Kinetics describes the rate and path by which a reaction occurs. In contrast, thermodynamics deals with state functions and describes the properties, behavior, and components of a system. It is not concerned with the path taken by the process and cannot address the rate at which a reaction occurs. Although it does provide information about what can happen during a reaction process, it does not describe the detailed steps of what appears on an atomic or a molecular level. On the other hand, kinetics provides information on an atomic or molecular level. In short, thermodynamics focuses on the energetics of the products and the reactants, whereas kinetics focuses on the pathway from reactants to products. Industrial processes where the value of ΔG is negative and the corresponding value of K is greater than one are too slow to be economically profitable. In such cases, a thermodynamically non-spontaneous reaction can be made to occur spontaneously by changing the reaction conditions, such as varying the pressure or temperature, supplying an external source of energy in the form of electricity, etc.

Atoms, molecules, or ions must collide before they can react with each other. Atoms must be close together to form chemical bonds. This premise is the basis for a theory that explains many observations regarding chemical kinetics, including factors affecting reaction rates. The collision theory is based on the postulates that (i) the reaction rate is proportional to the rate of reactant collisions, (ii) the reacting species collide in an orientation allowing contact between the atoms that become bonded together in the product, and (iii) the collision occurs with adequate energy to permit mutual penetration of the reacting species’ valence shells so that the electrons can rearrange and form new bonds (and new chemical species). When reactant species collide with both the correct orientation and enough activation energy, they combine to form an unstable species called an activated complex or a transition state. These species are short-lived and usually undetectable by most analytical instruments. In some cases, sophisticated spectral measurements can observe transition states. Collision theory explains why most reaction rates increase as temperature increases; with an increase in temperature, the frequency of collisions increases. More collisions mean a faster reaction rate, assuming the energy of the collisions is adequate.