Electrolysis

In a galvanic cell a reaction makes electricity. In an electrolytic cell it is the other way round: we push electricity in to make a reaction happen that would not go by itself. This is electrolysis.

An outside battery is joined to two electrodes in one vessel. Here the cathode is the electrode joined to the battery's negative end, and the anode to its positive end. Reduction still happens at the cathode and oxidation at the anode. Positive ions travel to the cathode and negative ions to the anode.

Take molten sodium chloride: on its own it will never give up its sodium or chlorine, but pass a current and sodium metal forms at the cathode and chlorine gas at the anode.

With water present the products can change, because water itself can be reduced (to hydrogen) or oxidised (to oxygen). Whichever reaction is easiest wins, and the electrodes matter too. These products follow the book's examples and have to be learnt.

Faraday's laws tell us how much is produced.

• First law: the mass of substance deposited is proportional to the charge passed. The charge is Q = I × t (current in amperes × time in seconds).
• Second law: if the same charge passes through different electrolytes, the masses deposited are proportional to (molar mass ÷ electrons needed per ion).

One faraday (1 F) is the charge on one mole of electrons, 96487 C. Read the half-reaction to see how much charge is needed: Cu²⁺ + 2e⁻ → Cu needs 2 F per mole of copper; an ion that takes 6 electrons needs 6 F per mole.

Putting it together: mass deposited = (Q ÷ nF) × molar mass.

When cells are joined in series, the same current flows through each, so each receives the same charge. The masses then simply follow the second law: for copper and silver, they are in the ratio 63.5/2 : 108/1.