electrochemists do- not believe in the ionization
of carbon there are certain experiments which tend to show the
contrary. Papasogli and Bartoli noticed that the passage of an
electric current between carbon electrodes immersed in dilute sul-
phuric acid was accompanied by certain marked chemical changes.
The carbon serving as anode is believed to go into solution for the
reason that carbon monoxide and carbon dioxide gases appear as
anode products together with the oxygen. Coehn, in working
along these lines was able, under certain conditions, to effect the con-
sumption of carbon in dilute sulphuric acid, with the evolution of
carbon monoxide and carbon dioxide gases, in the proportion of
30 per cent carbon monoxide and about 70 per cent carbon dioxide,
an evolution of only about i per cent of free oxygen being found.
The acid assumed a reddish-brown color believed to be due to the
actual dissolution of the carbon. If the electrolysis is allowed to
proceed with the carbon as anode, after substituting a platinum
cathode, the latter becomes covered with a black deposit. This has
been analyzed and found to consist of carbon containing occluded
oxygen and hydrogen gases in the proportion in which they combine
to form water. Coehn also showed that an electrode of carbon and
one of lead peroxide in dilute sulphuric acid produced a constant
current until the lead peroxide was reduced to lead, or else the carbon
consumed. Of all the attempts to dissolve carbon to obtain the energy
of oxidation as electricity, none of them, however, have met with
any real success. As pointed out, there are numberless cases of
experimental research with the most varied type of cell for the
oxidation of carbon, but the electrical data obtainable are exceedingly
meagre. In most of the reports of investigators, only a statement of
electromotive force is given on open circuits, whereas, we know elec-
trical data pertaining to batteries are useless unless we have together
with the electromotive force a statement regarding the number of
coulombs the cell will furnish. As many of the cells have been
upon the general principle of Jacques, an illustration of one of them
2 S 8
EXPERIMENTAL ELECTROCHEMISTRY.
is deemed of interest here. Fig. 121 represents Edison's design,
where we have a furnace for heating the iron melting-pot. The
cover is of an insulating material and supports the carbon electrode.
The electrolyte is chosen from the oxidizing agents, like nitre, for
example, or even certain oxides, and, according to Edison, a reduction
of the compound takes place, the oxygen combining with the carbon
FIG. 121. Edison's Cell for Obtaining Electricity from Carbon.
in the formation of carbon monoxide, which may be piped off and
used for fuel. The residue resulting from the reduction of the
oxide may be used over again as the negative agent of the cell.
From what we have seen this is entirely wrong in principle, and
the electricity obtained is primarily of thermoelectric origin. The
problem to-day is one of possible solution, but no practical results
have as yet been obtained. To enumerate the various researches
upon this important problem would fill a book in itself.
CHAPTER XIX.
USEFUL PIECES OF APPARATUS.
As has been frequently indicated throughout this experimental
work, the writer believes that through the introduction of con-
venient and useful types of apparatus of various kinds, the student
often obtains valuable information in the way of suggestions
for certain lines of research. It is believed that in many cases the
research student receives much benefit from a diagram of apparatus,
especially if he possesses ingenuity and has initiative, so to speak,
for investigation. There will be many problems in electrochemistry
confronting the experimenter, some of which require special types
of cells in order that certain conditions of electrolysis may be
established. There are, in addition, pieces of apparatus useful in
saving time because of their most convenient design and application.
For rapidly comparing the conductivities of various electrolytes in
small quantities it would be difficult to design a more handy and
convenient device than that illustrated in Fig. 122. This cut repre-
sents a special form of pipette equipped with a ground glass stop-cock
in order that it may easily be held filled by closing the stop-cock.
