J. Phys. Chem. 1987, 91, 3059-3062
loo -
I
I
3059
are estimated from the film thickness dependence of the relative yield of the 630-nm (5,10,15,20-tetraphenylchlorinato)zinc(II) (ZnTPC) emission band in a ZnTPP film on Pt and that of the relative yield of the 664-nm fluorescence band in a H2TPP film on Pt, respectively. It has been pointed out that the quenching of sensitized chlorin emission in ZnTPP can be explained by Forster-type energy transfer, but the additional long-range effect in fluoresecence quenching of H2TPP film can be caused by effective energy migration due to the exciton diffusion within the film. The depletion layer thickness Woat the Al/porphyrin interface was determined by low-frequency capacitance measurementsis at zero applied bias according to the equation W = e,eo/C, where Cis the capacitance. The dielectric constant used here to evaluate the thickness of the depletion layer is 2.0, being estimated for an octaethylporphyrin film by capacitance m e a s ~ r e m e n talthough ,~ the value of 2.0 seems to be lower estimated compared with the dielectric constants of 3.42 and 4.0 obtained for a j3-phthalocyanine24and a (phthalocyaninato)copper(II) film.25 However, the obtained quantities are in fairly good agreement with the values of W,, represented above. Of course, it should be noted that the physical meaning of Wapis somewhat different from that of the depletion layer thickness Wo. For analysis of our results we assume that excitons are created by light absorption and dissociate into free carriers in proportion to the electric field at the barrier, in which the ionized impurities or traps are distributed exponentially. Certainly, the fluorescence spectroscopic study of porphyrin films has yielded satisfactory results to prove the existence of excitons as shown above. However, the direct band-to-band excitation cannot be denied absolutely in the carrier formation processes presented here to explain the results.
I
1
50 100 Film thicknorr of H2TPP/nm
Figure 7. Plot of the value of iAu/iAlas a function of the thickness of a HzTPP film. Circles are the observed values, and lines are obtained by calculation on the basis of the exponential field-dependent quantum efficiency of carrier generation at the Al/H*TPP barrier. (1) Wap= 60 nm, Le,, = 15 nm; (2) Wap = 50 nm, Le,,,,= 15 nm; and (3) Wap= 30 nm, Le,, = 15 nm.
In the case of ZnTPP films the best fit is obtained for Le,, = 3 f 2 nm and W,, = 14 f 1 nm (Figure 6). In a similar fashion, we could obtain Le,, = 15 f 10 nm and Wap = 39 f 2 nm for H2TPP (Figure 7). The values of Le,, and Wapevaluated for various prophyrin films are summarized with the absorption constants in Table I. It should be noted here that for n-type porphyrins such as T(3-Py)P and T(4-Py)P the porphyrin/Au contact is blocking, while the porphyrin/Al contact is Ohmic. The finding that the Le,, for ZnTPP is smaller than that for H2TPP is the same as observed for critical energy transfer distances h of 1.3 nm for ZnTPP and 2.4 nm for H2TPP.23 They
Acknowledgment. The authors are grateful to the Japanese Ministry of Education, Culture and Science for supporting this work through the Grant-in-Aid for Energy Research No. 61040040 and thank Dr. Tadayoshi Sakata for helpful discussions. ZnTPP, 14074-80-7. Registry No. H,TPP, 917-23-7; (23)Tanimura, K.; Kawai, T.; Sakata, T. J . Phys. Chem. 1980,84,751. (24)Popovic, 2.D.;Sharp, J. H. J . Ckem. Phys. 1977,66,5076. (25)Hoshino, Y.J . Appl. Phys. 1981,52, 5655.
