Question

What do we call the energies required, respectively, to melt and to vaporize 1 mol of a substance? Which of these energies is always larger for a given sub. stance? Why?

Short answer

The energies required to melt and vaporize 1 mol of a substance are called enthalpy of fusion and enthalpy of vaporization, respectively. For a given substance, the enthalpy of vaporization is always larger than the enthalpy of fusion because vaporization requires breaking all the remaining intermolecular forces between particles, enabling them to move independently in the gaseous state, which demands more energy than the melting process.

Step-by-step solution

Identify the Energies Involved in Melting and Vaporizing a Substance

The energy required to change a substance from its solid state (melting) to its liquid state is called the enthalpy of fusion, or heat of fusion. The energy required to change a substance from its liquid state (vaporizing) to its gaseous state is called the enthalpy of vaporization, or heat of vaporization.

Determine Which Energy is Larger

For a given substance, the enthalpy of vaporization is always larger than the enthalpy of fusion.

Explain Why the Enthalpy of Vaporization is Larger

The reason why the enthalpy of vaporization is larger than the enthalpy of fusion is due to the differences in the forces between particles in the various states of matter. In the solid state, particles are held together by strong intermolecular forces. When a substance melts, some of these intermolecular forces are broken, and the particles are allowed to move more freely in the liquid state. However, in the liquid state, there are still some intermolecular forces holding the particles together. When a substance vaporizes, the remaining intermolecular forces between the particles must be completely overcome, allowing the particles to move independently as a gas. This requires more energy than breaking the intermolecular forces during the melting process. As a result, the enthalpy of vaporization is always larger than the enthalpy of fusion.

Key concepts

Enthalpy of Fusion

The enthalpy of fusion, also known as the heat of fusion, is a thermodynamic quantity representing the amount of energy needed to change one mole of a substance from the solid phase to the liquid phase at constant pressure. It is an essential concept in understanding phase changes. This energy is used to overcome some intermolecular forces that hold the solid particles together, allowing them to move past one another more freely in the liquid state.

For water, for example, the enthalpy of fusion is 6.01 kJ/mol. During this phase change, the temperature of the substance does not increase despite the input of energy. This energy is instead used solely for the phase change. Understanding enthalpy of fusion is important not only in chemistry and physics but also in various industries, including food production and cryogenics.

Enthalpy of Vaporization

The enthalpy of vaporization is the energy required to convert one mole of a liquid into a gas at constant pressure. Compared to the enthalpy of fusion, this energy input is significantly higher. For instance, water's enthalpy of vaporization is approximately 40.65 kJ/mol, much greater than its enthalpy of fusion.

The reason for such a high energy requirement lies in the need to completely break the intermolecular forces that keep the liquid molecules closely packed. As these forces are entirely overcome, the molecules can move independently in the gaseous phase, resulting in a gas that has neither a definite volume nor shape. This phase change, like melting, occurs at a specific temperature called the boiling point.

Intermolecular Forces

Intermolecular forces are the forces of attraction or repulsion between neighboring particles (atoms, molecules, or ions). They are significantly weaker than the intramolecular forces that hold atoms together within a molecule but are crucial in determining the physical properties of substances, including their states of matter at given temperatures and pressures.

There are several types of intermolecular forces, such as hydrogen bonds, dipole-dipole interactions, and London dispersion forces. The strength of these forces directly influences the enthalpy of fusion and vaporization. In the transition from solid to liquid and from liquid to gas, varying amounts of energy are required to overcome these intermolecular forces, thus explaining why the enthalpy of vaporization is usually greater than the enthalpy of fusion.

States of Matter

Matter can exist in different states, primarily solid, liquid, and gas. These states are distinguished by the density of particles and how they move relative to each other. In a solid, particles are tightly packed in a fixed arrangement and can only vibrate in place. Solids have a fixed shape and volume. As we transition to a liquid, the particles have enough energy to move around each other but are still close together; liquids have a definite volume but no fixed shape.

Upon further heating, a liquid becomes a gas. The gas state is characterized by particles that are much further apart and move freely at high speeds. Gases do not possess a fixed volume or shape and will expand to fill the container they are in. The phase changes between these states of matter involve either an increase or decrease in energy, relating to the concepts of enthalpy of fusion and vaporization discussed previously.