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J. N. Spencer, R. S. Harner, and C. D. Penturelli
Supplementary Material Available, Table I, containing a description of the experimental procedure and the vapor pressure data for methanol-hexadecane and methanolN,N-diethyldodecanamide-hexadecane solutions a t 25, 35, and 45O, will appear following these pages in the microfilm edition of this volume of the journal. Photocopies of the supplementary material from this paper only or microfiche (105 X 148 mm, 24X reduction, negatives) containing all of the supplementary material for the papers in this issue may be obtained from the Business Office, Books and Journals Division, American Chemical Society, 1155 16th St., N.W., Washington, D.C. 20036. Remit check or money order for $4.00 for photocopy or $2.50 for microfiche, referring to code number JPC-75-2484.
References and Notes (1) Presented in part at the 169th National American Chemical Society Meeting, Philadelphia, Pa., April 9, 1975. (2) A. Hall and J. L. Wood, Spectrochim. Acta, Part A, 28, 2331 (1972). (3) M. Gruner and H. G. Hertz, Adv. Mol. Relax. Processes. 3, 75 (1972). (4) T. T. Nakashima. D. D. Traficante, and G. E. Maciel, J. fhys. Chem., 78, 124 (1974). (5) D. Baron and N. Lumbroso-Bader, J. Phys. Chem., 79,479 (1975). ( 6 ) K. B. Whetsel and R. E. Kagarise, Spectrochim.Acta, 18, 315 (1962).
(7) D. Clotman. D. van Lerberghe, and Th. Zeegers-Huyskens, Spectrochim. Acta, Part A, 26, 1621 (1970). ( 8 ) E. E. Tucker, Ph.D. Dissertation, Oklahoma University, 1969. (9) E. E. Tucker, S.B. Farnham, and S. D. Christian, J. fhys. Chem., 73, 3820 (1969). (10) E. E. Tucker and E. D. Becker, J. fhys. Chem., 77, 1763 (1973). (11) E. E. Tucker and S. D. Christian, J. Am. Chem. Soc., 97, 1269 (1975). (12) D. B. Henson and C. A. Swenson, J. fhys. Chem., 77, 2401 (1973). (13) We have previously used a similar model involving three equilibrium constants to fit data for methanol-tri-n-octylamine complexes.” It is possible that the present two constant model could adequately describe that system. Further work is in progress on amine-alcohol systems to resolve this question. When applicable, infinite series models such as these are useful for representing real systems with use of as few parameters as possible. It would be desirable to obtain a specific equilibrium constant for each particular complex but even with the precision of our vapor pressure data we feel that use of any model with more than three Ks would probably not be realistic. (14) L. J. Bellamy and R. J. Pace, Spectrochim. Acta, Part A, 27, 705 (1971). (15) J. P. Muller, G. Vergruysse, and Th. Zeegers-Huyskens, J. Chim. fhys., 89, 1439 (1972). (16) S. D. Christian and B. M. Keenan, J. fhys. Chem., 78, 432 (1974). (17) I. M. Klotz, Acc. Chem. Res., 7, 162 (1974). (18) See, for example, N. Tanaka, T. Yamane, T. Tsukihara, T. Ashida, and M. Kakudo, J. Biochem. (Jpn.),77, 147 (1975). (19) E. E. Tucker, S. D. Christian, and L. N. Lln, J. fhys. Chem., 78, 1443 (1974). (20) E. E. Tucker and E. Lippert, “High Resolution NMR Studies of Hydrogen Bonding”, in “Recent Advances in Hydrogen Bonding”, P. Schuster, G. Zundel and C. Sandorfy, Ed., North-Holland, Amsterdam, in press.
Solvation Effects on the Thermodynamics of Hydrogen Bonding Systems J. N. Spencer,* R. S. Harner, and C. D. Penturelli Department of Chemistry, Lebanon Valley College, Annville, Pennsylvania I7003
(Received June 13, 1975)
Publication costs assisted by Lebanon Valley College
The hydrogen bonding of phenol and guaiacol to dimethyl sulfoxide in the solvents cyclohexane, carbon tetrachloride, carbon disulfide, benzene, 1,2-dichloroethane, and chloroform was studied as a function of temperature by monitoring the hydroxyl stretching frequency at 3 k. Thermodynamic functions are reported and compared to dielectric theory and empirically assigned G values for each solvent. No correlation of the thermodynamic functions with either dielectric theory or G value was found. Specific interactions of phenol or dimethyl sulfoxide with the solvent seems to be primarily responsible for variation of the properties of the phenol systems from solvent to solvent. In chloroform and possibly 1,2-dichloroethane part of the deviations for the thermodynamic data found for the hydrogen bonded adducts is attributed to solvation of dimethyl sulfoxide. In all systems solvation of guaiacol and the guaiacol complex contributes to the observed thermodynamic differences in the various solvents.
Despite the extensive literature on the hydrogen bond, few systematic studies of solvent effects on the thermodynamic properties of hydrogen bonded complexes have been reported.l.2 There are two general approaches to the effects of solvents on hydrogen bonds:3 (1) the so-called dielectric theory which attempts to correlate solvent effects with bulk properties of the solvent such as the dielectric constant; and (2) the specific interaction theory, propounded chiefly by Bellamy and his coworkers,4-6 which offers an explanation of spectral shifts on the basis of specific solutesolvent interactions. Horak and P l i ~ a ,among ~ . ~ othersFJO have used a combination of dielectric and specific interactions to interpret certain hydrogen bond properties. Most investigators agree that the dielectric theory by itself is inThe Journal of Physical Chemistry, Vol. 79,No.23, 1975
adequate and even the opponents of specific interactions admit that dielectric theory is not generally valid.3 This study reports on the thermodynamics of the complexes of DMSO with phenol and guaiacol in various solvents. The solvents chosen have dielectric constants ranging from 2.02 to 10.3 and vary from the proton donating solvent, chloroform, to the proton accepting solvent, benzene. Three solvents, carbon disulfide, carbon tetrachloride, and cyclohexane, are considered inert or at least weekly interactive with proton donors or acceptors, while the solvents chloroform and 1,2-dichloroethane are known to form complexes with DMSO,ll the proton acceptor in this investigation. Benzene, the remaining solvent of this study, is known to form O-HWT bonds to proton donors.12Phenol
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Thermodynamics of Hydrogen Bonding Systems and guaiacol were chosen as proton donors for comparative purposes. The hydroxyl group of guaiacol is intramolecularly hydrogen bonded and consequently solvent effects on the uncomplexed guaiacol will be different from those operative on the free hydroxyl of uncomplexed phenol. Because the proton acceptor is the same for all solvent studies, a comparison of the thermodynamic functions in the various solvents will allow an assessment of the relative importance of specific solvent interactions with free and bonded hydroxyl groups. Experimental Section Baker analyzed reagent grade phenol and Baker practical grade guaiacol were purified as previously described.13 Mallinkrodt analytical reagent grade dimethyl sulfoxide was refluxed over CaO and fractionally distilled under dry nitrogen. Baker analyzed reagent grade carbon disulfide was stored over PzOb and distilled for use from PzO5. Fisher certified ACS spectranalyzed chloroform was shaken with concentrated sulfuric acid, followed.by washing with water. After drying over CaClz the reagent was distilled from P205. Baker analyzed reagent grade benzene was fractionally distilled from sodium. Fisher certified ACS spectranalyzed cyclohexane was distilled over LiAlH4. Fisher certified ACS spectranalyzed 1,2-dichloroethane was distilled over P20.5. The methods of calculation and experimental details have been previously described.13 The molar absorptivities of phenol and guaiacol were determined in all solvents as a function of temperature. Corrections for changes in concentration with temperature were made. Solvent absorption for benzene and cyclohexane was excessive and 2-mm path length cells were used for these solvents. Both 10- and 2-mm cells were used for measurements in carbon disulfide, 1,2-dichloroethane, and chloroform with no detectable differences. The phenol concentrations ranged from 0.005 to 0.01 M with DMSO concentrations ranging from 0.01 to 0.04 M . Guaiacol concentrations were about 0.002 M with DMSO concentrations about 0.1 M. The temperature range over which measurements were made was from 10 to 4OoC for all solvents except CSz for which a maximum temperature of 30°C was used. Frequency measurements were made in lo-, 2-, and 1-mm cells. Due to the broad absorption bands for the bonded species, it was not possible to obtain accurate frequency values in cyclohexane and chloroform and the values for the bonded systems in other solvents cannot be considered highly reliable. The accuracy of measurement is made more difficult by the large frequency shifts which place the bonded absorbance close to the 6-H absorption region. Results and Discussion Prior to the work of Allerhand and Schleyer3 it was commonly assumed that hydrogen bonded systems would be insensitive to solvent effects because the hydrogen atom which might associate with the solvent was already associated with a proton acceptor. These workers found a correlation between the bonded OH-0 frequency and an empirically assigned G value for various solvents. They went on to argue against specific interactions but later Bellamy et al.4 showed that specific interaction theory could be reconciled with Allerhand and Schleyer’s results if solvent interactions with the hydroxyl oxygen and with the hydrogen of the hydroxyl grdilp were considered. Osawa and Yoshida14
