Svante Arrhenius introduced the term activation energy in 1889. It is a term that refers to the amount of energy necessary in order for a chemical reaction to occur.Another way of defining activation energy is the minimum amount of energy needed for a chemical reaction to take place.EA is the activation energy of a reaction, and is measured as a number of kilojoules per mole.

In general, activation energy can be taken as the height of the potential barrier (sometimes called the energy barrier) that separates two minima of potential energy (the reactants and the products).For a chemical reaction to be noticeable, there should be a noticeable number of molecules with energy equal to or higher than the activation energy.

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Transition states

Activation energy broadly correlates with transition state barrier height in most reactions.A transition state along a reaction coordinate is where bond formation and bond breakdown are balanced.Several transition states can be involved in multiple-step reactions.According to the overall activation energy, the transition state barrier is the rate-limiting step.

For a large number of reactions, however (those with loose transition states, tunneling reactions, and barrierless reactions), the height of the highest barrier on the reaction path is not equal to the activation energy implied by the temperature dependence of the reaction rate (see Arrhenius equation).In these cases, an activation energy may be considered as the height of an effective barrier that would yield the same rate if these factors weren't present.

In their definition of activation energy, it should be noted that IUPAC does not refer to transition states (see external links).

Negative activation energy

In some cases, reaction rates decrease as temperature increases.In the case of an approximately exponential relationship to fit an Arrhenius expression, Ea has a negative value when following a nearly exponential relationship.Typically, reactions with these negative activation energies are barrierless reactions in which the molecules are captured in a potential well to initiate the reaction.If the temperature is raised, there is a decreased probability of molecules colliding with one another (with less glancing collisions causing reaction since the collisions' higher momentum carries them from the potential well), represented by a decrease in reaction cross sections with increasing temperature.A situation like this no longer lends itself to direct interpretations as the height of a potential barrier.

Temperature independence and the relation to the Arrhenius equation

In Arrhenius' equation, the activation energy is related quantitatively to the rate at which a reaction proceeds.As a result of the Arrhenius equation, the activation energy is


Here, A is the frequency factor of the reaction, R is the universal gas constant, and T is the temperature (in Kelvin).This equation suggests that the activation energy varies with temperature, but the temperature dependence of k cancels this effect in regimes in which the Arrhenius equation holds.The activation energy can therefore be calculated from the rate constant at any temperature (within the Arrhenius equation's validity).


A catalyst is a substance that lowers the activation energy of the transition state; an enzyme is a biological catalyst.A catalyst increases the rate of a reaction without being consumed by it.Moreover, the catalyst lowers the activation energy while keeping the energy of the original reactants and products the same.Instead, the reactant energy and the product energy remain unchanged, and only the activation energy is altered.

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