yannick mesoporous absorbtie

8
Mesoporous materials for the removal of SO 2 from gas streams Yannick Mathieu a , Michel Soulard a , Joël Patarin a, *, Michel Molière b a Equipe Matériaux à Porosité Contrôlée, Institut de Science des Matériaux de Mulhouse, LRC 7228 CNRS, Université de Haute Alsace, ENSCMu 3, rue Alfred Werner, 68093 Mulhouse Cedex, France b GE Energy, 20 avenue du Marechal Juin, 90007 Belfort Cedex, France abstract article info Article history: Received 17 August 2011 Received in revised form 23 November 2011 Accepted 11 February 2012 Available online 9 March 2012 Keywords: SO 2 adsorption Mesoporous MCM-41 CuO CeO 2 LiCl A MCM-41 sorbent containing two metal oxides (CuO, CeO 2 ) and an alkaline additive (LiCl) was prepared to remove SO 2 from gas streams at 673 K. Two impregnation methods i.e. the template-ion exchange and the two solvents impregnation method were employed yielding a high dispersed state of the metal oxides inside the porosity of the MCM-41 sorbent. The MCM-41 sorbent was found to have a high adsorption capacity of 130 mg SO 2 /g at 673 K. During the adsorption process, cerium oxide partially oxidized SO 2 into SO 3 which was further chemically adsorbed on active sites (CuO, Li 2 O). The formation of CuSO 4 and Li 2 SO 4 was observed by X-ray diffraction after the reaction. © 2012 Elsevier B.V. All rights reserved. 1. Introduction Sulfur oxide (SO x =SO 2 +SO 3 ) emissions stemming from human activities including fossil fuel combustion in thermal power plants, petroleum renery or on-road vehicles are major source of atmo- spheric pollution leading to acid rain, smog and health problems. Therefore, the development of new sorbents for the removal of SO x is of primary importance. SO x sorbents are numerous and can be clas- sied in different categories such as porous silica-based oxides, single oxides materials (CaO, MgO, etc.), alumina-supported oxides (mainly CuO/γ-Al 2 O 3 ), mixed oxides derived from spinels or hydrotalcites and oxides supported on carbonaceous materials. Among these sorbents, porous silica-based materials including zeolites, mesoporous mate- rials or clay minerals are interesting sorbents due to their high specif- ic surface area and fair thermal stability. Numerous zeolites including silicalite-1 [1,2],Y [3,4], clinoptilolite [5], mordenite or ZSM-5 [6,7] and clay minerals such as bentonite [8] or montmorillonite [9] have been investigated as potential SO 2 sor- bents. These studies showed that porous silica materials are promising sorbents for the removal of SO 2 at low temperatures within physical sorption processes. However, increasing the adsorption temperature often results in a strong decrease of the SO 2 removal capacity [1,2,4]. Thus, as the ue gas temperature is generally ranging from 523 to 873 K, it is necessary to combine the porous silica support with metal oxides to chemically adsorb SO 2 . Copper oxide was found to be one of the most active metal oxides to capture SO 2 or SO 3 as copper sulfates mainly with CuO/γ-Al 2 O 3 sorbents [10]. However, CuO/γ-Al 2 O 3 sor- bents have only textural porosity and moderate specic surface area (b 200 m 2 /g) in contrast to porous silica based materials. Furthermore, the γ-Al 2 O 3 support is not inert as aluminum sulfate species are often observed by X-ray diffraction after bulk sulfation [11]. Numerous studies were devoted to the use of copper oxide dis- persed in γ-Al 2 O 3 support as SO x sorbent [1214]. The SO x adsorption on CuO/γ-Al 2 O 3 leads to the formation of CuSO 4 and Al 2 (SO 4 ) 3 . As it was showed by Centi et al., the sulfation reaction involves the decom- position of CuSO 4 leading to the formation of SO 3 which is further chemically adsorbed on Al sites [10]. Thus, an oxidation catalyst such as cerium which promotes the oxidation of SO 2 into SO 3 can sig- nicantly enhance the sulfation process. A few studies mentioned a signicant increase of the SO x adsorption capacity for sorbents pre- pared in the presence of cerium oxide [15,16]. For instance, Palomares et al. [15] showed that Ce-doped mixed oxides derived from hydrotal- cites exhibits higher SO 2 removal efciency than the corresponding mixed oxides alone. Wey et al. [16] showed that activated carbon - bers (ACs) containing CuO and CeO 2 have improved performance in SO 2 adsorption in contrast to ACs that do not contain metal oxides. Additions of alkaline metals are known to increase the SO 2 adsorp- tion capacity. As it has been showed by Jeong et al. [17], the adsorp- tion capacity of CuO/γ-Al 2 O 3 materials is signicantly enhanced by alkaline salts, in particular LiCl. Two effects were clearly observed by the authors. First, the required temperature for bulk sulfation is down to about 353 K and the melting of LiCl increases the ionic mo- bility and reagents diffusion. The strategy followed in this study was to develop a sorbent for the removal of SO x which combines an inert silica mesoporous sup- port (MCM-41) having a high specic surface area (~ 1000 m 2 /g) Fuel Processing Technology 99 (2012) 3542 Corresponding author. Tel.: + 33 3 89 33 68 80. E-mail address: [email protected] (J. Patarin). 0378-3820/$ see front matter © 2012 Elsevier B.V. All rights reserved. doi:10.1016/j.fuproc.2012.02.005 Contents lists available at SciVerse ScienceDirect Fuel Processing Technology journal homepage: www.elsevier.com/locate/fuproc

