Solvation Effects on Hydrogen Bonded Systems
2237
Solvation Effects on the Thermodynamics of Hydrogen Bonded Systems. 3 J. N. Spencer,* Judy R. Swelgart, Mlchaei E. Brown, Ronald L. Bensing, Thomas L. Hassinger, William Kelly, Donna L. Housel, G. William Relsinger, Daniel S. Reifsnyder, Jeffrey E. Gleim, and J. Christopher Peiper Department of Chemlstry, Lebanon Valley Col/ege, Annvllle, Pennsylvania 17003 (Received September 22, 1976; Revlsed Manuscript Received July 18, 1977) Publlcatlon costs assisted by the Petroleum Research Fund
Thermodynamic functions for the hydrogen bonded complexes of guaiacol with triethylamine and dimethylformamide in cyclohexane, benzene, and CC14 solvents were determined by infrared spectroscopy. Heats of solution at infinite dilution for phenol, guaiacol, Me2S0,pyridine, and dimethylformamide were determined calorimetrically. Previously reported enthalpies of complex formation for phenol and guaiacol with MezSO and pyridine in cyclohexane, CC14, benzene, CS2,1,2-dichloroethane,and chloroform solvents were combined with the heat of solution data to calculate transfer enthalpies for acid, base, and complex between various solvent pairs. TABLE I: Frequency Data for the Complexes of Introduction Guaiacol with TEA and DMF (cm-l) Drago et al.l“ have described an elimination of solvation Bonded frequency procedure (ESP) for relating enthalpies of hydrogen bond Nonbonded formation in polar solvents to solvation minimized data Solvent frequency TEA DMF for nonpolar solvents such as C C 4 or hexane. These Cyclohexane 3561 3220 3274 investigators report general success but caution that exBenzene 3544 3220 3287 trapolation to new systems cannot be done with conficc1, 3557 3240 3286 dence. In certain cases it has been necessary to infer specific interactions between base and solvent in order to given in an earlier reportg were found to be in error due account for deviations between ESP and e ~ p e r i m e n t . ~ # ~to a transposition of two digits in the molecular mass of Christian et a1.6-s have proposed that the transfer enguaiacol. The corrected enthalpy data and equilibrium thalpy of a given complex between any two solvents is constants are given in Table 11. Woolley et al.15 have proportional to the sum of the transfer enthalpies of the pointed out that when a standard state based on molarity acid and the base. is used, a direct application of the van’t Hoff equation The present investigation was undertaken to determine cannot be made. The enthalpy change must be calculated if thermodynamic evidence for certain specific interactions from proposed by ESP adherents could be detected. From heat of solution data transfer enthalpies may be calculated for all species and used for an examination of Christian’s proposal. where a is the coefficient of thermal expansion of the Experimental Section solvent. At 298 K this correction is about 0.2 kcal/mol-’ Purifications of the phenols and solvents have been for the solvents of this work. All the enthalpies in Table previously rep~rted.~-llMallinckrodt SpectrAR grade I1 have been calculated according to eq 1. dimethylformamide was first “benzene dried”12 then The calorimeter design was similar to that reported by further dried over Drierite before distilling under a niArnett et al.16 The calorimeter vessel is a standard 250-mL trogen atmosphere. Eastman practical grade triethylamine silvered Dewar flask immersed in a thermostatted water was distilled under nitrogen from barium oxide. The index jacket. The temperature sensing elements are two YSI of refraction was used as a purity criterion. 44011 100-kQthermistors encased in No. 26 Teflon tubing. TEA is known to interact with CCl,.13J4 Therefore TEA The stirrer was driven by a Barber-Colman synchronous and CC4were distilled immediately prior to use. Care was motor. Actual time-temperature data were recorded by taken to exclude oxygen and light. Solutions were stable interfacing the temperature sensing circuit to a Hewfor periods of 2-3 h. After about 3 h a precipitate of thin lett-Packard Model 2100A microprogrammable computer. needles was observed, apparently similar to that reported The computer program used a least-squares criterion to by Drago et al.5 in solutions of pyridine or quinuclidine establish preinjection, post-injection, and heating base in methylene chloride. lines. Frequency shifts were obtained from spectra recorded Solid samples were injected into the Dewar by a delivery in 2-mm path length cells and are given in Table I. The system similar to that used by Arnett et a1.16 Liquid complex absorption bands are broad for all systems. The samples were injected from syringes immersed in the assignment of the complex frequency for DMF systems in solvent in the Dewar. Sample quantities of 1-10 mmol cyclohexane are less reliable due to solubility problems. were injected. Extrapolated heats obtained from repeated For the determination of the equilibrium constants, sets of successive injections were averaged and the preguaiacol concentrations were about 0.0015 M, TEA concision reported in Table I11 is the standard deviation of centrations were about 0.25 M. Higher concentrations a single measurement. were used to determine the frequency shifts. Results and Discussion The enthalpies of formation for certain complexes were taken from previously reported spectrophotometric The enthalpies of hydrogen bond formation and the st~dies.~JO The data for the guaiacol-pyridine complexes equilibrium constants for the complexes of guaiacol with The Journal of Physical Chemistry, Vol. 81, No. 24, 1977
Spencer et al.
