γ-Al2O3 NOX Storage

Jun 26, 2008 - Abstract. There has been recent debate regarding the role or influence of BaCO3 species on the ... Chemical Reviews 2009 109 (9), 4054-...
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10952

J. Phys. Chem. C 2008, 112, 10952–10959

Carbonate Formation and Stability on a Pt/BaO/γ-Al2O3 NOX Storage/Reduction Catalyst William S. Epling,* Charles H. F. Peden, and Ja´nos Szanyi Department of Chemical Engineering, UniVersity of Waterloo, Waterloo, Ontario, N2L 3G1 Canada, and Institute for Interfacial Catalysis, Pacific Northwest National Laboratory, P.O. Box 999, Richland, Washington 99352 ReceiVed: December 31, 2007; ReVised Manuscript ReceiVed: April 28, 2008

There has been recent debate regarding the role or influence of BaCO3 species on the performance or operation of Pt/BaO/Al2O3 model NOX storage/reduction (NSR) catalysts. This influence is primarily regarded as negative, but the extent of its impact is not clear. For this reason, the formation and stability of barium carbonate species on a Pt/BaO/Al2O3 model NSR catalyst were characterized using Fourier transform infrared (FTIR) spectroscopy. The catalyst sample was exposed to CO2, CO, and CO + O2 at various temperatures, from 300 to >500 K. Bidentate carbonate species readily form under all conditions, while at higher temperatures, unidentate species were also observed and likely formed from bidentate species as a result of a change in their coordination to the oxide surface. Reaction of COX species with residual hydroxide species on the catalyst led to the formation of bicarbonates, and when the sample was exposed to CO at low temperature, formate species were also formed. These formate species decomposed at elevated temperatures and contributed to the formation of carbonates. H2O exposure resulted in the agglomeration of various COX-containing phases to larger particles. Introduction NOX storage/reduction (NSR) catalysts have become a potential solution to meet the upcoming U.S. EPA NOX emission regulations for on-road diesel engines. These catalysts operate on a cycle, between lean (in the presence of excess oxygen) and reductant-rich conditions. During lean operation, the engine is run in its normal operating mode, and NOX is stored as nitrites and nitrates on the catalyst. An intermittent rich phase, in excess reductant relative to oxygen, reduces the stored NOX species to N2 and regenerates the catalyst surface, completing the cycle.1–4 Numerous studies have evaluated the trapping and regeneration reactions; however, typically the test conditions vary considerably. The inclusion of CO2 and/or H2O in the reaction stream is one such methodological inconsistency. Effects of these two exhaust gas components during lean/rich cycling have been studied5–16 and there is a general consensus that individually they have a negative impact, while the addition of H2O to a CO2-containing test mixture has been shown to increase or decrease overall performance depending on the test conditions.8,10 Suggested changes induced by the presence of H2O alone include: • blocking NOX adsorption onto the alumina support material via hydroxyl formation;9,10,12 • extensive nitrate morphology changes for Ba/Al2O3 via H2Oinduced, bulk-like Ba nitrate particle formation from dispersed surface nitrate species,11,14 which thereby reduces the Pt/Ba interaction which has been proposed as important for high performance; and • the formation of Ba(OH)2 species on a Pt/BaO/Al2O3 catalyst, which are less reactive than BaO.8,17 BaCO3 species have also been observed on model Pt/BaO/ Al2O3 NSR catalyst formulations. Furthermore, several proposed trapping mechanisms involve different types of carbonate species * Corresponding author. E-mail: [email protected].

having different reactivities.18,19 These differences have been attributed to: • slower reactions or diffusion limitations in a bulk-like BaCO3 particle versus BaCO3 surface species8,10,16 and • differing BaCO3 thermal stabilities that are related to its proximity to the Pt, Al2O3, or even BaO.20–22 Overall, the negative impact of the CO2 has been specifically attributed to the formation of a more thermally and chemically stable type of BaCO3, often described as bulk-like, that will not readily form nitrates during the trapping phase when CO2 is also present. The effects of CO2 and H2O change in significance at different operating temperatures and, as already noted above, the presence of H2O can affect the impact of CO210 and vice versa. Understanding their separate and combined effects on NOX adsorption and nitrate formation is critical for modeling NSR catalysts. In this study, the stability of carbonate species as a function of temperature in the presence and absence of H2O was studied using infrared (IR) spectroscopy and temperature programmed desorption (TPD). The carbonate species were formed using CO2, CO and O2, or CO, with the latter representing possible COX surface species forming pathways during the regeneration phase of the cycle. These experiments determine the stability of the various surface species formed as a function of temperature and gas composition for a model Pt/ BaO/Al2O3 catalyst. While the FTIR spectra obtained in this study are complex, their interpretation is, for the most part, relatively straightforward using prior literature assignments. We do, however, find that the assignment for one small feature observed in these spectra (at 1741-1747 cm-1 in the figures below) cannot be usefully made based on prior literature. Experimental Section The catalysts were in powder form and were prepared by the incipient wetness method, using aqueous Ba(NO3)2 (Aldrich) and dinitrodiammine platinum (Aldrich) solutions to impregnate

10.1021/jp712180q CCC: $40.75  2008 American Chemical Society Published on Web 06/26/2008

Pt/BaO/γ-Al2O3 NOX Storage/Reduction Catalyst

J. Phys. Chem. C, Vol. 112, No. 29, 2008 10953 TABLE 1: Assignments of IR Peak Positionsa peak position (cm-1) 1232-1238 1340-1360 1357-1398 1400-1403 1448-1454 1585-1587 1530-1598 1650-1656 1741-1747

Figure 1. IR spectra obtained from Pt/BaO/Al2O3 powder after exposure to (a) 1 Torr of CO2 for 1 min at 300 K, (b) 1 Torr of CO2 for 5 min at 300 K, (c) 1 Torr of CO2 at 400 K, (d) 1 Torr of CO2 at 500 K, (e) 1 Torr of CO2 at 600 K, (f) 1 Torr of CO2 at 673 K, and (g) 1 Torr of CO2 at 300 K after the heat ramp.

the γ-alumina support (Condea, 200 m2/g). Due to the limited solubility of Ba(NO3)2, multiple impregnations of the alumina support with the Ba(NO3)2 solution were carried out first, followed by a single impregnation step with the Pt-containing solution. The final loading was 2 wt % Pt and 20 wt % BaO. After the impregnations, the catalyst was dried at 395 K and then calcined at 773 K in a flowing 5% O2/He gas mixture for 2 h. This procedure ensured the decomposition of essentially all of the precursor Ba(NO3)2 phase into BaO. The Pt dispersion (as determined by H2 chemisorption and TEM) was ∼45-50%. Fourier transform infrared (FTIR) spectroscopic measurements were carried out in transmission mode, using a Nicolet Magna-IR 750 spectrometer operating at 4 cm-1 resolution. Catalyst samples were mounted into an IR cell consisting of a 2.75 in six-way stainless steel cube equipped with CaF2 windows. This cell was connected to a gas handling/pumping station and through both leak and gate valves to a UTI 100C mass spectrometer. The powder catalyst sample was pressed onto a fine tungsten mesh, which was mounted onto a copper sample holder assembly attached to ceramic feedthroughs of a 2.75 in flange. The sample temperature was monitored through a chromel/alumel thermocouple spot-welded to the top center of the tungsten mesh. After the Pt/BaO/Al2O3 powder sample was inserted into the vacuum chamber, it was cycled four times between approximately 10 Torr of NO2 and 20 Torr of H2 at 573 and 773 K, respectively. Before each experiment, the sample was exposed to NO2 at 673 K, reduced in H2 at 673 K, and annealed in vacuum at 773 K. Before acquisition of each spectral series, a background spectrum of the clean, adsorbate-free sample was obtained. Results and Discussion CO2 Exposure. The first experiment consisted of introducing 1 Torr of CO2 into the chamber at 300 K and ramping the sample temperature to 673 K in the presence of CO2, with IR spectra obtained every 50 K. Selected spectra obtained as a function of temperature are shown in Figure 1. As shown, multiple spectral features are apparent between 1200 and 1700 wavenumbers. After 1 min of CO2 exposure at room temperature (Figure 1a), peaks at 1232, 1349, 1448, 1598, 1650, 2030, and 2347 cm-1 were evident. A summary of the peak assignments used

2000-2067 2169-2170 2347

assignment bicarbonate (with 1448-1454 and 1650-1656 cm-1) bidentate carbonate (with 1530-1598 cm-1) formate (with 1586-1587 cm-1) unidentate carbonate (with 1741-1747 cm-1) bicarbonate (with 1232-1238 and 1650-1656 cm-1) bulk carbonate formate (with 1357-1398 cm-1) bidentate carbonate (with 1340-1360 cm-1) bicarbonate (with 1232-1238 and 1448-1454 cm-1) physisorbed H2O unidentate carbonate (with 1400-1403 cm-1) Pt-bound CO Pt-bound CO isocyanate adsorbed CO2

a

These assignments are primarily based on previous literature results with references listed in the text.

in this manuscript are listed in Table 1. As mentioned above, these assignments are primarily based on literature values, with references cited in the text. The features at 1232, 1448, and 1650 cm-1 are assignable to surface Al- and/or Ba-bicarbonate species.12,23–30 The presence of these features during this experiment indicates that surface hydroxyl species were present, and not removed by the cleaning process described above. This is not unexpected since H2 was used as a reductant in a pretreatment step just prior to a final vacuum anneal step to 773 K, and Al-hydroxide species can withstand decomposition at temperatures above 1000 K.31 Baltrusaitis et al. have recently published details on the interactions between H2O and CO2 on oxide surfaces, which include bicarbonate formation.32 The peak at 2347 cm-1 is due to adsorbed CO223,33 and the small feature at 2030 cm-1, which can be attributed to CO bound to Pt,27,34,35 must originate from CO2 dissociation on well-dispersed Pt particles. The peaks at 1349 and 1598 cm-1 are assigned to bidentate carbonate species,12,23,26,28,30 in this case likely associated with the Ba component of the catalyst. Further exposure at room temperature, for an additional 5 min (Figure 1b), caused no observable changes in the spectrum collected in comparison to that obtained after 1 min. Heating the sample from room temperature to 400 K in 1 Torr of CO2 (Figure 1c) resulted in the loss of the adsorbed CO2 feature and a decrease in the relative intensity of the 1650 and 1232 cm-1 features relative to the 1598 cm-1 peak. Since the 1448 cm-1 feature did not decrease in relative intensity, this indicates that more than one type of species could be represented by this feature, assuming that the three assignments for bicarbonate species above are indeed correct. Note also that further heating (Figure 1d,e) results in the disappearance of the 1232 and 1650 cm-1 features, but still some intensity remains near 1448 cm-1. In a recent IR study, a 1450 cm-1 band was associated with a more crystalline form of Ba-carbonate, i.e., a bulk-like feature.36 Thus, the residual signal observed at 500 K and above, after the complete loss of the 1232 and 1650 cm-1 features, indicates that this more bulk-like species also formed upon CO2 exposures, perhaps requiring elevated temperatures. A similar observation has been reported for the formation of bulk Ba(NO3)2 at elevated temperatures in the presence of gasphase NO2.37 A new feature also became apparent at approximately 1747 cm-1, a peak previously attributed to an illdefined bridged “organic-like” carbon species on alumina.26,38 This latter feature seemed to appear and grow as the bicarbonate

