activation energy

activation energy, in chemistry, minimum energy needed to cause a chemical reaction. A chemical reaction between two substances occurs only when an atom, ion, or molecule of one collides with an atom, ion, or molecule of the other. Only a fraction of the total collisions result in a reaction, because usually only a small percentage of the substances interacting have the minimum amount of kinetic energy a molecule must possess for it to react. When the reactants collide, they may form an intermediate product whose chemical energy is higher than the combined chemical energy of the reactants. In order for this transition state in the reaction to be achieved, some energy must enter into the reaction other than the chemical energy of the reactants. This energy is the activation energy. Once the intermediate product, or activated complex, is formed, the final products are formed from it. The path from reactants through the activated complex to the final products is known as the reaction mechanism. (Reaction mechanisms for complex reactions may involve several steps analogous to that described here.) Because the heat energy of a substance is not uniformly distributed among its atoms, ions, or molecules, some may carry enough heat energy to react while others do not. If the activation energy is low, a greater proportion of the collisions between reactants will result in reactions. If the temperature of the system is increased, the average heat energy is increased, a greater proportion of collisions between reactants result in reaction, and the reaction proceeds more rapidly. A catalyst increases the reaction rate by providing a reaction mechanism with a lower activation energy, so that a greater proportion of collisions result in reaction. The activation energy and rate of a reaction are related by the equation k = A exp( - E a / RT ), where k is the rate constant, A is a temperature-independent constant (often called the frequency factor), exp is the function ex, E a is the activation energy, R is the universal gas constant, and T is the temperature. This relationship was derived by Arrhenius in 1899. Because the relationship of reaction rate to activation energy and temperature is exponential, a small change in temperature or activation energy causes a large change in the rate of the reaction. Activation energies are usually determined experimentally by measuring the reaction rate k at different temperatures T, plotting the logarithm of k against 1/ T on a graph, and determining the slope of the straight line that best fits the points.


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