What you'll learn
Giant covalent structures are substances in which every atom is joined to its neighbours by strong covalent bonds extending throughout the entire structure, with no separate molecules at any point. This topic sits within AQA GCSE Chemistry's bonding unit and is examined heavily, because it provides the clearest demonstration of the course's central principle: properties follow from structure and bonding. Diamond and graphite are made of identical atoms yet behave completely differently, and explaining why requires you to reason from bonding alone. By the end of this guide you should be able to define a giant covalent structure and distinguish it from a simple molecular one, describe and explain the structures and properties of diamond, graphite, graphene, fullerenes and silicon dioxide, account for the very high melting points of all giant covalent substances, explain why graphite conducts electricity while diamond does not, and identify a giant covalent substance from data.
Key terms and definitions
Giant covalent structure — a structure in which all the atoms are linked by strong covalent bonds throughout, also called a macromolecular structure
Macromolecule — a very large molecule; in giant covalent substances the whole structure is effectively one molecule
Covalent bond — a shared pair of electrons between two non-metal atoms
Delocalised electron — an electron that is not confined to one atom or bond and is free to move through the structure
Allotrope — a different structural form of the same element, such as diamond and graphite, which are both carbon
Fullerene — a molecule of carbon atoms arranged in a hollow shape, based on rings of carbon atoms
Nanotube — a cylindrical fullerene with a very high length-to-diameter ratio
Graphene — a single layer of graphite, one atom thick
Tensile strength — resistance to being pulled apart
Lubricant — a substance that reduces friction between surfaces
Core concepts
What makes a structure giant covalent
In a simple molecular substance such as methane or water, strong covalent bonds hold atoms together inside each molecule, but only weak intermolecular forces act between molecules. Melting or boiling overcomes those weak forces, which is why such substances have low melting and boiling points.
In a giant covalent structure there are no separate molecules and therefore no intermolecular forces at all. Every atom is covalently bonded to its neighbours in a continuous network. Melting the substance therefore requires breaking a very large number of strong covalent bonds.
This single difference explains why all giant covalent substances have very high melting and boiling points, and it is the explanation examiners want stated in those terms: many strong covalent bonds must be broken, and this requires a large amount of energy.
Giant covalent substances are also generally insoluble in water and, with the important exception of graphite, do not conduct electricity, because there are no free electrons or ions to carry charge.
Diamond
In diamond, each carbon atom forms four covalent bonds with four other carbon atoms, arranged in a rigid three-dimensional tetrahedral lattice.
Every one of carbon's four outer electrons is used in bonding, so no electrons are delocalised. Diamond therefore does not conduct electricity.
Because the whole structure is a rigid network of strong covalent bonds with no weak points, diamond is extremely hard — the hardest natural substance — and has a very high melting point of about 3,550 degrees Celsius. Its hardness explains its use in cutting tools, drill bits and abrasives.
Diamond does conduct thermal energy well, which is worth knowing because it is sometimes assumed that a non-conductor of electricity cannot conduct heat. The rigid lattice transmits vibrations efficiently.
Graphite
In graphite, each carbon atom forms only three covalent bonds, producing flat layers of hexagonal rings. Within each layer the bonding is strong, but there are no covalent bonds between the layers, which are held together only by weak intermolecular forces.
Two consequences follow, and both are examined constantly.
First, the layers can slide over one another because only weak forces resist the movement. Graphite is therefore soft and slippery, which is why it is used as a lubricant and in pencil leads, where layers are left behind on the paper.
Second, because each carbon atom uses only three of its four outer electrons for bonding, one electron per atom is delocalised and free to move through the layers. Graphite therefore conducts electricity and thermal energy, which is why it is used for electrodes in electrolysis. It is the only non-metal element you meet at GCSE that conducts electricity.
Graphite's melting point is also very high, around 3,600 degrees Celsius, because melting requires breaking the strong covalent bonds within the layers, not merely separating the layers.
Note the structural comparison that produces so many marks: diamond has four bonds per carbon and no delocalised electrons; graphite has three bonds per carbon, one delocalised electron per atom, and layers with no covalent bonds between them.
Graphene
Graphene is a single layer of graphite, one atom thick, consisting of carbon atoms in hexagonal rings.
Because each carbon still forms three covalent bonds, graphene has one delocalised electron per atom and conducts electricity extremely well. The strong covalent bonding within the layer gives it very high tensile strength relative to its mass, and it is almost transparent.
These properties make graphene a candidate for use in electronics, in touchscreens, and as a reinforcing additive in composite materials.
Fullerenes
Fullerenes are molecules of carbon atoms with hollow shapes. Their structures are based mainly on hexagonal rings of carbon atoms, but they may also contain rings with five or seven carbon atoms, which is what allows the sheet to curve into a closed shape.
The first fullerene to be discovered was Buckminsterfullerene, with the formula C60, which has a roughly spherical shape.
Fullerenes are being investigated for delivering drugs into the body, because a molecule can be enclosed within the hollow cage; as lubricants, because the spherical molecules roll; and as catalysts, because they have a very large surface area relative to their volume.
Carbon nanotubes are cylindrical fullerenes with a very high length-to-diameter ratio. They have very high tensile strength, conduct electricity and thermal energy well, and are used in nanotechnology, in electronics and to reinforce materials such as tennis rackets and building composites.
An important classification point: fullerenes are simple molecular substances, not giant covalent structures, because they consist of separate molecules with intermolecular forces between them. This is why C60 has a relatively low melting point compared with diamond or graphite, despite being made of carbon. Questions test this distinction.
