Metallic bonding and giant structuresEdexcel A-Level Chemistry: Revision notes

A giant metallic lattice is a regular array of positive metal ions surrounded by a sea of delocalised electrons (the outer electrons of the metal atoms, free to move through the structure). Metallic bonding is the strong electrostatic attraction between the metal ions and the delocalised electrons.

This explains the properties of solid metals:

• Electrical conductivity: delocalised electrons are free to move and carry charge, in both solid and liquid
• High melting and boiling temperatures: a lot of energy is needed to overcome the strong attraction between ions and electrons
• Malleable and ductile: layers of ions slide over each other, but the electrons keep the ions bonded together

The attraction is stronger when the ion has a higher charge, a smaller radius and contributes more delocalised electrons, so Mg melts at a higher temperature than Na.

In a giant ionic lattice (e.g. NaCl) oppositely charged ions are held by strong electrostatic attractions acting in all directions. It has a high melting temperature, is brittle, dissolves in water (polar water molecules can attract the ions) and conducts only when molten or dissolved, because the ions are then free to move. It does not conduct when solid because the ions are fixed.

In a giant covalent lattice (macromolecular structure) atoms are joined by a network of strong covalent bonds throughout the crystal. Examples are diamond, graphite and silicon(IV) oxide, SiO₂. Each silicon is bonded to four oxygen atoms and each oxygen to two silicon atoms. Melting requires many strong covalent bonds to be broken, so SiO₂ melts at 1710 °C. It is hard, insoluble in water and does not conduct, as all its electrons are held in bonds.

Diamond: each carbon is bonded to four others by covalent bonds in a tetrahedral arrangement. All outer electrons are used in bonding, so diamond is very hard, has a very high melting temperature and does not conduct electricity.

Graphite: each carbon is bonded to three others in flat hexagonal layers. The fourth outer electron is delocalised along the layers, so graphite conducts electricity. The layers are held by weak London forces, so they slide over each other and graphite is soft and slippery (a lubricant). Its melting temperature is very high because strong covalent bonds within each layer must be broken.

Graphene: a single layer of graphite, one atom thick. Each carbon forms three covalent bonds and has one delocalised electron, so it conducts electricity very well and is exceptionally strong.

In a simple molecular substance, small molecules are held together in a lattice by intermolecular forces (London forces, permanent dipole forces or hydrogen bonds). The covalent bonds inside each molecule are strong, but they are not broken on melting or boiling.

Iodine, I₂: non-polar molecules held by London forces, which are weak, so iodine is a soft solid that sublimes easily, is insoluble in water and does not conduct.

Ice, H₂O: molecules held in an open hexagonal lattice by hydrogen bonds. Ice has a low melting temperature, does not conduct as a solid, and is less dense than liquid water because the lattice holds the molecules further apart.

Simple molecular substances do not conduct because there are no ions or delocalised electrons. Larger molecules with more electrons have stronger London forces.

Use the evidence from a set of properties to decide the structure:

• Conducts when solid and high melting temperature: metallic (or graphite/graphene)
• High melting temperature, conducts only when molten or in solution, often soluble in water: giant ionic
• Very high melting temperature, hard, does not conduct, insoluble: giant covalent
• Low melting temperature, does not conduct (solubility varies): simple molecular

Check each property in turn and say what it shows about the particles. A single property is rarely enough, for example both a metal and graphite conduct when solid and have high melting temperatures.

To predict a property, state the particles, the forces between them and what must happen for the property to occur.

• Melting and boiling temperature: high if strong bonds (ionic, metallic, covalent network) must be overcome; low if only weak intermolecular forces
• Electrical conductivity: needs mobile charge carriers, either delocalised electrons or free ions
• Solubility in water: ionic compounds dissolve because water molecules attract the ions strongly enough to overcome the lattice; non-polar simple molecules (I₂) cannot form strong enough attractions with water, which is hydrogen-bonded; metals and giant covalent lattices are insoluble

Ionic lattices with higher charges and smaller ions (MgO compared with NaCl) have stronger attractions and higher melting temperatures.