A catalyst speeds up a chemical reaction without being consumed. It provides an alternative pathway with lower activation energy. The catalyst may participate in the reaction, forming intermediates, but it is regenerated at the end. A small amount can catalyze a large amount of reaction. Catalysts are essential in industry, biology, and environmental protection.

Catalysts come in several types. Homogeneous catalysts are in the same phase as the reactants, usually in solution. Acid catalysts for esterification are an example. Heterogeneous catalysts are in a different phase, usually solid catalysts with gaseous or liquid reactants. Platinum in catalytic converters, iron in the Haber-Bosch process, and zeolites in petroleum cracking are examples. Enzymes are biological catalysts, highly specific and efficient. They catalyze reactions at mild temperatures and pressures.

Catalysts do not change the equilibrium position of a reaction. They only change the rate at which equilibrium is reached. They also do not change the overall energy change. They lower the activation energy, making the reaction faster at a given temperature. Some catalysts are poisoned by impurities, which bind to the active sites and block them. That is why catalytic converters require unleaded gasoline. Lead poisons the catalyst.

Catalysts are used in about 90 percent of chemical processes. They make reactions faster, cheaper, and more selective. They reduce energy consumption and waste. They are one of the most important tools in green chemistry.