Metallic bonding

Metallic bonding is the electrostatic attraction between a regular lattice of positive metal ions (cations) and a sea of delocalized electrons.

When metal atoms pack together, each atom loses its valence (outer-shell) electrons to a shared pool. What remains is a 3D array of positive ions, and the released electrons spread out across the whole structure as a mobile 'sea'.

• The electrons are delocalized - not owned by any one atom, but free to move through the lattice.
• The bonding is the attraction between the fixed positive ions and these mobile negative electrons.
• Metallic bonding is non-directional: the attraction pulls equally in every direction, so the ions pack into close, regular arrangements.

A useful picture is positive ions sitting in a 'glue' of negative electrons. Because the electrons hold all the cations together at once, the structure is a giant metallic lattice - there are no individual molecules.

This model immediately explains several properties. The mobile electrons carry charge (so metals conduct electricity) and energy (so they conduct heat). The strong, non-directional attraction explains why metals are typically strong and have high melting points, while still being able to deform without shattering.

The strength of metallic bonding depends on three linked factors:

• Ionic charge - a higher charge on the cation means more electrons donated to the sea and a stronger attraction.
• Ionic radius - smaller ions let the electron sea sit closer to the nuclei, strengthening the bond.
• Number of delocalized electrons - more electrons per atom means more 'glue'.

Across Period 3, $\text{Na} \rightarrow \text{Mg} \rightarrow \text{Al}$ donate 1, 2 and 3 electrons respectively, the ionic radius falls, and melting point rises sharply - showing stronger bonding.

Key properties explained by the model:

• Electrical conductivity - delocalized electrons move freely to carry current (in solid and molten states).
• Thermal conductivity - mobile electrons transfer kinetic energy quickly.
• Malleability and ductility - layers of cations can slide over one another; the electron sea simply flows around them and keeps bonding intact, so the metal bends rather than cracks.
• High melting/boiling points - strong attractions need lots of energy to overcome.
• Lustre (shine) - the surface electron sea reflects light.

Contrast with ionic solids, which shatter when layers shift because like charges then repel.

Transition elements form particularly strong metallic bonds, giving them high melting and boiling points, high density, and great strength compared with s-block metals.

The reason is the involvement of 3d electrons. In an s-block metal such as sodium, only the outer 3s valence electron of sodium joins the delocalized sea. In a transition element, both the 4s electrons and some of the 3d electrons can be delocalized.

• This means more delocalized electrons per atom contribute to the bonding.
• More electrons in the sea = a stronger attraction between the cations and the electron cloud.
• The atoms are also relatively small and densely packed, reinforcing the effect.

Consequences:

• High melting points - e.g. iron, titanium, tungsten melt far above sodium ($\text{Na}$ melts near $98\,^{\circ}\text{C}$, whereas $\text{Fe}$ melts above $1500\,^{\circ}\text{C}$).
• High density - small atoms packed closely with strong bonding.
• Mechanical strength and hardness - useful for structural and engineering uses.

So the extra delocalized d-electrons are the key feature distinguishing transition-metal bonding from that of the more reactive, softer, lower-melting s-block metals.