Metallic Bonding

This lesson covers metallic bonding and the properties of metals as required by the Edexcel GCSE Combined Science specification (1SC0). You need to understand the model of metallic bonding — a lattice of positive metal ions surrounded by a sea of delocalised electrons — and be able to explain the typical properties of metals using this model.

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What Is Metallic Bonding?

Metallic bonding is the strong electrostatic attraction between positive metal ions (cations) and a sea of delocalised electrons.

How Metallic Bonding Works

1. Metal atoms are packed closely together in a regular arrangement (a lattice).
2. The outer-shell electrons of each metal atom become delocalised — they detach from their atoms and move freely throughout the entire structure.
3. The metal atoms become positive ions (they have lost their outer electrons).
4. The delocalised electrons form a "sea of electrons" that surrounds all the positive ions.
5. There is a strong electrostatic attraction between the positive ions and the negative delocalised electrons — this holds the structure together.

graph TD
    A["Metallic Bonding"] --> B["Metal atoms lose<br/>outer electrons"]
    B --> C["Positive metal ions<br/>in a lattice"]
    B --> D["Delocalised electrons<br/>(sea of electrons)"]
    C --> E["Strong electrostatic<br/>attraction between<br/>ions and electrons"]
    D --> E

    style A fill:#2c3e50,color:#fff
    style B fill:#2980b9,color:#fff
    style C fill:#27ae60,color:#fff
    style D fill:#e67e22,color:#fff
    style E fill:#c0392b,color:#fff

Exam Tip: The definition of metallic bonding is: "the strong electrostatic attraction between positive metal ions and a sea of delocalised electrons." Learn this word-for-word — it is a common 2-mark definition question.

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Properties of Metals

The properties of metals can be explained using the metallic bonding model.

1. High Melting and Boiling Points

Most metals have high melting points and high boiling points.

Explanation:

• The electrostatic attraction between the positive metal ions and the delocalised electrons is strong.
• A large amount of energy is needed to overcome these strong metallic bonds.
• The more delocalised electrons and the greater the charge on the metal ion, the stronger the metallic bond and the higher the melting point.

Metal	Group	Charge on Ion	Melting Point (°C)
Sodium (Na)	1	+1	98
Magnesium (Mg)	2	+2	650
Aluminium (Al)	3	+3	660
Iron (Fe)	Transition	+2/+3	1538

Exam Tip: Group 1 metals (like sodium and potassium) have relatively low melting points for metals because they only have one delocalised electron per atom, so the metallic bond is weaker. You may be asked to explain this.

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2. Good Electrical Conductivity

Metals are excellent conductors of electricity.

Explanation:

• The delocalised electrons are free to move through the metal structure.
• When a potential difference (voltage) is applied, the delocalised electrons flow in one direction — this is an electric current.

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3. Good Thermal Conductivity

Metals are good conductors of heat.

Explanation:

• When one part of a metal is heated, the delocalised electrons gain kinetic energy.
• These electrons move rapidly through the structure, transferring energy to other parts of the metal.
• The closely packed positive ions also vibrate and pass on energy to neighbouring ions.

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4. Malleability and Ductility

Metals are malleable (can be hammered into shape) and ductile (can be drawn into wires).

Explanation:

• When a force is applied, the layers of positive ions can slide over each other without breaking the metallic bond.
• The delocalised electrons can move and readjust to maintain the attraction, holding the structure together in its new shape.
• Unlike ionic compounds, there is no repulsion between layers because the ions are all positive and the electrons can redistribute.

graph LR
    A["Force applied<br/>to metal"] --> B["Layers of ions<br/>slide over<br/>each other"]
    B --> C["Delocalised electrons<br/>readjust position"]
    C --> D["Metallic bond<br/>maintained →<br/>metal does not<br/>shatter"]

    style A fill:#2c3e50,color:#fff
    style B fill:#2980b9,color:#fff
    style C fill:#27ae60,color:#fff
    style D fill:#e67e22,color:#fff

Exam Tip: Contrast metals with ionic compounds here. Ionic compounds are brittle (shatter) because displacing layers causes like-charged ions to repel. Metals are malleable because the delocalised electrons can redistribute and maintain the bond as layers slide.

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Alloys

An alloy is a mixture of a metal with one or more other elements (usually other metals or carbon).

Why Are Alloys Harder Than Pure Metals?

• In a pure metal, the layers of ions are all the same size and can easily slide over each other.
• In an alloy, the atoms of different elements are different sizes.
• The differently-sized atoms disrupt the regular arrangement of the layers.
• This makes it harder for the layers to slide, so alloys are harder and less malleable than pure metals.

Pure Metal	Alloy	Added Element(s)	Use
Iron	Steel	Carbon	Construction, vehicles
Copper	Bronze	Tin	Statues, medals
Copper	Brass	Zinc	Musical instruments
Gold	9-carat gold	Copper, silver	Jewellery
Aluminium	Duralumin	Copper, magnesium	Aircraft

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Summary

• Metallic bonding is the strong electrostatic attraction between positive metal ions and a sea of delocalised electrons.
• Metals have high melting points because of the strong metallic bonds.
• Metals conduct electricity because delocalised electrons are free to move and carry charge.
• Metals conduct heat because delocalised electrons transfer kinetic energy.
• Metals are malleable and ductile because layers of ions can slide without breaking the metallic bond.
• Alloys are harder than pure metals because differently-sized atoms disrupt the regular layers, preventing them from sliding.

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Worked Example — Explaining Why Magnesium Has a Higher Melting Point Than Sodium

Both are metals with metallic bonding, but: