Metallic Bonding and Giant Structures
Metallic Bonding
Metallic bonding is the strong electrostatic attraction between positive metal ions (cations) and a sea of delocalised electrons.
In a metal:
1. Metal atoms lose their outer shell electrons
2. These electrons become delocalised — they are free to move throughout the entire structure
3. The positive metal ions are arranged in a regular lattice
4. The delocalised electrons drift between the ions, holding them together
This is sometimes described as a "sea of electrons" model.
Properties of Metals Explained by Metallic Bonding
High Melting and Boiling Points
The electrostatic attraction between the positive ions and the sea of delocalised electrons is very strong. A large amount of energy is needed to overcome these forces, so metals generally have high melting and boiling points.
Good Electrical Conductivity
The delocalised electrons are free to move through the structure. When a potential difference is applied, these electrons flow in one direction, carrying charge. This is why all metals are good conductors of electricity.
Good Thermal Conductivity
Delocalised electrons can transfer kinetic energy rapidly through the structure when one end is heated. The electrons gain energy and move faster, colliding with other electrons and ions, spreading the energy quickly.
Malleability and Ductility
Metals can be hammered into shapes (malleable) and drawn into wires (ductile). This is because the layers of ions can slide over each other without disrupting the bonding — the sea of delocalised electrons can redistribute around the ions in their new positions.
This is a key difference from ionic compounds, which shatter when layers are forced to slide because ions of the same charge end up next to each other and repel.
Alloys
An alloy is a mixture of two or more elements, at least one of which is a metal. Examples include:
| Alloy | Components | Property advantage |
|---|---|---|
| Steel | Iron + carbon (+ others) | Harder than pure iron |
| Bronze | Copper + tin | Harder than copper, resistant to corrosion |
| Brass | Copper + zinc | Harder than copper |
| Gold (e.g. 9ct) | Gold + copper/silver | Harder than pure gold |
Why Alloys Are Harder Than Pure Metals
In a pure metal, the atoms are all the same size and arranged in regular layers, so the layers slide easily.
In an alloy, atoms of different sizes disrupt the regular arrangement. This makes it harder for layers to slide over each other, making the alloy harder and stronger.
Giant Metallic Structures
Metals form giant metallic structures — the lattice of ions and delocalised electrons extends in all directions with no defined boundary. There are no individual molecules, which is why the formula for a metal is simply its symbol (e.g. Fe, Cu, Na).
Comparing the Three Types of Giant Structure
| Property | Giant ionic | Giant covalent | Giant metallic |
|---|---|---|---|
| Melting point | High | Very high | High (usually) |
| Conduct electricity (solid) | No | No (except graphite) | Yes |
| Conduct electricity (liquid/dissolved) | Yes | No | Yes |
| Malleable | No (brittle) | No | Yes |
Exam Tips
- Always describe metallic bonding as the attraction between positive ions and delocalised electrons — never just "atoms"
- When explaining properties, always link the property to the bonding: name the force, say why it is strong or weak
- Know why alloys are harder than pure metals — the different-sized atoms argument
- Be ready to compare metallic bonding with ionic and covalent bonding