A student predicts that the bond polarity order for three single bonds is N–F > O–Cl > S–Cl, based solely on electronegativity differences. Infrared spectroscopy, however, shows that the O–Cl bond in a series of oxochloride compounds produces a more intense absorption band (proportional to bond polarity) than the N–F bond in analogous fluoride compounds.
The following additional information is available:
• Average bond lengths: N–F = 1.36 Å, O–Cl = 1.70 Å, S–Cl = 2.05 Å.
• Pauling electronegativities: N = 3.0, O = 3.5, S = 2.5, F = 4.0, Cl = 3.1.
(a) Based on the given electronegativity values, calculate ΔEN for each bond and provide the theoretical polarity order predicted by ΔEN alone.
(b) Propose a qualitative explanation, incorporating bond length considerations, for why the experimental polarity of the O–Cl bond is greater than that of the N–F bond despite having a smaller ΔEN.
(c) Evaluate the student’s original claim in light of the experimental evidence and your explanation in part (b). State whether the claim is fully supported, partially supported, or refuted, and justify your conclusion.
(d) Design an additional experimental or computational approach that would help clarify whether bond length or electronegativity difference plays the dominant role in determining bond polarity across this set of compounds. Outline your proposed method and the type of data you would collect.