HOWARD S. SHERRY
4086
The Ion-Exchange Properties of Zeolites.
IV.
Alkaline Earth Ion
Exchange in the Synthetic Zeolites Linde X and Y by Howard S. Sherry Mob2 Research and Development Corporation, Research Department, Central Research Division, Princeton, New Jersey 08640 (Received April 22, 1068)
The exchange of Ca2+,Sr2+,and Ba2+ions for Na+ ions has been studied in Linde X and Y. Replacement of 16 Na+ ions per unit cell by Ca2+and Ba2+ions is slow in NaX. In the case of Ca2+ions, this necessitates equilibration times of 1 week or longer. However, Ba2+will not replace these cations at all at 25”. The replacement of 16 Na+ ions by Ca2+,Sr2+,and Ba2+ions is also slow in Nay. In the case of the latter zeolite, no exchange with these 16 cations takes place at 25”. The Na+ ions that are difficult to replace are in the small cavities of these two zeolites. At 50°1Ba2+ions do exchange with all the cations in all the cages of NaX. Measurement of the free energy of exchange of the Na+ by Ba2+ions in the large cages at low temperature and in all the cages at higher temperatures has permitted calculation of the individual equilibrium constants for the ion exchange of the networks of large and small cages. The selectivity data indicate no binding of Sr2+ and Ba2+ions in the large cages of Linde Y and partial or complete binding in the other systems studied. The standard enthalpies and entropies of reaction indicate that most of these reactions are entropy directed] and in the case of Ca2+and Sr2+ions the most important interactions are ion-water interactions in the solution phase.
The synthetic zeolites Linde X and Y have the biggest cavities and cavity entrances of any known zeolites. The largest cavities, sometimes called supercages, have a free diameter of 12 A. The entrances to these cavities are rings of 12 silicon and alupinum containing tetrahedra with a free diameter of 9 A. There are also smaller cavities called sodalite cages that have a free diameter of 7 A. The sodalite cages are connected to the supercages by rings of six tetrahedra having a free diameter of 2.5 A.l I n the sodium form of hydrated NaX, there are 16 Na+ ions per unit cell in the hexagonal prisms, one in each of the 16 hexagonal prisms in a unit cell.’ I n order to replace these cations by ion exchange, diffusion of ions into the sodalite cages and then into the hexagonal prisms must octur. Ions with Pauling radii much greater than 1.25 A should not be able to replace the Na+ ions that are in the hexagonal prisms. It is possible that ions slightly greater than 1.25 A in radius may do so because of the thermal vibrations of the crystal lattice and the “softness” of the diffusing ions and lattice oxygen atoms. The results of our investigation of univalent ion excchange in Linde X and Y 2 and our initial work on polyvalent ion exchange3led to the conclusion that on a unit-cell basis there are also 16 sodium ions in the network of small cavities of NaY (the network of small cavities consists of 8 sodalite cages 16 hexagonal prisms per unit cell). This conclusion was based on the fact that Cs+, Rb+, T1+12and La3+ could not replace 16 of the 51 sodium ions in a unit cell of the NaY used in this study. A more extended discussion of the structures of NaX and NaY can be found in ref 2.
+
The Journal of Physical Chemistry
Polyvalent ion exchange in these zeolites is interesting because they resemble more closely than other zeolites the ion-exchange resins in their porosity and water content. Moreover, the selectivity of NaY for univalent ions2 is the same as that exhibited by sulfonated polystyrene resins with reasonable degrees of divinylbenzene crosslinking at moderate ionic ~ t r e n g t h . ~The rigidity of zeolite “backbones” relative to resin backbones affords the opportunity to sort out the role that changes in water content, ion content, ion-exchange capacity, and polymer configuration play in determining the ion selectivity exhibited by ion exchangers. NaX and Nay, in particular] afford an opportunity to study polyvalent ion exchange as a function of ionexchange capacity at constant polymer configuration and very nearly constant water content because these two zeolites are isostructural. They differ only in their aluminum content and, therefore, in their ionexchange capacity. It was already known to us that most zeolites appear to become increasingly selective for alkaline earth ions over sodium ions with increasing temperature] indicating that the standard enthalpy change that occurs in transforming the sodium form of the synthetic zeolite to the alkaline earth form is p o s i t i ~ e . ~I,n~Dowex-50, (1) L. Broussard and D. P. Shoemaker, J . Amer. Chem. Soc., 82, 1041 (1960). (2) H. 8. Sherry, J. Phys. Chem., 7 0 , 1158 (1966). (3) H. S. Sherry, J . Colloid Interfac. Sci., in press. (4) F. Helfferioh, “Ion Exchange,” McGraw-Hill Book Co., Inc., New York, N. Y., 1962. (5) R. M.Barrer, L. V. C. Rees, and D. J. Ward, Proc. Roy. Soc., A273, 180 (1963).