provided considerable support for specific interactions through their studies on hydrogen bonded systems. They followed Bellamy et al.4 and considered the interactions to be of a type I or type I1 nature, according to the site of solS---0-H---OR
I
OH---O-H---S
I1
vent interaction, where S is a solvent molecule. A type 11 interaction will be considered to be any interaction between a free hydroxyl hydrogen and solvent for the present work. Osawa and Yoshida14 also provided an interpretation of the G values of Allerhand and Schleyer by pointing out that a close correlation existed between G and the Taft inductive constant n*. Because (r* is a measure of the bond dipole moment, the correlation with G implies that G is a measure of the dipole-dipole interactions. Horak and his coworkers738 had previously advanced much the same ideas for solute-solvent interactions. They considered that dipole-dipole and dipole-induced dipole interactions between solvent and solute could occur and that even solvents which are considered nonpolar form weak complexes with polar solutes through collisions in the liquid state. For hydrogen bonded systems only type I interactions are possible. The correlation between the frequency of absorption of such species and G value found by Allerhand and Schleyer3 would be expected on the basis of Osawa and Yoshida’s14 interpretation because the G value is considered by them to be a measure of the type I interaction. Intramolecularly hydrogen bonded frequencies should also correlate with G and a linear relationship between G and intrabonded frequency was found by Allerhand and S ~ h l e y e r .For ~ those solutes not inter- or intramolecularly hydrogen bonded, a correlation of free hydroxyl absorption frequency is possible only if type I1 interactions with the solvent are weak. Allerhand and Schleyer suggest that free hydroxyl frequencies might correlate with G only for truly “inert” solvents. Table I lists the absorption frequencies found for uncomplexed phenol and guaiacol as well as those for the bonded complexes. The bonded absorption frequencies cannot be considered highly reliable for reasons previously advanced but the trend with G is similar to that observed by othe r ~ . ~No, correlation ~ , ~ ~ of the uncomplexed phenol absorption frequency with G is anticipated or found because type I1 interactions predominate for this solute. The absorption frequency of guaiacol decreases with increasing G value to the proton-donating solvent, chloroform. Because a type I interaction alone is possible for guaiacol, the higher frequency in chloroform must be attributed to a strong interaction of the chloroform with the hydroxyl oxygen according to specific interaction theory? The reversal of the trend of frequency with G value has also been reported for the phenol-dioxane and the phenol dimer bonded absorption in chloroform by Bellamy et al.4 and by Allerhand and Schleyer for the diphenylcarbinol dimer.3 Unfortunately it was not possible to obtain accurate absorbance frequencies for the phenol-DMSO and the guaiacol-DMSO bonded complexes in chloroform but due to the strength of the hydrogen bond formed with DMSO the oxygens of the bond would be expected to be quite basic and provide a‘ strong interaction site for chlor~form.~ The bonded frequency may then behave much as the uncomplexed guaiacol showing a reversal with increasing G value. Also included in Table I is the KBM15J6 relation for the dielectric constant of the medium. There is no satisfactory The Journal of Physical Cherni$try, Vol. 79, No. 23, 1975
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J. N. Spencer, R. S. Harner, and C. D. Penturelli
TABLE I: Frequency Data for Phenol and Guaiacol DMSO Complexes in Various Solventsa Von...o
uOH
Solvent Cyclohexane cc1,
cs2
Benzene 1,2-Dichloroethane CHC1, a Frequencies in c m - l reported.
