Understanding Arrhenius Plots: ln(rate) vs 1/T for Determining Activation Energy

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Arrhenius plots are commonly used to determine activation energy in chemical reactions, typically plotting ln(k) against 1/T. However, some discussions highlight that ln(rate) can also be plotted against 1/T to find activation energy, particularly in heterogeneous catalysis. This raises questions about the terminology, as the rate constant (k) is often referred to simply as "rate." The confusion stems from the fact that the Arrhenius equation, k=A*exp(-Ea/(R*T), applies to the rate constant, not the reaction rate itself. For the equation to hold, the rate must not be constant, suggesting that the use of "rate" in this context may lead to misunderstandings.
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I have been noticing that in some cases that Arrhenius plots are sometimes plotted ln(rate) vs 1/T as opposed to ln(k) vs 1/T to find activation energy especially in hetereogeneous catalysis.

Isn't the Arrhenius law k=A*exp(-Ea/(R*T))? why can ln(rate) also be plotted vs T^-1 to find activation energy?
 
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pwnzorz said:
I have been noticing that in some cases that Arrhenius plots are sometimes plotted ln(rate) vs 1/T as opposed to ln(k) vs 1/T to find activation energy especially in hetereogeneous catalysis.

Isn't the Arrhenius law k=A*exp(-Ea/(R*T))? why can ln(rate) also be plotted vs T^-1 to find activation energy?
Is it possible that some people refer to the rate constant simply as "rate?" Otherwise, the rate would have to be constant for the equation to apply.

Chet
 

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