Intermolecular Forces and Properties
Learn and apply Intermolecular Forces and Properties in the published Chemistry course sequence.
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The core idea
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Intermolecular Forces and Properties: Orientation
Intermolecular force questions are property questions: you must name the dominant force (London / dipole–dipole / hydrogen bonding) and then justify its strength using electrons, surface area, and polarity.
Treat this as an extension of Atomic Structure (A Level), then use the Chemical Bonding hub to compare models across the topic.
Definitions (Must Know)
A. London (induced dipole–induced dipole)
London forces are attractions caused by instantaneous dipoles inducing dipoles in neighbouring molecules.
B. Permanent dipole–dipole forces
Permanent dipole–dipole forces are attractions between the δ⁺ and δ⁻ ends of polar molecules.
C. Hydrogen bonding
Hydrogen bonding is the electrostatic attraction between a δ⁺ hydrogen atom in N-H, O-H, or F-H and a lone pair on N, O, or F in a neighbouring molecule.
Detailed Explanations
A. What each force actually is (beginner explanation)
- London forces: electrons move randomly, so a molecule can have a temporary (instantaneous) dipole; this induces a dipole in a neighbour; the opposite ends attract.
- Permanent dipole–dipole: polar molecules have permanent δ⁺ and δ⁻ ends, so molecules align and attract.
- Hydrogen bonding: a strong dipole–dipole attraction when H is bonded to N/O/F, and the δ⁺ H is attracted to a lone pair on N/O/F in a neighbour.
B. Choosing the dominant intermolecular force (workflow)
- Does it have N-H, O-H, or F-H? → hydrogen bonding matters.
- If not, is the molecule polar? → permanent dipole–dipole matters.
- Always compare London forces using electrons and surface area (branching).
Mini example:
- CH₃CH₂OH (ethanol) has O-H → hydrogen bonding.
- CH₃OCH₃ (dimethyl ether) has no O-H → no hydrogen bonding (but it is polar, so dipole–dipole + London).
Required reference examples:
- CHCl₃(l) is polar, so permanent dipole–dipole attractions act as well as London forces.
- Br₂(l) is non-polar, so its intermolecular attraction is induced dipole–induced dipole (London) forces.
- Liquid noble gases consist of atoms rather than molecules, but instantaneous dipoles still induce dipoles in neighbouring atoms and allow condensation.
Data table
| Hydrogen halide | Boiling point |
|---|---|
| HF | 20 |
| HCl | -85 |
| HBr | -67 |
| HI | -35 |
C. Why Mᵣ and branching matter for London forces
- More electrons / higher Mᵣ → electron cloud is more polarisable → stronger London forces.
- Less branching → larger surface contact between molecules → stronger London forces.
Data table
| Isomer | Boiling point |
|---|---|
| n-pentane | 36 |
| isopentane | 28 |
| neopentane | 10 |
D. What happens on boiling
- Boiling separates molecules, so it overcomes intermolecular forces, not covalent bonds inside molecules.
E. Why hydrogen bonding matters in water and ice
- Water has an unusually high melting point and boiling point for its small molecular size because substantial energy is needed to overcome its hydrogen-bond network.
- In liquid water, hydrogen bonds continually break and reform, allowing molecules to pack more closely than in ice.
- In ice, hydrogen bonds hold water molecules in an open lattice. The molecules are further apart on average, so ice is less dense than liquid water and floats.
For “why ice floats”, do not write only “hydrogen bonding”. State that hydrogen bonds maintain an open lattice, placing water molecules further apart and making ice less dense than liquid water.
Worked Examples
Modelled example 1
Compare Water and Hydrogen Sulfide
Problem
Explain why H₂O has a much higher boiling point than H₂S.
Study the worked solution
Identify water's strongest attraction
Method
State that water molecules form hydrogen bonds.Reason
Hydrogen is directly bonded to oxygen, and oxygen has lone pairs.Working
O-H\bond{...}O hydrogen bonding between water molecules.Contrast hydrogen sulfide
Method
State that H₂S does not form significant hydrogen bonds.Reason
Its intermolecular attractions are weaker than the hydrogen-bond network in water.Working
H₂S: weaker intermolecular forces.Link force to boiling point
Method
Conclude that more energy is required to separate water molecules.Reason
Boiling overcomes intermolecular attractions, not the covalent O–H bonds.Working
bp(H₂O)≫ bp(H₂S).
