CO2 adsorption–desorption properties of zeolite beta prepared from OSDA-free synthesis

Microporous and Mesoporous Materials 219 (2016) 125e133

Contents lists available at ScienceDirect

Microporous and Mesoporous Materials journal homepage: www.elsevier.com/locate/micromeso

CO2 adsorptionedesorption properties of zeolite beta prepared from OSDA-free synthesis Yoshihiro Kamimura, Marie Shimomura, Akira Endo* National Institute of Advanced Industrial Science and Technology (AIST), AIST Central 5-2, 1-1-1 Higashi, Tsukuba, Ibaraki, 305-8565, Japan

a r t i c l e i n f o

a b s t r a c t

Article history: Received 1 July 2015 Received in revised form 27 July 2015 Accepted 28 July 2015 Available online 12 August 2015

Recent research has demonstrated how the economically and environmentally benign organic structuredirecting agent (OSDA)-free synthesis of zeolite betadbased on the seed-assisted methoddexhibits superior properties in applications as antibacterial agents, catalysts and adsorbents in comparison with conventional zeolite beta. One possible way to take advantage of OSDA-free beta is to apply the material in energy-efficient capture of carbon dioxide (CO2) from a variety of large emission sources. For this reason, we have investigated the CO2 adsorptionedesorption properties of OSDA-free beta for the first time. The results of elemental analysis, N2 and CO2 adsorptionedesorption, evaluation of the isosteric heat of adsorption and solid-state NMR measurements have shown that OSDA-free beta exhibited higher crystallinity, larger pore volume and higher aluminum (Al) as well as sodium (Na) cation content with enhanced homogeneity in Al distribution within the *BEA framework as compared to the conventional TEA-directed beta synthesized using tetraethylammonium hydroxide (TEAOH). These unique properties of OSDA-free beta resulted in significant improvements of CO2 adsorptionedesorption capability over a wide pressure region under moderate temperature in comparison with conventional zeolite beta. © 2015 Elsevier Inc. All rights reserved.

Keywords: Zeolite beta Organic structure-directing agent (OSDA)free synthesis CO2 capturing Al distribution Isosteric heat of adsorption

1. Introduction Aluminosilicate zeolites are crystalline microporous materials composed of tetrahedrally-coordinated SiO4 and AlO4 units. Owing to their well-defined micropore architectures with intricate multidimensional channel systems and charge compensating cations located within cavities, zeolites have been widely applied in the fields of molecular sieving, adsorption, ion-exchange processes and shape-selective catalysis [1e7]. The state-of-the art synthesis of recent zeolites often employs commercial or designed organic structure-directing agents (OSDAs), which successfully broadened the framework topology of zeolites with promising catalytic performance and/or adsorptionedesorption properties [3,4]. However, a major drawback is that the use of OSDA makes commercialization of zeolites extremely expensive and additionally adds extra complex synthesis stages [8,9], which strictly limits the industrial usage of OSDA-directed zeolites. In response to this issue, OSDA-free syntheses of zeolites with the aid of seed crystals have allowed alternative routes to simple, robust and low-cost production of

* Corresponding author. Tel.: þ81 29 861 4864; fax: þ81 29 861 4660. E-mail address: [email protected] (A. Endo). http://dx.doi.org/10.1016/j.micromeso.2015.07.033 1387-1811/© 2015 Elsevier Inc. All rights reserved.

industrially valuable zeolites such as beta [10e14], RUB-13 [15,16], ZSM-11 [17], ZSM-12 [18e20], TON-type [21], MAZ-type [22], PAUtype [17] and MCM-68 [23]. In particular, zeolite beta (having the framework topology code of *BEA as defined by the Structure Commission of the International Zeolite Association (IZA-SC) [24]), which possesses unique threedimensionally intersected 12 membered-ring channel systems (6.6
6.7 Å along a- and b-axes and 5.6
5.6 Å along c-axis) [25,26], is one of the most important synthetic zeolites that has been extensively studied and attracted a considerable attention from both fundamental and industrial standpoints [25e30]. The conventional synthesis of beta required the use of OSDAs until Xiao and coworkers [10] reported the first OSDA-free synthesis of beta in which calcined beta seed was added to a Naealuminosilicate gel in the absence of OSDA. Furthermore, Okubo and Itabashi [12,13,17] have extensively investigated the detailed effects of various synthetic parameters to optimize the synthesis conditions to produce OSDA-free beta with high reproducibility. Moreover, they have discovered that OSDA-free beta possessed high Al content in the *BEA framework (Si/Al <7) with a greater degree of well-defined microstructural features (crystallinity, micropore volume and morphology) as compared with conventional beta synthesized using OSDAs [12,13,31]. Since the reporting of OSDA-free synthesis,

126

Y. Kamimura et al. / Microporous and Mesoporous Materials 219 (2016) 125e133

several important results have been published, in which OSDA-free beta showed enhanced performance in antibacterial [32], N2O decomposition [33], benzene alkylation [34], hexane cracking [23], selective catalytic reduction (SCR) of NOx [35] and NO sorption [36] applications when compared with conventional beta and other common zeolites. Although the key factors for understanding the aforementioned improvements remain unclear, such results lead us to believe that Al-rich OSDA-free beta having high microporosity should be further studied in a wider range of applications. One such application to determine if the resulting properties of OSDA-free beta are advantageous is within the energy-efficient capture of CO2 based on pressure swing adsorption (PSA). Conventional beta has already been applied as practical adsorbents [37e42] and catalysts [43e45] for energy-efficient management (sequestration and utilization) of CO2 emitted from large stationary sources [46e49]. Therefore, with higher micropore volumes in addition to increased Al content in OSDA-free beta it is feasible to assume enhanced CO2 adsorption capacity when compared with conventional beta under equal pressure swing conditions. However, to the best of our knowledge there are no reports hitherto on CO2 adsorptionedesorption properties for OSDA-free beta. For this reason, understanding the CO2 capturing properties of OSDA-free beta over a wide pressure region is essential not only in this process, but also providing general insights into the field of adsorption and catalyst applications. In the present work, we have investigated the CO2 adsorptionedesorption behavior of OSDA-free beta for the first time and correlated the results with the characteristic features of both OSDAfree and conventional beta to help determine the key factors to further broaden its applicability. 2. Experimental 2.1. Materials The following raw materials were used as provided: Cab-O-Sil® (Grade M5, Cabot) as a silica source, sodium aluminate (NaAlO2, Wako Pure Chemical Industries, Ltd.) as an aluminum source and sodium hydroxide solution (NaOH, 50% w/v in water, Wako Pure Chemical Industries, Ltd.) as an alkali source. For the synthesis of conventional beta products, tetraethylammonium hydroxide (TEAOH, 35 wt.% in water, SigmaeAldrich Co. LLC) was used as an OSDA. Sodium nitrate (NaNO3, 99.0 wt.%, Wako Pure Chemical Industries, Ltd.) was used for Na ion-exchange. Ultra-purified water was used for all experiments.

