Crystal Structures: Ionic, Metallic & Molecular (Part 1/2) Revision Notes | A Level AQA Chemistry

1. Why Crystal Structure Controls Everything

Pick up a lump of sodium chloride and a piece of iodine. Both are crystalline solids at room temperature — yet one conducts electricity when dissolved in water and melts above 800 °C, while the other sublimes gently at room temperature and is a poor conductor in every state. The difference is not the atoms themselves; it is the way those atoms are bonded and arranged into a crystal structure.

AQA identifies four types of crystal structure: ionic, metallic, molecular, and giant covalent (covered in Part 2). Each type has a characteristic arrangement of particles, a characteristic type of bonding, and — crucially — characteristic physical properties that examiners test directly. Master the link between structure, bonding, and properties and you unlock a large chunk of the bonding topic.

2. Ionic Crystals: Sodium Chloride as the Model

Sodium chloride ($\text{NaCl}$) is the textbook example of an ionic crystal, and for good reason — its structure is beautifully regular and its properties follow directly from that regularity. In the solid state, $\text{Na}^+$ and $\text{Cl}^-$ ions pack into a three-dimensional lattice in which every sodium ion is surrounded by six chloride ions, and every chloride ion is surrounded by six sodium ions.

The electrostatic attraction between oppositely charged ions is strong and acts in all directions simultaneously. This means a very large amount of energy is required to separate the ions from one another — which is why ionic compounds have high melting points. For $\text{NaCl}$, the melting point is 801 °C; for $\text{MgO}$, which has doubly charged ions ($\text{Mg}^{2+}$ and $\text{O}^{2-}$), the melting point rises to around 2852 °C because the greater charge density produces a stronger electrostatic attraction.

Conductivity in ionic compounds depends entirely on whether the ions can move. In the solid state, ions are locked in fixed lattice positions — so solid $\text{NaCl}$ does not conduct electricity. When melted or dissolved in water, the ions become free to move and carry charge, so molten or aqueous $\text{NaCl}$ conducts electricity. This is a key distinction that AQA tests regularly.

When drawing an ionic structure for AQA, you need to show a minimum of about 8–12 ions in a 2D grid, clearly alternating between $\text{Na}^+$ and $\text{Cl}^-$, with the ions touching (they are in contact in the real lattice). Label the ions and indicate that the pattern extends in all directions. The key mark-winning phrase is that the structure is a giant ionic lattice with strong electrostatic attractions in all directions.

3. Metallic Crystals: Magnesium and the Sea of Electrons

Metals have a completely different type of giant structure. In magnesium — a typical metallic crystal — each atom loses its outer electrons to form a $\text{Mg}^{2+}$ ion. These positive ions arrange themselves into a regular, close-packed lattice, and the electrons that were released form a delocalised 'sea' that surrounds and permeates the entire lattice.

The strength of metallic bonding — and therefore the melting point — depends on two factors: the charge on the metal ion and the number of delocalised electrons per atom. Magnesium ($\text{Mg}^{2+}$, two delocalised electrons per atom) has a higher melting point (650 °C) than sodium ($\text{Na}^+$, one delocalised electron per atom, melting point 98 °C) because the greater charge and higher electron density produce a stronger electrostatic attraction between the ions and the electron sea.

The delocalised electrons explain two of the most important properties of metals. First, electrical conductivity: because the electrons are free to move throughout the lattice, they can carry charge when a potential difference is applied — metals conduct in both solid and liquid states. Second, thermal conductivity: the mobile electrons transfer kinetic energy rapidly through the structure. These properties are absent in ionic and molecular crystals, which is a reliable way to distinguish metallic structures in exam questions.

4. Molecular Crystals: Iodine and Ice

Molecular crystals are structurally very different from the giant structures above. Here, the particles that occupy the lattice positions are not ions or atoms — they are discrete, covalently bonded molecules. The molecules themselves are held together by intermolecular forces, which are considerably weaker than ionic or metallic bonds. This single fact explains almost every property of molecular crystals.

Iodine ($\text{I}_2$) is the AQA standard example of a simple molecular crystal. Each $\text{I}_2$ molecule consists of two iodine atoms joined by a strong covalent bond. In the solid, these molecules pack into a regular lattice held together by weak van der Waals (London dispersion) forces between the molecules. Because these intermolecular forces are weak, relatively little energy is needed to overcome them — so iodine has a low melting point (114 °C) and has a measurable vapour pressure at room temperature, subliming slowly.

A critical point for the exam: the covalent bonds within each $\text{I}_2$ molecule are strong, but they are not broken when iodine melts or boils. It is the weak van der Waals forces between molecules that are overcome during melting. This distinction between intramolecular bonds and intermolecular forces is one of the most commonly tested ideas in this topic.

