SO2 Adsorption and Transformations on γ-Al2O3 Surfaces: A Density

May 17, 2010 - Datta , A.; Cavell , R. G.; Tower , R. W.; George , Z. M. J. Phys. Chem. 1985, 89, 443 ..... Methfessel , M.; Paxton , A. T. Phys. Rev...
2 downloads 0 Views 3MB Size
Download PDF
10444

J. Phys. Chem. C 2010, 114, 10444–10454

SO2 Adsorption and Transformations on γ-Al2O3 Surfaces: A Density Functional Theory Study John M. H. Lo,*,†,‡ Tom Ziegler,‡ and Peter D. Clark†,‡ Alberta Sulphur Research Ltd., UniVersity of Calgary, UniVersity Research Center, Unit 6-3535 Research Road N.W., Calgary, Alberta, Canada T2L 2K8, and Department of Chemistry, UniVersity of Calgary, 2500 UniVersity DriVe N.W., Calgary, Alberta, Canada T2N 1N4 ReceiVed: NoVember 16, 2009; ReVised Manuscript ReceiVed: March 21, 2010

The processes of physical and chemical adsorption of SO2 on clean (100), dehydrated (110), and hydrated (110) surfaces of γ-Al2O3 have been investigated using periodic density functional theory. In total, 18 stable forms of adsorbed SO2 have been identified on the three types of γ-Al2O3 surfaces. The computed binding energies of SO2 on these surfaces span the range of 15-70 kcal/mol, which agrees well with the experimental heat of SO2 adsorption determined using thermogravimetric methods. Among these surfaces, SO2 shows a preference to adsorb to the dehydrated surface, and the transformation into surface sulfite was observed. Theoretical vibrational frequencies of these species have been computed, and a good agreement was found with the experimental infrared spectra. It was shown that the characteristic 1060 cm-1 band on the IR spectra could be attributed, in addition to the proposed sulfate species SO4, to the HSO3 species on the hydrated (110)C surface and the SO3 species on both dehydrated and hydrated (110)C surfaces. The transformations of adsorbed SO2 to SO3/HSO3 were found to be highly exothermic with only moderate kinetic barriers on all the three surfaces. 1. Introduction SO2 is a product emitted in the combustion of coal and oils and is thought to be the major cause of acid rain. Because of its industrial, economical, and environmental importance, the process of SO2 conversion to other sulfur-containing species of lower risk is highly desirable. Among various possible routes, the Claus reaction is one of the most widely adopted industrial approaches; in this process, SO2, which is formed from the oxidation of sulfur impurities present in crude oil, is reduced to elemental sulfur by H2S over a γ-alumina catalyst.1 This method, however, suffers the drawback of catalyst deactivation by the formation of surface sulfate, which significantly impedes the reaction of SO2 with H2S and the hydrolysis of COS and CS2.2 A large number of experimental and theoretical studies have been performed to acquire insights into the adsorption and transformation mechanisms of SO2 on γ-alumina. A number of active sites on γ-alumina available for SO2 adsorption and possible adsorption modes have been proposed with the aid of IR,3-8 EPR,9,6 and XPS.10,11 Thermogravimetric analysis (TGA) using temperature-programmed desorption (TPD) and reduction (TPR) revealed that the adsorption enthalpy of SO2 varies from 13 kcal/mol,7,12 which corresponds to weakly chemisorbed SO2 molecules, to 85 kcal/mol,13 which is attributed to the strongly adsorbed surface bisulfite14 and sulfate.7,15 These sulfite-like species were proposed to be responsible for the characteristic 1060 cm-1 bands observed in TPD and FT-IR spectra;6 this proposition was supported by a recent work of Wu, Gao, and He using in situ IR and density functional theory (DFT) calculations.16 * To whom correspondence should be addressed. E-mail: [email protected]. † Alberta Sulphur Research Ltd. ‡ Department of Chemistry.

Up to this moment, the interaction mechanisms between SO2 and γ-alumina, in particular, the formation of sulfate, are still not fully understood because of the complex nature of the active Al and O sites on γ-alumina. This work, therefore, aimed at acquiring a better understanding of the geometrical and electronic factors of γ-alumina that influence the SO2 adsorption by means of DFT calculations. In particular, an extensive geometry search was performed to identify all possible stable adsorbed SO2 configurations and explore their correlation to the experimental IR spectra. Furthermore, the present work focused on the effects of different surface morphologies of γ-alumina (surface orientation and degree of hydration) on the adsorption enthalpies of SO2. The influences of the presence of OH groups on the γ-Al2O3 surfaces to the adsorption energies of SO2 and their band shifts in the IR spectra were thoroughly investigated. Finally, the energetics regarding the interconversion of adsorbed SO2 species to surface sulfite on the (100)E and dehydrated and hydrated (110)C surfaces were computed. To the authors’ knowledge, no information pertaining to the effects of OH on the chemisorption of SO2 and their surface reactions has been previously reported. 2. Computational Methods The calculations were performed with the Vienna Ab-initio Simulations Package (VASP)17-20 employing density functional theory and plane-wave basis sets. The one-electron wave function was expanded using a plane-wave basis set with an energy cutoff of 400 eV, in conjunction with the projector augmented wave (PAW) method,21 which describes the ion-electron interactions. The Perdew-Burke-Ernzerhof functional22 in a generalized gradient approximation was used to calculate the exchange and correlation energies. Spin polarization was applied in all calculations, and a Gaussian smearing of 0.1 eV was used to improve the convergence with respect to electron occupancy.23 Geometry optimization was performed

10.1021/jp910895g  2010 American Chemical Society Published on Web 05/17/2010

SO2 Adsorption and Transformations on γ-Al2O3

J. Phys. Chem. C, Vol. 114, No. 23, 2010 10445

using the conjugate-gradient algorithm24 with the force tolerance of 0.03 eV/Å. The k-point sampling was performed only at the Γ-point in all calculations. This approach may sacrifice to a small extent the accuracy of energies and structural parameters (e.g., ∼0.2 eV in total energy when k-point sampling is increased to 4 × 4 × 1) but, in exchange, is beneficial in terms of computational efficiency, especially for relatively large systems, such as those described in this work. A detailed description of the unit cell and slab model will be given in section 3. The adsorption enthalpy was computed according to the following definition

