• The?standard free energy change?can be calculated using the standard cell potential of an electrochemical cell

ΔG^??= - n x?E_cell^??x?F

ΔG^??= standard Gibbs free energy

n?= number of electrons transferred in the reaction

E_cell^??= standard cell potential (V)

F?= Faraday constant (96 500 C mol^-1)

Worked Example: Calculating the standard Gibbs free energy change

Answer

• Step 1:?Determine the two half-equations and their?E^??using the Data booklet

Fe³⁺?(aq) + e^-?? Fe²⁺?(aq)? ? ? ??E^??= +0.77 V

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

• Step 2 :?Calculate the?E_cell^?

E_cell^??=?E_red^??-?E_ox^?

= (+0.77) - (+0.34)

= +0.43 V

• Step 3:?Determine the number of electrons transferred in the reaction

The Cu²⁺/Cu has a smaller?E^??value which means that it gets?oxidised

It?transfers?two electrons to??two?Fe³⁺?ions

Each Fe³⁺?ion accepts one electron so the total number of electrons transferred is?two

• Step 4:?Substitute the values in for the standard Gibbs free energy equation

ΔG^??= - n x?E_cell^??x?F

= -2 x (+0.43) x 96 500

= -82 990 J mol^-1

= -83 kJ mol^-1