Group Six
Catalysts and Reaction Rate
A catalyst is a substance that speeds up a chemical reaction without itself being
permanently altered. For example, when a test tube containing solid KClO3 is heated, O2
gas is given off. Another solid, KCl, remains in the tube. The reaction takes place much
more rapidly if solid MnO2 is mixed with the KClO3 before heat is applied. When the
reaction ends and no more oxygen is given off, none of the KClO3 will be left in the tube,
but all of the MnO2 will remain. We can summarize the reaction as follows:
2 KClO3
MnO2, heat
2 KCl + 3 O2
The MnO2 speeds up the reaction, so it must play some part. However, the change it
undergoes is only temporary, because all of it can be recovered in its original form at the
end of the reaction.
Heterogeneous vs. Homogeneous Catalysts
A reaction in which the reactants and catalysts are not in the same phase is a
heterogeneous reaction. The catalyst is called a heterogeneous catalyst. This kind of
catalyst has a surface on which the substances can react. Platinum and other finely
divided materials and metallic oxides are examples of this kind of catalyst. Most
heterogeneous catalysts work by adsorbing (note spelling) one of the reactants.
Adsorption is the sticking of one substance to the surface of another. Catalytic converters
on automobile exhaust systems use a heterogeneous catalyst (Figures 11 and 12).
Figure 11
Figure 12
One of the reactions that takes place in a catalytic converter is the decomposition of NO
to N2 and O2 gas (Figure 13). The platinum-rhodium catalyst provides a surface on which
the NO decomposes. The catalyst surface is large because it coats the surface of a
honeycomb structure seen in the cutaway view of a catalytic converter in Figure 12.
A homogeneous catalyst exists in the same phase as the reactants. This kind of catalyst
does enter into the reaction, but is returned unchanged in the end. It forms an
intermediate compound or compounds that react more readily than the uncatalyzed
reactants. The hydrolysis of sucrose (table sugar) into glucose and fructose is an example:
C12H22O11 (aq) + H2O (l)  C6H12O6 (aq) + C6H12O6 (aq)
(sucrose)
(glucose)
(fructose)
This reaction is normally very slow. If, however, the solution is made acidic, the presence
of the acid causes the reaction to occur readily. In the reaction, all substances are in
aqueous solution (the same phase) and thus it is homogeneous in nature.
Catalysts in Industry
Both homogeneous and heterogeneous catalysts are extremely important in industry. The
presence of a catalyst usually allows the reaction to be carried out at a lower temperature.
For chemical companies this can mean thousands of dollars in annual savings on fuel
bills. At the Cominco Ltd. smelter in Trail, B.C., zinc metal is produced from zinc sulfide
ore. The first step in this new process involves the reaction of zinc sulfide with sulphuric
acid and oxygen gas from the atmosphere.
2 ZnS (s) + 2 H2SO4 (aq) + O2 (g)  2 ZnSO4 (aq) + 2 S (s) + 2 H2O (l)
With pure zinc sulfide this reaction is extremely slow. Fortunately, the ore also contains
some iron (II)sulfide. This reacts in a similar way to the zinc sulfide to produce iron (II)
sulfate:
2 FeS (s) + 2 H2SO4 (aq) + O2 (g)  2 FeSO4 (aq) + 2 S (s) + 2 H2O (l)
The iron (II) sulfate is a catalyst for the reaction of the zinc sulfide. First the iron (II)
sulfate reacts with more sulphuric acid and oxygen to produce iron (III) sulfate. Then,
iron (III) sulfate reacts with zinc sulfide to produce zinc sulfate and return iron to the iron
(II) ion:
4 FeSO4 (aq) + 2 H2SO4 (aq) + O2 (g)  2 Fe2(SO4)3 (aq) + 2 H2O (l)
2 Fe2(SO4)3 (aq) + 2 ZnS (s)  2 ZnSO4 (aq) + 4 FeSO4 (aq) + 2 S (s)
Notice that the iron (II) sulfate DID take part in the reaction sequence: it changed to iron
(III) sulfate, but then it returned to iron (II) sulfate. It is then, a homogeneous catalyst.
One of the most important industrial chemicals is sulphuric acid. This was once produced
by the complex lead chamber process. It has long been known that simpler route would
involve the reaction of SO2 (a common by-product of smelting) with O2 from the air to
produce SO3. Sulfuric acid (H2SO4) could easily be obtained from SO3. But the problem
was the slowness of the reaction between SO2 and O2. A laboratory discovery in 1929
showed that a relatively cheap compound, vanadium (V) oxide (V2O5), increased the rate
of this reaction substantially. Almost all commercial sulphuric acid production now uses
this contact process involving the V2O5 catalyst.
Hydrogenation is another important industrial reaction requiring the use of a catalyst.
This reaction involves adding hydrogen to some of the carbon-carbon double bonds
(C=C) found in unsaturated fats. The process converts these liquid oils into highermelting solid fats, such as those found in margarine.
Uncatalyzed hydrogenation is a slow process. The
overall reaction rate is greatly increased by the use
of a metal catalyst such as platinum (Pt), palladium
(Pd) or nickel (Ni). The metal provides an
adsorption surface for both reactants and acts as a
heterogeneous catalyst (Figure 14). The attractive
forces between the metal and hydrogen weaken the
H-H bonds in H2 molecule and changes the shape of
the ethene (C2H4) molecule. The hydrogen atoms
then attach themselves to the C2H4 molecule to give
ethane (C2H6). The ethane molecule then leaves the
surface of the metal catalyst.
Figure 14
Catalysts are also crucial for the proper operation of all living things. Biological catalysts
called enzymes allow chemical reactions that take place in our body, for example, to
proceed at a faster rate without the addition of heat or altering our delicate internal
environment. The hydrolysis of sucrose into glucose and fructose (noted above) and
thousands of other chemical reactions take place efficiently and effectively as the result
of enzymes. Enzymes are extremely efficient catalysts, increasing the rate of a reaction
by factors varying from 106 – 1012.
The catalytic ability of an enzyme is dependent on a well-defined region of the
enzyme molecule known as the active site. For catalysis to occur, a reactant molecule
must fit into the active site, which is usually a groove or fold in the surface structure of
the enzyme. This explain why each enzyme is specific for catalyzing one specific
reaction or group of reactions. Only the particular reactant molecule for the enzyme
matches the shape of the active site – like a lock and key. Unlike locks and keys
however, neither active sites nor reactants are rigid structures. It is believed that the
active site adjust its’ shape as the two fit together.
Figure 16, Bridge to Biology, explains the role of enzymes in more detail.
Other substances, called inhibitors, are also used to affect reaction rates. The substances
‘tie up’ a reactant or catalytic substance in a complex, so that it will no react. Since they
inhibit the reaction from taking place, they have the opposite effect of a catalyst.
Preservatives such as BHA (butylated hydroxyanisole) and BHT (butylated
hydroxytolune) are used in foods, paint, natural oils, some rubber and plastics, medical
preparations and other products to preserve their freshness.
Demonstration:
Demo: H2O2 - hydrogen peroxide decomposes over time to water and oxygen. Under
normal conditions this is very slow, but with a catalyst occurs much faster.
Safety: Hydrogen peroxide is an oxidant, care should be taken in handling.
Materials:
-
Erlenmeyer flask
Graduated Cylinder
Hydrogen Peroxide (H2O2)
Potassium Iodide (KI)
Large Beaker
Gas collection tube
Scoop
Set up:
- Ask your teacher to demonstrate how to set up the apparatus.
The demo:
- Place about 25 mL of the H2O2 into the Erlenmeyer flask.
- Explain to students that hydrogen peroxide decomposes into water and
oxygen and so we are able to collect the gas
- Set up gas collection apparatus – wait (nothing happens)
- Explain that under normal situations this reaction is very slow.
- To speed it up we can add a catalyst
- Add a small amount (the end of a scoop) of KI to the H2O2 and quickly
replace the gas collection tube.
2H2O2(aq)--> 2H2O(l) + O2(g)
Explanation:
i. The reaction you see is the decomposition of hydrogen peroxide.
ii. Normally this reaction takes years.
iii. The MnO2 is a catalyst.
iv. The catalyst supplies an alternate pathway that does not require as much energy to
overcome; lowers the energy of activation required for the reaction to complete.
v. It is not used in the reaction.
vi. Since a catalyst is not used in the reaction, it can be recovered when the reaction is
complete.
Sample Questions:
1. Define Catalyst
2. How do catalysts work, state the 2 main theories
3. What effect does a catalyst have on the following?
a. Enthalpy change of the reaction
b. Energy of the reactants
c. Pathway of the reaction
4. Why can manganese dioxide (MnO4) be classified as a catalyst when it is
used with potassium chlorate (KClO3) in the preparation of oxygen?
5. The reaction of ethyl acetate (a component of nail polish remover) with
acid can be written as follows:
CH3CO2C2H5 (aq) + H+1 (aq) + H2O (l) 
CH3CO2H (aq) + C2H5OH (aq) + H+1 (aq)
(a) What is the catalyst in this series of reactions?
(b) Is this a homogeneous or heterogeneous catalyst?
6. How does the action of a catalyst differ from that of an inhibitor?
7. (a) What is an enzyme? Give two examples.
(b) What controls the amount of enzymes found in cells?