Ead ) (Esurf + ESO2) - Esurf+SO2

(1)

where Esurf and ESO2 are the energies of the surface and an isolated gas-phase SO2 molecule, respectively, whereas Esurf+SO2 is the energy of the surface with an adsorbed SO2. A positive value of Ead represents a stable adsorption of SO2. Calculations considered the adsorption of SO2 on only one side of the slab. A vacuum layer of 12 Å was inserted between slabs. To eliminate the induced dipole-dipole interaction between two unit cells, the dipole correction along the surface normal was applied. The vibrational frequencies for all optimized geometries of adsorbed SO2 were computed using the harmonic approximation and double finite difference method with a step size of 0.015 Å. The thermodynamic energy profiles and transition-state structures associated with the processes of surface migration and transformation of adsorbed SO2 were determined using the climbing image nudged elastic band (ciNEB) method25-27 and RMM-DIIS algorithm28 as implemented in VASP. 3. Results and Discussion 3.1. Cell and Surface Optimizations. It is well-known that γ-Al2O3 exists as a defective spinel derived from fcc Mg8Al16O32. To fulfill the stoichiometry, a unit cell of γ-alumina contains an average of 22/3 cation vacancies.29 However, the positions of cation vacancies have been debated for a long time because of the conflicting evidence obtained using various experimental and theoretical probes.30 Some XRD,31 NMR,32 and TEM33 measurements suggested that the vacancies reside at the octahedral cation sites, whereas some electron diffraction34 and molecular dynamics simulations35,36 found that the tetrahedral vacancies are more energetically favorable. In this work, the model structure of γ-Al2O3 where vacancies are located at octahedral sites was adopted. This choice was justified by a large number of high-level theoretical investigations in which the presence of octahedral holes is energetically more favorable than tetrahedral holes.37-46 The vacancies in γ-Al2O3 are dominated by octahedral holes at low temperature, whereas the concentration of the tetrahedral holes increases to approximately 25% at high temperatures. It is, therefore, believed that the present model should be able to represent the majority of the active γ-Al2O3 surfaces. From static calculations using an 8-layer triclinic (2 × 2) unit cell composed of 32 Al and 48 O atoms, it was found that the vacancies are distributed uniformly among the layers containing octahedral Al ions, with the largest separation of 7.54 Å; this phenomenon has also been observed previously in DFT and MD calculations.42,44 Because of the breaking of long-range symmetry due to the artificial assignment of vacancy sites, the lattice parameter of the orthogonal unit cell of γ-Al2O3 can be determined only by statistical averaging. Using the approach of Vijay et al.,42 full

TABLE 1: Calculated Surface Energies of γ-Al2O3 Surfacesa (in J/m2)

a

surfaces

present work

literature44

(100)E (100)F (110)C (110)D

0.85 1.80 1.53 1.53

1.05 1.53 1.53

The definitions of C, D, E, and F follow the work of Lippens.49

relaxation calculations of the unit cell with (4 × 4 × 4) k-point sampling yielded the lattice vector of 7.9725 Å, which agrees fairly well with the experimental value of 7.90 Å.33 One of the more popular approaches of synthesizing γ-Al2O3 is the calcination of boehmite (AlOOH) at about 700 K. Because of the topotactic nature of this transformation, the resulting γ-Al2O3 shows the same surface morphology as the parent boehmite.47 A recent DFT investigation by Digne et al.48 revealed that the most exposed surfaces of the formed γ-Al2O3 are in the (100) and (110) orientations, although the (111) surface also contributes to approximately 10% of the overall surface area. Accordingly, only the (100) and (110) surfaces of γ-Al2O3 were considered in this work. There are two possible faces for each surface, depending on the density of surface atoms. The surface energies of these surfaces have been computed, and it was found that the high-density surfaces are more stable (see Table 1). Various forms of hydroxyl groups on γ-Al2O3 surface have been identified by IR50 and assigned according to the local environments.51,52 The presence of OH groups significantly alters the acid-base properties of the γ-Al2O3; accurate studies of SO2 adsorption on γ-Al2O3 thus require a proper description of the OH coverage on these surfaces. On the basis of the DFT surface energy calculations by Digne et al.,48 the optimal concentrations of OH on the (100) and (110) γ-Al2O3 surfaces under Claus reactor conditions (∼600 K) are 0 and 8.9 OH/ nm2, respectively. Accordingly, the present work adopted the model surfaces as described by Digne and co-workers48 where the slab contains 8 layers and 16 units of Al2O3. The hydrated (110)C surface is covered with eight OH groups, whereas the (100)E surface is not hydroxylated. For comparison purposes, the dehydrated (110)C surface was also considered in this study to explore the effects of OH groups on the adsorption of SO2. 3.2. SO2 Adsorption on a γ-Al2O3(100) Surface. The optimized nonhydroxylated (100)E surface of γ-Al2O3 consists of 5-fold coordinated Al atoms and the µ3-O atoms connecting the surface Al atoms and the subsurface Al atoms alternatively at tetrahedral and octahedral sites along the [11j0] direction. Therefore, two types of adsorption sites are possible for SO2: acidic Al atoms, which are coordinatively unsaturated, and basic O atoms, which are able to provide lone pairs of electrons for dative bond formation. To verify the acidity of Al sites on the surface, the binding energies of SO2 adsorbed in parallel (CM1) and perpendicular (CM2) orientations on a Al site have been calculated; these adsorption modes are unfavorable by 1-3 kcal/ mol with respect to free SO2 (Table 2). A physisorption state (CM3) has been found when SO2 adsorbs on a µ3-O atom whose computed adsorption enthalpy is only 1.98 kcal/mol. The O · · · S distance is 2.915 Å, and the SO2 molecular plane is almost parallel to the surface. On the other hand, the interaction of SO2 via terminal O atoms to surface Al atoms results in a chemisorption state (CM4) where the binding energy is 23.88 kcal/mol and the Al-O bond distance is 2.123 Å. The SdO bond attached to the surface Al atom is slightly stretched to 1.470 Å, and the OSO angle is

10446

J. Phys. Chem. C, Vol. 114, No. 23, 2010

Lo et al.

TABLE 2: Adsorption Enthalpies of Various Adsorption Modes of SO2 on γ-Al2O3 Surfaces (in kcal/mol) dehydrated (100)E

hydrated (110)C

dehydrated (110)C

modesa

Ead

modesa

Ead

modesa

Ead

CM1 CM2 CM3 CM4 CM5 CM6 CM7 CM8

-0.98 -2.69 1.98 23.88 39.08 -21.36 15.20 45.24

HM1 HM2 HM3 HM4 HM5

20.45 25.34 31.06 17.53 34.79

DM1 DM2 DM3 DM4 DM5 DM6 DM7 DM8

26.64 24.07 36.62 48.18 56.31 34.03 70.82 57.75

a

The modes are defined in Figures 1-3.

