🤓 Based on our data, we think this question is relevant for Professor Lewis' class at UGA.

We need to use the ** two-point form of the Arrhenius Equation**:

$\overline{){\mathbf{ln}}{\mathbf{}}{\mathbf{k}}{\mathbf{=}}{\mathbf{-}}\frac{{\mathbf{E}}_{\mathbf{a}}}{\mathbf{R}}{\mathbf{}}\left(\frac{\mathbf{1}}{\mathbf{T}}\right){\mathbf{}}{\mathbf{+}}{\mathbf{}}{\mathbf{ln}}{\mathbf{}}{\mathbf{A}}}$

where:

**k** = rate constant

**E _{a}** = activation energy (in J/mol)

**R** = gas constant (8.314 J/mol • K)

**T** = temperature (in K)

**A** = Arrhenius constant or frequency factor

Hydrogen iodide, HI, decomposes in the gas phase to produce hydrogen, H_{2,} and iodine, I_{2}. The value of the rate constant, k, for the reaction was measured at several different temperatures and the data are shown here:

What is the value of the activation energy (in kJ/mol) for this reaction?

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