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<?xml version="1.0" ?> <tei> <teiHeader> <fileDesc xml:id="-1"/> </teiHeader> <text xml:lang="en"> <front>Equilibrium Constants between Boron<lb/> Trifluoride Etherate and Carbonyl<lb/> Compounds in Chloroform Solution<lb/> J. J. Gajewski* and P. Ngernmeesri<lb/> Department of Chemistry, Indiana UniVersity, Bloomington, Indiana 47401<lb/> gajewski@indiana.edu<lb/> Received June 20, 2000<lb/> ABSTRACT<lb/> Equilibrium constants for boron trifluoride complexation with carbonyl compounds relative to diethyl ether were determined in CDCl 3 . With<lb/> benzaldehyde the equilibrium constant is 0.208. There is a 1.28-fold preference for BF 3 bonding to benzaldehyde-D over benzaldehyde-H. The<lb/> G + value for complexation with substituted benzaldehydes is −2.0. Equilibrium constants with cyclohexanone and isobutyraldehyde were<lb/> found to be not very different from those predicted from the heats of addition of BF 3 relative to that for diethyl ether. The heats of addition<lb/> can be correlated with Taft's values and with ′ values.<lb/></front> <body>Since various additions to aldehyde carbonyl groups involve<lb/> Lewis acid catalysis and often the use of boron trifluoride<lb/> etherate, 1 it would appear to be appropriate to determine the<lb/> equilibrium constant for relevant aldehydes under more or<lb/> less synthetic chemistry reaction conditions. Of further<lb/> concern should be the rate of the equilibration, substituent<lb/> effects, and, because of interest in studying the mechanism<lb/> of these additions by determination of deuterium kinetic<lb/> isotope effects, the effect of deuterium at the carbonyl carbon<lb/> on the equilibrium. 2<lb/> The benzaldehyde-BF 3 complex has been prepared and<lb/> characterized by X-ray crystallography. 3 It has the BF 3<lb/> moiety bound to an oxygen lone pair in a transoid (E)<lb/> fashion. The same is true of 2-heptenal in methylene chloride<lb/> solution. 4 Calorimetric determination of the heats of com-<lb/>plexation of benzaldehyde and diethyl ether with BF 3 in<lb/> methylene chloride solution indicate an enthalpy difference<lb/> of ca. 1 kcal/mol favoring boron trifluoride etherate over<lb/> benzaldehyde-boron trifluoride. 5 However, entropy differ-<lb/>ences, while not expected to be large, could modify this<lb/> preference. Additions to other aldehydes and unhindered<lb/> ketones were also found to be only slightly endothermic<lb/> relative to addition to diethyl ether, 5 and similar questions<lb/> about entropy differences can be posed.<lb/> In the hope of detecting a change in the chemical shift of<lb/> either benzaldehyde or diethyl ether upon reaction assuming<lb/> rapid exchange, which is the case for BF 3 in excess diethyl<lb/> ether, 6 1 equiv of benzaldehyde was added to 0.23 M boron<lb/> trifluoride (BF 3 ) etherate in deuteriochloroform solution.<lb/> While the spectrum resembled benzaldehyde with only a<lb/> slight downfield shift of the ortho protons 4 and a modestly<lb/> downfield shifted triplet for the diethyl ether methyl protons,<lb/> the methylene protons appeared as a broad singlet at 3.98<lb/> ppm downfield from TMS. Under these same conditions,<lb/> the spectrum of boron trifluoride etherate alone consists of<lb/> a quartet at 4.20 ppm and a triplet at 1.4 ppm. For<lb/> comparison, the methylene quartet of diethyl ether itself<lb/> appears at 3.478 ppm in deuteriochloroform at room tem-<lb/>perature. It therefore appears that boron trifluoride rapidly<lb/> interchanges its position on ether and benzaldehyde at normal<lb/> probe temperatures. Confirmation that the broad signal is<lb/> due to rapid equilibration comes from the spectrum recorded<lb/> at -60 °C in which the broad singlet is replaced by two<lb/> quartets separated by 0.70 ppm. Calculation of the equilib-<lb/></body> <listBibl>(1) (a) Cornel, V. in Encyclopedia of Reagents for Organic Synthesis;<lb/> Paquette, L. P., Ed.; John Wiley & Sons: New York, 1995 (under boron<lb/> trifluoride). (b) Hosomi, A.; Sakurai, H. Tetrahedron Lett. 1976, 16, 1295.