Activation Energy

Activation energy is the minimum energy that colliding molecules must have for a chemical reaction to occur. It acts like a barrier between reactants and products, and it explains why many reactions that release energy still do not happen quickly.

To react, molecules must pass through a high-energy arrangement called the transition state, in which bonds are partly broken and partly formed. The energy needed to reach it from the reactants is the activation energy, usually written E_a and measured in kilojoules per mole. A reaction can release energy overall, as measured by the change in Gibbs free energy, and still be slow if this barrier is high.

The Swedish chemist Svante Arrhenius showed in 1889 that the rate constant is proportional to e^−*E*_a/RT. Warming the mixture increases the fraction of collisions that have enough energy, as set by the Boltzmann constant and the thermal distribution, so a rise of about 10 degrees often doubles the rate of a reaction with a typical barrier. This explains why food spoils faster when warm and why enzymes work best near body temperature.

A catalyst provides a pathway with a lower activation energy without changing the overall energy change. Enzymes do this by binding the transition state tightly in the active site, holding reactants in the right orientation and sometimes providing acidic or basic groups. A reduction of a few tens of kilojoules per mole can speed a reaction by factors of millions. See also enthalpy.