Electrochemistry
This section covers Electrochemistry: conductance, molar conductivity, Kohlrausch's law, electrochemical cells, Nernst equation, EMF, Gibbs energy, electrolysis and corrosion. WBCHSE numericals commonly use the Nernst equation and ΔG = −nFE.
Core rules
- Conductivity decreases on dilution because fewer ions per unit volume; molar conductivity increases on dilution because dissociation increases.
- Kohlrausch's law: limiting molar conductivity of an electrolyte is the sum of independent ionic contributions.
- For a cell, E_cell = E°_cell − (0.0591/n) log Q at 25°C. ΔG = −nFE and ΔG° = −nFE°.
- Oxidation occurs at the anode; reduction at the cathode. In a galvanic cell the anode is negative; in electrolysis the anode is positive.
Worked example
For Zn(s) + Cu²⁺(aq) → Zn²⁺(aq) + Cu(s), E° = 1.10 V and n = 2. If [Zn²⁺] = 0.10 M and [Cu²⁺] = 0.010 M, Q = 0.10/0.010 = 10. E = 1.10 − (0.0591/2) log 10 = 1.10 − 0.0296 = 1.07 V. The cell still works because E is positive, but the EMF is lower than standard.
Traps
Do not reverse the Nernst log term; remember Q uses products over reactants; do not confuse conductivity with molar conductivity; check n from the balanced cell reaction; keep anode and cathode signs correct in galvanic versus electrolytic cells.