Galvanic Cells and Electrolysis
A galvanic cell (electrochemical cell) spontaneously converts the energy of a redox reaction into electrical energy. Conversely, an electrolytic cell uses an external electrical source to drive a non-spontaneous reaction. Both devices are the cornerstones of practical electrochemistry.
Half-cells and the Anode/Cathode Convention
Every electrochemical cell consists of two half-cells, each housing a redox couple Ox/Red. By universal convention:
- Anode: electrode where oxidation occurs (electrons lost). In a galvanic cell the anode is the negative terminal; in an electrolytic cell it is connected to the positive terminal of the power supply.
- Cathode: electrode where reduction occurs (electrons gained). In a galvanic cell the cathode is the positive terminal.
The salt bridge (KCl solution or agar gel) maintains electrical neutrality in both compartments by allowing ion migration without mixing the electrolyte solutions.
Electromotive Force (emf) and Thermodynamics
The emf E_cell of a galvanic cell is directly linked to the Gibbs energy of the overall reaction:
ΔG = −nFE_cell
where n is the number of electrons exchanged and F = 96 485 C·mol⁻¹ is Faraday's constant. For a spontaneous cell, ΔG < 0, so E_cell > 0.
The Nernst equation extends this to non-standard concentrations:
E = E° − (RT/nF) ln Q
At 298 K, RT/F ≈ 0.02569 V, giving the convenient form:
E = E° − (0.0592/n) log Q (in volts, at 25 °C)
Electrolysis: Faraday's Laws
During electrolysis, Faraday's first law states that the mass m of substance deposited (or dissolved) is proportional to the total charge Q = It:
m = (M / nF) × I × t
where M is the molar mass and I the current intensity. Faraday's second law states that equal charges deposit masses proportional to M/n across different substances.
Industrial applications: - Aluminium production (Hall-Héroult process, Carbon (C) cathode, cryolite bath). - Water electrolysis: 2 H₂O → 2 H₂ + O₂ (alkaline or PEM electrolysers). - Electroplating: depositing Copper (Cu) onto substrates.
Galvanic Cell vs. Electrolytic Cell — Summary
| Parameter | Galvanic cell | Electrolytic cell |
|---|---|---|
| Reaction ΔG | < 0 (spontaneous) | > 0 (non-spontaneous) |
| Energy source | Chemical reaction | External current |
| Anode | (−) terminal | (+) terminal of supply |
| Cathode | (+) terminal | (−) terminal of supply |
| Examples | Daniell cell, Li-ion battery | Water electrolysis, electroplating |
Mastering both devices is essential for understanding modern batteries, electrochemical corrosion, and electrolytic synthesis processes.