2404
Edwin E. Tucker and Sherril D. Christian
radiolysis observations to date, however, do not indicate two absorption bands in trapped-electron spectra for altohol-alkane solutions. In order to resolve this apparent discrepancy between the trapped-electron spectrum a t 77 K and that at room temperature, it would be necessary to make a more extensive optical absorption study of trapped electrons ifi alcohol-alkane solutions at temperature.
terscience, New York, N.Y.. 1968.Chapter 9. (2)J. E. Willard, in "Fundamental Processes in Radiation Chemistry", P. Ausloos, Ed., Wiley, New York, N.Y., 1968,Chapter 9. (3)L. Kevan, Adv. Radiat. Chem., 4, 181-305 (1974). (4) T. Ita, K. Fueki, A. Namiki, and H. Hase, J. Phys. Chem., 77, 1803 (1973). (5) T. Ito, S. Noda, K. Fueki, and z. Kuri, Can. J. Chem., 51,2801 (1973). (6)T. Sawai and W. H. Hamill. J. Phys. Chem., 73,3452(1969). ( 7 ) A. Ekstrom, R. Suenram, and J. E. Willard, J. Phys. Chem., 74, 1888
(1970). (8) J. R. Brandon and R. F. Firestone, J Phys. Chem., 78, 792 (1974) (9) B. J. Brown, N. T. Barker, and D.F Sangster, Aust. J. Chem., 26, 2089
References and Notes
(1973). (1) W H Hamill in "Radical Ions", E. T. Kaiser and L. Kevan, Ed., Wiley-In-
(IO) R. R. Hentz and G. Kenney-Wallace, J. Phys. Chem., 76,2931 (1972).
Suprabinary Hydrogen Bonded Complexes. The Methanol-N,N-Diethyldodecanamide System in n-Hexadecanel Edwin E. Tucker* and Sherrll D. Christian Department of Chemistry, University of Oklahoma, Norman, Oklahoma 73069 (Received May 30, 1975) Publication costs assisted by the National Science Foundation
High-precision vapor pressure studies of dilute methanol-N,N- diethyldodecanamide solutions in n-hexadecane a t 25, 35, and 45' are reported. The concentration range in which the 1:l hydrogen bonded methanol-amide complex is the predominant hetero complex is limited to extremely low alcohol concentration. Even at amide-alcohol concentration ratios greater than 1O:l the common assumption of the presence of only a 1:l complex is not justified. Accurate equilibrium constants and enthalpies for 1:l and higher order complexes obtained from analysis of a large data set support the presence of cooperative hydrogen bonding effects.
Introduction An overwhelming majority of the published work on hydrogen bonding in solution has been limited to consideration of only binary complexes. Familiar examples are studies reporting thermodynamic and spectral data for carboxylic acid dimers, alcohol dimers, and 1:l complexes. A number of recent publications have been based on the assumption that only binary complexes are important in hydroxylic systems even at rather high concentration^.^-^ On the other hand, evidence is accumulating which shows that even in quite dilut,e hydroxylic systems the presence of suprabinary complexes (those containing more than two molecules) simply cannot be A very recent study of methanol-tri- n-octylamine complexes has shown that the assumption of the existence of only a 1:l complex is not justified except in the limiting case of vanishingly small alcohol concentrati0n.l' The careful study of self-association and hetero-association in dilute hydroxylic systems to establish the extent to which suprabinary complexes occur is important. Noncritical use of the assumption that only binary complexes exist in a given system can lead to large errors in derived thermodynamic and spectroscopic parameters for 1:l complexes. Additionally, accurate equilibrium constants and enthalpies for formation of suprabinary complexes can offer some insight into hydrogen bond cooperativity effects The Journal of Physical Chemistry, Vol. 79, No. 23, 1975
and the consequent stabilization of particular hydrogen bonded complexes in dilute solution.EJ1 The present vapor pressure study of the methanol-N,Ndiethyldodecanamide-hexadecane system was undertaken to assess the extent of formation of suprabinary alcoholamide complexes and to derive reliable thermodynamic parameters for the formation of methanol-amide complexes in a relatively inert solvent. Experimental Section Vapor pressure measurements for solutions of methanol in n-hexadecane (Hx) and methanol in N,N-diethyldodecanamide (DEDA)-Hx mixtures were made as previously described.' Very briefly, the experimental measurement may be described as a process of volumetric addition of MeOH vapor to an evacuated system containing either pure Hx or a DEDA-Hx solution. A more complete description of the procedure is given in the supplementary material (see paragraph a t end of text regarding supplementary material). Methanol pressures were measured with either Texas Instruments or Mensor Corp. fused quartz Bourdon tube pressure gages. The minimum pressure resolution of both gages was 0.003 Torr absolute or lower. Both the nonvolatile components Hx (Aldrich) and DEDA (Eastman) had purity greater than 99% as received and were carefully vacuum distilled before use. Reagent
Suprabinary Hydrogen Bonded Complexes
".t
grade methanol was distilled from magnesium methoxide in a 20 plate column. The thermostat bath temperature was controlled to within f0.005'C at 25,35 and 45'.