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Page 1: YannicK Mesoporous absorbtie

Fuel Processing Technology 99 (2012) 35–42

Contents lists available at SciVerse ScienceDirect

Fuel Processing Technology

j ourna l homepage: www.e lsev ie r .com/ locate / fuproc

Mesoporous materials for the removal of SO2 from gas streams

Yannick Mathieu a, Michel Soulard a, Joël Patarin a,*, Michel Molière b

a Equipe Matériaux à Porosité Contrôlée, Institut de Science des Matériaux de Mulhouse, LRC 7228 CNRS, Université de Haute Alsace, ENSCMu – 3, rue Alfred Werner,68093 Mulhouse Cedex, Franceb GE Energy, 20 avenue du Marechal Juin, 90007 Belfort Cedex, France

⁎ Corresponding author. Tel.: +33 3 89 33 68 80.E-mail address: [email protected] (J. Patarin).

0378-3820/$ – see front matter © 2012 Elsevier B.V. Alldoi:10.1016/j.fuproc.2012.02.005

a b s t r a c t

a r t i c l e i n f o

Article history:Received 17 August 2011Received in revised form 23 November 2011Accepted 11 February 2012Available online 9 March 2012

Keywords:SO2 adsorptionMesoporous MCM-41CuOCeO2

LiCl

A MCM-41 sorbent containing two metal oxides (CuO, CeO2) and an alkaline additive (LiCl) was prepared toremove SO2 from gas streams at 673 K. Two impregnation methods i.e. the template-ion exchange and thetwo solvents impregnation method were employed yielding a high dispersed state of the metal oxides insidethe porosity of the MCM-41 sorbent. The MCM-41 sorbent was found to have a high adsorption capacity of130 mg SO2/g at 673 K. During the adsorption process, cerium oxide partially oxidized SO2 into SO3 whichwas further chemically adsorbed on active sites (CuO, Li2O). The formation of CuSO4 and Li2SO4 was observedby X-ray diffraction after the reaction.

© 2012 Elsevier B.V. All rights reserved.

1. Introduction

Sulfur oxide (SOx=SO2+SO3) emissions stemming from humanactivities including fossil fuel combustion in thermal power plants,petroleum refinery or on-road vehicles are major source of atmo-spheric pollution leading to acid rain, smog and health problems.Therefore, the development of new sorbents for the removal of SOx

is of primary importance. SOx sorbents are numerous and can be clas-sified in different categories such as porous silica-based oxides, singleoxides materials (CaO, MgO, etc.), alumina-supported oxides (mainlyCuO/γ-Al2O3), mixed oxides derived from spinels or hydrotalcites andoxides supported on carbonaceous materials. Among these sorbents,porous silica-based materials including zeolites, mesoporous mate-rials or clay minerals are interesting sorbents due to their high specif-ic surface area and fair thermal stability.

Numerous zeolites including silicalite-1 [1,2], Y [3,4], clinoptilolite[5], mordenite or ZSM-5 [6,7] and clay minerals such as bentonite [8]or montmorillonite [9] have been investigated as potential SO2 sor-bents. These studies showed that porous silica materials are promisingsorbents for the removal of SO2 at low temperatures within physicalsorption processes. However, increasing the adsorption temperatureoften results in a strong decrease of the SO2 removal capacity [1,2,4].Thus, as the flue gas temperature is generally ranging from 523 to873 K, it is necessary to combine the porous silica support with metaloxides to chemically adsorb SO2. Copper oxide was found to be one ofthe most active metal oxides to capture SO2 or SO3 as copper sulfates

rights reserved.

mainly with CuO/γ-Al2O3 sorbents [10]. However, CuO/γ-Al2O3 sor-bents have only textural porosity and moderate specific surface area(b200 m2/g) in contrast to porous silica based materials. Furthermore,the γ-Al2O3 support is not inert as aluminum sulfate species are oftenobserved by X-ray diffraction after bulk sulfation [11].

Numerous studies were devoted to the use of copper oxide dis-persed in γ-Al2O3 support as SOx sorbent [12–14]. The SOx adsorptionon CuO/γ-Al2O3 leads to the formation of CuSO4 and Al2(SO4)3. As itwas showed by Centi et al., the sulfation reaction involves the decom-position of CuSO4 leading to the formation of SO3 which is furtherchemically adsorbed on Al sites [10]. Thus, an oxidation catalystsuch as cerium which promotes the oxidation of SO2 into SO3 can sig-nificantly enhance the sulfation process. A few studies mentioned asignificant increase of the SOx adsorption capacity for sorbents pre-pared in the presence of cerium oxide [15,16]. For instance, Palomareset al. [15] showed that Ce-doped mixed oxides derived from hydrotal-cites exhibits higher SO2 removal efficiency than the correspondingmixed oxides alone. Wey et al. [16] showed that activated carbon fi-bers (ACs) containing CuO and CeO2 have improved performance inSO2 adsorption in contrast to ACs that do not contain metal oxides.

