CH4 on the synthesized NaA zeolite shaped with montmorillonite clay in natural gas purification process

CH4 on the synthesized NaA zeolite shaped with montmorillonite clay in natural gas purification process

Accepted Manuscript Adsorption separation of CO2/CH4 on the synthesized NaA zeolite shaped with montmorillonite clay in natural gas purification proce...

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Accepted Manuscript Adsorption separation of CO2/CH4 on the synthesized NaA zeolite shaped with montmorillonite clay in natural gas purification process A. Arefi Pour, S. Sharifnia, R. NeishaboriSalehi, M. Ghodrati PII:

S1875-5100(16)30810-1

DOI:

10.1016/j.jngse.2016.11.006

Reference:

JNGSE 1916

To appear in:

Journal of Natural Gas Science and Engineering

Received Date: 29 December 2015 Revised Date:

28 October 2016

Accepted Date: 3 November 2016

Please cite this article as: Pour, A.A., Sharifnia, S., NeishaboriSalehi, R., Ghodrati, M., Adsorption separation of CO2/CH4 on the synthesized NaA zeolite shaped with montmorillonite clay in natural gas purification process, Journal of Natural Gas Science & Engineering (2016), doi: 10.1016/ j.jngse.2016.11.006. This is a PDF file of an unedited manuscript that has been accepted for publication. As a service to our customers we are providing this early version of the manuscript. The manuscript will undergo copyediting, typesetting, and review of the resulting proof before it is published in its final form. Please note that during the production process errors may be discovered which could affect the content, and all legal disclaimers that apply to the journal pertain.

ACCEPTED MANUSCRIPT

Adsorption separation of CO2/CH4 on the synthesized NaA zeolite shaped with montmorillonite clay in natural gas purification process A. Arefi Pour, S. Sharifniaa,*, R. NeishaboriSalehi, M. Ghodrati

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Catal. Res. Cen., Dept. Chem. Eng., Razi Univ., Kermanshah 67149-67246, Iran.

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Abstract

In this study, the effects of parameters such as crystallization time and temperature on the

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phase and morphology of NaA zeolite prepared by hydrothermal method were investigated. The textural properties of the synthesized zeolites were characterized using XRD, SEM, FTIR and BET analyses. The findings revealed that the optimum time and temperature of NaA preparation to be 20 h and 363 K, respectively. The adsorption experiments of CO2 and CH4 by synthesized and commercial NaA zeolites were carried out at three temperatures (277, 290 and 310 K) for pressures up to 10 bar. The experimental data were analyzed using

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Langmuir and Sips equations, and the Sips model better described the experimental data. It was found that the synthesized NaA had BET specific surface area of 222.8 m2g-1 and pore volume of 0.096 cm3g-1. Also, the adsorption capacities of 5.2 and 2.6 mmolg-1 were obtained for CO2 and CH4 by the synthesized samples, respectively. Comparing with commercial

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sample, the synthesized zeolite showed higher adsorption capacities for CO2 and CH4 under the same condition. The heat of adsorption at zero coverage of CO2 and CH4 on the

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synthesized NaA zeolite were found to be 48.5 and 24.4 kJmol-1, respectively. The ideal CO2/CH4 selectivity of the synthesized and commercial zeolites were 7.1 and 6.4, respectively, at atmospheric pressure and 290 K. The dynamic adsorption experiments of CO2/CH4 gas mixture confirmed that the synthesized NaA zeolite had more promising performance for the separation of CO2/CH4 compared to the commercial zeolite.

Keywords: NaA zeolite, hydrothermal synthesis, CO2 adsorption, natural gas purification. *

Corresponding author: Tel.:+98 83 34274530; Fax: +98 83 34274542; Email: [email protected]

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ACCEPTED MANUSCRIPT 1. Introduction Nowadays, natural gas is one of the most applicable and cleanest fuel in comparison with other fossil fuels such as oil and coal (Avila et al., 2011). Natural gas is a gaseous mixture of different hydrocarbons such as ethane, propane (but it consists mainly of methane), and

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somewhat impurities like carbon dioxide, nitrogen, hydrogen sulfide, water, and very trace amounts of helium and mercury (Tagliabue et al., 2012; Tagliabue et al., 2009). Carbon dioxide is one of the major contaminants in natural gas, which besides reducing the heating

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value of natural gas, it also, in the presence of water and other contaminants, can turn into corrosive gases and damage to pipelines. The heating value of clean natural gas is about 11

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kWh/m3, while it is 6.5 kWh/m3 for natural gas containing carbon dioxide. These may be the main reasons that why carbon dioxide removal from natural gas has become a very serious issue (Doroudian Rad et al., 2012; Li et al., 2013; Gholipour and Mofarahi, 2016). Adsorption, cryogenic distillation, and membrane separation are technologies that have

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been used for purification and removal of CO2 from gas streams (Arefipour et al., 2015; Doroudian Rad et al., 2012; Hyun et al., 1999; Mulgundmath et al., 2012; Munusamy et al., 2012; Yu et al., 2013a, 2013b). The conventional technique to CO2 capture from natural gas

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is amine absorption processes, but the energy required to regenerate these processes

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dramatically is high, compared to the adsorption process (Jiang et al., 2013; Rufford et al., 2012). Adsorption process is the result of transferring some of the molecules of gas or liquid phases to surface of adsorbent. In this process, adsorption takes place based on the affinity between fluid phase and adsorbent surface. In other word, one species of the fluid phase can be adsorbed that has a higher affinity towards the adsorbent (Kulprathipanja, 2010; Shirani et al., 2010). In the adsorption process, high adsorption capacity and selectivity are two key parameters for effective removal of carbon dioxide from gas stream (Doroudian Rad et al., 2012; Shao et al., 2009). Metal organic frameworks (MOFs), molecular sieves like zeolites,

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ACCEPTED MANUSCRIPT activate carbons, clay (Araki et al., 2012; Li et al., 2013), limestone, lithium zirconate, lithium silicate, and amine-modified mesoporous silica are the most common CO2-adsorbent (Araki et al., 2012; Shen et al., 2010). Zeolites are natural or synthetic compounds, having hydrated alumina-silica structures of

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alkaline and alkaline-earth metals, along with frameworks made of SiO4 and AlO4 tetrahedral, bridging together by sharing oxygen atoms. Incorporation of Al3+ in place of Si4+ creates a negative charge on the lattice, which is compensated by extra-framework cations such as K+,

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Na+, Mg2+, Sr2+ and Ba2+. Extra-framework cations are enough mobile, so that they are easily replaced with other cations (Elaiopoulos et al., 2010; Gougazeh and Buhl, 2014;

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Kulprathipanja, 2010; Melo et al., 2012; Rožic et al., 2009; Ugal et al., 2010). The framework structure of zeolite consists intra-crystalline channels, which may be one, two or threedimensional (Kulprathipanja, 2010). Compared to natural zeolites, the synthetic zeolites such as X, Y, and A are often more applicable because of high purity and uniform particle size

