1.5.4Physical Chemistry
Intermolecular Forces and Physical Properties — Edexcel AS Chemistry Revision Notes & Topical Questions
5 min read 5 parts Lesson 20 of 42
Part 1 of 5Linking Forces to Properties
Simple molecular substances have low melting and boiling points because only weak forces between molecules must be overcome.
London forces
- Present between all molecules;
- they grow stronger as the number of electrons rises, so larger molecules have higher boiling points.
Permanent dipole forces
- Polar molecules attract each other through their partial charges, adding to London forces and raising boiling points a little.
Hydrogen bonds
- The strongest intermolecular force, found when H is bonded to N, O or F with a lone pair on a neighbouring molecule.
Energy needed
- Boiling or melting a simple molecular substance only overcomes intermolecular forces;
- the covalent bonds inside molecules stay intact.
Part 2 of 5Boiling Point Trends in Hydrides
Boiling Point Trends in Hydrides
Group 4 hydrides
- CH₄, SiH₄, GeH₄ and SnH₄ are non-polar, so boiling point rises smoothly down the group as London forces strengthen.
Group 5, 6 and 7 anomalies
- NH₃, H₂O and HF boil far higher than expected from their group trend because they form hydrogen bonds.
Why water is special
- Each water molecule can form up to four hydrogen bonds, two through its H atoms and two through its lone pairs.
Later members of the group
- H₂S, H₂Se and H₂Te follow the London force pattern and rise in boiling point down the group.
| Compound | Main forces between molecules | Boiling point trend |
|---|---|---|
| CH₄ | London only | Very low |
| HCl | London + permanent dipole | Low |
| H₂O | London + hydrogen bonds | Unusually high |
Part 3 of 5Hydrogen Bonding in Water and Ice
1
In liquid water, molecules are close together and hydrogen bonds constantly break and re-form as molecules move.
2
On cooling, molecules slow down and hydrogen bonds hold neighbouring molecules in fixed positions.
3
In ice, each molecule forms four hydrogen bonds, giving an open, hexagonal, regular lattice.
4
The open lattice contains gaps, so ice is less dense than liquid water and floats on it.
5
Solid water melts at a higher temperature than other group 6 hydrides because extra energy is needed to break these hydrogen bonds.
Part 4 of 5Solubility and Volatility
Solubility and Volatility
Polar solvents
- Water dissolves substances that can form hydrogen bonds or ion-dipole attractions, such as ethanol and sodium chloride.
Non-polar solutes
- Hydrocarbons are insoluble in water because they cannot replace the strong hydrogen bonds between water molecules.
Volatility
- Substances with weak intermolecular forces, such as methane, evaporate easily and have low boiling points.
Electrical conduction
- Simple molecular substances do not conduct because they have no mobile charged particles.
Part 5 of 5Exam Checklist & Tips
1.5.4 – Key Points for the Exam
- Intermolecular force An attraction between neighbouring molecules, much weaker than a covalent bond.
- Hydrogen bond Attraction between a δ+ hydrogen on N, O or F and a lone pair on N, O or F in another molecule.
- London force Attraction caused by temporary and induced dipoles from electron movement.
- Permanent dipole Uneven charge distribution in a polar molecule due to a difference in electronegativity.
- Volatile Easily turned into a vapour because intermolecular forces are weak.
Common Mistakes
- Saying covalent bonds break when a simple molecular substance boils; only intermolecular forces are overcome.
- Forgetting the lone pair on the acceptor atom when describing a hydrogen bond.
- Stating that hydrogen bonds are stronger than covalent bonds; they are much weaker.
- Claiming HCl has hydrogen bonds because it contains hydrogen; chlorine is not N, O or F.
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