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
Even though H₂S is heavier, H₂O has the much higher boiling point because of hydrogen bonding.
Key IB Chemistry Examples
1. Water
Water’s unusually high boiling point is due to:
- Two O–H bonds
- Two lone pairs on oxygen
- Ability to form up to four hydrogen bonds per molecule
This extensive hydrogen bonding network requires large amounts of energy to break.
2. Alcohols
Alcohols contain an –OH group capable of hydrogen bonding.
As the hydrocarbon chain length increases, boiling point rises because:
- Hydrogen bonding occurs at the –OH
- Dispersion forces increase with chain length
3. Carboxylic Acids
Carboxylic acids form dimers (pairs) via two hydrogen bonds between molecules, giving them even higher boiling points.
4. Amines
Amines can hydrogen bond through N–H groups, though weaker than O–H because nitrogen is less electronegative.
Understanding these patterns helps answer data trend questions in exams.
Comparison With Other Intermolecular Forces
Hydrogen bonds vs. Dipole–Dipole Forces
- Hydrogen bonds are stronger
- Produce much higher boiling points
- More structured intermolecular networks
Hydrogen bonds vs. London Dispersion Forces
- Hydrogen bonding dominates unless the molecule is extremely large
- Explains why small molecules like H₂O and HF boil higher than much larger molecules without H-bonding
Hydrogen bonds vs. Ionic Forces
- Ionic forces are much stronger
- Ionic compounds have very high melting/boiling points because of lattice structures
Hydrogen bonding sits between these on the strength scale.
How Hydrogen Bonds Influence Trends in Groups
Group 16 Hydrides (O, S, Se, Te)
- H₂O is the highest boiling point
- Others follow expected mass trends
Hydrogen bonding disrupts the normal “increasing boiling point with mass” pattern.
This is a classic IB exam trend question.
FAQs
Why does hydrogen bonding require N, O, or F?
These atoms are highly electronegative and small, allowing strong partial charges and efficient overlap with hydrogen’s exposed nucleus.
