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Industrial chemists need catalysts to make everything from polythene and
painkillers, to fertilizers and fabrics. Without these magic ingredients 鈥� and
their biological equivalents 鈥� to speed up reactions, chemistry and life would
grind to a halt

METHANOL is a rather boring, colourless, liquid. You could leave it in a
bottle for a hundred years and it would not have changed. And yet chemists can
turn methanol into petrol just by passing it over a porous mineral called a
zeolite. New Zealand makes a large proportion of its motor fuel in exactly
this way.

The zeolite is rather special. It is a compound made up of aluminium,
silicon and oxygen with the code name ZSM-5. It speeds up a reaction that
removes the elements of water (two atoms of hydrogen and one of oxygen) from
the methanol molecules. This reaction leaves behind hydrocarbons 鈥� compounds
of carbon and hydrogen. Chemists would write this reaction:

CH₃OH 鈫� (CH₂)_n + H₂O

methanol 鈫� hydrocarbon + water

But at the end of this reaction, ZSM-5 is unchanged.

How does ZSM-5 bring this about? And how does it manage to keep doing it
time and time again without itself being changed in the process? And why does
methanol not turn spontaneously into hydrocarbons and water, without the
zeolite? The answers to these questions take us into the almost magical world
of catalysis.

Catalysts are an essential part of the chemicals industry. If catalysts did
not exist, many chemical processes would go so slowly as to be uneconomic.
Many catalysts are metal compounds which chemists discovered by chance or by
testing substances that seem to have the right qualities. Today, chemists have
a range of techniques that help them to determine how catalysts work. Once
they know a catalyst鈥檚 secret they may be able to improve it. The ZSM-
5/methanol reaction is one such story.

Zeolites occur naturally in volcanic rocks and clay-like deposits. They are
named after the Greek words for 鈥渂oiling stone鈥� because this is what they
appear to do when heated on a fire. Zeolites have channels running through
them like a sponge. In the natural state, the channels are full of water
molecules. Heating the zeolite drives out the water, leaving behind a highly
potent catalyst. In fact, ZSM-5 is a synthetic zeolite with larger-than-
average pores. They have to be large to prevent the hydrocarbon molecules that
form inside them remaining trapped. Charles Plank and Edward Rosinski of the
US Mobil Oil company were the first to make ZSM-5.

ZSM-5 works like this: when methanol enters the zeolite it first reacts to
form dimethyl ether (methoxymethane). In this reaction two molecules of
methanol come together:

2CH₃OH 鈫� CH₃OCH₃ + H₂O

methanol 鈫� dimethyl ether + water

The dimethyl ether then meets up with a very reactive methyl radical
(CH₃*) which attacks it and adds a CH₂ group to the
ether. The methyl radical comes from methanol but no one knows quite how. All
we know is that methanol loses its hydroxyl (OH) group to the zeolite, leaving
behind a methyl radical.

As more methyl radicals attack, adding more CH₂ groups, the
hydrocarbon chain grows longer. The mixture of compounds that finally emerges
can be refined into petrol and other products.

The story of the ZSM-5 catalyst is an interesting one. It brought together
all kinds of chemists to create a material that was better than the natural
zeolites. At the same time, it solved New Zealand鈥檚 fuel problem. That country
has vast reserves of methane gas, 3.5 trillion cubic feet, but very little
natural oil. Chemists turn this methane into methanol, and from this they make
New Zealand鈥檚 petrol.

Chemical processes, whether those in our bodies, for example the regulation
of energy, or those in industry, for example the making of ammonia, have two
fundamental limits. One is the extent to which a reaction occurs at all, the
position of equilibrium. The other is the rate at which it reaches this
equilibrium. Catalysts increase the rate of some reactions that would
otherwise go along infinitesimally slowly. But a catalyst will not change the
position of equilibrium between reactant and product. If the balance of a
reaction lies heavily in favour of the reactant we also need to find some way
of tipping the balance towards the products.

Catalytic do鈥檚 and don鈥檛s

Across the energy barrier

CATALYSTS are very important. Life would be much slower, or may not even
happen at all, without the help of nature鈥檚 catalysts, enzymes. Life would
have evolved very differently if different catalysts had been present in the
Earth鈥檚 primative oceans. Industrial processes and our material comfort would
be different if chemists had not happened upon the particular catalysts they
did. One fascinating thing about catalysts is that the same chemical
principles apply to all of them. Catalysis by single chemical elements (such
as metals) or ions (such as hydrogen) has many features in common with
catalysis by the complex enzyme macromolecules in nature. The first principle
of catalysis is that after conversion of reactant to product, the catalyst
must be regenerated. A second principle is that the catalyst must speed up the
rate at which the reaction reaches equilibrium.

