Hydrogen bond

• 1. Introduction

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  A hydrogen bond holds water together, and it does so with a force somewhere between 1 and 40 kcal/mol. That range tells a strange story. At its weakest, the interaction barely outdoes the faint pull of van der Waals forces. At its strongest, in the bifluoride ion HF2-, it reaches 161.5 kJ/mol and starts to behave almost like a covalent bond. This is one molecular interaction wearing many faces. It is the reason ice floats on the water it came from. It is the reason a strand of DNA can find its partner and split apart again to copy itself. And it is the reason a washed wool garment can lose its shape forever. The notation chemists use looks deceptively simple: Dn-H...Ac, a solid line for a polar covalent bond and three small dots for the bond in question. But behind those three dots lies a decades-long argument about what the bond actually is. Is it electrostatic, a matter of opposite charges drawn together? Or is it partly covalent, a genuine sharing of electrons? Who first noticed it, and why was it once so hard to believe in? How can one bond span the gap from liquid water to a fibre stiff enough to stop a blade? And why does counting these bonds, even in a single glass of water, turn out to be so difficult that careful scientists cannot agree on the number?

• 2. More Than Electrostatics

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  Charge transfer is what separates a hydrogen bond from a simple dipole-dipole attraction. The interaction described as nB to sigma-star-AH means electrons flow from the acceptor into an antibonding orbital, a feature of orbital overlap rather than mere opposite charges meeting. A hydrogen bond arises from a blend of three contributions: electrostatics, including multipole-multipole and multipole-induced multipole effects; covalency, the charge transfer through orbital overlap; and dispersion, the London forces. This makes it a resonance-assisted interaction, not a purely electrostatic pull. The hydrogen donor carries a protic hydrogen attached to an electronegative atom such as nitrogen, oxygen, or fluorine. That hydrogen can act as a Lewis acid, while the acceptor, bearing a lone pair, acts as the Lewis base. Acceptors include amines, carboxylates, and water itself. The IUPAC nomenclature names the electronegative atom not attached to the hydrogen the proton acceptor, and the one covalently bound to the hydrogen the proton donor. Liquids that show this behaviour, water among them, are called associated liquids. The resonance-assisted hydrogen bond, abbreviated RAHB, is a strong variety distinguished by pi-delocalization involving the hydrogen atom, a feature the electrostatic model alone cannot capture. It was proposed to explain unusually short distances that ordinary descriptions could not.

• 3. A Continuum Of Strength

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  Sulfur and chlorine mark the weaker end of the hydrogen-bonding spectrum. Even carbon can serve as a donor, particularly when the carbon or one of its neighbours is electronegative, as in chloroform, aldehydes, and terminal acetylenes. These non-traditional interactions are about 1 kcal/mol, weak yet ubiquitous, and they shape the structures of many materials. The term hydrogen bond is reserved for well-defined, localized interactions with significant charge transfer, such as those in DNA base pairing or ice. The phrase hydrogen-bonding interactions covers the weaker, more dynamic, or delocalized cases, such as liquid water, lipid membranes, protein-protein interactions, or faint C-H...O contacts. Typical enthalpies in vapor span a wide range: 161.5 kJ/mol in HF2-, 29 kJ/mol for water-ammonia, 21 kJ/mol for water-water and alcohol-alcohol, 13 kJ/mol for ammonia-ammonia, and 8 kJ/mol for water-amide. One classification scheme calls bonds of 15 to 40 kcal/mol strong, 5 to 15 kcal/mol moderate, and 0 to 5 kcal/mol weak. The strength of intermolecular bonds is most often judged from equilibria in solution, while intramolecular bonds are studied through equilibria between conformers. Crystallography is the most important identification method, sometimes joined by NMR spectroscopy. A donor-acceptor distance smaller than the sum of the van der Waals radii signals a stronger bond. Bonds involving C-H are both very rare and weak.

• 4. Reading The Bond

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  An acidic proton in the enol tautomer of acetylacetone appears at 15.5 in the 1H NMR spectrum, roughly 10 ppm downfield of a conventional alcohol. Strong hydrogen bonds reveal themselves through such downfield shifts. In the infrared spectrum, hydrogen bonding shifts the stretching frequency to lower energy, a decrease that reflects a weakening of the bond. Certain so-called improper hydrogen bonds defy this pattern, showing instead a blue shift of the stretching frequency and a shortening of the bond length. The bond can also be measured through vibrational mode shifts of the acceptor. The amide I mode of backbone carbonyls in alpha-helices moves to lower frequencies when those carbonyls form bonds with side-chain hydroxyl groups. Geometry varies with the donor. For a hydrofluoric acid donor, experiments give a linear 180-degree angle at a C-triple-bond-N acceptor, 120 degrees at a trigonal planar C=O, 46 degrees at a pyramidal H-O, 89 degrees at a pyramidal H-S, and 145 degrees at a trigonal S=O. Structural details place the donor-hydrogen distance near 110 pm, while the hydrogen-acceptor distance runs from about 160 to 200 pm. In water, the typical bond length is 197 pm. Variable-temperature infrared spectroscopy can track how these bonds respond to heat, including in protic organic ionic plastic crystals, where the sudden weakening of bonds during a solid-solid phase transition appears coupled to the onset of rotational disorder of the ions.

