Hydrogen Bonding

Hydrogen bonding is a Supplement topic in the Cambridge 0620 syllabus. It explains why water has unexpectedly high melting and boiling points, why ice floats, and why many biological molecules have specific shapes. The concept builds on electronegativity and intermolecular forces.

What is a hydrogen bond?

A hydrogen bond is a strong intermolecular force that occurs when:

1. A hydrogen atom is covalently bonded to a highly electronegative atom (oxygen, nitrogen, or fluorine)
2. The highly electronegative atom pulls the shared electrons towards itself, leaving the hydrogen with a strong partial positive charge (delta+)
3. This delta+ hydrogen is attracted to a lone pair of electrons on an electronegative atom in a neighbouring molecule

The hydrogen bond forms between the delta+ hydrogen and the lone pair:

O-H…O (the dotted line represents the hydrogen bond)

Hydrogen bonds are much stronger than other intermolecular forces (van der Waals forces) but much weaker than covalent bonds or ionic bonds.

Why only O, N, and F?

These three elements have both:

• Very high electronegativity — they strongly pull the bonding electrons away from hydrogen
• Small atomic radius — the lone pairs are close to the nucleus and accessible
• Lone pairs available — to attract the delta+ hydrogen of a neighbouring molecule

Chlorine and sulfur are not electronegative enough (and are too large) to form effective hydrogen bonds.

Hydrogen bonding in water

Each water molecule can form up to four hydrogen bonds:

• Two through its two O-H bonds (hydrogen acts as the delta+ donor)
• Two through the two lone pairs on oxygen (accepting hydrogen bonds from neighbouring molecules)

This extensive hydrogen bonding network explains water’s unusual properties.

Properties of water explained by hydrogen bonding

High boiling point (100 °C)

Without hydrogen bonding, water (Mr = 18) would be expected to boil at about -80 °C, similar to H₂S. The actual boiling point of 100 °C is remarkably high because a large amount of energy is needed to break the extensive hydrogen bonding network.

High specific heat capacity

Water absorbs a large amount of heat before its temperature rises significantly. This makes it an excellent coolant (in car engines, industrial processes) and gives coastal areas in Malaysia a more moderate climate than inland regions.

Ice is less dense than liquid water

When water freezes, hydrogen bonds hold the molecules in an open, regular hexagonal arrangement. This structure has gaps, making ice less dense than liquid water. This is why ice floats — a property critical for aquatic life, as the floating ice insulates the water beneath.

High surface tension

Hydrogen bonds at the surface of water create a “skin” effect. Small insects can walk on water because of this high surface tension.

Hydrogen bonding in other molecules

Molecule	Formula	Hydrogen bonding?	Consequence
Ammonia	NH₃	Yes (N-H…N)	Higher boiling point than expected (−33 °C vs ~−80 °C expected)
Hydrogen fluoride	HF	Yes (F-H…F)	Higher boiling point than HCl
Ethanol	C₂H₅OH	Yes (O-H…O)	Higher boiling point than ethane (similar Mr)
Methane	CH₄	No (C is not electronegative enough)	Low boiling point (−161 °C)
Hydrogen sulfide	H₂S	No (S is not electronegative enough)	Lower boiling point than water despite higher Mr

Comparing intermolecular forces

Force type	Strength	Occurs in	Example
Hydrogen bond	Strongest intermolecular force	O-H, N-H, F-H groups	Water, ethanol, ammonia
Permanent dipole-dipole	Moderate	Polar molecules	HCl
Van der Waals (London)	Weakest	All molecules	CH₄, I₂

All three are much weaker than covalent or ionic bonds.

Worked exam question

Explain why water (H₂O, Mr = 18) has a much higher boiling point than hydrogen sulfide (H₂S, Mr = 34), even though H₂S has a larger relative molecular mass. [3]

Water molecules form hydrogen bonds between the O-H groups on one molecule and the lone pairs on oxygen of neighbouring molecules [1]. Oxygen is highly electronegative, so the O-H bond is very polar, making these hydrogen bonds strong [1]. H₂S cannot form hydrogen bonds because sulfur is not electronegative enough, so its intermolecular forces are weaker and less energy is needed to separate the molecules [1].

Common exam mistakes

1. Calling hydrogen bonds “covalent bonds” — hydrogen bonds are intermolecular forces, not covalent bonds. They are much weaker.
2. Saying hydrogen bonding occurs in all molecules containing hydrogen — it only occurs when H is bonded to O, N, or F.
3. Confusing hydrogen bonds with the H-H bond in H₂ — the H₂ molecule has a covalent bond, not a hydrogen bond.
4. Not mentioning lone pairs — hydrogen bonds form between a delta+ hydrogen and a lone pair on an electronegative atom. Both parts are needed for the explanation.