Transcript What is the “rate" of a reaction?

What is the “rate" of a reaction?
• The rate of a reaction is the speed of the
reaction. It is not “how much” of a product
is made, but instead “how quickly” a
reaction takes place.
How can we measure the rate?
• If we consider a reaction
zinc + hydrochloric acid —> zinc chloride + hydrogen
• then there are two possible ways of measuring the rate:
1) measure how quickly one of the products (e.g. the
hydrogen) is made
2) measure how quickly one of the reactants (e.g. the
zinc) disappears
• So we could, for example, measure the volume (in ml.)
of hydrogen made every 10 seconds or the loss in mass
(of the zinc and hydrochloric acid as they change into
hydrogen gas escaping from a beaker) every 10
seconds.
Measuring a rate of reaction
• There are several simple ways of
measuring a reaction rate.
• For example, if a gas was being given off
during a reaction, you could take some
measurements and work out the volume
being given off per second at any
particular time during the reaction.
• A rate of 2 cm3 s-1 is obviously twice as
fast as one of 1 cm3 s-1.
Note: Read cm3 s-1 as "cubic centimetres per second".
Measuring a rate of reaction
• However, for this more formal and mathematical look at
rates of reaction, the rate is usually measured by looking
at how fast the concentration of one of the reactants is
falling at any one time.
• For example, suppose you had a reaction between two
substances A and B. Assume that at least one of them is
in a form where it is sensible to measure its
concentration - for example, in solution or as a gas.
• For this reaction you could measure the rate of the
reaction by finding out how fast the concentration of,
say, A was falling per second.
Measuring a rate of reaction
• You might, for example, find that at
the beginning of the reaction, its
concentration was falling at a rate of
0.0040 mol dm-3 s-1.
• This means that every second the
concentration of A was falling by
0.0040 moles per cubic decimetre.
This rate will decrease during the
reaction as A gets used up.
Note: Read mol dm-3 s-1 as "moles per cubic decimetre
(or litre) per second".
Reactions involving collisions
between two species
Two species (molecule, ion, or free
radical) 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.
ineffective collisions
Reactions usually require
collisions between reactant
molecules or atoms. The
formation of bonds requires
atoms to come close to one
another. New bonds can
form only if the atoms are
close enough together to
share electrons. Some
collisions are not
successful. These are called
ineffective collisions. The
particles simply hit and then
rebound. This animation
illustrates what happens in
an ineffective collision.
effective collisions
Collisions that lead to
products are called
effective collisions.
An effective collision
must happen with a
great enough speed,
energy and force to
break bonds in the
colliding molecules.
The animation
illustrates an effective
collision between two
diatomic molecules.
The two product
molecules formed fly
outwards.
Reactions involving collisions
more than two species
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
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 orientation of collision
Of the collisions shown in the diagram, only collision 1 may
possibly lead on to a reaction.
The orientation of collision
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.
Proper Orientation of the colliding
molecules
The energy of the collision
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
It is the minimum energy required before a
reaction can occur.
You can show this on an energy profile for a
simple over-all exothermic reaction, the energy
profile looks like this:
Activation Energy
Activation Energy
You can think if the particles collide with
less energy than the activation energy,
nothing important happens. They bounce
apart.
Only those collisions which have energies
equal to or greater than the activation
energy result in a reaction.
Activation Energy
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.
Activation Energy
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.
The Maxwell-Boltzmann
Distribution
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 MaxwellBoltzmann Distribution which is a plot of the
number of particles having each particular
energy.
The Maxwell-Boltzmann Distribution
The area under the curve is a measure of the total number of particles
present. The graph only applies to gases, but the conclusions that we
can draw from it can also be applied to reactions involving liquids.
The Maxwell-Boltzmann Distribution
and activation energy
Remember that for a reaction to happen, particles must collide with
energies equal to or greater than the activation energy for the reaction:
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.
You can change the shape of the curve by changing the temperature of
the reaction. You can change the position of the activation energy by
adding a catalyst to the reaction.
THE EFFECT OF SURFACE AREA ON
REACTION RATES
This effect applies to reactions:
involving a solid and a gas, or
a solid and a liquid.
It includes cases where the solid is acting as
a catalyst.
