Lesson 1 of 12 · 9 min
Kössel-Lewis approach and the octet rule
NCERT § "Kössel-Lewis Approach to Chemical Bonding"
Sodium is a soft metal that fizzes in water; chlorine is a choking green gas. Together they make the salt Riya just stirred into her glass. What did each atom get out of the deal?
The story this chapter follows: A pinch of salt in water
The lesson in notes
In short
Kössel and Lewis explained bonding as atoms reaching a stable noble-gas-like octet, either by transferring electrons (ionic bond) or by sharing them (covalent bond).
Lewis symbols show only the valence electrons of an atom as dots around its symbol; the group valence is usually either the number of dots or 8 minus that number.
In a covalent bond each shared pair counts towards the octet of both atoms; one shared pair is a single bond, two a double bond and three a triple bond.
Formal charge on an atom in a Lewis structure = (valence electrons in the free atom) − (non-bonding electrons) − ½(bonding electrons).
When several Lewis structures can be drawn, prefer the one that keeps formal charges smallest, since it is normally the most stable; ozone, for example, carries formal charges of +1 and −1 on two of its oxygens.
Limitation 1, incomplete octet: in LiCl, BeH₂ and BCl₃ the central atom has fewer than eight electrons.
Limitation 2, odd-electron molecules: NO and NO₂ cannot give every atom an octet.
Limitation 3, expanded octet: elements from period 3 onwards can use d orbitals and hold more than eight electrons, as in PF₅, SF₆ and H₂SO₄.
The octet theory also fails to explain the shapes of molecules and their relative stability, and it ignores the fact that some noble gases (xenon, krypton) do form compounds.