J. Phys. Chem. 1984,88, 2687-2689
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Xenon as a Supercritical Solvent Val J. Krukonis, Phasex Corporation, Nashua, New Hampshire 03060
Mark A. McHugh,* and Andrew J. Seckner Department of Chemical Engineering, University of Notre Dame, Notre Dame, Indiana 43556 (Received: May 5, 1984)
The solubility of naphthalene in supercritical xenon is determined at temperatures of 34.9,39.8, and 45.0 "C and at pressures ranging from 100 to 225 atm. A section of the P-T trace of the xenon-naphthalene solid-liquid-gas line is also determined. At high pressures xenon depresses the melting point of naphthalene 37.5 O C , indicating that xenon is very soluble in liquid naphthalene. We believe that this is the first study to demonstrate that xenon, a noble gas, exhibits enhanced supercritical solvent properties comparable to the solvent properties of supercritical methane, ethane, ethylene, and carbon dioxide.
Introduction Numerous experimental studies have been performed to determine the solubility of gaseous xenon in a wide variety of organic substances ranging from normal alkanes to olive oil.' More recently, a number of investigators have studied the solubility behavior of various organic materials in liquid x e n ~ n . ~Certain .~ polymers3 and large organic hydrocarbons2 were found to dissolve in liquid xenon at temperatures ranging from 200 to 287 K. Everett and Stageman3 were the first to demonstrate that liquid xenon is a suitable solvent for certain polymers (poly(dimethy1siloxane), molecular weight of approximately 100 000) and that liquid xenon can be a suitable solvent for stable colloidal dispersions. In this work experimental evidence is presented which shows that xenon at conditions above its critical temperature and critical pressure can be a good solvent for organic solids. Although a noble gas, xenon has very convenient critical properties (T, = 289.8 K, P, = 58.0 atm). Xenon also possesses a polarizability of 4.01 X 10-24 cm3 14 which is relatively large when compared to the other noble but which is very close to the polarizability of the lower alkanes3 such as ethane3>l4and p r ~ p a n e .In ~ fact, the solvent capabilities of liquid xenon are attributed to its high polarizabilit~.~,~ In this study the solubility of naphthalene (T, = 80.1 "C) in supercritical xenon is measured at several temperatures above the critical temperature of xenon and over a range of pressures from 100 to 225 atm. A section of the pressure-temperature (P-T) trace of the three-phase solid-liquid-gas (S-L-G) freezing point depression curve for the naphthalene-xenon system is also determined. A schematic representation of the P-T region investigated in these experiments is shown in Figure 1 . Experimental Section Both types of experiments, the solubility measurements and the S-L-G line determination, are performed with the same experimental apparatus: a variable-volume, high-pressure view cell. The apparatus and the experimental procedures, which are described in detail el~ewhere,~ are similar to those of Liphard and Schneiders and Van Welie and Diepen6 So that solubility information could be obtained, a measured amount of naphthalene and xenon is charged to the view cell. At a fixed composition the pressure of the mixture is increased isothermally until all of the naphthalene is dissolved by the xenon. At this point a clear fluid (1) For a review of xenon solubility in a large variety of liquids see Solubility Data Ser., 2 (1979). (2) P. M. Rentzepis and D. C. Douglas, Nature (London), 293, 165 (1981). (3) D. H. Everett and J. F. Stageman, Trans. Faraday SOC.,230 (1978). (4) M. A. McHugh, A. J. Seckner, and T. J. Yogan, Ind. Eng. Chem. Fundam. in press. (5) K. G. Liphard and G. M. Schneider, J. Chem. Thermodyn., 7, 805 (1975) \ - - . - I .
(6) G. S.A. Van Welie and G. A. M. Diepen, R e d . Trau. Chim. Pays-Bas, 80, 659 (1961).
