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Redox Potential Calculator

Calculate cell potential, Gibbs free energy, and non-standard potentials

Oxidation occurs here
Reduction occurs here

Common Standard Potentials

Half-ReactionE° (V)
Li+ + e- → Li-3.04
Zn2+ + 2e- → Zn-0.76
2H+ + 2e- → H20.00
Cu2+ + 2e- → Cu+0.34
Ag+ + e- → Ag+0.80
Au+ + e- → Au+1.69

What is Redox Potential?

Redox potential, also called electrode potential or reduction potential, measures the tendency of a chemical species to gain electrons (undergo reduction) or lose electrons (undergo oxidation). It is a fundamental quantity in electrochemistry that determines the direction and driving force of electron transfer in redox reactions. Electrode potentials are measured relative to the Standard Hydrogen Electrode (SHE), which is assigned a potential of exactly 0.00 V.

Standard electrode potentials (E°) are measured under standard conditions: 1 M concentration for all dissolved species, 1 atm pressure for gases, and 25°C (298 K). Species with large positive E° values are strong oxidizing agents (they readily accept electrons), while species with large negative E° values are strong reducing agents (they readily donate electrons). The difference between two electrode potentials determines the cell potential of an electrochemical cell.

Under non-standard conditions, the actual electrode potential depends on concentration, temperature, and pressure through the Nernst equation. This equation corrects standard potentials for real-world conditions where concentrations deviate from the standard 1 M. The Nernst equation is essential for understanding batteries at different states of charge, corrosion behavior in different environments, and biological electrochemistry where ion concentrations vary significantly from standard conditions.

The Nernst Equation

The Nernst equation relates the actual electrode potential to the standard potential, temperature, and concentrations of reactants and products. It accounts for the effect of concentration on the thermodynamic driving force of a reaction.

At 25°C, the Nernst equation simplifies using the factor 0.0592 V (which equals RT/F × ln(10)), allowing easy calculation using base-10 logarithms. The reaction quotient Q is the ratio of product concentrations to reactant concentrations, each raised to the power of their stoichiometric coefficients.

Nernst Equation

E = E° − (RT/nF) × ln(Q)

Where:

  • E= Actual electrode potential under given conditions (V)
  • E°= Standard electrode potential (V)
  • R= Gas constant (8.314 J/mol·K)
  • T= Temperature in Kelvin
  • n= Number of electrons transferred
  • F= Faraday constant (96,485 C/mol)
  • Q= Reaction quotient [products]/[reactants]

Standard Cell Potential

The standard cell potential determines the thermodynamic favorability of a redox reaction under standard conditions. It is calculated from the difference between the cathode and anode reduction potentials.

The relationship between cell potential, Gibbs free energy, and equilibrium constant provides three interconnected ways to describe reaction thermodynamics. All three quantities must be consistent—a positive cell potential corresponds to a negative ΔG° and a K greater than 1.

Cell Potential Relationships

E°cell = E°cathode − E°anode ; ΔG° = −nFE°cell ; ln K = nFE°cell / RT

Where:

  • E°cell= Standard cell potential (V)
  • ΔG°= Standard Gibbs free energy (kJ/mol)
  • K= Equilibrium constant
  • n= Moles of electrons transferred
  • F= Faraday constant (96,485 C/mol)

How to Use This Calculator

This calculator supports two modes for electrochemical calculations:

  1. Standard Cell Potential Mode: Enter E°anode (oxidation) and E°cathode (reduction) voltages along with the number of electrons transferred. The calculator computes E°cell, ΔG°, and K.
  2. Nernst Equation Mode: Enter the standard potential E°, concentrations of oxidized and reduced species, and the number of electrons. The calculator computes the actual potential E under non-standard conditions.

Both modes use the temperature (default 298 K) and display intermediate calculations including the reaction quotient Q and the Nernst correction factor. The results indicate whether the reaction is spontaneous and provide the thermodynamic quantities that describe the reaction's driving force.

Real-World Applications

Redox potential measurements are essential for battery design and management. The voltage of a battery cell equals the difference between cathode and anode potentials. As the battery discharges, concentrations change, and the Nernst equation predicts how the cell voltage evolves with state of charge. Battery management systems in electric vehicles use Nernst-based models to estimate remaining capacity.

In biology, redox potentials govern electron transport in mitochondria and chloroplasts. The sequence of redox couples in the electron transport chain (NADH/FADH₂ to O₂) creates a cascade of decreasing potentials that drives ATP synthesis. Measuring redox potentials of biological couples helps understand metabolic pathways and design bioelectrochemical systems. In environmental science, redox potential determines the speciation and mobility of heavy metals in soil and groundwater, influencing remediation strategies for contaminated sites.

Worked Examples

Standard Cell Potential (Daniel Cell)

Problem:

Calculate E°cell, ΔG°, and K for the Daniel cell: Zn|Zn²⁺||Cu²⁺|Cu with n = 2.

