A phase diagram is a graphical tool used across chemistry and physics to map out the stable physical states, or phases, of a substance under varying conditions. This two-dimensional chart visually summarizes the complex relationship between temperature, pressure, and the form a substance takes (solid, liquid, or gas). Understanding how to read this diagram is necessary for predicting a substance’s state. The phase diagram shows that the boiling point is not a fixed number but a variable condition dependent on the external pressure applied to the system.
Interpreting the Axes and Phase Regions
A standard phase diagram for a pure substance is a Pressure-Temperature (P-T) plot, featuring pressure on the vertical axis and temperature on the horizontal axis. These two variables are the primary external factors that determine the physical state of the material. The graph is divided into three large, distinct areas, each representing a single stable phase: solid, liquid, and gas, or vapor.
The solid region occupies the upper-left area, found at high pressures and low temperatures. Conversely, the gas or vapor region dominates the bottom-right, situated at low pressure and high temperature. The liquid region lies between these two extremes, generally found at moderate pressures and temperatures.
Any point plotted within one of these three regions signifies conditions where the substance exists entirely in that single, stable phase. For instance, slight changes in pressure or temperature deep within the liquid region will not cause a phase change. The boundaries separating these regions, however, represent conditions where two phases can stably coexist in equilibrium.
Locating the Boiling Point: The Vaporization Curve
The boiling point is not represented by a single point on the diagram, but rather by an entire curve known as the vaporization curve, or the liquid-vapor boundary. This line separates the liquid region from the gas/vapor region and represents all combinations of pressure and temperature at which the liquid and gaseous phases of the substance are in dynamic equilibrium. As a liquid is heated, it will begin to boil precisely when its vapor pressure equals the external pressure exerted on it.
A specific point on this curve, known as the normal boiling point, is found by tracing the line corresponding to a standard pressure, such as one atmosphere (1 atm), across to the vaporization curve. For water, this intersection occurs at 100 degrees Celsius, which is the definition of its normal boiling point. Any point along the vaporization curve indicates the boiling temperature for that corresponding pressure.
Moving up the vaporization curve means increasing the pressure, which necessitates a higher temperature to reach the boiling point. This relationship explains why water boils at a lower temperature at high altitudes, where the atmospheric pressure is lower. Conversely, increasing the pressure requires a higher thermal input for the liquid to transition into a gas.
Defining the Boundary Limits: Triple and Critical Points
The vaporization curve is not infinite; it is precisely bounded by two unique and highly significant thermodynamic markers known as the triple point and the critical point. The triple point is the singular pressure and temperature combination where the solid, liquid, and gas phases of a substance all coexist in stable equilibrium. This point is where the vaporization curve, the fusion (melting) curve, and the sublimation curve all intersect.
Because the triple point represents a fixed, reproducible condition for a pure substance, it serves as a reference in thermometry; for example, the triple point of water is used in the definition of the Kelvin temperature scale. At pressures below this point, a substance cannot exist as a liquid, and heating the solid phase will cause it to sublime directly into a gas.
On the opposite end of the vaporization curve lies the critical point, defined by the critical temperature and critical pressure. This is the absolute limit of the liquid-gas boundary. Beyond this point, the distinction between the liquid and gas phases disappears entirely. The substance enters a state known as a supercritical fluid, which possesses properties intermediate between a gas and a liquid (e.g., the ability to diffuse like a gas and dissolve materials like a liquid).