Additions of alkaline metals are known to increase the SO2 adsorp-tion capacity. As it has been showed by Jeong et al. [17], the adsorp-tion capacity of CuO/γ-Al2O3 materials is significantly enhanced byalkaline salts, in particular LiCl. Two effects were clearly observedby the authors. First, the required temperature for bulk sulfation isdown to about 353 K and the melting of LiCl increases the ionic mo-bility and reagents diffusion.

The strategy followed in this study was to develop a sorbent forthe removal of SOx which combines an inert silica mesoporous sup-port (MCM-41) having a high specific surface area (~1000 m2/g)

Page 2: YannicK Mesoporous absorbtie

36 Y. Mathieu et al. / Fuel Processing Technology 99 (2012) 35–42

with two metal oxides (CuO, CeO2) and an alkaline additive (LiCl).Metal loaded MCM-41 sorbents for the removal of SOx can be pre-pared following two methods, i.e. the template-ion exchange (TIE)and the two solvents impregnation (TSI) methods, respectively. TheTIE is a recognized method to provide a high dispersed state ofmetal nanoparticles inside the porosity of mesoporous materials,which is useful, for instance to prepare Cr-, Ni- or Co-MCM-41[18–20]. In this method, a cationic exchange occurs in which theC16TMA+ surfactant molecules located inside the mesopores of theas-synthesized MCM-41 are substituted by the metallic cations inthe precursor solution. In the TSI method, a volume of a highly con-centrated aqueous solution of the metal salt (typically 2 or 3 M) isset equal to the pore volume of the silica support previously deter-mined by N2 sorption. The calcined and dehydrated silica support ispreviously dispersed in an organic solvent poorly miscible withwater before adding the inorganic precursors. During the dryingstep performed at room temperature, the inorganic precursors fillthe mesopores by capillarity. This method was proposed by Davidsonet al. to incorporate metal oxides such as β-MnO2 [21], Co3O4 [22] orγ-Fe2O3 [23] inside the mesopores of SBA-15 materials.

In our study, the TIE and TSI methods were used to prepare theMCM-41 sorbent containing the metal oxides (CeO2, CuO) and the al-kaline additive (LiCl). The samples were fully characterized by X-raydiffraction (XRD), transmission electron microscopy (TEM), thermalanalysis and nitrogen adsorption. A main part of this work was devot-ed to the study of SO2 adsorption of the different samples at a highspace velocity (SV) of 25,000 h−1, in particular the adsorption capac-ity (milligram of SO2 captured by gram of sorbent). The sulfated sam-ples were also characterized by XRD and temperature programmeddesorption (TPD).

2. Experimental

2.1. Synthesis of MCM-41

The synthesis procedure of MCM-41 includes the dissolution of3.35 g of C16TMABr (hexadecyltrimethylammonium bromide, Al-drich, ≥99%) in 70 g of distilled water at 313 K (=solution A) andthe dissolution of 10.50 g of sodium silicate (Riedel de Haën,27 wt.% SiO2, 10 wt.% NaOH) in 11 g of distilled water at room tem-perature (=solution B). After cooling solution A at room tempera-ture, solutions A and B were mixed leading to the formation of awhite precipitate and the resulting mixture was mechanically stirredat room temperature for 5 min. An aqueous solution of hydrogenchloride (1 M) was added slowly to decrease the pH up to 8.5 andthe mixture was stirred for 5 more minutes. The resulting mixturewas then transferred in a polypropylene bottle and heated under stat-ic conditions at 363 K for 24 h. The solid obtained after the hydrother-mal treatment was recovered by filtration, washed three times with500 ml of distilled water and dried at 343 K for one night.

2.2. Preparation of CeO2-MCM-41 (Template Ion Exchange)

The incorporation of cerium was performed following the TIEmethod. A mass of 2.70 g of cerium nitrate hexahydrate (Aldrich,99.9%) was dissolved in a solution containing 60 ml of absolute etha-nol and 300 ml of distilled water. After dissolution, 3 g of as-synthesized MCM-41 was added and the mixture was stirred at333 K for 24 h. The Ce-MCM-41 material was then recovered by filtra-tion, washed three times with 500 ml of distilled water and dried atroom temperature for 24 h. The Ce-MCM-41 material was further cal-cined at 873 K for 6 h (heating rate of 1.7 K/min) yielding the forma-tion of CeO2-MCM-41. The color of the CeO2-MCM-41 sample wasslightly yellow.

2.3. Preparation of the MCM-41 sorbent (Two Solvents Impregnation)

The incorporation of copper and lithium chloride was performedby the TSI method. The anhydrous CeO2-MCM-41 (1.70 g) was dis-persed in 100 ml of n-hexane and stirred for 20 min. Then, 700 μl ofa 3 M aqueous solution of copper nitrate (Carlo Erba, 99.5%) wasadded and the mixture was stirred for 10 min. The powder wasthen recovered by filtration, crushed in a mortar for 10 min anddried 24 h at ambient temperature. The material was further calcinedat 873 K for 6 h (heating rate of 1.7 K/min) leading to the formation ofCuO/CeO2-MCM-41. A pronounced gray color was observed. A similarprocedure was applied for the incorporation of lithium chloride (Al-drich, ≥99%). A mass of 1.80 g of CuO/CeO2-MCM-41 was dispersedin 100 ml of n-hexane and stirred for 20 min. Then, 680 μl of an aque-ous solution of lithium chloride 5 M was added and the mixture wasstirred for 10 min. Finally, the powder was recovered by filtration,crushed in a mortar for 10 min and dried 24 h at room temperatureleading to the formation of the sorbent for the removal of SOx. TheMCM-41 sorbent containing CeO2, CuO and LiCl will be designatedin this work as Li-doped MCM-41 sorbent.