Solution Steps:

  1. 1E°anode (Zn²⁺/Zn) = −0.76 V, E°cathode (Cu²⁺/Cu) = +0.34 V
  2. 2E°cell = E°cathode − E°anode = 0.34 − (−0.76) = 1.10 V
  3. 3ΔG° = −nFE°cell = −2 × 96,485 × 1.10 = −212.3 kJ/mol
  4. 4K = exp(nFE°cell / RT) = exp(85.7) ≈ 1.5 × 10³⁷

Result:

E°cell = 1.10 V, ΔG° = −212.3 kJ/mol, K ≈ 1.5 × 10³⁷

Nernst Equation Application

Problem:

For a Cu²⁺/Cu half-cell with E° = 0.34 V, [Cu²⁺] = 0.01 M at 298 K, n = 2, find the actual potential.

Solution Steps:

  1. 1Apply Nernst equation: E = E° − (RT/nF) × ln(1/[Cu²⁺])
  2. 2At 25°C: E = 0.34 − (0.0592/2) × log₁₀(1/0.01)
  3. 3Calculate: E = 0.34 − 0.0296 × log₁₀(100) = 0.34 − 0.0296 × 2
  4. 4E = 0.34 − 0.0592 = 0.2808 V

Result:

E = 0.2808 V (lower than standard due to lower Cu²⁺ concentration)

Determining Spontaneity

Problem:

E°anode = 0.00 V (SHE), E°cathode = 1.36 V (Cl₂/Cl⁻), n = 2. Is the reaction spontaneous?

Solution Steps:

  1. 1Calculate: E°cell = 1.36 − 0.00 = 1.36 V
  2. 2Positive E°cell → spontaneous reaction
  3. 3ΔG° = −2 × 96,485 × 1.36 = −262.4 kJ/mol (negative → spontaneous)
  4. 4K = exp(109.1) ≈ 3.4 × 10⁴⁷ (extremely product-favored)

Result:

E°cell = 1.36 V, spontaneous, ΔG° = −262.4 kJ/mol, K ≈ 3.4 × 10⁴⁷

Tips & Best Practices

  • ✓Use the simplified Nernst factor (0.0592/n V at 25°C) for quick calculations without the full RT/nF expression.
  • ✓Remember that E° is an intensive property—multiplying a half-reaction by a coefficient does not change E°.
  • ✓When combining half-reactions, always use reduction potentials and apply E°cell = E°cathode − E°anode.
  • ✓A large positive E°cell guarantees a large K and negative ΔG°—all three must be thermodynamically consistent.
  • ✓For non-standard conditions, measure or estimate concentrations before applying the Nernst equation.
  • ✓Temperature changes affect electrode potentials through the RT/nF term—higher T amplifies concentration effects.

Related Calculators

Redox Calculator

Balance redox reactions and find cell potential

Nernst Equation Calculator

Calculate non-standard potentials

Cell EMF Calculator

Determine cell electromotive force

Gibbs Free Energy Calculator

Calculate thermodynamic spontaneity

Frequently Asked Questions

A negative E° value means the species is a stronger reducing agent than H₂—it tends to lose electrons rather than gain them. For example, Zn²⁺/Zn has E° = −0.76 V, meaning zinc metal readily oxidizes to Zn²⁺ while reducing H⁺ to H₂. The more negative the E°, the stronger the reducing tendency of the reduced form.
According to the Nernst equation, increasing the concentration of the oxidized form increases the potential (makes reduction more favorable), while increasing the reduced form decreases the potential. At 25°C, a tenfold change in concentration ratio shifts the potential by 0.0592/n V, where n is the number of electrons transferred.
The Nernst equation calculates electrode potentials under non-standard conditions, accounting for concentration and temperature effects. It is used to predict battery voltages at different states of charge, determine ion concentrations from measured potentials (potentiometry), and calculate equilibrium positions for electrochemical reactions under real conditions.
The SHE is used as the universal reference because its potential is defined as exactly 0.00 V by convention. It involves the simple, reversible reaction 2H⁺ + 2e⁻ ⇌ H₂, which can be precisely reproduced under standard conditions. All other electrode potentials are measured relative to SHE, creating a consistent scale for comparing different redox couples.
No. Electrode potential is a thermodynamic quantity that predicts whether a reaction can occur spontaneously and how far it will proceed, but not how fast it will occur. Reaction rate depends on kinetic factors like activation energy, overpotential, and electrode surface area. A reaction with a large positive E° may still be very slow if it has a high activation energy barrier.

Sources & References

Last updated: 2026-06-06

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Editorial Note

MyCalcBuddy Editorial Team

This page is maintained as an educational calculator reference.

Source

Formula Source: Chemistry: The Central Science

by Brown, LeMay, Bursten

UpdatedLast reviewed: May 2026
CheckedFormula checks are based on standard references and internal QA review.