How Metallic Bonding Explains Conductivity, Lustre and Malleability

Metals are shiny, they conduct electricity and heat, and they can be hammered into sheets or drawn into wires without breaking. Most have high melting points. These properties are so familiar that we rarely ask why — but every one of them follows from a single idea about how metal atoms are held together: metallic bonding. This article explains the model and shows how it accounts for each property, with real data.

The sea-of-electrons model

Metal atoms have few outer electrons (often one, two or three), and they hold them relatively loosely — their ionisation energies are low. When huge numbers of metal atoms pack together in a solid:

1. Each atom releases its outer electrons into a shared pool.
2. These electrons are delocalised: they don’t belong to any one atom but move freely throughout the whole structure. This is often called a “sea” of electrons.
3. What remains are positive metal ions (sometimes called cations or atomic cores), arranged in a regular lattice.

Metallic bonding is the strong electrostatic attraction between the positive metal ions and the sea of delocalised electrons (see metallic bonding).

Two features make this bonding different from ionic and covalent bonding:

• It’s non-directional: each ion is attracted to the electron sea all around it, not to particular neighbours.
• The electrons are mobile.

Keep these two features in mind — they explain everything below.

Property 1: electrical conductivity

Electricity is a flow of charge. In a metal, the delocalised electrons are free to move. When a voltage is applied across a metal wire, electrons drift towards the positive end, carrying a current — and they’re immediately replaced by electrons entering from the other end.

Metals conduct in both solid and liquid states, because the electron sea is present in both. That’s different from ionic compounds, which conduct only when molten or dissolved (see properties of ionic compounds).

The best electrical conductors at room temperature:

Metal	Relative conductivity (silver = 100)
Silver	100
Copper	about 95
Gold	about 70
Aluminium	about 60
Iron	about 16

Copper is used for most electrical wiring because it combines excellent conductivity with low cost and good ductility (see copper). Aluminium, lighter and cheaper, is used in overhead power lines.

Temperature effect: as a metal gets hotter, its ions vibrate more, getting in the way of the moving electrons. So a metal’s electrical resistance increases with temperature — which is why the filament in an old light bulb had much higher resistance when glowing than when cold.

Property 2: thermal conductivity

Metals feel cold to the touch because they conduct heat away from your hand quickly. Heat travels through a metal in two ways:

1. The delocalised electrons pick up kinetic energy in the hot region and carry it rapidly through the metal as they move.
2. Vibrations of the ions pass from one to the next through the closely packed lattice.

The first mechanism is so efficient that good electrical conductors are almost always good thermal conductors (the relationship is called the Wiedemann–Franz law). Copper-bottomed saucepans and aluminium heat sinks in computers use this property.

Property 3: malleability and ductility

Metals are malleable (can be hammered or pressed into shape) and ductile (can be drawn into wires). Gold is the extreme case: a single gram can be beaten into a sheet about a square metre in area, thin enough for light to pass through.

Why don’t metals shatter like salt crystals? In a metal, the positive ions are arranged in layers. When a force is applied, whole layers can slide over one another. Because the bonding is non-directional, the electron sea simply moves with the ions and keeps holding them together in their new positions. No bonds are “broken” in the sense of creating repulsion.

Compare an ionic crystal: sliding a layer brings ions of the same charge next to each other, and the resulting repulsion splits the crystal (see ionic lattices).

In real metals, the sliding happens through tiny defects in the crystal called dislocations, which move through the lattice more easily than whole planes could. This is also the key to understanding why alloys are harder than pure metals (see alloys: why mixing metals makes them stronger).

Property 4: lustre (shininess)

Freshly cut or polished metals are shiny. When light hits a metal surface, the delocalised electrons absorb it and immediately re-emit it, reflecting most of the incoming light. Because the electrons can absorb and re-emit across the whole visible spectrum, most metals look silvery-grey.

Copper (reddish) and gold (yellow) are exceptions: they absorb some blue and violet light more strongly, so the reflected light is shifted towards red and yellow. In gold, this effect is due partly to relativistic effects on its electrons — its heavy nucleus makes the inner electrons move so fast that relativity changes orbital energies.

Many metals quickly lose their shine in air, because their surface reacts with oxygen to form a dull oxide layer (the most reactive metals tarnish fastest — see the most reactive metal).

Property 5: high melting and boiling points (mostly)

Most metals have high melting and boiling points, because the attraction between the ions and the electron sea is strong, and it extends throughout the whole structure.

The strength of metallic bonding increases with:

1. More delocalised electrons per atom — a bigger, denser electron sea.
2. Higher charge on the ions.
3. Smaller ions, so the electron sea is closer to the nuclei.

Metal	Delocalised electrons per atom	Melting point
Sodium	1	98 °C
Magnesium	2	650 °C
Aluminium	3	660 °C
Potassium	1	63 °C
Iron	several (including d electrons)	1,538 °C
Tungsten	several	3,422 °C

Down group 1, melting points fall (lithium 181 °C, sodium 98 °C, potassium 63 °C, caesium 28 °C), because the ions get larger and the electron sea is further from the nuclei (see alkali metals). Across period 3, from Na to Mg to Al, melting points rise as the number of delocalised electrons rises.

Transition metals, which can contribute d electrons to the electron sea, tend to have very high melting points. Tungsten has the highest melting point of any metal, which is why it was used for light-bulb filaments (see the highest boiling point element). Mercury is the famous exception: it’s a liquid at room temperature (melting point −39 °C), partly because its outer electrons are held unusually tightly (again, a relativistic effect).

Property 6: high density (mostly)

Metal atoms usually pack very efficiently — often in close-packed arrangements where each atom has 12 nearest neighbours. Combined with heavy nuclei, this gives many metals high densities. Osmium and iridium, at about 22.6 g cm⁻³, are the densest elements. The group 1 metals are exceptions: lithium, sodium and potassium are less dense than water.

Limitations of the model

The sea-of-electrons model is simple and explains the main properties well, but it’s not the whole story. It doesn’t easily explain why some elements are semiconductors, or why metals have particular magnetic properties. A more complete picture, band theory, extends molecular orbital theory to the enormous number of atoms in a solid (see valence bond vs molecular orbital theory).

Common misconceptions

• “Metals conduct because they contain ions.” It’s the delocalised electrons that move; the ions stay in the lattice.
• “Metallic bonds are between pairs of atoms.” The bonding is non-directional, involving the whole structure.
• “All metals have high melting points.” Group 1 metals and mercury are low.
• “Metals bend because their bonds are weak.” They bend because the non-directional bonding survives layers sliding.

Key takeaways

• Metallic bonding = attraction between positive metal ions and a sea of delocalised electrons.
• Conductivity (electrical and thermal): mobile delocalised electrons carry charge and energy.
• Malleability and ductility: layers slide; non-directional bonding holds them together.
• Lustre: electrons absorb and re-emit light.
• Melting points rise with more delocalised electrons, higher ion charge and smaller ions.