1. The four structure types
To explain a property, first name the structure, then the particles, then the force that must be overcome.
| Structure | Particles | Force holding the structure together |
|---|---|---|
| Giant ionic lattice | Positive and negative ions | Ionic bonds: electrostatic attraction between oppositely charged ions |
| Giant metallic lattice | Positive metal ions and delocalised electrons | Metallic bonding: attraction between ions and delocalised electrons |
| Macromolecular (giant covalent) | Atoms | Covalent bonds throughout the whole network |
| Simple molecular | Discrete molecules | Intermolecular forces between molecules (covalent bonds inside each molecule are not broken on melting) |
2. The six required crystals
Sodium chloride (ionic)
A giant lattice of Na+ and Cl− ions, each surrounded by six ions of the opposite charge. Melting needs a lot of energy to overcome the many strong attractions between oppositely charged ions, so the melting point is high. Solid NaCl does not conduct because its ions are fixed in place; molten or dissolved, the ions are free to move and carry charge.
Magnesium (metallic)
A giant lattice of Mg2+ ions with delocalised electrons. The strong attraction gives a high melting point. Delocalised electrons can move through the structure, so magnesium conducts as a solid and a liquid.
Diamond (macromolecular)
Each carbon atom forms four covalent bonds in a tetrahedral arrangement, making a 3D network. Melting means breaking many strong covalent bonds, so the melting point is very high. Diamond is very hard and does not conduct: every outer electron is in a bond, so there are no delocalised electrons or ions to carry charge.
Graphite (macromolecular)
Each carbon atom forms three covalent bonds in flat layers of hexagons. The fourth outer electron from each carbon is delocalised across the layer, so graphite conducts along the layers. The layers are held together by weak London forces, so they can slide over each other, which makes graphite soft and slippery. Its melting point is still very high, because melting requires breaking the strong covalent bonds within the layers.
Iodine (simple molecular)
Iodine is made of discrete I2 molecules in a regular lattice, held together by London forces. When iodine melts or sublimes, it is these intermolecular forces that are overcome, not the I–I covalent bond. The forces are weak, so the melting point is low compared with giant ionic or covalent structures. There are no ions or delocalised electrons, so it does not conduct.
Ice (simple molecular)
Ice is made of H2O molecules held together by hydrogen bonds (plus London forces) in an open, regular lattice. Melting overcomes some of the hydrogen bonds, not the O–H covalent bonds. Hydrogen bonding and the open lattice are covered in Part 4.
Extra example: silicon dioxide
Silica (SiO2) is also macromolecular: each silicon is covalently bonded to four oxygens, and each oxygen to two silicons, in a 3D network. It has a very high melting point and does not conduct.
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Six crystal structures (one card each)
Labels to include:
- Sodium chloride: alternating Na⁺ and Cl⁻ in a cubic lattice; each ion has 6 opposite neighbours
- Magnesium: regular Mg²⁺ ions with delocalised electrons
- Diamond: each C bonded to 4 C, tetrahedral, 3D network
- Graphite: hexagonal layers, 3 bonds per C, delocalised electrons in layers, weak forces between layers
- Iodine: I₂ molecules in a regular lattice, London forces between molecules
- Ice: H₂O molecules, hydrogen bonds in an open tetrahedral network
Each card should show the particles and label the force that is overcome on melting. The written descriptions above contain everything the diagrams would show.
3. Comparison table
| Structure | Melting point | Conducts as solid? | Conducts when molten? | Examples |
|---|---|---|---|---|
| Ionic | High | No (ions fixed) | Yes (ions mobile); also in solution if soluble | NaCl, MgO |
| Metallic | Generally high (varies) | Yes (delocalised electrons) | Yes | Mg, Na, Cu |
| Macromolecular | Very high | Diamond/silica: no; graphite: yes | — | Diamond, graphite, SiO2 |
| Simple molecular | Low | Generally no | Generally no: no appreciable concentration of mobile charged particles | I2, ice, CH4 |
4. Explaining melting and boiling
Build the answer as a chain:
- Name the structure.
- Name the force and the particles it acts between.
- Say how strong the force is (and why, if comparing).
- Link to the amount of energy needed to overcome it.
Cooling is the reverse: as the forces re-form, energy is released (exothermic).
5. Conductivity
Conduction needs mobile charged particles: ions or delocalised electrons.
- Ionic: no in the solid; yes when molten, or in aqueous solution if the compound dissolves.
- Metallic: yes as solid and liquid.
- Graphite: yes along the layers; diamond and silica no.
- Simple molecular: generally no in the pure solid or liquid: no appreciable concentration of mobile charged particles. Avoid saying their solutions never conduct: some (such as HCl) react with water to form ions.
6. Solubility, brittleness and malleability
Solubility: many ionic compounds dissolve in water, but not all (e.g. calcium carbonate is insoluble), so avoid a blanket rule. Simple molecular substances vary: some dissolve, some react with water, many are insoluble. Giant covalent structures generally do not dissolve in water; metals do not simply dissolve as intact atoms, although some react with water or aqueous reagents.
Ionic solids are brittle. If a force shifts one layer of the lattice, ions of the same charge end up next to each other. They repel, and the crystal splits.
Metals are malleable and ductile. Layers of metal ions can slide over each other while the delocalised electrons keep attracting them, so the bonding is kept.
7. Identifying an unknown structure
Use evidence carefully. Conductivity in the solid alone does not prove a metal, because graphite also conducts. Look at several properties together.
8. Quick checks
Finesse practice: indicative answers to check your reasoning, not official mark allocations.
Q1. Why does graphite conduct electricity but diamond does not?Show answer
In graphite, each carbon forms three bonds, so one electron per carbon is delocalised and can move along the layers. In diamond, all four outer electrons are in bonds, so there are no mobile charge carriers.
Q2. Why does solid sodium chloride not conduct, but molten sodium chloride does?Show answer
In the solid, the ions are fixed in the lattice. When molten, the ions are free to move and carry charge.
Q3. What is overcome when iodine sublimes?Show answer
The London forces between I2 molecules. The I–I covalent bonds stay intact.
Q4. Explain why ionic crystals are brittle.Show answer
When a layer shifts, ions with the same charge line up next to each other. They repel, so the crystal breaks apart.
Q5. A solid melts at 98 °C and conducts as a solid. Which structure is it, and why?Show answer
Metallic (sodium, for example). It conducts as a solid, which suggests delocalised electrons; graphite also conducts but has a very high melting point, so a low-melting conductor points to a metal with weak metallic bonding.
9. Sources
Sources and examiner guidance (reviewed 1 October 2026)
- Chemrevise — AQA 1.3 Bonding revision guide (N. Goalby) — primary content reference (structures and properties)
- AQA 7405 specification — 3.1.3 Bonding — 3.1.3.4 Bonding and physical properties
- AQA 7404/1 mark scheme, June 2023 — Q02.1: structure named with the bonding
- AQA 7404/1 examiner report, June 2022 — Q4: covalent bonds wrongly said to break on boiling
Finesse Tuition is not endorsed by AQA or Chemrevise. All explanations and examples here are our own.