Electrochemistry
Electrochemistry
This domain tests galvanic cells, the Nernst equation, conductance, electrolysis and corrosion.
Core rules
- E°cell = E°cathode − E°anode; a positive E°cell means ΔG° is negative and K>1.
- ΔG° = −nFE° = −RT ln K; at 298 K, E° = (0.059 V/n) log K.
- Nernst equation: E = E° − (0.059/n) log Q at 298 K.
- Λm = 1000κ/M; Λm increases on dilution while κ decreases.
- Kohlrausch law: Λ°m of a weak electrolyte equals the sum of limiting ionic conductivities.
- Faraday: Q=It; moles deposited = Q/(nF).
Worked example Zn|Zn²⁺(0.10 M)||Cu²⁺(0.010 M)|Cu, E°=1.10 V, n=2. Q=[Zn²⁺]/[Cu²⁺]=0.10/0.010=10. E=1.10 − (0.059/2)log10 = 1.10 − 0.0295 = 1.07 V. Increasing Zn²⁺ or decreasing Cu²⁺ lowers the emf, as Le Chatelier predicts.
Traps
- Reversing the anode sign or using E°oxidation instead of reduction.
- Using 0.059 V outside 298 K; use 2.303RT/F otherwise.
- Conductivity units: κ in S cm⁻¹, Λm in S cm² mol⁻¹.
- In electrolysis, the product depends on electrode material and overpotential, not just E°.