What you'll learn
Metals conduct electricity, can be hammered into shape, and mostly have high melting points — and all of these properties come from one type of bonding. Metallic bonding is the strong attraction between positive metal ions and a "sea" of delocalised electrons. For AQA GCSE Chemistry you need to describe metallic bonding, explain how it accounts for conductivity, malleability and high melting points, and explain why alloys are harder than pure metals. This guide covers the structure of a metal, each property in turn with the reasoning that earns marks, and the structure and uses of alloys. By the end you should be able to explain any metal property by linking it back to the delocalised electrons and the layers of ions.
Key terms and definitions
Metallic bonding — The strong electrostatic attraction between positive metal ions and the delocalised electrons that surround them.
Delocalised electron — An outer-shell electron that is free to move throughout the whole metal structure and is not attached to any one atom.
Lattice — The regular, repeating arrangement of particles; in a metal, a lattice of positive ions.
Malleable — Able to be hammered or rolled into shape without breaking.
Ductile — Able to be drawn out into wires.
Alloy — A mixture of a metal with one or more other elements, usually another metal.
Conductor — A substance that allows electricity or heat to pass through it.
Electrostatic attraction — The force of attraction between oppositely charged particles.
Core concepts
The structure of a metal
In a metal, the atoms are packed closely together in a regular lattice. Each metal atom loses its outer-shell electrons, which become delocalised and free to move throughout the whole structure. This leaves behind a lattice of positive metal ions surrounded by a "sea" of delocalised electrons. The metallic bond is the strong electrostatic attraction between these positive ions and the negative delocalised electrons. This attraction acts in all directions throughout the structure, which is why metallic bonding is so strong.
Why metals conduct electricity and heat
Metals are good conductors of electricity because the delocalised electrons are free to move through the structure. When a voltage is applied, these electrons flow through the metal, carrying charge from one end to the other. Metals also conduct heat well, for two reasons: the delocalised electrons move and transfer energy quickly through the metal, and the closely packed ions can pass on vibrations. The key phrase in any answer is that the electrons are free to move and carry charge or energy.
Why metals have high melting and boiling points
Most metals are solid at room temperature with high melting points. This is because the electrostatic attraction between the positive ions and the delocalised electrons is very strong, so a large amount of energy is needed to overcome these forces and separate the particles. The stronger the metallic bonding, the higher the melting point. Metals with ions of higher charge tend to have stronger bonding and therefore higher melting points, because there is greater attraction to the electron sea.
Why pure metals are malleable
In a pure metal, the ions are all the same size and are arranged in neat layers. When a force is applied, these layers can slide over one another without breaking the metallic bonds, because the delocalised electrons continue to hold the structure together as it changes shape. This is why pure metals are malleable (can be hammered into shape) and ductile (can be drawn into wires). The metal deforms rather than shattering.
Alloys and why they are harder
An alloy is a mixture of a metal with one or more other elements. The added atoms are a different size from the original metal atoms, so they disrupt the neat, regular layers. Because the layers are no longer uniform, they cannot slide over one another easily, so the alloy is harder and less malleable than the pure metal. This is why alloys are usually used in place of pure metals when strength matters — for example, steel (iron mixed with carbon) is much harder than pure iron, and bronze (copper and tin) is harder than pure copper.
Common alloys and their uses
| Alloy | Made from | Property | Use |
|---|---|---|---|
| Steel | Iron + carbon | Hard, strong | Buildings, cars, tools |
| Stainless steel | Iron + chromium + nickel | Resists corrosion | Cutlery, sinks |
| Bronze | Copper + tin | Hard, resists corrosion | Statues, bells |
| Brass | Copper + zinc | Hard, attractive finish | Instruments, fittings |
Metals as giant structures
It is important to see a metal as a giant structure, not as separate molecules. The lattice of ions and the electron sea continue throughout the whole piece of metal, which is why the bonding is described as strong and why properties like conductivity apply throughout the material rather than to individual particles.
Worked examples
Example 1: Explaining conductivity
Explain why copper is used for electrical wiring. Copper is a metal with delocalised electrons that are free to move throughout its structure. When a voltage is applied, these electrons flow and carry charge, so copper conducts electricity well. It is also ductile, so it can be drawn into long wires.
Example 2: Explaining a high melting point
Explain why magnesium has a high melting point. Magnesium has metallic bonding, with a strong electrostatic attraction between the positive magnesium ions and the delocalised electrons. A large amount of energy is needed to overcome these strong forces, so the melting point is high.
Example 3: Comparing an alloy with a pure metal
Explain why steel is harder than pure iron. In pure iron, the atoms are the same size and arranged in regular layers that can slide over each other. In steel, carbon atoms of a different size disrupt the regular layers, so the layers cannot slide easily. This makes steel harder and less malleable than pure iron.
Example 4: Choosing a material
A manufacturer needs a material for bicycle frames that is strong but not too heavy. Should they choose a pure metal or an alloy, and why? They should choose an alloy. Alloys contain atoms of different sizes that stop the layers sliding, making the material harder and stronger than the pure metal, which is important for a frame that must resist bending.
Common mistakes and how to avoid them
The most common error is saying metals conduct because of "free ions". They do not — the ions in a metal lattice are fixed in place. Conduction comes from the delocalised electrons being free to move. Always name the electrons.
Students often describe metallic bonding as "atoms attracted to electrons". Be precise: it is the attraction between positive metal ions and delocalised electrons. Using "atoms" instead of "ions" loses marks.
When explaining malleability, do not say the bonds break as the layers slide. The metallic bonds are maintained because the delocalised electrons move with the structure — that is exactly why the metal bends instead of shattering.
A frequent alloy mistake is saying the added atoms make the metal "stronger because they are stronger atoms". The real reason is that they are a different size, which disrupts the regular layers and stops them sliding.
Finally, do not confuse malleable and ductile. Malleable means it can be hammered into shape; ductile means it can be drawn into wires.
Exam technique for "Metallic bonding and properties of metals and alloys"
Almost every question here asks you to link structure to property, so build a chain in your answer: state the structure (lattice of positive ions in a sea of delocalised electrons), then the reason (electrons free to move, or layers able to slide, or strong electrostatic attraction), then the property.
For alloy questions, the marking points are almost always: different sized atoms, layers disrupted, layers cannot slide, so harder. Learn that sequence as a set.
Diagrams of metallic structure are useful, but the marks come from the written explanation — label positive ions and delocalised electrons if you do draw one. Six-mark questions often ask you to compare a pure metal with an alloy, or to explain several metal properties at once, so practise writing conductivity, melting point and malleability explanations back to back.
Quick revision summary
- Metallic bonding is the strong electrostatic attraction between a lattice of positive metal ions and a sea of delocalised electrons.
- Metals conduct electricity and heat because delocalised electrons are free to move and carry charge or energy.
- Metals have high melting points because a lot of energy is needed to overcome the strong attraction between ions and electrons.
- Pure metals are malleable and ductile because regular layers of same-sized ions slide over one another.
- Alloys contain different-sized atoms that disrupt the layers, so they cannot slide — making alloys harder than pure metals.
- Examples: steel (iron + carbon), bronze (copper + tin), brass (copper + zinc).