S. M.R. SPECTRA OF SILICOX HYDRIDE DERIVATIVES
April, 1963
with the limiting stage the transition of the a-complex into the carbonium ion, equation 7. Since reaction 6 is presumed t o be in equilibrium, it is easy to show
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(9) where r represents the net product of all the activity coefficients(non-idealities) in the system. According to the Hammett hypothesis, if the transition state is formed simply by the addition of a proton to the substrate, then the reaction velocity will be proportional to the acidity function ho,but if the transition state also contains a water molecule, then the reaction velocity will increase less rapidly than ho and perhaps may be approximately proportional to the hydrogen ion concentration. The use of the Hammett acidity function to linearly correlate J C ~ ~data , ~ in the absorption of ethylene in solutions of S0-9570 ?&SO4 (and water) has been demonstrated by Korovina, Entelis and Chirkovl’ and Lumbroso, Hellin, and Coussemant. l 8 On the other hand Gel’bshtein arid Ternkin*were unable to find such a linear correlation from their data. I n any case, the pertinency of these results t o this investigation is not clear, since the diethyl sulfate and ethyl hydrogen sulfate concentrations in their solutions were small, and the water concentrations were substantial-in contrast to the present work which was essentially carried out in non-aqueous solutions. Furthermore, Melander and Myhrelg point out that the direct proportionality between the acidity function ho and the reaction rate is pot proof of the mechanism proposed by Taft and (17) G.V. Korovina, S. 0.Entelis, and N. M. Chirkov, J . A p p l . Chem. (English Transl.), 81,597 (1958). (18) D . Lumbroso, M. Hellin, and F. Coussemant, Compt. rend. Congr. Intern. Chim. 81’ Liege, 1068, 1, b24 (1959). (19) L. Melander and P. 0.Myhre, Arlcivfdr Kemi, 13, 507 (1959).
805
co-workers. TaftZ0postulated that the first step in the hydration of olefins has a rapid proton transfer to the a-bond, and that the rate determining step came later in the migration of the proton to one of the carbon atoms forming a carbonium ion. Melander and Myhre’s argument indicates that the kinetic evidence alone cannot establish the existence or non-existence of the intermediate a-complex. How well do the rate constants determined in this work correlate, and what is their significance in connection with the mechanism indicated by equation 91 If log IC is plotted against H o or -log CH+, neither plot gives a straight line with slope of unity; neither do they give straight lines with other slopes. Attempts to obtain a linear correlation with other proton indices were also unsuccessful. One of the basic problems is how to evaluate realistically both the proton concentration and the activity coefficients in diethyl sulfate solutions. One might adjust the activity coefficients so the calculated data fit the proposed mechanism, or modify the mechanism assuming the activity coefficient product is essentially constant (or do both a t the same time). Figure 2 shows the highly nonlinear variation of the rate constant; adjustment of the abscissa scale by one scheme or another to straighten out the curve does not appear to be fruitful. The “true” mechanism must be considerably more complex than that indicated by Taft, perhaps involving orders of reaction greater than unity. One must conclude that considerable caution is needed in interpreting kinetic data in essentially non-aqueous solutions based on niodels proposed for aqueous solutions. Acknowledgment.-This research was supported by a grant from The Petroleum Research Fund administered by the American Chemical Society. Grateful acknowledgment is hereby made to the donors of said fund. (20) R. W. Taft, Jr., etal., J . Am. Chem. Soc., 74, 5372 (1952); 77, 1584 (1955); 78, 5807, 5Sll (1956).
N.M.R. SPECTRA aF SILICON HYDRIDE DERIVATIVES: 11. CHEMICAL SHIFTS IN SOME SIMPLE DERIVATIVES’ BY E. A. V. EBSWORTH ASD J. J. TURNER University Chemical Laboratory, LensJield Road, Cambridge, England Received September 15, 1962 The proton and fluorine chemical shifts in a number of simple compounds containing SiH bonds are presented; the SiH chemical shifts are much less sensitive to the nature of the substituents than are the CH chemical shifts in analogous carbon compounds. The SiH resonance shifts to high field from fluorosilane to trifluorosilane. The 19Fchemigal shifts in fluorosilanes are consistent with a recently developed interpretation of lgF chemical shifts in fluorocarbon derivatives.