2.4. Characterization

X-ray powder diffraction was performed in a reflection mode on aPhilips X'Pert diffractometer (Cu Kα radiation, λ=1.54 Å) equippedwith a X'Celerator detector. Diffractograms were recorded with a 2θscan range of 3–70° with a step of 0.02° per 20 s at 35 kV tube voltageand 55 mA tube current. A glass support was used to record the XRDat low angles. Thermal (TG/DTA) analyses were performed under airon a Setaram Labsys thermoanalyzer with a heating rate of 5 K/minup to 1273 K. Chemical analyses were performed on the MCM-41 sor-bent using an X-ray fluorescence apparatus (PHILIPS MagiX). Nitro-gen adsorption measurements were carried out on a MicromeriticsTristar surface area analyzer after outgassing the samples for 15 h at423 K. Specific surface areas and pore size distributions were deter-mined by the BET and BJH methods, respectively [24]. The adsorptionbranch was used for pore size calculations. Transmission electron mi-croscopy images were collected on a Philips CM200 microscopeequipped with a LaB6 filament. The accelerating voltage was 200 kV.Samples were previously dispersed in water and sonicated for10 min. Then, five to ten drops of the solution were deposited ontoCu grids coated with a thin (5 nm thickness) holey carbon supportfilm.

The SO2 adsorption tests were performed on the metal oxides con-taining MCM-41 samples. The materials were previously pelletized,crushed and sized to obtain grains with a size ranging from 250 to450 μm. The experimental set-up was carried out using a fixed-bedflow quartz reactor (16.5 mm inner diameter) designed to allow ex-periments over a wide range of space velocities. The experimentwas performed at 673 K using a gas mixture of 250 ppm SO2, 10 v.%O2 and nitrogen as the carrier gas. The gasses were preheated beforeentering the reactor. The gas flow rate was 53 Nl/h and the mass ofthe material deposited over the glass-frit was 1.0 g yielding a SV of25,000 h−1. The SO2 outlet concentration was continuously moni-tored by a Rosemount NGA 2000 analyser based on UV absorption(range 0–1000 ppm). Before the SO2 experiment, in order to get adry sample, the material was heated from 298 to 673 K at a heatingrate of 10 K/min under pure N2 (gas flow rate of 100 Nl/h) and thetemperature was maintained at 673 K for 1 h. The adsorption capacity(Cads) was calculated by integration of the breakthrough curve via thefollowing formula:

Cads ¼ ∫ CSO2in −CSO2

outð Þdtm where Cin

SO2and CoutSO2are the inlet and outlet con-

centrations of SO2 and m is the mass of the hydrated sample. Theamount of water in the hydrated sample was carefully determined

Page 3: YannicK Mesoporous absorbtie

Table 1Structural and textural properties of the calcined MCM-41 and metal oxides containingMCM-41 samples.

37Y. Mathieu et al. / Fuel Processing Technology 99 (2012) 35–42

by TG measurements. The temperature programmed desorption wasperformed under pure N2 by heating the sulfated samples fromroom temperature to about 1173 K at a heating rate of 10 K/min.

Sample 2θ (100)(°)

d100

(nm)a0a

(nm)SBET(m2/g)

Porevolumeb

(cm3/g)

Poresizeb

(nm)

MCM-41 2.41 3.66 4.23 1020 0.87 2.4–2.5CeO2-MCM-41 2.45 3.60 4.15 915 0.66 2.0–2.2CuO/CeO2-MCM-41 2.55 3.45 3.98 465 0.29 2.0–2.1Li-doped MCM-41 sorbent 2.57 3.44 3.97 453 0.28 1.9–2.1

a Wall thickness calculated as a0=2×d100/ffiffiffi

3p

.b Pore-size distributions and pore volumes were determined from N2 adsorption iso-

therms at 77 K.

3. Results and discussion

3.1. Characterization of the sorbent

The calcined MCM-41 (reference) and metal oxides loaded MCM-41 samples were first characterized by low angle X-ray diffraction.The XRD pattern of the calcined MCM-41 is a characteristic of a wellordered hexagonal mesoporous structure with the main diffractionpeak (100) at 2.4° 2θ and other diffraction peaks (110), (200) and(210) having lower intensities between 4 and 6.5° 2θ (Fig. 1a). The in-tensities of the (100) reflection is strongly decreased for the metal ox-ides containing MCM-41 samples (Fig. 1b–d). In addition, the otherreflections (110), (200) and (210) disappear suggesting a loss oflong range crystallographic order. The formation of metal oxide parti-cles inside the mesopores after calcination of the Ce-MCM-41 and Cu/CeO2-MCM-41 samples may also explain this effect. The interplanarspacing d100 and the hexagonal lattice parameter a0 of the calcinedMCM-41 and metal oxides containing MCM-41 samples are listed inTable 1. It is observed that, the diffraction peak (100) is shifted from2.41 (d100=3.66 nm) to 2.57° 2θ (d100=3.44 nm) leading to a de-crease of the unit cell size from 4.23 to 3.97 nm. Such a decreasemay be explained by the successive calcination steps. The X-ray dif-fraction patterns at wide angles of the metal oxides containingMCM-41 samples are shown in Fig. 2. After calcination of the Ce-MCM-41 material, large diffraction peaks suggesting small crystallitesize and corresponding to the cerium oxide are observed at 28.4,47.5 and 56.8° 2θ (Fig. 2a). The Scherrer equation was used to calcu-late the size of the CeO2 crystallites using the diffraction peaks (111),(220) and (311), respectively (Table 2). The calculated values areconsistent with CeO2 particles located inside MCM-41 mesopores.Other sharp peaks are attributed to copper oxide which is formedduring the calcination of the Cu/CeO2-MCM-41 sample (Fig. 2b). Atypical X-ray diffraction pattern of the Li-doped MCM-41 sorbent isshown in Fig. 2c. After incorporation of lithium chloride, diffractionpeaks corresponding to copper oxide are less intense in the Li-doped MCM-41 sorbent than in the CuO/CeO2-MCM-41 sample indi-cating that chloride anions from LiCl reacted with copper oxide yield-ing the formation of a hydrated copper chloride which was identifiedas Cu2(OH)3Cl (atacamite). It is also important to note that no lithiumchloride peaks are observed by XRD.

