Why does salt melt at 801 °C while oxygen is a gas? The answer is in how atoms connect. Learn the three types of bonding and what they mean for properties.
Atoms rarely exist on their own. Almost every substance you meet is made of atoms joined together, and the way they are joined decides whether the substance is a hard crystal, a flexible metal or a gas. This article explains the three main types of bonding, how to draw them, and how to link bonding to properties.
Most atoms are more stable when their outer shell is full, as in the noble gases. Atoms bond by losing, gaining or sharing electrons to reach a full outer shell. Which of these happens depends on the types of atoms involved.
| Bonding type | Between | What happens to electrons |
|---|---|---|
| Ionic | Metal and non-metal | Electrons are transferred, forming positive and negative ions |
| Covalent | Non-metal and non-metal | Electrons are shared in pairs |
| Metallic | Metal atoms only | Outer electrons are delocalised, shared in a sea around positive ions |
When a metal reacts with a non-metal, the metal atom loses its outer electrons to form a positive ion, and the non-metal atom gains them to form a negative ion. The strong electrostatic attraction between the oppositely charged ions is the ionic bond.
Ionic compounds do not form separate molecules. Instead, each ion is surrounded by ions of the opposite charge in a regular, repeating arrangement called a giant ionic lattice.
The total positive charge must equal the total negative charge, so the compound is neutral overall. Use the charges of the ions (from the group number): Group 1 gives +1, Group 2 gives +2, Group 13 (aluminium) gives +3, Group 16 gives −2 and Group 17 gives −1.
Question: Write the formula of aluminium oxide.
Check: 2 × (+3) + 3 × (−2) = 0.
Question: Write the formula of magnesium chloride.
When two non-metal atoms meet, neither can easily lose electrons, so they share them. A shared pair of electrons is a covalent bond, held between the two positive nuclei. Sharing one pair makes a single bond, two pairs a double bond, and three pairs a triple bond. Each atom counts the shared electrons as part of its outer shell.
| Molecule | Formula | Bonds | Notes |
|---|---|---|---|
| Hydrogen | H₂ | 1 single bond | Each H has 2 electrons (a full first shell) |
| Chlorine | Cl₂ | 1 single bond | Each Cl has 3 lone pairs |
| Oxygen | O₂ | 1 double bond | Two shared pairs |
| Methane | CH₄ | 4 single bonds | Carbon shares 4 pairs |
| Carbon dioxide | CO₂ | 2 double bonds | O=C=O |
| Nitrogen | N₂ | 1 triple bond | Very strong, so nitrogen is unreactive |
Substances such as water, methane, oxygen and carbon dioxide are made of separate small molecules. The covalent bonds inside each molecule are strong, but the forces between molecules (intermolecular forces) are weak. When such a substance melts or boils, only these weak forces are overcome, not the covalent bonds.
Some covalent substances form huge networks of atoms joined by covalent bonds, with no separate molecules. They have very high melting points because many strong bonds must be broken.
| Substance | Structure | Properties and uses |
|---|---|---|
| Diamond | Each carbon bonded to 4 others in a rigid 3-D tetrahedral network | Extremely hard, does not conduct; cutting tools, jewellery |
| Graphite | Each carbon bonded to 3 others in flat layers; one electron per atom is free to move | Soft and slippery (layers slide), conducts electricity; pencils, electrodes, lubricants |
| Silicon dioxide (silica) | Each silicon bonded to 4 oxygen atoms in a 3-D network | Hard, high melting point; sand and quartz, used in glass |
Diamond and graphite are both pure carbon, but their different structures give them completely different properties. Carbon can also form other structures, such as graphene (a single layer of graphite) and fullerenes (cage-like molecules), which are important in current materials research.
In a metal, atoms lose their outer electrons, which become delocalised and move freely among a lattice of positive ions. The attraction between the ions and the electron sea is the metallic bond. This explains why metals conduct electricity and heat and can be bent and drawn into wires (the layers of ions slide while remaining bonded), as discussed in the Metals and Non-Metals article.
| Property | Giant ionic | Simple molecular | Giant covalent | Metallic |
|---|---|---|---|---|
| Melting point | High | Low | Very high | Usually high |
| Conducts when solid? | No | No | No (graphite yes) | Yes |
| Conducts when molten? | Yes | No | No | Yes |
| Hardness | Hard, brittle | Soft or gas/liquid | Very hard (graphite soft) | Malleable and ductile |
| Example | NaCl | H₂O, CO₂ | Diamond, SiO₂ | Fe, Cu |
When answering a question about properties, use the pattern: name the structure, name the particles, say what forces must be overcome or what is free to move. For example: “Sodium chloride has a high melting point because it has a giant lattice with strong electrostatic attraction between oppositely charged ions that need much energy to break.”
Question: Explain why methane (CH₄) has a very low boiling point, although it contains strong covalent bonds.
Answer: Methane is a simple molecular substance. When it boils, only the weak intermolecular forces between molecules are overcome, not the strong covalent bonds inside the molecules. Little energy is needed, so the boiling point is very low.
You can usually decide the bonding type from the elements, using the periodic table:
Question: Classify the bonding in (a) potassium bromide, KBr, (b) sulfur dioxide, SO₂, (c) copper. Say whether each conducts when solid.
Water, H₂O, is a simple molecular substance, yet it boils at 100 °C, while methane, CH₄, a molecule of similar size, boils at about −162 °C. The reason is that the oxygen atom in water pulls the shared electrons towards itself, so the oxygen end of the molecule is slightly negative and the hydrogen ends are slightly positive. Neighbouring water molecules attract each other through these charges in a special strong type of intermolecular force called a hydrogen bond.
Hydrogen bonds hold ice molecules in an open pattern, which is why ice is less dense than liquid water and floats. They also help explain why water is a good solvent for many ionic compounds and why it takes a lot of heat to warm up. This shows that the forces between molecules can be very important, even when the bonds within molecules are covalent.
Try each question first, then tap it to reveal the answer.
Answer: Ionic: it is a metal combined with a non-metal, with electron transfer.
Answer: Na₂O (two Na⁺ for one O²⁻).
Answer: Four (two double bonds).
Answer: In the solid the ions are fixed; when molten they can move and carry charge.
Answer: Each carbon uses three electrons in bonds, leaving one delocalised electron per atom free to move between layers.
Answer: Each carbon atom forms four strong covalent bonds in a rigid 3-D network.
Answer: Only weak forces between molecules need to be overcome.
Suggested use: have students build ionic lattices and molecules with modelling kits or marshmallows and toothpicks, and draw dot-and-cross diagrams for each molecule they make. Practise the three-part answer pattern for property questions until it is automatic. Parents can quiz students on formulae using cards of ions. Please follow your own syllabus for which diagrams (dot-and-cross, outer shells only) are expected.
These reputable sources were consulted to verify the facts in this article:
Copyright note: This article is original EDUSAMBAM educational writing. The diagrams are original inline SVG schematics created for this page and are not copied from any source. Links open external websites; the sources were used to check facts and are suggested for further study.
Last updated: 21 October 2026. Written for students up to age 18 and checked against standard school-level chemistry references. This is general education, not a substitute for your teacher, textbook or syllabus.
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