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Phase boundaries on a phase diagram indicate the conditions of pressure and temperature under which two phases of a substance are in equilibrium. These lines separate different phases on the diagram.

The triple point in a phase diagram represents the unique set of conditions (pressure and temperature) at which three phases of a substance (solid, liquid, and gas) can coexist in equilibrium.

The critical point marks the end of the liquid-gas boundary on the phase diagram. Beyond this point, the liquid and gas phases become indistinguishable, and the substance exists as a supercritical fluid.

The Gibbs Phase Rule, described by the equation $F = C - P + 2$, where $F$ is the degrees of freedom, $C$ is the number of components, and $P$ is the number of phases, helps determine the number of variables that can be changed independently without altering the number of phases present.

The areas on the phase diagram labeled as solid, liquid, or gas indicate the range of pressures and temperatures over which that phase is stable and exists as the only phase.

The slope of the solid-liquid boundary on a phase diagram indicates whether the solid phase is more or less dense than the liquid phase. A positive slope means the solid is denser than the liquid, while a negative slope suggests the reverse.

The fusion curve represents the boundary on a phase diagram that separates solid and liquid phases. It shows how the melting point of a substance changes with pressure.

A eutectic point represents a fixed combination of temperature and pressure where a mixture of substances can coexist as solids and a liquid, resulting in the lowest melting point for the mixture.

On a phase diagram, the supercritical fluid region lies beyond the critical point. In this region, the substance doesn’t have distinct liquid or gas phases. Conditions here allow the supercritical fluid to exhibit unique properties.