Arrhenius' law is a law that describes the kinetics of a chemical reaction as a function of temperature. It was enunciated by Svante August Arrhenius in 1889. However, this law is an empirical law even though it has been verified experimentally on numerous occasions.
k, being the speed coefficient
T, the temperature in degrees Kelvin (K)
R, the perfect universal gas constant (i.e. 8,314 J.mol-1.K-1).
Ea, the activation energy in joules per mole (J.mol-1)
This law can be simplified as follows when the activation energy is not dependent on temperature.
This law makes it possible to conclude the following points:
The higher the temperature, the greater the kinetics of a reaction. This evolves even exponentially, generally speaking, an increase of 10 K (or 10°C, it is equivalent in this case), can multiply the kinetics of the reaction by 2 or 3.
The lower the activation energy of a chemical reaction, the faster the kinetics of the reaction.
At absolute 0, the kinetics of the reaction is zero and it does not occur.