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J. Phys. Chem. C 2009, 113, 716–723
BaO/Al2O3/NiAl(110) Model NOx Storage Materials: The Effect of BaO Film Thickness on the Amorphous-to-Crystalline Ba(NO3)2 Phase Transition Cheol-Woo Yi† and Ja´nos Szanyi* Institute for Interfacial Catalysis, Pacific Northwest National Laboratory, P.O. Box 999, MSIN: K8-80, Richland, Washington 99352 ReceiVed: October 03, 2008; ReVised Manuscript ReceiVed: NoVember 12, 2008
The reaction of NO2 with BaO (0.15-2 ML and >30 ML)/Al2O3(12 ML)/NiAl(110) model NOx storage materials was studied. A thick (∼12 ML), ordered Al2O3 film was prepared as the support oxide on a NiAl(110) substrate to minimize the effect of the intermixing between the two oxide phases (BaO and Al2O3) on the NOx chemistry of BaO. The growth of a thick alumina film, prepared by atomic oxygen deposition onto NiAl(110), follows a layer-by-layer growth mode, and the resulting film is much more stable when exposed to NO2 than the ultrathin alumina films studied before. The interaction of NO2 with the model NOx storage systems at low coverages of BaO shows behaviors fundamentally different from a thick BaO film, as nitrite species form at low exposures of NO2, followed by nitrate formation at high NO2 exposures. In contrast, on the thick BaO layer, nitrite-nitrate ion pairs form at 300 K under UHV conditions (PNO2 ∼ 1 × 10-9 Torr). However, at elevated NO2 pressures (g1 × 10-5 Torr), the thick BaO film is gradually converted into amorphous Ba(NO3)2 at 300 K. Raising the temperature of the samples with ΘBaO > 1 ML after NO2 exposure (in the absence of gas phase NO2) leads to the phase transformation of the amorphous Ba(NO3)2 layer into crystalline Ba(NO3)2 particles in the temperature range of 500-600 K. No phase transformation is observed in samples with ΘBaO < 1 ML. Introduction Three-way catalysts (TWC), which can convert NOx, COx, and hydrocarbons to N2, CO2, and H2O with high efficiency, have widely been employed to treat exhaust streams originating from internal combustion engines. However, the operation of TWCs is limited to stoichiometric air-to-fuel ratios, at which they can perform simultaneous oxidizing and reducing functions.1 In contrast, in lean-burn gasoline and diesel engine applications (i.e., under net oxidizing conditions), these catalysts are ineffective for NOx reduction, and therefore, other types of catalysts have to be developed to overcome this limitation. Today, one of the most promising technologies being considered to reduce NOx emission from automobile engines operating under lean conditions is NOx storage/reduction catalysis, in which the complex catalyst material contains alkali or alkaline (or both) earth oxides, for example, barium oxide, as the NOx storage component.1-11 The operation of these systems is based on the storage of NOx into the storage component of the catalyst during the lean engine operation phase and the release and subsequent reduction of NOx during the brief rich burn engine operation period.11 Despite the considerable research efforts, the details of these storage and release processes are not wellunderstood due, in part, to the difficulties of constructing model systems that accurately mimic the properties of real catalysts.1-8 Model NOx storage systems (BaO/Al2O3/NiAl(100),12 BaO/Al2O3/NiAl(110),9,10,13-15 BaO/Pt(111),16,17 and BaO/ Cu(111) 18) have been prepared and characterized, and their NOx storage properties have been studied in considerable detail. However, there have been two major problems that have prevented the unambiguous understanding of NOx chemistry * Corresponding author. E-mail:
[email protected]. † Present address: Department of Chemistry and Institute of Basic Science, Sungshin Women’s University, Seoul 136-742, Korea (ROK)
over these BaO-containing model storage materials: (i) on the alumina-based systems,9,10,12-15 the thickening of the alumina film upon its exposure to NO2; and (ii) on all substrates,9,10,12-18 the reaction of BaO with the underlying unstable metal oxides. We have recently reported that extended and point defect sites present in ultrathin alumina films were responsible for the increase in the alumina film thickness when these films were exposed to oxidants. In fact, our water adsorption/desorption study on an ultrathin Al2O3 film grown on a NiAl(110) substrate has revealed an ∼20% increase in the alumina film thickness (estimated from X-ray photoelectron spectroscopy (XPS) measurements) after 10 runs of temperature-programmed desorption (TPD) experiments.19 The significant increase in the alumina film thickness upon NO2 adsorption and desorption cycles can clearly be seen from the XPS data shown in Figure 1. Spectra of the Al 2p and Ni 3p regions for an as-prepared ultrathin Al2O3 film and of the alumina film after three cycles of NO2 adsorption/ desorption are displayed. After three cycles of NO2 adsorption/ desorption, the intensities of the XPS features of both metallic aluminum and nickel are reduced. Concomitantly, the intensity of the ionic aluminum feature increased, suggesting a significant thickening of the alumina film. The estimated thickness of the alumina film after three cycles of NO2 adsorption/desorption was 10.6 Å, a two-fold increase in film thickness with respect to the pristine ultrathin alumina film (∼5 Å). Recent studies have reported a characteristic defect structure for the ultrathin alumina films formed on NiAl(110). These studies reveal that these defects consist of network line defects, including domain boundaries and substrate steps, as well as a small number of point defects which are observed by scanning tunneling microscopy (STM) and electron paramagnetic resonance.20-22 These defect sites on the oxide film are the primary sources of water and NO2 decomposition and play a role in the increase in the oxide film thickness by the reaction of aluminum in the
10.1021/jp808766n CCC: $40.75 2009 American Chemical Society Published on Web 12/18/2008
BaO/Al2O3/NiAl(110) Model NOx Storage Materials
J. Phys. Chem. C, Vol. 113, No. 2, 2009 717 with BaO/Al2O3/NiAl(110) model NOx storage materials with a wide range of BaO coverages (0.15 - 30 ML) was investigated. The active NOx storage phase (i.e., BaO) was deposited onto a thick alumina film formed on the NiAl(110) substrate to minimize the effect of the alumina film on the NO2 chemistry discussed above. The focus of this study is to contrast the NO2 chemistry of submonolayer and bulklike BaO model NOx storage systems using surface spectroscopic techniques. Experimental Section
Figure 1. XPS spectra of the Al 2p and Ni 3p regions prior to and after NO2 adsorption and desorption experiments on an ultrathin (∼5 Å) Al2O3 film on NiAl(110).
