What you'll learn

• How atoms form covalent bonds by sharing pairs of electrons.
• How to draw dot-and-cross diagrams for common covalent molecules.
• Why simple molecular substances usually have low melting and boiling points.
• How diamond, graphite and $C_{60}$ fullerene get their different properties from their structures.

Before covalent bonding: outer electrons

Atoms have a small positive nucleus surrounded by negative electrons in shells. The outer-shell electrons are the electrons in the highest occupied shell, and these are the electrons involved in chemical bonding.

For the atoms you meet here, the group number in the Periodic Table tells you the number of outer-shell electrons:

• Group 1 atoms have 1 outer electron.
• Group 4 atoms, such as carbon, have 4 outer electrons.
• Group 5 atoms, such as nitrogen, have 5 outer electrons.
• Group 6 atoms, such as oxygen, have 6 outer electrons.
• Group 7 atoms, the halogens, have 7 outer electrons.

Most non-metal atoms become more stable by getting a full outer shell. Hydrogen is full with 2 electrons; carbon, nitrogen, oxygen and the halogens are usually full with 8 electrons in these IGCSE examples.

What is a covalent bond?

Covalent bonding usually happens between non-metal atoms. Instead of one atom transferring electrons to another, the atoms share electrons.

A shared pair of electrons counts towards the outer shell of both bonded atoms. One shared pair makes a single covalent bond. Two shared pairs make a double covalent bond. Three shared pairs make a triple covalent bond.

The electrostatic attraction in a covalent bond

An electrostatic attraction is an attraction between opposite charges. In a covalent bond, the shared pair of negative electrons is attracted to the positive nuclei of both atoms.

So, the “sharing” idea is helpful, but the force holding the atoms together is still an attraction between positive and negative charges.

Dot-and-cross diagrams

A dot-and-cross diagram shows the outer electrons in a covalent molecule. Dots and crosses are just a drawing convention: they show which atom the electrons originally came from.

To draw a dot-and-cross diagram:

1. Count the outer electrons on each atom.
2. Arrange atoms so that each bond contains one shared pair of electrons.
3. Add any non-bonding outer electrons as lone pairs, which are pairs of electrons not used in bonding.
4. Check that hydrogen has 2 electrons and most other atoms have 8 electrons in their outer shell.

The diagram below shows common patterns for single, double and triple covalent bonds.

Diatomic molecules

A diatomic molecule contains two atoms.

You should be able to represent:

• Hydrogen, $H_2$: one shared pair.
• Oxygen, $O_2$: two shared pairs, so a double bond.
• Nitrogen, $N_2$: three shared pairs, so a triple bond.
• Halogens, such as $F_2$, $C{l}_2$, $B{r}_2$ and $I_2$: one shared pair; each halogen also has three lone pairs.
• Hydrogen halides, such as $HF$, $HCl$, $HBr$ and $HI$: one shared pair between hydrogen and the halogen.

Common inorganic molecules

You also need these patterns:

• Water, $H_{2}O$: oxygen forms two single bonds and has two lone pairs.
• Ammonia, $N{H}_3$: nitrogen forms three single bonds and has one lone pair.
• Carbon dioxide, $C{O}_2$: carbon forms two double bonds, giving the structure O=C=O.

Organic molecules with up to two carbon atoms

An organic molecule is a carbon-based molecule, usually containing carbon bonded to hydrogen and sometimes other atoms such as halogens.

Key examples:

• Methane, $C{H}_4$: carbon forms four single C–H bonds.
• Ethane, $C_2{H}_6$: the two carbon atoms have a single C–C bond; each carbon also forms three C–H bonds.
• Ethene, $C_2{H}_4$: the two carbon atoms have a C=C double bond; each carbon also forms two C–H bonds.
• Molecules containing halogen atoms, such as chloromethane, $C{H}_{3}Cl$: carbon still forms four single bonds, and chlorine forms one single bond plus three lone pairs.

Simple molecular structures

Many covalent substances are made of small, separate molecules. These are called simple molecular substances.

Inside each molecule, the atoms are joined by strong covalent bonds. Between separate molecules, there are much weaker attractions.

Simple molecular substances are often gases or liquids at room temperature. If they are solids, they usually have low melting points compared with ionic or giant covalent substances.

This is because melting and boiling only require enough energy to overcome the weak intermolecular forces between molecules. The strong covalent bonds inside the molecules are not broken.

Relative molecular mass and boiling point

The relative molecular mass, $M_r$, is the sum of the relative atomic masses, $A_r$, of all the atoms in a molecule.

In general, as $M_r$ increases, the intermolecular forces between molecules become stronger. More energy is then needed to separate the molecules, so melting and boiling points increase.

Giant covalent structures

Not all covalent substances are simple molecular. Some have giant covalent structures.

Giant covalent structures have very high melting and boiling points because many strong covalent bonds must be broken to separate the atoms. This takes a lot of energy.

The comparison below shows how the structures of diamond, graphite and $C_{60}$ fullerene affect their properties.

Diamond

Diamond is a giant covalent structure. Each carbon atom forms four covalent bonds to other carbon atoms in a three-dimensional network.

This makes diamond:

• very hard, because the atoms are held in a rigid network;
• very high melting, because many strong covalent bonds must be broken;
• unable to conduct electricity, because it has no mobile charged particles.

Graphite

Graphite is also a giant covalent structure, but its bonding is different. Each carbon atom forms three covalent bonds in flat layers of hexagons. The layers are held together by weak forces.

Graphite is:

• soft and slippery, because layers can slide over each other;
• high melting, because strong covalent bonds exist within the layers;
• able to conduct electricity along the layers, because it has delocalised electrons that can move and carry charge.

$C_{60}$ fullerene

$C_{60}$ fullerene is made of molecules containing 60 carbon atoms arranged in a hollow sphere. Covalent bonds hold the atoms together within each molecule, but weak intermolecular forces act between separate $C_{60}$ molecules.

So $C_{60}$ fullerene has a much lower melting point than diamond or graphite. Pure $C_{60}$ does not usually conduct electricity because it does not have charged particles free to move throughout the solid.

Do covalent substances conduct electricity?

Covalent compounds do not usually conduct electricity. This is because they usually do not contain ions or delocalised electrons that are free to move.

For a substance to conduct electricity, it needs mobile charged particles:

• mobile ions, as in molten or dissolved ionic compounds;
• or mobile electrons, as in metals and graphite.