reduced to 115°. Interestingly, the S atom orients toward a neighboring µ3-O atom; the S · · · O distance of 2.914 Å infers a possible secondary weak interaction between the adsorbed SO2 molecule and the surface (Figure 1). This work also explored the influences of subsurface vacancies on the adsorption enthalpy of SO2. A stable state (CM5) was identified when SO2 adsorbs on the oxygen atom in the vicinity of an octahedral vacancy (i.e., the O atom is only 2-fold coordinated). Owing to a smaller geometric constraint of the lone pairs of electrons, the O atom is more basic, thus resulting in a stronger interaction with the adsorbed SO2 molecule. The calculated bond energy is 39.08 kcal/mol, and the O-S distance is only 1.645 Å, which is 44% shorter and almost 20-fold stronger than the corresponding O-S bond of CM3. Starting from CM3 and CM5, two different types of surface sulfite with a Al-O-SO2 configuration could be obtained. Breaking two of the µ3-O-Al bonds of CM3, followed by the migration of the thus formed SO3 to a neighboring Al site, leads to CM6. This structure has a rather short Al-O bond (1.780 Å) and a long S-O bond (1.571 Å); however, geometry optimization revealed that this structure is energetically unfavorable with the binding energy of -21.36 kcal/mol. A similar transformation converts CM5 to CM7, whose adsorption energy is, however, 15.20 kcal/mol. Structurally, CM7 is close to CM6 except for a more stretched Al-O bond (1.910 Å) and shorter S-O bond (1.526 Å). The noticeable difference in relative stability of CM6 and CM7 is possibly attributed to the creation and exposure of a subsurface, 5-fold Al atom, which is highly unstable, when CM3 is transformed into CM6. This problem is absent in the CM5 f CM7 process. One of the SdO bonds of CM7 can attack the neighboring 4-fold coordinated Al site to form a cyclic sulfite (CM8). This configuration has the highest adsorption energy (45.24 kcal/ mol) among all adsorption modes of SO2 on the (100)E surface and possesses a plane of symmetry that contains the SdO bond and bisects the molecule (Al-O bonds, 1.913, 1.919 Å; S-O bonds, 1.538, 1.541 Å; SdO bond, 1.456 Å). 3.3. SO2 Adsorption on a Hydrated γ-Al2O3(110) Surface. On the basis of the results obtained in FTIR and EPR studies, Datta et al. proposed five possible adsorption modes of SO2 on a hydroxylated γ-Al2O3 surface.6 Meanwhile, Babaeva et al.53 and Saur et al.15 suggested a surface bisulfite species on γ-Al2O3 in the presence of water. Therefore, these six models were taken as the initial configurations of adsorbed SO2 on the hydrated (110)C surface in the subsequent calculations. The calculated adsorption enthalpies for these species are shown in Table 2. Full geometry optimizations, including all atoms of the slab model, identified five stable configurations (Figure 2). Two physisorption modes were found that have the SO2 molecule coordinated via its S atom to a surface hydroxyl group,

corresponding to the model I proposed by Datta et al.6 The first mode (HM1), which is coordinated to the OH group on a 5-fold coordinated Al atom, has a S · · · O distance of 2.277 Å and forms two hydrogen bonds (1.769 and 1.980 Å) with two surface hydroxyl groups. The second mode (HM2) results from the interaction of SO2 to the OH group on a 4-fold coordinated Al atom, possessing a slightly longer S · · · O bond (2.520 Å) than HM1. Nevertheless, it forms three hydrogen bonds (1.761, 2.045, and 2.204 Å) with three other surface OH groups. Consequently, the computed adsorption energy of HM2 is about 5 kcal/mol larger than that of HM1. Contrary to the dehydrated (100)E surface, the direct chemisorption of SO2 on a bridging O (either µ2- or µ3-type) is not favorable, possibly because of the steric interaction with neighboring OH groups and the unavailability of free bridging O atoms. Similarly, no stable configurations of SO2 adsorbed via S on Al sites, which correspond to the models III-V of Datta et al., could be found in the present study. Instead, a chemisorption mode was noticed in which SO2 adsorbs on a 5-fold coordinated Al atom using its terminal oxygen. This mode (HM3) is energetically favorable with the binding energy of 31.06 kcal/mol and a short Al-O bond (1.997 Å), but the adsorbed SO2 does not form any hydrogen bonds with the surrounding OH groups (the shortest O · · · H distance is about 2.8 Å). A hydrogen transfer from the OH group to a SdO group of HM1 leads to the formation of a surface bisulfite (HM4). The structure of this species resembles HM1 closely; the two hydrogen bonds remain between SdO of HM4 and nearby OH groups (1.875 and 1.881 Å, respectively), while an additional hydrogen bond is formed between the O atom of HM4 bonded to Al and a neighboring µ3-OH whose H atom points toward HM4. Despite a more extensive H-bond network, HM4 was found to be less stable than HM1 by about 3 kcal/mol. In the presence of an unprotonated µ2-O bridge and a coordinatively unsaturated Al site in close proximity, SO2 can adsorb simultaneously on the O and Al atoms, forming a very stable surface sulfite (HM5). This species is characterized by its SdO bond; the bond is slightly stretched (1.479 Å) and is weakly hydrogen-bonded to a µ2-OH group (1.917 Å). The formation of HM5 is thermodynamically very favorable with the adsorption enthalpy of 34.79 kcal/mol; the preference of SO3 formation has also been observed in previous IR and thermogravimetric studies of SO2 adsorption on γ-Al2O3.4,6-8 3.4. SO2 Adsorption on a Dehydrated γ-Al2O3(110) Surface. The adsorption of SO2 is remarkably different on a dehydrated (110)C surface compared with the surface loaded with water. Because of the availability of a large number of highly unsaturated 3-fold and 4-fold Al sites, more adsorption modes and a stronger interaction of SO2 on the dehydrated (100)C surface are expected. As shown in Table 2, the calculated SO2 binding energies for this surface are generally higher than those for the hydrated counterpart. Eight stable configurations have been found for adsorbed SO2 on this surface; their structures are depicted in Figure 3. Two physisorption states were identified for SO2; the former one (DM1) has the S atom coordinated simultaneously to two 4-fold coordinated Al atoms, whereas the latter one (DM2) is adsorbed on a 3-fold coordinated Al atom. Both species are less stable compared with other possible modes, but their adsorption enthalpies (26.64 and 24.07 kcal/mol, respectively) are comparable to the values for HM1 and HM2 yet much higher than those for CM1 and CM2. Explanations for the exceptionally high adsorption energies of DM1 and DM2 are not clear, but

SO2 Adsorption and Transformations on γ-Al2O3

J. Phys. Chem. C, Vol. 114, No. 23, 2010 10447

Figure 1. Optimized geometries of CM3-CM5, CM7, and CM8 on the dehydrated (100)E γ-Al2O3 surface (green ) Al, white ) O, blue ) S, red ) O attached to S).

