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Chemistry · Textbook solutions

Chemical Bonding and Molecular Structure

Every solved example, exercise, and miscellaneous question — in the order the textbook teaches them. · 44 questions

Worked Examples

4 q

Solved Examples

Worked · 4
  1. Eg 4.1
    Write the Lewis dot structure of CO molecule.
  2. Eg 4.2
    Write the Lewis structure of the nitrite ion, NO2−\text{NO}_2^-.
  3. Eg 4.3
    Explain the structure of CO32−\text{CO}_3^{2-} ion in terms of resonance.
  4. Eg 4.4
    Explain the structure of CO2\text{CO}_2 molecule.

Exercises

40 q
  1. Ex 4.1
    Explain the formation of a chemical bond.
  2. Ex 4.2
    Write Lewis dot symbols for atoms of the following elements: Mg, Na, B, O, N, Br.
  3. Ex 4.3
    Write Lewis symbols for the following atoms and ions: S and S2−\text{S}^{2-}; Al and Al3+\text{Al}^{3+}; H and H−\text{H}^-
  4. Ex 4.4
    Draw the Lewis structures for the following molecules and ions: H2S\text{H}_2\text{S}, SiCl4\text{SiCl}_4, BeF2\text{BeF}_2, CO32−\text{CO}_3^{2-}, HCOOH
  5. Ex 4.5
    Define octet rule. Write its significance and limitations.
  6. Ex 4.6
    Write the favourable factors for the formation of ionic bond.
  7. Ex 4.7
    Discuss the shape of the following molecules using the VSEPR model: BeCl2\text{BeCl}_2, BCl3\text{BCl}_3, SiCl4\text{SiCl}_4, AsF5\text{AsF}_5, H2S\text{H}_2\text{S}, PH3\text{PH}_3
  8. Ex 4.8
    Although geometries of NH3\text{NH}_3 and H2O\text{H}_2\text{O} molecules are distorted tetrahedral, bond angle in water is less than that of ammonia. Discuss.
  9. Ex 4.9
    How do you express the bond strength in terms of bond order?
  10. Ex 4.10
    Define the bond length.
  11. Ex 4.11
    Explain the important aspects of resonance with reference to the CO32−\text{CO}_3^{2-} ion.
  12. Ex 4.12
    H3PO3\text{H}_3\text{PO}_3 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\text{H}_3\text{PO}_3? If not, give reasons for the same.
  13. Ex 4.13
    Write the resonance structures for SO3\text{SO}_3, NO2\text{NO}_2 and NO3−\text{NO}_3^-.
  14. Ex 4.14
    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. Ex 4.15
    Although both CO2\text{CO}_2 and H2O\text{H}_2\text{O} are triatomic molecules, the shape of H2O\text{H}_2\text{O} molecule is bent while that of CO2\text{CO}_2 is linear. Explain this on the basis of dipole moment.
  16. Ex 4.16
    Write the significance/applications of dipole moment.
  17. Ex 4.17
    Define electronegativity. How does it differ from electron gain enthalpy?
  18. Ex 4.18
    Explain with the help of suitable example polar covalent bond.
  19. Ex 4.19
    Arrange the bonds in order of increasing ionic character in the molecules: LiF, K2O\text{K}_2\text{O}, N2\text{N}_2, SO2\text{SO}_2 and ClF3\text{ClF}_3.
  20. Ex 4.20
    The skeletal structure of CH3COOH\text{CH}_3\text{COOH} as shown below is correct, but some of the bonds are shown incorrectly. Write the correct Lewis structure for acetic acid.
  21. Ex 4.21
    Apart from tetrahedral geometry, another possible geometry for CH4\text{CH}_4 is square planar with the four H atoms at the corners of the square and the C atom at its centre. Explain why CH4\text{CH}_4 is not square planar?
  22. Ex 4.22
    Explain why BeH2\text{BeH}_2 molecule has a zero dipole moment although the Be–H bonds are polar.
  23. Ex 4.23
    Which out of NH3\text{NH}_3 and NF3\text{NF}_3 has higher dipole moment and why?
  24. Ex 4.24
    What is meant by hybridisation of atomic orbitals? Describe the shapes of spsp, sp2sp^2, sp3sp^3 hybrid orbitals.
  25. Ex 4.25
    Describe the change in hybridisation (if any) of the Al atom in the following reaction. AlCl3+Cl−→AlCl4−\text{AlCl}_3 + \text{Cl}^- \rightarrow \text{AlCl}_4^-
  26. Ex 4.26
    Is there any change in the hybridisation of B and N atoms as a result of the following reaction? BF3+NH3→F3B ⁣⋅ ⁣NH3\text{BF}_3 + \text{NH}_3 \rightarrow \text{F}_3\text{B}\!\cdot\!\text{NH}_3
  27. Ex 4.27
    Draw diagrams showing the formation of a double bond and a triple bond between carbon atoms in C2H4\text{C}_2\text{H}_4 and C2H2\text{C}_2\text{H}_2 molecules.
  28. Ex 4.28
    What is the total number of sigma and pi bonds in the following molecules? (a) C2H2\text{C}_2\text{H}_2 (b) C2H4\text{C}_2\text{H}_4
  29. Ex 4.29
    Considering x-axis as the internuclear axis which out of the following will not form a sigma bond and why?
    1. A.
      1s1s and 1s1s
    2. B.
      1s1s and 2px2p_x
    3. C.
      2py2p_y and 2py2p_y
    4. D.
      1s1s and 2s2s
  30. Ex 4.30
    Which hybrid orbitals are used by carbon atoms in the following molecules? CH3–CH3\text{CH}_3\text{--CH}_3; (b) CH3–CH=CH2\text{CH}_3\text{--CH}=\text{CH}_2; (c) CH3-CH2-OH\text{CH}_3\text{-CH}_2\text{-OH}; (d) CH3-CHO\text{CH}_3\text{-CHO} (e) CH3COOH\text{CH}_3\text{COOH}
  31. Ex 4.31
    What do you understand by bond pairs and lone pairs of electrons? Illustrate by giving one exmaple of each type.
  32. Ex 4.32
    Distinguish between a sigma and a pi bond.
  33. Ex 4.33
    Explain the formation of H2\text{H}_2 molecule on the basis of valence bond theory.
  34. Ex 4.34
    Write the important conditions required for the linear combination of atomic orbitals to form molecular orbitals.
  35. Ex 4.35
    Use molecular orbital theory to explain why the Be2\text{Be}_2 molecule does not exist.
  36. Ex 4.36
    Compare the relative stability of the following species and indicate their magnetic properties; O2\text{O}_2, O2+\text{O}_2^+, O2−\text{O}_2^- (superoxide), O22−\text{O}_2^{2-} (peroxide)
  37. Ex 4.37
    Write the significance of a plus and a minus sign shown in representing the orbitals.
  38. Ex 4.38
    Describe the hybridisation in case of PCl5\text{PCl}_5. Why are the axial bonds longer as compared to equatorial bonds?
  39. Ex 4.39
    Define hydrogen bond. Is it weaker or stronger than the van der Waals forces?
  40. Ex 4.40
    What is meant by the term bond order? Calculate the bond order of: N2\text{N}_2, O2\text{O}_2, O2+\text{O}_2^+ and O2−\text{O}_2^-.