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Catalytic Decomposition of Ba(NO3)2 on Pt(111) Kumudu Mudiyanselage,† Jason F. Weaver,‡ and Janos Szanyi*,† †
Institute for Interfacial Catalysis, Pacific Northwest National Laboratory. P.O. Box 999, MSIN: K8-87, Richland, Washington 99352, United States ‡ Department of Chemical Engineering, University of Florida, Gainesville, Florida 32611, United States ABSTRACT: The decomposition of Ba(NO3)2 formed on BaO/Pt(111) (Pt(111) surface is partially covered by BaO) in the presence of CO was studied using temperature-programmed desorption, infrared reflection absorption, and X-ray photoelectron spectroscopies. The exposure of BaO/Pt(111) to elevated NO2 pressure (1.0 10-4 torr) at 450 K leads to the formation of Ba(NO3)2, chemisorbed O (OPt), and Pt-oxide-like domains. During TPD, the Ba(NO3)2 begins to thermally decompose near 490 K, releasing NO and NO2 with the maximum NOx desorption rate seen at 605 K. The OPt species formed following the exposure of BaO/Pt(111) to NO2 react with CO to release CO2 at 450 K. The consumption of OPt during CO oxidation initiates the migration of O from the Pt-oxidelike domains to the chemisorbed phase, where the CO oxidation reaction occurs. Therefore, the removal of OPt by CO leads to the reduction of oxidized Pt and to the formation of metallic Pt(111) domains, where, subsequently, catalytic decomposition of Ba(NO3)2 can take place. The Pt-catalyzed decomposition of Ba(NO3)2 occurs readily at 450 K, a temperature much lower than the onset of the decomposition temperature of Ba(NO3)2 in the presence of oxidized Pt.
1. INTRODUCTION One of the main steps in the NOx storage and reduction (NSR) cycle is the regeneration of the catalyst by releasing NOx from the stored nitrates, a prerequisite to its reduction to N2. Under practical conditions it is not straightforward to decouple these two sequential steps (release and reduction of NOx) in order to unambiguously characterize the NOx release process.1 However, a few studies with model catalyst systems have focused on the NOx release step of the NSR process, and mainly three primary driving forces are believed to be responsible for the NOx release from the stored nitrates:1 (i) Generation of heat from oxidation of the reductants by the oxygen present in the gas stream or stored on the surface of the catalyst. This heat generation due to the exothermic reactions (oxidation of reductants) under fuel-rich conditions increases the surface temperature of the catalyst, which leads to the release of NOx as the thermal stability of the stored nitrates decreases with increasing temperature. (ii) Creation of a net reducing environment under fuel-rich conditions. The NOx trapping reaction is believed to be partially driven by equilibrium.1 Therefore, under reduction conditions (low or zero oxygen partial pressures) in the fuel-rich period, the equilibrium stability of nitrate species is reduced. This drives the equilibrium toward the decomposition of nitrates and the release of NOx from the catalyst surface according to the following reactions: 2BaðNO3 Þ2ðsÞ T 4NO2ðgÞ þ O2ðgÞ þ 2BaOðsÞ BaðNO3 Þ2ðsÞ T NO2ðgÞ þ NOðgÞ þ O2ðgÞ þ BaOðsÞ 2BaðNO3 Þ2ðsÞ T 4NOðgÞ þ 3O2ðgÞ þ 2BaOðsÞ Liu et al. postulated that the removal of O2 from the gas phase destabilizes the nitrate species and induces NOx release.2 They also suggested that CO2 and H2O formed in the reactions r 2011 American Chemical Society
between the reductant and surface or gaseous oxygen decreased the stability of nitrates by their conversion to hydroxides or carbonates. However, this is believed to be a minor pathway. (iii) Reduction of the oxidized Pt under fuel-rich conditions. The oxidized Pt formed during the lean (air-rich) period is reduced under fuel-rich conditions, which leads to the decomposition of Ba(NO3)2 as Pt catalyzes this process.3-5 Pt-catalyzed barium nitrate decomposition has been substantiated experimentally on Pt-impregnated barium nitrate4 and with a high surface area model catalyst.4,5 Under flow conditions at atmospheric pressure, the Pt-catalyzed barium nitrate decomposition occurs at ∼650 K, whereas the thermal decomposition (i.e., in the absence of metallic Pt) only begins above 750 K. Coronado et al. also reported the Pt-facilitated nitrate decomposition by comparing nitrates formed on BaCl2/SiO2 and Pt/BaO/SiO2.6 Under identical conditions, heating at 573 K in a static vacuum, nitrate decomposition was observed on Pt/Ba(NO3)2/SiO2 while decomposition was not seen with Ba(NO3)2/SiO2. However, a physical mixture of Pt/SiO2 and BaCl2/SiO2 behaves like Pt-free samples indicating that the decomposition is easier when Pt and Ba are in close proximity. Therefore, they proposed that the decomposition of nitrates might take place at the interface between the barium compound and Pt, leading to the release of NO2 and to the formation of a surface oxygen layer on Pt according to the following reaction scheme: 3Pt þ 2NO3 - f 2ON-Pt-O þ PtO þ O2f 2Pt þ Pt-O þ 2NO2 þ O2They also suggested that this process would hinder the progress of nitrate decomposition in the absence of a reducing Received: December 15, 2010 Revised: February 3, 2011 Published: March 08, 2011 5903
