Steam is water in its gaseous phase. For pure water at standard atmospheric pressure, steam begins to form at a temperature of 100°C (212°F). This transformation from liquid water to gaseous steam represents a physical change of state.
The Boiling Point of Water
Boiling is a process where a liquid rapidly converts into vapor throughout its entire volume, not just at the surface. This occurs when the vapor pressure of the liquid equals the surrounding atmospheric pressure. For pure water at standard sea-level pressure, this temperature is 100°C (212°F).
Unlike evaporation, which occurs only at the liquid’s surface, boiling involves the formation of bubbles within the liquid. These bubbles, filled with water vapor, rise to the surface and release the gas into the atmosphere. The water maintains this constant temperature of 100°C as long as it continues to boil and convert into steam.
Beyond Boiling: Superheated Steam
Once water has transformed into steam at its boiling point, it can be heated further, resulting in “superheated steam.” Superheated steam differs from saturated steam (steam at its boiling point) because it no longer contains liquid water droplets and is considered a completely dry gas.
The properties of superheated steam are similar to a perfect gas, exhibiting higher energy content and lower density compared to saturated steam. This form of steam is typically invisible and can exist at a wide range of temperatures above the boiling point for a given pressure. Superheated steam finds widespread applications in various industries, particularly in power generation, where its high energy drives turbines.
Pressure’s Role in Steam Production
The temperature at which water produces steam is significantly influenced by external pressure. Higher pressure raises the boiling temperature, while lower pressure reduces it. This occurs because changes in pressure affect how much energy water molecules need to escape the liquid phase and become vapor.
For instance, in a pressure cooker, the sealed environment allows pressure to build up, raising the boiling point of water to temperatures around 121°C (250°F). This elevated temperature enables food to cook much faster. Conversely, at high altitudes, where atmospheric pressure is lower, water boils at a reduced temperature. On Mount Everest, for example, water boils at approximately 71°C (160°F) due to the significantly lower atmospheric pressure.
The Hidden Energy of Phase Change
During boiling, despite continuous heat input, the water’s temperature remains constant at its boiling point until all the liquid has converted to steam. This phenomenon is explained by “latent heat of vaporization.” Latent heat refers to the energy absorbed or released during a phase change without a corresponding change in temperature.
For water, the latent heat of vaporization is the specific amount of energy required to break the intermolecular bonds holding water molecules in their liquid state. Instead of increasing the kinetic energy of the molecules (which would raise the temperature), the added heat energy is used to overcome these attractive forces, allowing the molecules to transition into the gaseous phase. Only after all the liquid has turned into steam will further heat cause the steam’s temperature to rise.