2 4 6 8 2θ(CuKα)

Inte

nsity

(a.

u.)

(100)

(110) (200) (210)

(b)

(c)

(d)

(a)

Fig. 1. Low angle XRD patterns of calcined (a) MCM-41, (b) CeO2-MCM-41, (c) CuO/CeO2-MCM-41 and (d) Li-doped MCM-41 sorbent.

Further, the CeO2-MCM-41 sample and the Li-doped MCM-41 sor-bent were studied by TEM. The two samples show an ordered meso-porous structure. For the CeO2-MCM-41 sample, CeO2 nanoparticlesof highest electron density appear as black areas on the silica grains(Fig. 3a and b). As it can be seen, empty and filled mesopores withCeO2 nanoparticles are observed. Furthermore, no isolated or externalCeO2 nanoparticles attached onMCM-41 grains are evidenced. For theLi-dopedMCM-41 sorbent, filled and empty mesopores are also foundto coexist (Fig. 3c and d). No significant modifications of the meso-porous structure are observed.

The nitrogen adsorption/desorption isotherm of the calcinedMCM-41 is a typical type-IV isotherm with a vertical step at a relativepressure around P/P0=0.30 which is characteristic of capillary con-densation in mesopores having a diameter lower than 35 Å (Fig. 4a)[25]. A narrow and uniform pore size distribution around 24–25 Å isobserved (Fig. 5a). After the loading of CeO2, the shape of the iso-therm is slightly modified (Fig. 4b). However, a less pronounced ver-tical step slightly shifted toward lower P/P0 values (0.25–0.30) isobserved indicating a decrease of the mesoporous organization. A de-crease of the porous volume and the specific surface area is also ob-served. The corresponding BJH pore-size analysis shows a pore sizedistribution around 20–22 Å (Fig. 5b). For the CuO/CeO2-MCM-41and the Li-doped MCM-41 sorbent, no inflection point is visible (Fig.4c–d). The specific surface area and the pore volume are found to de-crease from 1020 to 453 m2/g and from 0.87 to 0.28 cm3/g, respec-tively throughout the preparation process (Table 1). As it wasshown by TEM observations, no metal oxide particles were found toplug the mesopore's entrance. Thus, the decrease of specific surfacearea and pore volume is attributed to the filling of the pores bymetal oxide particles (CeO2 and CuO) and atacamite. Accordingly,broad pore size distributions centered around 20 Å, characteristicsof a secondary microporosity, are observed for the CuO/CeO2-MCM-41 and the Li-doped MCM-41 sorbent (Fig. 5c–d).

The as-synthesized MCM-41, the Ce-MCM-41 sample and the Li-doped MCM-41 sorbent were characterized by TG/DTA. For the as-

10 20 30 40 50 60

*: CuO+: Cu2(OH)3Cl : CeO2

Inte

nsity

(a.u

.)

(a)

(b)

(c)

(111)

(220) (311)

2θ(CuKα)

Fig. 2. XRD patterns of (a) CeO2-MCM-41, (b) CuO/CeO2-MCM-41 and (c) Li-dopedMCM-41.

Page 4: YannicK Mesoporous absorbtie

Table 2Crystallite size L (Å) in different directions evaluated from the width at half-maximumintensity of the (111), (220) and (311) diffraction peaks of the CeO2-MCM-41 sample.

Diffraction peak Position (°2θ) L (Å)

(111) 28.2 25.4(220) 47.2 24.3(311) 56.3 24.5

0.1 0.2 0.3 0.4 0.5 0.6 0.7 0.8 0.9 100

100

200

300

400

500

600(a)

(b)

(c)

(d)

P/P0

Vad

s (c

m3 .

g-1 S

TP

)

Fig 4. Nitrogen adsorption–desorption isotherms for the calcined (a) MCM-41,(b) CeO2-MCM-41, (c) CuO/CeO2-MCM-41 and (d) Li-doped MCM-41 sorbent.

