Polarity of Bonds and Molecules - UNSOLVED PRACTICE SET

Q7. Define electronegativity. How does electronegativity difference determine the polarity of a bond?

Q8. Why is CO₂ a non-polar molecule despite having polar C=O bonds?

Q9. Define dipole moment. What is its SI unit?

Q10. Arrange the hydrogen halides (HF, HCl, HBr, HI) in order of increasing polarity. Give a reason. 

Q11. Your mother uses a microwave oven to heat food. She explains that water molecules in food rotate because they are polar, and this rotation generates heat. Using the concept of molecular polarity, explain why water is polar while carbon dioxide is not, even though both contain polar bonds.

Q12. What is percentage ionic character? How is it calculated from dipole moment data? 

Q13. Explain the concept of bond polarity and molecular polarity. Discuss the factors that determine whether a molecule is polar or non-polar. Predict the polarity of:

(a) NH₃

(b) BF₃

(c) CHCl₃

(d) C₂H₄

Illustrate with diagrams showing dipole moments.

Q14. Describe the relationship between electronegativity, dipole moment, and percentage ionic character. Explain why:

(a) HF has a higher dipole moment than HCl despite having a shorter bond length

(b) CO₂ is non-polar while SO₂ is polar

(c) CCl₄ is non-polar while CHCl₃ is polar

(d) The percentage ionic character of HCl is about 17% even though the electronegativity difference suggests it should be 100% ionic

Q15. During a chemistry lab session, you are given four liquids: water, carbon tetrachloride, ethanol, and hexane.

(a) You mix water and ethanol — they dissolve in each other completely. You mix water and CCl₄ — they form two layers. Explain using polarity concepts.

(b) Your teacher asks you to predict whether NH₃ or NF₃ has a larger dipole moment. Both have pyramidal shapes. Explain your reasoning considering both bond polarity and lone pair effects.

(c) The dipole moment of CO is 0.12 D, which is unusual for a molecule with a small electronegativity difference. Explain this using the concept of back-bonding or lone pair donation.

(d) Why does the presence of a lone pair on the central atom generally increase the dipole moment of a molecule?

Q16. The dipole moment of HCl is 1.03 D and the bond length is 127 pm.

(a) Calculate the percentage ionic character in HCl. (1 D = 3.336 × 10⁻³⁰ C·m, e = 1.6 × 10⁻¹⁹ C)

(b) If HCl were 100% ionic, what would be its dipole moment?

(c) Why is the actual dipole moment less than the calculated ionic dipole moment?

Q17. The dipole moments of the following molecules are given:

HF: 1.82 D

HCl: 1.03 D

HBr: 0.79 D

HI: 0.38 D

(a) Arrange these in order of decreasing bond polarity.

(b) The bond lengths are: HF = 92 pm, HCl = 127 pm, HBr = 141 pm, HI = 161 pm. Calculate the charge separation (δ) for each molecule using μ = δ × d.

(c) Why does the charge separation decrease down the group despite the increasing bond length?

(d) Predict which hydrogen halide would be most soluble in water and explain why.

Q18. Consider the following molecules and their shapes:

(a) BF₃ (trigonal planar, μ = 0 D) vs. NF₃ (trigonal pyramidal, μ = 0.23 D). Both have polar bonds. Explain why BF₃ is non-polar while NF₃ is polar.

(b) CO₂ (linear, μ = 0 D) vs. SO₂ (bent, μ = 1.62 D). Explain the difference in polarity.

(c) CH₄ (tetrahedral, μ = 0 D) vs. CHCl₃ (tetrahedral, μ = 1.15 D). Explain why replacing one H with Cl creates a polar molecule.

(d) Calculate the resultant dipole moment for a hypothetical molecule with two bond dipoles of 1.5 D each, oriented at 109.5° to each other (like in a tetrahedral molecule with two polar bonds).