Catalyst on Chemical Equilibrium

The function of a catalyst can be visualized using a reaction energy profile diagram. The catalyst lowers the energy of the transition state, which is the peak of the energy barrier between reactants and products. Crucially, it lowers this barrier by the same amount for both the forward and reverse reactions. This means the activation energy for the forward reaction ([latex]E_{a,fwd}[/latex]) and the activation energy for the reverse reaction ([latex]E_{a,rev}[/latex]) are both reduced, but the difference between them, which corresponds to the overall enthalpy change of the reaction ([latex]\Delta H = E_{a,fwd} – E_{a,rev}[/latex]), is not affected.

Since the equilibrium constant K is related to the standard Gibbs free energy change ([latex]\Delta G^{circ} = -RT \ln K[/latex]), and catalysts do not alter the thermodynamics ([latex]\Delta G^{circ}[/latex], [latex]\Delta H^{circ}[/latex]) of the overall reaction, they cannot change the value of K. Their role is purely kinetic, enabling the system to achieve the thermodynamically favorable equilibrium state in a shorter amount of time. This is vital in industrial processes where reaching equilibrium quickly is economically necessary, as many thermodynamically favorable reactions are too slow to be practical without catalysis.