of the anode copper relatively to the solution is raised, that of the cathode lowered (Fig. 35), where the dotted line shows the level of potential before electrolysis. On the whole, then, a fall of poten- Anode FIG. 35. Cathode will of course be reversed, so that it is always a force opposing the battery, and involving a waste of energy. It plays an especially important part in the action of the accumulator. Such changes of concentration occur in all electrolytic processes, so that the electromotive force is always modified by the process itself, if any appreciable amount of chemical action is produced, even if no polarisation of the electrodes in the ordinary sense takes place. Herein lies the cause of the influence that current density has on the course of an electrolytic process. We saw, in a general way (p. 26), that as in an electrolytic cell there is usually more than one way of effecting discharge, the ion which is most easily discharged will come down first; but that if the process is pushed, there will be a deficiency of such ions, and others, more tenacious of their charges, will be discharged. We can now, with the conception of voltage, give quantitative expression to this fact. electrode potential of any ion depends partly on its chemical nature, partly on the concentration at which it exists in the solution. If, then, there be two cations (or anion) present together, their electrode potentials will in general be different ; on electrolysis, that which has the higher potential (for an anion, the lower) will be discharged; but in this way its concentration will be reduced, and its potential lowered (for an anion, raised) till it becomes equal to that of the other. The two will then be discharged together. The To take an illustration of this—for convenience, one to which Nernst's logarithmic rule is fairly applicable—suppose a solution containing gram-equivalent of Pb(NO), and gram-equivalent of Ni(NO3)2 to be electrolysed (say with a lead cathode and platinum anode). The potential between the cathode and the lead salt alone would be o'129-0029 0*100 volt; with respect to the nickel salt alone, o'049 ・ 0'029 = 0'020 volt, showing a difference of o'080 volt; i.e. the lead has this much more tendency to deposit than the nickel. Now electrolyse this with a weak current (see below), such as is used for electrolytic analysis; lead alone will be deposited, until its concentration has fallen so much that it takes o'080 volt more to bring it down than was the case to start with. As a ratio of 10 changes the potential by o'029 volt, it will require a ratio of rather less than 1: 1000 to cause this change; i.. when the lead salt has fallen to 100 normal concentration, its potential (o'129-4 X o'029) will be about the same as that of the decinormal nickel. Beyond this point any current, however weak, will bring down the two metals together. It is, therefore, not possible to separate the remaining thousandth part of the lead from nickel electrolytically.1 1 But now suppose a stronger current to be used. Lead is at first deposited, in the way described, but of course the ions discharged are those next to the electrode, and the weakening of the solution takes place there: it is not the average concentration through the liquid on which the potential depends, but that of the layer actually in contact with the electrode, so that even while no considerable change occurs in the mass of the liquid, an important difference may happen in the contact layer, if the processes of diffusion, etc., are insufficient to make up for the local depletion of lead ions. It is, in fact, a question of balance between the consumption of ions by electrolysis and the supply by diffusion and convection of the liquid, so that if a strong current be used, unless corresponding precautions are taken to keep the liquid well stirred, the working layer in contact with the electrode may be so weakened in lead ions as to allow nickel ions to be deposited too. It is for this reason that the voltage required to electrolyse a given salt is This method is not used in practice, owing to secondary difficulties, such as oxidation of the lead deposit. not quite a fixed quantity; the numbers given in the tables give the least voltage required, or, in other words, the voltage required to perform the electrolysis very slowly, without introducing any disturbing factor. If, on the contrary, the process be pushed fast, by the application of a large current, concentration polarisation always occurs, and a greater voltage must be used. § 6. CHEMICAL POLARISATION The reversibility of any electrode is limited in practice by the concentration polarisation that occurs when any appreciable current is passed through it; but in very many cases a more important cause of irreversibility lies in the formation of new chemical materials by the current: this is polarisation in the usual sense of the word. Polarisation due to formation of a gas is a familiar and well-marked case of the phenomenon, and may serve as a typical instance. If current be passed through dilute sulphuric acid between platinum plates, oxygen is produced at the anode, hydrogen at the cathode; these substances appear as gas bubbles, clinging to the plates, and escaping into the air; but even before any visible bubbles have formed, a certain minute