Giant covalent structure — a 3-D network
A giant covalent structure is a continuous 3-D lattice where every atom is joined to its neighbours by strong covalent bonds, with no separate molecules; the solids have very high melting points. The three examples are diamond (each carbon bonded to 4 carbons; non-conductor), graphite (each carbon bonded to 3 carbons plus one delocalised electron; conducts) and silicon(IV) oxide, SiO2 (each Si bonded to 4 O; non-conductor). Diamond and graphite are allotropes of carbon.
Diamond and SiO2 — rigid 3-D networks
Diamond: each carbon forms four covalent bonds in a tetrahedral 3-D network, so it is very hard, has a very high melting point (much energy breaks the many strong bonds) and does not conduct (all four outer electrons are bonded). Uses: cutting tools, drill bits, jewellery. Silicon(IV) oxide has an analogous network — each Si bonded to four O, each O bridging two Si — so it is hard, melts near 1700 °C and does not conduct. Uses: glass, optical fibres.
Graphite — a conducting layered solid
Graphite has each carbon bonded to three others in flat hexagonal layers held by weak forces; the layers slide, so it is soft, slippery and a lubricant. Its high melting point comes from the strong covalent bonds within each layer, not the weak interlayer forces. Each carbon's fourth outer electron is delocalised. (Extended) These move through the layers and carry charge, so graphite conducts. Uses: electrodes, pencils.
Drawn from real examiner reports.
(Extended) "Free electrons" is not enough for graphite
To explain graphite's conductivity, Cambridge requires the electrons to be delocalised and free to move, carrying charge. "Graphite has free electrons" alone gains no full credit. Full answer: each carbon uses only three of its four outer electrons in covalent bonds; the fourth is delocalised and free to move. The same precision applies to metallic bonding.
W22 Paper 41 Q2(a); S22 Paper 31 Q3(a)(i)
No intermolecular forces in diamond or SiO2
Diamond and silicon(IV) oxide are giant covalent, not molecular, so explaining their high melting points by "strong intermolecular forces" is wrong. The force broken on melting is the covalent bond itself, throughout the 3-D network, needing a very large amount of energy. Only graphite has weak forces, and only between its layers, never within a layer.
Diamond has 4 bonds, graphite has 3
The bonds per carbon distinguish the allotropes. Diamond: each carbon bonds to four others, a rigid 3-D network that is hard and non-conducting. Graphite: each carbon bonds to only three others in flat layers, leaving one delocalised electron, so it is soft and conducts. Claiming all four of graphite's electrons are bonded wrongly predicts no conductivity.
Graphite's high m.p. is not the weak forces
A common error is "graphite has a high melting point because its layers are held by weak forces". That is backwards: the weak inter-layer forces make graphite soft and slippery. Its high melting point comes from the strong covalent bonds within each layer, which take much energy to break. Weak forces explain softness; strong in-layer bonds explain the high melting point.
Silicon(IV) oxide is covalent, not ionic
Silicon(IV) oxide, SiO2, is giant covalent, not a giant ionic lattice, so it contains no ions. Describing it as ionic — or explaining its high melting point by "forces between ions" — is wrong. The bonds are covalent Si–O bonds throughout a 3-D network. It does not conduct because there are no free electrons and no mobile ions.
SiO2 is not an allotrope of carbon
Allotropes are different structural forms of the same element. Diamond and graphite are allotropes of carbon — one element arranged two ways. Silicon(IV) oxide is not an allotrope of carbon: it is a compound of silicon and oxygen. Do not call SiO2 an allotrope, and do not treat diamond and graphite as different compounds.
Structure to bonding to energy to property
Explain-the-property questions need a chain, not just the property: name the structure (giant covalent network), then the bonding (strong covalent bonds throughout), then the energy needed (to break the many bonds), then the property (very high melting point).
Match State/Describe against Explain
Read the command word first. "State" or "Describe" wants only the structural fact — each carbon bonded to four others in a 3-D network. "Explain" wants the full structure-bonding-energy-property chain; a bare fact there loses the reasoning marks.
Link each use to a property
"State a use" expects the use tied to the property behind it: diamond in cutting tools (very hard, rigid network); graphite as electrodes (conducts, withstands heat); graphite as a lubricant (layers slide). A use named without its property often loses the mark.
A giant covalent structure (also called a macromolecular structure) is a huge, continuous three-dimensional network of atoms joined by strong covalent bonds throughout. There are no separate molecules — the formula given (e.g. ) is only the simplest empirical ratio of atoms, not a molecular formula.
Key consequence of the giant structure: A very large amount of energy is needed to break the many strong covalent bonds → very high melting and boiling points for all giant covalent solids.
The three giant covalent substances required for Cambridge 0620:
| Substance | Formula | C bonds per atom | Conducts? | Key use |
|---|---|---|---|---|
| Diamond | C | 4 (to 4 C atoms) | No | Cutting tools, jewellery |
| Graphite | C | 3 (to 3 C atoms) + 1 delocalised | Yes | Electrodes, pencils, lubricant |
| Silicon(IV) oxide | 4 O per Si (each O bridges 2 Si) | No | Glass, semiconductors |
Define a giant covalent structure.
Draw a diagram to represent the layer structure of graphite.
Your diagram should show:
