Why do reaction rates differ in homogeneous and heterogeneous reactions?

Reaction rates differ in homogeneous and heterogeneous reactions due to the differences in their physical states and contact between reactants.

In homogeneous reactions, all reactants are in the same phase, such as all gases or all liquids. This allows for a uniform distribution of reactants, leading to a higher probability of collisions between reactant molecules, and thus a faster reaction rate. The reactants are uniformly distributed throughout the reaction mixture, which means that every molecule has an equal chance of colliding with another molecule. This increases the likelihood of successful collisions, where the molecules have enough energy to overcome the activation energy and react.

On the other hand, in heterogeneous reactions, the reactants are in different phases, such as a solid in a liquid or a gas on a solid. The reaction can only occur at the interface between the two phases, which limits the contact between reactant molecules. This results in a lower probability of successful collisions and thus a slower reaction rate. The rate of reaction in heterogeneous reactions is often limited by the surface area of the solid or the boundary between the two phases. The larger the surface area, the more opportunities there are for collisions to occur, and the faster the reaction rate.

Furthermore, in heterogeneous reactions, factors such as the dispersion of the solid phase, the porosity of the solid, and the solubility of the gas in the liquid phase can also affect the reaction rate. These factors can influence the availability of reactant molecules at the interface, and thus the number of successful collisions.

In summary, the difference in reaction rates between homogeneous and heterogeneous reactions is primarily due to the differences in the physical states of the reactants and the contact between them.