pubs.acs.org/Langmuir © 2009 American Chemical Society
Reactivity of a Thick BaO Film Supported on Pt(111): Adsorption and Reaction of NO2, H2O, and CO2 Kumudu Mudiyanselage,† Cheol-Woo Yi,‡ and Janos Szanyi*,† †
Institute for Interfacial Catalysis, Pacific Northwest National Laboratory, P.O. Box 999, MSIN: K8-87, Richland, Washington 99352, and ‡School of Biological Sciences and Chemistry and Institute of Basic Science, Sungshin Women’s University, Seoul 136-742, Korea (ROK) Received April 17, 2009. Revised Manuscript Received June 19, 2009 Reactions of NO2, H2O, and CO2 with a thick (>20 monolayer equivalent (MLE)) BaO film supported on Pt(111) were studied with temperature programmed desorption (TPD) and X-ray photoelectron spectroscopy (XPS). NO2 reacts with a thick BaO layer to form surface nitrite-nitrate ion pairs at 300 K, while only nitrates form at 600 K. In the thermal decomposition process of nitrite-nitrate ion pairs, first nitrites decompose and desorb as NO. Then nitrates decompose in two steps: at lower temperature with the release of NO2 and at higher temperature, nitrates dissociate to NO þ O2. The thick BaO layer converts completely to Ba(OH)2 following the adsorption of H2O at 300 K. Dehydration/ dehydroxylation of this hydroxide layer can be fully achieved by annealing to 550 K. CO2 also reacts with BaO to form BaCO3 that completely decomposes to regenerate BaO upon annealing to 825 K. However, the thick BaO film cannot be converted completely to Ba(NOx)2 or BaCO3 under the experimental conditions employed in this study.
1. Introduction Lean-burn engines, which are operated under oxygen-rich conditions, provide high fuel efficiency for motor vehicles. However, under these highly oxidizing conditions, the reduction of nitrogen oxides (NOx) from the exhaust gas emission becomes difficult and conventional three-way catalysts are not effective in reducing NOx. Therefore, the development of new alternative strategies, such as NOx storage and reduction (NSR) and selective catalytic reduction (SCR) are necessary to replace the conventional three-way catalysts to remove NOx from lean-burn engines. One of the most promising approaches for lean-NOx removal process is NSR catalysis,1 which operates under cyclic (lean and rich) conditions. First, under a highly oxidizing environment (lean conditions), NO from the exhaust emission reacts with oxygen over Pt to form NO2, which can subsequently adsorb on the BaO storage material to form Ba(NO3)2. Then the reaction conditions are changed to a reducing environment (fuel-rich conditions) for a short period and the stored nitrate species are released as NOx, which subsequently react with reducing agents over Pt, producing N2, H2O, and CO2 (depending on the reducing agent). However, tightening emission standards and limitations of current NSR catalysts require improvements in their performance. To improve the efficiency of the NOx removal/conversion process, understanding of the elemental processes in NSR catalysis is essential. Although there have been a number of fundamental studies on NSR catalysts2-6 since they were introduced in 1990s,1 only *Corresponding author. E-mail:
[email protected]. (1) Takahashi, N.; Shinjoh, H.; Iijima, T.; Suzuki, T.; Yamazaki, K.; Yokota, K.; Suzuki, H.; Miyoshi, N.; Matsumoto, S.-i.; Tanizawa, T.; Tanaka, T.; Tateishi, S.-s.; Kasahara, K. Catal. Today 1996, 27, 63. (2) Amberntsson, A.; Skoglundh, M.; Ljungstrom, S.; Fridell, E. J. Catal. 2003, 217, 253. (3) Epling, W. S.; Campbell, L. E.; Yezerets, A.; Currier, N. W.; Parks, J. E. Cat. Rev. - Sci. Eng. 2004, 46, 163. (4) Epling, W. S.; Yezerets, A.; Currier, N. W. Catal. Lett. 2006, 110, 143. (5) Epling, W. S.; Yezerets, A.; Currier, N. W. Appl. Catal., B. 2007, 74, 117. (6) 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. (7) Bowker, M.; Stone, P.; Smith, R.; Fourre, E.; Ishii, M.; de Leeuw, N. H. Surf. Sci. 2006, 600, 1973.
