Electrolysis and Faraday's laws

In an electrolytic cell an outside voltage drives the reaction. Two copper strips in CuSO₄ solution: Cu²⁺ + 2e⁻ → Cu deposits on the cathode (negative), and Cu → Cu²⁺ + 2e⁻ dissolves the anode.

That is how copper is refined: impure copper is the anode, pure copper grows on the cathode.

Metals with no suitable chemical reducing agent are won by electrolysis: sodium and magnesium from their fused chlorides, aluminium from Al₂O₃ dissolved with cryolite.

Faraday's first law: how much substance reacts at an electrode is directly proportional to the charge that has flowed through the electrolyte, whether a solution or a melt.

Faraday's second law: the same quantity of electricity liberates different substances in proportion to their chemical equivalent weights (atomic mass of the metal ÷ electrons needed to reduce its cation).

Charge Q = I t (coulombs = amperes × seconds). One mole of electrons carries 1 F = 96487 C mol⁻¹, about 96500 C mol⁻¹ for rough work.

Ag⁺ + e⁻ → Ag needs 1 F per mole; Mg²⁺ + 2e⁻ → Mg needs 2 F; Al³⁺ + 3e⁻ → Al needs 3 F. Industrial cells run up to 50,000 A, about 0.518 F each second.

Worked: 1.5 A through CuSO₄ for 10 min passes 900 C; Cu needs 2F per mole, so mass = 63 × 900/(2 × 96487) = 0.2938 g of copper.