By convention, redox half reactions are generally tabulated in textbooks and other reference works as reduction reactions with the oxidized form on the left side and the reduced form on the right side, but it is understood that the reaction may occur in either direction.

### Standard Electrode Potential

Given that electrochemical half reactions can occur in either direction, they are often written using chemical equilibrium notation* as follows: $O + ne^- \rightleftharpoons R$

Each half reaction has an associated standard electrode potential ( $E^{{\circ}}$) which is a thermodynamic quantity related to the free energy associated with the equilibrium. Like many other standard thermodynamic quantities, the standard electrode potential corresponds to a given standard state. The standard state corresponds to a thermodynamic system where the chemical activities of O and R are unity (i.e., when all solution concentrations are $\text{1.0 mol/L}$, all gases are present at $\text{1.0 bar}$ partial pressure, and other materials are present as pure phases with unity activity).

### Nernst Equation

To account for the very likely possibility of non-unity activities, the Nernst equation (see below) can be used to express the equilibrium electrode potential ( $E_{NERNSTIAN}$) in terms of the actual activities ( $a_O$ and $a_R$) $E_{NERNSTIAN} = E^{\circ} + \left(\frac{RT}{nF}\right)ln\left[{\frac{a_O}{a_R}}\right]$

where $F$ is the Faraday constant ( $\text{F = 96485 C/mol}$), $R$ is the ideal gas constant ( $\text{R = 8.3145 J} \; mol^{-1} \; K^{-1}$), and $T$ is the absolute temperature ( $K)$. Usually, the activities of molecules or ions dissolved in solution are assumed to be the same as their molar concentrations.