The big idea: Metallic bonding is the electrostatic attraction between a lattice of positive metal cations and a 'sea' of delocalised electrons.
Each metal atom loses its outer (valence) electrons. Those electrons are no longer tied to one atom — they are delocalised, free to move throughout the whole structure. What is left behind is a regular lattice of cations, held together by their attraction to the shared electron sea.
Fixed metal cations (Na⁺) sit in a regular lattice; the valence electrons are delocalised into a 'sea' that flows between them.
Interactive diagram
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What is delocalised?: Delocalised means the electrons are not fixed to one atom — they are spread out and free to move through the whole lattice.
It is this attraction between the fixed positive cations and the mobile negative sea that is the metallic bond.
Every typical property of a metal can be traced back to the model: a lattice of cations in a sea of mobile, delocalised electrons. Always explain a property by saying what the electrons or cations do.
Property → explanation
- Electrical conductivity — the delocalised electrons are free to move, so they carry charge through the metal when a voltage is applied. Metals conduct as a solid and when molten.
- Thermal conductivity — the mobile electrons (and vibrating cations) transfer kinetic energy quickly through the lattice.
- Malleability and ductility — the bonding is non-directional, so layers of cations can slide over one another. The electron sea simply flows around them and keeps the lattice held together, so the metal bends instead of shattering.
- High melting point — the strong electrostatic attraction between cations and the electron sea takes a lot of energy to overcome.
- Lustre (shiny) and high density — the delocalised electrons reflect light, and the cations pack closely together.
Why metals don't shatter: In an ionic lattice, sliding the layers brings like charges next to each other; they repel and the crystal cracks — so ionic solids are brittle.
In a metal, the electron sea is non-directional, so when layers slide there is no layer of like charges to repel. The bonding re-forms in the new position, so the metal deforms without breaking.
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Metallic bond strength (and so the melting point and hardness) depends on how strongly the cations attract the electron sea. Two factors decide this:
Cation charge
- higher charge → stronger bonding
- more charge attracts the sea more strongly
- each atom also donates more electrons → denser sea
- Na⁺ (1+) < Mg²⁺ (2+) < Al³⁺ (3+)
Cation radius
- smaller radius → stronger bonding
- the sea sits closer to the nucleus
- down a group radius grows → bonding weakens
- Li > Na > K (Li strongest)
Putting it together
- small, highly-charged cation = strongest
- across period 3: Na < Mg < Al
- more delocalised electrons also helps
Worked example — comparing two metals
Explain why magnesium has stronger metallic bonding than sodium.
Solution
- Charge: sodium forms Na⁺ but magnesium forms Mg²⁺, so each Mg cation has a higher (2+) charge and attracts the electron sea more strongly.
- More electrons: each magnesium atom donates two electrons to the sea (sodium donates one), so the electron sea is denser.
- Radius: the Mg²⁺ cation is also smaller than Na⁺, so the sea sits closer. All three effects make the bonding stronger — hence magnesium's higher melting point.
Final answer
Mg²⁺ has a higher charge, donates more electrons and is smaller than Na⁺, so it attracts the delocalised electron sea more strongly.
Magnesium loses two electrons per atom, so the sea is denser (2 e⁻ per cation) and the 2+ cations pull on it harder — stronger metallic bonding than sodium.
Interactive diagram
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How this is tested: Metallic bonding shows up as a short explain question on Paper 2 and as Paper 1A MCQs.
- 'Explain why metals conduct electricity / are malleable' — every mark needs you to name delocalised electrons and say what they do. - 'Which metal has the strongest metallic bonding?' — decide using cation charge (higher = stronger) and radius (smaller = stronger).
Markers want the cause linked to the property, never just 'because of the structure'.
Common trap: Don't say metals conduct because 'ions move'. In a solid metal the cations stay fixed — it is the delocalised electrons that move and carry the charge. (Contrast this with a molten ionic compound, where the ions are the carriers.)
IB-style question — conductivity and malleability (a)
(a) Using the metallic bonding model, explain why aluminium is both a good electrical conductor and malleable. [3]
How to score the marks
- Mark 1 — the model. Aluminium is a lattice of Al³⁺ cations in a sea of delocalised electrons.
- Mark 2 — conductivity. The delocalised electrons are free to move, so they carry charge (current) through the metal.
- Mark 3 — malleability. The bonding is non-directional, so layers of cations can slide over one another (the electron sea keeps the lattice bonded), so the metal bends rather than shatters.
Final answer
A lattice of Al³⁺ cations in a sea of delocalised electrons; the mobile delocalised electrons carry charge (conductivity); layers of cations slide over each other because the bonding is non-directional (malleability).
IB-style question — comparing bond strength (b)
(b) Predict, with a reason, whether calcium or potassium has the stronger metallic bonding. [2]
How to score the marks
- Mark 1 — the prediction. Calcium has the stronger metallic bonding.
- Mark 2 — the reason. Calcium forms a 2+ cation (potassium only 1+) and the Ca²⁺ ion is smaller than K⁺, so it attracts the delocalised electron sea more strongly.
Final answer
Calcium — it forms a higher-charged (2+) and smaller cation than potassium, so it attracts the electron sea more strongly.