THE ROLE OF CATALYSTS IN MANUFACTURE CHEMICAL

By GEORGE W. BRIDGER

Research Department, Imperial Chemical Industries, Billingham. The Certificated Engineer May 1965.

World consumption of 120 000 tons a year may seem a great deal for substances which one is taught emerge unchanged from the chemical reaction. The author explains why these curious substances are so vital in industrial operations.

Reproduced from New Scientist (Vol. 24, No. 412 - 8th October 1964) with kind permission of the author and the publishers.

A catalyst is involved at some stage or other in the manufacture of almost all chemical products today, from margarine to artificial textiles, from chemical fertilisers to plastics and petrol. Most industrial catalysts are man-made materials but some, such as enzymes, also occur in nature, and play an important part in biochemical processes. But in this article I shall discuss neither these natural catalysts nor those which act in solution (the so-called 'homogeneous' catalysts). What I do propose to discuss in the field to which most industrial catalysts belong, that of 'heterogeneous' catalysts, where a solid assists in converting gaseous or liquid reactants into a required product.

At ordinary temperatures and pressures many pairs of substances-for example, petroleum hydrocarbons and steam, or nitrogen and hydrogen-show little inclination to combine. If we raise the temperature and perhaps apply extremely high pressure too, they can often be persuaded to unite, but these methods are usually prohibitively expensive on a large scale. However, by introducing a third substance, the catalyst, which may be as ordinary as metallic nickel or iron, the chemical reaction can often be induced to take place at reasonably low temperatures and pressures. Normally, only small amounts of a catalyst are needed because it is not consumed in the reaction (and therefore does not appear in the product-unless something goes wrong).

Just how important catalysis has become to the chemical industry can best be illustrated by means of figures. The world consumption of solid catalysts is about 120 000 tons a year, valued at ￡40 million pretty imposing figures for something 'not consumed in the reaction.' The petroleum industry alone uses 100 000 tons of catalysts, valued at ￡25 million.

What they can do. In extreme cases catalysts appear to make reactions take place which otherwise would not occur at all. The fact is, however, catalysts do not achieve the impossible, and they break no chemical or thermodynamic laws. They function simply by making chemical reactions go faster. Reactions that appear not to take place at ordinary temperatures in the absence of a catalyst do, in fast, occur-but so slowly that no change is detectable in a reasonable time. With a catalyst present, however, the rate of combination is so increased that the reaction can be observed to take place.

Time is the crucial factor, of course. A catalyst takes a reaction near to completion at technically acceptable temperatures and pressures, in a time that is commercially attractive. Figure 1 shows how far one sort of chemical reaction can go at different temperatures. The line AB represents completion; it is the conversion of reactants which, whether a catalyst is present or not, could theoretically be obtained at various temperatures given infinite time; that is, the equilibrium conversion. The other lines show the conversion that is obtained in a realistic time, with and without a catalyst. Plainly the catalyst initiates and completes the reaction at a lower temperature where, in this particular example, a higher yield is also obtained.

Catalysts are found to be very specific in their functions, each being suited to a particular type of reaction, or even to a specific reaction. The iron catalyst used for the synthesis of ammonia cannot be used in place of the vanadium catalyst which catalyses the oxidation of sulphur dioxide. This selective behaviour can be turned to an advantage. often the reactants can form several different products, but by choosing an appropriate catalyst the desired product can be obtained at the expense of the others.

For a particular reaction and product the right chemical composition has to be found for the catalyst, and this is complicated by the fact that impure substances are often better catalysts than pure ones. The 'impurities,' usually added deliberately in amounts up to perhaps 5 or 10 percent, are called 'promoters.' Some catalysts contain more than one promoter, and t present the best ones are found only by trial and error.

Catalysis is a property of the solid surface and the effectiveness of some catalysts is directly related to the area of their surface. Many catalysts are highly porous and their surface area may be very large 100 square meters a gram (40 square miles a ton) is a common figure. Careful control during the manufacture of the catalyst and in some cases the addition of a promoter is needed to secure so large a surface. Sometimes the catalytic material itself is present only in small amounts dispersed throughout a non-catalytic solid of high surface areas, such as carbon, which known as a support or a 'carrier.'

Catalysts might be expected to last forever because they are not consumed in the process, but in practice they eventually cease to function. Some will last for several years, others only for a few weeks. A catalyst may simply disintegrate, or high temperatures may slowly reduce the surface area (a process which can be retarded by the addition of promoters).

