Activation Energy (redirected from "energy of activation")

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=Aexp(−E_a/RT), where k is the rate constant, A is a temperature-independent constant (often called the frequency factor), exp is the function e^x, 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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activation energy

Minimum amount of energy (heat, electromagnetic radiation, or electrical energy) required to activate atoms or molecules to a condition in which it is equally likely that they will undergo chemical reaction or transport as it is that they will return to their original state. Chemists posit a transition state between the initial conditions and the product conditions and theorize that the activation energy is the amount of energy required to boost the initial materials “uphill” to the transition state; the reaction then proceeds “downhill” to form the product materials. Catalysts (including enzymes) lower the activation energy by altering the transition state. Activation energies are determined by experiments that measure them as the constant of proportionality in the equation describing the dependence of reaction rate on temperature, proposed by Svante Arrhenius. See also entropy, heat of reaction.

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activation energy [‚ak·tə′vā·shən ′en·ər·jē]

(physical chemistry)

The energy, in excess over the ground state, which must be added to an atomic or molecular system to allow a particular process to take place.

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Energy, Activation 

the difference between the values of the average energy of particles (molecules, radicals, ions) that participate in an elementary act of a chemical reaction and the average energy of all the particles in the reacting system.

Activation energy varies widely for different chemical reactions, from a few joules per mole to about 10 J/mole. For the same chemical reaction, the value of the activation energy depends on the form of the distribution functions of the molecules with respect to the energies of their translational motion and with respect to the internal degrees of freedom (electronic, vibrational, rotational). Activation energy should be differentiated as a statistical quantity from the threshold energy, the minimum energy that must be attained by a pair of colliding particles for a given elementary reaction to occur.

According to concepts of the theory of absolute reaction rates, activation energy is the difference between the average energy of activated complexes and the average energy of the initial molecules.

The idea of activation energy originated in the 1870’s and 1880’s as a result of research by J. Van’t Hoff and S. Arrhenius regarding the influence of temperature on the rate of a chemical reaction. The constant k of the rate of a reaction is related to the activation energy E by the Arrhenius equation:

k = k₀e^–E/RT

where R is the gas constant, T is the absolute temperature in °K, and k₀ is a constant called the pre-exponential, or frequency, factor of the rate constant. The equation, based on the molecular-kinetic theory, was later derived in statistical physics with consideration of a number of simplifying assumptions, one being that activation energy is independent of temperature. This assumption is valid for practical purposes and for theoretical calculations over comparatively narrow temperature ranges.

Activation energy can be found from experimental data by several methods. In one technique, the kinetics of a reaction are studied at different temperatures (for information on the various methods, seeREACTION RATE, CHEMICAL). A graph is plotted in coordinates of In k versus 1/T; according to the Arrhenius equation, the slope of the straight line on the graph is equal to E. For single-space reversible reactions, the activation energy of the reaction in the forward or reverse direction can be calculated if the activation energy of the reaction in the other direction is known, and the temperature dependence of the equilibrium constant is obtained from thermodynamic data. The temperature dependence of activation energy must be considered for more exact calculations.

The activation energy of complex reactions is a combination of the activation energies of the elementary stages. Sometimes the concept of “apparent” activation energy is used in addition to the true activation energy, determined from the Arrhenius equation. For example, if the rate constants of heterogeneously catalytic reactions are determined from changes in the volumetric concentrations of the initial substances and the products, then the apparent activation energy differs from the true value by the magnitude of the thermal effects that accompany processes of adsorption and desorption of the reacting substances on the surface of the catalyst. Determination of the activation energy is a very complex problem in nonequilibrium systems, for example, in plasma chemical systems. However, the formal application of the Arrhenius equation is possible in some cases.

The concept of activation energy is the most important concept of chemical kinetics; values are listed in special handbooks, and they are used in chemical technology to calculate the rate constants of reactions under various conditions.

IU. A. KOLBANOVSKII