Tm EFFECTOF AROMATICSOLVENTS ON P.M.R.SPECTRA
Similar desorption experiments were carried out for the propylene-silica-alumina system. Above 80°, the activation energies for the formation of individual products, as well as the distribution of the products, were found to be almost the same as those obtained with butene-l. This suggests that above 80' identical chemisorbed layers are formed in both cases.
2679
Acknowledgment. The author wishes to express his sincere thanks to professor P. H. Emmett of Johns Hopkins University for his helpful comments. It is also a pleasure for the author to acknowledge the help of the American Chemical Society in financing this work in the form of a Petroleum Research Fund Grant to Professor P. H. Emmett.
The Effect of Aromatic Solvents on Proton Magnetic Resonance Spectra
by Theodore L. Brown and Kurt Stark Noyes Chemical Laboratory, University of Illinois, Urbana, Illinois
(Received March 1, 1966)
The effect of aromatic solvents on chemical shift in the proton magnetic resonance spectrum where M is a group IV elehas been observed for compounds of the form (CH3).M&,, ment, and X is a halogen, usually chlorine. Chemical shifts relative to tetramethylsilane (TMS) were determined in carbon tetrachloride, benzene, toluene, and mesitylene. The results are not in good accord with the hypothesis that the shifts are due mainly to solvent orientation resulting from hydrogen bonding. They are in better agreement with a model in which dipole-induced dipole interactions are responsible for solvent orientation. The quantity (was- 7 ~ ~ 1 correlates ,) very well with solute dipole moment for compounds of related geometry. The aromatic solvent shift may have value in certain situations as a method for estimating dipole moments.
The effect of an aromatic solvent on proton magnetic resonance chemical shifts has been discussed by several workers.1-11 Solute protons generally experience a high-field shift, but its magnitude varies widely from one solute to another. It is commonly accepted that the shift arises predominantly from the large diamagnetic anisotropy of aromatic solvent molecules, which in turn owes its origin to the relatively free circulation of a-electrons in the molecular plane. The magnitude of shift predicted theoretically for complete averaging over the solvent molecule surface is small in comparison with many observed shiftsa6 Other nonspecific effects may contribute to a larger value for but the magnitude of the observed effects for many solutes requires that the time-averaged location of the solute protons relative to the aromatic solvent molecules is preferentially along the sixfold axis normal
to the plane of the ring. Hydrogen bonding of solute protons to the ring has been suggested to account for (1) A. A. Bothner-By and R. E. Glick, J . C h m . Phys., 26, 1651 (1957). (2) L. W.Reeves and W. G. Schneider, Can. J . Chem., 35,251 (1957). (3) T. Schaefer and W. G. Schneider, J. Chem. Phys., 32, 1224 (1960). (4) A. D.Buckingham, T. Schaefer, and W. G. Schneider, &id., 32, 1227 (1960). (5) A. D. Buckingham, T. Schaefer, and W. G. Schneider, 8id., 34, 1064 (1961). (6) R.J. Abraham, Mol. Phys., 4, 369 (1961). (7) R.J. Abraham, J. Cha. Phys., 34, 1062 (1901). (8) T.Schaefer and W. G. Schneider, tiid., 32, 1218 (1960). (9) J. V. Hatton and W. G. Schneider, Can. J. Chem., 40, 1285 (1962). (10) J. V. Hatton and R. E. Richards, Mol. Phys., 5, 139 (1962). (11) (a) W. G. Schneider, J. Phys. Chem., 66, 2653 (1962); (b) P. Diehl, J . chim. phys., 61, 199 (1964).
