Why does the solid-liquid line for water slope to the left, while for most substances it slopes to the right?

This anomalous slope is due to water's expansion upon freezing. The Clapeyron equation slope dP/dT = ΔH/(TΔV) depends on the volume change ΔV. For most substances, ΔV is positive (liquid less dense than solid), giving a positive slope. For water, the solid (ice) is less dense, so ΔV is negative, resulting in a negative slope. This means increasing pressure on ice lowers its melting point.

What exactly happens at the critical point?

At the critical point, the distinction between liquid and gas vanishes. Properties like density, enthalpy, and refractive index of the two phases become identical. Beyond this point, the substance becomes a supercritical fluid, which has properties of both a gas and a liquid, such as high density and excellent solvent power combined with low viscosity and high diffusivity.

Can this simulator show what happens if I heat a substance at constant pressure?

Yes. A horizontal line at a fixed pressure represents an isobaric process. Moving from left to right along this line shows the substance transitioning through its phases as temperature increases. For a pressure below the triple point, you would see a direct transition from solid to gas (sublimation). Above the triple point, you cross the solid-liquid line (melting) and then the liquid-gas line (vaporization).

Does the model show metastable states, like superheated ice or supercooled water?

No. This is a key simplification. The diagram shows only the most thermodynamically stable phase for a given (T,P) condition. In reality, phases can persist outside their stable regions (metastability), but these states are not at global equilibrium and are not represented on a standard equilibrium phase diagram.

Phase Diagram

Temperature-pressure diagram with phase transitions.

About this model

Phase diagrams map the stable states of matter—solid, liquid, and gas—as a function of temperature and pressure. This interactive model visualizes the classic phase diagram for a pure substance, such as water or carbon dioxide. The core physics is governed by the Clapeyron equation, dP/dT = ΔS/ΔV = ΔH/(TΔV), which describes the slope of the coexistence lines where two phases are in equilibrium. The simulator prominently features the triple point, where solid, liquid, and gas coexist in a unique thermodynamic equilibrium, and the critical point, beyond which the distinct liquid and gas phases merge into a supercritical fluid. The model simplifies real-world behavior by assuming a pure, single-component system with no metastable states (like supercooled liquid) shown. It also typically depicts the solid-liquid line with a negative slope for water (anomalous expansion) or a positive slope for most other substances. By manipulating temperature and pressure, students directly explore how these intensive variables dictate phase stability, reinforcing concepts from Gibbs free energy and the conditions for phase equilibrium. The simulator concretely illustrates why increasing pressure can melt ice (for water) or why boiling point changes with altitude, linking abstract thermodynamic principles to observable phenomena.

Who it's for: Undergraduate students in chemistry, physics, or engineering taking introductory thermodynamics or physical chemistry courses.

Key terms

Phase Diagram
Triple Point
Critical Point
Coexistence Curve
Clapeyron Equation
Phase Transition
Supercritical Fluid
Gibbs Free Energy

How it works

A schematic pressure–temperature diagram: solid, liquid, gas, and supercritical regions are separated by fusion (melting), vaporization (boiling), and sublimation curves meeting at a triple point. The dashed curves are guides; coordinates T̂, P̂ are normalized for teaching, not experimental water data. Drag the sliders to move the yellow dot and read off the phase.

Key equations

Along coexistence curves: chemical potentials equal (µ₁ = µ₂)

Clapeyron: dP/dT = ΔS/ΔV along a boundary

Frequently asked questions

Why does the solid-liquid line for water slope to the left, while for most substances it slopes to the right?: This anomalous slope is due to water's expansion upon freezing. The Clapeyron equation slope dP/dT = ΔH/(TΔV) depends on the volume change ΔV. For most substances, ΔV is positive (liquid less dense than solid), giving a positive slope. For water, the solid (ice) is less dense, so ΔV is negative, resulting in a negative slope. This means increasing pressure on ice lowers its melting point.
What exactly happens at the critical point?: At the critical point, the distinction between liquid and gas vanishes. Properties like density, enthalpy, and refractive index of the two phases become identical. Beyond this point, the substance becomes a supercritical fluid, which has properties of both a gas and a liquid, such as high density and excellent solvent power combined with low viscosity and high diffusivity.
Can this simulator show what happens if I heat a substance at constant pressure?: Yes. A horizontal line at a fixed pressure represents an isobaric process. Moving from left to right along this line shows the substance transitioning through its phases as temperature increases. For a pressure below the triple point, you would see a direct transition from solid to gas (sublimation). Above the triple point, you cross the solid-liquid line (melting) and then the liquid-gas line (vaporization).
Does the model show metastable states, like superheated ice or supercooled water?: No. This is a key simplification. The diagram shows only the most thermodynamically stable phase for a given (T,P) condition. In reality, phases can persist outside their stable regions (metastability), but these states are not at global equilibrium and are not represented on a standard equilibrium phase diagram.

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