Understanding Polar and Nonpolar Molecules

The question asks us to identify polar molecules from a given list. Molecules can be classified as polar or nonpolar based on the distribution of electron density within the molecule. This distribution is influenced by the polarity of the individual bonds and the overall shape of the molecule.

A polar molecule is one where there is an uneven distribution of electron density, creating a partial positive charge ($\delta+$) on one side and a partial negative charge ($\delta-$) on the other. This uneven distribution results in a net dipole moment. Polar molecules are typically formed when there are polar bonds (bonds between atoms with different electronegativities) and the molecular geometry is such that the bond dipoles do not cancel each other out.

A nonpolar molecule is one where the electron density is evenly distributed. This can happen if all bonds are nonpolar (between atoms of the same electronegativity) or if the molecule's symmetrical shape causes the individual bond dipoles to cancel each other out, resulting in a zero net dipole moment.

Analyzing the Polarity of Each Molecule

Let's examine each molecule provided:

• A. Solid SO$_2$ (Sulfur Dioxide): The SO$_2$ molecule has a central sulfur atom bonded to two oxygen atoms. Sulfur and oxygen have different electronegativities, making the S=O bonds polar. The molecule has a bent shape (due to a lone pair on sulfur). Because of this bent shape, the bond dipoles do not cancel out, resulting in a net dipole moment. Therefore, SO$_2$ is a polar molecule. Its polarity is the reason solid SO$_2$ can be held together by dipole–dipole interactions.
• B. NH$_3$ (Ammonia): The NH$_3$ molecule has a central nitrogen atom bonded to three hydrogen atoms, with one lone pair on the nitrogen. Nitrogen is more electronegative than hydrogen, making the N-H bonds polar. The molecule has a trigonal pyramidal shape. The bond dipoles and the dipole moment from the lone pair are directed towards the nitrogen, and they do not cancel out due to the pyramidal shape. Therefore, NH$_3$ is a polar molecule.
• C. Cl$_2$ (Chlorine): The Cl$_2$ molecule is diatomic, consisting of two identical chlorine atoms. The bond between two identical atoms is nonpolar because the electronegativity difference is zero. The electron density is evenly distributed. Therefore, Cl$_2$ is a nonpolar molecule.
• D. HCl (Hydrogen Chloride): The HCl molecule is diatomic, consisting of a hydrogen atom and a chlorine atom. Chlorine is significantly more electronegative than hydrogen, making the H-Cl bond polar. There is no symmetry to cancel this dipole moment in a diatomic molecule. Therefore, HCl is a polar molecule.
• E. He (Helium): Helium is a monatomic element (a single atom). It has a spherical electron cloud distribution and no bonds. It is inherently nonpolar.

Identifying the Polar Molecules

Based on our analysis, the polar molecules from the list are:

• SO$_2$ (A)
• NH$_3$ (B)
• HCl (D)

These molecules possess a net dipole moment due to uneven electron distribution caused by polar bonds and/or molecular geometry.

Conclusion

The polar molecules are SO$_2$, NH$_3$, and HCl, which correspond to options A, B, and D. This aligns with the understanding that polar molecules can experience dipole–dipole interactions in the solid state.

Revision Table: Molecular Polarity Summary

Additional Information: Intermolecular Forces and Polarity

The polarity of molecules plays a crucial role in determining the types of intermolecular forces present between them. Intermolecular forces are attractive forces between molecules.

• Dipole–Dipole Interactions: These occur between polar molecules. The partial positive end of one molecule is attracted to the partial negative end of another molecule. These interactions are stronger than London dispersion forces but weaker than covalent or ionic bonds. The question mentions that molecules in solids are held together by dipole–dipole interactions, which implies the solid being discussed is composed of polar molecules.
• London Dispersion Forces (LDF): These forces are present in all substances (polar and nonpolar) and arise from temporary fluctuations in electron distribution, creating instantaneous dipoles. These are the weakest type of intermolecular force but are significant in nonpolar molecules and noble gases.
• Hydrogen Bonding: A special type of dipole–dipole interaction that occurs when hydrogen is bonded to a highly electronegative atom (like N, O, or F). These are particularly strong intermolecular forces. NH$_3$ is an example of a molecule that can form hydrogen bonds.

Nonpolar molecules like Cl$_2$ and He are primarily held together by London dispersion forces in their condensed phases (liquid and solid).