Catalysis and Catalytic Reactions
8.1 CATALYSIS AND CATALYSTS .1 Nature and Concept
8.1.3 General Aspects of Catalysis
8.1.3.1 Catalytic Sites
Central to catalysis is the notion of the catalytic “site.” It is defined as the catalytic center involved in the reaction steps, and, in Figure 8.1, is the molybdenum atom where the reactions take place. Since all catalytic centers are the same for molec- ular catalysts, the elementary steps are bimolecular or unimolecular steps with the same rate laws which characterize the homogeneous reactions in Chapter 7. How- ever, if the reaction takes place in solution, the individual rate constants may de- pend on the nonreactive ligands and the solution composition in addition to tempera- ture.
For catalytic reactions which take place on surfaces, the term “catalytic site” is used to describe a location on the surface which bonds with reaction intermediates. This involves a somewhat arbitrary division of the continuous surface into smaller ensembles of atoms. This and other points about surface catalysts can be discussed by reference to the rather complex, but typical, type of metal catalyst shown in Figure 8.2. In this example, the desired catalytic sites are on the surface of a metal. In order to have as many surface metal atoms as possible in a given volume of catalyst, the metal is in the form of small crystallites (to increase the exposed surface area of metal), which are in turn supported on an inert solid (to increase the area on which the metal crystallites reside). In the electron micrograph in Figure 8.2(a), the metal crystallites show up as the small angular dark particles, and the support shows up as the larger, lighter spheres.
Such a material would be pressed (with binders) into the form of a pellet for use in a reactor. Figure 8.2(b) is a closeup of several of the metal particles (showing rows of atoms). A schematic drawing of the atomic structure of one such particle is shown in Figure 8.2(c).
(a)
(b)
Metal atom:
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i n g ,ites
Figure 8.2 (a) Electron micrograph of a supported metal catalyst (Rh-SiO,); (b) closeup of metal particles ((a) and (b) courtesy of Professor A. Datye); (c) schematic drawing of the atomic structure of a metal crystallite
Figures 8.2(b) and (c) illustrate two important aspects of surface catalysis that distin- guish it from molecular catalysis:
(1) A distribution of “sites” exists on surfaces. By contrast with homogeneous and/or molecular catalysts in which all the sites are the same, the catalytic sites on solid surfaces can have a distribution of reactivities. The metal crystallites (which are the molecular catalytic entity) are of different sizes. They also have several dif- ferent types of surface metal atoms available for catalytic reactions. The metal atoms are in a hexagonal packing arrangement on one face, while other faces consist of the metal atoms arranged in a square pattern. The bonding of reaction intermediates to these two surfaces is different. Further variety can be found by considering the atoms at the edges between the various faces. Finally, as dis- cussed in Chapter 6, bonding of an intermediate to a site can be influenced by the bonding to nearby intermediates. Reaction mechanisms on surfaces are not usually known in sufficient detail to discriminate among these possibilities. Nev- ertheless, the simplifying assumption that a single type of site exists is often made despite the fact that the situation is more complex.
(2) Intermediates on adjacent sites can interact because of the extended nature of the surface. This option is not available to the isolated molecular catalytic entities.
This allows more possibilities for reactions between intermediates.
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Figure 8.3 Proposed reaction mechanism for methanol synthesis on Pd and comparison with gas-phase mechanism; surface inter- mediates are speculative and associated energies are estimates
8.1 Catalysis and Catalysts 181 8.1.3.2 Catalytic Effect on Reaction Rate
Catalysts increase the reaction rate by lowering the energy requirements for the re- action. This, in turn, results from the ability of the catalyst to form bonds to reaction intermediates to offset the energy required to break reactant bonds. An example of a catalyst providing energetically easier routes to products is illustrated in the multi- step reaction coordinate diagram in Figure 8.3, for the methanol-synthesis reaction, CO + 2H2 + CHsOH. The energies of the intermediate stages and the activation en- ergies for each step are indicated schematically.
For this reaction to proceed by itself in the gas-phase, a high-energy step such as the breakage of H-H bonds is required, and this has not been observed. Even with H, dis- sociation, the partially hydrogenated intermediates are not energetically favored. Also, even if an efficient radical-chain mechanism existed, the energetic cost to accomplish some of the steps make this reaction too slow to measure in the absence of a catalyst.
The catalytic palladium metal surface also breaks the H-H bonds, but since this reaction is exoergic (Pd-H bonds are formed), it occurs at room temperature. The exact details of the catalytic reaction mechanism are unclear, but a plausible sequence is indicated in Figure 8.3. The energy scale is consistent with published values of the energies, where available. Notice how the bonding to the palladium balances the bonding changes in the organic intermediates. A good catalyst must ensure that all steps along the way are energetically possible. Very strongly bonded intermediates are to be avoided. Al- though their formation would be energetically favorable, they would be too stable to react further.
In general, the reaction rate is proportional to the amount of catalyst. This is true if the catalytic sites function independently. The number of turnovers per catalytic site per unit time is called the turnover frequewy;--The reactivity of a catalyst is the product of the number of sites per unit mass or volume and the turnover fre- quency.
8.1.3.3 Catalytic Control of Reaction Selectivity
In addition to accelerating the rates of reactions, catalysts control reaction selectivity by accelerating the rate of one (desired) reaction much more than others. Figure 8.4 shows schematically how different catalysts can have markedly different selectivities.
Nickel surfaces catalyze the formation of methane from CO and HZ but methanol is the major product on palladium surfaces. The difference in selectivity occurs be- cause CO dissociation is relatively easy on nickel surfaces, and the resulting carbon and oxygen atoms are hydrogenated to form methane and water. On palladium, CO dissociation is difficult (indicated by a high activation energy and unfavorable energetics caused by weaker bonds to oxygen and carbon), and this pathway is not possible.
8.1.3.4 Catalyst Effect on Extent of Reaction
A catalyst increases only the rate of a reaction, not the thermodynamic affinity. Since the presence of the catalyst does not affect the Gibbs energy of reactants or prod- ucts, it does not therefore affect the equilibrium constant for the reaction. It follows from this that a catalyst must accelerate the rates of both the forward and reverse re- actions, since the rates of the two reactions must be equal once equilibrium is reached.
From the energy diagram in Figure 8.4, if a catalyst lowers the energy requirement for the reaction in one direction, it must lower the energy requirement for the reverse reaction.
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Figure 8.4 Hypothetical reaction coordinate diagrams for CO hydrogena- tion on Pd and Ni; the dissociation of CO is more difficult on Pd, making methanol synthesis more favorable than methane formation, which requires C-O dissociation, and is the preferred pathway on Ni