Giant Covalent Structures
A concise revision guide to giant covalent structures: how diamond, graphite and graphene are bonded, why their properties differ so much, and how the structure explains melting point, hardness and electrical conductivity.
- 4.2.1(c)-i
- 4.2.1(c)-ii
- 4.2.1(c)-iii
- 4.2.2-a
- 4.2.2-c
What these spec points say
- 4.2.1(c)-i describe, in simple terms, the lattice structure of a crystalline solid which is giant molecular, including silicon(IV) oxide
- 4.2.1(c)-ii describe, in simple terms, the lattice structure of a crystalline solid which is giant molecular, including graphite
- 4.2.1(c)-iii describe, in simple terms, the lattice structure of a crystalline solid which is giant molecular, including diamond
- 4.2.2-a describe, interpret and predict the effect of different types of structure and bonding on the melting point of substances
- 4.2.2-c describe, interpret and predict the effect of different types of structure and bonding on the electrical conductivity of substances
Giant Covalent Structures: Diamond, Graphite, Graphene and SiO₂
Choose a step to watch each giant covalent structure build out from one atom and its bonds, and see how the bonding explains its properties.
© Dr. Mohammed Al-Fatah – onlinelearningsystem.net
GCSE Recap: Covalent Bonds and Small Molecules
Before you start, check three GCSE ideas that this page builds on.
What Is a Giant Covalent Structure?
In a giant covalent structure (also called a macromolecular or giant molecular structure) every atom is joined to its neighbours by strong covalent bonds that extend throughout the whole crystal.
There are no separate molecules: a single crystal of diamond is, in effect, one enormous molecule.
Compare this with a simple molecular substance such as iodine. Strong covalent bonds hold each I₂ molecule together, but only weak intermolecular forces act between the molecules.
| Feature | Giant covalent (diamond, graphite, graphene) | Simple molecular (iodine, ice, carbon dioxide) |
|---|---|---|
| Particles | Atoms in a continuous network | Discrete molecules |
| Forces broken on melting | Strong covalent bonds | Weak intermolecular forces |
| Melting point | Very high (diamond sublimes above 3600 °C) | Low |
| Solubility in water | Insoluble | Usually low; some polar molecules dissolve |
| Electrical conductivity | Usually none; graphite and graphene conduct | None |
Key idea: Melting a giant covalent structure means breaking covalent bonds, which needs a great deal of energy, so these substances have very high melting points and are hard or, in the case of graphite, strongly bonded within each layer.
In a giant covalent structure every atom is held by strong covalent bonds, so melting means breaking those bonds, unlike a simple molecular substance.
Quick Check: Identify the Structure from Data
Use the data to decide which type of structure the solid has.
Diamond
In diamond each carbon atom forms four single covalent bonds to four other carbon atoms arranged tetrahedrally around it, with bond angles of 109.5°. The result is a rigid three-dimensional network in which every atom is locked in place by strong C-C bonds.
| Property | Explanation from the structure |
|---|---|
| Very high melting and sublimation point | Many strong covalent bonds must be broken |
| Extremely hard | The rigid tetrahedral network resists any distortion; diamond is used in cutting tools and drill tips |
| Does not conduct electricity | All four outer electrons of each carbon are used in bonding, so there are no delocalised electrons |
| Insoluble in all solvents | No solvent interaction can compete with the covalent bonds |
Silicon has the same structure as diamond, which is why it is also a hard, high-melting solid, although its Si-Si bonds are weaker than C-C bonds.
Silicon atoms are larger than carbon atoms, so the Si-Si bond is longer and the shared pair is further from both nuclei. This is why silicon melts at a lower temperature than diamond.
Exam sentence: Diamond has a giant covalent structure in which each carbon atom is bonded to four others by strong covalent bonds, so a large amount of energy is needed to break the bonds and it has a very high melting point.
Quick Check: Predict the Properties of Silicon Carbide
Use the structure described to predict the properties of a compound that is not on this page.
Quick Check: Explain Silicon and Phosphorus
Write a short explanation, then compare it with the mark points and the model answer.
Watch: graphite's fourth electron
In graphite each carbon makes only three bonds. What happens to its fourth electron?
Graphite
In graphite each carbon atom forms three covalent bonds to three other carbon atoms, giving flat hexagonal layers with bond angles of 120°.
The fourth outer electron of each carbon is not used in a localised bond. It sits in a p orbital above and below the layer and becomes delocalised across the whole sheet.
The layers are held to each other only by weak instantaneous dipole-induced dipole (id-id) forces, also called London dispersion forces, and are relatively far apart.
| Property | Explanation from the structure |
|---|---|
| Very high melting point | The strong covalent bonds within the layers must be broken |
| Soft and slippery | The weak forces between layers let them slide over each other; graphite is used as a lubricant and in pencils |
| Conducts electricity | The delocalised electrons can move along the layers when a voltage is applied. Graphite conducts well along the layers but very poorly at right angles to them, because the electrons cannot move easily from one layer to the next |
| Low density compared with diamond | The layers are widely spaced |
| Insoluble | Covalent bonding throughout the layers |
Exam focus: Graphite conducts because it has delocalised electrons; it is soft because the layers are held together only by weak id-id forces. Do not say the layers are held by covalent bonds, and do not say graphite conducts because it is a metal.
Quick Check: Which Way Does Graphite Conduct?