• 5. The Long Argument

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  Linus Pauling proposed that hydrogen bonds had a partial covalent nature, and the idea stayed controversial for years. The resolution came when NMR techniques demonstrated information transfer between hydrogen-bonded nuclei, something possible only if the bond carried some covalent character. Interpretations of the anisotropies in the Compton profile of ordinary ice also claimed the bond was partly covalent, though that reading was challenged and later clarified. The concept itself was once hard to accept. Pauling credited T. S. Moore and T. F. Winmill with the first mention of the hydrogen bond in 1912. Moore and Winmill invoked it to explain why trimethylammonium hydroxide is a weaker base than tetramethylammonium hydroxide. The description in its better-known setting, water, arrived in 1920 from Latimer and Rodebush. In that paper they credited a colleague at their laboratory, Maurice Loyal Huggins, writing that Mr. Huggins of this laboratory in some work as yet unpublished, has used the idea of a hydrogen kernel held between two atoms as a theory in regard to certain organic compounds. The definition kept widening over the following century. In 2011 an IUPAC Task Group recommended an evidence-based definition, published in Pure and Applied Chemistry, defining the bond as an attractive interaction between a hydrogen atom from a molecule in which X is more electronegative than H and an atom or group in which there is evidence of bond formation.

• 6. Why Ice Floats

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  Four is the magic number for a water molecule. Its oxygen carries two lone pairs and two hydrogen atoms, so each molecule can bond to as many as four neighbours, two through its lone pairs and two through its hydrogens. The simplest case is the water dimer, a single bond between two molecules, often used as a model system. In bulk liquid the bonds proliferate and build a network. This network gives ice an open hexagonal lattice, which is why the solid phase is less dense than the liquid and floats on it, unlike most substances. The same abundance of bonds, set against water's low molecular mass, drives its anomalously high boiling point, along with its high melting point and viscosity compared to similar liquids that lack such bonds. The count is restless. From TIP4P simulations at 25 degrees C, each molecule averages 3.59 hydrogen bonds; at 100 degrees C the figure falls to 3.24 as motion increases and density drops, and at 0 degrees C it rises to 3.69. Another study found a much smaller average of 2.357 at 25 degrees C, a reminder that defining and counting these bonds is not straightforward. Each bond between water molecules lasts on average about 10 picoseconds, or 10 to the minus 11 seconds. A single hydrogen atom can even split its allegiance in a bifurcated bond, two-forked between two acceptors, a configuration suggested as an essential step in water reorientation. Hydrogen fluoride faces the opposite shortage: three lone pairs on fluorine but only one hydrogen, so it can form only two bonds, while ammonia has three hydrogens and only one lone pair.

• 7. The Architecture Of Life

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  DNA owes its double helix largely to hydrogen bonding between base pairs, alongside pi stacking, the bonds linking one complementary strand to the other and enabling replication. The bonds are not all equal. Quantum chemical calculations of compliance constants reveal that the central interresidue bond between guanine and cytosine is much stronger than the one between adenine and thymine. In proteins, hydrogen bonds form between backbone oxygens and amide hydrogens, and the spacing of the residues decides the shape. A regular spacing between positions i and i + 4 builds an alpha helix; a tighter spacing between i and i + 3 builds a 3-10 helix; and two strands joined through alternating residues build a beta sheet. These bonds also help shape tertiary structure through interactions of R-groups. Bifurcated systems are common in alpha-helical transmembrane proteins, where a backbone amide acceptor on residue i meets two donors from residue i + 4. The energy preference of the bifurcated hydroxyl system is -3.4 kcal/mol and the thiol system 40 kcal/mol, giving polar side-chains such as serine, threonine, and cysteine a bonding partner inside hydrophobic membranes. Because water competes for donors and acceptors, bonds within or between dissolved solutes are almost always unfavourable compared to bonds with water. Protective osmolytes such as trehalose and sorbitol shift the folding equilibrium toward the folded state, and molecular dynamics simulations suggest they work by modifying bonds in the protein hydration layer. A backbone bond left incompletely shielded from water is called a dehydron, which promotes the removal of water and so strengthens the electrostatic interaction it harbours.

• 8. Strength In Fibres

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  Wool recoils when stretched because hydrogen bonds hold the protein fibre together, yet washing at high temperatures can break those bonds permanently and rob a garment of its shape. The same family of bonds, perhaps only 5 percent as strong as the covalent bonds in a polymer backbone, governs the behaviour of many materials. That weakness is useful: the bonds can be broken by chemical or mechanical means while the backbone survives intact, a hierarchy in which covalent beats hydrogen beats van der Waals. Cellulose and its derived fibres, cotton and flax, depend on them. In nylon, bonds between the carbonyl and the amide NH link adjacent chains and supply mechanical strength. In aramid fibre, bonds stabilize the linear chains laterally while the chain axes align along the fibre axis, making the fibres extremely stiff and strong. The same networks make both polymers sensitive to atmospheric humidity, since water can diffuse in and disrupt them, and nylons prove more sensitive than aramids, with nylon 6 more sensitive than nylon-11. At the extreme strong end sits the symmetric hydrogen bond, where the proton rests exactly halfway between two identical atoms. It is a three-center four-electron bond with an effective bond order of 0.5, comparable to a covalent bond, seen in high-pressure ice, in anhydrous acids such as hydrofluoric and formic acid under pressure, and in the bifluoride ion. The hydrogen bond's reach even extends into medicine: Lipinski's rule of five holds that most orally active drugs have no more than five hydrogen bond donors and fewer than ten acceptors, though many useful drugs break the rule.