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LearnCambridge97012025–2027Paper 4 · A Level Structured QuestionsElectrochemistry

Paper 4 · Module

Electrochemistry

2 subtopics≈ 95 min in total

Start the module: Electrode potentials and cells Open graphing calculator

Electrode potentials and cells
A half-cell's standard electrode potential E⦵ is measured against the standard hydrogen electrode (SHE, 0.00 V) at 298 K, 1 mol dm⁻³ and 100 kPa. Half-equations are written as reductions; the more positive E⦵, the stronger the oxidising agent (left side). $E^\circ_\text{cell}=E^\circ_\text{(positive, reduction)}-E^\circ_\text{(negative, oxidation)}$ . A reaction is feasible if $E^\circ_\text{cell}>0$ (but may be slow, or conditions may be non-standard). Changing concentration shifts E (Nernst: $E=E^\circ+\frac{0.059}{z}\log\frac{[\text{ox}]}{[\text{red}]}$ ). $\Delta G^\circ=-nFE^\circ_\text{cell}$ . Fuel cells and batteries apply this.
≈ 55 min
Electrolysis calculations
Charge $Q=It$ (coulombs = amps × seconds). The Faraday constant $F=96\,500$ C mol⁻¹ is the charge on 1 mol of electrons, and $F=Le$ (Avogadro constant × electron charge). Moles of electrons $=Q/F$ ; then use the half-equation: $\mathrm{Cu^{2+}+2e^-\rightarrow Cu}$ needs 2 mol e⁻ per mol Cu. Products in aqueous electrolysis are predicted from E⦵: at the cathode, the ion with the more positive E⦵ is reduced (Cu²⁺, Ag⁺ before H₂O); at the anode, concentrated halide gives halogen, otherwise O₂ forms. The Avogadro constant can be found by electrolysis: $L=F/e$ .
≈ 40 min