[1] See J.A. Fleming, _The Alternate Current Transformer_, vol. i. p.
519.
[2] See Maxwell, _Electricity and Magnetism_, vol. ii. chap. ii.
[3] See Maxwell, _Electricity and Magnetism_, vol. ii. 642.
[4] _Experimental Researches_, vol. i. ser. 1.
[5] See Maxwell, _Electricity and Magnetism_, vol. ii. S 542, p. 178.
[6] See W.G. Rhodes, _An Elementary Treatise on Alternating Currents_
(London, 1902), chap. vii.
[7] See J.A. Fleming, "Problems on the Distribution of Electric
Currents in Networks of Conductors," _Phil. Mag_. (1885), or Proc.
Phys. Soc. Lond. (1885), 7; also Maxwell, _Electricity and Magnetism_
(2nd ed.), vol. i. p. 374, S 280, 282b.
[8] See _Berl. Acad. Ber._, 1876, p. 211; also H.A. Rowland and C.T.
Hutchinson, "On the Electromagnetic Effect of Convection Currents,"
_Phil. Mag._, 1889, 27, p. 445.
ELECTROLIER, a fixture, usually pendent from the ceiling, for holding
electric lamps. The word is analogous to chandelier, from which indeed
it was formed.
ELECTROLYSIS (formed from Gr. [Greek: lyein], to loosen). When the
passage of an electric current through a substance is accompanied by
definite chemical changes which are independent of the heating effects
of the current, the process is known as _electrolysis_, and the
substance is called an _electrolyte_. As an example we may take the case
of a solution of a salt such as copper sulphate in water, through which
an electric current is passed between copper plates. We shall then
observe the following phenomena. (1) The bulk of the solution is
unaltered, except that its temperature may be raised owing to the usual
heating effect which is proportional to the square of the strength of
the current. (2) The copper plate by which the current is said to enter
the solution, i.e. the plate attached to the so-called positive terminal
of the battery or other source of current, dissolves away, the copper
going into solution as copper sulphate. (3) Copper is deposited on the
surface of the other plate, being obtained from the solution. (4)
Changes in concentration are produced in the neighbourhood of the two
plates or electrodes. In the case we have chosen, the solution becomes
stronger near the anode, or electrode at which the current enters, and
weaker near the cathode, or electrode at which it leaves the solution.
If, instead of using copper electrodes, we take plates of platinum,
copper is still deposited on the cathode; but, instead of the anode
dissolving, free sulphuric acid appears in the neighbouring solution,
and oxygen gas is evolved at the surface of the platinum plate.
With other electrolytes similar phenomena appear, though the primary
chemical changes may be masked by secondary actions. Thus, with a dilute
solution of sulphuric acid and platinum electrodes, hydrogen gas is
evolved at the cathode, while, as the result of a secondary action on
the anode, sulphuric acid is there re-formed, and oxygen gas evolved.
Again, with the solution of a salt such as sodium chloride, the sodium,
which is primarily liberated at the cathode, decomposes the water and
evolves hydrogen, while the chlorine may be evolved as such, may
dissolve the anode, or may liberate oxygen from the water, according to
the nature of the plate and the concentration of the solution.
_Early History of Electrolysis._--Alessandro Volta of Pavia discovered
the electric battery in the year 1800, and thus placed the means of
maintaining a steady electric current in the hands of investigators,
who, before that date, had been restricted to the study of the isolated
electric charges given by frictional electric machines. Volta's cell
consists essentially of two plates of different metals, such as zinc and
copper, connected by an electrolyte such as a solution of salt or acid.
Immediately on its discovery intense interest was aroused in the new
invention, and the chemical effects of electric currents were speedily
detected. W. Nicholson and Sir A. Carlisle found that hydrogen and
oxygen were evolved at the surfaces of gold and platinum wires connected
with the terminals of a battery and dipped in water. The volume of the
hydrogen was about double that of the oxygen, and, since this is the
ratio in which these elements are combined in water, it was concluded
that the process consisted essentially in the decomposition of water.
They also noticed that a similar kind of chemical action went on in the
battery itself. Soon afterwards, William Cruickshank decomposed the
magnesium, sodium and ammonium chlorides, and precipitated silver and
copper from their solutions--an observation which led to the process of
electroplating. He also found that the liquid round the anode became
acid, and that round the cathode alkaline. In 1804 W. Hisinger and J.J.
Berzelius stated that neutral salt solutions could be decomposed by
electricity, the acid appearing at one pole and the metal at the other.
