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J. Phys. Chem. C 2009, 113, 7336–7341
Effects of Sulfation Level on the Desulfation Behavior of Presulfated Pt-BaO/Al2O3 Lean NOx Trap Catalysts: A Combined H2 Temperature-Programmed Reaction, in Situ Sulfur K-Edge X-ray Absorption Near-Edge Spectroscopy, X-ray Photoelectron Spectroscopy, and Time-Resolved X-ray Diffraction Study Do Heui Kim,*,† Janos Szanyi,† Ja Hun Kwak,† Xianqin Wang,‡ Jonathan C. Hanson,§ Mark Engelhard,† and Charles H. F. Peden† Institute for Interfacial Catalysis, Pacific Northwest National Laboratory, Richland, Washington 99354, Department of Chemical, Biological and Pharmaceutical Engineering, New Jersey Institute of Technology, Newark, New Jersey 07102, and Chemistry Department, BrookhaVen National Laboratory, Upton, New York 11973 ReceiVed: January 12, 2009; ReVised Manuscript ReceiVed: February 26, 2009
Desulfation by hydrogen of presulfated Pt (2 wt %)-BaO(20 wt %)/Al2O3 with various sulfur loading (S/Ba ) 0.12, 0.31, and 0.62) were investigated by combining H2 temperature programmed reaction (TPRX), X-ray photoelectron spectroscopy (XPS), in situ sulfur K-edge X-ray absorption near-edge spectroscopy (XANES), and synchrotron time-resolved X-ray diffraction (TR-XRD) techniques. We find that the amount of H2S desorbed during the desulfation in the H2 TPRX experiments is not proportional to the amount of initial sulfur loading. The results of both in situ sulfur K-edge XANES and TR-XRD show that at low sulfur loadings, sulfates were transformed to a BaS phase and remained in the catalyst rather than being removed as H2S. On the other hand, when the deposited sulfur level exceeded a certain threshold (at least S/Ba ) 0.31) sulfates were reduced to form H2S, and the relative amount of the residual sulfide species in the catalyst was much less than at low sulfur loading. Unlike samples with high sulfur loading (e.g., S/Ba ) 0.62), H2O did not promote the desulfation for the sample with S/Ba of 0.12, implying that the formed BaS species originating from the reduction of sulfates at low sulfur loading are more stable to hydrolysis. The results of this combined spectroscopy investigation provide clear evidence to show that sulfates at low sulfur loadings are less likely to be removed as H2S and have a greater tendency to be transformed to BaS on the material, leading to the conclusion that desulfation behavior of Pt-BaO/Al2O3 lean NOx trap catalysts is markedly dependent on the sulfation levels. 1. Introduction Internal combustion engines operating under highly oxidizing conditions, such as diesel and lean burn gasoline engines, exhibit high fuel efficiency and, therefore, lowered emissions of green house gases such as CO2. Removal of harmful nitrogen oxides (NOx) from the exhaust in the presence of excess oxygen, however, presents a great challenge to the catalysis community because traditional three-way catalysts (TWC) can not reduce NOx under these conditions. Among the approaches to remove NOx from lean-burn engines, urea and hydrocarbon selective catalytic reduction (SCR) and lean-NOx traps (LNTs, also known as NOx storage/reduction (NSR) catalysts or NOx adsorbers) are three promising technologies.1 In the LNT catalyst, which consists of a precious metal (Pt), a NOx storage element (alkali and/or alkaline earth oxides) and an alumina support, the active oxide material (typically BaO) takes up NOx under lean engine operation conditions and stores them as nitrates.2,3 In a brief rich cycle, these nitrates are released from the active storage phase, and NOx is then subsequently reduced to N2 on the precious metal component of the catalyst. * To whom correspondence should be addressed. E-mail: do.kim@ pnl.gov. † Pacific Northwest National Laboratory. ‡ New Jersey Institute of Technology. § Brookhaven National Laboratory.