0022-3654/84/2088-2687$01.50/0
TABLE I: P-T Trace of the Xenon-Naphthalene S-L-G Line Obtained in This Study T, OC P,atm T, OC P, atm 42.55 42.55 45.55 46.29 48.32 51.02
f 0.03 f 0.03 f 0.03 f 0.05 f 0.03 f 0.06
123.0 f 0.4 124.2 f 0.6 96.9 f 0.7 100.7 f 1.0 90.1 i 0.7 83.3 f 0.7
54.74 57.52 62.97 69.26 73.33
f 0.05
f 0.04
f 0.06 i 0.08 f 0.03
71.6 f 0.5 63.1 f 0.5 46.7 f 0.5 28.8 f 0.7 16.0 f 0.7
phase exists in the view cell. The pressure is then slowly decreased until solid (or liquid) naphthalene precipitates from solution. The solubility point is in the pressure interval between the single supercritical fluid phase state and the two-phase supercritical fluid-solid (or liquid) naphthalene state. The P-T trace of the S-L-G freezing point depression curve is obtained in the following manner. Again naphthalene and xenon are charged to the view cell. While under isothermal conditions, the pressure in the cell is increased until the solid naphthalene begins to liquify. At this point the pressure is decreased until the liquid begins to solidify. Hence, the S-L-G line exists in the pressure interval between the fluid-solid and the fluid-liquid condition in the view cell. (In these initial melting point depression determinations no attempt is made to determine the concentration of the equilibrium liquid and gas phase on the S-L-G line.)
Results and Discussion Listed in Table I and shown in Figure 2 are the results of the determination of the P-T projection of the naphthalene-xenon S-L-G line. On the basis of measurements reported here, the S-L-G line exhibits a temperature minimum very close to 42.5 OC. As the pressure is increased the S-L-G line, which originated at the melting point of naphthalene (80.1 "C), begins with a negative slope (ie., d P / d T < 0), passes through a temperature minimum, and then continues with a positive slope to the upper critical end point (UCEP). The melting point depression of naphthalene with supercritical methane,' ethane,* ethylene: and carbon d i o ~ i d e ' ~has J ~also been (7) J. A. M. Van Hest and G. A. M. Diepen, Symp. SOC.Chem. Ind. London, 10 (1963). (8) G . S. A. Van Welie and G. A. M. Diepen, J . Phys. Chem., 67, 755 ( 1 963). (9) C. A. Van Gunst, \ - - - - I
F. E. C. Scheffer, and G. A. M. Diepen, J . Phys. Chem., 57, 578 (1953). (10) M. A. McHugh and T. J. Yogan, J . Chem. Eng. Data, 29, 112 (19R4) \-__. (1); M. A. McHugh and M. E. Paulaitis, J. Chem. Eng. Data, 25, 326 (1980). (12) Yu.Y. Tsekhanskaya, M. B. Iomtev, and E. V. Mushinka, Russ. J . Phys. Chem., 38, 1173 (1964). (13) Yu. Y. Tsekhanskaya, M. B. Iomtev, and E. V. Mushinka, Rum. J . Phys. Chem., 36, 1177 (1962). (14) G. W. Castellan, "Physical Chemistry", 2nd ed, Addison-Wesley, Reading MA, p 625.
0 1984 American Chemical Society
The Journal of Physical Chemistry, Vol. 88, No. 13, 1984
2688
Letters TABLE 11: Xenon-Naphthalene Solubility Data Obtained in This Study P,atm naDhthalene. wt % T = 25.2 O C 216.6
* 1.7
T = 34.9
3.21 O C
119.4 & 1.4 123.5 f 1.4 169.4 f 1.3
2.83 3.21 3.85
T = 39.8 OC
116.0 i1.4 127.9 f 1.0 146.8 f 1.0 150.7 f 2.8
T = 45.0
'CI
Figure 1. The cross-hatched area represents the P-T region investigated in this study. CD and MH are pure component vapor pressure curves, MN the heavy component melting curve, and EM the heavy component sublimation curve. The dashed lines represent critical mixture curves.