the O atoms bonded to the Al atom where SO2 resides might play a role, as inferred by their noticeable outward relaxation. The monodentate coordination of SO2 to 3-fold Al sites, such as HM3, was not stable on the dehydrated (110)C surface, likely due to the lack of hydrogen-bond stabilization. Instead, SO2 prefers to coordinate to two Al sites simultaneously to form a ring configuration. This results in the formation of DM3, where SO2 is bonded to two 4-fold Al atoms, and DM4, where SO2 is bonded to a 3-fold Al atom and a 4-fold Al atom. The extra stability of DM4 (∼12 kcal/mol) can be attributed to the completion of an octet configuration of the more acidic 3-fold Al atom. No direct coordination of SO2 on the O bridge through the S atom was found; instead, geometry optimizations showed that the three atoms of SO2 are simultaneously involved in the surface-SO2 interactions, forming one S-O bond (1.690 Å) and two Al-O bonds (1.859 and 1.949 Å). Consequently, this state (DM5) possesses an exceptionally high adsorption enthalpy relative to other adsorption modes. A closely related adsorption mode (DM6) has also been found where the Al-O bond between the 3-fold coordinated Al atom and SO2 is broken. Meanwhile, the surface Al-O-Al linkage is broken, and the 3-fold coordinated Al atom relaxes laterally toward a neighboring subsurface oxygen, forming an additional Al-O bond to acquire 4-fold coordination. Because of the substantial surface

reconstruction and the scission of a surface Al-O bond, this sulfite species DM6 is much less stable than DM5, as shown in Table 2. Two sulfite species possessing a tripod configuration could be derived from DM6 by coordinating its SdO group to nearby Al sites. The first one (DM7) involves the formation of the third Al-O bond with a 4-fold Al atom. In this configuration, the three S-O bonds have almost the same length (1.57-1.58 Å). Interestingly, it is seen from Figure 3 that DM5 can transform directly into DM7 by splitting the surface 3-fold Al-O bond and undergoing structural relaxation. The resulting structure is energetically more favorable than DM5; the enhanced stability arises possibly from the formation of bonds between the laterally relaxed Al atom and subsurface oxygen atoms. The second one (DM8) results when the SdO bond of DM6 flips and coordinates to a subsurface O atom. This process, however, leads to a significant reconstruction of the surface; the subsurface O atom bonded to SO2 becomes 4-fold coordinated and moves upward to the surface. Consequently, the binding energy associated with this adsorption mode is much lower compared with that of DM7 (∆Ead ∼ 13 kcal/mol) (Figure 3 and Table 2). 3.5. Theoretical Vibrational Spectra: Comparison with Experiments. In addition to the thermochemical methods, such as TPD and TGA, IR spectroscopy serves as a direct experi-

10448

J. Phys. Chem. C, Vol. 114, No. 23, 2010

Lo et al.

Figure 2. Optimized geometries of HM1-HM5 on the hydrated (110)C γ-Al2O3 surface (green ) Al, white ) O, blue ) S, pink ) H, red ) O attached to S).

mental probe that provides useful information regarding the geometries of adsorbed SO2 on the γ-Al2O3 surfaces. Various adsorption modes can be distinguished by the appearance of distinct absorption bands within the 1400-1000 cm-1 region in the IR spectrum. An early work of Chang disclosed that two intense bands at 1100 and 1400 cm-1 are observed during SO2/O2 adsorption on γ-Al2O3, which only disappear when the desorption process is performed above 800 °C. In the absence of O2, two bands at 1060 and 1326 cm-1, arising from a sulfite species, appear instead, which can be removed at 600 and 100 °C, respectively.4 He ascribed the former two bands to the surface sulfate formed from oxidation while the latter two by the presence of sulfite. The 1060 cm-1 band has also been detected in the subsequent studies by Datta et al.,6 Nam et al.,7 and Mitchell et al.;8 it was concluded in these studies that Al-SO3 is responsible for this intense absorption band. Meanwhile, the band at 1326 cm-1 comes possibly from a physisorbed SO2 on a hydroxyl group.6,8 On the other hand, the pair of bands at 1400 and 1100 cm-1 are constantly observed in the FTIR and TGA experiments concerning the oxidative adsorption of SO2 on alumina,15,7,8,16 although the absolute positions of these bands are dependent upon the experimental conditions. The DFT calculations by Wu

et al. employing a cluster model of alumina confirmed the assignment of the 1400 cm-1 band to the surface sulfates in bidentate and tridentate configurations.16 The computed vibrational frequencies for various adsorption species (CM3-CM8, HM1-HM5, DM1-DM8) are summarized in Table 3. As can be seen, the experimentally observed IR absorption bands associated with adsorbed SO2 can be well accounted for by the models considered in the present work. In the comprehensive work of Datta et al., five groups of bands were found and were assigned to five different types of adsorbed SO2.6 The characteristic 1135 and 1065 cm-1 bands were designated to surface sulfite bonded through S to an Al site. The present work shows, however, that both bisulfite (HM4) and sulfite (HM5), which possess a SdO group, give rise to the 1060 cm-1 band. Sulfite CM7 on the dehydroxylated (100)E also shows a band at about 1060 cm-1 due to the SdO group, but its contribution is small relative to HM4 and HM5. On the other hand, the shoulder band at 1135 cm-1 can be traced to the S-O-H bending of HM4 (1119 cm-1) and the SdO stretching of CM8 (1139 cm-1). The low intensity of this shoulder band can be ascribed to the relatively low thermodynamic stability of HM4, as well as the less exposure of the (100)E surface compared with the (110)C surface.

SO2 Adsorption and Transformations on γ-Al2O3

J. Phys. Chem. C, Vol. 114, No. 23, 2010 10449

Figure 3. Optimized geometries of DM1-DM8 on the dehydrated (110)C γ-Al2O3 surface (green ) Al, white ) O, blue ) S, red ) O attached to S).

The band at 1189 cm-1, according to Datta et al., originates from the SO2 molecule adsorbed on an acidic Al site. As discussed in previous sections, no adsorption modes involving direct interaction between Al and S could be found except for the dehydrated (110)C surface (i.e., DM2); this is not unexpected as S, which is bonded to electronegative O in SO2, is not strongly basic to favorably form a dative bond with an acidic Al atom on γ-Al2O3. Notice also that, for the hydrated (110)C surface, OH groups block all the reactive 3-fold Al sites and part of the 4-fold Al sites, therefore making the Al-SO2 adsorption less feasible. The asymmetric SdO stretching mode of DM2 was found to be 1304 cm-1; this value is close to that of a gaseous SO2 molecule (1361 cm-1)54 but substantially different from the 1189 cm-1 band. According to the results obtained in this work, this band could be assigned to HM2 or DM6. The former configuration corresponds to a physisorbed SO2 that is simultaneously hydrogen-bonded by neighboring hydroxyl groups, thereby

reducing its asymmetric stretching mode to 1186 cm-1. The latter structure can be considered as a surface sulfite where two of the three S-O groups are bonded to Al atoms. Because of the induced steric strain, the SdO stretching mode is blue shifted to 1192 cm-1. The 1189 cm-1 band could also come from CM5 where SO2 is adsorbed on an O bridge above a vacancy. This structure shows an asymmetric SdO mode at 1168 cm-1. Nevertheless, its contribution to the IR spectrum should be insignificant owing to the small percentage of the dehydrated (100)E plane to the overall γ-Al2O3 surface and the low concentration of subsurface octahedral vacancies. Datta et al. assigned the 1255 cm-1 band to a SO2 molecule adsorbed simultaneously to several Al sites. This structure resembles the DM1 model on the dehydrated (110)C surface investigated in the present study; the calculated asymmetric SdO stretching mode is 1262 cm-1. No similar modes were found for the corresponding hydrated (110)C surface. This surface preference is in agreement with the experimental

