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desorption property of TiO2–ZrO2 having tolerance to SO2 poisoning
Applied Catalysis B: Environmental 74 (2007) 137–143 www.elsevier.com/locate/apcatb
NO adsorption/desorption property of TiO2–ZrO2 having tolerance to SO2 poisoning Kazuhiro Ito, Shuichi Kakino, Keita Ikeue, Masato Machida * Department of Nano Science and Technology, Graduate School of Science and Technology, Kumamoto University, 2-39-1 Kurokami, Kumamoto 860-8555, Japan Received 20 June 2006; received in revised form 10 December 2006; accepted 13 February 2007 Available online 16 February 2007
Abstract The effect of SO2 on the adsorption and desorption properties of NO was studied over TiO2–ZrO2 and g-Al2O3. The microporous binary oxide, TiO2–ZrO2, exhibited a large amount of sorptive NOx uptake (ca. 1 mmol g 1) at 25 8C in the presence of O2 and SO2. The irreversible chemisorption as nitrate (NO3) is a main cause for the large NO uptake. Desorption of NOx from TiO2–ZrO2 occurred at lower temperature of ca.450 8C, compared to 500 8C required for that from g-Al2O3. A similar trend was observed for the SOx desorption, because SO2 adsorbed on gAl2O3 is a stable sulfate-type (SO4), whereas less stable sulfite (SO3) is dominant species on the surface of TiO2–ZrO2. The SOx desorption could further be facilitated in the presence of reducing agents like CO. The thermal swing adsorption/desorption cycle experiment demonstrated that TiO2–ZrO2 can be used as an NOx adsorbent having improved tolerance to SO2. The SO2 tolerance can be explained by the weak basicity of TiO2– ZrO2 and the low activity to oxidize SO2 to form sulfate species on the surface. # 2007 Elsevier B.V. All rights reserved. Keywords: TiO2–ZrO2; Nitric oxide; Oxidative adsorption; SOx-tolerance
1. Introduction Interest in the application of sorptive NOx removal to emission control has increased markedly in the past decades. One excellent example can be seen in the automotive NOx storage and reduction (NSR) catalyst [1–4]. Although threeway catalysts for the gasoline engines operating at stoichiometric fuel:air ratio are able to purify HC, CO and NOx simultaneously, it is difficult to abate NOx emission from diesel or lean-burn gasoline engines. To solve this problem, the NSR catalyst contains NOx storage components (Ba or other alkaline-earth metals) in the vicinity of noble metal catalysts (Pt) on a support oxide (g-Al2O3). NO in the exhaust is first oxidized to NO2 and stored as Ba–NO3 under lean-burn conditions. By temporarily turning to rich condition, stored NOx is then released and simultaneously reduced by CO and hydrocarbons to N2 on noble metal catalysts. One serious drawback of the NSR catalyst is the deactivation caused by SO2
* Corresponding author. Fax: +81 96 342 3651. E-mail address: [email protected] (M. Machida). 0926-3373/$ – see front matter # 2007 Elsevier B.V. All rights reserved. doi:10.1016/j.apcatb.2007.02.008
originating from fuel [5]. Acidic gases, NO2 and SO2, should compete for the same adsorption sites and, generally, SO2 adsorption is stronger. When the adsorbed SO2 yields sulfate, the NOx storage should totally be lost. Therefore, the development of SO2-tolerant NOx-sorbing materials is now strongly requested. Another important requirement for NOx storage components is to improve NOx removal at lower temperatures. Increasingly stringent regulations of NOx emissions require efficient sorbents working even near ambient temperature. Recently, several researchers challenged to develop the SO2tolerant NOx adsorbent by employing non-alkaline solid as the storage components. Eguchi et al. [6] have reported the oxidative NO adsorption onto MnOx–ZrO2 at 200 8C. Thanks to the weaker basicity, the thermal desorption of NOx takes place at low temperatures of 400 8C, compared to 600 8C requested for the Ba-based adsorbents. A similar NO adsorbability can be demonstrated for MnOx–CeO2 in our previous studies [7], but their tolerance to SO2 should not be enough [8]. Centi et al. [9] reported that the resistance to deactivation by SO2 could be promoted in hydrotalcite-based supported Pt–Cu catalysts. It is well-known that TiO2 is tolerant
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to SO2 poisoning. Hachisuka et al. [10] and Huang et al. [11] studied TiO2 as an NOx storage component. Because of instability of sulfate on TiO2, the storage components after exposure to SO2 can easily be regenerated. Unfortunately, however, weak basicity of TiO2 led to a much smaller NOx storage capacity than that of Ba-containing sorbents. Eguchi and coworkers [12] recently reported the effect of additives (Li, Na, K, Cs, Sr, Ba, and La) on the NO uptake by Pt/TiO2 in the presence of SO2. The addition of Li caused the uptake of NO as well as SO2, but adsorbed SO2 was released at the lowest temperature. The different behavior of these additives can be explained by weak basicity of Li. As was evident from these previous studies, it is very difficult to achieve both NOx storage and SOx-tolerance simultaneously. For the design of SOx-tolerant NOx-sorbents, the precise control of the basicity is required not only to increase NOx uptake but also to liberate the adsorbed SO2 at lowest possible temperatures. As a candidate of such materials, TiO2-based oxides are promising. We have previously studied TiO2–ZrO2 as a support of Pt catalyst for NO–H2–O2 reactions at low temperatures of 100 8C [13–15]. Due to a number of Lewis base sites formed in accordance with the TiO2/ZrO2 ratio, a large amount of NO (0.4 mmol g 1 at 90 8C) was found to be stored during catalytic reactions. This is indicative of another possible application of TiO2–ZrO2 as an NOx-sorbing material, which can be used in the presence of SO2 at low temperatures. The present work aims to study the sorptive NO removal and SO2-tolerance of the TiO2–ZrO2 system at ambient temperature. Adsorption properties of NO and SO2 onto TiO2–ZrO2, g-Al2O3, SiO2, TiO2 and ZrO2 were compared to study the effect of surface basicity and the chemical state of adsorbates. The adsorption/ desorption cycles were performed using the temperature swing technique in the presence of excess O2 and SO2. 2. Experimental 2.1. Preparation and characterization Binary oxides, (n)TiO2–(1 n)ZrO2 (0 n 1), were prepared by the coprecipitation method [13]. Aqueous ammonia was added dropwise to a solution of ZrO(NO3)2 (97.0%, Wako) and [(CH3)2CHO]4Ti (95.0%, Wako) with vigorous stirring. The obtained precipitate was washed several times with deionized distilled water, followed by drying for overnight at 100 8C. The resulting solid product was calcined at 450 8C for 5 h in air. The other NOx adsorbents, g-Al2O3 (JRC, ALO-4) and SiO2 (JRC, SIO-5) were supplied by the Catalysis Society of Japan. The crystal structure of prepared materials was identified by X-ray diffraction using Cu Ka radiation (Rigaku Multiflex). The BET surface area was measured by a nitrogen isotherms adsorption system at 196 8C (BEL Japan, Belsorp). 2.2. NOx adsorption and desorption Sorptive removal of NOx was evaluated in a conventional flow system at atmospheric pressure. A granular sample