Phenol
Gu ai aco1
Phenol
Guaiacol
(E -
l)/(2e
+ l)b
G
3617 3564 3266 0.203 49 3610 3556 3256 3251 0.226 69 3595 3551 3243 3238 0.289 74 3557 3540 3231 3238 0.230 80 3564 3530 3210 3206 0.430 95 3602 3547 0.358 106 Dielectric constant at 20°C for all solvents except 1,2-dichloroethanefor which the dielectric constant at 25°C is
TABLE 11: Comparison of D a t a for Phenol and Guaiacol Hydrogen Bonding Equilibria to DMSO in Various Solvents Equilibrium constant' Solvent
Phenol
Guaiacol
ASoC
APb
Phenol
Gu ai aco 1
Phenol
Guaiacol
(E - 1)/
( 2 +~ 1)
G
811 8.69 -8.96 * 0.25 -3.77 i 0.37 -17.1 -8.55 0.203 49 215 4.10 -6.33 f 0.07 -3.04 * 0.14 -10.8 -7.57 0.226 69 381 5.51 -7.08 f 0.24 -4.50 f 0.05 -12.3 -12.4 0.289 74 cs2 Benzene 106 4.09 -5.14 * 0.12 -3.88 0.06 -8.25 -10.4 0.230 80 1 , 2 - Dichloroethane 33.0 1.86 -6.07* 0.14 -3.43 * 0.09 -10.5 0.430 95 -13.8 C HC 1, 19.4 0.86 -3.12 rt 0.03 -2.09 5 0.18 -4.73 -7.39 0.358 106 a Equilibrium constant at 20°C. Units are kcal mol-I. The error reported is the error in the least-squaresslope of a plot of log K vs. T - l . C Units are cal mol-l deg-1. Data from ref 25.
Cyclohexane CC14d
*
correlation between the bonded frequency and this function. According to specific interaction theory, as the strength of the type I interaction increases, the acidity of the hydrogen of the bonded systems should increase with consequent decreased absorption frequency.14 Thus the absorption frequency of uncomplexed guaiacol in chloroform should be lower not higher than that in the other solvents. Also according to this interpretation, the intrinsic hydrogen bond strength should be greater in solvents with higher G value.14 Measurements of the intrinsic hydrogen bond strength in various solvents however are not possible unless the magnitude of the solvation are known and until these parameters are determined the actual effect of solvent on intrinsic hydrogen bond strength will remain ambiguous. Table I1 lists the thermodynamic functions for the DMSO complexes of phenol and guaiacol. The enthalpy changes and the equilibrium constants for the guaiacol complexes are less than those for the phenol complexes in all cases due to the enthalpy required to disrupt the guaiacol intramolecular hydrogen bond. The relative consistency of the enthalpy changes for the guaiacol complexes indicates that the intramolecular bond persists in all solvents studied. There is no trend with G value or with dielectric constant for the thermodynamic functions for either the phenol or guaiacol complexes. If specific type I1 solvent interactions with phenol occur, no general trend of the enthalpy of equilibrium constant with G value would be expected for the phenol systems. Hydrogen bond formation between phenol and the .rr cloud of benzene is well known12 and the enthalpy change for the formation of the phenol-DMSO adduct in benzene is expected to be lower than that in solvents which undergo The Journal of Phy3ical Chemistry, Vol. 79, No. 23, 1975
a weaker solvation interaction. The enthalpy change for the phenol complex in benzene is nearly 4 kcal lower than that found in cyclohexane, while the enthalpy changes for the guaiacol-DMSO complexes in cyclohexane and benzene, for which no type I1 interaction is possible, are nearly identical. The entropy change for phenol in benzene is much lower than that in cyclohexane while the reverse is true for guaiacol. Because no specific solvation effects of type 11 are possible with guaiacol systems, it would be expected that a trend for the enthalpy change with G value would be found if the interpretation of Osawa and Yoshida is correct. The absence of any trend indicates that either G does not reflect type I interactions or that other solvation effects are more important. An assessment of the terms contributing to the enthalpy changes found for the various solvents may be made by consideration of the following cycle:
where D and A are the proton donor and acceptor, respectively, and (A) and (B) refer to different solvents. For the transfer of gaseous phenol to cyclohexane, cc14, and benzene, AH has been reported by Woolley and Hepler17 to be -8.8, -10.1, and -11.7 kcal mol-l, respectively. Arnett et a1.2 have compiled data for the enthalpy change for the bonding of phenol to six different proton acceptors in cyclohexane and CCld. The average difference between the enthalpy changes in the two solvents is 1.4 kcal, which is in good agreement with the solvation difference for phe-
2491
