Giant Structures
How atoms bond together in huge three-dimensional networks determines whether a material is hard, conducts electricity, melts at high temperature, or can be bent — and you can explain all of it by understanding structure.
Chapter Overview
🔷 Diamond & Graphite
- Both pure carbon (allotropes)
- Different bonding = completely different properties
- Diamond: hard, no conductor
- Graphite: soft, conducts electricity
⚪ Silicon(IV) Oxide
- SiO₂ — sand and quartz
- Similar structure to diamond
- Very hard, high melting point
- Used in sandpaper, furnace linings
✨ Metallic Bonding
- Metal atoms lose outer electrons
- Creates "sea of electrons"
- Explains why metals conduct, are malleable
- Works in pure metals and alloys
Diamond & Graphite: Two Forms of Carbon
What Are Allotropes?
Carbon can exist in multiple forms — diamond, graphite, and graphene. These are called allotropes: different physical forms of the same element. The key insight is that even though they're made of the same atoms, their bonding arrangements are totally different, so their properties are completely different.
💡 Key Concept
Structure determines properties. Two substances made of identical atoms can behave like completely different materials if their internal bonding is arranged differently.
Graphite: Layered Structure
In graphite, each carbon atom forms three covalent bonds with three neighbouring carbon atoms, creating flat layers of hexagons. Think of it like a honeycomb — if you look at one layer, it's six-sided rings connected endlessly in a sheet.
Graphite Structure (side view):
C—C—C—C (Strong covalent bonds
| | | | within each layer)
C—C—C—C
[Weak forces
between layers]
C—C—C—C
| | | |
C—C—C—C
Crucially: each carbon atom has one unpaired electron in its outer shell. This electron is not locked into a bond with any particular neighbouring atom — it becomes delocalised, meaning it's free to move between the layers.
Delocalised Electrons
Electrons that don't belong to a specific atom but can move freely through the structure. In graphite, these sit between the layers.
The layers themselves are held together only by very weak intermolecular forces (van der Waals forces). Covalent bonds within each layer are extremely strong, but the bonds between layers are not.
Q1. In graphite, each carbon atom forms how many covalent bonds?
Diamond: 3D Tetrahedral Network
In diamond, each carbon atom forms four covalent bonds with four neighbouring carbon atoms, arranged in a 3D tetrahedral pattern. Every single carbon atom is bonded to every carbon around it in all directions — there are no layers, no free electrons, just one continuous three-dimensional lattice of strong covalent bonds.
Diamond Structure (simplified):
Each C atom bonded to 4 other C atoms
in a tetrahedral arrangement:
C
|
C—C—C (all directions)
|
C
Result: A rigid 3D lattice
with NO free electrons
no free electrons
Q2. Why does diamond have a very high melting point, but graphite is easier to melt?
What to Memorise