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In metallic bonding, metal atoms release their valence electrons into a shared 'sea' of delocalized electrons that flows freely throughout the solid structure. The resulting lattice of positive metal cations is held together by electrostatic attraction to this mobile electron sea. This model explains characteristic metallic properties: electrical and thermal conductivity (mobile electrons carry charge and energy), malleability and ductility (cation layers can slide without breaking the delocalized bonding), and metallic luster (free electrons interact with and reflect visible light).
Compare metallic bonding to ionic and covalent bonding by contrasting their properties — conductivity, hardness, melting point. Use the trend in metallic properties across the d-block to see how electron count affects bond strength.
You already know from periodic trends that metals sit on the left and center of the periodic table and tend to have low ionization energies — they give up valence electrons readily. From ionic bonding, you learned how electrons can transfer entirely from one atom to another. Metallic bonding represents a third possibility: instead of transferring electrons to a specific partner, metal atoms collectively release their valence electrons into a shared pool that belongs to no individual atom. The result is a lattice of positively charged metal cations immersed in a "sea" of delocalized electrons that flows freely throughout the entire solid.
This electron sea model explains why metals behave so differently from ionic or covalent solids. In an ionic crystal like NaCl, each ion is locked in place by directional electrostatic attraction to its specific neighbors — if you try to shift one layer, like charges suddenly face each other and the crystal shatters. In a metal, the bonding is non-directional: the electron sea glues the cations together regardless of their exact positions. When you hammer a metal, the cation layers slide past one another, but the delocalized electrons simply redistribute to maintain bonding in the new configuration. This is why metals are malleable (can be hammered into sheets) and ductile (can be drawn into wires), while ionic solids are brittle.
Electrical conductivity follows directly from electron delocalization. When a voltage is applied across a metal wire, the free electrons drift toward the positive terminal — they are already mobile and require no energy to break free from individual bonds. This is fundamentally different from ionic conduction, which requires ions to physically migrate through a liquid or molten salt. Thermal conductivity works similarly: the mobile electrons efficiently transfer kinetic energy from hotter regions to cooler ones, supplementing the slower vibration-based heat transfer through the cation lattice. Metallic luster arises because free electrons can absorb and re-emit photons across a broad range of visible wavelengths, giving metals their characteristic reflective appearance.
The strength of metallic bonding varies systematically across the periodic table. Metals with more valence electrons and smaller atomic radii form stronger metallic bonds — the electron sea is denser and the cations are closer together. This is why transition metals in the middle of the d-block (like tungsten and chromium) generally have higher melting points and greater hardness than alkali metals like sodium, which contribute only one electron each to a sea spread across large, widely spaced cations. Alloying — mixing two or more metals — works because the electron sea accommodates different-sized cations, and the size mismatch can actually strengthen the material by disrupting the regular sliding of cation layers.