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limited work has been reported on single-crystal-based model systems.7-15 These systems can provide detailed microscopic information on the elementary processes in NSR catalysis without having to deal with the complexity of real catalysts. Another advantage of these model systems is the applicability of a wide variety of surface science techniques to interrogate their chemical and structural properties. Reactions of NO2 with BaO nanostructures deposited on Al2O3/NiAl(110) have been studied extensively.8,9,14,16-18 Libuda and co-workers showed that NO2 reacts with BaO on Al2O3/NiAl(110) to form initially surface nitrites at 300 K which then converts slowly to surface nitrates at 300 K. At 500 K, the conversion of nitrites to surface nitrates is rapid and is followed by formation of ionic nitrates.8 They also investigated the coordination geometries of nitrates with the combination of density functional theory (DFT) and infrared (IR) spectroscopy, and confirmed the formation of both bridging and monodentate nitrates on BaO/Al2O3/NiAl(110).9 Yi et al.14 provided the first experimental confirmation of the theoretically predicted formation of nitrite-nitrate ion pairs at the initial stage of NO2 uptake on BaO, and very recently Libuda and co-workers also reported the simultaneous formation of nitrites and nitrates at 100 K.19 However, Libuda et al. did not observe nitrite-nitrate ion pair formation at 300 K, in contrast to Yi and Szanyi who (8) Desikusumastuti, A.; Happel, M.; Dumbuya, K.; Staudt, T.; Laurin, M.; Gottfried, J. M.; Steinruck, H.-P.; Libuda, J. J. Phys. Chem. C 2008, 112, 6477. (9) Desikusumastuti, A.; Staudt, T.; Gr€onbeck, H.; Libuda, J. J. Catal. 2008, 255, 127. (10) Ozensoy, E.; Peden, C. H. F.; Szanyi, J. J. Catal. 2006, 243, 149. (11) Ozensoy, E.; Peden, C. H. F.; Szanyi, J. J. Phys. Chem. B 2006, 110, 8025. (12) Stone, P.; Ishii, M.; Bowker, M. Surf. Sci. 2003, 537, 179. (13) Tsami, A.; Grillo, F.; Bowker, M.; Nix, R. M. Surf. Sci. 2006, 600, 3403. (14) Yi, C.-W.; Kwak, J. H.; Szanyi, J. J. Phys. Chem. C 2007, 111, 15299. (15) Bowker, M. C., M.; Hall, M.; Fourre, E.; Grillo, F.; McCormack, E.; Stone, P.; Ishii, M. Top. Catal. 2007, 42-43, 341. (16) Desikusumastuti, A.; Laurin, M.; Happel, M.; Qin, Z.; Shaikhutdinov, S.; Libuda, J. Catal. Lett. 2008, 121, 311. (17) Staudt, T.; Desikusumastuti, A.; Happel, M.; Vesselli, E.; Baraldi, A.; Gardonio, S.; Lizzit, S.; Rohr, F.; Libuda, J. J. Phys. Chem. C 2008, 112, 9835. (18) Yi, C.-W.; Szanyi, J. J. Phys. Chem. C 2009, 113, 2134. (19) Desikusumastuti, A.; Staudt, T.; Happel, M.; Laurin, M.; Libuda, J. G. J. Catal. 2008, 260, 315.
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reported the formation of these ion pairs in the reaction of NO2 with a thick BaO film (>30 monolayer equivalent (MLE)) on Al2O3/NiAl(110) up to 300 K sample temperature.18 Studies on metal supported BaO model systems are rare, and only a limited amount of data has been previously reported on NOx chemistry of BaO/Pt(111)7,12,15 Bowker and co-workers first characterized the surface structure of BaO thin films deposited on Pt(111) by scanning tunneling microscopy (STM).7,15 They showed the presence of a (2 2) reconstructed BaO(111) surface, and they also prepared a BaO2 layer with (100) termination. They also applied STM to investigate the reactivity of BaO on Pt(111) toward NO and O2, and found that a mixture of O2 and NO reacted with BaO to form Ba(NO3)2.12,15 The preparation and reactivity of BaO layers on Cu(111) have been studied by Tsami et al.13 The BaO films prepared were characterized by low energy electron diffraction (LEED) and identified as BaO(100). They found that the BaO layer on Cu(111) did not react with NO or NO/O2 mixture but readily reacted with NO2 to form Ba(NO2)2. They also extended their investigation toward the reactivity of a BaO layer with H2O and CO2 and reported on the relative stabilities and decomposition mechanisms of Ba(OH)2 and BaCO3 phases. This is the only single-crystal-based model study under ultrahigh vacuum (UHV) conditions, to the best of our knowledge, which has investigated the reactivity of BaO with H2O and CO2 in a metal single-crystal supported model system. However, the interpretation of the data on the chemistry of the BaO/Cu(111) model system is complicated by the fact that the Cu substrate could easily be oxidized and then intermixed with the BaO phase. The very strong influence of mixed oxide formation on the NOx chemistry in BaO-based model systems was first realized by Ozensoy et al.10 and Yi and Szanyi,20 and later confirmed by Libuda and co-workers,17,19 who systematically investigated the reactivity of Ba-aluminates with NO2. The understanding of the chemistry of BaO with CO2 and H2O is important because they are both present in large amounts under real operating conditions of NSR catalysts. Even though they may not be directly involved in the NOx storage process, they influence the nature of the active storage sites or phase (BaO, Ba(OH)2, or BaCO3). Some experimental studies have already indicated that the storage compound is indeed BaCO3 initially.21-23 The competitive storage of NO2 and CO2 has also been suggested.22 Furthermore, both CO2 and H2O were found to decrease the storage capacity of NSR catalysts, with H2O having a stronger effect than CO2.24 However, in the presence of H2O, promoting effects on the NOx storage at lower temperatures and inhibiting effects at higher temperatures were also reported.21 Despite the extensive studies in the effect of CO2 and H2O on the NOx storage of NSR catalysts, the nature of the active storage sites as well as mechanisms of the NOx storage and removal process in the presence of CO2 and H2O are far from being well understood. Therefore, understanding the reactivities of individual components of the exhaust emission with the active storage phase (BaO) is important. Such information will be useful to understand the mechanisms of the NOx storage processes in the simultaneous presence of NO2, H2O, and CO2. In order to develop highly efficient NSR catalysts, understanding the effect of the interactions between BaO and Pt and (20) Yi, C.-W.; Szanyi, J. J. Phys. Chem. C 2009, 113, 716. (21) Lietti, L.; Forzatti, P.; Nova, I.; Tronconi, E. J. Catal. 2001, 204, 175. (22) Balcon, S.; Potvin, C.; Salin, L.; Tempere, J. F.; Djega-Mariadassou, G. Catal. Lett. 1999, 60, 39. (23) Cant, N. W.; Patterson, M. J. Catal. Today 2002, 73, 271. (24) Hendershot, R. J.; Vijay, R.; Snively, C. M.; Lauterbach, J. Appl. Surf. Sci. 2006, 252, 2588.