We have recently recorded the chemical shifts for protons and fluorine nuclei in a number of simple compounds containing SIH bonds; the results are presented in Tables 1-111. The values of the H-H, 2gSSiH, H-F, and 29Si-F coupling constants, which were also measured, are discussed elsewhere. Experimental The compounds mere prepared by standard methods, and were studied in the liquid phwe in concentrated (ca. 95%) and dilute (1) P a i t I, J . Chem. Phge., 86,2028 (1962).
solutions, using cyclohexane or tetramethylsilane (for proton resonances) or trichlorofluoromethane (for fluorine resonances) as solvent and internal standard. T h e samples were held in Pyrex tubing of 5 mm. 0.d.; the spectra were recorded using a Varian Associates V4300B spectrometer and 12 in. electromagnet, operating at 40 Mc./sec., with flux stabilization and sample spinning; chemical shifts were measured using side-bands generated by a Muirhead-Wigan D695A decade oscillator. Errors quoted for proton resonances are derived from extrapolations, taking at least six meaaurements a t each concentration, unless otherwise stated.
806
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"O
5.0
1
o
H
1
I
2.0
2.5
3.0
3.5
4.0
Electronegativity, x.
Fig. 1.-Diagram showing variation in proton chemical shift of SiH8-compounds with the Pauling electronegativity of the substituent. We use the "best valves" of H. 0. Pritchard and H. A. Skinner, Chem. Rev., 55, 745 (1955).
TABLE I PROTOX CHEMICAL SHIFTSIN SIMPLESILANEDERIVATIVES, IN T UNITS (CsHiz = 8.56), EXTRAPOL.4TED TO INFINITE DILUTION Compound
(hO.01 unless otherwise stated), p,p.m.
SiHl 6.80 SiHJ 6.56 i 0.03 5.97 SiH& SiHaBr 5.83 4.83 SiHzBrz 5.41 SiHaCl 4.60 SiH,C12 SiHClaa 3.93 5.24 SiHaF 5.29 i 0.02 SiH2F2 5.49 SiHF3 6.18 SiH3CSb 5.39 (SiHdzO 5.65 (SiHdzS 5.58 ( SiH3)SCFsC 5.88 A 0.04 (SiH,)zSe 5.70 SiHaSeCFad 5.71 SiH&eC3F7d 5.56 f 0.02h ( SiH3),NB ( SiH3)2NCHa6 5.56 i 0.02h SiHs?J(CHs)2' 5.66 f 0.04h 5.55 (SiH&CNz' 5.58 f 0 . 0 2 SiH3?TCOg 5.54 f 0.02 SiH3SCSf 5.51 =I= 0.02 SiHaNCSeg E. A. V. Ebsworth and S. G. Frankies, a See also ref. 9. J . Chem. Soc., in press. See ref. 7. %ee ref. 8. e See ref. 4. 1 See ref. 5. 0 See ref. 6. From 3 measurements or less.
Discussion It is interesting to consider the chemical shifts given in TabIe I with those of the analogous derivatives of methane. The most striking difference is that the SiH chemical shift is much the less sensitive to changes in the nature of the substituent a t the central atom; thus although the proton resonance for monosilane
appears to low field of that in methane (a point that will be discussed in a later paper),2 trichlorosilane gives a resonance substantially to high field of chl~roform.~ In the SiH3- compounds so far studied the chemical shift depends largely on the nature of the substituent atom bound to silicon, and is relatively little affected by changes in the rest of the molecule. Thus out of seven SiH&S- compounds, six give chemical shifts in the range 5.54 A 0.034-6 and the seventh4 has T = 5.66 f 0.04, In disilyl sulfide' and disilyl selenide,8 replacement of one SiH3- group by CF3- changes T by 0.07 and 0.18 p.p.m., respectively. The SiH- resonance is also relatively insensitive to dilution. The biggest dilution shift that we have observed is only 0.2 p.p.m. (for SiHaNCSe),6and the mean dilution shift is only 0.05 f 0.03 p.p.m.; it has previously been noted that the proton chemical shift in trichlorosilane is not very sensitive to changes in solvent.g This insensitivity to dilution is of some interest where t,he fluorosilanes are concerned. It has been suggested that these compounds are associated in the liquid phaselo; any such association would be expected to lead to unusually large dilution shifts in both proton and fluorine n.m.r. spectra. The observed dilution shifts are given in Table 11. I n tetramethylsilane the TABLE I1 DILUTIOSSHIFTSIN FLUOROSILANES (p.p.m.j Compound
I1 resonance (solvent SiMea)