(10–20 mesh, 0.05 g) was fixed in a quartz tube (4 mm i.d.) by packing quartz wool at both ends of the adsorbent bed. The adsorbents were pretreated in a stream of 10% O2 and He balance at 400 8C for 60 min to remove adsorbed gases. After subsequent cooling to 25 8C, gaseous mixtures of 0.16% NOx, 0–10% O2, and 0–50 ppm SO2 in He were fed to the sample at W/F = 0.06 g s cm 3. The NOx desorption from adsorbents was observed by means of temperature-programmed desorption (TPD) measurement. NOx-adsorbed samples were heated to 800 8C at the rate of 10 8C min 1 in a stream of 10% O2 and He balance (50 cm3 min 1). The effluent gases were analyzed by a chemiluminescence NOx analyzer (Shimadzu NOA-7000) and an on-line quadruple mass spectrometer (Pfeiffer, Omnistar). 2.3. SOx adsorption and desorption The pretreated sample was placed at 25 8C in a flowing gas mixture of 50–100 ppm SO2, 10% O2 in He supplied at W/ F = 0.06 g s cm 3 for 60 min. After the treatment at various temperatures, the uptake of SO2 was determined by use of an Xray fluorescence analyzer (Horiba MESA-500W). The treatment by SO2 was followed by heating at 400–800 8C for 60 min in the oxidizing atmosphere (10% O2 and He balance) or reducing atmosphere (0.04% CO and He balance) to obtain the desorption property of SO2. 2.4. FT-IR spectra of adsorbed NOx/SOx FT-IR spectra of NOx/SOx species adsorbed on TiO2–ZrO2 were recorded on a Jasco FT-IR-610 spectrometer. The sample was placed in a temperature-controllable diffuse reflectance reaction cell (Jasco DR600A), which was connected to a gas flow system and a vacuum line. The sample was first treated in a stream of 10% O2 and He balance at 400 8C and next exposed to gas mixtures of 0.04% NO or 0.08% SO2, 10–20% O2, and He balance at 25 8C. This was followed by spectra measurement in a flowing He. All spectra, thus obtained were transformed into absorption spectra by the use of Kubelka–Munk function and referenced to those taken just before NO/SO2 admission. 2.5. Temperature swing adsorption/desorption cycles Temperature swing NOx adsorption/desorption cycles in the presence or absence of SO2 were performed by switching two gas feeds between 0.08% NO, 10% O2, 0 or 100 ppm SO2 in He and 0.04% CO in He. First, an NO adsorption step was carried out in a stream of 0.08% NO, 10% O2, 0 or 100 ppm SO2 in He (W/F = 0.06 g s cm 3) for 20 min at 25 8C. This was followed by a regeneration step in a stream of 0.04% CO in He (W/ F = 0.06 g s cm 3) at 600 8C for 10 min. The end of a cycle was followed by the rapid cooling to ambient temperature, where the NO concentration in the gas feed was measured before going to a next cycle. A water-cooled infrared image furnace (ULVAC RHL-E25) was used for alternating the reaction temperature between adsorption and desorption steps.
K. Ito et al. / Applied Catalysis B: Environmental 74 (2007) 137–143
Table 2 BET surface area and sorptive NO/NO2 uptake of TiO2–ZrO2, g-Al2O3 and SiO2
3. Results and discussion 3.1. Adsorptive uptake of NO/SO2 Table 1 summarizes phases, BET surface areas and NO uptakes of binary oxide, (n)TiO2–(1 n)ZrO2 (0 n 1), after calcination at 450 8C. Although ZrO2 (n = 0) and TiO2 (n = 1.0) were composed of crystalline phases, monoclinic/ tetragonal fluorite-type and anatase, respectively, the mixed oxide (n = 0.25 and 0.5) yielded an amorphous phase. The two crystalline oxides (n = 0 and 1.0) exhibited low surface areas of <90 m2 g 1, but amorphous-like oxides (n = 0.25 and 0.5) showed larger surface area of approximately 200 m2 g 1. TEM observation suggested the amorphous product was metastable mixtures of fine particles (ca. 2 nm in diameter) of TiO2 and ZrO2 having a low crystallinity. The NO breakthrough curves were measured in a flowing gas mixture of 0.16% NO, 10% O2, He balanced at 25 8C. Because neither N2O nor N2 were detected in the effluent gas, abatement of NOx attributes to the adsorptive uptake by solids. As shown in Table 1, the NOx uptake increased with an increase of the surface area and the maximum amount of the uptake was observed at n = 0.25–0.5. The NO uptake normalized by the surface area also showed the maximum at n = 0.25–0.5, suggesting that the NO adsorption site was created by combining TiO2 and ZrO2. According to work done by Arata et al. [16,17], this can be explained by a noncrystalline porous structure having hetero-junctions at the boundary of the two different oxide components having different oxygen coordination numbers, i.e., 6 for Zr and 8 for Zr. Due to excess positive and negative charge formed at the boundary, Zr acts like a Lewis acid while Ti–O acts as a base that causes the oxidative adsorption site of NO [14,15]. Table 2 compares the NO/NO2 uptakes for TiO2–ZrO2, gAl2O3, and SiO2 at 25 8C. These three metal oxides showed similar surface area in the range 177–207 m2 g 1. Nevertheless, their uptakes of NO/NO2 were quite different depending on the surface property. SiO2 exhibited the smallest NO uptake (0.06 mmol g 1) because of the very weak basicity of the surface. By contrast, TiO2–ZrO2 as well as g-Al2O3 exhibited ca. 10–15 times larger NO uptakes. Larger uptakes for NO2 compared to NO can be explained by the fact that NO2 is condensable and may cause stronger interactions with the base site on the solid. The NO uptake of TiO2–ZrO2 decreased with decreasing O2 concentration from 1.06 mmol g 1 (10% O2) to 0.13 mmol g 1 (0% O2). As thermodynamics predicts, NO Table 1 BET surface area and sorptive NO uptake of (n)TiO2–(1
Sample
BET surface area (m2 g 1)
NO uptake (m2 g 1)a
NO2 uptake (m2 g 1)a
g-Al2O3 SiO2 TiO2–ZrO2
177 192 207
0.61 0.06 1.06
1.05 0.23 1.19
a
0.16% NO or NO2, 10% O2, He balance, W/F = 0.06 g s cm 3, 25 8C.