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the role of the BaO/Pt interface on elementary processes in NSR catalysis is important. Libuda and co-workers have recently shown the dependence of the surface nitrite formation on the Pd/BaO interactions in a BaAl2O4/Pd/Al2O3 model system.25 The findings of this study also indicate the importance of understanding and controlling of Pt/BaO interactions in nanostructured systems. Model systems consisting of BaO nanostructures on a Pt(111) substrate allow us to study the elementary processes occurring in NSR catalysts and their dependence on BaO/Pt interactions, as well as the role of the BaO/Pt interface on those processes. This is the first study in an ongoing investigation of the adsorption and reactions of NO2 with different BaO/Pt(111) systems. As a first step, in order to understand the chemistry of the pure BaO with NO2 without any influence from the Pt(111) substrate, we prepared a thick BaO (>20 MLE) layer on the Pt(111) single crystal. (NO2 itself is not a major component of a lean exhaust gas stream, like NO, CO2, and H2O. NO from the exhaust emission is oxidized over Pt to form NO2 during the storage cycle of NSR catalysis. However, in this study, the Pt sites are not accessible because they are completely covered by a thick BaO layer.) Currently, we are investigating the chemistry of NO2 on thin (∼3 MLE) and submonolayer coverages (20 MLE) BaO film, first the desired amount of Ba was deposited onto a N2O4 multilayer on Pt(111) crystal (kept at 90 K) by reactive layer-assisted deposition (RLAD) using a resistively heated Ba doser, and then the thus formed BaNxOy layer was thermally decomposed by annealing to 1000 K. The BaO film was characterized by XPS. The chemistry of NO2, CO2, and H2O on this thick BaO film on Pt(111) was studied by TPD and XPS. NO2, CO2, and H2O were introduced into the UHV chamber by pinhole dosers and delivered to the sample surface through collimating tubes. In all the TPD experiments, a sample heating rate of 2 K s-1 was used. (25) Desikusumastuti, A.; Staudt, T.; Qin, Z. H.; Happel, M.; Laurin, M.; Lykhach, Y.; Shaikhutdinov, S.; Rohr, F.; Libuda, J. ChemPhysChem 2008, 9, 2191.
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Figure 1. XP spectra of (a) N 1s and (b) O 1s obtained following the saturation exposure of the thick BaO film to NO2 at 300 K.
3. Results and Discussion 3.1. Adsorption and Reaction of NO2 with BaO/Pt(111). The XP spectra of N and O 1s obtained following saturation exposure of thick BaO film to NO2 at 300 K sample temperature are shown in panels a and b of Figure 1, respectively. Figure 1a shows two peaks for N 1s indicating the presence of two different nitrogen-containing species. Based on previously reported binding energies of nitrites and nitrates formed on barium oxide,8,13,14,17,26 the N 1s peaks at 404.3 and 408.1 eV can be assigned to nitrogen atoms in nitrite and nitrate species, respectively. In addition, the approximately identical intensities of these two N 1s peaks indicate the presence of equal amounts of nitrites and nitrates on the surface, suggesting the formation of nitritenitrate ion pairs by a cooperative NO2 adsorption mechanism, in accordance with the predictions of theoretical studies27-29 and its experimental confirmation by Yi et al.14 We also observed similar intensities of these two N 1s peaks in the XP spectrum obtained after exposing a thick BaO layer to NO2 at lower exposures than that needed to saturate the BaO film at 300 K sample temperature. This ion pair formation also occurs even at 90 K sample temperature.14 Theoretical calculations also proposed two possible mechanisms for the formation of nitrite-nitrate ion pairs. In one mechanism, the ion pair formation is oxide surface mediated. In this mechanism, first a NO2 molecule adsorbs and creates an electronic defect by forming a surface NO2- species. Then the (26) (27) 137. (28) (29)
Schmitz, P.; Baird, R. J. Phys. Chem. B 2002, 106, 4172. Broqvist, P.; Panas, I.; Fridell, E.; Persson, H. J. Phys. Chem. B 2002, 106, Gr€onbeck, H.; Broqvist, P.; Panas, I. Surf. Sci. 2006, 600, 403. Schneider, W. F. J. Phys. Chem. B 2004, 108, 273.
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second NO2 molecule heals this defect by adsorbing on the oxide surface (donating an electron to the surface), and this leads to the formation of a nitrite-nitrate ion pair.28 In the other mechanism, the formation of ion pairs is through some combination of electron transfer (redox) and acid-base interactions, where a transfer of an electron between two adsorbed NO2 molecules at neighboring surface sites yields a nitrite-nitrate ion pair.29 Theoretical calculation further predicted the formation of two kinds of nitrite-nitrate ion pairs, [NO2--Ba2þ-[OsNO2]-] and [NO3--Ba2þ-[Os-NO]-] (where OS is a surface O).27 These two forms can be interchangeable by transferring O between them. For the first time, this ion pair formation was observed experimentally by Yi et al. in the study of NO2 adsorption on thick BaO (>30 MLE)/Al2O3/NiAl(110).14,18 However, recently, Libuda and co-workers also reported the simultaneous formation of nitrites and nitrates on the mixed barium aluminum oxide nanoparticles (BaAl2xO1þ3x) on Al2O3/NiAl(110) at 100 K and confirmed the existence of a cooperative adsorption mechanism leading to nitrite-nitrate ion pair formation.19 They concluded that the reaction mechanism does not depend on the formation of BaAl2xO1þ3x but instead it is controlled by the temperature. They associated this low temperature cooperative adsorption mechanism with the presence of N2O4 dimers on the surface at 100 K and proposed a dimerization (N2O4) mediated mechanism, which is different from that predicted by theoretical studies. In their proposed mechanism, first the symmetric O2N-NO2 (N2O4 dimer) isomerizes to asymmetric dimer (ONO-NO2) and then nitrosonium nitrate (NOþNO3-) forms. Finally, this Lewis acid/ base pair directly reacts with Lewis acid/base site on the BaAl2xO1þ3x particles. However, the absence of N2O4 at 300 K under UHV conditions should completely rule out the involvement of this mechanism in the formation of nitrite-nitrate ion pairs in our study. Therefore, we conclude that the ion pair formation in our study is controlled by the nature of the BaO film (thickness > 20 MLE) and not the sample temperature. Figure 1b shows two peaks for O 1s at 527.8 and 533.5 eV. The peak at 527.8 eV can be assigned to O of BaO, while the peak at 533.5 eV is due to the O of both nitrite and nitrate species. However, we do not have sufficiently high resolution to observe two resolved O 1s peaks for nitrite and nitrate species with our XP spectrometer. We performed thermal desorption experiments to investigate the relative stabilities and decomposition mechanisms of nitrites and nitrates of the ion pairs formed at 300 K. Figure 2 shows a series of TPD spectra obtained after NO2 adsorption at 300 K for mass fragments 30, 32, and 46. At the lowest NO2 exposure, a broad peak is observed between 460 and 650 K in the NO TPD spectrum. The intensity of this broad peak increases with increasing NO2 exposure. At higher exposures, a new feature appears around 669 K and the lower temperature peak shifts to higher temperatures. The NO TPD spectrum obtained after saturation exposure of the thick BaO film to NO2 shows two resolved features at 590 and 669 K. At the lowest NO2 exposure, no O2 TPD peak was observed as shown in Figure 2b. Increasing the NO2 exposure leads to the appearance of a broad high temperature desorption feature with a maximum desorption rate at around 860 K. Further increasing the NO2 exposure results in the saturation of this high temperature desorption peak and leads to the development of a lower temperature desorption feature. At saturation, these two features exhibit maximum desorption rates at 669 and 864 K. At the lowest three NO2 exposures, no desorption of NO2 was observed. Then the TPD spectra for NO2 show a peak at 577 K, which shifts to higher temperatures with increasing NO2 exposure, and at saturation the maximum desorption rate is observed at 596 K. It is interesting to note that Langmuir 2009, 25(18), 10820–10828
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Figure 2. TPD spectra for NO2 adsorption on a thick BaO film at 300 K with increasing NO2 exposures for mass fragments (a) 30, (b) 32, and (c) 46.
the NO2 exposure regime where no NO2 desorption was seen coincides with the absence of the lower temperature O2 desorption feature. Even though the NO TPD spectrum obtained after maximum NO2 exposure shows only two resolved features, this TPD curve can be deconvoluted into three peaks as shown in the inset of Figure 2a. This peak deconvolution is based on the cracking pattern of NO2 on the mass spectrometer, the TPD spectrum of O2, and the XPS data which indicates the presence of nitritenitrate ion pairs (equal amount of nitrites and nitrates on the surface). First the NO2 peak was multiplied by a factor based on the cracking pattern of NO2 in our mass spectrometer to get the peak at 596 K. Then the O2 TPD peak was multiplied by a factor to match its intensity with the high temperature NO desorption peak at around 669 K. By subtracting these two NO features (obtained from the NO2 cracking and from the O2 peak) from the measured total NO peak, we obtained the feature centered at 583 K. Consistent with and based on this TPD data, we propose a three-step process for the thermal decomposition and desorption of nitrite-nitrate ion pairs. First, nitrites release NO as shown by the peak at 583 K in the inset in Figure 2a. Then nitrates Langmuir 2009, 25(18), 10820–10828
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decompose in two steps: at lower temperature, nitrates dissociate and release NO2 only as shown by the peak at 596 K in Figure 2c, and cracking of this NO2 gives a NO peak at 596 K as shown by the inset in Figure 2a. At higher temperature, nitrates decompose into NO þ O2, giving peaks at 669 K as shown in panels (a) and (b). Here, we denote these two types of nitrates, which dissociate at lower and higher temperatures, as type 1 and type 2 nitrates, respectively. At the lowest NO2 exposures, desorption features of NO and O2 that correspond to type 2 nitrates were not observed. At this NO2 exposure, the higher temperature desorption feature of O2 was also not detected as shown in Figure 2b. After achieving a certain coverage of this O, we can observe desorption features of NO and O2 due to the dissociation of type 2 nitrates. In this study, we could not determine the nature of these two nitrate species from the available experimental data. Even with the addition of the vibrational spectroscopy data, we were not able to unambiguously differentiate between the two nitrate species in our previous studies.14,18 One possible explanation is that, after the formation of nitrite-nitrate ion pairs, the bulk of the BaO layer begins to slowly convert into barium nitrate, resulting in the existence of two different types of nitrates (i.e., surface and bulk). These two nitrates, however, show different thermal decomposition properties. One of them directly decomposes to NO2, while the other one (at higher temperature) decomposes as NO þ O2. These decomposition processes may be further influenced by the presence of BaO2 that forms during the initial surface nitrite decomposition. The high temperature O2 desorption feature at 864 K is most likely due to the recombinative O desorption originating from the thermal decomposition of BaO2. The formation of BaO2 was evidenced by the presence of an O 1s XPS feature centered at 531.1 eV, which is comparable to that reported previously for barium peroxide (530.8 eV).30 However, this binding energy value is not consistent with the reported value of 533 eV for BaO2 on Pt(111), although the BaO film in that reference was much thinner, and the BaO2/Pt(111) electronic interaction may have influenced the binding energy value measured.15 The nitrite-nitrate ion pairs are formed after exposing the thick BaO film to NO2 at 300 K. In order to identify the species formed after exposing the thick BaO film to NO2 at elevated temperatures, we performed a series of TPD experiments, and the obtained results are shown in Figure 3. NO2 exposures enough to saturate the thick BaO film at 300 K were applied in each experiment at the indicated temperatures. The amount of NOx desorbed increases with increasing NO2 adsorption temperature up to 500 K. The maximum amount of NOx desorption was observed after NO2 adsorption at 500 K. At higher temperatures (above 500 K), the TPD peak area decreases with increasing sample temperature. This decrease in the NOx amount desorbed from the sample above 500 K is due to the sample temperature during NO2 exposure being higher than that at the onset of nitrate and nitrite decomposition. Figure 3b shows that the O2 TPD peak intensity increases with increasing adsorption temperature up to 575 K and then starts to decrease due to the dissociation of type 2 nitrates. Similarly, Figure 3c shows that the intensity of the NO2 peak increases with increasing sample temperature during NO2 adsorption up to 550 K and then it decreases. However, in contrast to the constant peak temperature of the O2 TPD feature, the temperature of maximum desorption rate of NO2 increases from 596 to 636 K with increasing NO2 adsorption temperature up to 575 K. The increase in TPD peak temperature indicates the enhancement of the stability of type 1 nitrates at NO2 adsorption temperatures higher than 300 K. In addition to the main peak at (30) Gauzzi, A.; Mathieu, H. J.; James, J. H.; Kellett, B. Vacuum 1990, 41, 870.