Catalysts or also very sensitive to impurities, some of which - the promoters - are beneficial whereas others are detrimental. The latter, known as' 'poisons,' can include such elements as sulphur, chlorine, and arsenic, when associated with certain catalysts. Obviously the 'poisons' must be eliminated from such a catalyst during its manufacture and removed from the reactants before these come into contact with it.

Another cause of failure is the blocking-up of the fine pores with a by-product of the main reaction. Here, the obstruction can often be removed for example, by roasting It in air, in the case of carbon contamination.

How they work. Having considered what catalysts are and what they do, let us now see how they function. That theory lags behind the practice in industrial catalysis is no reflection on the fundamental research effort, which has been considered in the last 20-30 years. Rather it is indicative of the complexity of catalyst systems.

We have seen that catalysis is brought about by a solid surface. The atoms at the surface of any solid are in a special situation, for they are not completely surrounded by other similar atoms; the surface, therefore, has properties very different from those of the solid in bulk.

In order to react, molecules must come within a certain distance of each other, and to do so they need to possess an energy that exceeds a critical value, the so-called 'activation energy.' If its value is high, as it may be in an uncatalysed reaction, very few molecules will react; then we have the reaction I referred to earlier which 'does not take place.'

According to what we might call the classical theories of catalysis, a catalyst lowers this energy barrier by providing the reactants with an easier reaction path. The reaction can now take place in steps, each of which involves activation energy lower than that of the uncatalysed reaction. More molecules now have the necessary energy, and so more participate; in other words, the reaction rate has been increased. The sequence the reaction follows is probably, first, adsorption of the reactant molecules on to the catalyst, where they combine or react with other gaseous molecules to form the desired molecules, which then leave the surface, allowing the cycle to be repeated (Fig. 2).

Many elegant experiments have demonstrated conclusively that adsorption of one or all of the reactants is an important part of catalysis. The adsorption is strong, involving chemical bonds, but it takes place more readily on some parts of the catalyst than others. This discovery has led to the concept of 'active' centers-areas which are held responsible for catalytic activity even though they represent only a very small proportion of the total surface of a catalyst.

The sensitivity of catalysts to small amounts of poisons is evoked as evidence for the existence of these active centers, which are believed to be associated with defects or irregularities in the surface of the catalyst's crystal lattice. At these dislocations or corners in the lattice the valency or combining forces of the surface atoms will be available for strong adsorption of the reactants. The defects in the crystal structure may arise through the presence of impurities -which explains, of course, the action of some promoters.

This picture is relevant for many kinds of catalysts and, although it has been faulted by recent experiments, it still forms a good foundation for modern theories. It can be developed to take into consideration the geometric arrangement of the atoms in the solid surface. Some reactions are only catalysed by metals having an atomic spacing that corresponds to the inter-atomic distances of one of the reacting molecules. Adsorption of the molecule is thus assisted: and such a catalyst will be specific to reactions involving that particular reactant. The geometric theory can be pushed too far, but it leaves us with the valuable concept that a reactant molecule can actually be positioned on the catalytic surface, and thus favourably located in relation to neighbouring adsorbed molecules.

The modern theories, developed since the Second World War, which consider the electronic structure of the surface atoms do not conflict basically with the classical ideas of active centers and surface geometry. Rather they attempt to explain their assumptions. There is undoubtedly s me relationship between catalytic activity and a substance's electronic properties. The availability of electrons at the solid surface or the ability of the surface to receive electrons clearly must affect catalytic activity because the adsorption of the reactants, which involves chemical bonds and electrons, is a vital step in catalysis.

Electrical conductivity is influenced similarly by this same availability of electrons" and we find that catalysts can be classified very broadly by their conductivity. The catalytic metals (such as iron, nickel, and platinum) which are good conductors are often found to catalyse reactions involving hydrogen. Many of these metals have similarities in their electronic structure. Some oxides and sulphides, such as the oxides of nickel, zinc, and copper, classified as semi-conductors because they have a small conductivity, catalyse reactions involving oxygen. Other oxides such as silica, alumina, and magnesia are electrical insulators, and these catalyse reactions involving water.

For completeness I should mention a fourth group of oxides which lies outside this electronic classification, the so-called 'acid' catalysts. They include the alumina-silica complexes which catalyse reactions involving hydrocarbon molecules by breaking up or 'cracking' them, or by transferring parts of the molecules, reactions known as isomerisation and polymerisation. How they are used. Let us now consider briefly the use of catalysts in industry. In developing a new process all of the properties and limitations of catalysts have to be borne in mind. In the search for a new catalyst one is guided both by experience and by the theories which we have discussed. Many laboratory and large-scale tests will be made to single out the 'best' catalyst, and to establish the rate of reaction and the yield we can expect, so that the amount of catalyst, the size of the reactor, the process conditions, and so on, can be settled.