THE EFFECT OF SURFACE AREA ON
REACTION RATES
The more finely divided the solid is, the faster
the reaction happens.
A powdered solid will normally produce a
faster reaction than if the same mass is present
as a single lump.
The powdered solid has a greater surface area
than the single lump.
Why normally? What exceptions
can there be?
Imagine a case of a very fine powder
reacting with a gas. If the powder was in
one big heap, the gas may not be able to
penetrate it.
That means that its effective surface area is
much the same as (or even less than) it
would be if it were present in a single lump.
A small heap of fine magnesium powder
tends to burn rather more slowly than a
strip of magnesium ribbon, for example.
Calcium carbonate and hydrochloric acid
In the lab, powdered calcium carbonate
reacts much faster with dilute hydrochloric
acid than if the same mass was present as
lumps of marble or limestone.
The catalytic decomposition of hydrogen
peroxide
Solid manganese(IV) oxide is often used as the
catalyst. Oxygen is given off much faster if the
catalyst is present as a powder than as the same
mass of granules.
The explanation:
You are only going to get a reaction if the
particles in the gas or liquid collide with the
particles in the solid.
Increasing the surface area of the solid
increases the chances of collision taking place.
Imagine a reaction between magnesium metal
and a dilute acid like hydrochloric acid. The
reaction involves collision between magnesium
atoms and hydrogen ions.
Increasing the number of collisions per second increases the
rate of reaction.
Catalytic converters
Catalytic converters use metals like platinum, palladium and
rhodium to convert poisonous compounds in vehicle
exhausts into less harmful things. For example, a reaction
which removes both carbon monoxide and an oxide of
nitrogen is:
Because the exhaust gases are only in contact with the
catalyst for a very short time, the reactions have to be very
fast. The extremely expensive metals used as the catalyst are
coated as a very thin layer onto a ceramic honeycomb
structure to maximize the surface area.
The effect of particle size
• Solids with a smaller
particle size (e.g.
powders or small
chips) react more
quickly than solids
with a larger particle
size (e.g large chips).
Here is why:
• Look at this diagram
The perimeter of the large chip is 12 units. The acid
particles can only collide with the edge of the chip.
However, if we break up the large chip into 9 smaller
chips:
The effect of particle size
• However, if we break up
the large chip into 9
smaller chips:
• then the perimeter around
each chip is 4 units, but
there are 9 of them so the
total perimeter is 4 x 9 = 36
units. Notice how the acid
in the second diagram can
reach what used to be the
centre of the large chip.
Reducing the size of particles increases the rate of
a reaction because it increases the surface area
available for collisions to take place. This
increases the number of collisions. It has no effect
on the energy of the particles.
THE EFFECT OF CONCENTRATION ON
REACTION RATES
For many reactions involving liquids or gases, increasing
the concentration of the reactants increases the rate of
reaction.
In a few cases, increasing the concentration of one of the
reactants may have little noticeable effect of the rate.
Don't assume that if you double the concentration of one of
the reactants that you will double the rate of the reaction. It
may happen like that, but the relationship may well be more
complicated.
The mathematical relationship between concentration and
rate of reaction is related with the orders of reaction.
Some examples to the concentration
effect:
Zinc and hydrochloric acid
In the lab, zinc granules react fairly slowly with dilute
hydrochloric acid, but much faster if the acid is concentrated.
The catalytic decomposition of hydrogen peroxide
Solid manganese(IV) oxide is often used as a catalyst in this
reaction. Oxygen is given off much faster if the hydrogen
peroxide is concentrated than if it is dilute.
The reaction between sodium thiosulphate
solution and hydrochloric acid
When a dilute acid is added to sodium thiosulphate
solution, a pale yellow precipitate of sulphur is formed.
As the sodium thiosulphate solution is diluted more and
more, the precipitate takes longer and longer to form.
The explanation of the concentration effect:
Collisions involving two particles
 In order for any
reaction to happen,
those particles must
first collide.
 This is true whether
both particles are in
solution, or whether
one is in solution and
the other a solid.
 If the concentration
is higher, the chances
of collision are greater.
The explanation of the concentration effect:
Reactions involving only one particle
If a reaction only involves a single
particle splitting up in some way,
then the number of collisions is
irrelevant.
what matters now is how many of the
particles have enough energy to react at
any one time.
