J. Phys. Chem. 1985,89, 1888-1891
1888
Solvent Effects on Hydrogen-Bond Formation. 2 J. N. Spencer,* C. L. Campanella, E. M. Harris, and W. S. Wolbach Department of Chemistry, Franklin and Marshall College, Lancaster, Pennsylvania I7604 (Received: July 26, 1984)
Thermodynamic parameters for the formation of hydrogen-bonded complexes of phenol and pyrrole with various bases in the solvents cyclohexane, CCl,, and benzene were determined by calorimetric and spectroscopicanalysis. These parameters are strongly influenced by solvent. The enthalpy changes and equilibrium constants for the phenol-DMA complex in cyclohexane and benzene are -9.0 kcal mol-' and 433 and -6.3 kcal mol-' and 71, respectively. A simple model for solvation effects has been proposed. The model requires that the solutesolvent interactions that occur at those binding sites used in the formation of the complex be overcome before the complex can be formed. These interaction site enthalpies have been determined by calorimetric analysis and used to test the model.
Introduction
Even though solvation effects on thermodynamic parameters for complex formation can be quite large, systematic studies of these effects are lacking. A recent compilation of the thermodynamic parameters for hydrogen-bonded complexes' lists results of some 600 investigations of phenol adducts, but only 80 of these studies were carried out in solvents other than CCl,. An understanding of solvent effects has been hampered not only by lack of data but also by the lack of agreement between existing data; for example, the reported enthalpy changes for phenol-pyridine complex formation in CC14range from -5.6 to -9.0 kcal mol-'.' The present study was undertaken to provide data for three solvent systems and to determine by calorimetric and spectroscopic methods the thermodynamic parameters for hydrogen-bond formation. These data are then used to test an interaction site approach to the understanding of solvation phenomena. Experimental Section
Reagents were purified as follows: CCl,, cyclohexane, benzene, N,N-dimethylacetamide (DMA), and pyridine were refluxed over P205, fractionally distilled under dry N,, and collected over molecular sieves. Pyrrole was dried over NaOH, fractionally distilled under dry N2, and collected over sieves. N,N-dimethylformamide (DMF) was fractionally distilled with benzene (which had been dried over P20,) in a 10%v/v ratio under dry N2 and collected over sieves. Dioxane was refluxed over NaOH, dried with CaCl,, fractionally distilled, and collected over sieves. Cyclohexanone and tetrahydropyran (THP) were dried with CaSO,, fractionally distilled under dry N2, and collected over sieves. The calorimeter and calorimetric methods of data collection have been d e s ~ r i b e d . ~Two . ~ methods of analysis of the calorimetric data for complex formation were employed. The Bolles-Drago method4 uses the heat change, Q,corrected for the heat of solution of the acid, the initial concentrations of acid and base, COAand COB, and the volume, V, of the solvent in the following equation:
