A Brief History of the Hydrogen Bond - ACS Publications - American

Chapter 1

A Brief History of the Hydrogen Bond 1

Douglas A. Smith

Downloaded by GEORGE MASON UNIV on July 26, 2012 | http://pubs.acs.org Publication Date: May 5, 1994 | doi: 10.1021/bk-1994-0569.ch001

Department of Chemistry, University of Toledo, 2801 West Bancroft Street, Toledo, OH 43606-3390

Between the beginning of the twentieth century and the late 1930's many descriptions of interactions between hydrogen atoms in molecules and electronegative atoms to which these hydrogens were not covalently bound appeared in the literature. This was part of the general explosion of physical, structural and quantum chemistry that was providing a descriptive and explanatory underpinning for chemical phenomena as a whole. In 1920, Latimer and Rodebush, working on the structure and properties of water with G. N. Lewis at UC Berkeley, proposed: "[A] free pair of electrons on one water molecule might be able to exert sufficient force on a hydrogen held by a pair of electrons on another water molecule to bind the two molecules together.... Such an explanation amounts to saying that the hydrogen nucleus held between 2 octets constitutes a weak 'bond' (1)." This description, based on the Lewis dot formalism, is the first to truly call this interaction a bond. It was in 1939, however, with the publication of the first edition of Pauling's "The Nature of the Chemical Bond" (2) and its seminal chapter The Hydrogen Bond, that this phenomenon really reached acceptance in the main stream. In this treatise was the following description: "[U]nder certain conditions an atom of hydrogen is attracted by rather strong forces to two atoms, instead of only one, so that it may beconsidered to be acting as a bond between them. This is called the hydrogen bond" (pp. 449, Third Edition) Furthermore, Pauling put the hydrogen bond on firm quantum mechanical grounds: "[A] hydrogen atom, with only one stable orbital, cannot form more than one pure covalent bond and that the attraction of two atoms observed in hydrogen-bond formation must be due largely to ionic forces." (pp. 450-1, Third Edition) 1

Current address: Concurrent Technologies Corporation, 1450 Scalp Avenue, Johnstown, PA 15904 0097-6156/94A)569-0001$08.00A) © 1994 American Chemical Society In Modeling the Hydrogen Bond; Smith, D.; ACS Symposium Series; American Chemical Society: Washington, DC, 1994.

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MODELING THE HYDROGEN BOND

As an ionic interaction, hydrogen bonding was limited to occurring only with the most electronegative atoms. It remained for Pimentel and McClellan provided a broader, operational definition in their 1960 book, "The Hydrogen Bond" (5):


"A H Bond exists between a functional group A-H and an atom or a group of atoms Β in the same or a different molecule when (a) there is evidence of bond formation (association or chelation), (b) there is evidence that this new bond linking A-H and Β specifically involves the hydrogen atom already bonded to A." This rather broad definition leaves open the possibility of hydrogen bonding to groups of atoms (or the electrons associated with those atoms) which are not themselves highly electronegative, or the involvement of hydrogen atoms covalently bound to non-electronegative atoms such as carbon. As both experiment and theory have shown, this is indeed a more correct definition. In current practice, the terms acceptor and donor refer to those atoms that either accept or donate the hydrogen of the hydrogen bond, respectively. Typical donors include but are not limited to electronegative atoms with at least one attached hydrogen, such as oxygen, nitrogen, sulfur and the halogens. Even carbon, silicon and phosphorus, however, can behave as donors. More generally, donors are simply Bronsted acids. Acceptor atoms or groups are Lewis bases which range from simple electronegative atoms with lone electron pairs to the π electrons of an unsaturated or aromatic system, as well as ionic species such as oxides and halides. Other attached atoms or functional groups influence the capabilities of both the donor and acceptor. Even though a description of intramolecular hydrogen bonding appeared as early as 1906 in azo compounds (4) and 1913 for the interaction between a hydroxyl hydrogen and a carbonyl oxygen (5), hydrogen bonds were considered a linear phenomenon based primarily on experimental data from x-ray and neutron diffraction data (2). Recent experimental and theoretical data support this viewpoint, at least in some systems. For example, only linear hydrogen bonding is observed experimentally for water (6), and theoretical calculations predict a linear angle in the case of water dimers (7). Such evidence has led, unfortunately, to a common misconception that hydrogen bonds cannot occur unless the angle of association is close to 180°. There is, however, a significant body of evidence to the contrary. A recent survey of 1357 crystallographically observed intermolecular N-H—0=C bonds indicates that the donor angle in about 28% of the cases falls in the range 170° -180°, and 38% fall in the range of 160°-170°. The mean angle is 161.2°. For O - H - 0 bonds, the distribution is similar and the mean angle is 163.1° (8). A similar survey of 152 crystallographically observed intramolecular N-H--0=0 hydrogen bonds show that they are significantly less linear, with a mean angle of 132.5° (9). An Historical Interlude: Ljubljana. Yugoslavia. 1957 From 29 July through 3 August 1957, some of the most important and influential names in bonding and theory, including Pauling, Pople, Pimentel, and Coulson, among others, met in Ljubljana at the Symposium on Hydrogen Bonding. Perhaps the most important event of that meeting was the paper presented by Coulson, simply

