Interaction of NO2 with BaO: From Cooperative Adsorption to Ba (NO3

In order to ensure the complete .... temperature desorption feature) after the entire BaO surface is ..... 20 wt % BaO/γ-Al2O3 is the complete agreem...
0 downloads 0 Views 310KB Size
Download PDF
J. Phys. Chem. C 2007, 111, 15299-15305

15299

Interaction of NO2 with BaO: From Cooperative Adsorption to Ba(NO3)2 Formation Cheol-Woo Yi, Ja Hun Kwak, and Ja´ nos Szanyi* Institute for Interfacial Catalysis, Pacific Northwest National Laboratory, P.O. Box 999, MSIN: K8-80, Richland, Washington 99352 ReceiVed: May 30, 2007; In Final Form: August 9, 2007

The adsorption and reaction of NO2 on a thick (>30 ML), pure BaO film deposited onto an Al2O3/NiAl(110) substrate were investigated using temperature-programmed desorption, infrared reflection absorption spectroscopy, and X-ray photoelectron spectroscopy techniques. For the first time, it was clearly demonstrated that BaO reacts with NO2 to initially form nitrite-nitrate ion pairs by the cooperative adsorption mechanism predicted by theoretical calculation. These NO2-/NO3- ion pairs readily form even at 90 K. In the decomposition process of these ion pairs, first the nitrite species release an NO molecule and then form BaO2. At higher temperatures, nitrate species decompose in two steps: at lower temperature as NO2 only, then, at higher temperature, as NO + O2. The results of NO2 adsorption/reaction on this model system are identical to those we have found on a high surface area 20 wt % BaO/γ-Al2O3 sample with the exception of surface nitrates that were only observed on the high surface area material.

Introduction The interaction of NOx molecules with basic oxides is of both fundamental scientific and technological interest. Both NO and NO2 contain an unpaired electron that plays a significant role in their bindings to oxide surfaces. The combined photoemission and density functional studies of Rodriguez et al.1,2 on the interaction of NO and NO2 on MgO(100) indicated that NO adsorbed very weakly on the defect-free MgO(100) surface. On the other hand, NO2 was found to be very reactive on the MgO(100), readily forming nitrates. They attributed this large difference in reactivity of NO and NO2 to the fact that NO2 was a good electron acceptor. Density functional theory (DFT) calculations of Schneider and co-workers show that upon their interaction with the basic MgO surface NOx molecules obtain an electron from the surface to form NOx- ions and simultaneously create an electronic defect in the oxide.3-6 This oxide defect is then healed by its interaction with another NOx molecule when an electron is transferred to the oxide surface and NOx+ forms. The adsorption of these ion pairs (Lewis acid, NOx+, and Lewis base, NOx-) is energetically much more favorable than the adsorption of neutral NOx molecules. The cooperative adsorption of these ion pairs can overcome the energy cost required for the charge separation. Similarly to the case of NOx on MgO, the calculations of Gro¨nbeck et al. show that the nitrite/nitrate ion pairs formed in the adsorption of NO2 on BaO interacts through the surface, while direct lateral interactions are insignificant.7 The adsorption of NOx molecules (NO and NO2) has been extensively studied on both model systems8-10 and on high surface area materials.11-18 In particular, systems containing BaO received most of the attention due to its potential use in lean NOx abatement technologies.17 Schmitz and Baird examined the adsorption of both NO and NO2 on a BaO/oxidized aluminum sheet model NOx storage system using exclusively X-ray photoelectron spectroscopy (XPS) technique. They found that both NO and NO2 adsorbed readily on the BaO film. Upon NO exposure, nitrites formed by molecular adsorp* Corresponding author. E-mail: [email protected].

