The presence of a catalyst affects the activation energy and rate of a chemical reaction by providing an alternative reaction pathway with a lower activation energy. This allows the reaction to proceed more quickly, increasing the rate of the reaction.Activation energy is the minimum amount of energy required for reactants to form products in a chemical reaction. In the absence of a catalyst, the reactants must overcome a higher energy barrier to form products. When a catalyst is present, it lowers the activation energy by stabilizing the transition state or forming intermediate species that have lower energy barriers.As a result of the lower activation energy, a greater proportion of the reactant molecules have sufficient energy to undergo the reaction at a given temperature. This leads to an increase in the rate of the reaction, as more reactant molecules can effectively collide and form products in a shorter amount of time.It is important to note that catalysts do not change the overall energy change enthalpy of the reaction, nor do they get consumed in the reaction. They simply provide an alternative pathway for the reaction to occur more quickly.