Download THE COLLISION THEORY OF REACTION RATES

THE COLLISION THEORY OF REACTION RATES
This page describes the collision theory of reaction rates. It concentrates on the key things which decide
whether a particular collision will result in a reaction - in particular, the energy of the collision, and whether or
not the molecules hit each other the right way around (the orientation of the collision).
The individual factors which affect the rate of a reaction (temperature, concentration, and so on) are discussed
on separate pages. You can get at these via the rates of reaction menu - there is a link at the bottom of the
page.
We are going to look in detail at reactions which involve a collision between two species.
Species: This is a useful term which covers any sort of particle you like - molecule, ion, or free radical
Reactions where a single species falls apart in some way are slightly simpler because you won't be involved in
worrying about the orientation of collisions. Reactions involving collisions between more than two species are
going to be extremely uncommon (see below).
Reactions involving collisions between two species
It is pretty obvious that if you have a situation involving two species they can only react together if they come
into contact with each other. They first have to collide, and then they may react.
Why "may react"? It isn't enough for the two species to collide - they have to collide the right way around, and
they have to collide with enough energy for bonds to break.
(The chances of all this happening if your reaction needed a collision involving more than 2 particles are
remote. All three (or more) particles would have to arrive at exactly the same point in space at the same time,
with everything lined up exactly right, and having
enough energy to react. That's not likely to
happen very often!)
The orientation of collision
Consider a simple reaction involving a collision
between two molecules - ethene, CH2=CH2, and
hydrogen chloride, HCl, for example. These react
to give chloroethane.
Diagram 1
As a result of the collision between the two
molecules, the double bond between the two
carbons is converted into a single bond. A
hydrogen atom gets attached to one of the
carbons and a chlorine atom to the other.
The reaction can only happen if the hydrogen end
of the H-Cl bond approaches the carbon-carbon
double bond. Any other collision between the two
molecules doesn't work. The two simply bounce
off each other.
Of the collisions shown in the diagram, only collision 1 may possibly lead on to a reaction.
If you haven't read the page about the mechanism of the reaction, you may wonder why collision 2 won't work
as well. The double bond has a high concentration of negative charge around it due to the electrons in the
bonds. The approaching chlorine atom is also slightly negative because it is more electronegative than
hydrogen. The repulsion simply causes the molecules to bounce off each other.
In any collision involving unsymmetrical species, you would expect that the way they hit each other will be
important in deciding whether or not a reaction happens.
Diagram 2
The energy of the collision
Activation Energy
Even if the species are orientated properly, you still won't
get a reaction unless the particles collide with a certain
minimum energy called the activation energy of the
reaction.
Activation energy is the minimum energy required before
a reaction can occur. You can show this on an energy
profile for the reaction. For a simple over-all exothermic
reaction, the energy profile looks like this:
If the particles collide with less energy than the activation
energy, nothing important happens. They bounce apart.
You can think of the activation energy as a barrier to the reaction. Only those collisions which have energies
equal to or greater than the activation energy result in a reaction.
Any chemical reaction results in the breaking of some bonds (needing energy) and the making of new ones
(releasing energy). Obviously some bonds have to be broken before new ones can be made. Activation energy
is involved in breaking some of the original bonds.
Where collisions are relatively gentle, there isn't enough energy available to start the bond-breaking process,
and so the particles don't react.
The Maxwell-Boltzmann Distribution
Because of the key role of activation energy in
deciding whether a collision will result in a reaction, it
would obviously be useful to know what sort of
proportion of the particles present have high enough
energies to react when they collide.
In any system, the particles present will have a very
wide range of energies. For gases, this can be
shown on a graph called the Maxwell-Boltzmann
Distribution which is a plot of the number of particles
having each particular energy.
Diagram 3
The area under the curve is a measure of the total number of particles present.
The Maxwell-Boltzmann Distribution and activation energy
Diagram 4
Remember that for a reaction to happen,
particles must collide with energies equal to or
greater than the activation energy for the
reaction. We can mark the activation energy on
the Maxwell-Boltzmann distribution:
Notice that the large majority of the particles
don't have enough energy to react when they
collide. To enable them to react we either have to
change the shape of the curve, or move the
activation energy further to the left. This is
described on other pages.
THE EFFECT OF
CATALYSTS ON REACTION
RATES
This page describes and explains the way that adding a catalyst affects the rate of a reaction. It
assumes that you are already familiar with basic ideas about the collision theory of reaction rates, and
with the Maxwell-Boltzmann distribution of molecular energies in a gas.
