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The activation energy may be defined as the extra energy which the molecules of the reactants have to absorb so that their energy becomes equal to the threshold energy.

The highest energy state in a molecular collision that leads to reaction is the transition state. Transition state is most unstable. The transition state or activated complex gets converted into product molecules. During the formation of this complex, old bonds start breaking and the new bonds start making.

It may be defined as a species that exists for a finite amount of time, although it might be short but it has some stability.

To understand the concepts of activation energy and transition state graphically, reaction energy diagram is used. The vertical axis of the diagram represents the total potential energy of all the species present in the reaction. The horizontal axis represents the reaction coordinate that gives the progress of the reaction, proceeding from reactants on the left to products on the right. The highest point on the graph is the transition state and the activation energy is the difference in energy between the reactants and the transition state.

It may be defined as the slowest step of a chemical reaction that determines the rate at which the overall reaction takes place.

For endothermic reactions, the potential energy of the products is higher than that of the reactants.

reaction-energy diagram for a two-step endothermic reaction with a rate-limiting step