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ELECTROLYSIS (hl-ek-troll-sis). The passage of an elec tric current through a substance, accompanied by definite chemical changes which are independent of the heating effects of the current, is known as electrolysis (formed from Gr. Xbetv, to loosen), and the substance is called an electrolyte. An example of this is a solution of a salt such as copper sulphate in water, through which an electric current is passed between copper plates, when the following phenomena are observed : 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 at tached 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 this case 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 solu tion. If instead of copper electrodes plates of platinum are em ployed, copper is still deposited on the cathode, but instead of the anode dissolving, free sulphuric acid appears in the neigh bouring solution, and oxygen gas is evolved at the surface of the platinum plate. With other electrolytes similar phenomena ap pear, 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.

The distinction between electrons associated with gases and with liquids is that, in gases the electrons sometimes travel alone, but in liquids they are always attached to matter and their motion involves the movement of chemical atoms or groups of atoms. An atom with an extra corpuscle is a univalent negative ion, an atom with one corpuscle detached is a univalent positive ion. In metals the electrons can slip from one atom to the next, since a current can pass without chemical action. When a current passes from an electrolyte to a metal, the electron must be detached from the atom it was accompanying and chemical action be mani fested at the electrode.

Alessandro Volta of Pavia discovered the electric battery in the year 1800, and thus placed the means of maintaining a steady tric 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 tially 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 cess consisted essentially in the decomposition of water. These observers also noticed that a similar kind of chemical action went on in the battery itself. Soon afterwards William shank decomposed the magnesium, sodium and ammonium chlorides, and precipitated silver and copper from their solutions, an achievement 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. zelius stated that neutral salt solutions could be decomposed by electricity, the acid appearing at one pole and the metal at the other. 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 decomposi tion of water could be effected even if 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 mois tened solids, and thus isolated the metals potassium and sodium. Faraday's Laws.—The first exact quantitative study of elec trolytic phenomena was made about 1830 by Michael Faraday (Experimental Researches, 1833), who examined the relation be tween the flow of electricity round the circuit and the amount of chemical decomposition. He found that, for the same current, the amount of chemical action was independent of the size of the elec trodes and proportional to the time that the current flowed. The results of all these experiments may be summed up in the state ment 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 differ ent electrolytes in series. He found that the amounts of the sub stances 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 I gm. of hydrogen and 31.8 gm. of copper unite chemically with the same weight of any acid radical such as chlorine or the sulphuric group, 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 Fara day's laws :—The mass of substance liberated from an electrolyte by the passage of a current is proportional (I) to the total quan tity 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 electro magnetic units. The mean result of the best determinations shows that when a current of I ampere is passed for I second, a mass of silver is deposited equal to 0.001 118 gm. So accurate and conven ient is this determination that it is now used conversely as a prac tical definition of the ampere, which (though defined theoretically in terms of magnetic force) is defined practically as the current which in one second deposits 1.118 mg. of silver. If, as is now usual, we take the equivalent weight of oxygen as our standard and call it 16, the equivalent weight of silver is 107.88, that of hy drogen being 1.008, and its electrochemical equivalent 1.036X The electrochemical equivalent of any other substance may be found by multiplying its chemical equivalent by 1.036g X If, instead of the ampere, we take the c.g.s. electromagnetic unit of current, this number becomes 1.036X It was pointed out by Helmholtz in 1881 that Faraday's law forms one of the most important reasons for assuming that elec tricity has an atomistic structure since it must be assumed that every valence bond in an atom or radical involves a definite quantity of electricity, just as in the formation of a chemical compound several atoms or radicals are able to unite with one another. Further according to Helmholtz if the assumption is made that simple substances are composed of atoms, the conclu sion cannot be avoided that electricity, positive as well as nega tive, decomposes into definite elementary particles which behave as electrical atoms. If Avogadro's number N (the number of molecules in a gram-molecule) is known, the size of the ele mentary charge of electricity is given in accordance with Fara onds) = 9'650 absolute electromagnetic units (e.m.u) = 28 95 X Iois N N absolute electrostatic units (e.s.u.).

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 deter mined 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 electro motive 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 volt age 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 second ary 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.

Early Theories of Electrolysis.

The obvious phenomena to be explained by any theory of electrolysis are the liberation of the products of chemical decomposition at the two electrodes while the intervening liquid is unaltered. To explain these facts T. Grotthus (1785-1822) in 1806 put forward an hypothesis which supposed that the opposite chemical constituents of an electrolyte inter changed partners all along the line between the electrodes when a current passed. Thus, if the molecule of a substance in solution is represented by AB, Grotthus considered a chain of AB molecules to exist from one electrode to the other. Under the influence of an applied electric force, he imagined that the B part of the first molecule was liberated at the anode, and that the A part thus isolated united with the B part of the second molecule, which, in its turn, passed on its A to the B of the third molecule. In this manner, the B part of the last molecule of the chain was seized by the A of the last molecule but one, and the A part of the last molecule liberated at the surface of the cathode. Chemical phe nomena throw further light on this question. If two solutions containing the salts AB and CD be mixed, double decomposition is found to occur, the salts AD and CB being formed till a certain part of the first pair of substances is transformed into an equiva lent amount of the second pair. A freedom of interchange is thus indicated between the opposite parts of the molecules of salts in solution, and it follows reasonably that with the solution of a single salt, say sodium chloride, continual interchanges go on between the sodium and chlorine parts of the different molecules.