4087
THEION-EXCH.ANQE PROPERTIES OF ZEOLITES the enthalpy of exchange is also positive, except when the univalent ion is H+.7 Thus when Na-form zeolites are converted to polyvalent ionic forms by ion exchange solution with aqueous solutions, the system zeolite increases in energy, This can be rationalized on simple electrostatic consideration^.^^^ I n the case of alkaline earth ion exchange of NaA, the increase in the entropy of the system that occurs when the exchanger is converted to the alkaline earth form6vs is large enough to more than compensate for the enthalpy increase. We will see later that this statement is also true in most of the systems we have investigated herein.
+
Experimental Section The reagents; used in this research were all Baker Analyzed reagents. The radioactive isotopes used were 46Ca and s6Sr purchased from the New England Nuclear Corp. and 133Ba purchased from the Oak Ridge National Laboratory. The zeolites used were NaX (lot 262) purchased from Union Carbide Corp. and N a y synthesized in our laboratories. They were preconditioned, by slurrying 100 g three times for 1 hr each time in 0.1 N NaCl (10% weight slurry), filtered, and washed. The washing was done on the filter using four 200-ml portions of deionized water. This procedure eliminates H+ ion exchange of the zeolite, which can take place when zeolites are extensively washed. The contents of an anhydrous unit cell of the NaX and NaY used is given by the formulas [(AlO~)~(Si0~)1071 and Naal [(AlOJsl(SiOJ1411. Phase equilibration was accomplished by weighing out a suitable quantity of the sodium form of the eeolite and pipetting a suitable volume of the salt solution into :2-oz polyethylene bottles. The total salt concentration was always 0.100 N , unless otherwise specified, and chloride salts were used. In some cases, the solution clontained only the ingoing cation; in other cases the solution contained both of the exchanging cations. The equilibrations were carried out in a shaking water bath in which the temperature was controlled to within *0.5". When conventional wet chemical analyses were used to obtain the phasedistribution data at 25", rapid filtration was used to separate the phases. When the ingoing ion was traced with a radioactive isotope by tagging the solution used for the ionexchange reaction, the shaker was stopped and the zeolite was allowed to settle. A 5-ml sample of the supernatant liquid was then pipetted into a test tube and allowed to stand a t room temperature for 1 dav. This was done to permit any zeolite particles in the 5-ml The radioactive content of 2-ml sample to aliquots from the 5-ml samples was assayed along with ml of the tagged stock solution- The ratio of the two counts, along with a knowledge of the initial weight and Of the and the and analysis of the solution phase, allowed US to Calculate
the composition of the two phases a t equilibrium. Checks were made by nonradioactive equilibrations and complete chemical analysis of both phases to ensure that the expected ion-exchange reaction was taking place, i.e., that no precipitation of polyvalent ion or hydrolysis of the zeolite (H+ ion exchange) took place. Powder diffraction patterns of nonradioactive samples were also obtained using a Siemens X-ray diffractometer to ensure that the crystals retained their structure. The p particle emitting "Ca isotope was counted using a liquid scintillation system and a Packard TriCarb Model 3375 scintillation spectrometer. The y emitting isotopes la3Ba, 86Sr, and 22Nawere counted using a Packard Model 5020 automatic sample changer, detector, and high-voltage supply system operated in tandem with the Model 3375 Tri-Carb scintillation spectrometer. Results Ion-Exchange Isotherms. The ion-exchange isotherms are presented in Figures 1-11. In these figures, the abscissa, S, is the equivalent fraction of the ingoing ion that is in the solution and the ordinate, 2, is the equivalent fraction of the ingoing ion in the zeolite Zhf
=
equivalents of M"+ in zeolite gram-atoms of A1 in zeolite
= equivalent fraction of Mm+ in zeolite S M
=
normality of M"+ in solution total normality of cations in solution
= equivalent fraction of
M"+ in solution
Using these coordinates, the termination point of an isotherm clearly gives the ion-exchange capacity of the zeolite for the pair of ions being considered. The ion-exchange isotherms (Figure 1-11) show that 100% exchange (termination point of the isotherm at the point S = 1, 2 = 1) was achieved in some cases and not in others. Complete replacement of all the Na+ ions in NaX was achieved with Ca2+ ions and Sr2+ ions at 25 and 50" (Figures 1, 2, 5, and 6) and with Ba2+ions only at 50" (Figure 10). At 5 and 25" we found that even when a 1000-fold excess of Ba2+ ions is used, only 82% of the Na+ ions are replaced over a period of 4 weeks. In zeolite Y, at temperatures up to 50", only 68% of the Na+ ions can be replaced in a reasonable time. (6) €3. 9.Sherry and H. F. Walton, J. Phw. Chem., 71, 1457 (1967). (7) I. Gamalinda, L. A. Schoemer, H. S. Sherry, and H. F. Walton, ibid., 71, 1622 (1967). (8) A. H. Truesdell and C. L. Christ, "Glass Electrodes for Hydrogen and Other Cations," G. Eisenman, Ed., Marcel Dekker, Inc., New York, N. Y., 1967. (9) H. S. Sherry in ,'Ion Exchange,,? vel. 11, J. A, Marinsky, Ed., Marcel Dekker, Ino., New York, N. Y., in press.
Volume 72, Number 12 November 1968
HOWARD 8. SHERRY
4088
0
I 0.2
I 0.6
I
0.4
I
0.8
I
1.0
-
0
0
0.2
0.4
0.6
0.8
1.0
SCa
SCa
Figure 1. The ion-exchange isotherm for the Ca-Na-X system a t 25" and 0.100 N total normality.
Figure 4. The ion-exchange isotherms for the Ca-Na-Y system at 50": 0,0.103 N total normality; E!, 0.051 N total normality.
t
0.2
-
SCa
1
0
0.2
0.4
0.6
0.8
I
Figure 5. The ion-exchange isotherm for the Sr-Na-X system at 25" and 0.1 N total normality.
I
0.8
0
0
SS*
Figure 2. The ion-exchange isotherms for the Ca-Na-X system a t 50": 0, 0.103 total normality; m, 0.051 total normality.
-
0
0.2
0.4
0.6
0.8
1.0
SCa
0 0
0.2
0.4
0.6
0.8
1.0
S Sr
Figure 3. The ion-exchange isotherm for the Ca-Na-Y system at 25" and 0.1 N total normality.
Figure 6. The ion-exchange isotherm for the Sr-Na-X system at 50" and 0.100 N total normality.
The isotherms for the Sr-Na-X system a t 25 and 50" (Figures 5 and 6) have a very unusual shape in the region of 5 5 4 0 % Sr2+ loading. This result cannot be experimental error because these points are reproducibly "bad."
The 25" isotherm for the Ca-Na-X system was also obtained by using a 24-hr equilibration time. It is similar to the Ba-Na-X isotherm at 25" in that it indicates that only 82% of the Na+ ions are replaced by these divalent ions. After a few hours, the termi-
The Journal of Physical Chemistry
THEION-EXCH.ANGE PROPERTIES OF ZEOLITES
4089
B
N
I
0,4!
Ob
012
014
SSr
Oi6
Ole
1.b
Ba
Figure 7. The ioin-exchange isotherm for the Sr-Na-Y system at 25" and 0.100 N total normality.
Figure 10. The ion-exchange isotherm for the Ba-Na-X system a t 50" and 0.100 N total normality.
8
0.2,'
Y
O
0
0.2
0.4
0.6
0.8
1.0
Figure 8. The ion-exchange isotherm for the Sr-Na-Y system at 50" and 0.100 AT total normality.
1
-/
m
N
I
0 . 9
-0OL
0.2
0.4
1
0.2
I
0.4
I
0.6
I
0.8
I
1,O
SBa
SSr
0.8
01 0
0.6
0.8
1.0
SBa
Figure 9. The ion-exchange isotherm for the Ba-Na-X system at 25" andl 0.100 N total normality.
nation point is independent of time when Ba2+ is the ingoing ion, whereas that of the Ca-Na-X isotherm is time dependent. The 24-hr isotherm is shown under the equilibrium isotherm as a solid line in Figure 1. The effect of varying the total normality at 50" in the Ca-Na-X and Ca-Na-Y systems is shown in Fig-
Figure 11. The ion-exchange isotherm for the Ba-Na-Y system a t 25" and 0.100 N total normality.