subsurface with the atomic oxygen originating from the decomposition of water and NO2. The thickening of the alumina film during NO2 adsorption experiments is an issue that makes some of the results obtained in TPD and XPS studies in previous works questionable. To date, the number of studies on BaOcontaining model systems is very limited, and even in some of these studies, the reproducibility of the results for NO2 adsorption was poor. For example, the NO2 desorption profiles from a 5-6-monolayer (ML)-thick BaO film on Cu(111), studied by Tsami et al.,18 were shown to vary significantly with the number of adsorption/desorption cycles. They concluded that the NOx chemistry on this particular model system strongly depended on the history of the sample. These variations were attributed to the reaction between BaO and the underlying CuOx during NOx adsorption and desorption processes that resulted in the formation of barium cuprate phases. In our recent studies,10,13 a thick BaO film was prepared to minimize the effect of the underlying thin alumina film and to allow the understanding of the details of the reaction of NO2 with pure BaO. The results (XPS, infrared reflection absorption spectroscopy (IRAS) and TPD) obtained on this model system have clearly demonstrated that impurity-free BaO reacted with NO2 to initially form nitrite-nitrate ion pairs by a cooperative adsorption mechanism. These ion pairs formed readily at room temperature and even at 90 K. During thermal decomposition, the nitrite species disappeared first by the release of an NO molecule, followed by the decomposition of nitrate species in two steps: at lower temperature, as NO2; and at high temperature, as NO and O2. These results are in good agreement with those obtained from high-surface-area NOx storage materials in which high loadings of BaO were supported on γ-alumina.10 In a study of model BaO/Al2O3/NiAl(110) systems with BaO loadings similar to those of practical catalysts,9 fundamentally different NOx chemistry was observed in comparison to the bulklike BaO systems (thick BaO layer). At low BaO coverages (ΘBaO e 0.75 ML), due to the strong active metal oxide (BaO) support (Al2O3) interaction, a barium aluminate-like surface phase forms. Upon NO2 exposure, initially nitrite species form on this Ba-alumina phase, and then, as the NO2 exposure increases, Ba ions are pulled out from the surface (the Ba aluminate surface phase decomposes), resulting in the formation of barium nitrates.9 In the present study, the interaction of NO2
The experimental setup and the data acquisition procedures employed in the current study have been discussed in detail elsewhere.9,10,19 All the experiments were performed in an ultrahigh vacuum (UHV) surface analysis chamber equipped with XPS, Auger electron spectroscopy (AES), low-energy electron diffraction (LEED), low-energy ion scattering spectroscopy, and quadrupole mass spectroscopy and connected to a high-pressure cell with CaF2 windows for IRAS measurements and pressure-dependence studies. The base pressure in both chambers was less than 2 × 10-10 Torr. The NiAl(110) single crystal used in these experiments (Princeton Scientific Corp., diameter ) 10 mm and thickness ) 2 mm) was spot-welded onto a U-shaped Ta wire (0.030 in. diameter), and the sample temperature was measured by a C-type thermocouple spotwelded to the backside of the single crystal. The NiAl(110) crystal was cleaned by repeated cycles of Ar ion sputtering and annealing at 1200 K in UHV, and the cleanliness of the surface was verified with AES, XPS, and LEED. An ultrathin alumina film was grown epitaxially by oxidation of a NiAl(110) single crystal, as described previously,10,23-25 after a clean NiAl(110) substrate was prepared. The ordered, ultrathin oxide film was formed by 1200 L of O2 adsorption on the clean NiAl(110) surface at 540 K and successive annealing at 1070 K in UHV. The alumina film thickness was estimated to be around 2 monolayers. The thick alumina film was prepared by the deposition of atomic oxygen using a hot filament method onto an ultrathin Al2O3 film on NiAl(110) at 540 K and subsequent annealing at 1070 K. The film thickness, calculated by the attenuation of the substrate XPS feature26 and by the ratio of metallic to ionic Al 2p XPS features,9 was 29 ( 2 Å (∼12 ML Al2O3). The inelastic mean free paths of photoelectrons were obtained from reference 27. The desired amount of Ba was deposited onto the thick alumina film by reactive layer-assisted deposition (RLAD) using a resistively heated Ba doser (SAES Getters Inc.) onto a N2O4 ice as a reactive layer at 90 K. The thus formed BaNxOy was thermally decomposed to give the BaO film. Infrared spectra were collected at 4 cm-1 resolution using a grazing angle of approximately 85° to the surface normal. All the IR spectra collected were referenced to a background spectrum acquired from the clean sample prior to