The shape and size of the reactor can vary enormously in practice, depending on such considerations as to whether the reaction is fast or slow and whether it evolves heat. Thus catalysts can be used as wide, thin layers, as deep beds, or as a solid suspended in a liquid reactant. The catalysts themselves, too, vary widely in shape and size from large irregular lumps or shaped pellets, pieces, granules, or fine powders to strips of metal and woven wire gauzes. A few of the catalysts used in chemical engineering are shown in Fig. 3. They include catalysts which are milestones in the progress of industrial catalysis and which illustrate some of the features I have discussed.

One of the first large-scale applications of catalysis was in the oxidation of sulphur dioxide in the manufacture of sulphuric acid, an extremely important 'raw material' for many other manufacturing processes. A vanadium catalyst is now employed as it is less sensitive to poisons than the platinum catalyst used when the process was first introduced around 1900. The vanadium catalyst, which is in the form of cylindrical pellets (Fig. 3), contains potassium to improve its activity and silica which functions mainly as a mechanically strong carrier.

In 1908 the process for the catalytic oxidation of ammonia for the production of nitric acid was brought into use. The reaction is very fast, being complete in about one-hundredth of a second so that it needs only a thin layer of platinum-rhodium wire gauzes as a catalyst. Another rapid reaction, the oxidation of methanol to formaldehyde, uses a layer of silver granules (Fig. 3).

The next milestone was in 1913 when the high-pressure synthesis of ammonia was introduced. It was the first process to be evolved by the proper application of thermodynamic principles, and the catalyst was the first to demonstrate the value of promoters. The catalyst is basically unchanged to this day-iron containing several promoters such as potassium, alumina, magnesia, and other oxides, each having its own job to do; one increases the surface area, another enhances the activity of the iron surface, and others protect it from gaseous impurities. It is one of the few catalysts made by fusing the ingredients together, producing the hard pieces shown in Fig. 3. The converter in which it is used as a deep bed is a heavy steel forging.

The high-pressure process for the synthesis of methanol (methyl alcohol) from hydrogen and carbon monoxide followed in 1923, and used a pelleted catalyst (Fig. 3) of zinc oxide and chromium oxide; the latter helped to preserve the surface area of the catalyst. The same two reactants, hydrogen and carbon monoxide, can be made to form hydrocarbons instead of methanol by operating at a lower pressure and using a different catalyst, such as cobalt-containing other oxide promoters. This process, the Fischer-Tropsch process, came into use in the 1930s.

At much the same time the first plants were in tabled to react methane with steam to produce hydrogen. The heat has to be supplied, and the catalyst is used at 700°-800°C in heated tubes. The catalyst here is nickel supported on a carrier of other oxides such as alumina and magnesia and is commonly made in the form of rings (Fig. 3) to keep the resistance to the flow of reactants down to a minimum. This process is the forerunner of the modern one for the manufacture of hydrogen in which the higher petroleum hydrocarbons such as naphtha, which are more readily available than methane in Britain, are reacted with steam at high pressure.

The year 1935 saw the introduction of the first catalytic process for petroleum refining, with the introduction of catalytic 'cracking' for the production of petroleum hydrocarbons from crude oils. The catalyst is an alumina-silica mixture, a synthetic version of the natural clays which were used for the earliest processes. Carbon forms on the catalyst and has to be burnt off periodically in a 'regenerator.' The catalyst is usually in the form of beads, shown in Fig. 3, which are shuttled back and forth between the reactor and the regenerator continuously. Sometimes the beads are as small as one-tenth of a millimeter in diameter. Another big advance came in the 1940s with the addition of platinum to this alumina-silica catalyst; this extended its effect on the crude oils and produced higher-grade petrol's.

I have mentioned but a handful of the catalysts which are now employed by the chemical industry. The average cost of these catalysts is about ￡350 per ton, a comparatively high figure which reflects the value of the knowledge and attention to detail which goes into their manufacture more than the intrinsic cost of the materials. The amount of these expensive catalysts which are used may seem considerable, but 120 000 tons is very small in comparison with the millions of tons of chemicals for the manufacture of which they are responsible. Catalysts have played no small part in the rapid expansion of the chemical manufacturing industry in the first half of the 20th century. They can truly be said to have catalysed the industry's growth as well as its individual process.