Tuning the Adsorption Properties of Zeolites as Adsorbents for CO2

Article pubs.acs.org/IECR

Tuning the Adsorption Properties of Zeolites as Adsorbents for CO2 Separation: Best Compromise between the Working Capacity and Selectivity Edder J. García,† Javier Pérez-Pellitero,† Gerhard D. Pirngruber,*,† Christian Jallut,‡ Miguel Palomino,§ Fernando Rey,§ and Susana Valencia§ †

IFP Energies nouvelles, Rond Point de l’échangeur de Solaize, 69360 Solaize, France Laboratoire d’Automatique et de Génie des Procédés, UMR UCBL/CNRS 5007, ESCPE, Université de Lyon, Université Claude Bernard Lyon 1 (UCBL), 43 Bd du 11 Novembre 1918, 69622 Villeurbanne Cedex, France § Instituto de Tecnología Química, Consejo Superior de Investigaciones Científicas, Universidad Politécnica de Valencia. Avenida de los Naranjos s/n, 46022 Valencia, Spain ‡

S Supporting Information *

ABSTRACT: The choice of an appropriate adsorbent for CO2 separation by pressure-swing adsorption remains a field of intense research. In this work, several FAU and LTA zeolites with different Na contents (Si/Al ratios) are studied for the separation of CO2 from mixtures of CO2, CO, and CH4 by means of breakthrough experiments. The breakthrough experiments were carried out between 1 and 5 bar at 303 K using two feed mixtures: 50/50 (v/v) CO2/CH4 and 75/15/15 (v/v/v) CO2/ CH4/CO. The most polar zeolites, i.e., those with high Na content, exhibit the highest adsorption capacity and selectivity for CO2, but their regeneration is difficult; hence, their working capacity is low. The opposite is true for the least polar zeolites, i.e., those with low Na content. In order to quantify the trade-off between the selectivity and working capacity, the Ruthven statistical model (RSM) was used. It satisfactorily reproduced the experimental trends. We, therefore, used the RSM to identify the properties of the adsorbent that lead to an optimal compromise between the working capacity and separation factor. The critical parameter is the concentration of extraframework cations, which, in turn, depends on the framework charge of the zeolites FAU and LTA. The optimal trade-off zone is defined in terms of the Henry constant of CO2 (KCO2). It is found that this zone is placed between KCO2 = 5 × 10−3 and 50 × 10−3 molecules·bar−1·Å−3. This interval corresponds to a heat of adsorption of CO2 at zero coverage between 27 and 32 kJ·mol−1. In our study, this optimal range of Henry constants was achieved for the zeolites Na-USY, SAPO-37, LTA (Si/Al = 5), and EMC-1.

1. INTRODUCTION The possibility of exploiting nonconventional natural gas resources is expected to completely change the energy landscape in the near future. In this scenario, methane will substitute for oil derivatives in many uses. It is expected that the production of natural gas in the United States will reach 33.1 trillion of cubic feet (tcf) in 2040, growing from 24 tcf in 2011. In practice, this increase in methane production is coming from the exploitation of shale gas reserves (from 8.1 tcf in 2012 to 16.7 tcf in 2040).1 However, this must be necessarily accompanied by an improvement of the technologies of separation and purification of methane from the rest of natural gas.2 The main objective in natural gas purification is to reduce the concentration of CO2, H2S, water, and other impurities to achieve transport specifications. The limit of the CO2 concentration in most commercial pipelines depends on domestic regulations but usually is less than 2%. Several technologies are available to remove CO2 from natural gas. The most important are cryogenic distillation, solvent extraction by alkaloamines, and adsorbent-based processes. The choice among these technologies depends on the operational cost and specifications of the feed gas and final product. Cryogenic distillation and solvent extraction are well© 2014 American Chemical Society

known for having a higher operational cost than adsorbentbased technologies. Nevertheless, these technologies are appropriate in the following cases: (1) when other impurities like water or H2S must be removed at the same time; (2) when the final product is liquefied natural gas; (3) when large amounts of gas need to be treated. For intermediate amounts of gas, adsorbent-based technologies become economically attractive. Pressure-swing adsorption (PSA) is an adsorbent-based process in which an impurity, in this case CO2, is selectively adsorbed in a solid adsorbent. The adsorbent is then regenerated by decreasing the pressure. CO2-PSAs have been studied by experimental and simulation methods.3−6 It has been shown that the properties of the adsorbent such as its polarity, pore diameter, and chemical composition have a dominant effect on the process.7 Many new adsorbent materials have been proposed in the literature: metal−organic frameworks (MOFs),8,9 zeolitic imidazolate frameworks, new Received: Revised: Accepted: Published: 9860

January 17, 2014 May 16, 2014 May 16, 2014 May 16, 2014 dx.doi.org/10.1021/ie500207s | Ind. Eng. Chem. Res. 2014, 53, 9860−9874

Industrial & Engineering Chemistry Research

Article

Figure 1. Schematic representation of a PSA process. Concentration profiles of heavy (solid line) and light components (dotted line).

interacting or “light” components of the gas mixture are hardly retained and flow through the adsorbent material. Hence, the column effluent, which is called the raffinate, is strongly enriched with these “light” components and depleted of the “heavy” components. Once the adsorbent material becomes saturated mainly with the heavy components, the former (heavy) components begin to break through at the column exit and the adsorption step is stopped. In order to recover the heavy components and regenerate the adsorbent, the column pressure is reduced and the flow direction is reversed. Desorption may be assisted by the use of a purge gas. The heavy components are then recovered in an extract flow. In addition, a complete elementary PSA cycle includes pressurization or depressurization steps of the column. In order to achieve a quasi-continuous process, many columns can be operated in parallel and flows may be recycled from one column to another. Because of its dynamic and cyclic character, the optimization of a PSA system is a complex problem that can be treated by numerical simulation (for the purpose of finetuning the PSA cycle), but first of all, a well-suited adsorbent has to be chosen. A rational approach toward the selection of the best-suited adsorbents by efficient screening methods, prior to PSA experiments or simulations, is therefore very important. Our interest lies in constructing a strategy for screening adsorbents for CO2 capture mostly from natural gas, but also from synthesis gas, biogas, flue gas, etc. In order to address this problem, we first discuss the way to roughly characterize the performance of a PSA process.21 2.2. Principle of the Screening: Definition of Adsorbent Selection Criteria. The performance of a PSA process can be characterized by its productivity as well as by the purity and recovery of the desired component denoted i. The productivity (Prod) is the amount of feed that can be treated by a given amount of adsorbent (or by a given volume of the column) during one cycle:

zeolites,10,11 micro/mesoporous carbon materials with or without functionalization,12,13 etc. Most full-fledged PSA experiments or simulations are still carried out either with activated carbons or with zeolite 13X (NaX).14−16 Activated carbons are reputed to be “apolar”, with a fairly weak affinity for CO2. Hence, they are easily regenerable but poorly selective for adsorption of CO2 versus other apolar components (CH4, N2, etc.). On the contrary, zeolite 13X is a highly “polar” adsorbent. It has a strong affinity for adsorption of CO2 and a high selectivity for CO2 versus apolar compounds, but as a consequence, its regeneration becomes more difficult. The ideal CO2 adsorbent would be a material with an intermediate CO2 affinity that combines a fairly high adsorption capacity with good selectivity and easy regeneration.17,18 The CO2 affinity can be tuned via the polarity of the adsorbent. Palomino et al.19 have shown that the polarity of a zeolite can be tuned via its Si/Al ratio and have demonstrated the impact on the adsorption of CO2/CH4 mixtures. Increasing the polarity, i.e., reducing the Si/Al ratio, improved the selectivity toward CO2, but regeneration became difficult because of the strong CO2− zeolite interaction. The present contribution follows up on this work, with the objective of providing a quantitative definition of the optimal polarity of an adsorbent material for CO2 capture by PSA. In order to explain our research strategy, we start out with an introductory section on the principles of PSA separations and on how the PSA performance can, in a first approximation, be captured by a few parameters that can be rather easily determined by experimental screening methods.

