


In order for chemicals to react, the species involved in any reaction must undergo a rearrangement of chemical bonds.

The slowest step in the bond rearrangement produces what is termed a transition state - a chemical species that is neither a reactant nor a product, but is an intermediate between the two. Energy is required to form the transition state. This energy is called the Energy of Activation, or Ea.

Reactants with energy lower than Ea cannot pass through the transition state to react and become products.

A catalyst works by providing a different route, with lower Ea, for the reaction. In any given time interval, the presence of a catalyst allows a greater proportion of the reactant species to acquire sufficient energy to pass through the transition state and become products.

Catalysts cannot shift the position of a chemical equilibrium - the forward and backward reactions are both accelerated so that the equilibrium constant Keq is unchanged. However, by removing products from the reaction mixture as they form, the overall rate of product formation can in practice be increased.

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