Covalent bonding occurs when two non-metal atoms each contribute an electron to form a shared pair, with both nuclei attracted to the shared electrons. No electrons are transferred — the atoms stay neutral, held by electrostatic attraction to the shared pair. The same principle produces everything from methane gas to diamond.
How does a covalent bond form?
A covalent bond is a shared pair of electrons between two atoms. Each atom contributes one electron to the shared pair (or in a dative/coordinate bond, one atom contributes both, but that is beyond basic GCSE). Both nuclei are attracted to the negatively charged shared electrons, and this attraction holds the atoms together.
Covalent bonds form between non-metal atoms — including within hydrogen molecules (H–H), between two different non-metals (H–Cl), and within carbon-based compounds. The number of covalent bonds an atom can form depends on how many electrons it needs to complete its outer shell:
| Element | Outer electrons | Electrons needed | Covalent bonds formed |
|---|---|---|---|
| Hydrogen (H) | 1 | 1 | 1 |
| Oxygen (O) | 6 | 2 | 2 |
| Nitrogen (N) | 5 | 3 | 3 |
| Carbon (C) | 4 | 4 | 4 |
| Chlorine (Cl) | 7 | 1 | 1 |
What are single, double, and triple covalent bonds?
| Bond type | Shared pairs | Example | Bond strength |
|---|---|---|---|
| Single (–) | 1 | H₂, H₂O, CH₄ | Weakest (per bond) |
| Double (=) | 2 | O₂, CO₂, C₂H₄ (ethene) | Stronger |
| Triple (≡) | 3 | N₂, C₂H₂ (ethyne) | Strongest (per bond) |
More shared pairs between the same two atoms means a stronger, shorter bond. The N≡N triple bond in nitrogen (N₂) is exceptionally strong (945 kJ mol⁻¹), which makes nitrogen gas very unreactive and requires extreme conditions to break it in the Haber process.
What are simple molecular substances?
Most small covalent compounds form simple molecular substances — discrete molecules held together within the molecule by strong covalent bonds, but with only weak intermolecular forces (Van der Waals/London dispersion forces and dipole-dipole forces) between molecules.
Properties of simple molecular substances:
- Low melting and boiling points — only weak intermolecular forces need to be overcome, not covalent bonds; e.g. water (100 °C) is far lower than NaCl (801 °C).
- Do not conduct electricity — no charged particles free to move; molecules are neutral.
- Often gases or liquids at room temperature — small molecules have very weak intermolecular forces.
Common examples
| Molecule | Formula | Displayed structure | State at room temp |
|---|---|---|---|
| Water | H₂O | H–O–H (two single bonds, 104.5° angle) | Liquid |
| Carbon dioxide | CO₂ | O=C=O (two double bonds) | Gas |
| Methane | CH₄ | C with 4 H atoms | Gas |
| Ammonia | NH₃ | N with 3 H atoms (+ lone pair) | Gas |
| Hydrogen chloride | HCl | H–Cl | Gas |
What are giant covalent structures?
Some covalent substances do not exist as small molecules. Instead, the atoms are bonded covalently in a vast three-dimensional network called a giant covalent structure (also called a macromolecular structure). Breaking these structures requires breaking many strong covalent bonds, giving them very high melting points.
| Substance | Structure | Key property |
|---|---|---|
| Diamond | Each C bonded to 4 others in tetrahedral arrangement | Hardest natural substance; does not conduct electricity |
| Graphite | Each C bonded to 3 others in layers; delocalised electrons between layers | Conducts electricity; layers can slide (lubricant) |
| Silicon dioxide (SiO₂) | Each Si bonded to 4 O atoms; each O to 2 Si | Very high melting point; does not conduct |
Diamond vs graphite — a classic comparison: Both are giant covalent structures of pure carbon, but their properties are completely different because of their different structures. Diamond is extremely hard (every carbon is held rigidly by 4 strong bonds) and does not conduct electricity (no free electrons). Graphite has delocalised electrons between layers that can carry charge, so it conducts; the layers slide over each other easily (weak forces between layers), making it a lubricant and useful in pencil leads.
What are the properties of covalent substances and why?
| Property | Simple molecular | Giant covalent |
|---|---|---|
| Melting/boiling point | Low | Very high |
| Electrical conductivity | None (usually) | None (diamond, SiO₂); yes (graphite) |
| Solubility in water | Variable | Generally insoluble |
| Hardness | Soft (solid form) | Hard |
Frequently asked questions
Why do simple covalent molecules have low melting points?
When a simple molecular substance melts, the molecules themselves do not break apart — the covalent bonds within the molecules remain intact. What breaks are the weak intermolecular forces (principally Van der Waals forces) between separate molecules. These intermolecular forces are much weaker than covalent bonds, so only a small amount of energy is needed to separate the molecules. This is why methane, oxygen, and water all have much lower melting points than ionic or giant covalent substances, where primary bonding must be broken.
Why does graphite conduct electricity but diamond does not?
In graphite, each carbon atom forms three covalent bonds, leaving one electron per carbon atom that is not directly involved in bonding. These electrons become delocalised — free to move between the layers of graphite sheets. When a voltage is applied, these electrons carry charge, so graphite conducts electricity. In diamond, every electron is used in a covalent bond (each carbon forms four single bonds), leaving no free electrons. Without mobile charge carriers, diamond cannot conduct electricity.
How do you draw a dot-and-cross diagram for water (H₂O)?
Oxygen has 6 outer electrons (represented as dots) and needs 2 more to complete its shell. Each hydrogen has 1 electron (represented as crosses). Draw the oxygen atom with 6 electrons. Draw each hydrogen atom sharing one electron with oxygen. The shared pairs are shown as one dot and one cross in the overlap region — these are the two single O–H covalent bonds. Oxygen will then have 4 electrons in bonding pairs and 4 electrons in lone pairs (2 lone pairs), giving a total of 8 electrons in its outer shell.
What is the difference between a covalent bond and an intermolecular force?
A covalent bond is the strong attraction between two nuclei and the shared electron pair between them — it holds atoms together within a molecule. An intermolecular force is a much weaker attraction between separate molecules (such as Van der Waals forces, dipole-dipole interactions, or hydrogen bonds). When a simple molecular substance melts or boils, it is the intermolecular forces that are overcome, not the covalent bonds. This distinction is critical for explaining why water can be boiled at 100 °C even though the O–H bond energy is 460 kJ mol⁻¹.
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