Final answer:
A catalyst speeds up the reaction rate by lowering the activation energy and providing an alternative reaction pathway, resulting in more effective collisions between reactants. It does not get consumed and does not alter the reaction equilibrium or product stability.
Step-by-step explanation:
The effect of a catalyst on a reaction is to increase the rate at which the reaction takes place. A catalyst achieves this by providing a new reaction pathway with a lower activation energy, thereby lowering the energy barrier for the reaction. It is crucial to note that while a catalyst speeds up the reaction, it is not consumed by the reaction and does not change the equilibrium position or make the products more stable.
Catalysts operate by bringing reactants together more quickly than they would on their own, which is possible due to the catalyst's affinity for its substrates. By lowering the overall activation energy, a greater percentage of reactant molecules can have effective collisions, which consequently increases the reaction rate. Even though the reaction rate is enhanced, the catalyst does not get used up in the process and can continue to facilitate additional reactions.