Thermodynamics of Hydrogen Bonding Systems no1 of 1.3 kcal in these solvents as calculated from Woolley and Hepler’s17 data. This implies that (AHA AHDA)is unaffected by these two solvents and further that AHAand AHDAmust largely cancel. Drago et al.l8J9 have advanced the same conclusions, Le., for all adducts of a given acid for a given pair of solvents ( A H A + AHDA) should be constant if no strong specific interactions occur. Then for the phenol complexes of this work, AH(cyc1ohexane) should differ from AH(CC14) by about 1.3 kcal. However, a difference of 2.6 kcal is found. Gramstad20 has reported the difference between m(cc14) and AH(cyc1ohexane) to be 2.3 kcal mol-’ for the formation of phenol trimer and rationalizes the difference by ascribing a bond enthalpy of 2.3 kcal mol-’ to a phenol-CCl4 hydrogen bonded complex. While have ascribed deviations seen in alcohol-CC14 systems to hydrogen bonding between donor and solvent, Drago and Nozarilg claim that such an interaction does not take place and for the pyridine-m-fluorophenol-CCl4 system argue for a pyridine-CC14 interaction of 0.9 kcal mol-l. If the data given by Woolley and Hepler17 are accurate, and they seem to be, based on the agreement with the enthalpy differences calculated from the tabulation of Arnett et al.,* the discrepancy between AH(CC14) and AH(cyc1ohexane) for the phenol-DMSO adducts cannot be attributed solely to the solvation of phenol by C c 4 . If an interaction between DMSO and cc14 occurs, as postulated by Drago for pyridine and CC14, AH in CC14 would be lower than that found in cyclohexane for phenol and guaiacol complexes. The data support this view. However, Dragolg finds no evidence for such an interaction. Drago and Nozarilg have found that when the oxygen acceptor N,N-dimethylacetamide was studied in cyclohexane and CC4, the enthalpy was much larger in cyclohexane. Because the enthalpy change in cyclohexane was close to that measured in the pure polar base, they postulated that aggregation of the proton acceptor about the complex gave rise to a solvation contribution. Similar behavior of the enthalpy was found for other oxygen acceptors. The extension to the present systems is that the aggregation of DMSO about the phenol or guaiacol complexes produces a solvation effect in cyclohexane altering AHDA, From Woolley and Hepler’s data, AHDfor the transfer of phenol from cc4 to benzene may be calculated to be -1.6 kcal and this largely accounts for the difference in enthalpy observed in these two solvents for the phenol adducts. For the transfer of phenol from cyclohexane to benzene, AHD is calculated to be -2.9 kcal which when added to AH(benzene) misses AH(cyc1ohexane) by about 0.9 kcal. Presumably this difference would be attributed to aggregation of DMSO about the complex in cyclohexane. However for the guaiacol adducts in cyclohexane and benzene, the enthalpy changes are essentially equal. It may be argued that because both guaiacol and the guaiacol-DMSO complex are hydrogen bonded that clustering of DMSO about both species occurs which tends to minimize solvation differences. This argument does not appear valid however because similar behavior would be required for C C 4 solvent. Because AHDand AHDAfor tho guaiacol systems refer to hydrogen bonded systems, it is reasonable to anticipate that these terms may be relatively solvent insensitive. Then for the formation of the guaiacol complexes AH(A) should differ from AH(B) by
+
AH(A)
+
- AH(B) = AHA +
k’
where k’ is (AHD AHDA).Drago et al.lSJg have deter-
mined AHA for the transfer of DMSO between various solvents. Table I11 lists the pertinent data and the calculated k’. Because k’ is not constant, interactions with the hydrogen bonded species must be important. This would imply that if G is a measure of the type I interaction there should be a correlation between AH and G for these systems. No correlation is found and this suggests that if solvent-solute interactions occur through the type I mechanism postulated by specific interaction theory, the effect on the enthalpy is small. The difference between AH for the phenol complexes in the solvents CC14 and benzene seems to be largely determined by AHDwhile solvation effects attributable to AHA and AHDAare more important for the other solvents. For the transfer to 1,2-dichloroethane and chloroform, AHA probably plays a more pronounced role. Chloroform is known to form a hydrogen bonded complex with and frequency shifts have been reported for this complex and that between DMSO and 1,2-dichloroethane.” The enthalpy change for the formation