The effect of a catalyst on a chemical reaction is to increase the rate of that reaction by reducing its activation energy. This activation energy is the energy required to move chemical reactants into a transition state, which is intermediate between reactant and product.
Catalysts reduce the activation energy of reactions by creating intermediate products with the reactants. Each reaction that occurs with a catalyst requires less activation energy than a direct reaction between the reactants would. The end product of the reaction is the same, with the same amount of energy consumed or released, but it is easier to get the reaction to start and continue using catalysts. The equilibrium point, that is, the point at which a chemical reaction and its reverse occur at the same rate, is also unchanged by the presence of a catalyst.
To be a true catalyst, a chemical must not be directly consumed by the reaction, although it can be altered by other effects of the reaction, such as heating. While the catalyst reacts with the reactants, the end product of the reaction does not incorporate it and the catalyst returns to its original state. Thus, very little catalyst is generally needed to accelerate a reaction, even if the amounts of reactants are relatively large.