Covalent Bonding and Molecular Substances
In 1926, Gilbert Lewis at UC Berkeley mapped the first covalent bond, showing two atoms sharing electrons rather than transferring them, explaining why water boils at 100 °C not −80 °C.
Printable Worksheets
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Q1 · Water, oxygen, and glucose are all essential for life, what do you think these molecules have in common at the atomic level, and how might their atoms be held together?
Q2 · If ionic compounds form when electrons are transferred between atoms, why do you think some atoms prefer to share electrons instead?
● Know
- How covalent bonds form by sharing electron pairs
- The difference between single, double, and triple bonds
- The properties of molecular substances (low melting points, poor conductors)
● Understand
- Why molecular substances have LOW melting points (intermolecular forces break, not covalent bonds)
- Why covalent compounds generally do not conduct electricity
- The difference between intra- and inter-molecular forces
● Can do
- Draw dot-and-cross diagrams for simple molecules (H₂, O₂, H₂O, CH₄)
- Predict properties of molecular substances
- Explain the melting misconception for covalent compounds
A covalent bond forms when two non-metal atoms share a pair of electrons, with each atom contributing one electron to the shared pair. Unlike ionic bonding (which transfers electrons from a metal to a non-metal), covalent bonding shares them. Each atom still reaches a stable, full outer shell by counting both its own electrons and the shared pair as its own. A single bond involves one shared pair; a double bond involves two shared pairs; a triple bond involves three shared pairs. Dot-and-cross diagrams show the electrons from each atom with different symbols (a dot and a cross) so you can track where each electron came from.
Water (H₂O) is a classic example: oxygen has 6 valence electrons and needs 2 more, and each hydrogen has 1 electron and needs 1 more. Oxygen forms a single covalent bond with each of two hydrogen atoms, and each shared pair gives both atoms an extra electron to count. Oxygen ends up with 8 outer electrons and each hydrogen with 2, so all three atoms reach a full outer shell. The shared electrons sit in the region between the bonded atoms, where they are attracted to both positive nuclei at once, which is what holds the molecule together.
Nitrogen gas (N₂): each nitrogen has 5 valence electrons and needs 3 more to reach a full outer shell. Two nitrogen atoms share 3 pairs of electrons, a triple bond. This triple bond is extraordinarily strong (945 kJ/mol to break), which is why nitrogen gas is so unreactive and makes up 78% of Australia's atmosphere without reacting with oxygen.
Liquid nitrogen (N₂, covalently bonded) is produced in large quantities at BOC Gas's Australian plants for food freezing, medical procedures, and laboratory work. Its covalent triple bond makes it chemically inert, safe to handle and use in proximity to reactive chemicals, while its very low boiling point (−196 °C) makes it an ideal coolant.
Simple covalent molecules have definite 3D shapes, and the shape depends on how many atoms are joined to the central atom. You do not need to calculate exact angles in Year 9, but it helps to recognise a few common shapes. Water (H₂O) is bent, the two hydrogen atoms sit on the same side of the oxygen rather than in a straight line. Carbon dioxide (CO₂) is linear, the two oxygen atoms sit on opposite sides of the carbon in a straight line. Methane (CH₄) has its four hydrogen atoms spread evenly around the central carbon, giving a three-dimensional shape rather than a flat one.
Shape matters because it affects how molecules behave. Water's bent shape gives it a slightly negative end (near the oxygen) and a slightly positive end (near the hydrogens), so the molecule as a whole is polar, like a tiny magnet with two ends. This is one reason water is such a good solvent and can dissolve ionic compounds like salt. Other molecules, such as carbon dioxide, are symmetrical and behave as non-polar molecules instead. The key idea for now is simply that the same atoms arranged in different shapes produce molecules with different properties.
Water (H₂O) and carbon dioxide (CO₂) both contain three atoms, but water is bent while carbon dioxide is linear. That difference in shape is one reason water is a liquid at room temperature and dissolves many substances, while carbon dioxide is a gas that does not dissolve ionic compounds well.
Water's bent shape, and the polarity that results, is the reason Sydney's drinking water supply can be treated using alum (aluminium sulfate), an ionic compound. The polar water molecules dissolve it completely, allowing the dissolved ions to clump suspended particles together and clarify billions of litres of drinking water per day.
Simple covalent molecules (H₂O, CO₂, CH₄, O₂) are small. The covalent bonds within each molecule are strong, but the forces between molecules (intermolecular forces) are weak. This is why simple molecules have low melting and boiling points, it only takes a small amount of energy to separate the molecules from each other, even though the bonds within each molecule remain intact. CO₂ sublimes at −78.5 °C because its molecules are so weakly attracted to each other.
In contrast, giant covalent structures (diamond, graphite, silicon dioxide/quartz) are not individual molecules, they are entire lattices of covalently bonded atoms extending in three dimensions with no definite end. Diamond has every carbon bonded to 4 others in a continuous 3D network. To melt diamond you must break actual covalent bonds, hence its melting point above 3500 °C. Silicon dioxide (quartz, SiO₂) is similar: each silicon bonds to 4 oxygens in a giant lattice, making quartz sand the most durable common mineral on Earth's surface.
Dry ice is solid CO₂, a simple covalent molecule that sublimes at −78.5 °C under atmospheric pressure. Diamond is giant covalent carbon, melts above 3500 °C. Both are covalent carbon compounds, but the structural scale (small molecule vs giant lattice) differs by a factor of roughly 10,000 in melting point.
Synthetic diamond, made by CSIRO researchers and companies like Element Six, is used in drill bits for Australian mineral exploration. The giant covalent structure of synthetic diamond gives it hardness of 10 Mohs, the maximum possible, allowing drill bits to cut through the hardest rock formations in WA and Queensland mines.
Simple covalent molecules like carbon dioxide have low melting points because the forces molecules are weak. The covalent bonds within each molecule are , but little energy is needed to separate the molecules. Giant covalent structures such as diamond and are lattices of atoms with no definite end. In diamond, every carbon atom is covalently bonded to others in a 3D network. Melting diamond requires breaking actual bonds, so its melting point is extremely high.
At the start of this lesson, you heard about copper wire, which bends, stretches, and hammers into shape without breaking, versus a ceramic cup that shatters when you drop it. Both are solid and strong, but their bonding is completely different. Metallic bonding explains why metals are ductile and conductive in a single elegant model.
Now that you've worked through the lesson, can you explain what "sea of delocalised electrons" means and how it accounts for both the conductivity and the bendability of metals? How does this compare to the ionic and covalent bonding you studied in earlier lessons?
Q1. Explain how a covalent bond forms between two hydrogen atoms. Draw a dot-and-cross diagram to support your explanation.
Q2. A student claims that when ice melts, covalent bonds in water molecules are broken. Evaluate this claim and provide a correct explanation of what actually happens.
Q3. Compare the structure and bonding of sodium chloride (ionic) and water (covalent). Use this comparison to explain why these substances have such different melting points.
Model answers (click to reveal)
SAQ 1 (2 marks)
Model answer: Each hydrogen atom has one electron in its outer shell but needs two electrons to be stable (a full first shell). When two hydrogen atoms come together, they share their electrons so that one pair of electrons sits between the two nuclei. This shared pair is the covalent bond, and the electrostatic attraction between the two positive nuclei and the shared negative electrons holds the atoms together. By sharing, each hydrogen atom effectively has two electrons, so both reach a stable, full outer shell. Dot-and-cross diagram: write H on the left and H on the right with one dot and one cross between them, H(•×)H, where the dot is the electron from one atom and the cross is the electron from the other, showing the single shared pair.
SAQ 2 (3 marks)
Model answer: The student's claim is incorrect. When ice melts, the covalent bonds within each water molecule (the strong bonds holding each H to the O atom) are not broken, because each molecule stays intact as H2O in both the solid and the liquid. What actually happens is that melting overcomes the much weaker intermolecular forces, the attractions between separate water molecules, allowing the molecules to move past one another and flow as a liquid. Because these intermolecular forces are weak compared with covalent bonds, only a small amount of energy is needed, which is why ice melts at the low temperature of 0 °C rather than the very high temperature that breaking covalent bonds would require.
SAQ 3 (3 marks)
Model answer: Sodium chloride is a giant ionic lattice made of positive Na+ ions and negative Cl- ions arranged in a regular repeating pattern, held together by strong electrostatic forces of attraction between the oppositely charged ions throughout the whole structure. Water is a molecular substance made of separate H2O molecules; the H and O atoms inside each molecule are joined by strong covalent bonds, but the molecules are attracted to each other only by weak intermolecular forces. To melt sodium chloride, a very large amount of energy is needed to overcome the strong ionic attractions across the entire lattice, so it has a high melting point (about 801 °C). To melt ice, only the weak intermolecular forces between water molecules need to be overcome, so water has a very low melting point (0 °C). The difference in melting points comes from the strength of the forces that must be broken, not from the strength of the bonds inside each molecule.