Electrochemistry
Electrochemistry tests galvanic and electrolytic cells, standard electrode potentials, the Nernst equation, conductance and Faraday’s laws. Cell notation and EMF numericals are common.
Core rules
- In a galvanic cell, oxidation occurs at the anode (negative) and reduction at the cathode (positive).
- E°cell = E°cathode - E°anode.
- Nernst: E = E° - (0.059/n) log Q at 298 K.
- Molar conductivity is lambda-m = kappa/c; it increases on dilution for weak and strong electrolytes.
- Kohlrausch’s law: limiting molar conductivity is the sum of ionic contributions.
- Faraday’s first law: mass deposited is proportional to charge; m = (Q/F)(M/z).
Worked example
For Zn|Zn2+||Cu2+|Cu, E° = 1.10 V. Calculate E when [Zn2+] = 0.10 M and [Cu2+] = 0.010 M. Reasoning: the cell reaction is Zn + Cu2+ -> Zn2+ + Cu, so Q = [Zn2+]/[Cu2+] = 0.10/0.010 = 10. n = 2. E = 1.10 - (0.059/2) log 10 = 1.10 - 0.0295 = 1.07 V.
Traps
- Do not reverse the anode and cathode in E°cell.
- E° is intensive; multiplying a half-reaction does not multiply E°.
- In lambda-m = kappa/c, use consistent units for kappa and c.
- Corrosion of iron is an electrochemical process; the anodic reaction is Fe -> Fe2+ + 2e-.