Explain the concept of activation energy and how a catalyst modifies the energy profile of a reaction.

• A. Activation energy is the energy barrier to reaction; a catalyst provides an alternative path with lower Ea, increasing rate without changing overall ΔH.
• B. Activation energy is the total energy released by the reaction.
• C. Activation energy is the rate constant for the reaction; catalysts change the rate constant.
• D. Activation energy is the energy absorbed from the surroundings.

Answer: (C)

Activation energy is the energy barrier that must be overcome for reactants to reach the transition state and be converted into products. It corresponds to the difference in energy between the reactants and the highest point on the reaction path. A catalyst provides an alternative mechanism with a lower-energy transition state, so the energy peak on the reaction profile is reduced. Because of this lower barrier, more molecules have enough energy at a given temperature to reach the transition state, which speeds up the reaction.

Importantly, the catalyst does not change the overall enthalpy change (the energy difference between products and reactants) of the process. The amount of energy released or absorbed remains the same. The rate at which the reaction proceeds is governed by the rate constant, which in the Arrhenius picture depends on the activation energy: lowering Ea increases the rate constant, making the reaction faster. Activation energy is thus the barrier to reaction, not the rate constant itself.