Hybridisation

Full Notes

Introduction to Hybridisation

Hybridisation is the concept of mixing two or more atomic orbitals of comparable energy to form new equivalent hybrid orbitals. These new orbitals:

• Have a different shape and energy than the original atomic orbitals.
• Are oriented in space to minimise repulsion, thus explaining molecular geometry.

Introduced by Linus Pauling, this theory explains the observed shapes of molecules like CH₄, NH₃, H₂O, and others which cannot be justified by the orientation of pure s or p orbitals alone.

For example, carbon is often used to illustrate the formation of hybrid orbitals as it reacts. The different types of hybrid orbitals formed are combinations of atomic orbitals in the unhybridized carbon atom.

Main Features of Hybridisation

1. The number of hybrid orbitals is equal to the number of atomic orbitals that get hybridised.
2. The hybridised orbitals are always equivalent in energy and shape.
3. The hybrid orbitals are more effective in forming stable bonds than pure atomic orbitals.
4. These hybrid orbitals are directed in space to minimise electron pair repulsion, explaining molecular geometry.

Important Conditions for Hybridisation

• The orbitals must be from the valence shell and have comparable energy (nearly or close to equal energies).
• It is not necessary that only half-filled orbitals participate – sometimes even filled orbitals of the valence shell take part in hybridisation.

4.6.1 Types of Hybridisation

The type of hydridisation that occurs depends on the number and type of orbitals mixing together.

1. sp Hybridisation

• One s-orbital mixes with one p-orbital → forms 2 sp hybrid orbitals.
• Orbitals are linearly arranged at 180° to each other.
• Each hybrid orbital forms a σ-bond.
• Example: BeCl₂ (beryllium chloride)
  Be: 1s² 2s² → excited: 2s¹ 2p¹ → hybridises → two sp orbitals.

2. sp² Hybridisation

• One s-orbital and two p-orbitals mix → forms 3 sp² orbitals.
• Trigonal planar arrangement at 120° angles.
• Example: BCl₃ (boron trichloride)
  B: 1s² 2s² 2p¹ → excited → 2s¹ 2p² → hybridises.

3. sp³ Hybridisation

• One s-orbital and three p-orbitals → form 4 sp³ orbitals.
• Tetrahedral geometry with bond angles of 109.5°.
• Example: CH₄ (methane)
• In NH₃ and H₂O, geometry is affected by lone pairs:
  • NH₃: Trigonal pyramidal (~107°)
  • H₂O: Bent/angular (~104.5°)

4.6.2 Other Examples of sp³, sp², and sp Hybridisation

• sp³ Examples: CH₄, NH₃, H₂O, C₂H₆
  Bond angle reduces from 109.5° due to lone pair–bond pair repulsion.
• sp² Examples: C₂H₄ (ethene), SO₃, NO₃⁻
  Bond angles ~120°, π-bond present in multiple bonds.
• sp Examples: C₂H₂ (ethyne), BeF₂, CO₂
  Bond angle: 180°, linear geometry.

These molecules demonstrate how hybridisation aligns orbitals to match experimental geometry, accounting for lone pairs and multiple bonds.

4.6.3 Hybridisation of Elements Involving d Orbitals

d-orbitals can get involved in hybridisation when the central atom can expand its octet – typically in elements from the 3rd period or beyond.

sp³d Hybridisation

• Involves: one s, three p, one d orbital.
• Forms 5 hybrid orbitals.
• Trigonal bipyramidal geometry.
• Example: PCl₅

sp³d² Hybridisation

• Involves: one s, three p, two d orbitals.
• Forms 6 hybrid orbitals.
• Octahedral geometry.
• Example: SF₆

sp³d³ Hybridisation

• Involves: one s, three p, three d orbitals.
• Forms 7 hybrid orbitals.
• Pentagonal bipyramidal geometry.
• Example: IF₇

Summary

• Hybridisation involves mixing atomic orbitals of similar energy to form equivalent hybrid orbitals.
• sp, sp², and sp³ hybridisations explain linear, trigonal planar, and tetrahedral geometries respectively.
• d-orbital participation leads to sp³d, sp³d², and sp³d³ hybridisations producing trigonal bipyramidal, octahedral, and pentagonal bipyramidal geometries.
• Hybridisation helps explain bond strength, shape, and angles observed in covalent molecules.