Questions: States of Matter and Phase Changes: Melting, Boiling, and Sublimation

Question 1

You are heating ice at a constant rate. When the temperature reaches 0°C, it stays constant for several minutes even though you continue adding heat. What explains this?

• A. The heat source loses efficiency as it warms the surrounding environment, reducing effective heat transfer
• B. Ice requires extra energy to increase its temperature once partial melting has begun
• C. The added energy is consumed breaking intermolecular attractions in the crystal lattice rather than increasing molecular kinetic energy
• D. The thermometer reads incorrectly near the freezing point due to the latent heat effect

Answer: (C)

During a phase change, added energy goes entirely into breaking intermolecular bonds (overcoming the crystal lattice in this case), not into increasing the speed of molecules. Since temperature measures average kinetic energy, and kinetic energy isn't increasing, temperature stays constant. This energy is the heat of fusion (ΔH_fus). Only once all the solid has melted — all lattice bonds broken — does further heating increase molecular speed and temperature again.

Question 2

Ethanol (which forms hydrogen bonds) boils at 78°C, while propane (similar molar mass, only London dispersion forces) boils at −42°C. What is the correct explanation for this 120°C difference in boiling point?

• A. Propane molecules move faster at any given temperature because they are slightly lighter
• B. Ethanol's stronger intermolecular forces require more kinetic energy to overcome before molecules can escape the liquid phase
• C. Propane's smaller molecular volume allows molecules to escape the liquid surface more easily
• D. Hydrogen bonds in ethanol prevent sublimation, forcing the substance to pass through the liquid phase and raising its boiling point

Answer: (B)

Boiling occurs when molecules gain enough kinetic energy to escape the collective pull of intermolecular forces. Ethanol's hydrogen bonds are much stronger than propane's London dispersion forces, requiring a higher temperature (more kinetic energy) before enough molecules can escape to maintain a vapor phase at atmospheric pressure. This relationship is the predictive core of this topic: stronger intermolecular forces → higher boiling point. Molecular weight differences are a factor but secondary to force strength here.

Question 3

The heat of vaporization for a substance is always larger than its heat of fusion because converting liquid to gas requires completely overcoming intermolecular forces, while melting only partially disrupts them.

• T. True
• F. False

Answer: True

In the solid-to-liquid transition, molecules break free of fixed lattice positions but remain within the collective pull of nearby molecules — they can still slide past each other but don't escape entirely. In the liquid-to-gas transition, molecules must overcome all remaining intermolecular attractions to fly freely. Because complete escape requires breaking more interactions than partial disordering, ΔH_vap is always substantially larger than ΔH_fus. For water, ΔH_vap ≈ 40.7 kJ/mol vs. ΔH_fus ≈ 6.0 kJ/mol — nearly a 7-fold difference.

Question 4

During sublimation, a solid converts directly to gas without passing through the liquid phase, which means it bypasses the energy cost of overcoming intermolecular forces.

• T. True
• F. False

Answer: False

Sublimation requires overcoming intermolecular forces just as thoroughly as boiling — in fact, solid molecules must escape from a more tightly ordered lattice directly into the gas phase. The total enthalpy of sublimation approximately equals ΔH_fus + ΔH_vap because all intermolecular interactions must be broken. Sublimation occurs not because less energy is needed, but because surface molecules gain sufficient energy to escape without the intermediate liquid state being thermodynamically stable under those conditions of temperature and pressure.

Question 5

Why does temperature remain constant during a phase change even when heat is continuously being added to the system?

Model answer: Temperature measures the average kinetic energy of molecules. During a phase change, added energy is used entirely to break intermolecular bonds (e.g., overcoming lattice attractions during melting, or breaking all remaining intermolecular attractions during boiling) rather than increasing molecular speed. Since kinetic energy isn't increasing, temperature doesn't increase. Only once the phase transition is complete — all the bonds that needed to be broken have been broken — does further energy input increase molecular kinetic energy and raise the temperature again.

This is why phase transitions appear as flat plateaus on heating curves. The energy going in during the plateau is the latent heat (heat of fusion or vaporization). It's 'hidden' in the sense that it doesn't register as a temperature change. This also explains why steam burns are more severe than boiling water burns at the same temperature: steam releases both sensible heat (cooling to 100°C) and the full heat of condensation as it returns to liquid.