Transition states
The transition state of a chemical reaction is a particular configuration along the reaction coordinate. It is defined as the state corresponding to the highest potential energy along this reaction coordinate. At this point, assuming a perfectly irreversible reaction, colliding reactant molecules always go on to form products. It is often marked with the double dagger ‡ symbol.
As an example, the transition state shown below occurs during the SN2 reaction of bromoethane with a hydroxyl anion:
The activated complex of a reaction can refer to either the transition state or to other states along the reaction coordinate between reactants and products, especially those close to the transition state.
The concept of a transition state has been important in many theories of the rates at which chemical reactions occur. This started with the transition state theory (also referred to as the activated complex theory), which was first developed around 1935 by Eyring, Evans and Polanyi, and introduced basic concepts in chemical kinetics that are still used today.
A collision between reactant molecules may or may not result in a successful reaction. The outcome depends on factors such as the relative kinetic energy, relative orientation and internal energy of the molecules. Even if the collision partners form an activated complex they are not bound to go on and form products, and instead the complex may fall apart back to the reactants.
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