How does structure and bonding explain the physical properties of substances?
Relate bonding type and intermolecular forces to melting point, boiling point, solubility and conductivity.
Ionic, covalent and metallic bonding, dispersion forces, dipole-dipole forces and hydrogen bonding, and how they explain melting point, solubility and conductivity.
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What this dot point is asking
You must explain physical properties in terms of bonding and intermolecular forces.
Types of bonding
Ionic bonding is the electrostatic attraction between oppositely charged ions in a lattice, giving high melting points, brittleness, and 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 have a continuous lattice of strong 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 conductivity, malleability 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. 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 N, O or F of another molecule.
Explaining properties
Melting and boiling points reflect the energy needed to overcome the forces. Stronger forces mean higher melting and boiling points.
Solubility follows like dissolves like. Polar and hydrogen-bonding substances dissolve in polar solvents like water; non-polar substances dissolve in non-polar solvents.
Electrical conductivity needs mobile charge carriers. Metals conduct via delocalised electrons; ionic compounds conduct only when molten or dissolved because the ions become free to move; molecular substances generally do not conduct.
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 practice questions
Practice questions written in the style of TASC exam questions on this dot point, with worked answer explainers. The year tag is the paper they imitate, not the source.
2022 TASC2 marksEthane has a boiling point of -89 degrees C and ethanol has a boiling point of 78 degrees C. Explain why ethane and ethanol have such different boiling points.Show worked answer →
Both have similar molar masses and sizes, so the difference is due to the type of intermolecular forces.
Ethane (C2H6) is a non-polar molecule, so the only forces between its molecules are weak dispersion (London) forces, which need little energy to overcome - hence its very low boiling point.
Ethanol (C2H5OH) has a polar hydroxyl (-OH) group, so its molecules can form hydrogen bonds with one another (in addition to dispersion forces). Hydrogen bonding is much stronger than dispersion forces alone, so much more energy is needed to separate the molecules, giving ethanol a far higher boiling point. (2 marks: identify hydrogen bonding in ethanol versus dispersion only in ethane, and link to boiling point.)
2021 TASC3 marksOne reason for using nitrogen rather than air to inflate a tyre is that oxygen molecules are smaller than nitrogen molecules. Explain why oxygen molecules are smaller than nitrogen molecules even though oxygen molecules have a higher molecular mass.Show worked answer →
Molecular size is determined by bond length and atomic radius, not by mass.
A nitrogen molecule has a triple bond (N(triple)N), whereas an oxygen molecule has a double bond (O=O). A triple bond involves more shared electron pairs, which pulls the two nuclei closer together; even so, the key factor is atomic radius across the period.
Oxygen is to the right of nitrogen in Period 2, so an oxygen atom has one more proton in its nucleus while its electrons are added to the same shell. The greater nuclear charge pulls the electron cloud in more tightly, so the oxygen atom (and hence the O2 molecule) has a smaller radius than nitrogen, despite oxygen's higher mass. (3 marks: mass does not set size; greater nuclear charge across the period contracts the atom.)