ION-EXCHANGE PROPERTIES OF ZEOLITES
1457
The Ion-Exchange Properties of Zeolites. 11. Ion Exchange in the Synthetic Zeolite Linde 4-A
by Howard S. Sherry Central Research Division Laboratory, Mobil Oil Corporation, Princeton, New Jersey
and Harold F. Walton Department of Chemistry, University of Colorado, Boulder, Colorado
(Received September 89, 1966)
An investigation of ion exchange in Linde A has been completed. The preliminary work involved characterizing the lot of Linde 4-A used and establishing the equilibrium conditions. Ion-exchange isotherms for the T1+ and Ag+ exchange for Na+ demonstrate that an average of 0.36 NaAlOz is occluded in each sodalite cage in the crystal. It is shown that a perfectly crystalline hydrated Ba-A can be prepared from Na-A. The maximum in plots of selectivity coefficient vs. per cent loading of alkaline earths, reported for the Ca-Ea-A system by other investigators, has been found to be a result of not reaching equilibrium a t low loadings. When equilibrium was reached a t all loadings, the selectivity coefficient continuously decreased with increasing per cent loading of Ca2+, Sr2+,and Ba2+. Free energies have been measured a t 278, 298, and 350°K. Enthalpies and entropies of exchange are derived from these data and are presented. The thermodynamic data lead to the conclusion that there is much ion-pair formation between cations and the fixed anionic sites in Ag-A and Ca-A, less in Sr-A, and almost none in Ba-A.
In the first paper of this series’ it was pointed out that, because molecular sieve-type zeolites are crystalline, study of the ion-exchange process in these materials affords a unique opportunity to understand interactions between the countercations and the anionic framework in terms of the local environment in which the cations exist. This local environment is known if the crystal structure has been completely determined. I n the case of Linde 4-A (Na-A) the structure of the anionic framework and the positions of some of the counterions are known.2 The unit cell is a cube 12.3 A on edge. The framework consists of sodalite cages connected by linking adjacent rings of four SiOz and AIOz groups to form square prisms between sodalite cages. Because the so-called four-membered rings of the sodalite cage are at the corners of an octahedron whose origin coincides with the center of the sodalite cage, the sodalite cages form an octahedral lattice (whose point symmetry is Oh) in a simple cubic symmetry. There is, therefore, one sodalite cage per unit
cell. The unit cell contains 12 Na, 12 Si, 12 Al, and 48 0 atoms because 24 SiOz or AIOz groups are required to construct a sodalite cage and the Si/Al atom ratio in Linde 4-A is unity (a complication that can make this ratio less than unity will be discussed later). The interstices within the cubic array of sodalite cages form a simple cubic array of much larger cages that interconnect by sharing rings of 8 tetrahedra with free diameters of 4.0 A. These large cages are connected to the sodalite cages by rings of 6 tetrahedra with free diameters of about 2.5 A. Eight of the 12 sodium ions have been located in the six-membered rings separating the large and small cavities-one in each of the 8 rings in a unit cell. The other 4 sodium ions have not been located and are presumably “dissolved” in the zeolitic water. (I) H.S.Sherry, J . Phys. C h m . , 70, 1158 (1966). (2) L. Broussard and D. P. Shoemaker, J . A m . Chem. SOC.,8 2 , 1041 (1960).
Volume 7 1 , Number 6 April 1967
HOWARD S. SHERRY AND HAROLD F. WALTON
1458
All of the cations have been located in Linde 5-A2 (roughly 67y0 are calcium and 33% are sodium). The positions of the ions used in this work-Ag+, Tl+, Ca2+, Sr2+, and Ba2+ ions-within the crystals of Linde A have never been determined. However, an ion-exchange study and a knowledge of the configuration of the anionic framework as well as the positions of Na+ ions in Na-A and Ca2+ ions and Na+ ions in a mixed Ca-Na-A (Linde 5-A) should make it possible to make very reasonable deductions concerning cation positions and cation interactions within the anionic framework. Having once correlated the ion-exchange properties of crystals with structural factors, it may be possible to make statements about the local environment of a counterion in an amorphous ion exchanger or in a crystalline exchanger of unknown structure. Barrer and Falconera have proposed a theory of ion selectivity to predict qualitatively the energy of exchange. In their model the water and zeolite phases are assumed to be continuous dielectric media, the zeolite phase having a lower dielectric constant than the solution phase. The energy of a cation is assumed to be the energy of a hard, charged sphere of radius r+ in a dielectric medium
where E is the internal energy of the cation, Z the charge of the cation, and e the dielectric constant of the medium. If alkaline earth ions are replacing Na+, the energy of the reaction is
AE =
1/2EAz+Zeo1
+
ENaaq
- EN?'
