Describe covalent bonding as the sharing of electron pairs, draw dot-and-cross diagrams for simple molecules, and contrast the properties of simple molecular substances with giant covalent structures

What this dot point is asking

SEAB wants you to describe covalent bonding as the sharing of pairs of electrons between non-metal atoms so each reaches a full outer shell, draw dot-and-cross diagrams for simple molecules, and contrast the two kinds of covalent substance: simple molecular (such as water and methane) and giant covalent (such as diamond and graphite). The recurring exam skill is explaining the very different properties from the structure.

The answer

Covalent bonding as sharing electrons

When two non-metal atoms bond, neither wants to lose electrons, so instead they share pairs of electrons. Each shared pair is a covalent bond. By sharing, both atoms count the shared electrons as part of their outer shell and so reach a stable full outer shell (a noble-gas configuration).

A single bond is one shared pair; a double bond is two shared pairs (as in oxygen, $\text{O}_2$, or carbon dioxide, $\text{CO}_2$). Electrons in the outer shell that are not in a bond are called lone pairs.

Dot-and-cross diagrams for simple molecules

Show only the outer-shell electrons, one atom's as dots and the other's as crosses, with the shared pairs in the overlap between the atoms:

• Hydrogen ($\text{H}_2$): one shared pair; each H has a full shell of $2$.
• Water ($\text{H}_2\text{O}$): oxygen shares one electron with each of two hydrogens, giving two bonding pairs and two lone pairs.
• Methane ($\text{CH}_4$): carbon shares with four hydrogens, four bonding pairs.
• Carbon dioxide ($\text{CO}_2$): carbon forms a double bond to each oxygen.

Simple molecular substances

In a simple molecular substance the atoms within each molecule are joined by strong covalent bonds, but the separate molecules are attracted to one another by only weak intermolecular forces. This gives the typical properties:

• Low melting and boiling points: only the weak forces between molecules need to be overcome (not the strong covalent bonds), so little energy is needed. Many are liquids or gases at room temperature.
• Do not conduct electricity: there are no free ions or free electrons to carry charge.
• Usually insoluble in water (but soluble in organic solvents).

Giant covalent structures

Some covalent substances form a giant covalent structure: a huge network of atoms all joined by strong covalent bonds. Diamond and graphite (both forms of carbon) and silicon dioxide are examples:

• Diamond: each carbon bonded to four others in a rigid 3D network, making it extremely hard with a very high melting point; all outer electrons are in bonds, so it does not conduct.
• Graphite: each carbon bonded to three others in flat layers, with one delocalised electron per atom between the layers. The layers slide (so graphite is soft and slippery, used as a lubricant), and the free electrons let it conduct electricity.

Giant covalent substances have very high melting points because melting them means breaking a vast number of strong covalent bonds.

Examples in context

Example 1. Why graphite is in pencils and electrodes. Graphite's layers slide over one another, so it leaves a mark on paper and acts as a lubricant, while its delocalised electrons let it conduct, so it is used as inert electrodes in electrolysis. Both uses come directly from its layered giant covalent structure.

Example 2. Why oxygen and nitrogen are gases. The air is mostly simple molecular oxygen ($\text{O}_2$) and nitrogen ($\text{N}_2$), held together only by weak forces between molecules, so they are gases at room temperature. Their low boiling points are exactly what the simple molecular model predicts.

Try this

Q1. State what a covalent bond is. [1 mark]

• Cue. A shared pair of electrons between two (non-metal) atoms.

Q2. Explain why methane has a low boiling point. [2 marks]

• Cue. Methane is simple molecular; only the weak forces between molecules need to be overcome to boil it, not the strong covalent bonds, so little energy is needed.

Q3. Explain why diamond is hard but graphite is soft, even though both are made only of carbon. [3 marks]

• Cue. Diamond has each carbon bonded to four others in a rigid 3D network, so it is hard; graphite has layers (three bonds per atom) held together only weakly, so the layers slide and it is soft.