10450

J. Phys. Chem. C, Vol. 114, No. 23, 2010

TABLE 3: Calculated Vibrational Frequencies of Adsorbed SO2 on γ-Al2O3 Surfaces (in cm-1) surfaces

adsorption modes

frequencies

dehydrated (100)E

CM3 CM4 CM5 CM7 CM8 HM1 HM2 HM3 HM4 HM5 DM1 DM2 DM3 DM4 DM5 DM6 DM7 DM8

1238, 1071, 494 1235, 1035, 538 1168, 1013, 692, 609, 543 1200, 1062, 881, 848, 573 1139, 898, 851, 589, 551 1221, 1050, 537 1186, 1041, 532 1113, 944, 499 1119, 1057, 665, 599, 519 1073, 970, 637, 517, 505 1262, 1080, 493 1304, 1082, 503 1032, 986, 490 998, 923, 489 982, 891, 727, 674, 528 1192, 891, 756, 557, 523 915, 829, 788, 583, 555 988, 895, 811, 574, 511

hydrated (110)C

dehydrated (110)C

observation that the 1255 cm-1 band was detected only for the alumina sample activated at 700 °C.6 There are two other structures that give rise to an IR absorption band at about 1240 cm-1; they correspond to the asymmetric SdO stretching mode of a physisorbed SO2 (CM3) and a SO2 adsorbed on an Al site via its terminal O (CM4). The calculated values (1238 cm-1 for CM3 and 1235 cm-1 for CM4) differ by about 20 cm-1 from the experimental band; however, they should have a slightly higher contribution relative to DM1 due to the fact that the dehydrated (100)E plane is more exposed than the dehydrated (110)C plane on γ-Al2O3. In Datta’s model, the chemisorbed SO2 on an O bridge are responsible for the IR bands at 1322 and 1140 cm-1 that disappear after evacuation at 100 °C.6 This adsorption mode has also been found, in this work, on the dehydrated (110)C surface (DM1). Nevertheless, the vibrational frequency calculations demonstrated that the SdO stretching frequencies are significantly red shifted to about 1260 cm-1. On the other hand, the other physisorption mode of SO2 (DM2) shows, in normalmode analysis, the asymmetric SdO stretching mode at 1304 cm-1 that is only slightly lower than the experimental value (1322 cm-1). The reasons explaining the red shift of these bands are not certain, but its low relative intensity is well correlated with the small exposure of the dehydrated (110)C surface and the occupation of 3-fold Al sites by surface hydroxyl groups. Being the first that disappear in the evacuation experiments, the set of bands at 1334 and 1148 cm-1 were assigned by Datta et al.6 and Mitchell et al.8 to the SdO vibrations of SO2 physically adsorbed on hydroxyl groups. The corresponding structure (HM1) found in the present study does not yield consistent SdO stretching frequencies; the normal mode calculations predicted the values of 1221 and 1050 cm-1, both of which are red shifted by approximately 100 cm-1. The large discrepancy between the experimental and theoretical frequencies can be rationalized by considering the fact that HM1 is extensively hydrogen bonded by surrounding hydroxyl groups, which significantly weakens the SdO bonds, thus lowering the SdO stretching frequencies. The same scenario, however, does not likely occur in practice because this physisorption mode appears at the end of the adsorption cycle where most of the surface hydroxyl groups have already been occupied; no free hydroxyl groups in close proximity would be available for forming hydrogen bonds with the SO2 molecule physisorbed on a OH group. Moreover, the experimentally observed bands

Lo et al. TABLE 4: Assignments of IR Bands for SO2 on γ-Al2O3 Surfacesa observed bands

Datta’s model6

1060, 1135

Al-SO3

1189

Al-SO2

1255

Aln-SO2

1140, 1322 1148, 1334 HM3 > HM2 > HM4 > HM1. This order is perfectly in line with the desorption temperatures associated with the experimental IR bands and their assignments based on the present study except for HM4, which is the most weakly bound species on the hydroxylated (110)C surface. However, this deviation is not influential as the surface concentration of HM4, at equilibrium, is several orders of magnitude lower than those of HM3 and HM5. The corresponding trends for the less abundant CM3-CM8 and DM1-DM8 are also consistent with the IR band assignments, as shown in Table 4. The more persistent band (i.e., (1322, 1140) < (1255) < (1189)) is given an assignment to a more thermally stable species (i.e., DM2 < DM1 < DM6 and CM3 < CM4 < CM5, respectively).

SO2 Adsorption and Transformations on γ-Al2O3

J. Phys. Chem. C, Vol. 114, No. 23, 2010 10451

Figure 4. Energy profile associated with the transformation of adsorbed SO2 on the dehydrated (100)E γ-Al2O3 surface. Energy values (in kcal/ mol) are reported with respect to a clean alumina surface and free SO2.

The assignment of the 1334 cm-1 band to HM1 is apparently inconsistent with the observed high volatility in the TPD experiments. The computed binding energy of HM1 is about 20 kcal/mol, which is 30% larger than the lower limit of the heat of adsorption derived from the SO2 adsorption isotherm.7 Note that this discrepancy can be attributed to the over stabilization of HM1 at low surface coverage by forming hydrogen bonds with surrounding OH groups. In FTIR and TPD experiments, the 1334 cm-1 band is the last that appears during an adsorption cycle; at this stage, with a high surface concentration of SO2, most of the surface OH groups have already been involved in SO2 chemisorption and are unavailable for forming additional hydrogen bonds with physisorbed SO2. In addition, lateral interaction between adsorbed species also weakens the strength of physisorption of SO2 on OH groups. Consequently, the resulting experimental adsorption enthalpy of physisorbed SO2 is much lower than the theoretical value. 3.6. EPR Measurements. The electron paramagnetic resonance spectroscopy (EPR) has been used to probe the possible adsorption sites on γ-Al2O3 for SO2;9,55-57 their measurements have shown the possible appearance of the SO2- radical on γ-Al2O3 surfaces, although no direct correlation between the EPR signal and major IR band intensities could be verified. The later work by Datta et al. also confirmed these observations; they reported that noticeable signals were detected only when the system was heated above 450 °C and the EPR signals should arise from only a small portion of SO2 adsorbed at some particular adsorption sites of γ-Al2O3.6 The present work investigated the spin densities of all the proposed adsorption modes of SO2 on the γ-Al2O3 surfaces. Among the three surfaces, the five stable adsorption modes on the dehydrated (100)E surface (i.e., CM3-CM5, CM7, CM8) are EPR-active, whereas the adsorption modes on the (110)C surfaces are EPR-inactive, with the only exceptions of DM3 and DM4, which show weak EPR activity. These results demonstrate a good agreement with the experimental observations. Note that the surface coverage of the (100)E plane on γ-Al2O3 is relatively small compared with the (110)C plane; therefore weak EPR signal intensities are expected. In addition, the low abundance of these adsorption sites partly accounts for the apparently uncorrelated EPR and IR band intensities. On the basis of Table 3, it can be seen that CM4, CM5, and CM8 show IR bands in the 1000-1200 cm-1 region. Because of their low concentrations on the γ-Al2O3 surfaces, these bands

