Metallic Bonding

1. Definition of Metallic Bonding

• Metallic bonding: The electrostatic attraction between positive metal ions (cations) arranged in a giant metallic lattice and a sea of delocalised electrons.
• The delocalised electrons come from the outermost shell of the metal atoms; they are free to move throughout the lattice.
• This bonding is non-directional, meaning the attractive force acts equally in all directions around the metal cations.
• The metallic bond is strong, requiring significant energy to break.

2. Structure of a Metallic Lattice

• Arrangement of atoms: Metals consist of a giant structure of closely packed positive ions in a regular arrangement.
• Sea of electrons: The outer shell electrons are delocalised — not bound to any particular atom but free to move throughout the whole structure.
• Electrostatic attraction: The strong attraction between the delocalised electrons and the positive metal ions holds the structure together.
• The number of delocalised electrons depends on the metal’s position in the Periodic Table (e.g., Group 1 metals have 1 delocalised electron per atom, Group 2 metals have 2).

3. Formation of Metallic Bonding

1. Metal atoms lose their outer shell electrons.
2. These electrons become delocalised, moving freely in the space between ions.
3. The atoms become positively charged ions.
4. Strong electrostatic attraction occurs between the positive ions and the sea of electrons, binding the metal atoms together in a lattice.

Example for magnesium (Mg):

• Each Mg atom loses 2 electrons → becomes Mg²⁺ ion.
• These electrons move freely throughout the structure → form a sea of electrons.
• Attraction between Mg²⁺ ions and electrons forms metallic bonds.

4. Properties of Metals and Their Explanation in Terms of Structure and Bonding

(a) Good Electrical Conductivity

• Reason: The delocalised electrons can move freely through the lattice, carrying electric charge when a potential difference is applied.
• Example: Copper is used extensively in electrical wiring due to its high electrical conductivity and flexibility.
• In molten form: Metals still conduct electricity because the lattice is broken but electrons remain free to move.

(b) Malleability and Ductility

• Malleability: Metals can be hammered or rolled into thin sheets without breaking.
• Ductility: Metals can be drawn into wires.
• Reason:
  • The layers of positive metal ions can slide over one another without breaking the metallic bond.
  • This is possible because the delocalised electrons act as a ‘glue’ holding the ions together regardless of position.
• Examples:
  • Aluminium is malleable → used in making foil and cans.
  • Copper is ductile → used in electrical wires.

5. Additional Properties of Metals (Related to Metallic Bonding)

• High melting and boiling points:
  • Large amounts of energy are required to break the strong electrostatic attractions between ions and delocalised electrons.
  • Metals like tungsten have extremely high melting points due to strong metallic bonding.
• Thermal conductivity:
  • Delocalised electrons can transfer thermal energy rapidly throughout the metal.
• Lustre (shiny appearance):
  • Delocalised electrons reflect light, giving metals their characteristic shiny surface when polished.

6. Factors Affecting Strength of Metallic Bonding

1. Number of delocalised electrons per atom:
   • More delocalised electrons → stronger attraction → stronger metallic bonding.
   • Example: Group 2 metals have stronger metallic bonds than Group 1 metals because they release 2 electrons per atom instead of 1.
2. Charge of metal ion:
   • Higher positive charge → stronger electrostatic attraction to delocalised electrons.
   • Example: Al³⁺ forms stronger metallic bonds than Na⁺.
3. Size of metal ion:
   • Smaller ions → closer packing → stronger metallic bonding.
   • Example: Mg²⁺ ions are smaller than Na⁺ ions → stronger metallic bonds in magnesium.

7. Summary Table