At the triple point water’s solid, liquid and gas phases coexist in equilibrium, so under the precise conditions of 0.01 °C and 611.657 pascals a sample of pure water can simultaneously exhibit ice, liquid water and water vapour.
Overview
The triple point is a unique thermodynamic state for a pure substance where three phases meet on the phase diagram. For water this point is defined by an exact temperature and pressure at which freezing, melting, evaporation and condensation processes balance so that ice, liquid water and vapour are present together and no net phase change occurs.
How it works
Phase behaviour depends on both temperature and pressure. At the triple point the energy exchanges associated with melting and vaporisation are in equilibrium: molecules escape the liquid as vapour while other vapour molecules condense, and solid and liquid exchange mass through melting and freezing. Small changes in temperature or pressure shift the balance, favouring one or two phases instead of all three.
Conditions and experiment
Reaching the triple point requires pure water and careful control of pressure and temperature. In laboratory demonstrations a sealed cell with a tiny amount of water is adjusted to the precise pressure and temperature, revealing ice, liquid and vapour in a stable coexistence. The triple point is also a reference standard in thermometry because it is reproducible for a pure substance.
Implications
Understanding the triple point clarifies how phase diagrams predict material behaviour, why boiling and freezing temperatures change with pressure, and why everyday experience (atmospheric pressure) differs from controlled laboratory conditions. It highlights that terms like \"boil\" and \"freeze\" describe processes that depend on the surrounding pressure as well as temperature.
Quick related facts
- Triple point of water: 0.01 °C and 611.657 pascals
- Phases present: solid (ice); liquid (water); gas (vapour)
- Use: reference point in precise temperature calibration