10954 J. Phys. Chem. C, Vol. 112, No. 29, 2008 peaks disappeared, and may be associated with decomposition of bicarbonate or other surface COX species into this more stable species as will be discussed below. The peak at 1448 cm-1 dropped in relative intensity after heating to 500 K (Figure 1d); however, it was still present up to the highest temperature, but as more of a shoulder to the lower wavelength peak. The intensity of the Pt-CO band remained relatively the same, or even increased slightly, as the temperature was raised. This is similar to previous findings in a study of CO2 interactions with Pt/K systems where a CO-surface bond was observed with HREELS and did not diminish significantly until temperatures above 650 K were reached.39 Increasing the temperature in the presence of CO2 also caused changes in the IR features at 1349 and 1598 cm-1 that are associated with Ba-carbonate species. Of particular note is that the 1598 cm-1 IR peak shifted to lower wavenumber values with increasing temperature. This was due, at least in part, to the increased background toward low wavenumbers in the spectra with increasing temperature. The 1349 cm-1 peak resolved into two features as the temperature was increased. This second, new peak at 1402 cm-1 has previously been assigned to the formation of unidentate carbonates.9,23,28,30,38 These species have been reported as both more stable23 and less stable30 than their bidentate counterpart. The temperature dependence of the 1402 cm-1 peak observed in this study seems to indicate that the unidentate species is more stable as its amount increased relative to the bidentate species as the temperature increased. Another possibility is that the presence of CO2 during the heating process may drive further carbonate formation, and therefore the unidentate species form after the bidentate species saturate their sites under these specific conditions. Finally, the bidentate species may partially decompose to the unidentate form resulting in the increased relative intensity. This was also true of the bicarbonate decomposition to a more stable species mentioned above in association with the development of the 1747 cm-1 peak. These data therefore indicate that the growth of the 1747 cm-1 feature was associated with the development of the 1402 cm-1 unidentate carbonate feature. These trends will be further discussed below in conjunction with the analysis of thermal stability in the absence of gas-phase COX species. The sample was ultimately heated to 673 K in the CO2 background (Figure 1f), and then cooled back to room temperature in this same CO2 background (Figure 1g). As shown in Figure 1g, the spectrum obtained after cooling back to room temperature closely resembles that obtained at room temperature before the heat treatment. The bicarbonate species reappeared, indicating that hydroxyls still reside on the sample surface. The bidentate carbonate features of Figure 1g are larger in absolute peak magnitude when compared to those of Figure 1a, indicating that the concentration of carbonates on the surface increased during the thermal treatment in CO2. After the sample was cooled to room temperature, the chamber was evacuated and the spectrum shown in Figure 2a was obtained. The only difference in the spectrum collected before removing the CO2 (Figure 1g) and that after the removal is the loss of the 2347 cm-1 feature assigned to weakly adsorbed CO2 on Pt. Increasing the sample temperature in the vacuum environment resulted in the loss of the bicarbonate species by 400 to 500 K (i.e., loss of 1656 and 1450 cm-1 peaks), which is similar to their disappearance even in the presence of CO2. Furthermore, just as in the presence of CO2, the band at 1353 cm-1 evolved into two features as the temperature was increased, coincident with the formation of the 1750 cm-1 feature, which

Epling et al.

Figure 2. Selected IR spectra obtained during a TPD experiment (in vacuum) from a Pt/BaO/Al2O3 powder after exposure to 1 Torr of CO2 at 300 K through 673 K. Spectrum (a) was obtained at 300 K, (b) at 400 K, (c) at 500 K, (d) at 600 K, (e) at 700 K, (f) at 773 K, and (g) at 300 K after the completion of the TPD run.

indicates that the unidentate carbonate formation was not due to new site availability at higher temperature, since no CO2 reactant was available in the experiment associated with Figure 2. This strongly suggests that it formed via bicarbonate and/or bidentate carbonate conversion. Overall, the same trends were observed as those with CO2 present in the background gas, which suggests that the relatiVe stabilities of these surface species are not a function of CO2 availability. However, there was an obvious loss in carbonate band intensity as the temperature was increased, demonstrating that the overall stability of the various species was dependent on the presence or absence of CO2 in the gas phase. This is more evident in comparing the spectrum obtained after cooling back to room temperature (Figure 2g) to that obtained before the heat treatment (Figure 2a). In comparing these room temperature data, it is apparent that the primary differences are a loss in carbonate species concentration and an irreversible removal of bicarbonate species when no CO2 was present during the annealing experiments. Mass spectrometer data (not shown) collected simultaneously with the FTIR spectra shown in Figure 2 demonstrated CO2 evolution maxima at 375 and 723 K, but the CO2 signal did not return to its baseline value throughout that temperature range, indicating a slow, but continuing release as the temperature was increased. Comparing this to the IR data, the first CO2 evolution maximum was likely due to the bicarbonate loss while the remainder to carbonate decomposition. After 5 min at 773 K, there were still carbonates on the sample surface, demonstrating their stability to this temperature even in vacuum. Also, note that the relative intensity of the bulklike carbonate feature (∼1450 cm-1) was higher after the hightemperature treatment than that after room-temperature adsorption. CO + O2 Exposure. A second set of experiments was performed where the sample was exposed to CO and O2 and the temperature ramped in this mixture. This simulates a possible reaction pathway to explain, for example, barium carbonate formation during the regeneration period of a NSR catalyst if residual O2 is present in the gas phase or the catalyst contains reactive oxygen-storage components. The FTIR spectra collected after exposure of the catalyst to CO (a) and then to CO + O2 (b) at room temperature are shown in Figure 3. The most prominent feature after exposing the sample to only CO is a Pt-CO band at approximately 2067 cm-1. Small features assigned to the presence of bidentate carbonate species were

Pt/BaO/γ-Al2O3 NOX Storage/Reduction Catalyst

J. Phys. Chem. C, Vol. 112, No. 29, 2008 10955

Figure 3. IR spectra obtained from a Pt/BaO/Al2O3 powder after exposure to (a) 1 Torr of CO at 300 K, (b) 1 Torr CO and 1 Torr O2 at 300 K, (c) the CO + O2 mixture at 330 K, (d) the CO + O2 mixture at 390 K, (e) the CO + O2 mixture at 450 K, (f) the CO + O2 mixture at 510 K, and (g) after (f) and cooled to 300 K in vacuum.

