No, the evaporation of water is not an exothermic process; it is an endothermic process.
Understanding Endothermic and Exothermic Processes
To understand why evaporation is endothermic, let's define these terms:
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Exothermic Process: A process that releases heat to the surroundings. The system loses energy, and the surroundings become warmer. Examples include combustion (burning) and freezing.
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Endothermic Process: A process that absorbs heat from the surroundings. The system gains energy, and the surroundings become cooler. Examples include melting and boiling/evaporating.
Evaporation as an Endothermic Process
Evaporation, specifically the process of water changing from a liquid to a gaseous state (water vapor), requires energy input. This energy is used to overcome the intermolecular forces (hydrogen bonds) holding the water molecules together in the liquid phase. Since heat is absorbed during this process, it is endothermic.
Think about sweating. When sweat evaporates from your skin, it absorbs heat from your body, which is why you feel cooler. The water needs to absorb energy (heat) to change its state.
Spontaneity vs. Endothermicity
It's important to note that while evaporation is endothermic, it can still be spontaneous under certain conditions (e.g., at temperatures above the boiling point or when the partial pressure of water vapor in the air is low). Spontaneity is governed by Gibbs Free Energy (ΔG), which considers both enthalpy (ΔH, heat absorbed or released) and entropy (ΔS, change in disorder). Evaporation increases entropy (gas is more disordered than liquid). Even though ΔH is positive (endothermic), if the TΔS term (temperature times the change in entropy) is large enough, ΔG can be negative, making the process spontaneous. However, it does not change the fact that the evaporation process itself requires energy input, making it endothermic.
