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The standard redox potential, often denoted E°, is the equilibrium potential of a redox couple under standard conditions. It is defined relative to the standard hydrogen electrode (SHE), which is assigned a potential of 0.00 volts. Standard conditions usually mean a temperature of 25°C (298 K), activities of all solutes equal to 1 (approximately 1 M for aqueous species), and a pressure of 1 atmosphere for any gases involved. E° is typically reported as a reduction potential, indicating the tendency of a species to gain electrons.

E° values are measured for half-reactions and are used to compare the driving force of different redox couples. A higher E° means a stronger tendency to be reduced (to gain electrons), while a lower E° indicates a tendency to be oxidized. The overall cell potential, E°cell, is obtained from the difference between the cathode and anode standard potentials: E°cell = E°cathode − E°anode. Spontaneity under standard conditions corresponds to a positive E°cell and a negative standard Gibbs free energy change, ΔG° = −nF E°cell, where n is the number of electrons transferred and F is the Faraday constant.

The Nernst equation relates E° to the actual potential under non-standard conditions: E = E° − (RT/nF) ln Q, where Q is the reaction quotient. At 25°C, this can be written as E = E° − (0.05916/n) log10 Q. The standard potential is temperature-dependent and can be affected by system pH, especially for redox couples that involve protons. In biological and environmental contexts, potentials are sometimes reported relative to pH-dependent scales (for example, E°′ at pH 7) rather than the strict standard state.

Standard redox potentials are tabulated for many common couples and underpin predictions of reaction direction, battery chemistry, corrosion processes, and metabolic electron transfer.