Chemical bonding and structure
Chemical bonding and structure tests ionic, covalent and metallic bonding and how structure explains physical properties.
Ionic bonding
Ionic bonding occurs between metals and non-metals. Electrons are transferred to form positive and negative ions, held in a giant ionic lattice by strong electrostatic attraction. Ionic compounds have high melting and boiling points, are usually soluble in water and conduct electricity when molten or aqueous because ions are free to move. Solid ionic compounds do not conduct because the ions are fixed.
Covalent and metallic bonding
Covalent bonds form when non-metal atoms share electron pairs. Simple molecular substances have weak intermolecular forces between molecules, so they have low melting points and do not conduct. Giant covalent structures such as diamond, graphite and silicon dioxide have many strong covalent bonds and very high melting points; graphite also conducts because each carbon atom contributes a delocalised electron. Metallic bonding is a lattice of positive ions in a sea of delocalised electrons, giving conduction, malleability and ductility.
Worked example
Explain why sodium chloride conducts electricity when molten but not when solid. In the solid lattice, Na+ and Cl− ions are held in fixed positions and cannot move, so there is no mobile charge. On melting, the lattice breaks down and the ions become free to move, carrying current to the electrodes. The ions are the charge carriers, not electrons.
Traps
Do not say covalent bonds break when a simple molecular solid melts; only the weak forces between molecules are overcome. Metals conduct because of delocalised electrons, not ions. Graphite conducts because of delocalised electrons, but diamond does not. Ionic substances do not conduct in the solid state, even though they contain ions.