Chemistry 2e · Liquids and Solids
Phase Diagrams
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A phase diagram Map of which phase (solid/liquid/gas) is stable at each temperature and pressure Full entry → is a map of a pure substance's states: a graph with temperature on the horizontal axis and pressure on the vertical axis, divided into regions where the substance is solid, liquid, or gas. Every point on the diagram answers one question: "At this temperature and pressure, which phase is stable?" The lines between regions show where two phases coexist in equilibrium — exactly the melting, freezing, boiling, and sublimation conditions studied in the previous topic.
This topic explains how to read a phase diagram, highlights the two special points on every diagram — the triple point Unique (T, P) where solid, liquid, and gas coexist Full entry → and the critical point (T, P) where the liquid–gas boundary ends Full entry → — and compares the phase diagrams of water and carbon dioxide, whose differences explain everyday phenomena like dry ice sublimation and ice melting under pressure.
Why this matters
Phase diagrams are the operating manuals for industrial chemistry. Food processing uses supercritical CO₂ for decaffeination; power plants run steam at pressures above water's critical point; the pharmaceutical industry freeze-dries drugs by controlling pressure and temperature; geologists explain metamorphic rocks with pressure–temperature maps. In everyday life, water's phase diagram explains why cooking at altitude takes longer, why pressure cookers work, and how a skate blade can glide on ice. Knowing how to read one diagram lets you reason about all of these.
The college version
Core Concepts
The anatomy of a phase diagram
A typical phase diagram has:
- Axes: temperature (K or °C) on the horizontal axis, pressure (atm, kPa, or MPa) on the vertical axis. Both axes are often nonlinear (log pressure is common) to fit huge ranges.
- Three regions: solid (low T, high P), gas (high T, low P), liquid (middle) — one region per stable phase.
- Three coexistence lines separating the regions, where two phases exist together:
- Sublimation/deposition curve: solid–gas boundary at low pressure.
- vapor pressure curve Liquid–gas coexistence line on a phase diagram Full entry → (vaporization/condensation): liquid–gas boundary; it ends at the critical point.
- Melting/freezing curve: solid–liquid boundary; it begins at the triple point and extends upward, often nearly vertical.
- Triple point: the single (T, P) combination where all three phases coexist.
- Critical point: the end of the liquid–gas line, beyond which liquid and gas become indistinguishable.
Reading the map
To identify the phase at a given (T, P): locate the point, find which region it falls in, and read the phase. To predict a transition: move along a path on the diagram and note every boundary crossed. At a boundary, the substance is at its melting or boiling point for that pressure; on a line, the two adjacent phases are in equilibrium and can interconvert with no temperature change.
The triple point
The triple point is the unique temperature and pressure where solid, liquid, and gas coexist. For water it is 0.01 °C (273.16 K) at 611.7 Pa (0.006 atm); for CO₂ it is −56.6 °C at 5.11 atm. Below the triple-point pressure, the liquid phase cannot exist: a solid heated at lower pressure goes straight to gas (sublimation) — which is exactly why dry ice at 1 atm sublimes rather than melts.
The critical point and supercritical fluids
The critical point is the highest temperature at which a distinct liquid phase can exist (the critical temperature Tc) at the corresponding critical pressure (Pc). Above Tc and Pc, the substance is a supercritical fluid Substance above Tc and Pc; liquid-like density, gas-like flow Full entry →: it has gas-like ability to fill a container and diffuse, yet liquid-like density and solvent power. For CO₂, Tc = 31.0 °C, Pc = 72.8 atm; for water, Tc = 374 °C, Pc = 217.7 atm. Supercritical CO₂ is a widely used "green" solvent — it decaffeinates coffee and extracts flavors because its solvent strength can be tuned by changing pressure, and it leaves no solvent residue.
Water's unusual phase diagram
Most substances freeze to a denser solid, so their solid–liquid boundary slopes to the right (higher pressure → higher melting point). Water is different: its solid–liquid line slopes left — the melting point of ice decreases slightly as pressure rises. The reason is that ice is less dense than liquid water (its open hydrogen-bonded lattice), so increasing pressure favors the denser phase, liquid water. Consequences: ice floats, lakes freeze from the top down, and pressure can locally melt ice under a skate blade.
How It Works / Step-by-Step Process
To use a phase diagram:
- Identify the substance's diagram (each substance has its own).
- Locate the (T, P) point of interest on the axes.
- Determine which region the point falls in — that is the stable phase.
- To find a transition temperature: find where a vertical (constant-P) or horizontal (constant-T) path crosses a boundary line.
- Read the normal melting/boiling points where the 1-atm line crosses the boundaries.
- For sublimation questions: check whether the pressure is below the triple-point pressure.