ures 2 and 4. Ca2+ ion specificity increases with decreasing total normality-the expected electroselective effect (for example, see ref 4). These isotherms for the Ca-Na-X system at 50" (Figure 2) are extrapolated to the point S = 1, 2 = 1, because complete Ca exchange was achieved in NaX at 25". The phase-distribution data we obtained have been used to calculate the standard free energies, enthalpies, and entropies of exchange by plotting the logarithm of the corrected selectivity coefficient vs. the normalized equivalent fraction of the ingoing ion in the zeolite, Z'M, and using a simplified version of the equation of Gaines and Thomaslo 1
I n this working equation, which neglects salt imbibement and changes in water activity, the corrected selectivity coefficient, K 'M , N ~ is , defined as
(10) G. L. Gaines, Jr., and H.C . Thomas, J. Chem. Phys., 21, 714 (1953). Vo'olume 72,Number 12
November 1968
HOWARD S. SHERRY
4090 Table I : Standard Free Energies, Enthalpies, and Entropies of Reaction
Reaction
+ + + + + Ba2f + 2NaY
Ca2+ 2NaX Ca2+ 2NaY Sr2+ 2NaX Sr2+ 2NaY Ba2+ 2NaX
278OK
-370 f 15
- 1250" f 30
--
A P T , cal/equiv 298OK
-320 -420 -740 -510 -1310"
323OK
f 10 f 10 f 20 i 20 f 30
-460 f 20 -470 i 20 -850 i 30 -680 f 20 -940b i.30 1080' i 30 -840 f 20
-870 i 20
AS0,
AHo, cal/equiv
eu/equiv
1200 f 100 640 f 50 5 3 0 i 50 1540 AZ 150 -430" 80
5 . 1 f0 . 5 3.6 f 0.4 4 . 3 f 0.4 6.9 f 0.7 3 . 0 i 0.6
-350
1 . 8 f 0.2
30
" Refers to complete exchange of the large cages and no replacement of Naf ions in the small cages. Refers to complete exchange The free energy for the replacement of only the N a + ions in the small cages. of all the N a + ions in all the cavities.
where m M and WLN* are the molalities of ions Ma+ and Naf, respectively, in the aqueous solution, y~ and yxa are the single-ion activity coefficients in the solution, and Z'M is the normalized equivalent fraction of the ion h12+in the zeolite phase. The equivalent fracin the zeolite phase is normalized by tion of the ion M2+ dividing the values of ZM plotted in the ion-exchange isotherms (Figures 1-1 1) by the maximum fractional loading that can be achieved-thus Z'Y now varies from 0 to 1. It is worth noting that the corrected selectivity coefficient is independent of total normality. Thus the values for the corrected selectivity coefficient calculated using data obtained at 0.051 and 0.103 N total normality and 50" in the Ca-Ea-X system (Figure 2) should result in one smooth curve. The corrected selectivity coefficients plotted in Figure 12 at 50" do indeed fall on one curve. A similar result was obtained using the data at 0.051 and 0.103 total normality in the Ca-Xa-Y system at 50". This is shown in Figure 13. We have tabulated values of the standard free energies, enthalpies, and entropies of reaction calculated from the equilibrium constants and their temperature dependence in Table I. I n this table, AGO* is the standard free energy of exchange in calories per equivalent of zeolite, AHo is the standard enthalpy of exchange per equivalent of zeolite, AS" is the standard entropy of exchange per equivalent of zeolite, and T is the absolute temperature in degrees Kelvin. When the standard free energy was obtained at three temperatures, two values of AH" were calculated using adjacent temperatures and the arithmetic mean reported. The free energy, enthalpy, and entropy changes reported are the difference between that for 1 equiv of the alkaline earth form of the zeolite (exchanged to the maximum level at that temperature), in equilibrium with pure water and 1 equiv of the sodium form of the zeolite in equilibrium with pure water, where 1 equiv of zeolite is defined as the weight, in grams, of zeolite containing 1 equiv of exchangeable cations. The Journal of Physical Chemistry
- 0.8
-'
0
0.2
0.4
0.6
0.8
I
&a
Figure 12. Corrected seleotivitv coefficient for the Ca-Na-X system as a function of zeolite composition: 0, 25" and 0.103 N total normality; El, 50" and 0.051 N total normality; A, 50" and 0.103 N total normality.
8
.
