Figure 4.1.1 schematically depicts the energy changes that occur during the progression of a simple reaction. In the figure, the energy differences during the reaction are compared for a catalyzed (plot on the right) and an uncatalyzed reaction (plot on the left). Notice that the reactants start at the same energy level for both conditions and that the products end at the same energy for both as well. Thus, the difference in energy between the energy of the ending compounds and the starting compounds is the same in both cases. This is the first important rule to understand any kind of catalysis – catalysts do not change the overall energy of a reaction. Given enough time, a non-catalyzed reaction will get to the same equilibrium as a catalyzed one.
Figure 4.1.1: Energetic considerations of catalysis
Another feature to note about catalyzed reactions is the reduced energy barrier (also called the activation energy or free energy of activation) to reach the transition state of the catalyzed reaction. This is the second important point about catalyzed reactions – catalysts work by lowering activation energies of reactions and thus molecules more easily reach the energy necessary to get to the point where the reaction occurs. Note that these reactions are reversible. The extent to which they will proceed is a function of the size of the energy difference between the product and reactant states. The lower the energy of the products compared to the reactants, the larger the percentage of molecules that will be present as products at equilibrium. At equilibrium, of course, no change in concentration of reactants and products occurs because at this point, the forward and reverse reaction rates are the same.