Metallic Bonding

This lesson covers metallic bonding and the properties of metals, as required by the AQA GCSE Chemistry specification (4.2.3). Metallic bonding is the third and final type of bonding you need to understand. You must be able to describe the structure of metals, explain their properties in terms of metallic bonding, and explain why alloys are harder than pure metals.

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What Is Metallic Bonding?

Metallic bonding is the strong electrostatic force of attraction between positively charged metal ions and a sea of delocalised electrons.

In a metal:

1. Metal atoms are arranged in a regular pattern called a giant metallic lattice.
2. The outer shell electrons of each metal atom are delocalised — they are no longer associated with any particular atom and are free to move throughout the structure.
3. This leaves behind positively charged metal ions (cations) arranged in rows and layers.
4. The delocalised electrons form a "sea" of electrons that surrounds the metal ions.
5. The strong electrostatic attraction between the positive ions and the negative sea of electrons holds the structure together — this is the metallic bond.

Exam Tip: The definition of metallic bonding is: "The strong electrostatic force of attraction between positively charged metal ions and a sea of delocalised electrons." You MUST use the terms "delocalised electrons" and "metal ions" for full marks. Simply saying "positive and negative charges" is not enough.

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Structure of Metals

graph TD
    A["Metallic Structure"] --> B["Positively Charged<br/>Metal Ions (cations)"]
    A --> C["Sea of Delocalised<br/>Electrons"]
    B --> D["Arranged in a regular<br/>giant metallic lattice"]
    C --> E["Free to move throughout<br/>the entire structure"]
    D --> F["Strong electrostatic<br/>attraction between<br/>ions and electrons"]
    E --> F
    F --> G["Metallic Bond"]

    style A fill:#2c3e50,color:#fff
    style B fill:#e74c3c,color:#fff
    style C fill:#3498db,color:#fff
    style G fill:#27ae60,color:#fff

The key features of the metallic structure are:

• Regular arrangement of positive metal ions.
• Delocalised electrons that are free to move.
• Strong attraction between the ions and electrons in all directions.

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Properties of Metals

1. High Melting and Boiling Points

Most metals have high melting and boiling points because there are strong electrostatic forces of attraction between the metal ions and the delocalised electrons. A large amount of energy is needed to overcome these forces.

Metal	Melting Point (degrees C)	Boiling Point (degrees C)
Iron (Fe)	1538	2861
Copper (Cu)	1085	2562
Aluminium (Al)	660	2519
Sodium (Na)	98	883
Mercury (Hg)	-39	357

Note: Some metals, like sodium, have relatively low melting points because they have fewer delocalised electrons (only 1 per atom) and larger ionic radii, resulting in weaker metallic bonds. Mercury is the only metal that is liquid at room temperature.

2. Good Electrical Conductivity

Metals are excellent conductors of electricity because the delocalised electrons are free to move through the structure. When a potential difference (voltage) is applied, the delocalised electrons flow in one direction, carrying charge through the metal — this is an electric current.

Exam Tip: When explaining why metals conduct electricity, always say "delocalised electrons are free to move and carry charge." Do not say "ions move" — the metal ions are in fixed positions in the lattice. Only the electrons move. This is the opposite of ionic conductivity, where it IS the ions that move.

3. Good Thermal Conductivity

Metals are also good conductors of heat because the delocalised electrons can transfer kinetic energy rapidly through the structure. When one end of a metal is heated, the electrons gain kinetic energy and move faster, colliding with other electrons and ions, passing the energy through the metal quickly.

4. Malleability and Ductility

Metals are malleable (can be hammered into shape) and ductile (can be drawn into wires). This is because the layers of metal ions can slide over each other without the metallic bond breaking. The delocalised electrons can move with the ions, maintaining the electrostatic attraction even when the layers shift.

This is a key difference from ionic compounds, which are brittle because sliding layers brings like charges together, causing repulsion.

Exam Tip: If asked why metals are malleable but ionic compounds are brittle, explain that in metals the delocalised electron sea can move with the ions, maintaining attraction. In ionic compounds, sliding layers brings same-charged ions adjacent, causing repulsion and shattering.

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Alloys

An alloy is a mixture of two or more elements, at least one of which is a metal. Common examples include:

Alloy	Components	Uses
Steel	Iron + carbon (+ other metals)	Construction, vehicles, bridges
Bronze	Copper + tin	Statues, medals, bearings
Brass	Copper + zinc	Musical instruments, fittings
Stainless steel	Iron + chromium + nickel	Cutlery, surgical instruments
Gold alloys	Gold + silver / copper / zinc	Jewellery (pure gold is too soft)

Why Are Alloys Harder Than Pure Metals?

In a pure metal, the atoms are all the same size and are arranged in regular layers. These layers can easily slide over each other, making the metal relatively soft.

In an alloy, atoms of different sizes are introduced into the lattice. These differently sized atoms disrupt the regular arrangement, making it more difficult for the layers to slide over each other. This makes the alloy harder and stronger than the pure metal.

graph LR
    A["Pure Metal"] --> B["Regular layers<br/>of same-sized atoms"]
    B --> C["Layers slide easily<br/>= softer metal"]
    D["Alloy"] --> E["Irregular arrangement<br/>with different-sized atoms"]
    E --> F["Layers cannot slide<br/>= harder metal"]

    style A fill:#3498db,color:#fff
    style C fill:#e74c3c,color:#fff
    style D fill:#27ae60,color:#fff
    style F fill:#27ae60,color:#fff