Valence Bond Theory - UNSOLVED PRACTICE SET

Q7. State the basic postulates of valence bond theory.

Q8. Differentiate between sigma and pi bonds on the basis of:

(a) Type of orbital overlap

(b) Bond strength

(c) Ability to rotate

Q9. Why is a sigma bond stronger than a pi bond? 

Q10. Explain the formation of H₂ molecule using valence bond theory. 

Q11. Your dance teacher explains that when two dancers face each other and hold hands in front (end-to-end), they form a strong, stable connection. But when they stand side by side and link arms (sideways), the connection is weaker and they can't spin as easily. How is this exactly analogous to sigma and pi bonds in valence bond theory?

Q12. Explain the formation of N₂ molecule using valence bond theory, identifying the number and types of bonds. 

Q13. Explain valence bond theory and its main postulates. Describe the formation of:

(a) H₂ molecule (overlap of two 1s orbitals)

(b) F₂ molecule (overlap of two 2p orbitals)

(c) HF molecule (overlap of 1s and 2p orbitals)

(d) O₂ molecule (overlap of p orbitals forming one σ and one π bond)

Illustrate with diagrams showing orbital overlap.

Q14. Compare sigma and pi bonds in detail. Discuss:

(a) The nature of orbital overlap in each

(b) Their relative strengths and why

(c) Their symmetry about the internuclear axis

(d) Why pi bonds restrict rotation around the bond axis

Explain why a C=C double bond is not twice as strong as a C-C single bond.

Q15. During a chemistry visualization activity, you are asked to model orbital overlap using foam balls or computer simulations.

(a) You bring two spherical foam balls together to represent two hydrogen atoms forming H₂. What type of bond forms? Why is the electron density maximum between the two nuclei?

(b) You try to model the F₂ molecule using two dumbbell-shaped p orbitals. Why must the overlap be end-to-end for a sigma bond, and what happens if you try sideways overlap first?

(c) For the O₂ molecule, you need to show one sigma and one pi bond. Explain why the two p orbitals that form the pi bond must be perpendicular to the one that forms the sigma bond.

(d) Valence bond theory explains why O₂ is paramagnetic, but only if you consider the molecular orbital picture. Why does the simple VB theory fail to explain the paramagnetism of O₂?

Q16. The bond dissociation energies are given as:

C-C single bond: 347 kJ/mol

C=C double bond: 614 kJ/mol

C≡C triple bond: 839 kJ/mol

(a) Calculate the ratio of the double bond energy to the single bond energy. Is it exactly 2:1?

(b) Calculate the ratio of the triple bond energy to the single bond energy. Is it exactly 3:1?

(c) Explain why a double bond is not twice as strong as a single bond, using the concept of sigma and pi bond strengths.

Q17. Consider the following molecules and their bond types:

(a) C₂H₂ (ethyne): Identify the number of sigma and pi bonds. Calculate the total bond energy if C-H = 413 kJ/mol, C≡C = 839 kJ/mol.

(b) C₂H₄ (ethene): Identify the number of sigma and pi bonds. Calculate the total bond energy if C=C = 614 kJ/mol.

(c) C₂H₆ (ethane): Identify the number of sigma bonds. Calculate the total bond energy if C-C = 347 kJ/mol.

(d) Compare the total bond energies and explain why C₂H₂ is most thermodynamically stable per carbon atom.

Q18. The nitrogen molecule (N₂) has a bond dissociation energy of 945 kJ/mol, while the oxygen molecule (O₂) has 498 kJ/mol.

(a) Using valence bond theory, explain why N₂ has a triple bond while O₂ has a double bond.

(b) Calculate the average bond energy per electron pair for N₂ and O₂.

(c) The N≡N triple bond is one of the strongest in chemistry. Explain why this makes nitrogen gas very unreactive at room temperature.

(d) In contrast, the O=O double bond is relatively weaker. Explain why oxygen is more reactive than nitrogen, and why this reactivity is essential for life.