Hydrogen bonding is one of the strongest types of intermolecular forces, and it plays a major role in determining physical properties such as boiling point, melting point, solubility, and viscosity. In IB Chemistry, understanding how hydrogen bonds affect boiling point helps explain trends in water, alcohols, amines, and biological molecules. This article explains the connection clearly and gives you the reasoning needed for exam questions.
What Is Hydrogen Bonding?
A hydrogen bond is an intermolecular attraction between a hydrogen atom covalently bonded to a highly electronegative element (N, O, or F) and a lone pair on another electronegative atom.
To form hydrogen bonds, three features must be present:
- Hydrogen bonded to N, O, or F (very electronegative atoms)
- Large difference in electronegativity
- Lone pairs available on adjacent molecules
Because hydrogen has only one electron, when it bonds to a highly electronegative atom, its nucleus becomes unusually exposed. This allows strong attraction to a lone pair on another molecule.
Hydrogen bonding is much stronger than dipole–dipole interactions but weaker than covalent bonds.
Why Hydrogen Bonds Increase Boiling Point
Boiling requires molecules to separate from each other and enter the gas phase.
To achieve this separation, intermolecular forces must be overcome.
Hydrogen bonds are significantly stronger than other forces such as:
- London dispersion forces
- Dipole–dipole interactions
Because hydrogen bonds require more energy to break, substances with hydrogen bonding have higher boiling points than those without.
For example:
- Water (H₂O) boils at 100°C
- Hydrogen sulfide (H₂S), which cannot form hydrogen bonds, boils at –60°C
