Thermodynamics of Transfer
The Journal of Physical Chemistry, Vol. 82, No. 6, 1978 675
The Thermodynamics of Transfer of Phenol and Aniline between Nonpolar and Aqueous Environmentst Kenneth J. Breslauer,' Lucia Witkowski, and Krlstina Bulas Department of Chemistry, Douglass College, Rutgers, The State University, New Brunswick, New Jersey 08903 (Received November 9, 1977) Publication costs assisted The Rutgers Research Council
A recently developed calorimetric technique is employed to directly determine the enthalpy change accompanying the transfer of phenol and aniline from nonpolar to aqueous environments. In conjunction with equilibrium studies, complete thermodynamic profiles are obtained for the transfer of phenol from octane, toluene, and octanol to water and for the transfer of aniline from toluene to water. The data are interpreted in terms of solute-organic solvent and solute-water interactions. These results are discussed in the light of currently accepted views concerning the nature of hydrophobic and hydrophilic forces.
Introduction Several years ago we described the development of a new calorimetric technique that allows the direct determination of the enthalpy change accompanying the transfer of a molecule from a nonpolar to an aqueous environment.' Our work in this area was stimulated by the fact that many investigators were reporting thermochemical data on complex biochemical systems that they were unable to explain in terms of currently recognized molecular interactions. At the same time, several laboratories were suggesting that important contributions to the energetics of many biochemical processes arose from the transfer of groups from nonpolar to aqueous environments, or the reverse.24 For example, in the denaturation of a globular protein those amino acid side chains originally embedded in the relatively nonpolar interior of the macromolecule, upon unfolding, become exposed to the highly polar, aqueous environment of the solvent. In contrast, the binding of an inhibitor or substrate to a macromolecule was frequently envisioned as involving a change in the environment of the binding species from a purely polar, aqueous environment to the relatively nonpolar surroundings of the binding sites5 In general, it was becoming clear that if one was to explain the molecular origins of the thermochemical data obtained on complex biochemical systems, one had to be able to define the thermodynamic contribution arising from such polar-nonpolar environmental changes. Unfortunately, these data are quite difficult to obtain by studying the biological systems themselves since these environmental changes invariably occur concurrently with charge-charge interactions as well as conformational changes. In order to isolate the thermodynamic changes associated purely with the transfer of a macromolecular component from one environment to another, one can study the thermodynamics accompanying environmental changes for small molecules that can serve as models for the individual troups that make up these biopolymers. In this connection, the free energy change associated with the transfer of a small molecule from a nonpolar to an aqueous medium would be of interest. Such determinations can and have been carried out by studying the equilibrium distribution of various compounds between water and some 'This work was supported by grants from the Rutgers Research Council, the Charles and Johanna Busch Memorial Fund,and the Research Corporation. 0022-3654/78/2082-0675$01 .OO/O
relatively nonpolar organic solvent.6 Alternatively, some workers have used solubility measurements to calculate partition coefficient^.^ The results of such experiments have been used to construct free energy tables (hydrophobicity scales) which reflect the relative affinity of various compounds and molecular groups for the organic and the aqueous phases. However, to begin to understand on a microscopic level what constitutes a hydrophobic or hydrophilic interaction as well as to explain much of the accumulated thermochemical data, one must dissect the free energy term into its constituent parts. That is to say, one must expand the model to include the determination of values for the enthalpy, entropy, and heat capacity changes accompanying these distribution experiments. Unfortunately, very few data exist for the enthalpy change accompanying the transfer of a molecule from a nonpolar to an aqueous environment. This probably has been due to the fact that no accurate and convenient method for measuring this parameter has in the past been available. In an attempt to alleviate this situation, we initiated a program that has resulted in the development of a direct and accurate calorimetric technique for measuring the enthalpy change accompanying the transfer of a molecule from a nonpolar to an aqueous medium. In this paper we present the results obtained by application of this technique to several molecules of general and biological interest.
Experimental Section Chemicals. Phenol and Aniline were obtained from Fisher Scientific Co., Fair Lawn, N.J. and were used without further purification. Chromatoquality toluene and octane were purchased from Matheson Coleman and Bell. Grade 1 99% pure octanol was obtained from Sigma Chemical Co., St. Louis, Mo. Method The Calorimeter. The experimental technique makes use of an extremely sensitive flow microcalorimeter developed by Sturtevant in collaboration with Beckman Instruments, Palo Alto, Calif.' The main components of the instrument are a precision fluid delivery system and a 10000 junction thermopile which is enclosed within a massive aluminum heat sink. The fluid delivery system consists of two glass syringes equipped with gas-tight Teflon-tipped plungers which are independently driven by 0 1978 American Chemlcal Soclety
676
The Journal of Physical Chemistry, Vol. 82, No.