2. SCREENING AND OPTIMAL DESIGN OF ADSORBENTS FOR PSA PROCESSES 2.1. Basic PSA Cycle. The principle of the basic cyclic PSA process is well-known (see Figure 1).20 The granular adsorbent material is placed in a column that is fed with a gas mixture. The “heavy” or most strongly adsorbed components are preferentially retained in the adsorbent, while the weakly 9861

dx.doi.org/10.1021/ie500207s | Ind. Eng. Chem. Res. 2014, 53, 9860−9874


Prod =

Ffeed Vcoltcycle

Article

SFij =

(1)

Ffeed is the flow of the feed, Vcol is the volume of the column, and tcycle is the time of each cycle. The productivity depends on the quantity of component i (in our case CO2) that the adsorbent material can retain during each adsorption cycle. This parameter is called the cyclic working capacity Δni22 and is defined as the difference between the adsorbed amounts during the adsorption and desorption steps. The higher this cyclic working capacity Δni is, the higher the feed flow that can be treated. We must note, however, that increasing Ffeed/Vcol decreases the contact time. At low contact times, the masstransfer rate may become a limiting factor. If the characteristic time of mass transfer between the gas phase and adsorbent is higher than the contact time, the full equilibrium cyclic capacity of the adsorbent cannot be exploited. The purity of component i in the extract is a function of the ratio of the cyclic working capacities of component i to the other component(s) (j) of the gas mixture. This ratio is called the separation factor (SF):23

SFij =

Δni Δnj

ni Xj nj Xi

(4)

Equation 4 shows that the separation factor is proportional to the product of the selectivity and Δni/ni. Δni/ni is the regenerability of the adsorbent with respect to the desired component i. 2.3. Limits of the Method. In the equilibrium approach, we assume that the entire column is in equilibrium with the feed at the end of the adsorption step. Because the concentration fronts are not infinitely sharp, a small fraction of the heavier component breaks with the light component. Hence, the recovery in a real PSA column is lower than the amount predicted by the equilibrium approach. As mentioned above, the shape of the breakthrough curve depends on the slope of the multicomponent adsorption isotherm. For steeper isotherms, the breakthrough front is sharp; hence, recovery will be higher than that for adsorbents with shallow isotherms and disperse breakthrough fronts. We also assumed that the adsorption process is isothermal. This condition will not be fulfilled in a real PSA column, which is adiabatic rather than isothermal.6 Adsorbents with high heat of adsorption lead to a high temperature rise in the column during the adsorption step, which lowers the adsorbed amount. Yet, high heats of adsorption are usually associated with steep isotherms. Thus, the two above-mentioned errors will partially compensate for each other,24 but unfortunately we cannot easily predict to which extent that will happen. Therefore, our method has to be regarded as a guide for preselecting the potentially most interesting adsorbent materials. For the final choice of the best solid, a series of PSA experiments and simulations are necessary in completing the investigation. Economic aspects (cost of adsorbent) must also be taken into account. 2.4. Best Adsorbent. The ideal adsorbent for a PSA process exhibits a very high cyclic working capacity and a very high separation factor, i.e., significant adsorption of the desired component i along with negligible adsorption of the other components j. Also, it must be very easy to regenerate. Unfortunately, high selectivity and easy regeneration can hardly be combined in one material. High selectivity for adsorption of component i is necessarily associated with its strong adsorption. Strong adsorption typically illustrates a steep slope, rendering adsorbent regeneration difficult, which leads to a low cyclic capacity. Hence, we necessarily have to make a compromise between the selectivity and cyclic capacity. The challenge in adsorbent selection is to define the optimal compromise. Let us illustrate the situation by considering the separation of CO2 from a CO2/CH4 mixture. CH4 is basically an apolar molecule that interacts with adsorbents mainly via dispersion− repulsion forces. The strength of the adsorption of CH4 mainly depends on the pore size of the adsorbent because dispersion forces are amplified by a strong confinement.28−30 The adsorption of CO2 via dispersion−repulsion forces is stronger than that of CH4, but the selectivity that can be achieved via this adsorption mechanism is low. However, CO2 has a permanent electric quadrupole moment that can interact with the electric field generated by the charge distribution of the adsorbent material. The additional electric forces enhance the selectivity for CO2 adsorption. We can, thus, tune the strength of the CO2 interaction and the selectivity for CO2 by changing the electric field of the adsorbent. Palomino et al.19 show that this can be done by changing the Si/Al ratio and, thus, the concentration of extraframework

(2)

Figure 1 illustrates that Δni is maximized when the whole column is saturated with the heavy component. Because the concentration front is not infinitely sharp, a partial breakthrough of the heavy component during the adsorption step must be accepted, at the expense of decreasing the recovery. The extent of breakthrough depends on the sharpness of the concentration front, which, in turn, depends on the slope of the adsorption isotherm (a steep isotherm leads to a sharp concentration front) and on the mass-transfer rate.24 As far as small molecules are concerned and because we do not consider kinetic separation processes here (i.e., carbon molecular sieves or narrow-pore zeolites10,25−27), we assume that the separation process is not limited by mass transfer. If the adsorption process is considered to occur isothermally, the equilibrium theory of adsorption columns can be used to estimate the performance indexes only from the multicomponent adsorption isotherms of the material. In other words, we assume that the entire adsorption column is in equilibrium with the gas phase at the end of the adsorption and desorption steps. For the adsorption step, this means that the column is in equilibrium with the feed at adsorption pressure. Because the objective in CO2 capture is to recover pure CO2, the desorption is usually carried out by depressurization, without using a purge gas that would dilute CO2. Hence, we consider that the column is in equilibrium with pure CO2 at the pressure of desorption. In the equilibrium model, the separation factor depends on the selectivity of the adsorbent with respect to component i, which is defined by αij =

Δni Xi αij ni Xj

(3)

where ni/j are the adsorbed amounts of components i and j at equilibrium with a gas phase defined by the molar fractions Xi/j. If the light component j is fully desorbed during the regeneration step, i.e., nj = Δnj, the separation factor is given by 9862



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Table 1. Characterization of the Adsorbents

a

material

Si/Al

% Na

SBET, m2·g−1

Vmicro,a mL·g−1

grain density, kg·m−3

NaX EMC-1 SAPO-37-Na Na-USY DAY AlPO4-42 LTA-5 LTA-3.5 LTA-2

1.29 3.89 0.91 8.6 >100 n.d. 5.5b 3.5b 1.9b

9.68 6.02 2.02 1.3 n.d. n.d. 1.0 4.3 7.6

816 780 696 n.d.c 642 510 654b 642b 583b

0.34 0.33 0.26 0.27 0.25 0.23 0.27b 0.30b 0.28b

1515 1390 1200 1348 1326 1200 1498 1527 1590

Vmicro is the microporous volume calculated using the t-plot method. bReported in ref 19. cn.d. = not determined. Vadsorbent and Vcrucible were calculated using the helium isotherm method.36 The contribution of the adsorbed phase volume (Vadsorbed phase) was approximated by the microporous volume of the samples. In order to obtain the saturation capacity, the single-component adsorption isotherms were fitted using the Langmuir model.

cations in a series of zeolite A materials. In the present paper, we extend the approach to faujasite zeolites. Several faujasitetype zeolites with different concentrations of extraframework Na+ cations were used for the adsorption of CO2, CO, and CH4 (three major components of synthesis gas) and for the separation of CO2 from binary CO2/CH4 and ternary CO2/ CH4/CO mixtures. Four solids with LTA topology were included in the comparison. The experimentally measured cyclic capacities and separation factors were compared with the results of a theoretical equilibrium model. The theoretical model is used to quantify the properties of the optimal CO2 adsorbent for a given set of operating conditions.

n=

(6)

where n is the absolute adsorbed amount, nsat is the saturation capacity, and KL is the Langmuir constant. The Henry constant (K) was calculated from the single adsorption isotherms using the virial plot method. This approach is applied to determine K for adsorbents that strongly interact with adsorbate molecules. For this system, it is difficult to obtain experimental adsorption data at low coverage. Therefore, K cannot be directly determined from the initial slope of the isotherm, as exemplified by the Langmuir model. This is the case of the adsorption of CO2 on cationic zeolites. The virial isotherm, eq 7, assumes that the adsorbed-phase state can be described by the virial equation of state.20

3. EXPERIMENTAL DETERMINATION OF PERFORMANCE INDEXES 3.1. Adsorbent Materials. In order to cover a large range of faujasite-type zeolites with different framework charges and therefore different concentrations of extraframework cations, we used the following five samples: NaX (Si/Al = 1.2), produced by IFPEN, EMC1 (Si/Al = 3.8), synthesized at IFPEN, SAPO-37, synthesized at IFPEN, USY CBV712, purchased from Zeolyst, and DAY (Si/Al > 100), purchased from Degussa. The properties of these zeolites are compiled in Table 1. EMC-1 was prepared following the protocol of Delprato et al.31 SAPO-37 was prepared by using the method described by Briend et al.32 The commercial USY CBV712 contains NH4+ as the extraframework cation. Therefore, a 3-fold ion exchange with a 0.5 M solution of NaNO3 was carried out to introduce Na + ions. Three zeolite A samples with different Si/Al ratios were prepared at ́ the Instituto de Tecnologiá Quimica de Valencia, according to the protocols described in the literature.19,33,34 The fourth sample was an aluminophosphate with LTA topology, i.e., AlPO4-42. The solid was prepared at IFPEN, following the modus operandi of Fayad et al.35 Elemental analysis of the solids was carried out by X-ray fluorescence (for the Si/Al ratio), atomic absorption spectroscopy (AAS; for the Na content), or inductively coupled plasma (ICP; in the case of the LTA samples). For ICP analysis, the samples were dissolved in 1/1/1 HF/HNO3/HCl (total acid: 5 mL), in a ratio of 100 mg of sample per 200 mL of acid solution. The Si, Al, and Na contents were analyzed simultaneously from the same solution. 3.2. Adsorption Isotherms. Single-component equilibrium isotherms of CO2, CH4, and CO were measured at 303 K by using a Rubotherm magnetic suspension balance. The mass of the sample was approximately 1−2 g. Before pressure measurements were carried out, all of the samples were outgassed under a high vacuum at elevated temperature. The conditions of outgassing are given in the Supporting Information. After stabilization of the mass, temperature, and pressure, the change in the mass was recorded. This allowed a direct measurement of the reduced mass (τ). To calculate the absolute adsorbed quantity (n), correction of the buoyancy effect was taken into account. The absolute adsorbed quantity was calculated by n = τ + ρbulk (Vadsorbent + Vcrucible + Vadsorbed phase)

nsatKLp 1 + KLp

⎛ Kp ⎞ 3 ln⎜ ⎟ = 2vc1n + vc 2n2 + ... ⎝ n ⎠ 2

(7)

where vc1, vc2, ... are the virial coefficients. At low pressure, the contribution of the virial coefficients of order higher than 1 can be neglected; thus, the model becomes

⎛ p⎞ ln⎜ ⎟ = 2vc1n − ln K ⎝n⎠

(8)

Therefore, a plot of ln(p/n) versus n is linear as the adsorbed amount goes to zero. The intersection of the plot gives ln K. This method is useful in the case of strong adsorption because a linear plot is found way above the Henry regime. 3.3. Breakthrough Measurements. Breakthrough curves of gaseous binary [50/50 (v/v) CO2/CH4] and ternary [70/15/15 (v/v/ v) CO2/CH4/CO] mixture feeds were carried out on FAU and LTA zeolites using a homemade apparatus. A schematic diagram of this apparatus is shown in the Supporting Information. The gaseous flows were controlled by means of mass-flow controllers. The total flow was kept at 4 NL·h−1. The pure compound flows were adjusted to generate a mixture of the desired composition. The gas concentrations at the system outlet were measured using an online mass spectrometer. In order to maintain the outlet gas flow rate at a constant value, the latter was diluted using helium with a 1/25 (v/v) ratio. This is important for accurate quantification of the components. The adsorbent was placed in powder form in a stainless steel column with an inner diameter of 10 mm and a length of 80 mm. The adsorbent mass was approximately 2 g. The adsorbents were tested as powder and activated under a helium flow at the appropriate temperature (see the Supporting Information) with a heating rate of 1 K·min−1. After cooling, the pressure in the column was raised to the desired value in a flow of helium. The pressure drop of the column was less than 0.05 bar; therefore, it was considered negligible. At the same time, the CO2/

(5) 9863



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multicage system instead; i.e., smax was arbitrarily fixed to 100 (see ref 30 for more details). We have shown30 that the multicage description is equivalent to the single-cage description if the number of adsorbed molecules per cage is large (which is the case for CO2 adsorbed in FAU and LTA zeolites). The adapted single-component RSM equation is

CH4 or CO2/CH4/CO mixture was pressurized in the bypass line. Then, the mixture was injected into the column using an automatic valve, and the breakthrough curve was recorded via the mass spectrometer. The breakthrough curves of the more weakly adsorbed components CH4 and CO are characterized by a so-called roll-up. Because the roll-up increases the experimental uncertainty on the calculated uptake, a supplementary experiment was carried out. In this case, the column was saturated initially using CO2 followed by injection of the mixture into the column. This second experiment allows measurement of the breakthrough curves of CH4 and CO without roll-up. Last, desorption experiments were conducted upon passing a flow of helium (4 NL·h−1) through the column at pressure. The absolute adsorbed quantity at equilibrium was calculated from the first moment of the breakthrough curve (μ), according to

⎛ ⎞ C m ni = i ⎜⎜μi Vf − Vcol + ads ⎟⎟ ρgrain ⎠ mads ⎝

⎧⎡ ⎢Kp(bsmax ) θ = (1/smax )⎨ ⎪ ⎩⎢⎣ ⎪

smax

+

s=2

μi =

∫

(9)

smax

+

∑ s=2

s ⎛ s 2 ε ⎞⎤⎫ (Kp)s (bsmax ) ⎛ s ⎞ ff ⎟⎥⎬ ⎜1 − ⎟ exp⎜ ⎪ s! smax ⎠ ⎝ ⎝ smax kT ⎠⎥⎦⎭ ⎪

(12)

where θ is the coverage, i.e., the fractional loading of the pore volume. The RSM can be readily extended to multicomponent adsorption.38 In the case of a binary mixture, the coverage of component 1 is given by

(10)

where Fi,0 and Fi are the molar flows of component i at the inlet and outlet of the column, respectively. A numerical integration based on the trapezoidal rule was performed in order to calculate μi. The interval of integration was taken from the time of injection (t = 0), after correction for the dead time, until the time where the Fi/Fi,0 ratio is stable and close to 1. The selectivities were determined from the experimental breakthrough curves by using eq 2. 3.4. Experimental Determination of the Working Capacity of CO2. The experimental working capacity of CO2 was obtained by calculating the difference between the adsorbed amounts of CO2 under adsorption and desorption conditions. The adsorption and desorption pressures were set at 5 and 1 bar, respectively. The volumetric composition of the feed was CO2/CH4 = 50/50 for the binary mixture and CO2/CH4/CO = 70/15/15 for the ternary mixture. The former mixture is roughly representative of a biogas and the latter of a synthesis gas. The desorption pressure was set at 1 bar. In CO2 capture, the objective is to recover pure CO2. Therefore, the amount of CO2 that remains adsorbed under desorption conditions was determined from a breakthrough curve of pure CO2 at 1 bar.

⎡ ⎧ ⎪ θ1 = (1/smax )⎨⎢K1pX1(B1smax ) ⎪⎢ ⎩⎣ jmax smax

+

∑∑ j=2 s=2

(K1pX1)s (K 2pX 2) j (b1smax )s (b2jmax ) j (s − 1) ! j!

⎤ Zff, sj ⎥ ⎥⎦

⎡ ⎢1 + K pX (b s ) + K pX (b j ) 1 1 1 max 2 2 2 max ⎢⎣ jmax smax

+

∑∑ j=2 s=2

(K1pX1)s (K 2pX 2) j (b1smax )s (b2jmax ) j s ! j!

⎤⎫ ⎪ Zff, sj ⎥⎬ ⎥⎦⎪ ⎭ (13)

where j + s ≥ 2, b1 and b2 are the excluded volumes of the adsorptive molecules 1 and 2, and X1 and X2 are the molar fractions of molecules 1 and 2, respectively. The term Zff,sj is the configurational integral of the fluid−fluid interactions in a system with s molecules of type 1 and j molecules of type 2. It is given by

4. ESTIMATION OF PERFORMANCE INDEXES VIA AN ADSORPTION MODEL We chose the Ruthven statistical model (RSM)37 for this purpose (the reasons for this choice are outlined in ref 30). The RSM is derived from statistical thermodynamics and describes the adsorption isotherm of a set of freely mobile adsorbate molecules in a cage. The main parameter of the RSM is the Henry constant K, which describes the strength of the adsorbate−adsorbent interaction. Adsorbate−adsorbate interactions are accounted for via a mean-field Sutherland potential, characterized by the depth of the potential well, εff. The saturation of the pore volume, smax, is given by the ratio between the excluded volume of the adsorbate molecules, b, and the pore volume, v. smax = v /b

s ⎛ s 2 ε ⎞⎤ (Kp)s (bsmax )s ⎛ s ⎞ ff ⎟⎥ ⎜1 − ⎟ exp⎜ smax ⎠ (s − [1)! ⎝ ⎝ smax kT ⎠⎥⎦

⎡ ⎢1 + Kp(bsmax ) ⎢⎣

where ni is the absolute adsorbed amount of component i, Ci is the gasphase concentration, mads is the adsorbent mass, Vf is the total volumetric flow, Vcol is the column volume, and ρgrain is the adsorbent grain density. The first moment for component i can be calculated from

⎛ F ⎞ ⎜⎜1 − i ⎟⎟ dt Fi,0 ⎠ ⎝

∑

i+j ⎛ ⎡β jb ⎞ sb Zff, sj = ⎜1 − 1 − 2 ⎟ exp⎢ (s 2b1εff,1 + sjb12εff,12 ⎣v ⎝ v v ⎠ ⎤ + j 2 b2εff2)⎥ ⎦ (14)

where β = 1/kT. The crossed terms εff,12 and b12 are obtained from Lorentz−Berthelot combining rules. At low pressure, the selectivity calculated by the RSM tends toward the ratio of the two Henry constants K1/K2. If the intermolecular interactions are weak, then the selectivity remains constant as a function of the pressure. If εff,1 > εff,2, then the adsorption selectivity for component 1 will increase with the pressure because of its stronger intermolecular interactions.

(11)

In Ruthven’s original work, the pore volume corresponds to the volume of a single zeolite cage and smax to the number of molecules that fits into a single cage. We use here an arbitrary 9864



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Figure 2. CO2, CH4, and CO adsorption isotherms at 303 K on (A) NaX, (B) EMC-1, (C) Na-USY, and (D) SAPO-37. The line is the best fit for the Langmuir adsorption model.

adsorbed amount was transformed to moles per volume of adsorbent using the grain density. All of the isotherms for the FAU series present a type I shape according to IUPAC classification.40 The steepest CO 2 isotherm is found for the NaX sample, followed by EMC-1. While the shape of the CO2 isotherms clearly depends on the polarity of the faujasite zeolite, the CH4 isotherms (and to a lesser extent the CO isotherms) are quite similar for the cases of the four faujasite zeolites shown in Figure 2. The best-fitting parameters of the Langmuir model for the FAU series are given in the Supporting Information. In the case of the DAY zeolite, the values were calculated from the experimental data reported by Maurin et al. at 300 K.41,42 The CO2 Langmuir constant decreases according to the following order: NaX > EMC-1 > SAPO-37 > Na-USY > DAY. The saturation capacity roughly follows the order of the pore volumes. Table 2 summarizes the Henry constants at 303 K of FAU zeolites using the virial plot method. The Henry constants were normalized with respect to the available volume for adsorption,

For our model calculations, we used the Henry constant as the only free parameter. The molecular volumes of CH4 and CO2 were set at 62 and 60 Å3·molecule−1, respectively. The fluid−fluid attraction parameters were the same as those used in the previous work.30 The temperature was fixed at 303 K, as in the experimental adsorption measurements. The RSM calculates a fractional loading of the pore volume θ. In order to convert the fractional loading into an adsorption capacity per volume of adsorbent (n), θ must be multiplied by the porosity of the adsorbent φ. The porosities of the crystal structures of FAU- and LTA-type zeolites are similar (differences in the experimentally measured porosities rather arise from differences in the degree of crystallinity). As a consequence, to establish a comparison between the different FAU- and LTA-type samples, we could, therefore, directly use the fractional loading.