Note that this is only a preliminary look at catalysis as far as it affects rates of reaction. If you are
looking for more detail, there is a separate section dealing with catalysts which you can access via a
links at the bottom of the page.
The facts
What are catalysts?
A catalyst is a substance which speeds up a reaction, but is chemically unchanged at the end of the
reaction. When the reaction has finished, you would have exactly the same mass of catalyst as you
had at the beginning.
Some examples
Some common examples which you may need for other parts of your syllabus include:
reaction
Decomposition of hydrogen peroxide
catalyst
manganese(IV) oxide,
MnO2
Nitration of benzene
concentrated sulphuric
acid
Manufacture of ammonia by the Haber Process
iron
Conversion of SO2 into SO3 during the Contact
Process to make sulphuric acid
vanadium(V) oxide, V2O5
Hydrogenation of a C=C double bond
nickel
The explanation
The key importance of activation energy
Diagram 5
Collisions only result in a reaction if the
particles collide with enough energy to get the
reaction started. This minimum energy
required is called the activation energy for the
reaction.
You can mark the position of activation energy
on a Maxwell-Boltzmann distribution to get a
diagram like this:
Only those particles represented by the area to
the right of the activation energy will react
when they collide. The great majority don't
have enough energy, and will simply bounce
Diagram 6
apart.
Catalysts and activation energy
To increase the rate of a reaction you need to
increase the number of successful collisions.
One possible way of doing this is to provide an
alternative way for the reaction to happen which
has a lower activation energy.
In other words, to move the activation energy on
the graph like this:
Adding a catalyst has exactly this effect on
activation energy. A catalyst provides an
alternative route for the reaction. That
alternative route has a lower activation energy.
Showing this on an energy profile:
A word of caution!
Be very careful if you are asked about this in an
exam. The correct form of words is
"A catalyst provides an alternative route for
the reaction with a lower activation energy."
It does not "lower the activation energy of the
reaction". There is a subtle difference between the
two statements that is easily illustrated with a simple analogy.
Suppose you have a mountain between two valleys so that the only way for people to get from one
valley to the other is over the mountain. Only the most active people will manage to get from one
valley to the other.
Now suppose a tunnel is cut through the mountain. Many more people will now manage to get from
one valley to the other by this easier route. You could say that the tunnel route has a lower activation
energy than going over the mountain.
But you haven't lowered the mountain! The tunnel has provided an alternative route but hasn't
lowered the original one. The original mountain is still there, and some people will still choose to climb
it.
In the chemistry case, if particles collide with enough energy they can still react in exactly the same
way as if the catalyst wasn't there. It is simply that the majority of particles will react via the easier
catalysed route.
Articles from…
http://www.chemguide.co.uk/physical/basicrates/catalyst.html
http://www.chemguide.co.uk/physical/basicrates/introduction.html
accessed on 2/23/09
AP Chemistry
Chemical Kinetics Introduction
Name ____________________________________
Questions about The Collision Theory of Reaction Rates
1. First a bit of review. What are the four factors that can affect the rate of a chemical reaction?
2. In the middle of the first page, the statement is made that molecules “first have to collide, and then
they may react.” What two things must be true about that collision in order for the molecules to
react?
3. Reactions involving two molecules colliding in order to react are said to be bimolecular. Reactions
involving three molecules colliding in order to react are called termolecular and are exceptionally
rare. Why?
4. Sketch and label diagram 2. It’s kind of an important diagram, so it’s worth knowing it well.
5. Define the term activation energy.
6. What happens if particles collide without sufficient activation energy?
7. Which types of reactions require activation energy to proceed?
8. If we could calculate the average energy of a molecule in diagram 3, what would that value be
called?
A Very Few Questions about The Effect of Catalysts on Reaction Rates
9. Copy and label diagram 5.
10. Adjust your just-drawn diagram to show what happens when a catalyst is added to a reaction.
11. The statement that a catalyst lowers the activation energy of a reaction is incorrect. What is the
more correct but related statement about what a catalyst does?
12. Explain – in your own words, of course – the mountain/tunnel analogy used on the last page of
the article and how it relates to catalysis.
From neither article…
13. Most reactions do not happen in one step. Instead, they happen in multiple steps. For example in
the upper atmosphere, ozone is destroyed in this process.
step 1
step 2
Cl + O3 à ClO + O2
ClO + O à Cl + O2
What is the overall reaction taking place?
14. In the reaction from #13, what substance is the catalyst - the thing that is consumed then remade
at the end?
15. What substance is an intermediate – being made in one step then consumed in a later step?