Vapor Pressure and Boiling gave one curve: where liquid and vapor can coexist. A phase diagram adds the curves for solid–liquid and solid–gas, so you can read the stable phase at any temperature and pressure. It is the map behind the Heating Curves and Phase Changes: a heating curve is a walk across it at constant pressure.
Key idea
Each region is one phase; the phase at a point is the region the point is in. Each line is a set of conditions where two phases coexist. At the triple point all three coexist. At the critical point the liquid–gas line ends: beyond it, at higher temperature and pressure, liquid and gas can no longer be told apart (a supercritical fluid).
Why it works
Higher pressure squeezes molecules together, favoring denser phases; higher temperature gives them energy to break free, favoring the gas. A line is where two phases are equally stable, so a sample on it can be part one and part the other. Above the critical temperature, molecules move too fast for any pressure to hold them as a separate liquid with a surface. Water has a rare feature: its solid–liquid line leans slightly to the left, because ice is less dense than liquid water, so squeezing ice can melt it. For carbon dioxide, the triple point is at 5.11 atm, above normal air pressure, so at 1 atm solid $\mathrm{CO_2}$ (dry ice) turns straight to gas.
Worked example: Read a point
What is the phase of water at $50^\circ\text{C}$ and $0.050\ \text{atm}$? Water's vapor pressure at $50^\circ\text{C}$ is $0.121\ \text{atm}$.
- At $50^\circ\text{C}$ the liquid–gas line sits at 0.121 atm.
- Our point, 0.050 atm, is below the line: lower pressure than the vapor pressure.
- Below the liquid–gas line lies the gas region.
Answer: Gas (water vapor).
Worked example: Follow a path
Dry ice at $-100^\circ\text{C}$ and 1 atm is warmed to $25^\circ\text{C}$ at constant pressure. What happens?
- 1 atm is below $\mathrm{CO_2}$'s triple-point pressure (5.11 atm), so the horizontal path never crosses the liquid region.
- It crosses only the solid–gas line, at $-78.5^\circ\text{C}$.
Answer: The solid sublimes to gas at $-78.5^\circ\text{C}$, then the gas warms.
Worked example: Squeeze a gas
Carbon dioxide gas at $20^\circ\text{C}$ is compressed slowly. Its vapor pressure at $20^\circ\text{C}$ is 56.5 atm. What happens as the pressure rises past 56.5 atm? What if the temperature were $40^\circ\text{C}$?
- At $20^\circ\text{C}$ the vertical path crosses the liquid–gas line at 56.5 atm: the gas condenses to liquid.
- At $40^\circ\text{C}$, above the critical temperature ($31.0^\circ\text{C}$), there is no liquid–gas line to cross: the gas just grows steadily denser, and above 72.8 atm it is a supercritical fluid. It never condenses.
Answer: It liquefies at 56.5 atm; at $40^\circ\text{C}$ it never does.
Common mistake
Reading the phase from the nearest line instead of the region the point is in, or mixing up the two special points. A point near the liquid–gas line can be firmly in the liquid region. The triple point is where three lines meet (all three phases); the critical point is where one line simply stops. Check: find which side of each line the point lies on.