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(Gougazeh and Buhl, 2014). The adsorption properties of zeolites are dependent on several factors such as structure, number and nature of extra framework cations, the silicon to aluminum ratio, and pore size (Araki et al., 2012; Sayari et al., 2011). The pore size is one of

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the most effective factors in gas separation processes by microporous adsorbents. When the

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pore diameter of adsorbent is between the diameters of two target compounds, the molecular sieve effect (steric effect) happened and the smaller compound can be separated. Being the pore size of adsorbent just a little larger than the diameters of two target compounds leads to kinetic separation of them. Kinetic separation is based on differences in the diffusion rates. In this case, the smaller component diffuses faster than the larger component. At sufficiently large pore size, both components can easily diffuse inside the pores and separating takes place through difference in adsorption equilibrium (equilibrium selectivity) (Gholipour and Mofarahi, 2016). Researches show that zeolites such as LTA (Linde type A), 13X, MOR, and

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ACCEPTED MANUSCRIPT MFI have been frequently used in the adsorption process of carbon dioxide (Mulgundmath et al., 2012; Montanari et al., 2011; You et al., 2013). A-type zeolites have Na12[(AlO2)12(SiO2)12].27H2O chemical formula, and their hydrated forms contain approximately 20% water (Ugal et al., 2010; Yuan et al., 2011). The LTA-type framework

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can be formed by connecting sodalite cages via double four-rings, making an alpha cavity in the center of the unit cell (Kulprathipanja, 2010). The unit cell of A zeolite consists of 24 tetrahedral, 12 AlO4 also 12 SiO4 (Yang, 2003). This zeolite is an aluminum rich zeolite

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which possesses a Si/Al ratio around unity. The synthetic A zeolites are categorized into three types, KA (3A), NaA (4A), and CaA (5A), all of which have the same chemical formula and

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the pore sizes of 3, 4, and 5Å, respectively (Auerbach et al., 2003; Gougazeh and Buhl, 2014; Ugal et al., 2010). This zeolite primarily has been used as adsorbent for purification of gases such as natural gas, nitrogen, methane, and air, or as ion exchanger in detergent industries, or catalyst in petroleum refining and petrochemistry (Loiola et al., 2012; Montanari et al., 2008).

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Zeolites are commonly synthesized via hydrothermal method, from a reactive mixture consists of an alumina source and a silica source or other alumosilicates such as industrial residue, bentonite, clay, kaolin, fly ash, halloysite, and rice hunk, and mineralizing agent of

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OH-, at temperature of 300-473 K and different times from a few hours up to several days

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(Anuwattanaand Khummongkol, 2009; Auerbach et al., 2003; Ma et al., 2014; Melo et al., 2012). Many studies have been performed on zeolites for separation of CO2 from CH4 in the adsorption process. Table 1 lists the results of some researches dealing with adsorption isotherms for pure CO2, CH4, and N2 on different zeolites at different pressures and temperatures. The pure and binary adsorption equilibrium data for CO2 and CH4 on zeolite 13X at different temperatures and pressures up to 10 bar were reported in the literature (Gholipour and Mofarahi, 2016). It was found that the experimental selectivity of CO2/CH4 changes between 26 and 2 at different pressure and temperatures. In another study, the binary

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ACCEPTED MANUSCRIPT adsorption CO2 and CH4 in binderless beads of 13X zeolite was examined at 313, 373 and 423 K and total pressure up to 5 atm. The extended Fowler isotherm was used to predict binary adsorption equilibrium (Silva et al., 2014). The results revealed the experimental selectivity of CO2/CH4 from 37 at pressure of 0.667 atm to ~5 at temperature of 423 K. Since

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there are no studies on adsorption separation of CO2/CH4 gas mixture by using synthesized zeolite shaped with montmorillonite binder, the purpose of this study was: (I) the hydrothermal synthesis of NaA zeolite with high purity, surface area and crystallinity as a

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proper adsorbent for CO2/CH4 separation, (II) a survey on the effects of crystallization time and temperature on the crystalline products, (III) attaining the best product according to the

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crystallinity and shaping with montmorillonite, and (IV) comparing adsorption performances of the shaped synthesized and commercial NaA zeolites for CO2 and CH4 at different pressure

2. Experimental

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and temperature.

The spherical beads form commercial NaA zeolite with an average diameter of 3 mm was

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purchased from Zeochem Co., Switzerland. The commercial NaA zeolite is referred to as

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NaA-Zeo. Montmorillonite K-10 clay as binder was supplied from Fluka.

2.1. Synthesis of NaA zeolite Sodium hydroxide (NaOH, 98 wt%, Merck), sodium aluminate (54 wt% Al2O3,41 wt% Na2O; impurities ≤0.05% Fe (as Fe2O3) Sigma-Aldrich), sodium metasilicate pentahydrate (Na2SiO3+5H2O, assay ≥95.0% (T) Sigma-Aldrich) and deionized water were used as starting materials in the hydrothermal synthesis of NaA zeolite. The methodology used to synthesize NaA zeolite was based on the research of Thompson and Huber (Robson, 2001), with changes in the crystallization time and temperature. Final hydrogel had a molar ratio of 3.165 Na2O: 5

ACCEPTED MANUSCRIPT Al2O3: 1.926 SiO2: 128 H2O. This gel is obtained by combining aluminate and silicate solutions. In the present study, 0.723 g sodium hydroxide was added to 80 mL deionized water and stirred until sodium hydroxide was fully dissolved. This solution was divided into two equal parts. Silicate and aluminate solutions were prepared by adding 15.480 g sodium

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silicate to the one part of the solution and 8.258 g sodium aluminate to the second part of the sodium hydroxide solution, respectively. Then, both of these solutions were mixed and stirred until a homogeneous gel was obtained. After 3 h aging at 333 K, the resulting gel was

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transferred into a sealed Teflon-lined autoclave, and crystallized in an oven at 363, 373, and 393 K for different periods of time from 8 to 24 h. After crystallization, the resulting products

products were dried at 473 K for 3 h.

2.2. Characterization

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were filtered and washed with deionized water until the solution pH dropped to 7. Finally, the

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This work was focused on the synthesis of NaA zeolites with the highest purity and crystallinity, and BET surface area to separate the CO2 from CH4. The reason for this can be explained in such a way that in crystalline material is present more sites and higher

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availability to adsorb the gas molecules and also specific surface area is directly related to the

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number of gas molecules that adsorbs onto an adsorbent surface (Mohamed et al. 2009; Ismail et al., 2010). Therefore, the phase identification and crystallinity of synthesized products were determined using X-ray diffraction technique. The XRD patterns of the productswere obtained using EQUINOX diffractometer (Inel Company) with CuKα radiation (λ=1.5406 Å), operated at 20 mA and 30 kV. The diffraction patterns were recorded in the 2θ range of 5-40o with 0.05 step size. The relative crystallinity of synthetic zeolites was determined from the ratio of the sum of height of the characteristic peaks of NaA zeolite in the 2θ range of 5-40°, to a fully crystalline of NaA zeolite as standard sample (Liu et al.,

