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Activation energyThe activation energy in chemistry is the energy needed by a system to initiate a particular process. Activation energy is often used to denote the minimum energy needed for a specific chemical reaction to occur. For a reaction to occur between two colliding molecules they must collide in the correct orientation and possess a certain minimum amount of energy. As the molecules approach their electron clouds repel. This requires energy - activation energy - and comes from the heat of the system i.e., the translational, vibrational etc... energy of each molecule. If there is enough energy available, this repulsion is overcome and the molecules get close enough for attractions between the molecules to cause a rearrangement of bonds. The Arrhenius equation gives the quantitative basis of the relationship between the activation energy and the rate at which a reaction proceeds. The study of reaction rates is termed chemical kinetics.
The transition state in a reaction is the point at which the original bonds have stretched to their limit. Transition states are only in existence for extremely brief (10-15 sec) periods of time. The energy required to reach the transition state is equal to the activation for that reaction. Multi-stage reactions involve a number of transition points, here the activation energy is equal to the one requiring the most energy. After this time either the molecules move apart again with original bonds reforming, or the bonds break and new products form. This is possible because both possibilities result in the release of energy (shown on the enthalpy profile diagram, Fig. 1, as both positions lie below the transition state). A substance that modifies the transition state to lower the activation energy is termed a catalyst; a biological catalyst is termed an enzyme.
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