This equation has two unknowns, the equilibrium constant, K , and the enthalpy change for the reaction, AH.Q is obtained for several different acid and base concentrations. Arbitrary values of A H are chosen, and using Q, V, eoA, and COBwe obtained a series of solutions to eq 1. K-' is plotted vs. the assumed AH, and a curve (1) Joesten, M. D.; Schaad, L. J. "Hydrogen Bonding"; Marcel Dekker: New York, 1974. (2) Spencer, J. N.; Sweigart, J. R.; Brown, M. E.; Bensing, R. L.; Hassinger, T. L.; Kelly, W.; Housel, D. L.; Reisinger, G. W.; Reifsnyder, P. S.; Gleim, J. E.; Peiper, J C. J. Phys. Chem. 1977,81, 2237. (3) Spencer, J. N.; Gleim, J. E.; Hackman, M. L.; Blevins, C. H ; Garrett, R. C. J. Phys. Chem. 1978,82, 563. (4) Bolles, T F.; Drago, R. S J Am. Chem. Soc. 1965, 87, 5015
0022-3654/85/2089-1888$01.50/0
is obtained for eacb set of concentrations; the intersections of these curves give graphical solutions to the simultaneous equations. The second method uses a least-squares technique for data a n a l y ~ i s . ~In general, the heat change for n reactions in the reaction vessel can be given by n
Q = 1C= 1 n , , p 4
(2)
where n,,p is the number of moles of product i and is a function of the equilibrium constant for reaction i . The best values of K, and AH, are calculated by a least-squares analysis of eq 2. The error square sum over the m data points is given by U(K,,AHJ =
(Qp
p=l
- h,,pAHl))2 ,=l
where the subscript p is over all the data points and i is over all reactions. The best values for K and AH for a given determination are those which minimize U(K,,AH,).The complete solution of the least-squares equations involves five steps: (1) assumption of initial K values, (2) calculation of the concentration of each species using the assumed K , ( 3 ) calculation of the best AH corresponding to each K , (4)evaluation of K and AH values to establish how well the data are fit, and ( 5 ) recalculation of steps (2), (3), and (4)using new K values until the best set of K and AH values is found. In addition to the calorimetric analysis, equilibrium constants were determined through near-IR spectroscopy by application of Beer's law. In those systems for which it was possible to determine the variation of the equilibrium constant with temperature, the enthalpy changes were also calculated from the variation of the equilibrium constant with temperature.* A combination of calorimetric and spectroscopic data was also used to determine the thermodynamic parameters for hydrogenbond formation. The heat of interaction, Q, is determined calorimetrically by injection of acid into a dilute solution of the base in the solvent of interest. Q is related to the enthalpy of hydrogen-bond formation, AH, by Q = AHC,V (4),where Vis the volume of the reaction mixture and C, is the equilibrium concentration of the adduct. C, is found from the equilibrium constant determined by near-IR spectroscopy, and A H is calculated directly from C,, Q, and V. Results
Table I lists the enthalpy changes for the formation of the hydrogen-bonded complexes in cyclohexane, CCl,, and benzene solvents. The least-squares and Bolles-Drago analyses of the calorimetric data are given in columns four and five. The enthalpies determined from the temperature variation of the equilibrium constant as determined by near-IR spectroscopy are ( 5 ) Christensen, J. J.; Ruckman, J.; Eatough, D. J.; Izatt, R. M. Thermochim. Acta 1972, 3, 203, 219, 233.
0 1985 American Chemical Society
The Journal of Physical Chemistry, Vol. 89, No. 10, 1985 1889
Solvent Effects on Hydrogen-Bond Formation TABLE I: Enthalpies of Complex Formation
-AH, kcal mol-’
acid phenol
base DMA
solvent
L-B
B-D~
SpeCC
spec-cald
9.0 7.2 6.3 9.6 6.9 5.8 6.1 5.2 4.2 5.9 5.3 4.8 6.8 5.2 4.3 5.4 3.9 3.1
