Redox Reactions and Electrochemistry
Redox Reactions and Electrochemistry tests oxidation numbers, balancing redox equations, galvanic and electrolytic cells, standard electrode potentials, the Nernst equation, and Faraday's laws. You must keep anode/cathode and oxidation/reduction paired correctly.
Core rules
- Oxidation is loss of electrons; reduction is gain. Anode: oxidation. Cathode: reduction.
- In a galvanic cell, E°cell = E°cathode − E°anode, using reduction potentials. A positive E°cell means spontaneous reaction.
- ΔG° = −nFE°cell. F = 96485 C/mol.
- Nernst equation at 298 K: E = E° − (0.059/n) log Q.
- Cell notation: anode | anode solution || cathode solution | cathode.
- Faraday's laws: moles of product relate to charge through nF.
Worked example
For Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s), E°cell = E°Cu2+/Cu − E°Zn2+/Zn = +0.34 − (−0.76) = +1.10 V. ΔG° = −nFE° = −2 × 96485 × 1.10 = −212,000 J/mol, or about −212 kJ/mol. The negative ΔG° confirms that the reaction is spontaneous under standard conditions.
Traps
Do not reverse the cathode and anode potentials. Do not multiply E° when balancing a cell reaction; E is intensive, while ΔG depends on n. In cell notation, oxidation occurs at the left electrode. For electrolysis, the applied voltage must exceed the back emf, and overpotential may matter.