38 Y. Mathieu et al. / Fuel Processing Technology 99 (2012) 35–42

synthesized MCM-41, two weight losses can be seen (Fig. 6a). Thefirst one of about 4.7 wt.% occurring between 298 and 423 K iscoupled with a weak endothermic signal which is attributed to theelimination of physisorbed water. The second one (~37 wt.%) takesplace in the range 423–773 K and corresponds to the C16TMA+ de-composition and is associated with two exothermic peaks around573 and 773 K, respectively. A larger loss (~17.7 wt.%) of physisorbedwater is seen for the Ce-MCM-41 after the drying and rehydration ofthe sample elaborated via the TIE procedure (Fig. 6b). The secondweight loss (~1.1 wt.%) coupled with a weak exothermic peak is at-tributed to residual C16TMA+. The residual C16TMA+ comes from anon complete exchange of the C16TMA+ located inside the mesoporesby the Ce3+. Strydom and Van Vuuren showed that the decomposi-tion of cerium nitrate leads to the formation of CeO2 in the tempera-ture range 523–573 K [26]. However, it is difficult to evaluate exactlythe temperature of the CeO2 formation in the MCM-41 since the de-composition of residual C16TMA+ occurs in the same temperaturerange (473–573 K). For the Li-doped MCM-41 sorbent, a loss of phy-sisorbed water about 8.1 wt.% associated with an endothermic peakoccurs between 323 and 513 K (Fig. 6c). As it was showed by Sharkeyand Lewin, atacamite was found to decompose in two steps includingdehydration (553–563 K) and the loss of halogen (713–723 K)according to the following reactions [27]:

2Cu2ðOHÞ3Cl→CuO·CuCl2 þ 2CuO þ 3H2O ð1Þ

Fig. 3. TEM micrographs of (a, b) CeO2-MCM-41 s

CuO·CuCl2 þ 1

/

2O2→2CuO þ Cl2 ð2Þ

Therefore, the broad endothermic peak observed around 593 Kshould correspond to the loss of water coming from atacamite. Theinterpretation of the exothermic signals above 673 K is more difficult.A XRD performed on the calcined sorbent at 923 K revealed the pres-ence of small amounts of quartz and lithium silicate (Li2SiO3). Thus,the exothermic signals visible above 673 K might correspond to theloss of halogen and/or the formation of quartz and lithium silicatedue to the loss of order of the MCM-41 sorbent but it is difficult toclearly assign them. Moreover, no sharp endothermic peak corre-sponding to the melting of LiCl is observed around 873–883 K [28].It is worth mentioning that the structure of a classical MCM-41 with-out any metal oxides is thermally stable up to 1073 K.

Table 3 summarizes the chemical analyses for the Li-doped MCM-41 sorbent. The Si/Ce molar ratio is found to strongly increase from4.8 (TIE procedure) to 35 as cerium nitrate is introduced in large ex-cess during the TIE procedure. The mother liquid recovered by

ample and (c, d) Li-doped MCM-41 sorbent.

Page 5: YannicK Mesoporous absorbtie

-5

0

5

10

15

20

-5

0

5

10

15

20

Exo

TG

DTA

ExoTG

DTA

350 450 550 650 750 850 950 1050Temperature(K)

350 450 550 650 750 850 950 1050Temperature(K)

350 450 550 650 750 850 950 1050Temperature(K)

ExoTG

DTA

-5

0

5

10

15

20

25

30

35

40

45

Wei

ght l

oss(

%)

Wei

ght l

oss(

%)

Wei

ght l

oss(

%)

b

c

a

Fig. 6. TG/DTA curves of (a) the as-synthesized MCM-41, (b) Ce-MCM-41 sample and(c) Li-doped MCM-41 sorbent.

Table 3Chemical composition of the Li-doped MCM-41 sorbent.

Chemical composition (in wt.%) Weight loss %from TGb

Molar ratio

Si/Ce Si/Cu Si/Li

Si O Cu Ce Lia Cl 18.3 35 11.5 11.336.5 44.6 7.2 5.2 0.8 5.7

a Determined by ICP (inductively coupled plasma).b Under air up to 1073 K.

Fig. 5. BJH pore size curves for (a) the calcined MCM-41, (b) CeO2-MCM-41, (c) CuO/CeO2-MCM-41 and (d) Li-doped MCM-41 sorbent.

39Y. Mathieu et al. / Fuel Processing Technology 99 (2012) 35–42

filtration of the Ce-MCM-41 sample after the TIE procedure was ana-lyzed by 1H liquid NMR spectroscopy (spectra not shown here) andconfirmed the presence of C16TMA+. Thus, a cationic exchange be-tween C16TMA+ molecules and Ce3+ cations was found to occur dur-ing the TIE procedure. Some preliminary experiments wereperformed to adjust the duration of the TIE procedure in order toreach a high cerium content in the MCM-41. The optimal time wasfound to be between 12 and 24 h. The Si/Li molar ratio is close to 11even though no lithium phase was observed by XRD.

3.2. SO2 adsorption

The CeO2-MCM-41 sample exhibits no particular SO2 removal ca-pacity (breakthrough curve not presented here) as the SO2 adsorption

0

50

100

150

200

250

0 3000 6000 9000 12000 15000 18000Time (s)

[SO

2] (

ppm

)

(a) (b)

Fig. 7. SO2 breakthrough curves of (a) the CuO/CeO2-MCM-41 sample and (b) Li-dopedMCM-41 sorbent. Conditions: T=673 K, SV=25,000 h−1, 53 Nl/h, gas composition:250 ppm SO2, 10% O2 in N2. The breakthrough curve of the CeO2-MCM-41 sample isnot shown since the SO2 adsorption capacity is very small (b5 mg SO2/g). Accordingto this result, the breakthrough curves of the MCM-41 sample and the empty reactorwere not performed.

