A transition state, also referred to as an activated complex, is a critical concept in chemical reactions as it represents the point of highest energy along the reaction pathway.

To elaborate, the transition state is a temporary, high-energy arrangement of atoms located at the peak of the reaction coordinate diagram, which illustrates the progress of a chemical reaction. At this stage, the reactants are in the process of transforming into products, although the conversion is not yet complete.

The energy required to reach the transition state from the reactants is termed the activation energy, denoted as $E_a$. This is the minimum amount of energy that must be supplied for a reaction to occur. Generally, the higher the activation energy, the slower the reaction rate, because fewer molecules possess the necessary energy to attain the transition state.

Understanding the transition state is also essential for grasping the role of catalysts in chemical reactions. Catalysts function by offering an alternative reaction pathway that has a lower activation energy, thereby enabling a greater number of molecules to reach the transition state and subsequently increasing the reaction rate.

Moreover, transition state theory, developed by Henry Eyring, provides a framework for calculating reaction rates. According to this theory, the reaction rate is proportional to the number of molecules that can attain the transition state. This theoretical approach has been pivotal in enhancing our understanding and predictions regarding the behavior of chemical reactions.

In summary, the transition state or activated complex is a fundamental concept in chemical kinetics, crucial for understanding reaction rates, the influence of catalysts, and the theoretical predictions of reaction behavior.