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gas such as propene due to the accumulation of adsorbed oxygen on the Pt surface.6 Even though the mechanisms of NOx release are relatively well accepted, decoupling the individual processes is challenging due to several nitrate decomposition pathways simultaneously in operation under practical conditions. Therefore, in this study we performed experiments on a BaO/Pt(111) model NOx storage system under well-controlled conditions, which allowed us to investigate the Pt-catalyzed decomposition of nitrates. Although many studies have investigated the effects of oxygen (surface or gaseous) on the decomposition of stored nitrates on high surface area model catalysts,1 to the best of our knowledge, such studies on single-crystal-based model systems have not been reported previously. Here we present the results of an experimental investigation aimed at understanding the Pt-catalyzed nitrate decomposition in the presence of gas phase CO.
2. EXPERIMENTAL SECTION All the experiments were performed in a combined ultrahigh vacuum (UHV) surface analysis chamber and elevated-pressure reactor/infrared reflection absorption spectroscopy (IRAS) cell system with a base pressure of less than 2.0 10-10 torr (1 torr = 1.3332 mbar). The UHV chamber is equipped with X-ray photoelectron spectroscopy (XPS), Auger electron spectroscopy (AES), low-energy electron diffraction (LEED), and temperature-programmed desorption (TPD) techniques. The elevatedpressure reactor cell is coupled to a commercial Fourier transform infrared (FT-IR) spectrometer (Bruker, Vertex 70) for IRAS experiments. The Pt(111) single-crystal (10 mm diameter, 2 mm thick, Princeton Scientific) used in these experiments was spot-welded onto a U-shaped Ta wire, and the sample temperature was measured by a C-type thermocouple spot-welded to the backside of the crystal. The Pt(111) crystal was cleaned by repeated cycles of Arþ ion sputtering and annealing in O2 at 800 K. The cleanliness of the surface was verified with AES, XPS, and LEED. The BaO clusters were prepared by reactive layer-assisted deposition (RLAD); first the desired amount of Ba was deposited onto a N2O4 multilayer on the Pt(111) crystal at 90 K by physical vapor deposition using a resistively heated Ba doser (SAES Getters), and then the thus-formed BaNxOy layer was thermally decomposed by annealing to 1000 K. In order to determine the fraction of the Pt(111) surface covered by BaO clusters, CO TPD experiments following room temperature adsorption were performed on BaO/Pt(111). The integrated CO desorption peak areas from the BaO/Pt(111) systems were compared to that obtained from Pt(111) to estimate the fraction of BaO-free Pt(111) surface as described previously.7 This comparison showed that approximately 50% of the Pt(111) surface of this model system was covered with BaO. The reactants were introduced into the UHV chamber through pinhole dosers and delivered to the sample surface through collimating tubes. The same gas dosing system was set up in the elevated-pressure reactor/IRAS cell. This setup allows us to expose the sample to the desired amount of gas by adjusting the pressure in the gas manifold (back pressure) and/or the exposure time. A precision leak valve was used to introduce NO2 gas for the elevated pressure experiments. IR spectra were collected at 4 cm-1 resolution using a grazing angle of approximately 85° to the surface normal. Each spectrum presented is the average of 1024 scans, requiring a spectral acquisition time of 80 s. All the IR spectra collected were referenced to a background spectrum
Figure 1. (a) IR spectrum obtained following the exposure of BaO/ Pt(111) to elevated NO2 pressure (1.0 10-4 torr for 180 s) at 450 K. (b) TPD spectra for O2, NO, and NO2 corresponding to the system in Figure 1a (the NO2 spectrum is multiplied by a factor of 5). The O2 TPD spectrum obtained following the exposure of Pt(111) to NO2 under the same conditions that applied to BaO/Pt(111) is also displayed for comparison.