tion; while in NO2 adsorption, nitrates were the primary species and were proposed to form through nitrite intermediates.9 Contrary to these findings, Tsami et al.10 showed that a BaO film (5-6 ML) deposited onto a Cu(111) substrate did not readily store NO. However, they agreed with Schmitz and Baird9 that the initial step in NO2 adsorption was nitrite formation followed by the oxidation of nitrites to nitrates. They conclude that “there is no direct formation of nitrate as a result of exposure of barium oxide to NO2”. These findings seem to corroborate with those obtained on high surface area BaO/Al2O3 and Pt/ BaO/Al2O3 systems studied primarily by Fourier transform infrared (FT-IR) spectroscopy.11-13,15,16 From the available literature data we can draw the following conclusions: (i) NO2 is much more reactive than NO on BaO; (ii) both nitrite and nitrate species are formed upon the exposure of BaO to NO2; and (iii) upon saturation of BaO with NO2 mostly nitrates are formed. The critical issues that are still debated are the nature of NOx species formed upon the initial exposure of BaO to NO2, the formation mechanisms of NOx species under different conditions, and the decomposition mechanism of NOx species formed in the NOx trapping process. There also seems to be a large gap between the pictures emerging from theoretical calculations and from experimental work regarding the adsorption of NO2 on base metal oxide surfaces. In this study, in order to exclude the effect of the underlying substrate, we focus on the interaction of NO2 with a thick (>30 ML) BaO film deposited onto an Al2O3/NiAl(110) substrate using temperature-programmed desorption (TPD), XPS, low-energy ion scattering spectroscopy (LEISS), and infrared reflection absorption spectroscopy (IRAS) techniques. Our results provide the first experimental confirmation of the predictions of theoretical calculations regarding the pairwise, cooperative adsorption of NO2 on BaO at low exposures and the subsequent conversion of nitrites to nitrates at elevated temperatures and/or high NO2 exposures. The second part of this study addresses the comparison of the results obtained in NO2 adsorption experiments on a thick BaO film and a high surface area 20 wt % BaO/γ-Al2O3 catalyst. Our findings also

10.1021/jp074179c CCC: $37.00 © 2007 American Chemical Society Published on Web 10/03/2007

15300 J. Phys. Chem. C, Vol. 111, No. 42, 2007 enable us to provide probable explanations for the widely differing conclusions regarding the interaction of NO2 with BaO. Experimental Section All the results presented here were obtained in a combined ultrahigh vacuum (UHV) surface analysis chamber and elevated pressure reactor system with a base pressure less than 2.0 × 10-10 Torr. The surface analysis chamber is equipped with XPS (using Al KR X-ray source, and a hemispherical electron energy analyzer from Omicron), Auger electron spectroscopy (AES), LEISS, and low-energy electron diffraction (LEED). Capabilities for metal and gas dosing, as well as a quadrupole mass spectrometer (QMS) for gas analysis (UTI 100C) are also included. Both the elevated pressure cell and the gas dosing system of the UHV chamber are connected to a gas handling system for reproducible gas introduction. The reactor cell is also equipped with an IRAS setup (Bruker, Equinox 55). The substrate used in this study was a thin alumina film prepared on a NiAl(110) metal alloy single crystal (Princeton Scientific Corp., 10 mm × 2 mm disc), and it was spot-welded onto a U-shaped Ta wire. A C-type thermocouple was spotwelded to the back side of the crystal for temperature measurement. The substrate was cleaned by alternating cycles of Ar+ ion sputtering and prolonged annealing at 1200 K. Sample cleanliness and surface order were checked by AES, XPS, and LEED. After a clean NiAl(110) substrate was prepared, an ultrathin alumina film was formed by following the procedures of references 19 and 20. The alumina film thickness was estimated to be around 2 monolayer (ML), and the surface order was verified by LEED. Ba deposition was carried out by reactive layer-assisted deposition (RLAD) using a resistively heated Ba doser (SAES Getters Inc.) onto a thick (>30 ML) N2O4 multilayer as a reactive layer. TPD experiments were conducted in the UHV chamber after dosing NO2 onto the BaO film through a pinhole doser. To minimize reactions of NO2 with the gas lines during actual dosing experiment, prior to these NO2 adsorption experiments the metal parts of the dosing system were passivated under NO2 flow for an extended period of time. The conditioning was stopped when the 46 amu signal (NO2) intensity in the mass spectrometer reached a constant level. NO2 dosing was carried out in a wide sample temperature range of 90-300 K. The sample was placed in front of (∼1 mm away) the long (∼12 in.) metal tube ensuring the collimation of the gas beam. The dosing amount was controlled by either changing the deposition time or the NO2 back pressure. After adsorption the sample was moved in front of a QMS that was sitting in a metal housing negatively biased (-70 V) to prevent the interaction between the adsorbates and the electrons from the ionization chamber. Then the temperature of the sample was raised linearly to 900 K with a heating rate of either 1 or 2 K/s, and the mass spectrometer signal intensities of multiple masses were recorded as a function of sample temperature. The same gas dosing system was set up on the elevated pressure reactor cell. IRAS spectra were collected in fast scan mode allowing the collection of 4096 spectra in less than 5 min. All the IR spectra collected were referenced to a background spectrum acquired from the clean sample prior to gas adsorption. Results and Discussion Before we discuss any of the results of the NO2 adsorption experiments, we need to address two key issues about our sample, namely the thickness and the preparation method of the BaO film. First of all, a thick (>30 ML) BaO film was

