States of Matter and the Gas Laws

Gases are invisible, yet they obey beautifully simple rules. Squeeze a gas and its pressure rises. Heat it and it expands. These relationships, the gas laws, were discovered through careful experiment and combine into one elegant equation. They explain everything from how a balloon behaves to how an engine works, and are a core part of chemistry.

Everything that has mass and takes up space is matter: the air you breathe, the water you drink, the chair you sit on. All matter is made of very small particles. Long ago, thinkers in India and Greece argued that matter could be divided only so far, into a smallest unit. Modern science confirms this. The particles of matter share three key features:

• Spaces between them: one substance can dissolve into another, like sugar in water, because particles slip into these gaps.
• Continuous motion: particles have kinetic energy, and that energy increases with temperature. This is why a smell spreads across a room (diffusion).
• Mutual attraction: a force of attraction holds particles together, strong in solids and weak in gases.

Matter exists in three common states (solid, liquid and gas), which differ in how their particles are arranged:

• Solids have particles packed very closely in a fixed order, with strong forces between them. So a solid has a fixed shape and fixed volume and is rigid.
• Liquids have particles a little farther apart and able to slide past one another. A liquid has a fixed volume but no fixed shape: it takes the shape of its container and can flow.
• Gases have particles far apart, moving fast and freely, with very weak forces. A gas has neither fixed shape nor fixed volume and can be compressed easily.

Matter can change from one state to another by changing the temperature or the pressure. Adding heat makes the particles move faster until they overcome their attractions. Removing heat does the reverse.

• Melting: solid to liquid, at the melting point.
• Boiling: liquid to gas, at the boiling point.
• Condensation: gas to liquid.
• Freezing: liquid to solid.
• Sublimation: a few substances (like camphor) change directly from solid to gas.

A key fact: while a substance is changing state, its temperature stays constant, even though heat is being supplied. This hidden heat is used to change the state rather than raise the temperature. It is called latent heat.

Evaporation is the change of a liquid into vapour below its boiling point (for example, wet clothes drying in the shade). It happens only at the surface of the liquid, as the faster particles escape.

Evaporation increases with:

• a higher temperature,
• a larger surface area,
• lower humidity (less water vapour already in the air),
• and a breeze that carries vapour away.

The escaping particles take energy with them, so evaporation causes cooling. This is why we feel cool when sweat evaporates, and why water stays cool in an earthen pot.

Boyle's Law describes how a gas responds to pressure at constant temperature. It states that, at a fixed temperature, the pressure of a gas is inversely proportional to its volume:

P × V = constant

So if you halve the volume of a gas, its pressure doubles. This is why a syringe with its outlet blocked becomes hard to push: squeezing the trapped air into a smaller space raises its pressure.

Charles's Law describes how a gas responds to temperature at constant pressure. It states that, at fixed pressure, the volume of a gas is directly proportional to its absolute (kelvin) temperature:

V ÷ T = constant

So heating a gas makes it expand. Cooling it makes it contract. This is why a balloon shrinks in the cold. A hot-air balloon rises because warm air expands and becomes less dense.

The separate gas laws can be combined into a single, powerful relationship: the ideal gas equation:

PV = nRT

Here P is pressure, V is volume, n is the amount of gas (in moles), T is the absolute temperature, and R is the universal gas constant. This one equation captures how pressure, volume, temperature, and amount of a gas are all linked. It applies (very nearly) to all gases under ordinary conditions. A gas that obeys it exactly is called an ideal gas. Real gases behave most like an ideal gas at high temperature and low pressure.

Liquids have their own characteristic behaviour at surfaces. The molecules of a liquid attract one another, a pull called cohesion. Molecules at the surface are pulled inward by the molecules below them, so the surface behaves like a stretched elastic skin. This effect is surface tension. It is why small water drops are spherical and why some insects can walk on water. Liquids also attract the molecules of other materials they touch, a pull called adhesion.

Capillarity (capillary action) is the rise or fall of a liquid inside a very narrow tube or porous material. When adhesion to the tube wall is stronger than the liquid's own cohesion, surface tension pulls the liquid up the tube against gravity. The narrower the tube, the higher the liquid climbs. Many everyday effects depend on capillarity:

• Lamp wick: kerosene rises through the fine fibres of the wick to feed the flame, so a kerosene lamp would not work without capillarity.
• Blotting paper: ink is drawn into the narrow pores of the paper by capillary action.
• Tall trees: capillarity in the fine vessels of the trunk helps draw water from the roots up to the leaves, so very tall trees could not survive without it.
• Towels and soil: a towel soaks up water, and water spreads through soil to plant roots, by the same mechanism.

One common trap must be avoided. Drinking through a straw has nothing to do with capillarity. Sucking on the straw lowers the air pressure inside it, and the higher atmospheric pressure on the drink's surface pushes the liquid up. A straw works by pressure difference, not by surface tension.