Two simple platinum disks welded to stout platinum wires are
sealed into the glass, facing one another, as shown, from opposite
sides. Two bent glass tubes fused on the outside of the bulb, are de-
signed to receive the mercury into which protrude the platinum wires
passing through the sides of the glass. It will be evident that we
can, without danger of breaking off any platinum terminals, make
and break electrical connection as often as we see fit by simply
immersing our conductors in the tubes and allowing them to dip into
the mercury. Another design of cell for conductivity determination
is shown in Fig. 123 where we have a special U-shape design of
tube with platinum disks, horizontally hung from platinum wires
259
260
EXPERIMENTAL ELECTROCHEMISTRY.
sealed into the bottoms of glass tubes, which in turn are supported
by turned wooden stoppers. These tubes also receive a small
quantity of metallic mercury, by means of which contact is made
with the electric battery or other source of current.
Fig. 124 shows still another form of cell for conductivity deter-
D
H
\
FIG. 122.
FIG. 124.
minations, and has the advantage of allowing for temperature deter-
mination being made at both anode and cathode. For this purpose
it is only necessary to remove the glass stoppers and insert the
thermometers. Such a conductivity cell may be separated into
two useful parts by cutting through the glass connecting neck
with a sharp file and joining them at almost any distance from
one another by inserting a glass tube and rubber connectors. We
will then have a piece of apparatus enabling us to obtain anode
USEFUL PIECES OF APPARATUS.
261
and cathode temperatures, to measure absolute velocity of the ions,
etc., etc.
Fig. 125 shows a conductivity cell for very accurate determina-
tions, as we have here means not only for making a careful tempera-
ture observation and correction, for keeping the temperature constant
by immersing the cell within a calorimeter receptical, or glass cylin-
der, which may be packed with ice.
FIG. 125.
FIG. 126.
Fig. 126 illustrates still another type of conductivity cell, which
consists of a cylindrical glass vessel containing two large platinum
disks fitted loosely within the interior and supported respectively
by two glass tubes, through which platinum stems are fused. These
tubes, like those in the previous chapter, are designed to be filled,
or partly filled, with mercury for the purpose of electrical contact,
and the turned wooden top with holes to receive these tubes with a
tight fit completes the equipment. The center tube carries an ordi-
nary disk, as shown, suspended by a platinum stem from its center,
but the tube on the right has its disk suspended from a point eccentric,
262
EXPERIMENTAL ELECTROCHEMISTRY.
the center of the disk being punched out to allow the middle glass
tube carrying the lower disk to pass through. By means of such an
arrangement the disks may be brought near together or separated,
and apart from serving the purpose of a useful conductivity cell,
may be employed also as a very handy and desirable electrolytic
rheostat for delicate work. Every electrolytic laboratory should
be equipped with a variety of ready decomposition tubes of various
shapes and patterns, to be to the electrochemist what the ordinary
test-tube is to the general chemist.
Fig. 127 illustrates a simple type of "electrolytic test-tube" of
FIG. 127.
FIG. 128.
V shape, equipped with platinum electrodes. Test-tubes of this pat-
tern are exceedingly useful when made from glass, only a few milli-
meters in diameter and 2 or 3 centimeters in height. There should
also be test tubes of this type several centimeters in diameter and 15
or 20 centimeters in height, for the tube is of such general utility
that it should be at hand in several sizes.
Fig. 128 is a useful design of cell not only for collecting gases
liberated at the electrodes, but also for solids which become detached
and fall to the bottom. It consists merely of a U tube equipped
with electrodes, and a glass bulb blown upon the lower extremity.
USEFUL PIECES OF APPARATUS.
263
Such a U tube equipped with a bulb serves a useful experimental pur-
pose when employed in the electrolysis of chloride of zinc where it is
present in a concentrated aqueous solution. When this is connected
with our lamp-bank and the current allowed to flow for a sufficiently
long time, the bulb will be filled with beautiful crystals of metallic
zinc, whereas, if we use a small U tube the branch-like growth would,
instead of breaking off and falling into the bulb, extend across to
the positive electrode and short-circuit the cell. It will, therefore,
be seen that such a U tube with bulb serves as a receptical for
certain electrode products, and therefore fills a useful purpose.