gas adsorption. Results and Discussion We have recently reported that the presence of extended and point defect sites19-22,28,29 in ultrathin alumina films (film thickness ∼ 5 Å, 2 ML Al2O3) prepared by the oxidation of NiAl(110) single crystal led to increases in the alumina film thicknesses19 when these films were exposed to H2O19 and NO2. Recent studies14,30 by Libuda and co-workers also confirmed the thickening of the ultrathin alumina film grown on NiAl(110) single crystal upon NO2 adsorption, and they observed the same phenomenon, even with relatively thick BaO (14 Å thickness) supported on Al2O3 (5 Å thickness)/NiAl(110) model systems. They also claimed that the growth of the alumina film on NiAl(110) stopped at high NO2 exposures.30 This latter result
718 J. Phys. Chem. C, Vol. 113, No. 2, 2009 indicates that the formation of atomic oxygen is reduced by the decrease in the number of defect sites where the dissociation of adsorbates can take place as the alumina film thickness increases, or the thickening of the alumina film is limited by the diffusion rate of atomic oxygen through the metal oxide overlayers as the alumina film becomes thicker. In addition, we have already pointed out the effect of the formation of mixed oxide phases formed by the reaction of BaO with underlying ultrathin alumina film support material10 on the NOx chemistry of the BaO film. These problems have caused difficulties in explaining some of the experimental NO2 uptake data as well as poor reproducibility. Therefore, to minimize the effects the ultrathin alumina films present, a thick, ordered alumina film (12 ML) was prepared as the oxide support for lower-coverage (0.15-2.0 ML) BaO model storage materials. A thick BaO film (>30 ML) was also prepared on the thin alumina film to obtain a pristine BaO phase that allowed the investigation of the reaction of NO2 with BaO only, without any influence from the underlying oxide support. First, we are going to discuss the properties of the thick (12 ML) Al2O3 film prepared on NiAl(110) and then the NO2 chemistry on BaO deposited onto the thick alumina film at coverages of 0.15, 0.75, and 2.0 ML. Finally, we will compare the results of NO2 adsorption on a thick (>30 ML) BaO layer to those obtained from the thin BaO films, under a wide range of NO2 pressures (1 × 10-9-1 × 10-4 Torr). 1. Preparation of a Thick, Ordered Alumina Film. Wellordered, ultrathin alumina films can be prepared on NiAl(110) by a typical procedure consisting of two steps: oxidation at 540 K and annealing at 1070-1100 K.10,23-25 The thickness of the resulting alumina film is ∼5 Å (∼2 ML). Yoshitake and Lay24,31 reported that the O 1s XPS intensity saturated at ∼600 L of O2 dose when the oxidation was performed in the temperature range of 570-770 K, and the intensity ratio of ionic to metallic aluminum 2p XPS features did not change upon annealing after additional oxygen exposures. These results suggest that the thickness of the alumina film did not change, even though further O2 was applied in the indicated temperature range. Hence, ultrathin alumina films on NiAl(110) are quite stable in molecular oxygen atmosphere, and thickening of these films was not observed after the formation of the ultrathin alumina film. However, recent studies9,10,14,19,30,32 have shown that these films were not chemically inert or stable in the presence of water19 or NO2 (as shown in Figure 1). To overcome the problems originating from the reactivity of the ultrathin alumina film discussed above, we prepared a thick alumina film on the NiAl(110) substrate by utilizing atomic oxygen generated by the dissociation of O2 on a hot tungsten filament. Methods for the preparation of ordered, thick alumina films on NiAl(110) substrates have not yet been reported, since in all the published reports, the oxidation step of the alloy surface was conducted in the presence of molecular oxygen. Thick (up to ∼8 ML), continuous, ordered alumina films have already been prepared on Re(0001) by Goodman and co-workers33 by evaporation of metallic Al under oxygen atmosphere. The thus prepared alumina films were characterized by XPS, STM, and LEED.33 Figure 2A shows a series of XPS spectra for the Al 2p and Ni 3p regions as a function of exposure of atomic oxygen of a clean NiAl(110) surface. The XPS spectrum (black line in Figure 2A) obtained after the reaction of molecular oxygen with the clean NiAl(110) substrate consists of a high intensity feature of the metallic aluminum and a small shoulder of the ionic aluminum, indicating the formation of the typical ultrathin alumina film. The XPS intensities of either the Al 2p or the O
Yi and Szanyi
Figure 2. The growth of an Al2O3 film as a function of atomic oxygen deposition time: (A) XPS spectra of the Al 2p and Ni 3p regions and (B) the calculated Al2O3 film thickness prepared by atomic and molecular oxygen deposition as a function of exposure time.