. Hx
100
Hx
Data Treatment a n d Results The original vapor pressure data are presented in Table I (supplementary material). Figure 1 shows the type of information obtained from the vapor pressure experiment. The primary quantities of interest are the pressure (in Torr) of MeOH above a solution of Hx or Hx + B(DEDA) and AfA which is the excess methanol dissolved in the DEDA-Hx solution compared with the MeOH concentration in Hx at the same temperature and methanol pressure or activity. Correction for MeOH self-association is no problem since any model which closely reproduces the curve for a MeOHHx solution will lead to the proper values of AfA. Equally, with closely spaced points on the MeOH-Hx curve a graphical interpolation should also lead to the same AfA values. The symbols used in Figure 1 do not represent the experimental precision. In particular, we note that the minimum resolution of the pressure gages used is approximately 500 times smaller than the diameter of the plotted symbols. Our experimental precision in A f A is expected to be no better than our reproducibility of the MeOH-Hx curve. In practice, duplicate measurements of the MeOH-Hx curve can be combined and fit to a power series with a RMSD (root mean square deviation) in MeOH concentration of about 0.00005 M . In accordance with previous work we assume that each molecular solute species (monomer or aggregate) obeys Henry's law and that no self-association of amide occurs a t the low concentration levels used. Dipolar association of amides with no active hydrogens can occur but a t the present amide concentrations (10.13 M ) there appears to be little or no self-association of N,N-dialkyl amides.12 The monomer methanol molarity in solution is expressed by PA CA = K D -where
RT
PA =C RT
~(1) ~
..
~ K ~ I C A 4~C ~B K ~ ~ C A 4- ~ .C B (2)
CA and CB are monomer MeOH and amide concentrations, respectively. Similarly, the total amide concentration (fB) is given generally by fB = CB
(torr) e.' 50
40
30
Figure 1. Comparison plot of methanol pressure vs. total methanol concentration for methanol-hexadecane ( 0 )and methanol-N,Ndiethyldodecanamide (0.083 M)-hexadecane (*) at 25'.
34
30
-
+ K ~ I C A C+B K ~ ~ C A +~ KC B~ ~ C A +~ C. .B.
(3)
We notice an extremely important relationship here. For any fixed value of CA (or pressure) the ratio AfAlfB must be a constant if our Henry's law assumption (and our assumption of no base self-association) is correct. That is, if the various equilibrium constants do not change with concentration in these dilute solutions, AfA and fB are both proportional to CB a t any given CA value. Consequently, if we define an apparent equilibrium constant by Kapp = AfA/(CA)(fB) (4) then a plot of Kappvs. CA (or pressure) should result in a
i 0
26.-
d 2s" 0 *
Y'W Kwl
0
*
18
-
35"
*=
0
0
I
*I
.O
4sao* '
I*" 0
or I
*
0
* I
o*.*
1D* .21
*
d I
I
I
14;om*
6.