<lb/> (c) Yamamoto, Y. Aldrichimca Acta 1987, 20, 45.<lb/> (2) Gajewski, J.J.; Bocian, W.; Harris, N. J.; Olson, L. P.; Gajewski, J.<lb/> P. J. Am. Chem. Soc. 1999, 121, 326.<lb/> (3) Reetz, M. T.; Hullmann, M.; Massa, W.; Berger, S.; Rademacher, P.<lb/> Heymanns, P. J. Am. Chem. Soc. 1986, 108, 2405.<lb/> (4) Denmark, S. E.; Almstead, N. G. J. Am. Chem. Soc. 1993, 115, 3133.<lb/> (5) Gal, J.-F.; Maria, P.-C. Prog. Phys. Org. Chem. 1990, 17, 159.<lb/> (6) Fratiello, A.; Onak, T. P.; Schuster, R. E. J. Am. Chem. Soc. 1968,<lb/> 90, 0, 1194. See also the following: Gillespie, R. J.; Hartman, J. S. Can.<lb/> J. Chem. 1968, 46, 2147. Childs,. R. F.; Mulholland, D. L.; Nixon, A. Can.<lb/> J. Chem. 1982, 60, 801.<lb/> </listBibl> <front>ORGANIC LETTERS<lb/> 2000 Vol. 2, No. 18<lb/> 2813-2815<lb/> </front> <front>10.1021/ol000164k CCC: $19.00 © 2000 American Chemical Society<lb/> Published on Web 08/09/2000<lb/></front> <body>rium constant from the observed chemical shift relative to<lb/> those of complexed and free diethyl ether at room temper-<lb/>ature for reaction 1 gave a value of 0.208 ( 0.011, resulting<lb/> from eight separate experiments. The relative heats of<lb/> reaction of BF 3 with both benzaldehyde and diethyl ether in<lb/> methylene chloride solution would suggest an equilibrium<lb/> constant of 0.22 if there were no entropy difference and no<lb/> solvent effect, and this is remarkably close to the value<lb/> observed directly in chloroform.<lb/> Similar observations were made with benzaldehyde-<lb/>carbonyl-D. However, the methylene protons of diethyl ether<lb/> appeared at slightly higher field: 3.95 ppm. When seven<lb/> experiments were averaged, the equilibrium constant was<lb/> determined to be 0.266 ( 0.020. This allowed calculation<lb/> of the equilibrium isotope effect or fractionation factor for<lb/> deuterium between protio and deuterio benzaldehyde (reac-<lb/>tion 2) to be 0.78 ( 0.07. Thus, complexation of deuterio<lb/> benzaldehyde is favored over protio benzaldehyde by a factor<lb/> of 1/0.78 or 1.28.<lb/> Confirming the magnitude and direction of the fraction-<lb/>ation factor for reaction 2 is a 6-31G*/MP2 geometry<lb/> optimization and harmonic frequency calculation 7a for ac-<lb/>rolein-H and -D and acrolein-H and -D BF 3 . 7b When these<lb/> unweighted frequencies where inserted into the Bigeleisen<lb/> equation characterizing a fractionation factor equilibrium, 7c<lb/> a value of 0.86 was obtained. The major contributor to this<lb/> equilibrium isotope effect appears to be the change in the<lb/> carbonyl-H stretching frequency (from 2997 cm -1 in acrolein<lb/> to 3181 cm -1 , uncorrected, in the complex). The recently<lb/> determined secondary deuterium kinetic isotope effects for<lb/> hydride reductions and Grignard and lithium reagent addition<lb/> to benzaldehyde assumed that rate-determining complexation<lb/> of the Lewis acid portion of the these reagents would result<lb/> in no large effect. 2 Clearly this is not the case with the strong<lb/> Lewis acid examined here. Whether or not a large effect<lb/> would be observed with metal cations has yet to be<lb/> determined.<lb/> Potentially complicating the analyses described above is<lb/> the possibility that there may be complexation of BF 3 with<lb/> deuteriochloroform solvent because the methylene protons<lb/> of diethyl ether appear at approximately 0.02 ppm lower field<lb/> in 1. M BF 3 -etherate in CDCl 3 . When 1 equiv of benz-<lb/>aldehyde was added, the shift observed allowed calculation<lb/> of an equilibrium constant which was 0.188. In effect, there<lb/> was less free diethyl ether produced both before and after<lb/> addition of benzaldehyde in the more concentrated solution.