Question

Polar molecule among the following is:

• A. $\mathrm{BF_3}$
• B. $\mathrm{XeF_4}$
• C. $\mathrm{CCl_4}$
• D. $\mathrm{NH_3}$

Hint:

Key points to remember:

• Lone pairs often create polarity
• Symmetrical shapes cancel bond dipoles
• $\mathrm{NH_3}$'s pyramidal shape makes it polar
• All given molecules except $\mathrm{NH_3}$ are symmetrical

Answer 1

The Correct Option is D

Step 1: Understand molecular polarity
A polar molecule must have:

• Polar bonds (electronegativity difference)
• Asymmetric geometry (net dipole moment)

Step 2: Analyze each option
\begin{tabular}{lll} Molecule & Geometry & Polarity
$\mathrm{BF_3}$ & Trigonal planar & Nonpolar (symmetrical)
$\mathrm{XeF_4}$ & Square planar & Nonpolar (symmetrical)
$\mathrm{CCl_4}$ & Tetrahedral & Nonpolar (symmetrical)
$\mathrm{NH_3}$ & Trigonal pyramidal & Polar (asymmetric)
\end{tabular}
Step 3: Verify $NH_3$ polarity

• N-H bonds are polar (EN difference)
• Lone pair creates asymmetric charge distribution
• Net dipole moment exists

Step 4: Why others are nonpolar

• $\mathrm{BF_3}$: Symmetric despite polar B-F bonds
• $\mathrm{XeF_4}$: Symmetric square planar cancels dipoles
• $\mathrm{CCl_4}$: Symmetric tetrahedral cancels dipoles