Hydrogen Bonding: One Interaction, Many Consequences

The three conditions

A hydrogen bond forms when all of the following hold. Missing any one of them means the interaction is an ordinary dipole–dipole attraction, not a hydrogen bond, and questions are set to test exactly this distinction.

• Hydrogen is covalently bonded to a highly electronegative atom — in practice nitrogen, oxygen or fluorine.
• That atom is small, so the resulting charge density is high.
• The acceptor atom carries a lone pair available to interact with the exposed hydrogen.

Chlorine has almost the same electronegativity as nitrogen, yet HCl shows no significant hydrogen bonding. The reason is size: chlorine is much larger, its charge is spread over a bigger volume, and its lone pairs are diffuse. Electronegativity alone is not enough, and this comparison is a standard two-mark question.

Intermolecular versus intramolecular

An intermolecular hydrogen bond links two separate molecules, so it raises boiling point, melting point, viscosity and solubility in water. An intramolecular hydrogen bond forms within a single molecule, usually closing a five- or six-membered ring, and it does the opposite: the hydrogen is now occupied internally and is unavailable to bond with neighbours.

Why ice floats

In liquid water each molecule hydrogen bonds to its neighbours in a constantly rearranging network, with molecules packing reasonably close. On freezing, every water molecule forms exactly four hydrogen bonds arranged tetrahedrally, and that geometric requirement produces an open hexagonal lattice with large empty channels.

The open structure is less dense than the liquid, so ice floats. Water is one of very few substances where the solid is less dense than the liquid, and the reason is directional bonding, not weak bonding. The consequence is that ponds freeze from the top down, insulating the water beneath.

The same argument explains why water has its maximum density at about 4°C rather than at its freezing point: below that temperature the open network is already forming.

Other consequences worth citing in answers

• Anomalous boiling points in groups 15, 16 and 17. NH₃, H₂O and HF all boil far above the trend set by their heavier congeners. Plotting boiling point against period gives the classic anomaly graph, and being able to explain it is worth more than being able to draw it.
• Solubility of lower alcohols, amines and carboxylic acids. These dissolve in water because they hydrogen bond with it. Solubility falls as the hydrocarbon chain grows and the non-polar part dominates.
• Carboxylic acid dimers. In non-polar solvents and in the vapour phase, carboxylic acids form cyclic dimers held by two hydrogen bonds, which is why their measured molar mass comes out at roughly double the formula value.
• Biological structure. The double helix of DNA is held by hydrogen bonds between complementary bases, three between guanine and cytosine and two between adenine and thymine, and the alpha helix of proteins is held by hydrogen bonds along the backbone.
• The high specific heat and high surface tension of water both arise from the energy needed to disrupt the hydrogen-bonded network.

How to answer a physical-property comparison

1. Check for O–H, N–H or F–H. Without one of these there is no hydrogen bond donor, whatever the electronegativity of the other atoms.
2. Ask whether the bond can be internal. If the two groups are positioned to close a five- or six-membered ring, expect chelation and reverse the usual prediction.
3. Only then compare molar masses and dispersion forces, which decide the outcome when hydrogen bonding is absent or equal on both sides.
4. Name the specific effect in the answer — higher boiling point, lower volatility, greater solubility — rather than stating that hydrogen bonding is present and leaving the consequence implied.

Frequently asked questions

Why does HF have a higher boiling point than HCl even though HCl has a greater molar mass?

Fluorine is small and highly electronegative, so HF forms strong hydrogen bonds and extra energy is needed to separate the molecules. Chlorine is too large for effective hydrogen bonding, so HCl relies on weaker dipole and dispersion forces despite its greater mass.

Why is ortho-nitrophenol more volatile than para-nitrophenol?

The ortho isomer forms an intramolecular hydrogen bond closing a six-membered ring, which uses up the O–H and prevents association between molecules. The para isomer has no such option and forms intermolecular hydrogen bonds, so it is less volatile and higher boiling.

Can a molecule accept a hydrogen bond without donating one?

Yes. Ethers, ketones and tertiary amines have lone pairs and accept hydrogen bonds, but have no hydrogen attached to N, O or F and so cannot donate. This is why acetone dissolves in water but boils lower than an alcohol of similar mass.

Why is ice less dense than water?

Each molecule in ice forms four hydrogen bonds in a tetrahedral arrangement, and that geometry forces an open hexagonal lattice containing empty space. Liquid water has a partially broken, more compact network, so it is denser.

Is a hydrogen bond a real chemical bond?

It is an intermolecular attraction, not a covalent bond. It is far stronger than ordinary dipole–dipole or dispersion forces, typically 10 to 40 kJ per mole, but roughly a tenth of a covalent bond, so it breaks and reforms readily at ordinary temperatures.