This page supplements the second handout on the
"clock reaction." You should have your copy of the
handout available as you look over this page. If
you need a spare, or just want an electronic (pdf)
version, click the link below.
A "Clock"
Reaction: Temperature effect
MAIN
IDEAS: Activation energy and the effect of
temperature
While it is somewhat intuitive that chemical
reactions are faster at higher temperatures, a real
understanding of why this is true and how to
express "faster" mathematically was a long time
coming (and in fact some reactions are slower at
higher temperature, but they are pretty rare). The
first sucessful theory of the effect of temperature
on reaction rates was put forward in 1889 by the
Swedish physical chemist Svante Arrhenius, who won
the Nobel Prize in Chemistry in 1903. (He won for
his earlier work on ionic dissociation in aqueous
solutions, work that established the existence of
ions in solution. He was also the first to propose
that atmospheric CO2 would have an
effect on the Earth's climate, what we now call the
Greenhouse Effect. You can click on his picture to
read more about him.)
Photo © 1999 The Nobel
Foundation
But to get back to the clock reaction, recall
that we determined in class the reaction rate law
and its rate constant, k, at one temperature.
The effect of temperature on reaction rates is
almost always entirely contained in k. (We say
"almost always," because rates also depend on
concentrations, and concentrations can change with
temperature, too. In solution, the change is way
too small to be important, but in the gas phase, if
we keep the pressure constant, then gas
concentrations will change with temperature.
Remember the Ideal Gas Equation, PV = nRT? Written
as P = (n/V)RT, we see that if P is to stay
constant while T changes, then the concentration,
n/V, must change, too.)
The Arrhenius Equation states
k = A exp(–Ea /
RT)
where A is a constant with the same
dimensions as k called the pre-exponential
factor, Ea is a constant with
dimensions of energy per mole called the
activation energy, and R and T
are the Universal gas constant and the
absolute temperature, respectively. This
equation predicts that a graph of the natural
log of the rate constant plotted against the
reciprocal of the absolute temperature
should be a straight line with a slope given by
slope = –Ea / R
We will measure the rate constant in class at
three temperatures and construct such a graph. Data
from a previous year gave the results shown below.
The data are the blue dots, and the purple line is
there to allow you to measure a slope and, from the
slope, an activation energy. (You will do this
again in the second week of the kinetics lab, so
now is a good time to practice. You may even get a
third chance to interpret such a graph on the first
exam...)
When I estimated the activation energy from this
graph, I got a value around 56 kJ
mol–1. If you are having trouble
getting a value anywhere close to this, see me.
Click
here for
the results, plus some extra commentary [to be
posted after the demo].
Finally, here's a page
that discusses the snowy tree cricket, a
chirping, common cricket that follows the Arrhenius
expression in a macroscopic way!
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