Chemistry: Chemical Bonding: Modern Theories of Chemical Bonding (For CBSE, ICSE, IAS, NET, NRA 2022)

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Modern Theories of Chemical Bonding

  • The first chemical bonding theory was proposed in 1916 by Kossel and Lewis
  • This theory doesn՚t involve the wave mechanical or quantum mechanical principles.
  • After the development of quantum mechanical description of atomic structure two more theories were proposed to explain the bonding between atoms.
  • The modern theories of chemical bonding are Valence Bond Theory (VBT) and Molecular Orbital Theory (MOT) .

Valence Bond Theory

  • The valence bond theory was first proposed by Heitler in London in 1927.
  • It describes the formation of hydrogen molecule from its atoms.
  • Later, Linus Pauling and others developed it by the process of chemical bond formation can be visualized as the overlapping of atomic orbitals of the two atoms.
  • The strength of the bond depends on the effectiveness of the overlapping. So, strong bond is formed when the overlapping of the orbitals is high. Example: Hydrogen atom
  • The two hydrogen atoms are at infinite distance from each other.
  • Their electrons are in 1s orbitals and they are influenced by the corresponding nuclei.
  • As the two atoms approach each other their 1s orbitals begin to overlap which lead to decrease in energy.
  • At a distance equal to the bond length the overlapping is maximum, and the energy is minimum.
  • The electrons occupying the shared region of orbitals are influenced of both the nuclei.
  • This simple approach can be explained in the bonding in diatomic molecules like HF, F2 etc.
  • The bonding in molecules containing more than two atoms are further explained by some additional concepts like excitation and hybridization.


  • Hybridization is the process of forming new orbitals by mixing two atomic orbital in a particular atom.
  • The new hybrid orbitals that are formed have same orbitals and energy.
  • Two main characteristics of hybridization are:
    • The number of hybrid orbitals formed is the same as the number of atomic orbitals which undergoes hybridization.
    • All the new hybrid orbitals that are formed are identical in their shapes and energy.
  • Beryllium hydride (BeH2) is taken as example to understand the concept of hybridization of orbitals.
  • The atomic number of beryllium is 4 and its electronic configuration is 1s2 2s2.
  • In order to form bonds with two hydrogen atoms the valence electrons (2s2) of beryllium atom must overlap with the 1s electrons of the two hydrogen atoms.
  • But, the valence shell of beryllium atom contains electrons in the same orbital (i.e.. , 2s) so it cannot overlap with the 3p orbital of hydrogen atoms.
  • Pauling suggest that in the process of bond formation an electron from the 2s orbital of beryllium atom gets excited to the empty 2p orbital as shown in figure below.
Hybridization Ground and Excited State
  • Now the two valence electrons can overlap with the 1s orbitals of the two hydrogen atoms and form two bonds of different nature.
  • One of these would involve overlapping of 2s orbital of beryllium with 3p orbital of chlorine while the other would involve overlapping of 2p orbital of beryllium with 2p orbital of chlorine.
  • According to this two or more nonequivalent orbitals of comparable energies hybridize and give rise to an equal number of equivalent hybrid orbitals.
  • In case of BeCl2 the two singly occupied orbitals (2s and 2p) hybridize to give two sp- hybrid orbitals. This is called sp hybridization.
  • These hybrid orbitals lie along the z- direction and point in opposite directions.
Hybrid Orbitals Lie

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