The collision theory states that molecules must collide before they react.

Only those molecules having energy above a certain minimum value will react when they collide. This minimum energy value is called the activation energy. The activation energy can be thought of as an energy barrier to the progress of the reaction.

Energy-profile diagrams:

The existence of an energy barrier explains why many energetically favourable reactions are not spontaneous at room temperature; there may be no (or very few) molecules with energy greater than the activation energy. The rate of reaction is directly related to the size of the activation energy.

The spread of energies in a sample of gas molecules is given by the Maxwell-Boltzmann distribution curve. Frequent collisions between molecules results in a spread of high and low energy molecules.

At the lower temperature T1, the number of molecules with energy greater than activation energy is given by the purple shaded area. An increase in temperature to T2 results in a large increase in the number of such molecules.

This increase is represented by the light blue shaded area. Therefore, the number of 'productive collisions' will be much greater at higher temperatures and so the rate will not be much faster.

This is a more sophisticated theory as it explains the nature of the energy barrier. It proposes that an immediate 'activated complex' or transition state is formed during the reaction.

This is a high energy species in which old bonds are partially broken and new bonds are partially formed.