Introduction
By now you know that molecules can attract each other through intermolecular forces — but not all of these forces are equally strong. The hydrogen bond is a special, relatively strong intermolecular force that explains many unusual properties of substances like water, ammonia, and hydrogen fluoride. Today we'll see exactly what a hydrogen bond is, when it forms, and why it matters.
What is a hydrogen bond?
A hydrogen bond is an electrostatic attraction between:
- a hydrogen atom that is covalently bonded directly to a highly electronegative atom (N, O, or F), and
- a lone pair of electrons on the N, O, or F atom of a neighbouring molecule (Talbot, p. 171).
This only happens between molecules (or between parts of the same large molecule), not within a single covalent bond — it's an intermolecular force, not a new type of bond between atoms in a molecule.
Why only N, O, and F?
Nitrogen, oxygen, and fluorine share two key features:
- Small atomic radius — this lets the hydrogen atom get very close to the lone pair.
- High electronegativity — this creates a large dipole, giving the hydrogen atom a strong partial positive charge (δ+) and the N/O/F atom a strong partial negative charge (δ−) (Talbot, p. 171).
The combination of a small, highly charged hydrogen atom approaching a close, exposed lone pair produces an unusually strong attraction compared with typical dipole-dipole forces.
Worked Example 1: Identifying hydrogen bonding
Question: Which of these substances can form hydrogen bonds between their molecules? CH₄, NH₃, CH₃OH, CO₂
Answer:
- CH₄ — No. Hydrogen is bonded only to carbon (not N, O, F).
- NH₃ — Yes. H is bonded directly to N, which has a lone pair.
- CH₃OH — Yes. The O–H hydrogen can bond to a lone pair on another molecule's oxygen.
- CO₂ — No. There is no H atom at all.
Worked Example 2: Comparing strength
Question: Explain why water has stronger overall hydrogen bonding than hydrogen fluoride, even though F is more electronegative than O.
Answer: Although the H–F bond is more polar, a water molecule has two lone pairs on its oxygen atom and two O–H hydrogens, so each H₂O molecule can form up to four hydrogen bonds with neighbours. HF only has one lone pair available per hydrogen in a simple chain arrangement, giving fewer hydrogen bonds per molecule. More hydrogen bonds overall means a greater collective attractive force in water (Talbot, p. 171).
Worked Example 3: Hydrogen bonding in ammonia
Question: Why is hydrogen bonding in NH₃ weaker than in HF?
Answer: Fluorine is more electronegative than nitrogen, so the N–H bond is less polar than the H–F bond. This means the δ+ charge on hydrogen in NH₃ is smaller, producing a weaker electrostatic attraction to a neighbouring lone pair (Talbot, p. 171).
Why hydrogen bonds matter
Hydrogen bonding explains many real-world observations, such as:
- Water's unusually high boiling point compared to other similar-sized molecules (e.g., H₂S).
- Ice floating on liquid water, because hydrogen bonds hold molecules in a more open, less dense structure.
- The shape and stability of DNA and proteins, which rely on hydrogen bonds between bases and chains.
Key takeaways
- A hydrogen bond forms when H, bonded directly to N, O, or F, is attracted to a lone pair on N, O, or F in a neighbouring molecule.
- It is the strongest type of intermolecular force, but still much weaker than a covalent or ionic bond.
- Strength depends on both electronegativity (bond polarity) and the number of lone pairs/hydrogens available per molecule.
- Water has stronger overall hydrogen bonding than HF because it can form more hydrogen bonds per molecule.
- Hydrogen bonding explains key properties of water, ice, and biological molecules like DNA.