Activation energy

Activation Energy is a fundamental concept in the field of chemical kinetics and physical chemistry. It refers to the minimum amount of energy that must be supplied to a chemical reaction for the reaction to proceed. The concept was first proposed by the Swedish scientist Svante Arrhenius in the late 19th century.

Definition

The activation energy (Ea) of a reaction is defined as the difference in energy between the reactants and the transition state. The transition state, also known as the activated complex, is the highest-energy state of the system during the reaction. It is a transient state that exists for a very short time before the reactants are converted into products.

Determination

The activation energy of a reaction can be determined experimentally by measuring the rate of the reaction at different temperatures and applying the Arrhenius equation. This equation, named after its developer, Svante Arrhenius, relates the rate constant of a reaction to the temperature and the activation energy.

Role in Reactions

The activation energy plays a crucial role in determining the rate of a chemical reaction. A reaction with a high activation energy will proceed slowly, as only a small fraction of the reactant molecules will have enough energy to reach the transition state. Conversely, a reaction with a low activation energy will proceed quickly, as a larger fraction of the reactant molecules will be able to reach the transition state.

Catalysts and Activation Energy

Catalysts are substances that can lower the activation energy of a reaction, thereby increasing the reaction rate. They achieve this by providing an alternative reaction pathway with a lower activation energy. Importantly, catalysts do not change the overall energy change of the reaction; they only make the reaction proceed faster.

See Also

• Arrhenius equation
• Chemical kinetics
• Catalysis
• Transition state theory