What is the effect of increasing activation energy on a chemical reaction?

Increasing the activation energy of a chemical reaction slows down the reaction. Activation energy is the minimum amount of energy required for reactants to undergo a chemical transformation into products. When the activation energy is set higher, fewer molecules have sufficient energy to overcome this barrier when they collide. As a result, the frequency of successful collisions that result in reactions decreases, leading to a slower overall reaction rate.

This principle is significant in understanding reaction kinetics, as it directly links temperature, energy, and rate. Lowering the activation energy—through catalysts, for instance—would increase the probability that reactants have enough energy to react, thereby speeding up the reaction. However, when activation energy increases, the opposite effect occurs, leading to a decrease in the reaction rate.