2238
TABLE 11: Enthalpies of Hydrogen Bond Formationa
AH" 298, kcal mol-' Me,SO Solvent
Pyridine
Phenol
Cyclohexane Carbon disulfide cc1, Benzene 1,2-Dichloroethane Chloroform
-9.18 -7.29 -6.55 -5.36 -6.28 -3.34
i. i: i. i. i. t
Guaiacol
0.25 0.24 0.07 0.12 0.14 0.03
Phenol
-3.99 i 0.25 -4.71 t 0.05 -3.26 t 0.14 -4.10 t 0.06 -3.64 t 0.09 -2.31 i 0.18
-7.49 -6.18 -5.92 -5.25 -5.76 -5.32
K293
GuaiacolC
0.34 * 0.41 t 0.29 i 0.32 i. 0.15 t 0.12
* 0.07
-3.35 -3.95 -3.29 -4.33 -3.33 -2.75
?
t t i t t
Data from ref 9 and 10. Corrected equilibrium constants erroneously reported in ref 9, vide infra. thalpies erroneously reported in ref 9, vide infra.
(z kcal , mol-') ,at 25
Solvent
Phenol
Guaiacol
Cyclohexane Carbon disulfide cc1, Benzene 1,2-Dichloroethane Chloroform
t 7 . 6 2 5 t 0.02a t 6 . 9 0 t 0.06 t 6 . 2 7 i: 0.07b t 4 . 7 2 t O.Olb t 4 . 5 9 i 0.18 t 4 . 4 7 * 0.04
t 3 . 5 6 i 0.05 t 2 . 6 0 t 0.06 t 2 . 0 5 0.03 t 1 . 3 5 i. 0.07 t 1 . 0 1 i. 0.12 t 0 . 1 5 t 0.06
a
Corrected en-
C
Me,SO t4.18 t t 3 . 0 2 i. t1.83c 1-1-15i -0.31e -3.30 t
*
2.71 2.67 1.44 1.72 2.10 1.90
0.13 0.38 0.15 0.41 0.30
a
TABLE 111: Partial Molar Heats of Solution
guaiacolpyridine
0.21 0.02 0.03d
Pvridine
DMF
t2.01e t 1 . 2 5 t 0.01 t 0.36 t 0.01f 0.00 t 0.01f
t 3.65 i 0.07
t 0.26
t 0.76 tO.09
- 0.64 i. 0.08g
I.
t
TEA
0.029 0.02
-0.54
i
t
0.04b 0.04d
t 0.01e
0.01
-1.95
a Reference 20. Reference 21. Average of data from ref 23 and 2. precision not given. f Reference 1. g Reference 23.
i.
0.01
Reference 2.
e
Calculated from ref 3,
TABLE IV: Thermodynamic Data for the Complexes of TEA and DMF with Guaiacol - A Ha,a kcal mol-'
Solvent
DMF
Cyclohexane Benzene CCl,
4.64 * 0.08 3.50 t 0.25 3.74 i. 0.18
- A S " , b cal deg'l mol-'
TEA 4.92 6.00 5.39
t
* ?I
0.11 0.16 0.16
KlOOC
DMF
TEA
DMF
TEA
- 13.4
- 14.9 - 19.0
3.40 1.06 1.38
2.06 1.56
-11.8 - 12.1
-17.5
a The error given was obtained from the slope of a least-squares plot of In K vs. T-'. T A S " at 20 "C.
