Molecular Orbital Theory - UNSOLVED PRACTICE SET

Q7. State the main postulates of molecular orbital theory. How does it differ from valence bond theory? 

Q8. What are bonding and antibonding molecular orbitals? How are they represented?

Q9. Write the molecular orbital configuration for N₂ molecule. Calculate its bond order. 

Q10. Why does O₂ exhibit paramagnetic behaviour according to molecular orbital theory?

Q11. Your music teacher explains that when two singers sing the same note in phase, the sound is louder (constructive interference). When they sing out of phase, the sound cancels out (destructive interference). How is this exactly analogous to the formation of bonding and antibonding molecular orbitals from atomic orbitals?

Q12. Compare the relative energies of σ2p and π2p molecular orbitals for O₂ and N₂. Why is the order different? 

Q13. Explain molecular orbital theory and its main postulates. Describe the formation of molecular orbitals from atomic orbitals using the LCAO method. Draw energy level diagrams for:

(a) O₂

(b) N₂

Write their molecular orbital configurations and calculate bond orders.

Q14. Compare valence bond theory and molecular orbital theory. Explain why MOT is superior to VBT in explaining:

(a) The paramagnetism of O₂

(b) The non-existence of He₂ molecule

(c) The bond order of species like O₂⁺, O₂⁻, and O₂²⁻

(d) The stability of H₂⁺ ion

Q15. During an advanced chemistry discussion, your teacher presents the following diatomic species: H₂, He₂, Li₂, Be₂, B₂, C₂, N₂, O₂, F₂, Ne₂.

(a) Which of these exist as stable molecules according to MOT? Explain using bond order calculations.

(b) B₂ is found to be paramagnetic with two unpaired electrons. Explain this using MOT and the σ2p-π2p energy level ordering.

(c) C₂ has a bond order of 2 but no traditional sigma bond between the carbons. Explain this unusual bonding using MOT.

(d) Compare the bond orders of O₂, O₂⁺, O₂⁻, and O₂²⁻. Arrange them in order of increasing bond length and explain your reasoning.

Q16. Write the molecular orbital electronic configurations for the following species and calculate their bond orders:

(a) H₂

(b) He₂⁺

(c) Li₂

(d) Be₂

(e) B₂

(f) C₂

(g) N₂

(h) O₂

(i) F₂

(j) Ne₂

For each, predict whether the species is stable, and if stable, whether it is diamagnetic or paramagnetic.

Q17. The following data is given for oxygen species:

(a) Complete the table by calculating the bond order for each species using MOT.

(b) Arrange the species in order of increasing bond length.

(c) Arrange the species in order of increasing bond energy.

(d) O₂⁻ is called the superoxide ion and O₂²⁻ is the peroxide ion. Explain how the bond order relates to their reactivity and the O-O bond length in each.

Q18. Consider the following molecules and their properties:

(a) For N₂, write the complete molecular orbital configuration (up to σ2p and π2p levels). Calculate the bond order. Why is N₂ so chemically inert?

(b) For NO (nitric oxide), which has 15 electrons, write the molecular orbital configuration. Calculate the bond order and predict whether it is paramagnetic or diamagnetic.

(c) The bond dissociation energy of N₂ is 945 kJ/mol while that of N₂⁺ is 842 kJ/mol. Explain why removing an electron from N₂ weakens the bond, using MOT.

(d) CO and N₂ are isoelectronic (both have 14 electrons). Compare their bond orders, bond lengths, and magnetic properties using MOT.