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Catalyst are substances which increases the rate of the reaction without getting used up by themselves while decreasing the activation energy of the reaction. Activation energy is the energy difference between the initial reagents and the transition state.

• Catalyst increases the rate of reaction.
• Catalyst accelerates the rate of a reaction by decreasing the activation energy.
• Catalysts may provide different reaction mechanisms involving extra steps in the reaction and lower the activation energy simultaneously increasing the reaction rate.
• Catalyst is regenerated during the reaction.

We know that, activation energy is the energy difference between the initial reagents and the transition state (it is the state where the energy is maximum on the reaction coordinate diagram). In the reaction (a), the reagents are at 35 kJ and the transition state is at 45 kJ, so the activation energy can be calculated as: \({{\rm{E}}_{\rm{a}}}{\rm{ = 45 kJ - 35 kJ = 10 kJ}}\)

Activation energy is the energy difference between the initial reagents and the transition state (it is the state where the energy is maximum on the reaction coordinate diagram).

In the reaction (a), the reagents are at 35 kJ and the transition state is at 45 kJ, so the activation energy can be calculated as:\({{\rm{E}}_{\rm{a}}}{\rm{ = 45 kJ - 35 kJ = 10 kJ}}\)Thus, the activation energy\(\left( {{E_a}} \right)\)of the reaction (a) a is 10 kJ and reaction (b) is 10 kJ.