Implicit in the discussion above is the fact that the phase in which a substance finds itself is a function of external conditions. On a plot of pressure versus temperature, one can imagine dividing the area of the quadrant into three sections, one corresponding to each of the three phases. The x and y values falling within each section are all the combinations of the pressure and temperature values at which the substance will be in that phase. In general, the gas phase is found at high temperature and low pressure; the solid phase is found at low temperature and high pressure; and the liquid phase is found at high temperature and high pressure. A typical phase diagram is shown below:
The three phases are demarcated by lines indicating the temperatures and pressures at which two phases are in equilibrium. Along line A, the solid and liquid phases are in equilibrium; along line B, 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. H2O 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 H2O has a solid-liquid equilibrium curve that slopes to the left.