Volume 69, Number 8 August 1966
THEODORE L. BROWN AND KURTSTARE
2680
the solvent shift in many The effect of solute molecular shape4 and the orienting influence of solute dipole rnornentll have also been mentioned as possible factors. We have examined the aromatic solvent shift in a series of closely related compounds of the general form (CHS).MXh, where M is a group IV element, and where X is a halogen, usually chlorine. This series of substances is well suited to a test of the hypothesis that hydrogen bonding accounts for extraordinary aromatic solvent shifts; the symmetries are well defined and the compounds are closely similar in over-all shape and geometry. Further, useful proton coupling constant data are readily obtained. The chemical shifts of all compounds investigated were measured relative to tetramethylsilane (TMS) in carbon tetrachloride, benzene, toluene, and mesitylene as solvents. The use of TMS as an internal standard is particularly apt for the purposes of this work, as effects other than those arising from solutesolvent interactions are all but eliminated in the comparison. All chemical shifts reported are essentially infinite dilution values.
Table I : Chemical Shift Data for Organomethyl Compounds in Various Solvents
ccl4
CsHs
8.353 8.835 9.368 8.148 7.68 9.292 8.86 9.249 9.578 7.257 7.828 8.404 9.090 8.817 8.750 9.714
1.432 0.825 0.402 1.228 1.02 0.091 0.74 0.396 0.244 0.594 0.407 0.226 0.018 1.158 1.104 0.153
-A CHsCaHr (CHdrCsHa
1.327 0.763 0.368 1.118 0.90 0.032 0.67 0.367 0.224 0.588 0.350 0.224 0.018
Jiac- H (nest)
1.126 0.635 0.312 1.042 0.86 0 0.59 0.311 0.201 0.502 0.290 0.208 0
... ... ...
PJ
D.
143 137.8 133 141.1 140.1 128 126 124.5 120.5 134.2 131.8 127 124
3.6 4.2 3.5 3.2 2.6 0 1.9 2.3 2.1 1.5 2.2 2.2 0 3.1 2.9 0
..
.. ..
...
...
130
I
I
Results and Discussion Hydrogen bonding has been proposed as an explanation for most large diamagnetic (upfield) shifts of solute proton resonances relative to the shifts exhibited by a nonpolar reference compound such as neopentane, cyclohexane, or TMS. We begin, therefore, by examining the data in Table I in terms of the hydrogenbonding hypothesis. The 13C-H coupling constants provide a measure of the acidity of hydrogen bound to carbon. This follows from the relationship between coupling constant and fractional s character in the carbon orbital on the one hand,12and the increase in electronegativity of the carbon orbital with increasing s character on the other.la There should be a monotonic relationship between J l c - ~ and the acidity of the proton. If hydrogen-bonding propensity is related to acidity of the proton, as is generally supposed, the degree of hydrogen bonding to aromatic solvents should increase with increase in J l c - ~ . Figure 1 shows the chemical shifts of the solute protons in benzene, relative to carbon tetrachloride (all related to TMS) graphed vs. J l c - ~for the methyl-group protons. A regularity in the relationship is discernable for compounds other than the chlorosilanes. Whatever interpretation is put on the low values of Jlc-= for the latter compounds, the failure of the relationship between J l c - ~ and AT in these instances is not in good accord with the Tha Jound of Physical Chemirtry
I
9
I
0.5
I
I
1.0
1.5
I
Ai
Figure 1. W-H coupling constant us. AT, the diamagnetic solvent shift in benzene as compared with carbon tetrachloride solution (tetramethylsilane as reference in both solvents). The chlorosilane data are represented by shaded circles.
notion that AT is due predominantly to hydrogen bonding. It is true for every solute in Table I that AT decreases in the order benaene > toluene > mesitylene. This is just the reverse of the order of solvent basicity, and therefore of the expected order of enthalpies for hydrogen bonding. The observed AT values represent (12) (a) N. Muller and D. E. Pritchard, J . Chem. Phya., 31, 708, 1471 (1959); (b) J. N. Schoolery, aid.,31, 1427 (1959); (0) C.Juan and H. S. Gutowsky, &id., 37, 2198 (1962). (13) J. Hinze and H. H. J d Q ,J. Am. Chem. Soe., 84,540 (1962).
