Phase Diagram & Heating Curve Explorer

Select a substance and explore its P-T phase diagram. Click anywhere on the diagram to identify the phase at that point. Adjust pressure to see how melting and boiling points shift, view the heating curve at constant pressure, and explore colligative effects of dissolved solutes.

Controls

Substance

Presets

Heating Curve Pressure1.00 atm

0.001 atm300 atm

Colligative Effects

Solute Molality (m)0.00 mol/kg

van't Hoff Factor (i)1.0

Results

Liquid

at 25.0°C, 1.00 atm

Transition Points at 1.00 atm

Melting point 0.0°C

Boiling point 100.0°C

Water (H₂O) Data

Triple point0.01°C, 0.0060 atmCritical point374°C, 218 atmΔH_fus6.01 kJ/molΔH_vap40.67 kJ/mol

Heating Curve at 1.00 atm

Solid heating1.90 kJ/mol

Melting (fusion)6.01 kJ/mol

Liquid heating7.53 kJ/mol

Boiling (vaporization)40.67 kJ/mol

Gas heating3.36 kJ/mol

Total energy59.47 kJ/mol

Clausius-Clapeyron Equation

\ln\!\left(\frac{P_2}{P_1}\right) = \frac{\Delta H_{\text{vap}}}{R}\left(\frac{1}{T_1} - \frac{1}{T_2}\right)

where T is in Kelvin and ΔH_vap = 40.67 kJ/mol

P-T Phase Diagram

Click anywhere on the diagram to probe the phase at that point.

SolidLiquidGasTriple PointCritical PointHeating P

Heating Curve at 1.00 atm

Temperature vs energy added for 1 mole of substance at constant pressure.

Solid heatingMelting (fusion)Liquid heatingBoiling (vaporization)Gas heating

Reference Guide

Phase Diagram Regions

A P-T phase diagram maps pressure vs temperature, dividing space into regions where each phase is stable. Three boundary curves separate these regions.

Solid-Liquid (melting/freezing curve). Nearly vertical for most substances. Water is anomalous with a negative slope.
Liquid-Gas (boiling/condensation curve). Ends at the critical point.
Solid-Gas (sublimation/deposition curve). Below the triple point.

Above the critical point, the substance exists as a supercritical fluid with properties of both liquid and gas.

Triple and Critical Points

The triple point is the unique temperature and pressure where all three phases coexist in equilibrium. For water, this is at 0.01°C and 0.006 atm.

The critical point marks the highest temperature and pressure at which a liquid-gas boundary exists. Beyond this point, heating a liquid does not produce a visible phase transition. For water, this is 374°C and 218 atm.

CO₂ has a triple point at 5.18 atm, which is why dry ice sublimes at normal atmospheric pressure. You need at least 5.18 atm to see liquid CO₂.

Clausius-Clapeyron Equation

The Clausius-Clapeyron equation describes how vapor pressure varies with temperature along a phase boundary.

\ln\!\left(\frac{P_2}{P_1}\right) = \frac{\Delta H_{\text{vap}}}{R}\left(\frac{1}{T_1} - \frac{1}{T_2}\right)

where $R = 8.314$ J/(mol K) is the gas constant and $\Delta H_{\text{vap}}$ is the enthalpy of vaporization. This explains why water boils at about 93°C in Denver (elevation 1,600 m, ~0.83 atm) instead of 100°C.

Colligative Effects on Phase Boundaries

Dissolving a solute in a solvent shifts the phase boundaries. The freezing point drops and the boiling point rises by amounts that depend only on the number of dissolved particles.

\Delta T_f = -i \cdot K_f \cdot m

\Delta T_b = +i \cdot K_b \cdot m

For water, $K_f = 1.86$ °C/m and $K_b = 0.512$ °C/m. Adding 1 mol/kg of NaCl ( $i = 2$ ) lowers the freezing point by 3.72°C, which is why salt melts ice on roads.