Table 1 The Si/Al, Na/Al and textural properties of BEATEA and BEAOSDAF samples. Here, BEATEA denotes the calcined product of conventional TEA-directed synthesis, and BEAOSDAF denotes the as-made product (non-calcined) of seed-assisted, OSDA-free synthesis respectively. Sample

Si/Ala

Na/Ala

Vmicro [cm3 g1]b

Vtotal [cm3 g1]c

BEATEA no. 2 BEATEA no. 2-IEd BEAOSDAF

9.3 7.6 6.8

0.18 0.53 0.98

0.18 0.19 0.22

0.32 0.35 0.37

a

Si/Al and Na/Al ratios of the product determined by chemical analysis (ICP-AES). Micropore volume (Vmicro) of products evaluated by the t-plot method. c Total pore volumes (Vtotal) of products estimated from N2 adsorption isotherms at a relative pressure of P P0 1 ¼ 0.95 with the assumption of bulk liquid density of N2 ¼ 0.808 cm3 g1 (Gurvich rule). d Na ion-exchanged (IE) form of BEATEA no. 2. b

H-form, Na ion-exchange of BEATEA no. 2 was performed by following a common procedure; 1 g of zeolite sample was added to 100 mL of a 1 M of sodium nitrate solution and stirred at 333 K for 24 h. After filtration and washing with ultra-purified water, the same ion-exchange procedure was repeated four times. Finally, the sample was dried at 333 K to prepare Na ion-exchanged beta sampleddenoted as BEATEA no. 2-IE hereafter. The OSDA-free synthesis of beta was performed according to a previously reported method [12,13], in which hydrothermal treatment of a Na-aluminosilicate gel (chemical composition of SiO2:0.3 Na2O:0.01 Al2O3:20 H2O) in the presence of 10 wt.% of calcined beta seed (BEATEA no. 1 with Si/Al ¼ 11 and Na/Al ¼ 0.06) relative to the silica source. The reason why beta with Si/Al ¼ 11 (BEATEA no. 1) was chosen for the seed is because higher Si/Al ratio in seeds can give faster crystallization of OSDA-free beta. This is supported by the previous report [12], in which the addition of beta seed with lower Si/Al ratios (9.2 and 7) to Na-aluminosilicate gel resulted in much longer crystallization periods of OSDA-free beta as compared to the case of using beta seed with Si/Al ¼ 12. Thus, use of beta seed with highest Si/Al in this study is beneficial from a viewpoint of reducing the hydrothermal reaction period. The seed-embedded gel was then transferred to a 23 mL stainless-steel autoclave and subjected to hydrothermal treatment at 423 K for 48 h under static conditions. After hydrothermal treatment the autoclave was quenched and solid products were recovered by filtration, washed carefully with ultra-purified water until neutral pH and subsequently dried at 333 K in an oven overnight. In the following section, OSDA-free beta prepared by the seed-assisted synthesis was denoted as BEAOSDAF.

2.2. Synthesis procedures

2.3. Characterization

Two different TEA-directed beta samples, denoted as BEATEA nos. 1 and 2 (Table 1), were prepared by the hydrothermal treatment of Na-aluminosilicate gels in the presence of TEAOH with the following chemical composition: SiO2:0.0357 Na2O:0.175 TEA2O:xAl2O3:14 H2O, where x ¼ 0.0286 (BEATEA no. 1) or 0.04 (BEATEA no. 2). The aluminosilicate gels containing TEAOH were transferred to 23 mL Teflon-lined stainless-steel autoclaves and subjected to hydrothermal treatment at 438 K for 96 h under agitation. After quenching the autoclaves, products were collected by filtration and washed carefully with ultra-purified water until neutral pH prior to drying at 333 K in an oven overnight. All TEAdirected beta samples were calcined at 823 K for 10 h to remove residual TEAþ from the cavities. In this study, BEATEA no. 1 was used as the seed for OSDA-free synthesis. BEATEA no. 2 followed by subsequent Na ion-exchange was subjected to the comparison experiment for evaluation of the CO2 adsorptionedesorption properties. Since calcined form of TEA-directed beta was mainly in

Wide-angle powder X-ray diffraction (XRD) patterns of the solid products were collected by a powder X-ray diffractometer (D8 Advance, Bruker AXS Corp.) with Cu Ka radiation (l ¼ 0.15406 nm, 40 kV, 40 mA) at 2q values of 5e40 with a scanning step of 0.04 at a scanning speed of 2 min1. Morphology and crystal size of products was observed by a field-emission scanning electron microscope (FE-SEM, SU9000, Hitachi High-Technologies Corp.). Elemental analysis was performed by an inductively coupled plasma-atomic emission spectrometer (ICP-AES, Optima-8300, PerkineElmer, Inc.). Solid-state 27Al and 29Si NMR spectra were recorded on a Bruker AVANCE-400 spectrometer equipped with a 4 mm magic angle-spinning (MAS) probe. 27Al dipolar-decoupling (DD)-MAS NMR spectra were recorded at a resonance frequency of 104.3 MHz, with a pulse width of 0.5 ms and a recycle time of 1.5 s. 29 Si DD-MAS NMR spectra were acquired at a resonance frequency of 79.5 MHz, with a pulse width of 3 ms and a recycle time of 60 s. All the samples were spun at 12.5 kHz in standard 4 mm zirconia

Y. Kamimura et al. / Microporous and Mesoporous Materials 219 (2016) 125e133

rotors. 27Al and 29Si chemical shifts were referenced to 1 M aluminum nitrate solution (dAl ¼ 0 ppm) and polydimethylsilane (dSi ¼ 34.16 ppm from TMS) respectively. To avoid misleading the presented data normalization of the 27Al and 29Si MAS NMR spectra allows for clear representation. Textural properties of the samples were obtained by performing N2 adsorptionedesorption measurements using an automatic gas adsorption apparatus (BELSORPmax and BELSORP-mini, Microtrac-BEL Japan, Inc.) at 77.36 K. Prior to the gas (N2 and CO2) adsorptionedesorption measurements, all the beta samples were degassed at 673 K for 8 h under vacuum. Micropore volumes (Vmicro) of products were calculated via the tplot analysis method. Total pore volumes (Vtotal) of the products were estimated from N2 adsorption isotherms at a relative pressure of P P0 1 ¼ 0.95 with the assumption of bulk liquid density of N2 at 77.36 K (0.808 cm3 g1) [50]. CO2 adsorptionedesorption measurements across low pressure regions (ca. 0.01 kPa) to highly pressurized regions (ca. 3200 kPa) were performed at moderate temperatures at 298.15 K, using an automatic adsorption measurement apparatus (BELSORP-HP, Microtrac-BEL Japan, Inc.). CO2 adsorptionedesorption measurements from very low pressure regions (ca. 1
105 kPa) to atmospheric pressure (ca. 100 kPa) were performed at moderate temperatures at 288.15, 293.15 and 298.15 K, using an automatic adsorption measurement apparatus (BELSORP-max, Microtrac-BEL Japan, Inc.). 3. Results and discussion 3.1. Product characterization Fig. 1(a, b, and c) shows the XRD patterns of BEATEA nos. 1, 2 and 2-IE, which are identified as having the *BEA topology. In BEATEA no. 2 and 2-IE, some unknown impurities were detected, although after ion exchange the level of impurities reduced when compared with BEATEA no. 2. In addition, BEATEA no. 2-IE exhibited sufficient micropore as well as total pore volume, and impurity phases were not detected by FE-SEM observations (see following sections). Thus, we considered that the degree of impurities in BEATEA no. 2-IE to be negligible, and subjected the material to subsequent CO2

Fig. 1. Wide-angle powder XRD patterns of (a) BEATEA no. 1 (used as the seed in OSDAfree synthesis of BEAOSDAF), (b) BEATEA no. 2 (before Na ion-exchange), (c) BEATEA no. 2IE (after Na ion-exchange) and (d) BEAOSDAF (product of OSDA-free synthesis). Square and circle symbols in the Figure represent impurities of unknown phases.

