Ind. Eng. Chem. Res. 2003, 42, 1451-1461
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Measurement of Adsorption Equilibrium by the Zero Length Column (ZLC) Technique Part 1: Single-Component Systems Federico Brandani,† Douglas Ruthven,*,† and Charles G. Coe‡ Department of Chemical Engineering, University of Maine, Orono, Maine 04469-5737, and Air Products and Chemicals, Inc., Allentown, Pennsylvania 18195
This paper reports the development and experimental validation of a new technique to measure adsorption equilibrium isotherms, based on measuring ZLC desorption curves at low flow rate under carefully controlled conditions. The validity of the method is confirmed by replication of the equilibrium isotherms for representative systems (CO2-CaA, CO2-NaLSX) measured by a conventional volumetric/piezometric method. The new technique has been used to measure equilibrium isotherms for CO2 and other sorbates on a range of zeolitic adsorbents over a wide range of temperature and loading. The data from these measurements are discussed and interpreted in relation to available structural information concerning cation locations in various X-type zeolites. Adsorption equilibrium measurements have traditionally been carried out by gravimetric, volumetric, or piezometric methods. These techniques are straightforward and accurate, but they are time-consuming and are therefore not well-suited to adsorbent screening studies, which require the rapid (approximate) characterization of many different materials. For initial screening, chromatographic methods are generally more convenient. Such methods have been widely used for the approximate measurement of Henry constants (see, for example, Choudhary and Mayadevi1,2) and have also been adapted to the measurement of complete isotherms for both single components and binary mixtures.3-5 The zero length column (ZLC) technique was introduced, some years ago, as a simple and rapid approach to the study of sorption kinetics.6,7 Approximate equilibrium constants are also obtained and have been shown to be consistent with directly measured values.8 In a traditional ZLC experiment, the measurements are carried out at high flow rates such that the effluent concentration is determined by the rate at which the sorbate diffuses out of the adsorbent particles. If the same kind of experiment is performed at a sufficiently low flow rate, the desorption rate will be determined by convection under equilibrium conditions, rather than by the desorption kinetics. The effluent concentration history then directly yields the equilibrium isotherm. The possibility of using this approach to measure Henry constants and equilibrium isotherms was suggested at a recent conference.9 The purpose of the present research was to explore, in greater detail, the potential and limitations of this approach for the study of several hydrophilic adsorbents of practical interest. Theory In a ZLC experiment, a small sample of the adsorbent is equilibrated with the sorbate at a known partial pressure in an inert carrier stream (generally He). At time zero, the feed is switched to the pure carrier (at * Corresponding author. † University of Maine. ‡ Air Products and Chemicals, Inc.
constant flow rate), and the composition of the effluent stream is monitored using a sensitive detector. The “column” is short enough (only a few layers of adsorbent) to be treated as a well-mixed cell. [Axial mixing in a packed column is controlled by the reciprocal axial Peclet number (DL/vL). As L f 0, (DL/vL) f ∞, so the system behaves as a well-mixed cell.] The differential mass balance is therefore given by
Vs
dq dt
+ Vg
dc + Fc ) 0 dt
(1)
For a trace system, the flow rate (F) remains constant, and the integration of eq 1 is straightforward. Linear Equilibrium. If the purge flow rate is sufficiently low that equilibrium is closely approached and the isotherm is linear, then
q ) q* ) Kc
(2)
With the initial condition
t ) 0, q ) q0 ) Kc0
(3)
the concentration response curve is given by
ln(c/c0) )
-Ft KVs + Vg
(4)
A plot of ln(c/c0) vs time (or the product Ft) should therefore yield a straight line through the origin with slope -1/(KVs + Vg). If the sorbate is strongly adsorbed (K large) such that KVs . Vg, this relation directly yields the value of K (the dimensionless Henry constant). For weakly adsorbed species, for which KVs and Vg are of similar magnitude, it is necessary to correct for the dead volume of the cell (Vg). This correction is easily obtained by repeating the experiment with no adsorbent present in the cell. Langmuir Equilibrium. Beyond the Henry’s Law region, the isotherm can commonly be approximated by the Langmuir expression
10.1021/ie020572n CCC: $25.00 © 2003 American Chemical Society Published on Web 03/04/2003
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q* qs
)
bc 1 + bc
(5)
To show the qualitative effect of the isotherm shape on the ZLC response, we integrate eq 1 under equilibrium conditions [dq j /dt ) (dq*/dc)(dc/dt)] with dq*/dc given by the differential of eq 5. This yields
ln
()
-Ft c ) c0 KVs + Vg KVs 1 + bc0 1 1 + ln KVs + Vg 1 + bc 1 + bc0 1 + bc
[
(
)]
(6)
