Metallic bonding is the force of attraction between valence electrons and the metal ions. It is the sharing of many detached electrons between many positive ions, where the electrons act as a "glue" giving the substance a definite structure.
The electrons and the positive ions in the metal have a strong attractive force between them. Therefore, metals often have high melting or boiling points. The principle is similar to that of ionic bonds.
Because the electrons move freely, the metal has some electrical conductivity. It allows the energy to pass quickly through the electrons, generating a current. Metals conduct heat for the same reason: the free electrons can transfer the energy at a faster rate than other substances with electrons that are fixed into position. There also are few non-metals which conduct electricity: graphite (because, like metals, it has free electrons), and ionic compounds that are molten or dissolved in water, which have free moving ions.   
Metal bonds have at least one valence electron which they do not share with neighboring atoms, and they do not lose electrons to form ions. Instead the outer energy levels (atomic orbitals) of the metal atoms overlap. They are similar to covalent bonds. Not all metals exhibit metallic bonding. For example, the mercurous ion (Hg2+
2) forms covalent metal-metal bonds.
An alloy is a solution of metals.