Uncatalyzed reactions typically proceed at a slower rate than their catalyzed counterparts due to the role of catalysts in reducing the activation energy required for the reaction to occur.

In any chemical reaction, reactants must overcome a specific energy barrier known as the activation energy, denoted as $E_a$. This activation energy represents the minimum energy necessary for the reactants to transform into products. In the case of uncatalyzed reactions, the reactants must inherently possess this energy, which can often be quite substantial. As a result, these reactions tend to occur slowly, as only a small fraction of the reactants have sufficient energy to react at any given moment.

Conversely, catalysts facilitate a different reaction pathway that requires a lower activation energy. They do this by forming temporary bonds with the reactants, resulting in the creation of intermediate species that can convert to products with less energy input. Consequently, a greater proportion of the reactants possess the necessary energy to undergo the reaction at any given time, significantly accelerating the reaction rate.

Moreover, catalysts are not consumed during the reaction process, allowing them to catalyze multiple reactions over time. This characteristic further contributes to the increased speed of catalyzed reactions compared to uncatalyzed ones.

In summary, catalysts enhance the reaction rate by lowering the activation energy, providing an alternative reaction pathway, and remaining available for numerous reactions over time. This is why uncatalyzed reactions, which do not benefit from these advantages, are often slower.