Chemical Kinetics - SS2 Chemistry Past Questions and Answers - page 6

Collision theory is a fundamental concept that explains the rate of chemical reactions based on the collision of reactant particles. According to collision theory, for a reaction to occur, reactant particles must collide with sufficient energy and proper orientation.

In the case of concentration, an increase in reactant concentration leads to a higher reaction rate. This is because a higher concentration results in a greater number of reactant particles within a given volume, increasing the probability of collisions. According to collision theory, an increase in concentration increases the collision frequency, resulting in more successful collisions and a faster reaction rate.

Regarding temperature, collision theory explains that raising the temperature increases the reaction rate. This is because higher temperatures provide more energy to reactant particles, increasing their kinetic energy and collision frequency. The increased kinetic energy leads to more frequent and energetic collisions, increasing the chances of successful collisions. Therefore, collision theory predicts that an increase in temperature accelerates reaction rates.

Collision theory also emphasises the importance of proper orientation during collisions. For a reaction to occur, reactant particles must collide in a specific way that allows the necessary bond-breaking and bond-forming processes. Proper orientation ensures that reactant molecules collide in a manner that facilitates the formation of products. Collision theory explains that the frequency of proper orientation increases with higher concentrations and temperatures, leading to an increased reaction rate.

Activation energy is the minimum energy required for a chemical reaction to occur. It represents the energy barrier that reactant particles must overcome to initiate the reaction. The activation energy is needed to break the existing bonds in the reactant molecules, allowing the formation of new bonds and the conversion of reactants into products.

The activation energy influences the rate of a chemical reaction. A higher activation energy generally results in a slower reaction rate because fewer reactant particles possess the necessary energy to surpass the energy barrier. As a result, the reaction occurs less frequently.

Conversely, a lower activation energy allows a larger fraction of reactant particles to overcome the barrier and proceed to form products. This results in a faster reaction rate since more collisions have sufficient energy to react.

Catalysts play a critical role in modifying the activation energy of a reaction. Catalysts lower the activation energy by providing an alternative reaction pathway with a lower energy barrier. This reduction in activation energy enables more reactant particles to reach the activation energy threshold, leading to an increased rate of reaction.

Overall, activation energy determines the rate of chemical reactions by influencing the frequency of successful collisions and the formation of activated complexes. It serves as a kinetic barrier that reactant particles must surpass to proceed to product formation.