The triple point on a phase diagram is the single combination of temperature and pressure where a substance’s solid, liquid, and gas phases all exist together in stable equilibrium. It appears as the spot where three boundary lines meet: the sublimation curve (solid-gas), the melting curve (solid-liquid), and the vaporization curve (liquid-gas). Unlike points along those boundary lines, where only two phases coexist, the triple point is the only place on the entire diagram where all three phases are present at once.
How the Triple Point Works
A phase diagram plots pressure on the vertical axis and temperature on the horizontal axis. Each region of the diagram represents conditions where one phase is stable. The lines separating those regions are equilibrium curves, meaning two phases can coexist anywhere along a given line. The triple point is where all three of those lines converge into a single intersection.
What makes this point special is that it has zero degrees of freedom. That idea comes from the Gibbs Phase Rule, a formula that tells you how many variables (like temperature or pressure) you can change independently without losing a phase. The rule is F = C − P + 2, where C is the number of chemical components and P is the number of phases. For a single substance with three phases present, the math gives F = 1 − 3 + 2 = 0. Zero degrees of freedom means you can’t adjust either temperature or pressure without causing at least one phase to disappear. The triple point is locked to one exact temperature and one exact pressure for every substance.
This also explains why you’ll never see four phases of the same substance coexisting in equilibrium. Three phases already use up all available degrees of freedom, so adding a fourth phase is thermodynamically impossible for a single-component system.
Water’s Triple Point
For water, the triple point occurs at 0.01 °C (273.16 K) and a pressure of about 611 Pa, which is roughly 4.6 torr. That pressure is far below normal atmospheric pressure (101,325 Pa), so you’d never encounter all three phases of water coexisting under everyday conditions. To observe it, you’d need to seal water inside a container and pump the air pressure down to that narrow window.
This particular value played a starring role in measurement science for decades. Until 2019, the Kelvin, the base unit of temperature, was defined as exactly 1/273.16 of the thermodynamic temperature of water’s triple point. Scientists chose it because the triple point is extraordinarily reproducible: prepare a sealed cell of highly purified water, and it will always reach that same temperature when all three phases are present. In May 2019, the Kelvin was redefined using the Boltzmann constant, a fundamental quantity from physics, but the triple point of water remains a practical calibration reference in laboratories worldwide.
Why Dry Ice Skips the Liquid Phase
Comparing water and carbon dioxide illustrates why the triple point matters in everyday life. Carbon dioxide’s triple point sits at −56.6 °C and 5.1 atmospheres, meaning liquid CO₂ can only exist above roughly five times normal atmospheric pressure. At the 1 atmosphere of pressure we live in, the phase diagram shows only solid and gas regions. That’s why frozen CO₂ (dry ice) sublimes directly into vapor at −78.5 °C instead of melting into a puddle. It earned the name “dry ice” precisely because it never becomes liquid under normal conditions.
Water behaves differently because its triple point pressure is far below 1 atmosphere. At everyday pressures, you’re well above the triple point on the diagram, which means liquid water sits comfortably between the solid and gas regions. That’s why ice melts into water before it evaporates, the familiar sequence we take for granted.
The Triple Point vs. the Critical Point
Phase diagrams have another notable landmark: the critical point, which sits at the upper end of the vaporization curve. While the triple point is where all three phases meet, the critical point is where the distinction between liquid and gas disappears entirely. Above the critical temperature and pressure, a substance becomes a “supercritical fluid” with properties of both liquid and gas, and no visible phase boundary exists between them.
The two points mark opposite extremes of the liquid-gas boundary line. The triple point anchors the low end, where the liquid phase first becomes possible. The critical point caps the high end, where the liquid phase ceases to be distinguishable from gas. Between them, you can boil or condense a substance at a well-defined temperature for any given pressure.
Practical Use in Freeze-Drying
One of the most common industrial applications of triple point knowledge is freeze-drying, or lyophilization. The goal is to remove water from food, pharmaceuticals, or biological samples by sublimation rather than evaporation. To make that happen, the pressure inside the freeze-dryer chamber must be brought below water’s triple point pressure of about 610 Pa. At that low pressure, ice in the product converts directly to vapor without ever passing through the liquid phase, which preserves the material’s structure and nutritional content far better than conventional drying.
In practice, freeze-dryers used for plant-based foods typically operate at 63 to 124 Pa and temperatures around −20 to −25 °C. Pharmaceutical applications push even lower, sometimes down to 5 to 20 Pa, to handle solutions containing sugars or proteins that lower the freezing point. The energy required is substantial: sublimating water takes about 2,885 kilojoules per kilogram, which is why freeze-drying is slower and more energy-intensive than simply heating something dry. But the quality of the end product, whether it’s instant coffee or a vaccine, depends on staying below that triple point pressure so liquid water never forms.