Electrochemistry
Electrochemistry tests galvanic cells, standard electrode potentials, the Nernst equation, conductance and Kohlrausch's law, Faraday's laws, commercial cells and corrosion.
Cells and potentials
Oxidation occurs at the anode, reduction at the cathode. E°cell = E°cathode − E°anode, with both values written as reduction potentials. The standard hydrogen electrode (Pt, H₂ at 1 bar, H⁺ at 1 M) is fixed at 0.00 V. A positive E°cell means ΔG° = −nFE° is negative, so the reaction is spontaneous.
Nernst equation
At 298 K, E = E° − (0.0591/n)·log Q. For Zn(s) + Cu²⁺ → Zn²⁺ + Cu, Q = [Zn²⁺]/[Cu²⁺] because pure solids have unit activity.
Conductance
Λm = κ × 1000/M gives S cm² mol⁻¹ when κ is in S cm⁻¹ and M in mol L⁻¹. Λm rises on dilution: sharply for weak electrolytes (degree of dissociation rises) and slightly for strong electrolytes (ion–ion attraction falls). Kohlrausch's law, Λ°m = ν₊λ°₊ + ν₋λ°₋, gives Λ° for weak electrolytes, hence α = Λm/Λ°m and Ka = Cα²/(1 − α).
Worked example
Zn|Zn²⁺(0.10 M)||Cu²⁺(0.010 M)|Cu with E° = 1.10 V. Q = 0.10/0.010 = 10 and n = 2. E = 1.10 − (0.0591/2)(log 10) = 1.10 − 0.0296 = 1.07 V. The cell still delivers current (E > 0), but the voltage is below the standard value because the products are relatively more concentrated than the reactants.
Traps that cost marks
- Reversing the sign convention: E°cell is always cathode minus anode.
- Using n = 1 when two electrons are transferred, or including a solid in Q.
- Mixing κ in S m⁻¹ with Λm in S cm² mol⁻¹ (1 S m² mol⁻¹ = 10⁴ S cm² mol⁻¹).
- Forgetting that in electrolysis the product at each electrode depends on preferential discharge, not on the cell diagram.