Metallic BondingCambridge IGCSE Chemistry: Revision notes
Section 1
What Is Metallic Bonding?
Metallic bonding is the strong electrostatic attraction between positive metal ions and a 'sea' of delocalised electrons, in a giant metallic lattice.
- When metal atoms bond together, each atom loses its outer shell electrons, which become delocalised (free to move throughout the whole structure) rather than belonging to one atom
- This leaves a regular, repeating lattice of positive metal ions
- The lattice of positive ions is held together by strong electrostatic attraction to the surrounding sea of negative delocalised electrons
- This attraction acts in all directions throughout the structure, giving a giant lattice
Think of a metal lattice like a sea of fixed positive 'buoys' (the metal ions) surrounded by freely flowing 'water' (the delocalised electrons) that holds them all together.
Section 2
How Does Metallic Bonding Explain the Properties of Metals?
The structure and bonding of metals directly explains their characteristic physical properties.
| Property | Explanation in terms of structure/bonding |
|---|---|
| Good electrical conductivity | The delocalised electrons are free to move through the structure, carrying charge |
| Good thermal conductivity | Delocalised electrons can transfer kinetic energy quickly through the lattice |
| Malleable and ductile | Layers of positive ions can slide over one another without breaking the metallic bond, since the delocalised electrons still surround the ions after they move |
| Generally high melting/boiling points | Strong electrostatic attraction between ions and delocalised electrons requires a lot of energy to overcome |
When explaining conductivity in an exam, always mention the delocalised electrons are 'free to move' — this is the key phrase examiners look for.
Section 3
How Does Metallic Structure Compare With Other Giant Structures?
Metals are one of several types of giant structure, along with giant covalent structures like diamond and silicon(IV) oxide. Comparing them shows why metallic bonding gives different properties.
- Diamond: a giant covalent structure where every carbon atom forms four strong covalent bonds to neighbouring carbon atoms; extremely hard, very high melting point, does not conduct electricity (no free electrons or ions)
- Silicon(IV) oxide (SiO2): a giant covalent structure similar to diamond, with each silicon atom bonded to four oxygen atoms; also very hard with a very high melting point, and does not conduct electricity
- Metals: a giant metallic lattice; unlike diamond and SiO2, metals do conduct electricity, because of the mobile delocalised electrons, and their ions can slide past each other (malleable), unlike the rigid covalent networks of diamond and SiO2
Don't say diamond or SiO2 conduct electricity — they have no delocalised electrons or free ions, unlike metals, so they are non-conductors even though all three are giant structures.
Must Know
- Metallic bonding is the electrostatic attraction between positive metal ions and a sea of delocalised electrons
- Metal atoms lose their outer electrons, which become delocalised and free to move through the whole structure
- Delocalised electrons explain good electrical and thermal conductivity in metals
- Layers of ions can slide over each other without breaking bonds, explaining malleability and ductility
- Diamond and silicon(IV) oxide are giant covalent structures — very hard, high melting point, but do not conduct electricity, unlike metals
- Strong electrostatic attraction throughout the metallic lattice gives metals generally high melting and boiling points
That's the notes covered.
Carry on to the next subtopic.