- l/2EA*+"
(2)
This model does not explicitly describe the factors determining the energy of reaction in terms of compositional or structural parameters, but, as we will show later, it does successfully predict a positive energy of complete exchange.
Experimental Section All reagents were reagent grade. The zeolite used was Linde 4-A, Lot No. 47077. The average of three analyses is given in Table I. The radioisotopes used were la3Ba, 89Srl 46Ca, and IlOmAg. The y-rayemitting isotopes 13aBa,85Sr,and l1OrnAgwere counted with a Model 510 Baird Atomic single-channel pulseheight analyzer, a Model 134 high-speed scaler, a Model 215 amplifier, a Model 312A high-voltage power supply, arid a Model 810 scintillation detector. The isotopes that emitted only p particles, 44Ca and 89Sr, were counted using a Model 3003 Packard Tricarb scintillation spectrometer. Equilibrations were performed essentially as preThe Journal of P h g s h l Chemistry
viously described,' with the exception that a Mode 76 shaking water bath was used for experiments done at 25 and 77". The temperature in this bath was controlled within 10.5". The equilibrations performed at 5" were done in a walk-in refrigerator whose temperature was controlled within 1".
*
Results and Discussion The first objective of this work was to characterize the lot of Linde 4-A to be used. This was necessary because Barrer and Meier demonstrated that Ag+ could replace all of the 13 Na+ ions in the unit cell of their lot of Linde 4-A, whereas TI+ could only replace 12 cation^.^*^ Their Na-A contained 13 NaA102 and 12 SiOzin the unit cell and they concluded that the only location for the thirteenth AlO2, which was consistent with the ion-exchange data, was in the sodalite cages. Barrer and Meier found one occluded NaAIOz group per unit cell. However, it is possible to have 0-1 AIOz group per unit cell, making it necessary to determine the exact amount. This was done by determining Ag-Na and T1-Na isotherms at 25". Figures 1 and 2 present the ion-exchange isotherms for the exchange of Ag+ and T1+ for Na+ in our lot of Type A. Figure 1 demonstrates the marked Ag+ selectivity Na-A exhibits and, more important, the fact that all the Na+ ions are replaced by Ag+ ions. Figure 2 demonstrates that only 97% of the Na+ ions are replaced by T1+ ions. This observation is consistent with the analysis of the zeolite given in Table I. One can therefore conclude that, in this lot of Na-A, there is an average of one-third of an NaA102molecule occluded in each sodalite unit. The thallium and Ag+ ion exchange for Na+ ions is very helpful in determining how many Na+ ions are in the small cavities of a zeolite. In many of the zeolites whose structures are known, the entrances of the small cavities are rings of 6 tetrahedra with free diameters ranging from 2.2 to 2.5 A. The Pauling radius of Ag+ and Tl+ is 1.26 and 1.40 A, respectively, so that the silver ion can pass quite readily through these openings and T1+ cannot. One important point to be verified was whether or not the zeolite retained its structural integrity during exchange. This work was done by measuring sorptive properties and by obtaining an X-ray powder diffraction pattern. A numerical value of the crystallinity was assigned by comparing the intensity of the 34.1" (3) R. M. Barrer and J. D. Falconer, Proc. Roy. SOC.(London), A236, 227 (1956). (4) R. M. Barrer and W. M.Meier, Trans. Fuvuduy SOC.,54, 1074 (1958). (5) R. M. Barrer and W. M. Meier, &id., 55, 130 (1959).
ION-EXCHANGE PROPERTIES OF ZEOLITES
1459
Table I : Composition of Some Zeolite Samples % wt" Samples
Na
Ba
SiOa
AliOa
Na
Linde 4-A, Lot No. 47077 Ba-A Na-A from Ba-A
16.0 2.14 15.5
... 31.8 0.78
41.3 32.6 41.8
36.1 28.4 35.8
12.2 2.05 11.6
" Composition of the anhydrous unit cell. atoms per unit cell.
