American Chemical Society (ACS) Chemistry Test 2025 – 400 Free Practice Questions to Pass the Exam

The unusual properties of water, such as its high specific heat, high heat of vaporization, and its ability to expand upon freezing, can be primarily attributed to hydrogen bonding. Water molecules have a bent shape, where the oxygen atom is more electronegative than the hydrogen atoms, creating a polar molecule with distinct positive and negative ends. Due to this polarity, each water molecule can form hydrogen bonds with nearby water molecules, where the hydrogen atom of one molecule is attracted to the electronegative oxygen atom of another.

These hydrogen bonds are relatively strong compared to other types of intermolecular forces, which result in significant energy requirements for processes like melting and boiling. This accounts for water's high boiling point relative to other similar-sized molecules. Moreover, the hydrogen bonding network contributes to the lower density of ice compared to liquid water, explaining why ice floats.

In contrast, the other options represent weaker forces or different interactions. Van der Waals forces include London dispersion forces, which play a role in nonpolar substances and are generally much weaker than hydrogen bonds. Dipole-dipole interactions occur between polar molecules but lack the unique strength that hydrogen bonds exhibit. Ionic interactions involve the attraction between charged ions and are not applicable to molecular compounds like water.