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Figure 4. XP spectra of N 1s obtained following exposure of the thick BaO film to NO2 at the indicated temperatures. Amount of NO2 dosed is the same in each experiment.
Figure 3. TPD spectra for NO, O2, and NO2 following the adsorption NO2 on the thick BaO film at the indicated temperatures. Amount of NO2 dosed is the same in each experiment.
669 K, the O2 TPD spectra show a weak feature around 876 K, most probably associated with the decomposition of BaO2 formed in the thermal decomposition of nitrite species, as we have discussed above. XPS experiments were also performed on the BaO film exposed to NO2 at elevated temperatures. Figure 4 shows N 1s XP spectra obtained after adsorption of NO2 on a thick BaO film at 550, 575, and 600 K. At 550 K, the amount of nitrites formed is less than nitrates as shown by the lower intensity of the N 1s peak. Adsorption of NO2 at 575 K leads to the formation of mainly nitrates with a small amount of nitrites, while at 600 K only nitrates form. We have carried out NO2 adsorption experiments (data not shown) in a cyclic fashion to find whether the BaO layer can be converted completely to Ba(NOx)2. In these experiments, the BaO film was exposed to NO2 at 300 K followed by annealing the sample to 575 K in the absence of NO2. The annealing temperature of 575 K was chosen, since it is below the decomposition temperature of type 2 nitrates, while most of the nitrites decomposed at this temperature. A small fraction of type 1 nitrates, however, also decompose at this temperature. We mainly focused on the more stable and strongly bound nitrates that are the major 10824 DOI: 10.1021/la901371g
species present under practical conditions. After 3, 6, and 10 cycles of 300 K adsorption of NO2/575 K anneal, TPD spectra were recorded. Results of these experiments showed the gradual conversion of BaO to Ba(NO3)2 at elevated temperatures and the storage of a large amount of NOx as nitrates. However, saturation was not achieved even after the 10th adsorption/anneal cycle and the BaO film was not completely converted to Ba(NOx)2 under employed experimental conditions, similarly to the thick BaO layer on an alumina film we have reported earlier.18 3.2. Adsorption and Reaction of H2O on BaO/Pt(111). Figure 5a shows XP spectra of O 1s obtained before and after saturation exposure of BaO to H2O at 90 K. The intensity of the O 1s peak of BaO at 527.8 eV decreased after exposure to H2O, while a new peak appeared at 530.7 eV. This value is close to the 531.0 eV previously reported for Ba(OH)2 on Cu(111)13 and is in good agreement with the value given by Verhoeven and Van Doveren for Ba(OH)2.31 Therefore, we assign the peak at 530.7 eV binding energy to O atoms in Ba(OH)2. The XPS data clearly indicate that the BaO layer reacts with H2O to form Ba(OH)2 at 90 K. The significant intensity of the O 1s peak at 527.8 eV after saturation exposure of the BaO film to H2O also reveals that the BaO layer has only partially been converted to Ba(OH)2. However, we cannot identify the chemical composition of the Ba(OH)2 layer from the XPS data only. Ba(OH)2 can exist in a number of polymorphs and can form several amorphous hydrates, such as Ba(OH)2 3 8H2O and Ba(OH)2 3 3H2O, crystalline hydrates R-, β-, and γ-Ba(OH)2 3 H2O, and anhydrous crystalline forms R- and β-Ba(OH)2.32 Thermal decomposition of Ba(OH)2 3 8H2O and Ba(OH)2 3 3H2O leads to the formation of β-Ba(OH)2 3 H2O, the stable form of the monohydrate at ambient temperature. Dehydration of β- and γ-Ba(OH)2 3 H2O results in the formation of anhydrous β-Ba(OH)2 at 378 and 388 K, respectively. In vacuum (∼1 Pa), anhydrous β-Ba(OH)2 can be prepared by dehydration of Ba(OH)2 3 8H2O at 413 K.32 In order to get more insights into the species formed by the reaction between H2O and the thick BaO film at 90 K, we performed TPD experiments following the exposure of the BaO film to various amounts of H2O at 90 K. TPD spectra recorded following exposure of a thick BaO film to H2O at 90 K are shown in Figure 5b. At the lowest exposure, desorption occurs in a single peak at 477 K, and as the exposure increases a second peak appears at 445 K. With increasing H2O dosage, both peaks shift to higher temperatures. First the higher (31) Verhoeven, J. A. T.; Van Doveren, H. Appl. Surf. Sci. 1980, 6, 225. (32) Lutz, H. D.; Eckers, W.; Christian, H.; Engelen, B. Thermochim. Acta 1981, 44, 337.
