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AP Chemistry
2.7 VSEPR and Hybridization
2.5 Lewis Diagrams
IntermediateMCQConceptual25.4k
Two molecular Lewis structures are shown side-by-side in black line art on a white background. On the left, labeled \(\text{BF}_3\), a central \(\text{B}\) atom is bonded by three single solid lines to three surrounding \(\text{F}\) atoms in a trigonal planar arrangement at \(120^\circ\) angles (one bond pointing straight up, two pointing downward and outward). Each \(\text{F}\) atom has exactly three pairs of electron dots (six dots total per \(\text{F}\)) on its non-bonding sides. The central \(\text{B}\) atom has zero lone pairs and no other dots. On the right, labeled \(\text{NH}_3\), a central \(\text{N}\) atom has one pair of electron dots directly above it and is bonded by three single solid lines to three surrounding \(\text{H}\) atoms angled downward in a trigonal pyramidal shape. Each \(\text{H}\) atom has zero lone pairs. No other particles, labels, charges, partial charge symbols, vectors, or annotations appear.
Lewis diagrams for \(\text{BF}_3\) and \(\text{NH}_3\)
The Lewis diagrams for \(\text{BF}_3\) and \(\text{NH}_3\) are shown below. Both molecules contain three single covalent bonds to peripheral atoms, but \(\text{NH}_3\) is polar with a nonzero dipole moment, whereas \(\text{BF}_3\) is nonpolar with a dipole moment of zero.

Which of the following statements best explains the difference in polarity between \(\text{BF}_3\) and \(\text{NH}_3\)?

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