TABLE I: Enthalpy of Transfer of Phenol and Aniline from Toluene to Water at 25 Cav
Scheme I: Determination of Enthalpies of Transfer R-OH(org) + OH-(aq) R-0-(aq) t H,O R-0-(aq) t H+(aq) R-OH(aq) H,O H+(aq) t OH-(aq)
AH, AH, AH,
R-OH( org)
A~trans
-+
-+
-+
R-OH(aq) AH-,= AH,
-+
-+
-
AH, t AH,
Scheme 11: Determination of Enthalpies of Transfer R-yH,(org) t H+(aq) R-hH,(aq) R-NH,(aq)-+ R-NH,(aq) + H+(aq)
AH, AH,'
R-NH,( org)
AHtmlla
-+
R-NH,(aq) A H b , = AH,
-f
K. J. Breslauer, L. Witkowskl, and K. Bulas
6, 1978
+ AH,'
two variable speed 10 rpm synchronous motors. The two fluids are separately delivered to the calorimeter system by passage through Teflon tubes. At the entrance to the heat sink, the Teflon tubing is changed to platinum tubing (1.0 mm id.) which is in good thermal contact with the heat sink. This ensures proper equilibration of the liquids prior to their entering the thermopile. Upon reaching the thermopile the two platinum tubes are brought together by means of a Y junction which serves to initiate the mixing process. The mixed liquids then pass through the platinum tubing which is pressed tightly against the inside surface of a 10 000 junction thermopile. Thus, any heat evolved or absorbed upon mixing the two solutions is quantitatively conducted through the thermopile to the massive aluminum heat sink. The resulting electrical output (which is proportional to the heat evolved or absorbed) is amplified and recorded. Enthalpies of Transfer. The technique employed involves the extraction of a compound of interest from an organic to an aqueous phase in the flow calorimeter described above. To avoid excessive heats of solvent mixing it is essential to use immiscible phases. The extraction is readily accomplished by flowing a dilute alkaline (or acidic) aqueous solution against an organic solution of the compound to be transferred. Obviously, such a compound must possess a site for facile protonation or deprotonation so that either an acidic or alkaline aqueous solution can be used to ensure complete extraction during the calorimetric residence time, Previous work has shown that the extraction is in fact complete before the solution exits from the ca1orimeter.l The overall extraction process results in an enthalpy change which not only includes the heat associated with the transfer of the molecule between the two phases, but also the heat of ionization (or protonation) of the compound transferred. In addition, one must correct for the heat of formation of water whenever an alkaline extracting solution is used, This latter enthalpy is well known and thus can be subtracted from the overall heat of the process. On the other hand, a separate experiment must be performed in order to determine the heat of protonation (or deprotonation) for each compound transferred. For the case of a molecule possessing an "active" hydrogen, the overall transfer process can be summarized as outlined in Scheme I. On the other hand, when dealing with compounds such as amines that are transferred (extracted) by means of aqueous acid, the overall set of reactions reduces to those shown in Scheme 11. Thus, simply by adding the three processes illus$rated in Scheme I or the two processes illustrated in Scheme I1 one obtains a value for the enthalpy change associated with transferring the uncharged molecule from a nonpolar to an aqueous environment.
Compd transferred Phenol Aniline
AH,, kcal mol-' -9.19 -8.48
AH2, kcal mol-'
AH,, kcal mol-'
-5.65 -7.28
+13.34
AH-,, kcal mol-' -1.50i: 0.13 -1.20 i 0.15
a See Schemes I and I1 for meaning of symbols. All enthalpies are averages of at least four determinations. The mean calculated enthalpies of transfer are listed in the last column along with the standard errors of the mean.
