A phase diagram looks intimidating — three curves splitting a graph into regions, two special dots, and a question asking what happens at some point in the middle. But it is just a map: tell it a pressure and a temperature, and it tells you whether a substance is solid, liquid, or gas. This guide walks the regions, the boundaries, and the two points you must know.
The Axes and the Three Regions
A phase diagram plots pressure on the vertical axis against temperature on the horizontal axis. Every point on the graph is one specific combination of P and T, and the diagram tells you the stable phase there.
The curves divide the plane into three regions:
- Solid — low temperature, generally toward the upper-left.
- Liquid — moderate temperature and higher pressure, the middle.
- Gas — high temperature and low pressure, toward the lower-right.
To use the diagram, find your temperature on the x-axis, your pressure on the y-axis, and see which region that point lands in. That single phase is the answer. The intuition holds: heat pushes a substance toward gas; pressure pushes it toward the denser condensed phases.
The Boundary Lines Are Equilibria
The curves between regions are not borders to avoid — they carry real meaning. A point on a line is a condition where two phases coexist in equilibrium. Three lines, three transitions:
- Solid–liquid line: melting and freezing coexist. This line is nearly vertical, because pressure has little effect on melting temperature.
- Liquid–gas line: boiling and condensing coexist. Follow it and you can read the boiling point at any pressure — it slopes up, which is why water boils below 100 °C at high altitude where pressure is lower.
- Solid–gas line: sublimation and deposition coexist, where a solid turns straight to gas without melting.
Cross a line and you have changed phase. The liquid–gas line is the one with a definite endpoint, which leads to the next section.
The Triple Point and the Critical Point
Two specific points on the diagram get tested constantly.
The triple point is where all three lines meet — the single unique pressure and temperature at which solid, liquid, and gas coexist together at equilibrium. For water it is 0.01 °C and 0.006 atm. There is exactly one such point per substance.
The critical point is where the liquid–gas line stops. Beyond it — above the critical temperature and critical pressure — liquid and gas are no longer distinct phases. The substance becomes a supercritical fluid, with properties of both: it fills its container like a gas but is dense like a liquid. For water the critical point is 374 °C and 218 atm. Past the critical point you cannot boil a liquid, because there is no longer a separate liquid to boil.
Tracing a Path Across the Diagram
The hardest exam question is not "what phase is point X" — it is "describe what happens as conditions change." That is a path across the diagram, and you read it by walking a straight line and noting every boundary you cross.
Heating at constant pressure is a horizontal line moving right. Start in the solid region and move right: you cross the solid–liquid line (the substance melts), continue through the liquid region, cross the liquid–gas line (it boils), and end in the gas region. Each line crossing is a phase change happening at a fixed temperature.
Compressing at constant temperature is a vertical line moving up. Depending on where you start, raising the pressure can turn a gas into a liquid, or a liquid into a solid. If your constant temperature is below the triple-point temperature, a vertical line can take you straight from gas to solid — deposition — without ever passing through liquid.
The trick: a path that does not cross any line stays in one phase, no matter how far the conditions move. Only crossing a boundary changes the phase.
The Water Exception
Most phase diagrams have a solid–liquid line that tilts slightly to the right — higher pressure raises the melting point. Water tilts the other way: its solid–liquid line leans left.
The reason is that ice is less dense than liquid water — it floats. For most substances the solid is denser than the liquid, so squeezing favors the solid. For water, squeezing favors the liquid, so increasing pressure can melt ice. That negative slope is the visual signature of water's phase diagram, and it explains why the line direction is a common exam question.
Getting Help
Phase behavior connects to the gas laws you use for the gaseous region of the diagram — see gas laws: which equation — and the broader General Chemistry study guides cover the related calculations.
Conclusion
Reading a phase diagram is a lookup: pressure on the y-axis, temperature on the x-axis, and the region your point lands in is the phase. The boundary lines mark two-phase equilibria — including the line whose slope gives you boiling point versus pressure. Memorize the triple point (all three phases coexist) and the critical point (where liquid and gas merge into a supercritical fluid), and remember water's solid–liquid line leans left because ice floats.