These views were applied to the theory of electrolysis by R. J. E. Clausius. He pointed out that it followed that the electric forces did not cause the interchanges between the opposite parts of the dissolved molecules but only controlled their direction. Interchanges must be supposed to go on whether a current passes or not, the function of the electric forces in electrolysis being merely to determine in what direction the parts of the molecules shall work their way through the liquid and to effect actual sepa ration of these parts (or their secondary products) at the elec trodes. This conclusion is supported also by the evidence supplied by the phenomena of electrolytic conduction. (See ELECTRICITY, CONDUCTION OF: In Liquids.) If we eliminate the reverse elec tromotive forces of polarization at the two electrodes, the con duction of electricity through electrolytes is found to conform to Ohm's law ; that is, once the polarization is overcome, the current is proportional to the electromotive force applied to the bulk of the liquid. Hence there can be no reverse forces of polarization inside the liquid itself, such forces being confined to the surface of the electrodes. No work is done in separating the parts of the molecules from each other. This result again indicates that the parts of the molecules are effectively separate from each other, the function of the electric forces being merely directive.

Migration of the

Ions.--The opposite parts of an electrolyte which work their way through the liquid under the action of the electric forces were named by Faraday the ions—the travellers. The changes of concentration which occur in the solution near the two electrodes were referred by W. Hittorf (1853) to the unequal speeds with which he supposed the two opposite ions to travel. It is clear that, when two opposite streams of ions move past each other, equivalent quantities are liberated at the two ends of the system. If the ions move at equal rates, the salt which is decomposed to supply the ions liberated must be taken equally from the neighbourhood of the two electrodes. But if one ion, say the anion, travels faster through the liquid than the other, the end of the solution from which it comes will be more exhausted of salt than the end towards which it goes. If we assume that no other cause is at work, it is easy to prove that, with non-dissolvable electrodes, the ratio of salt lost at the anode to the salt lost at the cathode must be equal to the ratio of the velocity of the cation to the velocity of the anion. This result may be illustrated by fig. 1.

The black circles represent one ion and the white circles the other. If the black ions move twice as fast as the white ones, their distribution after the passage of a current will be represented by the lower part of the figure. Here the middle part of the solution is unaltered and the number of ions liberated is the same at either end, but the amount of salt left at one end is less than that at the other. On the right, towards which the faster ion travels, five molecules of salt are left, being a loss of two from the original seven. On the left, towards which the slower ion moves, only three molecules remain—a loss of four. Thus, the ratio of the losses at the two ends is two to one—the same as the ratio of the assumed ionic velocities.

Transport Numbers.—It should be noted, however, that another cause would be competent to explain the unequal dilution of the two solutions. If either ion carried with it some of the unaltered salt or some of the solvent, concentration or dilution of the liquid would be produced where the ion was liberated. There is reason to believe that in certain cases such complex ions do exist, and interfere with the results of the differing ionic velocities. For certain concentrated solutions the transport number is found to be greater than unity ; thus for a normal solution of cadmium iodide its value is 1 • I 2. (One litre of a normal solution contains one gram equivalent of the dissolved substance ; see EQUIVALENT.) This is best explained by the for mation of structures represented by some such chemical formula as I . It is found that, in such cases as this, where it seems necessary to imagine the exist ence of complex ions, the trans port number changes rapidly as the concentration of the original solution is changed. Thus, dimin ishing the concentration of the cadmium iodide solution from normal to one-twentieth normal changes the transport number from 1.12 to 0•64. Hence it is probable that in cases where the transport number keeps constant with changing concentration the hypothesis of complex ions is un necessary, and we may suppose that the transport number is a true migration constant from which the relative velocities of the two ions may be calculated in the man ner suggested by Hittorf. This conclusion is confirmed by the re sults of the direct visual determination of ionic velocities (set ELECTRICITY, CONDUCTION OF : In Liquids), which, in cases where the transport number remains constant, agree with the values cal culated from those numbers. Many solutions in which the trans port numbers vary at high concentration often become simple at greater dilution.

F. Kohlrausch in 1879 provided the next important step in the theory of the subject. Kohlrausch formulated a theory of electro lytic conduction based on the idea that, under the action of the electric forces, the oppositely charged ions moved in opposite di rections through the liquid, carrying their charges with them. If we eliminate the polarization at the electrodes, it can be shown that an electrolyte possesses a definite electric resistance and therefore a definite conductivity. On the view of the process of conduction described above, the amount of electricity conveyed per second is measured by the product of the number of ions, (known from the concentration of the solution), the charge car ried by each of them and the velocity with which on the average they move through the liquid. The concentration is known, and the conductivity can be measured experimentally; thus the average velocity with which the ions move past each other under the existent electromotive force can be estimated, and is found to be equal to the sum of their individual velocities, which can there fore be calculated. From Hittorf's transport number, in the case of simple salts in moderately dilute solution, we have the ratio between the two ionic velocities. Hence the absolute velocities of the two ions can be determined, and we can calculate the actual speed with which a certain ion moves through a given liquid under the action of a given potential gradient or electromotive force.

current, solution, chemical, ions, electrodes, copper and electric