What you'll learn
Giant covalent structures are huge networks of atoms held together by strong covalent bonds, with no overall charge and no separate molecules. For AQA GCSE Chemistry you need to know three key examples — diamond, graphite, and silicon dioxide — and be able to explain how their structure gives them their properties, such as very high melting points and, in the case of graphite, the ability to conduct electricity. This guide explains what a giant covalent structure is, describes each example in detail, and links each property back to the bonding and arrangement of atoms. By the end you should be able to explain why diamond is hard, why graphite is soft and conducts, and why these substances have such high melting points.
Key terms and definitions
Covalent bond — A strong bond formed when two atoms share a pair of electrons.
Giant covalent structure — A structure containing a very large number of atoms joined by covalent bonds in a continuous network (also called a macromolecule).
Macromolecule — Another name for a giant covalent structure; "macro" means large.
Delocalised electron — An electron that is free to move through a structure and is not held between two particular atoms.
Allotrope — Different structural forms of the same element; diamond and graphite are allotropes of carbon.
Layer — A single sheet of covalently bonded atoms, as found in graphite.
Lattice — A regular, repeating arrangement of particles.
Core concepts
What makes a structure "giant covalent"
In a simple molecular substance such as water or carbon dioxide, atoms are joined by covalent bonds into small, separate molecules. In a giant covalent structure, the covalent bonds continue throughout the whole structure, joining millions of atoms into one enormous network. There are no separate molecules. Because you have to break many strong covalent bonds to melt or boil these substances, they all have very high melting and boiling points and are solids at room temperature.
Diamond
Diamond is a form of carbon. Each carbon atom forms four covalent bonds to four other carbon atoms, arranged in a rigid three-dimensional lattice. This makes diamond extremely hard and gives it a very high melting point, because a huge number of strong covalent bonds must be broken to separate the atoms. Every one of a carbon atom's four outer electrons is used in bonding, so there are no free electrons — this means diamond does not conduct electricity. Its hardness makes diamond useful in cutting tools.
Graphite
Graphite is another form of carbon, but with a completely different arrangement. Each carbon atom forms only three covalent bonds to three other carbon atoms, forming flat layers of hexagons. The fourth outer electron of each carbon atom is not used in bonding — it becomes delocalised and can move through the structure. This gives graphite two special properties: it conducts electricity (because the delocalised electrons can carry charge), and it is soft and slippery (because there are only weak forces between the layers, allowing them to slide over each other). Graphite is used in pencils and as a lubricant, and as electrodes because it conducts.
Silicon dioxide
Silicon dioxide (silica), found in sand and quartz, is a giant covalent structure in which each silicon atom is bonded to four oxygen atoms and each oxygen atom to two silicon atoms, forming a strong three-dimensional network. Like diamond, it has a very high melting point and is hard, because many strong covalent bonds must be broken to melt it. It does not conduct electricity because it has no delocalised electrons or free ions.
Graphene and fullerenes
Closely related to graphite are two other carbon structures you should recognise. Graphene is a single layer of graphite — one sheet of carbon atoms in hexagons, just one atom thick. It conducts electricity because of its delocalised electrons and is extremely strong for its mass, which makes it useful in electronics and composite materials. Fullerenes are molecules of carbon shaped like hollow balls or tubes; the first discovered was buckminsterfullerene, C₆₀, shaped like a football. Carbon nanotubes are cylindrical fullerenes that are very strong and conduct electricity, used to reinforce materials and in electronics. These are not giant covalent structures in the same way — graphene is a single giant layer, while fullerenes are large molecules — but they share the strong carbon–carbon covalent bonding.
Comparing the giant covalent structures
| Substance | Bonds per atom | Conducts? | Property | Use |
|---|---|---|---|---|
| Diamond | 4 (carbon) | No | Very hard | Cutting tools |
| Graphite | 3 (carbon) | Yes | Soft, slippery | Pencils, lubricant, electrodes |
| Silicon dioxide | 4 (silicon) | No | Hard, high melting point | Glass, sand |
Notice how the number of bonds per atom controls the properties: four bonds give a rigid, hard, non-conducting solid, while three bonds leave a spare electron that both allows conduction and, with weak forces between layers, makes the material soft.
Worked examples
Example 1: Explaining a high melting point
Why does diamond have a very high melting point? Diamond is a giant covalent structure in which every carbon atom is joined to four others by strong covalent bonds. To melt it, a very large number of these strong covalent bonds must be broken, which requires a great deal of energy, so the melting point is very high.
Example 2: Why graphite conducts but diamond does not
Explain why graphite conducts electricity but diamond does not. In graphite, each carbon atom forms only three bonds, leaving one delocalised electron per atom that is free to move and carry charge. In diamond, all four outer electrons of each carbon atom are used in bonding, so there are no free electrons to carry charge, and it does not conduct.
Example 3: Why graphite is soft
Explain, in terms of structure, why graphite is soft and can be used as a lubricant. Graphite is made of layers held together only by weak forces between the layers. These layers can slide over one another easily, which makes graphite soft and slippery, so it works well as a lubricant and in pencils.
Example 4: Identifying a structure
A solid has a very high melting point, is very hard, and does not conduct electricity. Is it likely to be giant covalent or simple molecular, and why? It is likely to be giant covalent. Simple molecular substances have low melting points because only weak intermolecular forces need breaking. A very high melting point and great hardness point to many strong covalent bonds in a giant network, and the lack of conduction rules out delocalised electrons (as in graphite) or free ions.
Common mistakes and how to avoid them
A very common mistake is saying that melting diamond or graphite "breaks the covalent bonds because they are weak". Covalent bonds are strong — that is exactly why the melting point is high. It is the large number of strong bonds that must be broken.
Students often forget that graphite conducts because of delocalised electrons, not free ions. Graphite is a form of an element, so it has no ions; the conduction comes from the one delocalised electron per carbon atom.
Another error is confusing the number of bonds: diamond has four bonds per carbon, graphite has three. Getting this wrong makes the explanations of hardness and conductivity fall apart.
When explaining why graphite is soft, do not say the covalent bonds within the layers are weak — they are strong. It is the forces between the layers that are weak, letting the layers slide.
Finally, remember that these substances have no separate molecules. Writing that giant covalent structures "have weak forces between molecules" is wrong — there are no molecules to have forces between.
Exam technique for "Giant covalent structures"
Exam answers on giant covalent structures almost always need you to link structure to property. Build your answer as a chain: state the structure (giant covalent, number of bonds per atom), then the consequence (many strong bonds, or delocalised electrons, or weak forces between layers), then the property.
For "explain the high melting point" questions, the key marking points are: giant covalent structure, strong covalent bonds, and a large amount of energy needed to break many bonds. For graphite, always mention the delocalised electron and the weak forces between layers separately, because they explain two different properties.
Diagrams can help you think, but the marks come from the written explanation, so practise writing these explanations in full sentences. Be ready to compare diamond and graphite side by side, as this is a favourite six-mark question.
Quick revision summary
- Giant covalent structures are huge networks of atoms joined by strong covalent bonds, with no separate molecules and very high melting points.
- Diamond: each carbon forms 4 bonds; very hard, high melting point, does not conduct (no free electrons).
- Graphite: each carbon forms 3 bonds, with one delocalised electron per atom; conducts electricity, soft and slippery because layers slide.
- Silicon dioxide: each silicon bonds to 4 oxygens; hard, high melting point, does not conduct.
- High melting points come from breaking many strong covalent bonds, not weak ones.
- Graphite conducts because of delocalised electrons, not ions.