In Modeling the Hydrogen Bond; Smith, D.; ACS Symposium Series; American Chemical Society: Washington, DC, 1994.


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titled "The Hydrogen Bond" (10). His paper has become a landmark in this field. Subtitled "A review of the present interpretations and an introduction to the theoretical papers presented at the Congress," it was a critical review of the status of theory in its attempts to understand and explain hydrogen bonding. In his paper, Coulson examined the electrostatic model of hydrogen bonding, which owed much to the work of Lennard-Jones and Pople (77) and discounted this model as insufficient to account for all of the phenomena associated with hydrogen bonding. Instead, Coulson divided the Coulombic attraction energy of the hydrogen bond into four causative factors: electrostatic, covalent, repulsive and dispersion contributions. Almost twenty years later, Kitaura and Morokuma developed a scheme within the framework of ab initio SCF theory in which they decompose the interaction energy of two molecules into five components: electrostatic (ES), polarization (PL), exchange repulsion (EX), charge transfer (CT), and coupling (MIX) (12,13). Application of this method to the origin of the hydrogen bond suggested that hydrogen bonds between neutral donors and acceptors were "strongly ES in nature, with a small but significant contribution of CT (14)." Furthermore, the uniqueness of hydrogen bonding was simply and unassumingly that "it always involves a moderately polar, short, and strong [Η-A] bond as the proton donor (14)" The major objection to the Morokuma method, that two of the intermediate wave functions used to calculate the ES and PL components violate the Pauli exclusion principle (13), has recently been addressed by an energy decomposition analysis based on the natural bond orbital (NBO) method (15,16). This natural energy decomposition analysis (NEDA) indicates that both ES and CT make significant and similarly strong contributions to the hydrogen bonding interaction energy, at least in the water dimer (77). The important molecular orbital interaction in hydrogen bonded systems is the two electron two orbital interaction between the antibonding σ* orbital corresponding to the donor Η-A bond and the highest occupied molecular orbital (HOMO) of the acceptor (18). This is the predominant contribution to the charge transfer component of the interaction energy (14,16) and corresponds to the dominant component of van der Waals forces at short range. Even though the maximum orbital overlap should occur for a colinear arrangement of A-H—B, the local symmetry of the σ* orbital is essentially spherical (ls ). Thus, the directionality of the hydrogen bonding interaction is not critical. Since 1957, a significant amount of theoretical work, compute cycles, and personnel time has been expended to examine the phenomenon of hydrogen bonding. Many empirical, semiempirical and ab initio molecular orbital studies have looked at hydrogen bonding systems and the nature of the hydrogen bond itself. Theoreticians have been examining many aspects of computational methods as applied to the study and accurate representation of hydrogen bonding. Recendy, for example, Del Bene has published extensively on the basis set effects (79) and Del Bene (20) and Davidson (27) have studied the electron correlation contribution to the computed hydrogen bond energies of many species. Parameterization of molecular mechanics force fields (22), inclusion of more accurate charges in molecular mechanics calculations (23) have been examined, and calibration of hydrogen bonding potentials based on organic crystal structures have appeared (24). Applications to specific molecular systems of interest, such as oxygen hydrogen bonding to hydrogen attached to carbon (25) and DNA base pairs (26), are becoming pervasive. Several H


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reviews have appeared, including excellent but now dated ones by Rao and Murthy (27) and Kollman and Allen (28). A recent book by Jeffry and Saenger, "Hydrogen Bonding in Biological Molecules" (29) gives an excellent compilation of more recent data, although it concentrates almost exclusively on crystallographic data. However, between 1957 and 1993, there has been no major symposium on modeling and theoretical methods for studying hydrogen bonding. The recent symposium on which this book is based, Modeling the Hydrogen Bond, sponsored by die Computers in Chemistry Division of the American Chemical Society, was put together in the hope that it would address many of the issues that have arisen since Ljubljana.


