6999 Pyrolysis of '"-Enriched
Phenyltrimeth! lsil) Idiazomethane.
13C-Enriched phenyltrimethylsilyldiazomethane was injected into a gas chromatograph under condition B. The product. I3C-enrichcd I.I-dimeth!I-I-silabenzoc)~clopcntene. was isolatcd in a cold trap d t -78 "C: Y.MR (CCIj) b 0.31 (6 H, 5 ) . 1.00 ( 2 H. t ) , 3.02 ( 2 H. t J . 6.95-7.51 ( 4 H. m ) ; mass spectrum (70 e V ) m/e (measured intensity) 163 (13.2). 162 (32.5). I20 (2.1). I19 ( 5 . 5 ) ; mass spectrum nile 119.0318 f 0.0006 (calcd for C7H7Si. nile 119.0317): I3C h M R from Me&i (intensit}) 153. I I (I24 992), 139.86 ( 1 I83 552). 131.93 (195 616), 129.20 (608 256), 1'25.56 ( I 042 816). 31.80 [288 640). I 1.47 [ I77 600), - 1.59 (230 336). Pyrol!sis of '"-Enriched Phenyl Trimeth!Isilyl Ketone Tosylh!drazone Salt. Enriched (12.8% I3C) salt was flash-p)rol!zed by method C to produce, after isolation b! gas chromatography under condition B. enriched ( I 3.090 I3C) 6.
References and Notes (1) We thank the donors of the Petroleum Research Fund, administered by the American Chemical Society, and the National Science Foundation (Grants GP 30797X and MPS 74-05690) for generous support of this research. In addition, we thank the Shinetsu Chemical Co. for gifts of silanes. (2) (a) Portions of this work were communicated p r e v i ~ u s l yand ~ , ~others are taken from the Ph.D. Theses of R.R.G. (Princeton, 1976) and J.A.K. (Iowa State. 1976) and the A.B. Thesis of A.J.R. (Princeton. 1975). (b) The University of Tsukuba. (c) Princeton University. (d) Iowa State University. (3) T. J. Barton, J. A. Kilgour, R . R . Gallucci, A. J. Rothschild, J. Slutsky, A. D. Wolf. and M. Jones, Jr., J. Am. Chem. SOC.,97, 657 (1975).
(4) W. Ando, A. Sekiguchi, T. Hagiwara, and T. Migita, J. Chem Soc., Chem. Common., 372 (1974). (5) For a review see L. E. Gusel'nikov, ACC. Chem. Res., 8, 18 (1975). (6) G. L.Delker, Y. Wang, G. D. Stucky, R. L. Lambert, Jr., C. K. Haas, and D. Seyferth, J. Am. Chem. SOC., 98, 1779 (1976), and references cited therein. (7) D. Seyferth, R . L. Lambert. Jr., and M. Massol, J. Organomet. Chem., 88, 255 (1975). (8) C. Wentrup, Top. Curr. Chem., 62, 173 (1976). (9) M. Jones, Jr., Acc. Chem. Res., 7,415 (1974). (10) A. G. Brook and P. F. Jones, Can. J. Chem., 47, 4353 (1969). (1 1) J. P. Picard. R . Calas, J. Dunogues, and N. Duffaut, J. Organomet. Chem., 26, 183 (1971). (12) G. C. Levy and G. L. Nelson, '113CNMR for Organic Chemists". Wiley-lnterscience, New York, N.Y., 1972, p 30 ff. (13) J. B. Stothers, "Carbon-I3 NMR Spectroscopy", Academic Press, New York, N.Y., 1972, p 158 ff. (14) D. Seyferth and D. C. Annarelii, J. Am. Chem. Soc., 97, 7162 (1975). (15) H. Shechter, private communication. See also footnote 13a in M. R . Chedekel, M. Skwlund, R. L. Kreeaer. and H.Shechter, J. Am. Chem. Soc., 98, 7846 (1976).(16) D. N. Roark and L. H. Sommer, J. Chem. Soc., Chem. Commun., 167 (1973). (17) T. J. Barton, E. A. Kline, and P. M. Darvey, Third International Symposium on Organosilicon Chemistry, Madison, Wis., 1972. (18) R . D. Bush, C. M. Golino, G. D. Homer, and L. H. Sommer, J. Organomet. Chem., 80, 37 (1974). (19) J. H.Fwd. C. D. Thompson. andC. S. Marvel, J. Am. Chem. Soc., 57, 2619 (1945). (20) G. L. Closs and R. A . Moss, J. Am. Chem. Soc., 86, 4042 (1964). (21) G. P. Newsoroff and S. Sternhell, Aust. J. Chem., 19, 1667 (1966). (22) G. M. Kaufman, J. A. Smith, G. G. Vanderstouw, and H. Shechter, J. Am. Chem. SOC.,87, 935 (1965).
Effects of Charge Delocalization on Hydrogen Bonding to Negative Ions and Solvation of Negative Ions. Substituted Phenols and Phenoxide Anions John B. Cumming, Margaret A. French, and Paul Kebarle* Contributionfrom the Department of Chemistry, Unicersity of Alberta, Edmonton, Canada T6G 2G2. Receiced March 24, 1977
Abstract: Measurements of gas phase ion equilibria with a pulsed electron beam high ion source pressure mass spectrometer lead to the AGO and AH" values for the reactions AHCI- = AH CI- and AHCI- = A- t HCI, where A H stands for substituted phenols. These results show that the hydrogen bond energies in AHCI- (for dissociation to CI- HA) increase linearly with the gas phase acidity of the phenols. The hydrogen bond energies for dissociation of AHCI- to A- t HCI increase linearly with increase of the (gas phase) basicity of the phenoxide ions A-. Comparison with investigations of the AHCI- complexes in acetonitrile by Kolthoff and Chantooni shows that the substituent effects in acetonitrile are strongly attenuated. I t is suggested that in acetonitrile the extent of proton transfer on formation of AHCI- from A H and CI- is much smaller than that occurring in the gas phase. Proton transfer leads to charge dispersal in AHCI' which would decrease the solvation of AHCI-. The decrease of hydrogen bonding in A--HCI with decreasing basicity of A-, i.e., increasing acidity of A H , is related to a n expected decrease of hydrogen bonding in A--(HOH), which is responsible for the attenuation of the acidities of phenols in aqueous solution.
