Activation Energy Definition in Chemistry

Activation energy is the minimum amount of energy required to initiate a reaction. It is the height of the potential energy barrier between the potential energy minima of the reactants and products. Activation energy is denoted by E_a and typically has units of kilojoules per mole (kJ/mol) or kilocalories per mole (kcal/mol). The term "activation energy" was introduced by the Swedish scientist Svante Arrhenius in 1889. The Arrhenius equation relates activation energy to the rate at which a chemical reaction proceeds:

k = Ae^-Ea/(RT)

where k is the reaction rate coefficient, A is the frequency factor for the reaction, e is the irrational number (approximately equal to 2.718), E_a is the activation energy, R is the universal gas constant, and T is the absolute temperature (Kelvin).

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How to Calculate Activation Energy

By Todd Helmenstine

From the Arrhenius equation, it can be seen that the rate of reaction changes according to temperature. Normally, this means a chemical reaction proceeds more quickly at a higher temperature. There are, however, a few cases of "negative activation energy", where the rate of a reaction decreases with temperature.

Why Is Activation Energy Needed?

If you mix together two chemicals, only a small number of collisions will naturally occur between the reactant molecules to make products. This is particularly true if the molecules have low kinetic energy. So, before a significant fraction of reactants can be converted into products, the free energy of the system must be overcome. The activation energy gives the reaction that little extra push needed to get going. Even exothermic reactions require activation energy to get started. For example, a stack of wood won't start burning on its own. A lit match can provide the activation energy to start combustion. Once the chemical reaction starts, the heat released by the reaction provides the activation energy to convert more reactant into product.

Sometimes a chemical reaction proceeds without adding any additional energy. In this case, the activation energy of the reaction is usually supplied by heat from the ambient temperature. Heat increases the motion of the reactant molecules, improving their odds of colliding with each other and increasing the force of the collisions. The combination makes it more likely bonds between reactant will break, allowing for the formation of products.

Catalysts and Activation Energy

A substance that lowers the activation energy of a chemical reaction is called a catalyst. Basically, a catalyst acts by modifying the transition state of a reaction. Catalysts are not consumed by the chemical reaction and they don't change the equilibrium constant of the reaction.

Relationship Between Activation Energy and Gibbs Energy

Activation energy is a term in the Arrhenius equation used to calculate the energy needed to overcome the transition state from reactants to products. The Eyring equation is another relation that describes the rate of reaction, except instead of using activation energy, it includes Gibbs energy of the transition state. The Gibbs energy of the transition state factors in both enthalpy and entropy of a reaction. Activation energy and Gibbs energy are related, but not interchangeable.