Anal. Chem. 1995, 67, 4354-4357
Evidence for Density-DependentChanges in Solute Molar Absorptivities in Supercritical COz: Impact on Solubility Determination Practices Jeanette K. Rice,+Emily D. Niemeyer, and Frank V. Bright* Department of Chemistry, Natural Sciences and Mathematics Complex, State University of New York at Buffalo, Buffalo, New York 74260-3000
Optimization of supercritical fluid extraction conditions requires accurate information on solute solubility under a given set of experimental conditions. In situ electronic absorbance spectroscopy is used commonly to determine solute solubilities, and one assumes that the solute molar absorptivity (E) is constant over a broad density range when monitoring at a particular wavelength. Using Wvis spectroscopy,we have found that c for anthracene and pyrene in supercritical C 0 2 is density dependent. Systematic increases in E of 1.3-2.7-fold (30-170%) are observed as C 0 2 density increases from 0.3 to 0.9 &an3. We account for the observed density dependence in terms of solute-solvent dielectric interactions. These results illustrate the pitfalls associated with using in situ spectroscopic techniques for the determination of solute solubilities in supercritical fluids. Many properties of supercritical fluids [e.g., density (p) and refractive index (n)] can be tuned continuously over a wide range by modest changes in system pressure or temperature. However, although such tunability makes supercritical fluids attractive solvents with wide applicability, it also introduces potential problems that are frequently overlooked. For example, supercritical fluid extraction (SFE), the selective separation of compounds with supercritical fluids based on their differential solubility, has developed into a powerful analytical tool with many practical appli~ations.l-~ However, optimization of analytical SFE requires that one know the actual analytical concentration of solute/extractant in the fluid phase at a given temperature and pressure. Moreover, the development of improved phase equilibria models also requires accurate information on solute solubility as a function of fluid density and/or fluid composition (if the fluid is cosolvent modified). The literature is ripe with reports on the determination of solute solubilities in neat and cosolvent-modified supercritical Current address: Department of Chemistry, Georgia Southern University, Landrum Box 8064, Statesboro, GA 30460. (1) Mch'ally, M. E. P. Anal. Chem. 1995,67, 308A. (2) McHugh. M.; Krukonis, V. J. Supercritical Fluid Extraction, 2nd ed.; Buttenvorths: Boston, 1994. (3) Ghonasgi. D.; Gupta, S.; Dooley, K. M.; Knopf. F. C. AIChE J. 1991,37, 944. (4) Larson, K. A; King, M. L. Biotechnol. Prog. 1986,2,73. (5) Williams, D. F. Chem. Eng. Sci. 1981,36, 1769. (6) Mulcahey, L. J.; Taylor, L. T. Anal. Chem. 1992,64, 981. (7) Ray, M. S. Sep. Sci. Technol. 1994,29,2203.
4354
Analytical Chemistry, Vol. 67, No. 23, December 7, 7995
fluids,s-20 and measurements based on in situ W-vis and IR absorbance spectroscopy are common.8~9~13~18~19 However, in order for an in situ absorbance-based scheme to yield accurate solubility data, there is the underlying assumption that the molar absorptivity, E, of a particular solute is constant with fluid density and fluid composition. Unfortunately, theory predicts (vide infra) that E is a function of solvent refractive index. Thus, because the refractive index of a supercritical fluid changes with fluid density, it is not unreasonable to expect that E is not constant with density in a supercritical fluid. Recently, IR absorbance studies have shown that E for a fundamental vibrational transition within a solute is not constant with fluid density. Specifically, Inomata et aLZ1studied naphthalene in supercritical COZ. By monitoring the intensity of the C-C (quadrant) vibrational mode at about 1600 cm-', these authors determined that E increases by a factor of 1.6 on changing the COS density from 0.3 to 0.9 g/cm3. Additional work by Franck and RothZ2showed that the molar absorptivity associated with the 0-D stretching mode of dilute HOD in water varied nearly 2 orders of magnitude over a density range of 0.02-0.90 g/cm3. Similar trends have also been reported for electronic absorbance transitions. For example, Kimura and YoshimuraZ3have used UV-vis absorbance to study the dimerization equilibrium of 2-methyl-2nitrosopropane in COS, CF3H, and CClF3. In the course of this work, the authors found that the molar absorptivity of the monomer increased systematically by a factor of 1.4 with increasing fluid density. (8) Cygnarowicz, M. L.; Maxwell, R J.; Seider, W. D. Fluid Phase Equilib. 