5. RESULTS 5.1. Characterization of the Solids. The Si/Al ratio was determined by X-ray fluorescence. The Na content was obtained by means of AAS. The textural properties were determined by N2 physisorption at 77 K. Table 1 summarizes the results of the characterization of tested materials. The grain density of the solids was calculated using their crystallographic data. 5.2. Adsorption Isotherms. FAU-Type Zeolites. Figure 2 shows the CO2, CH4, and CO adsorption isotherms at 303 K in FAU zeolites expressed in the absolute amount per volume of adsorbent. This representation is more suitable than the conventional adsorbed amount in moles per gram of adsorbent because the adsorption columns are usually designed in terms of the bed size rather than the mass of adsorbent.39 The

Table 2. Henry Constants at 303 K on FAU Zeolites

a

9865

solid

KCO2 × 103, molecules·bar−1· Å−3

KCH4 × 103, molecules·bar−1· Å−3

KCO × 103, molecules·bar−1· Å−3

NaX EMC-1 Na-USY SAPO-37 DAYa

252.96 43.07 9.85 21.94 1.07

1.53 0.70 1.53 0.56 0.56

3.69 1.06 0.56 0.80

Calculated from the experimental data at 300 K in refs 41 and 42. dx.doi.org/10.1021/ie500207s | Ind. Eng. Chem. Res. 2014, 53, 9860−9874


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i.e., the pore volume given in Table 1. The KCO2’s increase in the same order as that described above for the Langmuir constants. This sequence follows the order of the Na+ content of the samples (Table 1). LTA-Type Zeolites. The CO2 and CH4 single adsorption isotherms at 303, 283, and 273 K of the LTA-2, LTA-3.5, and LTA-5 samples used in this work have been previously reported.19 In the Supporting Information, these isotherms are compared with those of AlPO-42, i.e., an LTA-type zeolite without extraframework cations. The Henry constants at 303 K calculated using the virial method are shown in Table 3. For the LTA-2, LTA-3.5, and

from the breakthrough curves at 303 K for the faujasite materials. The adsorbed amounts of CH4 are averages of presaturated breakthrough curves and desorption curves. LTA-Type Zeolites. The breakthrough curves for the 50/50 CO2/CH4 mixture on the zeolite A materials are shown in Figure 5. The observed trends are similar to those described in the previous section. The quality of the separation and the volume of CO2 adsorbed increase as the framework charge in the sample is increased. Only the LTA-2 sample is an exception to this rule. Its CO2/CH4 separation is worse than that of LTA-3.5 because the breakthrough curve of CO2 on LTA-2 is highly dispersed. This is due to intracrystalline diffusion limitations. The high cation concentration in LTA-2 leads to the occupation of cation sites in the 8-membered ring, which strongly hinders the diffusion of CO2 through the window.43,44 Table 5 compiles the adsorbed amounts of CO2 and CH4 obtained from the breakthrough curves at 303 K for the LTA materials. The adsorbed amounts of CH4 are averages of presaturated breakthrough curves and desorption curves. Working Capacity and Selectivity. Tables 4 and 5 compile the results obtained for the working capacity and CO2/CH4 selectivity for the binary mixture. The working capacity between 5 and 1 bar was calculated by using the adsorption data of the mixture at 5 bar and of pure CO2 at 1 bar. In the case of samples strongly adsorbing CO2, the selectivity is affected by higher experimental uncertainties. This is because the amount of coadsorbed methane on NaX and EMC-1 is close to zero (within the experimental uncertainty), whereas for Na-USY, it rises up to ∼0.2 mmol·g−1 (at 5 bar). NaX and EMC-1 are highly selective, but their working capacities are low. The Na-USY and SAPO-37 samples present a lower selectivity but a higher working capacity. DAY has both poor selectivity and working capacity. For the LTA materials, the same trend is observed: the materials with a high framework charge have a low working capacity but high selectivity. The sample AlPO4-42, which does not incorporate extraframework cations, has both poor selectivity and working capacity. 5.4. Ternary Breakthrough Curves. FAU-Type Zeolites. Figure 6 shows the breakthrough curves for the CO2/CH4/CO mixture on the FAU samples. The curves reveal a good separation of CO2 from CH4 and CO on NaX and EMC-1 materials. The ternary mixture presents the same trends as those observed for the binary mixture concerning the quality of separation and quantity of CO2 adsorbed with respect to the amount of Na in the solid. An interesting feature that can be observed in the breakthrough curves is the fact that the order of elution changes from CO < CH4 < CO2 to CH4 < CO < CO2 as the framework electrostatic charge of the samples increases. This is because CO has a small dipole moment that interacts with the electric field of the adsorbent. In NaX, the adsorption of CO is sufficiently strong to produce a roll-up in the breakthrough curve of CH4.

Table 3. Henry Constants at 303 K on LTA Zeolites

a

solid

KCO2 × 103, molecules·bar−1· Å−3

KCH4 × 103, molecules·bar−1· Å−3

KCO × 103, molecules·bar−1· Å−3

AlPO4-42 LTA-5a LTA-3.5a LTA-2a

1.41 52.9 113.0 450.4

0.53 0.99 1.01 1.38

0.27

Calculated from the experimental data in ref 19.

LTA-5 samples, the Henry constants were taken from ref 19 and normalized with respect to the adsorption volume. The Henry constant of CO2 follows the order LTA-2 > LTA-3.5 > LTA-5 > AlPO4-42, which agrees with the order of decrease in the framework charge of the samples. KCH4 is rather constant for all of the LTA zeolites. These trends are in good agreement with those obtained for the FAU series. 5.3. Binary Breakthrough Curves. FAU-Type Zeolites. The breakthrough curves for equimolar CH4/CO2 mixtures at 303 K and 5 bar of the FAU series are shown in Figure 3. The V/Vcol ratio was expressed as a function of the ratio between the injected feed volume at a given time (V = F0t) and the volume of the column. Because this is a normalized representation, it allows a comparison of columns of different sizes. For all of the FAU samples, CH4 always breaks first because of its lower affinity with the adsorbent. The breakthrough curves of CH4 have a double roll-up. The main roll-up is due to the desorption of CH4 by incoming CO2. The superimposed second roll-up peak either is an artifact produced by the dead volume of the system or is provoked by the increase of the column temperature due to the exothermic adsorption of CO2. As can be judged by the separation between the CH4 and CO2 concentration fronts, the CO2/CH4 separation is excellent on NaX and EMC-1. The adsorbed quantity of CO2 is quite similar for these samples. The quality of the separation and the volume of CO2 adsorbed are reduced as the Na content (the framework charge) of the samples decreases. It should be noted that the samples with low Na content have a less steep breakthrough curve. The desorption curves are represented in Figure 4. Note that desorption at V/Vcol < 0 is due to the purging of the dead volume of the system. The true desorption begins at V/Vcol > 0. Figure 4 shows that the desorbed amount of CH4 is very low on NaX and EMC-1; i.e., these samples are highly selective. Moreover, it can be noted that the desorption of CO2 in NaX is slower because CO2 is more strongly adsorbed. In general, CO2 desorbs faster for samples with low Na content. Table 4 compiles the adsorbed amounts of CO2 and CH4 obtained

