States of Matter

The three states of matter

All matter is made of particles. The same particles — atoms, molecules, or ions — can sit in three quite different arrangements, and each gives the substance its familiar character. Recognising the three states is the foundation for everything that follows.

Solids

A solid has a fixed shape and a fixed volume, and is generally the densest of the three states. Its particles are packed closely together in a regular pattern, held in place by strong forces of attraction. They cannot move past one another — instead, they vibrate about fixed positions. This is why a solid keeps its shape: the particles, though jostling, never leave their seats.

Liquids

A liquid has a fixed volume but takes the shape of its container. Its particles are still close together but no longer in a regular arrangement — they are packed randomly. They can move past one another, sliding and tumbling, which is why a liquid flows and fills the bottom of a container. Liquids are usually less dense than solids, though water is a notable exception.

Gases

A gas has no fixed shape and no fixed volume; it fills whatever container holds it. Its particles are far apart, move quickly and randomly in all directions, and have very little attraction between them. The collisions of these particles with the container walls produce the pressure of the gas. Because the particles are so spread out, gases are easily compressed and have low density.

Changes of state

A change of state is a change in how particles are arranged, not in what they are. Energy added to a substance lets its particles move more — eventually enough to break free from the arrangement that held them. Removing energy reverses the process.

The five named changes

Melting is the change from solid to liquid, occurring at the melting point. Freezing is its reverse: liquid to solid, at the same temperature. Boiling is liquid to gas, occurring at a fixed boiling point and throughout the body of the liquid.

Evaporation is also liquid to gas, but it happens only at the surface and over a range of temperatures below the boiling point — higher temperatures and larger surface areas speed it up. Condensation is the reverse of boiling and evaporation: gas to liquid as the substance cools.

ExtendedHeating and cooling curves

When a solid is heated steadily, its temperature rises until it reaches the melting point. There the temperature stops rising, even though heat is still being added — the energy is being used to break the forces holding the particles in their regular arrangement, not to make them move faster. Once melting is complete, the temperature rises again until the boiling point, where another plateau appears for the same reason. A cooling curve is the mirror image: plateaus at the same temperatures as energy is released to the surroundings.

Gas behaviour: temperature & pressure

When you heat a gas, or squeeze it, you are changing what its particles do. Because a gas’s volume is so responsive to those changes, the gas state is where the kinetic particle theory is most visibly at work.

Effect of temperature

Increasing the temperature of a gas increases the average kinetic energy of its particles. They move faster and, on average, farther apart. If the container is flexible — like a balloon — the gas expands and the volume increases. If the container is rigid, the particles instead hit the walls more frequently and with more force, so the pressure rises.

The link runs through average kinetic energy: temperature is a measure of it, so raising temperature raises the speed of particle motion. Faster particles collide with the walls more often and more forcefully — the observable consequence is a rise in volume (flexible container) or pressure (rigid container).

Effect of pressure

If the pressure on a gas is increased — for example by pushing in a piston — the same number of particles is squeezed into a smaller space. The volume decreases. Reducing the pressure has the opposite effect: the particles spread into the larger space available.

At constant temperature, the particles still have the same average kinetic energy after the change in pressure. What changes is how densely packed they are. In a smaller volume, more collisions happen per second with each unit of wall area — which is why a high-pressure gas pushes outward more strongly than a low-pressure one.

Diffusion

Diffusion is the everyday consequence of the random motion described by kinetic particle theory. Wherever particles can move freely — in gases and in liquids — they spread out from where there are many of them to where there are few, until eventually they are mixed evenly.

The process

Diffusion occurs only in fluids — in gases and in liquids — because only there can the particles move past one another. In a solid the particles are locked in place, so no diffusion happens. Diffusion needs no external energy: the random kinetic motion of the particles already drives it. Raising the temperature does speed it up, because the particles then move faster.

A classic demonstration: a crystal of purple potassium manganate(VII) dropped into a beaker of still water. After a few hours, the purple colour has spread evenly through the whole beaker, even though no one has stirred it. The same demonstration with bromine vapour in air spreads much faster, because gas particles move faster than liquid particles.

ExtendedThe effect of molecular mass

At the same temperature, all gas particles have on average the same kinetic energy. But kinetic energy depends on both mass and speed — so heavier particles must move more slowly to share the same energy. The consequence: a gas with a lower relative molecular mass diffuses faster.

The standard demonstration uses ammonia, NH₃ (Mr = 17), and hydrogen chloride, HCl (Mr = 36.5). Soaked cotton wool plugs of each are placed at opposite ends of a long glass tube. Both gases diffuse along the tube and meet to form a white ring of solid ammonium chloride. Because ammonia is lighter, it travels faster — so the white ring forms closer to the hydrogen chloride end.