How does a catalyst increase the rate of reaction?

A catalyst increases the rate of a reaction by providing an alternative pathway with lower activation energy. Activation energy is the minimum energy required for a reaction to occur. When a catalyst is present, it lowers this threshold, allowing more reactant molecules to have sufficient energy to overcome the energy barrier and form products. This results in a faster reaction rate because more effective collisions can occur within a given time frame.

The alternative pathway facilitated by a catalyst does not change the overall energy change of the reaction; instead, it simply means that the energy required to reach the transition state is lessened. This is critical because it allows reactions to occur at lower temperatures or under milder conditions, making catalytic processes both efficient and effective.

In contrast, the other options do not accurately reflect how catalysts function. For instance, absorbing heat does not directly relate to altering the reaction mechanism. Similarly, a pathway with higher activation energy would slow down reactions rather than speed them up, and simply increasing the amount of reactants does not guarantee that the reaction will proceed faster without addressing the activation energy barrier.