Activation Energy (Molecular Biology)

The kinetic rates of chemical and enzyme-catalyzed reactions depend on temperature. The relationship for this dependence is known as the Arrhenius law:


where A and E are constants, R is the gas constant, and T is the absolute temperature in Kelvin units. The activation energy E is the height of the energy barrier that the reaction must exceed to pass from reactants to products. It is usually expressed in kilocalories per mole or in joules per mole. Equation 1 can be expressed as:


Therefore the magnitude of the activation energy can be obtained from the slope of a plot of the log of a rate constant for a reaction as a function of 1/T. Many chemical and enzyme-catalyzed reactions increase the rate of reaction by two- to threefold for each 10°C increase in temperature. Although this relationship is useful in explaining the temperature dependence of reactions, it does not explain the rate in the thermodynamic terms of enthalpy H entropy S, or free energy G. This analysis comes from transition state theory. The Arrhenius activation energy does correspond to the standard enthalpy of the reaction in the van’t Hoff equation (1).

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