The activation energy of a chemical reaction can be determined by:

Evaluating rate constants at two different temperatures.

Activation energy is a fundamental concept in chemistry, referring to the minimum amount of energy required for a chemical reaction to occur. Without sufficient activation energy, the reaction will not proceed. The concept of activation energy was introduced by the Swedish scientist Svante Arrhenius.

Arrhenius developed an equation that relates the activation energy (E_a) to the rate constant (k) of a reaction. This equation allows for the determination of activation energy by utilizing rate constants measured at two different temperatures. By comparing the rate constants at these temperatures, it is possible to calculate the activation energy.

Activation energy is also calculated by using rate constants known at two different temperatures. Let us consider the following equation for the above purpose,

\(\left(\frac{k_2}{k_1}\right) = \frac{−E_a}{R}\left[\frac{1}{T_1} − \frac{1}{T_2}\right]\)

In the presence of a catalyst, the activation energy gets lowered because the catalyst increases the rate of the chemical reaction, slower the chemical reaction the higher will be its activation energy. The release of heat also lowers the activation energy required by the reaction.