What effect do catalysts have on the activation energy of chemical reactions?

Catalysts play a crucial role in chemical reactions by facilitating the process without being consumed in the reaction. The primary effect of catalysts is to lower the activation energy required for a reaction to occur. Activation energy is the minimum energy needed for reactants to undergo a transformation into products. By decreasing this energy barrier, catalysts enable more reactant molecules to collide with enough energy to reach the transition state, thereby increasing the rate of the reaction.

Lowering the activation energy means that the reaction can proceed more easily and quickly, which is essential in many biological processes and industrial applications. This characteristic is vital because reactions that would normally occur very slowly or not at all under certain conditions can occur more readily in the presence of a catalyst.

The other options do not accurately describe the role of catalysts. They do not increase activation energy, do not leave activation energy unchanged, and do not alter the chemical products of the reaction; rather, they allow the reactants to be converted to the same products more efficiently.