Q62 of 68 Page 48

Briefly describe the valence bond theory of covalent bond formation by taking an example of hydrogen. How can you interpret energy changes taking place in the formation of dihydrogen?

The valence bond theory is based on the knowledge of atomic orbitals and electronic configurations of elements, overlap criteria of atomic orbitals and stability of the molecule.

The main points of the valence bond theory are:


(i) Atoms do not lose their identity even after the formation of the molecule.


(ii) The bond is formed due to the interaction of only the valence electrons as the two atoms come close to each other. The inner electrons do not participate in bond formation.


(iii) During the formation of a bond, only the valence electrons from each bonded atom lose their identity. The other electrons remain unaffected.


(iv) The stability of the bond is accounted for by the fact that the formation of the bond is accompanied by release of energy. The molecule has minimum energy at a certain distance between the atoms known as inter-nuclear distance. Larger the decrease in energy, stronger will be the bond formed.


Consider two hydrogen atoms A and B approaching each other having nuclei H and H and the corresponding electrons e- and e- respectively.


When atoms come closer to form molecules new forces begin to operate.


(a) The force of attraction between the nucleus of the atom and electron of another atom.


(b) The force of repulsion between two nuclei of the atom and electron of two atoms.


When two hydrogen atoms are at a farther distance, there is no force operating between them, when they start coming closer to each other, the force of attraction comes into play and their potential energy starts decreasing. As they come closer to each other potential goes on decreasing, but a point is reached when potential energy acquires minimum value.


(a) This distance corresponding to this minimum energy value is called the distance of maximum possible approach, i.e. the point which corresponds to minimum energy and maximum stability.


(b) If atoms come further closer than this distance of maximum possible approach, then potential energy starts increasing and the force of repulsion comes into play and molecules start becoming unstable.


More from this chapter

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60

(i) Discuss the significance/ applications of dipole moment.

(ii) Represent diagrammatically the bond moments and the resultant dipole moment in CO2 , NF3 and CHCl3.


 61

Use the molecular orbital energy level diagram to show that N2 would be expected to have a triple bond, F2, a single bond and Ne2, no bond.

 63

Describe hybridisation in the case of PCl5 and SF6. The axial bonds are longer as compared to equatorial bonds in PCl5 whereas in SF6 both axial bonds and equatorial bonds have the same bond length. Explain.

 64

(i)Discuss the concept of hybridization. What are its different types in a carbon atom?

(ii) What is the type of hybridisation of carbon atoms marked with star.



Comprehension given below is followed by some multiple choice questions. Each question has one correct option. Choose the correct option.


Molecular orbitals are formed by the overlap of atomic orbitals. Two atomic orbitals combine to form two molecular orbitals called bonding molecular orbital (BMO) and anti bonding molecular orbital (ABMO). Energy of anti bonding orbital is raised above the parent atomic orbitals that have combined and the energy of the bonding orbital is lowered than the parent atomic orbitals. Energies of various molecular orbitals for elements hydrogen to nitrogen increase in the order :


σσ < σ *1s < σ2s < (π2px = π2py) < σ2p2 < σ2pz < (π * 2px = π*2py) < σ* 2pz s ( p p) p ( * p * p) * p < and for oxygen and fluorine order of energy of molecular orbitals is given below :


σls < σ* 1s < σ2s < σ* 2s < σ2pz < (π2px = π2py) < (π* 2px = π * 2py) < σ* 2pz


Different atomic orbitals of one atom combine with those atomic orbitals of the second atom which have comparable energies and proper orientation. Further, if the overlapping is head on, the molecular orbital is called ‘Sigma’, (σ) and if the overlap is lateral, the molecular orbital is called ‘pi’, (π). The molecular orbitals are filled with electrons according to the same rules as followed for filling of atomic orbitals. However, the order for filling is not the same for all molecules or their ions. Bond order is one of the most important parameters to compare the strength of bonds.