Video Transcript
In this video, we will learn what a
catalyst does in terms of its effect on the rate of a chemical reaction. We will also learn about how a
catalyst works and the reasons why catalysts are used in industrial chemical
processes.
Chemical reactions usually happen
when reactor molecules are mixed together and collide with each other. For example, we could mix methane
gas with oxygen gas. The methane and oxygen gas
particles move around freely, and the molecules will collide with each other
frequently. But a chemical reaction does not
occur under normal conditions. In this case, we do not
automatically get combustion or burning taking place. We need to add a spark or a flame
to get the reaction started with a bit of a bang.
On the other hand, some chemical
reactions seem to start all by themselves. White phosphorus is a dangerous
element to leave lying around in the presence of air. It’s normally stored under
water. As it dries out in air, it will
first smolder and then ignite all by itself. It automatically reacts with oxygen
in the air to produce phosphorus oxide. In order to understand what makes
chemical reactions happen, we clearly need to explore the reasons behind why some
reactions start easily all by themselves, whereas others need heating, sometimes to
high temperatures, first.
To get a chemical reaction to
happen, we need reactant particles to collide with each other first. If we mix methane gas and oxygen
gas together, there’s no shortage of collisions between these molecules every second
because they’re both in the gas state. But we also need these particles to
collide with just enough energy to break chemical bonds within these reactant
molecules. In these collisions, the minimum
energy needed to break bonds and cause what we call a successful collision is known
as the activation energy. Activation energy is the minimum
energy needed for a reaction to occur.
In many reactions, including the
reaction of methane gas with oxygen gas, this activation energy is very high. So at normal temperatures, reactive
particles just bounce off each other when they collide. Bonds are not broken, and a
chemical reaction does not take place. Most collisions at normal
temperatures are not successful collisions. In the reaction of methane and
oxygen, the chemical energy stored in the reactants is greater than the chemical
energy stored in the products.
The difference between the energy
level of the reactants and the energy level of the products is called the energy
change for this reaction. The energy change is often seen
labeled as ΔH. This is a large triangle symbol
with a capital H. The energy change for the reaction
of methane with oxygen is negative as heat is released to the surroundings. A reaction where heat is released
to the surroundings is described as an exothermic reaction. All combustion reactions are
exothermic. The reaction of methane with oxygen
is a combustion reaction.
In the reaction of methane with
oxygen, which is an exothermic reaction releasing heat to the surroundings, the
activation energy is high enough to ensure that methane and oxygen cannot
automatically or spontaneously react with each other at normal temperatures. We need to put some additional
energy into the reactant mixture first to provide the activation energy to break
bonds and get the reaction started. This additional quantity of energy
is labeled as the activation energy on our energy level diagram here. The reaction then proceeds
automatically or spontaneously as it releases enough heat energy in the process to
maintain successful collisions.
The situation is similar for an
endothermic reaction, where heat is absorbed from the surroundings and the products
finish at a higher energy level relative to the reactants that we started with. Endothermic reactions, like all
chemical reactions, also require activation energy to get them started. We must raise the energy level of
the reactants above that of the products to provide this activation energy. This is labeled on the energy level
diagram.
So in summary at this point, some
reactions are harder to get started than others because they have a high activation
energy. Some reactions happen very slowly
or not at all under normal conditions because they have a high activation
energy. How can we make these reactions
with a high activation energy happen at all?
Hydrogen peroxide is a liquid
substance that can decompose to make water and oxygen gas. This reaction is very slow under
normal conditions. You wouldn’t see many bubbles of
oxygen gas forming in the liquid hydrogen peroxide. If we add a small amount of a black
powder called manganese dioxide to the hydrogen peroxide, we see a sudden increase
in the number of oxygen bubbles being produced every second. The manganese dioxide is increasing
the rate at which bubbles of oxygen are produced. It is increasing the rate of
reaction very dramatically in this case.
We have gone from a situation where
a few bubbles of oxygen gas are being produced every now and then to many bubbles of
oxygen gas being produced in a very short space of time. The manganese dioxide is behaving
as a catalyst. A catalyst is a substance that
increases the rate of a reaction without itself undergoing a permanent chemical
change. The energy profile diagram for the
decomposition of hydrogen peroxide shows us that the overall energy change is
negative. The decomposition of hydrogen
peroxide is therefore an exothermic reaction. This is the case if the reaction
goes ahead with or without a catalyst.
