If the activation energy of a reaction is lowered by a catalyst, what happens to the rate of reaction?

• A. Rate decreases
• B. Rate increases
• C. Rate remains unchanged
• D. Temperature decreases

Answer: (B)

Lowering the activation energy makes the energy barrier for the reaction smaller, so more collision events at a given temperature have enough energy to get over it. This increases the rate constant (k) because, from the Arrhenius relationship, a smaller Ea raises e^{-Ea/RT}. Since the reaction rate depends on this rate constant and the reactant concentrations, the overall rate increases. A catalyst does this by providing an alternative, lower-energy pathway and is not consumed in the reaction. It also doesn’t change the temperature or the final equilibrium position; it simply speeds up how quickly equilibrium is reached by increasing the forward (and, for a reversible reaction, the reverse) rates.