Yi et al. prepared and investigated to avoid any interference on the NOx chemistry by the underlying alumina and/or NiAl(110) substrate. We believe that most of the variations in NOx chemistry over the BaO-containing systems reported in the literature are caused by the influence of the substrate on which the base-metal oxide is supported. This effect was clearly shown in the study of NO2 adsorption on a BaO (5-6 ML)/Cu(111) system by Tsami et al.10 They observed that the chemistry significantly varied by the history of the sample, i.e., the TPD profiles following NO2 adsorption were very different from those of the freshly prepared samples and the samples that went through one or two cycles of adsorption/desorption. This was attributed to the reaction between BaO and CuOx. We also suspect that some of the results reported by Schmitz et al.9 were influenced by the reaction between the thin BaO and the aluminum oxide film which it was deposited onto, namely by the formation of a Ba-aluminate phase. We have clearly seen this effect in the BaO/Al 2O3/NiAl(110) system at lower BaO coverages, and the results obtained from these systems will be published in our subsequent contribution.21 We wanted to be absolutely sure that the chemistry we observed was characteristic of BaO itself and was not influenced by the underlying alumina film and/or NiAl(110) substrate in any way. The thick BaO film was prepared by deposition of Ba onto a thick N2O4 ice as reactive layer. During Ba deposition onto the N2O4 ice layer, BaNOx formed instantaneously and was converted to Ba(NO3)2 upon the thermal removal of the excess N2O4.(In order to ensure the complete conversion of the BaNOx layer to Ba(NO3)2, a thick N2O4 ice layer was deposited on top of the BaNOx film prior to annealing.) Annealing the sample to 1000 K resulted in the decomposition of Ba(NO3)2 and the formation of a thick BaO film (>30 ML). LEISS measurements following the preparation of the thick BaO film revealed the presence of only Ba and O on the surface of the sample (results are not shown for brevity), corroborating with the results of XPS analysis. A. NO2 Adsorption at 90 K. The adsorption of NO2 was investigated on the thick BaO film at 90 K sample temperature. The series of TPD spectra (30 amu fragment) is displayed in Figure 1 in the 100-900 K temperature range (heating rate ) 2 K/s). (Although Figure 1 only displays the TPD trace for the 30 amu signal, other mass fragments (14, 16, 28, 32, 44, 46) were also monitored simultaneously. This protocol is critical to identify the chemical nature of the species desorbed, as it is discussed below.) At the lowest NO2 exposure applied there is a broad, low-intensity desorption peak (two overlapping peaks) between 470 and 610 K. With increasing amount of NO2 doses the intensities of both of these features increased, and at a certain, higher NO2 dose a third feature started developing at an even higher temperature. The lower temperature peak (centered at about 540-575 K) originated from the desorption of NO only, while the higher temperature desorption feature (centered at 650 K) represented NO accompanied by desorption of O2. In between these two peaks, a NO2 desorption feature (centered at ∼610 K) was observed. The results of the deconvolution of TPD data is shown in the inset in Figure 1, suggesting the existence of three different surface species (spectral deconvolution was carried out following the determination of the cracking patterns of the NOx species in our mass spectrometer). As the intensities of all these three peaks increased, the maximum intensity of the lowest temperature feature gradually shifted to higher temperatures and overlapped with the second desorption peak. Note that no low temperature (800 K). These results suggest that bridge-bound nitrites (between two Ba2+ sites) decompose by releasing a NO molecule and leaving an O- on the surface bound to a Ba2+ ion, resulting in the formation of BaO2. The results presented in Figures 1-6 allow us to propose a mechanism for the reaction of NO2 with a clean BaO surface and also for the decomposition path of the ionic NOx species formed. Upon the exposure of a thick BaO film to NO2 at 90 K, via pairwise adsorption, NO2--NO3- ion pairs form and are stabilized by the cooperative bonding effect. The electron transfer between the two NO2 molecules results in a very strong Lewis acid (NO2+)/Lewis base (NO2-) pair adsorption on the base metal oxide (BaO) surface. The Lewis acid component of the ion pair (NO2+) strongly interacts with a surface O2- ion to form a surface nitrate. The most probable surface configuration of the thus-formed nitrite-nitrate ion pairs is NO2 bridge-bound to two Ba2+ ions through its two oxygen atoms with another NO2 molecule binding to a lattice O2- ion through its N atom. These proposed structures are identical to those that were predicted by Schneider for the cooperative adsorption of NO2 on MgO3 and also by Gro¨nbeck et al. for BaO.7 On the (001) plane of MgO the combined adsorption energy of this ion pair was calculated to be 15 kcal/mol higher than that estimated for the adsorption of two neutral NO2 molecules.3 In the case of the more basic BaO substrate a similar stabilization energy (“cooperative enhancement”) was estimated. However, the relative magnitude of this enhancement was much smaller in the case of BaO (∼20%) in comparison to MgO (∼100%). This is the result of the origin of the cooperative enhancement, the difference in the energy spent on charge separation to form the two charged adsorbates, and the adsorption energy gain resulting from the stronger adsorption of the ionic adsorbed species. All of our results suggest that initially the NO2 molecules impinging onto the clean BaO surface readily form the nitrite-nitrate ion pairs. However, the IRAS data also show that weakly adsorbed (physisorbed) N2O4 is present on the surface even before the entire surface is saturated with these ion pairs. This is the consequence of the low mobility of the NO2 molecules on the surface. As the BaO surface becomes partially covered with ionic NOx species, the probability of an incoming NO2 molecule to land on an empty adsorption site decreases. The surface becomes completely saturated with the ionic NOx species as the N2O4 multilayer covers the entire BaO surface. In the TPD experiment we do not see this due to the inherent nature of this technique; during the heating of the sample the weakly held molecular NO2 gains enough mobility to move to empty BaO sites and form the ionic NO2-/NO3- pairs before they can desorb. This was confirmed by temperature-dependent IRAS experiments for partially ionic NOx-covered BaO as shown in Figure 3. If the above-described mechanism for the cooperative adsorption of NO2 on BaO is true, the amount of NO2 adsorbed as nitrites should equal that of nitrates. In fact that is exactly

