Metal atoms form giant crytalline structures. The atoms are packed tightly together so the outer electrons get separated from the atom. The result is a lattice structure of positive ions in a sea of free electrons.
Elements exist in different forms - solid form, liquid form, and gaseous form. Solids are essentially metallic and non metallic. The structure of metals differ from those of other forms in which elements exist. The properties of metals are determined by their structure. Metals usually have the atoms arranged closely together in a compact form. It is this compactness that gives metals the different qualities such as strength, i.e. The atoms are bonded together very strongly. Weak bonds would make for weak structures. Basically, all metals have a compact arrangement of atoms, ensuring there is minimal space between them. While the strong bonding explains the strength that metals possess, how does one explain the other properties of metals, such as malleability, ductility, conductivity, etc? The fact that metals have these properties suggest a delocalized nature of bonding. The delocalized nature, complemented by the strong bonding is what gives metals their various properties. Basically, bonding in metals happen between atoms of low electronegativity, which means that there is not too strong an attraction between the valence electrons of the metal atom. The valence electrons are the outermost electrons among all in the atom, and since these have low attractively, they can be shared with the other atoms around them, thereby strengthening the bonds between the atoms themselves. Metallic bonding differs from other kinds of bonding in this respect - the valence electrons can be shared and are therefore considered free-form.