Bonding in Metals

A metallic bond is a bond between two metal atoms. Compared to nonmetals, metals have low ionization energies, allowing them to lose valence electrons easily. This gives the metallic bond distinct properties in contrast to ionic and covalent bonds. 

Metallic bonds and most of their properties can be explained using the simple electron sea model. Consider the metal potassium. Due to the low ionization energy, each potassium atom can easily lose its valence electron to become a cation. 

These potassium cations are held together in close-packing because of their attraction to the negatively charged sea of electrons. These electrons are not confined to any single ion but are evenly distributed and relatively free to move within the metal.  

The electron sea model accounts for several important characteristics of metals. For example, when a voltage difference is applied to a metal wire, like copper wire, the negatively charged electrons move freely towards the wire’s positive end, generating an electric current. This is why most metals are excellent conductors of electricity. 

In contrast, ionic compounds are nonconductors of electricity in their solid form but can conduct electricity when dissolved in water. This is because, in a crystalline ionic bond, electrons are transferred from the metal to the nonmetal, but remain localized to one ion. 

However, when dissolved in water, the cations and anions dissociate and can move when subjected to a potential difference, creating an electric current.

Metals are also excellent thermal conductors. According to the electron sea model, when heat is applied to one end of the metal, the electrons move freely and quickly disperse the heat throughout the metal.

Metals can easily be pounded into sheets due to their malleability or into wires due to the ductility property. Since there are no localized bonds in metals, the metal atoms can slide past each other allowing easy deformity. Electrons then flow into the new shape to accommodate the deformity.