Structure 2.3.1 — Metallic Bonding and Properties of Metals

Metallic Bonding:

A metallic bond is the strong electrostatic attraction between a regular lattice of positive metal ions (cations) and a ‘sea’ of delocalized electrons.

Structure of Metals:

• Metals form a giant metallic lattice.
• Atoms in a metal lose their outer shell electrons and become cations.
• These electrons become delocalized (free to move) and spread throughout the structure.

Diagram Description (visualize):

• Metal ions arranged in regular layers.
• Delocalized electrons move freely between them.

Properties of Metals and Their Explanations:

1. Electrical Conductivity:

• Metals conduct electricity in solid and molten states.
• This is because delocalized electrons are free to move and carry charge.

2. Thermal Conductivity:

• Metals conduct heat well due to:
• Delocalized electrons transferring kinetic energy quickly.
• Closely packed ions also vibrate and transfer heat via collisions.

3. Malleability and Ductility:

• Metals can be hammered into thin sheets (malleable) or drawn into wires (ductile).
  • 
• This is because layers of metal ions can slide over each other when a force is applied, without breaking the metallic bond.
  • 

Explanation of Why Metallic Bonds Are Not Broken During Deformation:

• The sea of electrons maintains attraction throughout the lattice even when layers shift.

Factors Affecting the Strength of Metallic Bonds:

• Number of delocalized electrons: More electrons → stronger bond (e.g. \( \text{Al} \) has 3).
• Charge on the metal ion: Higher charge → stronger electrostatic attraction.
• Size of the metal ion: Smaller ions → stronger attraction between nucleus and electrons.

Trends Across the Periodic Table:

• Bond strength increases across a period as number of delocalized electrons and charge density increase.

Relation to Uses: