Theory of Rate of Reaction

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Everyday experience demonstrates that, most of the time, the rate of a chemical reaction will increase with a rise in temperature. Food, for example, will spoil outside on a hot summer day much faster than it would in a refrigerator. A simple but remarkably accurate relationship for the temperature dependence of reaction rates was empirically found by the Swedish chemist Svante Arrhenius in 1889. The interpretation of the parameters in the Arrhenius equation leads to the development of the idea that when reactants convert into products, they must go through an activated state that requires a characteristic energy. This was the basis of two of the most important theories of reaction rates, collision theory and transition state theory. Collision theory, which only suffices for simple gas phase reactions, essentially views reactants as if they were particles with a certain kinetic energy. Reactions can only occur if two molecules collide with a minimum energy necessary for rearranging the bonds. Matter dynamic considerations play no role here. In transition state theory, a more comprehensive theory that can, in principle, be applied to every possible type of reaction, the rate coefficient is expressed in terms of a difference in chemical potentials between the reactants and a kind of “transition substance” (“ensemble” of all activated complexes), a so-called “potential barrier.” For a deeper understanding, the transition state can be interpreted on a molecular level with the help of potential energy surfaces and the “motion” of molecules through these surfaces.