Solvent Properties of Supercritlcal Xe and SF, R. D. Smith,* S. L. Frye, C. R. Yonker, and R. W. Gale Chemical Methods and Separations Group, Chemical Sciences Department, Battelle Pacific Northwest Laboratory,t Richland, Washington 99352 (Received: September 5, 1986; In Final Form: January 8, 1987)
The UV absorption solvatochromic shift for a probe solute (2-nitroanisole) has been used to explore the solvent properties of supercritical SF6, Xe, and C2H6. The A* polarity/polarizability scale established by Kamlet, Taft, and co-workers has been used to compare solvent properties with liquids and previously studied supercritical fluids (C02 and NH3). The A* measurements for Xe show that it is a less polar/polarizable solvent then C 0 2 at similar reduced densities. SF6 was found to have lower A* parameters than Xe, although it has a greater polarizability. A change in the dependence of A* upon density was observed at reduced densities of 0.70 f 0.15for all fluids studied.
Introduction
One of the most interesting and useful features of supercritical fluids is their variable solvating power. The solubility of naphthalene in c o 2 , for example, increases tremendously as the Operated by Battelle Memorial Institute.
0022-3654/87/2091-3059$01.50/0
pressure is increased near the critical point, much more than can be accounted for by the variation in naphthalene's vapor Pressure.' This phenomenon is generally observed for any solute in a supercritical solvent, and it is possible to adjust solvent properties (1) McHugh, M.; Paulaitis, M. E. J . Chem. Eng. Dura 1980,25, 236.
0 1987 American Chemical Society
3060 The Journal of Physical Chemistry, Vol. 91, No. 11, 1987
Smith et al.
over a wide range by manipulation of fluid temperature and density. The variable solvent properties are utilized, for example, in supercritical fluid chromatography: in chemical reaction media for chemical proce~ses,~ and for extraction or fractionation of mixture^.^ While it is known that fluid density has a profound impact on solvent-solute interactions in a supercritical solvent, there is little detailed information on the nature of the interactions or how they change with density. Recent reports by other authors have described solubility measurements in supercritical xenon and its application as a mobile phase in supercritical fluid ~hromatography.'*~In this paper we report solvatochromic measurements for supercritical sulfur hexafluoride and xenon as a function of fluid density. These fluids are of both fundamental and applied interest. Both are spherically symmetric molecules which should provide the simplest probe of solutesolvent interactions for supercritical fluid densities ranging from near the gas- and liquid-phase limits. In addition, results for ethane, another nonpolar solvent, and the various subcritical liquids are also presented and compared to previous results for carbon dioxide and ammonia.'
Experimental Section A Cary Model 14 spectrophotometer was used to measure the solvatochromic shift in the UV absorbance spectra as a function of fluid density by similar techniques as reported previously.' The high-pressure sample cell was constructed from stainless steel and used sapphire windows with a 1.25-cm thickness. A gas-tight seal was made between the window and the cell body by compression of a Teflon O-ring. The resulting path length was approximately 1 cm with a total sample volume of ca. 0.3 mL. The cell was equipped with fluid inlet and outlet lines. The cell body was wrapped with insulated nichrome wire for heating. Temperature control to f0.5 OC was provided by a West (East Greenwich, RI) Series 1400 single mode controller. Pressurized solvent was provided by a high-pressure syringe pump (High Pressure Equipment, inc., Erie, PA). Fluid pressure was monitored with a Serta Systems (Acton, MA) Model 204 