However, the stability against SO2 poisoning is considered to be biggest major drawback of the LNT technology. It is generally accepted that SO2 is oxidized to SO3 over Pt sites similarly to NO oxidation to NO2 reaction required for effective NOx storage. SO3 then reacts with BaO to form BaSO4, which is thermodynamically more stable than Ba(NO3)2. Thus, the formation of BaSO4 gives rise to the elimination of the active sites for NOx adsorption, consequently leading to deactivation of the catalyst.4 SO2 in low concentrations reduces the ability of the catalyst to store NOx, and it has been shown that the degree of deactivation closely follows the total amount of SO2 exposure to the catalyst.5 Desulfation of LNT materials (essentially decomposition and/or reduction of BaSO4) requires periodic high temperature excursion (e.g., 873 K and higher) that can eventually lead to irreversible deactivation due, primarily, to Pt sintering and/or BaAl2O4 formation via a solid state reaction between BaO and alumina.6 The catalyst after desulfation procedures, even at considerably high temperatures of up to 973 K, contains residual sulfur, resulting in the decrease of the overall NOx storage capacity.7 It has been shown that different types of sulfates (surface/ bulk barium sulfates and aluminum sulfates) are formed via the reaction of SO3, in turn produced by the Pt-catalyzed oxidation of SO2, with the respective adsorption sites, and that these various sulfate species can be differentiated with characterization tools such as X-ray diffraction (XRD)8 and IR.9 For example,
10.1021/jp900304h CCC: $40.75 2009 American Chemical Society Published on Web 04/03/2009
Presulfated Pt-BaO/Al2O3 Lean NOx Trap Catalysts the XRD results of Liu and Anderson8 demonstrated that the presence of Pt facilitated the formation of a bulk BaSO4 phase, implying the central role of Pt in the sulfation process. Understanding the desulfation mechanisms of the reactions of sulfates with reductants is of considerable practical importance, since these processes determine the required frequency of regeneration that, in turn, is directly related to the fuel economy penalty incurred with the use of the LNT technology. The desulfation of the NOx storage components strongly depends on the nature of reductants, among which hydrogen has been found to be superior to CO and hydrocarbons.8 In addition, while the regeneration of the sulfated catalyst leads to an efficient decomposition of sulfates, some sulfur remains trapped in the form of BaS which is very stable under rich conditions.10 Our group11 has reported that the desulfation processes of sulfated Pt-BaO/Al2O3 materials vary with baria loading and are more facile at lower loadings. This implies that the initial morphology of BaO/BaCO3 before sulfation is an important factor in determining the ease of sulfur removal from the catalyst. These findings were based on the results of a combined H2 temperature programmed reaction (TPRX), transmission electron microscopy (TEM) with energy dispersive spectroscopy (EDS), and in situ time-resolved X-ray diffraction (TR-XRD) study. In our previous study, although we demonstrated a relationship between the extent of desulfation and the initial morphology of the barium containing phase, the desulfation behavior as a function of sulfur loading has not been reported. Identifying the optimum conditions where sulfur removal is maximized will provide a path to enhance fuel economy of the overall LNT system because it can save the use of reductant by reducing the number and enhancing the effectiveness of the desulfation processes. In this contribution, we prepared three sulfated PtBaO(20)/Al2O3 samples with different sulfur loadings (the barium containing phase prior to sulfation was a mixture of BaO and BaCO3).12 To investigate the desulfation with H2 as a function of sulfur loading, we applied H2 TPRX. In addition, in situ S XANES and synchrotron TR-XRD were carried out under hydrogen flow during linear ramping, conditions that are essentially identical to those used for H2 TPRX. Such measurements provide complementary information on the chemical nature of residual sulfur, and the morphology changes that accompany the desulfation processes, respectively. 2. Experimental Section A 200 m2/g surface area γ-Al2O3 (Condea) was used as the support for all catalysts prepared in this study. 