Ice ($\text{H}_2\text{O}$(s)) is the second molecular crystal AQA requires you to know. Water molecules are held together in a lattice by hydrogen bonds — the strongest type of intermolecular force, arising from the interaction between the $\delta+$ hydrogen of one molecule and the lone pair on the $\delta-$ oxygen of a neighbouring molecule. Hydrogen bonds are significantly stronger than the van der Waals forces found in analogous small molecules, which is why ice has a higher melting point than expected for such a small molecule — for example, hydrogen sulfide ($\text{H}_2\text{S}$), which has a similar molecular mass, melts at −82 °C. Note that iodine (114 °C) actually has a higher absolute melting point than ice (0 °C); the point is that ice melts far higher than other small molecules of comparable mass.

The open lattice structure of ice — a consequence of the directional hydrogen bonds forcing molecules into a hexagonal arrangement — means ice is less dense than liquid water. This is an unusual property: most solids are denser than their liquid form. For AQA, you should be able to link this anomalous density to the hydrogen-bonded lattice structure.

5. Comparing All Three: Structure, Bonding, and Properties

Now that each structure type is established, the real exam skill is being able to compare them and justify differences in properties. The table below consolidates the key comparisons AQA expects you to make fluently.

Notice the pattern: giant structures (ionic and metallic) have high melting points because many strong bonds or attractions must be broken simultaneously. Molecular structures have low melting points because only weak intermolecular forces need to be overcome. Electrical conductivity requires mobile charge carriers — either free electrons (metallic) or free ions (ionic when molten/dissolved). Molecular crystals have neither, so they do not conduct in any state.

6. Drawing Crystal Structure Diagrams for AQA

AQA can ask you to draw diagrams of these structures showing a specified number of particles. Each structure type has specific conventions that earn marks — and specific errors that lose them. Here is what you need to produce for each type.

Worked example: Drawing an Ionic Lattice Diagram (NaCl)

Draw a diagram to represent the structure of sodium chloride, showing at least 8 ions. [3 marks]

1. Decide on a 3×3 or 4×3 grid of ions. You need at least 8 ions total. — A 3×3 grid gives 9 ions — enough to show the alternating pattern clearly. Larger grids are fine but take more time.
2. Place circles in the grid, alternating between Na+ and Cl-. If the top-left is Na+, the next position (right or down) must be Cl-. — [B1] Correct alternating arrangement of oppositely charged ions. Every Na+ must be surrounded only by Cl- ions and vice versa.
3. Write Na+ inside or beside the sodium circles and Cl- inside or beside the chloride circles. Make the Cl- circles slightly larger (Cl- is larger than Na+). — [B1] Correct labelling with ionic charges shown. Relative size difference is good practice and may earn credit.
4. Add dashed lines or '+...' notation at the edges of your diagram to show the structure continues in all directions. — [B1] Indicating the giant, repeating nature of the lattice. Without this, the diagram implies a finite molecule, which is incorrect.

Answer: A 3×3 (or larger) grid of alternating Na+ and Cl- circles, correctly labelled with ionic charges, with relative size difference shown, and edge notation indicating the lattice extends in all directions.

Worked example: Drawing a Metallic Structure Diagram (Mg)

Draw a diagram to represent the structure of magnesium metal, showing at least 9 positive ions. [3 marks]

1. Draw a regular grid (e.g. 3×3) of circles, each labelled Mg2+. Arrange them in rows and columns (or a close-packed arrangement if you prefer). — [B1] Regular arrangement of positive metal ions. The ions should all be the same size and evenly spaced.
2. Add small dots or minus signs (e-) scattered throughout the spaces between the Mg2+ ions, covering the whole diagram area. — [B1] Delocalised electrons shown throughout the structure — not attached to any single ion. This is the key feature that distinguishes metallic from ionic diagrams.
3. Label the diagram: 'Mg2+ ions in a regular lattice' and 'sea of delocalised electrons'. — [B1] Clear annotation distinguishing the two components of the metallic structure. Examiners award marks for correct identification of both the ions and the electron sea.

Answer: A regular grid of Mg2+ circles with small dots or e- symbols distributed throughout the spaces between them, clearly annotated to show the lattice of positive ions and the sea of delocalised electrons.

For molecular crystals, AQA diagrams show discrete molecules in a regular arrangement. For iodine, draw pairs of I atoms (each pair representing one $\text{I}_2$ molecule) arranged in a regular pattern, with the molecules clearly separated from one another. For ice, draw $\text{H}_2\text{O}$ molecules with dashed lines between them to represent hydrogen bonds, showing the open hexagonal arrangement.

7. Exam Technique: Scoring Full Marks on Structure and Bonding Questions

Structure and bonding questions follow predictable patterns in AQA papers. The most common errors come not from not knowing the content, but from imprecise language or missing the specific detail the mark scheme requires. The guidance below targets the highest-frequency pitfalls.

Part 2 of this lesson extends the crystal structure topic to giant covalent structures — including diamond and graphite — where the bonding is covalent throughout the lattice rather than ionic or metallic. The principles you have built here (giant structure = high melting point; mobile charge carriers = conductivity) apply directly, so the comparison becomes even more powerful once all four types are in place.