In the long-time region, c f 0, and a plot of ln(c/c0) vs Ft approaches a linear asymptote of slope -1/(KVs + Vg) with a negative intercept given by[KVs/(KVs + Vg)][1 - 1/(1 + bc0) + ln(1 + bc0)]. The Henry constant can therefore be found from the slope of the long-time asymptote. However, for highly nonlinear systems (bc0 . 1), this might not be practical, as the long-time asymptote will be buried in the baseline. Under equilibrium conditions, the quantity of sorbate desorbed depends only on the total volume of purge gas passed through the cell. The ZLC response curve plotted in terms of the product Ft is therefore invariant with flow rate. This is in contrast to the situation that prevails under conditions of kinetic control, where the response curve plotted in terms of absolute time should be invariant with flow rate. Comparison between the response curves (in terms of Ft) for different purge flow rates therefore provides a simple experimental check on the validity of the equilibrium assumption. The same test should also confirm the absence of significant heat effects because, under nonisothermal conditions, the response becomes flow-rate-dependent. Calculation of the Complete Isotherm. Implicit in the derivation of eqs 4 and 6 is the assumption that the flow rate remains constant. This is a valid approximation for a trace system (small mole fraction of sorbate in the carrier stream), but to extend the analysis to higher sorbate concentrations, it is necessary to allow for the variation of the effluent flow rate as desorption proceeds. Equation 1 can be rewritten in the more general form
FC
y 1-y
) -Vs
dq* dt
- VgC
dy
(7)
dt
where y represents the mole fraction of sorbate in the effluent gas stream. Integration from t ) 0, y ) y0 yields
q* )
y/y
y/y
0 dt ∫0∞1/y0 -0y/y0 dt - FC ∫t Vs 0 1/y0 - y/y0
FC Vs
Vg y c (8) V s 0 y0 Integration of the desorption curve (y/y0 vs t) thus yields the equilibrium isotherm q* vs c (because c/c0 ) y/y0 and c0 ) y0P/RT). The first integral represents q0*, the adsorbed-phase concentration at equilibrium with concentration c0 in the gas phase. To evaluate this integral, complete desorption of the sample is required. This can present a practical problem for strongly adsorbed species because, when K is large, the slope of the long-time asymptote (eq 6) is very small. The desorption curve then has a long tail that can be difficult to distinguish
Figure 1. Zero length column (ZLC) experimental setup.
from the baseline. This problem becomes increasingly severe as the strength of adsorption increases, thus restricting the practical usefulness of this technique to weakly or moderately strongly adsorbed species. The completeness of desorption can, however, be easily checked experimentally simply by increasing the temperature to carry out a thermal desorption at the end of the run. If no significant TPD peak is observed, then the isothermal desorption process must have been essentially complete. If there is a (small) TPD peak, then the quantity desorbed can simply be added to the quantity found by integration of the ZLC response as a correction to the initial loading. Experimental Section Figure 1 shows a simplified schematic diagram of the experimental system. The ZLC cell consists of a small sample of zeolite adsorbent (1-5 mg) sandwiched between two sintered disks and contained within a 1/8-in. Swagelok fitting. The switch valve is a standard hightemperature Valco valve, operable up to 250-300 °C. A small electric heater with its own temperature control surrounds the ZLC cell, thus allowing the adsorbent to be regenerated in situ at temperatures up to 450 °C without exposing the valve to these temperatures. The detector is an on-line quadrupole mass spectrometer (Ametek model MA200HDEF with a Balzers turbomolecular pump). The gas leaving the cell is sampled continuously through a silica capillary connected to the chamber of the mass spectrometer. The effluent concentration response curves are recorded in digital form for data processing. The flow rates of both the feed and purge streams are controlled by mass flow controllers. Preliminary experiments quickly revealed that, for hydrophilic adsorbents, it is important to pay careful attention to dehydration of both the feed and purge streams, as well as to regeneration of the adsorbent. Various arrangements for drying the gases were tried, but the system that worked best utilized individual drying columns packed with 5A zeolite in each of the gas lines. These columns were contained in a separate oven which allowed regeneration at up to 400 °C. During operation, the drying columns were immersed in ice to enhance their affinity for water vapor. A limited series of experiments using dry ice showed no significant difference. We concluded that, with well-regenerated
Ind. Eng. Chem. Res., Vol. 42, No. 7, 2003 1453 Table 1. Zeolite Adsorbents Used in This Study zeolite
Si/Al
cationic form
cations/ supercage
(molecules/cage)/ (mmol/g)
LiLSXa NaLSXa NaXa CaX CaA
1.0 1.0 1.25 1.25 1.0
Li+ Na+ Na+ Ca2+ Ca2+
12 12 10.6 5.3 6
1.8 2.0 1.95 1.65 1.69
a Containing about 18% clay binder. Other samples are binderless. All samples are fully cation exchanged.