1s (not shown here) do not change, even after additional dosages of molecular oxygen. This indicates that the film thickness does not change during the additional molecular oxygen deposition, consistent with other groups’ results.24,31 However, the increases in both ionic aluminum and oxygen XPS signals and the concomitant decreases in both the metallic aluminum and nickel signals observed as a function of the deposition time of atomic oxygen indicates the increase in the thickness of the alumina film. Figure 2B displays the estimated alumina film thickness as a function of atomic oxygen deposition time. The alumina film thickness was calculated by two methods,9 using both the attenuation of the Ni 2p3/2 XPS peak and comparison of the relative intensities of the ionic and metallic Al 2p XPS peaks.34 To calculate the film thickness, 14.9 Å was used as the inelastic mean free path (IMFP) of Ni photoelectrons in the alumina medium, according to reference 27, and IMFP values in metallic aluminum and aluminum oxide were ∼ 26 and 28 Å,34 respectively. The thickness of the alumina film gradually increases as a function of exposure of atomic oxygen (Figure 2B). The results also suggest a layer-by-layer growth of the alumina film on NiAl(110) with atomic oxygen deposition (at least up to the thickness limit of ∼30 Å in this study). The alumina film with ∼12 ML thickness provided a hexagonal
BaO/Al2O3/NiAl(110) Model NOx Storage Materials
Figure 3. A series of IRAS spectra at saturation coverages of NO2 on various BaO/alumina model catalysts at 300 K. The coverage of BaO are 0, 0.15, 0.75, and 2.0 on 12 ML Al2O3/NiAl(110). The NO2 saturation was carried out at 1 × 10-4 Torr NO2 pressure.
LEED pattern, in good agreement with the one shown by Goodman and co-workers for the ∼8.5 ML Al2O3 on Re(0001).33 This thick Al2O3 film showed no increase in its thickness upon exposure to NO2, in stark contrast to the ultrathin alumina films discussed above. 2. NO2 Adsorption on BaO (e2 ML)/Al2O3 (12 ML)/ NiAl(110) systems at 300 K. Using the inert and stable alumina film as support, model NOx storage systems with low coverages (0.15-2 ML) of BaO were prepared by the RLAD method we discussed previously,9,10 and their reactivities with NO2 were investigated. Figure 3 summarizes the IRAS spectra at saturation coverages of NO2 on various Ba-containing alumina filmsupported model NOx storage systems at 300 K (NO2 was dosed through a precision leak valve, PNO2 ) 1 × 10-4 Torr for 30 s at 300 K). As previously reported,9 upon the exposure of the Al2O3 film to NO2, IRAS features of nitrate species in bridging and chelating bidentate configurations were observed in the 1550-1700 cm-1 range, consistent with IR absorption features reported for high surface area γ-alumina (bridging nitrate: 1235, 1254, 1591, 1617, and 1646 cm-1; chelating bidentate nitrate: 1300 and 1570 cm-1)8 (the low frequency component of the split asymmetric nitrate stretch can be seen at 1250-1300 cm-1, although its intensity is very low). However, the IRAS features characteristic of nitrite species were not observed at 1220-1250 cm-1, although nitrite species on this model system were observed by XPS.9 These observations indicate that the nitrite species assume a geometry in which the N-O bond is parallel to the surface; therefore, the IR features for these species could not be observed due to the surface selection rules of IRAS. Furthermore, at low NO2 exposures (not shown here), no vibrational features in the IRAS spectrum were observed, whereas XPS analysis clearly indicated the formation of nitrite species. Upon high exposure of NO2, however, the formation of nitrate species was observed by both IRAS and XPS. These results are consistent with previously reported ones for NO2 adsorption on high-surface-area γ-alumina, which showed the formation of both nitrite and nitrate species. Grassian and coworkers35 reported the formation of nitrite species on a γ-alumina surface at very low exposures of NO2 and the dominant formation of nitrate species at elevated NO2 pressures using UV-vis spectroscopy. As BaO was added onto the thick alumina film model support, the IRAS spectra, displayed in Figure 3, clearly showed that at