Scheme 111: Enthalpy of Transfer of Phenol (PhOH) between Two Organic Solvents PhOH(octane) -+ PhOH(H,O) PhOH(H,O) -+ PhOH(to1uene) PhOH(octane)
-+
PhOH(to1uene)
A H = - 3.9 kcal mol-' A H = t l . 6 kcal mol-' A H = - 2.4 kcal mol-'
Results and Discussion Table I summarizes the enthalpy data obtained for the transfer of phenol and aniline from toluene to water. As described below, these data can be used to test the new calorimetric technique employed in this work. The enthalpy of transfer of phenol from toluene to water is exothermic by 1.5 kcal mol-1 (note: 1 calTh = 4.184 J). This value should be compared with the corresponding enthalpy of transfer that can be derived from previously published heats of solution. Parsons and co-workers found the heat of solution (25 " C ) of solid phenol in water to be +3.04 kcal mol-1.8 Arnett and co-workers determined the heat of solution (at 25 "C) of solid phenol in toluene to be +4.51 kcal m01-l.~From these data, one can derive a value for the enthalpy of transfer of phenol from toluene to water as shown in I. phenol (solid)
-AH, = - 3.04 kcal mol-'
AH^ = t 4.51 kcal mol-'
/
phenol (H,O)< phenol (toluene) AH, = - 1.47 kcal mol-'
(1)
The derived value of -1.47 kcal mol-l for the toluene to water transfer is in very good agreement with the 1.5 kcal mol-l determined by means of our direct calorimetric technique. This exothermicity probably reflects favorable hydrogen bonding between the phenolic hydroxyl group and water. The enthalpy of transfer of aniline from toluene to water was found to be -1.2 kcal mol-l at 25 " C (see Table I). Whetsel and Lady determined the enthalpy of complex formation between aniline and benzene to be -1.64 kcal mol-l.lo By combining these two sets of data, one can derive a lower limit of -2.84 kcal mol-l for the enthalpy of formation of a hydrogen bond between water and the aromatic amino group of aniline. Influence of the Organic Solvent. Clearly, different organic solvents will provide different "nonpolar environments" so that the enthalpy data reported above should not be interpreted exclusively in terms of solutewater interactions. To determine the degree to which solute-organic solvent interactions contribute to the observed enthalpies of transfer, phenol was transferred to water from several different organic solvents. The results of these studies are summarized in Table 11. These data clearly indicate that the AH," values are strongly dependent upon the specific organic solvent selected. Furthermore, by combining the data of Table I1 one can derive enthalpy values for the transfer of phenol
The Journal of Physical Chemistry, Vol. 82, No. 6, 1978 677
Thermodynamics of Transfer
TABLE 11: Enthalpies of Transfer of Phenol from Various Organic Solvents to Water at 25 'Cavb AH,, AH,, Transfer kcal kcal AHtrans, process mol-' mol-' kcal mol-' Octane- H,O -11.60 -5.65 +13.34 -3.91 i 0.11 -9.19 -5.65 t13.34 -1.50 f 0.13 Toluene H,O Octanol-. H,O -5.04 -5.65 t13.34 t2.65 2 0.18
-
AHl, kcal mol-'
See Schemes I and I1 for meaning of symbols. All enthalpies are averages of at least four determinations. The mean calculated enthalpies of transfer are listed in the last column along with the standard errors of the mean. a
TABLE 111: Enthalpies of Transfer of Phenol between Two Organic Solvents at 26 C
--
Transfer process Octane toluene Octane octanol Toluene -. octanol
~&ans,
kcal mol-' - 2.4 -6.5 -4.1
between two organic solvents as illustrated in Scheme 111. These data, which are summarized in Table 111, provide considerable, new insights into the nature of solute-organic solvent interactions and deserve further comment. We have suggested earlier that the 1.5 kcal mol-l exothermicity observed for the transfer of phenol from toluene to water reflects favorable hydrogen bonding between the phenolic hydroxyl group and water. We now see that for the octane-water solvent system phenol has a heat of transfer of -3.9 kcal mol-l (see Table 11). By combining these data we can conclude that phenol has an energetically favorable interaction with toluene (relative to octane) of 2.4 kcal mol-l (see Table 111). This result is in excellent agreement with a spectroscopic study carried out by Pimentel and co-workers.ll These investigators found a favorable interaction of some -2.3 kcal mol-I between phenol and benzene. Thus the derived value of -2.4 kcal mol-l for the transfer of phenol from octane to toluene (see Table 111)is in very good agreement with their findings. In connection with these studies, it is interesting to note that Arnett and co-workersg found no free hydroxyl frequency in the infrared spectrum of phenol in a base as weak as benzene. It should be emphasized that compilations such as Table I11 are useful in that they clearly demonstrate that the enthalpy changes determined for these transfer processes are not all simply a result of interactions with water. In addition, as illustrated above for phenol, we are able to glean further useful information concerning molecular interactions by looking at the derived enthalpies of transfer between two organic solvents. The data in Table I1 also indicate that the transfer of phenol from octanol to water is accompanied by an enthalpy change