Nernst Equation Calculator

The Nernst equation calculates the reduction potential of a half-cell or full electrochemical cell under non-standard conditions. It relates the measured cell potential to the standard electrode potential, temperature, and activities of the chemical species involved in the reaction.

Standard potential (E°) is the electrode potential measured under standard conditions (25°C, 1 M concentrations, 1 atm pressure). Cell potential (E) is the actual potential under real conditions, which varies with temperature and concentration according to the Nernst equation.

Temperature directly affects the RT/nF term in the Nernst equation. Higher temperatures increase this factor, making the potential more sensitive to concentration changes. The temperature must be in Kelvin (K = °C + 273.15).

The reaction quotient Q is the ratio of product concentrations to reactant concentrations, each raised to their stoichiometric coefficients. For a simple redox reaction, Q = [oxidized]/[reduced]. It determines how far the reaction is from equilibrium.

The number of electrons (n) is determined by balancing the half-reaction. Count the electrons on one side of the balanced equation. For example, Cu²⁺ + 2e⁻ → Cu has n = 2.

Concentrations should be in molarity (M) for the standard form of the Nernst equation. Some variations use activities instead of concentrations, but molarity is the most common unit for practical calculations.

The cell potential equals the standard potential when Q = 1, which occurs when the concentrations of oxidized and reduced species are equal (both 1 M under standard conditions). At this point, the ln(Q) term becomes zero.

A negative cell potential indicates that the reaction is not spontaneous under the given conditions. The reverse reaction would be spontaneous. For a reaction to be spontaneous, the cell potential must be positive.