acquired from the clean BaO/Pt(111) sample prior to gas adsorption. After the completion of IRAS experiments in the elevated-pressure reactor/IRAS cell, the sample was moved to the UHV chamber for XPS and TPD experiments to identify the species that remained on the sample surface and to determine their coverages. The heating rate used in all the TPD experiments was 2 K s-1.
3. RESULTS AND DISCUSSION 3.1. Formation and Thermal Decomposition of Ba(NO3)2. Figure 1a shows the IR spectrum obtained following the exposure of the BaO/Pt(111) system to an elevated NO2 pressure of 1.0 10-4 torr for 180 s at 450 K sample temperature. The observed IR features at 1332 and 1468 cm-1 can be assigned to amorphous Ba-nitrate species, as reported previously.8-14 These peaks originate from the splitting of the asymmetric stretching vibration of nitrates in the amorphous Ba-nitrate phase. The formation of nitrates is further confirmed by the observed N 1s XPS peak (Figure 2a) at 406.7 eV, which agrees with the binding energy reported for nitrate species formed on similar model systems.8,12,15-17 Therefore, these IR and XP spectroscopic data clearly indicate the presence of Ba-nitrates only (no nitrites) 5904
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Figure 2. (a) N and (b) O 1s XP spectra obtained following the formation of nitrates at elevated NO2 pressure (1.0 10-4 torr NO2 for 180 s) at 450 K and subsequent annealing to the indicated temperatures for 1 s.
following the exposure of BaO/Pt(111) to elevated NO2 pressures at 450 K. Figure 1b shows the TPD traces for NO (m/e = 30), O2 (m/e = 32), and NO2 (m/e = 46) corresponding to the system in Figure 1a. The O2 TPD spectrum obtained following the exposure of Pt(111) to NO2 under the same conditions applied to BaO/Pt(111) is also displayed for comparison. It is wellknown that NO2 dissociates on the Pt(111) surface to form adsorbed NO þ O.18,19 Atomic O coverages up to 0.75 ML on Pt(111) can be obtained under UHV conditions (∼2.0 10-9 torr) by exposure to NO2 at 400 K.20,21 The NO formed during NO2 dissociation at 400 K desorbs from the surface leaving a clean, adsorbed O layer on the surface. Weaver et al. also obtained atomic O coverages up to 0.75 ML on Pt(111) by exposing NO2 at 450 K.22 The O2 TPD curve obtained from the 0.75 ML O/Pt(111) sample, which was prepared by exposure of a clean Pt(111) surface to NO2 under UHV conditions at 400 or 450 K, shows three desorption features at ∼550, 620, and 700 K.20,22 However, the O2 TPD spectrum observed following the exposure of Pt(111) to elevated NO2 pressure (∼ 1.0 10-4 torr) at 450 K shows a relatively sharp single peak at 625 K followed by a broad trailing edge (Figure 1b). The shape of this O2 desorption feature is similar to the spectrum observed by Weaver et al. following the formation of Pt oxide domains on Pt(111) at 1.26 ML atomic O coverage, obtained by delivering an
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atomic O beam to the surface at 450 K.23 In that case, the sharp peak arises from the autocatalytic decomposition of Pt oxide domains, while the broad tail corresponds to recombinative desorption of oxygen atoms that are chemisorbed on metallic domains. The O coverage of 0.75 ML and the temperature of 625 K at the maximum desorption rate observed here are lower than those reported (1.26 ML and ∼670 K) by Weaver et al. Note that as the O coverage increases above 0.75 ML, the O2 desorption peak begins to shift to higher temperature and intensifies continuously