If a catalyst lowers the activation energy of a reaction, which of the following is most likely to occur?

• A. The equilibrium constant of the reaction changes.
• B. The reaction becomes non-spontaneous.
• C. The products of the reaction are altered.
• D. The reaction rate increases.

Answer: (D)

Verified step by step guidance

1. Understand that a catalyst works by providing an alternative reaction pathway with a lower activation energy, which affects the kinetics of the reaction but not the thermodynamics.
2. Recall that the activation energy is the minimum energy required for reactants to transform into products, and lowering it increases the number of effective collisions per unit time.
3. Recognize that the equilibrium constant, which depends on the free energy difference between reactants and products, remains unchanged because a catalyst does not alter the overall energy change of the reaction.
4. Note that the spontaneity of the reaction is determined by the Gibbs free energy change (\$\Delta G\$), which is unaffected by the catalyst, so the reaction does not become non-spontaneous.
5. Conclude that since the catalyst does not change the products formed, the main effect is an increase in the reaction rate due to the lowered activation energy.