Last updated: 30 September 2026
Copper can be drawn into long electrical wires, aluminium can be rolled into thin sheets, and steel can support enormous structures. These uses seem very different, yet they all begin with the same microscopic question: what is happening to the particles inside a metal?
Metallic bonding gives us a particle model that answers this question. Once you can picture the positive metal ions and the mobile electrons between them, properties such as electrical conductivity, high melting point and malleability no longer need to be memorised as separate facts.
The most useful habit in this topic is therefore to connect structure → particle behaviour → property. That same chain of reasoning will also explain why an alloy can behave differently from the pure metal from which it is made.
Try these five questions before reading the chapter. They check the particle ideas you will need when explaining metals and alloys.
- 5 short questions
- Instant feedback for every option
- Spot misconceptions before you begin
The Giant Metallic Structure
A piece of metal is not made of separate molecules. Instead, it forms one enormous, continuously bonded structure called a giant metallic lattice.
To understand this structure, begin with the outer electrons of the metal atoms. These electrons become separated from individual atoms and are able to move throughout the entire metallic lattice. We describe them as delocalised electrons.
Once the outer electrons have become delocalised, the particles left in the regular lattice are positively charged. At O-Level, we describe them as positive metal ions.
It consists of a regular giant lattice of positive metal ions surrounded by delocalised electrons that can move throughout the structure.
Why do we say positive metal ions rather than metal atoms?
A neutral atom contains equal amounts of positive and negative charge. In the metallic model, however, the outer electrons are no longer associated with one particular atom. They have become delocalised through the lattice.
The remaining particle therefore has more protons than electrons associated with it and is represented as a positive metal ion.
Before considering metallic bonding: we may begin with neutral metal atoms.
Inside the metallic lattice: describe the structure as positive metal ions surrounded by delocalised electrons.
What exactly does “delocalised” mean?
A localised electron is associated with a particular atom or bond. A delocalised electron is not restricted to one particular metal ion. It belongs to the structure as a whole and can move through the lattice.
Students sometimes say that the electrons are “free”. This can be a useful mental picture, but delocalised is the more precise word. The electrons have not escaped from the metal; they remain inside it and move throughout the metallic structure.
Metallic Bonding: What Holds the Metal Together?
The metallic structure contains oppositely charged particles. The metal ions are positive, while the delocalised electrons are negative. Opposite charges attract one another.
This electrical attraction extends throughout the giant structure and holds the lattice together.
Metallic bonding is the strong electrostatic attraction between positive metal ions and delocalised electrons.
Why is the attraction called electrostatic?
Electrostatic attraction means attraction between particles carrying opposite electrical charges. The positively charged metal ions are strongly attracted to the negatively charged electrons surrounding them.
It is important not to imagine one metallic bond as a single line between two neighbouring particles. The attraction acts throughout the giant lattice. Each positive ion experiences attraction to the delocalised electrons around it.
Metallic structure describes how the particles are arranged: positive metal ions surrounded by delocalised electrons.
Metallic bonding describes the force holding those particles together: electrostatic attraction between the positive ions and the electrons.
How Metallic Bonding Explains the Properties of Metals
In examination questions, stating that a metal has a particular property is usually only the beginning. The important marks come from explaining why the metallic structure produces that property.
Describe the relevant particles first, explain what those particles can do, and then connect that behaviour to the physical property being asked about.
Why do metals usually have high melting and boiling points?
The positive metal ions are strongly attracted to the surrounding delocalised electrons. When a metal is heated, thermal energy is supplied to the structure.
A large amount of energy is generally needed to overcome the strong electrostatic attractions in the giant metallic lattice. This is why most metals have relatively high melting and boiling points.
The key phrase is not simply “strong bonds”. A stronger answer identifies what is attracting what: positive metal ions are strongly attracted to delocalised electrons.
Why do metals conduct electricity?
An electric current requires charged particles that are able to move. Metals contain two kinds of charged particles: positive metal ions and negatively charged delocalised electrons. However, only the delocalised electrons are able to move throughout the solid metallic structure.
When a potential difference is applied across a metal, the delocalised electrons move through the structure and carry electrical charge.
Metals conduct electricity because they contain mobile delocalised electrons that can move through the structure and carry electrical charge.
The positive metal ions are also charged, but in a solid metal they are held in fixed positions in the lattice. They may vibrate about these positions, but they cannot move from one part of the metal to another to carry electrical charge.
Why do the electrons conduct but the positive metal ions do not?
Delocalised electrons: they are negatively charged and mobile. When a potential difference is applied, they can move through the metallic structure and carry electrical charge.
Positive metal ions: they are charged, but in a solid metal they are held in fixed lattice positions. They cannot move through the solid to carry charge.
Therefore, being charged is not enough for electrical conduction. The charged particles must also be able to move.
Why can metals conduct electricity when solid?
This point is particularly useful when comparing metallic and ionic substances. In a solid metal, the positive ions remain in fixed positions, but the delocalised electrons are still mobile.
A solid metal can therefore conduct electricity. By contrast, the ions in a solid ionic compound are fixed in the ionic lattice and cannot move from place to place to carry charge.
