Metallic Bonding

Summary: Metallic bonding is the electrostatic attraction between a lattice of positive metal ions and a sea of delocalised electrons, giving metals their characteristic properties: electrical and thermal conductivity, malleability, ductility, and high melting points. Alloys are harder because different-sized atoms disrupt the regular layers. Tags: igcse chemistry bonding Created: 2026-07-14 Last Updated: 2026-07-18

Metallic bonding is the model chemists use to explain how metal atoms are held together in the solid and liquid states, and it underpins nearly every characteristic property of metals. In a metal, each atom releases its outer-shell electrons into a shared pool, leaving behind positive ions packed into a regular, repeating Giant Metallic Lattice. Each freed Delocalised Electron no longer belongs to any single atom; instead, it moves freely throughout the whole structure, forming the famous “sea of electrons”. The strong Electrostatic Attraction between these mobile negative electrons and the fixed positive ions is the metallic bond itself, and because it acts in all directions it is described as non-directional. This bonding model directly explains why metals conduct electricity and heat, why they show Malleability when hammered and Ductility when drawn into wires, and why alloying changes their hardness. Understanding metallic bonding is therefore essential for explaining the many practical Uses of Metals, from copper wiring to steel construction and aluminium aircraft alloys.

What Is Metallic Bonding?

Definition: Metallic bonding is the electrostatic attraction between positive metal ions (cations) and the sea of delocalised electrons that surround them.

Each metal atom in the lattice loses its outer-shell electrons, becoming a positive ion. These outer-shell electrons become delocalised — they are no longer attached to any one particular atom but are free to move throughout the entire metallic structure.

The structure is described as:

“A giant metallic lattice of positive metal ions arranged in a regular pattern, surrounded by a sea of delocalised electrons.”

Why do metal atoms lose their outer electrons? Metal atoms have few outer-shell electrons (typically 1, 2, or 3). They have low ionisation energies, meaning they readily lose these electrons. The released electrons move freely among the resulting positive ions.

Structure of a Giant Metallic Lattice

Key features of the giant metallic lattice:

• Positive ions are arranged in a regular, closely-packed, repeating pattern (like spheres packed as tightly as possible).
• Delocalised electrons move freely throughout the entire structure, not bound to any specific ion.
• The electrostatic attraction between the positive ions and the negative delocalised electrons holds the lattice together — this is the metallic bond.
• Metallic bonding is non-directional — the attraction acts in all directions, unlike covalent bonds which are between specific pairs of atoms.

Comparison of bonding types:

Why Metals Conduct Electricity

Metals are excellent electrical conductors.

Explanation:

• The delocalised electrons are free to move throughout the lattice.
• When a potential difference (voltage) is applied across a metal, the delocalised electrons drift towards the positive terminal, creating an electric current.
• Because the electrons can move freely throughout the entire structure, the current flows easily.

Comparing with ionic compounds:

• Ionic compounds conduct electricity only when molten or in aqueous solution — because the ions are free to move.
• Ionic compounds do NOT conduct when solid — ions are locked in fixed positions.
• Metals conduct in ALL states (solid or liquid) because the delocalised electrons are always free to move.

Best conductors: Silver > Copper > Gold > Aluminium. Copper is used for most wiring because it is nearly as good as silver but much cheaper.

Why Metals Conduct Heat

Metals are excellent thermal conductors (they feel cold to the touch because they conduct heat away from your hand quickly).

Explanation:

• When a metal is heated at one end, the delocalised electrons near the heat source gain kinetic energy.
• These fast-moving electrons diffuse throughout the lattice, transferring kinetic energy to other electrons and to positive ions via collisions.
• This transfers thermal energy through the metal much faster than in non-metals, where only lattice vibrations (phonons) can transfer heat.

Why Metals Are Malleable and Ductile

• Malleable: can be hammered or pressed into shape (e.g., aluminium foil, gold leaf).
• Ductile: can be drawn into wires (e.g., copper wiring).

Explanation using the lattice model:

In the giant metallic lattice, the positive ions are arranged in regular layers (close-packed planes).

• When a force is applied, the layers of ions slide over each other.
• As the layers shift, the delocalised electrons move with the ions.
• The metallic bonding is maintained throughout the sliding because the delocalised electrons are non-directional — they continue to attract the ions regardless of their new positions.
• No bonds are broken; the metal simply changes shape without fracturing.

Contrast with ionic compounds:

• Ionic compounds are brittle, not malleable.
• When layers of ions slide in an ionic lattice, like-charged ions become adjacent (+ next to +, or - next to -).
• Strong repulsion between like charges causes the crystal to shatter.

Why Alloys Are Harder Than Pure Metals

In a pure metal: All atoms are the same size. The layers of ions are completely regular and can slide over each other easily when a force is applied. This makes pure metals soft and malleable.

In an alloy: An alloy contains atoms of different sizes (e.g., carbon atoms in iron, zinc atoms in copper).

• The different-sized atoms disrupt the regular arrangement of the layers.
• When a force tries to slide the layers, the disrupting atoms act as barriers.
• The layers can no longer slide easily past each other — they lock together.
• Therefore the alloy is harder and stronger than the pure metal.

This is why:

• Pure iron is too soft for most uses — adding carbon makes steel.
• Pure gold (24 carat) is too soft for jewellery — adding copper makes it harder (18 carat gold).
• Pure aluminium is quite weak — alloying with Cu, Mg, Mn makes duralumin strong enough for aircraft.

