Metallic Bonding and Giant Structures

This lesson covers metallic bonding, the properties of metals and alloys, and the giant covalent structures of diamond, graphite, graphene and silicon dioxide. Understanding these structures is central to AQA specification 3.1.3.

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Metallic Bonding

Key Definition: A metallic bond is the strong electrostatic attraction between the positive metal ions (cations) and the delocalised electrons (the "sea" of electrons).

In a metal:

• Metal atoms lose their outer-shell electrons to form positive ions (cations)
• These electrons become delocalised — they are not associated with any particular metal ion
• The delocalised electrons form a "sea" or "cloud" of electrons that surrounds the cations
• The attraction between the cations and the sea of delocalised electrons holds the structure together

    +   e⁻  +   e⁻  +
   e⁻  +   e⁻  +   e⁻
    +   e⁻  +   e⁻  +
   e⁻  +   e⁻  +   e⁻
    +   e⁻  +   e⁻  +

Factors Affecting Metallic Bond Strength

The strength of metallic bonding increases with:

Factor	Reason	Example
Higher charge on cation	Greater electrostatic attraction between cations and electrons	Mg²⁺ stronger than Na⁺
Smaller ionic radius	Delocalised electrons are closer to the nucleus	Na stronger than K
More delocalised electrons	Greater attraction	Al (3 e⁻) stronger than Mg (2 e⁻)

Evidence: Melting Points of Period 3 Metals

Metal	No. of delocalised e⁻	Ionic radius (pm)	Melting Point (°C)
Na	1	95	98
Mg	2	65	649
Al	3	50	660

Melting point increases from Na to Al because: more delocalised electrons, higher cation charge, and smaller cation radius all strengthen the metallic bonding.

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Properties of Metals

Property	Explanation
High melting and boiling points	Strong metallic bonding requires a large amount of energy to overcome
Good electrical conductivity	Delocalised electrons are free to move through the structure and carry charge
Good thermal conductivity	Delocalised electrons can transfer kinetic energy rapidly through the structure
Malleable (can be hammered into shape)	Layers of cations can slide over each other without breaking the metallic bond because the delocalised electrons can move with them
Ductile (can be drawn into wires)	Same as malleability — layers slide, and the sea of electrons adjusts
Shiny (lustrous)	Delocalised electrons can absorb and re-emit light of many wavelengths

Key Contrast with Ionic Compounds: Metals are malleable and ductile because the delocalised electrons can adjust as layers slide. Ionic compounds are brittle because displacing layers brings like-charged ions together, causing repulsion.

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Alloys

Key Definition: An alloy is a mixture of two or more elements, at least one of which is a metal, with metallic properties.

Why Alloys Are Harder Than Pure Metals

In a pure metal, the layers of identically sized cations can slide easily over each other. In an alloy, atoms of a different size disrupt the regular arrangement, preventing the layers from sliding.

Pure metal:           Alloy:
○ ○ ○ ○ ○           ○ ○ ○ ○ ○
 ○ ○ ○ ○ ○           ○ ● ○ ○ ○    (● = different-sized atom)
○ ○ ○ ○ ○           ○ ○ ○ ● ○
 ○ ○ ○ ○ ○           ○ ○ ○ ○ ○
Layers slide easily   Layers cannot slide

Examples:

• Steel = iron + carbon (and sometimes other metals). Carbon atoms are smaller than iron atoms and fit between the layers, preventing sliding.
• Brass = copper + zinc. Zinc atoms are a different size to copper atoms.
• Bronze = copper + tin.

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Giant Covalent Structures

Giant covalent structures (also called macromolecular structures) contain many atoms bonded together by strong covalent bonds in a continuous network. They have very high melting points because many strong covalent bonds must be broken.

Diamond

Structure:

• Each carbon atom is bonded to 4 other carbon atoms in a tetrahedral arrangement
• All bonds are strong C—C single covalent bonds (bond enthalpy 347 kJ mol⁻¹)
• The structure extends in three dimensions as a giant lattice
• Bond angle: 109.5° (tetrahedral)