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ACCEPTED MANUSCRIPT 2013). The morphologies of synthesized samples were investigated by scanning electron microscopy (SEM) Model XL-30 Philips equipment, operated at 10 kV. The nitrogen adsorption and desorption of samples at 77.3 K were obtained using a Quantachrome Nova instrument. In order to determine the textural properties of samples at first, approximately

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0.08 g of samples were degassed at 573 K for 6 h. The pore textural properties such as specific surface area was calculated via the BET method, the external surface area and the micropore volume were measured by the t-plot method and finally, the total pore volume was

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obtained from the amount of nitrogen adsorbed at (P/Po)=0.99 (Dey et al., 2013; Salama et al., 2009). The BET surface area and pore volume of NaA-Zeo were 190.2 m2/g and 0.085

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cm3/g, respectively. The pore size distribution of synthesized samples was determined from the desorption branch using the Barrett–Joyner–Halenda (BJH) model.Infrared spectra (FTIR) of samples were conducted by using a Bruker ALPHA-FTIR spectrometer in the 4004000 cm-1 wave number range.The chemical analyses of the loaded adsorbents in the

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adsorption column were determined by X-ray fluorescence technique (XRF; SPECTRO XLabPro) are reported in Table 2. The SiO2/Al2O3 ratio of NaA-Zeo and shaped synthesized zeolite with montmorillonite(referred to as NaA-S), were 2 and 2.16, respectively, which were

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agreement with a SiO2/Al2O3 ratio reported for NaA zeolite (SiO2/Al2O3=2).

2.3. Adsorption measurements and breakthrough experiments To evaluate the adsorption performances of the synthesized NaA zeolite, the montmorillonite binder was added with synthesized NaA zeolite powder (80 wt.% synthesized NaA zeolite powder and 20 wt.% montmorillonite binder), then the resulting mixture shaped into granules with a granule size of 3 mm. Finally, the products were dried in an oven at 573 K for 3 h and then the obtained adsorbent loaded in the adsorption column. Adsorption experiments were conducted in the apparatus displayed in Fig. 1. It contains three

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ACCEPTED MANUSCRIPT sections: i) Gas cylinder preparation; ii) an adsorption column; and ii) a Gas Chromatograph with a TCD detector and helium carrier gas. To prepare the gas cylinder first the air removed to create a vacuum. Then, the cylinder was filled with feed gas containing 60% methane, 20% carbon dioxide, and 20% helium to a pressure of 10 bar. The cylinder was equipped with a

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pressure gauge to measure gas pressure. The apparatus was checked to be leak-proof. The adsorption column was constructed from a stainless-steel tube, with an inner diameter and length of 9 mm and 62 cm, respectively. To minimize dead volume and dilution of the bed,

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after loading the adsorbent in the middle of the column, height equal 8 cm from the top and bottom of the adsorption column was filled with glass beads and adsorbent was placed in the

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middle column (Doroudian Rad et al., 2012; Munusamy et al., 2012). The column was equipped with an electric heating jacket to accurate temperature control of the adsorption column. Temperature was measured using a thermocouple embedded in the middle of the column. A pressure transducer was used for measurement of pressure in the adsorption

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column. The flow rate of the feed gas was controlled using a mass flow controller. Adsorption isotherms were obtained at three temperatures (277, 290 and 310 K) for pressures up to10 bar. Adsorption experiments at 277 K were conducted with the help of a water bath

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connected to refrigerated circulating bath, also temperature at 290 and 310 K was kept

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constant by a circulating water bath with a constant temperature. Before every measurement, the samples were outgassed with helium at a flow rate of 10 mL/min at 673 K for 12 h, and then cooled to ambient temperature. After adjusting the temperature of the adsorption column (277, 290 and 310 K at constant pressure of 10 bar for run 1), the adsorption experiment was initiated by opening the valves between the cylinder and adsorption column. The equilibrium state was reached when the pressure of the adsorption column remained at the constant level (the pressure change was less than 0.1%) within 5 min. The concentration of the outlet gas stream from the adsorption column was analyzed by using a gas chromatograph (Schroth

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ACCEPTED MANUSCRIPT Compact GC-CGA-1, 230 V, 700 W) equipped with a thermal conductivity detector (TCD). By doing breakthrough tests under gas mixture (CH4/CO2/He, 60:20:20 V/V/V) the separation performance of the adsorbents was studied. In a number of researches to analyze the dynamic performance of adsorbates on adsorbent in a fixed-bed column, kinetic studies

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were performed by breakthrough experiments (Li et al., 2004; Mosca et al., 2010; Sigrist et al., 2011; Yi et al. 2012). In this study, breakthrough experiments were carried out in the packed column of NaA-Zeo and NaA-S samples at 290 K and constant pressure of adsorption

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column at 1.0 bar with 10 mL/min total flow rate. Similarly, concentration at the bed exit was measured using GC. The gases (pure) used in the adsorption experiments were CO2

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(99.99%), CH4 (99.95%), and He (99.999%).

2.4. Theory and method

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2.4.1. Isotherm equations

Many models have been presented to describe the adsorption equilibrium data, such as, Sips, Toth, Langmuir, and UNILAN. In this study, Langmuir and Sips models were

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employed to fit equilibrium data. The Langmuir model is suitable to describe the monolayer adsorption on ideal surfaces. The Langmuir and Sips equations are presented as Eq.’s (1) and

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(2), where P(kPa) and q(mmolg-1) are

pressure and adsorption amount in natural gas

stream,respectively.qm is the maximum adsorption capacity, K and b are the affinity constant of the Sips and Langmuir models, respectively, and n in the Sips model is the heterogeneity parameter of a system (Deng et al., 2012; Do, 1998; Shao et al., 2009).

q=

qm bp 1 + bp

(1)

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ACCEPTED MANUSCRIPT qm Kp n q= 1 + Kp n

(2)

The parameters of Langmuir and Sips models, and correlation coefficients were obtained

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with MATLAB 7.0.The reason to choose MATLAB, in this work, is various applications of this software in the adsorption processes (Cavenati et al., 2004; Huang et al., 2010; Yi et al., 2012).

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2.4.2. Henry’s law constant and selectivity

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Henry's law constant, determining the affinity of adsorbate molecules toward adsorbent is correlated with the interaction between adsorbate molecules and adsorbent surface. So, accurate calculation of Henry's constant is important in the adsorption process. Henry's constant was obtained by using the Langmuir equation (Deng et al., 2012; Ridha and Webley, 2009).

linear form defined by:

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q lim   = bq m = K H  p

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The Langmuir equation at very low pressures, i.e., in Henry's law region, is reduced to

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p → 0

(3)

Where KH is Henry’s law constant. In the Langmuir equation, values of b and qm are obtained from a plot of (1/q) against (1/P). The relationship between temperature and Henry's constant is described by van’t Hoff equation, given as Eq. (4).

