Galvanic and Electrolytic Cells: Differences & Uses

A galvanic (voltaic) cell converts chemical energy into electrical energy through a spontaneous redox reaction — this is what a battery does as it discharges. An electrolytic cell does the reverse: it uses electrical energy from an outside source to force a non-spontaneous reaction, as in electroplating or recharging a battery.

The two cells are mirror images, and students constantly mix up their terminals. The reliable rule is functional, not positional: oxidation always happens at the anode, reduction always at the cathode — in both cells. What flips between them is which terminal is positive.

Key Concepts

• Galvanic cell: anode = oxidation = negative terminal; cathode = reduction = positive terminal. Electrons flow through the external wire from anode to cathode; a salt bridge carries ions (not electrons) to complete the circuit. E°cell > 0, ΔG < 0.
• Electrolytic cell: an external DC supply drives the reaction; the anode (still oxidation) is now the positive terminal and the cathode (still reduction) is negative. E°cell < 0, ΔG > 0.
• Daniell cell: Zn(s)|Zn2+(aq)||Cu2+(aq)|Cu(s) — zinc anode dissolves, copper cathode plates out, E°cell = 1.10 V.
• Cell notation lists the anode on the left, cathode on the right, with || for the salt bridge.

Formulas

• E°cell = E°cathode − E°anode (standard reduction potentials).
• ΔG° = −nFE°, with F = 96485 C/mol and n = moles of electrons transferred.

Worked Example

Analyse the Daniell cell: Zn → Zn2+ + 2e− (E° = −0.76 V, anode) and Cu2+ + 2e− → Cu (E° = +0.34 V, cathode).

1. E°cell = E°cathode − E°anode = 0.34 − (−0.76) = 1.10 V.
2. n = 2 electrons; ΔG° = −nFE° = −2 × 96485 × 1.10 = −212,267 J/mol ≈ −212 kJ/mol.
3. ΔG° < 0 confirms the cell is galvanic and spontaneous as written.
4. To run it backwards (electrolytic copper refining in reverse), you must apply more than 1.10 V from an external supply.

Common Misconceptions

• "The anode is always negative." Only in a galvanic cell. In an electrolytic cell the anode is positive — the anode is defined by oxidation, not by charge sign.
• "Electrons flow through the salt bridge." Electrons travel through the external wire; the salt bridge carries ions to balance charge.
• "Electrolytic cells violate spontaneity rules." They don't — the external power supply provides the positive ΔG the reaction needs.

Real-World Applications

• Batteries and fuel cells are galvanic cells engineered for portable power.
• Electroplating (chrome bumpers, gold jewellery) and aluminium extraction are electrolytic processes.
• Rechargeable batteries run as galvanic cells when discharging and electrolytic cells when charging.

Remember the unchanging core — anode oxidises, cathode reduces — and let the energy direction tell you the signs: spontaneous chemistry makes its own voltage (galvanic), while imposed voltage forces unwilling chemistry (electrolytic).