20 wt % BaO/ Al2O3 samples were prepared by the incipient wetness method using an aqueous Ba(NO3)2 solution and were then followed by drying in an oven at 393 K. Because of the limited solubility of Ba(NO3)2 in aqueous solution, multiple impregnation steps were performed to reach the desired Ba loading. An aqueous Pt(NH3)4(NO3)2 solution was applied to the Ba(NO3)2/Al2O3 samples, to prepare Pt-BaO/Al2O3 with Pt loading of 2 wt %. Following another drying step, this material was then subjected to a calcination process in flowing 5% O2/He. The furnace temperature was raised to 773 K at a rate of 1 K/min and maintained at that temperature for 5 h. Information about Pt particle size and dispersion have been reported previously.13 Prior to sulfation, a small amount of sample (∼0.15 g) was placed into a quartz reactor and treated with 10% O2 in He at 773 K for an additional 2 h. Typical sulfations were performed at 575 K in a gas mixture of 50 ppm of SO2 and 10% O2, balanced with helium (total flow rate of 150 cm3/min). To obtain samples with varying extents of sulfation, the exposure time of
J. Phys. Chem. C, Vol. 113, No. 17, 2009 7337 the clean samples to the SO2 containing gas mixture was varied. Sulfation levels, labeled as S/Ba throughout the text, were calculated on the bases of exposure time, resulting in S/Ba of 0.12, 0.31, and 0.62. The sulfated samples are designated as Pt-BaO/Al2O3-x, where x is S/Ba ratio. Note that these sulfation levels were a simple function of the SO2 exposures since no SO2 breakthrough was observed with an MKS Minilab mass spectrometer (MS) during sulfation. H2 TPRX (temperature programmed reaction) experiments were performed for the sulfated samples in the same reactor by raising the temperature to 1073 K at a rate of 8 K/min under flowing 20% H2/He. The desorbed gases were detected with the same mass spectrometer. In addition, in order to investigate the effect of H2O on the desulfation of Pt-BaO/Al2O3 sample with S/Ba ) 0.12, we performed H2 TPRX in the presence of 4% H2O and 10% H2 in He. The TR-XRD experiments and sulfur K-edge XANES were carried out at beamlines X7B and X19A of the National Synchrotron Light Source at Brookhaven National Laboratory, respectively. The detailed experimental setup of the TR-XRD has been described elsewhere.14 A small amount of sulfated sample was placed in a sapphire capillary tube and heated at 10 K/min from 300 to 1073 K while continuously flowing a 5% H2 in He gas mixture. XRD patterns were collected in situ eVery two minutes during the temperature ramping. The X19A beamline used to collect the S K-edge XANES data is equipped with a double-crystal Si(111) monochromator. Fluorescence measurements were performed in an in situ quartz reactor using a Lytle detector. The design of the reactor was adapted from a previous publication15 of Bare et al. A pellet of the sample (diameter 12 mm) was placed in the homemade sample holder, which is located inside the quartz reactor. The sample was heated at 10 K/min from 300 to 573 K, followed by heating at 2 K/min from 573 to 973 K while continuously flowing a 5% H2 in He gas mixture. Sulfur K-edge XANES spectrum was collected in situ every 9.5 min (thus, every 19 K) during the temperature ramping. The X-ray energy for the K-edge of elemental sulfur (Aldrich) was used for calibrating the energyscale. Its peak position was assigned an energy value of 2472 eV, and scans ranging from 20 eV below to 50 eV above the absorption edge were collected for all subsequent samples with a step size of 0.08 eV. The statistical noise at X19A is better than 0.001%. The white line intensity of the sulfate peak in the XANES spectra was obtained by the difference in the absorption intensities between the baseline at 2465 eV and the sulfate peak at ∼2482 eV. X-ray photoelectron spectroscopy (XPS) experiments were carried out for selected samples in the analysis chamber of a Physical Electronics Instruments Quantum 2000, using Al KR X-rays and a pass energy of 71 eV. The position and intensity of the Al 2s peak at 119.2 eV were used as references. 3. Results and Discussion In a preceding study, we found that desulfation behavior of model LNT catalysts was strongly dependent on barium loading which, in turn, determines the surface-to-bulk ratio of BaO species in BaO/Al2O3.11 Pt-BaO/Al2O3 catalysts with low (8 wt %) BaO loading contain predominantly surface BaO/BaCO3 species prior to sulfation and exhibit more facile desulfation during reduction with H2 at lower temperatures than Pt-BaO/ Al2O3 samples with high BaO loading (20 wt %) where a large fraction of the BaO is present as a bulk particulate phase. These results suggest that the nature of the BaO phase morphology (surface vs bulk) prior to sulfation has a strong impact on the
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Kim et al.
Figure 1. H2 TPRX data for Pt-BaO(20)/Al2O3 with varying sulfation levels, quantified as S/Ba 0.12, 0.31, and 0.62.