Why are metals good conductors of thermal energy?
When one end of a metal is heated, the delocalised electrons in that region gain kinetic energy. Because these electrons are mobile, they can move through the metal and transfer energy rapidly to other particles.
Thermal energy therefore spreads efficiently from the hotter region to cooler parts of the metal.
Why are metals malleable and ductile?
Malleable means that a material can be hammered or rolled into sheets. Ductile means that it can be drawn into wires.
In a metal, the positive metal ions are arranged in regular layers. When a force is applied, one layer of ions can slide past another.
The delocalised electrons remain throughout the structure as the layers move. Therefore, the positive metal ions remain attracted to the surrounding electrons, so metallic bonding is maintained instead of the structure immediately fracturing.
Metal: layers of positive ions can slide while remaining attracted to delocalised electrons, so the material can change shape.
Ionic crystal: if layers are forced to shift, ions with the same charge may be brought next to one another. Their strong repulsion can cause the crystal to crack.
| Property of a metal | Particle-level explanation |
|---|---|
| Usually high melting and boiling points | Strong electrostatic attraction between positive metal ions and delocalised electrons requires a large amount of energy to overcome. |
| Electrical conductivity | Mobile delocalised electrons move through the structure and carry electrical charge. The positive metal ions remain fixed in lattice positions in the solid. |
| Thermal conductivity | Mobile delocalised electrons gain kinetic energy and transfer energy rapidly through the metal. |
| Malleability | Layers of positive ions can slide while remaining attracted to the delocalised electrons. |
| Ductility | The ability of layers to move while metallic bonding is maintained allows the metal to be drawn into wires. |
Alloys: Why Changing the Structure Changes the Properties
The ability of layers to slide is useful when we want to shape a pure metal, but it can also make some pure metals too soft for a particular purpose.
One way to change a metal's physical properties is to mix it with another element or other elements. The resulting material is called an alloy.
An alloy is a mixture containing a metal and one or more other elements.
The additional element does not have to be another metal. For example, steel contains iron mixed with a small amount of carbon, and carbon is a non-metal.
Why can pure metals be relatively soft?
When comparing the hardness of pure metals and alloys, particle diagrams are often simplified to show the atoms as circles. The purpose of this model is to focus on particle size and the movement of layers.
In a pure metal, the atoms are similar in size and can form regular layers. These layers can slide over one another comparatively easily when a force is applied.
Why are many alloys harder than pure metals?
An alloy contains atoms of another element or other elements. These atoms may have different sizes from the atoms of the main metal.
The differently sized atoms disturb the regular arrangement of the layers. The layers can therefore no longer slide past one another as easily.
A greater force is needed to change the shape of the material, so many alloys are harder than the corresponding pure metal.
Pure metal: similarly sized atoms form regular layers that can slide over one another relatively easily.
Alloy: differently sized atoms disrupt the regular arrangement and make it more difficult for the layers to slide.
Common examples of alloys
| Alloy | Main components | Useful property |
|---|---|---|
| Brass | Copper and zinc | Harder than pure copper and useful for fittings, instruments and decorative objects. |
| Steel | Iron and a small amount of carbon | Generally harder and stronger than pure iron. |
| Stainless steel | Iron with chromium and other elements such as nickel and carbon | Strong and resistant to corrosion. |
| Gold alloys | Gold mixed with metals such as copper or silver | Harder and more durable than pure gold for everyday jewellery. |
Is an alloy a compound?
No. An alloy is a mixture. The different elements do not have to be present in one fixed ratio, and different compositions can be chosen to produce different physical properties.
This is why there are many different types of steel. By changing the proportions of carbon and other elements, manufacturers can produce materials suited to different uses.
Building Complete Examination Explanations
Metallic bonding questions are often short, but good answers still need a complete chain of reasoning. Avoid jumping straight from the name of a structure to the final property.
Example 1: Explain why copper is suitable for electrical wiring
A weak answer might simply say, “Copper conducts electricity.” That states a property but does not explain it.
Better answer: Copper contains mobile delocalised electrons that can move through the metallic structure and carry electrical charge. The positive copper ions remain in fixed lattice positions and therefore do not carry the current through the solid. Copper is also ductile, so it can be drawn into long, thin wires.
Example 2: Explain why steel is harder than pure iron
Answer: In pure iron, similarly sized atoms are arranged in regular layers that can slide over one another relatively easily. In steel, carbon atoms disrupt this regular arrangement and make it more difficult for the layers to slide. Steel is therefore harder than pure iron.
Example 3: Explain why a metal can have a high melting point
Answer: There is strong electrostatic attraction between the positive metal ions and the delocalised electrons in the giant metallic lattice. A large amount of thermal energy is needed to overcome these attractions, so the metal has a high melting point.