This observation showed that nascent hydrogen was not, as had been
supposed, the primary cause of the separation of metals from their
solutions, but that the action consisted in a direct decomposition into
metal and acid. During the earliest investigation of the subject it was
thought that, since hydrogen and oxygen were usually evolved, the
electrolysis of solutions of acids and alkalis was to be regarded as a
direct decomposition of water. In 1806 Sir Humphry Davy proved that the
formation of acid and alkali when water was electrolysed was due to
saline impurities in the water. He had shown previously that
decomposition of water could be effected although the two poles were
placed in separate vessels connected by moistened threads. In 1807 he
decomposed potash and soda, previously considered to be elements, by
passing the current from a powerful battery through the moistened
solids, and thus isolated the metals potassium and sodium.
The electromotive force of Volta's simple cell falls off rapidly when
the cell is used, and this phenomenon was shown to be due to the
accumulation at the metal plates of the products of chemical changes in
the cell itself. This reverse electromotive force of polarization is
produced in all electrolytes when the passage of the current changes the
nature of the electrodes. In batteries which use acids as the
electrolyte, a film of hydrogen tends to be deposited on the copper or
platinum electrode; but, to obtain a constant electromotive force,
several means were soon devised of preventing the formation of the film.
Constant cells may be divided into two groups, according as their action
is chemical (as in the bichromate cell, where the hydrogen is converted
into water by an oxidizing agent placed in a porous pot round the carbon
plate) or electrochemical (as in Daniell's cell, where a copper plate is
surrounded by a solution of copper sulphate, and the hydrogen, instead
of being liberated, replaces copper, which is deposited on the plate
from the solution).
[Illustration: FIG. 1.]
_Faraday's Laws._--The first exact quantitative study of electrolytic
phenomena was made about 1830 by Michael Faraday (_Experimental
Researches_, 1833). When an electric current flows round a circuit,
there is no accumulation of electricity anywhere in the circuit, hence
the current strength is everywhere the same, and we may picture the
current as analogous to the flow of an incompressible fluid. Acting on
this view, Faraday set himself to examine the relation between the flow
of electricity round the circuit and the amount of chemical
decomposition. He passed the current driven by a voltaic battery ZnPt
(fig. 1) through two branches containing the two electrolytic cells A
and B. The reunited current was then led through another cell C, in
which the strength of the current must be the sum of those in the arms A
and B. Faraday found that the mass of substance liberated at the
electrodes in the cell C was equal to the sum of the masses liberated in
the cells A and B. He also found that, for the same current, the amount
of chemical action was independent of the size of the electrodes and
proportional to the time that the current flowed. Regarding the current
as the passage of a certain amount of electricity per second, it will be
seen that the results of all these experiments may be summed up in the
statement that the amount of chemical action is proportional to the
quantity of electricity which passes through the cell.
Faraday's next step was to pass the same current through different
electrolytes in series. He found that the amounts of the substances
liberated in each cell were proportional to the chemical equivalent
weights of those substances. Thus, if the current be passed through
dilute sulphuric acid between hydrogen electrodes, and through a
solution of copper sulphate, it will be found that the mass of hydrogen
evolved in the first cell is to the mass of copper deposited in the
second as 1 is to 31.8. Now this ratio is the same as that which gives
the relative chemical equivalents of hydrogen and copper, for 1 gramme
of hydrogen and 31.8 grammes of copper unite chemically with the same
weight of any acid radicle such as chlorine or the sulphuric group, SO4.
Faraday examined also the electrolysis of certain fused salts such as
lead chloride and silver chloride. Similar relations were found to hold
and the amounts of chemical change to be the same for the same electric
transfer as in the case of solutions.
We may sum up the chief results of Faraday's work in the statements
known as Faraday's laws: The mass of substance liberated from an
electrolyte by the passage of a current is proportional (1) to the total
quantity of electricity which passes through the electrolyte, and (2) to
the chemical equivalent weight of the substance liberated.
Since Faraday's time his laws have been confirmed by modern research,
and in favourable cases have been shown to hold good with an accuracy of
at least one part in a thousand. The principal object of this more
recent research has been the determination of the quantitative amount of
chemical change associated with the passage for a given time of a
current of strength known in electromagnetic units. It is found that the
most accurate and convenient apparatus to use is a platinum bowl filled
with a solution of silver nitrate containing about fifteen parts of the
salt to one hundred of water. Into the solution dips a silver plate
wrapped in filter paper, and the current is passed from the silver plate
as anode to the bowl as cathode. The bowl is weighed before and after
the passage of the current, and the increase gives the mass of silver
deposited. The mean result of the best determinations shows that when a
current of one ampere is passed for one second, a mass of silver is
deposited equal to 0.001118 gramme. So accurate and convenient is this
determination that it is now used conversely as a practical definition
of the ampere, which (defined theoretically in terms of magnetic force)
is defined practically as the current which in one second deposits 1.118
milligramme of silver.
