Valence bond theory of a covalent bond

Defects of Lewis electron-dot model : It can not explain

a) Stability of molecules which do not obey octet rule
b) Shape of molecules
c) Bond characters 


Valence bond theory (VBT) 

1)It (VBT) explains the directional nature and strength of covalent bond.
2)Covalent bond is formed by the overlapping of half filled orbitals of two bonding atoms.
3)The electrons involved in the bond formation must have opposite spin. 
4)Strength of covalent bond is directly proportional to the extent of overlapping of orbitals.
5)Covalent bond is formed in the direction in which the bonding orbital is concentrated to a maximum extent.


TYPES OF COVALENT BOND

·      Overlappingof orbitals are of two types.
a) Sigma bond  
·      A bond formed by the axial or end-on over lapping
·      Electron density of -bond is concentrated along the inter nuclear axis
b) Pi-Bond
·      A bond formed by the lateral or side wise overlapping
·      Electron density of pi bond is concentrated above and below the inter nuclear axis .
·      The extent of lateral overlapping is less than that of axial overlapping
·      Pi- bond is weaker than sigma bond
·      A double bond consists of one sigma- bond and one pi bond.
·      A triple bond consists of one sigma- bond and two pi bonds.


Example 1: Hydrogen molecule
·      Electron configuration of hydrogen is 1s1
·      1s orbitals of two hydrogen atoms overlap each other
·      The bond electron pair is shared equally by two nuclei
·      The electron cloud distributes symmetrically about both the nuclei
·      The bond in H2 is known sigma s-s




 Example 2: OXYGEN MOLECULE (O2)

·      Atomic number of oxygen is 8
·      Electron configuration is 1s2 2s2 2Px2 2Py1 2Pz1
·      It has two unpaired electrons in 2Py & 2Pz orbitals
·      2Py orbital of one oxygen overlaps axially with 2Py orbital of the another oxygen forms sigma p-p bond.
·      2Pz orbitals overlap laterally to form pip-p bond.
·      Oxygen molecule has a double bond (1 sigma- 1pi)


Example 3: Nitrogen molecule (N2)

·      Atomic No. of ‘N’ is 7, Electron configuration is 1s2 2s2 2Px2 2Py2 2Pz2
·      3 unpaired electrons in 2Px, 2Py and 2Pz orbitals.
·      sigma p-p bond is formed by the axial overlapping of . 2Px orbitals of 2 nitrogen atoms.
·      Lateral overlapping of 2Py and 2Pz orbitals of two ‘N’ atoms, two piP-P bonds are formed.
·      N2 has a triple bond (1 sigma and 2 pi bonds)


 Example 4: Hydrogen chloride (HCl)

·      Electron configuration of H is 1s1
·      Electron configuration of Cl is 1s2 2s2 2p6 3s2 3Px2 3Py2 3Pz2
·      1s orbital H-atom axially overlaps with the 3P31 orbital of Cl-atom forming a  s-p bond.
·      Electro negativity of Cl-atom is more than that of H-atom.


Hydrogen chloride (HCl) 
·      The bond electron pair is attracted to a greater extent by the chlorine.
·      The electron cloud of the bonding orbital is concentrated more at the Cl-atom.
·      ‘Cl’ gets a partial negative charge
·      ‘H’ gets a partial positive charge
·      HCl is a polar molecule 


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