+
Earlier studies of the gas phase equilibrium BHR- = B-
+ HR
(1)
executed in this l a b o r a t ~ r y l -showed ~ that the strength of the hydrogen bond (Le., AH,' and AGIO) increased with the (gas phase) acidity of HR and the basicity of B-. For example, study of the series3 where B- = C1- and H R = HOH, C H 3 0 H , (CH3)3COH, C13CH, C ~ H S O H ,CH3COOH, HCOOH gave AH!' and AGIO, values which were found to increase in the order shown above, which is also the order of increasing gas phase acidity.6 In another series4 where H R = H 2 0 was kept constant and B- was changed, it was found that the hydrogen bond strengths increased almost linearly with the basicity of B- = I-, Br-, NO3-, Cl-, NOz-, F-, OH-.
+
Results involving positive ions BlH+ like H3O+, CH3OH2+, (CH3)20H+, NH4+, CH3NH3+, etc., and molecules B2 like H20, C H l O H , (CH3)20, NH3, CH3NH2, etc., also showed7-I0 that the strength of the hydrogen bond in B I HB2+ increases with the acidity of B I H + and the basicity of B2. The existence of homoconjugated (AHA-) and heteroconjugated (BHA-) complexes in solution has been long known. These complexes are particularly stable in aprotic solvents and have been subjected to numerous studies' I in these media. On the whole, considerable parallels exist regarding stabilities of the complexes in the gas p h a ~ e l and - ~ in aprotic solvents. Thus, GordonI2 found that the stabilities of the complexes AHA- (log K values i n acetonitrile, AH benzoic acids) increased linearly with the aqueous acidities of the acids
Cumming, French, Kebarle
/
Effects of Charge Delocalization on Hydrogen Bonding
7000 A H . Other s t ~ d i e s l ~have , ~ ~shown , ~ ~ that the stability of B-HA generally increases with the basicity of B- and the acidity of HA. However, the results in aprotic solutions are considerably more complex than those in the gas phase. This is caused by specific solvent differences and ion pairing, Le., dependence on the nature of the cation used. Furthermore since the acidity orders of the proton donors A H may be reversed in different solvents, a clear correlation between acidity of A H and stability of B--HA can often not be established. The same holds true for the relationship with the basicity of the proton acceptor B-. The systems studied in the present work involve the interactions of CI- with H A where H A stands for substituted phenols. The acidities in substituted phenols in the gas phase were recently determined in this 1aborat0ry.l~Furthermore the acidity order of substituted phenols in the gas phase, aqueous solution, and the aprotic solvent dimethyl sulfoxide is essentially the sameI5.l6so that a good comparison can be made between the gas phase results for the complexes CI-HA and recent quantitative results for the same system in acetonitrile.'' As was pointed o u t earlier,ls-18 results on the stability of complexes B--HA are of interest not only from the standpoint of hydrogen bonding interactions. They also have a bearing on ionic solvation by protic solvents. As an example we can consider the solvation of substituted phenoxide anions. The recent measurement of the gas phase acidities of substituted phenols15 allowed a detailed comparison of the gas phase and aqueous acidities to be made. It was foundl5 that a linear relationship between gas phase and aqueous acidities of the phenols existed. However, the substituent effects in aqueous solution were attenuated by a factor of seven. Evidently a linear relationship with such a large attenuation can result only if there is a proportional relationship between the intrinsic, electronic effect of the substituent and an opposing solvation effect caused by the same substituent. It has been long recognized that the unfavorable solvation effect of an acidifying substituent lies mostly in the reduction of the solvation energy of the substituted anion. For the present case, a Born cycle analysis involving the gas phase and aqueous acidities of phenols done by ArnettI9 has shown that a substituent that increases the gas phase acidity decreases the solvation energy of the substituted phenoxide ion by a similar amount. It is natural to connect this cancellation effect with the hydrogen bonding interactions discussed above. A substituent that increases the acidity of BH decreases the basicity of the conjugate base B-. Therefore the hydrogen bond with the first and subsequent 2o protic solvent molecules R H will be weaker. This means the solvation energy of B- will be lower. This decrease of hydrogen bonding energies between B- and (HR)n with increase of acidity of BH must be providing the automatic mechanism for the large attenuation of the substituent effect. Results on bonding between PhO- HCI obtained in the present work allow a further exploration of this relationship.
When more than one phenol is present also the exchange reaction 3 occurs. A,HCI-
+ A,H
= A,H
+ A,HCI-
(3)
Most of the ions diffuse to the walls of the ionization chamber where they become discharged. Some of the ions come to the vicinity of the ion exit slit through which they escape into an evacuated region where they are accelerated by electric fields, magnetically separated according to mass and detected with an ion-counting multiscaler system. The ion intensities can be observed for some 10 ms after the electron pulse. For the stated concentration conditions, the initial, kinetic stage in which the relative ion concentrations change, is short (l6only by a factor of
