Arrhenius studied the effect of temperature on the rate of chemical reactions and suggested that rate of reactions are exponentially influenced by temperature, according to Arrhenius;
$\mathrm{k}={\mathrm{Ae}}^{-\mathrm{Ea}/\mathrm{RT}}$
Where, k = Rate constant
A = Frequency factor, which is the product of collision frequency and probability factor.
${\mathrm{e}}^{-\mathrm{Ea}/\mathrm{RT}}=$Fraction of molecular collisions giving the reaction.
By applying the above Arrhenius equation, activation energy can be determined by measuring the rate constant at two different temperatures;
$\log \left(\frac{{\mathrm{k}}_2}{{\mathrm{k}}_1}\right)=\frac{\mathrm{Ea}}{2.303\mathrm{R}}(\frac{1}{{\mathrm{T}}_1}-\frac{1}{{\mathrm{T}}_2})$
where k₁ and k₂ are the rate constants at T₁ and T₂ and R is universal gas constant.