Heat is the transfer of thermal energy between systems of different temperatures. As a liquid absorbs heat, this energy transfer increases the kinetic energy of its molecules, causing them to move more vigorously. This rise in the average kinetic energy registers externally as an increase in the liquid’s temperature. The added energy weakens the intermolecular forces holding the liquid structure together, preparing it for a change in physical state.
The Initial Phase Transition: Boiling and Vaporization
As the temperature increases, molecules with enough kinetic energy escape the surface, a process known as evaporation. Evaporation is a surface phenomenon that occurs at any temperature below the boiling point. This slow vaporization uses the internal energy of the liquid, which has a slight cooling effect on the remaining liquid.
When heating is rapid and sustained, the liquid reaches its boiling point. This is the specific temperature at which the vapor pressure of the liquid equals the surrounding atmospheric pressure. At this point, vaporization is no longer confined to the surface but becomes a bulk phenomenon, occurring throughout the entire volume of the liquid, forming vapor bubbles that rise and burst.
Once the boiling point is reached, the liquid’s temperature stops rising, even as more heat energy is added. This additional energy is the latent heat of vaporization, required to completely break the remaining intermolecular bonds and facilitate the phase change to gas. For water at sea level, this constant temperature is 100°C (212°F). The conversion from liquid to vapor is accompanied by a massive increase in volume, such as one cubic meter of water converting into approximately 1,600 cubic meters of steam.
The Consequences of Containment: Vapor Pressure and Expansion
The conversion of a liquid to a gas phase has significant consequences, particularly in a closed system. As the liquid is heated, its vapor pressure—the pressure exerted by the gas molecules above the liquid—increases exponentially with temperature. This occurs because higher temperatures mean more molecules have the kinetic energy necessary to escape the liquid phase and contribute to the pressure.
In a sealed container, this dramatic rise in vapor pressure can quickly exceed the structural limits of the vessel. For example, a pressure cooker’s sealed lid allows vapor pressure to build up, which raises the boiling point of the liquid inside.
Another effect of heating is thermal expansion, where the liquid expands slightly as its molecules gain kinetic energy and occupy a larger volume. While this expansion is small for the liquid phase, it generates substantial pressure in systems completely filled with liquid and lacking gas space. Heating water trapped in a rigid pipe by just 20°C can result in a pressure increase of nearly 50 bar (725 psi). This hydraulic thermal expansion is a hazard if pressure relief mechanisms are blocked.
Reaching the Limits: Superheating and the Critical Point
Beyond normal boiling, heating a liquid can lead to superheating, where the liquid is heated above its boiling point without actually boiling. This unstable state happens when a liquid lacks nucleation sites, such as scratches or gas bubbles, where vapor bubbles can easily form. If a superheated liquid is disturbed, it can lead to a flash-boil known as bumping.
If the liquid is heated in a sealed container under extremely high pressure, the distinction between the liquid and gas phases can disappear. This occurs at the critical point, defined by a substance’s unique critical temperature and critical pressure. For water, these values are 374°C and 218 atmospheres, respectively.
Above the critical point, the substance enters the state of a supercritical fluid, a homogeneous phase that is neither a true liquid nor a true gas. A supercritical fluid combines the high density and solvent power of a liquid with the low viscosity and high diffusivity of a gas. This unique state is not governed by typical phase change rules, as there is no boundary separating the liquid and gas phases.