should react with O2 to form NO2 at ambient temperature, but, at low NO concentration as in the present case (0.16%), the NO oxidation is very slow in the gas phase. However, this is not the case on the surface of metal oxides, where NO adsorption in the presence of O2 takes place via smooth oxidation of NO to NO2, presumably because of interactions with Ti–O species on the base site. The formed NO2 further reacts with the basic oxygen site to produce more stable NO3 species as detected by infrared measurement (see Fig. 3). Fig. 1 shows the amount of SO2 adsorption as a function of adsorption temperature. The SO2 uptake at 25 8C was not consistent with that of NOx uptake shown in Tables 1 and 2. The largest SO2 uptake observed for g-Al2O3 at 25 8C, 0.12 mmol g 1, was almost equal to the total amount of SO2 supplied in the gas feed. The SO2 uptake decreased with an increase of adsorption temperature, but g-Al2O3 retained almost constant uptake of ca. 0.08 mmol g 1 at 200 8C. Interestingly, the SO2 uptake onto TiO2–ZrO2 at 25 8C was much smaller (0.05 mmol g 1) in contrast to the largest NO/ NO2 uptakes (Table 2). Above 400 8C, the SO2 uptake became negligible on TiO2–ZrO2. Generally, the base site on the solid surface was occupied by NOx as well as SOx. However, this is not the case for TiO2–ZrO2, probably because SOx adsorbed on TiO2–ZrO2 are unstable than those on g-Al2O3. The effect of coexisting SO2 on adsorptive NO uptake was next studied. Fig. 2 compares the NO/SO2 uptakes onto gAl2O3 and TiO2–ZrO2 in a flowing 0.16% NO, 10% O2 and He balance (a) in the absence and (b) presence of 50 ppm SO2 at 25 8C. The amount of SO2 uptake by g-Al2O3, 0.11 mmol g 1, agreed with the accumulation of almost all SO2 supplied in the gas feed. The NO adsorption was suppressed to a significant extent in a presence of SO2. In accordance with Fig. 1, TiO2– ZrO2 showed the smaller SO2 uptake (0.07 mmol g 1). In both
n)ZrO2 at 25 8C
n
Phase
BET surface area (m2 g 1)
NO uptake (mmol g 1)a
0 0.25 0.50 0.75 1.00
Monoclinic/tetragonal ZrO2 Amorphous Amorphous Anatase Anatase
85 201 207 179 58
0.08 1.14 1.06 0.74 0.15
a
139
0.16% NO, 10% O2, He balance, W/F = 0.06 g s cm 3, 25 8C.
Fig. 1. SOx adsorption onto (n)TiO2–(1 n)ZrO2 and g-Al2O3 at different temperature. 50 ppm SO2, 10% O2, He balance, W/F = 0.06 g s cm 3, 60 min.
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Fig. 2. Coadsorption of NO and SO2 onto g-Al2O3 and TiO2–ZrO2 at 25 8C. (a) 0.16% NO, 10% O2, He balance, W/F = 0.06 g s cm 3 and (b) 0.16% NO, 50 ppm SO2, 10% O2, He balance, W/F = 0.06 g s cm 3. Total SO2 supplied: 0.12 mmol g 1.
cases, the amount of NO uptake decreased in the presence of SO2, whereas the sum of NO and SO2 uptake was kept constant. The result simply means that NO and SO2 compete to adsorb on the same surface site. The SO2/NO ratios on the surface were larger than those expected from their concentration in the gas phase, suggesting that SO2 should be bound to the surface more strongly than NO. Nevertheless, the NO adsorption on TiO2– ZrO2 is less sensitive to SO2. Fig. 3a shows the differential FT-IR spectra of NO adsorbed onto g-Al2O3 and TiO2–ZrO2, which are referenced to those taken just before NO admission. On TiO2–ZrO2, a set of sharp double bands at 1575 and 1280 cm 1, and a weak band at 1023 cm 1, which can be ascribed to NO stretching modes of bidentate nitrate (NO3), appeared. Because these peaks were intensified with an increase of O2 concentration, oxidative NO adsorption to form NO3 could be proved. The spectra of NO adsorbed on g-Al2O3 was almost same, but the peaks due to nitrate was less intense. Fig. 3b shows the differential spectra taken after exposure of g-Al2O3 and TiO2– ZrO2 to a mixture of 0.08% SO2, 20% O2 in He for 6 h at 25 8C. After exposing g-Al2O3 to the gas mixture, the adsorption band appeared at 1150 cm 1, which can be assigned to the S–O asymmetric stretching mode of sulfate (SO4). By contrast, TiO2–ZrO2 exhibits bands centered at 951 and 1016 cm 1, which are characteristic for asymmetric and symmetric stretching sulfite (SO3) vibration. The result suggested that SO2 is readily oxidized to the most stable species, SO4, over g-Al2O3, but less-oxidized species (SO3) are dominant on TiO2–ZrO2. Such different SOx species should be associated with the SO2 adsorption behavior (Fig. 1) as well as the tolerance to SO2 (Fig. 2) of the present system. The reason for the lack of ability to oxidize SO2 is not clear at the present stage, but it seems to be consistent with the stability of sulfate, i.e., the sulfates of Ti and Zr decompose at 450 8C, compared to more than 1000 8C requested for the decomposition of Al2(SO4)3.
Fig. 3. FT-IR spectra of (a) NOx adsorbed on g-Al2O3 and TiO2–ZrO2 after exposure to 0.04% NO, 10% O2, He balance at 25 8C for 1 h and (b) SOx adsorbed on g-Al2O3 and TiO2–ZrO2 after exposure to 0.08% SO2, 20% O2, He balance at 25 8C for 6 h.
