A catalyst and an intermediate are both important concepts in the realm of chemical reactions, but they serve distinct roles.
A catalyst is a substance that speeds up a chemical reaction without being consumed in the process. It lowers the activation energy needed for the reaction to occur, enabling the reactants to convert into products more quickly. Importantly, a catalyst is present at the beginning of the reaction and is regenerated at the end, meaning its amount remains the same throughout the process.
On the other hand, an intermediate is a species that forms during the reaction and is consumed in subsequent steps. Intermediates exist only transiently and are often less stable than the reactants or products. They can be identified during the reaction, but they do not appear in the overall balanced equation, as they are not present at the reaction’s start or end.
To tell them apart, consider their roles in the reaction mechanism. If a substance is consumed and produced within the reaction mechanism and does not appear in the overall equation, it is an intermediate. Conversely, if a substance facilitates the reaction without undergoing any permanent change and is not consumed, it is a catalyst. Understanding these differences is crucial for studying reaction kinetics and mechanisms in chemistry.