I,
Atoms per unit cell" Ba Si
... 5.12 0.101
12.0 12.0 12.0
7
Alb
12.4 12.3 12.1
Composition of the unit cell is based on the assumption that there are 12 silicon
I .o
SAP
Figure 1. The ion-exchange isotherm for the Ag-Na-A system at 0.1 total normality and 25": 0 , Ag.+ Nas+.
+
are shown in Table I. The sorptive properties and relative crystallinities of the Na-A starting material and the Ba-A prepared from it are tabulated in Tables I1 and 111, respectively. The samples were activated for sorption by calcining in a stream of dry air at 350" and then brought into contact with water or nhexane a t 20 mm partial pressure and room temperature. The sample of Na-A sorbed no n-hexane, which is in agreement with other investigators.6 The sample of Ba-A was expected to sorb n-hexane, as reported by others for Ca-A;6 however it did not. Furthermore, its capacity to sorb water is essentially zero. X-Ray powder diffraction patterns and relative crystallinity data obtained using these two samples indicate that the Linde 4-A was quite crystalline whereas the calcined and rehydrated Ba-A was amorphous. The relative crystallinity data are shown in Table 111.
I .o
Table II: Sorptive Properties of Zeolite Samples
0.8
Grams of HtO sorbedo/ e of
6
0.6
Sample
0.4
0.2
0
I
I
0.2
0.4
I 0.6
I
0.8
I 1.0
activated zeolite
Grams of nhexane sorbed4/ I3 of
activated zeolite
Linde PA, Lot No. 47077 0.237 ... 0.003 0.011 Ba-A Na-A reexchanged from 0.237 Ba-A 0 At room temperature and 20 mm partial pressure of sorbate.
STI
Figure 2. The ion-exchange isotherm for the T1-Na-A system at 0.1 total normality and 25': 0 , T1.+ Na,+.
+
(28) peak, measured with a standard X-ray dsractometer, to the intensity of the peak obtained using Linde 4-A, Lot No. 4353 (for Linde 4-A), and Linde 5-A, Lot No. 5104 (for Ba-A), as standards. A sample of Ba-A was prepared that was 83.2% in the Ba form. The remainder of the cations was Na+. The analyses of this material and the material from which it was prepared (Linde 4-A, Lot No. 4077)
We believed that dehydration under -the conditions used was responsible for the loss in crystallinity of the Ba-A crystals. For verification of this suspicion, an X-ray powder diffraction pattern was taken of a sample of Ba-A that had never been dehydrated. The pattern was very weak and the relative crystallinity (Table 111) was correspondingly low. Despite the low intensity of all the lines, every major line of the Ca-A2 (6) D. W. Breok, W. C. Eversole, R. M. Milton, T. B. Reed, and T. L. Thomas, J . Am. Chem. Soc., 78, 5963 (1956).