I.'resonance
SiHaF SiHZF3 SiHF3
0.03 i 0.01 0 . 0 4 i .02 0 . 0 2 i : .02
3 i t 3 3 . 8 f 2.0 0 . 3 =I= 0 . 5
proton dilution shifts are all small; in the fluorine resonance spectra the dilution shift of difluorosilane is apparently larger, but the uncertainty in the measurement makes the significance of the shift very doubtful. Thus the n.m.r. measurements do not support the idea that the fluorosilanes are associated in the liquid phase. I n the four series of halomethanes of general formula CH4-,X,, where X is fluorine, chlorine, bromine, or iodine, the proton resonance moves to lower fields as n increases from 1to 33011; the same is true for the chlorosilanes, while dibromo- and diiodosilane both give resonances well to low field of the analogous monohalosilanes. In the fluorosilanes, on the other hand, increasing fluorine substitution leads to a small increase in T . This may be connected with the anomalous effect of F, 0, or N substitution upon the H-H coupling'; in this context, the effect of increasing 0 or N substitution upon the SiH chemical shift would be of interest. There is a general decrease in T for SiH3- compounds as the Pauling electronegativity of the substituent increases. If, however, the chemical shifts in SiH3compounds, together with the point for SiH4,2are plotted against the Pauling electronegativity of the (2) E. A. V. Ebsworth and J. J. Turner, t o be published. (3) See L. M . Jackman, "Nuclear Magnetic Resonance," Pergamon Press, New York. N. Y . , 1958. (4) E. A. V. Ebsworth and N. Sheppard, J . Inorg. A'ucl. Chem., 9, 98 (1959). ( 5 ) E. A . V. Ebsworth and If.J. Mays, J . Chem. Soc., 4879 (1961). (6) E. A. V. Ebsworth and M.J. Mays, i b i d . , 4844 (1962). (7) A. J. Downs a n d E. A. V. Ebsworth, ibid., 3516 (1960). (8) E. A. V. Ebsworth, 11.J. EmBleus, and N. Welcman, i b i d . , 2290 (1962). (9) C. RZ. Huggins and D. R. Carpenter, J . Phys. Chem., 68, 238 (1959). (10) H. J. Ernelkus, A. G. Maddock, and C. Reid, J. Chem. Soc., 293 (1944). (11) S. G. Frankiss, J. r h u s . Chem., 67, 752 (1963).
J. Phys. Chem. 1963.67:805-807. Downloaded from pubs.acs.org by UNIV OF NEBRASKA-LINCOLN on 08/26/15. For personal use only.
April, 1963
SPECTRA OF 0-, vn-,
substituent (see Fig. l), the pattern obtained is rather different from that given by the analogous CHIderivatives.lZ For the CI&-compounds, the points for F, 0, N, and 13 lie on a straight line; the points for I, Br, C1, and C I lie on a second h e , whose gradient I (drldz) I is less, aiid thc failure of the two lines to coincide has been explained in terms of differences in the magnetic anisotropies of the (C-halogen) bonds. Icor the SiH3-eompouiids, the points for IC, 0, and N lie 011 a straight line, but the line does not pass through the point for H. AIoreover, a line through the points for I, Br, and C1 does not pass through the point for F and is much steeper than the (F,0, N) line. Though these differenccs between CH3 and SiHo derivatives may be partly due to the different effectsof magnetic anisotropy, this sort of explanation is unlikely to account for the high-field shift produced by additional fluorine substitution a t silicon. The chemical shifts of protons bound to silicon arc likely to be affected by (p+d)n-bonding, and by any contraction in the d-orbitals of silicoii that the substituent may cause; the only compound considered here in which there can be no (p+d) T - bonding is SiH4. I n the fluoromcthanes, the fluorine resonance moves to successively lower field values with each additional fluorine substituent.l3 The fluorine rcsonaiicc in the fluorosilane moves to lower fields as the number of fluorine atoms incrcascs from 1 to 3, but C#I for tetrafluorosilane is closed to C#I for difluorosilanc (see Table 111). Changes in molecular symmetry may be at least partly responsible for this irregularity; a similar irregularity (12) €1. Spieseckc a n d W. G. Schneider, J . Chem. Phys., 36,722 (1961).
