What is an effect of a catalyst on the reaction's activation energy?

What is an effect of a catalyst on the reaction's activation energy?

• A. It decreases the activation energy
• B. It increases the activation energy
• C. It has no effect on the activation energy
• D. It eliminates the need for activation energy

Answer: (A)

A catalyst functions by providing an alternative pathway for a reaction with a lower activation energy compared to the uncatalyzed pathway. Activation energy is the minimum energy required for the reactants to reach the transition state and proceed to products. When a catalyst is present, it stabilizes the transition state or brings the reactants into a more favorable orientation for reaction, thereby allowing the reaction to occur more easily. Consequently, the presence of a catalyst decreases the activation energy needed for the reaction to proceed. This reduction in activation energy does not change the overall energy difference between reactants and products but simply facilitates the reaction by allowing more reactant molecules to possess enough energy to surpass the lower activation energy barrier. By decreasing activation energy, the rate of reaction is increased, leading to faster product formation under the same conditions.

A catalyst functions by providing an alternative pathway for a reaction with a lower activation energy compared to the uncatalyzed pathway. Activation energy is the minimum energy required for the reactants to reach the transition state and proceed to products. When a catalyst is present, it stabilizes the transition state or brings the reactants into a more favorable orientation for reaction, thereby allowing the reaction to occur more easily. Consequently, the presence of a catalyst decreases the activation energy needed for the reaction to proceed.

This reduction in activation energy does not change the overall energy difference between reactants and products but simply facilitates the reaction by allowing more reactant molecules to possess enough energy to surpass the lower activation energy barrier. By decreasing activation energy, the rate of reaction is increased, leading to faster product formation under the same conditions.