Molecular Heterogeneous Catalysis, Wiley (2006), 352729662X

216 Chapter 5

For a polar surface, there is an overall excess charge. The energy costs required for creating polar surfaces are quite high because of the large energy penalty one must pay in separating the strong electrostatic interactions between the positively and negatively charged surfaces.

Whereas the (111) surface of MgO is polar and hence unstable, the (100) surface is neutral and thus fairly stable.

Figure 5.1. Coordination of Mg2+ on the MgO (100) surface.

Brønsted acidity or basicity develops when a molecule of water dissociates on such an oxide surface.

The question of whether H2O dissociates or not on the (100) surface of MgO was debated for some time. First-principle DFT calculations were carried out on slab models for the MgO surface and showed that H2O alone will not dissociate on the MgO (100) surface[3] . The activation of H2O on MgO leading to surface hydroxylation requires coordinatively unsaturated surface sites that can arise as surface vacancies or step edge sites. In many cases, significant reconstruction can occur, ultimately leading to the formation of a surface that is close to that of Mg(OH)2.

Here we will use the results of quantum-chemical studies of H2O dissociation on the surfaces of alumina and titania, materials that are widely used as catalytic support sur-

faces and typically show limited reconstruction. The non-hydroxylated and hydroxylated (100) surfaces of γ-Al2O3 are shown in Fig. 5.2[4].

The bulk Al2O3 contains Al octahedra that share oxygen atoms. The Pauling ionic bond strength of an Al–O bond with respect to Al is

SAl+ = +36 = 12

Each oxygen atom has four alumina cation neighbors so that SO− = SAl+ , which is consistent with a stable oxide.

The Al cation in the (100) surface (shown in Fig. 5.2) has a coordination number of five, thus

e+Al = + 12

Catalysis by Oxides and Sulfides 217

Figure 5.2. Relaxed configurations of γ-Al2O3 (100) surface for di erent hydroxyl coverages (θ in OH nm−2. The most relevant surface sites are quoted. Aln stands for aluminum atoms surrounded by n oxygen atoms, and HO–µm for OH groups linked to m aluminum atoms. Oxygen atoms are black, aluminum atoms are shown in gray whereas hydrogen atoms are white[4].

The surface oxygen atoms each have three Al neighbor atoms, one of which is below the surface, leading to an excess ion charge of:

e−O = −12

Scheme 5.1 Schematic highlighting of the products from the dissociation of water on alumina.

The charge excess here makes the surface more reactive. Water dissociation can therefore proceed heterolytically as shown in Scheme 5.1 to form the two di erent hydroxyl intermediates labeled (1)and (2).

As illustrated in this scheme, hydroxylated surfaces can contain various di erent hydroxyl groups, each of which may present di erent vibrational frequencies (see Fig. 5.3). The high-frequency bond is usually assigned to hydroxyl groups coordinated to a single aluminum ion, and labeled type 1. Their chemical properties are usually Brønsted basic. The low frequency O–H stretch is usually assigned to the hydroxyl coordinated to several cations. As we will see, its chemical reactivity is Brønsted acidic.

Figure 5.2 shows the appearance of the two di erent hydroxyl intermediates, as well as the stabilization of adsorbed H2O, which is often linked to hydroxyl (1) through the formation of a hydrogen bond.

The Brønsted acidity or basicity of the hydroxyl groups in Scheme 5.1 can be deduced from the Pauling ionic charge excesses:

−(O1) = −2 + 1 + 12 = −12−(O2) = −2 + 1 + 32 = + 12

Catalysis by Oxides and Sulfides 221

The Brønsted acidity of hydroxylated anatase is due to the additional stabilization of the weakly acidic end-on OH groups that form due to hydrogen bonding. Such e ects have been experimentally reported for analogous hydrogen-bonded silanol nests[7] .

The (110) rutile surface has two di erent types of oxygen present. The four oxygen atoms that surround the coordinatively unsaturated Ti cation (5f, Fig. 5.4b) have no charge excess. The apical bridging oxygen atoms (θO , Fig. 5.4b), however, have an ion charge excess −(O) = −23 .

Figure 5.6. Snapshots of the ionic configuration taken from a dynamic simulation of water dissociation.

(a) Initial configuration in which the water molecule lies in the (110) plane. The large gray, small white and small gray speres represent oxygen, titanium and hydrogen, respectively[6b].

In agreement with our prediction, H2O dissociates, such that the proton attaches to an apical bridging oxygen atom, θO , and OH− to the coordinatively unsaturated Ti atom. Figure 5.6 shows that this indeed happens, with little relaxation of the surface. Because of the large distances, there is no hydrogen bonding between the two hydroxyls. The larger

Brønsted acidity of rutile has to be ascribed to the charge excess of proton-coordinated θO : −(O, θO ) = + 13 .

Pacchioni[8] summarized recently the results of quantum-chemical studies and experiments aimed at characterizing the surfaces of ionic solids such as MgO and related oxides. We present a number of the salient results from this study here. A schematic representation of the MgO surface which displays a surface terrace, step and edge sites and their respective coordination numbers is shown in Fig. 5.7.

