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Further Explanation
In physical chemistry, chemical kinetics or reaction kinetics studies the reaction rate in a chemical reaction. Analysis of the effect of various reaction conditions on the reaction rate provides information about the reaction mechanism and the transition state of a chemical reaction. In 1864, Peter Waage pioneered the development of chemical kinetics by formulating the law of mass action, which states that the speed of a chemical reaction is proportional to the quantity of the reacting substance.
The Arrhenius equation provides the basic value of the relationship between the activation energy and the rate of the reaction process.
In chemistry, activation energy is a term introduced by Svante Arrhenius, which is defined as energy that must be surpassed for chemical reactions to occur. Activation energy can also be interpreted as the minimum energy needed for certain chemical reactions to occur. The activation energy of a reaction is usually denoted as Ea, in kilojoules per mole (KJ / mol).
The chemical reaction is a natural process that always produces interchange of chemical compounds. The initial compounds or compounds involved in the reaction are called reactants. Chemical reactions are usually characterized by chemical changes and will produce one or more products that usually have different characteristics from reactants. Classically, chemical reactions involve changes involving the movement of electrons in the formation and breaking of chemical bonds, although the general concept of chemical reactions can also be applied to the transformation of elementary particles such as in nuclear reactions.
Collision theory is a theory independently proposed by Max Trautz in 1916 and William Lewis in 1918, which qualitatively explains how chemical reactions occur and how reaction rates differ for different reactions. Collision theory states that when the corresponding reactant particles collide with each other, only a certain percentage of collisions cause real or significant chemical changes; this successful change is called a successful collision. A successful collision has enough energy, also known as activation energy when the collision is to break the pre-existing bond and form all new bonds. This results in a reaction product. Increasing the concentration of reactant particles or raising the temperature, giving rise to more collisions and therefore more successful collisions increase the rate of reaction.
Collision theory is very closely related to chemical kinetics.
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definition of chemical kinetics
definition of The Arrhenius equation
definition of The chemical reaction
definition of Collision theory
Details
Grade: High School
Subject: Chemistry
keywords: chemical kinetics
