Activation+Energy


 * activation energy** is a term introduced in 1889 by the Swedish scientist [|Svante Arrhenius] that is defined as the energy that must be overcome in order for a [|chemical reaction] to occur. Activation energy may also be defined as the minimum energy required to start a chemical reaction. The activation energy of a reaction is usually denoted by //Ea//, and given in units of [|kilojoules per mole].

The least amount of energy needed for a chemical reaction to take place. Some elements and compounds react together naturally just by being close to each other, and their activationenergy is zero. Others will react together only after a certain amount of energy is added to them. Striking a match on the side of a matchbox, for example, provides the activation energy (in the form of heat produced by friction) necessary for the chemicals in the match to ignite. Activation energy is usually expressed in terms of joules per mole of reactants.

Activation energy can be thought of as the height of the [|potential barrier] (called the energy barrier) separating two [|minima] of [|potential energy] (of the reactants and products of a reaction). For a chemical reaction to proceed at a reasonable rate, there should exist an appreciable number of molecules with energy equal to or greater than the activation energy. At a more advanced level, the Arrhenius Activation is energy term from the Arrhenius equation is best regarded as an experimentally determined parameter that indicates the sensitivity of the reaction rate to temperature. There are two objections to associating this activation energy with the threshold barrier for an elementary reaction. First, it is often unclear as to whether or not reaction does proceed in one step; threshold barriers that are averaged out over all elementary steps have little theoretical value. Second, even if the reaction being studied is elementary, a spectrum of individual collisions contributes to rate constants obtained from bulk ('bulb') experiments involving billions of molecules, with many different reactant collision geometries and angles, different translational and (possibly) vibrational energies - all of which may lead to different microscopic reaction rates.



__**Negative Activation Energy**__
Rates of reaction decrease with increasing temperature. An Arrhenius expression results in a negative value of //Ea//. Increasing the temperature leads to a reduced probability of the colliding molecules capturing one another, expressed as a reaction that decreases with increasing temperature. Negative activation energy: In some cases, rates of reaction //decrease// with increasing temperature. When following an approximately exponential relationship so the rate constant can still be fit to an Arrhenius expression, this results in a negative value of //Ea//. Elementary reactions exhibiting these negative activation energies are typically barrierless reactions, in which the reaction proceeding relies on the capture of the molecules in a potential well. Increasing the temperature leads to a reduced probability of the colliding molecules capturing one another (with more glancing collisions not leading to reaction as the higher momentum carries the colliding particles out of the potential well), expressed as a reaction cross section that decreases with increasing temperature. Such a situation no longer leads itself to direct interpretations as the height of a potential spot.