10 20 30 40 50 60

Inte

nsity

(a.u

.)

+++

+ + +++ +

++++ + +

O

*

*

*

* * * * **

(a)

(b)

2 4 6 8 2θ (CuKα)

Inte

nsity

(a.u

.)

(a)

(b)

2θ(CuKα)

Fig. 8. XRD pattern of (a) the sulfated CuO/CeO2-MCM-41 sample and (b) Li-dopedMCM-41 sorbent. (Inset) Corresponding low angle XRD patterns.

Page 6: YannicK Mesoporous absorbtie

Fig. 9. TEM micrographs of the sulfated Li-doped MCM-41 sorbent.

40 Y. Mathieu et al. / Fuel Processing Technology 99 (2012) 35–42

capacity was found lower than 5 mg SO2/g. The SO2 breakthroughcurve of the CuO/CeO2-MCM-41 sample is shown in Fig. 7a. The sor-bent is found to completely adsorb SO2 for 460 s. The SO2 adsorptioncapacity measured by integration of the breakthrough curve is 15 mgSO2/g. As it can be seen, the saturation is reached around 220 ppmafter 6000 s. However, the inlet SO2 concentration of 250 ppm is notretrieved after the saturation of the sorbent. This suggests that a frac-tion of SO2 reacts with oxygen and/or is oxidized by CeO2 yielding theformation of SO3 which is not seen by the UV analyzer. A significantincrease of the SO2 removal performance is observed for the Li-doped MCM-41 sorbent (Fig. 7b). The adsorption capacity is 96 mgSO2/g which is more than 6 times higher than the CuO/CeO2-MCM-41 sample and all the SO2 is completely adsorbed for 7500–8000 s.The saturation is reached after 18,000 s with around 220 ppm SO2

detected at the reactor outlet. This result clearly shows that the com-bination of copper oxide, cerium oxide, atacamite and lithium ap-pears to be necessary to reach a high SOx removal activity.

0

200

400

600

800

1000

1200

0 1000 2000 3000 4000 5000 6000 7000Time (s)

Tem

pera

ture

(K

)

0

200

400

600

800

1000

1200

[SO

] (

ppm

)2

(a)(b)

Fig. 10. TPD of (a) the sulfated CuO/CeO2-MCM-41 sample and (b) Li-doped MCM-41sorbent performed under N2 (Dv=53 Nl/h).

3.3. Characterization of the sulfated sorbents

The XRD pattern of the sulfated CuO/CeO2-MCM-41 sample ex-hibits no particular modifications as characteristic peaks of CeO2

and CuO are still observed (Fig. 8a). This shows a lack of reactivityof the metal oxides towards SOx in this sample under these experi-mental conditions (673 K, SV=25,000 h−1). However, for the sulfat-ed Li-doped MCM-41 sorbent, several peaks of chalchantite(CuSO4·5H2O) and lithium sulfate (Li2SO4·H2O) are observed (Fig.8b). The presence of hydrated sulfates could be explained by the ex-position of the sample to moisture after the tests. Furthermore, theCeO2 and CuO peaks completely disappear suggesting a high activityof the metal oxides towards SOx in this sorbent. No cerium sulfatesare observed confirming that CeO2 only acts as an oxidation promot-er. The mesoporous organization is not affected confirming the stabil-ity of both the CuO/CeO2-MCM-41 sample and Li-doped MCM-41sorbent at 673 K after sulfation (Fig. 8, inset).

The sulfation reaction induces some morphological changes in theLi-doped MCM-41 sorbent as several grains are surrounded by

10 20 30 40 50 60

: SiO2

: Li2SiO3 : CeO2

2 4 6 8 2θ (CuKα)

inte

nsity

(a.

u.)

inte

nsity

(a.

u.)

2θ (CuKα)

Fig. 11. XRD pattern of the Li-doped MCM-41 sorbent after TPD. (Inset) Correspondinglow angle XRD patterns.

Page 7: YannicK Mesoporous absorbtie

Table 4Crystalline phases observed by XRD before and after sulfation and TPD peaks of the dif-ferent MCM-41 samples.

Sample XRD TDP peaks Cadsb

(mg SO2/g)Fresh Aftersulfation

T (K)a Max(K)

[SO2]max

(ppm)

CeO2-MCM-41 CeO2 –c 773–1073 933 35 b10CuO-MCM-41 CuO – 773–973 863 160CuO/CeO2-MCM-41

CeO2

CuOCeO2

CuO793–973 873 120 15

CuO/LiCl-MCM-41d

CuOatacamite

Li2SO4 673–1013 873 225 80>1073 /e 265

Li-doped MCM-41 sorbent

CeO2

CuOatacamite

CuSO4

Li2SO4

673–1013 953 700 130>1073 / 510

a Temperature range of the desorption peak.b Adsorption capacity measured by integration of the desorption peaks.c No XRD peaks observed.d Chemical composition of 4.1 wt.% Cu, 0.8% Li.e Not determined as the max temperature for the TPD experiment was around

1123 K.

41Y. Mathieu et al. / Fuel Processing Technology 99 (2012) 35–42

globular particles which may be attributed to the sulfated species(Fig. 9). The aspect of the majority of MCM-41 grains is howevernot changed.

