A phase diagram is used to show how the phases and phase changes of a substance vary with both temperature and pressure. FIGURE 18.6 shows the phase diagrams for water and carbon dioxide. You can see that each diagram is divided into three regions corresponding to the solid, liquid, and gas phases. The boundary lines separating the regions indicate the phase transitions. The system is in phase equilibrium at a pressure-temperature point that falls on one of these lines.

FIGURE 18.6 Phase diagrams (not to scale) for water and carbon dioxide.

Phase diagrams contain a great deal of information. Notice on the water phase diagram that the dashed line at $p=1\text{atm}$ crosses the solid-liquid boundary at $0\text{°}\text{C}$ and the liquid-gas boundary at $100\text{°}\text{C}$. These well-known melting and boiling point temperatures of water apply only at standard atmospheric pressure. You can see that in Denver, where $p_{\text{atmos}}<1\text{atm}$, water melts at slightly above $0\text{°}\text{C}$ and boils at a temperature below $100\text{°}\text{C}$. A pressure cooker works by allowing the pressure inside to exceed 1 atm. This raises the boiling point, so foods that are in boiling water are at a temperature above $100\text{°}\text{C}$ and cook faster.

Crossing the solid-liquid boundary corresponds to melting or freezing while crossing the liquid-gas boundary corresponds to boiling or condensing. But there’s another possibility—crossing the solid-gas boundary. The phase change in which a solid becomes a gas is called sublimation. It is not an everyday experience with water, but you probably are familiar with the sublimation of dry ice. Dry ice is solid carbon dioxide. You can see on the carbon dioxide phase diagram that the dashed line at $p=1\text{atm}$ crosses the solid-gas boundary, rather than the solid-liquid boundary, at $T=-78\text{°}\text{C}$. This is the sublimation temperature of dry ice.

Liquid carbon dioxide does exist, but only at pressures greater than 5 atm and temperatures greater than $-56\text{°}\text{C}$. A $\text{CO}{}_2$ fire extinguisher contains liquid carbon dioxide under high pressure. (You can hear the liquid slosh if you shake a $\text{CO}{}_2$ fire extinguisher.)

One important difference between the water and carbon dioxide phase diagrams is the slope of the solid-liquid boundary. For most substances, the solid phase is denser than the liquid phase and the liquid is denser than the gas. Pressurizing the substance compresses it and increases the density. If you start compressing $\text{CO}{}_2$ gas at room temperature, thus moving upward through the phase diagram along a vertical line, you’ll first condense it to a liquid and eventually, if you keep compressing, change it into a solid.

Water is a very unusual substance in that the density of ice is less than the density of liquid water. That is why ice floats. If you compress ice, making it denser, you eventually cause a phase transition in which the ice turns to liquid water! Consequently, the solid-liquid boundary for water slopes to the left.

The liquid-gas boundary ends at a point called the critical point. Below the critical point, liquid and gas are clearly distinct and there is a phase change if you go from one to the other. But there is no clear distinction between liquid and gas at pressures or temperatures above the critical point. The system is a fluid, but it can be varied continuously between high density and low density without a phase change.

The final point of interest on the phase diagram is the triple point where the phase boundaries meet. Two phases are in phase equilibrium along the boundaries. The triple point is the one value of temperature and pressure for which all three phases can coexist in phase equilibrium. That is, any amounts of solid, liquid, and gas can happily coexist at the triple point. For water, the triple point occurs at $T_3=0.01\text{°}\text{C}$ and $p_3=0.006\text{atm}$.