Standard electrode potential (data page)

The data below tabulates standard electrode potentials (E°), in volts relative to the standard hydrogen electrode, at:

• Temperature 298.15 K (25.00 °C; 77.00 °F);
• Effective concentration 1 mol/L for each aqueous or amalgamated (mercury-alloyed) species;
• Unit activity for each solvent and pure solid or liquid species; and
• Absolute partial pressure 101.325 kPa (1.00000 atm; 1.01325 bar) for each gaseous reagent — the convention in most literature data but not the current standard state (100 kPa).

The Nernst equation adjusts for general concentrations, pressures, or temperatures.

Simultaneous half-reactions do not in general add voltages, but instead add Gibbs free energy change: the product of the voltage and the number of electrons transferred, typically the Faraday constant. For example, from Fe²⁺ + 2 e⁻ ⇌ Fe(s) (–0.44 V), the energy to create one neutral atom of Fe(s) from one Fe²⁺ ion and two electrons is 2 × 0.44 eV = 0.88 eV, or 84 895 J/(mol  e⁻). That value is also the standard formation energy for an Fe²⁺ ion, since  e⁻ and Fe(s) both have zero formation energy.

Data from different sources may cause table inconsistencies. For example:

$${\displaystyle {\begin{alignedat}{4}&{\ce {Cu+ + e-}}&{}\rightleftharpoons {}&{\ce {Cu}}(s)&\quad \quad E_{1}=+0.520{\text{ V}}\\&{\ce {Cu^2+ + 2e-}}&{}\rightleftharpoons {}&{\ce {Cu}}&\quad \quad E_{2}=+0.337{\text{ V}}\\&{\ce {Cu^2+ + e-}}&{}\rightleftharpoons {}&{\ce {Cu+}}&\quad \quad E_{3}=+0.159{\text{ V}}\end{alignedat}}}$$

$${\displaystyle E_{3}={\frac {2\cdot E_{2}-1\cdot E_{1}}{1}}=0.154{\text{ V,}}}$$