Final answer:
A catalyst lowers the activation energy required for a reaction, thereby increasing the rate without changing the equilibrium constant or the free energy of the products.
Step-by-step explanation:
The effect of a catalyst on a reaction is to lower the energy of activation. This doesn't change the free energy of the products or make the reactants less stable, nor does it change the equilibrium constant of the reaction. What a catalyst does is provide an alternative reaction pathway with a lower activation energy. This makes it easier for reactant molecules to have effective collisions, thus increasing the reaction rate without affecting the overall energy balance of the reactants and products.
A catalyst essentially reduces the energy barrier that reactants need to overcome to convert into products. As a consequence, the rate at which the reaction approaches equilibrium is faster, but the position of equilibrium, represented by the equilibrium constant, remains unchanged. Because the catalyst offers a new mechanism with a lowered energy transition state, more reactant molecules can successfully interact and form products under given conditions.