We thus have the H.s arriving at the kathode and the O.s at the anode; while the liquid between, being in the same condition as to average constitution as it was before we tried to pass a current, appears to be unaffected.

This directed travelling of the groups has received the name of the migration of the ions.

Experiment illustrating Grothüss's hypothesis.-The following experiment illustrates how the electro + and electro

groups (or kathions and anions)

B

FIG. i.

wick, and in A is the kathode, and in C is the anode. If two Leclanché cells be employed, and the whole be left for a day or so, it will be found that A has become alkaline, and C is acid; while B-through which have passed both the alkaline and acid groups. is still unaltered.

We will now represent in chemical symbols the interchanges that are continually taking place in an electrolyte when the current is passing, choosing a few typical cases. The upper brackets represent the condition of the chain of molecules before interchange; the lower brackets give the new grouping after interchange, with the ions set free at the two electrodes respectively.

A

Pt

I. The case of water.-Here we see that initially we have groups of OH, on'y, while finally we have an odd H, at the kathode, an odd O at the anode, with groups of OH, between, as before.

2

FIG. ii.

Pt

II. The Daniell's cell.-Here we will consider the case of a battery-cell itself, since all we have said of electrolysis applies to the battery-cell equally with the electrolytic-cell, to chemical combination equally with chemical decomposition. In the battery

cell the copper plate is the kathode, and the zinc plate is the anode.

Cu

The 'Zn' against the zinc plate represents an atom of the originally undissolved zinc. The arrangement of brackets represents that when the circuit is closed this Zn takes SO from the nearest H2SO4; this leaves H2 to take SO, from the next HSO4 and SO on; H, then passing through the porous cell takes SO, from the next CuSO1, and

FIG. iii.

so on, until finally Cu is deposited on the copper plate (the kathode).

As explained earlier, no change due to this interchange will be observable saving at the surfaces of the battery plates.

§ 5. Primary and Secondary Decompositions. We will now complete the theory of electrolysis, explaining the results of decomposition noticed in § 3, (ii.), (iii.), (v.), &c.

There is every reason for, and no reason against, the view that in electrolysis we have primarily the metallic, and the nonmetallic, groups set free against the kathode and anode respectively.

4

But for reasons of 'chemical affinity' such a condition of things may not be stable. Thus if Na,SO, be decomposed into Na, and SO, the Na, will decompose water and give 2 NaHO+H2, while the SO, with water will give H2SO, +0.

Thus we have by primary decomposition,

The reader will notice that, in both cases, what we have stated as to metallic groups moving with, and non-metallic groups moving against, the current, still holds good.

In the case of the decomposition of CuSO4, the Cu is de

posited at the kathode whether that be of copper or of platinum, Cu being stable. But if the anode be platinum, the SO, gives H2SO4+0; while if it be of copper it will act on this and give simply CuSO4

So in the case of other salts; the general principle of electrolysis determines the primary decomposition, and then a knowledge of chemistry enables us to predict the nature of any secondary action that may occur.

§ 6. Simultaneous Decompositions. When the strength of current per unit area of the electrode is great, we find that we get in many cases a simultaneous decomposition of the salt and of the water in which it is dissolved.

2

Thus, if we use very small electrodes of Cu in a solution of CuSO,, and drive a large current through, we get H, set free with Cu at the kathode, and O set free as well as Cu dissolved at the anode.

Such is a case of what is called simultaneous decomposition.

§ 7. Faraday's Laws of Electrolysis.-The laws of electrolysis were experimentally investigated, and enunciated, by Faraday. We have somewhat forestalled them in what has preceded; as it was impossible to discuss the action in cells and the nature of electrolysis in a purely historical order.

We have already (Chapter XI. § 8) stated that in the galvanometer we have a means of measuring current-strength, or quantity of electricity passing per second. Hence we have a means of investigating the relation between the amount of chemical action per second, and the strength of the current.

Faraday found the following two laws to hold.

