3352
J. Phys. Chem. 1981, 85, 3352-3354
Hydrogen Bonding between Pentafluorobenzene and Pyridlne-d,‘ Joseph R. Murdoch’ and Andrew Streitwieser, Jr.” Department of Chemistry, University of California, Berkeley, California 94720 (Received: May 22, 198 1; In Final Form: JuV 28, 198 1)
Pentafluorobenzene forms a relatively strong hydrogen bond to pyridine-d, as measured by infrared (Av = 79 cm-’) and NMR techniques; the formation constant is 0.36 M-’ in CC14. The infrared Av values apparently correlate with equilibrium acidities for halogenated benzenes but not for other groups of compounds; in particular, the correlation does not apply to vinyl C-H compounds.
Introduction Many C-H bonds are now known to form hydrogen bonds, and the subject has been thoroughly r e ~ i e w e d . ~ Most of the C-H bonds involved in these examples are those of acetylenes or halogenated methanes, compounds that are also relatively strong carbon acids. Few benzenoid C-H bonds have been involved in such hydrogen-bonding studies, and apparently only a single limited study4has included pentafluorobenzene (PFB), a rather strong carbon acid; for example, a recent survey of C-H proton donors did not include PFBm6Infrared spectroscopy and NMR spectroscopy are among the most useful techniques for measuring hydrogen b~nding.~,’-’~ We have used both in a study of hydrogen bonding between PFB and pyridine-d, (PD5). We report the frequency shift, the half-bandwidths of the C-H absorptions, and the ratio of the integrated intensities for the associated and the free C-H stretching modes (RAF).In addition, the equilibrium constant for 1:l complex formation has been calculated from IR and NMR measurements in CC1, solution at 35 “C. Experimental Section Purification of Materials. CC1,. Reagent-grade (Baker) carbon tetrachloride was dried over P205for 24 h and vacuum distilled onto molecular sieves which had been previously dried and degassed. The container was kept under positive argon pressure, and solvent was removed by using a gas-tight syringe and an argon purge. Pyridine-d,. The pyridine-d6 (PD5) was commercially available and had a deuterium content of -99%. The material was vacuum transferred from dried, degassed molecular sieves and stored in the original container over molecular sieves. Purity of the substance was verified by GC (Carbowax and SE-30). Pentafluorobenzene. Pentafluorobenzene (Aldrich) was purified and stored in a similar manner. Purity was ver(1) Dedicated to Professor Joel Hildebrand in honor of his 100th birthday. (2) Department of Chemistry, University of California, Los Angeles, CA 90024. (3) Green, R. D. “Hydrogen Bonding by C-H Groups”; Wiley: New York, 1974. (4) Dale, A. J. Acta Chem. Scand. 1970,24, 3403. (5) Streitwieser, A., Jr.; Nebenzahl, C. N.; Scannon, P. J.; Niemeyer, H . M.J. Am. Chem. SOC.,submitted. (6) Slasinski, F. M.; Tustin, J. M.; Sweeney, F. J.; Armstrong, A. M.; Ahmed, Q . A,; Lorand, J. P. J. Org. Chem. 1976,41, 2693. (7)Pimentel, G. C.; McClellan, A. L. “The Hydrogen Bond”; W. H. Freeman: San Francisco, 1960. (8) Hamilton, W. C.; Ibers, J. A. “Hydrogen Bonding in Solids”; W. A. Benjamin: New York, 1968. (9) Murthy, A. S. N.; Rao, C. N. R. AppZ. Spectrosc. Reu. 1968,2,69. (IO) Vinogradov, S. N.; Linnell, R. H. “Hydrogen Bonding”; Van Nostrand-Reinhold New York, 1971. (11) Pimentel, G. C.;McClellan, A. L. Adu. Phys. Chem. 1971, 347. (12) Joesten, M. D.; Schaad, L. J. “Hydrogen Bonding”; Marcel Dekker: New York, 1974.
TABLE I: Infrared Spectral Data
pentafluorobenzene C-H stretch V , cm-I
free
3095 * 3
Av, cm-’
v ~ , ~cm-’ :
R*Fb
17 f 2 1.0
hydrogen bonded
3016 i 3 79 65 * 3 2.56 f 0.30
a Half-bandwidth. Relative areas of C-H absorption of associated species and free species, normalized to unit concentration and path length.
ified by both NMR and GC (Carbowax and SE-30). All compounds were stored in a desiccator in the dark at 25 “C. Infrared measurements were carried out on a PerkinElmer Model 421 grating spectrophotometer and a Beckman IR-12. Spectra were generally run without solvent blanks, and, if correction for absorption due to PD5 was necessary, this was accomplished by subtracting the appropriate absorbance, making proper allowance for the concentration of PD5, the cell path length, and the extinction coefficients of PD5. Measurements were carried out by using a cell with a fixed path length (1.584 mm) and a variable path length cell with path lengths of either 0.150 or 1.000 mm. The temperature was measured by a copper-constantan thermocouple attached to the cell and referenced to ice water. NMR measurements were carried out on a Varian A-60. Chemical shifts are accurate to -0.5 Hz and were measured relative to cyclohexane which was added to the solutions in a trace amount. The sample temperature was measured by using a copper-constantan thermocouple. All IR and NMR measurements were made at 35 f 2 OC. The concentration of each component in mol/L was calculated from the density of the solution and the weight of each component. No correction has been applied to account for expansion of the solution at 35 “C. Results and Discussion When gradually increasing amounts of PD5 are added to a solution of PFB in CC14,a new band appears at lower frequency (3016 cm-l) and increases in size, while the high-frequency band (3095 cm-l) decreases. The low-frequency band is significantly broader and has a greater integrated intensity than the band at 3095 cm-‘ (normalized to unit concentration and path length). Other than the decrease in intensity and possible slight broadening as increasing amounts of PD5 are added, no change is observed in the appearance of the high-frequency band. The pertinent spectra are illustrated in Figure 1. From this behavior it is apparent that PFB forms a hydrogen-bonded complex with pyridine. Furthermore, it was found that in pure pyridine essentially all of the PFB could be converted to the hydrogen-bonded form. The IR
0022-3654/81/2085-3352$01.25/00 1981 American Chemical Society
The Journal of Physical Chemistry, Vol. 85, No. 22, 198 1 3353
Hydrogen Bonding between PFB and PD5
TABLE 11: Hydrogen Bonding for PD5 and PFB ~
I V I
[ I
3120
3080
I
I
3040
I
I
3000
I
lJ 2960
c m-'
Flgure 1. Effect of pyridine on C-H stretching mode of PFB: (heavy line) PFB in CCl, (1.59 M), 0.15-mm path length; (medium line) PFB (0.88 M) and PO5 (4.87 M) in CCI,, 0.15-mm path length; (light line) PFB in pyridine-d, (PD5) (0.183 M), 1.0-mm path length.
spectral data are listed in Table I. Equilibrium constants were calculated from the IR data by assuming that the relative extinction coefficients (at 3095 and 3016 cm-l) of hydrogen-bonded PFB and free PFB in pure PD5 (11.75 M PD5,0.183 M PFB) were the same as those in dilute PD5/CC14 mixtures. The concentrations of free pentafluorobenzene and complexed pentafluorobenzene were calculated from the extinction coefficient data obtained in pure CC14and pure PD5, the relative absorbancesat 3095 and 3016 cm-l, and the formal concentration of PD5 and PFB. NMR chemical shifts are always relative to internal cyclohexane. It was found that pentafluorobenzene gave a limiting chemical shift in concentrated solutions of PD5. For example, the difference in chemical shift of the PFB proton in pure CC14and in 0.340 M PFB/11.574 M PD5 was 15.5 Hz. An identical value was found in 0.936 M PFB and 10.954 M PD5. The chemical shift of PFB in CC14was also found to be concentration independent up to at least 1.7 M PFB in C C 4 solution. The fraction of complexed PFB (FA) was calculated' from F A = (W - WB)/( WA - WB), where W is the proton resonance frequency in a given mixture of PFB/PD5/CC14, WB is the frequency of the free species, and WAis the frequency of the hydrogen-bonded species. WB is taken to be the C-H resonance frequency in pure CCI, relative to internal cyclohexane, while WA is assumed to be identical with the limiting frequency observed in pure PD5, relative to internal cyclohexane. This latter assumption is justified provided that the concentration of hydrogenbonded complex greatly exceeds that of the free species in concentrated solutions of PD5. This appears to be valid, judging by the IR data as well as by the relatively large ratio of PFB/PD5 at which the limiting chemical shift was measured. Equilibrium constants were then calculated from the fraction of complexed PFB and the formal concentrations of PFB and PD5 by using eq 1. The results are summarized in Table 11. FAIPFBIO
K=
(1)
([PFBlo - F.A[PFBIO)([PD~IO - FAU'FBIO) Except for the first two runs in Table 11, all of the K's are remarkably constant. The average equilibrium constant is 0.36 0.12 M-' for runs spanning a sixfold variation in [PD5Ioand a threefold variation in [PFBIo. The constancy plus the excellent agreement between the NMR and IR techniques provide compelling support for the assumption of a 1:l complex and also support the assumptions that the relative extinction coefficients of the
*
~ o,
run no.
