Gibbs Free Energy: Temperature Dependence?

The chemical reaction that causes iron to rust in air is given by

4Fe + 3O? → 2Fe?O?

in which

ΔHrxn? =-1648.4 kJ/mol

ΔSrxn? = -543.7 J/mol*K

What is the standard Gibbs free energy for this reaction? In kJ/mol

What is the Gibbs free energy for this reaction at 3200 K? Assume that ΔH and ΔS do not change with temperature. In kJ/mol

At what temperature T_eq do the forward and reverse rusting reactions occur in equilibrium? In Kelvin


dG = dH-TdS

dG = -1648.4kJ/mol - (298K)(-.5437kJ/moleK) = -1486.4kJ/mole

dG at 3200K = -1648.4kJ/mole - (3200K)(-.5437kJ/moleK) = 1739.8kJ/mole (non-spontaneous)

at equilibrium dG = 0 so

-dH = -TdS

-1648.4 kJ/mole= T(-.5437kJ/moleK)

T = 1648.7 / 0.5437K = 3032K