Environ. Sci. Technol. 1987, 2 1 , 925-928
Comment on “Aqueous Solubilities of Six Polychlorinated Biphenyl Congeners at Four Temperatures” SIR: In their contribution to this journal, Dickhut et al. (1) presented a set of important data about the relationship between experimental temperatures and aqueous solubilities of six polychlorinated biphenyls. Although the reported data may enable investigation of the full thermodynamics of the dissolution of some extremely hydrophobic chemicals, the authors have limited their discussion to the enthalpies of solution of the solid test compounds (AH,,). This is rather unfortunate, since it has often been suggested that entropy changes rather than enthalpy changes may dominate the thermodynamics of solution of such chemicals (2-4). With the presented solubility data of the six PCBs this can be investigated readily, since (5) -RT In x = R T In y + A S f ( T , - 5”) ACpT In ( T m / T ) - ACp(Tm- T ) (1)
+
and
R T In y = AH, - T A S , = AG,
(2)
x and y denote the solute’s mole fraction solubility and aqueous activity coefficient, AS, denotes the entropy of fusion a t the melting point, T is the experimental temperature, T , is the melting point, AC, is the differential heat capacity between the solid and super-cooled liquid phase of the solute, and AH,, AS,, and AG, are the enthalpy, entropy, and free energy of solution of the supercooled liquid solute. With the solubility data reported by Dickhut et al. ( 1 ) and Opperhuizen et al. (3),the AG,, AH,, and TAS, values of several PCBs are calculated for T = 298 K and listed in Table I. From the data it may be clear that for these PCBs the entropy contributions to the free energy of solution are more important than the enthalpy contributions. Perhaps the authors did not intend to discuss the full thermodynamics of solution but only wished to describe and predict the aqueous solubilities of the polychlorinated biphenyls a t various experimental temperatures. If so, however, it should be noted that they presuppose that the relationship between T and In x is linear (the van’t Hoff equation), with AHs, being the slope of this relationship. For solid solutes it has been argued that this assumption is debatable (3,5-7). This since AH,, is composed of the enthalpy changes associated with fusion of the solute’s molecules (AHf) and the enthalpy of solution of the (super-cooled liquid solute (AH,). For the former enthalpy change it has been argued that its value may vary with the experimental temperature, the variation being dependent
-Ln X 25
26 36
3.3
3.5
3.7
Figure 1. Relationship between In x and 1IT for 3,3’,4,4’-tetrachlorobiphenyl.
on the differential heat capacity of the solid and supercooled liquid state of the test compound (eq 1). Miller et al. (8) report that the differential heat capacity of PCBs is almost zero. Other data, however, do suggest that AH,, is dependent on the temperature, which may be explained by the differential heat capacity not being zero (3, 9, 10). The solubility data of the six PCBs reported by Dickhut et al. (1) support the latter suggestion, because there is no linearity between In x and 1/T for the PCBs. (See, for example, the plot of In X vs. 1 / T for 3,3’,4,4’-tetrachlorobiphenyl in Figure 1.) The observed increasing deviation from linearity with increasing experimental temperature ranges is consistent with the results we found for the aqueous solubility of penta- and hexachlorobenzenes (9). Finally, the authors argue that if one