8.9 f 0.2 7.2 f 0.1 6.3 f 0.2 9.7 f 0.4 6.8 0.1 5.8 f 0.1 6.0 f 0.5 5.2 f 0.2 4.3 f 0.4 5.8 f 0.2 5.3 f 0.4 4.7 f 0.7 6.8 f 0.4 5.0 f 0.5 4.4 f 0.4 5.4 f 0.5 3.8 f 0.4 3.1 f 0.1
9.3 f 0.7 7.2 f 0.3 6.2 f 0.5 8.1 f 0.2 6.4 f 0.3 5.3 f 0.6 3.3 f 0.1 2.6 f 0.1
8.8 f 0.4 7.0 f 0.2 6.2 f 0.1 9.1 f 0.5 6.8 f 0.1 5.7 f 0.1 5.7 f 0.2 4.6 f 0.3 3.9 f 0.2 5.2 f 0.3 4.8 f 0.3 4.4 f 0.3 6.4 f 0.2 4.5 f 0.5 4.1 f 0.1 5.9 f 0.4 4.7 f 0.4 3.2 f 0.0
C6H I2
CC14 C,H, DMF
C&;2
CC14 C6H6
dioxane
C6H12
THP
C6H6 C6H12
cc14 cc14 C6H6
cyclohexanone pyrrole
pyridine
C6HIZ
CCI4 CnH, GH;2 CCI4 C6H6
*
5.2 f 0.4 3.1 f 0.4 3.6 0.2 4.3 0.5 3.2 f 0.1 3.4 f 0.2
* *
“Least-squaresdata analysis of calorimetric data. The average standard error is 3%. * Bolles-Drago data analysis of calorimetric data. cNear-IR spectroscopic analysis. “Combined calorimetric-spectroscopic analysis. TABLE 11 Eauilibrium Constants at 298 K
K acid phenol
base DMA
solvent
L-Sb
B-D
spec
Av,O cm-I
C6H12
316 47 68 45 43 32 13 6.9 4.6 21 10.2 4.6 23 11 5.3 4.9 2.9 1.o
292 f 109 49 f 7 60 f 2 38 f 15 64 f 35 34 f 6 13 f 5 7.2 f 0.9 4.3 f 0.7 23 f 4 11 f 3 4.6 f 1.2 25 f 9 11f4 4.8 f 1.0 4.7 f 1.0 2.9 f 0.5 1.0 f 0.06
433 f 42 8 3 f 11 71f3 181 f 7 49 f 2 38 f 2 21 f 1 13fl 5.4 f 0.9 38 f 2 16.1 f 0.5 5.3 0.2 44 f 3 19 f 2 5.9 0.1 3.9 f 0.1 2.0 f 0.1 0.93 f 0.05
337 342 305 211 286 260 225 238 209 27 1 287 260 239 243 217 235 239 218
CC14 C6H6
DMF
C6H12
CC14 dioxane
C6H6 C6H12
CC14 THP
C6H6 C6H12
CC14
cyclohexanone
C6H6 C6H12
CC14
pyrrole
pyridine
C6H6 C6H12
CCl4 C6H6
*
*
“Frequency shift relative to the free N-H or 0-H stretch. In the solvents C6HI2,CCl,, and benzene, for phenol the free 0-H stretch is at 3621, 3614, and 3559 cm-I and for pyrrole the N-H stretch is at 3502, 3498, and 3461 cm-’, respectively. bThe standard error is 11%. given in column six. The combined calorimetric and spectroscopic data have been used in eq 4 to obtain the enthalpies given in column seven of Table I. Equilibrium constants determined by calorimetry and near-IR spectroscopy are given in Table 11. For those systems with small equilibrium constants and small enthalpy changes the spectroscopic data do not give thermodynamic parameters that are in good agreement with the calorimetric or combined calorimetric-spectroscopic data. The difficulty of obtaining reliable thermodynamic information from the temperature variation of the equilibrium constant is well-known,6 and for the determination of enthalpy changes the calorimetric procedures are preferred. If, for example, the relative error in the equilibrium constant is f10% at each end of the temperature intefval over which K was varied and if the mean temperature of that interval is 298 K and the interval 40 K, then the error in A H will be 0.6 kcal mol-’. If K > 25, AH may be determined from the calorimetricspectroscopic method to within f0.2 kcal mol-’ for a 10% error in K under the conditions used in this work. If K > 25, the combined calorimetricspectroscopic analysis can give reliable results for enthalpy changes. The equilibrium constants determined spectroscopically and calorimetrically are considerably different for several systems. The spectroscopic (6) Benson, S. W. “Thermochemical Kinetics”; Wiley: New York, 1976; pp 7-8.
equilibrium constant determined at 298 K may be more reliable than the calorimetrically determined equilibrium constant provided K is large enough to ensure sufficient concentrations of complex are formed.