<lb/> However, the equilibrium constants determined at the two<lb/> concentrations are within experimental error which is<lb/> substantially magnified relative to the difference in the<lb/> chemical shifts observed because of the squaring of the ratios<lb/> to obtain the equilibrium constants. Further, the heat of<lb/> addition of BF 3 to methylene chloride was found 5 to be -2.4<lb/> kcal/mol which is much less than the -18.8 kcal/mol for<lb/> the addition to diethyl ether in methylene chloride solvent<lb/> so the difference in equilibrium constants is on the order of<lb/> 7 × 10 11 ; therefore, even a concentration difference of a<lb/> factor of 1000 should have little effect on pushing the<lb/> equilibrium toward a CDCl 3 -BF 3 complex.<lb/> The Hammett F + value was determined to be -2.0 (r )<lb/> 0.999) from the equilibrium constants for p-anisyl (7.24),<lb/> p-tolyl (0.838), and p-chloro (0.112) benzaldehyde at the<lb/> same concentration as benzaldehyde itself in chloroform. The<lb/> equilibrium constant for the p-nitro derivative was also<lb/> determined, but its very small value (0.009) is subject to<lb/> very large error so it was excluded from the analysis (its<lb/> inclusion provides a F + value of -1.86; r ) 0.998). The F +<lb/> value is relatively small compared with that for solvolysis<lb/> of cumyl derivatives (-4 to -5). 8 This is not inconsistent<lb/> with the notion that the carbonyl carbon is strongly positively<lb/> charged even before complexation.<lb/> The effect of methylene chloride and benzene solvent on<lb/> the equilibrium constant for benzaldehyde complexation with<lb/> boron trifluoride etherate was determined at both 0.23 and<lb/> 1.0 M. The equilibrium constants determined in methylene<lb/> chloride-D 2 were 0.16 ( 0.015 and 0.14 ( 0.01, respectively,<lb/> and the equilibrium constants determined in benzene-D 6 were<lb/> 0.15 ( 0.05 and 0.10 ( 0.03, respectively. The potential<lb/> origins of these differences are multitudinous; however, the<lb/> actual changes in chemical shift are not large, indicating that<lb/> the relative concentrations do not change by more than 10%,<lb/> i.e., 70:30 vs 75:25 so it is not profitable to pursue these<lb/> effects at this time.<lb/> Finally, the equilibrium constants for complexation of<lb/> boron trifluoride etherate with cyclohexanone and iso-<lb/>butyraldehyde were determined in chloroform-D to be 0.276<lb/> and 0.204, respectively. These equilibrium constants are not<lb/> very different than those suggested from the heats of addition<lb/> of BF 3 in methylene chloride solution to acetone and<lb/> acetaldehyde relative to that of diethyl ether: + 0.6 kcal/<lb/> mol and +2.1 kcal/mol, respectively, which lead to equi-<lb/>librium constants of 0.36 and 0.029, respectively. Therefore,<lb/> the data provided in ref 5 would appear to be a reasonable<lb/> source for relative complexation preferences with BF 3 in<lb/> solution. Also of interest is the fact that the relative<lb/></body> <listBibl>(7) (a) Gaussian 98. Frisch, M. J.; Trucks, G. W.; Schlegel, H. B.;<lb/> Scuseria, G. E.; Robb, M. A.; Cheeseman, J. R.; Zakrzewski, V. G.;<lb/> Montgomery, J. A., Jr.; Stratmann, R. E.; Burant, J. C.; Dapprich, S.; Millam,<lb/> J. M.; Daniels, A. D.; Kudin, K. N.; Strain, M. C.; Farkas, O.; Tomasi, J.;<lb/> Barone, V.; Cossi, M.; Cammi, R.; Mennucci, B.; Pomelli, C.; Adamo, C.;<lb/> Clifford, S.; Ochterski, J.; Petersson, G. A.; Ayala, P. Y.; Cui, Q.;<lb/> Morokuma, K.; Malick, D. K.; Rabuck, A. D.; Raghavachari, K.; Foresman,<lb/> J. B.; Cioslowski, J.; Ortiz, J. V.; Baboul, A. G.; Stefanov, B. B.; Liu, G.;<lb/> Liashenko, A.; Piskorz, P.; Komaromi, I.; Gomperts, R.; Martin, R. L.;<lb/> Fox, D. J.; Keith, T.; Al-Laham, M. A.; Peng, C. Y.; Nanayakkara, A.;<lb/> Gonzalez, C.; Challacombe, M.; Gill, P. M. W.; Johnson, B.; Chen, W.;<lb/> Wong, M. W.; Andres, J. L.; Gonzalez, C.; Head-Gordon, M.; Replogle,<lb/> E. S.; Pople, J. A. Gaussian, Inc., Pittsburgh. PA, 