The explanation of the concentration effect:
Reactions involving only one particle
Suppose that at any one time 1 in a million
particles have enough energy to equal or
exceed the activation energy. If you had 100
million particles, 100 of them would react. If
you had 200 million particles in the same
volume, 200 of them would now react. The rate
of reaction has doubled by doubling the
concentration.
Cases where changing the concentration
doesn't affect the rate of the reaction
Suppose you are using a small amount of a solid
catalyst in a reaction, and a high enough
concentration of reactant in solution so that the
catalyst surface was totally cluttered up with
reacting particles.
Increasing the concentration of the solution even
more can't have any effect because the catalyst is
already working at its maximum capacity.
In certain multi-step reactions
Suppose you have a reaction which happens in a series of
small steps. These steps are likely to have widely different
rates - some fast, some slow.
For example, suppose two reactants A and B react
together in these two stages:
The overall rate of the reaction is going to be governed by
how fast A splits up to make X and Y. This is described as
the rate determining step of the reaction.
In certain multi-step reactions
If you increase the concentration
of A, you will increase the
chances of this step happening
for reasons we've looked at
above.
If you increase the concentration
of B, that will undoubtedly speed
up the second step, but that
makes hardly any difference to
the overall rate.
You can picture the second step
as happening so fast already that
as soon as any X is formed, it is
immediately pounced on by B.
That second reaction is already
"waiting around" for the first one
to happen.
In certain multi-step reactions
• The overall rate of reaction isn't
entirely independent of the
concentration of B. If you lowered
its concentration enough, you will
eventually reduce the rate of the
second reaction to the point
where it is similar to the rate of
the first. Both concentrations will
matter if the concentration of B is
low enough.
• However, for ordinary
concentrations, you can say that
(to a good approximation) the
overall rate of reaction is
unaffected by the concentration
of B.
Orders of reaction
• Orders of reaction are always found by
doing experiments. You can't deduce
anything about the order of a reaction
just by looking at the equation for the
reaction.
• So let's suppose that you have done
some experiments to find out what
happens to the rate of a reaction as the
concentration of one of the reactants, A,
changes. Some of the simple things that
you might find are:
Orders of reaction
• One possibility: The rate of reaction is
proportional to the concentration of A
• That means that if you double the
concentration of A, the rate doubles as well.
If you increase the concentration of A by a
factor of 4, the rate goes up 4 times as well.
• You can express this using symbols as:
Writing a formula in square
brackets is a standard way
of showing a concentration
measured in moles per
cubic decimetre (litre).
Orders of reaction
• You can also write this by getting rid of the
proportionality sign and introducing a
constant, k.
Orders of reaction
• Another possibility: The rate of reaction is
proportional to the square of the
concentration of A
• This means that if you doubled the
concentration of A, the rate would go up 4
times (22). If you tripled the concentration
of A, the rate would increase 9 times (32). In
symbol terms:
Orders of reaction
• Generalising this
• By doing experiments involving a reaction between
A and B, you would find that the rate of the reaction
was related to the concentrations of A and B in this
way:
• This is called the rate equation for the reaction.
• The concentrations of A and B have to be raised to
some power to show how they affect the rate of the
reaction. These powers are called the orders of
reaction with respect to A and B.
Orders of reaction
• If the order of reaction with respect
to A is 0 (zero), this means that the
concentration of A doesn't affect the
rate of reaction. Mathematically, any
number raised to the power of zero
(x0) is equal to 1. That means that
that particular term disappears from
the rate equation.
overall order of the reaction
• The overall order of the reaction is found
by adding up the individual orders. For
example, if the reaction is first order with
respect to both A and B (a = 1 and b = 1),
• R= k[A] [B]
• the overall order is 2. We call this an
overall second order reaction.
Some examples
• Each of these examples involves a
reaction between A and B, and each rate
equation comes from doing some
experiments to find out how the
concentrations of A and B affect the rate
of reaction.
Example 1
• Rate = k[A] [B]
•In this case, the order of reaction with
respect to both A and B is 1. The overall
order of reaction is 2 - found by adding up
the individual orders.
Note: Where the order is 1 with respect to one of the
reactants, the "1" isn't written into the equation. [A]
means [A]1.
Example 2
•Rate = k [B]2
• This reaction is zero order with respect
to A because the concentration of A
doesn't affect the rate of the reaction.
