bond: A fascination forever!
In 1920, Latimer and
Rodebush conceptualized the hydrogen bond (H bond)1. Almost eight decades have
passed and the H bond is continuing to fascinate researchers in the fields of physics,
chemistry, biology and materials sciences. Literally, thousands of papers are published
every year on H bonds (A Chemical Abstract search for the year 1998 with
hydrogen bond as the key word resulted in 3707 documents and 13436
occurrences!). Some of the recent publications deserve a closer look as they address some
fundamental issues on H bonds.
Hydrogen bond has been generally thought of as
originating from electrostatic interactions such as interaction between two dipoles.
Electron sharing or covalent nature was considered not to be important. This view was
strengthened by the fact that electrostatic models, such as the one by Buckingham and
Fowler, were able to explain/predict the H bond geometry fairly accurately2.
However, in 1949 Pauling3 had empirically calculated that H bond would have 5%
covalent character in an OH-----O bond. It has taken almost 50 years for
experimentalists to verify Paulings empirical estimate. Two techniques, NMR and
Compton scattering, have recently given convincing evidence that Pauling was right! The
NMR experiment can, in principle, give some estimate of covalent nature if one could see
the spinspin coupling between the two other nuclei involved in the H bond. It has
been done now.
Dingley and Grzesiek4
used a novel pulse sequence in their NMR experiment to observe J coupling between two 15N
nuclei (2JNN) connected through an H bond, NH-----N, in
nucleic acidbase pairs. Spinspin or scalar coupling, also known as J coupling,
is generally observed between nuclei that are connected by covalent bonds. The magnitude
of the J coupling depends on the extent of the orbital overlap in a bond. It was widely
believed that the covalent nature of the H bonds would be too insignificant and it may at
best produce a small splitting of << 1 Hz due to J coupling5.
Dingley and Grzesiek observed a splitting of 7 Hz! This was the first time J coupling
could be observed between two nuclei through an H bond, which implies that the H bond has
some covalent character (Figure 1). The key to their experimental success was the fact
that they were using a uniformly 13C/15N enriched sample. Since this
observation, several groups have found similar and higher J coupling between nuclei
connected through an H bond. Golubev et al.6 have observed a splitting
of 96 Hz between 19F and 15N in an
Figure 1. Schematic representations
of the hydrogen bond. a, The s bond is depicted as the overlap of the
hydrogen s-orbital (represented by the circle) and a nitrogen sp3
orbital. The H bond is depicted as the interaction between the positive end of the dipole
in NH and the negative end of the other N. b, The lone pair electrons
are shown spending a non-negligible amount of time in the vicinity of the hydrogen, though
the electrostatic interaction is still present and dominant. Though it is a small effect,
it is significant as it leads to an essentially continuous wave function between the two N
atoms. This causes the spinspin coupling in the NMR experiment and the two
additional peaks in the Compton scattering experiment. (Figure adapted from ref. 9 with
FH-----N bond of an acidbase complex.
Grzesieks group7 has observed a coupling of 15 Hz through 3 bonds, 3JNC,
in proteins with NH-----OC hydrogen bond. These observations have not yet been
utilized to determine the extent of covalent character in an H bond. However, such an
estimate has been made with Compton scattering data in common ice.
A Compton scattering experiment by US, French and
Canadian physicists, utilizing the European Synchrotron Radiation Facility in Grenoble has
been recently reported8. Just as Rayleigh scattering of UV-visible radiation,
we have Thomson scattering of X-rays which forms the basis of standard X-ray
crystallography. Compton scattering is an X-ray analog of the Raman scattering, which is
more familiar to chemists (To be fair, Raman discovered the optical analog of Compton
scattering for which he won the Nobel Prize two years after Compton did). Compton
scattering leads to the momentum space profile of the wave functions of the valence
electrons in a molecule. One can Fourier transform this to find how the electron density
is distributed in the molecule. In ice, the authors found three peaks at 0.89, 1.72 and
2.85 ┼. These distances give extents to which wave functions of the bonding
electrons are distributed. For example, the peak at 0.89 ┼ corresponds to the
OH bond and indicates that wave functions of the electrons involved in this bond
extend over this distance. The 2.85 ┼ peak corresponds to the nearest neighbour
OO distance and 1.75 ┼ peak corresponds to H bond distance. These two peaks
are new and significant, as they indicate that there are bonding wave functions extending
in space between OO, with part of it in the H bond, as in Figure 1. This clearly
demonstrates that there is electron sharing in the H bond. The results indicate ╗
10% covalent character in the OH----O hydrogen bond. This is in
remarkable agreement with Paulings estimate, considering the fact that Pauling used
empirical relations based on intuition. This estimate is as yet unpublished but reported
in a news article in Nature: Structural Biology9.
Recently, Hobza and coworkers from Heyrovsky
Institute of Physical Chemistry, Czech Republic, claim to have observed what they call an
anti H bond10. In general, H bond formation leads to lengthening of the donor
XH bond and a concomitant red shift in the XH stretching
frequency. However, for two complexes namely flouroform-ethylene oxide11 and
chloroform-fluoro-benzene12, the CH bond which is involved in the H bond
appears to show a blue shift and bond contraction! The authors have been successful in
using ab initio calculations that support their experimental observations. They
conclude that dispersion forces play a major role in these interactions. In a normal H
bond, the bond lengthening in the donor leads to enhancement of dipole moment and more
stability in the dipoledipole interaction. Dispersion energy is proportional to the
higher power of reciprocal distance between the centers of mass of both sub systems. To
minimize the distance and maximize the attraction, it is advantageous to contract the
CH bond of the proton donor. This contraction, they claim, can lead to shorter
intermolecular distances and larger dispersion energy. According to these workers, the
anti-H bond is not limited to the gas phase complexes. The CH----O contact in
adeninethymine base pair also exhibits such character13. The T-shaped
benzene dimer too has anti H bond! This observation, if true, is very important as the
benzene dimer is the simplest model for aromaticaromatic interactions in biological
Ever since the H bond was conceptualized, chemists
(knowing their love for Periodic Table) must have wondered about the possibilities of
other atoms exhibiting similar interactions. Two recent papers in the Angewandte Chemie
appear to give convincing evidence for halogen bonds! Resnati and coworkers14
have reported single crystal X-ray studies on (S)-1,2-dibromohexafluoro-propane which was
crystallized in 100% enantiomeric excess from a solution of racemate with ()
sparteine hydrobromide. They suggest that the interaction between the compounds through a
halogen bond, analogous to an H bond, facilitates the preferential crystallization of one
form. More importantly, Legon of University of Exeter has reviewed15 rotational
spectroscopic studies on a series of HX and XY (X and Y halogens) complexes
with Lewis bases, B. These complexes were formed in gas
Figure 2. Experimentally observed
structures (C3v, not to scale) for H3N---HCl and H3N---Cl2.
The intermolecular stretching force constant for HCl complex is 17.6 N m1
and that for Cl2 complex is 12.7 N m1. In the gas
phase H3N---HCl exists as an H bonded complex and not as ammonium chloride. The
N---H distance is about 1.9 ┼ compared to the NH distance of 1 ┼ in
ammonia. The distance between N and the bonded chlorine is 2.7 ┼ in NH3Cl2
complex, which is about 0.6 ┼ shorter than the sum of van der Waals radii for N and
Cl atoms. Typical NCl bond distance, for example in NCl3, is about
1.7 ┼. The HCl bond forms an angle of 13.4░ with respect to the
intermolecular bond while the ClCl bond is at 7.5░ from the intermolecular bond.
