2
298.15
K
1
atm
25.01
°C
2
298.15
K
1
atm
25.01
°C
The Phase Diagram Calculator determines the phase of water (solid, liquid, gas, or supercritical fluid) at any given temperature and pressure combination. Based on the phase diagram of water — incorporating the triple point (0.01°C, 0.612 kPa), critical point (373.95°C, 22,064 kPa), and vapor pressure curve — this tool identifies which phase is thermodynamically stable at your specified conditions.
Phase diagrams are essential tools in chemistry, geology, materials science, and engineering. They map the regions of temperature-pressure space where each phase is the most stable, showing the boundaries where phase transitions occur.
The phase of water is determined by comparing the given temperature and pressure to the phase boundaries. The key reference points for water are:
$$\text{Triple Point: } T_t = 0.01°C, \; P_t = 0.611657 \text{ kPa}$$
$$\text{Critical Point: } T_c = 373.946°C, \; P_c = 22,064 \text{ kPa}$$
The liquid-vapor boundary (vaporization curve) follows the Antoine equation:
$$\log_{10}(P_{mmHg}) = 8.07131 - \frac{1730.63}{233.426 + T}$$
The phase determination logic is: if the temperature and pressure exceed the critical point, the substance is a supercritical fluid. Below the triple point temperature and at sufficient pressure, ice is stable. Below the vaporization curve, gas/vapor is stable. Otherwise, the substance is in the liquid phase. The solid-liquid boundary for water has a negative slope (unique to water) due to ice being less dense than liquid water, but this effect is significant only at very high pressures (hundreds of MPa).
Phase numbers: 1 = Solid (Ice), 2 = Liquid (Water), 3 = Gas (Steam/Vapor), 4 = Supercritical Fluid.
The phase output (1-4) indicates: 1 = Solid (ice), 2 = Liquid (water), 3 = Gas (water vapor/steam), 4 = Supercritical fluid. The distance from the triple point gives a rough measure of how far the conditions are from the unique point where all three phases coexist. Near phase boundaries, small changes in temperature or pressure can cause phase transitions. Note that this simplified model is most accurate for the liquid-vapor boundary; the solid-liquid and solid-vapor boundaries use approximations.
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At 25°C and 1 atm (standard room conditions), water is in the liquid phase. The pressure is well above the vapor pressure of water at 25°C (about 3.17 kPa), so liquid is the stable phase.
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At 25°C and only 1 kPa (about 0.01 atm), the pressure is below water's vapor pressure. Water would be in the gas phase — it would boil and evaporate under these near-vacuum conditions.
A phase diagram is a graphical representation showing the conditions of temperature and pressure under which different phases (solid, liquid, gas) of a substance are thermodynamically stable. The lines on the diagram represent phase boundaries where two phases coexist in equilibrium.
The triple point of water is 0.01°C (273.16 K) and 611.657 Pa (0.00604 atm). It is the unique temperature and pressure where ice, liquid water, and water vapor all coexist in thermodynamic equilibrium. The triple point of water is used as a fundamental reference point for the Kelvin temperature scale.
The critical point of water is 373.946°C (647.096 K) and 22,064 kPa (217.7 atm). Above this temperature and pressure, the distinction between liquid and gas disappears — the substance becomes a supercritical fluid with properties intermediate between liquid and gas.
A supercritical fluid exists above the critical temperature and pressure. It has liquid-like density and dissolving power but gas-like viscosity and diffusivity. Supercritical water is used in power plants and waste treatment. Supercritical CO₂ is widely used for decaffeination, extraction, and green chemistry applications.
Water's solid-liquid boundary has a negative slope, meaning the melting point decreases with increasing pressure. This is because ice is less dense than liquid water (ice floats). Most other substances have a positive slope — their solid phase is denser than the liquid. This anomaly is due to hydrogen bonding creating an open crystal structure in ice.
At a phase boundary, two phases coexist in equilibrium. For example, along the vaporization curve, liquid and vapor coexist at the same temperature and pressure. Moving slightly off the boundary favors one phase. The Gibbs phase rule (F = C - P + 2) determines the degrees of freedom at each point.
Yes, through supercooling. Pure water can remain liquid down to about -48°C under careful conditions. Also, at high pressures, the liquid phase is stable below 0°C because the solid-liquid boundary has a negative slope. This is why pressure can cause ice to melt.
Sublimation is the direct transition from solid to gas, bypassing the liquid phase. On the phase diagram, it occurs along the solid-gas boundary, which lies below the triple point pressure. Freeze-drying (lyophilization) exploits this by reducing pressure below the triple point so ice sublimes directly to vapor.
Most substances have phase diagrams similar to water but with a positive solid-liquid slope (solid is denser than liquid). CO₂ has a triple point at -56.6°C and 5.11 atm, so it cannot exist as a liquid at 1 atm — it sublimes directly from solid to gas (dry ice). Each substance has unique triple and critical points.
The Gibbs phase rule states F = C - P + 2, where F is degrees of freedom, C is the number of components, and P is the number of phases present. For pure water (C=1): at the triple point (P=3), F=0 (fixed T and P); on a boundary (P=2), F=1 (one variable free); in a single phase (P=1), F=2 (both T and P free).
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