with O coverage.23,24 Our results seem to indicate that the O phase formed at elevated NO2 pressure is different form that observed following the exposure of Pt(111) to NO2 under UHV conditions at 450 K, although very similar O coverages (∼0.75 ML) were achieved in both cases. Considering the similarity between the O2 desorption spectrum observed in this study and those reported previously for partially oxidized Pt(111), we conclude that the O phase formed here contains Pt-oxide-like domains (OPtO). The formation of this Pt-oxidelike phase only at elevated NO2 pressures indicates that the NO2 pressure greatly influences the nature of the O phase formed on Pt(111). NO2 also dissociates to form atomic O on the BaO/Pt(111) system due to the presence of bare Pt(111) sites for the adsorption of NO2 because ∼50% of the surface exposes Pt(111) domains, as described in the Experimental section. Therefore, the exposure of BaO/Pt(111) to elevated NO2 pressure at 450 K leads to the formation of a system with Pt-related O species (chemisorbed (OPt) and oxide-like (OPtO)) as well as Ba(NO3)2 particles with oxygen at the boundary of nitrate/Pt(111) [Ba(NO3)2/O/Pt(111)]. The thermal decomposition of these Ba-nitrate particles produces BaO, NOx, and O2. Desorption of the NOx species initiates at about 490 K and gives a peak centered at 605 K as shown in Figure 1b. Since most of the NOx desorbs before O2 desorption occurs appreciably, we conclude that the Ba-nitrate particles are slightly less stable than the coexisting Pt-oxide-like domains and thus decompose to a significant extent in the presence of the surrounding Pt-oxidelike phase. The desorption of O2 from Ba(NO3)2/O/Pt(111) occurs in at least three steps, with maximum desorption rates observed at 625, 721, and 834 K. A comparison of O2 desorption curves obtained from O/Pt(111) and Ba(NO3)2/O/Pt(111) (prepared under identical conditions) suggests that OPt and OPtO on the bare Pt(111) sites of BaO/Pt(111) contribute to the desorption feature observed at 625 K. Oxygen atoms originating from the decomposition of Ba(NO3)2 can adsorb: (i) on the bare Pt(111) sites (OPt), (ii) at the boundary between BaO/Pt or the perimeter of BaO particles (Oint) (around the BaO particles), and (iii) at the interface between Pt and BaO (OBaO/Pt) (under the BaO particles). The oxygen atoms derived from the decomposition of Ba(NO3)2 may also contribute to the sharp O2 desorption peak at 625 K as well as the trailing edge of the O2 TPD spectrum. The O2 desorption feature at 721 K can be assigned to adsorbed O at the boundary of BaO/Pt or perimeter of BaO particles (Oint). Our previous study has shown that O atoms at the boundary of BaO/Pt or at the perimeter of BaO particles are bound to the surface more strongly than to the bare Pt(111) sites.7 However, the temperature of 721 K at the maximum desorption rate observed here is lower than that observed in our previous study (i.e., ∼760 K). Note that the conditions applied in these two studies were very different (O2 vs NO2 as oxidants, UHV vs elevated pressure conditions, and 300 5905
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Figure 3. Changes in the 28 (CO), 30 (NO), 32 (O2), and 44 (CO2) m/e signals as a function of CO exposure time when the Ba(NO3)2/O/ Pt(111) sample kept at 450 K temperature was exposed to a constant CO flux (Pco(back) = 4.0 torr). The Ba(NO3)2/O/Pt(111) system was prepared by following the exposure of BaO/Pt(111) to elevated NO2 pressure (1.0 10-4 torr for 180 s) at 450 K.