Volume 71, Number 6 April 1967
HOWARD S. SHERRYAND HAROLD F. WALTON
1460
Table III : Relative X-Ray Crystallinity of Zeolite Samples % cryrrtallinity
Sample
Linde 4-A, Lot No. 47077, never calcined Ba-A, calcined and rehydrated Ba-A, never calcined Na-A, reexchanged from uncalcined Ba-A, calcined, and rehydrated Na-A, reexchanged from uncalcined Ba-A, and never calcined
90"
Amorphousb 1Sb 85" 90" '0
Ratio of intensity of the 34.1" (28) peak of the sample to that of Linde 4-A, Lot No. 4353, times 100%. Ratio of intensity of the 34.1' (28) peak of the sample to that of Linde 5-A, Lot No. 5104. a
'
pattern was present. The low intensity of the pattern obtained from this sample was probably due to the high absorption coefficient barium has for X-rays. As a check, we then took a sample of Ba-A that had never been dehydrated and reexchanged it with NaCl solution until it was essentially in the pure sodium form (Table I). The analysis, water absorption data, and relative crystallinities shown in Tables I, 11, and I11 indicate that the Na-A obtained by reexchange of Ba-A is identical with the Linde 4-A starting material. Powder diffraction patterns of both materials are identical. It is therefore concluded that barium ion exchange of Linde 4-A does not cause the crystals to become amorphous. The very calcination required to activate the zeolite for sorption causes the Ba-A crystals to collapse and become amorphous to X-rays. This structural collapse is reflected in the lack of water sorption capacity (shown in Table I1 for Ba-A). The loss in stabiljty of the dehydrated Linde A structure in the barium form is probably due to the large internuclear distances that must be expected between the barium ion and the oxygen atoms of the lattice and to the concomitant large charge separation. Previous reports of the decomposition of Linde A caused by barium e ~ c h a n g e must ~ ? ~ have been based on X-ray studies of calcined samples. Although calcination does destroy the crystal lattice, any ion-exchange studies involving Linde 4-A and BaClz a t moderate temperatures in aqueous solution will be valid because the system is reversible. After characterizing our lot of Linde 4-A and demonstrating the reversibility of the systems of interest, these systems were studied a t 5, 25, and 77"-all at 0.1 total normality. The ion-exchange isotherms for the Ca, Sr, and Ba systems a t 25" are presented in Figures 3, 4, and 5. The Ca and Sr isotherms are The Journal of P h y e k d Chemistry
0.2
0.4
0.6
0.8
1.0
sca Figure 3. The ion-exchange isotherm for the Ca-Na-A system st 25" and 0.1 total normality: 0 , Chat 2Na,+, tracer used; 0, CaIp+ 2Na,+, no tracer used.
+
+
0.8
0
3i
N
0.
0
Figure 4. The ion-exchange isotherm for the Sr-Na-A system at 0.1 total normality and 25": 0, Sr.*+ 2Na,+, no tracer used; m, 2Na,+ Sr2+, no tracer used; A, Sr,*+ 2Na+, tracer used.
+
+
+
N 0.4
o*2 ' 0
0.2
0.4
0.6
0.8
1.0
Ba
Figure 5. The ion-exchange isotherm for the Ba-Na-A system a t 0.1 total normality and 25": 0 , Bas4+ 2Na,+, Be.*+, no no tracer used; 0, 2Na.+ tracer used; A, B a P 2Na*+, tracer used.
+
+
+
ION-EXCHANGE PROPERTIES OF ZEOLITES
1461
extrapolated through the point S = 1, Z = 1, whereas the Ba isotherm is allowed to terminate a t the point S = 1, Z = 0.97. Because the curves for all three temperatures lie so close together, the data taken at 5 and 77" are not shown on these figures. The rational selectivity coefficients N K ~ s M were calculated for all the points collected. This quantity is defined by the chemical reaction and equations
+ uB,b+
bASa+
bA,"+
-
8P 0 Y
8
d
+ uB,b+
t
240
I .60
where fA and f B are the rational single-ion activity coefficients of ions A and B in the zeolite phase, YA and YB the molal single-ion activity coefficients in the solution phase, ZA and ZB the equivalent fraction of ions A and B in the zeolite phase, mA and m B the molalities of ions A add B in the solution phase, K the thermodynamic equilibrium constant for the ion-exchange reaction, s and z the subscripts which identify the solution and zeolite phases, and NKBAthe rational selectivity coefficient or concentration quotient. The correct'ed rational selectivity coefficient, KC, is defined as
(4)
For uni-univalent ion exchange, the quotient containing the mean molal activity coefficients was evaluated by assuming that the ratio of mean molal activity coefficients is constant at constant ionic strength. For di-univalent ion exchange, constant ionic strength is not maintained as an isonormal ion-exchange isotherm is traversed. However, since most of the loading of Na-A with alkaline earth ions occurs over a very restricted range of solution compositions (in almost pure 0.1 N NaCl solution), the ionic strength does not vary too much. Mean molal activity coefficients a t 0.1 ionic strength were used to evaluate the quotient of mean molal activity coefficients in eq 5. A much poorer assumption was made when the quotient of mean molal activity coefficients, evaluated a t 25", was assumed to remain constant with temperature. Despite this approximation, it will be seen later that this was satisfactory because reasonable agreement was obtained between the standard enthalpy of reaction for the Ca-Na-A system measured in this work and that measured calorimetricadly a t 25" by other investigators. Plots of log Ko us. zeolite composition are shown in Figures 6-10. One of the salient features of these
0.4
0.8
0.6
I
ZAg
Figure 6. Corrected selectivitv coefficient for the Ai-Na-A system a t 0.1 total normality as a function Na,+ a t 5", of zeolite composition: A, Ag,+ tracer added; 0, Ag.+ Na.+ a t 25", tracer added; a, Ag,+ Na. a t 77", tracer added.