The U tube, as depicted in Fig. 129, is of special construction
FIG. 129.
and is designed to allow a removable porous partition to be placed
between the electrodes. This partition or diaphragm may be of
filter-paper, parchment, or even a disk of baked porous material,
as the requirements may dictate. The lower ends of the U tube
terminate in a bell-shape mouth, over which a clamp may be fixed
to draw them into close contact with the separating membrane.
Such a piece of apparatus will serve only in special cases, but it has
its application in research work and should be included in an equip-
ment.
Fig. 130 illustrates the front and side view of a Hoffmann appara-
tus, blown from one piece of glass, which deserves a special place of
264
EXPERIMENTAL ELECTROCHEMISTRY.
honor as a device of general utility. Although small pieces of
Hoffmann apparatus may be ordered from almost any dealer in
chemical ware, when necessary to employ such a cell on a large
scale, it will be convenient and easy to improvise one by using a long
length of glass combustion tubing of large bore. If a Woulf's bottle
with three wide necks is available, the assembling of a large Hoff-
r
Q
>___
'
A
a
T~ 1
T
T
FIG. 130.
mann apparatus is a comparatively simple matter, as will be seen
by any one with aptitude for construction. In place of the Woulf
bottle a large Hoffmann apparatus for the accumulation of large
quantities of electrode gas may be assembled from the combustion
tubing and a large T joint of glass with rubber connectors.
In closing this chapter a few words relative to the Wenhelt inter-
rupter and aluminum rectifier will perhaps be of interest and value
to us. Wenhelt's electrolytic interrupter is particularly useful
USEFUL PIECES OF APPARATUS.
265
in connection with the operation of induction-coils, for it serves
to take the place of the ordinary vibrator.
Fig. 131 shows a simple form of electrolytic interrupter, which
consists simply of a beaker glass containing dilute sulphuric acid,
into which is immersed strips of platinum to serve
as anode, while the cathode consists of a short
piece of platinum wire fused into the lower end
of a stout glass tube, bent as shown in the illustra-
tion. This tube is filled with mercury, both for
the purpose of making an electrical contact with
the platinum wire cathode and for conducting
away the heat which is generated there. When
interposed in an electrical circuit this device
serves to rapidly make and break the electric cur- FIG. 131.
rent. The principle of this interrupter is based
upon the rapid formation and discharge of bubbles of hydrogen from
the platinum point or cathode of the cell which, because of the high
current density existing, becomes quite hot. Such an interrupter
operates in a most satisfactory manner an induction-coil and other
devices where an intermittent current is desired. The adjustment
is brought about by varying the distance between the platinum strip
anode and the platinum cathode, together with changing the concen-
tration of the sulphuric acid solution. This apparatus is introduced
here as of scientific interest from an electrochemical standpoint, in the
belief that the student will profit by conducting experiments there-
with. Another electrolytic cell of remarkable performance is the
aluminum rectifier, which, when placed in series in the circuit of
an alternating current, converts the alternating current into an inter-
mittent direct current. This may be regarded as a species of elec-
trical " check- valve," allowing the impulses in one direction to pass
through and preventing the impulses in the opposite direction from
getting past. Such a rectifier consists simply of a beaker glass con-
taining a suitable electrolyte, into which is immersed an aluminum
anode and a platinum cathode. With such an equipment, using
disodium phosphate as an electrolyte, a potential difference as high
as 300 volts may be established between the electrodes, while only
a few hundredths of an ampere will flow through the cell in one
direction, because of its high resistance, whereas the resistance is
266 EXPERIMENTAL ELECTROCHEMISTRY.
comparatively zero in the reverse direction. Such a cell interposed
in an alternating current circuit with only no volts pressure is of
exceedingly great interest. It is necessary to keep these cells cool
by the circulation of water, and for this purpose the aluminum elec-
trode has been employed in the form of a U tube, through which cold
water may be made to flow. The principle of the rectifier is believed
to be based upon the formation of a non-conducting film upon the
aluminum, and which allows a large current to pass through in one
direction, but only an exceedingly small one in the other.
CHAPTER XX.
