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

1. Picking Up From Part 1: The Four Crystal Types

In Part 1 you met the four categories of crystal structure: ionic, metallic, giant covalent, and molecular. Now the focus shifts to the specific examples AQA expects you to know in detail — and more importantly, to the why behind each substance's properties.

The central question running through this lesson is: how does the arrangement of particles and the nature of the forces between them determine what a substance can and cannot do? Every property — melting point, electrical conductivity, hardness, solubility — traces back to structure.

2. Giant Covalent Structures: Diamond, Graphite, Graphene, and Fullerenes

All four of these substances are allotropes of carbon — different structural forms of the same element. The fact that they share the same atoms yet have wildly different properties is one of the most striking demonstrations of how structure governs behaviour.

2a. Diamond

In diamond, each carbon atom forms four strong covalent bonds arranged tetrahedrally to four neighbouring carbon atoms. This pattern repeats throughout the entire crystal, creating a rigid, three-dimensional giant covalent lattice with no weak points.

Because every bond in the structure is a strong, short C–C covalent bond, an enormous amount of energy is needed to break the lattice apart. This explains diamond's exceptionally high melting point (above 3500 °C). The rigidity of the tetrahedral network also makes diamond the hardest naturally occurring substance.

Diamond does not conduct electricity. All four of each carbon's outer electrons are used in bonding — there are no delocalised electrons free to carry charge. This is a key contrast with graphite and metals.

2b. Graphite

Graphite has a layered structure. Each carbon atom forms three covalent bonds to three neighbours in the same plane, creating flat hexagonal rings that extend into large sheets. The fourth outer electron from each carbon atom is delocalised across the entire layer.

The layers themselves are held together only by weak van der Waals forces. This means layers can slide over one another easily, which is why graphite feels slippery and is used as a lubricant. It also explains why graphite cleaves into flakes — you are separating layers, not breaking covalent bonds.

The delocalised electrons within each layer are free to move and carry charge, so graphite conducts electricity along the layers. This makes graphite and graphene the only non-metal allotropes of carbon that are good electrical conductors at room temperature, and graphite is used as electrodes in electrolysis cells.

2c. Graphene

Graphene is a single layer of graphite — one atom thick, arranged in a flat hexagonal lattice. You can think of it as the fundamental building block from which graphite is assembled by stacking many such layers.

Because it is a single layer, graphene has an exceptionally high surface area relative to its mass. The delocalised electrons are present across the whole sheet, giving graphene outstanding electrical conductivity. Its covalent bonds make it remarkably strong for its thickness — stronger than steel by weight.

Graphene is a nanomaterial with potential applications in electronics, composite materials, and filtration membranes. AQA expects you to know that it is a single layer of carbon atoms in a hexagonal arrangement and that it conducts electricity due to delocalised electrons.

2d. Fullerenes

Fullerenes are closed-cage molecules made entirely of carbon. The most well-known is buckminsterfullerene, $\text{C}_{60}$, which consists of 60 carbon atoms arranged in 20 hexagons and 12 pentagons — the same pattern as a football. Other fullerenes include $\text{C}_{70}$ and carbon nanotubes.

Each carbon in $\text{C}_{60}$ forms three covalent bonds, with the remaining electrons participating in delocalised pi bonding within the cage. However, unlike graphite or graphene, $\text{C}_{60}$ does not conduct electricity under normal conditions — the delocalisation is not sufficient to allow free movement of charge, and fullerenes behave as insulators or semiconductors. Unlike diamond or graphite, fullerenes are discrete molecules — they are not giant structures. This means the forces between $\text{C}_{60}$ molecules are only weak van der Waals forces, giving fullerenes relatively low melting points compared with diamond or graphite.

Carbon nanotubes are cylindrical fullerenes. Their structure gives them exceptional tensile strength along the tube axis and electrical conductivity along the tube, making them candidates for use in nanotechnology and composite materials.

3. Ionic Crystal Structures: Sodium Chloride and Magnesium Oxide

Ionic crystals form when metal atoms transfer electrons to non-metal atoms, producing oppositely charged ions that arrange themselves into a regular lattice. The electrostatic attraction between every positive ion and every surrounding negative ion — and vice versa — acts in all directions simultaneously. This is what makes ionic lattices so stable.

3a. Sodium Chloride (NaCl)

In the sodium chloride lattice, each $\text{Na}^+$ ion is surrounded by six $\text{Cl}^-$ ions, and each $\text{Cl}^-$ ion is surrounded by six $\text{Na}^+$ ions. This 6:6 coordination gives the structure its characteristic cubic symmetry. There are no discrete $\text{NaCl}$ molecules — the formula unit represents the simplest ratio of ions.

Sodium chloride has a high melting point (801 °C) because a large amount of energy is needed to overcome the many strong electrostatic attractions throughout the lattice. In the solid state it does not conduct electricity because the ions are fixed in position. When molten or dissolved in water, the ions become free to move and the substance conducts.

3b. Magnesium Oxide (MgO)

Magnesium oxide has the same 6:6 coordination structure as sodium chloride. However, its melting point (2852 °C) is far higher. The reason lies in the charges: $\text{Mg}^{2+}$ and $\text{O}^{2-}$ carry twice the charge of $\text{Na}^+$ and $\text{Cl}^-$. Greater ionic charge means a stronger electrostatic attraction between ions, so more energy is needed to break the lattice.

The ionic radius also matters: $\text{Mg}^{2+}$ is smaller than $\text{Na}^+$, so the ions in MgO are closer together, further increasing the lattice enthalpy. These two factors — higher charge and smaller ionic radius — combine to make MgO one of the highest-melting ionic compounds and a useful refractory material in furnace linings.

4. Metallic Structures

Metals consist of a regular lattice of positive ions (cations) surrounded by a 'sea' of delocalised electrons — electrons that have left their parent atoms and are free to move throughout the entire structure. This model, sometimes called the sea of electrons model, accounts for the characteristic properties of metals.

The electrostatic attraction between the positive ions and the surrounding delocalised electrons holds the structure together. Melting a metal requires enough energy to disrupt this attraction. Metals with higher ionic charge or smaller ionic radius tend to have stronger metallic bonding and therefore higher melting points — compare sodium (98 °C) with magnesium (650 °C), where the extra charge on $\text{Mg}^{2+}$ and the greater number of delocalised electrons per atom both strengthen the bonding.

The delocalised electrons explain two key properties simultaneously. First, they carry charge when a potential difference is applied, giving metals their electrical conductivity. Second, they carry kinetic energy through the lattice, giving metals their thermal conductivity. When a metal is deformed, the layers of ions can slide past one another while the electron sea maintains the bonding — this is why metals are malleable and ductile rather than brittle.

5. Molecular Crystal Structures: Iodine and Ice

Molecular crystals are built from discrete molecules held in a regular lattice by intermolecular forces. The key distinction from giant covalent or ionic structures is that the forces holding the lattice together are intermolecular — between molecules — rather than the strong intramolecular covalent or ionic bonds within or between ions.

5a. Iodine ($\text{I}_2$)

Solid iodine consists of $\text{I}_2$ molecules arranged in a regular lattice. Within each molecule, the two iodine atoms are held together by a strong covalent bond. Between molecules, the only forces are induced dipole–induced dipole interactions (London dispersion forces, a type of van der Waals force).

Because these intermolecular forces are relatively weak, iodine has a low melting point (114 °C) and has a tendency to sublime, especially when gently heated. It does not conduct electricity because there are no ions and no delocalised electrons — all electrons are localised within the covalent bonds or as lone pairs on the iodine atoms.

The van der Waals forces in iodine are stronger than those in, say, chlorine ($\text{Cl}_2$) because iodine molecules are larger and have more electrons, making them more polarisable. This is why iodine is a solid at room temperature while chlorine is a gas.

5b. Ice ($\text{H}_2\text{O}$)

Water molecules contain two O–H bonds and two lone pairs on the oxygen atom. The large electronegativity difference between oxygen and hydrogen makes each O–H bond highly polar, and the bent shape of the molecule means the dipoles do not cancel. This gives water a permanent dipole.

In ice, each water molecule forms up to four hydrogen bonds — two through its lone pairs acting as hydrogen bond acceptors, and two through its O–H bonds acting as hydrogen bond donors. This creates an open, tetrahedral network structure. The hydrogen bonds are directional and hold the molecules further apart than in liquid water, which is why ice is less dense than liquid water and floats.

Hydrogen bonds are significantly stronger than the van der Waals forces in iodine, which is why water has a much higher melting point (0 °C) than you would predict from its molar mass alone. However, hydrogen bonds are still far weaker than covalent or ionic bonds, so ice melts at a much lower temperature than any giant covalent or ionic solid.

6. Comparing All Four Crystal Types: A Structured Overview

Now that you have studied each type in detail, it is worth stepping back to see the full picture. The table below consolidates the key structural features and properties across all four crystal types. Use it as a revision anchor — but make sure you can explain each entry, not just recall it.

Worked example: Applying Structure to Predict Properties: A Worked Exam Question

Substance X has a melting point of 2072 °C, does not conduct electricity as a solid or when molten, and is insoluble in water. Identify the most likely crystal type for substance X and justify your answer. [3 marks]

1. Melting point = 2072 °C — this is very high. — Very high melting points are characteristic of giant covalent structures or ionic structures. Molecular crystals and most metals melt at far lower temperatures.
2. No conductivity as solid OR when molten. — Ionic substances conduct when molten (ions become mobile). Metals conduct in both states. The fact that X does not conduct even when molten rules out both ionic and metallic structures. This points strongly to a giant covalent structure.
3. Molecular crystals: ruled out by very high melting point. Ionic: ruled out by no conductivity when molten. Metallic: ruled out by no conductivity. Giant covalent: consistent with both observations. — Each piece of data eliminates one or more crystal types. Giant covalent is the only type consistent with all three observations.
4. Giant covalent structure [B1]. Very high melting point because strong covalent bonds throughout the lattice require large amounts of energy to break [B1]. No electrical conductivity in any state because there are no ions and no delocalised electrons [B1]. — Each marking point requires both the observation and the structural explanation. Stating 'giant covalent' alone scores only one mark.

Answer: Giant covalent structure. The very high melting point indicates strong covalent bonds throughout the lattice. The absence of electrical conductivity in both solid and molten states confirms there are no free ions or delocalised electrons.

7. Exam Technique: Scoring Full Marks on Structure Questions

AQA structure questions follow predictable patterns. The most common errors are not factual — they are presentational. Students know the right idea but phrase it in a way that does not match the mark scheme.

When asked to explain a property, you must link the structural feature to the property using causal language. 'Graphite conducts electricity because it has delocalised electrons that are free to move and carry charge' scores the mark. 'Graphite has free electrons' does not — it is incomplete because it does not state what those electrons do.

When asked to compare NaCl and MgO, you must address both substances and state the direction of the difference. 'MgO has a higher melting point than NaCl because the $\text{Mg}^{2+}$ and $\text{O}^{2-}$ ions have a greater charge, so the electrostatic attraction is stronger' is a complete comparative answer. Describing only one substance scores at most half the available marks.