Key ideas
Electrochemistry connects chemical reactions to electrical energy. In a galvanic cell a spontaneous redox reaction produces voltage, and in electrolysis an applied current drives a reaction.
- Oxidation happens at the anode and reduction at the cathode. The cell potential is E°_cell = E°_cathode − E°_anode, and a positive value means a spontaneous cell.
- The free energy is ΔG° = −nFE°, with F = 96 485 C/mol and n the number of electrons transferred.
- Away from standard conditions the Nernst equation gives E = E° − (0.0592 ÷ n) log₁₀ Q at 25 °C.
- Faraday’s law for electrolysis: moles of electrons = It ÷ F, and the mass deposited is m = ItM ÷ (nF).
Key equations
E°_cell = E°_cathode − E°_anodeΔG° = −nFE°E = E° − (0.0592 ÷ n) log₁₀ Q (25 °C)m = ItM ÷ (nF)F = 96 485 C/mol
Worked example
For Zn | Zn²⁺ ‖ Cu²⁺ | Cu, E° = 0.34 − (−0.76) = 1.10 V, and ΔG° = −2 × 96 485 × 1.10 = −212 kJ/mol. With [Zn²⁺] = 1.0 M and [Cu²⁺] = 0.010 M, Q = 100 and E = 1.10 − (0.0592 ÷ 2) × log 100 = 1.10 − 0.059 = 1.04 V. Plating copper at 2.0 A for 30 min deposits m = 2.0 × 1800 × 63.5 ÷ (2 × 96 485) = 1.2 g.
Common mistakes
- Subtracting the potentials in the wrong order or changing E° when you multiply a half-reaction by a coefficient.
- Forgetting that n in the Nernst and ΔG formulas is the number of electrons in the balanced equation.
- Mixing up anode and cathode. Oxidation is always at the anode, in both galvanic and electrolytic cells.
Practice
Test yourself with practice questions on this topic: Electrochemistry