Table 2. Summary of van’t Hoff Constants Giving Temperature Dependence of Dimensionless Henry Constants According to Eq 9 sorbate CO2
N2 CO CH4
sorbate silicalite Graham et al.10 CaA CaA (Air Products) 5A (Haq and Ruthven11) CaX NaX NaLSX LiLSX 5A11 5A12 CaA (Air Products) CaX NaX
temp K0 × 103 -∆U0 range (K) dimensionless (kJ/mol) 298-373 273-343 323-448 348-448 420-465
59.2 33.0 10.4 7.0 6.5
19.8 20.4 38.4 40.4 32.8
323-423 312-398 293-423 308-448 293-360 200-300 263-303 308-398 195-273
0.6 1.2 2.4 0.5 4.6 4.4 2.2 0.1 0.1
50.5 40.6 39.9 48.6 20 18.8 23.8 40.9 26.0
0.15 mL. Several different zeolite-based adsorbents were studied; details are given in Table 1.
Figure 2. Blank ZLC response curves for (a) CO2-He and (b) CH4-He.
drying columns, operation at 0 °C is sufficient to ensure satisfactory dehydration of the gas stream. Measurements were carried out either using a premixed feed stream (as indicated in Figure 1) or with online mixing of the sorbate and carrier. The former arrangement requires careful preequilibration of the feed drying column, which can be time-consuming. The latter arrangement requires an additional flow controller and a drying column for the sorbate but is perhaps more satisfactory because this avoids the lengthy preequilibration of the drying column. Figure 2 shows the results of representative blank experiments, carried out with no adsorbent present in the cell. The response curves plotted against the product Ft are clearly independent of flow rate or the nature of the sorbate and show the expected exponential decay, in conformity with eq 4. The dead volume derived from the slope is about
Results and Discussion Linear Systems/Henry Constants. Representative response curves showing plots of ln(c/c0) vs Ft are displayed in Figure 3. In conformity with eq 4, in these coordinates, the response curves are essentially linear and independent of flow rate, thus confirming both the validity of the equilibrium approximation and the linearity of the isotherms. The small curvature in the initial region probably arises from uncertainty in the zero time resulting from the small delay in the gas sampling system. Measurements were made for several different systems over a range of temperature. Examples of the van’t Hoff plots showing the temperature dependence of the dimensionless Henry constant are presented in Figure 4. The resulting parameters, calculated from the van’t Hoff expression
K ) K0 exp(-∆U0/RT)
(9)
are summarized in Table 2. The ZLC Henry constants for CO2-silicalite (Figure 4a) are generally consistent with, although slightly higher than, the gravimetric data reported by Graham et al.10 The ZLC data for N2-5A (Figure 4b) are also broadly consistent with, although slightly higher than, previously reported data for this system and agree well with the values derived from the equilibrium isotherms measured on samples of the same zeolite in the Air Products laboratory. The strong effect of traces of water is evident from the CO2-CaA data shown in Figure 4c. Early ZLC measurements yielded Henry constants slightly larger than the values derived from the chromatographic measurements of Haq and Ruthven.11 However, careful
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Figure 3. ZLC response curves for (a,b) CO2-silicalite at 25 °C and (c,d) CH4-NaLSX at 0 °C plotted against time and against the product Ft.
measurements carried out in the Air Products laboratory (by the volumetric/piezometric method) yielded substantially higher K values, which were replicated by the ZLC data only after incorporation of the drying columns (at 0 °C) in the purge and feed streams and provision for in situ regeneration of the adsorbent at elevated temperature. The good agreement finally achieved provides convincing evidence of the validity of the ZLC approach. The van’t Hoff plots for CH4, CO, and CO2 on various different forms of zeolite X are shown in Figure 5; the derived parameters are included in Table 2. The ∆U0 values for CO2 in A and X zeolites show clearly how the difference in the strength of the most favorable adsorption sites depends on the nature of the cation. The sequence, for the X zeolites, correlates well with structural data relating to cation positions, as shown in Table 3. However, the large difference in ∆U0 between CaX and CaA seems surprising, because the cation sites (type II) in these two materials are similar. It can be seen that the adsorption energies in these (cationic) zeolites are all substantially greater than those in silicalite
(nonpolar), thus providing convincing evidence of the importance of electrostatic forces in these systems. From Figure 5a, it is clear that the Henry constants and adsorption energies for methane on NaX and NaLSX are almost identical. Because methane is nonpolar and has a rather small polarizability, the adsorption is dominated by the van der Waals interaction with the framework; the cations have very little effect. Because NaX and NaLSX have the same framework and differ only in the occupancy level of the cation sites, the similarity in their adsorption behavior is understandable. Nonlinear Systems. As an example of the behavior of a nonlinear system, Figure 6 shows the response curves for CO2-CaX at 75 °C at three different partial pressures of CO2. The response curves, plotted in terms of Ft, are independent of flow rate, showing that, in all cases, the equilibrium assumption is a valid approximation. As the initial pressure (and therefore the degree of nonlinearity of the isotherm) increases, the response curves show an increasingly strong deviation from linearity and a corresponding decrease in the intercept
Ind. Eng. Chem. Res., Vol. 42, No. 7, 2003 1455
Figure 4. van’t Hoff plots showing temperature dependence of the Henry constants for (a) CO2-silicalite, (b) N2-5A (CaA), and (c) CO2-5A (CaA).
Figure 5. van’t Hoff plots showing the temperature dependence of the Henry constants for (a) CH4 in NaX and NaLSX, (b) CO2 and CO in CaX, and (c) CO2 on various forms of X zeolite.
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Figure 6. ZLC response curves, plotted against Ft, for CO2-CaX at 75 °C for (a) pCO2 ) 12 Torr, (b) pCO2 ) 24 Torr, and (c) pCO2 ) 47 Torr showing lack of dependence on flow rate and conformity with the Langmuir model (qs ) 0.91 m‚mol/g, b ) 17 atm-1). Table 3. Variation of Limiting Adsorption Energies for CO2 at Low Loading Levels in Different Cationic Forms of Zeolite X material
-∆U0 (kJ/mol)
dominant site
CaX LiLSX NaX NaLSX silicalite10
50.5 48.6 40.6 39.9 20.0
Ca2+ in site II Li+ in site III Na+ in site II Na+ in site III framework
of the long-time asymptote. All of the response curves can be adequately represented by eq 6 with a single pair of Langmuir constants (qs ) 0.91 m‚mol/g, b ) 17 atm-1). These runs were carried out early in the study with an incompletely dehydrated adsorbent. As a result, the value of b is lower than that obtained later for the CO2-CaX system under dry conditions (∼350 atm-1). It should be noted that the Langmuir model is used here merely as a qualitative guide to show the effect of nonlinearity on the form of the ZLC response curves. In the detailed analysis of the experimental data, the response curves were all processed by direct integration (eq 8). Complete Isotherm Determination. Figure 7a-c shows examples of the complete isotherms for CO2 on CaA and on NaLSX calculated by integration of the ZLC response curves. For these systems, the isotherms were also measured by the piezometric/volumetric method, using the same adsorbent samples, at the Air Products laboratory. In all cases, the ZLC isotherms lie close to the piezometric/volumetric data, thus confirming the validity of the ZLC approach. For CO2-NaLSX, the recently reported experimental data of Siperstein and Myers13 and Shen et al.14 are also shown. The Shen and Siperstein isotherms are close to both the ZLC and the Air Products data, suggesting that their NaX adsorbent must have been very similar to our NaLSX material. The isotherm for our regular NaX sample (at 75 °C) is substantially lower.
