Thermodynamics of Redox Reactions

The thermodynamic criterion for determining the spontaneity of a reaction is ∆G, Gibbs free energy,
the maximum amount of useful work produced by a chemical reaction. In an electrochemical cell,
the work done is dependent on the number of coulombs and the energy available. Thus, ∆G and
EMF are related as follows:

∆G  =   −nFEcell

where n is the number of moles of electrons exchanged, F is Faraday’s constant, and Ecell is the EMF
of the cell. Keep in mind that if Faraday’s constant is expressed in coulombs (J/V), then ∆G must be
expressed in J, not kJ.

If the reaction takes place under standard conditions, then the ∆G is the standard Gibbs free energy
and Ecell is the standard cell potential. The above equation then becomes:

∆G° =   −nFE°cell

Recall that in the chapter on thermochemistry we derived the following equation:

∆G° =   −RT ln  Keq

where R is the gas constant 8.314 J/(K•mol), T is the temperature in K, and Keq is the equilibrium
constant for the reaction. Combining this with the equation above, we get:

∆G° =   −nFE°cell   =   −RT ln  Keq

or simply: