Why are uncatalysed reactions often slower than their catalysed counterparts?

Uncatalysed reactions are often slower than their catalysed counterparts because catalysts lower the activation energy required for the reaction.

In a chemical reaction, the reactants need to overcome a certain energy barrier known as the activation energy to form the products. This activation energy is the minimum energy required for the reaction to occur. In an uncatalysed reaction, the reactants need to possess this energy naturally, which can often be quite high. This means that the reaction will proceed slowly, as only a small proportion of the reactants will have enough energy at any given time.

Catalysts, on the other hand, provide an alternative reaction pathway with a lower activation energy. They achieve this by forming temporary bonds with the reactants, creating an intermediate species that requires less energy to transform into the products. This means that a larger proportion of the reactants will have enough energy to react at any given time, thus speeding up the rate of the reaction.

Furthermore, catalysts are not consumed in the reaction, meaning they can continue to catalyse many reactions over time. This is another reason why catalysed reactions are often faster than uncatalysed ones.

In summary, catalysts increase the rate of a reaction by lowering the activation energy, providing an alternative reaction pathway, and being able to catalyse many reactions over time. This is why uncatalysed reactions, which lack these advantages, are often slower.