K vs 450 K sample temperature during exposure to oxidant). In addition, in the present study, the BaO clusters were derived from the decomposition of Ba(NO3)2 to obtain an adsorbed O layer at the boundary of BaO/Pt, in contrast to the previous study where BaO clusters were already present before the adsorption of O2. A fraction of this Oint may also contribute to the feature centered at 625 K. The O2 desorption feature at 834 K is most likely due to desorption of oxygen from the interface between Pt and BaO (OBaO/Pt). This assignment is further supported by the low reactivity of OBaO/Pt with CO as described in the following section. Figure 2 shows N and O 1s XP spectra obtained following the formation of Ba-nitrates under elevated NO2 pressure at 450 K and subsequent annealing to the indicated temperatures. At 450 K, the N 1s XP spectrum shows only one feature at 406.7 eV assigned to nitrate species.8,12,15-17 The corresponding O 1s spectrum shows two peaks at 532.2 and 529.3 eV indicating the presence of at least two types of O. The O and N 1s features at 532.2 and 406.7 eV, respectively, disappear completely after annealing to 650 K. During the anneal to 650 K, nitrates decompose completely, releasing NOx as shown by the TPD spectra in Figure 1b. Therefore, the O 1s peak at 532.2 eV can be assigned to the O in nitrate ions. The O 1s peak observed at 529.3 eV after the formation of nitrates at 450 K is due to the OPt and OPtO and may be adsorbed O at the Pt-Ba(NO3)2 boundary and the perimeter of Ba(NO3)2 particles. (The O 1s XP peak observed following the exposure of clean Pt(111) to elevated NO2 pressure (not shown) is centered at 529.5 eV.) After annealing to 650 K, the O 1s peak at 529.3 eV shifts to 528.9 eV, representing oxygen atoms of Oint, OBaO/Pt, O2- of BaO, and a small fraction of OPt or OPtO that still remained on the surface. Annealing to 850 K results in the desorption of all the adsorbed O species, and the O 1s peak observed at 528.9 eV is due to the O2of BaO. The data in Figures 1 and 2 indicate that the thermal decomposition of nitrates begins near 490 K with a maximum NOx desorption rate at 605 K. 3.2. Isothermal Decomposition of Ba(NO3)2 in the Presence of CO at 450 K. Figure 3 shows the mass spectrometer signals for the 28 (CO), 30 (NO), 32 (O2), and 44 (CO2) amu as
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a function of CO exposure time when the Ba(NO3)2/O/Pt(111) sample kept at 450 K temperature was exposed to a constant CO flux. This Ba(NO3)2/O/Pt(111) system was prepared by exposure of the BaO/Pt(111) system to an elevated NO2 pressure (1.0 10-4 torr) at 450 K (the same way we have prepared the sample for the thermal decomposition study; results shown in Figures 1 and 2). The data in Figure 3 shows the release of NO and CO2 and the consumption of CO simultaneously during the exposure of Ba(NO3)2/O/Pt(111) to CO. The production of CO2 during this experiment arises from the oxidation of CO on the Pt domains since the Ba-nitrate particles are inactive toward CO oxidation. As seen in Figure 3, the CO oxidation rate is initially very low and steady during the CO exposure, indicating negligible CO2 production during this time. After the CO exposure has continued for ∼100 s, the CO2 production rate increases sharply toward a maximum before decaying back to the baseline. This behavior is analogous to that reported in a prior study of CO oxidation on partially oxidized Pt(111) in which the onset of reaction was seen to occur well after the CO exposure to the surface was initiated.24 In this prior study, an autocatalytic mechanism for CO oxidation was proposed to explain the time delay prior to the rate maximum. This model considers that CO adsorption and oxidation occur only on metallic Pt domains that contain chemisorbed oxygen atoms, and that the coexisting Pt oxide domains supply oxygen atoms to the metallic domains. The measured rate of CO oxidation is thus proportional to the fractional area of the surface that is covered by the metallic phase. During the initial CO exposure, the oxide phase replenishes the metallic phase with oxygen atoms more quickly than CO oxidation removes the chemisorbed oxygen. As a result, the observed CO oxidation rate remains low because