MCQ Detailed View

London dispersion forces are a type of van der Waals force present in all molecules. They arise due to temporary fluctuations in the electron cloud, creating instantaneous dipoles that attract neighboring molecules. The strength of these forces depends on molecular size, number of electrons, and polarizability.

Water (H₂O) is a polar molecule with a bent structure. Each water molecule has two hydrogen atoms bonded to oxygen, creating a partial positive and negative charge. While hydrogen bonding is the dominant intermolecular force in water, it also exhibits London dispersion forces. The presence of both hydrogen bonding and dispersion forces makes water’s overall intermolecular interactions stronger than in simple nonpolar molecules like hydrogen gas (H₂) or bromine (Br₂).

Hydrogen gas has very few electrons and is nonpolar, so its dispersion forces are extremely weak. Bromine is larger and nonpolar, so it has stronger dispersion forces than H₂, but still weaker than water because water’s hydrogen bonds amplify intermolecular attractions. Hydrochloric acid (HCl) is polar and experiences dipole-dipole interactions, but its dispersion forces are weaker than in water.

Understanding London dispersion forces alongside other intermolecular forces helps explain physical properties like boiling point, viscosity, and solubility. Water’s high boiling point relative to its molecular weight demonstrates the cumulative effect of strong intermolecular attractions, including both hydrogen bonding and London dispersion forces.

Among the given molecules, the strongest London dispersion forces are observed in water, due to its combination of polar nature and extensive hydrogen bonding, which enhances intermolecular attraction compared to the other options.