number for the enthalpy of dissolution has to be proposed that can be used as an estimate for the actual values for all PCBs 44.6 kJ/mol appears to be a more realistic approximation than 29.3 kJ/mol as was proposed previously (11). It is, however, debatable whether one of either values is realistic at all for solid solutes. This since, as was mentioned above, AH€contributes to AH,,, AHf being dependent on the melting point of the solid test compound. If it is assumed that AC, = 0 and thus AHf is a constant as is assumed by Dickhut et al., then the enthalpy of fusion at the experimental temperature is equal to that at the melting point. At the melting point AHf = T,ASf, enabling AHf to be calculated if AS, is known. It has often been argued that, according to the Walden rule, the entropy of fusion of rigid organic chemicals may be approximated as being -58 J mol-l K-l (12). If so, the enthalpies of fusion for different solid solutes must be proportional to their melting points. Since the melting points of the six PCBs for which the aqueous solubilities have been measured are significantly different, it may be
Table I. Enthalpy and Entropy Contributions to the Thermodynamics of Dissolution of Some Extremely Hydrophobic Chemicals in Water (All Values in kJ mol-’) compound biphenyl 4-chlorobiphenyl 3,3’,4,4’-tetrachlorobiphenyl 2,2’,4,5,5’-pentachlorobiphenyl 2,2’,3,3’,6,6’-hexachlorobiphenyl 2,2’,3,3’,5,5’,6,6’-octachlorobiphenyl 2,2’,3,3’,4,4’,5,5’,6-nonachlorobiphenyI decachlorobiphenyl 2,2’,4,4’-tetrachlorobiphenyl 2,2’,4,4’,5,5’-hexachlorobiphenyl pentachlorobenzene hexachlorobenzene 0013-936X/87/0921-0925$01.50/0
Tm
AG,
344 350 453 350 385 435 479 578 316 386 358 501
31.8 39.8 50.6 48.7 49.9 56.0 57.9 61.4 47.3 58.1 44.7 54.8
0 1987 American Chemical Society
33.0 28.5 50.7 31.9 45.6 50.7 49.8 66.6 32.0 30.9 31.3 36.9
A“,
f i e
20.6 20.9 27.1 20.9 23.0 26.0 28.7 34.6 19.0 20.9 16.1 18.7
12.4 7.6 23.6 11.0
22.6 24.7 21.1
32.0 13.0 10.0 15.2 18.2
-TAS,
ref
19.3 32.2 27.0 37.7 27.3 31.3 37.8 29.4 34.4 48.1 28.6 36.6
1 1 1
1 1 1 1 1
3 3 9 9
Environ. Sci. Technol., Vol. 21, No. 9, 1987
925
clear that there is no basis to assume that their AHH,values can be approximated by one value. However, if it is assumed that AC, = 0, subtraction of AHf from the reported AH,,values gives AH,values, which are listed in Table I. As is shown, these enthalpies of solution of super-cooled liquid solutes show less variation than the AH,, values for the six PCBs investigated by Dickhut et al. (1). In addition, the AHs values of two PCBs investigated by Opperhuizen et al. (3),as well as with those for penta- and hexachlorobenzene, are comparable. These data suggest that the AHs values of hydrophobic chemicals vary little with chemical structure. While AHf values of various solid hydrophobic chemicals vary with the melting point, AHBs values will be different for the different chemicals. As a consequence, variations of the aqueous solubility with temperature are not similar for different chemicals. Hence, it can be concluded that the dissolution of extremely hydrophobic chemicals is associated with a relatively constant endothermic enthalpy of solution (from the super-cooled liquid state) and an endothermic enthalpy of fusion that is proportional to the solute’s melting point. More important, however, is that the presented data show clearly that unfavorable entropy changes dominate dissolution processes of extremely hydrophobic chemicals. Registry No. Water, 7732-18-5.