oxygen migration maintains a high coverage of chemisorbed oxygen atoms on the metallic phase which suppresses CO adsorption and hence reaction. The CO oxidation rate begins to increase once the metallic surface area reaches a critical value, and the rate of oxygen migration to the metallic phase drops below the rate of CO adsorption and oxidation on this phase. At this point, the oxygen coverage within the metallic domains decreases below saturation, thus allowing CO to adsorb and react more readily. The mechanism is autocatalytic because the production of metallic areas promotes further CO oxidation until nearly all of the coexisting Pt oxide is consumed. A similar mechanism appears to govern the CO oxidation kinetics observed in the Ba(NO3)2/O/Pt(111) system as well. Most interestingly, the evolution of NO and CO2 are observed at the same time during the isothermal CO oxidation experiment. The release of NO demonstrates that the nitrates can decompose at ∼450 K in the presence of gaseous CO, which is ∼40 K lower than the onset of nitrate decomposition observed during TPD. Furthermore, we observe negligible NO2 desorption during the CO oxidation experiment, whereas an appreciable amount of NO2 evolves during thermal decomposition of the Ba-nitrate particles. These differences may also point toward distinct mechanisms for the thermal vs CO-induced decomposition of the Ba-nitrate particles. As mentioned above, CO does not directly react with Ba-nitrates, but rather facilitates nitrate decomposition by removing oxygen from the Pt domains and thus creating metallic Pt sites that can catalyze the decomposition of Ba-nitrates presumably through a spillover phenomenon. Ptcatalyzed decomposition of Ba-nitrates has also been suggested by previous studies on barium nitrate-impregnated Pt4 and on 5906
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Figure 4. Series of (a) O2 and (b) NO TPD spectra obtained following the exposure of Ba(NO3)2/O/Pt(111) to CO at 450 K as a function of CO exposure (Pco(back) = 10.0 torr ; exposure times are indicated in the figure). The Ba(NO3)2/O/Pt(111) system was prepared by following the exposure of BaO/Pt(111) to elevated NO2 pressure (1.0 10-4 torr for 180 s) at 450 K.
high surface area model NSR catalysts.4,5 TPD spectra obtained following the isothermal CO exposure experiment showed only a CO2 desorption feature at ∼725 K, assigned to the decomposition of BaCO3 formed from the reaction between CO2 formed in the CO oxidation reaction and BaO. The release of NO during the decomposition of nitrates in the CO flux also indicates the incomplete reduction of NOx by CO under the experimental conditions applied, and more experiments are needed to quantitatively characterize the reduction of NOx by CO on this model system. However, CO has been shown to be a less efficient reductant compared with H2.25,26 Furthermore, Pt is not a very effective catalyst for the dissociation of NO (formed in the decomposition of nitrates) that could ultimately lead to the recombinative desorption of N2 from Nad. Bowker et al. have shown the promotional effect of CO on nitrate decomposition, but not on the reduction of NOx.4 On the other hand, H2 was shown to facilitate both nitrate decomposition and NO reduction with excellent selectivity toward N2 formation.4 They attributed the difference in reduction efficiency between these two reductants (i.e., CO and H2) to the binding and consequent poisoning of metallic Pt by CO. H2 has also been proved to be a better reductant than others, such as CO and hydrocarbons, at low temperatures ( OBaO/Pt at 450 K. A removal of OPt by CO leads to the reduction of oxidized Pt and, thus, to the formation of metallic Pt, which catalyzes the decomposition of Ba(NO3)2. The Pt-catalyzed Ba(NO3)2 decomposition occurs at 450 K, which is ∼40 K lower than the onset temperature of the Ba(NO3)2 decomposition in the presence of oxidized Pt. A fraction of CO2 formed from CO and OPt during CO exposure on Ba(NO3)2/O/Pt(111) reacts with BaO to form BaCO3. Although CO facilitates the decomposition of Ba(NO3)2, the complete reduction of NOx to N2 by CO was not observed under the experimental conditions applied. ’ AUTHOR INFORMATION Corresponding Author
*E-mail
[email protected].