are interfered and masked by the IR bands arising from more abundant HM2 and HM3. Consequently, the variations of these IR bands with respect to the changes in the EPR signals intensities could not be easily seen. On the other hand, the observation that a high activation temperature (∼700 °C) is required to generate the particular alumina sites responsible for the formation of adsorbed SO2radicals is consistent with the results obtained in the present study. The theoretical studies by Digne et al.48 have revealed that the OH concentration on the (110)C surface is reduced from 8.9 to 3.0 OH/nm2 when the system is heated at about 700 °C. It is inferred from these observations that the presence of OH groups likely quenches the EPR signals or suppresses the formation of SO2- radicals. In the present calculations, no EPRactive adsorption sites could be found for the hydrated (110)C surface, whereas a number of EPR-active adsorption sites (DM3 and DM4) were identified when the surface is dehydrated. At higher temperatures, the surface coverage of OH is reduced, and more DM3 and DM4 species would be present on the surface, thus leading to an enhanced EPR signal. 3.7. Transformations of SO2 on γ-Al2O3 Surfaces. Figure 4 illustrates the kinetic energy profile describing the adsorption and transformation of adsorbed SO2 on the dehydrated (100)E surface of γ-Al2O3. The initial adsorption of SO2 takes place readily because the reaction barrier is only 2 kcal/mol leading to CM4; the formation of physisorbed CM3 is barrierless. Both CM3 and CM4 transform into CM5, via TS2 and TS3 respectively, but the former process proceeds more rapidly because the associated activation enthalpy is only 0.2 kcal/mol. The corresponding activation energy for CM4 f TS3 is relatively large (∼14 kcal/mol), but the energy gain due to SO2 adsorption (-24 kcal/mol) is sufficient to compensate for this cost. The subsequent transformations from CM5 are energetically demanding. Because of a transition state (TS4) possessing an oxygen vacancy at a surface octahedral site, which is highly unstable, a remarkable reaction barrier (∼27 kcal/mol) has to be overcome in the process CM5 f TS4 f CM7. The barrier for the reverse process CM5 f TS3 f CM4 is even higher (∼30 kcal/mol). Therefore, these processes are mainly kinetically controlled, and heat has to be provided to trigger the transformations. These results are consistent with the experimental observations that the 1189 cm-1 band, which is assigned to CM5 in the present study, disappears only if the evacuation

10452

J. Phys. Chem. C, Vol. 114, No. 23, 2010

Lo et al.

Figure 5. Energy profile associated with the transformation of adsorbed SO2 on the hydrated (110)C γ-Al2O3 surface. Energy values (in kcal/ mol) are reported with respect to a clean alumina surface and free SO2.

is performed at above 200 °C, accompanied by the appearance of the 1130 cm-1 band, which is assigned to CM8.6 Depending on the locations where SO2 adsorbs, three different processes of SO2 adsorption can occur on the hydroxylated (110)C surface concurrently, resulting in HM1, HM2, and HM5. As shown in Figure 5, these three processes, based on the ciNEB calculations, contain no transition state; therefore, HM1, HM2, and HM5 are formed very rapidly when the surface is first exposed to SO2. This is in good agreement with the observation by Chang4 and Datta et al.6 that the 1322, 1189, and 1050 bands appeared when γ-Al2O3 was treated with a small dose of SO2 (about 0.03 mmol) at room temperature. Upon adsorption on a surface hydroxyl group (HM1), SO2 may undergo either hydrogen abstraction from the OH group it coordinates to become bisulfite (HM4) or surface hopping to a coordinatively unsaturated Al site in near proximity to form a chemisorbed SO2 (HM3). Both steps are kinetically feasible, with an activation energy below 10 kcal/mol, but the first one is marginally thermodynamically unfavorable. As explained in earlier sections, this may not represent the true kinetic picture of this transformation because the current calculations overestimated the hydrogen-bond stabilization on HM1. It is, therefore, very likely that the process HM1 f TS7 f HM4 is favorable from both thermodynamic and kinetic perspectives. On the other hand, the transformation of HM1 to HM3 is highly favorable with an energy gain of 11 kcal/mol. Hence, it is believed that HM1, once it is formed, quickly turns to HM3, leading to the accumulation of HM2, HM3, and HM5 on the hydrated (110)C surface, and this accounts for the extremely low intensity of the 1334 cm-1 (HM1) band relative to the 1189 (HM2), 1135 (HM3), and 1065 cm-1 (HM5) bands observed in Datta’s work.6 The enthalpy diagram regarding the interconversions of DM1-DM8 is depicted in Figure 6. It is expected from the structures of DM1-DM8 that many of these species could be obtained by direct adsorption of SO2. Interestingly, only two pathways of direct adsorption of SO2 have been located, and both channels possess no activation barrier. The first path is the adsorption of SO2 on an Al-O-Al bridge, resulting in the formation of DM5, with the associated enthalpy change of 56 kcal/mol. The second path involves the coordination of SO2 asymmetrically on a 3-fold Al and a 4-fold Al site simultaneously (DM4); this path is less favorable than the first one by about 8 kcal/mol.

Figure 6. Energy profile associated with the transformation of adsorbed SO2 on the dehydrated (110)C γ-Al2O3 surface. Energy values (in kcal/ mol) are reported with respect to a clean alumina surface and free SO2.

