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Intermolecular forces questions reward a clear, systematic approach: identify the forces present, compare their strengths, and connect them to the property asked about. These 15 questions cover identifying forces, ranking boiling points, explaining solubility and interpreting data. Answer each one fully in writing before checking the key — the explanations matter as much as the final answers. For background, see intermolecular forces.
Quick reference
| Force | Present when | Relative strength |
|---|---|---|
| London (dispersion) | Always, in all molecules | Grows with electron count and contact area |
| Permanent dipole–dipole | Polar molecules | Moderate |
| Hydrogen bonding | H bonded to N, O or F, plus a lone pair on N, O or F | Strongest between neutral molecules |
| Ion–dipole | An ion and a polar molecule (e.g. in solutions) | Stronger still |
Questions
Q1. State all the types of intermolecular force present between molecules of: (a) CH₄ (b) HCl (c) CH₃OH (d) CO₂ (e) CH₃Cl
Q2. Explain why the boiling points of the noble gases increase from helium to radon.
Q3. Put these in order of increasing boiling point and explain: F₂, Cl₂, Br₂, I₂.
Q4. Which has the higher boiling point, pentane (straight chain) or 2,2-dimethylpropane? Both are C₅H₁₂. Explain.
Q5. Propane (M = 44) boils at −42 °C and ethanal (M = 44) at 20 °C. Explain the difference.
Q6. Explain why water (M = 18) has a higher boiling point than hydrogen sulfide (M = 34).
Q7. The boiling points of the group 16 hydrides are: H₂O 100 °C, H₂S −60 °C, H₂Se −41 °C, H₂Te −2 °C. Explain the trend from H₂S to H₂Te, and why water doesn’t fit it.
Q8. Ethanol (C₂H₅OH) and methoxymethane (CH₃OCH₃) have the same molecular formula, C₂H₆O. Ethanol boils at 78 °C; methoxymethane at −24 °C. Explain.
Q9. Explain why ethanol dissolves in water in all proportions, but hexane doesn’t dissolve in water.
Q10. Draw a diagram (or describe in words) showing a hydrogen bond between two ammonia molecules. Include partial charges and lone pairs.
Q11. Explain why sodium chloride dissolves in water but not in hexane.
Q12. Explain why ice is less dense than liquid water.
Q13. HF boils at 20 °C, but HCl boils at −85 °C. Yet HI (−35 °C) boils higher than HCl. Explain both observations.
Q14. Iodine is a solid at room temperature, but chlorine is a gas. Both are non-polar diatomic molecules. Explain.
Q15. A student says: “When water boils, the O–H bonds break, which is why it needs so much energy.” Explain what’s wrong and give a correct statement.
Answer key
A1. (a) CH₄: London forces only (non-polar). (b) HCl: London + permanent dipole–dipole (polar, but H isn’t bonded to N, O or F). (c) CH₃OH: London + dipole–dipole + hydrogen bonding (O–H group). (d) CO₂: London only — polar C=O bonds, but the linear molecule is non-polar (see how to tell if a molecule is polar). (e) CH₃Cl: London + dipole–dipole.
A2. Noble gases have only London forces. Down the group, atoms have more electrons and are larger, so their electron clouds are more polarisable. Larger temporary and induced dipoles form, giving stronger London forces, which need more energy to overcome (see London dispersion forces).
A3. F₂ < Cl₂ < Br₂ < I₂. All are non-polar with only London forces. The number of electrons increases (18, 34, 70, 106), increasing polarisability and London forces, so boiling points rise (−188, −34, 59, 184 °C).
A4. Pentane. Both have the same number of electrons, but pentane’s straight chain has a larger surface area in contact with neighbouring molecules, giving stronger London forces. 2,2-Dimethylpropane is compact and nearly spherical, with less contact. (Boiling points: 36 °C and 10 °C.)
A5. Similar molar mass and electron numbers → similar London forces. Ethanal contains a polar C=O group, so it also has permanent dipole–dipole forces; propane is almost non-polar. More energy is needed to separate ethanal molecules (see dipole–dipole forces).
A6. Water forms hydrogen bonds between molecules (H bonded to highly electronegative O, and lone pairs on O). H₂S can’t form significant hydrogen bonds because S is much less electronegative (2.58 vs 3.44). Although H₂S has more electrons (stronger London forces), water’s hydrogen bonds are much stronger overall (see why water has such a high boiling point).
A7. From H₂S to H₂Te, the molecules have more electrons, so London forces increase and boiling points rise. Water is far higher than the trend predicts because of hydrogen bonding, which the other hydrides don’t form significantly. Extrapolating the H₂S–H₂Te trend backwards would put water’s boiling point somewhere around −80 °C.
A8. Ethanol has an O–H group, so it forms hydrogen bonds between molecules. Methoxymethane has an oxygen with lone pairs but no O–H — its hydrogens are all bonded to carbon — so it can’t hydrogen-bond with itself; it has only London and dipole–dipole forces. Same formula, very different boiling points.
A9. Ethanol forms hydrogen bonds with water (its O–H group can donate and accept), so it mixes readily, replacing water–water hydrogen bonds with ethanol–water hydrogen bonds. Hexane is non-polar and can’t hydrogen-bond; mixing it in would break water–water hydrogen bonds without replacing them, which is unfavourable (see “like dissolves like”).
A10. In molecule 1, an N–H hydrogen carries δ+. In molecule 2, the nitrogen carries δ− and has a lone pair. The hydrogen bond (drawn as a dashed line) runs from the δ+ H of molecule 1 to the lone pair on the N of molecule 2, with the N–H···N atoms roughly in a straight line.
A11. In water, the ions are attracted to polar water molecules by ion–dipole forces, forming hydration shells; the energy released helps overcome the lattice energy. Hexane is non-polar and can’t form ion–dipole attractions, so nothing compensates for breaking the lattice (see ion–dipole forces).
A12. In ice, each water molecule forms four hydrogen bonds in a regular tetrahedral arrangement, building an open hexagonal lattice with lots of empty space. When ice melts, some hydrogen bonds break and molecules can pack more closely, so liquid water is denser (see why ice floats).
A13. HF forms hydrogen bonds (H bonded to F), which are much stronger than the forces between HCl molecules, so HF boils much higher. From HCl to HI, London forces increase with the number of electrons (18 → 54), outweighing the decrease in dipole–dipole forces as polarity falls, so HI boils higher than HCl.
A14. Both have only London forces. Iodine has far more electrons (106 vs 34), so its electron cloud is much more polarisable, giving much stronger London forces — enough to hold the molecules in a solid at room temperature.
A15. Boiling water overcomes the hydrogen bonds between molecules (intermolecular forces); the covalent O–H bonds within each molecule don’t break. Steam is still H₂O molecules. Water needs a lot of energy to boil (relative to its size) because its intermolecular hydrogen bonds are unusually strong (see intramolecular vs intermolecular forces).
Scoring
- 13–15: excellent — move on to ranking boiling points using intermolecular forces.
- 9–12: good; review the questions involving competing factors (Q7, Q13).
- Below 9: revisit the explainers on each type of force.
Exam tips
- Always name all the forces present, not just the strongest.
- Compare like with like: when molecules are similar in size, polarity and hydrogen bonding decide; when sizes differ a lot, London forces may dominate.
- Link forces to energy: “stronger forces → more energy needed to overcome → higher boiling point”.
- Never say bonds break when a molecular substance boils or melts.
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