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(4) James, D.; Fourre, E.; Ishii, M.; Bowker, M. Appl. Catal., B 2003, 45, 147. (5) Poulston, S.; Rajaram, R. R. Catal. Today 2003, 81, 603. (6) Coronado, J. M.; Anderson, J. A. J. Mol. Catal. A: Chem. 1999, 138, 83. (7) Mudiyanselage, K.; Mei, D. H.; Yi, C.-W.; Weaver, J. F.; Szanyi, J. J. Phys. Chem. C 2010, 114, 20195. (8) Desikusumastuti, A.; Happel, M.; Dumbuya, K.; Staudt, T.; Laurin, M.; Gottfried, J. M.; Steinruck, H.-P.; Libuda, J. J. Phys. Chem. C 2008, 112, 6477. (9) Desikusumastuti, A.; Laurin, M.; Happel, M.; Qin, Z.; Shaikhutdinov, S.; Libuda, J. Catal. Lett. 2008, 121, 311. (10) Desikusumastuti, A.; Staudt, T.; Gr€onbeck, H.; Libuda, J. J. Catal. 2008, 255, 127. (11) Desikusumastuti, A.; Staudt, T.; Happel, M.; Laurin, M.; Libuda, J. J. Catal. 2008, 260, 315. (12) Yi, C.-W.; Kwak, J. H.; Szanyi, J. J. Phys. Chem. C 2007, 111, 15299. (13) Yi, C.-W.; Szanyi, J. J. Phys. Chem. C 2009, 113, 2134. (14) Yi, C.-W.; Szanyi, J. J. Phys. Chem. C 2009, 113, 716. (15) Schmitz, P.; Baird, R. J. Phys. Chem. B 2002, 106, 4172. (16) Tsami, A.; Grillo, F.; Bowker, M.; Nix, R. M. Surf. Sci. 2006, 600, 3403. (17) Staudt, T.; Desikusumastuti, A.; Happel, M.; Vesselli, E.; Baraldi, A.; Gardonio, S.; Lizzit, S.; Rohr, F.; Libuda, J. J. Phys. Chem. C 2008, 112, 9835. (18) Bartram, M. E.; Windham, R. G.; Koel, B. E. Surf. Sci. 1987, 184, 57. (19) Bartram, M. E.; Windham, R. G.; Koel, B. E. Langmuir 1988, 4, 240. (20) Mudiyanselage, K.; Yi, C. W.; Szanyi, J. J. Phys. Chem. C 2009, 113, 5766. (21) Parker, D. H.; Bartram, M. E.; Koel, B. E. Surf. Sci. 1989, 217, 489. (22) Devarajan, S. P.; Hinojosa, J. A., Jr.; Weaver, J. F. Surf. Sci. 2008, 602, 3116. (23) Weaver, J. F.; Chen, J.-J.; Gerrard, A. L. Surf. Sci. 2005, 592, 83. (24) Gerrard, A. L.; Weaver, J. F. J. Chem. Phys. 2005, 123, 224703. (25) Szailer, T.; Kwak, J. H.; Kim, D. H.; Hanson, J. C.; Peden, C. H. F.; Szanyi, J. J. Catal. 2006, 239, 51. (26) Nova, I.; Lietti, L.; Forzatti, P.; Prinetto, F.; Ghiotti, G. Catal. Today 2010, 151, 330. (27) Abdulhamid, H.; Fridell, E.; Skoglundh, M. Top. Catal. 2004, 30-31, 161.
’ 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. J.F.W. gratefully acknowledges financial support provided by the Department of Energy, Office of Basic Energy Sciences, Catalysis Science Division through grant number DE-FG02-03ER15478. ’ REFERENCES (1) Epling, W. S.; Campbell, L. E.; Yezerets, A.; Currier, N. W.; Parks, J. E. Cat. Rev.-Sci. Eng. 2004, 46, 163. (2) Liu, Z.; Anderson, J. A. J. Catal. 2004, 224, 18. (3) Bowker, M. Chem. Soc. Rev. 2008, 37, 2204. 5909
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