As with all chemical reactions,
there is an activation energy associated with this decomposition process. If we add a catalyst to the
reactant, we find that the activation energy is much smaller than it is without a
catalyst. Catalysts speed up the rate of a
chemical reaction. They do this by offering an
alternative reaction pathway or route for the reaction to take place. Most importantly, this alternative
reaction pathway has a reduced or lower activation energy. With a lower activation energy, a
much greater percentage of the collisions involving reactant particles will lead to
bonds breaking and a reaction occurring. More successful collisions occur
every second, and the rate of reaction will be increased.
This situation is similar to going
on a journey over Mont Blanc, the highest mountain in France. We could expend a lot of energy
climbing over the top of Mont Blanc and visiting Italy on the other side. This is the high energy route. Alternatively, we could pass
through the tunnel quite easily instead. This would be a lower energy route
and would be much faster.
We now know that catalysts speed up
chemical reactions by providing an alternative reaction pathway with a lower
activation energy. So how will the catalysts become
involved in a chemical reaction? Although catalysts remain
chemically unchanged at the end of the reaction that they are catalyzing and they
can be recovered, it will be incorrect to say that catalysts do not become
chemically involved or changed during the reaction that they are catalyzing.
Catalysts can concentrate reactant
particles on their surface. These reactant particles would be
moving around randomly and occasionally colliding with each other in the
process. By doing this, catalysts weaken
bonds in the reactant particles. This is how the activation energy
for the reaction is reduced. At the end of the reaction,
products are released and the catalyst is chemically unchanged. We could say that it is chemically
regenerated. It’s not used up or consumed as a
reactant normally is during the chemical reaction. For this reason, the overall
equation for the chemical reaction will not usually include the catalyst. The catalyst may be written above
the arrow instead.
As catalysts are not consumed
during a chemical reaction, only a small quantity of catalyst need be used for a
given situation. In a catalytic converter found
within a car exhaust pipe close to the engine, toxic gases such as carbon monoxide
and oxides of nitrogen enter the catalytic converter. These toxic gases are chemically
changed to safer gases such as carbon dioxide and nitrogen, which exit the
tailpipe.
In many situations, such as the
catalytic converter in car exhaust pipes, a small amount of catalyst is spread over
a large surface area. At normal operating temperatures,
it will be very hot. It’s important to remember here
that raising the temperature of a reaction increases the rate much more than using a
larger quantity of catalyst. Both of these factors, surface area
and temperature, help to improve the efficiency of the catalyst whilst keeping the
amount used to a minimum.
Catalysts often contain transition
metals or elements located in the central block of the periodic table. Some of these metals are highly
expensive, such as platinum and palladium, which are used in the catalytic
converter. These metals can, of course, be
recovered at the end of the car’s life. Although some common industrial
catalysts contain precious metals, many catalysts contain more common metals and
they’re relatively cheap.
In the Haber process, ammonia is
produced. Ammonia is an important starting
material for fertilizers, which help to improve crop yields. In the Haber process, nitrogen gas
is reacted with hydrogen gas. Nitrogen and hydrogen do not
normally react with one another unless high temperature and very high pressures are
used. In the Haber process, iron is used
as a catalyst to lower the activation energy and allow the reaction to proceed at
lower temperatures with a reasonable rate of reaction.
The situation is complicated by the
fact that this reaction is reversible. The iron catalyst speeds up the
forward and reverse reaction rates equally, allowing the mixture to reach
equilibrium in the reactor vessel in less time. Some ammonia is obtained in less
time, regardless of the yield for this reversible process. Any ammonia formed in the reactor
vessel is cooled, liquefied, and collected. Unreacted nitrogen and hydrogen are
recycled and returned to the reactor vessel.
Catalysts save vast amounts of
energy in industry by allowing reactions to proceed at lower temperatures with a
reasonable rate. This saves money, in terms of the
expensive energy saved, and time, since the reaction will be taking place at a
faster rate than it would without a catalyst. The economic benefit of using a
catalyst is substantial. And it will often outweigh the cost
of the catalyst.
