| # | Anode | Cathode | E°cell | E_N | ΔG |
|---|---|---|---|---|---|
| No data yet | |||||
The standard electrode potential measures the tendency of a species to be reduced at 25 °C, 1 atm, with all ions at 1.00 M concentration. It is measured vs. the Standard Hydrogen Electrode (SHE), defined as 0.00 V.
A more positive E° means the species is a stronger oxidising agent and preferentially undergoes reduction. The half-cell with the highest E° acts as the cathode (reduction), while the lower E° half-cell acts as the anode (oxidation).
Galvanic (voltaic) cells convert spontaneous chemical energy into electrical energy. They require E°cell > 0 for spontaneous operation (ΔG < 0).
Electrolytic cells use external electrical energy to drive non-spontaneous reactions. An external power supply forces electrons against the natural potential gradient — key for electroplating and electrolysis of water.
At non-standard conditions the cell potential is given by the Nernst equation:
Where: R = 8.314 J mol⁻¹ K⁻¹, T in Kelvin, n = moles of electrons, F = 96485 C mol⁻¹, Q = reaction quotient.
At 25 °C this simplifies to:
When Q < 1, ln Q < 0, so E > E°. As Q increases, the cell potential decreases until equilibrium (E = 0, Q = K).
For a concentration cell (same metal/ion, E° = 0):
First Law: The mass of substance deposited is proportional to the total charge passed.
Second Law: The mass deposited is inversely proportional to the equivalent weight (M/n) of the substance.
Where F = 96485 C mol⁻¹ (Faraday's constant), n = oxidation state change, M_r = molar mass.
For a spontaneous reaction: E°cell > 0 → ΔG° < 0.
For a non-spontaneous reaction: E°cell < 0 → ΔG° > 0 (requires energy input).