Electrochemistry
What this domain tests
Conductance and molar conductivity, Kohlrausch's law, the Nernst equation, EMF and Gibbs energy, equilibrium constants, commercial cells and corrosion.
Core rules
- Conductance κ = G × (l/A); the ratio l/A is the cell constant, in cm⁻¹.
- Molar conductivity Λm = 1000κ/M, with κ in S cm⁻¹ and M in mol L⁻¹, giving S cm² mol⁻¹.
- Kohlrausch: Λm° = ν₊λ₊° + ν₋λ₋°. The Λm° of a weak acid is obtained by combining strong electrolytes whose values are measurable.
- Nernst at 298 K: E = E° − (0.0591/n) log Q, where n is the number of electrons in the balanced cell reaction and Q uses products over reactants.
- ΔG° = −nFE° and log Kc = nE°cell/0.0591.
- Rusting is electrochemical: iron dissolves at the anodic site, while O2 + 2H2O + 4e⁻ → 4OH⁻ occurs at the cathodic site. Galvanising and sacrificial anodes protect by making iron cathodic.
Worked example
For Daniell cell Zn|Zn²⁺||Cu²⁺|Cu, E° = 1.10 V. With [Zn²⁺] = 0.1 M and [Cu²⁺] = 0.01 M, Q = [Zn²⁺]/[Cu²⁺] = 10 and n = 2. E = 1.10 − (0.0591/2) log 10 = 1.10 − 0.0296 ≈ 1.07 V. Increasing the anode ion concentration lowers the EMF, as the equation predicts.
Traps that cost marks
Use log₁₀ with 0.0591, not ln. Reversing the cell reverses the sign of E°cell. On dilution, Λm rises slightly for strong electrolytes and steeply for weak ones; do not assume the same shape. Check which medium is specified for a redox couple, since n changes. Note that a salt bridge maintains electrical neutrality, not electron flow.