Figure 7. Comparison of isotherms derived by integration of ZLC response curves with independently measured volumetric/ piezometric data. (a) CO2-CaA showing comparison of ZLC data (filled symbols) with volumetric/piezometric data obtained by Air Products for the same adsorbent sample (open symbols). (b) CO2-NaLSX at 20 °C showing comparison with volumetric/ piezometric data obtained by Air Products for the same adsorbent sample and with reported data of Siperstein and Myers13 for a different NaX sample. (c) CO2-NaLSX and NaX at 75 °C showing comparison with the Shen and Siperstein isotherm.14
Ind. Eng. Chem. Res., Vol. 42, No. 7, 2003 1457
Figure 8. Equilibrium isotherm for CH4-NaX at subambient temperature measured by integration of the ZLC response curves. Also shown are the high-pressure ambient-temperature data reported by Hayhurst et al.15 and Kobayashi.16
As an example of isotherm measurement over a very wide loading range, Figure 8 shows the data set obtained for CH4-NaX at subambient temperatures. The 195 K isotherm extends from the Henry’s law region almost to the saturation limit. Also shown are the highpressure isotherms obtained at 298 K by Zheng et al.15 and Rolniak and Kobayashi,16 which appear to be broadly consistent with the present data and show a similar saturation limit. Analysis and Interpretation of CO2 Isotherms. Figures 9 and 10 show sets of isotherms for CO2 on NaX, CaX, NaLSX, LiLSX, and CaA measured at several temperatures by the ZLC method. The shapes of the isotherms for these adsorbents are obviously different. It seems reasonable to assume that, at least in the low to intermediate loading region, the quadrupolar CO2 molecules are adsorbed mainly at the cation sites within the supercage. The slopes of the isotherms thus reflect differences in both the numbers and positions of the supercage cations. In zeolite X, the cations within the supercage occupy two different types of sites (II and III in the standard notation),17 so the simplest isotherm model that can be expected to capture the essential features of the physical system is the double Langmuir expression
q)
b1qs1 p b2qs2 p + 1 + b1 p 1 + b2 p
(10)
where b1, qs1 and b2, qs2 are the characteristic parameters for sites of types 1 and 2, respectively. The first term is assumed to represent adsorption on the “strong” cation sites within the supercage, while the second term represents adsorption on weak cation sites plus any additional contribution from delocalized adsorption or adsorption on the framework. If there is a large difference in affinity between sites 1 and 2, one can expect that the most favorable (type 1) sites will be predominantly occupied in the lowconcentration region, with significant occupation of the type 2 sites only at higher loading. To interpret the experimental data, we therefore fitted the low-concentration region of the isotherm (from zero to just beyond the linear region) to the single Langmuir model (eq 5)
Figure 9. Equilibrium isotherms for CO2 on Na+ and Li+ zeolites showing comparison of experimental data and model isotherms calculated with parameters given in Table 4. (a) NaLSX, ideal Langmuir model (eq 5); (b) LiLSX, dual-site model (eq 10); (c) NaX, ideal Langmuir model (eq 5).
subject to the restrictions that, for a given sorbent, qs1, which represents the density of strong cation sites, must be the same at all temperatures and the product bqs must coincide with the Henry constant (with appropriate correction for units). For the systems that showed substantial deviations from the simple Langmuir model, these values (b1, qs1) were then substituted into eq 10, and the parameters b2 and qs2, corresponding to the “weak” sites, were estimated by fitting the complete isotherms. Minor readjustments were then carried out to optimize the fit of the entire data set with constant
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Table 4. Langmuir Constants and Saturation Capacities for CO2 Adsorption on Zeolite X NaLSX (atm-1)
T (K)
b1
293 308 312 323 328 348 373 398 423 448
138.2 68 21.5 7.9 3.2 1.6 -
-∆H1a (kJ/mol) b01 (atm-1) qs1 (m‚mol/g) qs2 (m‚mol/g)
45.5 9.0 × 10-6 2.75
LiLSX b1
b2
1462.5 312.5 100.0 35.6 15.0 5.0 3.8
NaX (atm-1) 63.5 5.4 0.3 -
48.3 6.0 × 10-6 1.6 2.3
b1
(atm-1)
80.0 39.3 18.5 6.7 2.9 39.9 2.0 × 10-5 1.5
CaX b1
b2
1606.0 321.4 105.9 35.7 14.1 -
CaA (atm-1) 11.4 3.0 1.1 0.3 0.1 -
53.3 4.0 × 10-6 1.4 2.2
b1
b2 (atm-1)
800.0 250.0 83.0 37 16.0 8.6 -
25.0 5.0 1.25 0.36 -
43.9 6.0 × 10-5 0.52 2.8
a ∆H values given in Table 4 are derived from the temperature dependence of b . These values are similar, but not identical, to the 1 1 values of ∆H0 derived from the values of ∆H0 given in Table 2 (∆H0 ) ∆U0 - RT), which were obtained directly from the temperature dependence of the Henry constants.
