Each half-cell has a standard reduction potential E° (25 °C, 1 M, from a real table). At non-standard concentration, the Nernst equation corrects it:
Mⁿ⁺ + ne⁻ → M
E = E° + (0.05916/n)·log₁₀[Mⁿ⁺] (25 °C, base-10 form)
E°cell = E°cathode − E°anode
Ecell = Ecathode − Eanode (Nernst-corrected)
The electrode with the higher effective potential is the cathode (reduction, gains mass); the other is the anode (oxidation, loses mass). If Ecell > 0 the reaction is spontaneous — a galvanic cell that drives current through the external wire from anode to cathode on its own, exactly as read on the voltmeter.
In electrolytic mode an external power supply forces current the other way. That only happens once the applied voltage exceeds the cell's own back-EMF (Ecell): net driving voltage = Vapplied − Ecell. Below that threshold nothing flows — this is the real reason electrolysis needs "enough" voltage, not just any voltage.
- Yellow dots on the wire — electrons; they always flow from the anode (oxidation, releasing e⁻) to the cathode (reduction, consuming e⁻), reversing on the wire when electrolytic mode overcomes E°cell.
- Salt bridge dots — spectator ions maintaining electroneutrality: anions drift toward the anode compartment, cations toward the cathode compartment.
- Electrode glow — brighter on the electrode currently consuming ions (cathode, metal deposits) vs. the one dissolving (anode).