CHEM: 1210 Liquids and solids

Q: You are given ethanol and dimethyl ether, both with the same molecular formula (C₂H₆O). If you were asked to choose one as a fuel for a system where minimal evaporation loss is desired, which would you pick and why?

A: Ethanol would be the better choice. It has stronger hydrogen bonding, resulting in lower vapor pressure compared to dimethyl ether. This means it evaporates less easily, minimizing evaporation losses.

Q: You spill some water and mercury on a clean glass surface. Which liquid will spread out more and why?

A: Water will spread out more because it has stronger adhesive forces with glass. Mercury has stronger cohesive forces, causing it to bead up.

Q: If a substance has a steep slope in a Clausius-Clapeyron plot (ln P vs. 1/T), what can you infer about its heat of vaporization (ΔHᵥₐₚ)?

A: A steep negative slope indicates a high ΔHᵥₐₚ, meaning the substance requires more energy to vaporize.

Q: You need a substance that boils at room temperature under low pressure for use in a vacuum system. What property should this substance have and why?

A: It should have a high vapor pressure at room temperature, which implies weaker intermolecular forces and easy evaporation at low pressure.

Q: If a liquid has a concave meniscus in a glass capillary, what does that tell you about its adhesive and cohesive forces?

A: Adhesive forces (to the glass) are stronger than cohesive forces within the liquid, causing the liquid to climb up the sides.

Q: Why would a substance with high surface tension also tend to have a high heat of vaporization?

A: Both properties are related to strong intermolecular forces. High surface tension indicates molecules resist separation, which also means more energy is needed to vaporize the substance.

Q: Why does water's solid-liquid boundary in the phase diagram slope negatively, while most substances have a positive slope?

A: Because ice is less dense than liquid water. Increasing pressure favors the denser phase—liquid—causing the melting point to decrease under pressure.

Q: You notice that alcohol feels cool when it evaporates from your skin. What does this tell you about its heat of vaporization?

A: It requires energy to evaporate—heat is absorbed from your skin, indicating a relatively high heat of vaporization.

Q: Why is sweating an effective method of cooling the human body?

A: Because water’s high heat of vaporization means it absorbs a lot of body heat as it evaporates, effectively cooling the skin.

Q: If a liquid has a very low boiling point at standard pressure, what can you say about its intermolecular forces?

A: They are weak—less energy is needed to separate the molecules into the gas phase.

Q: You observe a substance that instantly sublimates at room temperature. What does this suggest about its vapor pressure and intermolecular forces?

A: It has a high vapor pressure and very weak intermolecular forces.

Q: A phase diagram shows a substance with a triple point at low pressure and temperature. What does this suggest about the stability of its solid and gas phases?

A: Both the solid and gas phases are more stable at low pressures—indicating the substance easily sublimates.

Q: Why is ice less dense than water, even though most solids are denser than their liquid forms?

A: Hydrogen bonding in ice forms an open hexagonal structure with more space between molecules than in liquid water.

Q: If a liquid climbs up a narrow glass tube, what does that tell you about the balance of forces?

A: Adhesive forces to the glass are stronger than cohesive forces within the liquid, enabling capillary action.