A catalyst is a substance that increases a reaction rate without being used up by the reaction. It does this by lowering the activation energy for the reaction, thus allowing the reaction to proceed at a lower temperature. (See the help page on the Bolztmann factor for more details.) In the diagram below, the activation energy E_a for the reaction is much higher than the activation energy for the reaction with a catalyst, E_cat

There are two basic kinds of catalysts:

• Hetereogenous
• Homogenous

Heterogeneous catalysts exist in a different phase than the reactants. These are quite commonly metals: a classic example is hydrogen, oxygen and platinum. The reaction 2H₂ + O₂ ->2H₂O is highly exothermic, but the activation energy is too high for the reaction to proceed at room temperature. This is why you can have a mixture of hydrogen and oxygen and not have an instant reaction

However, adding a small amount of platinum metal to the mixture lowers the activation energy of the reaction enough for it to proceed at room temperature. The platinum is not used up in the reaction, but as soon as the hydrogen contacts the platinum the reaction proceeds rapidly.

A homogeneous catalyst is one that is in the same phase as the reactants. Often these are species that are both created and destroyed during a multiple step reaction, so that there is no overall change in the amount of the species.

For example, the reduction of Ce⁺⁴ by Tl⁺ is very slow, however, adding some silver ion results in a great speed up of the reaction. It is a three step process:

1. Ag⁺ + Ce⁺⁴ -> Ag⁺² + Ce⁺³+
2. Tl⁺ + Ag⁺² -> Tl⁺² + Ag⁺+
3. Tl⁺² + Ce⁺⁴ -> Tl⁺³ + Ce⁺³+

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4. Tl⁺ + 2Ce⁺⁴ -> Tl⁺³ + 2Ce⁺³

Information provided by: http://learn.chem.vt.edu