Which statement about activation energy in catalyzed reactions is true?

Which statement about activation energy in catalyzed reactions is true?

• A. It is always zero
• B. It is always lower than the uncatalyzed case
• C. It is always higher than the uncatalyzed case
• D. It cannot be determined

Answer: (B)

Activation energy is the energy hill reactants must climb to reach the transition state. A catalyst provides an alternative pathway with a different, lower-energy transition state or set of steps, often via stable intermediates. Because the barrier is lower, the rate constant increases (via the Arrhenius relationship), so the reaction proceeds faster at the same temperature. The catalyst doesn’t change the overall energy change of the reaction or its equilibrium; it only changes the pathway and lowers the barrier. So, the activation energy for the catalyzed route is lower than for the uncatalyzed route. It’s not zero in general, and it’s not higher than the uncatalyzed case.

Activation energy is the energy hill reactants must climb to reach the transition state. A catalyst provides an alternative pathway with a different, lower-energy transition state or set of steps, often via stable intermediates. Because the barrier is lower, the rate constant increases (via the Arrhenius relationship), so the reaction proceeds faster at the same temperature. The catalyst doesn’t change the overall energy change of the reaction or its equilibrium; it only changes the pathway and lowers the barrier. So, the activation energy for the catalyzed route is lower than for the uncatalyzed route. It’s not zero in general, and it’s not higher than the uncatalyzed case.