Ionic, covalent, dative covalent and metallic bonding, the four crystal structures, electron pair repulsion theory and molecular shapes, bond polarity and electronegativity, and the forces between molecules including van der Waals, dipole-dipole and hydrogen bonding.

What this dot point is asking

AQA wants you to describe ionic, covalent, dative and metallic bonding, link the four crystal structures to physical properties, predict molecular shapes using electron pair repulsion, explain electronegativity and polarity, and identify the intermolecular forces between molecules.

The four types of bonding

Crystal structures and properties

Physical properties follow directly from structure and from which particles attract one another:

• Ionic (e.g. NaCl): a giant lattice of alternating ions; high melting point (strong electrostatic forces between ions); conducts when molten or aqueous (ions free to move) but not as a solid; brittle, because shifting a layer brings like charges together and the lattice repels and splits.
• Macromolecular covalent (e.g. diamond, graphite, $\text{SiO}_2$): very high melting point because melting means breaking many strong covalent bonds. Diamond does not conduct (all four outer electrons localised in bonds); graphite conducts and is slippery because each carbon uses only three bonds, leaving one delocalised electron per atom and weak forces between sheets.
• Molecular covalent (e.g. $\text{I}_2$, ice): low melting point because only weak intermolecular forces (not the covalent bonds) break on melting; does not conduct (no free charges).
• Metallic (e.g. Mg): a lattice of positive ions in a sea of delocalised electrons; high melting point, good electrical and thermal conductor (mobile electrons), and malleable because layers of ions can slide without breaking the bonding.

Molecular shapes

Electron-pair repulsion theory states that the electron pairs around a central atom repel one another and arrange themselves as far apart as possible to minimise repulsion, which fixes the shape. Lone pairs are held closer to the central atom and repel more strongly than bonding pairs, so each lone pair reduces the bond angle between the remaining bonding pairs by about $2.5^\circ$. The order of repulsion strength is lone pair to lone pair, then lone pair to bonding pair, then bonding pair to bonding pair.

Polarity and intermolecular forces

A bond is polar when the two atoms differ in electronegativity (the power of an atom to attract the bonding electron pair), producing a permanent dipole written $\delta+$ and $\delta-$. A molecule can contain polar bonds yet be non-polar overall if the dipoles are arranged symmetrically and cancel: carbon dioxide ($\text{O}=\text{C}=\text{O}$, linear) and tetrachloromethane ($\text{CCl}_4$, tetrahedral) are non-polar despite polar bonds, whereas water is polar because its bent shape leaves a net dipole.

The three intermolecular forces, weakest first, are van der Waals (London) forces (instantaneous induced dipoles, present in all molecules and stronger with more electrons), permanent dipole-dipole forces (between polar molecules), and hydrogen bonding (between an H atom bonded to N, O or F and a lone pair on the N, O or F of a neighbouring molecule). Hydrogen bonding is the strongest of the three and explains water's anomalously high boiling point, ice being less dense than liquid water, and the high boiling points of alcohols and carboxylic acids.