Covalent Bonds
Section: Chemistry | Syllabus: Cambridge IGCSE Combined Science (0653)
What is a Covalent Bond?
A covalent bond forms when a pair of electrons is shared between two atoms, so that both atoms gain a stable, noble-gas electronic configuration. Covalent bonding occurs between non-metal atoms, since neither atom can easily give up electrons to the other - sharing lets both fill their outer shells at once.
Shared electron pairs sit in the overlap between two atoms and count towards both atoms' outer shells at once
Core Examples: H₂, Cl₂, H₂O, CH₄, NH₃, HCl
| Molecule | Bonds formed | Structure |
|---|---|---|
| Hydrogen, H₂ | 1 single bond (H shares its 1 electron) | H-H |
| Chlorine, Cl₂ | 1 single bond (each Cl shares 1 of its 7 outer electrons) | Cl-Cl |
| Water, H₂O | O forms 2 single bonds, one to each H | H-O-H |
| Methane, CH₄ | C forms 4 single bonds, one to each H | H around a central C (tetrahedral) |
| Ammonia, NH₃ | N forms 3 single bonds, one to each H | H-N-H with a third H and a lone pair on N |
| Hydrogen chloride, HCl | 1 single bond | H-Cl |
Each element tends to form a fixed number of bonds, based on how many more electrons it needs to complete its outer shell: H = 1 bond, O = 2 bonds, N = 3 bonds, C = 4 bonds, Cl = 1 bond.
Worked Example: Methane, CH₄
- Question: carbon has the electronic configuration 2.4. Explain how carbon forms covalent bonds with four hydrogen atoms in methane, CH₄. [2]
- Answer: carbon has 4 outer electrons and needs 4 more to reach a stable configuration of 8, so it shares one electron with each of 4 hydrogen atoms, forming 4 single covalent bonds [1]; each hydrogen shares its 1 electron with carbon, giving each hydrogen a full outer shell of 2 [1]
Supplement Examples: O₂, CO₂, N₂, CH₃OH, C₂H₄
A double bond is 2 shared pairs of electrons (=); a triple bond is 3 shared pairs (≡).
- Oxygen, O₂: each O atom (6 outer electrons) needs 2 more, so the two atoms share 2 pairs → O=O (double bond)
- Carbon dioxide, CO₂: carbon forms a double bond to each oxygen → O=C=O
- Nitrogen, N₂: each N atom (5 outer electrons) needs 3 more, so the two atoms share 3 pairs → N≡N (triple bond)
- Methanol, CH₃OH: carbon forms single bonds to 3 hydrogens and to the oxygen; the oxygen forms a second single bond to the remaining hydrogen (H₃C-O-H)
- Ethene, C₂H₄: the two carbon atoms share a double bond (C=C), and each carbon also forms a single bond to 2 hydrogen atoms
In ethene the two shared pairs of the C=C double bond count towards both carbons' outer shells, alongside their single C-H bonds
Common Mistakes
- ❌ Thinking a double bond means two entirely separate single bonds between different atoms
- ✓ a double bond is 2 shared pairs of electrons between the same two atoms, not two bonds to two different atoms
- ❌ Forgetting that a shared pair of electrons counts towards both atoms' outer shells at the same time
- ✓ each electron pair in a covalent bond is counted twice - once for each atom sharing it - when checking that an atom has a full outer shell
Properties of Simple Molecular Compounds (Supplement)
Simple molecular substances have two very different types of force acting on them, and it is essential not to confuse the two:
- Covalent bonds WITHIN each molecule: strong - these hold the atoms of one molecule together and do not break on melting or boiling
- Intermolecular forces BETWEEN separate molecules: weak - these hold molecules to their neighbours, and are what breaks when a substance melts or boils
You are not required to know the names of specific types of intermolecular force - only that they exist and are weak compared to covalent bonds.
When water boils, only the weak dashed forces between molecules break - the strong O-H bonds within each molecule stay intact, so steam is still made of H₂O molecules
- Low melting and boiling points: only the weak intermolecular forces need to be overcome, so simple molecular substances are often gases or liquids at room temperature (e.g. H₂O: 0-100°C, CH₄: -182°C) - this contrasts with ionic compounds like NaCl (801°C), where every ion is held by strong electrostatic forces throughout a giant lattice
- Poor electrical conductivity: the molecules are electrically neutral overall, and there are no ions and no free electrons available to carry charge
Worked Example: Comparing Structures
- Question: methane (CH₄) is a gas at room temperature, while sodium chloride (NaCl) is a solid with a melting point of 801°C. Explain this difference in terms of structure and bonding. [3]
- Answer: methane is a simple molecular substance - only weak intermolecular forces exist between separate CH₄ molecules, which need little energy to overcome [1]; sodium chloride is a giant ionic lattice - strong electrostatic forces of attraction exist between Na⁺ and Cl⁻ ions throughout the whole structure [1]; overcoming the ionic forces to melt NaCl requires far more energy than overcoming the intermolecular forces in methane [1]
Common Mistakes
- ❌ Thinking that when a simple molecular substance melts or boils, the molecules themselves break apart into separate atoms
- ✓ only the weak forces between molecules break; the strong covalent bonds within each molecule remain intact, so the molecules stay whole
Interactive revision notes, videos and practice questions load below.