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A critical point in a phase diagram refers to the unique combination of temperature and pressure at which the distinction between liquid and gaseous phases vanishes. At this point, the liquid phase and the gas phase become identical, and their physical properties like density and viscosity become the same. This phenomenon is known as phase coexistence. Beyond the critical point, we can only observe a single-phase called the supercritical fluid. Supercritical fluids exhibit properties of both liquids and gases, which is useful in various industries as solvents for chemical reactions.

The line that separates the gas and liquid phases in a phase diagram is called the vaporization curve or vapor-liquid equilibrium curve. This curve helps us understand the equilibrium behavior of substances when they are subjected to different pressure and temperature conditions. When we move along this curve, the system is at equilibrium, which means that both liquid and gas phases coexist. As we reach the critical point, phase boundaries disappear due to the fusion of gas and liquid phases into a supercritical fluid. This occurrence is caused by the change of the substance's properties, particularly the van der Waals forces which hold particles together in a liquid phase. At the critical point, the attractive forces are not strong enough to keep particles in a liquid state exclusively, but spatial proximity allows the particles to move closer and remain in a fluid state. The critical point represents the saturation pressure and temperature at which gas and liquid phases can no longer be distinguished. As a result, the line separating the gas and liquid phases ends at the critical point, and beyond that point, the substance exists solely as a supercritical fluid.