How can the boundary between a liquid and a gas disappear?

Liquid and gas seem like sharply different states: a liquid has a visible surface, while a gas fills the available space. Yet that distinction does not persist at every temperature and pressure. For a pure fluid, the line separating coexisting liquid and vapor ends at a critical point. Beyond it, there can be one supercritical fluid rather than a liquid layer beneath a gas layer.

Below the critical temperature, liquid and vapor can coexist at a particular pressure for each temperature. Their densities differ, giving a recognizable interface. Approaching the critical point along this coexistence curve, the density difference decreases. At the endpoint, the two phases become indistinguishable. The diagram is schematic; the drawn curve is not a formula for calculating saturation pressure.

For carbon dioxide, the critical temperature is about 31.0°C and the critical pressure about 7.38 megapascals. The supplied point at 40°C and 10 megapascals is above both, so it is supercritical. Temperature alone is not enough for that label: the point at 40°C and 5 megapascals is below the critical pressure. A supercritical fluid is not a mixture of little liquid drops and gas bubbles.

Changing pressure and temperature in the single-phase region can strongly change density and solvent behavior without crossing a liquid–vapor coexistence line. This makes supercritical carbon dioxide useful in some extraction processes. It does not mean that every supercritical fluid has identical properties or can dissolve every substance. Exact behavior requires measured data or an appropriate equation of state. The vessel sketch explains phase structure, not high-pressure equipment or instructions for producing it.