What you'll learn
Metallic bonding explains why metals are strong, conduct electricity and heat, and can be bent and shaped. In this guide you will learn how metallic bonds form, the structure of a giant metallic lattice with delocalised electrons, how that structure explains the characteristic properties of metals, and how alloys differ from pure metals. These ideas connect the bonding models to the everyday usefulness of metals in construction, wiring and manufacturing.
Key terms and definitions
Metallic bond — the strong electrostatic attraction between positive metal ions and a sea of delocalised electrons.
Delocalised electrons — outer-shell electrons that are free to move throughout the metal structure.
Giant metallic lattice — a regular arrangement of metal ions surrounded by delocalised electrons.
Alloy — a mixture of a metal with one or more other elements.
Malleable — able to be hammered or rolled into shape.
Ductile — able to be drawn into wires.
Core concepts
How metallic bonding works
In a metal, the atoms lose their outer electrons, which become delocalised — free to move throughout the structure. This leaves a regular lattice of positive metal ions held together by their attraction to the sea of delocalised electrons. The metallic bond is the strong electrostatic attraction between these positive ions and the negative delocalised electrons, acting in all directions throughout the giant lattice.
Why metals conduct electricity and heat
Because the delocalised electrons are free to move, they can carry electrical charge through the metal, making metals good electrical conductors. The same mobile electrons transfer kinetic energy quickly through the structure, making metals good thermal conductors (good at conducting heat).
Why metals have high melting points
The electrostatic attraction between the metal ions and the delocalised electrons is strong, so a large amount of energy is needed to overcome it. This gives most metals high melting and boiling points, and they are solids at room temperature (mercury is an exception).
Why metals are malleable and ductile
The metal ions are arranged in layers that can slide over each other without breaking the metallic bonding, because the delocalised electrons continue to hold the structure together. This means metals can be bent, hammered into shape (malleable) and drawn into wires (ductile).
Alloys
Most metals in everyday use are alloys — mixtures of a metal with other elements. Pure metals are often too soft because their layers slide easily. In an alloy, the different-sized atoms distort the regular layers, making it harder for them to slide over each other, so alloys are usually harder than the pure metal. Examples include steel (iron with carbon), bronze (copper and tin) and brass (copper and zinc).
Worked examples
Example 1: Explaining conductivity
Why do metals conduct electricity?
Metals have delocalised electrons that are free to move through the structure. These electrons carry electrical charge, so the metal conducts.
Example 2: Explaining malleability
Explain why copper can be drawn into wires.
Copper's ions are in layers that can slide over one another while the delocalised electrons keep the structure bonded. This makes copper ductile, so it can be drawn into wires.
Example 3: Why alloys are harder
Explain why steel is harder than pure iron.
Steel contains carbon atoms of a different size, which distort the layers of iron ions. This stops the layers sliding easily, so steel is harder than pure iron.
Common mistakes and how to avoid them
Forgetting to mention delocalised electrons. They are the key to both conductivity and the bonding itself.
Saying ions move to carry current. In a solid metal the ions stay fixed; the electrons move.
Not explaining malleability by sliding layers. Bonding stays intact as layers slide, so the metal bends rather than shatters.
Thinking alloys are compounds. Alloys are mixtures, not chemically combined compounds.
Saying alloys are harder because they are stronger metals. It's the different-sized atoms distorting layers that increases hardness.
Exam technique for Metallic Bonding
Define the metallic bond fully — attraction between positive ions and a sea of delocalised electrons.
Use delocalised electrons to explain electrical and thermal conductivity.
Explain high melting points by the strong electrostatic attraction needing lots of energy to overcome.
Explain malleability/ductility using sliding layers held together by delocalised electrons.
Explain alloy hardness with different-sized atoms distorting the layers.
Quick revision summary
Metallic bonding is the strong electrostatic attraction between a lattice of positive metal ions and a sea of delocalised electrons (the outer electrons that have become free to move). This structure explains all the key properties of metals. The delocalised electrons are free to move, so metals are good conductors of electricity and heat. The strong attraction between ions and electrons needs a lot of energy to overcome, giving high melting and boiling points. The metal ions lie in layers that can slide over each other while staying bonded, so metals are malleable and ductile. Pure metals are often soft because the layers slide easily; alloys mix in different-sized atoms that distort the layers and stop them sliding, making alloys harder (steel, bronze, brass). When answering, always name the delocalised electrons, link them to conductivity, and use sliding layers to explain shaping and alloy hardness.