• The order with respect to B is 2 - it's a
second order reaction with respect to B.
• The reaction is also second order
overall (because 0 + 2 = 2).
Example 3
Rate = k[A]
• This reaction is first order with
respect to A and zero order with
respect to B, because the
concentration of B doesn't affect the
rate of the reaction.
• The reaction is first order overall
(because 1 + 0 = 1).
Orders of reaction
• What if you have some other number
of reactants?
•It doesn't matter how many reactants there are.
• The concentration of each reactant will occur
in the rate equation, raised to some power.
• Those powers are the individual orders of
reaction.
•The overall order of the reaction is found by
adding them all up.
THE EFFECT OF PRESSURE ON
REACTION RATES
 Up to know we dealt with concentration effect in
the reactions that occure in aqueous solutions.
 Changing the concentration of a gas is achieved
by changing its pressure.
 Increasing the pressure on a reaction involving
reacting gases increases the rate of reaction.
Changing the pressure on a reaction which
involves only solids or liquids has no effect on the
rate.
An example to pressure effect on
reaction rate
• In the manufacture of ammonia by the Haber
Process, the rate of reaction between the
hydrogen and the nitrogen is increased by the
use of very high pressures.
• In fact, the main reason for using high
pressures is to improve the percentage of
ammonia in the equilibrium mixture, but there
is a useful effect on rate of reaction as well.
The explanation of the relationship
between pressure and concentration
• Increasing the pressure of a gas is exactly
the same as increasing its concentration. If
you have a given mass of gas, the way you
increase its pressure is to squeeze it into a
smaller volume. If you have the same mass
in a smaller volume, then its concentration
is higher.
• You can also show this relationship
mathematically if you have come across the
ideal gas equation:
The explanation of the relationship
between pressure and concentration
 Rearranging PV = nRT gives:
 Because "RT" is constant as long as the
temperature is constant, this shows that the
pressure is directly proportional to the
concentration. If you double one, you will also
double the other.
The effect of increasing the pressure on
the rate of reaction
Collisions involving two particles
• The same argument applies whether the
reaction involves collision between two
different particles or two of the same
particle.
• In order for any reaction to happen, those
particles must first collide. This is true
whether both particles are in the gas state,
or whether one is a gas and the other a
solid. If the pressure is higher, the
chances of collision are greater.
The effect of increasing the pressure on
the rate of reaction
Collisions involving two particles
The effect of increasing the pressure on
the rate of reaction
• Reactions involving only one particle
• If a reaction only involves a single
particle splitting up in some way, then
the number of collisions is irrelevant.
What matters now is how many of the
particles have enough energy to react at
any one time.
THE EFFECT OF TEMPERATURE ON
REACTION RATES
• As you increase the temperature the rate of
reaction increases. As a rough approximation,
for many reactions happening at around room
temperature, the rate of reaction doubles for
every 10°C rise in temperature.
• You have to be careful not to take this too
literally. It doesn't apply to all reactions. Even
where it is approximately true, it may be that
the rate doubles every 9°C or 11°C or whatever.
The number of degrees needed to double the
rate will also change gradually as the
temperature increases.
THE EFFECT OF TEMPERATURE
ON REACTION RATES
• Some reactions are virtually instantaneous - for
example, a precipitation reaction involving the
coming together of ions in solution to make an
insoluble solid, or the reaction between
hydrogen ions from an acid and hydroxide ions
from an alkali in solution. So heating one of
these won't make any noticeable difference to
the rate of the reaction.
• Almost any other reaction you care to name
will happen faster if you heat it - either in the
lab, or in industry.
The explanation of the temperature effect
Increasing the collision frequency
• Particles can only react when they collide. If you
heat a substance, the particles move faster and so
collide more frequently. That will speed up the rate
of reaction.
• That seems a fairly straightforward explanation
until you look at the numbers!
• It turns out that the frequency of two-particle
collisions in gases is proportional to the square
root of the kelvin temperature.
The explanation of the temperature effect
Increasing the collision frequency
•If you increase the temperature from 293 K to
303 K (20°C to 30°C), you will increase the
collision frequency by a factor of:
• That's an increase of 1.7% for a 10° rise. The
rate of reaction will probably have doubled for
that increase in temperature - in other words, an
increase of about 100%. The effect of increasing
collision frequency on the rate of the reaction is
very minor. The important effect is quite
different . . .