Halogen bonds are in general more collinear than the H bond. (Figure adapted from ref. 15
phase by supersonic expansion and observed by
Fourier transform microwave spectrometer. Thus they are direct observation of the halogen
bond. Legons group has done systematic studies on these complexes with
B = CO, C2H2, C2H4, HCN, H2S,
NH3, pyridine, furan, thiophen, cyclopropane, methylene cyclopropane, etc. and
HX/XY = HF, HCl, HBr, HI, F3, Cl2, Br2, BrCl,
ClF and ICl. The B---HX and B---XY complexes were found to be isostructural in all the
cases. Figure 2 shows the geometry of NH3---HCl and NH3---Cl2
complexes, both having C3v point group. Ironically, it has been found that the
B---XY complexes are more collinear than the H bonded B---HX. For example, in
thiiraneClF complex, the deviation from linearity is found to be 3.5 ▒
2░ while that in thiiraneHCl complex is found to be 21 ▒ 5░ .
Also, in several cases the B---XY bond has been found to be stronger than B---HY bond.
This extensive study clearly demonstrates the presence of a halogen bond.
Pimentel and McLellan wrote the first definitive
book on hydrogen bonds16. They have given a very practical definition of the
hydrogen bond which follows: A hydrogen bond is said to exist when 1) there is evidence
of a bond and 2) there is evidence that this bond specifically involves a hydrogen atom
already bonded to another atom. According to this definition, we see that both the H
bond and the anti-H bond can be classified as just H bond! Also, it justifies the halogen
bond description discussed above. What would be the extent of covalent character in H
bonds observed in various surroundings? What would be the nature of the halogen bond? It
is clear that in the next few decades, we will continue to see growing literature on the H
bond or intermolecular interactions, in general!
Note added in proof: Very recently, Scheurer
and BrŘschweiler17 have reported quantum chemical studies, using density
functional theory (DFT), on the spinspin couplings across H bonds. They conclude
that the Fermi contact terms, originating from orbitals with s bond character across the H
bond, make the major contributions to the spinspin couplings experimentally
- Latimer, W. M. and Rodebush, W. H., J. Am. Chem. Soc., 1920, 42,
- Buckingham, A. D. and Fowler, P. W., J. Chem. Phys., 1983, 79,
- Pauling, L., The Nature of the Chemical Bond, Cornell
University Press, Ithaca, NY, 1960, p. 453.
- Dingley, A. J. and Grzesiek, S., J. Am. Chem. Soc., 1998, 120,
- Cornilescu, G., Hu, J-S. and Bax, A.,
J. Am. Chem. Soc., 1999, 121, 2949; Borman, S., Chem. Eng. News, 10
May 1999, p. 36.
- Golubev, N. S. et al., Chem. Eur. J., 1999, 5,
- Cordier, F. and Grzesiek, S., J. Am. Chem. Soc., 1999, 121,
- Isaacs, E. D. et al., Phys. Rev. Lett., 1999, 82,
- Martin, T. W. and Derewenda, Z. S., Nature Struct. Biol.,
1999, 6, 403.
- Cubero E. et al., J. Phys. Chem., 1999, A103,
- Hobza, P. and Havlas, Z., Chem. Phys. Lett., 1999, 303,
- Hobza, P. et al., Chem. Phys. Lett., 1999, 293,
- Hobza, P. et al., J. Phys. Chem., 1998, A102,
- Resnati, G. et al., Angew. Chem. Int. Ed. Engl., 1999, 38,
- Legon, A. C., Angew. Chem. Int. Ed. Engl., 1999, 38,
- Pimentel, G. C. and McLellan, A. L., The Hydrogen Bond, W. H.
Freeman and Co., San Francisco, 1960.
- Scheurer, C. and BrŘschweiler, R.,
J. Am. Chem. Soc., 1999, 121, 8661.
ACKNOWLEDGEMENTS. I thank Profs.
K. L. Sebastian and A. G. Samuelson for stimulating discussions. I thank Prof.
P. Balaram for bringing refs 9 and 17 to my attention.
E. Arunan is in the Inorganic and Physical
Chemistry Department, Indian Institute of Science, Bangalore 560 012, India