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How does a catalyst increase the rate of a chemical reaction?

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A graph showing a potential energy curve for a chemical reaction, with the y-axis labeled potential energy and the x-axis labeled progress of reaction. The curve rises from the initial state (reactants) to a peak labeled transition state (activated complex), indicating activation energy, then drops to the final state (products). — chemical kinetics In the Arrhenius equation, the activation energy (E) represents the minimum amount of energy required to transform reactants into products in a chemical reaction. On a potential energy curve, the value of the activation energy is equivalent to the difference in potential energy between particles in an intermediate configuration (known as the activated complex, or transition state) and particles of reactants in their initial state. The activation energy thus can be visualized as a barrier that must be overcome by reactants before products can be formed.

A catalyst increases the rate of a chemical reaction by lowering the activation energy—the minimum amount of energy required for a reaction to occur. By reducing this energy barrier, the catalyst makes it easier for reactant molecules to reach the transition state and transform into products.123

Catalysts achieve this by facilitating a specific reaction mechanism, often by forming unstable intermediate compounds with the reactants. Although the catalyst may chemically combine with reactants during the process, it is ultimately regenerated, meaning it is not consumed and remains unchanged in concentration. Because they are not consumed, a small amount of catalyst can facilitate the transformation of a large number of reactant molecules.421

It is important to note that while a catalyst accelerates the rate at which a reaction reaches chemical equilibrium, it does not change the position of that equilibrium. Furthermore, a catalyst accelerates the reaction in both the forward and reverse directions equally.12