Metallic Structure and Properties

Metallic structure and bonding is a core topic in the Cambridge 0620 syllabus. Paper 4 regularly asks candidates to explain metal properties in terms of structure and bonding — vague answers that simply say “metals are strong” without referencing delocalised electrons and ion layers score poorly.

The metallic structure

A metal consists of:

• Positive metal ions (cations) arranged in regular layers
• A sea of delocalised electrons that moves freely between and around the ions

The outer-shell electrons of each metal atom become detached and are shared across the entire structure. These are the delocalised electrons. The metal atoms, having lost electrons, become positive ions.

The metallic bond is the strong electrostatic attraction between the positive metal ions and the sea of delocalised electrons. This definition must be stated precisely for full marks.

Properties explained

High melting and boiling points

The metallic bonds (attraction between positive ions and delocalised electrons) are strong. A large amount of energy is required to overcome these attractions and separate the ions, resulting in high melting points.

The strength of metallic bonding depends on:

• Number of delocalised electrons per atom (more electrons = stronger bonding)
• Size of the ion (smaller ions pack closer = stronger attraction)
• Charge on the ion (higher charge = stronger attraction)

This explains why sodium (1 delocalised electron, large ion) melts at 98 °C, while iron (more delocalised electrons, smaller ion) melts at 1538 °C.

Electrical conductivity

Delocalised electrons are free to move throughout the metallic structure. When a voltage is applied:

• Electrons drift towards the positive terminal
• This movement of charge is an electric current

Metals conduct electricity in both the solid and liquid state because the delocalised electrons remain free to move in both states.

This is different from ionic compounds, which conduct only when molten or dissolved (ions must be free to move).

Thermal conductivity

When one end of a metal is heated, the delocalised electrons gain kinetic energy and move faster. They transfer this energy rapidly through the structure by colliding with other electrons and ions. This is why metals are good conductors of heat.

Malleability and ductility

When a force is applied to a metal:

1. Layers of positive ions slide over one another
2. The delocalised electrons shift with the ions
3. The metallic bond is maintained throughout — it reforms in the new position
4. The metal changes shape without breaking

This is why metals are malleable (can be hammered) and ductile (can be drawn into wire).

Contrast with ionic compounds

In giant ionic structures, if layers are forced to slide, ions of the same charge align next to each other, causing strong repulsion. The structure shatters — ionic compounds are brittle, not malleable.

Alloys

An alloy is a mixture of two or more elements, at least one of which is a metal. The differently-sized atoms in an alloy disrupt the regular arrangement of layers.

Property	Pure metal	Alloy
Layer arrangement	Regular, uniform	Disrupted by different-sized atoms
Sliding of layers	Easy	Difficult
Hardness	Softer	Harder
Strength	Lower	Higher

Examples:

• Steel = iron + carbon (stronger than pure iron)
• Brass = copper + zinc (harder than pure copper)
• Bronze = copper + tin (harder, more durable)
• Stainless steel = iron + chromium + nickel (resists corrosion)

Metallic bonding vs other bonding types

Feature	Metallic	Ionic	Covalent (simple molecular)
Particles	Positive ions + delocalised electrons	Positive and negative ions	Molecules
Melting point	Usually high	High	Low
Conducts solid?	Yes	No	No
Conducts liquid?	Yes	Yes	No
Malleable?	Yes	No (brittle)	Not applicable

Worked exam question

Copper is used for electrical wiring and saucepans. Explain, in terms of structure and bonding, why copper is a good conductor of both electricity and heat. [4]

Copper has a metallic structure with positive copper ions in layers surrounded by a sea of delocalised electrons [1]. The delocalised electrons are free to move throughout the structure [1]. When a voltage is applied, the electrons move/drift carrying electrical charge [1]. When heated, the electrons gain kinetic energy and transfer thermal energy rapidly by colliding with ions and other electrons throughout the structure [1].

Common exam mistakes

1. Saying “metals have free electrons” without explaining that they are delocalised from the outer shells of metal atoms.
2. Writing “atoms slide over each other” — it is the layers of positive ions (not atoms) that slide. The atoms have become ions by losing electrons.
3. Saying metals conduct “because they have electrons” — all substances have electrons. The key is that metals have delocalised electrons that are free to move.
4. Confusing metallic bonding with ionic bonding — metallic bonding involves delocalised electrons shared across many ions, not the transfer of electrons from one atom to another.