n. (context chemistry physics English) The energy required to initiate a reaction. For example, the flame from the fuse of a firecracker provides a small initial amount of energy, after which the explosive reaction proceeds by itself, releasing a considerably larger quantity of energy. A small push given to a stable but top-heavy object may cause it to fall over; the potential energy released during the fall was present in the system all along but could not be realized as long as the object was upright and balanced.
n. the energy that an atomic system must acquire before a process (such as an emission or reaction) can occur; "catalysts are said to reduce the energy of activation during the transition phase of a reaction" [syn: energy of activation]
In chemistry, activation energy is a term introduced in 1889 by the Swedish scientist Svante Arrhenius to describe the minimum energy which must be available to a chemical system with potential reactants to result in a chemical reaction. Activation energy may also be defined as the maximum energy required to start a chemical reaction. The activation energy of a reaction is usually denoted by E and given in units of kilojoules per mole (kJ/mol) or kilocalories per mole (kcal/mol).
Activation energy can be thought of as the height of the potential barrier (sometimes 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 translational energy equal to or greater than the activation energy.
Usage examples of "activation energy".
At low temperatures, certain classes of molecules require very little activation energy to undergo chemical reactions, but because our laboratories are at room temperature and not, say, at the temperature of Neptune's satellite Triton, our knowledge of those molecules may well be inadequate.