GIANT COVALENT STRUCTURES (also called macromolecular or network structures) are substances where a VERY LARGE NUMBER of atoms are all joined together by COVALENT BONDS throughout the entire structure.
Unlike simple molecules (discrete, small), giant covalent structures have no individual molecules — the covalent bonding extends in all directions throughout the solid.
Because the structure consists entirely of strong covalent bonds:
VERY HIGH melting and boiling points (must break many strong covalent bonds).
Very HARD — rigid structure.
Generally do NOT conduct electricity (no free electrons or ions).
Examples of giant covalent structures:
Diamond (carbon)
Graphite (carbon)
Silicon dioxide (SiO₂) — quartz/sand
Diamond
DIAMOND is a form (allotrope) of CARBON.
Structure:
Each carbon atom is bonded to FOUR other carbon atoms — in a TETRAHEDRAL arrangement.
Every carbon uses all 4 outer electrons for bonding.
The result is an extremely strong, rigid 3D lattice of covalent bonds.
No free electrons — all electrons are in bonds.
Properties and explanations:
VERY HIGH MELTING POINT: must break many strong C-C covalent bonds throughout the lattice.
VERY HARD — hardest natural substance on Earth: rigid 3D lattice, all atoms held firmly in all directions.
DOES NOT CONDUCT ELECTRICITY: all 4 outer electrons are used in covalent bonds — no free electrons to carry charge.
TRANSPARENT: interacts uniquely with light due to its bonding structure.
GRAPHITE is another allotrope of CARBON — same element, completely different structure and properties.
Structure:
Each carbon atom is bonded to THREE other carbon atoms — forming flat HEXAGONAL RINGS.
These hexagons form large flat SHEETS (layers).
The layers are held together only by WEAK intermolecular forces — they can slide over each other.
The FOURTH outer electron from each carbon is NOT used in bonding — these electrons are DELOCALISED between the layers.
Properties and explanations:
HIGH MELTING POINT: strong covalent bonds within each layer must be broken.
SLIPPERY / SOFT: layers can slide over each other easily (weak forces between layers).
CONDUCTS ELECTRICITY: delocalised electrons between layers are free to move and carry charge.
GREY/BLACK and OPAQUE: absorbs light.
USES: pencil leads (layers slide off onto paper), lubricants, electrodes in electrolysis (conducts, unreactive), electrical contacts.
⚠️ Common Mistake
GRAPHITE conducts electricity — it is the exception to the rule that giant covalent structures don't conduct. Diamond does NOT conduct. The difference: in graphite, each carbon uses only 3 of its 4 outer electrons for bonds — the 4th electron is delocalised and free to move. In diamond, all 4 electrons are used in bonds — none are free.
📌 Key Note
Giant covalent: many atoms all bonded by covalent bonds — very high MP, very hard. Diamond: each C bonds to 4 others, 3D lattice, does NOT conduct. Graphite: each C bonds to 3 others in layers, 4th electron delocalised — DOES conduct, layers slide (soft/lubricant). Both are carbon — different allotropes.
🎯 Matching Activity — Diamond or Graphite?
Sort each property into diamond or graphite. — drag the symbols on the right to match the component names on the left.
Diamond
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Graphite
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Graphite
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Diamond
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Graphite
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Diamond
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Used for cutting tools and as a gemstone — hardest natural substance
Conducts electricity — 4th outer electron is delocalised between layers
Does NOT conduct electricity — all 4 outer electrons are used in bonds
Each carbon bonded to 3 others in flat hexagonal sheets — layers can slide
Used as pencil lead and as electrodes in electrolysis
Each carbon bonded to 4 others in a 3D tetrahedral lattice — very hard
⭐ Higher Tier Only
Graphene: a single layer of graphite — one atom thick, conducts electricity (delocalised electrons), extremely strong. Fullerenes: C₆₀ (buckminsterfullerene) — hollow spherical cage, used in drug delivery and as lubricant. Carbon nanotubes: cylindrical graphene — very strong along axis, conduct electricity. All based on hexagonal carbon networks.
🎯 Test Yourself
Question 1 of 2
1. Why does graphite conduct electricity but diamond does not?
2. Why is graphite used as a lubricant?
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