Figure 10. Equilibrium isotherms for CO2 on Ca2+ zeolites showing comparison of experimental data and dual-site model isotherms (eq 10) calculated with parameters given in Table 4. (a) CaX, (b) CaA.
values of qs1 and qs2. The resulting parameters are summarized in Table 4, and comparisons between the experimental data and the model isotherm are shown in Figures 9 and 10. It can be seen that the data for NaX and NaLSX can be reasonably well approximated by the ideal Langmuir model (eq 5) with qs independent of temperature. In contrast, the ideal Langmuir model cannot adequately fit the data for LiLSX, CaX, and CaA. The isotherms for these sorbates are, however, reasonably well represented by the dual-site model (eq 10). If the values of qs1 and qs2 are allowed to vary with temperature, it is possible to obtain much better fits of all of the isotherms, but that approach would violate the physical basis of the dual-site Langmuir model.
According to the dual-site model the Henry constant is given by K ) b1qs1 + b2qs2 (with due correction for units). The Henry constants calculated in this way are close to but not identical to the values given in Table 2 (and Figures 4 and 5) which were derived directly from the ZLC response curves. The locations of the cations in zeolite X have been studied extensively by X-ray diffraction, neutron diffraction, and other techniques.17,24 The available information is summarized in Table 5. In principle, it is possible to compare the qs values derived from the CO2 isotherms with the values that might be anticipated from structural considerations (the number of cations in the most favorable sites). However, it is well-known that the distribution of cations between the available sites depends not only on the nature of the cation and the Si/Al ratio (which determines the total number of cations) but also on details of the sample dehydration and regeneration procedures,25 so such comparisons are not unambiguous. Further structural information can be obtained from detailed calorimetric studies,14,25-27 by molecular modeling,28,29 and MAS NMR spectroscopy.30,31 As for the diffraction data, the results of the calorimetric studies are somewhat ambiguous. For LiX, Shen and Bulow25 found that, in the low-loading region (up to 2.5 molecules/ cage), the heat of adsorption is constant. This is consistent with structural information that suggests that, for this sample (Si/Al ) 1.28) with the idealized cation distribution, there should be 2.5 Li+ cations per cage in the exposed type III sites. However, Barrer and Gibbons,27 for an apparently similar LiX sample, found a continuous decline of the heat of adsorption with CO2 loading (at least above 1.7 molecules per cage), suggesting a more random distribution of cations among the available sites. Similarly, for NaX and NaLSX, most studies report a continuous decline in heat of sorption with loading,14,25-27 although a few of the studies suggest that the heat of sorption approaches a constant value at low loadings.25 Information concerning the accessibility of the cations to the guest molecules has been obtained from MAS NMR spectra29,30 with both Li6 and Li7 and from infrared spectroscopic studies.31,32 It seems clear from these results that Li+ in site II is essentially buried (at the center of the oxygen six ring) and is inaccessible to either N2 or O2 (and presumably also to CO2). Li+ occupies site III only at high degrees of Li+ exchange
Ind. Eng. Chem. Res., Vol. 42, No. 7, 2003 1459 Table 5. Structural Data, Cation Positions, and Adsorption Capacities for X-Type Zeolites zeolite
LiLSX
LiX
NaLSX
NaX
NaX
NaX
NaX
CaX
CaX
Si/Al cations/cage
1.0 12
1.28 10.5
1.0 12
1.1 11.5
1.1 11.5
1.2 11
1.25 10.6
1.0 6
1.25 5.3
site Ia,b I′ a,b II′ a,b IIb,c III + III′b,c ref
0 4 0 4 4 Plevert22,23
0 4 0 4 2.5 Shen25
0 4 0 4 4 -
0 4 0 4 3.5 Zhu20
0.36 3.64 0 3.85 3.65 Olson19
4 0 4 3 Vitale21
0 4 0 4 2.6 -
1.66 0.63 0.75 3.1 0 Breck17
1.66 0.63 0.75 2.3 0 d
qs1e qs2e qs1 + qs2e
2.9 4.1 7.0
5.5
2.9 (4) 6.9
2.3 3.6 5.9
a Site I, hexagonal prism; sites I′ and II′, inside sodalite cage near the 6-rings. All inaccessible to gas. b See Breck17 or Mortier18 for a more detailed description of the cation sites. c Site II, within supercage at the 6-rings; sites III and III′, within supercage at the 4-rings. All accessible to gas. d Calculated assuming that the occupations of sites I, I′, and II′ are the same as for the lower Si/Al sample. e qs is expressed in molecules/cage. Conversion factors (molecules/cage)/(mmol/g) are given in Table 1.
