The Activation Energy

Definition

Activation energy, denoted by Ea, is defined as the minimum amount of energy necessary for a chemical reaction to occur.

Every chemical reaction requires a certain minimum energy to break the bonds in the reactants, which is provided by activation energy.
It is the energy barrier that must be overcome for reactants to be converted to products.
Not all collisions between reactant molecules result in a reaction, the colliding molecules must also possess energy equal to or greater than the activation energy.

The higher the activation energy, the slower the chemical reaction.
A low activation energy means that more molecules will have the required energy, hence the reaction will occur faster.
Lowering the activation energy increases the rate of reaction, which is what catalysts do.

The activation energy can be overcome by heating the reactants. This gives the particles more energy, so they move faster and collide more often and with greater energy.
When the temperature is increased, the number of molecules that possess the required activation energy also increases. Therefore, the rate of reaction increases.

Catalysts are substances that decrease the activation energy required for a reaction, thereby increasing the reaction rate.
Catalysts achieve this by providing an alternative reaction pathway with a lower activation energy.
Importantly, catalysts are not used up in the reaction. They remain chemically unchanged after the reaction.

The activation energy can be represented on an energy profile diagram.
On these diagrams, the activation energy is the difference in energy between the reactants and the highest point on the curve (the transition state).
Exothermic reactions result in a net release of energy, so the products are at a lower energy level than the reactants.
In endothermic reactions, more energy is required to break the bonds in the reactants than is released by forming the products, so the energy of the products is higher than that of the reactants.