Figure 4. IR spectra obtained from a Pt/BaO/Al2O3 powder after exposure to 1 Torr of CO at (a) 300 K, (b) 330 K, (c) 360 K, (d) 420 K, (e) 480 K, (f) 523 K, (g) 523 K and cooled back to 300 K in CO, and (h) after evacuation at 300 K.

also evident with CO exposure, and increased in intensity when O2 was also available. The 1741 cm-1 feature was also evident at this room-temperature exposure, unlike with the CO2 exposure where it was only observed at higher temperatures. Although this again coincides with the presence of the 1402 cm-1 feature, but less evident in spectrum a or b of Figure 3 due to the size of the Pt-CO band, this feature appears broader than that observed in Figure 1. As shown in this figure and those associated with CO below, this broader feature only occurs with CO dosing, not CO2. Furthermore, it follows the same trends as the feature at ∼2050 cm-1, indicating that it most probably is associated with CO bound to Pt, possibly influenced by the presence of Ba. Although only at room temperature, it is clear that the addition of gas-phase O2 resulted in an increase in the carbonate species, as could be expected solely on the basis of reactant availability. Their formation prior to O2 addition must include reaction with surface-bound oxygen species. A distinct difference between the spectra obtained with CO2 and CO as the gas-phase species is the ratio of the carbonate feature assignments in the ranges 1330-1360 and 1570-1600 cm-1. Iordan et al. have observed a similar difference in relative intensities of these two peaks in analogous experiments with K/Al2O3 samples and assigned a feature at 1591 cm-1 to a chelating formate species.28 Such an assignment cannot be ruled out in this study due to the presence of small amounts of hydroxyl groups evident throughout the experiments, and the presence of formates would thus account for the relative intensities of the peaks observed here. Furthermore, features at 2821 and 2721 cm-1, identified in studies of Ba formate crystal,40 also evolve starting at 360 K. As shown in spectra c-f of Figure 3, increasing the sample temperature from 330 to 510 K, while keeping CO and O2 present in the gas phase, resulted in increased carbonate species on the surface. The heating also caused a decrease in the higher wavenumber peak relative to the 1346 cm-1 feature, indicating the formates were decomposing in this environment as the temperature was raised (this is clearly evidenced by the disappearance, in spectrum e, of the “sharp” formate band, sitting on top of the broad carbonate feature as in spectra c and d). Otherwise, increasing the test temperature induced further carbonate formation. The peak associated with bulk-like carbonate species also grew with the increase in temperature to approximately 450 K. This is evident in evaluating the spectrum obtained after cooling the sample back to room temperature

(spectrum g of Figure 3) where the intensity of the 1454 cm-1 feature is quite high. This peak is assigned to bulk carbonate species, with possible contributions of bicarbonates ruled out due to the absence of ∼1650 and 1238 cm-1 peaks through Figure 3f. This experiment was ended at 510 K, as it appeared that the same trends as observed with just CO2 in the stream were being followed after this point. This is expected due to the facile oxidation of CO to CO2 on the metallic Pt particles above ∼500 K. Recooling to room temperature in CO + O2 does result in bicarbonate formation (reappearance of ∼1656, 1454, and 1238 cm-1 peaks). A subsequent TPD experiment in vacuum to 773 K (data not shown) gave identical results as those shown in Figure 2. CO Exposure. A similar set of experiments was run with just CO in the gas phase to determine what species might be formed during regeneration with no O2 present, and the data are shown in Figure 4. Again, after room-temperature CO exposure, the prominent IR feature was at 2060-2070 cm-1, due to the formation of Pt-CO. Very small features near 1330, 1590, and 1650 cm-1 were also observed. These correspond to the formation of formate, carbonate, and bicarbonate species, with formate being the primary species. Another feature, at 2170 cm-1, becomes evident with exposure to CO at temperatures above 450 K, for example, in spectrum e of Figure 4, which was taken at 480 K. This IR peak has previously been assigned to an isocyanate species on either Ba or Pt.41,42 The presence of this feature unfortunately indicates that not all of the NOX species were eliminated by the reduction with H2 during the sample cleaning process (as a reminder, after each experiment the sample was exposed to 10 Torr of NO2, then 20 Torr of H2 at 673 K, and then annealed in vacuum at 773 K). We do not believe these residual NOX species affect any of the data interpretation in this study, as it appears to be a small amount after the cleaning procedure, on the order of 1-5% at most. It is however possible that NCO species may contribute to the formation of carbonates upon decomposition at elevated temperatures. Indeed, after the sample was exposed to CO up to 523 K with a subsequent TPD experiment, to be discussed in the following paragraph, no nitrate/isocyanate species were evident, but a small amount of residual carbonates were. This indicates that the CO treatment followed by the annealing cycle was sufficient to reduce the residual nitrate species signal to below detection. Analysis of the data obtained prior to each experiment that showed residual

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Epling et al.

Figure 5. IR spectra obtained during a TPD experiment (in vacuum) from a Pt/BaO/Al2O3 powder after exposure to 1 Torr of CO at room temperature through 523 K. Spectrum (a) was obtained at 300 K, (b) at 390 K, (c) at 450 K, (d) at 510 K, (e) at 600 K, (f) at 660 K, (g) at 773 K, and (h) at 300 K after completion of the TPD run to 773 K.

