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Enthalpy change, symbolized by \( \Delta H \), is a critical concept in thermodynamics that describes the heat exchange in a system under constant pressure. In the process of condensation, like that of water vapor turning into liquid water at \(100^\circ \mathrm{C}\), we observe this change directly. The enthalpy change is negative (\(\Delta H < 0\)) for exothermic processes, like condensation, meaning that heat flows out of the system and into the surroundings. This release of heat is a result of the loss of kinetic energy as the water molecules move closer together, forming the more structured arrangement of liquid water.
The enthalpy change is also closely related to the disorder in a system. As water shifts from a gas to a liquid, there is a decrease in energy and disorder, indicating a more stable and orderly state. Thus, condensation leads to a loss of both energy (heat) and entropy (disorder), confirmed by the negative signs of \(\Delta H\) and \(\Delta S\) in the process.

Condensation, a common natural phenomenon, is the phase change from a gas to a liquid. At the molecular level, condensation involves gas particles losing energy and slowing down, clustering together to form a denser phase—the liquid. This process occurs when the temperature drops to the dew point or when water vapor comes into contact with a colder surface.
During condensation, the attractive forces between water molecules overcome the kinetic energy that keeps them apart in the gaseous state. This results in a phase change which is marked by noticeable physical changes, such as the appearance of fog or dew. In the given exercise, water vapor condenses to liquid water at \(100^\circ \mathrm{C}\), showcasing a fundamental thermodynamic transformation which is both practically and theoretically significant for understanding weather patterns, the water cycle, and various industrial processes like distillation.

A phase change is the transformation from one state of matter to another, such as solid to liquid (melting), liquid to gas (evaporation), gas to liquid (condensation), or liquid to solid (freezing). Each phase change occurs at specific temperature and pressure conditions that are characteristic of the substance.
When a phase change happens, the substance generally absorbs or releases energy without changing temperature—a property we see as latent heat. During the transition, bonds between molecules are made or broken, changing the structure and arrangement of molecules. The exercise we focus on presents water vapor, an energetic and less structured phase, condensing to form liquid water, which is denser and has a fixed volume. This is an exothermic phase change, one in which energy is released into the environment, and is also accompanied by changes in entropy.

Disorder in systems, or entropy, is a fundamental concept in the second law of thermodynamics and is denoted by the symbol \(S\). In the context of thermodynamics, disorder refers to the number of microscopic configurations that are available to a system. A highly ordered system, with fewer configurations, has low entropy, while a disordered system, with many possible configurations, has high entropy.
When we observe a phase change such as condensation, entropy typically decreases because the molecules in the liquid state are more ordered than those in the gas state. In thermodynamics, increasing order is synonymous with decreasing entropy, which is precisely the scenario depicted in our exercise: as water vapor condenses at \(100^\circ \mathrm{C}\), it transitions to a more orderly liquid state, and thus, the entropy change (\(\Delta S\)) is negative. Understanding how entropy changes during phase transitions helps us comprehend not only physical and chemical processes but also the flow of energy in different systems.