2.58
Calculated from A G " = AH"
-
TABLE V: Enthalpies of Complex Formation for Various Bases with Guaiacol - AH" , kcal mol-'
Solvent Cyclohexane Benzene cc1,
DMF 4.64 3.50 3.74
t t t
0.08 0.25 0.18
TEA 4.92 6.00 5.39
DMF and TEA in the solvents of this study are given in Table IV. Enthalpies previously r e p ~ r t e d ~for ~ ' ~the ~~' pyridine and MezSO complexes with phenol and guaiacol are given in Table 11. According to ESP, pyridine interacts with CCl,, benzene, and o-dichlorobenzene and TEA interacts with CCl, but not with aromatic solvent^.^ ESP workers have also suggested that oxygen containing proton acceptors form aggregates in cyclohexane which solvate the adduct causing the enthalpy of hydrogen bond formation to be higher in cyclohexane than in CC14.3,13,18 According to Drago and Nozari3 neither TEA nor pyridine aggregate significantly in cyclohexane but because of their interaction with CC14, the enthalpy of complex formation with these bases is lower in CC14 than in cyclohexane. Several investigator^^^-^^ have proposed an interaction of the phenolic hydroxyl proton with C C 4 to account for differences found in enthalpies determined in CC4 and other solvents. Because specific interaction of the hydroxyl hydrogen with the solvent is minimized by the intramolecular bond of guaiacol, the bulk of the solvent effects in guaiacol systems should be due to interactions of the solvent with base, adduct, or both. The data in Table V, show that the enthalpy is larger for the MezSO and DMF complexes in cyclohexane than in CC4. According to Drago et al.,2,13,18 aggregation of these The Journal of Physical Chemktty, Vol. 81, No. 24, 1977
f
i i.
Pyridine 3.35 i 0.07 4.33 f 0.15 3.29 i: 0.38
0.11 0.16 0.16
Me,SO 3.99 4.10 3.26
* 0.37 t i
0.06 0.14
oxygen containing donors about the complex could cause the enthalpy in cyclohexane to be larger than that in CC14. If aggregation of DMF occurs only in cyclohexane and if other solvent effects are not thermodynamically significant, the enthalpy of complex formation with DMF should be nearly the same in benzene and CC14 Within error limits, the enthalpies are the same in these latter two solvents. If aggregation of MezSO in cyclohexane is to be accepted, the solvation of MezSO, its adduct, or both by benzene must be different from that of DMF. Only for DMF in cyclohexane and in the comparison of TEA complexes in benzene and CC4do the predided base interactions of ESP seem to be supported. Thus other solvation effects must be supposed in order to provide a generally satisfactory treatment for these complexes. In order to attempt to identify these effects, partial molar heats of solution were determined and combined with existing literature values so that enthalpies of transfer for the various complexes could be calculated. The following thermodynamic cycle was used to calculate the transfer enthalpies from solvent I to solvent 11. A(1)
. H 1 A(I1)
+
1%
B(1)
t B(I1)
+ A.B(I) ]AHc FZ
A*B(II)
(2)
Solvation Effects on Hydrogen Bonded Systems
2239
TABLE VI: Heats of Transfer from Cyclohexane to Other Solvents (kcal mol-') Solvent Phenol Guaiacol Pyridine Me,SQ Phenol-pyridine Phenol-Me,SQ Guaiacol-pyridine Guaiacol-Me,SO TEA DMF Guaiacol-TEA Guaiacol-DMF a
CS, -0.73 i 0.21 -0.96 i 0.11 -0.76 i: 0.02' -1.16 i 0.23 -0.18 i 0.97 t 0.01 +- 0.93 - 2.32 i 0.32 -2.84 c 0.85
cc1, -1.36 i 0.09 -1.51 i: 0.08 -1.65 c 0.02' -2.35 i 0.22' -1.44 c 0.73 -1.08 i 0.62 -3.10 i 0.54 -3.13 i 0.51 -0.90 i 0.12 -2.89 c 0.09 -2.88 i 0.47 -3.50 i 0.43
C6 H6 -2.91 i 0.03 -2.21 i 0.12 -2.01 i 0.02' -3.03 i 0.24 -2.68 c 0.93 -2.12 t 0.65 -5.20 i 0.35 -5.35 i 1.02 t 0.28 i 0.08 -3.46 i 0.09 -3.01 i 0.47 -4.63 i 0.54
C,H,Cl, -3.04 i 0.20 -2.55 i 0.17 -2.00 i 0.02' -4.49 i 0.22a -3.31 i 0.69 -4.63 f 0.80 -4.53 f 0.65 -6.69 i 0.84
CHC1, -3.16 i 0.06 -3.41 i 0.11 -3.96 i 0.02' -7.48 * 0.22 -4.95 f 0.53 -4.80 i 0.56 -6.77 i 0.49 -9.21 * 0.88
Estimated error.