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267
268 EXPERIMENTAL ELECTROCHEMISTRY.
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276 EXPERIMENTAL ELECTROCHEMISTRY.
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278 EXPERIMENTAL ELECTROCHEMISTRY.
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INDEX.
Accumulator, 237
Acetylene, synthesis of, 211
Acids, ionization of, II
Acids, sulphuric, electrolysis of, 108
Acids, nitric, preparation of from atmos-
phere, 150
Acids and bases, neutralization of, 36
Acids and carbonates, inactive when dry,
43
Affinity, chemical, 49
Agassiz, 95
Alcohol as dissociant, 43
Alternating currents, rectification of, 265
Alternating currents and electrolytes, 63
Aluminum, isolation of, 174
Aluminum bronze, 184
Aluminum rectifier, 265
Ampere hour meter, 89
Andrews, 140
Anode, 9
Anode, definition of, 16
Anode, rotating, 112
Apparatus, useful, 259
Arc, temperature of, 265
Arrhenius, 5, 17
Artificial carnalite, preparation of, 105
Atomic heat, table of, 87
Attraction, chemical, 10
Avogadro's law, 20
Ayrton, 243
Azobenzene, 2l6
Bases, neutralization of, 36
Battery, chemistry and electrochemistry
of Daniell's, 234
Battery, chemistry of .storage, 243
Battery, density changes in storage, 242
Battery, depolarization in, 238'
Battery, discharge curve of storage, 241
Battery, polarization in, 237
Battery, primary, thebry of, 219
Battery, secondary, theory of, 237
Battery; thermo-electric, 252
Bauxite, 174
Becaria, I
Beckmann's thermometer, 33
Becquerel, 140, 254
Berthelot, 212
Bertoli, 257
Berzilius, 3
Bibliography, 267
Bladder, experiment with, 22
Bloxam, 145
Boiling-point, elevation of, 34
Boiling-point, influence by electrolytes, 34
Boiling-point, influence by non-electro-
lytes, 34
Borscher, 168, 170, 177, 191
Boyle, 20
Boyle's law, 20
Bradley, 158
Bredig, 75
Brugnatelli, 4
Buckling in storage-batteries, 246
Bumping, how to prevent, 36
Bunsen, 177, 187
Cadmium yellow, 126
Calcium, isolation of, 186
Calcium carbide, preparation of, 205
Calorie, 49, 102
Calorimeter for measuring neutralization
heats, 37
Carbide, 205
Carbon, electricity from, 250
Carbon, ionization of, 257
Carbon, solution of, 257
Carbon disulphide, preparation of, 213
Carlisle, 3
Carnalite, preparation of artificial, 105
Castner, 166
Cathode, 9
Cathode, definition of, 16
Caustic soda and chlorine from salt 121
Cavendish, 151
Cells, conductivity, 159
'Chemical action at a distance, " 224
Chemically active bodies, 18
Chemical activity and dissociation, 18
Chemical am'mty, 49
2/9
280
INDEX.
Chemical attraction, 10
Chemical equivalents, 83
Chemically inactive bodies, 18
Chemical indicator, reflecting galva-
nometer as, 63
Chloride of lime, 121
Chloroform, electrolytic preparation of,
211
Chronograph, diagram of connections -of =
Hopkins, 69
Chronograph, use in studying electro-
lytes, 68
Coehn, 257
Collins, 81
Conduction, electrodeless, 61
Conductivity, ^molecular, 49
Conductivity cell, 259
Conductivity and dissociation, 46
Conductors, first class, 17
Conductors, first and second class com-
pared, 68
Conductors, second class, 17
Conservation of energy, 101
Converter, rotary, 16
Copper ferrocyanide, use for semi-per-
meable membranes, 23
Coulomb, 92
Coulombmeter. See Voltameter.