6. DISCUSSION 6.1. Correlation of the Henry Constant with the Selectivity and Working Capacity. A correlation between the Henry constant of cationic zeolites and the Na content is shown in Figure 7a,b for FAU- and LTA-type zeolites, respectively. The noncationic samples were excluded from the plot, i.e., DAY and AlPO4-42. For both zeolite types, there is a nearly linear relationship between the ln K value of CO2 and 9866



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Figure 3. 50/50 CO2/CH4 breakthrough curves at 5 bar and 303 K on (A) NaX, (B) EMC-1, (C) NaUSY, (D) SAPO-37, and (E) DAY.

pressure. However, they have a large selectivity. (ii) The samples with a low affinity for CO2, i.e., AlPO4-42 and DAY, have a low working capacity and selectivity. Finally, there is a trade-off zone where the adsorbent presents the best compromise between the selectivity and working capacity. Four samples are in this zone: the three FAU zeolites SAPO-37, EMC-1, and Na-USY and LTA-5. These FAU zeolites have Na concentrations between 1 and 6 wt % (0.035−0.2 Na cations per T atom). The experimental uncertainties in the separation of the ternary mixture were unfortunately too high to construct a similar figure. 6.2. Definition of the Optimal Henry Constant via the RSM. RSM versus Experiments. The experimental results discussed in the previous section perfectly reproduce the tradeoff between the cyclic capacity and selectivity that had been predicted. In order to formulate a mathematical description of this trade-off, we calculated the cyclic capacity and separation factor from the RSM. The fractional loading, according to the RSM, is

the extraframework cation concentration. At the same Na content, LTA solids exhibit a higher Henry constant than FAU solids. This is probably due to the different geometric arrangements of cations in the two structures. In the case of CH4, the Henry constant remains almost constant, i.e., K × 103 = 1.0 molecules·bar−1·Å−3, irrespective of the Na concentration and zeolite type. This is because of the weak electrostatic interaction between CH4 molecules and Na+ cations. CH4 interacts with the zeolites mainly via van der Waals interactions (dispersion−repulsion forces). These only depend on the sizes of the zeolite cages, which are similar for FAU- and LTA-type zeolites. The Henry constant of CO slightly increases with the number of cations in the unit cell because of the contribution of the dipole−charge interaction. This contribution is weak because the permanent dipole in the CO molecule is also very small. Figure 8 shows a plot of the CO2 working capacity for a cyclic process between 1 and 5 bar at 303 K as a function of the CO2 Henry constant. Three important trends are noted: (i) For samples that strongly adsorb CO2, like NaX and LTA-2, the working capacity is small because they tend to saturate at low 9867



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Figure 4. CO2/CH4 desorption curves at 5 bar and 303 K on (A) NaX, (B) EMC-1, (C) Na-USY, (D) SAPO-37, and (E) DAY.

Table 4. Adsorbed Amounts Measured Using Breakthrough Curves at 303 K for the 50/50 CO2/CH4 Mixture on FAU Zeolites with CO2/CH4 Selectivity at 5 bar and Working Capacity between 5 and 1 bar solid NaX EMC-1 NaUSY SAPO-37 DAY

pure CO2 nCO2 at 1 bar, mmol·g−1 4.3 4.8 1.2 2.8 0.5

± ± ± ± ±

0.1 0.2 0.1 0.3 0.1

mixture nCO2 at 5 bar, mmol·g−1 4.6 5.2 1.9 3.8 0.9

Δθ = θCO2(pads , XCO2,ads) − θCO2(pdes , XCO2,des)

± ± ± ± ±

0.4 0.2 0.1 0.2 0.1

mixture nCH4 at 5 bar·mmol·g−1

ΔnCO2 mmol·g−1

ΔnCO2, mol·m−3

αCO2/CH4

± ± ± ± ±

0.3 0.4 0.7 1.0 0.4

455 556 944 1200 530

61 41 12 14 5

0.07 0.1 0.16 0.3 0.21

0.10 0.1 0.04 0.1 0.04

first compared the calculated working capacity for the separation of CO2/CH4 with the experimental measurements. For this purpose, the experimental working capacity was transformed into a fractional capacity according to

(15)

where the θCO2(pads,Xads) is the CO2 coverage at the end of the adsorption step and θCO2(pdes,Xdes) is the CO2 coverage at the end of the desorption step. The experimental results had shown that changing the concentration of the extraframework cations affects the Henry constant of CO2 but not the Henry constant of CH4. For our modeling, we therefore varied the Henry constant of CO2 but fixed the Henry constant of CH4 to the experimental value found for the series of faujasite and zeolite A samples, i.e., 1 × 10−3 molecules·bar−1·Å−3. In order to validate the model, we

Δθ =

Δn nsat

(16)

where nsat is the CO2 saturation capacity obtained in the Langmuir fitting. Figure 9 shows that the RSM reproduces the general trend that has been found experimentally, although the agreement between the model and experiments lacks statistical significance 9868



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Figure 5. 50/50 CO2/CH4 breakthrough curves at 5 bar and 303 K on (A) LTA-2, (B) LTA-3.5, (C) LTA-5, and (D) AlPO4-42.

Table 5. Adsorbed Amounts Measured Using Breakthrough Curves at 303 K for the 50/50 CO2/CH4 Mixture on LTA Zeolites with CO2/CH4 Selectivity at 5 bar and Working Capacity between 5 and 1 bar solid LTA-2 LTA-3.5 LTA-5 AlPO4-42

single CO2 nCO2 at 1 bar, mmol·g−1 4.8 4.8 2.9 1.0

± ± ± ±

0.2 0.2 0.2 0.1

mixture nCO2 at 5 bar, mmol·g−1 4.7 5.0 3.6 1.25

± ± ± ±

0.3 0.1 0.1 0.1

mixture nCH4 at 5 bar·mmol·g−1

ΔnCO2 mmol·g−1

ΔnCO2, mol·m−3

αCO2/CH4

± ± ± ±

−0.1 0.2 0.7 0.2

−159 305 1049 240

306 19 20 6

0.02 0.3 0.2 0.2

0.1 0.1 0.2 0.1

In order to select the optimal adsorbent, we also have to calculate the separation factor. The RSM separation factor can be defined in terms of the coverage as follows:

because of the large error bars of the experimental working capacities. In order to calculate the separation factor, we also have to evaluate the separation selectivity from the RSM. The RSM predicts a selectivity that is close to the ratio of the Henry constant. Figure 10 shows the CO2/CH4 selectivity at 5 bar as a function of the ratio between the CO2 and CH4 Henry constants. It can be seen that the experimental coadsorption selectivity is in rough agreement with the theoretical selectivity, calculated from the ratio of the Henry constants. Optimal Henry Constant. Considering that the RSM reproduces the experimental trends, we can use it to theoretically define the optimal Henry constant for our CO2/ CH4 separation problem (making abstraction of the experimental uncertainties). In Figure 9, we have already seen that the cyclic capacity, under the given operating conditions, is maximal for a Henry constant of 5 × 10−3 molecules·bar−1·Å−3. On the left side of the maximum, the regeneration capacity is good, but the absolute adsorption capacity at 5 bar is too low. On the right side of the maximum, the adsorption capacity at 5 bar approaches the maximum saturation capacity nsat, but the regeneration capacity decreases.

