How does increasing the surface area of a solid affect the rate of reaction according to particle collision theory?

Increasing the surface area of a solid enhances the rate of reaction because it increases the exposure of particles and the frequency of collisions. According to particle collision theory, for a reaction to occur, reactant particles must collide with sufficient energy and proper orientation. When the surface area of a solid is larger, more particles are available to come into contact with reactant particles, leading to more frequent collisions.

This higher collision frequency increases the likelihood that reactant particles will collide with the necessary energy to overcome the activation energy barrier, thus promoting faster reactions. For example, powdered solids react more quickly than larger chunks of the same substance because the powdered form has a greater total surface area available for collisions.

In contrast, simply reducing the number of collisions, decreasing the energy of collisions, or suggesting that there is no effect on the reaction rate would not accurately represent the relationship between surface area and reaction kinetics as described by particle collision theory.