IB Chemistry home > Syllabus 2016 > Energetics > Gibbs' free energy and Standard Electrode Potentials

Syllabus ref: 15.2

Although the link between Gibbs Free Energy and Standard Electrode Potentials is not specifically stated in the syllabus, both concepts are used to predict spontaneity and should be considered together.

Gibbs free energy

Gibbs free energy change is used to predict spontaneity of thermodynamic processes according to the equation:


If ΔG is negative then the process is spontaneous. Remember that this does not mean that the process will occur, simply that it can occur, it is thermodynamically feasible.


Standard electrode potentials

Standard electrode potentials can be used to predict the spontaneity of a proposed reaction:

If the difference between the E of the component being reduced and the E of the component being oxidised is positive (and greater than 0.3V) then the proposed reaction is spontaneous.

E(reaction) = E(reduced state) - E(oxidised state)

For example, in the following proposed reaction between zinc metal and copper ions in solution, the equation would be:

Zn(s) + Cu2+(aq) Zn2+(aq) + Cu(s)

Inspection shows that the copper ions are being reduced and the zinc metal is being oxidised. The half-equations for the standard electrode potentials are:

Zn2+(aq) + 2e Zn(s)E = -0.76V

Cu2+(aq) + 2e Cu(s)E = +0.34V

Applying the equation E(reaction) = E(red) - E(ox) gives E(reaction) = +0.34 - (-0.76) = +1.10V.

The value, also called the 'cell potential', is positive, therefore the reaction is spontaneous as written.


The relationship between SEP and ΔG

From the above sub-sections we can see that spontaneity results when Gibbs free energy change is negative and when the cell potential is positive.

Gibbs Free Energy Cell Potential Reaction
negative positive spontaneous
positive negative non-spontaneous

Therefore ΔG -E

The actual proportionality depends on the number of electrons transferred and the charge on one mole of electrons.

ΔG = -nFE

You are expected to be familiar with both of the concepts and it is not beyond the bounds of possibility that you could be asked to discuss them together.