Crookes, Sir William, 153
Cruikshank, 3
Cryolite, 174
Current, direction of, II
Current density, 117
Current indicator, 12
Daniell cell, chemistry and electrochem-
istry of, 235
Daniell cell, electromotive force of, 236
Davy, Sir Humphry, 4, 166
De la Rive, 140, 220
Depolarization in primary batteries, 238
Deville, 175, 177
Diaphragm, semi-permeable, 22
Diffusion, 22
Dissociation and chemical activity, 18
Dissociation and conductivity, 46
Dissociation theory and the primary cell,
222
Dissociant, 43, 49
Dissociants, table of, 45
Distilled water as conductor, 18
Double layer, Helmholtz's, 229
Dulong, 87
Dye, yellow, preparation of, 215
Edison, 258
Elbs, 136
Electric furnace, 194
Electric furnace, Hopkin's "series car-
bon," 201
Electric-lighting current, how to adapt
for electrolytic work, 14
Electrical energy, correlation of, 253
Electricity from carbon, 250
Electricity, thermo, 252
Electrochemical equivalent, 8 1
Electrochemical equivalents, table of, 84
Electrochemical equivalent of electricity,
82
Electrochemical order of elements, 10
Electrode, definition of, 16
Electrode, normal, 230
Electrode, negative, 9
Electrode, positive, 9
EJgc_trode area, influence of. See Cur-
rent Density.
Electrode tension in separation of metals,
in
Electrodeless conduction, Hopkins' ex-
periment to show, 6 1
Electrolysis, conditions governing success
in, 117
Electrolysis, definition of, 16
Electrolysis, energy absorbed in, 92, 100
Electrolytes, n
Electrolytes, definition of, 1 6
Electrolytes, alternating currents with
change of frequency, 63
Electrolytes, frozen, 94
Electrolytes, fused, 102
Electrolytes, law of conductivity, 75
Electrolytes, lowering of freezing-point
by, 32
Electrolytes, table of, 18
Electrolytes, Wheatstone's bridge ap-
plied to, 73
Electrolytes and alternating currents. 63
Electrolytes and non.electrolytes, table
of, 18
Electrolytic conductivity, definition of, 16
Electrolytic dissociation, definition of, 16
Electrolytic dissociation, evidence for, 17
Electrolytic dissociation, experiment in
support of, 40
Electrolytic dissociation and gas laws, 20
Electrolytic dissociation and heat neu-
tralization, 38
Electrolytic induction, 50
Electrolytic induction, effect on magnetic
needle, 59
Electrolytic oxidation, 215
Electrolytic reduction, 216
Electrolytic separation of metals, no
Electro-magnetic deportment of sub-
stances in solution, 50
Electromotive force of Daniell cell, 236
Electromotive force, electrolysis, mini-
mum required in, 103
Electromotive force, method of calculating
in primary cells, 235
Electromotive force and solution tension,
223
Electrostatic deportment of substances in
solution, 50.
INDEX.
281
Elements, electrochemical order of, 10
Elements, negative, 10
Elements, positive, 10
Energy absorbed in electrolysis, 92, 100
Energy, conservation of, 101
Energy, correlation of, 253
Energy, electrical, correlation of, 253
Energy, heat, correlation of, 253
Energy, mechanical, correlation, 253
Ethane, production of, 209
Ethyl alcohol as dissociant, 45
Experiment, absolute velocity of ions,
Whetham's, 78
Experiment, absolute velocity of hydro-
gen ions, Lodge's, 76
Experiment, alternating currents at differ-
ent frequencies, Hopkins', 64
Experiment, bladder to show osmosis, 22
Experiment, chronograph in electrolytic
conduction, Hopkins', 68
Experiment, Davy's classic, Sir Hum-
phry, 4
Experiment, with dry gases, 41