Literature Cited 1. Latimer and Rodebush, J. Am. Chem. Soc. 1920, 42, 1419. 2. Pauling, L. "The Nature of the Chemical Bond and the Structure of Molecules and Crystals: An Introduction to Modern Structural Chemistry," 3rd ed., Cornell University Press, Ithaca, NY, 1960, pp. 449 - 504. 3. Pimentel, G. C. and McClellan, A. L. "The Hydrogen Bond" W. H. Freeman and Co., San Francisco, 1960. 4. Oddo, G.; Puxeddu, E. Gazz. 1906, 36, 1. 5. Pfeiffer, P. Ann., 1913, 398, 137. 6. Dyke, T. R.; Muenter, J. S. J. Chem. Phys. 1974, 60, 2929. 7. Harrell, S. Α.; McDaniel, D. H. J. Am. Chem. Soc. 1964, 86, 4497. Morokuma, K.; Pedersen, L. J. Chem. Phys. 1969, 51, 3286. 8. Vinogradov, S. N.; Linnell, R. H. "Hydrogen Bonding," Van Nostrand Reinhold, New York, 1971. 9. Taylor, R.; Kennard, O.; Versichel, W. Acta Crystallogr. Β 1984, B40, 280. 10. Coulson, C. A. "The Hydrogen Bond. A review of the present interpretations and an introduction to the theoretical papers presented at the Congress" in Hydrogen Bonding, D. Hadzi and H. W. Thompson, Eds., Pergamon Press, London, 1959, pp. 339. 11. Lennard-Jones, J.; Pople, J. A. Proc. Roy. Soc. (London) 1951, 205A, 155, 163. 12. Morokuma, K. J. Chem. Phys. 1971, 55, 1236. Kitaura, K.; Morokuma, K. Int. J. Quantum Chem. 1976, 10, 325. 13. Morokuma, K. Acc. Chem. Res., 1977, 10, 294. 14. Umeyama, H.; Morokuma, K. J. Am. Chem. Soc. 1977, 99, 1316. 15. Foster, J. P.; Weinhold, F. J. Am. Chem. Soc. 1980, 102, 7211. 16. Reed, A. E.; Weinstock, R. B.; Weinhold, F. J. Chem. Phys. 1985, 83, 735. 17. Glendening, E. D.; Streitwieser, A. J. Chem. Phys. 1994, 100, 2900. 18. Rauk, A. "Orbital Interaction Theory of Organic Chemistry," John Wiley & Sons, Inc., New York, 1994, pp. 160-2. 19. Del Bene, J. E. Int. J. Quanum Chem.: Quantum Biol. Symp. 1987, 14, 27. 20. Del Bene, J. E.; Shavitt, I. Int. J. Quanum Chem.: Quantum Biol. Symp. 1989, 23, 445. 21. Racine, S. C.; Davidson, E. R. J. Phys. Chem. 1993, 97, 6367. 22. See, for example, Damewood, J. R., Jr.; Kumpf, R. Α.; Muhlbauer, W. C. F.; Urban, J. J.; Eksterowicz, J. E. J. Phys. Chem. 1990, 94, 6619.


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23. See, for example, Cornell, W. D.; Cieplak, P.; Bayly, C. I.; Kollman, P. A. J. Am. Chem. Soc. 1993, 115, 9620. 24. Gavezzotti, Α.; Filippini, G. J. Phys. Chem. 1994, 98, 4831. 25. See, for example, Turi, L.; Dannenberg, J. J. J. Phys. Chem. 1993, 97, 7899. 26. Gould, I. R.; Kollman, P. A. J. Am. Chem. Soc. 1994, 116, 2493. Del Bene, J. E. Int. J. Quanum Chem.: Quantum Biol. Symp. 1988, 15, 119. 27. Murthy, A. S. N.; Rao, C. N. R. J. Mol. Struct. 1970, 6,253. 28. Kollman, P. Α.; Allen, L. C. Chem. Rev. 1972, 72, 283. 29. Jeffrey, G. Α.; Saenger, W. "Hydrogen Boding in Biological Structures," Springer-Verlag, Berlin, 1991. 1994


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A Brief History of the Hydrogen Bond - ACS Publications - American

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