Covalent Bonding
Formation of Covalent Bonds
A covalent bond is formed when two non-metal atoms share a pair of electrons. Each atom contributes one electron to the shared pair.
The bond is the result of a strong electrostatic attraction between the positively charged nuclei of both atoms and the shared (negatively charged) pair of electrons between them.
Sharing electrons allows each atom to achieve a full outer shell, giving it the stable electronic configuration of a noble gas.
Covalent bonds are very strong. Covalently bonded substances can have simple molecular structures or giant covalent structures.
Dot-and-Cross Diagrams
Dot-and-cross diagrams for covalent molecules show the electrons from different atoms as dots or crosses. Shared pairs of electrons (bonding pairs) appear between the two atoms. Unshared pairs (lone pairs) are shown on the outer atom.
Diatomic Molecules
| Molecule | Formula | Bonds |
|---|---|---|
| Hydrogen | H2 | Single bond (1 shared pair) |
| Chlorine | Cl2 | Single bond |
| Oxygen | O2 | Double bond (2 shared pairs) |
| Nitrogen | N2 | Triple bond (3 shared pairs) |
| Hydrogen chloride | HCl | Single bond |
Inorganic Molecules
| Molecule | Formula | Bonds |
|---|---|---|
| Water | H2O | 2 single bonds (2 lone pairs on O) |
| Ammonia | NH3 | 3 single bonds (1 lone pair on N) |
| Carbon dioxide | CO2 | 2 double bonds |
Organic Molecules
| Molecule | Formula | Description |
|---|---|---|
| Methane | CH4 | 4 single C-H bonds |
| Ethane | C2H6 | C-C single bond + 3 H on each C |
| Ethene | C2H4 | C=C double bond + 2 H on each C |
Exam Tip
Each line in a structural formula represents one shared pair of electrons. A double bond (=) means 4 electrons are shared. In O2, two pairs are shared between the two oxygen atoms.
Simple Molecular Structures
Simple molecular structures consist of small molecules (e.g. H2O, CO2, CH4) held together by covalent bonds within each molecule.
Forces between molecules (intermolecular forces) are much weaker than the covalent bonds within molecules.
Properties of Simple Molecular Structures
- Low melting and boiling points: little energy is needed to overcome the weak intermolecular forces between molecules. The strong covalent bonds within molecules are NOT broken during melting or boiling.
- Usually gases or liquids at room temperature (or solids with low melting points).
- Do not conduct electricity: molecules carry no overall electric charge, so there are no freely moving charged particles to carry a current.
Effect of Molecular Mass on Melting/Boiling Points
As the relative molecular mass of a substance increases, its melting and boiling points generally increase. Larger molecules have more electrons, leading to stronger intermolecular forces that require more energy to overcome.
Giant Covalent Structures
Giant covalent structures (macromolecules) contain a huge number of non-metal atoms all bonded together by strong covalent bonds in a continuous network.
Properties of Giant Covalent Structures
- Very high melting and boiling points: the large number of strong covalent bonds throughout the structure require a lot of energy to break.
- Solid at room temperature.
Diamond
- Each carbon atom is bonded to 4 other carbon atoms by strong covalent bonds in a tetrahedral arrangement.
- This forms an extremely rigid, three-dimensional lattice.
- Very hard (hardest natural mineral) because of the strong covalent bonds in all directions.
- Very high melting point for the same reason.
- Does not conduct electricity: all electrons are used in covalent bonds; there are no free electrons to carry charge.
- Uses: cutting tools, drill bits, jewellery.
Graphite
- Each carbon atom is bonded to 3 other carbon atoms in layers of hexagonal rings.
- The layers are held together by weak intermolecular forces (not covalent bonds), so they can slide over each other.
- Each carbon atom has one delocalised electron (not used in bonding) that is free to move between the layers.
- Soft and slippery (layers slide): used as a lubricant, in pencil leads.
- High melting point: the covalent bonds within layers are strong.
- Conducts electricity: the delocalised electrons can carry charge. This is unusual for a non-metal.
C60 Fullerene (Buckminsterfullerene)
- 60 carbon atoms joined together forming 20 hexagons and 12 pentagons in a hollow sphere shape (like a football).
- Each carbon is bonded to 3 others; one electron per carbon is delocalised within the molecule.
- Simple molecular structure (not giant covalent): there are weak intermolecular forces between individual buckyballs.
- Does not conduct electricity: delocalised electrons are confined within one molecule and cannot migrate between molecules.
- Relatively low melting point (for a carbon allotrope): little energy needed to overcome weak intermolecular forces.
- Uses: drug delivery (hollow cage can trap molecules), lubricants, catalyst supports.
Graphene
- A single layer of graphite; a single layer of hexagonally arranged carbon atoms.
- Properties similar to graphite: strong, conducts electricity, very thin and lightweight.
- Useful in electronics and composite materials.
| Structure | Bonds | Hard/Soft | Conducts? | Melting Point |
|---|---|---|---|---|
| Diamond | 4 covalent bonds per C | Very hard | No | Very high |
| Graphite | 3 covalent bonds per C; layers with delocalised e- | Soft/slippery | Yes | Very high |
| C60 fullerene | 3 covalent bonds per C; simple molecular | Soft | No | Relatively low |
| Graphene | 3 covalent bonds per C | Very strong | Yes | Very high |
Conductivity of Covalent Compounds
Most covalent compounds do not conduct electricity in any state because they have no freely moving charged particles (no ions, no delocalised electrons) to carry a current.
Exceptions: graphite and graphene conduct because they have delocalised electrons.
Exam Questions
Q: Explain why substances with simple molecular structures have low melting points.
Simple molecular substances have weak intermolecular forces between the molecules. These forces require little energy to overcome, so the melting point is low. Note: the strong covalent bonds within the molecules are not broken on melting.
Q: Explain why diamond is hard but graphite is soft.
In diamond, each carbon atom is bonded to four others by strong covalent bonds in a rigid three-dimensional lattice, making it very hard. In graphite, each carbon is bonded to only three others in flat layers; the layers are held together by weak intermolecular forces only, allowing them to slide over each other, making graphite soft and slippery.
Q: Explain why graphite conducts electricity but diamond does not.
In graphite, each carbon atom forms only 3 covalent bonds, leaving one electron per atom delocalised and free to move through the structure, enabling electrical conduction. In diamond, all 4 outer electrons of each carbon atom are used in covalent bonds, leaving no free electrons to carry charge.
Q: Explain why the boiling point increases from methane (CH4) to butane (C4H10).
Larger molecules have more electrons, leading to stronger intermolecular forces. More energy is needed to overcome these forces, so the boiling point is higher.
Exam Tip
When explaining the low melting or boiling point of a simple covalent substance, always specify that it is the WEAK INTERMOLECULAR FORCES that are broken, NOT the covalent bonds. Confusing these is one of the most common errors in exam answers on this topic.