Overview
What you’ll learn
- Describe reversible reactions with the ⇌ symbol, using hydrated and anhydrous copper(II) sulfate (and cobalt(II) chloride) as examples.
- State that dynamic equilibrium exists only in a closed system, where forward and reverse rates are equal and concentrations stay constant.
- State the Haber process equation and conditions (about 450 °C, about 200 atm, iron catalyst) and explain each condition as a compromise between yield, rate and cost.
- Describe the uses of ammonia, including the manufacture of nitric acid and nitrogenous fertilisers.
- Name the three essential elements in NPK fertilisers and write equations for making ammonium salts from ammonia.
- Outline how over-use of fertilisers can cause water pollution (eutrophication).
In 30 seconds
- 1A reversible reaction runs in both directions: CuSO₄·5H₂O ⇌ CuSO₄ + 5H₂O (heat drives right, water drives left).
- 2Dynamic equilibrium needs a closed system: forward and reverse reactions never stop, but run at equal rates, so concentrations stay constant (not equal).
- 3Haber equation: N₂ + 3H₂ ⇌ 2NH₃, exothermic forward (ΔH negative); 4 volumes of gas become 2, so high pressure favours ammonia.
- 4Haber conditions: about 450 °C, about 200 atm, iron catalyst — each one a compromise: low temperature gives more NH₃ but too slowly; high pressure is costly and dangerous; the catalyst speeds both directions equally.
- 5Ammonia is manufactured, then oxidised to nitric acid; ammonia + nitric acid → ammonium nitrate, the classic nitrogenous fertiliser.
- 6NPK = nitrogen, phosphorus, potassium — the three elements plants need most; examples: ammonium nitrate, ammonium sulfate, potassium chloride blends.
- 7Ammonium salts are made by neutralising NH₃ solution with the right acid: NH₃ + HNO₃ → NH₄NO₃; 2NH₃ + H₂SO₄ → (NH₄)₂SO₄.
- 8Over-using fertilisers washes nitrates into rivers → eutrophication: algae bloom, oxygen drops, aquatic life suffocates.
Notes · 1 of 5
Reversible reactions
Some reactions run both ways — heat one way, water the other.
Heat blue copper(II) sulfate crystals and they spit out their water of crystallisation, turning white: CuSO₄·5H₂O → CuSO₄ + 5H₂O. Drip water back on and the blue returns. The pair is written with the double arrow: CuSO₄·5H₂O ⇌ CuSO₄ + 5H₂O — heating is the forward drive, water the reverse. (Blue-to-pink cobalt(II) chloride paper works the same way and is the standard test for water.)
The same double arrow governs the biggest reaction in fertiliser chemistry: N₂ + 3H₂ ⇌ 2NH₃. Left-to-right builds ammonia; right-to-left tears it down. Which direction wins depends on the conditions — which is exactly what the next section is about.
Notes · 2 of 5
Dynamic equilibrium
A closed system where both reactions run flat out — and cancel exactly.
Leave a reversible reaction in a closed system (nothing in, nothing out) and it reaches dynamic equilibrium: the forward and reverse reactions continue at equal rates, so the concentrations of everything stay constant. Dynamic because molecules keep reacting both ways; equilibrium because nothing visible changes.
| Statement | True? | Why |
|---|---|---|
| Forward and reverse reactions have stopped | No | Both run continuously — that is the "dynamic" part |
| Forward and reverse rates are equal | Yes | The definition of dynamic equilibrium |
| Concentrations of reactants and products are equal | No | They are constant, not equal — the position can sit anywhere |
| It can be reached in an open beaker | No | Gas escapes; a closed system is required |
Notes · 3 of 5
The Haber process
Nitrogen from air, hydrogen from natural gas, and three carefully chosen compromises.
The Haber process fixes atmospheric nitrogen into ammonia: N₂(g) + 3H₂(g) ⇌ 2NH₃(g), with the forward reaction exothermic. Purified gases flow over trays of iron catalyst at about 450 °C and 200 atmospheres; the ammonia is cooled and liquefied out of the mix, and unreacted gases are recycled — so nothing is wasted even though each pass converts only a fraction.
| Condition | Effect on yield | Effect on rate / cost | The compromise |
|---|---|---|---|
| 450 °C (not colder) | Lower temperature would give MORE NH₃ (forward is exothermic) | Too cold = far too slow | Warm enough for a fast rate, cool enough for usable yield |
| 200 atm (not higher) | Higher pressure would give MORE NH₃ (4 gas volumes → 2) | Very high pressure = expensive pumps, thick walls, danger | High but affordable pressure |
| Iron catalyst | No effect on yield | Speeds up BOTH directions equally — equilibrium reached sooner | Faster production per day, same yield |
Notes · 4 of 5
Ammonia: properties, nitric acid and uses
An alkaline gas you already know — now its industrial life story.