3.2. Desorption of NO/SO2 The TPD profiles of NO after adsorption in the absence of SO2 are shown in Fig. 4. Although adsorption was carried out at ambient temperature, heating above 400 8C was required to complete the NOx desorption. This is indicative of the fact that NOx species are strongly chemisorbed as nitrate on the solid surface. Especially, the NOx desorption from g-Al2O3 and ZrO2 took place at the highest temperature (500 8C). The addition of TiO2 to ZrO2 shifted the desorption peak to lower temperatures and the minimal desorption temperature was attained for TiO2–ZrO2 (n = 0.5), which exhibited the largest NO uptake (Table 1). This result implies that the combination between TiO2 and ZrO2 increased the number of base site for NO adsorption, but the basicity of each site became rather weak. The thermal desorption of SOx can be estimated by measuring the residual SOx uptake after exposure to a stream of 100 ppm SO2, 10% O2, and He balance at 25 8C for 60 min and subsequent heating in a stream of 10% O2 and He balance at elevated temperatures for 60 min (Fig. 5). The initial SOx uptake was 0.18 mmol g 1 (g-Al2O3) and 0.11 mmol g 1 (TiO2–ZrO2), respectively. The thermal desorption of SOx from g-Al2O3 only was observed at 600 8C and did not complete even at 800 8C. The SOx desorption therefore requires
K. Ito et al. / Applied Catalysis B: Environmental 74 (2007) 137–143
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Fig. 6. Residual SOx uptake after heating in a reducing gas at elevated temperatures. SO2 adsorption: 100 ppm SO2, 10% O2, He balance, at 25 8C for 60 min, W/F = 0.06 g s cm 3. Desorption condition: 0.04% CO, He balance at 400–800 8C for 60 min, W/F = 0.06 g s cm 3.
The thermal SOx desorption was evaluated in a stream of 0.04% CO and He balance as shown in Fig. 6. In such a reducing atmosphere, the SOx desorption is expected to be facilitated because the adsorbed SOx species should become less stable. The SOx desorption from TiO2–ZrO2 started at 400 8C and finished at 700 8C. The SOx desorption from gAl2O3 also took place at lower temperatures, but it was still incomplete at 800 8C. These results indicate that not only NOx but also SOx can be liberated from TiO2–ZrO2 at lower temperature than those from g-Al2O3. 3.3. Thermal swing NOx adsorption/desorption cycles Fig. 4. NOx-TPD profiles of g-Al2O3 and (n)TiO2–(1 n)ZrO2 after adsorption in 0.16% NO, 10% O2, He balance at 25 8C. 10% O2, He balance, heating rate 10 8C min 1.
temperatures higher than those for NOx desorption as shown in Fig. 4. This is indicative of the higher stability of adsorbed SOx (sulfate) compared to adsorbed NOx (nitrate) on the surface of g-Al2O3. On the contrary, the SOx desorption from TiO2–ZrO2 started at lower temperature of 400 8C and finished at 800 8C. Again, the result can be explained by the fact that the chemisorption of SO2 onto TiO2–ZrO2 as sulfite (SO3) is weak compared to sulfate (SO4) bound to g-Al2O3.
Fig. 5. Residual SOx uptake after heating in an oxidizing gas at elevated temperatures. SO2 adsorption: 100 ppm SO2, 10% O2, He balance, at 25 8C for 60 min, W/F = 0.06 g s cm 3. Desorption condition: 10% O2, He balance at 400–800 8C for 60 min, W/F = 0.06 g s cm 3.
The stability of NOx adsorption/desorption cycles for the TiO2–ZrO2 system was evaluated using thermal swing technique. Figs. 7 and 8 display the change of effluent NOx concentration during the thermal swing adsorption–desorption cycles over gAl2O3 and TiO2–ZrO2, respectively. The adsorption was carried out at 25 8C for 20 min in a flowing mixture of 0.08% NO, 10% O2, and He balance (a) in the absence of SO2 or (b) in the presence of 100 ppm SO2. At the end of each adsorption step, the temperature was raised to 600 8C, where a mixture of 0.04% CO and He balance was supplied for 10 min. The adsorbed NOx species were then immediately released to give rise to a sharp peak in a NOx concentration profile. Here, CO was used as a gas feed to facilitate the desorption of NOx and SOx. In the absence of SO2 (a), the observed NO uptakes in each absorption steps were almost constant at 0.30 mmol g 1 (g-Al2O3) and 0.44 mmol g 1 (TiO2–ZrO2), indicating that stable NOx adsorption and desorption are possible without thermal degradation. However, when 100 ppm SO2 was contained in the gas feed, the total NOx uptake was gradually decreased with the progress of adsorption/ desorption cycles. Fig. 9 plots the relative NOx uptakes in each adsorption step that are normalized by the NOx uptakes in the first cycle. Although the amounts of NOx uptake decreased gradually with an increase of cycle number, almost constant uptake was obtained after six cycles. After this period, TiO2–ZrO2 lost only 13% of the first cycle NOx uptake, compared to more than 30% NOx uptake lost for g-Al2O3. These results indicated the higher stability of TiO2–ZrO2 to SO2 poisoning compared to g-Al2O3 in the thermal swing cyclic NOx adsorption/desorption processes.
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Fig. 9. Effect of SOx on sorptive removal of NOx in thermal swing cycle experiments. (*) g-Al2O3 and (~) TiO2–ZrO2 in adsorption/desorption cycle (a) in Figs. 7 and 8. (&) g-Al2O3 and (^) TiO2–ZrO2 in adsorption/desorption cycle (b) in Figs. 7 and 8.
4. Conclusion
Fig. 7. Thermal swing NOx adsorption–desorption cycles over g-Al2O3 in the absence (a) and presence (b) of SO2. Adsorption step: (a) 0.08% NO, 10% O2, He balance, (b) 0.08% NO, 100 ppm SO2, 10% O2, He balance at 25 8C for 20 min, W/F = 0.06 g s cm 3. Desorption step: (a), (b) 0.04% CO, He balance, at 600 8C for 10 min, W/F = 0.06 g s cm 3.
The following conclusions have been emerged from this study. The adsorptive NO uptake of (n)TiO2–(1 n)ZrO2 became maximum at n = 0.25–0.5, where a largest surface area was attained with an amorphous-like microporous structure. The present adsorbent was found to be effective not only for the oxidative adsorption of NO at ambient temperature, but also thermal NOx desorption at low temperatures (500 8C) in the presence of SO2. TiO2–ZrO2 exhibited stable temperatureswing NO adsorption/desorption cycles in a stream of 0.16% NO, 10% O2, 100 ppm SO2 and He balance. By contrast, gAl2O3 exhibited a significant deterioration by SO2, because strong chemisorption of sulfate species block the base sites from NOx adsorption. The tolerance of TiO2–ZrO2 to SO2 should be associated with weak base sites and instability of sulfate (SO4) species on the surface. Acknowledgement This study was financially supported by Grant-in aid for Scientific Research from Japan Society for Promotion of Science. References
Fig. 8. Thermal swing NOx adsorption–desorption cycles over TiO2–ZrO2 in the absence (a) and presence (b) of SO2. Adsorption step: (a) 0.08% NO, 10% O2, He balance and (b) 0.08% NO, 100 ppm SO2, 10% O2, He balance at 25 8C for 20 min, W/F = 0.06 g s cm 3. Desorption step: (a), (b) 0.04% CO, He balance, at 600 8C for 10 min, W/F = 0.06 g s cm 3.