of phenol complexes in these two solvents will be lower than that in more “inert” solvents such as cyclohexane due to the enthalpy required to break the DMSO-solvent complex. Chloroform forms a stronger complex with DMSO than 1,2-dichloroethane and this is reflected in the smaller enthalpy and entropy changes for the phenol-DMSO complex in CHC13. Similar behavior is seen for the guaiacol-DMSO complex in chloroform. Comparison of AH for the guaiacol-DMSO complex in CHCls to that found for the other solvents suggests that the strength of the DMSO-CHClz hydrogen bond is about 1.5 kcal. A comparison of the enthalpy and entropy changes for the formation of the phenol and guaiacol adducts with DMSO in chloroform shows that not all the solvation effects can be attributed to the DMSO-CHC1:j interaction. Both the enthalpy and entropy for the phenol complex show a greater change in chloroform relative to the other solvents than the corresponding functions for guaiacol. For a given solvent, if the enthalpy change for the formation of the phenol-DMSO complex is subtracted from that for the guaiacol-DMSO complex, the resultant enthalpy refers to the process:
H---DMSO
0 ’
For this displacement reaction, DMSO does not appear so that complications due to solvation of DMSO have been eliminated. If the intrinsic hydrogen bond strengths of the phenol and guaiacol adducts are about the same as is indicated by nearly identical absorption frequencies for both complexes, the enthalpy change found for eq 1 gives the enthalpy required to break the intramolecular bond and the difference in solvation of the reactants and products. The enthalpy changes for eq 1 are given in Table IV and it is noted that with the exception of 1,2-dichloroethane a trend with G is seen. Drago and NozariIg have also found disThe Journal of phvsical Chemistry, Vol. 79, No. 23, 1975
2492
J. N. Spencer, R. S.Harner, and C. D. Penturelli
TABLE 111: Transfer Enthalpies for Guaiacola
- -
AH( A) AH( B)
AH,^ k‘ a
cc4
CCh
c6H6
CZH4C12
CZH4C1.2
-3 -04 -3.88 -0 -74 +1.58
-3.04 -3.43 -2.20 +2.59
-3.88 -3.43 -1 -46 +1.01
Units are kcal mol-I.
c6&
Data from ref 19.
TABLE IV: Enthalpy Changes and Transfer Enthalpiesa for the Displacement Reaction of Ea 1 Solvent
AHb
--C6H12C
CC14
‘gH12
CC14
CgHg
‘sHs
-~~
Cyclohexane
cc14 cs2
Benzene 1,2-Dichloroethane CHC1,
+5.19 -1.90 +3.29 (-1.3)d 1-2.58 +1.26 +2.64 +LO3
-2.03 (-l.6)d
~~
~
-3.93 (-2.9)d
Units are kcal mol-1. 6 Enthalpy change for the reaction given by eq 1. Enthalpy difference for the reaction of eq 1 in the two solvents listed. d Solvation difference of phenol between the pair of solvents. a
crepancies with displacement reactions in 1,Z-dichloroethane and suggest that this solvent should be avoided because of “extensive and unpredictable” solvation contributions. Others,2,26however, disagree with Drago and Nozari. For the solvents cyclohexane, CC14, and benzene the enthalpy change for eq 1 decreases with increasing solvation of phenol as calculated from the data of Woolley and Hepler. Also given in Table IV are the enthalpy changes for the transfer of the reaction of eq 1 from one solvent to another. Because all the species of eq 1 are engaged in hydrogen bonding with the exception of phenol, the enthalpy difference observed for the transfer will be indicative of the solvation of these species if the solvation of phenol is taken into account. The numbers in parentheses in Table IV give the enthalpy difference due to solvation of phenol and it is seen that in all cases the enthalpy change for eq 1 is greater than can be accounted for the solvation of phenol. This implies that solvation effects on the hydrogen bonded species are not negligible. Because of the similarity of structure of the two intermolecularly hydrogen bonded species of eq 1, solvation effects due to these species should be minimized so that the bulk of the solvation effects are due to the difference in solvation between phenol and guaiacol in a given solvent. The variation of the equilibrium constant (Table 11) for the phenol complexes in cyclohexane, CS2, cc14, and benzene follows the trend in AH and reflects increasing stabilization of the uncomplexed phenol by solvent interactions. The smaller equilibrium constants in 1,2-dichloroethane and CHC13 result at least partially from stabilization of uncomplexed DMSO by solvent interaction. The equilibrium constants for the guaiacol-DMSO complexes in these two solvents are much smaller than those in the other solvents and some measure of the stabilization of uncomplexed DMSO by these solvents may be made by comparison with the guaiacol equilibrium constants in the other solvents. The equilibrium constant for guaiacol-DMSO in cycloThe Journal of Physical Chemistry, Vol. 79, No. 23, 1975
