Question

The enthalpy \((\Delta H)\) of vaporization of water is about seven times larger than water's enthalpy of fusion \((41 \mathrm{~kJ} / \mathrm{mol}\) vs. \(6 \mathrm{~kJ} / \mathrm{mol}\) ). What does this tell us about the relative similarities among the solid, liquid, and gaseous states of water?

Short answer

The enthalpy of vaporization of water being 7 times larger than its enthalpy of fusion (\(41 \mathrm{~kJ/mol}\) vs. \(6 \mathrm{~kJ/mol}\)) indicates that the intermolecular forces between water molecules are stronger in the liquid state than in the solid state. This is due to the hydrogen bonding between water molecules, which results in a unique situation where the forces are stronger in the liquid state than in the solid state. In the solid state, water molecules are organized in a lattice structure, creating spaces between molecules and reducing the overall strength of hydrogen bonding. In contrast, liquid water molecules are closer together, allowing for stronger hydrogen bonding. The gaseous state has the weakest intermolecular forces as water molecules are far apart and there is minimal hydrogen bonding.

Step-by-step solution

Understand Enthalpy of Fusion and Enthalpy of Vaporization

Enthalpy of fusion is the amount of energy required to change a substance from a solid to a liquid state at its melting point, while the enthalpy of vaporization is the amount of energy required to change a substance from a liquid to a gaseous state at its boiling point. Both these enthalpies are measures of the energy required to overcome the intermolecular forces holding the particles together.

Relate Given Enthalpies to Intermolecular Forces

In this case, we are given that the enthalpy of vaporization of water is 7 times larger than its enthalpy of fusion (\(41 \mathrm{~kJ/mol}\) vs. \(6 \mathrm{~kJ/mol}\)). This difference implies that the amount of energy required to overcome the forces holding water molecules together in the liquid state is significantly larger compared to the energy required to overcome the forces in the solid state. In other words, the intermolecular forces between water molecules are stronger in the liquid state than in the solid state.

Discuss the Relative Similarities of Water States

As a general rule, intermolecular forces are strongest in the solid state, weaker in the liquid state, and weakest in the gaseous state. However, in the case of water, the hydrogen bonding between molecules results in a unique situation where the forces are stronger in the liquid state than in the solid state. This is because, in the solid state, the water molecules are organized in a lattice structure, which creates spaces between molecules and reduces the overall strength of the hydrogen bonding. In the liquid state, the water molecules are closer together, allowing for stronger hydrogen bonding. Given this information, we can conclude that the relative similarities among the solid, liquid, and gaseous states of water are as follows: 1. The solid state has weaker intermolecular forces than the liquid state due to lattice structure. 2. The liquid state has stronger intermolecular forces compared to the solid state and gaseous state because of closer proximity and more extensive hydrogen bonding. 3. The gaseous state has the weakest intermolecular forces, as the water molecules are far apart and there is minimal hydrogen bonding. In summary, the given enthalpies of fusion and vaporization for water tell us that water's liquid state has stronger intermolecular forces compared to its solid state, resulting in unique properties not typically observed in other substances.

Key concepts

Intermolecular Forces

Intermolecular forces are the forces of attraction or repulsion that act between neighboring particles (atoms, molecules, or ions). These forces are critical in determining the physical properties of substances, such as their boiling points, melting points, and solubility. In general, intermolecular forces can be categorized into three main types: dipole-dipole interactions, London dispersion forces, and hydrogen bonds.

While all three types play roles in different scenarios, water is particularly influenced by hydrogen bonding. This form of intermolecular force is relatively strong and significantly affects how water behaves in different states. The differences in enthalpy of fusion and vaporization of water highlight the variations in intermolecular forces.

• In the solid state, the molecules are in a fixed, lattice structure that allows for hydrogen bonds but provides less flexibility.
• In the liquid state, water molecules move more freely and are closely packed. This proximity enhances hydrogen bonding.
• In the gaseous state, molecules are far apart, resulting in weak intermolecular forces.

Understanding these forces can help explain why water requires more energy to transition from liquid to gas than from solid to liquid.

Hydrogen Bonding

Hydrogen bonding is a special type of intermolecular force that occurs when a hydrogen atom is bonded to a strongly electronegative atom like oxygen, nitrogen, or fluorine. In the case of water, it is oxygen that creates hydrogen bonds. These bonds are essential to the unique properties of water.

Despite being a type of dipole-dipole interaction, hydrogen bonds are much stronger. They are responsible for water's high surface tension and its role as a universal solvent. In the liquid state, hydrogen bonds hold water molecules together, but they can break and reform, allowing water to flow.

Interestingly, hydrogen bonds in solid water (ice) create a hexagonal lattice, which results in less density compared to liquid water. This unusual feature means that ice floats on water—a property important for aquatic life.

• Hydrogen bonds create significant cohesion and adhesion forces.
• They contribute to water's higher enthalpy of vaporization than fusion.
• These bonds can adjust dynamically in the liquid state, allowing water to adapt to environmental changes.

This adaptability of hydrogen bonds explains why the liquid state has stronger intermolecular forces than the solid state in water.

Enthalpy of Fusion

Enthalpy of fusion is the amount of energy required to change a substance from a solid state to a liquid state at constant pressure. It's a measure of the energy needed to overcome the intermolecular forces present in a solid. When a substance melts, it absorbs energy, which is used to break some of the intermolecular bonds holding it in a fixed position.

For water, the enthalpy of fusion is relatively low compared to its enthalpy of vaporization. This indicates that less energy is needed to go from the solid state (ice) to the liquid state because the hydrogen bonds in the lattice structure of ice don't fully restrict movement; they are primed to break under slight increases in energy.

• In solids like ice, molecules have restricted movement due to strong binding forces.
• Fusion requires less energy as these forces only need to be partially overcome to gain some fluidity.
• This process differs greatly from vaporization, where almost all intermolecular forces must be completely overcome for a substance to become gaseous.

Understanding the enthalpy of fusion helps clarify why certain substances, like water, transition between states at different rates, indicating significant changes in intermolecular force requirements.