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Figure 5. (a) XP spectra of O 1s obtained before and after the saturation exposure of BaO to H2O at 90 K. (b) TPD spectra obtained after exposing the thick BaO film to increasing exposure of H2O at 90 K.
temperature peak saturates, and a further increase in H2O exposure leads to the increase in the intensity of the lower temperature peak. At saturation, a single peak is observed in the TPD spectrum at 470 K with a higher temperature shoulder at around 490 K. The low temperature desorption features at 147 and 173 K are due to the desorption of weakly bound multi- and monolayer H2O, respectively, which is adsorbed to the surface of the Ba(OH)2. Since the desorption of H2O is observed above 400 K (TPD spectra in Figure 5b) while the dehydration of the hydrated forms of Ba(OH)2 discussed above occurs below 400 K, under the experimental conditions of our study, most likely anhydrous Ba(OH)2 forms when the thick BaO film is exposed to H2O. Anhydrous crystalline Ba(OH)2 also exists as two forms, R and β, and they show different IR spectra.33 The phase transformation from β-Ba(OH)2 to R-Ba(OH)2 occurs at 526 K.33 However, TPD spectra in Figure 5b show that Ba (OH)2 completely decomposes below 500 K. Based on the TPD data in our study, we can confirm the formation of anhydrous Ba (OH)2 by adsorbing H2O on thick BaO film at 90 K. Therefore, the main peak in the TPD spectrum, which was obtained after saturation H2O exposure, at 470 K is assigned to the desorption of H2O from the thermal decomposition of anhydrous Ba(OH)2 phase. The desorption feature observed at higher temperature in the TPD spectrum is most likely due to the surface OH groups that are formed during the dissociation of H2O on the BaO surface as has been suggested previously by Tsami et al.13 They (33) Cordfunke, E. H. P.; Booij, A. S.; Konings, R. J. M.; van der Laan, R. R.; Smit-Groen, V. M.; van Vlaanderen, P. Thermochim. Acta 1996, 273, 1.
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Figure 6. (a) TPD spectra obtained after exposing the thick BaO film to increasing exposure of H2O at 300 K. (b) XP spectra of O 1s obtained after saturation exposure of BaO to H2O at 300 K.
observed similar desorption features when a BaO film on Cu(111) was exposed to H2O. The presence of the higher temperature shoulder on the main peak at saturation indicates that the hydroxide phase is transformed back to an OH terminated oxide phase during the thermal decomposition process. Recently, theoretical calculation also showed the formation of hydroxide pairs following the adsorption of H2O on BaO(001).34,35 In order to investigate the reactivity of the thick BaO layer with H2O at 300 K, we performed TPD and XPS experiments. TPD spectra recorded following the exposure of the thick BaO film to various amounts of H2O at 300 K are shown in Figure 6a. At the lowest exposure, the TPD spectrum shows two resolved features at 446 and 490 K similar to the spectra obtained following the low exposure of H2O at 90 K as shown in Figure 5b. With increasing H2O exposure, desorption occurs in a single peak and the maximum desorption rate gradually shifts from 463 to 494 K. These TPD spectra follow a common leading edge. (The TPD spectrum obtained following the saturation exposure of the thick BaO layer to H2O at 90 K is also displayed in this figure for comparison.) Upon a certain H2O exposure, the TPD profile abruptly changes and displays two clearly resolved features. At saturation, the TPD spectrum shows maximum desorption rates at ∼482 and ∼520 K. Note that the TPD spectrum collected after the highest H2O exposure shows a small desorption feature with a maximum desorption rate at 420 K and can be assigned to the desorption of hydrating water molecules. Similar TPD spectra were observed by Yi and Szanyi following the adsorption of D2O (34) Gronbeck, H.; Panas, I. Phys. Rev. B: Condens. Matter Mater. Phys. 2008, 77, 245419. (35) Carrasco, J.; Illas, F.; Lopez, N. Phys. Rev. Lett. 2008, 100, 016101.
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Figure 7. XP spectra of (a) O 1s and (b) C 1s obtained before and after saturation exposure of the thick BaO film to CO2 at 300 K.
on thick BaO (>30 MLE) on Al2O3/NiAl(110).36 They also performed infrared reflection absorption spectroscopy (IRAS) experiments to identify the hydroxide phases that can form under different experimental conditions and to rationalize the TPD results. They showed that, upon D2O exposure of a thick BaO film, amorphous barium hydroxide formed at room temperature that readily converted to crystalline Ba(OD)2 phases during annealing in UHV. The formation of crystalline hydroxide phases was dependent on the initial D2O exposure at 300 K. Following low D2O exposures at 300 K results in the formation of amorphous Ba(OD)2 with no hydrating water, and only the R-Ba(OD)2 crystalline phase was observed after 400 K annealing. The sample that was exposed to D2O extensively (i.e., hydrated amorphous Ba(OD)2 formed) showed a series of phase transformations as the sample was annealed to increasingly higher temperatures: amorphous Ba(OD)2 3 x(D2O) f β-Ba(OD)2 3 yD2O f β-Ba(OD)2 f R-Ba(OD)2. We also observed similar IR spectra in this study (not shown) as reported by Yi et al. following the adsorption of H2O at 300 K. Therefore, we confirmed the results reported by Yi et al. and concluded that at lower exposures anhydrous amorphous Ba (OH)2 is formed at 300 K and then it converts to crystalline R-Ba (OH)2 following annealing to 400 K. However, at higher exposures, hydrated amorphous Ba(OH)2 is formed at 300 K, and it converts to β-Ba(OH)2 3 1H2O and β-Ba(OH)2 after annealing to 400 K. This β-Ba(OH)2 3 H2O releases its hydrating H2O at ∼480 K, as shown in Figure 6a, and converts to β-Ba(OH)2. Further annealing leads to the conversion to R-Ba(OH)2, and dehydration/
(36) Yi, C.-W.; Szanyi, J. J. Phys. Chem. C 2009, accepted.