Figure 6. IR spectra obtained from a Pt/BaO/Al2O3 powder after exposure to 1 Torr of CO during a temperature ramp from 300 to 523 K, followed by cooling to 300 K, and maintaining 300 K for 10 more minutes in CO. The other spectra were obtained after evacuating the CO and then exposing the sample to 1 Torr of H2O at room temperature for 1 min, 9 min, and 21 min.

nitrate species also demonstrated that the addition of CO caused the nitrate features to ultimately disappear. Overall, however, these data support the building evidence of isocyanate intermediates during the regeneration phase of the NSR cycle.41,43 The most important conclusion that we can draw from this series of IR spectra is that, in the presence of CO only at elevated temperatures, the primary surface species formed are Baformates (1359, 1369, and 1585 cm-1). The formation of carbonates takes place to a much lower extent (in comparison with the CO2 only, and the CO + O2 cases), due to a lack of significant amounts of CO2 required for that reaction to occur. The source of CO2 could reflect impurity levels or perhaps was due to the reaction between the small amount of residual nitrates left behind from the reduction step and the large excess of CO present. After CO exposure of the sample at the highest temperature, the carbonate that formed was mostly of bulk type (1454 cm-1). This series of IR spectra also suggests that Baformate species formed in the course of this heating experiment are very stable up to the maximum temperature of this study (523 K), and its decomposition did not contribute significantly to the formation of carbonates. Figure 5 displays FTIR spectra obtained during temperature programmed desorption in vacuum following the Figure 4 CO dosing experiments. As the temperature was increased from 300 to 510 K, the most prominent change in the IR spectra was the decrease in the intensity of the Pt-CO feature, while the IR bands associated with the Ba-formates and carbonates remained essentially unchanged. The most important change in the intensities of these latter IR features occurred as the temperature was raised from 510 to 600 K. The comparison of spectra d and e of Figure 5 clearly shows the dramatic drop in the intensity of the Ba-formate-related IR bands (1585, 1371, 1357 cm-1) and, concomitantly, the large increase in the intensities of IR band associated with bidentate carbonates (1530-1590, and 1360 cm-1). The IR band centered at 1360 cm-1 in spectrum d represents primarily formate species, while that in spectrum e mostly bidentate carbonate as evidenced by changes in the peak shape of this region in the two spectra, and also by the large drop in the intensity of the other formate-related feature at 1585 cm-1. Heating the sample to 660 K (Figure 5f) results in the complete disappearance of IR features associated with Baformates. At this temperature, only bidentate (∼1542 and ∼1360 cm-1), unidentate (1403 cm-1 and the sharp 1745 cm-1 feature),

and “bulk” (∼1452 cm-1) carbonates are present, along with some residual adsorbed NCO and CO species. Both the isocyanate (2169 cm-1) and adsorbed CO (∼2000 cm-1) peaks decrease to very low levels only at the highest temperature (773 K). By comparison of spectrum g of Figure 5 to spectrum f of Figure 2, it is apparent that at high temperature (above 660 K) the same species (surface and bulk carbonates) are present regardless of the initial exposure of the catalyst sample to CO or CO2. What we wish to emphasize here is that although we end up with the same carbonate species in both experiments (following the TPD after CO and CO2 exposure), the way the carbonates were formed is fundamentally different in the two experiments. Upon CO2 exposure, carbonate formation is straightforward, while in the experiments with CO, carbonates form primarily by the decomposition of Ba-formates. At high temperature, it does not matter whether we expose the catalyst to CO or CO2; we always end up with carbonates. However, at lower temperatures, in the absence of oxygen, CO will form almost exclusively formates; therefore, we may need to consider the effect of these formates on the NSR performance of the catalyst. H2O + CO Exposure. The effect of H2O on the COX species formed and their stability was also examined. In the first experiment, the sample was again exposed to 1 Torr CO, and the temperature of the sample ramped to 523 K in the CO background, then cooled to room temperature where the CO was removed. The spectrum plotted in Figure 6a was then obtained. This spectrum is quite similar to that shown in Figure 4h acquired after the same experiment (the spectral range of analysis shown here was decreased for this figure to highlight changes in the 1300-1700 cm-1 region). 1 Torr of H2O was then added to the vacuum chamber, and after one minute of exposure another spectrum was obtained (dashed line spectrum of Figure 6). The notable changes were (1) the intensity of the peak at 1587 cm-1 increased and (2) the intensity of the shoulders toward the lower wavenumbers of the ∼1360 and 1587 cm-1 features decreased significantly. With further exposure to H2O, the peaks previously assigned to the formation of bicarbonate species, at 1450 and 1654 cm-1, were observed along with a further increase in the peak at 1587 cm-1. However, the intensity of the 1652 cm-1 feature increased in greater proportion to that of the 1450 cm-1 feature, while

Pt/BaO/γ-Al2O3 NOX Storage/Reduction Catalyst

Figure 7. IR spectra obtained from a Pt/BaO/Al2O3 powder after (a) exposure to 1 Torr of CO at 300 K for 10 min, followed by evacuation and then exposing the sample to 1 Torr of H2O at 300 K for 21 min; (b) heating in vacuum to 410 K; (c) heating to 410 K in vacuum and then exposure to 1 Torr of H2O for 5 min at 410 K; (d) heating in vacuum to 470 K; (e) heating to 470 K in vacuum and then exposure to 1 Torr of H2O for 9 min at 470 K.