A (molarity) ~ ~
where P A is monomer alcohol pressure (total pressure corrected for small vapor nonideality effects) and K D is the ratio of alcohol concentration in solution to vapor alcohol concentration at infinite dilution and a 1.0 M standard state. At any given methanol activity, the total heterocomplexed MeOH (AffA)in molarity units can be generally expressed by AfA = KIICACB4-
' +B
I o 0 "
O *I
I
I
I
I
I
I
~
unique curve which will be independent of initial amide concentration. Note that the use of f B in the denominator of eq 4 will strongly repress curvature. For example, if only a 1:l complex existed Kappwould decrease with increasing concentration. Figure 2 presents a plot of Kappvs. CA for our MeOHDEDA-Hx data. It is obvious that within about 1%in Kapp the three data sets (for different initial amide concentrations) do fall on the same curve a t a given temperature. As stated previouslyll such coincidence gives strong support to our Henry's law assumption. Additionally, these curves imply that the Henry's law assumption is also valid for the self-associated MeOH polymers over the present concentration range. Of course, we cannot exclude the possibility that two or more activity coefficient effects of opposite direction could be present but their near-exact cancellation (over a range of more than a factor of 2 in amide concentration) would be extremely fortuituous. We wish to point out that if significant base self-association were occurring, different intercepts would be obtained at a single temperature The Journal of Physical Chemistry, Vol. 79, No. 23, 1975
2486
Edwln E. Tucker and Sherril D. Christian
TABLE 11: Least-Squares Values of K11 a n d K, for Methanol-N,N-DiethyldodecanamideComplexesa
f, = 0.056 M T,OC
Kl1
25
12.81 0.04
*
fB
K,
13.70 f 0.02
= 0.083 M
K11
f, = 0.123 M K,
13.16 .I: 0.02 13.12 f 0.03 9.86 f 0.03
Kl 1
13.75 * 0.01 13.72 rt 0.01 10.35 * 0.02
K,
13.11 f 0.03 13.75 f 0.01 13.18 f 0.02 13.73 0.01 35 9.67 * 0.04 10.22 f 0.01 9.81 f 0.04 10.25 * 0.01 9.84 f 0.03 10.22 f 0.01 45 7.71 f 0.05 7.68 rt 0.02 7.72 f 0.02 7.74 0.01 7.58 0.02 7.79 0.01 a All K units are (liter/mole). The RMSD in A f A for all fits is in the range from 0.00007 to 0.00031 M with an average value for the 12 data sets of 0.00018 M. The average maximum A f A value for all 12 data sets is 0.1165 M.
*
for data sets with different initial amide concentrations. This is not observed. Additionally, if the base does not selfassociate but the heteroequilibrium constants change rapidly with increasing concentration, the Kappvalues a t a single temperature for different base concentrations would diverge with increasing alcohol concentration. This effect is also not observed. In agreement with the results for the MeOH-trioctylamine-Hx system'l Figure 2 shows very dramatically the importance of suprabinary complexes at extremely low MeOH concentration. We particularly note that even in the range of CA values below 0.01 M (where the total amide concentration is as large as an order of magnitude in excess of the total methanol concentration) the assumption of the presence of only a 1:1 complex is simply not justified. It can readily be shown by dividing both sides of eq 2 by the product CACB that the degree to which the apparent Kl1 deviates from the true K11 is only a function of the monomer MeOH concentration and the overall K's for forming complexes larger than 1:l. A t a fixed monomer MeOH concentration the Kapp is independent of base concentration. This observation indicates the fallacy of assuming that having a base concentration greatly in excess of the hydroxylic acid concentration automatically ensures the effective presence of only a 1:l complex. The curves in Figure 2 provide a graphic demonstration of the degree of formation of heterocamplexes which is essentially parameter free. The behavior of these curves suggests the existence of complexes containing more than two MeOH molecules. If the stepwise equilibrium constants for adding a MeOH molecule to the amide or to an existing MeOH-amide complex are not all equal, then the next most simple model involves two equilibrium constants where it is assumed A+B=AB
