Chemistry: Atoms First 2e · Liquids and Solids
Phase Diagrams
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A phase diagram Map of stable phase (solid/liquid/gas) at every temperature and pressure Full entry → is a map of a pure substance: temperature on the horizontal axis, pressure on the vertical axis, with regions labeled solid, liquid, and gas. Every point answers one question — "at this temperature and pressure, which phase is stable?" — and the lines between regions show where two phases coexist in equilibrium. Two special points complete the map: the triple point Unique (T, P) where solid, liquid, and gas coexist, where all three phases meet, and the critical point End of the liquid–gas line at Tc, Pc Full entry →, where the liquid–gas boundary ends.
The diagram summarizes the previous three topics: the sublimation, vapor-pressure, and melting curves are the solid–gas, liquid–gas, and solid–liquid lines. This topic teaches you to read the map, explains the triple and critical points, and compares water and CO₂ — explaining dry ice, ice skating, and pressure cooking.
Why this matters
Phase diagrams are operating manuals for real processes: decaffeinating coffee with supercritical CO₂, supercritical steam in power plants, freeze-drying drugs by subliming ice, reading rock history from pressure–temperature maps. Water's diagram explains high-altitude cooking, pressure cookers, and ice skating — and why liquid CO₂ can't exist in a fire extinguisher at room pressure.
The college version
Core Concepts
Reading the map
- Axes: temperature (K or °C) horizontal, pressure (atm, kPa, or MPa) vertical (often logarithmic).
- Three regions: solid (low T, high P), liquid (middle), gas (high T, low P).
- Three coexistence lines, where two phases are in equilibrium:
- Sublimation/deposition curve (solid–gas) at low pressure.
- vapor pressure curve Liquid–gas coexistence line Full entry → (liquid–gas), ending at the critical point.
- Melting/freezing curve (solid–liquid), starting at the triple point.
To find the stable phase at any (T, P), locate the point and read the region; to predict heating, follow a path and note every boundary crossed.
The triple point
The triple point is the unique temperature and pressure where solid, liquid, and gas coexist: 0.01 °C (273.16 K) at 611.7 Pa for water; −56.6 °C at 5.11 atm for CO₂. Its pressure sets a hard rule: below it, the liquid phase cannot exist — a solid heated at lower pressure goes straight to gas. That is why dry ice at 1 atm sublimes instead of melting.
The critical point and supercritical fluids
The critical point marks the end of the liquid–gas line: the highest temperature (Tc) and pressure (Pc) at which a distinct liquid can exist. Above both, the substance is a supercritical fluid Substance above Tc and Pc; liquid-like density, gas-like flow — gas-like in filling a container, liquid-like in density and dissolving power. For CO₂: Tc = 31.0 °C, Pc = 72.8 atm; for water: Tc = 374 °C, Pc = 217.7 atm. Supercritical CO₂ is a tunable "green" solvent that leaves no residue.
Water versus carbon dioxide
Water's solid–liquid line slopes left: pressure lowers ice's melting point because ice is less dense than liquid water and pressure favors the denser phase. Consequences: ice floats, lakes freeze top-down, and pressure can locally melt ice under a skate blade. CO₂'s line slopes right, like most substances; same layout, different numbers.
Normal melting and boiling points
Read the normal melting point and normal boiling point Boiling temperature at exactly 1 atm where the 1-atm line crosses the solid–liquid and liquid–gas boundaries — for water, 0 °C and 100 °C. At altitude the 1-atm line moves down and the liquid–gas crossing slides lower — high-altitude cooking in picture form.
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.
- Read the region the point falls in — that is the stable phase.
- To predict a transition, follow a constant-pressure or constant-temperature path, noting every boundary crossed.
- Read normal melting/boiling points where the 1-atm line crosses the boundaries.
- For sublimation, compare pressure with the triple-point pressure: below it, no liquid.
Common Confusions
| Common Confusion | Correct Understanding |
|---|---|
| "Lines show where the substance changes." | Lines show where two phases coexist; crossing one changes the phase. |
| "Liquid can exist at any pressure." | Below the triple-point pressure, only solid and gas exist. |
| "Water boils at 100 °C everywhere." | That is the normal boiling point at 1 atm; elsewhere it follows the vapor pressure curve. |
| "A supercritical fluid is just very hot gas." | It has liquid-like density and dissolving power with gas-like flow — a distinct state above Tc and Pc. |
| "Ice skating works because pressure melts ice dramatically." | Pressure lowers the melting point, but the effect is small; skating depends mostly on friction-generated heat and a surface film. |
| "All substances have identical phase diagrams." | Same layout, different numbers — water's left-sloping solid–liquid line is unusual (most slope right). |
| "The triple point of water is 0 °C at 1 atm." | It is 0.01 °C at 0.006 atm; the everyday 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 a weather map for a chemical: it tells you whether the chemical is ice, water, or steam based on how cold it is and how hard it's squeezed. Two special places exist: the triple point, where all three show up at once, and the critical point, past which water and steam merge into a "super" fluid.
