Properties of Metallic Substances (AQA A Level Chemistry)

• Metals form giant metallic lattices in which the metal ions are surrounded by a ‘sea’ of delocalised electrons
• The metal ions are often packed in hexagonal layers or in a cubic arrangement
• This layered structure with the delocalised electrons gives a metal its key properties

Layers of copper ions (the delocalised electrons are not shown in the diagram)

• If other atoms are added to the metal structure, such as carbon atoms, this creates an alloy
• Alloys are much stronger than pure metals, because the other atoms stop the layers of metal ions sliding over each other easily

• The strength of the metallic attraction can be increased by:
  • Increasing the number of delocalised electrons per metal atom
  • Increasing the positive charges on the metal centres in the lattice
  • Decreasing the size of the metal ions

• Due to the delocalised ‘sea’ of electrons, metallic structures have some characteristic properties:

Malleability

• Metallic compounds are malleable
• When a force is applied, the metal layers can slide
• The attractive forces between the metal ions and electrons act in all directions
• So when the layers slide, the metallic bonds are re-formed
• The lattice is not broken and has changed shape

How metals are malleable diagram

Atoms are arranged in layers so the layers can slide when force is applied

Strength

• Metallic compounds are strong and hard
  • Due to the strong attractive forces between the metal ions and delocalised electrons

Electrical conductivity

• Metals can conduct electricity when in the solid or liquid state
  • In the solid and liquid states, there are mobile electrons which can freely move around and conduct electricity
• When a potential difference is applied to a metallic lattice, the delocalised electrons repel away from the negative terminal and move towards the positive terminal
  • As the number of outer electrons increases across a period, the number of delocalised charges also increases:
    • Sodium = 1 outer electron
    • Magnesium = 2 outer electrons
    • Aluminium = 3 outer electrons
  • Therefore, the ability to conduct electricity also increases across a period

How metals conduct electricity diagram

The delocalised electrons move towards the positive terminal when a potential difference is applied

• Since the bonding in metals is non-directional, it does not really matter how the cations are oriented relative to each other

Thermal conductivity

• Metals are good thermal conductors due to the behaviour of their cations and their delocalised electrons
  • When metals are heated, the cations in the metal lattice vibrate more vigorously as their thermal energy increases
    • These vibrating cations transfer their kinetic energy as they collide with neighbouring cations, effectively conducting heat
  • The delocalised electrons are not bound to any specific atom within the metal lattice and are free to move throughout the material
    • When the cations vibrate, they transfer kinetic energy to the electrons
    • The delocalised electrons then carry this increased kinetic energy and transfer it rapidly throughout the metal, contributing to its high thermal conductivity.

Melting and boiling point

• Metals have high melting and boiling points
  • This is due to the strong electrostatic forces of attraction between the cations and delocalised electrons in the metallic lattice
  • These require large amounts of energy to overcome 
  • As the number of mobile charges increases across a period, the melting and boiling points increase due to stronger electrostatic forces