hexane, cc4, CS2, and benzene shows no trend with G value. If type I interactions were significant in determining the equilibrium position, the equilibrium constant should increase with G. Solvent interaction with the hydroxyl oxygen would be expected to be greater for the complex than the uncomplexed guaiacol because the basicity of the bonded oxygens of the complex will be greater. If type I interactions are not significant, the equilibrium constants for guaiacol complexes should be relatively solvent insensitive for those solvents which do interact strongly with the acceptor. The largest equilibrium constant found for guaiacol-DMSO is in cyclohexane and implies that solvation effects in the other solvents stabilize the reactants. Conclusions It has recently been commented that extrasensory perception is needed in order to know when solvents interact specifically with a solute.ls It now seems that just as there is a solvent effect on the absorbance frequencies of hydrogen bonded species, there is also a solvation contribution to the thermodynamic properties of hydrogen bonded adducts. There appears to be a limited correlation between G values and absorption frequencies but whether this is an inductive effect as favored by Allerhand and Schleyer3 or specific as suggested by Bellamy et ale4and Osawa and Yoshida14 is not clear. It does seem clear, however, that, if G is a measure of dipole-dipole interactions, the effect on the thermodynamic properties of hydrogen bonded systems is small compared to other effects. When the diversity of solvation interactions is considered, it is not surprising that a general method of correlating hydrogen bonded properties has not yet been produced. Even for guaiacol equilibria, where specific solvation effects should be minimized, the thermodynamic functions are sensitive to solvent. In those cases where solvation effects on the proton acceptor can be taken into account, thermodynamic differences, which must be attributed to solvation of the bonded species, are found. Solvents such as cyclohexane and carbon tetrachloride do not permit a simple correlation of thermodynamic data and aggregation about the complex in cyclohexane and specific interactions with carbon tetrachloride have been invoked to explain thermodynamic differences in the two media. For the guaiacol systems of the present study, carbon tetrachloride appears to be less like cyclohexane in its solvating interactions than is benzene. It is debatable whether or not a “truly inert” solvent exists. In this respect, the conclusion of Arnett et a1.2 that, “there will always be some interactions which are idiosyncratic to any solvent-solute pair; our problem is not so much of identifying specific interactions but of deciding on how large they must be to deserve special attention”, is of principle concern to the study of hydrogen bonding. Acknowledgment. Support for this work was provided by Research Corporation through the Cottrell College Science Grants program. The authors also thank Judy Sweigart and Jerry Steiner for many valuable contributions to this study. References and Notes (1) A. S. N. Murthy and C. N. R. Rao. Appl. Spectrosc. Rev., 89, 191 (1968). (2) E. M. Amett, E. J. Mitchell, and T . S. S. R. Murty, J. Am. Chem. Sm., 96, 3875 (1974). (3) A. Allerhand and P. v. R. Schleyer, J. Am. Chem. SOC.,95, 371 (1963). (4) L. J. Bellamy, K. J. Morgan, and R. J. Pace, Spectrochim. Acta, 22, 535 (1966). ( 5 ) H. E. Hallam, J. Mol. Struct., 3, 43 (1969).
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Viscosities and Partial Molal Volumes of Tetramethylcarboxamides L. J. Bellamy, H. E. Hallam. and R . L. Williams, Trans. Faraday SOC.,54,
1120 (1958). M. Horak and J. Pliva, Spectrochim. Acta, 21, 911 (1965). M. Horak, J. Moravec, and J. Pliva, Spectrochim. Acta, 21, 919 (1965). L. J. Bellamy and R . L. Wllliams, Proc. R. SOC.(London), Ser. A, 255, 22
(1960). G. L. Caldow and H. W. Thompson, Proc. R.SOC.(London), Ser. A, 254, 1 (1960). A. Allerhand and P. v. R. Schleyer, J. Am. Chem. Soc., 85, 1715
(1963). D. Bloser, J. Murphy, and J. N. Spencer, Can. J. Chem., 49, 3913
(1971). J. N. Spencer, R. A. Heckman. R. S. Harner, S. L. Shoop, and K. S. Robertson, J. Phys. Chem., 77,3103(1973). E. Osawa and Z. Yoshida, Spectrochim. Acta, Part A, 23, 2029 (1967).