Now let us look at a question to
test our understanding of catalysts.
Manganese dioxide is used as a
catalyst in the decomposition of hydrogen peroxide to form water and oxygen. Which statement is untrue when
using MnO2 as a catalyst? 2H2O2 aqueous produces 2H2O liquid
plus O2 gas. (A) The oxygen will be formed more
quickly. (B) More oxygen will be
produced. (C) The mass of MnO2 before and
after the reaction will be the same. (D) An alternative reaction pathway
is provided by the catalyst. (E) The catalyst remains unchanged
at the end of the experiment.
In this question, we’re being asked
about how a catalyst, which is manganese dioxide, behaves in a decomposition
reaction. In this decomposition reaction, a
single reactant, the hydrogen peroxide, decomposes to form new products. The products are water and oxygen
gas. This decomposition reaction is very
slow under normal conditions at room temperature. We would not observe many bubbles
of oxygen gas coming from the hydrogen peroxide solution at all. This situation would change rapidly
if a small amount of the solid catalyst, manganese dioxide, were added to the
hydrogen peroxide solution.
Rapid fizzing or effervescence
would be observed as soon as the black powder is added to the hydrogen peroxide. The manganese dioxide catalyst will
increase the rate of decomposition of the hydrogen peroxide. More oxygen gas, seen as bubbles,
will be produced per unit of time as the reaction rate has been increased. Oxygen gas will certainly be
produced more rapidly. This statement is true, so it’s not
the correct answer. Remember, in this question, we’re
looking for an untrue statement.
Notice that the oxygen gas, which
is one of the products, originates from the hydrogen peroxide molecules. According to the balanced equation,
two molecules of hydrogen peroxide are required to produce one molecule of oxygen
gas. If we have a fixed amount of
hydrogen peroxide molecules at the start of the reaction, we can only produce a
fixed amount of oxygen molecules during the decomposition reaction. Adding the manganese dioxide
catalyst does not change the amount of oxygen gas produced. It simply increases the rate of
reaction.
The same amount of oxygen gas is
produced in much less time. The manganese dioxide does not
appear in the overall reaction equation. The amount of oxygen gas obtained
will be the same with or without the catalyst present. This means that the statement that
more oxygen gas is produced is untrue. Therefore, it’s likely to be the
correct answer.
Let us consider the other possible
responses first before committing to this response. We could take one gram of manganese
dioxide catalyst, record its mass, and add it to the hydrogen peroxide at the start
of the reaction. When the reaction is complete, the
beaker will contain only pure liquid water and manganese dioxide as a mixture. The manganese dioxide could be
filtered off, dried, and its mass recorded on an accurate balance. We would find that it is chemically
unchanged and, in fact, the mass recovered will be exactly the same as the mass used
at the start.
The catalyst MnO2 will be unchanged
at the end of the experiment, and the mass of MnO2 before and after the reaction
will be the same. Remember that a catalyst becomes
involved with a chemical reaction. It may be chemically changed during
this process, but it is regenerated at the end of the reaction. Both of these responses are
true. And they are therefore not correct
answers.
The decomposition of hydrogen
peroxide is an exothermic reaction. We can view this in an energy level
diagram. And the energy change for this
reaction is negative. The MnO2 catalyst increases the
rate of the chemical reaction by providing an alternative reaction pathway with a
lower activation energy. By reducing the activation energy
for the reaction, a greater proportion of reactant molecular collisions are
successful. Successful collisions are those
where chemical bonds are broken and they lead to a reaction occurring. Hence, we see more successful
collisions per second and a faster rate of reaction. The final statement that the
catalyst provides an alternative reaction pathway is true and, therefore, not the
correct answer here. We’re looking for a statement that
is untrue.
The statement that more oxygen will
be produced is, therefore, the correct answer to this question.
To summarize, let us review the key
points from this lesson. Catalysts increase the rate at
which products are formed in a chemical reaction. Catalysts increase the rate of
reaction by offering an alternative reaction pathway with a lower activation
energy. Catalysts are not consumed in the
reaction. And catalysts save time and money
in industrial reactions.
