Metallic Bond: Definition, Properties, and Examples

A metallic bond is a type of chemical bond formed between positively charged atoms in which the free electrons are shared among a lattice of cations. In contrast, covalent and ionic bonds form between two discrete atoms. Metallic bonding is the main type of chemical bond that forms between metal atoms.

Metallic bonds are seen in pure metals and alloys and some metalloids. For example, graphene (an allotrope of carbon) exhibits two-dimensional metallic bonding. Metals, even pure ones, can form other types of chemical bonds between their atoms. For example, the mercurous ion (Hg₂²⁺) can form metal-metal covalent bonds. Pure gallium forms covalent bonds between pairs of atoms that are linked by metallic bonds to surrounding pairs.

How Metallic Bonds Work

The outer energy levels of metal atoms (the s and p orbitals) overlap. At least one of the valence electrons participating in a metallic bond is not shared with a neighbor atom, nor is it lost to form an ion. Instead, the electrons form what may be termed an "electron sea" in which valence electrons are free to move from one atom to another.

The electron sea model is an oversimplification of metallic bonding. Calculations based on electronic band structure or density functions are more accurate. Metallic bonding may be seen as a consequence of a material having many more delocalized energy states than it has delocalized electrons (electron deficiency), so localized unpaired electrons may become delocalized and mobile. The electrons can change energy states and move throughout a lattice in any direction.

Bonding can also take the form of metallic cluster formation, in which delocalized electrons flow around localized cores. Bond formation depends heavily on conditions. For example, hydrogen is a metal under high pressure. As pressure is reduced, bonding changes from metallic to nonpolar covalent.

Relating Metallic Bonds to Metallic Properties

Because electrons are delocalized around positively charged nuclei, metallic bonding explains many properties of metals.

Electrical conductivity: Most metals are excellent electrical conductors because the electrons in the electron sea are free to move and carry charge. Conductive nonmetals (such as graphite), molten ionic compounds, and aqueous ionic compounds conduct electricity for the same reason—electrons are free to move around.

Thermal conductivity: Metals conduct heat because the free electrons are able to transfer energy away from the heat source and also because vibrations of atoms (phonons) move through a solid metal as a wave.

Ductility: Metals tend to be ductile or able to be drawn into thin wires because local bonds between atoms can be easily broken and also reformed. Single atoms or entire sheets of them can slide past each other and reform bonds.

Malleability: Metals are often malleable or capable of being molded or pounded into a shape, again because bonds between atoms readily break and reform. The binding force between metals is nondirectional, so drawing or shaping a metal is less likely to fracture it. Electrons in a crystal may be replaced by others. Further, because the electrons are free to move away from each other, working a metal doesn't force together like-charged ions, which could fracture a crystal through the strong repulsion.

Metallic luster: Metals tend to be shiny or display metallic luster. They are opaque once a certain minimum thickness is achieved. The electron sea reflects photons off the smooth surface. There is an upper-frequency limit to the light that can be reflected.

The strong attraction between atoms in metallic bonds makes metals strong and gives them high density, high melting point, high boiling point, and low volatility. There are exceptions. For example, mercury is a liquid under ordinary conditions and has a high vapor pressure. In fact, all of the metals in the zinc group (Zn, Cd, and Hg) are relatively volatile.

How Strong Are Metallic Bonds?

Because the strength of a bond depends on its participant atoms, it's difficult to rank types of chemical bonds. Covalent, ionic, and metallic bonds may all be strong chemical bonds. Even in molten metal, bonding can be strong. Gallium, for example, is nonvolatile and has a high boiling point even though it has a low melting point. If the conditions are right, metallic bonding doesn't even require a lattice. This has been observed in glasses, which have an amorphous structure.