Common Mistakes with Metallic Bonding
| Common mistake | Why it is wrong | Correct approach |
|---|---|---|
| Describing the metallic lattice only as neutral metal atoms | In the metallic-bonding model, the outer electrons are delocalised. | Describe the structure as positive metal ions surrounded by delocalised electrons. |
| Saying that the delocalised electrons have escaped from the metal | They are no longer associated with individual atoms but remain within the metallic structure. | Say that the electrons are mobile throughout the metal. |
| Saying that solid metals conduct because the positive ions move | The positive ions are charged, but they remain in fixed lattice positions in the solid. | Identify the mobile delocalised electrons as the charge carriers. |
| Assuming that any charged particle automatically conducts electricity | A charged particle must also be mobile in order to carry charge through the substance. | Ask two questions: Is the particle charged? and Can it move? |
| Explaining a high melting point using “strong intermolecular forces” | A metal is a giant metallic structure rather than a simple molecular substance. | Refer to strong electrostatic attraction between positive metal ions and delocalised electrons. |
| Saying that metals are malleable because metallic bonds are weak | Metallic bonding can be strong. Malleability results from the bonding being maintained as layers move. | Explain that layers of positive ions can slide while remaining attracted to delocalised electrons. |
| Calling an alloy a compound | Alloys are mixtures and their compositions can vary. | Define an alloy as a mixture containing a metal and one or more other elements. |
| Saying that an alloy is harder simply because it contains more atoms | The important change is disruption of the regular layers. | Explain that differently sized atoms make it more difficult for the layers to slide. |
| Saying that every alloy contains only metals | Some alloys contain non-metallic elements. Steel contains carbon. | Remember that an alloy contains a metal plus one or more other elements. |
Metallic Bonding Checklist
Describe the structure
- State that the metal has a giant metallic structure.
- Identify the lattice particles as positive metal ions.
- Identify the surrounding electrons as delocalised.
- Remember that the electrons remain within the metal.
Define the bonding
- Use the words electrostatic attraction.
- Name the two attracted particles.
- Positive metal ions are attracted to delocalised electrons.
- Do not describe metallic bonding as an intermolecular force.
Explain the properties
- Melting point → strong attraction requires much energy to overcome.
- Electrical conductivity → mobile electrons carry charge.
- Positive ions are charged but fixed in the solid lattice.
- Thermal conductivity → mobile electrons transfer energy.
- Malleability and ductility → ion layers slide while bonding is maintained.
Explain alloys
- Define an alloy as a mixture.
- Compare regular layers with disrupted layers.
- Refer to differently sized atoms.
- Explain that the disrupted layers cannot slide as easily.
Now apply the ideas from the chapter. These ten questions focus on explanation, comparison and unfamiliar situations rather than simple definition recall.
- 10 application questions
- Explanation for every answer option
- Practise structure → behaviour → property reasoning
Frequently asked questions
Questions Students Often Ask About Metallic Bonding
These questions clarify the particle model and help separate metallic bonding from ideas that belong to ionic or covalent structures.
Are delocalised electrons completely free?
They are free to move throughout the metallic structure, but they have not escaped from the metal. The word delocalised means that they are not associated with one particular metal ion.
Why do the positive metal ions not simply repel one another?
The metal ions are surrounded by negatively charged delocalised electrons. Strong electrostatic attraction between the ions and these electrons holds the giant structure together.
Do metals consist of molecules?
No. A typical metal forms a giant metallic lattice extending throughout the solid. It is not divided into separate molecules.
Why do the delocalised electrons conduct electricity instead of the positive metal ions?
Both are charged, but only the delocalised electrons are mobile throughout a solid metal. The positive metal ions are held in fixed lattice positions and cannot travel through the solid to carry charge. Electrical conduction requires particles that are both charged and mobile.
Why can both solid and molten metals conduct electricity?
Delocalised electrons remain mobile in both states. In the solid, the positive ions occupy lattice positions while the electrons move. When the metal melts, mobile electrons are still present and can carry charge.
Does bending a metal mean that all its metallic bonds are broken?
No. When layers of positive ions move, they remain surrounded by and attracted to the delocalised electrons. Metallic bonding continues to hold the structure together.
What is the difference between malleability and ductility?
Malleability is the ability to be hammered or rolled into sheets. Ductility is the ability to be drawn into wires.
Can an alloy contain a non-metal?
Yes. An alloy must contain a metal, but the additional element does not have to be metallic. Steel, for example, contains iron and carbon.
Why do alloy diagrams sometimes show atoms rather than positive ions and electrons?
The alloy-hardness diagram is a simplified particle model designed to show differences in particle size and how easily layers can slide. The metallic bonding within the alloy is still described using positive metal ions and delocalised electrons.
Are all alloys harder than pure metals?
Not necessarily. Many common alloys are designed to be harder or stronger, but alloys can also be made to improve other properties such as corrosion resistance. Their exact properties depend on their composition.
What should I write when asked to define metallic bonding?
State that metallic bonding is the strong electrostatic attraction between positive metal ions and delocalised electrons.
Metallic bonding gives one connected explanation for the behaviour of metals. Strong attraction between positive ions and delocalised electrons holds the giant structure together; mobile electrons carry charge and thermal energy; and the same attraction allows layers of ions to slide without the metal immediately breaking. Change the regular arrangement by forming an alloy, and the physical properties can change as well. Once you see this structure–property link, metallic bonding becomes a reasoning topic rather than a list of facts to memorise.