One way of looking at a chemical reaction is to imagine it as a pathway
over a hill from reactant to product. Before the reactants become products,
they must cross an energy barrier. The top of the hill or energy barrier is
called a transition state. It is neither one thing nor the other 鈥� some old
bonds are breaking and other new ones are being formed at this stage.

The energy we are talking about is called the Gibbs energy. How large the
Gibbs energy is depends partly on the complexity of the molecule.

Before a molecule can react, it needs enough Gibbs energy to climb the
barrier to the transition state. At normal temperatures only a minute fraction
of the reactant molecules have enough energy. At a higher temperature, a
greater fraction of the molecules will have this energy. This is why heating
is one way to speed up a reaction.

The position of the chemical equilibrium is determined by the difference in
Gibbs energy between reactant and product. If the products are of much higher
Gibbs energy than the reactant the equilibrium will favour the reactant. But
if the products are of lower energy, the equilibrium favours the products.
Then the reaction is an overall downhill journey from reactant to product,
even then the molecules must climb over an energy barrier on the way. A
catalyst cannot alter the Gibbs energy of the reactants and products, so it
cannot alter the equilibrium. What it can do is to provide an easier pathway
between them.

The rate of a chemical reaction depends on the height of the energy
barrier. The height of the energy barrier is the difference between the Gibbs
energies of the transition state and the reactant. The larger this difference,
the slower the reaction will be. The catalysed reaction goes faster than the
uncatalysed one because the highest energy point on the catalysed pathway is
easy to climb over.

Chemists divide catalysts into two types: homogeneous and heterogeneous.
Homogeneous catalysts are in the same state of matter as the reactants. For
example, in the production of polythene the reactant (ethene) and the catalyst
are both dissolved in a solvent, so they are both in the same state 鈥� liquid.
When we talk about the destruction of the ozone layer in the upper atmosphere
we are talking about homogeneous catalysis. This time, in the gaseous state
chlorine atoms, Cl, are the catalysts.

The source of chlorine atoms is decomposition of chlorofluorocarbons by
sunlight. Chlorine atoms destroy ozone, O₃:

O₃ + Cl 鈫� ClO + O₂

and are regenerated by reaction of oxygen atoms, to repeat the cycle:

ClO + O 鈫� Cl + O₂

In this way, one chlorine atom can destroy many ozone molecules.

In heterogeneous catalysis, the catalyst and the reactants are in different
states. For example, the catalyst may be a solid, and the reactants gases or
liquids. Since catalysis takes place on the solid catalyst, its surface area
must be as large as possible. In the conversion of methanol to petrol, the
reactant (methanol) is a liquid and the zeolite catalyst is a solid.

Catalytic convertors that clean up the exhaust gases of motors contain a
thin coating of rhodium and platinum metals on a solid honeycomb support. Full
of tiny pores, these catalysts offer a large surface area. They turn obnoxious
mixtures of unburnt fuel, carbon monoxide, nitrogen oxides and air, into
carbon dioxide, water, and nitrogen.

Catalysts are used in strong glues such as epoxy resins. A base catalyses
the reaction that links two different molecules into chains and makes the glue
set.

Industrial chemists prefer to use heterogeneous catalysts because these do
not need to be separated from products of the reaction. They make sulphuric
acid, the manufacturing world鈥檚 most important chemical, from sulphur trioxide
(SO₃). Sulphur trioxide is produced from sulphur dioxide and oxygen
in a reaction catalysed by vanadium pentoxide and known as the contact
process.

With a reaction catalysed by iron, and known as the Haber process,
manufacturers can produce another bulk chemical, ammonia, from hydrogen and
nitrogen gases. With ammonia and oxygen, they can make nitric acid using a
catalyst made from platinum and rhodium. Industrial chemists use methanol as
the starting material for many products from headache cures to Perspex, as
well as for the ZSM-5 process. To make methanol from CO₂ and
H₂, they use a copper and zinc oxide catalyst. Every year the
chemicals industry throughout the world produces more than 10 million tonnes
of each of these chemicals.

Catalysts are big business.

Nature鈥檚 catalysts

Enzymes do it better

LIFE has evolved in such a way that the chemical reactions that go on in
biological systems are catalysed by proteins called enzymes. A particular
enzyme will often catalyse only one reaction. This is known as enzyme
specificity. The living cell is an incredibly complex mixture of chemicals. An
enzyme has to pick out a particular reaction. This may be one step in a
sequence of reactions which make up an important biological process. A typical
chemical catalyst, such as an acid, would not be this specific. It would
catalyse a whole range of reactions, and thus cause havoc to the processes of
life.

Enzymes also work under extremely mild conditions. Chemical catalysts often
work only under severe conditions such as high temperatures and pressures. For
nature鈥檚 reactions, chemical catalysts cannot match the enzyme in the
magnitude and skill of its catalytic effect.

Enzymes have evolved as huge protein molecules made up of many thousands of
atoms and some metal ions. Many enzymes possess a cleft or pocket into which
the reactant molecule is manoeuvred. It is here that it is converted to a
product molecule. Enzymes hold the reactant in the pocket in various ways.
These include hydrogen bonds and electrostatic forces between groups of atoms
with opposite charges. As the reactant is transformed through a series of
steps into the product, these interactions change so that after each step, the
molecule, is stabilised. These multiple interactions are what make enzymes so
specific.

The enzyme thermolysin catalyses the break up of peptides 鈥� components of
proteins 鈥� in bacterial cells. The enzyme is a sort of test tube containing
several chemicals that take part in the reaction. The zinc ion of the enzyme
attracts the oxygen of the carbonyl (C=O) group of the peptide. This makes it
more likely to be attacked by water. Acids and bases also contribute at
various stages. Today, biologists and chemists are cooperating in many
exciting experiments to find out how enzymes work.

Simplest chemical catalyst

ACIDS catalyse chemical reactions in solution by supplying hydrogen ions,
H⁺. Meat or fish can be tenderised by acids such as lemon
juice because the acids help to break down the peptide chains that make up
proteins 鈥� a process called hydrolysis. Chemically, this reaction is the
attack of water on an amide group (-CONH-). Adding a hydrogen ion from an acid
to a peptide accelerates the reaction with water. The H⁺
places some of its positive charge on the oxygen to which it is attached. Then
the oxygen off-loads some on to the carbon atom next to it in the peptide
chain. This carbon is more easily attacked by water because of its partial
positive charge. The peptide link (C-N bond) breaks and the H⁺
is free to find another -CONH- group.

Polythene story

On 24 March 1933, Reginald Gibson and Eric Fawcett first prepared
polythene. They polymerised ethene gas to a colourless wax by heating it at
170掳C under a pressure of 1900 atmospheres (190 Mpascals). The polymer was
very stable and was a good insulator. Still, their method of preparation was
not an economical proposition. The chemical reaction is:

nCH₂ = CH₂ 鈫� -CH₂CH₂CH₂CH₂CH₂CH₂鈥�

ethene 鈫� polythene chain

The work continued into 1935, sometimes with explosive results. The process
became less dangerous when chemists realised that a trace of oxygen gas in the
ethene acted as a homogeneous catalyst for the reaction. Oxygen reacts with
ethene to form a free radical, which then can attack another ethene, and so
on.

Polythene played a vital part in the success of airborne radar. By 1945
production had reached 5000 tonnes 鈥� all made by the high-pressure method.
Today, the world鈥檚 chemicals industries produce more than 20 million tonnes
every year, most of it made at much lower pressures, thanks to the discovery
of metal catalysts. Again, the reaction is homogeneous. It takes place in a
solvent in which both the ethene and the catalyst dissove.

In the early 1950s, Karl Ziegler of the Max Planck Institute at
M眉lheim found that triethylaluminium [Al(CH₂CH₃)
₃] and titanium tetrachloride [TiCl₄] would catalyse the
polymerisation under gentler conditions, 50掳C and 10 atmospheres (1
Mpascal) pressure. The Italian chemist Giulo Natta also added to the
development of these catalysts which today are called Ziegler-Natta catalysts
in their honour. In 1963, they won the Nobel Prize for Chemistry for their
work. The polymers produced by these methods have a higher density and are
stiffer than those produced by the high-pressure method. This high-density
polythene is better for such things as buckets and crates.

The ammonia story

CHEMISTS use the equilibrium (1) to turn nitrogen from the Earth鈥檚
atmosphere (N₂ into ammonia gas (NH₃).

