The heat of vaporization and the boiling point are fundamentally different properties of a substance, though they are closely related during phase change. The distinction is simple: one is a measure of temperature (a condition), and the other is a measure of energy (a quantity). Clearly defining and separating these two concepts is necessary to understand how a liquid turns into a gas.
Boiling Point: The Temperature Threshold
The boiling point is the specific temperature at which a liquid rapidly changes into a gas throughout its volume. This temperature is reached when the liquid’s vapor pressure equals the external pressure, such as atmospheric pressure. When this equality is reached, the liquid forms vapor bubbles internally, a process known as boiling.
The boiling point is not fixed for every substance; it is directly dependent on the external pressure. For example, water boils at \(100^\circ\text{C}\) at standard sea-level pressure. At higher altitudes, where atmospheric pressure is lower, water boils at a lower temperature.
Heat of Vaporization: The Energy Requirement
The heat of vaporization, formally known as the enthalpy of vaporization (\(\Delta H_{vap}\)), is the specific amount of thermal energy required to convert a set amount of liquid into a gas. This energy is generally quantified for one mole or one kilogram of the substance and is measured at its boiling point.
This energy is absorbed without causing an increase in temperature and is referred to as “latent heat.” Instead of increasing the kinetic energy of the molecules, this energy is used entirely to overcome the strong attractive forces holding the liquid molecules together. The magnitude of the heat of vaporization is thus a direct measure of the strength of the intermolecular forces within that liquid.
Distinguishing Temperature and Energy in Phase Change
The key difference between the boiling point and the heat of vaporization lies in what each one measures during the phase change. The boiling point is the temperature at which liquid and gas coexist in equilibrium, acting as the temperature gateway for the transition. Once a liquid reaches this temperature, additional heat supplied does not increase the temperature further until all the liquid converts to gas.
This added heat is the heat of vaporization, representing the energy cost of crossing that gateway. The energy is consumed to break the remaining intermolecular bonds, increasing the potential energy of the molecules. Since the energy is not translated into faster molecular motion, the temperature remains constant throughout the vaporization process. The transition is a two-step thermal process: first, reaching the boiling point, and second, absorbing the heat of vaporization to break bonds at a constant temperature.
Practical Applications of Latent Heat
The high heat of vaporization of certain substances, particularly water, has wide-ranging practical implications, primarily in cooling systems. When liquid water turns into water vapor, it pulls a massive amount of heat from its surroundings due to its high \(\Delta H_{vap}\). This principle is the basis for evaporative cooling, which is how the human body regulates its temperature through sweating.
As sweat evaporates from the skin’s surface, it draws latent heat directly from the body, leading to a significant cooling effect. Conversely, the same large energy quantity is released during condensation. This explains why steam causes severe burns; when steam contacts cooler skin, it rapidly condenses back into liquid water and releases its high latent heat into the tissue. Furthermore, the vapor-compression refrigeration cycle used in air conditioners and refrigerators relies on the latent heat of vaporization, as a refrigerant absorbs heat from the interior space when it evaporates and then releases that heat outside when it condenses.