15304 J. Phys. Chem. C, Vol. 111, No. 42, 2007 what we observed. The intensities of N 1s XPS signals of nitrites and nitrates after the 200 K anneal are practically identical. Furthermore, the area of the lowest temperature desorption peak in TPD (nitrite decomposition) is very close to the sum of the two higher temperature desorption features (nitrate decomposition, see inset of Figure 1). Interestingly the thermal stabilities of the two ionic NOx species are different. During adsorption nitrites and nitrates form simultaneously in a pairwise manner. However, in the desorption experiments first nitrites decompose, releasing an NO molecule and forming BaO2 with the extra O- ion. Then nitrate decomposition takes place at higher temperatures in two different ways. At lower temperatures nitrates decompose as NO2, while at higher temperatures it proceeds through the formation of NO + O2. In this decomposition mechanism the nitrate species that are still present on the surface seem to be stabilized by the O- ion left behind the nitrite decomposition by facilitating the charge separation necessary for the cooperative enhancement. We should mention here that although O 1s XPS data clearly show that the BaO2 decomposes at high temperatures (after all the NOx are gone), in our TPD experiments we did not observe a corresponding O2 desorption feature. One plausible explanation for this observation is the migration of the oxygen toward the BaO/Al2O3 interface and its reaction with the Al atoms in the substrate NiAl alloy. The result of this process should be the increase in the Al2O3 film thickness. In fact, a significant increase in alumina film thickness was observed by XPS and LEISS when the thick BaO film was sputtered by Ar+ ions following completion of the NOx uptake experiments (data not shown for brevity). B. Comparison to High Surface Area NOx Storage Systems. In order to compare the results of the initial NO2 uptake on BaO studied here, and a practically more relevant NOx storage system, we investigated the adsorption of NO2 on a high surface area 20 wt % BaO/γ-Al2O3 sample at cryogenic temperatures. The preparation method and characterization of this sample have been described in detail in our previous publications.18 Here we focus on the identification of NOx species formed on the powder BaO/γ-Al2O3 sample using transmission FTIR spectroscopy (for experimental details, see ref 18). After preparation of 20 wt % BaO/γ-Al2O3, the sample was cooled with liquid nitrogen to 150 K, and then NO2 was introduced into the IR cell through a leak valve. Due to this background type of dosing, most of the NO2 introduced condensed onto the sample holder rod containing the liquid nitrogen, and only a very small fraction of the NO2 adsorbed onto the sample. However, even at the lowest sample temperature (150 K) we can see the appearance of IR features characteristic for both nitrates and nitrites. Following spectral collection at 150 K we started purging the liquid-nitrogen-cooled sample holder rod with N2, which resulted in the desorption of NO2 from the rod (that practically served as a cryostat) and adsorption onto the sample, which was now at a lower temperature than the rod. During this process we acquired IR spectra at every 15-20 K temperature increment. The series of IR spectra obtained in this experiment is displayed in Figure 7. As NO2 desorbed from the sample holder rod, it condensed onto the sample, resulting in large intensity gains of the IR features characteristic of N2O4 ice. Heating the sample to higher temperatures resulted in the loss of intensity of the N2O4 vibrational features, while those of the nitrite-nitrate species formed increased. At the highest temperature studied (245 K), all the N2O4 ice desorbed from the sample, and only the strongly held ionic NOx species were seen. The nitrate and nitrite species

