Metallic bonding theory

Metallic bonding


Metallic bonding (metal-metal):

This type of bond is given in all pure metals and most alloys.
In a metal, atoms are very close to each other forming what is called a “compact packaging”; each atom around it has eight or twelve nearest neighbors. Thus, the valence orbitals of the metal atoms widely overlap each other resulting in a molecular orbital, which extends throughout the metal. The valence electrons that occupy this molecular orbital, cannot be said to belong to a particular atom, but rather to the whole being, as they say, completely delocalized.
The metal structure can be considered, therefore, it consists of a set of positive ions, tightly packed and surrounded by a number of free electrons (valence electrons), which form a kind of electronic fluid called sometimes sea electrons. These electrons can move relatively easily through all metal, which explains the high electrical and thermal conductivity.
Ions that form the metal net are all the same, so they can be moved from a position to other, relatively easily, which also explains the ductility and malleability of the metals, as shown below.
Any theory of the metallic bond should give explanation to this freedom of movement and to address the remaining properties of metals. These are two: the so-called theory of the electron gas or the sea of ​​electrons and band theory.


In electronic gas theory (also called sea of ​​electrons or electron cloud), metal atoms lose their valence electrons and form a dense network of cations.
For example, in the case of sodium, whose electronic configuration is 2s2 1s2 2p6 3s1;
The Na + cations, formed by atomic nuclei and electrons of the inner layers, are packaged and the valence electrons move freely. These electrons don’t belong to individual atoms, but are common to all atoms that form the network. It is said that electrons are delocalized.
Depending on the number of valence electrons that the metal has, there will be as many electrons delocalized as atoms or even more. For example, in sodium, which loses an electron, there will be many electrons as sodium atoms, but magnesium, which has two valence electrons and loses both, there will be twice electron nuclei Mg (2+) .
Thus, the cations are arranged forming a metallic compact or packaging crystal lattice and each cation surrounds the maximum number of neighbors cations. Valence electrons move freely through the interstices of the grid, forming the electron gas and also acting as a cushion which prevents repulsion between different cations.
Given the freedom of movement of valence electrons, this theory explains the metallic bond very well many of the metallic properties, such as high electrical and thermal conductivity properties. It also explains the ductility and malleability or resistance to deformation, because the layers of cations can slide over each other, keeping the type of structure and the bond strength.





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