0
0
0.2
0.4
0.6
0.8
1.0
Z'Ca
Figure 13. Corrected selectivity coefficients for the Ca-Na-Y system as a function of zeolite composition: 0, 25" and 0.103 total normality; U, 50" and 0.051 i V total normality; A, 50' and 0.103 N total normality.
THEION-EXCHANGE PROPERTIES OF ZEOLITES
Discussion The isotherms in Figures 1-11 indicate that 18 f 2% of the Na+ ions in NaX and 32 f 2% of those in NaY are difficult to replace with divalent cations. A similar result was obtained using La*+ ions and has already been r e p ~ r t e d . ~ In the case of NaX, 18% of the Na+ ions represents very close to 16 Na+ ions per unit cell (there are 85 Na+ ions in a unit cell) and in the case of Nay, 32% also represents very close to 16 Na+ ions per unit cell (there are 51 Naf ions in a unit cell). Thus we must conclude that the 16 Na+ ions per unit cell that are in the hexagonal prisms of hydrated NaX' are difficult to replace with divalent ions. I n hydrated N a y , we must also conclude that there are 16 ??a+ ions per unit cell in the small cavities and that they are difficult to replace, because we cannot imagine divalent ions experiencing difficulty in replacing Na+ ions in the large cavities of synthetic faujasites. Moreover, we have studied the rate of exchange of pollyvalent ions for Na+ ions and the replacement of 82 and 68% of the Na+ ions in NaX and NaY is very ra,pid (minutes); whereas many days of exchange at 100" are required to replace the last 16 Na+ ions in Nay, 4 days are required with Ca2+ions a t 25" ip NaX and 1 week is required with Ba2+ions at 50" in NaX. We must conclude that the slow rate of Ca2+ ion exchange of the network of small cavities in the synthetic faujasites, NaX and Nay, is primarily attributable to the large hydration energy of the Ca2+ion. This ion has about t,he same Pauling radius as Na+ ion" (0.99 vs. 0.95 A). Sodium isotope exchange experiments in NaX and NaY indicate that 22Na+ions exchange very rapidly with all the Na+ ions in the zeolitesg Thus, based on ionic radius, Ca2+should diffuse quite freely into the small cavities of these zeolites. It does not do so because the hydrated ion is too large to diffuse through the 2.5-A entrances to the small cages, and energy must be supplied to at least partially strip the primary hydration shell of the Ca2+ions. The hydration energy is 140.2 kcal/g-ion,12 and even if only part of this energy must be supplied, the potential energy barrier to Ca2+ ion diffusion from the large cages into the small cages would be large. The hydration energy of the Ba2+ion is smaller than that of the Ca2+ ion,12 and on this basis a more rapid rate of exchange of Ba2+ ions is anticipated. However, a t 25" the rate of replacement of the 16 Na+ ions in the hexagonal prisms of NaX by Ba2+ ions is immeasurably slow. The Pauling radius of the Ba2+ion is 1.35 A, and this ion must pass through a 2.5-A window in order to replace the Na+ ions that are in the small hexagonal prisms-on the basis of bare ion size, slow diffusion is expected. Bare ion size cannot be the sole cause of slow Ba2+ ion diffusion into the small cages of KaX and N a y because K + ion, which has a Pauling radius of 1.33 A,
4091
rapidly replaces all the Na+ ions in NaX and Undoubtedly, the dehydration energy of Ba2+ ions in the large cavities also contributes to the potential energy barrier to diffusion from the large cavities into the small ones. I n the alkaline earth series, hydration energies decrease with increasing atomic number and, of course, ionic radii increase. An optimum combination of hydration energy (or size of the hydrated ion) and bare ion size occurs at Sr, and complete replacement of all the Na+ in NaX by Sr2+ions can be accomplished in 1 hr at 25°.9 When the equilibrium