Giant covalent structures- CIE iGCSE Chemistry Notes - New Syllabus
Giant covalent structures for iGCSE Chemistry Notes
Core Syllabus
- Describe the giant covalent structures of graphite and diamond
- Relate the structures and bonding of graphite and diamond to their uses, limited to:
(a) graphite as a lubricant and as an electrode
(b) diamond in cutting tools
Supplement Syllabus
- Describe the giant covalent structure of silicon(IV) oxide, SiO₂
- Describe the similarity in properties between diamond and silicon(IV) oxide, related to their structures
Giant Covalent Structures of Graphite and Diamond
Giant Covalent Structures of Graphite and Diamond
A giant covalent structure (also called a macromolecular structure) is a structure in which a huge number of atoms are bonded together by strong covalent bonds in a continuous network. These structures have very high melting points and strong physical properties.
1. Diamond
- Each carbon atom is covalently bonded to 4 other carbon atoms in a tetrahedral arrangement.
- This forms a 3D lattice of strong covalent bonds throughout the structure.
- There are no free electrons or layers – all electrons are used in bonding.
Key Features of Diamond:
- Hard and rigid – due to the strong network of covalent bonds.
- Very high melting point – a large amount of energy is needed to break the many covalent bonds.
- Does not conduct electricity – no free electrons or ions.
2. Graphite
- Each carbon atom is bonded to 3 other carbon atoms in flat hexagonal layers.
- The carbon atoms form layers of hexagons, with weak forces between layers.
- The 4th electron in each carbon atom is delocalised and free to move between layers.
Key Features of Graphite:
- Slippery and soft – layers can slide over each other easily due to weak intermolecular forces between layers.
- Conducts electricity – because of the delocalised electrons that can move freely.
- High melting point – due to strong covalent bonds within the layers.
| Property | Diamond | Graphite |
|---|---|---|
| Type of bonding | Each carbon makes 4 covalent bonds (tetrahedral) | Each carbon makes 3 covalent bonds (hexagonal layers) |
| Structure | 3D lattice | Layers of hexagons |
| Electrical conductivity | No (no free electrons) | Yes (delocalised electrons) |
| Hardness | Very hard | Soft and slippery |
| Melting point | Very high | Very high |
Uses of Graphite and Diamond
1. Diamond – Used in Cutting Tools
- Structure: Each carbon atom is bonded to 4 others in a rigid 3D tetrahedral arrangement.
- Bonding: All covalent bonds are strong and extend throughout the structure.
Diamond is extremely hard due to its strong covalent bonds throughout the structure. This makes it ideal for cutting, grinding, or drilling hard materials. It also has a very high melting point, so it doesn’t break down easily under heat.
2. Graphite – Used as a Lubricant and as an Electrode
- Structure: Carbon atoms are bonded in layers of hexagons, with weak forces between layers.
- Bonding: Each atom is bonded to 3 others; the 4th electron is delocalised.
Layers in graphite can slide over each other easily due to weak forces between them. This makes graphite feel slippery – ideal for use in lubricants or pencil ‘lead’.
The delocalized electrons in graphite can move freely between layers. This allows graphite to conduct electricity, making it suitable for use in electrodes in electrolysis. It also has a high melting point and is chemically stable.
Example
Explain why diamond is used in cutting tools and graphite is used as an electrode.
▶️Answer/Explanation
Diamond – Cutting Tools:
Diamond has a giant covalent structure where each carbon atom is covalently bonded to 4 others in a rigid 3D tetrahedral arrangement. All the covalent bonds are extremely strong and extend throughout the structure. This makes diamond very hard and rigid, with an exceptionally high melting point. Because of these properties, diamond can cut or grind even the hardest materials without breaking or wearing down easily. It is therefore used in industrial cutting, drilling, and grinding tools.
Graphite – Electrodes and Lubricants:
Graphite has layers of carbon atoms arranged in hexagons. Each carbon atom is bonded to 3 others, leaving the 4th electron delocalised. These delocalised electrons can move freely between layers, allowing graphite to conduct electricity. This makes graphite ideal for use as electrodes in electrolysis or in batteries.
The layers in graphite are held together by weak intermolecular forces, so they can slide over each other easily. This gives graphite a soft and slippery texture, making it useful as a lubricant and as pencil ‘lead’.