(15) W. West and R. T. Edwards, J. Chem. Phys., 5 , 14 (1937). (16)E. Bauer and M. Magat, J. Phys. Radium, 9,319 (1938). (17)E. M. Woolley and L. G. Hepler, J. Phys. Chem., 78, 3058 (1972). (18) R. M. Guidry and R. S.Drago, J. Phys. Chem., 78, 454 (1974). (19)M. S.NotariandR. S.Drago, J. Am. Chem. SOC.,94, 6877 (1972). (20)T. Gramstad and 0. Mundheim, Spectrochim. Acta, Part A, 28, 1405 (1972). (21)W. C.Duer and G. L. Bertrand, J. Am. Chem. SOC.,92, 2587 (1970). (22)A. N. Fletcher and C. A. Heller, J. Phys. Chem.,71,3742 (1967). (23)E. M. Arnett, L. Joris, E. Mitchell, T. S.S.R . Murty, T. M. Gorrie, and P. v. R. Schleyer, J. Am. Chem. Soc., 92, 2365 (1970). (24)G. G. Hammes and P. J. Lillford, J. Am. Chem. SOC.,92, 7578 (1970). (25)J. N. Spencer, K. S. Robertson, and E. E. Quick, J. Phys. Chem., 78, 2236 (1974). (26)G. Olofsson and I. Olofsson, J. Am. Chem. SOC.,95,7231 (1973)
Viscosities and Partial Molal Volumes of Some Tetramethylcarboxamides at 25OC P. P. DeLuca' and T. V. Rebagay College of Pharmacy, University of Kentucky, Lexington, Kentucky 40506 (Received December 2, 1974: Revised Manuscript Received July 28, 1975) Publication costs assisted by the College of Pharmacy, University of Kentucky
The viscosities and densities of four tetramethyldicarboxamides, tetramethylpimelamide, tetramethylazelamide, tetramethylsuberamide, and tetramethylsebacamide, were measured throughout the entire aqueous solubility range of these compounds a t 25'C. An interesting feature of the partial molal volume behavior patterns of the amides is the presence of a minimum in the partial molal volume and a maximum in viscosity occurring at definite integral mole fractions for the odd C homologs, tetramethylpimelamide and tetramethylazelamide. Since these occurrences were not observed in the even C analogs, tetramethylsuberamide and tetramethylsebacamide, a stronger amide-water interaction is indicated for the odd C amides.
Introduction Substituted carboxylic acid amides have been found to exhibit potential as solubilizing agents for hydrophobic molecules. The solvent properties of amides can be attributed in part to the strong dipole of the amide function and to the good electron-donating properties of the carbonyl group. Previous studies on the physical-chemical properties of the tetramethyl-substituted amides of pimelamide (TMPA), suberamide (TMSuA), azelamide (TMAA), and sebacamide (TMSeA) revealed that the amides show strong association tendencies in very dilute solutions and suggested that further aggregation of these associated molecules leads to mice1lization.l In continuing investigations of the tetramethyl-substituted dicarboxylic acid amides possessing the following structure, where n = 5-8
This study was undertaken because viscosity and molal volume resulting from density data can provide valuable and interesting information on structure, molecular state, symmetry, and constitution. For example, Assarsson and Eirich2 reported large viscosity maxima in all mixtures of water and N,N-disubstituted amides with the maxima occurring a t specific mole ratios. Association and complex formation between the peptide dipole and water have been offered as possible explanations to account for this observed phenomenon. Herskovits and Kelly3 have used viscosity studies of a series of amides to discuss the water structure-forming and structural-breaking influence of the nonpolar hydrocarbon moieties and the polar amide portion of the amides. Since the possibility of utilizing these amides as solubilizers for poorly water-soluble drugs has been established1 it was also considered desirable to extend the work to investigate the effect of molecular size, aggregation, and composition on the equilibrium and transport properties of these amides in aqueous solutions. These amides form a useful homologous series for such study. Experimental Section
to gain more knowledge of these agents as solubilizing agents, densities and viscosities throughout the entire aqueous solubility range were measured.
The synthesis and purification of the amides have been reported e1sewhere.l The synthesis involved the conversion of the acids to the acyl halides by treatment with excess thionyl chloride. The acyl halides were then allowed to The Journal of Physical Chemistry, Voi. 79, No. 23, 1975