extent of desulfation and the temperature where it occurs. The question, however, that remained unanswered was the dependence of desulfation is on the initial sulfur loading at constant BaO loading (i.e., 20 wt %). Therefore, we extended our studies here to understand the desulfation behavior of Pt-BaO(20)/Al2O3 samples sulfated to varying extents by controlling the SO2 exposure times, resulting in sulfated samples with S/Ba ratio of 0.12, 0.31 and 0.62. H2 TPRX was applied to investigate the reactivity of sulfate species with hydrogen by monitoring the temperature where H2S evolved and quantifying the amount of H2S desorbed. In previous studies,16,17 this technique was used to identify the various sulfate species formed on alumina and on the surface and in the bulk of barium oxide particles. Figure 1 shows H2 TPRX data for the three sulfated Pt-BaO(20)/Al2O3 samples studied here, with the amounts of H2S desorbed during H2 TPRX up to 1173 K being drastically different. H2 TPRX spectrum of the Pt-BaO(20)/Al2O3-0.62 demonstrates two peaks with maximum at 673 and 873 K, which can be assigned to H2S evolution due to the desulfation of aluminum sulfates and barium sulfates, respectively, based on previous studies.11,17 However, no H2S evolution from the Pt-BaO(20)/Al2O3-0.12 sample was seen during the H2 TPRX, while the Pt-BaO(20)/Al2O3-0.31 sample showed only a trace amount of H2S evolution at ∼975 K. These results imply that sulfur containing species present for samples at low S/Ba ratios (873 K). As described in the experimental section, H2 TPRX was performed on the lightly sulfated Pt-BaO(20)/Al2O3-0.12 sample in the presence of 4% H2O and 10% H2. However, no H2S was detected by the mass spectrometer during H2 TPRX (not shown). This suggests that the promotional effect of H2O on the sulfur removal is absent for this low sulfur containing sample. In addition, XPS was used to estimate the residual sulfur level after the H2 TPRX experiments both in the absence and presence of H2O. The estimated sulfur concentrations obtained from XPS on these samples were 1.05 and 1.33% in the absence and presence of H2O, respectively, essentially identical levels for the two samples within the error range. Clearly, unlike the sample with high sulfur loading (S/Ba of 0.62), H2O does not aid the removal of the residual sulfur species from samples with low sulfur loadings. The fact that we did not see any observable H2S during H2 TPRX nor any differences in the amount of residual sulfur based on the XPS analysis of the two post-H2 TPRX samples suggests that BaS species formed on the sample with low sulfur loading is very resistant to reaction with H2O. We propose that the low reactivity of BaS phase thus formed on the low loaded sample is ascribed to their much larger crystallite size, thereby suppressing the reaction with water. It can be summarized that for the sample with low sulfur loading, the reaction (eq 1) proceeds quite readily; however, reaction 2 for desulfation hardly occurs at all due to the highly crystalline nature of BaS in this case. Therefore, the overall amount of sulfur before and after desulfation remains essentially constant on the catalyst, and the only change during the process is, in fact, the transformation from BaSO4 to BaS. On the other hand, when the sulfur loading exceeds some level (at least S/Ba of 0.31 according to H2 TPRX), the reaction (eq 2) occurs more significantly based on the H2 TPRX results, probably due to the formation of a more reactive BaS phase with increasing sulfur loading. Thus, these results show evidence about for different reactivities of sulfate species with H2 depending on the total sulfur loading. A practical implication of this work is to suggest when to regenerate sulfated catalysts and the frequency of the regeneration, an important factor in determining the fuel economy penalty of the overall LNT system. 4. Conclusions The desulfation behavior of presulfated Pt (2 wt %)-BaO (20 wt %)/Al2O3 as a function of sulfation loading was studied by using a combination of H2 TPRX, in situ sulfur K-edge XANES, XPS, and TR-XRD techniques. We found that the adsorbed sulfur species show different desulfation behavior depending on the degree of sulfation such that initially loaded sulfur species are less reactive with hydrogen and have a greater tendency to form highly crystalline BaS rather than being removed from the catalyst. On the other hand, sulfur species deposited above a certain level were more readily desorbed in the form of H2S, rather than remaining on the catalyst. Unlike the sample with high sulfur loading (S/Ba ) 0.62), H2O does not promote the desulfation for the sample with
