Electrolysis

The essentials

In an electrolytic cell, an external source of electricity drives a non-spontaneous reaction. Oxidation still happens at the anode and reduction at the cathode, but here the anode is connected to the positive terminal and the cathode to the negative terminal. Electrolysis is used to extract reactive metals, refine copper, and electroplate objects.

Key points

• Faraday's first law: the mass of substance deposited or liberated is proportional to the charge passed, w = Z × I × t, where Z is the electrochemical equivalent.
• Faraday's second law: for the same charge, masses of different substances deposited are proportional to their equivalent masses (atomic mass ÷ number of electrons).
• One faraday (F ≈ 96500 C) is the charge on one mole of electrons. Q = I × t.
• Ag⁺ + e⁻ → Ag needs 1 F per mole; Cu²⁺ + 2e⁻ → Cu needs 2 F; Al³⁺ + 3e⁻ → Al needs 3 F.
• Products depend on the electrolyte, electrodes and competing reactions. Molten NaCl gives Na at the cathode and Cl₂ at the anode. Aqueous NaCl gives H₂ at the cathode (water is reduced more easily than Na⁺) and Cl₂ at the anode (because of overpotential for O₂), leaving NaOH in solution.
• Dilute aqueous H₂SO₄ gives O₂ at the anode; concentrated H₂SO₄ can give S₂O₈²⁻ (peroxodisulphate).
• With an active copper anode in CuSO₄ solution, copper dissolves at the anode and deposits at the cathode.

Exam tips

• Example: 2 A for 965 s passes 1930 C = 0.02 mol of electrons; this deposits 0.01 mol Cu (0.635 g) or 0.02 mol Ag (2.16 g).

Common mistakes

• Assuming sodium metal forms when aqueous NaCl is electrolysed.
• Forgetting to divide moles of electrons by the charge on the ion.