Electrolysis rules become electrode potentials

In electrolysis of aqueous solutions, hydrogen is produced at the cathode unless the metal is less reactive than hydrogen; at the anode, a halogen is produced if a halide is present, otherwise oxygen. Learned as rules.

Cells are classified: voltaic (spontaneous, chemical → electrical) vs electrolytic (non-spontaneous, electrical → chemical), with oxidation always at the anode and reduction always at the cathode in both — but with opposite electrode polarities, which is the source of endless confusion (RSC explicitly recommends using "positive/negative electrode" language to avoid it). At HL, the standard hydrogen electrode defines zero (3.2.12), tabulated E⦵ values let you calculate E⦵cell = E⦵cathode − E⦵anode, and positive E⦵cell ⇔ spontaneous (3.2.13), linked to thermodynamics by ΔG⦵ = −nFE⦵cell (3.2.14). The GCSE electrolysis rules are then derived from competing electrode reactions including the oxidation and reduction of water (3.2.15).

From rules to a quantitative predictive framework tied to Gibbs energy. Concentration and overpotential effects mean the prediction isn't always what the E⦵ table says — IB flags "competing reactions".

No thermodynamics available at GCSE.

Anode/cathode confusion across cell types (documented as a major difficulty); assuming the electrode signs are the same in both cell types; subtracting E⦵ values in the wrong order; forgetting that E⦵ values are not multiplied by stoichiometric coefficients (they are intensive) — this is a very common error.