There are three possible transformation routes for DM5. It could be converted to DM3 via either a single-step (DM5 f TS8 f DM3) or a multiple-step (DM5 f TS9 f DM4 f TS11 f DM3) reaction. The barriers of the rate-determining step of these two channels are similar (∼24 kcal/mol), but the effective activation energy for the latter channel is larger; consequently, the single-step transformation of DM5 to DM3 should be more probable. However, it is noted that DM3 and DM4 should not constitute a significant portion of the surface species because of the endothermic nature of their transformations from DM5. DM5 can also undergo a conversion with essentially a zero barrier (∼0.5 kcal/mol) to the extremely stable DM7. The transition state TS10 resembles very much DM5; the surface 3-fold Al-O bond is stretched from 2.031 to 2.333 Å, while the 4-fold Al-O bond is compressed by 0.02 Å. The subsequent transformation of DM7 to DM6 is kinetically hindered because of the exceptional reaction barrier (∼37 kcal/mol), although DM6 is easily converted to DM8 and has a noticeable energy relief (∼24 kcal/mol). Accordingly, it is anticipated that DM3-DM8 would not show up with significant intensity in the 1500-1000 cm-1 region of the IR spectrum for adsorbed SO2 on γ-Al2O3. This proposition agrees with the IR band assignments in Table 4 where none of these species, except DM6, is associated with the five major bands above 1000 cm-1. It is worth mentioning that the initial adsorption of SO2 on γ-Al2O3 practically possesses no activation barrier. The only exception is the formation of CM4 where the activation energy is 2 kcal/mol. According to the transition-state theory, it is expected that the reaction rates associated with these steps should drop when the reaction temperature increases; in other words, the SO2 capacity of γ-Al2O3 should demonstrate a negative dependence upon temperature. Consistent experimental observations have been obtained by Nam and Gavalas in their TGA studies on the temperature effects on the SO2 adsorption on γ-Al2O3.7 In their work, it was observed that, in a 0.22% N2 stream at 1 atm pressure, the amount of SO2 adsorbed by γ-Al2O3 after 1 h of exposure decreases from 5.0 mol % of SO2 at 500 °C to 2.0 mol % of SO2 at 700 °C.

SO2 Adsorption and Transformations on γ-Al2O3 4. Conclusions The process of SO2 adsorption on γ-Al2O3, which is an essential part of the Claus process, has been investigated using density functional theory in conjunction with the PAW method. In this work, three different surfaces of γ-Al2O3 have been employed, namely, dehydrated (100)E, dehydrated (110)C, and hydrated (110)C surfaces. For each surface, the structures and energetics of various adsorbed SO2 species were optimized. It was found that the calculated adsorption enthalpies of adsorbed SO2 on these surfaces fall within the experimental adsorption enthalpy (13-85 kcal/mol). Among these surfaces, SO2 preferentially adsorbs to the (110) surface if it is water-free. This observation agrees with the fact that the SO2 capacity increases when γ-Al2O3 is activated at higher temperatures. When the surface is hydroxylated, the adsorption energies of SO2 drop dramatically from 24-71 to 17-34 kcal/mol; the SO2 adsorption then shows no selectivity between the dehydrated (100)E and hydrated (110)C surfaces. Interestingly, the presence of surface OH groups assists the physisorption of SO2 by means of forming extensive hydrogen bonds, resulting in an overestimation of the computed adsorption energies of HM1 and HM2. To verify the structures of adsorbed SO2, normal-mode analysis was performed and the results were compared with the experimental IR spectra of SO2 adsorbed on γ-Al2O3. A good agreement between the computed and experimental SdO stretching modes was observed; the present calculations reproduced the five major groups of bands in the 1400-1000 cm-1 region, as described by Datta and co-workers. Nevertheless, it is worth noting that the present assignments of bands are partly different from Datta’s propositions. On one hand, as suggested by Datta et al., the (1334 and 1148 cm-1) and (1322 and 1140 cm-1) bands are assigned to a physisorbed SO2 on hydroxyl (HM1) and a chemisorbed SO2 on an O bridge (DM2), respectively, although a red shift was observed for both sets in the present calculations. On the other hand, instead of assigning the 1255 and 1189 cm-1 bands to SO2 bound to surface Al via S, the present study suggested that they likely arise from the species where S forms strong bonds with surface oxygen atoms. It was also found that surface sulfite and bisulfite in monodentate (CM7 and HM4) and bidentate (CM8 and HM5) coordination modes are responsible for the 1135 and 1065 cm-1 bands that survive in evacuation experiments up to 600 °C. It has to be emphasized that the sulfate species (SO4) was not considered in this work, and thus, its contributions to the experimental IR spectra cannot be excluded. It has been demonstrated that bidentate sulfate species formed on γ-Al2O3 give rise to the 1214 and 1349 cm-1 bands. Finally, the present work has explored the transformations of various SO2 species on the three different surfaces of γ-Al2O3. It was noticed that the nondissociative adsorption of SO2 on γ-Al2O3 is effectively barrierless, regardless of the surface orientations; this accounts for the negative temperature dependence of the SO2 adsorption observed in the TGA measurements of Nam and Gavalas. In addition, the computed kinetic energy profile for CM3-CM8 provides a rationale behind the shift of the band at 1189 cm-1 (CM5) to 1130 cm-1 (CM8) at 200 °C; the process is mainly governed by the exceptionally high reaction barrier that can be overcome only when sufficient energy is provided. Acknowledgment. J.M.H.L. would like to acknowledge the Industrial Research and Development Fellowship by the Natural Sciences and Engineering Research Council of Canada (NSERC) and Alberta Sulphur Research Ltd. All calculations were