1998. </listBibl> <note place="footnote">(b) Previous work<lb/> in our laboratory (Harris, N. J. J. Phys. Chem. 1995, 99, 14689) revealed<lb/> that this level of theory was necessary to reproduce the fractionation factors<lb/> for deuterium between small molecules, and it became less accurate with<lb/> larger molecules. Therefore, our calculations were performed at as high a<lb/> level as possible on as small a model system as possible rather than<lb/> performing a suboptimal calculation on benzaldehyde and benzaldehyde-<lb/>BF 3 with the resources available. </note> <listBibl>(c) Lowry, T. H.; Richardson, K. S.<lb/> Mechanism and Theory in Organic Chemistry, 3rd ed.; Harper Collis<lb/> Publishers: New York, 1987; Chapter 2, Appendix 2.<lb/> </listBibl> <note place="footnote">(8) Reference 7, Chapter 4, Table 4.19.<lb/></note> <body>PhCHdO + BF 3 ‚OET 2 / PhCHdO‚BF 3 + OEt 2 (1)<lb/> PhCHO + PhCDO‚BF 3 / PhCHO‚BF 3 + PhCDO (2)<lb/> <page>2814</page> <note place="footnote">Org. Lett., Vol. 2, No. 18, 2000<lb/> </note> complexation enthalpies for a wide range of solvent mol-<lb/>ecules besides methylene chloride, nitromethane, benzonitrile,<lb/> acetonitrile, methyl acetate, acetone, diethyl ether, 1,4-<lb/>dioxane, tetrahydrofuran, DMSO, DMF, and HMPA, re-<lb/>ported in ref 5 can be correlated with the Taft values alone<lb/> (r ) 0.94) 9 and with ′ values (r ) 0.894). 10 The fact that<lb/> Taft values provide a reasonable correlation is probably<lb/> the result of the fact that these parameters were obtained<lb/> from the spectroscopy of single molecule-base interactions.<lb/> Single molecule-BF 3 interactions are being measured in the<lb/> calorimetry experiment as opposed to a bulk solvent interac-<lb/>tion with BF 3 . More remarkable is the correlation with ′<lb/> values which were derived from potassium ion transfer data<lb/> between bulk solvents. 11<lb/> </body> <div type="acknowledgement">Acknowledgment. We thank the Department of Energy<lb/> for support of this work. P.N. thanks the Thai Government<lb/> for an Undergraduate Scholarship to Indiana University. We<lb/> also thank Marty Pagel, Ulrike Werner-Zwanziger, and Jeff<lb/> Frey in the NMR facility of Indiana University for their help.<lb/> OL000164K<lb/></div> <listBibl>(9) For an excellent summary, see: Abraham, M. H.; Grellier, P. L.;<lb/> Abboud, J.-L. M.; Doherty, R. M.; Taft, R. W. Can, J. Chem. 1988, 66,<lb/> 2673.<lb/> (10) Gajewski, J. J. J. Org. Chem. 1992, 57, 5500. Gajewski, J. J.;<lb/> Brichford, N. J. In Structure and ReactiVity in Aqueous Solution; Cramer,<lb/> C. J., Truhlar, D. G., Eds.; ACS Symposium Series #568; American<lb/> Chemical Society: Washington, DC, 1994; p 229.<lb/> </listBibl> <note place="footnote">(11) It may be of interest to note that complexation of hydrogen ion<lb/> with individual solvent molecules in the gas phase reported in ref 5 can be<lb/> well-correlated with the combination of ′ and the Kirkwood-Onsager<lb/> function of dielectric, i.e., ( -1)/(2 + 1) and the polarizability parameters<lb/> of ref 7, Chapter 2, Table 2.15, footnote c. However, complexation of lithium<lb/> and potassium cation with individual solvent molecules in the gas phase<lb/> (ref 5) can be well-correlated with only the combination of the Kirkwood-<lb/>Onsager function and the polarizability parameters, that is, or ′ factors<lb/> do not contribute in these two cases.<lb/></note> <note place="footnote">Org. Lett., Vol. 2, No. 18, 2000<lb/> </note> <page>2815</page> </text> </tei>