Guided practice 2
Explain Why Ice Is Less Dense
Problem
Explain why ice is less dense than liquid water.
Try this before viewing the solution
Hints
Hint 1: describe solid water
Hint 2: connect volume to density
View solution step by step
Describe the ice structure
Method
State that hydrogen bonds maintain an open lattice.Reason
The ordered hydrogen-bond arrangement leaves more empty space than the liquid structure.Working
Ice: open hydrogen-bonded lattice.Compare molecular separation
Method
State that water molecules are farther apart on average in ice.Reason
Some of the open lattice collapses when ice melts.Working
Same number of molecules occupies more volume in ice.Deduce density
Method
Use ρ = m/V for the same mass.Reason
A larger volume at constant mass gives a lower density.Working
ρ_ice < ρ_(liquid water).
Common misconception 3
Correct an Isomer Boiling-Point Claim
Learner claim
Both (CH₃)₄C and CH₃CH₂CH₂CH₂CH₃ have formula C₅H₁₂. A learner says they must have the same boiling point because they have the same electron count. Explain why (CH₃)₄C actually has the lower boiling point.
Identify the structural factor
View solution step by step
Hold electron count constant
Method
State that both non-polar isomers have the same formula and electron count.Reason
London forces dominate for both, so the difference is not a new force type.Working
Both substances: London forces.Compare molecular contact
Method
Identify (CH₃)₄C as more branched and compact.Reason
Its smaller contact surface produces weaker total London attractions between neighbouring molecules.Working
More branching → less surface contact.Deduce boiling point
Method
State that less energy is needed to separate the branched molecules.Reason
Their London attractions are weaker.Working
bp((CH₃)₄C) < bp(CH₃CH₂CH₂CH₂CH₃).
Examiner practice 4
Compare Ethanol and Dimethyl Ether
Problem
Explain why ethanol, CH₃CH₂OH, has a higher boiling point than dimethyl ether, CH₃OCH₃. [3 marks]
Try this before viewing the solution
View solution step by step
Identify ethanol's force
1 markMethod
State that ethanol molecules form hydrogen bonds with one another.Reason
Ethanol contains O–H bonds and oxygen lone pairs.Working
CH₃CH₂O-H\bond{...}O(H)CH₂CH₃.Contrast the ether
1 markMethod
State that dimethyl ether cannot hydrogen-bond with itself.Reason
It has oxygen lone pairs but no O–H bond to donate; its molecules rely on dipole–dipole and London forces.Working
CH₃OCH₃: no intermolecular self hydrogen bonding.Link to boiling point
1 markMethod
Conclude that ethanol’s stronger attractions require more energy to overcome.Reason
Boiling separates molecules.Working
bp(ethanol) > bp(dimethyl ether).
Self-mark with the mark scheme
Compare your response with each mark point. Select a point only when your response contains that evidence.
Credit ethanol hydrogen bonding, the ether contrast and the energy/boiling conclusion.
Challenge 5
Transfer Intermolecular Forces to Miscibility
Problem
Explain why ethanol, CH₃CH₂OH, is miscible with water but hexane, C₆H₁₄, is not.
Try this before viewing the solution
Hints
Hint 1: test ethanol–water attraction
Hint 2: test hexane–water attraction
View solution step by step
Analyse ethanol with water
Method
Identify hydrogen bonds between ethanol and water molecules.Reason
Ethanol’s O–H group and oxygen lone pairs interact strongly with water.Working
Ethanol–water hydrogen bonding → miscible.Analyse hexane with water
Method
Identify hexane as non-polar and unable to hydrogen-bond.Reason
Hexane–water attractions are too weak to compensate for disrupting water’s hydrogen-bond network.Working
Hexane–water: weak interactions.State the contrast
Method
Conclude that ethanol mixes throughout water while hexane separates.Reason
Miscibility is favoured when new solute–solvent attractions are sufficiently strong.Working
Ethanol is miscible; hexane is not.
Mind Stretchers
Mind stretcher 1Extension
Rank the boiling points of HF, HCl, and HI and explain your order.
Show Answer
Mark scheme:
- HF has hydrogen bonding → strongest intermolecular forces → highest boiling point.
- HCl and HI do not hydrogen bond; London forces dominate.
- HI has more electrons / higher Mᵣ than HCl → stronger London forces → higher boiling point than HCl.
- Order: HF > HI > HCl.