Yi et al.

Figure 7. IR data collected during low-temperature NO2 adsorption on high surface area 20 wt % BaO/γ-Al2O3 catalysts.

formed were very similar to those we have seen for the thick BaO film. Comparing the IR data presented in Figures 2 and 7, we can conclude that the same type of nitrite-nitrate pairs formed on both the high surface area material and the BaO film. As we have already mentioned briefly, the main difference between the two sets of data is the presence of IR features attributed to surface nitrates on the high surface area sample. These bands are located at around 1570 and 1300 cm-1. The 1570 cm-1 peak is completely absent in the spectra collected from the NO2-exposed BaO film, consistent with the absence of the exposed BaO/alumina interface where the BaO monolayer forms and adsorbs NO2 as surface nitrates. Also, the intensity of the broad peak centered around 1320 cm-1 is much higher on the high surface area material than on the BaO film, due to the overlap of the surface and bulk nitrate features. The main conclusion regarding the initial adsorption of NO2 on BaO and 20 wt % BaO/γ-Al2O3 is the complete agreement on the formation mechanism of strongly held NOx species. NO2 exposure of both samples results in the formation of nitritenitrate pairs by the cooperative adsorption mechanism. This conclusion contradicts the findings of a number of previously published works, in particular the one of Tsami et al.10 In their study on a thin (5-6 ML) BaO film supported on Cu(111), they found that NO2 exposure resulted in the formation of nitrites exclusively, but they found no evidence for direct nitrate formation. We argue that these findings were probably strongly influenced by the underlying Cu substrate, just as they claim for the decomposition of the thus-formed nitrites. Any level of Cu contamination of this very thin BaO film will result in significant changes in the NOx chemistry, just as we have seen for the BaO/Al2O3 system.21 Their TPD and XPS results clearly show the instability of the system studied and the irreproducibility of the results in subsequent NO2 exposures. The results of Schmitz and Baird9 are also inconsistent with our findings, as they invoke the presence of nitrites only on the BaO surface initially and its subsequent oxidation to nitrates by NO2. Since they use a thin BaO film (6-8 ML) on an oxidized aluminum sheet, the formation of Ba-aluminate is very likely. Indeed, we have observed the formation of Ba-aluminate for thin BaO films on alumina that completely changed their NO2 uptake properties. In particular, they always exhibited the

Interaction of NO2 with BaO initial formation of nitrites only, instead of the nitrite-nitrate pairs that we have proven to form on the pure BaO.21 Conclusions One of the most important aspects of this work is the confirmation of the cooperative NO2 adsorption mechanism on a clean, thick BaO film supported on Al2O3/NiAl(110) substrate. Exposure of BaO to NO2 at 90 K resulted in the formation of nitrite-nitrate pairs, suggesting that the activation energy of the ion pair formation is very low. Upon heating the nitritenitrate pair-covered BaO, nitrites decompose first, leaving one O atom behind that takes part in the formation of BaO2. The thus-formed BaO2 is present even after all the nitrates are decomposed. Nitrates decompose in two steps: at low temperature as NO2 only, and at higher temperature as NO + O2. In addition, the results of NO2 adsorption/reaction on this model system (thick BaO film supported on Al2O3/NiAl(110)) are in complete agreement with those we have found on a high surface area 20 wt % BaO/γ-Al2O3 sample with the exception of surface nitrates that were only observed on the high surface area material. Acknowledgment. We gratefully acknowledge the U.S. Department of Energy (DOE), Office of Basic Energy Sciences, Division of Chemical Sciences, for the support of this work. The research described in this paper was performed at the Environmental Molecular Sciences Laboratory (EMSL), a national scientific user facility sponsored by the DOE Office of Biological and Environmental Research and located at Pacific Northwest National Laboratory (PNNL). PNNL is operated for the U.S. DOE by Battelle Memorial Institute under contract number DE-AC05-76RL01830. We also thank Drs. B. D. Kay and R. S. Smith of PNNL for lending us the FT-IR spectrometer.

