Nernst Equation

A cell's potential falls as its reactants are used up: E = E° − (RT ÷ nF)·ln Q, with Q the reaction quotient, n the electrons transferred and F Faraday's constant. In base-10 form the factor 2.303RT ÷ F is the famous 0.0592 V per decade — but only at 25 °C. At 60 °C it is 0.0661, and the page recomputes it for the temperature you give rather than repeating the textbook number. Every tenfold change in Q moves the potential by that slope divided by n.

The quotient is products over reactants, each raised to its stoichiometric coefficient, with pure solids, pure liquids and the solvent left out at unit activity; gases use partial pressures in bar. Concentrations stand in for activities, which is good to a few percent in dilute solution and poor above about 0.1 M. At equilibrium E = 0 and Q = K, which gives K = exp(nFE° ÷ RT) from the standard potential alone. The free energy follows as ΔG = −nFE: negative means the reaction as written runs, and its magnitude is the electrical work the cell can do.

The same equation with one ion crossing a membrane gives E = (RT ÷ zF)·ln(c_out ÷ c_in), the voltage at which an ion's electrical and chemical driving forces balance — about −91 mV for a 4 / 140 mM potassium ratio at body temperature, and the basis of every pH and ion-selective electrode. A real membrane potential is a weighted mixture of several ions (Goldman–Hodgkin–Katz) and is not computed here. Standard potentials come from your own table; none is stored, and they shift with temperature themselves in ways this equation does not model. Nothing leaves the browser; the same four anonymous usage counts as the rest of the site apply.