Chapter 4 Solutions — Chemical Bonding and Molecular Structure

Exercises

1. 1.
   Explain the formation of a chemical bond.
2. 2.
   Write Lewis dot symbols for atoms of the following elements: Mg, Na, B, O, N, Br.
3. 3.
   Write Lewis symbols for the following atoms and ions: S and S²⁻; Al and Al³⁺; H and H⁻
4. 4.
   Draw the Lewis structures for the following molecules and ions: H₂S, SiCl₄, BeF₂, CO₃²⁻, HCOOH
5. 5.
   Define octet rule. Write its significance and limitations.
6. 6.
   Write the favourable factors for the formation of ionic bond.
7. 7.
   Discuss the shape of the following molecules using the VSEPR model: BeCl₂, BCl₃, SiCl₄, AsF₅, H₂S, PH₃
8. 8.
   Although geometries of NH₃ and H₂O molecules are distorted tetrahedral, bond angle in water is less than that of ammonia. Discuss.
9. 9.
   How do you express the bond strength in terms of bond order?
10. 10.
    Define the bond length.
11. 11.
    Explain the important aspects of resonance with reference to the CO₃²⁻ ion.
12. 12.
    H₃PO₃ can be represented by structures 1 and 2 shown below. Can these two structures be taken as the canonical forms of the resonance hybrid representing H₃PO₃? If not, give reasons for the same.
13. 13.
    Write the resonance structures for SO₃, NO₂ and NO₃⁻.
14. 14.
    Use Lewis symbols to show electron transfer between the following atoms to form cations and anions:

    • (a) K and S
    • (b) Ca and O
    • (c) Al and N.
15. 15.
    Although both CO₂ and H₂O are triatomic molecules, the shape of H₂O molecule is bent while that of CO₂ is linear. Explain this on the basis of dipole moment.
16. 16.
    Write the significance/applications of dipole moment.
17. 17.
    Define electronegativity. How does it differ from electron gain enthalpy?
18. 18.
    Explain with the help of a suitable example polar covalent bond.
19. 19.
    Arrange the bonds in order of increasing ionic character in the molecules: LiF, K₂O, N₂, SO₂ and ClF₃.
20. 20.
    The skeletal structure of CH₃COOH as shown below is correct, but some of the bonds are shown incorrectly. Write the correct Lewis structure for acetic acid.
21. 21.
    Apart from tetrahedral geometry, another possible geometry for CH₄ is square planar with the four H atoms at the corners of the square and the C atom at its centre. Explain why CH₄ is not square planar.
22. 22.
    Explain why BeH₂ molecule has a zero dipole moment although the Be-H bonds are polar.
23. 23.
    Which out of NH₃ and NF₃ has higher dipole moment and why?
24. 24.
    What is meant by hybridisation of atomic orbitals? Describe the shapes of sp, sp², sp³ hybrid orbitals.
25. 25.
    Describe the change in hybridisation (if any) of the Al atom in the following reaction: AlCl₃ + Cl⁻ -> AlCl₄⁻
26. 26.
    Is there any change in the hybridisation of B and N atoms as a result of the following reaction? BF₃ + NH₃ -> F₃B-NH₃
27. 27.
    Draw diagrams showing the formation of a double bond and a triple bond between carbon atoms in C₂H₄ and C₂H₂ molecules.
28. 28.
    What is the total number of sigma and pi bonds in the following molecules?

    • (a) C₂H₂
    • (b) C₂H₄
29. 29.
    Considering x-axis as the internuclear axis which out of the following will not form a sigma bond and why?

    • (a) 1s and 1s
    • (b) 1s and 2px;
    • (c) 2py and 2py
    • (d) 1s and 2s.
30. 30.
    Which hybrid orbitals are used by carbon atoms in the following molecules?

    • (a) CH₃-CH₃;
    • (b) CH₃-CH=CH₂;
    • (c) CH₃-CH₂-OH;
    • (d) CH₃-CHO
    • (e) CH₃COOH
31. 31.
    What do you understand by bond pairs and lone pairs of electrons? Illustrate by giving one example of each type.
32. 32.
    Distinguish between a sigma and a pi bond.
33. 33.
    Explain the formation of H₂ molecule on the basis of valence bond theory.
34. 34.
    Write the important conditions required for the linear combination of atomic orbitals to form molecular orbitals.
35. 35.
    Use molecular orbital theory to explain why the Be₂ molecule does not exist.
36. 36.
    Compare the relative stability of the following species and indicate their magnetic properties: O₂⁺, O₂, O₂⁻ (superoxide), O₂²⁻ (peroxide)
37. 37.
    Write the significance of a plus and a minus sign shown in representing the orbitals.
38. 38.
    Describe the hybridisation in case of PCl₅. Why are the axial bonds longer as compared to equatorial bonds?
39. 39.
    Define hydrogen bond. Is it weaker or stronger than the van der Waals forces?
40. 40.
    What is meant by the term bond order? Calculate the bond order of: N₂, O₂, O₂⁺ and O₂⁻.