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NO adsorption/desorption property of TiO2–ZrO2 having tolerance to SO2 poisoning Kazuhiro Ito, Shuichi Kakino, Keita Ikeue, Masato Machida * Department of Nano Science and Technology, Graduate School of Science and Technology, Kumamoto University, 2-39-1 Kurokami, Kumamoto 860-8555, Japan Received 20 June 2006; received in revised form 10 December 2006; accepted 13 February 2007 Available online 16 February 2007
Abstract The effect of SO2 on the adsorption and desorption properties of NO was studied over TiO2–ZrO2 and g-Al2O3. The microporous binary oxide, TiO2–ZrO2, exhibited a large amount of sorptive NOx uptake (ca. 1 mmol g 1) at 25 8C in the presence of O2 and SO2. The irreversible chemisorption as nitrate (NO3) is a main cause for the large NO uptake. Desorption of NOx from TiO2–ZrO2 occurred at lower temperature of ca.450 8C, compared to 500 8C required for that from g-Al2O3. A similar trend was observed for the SOx desorption, because SO2 adsorbed on gAl2O3 is a stable sulfate-type (SO4), whereas less stable sulfite (SO3) is dominant species on the surface of TiO2–ZrO2. The SOx desorption could further be facilitated in the presence of reducing agents like CO. The thermal swing adsorption/desorption cycle experiment demonstrated that TiO2–ZrO2 can be used as an NOx adsorbent having improved tolerance to SO2. The SO2 tolerance can be explained by the weak basicity of TiO2– ZrO2 and the low activity to oxidize SO2 to form sulfate species on the surface. # 2007 Elsevier B.V. All rights reserved. Keywords: TiO2–ZrO2; Nitric oxide; Oxidative adsorption; SOx-tolerance
1. Introduction Interest in the application of sorptive NOx removal to emission control has increased markedly in the past decades. One excellent example can be seen in the automotive NOx storage and reduction (NSR) catalyst [1–4]. Although threeway catalysts for the gasoline engines operating at stoichiometric fuel:air ratio are able to purify HC, CO and NOx simultaneously, it is difficult to abate NOx emission from diesel or lean-burn gasoline engines. To solve this problem, the NSR catalyst contains NOx storage components (Ba or other alkaline-earth metals) in the vicinity of noble metal catalysts (Pt) on a support oxide (g-Al2O3). NO in the exhaust is first oxidized to NO2 and stored as Ba–NO3 under lean-burn conditions. By temporarily turning to rich condition, stored NOx is then released and simultaneously reduced by CO and hydrocarbons to N2 on noble metal catalysts. One serious drawback of the NSR catalyst is the deactivation caused by SO2
* Corresponding author. Fax: +81 96 342 3651. E-mail address: [email protected] (M. Machida). 0926-3373/$ – see front matter # 2007 Elsevier B.V. All rights reserved. doi:10.1016/j.apcatb.2007.02.008
originating from fuel [5]. Acidic gases, NO2 and SO2, should compete for the same adsorption sites and, generally, SO2 adsorption is stronger. When the adsorbed SO2 yields sulfate, the NOx storage should totally be lost. Therefore, the development of SO2-tolerant NOx-sorbing materials is now strongly requested. Another important requirement for NOx storage components is to improve NOx removal at lower temperatures. Increasingly stringent regulations of NOx emissions require efficient sorbents working even near ambient temperature. Recently, several researchers challenged to develop the SO2tolerant NOx adsorbent by employing non-alkaline solid as the storage components. Eguchi et al. [6] have reported the oxidative NO adsorption onto MnOx–ZrO2 at 200 8C. Thanks to the weaker basicity, the thermal desorption of NOx takes place at low temperatures of 400 8C, compared to 600 8C requested for the Ba-based adsorbents. A similar NO adsorbability can be demonstrated for MnOx–CeO2 in our previous studies [7], but their tolerance to SO2 should not be enough [8]. Centi et al. [9] reported that the resistance to deactivation by SO2 could be promoted in hydrotalcite-based supported Pt–Cu catalysts. It is well-known that TiO2 is tolerant
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to SO2 poisoning. Hachisuka et al. [10] and Huang et al. [11] studied TiO2 as an NOx storage component. Because of instability of sulfate on TiO2, the storage components after exposure to SO2 can easily be regenerated. Unfortunately, however, weak basicity of TiO2 led to a much smaller NOx storage capacity than that of Ba-containing sorbents. Eguchi and coworkers [12] recently reported the effect of additives (Li, Na, K, Cs, Sr, Ba, and La) on the NO uptake by Pt/TiO2 in the presence of SO2. The addition of Li caused the uptake of NO as well as SO2, but adsorbed SO2 was released at the lowest temperature. The different behavior of these additives can be explained by weak basicity of Li. As was evident from these previous studies, it is very difficult to achieve both NOx storage and SOx-tolerance simultaneously. For the design of SOx-tolerant NOx-sorbents, the precise control of the basicity is required not only to increase NOx uptake but also to liberate the adsorbed SO2 at lowest possible temperatures. As a candidate of such materials, TiO2-based oxides are promising. We have previously studied TiO2–ZrO2 as a support of Pt catalyst for NO–H2–O2 reactions at low temperatures of 100 8C [13–15]. Due to a number of Lewis base sites formed in accordance with the TiO2/ZrO2 ratio, a large amount of NO (0.4 mmol g 1 at 90 8C) was found to be stored during catalytic reactions. This is indicative of another possible application of TiO2–ZrO2 as an NOx-sorbing material, which can be used in the presence of SO2 at low temperatures. The present work aims to study the sorptive NO removal and SO2-tolerance of the TiO2–ZrO2 system at ambient temperature. Adsorption properties of NO and SO2 onto TiO2–ZrO2, g-Al2O3, SiO2, TiO2 and ZrO2 were compared to study the effect of surface basicity and the chemical state of adsorbates. The adsorption/ desorption cycles were performed using the temperature swing technique in the presence of excess O2 and SO2. 2. Experimental 2.1. Preparation and characterization Binary oxides, (n)TiO2–(1 n)ZrO2 (0 n 1), were prepared by the coprecipitation method [13]. Aqueous ammonia was added dropwise to a solution of ZrO(NO3)2 (97.0%, Wako) and [(CH3)2CHO]4Ti (95.0%, Wako) with vigorous stirring. The obtained precipitate was washed several times with deionized distilled water, followed by drying for overnight at 100 8C. The resulting solid product was calcined at 450 8C for 5 h in air. The other NOx adsorbents, g-Al2O3 (JRC, ALO-4) and SiO2 (JRC, SIO-5) were supplied by the Catalysis Society of Japan. The crystal structure of prepared materials was identified by X-ray diffraction using Cu Ka radiation (Rigaku Multiflex). The BET surface area was measured by a nitrogen isotherms adsorption system at 196 8C (BEL Japan, Belsorp). 2.2. NOx adsorption and desorption Sorptive removal of NOx was evaluated in a conventional flow system at atmospheric pressure. A granular sample