Experiment, electrical conduction, in-
stantaneous, 67
Experiment, electrodeless conduction,
Hopkins', 62
Experiment, electrolysis of magnesium
chloride, 104
Experiment, electrolysis sodium hydrox-
ide, 9
Experiment, electrolytes effect on mag-
net, Hopkins', 56
Experiment, electrolytes frozen, Hop-
kins', 95
Experiment, electrolytes, magnetic pull
of, Hopkins', 61
Experiment, electrolytes, quantitative
effect of magnet on, Hopkins', 58
Experiment, electrolytes from non-elec-
trolytes, to distinguish, 19
Experiment, electrolytic conduction, effect
on magnetic needle, Hopkins', 59
Experiment, electrolytic dissociation, 47
Experiment, Faraday's law, to demon-
strate, 84
Experiment, free ions, Ostwald's, 51, 54
Experiment, free ions, Hopkins', 54
Experiment, glass as a conductor, Hop-
kins', 94
Experiment, heat convection in electro-
lytic conduction, Hopkins', 97
Experiment, matter, mechanical tran-
sport of, 93
Experiment, nitric acid from atmosphere,
160
Experiment, sulphuric acid, electrolysis
of, Hopkins', 108
Fabroni, 2
Faraday, Michael, 52, 80, 220
Faraday's law, 80
Faraday's law, experimental demonstra-
tion of, 84
Faraday's law and primary cell, 234
Feldspar, 174
Formic acid as dissociant, 45
Freezing-point, depression of, 32
Freezing-point, practical determination
of, 35
Fremy, 140
Frequency, in alternating currents ap-
plied to electrolytes, 63
Froelich, 142
Frozen electrolytes, 94
Furnace, electric, 194
Furnace, electric, Hopkins' series car-
bon, 201
Fused electrolytes, 102
Galvani, 2
Galvanometer, reflecting as chemical in-
dicator, 53
Gas laws, application to substances in
solution, 20
Gaugin, 53
Gay-Lussac's law, 20
Geber, 151
Gerboin, 3
Gladstone and tribe, 247
Glass, conductivity of, 93
Glauber, 151
Gore, 221
Grabau, 175 ^
Gram-molecular weight, 31
Gram-molecule, 48
Gunpowder incombustible when dry, 43
Hall, 180
Heat of combustion, 190
Heat convection in electrolytic conduc-
tion, 97
Heat energy, correlation of, 253
Heat of formation, 100
Heat of formation, table of, 102
Heat of formation of salts, 39
Heat of formation of water, 39
Heat, latent, 35, 48
Heat of neutralization, 36
Heat of neutralization and electrolytic
dissociation, 38
Heimrod, 8l
Helmholtz's double layer, 229
Henry, Joseph, 3, 250
Herault, 184
Hess, 100
Kissinger and Berzelius, 3
Hittorf, 114
Hoffman's apparatus, 216, 264
Hopkins' electric furnace, series carbon,
201
Hopkins' electrolysis of sulphuric acid,
107
283
INDEX.
Hopkins' electrode temperatures, method
for studying, 1 19
Hopkins' experiment to show dissociation
simultaneously in two ways, 47
Hopkins' experiment to show electrode-
less conduction, 62
Hopkins' experiment to show electrolytic
induction upon magnetic needle, 59
Hopkins' experiment in heat convection
in electrolytic conduction, 97
Hopkins' experiment to show velocity of
electrolytic conduction, 68
Hopkins' experiment to show deportment
of electrolytes with alternating currents
at various frequencies, 64
Hopkins' experiment with frozen electro-
lytes, 94
Hopkins' experiment to show effect of
magnet upon coil of electrolyte, 56
Hopkins' experiment to show magnetic
effects upon electrolytes, 58
Hopkins' experiment to show and meas-
ure pull of coiled electrolyte, 61
Hopkins' experiment with glass as con-
ductor, 93
Hopkins' experiment in magnetic induc-
tion, 56
Hopkins' experiment in static induction
to show free ions, 54