J. Phys. Chem. C, Vol. 114, No. 23, 2010 10453 performed on the computing facilities provided by the Western Canada Research Grid. Note Added after ASAP Publication. This article was published ASAP on May 17, 2010. On page B, the force tolerance of 0.03 eV was changed to 0.03 eV/Å. The correct version was published on May 19, 2010. Supporting Information Available: Fractional coordinates of the (100)E, dehydrated (110)C, and hydrated (110)C surfaces, as well as the adsorbed SO2 species (CM3-CM5, CM7-CM8, HM1-HM5, DM1-DM8). This material is available free of charge via the Internet at http://pubs.acs.org. References and Notes (1) Pie´plu, A.; Saur, O.; Lavalley, J.-C.; Legendre, O.; Ne´dez, C. Catal. ReV.-Sci. Eng. 1998, 40, 409. (2) Graulier, M.; Papee, D. Energy Process./Can. 1974, 1. (3) Deo, A. V.; Dalla Lana, I. G.; Habgood, H. W. J. Catal. 1971, 21, 270. (4) Chang, C. C. J. Catal. 1978, 53, 374. (5) Karge, H. G.; Dalla Lana, I. G. J. Phys. Chem. 1984, 88, 1538. (6) Datta, A.; Cavell, R. G.; Tower, R. W.; George, Z. M. J. Phys. Chem. 1985, 89, 443. (7) Nam, S. W.; Gavalas, G. R. Appl. Catal. 1989, 55, 193. (8) Mitchell, M. B.; Sheinker, V. N.; White, M. G. J. Phys. Chem. 1996, 100, 7550. (9) Karge, H. G.; Trevizan de Suarez, S.; Dalla Lana, I. G. J. Phys. Chem. 1984, 88, 1782. (10) Smirnov, M. Y.; Kalinkin, A. V.; Pashis, A. V.; Sorokin, A. M.; Noskov, A. S.; Bukhtiyarov, V. I.; Kharas, K. C.; Rodkin, M. A. Kinet. Catal. 2003, 44, 575. (11) Smirnov, M. Y.; Kalinkin, A. V.; Pashis, A. V.; Sorokin, A. M.; Noskov, A. S.; Kharas, K. C.; Bukhtiyarov, V. I. J. Phys. Chem. B 2005, 109, 11712. (12) Fellner, P.; Jurisova, J.; Khandl, V.; Sykorova, A.; Thonstad, J. Chem. ZVesti 2006, 60, 311. (13) Grass, R. W.; Ross, R. A. Can. J. Chem. 1972, 50, 2537. (14) Dalla Lana, I. G.; Karge, H. G.; George, Z. M. J. Phys. Chem. 1993, 97, 8005. (15) Saur, O.; Bensitel, M.; Mohammed Saad, A. B.; Lavalley, J.-C.; Tripp, C. P.; Morrow, B. A. J. Catal. 1986, 99, 104. (16) Wu, Q.; Gao, H.; He, H. J. Phys. Chem. B 2006, 110, 8320. (17) Kresse, G.; Furthmuller, J. Comput. Mater. Sci. 1996, 6, 15. (18) Kresse, G.; Furthmuller, J. Phys. ReV. B 1996, 54, 11169. (19) Kresse, G.; Hafner, J. Phys. ReV. B 1993, 47, 558. (20) Kresse, G.; Hafner, J. Phys. ReV. B 1994, 49, 14251. (21) Blo¨chl, P. E. Phys. ReV. B 1994, 50, 17953. (22) Perdew, J. P.; Burke, K.; Ernzerhof, M. Phys. ReV. Lett. 1996, 77, 865. (23) Methfessel, M.; Paxton, A. T. Phys. ReV. B 1989, 40, 3616. (24) Press, W. H.; Flannery, B. P.; Teukolsky, S. A.; Vetterling, W. T. Numerical Recipes; Cambridge University Press: New York, 1986. (25) Jo´nsson, H.; Mills, G.; Jacobsen, K. W. In Classical and Quantum Dynamics in Condensed Phase Simulations; Berne, B. J., Ciccotti, G., Coker, D. F., Eds.; World Scientific: River Edge, NJ, 1998; p 385. (26) Henkelman, G.; Uberuaga, B. P.; Jo´nsson, H. J. Chem. Phys. 2000, 113, 9901. (27) Henkelman, G.; Jo´nsson, H. J. Chem. Phys. 2000, 113, 9978. (28) Pulay, P. Chem. Phys. Lett. 1980, 73, 393. (29) Well, A. F. Structural Inorganic Chemistry, 3rd ed.; Clarendon Press: Oxford, U.K., 1962. (30) Sohlberg, K.; Pennycook, S. J.; Pantelides, S. T. Chem. Eng. Commun. 2000, 181, 107. (31) Wang, J. A.; Bokhimi, X.; Morales, A.; Novaro, O.; Lopez, T.; Gomes, R. J. Phys. Chem. B 1999, 103, 299. (32) Dupree, R.; Lewis, M. H.; Smith, M. E. Philos. Mag. A 1986, 53, L17. (33) Wang, Y. G.; Bronsveld, P. M.; DeHosson, J. T. M.; Djuricic, B.; McGarry, D.; Pickering, S. J. Am. Ceram. Soc. 1998, 81, 1655. (34) Jayaram, V.; Levi, C. G. Acta. Metall. 1989, 37, 69. (35) Alvarez, L. J.; Sanz, J. F.; Capitan, M. J.; Odriozola, J. A. Chem. Phys. Lett. 1992, 192, 463. (36) Alvarez, L. J.; Leon, L. E.; Sanz, J. F.; Capitan, M. J.; Odriozola, J. A. Phys. ReV. B 1994, 50, 2561. (37) Blonski, S.; Garofalini, S. H. Surf. Sci. 1993, 295, 263. (38) Mo, S. D.; Xu, Y. N.; Ching, W. Y. J. Am. Ceram. Soc. 1997, 80, 1193. (39) Streitz, F. H.; Mintmire, J. W. Phys. ReV. B 1999, 60, 773.

10454

J. Phys. Chem. C, Vol. 114, No. 23, 2010

(40) Krokidis, X.; Raybaud, P.; Gobichon, A.-E.; Rebours, B.; Euzen, P.; Toulhoat, H. J. Phys. Chem. B 2001, 105, 5121. (41) Gutie´rrez, G.; Taga, A.; Johansson, B. Phys. ReV. B 2001, 65, 012101. (42) Vijay, A.; Mills, G.; Metiu, H. J. Chem. Phys. 2002, 117, 4509. (43) Ahuja, R.; Osorio-Guillen, J. M.; Souza de Alemida, J.; Holm, B.; Ching, W. Y.; Johansson, B. J. Phys.: Condens. Matter 2004, 16, 2891. (44) Pinto, H. P.; Nieminen, R. M.; Elliott, S. D. Phys. ReV. B 2004, 70, 125402. (45) Sun, M.; Nelson, A. E.; Adjaye, J. J. Phys. B 2006, 110, 2310. (46) Maglia, F.; Gennari, S. J. Am. Ceram. Soc. 2008, 91, 283. (47) Nortier, P.; Fourre, P.; Mohammed Saad, A. B.; Saur, O.; Lavalley, J.-C. Appl. Catal. 1990, 61, 141. (48) Digne, M.; Sautet, P.; Raybaud, P.; Euzen, P.; Toulhoat, H. J. Catal. 2004, 226, 54.

Lo et al. (49) Lippens, B. C. Ph.D. Thesis, Delft, The Netherlands, 1961. (50) Morterra, C.; Magnacca, G. Catal. Today 1996, 27, 497. (51) Kno¨zlinger, H.; Ratnasamy, P. Catal. ReV.-Sci. Eng. 1978, 17, 21. (52) Tsyganenko, A. A.; Mardilovich, P. P. J. Chem. Soc., Faraday Trans. 1996, 92, 4843. (53) Babaeva, M. A.; Tsyganenko, A. A.; Filimonov, V. N. Kinet. Katal. 1984, 25, 921. (54) Person, W. B., Zerbi, G., Eds. Vibrational Intensities in Infrared and Raman Spectroscopy; Elsevier: Amsterdam, 1982. (55) Ono, Y.; Takagiwa, H.; Fukuzumi, S. J. Catal. 1977, 50, 181. (56) Khulbe, K. S.; Mann, R. S. J. Catal. 1978, 51, 364. (57) Karge, H. G.; Tower, R. W.; Dudzik, Z.; George, Z. M. Stud. Surf. Sci. Catal. 1981, 7, 643.

JP910895G