# Standard Electrode Potential Calculator

Calculate standard cell potential from half-reaction potentials. Find ΔG°, spontaneity, and equilibrium constant for electrochemical cells.

## What this calculates

Calculate the standard cell potential (E°cell) from cathode and anode half-reaction potentials. Also determines the standard Gibbs free energy change (ΔG°), whether the reaction is spontaneous, and the equilibrium constant (Keq).

## Inputs

- **Cathode Potential E°(cathode)** (V) — Standard reduction potential of the cathode half-reaction in volts.
- **Anode Potential E°(anode)** (V) — Standard reduction potential of the anode half-reaction in volts.
- **Electrons Transferred (n)** — min 1, max 20 — Number of electrons transferred in the balanced cell reaction.
- **Temperature** (K) — min 0.01 — Temperature in Kelvin (default 298.15 K = 25 °C).

## Outputs

- **E°cell** (V) — Standard cell potential.
- **ΔG°** (kJ/mol) — Standard Gibbs free energy change.
- **Spontaneity** — formatted as text — Whether the reaction is spontaneous under standard conditions.
- **Keq** — Equilibrium constant at the given temperature.

## Details

Electrochemical cells convert chemical energy to electrical energy (galvanic cells) or vice versa (electrolytic cells). The standard cell potential determines the voltage a cell can produce under standard conditions.

Key Equations

- E°cell = E°(cathode) - E°(anode): The cell potential is the difference between the reduction potentials of the two half-reactions. The cathode is where reduction occurs; the anode is where oxidation occurs.

- ΔG° = -nFE°cell: Links cell potential to thermodynamics. F = 96,485 C/mol (Faraday constant), n = electrons transferred. A positive E°cell gives a negative ΔG° (spontaneous).

- Keq = e^(nFE°/RT): Links cell potential to the equilibrium constant. Even small positive E°cell values correspond to very large Keq values.

Example: Daniell Cell (Zn-Cu)

Cathode: Cu²⁺ + 2e⁻ → Cu (E° = +0.34 V). Anode: Zn²⁺ + 2e⁻ → Zn (E° = -0.76 V). E°cell = 0.34 - (-0.76) = 1.10 V. ΔG° = -2(96485)(1.10)/1000 = -212.3 kJ/mol. The large positive E°cell and very large Keq confirm this reaction strongly favors products.

## Frequently Asked Questions

**Q: What is the difference between cathode and anode potentials?**

A: Both are standard reduction potentials from tables. The cathode is where reduction occurs (gains electrons), and the anode is where oxidation occurs (loses electrons). E°cell = E°cathode - E°anode. The half-reaction with the more positive E° is the cathode.

**Q: What does a negative cell potential mean?**

A: A negative E°cell means the reaction is non-spontaneous under standard conditions. The reverse reaction would be spontaneous. To drive the reaction forward, you would need to apply an external voltage (electrolysis). The minimum voltage required is |E°cell|.

**Q: What is the Faraday constant?**

A: The Faraday constant (F = 96,485 C/mol) is the electric charge per mole of electrons. It connects the macroscopic voltage of a cell to the microscopic transfer of electrons, allowing conversion between electrical energy and chemical energy.

**Q: How do I find standard reduction potentials?**

A: Standard reduction potentials are measured relative to the Standard Hydrogen Electrode (SHE, E° = 0.000 V). They are tabulated in chemistry references. Common values: Cu²⁺/Cu = +0.34 V, Zn²⁺/Zn = -0.76 V, Ag⁺/Ag = +0.80 V, Fe²⁺/Fe = -0.44 V.

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Source: https://vastcalc.com/calculators/chemistry/standard-electrode-potential
Category: Chemistry
Last updated: 2026-04-21