Solides métalliques

A metallic bond forms when electron clouds of neighboring metal atoms overlap, causing valence electrons to become delocalized throughout the solid. Unlike ionic or covalent bonds, metallic bonding creates a flexible lattice of stationary metal cations surrounded by a sea of mobile electrons. This unique arrangement allows metals to conduct electricity and heat exceptionally well while remaining malleable and ductile.

Metallic solids conduct electricity and heat due to their delocalized electrons, which move freely throughout the structure. These mobile electrons can transfer energy and charge rapidly across the material without being bound to specific atoms. This electron mobility is a direct consequence of metallic bonding and explains why metals like copper and aluminum are preferred for electrical wiring and heat transfer applications.

Hexagonal close-packing (hcp) arranges metal atoms in alternating ABAB layers, where the third layer repeats the first layer's position. Cubic close-packing (ccp) uses three repeating ABCABC layers, with the fourth layer aligning with the first. Both achieve 74% volume efficiency and give each atom a coordination number of twelve, but they differ in layer stacking patterns. Zinc exhibits hcp structure, while copper displays ccp structure.

Metallic solids can bend and deform because their flexible internal structure allows layers of atoms to slide past one another while maintaining metallic bonding. The delocalized electron sea accommodates atomic displacement without breaking bonds, unlike rigid ionic or molecular solids. This malleability makes metals valuable construction materials, as they absorb stress through deformation rather than fracturing.

Metallic bonding strength varies significantly among metals, directly affecting their melting points. Mercury is liquid at room temperature due to weak metallic bonds, while transition metals like tungsten melt above 1000°C because of strong metallic interactions. Alkali metals melt below 200°C, reflecting intermediate bond strengths. These differences in bonding strength determine how much thermal energy is required to disrupt the metallic lattice.

The coordination number in metallic solids is twelve, meaning each atom contacts six atoms in its own layer, three in the layer above, and three in the layer below. This high coordination number is significant because it maximizes attractive forces between atoms and minimizes total intermolecular energy, making close-packing arrangements the most stable and common configuration in metals.

Metallic solids exhibit high thermal and electrical conductivity, metallic luster, malleability, and variable hardness. Their malleability allows them to deform under pressure without shattering, making them ideal for construction. Many metals are also very strong, and their ability to conduct heat and electricity efficiently makes them suitable for diverse industrial applications from building frameworks to electrical systems.