Standard electrode potential (data page)

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

• Temperature 298.15 K (25.00 °C; 77.00 °F);
• Effective concentration (activity) 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).

Variations from these ideal conditions affect measured voltage via the Nernst equation.

Electrode potentials of successive elementary half-reactions cannot be directly added. However, the corresponding Gibbs free energy changes (∆G°) must satisfy

∆G° = –zFE°,

where z electrons are transferred, and the Faraday constant F is the conversion factor describing Coulombs transferred per mole electrons. Those Gibbs free energies can be added.

For example, from Fe²⁺ + 2 e⁻ ⇌ Fe(s) (–0.44 V), the energy to form one neutral atom of Fe(s) from one Fe²⁺ ion and two electrons is 2 × 0.44 eV = 0.88 eV, or 84 907 J/(mol  e⁻). That value is also the standard formation energy (∆G_f°) for an Fe²⁺ ion, since  e⁻ and Fe(s) both have zero formation energy.

Data from different sources may cause table inconsistencies. For example: $${\displaystyle {\begin{alignedat}{4}&{\ce {Cu+ + e-}}&{}\rightleftharpoons {}&{\ce {Cu(s)}}&\quad \quad E_{1}=+0.520{\text{ V}}\\&{\ce {Cu^2+ + 2e-}}&{}\rightleftharpoons {}&{\ce {Cu(s)}}&\quad \quad E_{2}=+0.337{\text{ V}}\\&{\ce {Cu^2+ + e-}}&{}\rightleftharpoons {}&{\ce {Cu+}}&\quad \quad E_{3}=+0.159{\text{ V}}\end{alignedat}}}$$ From additivity of Gibbs energies, one must have $${\displaystyle 2\cdot E_{2}=1\cdot E_{1}+1\cdot E_{3}}$$ But that equation does not hold exactly with the cited values.