What you'll learn
Covalent bonding explains how non-metal atoms join by sharing electrons to form molecules such as hydrogen, water, oxygen and methane. In this guide you will learn how covalent bonds form, how to represent them with dot-and-cross and structural diagrams, the difference between simple molecular substances and giant covalent structures, and how each type of structure explains physical properties such as melting point and electrical conductivity. Understanding covalent bonding is essential for explaining the behaviour of gases, water, diamond, graphite and polymers.
Key terms and definitions
Covalent bond — a shared pair of electrons between two non-metal atoms.
Molecule — a group of atoms held together by covalent bonds.
Simple molecular substance — a substance made of small molecules with weak forces between them.
Giant covalent structure — a huge lattice of atoms joined by covalent bonds throughout (e.g. diamond, graphite, silicon dioxide).
Intermolecular forces — the weak forces of attraction between molecules.
Double bond — two shared pairs of electrons between the same two atoms.
Core concepts
How covalent bonds form
When two non-metal atoms bond, they share pairs of electrons so that each atom gains a full outer shell. Each shared pair forms one covalent bond. For example, two hydrogen atoms share one pair of electrons to form H₂; oxygen atoms share two pairs to form a double bond in O₂. The shared electrons are attracted to the nuclei of both atoms, holding them together strongly.
Representing covalent bonds
Covalent molecules can be shown in several ways. Dot-and-cross diagrams show the shared pairs in the overlap of the outer shells. Structural formulae use a line for each bond (H–H, O=O). Displayed formulae show all atoms and bonds. You should be able to draw simple molecules such as H₂, Cl₂, O₂, N₂, HCl, H₂O, NH₃, CH₄ and CO₂.
Simple molecular substances
Most covalent substances are simple molecular — small molecules such as water, carbon dioxide and methane. Within each molecule the covalent bonds are strong, but the forces between molecules (intermolecular forces) are weak. Because only the weak intermolecular forces need to be overcome when melting or boiling, these substances have low melting and boiling points and are often gases or liquids at room temperature. They do not conduct electricity because they have no charged particles free to move. As molecules get bigger, intermolecular forces increase, so melting and boiling points rise.
Giant covalent structures
In giant covalent structures, huge numbers of atoms are joined by covalent bonds throughout the lattice. Examples include diamond, graphite, silicon dioxide and graphene. Because covalent bonds are strong and there are vast numbers of them, these substances have very high melting points. Most do not conduct electricity — except graphite and graphene, which do, because each carbon atom forms only three bonds, leaving delocalised electrons free to move.
Diamond and graphite compared
In diamond, each carbon forms four covalent bonds in a rigid 3D lattice, making it very hard with a very high melting point; it does not conduct electricity. In graphite, each carbon forms three bonds, creating layers that can slide over each other (making it soft and slippery, useful as a lubricant), and the spare delocalised electron per atom lets graphite conduct electricity and heat.
Worked examples
Example 1: Drawing a simple molecule
Describe the bonding in a water molecule.
Oxygen needs two more electrons; it shares one electron with each of two hydrogen atoms, forming two covalent bonds (O–H). Each hydrogen gains a full shell of 2, and oxygen gains a full shell of 8.
Example 2: Explaining a low boiling point
Why does carbon dioxide have a low boiling point?
CO₂ is a simple molecular substance. The covalent bonds within molecules are strong, but the intermolecular forces between molecules are weak and need little energy to overcome, so the boiling point is low.
Example 3: Why graphite conducts
Explain why graphite conducts electricity but diamond does not.
In graphite each carbon forms only three bonds, leaving one delocalised electron per atom free to move and carry charge. In diamond every carbon forms four bonds, so there are no free electrons.
Common mistakes and how to avoid them
Saying you break covalent bonds when boiling a simple molecular substance. You overcome the weak intermolecular forces, not the strong covalent bonds.
Thinking covalent substances conduct electricity. Simple molecular substances have no free charges; only graphite/graphene conduct among covalent structures.
Confusing the strength of covalent bonds with low melting points. The bonds are strong; the forces between molecules are weak — that's why melting points are low.
Drawing the wrong number of shared pairs. Check each atom achieves a full outer shell; use double bonds where needed (O₂, CO₂).
Mixing up diamond and graphite bonding. Diamond = 4 bonds, hard, non-conductor; graphite = 3 bonds, layers, conductor.
Exam technique for Covalent Bonding
Show shared pairs clearly in dot-and-cross diagrams, with full outer shells for every atom.
Distinguish intramolecular (covalent) from intermolecular forces when explaining melting and boiling points.
Link structure to properties — small molecules → low melting points and non-conducting; giant covalent → very high melting points.
Explain graphite's conductivity via delocalised electrons and its softness via sliding layers.
Use correct terms — molecule, covalent bond, intermolecular force, delocalised electron.
Quick revision summary
Covalent bonding is the sharing of electron pairs between non-metal atoms so each gains a full outer shell; each shared pair is one covalent bond, and atoms can share more than one pair (double bonds in O₂ and CO₂). Simple molecular substances (water, carbon dioxide, methane) have strong covalent bonds within molecules but weak intermolecular forces between them, giving low melting and boiling points and no electrical conductivity; melting points rise as molecules get larger. Giant covalent structures (diamond, graphite, silicon dioxide) have atoms covalently bonded throughout, giving very high melting points. Diamond (four bonds each) is hard and non-conducting; graphite (three bonds each) has sliding layers and delocalised electrons that let it conduct electricity and act as a lubricant. When explaining properties, always separate the strong covalent bonds from the weak intermolecular forces, link structure to behaviour, and draw shared pairs accurately with full outer shells.