+
+
+
7
*'O 1.6-
B
d
(5)
0.2
f
I
.
.
.
-
0.8.
0.4
.
'0
0.2
0.4
0.6
0.8
1.0
ZTl
Figure 7. Corrected selectivity coefficient for the T1-Na-A system a t 0.1 total normality and 25" as a function of zeolite composition: 0, T1,+ Nag+, no tracer added.
+
figures is that they do not show log Kc reaching a maximum when plotted as a function of zeolite composition, although such maxima have been reported by Barrer and co-worker~.~JInterestingly, Ames also reported no maxima in the same type of plots for the Ca-Na-A and Sr-Na-A systems a t 25°.8 His linear plot of log KC us. ZC, agrees quite well with our plot in Figure 8. His plot of log Kc us. Zsr has the same general shape as our curve in Figure 9, although we do not completely agree. When we made the first measurements by contacting the zeolite and solution phases for 24 hr, a maximum in Kc as a function of Zsr was (7) R. M. Barrer, L. V. C. Rees, and D. J. Ward, Proc. Roy. Soc. (London), A273. 180 (1963). (8) L. L.Ames, Jr., Am. Miwa.Zo&t, 49, 1099 (1964).
Volume 71,Number 6 April 1067
HOWARD S. SHERRY AND HAROLD F. WALTON
1462
2.4
2 .c
1.6 d o
02
y"
1.1
8
A
0.t I .o
0.4
0
0.2
0.4
0.8
0.6
1.0
ZBO C
fca Figure 8. Corrected selectivity coefficient for the Ca-Na-A system a t 0.1 total normality as a function of zeolite composition: A, C a F 2Na.+, 5 O , tracer added; 0, Ca, 2Na.+, 25", tracer added; m, Ca,2+ 2Na,, 25", Cas2+ 2Nas, 77", tracer added. no tracer added; 0,
+
+
+
+
2.1) 2.6
2.4
2.2 2.0
I .a I.6
I.4
$4
1.2
3
1.0
0.8
0.6
0.0 -0.2
-
0.4 0.6
o
0.2
0.4
4r
aa
0.8
1.0
Figure 9. Corrected selectivity coefficient for the Sr-Na-A system a t 0.1 total normality as a function of zeolite composition: A, Sr,g+ 2Na,+, 5 O , tracer added, equilibrium reached; 0, Sr,l+ 2Na,+, 25O, no tracer added, 24-hr contact time; D, Sr,%+ 2Na,+, 2 5 O , tracer added, 24hr contact time; V, Sr.'+ 2Na.+, 2 5 O , tracer added, equilibrium reached; 0, Sr.*+ 2Na.+, 77", tracer added, equilibrium reached.
+
+
+
+
+
found. A check of whether or not equilibrium waa reached was made by reexchanging SPA with NaCl The Journal of Physical Chembtrv
Figure 10. Corrected selectivity coefficient for the Ba-Na-A system a t 0.1 total normality: A, Baa2+ 2Naa+at 5 " ; 0, Ba,$+ 2Na, at 25'; m, Bas2+ 2Na, at 77'.
