It is the behaviour of the second class of bodies under the influence of the electric current that we have now to discuss. The physical side of the subject has already been considered in the article Electricity; so we shall principally confine our attention to the phenomena of electrolysis which bear on the laws and principles of chemistry. Before going further it will be necessary to introduce the technical terms which have now become familiar, and, in order to be definite, we will consider somewhat closely a particular instance of electrochemical decomposition of the simplest type.

The cell in which the action takes place consists of a wide tube of hard glass, bent into a V-shape; into this is introduced some silver chloride, which is kept fused during the experiment; into the liquid in one leg of the tube is dipped a platinum wire connected with the negative pole (zinc) of a battery of 3 or 4 Grove’s cells, and into that in the other a piece of graphite or gas carbon connected with the positive pole of the same battery. We will suppose a galvanometer introduced into the circuit, and that the current strength as indicated thereby is, roughly speaking, constant, so that the quantity of electricity which passes can be measured roughly by the time occupied in passing. After the circuit has been closed a short time, bubbles of chlorine will begin to come off from the carbon, while pure silver is deposited upon the platinum wire, but except at these points no alteration will take place at any part of the fluid. If the platinum wire with the attached silver be weighed at intervals, it will be found that the amount deposited after the current has become constant is proportional to the time, i.e., to the amount of electricity which has passed through the liquid.. The same will be true of the chlorine if collected in the other leg of the tube, due allowance being made for the small bubbles retained by the carbon, &c. And the amount of chlorine will be chemically equivalent to the amount of silver; thus for every 108 grammes of silver on the platinum there will be 35 5 grammes of chlorine set free in the other leg of the tube. Moreover if the current be varied by varying the number of battery cells, it will be found that the amount of decomposition in a given time is proportional to the current, that is, again, to the quantity of electricity which traverses the substance.

Faraday, who was the first to define the laws which hold in electrochemical decomposition, introduced, for the sake of precision, a system of nomenclature which has since been generally employed. Wishing to regard the terminals corresponding, in any similar case, to the carbon and platinum in the above experiment merely as the “doors” by which the electricity enters and leaves the liquid, he denominated them electrodes, and, comparing the “path” of the current to those of the currents which may produce terrestrial magnetism, and hence to the course of the sun, he called the homologue of the carbon (where the current, so to speak, “rose,” or entered) the anode, that of the platinum (where the current “set,” or left) the cathode. The component parts, no matter how complex, into which the liquid was decomposed, corresponding to the Ag and Cl of the above, received the name of “ions ”—that component which went down with the current to the cathode, and there either was set free or combined with the cathode or the surrounding liquid, beiug the cation, and that which went up against the current, and appeared or promoted some chemical action at the anode, the anion. Moreover, the substance decomposed was called an electrolyte, and the process itself electrolysis. (Faraday, Exp. Res., 662 sqq.)

The phenomena which occur at the electrodes when the ions there set free react upon the electrode or the surrounding fluid, so that the resulting products of electrolysis are not the ions themselves, are called secondary actions.

The anion and the cation are frequently called the negative and positive ion respectively. Similarly the cathode and anode are termed the negative and positive electrodes; Daniell denoted them the platinode and the zincode, but these terms have fallen into disuse.

Of the bodies which are capable of electrolytic conduction nearly all,if not all, are liquids. Faraday {Exp.Res., 433, 1340) apparently obtained some chemical decomposition iu sulphuret of silver and a few other salts when solid, but this did not alter his opinion that the mobility secured in the fluid state, either by fusion or by solution, was necessary to the phenomena of electrolysis; and his view, which he supported by experiments on ice and other solids that conduct when fused {Exp. lies., 380-397, 419-428), still obtains. Electrolytic action doubtless sometimes takes place in gases, but accurate investigation of the subject is difficult on account of the extreme mobility of the particles and the danger of confusing electrolytic effects with effects due to disruptive discharge by convection. Gases have, however, been decomposed by the silent discharge, as CO 2 into CO + O.

From Faraday’s time attempts have continually been made to classify strictly, according to their chemical composition or constitution, the liquids capable of electrolytic conduction, but hitherto without very much success. It must be remembered that, as the resistance of a liquid increases, the tests of electrolytic conduction become less and less sensitive. We cau consider a body an electrolyte if we can (1) collect the products of decomposition, or (2) demonstrate their presence on the electrodes by means of the return current due to polarization. If the resistance be very great the former method becomes evidently very difficult, and in the latter complications are introduced which cannot here be discussed (see Electricity). On the other hand, we might easily be misled into considering a body an electrolyte from the presence of mere traces of a foreign substance. Thus at one time water was regarded as the only electrolyte, but it is found that the purer the water is the less does it conduct electricity, and now Kohlrausch and Nippoldt have shown that the presence of one 10-millionth of II 2 SO 4 would be sufficient to account for its observed conducting power, so that the weight of evidence goes to show that water itself is not an electrolyte at alL