Table VI lists the enthalpies of transfer of acid, base, and adduct from cyclohexane to all other solvents. These transfer enthalpies relate the solvation of each species in the different solvents to that in the relatively inert cyclohexane. Thus Table VI is arranged in order of increasing solvation of each species. The relative solvation of the adducts increases in the same order as that of the acid and base. The MezSO complex of guaiacol in 1,2dichloroethane and chloroform is more solvated relative to cyclohexane than the guaiacol-pyridine complex. This effect appears to be due to the solvation of MezSO being greater than that of pyridine in 1,2-dichloroethane and chloroform. In the other three solvents, CSz, CC14, and benzene, the differences in solvation of MezSO and pyridine relative to cyclohexane are not as great as with 1,2-dichloroethane and chloroform. Consequently the difference between the solvation of the guaiacol-pyridine and the guaiacol-Me2S0 complexes is not as pronounced in these three solvents. Similar conclusions may be applied to the phenol complexes for CSz, CC14, benzene, and possibly 1,Zdichloroethane. The phenol-CHC13 data do not conform to this analysis. Relative to cyclohexane any difference between the transfer enthalpies of the phenol-pyridine and guaiacol-pyridine or the phenol-Me2S0 and guaiacol-Me2S0 adducts must be due to differences in solvation of the acid part of the adduct. With the exception of benzene, the enthalpies of transfer of phenol and guaiacol from cyclohexane to the other solvents are nearly equal, however, the transfer enthalpies of the phenol-MepSO and guaiacol-Mea0 adducts differ by more than would be expected if only the phenol and guaiacol transfer enthalpies are considered. The discrepancy can be attributed to the solvation of the guaiacol methoxy group which is freed upon complex formation providing an additional solvation site. Similar considerations apply to the phenol and guaiacol complexes with pyridine. In cyclohexane, CS2,CC4, and chloroform, the value of
-
AH, (phenol) -
z, (guaiacol)
is relatively constant at 4.23 f 0.11 kcal mol-l. If the intramolecular bond of guaiacol effectively squeezes out solvent interaction sites so that the number of sites for solvent interactions are about equal for phenol and guaiacol, the solvation differences between phenol and guaiacol would not be expected to differ greatly in the absence of strong specific effects. For benzene and 1,2dichloroethane, a B ( p h e n o l )- z8(guaiacol) is 3.37 and 3.58 kcal mol-l, respectively. In benzene solvent, this results from the well-known interaction of the phenol hydroxyl proton with the a cloud. Although a similar explanation for 1,2-dichloroethane would seem to be implied, Drago et al.3*5have argued that 1,2-dichloroethane
cannot behave as a basic solvent and have given a molecular interpretation to the unusual solvent effects observed in 1,2-dichloroethane. The heats of solution indicate that pyridine behaves similarly in 1,Zdichloroethane and in benzene. Phenol also appears to behave similarly toward these two solvents. Since both the acid and base show similar trends, 1,2-dichloroethaneappears to behave in an unusual manner, supporting Drago's contention that solvent interactions in 1,2-dichloroethane are nonspecific. Christian et ala6 have proposed that the transfer enthalpy of the adduct should be proportional to the enthalpies of transfer for acid and base
Afi!c= &(AHAt AH,)
(3)
where a is a constant not strongly dependent upon temperature or solvent pairs. Evaluation of a should allow inferences about the physical nature of the adduct and the adduct interaction with the solvent. Weak complexes should have a less than unity because at least one solvent molecule has been lost from each solvation shell about the acid and upon complex formation. An a value in excess of unity indicates a strong complex and a large interaction between the dipole of the complex and the medium. Drago et aL5have argued that one possible reason for the success of ESP is that a slight increase in the nonspecific solvation of the complex could compensate for the loss af solvent molecules when complexation occurs. According to eq 3,a is the slope of the plot of AHc vs. AHA + AHB. The data in Tables 11,111, and IV were used to calculate a for each acid-base pair by least-squares analysis. The least-squares plots were not forced through the origin. This value of a,along with AHAand AHBfrom Table 11,was then used in eq 2 for each acid-base pair to calculate AHc for each possible solvent transfer. Although the trends in solvation of the adduct followed those predicted by eq 3,agreement between AHc found from the cycle and that calculated by eq 3 was no better than about 0.5 kcal mol-l.
