Questions 39-41 refer to the following graph, which shows the heating curve for methane, CH4. [VISUA — Bonding Chemistry Question

STEPS:

1. Understand the relationship between vaporization and intermolecular forces (IMFs): Vaporization is the phase transition from liquid to gas, which requires completely overcoming the intermolecular attractions holding molecules together in the liquid phase. The stronger the intermolecular forces holding the molecules together, the more heat energy is required to vaporize the liquid, resulting in a higher enthalpy of vaporization.
2. Analyze the intermolecular forces in water ($\text{H}_2\text{O}$): Water is a highly polar bent molecule. Because a highly electronegative oxygen atom is covalently bonded directly to hydrogen, water molecules exhibit extremely strong dipole-dipole attractions known as hydrogen bonds.
3. Analyze the intermolecular forces in methane ($\text{CH}_4$): Methane is a symmetric, tetrahedral, nonpolar molecule. The electronegativity difference between carbon and hydrogen is negligible, making the $\text{C}-\text{H}$ bonds essentially nonpolar. Therefore, methane molecules cannot form dipole-dipole attractions or hydrogen bonds. They are held together in the liquid phase solely by weak London dispersion forces.
4. Compare the energy requirements: Because hydrogen bonding is a much stronger intermolecular force than London dispersion forces, it requires significantly more energy to separate water molecules during boiling than to separate methane molecules. This explains why the enthalpy of vaporization of methane ($8.2\text{ kJ/mol}$) is much less than that of water ($40.7\text{ kJ/mol}$), making Option A the correct answer.

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WHY_OTHERS_WRONG:

• Option B is incorrect: While methane does have weak dispersion forces, comparing dispersion forces alone is not the primary reason for the massive difference in their enthalpies of vaporization. The fundamental, dominant factor that distinguishes the physical properties of these two substances is that water possesses strong hydrogen bonds while methane does not.
• Option C is incorrect: Methane has a molar mass of $16\text{ g/mol}$, which is extremely close to water's molar mass of $18\text{ g/mol}$. This tiny difference of $2\text{ g/mol}$ in mass is far too small to account for water's enthalpy of vaporization being nearly five times larger than that of methane.
• Option D is incorrect: Density is a macroscopic physical property (mass per unit volume) that describes how closely packed a substance's molecules are under given conditions. It is a physical consequence of molecular weight and intermolecular forces, not the microscopic chemical cause of the energy required to overcome intermolecular attractions during vaporization.