It is not, then, surprising that views on the question of what constitutes an electrolyte have changed considerably. Davy and the older chemists, as mentioned above, considered water to be the only electrolyte; Faraday, by electrolysing fused chlorides, &c., dissipated these notions, but still regarded water as the electrolyte which was decomposed when acids were subjected to the electric current, and his general conclusion was that an electrolyte must be a compound consisting of an equal number of chemical equivalents of its elements, that is, in modern notation, must be of the type M∞R ’ ^ where λ∙ and y are the atomicities or valencies of the elements whose atomic ινeights are represented by M and R, and thus that two elements would by uniting form only one electrolyte (Exp. lies., C79-70l, 830). The oxygen salts for which Faraday assigned no law were included by Daniell in the same formula as binary compounds, of which the part R acting as anion was no longer an element but a compound; thus ZnSO 4 was shown to be split up by electrolysis into Zn and SO 4 ; in that case y would represent the basicity of the acid forming the salt.

This hypothesis lacks definiteness, on account of the variation of the atomicity of the elements, and falls through altogether in the case of copper and iron, which form each two chlorides, (CuCl 2 ,Cu 2 Cl 2 ), (FeCl 2 ,Fe 2 Cl c ), both electrolytes, and in consequence Wiedemann (Galv., Bd. i. §§ 295, 34Ga, 418 (5)) modifies the statement of the hypothesis, and considers that for a body to be an electrolyte it must be capable of formation by double decomposition from one of the simple binary electrolytes, the exchanging atoms or groups of atoms forming the ions of the new compound. Thus silver acetate gives, by double decomposition with sodium chloride, silver chloride and sodium acetate. Sodium acetate and silver chloride are therefore electrolytes of which Ag, Cl, Na, C 2 H 3 O 2 are the respective ions. This hypothesis may be illustrated by a great number of instances:— the case of the decomposition of uranium compounds, as UOC1 into UO and Cl, is a very good example. But Wiedemann’s view would indicate that a body, in order to be an electrolyte, need but be one of a “series of salts,” and we then see no reason for excluding the hydrogen salts from the class; thus H.,0 and HC1 can be easily formed by double decomposition, yet the former is, when pure, one of the worst liquid conductors, while the latter as liquefied gas is apparently not decomposed even by 5640 cells of De la Rue’s chloride of silver battery, but gives vibrations indicating very high resistance. Bleekrode has also shown that, of all the pure liquefied hydrogen acids, only HCN is an electrolyte. On the other hand, liquefied NH 3 , which is not formed, so far as we are aware, by double decomposition, is electrolysable by only a moderate battery of Bunsen's cells, giving a blue liquid at the cathode. Moreover, Buff (Ann. d. Chem. und Pharm,, Bd. cx.) has electrolysed molybdic and vanadic anhydrides after the manner MυO 3 = MυO 2 + O, but these bodies are not obtainable by double decomposition with a simple electrolyte.

Miller (Elements Chem., i. § 282 (v) ) considers that an electrolyte must be a combination of a conductor and a nonconductor, and so the majority of electrolytes are. But alloys behave to a certain extent as electrolytes when fused (see Wied., Galv., Bd. i. § 328), and SnCl 4 , though consisting of a conductor and a non-conductor, is not an electrolyte; so that this classification is not exclusive.

It would therefore appear that the condition does not lie in the chemical constitution of the body, but rather in its molecular state, and to this points the fact that two nonconductors, as 1I 2 O and HC1, on being mixed form a very good conductor. In addition to this, quantitative measurements of the resistance of electrolytes show that, iu the case of many salt and acid solutions, there is a point of concentration below saturation, for which the conductivity is a maximum. This would scarcely be the case if one alone of the bodies were the conductor.

The liquids which do not conduct are very various, including, besides oils and resins and other organic bodies, benzine, iodide of sulphur, carbon disulphide, glacial acetic acid, fused boracic anhydride, antimonic oxide and oxychloride, the higher halogen salts of tin, liquid sulphurous anhydride, pure water, and pure halogen acids. For others see article Electricity, p. 51.

In the description of the phenomena, in the typical case of electrolysis given above, it was stated that the amount of chemical decomposition in any time is proportional to the whole quantity of electricity which passes through the liquid in that time; this is true in all cases of electrolysis, and was established by Faraday (Exp. lies., v. 505, and ser. vii.). It forms part of the general law to which his name is attached, but we prefer to consider it separately for reasons that will appear when we discuss th i statement of that law. We may put it thus:—If W be he mass of an electrolyte, decomposed by the passage of a qb αntity E of electricity, then, as long as the ions remain of tl e same nature,

W = KE . ... ... (1), where K is a constant dependent only on the nature of the electrolyte, and therefore independent of the nature or size of the electrodes and of any secondary actions which may take place.