the 1236 cm-1 peak did not appear with the addition of water. Some of the intensity of the peak at 1652 cm-1 can also be attributed to physisorbed water, which is expected with the exposure at room temperature for prolonged periods of time. Therefore, with the absence of the 1236 cm-1 peak and the assignment of the 1650 cm-1 feature to physisorbed H2O in this case, the peak at 1450 cm-1 is assigned as the bulk carbonate species. However, we suggest below that other surface species may also lend some intensity to this 1650 cm-1 feature. The Ba-formate feature at 1587 cm-1 further increased with water exposure, and peaks centered at 1394, 1450, and 1653 also increased in intensity during the extended (9 and 21 min) water exposures (spectra in thin solid and shaded lines, respectively). Previous work with BaO/Al2O3 samples has demonstrated that the addition of H2O converts surface nitrates and dispersed carbonate species to larger, more bulk-like, phases.11,14 It is proposed here that this also occurs for formate and carbonate species on the Pt/BaO/Al2O3 sample. This should explain the observed “sharpening” and intensity increases of the 1585 and 1450 cm-1 peaks. The agglomeration of the Bacontaining phases could also result in greater exposure of the Al2O3 support surface, which might allow formation of Al-COX surface species. Thus, under the experimental condition applied here, Al-formate species could also form, and would account for the intensity of the composite features centered at 1398 and 1587 cm-1 (vibrational features for Al-formate have been reported at 1375, 1393, and 1598 cm-1 44). Growth of intensity near 1598 cm-1 would also contribute to the increase in the 1587 cm-1 peak. In turn, the Al-formate doublet (1375, 1393 cm-1) may be yielding some of the intensity of the broad feature at 1398 cm-1 in Figure 6. It is also conceivable that there is a significant contribution to the IR features at 1450 and 1650 cm-1 from Al-related formates, carbonates, or bicarbonates that could form in the presence of CO and H2O. Due the very strong overlap among the vibrational features of these species, it is very difficult to assign a number of these IR bands to specific surface species. The sample was then heated to 410 K in vacuum, and spectrum b shown in Figure 7 was obtained (for comparison, spectrum a of Figure 7 is a reproduction of spectrum d of Figure 6). The intensity of the 1650 cm-1 band decreased substantially

J. Phys. Chem. C, Vol. 112, No. 29, 2008 10957 as most of the physisorbed water desorbed, as would be expected under this condition. There was also a noticeable change in the 1450 cm-1 IR feature as its intensity decreased and its peak position shifted toward lower wavenumbers, likely due to the disappearance of the low-stability Al-carbonate or bicarbonate species. The temperature was then held at 410 K and 1 Torr H2O was again introduced, and after 5 min of H2O exposure, spectrum c was obtained. The peak intensity at 1587 cm-1 did not change relative to that in spectrum b. In the features below 1500 cm-1, the peak at approximately 1400 cm-1 increased in relative intensity when compared to those at 1350-1370 and 1450 cm-1. This result, along with the appearance of the small peak at 1747 cm-1, again indicates the formation of unidentate species, likely at the expense of the bidentate species as the surface concentration of adsorbed water increased. Further increasing the temperature to 470 K in vacuum resulted in no decrease in the intensities of IR features assigned to formate species (spectrum d), in agreement with the thermal stability of these species we have discussed above. The most prominent change in the IR spectra during this water exposure at elevated temperatures is the intensity gain of the bidentate carbonate species, which is evidenced by the development of a shoulder toward the low wavenumber side of the 1587 cm-1 formate band, as well as the change in shape (broadening) of the band centered at 1360 cm-1. Fully accounting for the formation of these carbonate species by formate decomposition is possible but unlikely, since the formate coverage only seems to decrease a little. Some contribution could also result from reaction of CO2 and BaO, with CO2 possibly originating from the hydrolysis reaction between the adsorbed NCO and the introduced water, a critical part of NOX reduction processes on a number of catalysts. With the addition of 1 Torr of H2O for 9 min at 470 K, the intensity of the IR band assigned to the Ba-formate species (1359, 1367, and 1585 cm-1) significantly decreased, while that of the carbonate features increased. This suggests that the formate species in the presence of water at elevated temperatures (g470 K) decompose, and this leads to the formation of both surface (1360 and 1570 cm-1) and bulk (1450 cm-1) carbonate species. In the TPD experiment after CO addition, but where no H2O was added, higher temperatures were required for observation of formate decomposition to carbonates. Therefore, the presence of H2O impacts the thermal stability of the formate species (probably inducing their hydrolysis and subsequent decomposition), causing a transition to the carbonates at lower temperatures than in the absence of water. H2O + CO2 Exposure. Finally, a similar set of experiments was performed to evaluate the effect of H2O on surface COX species under conditions where CO2 was used for the carbonate/ COX species formation rather than CO. The series of IR spectra obtained in this experiment are shown in Figure 8. Spectrum a was obtained after exposing the sample to 1 Torr of CO2 for 10 min, evacuating the cell and then exposing the sample to 1 Torr of H2O for 1 min, all at 300 K. In comparing this spectrum to that shown in Figure 1b, obtained after only CO2 exposure, little difference is observed, except for reduced bicarbonaterelated features in the latter case, although they are present. This demonstrates that bicarbonate formation does not easily occur from carbonate transformation, but must start with the surface hydroxyl species. In the presence of H2O, the surface hydroxyl groups are predominantly interacting with adsorbed water, suppressing the formation of surface bicarbonates. Furthermore, the H2O exposures at room temperature just after the CO2 exposure did not result in any observable change in the COX

10958 J. Phys. Chem. C, Vol. 112, No. 29, 2008

Figure 8. IR spectra obtained from a Pt/BaO/Al2O3 powder after (a) exposure to 1 Torr of CO2 at 300 K for 10 min followed by evacuation and exposing the sample to 1 Torr of H2O at 300 K for 1 min; (b) exposure to 1 Torr of H2O for 21 min at 300 K; (c) heating in vacuum to 410 K; (d) heating to 410 K in vacuum and then exposure to 1 Torr of H2O for 5 min at 410 K; (f) heating in vacuum to 470 K; (g) heating to 470 K in vacuum and then exposure to 1 Torr of H2O for 9 min at 470 K.

