Phase (matter)

In the physical sciences, a phase is a region of material that remains chemically uniform and physically distinct. Within a single phase, all physical properties stay essentially the same. These properties include density, chemical composition, magnetization, and the index of refraction. A phase is more than just a state of matter. For example, oil and water are both liquids, but they are immiscible. This means they do not mix and instead form two separate liquid phases.

Phase separation often occurs because of solubility. Solubility is the maximum amount of a solute that can dissolve in a solvent. If a substance cannot dissolve, it remains in a separate phase. This concept applies to many systems. In liquids, as many as eight immiscible phases have been observed together. These can include an aqueous phase of water, hydrophobic organic solvents, perfluorocarbons, silicones, and even certain metals or molten phosphorus.

Phases can also exist within a single state of matter. For instance, iron alloys contain several different solid and liquid phases. In solids, materials can form solid solutions or crystallize into different crystal phases. Some materials show polymorphism, which is the ability to exist in more than one crystal form. When this happens in pure chemical elements, it is called allotropy. A famous example is carbon, which can exist as diamond, graphite, or fullerenes.

Scientists use phase diagrams to map how these phases change. A phase diagram is a chart showing which phases are possible at specific temperatures and pressures. In a single-component system, the phases depend only on these two variables. These diagrams often feature a triple point. This is a specific intersection where all three phases can coexist in equilibrium.

Phase diagrams also show a critical point. This is a point where the boundary between liquid and gas ends. As temperature and pressure approach this point, the liquid and gas become more similar. At the critical point, they become indistinguishable. Beyond this point, the substance becomes a supercritical fluid. For water, the critical point occurs at approximately 647 K or 374 °C, at a pressure of 22.064 MPa.

Water has an unusual feature in its phase diagram. Most substances have a positive slope for the solid-liquid phase line. Water has a negative slope instead. This is because ice has a lower density than liquid water. Increasing the pressure can actually drive water into the higher density liquid phase, causing it to melt. This unique behavior is a key part of how water interacts with its environment.

When a substance moves from one phase to another, it undergoes a phase transition. These transitions involve the movement of energy. When water evaporates, it absorbs kinetic energy to break the attractive forces of the liquid. This process, known as the enthalpy of vaporization, actually cools the remaining liquid. Conversely, condensation releases heat. The energy required to change a solid to a liquid is called the enthalpy of fusion. The energy for a solid to change directly to a gas is the enthalpy of sublimation.

Finally, we must consider the interface between phases. Between two phases in equilibrium, there is a narrow interfacial region. In this thin zone, the properties are not exactly like either phase. This region can cause observable effects like surface tension. While most studies focus on thermal equilibrium, researchers also study phases out of equilibrium. This includes quantum many-body localized systems, which allow for new types of order through localization protected quantum order.