M
l-NMRb 7-NMRb 6-IR 11-NMR 10-NMR 9-NMR 4-IR 2-IR 8-NMR 5-IR
9.470 8.033 4.874 4.801 3.943 3.246 2.91 2.59 1.135 0.771
5[PFBI1 o, M K,M-' 1.757 2.821 0.882 0.556 0.388 0.603 0.308 0.521 0.574 0.662 K,=
range of Ka
1.8 0.8-3.7 1.2 0.7-2.6 0.35 0.2-0.6 0.4-0.9 0.63 0.31 0.2-0.4 0.30 0.2-0.4 0.39 0.2-0.6 0.36 0.2-0.5 0.25 0.1-0.4 0.28 0.1-0.4 0.36 k 0.12c
a This is the maximum and the minimum value of K that can be calculated from the estimated uncertainty (assuming normal error propagation of ex erimental quantities) in the fraction of complex, F A . No CCl,. First two entries are omitted. This equilibrium constant can be converted to mole fraction units (MF) by multiplying by a number of moles of all components per liter of solution; this factor is -10-11.
IR absorptions at 3095 and 3016 cm-l are solvent independent and that the chemical shift (relative to internal cyclohexane) of the proton in the hydrogen-bonded complex is solvent independent. The data suggest that the equilibrium constant may be larger in the absence of CC14 Since chloroform is known to be self-associatedg and equilibrium constants for hydrogen-bond formation are generally higher in cyclohexane than in CC14? it is possible that PFB is capable of complexing with CC4. A larger equilibrium constant in pure pyridine is also consistent with the observation that there is essentially no absorbance near 3095 cm-' that could be attributed to free pentafluorobenzene. In 11.5 M PD5 and with K = 0.36, the ratio of complexed to free PFB would be ca. 41, and the intensity of the C-H stretch at 3095 cm-' should be 20% of that observed if the PFB were in pure CC1,. If K = 1.7 (run 1,Table 11),then over 95% of the PFB would be hydrogen bonded. The close agreement of the IR and NMR measurements over a fairly wide concentration range probably rules out the existence of two or more distinct pyridine complexes, since it is highly unlikely that both would have similar dissociation constants and also lead to the same dramatic perturbations of the C-H stretching frequency. Furthermore, the visible spectra (380-700 nm) of PD5/PFB solutions do not reveal the development of any new absorption bands, and freshly prepared solutions are always colorless. In addition, the NMR spectra of PFB in PD5 and in CC14are very similar (except for the small frequency shift, WA - WB) and do not change with time. These latter observations suggest that nucleophilic addition of PD5 to PFB is not making a significant contribution to the observed spectral changes. Finally, if one assumes a stoichiometry other than 1:l (e.g., 1:2 or 2:1), the calculated values of K are no longer constant. Chloroform appears to be a weaker hydrogen donor than PFB with respect to pyridine. For example, Findlay13has found that the chloroform/pyridine association constant is 1.3 (MF)-l (40 " C ) compared with -3.6 (MF)-' found here (see footnote c, Table 11) for PFB/pyridine. However, Dale4has reported that chloroform is a stronger donor than PFB when the hydrogen acceptor is hexamethylphosphoric triamide. He reports that equilibrium constants (2.3614 (13) Findlay, T. J. V.; Keniry, J. S.; Kidman, A. D.; Pickles, V. A. Trans. Faraday SOC.1967, 63, 846. (14) Gramstad,T.;Mundheim, 0. Spectrochirn. Acta, Part A 1972,28, 1405.
3354
The Journal of Physlcal Chemlstty, Vol. 85, No. 22, 1981
TABLE 111: Infrared Frequency Shifts of Proton Donors with Pyridine compd ventafluorobenzene 1,3,54richlorobenzene 1,2,4,54etrabromobenzene 1,2,4,54etrachlorobenzene l-nitro-2,3,5,6tetrachlorobenzene 3,5d i n i t r o-2,4,6trichlorobenzene trans-CHC1= CHCl &CHCl=CHCl CF,=CFH CCl,=CClH CBr,=CBrH CCl,=C(CCl,)H HC=CH PhCECH t-BuC=CH BrCH,C=CH CDCI, CHC1, CHI, CHBr,
A U , cm-I ~
pKcsCHAh
7gb 34
25.6 (25.8) 33.1 ( 3 3 . 6 j
51
30.3
42
31.8 (31.1)
45
31.3
53
29.9
34 23 22 45 52 64 85c 97d 7 9e 110 36f 46
33.1 35.0
55 66
35.2 31.3 30.1 28.1 24.6 22.6 (23.2)g 25.6 (25.5)Z 20.4 32.8 31.1 29.6 27.8
This work. Reference 1 9 unless otherwise indicated. Shuvalova, E. V. Opt. Spectrosc. (Engl. Transl.) 1 9 5 9 , 6, 453. West, R.; Kraihanzel, C. S. J. Am. Chem. SOC. 1 9 6 1 , 8 3 , 765. e Gastilovich, E. A.; Shigorin, D. N. Russ. Chem. Rev. (Engl. Transl.) 1 9 7 3 , 4 2 , 611. f Lord, R. C.; Nolin, B.; Stidham, H. D. J. Am. Chem. SOC.1955, 77, 1365. ~ K L ~ C I ; I A . Calculated from eq 2; experimental values are given in parentheses.
