Questions 14-17 refer to the following information. Substance | Lewis Diagram | Boiling Point CH3OH — Bonding Chemistry Question

STEPS:

1. Understand the relationship between vapor pressure and intermolecular forces (IMFs):
Equilibrium vapor pressure is the pressure exerted by a gas in dynamic equilibrium with its liquid phase in a closed system. At a given temperature, a substance with stronger intermolecular forces holds its molecules more tightly in the condensed liquid phase, making it harder for them to escape into the gas phase. Consequently, stronger overall intermolecular forces result in a lower equilibrium vapor pressure.

2. Analyze the given boiling point data:
* Methanol ($\text{CH}_3\text{OH}$): Boiling point = $338\text{ K}$
* Ethanol ($\text{C}_2\text{H}_5\text{OH}$): Boiling point = $351\text{ K}$
* Because boiling point is a direct macroscopic indicator of the energy required to overcome intermolecular attractions, ethanol's higher boiling point confirms that it experiences stronger overall intermolecular forces than methanol. Thus, the equilibrium vapor pressure of ethanol at $300\text{ K}$ must be lower than that of methanol.

3. Compare the types of intermolecular forces present:
* Both methanol and ethanol possess a polar hydroxyl ($-\text{OH}$) group. This allows both substances to form strong intermolecular hydrogen bonds in addition to exhibiting dipole-dipole attractions and London dispersion forces (LDFs).
* Since both molecules form a highly comparable network of hydrogen bonds due to having exactly one hydroxyl group, we must look to the difference in their hydrocarbon tails to explain the difference in their IMF strengths.

4. Evaluate the difference in London dispersion forces:
* Ethanol ($\text{C}_2\text{H}_5\text{OH}$) possesses a larger ethyl group (two carbon atoms) compared to methanol's ($\text{CH}_3\text{OH}$) smaller methyl group (one carbon atom).
* Because of this larger size and additional electrons, ethanol has a larger, more polarizable electron cloud than methanol.
* A more polarizable electron cloud results in significantly stronger London dispersion forces.
* Therefore, the cumulative intermolecular attractions in ethanol are stronger because of its greater London dispersion forces, which results in a lower vapor pressure. This directly corresponds to Option C.

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WHY_OTHERS_WRONG:

• Option A is incorrect: Although both molecules can form hydrogen bonds, they do not have the same vapor pressure. This explanation completely ignores the influence of the hydrocarbon chains, which differ in size and yield different magnitudes of London dispersion forces.
• Option B is incorrect: This option correctly states that the London dispersion forces are greater in ethanol, but it reaches the wrong physical conclusion. Stronger intermolecular forces lead to a *lower* vapor pressure, not a higher one, because fewer molecules possess enough kinetic energy to escape into the vapor phase.
• Option D is incorrect: While this option correctly concludes that ethanol's vapor pressure is lower, the explanation is chemically incorrect. Both methanol and ethanol have exactly one hydroxyl ($-\text{OH}$) group per molecule. Therefore, they have the same capacity for hydrogen bonding; the difference in their physical properties is driven by the strength of their London dispersion forces.