Giant Covalent Structures

Theory · read, then commit

Giant covalent structures (also called macromolecular) are substances where a huge number of atoms are all joined by covalent bonds throughout the whole structure — there are no separate molecules. Examples: diamond, graphite and silicon dioxide (SiO₂).

Checkpoint · quick check

Describe what a giant covalent structure is.

Properties of giant covalent structures

🌡️Very high melting point

Melting means breaking millions of strong covalent bonds, which takes a lot of energy.

💎Very hard & rigid

A rigid 3-D network of strong bonds holds every atom firmly in place.

🚫Usually don't conduct

Most have no free electrons or ions to carry charge — graphite is the exception.

💎Diamond

Bonds per carbon4 — rigid 3-D tetrahedral lattice

Free electronsNone (all 4 used in bonds)

Conducts electricity?No

HardnessHardest natural substance

✏️Graphite

Bonds per carbon3 — flat hexagonal layers

Free electrons1 delocalised per carbon

Conducts electricity?Yes

HardnessSoft & slippery (layers slide)

Both are pure carbon. The spare 4th electron and the sliding layers are why graphite conducts and is soft, while diamond does neither.

Checkpoint · quick check

State why diamond and graphite can both be made of carbon yet have very different properties.

⚠️ Common Mistake

Students often assume that because diamond and graphite are both giant covalent structures made of carbon, they must have the same properties. They do not: graphite conducts electricity and is soft and slippery, while diamond does not conduct and is extremely hard. The difference is that in graphite each carbon bonds to only three others, leaving one delocalised electron per atom and weak forces between layers; in diamond every carbon bonds to four others with no free electrons.

Other forms of carbon

▦Graphene

A single layer of graphite, one atom thick — extremely strong, very light, and conducts electricity.

⚽Buckminsterfullerene (C₆₀)

60 carbon atoms in a hollow sphere ('buckyball') — can cage molecules; used in drug delivery and lubricants.

🛢️Carbon nanotube

A rolled-up graphene sheet forming a hollow tube — very strong along its length and conducts electricity.

🛠️Uses

Diamond: cutting tools, drill bits, gemstones, abrasives (very hard).

Graphite: pencil leads and lubricants (layers slide); electrodes in electrolysis (conducts, unreactive).

Retrieval · sort the properties

Test yourself · 8 questions, exam order

0 of 8 answered

Warm-up · 3 questions

Explain why graphite conducts electricity but diamond does not.

Explain why graphite is used as a lubricant.

Explain why diamond is very hard and has a very high melting point.

Exam standard · 3 questions

Explain why silicon dioxide (found in sand and quartz) has a very high melting point.

Graphene is a single layer of graphite. Explain why it conducts electricity and is very strong.

Explain why buckminsterfullerene (C₆₀) can be used to deliver drugs to specific parts of the body.

Stretch · 2 questions

Explain why graphite has a high melting point even though its layers can slide over each other easily.

Compare diamond and graphite in terms of the number of bonds each carbon forms and whether they conduct electricity.

Key note · cover it, say it, check it

Revision card · 4 steps Giant Covalent Structures

01Giant covalent: many atoms all bonded by covalent bonds — very high MP, very hard.
02Diamond: each C bonds to 4 others, 3D lattice, does NOT conduct.
03Graphite: each C bonds to 3 others in layers, 4th electron delocalised — DOES conduct, layers slide (soft/lubricant).
04Both are carbon — different allotropes.

Exam ladder · Giant Covalent Structures

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