(1) N₂ + 3H₂ &rlarr2; 2NH₃

Ammonia is the raw material for a large number of other useful chemical
products such as fertilizers, plastics, drugs and explosives. But a mixture of
nitrogen and hydrogen gases will not react even at a temperature of
1000掳C. Above this, equilibrium (2) gives a small amount of

(2) H₂ &rlarr2; H + H

atomic hydrogen (H), which can then react with nitrogen But even this
produces little ammonia because the reaction of hydrogen atoms with N
₂ is very slow. Only when the temperature reaches 3000掳C does
equilibrium (3) release atoms of nitrogen which can react with atoms of
hydrogen to form ammonia.

(3) N₂ &rlarr2; N + N

To make ammonia at lower temperatures, chemists need a catalyst that will
allow (1) to reach equilibrium rapidly. That catalyst is iron. But the
catalyst cannot change the unfavourable equilibrium (1) which prefers the
nitrogen and hydrogen side of the equation at high temperature. Chemists use
the trick of removing the ammonia as it is formed, which unbalances the
equilibrium so that more nitrogen and hydrogen atoms react.

One of the most important things about a heterogeneous catalyst is its
surface area. The larger this is the better. If you divide a material into
many pieces, its total surface area increases. One gram of the powdered iron
catalyst has an area of 50 square metres, which means that in a typical
ammonia converter the total amount of catalyst present may have a surface area
of 5000 square kilometres. Molecules of H₂ and N₂ stick
onto the surface of the iron catalyst by a process called adsorption. Here
they split up into atoms (N and H) which then recombine as NH₃. The
iron surface stabilises the atoms of N and H.

The dramatic effect of the catalyst in the ammonia process can be seen by
comparing the energy pathways for catalysed and uncatalysed reactions. The
difference is like going from Tibet to Nepal by two routes 鈥� one along a
valley and one over Mount Everest. With the catalyst, chemists can produce
ammonia at an economical rate with temperatures of 525掳C and a pressure of
20 atmospheres (2Mpascals). To make one gram of ammonia under these
conditions, but without using a catalyst, we would need a reactor 10 times the
size of the Solar System.

A typical industrial plant for manufacturing ammonia can make 1000 tonnes a
day. There are about 600 plants worldwide, manufacturing over 100 million
tonnes a year. A German chemist, Fritz Haber, was first to show that
equilibrium (1) could be used to make ammonia. His compatriot Carl Bosch, a
chemical engineer, showed that it would work on an industrial scale, The first
plant opened in Germany in 1913. By the next year an isolated Germany was able
to make good use of the Haber process in producing wartime explosives.

Chemists tried more than 2500 combinations of metals and metal oxides as
catalysts for the Haber process. Even though they found iron to be best, it
needed promoters to make it work effectively. To make the catalyst, chemists
melt together magnetite (Fe₃O₄) with a few per cent each
of potassium, calcium and aluminium oxides. Then they grind this mixture to a
fine powder and reduce the magnetite to iron by heating it in synthesis gas.
The result is a catalyst with a very porous surface. The promoters prevent
crystals of iron from forming, which would reduce the surface area.

Some catalysts have a limited working life. The iron catalysts that
manufacturers use to make ammonia last for between 5 and 10 years. The time
depends on whether or not the catalyst encounters substances that 鈥減oison鈥� the
surface of the catalyst. In the ammonia reaction sulphur presents the greatest
danger. This comes mainly from traces of hydrogen sulphide that are in the
natural gas used as the source of hydrogen.

Further reading

For a practical guide to catalysis, Catalyst Handbook, 2nd edition, edited
by M.V. Twigg (Wolfe Publishing Ltd., 1989) is excellent. Homogeneous
Transition-metal Catalysis 鈥� a Gentle Art, by C. Masters (Chapman and Hall,
1981) is a readable introduction to the subject from a chemist鈥檚 point of
view. Catalysis at Surfaces, by I.M. Campbell (Chapman and Hall, 1988) is its
heterogeneous equivalent. An interesting introduction to zeolites is 鈥淭he
zeolite cage structure鈥�, by J.M. Newsam, in the journal Science (volume 231,
p1093, 1989). Enzyme Structure and Mechanism, by A. Fersht (W.G. Freeman &
Co., 1985) reviews both enzyme catalysis and the catalysis of organic
reactions.