11.2 Phase Diagrams
PRE-HEALTH PREP
Phase Diagram Deep Dive
Typical vs CO₂ and H₂O
General Chemistry 11.2
Phase diagrams map the pressure and temperature conditions under which a substance exists as a solid, liquid, or gas, with key features including the triple point, critical point, and six named phase transitions. Carbon dioxide's triple point lies above one atmosphere, which is why dry ice undergoes sublimation rather than melting under normal conditions. Water's phase diagram is distinguished by a negative-slope solid-liquid equilibrium line—a consequence of hydrogen bonding that causes ice to be less dense than liquid water, enabling ice to float and life to thrive in aquatic environments.
What Is a Phase Diagram?
A Plot of Pressure vs. Temperature
A phase diagram is a plot of pressure versus temperature that shows where the three phases of matter — solid, liquid, and gas — exist for a given substance. Even for unusual compounds like CO₂ and H₂O, the three phases always occupy the same general regions of the diagram. As temperature increases from left to right, a substance progresses from solid to liquid to gas.
Lines of Equilibrium
The curves drawn on a phase diagram are called lines of equilibrium. Each line represents the set of pressure-temperature conditions where two phases coexist in equilibrium. If a point falls clearly within a region, you have a single phase. But if a point falls directly on a line, two phases exist simultaneously.
| PHASE DIAGRAMS | |
|---|---|
| Lines of Equilibrium | Lines at which two phases are in equilibrium |
| Triple Point | The single point at which all three phases are in equilibrium |
| Critical Point | The temperature/pressure beyond which there is no liquid/gas phase change |
| Normal m.p. and b.p. | The melting point and boiling point at 1atm |
Example: Reading a Phase Diagram
A point in the lower-left region represents a solid. A point in the upper-middle region represents a liquid. But a point sitting exactly on the curve between those two regions indicates a solid and a liquid coexisting in equilibrium. Similarly, the curve between liquid and gas regions indicates liquid-gas equilibrium, and the curve between solid and gas regions indicates solid-gas equilibrium.
What Is the Triple Point?
Where all three lines of equilibrium converge, you find one very special location called the triple point. This is the unique combination of pressure and temperature at which all three phases — solid, liquid, and gas — coexist in equilibrium simultaneously. That single defining fact is what you need to know about the triple point.
The 6 Phase Transitions
You should be able to identify all six phase transitions on a phase diagram. Each transition has a specific name depending on the direction of the change:
| From | To | Transition Name |
|---|---|---|
| Solid | Liquid | Fusion (melting) |
| Liquid | Solid | Crystallization (freezing) |
| Liquid | Gas | Vaporization |
| Gas | Liquid | Condensation |
| Solid | Gas | Sublimation |
| Gas | Solid | Deposition |
Common Mistake: Students frequently forget the names for the solid-gas transitions. Sublimation is the direct conversion of a solid to a gas, and deposition is the reverse — gas directly to solid. Don't forget to deposit that in your memory.
Exam questions often present a phase diagram with an arrow drawn between two regions and ask you to identify which phase transition it represents. Make sure you can name all six.
Critical Point and Supercritical Fluids Explained
What Is the Critical Point?
At the far end of the liquid-gas line of equilibrium sits another special location called the critical point. It corresponds to a specific critical pressure and critical temperature, and it is defined as the point beyond which no liquid-gas phase transitions occur.
Phase Transitions Below the Critical Point
Below the critical point, liquid-gas phase transitions are dramatic and unmistakable. Consider what happens when you take a gas and rapidly increase the pressure: the moment you cross the liquid-gas line of equilibrium, the substance condenses into a liquid. The transition is sudden and obvious. The same thing happens if you lower the temperature instead — as soon as you cross that boundary, the gas collapses into a liquid.
Example: The Imploding 55-Gallon Drum
A classic chemistry demonstration illustrates just how sudden these transitions can be. A professor fills a 55-gallon drum with superheated steam, seals it, and leaves it in the corner before class begins. As students file in and the lecture proceeds, the steam gradually cools. The moment the temperature drops below 100°C, the steam undergoes a sudden phase transition, condensing into liquid water. With far less gas pressure pushing outward on the walls — but atmospheric pressure still pushing inward — the entire drum implodes with a thunderous crash.
Going Around the Critical Point
Here is where things get fascinating. Instead of crossing the liquid-gas boundary directly, imagine taking a gas and raising its temperature first, then increasing the pressure, then lowering the temperature, and finally reducing the pressure — essentially traveling around the critical point. You end up at the exact same destination, but you never cross the liquid-gas line of equilibrium. There is no dramatic moment where the substance visibly transforms from gas to liquid.
Why does this work? When a substance is above the critical temperature, its molecules possess so much kinetic energy that no amount of pressure can force them close enough together for intermolecular attractive forces to trigger condensation. It does not matter how high you raise the pressure — the molecules simply have too much energy to condense.
Key Concept: The greater the intermolecular attractive forces present in a substance, the higher both its critical temperature and critical pressure will be. Stronger attractions require more kinetic energy (higher temperature) to overcome, and the molecules can withstand being forced closer together (higher pressure) before condensing.
What Is a Supercritical Fluid?
