Electrochemistry and Redox Reactions
What this domain tests
Electrochemistry items test redox bookkeeping, cell construction and potential calculations: assigning oxidation numbers, writing half-equations, computing E°cell and ΔG°, applying the Nernst equation, and using Faraday's laws in electrolysis.
Core rules
- Oxidation is loss of electrons at the anode; reduction is gain of electrons at the cathode. In a galvanic cell the anode is negative and the cathode positive; in an electrolytic cell the polarity is reversed but the electrode chemistry is not.
- E°cell = E°cathode − E°anode, using standard reduction potentials for both. E° is intensive: doubling a half-equation does not double its potential.
- ΔG° = −nFE°, with F = 96 485 C mol⁻¹. A positive E°cell means a spontaneous reaction.
- Nernst equation at 298 K: E = E° − (0.0592/n) log Q. As the cell discharges, Q rises and E falls to zero at equilibrium.
- Electrolysis: moles of electrons = It/F; the mass deposited follows from the half-equation's electron count.
Worked example
For Zn(s) + Cu²⁺(aq) → Zn²⁺(aq) + Cu(s), with E°(Cu²⁺/Cu) = +0.34 V and E°(Zn²⁺/Zn) = −0.76 V: E°cell = 0.34 − (−0.76) = 1.10 V. Two electrons are transferred, so ΔG° = −2 × 96 485 × 1.10 = −212 kJ mol⁻¹. Cell notation: Zn(s) | Zn²⁺(aq) ‖ Cu²⁺(aq) | Cu(s), with the anode on the left.
Traps that cost marks
- Reversing the sign of the anode potential instead of subtracting it.
- Multiplying E° by the stoichiometric coefficient when scaling a half-equation.
- Writing the cathode on the left in cell notation.
- Forgetting that a species with a more positive reduction potential is the better oxidising agent, not the better reducing agent.