1990, 59,57. (9) Zerda, T.W.; Wiegrand, B.; Jonas, J. Chem. Eng. Data 1980,3I,274. (10) Nakatani, T.; Kazunari, 0.;Katayama, T. Ind. Eng. Chem. Res. 1991,30, 1362. (11) Lemert, R M.; Johnston, K P. Fluid Phase Equilib. 1990,59,31. (12) Bartle, K. D.; Clifford, A. A: Jafar, S.A,]. Chem. Eng. Data 1990,35, 355. (13) Kajimoto, 0.; Futakami, M.; Kobayashi, T.; Yamasaki, K. J. Phys. Chem. 1988.92,1347. (14) Smith, R. D.; Udseth. H. R.; Wright, R. W.; Yonker, C. R. Sep. Sci. Technol. 1987,22,1065. (15) Johnston, K. P.; Ziger, D. H.; Eckert, C. A Znd. Eng. Chem. Fundam. 1982, 21,191. (16) Dobbs, J. M.; Johnston, K. P. Ind. Eng. Chem. Res. 1987,26,1476. (17) Dobbs, J. M.; Wong, J. M.; Johnston, K. P. J. Chem. Eng. Data 1986,31, 303. (18) Ebeling, H.; Franck. E. U. Bey. Bunsen-Ges. Phys. Chem. 1984,88,862. (19) Rossling, G. L.; Franck, E. U. Ber. Bunsen-Ges. Phys. Chem. 1983,87, 882. (20) Kwiatkowski, J.; Lisicki, Z.; Majewski, W. Ber. Bunsen-Ges. Phys. Chem. 1984,88,865. (21) Inomata, H.; Yago, Y.; Saito, M.; Saito, S.J. Supercn't. Fluids 1993,6,237. (22) Franck, E. U.; Roth, K. Discuss. Faraday SOC.1967,43, 108. (23) Kimura, Y.; Yoshimura, Y. J. Chem. Phys. 1992,96,3085.
0003-2700/95/0367-4354$9.00/0 0 1995 American Chemical Society
there is the potential for concomitant changes in the position of Together these data suggest that the of the solute's absorbance spectrum and the solute molar absorp acquiring or using solute molar absorptivities measured in liquid tivity. solvents as calibration points for in situ solubility determinations in supercritical fluids may yield grossly inaccurate solubilities. This EXPERIMENTAL SECTION problem is exacerbated further because it is well-known that Instrumentation. All absorbance measurements were redensity-dependent spectral shifts are frequently observed for corded on a Spectronic 1201 UV-vis spectrophotometer (Milton solutes dissolved in supercritical f l ~ i d s . ~Unfortunately, ~-~~ most Roy Co., Rochester, NY), which has been modified to accomin situ absorbance measurements are performed at a single, k e d modate the high-pressure optical cells. The cells were constructed wavelength and generally do not take spectral shifts into account in-house and have been described in detail e l ~ e w h e r e . ~Pres~,~~ for solubility determinations. sure was generated using a syringe pump (Isco, Model 260D, In this work, we report on the density dependence of the Lincoln, NE) and was monitored by an in-line analog Heise Gauge anthracene and pyrene molar absorptivity and absorbance spectra (flpsi). Temperature control was achieved using a temperature in supercritical COZ. These particular solutes were chosen circulating bath (Haake, Model A80) and monitored by a solidbecause they have been well-characterized spectroscopically.30-36 state thermometer (Cole-Parmer, Model 8517-00) to within *O.l They are nonpolar, so strong, specific solute-fluid interactions "C. (Note: Supercritical fluid work requires the use of high (e.g., hydrogen bonding) are minimal, and solubility data have pressures. Appropriate precautions should be taken to minimize been reported for pyrene37-39and anthracene.15~16~19~20~37~40~41 Thus, danger.) the magnitude of the bias reported here represents a best case Sample Preparation. Pyrene (99+%, Aldrich), anthracene scenario as solute-fluid interactions are relatively weak in (99%, Aldrich) , carbon dioxide (SFC Grade, Scott Specialty Gases), comparison to systems that can exhibit, for example, hydrogen n-hexane (99%,Aldrich), and ethanol (Pharmco, 200 Proof) were bonding. Carbon dioxide was chosen because it is the most used as received. Samples were prepared in the following manner. commonly used supercritical fluid. Stock solutions of pyrene and anthracene in ethanol (W3M)were prepared and stored in amber bottles, and an aliquot was pipetted THEORY directly into the high-pressure optical ~ e 1 1(internal ~ ~ , ~ volume ~ = It is well-known that interactions between a dissolved solute 5.0 mL) such that the final concentration in the cell would be I, and the solvent are responsible for shifts in a solute