+
+
+
for 24 hr. No hysteresis loop was observed in the ionexchange isotherm. The use of this technique did, however, limit the fractional loading of Na+ onto the Sr-A to high Sr2+ loadings because of the very unfavorable equilibrium (see Figure 4). When measurements were made a t low Sr2+loadings by loading Sr2+ onto Na-A, it was initially assumed that equilibrium was reached in 24 hr. (These points are shown in Figure 9,) When the uptake of radioactive *5Sr2+by the zeolite was later followed as a function of time, by assaying the liquid phase for its radioactive content, it was found that, a t low Sr2+loadings, it took 72 hr to reach equilibrium. At the high loadings the equilibrium state was attained in 24 hr, which is in agreement with our early work a t high Sr2+loadings. It is therefore concluded that there is no maximum in plots of selectivity coefficient us. ZA in the alkaline earthNa-A systems. This conclusion is probably valid for other zeolites as well. Another salient feature of the plots in Figures 6-10 is the presence or absence of parallel curves. Those for the Ag-Na system a t 5, 25, and 77" (Figure 6) are all parallel to one another, indicating that the heat of partial exchange remains constant with Ag+ loading. The selectivity of the zeolite for Ag+ is extremely high, indicating extensive binding of Ag+ by the negatively charged aluminosilicate network. The calcium system (Figure 8) shows a markedly nonuniform temperature dependence, consistent with the presence of calcium ions in different environments. It is possible that most of the Ca2+ forms ion pairs with the fixed anions and are located within the rings of six tetrahedra that line the large cavity of .Linde A and that a small number of Ca2+ions are hydrated or are "dissolved" in the zeolite water.
ION-EXCHANGE PROPERTIES OF ZEOLITES
1463
Table IV: Standard Free Energies, Enthalpies, and Entropies of Reaction Reaction
Ag,+
+ Na.+
+ Nan+ Ca,*+ + Na, Sr,*+ + Na. Ba,a+ + Na,
- AFOns,
- AFOws,
- AFOw,
AHo,
AS',
cal/equiv
cal/equiv
cal/equiv
cal/equiv
eu/equiv
3870 f 40
3930 f 40 2320 f 20 733 f 300 1010 f 10 1168 f 20
4200 f 40
-2780 f 450
3 . 9 f0 . 6
1320 =k 20 1314 f 15 1370 f 20
2700 f 300 500 f 200 0 f 200
11.5 f 1 5.1 I 2 3 . 9 f0 . 8
TI,+
495 f 10 970 f 10 1088 f 20
A smaller temperature dependence of Kc is observed in the Sr-Na system (Figure 9). However, the curves of log Kc us. Zsr at different temperatures definitely do cross, again indicating the presence of dissimilar cation positions, although in this system the energetic heterogeneity seems less than in the Ca system. Also, the bend in the curves after 40% Sr loading indicates that there are at least two different cation positions. It is probable that 50-6070 of the Sr2+forms ion pairs with the anions fixed in the framework and that the rest exist as fully hydrated ions. The Barium system (Figure 10) is interesting because there is no measurable temperature dependence of Kc over the temperature range studied. Not only is there no dependence, but the partial heats of exchange must be uniformly zero over the whole range of Ba2+loading, which may mean that the Ba2+ions are all bound equally strongly by the anionic surface. The small or zero dependence of the corrected selectivity coefficient indicates that electrostatic interactween Ba2+and the lattice were weak. All the barium ions are probably hydrated in the large cage and are bound only in the sense that the hydrated cation is localized by the electric field of the aluminosilicate matrix. The method of Gaines and Thomase was used to calculate standard free energies, standard enthalpies, and standard entropies of exchange from the corrected selectivity data given in Figures 6-10. We employed a modified form of their equation that neglects imbibement of electrolyte by the crystals and assumes the water activity to be unity. Our working equation is In K = (b
- a) +
1
In KcdZA
(6)
The first term on the right-hand side of eq 6 is the difference between the charge of the two ions. The second term on the right-hand side is graphically evaluated by calculating the area under the curves in Figures 6-10. The standard free energies, enthalpies, and entropies of reaction were then computed using the relations
A P T = -RT In K
(7)
d-l-n K - AH' dT RT2 and
AS" =
AH"
- AF'T T
(9)
The results are tabulated (Table IV) on the basis of 1equiv of exchanger. The standard enthalpy of exchange is positive for the complete exchange of each alkaline earth ion for Na+ in Linde A. The standard entropy of exchange is, however, in each case a large enough positive quantity that the standard free energy of exchange is negative. Thus it is the entropy function that is responsible for the zeolite's preferring alkaline earth ions t o sodium ions. The positive values of the enthalpy are predicted for this site spacing by Barrer and Falconer's simple theory. The data in Table IV show the danger of neglecting to consider the entropy contribution to the free energy of reaction. Other published data comparable with ours are those of Barrer and Meier? Barrer, Rees, and Ward,' and Ames.* Barrer and co-workers6~' investigated alkali metal ion exchange and Ca2+ exchange for Na+ in Linde 4-A. They measured heats of reaction calorimetrically and obtained, for the Ca-Na-A system, a standard enthalpy of reaction of +2100 f 50 cal/ equiv of zeolite. Our value of +2700 f 300, obtained by determining AG", as a function of temperature, is in reasonable agreement, although the direct calorimetric measurement of AH" is certainly preferable to the technique used in our work. The free energy of exchange reported by Barrer, Rees, and Ward' is -140 20 cal/equiv of zeolite at 298°K. We report -733 300 cal/equiv of zeolite. Barrer and co-workers also used the method of Gaines and Thomas8 to compute the free energy of the exchange reaction. However, their plot of log Kc us. Zc. has a maximum and
* *
(9) G. (1953).