THEEFFECTOF AROMATIC SOLVENTS ON P.M.R. SPECTRA
an average value for all solute species. Assuming for simplicity that there is a narrow range of AT values, centered at Arn, which characterizes nonhydrogenbonded solute protons, and another narrow range centered at Arb which characterizes the hydrogenbonded moieties, the observed quantity is ATO= FnArn FbArb. Letting A n = AT, e, then ATO= FnArn Fb(Arn E ) = AT, eFb. E is likely to be related to the enthalpy change, whereas Fb is a function of the free energy change. E should increase in the order benzene < toluene < mesitylene. The equilibrium constant for hydrogen-bond formation must therefore decrease in the order benzene > toluene > mesitylene. It is not easy to accept the proposition that steric factors could so affect Kequilfor hydrogen-bond formation for every solute studied. There is evidence that for many solutes Kequilis in the order mesitylene > toluene > benzene. 14, l5 We conclude, therefore, that specific hydrogen bonding is not responsible for a major share of the observed AT values in the solutes examined. The magnitudes of many of the AT values listed in Table I are as large as any yet reported in the literature (e.g., for chloroform2). The hypothesis that hydrogen bonding (Le., the formation of specific 1:l complexes) plays a major role in producing large diamagnetic shifts in any solute is open to serious question. It has been reported that group IV halides and related compounds such as the haloforms form chargetransfer complexes with aromatic solvents.l6 This kind of complex formation is possible for the more acidic of the solutes which we have studied, e.g., CH3SnC13. We have examined the llgSn-C-H coupling constants in all the tin compounds to determine whether this quantity shows a dependence on aromatic solvent. On the basis of previous work17~18 one expects that if discrete complex formation occurs, with consequent increase in the coordination about tin, JB,-C-H should change. There is evidence (Table 11) of a small increase in J s ~ - c - Hin mesitylene as compared with benzene for CH3SnC13,but the effect is slight, and is not seen in CH3SnBr3. We conclude that any charge-transfer complex formation which occurs does not significantly affect the coordination about tin, nor in all probability does it affect the chemical shift of the methyl-group protons. If specific hydrogen-bond formation is excluded as the major cause of the observed AT values, there remain solute-solvent interactions of a more general kind. For polar solutes, an orientation of solvent molecules through dipole-induced dipole interactions is the most important effect to be considered. All of
+ +
+
+
+
2681
Table I1 : 11%-C-H Coupling Constants” of Methyltin Compounds in Various Solvents Compd.
CClr
CsHa
Solvent CHaCiHr
(CHa)sCsHs
CH 8Sn C4 (CH$),SnClz (CHs)aSnC1 CHsSnBr$ CHsSnI8 (CHdSn
99.4 69.5 59.3 88.4 74.5 54.2
99.5 68.8 58.8 89.1
99.3 69.8 58.6 89.4
101.2 72.1 59.2 88.3
54.7
54.0
54.5
I
a
...
...
...
Neat
100.0
...
60.2 88.6 72.7 54.0
Estimated uncertainty is f 0 . 5 C.P.S.
the solutes examined here have an axis of symmetry along which the molecular dipole moment is oriented. The aromatic molecules are more polarizable in the molecular plane than normal to it (al = 6.35 QII - 12.3 8.9l9;further, a closer approach to the solute molecules can be made along the sixfold axis than in the molecular plane, because of the hindrance presented by hydrogens or methyl groups. These qualitative considerations suggest that there should be a time-averaged orientation of solvent molecules with the plane of the ring parallel to the solute molecular dipole moment. Protons of the solute methyl groups will, therefore, experience an averaged solvent environment which presents the protons with solvent molecules along the sixfold axes of the latter. The solute protons may not all be simultaneously exposed to the solvent environment. Figure 2 depicts solute molecules containing three, two, or one methyl groups, viewed in each case along the symmetry axis. Solvent molecules, assumed for simplicity to lie parallel to the solute dipole axis, are also shown schematically as oriented normal to the plane of the paper. The purpose of the figure is to show in a crude way how the fractional exposure of protons to solvent molecules varies with the number of methyl groups. The axis of the CH3 group is, of course, tilted differently with respect to the plane of the paper in the three cases. However, assuming free rotation about the C-M bond, the average angle between a C-H bond and the dipole axis is not greatly different in the three cases. If it j s assumed that only solvent orientations parallel (14) 2. Yoshida and E. Osawa, J. Am. Chem. SOC.,87, 1467 (1965). (15) M. R. Basila, E. L. Saier, and L. R. Cousins, ibid., 87, 1665 (1965). (16) F. Dorr and G. Buttgemit, Be?. Bunsenges., 67, 867 (1963). However, see W.B. Person, J. Am. Ch5m. SOC.,87, 167 (1965). (17) J. R. Holmes and H. D. Kaesz, &id., 83, 3903 (1961). (18) N.A. Matwiyoff and R. S. Drago, Inorg. Chem., 3, 337 (1964). (19) LandolbBornstein, “Zshlenwerte and Functionen,” Vol. I, Pt. 3, J. Springer-Verlag, Berlin, 1951,p. 510.