Temperature effects on rates and
activation energy diagram
• This illustration shows what
happens to an exothermic
reaction when the
temperature is changed.
• The dotted blue curve shows
the energy for a reaction
mixture that is heated. The
reactants are "part way" up
the energy barrier because
they are "hot".
• The dotted magenta curve
shows what cooling does to
the reactant energy. The
energy goes down and the
reaction happens with more
difficulty.
Temperature effects on rates and
activation energy diagram
NOTE: The energies of reactants and products have changed.
They both have different energies because they were either
heated or cooled. The heat of reaction is the slightly different.
The relative amounts of reactants and products are slightly
different because of the temperature changes.
The key importance of activation energy
• 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.
• 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 apart.
The key importance of activation energy
• To speed up the
reaction, you need to
increase the number
of the very energetic
particles - those with
energies equal to or
greater than the
activation energy.
Increasing the
temperature has
exactly that effect - it
changes the shape of
the graph.
The key importance of activation energy
• If you now mark the
position of the
activation energy,
you can see that
although the curve
hasn't moved very
much overall, there
has been such a
large increase in
the number of the
very energetic
particles that many
more now collide
with enough energy
to react.
The key importance of activation energy
• Remember that the
area under a curve
gives a count of the
number of particles.
On the last diagram,
the area under the
higher temperature
curve to the right of
the activation energy
looks to have at least
doubled - therefore at
least doubling the
rate of the reaction.
Summary of the temperature effect
Increasing the
temperature increases
reaction rates because
of the disproportionately
large increase in the
number of high energy
collisions. It is only
these collisions
(possessing at least the
activation energy for the
reaction) which result in
a reaction.
THE EFFECT OF CATALYSTS ON
REACTION RATES
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 of catalysts
reaction
catalyst
Decomposition of
hydrogen peroxide
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
Hydrogenation of a C=C
double bond
vanadium(V)
oxide, V2O5
nickel
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:
Catalysts and activation energy
• 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:
Catalysts and activation energy
Catalysts and activation energy
• 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.
Catalysts and activation energy
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.
Catalysts and activation energy with a
simple analogy.
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.
TYPES OF CATALYSIS
•
heterogeneous and homogeneous
•
autocatalysis - a reaction which is
catalysed by one of its products.
Types of catalytic reactions
• Catalysts can be divided into two main
types
• heterogeneous and homogeneous.
• In a heterogeneous reaction, the catalyst
is in a different phase from the reactants.
• In a homogeneous reaction, the catalyst is
in the same phase as the reactants.
What is a phase?
• If you look at a mixture and can see a
boundary between two of the components,
those substances are in different phases.
• A mixture containing a solid and a liquid
consists of two phases.
• A mixture of various chemicals in a single
solution consists of only one phase, because
you can't see any boundary between them.
What is a phase?
What is a phase?
• You might wonder why
phase differs from the term
physical state (solid, liquid
or gas). It includes solids,
liquids and gases, but is
actually a bit more general.
• It can also apply to two
liquids (oil and water, for
example) which don't
dissolve in each other. You
could see the boundary
between the two liquids.
If you want to be fussy about things, the diagrams actually show more phases
than are labeled. Each, for example, also has the glass beaker as a solid
phase. All probably have a gas above the liquid - that's another phase. We
don't count these extra phases because they aren't a part of the reaction.
Heterogeneous catalysis
• This involves the use of a catalyst in a
different phase from the reactants.
• Typical examples involve a solid catalyst
with the reactants as either liquids or
gases.
Note: It is important that you remember the difference between the two
terms heterogeneous and homogeneous.
hetero implies different (as in heterosexual). Heterogeneous catalysis
has the catalyst in a different phase from the reactants.
homo implies the same (as in homosexual). Homogeneous catalysis
has the catalyst in the same phase as the reactants.
How the heterogeneous catalyst works
• One or more of the reactants are adsorbed on to
the surface of the catalyst at active sites.
• Adsorption is where something sticks to a
surface. It isn't the same as absorption where one
substance is taken up within the structure of
another.
• An active site is a part of the surface which is
particularly good at adsorbing things and helping
them to react.
