Cell EMF Calculator (standard cell potential E°cell, ΔG° and K from half-reactions)

Understanding the cell EMF (standard cell potential) calculator — concept and principle

A galvanic cell produces a voltage from a redox reaction. This calculator finds the standard cell potential from two standard reduction potentials, along with the number of electrons transferred, the standard Gibbs energy and the equilibrium constant.

The cell voltage is the gap between the two reduction potentials: 0.34 − (−0.76) = 1.10 V, so it runs on its own.

How it is calculated

E°cell = E°cathode − E°anode, using both half-reactions written as reductions.
The anode reaction actually runs as an oxidation, but its E° is not multiplied when the equation is scaled.
n = least common multiple of the electrons in the two half-reactions.
ΔG° = −nFE°cell; K = e^(nFE° ÷ RT) at 298.15 K, with F = 96,485.33212 C/mol and R = 8.314462618 J/(mol·K).
A positive E°cell means a spontaneous (galvanic) cell.

The principle

Each half-reaction has a tendency to gain electrons, measured against the standard hydrogen electrode. In a cell, the half-reaction with the higher reduction potential takes electrons (the cathode) and forces the other to give them up (the anode). The difference in potentials is the driving force. Because potential is energy per charge, it does not scale with the amount of reaction, which is why E° is never multiplied. Free energy and the equilibrium constant follow from thermodynamics.

Worked example

Default inputs: cathode Cu²⁺ + 2e⁻ → Cu (+0.34 V); anode Zn²⁺ + 2e⁻ → Zn (−0.7618 V) (the Daniell cell).

E°cell = 0.34 − (−0.7618) = 1.1018 V.
n = 2.
ΔG° = −2 × 96,485 × 1.1018 ≈ −212.6 kJ/mol.
log₁₀ K = 2 × 1.1018 ÷ 0.05916 ≈ 37.25, so K ≈ 1.772 × 10³⁷.
E° is positive: spontaneous, a galvanic cell.

How to read the result

A huge K means the reaction goes essentially to completion: zinc reduces copper ions almost entirely. Real cells give less than E° when concentrations are not 1 M, current flows or the temperature differs; use the Nernst equation for non-standard conditions.

Common mistakes

Multiplying E° by stoichiometric coefficients.
Reversing the sign of the anode potential and then subtracting it again.
Using oxidation potentials from one table with reduction potentials from another.
Forgetting that standard values assume 25 °C, 1 M and 1 bar.

Keep in mind: Standard potentials vary slightly between tables. Real cell voltages under load are lower than E°.

Frequently asked questions

What is the EMF of a Daniell cell?

About 1.10 V under standard conditions.

How do I know which electrode is the cathode?

The half-reaction with the more positive reduction potential is the cathode in a galvanic cell.

Why isn't E° multiplied when balancing?

Potential is an intensive quantity, energy per unit charge, so it does not depend on the amount of reaction.

Sources: Bard, A. J., Parsons, R. & Jordan, J. (eds.), Standard Potentials in Aqueous Solution. IUPAC / Marcel Dekker (1985) · CRC Handbook of Chemistry and Physics, Electrochemical Series

Checked 9 Nov 2026. Standard reduction potentials: OpenStax, Chemistry: Atoms First 2e, Appendix L “Standard Electrode (Half-Cell) Potentials”. E°_cell = E°_cathode − E°_anode, ΔG° = −nFE°, K = exp(nFE°/RT): OpenStax Chemistry 2e §17.3–17.4.

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