L. Gaines and H. C. Thomas, J. C h m . Phys.,
21, 714
Volunw 71, Number 6 Aptdl 1967
HOWARD S. SHERRY AND HAROLD F. WALTON
1464
the area under the curve is less than the area under our curve a t 25". Although the free energy of exchange reported by Barrer, Rees, and Ward is significantly different from the value reported herein, the calorimetrically determined enthalpy of exchange may be expected to agree with that measured by us. Failure to reach equilibrium a t low Ca2+ loadings can result in a marked error in the value of the selectivity coefficient without the zeolite composition being noticeably in error. It occurs because a t low Ca2+ loadings the zeolite takes up Ca2+ almost quantitatively from solution. Toward the end of the reaction, the concentration of Ca2+sometimes reaches 5 X N and may reach 5 X N a t equilibrium. If the reaction is stopped too soon, the measured concentration of Ca2+ in solution may differ from the equilibrium value by a factor of 10. It makes a marked difference in the calculated value of the selectivity coefficient, but little difference in the zeolite composition, because the additional number of ions that must diffuse into the zeolite for the system to attain equilibrium is insignificant compared to the total number of Ca2+ in the zeolite. Because essentially all of the ions have exchanged into the zeolite even when the reaction is stopped too soon, the heat of reaction is correctly determined. The entropy of exchange will, however, be incorrect because of the error in the free energy. Amess has studied the Ca-Na-A and Sr-Na-A systems a t 25°.8 He reports values of -850 and - 1300 cal/equiv of zeolite for the free energy of exchange of Ca2+and Sr2+into Na-A. These values agree reasonably well with our values of -730 and -1010 (Table IV) . The observed increase in the entropy of the system must be considered as consisting of two contributionsone from the aqueous phase and one from the zeolite phase. In the aqueous phase one alkaline earth cation is replaced by two sodium ions. This frees water molecules and thereby increases the entropy of the solution phase for Ca2+and Sr2+because either one of these two cations is more hydrated than two Na+. For Ba2+, the replscement of one Ba2+by two Na+ results in an entropy decrease because the large Ba2+ disrupts the local water structure. These changes can be calculated using the data presented by Rosseinsky.'O In the zeolite phase one alkaline earth cation replaces two sodium cations. This standard entropy change can be calculated using the experimental values of the standard entropy of reaction and published standard entropies of hydration of the ions, l o from the relation AS" =
('/zSAZ+~'
+
SNae')
+(
3 ~ -2l/23A2thyd) ~ ~ (10)
The Journal of Physic& Chemistry
The values of the two terms on the right-hand side of eq 10 are given in Table V. Table V: The Zeolite and Solution Phase Contributions to the Of Reaction Ag
l/zSAa+*'
- SNae'
S N ~ ~ Y'/JA%@~'J ~
+
2.5 +1.4
Ca*
7.3 +4.2
Sr2
1.7 $3.4
Bat
5.85 -1.95
The data in Table V indicate that the entropy in the zeolite phase increases upon exchange of alkaline earth and silver cations for Na+ ions. This increase in entropy exceeds the one that occurs in solution. It decreases as one descends the alkaline earth series and then increases at the Ba2+ ion. The entropy increase of 2.5 eu/equiv of zeolite in the case of Ag+ ion exchange can only be accounted for by assuming that Ag+ ions coordinate directly to lattice oxygen atoms and thereby free the intracrystalline water. The extremely large increase in the entropy of the zeolite resulting from Ca2+ion exchange can be only partially accounted for by the increase in the number of ways of distributing 6 Ca2+ ions among 8 positions as compared to placing 8 Na+ ions into exactly 8 positions. Again, the only process we can imagine which will lead to an increase in entropy is the binding of Ca2+cations to the lattice oxygens and the concomitant release of water of hydration. This "free" intracrystalline water should possess more entropy than water coordinated to cations. In the case of Sr2+ion exchange the entropy change in the zeolite phase is much smaller and this, coupled with the evidence for site heterogeneity provided by the selectivity curves in Figure 9, indicates that perhaps half of the Sr2+ions coordinate to lattice oxygen atoms and half are "dissolved" in the zeolitic water. The large increase in entropy which result,s from Ba2+ ion exchange of Na-A probably is not caused by binding of Ba2+ ions to the negatively charged framework, thereby freeing water, because the enthalpy of exchange is zero or at least a very small number. It is probable that the water content of the Ba-A crystals is lower than that of the Na-A crystals. This means that during the ion-exchange reaction there is a net transfer of water molecules from the zeolite phase to the solution phase, a process which should lead to an entropy increase.
