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Some reactions can run backwards as easily as forwards. In a closed container they settle into a balance where both directions happen at once. Learning to shift that balance is how industry makes the ammonia behind half the world's food.
Recommended for MYP 5 · eAssessment priority · About 3 lessons · Criteria A, B, C and D
By the end of this topic you should be able to:
You will use these skills. If any feel shaky, review them first.
In a reversible reaction, products can react to re-form the reactants. If one direction is exothermic, the other is endothermic by the same amount: heating blue hydrated copper sulfate drives off water, and adding water back to the white solid releases heat.
In a closed system the forward reaction starts fast and slows as reactants are used up, while the reverse reaction speeds up as products build up. When the two rates are equal, the amounts stop changing: a dynamic equilibrium.
Colour changes (such as the cobalt chloride and nitrogen dioxide equilibria) show the position shifting when conditions change. Experiments with isotopes show that atoms keep moving between reactants and products even at equilibrium.
Industry must choose conditions that give enough product, fast enough, at an acceptable cost. The Haber process, developed by Fritz Haber and Carl Bosch in the early 1900s, makes ammonia for fertilisers that feed about half the world's population.
Equilibrium links to rates of reaction and energy changes in chemistry; to oxygen binding in blood and ocean carbon dioxide in biology and Earth science; and to economics and sustainability.
When conditions change, the equilibrium shifts to oppose the change:
The Haber process: N₂ + 3H₂ ⇌ 2NH₃ (forward reaction exothermic). About 450 °C (a compromise between yield and rate), about 200 atm (a compromise between yield and cost), an iron catalyst, and ammonia removed and unreacted gases recycled.
Pakistan and many other countries produce nitrogen fertilisers such as urea from ammonia made by the Haber process. Photochromic sunglass lenses darken in sunlight and clear indoors because of a reversible reaction. The oxygen carried by haemoglobin in your blood is released in tissues because its binding is reversible.
For 2SO₂(g) + O₂(g) ⇌ 2SO₃(g) (forward exothermic), predict the effect of raising the pressure.
5.00 g of hydrated salt leaves 3.20 g after heating to constant mass. Find the percentage of water.
Equilibrium questions test clear reasoning about direction and cause. Expect:
Common ways to lose marks: saying the reaction stops at equilibrium; saying there are equal amounts at equilibrium; saying a catalyst increases the yield; and forgetting to count gas molecules for pressure.
Quick questions on the ideas above. Try each one before using a hint.
Partially guided investigation · about 45 minutes · pairs
Safety. Wear eye protection and gloves. Copper sulfate is harmful if swallowed and an irritant to eyes and skin; avoid raising dust and wash hands afterwards. Dispose of solutions as your teacher directs, not down the sink.
Then evaluate: where was energy lost, and how could your method reduce this?
Harder problems in unfamiliar contexts. Plan before you calculate.
Five questions picked at random from the whole topic. Take a new set whenever you like.
Questions you got wrong on this device appear here so you can try them again. Answer one correctly and it leaves the list.
Tracked separately for each skill, on this device only.
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