Catalysts enable reactions to occur much faster or at lower temperatures because of changes that they induce in the reactants. Catalysts provide an alternative pathway with a lower activation energy, for a reaction to proceed. Molecules that would not have had the energy to react or that have such low energies that it is likely that they would take a long time to do so are able to react in the presence of a catalyst. A catalyst reduces the energy required for the reaction to occur.
The two main categories of catalysts are heterogeneous and homogeneous catalysts. Heterogeneous catalysts are present in different phases from the reactants in the reaction they are catalysing, whereas homogenous catalysts are in the same phase. A simple model for heterogeneous catalysis involves the catalyst providing a surface on which the reactants (or substrates) temporarily become adsorbed. Bonds in the substrate become weakened sufficiently for new products to be created. The bonds between the products and the catalyst are weaker, so the products are released.
Homogenous catalysts generally react with one or more reactants to form a chemical intermediate that subsequently reacts to form the final reaction product, in the process regenerating the catalyst. The following is a typical catalytic reaction scheme, where C represents the catalyst:
Use of "catalyst" in a broader cultural sense is in rough analogy to the sense described here.
Fig. 1. Enthalpy profile for catalysed and uncatalysed reactions. AU is the activation energy for an uncatalysed reaction, AC is the reduced activation energy for the same reaction when catalysed. I represents the point at which a chemical intermediate has been formed, which then reacts to form the products.