Electrochemistry
Electrochemistry covers galvanic and electrolytic cells, Nernst equation, conductance, Kohlrausch's law, Faraday's laws and commercial cells. Numerical EMF and conductivity questions are standard.
Core rules
- E°cell = E°cathode - E°anode (reduction potentials).
- Nernst: E = E° - (0.059/n) log Q at 298 K.
- Conductivity κ = (1/R) x cell constant; molar conductivity Λm = κ x 1000/M.
- Kohlrausch: Λm° = ν+λ+° + ν-λ-° for weak electrolytes.
- Faraday: mass deposited m = (MIt)/(nF), F = 96500 C mol^-1.
- In a galvanic cell, oxidation occurs at the anode and reduction at the cathode; electrons flow anode to cathode in the external circuit.
Worked example
Zn|Zn2+(0.1 M)||Cu2+(1 M)|Cu has E° = 1.10 V. Q = [Zn2+]/[Cu2+] = 0.1/1 = 0.1; n = 2. E = 1.10 - (0.059/2) log 0.1 = 1.10 + 0.0295 = 1.13 V. Reasoning: log 0.1 = -1, and a minus times a minus becomes positive, so the cell voltage rises as the product concentration falls.
Traps
Using ln instead of log with 0.059, wrong n, reversing Q, and forgetting that molar conductivity decreases on dilution for strong electrolytes while increasing for weak ones. Cell constant must be in cm^-1 if κ is in S cm^-1.