127

adsorptionedesorption measurements. Elemental analysis (ICPAES) reveals Si/Al ratios for BEATEA nos. 1, 2 and 2-IE as 11, 9.3 and 7.6, and Na/Al ratios as 0.06, 0.18, and 0.53 respectively. Hence, BEATEA nos. 1 and 2 were mainly in H-form, while BEATEA no. 2-IE was in (H,Na)-form where the degree of Na ion-exchange was 53%. Fig. 1(d) shows the XRD pattern of BEAOSDAF confirming the presence of phase pure beta. Elemental analysis (ICP-AES) gives Si/ Al and Na/Al ratios of BEAOSDAF as 6.8 and 0.98 respectively. In comparison with the results of BEATEA no. 2-IE (Si/Al ¼ 7.6 and Na/ Al ¼ 0.53), BEAOSDAF exhibited a similar but slightly lower Si/Al ratio (slightly higher Al content). The high Na/Al ratio in BEAOSDAF indicates that the material is mainly in the Na-form. The primary particle morphology of all beta products (Fig. 2) display welldefined crystals of truncated octahedral morphology analogous to natural counterparts of beta, such as Tschernichite [51]. The primary particle size for BEATEA no. 1 and BEAOSDAF (Fig. 2a and d) was no more than 120 nm, while that for BEATEA nos. 2 and 2-IE (Fig. 2b and c) was mostly less than 50 nm. These primary particles were aggregated to form larger secondary particles. Furthermore, no observations of any amorphous material or impurity phases were seen in any of the beta products. By comparing the synthesis methods demonstrated in this study, OSDA-free synthesis of beta can be achieved at lower hydrothermal temperature with short crystallization period (423 K for 48 h) as compared to the TEAdirected synthesis (438 K for 96 h). Although 10 wt.% of beta seed was used, such time savings and less energy input (heat) in OSDAfree synthesis is advantageous which will result in low-cost and environmentally friendly production of zeolite beta. In the present study, we intend to compare the CO2 adsorptionedesorption properties of BEATEA no. 2-IE with BEAOSDAF because these zeolite products showed *BEA topology with similar Si/Al ratio and cation form. Therefore, further detailed characterization mainly focused on BEATEA no. 2-IE with BEAOSDAF. Fig. 3 shows the 29Si DD-MAS NMR spectra for selected beta products. At a first glance, the 29Si DD-MAS spectra for BEATEA nos. 2 and 2-IE, and BEAOSDAF were similar to those reported in previous literature [35,52]. All beta products exhibited two resonances centered close to dSi ¼ 108 and 111 ppm, which can be assigned to two Q4 (4Si, 0Al) units respectively. The split of the corresponding resonances indicates that there is at least two nonequivalent (4Si, 0Al) configurations existing within the *BEA framework [31]. Additionally, a further two resonances centered close to dSi ¼ 101 and 95 ppm attributed to Q4 (3Si, 1Al) and Q4 (2Si, 2Al) units were observed. According to Dedecek et al. [53,54], each Q4 (4Si, 0Al), Q4 (3Si, 1Al) and Q4 (2Si, 2Al) unit presented in the 29Si DD-MAS spectra generally represent the presence of AleOe(SieO)n>2eAl, AleOe(SieO)2eAl and AleOeSieOeAl sequences in the *BEA framework. With no apparent change in the 29Si DD-MAS spectra of BEATEA nos. 2 and 2-IE (Fig. 3a and b), their Si and Al sequences were generally unaffected by the Na ion-exchange process. Fig. 4 shows the resulting 27Al DD-MAS NMR spectra of selected beta samples. Prior to ion-exchange, BEATEA no. 2 (Fig. 4a), is comprised of both tetrahedral and octahedral Al coordinated atoms as is evident by the corresponding peaks centered at approximately dAl ¼ 60 and 0 ppm respectively [55,56]. The signal close to dAl ¼ 0 ppm is typically related to extra-framework Al (dealuminated species)dejected from the framework during calcination of the occluded TEAþ [52]. Hence, a degree of dealumination is observed within BEATEA no. 2 prior to the Na ion-exchange. The intensity of the signal, dAl ¼ 0 ppm, after BEATEA no. 2 had been subjected to ion exchange, reduced signifying the removal of the extra-framework Al (BEATEA no. 2-IE, Fig. 4b, solid line) as a result of the acidic NaNO3 ion exchange media [57]. Conversely, the 27Al DDMAS NMR spectrum for the as-made BEAOSDAF material (Fig. 4c, solid line) exhibited a sole peak in the region of dAl ¼ 50e60 ppm

128

Y. Kamimura et al. / Microporous and Mesoporous Materials 219 (2016) 125e133

Fig. 2. FE-SEM micrographs of (a) BEATEA no. 1 (used as the seed in OSDA-free synthesis of BEAOSDAF), (b) BEATEA no. 2 (before Na ion-exchange), (c) BEATEA no. 2-IE (after Na ionexchange) and (d) BEAOSDAF (product of OSDA-free synthesis).

attributed to tetrahedral Al species [55]. No additional peaks attributed to octahedral Al (dAl ¼ 0 ppm) were observed. Thus, with no observation of octahedral Al sites present and with all Al atoms residing within the *BEA framework, the OSDA-free beta material possesses a higher degree of crystallinity in comparison with the TEA-directed beta. Previous reports of varying zeolite framework topologies [12,17e20] adopting a seed-assisted, OSDA-free approach also showed products in the absence of extra-framework Al. Furthermore, all 27Al DD-MAS NMR spectra were fitted with Gaussian functions to deconvolute overlapped peaks. BEATEA no. 2 and 2-IE together with BEAOSDAF exhibited two overlapping peaks close to dAl ¼ 58 and 50 ppm as given by the dotted lines in Fig. 4.

