HYBRIDISATION – Simplified & Advanced Part-1

Very Important Points:

(1) Lone pairs of electrons, if involved in resonance ( or back-bonding) i.e., if lone pairs are involved in “π” bonding formation, then they are “unhybridised”. e.g, in trisilylamine, (SiH3)3N, nitrogen atom is sp2 hybridised as its lone pair is involved in back bonding with vacant ‘d’ orbital of silicon atom. Hence, lone pair orbital of nitrogen is pure p-orbital.

(2) Unshared Odd electron orbitals (i.e, Single electron orbitals) are generally unhybridised e.g, in methyl free radical, carbon atom is sp2 hybridised.

(3) Unshared Odd electron orbitals (i.e, Single electron orbitals) are hybridised if the atom is bonded to more than one highly electro-negative atoms. e.g, In ClO2, chlorine atom has one lone pair orbital, one odd electron orbital and two σ- bonds hence it is sp3 hybridised.

“Hybridisation is a hypothetical process of intermixing of “valence” atomic (pure) orbitals of similar energy (degenerate orbitals) (so as to redistribute their energies) to form same number of hybrid orbitals which are equivalent in energy or shape (but having different orientation or direction).”

Tricky Features:

(1) Hybridisation is property of an atom(except Hydrogen) hence a molecule have different (maybe same as well) hybridisation of different atoms. [Hydrogen has only one electron hence it doesn’t  exhibit hybridisation). Also, atom of same element can have different hybridisation in different bonding situations e.g Carbon atom in ethane, ethene and ethyne are sp, sp2 and sp3 hybridised respectively.

(2) By default, hybridisation of a molecule is taken as hybridisation of central atom (if there is one). In case of molecules having no central atom or more than one central atoms, their hybridisation must be separately determined.

(3) Hybridised orbitals are involved in overlapping during formation of “σ-(sigma) bonds” and ‘lone pairs’ (unshared electron pairs). Half filled valence orbitals are involved in overlapping during formation of ‘σ-(sigma) bonds’ and Fully filled valence orbitals are involved in overlapping during formation of ‘lone pairs’.

(4) “π” bond are formed by overlap of pure p or pure d or pure f- atomic orbitals hence orbitals involved in formation of π- bonds are unhybridised.  [‘s’ orbitals are never involved in formation of “π”bonds]

(5) Hybridisation must involve mixing of at least two pure atomic orbitals (obviously) and hybridisation must always involve one and only one pure “s” orbital (as they are closer to nucleus and hence provide stability to hybrid orbitals).

(6) In case of trigonal bi-pyramidal structure and pentagonal bi-pyramidal structures, hybridised orbitals are not all equivalent. These involve two types of bonds- axial and equatorial.

(7) Hybridisation is a mathematical way of describing molecule formation by elementary mixing of atomic orbitals, and was able to explain ‘equivalency of bonds’ in compounds e.g,  all bonds in methane, Beryllium chloride etc. are equivalent, which cannot be explained by involvement of pure atomic orbitals only.

Note- Hybridisation is not something that actually happens. This concept has been further advanced to include more advanced concepts as in later theories e.g Molecular orbital theory, Crystal Field theory, Ligand Field theory.


Salient Features:

(1) The number of hybrid orbitals formed is equal to the number of the atomic orbitals taking part in hybridisation.

(2) Hybrid orbitals are more “directional” than pure atomic orbitals hence more effective in forming stable bonds and hence stable molecule.

(3) Promotion of electrons is not essential for hybridisation i.e, hybridisation can take place in ground state or any of the excited states of atom.

(4) Both Half filled valence orbitals as well as Fully filled valence orbitals can take part in hybridisation.

(5) Hybrid orbitals are equivalent hence, provides stability to the molecule. (Consider molecules having different strength of σ- bonds between same pair  of elements, then weaker bond would be broken easily hence molecule would be less stable)

(6) Hybridised orbitals are spatially oriented in some specific directions to provide maximum possible distance or greater bond angles to minimise repulsion between valence shell electron pairs and hence stable arrangement. [Imagine nitrogen atom in ammonia molecule using only pure orbitals in molecule formation then lone pair of nitrogen would be in ‘s’ orbital and bonds would involve pure p-orbitals of nitrogen hence bond angle would be 90 degree, much less than tetrahedral bond angle as expected by sp3 hybridisation)


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