Editing Arrhenius equation (section)

== Arrhenius plot ==
{{main|Arrhenius plot}}

[[Image:Arrhenius_plot_with_break_in_y-axis_to_show_intercept.svg|thumb|Arrhenius linear plot: ln&nbsp;''k'' against 1/''T''.]]
Taking the [[natural logarithm]] of Arrhenius equation yields:
<math display="block">\ln k= \ln A - \frac{E_\text{a}}{R} \frac{1}{T}.</math>

Rearranging yields:
<math display="block">\ln k = \frac{-E_\text{a}}{R}\left(\frac{1}{T}\right) + \ln A.</math>

This has the same form as an equation for a straight line:
<math display="block">y = a x + b,</math>
where ''x'' is the [[Multiplicative inverse|reciprocal]] of ''T''.

So, when a reaction has a rate constant obeying the Arrhenius equation, a plot of ln&nbsp;''k'' versus ''T''<sup>−1</sup> gives a straight line, whose slope and intercept can be used to determine ''E''<sub>a</sub> and ''A'' respectively. This procedure is common in experimental chemical kinetics. The activation energy is simply obtained by multiplying by (−''R'') the slope of the straight line drawn from a plot of ln&nbsp;''k'' versus (1/''T''):
<math display="block">E_\text{a} \equiv -R \left[ \frac{\partial \ln k}{\partial (1/T)} \right]_P.</math>