Giant Covalent Structures

This lesson covers giant covalent structures — substances where millions of atoms are bonded together by covalent bonds in a huge network. You need to know the structures and properties of diamond, graphite and silicon dioxide for the Edexcel GCSE Chemistry (1CH0) exam, and be able to explain their properties in terms of structure and bonding.

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What Are Giant Covalent Structures?

Giant covalent structures (also called macromolecular structures) are substances in which:

• All the atoms are bonded to each other by strong covalent bonds.
• The structure extends throughout the entire material — there are no individual molecules.
• Millions (or billions) of atoms are joined in a continuous network.

Because all the bonds are strong covalent bonds, giant covalent structures have very high melting and boiling points — a large amount of energy is required to break the many covalent bonds.

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Diamond

Diamond is an allotrope of carbon (a different structural form of the element carbon).

Structure

• Each carbon atom is bonded to 4 other carbon atoms by strong covalent bonds.
• The bonds are arranged in a tetrahedral shape (each bond is at 109.5° to the others).
• This creates a rigid, three-dimensional giant covalent lattice.
• There are no free electrons (all four outer electrons of each carbon are used in bonding).

Properties

Property	Value / Description	Explanation
Melting point	Very high (approx. 3550 °C)	Very many strong covalent bonds must be broken — requires a lot of energy
Hardness	Very hard (hardest natural substance)	Rigid 3D lattice with strong bonds in all directions
Electrical conductivity	Does not conduct	All four outer electrons are involved in covalent bonds — no free electrons (no delocalised electrons)
Solubility	Insoluble in water	The covalent bonds are too strong to be broken by water molecules

Exam Tip: When explaining why diamond is hard, say "each carbon is bonded to four others in a rigid tetrahedral arrangement by strong covalent bonds." When explaining why it doesn't conduct, say "there are no free/delocalised electrons."

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Graphite

Graphite is another allotrope of carbon.

Structure

• Each carbon atom is bonded to 3 other carbon atoms by strong covalent bonds.
• This forms flat layers (sheets) of hexagonally arranged carbon atoms.
• There are weak intermolecular forces (van der Waals forces) between the layers.
• Each carbon atom has 1 free (delocalised) electron — the fourth outer electron that is not involved in bonding.
• These delocalised electrons can move along the layers.

Properties

Property	Value / Description	Explanation
Melting point	Very high (approx. 3600 °C)	Very many strong covalent bonds within the layers must be broken
Hardness	Soft and slippery	Weak forces between layers allow layers to slide over each other
Electrical conductivity	Does conduct electricity	Delocalised electrons are free to move along the layers, carrying charge
Use as lubricant	Yes	Layers slide easily — reduces friction
Use in pencils	Yes (graphite mixed with clay)	Layers slide off onto paper

Exam Tip: The most common mistake is to say graphite conducts because "ions are free to move." Graphite has no ions — it conducts because of delocalised electrons that can move along the layers.

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Diamond vs Graphite — Comparison

graph TD
    subgraph Diamond
        D1["Each C bonded to 4 others"]
        D2["Tetrahedral, rigid 3D lattice"]
        D3["Very hard"]
        D4["Does NOT conduct electricity"]
        D5["No delocalised electrons"]
        D1 --> D2 --> D3
        D1 --> D5 --> D4
    end
    subgraph Graphite
        G1["Each C bonded to 3 others"]
        G2["Flat layers of hexagons"]
        G3["Soft and slippery"]
        G4["DOES conduct electricity"]
        G5["1 delocalised electron per C"]
        G1 --> G2 --> G3
        G1 --> G5 --> G4
    end

Property	Diamond	Graphite
Bonds per carbon atom	4	3
Shape	Tetrahedral 3D lattice	Flat hexagonal layers
Hardness	Very hard	Soft, slippery
Electrical conductivity	No	Yes
Delocalised electrons?	No	Yes (1 per carbon)
Melting point	Very high (~3550 °C)	Very high (~3600 °C)

Exam Tip: Both diamond and graphite have very high melting points because both contain strong covalent bonds. The difference in other properties arises from the different structures (3D vs layered).

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Silicon Dioxide (SiO₂)

Silicon dioxide (also called silica) has a structure similar to diamond.

Structure

• Each silicon atom is covalently bonded to 4 oxygen atoms.
• Each oxygen atom is bonded to 2 silicon atoms.
• This forms a giant covalent lattice extending in all directions.

Properties

Property	Description	Explanation
Melting point	Very high (~1710 °C)	Many strong covalent bonds must be broken
Hardness	Very hard	Rigid 3D covalent network
Electrical conductivity	Does not conduct	No free electrons or ions

Silicon dioxide is found naturally as quartz and sand.

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Why Giant Covalent Structures Have High Melting Points

To melt a giant covalent structure, you must break the strong covalent bonds throughout the entire structure. Because:

1. There are very many covalent bonds.
2. Each covalent bond is strong (high bond energy).
3. A very large amount of energy is needed to overcome all of these bonds.

This is different from simple molecular substances where only weak intermolecular forces need to be overcome.

Structure Type	What Is Broken on Melting	Energy Needed	Melting Point
Simple molecular	Weak intermolecular forces	Low	Low
Giant covalent	Strong covalent bonds	Very high	Very high

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Key Points

• Giant covalent structures have very high melting points because many strong covalent bonds must be broken.
• Diamond: each C bonded to 4 others, tetrahedral, very hard, does not conduct (no free electrons).
• Graphite: each C bonded to 3 others in layers, soft/slippery (layers slide), conducts electricity (delocalised electrons).
• Silicon dioxide: similar to diamond, very hard, very high melting point, does not conduct.
• In giant covalent structures, covalent bonds are broken during melting (unlike simple molecules where only intermolecular forces are overcome).

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Practice Questions

1. Explain why diamond has a very high melting point.
2. Explain why graphite can conduct electricity but diamond cannot.
3. Explain why graphite is used as a lubricant.
4. Both diamond and graphite are forms of carbon. Explain why they have such different properties.
5. Silicon dioxide has a very high melting point. Explain why, in terms of its structure and bonding.

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Worked Example: Why Diamond Is Hard but Graphite Is Soft

Question (6 marks): Diamond and graphite are both pure carbon but have very different physical properties. Diamond is extremely hard, while graphite is soft and slippery. Explain this difference in terms of their structures and bonding.

Answer: