States of matter

October 16, 2026

Matter comes in four everyday states: solid, liquid, gas and plasma. The same substance can take any of them. Water is ice in the freezer, liquid in the tap and steam over the kettle, and the water molecule is the same in all three. What changes is how the molecules are arranged and how freely they move. That depends on a contest between the attraction that pulls neighbouring particles together and the thermal motion that shakes them apart.

That contest explains why ice floats while almost every other solid sinks in its own liquid, why sweat cools you, and why steam scalds worse than boiling water. It also explains how a pressure cooker cooks faster, how a fridge moves heat out of its inside, and why pasta takes longer to cook on a mountain.

Temperature is motion

Every atom and molecule is always moving: vibrating in place in a solid, tumbling past its neighbours in a liquid, or flying freely in a gas. Temperature measures how much kinetic energy that random motion carries on average, so a hotter substance has faster-moving particles. At room temperature the nitrogen molecules in air average about 470 m/s, faster than sound, colliding with one another billions of times a second. At absolute zero, −273.15 °C or 0 kelvin, the motion is as small as quantum mechanics allows, and nothing can be colder.

Particles also attract their neighbours, through the bonds and weaker forces described in Molecules and chemistry. The state of a substance comes from comparing those two energies: how strongly the particles attract, against how much thermal energy they have to break free.

• Solid: the attraction wins. Each particle stays in one place and only vibrates about it, so a solid keeps its shape.
• Liquid: the two are evenly matched. Particles stay in contact but have enough energy to slide past each other, so a liquid keeps its volume but flows to take the shape of its container.
• Gas: thermal energy wins. Particles break free of each other completely and fly about, so a gas spreads to fill whatever space it has.

This is why substances with strong attractions between their particles melt and boil at high temperatures. Tungsten atoms are joined by metallic bonds and melt at 3,422 °C. Water molecules are held by hydrogen bonds and melt at 0 °C, and nitrogen molecules attract only through weak van der Waals forces and melt at −210 °C.

Solids

In a solid every particle has a fixed place, and the arrangement of those places decides most of the solid's properties. Most solids are crystals: their particles sit in a pattern that repeats in every direction. Salt, sugar, ice, metals and most rocks are crystalline. A glass is a solid whose particles froze in a jumbled arrangement, like a liquid's, before they could line up. Window glass is the familiar one, but many plastics and some metals can be made glassy too.

The difference shows when they melt. Every particle in a crystal is held the same way, so the whole crystal lets go at one sharp temperature: ice melts at exactly 0 °C. The particles in a glass are held in many slightly different ways, so the solid softens gradually over hundreds of degrees, which is what lets a glassblower shape it.

A crystal is one small block of particles, the unit cell, repeated in every direction. Metals almost all use one of a few cells, because metal atoms are round and pack like stacked balls.

The stacking affects how a metal behaves. Face-centred cubic metals such as copper, aluminium and gold have many planes along which layers of atoms can slide, so they are soft and can be drawn into wire. Body-centred cubic iron has fewer easy planes and turns brittle in extreme cold, which is part of why the steel hulls of some ships in the 1940s cracked in icy seas.

Why ice floats

Most substances are densest as solids, because their particles pack closest when they stop moving around, but water is an exception. In ice every water molecule makes four hydrogen bonds at fixed angles, about 109° apart, and those angles force the molecules into open hexagonal rings with empty space inside them.

When ice melts, some of the hydrogen bonds break and the rings collapse, so the molecules crowd about 9% closer. Liquid water is therefore denser than ice (1.000 g/cm³ against 0.917), and ice floats with about a tenth of its volume above water. Lakes freeze from the top down, and the ice layer insulates the water below, which is why fish survive winter.

Liquids

A liquid's particles touch, like a solid's, but have enough energy to keep changing neighbours. Because the particles are already touching, liquids barely compress: squeezing water with a pressure of 100 atmospheres shrinks it by only half a percent. That is why hydraulic systems can push a car lift or an excavator arm with oil: the force goes through the liquid without being lost to squashing it.

Because each particle is pulled equally by all its neighbours inside a liquid but only from below and the sides at the surface, the surface acts like a stretched skin. This surface tension pulls drops into spheres and lets insects stand on water. Viscosity is a liquid's resistance to flowing, and it is high when the particles are large or tangled or attract strongly: honey is several thousand times more viscous than water. Warming a liquid lowers its viscosity, because the particles have more energy to break away from their neighbours.

Gases

In a gas the particles are far apart. In air at room temperature the molecules are about 3.4 nm apart on average, ten times their own size, so about 99.9% of a gas is empty space. That's why gases compress easily and why they mix completely with each other.

A gas pushes on everything it touches, because its molecules keep hitting surfaces and bouncing off, and every bounce is a small push. Pressure is the total of those pushes per unit area. Air at sea level presses on every square centimetre of you with a force of about 10 N, and you don't feel it only because it pushes equally from inside and out.

