nernst equation calculator

Calculate electrode potential using the Nernst equation for a redox half-cell.

E = E° − (RT / nF) ln(Q)
At 25°C (298.15 K): E = E° − (0.05916 / n) log₁₀(Q)

Standard electrode potential, E° (volts)

Number of electrons transferred, n

Temperature

Enter Q directly Use [Red]/[Ox] concentrations

Reaction quotient, Q

Q must be greater than 0.

Enter values and click "Calculate Potential".

What the Nernst equation tells you

The Nernst equation links electrode potential to concentration (more accurately, activity), temperature, and the number of electrons transferred in a redox reaction. In plain language, it tells you how far a real electrochemical system is from standard conditions.

Under standard conditions (1 M concentrations, 1 atm gases, pure solids/liquids, usually 25°C), the potential is E°. Once concentrations shift, the actual potential E changes, and the Nernst equation gives you that corrected value.

Variables in the equation

E: electrode potential under the given conditions (V)
E°: standard electrode potential (V)
R: gas constant (8.314462618 J·mol^-1·K^-1)
T: absolute temperature (K)
n: moles of electrons transferred
F: Faraday constant (96485.33212 C·mol^-1)
Q: reaction quotient

How to use this calculator

Option 1: Enter Q directly

Use this when you already know the reaction quotient from a balanced half-reaction or full cell reaction. Enter E°, n, temperature, and Q, then calculate.

Option 2: Use concentrations

For a simple reduction half-cell representation, this calculator can estimate Q from concentration ratio: Q = [Red]/[Ox]. This is convenient for quick checks in class and lab.

Worked example

Suppose E° = 0.34 V, n = 2, T = 25°C, and Q = 100.

At 25°C, use E = E° − (0.05916/n)log₁₀(Q)
E = 0.34 − (0.05916/2)log₁₀(100)
E = 0.34 − 0.05916 = 0.28084 V

A larger Q usually lowers reduction potential for the half-reaction as written, while Q less than 1 increases it.

Common mistakes to avoid

Using temperature in °C directly in the full equation (convert to K first).
Entering Q ≤ 0 (logarithms require positive values).
Using the wrong electron count n from an unbalanced equation.
Mixing up reaction direction when forming Q.
For high-accuracy work, ignoring activities and using concentrations blindly.

Practical notes for students and lab users

In introductory chemistry, concentration-based Q is usually acceptable. In research or concentrated solutions, activity coefficients can matter significantly. If your system is non-ideal, treat this as a strong estimate rather than an exact value.

Also remember: the sign and magnitude of potential depend on how the half-reaction is written. If you reverse the reaction, the sign behavior changes accordingly.

Quick FAQ

Is this for full cells or half-cells?

The formula applies to both. This interface is centered on half-cell potential calculation. For full-cell potential, compute each half-cell (or use overall E° and Q correctly) and combine as needed.

Can I use this for pH-dependent reactions?

Yes, if hydrogen or hydroxide appears in the balanced reaction and is included in Q. Just be careful with stoichiometric exponents.

Why does potential change with concentration?

Because reaction driving force depends on chemical potential. The Nernst equation is the electrochemical expression of that thermodynamic shift.