The *BEA framework has been well studied and is known to contain nine T-sites [55] with the peaks centering at dAl ¼ 58 and 50 ppm assigned to Al occupying the T3eT9 and T1, T2 sites respectively [23,55,56] (schematic representation is shown in Supplementary information Fig. S1). Determining the ratios of the deconvoluted peak areas it was found that BEAOSDAF possesses more Al residing in the T1, T2 sites when compared with the BEATEA nos. 2 and 2-IE samples. On the basis of the 29Si and 27Al DD-MAS NMR spectra, even with similar Si/Al ratios between BEATEA no. 2-IE (Si/Al ¼ 7.6) and BEAOSDAF (Si/Al ¼ 6.8) the local distribution of Si and Al within the *BEA framework are different and depend on the synthetic methodology adopted (TEA-directed or OSDA-free).

Fig. 3. 29Si dipolar-decoupling (DD)-MAS NMR spectra for (a) BEATEA no. 2, (b) BEATEA no. 2-IE and (c) BEAOSDAF. Standard spectra and overlapped peaks (fitted with Gaussian functions) are presented with solid and dotted lines respectively.

Fig. 4. 27Al dipolar-decoupling (DD)-MAS NMR spectra for (a) BEATEA no. 2, (b) BEATEA no. 2-IE and (c) BEAOSDAF. Standard spectra and overlapped peaks (fitted with Gaussian functions) are presented with solid and dotted lines respectively.

Y. Kamimura et al. / Microporous and Mesoporous Materials 219 (2016) 125e133

Nitrogen adsorptionedesorption isotherms measured at 77.36 K allowed for the textural properties of the beta samples to be assessed (Fig. 5). All samples displayed Type I isotherms as defined in the International Union of Pure and Applied Chemistry (IUPAC) classification [58], and the uptake of adsorbate at low relative pressures (P P0 1 < 0.1) for the BEATEA nos. 2 and 2-IE samples was similar to conventional Al-rich beta zeolites (Si/Al ¼ 7.2) synthesized with TEAþ reported previously [59]. It is evident from the desorption branch of the isotherm that both BEATEA nos. 2 and 2-IE exhibit a narrow hysteresis loop, which closes at a relative pressure of P P0 1 ¼ 0.46. The type I isotherms show that it is mainly microporosity formed in all the beta products. Additionally, a relatively small amount of mesoporosity is present in the BEATEA nos. 2 and 2IE samples. With a narrow hysteresis curve present in the desorption branch of the BEATEA no. 2 sample (Fig. 5a), the level of mesoporosity displayed in BEATEA no. 2-IE is unlikely to originate from the Na ionexchange step. A plausible explanation for the formation of a relatively small amount of mesoporosity in BEATEA nos. 2 and 2-IE is related to the interparticle voids formed after aggregation of the smaller primary particles of sizes less than 50 nm (as observed in the FE-SEM micrograph Fig. 2b and c). Such mesoporosity has previously been reported by Camblor et al. [60], in which an Al-rich TEAdirected beta (Si/Al ¼ 7.4) with primary particle sizes as small as 10 nm gave rise to mesoporosity in the final product as a result of interstitial sites formed in between the crystallites. Table 1 summarizes the microporosity and total pore volume (Vmicro and Vtotal) of the beta products. BEAOSDAF shows larger microporosity as well as total pore volume (0.22 and 0.37 cm3 g1) than that of BEATEA nos. 2 and 2-IE (0.18e0.19 and 0.32e0.35 cm3 g1). The enhancement of pore volume is essential for CO2 capture applications as the capacity of zeolite materials to adsorb guest molecules increases proportionately [61]. The larger microporosity and total pore volume in the OSDA-free beta stem from the higher crystallinity as suggested by the abovementioned results of 27Al DD-MAS NMR and N2 adsorptionedesorption measurements.

129

from large emission sources, for example, from post-combustion flue gases emitted from coal-based power plants or industrial steel plants operating at low pressure (<100 kPa), or in high pressure applications, such as Integrated Gasification Combined Cycle (IGCC) plants operating in the range of 2000e3000 kPa, which emit high concentrations of pre-combustion CO2 [46,47]. Correlating the CO2 adsorptionedesorption behavior with that of the OSDA-free beta zeolite properties over a wide pressure range is of importance for real-life applications not only in the processes of capturing CO2 but in how best to optimize adsorbents and catalysts for applications in CO2 management. In the present study, CO2 adsorptionedesorption measurements on BEATEA no. 2-IE and BEAOSDAF were recorded via two experiments: i) from low pressure (ca. 1
102 kPa) to high pressures (ca. 3200 kPa) performed at 298.15 K; ii) from very low pressures (ca. 1
105 kPa) to atmospheric pressure (ca. 100 kPa) as a function of temperature at 288.15, 293.15 and 298.15 K (see Experimental section for more details). To the best of our knowledge, this is the first report detailing CO2 adsorptionedesorption isotherms for OSDA-free beta measured over a wide pressure region. 3.3. CO2 adsorptionedesorption up to 3200 kPa

The facile and inexpensive route to prepare highly crystalline OSDA-free beta can be utilized in energy-efficient capture of CO2

Fig. 6 compares the CO2 adsorptionedesorption isotherms at 298.15 K up to 3200 kPa for BEATEA no. 2-IE and BEAOSDAF. BEAOSDAF exhibited higher CO2 adsorption capacity than BEATEA no. 2-IE across all pressure regions up to 3200 kPa. Both materials display the following CO2 adsorptionedesorption behaviors: i) an initial adsorption resulting from rapid uptake within the available pores over a narrow pressure range in the region of 1
102e500 kPa, where the CO2 uptake at 500 kPa for BEATEA no. 2-IE and BEAOSDAF corresponds to 99 and 114 cm3 (STP) g1 (ca. 4.4 and 5.1 mmol g1), respectively; ii) in the pressure region above 500 kPa, the uptake of adsorbed CO2 was near to asymptote and increased only moderately up to 3200 kPa. At 3000 kPa the CO2 uptake for BEATEA no. 2-IE and BEAOSDAF corresponded to 125 and 133 cm3 (STP) g1 (ca. 5.6 and 6 mmol g1), respectively. Both beta samples displayed reversible desorption after being subjected to 3200 kPa at 298.15 K. Considering the CO2 molecules fill the available pore volume in the beta products up to 3200 kPa, the adsorbed uptake of CO2 (at 298.15 K) in Fig. 6 was converted into occupied volume (cm3 g1) by

Fig. 5. N2 adsorptionedesorption isotherms at 77.36 K for (a) BEATEA no. 2, (b) BEATEA no. 2-IE and (c) BEAOSDAF. Solid and open symbols in the Figure indicate adsorption and desorption points respectively. The isotherms for (b) BEATEA no. 2-IE and (c) BEAOSDAF are offset vertically by 400 and 800 cm3 (STP) g1 respectively.

Fig. 6. CO2 adsorptionedesorption isotherms at 298.15 K in the pressure range up to 3200 kPa for (a) BEATEA no. 2-IE (Si/Al ¼ 7.6) and (b) BEAOSDAF (Si/Al ¼ 6.8). Solid and open symbols in the Figure indicate adsorption and desorption points respectively. (Inset): The corresponding degree of pore filling of CO2 (FCO2 calculated based on total pore volume).