J. Phys. Chem. C, Vol. 113, No. 2, 2009 719 saturation coverage at 300 K (PNO2 ) 1 × 10-4 Torr), the intensities of the alumina-based nitrate features decreased (1550-1700 cm-1), whereas those of the Ba-related nitrate features (1300-1335 and 1400-1500 cm-1) significantly increased. Figure 3 also shows that with increasing BaO coverage (from 0.15 to 2 ML), the intensities of the alumina film-related nitrate features gradually decreased, whereas those of the BaO-related nitrate features concomitantly increased. In our studies of NO2 adsorption on high-surface-area BaO/γ-Al2O3 systems, the formation of two different types of nitrate species was indicated: specifically, surface and bulk nitrates.8 Surface nitrates were proposed to form on the BaO monolayer that strongly interacted with the alumina substrate or the nitrate species on the surface Ba aluminate layer. Their characteristic IR features appeared at ∼1300 and ∼1575 cm-1. In contrast, in the case of NO2 adsorption onto a pure, thick BaO film,10,13 no IRAS features characteristic of surface nitrates of the type discussed above were observed. However, bulk nitrate species formed on BaO particles on the alumina support or by the agglomeration of high-mobility barium nitrate species formed in the reaction of NO2 with the BaO layer. The IR features of these bulk-type nitrate species were observed at 1450 and 1310 cm-1 in high surface area catalysts. Practical NOx storage/reduction catalysts contain BaO as the active NOx storage phase in concentrations around 20 wt %. The surface coverage of BaO on γ-Al2O3 (surface area ) 200 m2/g) in the 20 wt % BaO/γ-Al2O3 system, assuming complete spreading of the BaO layer and using a distance between Ba and O of 2.77 Å,9 is 0.75 ML. The results of our NO2 adsorption experiments on systems with ΘBaO < 1 ML9 discussed briefly in the previous paragraph and those for a thick (>30 ML) BaO film10 revealed significant differences in the initial stages of NO2 uptake. On bulklike (thick) BaO films, nitrite-nitrate ion pairs formed initially by the cooperative adsorption mechanism at low NO2 exposures.10 Upon exposing the BaO/Al2O3/NiAl(110) samples with ΘBaO < 1 ML to NO2 at low NO2 pressures ( 1 × 10-6 Torr), we begin to see the formation of nitrates both with XPS (the appearance of the N 1s feature characteristic of nitrates at ∼407.5 eV), and IRAS. With further NO2 exposure at elevated NO2 pressures, the intensities of nitrate-related IR features increase, concomitant with the increase in the N 1s feature of the nitrate species in the XP spectra. It is also interesting to note that as the N 1s feature of the nitrate species gains intensity, that of the nitrite species loses intensity, and ultimately, the only N 1s feature observed is the one representing the nitrates. To bridge the gap between these two extreme BaO coverage regimes (i.e., submonolayer versus multilayer), a BaO film with
720 J. Phys. Chem. C, Vol. 113, No. 2, 2009
Yi and Szanyi
Figure 5. A series of IRAS spectra obtained after NO2 adsorption on a BaO (2 ML)/Al2O3(12 ML)/NiAl(110) model NOx storage material. The effect of NO2 pressure on the NOx species formed (bottom two spectra) and the amorphous-to-crystalline barium nitrate phase transition at elevated temperatures.
Figure 4. Series of RAIRS (A) and XPS (B) spectra collected during NO2 exposure of a 0.75 ML BaO/Al2O3/NiAl(110) sample at 300 K as functions of NO2 pressure.
a nominal coverage of 2 ML was deposited onto the thick Al2O3 support layer on the NiAl(110) substrate, and the NO2 adsorption properties of this sample was investigated. At low exposures of NO2, XPS analysis identified the presence of a small nitrite species only, but there was no IRAS feature due to the abovementioned surface IR selection rules. Similarly to the samples with ΘBaO < 1 ML, no nitrate species were detected under these conditions. With the increase in the NO2 pressure from 1 × 10-9 to 1 × 10-5 Torr, small, broad vibrational features (bottom spectrum in Figure 5) appear at 1320 and 1400-1500 cm-1 that can be assigned to ionic bulk nitrate species.8 Exposure of the sample to even higher pressures of NO2 (1 × 10-4 Torr) up to saturation coverage (no increase in the intensities of the IR bands) resulted in significant intensity increases in these IR features (red solid line in Figure 5) and in the shifts of their positions to 1335 and 1463 cm-1. These results suggest that the NO2 uptake mechanism in this thin BaO film (ΘBaO ∼ 2 ML) is the same as the one we have discussed for the samples with submonolayer BaO coverages. This behavior is attributed to the strong interaction between the thin BaO layer and the underlying Al2O3 film. The vibrational features at 1335 and 1463 cm-1 were previously assigned to ionic barium nitrate species,6,8 whereas the small shoulder appearing at 1570 cm-1 was assigned to surface nitrate species.9 At low pressures of NO2, this shoulder at 1570 cm-1 was not observed; however, it can clearly be seen at elevated pressures of NO2 at 300 K, as shown in Figure 5.
