How does a catalyst affect activation energy?

A catalyst significantly lowers the activation energy of a chemical reaction. The activation energy is the minimum amount of energy required for reactants to convert into products during a chemical reaction. By providing an alternative pathway with a lower energy barrier, a catalyst increases the rate at which the reaction occurs without altering the overall reaction or its products.

This lowering of activation energy means that more reactant molecules can overcome this energy barrier at a given temperature, thus leading to a higher frequency of successful collisions that produce the products. Importantly, a catalyst does not change the enthalpy or the thermodynamics of the reaction; it simply facilitates a quicker reaction by making it easier for the reactants to transition to the products.

In certain cases, it may appear that some options could be reasonable, but they do not accurately reflect the fundamental role of a catalyst in chemical kinetics. Thus, the correct understanding of a catalyst’s function is essential for grasping how reaction rates are affected in chemical processes.