surface species. The data shown in spectrum c were obtained after heating the CO2- and H2O-exposed samples to 410 K in vacuum. Again, an increase in a feature being resolved at approximately 1390-1400 cm-1 relative to the 1360 cm-1 feature is apparent and diagnostic of the appearance of unidentate carbonates. This result, coupled with results discussed above, indicates that the transition from bidentate carbonate species to unidentate carbonates is not affected by the gas composition conditions used in this study. Exposure to H2O for 5 min at 410 K caused a further increase in the higherwavenumber shoulder of the 1300-1420 cm-1 feature (spectrum d of Figure 8), as well as an increase in the intensity of the band associated with bulk carbonates (1450 cm-1). Because the sample was not exposed to CO2 during the heat to 410 K, nor during the exposure to H2O at 410 K, these spectral changes in COX species are likely induced by the presence of H2O. Heating the sample to 470 K in vacuum resulted in a further increase in unidentate species. Exposure to H2O at 470 K caused another significant increase in the intensities of the ∼1400 (unidentate carbonate) and 1450 (“bulk” carbonate) cm-1 IR features. TPD carried out following the CO2 and H2O exposure of the catalyst at elevated temperatures revealed that no carbonates were desorbed as a result of H2O addition onto the CO2-saturated sample, only the just-described conversions. In comparing the results obtained after exposing the Pt/BaO/ Al2O3 sample to CO and H2O (Figures 6 and 7) versus CO2 and H2O (Figure 8), it is clearly apparent that different chemistry occurred. When the sample was exposed to CO at low temperatures, formate species are formed readily, and almost exclusively. Carbonates form from the formate species in the presence of H2O and at higher temperatures. In contrast, exposing the sample to CO2 resulted in almost exclusive formation of carbonates, with only a small amount of bicarbonates via reaction with residual surface hydroxyls on the surface. Exposing this sample to H2O did not lead to larger quantities of bicarbonate species as might have been expected. Apparently, carbonates do not transform to bicarbonates by reaction with H2O. With many mechanistic models of this system including BaCO3 species as reactive sites, and assigning relative reac-

Epling et al. tivities to different types of carbonates, understanding their stability and formation under normal operating conditions is needed. The results of this study demonstrate that carbonates are readily formed on this model Pt/BaO/Al2O3 sample by exposures to CO2 or CO + O2, and they are stable to relatively high operating temperatures for these catalysts. With gas-phase COX species, their thermal stability is strengthened, but the relative stabilities of the species as a function of temperature observed are the same as in the absence of COX species. Exposing the sample to CO at low temperatures resulted in the formation of formate species, which were not as thermally stable as the carbonates, but were present in significant quantities in the lower operating temperature regime of NSR catalysts. Their decomposition may contribute to increased surface and bulk carbonate species. Furthermore, the presence of H2O had an effect on formate and carbonate formation, their decomposition, and possibly their crystallite size as well. These experimental studies of the relative stabilities of carbonate and formate species should be useful for the development of realistic mechanistic models. Conclusions A Pt/BaO/Al2O3 model NOX storage/reduction catalyst was characterized using Fourier transform infrared (FTIR) spectroscopy after exposure to CO2, CO, and CO + O2 at room temperature through >500 K. Barium carbonate species readily formed under all conditions, and were found to be stable in vacuum up to 773 K. At elevated temperatures, some of the bidentate carbonate species appear to transform to unidentate species. Residual hydroxide species on the catalyst, remaining on both the BaO and the alumina support and/or forming during cleaning with H2, resulted in the formation of bicarbonates. Formate species were observed to form when CO was used as the COX source, but decomposed at moderate temperatures, forming more carbonate species in the process. As has been suggested with nitrate species on BaO/Al2O3 samples, H2O exposure leads to phase changes and enhanced interconversions between multiple COX-containing species. Acknowledgment. W.S. Epling gratefully acknowledges the financial support of Natural Sciences and Engineering Research Council of Canada. Support for PNNL staff was provided by the U.S. Department of Energy (DOE), Office of Freedom Car and Vehicle Technologies. The work was performed in the Environmental Molecular Sciences Laboratory (EMSL) at Pacific Northwest National Laboratory (PNNL). The EMSL is a national scientific user facility and supported by the U.S. DOE Office of Biological and Environmental Research. PNNL is a multiprogram national laboratory operated for the U.S. Department of Energy by Battelle Memorial Institute under Contract DE-AC06-76RLO 1830. References and Notes (1) Campbell, L.; Danzinger, R.; Guth, E.; Padron, S. U.S. Patent 5,451,558, 1994. (2) Takahashi, N.; Shinjoh, H.; Iijima, T.; Suzuki, T.; Yamazaki, K.; Yokota, K.; Suzuki, H.; Miyoshi, N.; Matsumoto, S.; Tanizawa, T.; Tanaka, T.; Tateishi, S.; Kasahara, K. Catal. Today 1996, 27, 63. (3) Miyoshi, N.; Matsumoto, S.; Katoh, K.; Tanaka, T.; Harada, J.; Takahashi, N.; Yokota, K.; Sgiura, M.; Kasahara, K. SAE Tech. Pap. Ser. 1995, 950809. (4) Epling, W.; Campbell, L.; Yezerets, A.; Currier, N.; Parks, J. Catal. ReV. 2004, 46, 163. (5) Balcon, S.; Potvin, C.; Salin, L.; Tempere, J. F.; Djega-Mariadassou, G. Catal. Lett. 1999, 60, 39.