Worked example
Example 1: Reading water's phase diagram at four points
Identify water's phase at each condition:
| Condition | Reading the map | Phase |
|---|---|---|
| 1 atm, 25 °C | Between the 0 °C and 100 °C crossings | Liquid |
| 1 atm, 120 °C | Right of the 100 °C crossing | Gas (steam) |
| 0.006 atm, 0.01 °C | Exactly at the triple point | All three coexist |
| 250 atm, 400 °C | Above Tc and Pc | Supercritical fluid |
The last row is the key: at 400 °C no liquid can exist at any pressure, and 250 atm exceeds Pc, so the substance is supercritical. The 1 atm, 120 °C row is why a pot of water at sea level cannot exceed 100 °C until it has boiled away.
Example 2: Why dry ice sublimes — and an estimate of its sublimation temperature
CO₂'s triple point is −56.6 °C at 5.11 atm. At 1 atm — below the triple-point pressure — no liquid region exists, so solid CO₂ goes directly to gas. Estimate its sublimation temperature at 1 atm with the Clausius–Clapeyron equation, using the triple point as the known (P, T) pair and ΔHsub = 25,200 J/mol:
ln(P2P1) = -ΔHsubR(1T2 - 1T1)
Substitute P1 = 5.11 atm, T1 = 216.6 K, P2 = 1.00 atm:
ln(1.005.11) = -25,2008.314(1T2 - 1216.6)
-1.631 = -(3031 K)(1T2 - 0.004617 K-1)
1T2 = 0.000538 + 0.004617 = 0.005155 K-1 ⇒ T2 = 194 K ≈ -79 °C
The estimate (−79 °C) matches the accepted −78.5 °C at 1 atm and explains dry ice's "smoke": solid CO₂ converts directly to cold gas, and water vapor condenses in it into fog.
Key takeaways
- Axes: temperature (x), pressure (y); regions = solid, liquid, gas; lines = two-phase equilibrium.
- Triple point: unique (T, P) where all three phases coexist; below its pressure, no liquid.
- Critical point: end of the liquid–gas line; above Tc and Pc: supercritical fluid.
- Supercritical CO₂: Tc = 31.0 °C, Pc = 72.8 atm; water: Tc = 374 °C, Pc = 217.7 atm.
- Water's solid–liquid line slopes left (ice less dense than water); most substances slope right.
- Normal melting/boiling points = crossings with the 1-atm line.
- Dry ice sublimes at 1 atm because CO₂'s triple-point pressure (5.11 atm) exceeds 1 atm.
- Each substance has its own diagram — same layout, different numbers.
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
Triple point — unique (T, P) where all three phases coexist. Critical point — end of the liquid–gas line, above which only a supercritical fluid exists.
Why can liquid CO₂ never exist at 1 atm?
Show answer
Because CO₂'s triple-point pressure (5.11 atm) is above 1 atm: below it no liquid region exists, so solid CO₂ sublimes directly to gas.
Describe the path of solid CO₂ at 1 atm heated from −100 °C to 20 °C.
Show answer
Solid to about −78.5 °C, sublimation to gas, then gas upward. No liquid appears.
What is a supercritical fluid? Give one industrial use of supercritical CO₂.
Show answer
A state above Tc and Pc with liquid-like density and gas-like flow; supercritical CO₂ decaffeinates coffee.
How does water's solid–liquid line differ from most substances'? What property of ice causes it?
Show answer
Water's solid–liquid line slopes left (melting point falls with pressure) because ice is less dense than liquid water: the open hydrogen-bonded lattice makes the solid low-density, so pressure favors liquid.
Study tools & related lessonsKey vocabulary · Related
Key vocabulary
- phase diagram
- Map of stable phase (solid/liquid/gas) at every temperature and pressure
- coexistence line
- Boundary where two phases are in equilibrium
- triple point
- Unique (T, P) where solid, liquid, and gas coexist
- critical point
- End of the liquid–gas line at Tc, Pc
- supercritical fluid
- Substance above Tc and Pc; liquid-like density, gas-like flow
- sublimation curve
- Solid–gas coexistence line at low pressure
- vapor pressure curve
- Liquid–gas coexistence line
- normal boiling point
- Boiling 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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