of +2.65 kcal mol-l. This endothermicity is probably due to the formation of a hydrogen bond between the hydroxyl group of the solute and that of the organic solvent. Such a value for this sort of interaction can be compared with that derived from heats of solution reported in the literature. Parsons and co-workers found the heats of solution for phenol in methanol and water to be +0.88 and +3.04 kcal mol-l, respectively. From these data one can derive a value of +2.16 kcal mol-l for the enthalpy of transfer of phenol from methanol to water. Considering the differences in the solvent systems, this number agrees rather well with the +2.65 kcal mol-l determined for the octanol to water transfer. By combining this result with the data from the octane-water system, one can derive the AHt for the transfer
TABLE I V : Thermodynamics of Transfer of Phenol and Aniline between Various Organic Solvents and Water at 25 C O
AG,",
Octane H,O Toluene- H,O Octanol- H,O
kcal mol-' Phenol -1.14 +0.27 t 1.80
Toluene-. H,O
Aniline t1.32
Transfer process
-
A
q0,,
AH,", kcal mol-'
cal deg-' mol-'
-3.9 -1.5 t2.6
-9.2 -5.9 t 2.8
-1.2
-8.7
of phenol from octane to octanol. As seen in Table 111, a value of -6.5 kcal mol-l is obtained for the hydrogen bond assumed to be formed between phenol and octanol. Such a result is not unreasonable in light of the work of Nagakura12 and Arnettag These investigators studied hydrogen bond formations between phenol and a number of different proton acceptors (ether, dioxane, ethyl acetate, N-methylformamide, N,N-dimethylformamide) and found AH values ranging from -4 to -6.8 kcal mol-l. Free Energies of Transfer. Free energies of transfer were calculated as previously described by determining partition coefficients of the compounds between water and various organic so1vents.l These partition coefficients were found to be independent of solute concentration over the range of 10-1 to M. This allows us to conclude that our data are not influenced by aggregation of the solute molecules. The free energies of transfer, along with the calorimetrically determined enthalpies of transfer, allow the calculation of the entropy of transfer. As discussed below, knowledge of the sign and magnitude of this entropy change leads to further insights into the nature of the molecular interactions involved in a given transfer process. Table IV provides complete thermodynamic profiles for the transfer of phenol from several organic solvents to water and for the transfer of aniline from toluene to water. Significantly, the transfer of phenol from either octane or toluene to water is strongly entropy inhibited as is the transfer of aniline from toluene to water. Frank and Evans13 and Kauzmann2 have described models in which the transfer of nonpolar groups from organic to aqueous medium should result in a decrease in entropy due to the ordering of water molecules around the hydrophobic groups. The observed entropic inhibition to the transfer of phenol and aniline to water is consistent with their predictions. In contrast, the transfer of phenol from octanol to water is accompanied by an increase in entropy. This may well reflect a high degree of order required for the formation of a phenol-octanol hydrogen bond. Clearly, in this solvent system octanol does not provide an "inert" nonpolar environment. Conclusion We believe that we have demonstrated that a great deal of new, fundamentally important information concerning the nature of solute-solvent interactions can be obtained by investigating the thermodynamic changes accompanying the transfer of molecules between organic solvents and water. In the past, such studies were limited to the calculation of free energies of transfer from partition coefficients. The calorimetric technique described here provides a convenient method for determining enthalpies of transfer which in turn allows the calculation of entropies of transfer.
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The Journal of Physical Chemistry, Vol: 82, No. ti, 1978
We strongly believe that further application of this general technique to additional compounds will lead to the forces for the recognition Of the thermodynamic changes observed in many biochemical reactions.
References and Notes (1) K. J. Breslauer, B. Terrin, and J. M. Sturtevant, J. Phys. Chem., 78, 2363 (1974). (2) W. Kauzmann, Adv. Protein Chem., 14, 1 (1959). (3) C. Tanford, J . Am. Chem. Soc., 84, 4240 (1962).
Tanaka
et al.
(4) C. Tanford and Y. Nozaki, J. Bioi. Chem., 246 2211 (1971). (5) I. V. Berezin, A. V. Lerashov, and K. Martinek, F€BS Lett., 7, 20 (1970). ( 6 ) C. Hansch and T. Fujita, J. Am. Chem. Soc., 86, 1616 (1964). (7) J. M. Sturtevant and P. A. Lyons, J. Chem. T h e r w n . , 1,201 (1969). ( 8 ) G. H. Parsons, C. H. Rochester, and C. E. C. Wood, J. Chem. SOC. B, 533 (1971). (9) E. M. Amett, L. Joris, E. Mitchell, T. S. S. F. Marty, T. M. Gorrle, and P. v. R. Schleyer, J. Am. Chem. SOC.,92, 2365 (1970). (10) K. B. Whetsei and J. H. Lady, J . Phys. Chem., 69, 1596 (1965). (11) G.C. Pimentel and C. M. Huggins, J. Phys. Chem., 60, 1615 (1958). (12) S. Nagakura, J. Am. Chem. SOC.,76, 3070 (1954). (13) H. S. Frank and M. W. Evans, J. Chem. Phys., 13, 507 (1945).