Pt/BaO/γ-Al2O3 NOX Storage/Reduction Catalyst (6) Rodrigues, F.; Juste, L.; Potvin, C.; Tempere, J. F.; Blanchard, G.; Djega-Mariadassou, G. Catal. Lett. 2001, 72, 59. (7) Amberntsson, A.; Persson, H.; Engstrom, P.; Kasemo, B. Appl. Catal., B: 2001, 31, 27. (8) Lietti, L.; Forzatti, P.; Nova, I.; Tronconi, E. J. Catal. 2001, 204, 175. (9) Toops, T.; Smith, D.; Partridge, W.; Epling, W.; Parks, J. In Proceedings of the Third Joint Meeting of the U.S. Sections of the Combustion Institute; 2003. (10) Epling, W.; Campbell, G.; Parks, J. Catal. Lett. 2003, 90, 45. (11) Kim, D. H.; Chin, Y.-H.; Kwak, J. H.; Szanyi, J.; Peden, C. H. F. Catal. Lett. 2005, 105, 259. (12) Toops, T.; Smith, D.; Epling, W.; Parks, J.; Partridge, W. Appl. Catal., B: 2005, 58, 255. (13) Hendershot, R. J.; Vijay, R.; Snively, C. M.; Lauterbach, J. Appl. Surf. Sci. 2006, 252, 2588. (14) Szanyi, J.; Kwak, J.; Kim, D. H.; Wang, X.; Chimentao, R.; Hanson, J.; Epling, W.; Peden, C. H. F. J. Phys. Chem. C 2007, 111, 4678. (15) Broqvist, P.; Panas, I.; Gronbeck, H. J. Phys. Chem. B 2005, 109, 9613. (16) Nova, I.; Castoldi, L.; Lietti, L.; Tronconi, E.; Forzatti, P. Catal. Today 2002, 75, 431. (17) Cant, N.; Patterson, M. Catal. Lett. 2003, 85, 153. (18) Theis, J.; Gobel, U.; Kogel, M.; Kreuzer, T. P.; Lindner, D.; Lox, E.; Ruwisch, L. SAE Tech. Pap. Seri. 2002, 01–0057. (19) Tuttlies, U.; Schmeiber, V.; Eigenberger, G. Top. Catal. 2004, 3031, 187. (20) Piacentini, M.; Maciejewski, M.; Baiker, A. Appl. Catal., B: 2005, 59, 191. (21) Piacentini, M.; Maciejewski, M.; Baiker, A. Appl. Catal., B: 2006, 66, 126. (22) Piacentini, M.; Strobel, R.; Maciejewski, M.; Pratsinis, S. E.; Baiker, A. J. Catal. 2006, 243, 43. (23) Morterra, C.; Ghiotti, G.; Boccuzzi, F.; Coluccia, S. J. Catal. 1978, 51, 299. (24) Evans, J.; Whateley, T. Trans. Faraday Soc. 1967, 63, 2769.

J. Phys. Chem. C, Vol. 112, No. 29, 2008 10959 (25) Shen, J.; Cortright, R. D.; Chen, Y.; Dumesic, J. A. J. Phys. Chem. 1994, 98, 8067. (26) Krupay, B. W.; Amenomiya, Y. J. Catal. 1981, 67, 362. (27) Prinetto, F.; Ghiotti, G.; Nova, I.; Tronconi, E.; Forzatti, P. J. Phys. Chem. B 2001, 105, 12732. (28) Iordan, A; Zaki, M. I.; Kappenstein, C J. Chem. Soc., Faraday Trans. 1993, 89, 2527. (29) Phillip, R.; Fujimoto, K. J. Phys. Chem. 1992, 96, 9035. (30) Kantschewa, M.; Albano, E. V.; Ertl, G.; Knozinger, H. Appl. Catal. 1983, 8, 71. (31) Ballinger, T. H.; Yates, J. T. Langmuir 1991, 7, 3041. (32) Baltrusaitis, J.; Jensen, J. H.; Grassian, V. H. J. Phys. Chem. B 2006, 110, 12005, and references therein. (33) Lercher, J. A.; Colombier, C; Noller, H J. Chem. Soc., Faraday Trans. 1984, 80, 949. (34) Jacobs, G.; Keogh, R. A.; Davis, B. H. J. Catal. 2007, 245, 326. (35) Liu, Z.; Anderson, J. J. Catal. 2004, 224, 18. (36) Frola, F.; Prinetto, F.; Ghiotti, G.; Castoldi, L.; Nova, I.; Lietti, L.; Forzatti, P. Catal. Today 2007, 126, 81. (37) Szanyi, J.; Kwak, J. H.; Hanson, J.; Wang, C. M.; Szailer, T.; Peden, C. H. F. J. Phys. Chem. B 2005, 109, 7339. (38) Busca, G.; Lorenzelli, V. Mater. Chem. 1982, 7, 89. (39) Liu, Z.; Zhou, Y.; Solymosi, F.; White, J. Surf. Sci. 1991, 245, 289. (40) Liu, J.; Wang, Y.; Lan, G.; Zheng, J. J. Raman Spectrosc. 2001, 32, 1000. (41) Lesage, T; Verrier, C; Bazin, P; Saussey, J; Daturi, M Phys. Chem. Chem. Phys. 2003, 5, 4435. (42) Pihl, J.; Daw, S.; Toops, T. Presentation at 10th DOE Crosscut Workshop on Lean Emissions Reduction Simulation, May 2007, Dearborn, MI. (43) Szailer, T.; Kwak, J. H.; Kim, D. H.; Hanson, J. C.; Peden, C. H. F.; Szanyi, J. J. Catal. 2006, 239, 51. (44) Amenoiya, Y. J. Catal. 1979) , 57, 64.

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