Question

which of the following correctly shows the bond polarity? chapter 4 compounds and their bonds - ppt video online download a δ+s−siδ− b δ+o−pδ− c δ+br−fδ− d δ+cl−clδ−

Explanation:

• Option A: In the \(S - Si\) bond, silicon (\(Si\)) has a lower electronegativity (\(1.90\)) than sulfur (\(S\), \(2.58\)). So, the bond polarity should be \(Si^{\delta +}-S^{\delta -}\), not \(S^{\delta +}-Si^{\delta -}\).
• Option B: Oxygen (\(O\), \(3.44\)) has a higher electronegativity than phosphorus (\(P\), \(2.19\)). The bond polarity should be \(P^{\delta +}-O^{\delta -}\), not \(O^{\delta +}-P^{\delta -}\).
• Option C: Fluorine (\(F\), \(3.98\)) has a higher electronegativity than bromine (\(Br\), \(2.96\)). So, the bond polarity \(Br^{\delta +}-F^{\delta -}\) is correct.
• Option D: In the \(Cl - Cl\) bond, since both atoms are the same, the electronegativity difference is zero. So, the bond is non - polar, and there should be no \(\delta+\) and \(\delta-\) notation.

Answer:

C. \(^{\delta +}\text{Br}-\text{F}^{\delta -}\)