Electrochemistry and Redox Reactions
Electrochemistry questions connect redox bookkeeping to cell potential and to the free-energy change that drives a galvanic cell or an electrolysis. Expect oxidation numbers, cell notation, E°_cell, the Nernst equation, spontaneity and Faraday-type electrolysis calculations.
Core rules
- Anode is oxidation, cathode is reduction. Electrons flow from anode to cathode through the external circuit.
- E°_cell = E°_cathode - E°_anode, using reduction potentials for both; coefficients never multiply the potential.
- ΔG° = -nFE° with F = 96 485 C mol-1. A positive E°_cell means the standard reaction is spontaneous.
- Nernst: E = E° - (0.0592/n)·log Q at 25 °C, with Q taken from the balanced equation and pure solids and liquids omitted.
- Electrolysis: Q = It, moles of electrons = It/F, mass deposited = (It/nF) × M.
- Cell notation: Zn(s) | Zn2+(aq) || Cu2+(aq) | Cu(s), anode on the left.
Worked example
Zn | Zn2+(0.10 mol dm-3) || Cu2+(1.0 mol dm-3) | Cu has E°_cell = +1.10 V. The reaction Zn + Cu2+ → Zn2+ + Cu gives Q = [Zn2+]/[Cu2+] = 0.10 and n = 2. E = 1.10 - (0.0592/2)·log(0.10) = 1.10 + 0.0296 = 1.13 V. The dilute zinc solution makes oxidation easier, so the voltage rises slightly, which is a useful check that Q is the right way up. A concentration cell uses the same principle and has E° = 0.
Descriptive links
Rusting is electrochemical: oxygen reduction at cathodic regions and Fe → Fe2+ at anodic regions, with sacrificial anodes exploiting a more negative reduction potential. In electrolysis the anode is the positive terminal, the opposite sign convention from a galvanic cell.
Traps
- Reversing the anode sign twice, or doubling E° when balancing electron transfer.
- Treating potential as extensive: ΔG scales with n, E° does not.
- Inverting Q in the Nernst equation.
- Applying the galvanic sign convention to an electrolytic cell.
- Forgetting that molten-salt electrolysis has no water, so the discharge rules change.