transducer (f10 psi). A Rheodyne, Inc. (Cotati, CA), HPLC sample valve placed between the pump and the cell was used to introduce the solute probe (2-nitroanisole). Spectroscopic measurements were made under static conditions by stopping fluid flow through the cell once the presence of the solute probe was detected and adjusted to the desired concentration by dilution. The concentration of 2-nitroanisole probe was maintained at - 5 X lo4 M for all studies. CzH6,SF6,and Xe were obtained from Scott Speciality Gas (Plumsteadville, PA) and were used without further purification. 2-Nitroanisole (99+%) was obtained from Aldrich Chemical Co. (Milwaukee, WI) and was used as received. Results and Discussion A number of different solvent scales have been devised in order to quantify the differences between liquid solvents? The K* scale of Kamlet and Taftg was developed to correlate different solvent-solute interactions based upon the solvatochromic effect of the solvent on the K-K* electronic transition of suitable solute probes. The basic relationship contains terms to account for solvent polarity (i.e., dipolarity) and polarizability, solvent acidity (hy(2) Smith, R. D.; Chapman, E. G.; Wright, B. W. Anal. Chem. 1985,57, 2829-2836. (3) Subramaniam, B.; McHugh, M. A. Ind. Eng. Chem. Process Des. Dev. 1986, 25, 1-12. (4) McHugh, M. A.; Krukonis, V. J. Supercritical Fluid Extraction; Butterworths: Boston, MA, 1986. ( 5 ) Krukonis, V. J.; McHugh, M. A.; Seckner, A. J. J . Phys. Chem. 1984, 88, 2687. (6) French, S.B.; Novotny, M. Anal. Chem. 1986, 58, 164. (7) Yonker, C. R.; Frye, S. L.; Kalkwarf, D. R.; Smith, R. D. J. Phys. Chem. 1986, 90, 3022. (8) (a) Grunwald, E.; Winstein, S. J. Am. Chem. Soc. 1948,70, 846. (b) Smith, S. G.; Fainberg, A. H.; Winstein, S. J. Am. Chem. Soc. 1%1,83, 618. (c) Berson, A.; Hamlet, E.; Mueller, W. A. J.Am. Chem. Soc. 1962,84. 297. ( 9 ) Kamlet, M. J.; Abboud, J. L.; Taft, R. W. J. Am. Chem. Soc. 1977, 99, 6027.
100
0
300
200
400
500
6 0
PRESSURE (bar)
Figure 1. Wavelength of absorbance maximum of 2-nitroanisole vs. pressure for supercritical NH3 (145 "C),C02(35 "C), Xe (26 "C), C2H, (35 "C), and SF6 (50 "C) at reduced temperatures of 1.02.
-
__ 0.4I
-3
Xe
I
-08 0 00
1
00
2 0
DENSITY (giml)
Figure 2. Kamlet and Taft r* values vs. solvent density obtained by using the 2-nitroanisole absorbance peak solvatochromic shift. Densities were calculated from the Peng-Robinson equation of state for the pure supercritical fluid solvent.
drogen-bond donation), and solvent basicity (hydrogen-bond acceptance), as well as other solvent properties:I0 u = uo SK* aa bp ...
+
+
+ +
The wavelength of solute absorbance (u) is dependent upon the solvent polarity/polarizability (T*), the specific solute (s), hydrogen-bond donating ability (a),hydrogen-bond accepting ability (p),and other specific solvent properties. On the basis of these solvent parameters, Kamlet and Taft have successfully predicted solute solubility, chromatographic behavior, and many widely varied solvent properties." The Kamlet and Taft solvent T * scale was developed by selecting a group of solute probes which have electronic absorption spectra sensitive to solvent polarity/polarizability but which would not participate in hydrogen bonding. The K* value of cyclohexane was arbitrarily set to 0.0 and that of dimethyl sulfoxide to 1.0, so that the more positive the K* value of a solvent the more "polar/polarizable" the solvent. On this scale the solute in the dilute vapor phase indicates a K* value nenr -1 .O, corresponding (IO) Kamlet, M. J.; Abboud, J. L.; Abraham, M. H.; Taft, R. W. J. Org. Chem. 1983, 48, 2877. ( 1 1) (a) Taft, R. W.; Abraham, M. H.; Doherty, R. M.; Kamlet, M. J. Nature (London) 1985, 313, (b) Sadek, P. C.; Carr, P. W.; Doherty, R. M.; Kamlet, M. J.; Taft, R. W.; Abraham, M. H. Anal. Chem. 1985, 57,2971.
