Metallic Bonding Explained

Metallic bonding is a Supplement topic, but understanding it is essential for explaining the physical properties of metals that appear throughout the Core syllabus. Paper 4 frequently asks for 2-3 mark explanations linking the model to a specific property.

The model

In a metal, each atom loses its outer-shell electrons. These electrons become delocalised, meaning they are no longer associated with any one particular atom. They form a “sea” of electrons that moves freely throughout the metal structure.

The remaining positive ions (cations) are arranged in a regular lattice. The metallic bond is the strong electrostatic attraction between these positive metal ions and the sea of delocalised electrons.

Key components:

• Positive metal ions arranged in a regular lattice
• Delocalised electrons that move freely between the ions
• Electrostatic attraction between the ions and the electron sea

This is different from ionic bonding (attraction between oppositely charged ions) and covalent bonding (shared pair of electrons between specific atoms). In metallic bonding, the electrons are shared across the entire structure.

Properties explained by the model

High melting and boiling points

The electrostatic attraction between positive ions and delocalised electrons is strong and acts throughout the giant metallic structure. A large amount of energy is needed to overcome these forces, so metals have high melting points.

The strength of metallic bonding increases with:

• More delocalised electrons per atom (e.g. Mg donates 2 electrons vs Na donating 1)
• Smaller ionic radius (ions pack closer, stronger attraction)

This explains why magnesium (m.p. 650 C) has a higher melting point than sodium (m.p. 98 C). Magnesium has a 2+ ion with more delocalised electrons and a smaller ionic radius.

Electrical conductivity

Delocalised electrons are free to move through the structure. When a potential difference is applied, electrons drift toward the positive terminal. This movement of charged particles is an electric current.

Metals conduct using electrons, not ions. This is different from ionic conduction in electrolytes, where ions carry the charge.

Thermal conductivity

When one end of a metal is heated, the delocalised electrons gain kinetic energy and move faster. They collide with other electrons and ions, transferring energy rapidly through the structure. The free movement of delocalised electrons makes metals good thermal conductors.

Malleability and ductility

When a force is applied to a metal, layers of positive ions can slide over each other. The metallic bond does not break because the delocalised electrons can redistribute around the ions in their new positions, maintaining the electrostatic attraction.

This contrasts sharply with ionic compounds, where displacing layers brings ions of the same charge next to each other, causing repulsion and fracture. Ionic compounds are brittle; metals are malleable.

Property	Metallic bonding explanation
High melting point	Strong attraction between ions and delocalised electrons throughout giant structure
Electrical conductivity	Delocalised electrons free to move and carry charge
Thermal conductivity	Delocalised electrons transfer kinetic energy rapidly
Malleability	Layers of ions slide; delocalised electrons redistribute, maintaining bonding
Lustrous (shiny)	Delocalised electrons interact with light, reflecting it

Alloys and disrupted layers

An alloy is a mixture of a metal with one or more other elements (usually metals). In an alloy, atoms of different sizes disrupt the regular arrangement of layers. This makes it harder for layers to slide over each other, so alloys are harder and less malleable than pure metals.

Examples covered in alloys and uses:

• Steel: iron + carbon (and other elements)
• Brass: copper + zinc
• Bronze: copper + tin

The different-sized atoms act as obstacles to layer movement, increasing hardness and strength.

Comparing bonding types

Feature	Metallic	Ionic	Covalent (simple)
Particles	Metal ions + delocalised electrons	Positive and negative ions	Molecules
Forces overcome on melting	Metallic bonds (strong)	Ionic bonds (strong)	Intermolecular forces (weak)
Melting point	High	High	Low
Conducts as solid?	Yes (electrons)	No	No
Conducts when molten?	Yes (electrons)	Yes (ions)	No
Malleable?	Yes	No (brittle)	N/A

This comparison table is worth memorising. A 6-mark question asking for differences between bonding types can draw from every row.

Common exam errors

1. “Metallic bonds are weak” — they are strong; metals have high melting points.
2. “Metals conduct because ions move” — metals conduct via delocalised electrons, not ions.
3. “Atoms slide over each other” — the correct term is ions (or positive ions), not atoms. Once an atom has lost its outer electrons, it is an ion.
4. “Electrons are shared between two atoms” — that describes covalent bonding. In metallic bonding, electrons are delocalised across the whole structure.

Worked exam questions

Explain, in terms of its structure and bonding, why copper is a good electrical conductor. [2 marks]

• Copper has delocalised electrons [1]
• These electrons are free to move through the structure and carry charge [1]

Note: Simply saying “copper has free electrons” may earn 1 mark. For the second mark, link movement to carrying charge.

Explain why magnesium has a higher melting point than sodium. [3 marks]

• Both have metallic bonding: electrostatic attraction between positive ions and delocalised electrons [1]
• Magnesium has a 2+ ion and donates 2 electrons per atom, compared to sodium’s 1+ ion and 1 electron [1]
• The greater charge on the ion and more delocalised electrons give a stronger electrostatic attraction, requiring more energy to overcome [1]

Explain why metals are malleable but ionic compounds are brittle. [3 marks]

• In metals, layers of positive ions can slide over each other [1]
• The delocalised electrons redistribute, maintaining the metallic bond [1]
• In ionic compounds, when layers shift, ions of the same charge come adjacent, causing repulsion and the lattice fractures [1]