• There is some sort of interaction between the
surface of the catalyst and the reactant molecules
which makes them more reactive.
• This might involve an actual reaction with the
surface, or some weakening of the bonds in the
attached molecules.
How the heterogeneous catalyst works
• The reaction happens.
• At this stage, both of the reactant molecules
might be attached to the surface, or one might
be attached and hit by the other one moving
freely in the gas or liquid.
• The product molecules are desorbed.
• Desorption simply means that the product
molecules break away. This leaves the active
site available for a new set of molecules to
attach to and react.
How the heterogeneous catalyst works
• A good catalyst needs to adsorb the reactant
molecules strongly enough for them to react,
but not so strongly that the product molecules
stick more or less permanently to the surface.
• Silver, for example, isn't a good catalyst
because it doesn't form strong enough
attachments with reactant molecules. Tungsten,
on the other hand, isn't a good catalyst because
it adsorbs too strongly.
• Metals like platinum and nickel make good
catalysts because they adsorb strongly enough
to hold and activate the reactants, but not so
strongly that the products can't break away.
Examples of heterogeneous catalysis
• The hydrogenation of a carbon-carbon
double bond
• The simplest example of this is the reaction
between ethene and hydrogen in the presence of
a nickel catalyst.
• In practice, this is a pointless reaction, because
you are converting the extremely useful ethene
into the relatively useless ethane. However, the
same reaction will happen with any compound
containing a carbon-carbon double bond.
• One important industrial use is in the
hydrogenation of vegetable oils to make
margarine, which also involves reacting a carboncarbon double bond in the vegetable oil with
hydrogen in the presence of a nickel catalyst
The hydrogenation of a carboncarbon double bond
• Ethene molecules are
adsorbed on the surface of
the nickel.
• The double bond between
the carbon atoms breaks
and the electrons are used
to bond it to the nickel
surface.
The hydrogenation of a carbon-carbon
double bond
• Hydrogen molecules are
also adsorbed on to the
surface of the nickel.
• When this happens, the
hydrogen molecules are
broken into atoms.
• These can move around
on the surface of the
nickel.
The hydrogenation of a carboncarbon double bond
• If a hydrogen atom diffuses
close to one of the bonded
carbons, the bond between
the carbon and the nickel is
replaced by one between the
carbon and hydrogen.
•That end of the original
ethene now breaks free of
the surface, and
•eventually the same thing
will happen at the other
end.
The hydrogenation of a carboncarbon double bond
• As before, one of the
hydrogen atoms
forms a bond with the
carbon, and that end
also breaks free.
• There is now space on
the surface of the
nickel for new
reactant molecules to
go through the whole
process again.
Homogeneous catalysis
• This has the catalyst in the same phase as
the reactants.
• Typically everything will be present as a
gas or contained in a single liquid phase.
Examples of homogeneous catalysis
• The reaction between persulphate ions and iodide
ions
• Persulphate ions (peroxodisulphate ions), S2O82-, are very
powerful oxidising agents. Iodide ions are very easily oxidised to
iodine. And yet the reaction between them in solution in water is
very slow.
• If you look at the equation, it is easy to see why that is:
• The reaction needs a collision between two negative ions.
Repulsion is going to get seriously in the way of that!
• The catalysed reaction avoids that problem completely. The
catalyst can be either iron(II) or iron(III) ions which are added to
the same solution. This is another good example of the use of
transition metal compounds as catalysts because of their ability
to change oxidation state.
The reaction between persulphate ions and
iodide ions
•
•
For the sake of argument, we'll take the catalyst to be iron(II) ions.
As you will see shortly, it doesn't actually matter whether you use
iron(II) or iron(III) ions.
The persulphate ions oxidise the iron(II) ions to iron(III) ions. In the
process the persulphate ions are reduced to sulphate ions.
•
The iron(III) ions are strong enough oxidising agents to oxidise
iodide ions to iodine. In the process, they are reduced back to
iron(II) ions again.
•
Both of these individual stages in the overall reaction involve
collision between positive and negative ions. This will be much
more likely to be successful than collision between two negative
ions in the uncatalysed reaction.
What happens if you use iron(III) ions as the catalyst instead of
iron(II) ions? The reactions simply happen in a different order.
•