The Journal of Physical Chemistry, Vol. 91, No. 11, 1987 3061
Solvent Properties of Supercritical Xe and SF6
TABLE I: Constant Values for the Two Density Regions (Low and High) Shown in Figure 3 and T* Values for Subcritical Liquids” subcritical liquid solvent
TR
NH3
1.02 1.01 1.03 1.01 1.01
co2 Xe
C2H6 F6
a
-0.8 -0.98 -1 .o -1 .o
b 0.98 1.15 0.74 0.63
‘See ea 1: a and c are the intercepts and 6 and d are the slopes for
(for
co2j.
d
C
-0.37 -0.43 -0.63 -0.63 ?r*
to the absence of any substantial solute-solvent interactions and a completely nonpolar and nonpolarizable “ ~ o l v e n t ” .The ~ ~ ~“s”~ value for 2-nitroanisole, the probe molecule selected for this work, was determined to be -2.426 with a vo of 32 560 cm-’.I3 Figure 1 illustrates the effect of changing pressure on the wavelength of the absorption maximum of 2-nitroanisole in supercritical NH3, Co2, C2H6, SF6, and Xe at reduced temperatures of 1.02. Extensive solubility measurements and related results (including limited solvatochromic meas~rements’,’~*’~) have been reported for supercritical carbon dioxide; consequently, its solvent properties have been better defined than other supercritical fluids. All of the fluids examined show a similar trend: as the pressure is increased, the probe absorbance shifts to longer wavelengths. It can be seen that the most rapid change in the peak maximum occurs below the critical pressure of the pure fluid. (The SF6data are limited at lower pressures due to the limited solubility of the solvatochromic probe.) The a* values are plotted as a function of fluid density in Figure 2 and show that a* can vary considerably between the fluids at the same density. A more useful comparison of the properties of the various supercritical fluids can be made on the basis of reduced densities (pip,, where p; is the critical density). Figure 3 gives plots of the a* values for the five fluids as a function of reduced density. It is obvious that at equal reduced densities the various fluids have quite different a* values, indicating that there are large differences in their effective polarities/polarizabilities. N H 3 has the largest a* values, which supports the expectation that it is the most polar solvent. The most negative values are obtained for supercritical SF6, with C 0 2 , C&, and Xe falling between these extremes. The large a* values for NH3, however, may be incorrect and actually reflect a specific hydrogen-bonding contribution to the peak shift. Further studies with additional solvatochromic probes will be required to establish such contributions. The data in Figure 3 illustrate intrinsically different effective solvent polarities for the five fluids. Ammonia is clearly the most polar studied (assuming no hydrogen bonding with the solute probe), and at the highest density studied it has a a* value approximately equivalent to that of methanol (0.60) on the Kamlet and Taft ~ c a 1 e . l ~The measured a* values for SF6 are quite negative, which implies that it is an extremely nonpolar solvent. In fact, at high densities its a* polarity/polarizability is similar to those measured for the liquid perfluoroalkanes.I6 This is not surprising in view of the symmetry of SF6and the high degree of ionic character of the S-F b0nds.l’ The data for Xe are interesting in that they show that it is a “better” solvent than SF6, with a* values slightly greater than C2H6. It should also be noted that the polarizability of Xe is substantially greater than either C 0 2 or NH3.’* However, the solvent properties of Xe can be attributed solely to dispersion forces and the polarizability of the (12) Essfar, M.; Guiheneuf, G.; Abboud, J. L. J. Am. Chem. SOC.1982, 104,6786. (13) Kamlet, M. J.; Abboud, J. L.; Taft, R. W. Prog. Phys. Org. Chem. 1981, 13, 485. (14) Hyatt, J. A. J. Org. Chem. 1984, 49, 5097. (15) Sigman, M.E.; Lindley, S.M.; Leffler, J. E. J. Am. Chem. Soc. 1985 107, 1471. (16) Brady, J. E.; Carr, P. W. Anal. Chem. 1982, 54, 1751. (17) Reed, A. E.;Weinhold, F. J. Am. Chem. SOC.1986, 108, 3586. (18) Landolt-Borstein Zahlenwerte und Functionen, 6 Auflage; Springer-Verlag: Berlin, 1950; 3 Teil, pp 510-517.
0.173 0.05 1 0.145 0.067
T, K
PR
296 297
2.41 2.11
296 298
2.42 2.60
T*
0.78 0.04 -0.22 -0.36
vs. reduced density relationships for the regions below and above pR = 0.7
0.5-
Y
0.0-
0.5
.”
