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Most covalent substances are made of small, separate molecules — water, carbon dioxide, sugar — and they melt or boil at low temperatures. But a few elements and compounds form giant covalent structures (also called macromolecular or network covalent structures), in which covalent bonds link every atom to its neighbours throughout the whole solid. There are no separate molecules at all: a diamond is, in a sense, one enormous molecule. This article explains the main examples — diamond, graphite, graphene and silicon dioxide — and how their structures explain their remarkable properties.
What makes a structure “giant”?
In a simple molecular substance, strong covalent bonds hold atoms together within each molecule, but only weak intermolecular forces hold the molecules to each other (see simple molecular vs giant covalent substances).
In a giant covalent structure, strong covalent bonds extend in all directions (or across whole layers) through the entire solid. To melt it, you must break a huge number of strong covalent bonds, not just weak attractions. That’s why these substances have very high melting points.
The elements that form giant covalent structures are mainly in the middle of the periodic table — especially carbon, silicon and boron — which can form three or four strong bonds each.
Diamond
In diamond, every carbon atom (see carbon) forms four covalent bonds to four other carbon atoms, arranged tetrahedrally at 109.5° (each carbon is sp³ hybridised — see hybridisation explained). This creates a rigid 3D network extending in every direction. The C–C bond length is 154 pm.
Properties and explanations:
| Property | Explanation |
|---|---|
| Extremely hard — the hardest natural material (10 on the Mohs scale) | Every atom is locked in place by four strong bonds in a 3D network; there are no weak planes |
| Very high melting point (above about 3,500 °C; it tends to convert to graphite or burn first) | Melting requires breaking many strong C–C bonds |
| Does not conduct electricity | All four outer electrons of each carbon are in localised bonds; no free electrons or ions |
| Excellent conductor of heat — better than copper at room temperature | Vibrations pass efficiently through the stiff, light, strongly bonded lattice |
| Insoluble in all solvents | No solvent can overcome the covalent network |
| Transparent, high refractive index | Strong bonding and no free electrons to absorb visible light; light bends strongly, causing sparkle |
Uses: jewellery; cutting, drilling and grinding tools (diamond-tipped saws and drills); heat sinks for electronics.
Graphite
In graphite, each carbon forms three covalent bonds to three other carbons, in flat hexagonal layers (each carbon is sp² hybridised, with bond angles of 120°). The fourth outer electron of each carbon is in a p orbital perpendicular to the layer, and these join up into a delocalised π system across the whole layer (see sigma and pi bonds).
The layers are held together only by weak London forces, and they’re 335 pm apart — more than twice the C–C distance within a layer (142 pm).
Properties and explanations:
| Property | Explanation |
|---|---|
| Very high melting point (sublimes around 3,600 °C) | Strong covalent bonds within layers must be broken |
| Soft and slippery | Weak forces between layers let them slide over each other |
| Conducts electricity (along the layers) | Delocalised electrons move freely within the layers |
| Less dense than diamond (2.2 vs 3.5 g cm⁻³) | Large gaps between layers |
| Grey-black, shiny | Delocalised electrons absorb and re-emit light across the visible range |
Uses: pencil “leads” (mixed with clay — layers rub off onto paper), lubricants, electrodes in electrolysis and batteries (lithium-ion battery anodes are mostly graphite), and in carbon-fibre composites.
Interestingly, graphite, not diamond, is the more stable form of carbon at room temperature and pressure. Diamond converts to graphite only extremely slowly under normal conditions, because breaking diamond’s network takes so much energy. “Diamonds are forever” is really a statement about kinetics, not thermodynamics.
Graphene
Graphene is a single layer of graphite — one atom thick. It was first isolated in 2004 by Andre Geim and Konstantin Novoselov at the University of Manchester, famously by peeling layers off graphite with sticky tape. They received the 2010 Nobel Prize in Physics.
Graphene has extraordinary properties: it’s extremely strong for its thickness, very flexible, almost transparent, and conducts electricity and heat exceptionally well, thanks to its delocalised electrons. Researchers are exploring uses in electronics, sensors, composite materials and water filtration.
Related carbon forms include carbon nanotubes (graphene rolled into tubes) and fullerenes such as C₆₀ (“buckyballs”), which are actually simple molecular structures — separate molecules — rather than giant networks (see allotropes of carbon).
Silicon dioxide (silica)
Silicon dioxide, SiO₂, is the main compound in sand, quartz and flint. Each silicon atom is bonded tetrahedrally to four oxygen atoms, and each oxygen is bonded to two silicon atoms. This makes a 3D network similar in principle to diamond, with an oxygen bridging each pair of silicons.
The formula SiO₂ shows the ratio of atoms (one Si to two O), not a molecule — there are no separate SiO₂ molecules.
Properties: very hard (quartz is 7 on the Mohs scale), high melting point (about 1,700 °C), does not conduct electricity, insoluble in water.
Uses: glass making (sand is melted with sodium carbonate and limestone), concrete and construction, optical fibres, and quartz crystals in watches and electronics.
Why is CO₂ a gas but SiO₂ a solid?
Carbon and silicon are in the same group, so why are their dioxides so different? Carbon is small enough to form strong C=O double bonds (π bonds), so each carbon satisfies its bonding with just two oxygens, forming small, separate O=C=O molecules held together by weak forces. Silicon is larger; its p orbitals overlap poorly sideways, so it doesn’t form good Si=O double bonds. Instead, it forms four Si–O single bonds, linking into a giant network. A small difference in atom size leads to a gas in one case and rock in the other.
Other giant covalent substances
- Silicon itself has the diamond structure; it’s the basis of computer chips.
- Silicon carbide (SiC, carborundum) — very hard, used in abrasives and high-power electronics.
- Boron nitride — can form a soft, graphite-like structure or a very hard, diamond-like one.
Summary table
| Substance | Bonds per atom | Structure | Conducts? | Hardness |
|---|---|---|---|---|
| Diamond | C: 4 | 3D tetrahedral network | No | Extremely hard |
| Graphite | C: 3 | Layers of hexagons | Yes (along layers) | Soft, slippery |
| Graphene | C: 3 | Single layer | Yes, excellent | Very strong in-plane |
| Silicon dioxide | Si: 4, O: 2 | 3D tetrahedral network | No | Hard |
Common misconceptions
- “Melting graphite overcomes the weak forces between layers.” Its very high melting point shows that strong covalent bonds within layers must break.
- “Graphite conducts because it’s a metal.” It’s a non-metal; it conducts because of delocalised electrons.
- “SiO₂ is a molecule like CO₂.” It’s a giant network; the formula is a ratio.
- “Diamond is the most stable form of carbon.” Graphite is, at normal conditions.
Key takeaways
- In giant covalent structures, strong covalent bonds extend throughout the solid; there are no separate molecules.
- Diamond: 4 bonds per carbon, 3D network → extremely hard, high melting point, non-conductor.
- Graphite: 3 bonds per carbon, layers with delocalised electrons and weak forces between layers → conducts, soft, slippery.
- Silicon dioxide: SiO₄ tetrahedra linked in 3D → hard, high melting point.
- Structure explains properties and uses.
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