Oxidation-Reduction and Electrochemistry
Oxidation–Reduction and Electrochemistry covers oxidation numbers, balancing redox equations, galvanic cells, standard electrode potentials, the Nernst equation and electrolysis.
Core rules
- Oxidation numbers: free element 0; oxygen usually −2 (peroxides −1); hydrogen +1 with non-metals, −1 with metals; the sum equals the overall charge.
- Balance redox in acid by adding H₂O and H⁺; in base, balance in acid first, then add OH⁻ to both sides.
- E°cell = E°cathode − E°anode. A positive value means the reaction is spontaneous and ΔG is negative.
- Nernst at 25 °C: E = E° − (0.0592/n) log Q.
- Electrolysis: reduction at the cathode, oxidation at the anode; n(e⁻) = It/F with F = 96 500 C/mol.
Worked example
Balance MnO₄⁻ + Fe²⁺ → Mn²⁺ + Fe³⁺ in acid and find E°cell. Reduction: MnO₄⁻ + 8H⁺ + 5e⁻ → Mn²⁺ + 4H₂O (E° = +1.51 V). Oxidation: Fe²⁺ → Fe³⁺ + e⁻ (E° = +0.77 V). Equalise electrons: MnO₄⁻ + 5Fe²⁺ + 8H⁺ → Mn²⁺ + 5Fe³⁺ + 4H₂O. E°cell = 1.51 − 0.77 = +0.74 V, so the reaction is spontaneous as written and permanganate decolourises.
Traps
- Reversing the sign of the anode potential instead of subtracting; tabulated E° values are always reduction potentials.
- Forgetting to multiply the iron half-reaction by five - both atoms and charge must balance.
- Calling the more negative electrode the cathode in a galvanic cell; it is the anode.
- Using n from a single half-reaction in the Nernst equation instead of electrons transferred in the overall reaction.