GCSE Chemistry Equilibrium revision

A reversible reaction produces products that can react again to remake the reactants: A + B ⇌ C + D. The classic example is hydrated copper(II) sulfate: the blue crystals lose water on heating to leave white powder, and the white powder turns blue again when water is added. Equilibrium can only settle in a closed system — nothing added, nothing escaping. In an open beaker the escaping substance is lost and the reaction just runs one way.

Left alone in a closed system, the reaction reaches dynamic equilibrium: the forward and reverse reactions are still happening — that is the dynamic part — but at equal rates, so the amounts of reactants and products stop changing. Equal rates, not equal amounts: the equilibrium can sit far to one side. The position of equilibrium describes that balance — far right means mostly products, far left means mostly reactants.

Le Chatelier’s principle: when the conditions of an equilibrium change, the position shifts to oppose the change. Concentration: add more reactant and the equilibrium shifts towards the products to use it up. Pressure (gases only): compressing the mixture shifts the equilibrium towards the side with fewer moles of gas, because fewer molecules lowers the pressure. Temperature: heating shifts the equilibrium in the endothermic direction to absorb the extra heat; cooling shifts it in the exothermic direction. For an exothermic forward reaction (like the Haber process), that means heating reduces the product yield — an awkward truth industry must design around.

A catalyst speeds up the forward and reverse reactions equally, so equilibrium is reached faster but the position does not move — the equilibrium amount of product is unchanged. And a warning for predictions: pressure has nothing to do with solids or pure liquids, so count only gaseous (or dissolved) moles on each side.

The Haber process makes ammonia for fertilisers: N₂(g) + 3H₂(g) ⇌ 2NH₃(g), with the forward reaction exothermic (heat released). Purified nitrogen (from air) and hydrogen (usually from natural gas) are passed over an iron catalyst at about 450 °C and 200 atmospheres. Only about a third of the mixture reacts each pass, so the ammonia is condensed to a liquid (it has a far higher boiling point than nitrogen or hydrogen) and the unreacted gases are recycled back over the catalyst. Recycling means that overall, most of the nitrogen and hydrogen eventually become ammonia.

Why 450 °C? A lower temperature would give a higher yield (the exothermic direction is favoured when you take heat away) — but the rate would be impossibly slow, even with the catalyst. A higher temperature gives a fast rate but a poor yield. 450 °C is a compromise — fast enough to be economic, with an acceptable yield. Similarly, 200 atm rather than 1000 atm: pressure favours the ammonia side (4 gas moles → 2), but very high pressure means expensive pumps, thick-walled vessels and safety risks. Another compromise.

Percentage yield = (actual yield ÷ theoretical yield) × 100. The theoretical yield comes from the balanced equation — the maximum possible mass of product. Real reactions rarely hit 100%: the reaction may be reversible (some product turns back), reactants may be impure, product is lost in transfers and purification, or side reactions make unwanted products. Note the difference between yield (how much you collected, in practice) and atom economy (Chemistry only): a purely theoretical measure of how much of the reactant mass ends up in the molecule you want — atom economy = (Mr of desired product ÷ sum of Mr of ALL products) × 100, using the balanced equation. A single-product reaction is automatically 100%.

Industry weighs both numbers: a route with low atom economy throws mass away as waste (cost, disposal, environmental impact), so chemists design high atom economy routes and find uses for co-products — selling the leftovers instead of paying to dump them. Chemistry only: the same logic explains why the Haber process is attractive — it has a single product, so its atom economy is 100% even though its percentage yield per pass is nothing like it.

Plants need three main nutrients in fertiliser: nitrogen (N), phosphorus (P) and potassium (K). An NPK fertiliser is a formulation — a mixture blended to a deliberate recipe of several compounds, each chosen for its nutrient and its solubility, often with fillers that keep the prills free-flowing. Roots can only absorb nutrients dissolved in water, so every nutrient compound in the blend must be soluble.

1. 1Nitrogen salts

   Ammonia (from the Haber process) neutralises nitric acid to make ammonium nitrate, or sulfuric acid for ammonium sulfate: NH₃ + HNO₃ → NH₄NO₃. Both are soluble nitrogen sources.

2. 2Phosphorus salts

   Phosphate rock is mined but its phosphates are insoluble. Treating it with phosphoric acid gives soluble phosphates; nitric acid or sulfuric acid give other soluble phosphate/nitrate mixtures.

3. 3Potassium salts

   Potassium chloride and potassium sulfate are mined directly (as in the Yorkshire potash mines) and dissolve readily — no chemical conversion needed.

Making a fertiliser salt in the lab (ammonium sulfate is the usual example): a small batch, titrated to find the exact neutralisation point, with evaporation and crystallisation — slow and small-scale, but precise and safe. Industrial production is a continuous process: huge volumes of ammonia and acid stream through at controlled stoichiometry, making tonnes per hour far more cheaply per tonne. The lab optimises for care; the plant optimises for scale. (This chapter has no named required practical, but the fertiliser preparation is examined as a method question — learn the titrate-neutralise-evaporate sequence.)