Ranking Boiling Points Using Intermolecular Forces

“Arrange these substances in order of increasing boiling point and explain your answer” is one of the most common structure-and-bonding exam questions. It looks like it needs memorised data, but it doesn’t: with a clear method and an understanding of intermolecular forces, you can rank almost any set of substances and justify the order. This article gives that method and then works through seven examples of increasing difficulty.

Why boiling point tracks intermolecular forces

When a simple molecular liquid boils, its molecules separate from each other and escape as a gas. The covalent bonds inside the molecules don’t break; only the forces between molecules are overcome (see intramolecular vs intermolecular forces). So:

stronger intermolecular forces → more energy needed to separate molecules → higher boiling point

For substances that aren’t simple molecular — ionic compounds, metals and giant covalent structures — boiling requires overcoming much stronger forces or breaking bonds, so their boiling points are usually far higher (see simple molecular vs giant covalent).

The method

Step 1. Sort by structure type. Giant covalent, ionic and metallic substances have much higher boiling points than simple molecular ones. Put any of these at the top of the list first.

Step 2. For simple molecular substances, identify the forces in each.

• All molecules: London forces (see London dispersion forces).
• Polar molecules: + permanent dipole–dipole (see dipole–dipole forces).
• H bonded to N, O or F: + hydrogen bonding (see hydrogen bonding).

Step 3. Compare, using these rules of thumb:

• If molecules are similar in size, the one with hydrogen bonding boils highest, then polar molecules, then non-polar.
• If molecules have the same type of forces, the one with more electrons (and more surface contact) boils higher.
• If sizes are very different, a much larger molecule with only London forces can boil higher than a small hydrogen-bonded one.
• More hydrogen-bonding groups per molecule → higher boiling point.

Step 4. Write the explanation: name the forces, say which are stronger and why, and link to the energy needed.

Example 1: a homologous series

Rank: butane (C₄H₁₀), methane (CH₄), propane (C₃H₈), ethane (C₂H₆).

All are non-polar alkanes with London forces only. More carbons → more electrons and larger surface area → stronger London forces.

Order: methane (−162 °C) < ethane (−89 °C) < propane (−42 °C) < butane (−1 °C).

Example 2: isomers

Rank: pentane, 2-methylbutane, 2,2-dimethylpropane (all C₅H₁₂).

Same formula, same number of electrons, all non-polar. The difference is shape. More branching makes molecules more compact and spherical, reducing the surface area in contact, so London forces are weaker.

Order: 2,2-dimethylpropane (10 °C) < 2-methylbutane (28 °C) < pentane (36 °C).

Example 3: similar size, different forces

Rank: propan-1-ol (CH₃CH₂CH₂OH, M = 60), butane (C₄H₁₀, M = 58), propanone (CH₃COCH₃, M = 58).

All have similar sizes and electron counts (32–34 electrons), so London forces are comparable. The deciding factor is the other forces:

• Butane: London only.
• Propanone: London + dipole–dipole (polar C=O).
• Propan-1-ol: London + dipole–dipole + hydrogen bonding (O–H).

Order: butane (−1 °C) < propanone (56 °C) < propan-1-ol (97 °C).

Example 4: alcohols vs ethers (isomers again)

Rank: ethanol (C₂H₅OH) and methoxymethane (CH₃OCH₃), both C₂H₆O.

• Ethanol has an O–H group → hydrogen bonding between molecules.
• Methoxymethane has an oxygen but no O–H — all its hydrogens are on carbon — so it can’t hydrogen-bond with itself; it has dipole–dipole and London forces only.

Order: methoxymethane (−24 °C) < ethanol (78 °C). A difference of over 100 °C from the same atoms arranged differently.

Example 5: when size beats hydrogen bonding

Rank: water (H₂O, M = 18), octane (C₈H₁₈, M = 114), ethanol (C₂H₅OH, M = 46).

• Water: small, but extensive hydrogen bonding (up to four per molecule).
• Ethanol: hydrogen bonding (one O–H), plus a small hydrocarbon part.
• Octane: non-polar, London only, but a large molecule with 66 electrons and a long chain.

Order: ethanol (78 °C) < water (100 °C) < octane (126 °C).

Octane beats both hydrogen-bonded liquids because its much larger size gives very strong London forces. Water beats ethanol because each water molecule can form more hydrogen bonds (two H atoms and two lone pairs) than each ethanol molecule (one H–O).

Example 6: hydrides of groups 14–17

Rank within each group: CH₄, SiH₄, GeH₄ (group 14) and NH₃, PH₃, AsH₃ (group 15).

Group 14: all non-polar tetrahedral molecules with London forces only. Boiling point rises down the group as electrons increase: CH₄ (−162 °C) < SiH₄ (−112 °C) < GeH₄ (−88 °C).

Group 15: PH₃ (−88 °C) < AsH₃ (−62 °C), following the London-force trend, but NH₃ (−33 °C) is anomalously high, because it forms hydrogen bonds.

The same anomaly appears in group 16 (H₂O far above H₂S) and group 17 (HF far above HCl). In group 14, there’s no anomaly — CH₄ can’t hydrogen-bond. This pattern is one of the key pieces of evidence for hydrogen bonding (see why water has such a high boiling point).

Example 7: mixed structure types

Rank: sodium chloride, water, diamond, oxygen, copper.

Step 1 — sort by structure:

• O₂: simple molecular, non-polar → very low.
• H₂O: simple molecular, hydrogen bonding → low.
• NaCl: giant ionic → high.
• Cu: giant metallic → high.
• Diamond: giant covalent → extremely high (sublimes/converts rather than boils at normal pressure).

Order (boiling points): O₂ (−183 °C) < H₂O (100 °C) < NaCl (1,413 °C) < Cu (about 2,560 °C) < diamond (above about 3,600 °C, sublimes).

The exact order among the giant structures needs data, but the principle is clear: strong bonds throughout the structure put them far above any small-molecule substance.

Writing the explanation: a model answer

Question: Explain why propan-1-ol has a higher boiling point than butane, even though their molar masses are almost the same.

Answer: Both molecules have London forces of similar strength because they have similar numbers of electrons. Propan-1-ol also has an O–H group, so its molecules form hydrogen bonds with each other, in addition to London forces and dipole–dipole forces. Butane is non-polar and has London forces only. Hydrogen bonds are much stronger than London forces, so more energy is needed to separate propan-1-ol molecules, giving it a higher boiling point.

Common mistakes

• Saying bonds break when a simple molecular substance boils.
• Ignoring London forces in polar molecules — every molecule has them.
• Assuming hydrogen bonding always wins — very large molecules can boil higher (octane vs water).
• Using molar mass as the explanation — it’s the number of electrons and surface area that matter, not mass itself.
• Counting any hydrogen as able to hydrogen-bond — only H bonded to N, O or F.

Key takeaways

• Boiling point rises with the strength of intermolecular forces (for simple molecular substances).
• Sort by structure type first; giant structures boil far higher.
• For similar-sized molecules: hydrogen bonding > dipole–dipole > London only.
• For the same type of forces: more electrons and more surface contact → higher boiling point.
• Always name the forces, compare their strength and link to energy in your explanation.

Test yourself with the intermolecular forces practice questions.