Literature Cited (1) Dickhut, R. M.; Andren, A. A,; Armstrong, D. E. Environ. Sci. Technol. 1986, 20, 807-810. (2) Chothia, C.; Janin, J. Nature (London) 1975,256,705-708. (3) Opperhuizen, A.; Gobas, F. A. P. C.; Van der Steen, J. M. D.; Hutzinger, 0. Environ. Sci. Technol., in press. (4) Tomlinson, E.; Davis, S. S. J. Colloid Interface Sci. 1980, 76, 563-572. ( 5 ) Hildebrandt, J. H.; Prausnitz, J. M.; Scott, R. L. Regular and Related Solutions;Van Nostrand Reinhold New York, 1970. (6) Hollenbeck, G. E. J . Pharm. Sci. 1980, 69, 1241-1242. (7) Riebesehl, W. Ph.D. Thesis, University of Amsterdam, The Netherlands, 1984. (8) Miller, M. M.; Ghodbane, S.; Wasik, S. P.; Tewari, Y. B.; Martire, D. E, J . Chem. Eng. Data 1984, 20, 184-190. (9) Benecke, J. I.; M.Sc. Thesis, University of Amsterdam, The Netherlands, 1986. (10) Hafkenscheid, T. L.; Tomlinson, E. Znt. J. Pharm. 1981, 8, 331-335. (11) Mackay, D.; Mascarenhas, R.; Shiu, W. Y.; Valvani, S. C.; Yalkowsky, S. H. Chemosphere 1980, 9, 257-264. (12) Yalkowsky, S. H. Ind. Eng. Chem. Fundam. 1979, 18, 108-1 11.
Antoon Opperhuizen,” Jeanet I. Benecke John R. Parsons
Laboratory of Environmental and Toxicological Chemistry University of Amsterdam 1018 WV Amsterdam, T h e Netherlands
SIR: In their correspondence Opperhuizen et al. ( I ) raise the following issues regarding our article on the effect of temperature on the aqueous solubility of six PCB congeners (2). First, they suggest that we should not have limited our discussion of the dissolution thermodynamics to the enthalpies of solution for the solid test compounds. Second, they objected to our assumption of a linear relationship for a plot of In x vs. 1/T resulting in a constant 926
Environ. Sci. Technol., Vol. 21, No. 9, 1987
-18 -18
-20
x
I
MCB
..
-22
’.
-24
‘’
-28
’’
TCB
C
I
.982
r2
-32
‘
2. 8
= ,958
.
I
3. 1
3.4
3. 7
T-’ X IO3(OK-l) Flgure 1. van’t Hoff plots for selected PCB congeners (B = biphenyl, MCB = 4-chlorobipheny1, PCB = 2,2’,4,5,5’-pentachlorobiphenyl, HCB = 2,2’,3,3‘,6,6’-hexachlorobiphenyl, OCB = 2,2’,3,3’,5,5‘,6,6’-octachlorobiphenyl, NCB = 2,2’,3,3’,4,4’,5,5‘,6-nonachlorobiphenyl, DCB = decachlorobiphenyl).
AHs,. Last, they disagree with our suggestion of an alternative average value for the enthalpy of dissolution for PCBs. The intention of our study was to describe solubility variation with temperature for PCBs so that predictions of the effect of this variable on partitioning and transport processes (3, 4) could be validated. Hence, most of our data (2, 5 , 6) were measured in the range of 0-40 “C of environmental interest. We agree that an investigation of the full thermodynamics of dissolution for the PCBs will aid in understanding the nature of hydrophobic interactions, but reiterate that this was not the original purpose of our paper. To perform a rigorous analysis requires more information than was obtained in our study (Le,, vapor pressure and fusion data). The linearity of the van’t Hoff relationship for PCBs is evident from the data. Figure 1shows reciprocal plots for the PCB solubility-temperature data of ref 2 and 6-8. Correlation coefficients for linear fits to these data range from 0.958 to 1,000 (Figure 1). In these investigations and others (9, IO),where the temperature range is restricted to the environmental region, good linearity of In x vs. 1/T is observed. Presupposition that the relationship between T and In x is linear (the van’t Hoff equation) is tantamount to assuming that AH,, is constant with temperature within the region of interest, Given the linearity of the plots in Figure 1 and those of others (5, 9), we do not feel that the assumption is inappropriate when made for a limited temperature range. However, the preceding correspondence (1) cites that there is an “observed increasing deviation from linearity with increasing experimental temperature ranges” ( I 1). We do not dispute this. The PCB data do show increased curvature, although not greatly significant, when the experimental temperature range is large. This is evidenced by the biphenyl curve, temperature range 0.4-64.5 “C, and