Discussion When the attractive forces between a solvent and solute are greater than those which exist between solvent-solvent or solute-solute molecules, the solute is said to be solvated. These solutesolvent interactions may be classified as nonspecific, Le., due to dipolar and van der Waals forces, or as specific, in which case an orientation of the solute and solvent molecules exists for which donor and acceptor orbitals overlap to the greatest extent.’ The hydrogen bond is readily identifiable as a specific interaction of the donoracceptor type. Other solutesolvent interactions are not so clear-cut, and the distinction between specific and nonspecific interactions becomes blurred. Two of the solvents of this study, CC14, and benzene, have surprisingly strong electrophilic properties. These properties were apparently first recognized by Gutmann and co-workers? who (7) Drago, R. S.; Parr, L.9.; Chamberlain, C. S. J. Am. Chem. SOC.1977, 99, 3203. (8) Gutmann, V . ‘The Donor-Acceptor Approach to Molecular Interactions”; Plenum Press: New York, 1978; pp 27-33.
1890 The Journal of Physical Chemistry, Vol. 89, No. 10, 1985
assigned an empirical number called the acceptor number to various solvents. The acceptor number expresses the acceptor property of a given solvent relative to that of SbCIS. Both CC14 and benzene have higher acceptor numbers than THF and diethyl ether and can act as electron pair acceptors. Thus, a solute such as pyridine could undergo a donoracceptor interaction with CC14 or benzene. The nitrogen lone-pair interaction by pyridine and ammonia with CC14 seems well established and appears to be of the n,u type.739J0 The pyridine-benzene interaction may be A,* or n,r. Huyskens concludes the ?r,a type is most likely based on dipole moment measurements in cyclohexane, benzene, and CC4
solvent^.^ Carbonyl and ether oxygens have two lone pairs and provide two sites for donor-acceptor interactions. It is generally accepted that two hydrogen bonds to such oxygens can be formed, and in the case of water solvent it seems certain that carbonyl oxygens are solvated specifically by hydrogen bonds from two water molecules. A recent study of directional hydrogen bonding of hydrogen-bond donors to ketones and ethers" has shown a pronounced tendency for the concentration of hydrogen-bonded group to lay in the direction ascribed to lone pairs. Because hydrogen bonding is but one example of the broad class of donoracceptor interactions, it seems reasonable that other molecules capable of these interactions would show similar orientational selectivity in their interactions with lone pairs. Thus, the lone pairs of carbonyl and ether oxygens would be expected to undergo two specific solvation interactions with C C 4 and benzene similar to those for the nitrogen lone pair of pyridine. The first solvation shell of solutes consists of several solvent molecules each at the favored energy minima in the potential field about the solute. Theoretical studiesI2indicate that, even for water solvent, the completion of a second shell of tightly bound water molecules around a neutral polar solute is unlikely. Therefore, in a first analysis, for organic solvents, it seems reasonable that only the first solvation shell need be considered. These considerations allow a model to be constructed for solvation of pyridines and other molecules with carbonyl groups or ether linkages. Because of the large solvation effects on the thermodynamic parameters of hydrogen-bond formation, hydrcgen-bonded systems allow a convenient means of applying the principles outlined. The model assumes that solvation of solutes containing lone pairs may occur through donor-acceptor interactions a t specific sites on the solute molecules. It is possible that the dipole moment of the complex may be enhanced over that of the monomers comprising the complex. In such a case the solvation of the complex may be of significance. However, the limited studies of gas-phase reactions show that more exothermic enthalpy changes are found in the gas phase than in solution. This means that the solvation of the two monomers is larger than that of the com~1ex.l~ Hence, it is assumed for this study that no significant solvation of the complex due to increased dipole moment occurs. The acid and base are assumed to bind through selected sites; e.g., the phenol-pyridine complex binds