TABLE I11 FLUORINE CHEMICAL SHIETSIN FLUOROSILANES (p.p.m. RELATIVE TO CC13E') Compound
d
SiH3F 217 =k 3 SiHzFz 161 1.5 SiHF3 109.5 rt 0 . 6 SiF4" 163.3 a Corrected from the measurement in ref. 13 using the value for ICtsSiF given in ref. 13 and by G. Filipovitch and G. V. I). Tiers, J . Phys. Chem., 63, 761 (1959).
*
is also observed in the alkylfl~orosilaiies.~~ Explanatioiis in terms of changes in electron-density a t fluorine arc almost certainly inadequate, since it has been shown that fluorine chemical shifts are largely determined by changes in the paramagnetic term. The observed lowfield shift as hydrogen is replaced by alkyl groups or fluorine atoms in mono- or difluorosilancs agrees with a recent interpretation of the fluorine chemical shifts in a numhcr of C l ~ - d e r i v a t i ~ e sThis . ~ ~ indicates that a substituent a t the cai.bon atom will produce a low-field shift if it has more relatively low-lying excited states than the group it replaces; this is certainly true when methyl groups or fluorine atoms are compared with hydrogen. This type bf cxplaiiation also accounts for the large low-field shifts in vinyl- and phenyltrifluorosilanes as against trifluorosilanc itself. Acknowledgments.--We are grateful to Dr. N. Sheppard for his interest, in this work, and to the Wellcome Trustees, who lent the spectrometer to the department. (13) E.Sclinell and E. G. Rochow, J . Inorg. Nuclear Chem., 6 , 303 (1958). (14) E. Pitcher, A. D. Buckingham, and I'. G. A. Stone, J . Chem. Phys., 36, 124 (1962).
ULTRAVIOLET ABSORPTION SPECTRA OF 0-,m-, AND p-PHENYLENEDIAMINES AND THEIR RIONO-AXD DIHYDROCHLOIUDES I N AQUEOUS SOLUTION' BY P. K. GALLAGHER Bell Telephone Laboratories, Incorporated, Murray Hill, A'ew Jersey Received September 18, 1862 l m > p. The ultraviolet absorption spectra were measured for each of the nine species a t 26" and an ionic strength of 1.0 M . The spcctra agreed well with the data available in the literature. The spcctra of the ionic species, with the exception of the divalent orlho ion, indicate that the addition of a proton to the free electron pair on the nitrogen cancels the original effect of the amine group upon the benzene spectrum.
Introduction Although there has been considerable eff 01% devoted to the absorption spcctra of substituted benzenes, there are only meager data available on the effects of proton association or dissociation reactions on the ultraviolet absorption spectra of the substituted benzene species.2 It has been observed by numerous investigators that the spectrum of anilinium ion is very similar to that of benzene. Recently Semba3 has shown that the spectrum of the nitroanilinium ion is similar to that of (1) Presented in p a i t a t the 142nd National ACS Meeting, Atlantic City, New Jersey. Sept. 10, 1962. (2) hl. J. Kamlet, "Organic Elcrtronic Spectral Data," Vol. I and 11, Interscience Publishcrs, Now York, N. Y., 19GO. (3) K. Semba. Bull. Chem. SOC.J a p a n , 33, 1040 (10GO).
nitrobenzene. Kohlcr and Scheibe4 indicate that the spectrum of the monoprotonated m-phenylenediamine ion resembles that of aniline. The contention is that the addition of a proton to the free electron pair on the nitrogen atom cancels the original effect of the amine group upon the absorption spectrum. This investigation extends the work of Kiihler and Scheibe to cover all three isomeric phenylenediamines and their divalent ions as well. The following aqueous equilibria are of interest if the absorption spcctra arc to be measured for each of the riiiie spccics (4)
V. 11. Kijhler and S. Scllcibe, Z . anorg. allgem. Chem., 266, 221 (19.X).
