Chapter summary
Electrochemistry, at a glance
Measure and combine electrode potentials, predict cell reactions, and calculate what electrolysis produces.
A-Level 9476 (2026-2027)
Quick revision
Revisit the essentials, then return to an explanation when you need it.
| Term | Definition to write |
|---|---|
| Standard electrode potential, E° | The potential difference between a half-cell and the standard hydrogen electrode, measured under standard conditions: 298 K, 100 kPa for gases and 1 mol dm-3 for ions in solution. |
| Oxidation and reduction | Oxidation is the loss of electrons, or an increase in oxidation number. Reduction is the gain of electrons, or a decrease in oxidation number. |
| Task | Use | Essential check |
|---|---|---|
| Predict a standard cell | E°cell = E°cathode - E°anode. | Both values are reduction potentials; positive means favourable for the written reaction under standard conditions. |
| Relate voltage to energy | ΔG° = -nFE°cell. | n belongs to the balanced reaction; volts with coulombs give joules. |
| Combine half-reactions | Add their Gibbs-energy contributions, then convert back to E°. | Potentials are not directly additive or multiplied by stoichiometric coefficients. |
| Calculate electrolysis yield | Q = It; n(e-) = Q/F; use the half-equation. | Time in seconds, correct electron ratio, current efficiency and gas conditions. |
| Explain an industrial cell | Write both electrode reactions first. | Purification: impure Cu anode/pure Cu cathode. Anodising: aluminium object is the anode. |
Always: oxidation at the anode, reduction at the cathode. Galvanic discharge: anode negative, cathode positive. Electrolysis: anode positive, cathode negative. Standard-potential tables predict thermodynamic direction; composition and kinetic barriers explain why a real observation can differ.
Quick recall
Define standard electrode potential.