What is the effect of a catalyst on the activation energy of a reaction?

A catalyst plays a crucial role in chemical reactions by lowering the activation energy required for the reaction to proceed. Activation energy is the minimum energy required for reactants to collide and form products. By providing an alternative reaction pathway, a catalyst enables more molecules to have sufficient energy to overcome this barrier at a given temperature. This generally speeds up the reaction without being consumed in the process.

When a catalyst is introduced, it does not change the overall energy difference between the reactants and products; it simply makes it easier for the reactants to convert into products by lowering the energy barrier. Therefore, the presence of a catalyst can significantly increase the rate at which equilibrium is reached in both forward and reverse reactions, making them more efficient.

The other choices do not accurately reflect the behavior of catalysts in chemical reactions. A catalyst does not increase activation energy or have no effect on it; it actively reduces the activation energy. Furthermore, it does not eliminate the need for activation energy altogether, as some energy input is still necessary for the reactants to reach the transition state.