I70 O
r
E
a
loo-
h
E 150
W
w
W
a
E 140 3
3 v, 0
% a
------I
I60
h
+
3.21 3.85 5.23 5.44 6.58 9.56
*
I80
5
O C
114.1 f 2.6 121.4 f 0.7 125.0 f 0.6 1.0 126.2 132.0 f 1.0 156.4 f 1.0
TEMPERATURE
I
3.21 3.85 5.23 5.44
m
CI)
50-
$ 130 a 120
01
-
35
o T=45.0°C
I IO I
I
1
55 65 75 TEMPERATURE ("C)
45
85 104
I
I
I
I
I
I
Figure 2. Experimental P-T trace of the S-L-G line for the naphthalene-xenon system starting at the melting point of naphthalene. determined. Of these supercritical fluids ethylene causes the largest melting point depression of naphthalene (Le., from 80.1 to 52.1 O C ) . With supercritical xenon a 37.5 O C melting point depression is observed indicating that xenon is more soluble in liquid naphthalene than any of the other gases listed.16 Some interesting density phenomena are observed during the solubility and melting point investigations. Along the S-L-G line at pressures above approximately 80 atm, supercritical xenon becomes more dense than solid naphthalene (specific gravity 1.145), and solid naphthalene floats to the top of the equilibrium view cell. However, when solid naphthalene melts as the S-L-G line is crossed at a higher pressure, the resultant naphthalene-rich liquid phase is more dense than the xenon-rich gas phase. Now the naphthalene-rich liquid phase is on the bottom of the cell. At (15) M. A. McHugh, Ph.D. Thesis, University of Delaware, Newark, DE, 1981. (16) J. De Swaan Arons and G . A. M. Diepen, Red. Trau. Chim. PaysBas, 82, 249 (1963).
Figure 3. Solubility of naphthalene in supercriticalxenon at 34.9, and 45.0 'C.
39.8,
still higher pressures the xenon-rich gas phase becomes more dense than the naphthalene-rich liquid phase, and therefore the xenon-rich phase settles at the bottom of the view cell. This unusual behavior is a function of the partial molal volumes of both naphthalene and xenon in both the liquid and gas phases. Further work is being carried out to quantify these observations and to extend the S-L-G line to the upper critical end point (UCEP). The solubility data for naphthalene in supercritical xenon at 25.2, 34.9, 39.8, and 45.0 OC are reported in Table I1 and shown in Figure 3. As shown in Table 11, and as discussed above, some of the solubility data are for solid naphthalene dissolved in supercritical xenon, and because the S-L-G line is crossed in some of the tests, some of the data are liquid naphthalene solubilities in supercritical xenon. At 25.2 OC,more than 200-atm pressure is needed to solubilize 3.0% (w/w) naphthalene in supercritical xenon. At 34.9 O C , considerably less pressure is required to solubilize up to 3.8%
J. Phys. Chem. 1984, 88, 2689-2697 (w/w) naphthalene. As shown in Figure 3, at this temperature the solubility quickly reaches a limiting value at high pressures. At 39.8 OC, the solubility of solid naphthalene in supercritical xenon increases at a much faster rate than at the two previous temperatures, however, again the solubility quickly reaches a limiting value of approximately 5.4% (w/w) at high pressures. At 45.0 O C the solubility isotherm exhibits behavior which is significantly different from that of the other solubility isotherms. A large increase in naphthalene solubility occurs when the pressure is increased above only 120 atm. This large increase in naphthalene solubility is a consequence of operating very close to the mixture UCEP. At much higher pressures the solubility also reaches a limiting value, but now this limiting solubility (9.5% (w/w)) is much greater than the limiting solubility at any of the other isotherms. On the basis of solubility behavior depicted in
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Figure 3 the UCEP should be in the vicinity of 125 atm and 47 “C. Notice that the 45 OC isotherm represents liquid solubilities at low pressures and solid solubilities at high pressures. This solubility behavior verifies that the S-L-G line is crossed at two different pressures at this temperature. Hence, a temperature minimum must exist in the S-L-G line in the vicinity of 45 OC. As previously mentioned, the solubility of naphthalene has been determined in supercritical methane,’ ethane,* ethylene,6 and carbon On the basis of our findings, naphthalene solubility in supercritical xenon is higher on a mole fraction basis than any of the above-mentioned supercritical solvents at equivalent pressures and temperatures. We believe that this is the first time a noble gas has been shown to be a good supercritical solvent. Registry No. Xenon, 7440-63-3; naphthalene, 91-20-3.