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dehydroxilation is complete after annealing to temperatures higher than 550 K. Figure 6b shows the XP spectrum of O 1s obtained after the saturation exposure of BaO to H2O at 300 K. Unlike the spectrum obtained following the adsorption of H2O at 90 K (Figure 5a), this XP spectrum shows only one O 1s peak at 531.3 eV due to the presence of Ba(OH)2. The absence of the O 1s peak for BaO indicates that after the saturation exposure of BaO layer to H2O at 300 K the conversion to Ba(OH)2 is complete. In contrast, adsorption of H2O at 90 K converts the BaO layer to Ba(OH)2 only partially, since after surface hydroxylation the reaction stops due to the desorption of water before it could react with bulk BaO. 3.3. Adsorption and Reaction of CO2 with BaO/Pt(111). Figure 7 shows XP spectra of O 1s (panel a) and C 1s (panel b) obtained before and after saturation exposure of the thick BaO film to CO2 at 300 K. After the exposure to CO2, a new peak appears at 531.5 eV in the O 1s spectrum and is accompanied by a peak at 289.6 eV in the C 1s XP spectrum. These C and O 1s binding energies of the new compound are in good agreement with the values reported previously for BaCO3,13,31 indicating the formation of BaCO3 by adsorption/reaction of CO2 on the thick BaO film at 300 K. The presence of the O 1s peak of BaO with significant intensity at 527.8 eV after CO2 exposure indicates that the carbonate formation has proceeded only to a limited extent to form BaCO3 under the experimental conditions employed in this study. The comparison of the O 1s spectra before and after adsorbing H2O (at 90 K) and CO2 on BaO clearly indicates that the reaction of BaO with H2O proceeds much more efficiently than that with CO2 even at 90 K as shown in Figure 5a. It appears that BaCO3 forms a protective layer on the BaO phase that inhibits further reaction of the oxide with CO2 under the conditions applied. Nix and co-workers have also reported that BaCO3 formed on the BaO/Cu(111) system from the reaction between BaO and CO2 was localized to the surface rather than uniformly distributed throughout the film, that is, forming bulk carbonates.13 TPD spectra obtained following the exposure of thick BaO film to various CO2 exposures at 300 K are shown in Figure 8a. At the lowest CO2 exposure, decomposition occurs in a single peak at 788 K. As the exposure is increased, a second peak appears at 720 K and both peaks shift to higher temperatures. The higher temperature desorption peak saturates first, and with further increase in CO2 exposure the intensity of the lower temperature peak increases gradually. At the saturation CO2 exposure, two desorption features are observed at 748 and 825 K. In contrast to the phase transformation temperatures of Ba(OH)2 (discussed above), those of BaCO3 occur at higher temperatures. The occurrence of two phase transformations of BaCO3, that is, orthorhombic-to-hexagonal (1079 K) and hexagonal-to-cubic (1237 K), have been reported in an investigation of the decomposition of BaCO3 by thermoanalytical techniques.37,38 Therefore, we cannot expect any phase transformation of BaCO3 in the temperature range employed in this study. We assign these two peaks in the TPD spectra, similarly to the peak assignments in the TPD spectrum (following the adsorption of H2O at 90 K; Figure 5b) of Ba(OH)2 decomposition, to surface and bulk carbonate decomposition. The main CO2 peak in the TPD spectrum at 748 K is due to the thermal decomposition of bulklike BaCO3. The TPD feature observed at higher temperature is most likely due to the decomposition of surface carbonate groups that are formed during the adsorption of CO2 on the surface. (37) Arvanitidis, I.; Sichen, D.; Seetharaman, S. Metall. Mater. Trans. B 1996, 27, 409. (38) Lander, J. J. J. Am. Chem. Soc. 1951, 73, 5893.
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Figure 8. TPD spectra obtained (a) after exposing the thick BaO film to increasing exposure of CO2 at 300 K and (b) after exposing the thick BaO film to CO2 at the indicated temperatures.
Theoretical calculations also predicted the formation of a carbonate-type structure by chemisorbing CO2 on BaO(100).39,40 TPD spectra following the saturation exposure of the BaO film to CO2 at different temperatures are shown in Figure 8b. No significant changes in the TPD spectra were observed after adsorbing CO2 at 90 and 300 K, and only a small reduction of the TPD feature at 748 K was seen after exposing the thick BaO layer to CO2 at 600 K. Increasing the CO2 adsorption temperature to 750 K results in only a higher temperature desorption peak at 825 K. Due to the high sample temperature during CO2 adsorption (above the onset temperature for CO2 desorption), only surface carbonates formed after the exposure of the thick BaO layer to CO2 at 750 K.