and when the Si/Al ratio is sufficiently low, but these are clearly the strong adsorption sites in LiLSX or LiX. By contrast, Na+ and Ca2+ in site II appear to be accessible to the gas molecules. In fact, for Na+, there appears to be little difference in interaction energy for a guest molecule with cations in sites II and III. CO2)NaLSX (Figure 9a). The structural data suggest that this material should contain four Na+ cations in type III sites and four in type II. The fact that cations in both of these sites are accessible to the gas molecules with little energetic difference would suggest up to eight strong sites per cell. The experimental data are well represented by the ideal Langmuir model and suggest 5.5 equivalent (strong) sitessa much higher value than for LiLSX. The sorption data are thus seen to be qualitatively (although not quantitatively) consistent with the structural information. It seems possible that a fraction of the cations attributed to sites II and III in the idealized structure might, in fact, be located in the interior sites (I, I′, II′). CO2)LiLSX (Figure 9b). For the ideal cation distribution shown in Table 5, one would expect four strong sites per cage corresponding to Li+ in the type III sites. In fact, the isotherm data suggest about 2.9 strong and 4.1 weak sites per cage. This might be interpreted as suggesting that, in practice, the type III (and III′) sites are somewhat less fully populated than in the idealized structure. CO2)NaX (Figure 9c). For a given CO2 pressure and temperature, the equilibrium loading on NaX is substantially smaller than that for NaLSX, reflecting the lower cation density (see Figure 7c). Analysis of the experimental isotherms suggests about three strong sites and four weaker sites per cage (qs1 ) 2.9, qs2 ) 4.0). These values are close to the expected numbers of cation sites III (strong) and II (weak), respectively. Furthermore, the equilibrium constants for the strong sites (b1) for NaX and NaLSX are very similar, suggesting that the difference in affinity is related to the difference in site densities, rather than to differences in the intrinsic affinity of an individual site. These conclusions are broadly consistent with the structural data except that, in NaLSX, the strengths of the type II and type III cation sites appear to be similar, whereas in NaX,26 there is a measurable difference. CO2)CaX (Figure 10a). The experimental isotherms are well represented by the dual-site Langmuir model with qs1 ≈ 2.3 sites per cage and qs2 ≈ 3.6 sites per cage. This is reasonably consistent with the struc-
tural information (2.3 cations per cage in site II) if it is assumed that the occupancy of the internal sites is the same as for the lower-silica material. The data suggest some additional adsorption on sites II, but the b2 values are very much smaller (by 2 orders of magnitude) than the b1 values. This might suggest that, for this system, the second set of sites represents the framework rather than less favorable cation sites. CO2)CaA (Figure 10b). As for CaX, the experimental isotherms are poorly represented by the simple Langmuir model, but they are reasonably well represented by the dual-site model with qs1 ≈ 0.9 and qs2 ≈ 4.8 sites per supercage. For the idealized structure, in fully Ca2+-exchanged form, there should be six Ca2+ ions per supercage, so the total (qs1 + qs2) agrees well with the total number of accessible cations. Because the six Ca2+ cations are all located in equivalent six-ring sites, one might expect an ideal Langmuir isotherm with qs ) 6. However, more detailed consideration shows that the distribution of six Ca2+ cations between the eight equivalent six-ring sites actually yields two slightly different types of site, depending on whether all adjacent six-ring sites are occupied or whether there is a vacancy at one of these sites. Recent infrared spectroscopic studies, carried out at the Air Products laboratory, confirm the existence of two different sites with slightly different energies, as suggested by the isotherm data. If a random distribution of vacancies is assumed, the probability that any given cation will have all adjacent sites occupied is much smaller than the probability that one of the sites will be vacant, so the numerical imbalance between the site densities is understandable. Effect of Water Vapor. The major problem with the ZLC technique is the need to eliminate all traces of water vapor. With hydrophobic adsorbents such as silicalite or even with hydrophilic adsorbents such as CaA or NaX, this difficulty can be overcome relatively easily by careful attention to the obvious details of the experimental system. However, with strongly hydrophilic adsorbents such as LiLSX, for which the effect of part per million (or even part per billion levels) of water can be significant, the problem can be more severe. Figure 11 shows the effect on the CO2 isotherm of purging for several hours with cylinder He (without a molecular sieve dryer in the line). There is a perceptible decline in capacity within a few hours, and after 2 days, the capacity and affinity fall to very low levels characteristic of a noncationic system, suggesting that all cation sites have been deactivated by water adsorption.
