SOLVENT E ~ ~ n c lON c s THE FLUOROFORM NMRSPECTRA
497
Solvent Effects on the Fluoroform Nuclear Magnetic Resonance Spectra
’by William B. Smith* and Arthur M. Ihrig Department of Chemistry, Texas Christian University, Fort Worth, Texas 76189 (Received July $9, 1970) Publication costs assisted by T h e Robert A . Welch Foundation
Downloaded by FLORIDA STATE UNIV on September 8, 2015 | http://pubs.acs.org Publication Date: February 1, 1971 | doi: 10.1021/j100674a009
The proton and fluoride chemical shifts of fluoroform in a variety of solvents have been determined. These results are compared with those reported in the literature for chloroform. The chemical shifts for the proton, in general, move downfield with increasing polarity of the solvent while the fluorine chemical shift moves upfield. The operation of a solvent “reaction field” is indicated. Specific complexation of the fluoroform to benzene was suggested by the marked upfield shifts of both the proton and fluorine chemical shifts.
Recently, Lichter and Roberts have reported on the proton and ‘ 3 C chemical shifts of chloroform in a variety of solvents. It was their conclusion that the primary effect producing solvent-induced chemical shifts was the known liyt3rogen bonding propensities of the solute rather than such efflscts as the solvent “reaction field.” Their results in benzene were consistent with the idea of a 1: 1 complex in which the CH bond axis of the solute lies along the sixfold axis of the solvent molecule The investigation of solvent-induced chemical shifts has been a subject of considerable interest as a means of studying neak i~iteractions.~However, as Lichter and Roberts1 point out, few solvent studies involving nuclei other than protons have appeared. Because of the greater polarity of fluoroform (1.6 D as opposed to 1.1. D3) as a solute, the study of solvent effects on both proton and fluorine chemical shifts seemed a worthwhile ex tension of the chloroform work.
‘
Experimental Seetiom A sample of fluomlform was obtained from the MatheEton Company. Reagent and spectral grade solvents were used, and in no case were any impurities detected. Proton and fluorine spectra were taken on a Varian HA-100 spec6rometer operating at 100 and 94.1 MHz. Gaseous Auoroform was condensed and measured as a liquid and then distilled into an appropriate quantity of solvent to give cu. 5 mol % solution. These samples were degassed by lhe usual freeze-thaw techniques and aealed under T”XU.H~. Two sets of samples containing tetramethylsilelne and 1,1,2,2-tetrachlor0-3,3,4,4-tetrafluorocyclobut m e 4 as internal references were prepared for the pro1 on and Jluorine investigations, respectively. The proton spectra were calibrated in frequency sweep mode with tht: final line positions representing an average of at least five independent scans. The fluorine spectra were caliblrated in the HR mode by superimposing audiofrequency d e bands generated with a Hewlett-Packard 280 GD oscillator on each of the transitions in the fluorine spectra. The fluorine chemical shifts represent an average of at least five independent measiirementt, with an estimated accuracy of 3tO.5 Hz.
Results and Discussion The chemical shifts on the protons and fluorines in fluoroform in a variety of solvents are given in Table 1 as are the values of the H F coupling constant. It has been noted before5 that the fluoroform H F coupling Table I: Nmr Data for Fluoroform in Various Solvents
1. Cyclohexane Carbon tetrachloride Benzene Anisole Chloroform Methylene chloride Acetone Nitrobenzene Dimethylformamide Acetonitrile 11. Nitromethane 2. 3. 4. 5. 6. 7. 8. 9. 10.
6.25 6.46 5.31 5.89 6.47 6.54 7.04 6.83 7.32 6.75 6.74
35.96 35.69 35.73 35.78 35.45 35.25 34.70 35.35 34.63 34.39 34.83
79.31 79.23 79.27 79.26 79.25 79.35 79.40 79.32 79.37 79 38 79.38 I
constant is to a large extent solvent independent. That trend is confirmed in Table I and stands in contrast to the observations of Lichter and Roberts’ that the CH coupling of chloroform follows the proton chemical shift in a linear fashion. With the exceptions of benzene and anisole all proton chemical shifts are downfield from the value in the least polar solvent, cyclohexane. This observation is in accord with those of Lichter and Roberts1 for chloroform. However, in contrast to the carbon chemical (1) R. L, Lichter and J. D. Roberts, J . Phys. Chem., 74, 912 (1970). (2) (a) P. Laszlo, Progr. Nucl. Magn. Resonance Xpectrouc., 3 (1968);
(b) J. Ronayne and D. H. Williams, Ann. Rev. N M R (Nucl. Magn. Resonance) Spectrosc., 2 (1969). (3) A. L. McClellan, “Tables of Experimentd Dipole Moments,” W. H. Freeman, San Francisco, Calif., 1963. (4) R. W. Taft, E. Price, I. R. Fox, I. C. Lewis, K. K. Anderson, and G. T . Davis, J . Amer. Chem. Soc., SS> 709 (1963), have adequately established the suitability of the cyclobutane a6 an internal standard for fluorine chemical shifts. ( 5 ) R . H. Cox and S. L. Smith, J . Magn. Resonance, 1, 432 (1969). The Journal of Physical Chemistry, Vol. 76, N o . 4, 1971
