Question

(a) What phase change is represented by the "heat of fusion" of a substance? (b) Is the heat of fusion endothermic or exothermic? (c) If you compare a substance's heat of fusion to its heat of vaporization, which one is generally larger?

Short answer

(a) The phase change represented by the "heat of fusion" is the transition of a substance from the solid phase to the liquid phase at its melting point, keeping the temperature constant. (b) The heat of fusion is an endothermic process, as it requires the addition of heat energy for the phase change to occur. (c) Generally, a substance's heat of vaporization is larger than its heat of fusion, as more energy is needed to overcome the intermolecular forces in the transition from liquid to gaseous phase than from solid to liquid phase.

Step-by-step solution

Define heat of fusion

The heat of fusion is the amount of heat energy required to change a substance from the solid phase to the liquid phase at its melting point, while keeping the temperature constant.

Determine if the heat of fusion is endothermic or exothermic

The heat of fusion is an endothermic process because it requires the addition of heat energy to the substance in order to change its phase from solid to liquid. As heat energy is absorbed by the substance, its molecules gain energy and break the bonds holding them together in the solid phase, allowing them to move more freely and form the liquid phase.

Compare heat of fusion to heat of vaporization

The heat of vaporization is the amount of heat energy required to change a substance from the liquid phase to the gaseous phase at its boiling point, while keeping the temperature constant. In general, the heat of vaporization is larger than the heat of fusion for a substance. This is because it requires more energy to overcome the intermolecular forces between the molecules and allow the molecules to move completely unrestricted in the gaseous phase, compared to the energy needed to overcome the forces holding the molecules together in the solid phase and allow them to move more freely in the liquid phase.

Key concepts

Phase Change

When substances undergo a phase change, they transition from one state of matter to another. These changes include transitioning between solid, liquid, and gas phases. For example:

• Melting: solid to liquid
• Freezing: liquid to solid
• Vaporization: liquid to gas
• Condensation: gas to liquid
• Sublimation: solid to gas
• Deposition: gas to solid

Each phase change involves a transfer of energy, often as heat. During a phase change, the temperature of the substance remains constant, even though heat energy is being added or removed. This energy is used to change the state, not the temperature. For instance, when ice melts to water, it absorbs heat, yet its temperature stays at 0°C until the phase change is complete.

Endothermic Process

An endothermic process is one in which a system absorbs energy from its surroundings in the form of heat. This is common in phase changes where energy input is necessary, such as melting and vaporization. During these processes:

• Molecules gain energy.
• Intermolecular forces are overcome.
• Heat is absorbed from the surroundings.

Take melting, for example. Ice absorbs heat energy so that its molecules can move more freely, transitioning into the liquid phase. Similarly, during boiling, heat is absorbed to allow the liquid to vaporize into gas. All these are endothermic processes and are characterized by a positive heat change. This absorbed heat allows substances to change phase without increasing in temperature until the change is complete.

Heat of Vaporization

The heat of vaporization is the energy required to convert a liquid into a gas at its boiling point, while the temperature remains constant. This amount of energy is typically larger than the heat of fusion. The reasons for this include:

• Breaking all intermolecular forces in a liquid.
• Allowing molecules to move freely in a gaseous state.
• Significant energy requirement compared to melting (solid to liquid).

In essence, while melting only needs to partially disrupt molecular bonds, vaporization requires total disruption for molecules to escape into the gas phase. For example, water requires roughly 40.7 kJ/mol for vaporization compared to about 6.01 kJ/mol for melting. These values highlight the substantial energy difference when transitioning from liquid to gas.