vs. 0.534M-l at 35 “C) and AH (2.714 vs. 1.74kcal/mol) are both larger for chloroform. Taken at face value, these results suggest that complex formation cannot be described in terms of additive properties of the donor and the acceptor. Such a possibility is not surprising in view of Drago’s finding that the enthalpy of adduct formation for a number of acid-base complexes can be correlated with an equation of the form pH = CACB EAEB where C and E represent “covalent” and “electrostatic” contributions of the acid or the base.15 Drago’s parameters for pentafluorobenzene are not available, and a more quantitative analysis is not presently feasible. Moreover, it has been conclusively demonstrated that there is no general correlation between infrared shifts and equilibrium constants or AH for hydrogen bonding.6,16
+
Murdoch and Streitwieser
Within families of related compounds, however, the hydrogen-bonding basicity of acceptors tends to parallel proton-transfer basicities.17J8 Accordingly, we next explore related parallels between acidity and hydrogen bonding of proton donors. Allerhand and Schleye@ have reported IR Au for several halogenated benzenes with pyridine (Table 111). Ion-pair pK‘s (PKCsCHA)’ are available for PFB (25.8), 1,2,4,5C6C14H2(31.1), and 1,3,5-C&Cl, (33.6). These three compounds give a linear relationship between Av and PKC~CHA (eq 2). The pKcScmvalues calculated from obAu (cm-l) = 231 - 5.96pKCsCHA (2) served shifts as summarized in Table I11 show reasonable agreement over a pK range of 8 units. Such a correlation is perhaps not unexpected for such a group of closely related compounds in which the conjugate base in a highly localized carbanion. The correlation is defined by only three points, but the value derived for 1,2,4,5-C6Br4H2 is not unreasonable. Perhaps more surprising is the agreement shown in Table I11 for acetylenes since the C-H bond involved now has different hybridization.18 We note that the experimental acidities listed for acetylenes are actually Accordingly, this those for lithium as gegenion, pKLi~HA.20 agreement may well be fortuitous. For the halomethanes, on the other hand, the values derived from Au are substantially higher than pEcs estimated from kinetic acidities and a Bronsted correlation established for fluorinated bicyclics.21 These sp3 C-H bonds clearly do not belong on the correlation of eq 2. Finally, we inquire about the vinyl systems since the C-H bond involved is approxivalues derived from eq mately sp2. In fact, the pKCsCHA 2 are substantially higher (8-9 pK units) than those derived from kinetic acidities and Bronsted correlations.6tZ1 The presence of an a halogen in most of these cases may have a perturbing effect on Av. Nevertheless, the “family” of compounds to which eq 2 applies appears to be a rather restricted one, and measurements of hydrogen bonding have only limited application to determining carbon acidities. Acknowledgment. This research was supported in part by USPH grant GM-12855 and by NSF grant CHE7910814. (17)Taft, R. W.;Gurka, D.; Joris, L.; Schleyer, P. V. R.; Rakshys, J. W. J. Am. Chem. SOC.1969,91,4801.Gurka, D.;Taft, R. W. E d . 1969, 91,4794. (la)Kamlet, M. 3.; Solomonovici, A.; Taft, R. W. J.Am. Chem. Soc. 1979,101, 3734. (19)Allerhand, A.; von Rague Schleyer, P. J. Am. Chem. SOC.1963,
ati,im.
(15)Slejko, F.L.;Drago, R. S.; Brown, D. G. J. Am. Chem. SOC. 1972, 94,9210. (16)Arnett, E. M.;Mitchell, E. J.; Murty, T. S. S. R. J. Am. Chem. SOC.1974,96,3875.
(20)Streitwieser, A,, Jr.; Reuben, D. M. E. J. Am. Chem. SOC.1971, 93,1794. (21)Streitwieser,A., Jr.; Holtz, D.; Ziegler, G. R.; Stoffer, J. 0.; Brokaw, M. L.;Guib6, F. J. Am. Chem. SOC.1976,98,5229.
