Metallic Bonding and Properties of Metals
Metallic bonding explains every physical property of a metal, from why wires conduct to why metals can be hammered into shape. The definition uses one phrase that must appear: delocalised electrons.
1. The structure
A metal is a GIANT LATTICE of POSITIVE METAL IONS arranged in layers, surrounded by a “SEA” of DELOCALISED ELECTRONS.
How it forms: each metal atom loses its outer-shell electrons, becoming a positive ion. Those electrons are no longer attached to any one atom — they are delocalised and free to move throughout the structure.
A metallic bond is the strong ELECTROSTATIC ATTRACTION between the POSITIVE METAL IONS and the DELOCALISED ELECTRONS.
“Delocalised” is the mark-earning word. Say positive ions and delocalised electrons — “free electrons” is usually accepted, but “delocalised” is safest.
The lattice contains IONS, not atoms — the atoms have lost electrons.
2. Explaining the properties
Every property comes from either the delocalised electrons or the layers of ions.
Good electrical conductivity
The delocalised electrons are FREE TO MOVE through the structure and carry the charge.
Metals conduct as solids — unlike ionic compounds, which need to be molten or aqueous. The difference is that in a metal it is electrons that move; in an ionic compound it is ions.
Good thermal conductivity
The delocalised electrons move and transfer energy quickly through the metal. The vibrating ions also pass energy along.
Malleable and ductile
The LAYERS of ions can SLIDE over each other when a force is applied, without breaking the metallic bonding — because the delocalised electrons continue to hold the structure together wherever the ions move to.
Malleable — can be hammered into sheets. Ductile — can be drawn into wires.
This contrasts sharply with ionic compounds, which are brittle because sliding layers bring like charges together, which repel. In a metal, the electron sea means there are no like-charge repulsions to worry about.
High melting and boiling points
There are strong electrostatic attractions between the positive ions and the delocalised electrons throughout the lattice, requiring a lot of energy to overcome.
Other properties
- Shiny (lustrous) when polished
- High density — the ions are closely packed
- Sonorous — they ring when struck
3. Alloys
An ALLOY is a mixture of a metal with one or more other elements.
Why alloys are harder than pure metals:
The added atoms are a DIFFERENT SIZE, which DISRUPTS the regular arrangement of the layers. The layers can no longer slide over each other easily, so the alloy is harder and stronger.
Common alloys:
| Alloy | Contains | Used for |
|---|---|---|
| Brass | copper + zinc | instruments, fittings |
| Bronze | copper + tin | statues, bearings |
| Steel | iron + carbon | construction |
| Stainless steel | iron + chromium + nickel | cutlery, corrosion resistance |
Brass is copper and zinc. Confusing which metal is in brass was recorded.
The explanation for hardness is about DISRUPTED LAYERS — say that the different-sized atoms stop the layers sliding. Simply saying “it’s a mixture so it’s stronger” earns nothing.
4. Comparing the structures
| Metallic | Ionic | Simple molecular | |
|---|---|---|---|
| Particles | positive ions + delocalised electrons | positive and negative ions | molecules |
| Conducts as a solid | YES | no | no |
| Conducts when molten | yes | yes | no |
| Charge carrier | electrons | ions | none |
| Malleable or brittle | malleable | brittle | soft/weak |
| Melting point | high | high | low |
Metals conduct as solids; ionic compounds do not. This single contrast is the most useful way to identify a substance from its properties.
5. Mistakes that cost marks
Omitting “delocalised” from the definition.
Saying the lattice contains atoms rather than positive ions.
Explaining conductivity without saying the electrons MOVE.
Saying ions carry the charge in a metal.
Explaining malleability without “layers slide”.
Saying alloys are harder “because they’re mixtures” without the disrupted-layers reason.
Confusing which metals are in brass or bronze.
Saying metals are brittle — that is ionic compounds.
Frequently asked questions
What is metallic bonding? The electrostatic attraction between positive metal ions and a sea of delocalised electrons.
What is in the lattice? Positive metal ions — the atoms have lost their outer electrons.
Why do metals conduct electricity? The delocalised electrons are free to move and carry charge.
Do metals conduct as solids? Yes — unlike ionic compounds.
Why are metals malleable? The layers of ions slide over each other while the delocalised electrons keep the structure bonded.
Why do metals have high melting points? Strong electrostatic attraction between ions and delocalised electrons throughout the lattice.
What is an alloy? A mixture of a metal with other elements.
Why are alloys harder? Different-sized atoms disrupt the layers, so they cannot slide easily.
What is brass? Copper and zinc.
What carries the charge in a metal compared with an ionic compound? Electrons in a metal; ions in a molten or aqueous ionic compound.
Quick revision checklist
- I can define metallic bonding using positive ions and delocalised electrons
- I know the lattice contains ions, not atoms
- I can explain electrical conductivity via moving electrons
- I can explain thermal conductivity
- I can explain malleability via sliding layers
- I know why metals aren’t brittle but ionic compounds are
- I can explain high melting points
- I can define an alloy
- I can explain why alloys are harder with the disrupted-layers reason
- I know the composition of brass, bronze and steel
- I can compare metallic, ionic and simple molecular structures
These notes cover metallic bonding and the properties of metals in the Cambridge IGCSE Chemistry (0620) syllabus and are written for Grade 9–11 / Year 10–11 students. They are based on teaching patterns observed across a large set of one-to-one IGCSE Chemistry lessons, with particular attention to the errors students make most often and the wording examiners reward. Always check the current syllabus for your own exam series.
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