Hydrogenation of Dienes and the Selectivity for Partial Hydrogenation on a Molybdenum Disulfide Catalyst Toshio Okuhara, Hlroyukl Itoh, Koshiro Mlyahara, and Ken-ichl Tanaka" Research Institute for Catalysis. Hokkaido University, Sapporo, Japan (Received July 18, 1977) Publication costs assisted by the Research Institute for Catalysis
The hydrogenation reactions of conjugated and nonconjugated dienes on MoSz catalyst are brought about by the 1,2 addition of hydrogen, and the reaction with a mixture of Hz and D2 yields selectively do and dz adducts. It has been found that the hydrogenation of isoprene on heterogeneous catalysts such as MoSz, ZnO, CrZ03, P t black, Pd black, and Raney Ni yields more 2-methyl-1-butene than 3-methyl-l-butene, while homogeneous catalysts give the opposite selectivity. Upon deuteration of isoprene on MoS2,the deuterium molecular identity is maintained more strictly in 3-methyl-1-butene than in 2-methyl-l-butene, which may suggest a repulsive effect of the methyl group of isoprene on its adsorption. An extended Huckel MO calculation for the Mo atom or its ions as well as for a cluster of the Mo ions surrounded by four sulfur ions supports the repulsive effect of the methyl group on adsorption.
Introduction The hydrogenation of conjugated dienes on a heterogeneous catalyst may be brought about by either 1,2 addition1J4 or 1,4 addition2 In a previous paper,3it has been shown that the hydrogenation of butadiene on a MoSz catalyst occurs only through the sec-butenyl intermediate (a-allylic species) and that the a-allylic intermediate does not bring about interconversion to the n-allylic species. In the case of the hydrogenation of butadiene on ZnO catalyst, both 1,2 addition and 1,4 addition occur simultaneously through different intermediate^,^ and the two intermediates do not undergo mutual interconversion on ZnO during the hydrogenation reaction. In contrast to the hydrogenation of butadiene on MoSz and ZnO, it is well known that the a-allylic complex converts easily to the r-allylic form in Co(CN)l- and PdCl(2-methylallyl)PPh3as a function of the concentration of ligands, CN- and PPh3,6$6J0 and the prevalence of the a-allylic form yields 1,2 addition while the n-allylic form prefers 1,4 addition. In conformity with these facts, the hydrogenation of a series of dienes on MoSz catalyst has been performed, and the selectivity for the hydrogenation of dienes by 1,2 addition will be discussed. Experimental Section A commercial MoSz powder (Kanto Chemicals Co.) was purified by boiling in a HC1 solution for several hours and also in an NaOH solution for several hours, followed by washing with distilled water as well as with hot distilled water as described in the previous paper.' The impurity metals analyzed by atomic absorption were as follows (in 0022-365417612082-0678$01 .OO/O
percentage): Fe, 0.02; Mg, 0.0015; Ca, 0.0077; Na, 0.012; Mn, 0.0003; Cr < 0.0001; K < 0.1. Auger spectroscopic analysis of the evacuated MoSz at 450 "C for 4 h is shown in Figure 1,which indicates no appreciable segregation of impurity metals. X-ray diffraction of the MoS2powder showed that MoS2 has a typical 2-H (hexagonal) layer structure, and the surface area of the sample evacuated at ca. 450 "C for several hours was 15 m2/g by the BET method with nitrogen adsorbent. The reaction was carried out in a closed circulating system with a total volume of about 300 mL. Analysis of dienes and/or olefins was carried out by on-line gas chromatography. The deuterioolefins were preliminarily separated by gas chromatography and were subjected to mass spectrometric analysis with 11eV ionization voltage to obtain parent ions, from which the deuterium distribution was computed. The purification of dienes was performed by passing over a trap containing molecular sieve at room temperature. Hydrogen was purified by diffusing through a silver-palladium thimble, and deuterium from a commercial cylinder was used after passage through a liquid nitrogen temperature trap. Pretreatment of the catalyst consisted of evacuation at about 450 "C for 4-5 h, and the reactions were run at room temperature.
Results and Discussion The hydrogenation of butadiene on MoSz catalyst occurs through the sec-butenyl intermediate (a-allylic form), which does not bring about interconversion to the x-allylic intermediate (reaction l).3 As a result, the deuteration of butadiene on MoS2 catalyst gives 1-butene-3,4-d2se0 1978 American Chemlcal Soclety