J. Phys. Chem. C, Vol. 113, No. 17, 2009 7341 S/Ba of 0.12, demonstrating that the formed BaS species on this sample with low sulfur loading are more resistant to reaction with H2O. Acknowledgment. The authors would like to thank Dr. Wen Wen, Dr. Khalid Syed, and Nebojsa Marinkovic at the National Synchrotron Light Source (NSLS) for help with the TR-XRD and sulfur K-edge XANES spectroscopy measurements. Use of the NSLS at Brookhaven National Laboratory (BNL), was supported by the U.S. Department of Energy (DOE), Office of Science/Basic Energy Sciences, under Contract No. DE-AC02-98CH10886. The authors also give thanks to Dr. Simon Bare (UOP) for help with the design of our in situ S XANES reactor. Financial support was provided by the U.S. DOE, Office of Energy Efficiency and Renewable Energy/Vehicle Technologies Program. Many of the experiments were performed in the Environmental Molecular Sciences Laboratory (EMSL) at Pacific Northwest National Laboratory (PNNL). The EMSL is a national scientific user facility supported by the U.S. DOE, Office of Science/Biological and Environmental Research. PNNL is a multiprogram national laboratory operated for the U.S. DOE by Battelle Memorial Institute under Contract DE-AC06-76RLO 1830. J.C.H. was supported through Contract DE-AC02-98CH10086 with the U.S. DOE, Office of Science/Basic Energy Sciences, Division of Chemical Sciences. References and Notes (1) Liu, Z. M.; Woo, S. I. Catal. ReV. Sci. Eng. 2006, 48, 43. (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) Takeuchi, M.; Matsumoto, S. Top. Catal. 2004, 28, 151. (4) Lietti, L.; Forzatti, P.; Nova, I.; Tronconi, E. J. Catal. 2001, 204, 175. (5) Engstrom, P.; Amberntsson, A.; Skoglundh, M.; Fridell, E.; Smedler, G. Appl. Catal. B 1999, 22, L241. (6) Epling, W. S.; Campbell, L. E.; Yezerets, A.; Currier, N. W.; Parks, J. E. Catal. ReV. Sci. Eng. 2004, 46, 163. (7) Kim, D. H.; Chin, Y. H.; Muntean, G.; Yezerets, A.; Currier, N.; Epling, W.; Chen, H. Y.; Hess, H.; Peden, C. H. F. Ind. Eng. Chem. Res. 2007, 46, 2735. (8) Liu, Z. Q.; Anderson, J. A. J. Catal. 2004, 228, 243. (9) Su, Y.; Amiridis, M. D. Catal. Today 2004, 96, 31. (10) Rohr, F.; Peter, S. D.; Lox, E.; Kogel, M.; Sassi, A.; Juste, L.; Rigaudeau, C.; Belot, G.; Gelin, P.; Primet, M. Appl. Catal. B 2005, 56, 201. (11) Kim, D. H.; Szanyi, J.; Kwak, J. H.; Szailer, T.; Hanson, J.; Wang, C. M.; Peden, C. H. F. J. Phys. Chem. B 2006, 110, 10441. (12) Szanyi, J.; Kwak, J. H.; Kim, D. H.; Burton, S. D.; Peden, C. H. F. J. Phys. Chem. B 2005, 109, 27. (13) Kwak, J. H.; Kim, D. H.; Szailer, T.; Peden, C. H. F.; Szanyi, J. Catal. Lett. 2006, 111, 119. (14) Wang, X. Q.; Hanson, J. C.; Frenkel, A. I.; Kim, J. Y.; Rodriguez, J. A. J. Phys. Chem. B 2004, 108, 13667. (15) Bare, S. R.; Mickelson, G. E.; Modica, F. S.; Ringwelski, A. Z.; Yang, N. ReV. Sci. Instrum. 2006, 77, 1. (16) Elbouazzaoui, S.; Corbos, E. C.; Courtois, X.; Marecot, P.; Duprez, D. Appl. Catal. B 2005, 61, 236. (17) Wei, X. Y.; Liu, X. S.; Deeba, M. Appl. Catal. B 2005, 58, 41. (18) Solomon, D.; Lehmann, J.; Martinez, C. E. Soil Sci. Soc. Am. J. 2003, 67, 1721. (19) Dathe, H.; Jentys, A.; Lercher, J. A. Phys. Chem. Chem. Phys. 2005, 7, 1283. (20) Kim, D. H.; Kwak, J. H.; Szanyi, J.; Cho, S. J.; Peden, C. H. F. J. Phys. Chem. C 2008, 112, 2981. (21) Rodriguez, J. A.; Jirsak, T.; Freitag, A.; Hanson, J. C.; Larese, J. Z.; Chaturvedi, S. Catal. Lett. 1999, 62, 113. (22) Szanyi, J.; Kwak, J. H.; Hanson, J.; Wang, C. M.; Szailer, T.; Peden, C. H. F. J. Phys. Chem. B 2005, 109, 7339. (23) Szanyi, J.; Kwak, J. H.; Kim, D. H.; Wang, X. Q.; Hanson, J.; Chimentao, R. J.; Peden, C. H. F. Chem Commun. 2007, 984, 1. (24) Kim, D. H.; Kwak, J. H.; Wang, X. Q.; Szanyi, J.; Peden, C. H. F. Catal. Today 2008, 136, 183.
JP900304H
