Giant Covalent Structures: Diamond & Graphite Revision Notes | A Level AQA Chemistry

1. The Four Types of Crystal Structure

Before zooming into diamond and graphite, you need a map of the territory. AQA groups all crystalline solids into four types: ionic, metallic, simple molecular (simple covalent), and giant covalent (macromolecular). Each type has a characteristic bonding pattern, and that bonding pattern directly controls every physical property — melting point, conductivity, hardness, and solubility.

Diamond and graphite are both forms of pure carbon — they are allotropes of carbon — yet their properties could hardly be more different. That contrast is the central puzzle of this lesson, and the answer lies entirely in how the atoms are arranged and bonded.

2. Diamond — Structure and Bonding

Picture a carbon atom at the centre of a tetrahedron. Each carbon atom uses its four outer electrons to form four covalent bonds (shared electron pairs) with four neighbouring carbon atoms. Each of those neighbours does exactly the same thing, and so on, extending in all three dimensions without interruption. The result is a rigid, three-dimensional lattice that constitutes the entire crystal — there are no separate molecules, no weak intermolecular forces, and no free electrons.

In diamond, every carbon atom forms four equivalent covalent bonds arranged tetrahedrally, giving a bond angle of 109.5°. This geometry repeats throughout the crystal, locking every atom into place. There are no lone pairs and no unused electrons — every electron is committed to a bond. (Note: sp³ hybridisation is extension knowledge beyond the core AQA specification for this topic. For exam purposes, focus on the tetrahedral geometry and 109.5° bond angle.)

When drawing a diamond structure diagram, you need to show a minimum of five carbon atoms to represent the tetrahedral environment of the central atom correctly — one central carbon bonded to four others. AQA mark schemes often specify the number of particles required, so always check the question wording carefully.

2a. Properties of Diamond — Explained by Structure

Diamond has an exceptionally high melting point (approximately 3550 °C). To melt diamond, you must break covalent bonds throughout the entire lattice — not just weak intermolecular forces. Because covalent bonds are strong (C–C bond enthalpy ≈ 347 kJ mol⁻¹) and there are an enormous number of them per mole of carbon, an enormous amount of energy is required.

Diamond does not conduct electricity. Every one of its four outer electrons per carbon atom is used in a covalent bond, leaving no delocalised electrons and no mobile charge carriers. This makes diamond an excellent electrical insulator — a fact that is directly testable in AQA exam questions.

Diamond is the hardest naturally occurring substance. The three-dimensional network of covalent bonds means that any attempt to scratch or deform the crystal requires breaking bonds in multiple directions simultaneously — there are no planes of weakness to exploit.

3. Graphite — Structure and Bonding

Graphite is where the story gets genuinely interesting. It is also a giant covalent structure, but the arrangement of atoms is completely different from diamond — and that difference produces a completely different set of properties.

In graphite, each carbon atom forms three covalent bonds to three neighbouring carbons, arranged in a flat, hexagonal pattern. This produces extended layers of hexagonal rings — often described as a honeycomb arrangement. Each carbon is sp² hybridised, with bond angles of 120° within the layers.

Here is the critical detail: each carbon has four outer electrons, but only three are used in covalent bonds within the layer. The fourth electron from each carbon atom enters a system of delocalised π electrons that spread across the entire layer, above and below the plane of the hexagonal rings. These delocalised electrons are the key to graphite's most distinctive properties.

The layers themselves are held together by weak van der Waals forces. These forces arise from temporary fluctuating dipoles in the electron clouds of the layers. Because van der Waals forces are much weaker than covalent bonds, the layers can slide over one another relatively easily — which is why graphite feels slippery and acts as a lubricant.

3a. Properties of Graphite — Explained by Structure

Graphite sublimes at approximately 3650 °C at atmospheric pressure (it does not have a conventional melting point at 1 atm). This very high sublimation point reflects the energy needed to break the strong covalent bonds within the layers — and there are many of them. The weak van der Waals forces between layers are overcome at much lower temperatures, but the overall sublimation point is still very high because the covalent network within each layer is extensive.

Graphite conducts electricity. The delocalised π electrons are free to move along the layers when a potential difference is applied. This makes graphite unusual among non-metals — it is a good electrical conductor, and this property is directly linked to the one electron per carbon that is not used in a localised covalent bond.

Graphite is soft and slippery because the layers can slide over one another. The weak van der Waals forces between layers are easily overcome, allowing one layer to move relative to another. This is why graphite is used in pencils and as a dry lubricant. Within each layer, however, the covalent bonds are very strong — so graphite is actually quite strong in the direction parallel to the layers.

4. Comparing Diamond and Graphite

Now that both structures are clear, the comparison becomes a powerful exam tool. AQA frequently asks you to compare diamond and graphite — which means you must address both similarities and differences explicitly. A common error is to describe only one structure and assume the examiner will infer the comparison.

Notice that both structures share a very high melting point — this is the similarity that catches students out. Both require breaking strong covalent bonds to melt, so both have high melting points. The difference in conductivity, however, is stark and arises entirely from whether a fourth electron per carbon is delocalised or locked into a bond.

5. Drawing Diagrams for Diamond and Graphite

AQA exam questions regularly ask you to draw diagrams of these structures showing a specified number of particles. This is a skill that requires practice, not just recognition. The key rules for each structure are straightforward once you know what to look for.

Worked example: Drawing a Diamond Structure Diagram

Draw a diagram to represent the structure of diamond, showing at least 8 carbon atoms and the bonds between them.

1. Draw one carbon atom (C) at the centre of your diagram. — This will be the atom whose tetrahedral environment you show first.
2. Draw four C atoms around the central one: one above, one below-left, one below-right, and one behind (shown as a dashed bond). Connect each with a single bond line. — The tetrahedral arrangement has bond angles of 109.5°. On a 2D diagram, use solid lines for bonds in the plane, a wedge for a bond coming towards you, and a dashed line for a bond going away.
3. From each of the four neighbours, draw three more bonds to new carbon atoms (the fourth bond of each neighbour is already accounted for by the bond back to the central atom). — This shows the repeating nature of the lattice. Each new carbon you add also has four bonds — you are showing a fragment of an infinite network.
4. Count the carbon atoms shown. Aim for at least 8. Verify every interior carbon has four bonds drawn. — AQA mark schemes award marks for showing the correct number of particles AND for showing each carbon with four bonds. Atoms at the edge of your diagram may have fewer bonds shown — that is acceptable as they represent the continuation of the lattice.
5. Label: 'C' at each atom, '109.5°' for one bond angle, and 'covalent bond' for one bond. — Labelling demonstrates understanding and picks up additional marks in extended questions.

Answer: A correct diamond diagram shows carbon atoms in a tetrahedral arrangement, each bonded to four others, with bond angles of 109.5°, extending in three dimensions. No layers, no free electrons, and no molecules are present.

Worked example: Drawing a Graphite Structure Diagram

Draw a diagram to represent the structure of graphite, showing the layered arrangement and at least 10 carbon atoms.

1. Draw a regular hexagon. Label each vertex as C. Each side of the hexagon represents a covalent bond. — The hexagonal ring is the repeating unit within a graphite layer. Bond angles within the ring are 120°.
2. Share one edge of the first hexagon with a new hexagon drawn adjacent to it. Add the four new carbon atoms needed. — Fused hexagons share a bond — this is how the honeycomb layer extends. Each shared bond belongs to both rings.
3. Continue adding fused hexagons until you have shown at least 10 carbon atoms in one layer. — A fragment of three or four fused hexagons is sufficient to show the layered structure clearly.
4. Draw a second identical layer below the first, but shifted horizontally so that the atoms of the second layer sit above the centres of the hexagons in the first layer (ABAB stacking). — The offset stacking is a feature of graphite's structure. The layers are not directly on top of one another.
5. Draw dashed lines between the layers to represent the weak van der Waals forces. Do NOT draw solid bond lines between layers. — Solid lines between layers would imply covalent bonds — which do not exist between graphite layers. Dashed lines correctly represent the weak van der Waals forces.
6. Label: '120°' for one bond angle within a layer, 'covalent bond' for one bond within a layer, 'van der Waals forces' for the interlayer interaction, and 'delocalised electrons' with an arrow indicating the π system above/below the layer. — These labels directly address the mark scheme points for structure and bonding questions.

Answer: A correct graphite diagram shows flat hexagonal layers of carbon atoms with 120° bond angles, layers stacked in an offset arrangement, dashed lines between layers representing van der Waals forces, and a label indicating the delocalised electron system within each layer.

6. Relating Structure to Melting Point and Conductivity

AQA exam questions on this topic almost always ask you to explain a physical property in terms of structure and bonding. The command word 'explain' requires you to state the observation and give the reason using a scientific principle. Saying 'diamond has a high melting point because it has strong bonds' earns one mark; saying 'diamond has a very high melting point because a large amount of energy is required to break the many strong covalent bonds that extend throughout the giant covalent lattice' earns full marks.

The general principle is this: melting point reflects the strength and number of bonds or forces that must be broken. For giant covalent structures, those are strong covalent bonds — so melting points are very high. For simple molecular structures, those are weak van der Waals or hydrogen bonds — so melting points are low. This is why you can melt ice at 0 °C but cannot melt diamond at the same temperature.

Electrical conductivity depends on the presence of mobile charge carriers — either delocalised electrons or mobile ions. Diamond has neither, so it does not conduct. Graphite has delocalised π electrons that can move along the layers, so it conducts. Ionic compounds conduct when molten or dissolved (mobile ions) but not when solid (ions fixed in lattice). Metals conduct because of their delocalised electron sea. Knowing this framework lets you answer conductivity questions for any crystal type.

The bar chart above makes the pattern unmistakable: giant covalent structures sit at the top of the melting point scale, far above simple molecular substances. When an exam question gives you melting point data and asks you to identify the type of structure, a very high melting point (above ~1000 °C) strongly suggests either ionic, metallic, or giant covalent — and you then use conductivity data to distinguish between them.

7. Exam Technique and Common Mistakes

This topic is one of the most reliably examined areas of the bonding unit. The questions are predictable, which means the marks are genuinely available — but only if you use precise language. The following guidance is drawn directly from the patterns seen in AQA mark schemes.

To consolidate everything: diamond and graphite are both giant covalent structures built from carbon atoms, but the number of covalent bonds each carbon forms — four in diamond, three in graphite — determines every difference in their properties. That single structural distinction explains hardness, conductivity, and the ability of layers to slide. Carry that principle into any exam question on this topic and you will have the reasoning framework to score full marks.