GCSE Chemistry States & separation revision

The particle model treats every substance as tiny moving particles. All matter in the universe is built from particles too small to see, and the three states differ only in how those particles sit and move. In a solid the particles are packed in a regular pattern, touching their neighbours, and can only vibrate about fixed positions — which is why a solid keeps both its shape and its volume. In a liquid the particles are still mostly touching but in a random arrangement, free to move past one another — so a liquid keeps its volume but takes the shape of its container. In a gas the particles are far apart, random, and moving fast in every direction — so a gas fills whatever space it is given and can be compressed.

Which state a substance takes at room temperature depends on the strength of the forces between its particles. Strong forces lock particles into a lattice: ionic solids such as sodium chloride and giant covalent solids such as diamond are solid far above room temperature. Weaker forces are easier to overcome: ethanol’s particles slip past each other easily, making it a liquid, while the very weakly attracted particles of methane leave it a gas at −162 °C and above.

The six changes of state come in three reversing pairs: melting ⇄ freezing, boiling ⇄ condensing, and sublimation ⇄ deposition (a solid such as dry ice turning straight to gas and back). They are physical changes: the particles themselves are untouched, only their arrangement and spacing change, and no new substance forms. Melted ice is still water; condensed steam is still water. That is why every change of state can be reversed by heating or cooling — unlike chemical changes such as burning, which produce new substances.

Energy always moves with the change. Melting, boiling and sublimation take energy in from the surroundings (endothermic); freezing and condensing give energy out to the surroundings (exothermic). Heat a beaker of melting ice with a steady flame and a thermometer surprises many students: the reading sits stubbornly at 0 °C until the last crystal has melted, then climbs. The energy is going somewhere — it is being used to break the attractive forces holding particles in fixed positions, not to make them move faster. Only once melting is complete can the energy raise the temperature of the liquid.

Models are simplifications, and examiners want you to say where this one stops. The particle model does not show the forces between particles — yet those forces decide most of the behaviour you learn about. It gives particles a size on the page that wildly misrepresents reality: in solids and liquids the particles are much smaller than the gaps the drawings suggest, and in a gas the true empty space is almost embarrassing — real gas particles occupy a vanishing fraction of the container. The model also says nothing about a particle’s internal structure (electrons, protons, neutrons) because at this level the particle is treated as a point mass.

• Not shown: interparticle forces and how strongly they attract.
• Not shown: the relative sizes of particles versus the spaces between them.
• Not shown: internal structure — each particle is drawn as a featureless ball.

In the laboratory, pure has a strict meaning: a substance made of one single element or one single compound, and nothing else. Distilled water is pure; sea water, though 96% water, is a mixture because dissolved salts ride along with it. The everyday sense of “pure” (nothing artificial added) is looser, and confusing the two is a classic exam slip. Melting and boiling behaviour is the experimental test: a pure substance melts and boils at one sharp, fixed temperature — ice at exactly 0 °C, pure water boiling at exactly 100 °C — while a mixture melts and boils gradually over a range because each dissolved or blended component shifts the behaviour slightly. An impure solid usually also melts at a lower temperature than the pure substance.

Useful mixtures are everywhere, and chemists design them deliberately. A formulation is a mixture whose composition is worked out precisely to give a product the right properties: petrol blended to ignite cleanly in engines; an alloy such as stainless steel mixed to resist rust; fertiliser granules blended to an exact N-P-K ratio; a tablet combining the drug with binders and coatings; washing powder mixing surfactants, enzymes and fragrance. Formulations are made to a recipe with exact quantities — change the recipe and you change the product.

Filtration separates an insoluble solid from a liquid. Pour the mixture through filter paper in a funnel: the liquid (the filtrate) passes through the tiny pores; the solid (the residue) is trapped. It works only for solids that genuinely do not dissolve — sand from water, chalk dust from acid. A dissolved salt sails straight through filter paper, because its particles are individual ions or molecules swimming in the water.

1. 1Crystallisation — recovering a soluble solid

   To get pure crystals of a dissolved salt, first remove the insoluble dirt by filtration (for rock salt: grind it, stir with water to dissolve the salt, filter off the grit).

2. 2Evaporate partially

   Warm the filtrate in an evaporating basin until crystals just begin to form at the edge — stopping before all the water is gone prevents the salt spitting and decomposing.

3. 3Leave to cool

   As the concentrated solution cools, the salt’s solubility falls, so crystals grow. Slower cooling gives bigger, purer crystals.

4. 4Dry

   Pat the crystals dry with filter paper or leave them in a warm place (or a desiccator).

The choice matters: filtration alone cannot get salt out of salt water, and crystallisation cannot get sand out of sand water. Every separation begins with the question — what is dissolved, and what is not?

Simple distillation separates a solvent from a solution, keeping the liquid. The solution is heated; the solvent evaporates first, its vapour travels into a water-cooled condenser (cold water entering at the bottom, leaving at the top so the jacket stays full), and drips out as the distillate. The dissolved solid is left behind in the flask — sea water gives fresh water plus a crust of salt. The method is beautiful but slow and energy-hungry, which is why industrial fresh water is usually made by other means.

When two liquids are mixed — ethanol and water, or the hundred compounds in crude oil — simple distillation cannot separate them cleanly: both vapours reach the condenser. Fractional distillation adds a fractionating column above the flask, hot at the bottom and cooler towards the top. Vapours rise, condense on the packing, and re-evaporate many times; each cycle enriches the vapour in the lower-boiling liquid. The liquid with the lowest boiling point climbs out of the top first; higher-boiling liquids condense part-way up and run back. Read off the thermometer: the steady temperature at the outlet names the fraction you are collecting.

The same physics at industrial scale runs the two great separations of chemistry. Crude oil enters a tall tower hot at the base; long-chain, high-boiling molecules condense low down as thick fractions, short-chain, low-boiling ones rise to the top as gases. Liquid air (about −200 °C) is warmed slowly: nitrogen boils off first at −196 °C, argon at −186 °C, oxygen last at −183 °C.

In paper chromatography a spot of the mixture is placed near the bottom of absorbent paper and the paper’s end is dipped in a solvent. Two phases do the work: the stationary phase — the paper, which does not move — and the mobile phase — the solvent, which creeps up the paper by capillary action. Each dissolved substance is pulled upward by the solvent but held back by attraction to the paper. The balance differs for every substance: ones that love the solvent race near its front; ones that cling to the paper lag behind. The result is a chromatogram: separated spots, one per substance, and a solvent front marking how far the solvent climbed.

Interpreting a chromatogram is a matching exercise. Substances that are identical produce spots at the same height when run side by side, so an unknown dye next to reference dyes can be identified by comparison. A pure substance shows one spot; a mixture shows several. (How to turn spot heights into numbers — Rf values — is covered in Chapter 16, Chemical analysis.)