3.2. CO2 adsorptionedesorption measurements

130

Y. Kamimura et al. / Microporous and Mesoporous Materials 219 (2016) 125e133

using the bulk liquid density of CO2 (0.71 cm3 g1) [50]. Thereafter, the degree of pore filling of CO2 (FCO2) was determined by taking into account the total pore volume (Vtotal) evaluated by the N2 physisorption measurement (see Table 1). Fig. 6 (inset) displays FCO2total for the corresponding beta samples. The degree of pore filling at 100 kPa for BEATEA no. 2-IE and BEAOSDAF were ca. 59.2% and 72.5%. At the higher pressure region of 1000 kPa, the degree of pore filling with respect to BEATEA no. 2-IE and BEAOSDAF increased to 88.8% and 91.2% and further increased to 99% and 100% at 3000 kPa respectively. At 3000 kPa, almost all the pore filling was complete for both beta zeolites suggesting that the existing pores in both zeolites are accessible for CO2 molecules. 3.4. CO2 adsorption up to 100 kPa To give a clear representation for the detailed CO2 adsorption behavior at low pressure coverage, CO2 adsorption isotherms were measured for BEATEA no. 2-IE and BEAOSDAF at temperatures of 288.15, 293.15 and 298.15 K within the pressure range of 1
105e100 kPa (very low to atmospheric pressure). Fig. 7 shows a typical example of the CO2 adsorptionedesorption isotherm at 298.15 K for BEAOSDAF. The adsorption isotherm displays Type I character [58] where the uptake of adsorbed CO2 increased rapidly as a function of pressure up to ca. 20 kPa followed by a steady increase through to 100 kPa. The desorption branch of the isotherm shows that CO2 desorbs reversibly down to ca. 1 kPa at the same temperature of 298.15 K. No additional heating at high temperatures was necessary for the given desorption process. Fig. 8 compares the CO2 adsorption isotherms for BEATEA no. 2-IE and BEAOSDAF at 288.15, 293.15 and 298.15 K with all isotherms being similar in possessing Type I character [58]. BEAOSDAF exhibited higher CO2 adsorption capacity than BEATEA no. 2-IE in the pressure region of 1
103e100 kPa regardless of the adsorption temperatures. As seen in the adsorption isotherms at 298.15 K, BEATEA no. 2-IE and BEAOSDAF exhibited ca. 74 and 92 cm3 (STP) g1 at 100 kPa (CO2 capacities at 100 kPa for BEATEA no. 2-IE and BEAOSDAF are also expressed as ca. 3.3 and 4.1 mmol g1). When the adsorption temperature reduced from 298.15 to 293.15 or 288.15 K, the adsorbed uptake of CO2 in both BEATEA no. 2-IE and BEAOSDAF increased. Similar temperature dependence is evident in CO2 adsorption studies of other like zeolites, activated carbons and metal oxides as a function of varying temperature [47,61,62]. The isosteric heat of adsorption for BEATEA no. 2-IE and BEAOSDAF was estimated based on the ClausiuseClapeyron equation from the data

points of the CO2 adsorption isotherms shown in Fig. 8. Fig. 9 shows the isosteric heat of adsorption plotted against CO2 capacity (cm3 (STP) g1) and the number of adsorbed CO2 molecules in the *BEA unit cell. For BEAOSDAF, at the very low pressures (ca. 1
102 kPa) and low CO2 coverage, the heat of adsorption was 39.8 kJ mol1. Upon further loading of CO2, the heat of adsorption moderately decreased to 21.8 kJ mol1, which is very near to the heat of condensation ca. 22 kJ mol1 for the CO2 molecule [63]. Conversely, the heat of adsorption for BEATEA no. 2-IE is 45.4 kJ mol1 when at the very low pressures and low CO2 coverage. Increasing the CO2 coverage, the heat of adsorption moderately decreased to 28.8 kJ mol1. The isosteric heat of adsorption for both samples typically shows exothermic contribution by the physical adsorption [64]. Across all pressure points the isosteric heat of adsorption for BEATEA no. 2-IE (28.8e45.4 kJ mol1) was consistently higher than the corresponding BEAOSDAF (21.8e39.8 kJ mol1) sample. In a similar manner to the calculation presented in Fig. 6 (inset), Fig. 10 displays the degree of pore filling of CO2 (FCO2micro) related to the micropore volume (Vmicro) of the corresponding beta samples. FCO2micro is preferentially complete at the early stage of adsorption. Complete FCO2micro for BEAOSDAF occurs at ca. 32.9 kPa, with the corresponding BEATEA no. 2-IE sample having only 80% of the micropores filled at this pressure. Complete FCO2micro for BEATEA no. 2-IE only prevails when further increasing the pressure to 73.3 kPa. Data points given above the 100% threshold line represent pore filling in the mesopores and the interstitial sites between the smaller aggregated particles. 3.5. Differences in CO2 adsorption properties for OSDA-free and TEA-directed beta On the basis of the abovementioned results, an interesting question to discuss is why the CO2 adsorption properties between OSDA-free and conventional TEA-directed beta products are different across the wide pressure range. Plausible explanations relate to the unique properties of OSDA-free beta (BEAOSDAF) are suggested as follows: (i) larger pore volume, (ii) higher Al and Na cation content and (iii) different Al and cation distribution in the *BEA framework. Obviously the larger pore volume in BEAOSDAF should provide more capacity for ingress of accessible adsorbate molecules to the internal cavities and/or channels of the zeolite. The present study allows the adsorbed uptake of CO2 (cm3 (STP) g1) in relation to the pore volume to be calculated using the bulk liquid density of CO2

Fig. 7. Typical example of the CO2 adsorptionedesorption isotherms at 298.15 K in the pressure range below 100 kPa for BEAOSDAF (Si/Al ¼ 6.8). Here, the isotherms are expressed in (left) linear scale and (right) semi-logarithmic scale.

Y. Kamimura et al. / Microporous and Mesoporous Materials 219 (2016) 125e133

131

Fig. 8. CO2 adsorption isotherms at various temperatures at 288.15, 293.15 and 298.15 K in the pressure range up to 100 kPa for (a) BEATEA no. 2-IE (Si/Al ¼ 7.6) and (b) BEAOSDAF (Si/ Al ¼ 6.8). Here, the isotherms are expressed in (Left) linear scale and (Right) semi-logarithmic scale.