The appearances of both surface and bulk barium nitrate features in the IRAS spectra provide clear evidence for the formation of barium nitrate clusters, even at 300 K upon high NO2 exposures at elevated NO2 pressures (>1 × 10-5 Torr). Recently, we have reported that BaO strongly interacted with the alumina support and formed a barium aluminate-like surface phase at submonolayer coverages of BaO on alumina, and the reaction of NO2 with this surface phase resulted in the formation and agglomeration of barium nitrate clusters.9 The results reported here for the 2 ML BaO film combined with those obtained at submonolayer BaO covarages suggest that the barium nitrate phase that formed under elevated pressure NO2 exposures possesses sufficient mobility, even at 300 K, to form bulklike barium nitrate particles (clusters). Here, we need to emphasize the importance of NO2 pressure in this agglomeration process. When these samples were exposed to NO2 through a precision leak valve (i.e., the base pressure in the chamber stayed below 2 × 10-9 Torr), the formation of these barium nitrate clusters was extremely slow, hardly observable in the time frame of these experiments. This suggests that the rate of bulk barium nitrate formation is kinetically limited. With the increase in the NO2 pressure, as bulk barium nitrate forms, it agglomerates into larger clusters, opening the path for the formation of surface nitrates by the reaction of NO2 and BaO strongly interacting with the underlying alumina film, as evidenced by the results presented in Figure 5 for the 2 ML BaO film. 3. Amorphous-to-Crystalline Phase Transition in Ba(NO3)2 Particles. In the previous paragraphs, we have shown that at high NO2 exposures under elevated NO2 pressures, the BaO layers with coverages 0.15 ML e ΘBaO e 2 ML were converted to Ba(NO3)2 particles, even at 300 K sample temperature. The formation of these Ba(NO3)2 particles was very slow, even at high NO2 exposures (long dosing times) when the NO2 dosing was administered through a leak valve. After exposing a 2-ML-thick BaO layer to NO2 at a pressure of ∼1 × 10-4 Torr to saturation (the intensities of the IR features of nitrate species did not change anymore), IR absorption features for both surface and bulklike Ba(NO3)2 were observed (Figure 5, second spectrum from the bottom). The same results were obtained over a thick (>30 ML) BaO film, although the intensities of the IR features of the bulk nitrates were much
BaO/Al2O3/NiAl(110) Model NOx Storage Materials
Figure 6. The amorphous-to-crystalline phase transformation in the NO2-saturated, thick BaO/Al2O3/NiAl(110) sample (saturation at 300 K under 1 × 10-4 Torr of NO2).
higher in this sample (see Figure 6 spectrum obtained at 300 K (black spectrum)). The IRAS features of the nitrate species we report here (1332-1338 and ∼1465 cm-1) for BaO films (from 0.15 to 30 ML) deposited onto a thick alumina layer support are identical to those observed by Libuda and co-workers32 for BaO/Al2O3/NiAl(110) systems with BaO coverages of 0.125 and 2.5 MLE at 300 K sample temperature and saturation NO2 exposures. They attributed these features (1331 and 1465 cm-1) to the formation of surface nitrates only. When they exposed the 2.5 MLE BaO/Al2O3/NiAl(110) sample to NO2 at 500 K sample temperature, they observed the formation of two new, high-intensity IR features at 1386 and 1421 cm-1 that they assigned to the formation of ionic (bulk barium nitrates). They concluded that at 300 K sample temperature, the only nitrate species that can form are nitrates on the topmost layer of the BaO clusters when BaO particles (regardless of their size) are exposed to NO2. The formation of bulk barium nitrates (ionic nitrates) are kinetically controlled and can take place only at elevated sample temperatures. These results seem to contradict our findings in that we have observed the formation of bulklike nitrate particles in the entire range of BaO coverages studied here when the samples were exposed to NO2 at elevated pressures at 300 K. The particle sizes of the barium nitrates thus formed, however, must be dependent upon the nominal BaO coverage (small barium nitrates at 0.15 ML, and very large ones at >30 ML of BaO), as it has been shown in the STM images of Libuda et al.15 On the other hand, in our previous study on cyclic NO2 adsorption (NO2 saturation at 300 K using a pinhole doser followed by annealing to 575 K) on the 30 ML BaO film, IRAS features very similar to those reported by Libuda et al. were observed32 for a 2.5 MLE BaO film exposed to 2400 L of NO2 at 500 K. However, we observed these IRAS features of barium nitrates after the BaO films, exposed to NO2 at 300 K, were heated to 575 K in the absence of NO2 (i.e., no additional NOx uptake was possible during the annealing step). These results suggested that the observed changes in the IR spectra may not be related to the formation of a bulk-type (ionic) barium nitrate species, but rather, the transformation of the amorphous bulk barium nitrate phase formed at 300 K into a crystalline form at higher temperatures. To test this hypothesis, we conducted a series of annealing experiments on both the 2 ML and >30 ML BaO films after their saturation with NO2
