Metallic Bonding
Higher Tier
Metallic bonding is Chemistry-only content (spec points 1.52C to 1.54C). All content in this note is Higher Tier.
The Metallic Lattice
Metals consist of giant structures of atoms arranged in a regular, close-packed pattern called a metallic lattice.
In the metallic lattice, the outer (valence) electrons from each metal atom become delocalised: they are no longer associated with any individual atom but are free to move throughout the entire structure.
This creates an arrangement of positive metal ions surrounded by a “sea” of delocalised electrons.
Metallic bonding is the strong electrostatic attraction between the positively charged metal ions and the negatively charged delocalised electrons.
2-D Representation of a Metallic Lattice
A 2-D diagram shows:
- Positive metal ions (circles with + signs) arranged in a regular grid
- Delocalised electrons (small dots or e- symbols) distributed throughout the spaces between the ions
Properties of Metals
The metallic lattice and the nature of metallic bonding explain the characteristic properties of metals.
High Melting and Boiling Points
Most metals have high melting and boiling points because:
- There are strong electrostatic forces of attraction between the positive metal ions and the delocalised electrons throughout the lattice
- A large amount of energy is needed to overcome these forces
Electrical Conductivity
Metals are good conductors of electricity because:
- The delocalised electrons are free to move throughout the metallic lattice
- When a voltage is applied, these electrons move in one direction, constituting an electric current
Malleability and Ductility
Metals are malleable (can be hammered into shape) and ductile (can be drawn into wires) because:
- The atoms/ions in a metal are arranged in layers
- When a force is applied, these layers can slide over each other without breaking the metallic bonds
- The metallic bonding remains intact as the layers slide, so the metal does not shatter
Thermal Conductivity
Metals are good conductors of heat because:
- The delocalised electrons can transfer kinetic energy rapidly through the structure
Properties Summary Table
| Property | Explanation |
|---|---|
| High melting/boiling point | Strong electrostatic attractions between positive ions and delocalised electrons throughout the lattice; large energy needed to overcome them |
| Electrical conductivity | Delocalised electrons free to move and carry charge |
| Malleability/ductility | Layers of ions can slide over each other; metallic bonding maintained |
| Thermal conductivity | Delocalised electrons transfer kinetic energy through the structure |
Alloys
An alloy is a mixture of a metal with one or more other elements, usually other metals or carbon.
Most metals used in everyday applications are alloys, because pure metals are often too soft for practical use.
Examples of alloys:
- Steel: iron mixed with carbon (and sometimes other metals such as chromium and nickel)
- Brass: copper and zinc
- Bronze: copper and tin
- Gold alloys in jewellery: gold mixed with silver, copper and zinc
Why Alloys are Harder than Pure Metals
In a pure metal, all the ions are the same size and arranged in regular layers. The layers can slide over each other easily, making pure metals soft.
In an alloy, atoms of a different size disrupt the regular arrangement of the metal ions. This makes it harder for the layers to slide over each other, so the alloy is harder and stronger than the pure metal.
Exam Questions
Q: What is meant by the term metallic bonding?
Metallic bonding is the strong electrostatic attraction between positively charged metal ions and the delocalised electrons that surround them throughout the metallic lattice.
Q: Explain why metals can conduct electricity.
In metals, there are delocalised electrons that are free to move throughout the metallic lattice. When a voltage is applied, these electrons move in one direction, carrying charge and constituting an electric current.
Q: Explain why metals are malleable.
The ions in a metal are arranged in layers. When a force is applied, these layers can slide over each other without breaking the metallic bonds (the delocalised electrons rearrange to maintain bonding), so the metal bends rather than shattering.
Q: Explain why alloys are harder than the pure metals from which they are made.
In an alloy, atoms of different sizes are present. This disrupts the regular arrangement of metal ions in the lattice, preventing the layers from sliding over each other easily. Therefore, alloys are harder than pure metals.
Q: Why is pure iron too soft for most everyday uses?
In pure iron, the iron ions are all the same size and arranged in regular layers that can slide over each other easily, making pure iron too soft. Iron is alloyed with carbon to make steel, which is harder and stronger.
Exam Tip
When explaining malleability, you must refer to layers of ions, not just “atoms”. Also, make clear that the metallic bonds are NOT broken when layers slide; the delocalised electrons simply move with the ions. This distinguishes metallic materials from brittle ionic or covalent solids.