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The magic of catalysts
23 June 1990
From New Scientist Print Edition.
John Emsley and Frank Hibbert

Industrial chemists need catalysts to make everything from polythene and painkillers, to fertilizers and fabrics. Without these magic ingredients - and their biological equivalents - 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 - compounds of carbon and hydrogen. Chemists would write this reaction:

CH3OH (CH2)n + H2O

methanol 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's 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 "boiling stone" 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 CH3OCH3 + H2O

methanol 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's 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's 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's and don'ts

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's catalysts, enzymes. Life would have evolved very differently if different catalysts had been present in the Earth's 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 - 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 - 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 ClO + O2

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

ClO + O 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's 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's 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's 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 - components of proteins - 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 - 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°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 -CH2CH2CH2CH2CH2CH2-

ethene 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 - all made by the high-pressure method. Today, the world's 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 Mülheim found that triethylaluminium [Al(CH2CH3) 3] and titanium tetrachloride [TiCl4] would catalyse the polymerisation under gentler conditions, 50°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's atmosphere (N2 into ammonia gas (NH3).

(1) N2 + 3H2 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°C. Above this, equilibrium (2) gives a small amount of

(2) H2 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°C does equilibrium (3) release atoms of nitrogen which can react with atoms of hydrogen to form ammonia.

(3) N2 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 - one along a valley and one over Mount Everest. With the catalyst, chemists can produce ammonia at an economical rate with temperatures of 525°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 "poison" 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.