It is evident that if we cau prove the truth of this law for one electrolyte, with ions which do not vary with variations of electromotive force, we shall have a very convenient means of measuring the total amount of electricity which passes through any circuit in a given time by introducing such an electrolyte into the circuit, and measuring the amount of defcomposition in the given time. Faraday demonstrated the truth of the law in the case of dilute sulphuric acid by experiments with vessels in which the products of decomposition of the dilute acid between platinum electrodes could be conveniently collected, either separately or together, and measured [Exp. Res., 714— 728.) Such an instrument he called a volta-electrometer, and subsequently a voltameter. After demonstrating that the amount of decomposition was independent of the size of the electrodes, he connected up two voltameters A and B, in multiple arc, as in the accompanying diagram, and then passed the whole current through a third C, and found that the amount of decomposition in C was equal to the sum of the amounts in A and B. He therefore applied the voltameter to measure quantities of electricity in other cases.

Various forms of voltameter have been employed (see Wiedemann, Calvaιιismus, Bd. i. § 317-319). The most accurate is the silver voltameter of Poggendorff, which consists of a vertical rod of silver with the lower end immersed in a solution of silver nitrate contained in a platinum vessel; the silver is connected with the positive, the platinum vessel with the negative pole of the battery, and the amount of decomposition is ascertained by weighing the platinum vessel with the attached silver before and after the experiment. Buff directly proved the truth of equation (1) for such an instrument by electrolysing silver nitrate solutions of different strengths between silver electrodes. The currents employed were varied for different experiments, and were measured by a tangent galvanometer, and the quantity E of electricity was deduced by observing the time of passage of the current. {Ann. d. Chem. u. Pharm., xciv. 15.)

We have, then, in order to demonstrate generally the law expressed in equation (1), to measure the amount of the ions set free in any case of electrolysis, while the amount of electricity is measured at the same time by means of a voltameter included in the circuit. But the measurement of the amount of ions liberated is not always an easy task; in the great majority of cases secondary actions (see above, p. 106) occur, the primary results of electrolysis are obscured, and in order to determine the nature and amount of the ions special apparatus and further investigation are necessary.

Since the ions are liberated at the electrodes the products of secondary action will remain in the immediate neighbourhood if the action be not too long continued. We may therefore determine the ions by collecting any gaseous products, ascertaining the loss or gain in weight of the electrodes, and analysing the electrolyte in the immediate neighbourhood of the electrodes, taking care that the products at the two do not mix by gravitation, by diffusion, or otherwise.

For instance, if a fused chloride {e.g., PbCl 2 ) be electrolysed with platinum electrodes, no chlorine will be evolved at the anode, although Pb will be deposited at the cathode; but if the liquid round the anode be analysed, for every 414 grammes of lead at the cathode will be found 339 grammes of PtCl 4 round the anode. Now the platinum must have been derived from the anode, which will be found to have lost 197 grammes in weight, consequently the 142 grammes of Cl were derived by the electrolysis from the PbCl 2 , and hence PbCl 2 is electrolysed as Pb + Cl 2 .

In order to separate the fluids at the two electrodes, various forms of apparatus have been employed. For fused electrolytes a W-shaped tube, which can be divided after the fluid has solidified, is sufficient; with solutions, however, where the solvent introduces new complications the separation is more difficult, owing to the “migration of the ions” and other causes which will be considered below. Daniell and Miller {Phil. Trans., 1844) used a cylindrical glass vessel separated into three compartments by porous clay diaphragms, the two end compartments containing the electrodes, and having tubes for conducting away gaseous products; while Hittorf, in a classical series of experiments {Pogg. Ann., lxxxix. xcviii. ciii. cvi.), used a number of bell-shaped glass vessels fitted to each other with indiarubber washers, the electrodes being inserted in the bottom and top vessels respectively. The lower end of each bell was covered with membrane to prevent mixing of the products; the whole apparatus was filled with the electrolyte to be decomposed; and the products at the two electrodes were known to be separated if the composition of the fluid in one of the intermediate bells remained the same throughout the experiment.

Great numbers of experiments have been made by different experimenters in one or other of the ways mentioned, and they have thus proved that, whatever the electrodes, and whatever the electromotive force, the secondary action at the electrodes has no effect upon the amount of chemical decomposition, and therefore the law of equation (1) always holds.

We can give here but a few examples of secondary action. A very good account will be found in Wiedemann, Bd. i. § 326-385, with, however, the drawback of the use of an obsolete chemical notation.

(1.) The ions themselves are set free, but separate into component parts. That this is the case with oxygen salts, which are separated into the metal and a complex anion which is resolved into oxygen and an anhydride, was pointed out by Daniell CΡhil. Trans., 1839), who gave to the SO 4 , derived as electro-negative ion from sulphates, the name of oxysulphion, and so on. Many similar cases occur in electrolyses of organic compounds. Thus potassium acetate is electrolysed originally as KC 2 H 3 O 2 = K + C 2 H 3 O 2 ; hut the anion splits up (partly at least) thus: 2C 2 H 3 O 2 ≈ C 2 H 6 + 2CO 2 . All the potassium salts of the fatty acids behave similarly, so that this becomes a general method of preparing the normal paraffins.

(2.) The ions appear in an abnormal molecular state. The deposit of copper in Gladstone and Tribe’s ZnCu couple is a black crystalline powder (see p. 114). The...