Recaps.

Metallic Bonding
1)-The electromagnetic interaction between delocalized electrons, called conduction electrons, and the metallic nuclei within metals.
-Electropositive elements.
Understood as the sharing of "free" electrons among a lattice of positively-charged ions (cations), metallic bonding is sometimes compared with that of molten salts; however, this simplistic view holds true for very few metals.
2)In a more quantum-mechanical view, the conduction electrons divide their density equally over all atoms that function as neutral (non-charged) entities.
3)Metallic bonding accounts for many physical properties of metals,such as:
-High Melting and Boiling point,(Very strong electrostatic forces of attraction between sea of electrons and the positive metal ions.)

-Strength,

-Malleability and Ductility,(Due to orderly packing of metal atoms.The layers of atom can slide over each other easily whenever a force is applied.)

-Thermal and Electrical conductivity,(They have mobile electrons to carry electricity and heat energy.)

-Opacity, and

-Luster.

4)Although the term metallic bond is often used in contrast to the term covalent bond, it is more preferable to use the term metallic bonding, because this type of bonding is collective in nature and a single "metallic bond" does not exist.
5)The atoms of a metal contribute their valence electrons to form a 'sea' of electrons.
6)The sea of electrons helps to cement the positive ions together to give metals a giant metallic structure.

Physical Properties
1)The physical properties of a substance are determined by its structure and bonding.

STRUCTURE=>PARTICLES=>BONDING BETWEEN PARTICLES=>M.P
Simple Covalent->Simple molecules->Weak Van der waals F.O.A->Low
Macromolecules->Atoms->Strong covalent bonds->High
Metallic->+Ions in "sea" of e- =>Strong metallic Bonds->High
Ionic->Ions->Strong Ionic Bonds->High

The nature of metallic bonding
-The combination of two phenomena gives rise to metallic bonding: delocalization of electrons and the availability of a far larger number of delocalized energy states than of delocalized electrons. The latter could be called electron deficiency.

Delocalization

In 2D
Delocalization -bonding that involves more than one pair of atoms held together by one pair of electrons- is most familiar from the example of benzene C6H6, where six electrons from six carbon atoms are engaged in joint aromatic bonding. However, there are other examples like the three-center two-electron bond, prevalent in boron chemistry. The principle can easily be extended over larger aromatic molecules like naphthalene, anthracene and if the process is taken to its extreme: graphite. The latter is an example of a system delocalized in two dimensions. Interestingly, there is an isoelectronic analog of benzene, B3N3H6 (borazine) for which the same argument holds. It has very similar properties to benzene[5] When extended infinitely hexagonal boron nitride BN is obtained with a structure identical to that of graphite, apart from the alternation between boron and nitrogen in each ring. This material is a semiconductor, exemplifying that delocalization is a necessary but not sufficient requirement for conductivity. Electrical conductivity does occur in graphite, because the π and π*-like bands overlap, making it a semimetal, with partly filled bands, fulfilling the other requirement for conductivity.

In 3D
Metal aromaticity in metal clusters is another example of delocalization, this time often in three-dimensional entities. Metals take the delocalization principle to its extreme and one could say that a crystal of a metal represents a single molecule over which all conduction electrons are delocalized in all three dimensions. This means that inside the metal one can generally not distinguish molecules so that the metallic bonding is neither intra- nor intermolecular. 'Nonmolecular' would perhaps be a better term. Metallic bonding is mostly non-polar, because even in alloys there is little difference among the electronegativities of the atoms participating in the bonding interaction (and in pure elemental metals, none at all). Thus metallic bonding is an extremely delocalized communal form of covalent bonding. In a sense metallic bonding is not a 'new' type of bonding at all therefore and it only describes the bonding as present in a chunk of condensed matter, be it crystalline solid, liquid or even glass. Metallic vapors by contrast are often atomic (Hg) or at times contain molecules like Na2 held together by a more conventional covalent bond. This is why it is not correct to speak of a single 'metallic bond'.
The delocalization is most pronounced for s- and p-electrons. For cesium it is so strong that the electrons are virtually free from the cesium atoms to form a gas only constrained by the surface of the metal. For cesium therefore the picture of Cs+-ions held together by a negatively charged electron gas is not too inaccurate [6]. For other elements the electrons are less free, in that they still experience the potential of the metal atoms, sometimes quite strongly. They require a more intricate quantum mechanical treatment (e.g. tight binding) in which the atoms are viewed as neutral much like the carbon atoms in benzene. For d- and especially f-electrons the delocalization is not strong at all and this explains why these electrons are able to continue behaving as unpaired electrons that retain their spin, adding interesting magnetic properties to these metals.

Electron deficiency and mobility

Metal atoms contain few electrons in their valence shells relative to their periods or energy levels. They are electron deficient elements and the communal sharing does not change that. There remain far more available energy states than there are shared electrons. Both requirements for conductivity are therefore fulfilled: strong delocalization and partly filled energy bands. Such electrons can therefore easily change from one energy state into a slightly different one. Thus, not only do they become delocalized, forming a sea of electrons permeating the lattice, but they are also able to migrate through the lattice when an external electrical field is imposed, leading to electrical conductivity. Without the field there are electrons moving equally in all directions. Under the field some will adjust their state slightly, adopting a different wave vector. Consequently, there will be more moving one way than the other and a net current will result.

The freedom of conduction electrons to migrate also gives metal atoms, or layers of them, the capacity to slide past each other. Locally bonds can easily be broken and replaced by new ones after the deformation. This process does not affect the communal metallic bonding very much. This gives rise to metals' typical characteristic phenomena of malleability and ductility. This is particularly true for pure elements. In the presence of dissolved impurities the defects in the lattice that function as cleavage points may get blocked and the material becomes harder. Gold for example is very soft in pure form (24 kt), which is why for jewelry alloys of 18 kt or lower are preferred.

Metals are typically also good conductors of heat, but the conduction electrons only contribute partly to this phenomenon. Collective (i.e. delocalized) vibrations of the atoms known as phonons that travel through the solid as a wave, contribute strongly.
However, the latter also holds for a substance like diamond. It conducts heat quite well but not electricity. The latter is not a consequence of the fact that delocalization is absent in diamond, but simply that carbon is not electron deficient . The position of carbon in the middle of its period in the Periodic Table makes that there are precisely enough electrons to fill the energy states. Under a field electrons are not able to adopt a different wave vector because there are no empty states to move into. This makes a current impossible in this wide band gap semiconductor. However, as soon as charge carriers are introduced by doping the crystal with a suitable impurity the resulting charge carriers are as mobile as in a metal, though far fewer in number. Even without doping the vibrational motions (the phonons) are delocalized over the crystal explaining the heat conduction. Still the bonding in diamond is better described as covalent than as metallic if only because there is a very strong directional preference for tetrahedral stacking, producing a structure that is extremely hard to deform and by no means close packed.
Clearly, the electron deficiency is an important point in distinguishing metallic from more conventional covalent bonding. Thus, we should amend the expression given above into:
Metallic bonding is an extremely delocalized communal form of electron deficient[7] covalent bonding.