Class 11 Chemistry Chemical Bonding Structure Valence bond approach of Covalent bond

Valence bond approach of covalent bond

VSEPR theory does not tell us about bond parameters like directional nature, bond angle, repulsions etc.

To explain we have new theory:

  • Valence bond approach
  • Molecular orbital theory

 

  • Valence bond theory

Assumptions:

  1. According to this, the atom retains their identity even after bonding.
  2. Bond is formed due to interaction of valence electrons only.
  3. Only the valence electrons lose their identity whereas inner electrons do not participate.
  4. Stability of bond depends upon amount of energy released
  5. The molecule has minimum energy at a specific distance called inter-nuclear distance and at that the bond formation takes place.

Overlapping: sigma and pi bond

Overlapping can be defined as partial merging of bonded orbitals .More is the overlapping, stronger is the bond formed.

Types of overlapping

  • Head to head
  • Sidewise

Accordingly, the bond formed is sigma or pi bond.

Sigma bond:

  1. In this head to head overlapping occurs.
  2. More is the overlapping region more stable is the bond.
  3. The bond is stronger.
  4. This bond can exist independently.

Pi bond:

  1. It is formed by side wise overlapping of orbitals.
  2. It is not too strong as in this overlapping region is less.
  3. It is weaker bond as compared to sigma bond.
  4. It can’t exist independently, it exist along with sigma bond.

 Let’s make Sigma  bond :

 Class_11_Chemistry_Chemical_Bonding_Sigma_Bond

 

 

Delocalization of sigma electron do not take place ,whereas pi electron can be delocalized.

Directional properties of covalent bonds

Whenever the overlapping takes place it occurs in three ways:

  • Electron density High (-ve)
  • Electron density low (+ve)

If – and + lobes try to overlap then no Overlapping occurs.

 Class_11_Chemistry_Chemical_Bonding_pi_Bond

  •  Molecular orbital theory

It was developed by F.Hund and R.S Mullikan in 1932.

The features of this theory are:

  • The electrons in a molecule are present in various molecular orbital as the electrons of atom are present in different various shells.
  • The atomic orbitals of comparable energies and proper symmetry combine to form molecular orbitals.
  • In molecular orbital electrons are in influence of two or more nuclei.
  • The number of molecular orbitals formed is equal to number of atomic orbitals that combine.
  • The two orbitals formed due to combination are: Bonding and Anti -bonding.
  • The Bonding molecular orbital has lower energy and greater stability than Anti -bonding.
  • The electron probability distribution around group of nuclei is given by molecular orbital.
  • The molecular orbitals are filled in accordance to Aufbau principle, Pauli’s principle and Hund’s rule.

The linear combination of atomic orbitals to form molecular orbital takes place only if:

  • The combining orbitals must have same energy.
  • The combining orbitals must have same symmetry.
  • The combining orbitals must overlap to maximum extent.

The order to fill molecular orbital is:

Class_11_Chemistry_Chemical_Bonding_Molecular_Orbital_2

This order is for all, except Oxygen , Fluorine and Neon.

For  Oxygen , Fluorine and Neon the order is :

Class_11_Chemistry_Chemical_Bonding_Molecular_Orbital_1

Information conveyed by molecular orbital diagram :

For Oxygen like molecules :

Class_11_Chemistry_Chemical_Bonding_OxygenLike_Molecules

For Nitrogen like molecules:

Class_11_Chemistry_Chemical_Bonding_NitrogenLike_Molecules

  • Stability of molecule

         If Nb > Na Then molecule is stable

         If Nb < Na Then  molecule is un-stable

         If Nb = Na then molecule is un-stable

  • We can find the bond order : That is

         Bond Order 1 (single bond)

        Bond Order 2 (double bond)

        Bond Order 3 (triple bond)

      ( c)  Tell us about the type of bond :

  • Magnetic character :

          If unpaired electrons - paramagnetic

         If no unpaired electrons - dimagnetic

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