quantity of oxygen and hydrogen is liberated; this clings to the platinum, either in the form of a condensed layer on the surface, or actually in solution in the metal. (Much more markedly with a palladium electrode, for this metal, it is well known, dissolves hydrogen largely.) The platinum plates, neutral to start with, are thus transformed practically into electrodes of oxygen and hydrogen respectively, and come to have potentials corresponding to the position of those substances in the list of electro-affinities. That is, the oxygen is about II volt positive to the hydrogen. If now the decomposing current be stopped, and the cell examined with a voltmeter, it will be found that this difference of potential persists. We have, in fact, a voltaic combination O: H2SO: H capable of yielding a current, which will, of course, flow from the oxygen pole (+) to the hydrogen (-) through the outside circuit, ie. in the opposite direction to the original or charging current. The arrangement is, in fact, an accumulator," which can be charged and subsequently dischargedbut with this difference from the lead accumulator of practice, that its capacity is excessively small. The capacity, in fact, is only that due to the traces of gas absorbed by the plates. If the cell be made to yield a current, it does so by hydrogen going into solution at the negative, oxygen at the positive. Thus the "polarisation" by the gases is rapidly used up in production of current, and disappears. So far the electrode has behaved in a reversible manner; but as soon as the supply of gas has disappeared, no more of this reverse current will flow, unless it be driven by some external power, and then it will be carried, not by the same ions as before, but by a different polarisation, viz. by discharge of oxygen at the plate which was formerly coated with hydrogen, and vice versa. Thus the plate which was formerly at a high potential is now at a low; whereas in a reversible electrode (e.g. zinc in ZnSO, solution) the potential is the same whether it be used as anode or cathode (except for concentration polarisation). In other words, the polarisation always acts in opposition to the working electromotive forces of the circuit. The credit is due to Leblanc1 for first disentangling the rather complex phenomena of polarisation, and showing that they depend on the same principles of potential difference as those of voltaic cells. He showed that to electrolyse a solution it is simply necessary to apply a difference of potential equal to that between the substance to be liberated at the two poles (taking account of concentration, of course). Thus to decompose copper sulphate (normal-ionic) we require, according to the table, p. 159, 1'9 06 13 volts; to decompose zinc sulphate, 1'9 − (−0′5) = 2'4 volts. Again, the decomposition potentials for hydrochloric acid were given on p. 166, and their dependence on both electro-affinity and concentration is clear; and further, if the dilution be made greater than that given in the table, the potential difference will rise above 1'7 1 Zeitschr. phys. Chem., 8. 299-330 (1891); 12. 333-358 (1893). = = volts. But this is the amount ordinarily required to decompose water (vide infra, p. 177); hence, in accordance with the general principles stated, at this point oxygen will be liberated as well as chlorine. We have, indeed, assumed Leblanc's conclusions in the preceding sections, and made no distinction between the treatment of voltaic and electrolytic cells. But this account of the phenomena would be very incomplete without mention of several minor points that tend to obscure the simplicity of the fundamental relations. If an electrolytic cell containing dilute sulphuric acid between platinum plates be connected to an ammeter and voltmeter, and successively greater voltages be applied to it, the facts observed are as follows: When the voltage is small-say half a volt—at the first moment of turning it on a considerable deflexion of the ammeter is produced, which immediately drops to almost nothing. Before starting, the two electrodes, being quite similar, offer no back electromotive force, and consequently current flows in accordance with Ohm's law. This almost immediately produces enough oxygen and hydrogen on the plates to set up a back electromotive force of half a volt, and so stop the current. The platinums are, in fact, imperfectly converted into oxygen and hydrogen electrodes, and are said to be polarised to the extent of half a volt. The current does not altogether stop, however; oxygen and hydrogen, being both soluble in water, slowly pass off from the plates into the electrolyte; there they may either escape into the air, or, meeting, recombine. The polarisation of the plates, therefore, slowly disappears of its own accord—or would do so, but that it is kept up by renewed action of the applied electromotive force. The small steady current indicated by the ammeter is that required to keep up the polarisation that is being dissipated, and is usually called the "residual current." It is very small when oxygen and hydrogen are the products of electrolysis, for their solubility in water is very small; but when strong hydrochloric acid is electrolysed, chlorine, being freely soluble, diffuses over to the hydrogen on the cathode and recombines with it to a considerable |