Conclusion:
The different bonding in diamond and graphite explains their very different uses. Diamond’s strong 3D covalent network gives hardness (cutting tools), while graphite’s delocalised electrons and sliding layers allow electrical conduction and lubrication.
Giant Covalent Structures of Silicon(IV) Oxide
Covalent Structure of Silicon(IV) Oxide
Silicon(IV) oxide (also called silicon dioxide or silica) is a compound found naturally in sand, quartz, and some rocks. It has a giant covalent structure, which means the atoms are bonded together in a continuous 3D network using strong covalent bonds.
Structure of SiO2:
- Each silicon atom forms 4 single covalent bonds with 4 oxygen atoms.
- Each oxygen atom is bonded to 2 silicon atoms.
- This results in a tetrahedral arrangement around each silicon atom, similar to the structure of diamond.
- There are no small molecules – it is one large molecule held together by covalent bonds throughout.
Bonding in SiO2:
- All bonds in the structure are strong covalent bonds.
- There are no weak intermolecular forces, unlike simple molecular substances.
- The entire crystal is essentially a single molecule with millions of atoms joined together.
Physical Properties of SiO2:
- Very high melting and boiling points: A large amount of energy is needed to break the many covalent bonds throughout the structure.
- Very hard and rigid: The strong network of covalent bonds makes it extremely tough and resistant to physical damage.
- Insoluble in water: It does not dissolve in water or most solvents because of its strong covalent bonding.
- Does not conduct electricity: There are no free electrons or ions to carry charge.
Similarities between Diamond and Silicon(IV) Oxide
Although diamond is made of only carbon atoms and SiO2 is made of silicon and oxygen, both substances have very similar structural features that explain why their physical properties are alike.
| Feature | Diamond | Silicon(IV) Oxide (SiO2) |
|---|---|---|
| Type of structure | Giant covalent (macromolecular) | Giant covalent (macromolecular) |
| Bonding | Each carbon atom forms 4 covalent bonds with other carbon atoms | Each silicon atom forms 4 covalent bonds with oxygen atoms |
| Arrangement | Tetrahedral arrangement of atoms in 3D lattice | Tetrahedral arrangement around each silicon atom |
| Melting and boiling points | Very high – strong covalent bonds throughout | Very high – strong covalent bonds throughout |
| Hardness | Extremely hard – used in cutting tools | Hard – not as hard as diamond but still very tough |
| Electrical conductivity | Does not conduct electricity – no free electrons | Does not conduct electricity – no free electrons or ions |
| Solubility in water | Insoluble | Insoluble |
Example
Explain why \( \text{SiO}_2 \) (silicon(IV) oxide) has similar physical properties to diamond, despite being made of different elements.
▶️Answer/Explanation
1. Structure and Bonding:
\( \text{SiO}_2 \) has a giant covalent structure, just like diamond. Each silicon atom is covalently bonded to 4 oxygen atoms, and each oxygen atom is bonded to 2 silicon atoms. This creates a continuous 3D network of strong covalent bonds with a tetrahedral arrangement around each silicon atom. Similarly, diamond consists of carbon atoms, each covalently bonded to 4 others in a tetrahedral 3D lattice.
2. Melting and Boiling Points:
Both \( \text{SiO}_2 \) and diamond have very high melting and boiling points. A large amount of energy is needed to break the many strong covalent bonds throughout their structures. This is why sand (mostly \( \text{SiO}_2 \)) and diamond remain solid even at extremely high temperatures.
3. Hardness:
The strong covalent network makes both substances very hard and rigid. Diamond is harder, but \( \text{SiO}_2 \) is still very tough and used in glass, ceramics, and construction materials. The extensive bonding gives high resistance to scratching or breaking.
4. Electrical Conductivity:
Neither diamond nor \( \text{SiO}_2 \) has free-moving electrons or ions. All valence electrons are used in covalent bonding. As a result, they are electrical insulators, not conducting electricity in any state.
5. Solubility:
Both are insoluble in water and most solvents. The covalent network cannot be broken by polar or non-polar solvents, unlike ionic or simple molecular substances.
Conclusion:
Despite being composed of different atoms (carbon vs silicon and oxygen), diamond and \( \text{SiO}_2 \) share similar physical properties because both have giant covalent structures with tetrahedral arrangements and extensive strong covalent bonding throughout the lattice.