Silicon dioxide
Silicon dioxide, also called silica, is the main component of sand and quartz. Each silicon atom is covalently bonded to four oxygen atoms and each oxygen atom to two silicon atoms, producing a giant three-dimensional lattice similar in arrangement to diamond.
It therefore shares diamond's properties: a very high melting point of about 1,710 degrees Celsius, hardness, insolubility in water, and no electrical conductivity because there are no delocalised electrons or free ions.
Its high melting point and chemical resistance explain its use in furnace linings and in glassmaking.
Identifying a giant covalent substance from data
Examination questions frequently supply a table of properties and ask you to classify a substance. The reasoning runs as follows.
A very high melting point, of several hundred degrees or more, indicates a giant structure — either giant covalent, giant ionic or metallic.
If the substance does not conduct electricity when solid or when molten, it is giant covalent rather than ionic, since molten ionic compounds conduct.
If it does not conduct when solid but is not ionic, and is insoluble in water, graphite is the exception to check for.
Metallic substances conduct when solid, which distinguishes them from giant covalent substances immediately.
Worked examples
Example 1: Explaining the difference in hardness (4 marks)
Diamond and graphite are both made only of carbon atoms. Explain why diamond is very hard but graphite is soft.
In diamond each carbon atom forms four covalent bonds in a rigid three-dimensional lattice, so there are strong covalent bonds in every direction and no weak points along which the structure can give way. A great deal of force is needed to break any of these bonds, so diamond is extremely hard.
In graphite each carbon atom forms only three covalent bonds, producing layers. There are no covalent bonds between the layers, which are held together only by weak intermolecular forces. These weak forces are easily overcome, so the layers slide over one another readily and graphite feels soft and slippery.
Example 2: Explaining conductivity (4 marks)
Explain why graphite conducts electricity but diamond and silicon dioxide do not.
In graphite each carbon atom uses only three of its four outer electrons to form covalent bonds, leaving one electron per atom delocalised. These delocalised electrons are free to move through the structure and can therefore carry charge, so graphite conducts electricity.
In diamond every carbon atom uses all four of its outer electrons in covalent bonds, and in silicon dioxide all the outer electrons of both silicon and oxygen are used in bonding. Neither structure contains delocalised electrons or free ions, so there are no charged particles free to move and neither substance conducts.
Example 3: Classifying from data (4 marks)
A substance has a melting point of 1,610 degrees Celsius, does not conduct electricity as a solid or when molten, and is insoluble in water. Deduce its structure and bonding, explaining your reasoning.
The very high melting point shows that a large amount of energy is needed to separate the particles, which indicates a giant structure rather than simple molecules.
Because the substance does not conduct when molten, it cannot be ionic: molten ionic compounds contain ions free to move and therefore conduct. Because it does not conduct as a solid, it cannot be metallic, since metals conduct in the solid state through delocalised electrons.
The substance must therefore have a giant covalent structure, in which all atoms are held by strong covalent bonds throughout with no delocalised electrons and no free ions. Its insolubility in water is consistent with this classification.
Common mistakes and how to avoid them
The most frequently penalised error is stating that covalent bonds are broken when graphite's layers slide apart. The layers are separated by overcoming weak intermolecular forces; the covalent bonds within each layer remain intact.
Students often say that graphite's high melting point is due to the forces between layers. It is due to the strong covalent bonds within the layers, which must be broken to melt the substance.
Another common error is calling fullerenes giant covalent structures. They are simple molecular, which is why C60 melts at a far lower temperature than diamond.
Many answers describe delocalised electrons as free electrons without using the word delocalised, or fail to say that they can move through the structure. Both elements are usually needed.
Finally, candidates sometimes claim diamond cannot conduct thermal energy because it cannot conduct electricity. Diamond is in fact an excellent thermal conductor.
Exam technique for "Giant covalent structures"
Every explanation in this topic follows the same three-step chain: identify the structure, describe the bonding and any delocalised electrons, then link that to the property being asked about. Writing the chain in order collects the marks even for an unfamiliar substance.
When comparing diamond and graphite, always state the number of bonds per carbon atom. Four against three is the root of every difference between them, and examiners look for the figures.
For melting point explanations, use the phrase many strong covalent bonds must be broken, which requires a large amount of energy. Vague references to strong bonding score less.
In data questions, work through conductivity in both the solid and molten states before deciding. That single comparison separates giant covalent from ionic and metallic more reliably than melting point alone.
Quick revision summary
A giant covalent structure has all its atoms joined by strong covalent bonds throughout, with no separate molecules and no intermolecular forces, so melting requires breaking many strong covalent bonds and melting points are very high. Diamond has four covalent bonds per carbon in a rigid three-dimensional lattice, making it extremely hard with no delocalised electrons, so it does not conduct electricity though it does conduct heat. Graphite has three bonds per carbon in layers with no covalent bonds between them, so the layers slide and it is soft and used as a lubricant, and one delocalised electron per atom lets it conduct electricity and be used for electrodes. Graphene is a single graphite layer, strong, conducting and nearly transparent. Fullerenes such as C60 and nanotubes are hollow carbon molecules, useful for drug delivery, lubricants, catalysts and reinforcement — but they are simple molecular, not giant covalent. Silicon dioxide forms a diamond-like lattice with a very high melting point and no conductivity. Identify giant covalent substances from a high melting point combined with no conductivity when solid or molten.