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Activation energy
FIELDS OF STUDY
Physical Chemistry; Inorganic Chemistr y; Biochemistry
SUMMARY
The activation energy of a process is defined, and
its importance in chemical processes is elaborated.
Activation energy is a widely variable quantity in dif-
ferent reactions but is nevertheless characteristic of
any specific reaction process.
PRINCIPAL TERMS
Arrhenius equation: a mathematical function
that relates the rate of a reaction to the energy
required to initiate the reaction and the absolute
temperature at which it is carried out.
catalyst: a chemical species that initiates or speeds
up a chemical reaction but is not itself consumed
in the reaction.
chemical reaction: a process in which the mole-
cules of two or more chemical species interact with
each other in a way that causes the electrons in the
bonds between atoms to be rearranged,
resulting: in changes to the chemical identities of
the materials.
reaction rate: how much of a particular reaction or
reaction step occurs per unit time.
transition state: an unstable structure formed
during a chemical reaction at the peak of its poten-
tial energy that cannot be isolated and ultimately
breaks down, either forming the products of the
reaction or reverting back to the original reactants.
Activation Energy in Chemical Reactions
Activation energy can be thought of as a barrier that
the reactants in a chemical reaction must overcome if
the reaction is to proceed to the formation of prod-
ucts. The molecules that are involved must rearrange
to form either a transition state or an intermediate
that is higher in energy than the starting materials.
An intermediate is a stable chemical structure formed
during a reaction process that can often be captured
and isolated by chemical means. Once this structure
is formed, the reaction process can either progress to
form products or revert back to the original reactants.
Activation energy is the energy required for a spe-
cific chemical reaction to occur. In a reaction, two
reactant molecules contact each other with the en-
ergy of their ambient states.(The ambient state of a
material can be thought of as its “default” state—the
state it takes at one atmosphere of pressure and what
is commonly considered to be room temperature.) In
the case of a spontaneous reaction, the energy of the
collision is sufficient to initiate the formation of the
transition state or intermediate. In a nonspontaneous
reaction, the energy of the molecular collision is not
sufficient, and the two molecules will not interact.
The input of some additional energy is required to
drive the two molecules together so that the transi-
tion state or intermediate is formed and the reaction
can proceed. The energy released in the transforma-
tion of reactants into products is generally sufficient
to drive the reactions of other molecules in the reac-
tion mixture.
Another way to look at activation energy is to think
of it as the minimum amount of energy that two inter-
acting molecules must gain in order to weaken bonds
between atoms in both molecules so that those bonds
can be rearranged. Since chemical reactions are es-
sentially processes of breaking and making bonds,
having sufficient energy to overcome the strength of
the appropriate bonds is essential if there is to be any
reaction between the two molecules.
Activation Energy and Reaction Rates
Reaction rates can be related directly to their ac-
tivation energies. This relationship is defined by
the Arrhenius equation, formulated in 1884 by the
Swedish scientist Svante Arrhenius (1859–1927), who
received the Nobel Prize in Chemistry in 1903.The