absorbance 10, 20, or 50 pM. The liquid solvent was allowed to evaporate; and emission spectra.42 For the current systems, spectral shifts the cell was connected to the high-pressure pump, briefly arise because of dispersion forces between the solute and solvent evacuated, and then heated to 45.0 "C. Once the cell reached (polarizability driven) and the transition dipole moment of the thermal equilibrium, it was charged with COe to the starting solute that drives induced interactions between the solutes and pressure and sample stirring was initiated. Pressure was increthe fluid. In both cases, the electronic transition energy is known mented from low to high throughout the experiment, and we to depend on the solvent refractive The frequency of this typically allowed at least 15 min for equilibration between pressure electronic transition is also related to the solute oscillator strength, changes and spectral acquisition. The density range investigated its transition dipole moment, ground, and excited-state wavefuncwas from pr = 0.714-1.91 (pr = p,/p,; pr is reduced fluid density; tions, and ultimately (for this work) the solute's molar absorption pe is experimental density; pc is critical density). Carbon dioxide ~oefficient.~~ Thus, if the refractive index of the solvent changes, densities were estimated using commercially available software (SFSolver, Isco, Inc., Lincoln, NE). The absorbance center of (24) Yonker, C. R.; Frye, S. L.; Kalkwarf, D. R; Smith, R D.J Phys. Chem. 1986, 90,3022. gravity was recovered using computer software developed in(25) Sun, Y.-P.; Fox, M. A;Johnston, K. P.J Am. Chem. SOC.1992,114, 1187. house that is based on the protocol outlined in ref 46. (26) Kim, S.; Johnston, K P. AIChE J. 1987,33,1603. Calculation of Molar Absorptivity. Solute molar absorptivi(27) Kim, S.; Johnston, IC P. Znd. Eng. Chem. Res. 1987,26,1206. (28) Rice, J. K.: Niemeyer, E. D.: Dunbar, R. A,; Bright, F. V.J. Am. Chem. SOC. ties were determined from a series of density-dependent working 1995,117, 5832. curves of absorbance vs solute c~ncentration.~'In all cases, the (29) Rice, J. IC; Niemeyer, E. D.; Bright, F. V. J Phys. Chem., submitted for dissolved solute concentration in the high-pressure optical cell publication. (30) Winnik, F. Chem. Rev. 1993,93,587 and references cited therein. was well below the know solubility limits at the lowest fluid (31) Weinstein, Y. A; Sadovskii, N. A; Kuz'min. M. G. High Energy Chem. 1994, d e n ~ i t y . ~Thus, ~ , ~ ~at all fluid densities, at our highest solute 28,211. concentration, all solute loaded into the high-pressure cell is (32) Yorozu, T.; Hoshino, M.; Imamura, M. J Phys. Chem. 1982,86, 4426. (33) Bauer, R. K; de Mayo, P.: Ware, W. R.; Wu, K. C. J Phys. Chem. 1982,86, completely dissolved. Hence, we know the analytical concentra3781. tion of dissolved solute under all conditions. In the worst case, (34) de Mayo, P.: Natarajan. J. V.: Ware, W. R. Organic Phototransitions in Nonthe correlation coefficient for our linear regression was 0.997 and Homogeneous Media; ACS Symposium Series 278; Fox, M. A, Ed.; American Chemical Society: Washington, DC, 1985. the intercept was 0.002. (35) Amirav, A. Chem. Phys. 1988,124,163. (36) Hirayama, S.; Iuchi, Y.; Tanaka, F.; Shobatake, IC Chem. Phys. 1990,144, 401. (37) Van Alsten, J. G.: Eckert, C. A J . Chem. Eng. Data 1993,38, 605. 355. (38) Bartle. K D.; Clifford, A A.; Jafar, S. A.J. Chem. Eng. Data 1990,35, (39) Yu, E.; Richter, M.; Chen, P.; Wang, X.: Zhang, Z.; Tavlarides, L. L. Ind. Eng. Chem. Res. 1995,34,340. (40) Zerda, T. W.: Wiegand, B.; Jonas, J. J. Chem. Eng. Data 1986,31,274. (41) Murray, J. S.; Lane, P.; Brinck, T.; Politzer, P. J. Phys. Chem. 1993,97, 5144. (42) Suppan, P. J. Photochem. Photobiol. A 1990,50, 293. (43) Atkins, P. W. Physical Chemistry, 4th ed.: W. C. Freeman & Co.: New York, 1990; Chapter 17, pp 503-504.
RESULTS AND DISCUSSION The absorbance spectra of pyrene and anthracene in supercritical COz red-shift with increasing fluid density (Figure 1).As (44) Betts, T. A.; Bright, F. V. Appl. Spectrosc. 1990,44,1196. (45) Zagrobelney, J. Ph.D. Dissertation, State University of New York at Buffalo, 1992. (46) Lakowicz. J. R; Hogan, D. Biochemistty 1981,20, 1366. (47) Willard, H. H.; Memt, L. L., Jr.; Dean, J. A,: Settle, F. A,, Jr. Instrumental Methods ofAnalysis, 7th ed.; Wadsworth: Belmont. CA, 1988; Chapter 7, pp 159-196.