J. Phys. Chem. C, Vol. 111, No. 42, 2007 15305 References and Notes (1) Rodriguez, J. A.; Jirsak, J.; Kim, J. Y.; Larese, J. Z.; Maiti, A. Chem. Phys. Lett. 2000, 330, 475. (2) Rodriguez, J. A.; Jirsak, T.; Sambasivan, S.; Fischer, D.; Maiti, A. J. Chem. Phys. 2000, 112, 9929. (3) Schneider, W. F. J. Phys. Chem. B 2004, 108, 273. (4) Miletic, M.; Gland, J. L.; Hass, K. C.; Schneider, W. F. Surf. Sci. 2003, 546, 75. (5) Miletic, M.; Gland, J. L.; Hass, K. C.; Schneider, W. F. J. Phys. Chem. B 2003, 107, 157. (6) Schneider, W. F.; Hass, K. C.; Miletic, M.; Gland, J. L. J. Phys. Chem. B 2002, 106, 7405. (7) Gronbeck, H.; Broqvist, P.; Panas, I. Surf. Sci. 2006, 600, 403. (8) Ozensoy, E.; Peden, C. H. F.; Szanyi, J. J. Catal. 2006, 243, 149. (9) Schmitz, P. J.; Baird, R. J. J. Phys. Chem. B 2002, 106, 4172. (10) Tsami, A.; Grillo, F.; Bowker, M.; Nix, R. M. Surf. Sci. 2006, 600, 3403. (11) Broqvist, P.; Gronbeck, H.; Fridell, E.; Panas, I. J. Phys. Chem. B 2004, 108, 3523. (12) Fanson, P. T.; Horton, M. R.; Delgass, W. N.; Lauterbach, J. Appl. Catal. B 2003, 46, 393. (13) Hess, C.; Lunsford, J. H. J. Phys. Chem. B 2002, 106, 6358. (14) Mahzoul, H.; Brilhac, J. F.; Gilot, P. Appl. Catal. B 1999, 20, 47. (15) Nova, I.; Castoldi, L.; Lietti, L.; Tronconi, E.; Forzatti, P.; Prinetto, F.; Ghiotti, G. J. Catal. 2004, 222, 377. (16) Prinetto, F.; Ghiotti, G.; Nova, I.; Lietti, L.; Tronconi, E.; Forzatti, P. J. Phys. Chem. B 2001, 105, 12732. (17) Su, Y.; Amiridis, M. D. Catal. Today 2004, 96, 31. (18) Szanyi, J.; Kwak, J. H.; Kim, D. H.; Burton, S. D.; Peden, C. H. F. J. Phys. Chem. B 2005, 109, 27. (19) Franchy, R. Surf. Sci. Rep. 2000, 38, 195. (20) Stierle, A.; Renner, F.; Streitel, R.; Dosch, H.; Drube, W.; Cowie, B. C. Science 2004, 303, 1652. (21) Yi, C. W.; Kwak, J. H.; Peden, C. H. F.; Wang, C. M.; Szanyi, J. J. Phys. Chem. C, in press. (22) Sedlmair, C.; Seshan, K.; Jentys, A.; Lercher, J. A. J. Catal. 2003, 214, 308. (23) Westerberg, B.; Fridell, E. J. Mol. Catal. A: Chem. 2001, 165, 249. (24) Cotton, F. A.; Wilkinson, G. AdVanced inorganic chemistry: a comprehensiVe text, 4th ed.; Wiley: New York, 1980. (25) Ganguly, P.; Hegde, M. S. Phys. ReV. B 1988, 37, 5107.