In metallic bonds, the valence electrons are delocalized, meaning they are not confined to a specific atom or molecule. This is in contrast to other types of chemical bonds, such as covalent bonds and ionic bonds, where the valence electrons are localized between two or more atoms.

The delocalization of valence electrons in metallic bonds is due to the low ionization energy of metals. This means that it is relatively easy for metal atoms to lose their valence electrons, which then become free to move about the metal lattice. The free electrons are what give metals their characteristic properties, such as high electrical and thermal conductivity, as well as their shiny appearance.

The delocalization of valence electrons in metallic bonds also has an impact on the strength of the bond. Metallic bonds are generally weaker than covalent bonds and ionic bonds. This is because the delocalized electrons are not as strongly attracted to the positively charged metal ions as they would be to the nuclei of atoms in a covalent or ionic bond.

Despite their weaker strength, metallic bonds are still able to hold metals together in a solid state. This is because the large number of delocalized electrons creates a strong electrical attraction between the positively charged metal ions and the negatively charged electrons. This attraction is strong enough to overcome the repulsive forces between the positively charged metal ions, keeping the metal in a solid state.