Analytical Chemistty, Vol. 67, No. 23, December 1, 1995
4355
30700
-
AI
-
14000 12000
+
A
10000
E 8000 : v
6000
1 .o
0.5
1.5
2.0
27400
B
27350
0.50
1 .o
0.5
1.5
Figure 1. Absorbance center of gravity as a function of reduced fluid density for the lowest energy transition of pyrene (A) and anthracene (B) in supercritical CO2. Solute analytical concentration is 20 pM. T = 45.0 "C. 75000
i 4s000i 1 7
E
A
60000
-
:
v
JOOOO
1 15000 - 4 2.5
1.25
1.50
1.75
2.00
Figure 3. Molar absorptivity of anthracene in supercritical CO2 as a function of reduced fluid density using peak maxima (0)and using 2. = 370 nm (0) (A). Ratio of €(peak max) to E(>. = 370 nm) as a function of reduced fluid density (B). T = 45.0 "C. Table 1. Molar Absorptivity of Pyrene in Supercritical C02 at Absorbance Peak Maxima for Lowest and Highest Reduced Fluid Densities
peak m a U 311 326
at pr = 0.714
21 097 f 85 33 400 i. 565
peak maxu
cb.cat pr = 1.91
314 329
34 070 f 599 43 590 i 151
In nanometers. * In L/mol cm. Six points in the calibration.
B
:i
Table 2. Molar Absorptivity of Anthracene in Supercritical C02 at Absorbance Peak Maxima for Lowest and Highest Reduced Fluid Densities
1 .o
1-
0.5 0.50
1.00
Reduced Density
2.0
Reduced Density
0.75
0.75
1.00
1.25
1.50
1.75
2.00
Reduced Density
Figure 2. Molar absorptivity of pyrene in supercritical CO2 as a function of reduced fluid density using actual peak maxima (0)and using 2. = 329 nm (0)(A). Ratio of c(peak max) to €(I. = 329 nm) as a function of reduced fluid density (6).T = 45.0 "C.
mentioned previously, these shifts will (even if the molar absorp tivity at a given absorbance band were constant) bias any in situ solubility determinations based on spectroscopic measurements made at a single, fixed wavelength. One can easily see this effect by determining the molar absorptivity of each solute at a fixed wavelength and comparing it to the molar absorptivity determined at the actual peak maximum (which accounts for the spectral shift). Figures 2A and 3A present the molar absorptivities of pyrene and anthracene, respectively, using a single wavelength (pyrene, 329 nm; anthracene, 370 nm) (open symbols) and using the peak maximum for the lowest energy transition (solid symbols). (Note: The single wavelength used in Figures 2 and 3 is the wavelength where the peak maximum of the lowest energy absorbance transition occurs at the highest fluid density. Thus, the recovered absorptivities using the two methods approach the same value at the highest fluid densities.) The magnitude of the bias is clearly illustrated in Figures 2B and 3B where we present the ratio of the molar absorptivity at a fixed wavelength to the absorptivity of the lowest energy absorbance peak vs COZdensity. 4356 Analytical Chemistry, Vol. 67, No. 23, December 1, 1995
peak max'
cbscat pr = 0.714
peak m a U
317 332 349 367
3223 f 147 6433 f 490 9930 5 340 9187 i. 168
320 335 351 370
at pr = 1.91
4 773 i 31 9 100 f 113 14 150 f 181 13 793 i. 137
In nanometers. * In L/mol cm. S i points in the calibration.
At the lowest fluid density, when monitoring at a single wavelength, one sees that the actual peak absorptivity is greater by a factor of 2.0 (pyrene) and 1.8 (anthracene) compared to the value measured using a tixed Wavelength. Furthermore, over the density range investigated, the magnitude of 6 (recovered from peak maxima data) increases by a factor of 1.3 for pyrene and anthracene. However, E increases by a factor of 2.7 (pyrene) and 2.2 (anthracene) based on single wavelength values. This illustrates experimentally that (1) the solute molar absorptivity is indeed density dependent and (2) an additional bias can result from shifts in the actual absorbance spectra with density. These results clearly demonstrate the difficulties associated with recovering accurate solute solubility data using in situ absorbance measurements made at a single wavelength. Tables 1 and 2 collect the molar absorptivity at the peak absorbance maxima for pyrene and anthracene at pr = 0.714 and 1.91 (i.e., at the lowest and highest reduced densities studied). The actual wavelengths corresponding to the peak maxima are also reported. Changes in the solute molar absorptivity of 1.3-
Table 3. Molar Absorptivity of Pyrene in Liquid Ethanol and -Hexane at Absorbance Peak Maxima
peak m a U
ch ethanol
318 334
26 433 595 38 960 & 711
*
peak m a U
cb,