(0.71 cm3 g1) [50]. For example, an uptake of CO2 corresponding to a pore volume of 0.01 cm3 g1 results in 3.6 cm3 (STP) g1 of CO2 being adsorbeddequivalent of 0.16 mmol g1. As listed in Table 1, BEAOSDAF has 0.03 cm3 g1 more microporous volume compared with BEATEA no. 2-IE, therefore there will be at least 10.8 cm3 (STP) g1 (0.48 mmol g1) additional uptake of CO2 in the micropores when FCO2micro is 100%, which relates to pressures of 73.3 and 32.9 kPa for BEATEA no. 2-IE and BEAOSDAF respectively. Regarding the total pore volume there is an additional 0.02 cm3 g1 more volume in the BEAOSDAF samples over that of BEATEA no. 2-IE, hence an extra CO2 uptake of 7.2 cm3 (STP) g1 (0.32 mmol g1) when FCO2total ¼ ca. 100%, relating to a pressure close to 3000 kPa (see Fig. 6, inset). Therefore the larger pore volume in BEAOSDAF owing to the well-defined crystallinity of the zeolite framework brings advantages in higher available CO2 adsorption capacity in pressure swing operations across a wide pressure range. As can be seen in Fig. 6 (inset) and 10, it is clearly evident that the degree of pore filling (FCO2micro and FCO2total) in BEAOSDAF is much higher at equal equilibrium pressures across the wide pressure range than that of BEATEA no. 2-IE. Comparatively, there is not so great a difference in the Al content between BEAOSDAF and BEATEA no. 2-IE; however, there is considerably more Na in the former sample (Table 1). Na content is known to effect the uptake of

adsorbed CO2 at equivalent pressures because the Na cation possesses a higher electric field than a corresponding proton [61,65]. Additionally, CO2 has large polarizability together with an electric quadrupole moment [61,65]. Therefore CO2 is more prone to perform ionequadrupole interactions with Na cations located within the extra-framework, and thus, the higher the content of Na cations the higher CO2 capacity results at equivalent equilibrium pressuresdas is the case for the BEAOSDAF sample. Typical trends, when relating the isosteric heat of adsorption to the levels of Na cations, suggest that the more Na counter ions present the greater the interaction with CO2, and therefore one would assume that the BEAOSDAF sample would exhibit a higher isosteric heat of adsorption value because of the higher Na loading levels. However, in the present study, the results suggest the contrary (Fig. 9), with of BEATEA no. 2-IE having the higher value and thus interacting more strongly with CO2. It is assumed that the different CO2 adsorption character is concurrent with differences in the nature of the cation sites accompanied by variation in Al homogeneity. From Fig. 9 (Right), the ingress of one CO2 molecule adsorbed at a site-specific location within the *BEA unit cell gives an isosteric heat of adsorption for BEATEA no. 2-IE and BEAOSDAF as 42.9 and 38.6 kJ mol1 respectively. These values at very low coverage of CO2 are generally observed in a variety of Na-form

Fig. 9. Isosteric heat of adsorption of CO2 on (a) BEATEA no. 2-IE (Si/Al ¼ 7.6) and (b) BEAOSDAF (Si/Al ¼ 6.8). (Left) isosteric heat of adsorption plotted against CO2 capacity (expressed in linear scale), and (Right) isosteric heat of adsorption plotted against number of CO2 molecules in the *BEA unit cell (expressed in linear scale).

132

Y. Kamimura et al. / Microporous and Mesoporous Materials 219 (2016) 125e133

Fig. 10. Degree of pore filling of CO2 (FCO2 calculated based on micropore volume) at 298.15 K in the pressure range below 100 kPa for (a) BEATEA no. 2-IE (Si/Al ¼ 7.6) and (b) BEAOSDAF (Si/Al ¼ 6.8).

zeolites, for example, Na-SSZ-13 (ca. 44 kJ mol1) [61], Na-RHO (ca. 40 kJ mol1) [62] and Na-FER (42e47 kJ mol1) [66], while the isosteric heat of adsorption for H-form zeolites are mainly below 30 kJ mol1 [67,68]. Therefore, consideration is given that the majority of the adsorbed CO2 molecules in both the BEATEA no. 2-IE and BEAOSDAF samples at the point of immediate adsorption at verylow loadings/pressures are interacting with the extra-framework Na cations through ionedipole interactions [61]. In terms of equilibrium pressure, one CO2 adsorbed molecule relates to equilibrium pressures at ca. 2
102 kPa for BEAOSDAF and ca. 1
101 kPa for BEATEA no. 2-IE, whereby BEAOSDAF already exhibits higher CO2 adsorption capacity (Fig. 8 right) even with a lower isosteric heat of adsorption. One reasonable explanation for such unique adsorption character is the nature of the cation location and Al distribution in the *BEA framework. As shown by 29Si and 27Al DD-MAS NMR measurements (Figs. 3 and 4), zeolite beta obtained by the different synthesis methodologies (TEA-directed or OSDA-free) resulted in variation in the local distributions of Si and Al atoms (T-sites). Indeed there is difficulty to assign the exact crystallographic position of the Si and Al T-atoms by single pulse NMR onlydalthough different T-sites are claimed, especially for Al distribution, which can provide indirect evidence that the locations of the counter cations should at least be varied. Together with BEAOSDAF possessing a lower isosteric heat of adsorption with that of BEATEA no. 2-IE, such additional energy contribution in BEATEA no. 2-IE is presumably due to the predominant formation of dual cation sites, in which two neighboring Na cations complementary interact with one CO2 molecule. The possibility of CO2 interaction with dual cations was proposed by Nachtigall and Cejka [66,69,70], in which they controlled the homogeneity of Al distribution in Na-FER (model zeolite with Si/Al ¼ 15.7) by changing the type of organic structure-directing agents during the synthesis [66]. When Al sites were distributed homogeneously in the vicinity of Na cations, one Na cation can interact with an average of one CO2 molecule where the isosteric heat of adsorption is in the range of 42e47 kJ mol1 (depending on Al T-site location) [66]. On the other hand, when Al sites were less homogeneously distributed, two neighboring Na cations are preferred to interact with one CO2 molecule, and as a result, the isosteric heat of adsorption increases (46e56 kJ mol1) [66]. Thus, in accordance with the previous report, we assume that the reversed trend in the isosteric heat of adsorption behavior in BEAOSDAF (38.6 kJ mol1) and BEATEA no. 2-IE (42.9 kJ mol1) is mainly derived from the difference in the formation of single or

