2.2.8Additional lessons
Giant Covalent Structures — IB Chemistry SL Revision Notes & Topical Questions
5 min read 6 parts Lesson 29 of 98
Part 1 of 6What Is a Giant Covalent Structure
Some substances are a single enormous molecule of covalently bonded atoms.
Network
- Atoms are joined by covalent bonds in a continuous three-dimensional or layered network, with no separate molecules.
Strong bonds
- To melt the solid, many strong covalent bonds must be broken, so melting points are very high.
Insoluble
- They are insoluble in almost all solvents because solvent molecules cannot overcome the strong bonds.
Hard
- Atoms cannot easily be moved relative to one another, so most giant covalent solids are hard.
Part 2 of 6Carbon: Diamond and Graphite
Carbon: Diamond and Graphite
Diamond
- Each carbon is bonded to four others in a tetrahedral arrangement, forming a rigid structure that is extremely hard.
Diamond conduction
- All four outer electrons are used in bonds, so diamond does not conduct electricity.
Graphite
- Each carbon bonds to three others in flat hexagonal layers, with weak forces between layers.
Graphite conduction
- The fourth electron on each carbon is delocalised across the layer, so graphite conducts electricity along the layers.
| Property | Diamond | Graphite |
|---|---|---|
| Bonds per carbon | 4 | 3 |
| Structure | tetrahedral network | layers of hexagons |
| Hardness | very hard | soft, layers slide |
| Electrical conductivity | none | good along layers |
Part 3 of 6Silicon and Silicon Dioxide
1
Silicon has the same structure as diamond, with each atom bonded to four others in a tetrahedral network.
2
The network is strong, so silicon melts at 1414 °C and is hard.
3
Silicon is a semiconductor, so its conductivity is small but increases with temperature.
4
In silicon dioxide, SiO₂, each silicon atom bonds to four oxygen atoms and each oxygen links two silicon atoms.
5
Silicon dioxide melts at about 1700 °C, is hard and does not conduct, because all electrons are in localised bonds.
Part 4 of 6Explaining Differences
Layers in graphite
- Weak dispersion forces between layers allow them to slide, making graphite a lubricant.
Delocalised electrons
- These are free to move in graphite, which explains why it conducts but diamond does not.
Allotropes
- Diamond and graphite are allotropesdifferent forms of the same element in the same state.
Graphene
- A single layer of graphite is graphene, which conducts well because of its delocalised electrons.
Part 5 of 6Worked Example
Question Explain why graphite conducts electricity but diamond does not, although both contain only carbon.
- In diamond each carbon uses all four outer electrons in four covalent bonds, so no electrons are free to move.
- In graphite each carbon uses three outer electrons for bonding within a layer.
- The fourth electron from each atom is delocalised across the layer and acts as a mobile charge carrier.
Answer
- Graphite has delocalised electrons within its layers that can carry charge;
- diamond has none.
Part 6 of 6Exam Checklist & Tips
2.2.8 – Key Points for the Exam
- Giant covalent structure A network of atoms joined by covalent bonds throughout the whole crystal.
- Allotropes Different structural forms of the same element in the same physical state.
- Delocalised electrons Electrons not fixed to one atom or bond that are free to move.
- Semiconductor A material with electrical conductivity between a metal and an insulator.
Common Mistakes
- Saying that melting a giant covalent substance only breaks intermolecular forces; covalent bonds break.
- Stating that graphite conducts between its layers; it conducts mainly along them.
- Forgetting that graphite has weak forces between layers, not covalent bonds.
- Describing silicon dioxide as made of molecules of SiO₂.
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