In metallic bonding, metal atoms lose their outer electrons to form positive ions surrounded by a "sea" of delocalised electrons. The strong electrostatic attraction between these positive ions and the delocalised electron sea holds the structure together, and directly explains why metals are excellent conductors of both electricity and heat.
What is metallic bonding?
When metal atoms come together in a solid, their outer electrons are not held by any individual atom — instead, they leave their "parent" atoms entirely and become delocalised: free to move throughout the whole structure. The atoms, now positive ions (cations), are arranged in a regular lattice, and the delocalised electrons form a shared "sea" surrounding all the ions.
Metallic bonding is the strong electrostatic attraction between:
- the positive metal ions (regularly arranged in a lattice), and
- the sea of delocalised (free) electrons
This attraction acts in all directions and holds the structure together. It is this bond that must be overcome when a metal melts.
The "sea of electrons" model is a simplification, but it correctly predicts the characteristic properties of metals at GCSE level.
Why do metals conduct electricity?
Electrical conductivity requires mobile charged particles. In metallic bonding, the delocalised electrons are not fixed to any particular ion — they are free to move throughout the lattice. When a voltage (potential difference) is applied across a metal, these free electrons drift in one direction, creating a conventional electric current.
This is why all metals are good electrical conductors, and non-metals (which do not have delocalised electrons) are typically poor conductors. Graphite is the main exception among non-metals — its one delocalised electron per carbon atom makes it a conductor.
Why are metals good conductors of heat?
Metals conduct heat efficiently through two mechanisms:
- Electron conductivity: delocalised electrons in the hotter region of the metal have higher kinetic energy and rapidly transfer this energy to cooler regions as they move.
- Lattice vibration: the positive ions vibrate and transfer vibrational energy to neighbouring ions, but this mechanism is less efficient than electron transfer.
Because electrons spread energy so quickly, metals feel cold to the touch even at room temperature — they conduct heat away from your hand more quickly than non-metals at the same temperature.
Why are metals malleable and ductile?
Two important properties that metallic bonding explains are:
- Malleable: can be hammered into flat sheets without cracking
- Ductile: can be drawn into wires without breaking
In an ionic lattice (like sodium chloride), shifting the layers by even one ion position brings like charges next to each other, creating a strong repulsive force that shatters the crystal. This is why ionic compounds are brittle.
In a metallic lattice, the positive ions are all identical (in a pure metal). When layers of ions slide over each other under an applied force, each ion remains surrounded by the same sea of electrons — there is no electrostatic repulsion at any position. The ions simply slide to new positions while maintaining the bond. This is why metals can be deformed without fracturing.
Why do metals have high melting points?
To melt a metal, you must overcome the electrostatic attraction between the positive ions and the delocalised electron sea. This attraction is strong throughout the lattice — in all directions — and a large amount of energy is required to separate the ions from the electron sea.
Metals with:
- More electrons lost per atom (higher charge on the ion) → stronger attraction → higher melting point
- Smaller ions (electrons closer to the nucleus) → stronger attraction → higher melting point
| Metal | No. of outer electrons | Melting point (°C) |
|---|---|---|
| Sodium (Na) | 1 | 98 |
| Magnesium (Mg) | 2 | 650 |
| Aluminium (Al) | 3 | 660 |
| Iron (Fe) | 2–3 (d-block) | 1,538 |
| Tungsten (W) | 6 | 3,422 |
The trend is not perfectly linear because transition metals have complex electron arrangements, but the general pattern holds.
How do alloys relate to metallic bonding?
A pure metal has all atoms of the same size, so the lattice is regular and the layers of ions can slide easily — this makes pure metals relatively soft and deformable.
An alloy contains two or more metals (or a metal and a non-metal, such as iron and carbon in steel). The different-sized atoms distort the regular lattice:
- The mismatched atoms prevent layers from sliding freely over each other.
- More force is required to cause deformation.
- The result is a harder, stronger material.
Examples:
- Steel (iron + 0.1–2% carbon): much harder than pure iron; used in construction, tools, cars
- Bronze (copper + tin): harder than pure copper; used in tools, sculptures, bearings
- Brass (copper + zinc): harder than copper; used in musical instruments, fittings
- Duralumin (aluminium + copper + other metals): much stronger than pure aluminium; used in aircraft structures
The delocalised electron sea remains intact in alloys, so alloys still conduct electricity and heat well.
Frequently asked questions
What is the difference between metallic bonding, ionic bonding, and covalent bonding?
Ionic bonding involves the transfer of electrons from metal to non-metal atoms to form oppositely charged ions, held together by electrostatic attraction in a lattice — these compounds are brittle and conduct electricity only when dissolved or molten. Covalent bonding involves the sharing of electrons between non-metal atoms to form discrete molecules (or giant lattices like diamond), and most covalent substances do not conduct electricity. Metallic bonding involves metal atoms releasing electrons into a delocalised sea that holds positive ions together — this allows electrical conduction in the solid state and accounts for malleability, ductility, and thermal conductivity.
Why are alloys harder than pure metals?
In a pure metal, all atoms are the same size, and the regular lattice allows layers of ions to slide over each other easily when a force is applied, making the metal relatively soft. In an alloy, atoms of different sizes are randomly distributed in the lattice, creating local distortions. These distortions prevent the layers from sliding smoothly — a greater force is needed to cause deformation, making the alloy harder and stronger. The delocalised electron sea is not significantly disrupted, so alloys retain the conductivity and metallic appearance of pure metals.
Why does metallic bonding produce a shiny appearance?
The delocalised electrons in a metal respond to visible light by absorbing photons across a wide range of frequencies and immediately re-emitting them. This re-emission of light from the smooth metal surface produces the characteristic shiny, lustrous appearance of metals. Non-metals that lack delocalised electrons cannot do this — their electrons are fixed in bonds or lone pairs and can only absorb specific frequencies. This is why non-metals are typically dull, while metals are shiny.
Why is mercury a liquid at room temperature despite having metallic bonding?
Mercury (Hg) is unusual: it has a completely filled 5d electron shell, which makes its electrons particularly tightly bound and less available for strong metallic bonding compared with neighbouring transition metals. The result is a weaker metallic bond and a melting point of −39 °C — well below room temperature. Mercury is the only metal that is liquid at standard conditions (25 °C). This illustrates that metallic bond strength varies considerably across the periodic table depending on the number and availability of the delocalised electrons.
For Socratic GCSE chemistry with Professor Curie — building from the sea of electrons model upward to explain every observed property of metals — visit aitutors.me.