dual cations. It is still feasible that dual cation formation is present in BEAOSDAF since this zeolite contains high Al content. However, the implication is that a higher population of single cation sites exist in BEAOSDAF with enhanced homogeneity of the Al sites and Na cation location when compared with the BEATEA no. 2-IE. Furthermore, we speculate that the higher population of single Na cation sites effectively serves as an adsorption center for one CO2 molecule, whereas dual cation sites serve to adsorb one CO2 molecule from two Na cations. Such a unique feature in BEAOSDAF gives rise to a rapid uptake of CO2 and higher capacity at low pressures (ca. 1
101 kPa) in contrast to BEATEA no. 2-IE (Fig. 8 right). Consequently, homogeneous Al distribution in OSDA-free beta is an important parameter for optimized tailoring of the zeolites in catalysis and adsorbent applications [32e36]. In a related matter, Sazama et al. [35] reported that the Al distribution and cation position in Fe ion-exchanged OSDA-free beta (Si/Al ¼ 4.6) were different from Fe ion-exchanged TEA-directed beta (Si/Al ¼ 11.3), and claimed that the unique Al distribution in Fe-OSDA-free beta gave enhanced activity of N2O decomposition compared with the TEA-directed beta. Valtchev et al. [71] reported that synthesis of zeolite beta under fluoride media with an aid of beta seed resulted in different Al distribution (active sites), and such beta product exhibited higher performance in m-xylene transformation as compared to the TEA-directed beta even both samples possessed similar Si/Al ratio (ca. 33). This is one good example that implies and supports the essential role of seed in the formation of *BEA framework with different Al distribution. Further investigations are required to determine the Al framework distribution within the OSDA-free zeolite beta and the impact on CO2 adsorption behavior. 4. Conclusion In the present study, the detailed CO2 adsorptionedesorption properties of OSDA-free beta prepared by adopting a seed-assisted method was investigated. Detailed characterization of OSDA-free beta revealed that the material exhibited high crystallinity, high Al and Na content (Si/Al ¼ 6.8, Na/Al ¼ 0.98), large pore volume and different Al distribution in the *BEA framework when compared with conventional Na ion-exchanged TEA-directed zeolite beta (Si/ Al ¼ 7.6, Na/Al ¼ 0.53). OSDA-free beta showed reversible CO2 adsorptionedesorption behavior with higher CO2 adsorption capacity than the Na ion-exchanged TEA-directed beta across a wide pressure range, 1
105e3200 kPa, under moderate temperatures. Additionally, the pore filling of the CO2 adsorbate in the OSDA-free beta was higher at lower equivalent equilibrium pressures than that of the Na ion-exchanged TEA-directed beta. The overall change in the isosteric heat of adsorption for OSDAfree beta (21.8e39.8 kJ mol1) was lower than that of the Na ionexchanged TEA-directed beta (28.8e45.4 kJ mol1). Solid-state NMR analysis and the dual cation concept previously described, provide explanations as to the difference in the isosteric heat of adsorptiondattributed to the difference in the homogeneity of the Al (and counter cation). It is assumed that Al sites and single Na cations in OSDA-free beta are more homogeneously distributed when compared with the Na ion-exchanged TEA-directed beta, and a higher population of single Na cation sites effectively serves as an adsorption center for CO2 molecules. Consequently, the higher CO2 capacity in OSDA-free beta is mainly because of the larger pore volume, higher loadings of Na cation sites and the unique nature of the Al distribution (adsorption site) in the *BEA framework resulting from the synthesis route, as evidenced by the CO2 adsorptionedesorption, degree of Na ion-exchange, NMR measurements and evaluation of the isosteric heat of adsorption. These unique properties of OSDA-free beta significantly improved the CO2 adsorptionedesorption capacity over a wide

Y. Kamimura et al. / Microporous and Mesoporous Materials 219 (2016) 125e133

pressure range under moderate temperatures in comparison with the conventional zeolite beta. High quality OSDA-free beta obtained by a simple, low-cost and environmentally benign “green synthesis” will find applications not only in the capture of CO2 but generally across the fields of adsorbents and catalysis. Acknowledgments We gratefully acknowledge the Japan Society for the Promotion of Science (JSPS) for a Grant-in-Aid for Scientific Research (B)26820347. Appendix A. Supplementary data Supplementary data related to this article can be found at http:// dx.doi.org/10.1016/j.micromeso.2015.07.033. References [1] D.W. Breck, Zeolite Molecular Sieves, Wiley, New York, 1974. [2] R.M. Barrer, Hydrothermal Chemistry of Zeolites, Academic Press, London, 1982. [3] M.E. Davis, R.F. Lobo, Chem. Mater. 4 (1992) 756e768. [4] C.S. Cundy, P.A. Cox, Chem. Rev. 103 (2003) 663e702. [5] S.M. Csicsery, Pure Appl. Chem. 58 (1986) 841e856. [6] A. Corma, J. Catal. 216 (2003) 298e312. [7] S.C. Larsen, J. Phys. Chem. C 111 (2007) 18464e18474. [8] X. Meng, F.-S. Xiao, Chem. Rev. 114 (2014) 1521e1543. [9] K. Iyoki, K. Itabashi, T. Okubo, Micropor. Mesopor. Mater. 189 (2014) 22e30. [10] B. Xie, J. Song, L. Ren, Y. Ji, J. Li, F.-S. Xiao, Chem. Mater. 20 (2008) 4533e4553. [11] G. Majano, L. Delmotte, V. Valtchev, S. Mintova, Chem. Mater. 21 (2009) 4184e4191. [12] Y. Kamimura, W. Chaikittisilp, K. Itabashi, A. Shimojima, T. Okubo, Chem. Asian J. 5 (2010) 2182e2191. [13] Y. Kamimura, S. Tanahashi, K. Itabashi, A. Sugawara, T. Wakihara, A. Shimojima, T. Okubo, J. Phys. Chem. C 115 (2011) 744e750. [14] K. Iyoki, K. Itabashi, T. Okubo, Chem. Asian J. 8 (2013) 1419e1427. [15] T. Yokoi, M. Yoshioka, H. Imai, T. Tatsumi, Angew. Chem. Int. Ed. 48 (2009) 9884e9887. [16] M. Yoshioka, T. Yokoi, M. Liu, H. Imai, S. Inagaki, T. Tatsumi, Micropor. Mesopor. Mater. 153 (2012) 70e78. [17] K. Itabashi, Y. Kamimura, K. Iyoki, A. Shimojima, T. Okubo, J. Am. Chem. Soc. 134 (2012) 11542e11549. [18] K. Iyoki, Y. Kamimura, K. Itabashi, A. Shimojima, T. Okubo, Chem. Lett. 39 (2010) 730e731. [19] Y. Kamimura, K. Itabashi, T. Okubo, Micropor. Mesopor. Mater. 147 (2012) 149e156. [20] Y. Kamimura, K. Iyoki, S.P. Elangovan, K. Itabashi, A. Shimojima, T. Okubo, Micropor. Mesopor. Mater. 163 (2012) 282e290. [21] Y. Wang, X. Wang, Q. Wu, X. Meng, Y. Jin, X. Zhou, F.-S. Xiao, Catal. Today 226 (2014) 103e108. [22] A. Ogawa, K. Iyoki, Y. Kamimura, S.P. Elangovan, K. Itabashi, T. Okubo, Micropor. Mesopor. Mater. 186 (2014) 21e28. [23] Y. Kubota, K. Itabashi, S. Inagaki, Y. Nishita, R. Komatsu, Y. Tsuboi, S. Shinoda, T. Okubo, Chem. Mater. 26 (2014) 1250e1259. [24] Ch. Baerlocher, L.B. McCusker, D.H. Olson, Atlas of Zeolite Framework Types, sixth ed., Elsevier, Amsterdam, 2007 see also: http://www.iza-structure.org/ databases/. [25] M.M.J. Treacy, J.M. Newsam, Nature 332 (1988) 249e251. [26] J.B. Higgins, R.B. LaPierre, J.L. Schlenker, A.C. Rohrman, Zeolites 8 (1988) 446e452. [27] G. Bellussi, G. Pazzuconi, C. Perego, G. Girotti, G. Terzoni, J. Catal. 157 (1995) 227e234. [28] B. Wichterlova, J. Cejka, N. Zilkova, Micropor. Mater. 6 (1996) 405e414. [29] C. Perego, P. Ingallina, Catal. Today 73 (2002) 3e22.

