5.4.2. Enthalpy Changes during Phase Transformations

Phase transformations also involve energy changes. Ice, for example, requires heat for melting. Normally this melting takes place at constant pressure (atmospheric pressure) and during phase change, temperature remains constant (at 273 K).

Phase transformations, such as melting (from solid to liquid), involve changes in energy. For instance, when ice melts, it absorbs heat from its surroundings. Interestingly, during this phase change, while heat is absorbed, the temperature of the ice remains constant at 0°C (273 K) until all the ice has melted. This is because the added energy is used to break the intermolecular bonds in the ice rather than increasing the temperature.

Think of ice melting in a warm room. You might notice that even though the warm air around the ice keeps bringing in heat, the ice does not start getting hotter until it has completely melted. This is similar to how the temperature of a pot of water stays at 100°C until all the water has boiled away.

H2O(s) → H2O(l); ∆fusH° = 6.00 kJ mol–1 Here ∆fusH° is enthalpy of fusion in standard state. If water freezes, then process is reversed and equal amount of heat is given off to the surroundings.

The enthalpy of fusion is the amount of heat required to melt one mole of a solid at its melting point. For ice, the enthalpy of fusion (∆fusH°) is +6.00 kJ/mol. This positive value indicates that heat is absorbed when ice melts to form liquid water. Conversely, when water freezes, it releases the same amount of energy (6.00 kJ/mol) back to the surroundings, highlighting the conservation of energy during phase changes.

Imagine making ice cubes in a freezer. You fill an ice tray with water and place it in the freezer. As the water freezes, it gives off heat to its surroundings, which helps maintain the temperature of the freezer. This is why freezers can feel warm on the outside; they are releasing heat from the water as it turns into ice.

H2O(l) → H2O(g); ∆vapH° = +40.79 kJ mol–1 ∆vapH° is the standard enthalpy of vaporization.

The enthalpy of vaporization is the amount of heat required to convert one mole of a liquid into gas at its boiling point. For water, the enthalpy of vaporization is +40.79 kJ/mol, meaning that when water transitions from liquid to gas (steam), it must absorb a significant amount of heat. This process occurs at constant temperature and typically at atmospheric pressure.

Consider boiling water on a stove. While the water heats up to 100°C, once it reaches that temperature, it doesn't get any hotter. Instead, it starts turning into steam, and this requires energy. The energy provided by the stove is what allows the water to transition from liquid to gas, illustrating the enthalpy of vaporization.

Sublimation is the direct conversion of a solid into its vapor. Solid CO2 or 'dry ice' sublimes at 195K with ∆subH° = 25.2 kJ mol–1; naphthalene sublimes slowly and for this ∆subH° = 73.0 kJ mol–1.

Sublimation is a special phase transition where a solid changes directly into a gas without passing through the liquid phase. For example, dry ice, which is solid carbon dioxide (CO2), sublimates at temperatures below -78.5°C (195 K), requiring 25.2 kJ of energy per mole. Other compounds, like naphthalene, also sublime but do so at different rates and energy requirements. This energy change is referred to as the standard enthalpy of sublimation.

If you've ever seen dry ice at a party or Halloween event, you might notice it 'smoking' as it sublimates, creating fog-like effects. This is the solid turning into gas, showcasing sublimation in action, and that fog is water vapor condensing from the air.

The magnitude of the enthalpy change depends on the strength of the intermolecular interactions in the substance undergoing the phase transformations.

The amount of energy required for phase changes is influenced by how strongly the molecules are held together by intermolecular forces. For instance, water has strong hydrogen bonds between its molecules, which means it requires more energy to break these bonds during melting and vaporization compared to a substance like acetone, which has weaker interactions.

Imagine trying to pull apart two pieces of paper. If you use glue to stick them together, it takes more force (energy) to separate them than if you just tapped them lightly with some tape. Similarly, ice holds onto itself tightly because of hydrogen bonds, requiring more energy to break these bonds than substances with weaker molecular interactions.

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