What it does
The voltage of a battery, the potential of an electrode in a corroding pipe and the point where a redox titration turns all come from one equation. The Nernst equation corrects a standard potential E° for the real concentrations:
E = E° − (RT / nF) ln Q
At 25 °C that is E° − (0.05916 / n) log₁₀ Q, so the potential moves by 59.16 mV divided by the number of electrons for every tenfold change in Q.
The tool works in two ways:
- Cell (two electrodes). Pick the cathode (where reduction happens, the + terminal) and the anode (where oxidation happens, the − terminal). You get E°cell = E°cathode − E°anode, the cell voltage at your concentrations, the balanced overall reaction, the cell in line notation, ΔG = −nFE, the equilibrium constant K and the reaction quotient Q. A negative voltage means the cell would need power to run as written, and the tool suggests swapping the electrodes.
- One electrode. The potential of a single half-cell against the standard hydrogen electrode, also given against Ag/AgCl and SCE, with the Nernst slope and how far it has moved from E°.
A plot shows the voltage against log₁₀ Q with your conditions marked. A second card converts any potential between SHE, SCE, Ag/AgCl (saturated or 3 M KCl) and the reversible hydrogen electrode at your pH.
How to use it
- Choose the half-reactions from the lists. Each is written as a reduction with its standard potential at 25 °C. For a couple that is not in the table, choose “Your own half-reaction” and type E°, the number of electrons and Q.
- Type the concentrations in the table that appears: ions in mol/L, gases in bar, and H⁺ or OH⁻ as a pH. Solids, liquids and water have an activity of 1 and do not need a value.
- Change the temperature if you need to. The Nernst slope follows the temperature; E° stays at its 25 °C value.
- Read the result, check the plot, and download a PNG, PDF or Word report.
Worked example: a Daniell cell
Zinc in 1.0 M Zn²⁺ is joined by a salt bridge to copper in 0.010 M Cu²⁺.
- E°cell = +0.3419 − (−0.7618) = +1.1037 V.
- Cell reaction: Cu²⁺(aq) + Zn(s) → Cu(s) + Zn²⁺(aq), n = 2, written Zn(s) | Zn²⁺(aq) ‖ Cu²⁺(aq) | Cu(s).
- Q = [Zn²⁺]/[Cu²⁺] = 100, so Ecell = 1.1037 − (0.05916/2) × 2 = +1.0445 V.
- ΔG° = −2 × 96485 × 1.1037 = −213.0 kJ/mol and K ≈ 2 × 10³⁷: copper is plated out essentially completely.
For the oxygen electrode in neutral water open to air (pH 7, 0.21 bar O₂), E = 1.229 − 0.05916 × 7 − (0.05916/4) log(1/0.21) = +0.805 V, which is why oxygen corrodes iron so readily.
Questions
Why does the anode list show reductions? Tables of standard potentials are always written as reductions. The tool reverses the anode half-reaction for you and multiplies both halves so the electrons cancel. Do not multiply E° by these factors: potentials are intensive.
What about activity? Strictly, Q uses activities. Concentration is a good stand-in in dilute solutions; at 0.1 M and above, activity coefficients can shift the result by 10 mV or more.
Which reference values are used? SCE +0.241 V, Ag/AgCl in saturated KCl +0.197 V and in 3 M KCl +0.210 V, against SHE at 25 °C. RHE moves with pH: E(RHE) = E(SHE) + 0.05916 × pH.
Is anything sent anywhere? No. It all runs in your browser and works offline once the page has loaded.
Where the numbers come from
- Atomic weights: IUPAC/CIAAW standard atomic weights (abridged values).
- Standard reduction potentials: the CRC Handbook electrochemical series and Bard, Parsons and Jordan, Standard Potentials in Aqueous Solution, which you can type over.
- Physical constants: CODATA recommended values (NIST).
Standard scientific values, such as constants, atomic weights and melting points, are facts found in every textbook; we name where ours were checked. If you are an author or publisher and want a value, a credit or a table changed or removed, tell us and we will act quickly.
Last updated 2026-11-08.