Covalent bonding as shared pairs of electrons between non-metals, dot and cross diagrams for simple molecules, simple molecular substances, and giant covalent structures such as diamond, graphite and silicon dioxide.

What this dot point is asking

OCR wants you to explain covalent bonding as shared pairs of electrons between non-metal atoms, draw dot and cross diagrams for simple molecules, describe the properties of simple molecular substances, and describe and explain giant covalent structures such as diamond, graphite and silicon dioxide. Covalent bonding is the second of the three bond types.

How covalent bonds form

The number of covalent bonds an atom forms is usually the number of electrons it needs to complete its outer shell. For example, hydrogen forms one bond, oxygen forms two, nitrogen forms three and carbon forms four. A double bond is two shared pairs (as in $\text{O}_2$), and a triple bond is three shared pairs (as in $\text{N}_2$).

Dot and cross diagrams for molecules

A dot and cross diagram for a molecule shows the outer-shell electrons of each atom, with the shared pairs sitting in the overlap between the bonded atoms (one dot and one cross per bond).

Simple molecular substances

Many covalent substances exist as small, separate molecules (such as $\text{H}_2\text{O}$, $\text{CO}_2$, $\text{Cl}_2$ and $\text{CH}_4$). These are simple molecular substances. The covalent bonds within each molecule are strong, but the forces of attraction between the molecules (intermolecular forces) are weak.

Because only the weak intermolecular forces need to be overcome to melt or boil them, simple molecular substances have low melting and boiling points and are often gases or liquids at room temperature. They do not conduct electricity because the molecules are neutral overall and have no free electrons or ions. The intermolecular forces get stronger as the molecules get bigger, so larger molecules have higher boiling points.

Giant covalent structures

• Diamond. Each carbon forms four covalent bonds in a rigid 3D lattice. This makes diamond very hard with a very high melting point. It has no free electrons, so it does not conduct electricity.
• Graphite. Each carbon forms only three covalent bonds, giving layers of hexagonal rings. The layers are held by weak forces and can slide, so graphite is soft and slippery. Each carbon has one delocalised electron, so graphite conducts electricity and heat.
• Silicon dioxide (silica, the main part of sand). Each silicon is bonded to four oxygens and each oxygen to two silicons, in a rigid giant lattice, so it is hard with a very high melting point and does not conduct.

Graphene (a single layer of graphite) and fullerenes (such as $\text{C}_{60}$ and nanotubes) are also covalent forms of carbon with useful properties.