through the nitrogen lone pair and hydroxyl proton. Those solvent molecules engaged in specific interactions at the sites necessary for hydrogen bonding must be cleared before complex formation can occur. This requires an input of enthalpy equivalent to the solvent-solute interaction enthalpy. If the monomers are more strongly solvated by one solvent than by another, the enthalpy change for the formation of the complex in the more strongly solvating solvent will be more positive than that for the more weakly solvating medium. The three solvents chosen for this study provide a poorly solvating medium, cyclohexane; CC14 is moderately solvating, and benzene is more strongly solvating. The enthalpy change for complex formation in a given solvent will be determined by the ~
Spencer et al. TABLE III: Enthalpies of Solution at 298 K AH,.kcal mol-] solute DMA DMF dioxane THP cyclohexanone pyridine phenol pyrrole cyclohexane N-methylpyrrole
C6H12
3.20 f 0.12b 3.25 0.09b 1.5Y 0.52c 1.77 0.04 2.01" 7.625 & 0.02' 3.74 0.13b 0.00 1.85 f 0.06'
*
*
CCll
* *
C6H6
0.40 0.09' 0.22 f 0.08 0.85 0.04b 0.28 f 0.04 -0.04c -0.1 8c -0.52 0.08 -0.14 f 0.03 0.07 f 0.03 -0.13 0.07 0.36 f 0.01" 0.00 0.01' 6.27 0.07" 4.72 0.01" 2.22 0.06b 1.01 0.02 0.17c 0.92c 0.38 0.02' 0.17 0.06
**
*
*
*
"Reference 14. 'Reference 17. 'Calculated from data given in 'Handbook of Heats of Mixing"; Christensen, J. J., Hanks, R. W., Izatt, R. M., Eds.; Wiley: New York, 1982. TABLE I V Interaction Site Enthalpies Relative to CCl, compound/site solvent AH,, kcal mol-l phenol/hydroxyl proton CnH 1 2 -0.2 . . . . CiHL +0.9 DMA/carbonyl lone paif C 6 H I2 -1.3 DMF/carbonyl lone paif dioxane/oxygen lone pairb THP/oxygen lone pair" cyclohexanone/oxygen lone paif pyrrole/nitrogen proton pyridine/nitrogen lone pair m-fluorophenol/hydroxyl proton
C6H6 C6H12 C6H6 C6H12 C6H6 C6H12 C6H6 C6H12
C6H6 C6H12 C6H6 C6H12 C6H6 C6H12 C6H6
"Two lone pairs per carbonyl; AH,is s per lone pair. pairs; AH,is per lone pair.
+0.3 -1.1 +0.3
-0.5 +0.1
-0.6 c0.2 -0.9
+OS -0.1 +1.0 -0.9 +0.2 -0.2 +1.0
Four lone
extent the acid and base are solvated by that solvent. Benzene, for example, solvates hydrogen-bonding molecules by formation of an X-H-A bond and is more solvating for these molecules than CC14. A quantification of these effects is possible if the solvent interaction enthalpies are known for the sites involved in the complex formation. The enthalpies due to interactions at the solute sites have been evaluated by a calorimetric method previously described.I4 In brief, the analysis consists of choosing a reference solvent and determining the enthalpy of transfer of a solute from the reference solvent to a second solvent. The enthalpy of transfer is similarly determined for a model solute. The difference in the transfer enthalpies is attributed to interactions with the second solvent relative to the reference solvent. For example, from the enthalpy of solution data given on Table I11 the transfer enthalpy of pyridine from cyclohexane to CC14 can be found to be -1.65 kcal mol-' and that for the same transfer for benzene is -0.80 kcal m01-I.'~ If the difference between the benzene and pyridine transfer is attributed to interactions at the pyridine nitrogen by CC14,this interaction enthalpy relative to cyclohexane is -0.85 kcal mol-(. Nozari and DragoIS have used other methods to assign an interaction enthalpy of -0.9 kcal mol-' to the pyridine-CC14 interaction. Similar considerations with cyclohexane as the model compound were used to obtain interaction enthalpies for THP, cyclohexanone, and dioxane. The pyrrole N-H.-benzene interaction enthalpy was determined directly by the pure base calorimetric method.16 Interaction enthalpies for DMA, DMF, and
~
(9) Huyskens, P.; Mahillon, Ph. Bull. SOC.Chim. Belg. 1980, 89, 701. (IO) Datta, P.; Barrow, G . M. J . Am. Chem. SOC.1965, 87, 3053. (11) Murray-Rust, P.; Glusker, J. P. J . Am. ChemSoc. 1984,106, 1018. (12) Pullman, A. In -Quantum Theory of Chemical Reactions"; Daudel, R., Ed.; D. Reidel Publishing Co.: Boston, 1981; p 14. (13) Sandorfy, C Top. Curr. Chem 1984, 120, 64
(14) Spencer, J. N.; Holmboe, E. S.; Firth, D. W.; Kirshenbaum, M. R . J . Solution Chem. 1981, 10, 745. (15) Nozari, M. S.; Drago, R. S . J . Am Chem. SOC.1972, 94, 6877. (16) Spencer, J. N.; Gleim, J. E.; Blevins, C H.; Garrett, R. C.; Mayer, F J J Phys. Chem. 1979,83, 1249.
The Journal of Physical Chemistry, Vol. 89, No. 10, 1985 1891
Solvent Effects on Hydrogen-Bond Formation
TABLE V Calculated Enthalpies of Complex Formation Relative to CCb -AH, kcal mol-l acid phenol
p yr roIe m-fluorophenol a Reference
base DMA DMF THP cyclohexanone dioxane pyridineb pyridine pyridineb
CCl, exptl
C6H12 calcd
7.2 6.9 5.3 5.2 5.2 6.5 3.9 7.5
8.7 8.2 6.1 6.3 5.9 7.6 4.9 8.6
C6H6
exptl 9.0., 8.9., 9.3.. 8.8.. 8.5" 9.6,9.7,8.1,9.1 5.9,5.8,5.2 6.8,5.2,6.4 6.1,6.0,3.3, 5.7 7.5 5.4,4.3, 5.9,4.3: 5.OC 8.4
calcd
exptl
6.0 5.7 4.2 3.8 4.2 5.4 2.7 6.3
6.3.6.2. 6.2 5.8;5.3;5.7 4.8,4.7,4.4 4.3,4.4,3.6,4.1 4.3,3.9 5.3 3.1, 3.4,3.2 6.3
18. See ref 14. Reference 1.
The interaction phenol have been previously enthalpies and designated sites are given in Table IV. Table V lists the enthalpy changes calculated for the complexes in the solvents cyclohexane and benzene. CC14was chosen as the reference solvent, and in order to use consistently determined enthalpies, the enthalpies of complex formation in CCl, were taken to be those determined by the least-squares calorimetric analysis. As an example of how the interaction enthalpies of Table IV are used to calculate the enthalpy changes given in columns four and six of Table V, consider the phenol-DMA complex: the enthalpy change for this complex in C C 4 is -7.2 kcal mol-'. The formation of the same complex in benzene solvent requires a 0.3 kcal mol-' larger enthalpy input than in CC14 to remove the solvent molecule(s) binding to the one lone pair on the carbonyl group, and 0.9 kcal mol-' more is required to overcome the phenol hydroxyl proton specific hydrogen-bonded interaction. Thus, the enthalpy change for hydrogen-bond formation in benzene solvent should be 1.2 kcal mol-' more positive than that in C C 4 or -6.0 kcal mol-'. The experimental enthalpy changes found in this work are -6.3, -6.2, and 6.2 kcal mol-'. For cyclohexane solvent the enthalpy of formation of the phenol-DMA complex should be 1.5 kcal mol-' more negative (-0.2 for the phenol hydroxyl proton and -1.3 for the DMA carbonyl lone-pair interactions) than that in CC14 or -8.7 kcal mol-'. The experimental results from the present study are -8.8, -8.9, and -9.3 kcal mol-', and the value from the literature is -8.5 kcal mol-'.'* Other enthalpy changes similarly calculated and experimental results are given in Table V. Also included in Table V are two similar systems, phenolpyridine and m-fluorophenol-pyridine, for which sufficient data exist to provide a test for the interaction site model. For both systems agreement between experiment and model is excellent.
