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ChemBench
Standard electrode potentials tell you the starting point — the Nernst equation tells you what really happens.
Standard electrode potentials (E°) are measured at 25°C, 1 M concentrations, and 1 atm — real batteries and electrochemical cells almost never operate under those conditions. The Nernst equation adjusts E° for actual concentrations and temperature, predicting the real voltage a cell produces.
The Nernst equation connects electrochemistry to thermodynamics: E = E° − (RT/nF)ln(Q). When Q = K (equilibrium), E = 0 and the battery is dead. When Q < K, the reaction is spontaneous (positive E) and the cell produces electricity.
Molar mass of any chemical formula (e.g. H2O, C6H12O6, Ca(OH)2).
OpenPercent composition of an element within a chemical formula.
OpenSimplest whole-number ratio formula from elemental masses or percentages.
OpenPercent yield from actual and theoretical reaction yields.
OpenCell potential E
1.159
What you entered
ln(Q)
ln(0.01)= -4.6052RT ÷ (nF)
(8.314 × 298) ÷ (2 × 96485)= 0.0128E = E° − (RT/nF)·ln(Q)
1.1 − -0.0591= 1.1591 VResult
Cell potential E: 1.159
At 298 K with Q = 0.01, the cell potential is 1.159 V (standard E° = 1.1 V).