K H = K Ho exp(

− ∆H 0 ) RT

(4)

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ACCEPTED MANUSCRIPT In Eq. (4), KHo and

H0 (J.mol-1) are the parameter of the van’t Hoff equation and the heat of

adsorption, respectively.T (K) is the absolute temperature, and R is the gas constant (Pakseresht et al., 2002). Heat of asdsorption and pre-exponential factor of the van’t Hoff equation were determined from the temperature dependence of the Henry adsorption

temperature, the following expression is used.

K Hi K Hj

(5)

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S ei, j =

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constants (Denayer et al., 2008). In order to calculate equilibrium selectivity at a specific

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Where KHi and KHj are the Henry's constant of gases i and j, respectively (Deng et al., 2012).The ideal selectivity of gas i over gas j is determined using Eq. (6) (Doroudian Rad et

Si , j =

qi pi qj pj

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3. Results and discussion

(6)

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al., 2012).

3.1. Synthesis of NaA zeolite and characterization

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3.1.1. Effect of crystallization time Fig. 2 shows the XRD patterns of the synthesized samples at 363 K and different crystallization times of 8, 15, 20, and 24 h. The XRD patterns of all synthesized samples confirmed the formation of crystalline structure of NaA zeolite. It can be seen that by increasing the crystallization time from 8 to 24 h, the characteristic peaks of the prepared NaA zeolite became more intense, suggesting that the crystallization time of 8 h was not sufficient for complete crystallization of NaA zeolite. Table 3 gives information about the

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ACCEPTED MANUSCRIPT effects of crystallization condition on phase purity and crystallinity of NaA zeolite. As shown in Fig. 2 and Table 3, increasing crystallization time from 8 to 15 and 20 h led to increasing crystallinity of the final products, because of the increased number of nucleuses with increasing crystallization time. With further increase in crystallization time from 20 to 24 h,

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very weak peak of NaP zeolite as an impurity phase was appeared (see Fig. 2). So, the crystallization of NaA zeolite was completed within 20 h and longer crystallization time resulted in formation of NaP zeolite. The decreasing diffraction intensity of NaA zeolite with

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prolonging treatment time to 24 h has been reported by other researchers (Zhang et al., 2013). The NaA zeolite needs to a short crystallization time because of simple and small polymeric

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silicate unites (D4R) (Purnomo et al., 2012; Nibouet al., 2011).

The effect of crystallization time on the particle size and morphology of the synthesized samples are shown in Fig. 3. By 8 h crystallization time, the sample had amorphous nature with irregular shaped particles and very small cubic particles with size diameter below 100

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nm (Fig. 3(a)). This indicates that only a few zeolite particles had the nucleation opportunity, and crystallization process was incompletely done. With prolonging crystallization time to 15 h, the cubic-shaped crystals of NaA zeolite with beveled edges and uniform size of 300 nm

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were formed (Fig. 3(b)). By raising crystallization time to 20 h, the cubic-shaped crystals

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with sharp edges and mean crystal size around 400 nm were obtained (Fig. 3(c)). As shown in Fig. 3(d), further increasing crystallization time resulted in the formation of particles with different shapes and sizes. By comparing Fig. 3(b) and Fig. 3(c), it can be concluded that increasing in the crystallization time from 15 to 20 h led to growth of larger particle size. Results obtained from SEM images were confirmed by the XRD findings. As a result, it could be concluded that the time of 20 h was optimal for the preparation of NaA zeolite. FTIR spectra provide useful information to identify the structure of zeolites (Elaiopoulos et al., 2010). Fig. 4 shows FTIR spectra of samples crystallized from 8 to 24 h. The band

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ACCEPTED MANUSCRIPT appeared at 467 cm-1 is attributed to the internal vibration of (Si, Al)-O bending. The characteristic band at the 557 cm-1 is related to the external vibration of the double four rings (D4R) in the structure of NaA zeolite. By increasing the crystallization time from 8 to 15 and 20 h, the bands at 668 and 748 cm-1 were appeared, which are assigned to the internal

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vibration of (Si, Al)-O and Si-O-Al symmetric stretching, respectively. Another band at around 1002 cm-1 is ascribed to the internal vibration of (Si, Al)-O asymmetric stretching (Loiola et al., 2012; Nibou et al., 2011). With further increase in synthesis time to 24 h, these

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peaks were disappeared and a band at 701 cm-1 was detected, assigning to the symmetric stretch vibration of internal tetrahedron (Huo et al., 2012). The single bands observed at 3450

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and 1650 cm-1 correspond to the existence of (OH) groups and H2O, respectively. The presence of water in the zeolite structure shows that the zeolite had not been fully dehydrated (Zhao et al., 2010).

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3.1.2. Effect of crystallization temperature

To evaluate the effect of crystallization temperature on synthesis of NaA zeolite, the samples were prepared at different temperatures from 363 to 393 K for 20 h. Fig. 5 shows the

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XRD patterns of the synthesized samples at different temperatures. As shown in Fig. 3 and

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Table 3, changing crystallization temperature affected the crystallinity of the final products. By increasing crystallization temperature from 363 to 373 K, the characteristic peaks of pure phase of NaA zeolite was still detected in the XRD patterns without any other impurity phases, but relative crystallinity decreased (see Table 3). As previously mentioned, the NaA zeolite needs to a short synthesis time. In this study, the crystallization time was long enough, so that raising the temperature was not favorable for NaA zeolite preparation, and increasing this parameter provided the conditions for synthesizing other zeolite phases. The XRD analysis of the synthesized sample at 393 K confirmed the partially formation of hydroxyl

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ACCEPTED MANUSCRIPT sodalite. Therefore, the optimal crystallization temperature for preparation of NaA zeolite was 363 K. Fig. 6 illustrates the nitrogen adsorption and desorption isotherms and pore size distribution of the samples crystallized at different temperatures (363, 373 and 393 K). All

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three samples exhibited isotherm shape as H3 loops type, according to the IUPAC classification (Rouquerol et al., 1999). For all the synthesized samples in the low relative pressure region about (P/Po)<0.02, a steep nitrogen uptake appeared, which is assigned to the

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filling of micropores (Fig. 6(a)). Also, the hysteresis loops (0.15-0.9) emerged in all of the samples are ascribed to the filling of slit-like shaped mesopores. In other words, hysteresis

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loop is related to the filling and depleting of mesopores. Mesopores with slit-like geometry are created during the formation of zeolite crystals (Rouquerol et al., 1999). Moreover, next uptake nitrogen for all samples at a relative pressure (0.9-1) corresponds to the filling of macropores. The sample synthesized at 363 K due to the high crystallinity had maximum

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nitrogen adsorption capacity, in comparison with other samples. Fig. 6(b) shows pore size distribution of synthesized samples at different temperatures. All samples demonstrate one mesopore at about 3.6 nm and a larger mesopore in the range of 10-12 nm. So the major part

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of the pore volume can be attributed to the mesopores. In addition, the sample synthesized at

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393 K shows the highest mesoporosity. The textural characteristics of products are presented in Table 4. It can be concluded that for NaA zeolite, with increasing crystallization temperature from 363 to 393 K, the BET surface area, external surface area, micropore volume, and total pore volume were diminished. This could be attributed to the variation of crystal morphology and lower crystallinity of NaA zeolites synthesized at temperatures 373 and 393 K as indicated by XRD studies. The maximum BET surface area belonged to the sample crystallized at 363 K, because of its high crystallinity. The BET surface area and pore volume of the best product were 222.8 m2/g and 0.096 cm3/g, respectively (Table 4).

