Giant Covalent Structures

This lesson covers giant covalent structures — substances where millions of atoms are all bonded together by covalent bonds in a huge network — as required by the Edexcel GCSE Combined Science specification (1SC0). You need to know the structures and properties of diamond, graphite and silicon dioxide, and be able to explain their properties in terms of their bonding and structure.

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What Is a Giant Covalent Structure?

A giant covalent structure (also called a macromolecular structure) is a substance where all the atoms are linked by strong covalent bonds in a continuous network extending in all directions. There are no individual molecules — the entire structure is one giant lattice of atoms.

Examples include:

• Diamond (a form of carbon)
• Graphite (another form of carbon)
• Silicon dioxide (SiO₂, found in sand and quartz)

Exam Tip: The key difference between giant covalent structures and simple molecules is that giant covalent structures have no weak intermolecular forces to overcome — you must break strong covalent bonds to melt them. This is why they have very high melting points.

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Diamond

Diamond is a form of carbon in which every carbon atom is bonded to four other carbon atoms by strong covalent bonds arranged in a tetrahedral pattern.

Structure of Diamond

Feature	Detail
Each carbon bonded to	4 other carbon atoms
Bond type	Strong covalent bonds
Shape around each atom	Tetrahedral
Structure type	Giant covalent lattice
Free electrons	None — all 4 outer electrons are used in bonding

graph TD
    A["Diamond Structure"] --> B["Every C bonded to<br/>4 other C atoms"]
    A --> C["Tetrahedral<br/>arrangement"]
    A --> D["All bonds are<br/>strong covalent"]
    A --> E["No free electrons<br/>or ions"]

    style A fill:#2c3e50,color:#fff
    style B fill:#2980b9,color:#fff
    style C fill:#27ae60,color:#fff
    style D fill:#e67e22,color:#fff
    style E fill:#c0392b,color:#fff

Properties of Diamond

Property	Explanation
Very high melting point (3550 °C)	Many strong covalent bonds must be broken — requires a large amount of energy
Very hard	The rigid 3D network of strong covalent bonds in all directions makes it extremely hard
Does not conduct electricity	All outer electrons are used in covalent bonds — no free electrons to carry charge
Insoluble in water	The covalent bonds are too strong to be broken by water molecules

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Graphite

Graphite is another form of carbon (an allotrope of carbon). Its structure is very different from diamond.

Structure of Graphite

Feature	Detail
Each carbon bonded to	3 other carbon atoms
Bond type	Strong covalent bonds within layers
Structure	Flat hexagonal layers (sheets)
Between layers	Weak intermolecular forces (van der Waals forces)
Free electrons	Yes — one outer electron per carbon atom is delocalised

Properties of Graphite

Property	Explanation
Very high melting point (3652 °C)	Many strong covalent bonds within layers must be broken
Soft and slippery	Weak forces between layers allow them to slide over each other
Conducts electricity	Each carbon has one delocalised electron that is free to move along the layers and carry charge
Insoluble in water	The covalent bonds within layers are too strong to be broken by water

Exam Tip: Graphite is a common exam question because it has unusual properties for a giant covalent structure — it conducts electricity and is soft. Always explain these using the delocalised electrons (conductivity) and weak forces between layers (softness).

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Silicon Dioxide (SiO₂)

Silicon dioxide (also called silica) is found naturally as sand and quartz. It has a giant covalent structure similar to diamond.

Structure of Silicon Dioxide

• Each silicon atom is bonded to 4 oxygen atoms.
• Each oxygen atom is bonded to 2 silicon atoms.
• This creates a rigid 3D network of covalent bonds.

Properties of Silicon Dioxide

Property	Explanation
Very high melting point (1710 °C)	Many strong covalent bonds must be broken
Very hard	Rigid 3D covalent network
Does not conduct electricity	No free electrons or delocalised electrons
Insoluble in water	Strong covalent bonds cannot be broken by water

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Comparing Diamond, Graphite and Silicon Dioxide

Property	Diamond	Graphite	Silicon Dioxide
Melting point	Very high (3550 °C)	Very high (3652 °C)	Very high (1710 °C)
Hardness	Very hard	Soft and slippery	Very hard
Electrical conductivity	Does not conduct	Conducts	Does not conduct
Bonding	Each C bonded to 4 C	Each C bonded to 3 C in layers	Each Si bonded to 4 O
Free electrons	None	Yes (delocalised)	None

Exam Tip: You may be asked to compare diamond and graphite. Both are forms of carbon with high melting points due to strong covalent bonds. The key differences are: (1) graphite has layers that can slide (soft), (2) graphite has delocalised electrons (conducts electricity).

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Summary

• Giant covalent structures have many atoms bonded by strong covalent bonds in a large network — no individual molecules.
• Diamond: each C bonded to 4 others, very hard, very high MP, does not conduct.
• Graphite: each C bonded to 3 others in layers, soft and slippery, very high MP, conducts electricity due to delocalised electrons.
• Silicon dioxide: Si and O in a 3D network, very hard, very high MP, does not conduct.
• All three are insoluble in water and have very high melting points because many strong covalent bonds must be broken.

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Worked Example — Comparing Diamond and Graphite

Diamond and graphite are both allotropes of carbon — the same element arranged differently. Yet their properties are almost opposite.