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How is Activated Complex Theory different from Transition State Theory?
The Activated Complex Theory, also known as the Transition State Theory, is a chemical theory that explains the process of chemical reactions at a molecular level. It proposes that during a chemical reaction, reactant molecules must first reach a high-energy state, known as the activated complex, before forming the products.
The Activated Complex Theory states that the rate of a chemical reaction is determined by the number of molecules that have enough energy to form the activated complex. The higher the concentration of these high-energy molecules, the faster the reaction rate will be.
The formation of the activated complex is affected by several factors, including temperature, concentration, and the presence of a catalyst. An increase in temperature and concentration increases the number of high-energy molecules, while a catalyst lowers the activation energy required for the formation of the activated complex.
The Activated Complex Theory is a general theory that can be applied to most chemical reactions. However, it does not account for reactions that occur without the formation of an activated complex, such as some photochemical reactions.
The Arrhenius equation, which relates the rate constant of a reaction to temperature, can be derived from the Activated Complex Theory. It shows that as temperature increases, the rate constant and therefore the rate of the reaction also increases.