Hopkins' high speed chronograph, 68
Hopkins' reflecting hydrometer, 242
Hopkins' Soret apparatus, design of, 29
Hopkins' table of electrochemical condi-
tions to be noted, 120
Houzan, 145
Howies, 153
Hydrogen, occlusion in storage-batteries,
240
Hydrogen ion, absolute velocity of, 77
Hydrometer, reflecting, Hopkins, 242
Hysterisis of iron, 63
Indicator, reflecting galvanometer as
chemical, 53
Interrupter, Wenhelt's, 264
lodoform, electrolytic preparation of, 209
Ions, 8
Ions complex. II
Ions, conduction through an electrolyte,
instantaneous, due to, 67
Ions, definition of. 16
Ions, existence shown by experiments,
50, 56
Ions, heat, transported by, 97
Ions, hydrogen, absolute velocity of, 77
Ions, inertia of, 65
Ions, mechanical representation of migrat-
ing, Hopkins, 60
Ions, negative, 10
Ions, positive, 10
Ions, presence of free ions shown by
chronograph work, Hopkins, 68
Ions, simple, II
Ions, velocity of, absolute, 76, 97
Ions, velocity of, relative, 1 14
Ions, velocity, Whetham's method for
determining, 78
lonization of acids, hydrochloric, 1 1
lonization of acids, nitric, 1 1
lonization of acids, sulphuric, n, 108
lonization of carbon, 25 7
Isolation of aluminum, 174
Isolation of calcium, 186
Isolation of magnesium, 104
Isolation of sodium and potassium, 165
Jablochkoff, 254
Jacques, 255
Jahn, 209
Joule, 92, 102
Kahle, 81
Kanarin, electrolytic preparation of, 215
Kaolin, 174
Kohlrausch, 73, 8l
Kowalski, 158
Kuester, 54
Lamp-bank, application to lighting cir-
cuit, 13
Latent heat, 35, 48
Laws, Avogadro's, 20
Laws, Boyle's, 20
Laws, Dulong and Petit, 87
Laws, Faraday's, 80
Laws, Gay-Lussac's, 20
Laws, Hess's, 100
Laws, Ohm's, 14
Lead, white, electrolytic preparation of,
122
Le Blanc, in
Litmus paper unaffected by dry acids and
bases, 43
Lodge, 76
Lodge's method for determining absolute
velocity of hydrogen ion, 77
Lovejoy, 158
Lullius, 151
McDougall, 153
Magnesium chloride, electrolysis of, 104
Magnetic needle affected by electrolytic
induction, 59
Marignac, 140
Matthiessen, 187
Mechanical energy, correlation of, 253
Membranes, seinipermeable, 22
Membranes, semipermeable, method . of
preparing, 24
Methyl alcohol as dissociant, 45
Mercury as pole finder, 12
Mercury vermillion, 128
Metals, electrolytic separation of, no
Metals, table of tension series, 234
INDEX.
283
Metals and their salts, potentials between,
233
Meter, ampere hour, 89
Meter, calorimeter, 89
Meter, coulombmeter, 88
Meter, gas- type voltameter, 88
Meter, voltameter, 105
Mica, 174
Moisture, influence in reactions, 43
Molecular conductivity, 49
Molugram, 45, 48
Moscicki, 158
Motor-generator, 15
Motor- generator, diagram showing appli-
cation, 15
Nernst, Walter, 52, 87, 229
Neumann, 233
Neutralization of acids and bases, 36
Neutralization heats, calorimeter for de-
termining, 37
Newton, 153
Nicholson and Carlisle, 3
Nitric acid, ionization of, n
Nitric acid, production of from atmos-
phere, 150
Nitrobenzene, electrolysis of, 216
Non-electrolytes, freezing-point lowered
by, 32
Non-electrolytes, table of, 18
Normal electrode, 230
Normal solution, sugar, 27
Occlusion of hydrogen in secondary bat-
teries, 240
Oersted, 175
Ohm's law, 14
Organic compounds, preparation of, 207
Osmotic pressure, 21
Osmotic pressure, experiment with blad-
der, 22
Osmotic pressure, theory of, 26
Ostwald, Wilhelm, 50, 66, 87, 224, 256
Ostwald's experiment to show existence
of free ions, 50
