Worked example 1
Which of CH₃OCH₃ and CH₃CH₂OH hydrogen bonds to itself, and which boils higher?
Try it first: Look for an O-H bond in each structure.
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Types of Solids and IMFSIntro to IMFSContent level: Core 20 min
Hydrogen Bonding
What you'll be able to do: Recognise when hydrogen bonding is possible and use it to explain anomalous physical properties.
Best after: Dipole-Dipole and Ion-Dipole Forces
Introduction
Water should boil around -80 °C based on its size. It boils at 100 instead, and hydrogen bonding is the reason.
These are recommended, not required. You can start this lesson at any time.
Learning objectives
- State the conditions required for hydrogen bonding
- Distinguish hydrogen bond donors from acceptors
- Explain anomalous boiling points in groups 15 to 17
- Connect hydrogen bonding to the density of ice and to solubility
Lesson
The N, O, F rule
Hydrogen bonding needs a hydrogen atom covalently bonded to nitrogen, oxygen or fluorine, and a lone pair on a nearby N, O or F. Those three elements are small and highly electronegative, so the hydrogen is left almost bare and can approach a lone pair very closely.
A molecule with hydrogen and oxygen in it does not qualify unless H is bonded directly to the O.
Donors and acceptors
The molecule supplying the H-N, H-O or H-F hydrogen is the donor. The molecule supplying the lone pair is the acceptor. Acetone has no O-H bond, so it can only accept, which is why it mixes with water but does not hydrogen bond to itself.
The boiling point anomaly
Across group 16 the hydrides H₂S, H₂Se and H₂Te rise steadily with size, but H₂O sits far above the extrapolated line. NH₃ and HF show the same jump in their groups. Hydrogen bonding is the only difference.
Hydrogen bonds are around 5 to 30 kJ/mol, far weaker than the roughly 460 kJ/mol O-H covalent bond.
Ice, water and life
Each water molecule can form up to four hydrogen bonds. In ice those bonds lock molecules into an open hexagonal lattice that is less dense than liquid water, so ice floats. Hydrogen bonding also explains water high surface tension and its ability to dissolve alcohols and sugars.
Key ideas
Definition
Hydrogen bond
An attraction between an H bonded to N, O or F and a lone pair on another N, O or F.
Rule
Donor requirement
No H-N, H-O or H-F bond means the molecule cannot donate a hydrogen bond.
Rule
Strength band
Roughly 5 to 30 kJ/mol, stronger than ordinary dipole-dipole but far weaker than a covalent bond.
Key concept
Water anomaly
H₂O boils at 100 °C while the heavier H₂S boils at -60.
Worked examples
Worked example 2
Explain why HF boils at 20 °C while HCl boils at -85 despite chlorine being heavier.
Try it first: Check which hydride can hydrogen bond.
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Common mistakes
Claiming a hydrogen bond is a covalent bond.
Why it's wrong: It is an intermolecular attraction roughly twenty times weaker than an O-H bond.
Check instead: Say hydrogen bonds break on boiling while covalent bonds stay intact.
Allowing hydrogen bonding with Cl or S because they are electronegative.
Why it's wrong: They are too large and insufficiently electronegative to produce the effect.
Check instead: Restrict the rule to N, O and F.
Assuming any molecule containing O can hydrogen bond to itself.
Why it's wrong: Ethers and ketones lack an O-H bond and can only accept.
Check instead: Look for H bonded directly to N, O or F before answering.
Practice this skill
No practice questions are available for this topic yet. You can still practice the whole unit.
What you should now know
Hydrogen bonding is an unusually strong dipole-dipole attraction that occurs when hydrogen is bonded directly to N, O or F and interacts with a lone pair on another N, O or F. It explains the high boiling points of water, ammonia and hydrogen fluoride, and the fact that ice floats.
- Hydrogen bonding requires H bonded to N, O or F plus a lone pair acceptor
- Donors need an H-N, H-O or H-F bond; acceptors only need a lone pair
- It explains the anomalous boiling points of H₂O, NH₃ and HF
- The open hydrogen bonded lattice makes ice less dense than water
Sources and further reading
- Chemistry 2e, Section 10.1: Intermolecular ForcesOpenStax · Rice University · 10.1View source
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