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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?

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  • $\begingroup$ @Karl Sir I thought it was a famous relation from Potential Energy - Reaction Coordinate profile. $\endgroup$ – Tony Stark Aug 31 '20 at 6:56
  • $\begingroup$ see the answer here chemistry.stackexchange.com/questions/139196/… $\endgroup$ – porphyrin Aug 31 '20 at 7:59
  • $\begingroup$ @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. $\endgroup$ – Tony Stark Aug 31 '20 at 8:17
  • $\begingroup$ 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. $\endgroup$ – porphyrin Aug 31 '20 at 9:12
  • $\begingroup$ @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. $\endgroup$ – Tony Stark Aug 31 '20 at 10:36

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