Ki1
and An-lB
+ A = A,B
K, for n L 2
(5)
that K1l for 1:1 complex formation is unique and that subsequent additions of a manomer give the same free-energy change.13 Explicit expTessions for hfA and f B are then given by
and
(7) Using a value of KD determined from MeOH-Hx data, The Journal of Phxsical Chedstty, Vol. 79. No. 23, 1B75
*
*
*
least-squares fits of each of the 12 MeOH-amide data sets in Table I may be performed to determine the best values of K11 and K,. These values and their standard e.rrors are given in Table 11. The reproducibility of the calculated equilibrium constants is generally quite good and in fact remarkable in those cases where duplicate data sets were taken. Although the calculated K11 values a t 25 and 3 5 O for the lowest concentration amide data are slightly smaller than the K11 values for the other amide concentrations there appears to be no systematic change in the K values with amide concentration. The question may be raised (in view of the small differences between K11 and K , at the two lower temperatures) as to whether we are justified in using two equilibrium constants to fit these data. If the assumption is made that only a single equilibrium constant might describe the system then we possess enough information to directly calculate the value of this K at each data point without any data fitting. Equation 8 gives this K value ex-
plicitly. Evaluation of K from eq 8 a t each data point for all the 25 and 35" data shows that, instead of a random scatter about a mean value, K values are obtained which monotonically increase with increasing MeOH activity. Additionally, the RMSD values for all sets of data a t 25 and 3 5 O rise by about a factor of 5 upon fitting the data with only one equilibrium constant. We conclude that the model of minimum complexity which adequately describes this system contains two equilibrium constants. Enthalpies and entropies derived from fitting the entire set of 12 K11 and K, values (transformed to the mole fraction scale) are given in Table 111. The difference between the aHllo and AH," values is small but apparently real since analysis of any one set of data for a given amide concentration gives a comparable difference between AFZ11" and AH,".
Discussion and Conclusions No cyclic hydrogen bonded MeOH-DEDA complexes can exist but some alternative linear structures are possible. It is generally considered that the primary site of amide interaction with hydrogen bonding acids is at the amide carbonyl.12J4 Although MeOH bonding to the amide nitrogen cannot be totally excluded, this effect should be small in comparison to MeOH bonding to the carbonyl OXYgen due to the electronic structure of the amide. Some previous studies of phenol-carbonyl base complexes have been interpreted in terms of 1:1 and 2:l phenol-base complexes with the 2:l complex assigned one of
2487
Suprabinary Hydrogen Bonded Complexes TABLE 111: Enthalpies and Entropies for
Methanol-N,N-Diethyldodecanamide Complexes in n-Hexadecand -AS,,' = 10.0 * 0.3 = 5.23 i 0.08 kc al/mol cal deg" mol'* -AH,' = 5.58 f 0.04 -AS," = 11.1 * 0.1 a These enthalpies and entropies are based on K values transformed to the mole fraction scale. See E. M. Woolley, J. G. Travers, B. P. Erno, and L. G. Hepler, J. Phys. Chem., 75, 3591 -AHii"
(1971).
