In simple terms
A friendly intro before the formal notes — no formulas yet.
Cations in a Sea of Electrons
A metal is a giant, orderly array of positive ions sitting in a shared pool of electrons that are free to roam. Nothing owns those electrons and everything is attracted to them at once — and that one idea explains almost every property a metal has.
Picture a car park full of parked cars (the fixed positive cations) with water flooded across the whole surface (the delocalised electrons). The water touches every car and glues the scene together, but it can flow. Tip the ground and the water carries things along (electrical current). Nudge a row of cars sideways and the water simply flows into the new gaps, so nothing shatters (malleability). Try to lift a car right out of the flood and the water pulls it back hard (high melting/boiling point). Deeper water, or cars with more charge, grip harder — that is a stronger metallic bond.
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Metal atoms have low ionisation energies, so each atom releases its valence electrons and becomes a positive ion (cation).
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Those released electrons are delocalised — no longer tied to one atom, but free to move through the whole lattice as a 'sea'.
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The metallic bond is the electrostatic attraction between the fixed cations and the mobile electron sea; it is strong and non-directional.
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Read every property off this picture: mobile electrons give conductivity; a fluid sea lets layers slide (malleability); strong attraction gives high melting points.
Explore the concept
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Step 1
Metal atoms have low ionisation energies, so each atom releases its valence electrons and becomes a positive ion (cation).
Full topic notes
Formal explanation with the rigour you need for the exam.
The electron-sea model
Metal atoms sit on the left of the periodic table and have low ionisation energies, so they lose their outer (valence) electrons readily. In the solid, each atom gives up those electrons and becomes a positive ion, a cation. The released electrons are no longer attached to any one atom: they are delocalised, free to move throughout the entire piece of metal, forming a 'sea' of negative charge that washes over and between the fixed cations. The metallic bond is the strong, non-directional electrostatic attraction between this sea of delocalised electrons and the lattice of cations. Because the attraction is non-directional — pulling equally in all directions rather than along fixed bonds — the structure can be deformed without breaking, a point we will return to for malleability.
Metal atoms lose valence electrons to become a lattice of positive ions (cations).
The lost electrons are delocalised, forming a mobile 'sea of electrons'.
The metallic bond is the electrostatic attraction between the cations and the delocalised electrons.
The bond is strong and NON-DIRECTIONAL — this is what lets metals deform without shattering.
Reading the properties off the model
Almost every characteristic property of a metal is one of two features of the model in disguise: the electrons are MOBILE, and the attraction is STRONG. Name the feature, then link it to the property — that is the structure examiners reward.
Electrical conductivity. The delocalised electrons are mobile. Under a potential difference they drift towards the positive terminal, carrying charge as a current. Metals conduct as solids and when molten, because the electron sea stays mobile in both states.
Thermal conductivity. The same mobile electrons transfer kinetic energy rapidly through the lattice (the vibrating cations help too), so heat spreads quickly.
Malleability and ductility. When a force is applied, layers of cations slide over one another. The electron sea immediately flows to the new positions, keeping the attraction, so the metal changes shape without shattering — hammered into sheets (malleable) or drawn into wire (ductile).
High melting and boiling points. The electrostatic attraction between cations and the electron sea is strong, so a large amount of energy is needed to break the lattice apart. Most metals are solids with high melting points at room temperature.
Lustre. Surface delocalised electrons absorb and re-emit photons across many frequencies, so a fresh metal surface is shiny and reflective.
What makes one metallic bond stronger than another
Not all metallic bonds are equal, and comparing them is a favourite exam task. Three factors decide the strength of the attraction between the cations and the electron sea. First, the charge on the cation: a more highly charged cation attracts the electron sea more strongly. Second, the number of delocalised electrons contributed per atom: more electrons in the sea means more attraction. Third, the ionic radius: a smaller cation lets the electron sea sit closer to the positive nuclei, strengthening the bond. Strong metallic bonding shows up as high melting and boiling points, high density and greater hardness.
Charge of the cation — higher charge (e.g. Mg²⁺ vs Na⁺) means stronger attraction.
Number of delocalised electrons per atom — more electrons in the sea (e.g. 2 for Mg vs 1 for Na) means stronger bonding.
Ionic radius — a smaller cation holds the electron sea closer to the nuclei, strengthening the bond.
Down a group (Na → K): radius increases, so bonding weakens and melting point falls.
Across a period (Na → Mg → Al): charge and electron count rise while radius falls, so bonding strengthens sharply and melting point rises.
Alloys
An alloy is a mixture that contains at least one metal — either two or more metals, or a metal with a non-metal. Alloys are made because they usually improve on the pure metal: harder, stronger, or more corrosion-resistant. Brass is copper mixed with zinc; steel is iron with a little carbon; bronze is copper with tin. The reason alloying works comes straight from the sliding-layers picture. In a pure metal every ion is the same size, so the layers are neat and slide over each other relatively easily. Introduce atoms of a different size and the regular layers are distorted; the layers can no longer slide past one another smoothly, so the alloy resists deformation — it is harder and less malleable than the pure metal.
An alloy is a mixture containing at least one metal.
Alloys are usually harder and stronger than the pure metal.
Atoms of a different size distort the regular lattice layers.
Distorted layers cannot slide over each other easily, so the alloy resists being deformed.
