In giant covalent structures, millions of atoms are joined in a continuous lattice by strong covalent bonds. Diamond, graphite, and graphene are the three key examples in the GCSE specification. Because every bond must be broken to melt the solid, these substances have extremely high melting and boiling points and do not dissolve in water.
What is a giant covalent structure?
A giant covalent structure (also called a covalent network solid or macromolecular structure) is an arrangement in which a very large number of atoms are all bonded together by covalent bonds in a continuous, repeating three-dimensional (or two-dimensional) lattice. There are no separate molecules — the whole crystal is effectively one giant molecule.
Key properties that all giant covalent structures share:
- Very high melting and boiling points — a huge amount of energy is needed to break the many strong covalent bonds throughout the lattice
- Insoluble in water — there are no ions to be attracted away by water molecules; the covalent bonds are far too strong for water to break
- No free electrons (with an important exception — see graphite below) — so most do not conduct electricity
The three GCSE examples are all forms of carbon: diamond, graphite, and graphene. Silicon dioxide (SiO₂, found in sand and quartz) is another common example.
Why do giant covalent substances have such high melting points?
In a simple molecular substance (such as iodine or ice), melting only requires overcoming the weak intermolecular forces between molecules — the strong covalent bonds inside each molecule remain intact. This takes relatively little energy.
In a giant covalent structure, there are no separate molecules and no intermolecular forces. To melt the solid, you must break the actual covalent bonds throughout the entire lattice. Each bond has a bond enthalpy of several hundred kJ/mol, and there are an enormous number of them per mole of substance. The total energy required is immense, which is why diamond melts above 3,500 °C and graphite sublimes above 3,600 °C.
What is the structure of diamond and what are its properties?
In diamond, each carbon atom forms four covalent bonds to four neighbouring carbon atoms arranged in a tetrahedral shape. This creates a rigid, three-dimensional lattice that extends throughout the crystal.
| Property | Explanation |
|---|---|
| Hardest natural substance | Every C atom is bonded to four others; the lattice resists deformation in all directions |
| Very high melting point (>3,500 °C) | Enormous number of strong C–C covalent bonds to break |
| Does not conduct electricity | All four outer electrons are used in covalent bonds; no free electrons to carry charge |
| Transparent | No free electrons to absorb visible light |
| Does not dissolve in any common solvent | Covalent bonds are not attracted away by solvents |
Uses: cutting tools and drill bits (industrial diamonds), abrasives, gemstones, heat sinks in electronics.
What is the structure of graphite and why is it different from diamond?
In graphite, each carbon atom forms three covalent bonds to three neighbours, arranged in flat hexagonal rings. These rings extend into large flat sheets (layers). The fourth outer electron from each carbon atom is delocalised — free to move between the layers.
The layers themselves are held together only by weak intermolecular forces (van der Waals forces), which means they can slide over each other easily.
| Property | Explanation |
|---|---|
| Conducts electricity | One delocalised electron per C atom is free to move and carry charge |
| Conducts heat | Delocalised electrons also transfer kinetic energy efficiently |
| Soft and slippery | Layers held by weak forces only; can slide over each other |
| High melting point | Strong covalent bonds within layers are still many and strong |
| Grey/black, opaque | Delocalised electrons absorb visible light |
Uses: electrodes in electrolysis (conducts, withstands high temperatures), lubricant, pencil "lead" (layers slide off onto paper).
What is graphene and why is it remarkable?
Graphene is a single layer of graphite — one atom thick, arranged in a perfect hexagonal lattice. It was first isolated in 2004 by Andre Geim and Konstantin Novoselov at the University of Manchester, who received the 2010 Nobel Prize in Physics for this work.
Graphene's properties are extraordinary:
- Strongest material ever measured — about 200 times stronger than steel by mass
- Excellent electrical conductor — all electrons in graphene are delocalised in two dimensions; electron mobility is higher than in any other known material at room temperature
- Excellent thermal conductor
- Almost completely transparent — absorbs only ~2.3% of visible light
- Impermeable to all gases and liquids
- Extremely lightweight — a 1 m² sheet weighs about 0.77 milligrams
Potential uses: flexible transparent electrodes for touchscreens, ultra-strong composites, water filtration membranes, next-generation semiconductors.
Comparing diamond, graphite, and graphene
| Property | Diamond | Graphite | Graphene |
|---|---|---|---|
| Bonding | 3D lattice; each C bonded to 4 others | Layers; each C bonded to 3 others | Single sheet; each C bonded to 3 others |
| Free electrons | None | 1 per C atom (between layers) | 1 per C atom (in plane) |
| Electrical conductivity | None | Good (along layers) | Excellent |
| Hardness | Hardest natural substance | Very soft (layers slide) | Strongest material (tensile) |
| Appearance | Colourless/transparent | Grey/black opaque | Nearly transparent |
Frequently asked questions
Why does diamond not conduct electricity but graphite does?
In diamond, every carbon atom uses all four of its outer electrons to form four covalent bonds. There are no electrons left that are free to move through the structure, so diamond cannot conduct electricity. In graphite, each carbon atom forms only three covalent bonds, leaving one electron per atom delocalised — free to move between the layers throughout the structure. These mobile electrons can carry electric charge, so graphite conducts electricity. Graphene, being a single graphite layer, also has one delocalised electron per carbon atom and is an even better conductor.
Why can graphite be used as a lubricant but diamond cannot?
The layers of carbon in graphite are held together only by weak van der Waals forces (no covalent bonds between layers). When a shearing force is applied, the layers slide over each other easily, producing a low-friction surface — making graphite a useful dry lubricant. Diamond has strong covalent bonds in all directions (tetrahedral, 3D lattice) with no layers and no weak planes, so there is no easy direction for deformation. Instead of sliding, it resists mechanical stress in every direction, making it the hardest natural substance.
What is the difference between a giant covalent structure and a simple molecular structure?
In a simple molecular structure (such as water, carbon dioxide, or iodine), individual molecules are held together internally by strong covalent bonds, but the molecules are only weakly attracted to each other by intermolecular forces. Melting or boiling only breaks the weak intermolecular forces, not the covalent bonds, so simple molecular substances have low melting and boiling points. In a giant covalent structure, there are no separate molecules — the covalent bonds extend throughout the entire solid. Melting requires breaking actual covalent bonds, which needs enormous energy, explaining the very high melting and boiling points.
Why does silicon dioxide have a high melting point?
Silicon dioxide (SiO₂, found in quartz and sand) is a giant covalent structure in which each silicon atom is bonded to four oxygen atoms, and each oxygen atom bridges two silicon atoms, creating a three-dimensional network that extends throughout the whole crystal. Like diamond, there are no separate molecules and no weak intermolecular forces — only strong Si–O covalent bonds throughout. Breaking all these bonds to melt the solid requires enormous energy, giving SiO₂ a melting point of about 1,700 °C. This is why sand does not melt at normal temperatures, even though it feels quite insubstantial.
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