133

[30] M. Ogura, T. Okubo, S.P. Elangovan, Catal. Lett. 118 (2007) 72e78. [31] M. Matsukata, M. Ogura, T. Osaki, P.R.H.P. Rao, M. Nomura, E. Kikuchi, Top. Catal. 9 (1999) 77e92. [32] P. Saint-Cricq, Y. Kamimura, K. Itabashi, A. Sugawara-Narutaki, A. Shimojima, T. Okubo, Eur. J. Inorg. Chem. 21 (2012) 3398e3402. [33] H. Zhang, L. Chu, Q. Xiao, L. Zhu, C. Yang, X. Menga, F.-S. Xiao, J. Mater. Chem. A 1 (2013) 3254e3257. [34] T. De Baerdemaeker, B. Yilmaz, U. Müller, M. Feyen, F.-S. Xiao, W. Zhang, T. Tatsumi, H. Gies, X. Bao, D. De Vos, J. Catal. 308 (2013) 73e81. [35] P. Sazama, B. Wichterlova, S. Sklenak, V.I. Parvulescu, N. Candu, G. Sadovska, J. Dedecek, P. Klein, V. Pashkova, P. Stastny, J. Catal. 318 (2014) 22e33. [36] M. Ogura, K. Itabashi, J. Dedecek, T. Onkawa, Y. Shimada, K. Kawakami, K. Onodera, S. Nakamura, T. Okubo, J. Catal. 315 (2014) 1e5. [37] Z. Huang, L. Xu, J.-H. Li, G.-M. Guo, Y. Wang, J. Chem. Eng. Data 55 (2010) 2123e2127. [38] S.-T. Yang, J. Kim, W.-S. Ahn, Micropor. Mesopor. Mater. 135 (2010) 90e94. [39] M.R. Delgado, C.O. Arean, Energy 36 (2011) 5286e5291. [40] Z. Bao, L. Yu, T. Dou, Y. Gong, Q. Zhang, Q. Ren, X. Lu, S. Deng, J. Chem. Eng. Data 56 (2011) 4017e4023. [41] H.-S. You, H. Jin, Y.-H. Mo, S.-E. Park, Mater. Lett. 108 (2013) 106e109. [42] J. Yang, J. Li, W. Wang, L. Li, J. Li, Ind. Eng. Chem. Res. 52 (2013) 17856e17864. [43] R. Srivastava, D. Srinivas, P. Ratnasamy, Appl. Catal. A Gen. 289 (2005) 128e134. [44] C. Li, X. Yuan, K. Fujimoto, Appl. Catal. A Gen. 475 (2014) 155e160. [45] M. Fujiwara, H. Sakurai, K. Shiokawa, Y. Iizuka, Catal. Today 242 (2015) 255e260. [46] D.M. D'Alessandro, B. Smit, J.R. Long, Angew. Chem. Int. Ed. 49 (2010) 6058e6082. [47] A. Samanta, A. Zhao, G.K.H. Shimizu, P. Sarkar, R. Gupta, Ind. Eng. Chem. Res. 51 (2012) 1438e1463. [48] T. Sakakura, J. Choi, H. Yasuda, Chem. Rev. 107 (2007) 2365e2387. [49] T. Sakakura, K. Kohno, Chem. Commun. (2009) 1312e1330. [50] D. Cazorla-Amoros, J. Alcaniz-Monge, A. Linares-Solano, Langmuir 12 (1996) 2820e2824. [51] R.C. Boggs, D.G. Howard, J.V. Smith, G.L. Klein, Am. Miner. 78 (1993) 822e826. [52] R.B. Borade, A. Clearfield, Micropor. Mater. 5 (1996) 289e297. [53] J. Dedecek, D. Kaucky, B. Wichterlova, O. Gonsiorova, Phys. Chem. Chem. Phys. 4 (2002) 5406e5413. [54] J. Dedecek, Z. Sobalik, B. Wichterlova, Catal. Rev. Sci. Eng. 54 (2012) 135e223. [55] S.M. Maier, A. Jentys, J.A. Lercher, J. Phys. Chem. C 115 (2011) 8005e8013. [56] J.A. van Bokhoven, D.C. Koningsberger, P.J. Kunkeler, H. van Bekkum, A.P.M. Kentgens, J. Am. Chem. Soc. 122 (2000) 12842e12847. [57] E. Bourgeat-Lami, P. Massiani, F. Di Renzo, P. Espiau, F. Fajula, T. Des Courires, Appl. Catal. 72 (1991) 139e152. rol, [58] K.S.W. Sing, D.H. Everett, R.A.W. Haul, L. Moscou, R.A. Pierotti, J. Rouque T. Siemieniewska, Pure Appl. Chem. 57 (1985) 603e619. [59] M.A. Camblor, A. Corma, S. Valencia, J. Mater. Chem. 8 (1998) 2137e2145. rez-Pariente, S. Valencia, Stud. Surf. [60] M.A. Camblor, A. Corma, A. Mifsud, J. Pe Sci. Catal. 105 (1997) 341e348. [61] T.D. Pham, Q. Liu, R.F. Lobo, Langmuir 29 (2013) 832e839. [62] M.M. Lozinska, J.P.S. Mowat, P.A. Wright, S.P. Thompson, J.L. Jorda, M. Palomino, S. Valencia, F. Rey, Chem. Mater. 26 (2014) 2052e2061. [63] Website of National Institute of Standards and Technology (NIST), http:// webbook.nist.gov/chemistry/fluid/. [64] F. Rouquerol, J. Rouquerol, K. Sing, Adsorption by Powders & Porous Solids, Academic Press, London, 1999. [65] K.S. Walton, M.B. Abney, M.D. LeVan, Micropor. Mesopor. Mater. 91 (2006) 78e84. rez-Pariente, A.B. Pinar, A. Zukalc, J. Cejka, Phys. [66] P. Nachtigall, L. Grajciar, J. Pe Chem. Chem. Phys. 14 (2012) 1117e1120. [67] J. Pires, M.B. de Carvalho, F.R. Ribeiro, E.G. Derouane, J. Mol. Catal. 85 (1993) 295e303. [68] P.J.E. Harlick, F.H. Tezel, Micropor. Mesopor. Mater. 76 (2004) 71e79. [69] P. Nachtigall, M.R. Delgado, D. Nachtigallova, C.O. Arean, Phys. Chem. Chem. Phys. 14 (2012) 1552e1569. [70] L. Grajciar, J. Cejka, A. Zukal, C.O. Arean, G.T. Palomino, P. Nachtigall, ChemSusChem 5 (2012) 2011e2022. [71] Y. Kalvachev, M. Jaber, V. Mavrodinova, L. Dimitrov, D. Nihtianova, V. Valtchev, Micropor. Mesopor. Mater. 177 (2013) 127e134.