(HL) Transition metals
Transition metals sit at the strong end of the metallic-bonding scale, and this extra detail is HL-only. Unlike a Group 1 or Group 2 metal, a transition metal can delocalise electrons from both the 4s and the inner 3d sub-shell, so it typically contributes more delocalised electrons per atom to the sea. Combined with relatively small ions, this gives especially strong metallic bonding. The consequences are the properties transition metals are known for: high melting and boiling points, high density, and considerable hardness and strength — which is why iron, chromium, titanium and their alloys dominate structural and engineering uses. (SL students are responsible for the model and the three strength factors, but not this 3d/4s explanation.)
Transition metals can delocalise both 4s and 3d electrons, giving more delocalised electrons per atom.
Their ions are relatively small, so the electron sea is held close and strongly.
Result: strong metallic bonding → high melting/boiling points, high density, hardness and strength.
This 3d/4s reasoning is HL only; the general model and strength factors are common to SL and HL.
Common mistakes examiners penalise
Saying the ions carry the current — in a solid metal it is the DELOCALISED ELECTRONS that move and carry charge; the cations stay fixed in the lattice.
Confusing metallic with ionic bonding — metallic bonding has cations ONLY, delocalised over a mobile electron sea; there are no anions and no transferred electrons sitting on a partner ion.
Explaining malleability by 'the bonds break' — the metallic bond does NOT break; layers of cations SLIDE while the electron sea flows and keeps the attraction.
Comparing bond strength with only one idea — a full comparison should reference the cation charge, the number of delocalised electrons and/or the ionic radius, then link to a stronger attraction.
Forgetting that stronger attraction ⇒ more energy ⇒ higher melting point — state the whole chain; stopping at 'stronger bonds' without linking to energy or melting point can drop the conclusion mark.
Describing an alloy as adding electrons or forming new bonds — alloying works by DISTORTING the layers with different-sized atoms so they cannot slide, not by changing the bonding type.
Using 3d/4s transition-metal reasoning in an SL answer as if required — it is correct chemistry but HL-only detail; SL answers are marked on the model and the three strength factors.
Model answer — marked the way our engine marks it
On Paper 2 an 'explain' question like this is marked analytically (points-based): each distinct, valid chemical point is worth one mark, and the marks are independent of one another. There is no single 'right wording' — the engine credits any equivalent statement that makes the point, and error-carried-forward (ECF) means a reasoning mark you have earned is not taken away because a later conclusion slips. In a structured answer, method/reasoning marks (M) are earned for correct explanatory steps and answer/conclusion marks (A) for the final linked statement. Study how each mark below is tied to one specific point.
Where this leads
The metallic model completes the trio of bonding types — ionic, covalent and metallic — and the same habit of thought runs through all of them: describe the particles and the forces between them, then read the properties off that picture. In later topics the bonding continuum (S2.4) places metallic bonding alongside ionic and covalent on a triangle of bonding character, and the reactivity of metals in redox and electrochemistry (Reactivity 3) builds directly on how readily metals lose those valence electrons in the first place. Own the electron-sea model here and those topics become extensions of a picture you already carry.
Worked examples
See the formulas applied — reveal one step at a time, like the exam.
Explain, in terms of the metallic bonding model, why copper is used for electrical wiring and can also be bent into shape without snapping. [4]
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Copper consists of a lattice of Cu cations surrounded by a sea of delocalised electrons. [1]
Sodium melts at 98 °C and magnesium at 650 °C. Explain, in terms of metallic bonding, why magnesium has the much higher melting point. [3]
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In sodium each atom loses one electron to form Na⁺ and contributes one delocalised electron; in magnesium each atom loses two to form Mg²⁺ and contributes two delocalised electrons. [1]
Brass is an alloy of copper and zinc. Explain, in terms of structure, why brass is harder than pure copper. [3]
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In pure copper the cations are all the same size and form regular layers that can slide over one another fairly easily (so pure copper is soft/malleable). [1]
Explain, in terms of metallic bonding, why magnesium has a higher melting point than sodium and why both conduct electricity. [4]
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Model answer — full response.
How it all connects
The big idea sits in the middle — tap a linked idea to explore the link.
Tap a linked idea to see how it connects back to the main topic — that connection is what examiners reward.
Glossary
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Quick check
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Revision flashcards
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Metallic bond
The electrostatic attraction between a lattice of positive metal ions (cations) and a 'sea' of delocalised valence electrons. It is strong and non-directional.
Key takeaways
Review these before you close the topic — retrieval beats re-reading.
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Metal atoms lose valence electrons to become a lattice of positive ions (cations).
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The lost electrons are delocalised, forming a mobile 'sea of electrons'.
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The metallic bond is the electrostatic attraction between the cations and the delocalised electrons.
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The bond is strong and NON-DIRECTIONAL — this is what lets metals deform without shattering.
Practice — then mark it
The whole point: a real Cambridge question, marked mark-by-mark.
Get a Paper 2 explanation marked: compare metallic bond strength and explain conductivity with full reasoning
Get a Paper 2 explanation marked: compare metallic bond strength and explain conductivity with full reasoning
Extra simulations & links
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Frequently asked
Checkpoint
One marked question is worth ten re-reads — close the loop before you move on.
Reading it isn’t knowing it — prove it.
Before you move on: do Get a Paper 2 explanation marked: compare metallic bond strength and explain conductivity with full reasoning on paper, snap a photo, and get examiner-style feedback on exactly where you win and lose marks.