conditions can be established with respect to both the large and the small cages of NaX, we find that the ion-exchange isotherms for the replacement of all the Na+ ions by alkaline earth ions are sigmoidal (Figures 1 and 10). Their shape could be interpreted to mean that in the small cages the equilibrium constant is less than 1 and in the large cages it is greater than 1. I n the case of Ba2+ ions, these two equilibrium constants were determined. Let us define K , as the equilibrium constant for the replacement of the Na+ ions in the network of large cavities of synthetic faujasite and K , as the equilibrium constant for the ion-exchange reaction in the networklof small cages. At 5 and 25", we have measured K , because only the Naf ions in the large cages can be replaced at these temperatures (see Figure 9 for the 25" data), and at 50" we have measured K,+, ( = K a 0 J 2KpO.lS) because the Na+ ions in both networks of cavities can be replaced (see Figure 10). The equilibrium constants K , and KO can be calculated with reasonable accuracy from free energy data given in Table I by extrapolating to 50" the free energy of exchange of the large cages measured at 5 and 25". The value obtained is -1370 cal/g-equiv of Ba2+ ions. The equilibrium constant corresponding to this free energy change (K,) is 72.5, and the equilibrium constant for the replacement of all the Na+ ions in NaX with Ba2+ ions at 50" is 18.8 (K,+,). The equilibrium constant for the replacement of Na+ ions in the network of small cavities (K,) is 0.037, and, therefore, we conclude that the sigmoidal shape of the Ba-Na-X isotherm at 50" is due to the heterogeneous character of the ion-exchange sites (heterogeneous in the sense that cations are involved in different interactions with lattice oxygen atoms and water molecules in the different cages). We also recognize that even in the large cages the cations can occupy crystallographically different sites, but it is not possible to "sort out" the equilibrium constants for the exchange of these cations. The sigmoidal character of the Ca-Na-X isotherms at 25 and 50' is most probably also due to an equilibrium constant for the exchange of the large cages that is (11) L. Pauling, "Nature of the Chemical Bond," Cornel1 University Press, Ithaca, N. Y., 1960. (12) D. R. Rosseinsky, Chem. Rev.,65, 467 (1965).
Volume 72, Number 12 iVovember 1968
HOWARD S. SHERRY
4092 greater than unity and one for the small cages that is less than unity. We attempted to obtain an isotherm at 5" with the hope that ion exchange in the network of small cavities would be completely quenched. A much slower, but unfortunately finite, rate of exchange of these cavities was observed, and, therefore, we cannot study ion exchange in the large cavities alone. It is noteworthy that in the case of Ca2+and Sr2+ion exchange of NaX and Nay, the standard free energy of reaction is negative despite positive values of the standard enthalpies of reaction (Table I). These reactions are entropy directed-the standard entropy of reaction is positive. Moreover, we see from the data in Table I1 that most of the increase in the entropy of these two-phase systems occurs in the solution phase. These reactions are controlled by water-ion interactions in the aqueous solution phase.
Table 11: Entropy and Enthalpy Changes in Aqueous and Zeolite Phases i/,,y31it.hyd
M2+
CaX
CaY SrX SrY BaX BaY a
-
i/l&l\Iweol
-
i/,,yMz+hyd
-
-
i/z~lII~+zeol
SN&hYd,Q
SN&'~O',
"ahyd,a
"aeeo1,
eu/equiv of zeolite
eu/equiv of zeolite
cel/equiv of zeolite
cal/equiv of zeolite
-4.2 -4.2 -3.4 -3.4 1.95 1.95
1.0 -0.6 0.9 3.6 4.65 3.65
-93,400 -93,400 -75,700 -75,700 -58,700 -58,700
-92,200 -92,800 -75,000 -75,200 -59,130 -59,050
Values calculated from data in ref 11.
