Transition state

Transition state is a term used in chemistry and biochemistry to refer to the highest-energy state of a molecule during a chemical reaction. It is a state through which a molecule must pass as it is being transformed from reactants to products. The transition state is often symbolized by the double dagger ‡ in chemical equations.

Overview

The concept of the transition state is fundamental to the understanding of chemical kinetics and the rates of chemical reactions. It was first proposed by Henry Eyring, Michael Polanyi and Meredith Gwynne Evans in the 1930s, and is based on the principles of quantum mechanics.

In a chemical reaction, the transition state is the point at which the original bonds have stretched to their limit and are about to break, while the new bonds are forming. This state is very unstable and exists for only a very short time – typically on the order of femtoseconds (10^-15 seconds).

The energy required to reach the transition state from the reactants is known as the activation energy. The lower the activation energy, the faster the reaction will proceed.

Transition state theory

Transition state theory (TST), also known as activated complex theory, is a theory that explains the rates of chemical reactions. According to TST, the rate of a reaction is directly proportional to the number of molecules that have enough energy to reach the transition state.

TST provides a way to calculate the rate constant for a reaction, based on the energies of the reactants and the transition state. It also allows for the prediction of reaction rates under different conditions, such as changes in temperature or pressure.

See also

• Reaction rate
• Chemical kinetics
• Potential energy surface
• Arrhenius equation

References

<references />

Error creating thumbnail: This article is a medical stub. You can help WikiMD by expanding it!
Most recent articles Ongoing trials	Wikipedia