Volume 69, Number 8 August 1966
THEODORE L. BROWNAND KURTSTARK
2682
(4
(b)
(C
1
Figure 2. A schematic illustration of the orientation of solvent molecules (viewed along the molecular plane, and shown as rectangles) about solute molecules containing (a) one, ( b ) two, or (c) three methyl groups. The heavy lines represent the angular segment about each methyl group in which the protons are in contact with the solvent. The solute molecule in each case is viewed along the dipolar (symmetry) axis.
to the solvent molecule axis contribute to AT, then the relative exposure of solute protons to such oriented solvent molecules is estimated to be: (CHa)aMX, 1.0; (C€€&MX2,1.5; and CHaMXa, 3.0. The AT values observed in benzene for all compounds studied are graphed in Figure 3 vs. the molecular dipole moments.20 For compounds containing equal numbers of methyl groups the relationship is nicely regular, and is adequately approximated as linear. It should be noted that the graph includes the methylchlorosilane data, and that the line for the monomethyl compounds includes compounds other than those in the group IV series. The slopes of the lines are 0.115, 0.188, and 0.384 p.p.m./D. These are in the ratio 1:1.6:3.3, in remarkably good agreement with the estimated relative exposures of protons to solvent environment. The results shown in Figure 3 are in excellent accord with the hypothesis that dipole-induced dipole orientation of solvent molecules is responsible for the observed AT values. The diminished solvent effect in toluene, and still further in mesitylene, is ascribed to steric hindrance which impedes simultaneous packing of solvent molecules about a particular solute. A relationship between AT and p should be observed for any series of suitably related compounds. The relationship shown in Figure 3 for the monomethyl compounds can be expected to hold also for other compounds in which a methyl group is located on the ~ d e c u l a axis, r and which have roughb similar shape. Methyl iodide and acetonitrile, for example, with molecular dipole moments of 1.48 and 3.5 De,” respectively, exhibit corresponding AT values in benzene of 0.667 and 1‘016* The iodide datum ‘lose to the line in Figure 3, whereas the acetonitrile point is rather far off. But acetonitrile has a much smaller molecular volume than the other solutes; the packing The Journal of Physical Chemistry
3 ,d (DEBYE)
2
4
Figure 3. Dipole moment us. AT, the diamagnetic shift in benzene as compared with carbon tetrachloride solution.
of solvent molecules about the protons is probably not similar. The present resulta, as well as those reported by others,ll suggest that the aromatic solvent shift of proton magnetic resonances could be made the basis of a method for estimating dipole moments. It is obvious that the method would be largely empirical, and that it could not compete in accuracy with the conventional dielectric constant method. It would, on the other hand, possess the singular advantage that it could be applied to the components of mixtures, for estimation of dipole moment, for distinguishing isomers on the basis of polarity, and for other like purposes.
Experimental Materhls. AU solutes employed were purified before use. Methyltin tribromide and methyltin triiodide were synthesized; all other solute compounds were obtained from commercial sources. Methyltin tribromide was synthesized by the method described by Krause and Grosse.21 The triiodide was prepared by reaction of methyltin trichloride with potassium iodide in liquid
sot.
