1.7 Covalent BondingEdexcel IGCSE Chemistry: Revision notes
Section 1
What is a covalent bond?
A covalent bond is formed between two atoms (usually non-metals) by the sharing of a pair of electrons — one electron from each atom.
The shared pair of electrons is attracted to the nucleus of both atoms, and this electrostatic attraction between the positively charged nuclei and the shared, negatively charged electrons holds the atoms together.
Section 2
How are dot-and-cross diagrams used for covalent molecules?
Dot-and-cross diagrams show the outer-shell electrons of each atom (dots for one atom, crosses for the other) with shared pairs drawn between the atoms.
- Diatomic molecules: e.g. Cl₂, H₂, O₂ (O2 has a double bond, two shared pairs)
- Inorganic molecules: water (H₂O, two single O–H bonds), ammonia (NH₃, three single N–H bonds), carbon dioxide (CO₂, two C=O double bonds)
- Organic molecules with up to two carbons: e.g. methane (CH₄), ethane (C₂H₆), ethene (C₂H₄, C=C double bond)
Each atom in these diagrams typically ends up with a full outer shell (usually 8 electrons, or 2 for hydrogen), shown by the combination of its own and the shared electrons.
Section 3
Why do simple molecular substances have low melting/boiling points?
Substances with simple molecular structures (e.g. H₂O, CO₂, Cl₂) consist of individual molecules held together by weak intermolecular forces, even though the covalent bonds within each molecule are strong.
When melting or boiling occurs, it is these weak intermolecular forces between molecules that break — not the strong covalent bonds within the molecules. Because intermolecular forces are weak, only a small amount of energy is needed to overcome them, so simple molecular substances are typically gases, liquids or low-melting-point solids at room temperature.
As relative molecular mass increases, intermolecular forces generally get stronger (more electrons, larger surface area for attraction), so melting and boiling points generally increase with increasing relative molecular mass.
A very common error is saying covalent bonds break when a simple molecular substance melts or boils — it is the weak forces between molecules that break, not the covalent bonds within them.
Section 4
How do giant covalent structures differ?
Giant covalent structures consist of a huge number of atoms bonded together by strong covalent bonds throughout the whole structure, with no separate molecules.
Because breaking a giant covalent structure means breaking many strong covalent bonds, these substances are solids with very high melting and boiling points.
| Structure | Bonding | Property | Explanation |
|---|---|---|---|
| Diamond | Each carbon bonded to 4 others | Very hard, does not conduct | Rigid 3D lattice of strong covalent bonds; no free electrons or ions |
| Graphite | Each carbon bonded to 3 others in layers | Soft, slippery, conducts electricity | Layers held by weak forces so slide over each other; one delocalised electron per carbon can move along the layers |
| C₆₀ fullerene | Carbon atoms in a hollow ball (60 atoms) | Soft, low melting point relative to diamond/graphite, does not conduct well | Molecules held together by weak intermolecular forces between separate 'ball' molecules |
For graphite's conductivity, explain it in terms of delocalised electrons that are free to move along the layers — simply saying 'graphite conducts' without the electron reason won't score full marks.
Section 5
Do covalent compounds conduct electricity?
Covalent compounds do not usually conduct electricity, whether solid, liquid or (in most cases) dissolved, because they are made of neutral molecules with no free ions or delocalised electrons to carry charge.
Exception: graphite conducts because of its delocalised electrons, even though it is a giant covalent structure of a non-metal element rather than a compound.
Must Know
- A covalent bond is a shared pair of electrons, held by electrostatic attraction to both nuclei
- Simple molecular substances have weak intermolecular forces between molecules but strong covalent bonds within them
- Melting/boiling a simple molecular substance breaks intermolecular forces, not covalent bonds — this is why melting/boiling points are low
- Melting/boiling points generally rise with increasing relative molecular mass (stronger intermolecular forces)
- Giant covalent structures (diamond, graphite, silicon dioxide) have very high melting points due to many strong covalent bonds
- Graphite conducts due to delocalised electrons between layers; diamond does not conduct (no free electrons or ions)
- Covalent compounds do not usually conduct electricity
That's the notes covered.
Carry on to the next subtopic.