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The reaction involved the effective collision of two reactants to produce the desired products. Reactions can be natural, which occur in the surrounding environment, whereas it can be artificially done in the laboratory to form the desired product.

The reaction rate can be defined as the reaction speed to produce the products. The reaction rate can be slow, fast or moderate. The reaction can take less than a millisecond to produce products, or it can take years to produce the desired product.

'Activation energy' term was used by Svante Arrhenius, a Swedish scientist, in the year 1889. The activation energy can be defined as the threshold energy (which is called the minimum amount of energy) after which the reaction takes place. A chemical reaction can take place because of the effective collision of the two or more reactants, which produces energy more than the threshold energy (minimum energy).

The unit of the Activation energy is Joule or kcal/mole.

The rate of the reaction depends upon the Activation energy. As the activation energy is high, then the reaction rate is low, whereas if the activation energy is low, then the reaction rate is high. Activation energy can be defined as the 鈥渂arrier鈥� which is crossed by the reactant to produce products.