Conclusions The simple interaction site model correctly predicts the solvation trends for the enthalpies of complex formation in cyclohexane, CCl,, and benzene solvents. In most cases the enthalpy changes calculated from the model agree with experiment to within the uncertainties of experiment. the interaction site enthalpies for a given solvent generally increase with solute polarity as measured by the Reichardt ETnumbersIg and with the solute donor numbers defined by Gutmanne8 As the availability of electron pairs increases, solvation interactions also increase. It seems likely that for CCl, and benzene solvent these solvating interactions are specific, i.e., of a donor-acceptor type. These interactions are surprisingly large. The carbonyl group of DMA, for example, is stabilized enthalpically in benzene solvent by 3.2 kcal mol-' (two (17)Spencer, J. N.;Berger, S.K.; Powell, C. R.; Henning, B. D.; Furman, G. S.;Loffredo, W. M.; Rydberg, E. M.; Neubert, R. A.; Shoop, C. E.; Blauch, D. N. J. Solution Chem. 1981, 10, 501. (18) Huyskens, P. L.; Cleuren, W.; Van Brabant-Govaerts, H. M.; Vuylsteke, M.A. J. Phys. Chem. 1980,84, 2740. (19) Reichardt, C. "Solvent Effects in Organic Chemistry"; Verlag Chemie: New York, 1979; pp 242-43.
lone-pair interactions) relative to cyclohexane. Obviously the interaction site enthalpies have uses other than for anticipating solvent effects on enthalpy changes for complex formation. It has already been shown that these parameters correlate with absorption freq~encies'~ and that tautomeric equilibria depend on specific solvation.20 Solvent effects on reaction rates and on ESR and NMR spectra should be similarly affected by specific interactions. The use of model compounds to obtain interaction enthalpies seems adequate given the uncertainties which exist in the measurement of most parameters affected by specific interactions. Drago et al.15921have described an "elimination of solvation procedure" (ESP) to detect solvation interactions. The procedure involves the investigation of the enthalpy change for displacement reactions of the type AB
+ B' = AB' -I B
(3)
where A is the acid and B and B' are the bases. The complex AB is assumed not to undergo specific interactions with the solvent. If the active sites on B and B' undergo similar nonspecific interactions, the solvation of reactants and products for eq 3 should cancel. Drago's data show this to be true for a few systems. The data obtained in the present work provide only limited support for this approach. Inspection of the interaction enthalpies given in Table IV shows that the bases do not undergo similar interactions with benzene relative to cyclohexane. An extension of ESP theory allows the conclusion -AH(cyclohexane) = AH(benzene)
+S
where S is a constant for a given acid with a series of bases. The data of Table I may be used to find S for complexes of phenol with various bases. S is 2.6 and 2.4 for the bases DMA and DMF but varies from 0.8 to 2.3 for the bases dioxane, THP, and cyclohexanone. As Drago et al. have pointed out, if specific interactions are present, ESP theory will not give a constant value for S unless additional corrections are made. We conclude that it is unlikely that a single equation containing only one parameter will provide a fit to solvation data in general. The interactions between acid and solvent and between base and solvent are so complex and unique to each solvated species that the best hope for an understanding of solvation effects lies in evaluating each system separately.
Acknowledgment. Acknowledgement is made to the donors of the Petroleum Research Fund, administered by the American Chemical Society, and to Research Corporation through Cottrell College Science Grants program for support of this research. Registry No. DMA, 127-19-5; DMF, 68-12-2; THP, 142-68-7; phenol, 108-95-2; pyrrole, 109-97-7; p-dioxane, 123-91-1; cyclohexanone, 108-94-1;pyridine, 110-86-1. ~
~~
(20) Spencer, J. N.; Holmboe, E. S.;Kinhenbaum, M. R.; Firth, D. W.; Pinto, P. B. Can. J. Chem. 1982,60, 1178. (21) Drago, R. S.; Nozari, M. S.; Vogel, G. C. J. Am. Chem. SOC.1972, 94, 90.