(10–20 mesh, 0.05 g) was fixed in a quartz tube (4 mm i.d.) by packing quartz wool at both ends of the adsorbent bed. The adsorbents were pretreated in a stream of 10% O2 and He balance at 400 8C for 60 min to remove adsorbed gases. After subsequent cooling to 25 8C, gaseous mixtures of 0.16% NOx, 0–10% O2, and 0–50 ppm SO2 in He were fed to the sample at W/F = 0.06 g s cm 3. The NOx desorption from adsorbents was observed by means of temperature-programmed desorption (TPD) measurement. NOx-adsorbed samples were heated to 800 8C at the rate of 10 8C min 1 in a stream of 10% O2 and He balance (50 cm3 min 1). The effluent gases were analyzed by a chemiluminescence NOx analyzer (Shimadzu NOA-7000) and an on-line quadruple mass spectrometer (Pfeiffer, Omnistar). 2.3. SOx adsorption and desorption The pretreated sample was placed at 25 8C in a flowing gas mixture of 50–100 ppm SO2, 10% O2 in He supplied at W/ F = 0.06 g s cm 3 for 60 min. After the treatment at various temperatures, the uptake of SO2 was determined by use of an Xray fluorescence analyzer (Horiba MESA-500W). The treatment by SO2 was followed by heating at 400–800 8C for 60 min in the oxidizing atmosphere (10% O2 and He balance) or reducing atmosphere (0.04% CO and He balance) to obtain the desorption property of SO2. 2.4. FT-IR spectra of adsorbed NOx/SOx FT-IR spectra of NOx/SOx species adsorbed on TiO2–ZrO2 were recorded on a Jasco FT-IR-610 spectrometer. The sample was placed in a temperature-controllable diffuse reflectance reaction cell (Jasco DR600A), which was connected to a gas flow system and a vacuum line. The sample was first treated in a stream of 10% O2 and He balance at 400 8C and next exposed to gas mixtures of 0.04% NO or 0.08% SO2, 10–20% O2, and He balance at 25 8C. This was followed by spectra measurement in a flowing He. All spectra, thus obtained were transformed into absorption spectra by the use of Kubelka–Munk function and referenced to those taken just before NO/SO2 admission. 2.5. Temperature swing adsorption/desorption cycles Temperature swing NOx adsorption/desorption cycles in the presence or absence of SO2 were performed by switching two gas feeds between 0.08% NO, 10% O2, 0 or 100 ppm SO2 in He and 0.04% CO in He. First, an NO adsorption step was carried out in a stream of 0.08% NO, 10% O2, 0 or 100 ppm SO2 in He (W/F = 0.06 g s cm 3) for 20 min at 25 8C. This was followed by a regeneration step in a stream of 0.04% CO in He (W/ F = 0.06 g s cm 3) at 600 8C for 10 min. The end of a cycle was followed by the rapid cooling to ambient temperature, where the NO concentration in the gas feed was measured before going to a next cycle. A water-cooled infrared image furnace (ULVAC RHL-E25) was used for alternating the reaction temperature between adsorption and desorption steps.
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Table 2 BET surface area and sorptive NO/NO2 uptake of TiO2–ZrO2, g-Al2O3 and SiO2
3. Results and discussion 3.1. Adsorptive uptake of NO/SO2 Table 1 summarizes phases, BET surface areas and NO uptakes of binary oxide, (n)TiO2–(1 n)ZrO2 (0 n 1), after calcination at 450 8C. Although ZrO2 (n = 0) and TiO2 (n = 1.0) were composed of crystalline phases, monoclinic/ tetragonal fluorite-type and anatase, respectively, the mixed oxide (n = 0.25 and 0.5) yielded an amorphous phase. The two crystalline oxides (n = 0 and 1.0) exhibited low surface areas of <90 m2 g 1, but amorphous-like oxides (n = 0.25 and 0.5) showed larger surface area of approximately 200 m2 g 1. TEM observation suggested the amorphous product was metastable mixtures of fine particles (ca. 2 nm in diameter) of TiO2 and ZrO2 having a low crystallinity. The NO breakthrough curves were measured in a flowing gas mixture of 0.16% NO, 10% O2, He balanced at 25 8C. Because neither N2O nor N2 were detected in the effluent gas, abatement of NOx attributes to the adsorptive uptake by solids. As shown in Table 1, the NOx uptake increased with an increase of the surface area and the maximum amount of the uptake was observed at n = 0.25–0.5. The NO uptake normalized by the surface area also showed the maximum at n = 0.25–0.5, suggesting that the NO adsorption site was created by combining TiO2 and ZrO2. According to work done by Arata et al. [16,17], this can be explained by a noncrystalline porous structure having hetero-junctions at the boundary of the two different oxide components having different oxygen coordination numbers, i.e., 6 for Zr and 8 for Zr. Due to excess positive and negative charge formed at the boundary, Zr acts like a Lewis acid while Ti–O acts as a base that causes the oxidative adsorption site of NO [14,15]. Table 2 compares the NO/NO2 uptakes for TiO2–ZrO2, gAl2O3, and SiO2 at 25 8C. These three metal oxides showed similar surface area in the range 177–207 m2 g 1. Nevertheless, their uptakes of NO/NO2 were quite different depending on the surface property. SiO2 exhibited the smallest NO uptake (0.06 mmol g 1) because of the very weak basicity of the surface. By contrast, TiO2–ZrO2 as well as g-Al2O3 exhibited ca. 10–15 times larger NO uptakes. Larger uptakes for NO2 compared to NO can be explained by the fact that NO2 is condensable and may cause stronger interactions with the base site on the solid. The NO uptake of TiO2–ZrO2 decreased with decreasing O2 concentration from 1.06 mmol g 1 (10% O2) to 0.13 mmol g 1 (0% O2). As thermodynamics predicts, NO Table 1 BET surface area and sorptive NO uptake of (n)TiO2–(1
Sample
BET surface area (m2 g 1)
NO uptake (m2 g 1)a
NO2 uptake (m2 g 1)a
g-Al2O3 SiO2 TiO2–ZrO2
177 192 207
0.61 0.06 1.06
1.05 0.23 1.19
a
0.16% NO or NO2, 10% O2, He balance, W/F = 0.06 g s cm 3, 25 8C.
