In 1884, Jacobus van t Hoff proposed the Van t Hoff equation describing the temperature dependence of the equilibrium constant for a reversible reaction:
where ΔU is the change in internal energy, K is the equilibrium constant of the reaction, R is the universal gas constant, and T is thermodynamic temperature. Based on experimental work, in 1889, Svante Arrhenius proposed a similar expression for the rate constant of a reaction, given as follows:
Integration of this expression leads to the Arrhenius equation
where k is the rate constant. A was referred to as the frequency factor (now called the pre-exponential coefficient), and E is regarded as the activation energy. By the early 20th century many had accepted the Arrhenius equation, but the physical interpretation of A and E remained vague. This led many researchers in chemical kinetics to offer different theories of how chemical reactions occurred in an attempt to relate A and E to the molecular dynamics directly responsible for chemical reactions.
In 1910, French chemist René Marcelin introduced the concept of standard Gibbs energy of activation. His relation can be written as
At about the same time as Marcelin was working on his formulation, Dutch chemists Philip Abraham Kohnstamm, Frans Eppo Cornelis Scheffer, and Wiedold Frans Brandsma introduced standard entropy of activation and the standard enthalpy of activation. They proposed the following rate constant equation
However, the nature of the constant was still unclear.