the structures b e l o ~ . ~ JNo ~ Jdefinitive ~ evidence exists for the exclusive presence of either structure I or 11. From our
i
I
H
I
I1
thermodynamic data alone we do not expect to divide the stoichiometric concentration of (MeOH)Z-DEDA into numerical proportions of I and.11. However, the use of reasonable assumptions coupled with our thermodynamic data may suggest which structure is favored. Christian and Keenan have provided evidence which strongly suggests that 2:l HC1-ether complexes have structures analogous to I.I6 The equilibrium constant for adding a second HC1 to a 1:1 HC1-ether complex was significantly smaller than the K which would be predicted (K11/4) for addition of the second HC1 in the absence of interaction between the two lone pair sites of the ether oxygen. Formation of the 1:1 HC1ether complex apparently markedly reduces the electron density a t the second lone pair site on the ether oxygen.16 In the present case, we would expect that the K for adding a second MeOH molecule to the 1:1 MeOH-DEDA complex to form structure I would be much smaller than K11 for forming the 1:l complex. Also, we would expect the AH for this process to be smaller in absolute magnitude than AH11'. Since our K, and -AH,' are larger than K11 and -AH11' we conclude that structure I1 is present in excess over structure I. Our thermodynamic data for the MeOH-DEDA-Hx system support the presence of cooperative or nonadditive contributions to stabilization of structure I1 in close similarity to the previously determined enthalpies for MeOHamine complexes.8J1 The additional step of analyzing site equilibrium constants17 for formation of a mixture of structures I and I1 leads to the conclusion that the MeOH oxygen in the 1:l complex has become a better hydrogen bond acceptor than the carbonyl oxygen of the amide. This effective spacial extension of an electron-rich center could have interesting consequences for systems of biological interest. For example, a carbonyl group of a solute in aqueous solution would have one or more water molecules hydrogen bonded to it. The bonded OH group of a water (or other hydroxylic molecule) could effectively serve as a bridge between the carbonyl oxygen and an electrophilic moiety such as a metal ion or a proton donor. Such a bridging OH
group would presumably increase thermodynamic stability and also allow some modification of structural restraints in formation of a particular complex. The significance of such interactions is difficult to assess for aqueous solutions but such configurations appear to exist in crystal structures.18 It is probable that the relative stabilities of a 1:l complex and a 2:l complex with a structure analogous to I1 vary considerably depending on the nature of the R group in ROH, the hydrogen bonding acid. For example, when ROH is methanol8Jl or water8 both the free energy and enthalpy for adding a second ROH to a 1:l ROH-amine complex are more exothermic than the free energy and enthalpy for formation of the 1:l complex. This is also true for the present methanol-amide data. This situation may be contrasted with data for phenol-amine7 and phenol-tetramethyl~rea'~ complexes. Although no enthalpy data are available, the equilibrium constants reported by Zeegers-Huyskens et al.7J5 for adding a second phenol molecule to 1:l phenolbase complexes are generally much smaller than the equilibrium constant for forming the 1:l complex. It may be expected that, when R in ROH is more electron withdrawing than CH3 or H, K and AH for adding ROH to the heterocomplex ROH-B to form a sequential structure similar to I1 will be smaller in absolute value than K and AHllo for formation of the 1:l complex. Methanol and water may be relatively unique (with respect to more acidic proton donors) in their ability to form a heterocomplex similar to I1 which has an enthalpy of formation greater than twice that of the 1:l ROH-B complex. We have previously emphasized the importance of using an experimental technique for hydrogen bonding studies which results in an observable quantity directly proportional to monomer c o n c e n t r a t i ~ n . ~The ~ J ~effectiveness of vapor pressure measurements for the present study of MeOH-amide complexes is shown directly by our ability to demonstrate the formation of suprabinary complexes (Figure 2) without invoking a particular model involving several adjustable parameters. In addition, the direct measurement of methanol activity provides a simple test of our initial assumptions (no base self-association and no marked activity coefficient effects) for the dilute solutions employed. For complex hydrogen bonding systems of the present type the vapor pressure method is probably the most accurate technique for determining monomer concentration. Infrared measurements may certainly be used for this purpose and, in fact, would be superior for systems containing more than one volatile component. However, the interpretation of infrared data on hydroxylic systems is often limited by the extreme difficulty of obtaining both equilibrium constants and absorptivities for associated species from spectral data alone. The most effective use of infrared data would be realized through combining structural information (obtainable from infrared measurements involving isotopic substitution and cancellation studies) with accurate stoichiometric data from vapor pressure measurements. The sole use of nuclear magnetic resonance measurements for study of systems containing suprabinary complexes is unlikely to result in reliable parameters since no quantity proportional to monomer concentration can be measured.10!20 Acknowledgment. We wish to express our appreciation for support of this work by the National Science Foundation (Grant No. GP-43307). The Journal of Physicai Chemistry, Vol. 79, No. 23, 1975
2488
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
