Which type of reaction is more likely to require a catalyst for efficiency?

Reactions with a high activation energy are more likely to require a catalyst for efficiency because the catalyst serves to lower the activation energy barrier. Activation energy is the minimum energy required for a reaction to proceed. When this energy is high, the reaction occurs slowly because fewer molecules have enough energy to overcome this barrier.

By introducing a catalyst, the pathway for the reaction is altered, allowing a greater proportion of the reactant molecules to participate in the reaction at a given temperature. Consequently, catalysts increase the reaction rate without being consumed in the process. This is particularly important for reactions that do not occur readily under normal conditions and would benefit significantly from the increased rate provided by the catalyst.

Other types of reactions, such as those that are already fast, do not typically need a catalyst because they occur quickly without additional aid. Irreversible reactions, while they can benefit from catalysts, do not inherently require one, as their directionality is more about product formation than activation energy. Reactions with low energy requirements generally proceed efficiently without a catalyst, as the energy needed to initiate them is already easily accessible.