liquid and gas; and along line C, solid and gas. Crossing one of these lines represents a phase change
process: Crossing line B, for example, denotes either evaporation or condensation, depending on
the direction of travel. The intersection of the three lines is called the triple point. At this
temperature and pressure, unique for a given substance, all three phases are in equilibrium. Each
substance has its own characteristic phase diagram that describes its physical properties. The
reason why dry ice sublimes rather than melts, for example, is because the triple point of carbon
dioxide lies at a pressure above 1 atm. The process of raising its temperature in open air
(atmospheric pressure) thus occurs in the lower portion of the plot and the phase transition takes
the substance across line C, bypassing the liquid phase. If the external pressure is 8 atm, then
heating a block of dry ice would cause it to melt into the liquid state.
The liquid-gas equilibrium curve, line B, terminates at a point known as the critical point, beyond
which there are no distinct liquid and gas phases. Instead, the substance exists in a form known as a
supercritical fluid. On the other hand, the boundary between the solid and liquid phases continues
indefinitely (hence the arrowhead on line A), and for almost all substances leans to the right, which
means that as the pressure increases, a higher and higher temperature is needed to cause melting
(solid to liquid) to occur. This is because high pressure favors the typically denser solid phase over
the liquid one. H 2 O is unique in that its solid form is generally less dense than its liquid form (the
reason ice floats on water). As a result, the phase diagram for H 2 O has a solid-liquid equilibrium
curve that slopes to the left.
THINGS TO REMEMBER
General Properties of Liquids
Crystalline Solids
Amorphous Solids
Liquid Crystals
Dipole-Dipole Interactions
Dispersion Forces
Hydrogen Bonding
Gas-Liquid Equilibrium
Liquid-Solid Equilibrium
Gas-Solid Equilibrium