Phase transitions, such as water turning into ice or steam, are driven by changes in temperature or pressure. The phase of a substance at any given moment depends entirely on these two external conditions. The triple point is a unique and precise combination of temperature and pressure where a substance’s three main phases—solid, liquid, and gas—can all coexist in perfect thermodynamic equilibrium. At this specific juncture, the substance is simultaneously freezing, melting, boiling, and condensing, but the total amount of each phase remains constant. This singular point represents a fundamental physical property for every pure substance.
The Specific Conditions Required
The triple point is a single, exact state defined by a unique temperature and pressure value for any pure substance. This point can be visualized on a phase diagram as the intersection where the three phase-boundary curves meet: the sublimation, melting, and vaporization curves. These lines represent the conditions where two phases are in equilibrium, and their intersection is the only point where all three phases balance one another.
For water, the most familiar substance, this specific point occurs at a temperature of 273.16 Kelvin (0.01 degrees Celsius). At this temperature, the pressure must also be exactly 611.73 Pascals. This pressure is extremely low, less than one percent of the atmospheric pressure at sea level, which explains why we do not commonly observe water’s triple point in everyday life.
If the temperature or pressure moves even slightly away from these exact values, the three-phase equilibrium will be lost, and one phase will begin to dominate. For example, a tiny increase in temperature at this exact pressure will cause all the ice to melt and the vapor to condense into liquid water. Conversely, a slight decrease in pressure will cause the liquid water to freeze or the ice to sublimate directly into vapor.
Why Three Phases Exist Simultaneously
The coexistence of solid, liquid, and gas at the triple point is a dynamic equilibrium. This means phase changes are constantly occurring at the molecular level, but the rates of opposing changes are precisely equal, resulting in no net change in the overall composition of the system. For example, the rate of melting is exactly matched by the rate of freezing. Similarly, evaporation balances condensation, and sublimation balances deposition.
This complex, three-way molecular traffic jam is sustained because the energy state (Gibbs free energy) of the solid, liquid, and gas phases is identical at this one temperature and pressure. Because no one phase is thermodynamically favored over the others, the system remains perfectly stable. The slightest deviation in temperature or pressure, however, changes the energy balance, immediately favoring one phase and causing the others to disappear.
Its Role in Defining Temperature Scales
The triple point of water is profoundly important because it is an invariant physical constant of nature. Unlike normal boiling or freezing points, which change depending on atmospheric pressure, the triple point is fixed for a pure substance. This makes it an ideal, highly reproducible reference point for defining temperature.
Historically, the kelvin, the base unit of thermodynamic temperature in the International System of Units (SI), was defined using the triple point of water. The temperature of the triple point of water was precisely set at 273.16 K, and one kelvin was defined as exactly 1/273.16 of that temperature. This provided a universal and reliable standard for all temperature measurements.
While the definition of the kelvin was revised in 2019 to be based on the Boltzmann constant, the practical importance of the water triple point remains immense. Highly precise instruments called triple-point cells are still used in metrology laboratories worldwide to calibrate thermometers. These cells use extremely pure water to realize this fixed temperature, providing a gold standard for accuracy. The International Temperature Scale of 1990 (ITS-90) utilizes the triple points of several substances as fundamental fixed points. This allows scientists and engineers to achieve highly precise, reproducible temperatures for the most demanding calibration work, maintaining the integrity of global temperature standards.
Examples in Everyday Substances
While water’s triple point is the most famous, every pure chemical substance has its own unique triple point. For instance, carbon dioxide (dry ice) has a triple point that occurs at a much higher pressure and a lower temperature than water’s: 517 kilopascals (about 5.1 atmospheres) and -56.6 degrees Celsius.
Because the triple point pressure for carbon dioxide is significantly higher than normal atmospheric pressure, liquid carbon dioxide cannot exist under ordinary conditions. This is why solid dry ice sublimates, turning directly into gas without melting when left out in the open. Oxygen, another common substance, has a triple point at the extremely low temperature of 54.36 Kelvin (about -218.79 degrees Celsius).
These varying points underscore that the triple point is a characteristic physical fingerprint for each substance. The pressure and temperature coordinates are fixed solely by the substance’s inherent molecular properties. Understanding these unique points is necessary for chemical manufacturing, cryogenics, and any field that involves precise control over phase changes.