SFCO2,CH4 =

ΔθCO2 XCH4,ads θCH4,ads XCO2,ads

(17)

In Figure 11, both the working capacity and separation factor (as calculated from the RSM) are plotted as a function of the Henry constant of CO2. The Henry constant that maximizes the separation factor is 50 × 10−3 molecules·bar−1·Å−3. Lower Henry constants lead to low adsorption selectivities. Higher Henry constants lead to low working capacities because regeneration of CO2 is too difficult. The optimal Henry constant in terms of the working capacity is 5 × 10−3, while the optimal Henry constant in terms of the separation factor is 50 × 10−3 molecules·bar−1·Å−3. The best overall compromise between the working capacity and separation factor will be found in the region between both maxima, in good agreement with the experiments (Figure 8). If we give the same importance to the separation factor and working capacity, the maximum of the function (Δθ*SF) gives a Henry constant that maximizes both functions at the same 9869



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Figure 6. Breakthrough curves for a 70/15/15 CO2/CH4/CO mixture at 5 bar and 303 K on (A) NaX, (B) EMC-1, (C) Na-USY, (D) SAPO-37, and (E) DAY.

by the adsorbent pore volume, the heat of adsorption is given by

time. This maximum was calculated using the data in Figure 7, and it is located approximately at 25 × 10−3 molecules·bar−1· Å−3. From Figure 7, we can interpolate that this optimal Henry constant is reached for a Na+ cation concentration of roughly 3.45% (0.1 Na cations per T atom) for FAU zeolites and less than 1% (less than 0.025 Na cations per T atom) for LTA zeolites (by extrapolation). Optimal Heat of Adsorption of CO2 at Zero Coverage. Once the optimal Henry constant of CO2 is known, the optimal heat of adsorption at zero coverage (ΔH0) of CO2 can be calculated according to K=

⎛ ΔS° ⎞ ⎛ ΔH ⎞ 1 0 ⎟ ⎟⎟ exp⎜⎜ − exp⎜⎜ ⎟ p° ⎝ R gas ⎠ ⎝ R gasT ⎠

ΔH0 = ΔS°T − R gasT ln(Kp°v)

(19)

where v is the volume of the cavity where the molecules are adsorbed. For LTA and FAU zeolites, the pore volume is close to 0.32 cm3·g−1; therefore, for the all-silica frameworks, the volume of the supercage cavity is 766 Å3·supercage−1. As a first approach, one can assume that the supercage volume is rather constant with respect to the Si/Al ratio. Also, it should be noted that small changes in the cavity volume have a small effect on the optimal heat of adsorption because of the logarithmic function. Thus, the optimal heat of adsorption of CO2 for a maximal working capacity (KCO2 = 5 × 10−3 molecules·bar−1· Å−3) is −27.6 kJ·mol−1. On the other hand, in order to maximize the separation factor (KCO2 = 50 × 10−3 molecules· bar−1·Å−3) a value of ΔH0 of CO2 equal to −33.4 kJ·mol−1 should be reached. As stated before, the choice of an adsorbent depends on the relative importance given to the separation factor and working capacity. Giving the same importance to separation factor and

(18)

where ΔS° is the entropy change upon adsorption relative to a pressure p°. Here, the reference state was placed at 1 bar. The typical entropy change of CO2 in microporous solids is placed between −60 and −100 J·mol−1·K−1.45−47 We have chosen an intermediate value of −80 J·mol−1·K−1 to carry out the calculations.48 Because we have normalized the Henry constant 9870



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Figure 9. Working capacity versus Henry constant of CO2. Adsorption conditions: 5 bar, 50% CO2, 50% CH4 at 303 K. Desorption conditions: 1 bar, 100% CO2 at 303 K. Symbols denote the experimental results and the solid line the RSM.

Figure 7. Correlation between the Henry constant at 303 K and the Na+ content for (a) FAU and (b) LTA cationic zeolites.

Figure 10. Experimental CO2/CH4 selectivity as a function of the KCO2/KCH4 ratio.

Figure 8. Working capacity and CO2/CH4 selectivity at 5 bar and 303 K as a function of the Henry constant of CO2 at 303 K. The size of the circle is proportional to the CO2/CH4 selectivity.

working capacity, corresponds to an intermediate value of KCO2 = 25 × 10−3 molecules·bar−1·Å−3, which is reached if ΔH0 is equal to −31.7 kJ·mol−1. The reported enthalpies of adsorption of CO2 on NaX is 35−45.49,50 Therefore, this solid presents a heat of adsorption higher that the optimal values. These optimal heats of adsorption of CO2 are in good agreement with our recent finding for MOFs with coordinative unsaturated sites.51 For the separation of CO2 by PSA between 1 and 5 bar, the MOFs known as Cu-BTC (or HKUST-1)52 and CPO-27-Zn are interesting adsorbents because they have both good selectivities18 and working capacities. Cu-BTC shows isosteric heats of adsorption CO2 of 30−35 kJ·mol−1.47,53 In the case of CPO-27-Zn, the heat of adsorption of CO2 is around 31 kJ·mol−1.48,54 Effect of the Operating Conditions. We use the model to study the effect of the operating conditions on the optimal properties of the adsorbent. For this purpose, the pressure during the adsorption step at high pressure was varied between

Figure 11. Working capacity and separation factor as a function of the Henry constant of CO2 at 303 K.

2 and 25 bar. Figure 12 shows a contour plot of the working capacity (Δθ) as a function of the adsorption pressure of the mixture and the Henry constant of CO2. The optimal Henry constant is always in the region between 5 and 50 × 10−3 9871



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ASSOCIATED CONTENT

S Supporting Information *

Gravimetric experiments and breakthrough measurements. This material is available free of charge via the Internet at http:// pubs.acs.org.

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AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected]. Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS M.P., F.R., and S.V. are thankful for financial support by the Spanish Government (MAT2012-38567-C02-01, CTQ201017988/PPQ, Consolider Ingenio 2010-Multicat CSD-200900050, and Severo Ochoa SEV-2012-0267).

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Figure 12. Working capacity as a function of the adsorption pressure of the 50/50 CO2/CH4 mixture and the Henry constant of CO2 at 303 K and KCH4 = 1 × 10−3 molecules·bar−1·Å−3.

LIST OF SYMBOLS b molecule volume F flow K Henry constant KL Langmuir constant m mass n adsorbed amount p pressure Prod productivity Rgas universal gas constant S entropy s number of molecules SF separation factor T absolute temperature t time v volume of the cavity vc virial coefficient Vf volumetric flow X molar fraction in the gas phase Y molar fraction in the adsorbed phase Z configurational integral

molecules·bar−1·Å−3; i.e., the adsorption pressure does not have a strong impact on the choice of the ideal adsorbent.

7. CONCLUSIONS The separation of CO2 from CH4 was studied by means of breakthrough experiments and theoretical calculations using the equilibrium theory of adsorption columns based on the RSM. We have found good agreement between the experimental trends and the RSM. The experimental results show clearly how the shape of the isotherm, selectivity, and working capacity change with the Na content for LTA and FAU zeolites. Three different kinds of behavior can be identified for these materials: (i) Zeolite with low Na content: these solids present a smooth CO2 isotherm (low Henry constant) and poor selectivity and working capacity. (ii) Zeolites with intermediate Na content: these adsorbents have intermediate Henry constants of CO2 and a good compromise between the selectivity and working capacity. (iii) Zeolites with high Na content: these solids have a steep CO2 isotherm (high Henry constant), a high selectivity, but a low working capacity. These different behaviors are due to the tuning of the CO2−adsorbent interaction introduced by changing the Na content via the Si/Al ratio. Therefore, the performance of the adsorbent for CO2 separation by PSA can be optimized simply by changing the framework charge. The RSM allows identification of the optimal Henry constant of CO2 under typical PSA conditions. This zone is placed between 5 × 10−3 and 50 × 10−3 molecules·bar−1·Å−3. The trade-off zone corresponds to materials presenting heats of adsorption of CO2 at zero coverage in the range between 27 and 32 kJ·mol−1. The experimental results show that, among the adsorbents used here, EMC-1 (Si/Al = 4), SAPO-37, Na-USY (Si/Al = 8.6), and LTA-5 (Si/Al = 5) are in the optimal zones. By using our experimental correlation, it was shown that the optimum Na content is roughly 3.45 wt % for FAU zeolites (and less than 1 wt % for LTA zeolites). This corresponds to a ratio of Na cations to T atoms of 0.1 for FAU (less than 0.025 for LTA). We emphasize that the methodology shown in this work can be transposed to other separations of small molecules and/or different operating conditions, leading to an intelligent design of the adsorption processes.

Greek Letters

μ α β ε θ ρ τ φ

first moment of the breakthrough curve selectivity 1/kT depth of the potential well coverage density reduced mass porosity

Superscripts

° reference state Subscripts

0 ads bulk col cycle des feed ff i,j max 9872

at zero coverage adsorbed phase bulk phase column PSA cycle desorption feed gas mixture fluid−fluid component i or j maximal dx.doi.org/10.1021/ie500207s | Ind. Eng. Chem. Res. 2014, 53, 9860−9874

Industrial & Engineering Chemistry Research sat

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(21) Wiersum, A. D.; Chang, J. S.; Serre, C.; Llewellyn, P. L. An Adsorbent Performance Indicator as a First Step Evaluation of Novel Sorbents for Gas Separations: Application to Metal−Organic Frameworks. Langmuir 2013, 29, 3301. (22) Rege, S. U.; Yang, R. T. A Simple Parameter for Selecting an Adsorbent for Gas Separation by Pressure Swing Adsorption. Sep. Sci. Technol. 2001, 36, 3355. (23) Ackley, M. W. Multilayer Adsorbent Beds for PSA Gas Separation. U.S. Patent 6,152,991, 2000. (24) Ruthven, D. M. Principles of Adsorption and Adsorption Processes; John Wiley & Sons: New York, 1984. (25) Palomino, M.; Corma, A.; Jorda, J. L.; Rey, F.; Valencia, S. Zeolite Rho: a Highly Selective Adsorbent for CO2/CH4 Separation induced by a Structural Phase Modification. Chem. Commun. 2012, 48, 215. (26) Shang, J.; Li, G.; Singh, R.; Xiao, P.; Liu, J. Z.; Webley, P. A. Determination of Composition Range for Molecular Trapdoor Effect in Chabazite Zeolite. J. Phys. Chem. C 2013, 117, 12841. (27) Shang, J.; Li, G.; Singh, R.; Gu, Q.; Nairn, K. M.; Bastow, T. J.; Medhekar, N.; Doherty, C. M.; Hill, A. J.; Liu, J. Z.; Webley, P. A. Discriminative Separation of Gases by a Molecular Trapdoor Mechanism in Chabazite Zeolites. J. Am. Chem. Soc. 2012, 134, 19246. (28) Derycke, I.; Vigneron, J. P.; Lambin, P.; Lucas, A. A.; Derouane, E. G. Physisorption in Confined Geometry. J. Chem. Phys. 1991, 94, 4620. (29) Derouane, E. G.; Andre, J. M.; Lucas, A. A. A simple van der Waals Model for Molecule-Curved Surface Interactions in MolecularSized Microporous Solids. Chem. Phys. Lett. 1987, 137, 336. (30) Garcia, E. J.; Perez-Pellitero, J.; Jallut, C.; Pirngruber, G. D. Modeling Adsorption Properties on the Basis of Microscopic, Molecular, and Structural Descriptors for Nonpolar Adsorbents. Langmuir 2013, 29, 9398. (31) Delprato, F.; Delmotte, L.; Guth, J. L.; Huve, L. Synthesis of New Silica-Rich Cubic and Hexagonal Faujasites using Crown-Ether based Supramolecules as Templates. Zeolites 1990, 10, 546. (32) Briend, M.; Lamy, A.; Peltre, M. J.; Man, P. P.; Barthomeuf, D. Thermal Stability of Tetrapropylammonium (TPA) and Tetramethylammonium (TMA) Cations Occluded in SAPO-37 Molecular Sieves. Zeolites 1993, 13, 201. (33) Moscoso, J. G.; Lewis, G. J.; Gisselquist, J. L.; Miller, M. A.; Rohde, L. M. Crystalline Aluminosilicate Zeolitic Composition: UZM9. U.S. Patent 6,713,041 B1, 2004. (34) Wadlinger, R. L.; Rosinski, E. J.; Plank, C. J. Synthetic zeolite. U.S. Patent 3,375,205, 1968. (35) Fayad, E. J.; Bats, N.; Kirschhock, C.; Rebours, B.; Quoineaud, A. A.; Martens, J. A Rational Approach to the Ionothermal Synthesis of an AlPO4Molecular Sieve with an LTA-Type Framework. Angew. Chem., Int. Ed. 2010, 49, 4585. (36) Keller, J.; Staudt, R. Gas Adsorption Equilibria. Experimental Methods and Adsorption Isotherms; Springer: Berlin, 2005. (37) Ruthven, D. M.; Loughlin, K. F. Sorption of Light Paraffins in Type-A Zeolites. Analysis and Interpretation of Equilibrium Isotherms. J. Chem. Soc., Faraday Trans. 1 1972, 68, 696. (38) Ruthven, D. M. Sorption of Oxygen, Nitrogen, Carbon Monoxide, Methane, and Binary Mixtures of these Gases in 5A Molecular Sieve. AIChE 1976, 22, 753. (39) Kane, M. S.; Leavitt, F. W.; Ackley, M. W.; Notaro, F. Pressure Swing Adsorption Process and Apparatus. U.S. Patent 6,245,127, 2001. (40) Sing, K. S. W. Reporting Physisorption Data for Gas Solid Systemswith Special Reference to the Determination of SurfaceArea and Porosity. Pure Appl. Chem. 1982, 54, 2201. (41) Maurin, G.; Bell, R.; Kuchta, B.; Poyet, T.; Llewellyn, P. Adsorption of Non Polar and Quadrupolar Gases in Siliceous Faujasite: Molecular Simulations and Experiments. Adsorption 2005, 11, 331. (42) Maurin, G.; Llewellyn, P. L.; Bell, R. G. Adsorption Mechanism of Carbon Dioxide in Faujasites: Grand Canonical Monte Carlo Simulations and Microcalorimetry Measurements. J. Phys. Chem. B 2005, 109, 16084.

at saturation

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Tuning the Adsorption Properties of Zeolites as Adsorbents for CO2

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