How do collision theory and transition state theory differ in explaining reaction rates?

Collision theory explains reaction rates based on the frequency and energy of colliding molecules, while transition state theory focuses on the energy barrier to reactions.

Collision theory is a model used to explain the rates of chemical reactions. According to this theory, for a reaction to occur, not only must particles collide, but they must also have sufficient energy and the correct orientation. The rate of reaction is directly proportional to the number of effective collisions per unit time. This theory assumes that all collisions between reactant molecules result in a reaction, which is not always the case. Some collisions might not lead to a reaction, due to insufficient energy or incorrect orientation of the molecules.

On the other hand, transition state theory, also known as activated complex theory, provides a more detailed and accurate explanation of reaction rates. This theory proposes that molecules form a transition state or activated complex during the reaction. The transition state is a high-energy, unstable arrangement of atoms which exists momentarily at the peak of the energy barrier to the reaction. The rate of reaction is determined by the rate at which the transition state is formed. This theory takes into account the energy required to reach the transition state, which is the activation energy, and the orientation of the molecules.

In essence, while collision theory provides a basic understanding of reaction rates, it oversimplifies the process by assuming all collisions lead to a reaction. Transition state theory, however, provides a more comprehensive explanation by considering the energy barrier and the formation of an activated complex. Both theories are essential in understanding reaction rates, but transition state theory offers a more nuanced and detailed explanation.

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