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ACCEPTED MANUSCRIPT Finally, according to the results of XRD, SEM and BET analyses, the sample crystallized at 363 K for 20 h was considered as the best product for adsorption experiments, due to the

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high crystallinity, high surface area, and uniform particle size.

3.2. Adsorption isotherms

The adsorption isotherms of CO2 and CH4 on NaA-Zeo and NaA-S zeolites were obtained at 277, 290 and 310 K at pressures up to 10 bar. The CO2 and CH4 adsorption isotherms on

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NaA-Zeo and NaA-S zeolites are plotted in Fig. 7 and Fig. 8, respectively. Based on the

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IUPAC classification, all isotherms are categorized as Type-I (Rouquerol et al., 1999). NaA zeolite has 8-oxygen ring windows with the pore size 4.1 Å× 4.1 Å and 12 Na+ ions in the pseudo cell. Major cation sites in the NaA zeolite structure are site I, site II, and site III, which can accommodate 8, 3 and 1 cations, respectively (Kulprathipanja, 2010). CO2 is a polar molecule, which has the quadrupole moment, polarizability, and the kinetic diameter of

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4.30 ×10-26 esu.cm2, 26.5 ×10-25 cm3, and 3.3Å , respectively. CH4 is a non-polar molecule with polarizability and the kinetic diameter of 26.0 ×10-25 esu.cm3, and 3.8Å, respectively

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(Sawant et al., 2012). In this study, the pore size of NaA zeolite is a little larger than the kinetic diameter of the CH4 molecules. Thus, it can be concluded that the equilibrium

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selectivity didn’t limit the separation process and the molecular sieve effect was the dominant. It can be clearly observed that for both natural and synthetic adsorbents, the adsorption capacity of CO2 was considerably higher than that of CH4. The difference in adsorption capacity of CO2 and CH4 is due to differences in electrical properties and molecular size of them. Other than being more polar and having quadrupole moment, CO2 has also smaller molecular size than CH4 (Mofarahi and Gholipour, 2014). Also, at low-pressure regime, the slope of the adsorption capacity of CO2 was much steep, whereas the CH4 adsorption isotherms exhibited an almost linear behavior. This indicates that 15

ACCEPTED MANUSCRIPT CO2 was adsorbed markedly stronger than CH4 on the zeolites surface, and adsorbents showed less affinity for the adsorption of CH4. It is worth mentioning that NaA zeolite has the highest cations in its structure, creating an electric field. Thus, molecules like CO2, which possess a quadrupole moment, could strongly interact with adsorbent surface. In contrast,

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CH4, which is a non-polar molecule, may be adsorbed on the adsorbent surface only at sufficiently high pressure (Bao et al., 2011; Doroudian Rad et al., 2012; Mofarahi and Gholipour, 2014). Physical adsorption is always an exothermic phenomenon, so with

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increasing temperature from 277 to 310 K, adsorption capacity of CO2 decreased. Considering the direct relationship between temperature and kinetic energy of the molecules,

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this behavior can be explained. As a result of raising temperature during the adsorption process, the excess thermal energy makes the desorption process faster, allowing the system to reach equilibrium quickly (Mulgundmath et al., 2012; Munusamy et al., 2012). Also, it can be observed from Fig. 7 and Fig. 8 that the adsorption capacity of CO2 onto NaA-S zeolite

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was higher than that of NaA-Zeo zeolite. At 290 K and 10 bar, the adsorption capacity of CO2 on NaA-S zeolite was ~5.9 mmol/g, almost 10% higher than the CO2 adsorption capacity of NaA-Zeo. This is due to the larger surface area and pore volume of NaA-S zeolite. The BET

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surface area of the synthesized NaA zeolite was ca. 17% higher than the BET surface area of

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commercial NaA zeolite. In addition, sodium is the main cation in the structure of both adsorbents and the maximum amount of sodium belongs to the synthesized sample (Table 2). According to the literature (Siriwardane et al.,2003), carbon dioxide molecules prefer to be adsorbed on the zeolite with higher sodium content. Because, the presence of more cations in the structure of zeolite causes more interesting heterogeneous surface for adsorption of CO2 with quadrupole moment (Silva et al., 2012; Li and Tezel, 2007). In the next step, experimental data were fitted using Langmuir and Sips models. The parameters of Langmuir and Sips models for NaA-Zeo and NaA-S zeolites are listed in Table 5 and Table 6,

16

ACCEPTED MANUSCRIPT respectively. As it was stated earlier, the b and K parameters are the affinity constant of the Langmuir and Sips equations, respectively. These parameters express how strong the adsorbate molecules are adsorbed on the adsorbent surface. Whatever these parameters be larger, the adsorbate molecules have a stronger affinity towards the adsorbent. As can be seen

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in Table 5 and Table 6 that b and K values decreased with increasing temperature. This demonstrates that increasing temperature leads to adsorption of fewer molecules by the surface, and finally, a weaker affinity between adsorbent and adsorbate molecules emerges.

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The highest b and K values are related to the adsorption of carbon dioxide on the NaA-S zeolite, showing that carbon dioxide had a stronger affinity towards the NaA-S zeolite surface

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in comparison with methane. This can be assigned to the presence of higher sodium content of NaA-S sample relevant to NaA-Zeo. In the Sips equation, the n parameter represents the heterogeneity degree of system. When n gets larger, the system becomes more heterogeneous. Similarly, with increasing temperature from 277 to 310 K, the n values or the heterogeneity of

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the system decreased. It can be concluded that the system was less heterogeneous at higher

temperatures. As can be seen in Table 5 and Table 6, the n values for CO2 were higher than CH4 on both adsorbents at the same conditions, reflecting the more heterogeneity CO2/adsorbents

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system than that of CH4/adsorbents system that is associated with quadrupole moment of

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CO2. The values of this parameter for CO2/NaA-S system were higher than those of CO2/NaA-Zeo system, may be due to more sodium content of NaA-S. In addition, the values of qm in both models were reduced with increasing temperature. This implies that at low temperature, the adsorbate loading on the surface increased (Do, 1998; Doroudian Rad et al., 2012). At all temperatures, NaA-S zeolite has the higher qm values when compared to NaA-S zeolite may be due to the higher specific surface area of the synthesized adsorbent. The specific surface area is directly related to the number of gas molecules that adsorbs onto an adsorbent. Higher specific surface area provides more sites and higher availability to adsorb

17

ACCEPTED MANUSCRIPT the gas molecules (Mohamed et al. 2009; Ismail et al., 2010). The correlation coefficients of the Langmuir and Sips equations for CO2 and CH4 are given in Table 5 and Table 6. The values of this parameter for Sips model were higher than Langmuir model, suggesting the best fit of Sips model to experimental data. The Sips model with three adjustable parameters

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is able to describe heterogeneity of a system, hence it could better fit the data, as expected (Do, 1998).

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3.3. Henry’s constant and selectivity

The values of Henry’s constant for CO2 and CH4 were determined by using Langmuir

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model. This constant represents the degree of interaction between the adsorbate and adsorbent. Whatever Henry’s constant be greater, it is proved that there is a stronger affinity between the adsorbate and adsorbent surface (Deng et al., 2012). Henry’s constant of CO2 and CH4 and equilibrium selectivity of CO2/CH4 on NaA-Zeo and NaA-S zeolites are

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summarized in Table 7. It can be seen that at all temperatures, Henry’s constant of CO2 was higher than that of CH4 in both adsorbents. As explained previously, this is due to the quadrupole moment of carbon dioxide and non-polar nature of methane. CO2 had a stronger

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affinity towards the adsorbents surface, hence the Henry’s constant of CO2 was larger. At all

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temperatures, the value of the Henry’s constant was higher for CO2 on NaA-S zeolite, because of the presence of more cations in the structure of synthesized adsorbent. Also, from Table 7, it can be clearly seen that by raising temperature, the Henry’s constant of CO2 and CH4 decreased. The origin of this behavior may be the exothermic nature of the adsorption phenomenon (Li and Tezel, 2007). The capability of adsorbent for gas separation is related to the equilibrium selectivity, originating from the differences in affinity of adsorbent for the adsorption of target components (Tagliabue et al., 2009). In the separation processes, equilibrium selectivity is an

18

ACCEPTED MANUSCRIPT important factor for evaluating the adsorbent separation ability. In this study, the ratio of Henry’s constants of CO2 and CH4 at zero adsorption coverage was calculated to investigate the capability of NaA-S and NaA-Zeo zeolites for separation of CO2 and CH4 mixture (S ei, j ) . Since the isotherms of CO2 were steep on both adsorbents, thus Langmuir model cannot

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conveniently predict the experimental adsorption data in the all pressures studied. According to this explanations, to correlate isotherms of CO2, the parameters of the Langmuir model were derived by fitting the experimental data at pressures up to 1bar (based on low pressure

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or low concentration i.e. Henry’s Law region) and parameters of the Sips model were

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determined by fitting the experimental data at pressures up to 10 bar (Bao et al.,2011). The equilibrium selectivity values of CO2 over CH4 for the adsorbents are shown in Table 7. It is observed that the equilibrium selectivity decreased with increasing adsorption temperature for both adsorbents. One reason may be the relative magnitudes of the adsorption heats of the adsorbate gases. By combining Equations 4 and 5, it can be found that raising temperature

(

)

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causes decreasing the value of − ∆HCO2 −∆HCH4 RT, which reduces equilibrium selectivity (

S ei , j ) (Li and Tezel, 2007). Since the greatest Henry’s constant belonged to carbon dioxide

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onto NaA-S zeolite, it had greatest equilibrium selectivity for CO2/CH4 separation. This is because of the more heterogeneous surface of NaA-S zeolite and quadrupole moment of CO2.

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As a consequence, it can be concluded that NaA-S zeolite is able effectively separate CO2/CH4 gas mixture.

The ideal selectivity of CO2 over CH4 for both adsorbents at 290K is shown in Fig.9. The ideal selectivity of CO2 to CH4 at specific temperature and pressure was obtained from the ratio of CO2 and CH4 adsorbed. As observed, the highest ideal selectivity of CO2/CH4 on both adsorbents was obtained at ambient pressure and value for NaA-S zeolite was greater than NaA-Zeo zeolite. The ideal selectivity of CO2/CH4 for NaA-S and NaA-Zeo samples at

19

ACCEPTED MANUSCRIPT 290 K and ambient pressure are 7.1 and 6.4, respectively. In other pressures, it can be seen that the ideal selectivity of CO2/CH4 in both adsorbents almost coincide upon each other. This proves that CO2 removal from CO2/CH4 mixture in both adsorbents was done via a similar mechanism (Bao etal.,2011). For both adsorbents, with raising pressure, the selectivity of

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CO2 over CH4 decreased and then tend to be a constant value, as the adsorbents had higher tendency to adsorbed CO2 at low pressure, and they finally saturated at higher pressure (Doroudian Rad et al., 2012).

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The heat of adsorption at zero coverage, ∆Ho, was calculated using the Vant Hoff equation. The adsorption heat values for CO2 and CH4 were determined from the

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corresponding adsorption isotherms at 277, 290, and 310 K and plotting the Ln(KH) against (1/T) (Pakseresht et al., 2012), as shown in Fig. 10. The heat of adsorption at zero coverage and pre-exponential factor of CO2 and CH4 on both adsorbents are reported in Table 8. The adsorption heat reflects the strength of the interactions between adsorbate and the adsorbent

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and it is a quantity to measure the energetic heterogeneity of the adsorbent (Bao et al., 2011). As can be seen from Table 8, the highest heat of asdsorption at zero coverage belongs to adsorption of CO2 on NaA-S zeolite (48.5 kJmol-1). More heterogeneity of the adsorbent

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surface due to the presence of more sodium cations inside the cavities of the synthesized

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zeolite compared to commercial zeolite, and also the quadropole moment of CO2 may be the origin of this behavior (Silva et al., 2012; Li and Tezel, 2007). It can be concluded that CO2 molecules had higher affinity towards surface of the synthesized adsorbent and showed more significant interactions with this sample. The heat of asdsorption of CH4 on NaA-S zeolite was 24.4 kJmol-1. Due to the non-polar nature of CH4, the prepared sample had a lower affinity to adsorbed it.

20

ACCEPTED MANUSCRIPT 3.4. Dynamic adsorption The breakthrough curves of gases were obtained by using gaseous mixture (CH4:CO2:He=60:20:20 V:V:V) on NaA-S and NaA-Zeo zeolites at 290 K and ambient pressure. Fig. 11 shows the breakthrough curves of CO2 and CH4 on NaA-S and NaA-Zeo

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zeolites. At the beginning of the breakthrough experiments on the both adsorbents, outlet gas stream from the column was containing CH4 and He, whereas CO2 was preferentially adsorbed on the zeolites. This showed the stronger adsorption of CO2 in comparison to CH4

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(Doroudian Rad et al., 2012). This process was continued until the bed was saturated and it needed to be regenerated. Breakthrough times of CO2 and CH4 on NaA-S zeolite were near

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649 and 43 s, respectively. For NaA-Zeo sample, the breakthrough times of CO2 and CH4 were approximately 545 and 37 s, respectively. Longer breakhrough time of CO2 on NaA-S zeolite compared to NaA-Zeo zeolite confirms the higher capacity of the synthesized adsorbent to adsorb CO2. In addition, the larger difference between breakthrough time of CO2

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and CH4 on NaA-S zeolite exhibited better performance of this adsorbent for separation of CO2 from CH4/CO2 gas mixture. Since breakthrough fronts, for both adsorbents were very steep, it can be concluded that the flow pattern is close to ideal (virtually no axial dispersion)

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et al., 2013b).

