Introduction
Metals make up most of the periodic table, and they share some very recognisable properties: they conduct electricity, can be hammered into shapes (malleable), pulled into wires (ductile), and often have high melting points. All of these properties come from one type of bonding — the metallic bond. In this lesson we'll look at what a metallic bond actually is, what makes it stronger or weaker, and how alloys relate to this model.
What is a metallic bond?
A metal can be pictured as a regular lattice of positive metal ions surrounded by a "sea" of delocalised valence electrons. These electrons don't belong to any single atom — they are free to move throughout the whole structure. The metallic bond is the strong electrostatic attraction between the fixed positive ions and this mobile sea of negative electrons.
This model neatly explains typical metallic properties:
- Electrical conductivity – delocalised electrons can flow through the structure when a voltage is applied.
- Thermal conductivity – the mobile electrons transfer kinetic energy quickly.
- Malleability and ductility – layers of positive ions can slide over each other without breaking bonds, because the electron sea simply adjusts and keeps holding the ions together.
- High melting and boiling points – a lot of energy is needed to overcome the attraction between ions and electrons.
What makes metallic bonds stronger or weaker?
Two main factors control the strength of a metallic bond (Talbot, p.202):
- Number of valence (delocalised) electrons per atom – more delocalised electrons means a stronger "glue" holding the lattice together.
- Charge density of the metal ion – this is the ratio of the ion's charge to its volume. Smaller ions with a higher charge have a greater charge density, and attract the electron sea more strongly.
In short:
strength of metallic bond ∝ charge density ∝ number of valence electrons per ion
So across a period