The solvents employed were all reagent grade; benzene, toluene, and mesitylene were dried, distilled and stored over sodium wire. (Solutions of methyltin (20) All dipole moment data not determined in our laboratory (see Experimental) were taken from “Tables of Experimental Dipole Momenta,”A. L.M c C l h , Ed., W. H. Freeman and Co., San Francisco, 1963. The quality of the data of course varies from one solute to another. but the dinole moments are Drobablv relativelv correct to 0.2D. (21) (a) CHaSnBra waa prepared from CHaSnOOH by heating the latter with concentrated aqueous HBr: E. Krause and A. V. Grow, “Die Chemie der Metallorganische Vergindungen,” G. Borntraeger, B e r h 1937, P- 341; (b) CHsSnOOH wm obtained by reacting SnClz with CHsI in EtOH and precipitating the acid with COz: P. PfeiEer and R. Lehnardt, c h . BW.,36, 3028 (1903).
THEEFFECT OF AROMATIC SOLVENTSON P.M.R.SPECTRA
2683
trichloride in mesitylene turn a violet color during the P A to be 0.75 of the corresponding values for SnBr4 course of an hour or so after the solutions are made up. and Sn14.24325Proceeding similarly with PT data for The origin of this color was not fully investigated, but (CH3)2Hg,26HgI2, and HgBr,25we obtain the estimates it was ascertained that its appearance is not accomof PEand PAshown in Table 111. panied by any observable change in the proton magnetic resonance spectrum.) N.m.r. Spectra. The proton magnetic resonance Table III : Dipole Moments of Methyltin and spectra were obtained for the most part on a Varian Methylmercuric Halides in Benzene Associates Model A-60 spectrometer. Chemical shifts are accurate to kO.01 p.p.m.; coupling constants are PT, PE, PA, P (at Compd. P cm.8 cm.8 cm.8 26') accurate to within about 0.5%. The reported chemical 6 . 6 3.20 262 45.1 CHsSnBra 3.20 -0.80 shifts were obtained from measurements on solutions 2.64 59.5 10.9 212 1.70 -0.82 CHaSnIs of varying concentration (2-10%) and extrapolation 26.0 4.0 3.1 -0.9 231 CHsHgBr 3.80 to infinite dilution. Only small concentration d e 33.1 6.0 2.9 -1.0 210 CHaHgI 3.00 pendences of the J values were noted. a a = &/dw2) where w 2is weight fraction. ' j 3 = a Y / b w 2 1 where Dipole Moment Determinations. The dipole moments Y is specsc volume in cubic centimeters per gram. of methyltin tribromide and triiodide, and of methylmercuric bromide and iodide have not been reported heretofore. The dielectric constants of dilute solutions of these compounds in benzene were measured as deAcknowledgment. This research was supported by a scribed previously.22 The relevant results are sumresearch grant from the National Science Foundation. marized in Table 111. Total polarization is calcuK. S. expresses his appreciation for a Fullbright Travel lated by the method of Halverstadt and K U m l e ~ , ~ Grant. ~ with slight modification.22a Allowance for atomic and electronic polarization in these compounds presents (22) (a) T.L.Brown, J . Am. Chem. SOC.,81, 3232 (1959); (b) T.L. Brown, J. G. Verkade, and T. S. Piper, J. Phys. Chem., 65, 2051 some diaculties. For nonpolar SnBr4 and snI4, PE are 46.8 and 66.0 cm.a, r e ~ p e c t i v e l y . For ~ ~ ~ ~ ~(1961). (23) I. F. Helverstadt and W. D. Kumler, J. Am. Chem. SOC.,64, Sn(CH3)4, with d26 1.295, nD 1.4932, the molar refrac2988 (1942). tion is 40.1 ~ m .from ~ , which the SnC bond refraction (24) S. E. Coop and L. E. Sutton, J . Chem. SOC.,1269 (1938). is estimated to be 5.0 From these results we (25) Reference 19,p. 514. obtain PE for CH3SnBr3and CH3Sn13. We estimate (26) H. Sawatzky and G. F. Wright, Can. J . Chem., 36, 1555 (1958).
Volums 69,Number 8 August 1986
