Question

In which of the following pairs, both molecules possess dipole moment?

• A. \( CO_2, BCl_3 \)
• B. \( BCl_3, NF_3 \)
• C. \( CO_2, SO_2 \)
• D. \( SO_2, NF_3 \)

Hint:

Molecules with symmetric geometry (linear, trigonal planar) usually have zero dipole moment.

Answer 1

The Correct Option is D

Step 1: Identify Molecular Geometry and Dipole Moments 
- \( CO_2 \) and \( BCl_3 \) are symmetric → No dipole moment.
- \( SO_2 \) (bent) and \( NF_3 \) (pyramidal) → Have dipole moment.

Step 2: Correct Pair \[ SO_2, NF_3 \]

Answer 2

To determine whether molecules possess a dipole moment, we need to consider their molecular geometry and the electronegativity difference between atoms.

1. \( SO_2 \) (Sulfur dioxide):
- Molecular shape: Bent (angular)
- The S–O bonds are polar due to electronegativity difference.
- The bent shape means the bond dipoles do not cancel out.
- Hence, \( SO_2 \) has a net dipole moment.

2. \( NF_3 \) (Nitrogen trifluoride):
- Molecular shape: Trigonal pyramidal (like ammonia)
- N–F bonds are polar.
- The lone pair on nitrogen causes the bond dipoles to not cancel completely.
- Therefore, \( NF_3 \) has a net dipole moment, though smaller than ammonia.

Therefore, in the pair \( SO_2 \) and \( NF_3 \), both molecules possess dipole moments.

Hence, the correct answer is:
\( SO_2, NF_3 \)