FEATURE ARTICLE Ionic Fluids Kenneth S. Pitzer Department of Chemistry and Lawrence Berkeley Laboratory, University of California, Berkeley, California 94720 (Received: December 6, 1983)
The statistical mechanics and thermodynamics of ionic fluids are reviewed and discussed on a corresponding-states basis and with emphasis on areas of recent advances. First, the vapor-liquid and critical properties are considered for pure ionic fluids such as NaCl as well as for the primitive ionic model. The underlying reasons are discussed for differences between fluids of ions and those of neutral molecules. Then, systems involving a solvent are considered with emphasis on simple systems showing phase separation or with strong ion pairing. Estimates are made for the critical curve for the system NaC1-H20 to the critical point of pure NaCl at 3900 K.
Introduction Ionic fluids have long been of interest to physical chemists. The early interest was in very dilute aqueous electrolytes. More recently, concentrated aqueous electrolytes, including mixtures, have received attention including systems continuously miscible to the fused salt. Highly charged ions introduce new features which also appear for 1:l electrolytes in solvents of low dielectric constant. Water becomes such a solvent at very high temperature. The critical properties and phase relationships of a pure salt such as NaCl represent an even more extreme case with unit dielectric constant. In this paper I will discuss several of these topics where there have been recent advances starting with the case of the critical properties of pure NaCl and of the primitive, hard-core, ionic fluid model. With certain approximations, ionic fluids should follow the principle of corresponding states. Consequently, it is useful to consider the properties of various systems on the basis of the reduced variables T I = T / T, and VI = V / V,. Figure 1 presents an overview in terms of reduced temperature and volume with molar concentration also indicated. The solid curve shows the liquid and vapor properties of pure NaCl with a critical temperature of 3900 K and volume of 490 cm3.mol-1. The phase boundaries for a molecular fluid, argon, are shown as a dashed curve for comparison. For various ionic systems in solvents the expected relationships for the critical constants are T,” = T , ’ ( Z ’ ’ / Z ~ 2 ( D ~ ’ / D ’ ‘ U ’ ’ ) 0022-3654/84/2088-2689$01.50/0
V / = V,‘(a”/a’)3 where 2 and a are the number of charges and the collision diameter, respectively, for the solute and D is the dielectric constant (relative permittivity) of the solvent. The approximate reduced temperatures are indicated in Figure 1 for simple aqueous 1:l electrolytes at 300 and 573 K and for 2:2 electrolytes at 300 K. Since for ionic systems in solvents the equivalent temperature is given by the product DT, one obtains a high reduced temperature in real systems by raising D even at the expense of some reduction in T . An aqueous 1:l system at room temperature has a very high reduced temperature. Indeed, it is impossible to obtain simple ionic fluids at such high reduced temperatures because there is dissociation yielding electrons. The resulting plasmas are interesting systems but will not be considered in this paper. The saturated vapor of a pure ionic fluid is comprised almost entirely of neutral molecules. Ion pairs predominate with substantial concentrations of dimers M,X, and some larger clusters. It is the liquid rather than the vapor where there is a major difference between the ionic fluid and a fluid of neutral molecules. One difference is the much greater thermal expansion of the ionic liquid; this is seen by comparison of the solid and dashed curves for the liquid volume in Figure 1. The ionic fluid is very highly expanded at the critical point. These characteristics will be explained in terms of the Coulombic potential with its repulsive nature between ions of the same sign. With increase in reduced temperature at vaporlike dilution, the pattern gradually changes to a mixture of ions and ion pairs and 0 1984 American Chemical Society