J. Phys. Chem. C, Vol. 113, No. 2, 2009 721 under elevated NO2 pressures (1 × 10-4 Torr) at 300 K. The series of IRAS spectra obtained for the 2 and >30 ML BaO samples (after exposure to NO2) are displayed in Figures 5 and 6, respectively. For the NO2-saturated, 2-ML-thick BaO film, only very small changes were seen as this sample was heated from 300 to 450 K. As the sample was annealed to 500 K, however, an abrupt change in the IRAS spectrum was observed. Two new, sharp features appeared at 1390 and 1427 cm-1, concomitant with the large intensity decreases in both the 1335 and 1465 cm-1 features. The intensities of these new IRAS features reached their maxima as the sample was heated to 550 K. Heating to even higher temperatures resulted in the decrease in the intensities of these features at 600 K and in their subsequent disappearance at 650 K. At this temperature, the only IRAS features that were observed were those we have seen after the exposure of this sample to NO2 at 300 K. As the sample was heated to 700 K, these features also lost most of their intensities due to the almost complete decomposition of the nitrate species. Similar results were observed for the >30 ML BaO film, although the clear separation between the IRAS features obtained after annealing the NO2-saturated sample to 550 K was not observed here. The spectrum recorded on this sample after 550 K annealing is very similar to the ones obtained from the same sample in the cyclic NO2 adsorption/annealing experiments (mentioned above), as well as to that reported by Libuda et al.32 for NO2 adsorption on 2.5 MLE BaO at 500 K. The appearance of this broad, poorly resolved IRAS feature is the consequence of the presence of Ba(NO3)2 particles with wide size and shape distributions and the presence of surface, amorphous, and crystalline nitrate phases at the same time. (Note the appearance of the nitrite feature at ∼1250 cm-1 as the sample was annealed at 500 - 600 K. This is due to the adsorption of NO2 on the unreacted surface BaO sites that are created by the movement of the Ba(NO3)2 particles on top of the BaO film. This observation suggests the incomplete conversion of the thick BaO film into Ba(NO3)2, even though a high NO2 pressure was applied during the uptake step. Since all the IR spectra were collected at 300 K, NO2 desorbing from the walls of the IR cell after the high pressure NO2 exposures can react with the adsorbate-free BaO sites during the cooling period.) Phase transformation from amorphous to crystalline barium nitrate discussed here, has previously been reported during NO2 adsorption on a high-surface-area BaCl2/SiO2 system36 and on a BaO powder.1 IR spectra obtained after the exposure of BaCl2/ SiO2, calcined at 623 K, to NO2 at room temperature shows two broad features at 1368 and 1460 cm-1. However, when the sample was heated briefly at 563 K in vacuum, the sharpening of these broad IR bands and shifts in their vibrational frequencies to 1380 and 1418 cm-1 were observed.36 In another investigation, when a BaO powder was exposed to NO2 at 523 K, the formation of a broad IR feature in the 1400-1500 cm-1 range was observed at low exposure, whereas at high exposures of NO2, two IR features centered at 1380 and 1420 cm-1 were observed with high intensities.1 The aggregation of Ba(NO3)2 species formed upon NO2 exposure of both model and highsurface-area BaO/Al2O3 NOx storage materials at 300 K has also been verified in our previous publication.9 We found that Ba(NO3)2 species formed in the reaction of NO2 with BaO on/ in alumina or barium aluminates displayed high enough mobility to aggregate into small clusters. Although the formation and agglomeration of barium nitrate occur upon the exposure of the BaO/Al2O3 model catalyst to NO2 at 300 K, the IR spectrum of this NO2 saturated model catalyst does not exhibit any change up to 450 K annealing temperature, as shown in Figure 5.
722 J. Phys. Chem. C, Vol. 113, No. 2, 2009 Annealing to 500 K, however, results in dramatic changes in the IRAS spectrum, as discussed above. This proposed phase transformation is affected by two factors: the size of the barium nitrate clusters and the surface temperature. Model systems with submonolayer coverages of BaO did not show any changes in the IR spectra during annealing to 550 K, above which their intensities dropped as the temperature reached the onset of nitrate decomposition. Systems with low coverages of BaO (30 ML) BaO layer; that is, it displays this phase transformation while samples at ΘBaO < 1 ML do not. This is due to the high enough surface coverage of BaO that can form amorphous, highly mobile (even at room temperature) Ba(NO3)2 particles that can go through the phase transformation to form crystalline Ba(NO3)2 clusters. Conclusions In this study, the interaction of NO2 with BaO/Al2O3/ NiAl(110) model NOx storage systems was examined using surface-sensitive techniques to correlate the NOx uptake properties of BaO films of low BaO coverage with those of a thick BaO. A thick, ordered alumina film was grown on NiAl(110) by atomic oxygen deposition. Under these conditions, the alumina film growth follows a layer-by-layer growth mode. The film produced by this method is much more stable against further oxidation (film thickness increase) upon its exposure to NO2 in comparison with ultrathin alumina films prepared using molecular O2. The interaction of NO2 with model systems at low BaO coverages is significantly different from that of a thick BaO film. On submonolayer BaO-containing model systems, nitrite species form at low exposures of NO2, and then nitrates form at high NO2 exposures (at elevated NO2 pressures; i.e. >1 × 10-5 Torr). This is contrasted with a bulklike, thick (>30 ML) BaO film, which exhibits cooperative NO2 adsorption that results in the initial formation of nitrite-nitrate ion pairs at low NO2 exposures. The IRAS results on BaO/Al2O3/NiAl(110) model NOx storage systems with low BaO coverages (e2 ML) suggest that reaction of NO2 at elevated NO2 pressures (up to 1 × 10-4 Torr) of the barium aluminate-like surface phase results in the formation of barium nitrate clusters due to the high mobility of barium nitrate species formed even at 300 K. These barium nitrate species, however, are highly dispersed and amorphous.