-
.
1
0
2
REDUCED D E N W
Figure 3. Plot of
T*
vs. reduced density ( p / p c ) for the five fluids.
electron cloud. Surprisingly, SF6, which has no dipole or quadrupole moment, has a greater polarizability than Xe, but lower a* values. It is obvious that the effective solvent polarity/polarizability for a solute molecule is dependent on the fluid density, which is a function of both temperature and pressure. As the fluid density is increased, the a* values also increase for all of the solvents studied, consistent with empirical observations that it is possible to adjust the solvent properties of a supercritical fluid by changing the pressure or temperature.’” It is interesting to note that the rate of change with density is greater for the more polar solvents. It is also apparent that the rate of change of a* is not constant with density. At lower densities the a* dependence is greater while at reduced densities greater than 0.70 f 0.15 the change in solvent properties is more gradual. A more quantitative description of the change of solvent strength with density can be obtained by considering the slopes of the two regions at low and high reduced densities (pR). a* = b p R - a for a* < 0.7, a* = dpR - c for a* > 0.7 (1) Equation 1 describes both the low- and high-density region of the supercritical fluid spectra shown in Figure 3, where these constants for the slope ( b and d) and intercept ( a and c) of a* vs. pR for the supercritical fluids studied are given in Table I. The vapor-phase a* value for 2-nitroanisole of -0.99 was added to the data in the low-density region from which the constants a and b where regressed. Abboud et alet9report a a* value in the low-density region of -1.05 f 0.10; the values reported in Table I fall within the error limits of the respective studies. Carbon dioxide has been studied most extensively, and only a small fraction of the data is shown in Figure 3, all of which is consistent with presented data. The more positive value for N H 3 could be due to possible hydrogen-bond interactions noted earlier. N H 3 and C 0 2 show the largest rate in change in solvent strength with density at low densities. At high densities the change in a* with density ( d ) decreased by approximately a n order of magnitude. (19) Abboud, J.-L. M.; Guiheneuf, G.; Essfar, M.; Taft, R. W.; Kamlet, M. J. J. Phys. Chem. 1984.88, 4414.
3062
J. Phys. Chem. 1987, 91, 3062-3068
The change in s* vs. pR behavior is interesting since available data from solubility and chromatographic retention studies do not indicate any similar density dependence.20.21 Thus, while s* measurements appear consistent with relative fluid-phase solubilities, and effectively probe certain specific solute-solvent interactions, a quantitative relationship between s* values and supercritical fluid solubilities is not evident. The origin of these differences may be the fact that the s*value most strongly reflects the immediate solute environment rather than the effects of a larger solvation sphere. Data for the subcritical liquids of NH3,C o 2 , CzH6, and SF6 obtained at room temperature are also given in Table I. The x* values determined from the liquefied gas spectrum of C2H6and SF6for 2-nitroanisole correspond to the s* shown in Figure 3 for the supercritical fluid at similar values of reduced density. Therefore, the solvent power of a subcritical liquefied gas and a supercritical fluid are approximately equal at like densities. The present data also show that both supercritical SF6and Xe are solvents with properties comparable to other nonpolar solvents (e.g., ethane) with a very low effective polarity. Xe, which has
somewhat larger s* values than SF6, should prove to be a useful supercritical solvent for three reasoris. First, it has convenient critical properties (T,= 289.8 K, P, = 58.0 atm). Second, since it is a monatomic gas, it does not have specific solvent-solute interactions. Finally, for practical applications requiring detection or analysis, it provides an optimum spectroscopic window. The ability of liquid Xe to solubilize polymers22 and large hydroc a r b o n ~has ~ ~been attributed to its polarizability; however, the present s* measurements lead to the prediction that its solvent powers should be somewhat inferior to other supercritical fluids of common interest (e.g., C02). The relative s* values of Xe and sF6are surprising, however, and suggest the importance of specific interactions in these systems. Future studies will be aimed at understanding the specific solventsolute interactions that occur in supercritical solvents and the dependence of these interactions on the temperature and density.