Taking the chemical equivalent weight of silver, as determined by
chemical experiments, to be 107.92, the result described gives as the
electrochemical equivalent of an ion of unit chemical equivalent the
value 1.036 X 10^(-5). If, as is now usual, we take the equivalent
weight of oxygen as our standard and call it 16, the equivalent weight
of hydrogen is 1.008, and its electrochemical equivalent is 1.044 X
10^(-5). The electrochemical equivalent of any other substance, whether
element or compound, may be found by multiplying its chemical equivalent
by 1.036 X 10^(-5). If, instead of the ampere, we take the C.G.S.
electromagnetic unit of current, this number becomes 1.036 X 10^(-4).
_Chemical Nature of the Ions._--A study of the products of decomposition
does not necessarily lead directly to a knowledge of the ions actually
employed in carrying the current through the electrolyte. Since the
electric forces are active throughout the whole solution, all the ions
must come under its influence and therefore move, but their separation
from the electrodes is determined by the electromotive force needed to
liberate them. Thus, as long as every ion of the solution is present in
the layer of liquid next the electrode, the one which responds to the
least electromotive force will alone be set free. When the amount of
this ion in the surface layer becomes too small to carry all the current
across the junction, other ions must also be used, and either they or
their secondary products will appear also at the electrode. In aqueous
solutions, for instance, a few hydrogen (H) and hydroxyl (OH) ions
derived from the water are always present, and will be liberated if the
other ions require a higher decomposition voltage and the current be
kept so small that hydrogen and hydroxyl ions can be formed fast enough
to carry all the current across the junction between solution and
electrode.
The issue is also obscured in another way. When the ions are set free at
the electrodes, they may unite with the substance of the electrode or
with some constituent of the solution to form secondary products. Thus
the hydroxyl mentioned above decomposes into water and oxygen, and the
chlorine produced by the electrolysis of a chloride may attack the metal
of the anode. This leads us to examine more closely the part played by
water in the electrolysis of aqueous solutions. Distilled water is a
very bad conductor, though, even when great care is taken to remove all
dissolved bodies, there is evidence to show that some part of the trace
of conductivity remaining is due to the water itself. By careful
redistillation F. Kohlrausch has prepared water of which the
conductivity compared with that of mercury was only 0.40 X 10^(-11) at
18 deg. C. Even here some little impurity was present, and the
conductivity of chemically pure water was estimated by thermodynamic
reasoning as 0.36 X 10^(-11) at 18 deg. C. As we shall see later, the
conductivity of very dilute salt solutions is proportional to the
concentration, so that it is probable that, in most cases, practically
all the current is carried by the salt. At the electrodes, however, the
small quantity of hydrogen and hydroxyl ions from the water are
liberated first in cases where the ions of the salt have a higher
decomposition voltage. The water being present in excess, the hydrogen
and hydroxyl are re-formed at once and therefore are set free
continuously. If the current be so strong that new hydrogen and hydroxyl
ions cannot be formed in time, other substances are liberated; in a
solution of sulphuric acid a strong current will evolve sulphur dioxide,
the more readily as the concentration of the solution is increased.
Similar phenomena are seen in the case of a solution of hydrochloric
acid. When the solution is weak, hydrogen and oxygen are evolved; but,
as the concentration is increased, and the current raised, more and more
chlorine is liberated.
An interesting example of secondary action is shown by the common
technical process of electroplating with silver from a bath of
potassium silver cyanide. Here the ions are potassium and the group
Ag(CN)2.[1] Each potassium ion as it reaches the cathode precipitates
silver by reacting with the solution in accordance with the chemical
equation
K + KAg(CN)2 = 2KCN + Ag,
while the anion Ag(CN)2 dissolves an atom of silver from the anode,
and re-forms the complex cyanide KAg(CN)2 by combining with the 2KCN
produced in the reaction described in the equation. If the anode
consist of platinum, cyanogen gas is evolved thereat from the anion
Ag(CN)2, and the platinum becomes covered with the insoluble silver
cyanide, AgCN, which soon stops the current. The coating of silver
obtained by this process is coherent and homogeneous, while that
deposited from a solution of silver nitrate, as the result of the
primary action of the current, is crystalline and easily detached.
In the electrolysis of a concentrated solution of sodium acetate,
hydrogen is evolved at the cathode and a mixture of ethane and carbon
dioxide at the anode. According to H. Jahn,[2] the processes at the
anode can be represented by the equations