Predict what happens when the conductivity of a graphite crystal is measured in two directions.
Graphene
Graphene is a single layer of graphite: a sheet of carbon atoms one atom thick, each bonded covalently to three others in a hexagonal pattern, with one delocalised electron per carbon spread over the sheet. It is described as a two-dimensional giant covalent structure.
| Property | Explanation | Application |
|---|---|---|
| Extremely strong for its thickness | Strong covalent bonds throughout the sheet | Reinforcing composites and lightweight materials |
| Excellent electrical conductor | Delocalised electrons move freely across the sheet | Flexible electronics, touchscreens, sensors |
| Very good thermal conductor | Vibrations pass easily through the rigid covalent network | Heat management in electronic devices |
| Transparent and flexible | A single atomic layer absorbs little light and can bend without breaking bonds | Transparent conducting films |
Exam sentence: Graphene is a single layer of graphite in which each carbon atom is covalently bonded to three others; the delocalised electrons make it an excellent conductor and the covalent network makes it very strong.
Quick Check: From Graphene to Nanotubes
Apply what you know about graphene to two related materials, answering as quickly as you can.
Silicon(IV) Oxide
Silicon(IV) oxide, SiO2, is a giant molecular structure. Each silicon atom is bonded tetrahedrally to four oxygen atoms and each oxygen bridges two silicon atoms, giving an extended Si-O network.
It is very hard, has a high melting point (about 1700 °C), is insoluble in water and does not conduct, because every outer electron is used in Si-O bonds.
Exam sentence: Silicon(IV) oxide has a giant molecular structure in which strong Si-O covalent bonds extend throughout the lattice, so it has a high melting point and does not conduct electricity.
Comparing the Three Carbon Structures
| Diamond | Graphite | Graphene | |
|---|---|---|---|
| Bonds per carbon atom | 4 | 3 | 3 |
| Arrangement | Tetrahedral, 109.5°, three-dimensional | Hexagonal layers, 120° | One hexagonal layer, 120° |
| Delocalised electrons | None | One per carbon, within each layer | One per carbon, across the sheet |
| Electrical conductivity | None | Along the layers | Excellent |
| Hardness | Extremely hard | Soft, layers slide | Very strong sheet |
| Melting point | Very high | Very high | Very high |
Key idea: The same element gives three very different materials because the bonding arrangement differs: four localised bonds give hardness and no conduction; three bonds plus a delocalised electron give layers that conduct.
Diamond, graphite and graphene are all carbon, but four localised bonds per atom give hardness while three bonds plus a delocalised electron give conduction.
Quick Check: Two Forms of Boron Nitride
Drag the words into place to compare the two forms of another giant covalent compound.
Quick Check: The Whole Page in Five Rounds
Pick the one accurate statement in each round, from the carbon structures to silicon(IV) oxide.
Where This Sits in the Cambridge Syllabus
Cambridge places giant molecular structures in 4.2.1(c). You must describe the lattice structures of silicon(IV) oxide, graphite and diamond in simple terms and use them to explain melting point and conductivity (4.2.2).
Common Exam Mistakes
- Saying graphite is soft because its covalent bonds are weak. The covalent bonds are strong; it is the forces between layers that are weak.
- Saying diamond conducts electricity, or that graphite conducts because it contains ions or is metallic. Only delocalised electrons explain the conduction.
- Describing melting as breaking intermolecular forces. In a giant covalent structure the covalent bonds themselves must be broken.
- Forgetting the bond angle: 109.5° in diamond, 120° in graphite and graphene.
- Drawing a diamond fragment with carbon atoms that have only three bonds, or a graphite layer with four bonds per carbon.
Exam sentence: Giant covalent structures have very high melting points because many strong covalent bonds must be broken; only those with delocalised electrons, graphite and graphene, conduct electricity.
FAQs
These questions secure the structure and property links examiners test most on giant covalent lattices.
If graphite is soft, why is its melting point so high?
Softness and melting point depend on different forces. Graphite is soft because only weak id-id forces hold the layers together. Melting breaks the strong covalent bonds within the layers, so its melting point is very high, close to that of diamond.
Why does graphite conduct electricity but diamond does not?
Each carbon in graphite uses three outer electrons in bonds and the fourth is delocalised across the layer, so it can move when a voltage is applied. In diamond all four outer electrons are held in localised C-C bonds.
Is graphene just a very thin piece of graphite?
Graphene is a single layer of graphite, one atom thick. It has the same hexagonal bonding and one delocalised electron per carbon, but no layers above or below, so it is very strong for its thickness and an excellent conductor.
Why is SiO2 a hard solid when CO2 is a gas?
SiO2 is a giant covalent lattice: each silicon is bonded to four oxygen atoms and each oxygen bridges two silicon atoms, so melting breaks strong Si-O bonds. CO2 is a simple molecule with only weak forces between molecules.
Why does silicon melt at a lower temperature than diamond?
Silicon has the same tetrahedral structure as diamond, but its atoms are larger. The Si-Si bond is longer, so the shared pair is further from both nuclei and the bond is weaker. Less energy is needed to break the bonds.
Copyright notice: This OLS revision content, including the explanations, layout, diagrams, tables and embedded learning structure, is authored for Online Learning System by Dr. Mohammed Al-Fatah. It may not be copied, reproduced, redistributed or adapted without written permission.
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