Why Melting Points Vary Between Metals

The melting point of a metal depends on the strength of the metallic bond, which in turn depends on two factors:

Factor 1: Number of delocalised electrons per atom

• More delocalised electrons = stronger electrostatic attraction between ions and electrons = stronger metallic bond = higher melting point.
• Group 1 metals (Li, Na, K): only 1 delocalised electron per atom → relatively low melting points.
  • Na mp: 98 degrees C; K mp: 63 degrees C.
• Group 2 metals (Mg, Ca): 2 delocalised electrons per atom → higher melting points than Group 1.
  • Mg mp: 650 degrees C.
• Group 3 metals: Al has 3 delocalised electrons per atom → higher still.
  • Al mp: 660 degrees C.
• Transition metals (Fe, Cu, Cr): can lose electrons from both the outer s and inner d subshells, contributing multiple delocalised electrons → very high melting points.
  • Fe mp: 1538 degrees C.

Factor 2: Size of the metal ion (charge density)

• Smaller ions with higher charge → greater charge density → stronger electrostatic attraction → higher melting point.
• This explains why Group 1 melting points decrease DOWN the group: Li+ has a smaller ionic radius than Na+, so the attraction to delocalised electrons is stronger in Li than in Na.
  • Li mp: 181 degrees C > Na mp: 98 degrees C > K mp: 63 degrees C.

Comparison of Properties Explained by Bonding

Key Points

• Metallic bonding = electrostatic attraction between positive metal ions and delocalised electrons
• Giant metallic lattice: regular arrangement of positive ions in a “sea of delocalised electrons”
• Metals conduct electricity: delocalised electrons are free to move through the lattice
• Metals conduct heat: delocalised electrons transfer kinetic energy
• Metals are malleable/ductile: layers of ions slide over each other; delocalised electrons move with them; metallic bonding is maintained
• Alloys are harder: different-sized atoms disrupt the regular layers, preventing easy sliding
• Higher melting points: more delocalised electrons per atom AND/OR smaller ions = stronger metallic bond
• Group 1 metals: low mp (1 delocalised electron); transition metals: very high mp (multiple delocalised electrons)
• Metallic bonding is non-directional (unlike covalent bonding)

Key Concepts from Past Papers

• Metallic bonding: the electrostatic attraction between positive metal ions and the sea of delocalised electrons
• Alloy: a mixture of a metal with one or more other elements (usually other metals or carbon)
• Delocalised electrons: electrons that are not associated with any particular atom and are free to move throughout the structure
• Delocalised electrons can move freely throughout the lattice
• Metals are malleable because layers of ions can slide over each other and the delocalised electrons move with them
• An alloy is harder because different-sized atoms disrupt the layers, preventing them from sliding

Keywords from Past Papers

graphite, reacts, calcium, oxide, diamond, platinum, pure, giant, covalent, alloy, down, forms, silicon, silica, heat

Related Notes

• Atomic Structure
• Ions and Ionic Bonds
• Covalent Bonding
• Giant Structures
• Giant Metallic Lattice
• Delocalised Electron
• Electrostatic Attraction
• Malleability
• Ductility
• Uses of Metals
• Cambridge IGCSE Chemistry 0620 Syllabus (2023-2025)

Electron Sea Model

The electron sea model (also called the delocalised electron model) describes metallic bonding as a regular lattice of positive metal ions immersed in a ‘sea’ of delocalised electrons — the outer-shell electrons that are free to move throughout the entire metallic structure. The metallic bond is the electrostatic attraction between the positive ion cores and this mobile sea of negative electrons. This model explains all the characteristic properties of metals:

• Electrical conductivity: delocalised electrons can drift through the lattice carrying charge
• Thermal conductivity: electrons transfer kinetic energy rapidly through the structure
• Malleability and ductility: layers of ions can slide without breaking the metallic bond because the electron sea readjusts around the moving ions

Exam point: Metallic bonding = electrostatic attraction between positive metal ions in a lattice AND a sea of delocalised electrons.

Sources

• OpenStax Chemistry 2e — Chapter 10: Liquids and Solids, Rice University (free, CC BY 4.0)
• BBC Bitesize GCSE Chemistry — Metallic Bonding, BBC (free educational resource)
• Cambridge IGCSE Chemistry 0620 — Syllabus Section 2.7, Cambridge Assessment International Education
• CK-12 Chemistry for High School — Metallic Bonding, CK-12 Foundation (free, CC BY-NC 3.0)

Past Paper Sources

• 0620/32 Feb/March 2017: Q22(b)(iii) (1m), Q22(d)(i) (1m)
• 0620/32 Feb/March 2018: Q55(d)(ii) (1m)
• 0620/32 Feb/March 2020: Q22(b)(ii) (3m)
• 0620/32 Feb/March 2021: Q77(b)(i) (1m)
• 0620/32 Feb/March 2022: Q66(c)(i) (1m)
• 0620/32 May/June 2020: Q55(b)(ii) (1m)
• 0620/33 May/June 2021: Q88(b)(iii) (1m)
• 0620/33 May/June 2024: Q33(a)(ii) (1m)
• 0620/33 October/November 2015: Q44(b)(i) (2m)
• 0620/33 October/November 2017: Q44(f)(i) (0m)
• 0620/33 October/November 2023: Q55(c)(ii) (1m)
• 0971/31 May/June 2021: Q44(c)(i) (2m)

Common Misconceptions