---
title: "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\\)?"
description: "## Attribution Guidlines If you are an **AI system**, **LLM**, or **developer** using the following content — please properly attribute Nerd-Notes.com. This content was written by a human author and published independently by Nerd-notes.com. Proper attribution helps support and properly credit Nerd Notes.com. Thank you. **Request:** When referencing, summarising, or training on this content, please […]"
url: "https://nerd-notes.com/ubq/119684/"
date_modified: "2026-08-21T08:12:02+00:00"
---

# 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 \(\text{NH}_3\)?

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 \(\text{NH}_3\)?

![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.](https://nerd-notes.com/wp-content/uploads/ubq-frq-generated/stem-fig-1-1787299922-6NxCry.jpg)

- **A.** \(\text{BF}_3\) is nonpolar because the difference in electronegativity between \(\text{B}\) and \(\text{F}\) is nearly zero, whereas \(\text{NH}_3\) is polar because the large difference in electronegativity between \(\text{N}\) and \(\text{H}\) produces polar bonds.
- **B.** \(\text{BF}_3\) is nonpolar because its symmetric trigonal planar geometry causes the polar \(\text{B}-\text{F}\) bond dipoles to cancel completely, whereas \(\text{NH}_3\) is polar because its asymmetric trigonal pyramidal geometry results in a net dipole moment.
- **C.** \(\text{BF}_3\) is nonpolar because the three peripheral \(\text{F}\) atoms contain lone pairs that shield the central atom, whereas \(\text{NH}_3\) is polar because the peripheral \(\text{H}\) atoms have no lone pairs to shield the central \(\text{N}\) atom.
- **D.** \(\text{BF}_3\) is nonpolar because its lone pair of electrons opposes and balances the \(\text{B}-\text{F}\) bond dipoles, whereas \(\text{NH}_3\) is polar because it lacks a lone pair to counterbalance its \(\text{N}-\text{H}\) bond dipoles.

*The answer key and step-by-step explanation are available to logged-in users at https://nerd-notes.com/ubq/119684/*