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A catalyst is a substance which increases the speed of a reaction without itself undergoing any chemical change.

a) To understand the concept of Activation energy, first, we need to know about threshold energy.

The minimum amount of energy which the colliding molecules must possess as to make the chemical reaction to occur is known as threshold energy.

The minimum amount of energy required by the reactant molecules to participate in a reaction is called activation energy.

Activation energy = threshold energy - the average kinetic energy of the reactants

Note: Less is the activation energy, faster is the reaction or greater is the activation energy, slower is the reaction.

The function of a catalyst is to lower down the activation energy. The greater the decrease in the activation energy caused by the catalyst, higher will be the reaction rate. In the presence of a catalyst, the reaction follows a path of lower activation energy. Under this condition, a large number of reacting molecules are able to cross over the energy barrier and thus the rate of reaction increases.

b) The catalyst does not affect the value of ΔG (as the free energy of the reactants and products remain the same)

Note: Gibbs free energy is that thermodynamic quantity of a system, the decrease in whose value during a process is equal to useful work done by the system.