When I obtained the activation energy using the Arrhenius equation for a 10-minute glowstick (rapid, rigorous reaction), it was almost 4 times that of a 12-hour glowstick (slow, mild reaction).

However, I learned in chemistry that a reaction with a lower activation energy need less energy for successful collisions, therefore will have a higher reaction rate. This clearly does not hold true in my case, since a shorter 10-minute glowstick obviously has a higher reaction rate compared to that of a 12-hour glowstick(and thats why it only glows for 10 minutes).

Can anyone provide an explanation to this?

  • $\begingroup$ Well, the activation energy is not the only variable in the Arrhenius equation. $\endgroup$ Jan 16 '20 at 15:57
  • $\begingroup$ I don't know the composition of these glow sticks. But I assume that they cannot be compared, because the chemical reactions may be absolutely different in both sticks. $\endgroup$
    – Maurice
    Jan 16 '20 at 17:26
  • $\begingroup$ Activation energy is definitely always not proportional to the reaction rate, since rates are exponentionally related to energies. "Proportional" implies a linear relationship. $\endgroup$
    – Zhe
    Jan 16 '20 at 19:36

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