Chemical reactions and catalysis

To my great shame, I have to admit that the intention of this brief section about chemical reactions is more to avoid a bad conscience for not mentioning them at all rather than for the reader to learn a lot! Anyway, you might get the flavour. One of the major technical motivations for doing research on surfaces has always been the understanding of heterogeneous catalysis . So let's discuss a little what this actually is. In heterogeneous catalysis the presence of a solid (the catalyst) speeds up chemical reactions which are slow or impossible in the gas phase. Why? One obvious advantage of the catalyst's surface is that adsorbed molecules are much more likely to meet each other than in the gas phase. It is the pure presence of a surface which enhances the reaction speed. Therefore an industrial catalyst is always build such as to have a high surface area. There is also another reason for this: if the really active catalyst is a precious metal then it is much wiser to distribute small metal particles on a cheap oxide or ceramic support than to have a small surface enclosing a big, expensive and useless chunk of bulk metal. A more exotic variation of this theme is the following: in the interstellar medium (the pure presence) of dust particles is needed if one wants to form hydrogen molecules. The particles are needed to get the momentum / energy balance right when the molecule is formed. Another purpose of the catalyst is to be chemically involved in the reaction. Take for example the ammonia synthesis (Haber-Bosch-process)  . N +3H =2NH The catalyst must have chemical properties such that both H and N chemisorb on the surface and get dissociated. Then they must react to form ammonia and this has to desorb again (Fig. ). This is a complicated process involving not only the actual reaction but all the chemisorption (adsorption and desorption) processes mentioned above. It is easy to see that the whole thing will not work or will be very slow if there are problems with one single step in the reaction pathway. Surface Science experiments can help to understand all steps in such a pathway and one can try to find concepts to improve the catalyst. On the other hand, the experiments are typically performed under ideal conditions, in UHV on a single-crystal surface, and one has to pay attention when applying the results to real catalysts which are often just small metal particles dispersed on some support, working under high pressure and high temperatures. In the following we briefly discuss some central concepts in heterogeneous catalysis . The first question one has to ask is how the reaction pathway actually looks like. A classical example is the oxidation of carbon monoxide. Two pathways have been suggested:

1. Langmuir-Hinshelwood:   CO CO(ads), O 2O(ads), CO(ads.)+O(ads) CO
2. Eley-Rideal:   O 2O(ads), CO+O(ads) CO

  Issues like this can be solved by molecular beam studies. One adsorbs oxygen and directs a CO beam on the surface. Then one measures the CO desorption from the surface. The time difference between the CO hitting the surface and the CO being desorbed directly points to the reaction pathway. Other central concepts in the design of a catalyst are the activity   and the selectivity  . The activity describes the degree of acceleration for the desired chemical reaction. The selectivity describes how much the converter catalyses the desired reaction as opposed to other possible reactions which are unwanted. One can change all sort of parameters in a giant parameter space to influence both reactivity and selectivity. Two other important concepts are promotion   and poisoning   of a catalyst. Let us give an example for promotion. We consider the CO dissociation. Such a reaction is often promoted by some small amount of alkali atoms into the catalyst. We can make plausible why this is so by looking at Fig. . The CO bond is already weakened upon adsorption on the surface. Co-adsorption with alkali atoms will modify the electronic structure of the substrate such that the bond will be weakened even more. Poisoning is simply the opposite effect. The catalyst ``dies'' sooner or later by adsorbing "poison'' on its surface. The poison can have different effects. The trivial one is that it simply sticks to the surface without ever being desorbed again. In this way, it will sooner or later block the surface for the real, wanted reaction. But it could also have the opposite effect as shown above for the alkali atoms: it could in some way influence the electronic structure of the substrate such that a particular reaction becomes impossible. In the latter case only a very small amount of ``poison'' could do a big damage to the catalyst. The last concept we want to mention here is that of active sites  . The idea is the following: Suppose the reaction only takes place at some sort of defects on a metal surface. The defect could be such that the atoms at the defect have a smaller co-ordination than the other surface atoms and therefore a different chemical reactivity. Such a defect could be a step on a surface. In some way, the active site idea is the kiss to death for a surface science experiment where one deals (or tries to deal with) almost perfect surfaces. On the other hand, choosing different crystal faces of a material will give surface atoms with different co-ordination numbers and one might learn something about the influence of co-ordination. There are several ways out of the problem: one can prepare surfaces with a controlled number of steps such as to increase the active sites. One can use a local technique such as Scanning Tunnelling Microscopy to look only at one active site and study it. Or one can try to ``build'' something which looks much more like an industrial catalyst but which is still a very-well characterized system. An example would be the controlled adsorption of size-sorted metal clusters on a well-prepared oxide surface and catalytic studies on such systems. This is, admittedly, very difficult but not impossible.

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