Structure and bonding · GCSE Chemistry
Giant covalent structures
GCSE Chemistry revision on giant covalent structures: diamond, graphite, graphene and silicon dioxide, and why graphite conducts while diamond does not.
Diamond: four bonds, tetrahedral, hard, insulator. Graphite: three bonds, layers, delocalised electrons, conducts, layers slide. SiO₂: giant covalent oxide, high melting point, insulator.
The important bits
What you need to know
- 1
Giant covalent structures are networks of atoms joined by strong covalent bonds in all directions (or, in graphite, throughout each layer). They have very high melting points.
- 2
Diamond: each carbon atom is covalently bonded to four others in a tetrahedral lattice. It is very hard, has a very high melting point, and does not conduct (no delocalised electrons). Used in cutting tools.
- 3
Graphite: each carbon atom is bonded to three others in hexagonal layers. The fourth outer electron is delocalised, so graphite conducts electricity. Layers are held by weak forces and can slide — lubricant and pencil “lead”.
- 4
Graphene is a single layer of graphite: one atom thick, very strong for its mass, and it conducts because of delocalised electrons. Triple papers may also mention fullerenes (cages or tubes of carbon).
- 5
Silicon dioxide (silica, SiO₂) is a giant covalent structure similar to diamond, with each silicon bonded to four oxygen atoms. High melting point, hard, does not conduct. Sand and glass start here.
- 6
To melt diamond, graphite or SiO₂ you must break strong covalent bonds throughout the lattice, which takes a huge amount of energy — unlike boiling CO₂, which only separates simple molecules.
- 7
Graphite electrodes are used in electrolysis because graphite conducts and has a high melting point. Diamond cannot be used as an electrode: it does not conduct.
- 8
Allotropes of carbon are different structures of the same element. Same atoms, different bonding pattern, different properties. That is the comparison the mark scheme wants.
Quotations worth analysing
Short evidence. Real method.
“Diamond: four covalent bonds per carbon. Graphite: three covalent bonds per carbon, plus delocalised electrons.”
The extra bond in diamond explains hardness and insulation. The missing bond in graphite explains conduction and sliding layers. Do not mix those two explanations.
“C(diamond) and C(graphite) are allotropes of carbon.”
Allotropes are different structural forms of the same element in the same physical state. O₂ and O₃ are allotropes of oxygen; diamond and graphite are allotropes of carbon.
“SiO₂(s) is giant covalent; CO₂(g) is simple molecular.”
Both contain covalent bonds. The structure is different, so the melting points are wildly different. Never explain CO₂’s low boiling point by saying covalent bonds are weak.
Go deeper
Graphite’s two properties need two different sentences
Conductivity: each carbon uses three electrons in covalent bonds, so one electron per carbon is delocalised and can carry charge along the layers. Softness / lubricant: layers are held by weak intermolecular forces, so they slide over each other. Students write “graphite conducts because layers slide”. Sliding is not conduction. Diamond has no delocalised electrons because every outer electron is in a C–C bond, and it is hard because those bonds form a rigid 3D network. If a question asks for a use, attach it to the property: cutting tools for hardness; electrodes for conduction and high melting point; pencils for sliding layers that leave carbon on the paper.
Go deeper
Giant covalent versus simple molecular is the oxide trick
Carbon dioxide is a small molecule with two double bonds: O=C=O. Silicon dioxide is a giant lattice. Both are non-metal oxides with covalent bonding, but only SiO₂ has covalent bonds that must be broken to melt the substance. That is why sand does not boil away on a beach and why CO₂ is a gas at room temperature. A six-mark “compare diamond and graphite” or “compare CO₂ and SiO₂” is this page. Write structure, bonding, then one property and the particle reason. Repeat for the second substance. Do not wander into ionic bonding unless the question includes a salt.
Go deeper
Graphene and fullerenes are the same idea, scaled
Graphene is one hexagon sheet of graphite, so it still has delocalised electrons and covalent bonding in two dimensions. It is strong and a good conductor, which is why it appears in materials questions. Fullerenes such as C₆₀ are cages; nanotubes are rolled sheets. Large surface area and, in some cases, delocalised electrons lead to uses in catalysts, lubricants and electronics. On Combined Science you may only need diamond, graphite and silica. On Triple, name graphene or a fullerene and link a property to a use. Keep the carbon atom as the starting point: how many covalent bonds does each carbon make?
See the idea in action
Explain why graphite conducts electricity but diamond and silicon dioxide do not, even though all three have high melting points. All three are giant covalent, so strong covalent bonds must be broken to melt them. In graphite each carbon bonds to three others, leaving delocalised electrons that carry charge. In diamond each carbon bonds to four others, so there are no delocalised electrons. In SiO₂, silicon and oxygen atoms use their outer electrons in covalent bonds, so there are no mobile electrons or ions.
Exam technique
Turn knowledge into marks
For each allotrope, state the number of covalent bonds per carbon, then attach hardness or conduction separately. For SiO₂ versus CO₂, name giant covalent versus simple molecular before you mention melting point.
Common mistakes
Do not give these marks away
- 01
Explaining graphite’s conductivity with sliding layers, or diamond’s hardness with “strong intermolecular forces”.
- 02
Saying diamond conducts because it is carbon, or that graphite is a metal.
- 03
Treating CO₂ as giant covalent, or SiO₂ as a simple molecule like CO₂.
Why does graphite conduct electricity when diamond does not?
AGraphite has ionic bonds
BGraphite has delocalised electrons; in diamond every outer electron is in a covalent bond
CDiamond molecules are small and cannot move
DGraphite has a lower melting point so electrons are free
Show the answer
Graphite has delocalised electrons; in diamond every outer electron is in a covalent bond. Each carbon in graphite bonds to three others, leaving delocalised electrons. In diamond each carbon bonds to four others, so there is no mobile charge carrier.
Quick questions
If this is the bit you searched
Why is diamond hard GCSE Chemistry?
Each carbon atom is covalently bonded to four others in a rigid tetrahedral giant structure. Those strong bonds must be broken to scratch or melt diamond.
Why does graphite conduct electricity?
Each carbon is bonded to three others in layers, so one electron per carbon is delocalised and can carry charge. Sliding layers explain why graphite is soft, not why it conducts.
What is the structure of silicon dioxide?
SiO₂ is giant covalent: each silicon atom is covalently bonded to four oxygen atoms in a network with a very high melting point. It does not conduct.
What are allotropes of carbon?
Different structural forms of carbon in the same state, including diamond, graphite, graphene and fullerenes. They contain only carbon atoms arranged differently.