Metallic Bonding

Metallic bonding is the third of the three types of strong chemical bond. Which type you get depends entirely on the elements involved.

Bond type	Between	How the outer electrons behave
Ionic	Metal + non-metal	Transferred, forming charged ions
Covalent	Non-metal + non-metal	Shared in pairs between atoms
Metallic	Metallic elements and alloys	Delocalised across the whole structure

Metallic bonding occurs in metallic elements (such as copper, iron and sodium) and in alloys, which are mixtures containing metals.

All three bonds are held together by electrostatic attraction between positive and negative charges. In a metal, the positive charges are the metal ions and the negative charge is the sea of delocalised electrons.

The key idea, developed over the next slides, is that the outer electrons in a metal do not stay with their own atoms.

A metal is a giant structure: a regular, repeating 3D arrangement of atoms that extends throughout the whole piece of metal. This is why a lump of copper has no fixed number of atoms in a "molecule" — the structure just keeps repeating.

Within this structure, each metal atom loses its outer-shell electrons. These electrons leave the individual atoms and become delocalised, meaning they are free to move throughout the whole structure and are no longer held by any single atom.

Losing outer electrons turns each atom into a positive ion. So the structure becomes a regular array of positive metal ions, and the released electrons spread out among them.

"Delocalised" means the electron is not fixed to one atom. In a metal, the outer electrons belong to the whole structure, not to any one ion.

Because every atom donates its outer electrons, there is a large number of delocalised electrons shared across billions of positive ions.

The delocalised electrons form what is often described as a sea of electrons surrounding the positive metal ions. The arrangement is best pictured as rows of positive ions sitting in this shared sea.

The metallic bond is the strong electrostatic attraction between the positive metal ions and the negatively charged sea of delocalised electrons. Because this attraction pulls the ions and electrons together throughout the whole structure, metallic bonds are strong.

A simple way to picture it:

• Positive metal ions sit in fixed positions in a regular pattern.
• Delocalised electrons (one or more per atom) drift freely between them.
• The oppositely charged ions and electrons attract, holding the whole structure together.

Metallic bond = the electrostatic attraction between the positive ions and the sea of delocalised electrons. State both the ions and the delocalised electrons to score the mark.

The attraction acts in all directions throughout the giant structure, much like the forces in an ionic lattice, but here one of the "charges" is a mobile sea of electrons rather than fixed negative ions.

The spec asks you to recognise metallic structures from diagrams. A metallic bonding diagram has a distinctive look that separates it from ionic and covalent diagrams.

Look for these features:

• A regular grid of circles all labelled as positive ions (often written with a + sign).
• Small dots or a shaded region between the ions, labelled as delocalised electrons or a "sea of electrons".
• Every ion the same (a pure metal) or a mixture of sizes (an alloy).

You see...	Bonding type
Positive ions in a sea of loose electrons	Metallic
Alternating positive and negative ions in a lattice	Ionic
Atoms sharing pairs of electrons (dots and crosses between two atoms)	Covalent

The give-away for metallic bonding is loose delocalised electrons drawn between rows of positive ions. Ionic diagrams show fixed negative ions, not free electrons.

If the diagram shows charged ions but the negative charge is a shared cloud rather than separate ions, it is metallic.

The metallic bonding model is not just a picture — it explains what metals are actually like. Each property traces back to the same two features: strong bonding, and mobile delocalised electrons. (The properties themselves are covered in full in the metals and alloys lesson; here they show why the model matters.)

Property of metals	Explained by
High melting and boiling points	Strong electrostatic attraction between ions and delocalised electrons needs lots of energy to overcome
Good electrical conductivity	Delocalised electrons are free to move and carry charge through the metal
Good thermal conductivity	Delocalised electrons transfer energy quickly through the structure

The single feature doing most of the work is the delocalised electrons. Because they are free to move, they can both carry an electric current and transfer thermal energy, and because they bind strongly to the positive ions, they give metals high melting points.

One model, one set of features: positive ions plus a mobile electron sea. Almost every metallic property comes back to it.

1. Saying metals are held together by "shared electrons" like covalent bonds

Metallic bonding is not the same as covalent bonding. In a metal the electrons are delocalised across the whole structure, not shared as pairs between two specific atoms.

2. Forgetting to mention the delocalised electrons

A metallic bond is the attraction between positive ions and delocalised electrons. Writing only "attraction between positive ions" is incomplete and describes nothing that would hold them together.

3. Calling the particles "atoms" instead of "ions"

Once the outer electrons are delocalised, the particles left are positive ions, not neutral atoms. Use "positive metal ions" when describing the structure.

4. Confusing the metallic sea with negative ions

The mobile negative charge in a metal is a sea of loose electrons, not separate negative ions. Only ionic compounds contain fixed negative ions.

5. Not linking properties back to the model

An "explain" question wants the reason, not just the fact. Link conductivity to the free-moving delocalised electrons and high melting point to the strong electrostatic attraction.