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and the resistance to mass transfer in the both adsorbents is very low (Mosca et al., 2010; Yu

3. Conclusions

In this work, NaA zeolite was prepared via hydrothermal treatment. The effects of crystallization time and temperature on the synthesis of NaA zeolite were examined. It was found that at crystallization temperature of 363 K, NaA particles with well-shaped can be fully crystallized during 20 h. Then, the capability of the synthesized zeolite shaped with montmorillonite binder was evaluated for removal of carbon dioxide from carbon dioxide–

21

ACCEPTED MANUSCRIPT methane gaseous mixture in adsorption process and compared with commercial NaA zeolite. The adsorption tests of CO2 and CH4 on the adsorbents were performed at 277, 290 and 310 K, at pressures up to 10 bar. Then, the experimental data were fitted using Langmuir and Sips equations. It was found that the synthesized adsorbent had a BET specific surface area of

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222.8m2g-1, pore volume of 0.096 cm3g-1, and CO2 and CH4 adsorption capacities of 5.2 and 2.6 mmolg-1, respectively, at 310 K and 10 bar. Also, the Sips model better fitted the experimental data. The ideal selectivity of synthesized and commercial NaA zeolites were 7.1

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and 6.4 at atmospheric pressure and 290 K, respectively. The heats of adsorption at zero coverage on the synthesized adsorbent were 48.5 and 24.4 kJmol-1 for CO2 and CH4,

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respectively. Breakthrough time of CO2 on the synthesized and commercial NaA zeolites were near 649 and 545 s, respectively. According to these results, it can be stated that the synthesized NaA zeolite was a more promising candidate as adsorbent for separating carbon

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dioxide from methane in natural gas purification process.

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29

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Figure captions: Fig. 1: Schematic diagram of the laboratory set up used for obtaining adsorption equilibrium data. Fig. 2: XRD patterns of the samples obtained at different crystallization times (NaA ( ), and

RI PT

NaP( )).

Fig. 3: SEM images of the samples obtained at different crystallization times, (a) 8 h, (b)15 h, (c) 20 h, and (d) 24 h.

SC

Fig. 4: FTIR spectra of the samples obtained at different crystallization times.

Fig. 5: XRD patterns of the samples obtained at different temperatures (NaA( ), and

M AN U

sodalite( )).

Fig. 6: Nitrogen adsorption and desorption isotherms (a) and BJH pore size distribution (b)of the samples obtained at different crystallization temperatures. Fig. 7: Adsorption isotherms of CO2 and CH4 on NaA-Zeo at different temperatures and

TE D

pressures up to 10 bar.

Fig. 8: Adsorption isotherms of CO2 and CH4 on NaA-S at different temperatures and pressures up to 10 bar.

EP

Fig. 9: Ideal selectivity profiles of CO2 over CH4 by NaA-Zeo and NaA-S zeolites at 290 K

AC C

and different pressures.

Fig. 10: Plot of Ln(KH) versus(1/T) for adsorption of CO2 and CH4 on NaA-Zeo and NaA-S zeolites. Dashed lines are the models fit by Eq. (4).

Fig.11: CO2/CH4 mixture breakthrough curves for NaA-Zeo and NaA-S zeolites at 290 K and ambient pressure.

30

AC C

EP

TE D

M AN U

SC

RI PT

ACCEPTED MANUSCRIPT

Figure 1

31

ACCEPTED MANUSCRIPT

Intensity (a.u.)

RI PT

24 h

SC 15

M AN U

10

20



TE D

5

20 h

AC C

EP

Figure 2

32

25

15 h

8h 30

35

40

AC C

EP

TE D

M AN U

SC

RI PT

ACCEPTED MANUSCRIPT

Figure 3 33

100 nm

ACCEPTED MANUSCRIPT

Intensity (a.u.)

TE D

24h

20h

1900

H-O-H

701

900

AC C

1400

EP

1650

15h

8h

2400

Wavenumber (cm-1)

Figure 4

34

M AN U

3450

2900

SC

3400

RI PT

3900

O-H

Si-O, Al-O asy str.

1002

748

400

557 467

668

Si-O-Al sy str. Si-O, Al-O sy str.

D4R Si-O and Al-O bending

ACCEPTED MANUSCRIPT

RI PT

Intensity (a.u.)

393 K

10

15

20



Figure 5

AC C

EP

TE D

5

M AN U

SC

373 K

35

25

363 K 30

35

40

ACCEPTED MANUSCRIPT

RI PT

(a)

SC

363 K

393 K

(b)

8.00E-03

M AN U

373 K

4.00E-03

TE D

3.6 nm

11.2 nm

EP

6.00E-03

11.2 nm

AC C

Pore volume (cm3/g. Å)

3.6 nm

363 K 373 K 393K

2.00E-03

0.00E+00

0

20

40

60

80

100

120

Pore diameter (Å)

Figure 6 36

140

160

180

200

ACCEPTED MANUSCRIPT

P (bar) 0

2

4

6

8

10

7

RI PT

5 4 3

SC

q (mmol/g) [CO2]

6

Exp(T=277K) Exp(T=290K) Exp(T=310K) Sips model Langmuir model

2

M AN U

1 0 3.5

TE D

2.5 2 1.5

AC C

1

EP

q (mmol/g) [CH4]

3

EXP(T=277 K) Exp(T=290 K)

Exp(T=310 K)

0.5

Sips model Langmuir model

0

0

2

4

6

P (bar)

Figure 7

37

8

10

ACCEPTED MANUSCRIPT

0

2

4

P (bar)

6

8

10

7

RI PT

5 4 3

Exp(T=277K) Exp(T=290K) Exp(T=310K) Sips model Langmuir model

SC

q (mmol/g) [CO2]

6

2

M AN U

1 0 4 3.5

TE D

2.5 2 1.5

AC C

1

EP

q (mmol/g) [CH4]

3

Exp (T=277 K) Exp (T=290 K) Exp (T=310K) Sips model Langmuir model

0.5 0

0

2

4

P (bar)

6

Figure 8 38

8

10

ACCEPTED MANUSCRIPT

8

NaA-S

7

NaA-Zeo

RI PT

qCO2/qCH4

6 5 4 3

SC

2

0 1

2

3

M AN U

1

4

5

6

P (bar)

AC C

EP

TE D

Figure 9

39

7

8

9

10

ACCEPTED MANUSCRIPT (1/T) 103 3.2

3.3

3.4

3.5

3.6

3.7

0

-1

RI PT

Ln KH [CO2]

-0.5

-1.5

-2

SC

NaA-S

-2.5 -3

-4

-5

AC C

-5.5

3.4

3.5

3.6

TE D

Ln KH [CH4]

