Redox Reaction Calculator

In an electrochemical cell, the cathode is where reduction occurs (gain of electrons) and the anode is where oxidation occurs (loss of electrons). Both half-reactions are tabulated as reduction potentials measured against the standard hydrogen electrode.

The cell potential under standard conditions is: E_cell = E_cathode - E_anode. Both values are taken as reduction potentials. The minus sign accounts for the fact that the anode reaction runs in reverse (as an oxidation). If E_cell > 0, the reaction is spontaneous under standard conditions.

For the classic copper-zinc (Daniell) cell: copper cathode E = +0.34 V, zinc anode E = -0.76 V. Cell potential = 0.34 - (-0.76) = 1.10 V. This is the familiar battery voltage for a Zn/Cu cell.

The relationship to Gibbs free energy is: delta_G = -n * F * E_cell, where n is the number of electrons transferred and F is Faraday’s constant (96485 C/mol). A positive E_cell gives a negative delta_G, confirming spontaneity. delta_G is in joules; divide by 1000 for kilojoules.

The equilibrium constant K at temperature T is: K = exp(n * F * E_cell / (R * T)), where R = 8.314 J/(mol*K). At standard temperature (298.15 K), this simplifies to log(K) = n * E_cell / 0.05916 (in base 10). A large positive E_cell drives K toward very large values, meaning the reaction goes nearly to completion.

The Nernst equation adjusts E for non-standard concentrations: E = E_cell - (RT / (nF)) * ln(Q), where Q is the reaction quotient. At equilibrium, Q = K and E = 0.

Reduction potentials are intrinsic to the half-reaction, not the cell. You can always reverse a half-reaction to make it an oxidation, but you must flip the sign of its potential when you do.