Electrochemistry and Redox
What this domain tests
Redox balancing, cell potentials, predicting electrolysis products and Faraday calculations.
Core rules
- Oxidation is loss of electrons and reduction is gain; oxidation numbers track electron transfer.
- E°cell = E°cathode − E°anode, using standard reduction potentials. A positive E°cell means the reaction is feasible; ΔG = −nFE.
- In any cell, oxidation occurs at the anode and reduction at the cathode. The couple with the more positive reduction potential forms the cathode.
- Electrolysis: the cathode attracts cations. In aqueous solution, competing reductions (water/H+) and oxidations (water/OH−, electrode material) must be compared; concentrated halides favour halogen formation at the anode.
- Q = It and moles of electrons = Q/F, with F = 96 500 C mol−1.
Worked example
For Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s), E°(Cu2+/Cu) = +0.34 V and E°(Zn2+/Zn) = −0.76 V. E°cell = 0.34 − (−0.76) = 1.10 V. Zinc is oxidised at the anode; copper(II) ions are reduced at the cathode. The positive value confirms the reaction is spontaneous under standard conditions.
Traps
Reversing the sign of the anode potential instead of subtracting it; forgetting n in ΔG = −nFE; using 96 500 with current in mA or time in minutes without converting; assuming the most concentrated ion always discharges; confusing anode and cathode between galvanic and electrolytic cells.