I. The amount of chemical action per second is directly proportional to the current-strength.'

The physical meaning of this law would seem to be twofold; 1st, that the current does not pass through an electrolyte partly by electrolysis and partly by ordinary conduction, but only by electrolysis; 2nd, that an ion can only carry with it a fixed quantity of +, or of —, electricity, so that additional current requires a proportionally additional number of ions to carry it. The above law involves also the statement that a fixed current liberates per second a fixed mass of each ion respectively' (see § 9).

II. 'If the same or equal currents pass through several electrolytic cells, the weights of ions (which may be either atoms or molecular groups) set free at the several electrodes will be in the proportion of the chemical equivalents of these ions.'

An example will make this clear, our modern system of chemical notation rendering explanation an easy matter. Let the same current pass in succession through four cells containing copper sulphate, ammonium chloride, sodium sulphate solutions, and fused silver chloride, respectively. Then, taking secondary actions into account and considering the kathodes only, we get set free at the successive kathodes (i.) copper, (ii.) ammonia and hydrogen (which can be regarded as ammonium), (iii.) hydrogen, and (iv.) silver respectively. Now these are, when in chemical equivalents, represented by Cu, 2NH1, H2, and Ag, respectively; giving us as the numbers, to which the masses set free are proportional, (i.) 63 of copper, (ii.) 36 of ammonium, (iii) 2 of hydrogen, and (iv.) 216 of

silver.

The physical meaning of this law seems to be somewhat as follows. Each equivalent atom or group such as Cu, H, Ag2, SO1, 2(NO3), &c., carries with it the same unalterable quantity of electricity as it migrates to that electrode at which it is 'set free.' We have, in fact, a new significance attached to the expression 'chemical equivalent,' viz., that those groups or atoms which are chemical equivalents have also the same power of carrying electricity.

We may throw Law II. into the form

When the same or equal currents pass through several cells, there are the same number per second of equivalent atoms or groups set jree against each electrode.

$ 8. Further on Faraday's Laws of Electrolysis.

The case of simultaneous decompositions.--In this case (see § 6) the form of the law last given is the more useful.

Now if we have P grammes of copper and P grammes of silver, Р then the numbers of atoms in each are proportional to and 63 respectively. But, since copper is dyad and silver is monad,

Р

100

the numbers of equivalent atoms or molecules are proportional to

Р

Р

and respectively. If in successive cells we have set free at 63 216

the kathodes P grammes of copper, Q grammes of silver, and R grammes of hydrogen, respectively, then the second form of the Р Q R law gives us the result that =

63

=

216 2

Now if in one cell we get simultaneous decompositions, it is not hard to see what relation must hold. If, e.g., in the first cell we get both copper deposited, and water also decomposed setting free S grammes of hydrogen at the kathode, then we must have

But the proportion of S to P cannot be predicted; it depends upon the density of current and upon other conditions.

$ 9. Electro-chemical Equivalents. We shall in Chapter XIV. § 2 explain in what units we measure current-strength. We shall there find that the generally-employed unit of current is the ampère; this being a current in which a definite quantity of electricity called a coulomb passes across any section of the conducting circuit in one second of time.

The quantity, one coulomb, of electricity performs a definite amount of electrolysis in its passage through an electrolyte ; it liberates a definite mass of each ion. If we express this mass Z in grammes, the number is called the electro-chemical equivalent of that ion. We usually, however, give Z in milligrammes.

From Faraday's laws it is clear that if we know this number for H, we know it for all ions. Now, the chemical equivalent of silver has been of late experimentally determined with great care. And it seems certain that if we take 1118 milligrammes of silver to be set free by 1 ampère in 1 second, we shall be right to within part in 100 parts. In these experiments the strength of the current was of course measured independently of chemical action.

§ 10. Electro-plating. One case of electrolysis forms the basis of an important industry. It is the case where the electrolyte is a solution of some metallic salt, the anode is a plate of the same metal, and the kathode is a conducting surface that we desire to coat with this metal.