WILLIAMB. SMITHAND ARTHUR31. IHRICI
Downloaded by FLORIDA STATE UNIV on September 8, 2015 | http://pubs.acs.org Publication Date: February 1, 1971 | doi: 10.1021/j100674a009
498 shifts in the chloroform study, the chemical shifts of the Huoroform fluorines are all upfield in all solvents compared with ejdohexane. There is only a very rough linear correlation between the proton chemical shifts for fluoroforin and chloroform in the same solvents. Solvent-induced chemical shifts may be caused by the operation of n number of different effectsa2 Presumably the effecta of bulk susceptibility are eliminated hy the ute o f an internal standard. Similarly, one may suppose that van der Waals effects and anisotropy of randomly oriented solvent) molecules are of little importance. Lieinter and Roberts1 suggest that the major causative fwtor for solvent-induced chemical shift effects I excluding the specific complexation with aromatic solvenfs such as benzene) is the hydrogen bonding inteiractioin of chloroform with the solvent. They found no reason to propose the operation of the solvent “ r e a d o n field” effect. A similar conclusion had been reached previously by Kuntz and Johnson6 4 and polyhalides.’ for a series of a l k ~ halides However, the gcneral downfield shift of the proton and the conccimitant upfield shift of the fluorine chemical shifts in solvents of increasing polarity suggests the operation of the “reaction field” effect. Petrakis and Herristein8 have observed the proton and fluorine chemical shifts in the gas phase of fluoroform in gas mixtures of various polar “perturbers.” They concluded that the effects of the perturbing fields of the gaseous diluerirs operated in opposite senses on ends of the fluoroforno, dipole; i.e., with increasing polarity of the medium the proton will be deshielded while the fluorines will be shielded. Our results in solution concur with their vapor phase measurements. Furthermore, excepting the aromatic solvents, there are rough linear correkttions between such solvent parameters as (e - 1) / ( t b), as shown in Figure 1, and the solvent nicaia.~’transition energies ET of Dimroth, Reiehardt, S!epmann, and Bohlmann.$ No doubt hydrogen ’bonding and more specific complexation interactions operate with fluoroform, but this more polar solute :tllows the “reaction field” effect of the solvents to be S P ~ It is now gmerdly accepted that chloroform hydrogen bonds w ~ a k l gto benzene forming a 1 : l complex with the C-N bond along the sixfold symmetry axis of the bazene. Lichler and Roberts’ have calculated a separation of &out 3.8 8 for the difference between the center of the ring rznd the chloroform proton. Given the greater electronegaiivity of fluorine over chlorine and the greater polarity of fluoroform, one might expect
-+
The Journal of Physical C%,emistry, Yoliol. 76, No. 4, 1971
0 X X
0
0.4
0.6
0
x x
0.9
x
1.0
€-I Efl
c _ _
Figure 1. Plot of the proton (circles) and fluorine (x ) chemical shifts in various solvents (aromatics excluded) vs. the solvent dielectric constant function.
fluoroform to complex more closely to the benzene. Using Johnson-Bovey tables for the calculation of anisotropic effect of the solvent on the fluoroform proton yields a range of 3.6-3.8 ,& for the separation.10 As the confirmation of these values one can calculate that the fluorine should experience an upfield of ca. 0.23 ppm on going from cyclohexane to benzene,” a value in good agreement with the experimental value of 0.33 ppm and suggestive of the correctness of the geometry of the complex. Acknowledgnzent. This work ’was supported by the Robert A. Welch Foundation. We wish to express our gratitude to the Foundation. A . AI. 1. wishes to express his appreciation to the T. C. U. Research Foundation for a Postdoctoral Feilowship. (6) I. D. Kuntz and M. D. Johnson, J . Amrr. Chem. Hoe., 89, 6008 (1967). (7) A referee has pointed out that the behavior of the protons in our system is not to dissimilaI from that recently observed tor ~ l chloroacetonitrile in a series of solvents [lt.L. Bchmidt, E. S. Butler, and J . H. Goldstein, J . Phys. Chem., 73, 1117 (1969)l. Collision complexes were evoked as the best explanation for the latter system. (8) L. Petraltis and H. J . Rernstein, J . Chem. E‘hys., 38, 1562 (1963). (9) See C. Reichardt, Angezu. Chem., Tnt. Ed. Enol., 4, 29 (1965). (10) Lichter and Roberts’ chose a chemical shift v a l u e for the chloroform proton in a nonpolar-isotropic solvent from the linear relationship of the chemical shift and J13cR. Their value was 0.18 ppm upfield from the experimentd value in cyclohexane. Our calculations were made from the data in cyclohexane and with a 0 . 1 6 0 p ~ mcorrection added in. The latter gave the separation of 3.6 A. (11) The geometry of the fluoroform molecule was taken from R. A. Rerheim, D. J. Hoy, T. R. Krugh, and B. J. Lnvery, J , Chem. Phys., 50, 1350 (1969).
