The types of bonding and structure (covalent molecular, covalent network, ionic, metallic), the intermolecular forces including London dispersion forces, permanent dipole-permanent dipole interactions and hydrogen bonding, and how these explain physical properties.

What this key area is asking

The SQA wants you to identify the type of bonding and structure in a substance, describe the intermolecular forces (London dispersion forces, permanent dipole-permanent dipole interactions and hydrogen bonding), and use them to explain physical properties such as melting point, boiling point and solubility. "Explain fully" questions reward naming the exact force and saying which force is overcome when a substance changes state.

Types of structure

• Covalent molecular: discrete molecules held by weak intermolecular forces; low melting and boiling points (for example $CO_2$, $I_2$).
• Covalent network: a giant lattice of atoms joined by covalent bonds throughout; very high melting points (for example diamond and silicon dioxide $SiO_2$).
• Ionic: a lattice of oppositely charged ions held by strong electrostatic forces; high melting points, and conducts when molten or dissolved.
• Metallic: positive ions in a sea of delocalised electrons; conducts electricity and is malleable.

Intermolecular forces

London dispersion forces arise from the constant motion of electrons, which creates temporary dipoles that induce dipoles in neighbouring molecules. They are present in all substances and get stronger as the number of electrons (and so the molecular size) increases.

Permanent dipole-permanent dipole interactions occur between polar molecules, which have a permanent separation of charge because of an electronegativity difference between bonded atoms. A molecule is polar overall only if the bond dipoles do not cancel; in $CO_2$ the two bond dipoles point oppositely and cancel, so the molecule is non-polar despite having polar bonds.

Hydrogen bonding

Worked example: ranking boiling points

Explaining properties

The strength of the intermolecular forces, not the covalent bonds, sets the melting and boiling points of molecular substances. For molecules of similar size, those that can hydrogen bond boil at a higher temperature than those with only permanent dipole or dispersion forces. Solubility follows "like dissolves like": polar and hydrogen-bonding molecules dissolve in water, while non-polar molecules dissolve in non-polar solvents.

Examples in context

Hydrogen bonding is the single most important intermolecular force for life. It holds the two strands of DNA together and gives water its high specific heat capacity, which stabilises the climate and the temperature inside cells. In materials science, the contrast between covalent network and covalent molecular structure explains why a diamond-tipped drill (a covalent network of carbon) can cut through almost anything, while solid carbon dioxide (dry ice, a molecular solid) sublimes away to a gas at $-78 \,^{\circ}\text{C}$. Engineers exploit metallic bonding when they choose copper for wiring: the sea of delocalised electrons carries current freely, and the non-directional bonding lets the metal be drawn into thin wires without shattering.

Try this

Q1. Explain why water has a higher boiling point than hydrogen sulfide, $H_2S$. [2 marks]

• Cue. Water molecules form hydrogen bonds (O is highly electronegative), which are stronger than the permanent dipole and dispersion forces in $H_2S$.

Q2. State the strongest intermolecular force present in $CO_2$. [1 mark]

• Cue. London dispersion forces, because $CO_2$ is a non-polar molecule overall.

Q3. Explain why silicon dioxide has a far higher melting point than iodine. [2 marks]

• Cue. $SiO_2$ is a covalent network whose strong covalent bonds must be broken to melt; iodine is covalent molecular, held by weak dispersion forces only.