Metallic bonding is a type of chemical bonding that rises from the attractive electrostatic force between conduction electrons (in the form of an electron cloud of delocalized electrons) and positively charged metal ions. It may be described as the sharing of free electrons among a structure of positively charged ions (cations).
Metal atoms typically have a few valence electrons that are easily relocated into a lattice of positively charged metal atoms. One can visualize this type of bonding by imagining a metal as a lattice of positive ions held together by a cloud of electrons. In other words valence electrons are delocalized, that is not bound to a specific atom: they are distributed along the surface of the metal, as to form an “electronic cloud” that surrounds all the ions of the element.
The presence of this “electronic cloud” allows the maximum approach of the ions without repulsive forces that push them away, implying a high density. Mobile electrons also explain properties such as high thermal and electrical conductivity: because they are free to move, they facilitate the processes of heat and electrical energy diffusion. The metallic bond belongs to the family of strong bonds (along with covalent and ionic): the high binding energies (between ion and ion) imply high temperatures of fusion and boiling. It is non-directional, i.e. it has the property of being equal in the three coordinated directions; there are, therefore, no preferences related to the direction to establish the chemical bond between the ions of the element nor particular restrictions in the relative positions occupied by them. The immediate consequence of the directionality is the unalteration of the bonding interactions following the shift of two lattice planes.
If, for example, we exerted a shear force on a lattice plane, it would produce a deformation of the metal, but not a rupture of the crystal lattice. In chemical terms, the only effect realized by this force is a sliding of the lattice plane along which it is applied with respect to the lattice plane below. The metal deforms because of the sliding, but it doesn’t split because it doesn’t create repulsive forces (ions are all positive, valence electrons are always delocalized, the chemical state remains unchanged!); the bond remains. The shear force moves the lattice planes: the instantaneous effect of the sliding is a coincidence of negative and positive charges that exert a great repulsive force; the chemical bond is altered, the lattice is broken. This explains why metals are ductile and malleable, while amorphous solids like glass are notoriously brittle.
As in the case of ionic bonding, there are not real molecules but reticular aggregates of metal atoms held together by this electrostatic force.
This model explains some properties of metals such as their high electrical conductivity (in fact, since these electrons are not bound to any particular atom, they are extremely mobile) and thermal conductivity, their malleability and ductility. Heat conduction and their opacity and brightness are related to the mobility of valence electrons (delocalized electrons) that increase their kinetic energy while ductility and malleability are explained by the free mutual sliding of lattice planes (non-directed bonds), which does not cause the destruction of the crystalline building because the bond is not made by few localized electrons, but by all available electrons. The presence of strong bonds within the metal bond also explains other characteristics of metals themselves, namely high density, non-solubility and very high boiling and melting points.
Another model used to interpret the metallic bond is that of plane wave expansion, which consists in representing the wave function as a linear combination of plane waves, which produces a partial localization of free electrons, this model is applicable when there are not large variations of the crystalline potential
Metal bonding is also the bond formed in brazing processes between filler and base metals.