should react with O2 to form NO2 at ambient temperature, but, at low NO concentration as in the present case (0.16%), the NO oxidation is very slow in the gas phase. However, this is not the case on the surface of metal oxides, where NO adsorption in the presence of O2 takes place via smooth oxidation of NO to NO2, presumably because of interactions with Ti–O species on the base site. The formed NO2 further reacts with the basic oxygen site to produce more stable NO3 species as detected by infrared measurement (see Fig. 3). Fig. 1 shows the amount of SO2 adsorption as a function of adsorption temperature. The SO2 uptake at 25 8C was not consistent with that of NOx uptake shown in Tables 1 and 2. The largest SO2 uptake observed for g-Al2O3 at 25 8C, 0.12 mmol g 1, was almost equal to the total amount of SO2 supplied in the gas feed. The SO2 uptake decreased with an increase of adsorption temperature, but g-Al2O3 retained almost constant uptake of ca. 0.08 mmol g 1 at 200 8C. Interestingly, the SO2 uptake onto TiO2–ZrO2 at 25 8C was much smaller (0.05 mmol g 1) in contrast to the largest NO/ NO2 uptakes (Table 2). Above 400 8C, the SO2 uptake became negligible on TiO2–ZrO2. Generally, the base site on the solid surface was occupied by NOx as well as SOx. However, this is not the case for TiO2–ZrO2, probably because SOx adsorbed on TiO2–ZrO2 are unstable than those on g-Al2O3. The effect of coexisting SO2 on adsorptive NO uptake was next studied. Fig. 2 compares the NO/SO2 uptakes onto gAl2O3 and TiO2–ZrO2 in a flowing 0.16% NO, 10% O2 and He balance (a) in the absence and (b) presence of 50 ppm SO2 at 25 8C. The amount of SO2 uptake by g-Al2O3, 0.11 mmol g 1, agreed with the accumulation of almost all SO2 supplied in the gas feed. The NO adsorption was suppressed to a significant extent in a presence of SO2. In accordance with Fig. 1, TiO2– ZrO2 showed the smaller SO2 uptake (0.07 mmol g 1). In both
n)ZrO2 at 25 8C
n
Phase
BET surface area (m2 g 1)
NO uptake (mmol g 1)a
0 0.25 0.50 0.75 1.00
Monoclinic/tetragonal ZrO2 Amorphous Amorphous Anatase Anatase
85 201 207 179 58
0.08 1.14 1.06 0.74 0.15
a
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0.16% NO, 10% O2, He balance, W/F = 0.06 g s cm 3, 25 8C.
Fig. 1. SOx adsorption onto (n)TiO2–(1 n)ZrO2 and g-Al2O3 at different temperature. 50 ppm SO2, 10% O2, He balance, W/F = 0.06 g s cm 3, 60 min.
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Fig. 2. Coadsorption of NO and SO2 onto g-Al2O3 and TiO2–ZrO2 at 25 8C. (a) 0.16% NO, 10% O2, He balance, W/F = 0.06 g s cm 3 and (b) 0.16% NO, 50 ppm SO2, 10% O2, He balance, W/F = 0.06 g s cm 3. Total SO2 supplied: 0.12 mmol g 1.
cases, the amount of NO uptake decreased in the presence of SO2, whereas the sum of NO and SO2 uptake was kept constant. The result simply means that NO and SO2 compete to adsorb on the same surface site. The SO2/NO ratios on the surface were larger than those expected from their concentration in the gas phase, suggesting that SO2 should be bound to the surface more strongly than NO. Nevertheless, the NO adsorption on TiO2– ZrO2 is less sensitive to SO2. Fig. 3a shows the differential FT-IR spectra of NO adsorbed onto g-Al2O3 and TiO2–ZrO2, which are referenced to those taken just before NO admission. On TiO2–ZrO2, a set of sharp double bands at 1575 and 1280 cm 1, and a weak band at 1023 cm 1, which can be ascribed to NO stretching modes of bidentate nitrate (NO3), appeared. Because these peaks were intensified with an increase of O2 concentration, oxidative NO adsorption to form NO3 could be proved. The spectra of NO adsorbed on g-Al2O3 was almost same, but the peaks due to nitrate was less intense. Fig. 3b shows the differential spectra taken after exposure of g-Al2O3 and TiO2– ZrO2 to a mixture of 0.08% SO2, 20% O2 in He for 6 h at 25 8C. After exposing g-Al2O3 to the gas mixture, the adsorption band appeared at 1150 cm 1, which can be assigned to the S–O asymmetric stretching mode of sulfate (SO4). By contrast, TiO2–ZrO2 exhibits bands centered at 951 and 1016 cm 1, which are characteristic for asymmetric and symmetric stretching sulfite (SO3) vibration. The result suggested that SO2 is readily oxidized to the most stable species, SO4, over g-Al2O3, but less-oxidized species (SO3) are dominant on TiO2–ZrO2. Such different SOx species should be associated with the SO2 adsorption behavior (Fig. 1) as well as the tolerance to SO2 (Fig. 2) of the present system. The reason for the lack of ability to oxidize SO2 is not clear at the present stage, but it seems to be consistent with the stability of sulfate, i.e., the sulfates of Ti and Zr decompose at 450 8C, compared to more than 1000 8C requested for the decomposition of Al2(SO4)3.
Fig. 3. FT-IR spectra of (a) NOx adsorbed on g-Al2O3 and TiO2–ZrO2 after exposure to 0.04% NO, 10% O2, He balance at 25 8C for 1 h and (b) SOx adsorbed on g-Al2O3 and TiO2–ZrO2 after exposure to 0.08% SO2, 20% O2, He balance at 25 8C for 6 h.