In the case of Ba2+ ion exchange, the standard enthalpy of reaction is small but negative, instead of being positive, as was found for Ca2+and Sr2+,and this makes Ba2+ ion the most preferred alkaline earth ion. The negative standard enthalpies of reaction (Table I) more than compensate for the lower positive standard entropy of reaction. It is interesting to note that the standard entropy of reaction is positive (Table I) even though the solution phase decreases in entropy. A net transfer of water from the zeolite phase to the solution takes place during Ba2+ ion exchange. For NaX, 1 equiv contains 56.71 g of H20and 1 equiv of Ba82NalsX contains 50.45 g of HzO or 2.68 and 2.38 mol of HzO, respectively. Thus 1 equiv of NaX contains 0.30 mol more water than 1 equiv of Ba&alsX. Using values of 14 eu for the partial molal entropy of H20in NaX at 70°, taken from the data of Barrer and Bratt,13 and 17 eu for the molal entropy of liquid water at 2 5 O , I 4 it is estimated that the desorption of 0.30 mol of HzO/equiv of zeolite causes an increase in the entropy of the system of about 1 eu. This change accounts for at least part of the entropy increase that is attributed to the zeolite phase. The Journal of Physical Chemistry
The Sr-Na-X isotherms (Figures 5 and 6) and selectivity plot (Figure 14) have very unusual shapes. An X-ray diffraction study of this system16is being reported that shows that the end members of this system, NaX and SrX, exhibit limited solubility. There is a region of mutual insolubility extending from 71 to 87% Sr loading of NaX that completely accounts for the unusual selectivity behavior. In an earlier study of univalent ion exchange in zeolites X and Y,2 we concluded that the ions in the large cages of zeolite Y are not sited in hexagonal windows but, rather, they are mobile and "dissolved" in the water in these cages. The independence of the corrected selectivity coefficient on the degree of Sr2+ and Ba2+ion loading of NaY observed in the present study indicates that these ions also are not sited in the large cages-they too are "dissolved" in the water in the large cages. (See Figures 15-17.) When the corrected selectivity coefficient decreases with increasing alkaline earth loading, heterogeneous siting of ions is indicated. For example, in the Caand Ba-Na-X systems, we know that ion exchange takes place in both the network of large cavities and the network of small ones. Indeed, we have obtained the individual equilibrium constants for these two reactions in the second case at 50". The lack of ion siting in NaY is undoubtedly due to its low aluminum and high water content. I n the more aluminous NaX, stronger interactions with lattice oxygen atoms are possible and have been observed by X-ray difffacti0n.l In this zeolite, there are not sufficient water molecules to satisfy the coordination numbers of the counterions (there are about 3.3 water molecules/Na+ ion in the large cages), and some of these cations must interact with lattice oxygen atoms in the large cages. I n a unit cell, there are only 32 of the rings of 6 tetrahedra that separate the large and small cages (containing 2 or 3 AlOz- groups/ring), and, therefore, a maximum of 32 of the 69 Na+ ions in a unit cell can be, and are, sited in these positions.' The remaining 37 cations are presumed t o interact only with water molecules in the large cages. I n a unit cell of Ca-, Sr-, or BaX, there are, at most, 35 cations in the large cages, and there are about 6 water molecules/ cation. It is possible to have all the divalent ions sited in the hexagonal windows in the large cages, to have all of them "dissolved" in the zeolitic water in the large cages, or to have both sited and fully hydrated ions in the large cages. I n the last case, a decrease in the selectivity coefficient should be observed with increasing alkaline earth ion loading, even when ion exchange takes place only in the network of large cages, because (13) R. M.Barrer and R. C. Brett, J. Phys. Chem. Solids, 12, 146 (1959). (14) W. M. Latimer, "Oxidation Potentials," Prentice-Hall, Inc., Englewood Cliffs, N. J., 1952. (15) D.H. Olson and H. S. Sherry, J . Phys. Chem., 72,4095 (1968).
THEION-EXCHANGE PROPERTIES OF ZEOLITES
4093 3.2,
Figure 14. Corrected selectivity coefficients for the Sr-Na-X system a t 0.100 N total normality as a function of zeolite composition: D, 25"; 0, 50".
1.6
1.2
I -
4
B
- 1
0.8
v)
k
0
Figure 16. Corrected selectivity coefficients for the Ba-Na-X system at 0.100 N total normality as a function of zeolite composition: A, 5'; 0, 25'; D, 50". The data a t 50" are referred to complete replacement of the Naf ions from the large and small cages.
-
n
c-
I
1.6
-
0 0
-0.4--0.8--
t
I
I
I
I
0
0.2
0.4
0.6
0.8
J
1.0
zir
t
0.4
Figure 15. Corrected selectivity coefficients for the Sr-Na-Y system tit 0,100 N total normality as a function of zeolite composition: A, 5"; 0, 25"; D, 50".
