Chapter 1
Introduction to Catalysis

Ask the average person in the street what a catalyst is, and they will very likely tell you that it is part of a car. Indeed, the automotive exhaust converter represents a very successful application of catalysis; it effectively removes most of the pollutants that leave the engines of cars through exhausts. However, catalysis has a much wider scope of application than abating pollution alone [1-11].

Catalysis in Industry

Catalysts are the workhorses of chemical transformations in industry. Approximately 85-90% of the products of chemical industry are made in catalytic processes. Catalysts are indispensable in

• Production of transportation fuels in approximately 440 oil refineries all over the world.
• Production of bulk and fine chemicals in all branches of chemical industry.
• Prevention of pollution by avoiding formation of waste (unwanted byproducts).
• Abatement of pollution in end-of-pipe solutions (automotive and industrial exhaust).

A catalyst offers an alternative, energetically favorable mechanism for the noncatalytic reaction, thus, enabling processes to be carried out at industrially feasible conditions of pressure and temperature.

The economical importance is immediately evident from the estimate that industrial catalysis contributes about one quarter to the gross domestic product of developed countries. Around 2000, the catalyst world market was estimated to be about 10 billion US$, about equally distributed over refining, polymerization, chemicals, and environmental applications. The products of these processes, however, were valued at 200-300 times that of the catalyst [12].

Outside science and technology, the word catalyst is often used in the sense of 'enabler' for all kinds of processes in society.

Catalysis also plays a key role in nature [13]. Living matter relies on enzymes, which are the most abundant catalysts to be found. Photosynthesis generates sugars and oxygen from carbon dioxide and water by using chlorophyll as the catalyst and is probably the largest catalytic process in nature. Finally, the chemical industry relies heavily on catalysis, which is an indispensable tool in the production of bulk and fine chemicals, as well as fuels. In fact, energy technology also depends largely on catalysts, not only in the processing of conventional fuels, but also in emerging forms of renewable energy, for instance, where hydrogen is produced via photocatalytic or electrocatalytic routes [14, 15].

For scientists and engineers, catalysis is a tremendously challenging and highly multidisciplinary field. To better understand it, let us first see what catalysis is before we proceed to its importance and functionality for mankind.

1.1 What Is Catalysis?

In simple terms, a catalyst accelerates a chemical reaction. It does so by forming bonds with reacting molecules, allowing these to react to form a product, which then detaches from the catalyst (leaving it unaltered such that it is available for the next reaction). In fact, we can describe the catalytic reaction as a cyclic event in which the catalyst participates and is recovered in its original form at the end of the cycle.

Let us consider the catalytic reaction between two molecules A and B to a product P, see Figure 1.1. The cycle starts with the bonding of molecules A and B to the catalyst. The next step is that A and B react, by ways of the catalyst, to create product P. In the final step, product P separates from the catalyst, thus, leaving the reaction cycle in its original state.

In order to see how the catalyst accelerates the reaction, we need to look at the potential energy diagram in Figure 1.2. To fully appreciate such diagrams, it is important to realize that lower energies imply more stable situations. It compares the noncatalytic and the catalytic reaction. For the former, the figure is simply the familiar way to visualize the Arrhenius equation: the reaction proceeds when A and B collide with sufficient energy to overcome the activation barrier ?E in Figure 1.2. The change in Gibbs free energy between the reactants, A + B, and the product P is ?G, and the change in enthalpy is ?H.

Figure 1.1 Every catalytic reaction is a sequence of elementary steps, in which reactant molecules bind to the catalyst, where they react, after which the product detaches from the catalyst, liberating the latter for the next cycle.

The catalytic reaction starts by the bonding of reactants A and B to the catalyst, in a spontaneous reaction. Thus, the formation of this complex is exothermic, which in turn means that energy is lowered. Next comes the reaction between A and B while they are bound to the catalyst. This step is associated with an activation energy, ?Ecat. However, it is significantly lower than for the uncatalyzed reaction. Eventually, product P separates from the catalyst in an endothermic step.

The diagram of Figure 1.2 illustrates several important points, in terms of the enthalpy changes in the course of the reaction:

• The catalyst offers an alternative path for the reaction, which is obviously more complex, but energetically much more favorable.
• The activation energy of the catalytic reaction is significantly smaller than that of the uncatalyzed reaction; hence, the rate of the catalytic reaction is much larger (we explain this in greater detail in Chapter 2).
• The overall changes in enthalpy and also in free energy for the catalytic reaction equals that of the uncatalyzed reaction. Hence, the catalyst does not affect the equilibrium constant for the overall reaction of A + B to P. Thus, if a reaction is thermodynamically unfavorable, a catalyst cannot change this situation. A catalyst changes the kinetics but not the thermodynamics.
• The catalyst accelerates both the forward and the reverse reaction to the same extent. In other words, if a catalyst accelerates the formation of the product P from A and B, it will do the same for the decomposition of P into A and B.

In addition to the enthalpy, H, one should also consider the entropy, S. Together they determine the free energy, G = H - TS. Hence, we can also draw Figure 1.2 in terms of changes in free energy. While a reaction releasing heat is called exothermic (?H < 0), one that produces free energy is exergonic (?G < 0). Spontaneous processes are exergonic, and in catalysis, adsorption is an example. Conversely, a reaction that consumes heat is endothermic (?H > 0), while a reaction for which the free energy increases is endergonic (?G > 0). The final state of a process - for example of the reaction - is dictated by the equilibrium, and the process can be manipulated to become more exergonic by choosing optimal conditions, in case of the ammonia synthesis low temperature and high pressure. A catalyst only affects the way in which the reaction reaches its final state, and conversions can never lead to product concentrations exceeding those dictated by equilibrium.

Figure 1.2 Potential energy diagram of a heterogeneous catalytic reaction, with gaseous reactants and products and a solid catalyst. Note that the uncatalyzed reaction has to overcome a substantial energy barrier, whereas the barriers in the catalytic route are much lower.

Hence, in order to predict whether a reaction will proceed, one needs to consider the free energy, rather than the enthalpy. Doing this on the level of the elementary steps of a catalytic reaction requires that one knows both ?H and ?S, which has only recently become possible with the advent of computational chemistry. We discuss this in detail in Chapters 3 and 7.

With our current knowledge of catalytic reactions from Figure 1.2, we can already understand that there will be cases in which the combination of catalyst with reactants or products is not successful:

• If the bonding between reactants and catalyst is too weak, there will hardly be any conversion of A and B into products.
• If the bond between the catalyst and one of the reactants, say A, is too strong, the catalyst will mostly be occupied with species A, thus, not allowing B to form any product. Also, if both A and B form strong bonds with the catalyst, the intermediate situation with A or B on the catalyst may be so stable that reaction becomes unlikely. In terms of Figure 1.2, the second level lies so deep that the activation energy to form P on the catalyst becomes too high. The catalyst is said to be poisoned by (one of) the reactants.
• In the same way, the product P may be too strongly bound to the catalyst for separation to occur. In this case, the product poisons the catalyst.

Hence, we intuitively feel that the successful combination of catalyst and reaction is that in which the interaction between catalytic surface and reacting species is not too weak, but also not too strong. This is a loosely formulated version of Sabatier's Principle, which we encounter in a more precise form in Chapter 2.

The catalyst has thus far been an unspecified, abstract body, so let us now look at the different forms of catalysts that exist.

1.2 Catalysts Can Be Atoms, Molecules, Enzymes, and Solid Surfaces

Catalysts come in a multitude of...