# Arrhenius Equation : Interpretation

Consider the relation: $$E_{Activation}=E_{Threshold} - E_{Avg}$$

Here the $$E_{Avg}$$ refers to the Potential Energy of the reactants. Now in order of meet $$E_{Threshold}$$ , the molecules must have sufficient Kinetic Energy which is measured in terms of $$E_{Activation}$$.

Hence when we write Arrhenius equation: $$k=Ae^{-\frac {E_{activation}}{RT}}$$

The term $$exp(-E_{activation}/RT)$$ means the number of molecules having Kinetic Energy greater than Activation Energy which coherently measures the minimum Kinetic Energy reactants need to have for reaction

Do my arguments make sense?

• @Karl Sir I thought it was a famous relation from Potential Energy - Reaction Coordinate profile. – Tony Stark Aug 31 '20 at 6:56
• see the answer here chemistry.stackexchange.com/questions/139196/… – porphyrin Aug 31 '20 at 7:59
• @porphyrin Sir I have seen it. I just want to be reassured that whatever I have understood is correct. Please do the same if you feel my understanding is correct. – Tony Stark Aug 31 '20 at 8:17
• The activation energy is $E_T$ which you can split into parts if you want to to include a threshold plus average energy in the transition state, but not as you have written it above. – porphyrin Aug 31 '20 at 9:12
• @porphyrin Sir in the post that you referred to me chemistry.stackexchange.com/questions/139196/… ,the same expression is in the question part as used as in my first expression. I have updated my expression accordingly. – Tony Stark Aug 31 '20 at 10:36