3.2. Desorption of NO/SO2 The TPD profiles of NO after adsorption in the absence of SO2 are shown in Fig. 4. Although adsorption was carried out at ambient temperature, heating above 400 8C was required to complete the NOx desorption. This is indicative of the fact that NOx species are strongly chemisorbed as nitrate on the solid surface. Especially, the NOx desorption from g-Al2O3 and ZrO2 took place at the highest temperature (500 8C). The addition of TiO2 to ZrO2 shifted the desorption peak to lower temperatures and the minimal desorption temperature was attained for TiO2–ZrO2 (n = 0.5), which exhibited the largest NO uptake (Table 1). This result implies that the combination between TiO2 and ZrO2 increased the number of base site for NO adsorption, but the basicity of each site became rather weak. The thermal desorption of SOx can be estimated by measuring the residual SOx uptake after exposure to a stream of 100 ppm SO2, 10% O2, and He balance at 25 8C for 60 min and subsequent heating in a stream of 10% O2 and He balance at elevated temperatures for 60 min (Fig. 5). The initial SOx uptake was 0.18 mmol g 1 (g-Al2O3) and 0.11 mmol g 1 (TiO2–ZrO2), respectively. The thermal desorption of SOx from g-Al2O3 only was observed at 600 8C and did not complete even at 800 8C. The SOx desorption therefore requires
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Fig. 6. Residual SOx uptake after heating in a reducing gas at elevated temperatures. SO2 adsorption: 100 ppm SO2, 10% O2, He balance, at 25 8C for 60 min, W/F = 0.06 g s cm 3. Desorption condition: 0.04% CO, He balance at 400–800 8C for 60 min, W/F = 0.06 g s cm 3.
The thermal SOx desorption was evaluated in a stream of 0.04% CO and He balance as shown in Fig. 6. In such a reducing atmosphere, the SOx desorption is expected to be facilitated because the adsorbed SOx species should become less stable. The SOx desorption from TiO2–ZrO2 started at 400 8C and finished at 700 8C. The SOx desorption from gAl2O3 also took place at lower temperatures, but it was still incomplete at 800 8C. These results indicate that not only NOx but also SOx can be liberated from TiO2–ZrO2 at lower temperature than those from g-Al2O3. 3.3. Thermal swing NOx adsorption/desorption cycles Fig. 4. NOx-TPD profiles of g-Al2O3 and (n)TiO2–(1 n)ZrO2 after adsorption in 0.16% NO, 10% O2, He balance at 25 8C. 10% O2, He balance, heating rate 10 8C min 1.
temperatures higher than those for NOx desorption as shown in Fig. 4. This is indicative of the higher stability of adsorbed SOx (sulfate) compared to adsorbed NOx (nitrate) on the surface of g-Al2O3. On the contrary, the SOx desorption from TiO2–ZrO2 started at lower temperature of 400 8C and finished at 800 8C. Again, the result can be explained by the fact that the chemisorption of SO2 onto TiO2–ZrO2 as sulfite (SO3) is weak compared to sulfate (SO4) bound to g-Al2O3.
Fig. 5. Residual SOx uptake after heating in an oxidizing gas at elevated temperatures. SO2 adsorption: 100 ppm SO2, 10% O2, He balance, at 25 8C for 60 min, W/F = 0.06 g s cm 3. Desorption condition: 10% O2, He balance at 400–800 8C for 60 min, W/F = 0.06 g s cm 3.
The stability of NOx adsorption/desorption cycles for the TiO2–ZrO2 system was evaluated using thermal swing technique. Figs. 7 and 8 display the change of effluent NOx concentration during the thermal swing adsorption–desorption cycles over gAl2O3 and TiO2–ZrO2, respectively. The adsorption was carried out at 25 8C for 20 min in a flowing mixture of 0.08% NO, 10% O2, and He balance (a) in the absence of SO2 or (b) in the presence of 100 ppm SO2. At the end of each adsorption step, the temperature was raised to 600 8C, where a mixture of 0.04% CO and He balance was supplied for 10 min. The adsorbed NOx species were then immediately released to give rise to a sharp peak in a NOx concentration profile. Here, CO was used as a gas feed to facilitate the desorption of NOx and SOx. In the absence of SO2 (a), the observed NO uptakes in each absorption steps were almost constant at 0.30 mmol g 1 (g-Al2O3) and 0.44 mmol g 1 (TiO2–ZrO2), indicating that stable NOx adsorption and desorption are possible without thermal degradation. However, when 100 ppm SO2 was contained in the gas feed, the total NOx uptake was gradually decreased with the progress of adsorption/ desorption cycles. Fig. 9 plots the relative NOx uptakes in each adsorption step that are normalized by the NOx uptakes in the first cycle. Although the amounts of NOx uptake decreased gradually with an increase of cycle number, almost constant uptake was obtained after six cycles. After this period, TiO2–ZrO2 lost only 13% of the first cycle NOx uptake, compared to more than 30% NOx uptake lost for g-Al2O3. These results indicated the higher stability of TiO2–ZrO2 to SO2 poisoning compared to g-Al2O3 in the thermal swing cyclic NOx adsorption/desorption processes.
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Fig. 9. Effect of SOx on sorptive removal of NOx in thermal swing cycle experiments. (*) g-Al2O3 and (~) TiO2–ZrO2 in adsorption/desorption cycle (a) in Figs. 7 and 8. (&) g-Al2O3 and (^) TiO2–ZrO2 in adsorption/desorption cycle (b) in Figs. 7 and 8.
4. Conclusion
Fig. 7. Thermal swing NOx adsorption–desorption cycles over g-Al2O3 in the absence (a) and presence (b) of SO2. Adsorption step: (a) 0.08% NO, 10% O2, He balance, (b) 0.08% NO, 100 ppm SO2, 10% O2, He balance at 25 8C for 20 min, W/F = 0.06 g s cm 3. Desorption step: (a), (b) 0.04% CO, He balance, at 600 8C for 10 min, W/F = 0.06 g s cm 3.
The following conclusions have been emerged from this study. The adsorptive NO uptake of (n)TiO2–(1 n)ZrO2 became maximum at n = 0.25–0.5, where a largest surface area was attained with an amorphous-like microporous structure. The present adsorbent was found to be effective not only for the oxidative adsorption of NO at ambient temperature, but also thermal NOx desorption at low temperatures (500 8C) in the presence of SO2. TiO2–ZrO2 exhibited stable temperatureswing NO adsorption/desorption cycles in a stream of 0.16% NO, 10% O2, 100 ppm SO2 and He balance. By contrast, gAl2O3 exhibited a significant deterioration by SO2, because strong chemisorption of sulfate species block the base sites from NOx adsorption. The tolerance of TiO2–ZrO2 to SO2 should be associated with weak base sites and instability of sulfate (SO4) species on the surface. Acknowledgement This study was financially supported by Grant-in aid for Scientific Research from Japan Society for Promotion of Science. References
Fig. 8. Thermal swing NOx adsorption–desorption cycles over TiO2–ZrO2 in the absence (a) and presence (b) of SO2. Adsorption step: (a) 0.08% NO, 10% O2, He balance and (b) 0.08% NO, 100 ppm SO2, 10% O2, He balance at 25 8C for 20 min, W/F = 0.06 g s cm 3. Desorption step: (a), (b) 0.04% CO, He balance, at 600 8C for 10 min, W/F = 0.06 g s cm 3.
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