# Cell Potential Calculator - Standard Reduction Potentials, Nernst Equation

> Standard reduction potential table with cell reaction, notation, E°cell, ΔG° and K, plus Nernst equation, concentration cells and Faraday electrolysis.

- URL: https://qilingtools.com/en/math/electrochem
- Category: Chemistry Tools (https://qilingtools.com/en/chemistry)
- Tags: [Chemistry](https://qilingtools.com/en/tags/chemistry)
- Last updated: 2026-10-26
- Price: Free, no sign-up
- Privacy: Runs locally in the browser; files are never uploaded

A standard reduction potential E° measures how strongly a couple wants to gain electrons: the higher E°, the stronger the oxidizing agent. In a galvanic cell the half-reaction with the higher E° is the cathode (reduction, positive terminal) and the lower one is the anode (oxidation, negative terminal), so E°cell = E°cathode − E°anode. For the Daniell cell, Cu²⁺/Cu is +0.3419 V and Zn²⁺/Zn is −0.7618 V, giving E°cell = 1.104 V, n = 2, ΔG° = −nFE° ≈ −213 kJ/mol, K ≈ 10^37 and the cell notation Zn(s) | Zn²⁺(aq) ‖ Cu²⁺(aq) | Cu(s). This calculator holds more than 80 common half-reactions in acidic and basic solution (CRC Handbook and Bard values). Choose any two and it scales them to a common electron count, cancels species that appear on both sides, writes the balanced overall reaction and the cell notation, computes E°cell, ΔG° and K, and draws a cell diagram showing electron flow, the salt bridge and the electrode signs. The Nernst tab takes concentrations, gas pressures and temperature to give the reaction quotient Q, the actual E = E° − (RT/nF) ln Q and each electrode potential. There is also a concentration-cell calculator and a Faraday's-law electrolysis tab that solves m = ItM/(nF) for mass, time or current.

## How to use

1. On the Galvanic cell tab choose two half-reactions (or use Cathode / Anode buttons in the E° table); the higher E° becomes the cathode
2. Read the overall reaction, cell notation, E°cell, ΔG°, K and the diagram; switch to Nernst to enter concentrations and temperature
3. Use the concentration-cell or electrolysis tab to find a cell voltage, deposited mass, time or current

## Features

- 80+ standard reduction potentials (V vs SHE, 25 °C), searchable, incl. basic-solution couples
- Automatic cathode/anode assignment, half-reactions combined by LCM of electrons with spectator cancelling
- Line notation such as Zn(s) | Zn²⁺(aq) ‖ Cu²⁺(aq) | Cu(s)
- E°cell, electrons transferred n, ΔG° = −nFE° and K
- Nernst equation at any concentration, pressure and temperature, with each electrode potential
- Concentration cell voltage
- Faraday's law: solve for mass, time or current, with current efficiency and gas volume
- SVG galvanic cell diagram with electron flow, salt bridge and electrode signs

## FAQ

### How do I know which electrode is the cathode?

Compare the standard reduction potentials. The couple with the higher E° is reduced and is the cathode (positive terminal of a galvanic cell); the lower one is oxidized at the anode (negative terminal). Electrons flow through the wire from anode to cathode. The calculator does this automatically; switch off auto mode to assign them yourself — a negative E°cell means the reaction as written is not spontaneous.

### Why isn't E° multiplied when I multiply a half-reaction?

Electrode potential is an intensive property: it does not depend on how many times the reaction is written. The half-reactions are scaled only so the electrons cancel, and E°cell is still E°cathode − E°anode. ΔG° = −nFE° is extensive, so it does depend on the number of electrons n in the overall equation.

### How do I use the Nernst equation?

E = E° − (RT/nF) ln Q, which at 25 °C becomes E = E° − (0.05916/n) log Q. Q is the reaction quotient: products over reactants, each raised to its coefficient, with solids and pure liquids set to 1. In the Daniell cell with [Zn²⁺] = 0.1 M and [Cu²⁺] = 1 M, Q = 0.1 and E = 1.104 + 0.0296 ≈ 1.133 V.

### How do I calculate the mass deposited in electrolysis?

The charge passed is Q = It; moles of electrons = Q/F with F = 96485 C/mol; moles of product = moles of electrons ÷ n; multiply by the molar mass. Electrolysing CuSO₄ at 2 A for one hour gives Q = 7200 C, 0.0746 mol e⁻, 0.0373 mol Cu ≈ 2.37 g of copper.

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