Halve the volume of a gas and the molecules hit each wall twice as often, so the pressure doubles. Double its absolute temperature and the molecules move faster, hitting the walls both harder and more often, and the pressure also doubles. Both facts are combined in the ideal gas law:

Here p is the pressure, V the volume, n the number of moles of gas, R the gas constant (8.314 J per mole per kelvin) and T the absolute temperature in kelvin. The law holds well for most gases at everyday conditions, because their molecules are far apart and barely interact.

A mole of any gas at 0 °C and 1 atmosphere fills 22.4 litres, whatever the gas is, because the law doesn't depend on what the molecules are. The law also says why car tyres are checked cold: a tyre at 2.2 bar above atmosphere (3.2 bar absolute) at 10 °C reaches about 3.4 bar absolute at 30 °C, because 303 K is 7% more than 283 K.

Changing state takes energy

Melting a solid or boiling a liquid needs energy even though the temperature doesn't change while it happens. That energy goes into pulling particles away from their neighbours instead of speeding them up, so a pot of boiling water stays at 100 °C however high the flame is turned. This hidden energy is called latent heat.

Heating a substance without changing its state takes its specific heat capacity, c, for every kilogram and every degree. Changing its state takes the latent heat, L, for every kilogram:

For water, c is 4.18 kJ/kg·K, melting takes 334 kJ/kg, and boiling takes 2,257 kJ/kg. Boiling costs so much more than melting because melting only loosens the hydrogen bonds, while boiling breaks every one of them and frees each molecule completely.

Those numbers explain several everyday effects:

• Ice cools a drink far better than cold water. Melting 1 kg of ice absorbs 334 kJ, the same as warming 1 kg of water by 80 °C.
• Sweating cools you. Evaporating a gram of sweat takes about 2.4 kJ from your skin, and evaporating a litre of sweat removes about 2,400 kJ, more than a quarter of the energy in a day's food.
• Steam scalds worse than boiling water. When steam condenses on your skin it gives up its latent heat, 2,257 kJ/kg, on top of the heat from cooling.
• A fridge pumps heat with latent heat. Its refrigerant evaporates inside the fridge, taking heat from the food, and condenses in the coils at the back, releasing that heat into the room.

Boiling and pressure

Even below its boiling point, a liquid slowly evaporates, because a few of its fastest molecules at the surface always have enough energy to escape. In a closed container, escaping molecules build up a vapour above the liquid until as many return as leave. The pressure of that vapour is the liquid's vapour pressure, and it rises steeply with temperature, since more molecules have enough energy to escape.

A liquid boils when its vapour pressure reaches the pressure of the air above it. Then bubbles of vapour can form inside the liquid, without being crushed by the surrounding pressure. Water's vapour pressure reaches one atmosphere at 100 °C, so that is its boiling point at sea level. Lower the air pressure and water boils at a lower temperature, and raise it and water boils hotter.

A phase diagram maps which state a substance takes at every temperature and pressure. On the summit of Everest the air pressure is about a third of sea level's, so water boils at 71 °C, and pasta cooked there stays hard. A pressure cooker holds the steam in until the pressure inside is about twice atmospheric, so the water reaches 120 °C and food cooks in about a third of the time.

Two points on the diagram are special:

• The triple point (0.01 °C, 0.6% of atmospheric pressure) is the only temperature and pressure where ice, liquid water and vapour all exist together. Below that pressure, ice turns straight to vapour without melting, which is how freeze-drying works. Carbon dioxide's triple point is at 5.1 atmospheres, so at normal air pressure solid CO₂, dry ice, turns straight to gas.
• The critical point (374 °C and 218 atmospheres for water) is where the boiling line ends. Past it, the liquid and the gas become so alike in density that there is no boundary between them, and the substance is a supercritical fluid. Supercritical carbon dioxide, which only needs 31 °C and 73 atmospheres, flows through coffee beans like a gas and dissolves caffeine like a liquid, and that is how most decaffeinated coffee is made.

Plasma

Heat a gas enough and its collisions become violent enough to knock electrons off its atoms. The result is a plasma: a gas of positive ions and free electrons. Because its charges are free to move, a plasma conducts electricity and responds to magnetic fields, and as its electrons fall back onto ions it glows. Hydrogen needs about 10,000 °C to be mostly ionised by heat alone, but an electric field can make a plasma at room temperature by accelerating the few free electrons until they knock others loose.

Plasma is rare on Earth but makes up nearly all the visible matter in the universe, because stars are plasma.

Other states

A few states only appear in special conditions:

• Liquid crystals flow like a liquid, but their long molecules line up in the same direction like a crystal. An electric field can turn them, which changes how they pass light, and that is how an LCD screen works.
• Superfluid helium forms below 2.17 K. It flows with no viscosity at all, and will creep up and over the walls of its container.
• Bose–Einstein condensates, first made in 1995, form when certain atoms are cooled to a few billionths of a degree above absolute zero. Thousands of atoms fall into the same quantum state and behave as one giant wave.