⚡ Nernst Equation
Animate a galvanic cell and watch cell potential E change as you vary temperature, ion concentrations, and electrode material. Grounded in E = E° − (RT/nF) ln Q with live ΔG and equilibrium constant readouts.
Electrode Pair
Conditions
Live Readouts
About the Nernst Equation
Derivation from Gibbs Energy
For a half-cell or full cell reaction, the Gibbs free energy change is ΔG = ΔG° + RT ln Q. The electrical work done by the cell equals ΔG = −nFE, so E = E°−(RT/nF) ln Q. At 25 °C (298.15 K) and using log base 10, RT/F ≈ 0.02569 V, giving the common approximation E = E°−(0.0592/n) log Q.
The Daniell Cell
The historic Daniell cell (1836) uses a zinc anode in ZnSO&sub4; solution and a copper cathode in CuSO&sub4; solution, connected by a salt bridge. Zn oxidises (Zn → Zn²&sup+; + 2e−, E° = −0.76 V) and Cu²&sup+ reduces (Cu²&sup+; + 2e− → Cu, E° = +0.34 V), giving E° = 1.10 V. As Zn ion concentration rises and Cu ion concentration falls, Q increases and E decreases until equilibrium (E = 0).
Equilibrium and the EMF
At equilibrium ΔG = 0 and E = 0, so ln K = nFE°/RT. This connects the thermodynamic equilibrium constant K to the standard cell potential E°. A cell with E° > 0 has K > 1 (products favoured), while E° < 0 means K < 1. The Nernst equation quantifies how far Q is from K and hence how much remaining driving force the cell retains.
Temperature Dependence
The RT/nF prefactor increases linearly with T. At higher temperatures concentration effects are amplified: a tenfold change in Q shifts E by (RT/nF) ln 10 ≈ (0.0592 V/n) × T/298. Industrial electrochemical processes (aluminium smelting, chlor-alkali electrolysis, lithium-ion batteries) all rely on precise Nernst-equation calculations to optimise operating voltages and efficiencies.