Conclusions It has been shown that one-third of an Na+ ion per unit cell cannot be replaced by a T1+ ion at 25". Be(10) D. R.Rosseinsky, C h a . Rev., 65, 467 (1965).
THERMODYNAMICS OF THE SODIUM-BISMUTH SYSTEM
1465
cause this corresponds to the amount of excess aluminum atoms in the unit cell, we have verified the work of Barrer and Meier4*5in this respect. We were unable to verify the maximum in the plot of log Kc us. Zca obtained by these6 and other? workers. We have therefore concluded that they did not reach the equilibrium state at low Ca2+loadings. We have also rechecked the Ba2+ exchange of Linde 4-A and have concluded that it does not cause the crystal structure to collapse at 25' and 0.1 total normality.
The entropy and enthalpy changes lead to the conclusion that there is much ion-pair formation in Ca-A, less in Sr-A (perhaps only 40% of the Sr2+form ion pairs), and none in Ba-A.
Acknowledgments. H. S. S. wishes to thank the Mobil Oil Corp. for its support of part of this work and both of us wish t o thank the U. S. Atomic Energy Commission for its support under Contract No. AT(11-1)499.
Liquid-Vapor Phase Diagram and Thermodynamics of the Sodium-Bismuth System
by Albert K. Fischer, Stanley A. Johnson, and Scott E. Wood Argonne National Laboratory, Argonne, Illinois
(Received October 4, 1966)
Liquid-vapor equilibria in the Na-Bi system were studied by the transpiration method at 1173'K and by the quasi-static and boiling point methods of measuring total vapor pressure over a temperature range. The total pressure curves indicated the appearance of three-phase equilibria (involving vapor, liquid, and solid Na3Bi) below a pressure of about 240 torr. The melting point of Na&i was found to be 840'. A quasi-ideal solution treatment, assuming the presence of the compounds Na3Bi and NaBi as species and with respective formation equilibrium constants of 2 X lo5 and 300 at 1173'K, was able to fit the observed excess chemical potentials of the components.
Introduction As part of a research program on the thermodynamic properties of binary systems of fused metals, one component of which is an alkali metal, a study was made of liquid-vapor equilibria in the sodium-bismuth system. This study took the form of measuring the total vapor pressure as a function of temperature for a series of compositions and of measuring the vapor density and composition by means of transpiration studies at 1173'K for a series of compositions. The interpretation of the thermodynamic behavior of the system in terms of a quasi-ideal solution treatment was successful in fitting the data. Underlying this treatment was the
assumption of the existence of the compounds NaaBi and NaBi in the liquid state. While not a proof of the existence of the compounds, the success of the quasiideal approach does serve to indicate, in a perhaps oversimplified way, the likelihood of rather specific short-range interactions in these fused systems.
Experimental Section The total vapor pressures were measured by either the quasi-static (Rodebush-Dixon) method' or the boiling point method. Our adaptation of the Rodebush(1) W. H.Rodebush and A. L. Dixon, Phys. Rev., 26, 851 (1925).
Volume 7 1 , Number 6 April 1967
