Metallic Bonding and Properties of MetalsAQA GCSE Chemistry: Revision notes
Section 1
What is metallic bonding?
Metallic bonding occurs when metal atoms lose their outer electrons, forming a sea of delocalised electrons surrounding a lattice of positive metal ions.
- Metal atoms lose their outer electrons → become positive ions
- The lost electrons are delocalised — free to move throughout the whole structure
- Strong electrostatic attraction between the positive ions and the delocalised electrons = the metallic bond
- This attraction acts in all directions → giant metallic lattice
- All metal atoms contribute electrons to the shared pool
Picture a crowd of people (positive ions) standing in a swimming pool full of water (delocalised electrons). The water surrounds everyone and holds them all together — that is the metallic bond.
Section 2
Properties of metals — explained
Every property of a metal follows from the sea of delocalised electrons and the giant lattice of positive ions:
| Property | Behaviour | Reason |
|---|---|---|
| Electrical conductivity | Conducts well | Delocalised electrons carry charge through the lattice |
| Thermal conductivity | Conducts heat | Electrons transfer kinetic energy rapidly |
| High melting point | Very high (most) | Strong electrostatic forces throughout the lattice |
| Malleable/ductile | Bends, drawn into wire | Layers of ions slide past each other — metallic bonding reforms |
| Lustrous (shiny) | Reflects light | Free electrons absorb and re-emit light |
| Density | High (most) | Atoms packed closely in regular lattice |
For any conductivity question: metals conduct because delocalised electrons are free to move. For ionic compounds: ions must be free to move. Never mix these up — ionic compounds do NOT conduct when solid.
"Why can metals be bent without breaking?" — layers of positive ions slide past each other. The metallic bonds reform after sliding because the delocalised electrons adjust. This is what makes metals malleable, not ionic or covalent compounds.
Section 3
Alloys
An alloy is a mixture of a metal with one or more other elements (usually other metals or carbon).
- Pure metals have a regular lattice — layers can slide easily → soft
- In an alloy, different-sized atoms distort the lattice
- Distortion prevents layers from sliding → harder and stronger than pure metal
| Alloy | Composition | Use |
|---|---|---|
| Steel | Iron + carbon | Construction, cars |
| Stainless steel | Iron + carbon + chromium | Cutlery, medical instruments |
| Brass | Copper + zinc | Musical instruments, pipes |
| Bronze | Copper + tin | Statues, medals |
| Duralumin | Aluminium + copper | Aircraft |
- Alloys are harder because different sized atoms disrupt the regular pattern of the lattice
Why is steel harder than pure iron? Pure iron has a regular lattice where layers of ions slide easily. When carbon atoms are added to make steel, the smaller carbon atoms sit between the iron atoms and distort the lattice. This prevents layers from sliding, making steel much harder and stronger.
Saying alloys conduct less electricity than pure metals is not always needed — examiners rarely ask this. Focus on hardness and strength as the key property change in alloys.
Must Know
- Metallic bonding = positive metal ions in a sea of delocalised electrons — electrostatic attraction holds them together
- Conducts electricity and heat because delocalised electrons are free to move
- Malleable and ductile because layers of ions slide and metallic bonding reforms
- High melting point — strong electrostatic forces throughout the lattice
- Alloys are harder than pure metals — different sized atoms distort the lattice and prevent sliding
- Steel = iron + carbon; Brass = copper + zinc; Bronze = copper + tin
That's the notes covered.
Carry on to the next subtopic.