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Ind. Eng. Chem. Res., Vol. 42, No. 7, 2003
Figure 11. Isotherms for CO2-LiLSX at 35 °C showing deactivation with time on stream (1) for freshly regenerated sample, (2) after 5 h, (3) after 24 h, (4) after 27 h, and (5) after 48 and 72 h.
isotherms and Henry constants reported here should reflect only the physically (reversibly) adsorbed CO2. The limiting heats of sorption derived from the temperature dependence of the Henry constants, therefore, do not show the very high values (up to 100 kJ/mol) found from calorimetric measurements at very low loading. The low and intermediate concentration regions of the isotherm for these systems can be interpreted assuming that adsorption occurs predominantly on the cations in the supercage. The site densities derived by application of the simple Langmuir and dual-site Langmuir models to the CO2 isotherms are qualitatively consistent with the available information on cation site occupancy levels derived from structural data. For CaA, the isotherm data suggest 5.7 sites per supercage, which is close to the number of accessible Ca2+ cations (6 per supercage). For the NaX and CaX samples, the correspondence between the density of strong adsorption sites and the expected cation occupancies of sites III + III′ and site II is almost quantitative. For NaLSX and LiLSX, the agreement is not as close, although the cation occupancy levels and adsorption site densities are still similar. Notation
Figure 12. Replicate isotherms for CO2-LiLSX at 35 °C showing no decline in capacity after successive regeneration at 400 °C.
Heating under He purge at 400 °C for a relatively short period (30 min) restored the capacity and affinity to the original level for a water-free sample. Indeed, the isotherms determined after successive regenerations at 400 °C showed no evidence of any permanent decline in capacity (see Figure 12). Conclusions The validity and viability of the ZLC technique for measuring Henry constants and adsorption equilibrium isotherms has been confirmed, subject to the requirement that, for strongly hydrophilic adsorbents, it is essential to maintain a dry system. The Henry constants reported here for several hydrophilic systems (CO2CaA, CO2-CaX, CO2-NaLSX, CO2-LiLSX) are consistent with recent data for these systems obtained by conventional piezeometric/volumetric methods but substantially higher than most of the earlier reported values, suggesting that the earlier studies might have been impacted by the presence of traces of water vapor. There is considerable spectroscopic33,34 and calorimetric34,35 evidence for the formation of trace amounts of chemisorbed CO2, particularly in the sodium forms of zeolites A and X. Because the ZLC measurements are made in desorption and the strongly chemisorbed species are not desorbed except at high temperature, the
b ) Langmuir equilibrium constant (concentration-1 or atm-1) c ) sorbate concentration in gas phase (mol/mL) c0 ) initial (steady) value of c C ) total molecular density (mol/mL) q ) adsorbed-phase concentration (mol/kg) qs ) saturation capacity in the adsorbed phase for Langmuir equation (mol/kg) DL ) axial dispersion coefficient (cm2‚s-1) F ) effluent flow rate or inert flow rate (mL‚s-1) K ) Henry’s law constant (dimensionless or m‚mol/g‚atm) K0 ) preexponential factor van’t Hoff expression (eq 9) L ) “length” of ZLC p ) partial pressure of sorbate (atm) t ) time v ) interstitial gas velocity (cm‚s-1) Vs ) volume of solid in the ZLC cell Vg ) volume of gas in the ZLC cell y ) mole fraction in the gas phase
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Received for review July 29, 2002 Revised manuscript received January 23, 2003 Accepted January 23, 2003 IE020572N
