States of Matter

Put a thermometer in a beaker of ice and water and heat it steadily. The temperature climbs to 0 °C and then stops. It stays at 0 °C for a long time, while the burner keeps supplying energy at the same rate as before, and only starts rising again once the last of the ice has gone.

Where is that energy going? It is the question heating-curve problems are really testing, and the answer explains most of what states of matter are about.

The thermometer stops moving while the ice melts

Temperature measures the average kinetic energy of particles — how fast they are moving or vibrating. Energy supplied to a substance can go into two places: making the particles move faster, which raises the temperature, or pulling them further apart against the attractions holding them together, which does not.

During a phase change, all of it goes into the second. The energy is spent dismantling the arrangement rather than speeding anything up, so the thermometer reads a constant value until the job is finished. That energy is stored as potential energy and is returned when the substance changes back — which is exactly why freezing releases heat.

The two quantities involved have names. Enthalpy of fusion is the energy per gram needed to melt a substance at its melting point; enthalpy of vaporisation is the energy per gram needed to boil it at its boiling point. Neither involves any temperature change.

Heating 100 g of ice all the way to steam

The standard problem strings five stages together, and the numbers for water are worth seeing because of how lopsided they are.

Water's specific heat capacities are 2.09 J/g°C for ice, 4.18 for liquid water and 2.03 for steam. Its enthalpy of fusion is 334 J/g and its enthalpy of vaporisation 2,260 J/g.

Taking 100 g of ice from −20 °C to steam at 120 °C:

About 309 kJ, and boiling accounts for 226 kJ of it — 73% of the whole process in a single stage. Turning water at 100 °C into steam at 100 °C takes over five times as much energy as heating that water from freezing to boiling.

Two everyday facts fall straight out of that figure. A steam burn is far worse than a boiling-water burn of the same mass and temperature, because the steam releases all 226 J per gram back into your skin as it condenses. And sweating cools you efficiently for the same reason in reverse — each gram that evaporates takes a large amount of energy away with it.

Vaporisation costs so much more than fusion because melting only loosens a structure while keeping particles in contact, whereas boiling separates them completely.

Six changes, and the one everyone forgets

There are six transitions between the three familiar states, and they are properly paired:

• Melting (fusion) and freezing, between solid and liquid.
• Vaporisation and condensation, between liquid and gas.
• Sublimation and deposition, between solid and gas directly.

Deposition is the one that gets left out of lists, and it is not rare. Frost on a cold window is not frozen dew — it forms when water vapour turns directly to ice without ever being liquid, which is why frost crystals have that feathery branching shape rather than looking like frozen droplets.

Sublimation is familiar from solid carbon dioxide, and also explains why ice cubes left in a freezer for months slowly shrink, and why snow can disappear from a mountainside on a dry sunny day without ever visibly melting.

Reading a phase diagram

A phase diagram plots pressure against temperature and divides the plane into regions labelled solid, liquid and gas. The lines between regions are the conditions at which two states coexist — so the line between liquid and gas is the boiling point at each pressure, not a single number.

Two points on the diagram have names.

The triple point is where all three lines meet: the single combination of temperature and pressure at which solid, liquid and gas coexist in equilibrium. For water it is 0.01 °C and 611.7 Pa, a little under a hundredth of atmospheric pressure. It is so reproducible that it was used to define the kelvin until 2019.

The critical point is where the liquid–gas line simply stops. For water it is at 374 °C and about 218 atmospheres. Above that point the two states become indistinguishable — a supercritical fluid that fills its container like a gas but dissolves things like a liquid. Supercritical carbon dioxide is used industrially to decaffeinate coffee and to clean electronic components, because it behaves as a solvent and then vanishes as a gas when the pressure drops.

Why there is no liquid carbon dioxide at ordinary pressure

Carbon dioxide's triple point sits at −56.6 °C and 5.11 atmospheres, and that single fact explains dry ice completely.

The triple point is the lowest pressure at which the liquid can exist at all. Since 5.11 atm is well above the 1 atm of ordinary air, there is no temperature whatsoever at which carbon dioxide is a liquid under normal atmospheric pressure. Warm the solid and it goes straight to gas, subliming at −78.5 °C.

Water's triple point is far below atmospheric pressure, which is why water has a comfortable liquid range at sea level and carbon dioxide does not. Nothing about the molecules is special here — the difference is entirely in where the triple point falls relative to the pressure you happen to be standing in.

Water's backwards line

On almost every phase diagram, the boundary between solid and liquid leans slightly to the right: raising the pressure raises the melting point, because squeezing a substance favours its denser state, and for nearly everything the solid is denser than the liquid.

Water's line leans left. Ice is less dense than liquid water, so pressure favours the liquid, and increasing the pressure on ice lowers its melting point.

The consequence you can see is that ice floats, which is unusual enough that it shapes the whole planet: lakes freeze from the top down, insulating the water below, and the aquatic life in them survives the winter. Had water behaved like a normal substance, lakes would freeze solid from the bottom.

Boiling is not about bubbles

Two definitions get muddled, and the muddle produces wrong answers about altitude.

Evaporation happens at the surface, at any temperature. Some molecules in a liquid always have more than average energy, and those near the surface can escape. This is why a puddle dries without ever reaching 100 °C.

Boiling happens when the liquid's vapour pressure equals the surrounding pressure, at which point vapour bubbles can form and survive inside the liquid rather than being crushed. That is the actual definition, and it is a statement about pressure, not about bubbles.

So a boiling point is not a fixed property of a substance. Reduce the surrounding pressure and the boiling point falls with it: on the summit of Everest water boils at around 70 °C, which is too cool to cook food properly no matter how long it sits there.

One more piece of vocabulary. A vapour is the gaseous form of something that would be liquid or solid at that temperature. Water above 100 °C is a gas; water molecules in the air on a warm day are a vapour.

Beyond three: plasma, condensates, and no fixed count

Solid, liquid and gas are not an exhaustive list, and there is no agreed number of states of matter.

Plasma is a gas whose atoms have been ionised, so it contains free electrons and conducts electricity and responds to magnetic fields. It is not a rarity — stars are plasma, and by mass most of the ordinary matter in the visible universe is in this state.

A Bose–Einstein condensate forms when certain atoms are cooled to within billionths of a degree of absolute zero, at which point they collapse into a single shared quantum state and behave as one entity. Predicted in the 1920s, it was first produced in 1995 with rubidium atoms at 170 nanokelvin, work recognised with the 2001 Nobel Prize in Physics.

Beyond those sit superfluids, which flow without viscosity; liquid crystals, ordered in some directions and not others, which run every flat display panel; and degenerate matter inside white dwarf stars. Which of these count as separate states is a matter of where you draw lines, and the honest answer to "how many states of matter are there" is that the question has no fixed answer.

What decides which state a familiar substance is in comes down to how strongly its particles attract one another relative to the thermal energy available — a comparison set out in more detail under covalent bonding. It is a demanding threshold: of the 118 elements, only eleven are gases at room temperature and just two, bromine and mercury, are liquids.