When a substance exists above both its critical temperature and critical pressure, it enters a state known as a supercritical fluid. In this region, the substance is neither a liquid nor a gas — it occupies a distinct state with properties of both. The transition into the supercritical region is gradual, not abrupt, which is precisely why traveling around the critical point avoids a visible phase change.
What Are Normal Melting and Boiling Points?
Two more points you should be able to identify on a typical phase diagram are the normal melting point and the normal boiling point. Both are defined at a pressure of exactly one atmosphere — the "normal" designation simply means the measurement occurs at standard atmospheric pressure at sea level.
Example: Finding Normal Points on a Phase Diagram
Draw a horizontal line across the diagram at one atmosphere. Where that line intersects the solid-liquid line of equilibrium, you find the normal melting point. Where it intersects the liquid-gas line of equilibrium, you find the normal boiling point. These are the two additional points you may be asked to locate on an exam.
Why Does Dry Ice Sublime? The CO₂ Phase Diagram
What makes carbon dioxide unique is the location of one atmosphere relative to its triple point. For a typical substance, one atmosphere lies above the triple point in terms of pressure. For CO₂, however, one atmosphere lies below the triple point. This single difference has a profound consequence.
When you heat solid carbon dioxide — commonly known as dry ice — at atmospheric pressure, it does not melt. Instead, it sublimes, converting directly from solid to gas. Liquid carbon dioxide simply cannot exist at one atmosphere; it only exists at pressures above the triple point.
Why "Dry" Ice? Dry ice earns its name because it transitions straight from solid to gas at atmospheric pressure, skipping the liquid phase entirely. It never produces a liquid, so it never feels wet. That is the key fact about carbon dioxide: it sublimes rather than melts at standard atmospheric pressure.
Why Does Ice Float? Water's Phase Diagram
What Makes Water's Phase Diagram Unique?
Water's one-atmosphere line sits above the triple point, just like a typical substance — so that is not what makes it unusual. The distinguishing feature of water's phase diagram is the solid-liquid line of equilibrium. For virtually every other substance, this line has a positive slope (tilting to the right). For water, it has a negative slope (tilting to the left). This seemingly small difference has enormous consequences.
Why Ice Is Less Dense Than Liquid Water
The negative slope exists because water does something unusual when it freezes: it expands. Most substances contract and become denser when they solidify, but water behaves in the opposite way. The reason traces back to hydrogen bonding. When water molecules arrange themselves into the crystal lattice of ice, they space themselves out to maximize hydrogen bonding interactions. The result is a structure that occupies more volume than the same mass of liquid water.
Key Fact: Ice is less dense than liquid water because of hydrogen bonding. This is why ice floats — and why water's solid-liquid equilibrium line has a negative slope on the phase diagram.
Why This Matters for Life on Earth
Without this negative slope, you probably would not be alive today. Because ice is less dense than liquid water, it floats to the surface when bodies of water freeze. Lakes, rivers, and oceans freeze from the top down, not the bottom up. That layer of surface ice acts as insulation, protecting the liquid water underneath from the frigid air above. In many cases, the water never freezes all the way through.
If ice were denser than liquid water — as frozen solids are for most substances — bodies of water would freeze from the bottom up, potentially freezing solid every winter. Most aquatic life would perish, and the cycle of life as we know it could never have developed. The fact that ice floats is, quite literally, essential to life on Earth.
Ice Skating and Pressure Melting
The negative slope also explains something more playful: why ice skating works so well. Recall that pressure equals force divided by area. An ice skater's weight is concentrated over the extremely narrow blade of the skate, generating very high pressure directly beneath it. On water's phase diagram, increasing pressure on ice at constant temperature pushes you across the solid-liquid line of equilibrium — and the ice melts into a thin film of liquid water right under the blade.
A thin layer of water on ice reduces the coefficient of friction dramatically, allowing the skater to glide with remarkable speed. Slightly wet ice is far more slippery than dry ice (the frozen-water kind, not the CO₂ kind).
Example: What If a Hockey Rink Were Made of Dry Ice?
Imagine replacing the ice in a hockey rink with solid CO₂. A player steps onto the surface, and the skate blade generates the same high pressure beneath it. But CO₂ has a normal, positive-slope solid-liquid line — increasing pressure on a solid just keeps it solid (or pushes it further into the solid region). The dry ice would never melt under the blade, the surface would never become slippery, and the players would struggle to skate at anything close to their normal speed.
7 Key Takeaways
- A phase diagram plots pressure vs. temperature and shows the solid, liquid, and gas regions for a substance, separated by lines of equilibrium where two phases coexist.
- The triple point is where all three phases coexist in equilibrium; the critical point marks the end of the liquid-gas boundary, beyond which no distinct liquid-gas transition occurs.
- Above the critical point, a substance exists as a supercritical fluid — neither liquid nor gas — and traveling around the critical point avoids any abrupt phase change.
- There are 6 phase transitions: fusion, crystallization, vaporization, condensation, sublimation, and deposition. The normal melting point and normal boiling point are defined at one atmosphere.
- CO₂ is unique because one atmosphere lies below its triple point, so dry ice sublimes rather than melts at atmospheric pressure.
- H₂O is unique because its solid-liquid equilibrium line has a negative slope, a consequence of ice being less dense than liquid water due to hydrogen bonding.
- Water's negative slope means ice floats, lakes freeze from the top down (insulating aquatic life), and pressure from an ice skate can melt ice — enabling ice skating.