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Theory:

Electronic theory of valence or Octet rule:
 
In \(1916\), Kossel and Lewis proposed the Electronic theory of valence or Octet rule, based on the noble gas electronic configuration, to explain chemical combinations between atoms.

According to this, atoms of all elements except inert gases combine to form molecules because they have an incomplete valence shell and tend to achieve a stable electronic configuration similar to noble gases.
Atoms can combine by transferring valence electrons from one atom to another or sharing valence electrons to form the stable outer shell of eight electrons.
The tendency of atoms to have eight electrons in the valence shell is known as the ‘Octet rule’ or the ‘Rule of eight.’
Example:
1. Sodium with atomic number \(11\) will quickly lose one electron to attain neon’s stable electronic configuration.
 
14.png
Stable electronic configuration of Sodium  
 
2. Chlorine has an electronic configuration of \(2\), \(8\), \(7\). Therefore, one extra electron is required to reach the nearest noble gas (argon) configuration. As a result, chlorine readily accepts one electron from another atom and achieves a stable electronic configuration. We can conclude that the elements with a stable valence shell (eight electrons) tend to lose or gain electrons.
 
2.png
Stable electronic configuration of Chlorine 
Which atoms tend to lose electrons?
Which ones are more likely to gain electrons?
 
Atoms with \(1\), \(2\), \(3\) valence electrons tend to lose electrons, whereas atoms with \(5\), \(6\), \(7\) valence electrons prefer to gain electrons.
  
Unstable electronic configuration:
 
Element Atomic numberElectron distributionValence electrons
Boron
\(5\)
\(2\), \(3\)
\(3\)
Nitrogen
\(7\)
\(2\), \(5\)
\(5\)
Oxygen
\(8\)
\(2\), \(6\)
\(6\)
Sodium
\(11\)
\(2\), \(8\), \(1\)
\(1\)