Adsorption experiments with CO have conclusively shown that CO adsorbs weakly on the MgO terrace but more strongly on edges and corners sites. Using QM cluster calculations, Petterson et al.[10] predicted CO adsorption values of 8, 18 and 48 kJ/mol for the terrace, corner and edge sites, respectively. The adsorption of CO at terrace, edge, and kink sites was found to lead to an upward shift in the CO stretching frequency by +9, +27 and +50 cm−1, respectively. These results are consistent with the generally accepted experimental data on this system of Wichtendahl et al.[11].

Interestingly, the adsorption energy of CO to Mg2+ located at an MgO corner site, attached through three neighbor oxygen atoms, is close to the adsorption energy of CO bound to an Mg2+ ion, ion-exchanged into a zeolite and charge compensated by a zeolitic four ring sof tetrahedra (see Chapter 4 pages 179, 180). The upwards shift of the CO frequency is also similar. The charges on the Mg ions located on the terrace and corner sites are very similar. The Madelung constants, however, are proportional to the electrostatic potentials and, hence, are quite di erent, as is shown in the caption to Fig. 5.7. The

222 Chapter 5

Figure 5.7. Schematic representation of an MgO surface; Mg2+ ions are represented as small spheres, the O2− anions as large spheres. The subscript indicates the coordination number of each cation site. The Madelung constants computed for these sites are: 1.681 Mg5c (terrace), 1.591 4c (edge), 1.566 Mg4c

(step), 1.344 Mg3c (corner), 0.873 Mg3c (step). The Madelung constant for a bulk Mg6c is 1.747. From Sauer et al.[9].

Madelung constant M is the proportionality constant for the ionic energy of stabilization

of cation–anion pairs in an ionic solid: Eion = −M qr2 , where q is the ion charge and r the nearest-neighbor cation–anion distance. The reduced polarization of CO at terrace sites over those at corner sites is the main contribution to the reduced Mg–CO bond energy. This reduction is due to the larger number of O2− ions coordinated to Mg2+ at the terrace site. The interaction with the O2− anions reduces the electrostatic field at CO, because their negative field counteracts the positive field from the Mg2+ cation. The decreased e ective electrostatic local fields are reflected by the decreased Pauling charge excesses that we discussed above. The low reactivity of the MgO(100) for the dissociative adsorption of H2 as well as H2O, together with the observations of reactivity at step edges for these reactions, are consistent with these conclusions. The electrostatic basis to the reactivity of the MgO surface is further elucidated by a comparison with CaO[12].

Based on theoretical studies, it is concluded that the MgO(001) surface is unreactive towards CO2 and SO2. These surfaces, however, are known to contain defect sites such as step edges and kink sites. The defect sites can be quite reactive to CO2 and SO2, thus leading to the formation of surface carbonates and sulfites, respectively. The reactivities of CaO, SrO and BaO are significantly greater than that of MgO whereby both CO2 and SO2 can react at the (100) terraces of these surfaces. This increase in activity is the result of increasing basicity as one moves from MgO to CaO, SrO or BaO. This can be explained by the di erences in electrostatic potential between MgO and CaO, SrO or BaO. Oxygen is more easily donated by oxides which have smaller Madelung potentials and hence have weaker cation–anion interactions. This reduction is due to the larger cation–anion distances for oxides with larger cation radii. For instance, the Madelung constant for the bulk oxygen anion in MgO is 23.9 eV. CaO has the same cubic structure as MgO but has a much larger lattice constant (2.399 ˚A in CaO versus 2.106 ˚A in MgO). The Madelung constant for CaO is 20.2 eV, which is 3.7 eV lower than that for MgO.

Catalysis by Oxides and Sulfides 223

5.3 The Contribution of Covalency to the Ionic Surface Chemical Bond

In this section, we extend our treatment on ionic bonding to include covalent contributions and their relevance to oxidation catalysis. We provide a more detailed molecular orbital analysis of the properties of these oxides by relating the electronic structure of cations at the surface of the oxide with that found in corresponding organometallic cluster complexes. As such, we can use more classical hybridization schemes to understand their reactivity. Accurate calculations are available for the RuO2 system that have been used as input to dynamic Monte Carlo simulations of the CO oxidation reaction to be discussed in the next subsection.

5.3.1 CO Oxidation by RuO2

In order to begin to examine the e ects of covalent bonding, we move from the more classical ionic oxides such as MgO, Al2O3 and TiO2 to RuO2, which provides more covalent chemical bonding contributions. RuO2 is a natural extension for our discussions since both the rutile and the anatase surface structures have been extensively examined in the previous section. RuO2 is known to be e cient in the oxidation of CO down to as low as room temperature[14] . CO oxidation is catalyzed by a range of other transitionmetal catalysts such as Pt (see also Chapter 8). At low temperatures on metals such as Pt, however, the rate is greatly suppressed. CO is strongly bound to Pt and results in blocking of active sites, which subsequently prevents O2 dissociation. This is not the case on Ru since Ru readily forms RuO2, which turns out to be the catalytically active phase. The dissociative adsorption of oxygen and the adsorption of CO tend to compete on the RuO2 surface.

