Question

Which of the following affects the vapor pressure of a liquid? (a) Volume of the liquid, \((\mathbf{b})\) surface area, \((\mathbf{c})\) intermolecular attractive forces, (d) temperature, (e) density of the liquid.

Short answer

The factors that affect the vapor pressure of a liquid are \((\mathbf{c})\) intermolecular attractive forces and (d) temperature.

Step-by-step solution

Option (a): Volume of the liquid

The volume of the liquid does not affect the vapor pressure because vapor pressure depends on the equilibrium between the liquid and its vapor, not the amount of the liquid itself.

Option (b): Surface area

The surface area of the liquid also does not affect the vapor pressure, as it is not related to the equilibrium between the liquid and its vapor. A large liquid surface area can cause faster evaporation, but it does not affect the equilibrium vapor pressure at a given temperature.

Option (c): Intermolecular attractive forces

The intermolecular attractive forces have a direct impact on the vapor pressure of a liquid. Strong intermolecular forces make it difficult for molecules to escape from the liquid phase to the vapor phase. As a result, liquids with strong intermolecular forces have lower vapor pressures.

Option (d): Temperature

The temperature indeed affects the vapor pressure of a liquid. As the temperature increases, the vapor pressure of a liquid also increases. This is because as the temperature increases, more molecules have enough kinetic energy to overcome the intermolecular forces and escape from the liquid phase to the vapor phase.

Option (e): Density of the liquid

The density of a liquid does not have a direct impact on its vapor pressure. Density affects the mass per unit volume of a liquid, but it does not change the equilibrium between the liquid and vapor phases. In conclusion, the factors that affect the vapor pressure of a liquid are \((\mathbf{c})\) intermolecular attractive forces and (d) temperature.

Key concepts

Intermolecular Forces

Intermolecular forces are the attractions between molecules that influence the physical properties of a substance. These forces include hydrogen bonding, dipole-dipole interactions, and London dispersion forces. The strength of these forces affects how easily molecules can transition from the liquid phase to the vapor phase.
A liquid with strong intermolecular forces like hydrogen bonds will have a lower vapor pressure. This is because molecules need more energy to overcome these forces to evaporate. Conversely, substances with weaker intermolecular forces will have higher vapor pressures since it's easier for their molecules to escape into the vapor phase.
For example:

• Water, with significant hydrogen bonding, exhibits a lower vapor pressure than many other liquids.
• Noble gases, which primarily experience London dispersion forces, can have relatively higher vapor pressures.

Understanding the nature of intermolecular forces is crucial in predicting and explaining the vapor pressure of different liquids under various conditions.

Temperature Effects on Vapor Pressure

Temperature plays a significant role in determining the vapor pressure of a liquid. As temperature increases, so does the kinetic energy of the molecules within a liquid. This added energy allows more molecules to overcome any intermolecular forces and escape into the vapor phase, thus increasing the vapor pressure.
To visualize this, imagine a pot of water on a stove. As you heat the water, it begins to boil, increasing the number of molecules transitioning from liquid to gas. This increase in the vapor phase elevates the vapor pressure.
Key points include:

• At higher temperatures, thermal agitation overcomes the attractive intermolecular forces more effectively.
• The relationship between temperature and vapor pressure is exponential, as shown in the Clausius-Clapeyron equation.

This relationship explains why liquids tend to evaporate more readily at higher temperatures and why boiling points are higher at lower atmospheric pressures.

Equilibrium in Liquids

Equilibrium in liquids refers to the dynamic balance between the liquid phase and the vapor phase. When a liquid is in a closed container, molecules continually evaporate into the vapor phase and condense back into the liquid phase.
At equilibrium, the rate of evaporation equals the rate of condensation. This state results in a constant vapor pressure, which is independent of the amount of liquid or surface area. Therefore, changes in volume or surface area do not affect the equilibrium vapor pressure, as it is solely reliant on temperature and intermolecular forces.
To summarize:

• Equilibrium exists when both evaporation and condensation rates balance out.
• External factors like pressure and temperature shifts can disturb this equilibrium, leading to changes in vapor pressure.

Understanding liquid-vapor equilibrium helps in comprehending why certain conditions favor evaporation or condensation, making it critical for tasks involving boiling, distillation, and even climate studies.