-4.5

3.3

EP

3.2

M AN U

NaA-Zeo

NaA-Zeo NaA-S

-6

(1/T) 103

Figure 10

40

3.7

ACCEPTED MANUSCRIPT Time (s) 0

200

400

600

800

1.4

NaA -S 1.2

RI PT

C/Co

1 0.8 0.6

SC

0.4

CO2 0.2

M AN U

CH4

0 1.4 1.2

TE D

C/Co

1

NaA-Zeo

0.8 0.6

EP

0.4

CO2

0.2

AC C

CH4

0

0

200

400

Time (s)

Figure 11

41

600

800

ACCEPTED MANUSCRIPT

Table 1: Studies on CO2 and CH4 adsorption properties of some zeolites. CH4Adsorption capacity (mmol/g)

13X

298 K, 1000 kPa

6.52

3.06

T-type

288 K, 100 kPa

4.01

0.74

NaA

298.15 K, 300 Psi

4.80

_

5A

303 K, 1000 kPa

3.55

13X

313 K, 400 kPa

5.20

Beta

328 K, 1000 kPa

2.78

NaA

273 K, 100 kPa

4.46

NaA

298 K, 1000 kPa

4.47

NaKA

273 K, 101 kPa

3.47

NaA

273 K, 101 kPa

NaA

293 K, 60 kPa

References

2.13

Cavenati et al., 2004

5.42

Jiang et al., 2013

Siriwardaneet al., 2001

1.78

2.00

Pakseresht et al.,2002

1.2

4.33

Silva et al., 2012

1.14

2.43

Huang et al.,2010

_

_

Zukal et al., 2011

_

_

Akten et al., 2003

_

_

Cheung et al., 2013

3.77

_

_

Cheung et al., 2013

3.09

_

_

EP

TE D

M AN U

_

AC C

Adsorbent

Ideal Separation capacity (CO2/CH4)

RI PT

CO2Adsorption capacity (mmol/g)

SC

Adsorption Condition

Romero-Pérez and Aguilar Armenta, 2010

42

ACCEPTED MANUSCRIPT Table 2: Chemical analyses of the loaded adsorbents in the adsorption column. NaA-S

NaA-Zeo

SiO2

35.59

44.88

Al2O3

27.95

38.06

Na2O

23.21

Fe2O3

1.42

MgO

0.79

CaO

1.52

K2O P2O5

AC C

EP

TE D

SiO2/Al2O3(mol/mol)

0.18

0.47

SC

0.27

7.13

M AN U

TiO2

RI PT

Component (Wt.%)

43

0.53 1.26

1.44

1.23

0.12

0.26

2.16

2

ACCEPTED MANUSCRIPT Table 3: Effects of crystallization conditions on phase and relative crystallinity of NaA zeolite. Crystallization

Crystallization

Relative

Temperature (K)

Time (h)

Crystallinity (%)

363

8

38

363

15

77

363

20

100

363

24

56

373

20

73

A

393

20

57

A+SOD c

SC

RI PT

Aa

M AN U

aZeolite

NaA. NaP. cHydroxy-sodalaite.

AC C

EP

TE D

bZeolite

Product phase

44

A

A

A+P b

ACCEPTED MANUSCRIPT Table 4: The textural properties of NaA synthesized for 20 h.

SBET

SEternx

VMicropors

VTotal

Temperature (K)

(m2.g-1)

(m2.g-1)

(cm3.g-1)

(cm3.g-1)

363

222.8

40.58

0.087

0.17

373

97.2

36.50

0.029

0.16

393

66.8

24.65

0.020

0.08

AC C

EP

TE D

M AN U

SC

RI PT

Synthesis

45

ACCEPTED MANUSCRIPT

RI PT

Table 5: Langmuir and Sips parameters for CO2 and CH4 at 277, 290 and 310 K on NaA-Zeo zeolite.

qm

K*106

(mmol.g-1)

(Pa-1)

R2

(mmol.g-1)

(Pa-1)

N

R2

277

6.1153

84.280

0.9972

140.847

1.6287

0.9997

290

5.5586

45.534

0.9954

5.3948

64.212

1.4573

0.9990

310

5.1697

13.413

0.9938

4.9207

15.290

1.2431

0.9966

277

4.4216

2.512

0.9976

3.9539

2.730

1.1402

0.9984

290

4.2923

1.844

0.9981

3.7952

2.011

1.1195

0.9987

310

4.2631

0.9976

3.6677

1.332

1.0999

0.9980

5.9870

M AN U

(K)

TE D

CH4

b*106

EP

CO2

qm

1.192

AC C

Adsorbate

Temp.

Sips equation

SC

Langmuir isotherm

46

ACCEPTED MANUSCRIPT

Table 6: Langmuir and Sips parameters for CO2 and CH4 at 277, 290 and 310 K on NaA-S zeolite.

K*106

(K)

(mmol.g-1)

(Pa-1)

R2

(mmol.g-1)

(Pa-1)

n

R2

277

6.5084

103.490

0.9968

212.543

1.8723

0.9997

290

5.9893

54.861

0.9934

5.8013

91.759

1.6696

0.9986

310

5.7279

14.116

0.9889

5.3722

16.932

1.3819

0.9957

277

4.7178

2.692

0.9969

4.1171

2.959

1.1945

0.9985

290

4.3961

2.345

0.9976

3.8203

2.583

1.1762

0.9990

310

4. 3641

1.520

0.9966

3.7403

1.699

1.1290

0.9972

SC

qm

6.3755

M AN U

b*106

TE D

CH4

qm

EP

CO2

Temp.

AC C

Adsorbate

Sips equation

RI PT

Langmuir isotherm

47

ACCEPTED MANUSCRIPT Table 7: Henry’s constant and equilibrium selectivity for CO2 and CH4.

KH* 106 (mmol.g-1. Pa-1)

NaA-S

CH4

KH(CO2)/KH(CH4)

277

478.24

9.65

49.56

290

229.04

6.98

310

59.1

3.39

277

622.28

290

11.54

32.81

17.43

53.94

290.70

7.92

36.72

67.34

3.77

17.86

AC C

EP

TE D

310

RI PT

CO2

SC

NaA-Zeo

Temp. (K)

M AN U

Adsorbent

48

ACCEPTED MANUSCRIPT Table 8: Heat of adsorption and pre-exponential factors for CO2 and CH4 on NaA-Zeo and NaA-S zeolites.

- Ho (Kj/mol)

KHo (mmol.g-1.kPa-1) CH4

CO2

NaA-Zeo

1.301e-9

4.85e-7

45.55

22.91

NaA-S

4.787e-10

2.99e-7

48.50

24.40

M AN U TE D EP AC C 49

CH4

RI PT

CO2

SC

Adsorbent

ACCEPTED MANUSCRIPT The synthesis of NaA zeolite via the hydrothermal method was examined. The optimum conditions for synthesis of NaA zeolite were determined. The performance of the shaped synthesized NaA zeolite with montmorillonite clay for CO2/CH4 separation were investigated. The synthesized NaA zeolite showed a more promising performance compared

AC C

EP

TE D

M AN U

SC

RI PT

to commercial NaA zeolite for CO2 separation.