Yi and Szanyi Increasing the temperature of these NO2-saturated systems leads to the amorphous-to-crystalline phase transformation in the temperature range of 500-600 K. The formation of the crystalline barium nitrate phase was manifested in the appearances of sharp, intense IRAS features at 1390 and 1427 cm-1. At submonolayer coverages of BaO, however, this phase transformation did not take place due to the limited amount of BaO and strong interaction of the small Ba(NO3)2 clusters with the alumina surface after saturation with NO2. Acknowledgment. We gratefully acknowledge the U.S. Department of Energy (DOE), Office of Basic Energy Sciences, and 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. This work was also supported by the Sungshin Women’s University Research Grant of 2008. References and Notes (1) Broqvist, P.; Gronbeck, H.; Fridell, E.; Panas, I. J. Phys. Chem. B 2004, 108, 3523. (2) Fanson, P. T.; Horton, M. R.; Delgass, W. N.; Lauterbach, J. Appl. Catal., B 2003, 46, 393. (3) Hess, C.; Lunsford, J. H. J. Phys. Chem. B 2002, 106, 6358. (4) Mahzoul, H.; Brilhac, J. F.; Gilot, P. Appl. Catal., B 1999, 20, 47. (5) Nova, I.; Castoldi, L.; Lietti, L.; Tronconi, E.; Forzatti, P.; Prinetto, F.; Ghiotti, G. J. Catal. 2004, 222, 377. (6) Prinetto, F.; Ghiotti, G.; Nova, I.; Lietti, L.; Tronconi, E.; Forzatti, P. J. Phys. Chem. B 2001, 105, 12732. (7) Su, Y.; Amiridis, M. D. Catal. Today 2004, 96, 31. (8) Szanyi, J.; Kwak, J. H.; Kim, D. H.; Burton, S. D.; Peden, C. H. F. J. Phys. Chem. B 2005, 109, 27. (9) Yi, C. W.; Kwak, J. H.; Peden, C. H. F.; Wang, C.; Szanyi, J. J. Phys. Chem. C 2007, 111, 14942. (10) Yi, C. W.; Kwak, J. H.; Szanyi, J. J. Phys. Chem. C 2007, 111, 15299. (11) Matsumoto, S. CATTECH 2000, 4, 102. (12) Ozensoy, E.; Peden, C. H. F.; Szanyi, J. J. Catal. 2006, 243, 149. (13) Yi, C. W.; Szanyi, J. J. Phys. Chem. C 2008, DOI: 10.1021/ jp806854y. (14) 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. (15) Desikusumastuti, A.; Staudt, T.; Gronbeck, H.; Libuda, J. J. Catal. 2008, 255, 127. (16) Bowker, M.; Cristofolini, M.; Hall, M.; Fourre, E.; Grillo, F.; McCormack, E.; Stone, P.; Ishii, M. Top. Catal. 2007, 42-43, 341. (17) Bowker, M.; Stone, P.; Smith, R.; Fourre, E.; Ishii, M.; de Leeuw, N. H. Surf. Sci. 2006, 600, 1973. (18) Tsami, A.; Grillo, F.; Bowker, M.; Nix, R. M. Surf. Sci. 2006, 600, 3403. (19) Yi, C. W.; Szanyi, J. J. Phys. Chem. C 2007, 111, 17597. (20) Adelt, M.; Nepijko, S.; Drachsel, W.; Freund, H. J. Chem. Phys. Lett. 1998, 291, 425. (21) Frank, M.; Baumer, M. Phys. Chem. Chem. Phys. 2000, 2, 3723. (22) Frank, M.; Wolter, K.; Magg, N.; Heemeier, M.; Kuhnemuth, R.; Baumer, M.; Freund, H. J. Surf. Sci. 2001, 492, 270. (23) Kresse, G.; Schmid, M.; Napetschnig, E.; Shishkin, M.; Kohler, L.; Varga, P. Science 2005, 308, 1440. (24) Lay, T. T.; Yoshitake, M.; Song, W. Appl. Surf. Sci. 2005, 239, 451. (25) Franchy, R. Surf. Sci. Rep. 2000, 38, 195. (26) Yi, C. W.; Luo, K.; Wei, T.; Goodman, D. W. J. Phys. Chem. B 2005, 109, 18535. (27) Tanuma, S.; Powell, C. J.; Penn, D. R. Surf. Interface Anal. 1988, 11, 577. (28) Schmid, M.; Shishkin, M.; Kresse, G.; Napetschnig, E.; Varga, P.; Kulawik, M.; Nilius, N.; Rust, H. P.; Freund, H. J. Phys. ReV. Lett. 2006, 97. (29) Stierle, A.; Renner, F.; Streitel, R.; Dosch, H.; Drube, W.; Cowie, B. C. Science 2004, 303, 1652.
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