Phase changes: melting, boiling, and beyond

The six phase changes

Changing state means rearranging how particles are held together, and there are exactly six ways to do it:

• Melting (fusion): solid to liquid, as ice becomes water.
• Freezing (solidification): liquid to solid, as water becomes ice.
• Vaporization: liquid to gas, which includes both evaporation and boiling.
• Condensation: gas to liquid, as dew forms on cold grass.
• Sublimation: solid to gas with no liquid step, as dry ice becomes carbon dioxide vapor.
• Deposition: gas to solid with no liquid step, as frost forms on a window.

Each change has a reverse partner: melting reverses freezing, vaporization reverses condensation, sublimation reverses deposition. Heating favors the changes that pull particles apart, cooling favors the ones that pack them together. Whether a substance melts or sublimes depends on its own attractions and the surrounding pressure, not on any choice you make.

Why temperature pauses during a change

Here is the fact that surprises most students: while a pure substance changes state, its temperature does not change at all. Ice water stays at $0^\circ\text{C}$ until the last cube melts, and boiling water stays at its boiling point no matter how high you turn the burner.

The reason is energy accounting. Temperature measures the average kinetic energy of particles, but during a phase change the incoming heat is spent on a different job: breaking the attractions between particles. The particles do not speed up, so the thermometer does not move. Only when the change is finished does added heat raise the temperature again.

The energy required has specific names. The heat of fusion melts one gram of solid at its melting point; for water it is 334 J/g. The heat of vaporization boils one gram of liquid; for water it is 2260 J/g, nearly seven times larger, because pulling particles completely apart costs far more than loosening them.

A worked example: melting 50 g of ice at $0^\circ\text{C}$ needs $50 \times 334 = 16{,}700$ J. Warming that same 50 g of resulting water from $0^\circ\text{C}$ to $20^\circ\text{C}$ needs only $50 \times 4.184 \times 20 = 4{,}184$ J. The phase change swallows four times the energy of a 20-degree warm-up, which is why melting ice is such an effective coolant.

Evaporation, boiling, and vapor pressure

Vaporization has two personalities. Evaporation happens at the surface at any temperature: the fastest-moving liquid particles escape into the air one by one. This is why a puddle dries on a cool day. Boiling happens throughout the liquid, but only when the liquid's vapor pressure equals the surrounding pressure, which is why each liquid has a characteristic boiling point at a given pressure.

Vapor pressure is the pressure exerted by a liquid's own vapor in a closed container, and it rises steeply with temperature as more particles gain escape energy. When vapor pressure matches atmospheric pressure, bubbles of vapor can form inside the liquid without being crushed, and boiling begins.

A worked example shows why mountaintop cooking is slow. At sea level the atmosphere presses at 101.3 kPa, and water's vapor pressure reaches 101.3 kPa at $100^\circ\text{C}$. In Denver, about a mile up, atmospheric pressure is only about 83 kPa, and water's vapor pressure reaches 83 kPa at roughly $94^\circ\text{C}$. Pasta cooked in $94^\circ\text{C}$ water cooks more slowly because every chemical reaction in the pot runs cooler.

Sublimation, deposition, and phase diagrams

Sublimation skips the liquid entirely: solid turns directly to gas. Dry ice, solid carbon dioxide, sublimes at $-78.5^\circ\text{C}$ at ordinary pressure, which is why it never leaves a puddle. Mothballs slowly vanish the same way. Deposition is the reverse, gas to solid, and it builds frost on windows and snowflakes in clouds.

Whether a solid melts or sublimes depends on pressure. Carbon dioxide can only be a liquid above about 5 atm; at the 1 atm of everyday life, heating solid CO2 jumps straight to gas. This is visible on a phase diagram, a map with temperature on one axis and pressure on the other, divided into solid, liquid, and gas regions.

Phase diagrams also show the triple point, the single temperature and pressure where all three states coexist, and the critical point, beyond which liquid and gas become indistinguishable as a supercritical fluid. Supercritical carbon dioxide is used to decaffeinate coffee, a neat industrial payoff from reading the map correctly.

Phase changes at work: cooling and sweating

Phase changes are not just textbook curiosities; they are the working principle of the machines that heat and cool buildings. A refrigerator does not create cold. It vaporizes a refrigerant inside the fridge, absorbing heat from the food compartment, then condenses the vapor outside the fridge, dumping that heat into the kitchen. The phase change carries the energy; the compressor just keeps the cycle moving.

Sweating uses the same physics in reverse. Each gram of sweat that evaporates carries away about 2430 J at body temperature. A worked example: evaporating just 100 g of sweat, less than half a cup, removes $100 \times 2430 = 243{,}000$ J. That is roughly the energy needed to warm 1.5 liters of water from room temperature to boiling, pulled straight off your skin. No wonder humid days feel miserable: when the air is already saturated, sweat cannot evaporate and the cooling stalls.

Freeze-drying exploits sublimation under vacuum to preserve food and vaccines: frozen water leaves the product as vapor without ever becoming liquid, so delicate structures survive intact.

Common mistakes

Real-world applications

• Freeze-drying foods and medicines by subliming ice under vacuum

• Designing refrigeration cycles that move heat via vaporization and condensation

• Freeze-drying-free ice prevention: de-icing aircraft wings before takeoff

• Casting metals: controlling melting and freezing for uniform parts