Nernst Oxidation-Reduction Calculator
Use this tool to calculate electrode potential (E) for a redox half-reaction under non-standard conditions. Enter your values and click calculate.
E = E° − (RT / nF) ln(Q)For the reduction form: Ox + ne⁻ ⇌ Red, where Q = [Red]/[Ox]
What this oxidation reduction calculator does
This oxidation reduction calculator is built around the Nernst equation, one of the most useful equations in electrochemistry. It helps you predict how cell potential changes when concentrations and temperature move away from standard-state conditions.
In practical terms, this means you can estimate whether a redox process is more or less favorable in real solutions than it appears from table values of E°. Students use this for homework and exam prep, while researchers and engineers use the same idea in batteries, corrosion studies, sensors, and environmental chemistry.
Quick redox refresher
Oxidation and reduction
- Oxidation is loss of electrons.
- Reduction is gain of electrons.
- These always happen together in a coupled oxidation-reduction (redox) process.
Potential and spontaneity
For a reduction half-reaction written as Ox + ne⁻ ⇌ Red, a more positive reduction potential generally means stronger tendency for reduction. If the computed E for your conditions is positive, the reduction direction (as written) is favored. If E is negative, the reverse direction is favored.
How the calculator works
The calculator uses:
E = E° − (RT / nF) ln(Q)
with constants:
- R = 8.314462618 J·mol⁻¹·K⁻¹
- F = 96485.33212 C·mol⁻¹
- T in kelvin (K)
- n = electrons transferred
- Q = reaction quotient = [Red]/[Ox] for the half-reaction form shown above
It also estimates:
- ΔG from ΔG = −nFE (kJ/mol output)
- log₁₀K from E° at the entered temperature, using log form of the equilibrium relation
Input guide and best practices
1) Standard potential E°
Use the value from a reliable electrochemical table and make sure the half-reaction direction matches reduction form.
2) Electron count n
Enter the stoichiometric number of electrons in that half-reaction. This must be a positive number.
3) Concentrations
Use molar concentrations (M). Concentrations must be greater than zero, since logarithms are undefined for zero or negative values.
4) Temperature
The default is 25°C, but changing temperature lets you explore non-standard thermal conditions.
Worked example
Suppose a half-cell has E° = 0.34 V, n = 2, T = 25°C, [Ox] = 1.0 M, and [Red] = 0.10 M.
- Q = [Red]/[Ox] = 0.10
- ln(Q) is negative, so subtracting it increases E
- The calculated E becomes slightly larger than E°
This makes chemical sense: lower reduced-species concentration shifts the system toward reduction for the reaction written, increasing the instantaneous driving force.
Common mistakes to avoid
- Mixing up oxidation and reduction forms of the same half-reaction.
- Using Celsius directly in formulas instead of converting to kelvin.
- Setting concentration to zero.
- Forgetting that activities (not raw concentration) are the formal thermodynamic quantity in non-ideal solutions.
Where oxidation-reduction calculations are used
- Battery voltage prediction and state-of-charge estimation
- Corrosion tendency and metal protection analysis
- Electroplating and industrial electrolysis
- Water treatment and environmental redox chemistry
- Biological electron-transfer systems
Final note
This calculator is an educational and practical planning tool. For high-precision lab work, include activity coefficients, ionic strength effects, and full-cell balancing across both half-reactions.