Acknowledgment. The authors acknowledge the support of the
US.Department of Energy, Office of Basic Energy Science, under Contract DE-AC06-76RLO- 1830. (22) Everett, D. H.; Stagemen, J. F. J. Chem. SOC.,Faraday Trans. 2
(20) Yonker, C. R.; Gale, R. W.; Smith, R. D. J. Chromatogr. 1986,371, 83. (21) Gitterman, M.; Procaccio, I. J . Chem. Phys. 1983, 78, 2648.
1978, 230.
(23) Rentzepis, P. M.; Douglas, D. C. Nature (London) 1981, 293, 165.
Temperature of Maximum Density in Water at Negative Pressure Stephen J. Hendersont and Robin J. Speedy* Chemistry Department, Victoria University of Wellington, Wellington, New Zealand (Received: September 17, 1986)
The locus of the temperature of maximum density in stretched water is reported to pressures below -200 bar for H20, D,O, and an HDO mixture. The water samples were stretched in a fine helical capillary by the Berthelot tube principle. Pressure in the sample was measured by monitoring the unwinding of the helix, using the Bourdon tube principle.
Introduction A liquid which is under tension, or stretched, is metastable with respect to the liquid plus vapor system. Nucleation of the vapor phase results in the sudden loss of tension. Nucleation and growth of a vapor bubble is called cavitation and is accompanied by an audible click. The maximum negative pressure that a liquid can sustain without cavitating is its tensile strength. Experimental studies of water over the past three centuries' have yielded tensile strengths which vary by orders of magnitude, from about 1 to 277 bar. This range reflects both the difficulty of attaining tensions in liquids and of measuring them. Theoretical estimates of the tensile strength of water are considerably higher, varying from 2702to 6000 bar3 at 10 O C . Several reviews have been written on stretched Nucleation can occur homogeneously from spontaneous fluctuations in the local liquid structure which produces a viable bubble. However, the lack of reproducibility of nucleating tensions between different workers indicates that cavitation probably occurs at heterogeneous nuclei such as the solid surfaces in contact with the sample. The critical nucleating factors(s), though, have not been isolated. Indications are that cleanliness of both the water and solid surfaces in contact with it are very important. The importance of working with small samples and scrupulously clean conditions, which is well recognized in studies of superheated and Current address: Research School of Chemistry, Australian National University, Canberra, A.C.T., Australia.
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supercooled liquids, is often overlooked in work on stretched liquids and probably accounts for the wide range of reported nucleating tensions. Most previous work on stretched liquids has focussed on the problem of measuring the maximum tension that the sample will sustain before cavitation. Because of the difficulty of maintaining samples under tension, there have been few reports of the physical properties of stretched liquids. Water is the most important and most intensively studied liquid. The recent progress* which has been made toward mapping out its properties in the superheated and supercooled regions and the questions raised by its profoundly anomalous character highlight the need for information about its behavior at negative pressure. For example, water is anomalous in that it expands when cooled below 4 OC at atmospheric pressure. The temperature of maximum density (TMD) is the locus where the thermal expansivity of water is zero, Le., a = V ' ( I ~ V / I ~=T0.) , The TMD is a unique (1) Kell, G. S. Am. J . Phys. 1983, 51(11), 1038. (2) Kwak, H.; Panton, R. J . Phys. D . 1985, 18, 647. (3) Benson, S.; Gerjoy, E. J . Chem. Phys. 1949, 17, 914. (4) Hayward, A. T. J. Am. Sci. 1971, 59, 434. (5) Apfel, R.E. Sci. Am. 1972, 227(6), 58. (6) Trevena, D. H. Contemp. Phys. 1976, 17(2), 109. (7) Henderson, S. J.; Speedy, R. J., following paper in this issue.
(8) Angell, C. A. "Superheated and Supercooled Water" in Wuter and Steam: Their Properties and Current Industrial Applications: Pergamon: New York. 1980.
0 1987 American Chemical Society