Acknowledgment. Acknowledgment is made to the Donors of the Petroleum Research Fund, administered by the American Chemical Society, for partial support of this research and to the Cottrell College Science Grants Program of the Research Corporation. References and Notes (1) R. M. Guidry and R. S. Drago, J . Phys. Chem., 7 8 , 454 (1974). (2) R. S. Drago, M. S. Nozari, and G. C. Vogel, J. Am. Chem. Soc., 94, 90 (1972). (3) M. S. Nozarl and R. S. Drago, J. Am. Chem. Soc., 94, 6877 (1972). (4) M. S. Nozari, C. D. Jensen, and R. S.Drago, J . Am. Chem. Soc., 95. 3162 (1973). (5) R. S. Drago, A.'Nusz, and R. C . Courtright, J. Am. Chem. Soc., 96, 2082 (1974). The Journal of Physical Chemlstiy, Vol. 81, No. 24, 1977
J. Pacansky
2240 (6) S.D. Christian, J. R. Johnson, H. E. Affsprung, and P. J. Kilpatrick, J. Phys. Chem., 70, 3376 (1966). (7) J. Grundnes and S. D. Christian, Acta Chem. Scand., 23, 3583 (1969). (8) S.D. Christian and E. E. Tucker, J . Phys. Chem., 74, 214 (1970). (9) J. N. Spencer, J. R. Sweigart, M. E. Brown, R. L. Bensing, T. L. Hassinger, W. Kelly, D. L. Housel, and G.W. Reisinger, J . Phys. Chem., 80, 811 (1976). (10) J. N. Spencer, R. S. Harner, and C. D. Penturelli, J. Phys. Chem., 79, 2488 (1975). (1 1) J. N. Spencer, R. A. Heckman, R. S. Harner, S. L. Shoop, and K. S. Robertson, J. Phys. Chem., 77, 3103 (1973). (12) A. B. Thomas and E. C. Rochow, J. Am. Chem. SOC.,79, 1843 (1957). (13) R. S.Drago and T. D. Epley, J . Am. Chem. SOC.,91, 2883 (1969). (14) W. J. Lautenberger, E. N. Jones, and J. G. Miller, J. Am. Chem. SOC., 90, 1110 (1968).
(15) E. M. Woolley, J. G. Travers, 8. P. Erno, and L. G. Hepler, J . Phys. Chem., 75, 3591 (1971). (16) E. M. Arnett, W. G. Bentrude, J. J. Burke, and P. McC. Duggieby, J. Am. Chem. SOC.,87, 1541 (1965). (17) J. N. Spencer, K. S. Robertson, and E. E. Quick, J . Phys. Chem., 78, 2236 (1974). (18) W. Partenheimer, T. D. Epley, and R. S. Drago, J. Am. Chem. Soc., 90, 3886 (1968). (19) T. Gramstad and W. J. Fuglevik, Acta Chem. Scand., 18, 6 (1962). (20) W. C. Duer and G. L. Bertrand, J. Am. Chem. SOC.,92, 2587 (1970). (21) E. M. Arnett, L. Joris, E. Mitchell, T. S. S. R. Murty, T. M. Gorrie, and P. v. R. Schieyer, J. Am. Chem. SOC.,92, 2365 (1970). (22) A. N. Fletcher and C. A. Heller, J. Phys. Chem., 71, 3742 (1967). (23) E. M. Arnett, E. J. Mitchell, and T. S. S. R. Murty, J. Am. Chem. SOC.,96, 3875 (1974).
The Infrared Spectrum of a Molecular Aggregate. The HCN Dimer Isolated in an Argon Matrix J. Pacansky IBM Research Laboratory, San Jose, California 95 193 (Received June 9, 1977) Publlcation costs assisted by
IBM Research laboratory
Spectral properties of well-defined molecular aggregates are difficult to obtain. A study is presented which demonstrates that a well-defined molecular aggregate may be produced by photochemical methods. This is specifically shown by producing the HCN dimer from the photochemical decomposition of s-tetrazine. The infrared spectrum of the dimer clearly shows that the vibrational bands associated with the hydrogen bond are intensified and broadened in comparison to those not involved in the hydrogen bond.
Introduction The matrix isolation technique has amply demonstrated that spectroscopic properties of isolated molecules may be readily 0btained.l Here, and in subsequent studies, we show that spectroscopic properties of well-defined aggregates of molecules may also be observed utilizing the matrix isolation technique. This is demonstrated by showing the infrared spectrum of the matrix isolated HCN dimer. The unique feature of this spectrum and the manner in which the dimer was formed is that all of the sites in the matrix are occupied by only dimers. The spectra thus clearly reveal the spectroscopic changes which are manifestations of the rather strong hydrogen bond2 (AH= 3.8 kcal/mol) in the dimer. The method used to produce molecular aggregates consist of selecting a system that upon exposure to light photochemically decomposes to the particular aggregate desired. For example, tetrazine3i4has been studied in both the gas and solid phase and appears to quantitatively decompose when exposed to light with wavelength h