We analyze, here the interaction of CO with the RuO2 surface in detail, since CO is often used as a probe molecule to estimate the Lewis acidity (see Section 4.2.3.1) The chemisorption of CO on the Ru4+ ions of the rutile RuO2(110) surface [which has the same structure of rutile TiO2(110) (Fig. 5.4b)] was calculated to be 120 kJ/mol [13,14]. The oxygen atoms around the Ru ions are all coordinatively saturated since they are attached to three Ru neighbors. The binding energy of CO on the more stable RuO2(001) anatase type surface (Fig. 5.4a) is only 70 kJ/mol. The decreased interaction energy is due to the fact that two of the four surface oxygen atoms around Ru are now coordinatively unsaturated, thus leading to two stronger Ru–O bonds. Therefore, as follows from the application of bond order conservation, the Ru4+–CO interaction energy on the RuO2(001)surface is decreased compared with that on the RuO2(110) surface. The adsorption energies are substantially less than the predicted DFT values on the metallic Ru surface, where the adsorption energy is 180 kJ/mol atop Ru. This reduction in the CO adsorption energy is mainly due to the substantially reduced interaction of CO with the Ru metal s and p electrons on the oxide surface. In contrast the latter contribute significantly to the metal–oxygen bond strength. The adsorption energies of CO adsorbed on Ru4+ are relatively high compared with the interaction with a hard or soft cation as found in the zeolites (Section 4.2.3.1). This is due to the relatively high charge on Ru4+ and to the additional interactions with the d-electronic orbitals.

The electronic structure of RuO2 can be deduced by inspection of Fig. 5.8a and b. Figure 5.8a compares the LDOS on oxygen in the RuO2 bulk with that of oxygen

at the surface. The main contribution to the LDOS of oxygen is from the 2p atomic orbitals. In Fig. 5.8a one notes a gap in the LDOS at –2 eV below the Fermi level. The density of states below –2 eV is assigned to electrons localized mainly in bonding orbitals with oxygen. The density of states above –2 eV is due to electron density in antibonding

224 Chapter 5

Figure 5.8a. Electronic structure of RuO2, comparison of local density of states (LDOS) of oxygen in bulk and at the (110) surface. Adapted from Y.D. Kim et al.[15].

Figure 5.8b. Pseudo valence charge density contour plots of the (a) RuO2 (110) surface in comparison with (b) the RuO2(001) surface cut through the cus-Ru atoms. These plots are defined as the di erence between the total valence electron density and a linear superposition of radially symmetric atomic charge

˚3

densities. Contours of constant charge density are separated by 0.15 eV/A . Electron depletion and accu-

mulation are marked by dashed and solid lines, respectively. In addition, regions of electron accumulation are shadowed[15].

orbitals mainly localized on Ru. One deduces from Fig. 5.8a a higher electron density on the oxygen atoms at the bridging sites (br) and a lower electron density on the oxygen on the coordinatively unsaturated (cus) sites (see Fig. 5.9a for notation). The smaller number of cation neighbors leads to less electron donation. The decreased band gap also indicates a lower electron density on surface Ru. The Ru atoms have less than half of their atomic density unoccupied, in agreement with their high positive charge. This is confirmed by the charge density di erence plots of the RuO2 (110) and (001) surfaces, respectively (Fig. 5.8b). One notes the decrease in charge densities on Ru and increased densities on oxygen.

The mechanism for CO oxidation over the RuO2 (110) surface is fairly well understood

Catalysis by Oxides and Sulfides 225

Figure 5.9. a) Perspective view of the RuO2(110) surface, illustrating the two prominent adsorption sites in the rectangular surface unit cell. These sites are labeled as the br(bridge) and the cus(coordinatively unsaturated) site, and both are occupied with oxygen atoms. b) Regions of the lowest–energy structures in (µ0 O, µCO space. The additional axes give the corresponding pressure scales at T=300 and 600 K. In the hatched region gas phase CO is transformed into graphite. Regions that are particularly strongly a ected by kinetics are marked by white hatching. [16]

[16]. The surface thermodynamics and the kinetics have been studied by first-principle DFT calculations in combination with dynamic Monte Carlo simulations. Figure 5.9 shows the reactive O atoms and vacant adsorption sites on the RuO2(110) surface. Two Ru atoms are present at the surface. The Rucus atom has five oxygen atom neighbors and the other Ru atom is coordinated to six oxygen atoms. The four surface O atoms (Obr that coordinate to Rucus are all coordinatively saturated.

Table 5.1. DFT binding energies, Eb, for CO and O [with respect to (1/2)O2] at br and cus sites (Fig. 5.9a), di usion energy barriers, ∆Edib , to neighboring br and cus sites, and reaction energy barriers, ∆Ereacb , of neighboring species at br and cus sites. All values are in eV[16b]

Ocus sites terminate the oxygen octahedron that surrounds the coordinatively saturated Ru atom. During reaction Obr may exchange position with an adsorbing CO molecule (CObr ) and, both CO and O can adsorb to the coordinatively unsaturated Rucus atom