4. Similarities and Differences in the Interactions of BaO with NO2, CO2, and H2O In this study, we confirmed the previous experimental observation of nitrite-nitrate ion pair formation and also validated the theoretical prediction of a cooperative adsorption mechanism. We also concluded that the cooperative adsorption did not occur through the dimerization (N2O4) mediated mechanism under the experimental conditions applied in this study. This nitrite-nitrate ion pair formation is controlled by the nature of the BaO film; that is, this mechanism is always in effect when chemically pure BaO is exposed to NO2. Nix and co-workers have observed different TPD spectra after adsorption of NO2 on an as-prepared BaO (39) Tutuianu, M.; Inderwildi, O. R.; Bessler, W. G.; Warnatz, J. J. Phys. Chem. B 2006, 110, 17484. (40) Karlsen, E. J.; Nygren, M. A.; Pettersson, L. G. M. J. Phys. Chem. B 2003, 107, 7795.
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layer (pure BaO surface) and on BaO layers which were previously subjected to NO2 adsorption and desorption processes.13 They showed that the formation of nitrite was the first step in the production of Ba(NO3)2, and they claimed that no direct formation of nitrate occurred on BaO/Cu(111) upon its exposure to NO2. However, our TPD spectra were always reproducible in a large number of experiments after NO2 adsorption and thermal decomposition. No influence from the Pt substrate on the chemistries of NO2, H2O, and CO2 with BaO was observed in our study due to the higher thickness of the BaO film. The differences between the chemistry reported for the BaO film supported on Cu(111) and that for BaO on Pt(111) arise from the film thickness applied in the two studies. Nix et al. studied a BaO film approximately 6 MLE thick,13 while the BaO film in our study was thicker than 20 MLE. The combination of the lower film thickness and the high reactivity of the Cu(111) substrate produced a highly unstable system, that led to the modification of the chemistry in comparison to the pure BaO films. The results reported here clearly indicate the presence of reactions between the thick BaO layer and NO2, H2O and CO2 even at 90 K. Uptake of NO2, H2O, and CO2 resulted not only in adsorption but also in conversion of the BaO layer to produce new compounds (Ba(NOx)2, Ba(OH)2, and BaCO3). The thermal stabilities of these phases vary in the order of BaCO3 > Ba(NOx)2 > Ba(OH)2, in concert with the results of Nix and co-workers.13 Ba(OH)2 dissociates at the lowest temperature and releases H2O, and then Ba(NOx)2 releases NOx and finally BaCO3 releases CO2 at the highest temperature. Not only the thermal stabilities but also the formation processes of these three phases (Ba(NOx)2, Ba(OH)2, and BaCO3) are different. The BaO layer completely converts to Ba(OH)2 by the adsorption of H2O at 300 K, while the formation of bulk carbonates is completely stopped after the initial formation of a protective carbonate layer. The bulk carbonate formation is very slow even in experiments performed in a cyclic fashion, similarly to the NO2 adsorption/annealing cyclic experiments described earlier. At 300 K, NO2 reacts to form nitrite-nitrate ion pairs and then reaction stops; the conversion of bulk BaO to Ba(NOx)2 can only be achieved at higher sample temperatures, or at elevated NO2 pressures as we have shown for the BaO (>30 MLE)/Al2O3/NiAl(110) system.20 BaO gradually converts to Ba(NOx)2 in cyclic experiments (NO2 adsorption/ annealing cycles); however, complete saturation was not achieved even after 10 cycles under the conditions applied in this study, most probably due to the diffusion limitation of bulk nitrate formation in the thick BaO film. Both this study and that performed by Nix and co-workers13 showed similar results for the formation and thermal decomposition of Ba(OH)2 and BaCO3. The similarities observed for the H2O and CO2 chemistries and the large dissimilarity observed for the NO2 chemistry in these two studies also underscores the differences in reactivities of these reactants. NO2 is one of the strongest oxidizing agents, and therefore, during NO2 adsorption/desorption studies, it can easily aim the oxidation of the underlying support material (e.g., Cu). The thus-formed oxide under the conditions of these studies may intermix with the oxide layer, resulting in a significant modification of the chemical properties of the oxide of interest (e.g., the formation of a mixed BaO/CuO phase in ref 13). On the other hand, the much less reactive compounds of CO2 and H2O do not cause significant intermixing of the oxide phases, or this process is much slower.
5. Conclusions Exposure of a thick BaO film to NO2 leads to the formation of nitrite-nitrate ion pairs by cooperative adsorption at 300 K, DOI: 10.1021/la901371g
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while only nitrates form at 600 K. In the thermal desorption process of nitrite-nitrate ion pairs, first nitrites decompose and desorb as NO. Then nitrates decompose in two steps: at lower temperature, nitrates dissociate and desorb as NO2 only, and at higher temperature nitrates dissociate and desorb as NO þ O2. The thick pure BaO film is extremely sensitive to water and is converted partially to the hydroxides by exposure to H2O at 90 K (or during the TPD) and converted completely by exposure to H2O at 300 K. Dehydration can fully be achieved by heating the hydroxides to 550 K. The surface OH groups that are formed during the dissociation of H2O on the surface are thermodynamically more stable than the Ba(OH)2 phase with respect to dehydration, and the hydroxide phase is transformed back to OH terminated oxide during the thermal decomposition process. CO2 reacts with BaO to form BaCO3, and it decomposes completely to
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regenerate BaO by annealing to 825 K. The thick BaO film could not be converted completely to Ba(NOx)2 or BaCO3 under employed experimental conditions of this study. Acknowledgment. We gratefully acknowledge the U.S. Department of Energy (DOE), Office of Basic Energy Sciences, Division of Chemical Sciences for the support of this work. The research described in this paper was performed at the Environmental Molecular Sciences Laboratory (EMSL), a national scientific user facility sponsored by the DOE Office of Biological and Environmental Research and located at Pacific Northwest National Laboratory (PNNL). PNNL is operated for the U.S. DOE by Battelle Memorial Institute under contract number DEAC05-76RL01830. C.-W.Y. also acknowledges the support of this work by Sungshin Women’s University Research Grant of 2008.
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