The transformation of liquid water into a gas (water vapor or steam) is a fundamental change of state. Vaporization requires an input of energy to overcome the attractive forces holding the water molecules together in their liquid form. This change can occur through two distinct pathways: the slow, surface-level process of evaporation or the rapid, bulk process of boiling.
The Role of Molecular Kinetic Energy
Water molecules are held in a liquid state by cohesive forces, known as hydrogen bonds. When heat is added to the liquid, it increases the internal motion of the molecules. This molecular motion is referred to as kinetic energy, which is directly related to the water’s temperature.
As the temperature rises, the average speed of the water molecules increases. This increased motion puts a strain on the hydrogen bonds linking the molecules together. Once a molecule gains enough kinetic energy to completely break free, it can escape the liquid and transition into the gaseous phase as water vapor.
Evaporation: The Gradual Surface Process
Evaporation is the gradual conversion of liquid water to a gas that occurs at temperatures below the boiling point. Within a body of liquid, not all molecules possess the same amount of kinetic energy. The fastest-moving molecules located near the liquid’s surface are the ones that are most likely to possess sufficient energy to overcome the surface tension and escape into the air.
Evaporation is a relatively slow process that does not involve the formation of bubbles within the liquid. The rate is influenced by external factors, such as surface area and ambient temperature. A larger surface area provides more opportunity for molecules to exit, and warmer air accelerates the process by increasing energy transfer to the liquid.
Boiling: Reaching the Saturation Point
Boiling represents a rapid phase change that occurs not just at the surface but throughout the entire volume of the liquid. This rapid vaporization begins when the liquid reaches a specific temperature known as the boiling point. At standard atmospheric pressure, the boiling point of pure water is 100°C (212°F).
The boiling point is reached when the vapor pressure of the liquid becomes equal to the pressure of the surrounding atmosphere. Vapor pressure is the pressure exerted by the gas molecules escaping the liquid. At this point, the liquid can no longer resist the formation of gas bubbles internally. These bubbles are composed entirely of water vapor and rapidly rise and burst at the surface.
Changes in external atmospheric pressure directly affect the boiling point of water. At high altitudes, the atmospheric pressure is lower, meaning water requires less energy to reach the saturation point. Consequently, water boils at a lower temperature. Conversely, increasing the pressure, such as within a pressure cooker, raises the boiling point above 100°C.