The TPD curves of the sulfated CuO/CeO2-MCM-41 sample and Li-doped MCM-41 sorbent are shown in Fig. 10. As it can be seen, onlyone desorption peak is observed for the sulfated CuO/CeO2-MCM-41sample. The maximum of desorption (~100–120 ppm SO2) occurs be-tween 833 and 953 K. Mu and Perlmutter clearly showed that the de-composition of CuSO4·5H2O under N2 occurs in the sametemperature range [29]. Thus, the desorption peak in the sulfatedCuO/CeO2-MCM-41 sample corresponds to the decomposition ofCuSO4 which is in a highly dispersed state in the MCM-41 and istherefore not detectable by XRD. The adsorption capacity, calculatedby integration of the desorption peak, is about 21 mg SO2/g. Thisvalue is higher than the adsorption capacity calculated by integrationof the breakthrough curve (15 mg SO2/g). It is worth mentioning thatthe TPD is performed under a non oxidative atmosphere (no forma-tion of SO3 can occur). Therefore, it is likely that previously adsorbedSO3 molecules have been desorbed in the form of SO2 thus leading toa SO2 adsorption capacity which is higher than the one measured byintegration of the breakthrough curve. For the sulfated Li-dopedMCM-41 sorbent, two distinct desorption peaks are observed. Thefirst desorption peak corresponds to the decomposition of CuSO4

and starts at around 763 K with a maximum observed around 963 K(~700 ppm SO2) and finishes at 1023 K. The second desorption occursat higher temperature (~1153 K) and may correspond to the decom-position of Li2SO4 as reported in the literature [30]. The adsorption ca-pacity as determined in the desorption phase is therefore 130 mgSO2/g which is one of the highest values found in the current litera-ture for porous silica supports. Again the desorption process leadsto a higher adsorption capacity than the adsorption process. An XRDanalysis performed on the MCM-41 sorbent after the TPD confirmsthe complete transformation of the sorbent into SiO2 (quartz) andlithium silicate (Fig. 11). Several sharp peaks of CeO2 are alsoobserved.

3.4. Proposed mechanism

The main crystalline phases determined by XRD and inferred fromthe TPD peaks for the various samples involved in this study are sum-marized in Table 4. As it was shown, it is necessary to combine metaloxides (CuO, CeO2) in a highly dispersed state, atacamite and lithiumto reach a high SOx removal activity. However, the role of both ataca-mite and lithium was not clearly established during the sulfation pro-cess. Therefore, an intermediate MCM-41 sorbent type containingonly CuO and lithium chloride was synthesized to investigate boththe crystalline phases formed after the sulfation reaction and the de-sorption peaks developed during the TPD. The XRD of the fresh CuO/LiCl-MCM-41 sample exhibits CuO and atacamite crystalline phases.However, only lithium sulfate peaks are observed by XRD after sulfa-tion. Two characteristic desorption peaks attributed respectively toCuSO4 and Li2SO4 are observed by TPD between 673 and 1013 K andfor a temperature higher than 1073 K, respectively. The area of thedesorption peak attributed to CuSO4 is about three times lower thanin the Li-doped MCM-41 sorbent resulting in a lower adsorption ca-pacity of 80 mg SO2/g. This confirms that the highly dispersed CuO,and therefore not detected by XRD, did react with SOx during the sul-fation reaction yielding the formation of CuSO4 despite the absence ofthe corresponding XRD peaks. Thus, the presence of an oxidation pro-moter such as cerium oxide efficiently enhances the sulfation of CuO.According to the previous observations made by Centi et al. who ob-served that the sulfation process in CuO/Al2O3 material occurs via ad-sorption of SO3 on CuO sites [10]. The main following reactions mayoccur:

SO2 þ 1

/

2O2→SO3ðcatalyzed by CeO2Þ ð3Þ

CuO þ SO3→CuSO4 ð4Þ

Li2O þ SO3→Li2SO4 ð5Þ

CuO þ SO2 þ 1

/

2O2→CuSO4 ð6Þ

Li2O þ SO2 þ 1

/

2O2→Li2SO4 ð7ÞIt is noteworthy that Li2O is not a stable species in the test condi-

tions and is actually in the form of LiCl or LiOH before reacting withSOx. Furthermore, the decomposition of atacamite via reaction (2)may lead to the formation of highly active CuO sites.

4. Conclusion

A high specific surface area and thermally stable MCM-41 sorbentwas prepared for the removal of SOx at 673 K. The investigated for-mulation combined a SOx trap (copper oxide), an oxidation promoter(cerium oxide) and lithium chloride to assist the sulfation reaction.This type of sorbent exhibits characteristic XRD peaks of copper andlithium sulfate after the sulfation. A high adsorption capacity of130 mg SO2/g was obtained by integration of the TPD curve which isone of the highest values mentioned in the literature for porous silicabased sorbents. Further studies are underway, dealing with the re-generation of the sorbent; in particular a regeneration process involv-ing a reductive gaseous flow (hydrogen, methane) is likely tosignificantly decrease the decomposition temperature of copper andlithium sulfates and to avoid the collapse of the MCM-41 supportthat has been observed.

Acknowledgments

The authors would like to thank Sophie Dorge from the Labora-toire Gestion des Risques et Environnement for her help in SO2 ad-sorption experiments and interpretations.

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