The Nernst equation relates the electrode potential (E) of an electrochemical cell to the concentrations (or activities) of the reactants and products involved in the redox reaction occurring at the electrode. The equation is named after the German physical chemist Walther Nernst. The general form of the Nernst equation is as follows:

E = E° – (RT / nF) * ln(Q)

Where:

• E is the electrode potential (in volts)
• E° is the standard electrode potential (in volts), which is the potential of the electrode under standard conditions (usually at 25°C, 1 atm pressure, and 1 M concentration of all species)
• R is the ideal gas constant (8.314 J/(mol·K) or 0.0592 V/(mol·K))
• T is the temperature in Kelvin
• n is the number of moles of electrons transferred in the balanced redox equation
• F is Faraday’s constant (96,485 C/mol)
• ln(Q) is the natural logarithm of the reaction quotient, Q, which is the ratio of the concentrations (or activities) of the products to the reactants, each raised to their respective stoichiometric coefficients.

The Nernst equation allows the calculation of the electrode potential under non-standard conditions, where the concentrations of the species involved in the redox reaction are not at their standard state values. By considering the concentrations of the reactants and products, it provides a more accurate representation of the actual electrode potential in a real electrochemical cell.

It’s important to note that in some cases, the Nernst equation may be modified depending on the specific redox reaction and the type of electrode (e.g., for different types of electrodes like metal-metal ion electrodes or gas electrodes). These modifications take into account specific factors and considerations associated with those electrode types.