Bonding and intermolecular forces: TCE Chemistry (Tasmania)

Syllabus dot point

“Relate bonding type and intermolecular forces to melting point, boiling point, solubility and conductivity.”

10 min readUpdated 2 Jun 2026

Quick answer

Physical properties depend on the type of bonding and, for molecular substances, on the strength of the intermolecular forces. Dispersion forces, dipole-dipole forces and hydrogen bonding increase boiling point in that order, while ionic and covalent network solids have very high melting points because strong bonds extend throughout the lattice.

What this dot point is asking

TASC expects you to explain melting point, boiling point, solubility and electrical conductivity in terms of the bonding type and, for molecular substances, the specific intermolecular forces present.

Types of bonding

Ionic bonding is the electrostatic attraction between oppositely charged ions in a three-dimensional lattice. It gives high melting points, hardness with brittleness (the lattice shatters when like charges are forced together), and electrical conductivity only when molten or dissolved.
Covalent bonding shares electron pairs between atoms. Covalent molecular substances have strong bonds within molecules but weak forces between them, so they melt and boil at low temperatures. Covalent network solids such as diamond and silicon dioxide have a continuous lattice of strong covalent bonds and very high melting points.
Metallic bonding is the attraction between a lattice of positive ions and a sea of delocalised electrons, giving good electrical and thermal conductivity, malleability and ductility, and a range of melting points.

Intermolecular forces

For molecular substances, the forces between molecules, not the covalent bonds within them, are broken on melting and boiling:

Dispersion (London) forces act between all molecules and arise from temporary dipoles caused by the constant movement of electrons. They strengthen with more electrons, so larger molecules have higher boiling points.
Dipole-dipole forces act between polar molecules, where a permanent partial charge separation attracts neighbouring molecules.
Hydrogen bonding is an especially strong dipole interaction occurring when hydrogen is bonded to nitrogen, oxygen or fluorine and is attracted to a lone pair on the N, O or F of another molecule.

Order of strength

For molecules of similar size, boiling point rises in the order dispersion forces, then dipole-dipole forces, then hydrogen bonding. Hydrogen bonding explains why water and alcohols boil far higher than their size alone would suggest, and why $\text{HF}$ , $\text{H}_2\text{O}$ and $\text{NH}_3$ have anomalously high boiling points compared with the other hydrides in their groups.

Explaining the properties

Melting and boiling points reflect the energy needed to overcome the forces; stronger forces mean higher values. For molecular substances, compare the intermolecular forces; for ionic, network covalent and metallic substances, the lattice itself must be broken, which needs much more energy.

Solubility follows the rule like dissolves like. Polar and hydrogen-bonding substances dissolve in polar solvents such as water because favourable solute-solvent attractions replace the solute-solute and solvent-solvent attractions; non-polar substances dissolve in non-polar solvents. This is why short-chain alcohols mix with water while long-chain hydrocarbons do not.

Electrical conductivity requires mobile charge carriers. Metals conduct through delocalised electrons in any state; ionic compounds conduct only when molten or dissolved, because then the ions are free to move; molecular substances generally do not conduct because they have no free charges.

Comparing two boiling points

Identify the shared force

Explain why ethanol ( $\text{CH}_3\text{CH}_2\text{OH}$ , boiling point $78\ ^\circ\text{C}$ ) boils far above ethane ( $\text{CH}_3\text{CH}_3$ , boiling point $-89\ ^\circ\text{C}$ ), despite similar sizes. Both have dispersion forces of similar strength because their electron counts are comparable.

Identify the extra force in ethanol

Ethanol has an $\text{O}-\text{H}$ group, so its molecules form hydrogen bonds with one another; ethane is non-polar and has only dispersion forces.

Link to the property

Hydrogen bonds are much stronger than dispersion forces, so far more energy is needed to separate ethanol molecules, giving the much higher boiling point.

Saying covalent bonds break on boiling a molecular substance

When a molecular substance like water boils, the covalent $\text{O}-\text{H}$ bonds stay intact; only the intermolecular forces between molecules are overcome. Writing that boiling breaks covalent bonds is a serious error that confuses bond types. A second slip is forgetting that hydrogen bonding requires H bonded directly to N, O or F, not merely a polar molecule.

In the exam, identify the bonding type first, then for molecular substances name the specific intermolecular force, and link its strength directly to the property you are explaining.

Exam-style questions

Questions in the style of TASC exam questions on this dot point, each with a worked answer. They are written by ExamExplained unless tagged "Past paper"; the year shows the paper a question is modelled on.

Practice2 marks

Ethane has a boiling point of

-89\ ^\circ\text{C}

and ethanol has a boiling point of

78\ ^\circ\text{C}

. Explain why they have such different boiling points despite similar molar masses.

Show worked answer →

Both have similar molar masses and sizes, so the difference is due to the type of intermolecular forces. (1 mark)

Ethane, $\text{C}_2\text{H}_6$ , is non-polar, so the only forces between its molecules are weak dispersion (London) forces, needing little energy to overcome, hence its very low boiling point. Ethanol, $\text{C}_2\text{H}_5\text{OH}$ , has a polar hydroxyl ( $-\text{OH}$ ) group, so its molecules form hydrogen bonds in addition to dispersion forces. Hydrogen bonding is much stronger, so much more energy is needed to separate ethanol molecules, giving the far higher boiling point. (1 mark)

Practice3 marks

An oxygen molecule is smaller than a nitrogen molecule even though

\text{O}_2

has the higher molecular mass. Explain why oxygen molecules are smaller.

Show worked answer →

Molecular size is set by bond length and atomic radius, not by mass. (1 mark)

A nitrogen molecule has a triple bond ( $\text{N}\equiv\text{N}$ ) and oxygen a double bond ( $\text{O}=\text{O}$ ); more shared pairs pull nuclei closer, but the decisive factor is atomic radius across the period. (1 mark)

Oxygen lies to the right of nitrogen in Period 2, so an oxygen atom has one more proton while electrons are added to the same shell. The greater nuclear charge pulls the electron cloud in more tightly, so the oxygen atom and the $\text{O}_2$ molecule have a smaller radius than nitrogen, despite the higher mass. (1 mark)

Practise this

Sources & how we know this

Chemistry Level 4 (CHM415115) course document (2015 onwards) — TASC