the ingoing ions do not occupy equivalent positions. This behavior ia observed in the Ba-Na-X systems a t 5 and 25" (Figure 16). I n alkaline earth forms of zeolite Y there are, a t most, about 17 doub1,y charged ions and 224 water molecules in the 8 large cages in a unit cell-2 ions and 28 water molecules per large cage. Again, all, some, or none of these cations may be sited in the rings of 6 tetrahedra that separate the large cages from the small ones. The independence of the corrected selectivity coefficient for Sr2+and Ba2+ion exchange of NaY on zeolite composition indicates that these ions are not sited. Thus hydrated Na+ ions2 are replaced by hydrated alkaline earth cations. This behavior in the large cages of zeolite Y is easy to rationalize because the water content is high and the density of fixed charges is low; weak
cation-cation and cation-fixed anion interactions are expected. The only data to which we can compare ours are those of Barrer, Rees, and Shamsuzzoha (BRS),16Ames," and Rees and Williams (RW).ls Our selectivity curves for Ca2+and Ba2+ion exchange in NaX smoothly decrease with increasing loading of alkaline earth cations, where(16) R. M. Barrer, L. V. C. Rees, and M. Shamsuzzoha, J . Inorg. Nucl. Chem., 28, 629 (1966). (17) (a) L. L. Ames, Jr., Amer. Miner., 49, 1099 (1964); (b) L. L. Amea, Jr., J. Inorg. Nucl. Chem., 27, 885 (1965). (18) L. V. C. Rees and C. J. Williams, Trans. Faraday Soc., 61, 1481 (1965).
Volume 7.8, Number 1.8 November 1968
HOWARD S. SHERRY
4094 as BRS obtained a maximum in the selectivity curve for the Ca-Na-X system at 25", and both BRS and RW showed that the selectivity coefficient continually increased with loading of barium in the Ba-Na-X system at 25". Furthermore, the termination point of their isotherm for the Ba-Na-X system indicates that the maximum level of Ba2+ ion loading that can be achieved is 77%. (See ref 16 and 18.) We find a value of 82 f 2%. Our result agrees quite well with that expected, considering that 16 of the 85 cations in a unit cell of NaX are in the hexagonal prisms (16/85 = 0,188). We have obtained results similar to BRS and RW using a batch of zeolite that was partially hydrolyzed and was 7% in the Hs0+ (or H+) form and 93% in the Na+ form. The free energy of Ca2+, Sr2+, and Ba2+ ion exchange at 25" reported by BRS is -160, -710, and -1130 cal/equiv, whereas we find (Table I) -320, -740, and -1310 cal/equiv. It is only expected that the free energies of exchange do not agree because the shapes of the selectivity plots differ, Furthermore, BRS could not measure the heat evolved during the slow step in, for example, Ca2+ion exchange by their calorimetric method, because this step is so slow, The calorimetric method is a much more accurate way to measure AH" than is the use of the temperature dependence of AGO. However, in this case, 17-18% of the reaction cannot be studied calorimetrically, and it is not reasonable to compare our heats of complete exchange with the heats of partial exchange obtained by BRS. Ames" has studied Ca2+ and Sr2+ ion exchange of NaX, but we cannot compare our isotherms with those in ref 17a, because he did not publish either the isotherms or the raw data-only the derived thermodynamic data. He found a standard free energy of exchange of -500 and -900 cal/equiv that compares with our values of -320 and -740 cal/equiv. The Sr-Na-X isotherm at 25" and 1.0 N is presented in ref 17b. The isotherm is very similar in shape to the one
The Journal of Physical Chemistry
we obtained at 25" and 0.100 N (Figure 5), but the curve is incorrectly smoothed. We repeatedly have obtained the unusually shaped Sr-Na-X isotherm. It is difficult to account for the difference in our value for the free energy of exchange. However, again, we must state that without seeing the selectivity plots, it is not possible to compare free energy changes. I n the case of the Sr-Na-X isotherm presented in ref 17b, the smoothed version must certainly lead t o a different selectivity plot than that presented in Figure 14 and, therefore, t o a different value of the free energy of exchange. Summary and Conclusions This work demonstrates that small divalent ions can be sieved from the small cages of zeolites. This sieving action is probably due to the fact that these ions are highly hydrated. At least some water molecules must be stripped from the first hydration shell before hydrated cations can diffuse into the network of small cages. The enthalpy of exchange of divalent metal ions for Na+ is either a positive or a small negative number. It is the positive entropy of reaction that is primarily responsible for the divalent ion specificity exhibited by NaX and N a y . I n the cases of Ca2+and Sr2+ions, it is chiefly the entropy change that occurs in solution that is responsible for the observed ion specificity. The dependence of the corrected selectivity coefficients on zeolite composition indicates that there is partial siting of Ca2+ and Ba2+ cations in the large cages of zeolite X and no siting of Sr and Ba2+ions in zeolite Y.
Acknowledgment. We wish to thank the Mobil Research and Development Corporation for its support and encouragement of this work. I n particular, I thank Mr. Charles Adams who has obtained most of the equilibrium data, and Dr. David H. Olson for many stimulating and enlightening discussions.
