Activation energy

The activation energy, characterized in 1889 by Svante Arrhenius is an energy barrier that must be overcome in a chemical reaction of the reactants. In general, the lower is the activation energy, the faster the reaction. A high activation energy inhibits reactions that would be expected for reasons of energy, thereby preventing the recruitment of a ( thermodynamic ) chemical equilibrium. Thus, a mixture of methane and oxygen in the air at standard conditions exist almost unchanged ( ie, the reaction is immeasurably slowly ), although the combustion products carbon dioxide and water are thermodynamically more stable. Exceptions are, for example, Acid -base reactions.

Arrhenius activation energy is an empirical quantity which can be determined by the high temperature dependence of the rate of many chemical reactions.

Reaction kinetics

In the idea of physical chemistry has to be in the course of a chemical reaction is a rearrangement of the atoms from the arrangement of the reactants in the products take place, where old bonds broken and new bonds are formed. The reactants undergo this conversion whose products an activated state, the so-called transition state (see figure on the right, curve maximum ), the formation of which requires a certain energy (activation energy). The rate of a chemical reaction that is dependent on ( at constant temperature), not the energy of the reactants and products, but only by the energy difference between the reactants and the transition state. The higher the temperature of the reaction system, the higher is also the probability that the reactant to provide the activation energy required to overcome the energy barrier and further react with the product.

Temperature dependence of the reaction rate

The relationship between the rate constant K, the activation energy Ea and the thermodynamic temperature T can be described (neglecting the activation volume and at low pressures ) in many cases, by the Arrhenius equation:

Logarithm of the equation, we obtain:

Since the pre-exponential factor A is often sufficient, regardless of the temperature, the following applies:

Possible to determine the rate constants of irreversible reactions at different temperatures, ln ( k) against 1 / T are applied and EA are determined from the slope of the line; see Arrheniusgraph.

Two rate constants (K1 and K2 ) of a reaction at two temperatures (T1 and T2 ) are known, can be combined with

Calculate the activation energy. The equation is a difference of two logarithmic transformed Arrhenius equations ( one for each temperature). In many reactions in solution, the activation energy is in the range of 50 kJ · mol -1. A temperature rise of 290 K to 300 K leads to approximately a doubling of the rate constant (see RGT rule). However, it is always to be considered that as the activation energy ( that is, with increasing pitch in the Arrheniusgraph ) the effect of the temperature dependency is enhanced. Thus, reactions with low activation energy ( about 10 kJ · mol -1 ) is only slightly accelerated by increasing the temperature. The rate of reactions with large activation energy ( about 60 kJ · mol -1 ), however, increases strongly with increasing temperature. The values ​​for the molar activation energies of many common reactions are between 30 and 100 kJ · mol -1.

In some reactions, the temperature dependence of the rate constant does not follow the Arrhenius equation. These are, for example, reactions without activation energy, explosive type reactions, reactions with upstream equilibria and many enzymatic or straight catalytic reactions.

Theoretical background

In fact, the model of Arrhenius happening when a chemical reaction describes incomplete. The Arrhenius equation can be explained theoretically by the classical collision theory. The high effect of increasing the temperature on the rate of reaction due to the strong increase in the proportion of the particles have sufficient energy to overcome the barrier. Besides, grows with a temperature increase, the frequency of collisions ( the collision frequency ) of the reactants. However, the increase in the number of collisions leads practically to a very small increase in the rate of reaction and goes into the Arrhenius equation as a component in the " temperature-independent " pre-exponential factor A below. If the activation energy is small or zero determine the collision rate or the diffusion rate of the reaction rate. According to Eyring (see also theory of the transition state ), the free energy of activation is the determining factor for the reaction rate.

Catalysis

A catalyst is the activation energy for chemical reactions down, but does not change the free reaction enthalpy. It is believed that in the presence of a catalyst, a complex having a lower activation energy and is formed as the reaction probability increases.