Periodic Trends in Structure and Melting Point

Full Notes

Metallic Bonding

In metals, positive metal ions are arranged in a giant lattice and outer shell electrons are delocalised, moving freely through the structure.

Metallic bonding is the strong electrostatic attraction between these delocalised electrons and positive metal ions.

Metallic bonding explains the high electrical conductivity and malleability of metals.

Example Structure of Sodium (Na)

Each sodium atom loses one outer electron, forming Na⁺ ions. The lost electrons become delocalised, forming an electron cloud.

There is strong attraction between Na⁺ ions and the delocalised electrons, which holds the metal together.

Giant Covalent Structures

Some non-metal elements and compounds form giant covalent structures. These have no maximum size and all atoms are covalently bonded in a repeating pattern. Common examples include diamond, graphite, graphene and silicon.

Diamond (C)

In diamond, each carbon atom bonds to four others in a 3D tetrahedral lattice with no maximum size.

Key Properties:

• Very high melting point (~3550 °C) as covalent bonds require very high amounts of energy to break.
• Does not conduct electricity as there are no free electrons or ions.
• Very hard because of strong bonding in all directions.

Graphite (C)

In graphite, each carbon atom bonds to three others, forming hexagonal layers with delocalised electrons between the layers and weak forces of attraction between layers.

Key Properties:

• High melting point (~3700 °C) as covalent bonds require very high amounts of energy to break.
• Conducts electricity as delocalised electrons can move freely between layers.
• Soft and slippery, layers slide over each other due to weak forces of attraction.

Graphene

Graphene is a single layer of graphite.

Key Properties:

• High strength, flexibility, and excellent electrical conductivity.

Silicon

Silicon has a giant covalent structure similar to diamond, with silicon atoms bonded in a tetrahedral arrangement.

Key Properties:

• High melting point and poor conductivity.

Properties of Giant Lattices

• Giant Metallic Lattices:
  • Melting/Boiling Points: High (strong metallic bonds).
  • Electrical Conductivity: Good (delocalised electrons).
  • Solubility: Insoluble in water and organic solvents.
• Giant Covalent Lattices:
  • Melting/Boiling Points: Very high (strong covalent bonds).
  • Electrical Conductivity:
    • Diamond and silicon: no free electrons → non-conductors.
    • Graphite and graphene: conduct (delocalised electrons).
  • Solubility: Insoluble.

Melting Point Trends Across Periods

Melting points vary across a period in the periodic table due to different bonding and structures.

For example, for period elements (Na to Ar):

• Metals (Na, Mg, Al):
  High melting points due to strong metallic bonding. Metallic bonding increases in strength from Na → Al due to more delocalised electrons and greater positive charge of metal ions.


• Silicon (Si):
  Highest melting point due to its giant covalent structure with strong covalent bonds requiring a lot of energy to break.


• Non-metals (P₄, S₈, Cl₂, Ar):
  Lower melting points due to weak London Dispersion forces holding molecules together. S₈ has a higher melting point than P₄ and Cl₂ because it is a larger molecule, meaning stronger London Dispersion forces.

Matt’s exam tip

Don’t forget that sulfur has a slightly higher melting point than phosphorus. This is because sulfur exists as molecules of S₈, whereas phosphorus is most commonly found as P₄ molecules. S₈ molecules are larger than P₄ meaning stronger van der Waals forces and a higher melting point.

Summary

• Metals have giant metallic lattices with delocalised electrons.
• Non-metals like diamond, graphite, graphene and silicon form giant covalent structures.
• Giant lattices give high melting/boiling points, insolubility and conductivity variation.
• Across a period: melting points high for metals and silicon, lower for molecular elements.