Activation energy

In the Arrhenius model of reaction rates, activation energy is the minimum amount of energy that must be available to reactants for a chemical reaction to occur. The activation energy (Ea) of a reaction is measured in kilojoules per mole (kJ/mol) or kilocalories per mole (kcal/mol). Simplified:

Activation energy is the minimum energy barrier that reactant molecules must overcome to transform into products. A reaction occurs only if enough molecules have kinetic energy equal to or greater than this barrier, which usually requires sufficiently high temperature. The term "activation energy" was introduced in 1889 by the Swedish scientist Svante Arrhenius.

Other uses

Although less commonly used, activation energy also applies to nuclear reactions and various other physical phenomena.

Temperature dependence and the relation to the Arrhenius equation

The Arrhenius equation gives the quantitative basis of the relationship between the activation energy and the rate at which a reaction proceeds. From the equation, the activation energy can be found through the relation

k

=

A

e

−

E

a

/

(

R

T

)

{\displaystyle k=Ae^{{-E_{\textrm {a}}}/{(RT)}}}

where A is the pre-exponential factor for the reaction, R is the universal gas constant, T is the absolute temperature (usually in kelvins), and k is the reaction rate coefficient. Even without knowing A, Ea can be evaluated from the variation in reaction rate coefficients as a function of temperature (within the validity of the Arrhenius equation).

At a more advanced level, the net Arrhenius activation energy term from the Arrhenius equation is best regarded as an experimentally determined parameter that indicates the sensitivity of the reaction rate to temperature.