Common Confusions
| Common Confusion | Correct Understanding |
|---|---|
| "The lines on a phase diagram show where the substance changes." | They show where two phases coexist in equilibrium; a change occurs only when you cross a line. |
| "Liquids can exist at any pressure." | Below the triple-point pressure, the liquid phase cannot exist — only solid and gas. |
| "Boiling point is always 100 °C for water." | That is the normal boiling point at exactly 1 atm; at other pressures the boiling point follows the vapor pressure curve. |
| "Supercritical fluid is just very hot gas." | It has liquid-like density and solvent power with gas-like flow — a distinct state above Tc and Pc. |
| "Ice skating works because pressure melts ice dramatically." | Pressure does lower ice's melting point (left-sloping line), but the effect is small; skating depends mostly on friction-generated heat and surface films. |
| "All substances have the same phase diagram shape." | Same layout, but different numbers — and water's left-sloping solid–liquid line is unusual (most substances slope right). |
| "The triple point is where ice, water, and steam meet at 0 °C." | It is at 0.01 °C and 0.006 atm — the exact 0 °C/1 atm point lies on the melting line, where only ice and water coexist. |
Eli explains
The same idea, in plain words
Explain it like I’m 10
A phase diagram is like a weather map for a chemical. The map shows where the chemical is ice, water, or steam depending on how cold it is and how hard it's being squeezed. Two special places exist: the triple point, where all three show up at once, and the critical point, past which water and steam become one weird "super" fluid that can't tell them apart.
Worked example
Example 1: Why does dry ice sublime instead of melting?
Consider CO₂ at room pressure, 1 atm. Its triple point is −56.6 °C at 5.11 atm — the triple-point pressure is higher than 1 atm.
At 1 atm, follow the temperature axis: the solid region exists only below about −78.5 °C. The sublimation curve separates solid and gas, and the liquid region never appears at this pressure because the triple point sits above it. So solid CO₂ heated at 1 atm crosses the solid–gas boundary directly — it sublimes at −78.5 °C. To see liquid CO₂, you must compress the gas above 5.11 atm first. This is why dry ice "smokes" (sublimes to cold CO₂ gas) at room conditions instead of puddling into liquid.
Example 2: Estimating the boiling point of water at high altitude
Atmospheric pressure at 3000 m is about 0.70 atm. Use the Clausius–Clapeyron equation to estimate the boiling point of water there. Known: normal boiling point T1 = 373 K at P1 = 1.00 atm, ΔHvap = 40.65 kJ/mol = 40,650 J/mol, R = 8.314 J/(mol · K).
Start with the equation:
ln(P2P1) = -ΔHvapR(1T2 - 1T1)
Substitute P2 = 0.70 atm:
ln(0.70) = -40,6508.314(1T2 - 1373)
-0.357 = -(4889 K)(1T2 - 0.002681 K-1)
1T2 = 0.3574889 + 0.002681 = 0.0000730 + 0.002681 = 0.002754 K-1
T2 = 10.002754 = 363 K ≈ 90 °C
Water boils at about 90 °C at 3000 m — 10 degrees cooler than at sea level. This is why high-altitude recipes need longer cooking times: food simmers at a lower temperature.
Key takeaways
- Phase diagram axes: temperature (x) and pressure (y); regions are solid, liquid, gas.
- Coexistence lines = phase boundaries; crossing one means a transition.
- Triple point: all three phases coexist — a unique (T, P) pair.
- Critical point: end of the liquid–gas line; above it, supercritical fluid.
- Below the triple-point pressure, no liquid exists (sublimation only).
- Water's solid–liquid line slopes left (ice less dense than water); most substances slope right.
- Water: triple point 0.01 °C / 0.006 atm; critical point 374 °C / 217.7 atm. CO₂: triple point −56.6 °C / 5.11 atm; critical point 31.0 °C / 72.8 atm.
- Supercritical CO₂: tunable green solvent (decaffeination, extraction).
Check yourself
5 review questions from the chapter. Try each one, then open the answer.
What are the two special points on every phase diagram, and what happens at each?
Show answer
The triple point — unique (T, P) where solid, liquid, and gas coexist. The critical point — end of the liquid–gas line, above which only a supercritical fluid exists.
Why can liquid CO₂ never be seen at 1 atm?
Show answer
Because CO₂'s triple-point pressure is 5.11 atm, above 1 atm. At pressures below 5.11 atm the liquid region does not exist, so solid CO₂ sublimes directly to gas.
Describe the path of solid CO₂ at 1 atm heated from −100 °C to 20 °C — which phase changes occur?
Show answer
Solid from −100 °C to −78.5 °C, then sublimation to gas at −78.5 °C; gas from −78.5 °C upward. No liquid appears.
What is a supercritical fluid, and one industrial use of supercritical CO₂?
Show answer
A state above Tc and Pc with liquid-like density and gas-like flow. Supercritical CO₂ is used to decaffeinate coffee and extract flavors.
How does water's solid–liquid line differ from most substances', and what physical property of ice causes it?
Show answer
Water's solid–liquid line slopes left (melting point decreases with pressure) because ice is less dense than liquid water — the open hydrogen-bonded lattice makes the solid the low-density phase, so pressure favors liquid.
Study tools & related lessonsKey vocabulary · Related
Key vocabulary
- phase diagram
- Map of which phase (solid/liquid/gas) is stable at each temperature and pressure
- coexistence line
- A boundary where two phases are in equilibrium
- triple point
- Unique (T, P) where solid, liquid, and gas coexist
- critical point
- (T, P) where the liquid–gas boundary ends
- critical temperature, Tc
- Highest temperature at which a liquid phase can exist
- supercritical fluid
- Substance above Tc and Pc; liquid-like density, gas-like flow
- vapor pressure curve
- Liquid–gas coexistence line on a phase diagram
- normal melting/boiling point
- Transition temperature at exactly 1 atm
Sources & references
This lesson was adapted from the open educational references above; their licenses and attributions are preserved. See Copyright & Licensing.
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