Ostwald's and Nernst' s experiment to show
existence of free ions, 53
Oxidation, 118
Oxidation, electrolytic, 215
Oxidation and reduction, 1 18
Oxygen non-supporter of combustion in
moisture free substances, 43
Ozone, 139
Ozone, commercial production of, 146
Ozone, experimental production of, 139
Paetz and Van Troostvik, I
Papasogli and Bertoli, 257
Passive state, 8
Patterson, 81
Pellat, 8l
Petit, 87
Pfeffer, 23
Phenolphthalein as a chemical indicator,
46
Plante, 244
Polarization, III
Polarization in batteries, 237
Polarization currents, in
Pole finder, electrolytic, 9
Porous pots, use in preparing semiper-
meable membranes, 23
Portier, 8 1
Potassium, isolation of, 165
Potassium chlorate, 133
Potentials between metals and their salts,
233
Power, primary, 250
Primary cell, 219
Primary cell, theory of, 226
Primary cell and Faraday's law, 234
Prime movers, 250
Principle of Soret, 28
Prussian blue, 131
Ramsay, 158
Rauolt, 32
Rayleigh, Lord, 8l, 153
Reactions, chemical, moisture influence
in, 43
Rectification of alternating currents, 265
Rectifier, aluminum, 265
Reduction, 118
Reduction, electrolytic, 216
Reflecting galvanometer as a chemical
indicator, 54
Reflecting hydrometer, Hopkins, 242
Resistance, lamp-bank as, 13
Rheostat, useful, 262^
Richards, 81
Rose, 175
Rotary -converter, .16
Rotating anode, 112
Salts, table showing heat of formation, 39
Scheele's green, preparation of, 130
Schoenbein, 139
Secondary cell, 237
Secondary cell. See Storage battery
Sedgewick, Mrs., 8l
Seeback, 252
Sellon, 243
Semipermeable membrane, 22
Semipermeable membrane, method of
preparing, 24
Semipermeable membranes, use of porous
pots in preparing, 23
Siemen's tubes, 142
Silver, electrochemical equivalent of, 8 1
Sodium, isolation of, 165
Sodium acetate, electrolysis of. 207
Sodium hydroxide, electrolysis, minimum
voltage for. 172
Solution tension of metals, 222
284
INDEX.
Solution tension and electromotive force,
223
Soret, 28, 140
Soret, principle of, 28
Specific heat, 49
Specific inductive capacity, 49
Steam-engine, non-efficiency of, 25 1
Storage-battery, "boiling" in, 248
Storage-batteries, "buckling" in, 246
Storage-batteries, calculation for capacity,
248
Storage-batteries, care in charging, 247
Storage-batteries, chemistry of, 243
Storage-batteries, density changes in, 242
Storage-batteries, discharge curve of, 241
Storage-batteries, electromotive force,
calculation of, 248
Storage-batteries. See Secondary cells
Sugar, table showing osmotic pressure
of, 26
Sulphur trioxide, 137
Sulphuric acid, electrolysis of, 197
Sulphuric acid, ionization of, n, 108
Sulzer, I
Synthesis of acetylene, 21 1
Tait, 140
Taylor, 214
Temperature of electric arc, 205
Tension series of metals, 234
Tension solution of metals, 222
Tension, vapor, 222
Thermoelectric current, 252
Thermostat, air-bulb, 30
Thomson, J. J., 43
Trauve, Moritz, 23
Treadwell, 245
Tyndall, 95
Van't Hoff, 27, 30
Vapor tension of liquids, 222
Volt-coulomb. See Joule
Volta, Alexander, 2
Voltameter, metal and gas type, 88
Voltmeter, 105
Von Helmholtz, 80
Von Marum, 139
Von Troostvik, i
Wade, 248
Warburg, 148
Water as conductor, 18
Water as dissociant, 45
Water, feeezing-point lowered, 32
Water, heat of formation, 39
Watt, 1 66
Welhelt's interrupter, 264
Wheatstone's bridge, application to elec-
trolytes, 73
Whetham, 78
Wohler, 175
Wollaston, 3
Zosimus, I
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