You met ammonia in Sec 3: a pungent, alkaline gas that turns damp red litmus blue (NH₃ + H₂O ⇌ NH₄⁺ + OH⁻) and fumes with hydrogen chloride (NH₃ + HCl → NH₄Cl). Industrially it is far more than a gas bottle: ammonia is the feedstock for nitric acid (catalytic oxidation of NH₃ over platinum, then absorption in water) and for nearly all nitrogenous fertilisers.
Ammonium salts are straightforward neutralisations: bubbled into nitric acid, NH₃ + HNO₃ → NH₄NO₃ (ammonium nitrate); into sulfuric acid, 2NH₃ + H₂SO₄ → (NH₄)₂SO₄ (ammonium sulfate). Both are made as solutions, then evaporated to crystallise — titration-style precision, exactly as in Sec 3 Chapter 7.
Notes · 5 of 5
NPK fertilisers and water pollution
Three elements feed the world; too much of a good thing chokes the rivers.
Plants need three elements in largest amount: nitrogen (N) for leaves, phosphorus (P) for roots, potassium (K) for flowers and fruit — hence NPK fertilisers, blends chosen for the crop. Single-nutrient examples: ammonium nitrate and ammonium sulfate (N), potassium chloride (K), and phosphate salts (P). Ammonium salts are extra useful because the nitrate/ammonium ions are directly absorbable by roots.
The downside: rain washes excess fertiliser into rivers and lakes. Dissolved nitrates trigger algal blooms; when the algae die, decomposing bacteria multiply and strip the dissolved oxygen, so fish and other aquatic life suffocate — eutrophication. Careful dosing, slow-release formulations and buffer strips near waterways all limit the damage.
Revise
Key terms
The words examiners expect you to use precisely.
- Reversible reaction
- A reaction that proceeds in both directions at once, written with the ⇌ symbol.
- Dynamic equilibrium
- In a closed system, forward and reverse reactions running at equal rates, so concentrations stay constant.
- Closed system
- A set-up where no matter enters or escapes — required for true equilibrium.
- Haber process
- The industrial manufacture of ammonia: N₂ + 3H₂ ⇌ 2NH₃ at about 450 °C, 200 atm, over iron.
- Compromise conditions
- Chosen values balancing yield against rate and cost, since the ideal-yield option is too slow or expensive.
- Recycling (unreacted gases)
- Haber gases that did not react are cooled, separated and passed back over the catalyst.
Show 6 more key termsShow fewer key terms
- Nitrogenous fertiliser
- A fertiliser supplying nitrogen, e.g. ammonium nitrate or ammonium sulfate.
- NPK fertiliser
- A blend supplying nitrogen, phosphorus and potassium, the three elements plants need most.
- Ammonium nitrate
- NH₄NO₃, made from ammonia + nitric acid; the standard nitrogenous fertiliser.
- Eutrophication
- Fertiliser run-off feeding algal blooms whose decay strips dissolved oxygen from water, killing aquatic life.
- Exothermic (forward)
- The Haber forward reaction releases heat — hence lower temperatures favour ammonia yield.
- Catalytic oxidation
- Ammonia oxidised over platinum as the first step of nitric acid manufacture.
Revise
Common mistakes
Where students most often lose marks in this chapter — and what to do instead.
Mistake: Saying equilibrium means reactant and product concentrations are equal.
Equal is for rates, not amounts.
Instead: At equilibrium the concentrations are constant but generally unequal.
Mistake: Trying to reach equilibrium in an open beaker.
Escaping gas keeps the reverse reaction from proceeding.
Instead: Dynamic equilibrium requires a closed system.
Show 10 more mistakesShow fewer mistakes
Mistake: Claiming the catalyst increases the ammonia yield.
The catalyst speeds both directions equally.
Instead: Catalysts reach equilibrium faster; yield depends on temperature, pressure and the equilibrium itself.
Mistake: Using the highest possible pressure in the Haber process.
Yield would rise, but pumps, reactors and safety costs explode too.
Instead: About 200 atm is the economic compromise; quote cost as the limiting factor.
Mistake: Using a very low temperature for maximum yield.
Equilibrium would sit right, but the rate would be uselessly slow.
Instead: About 450 °C trades some yield for an industrially useful rate.
Mistake: Writing the Haber equation without the reversible arrow.
A single arrow claims complete conversion, which never happens.
Instead: N₂ + 3H₂ ⇌ 2NH₃ — the ⇌ is the heart of the topic.
Mistake: Preparing ammonium sulfate by the excess-solid method.
Both ammonia solution and sulfuric acid are soluble; there is no solid to filter.
Instead: Use titration to find neutral volumes, then crystallise.
Mistake: Saying ammonium nitrate supplies P or K.
Its formula carries only N, H and O.
Instead: Ammonium nitrate is a nitrogenous (N) fertiliser only; P and K come from other salts.
Mistake: Explaining eutrophication without the decomposers.
Algae do not suck out the oxygen themselves — the decay does.
Instead: Algal bloom → death → bacterial decomposition consumes dissolved oxygen → suffocation.
Mistake: Heating hydrated copper(II) sulfate and calling the whitening "reversible" with no water added.
Reversing needs the water back.
Instead: Forward = heating (whitens); reverse = adding water (blues) — quote both directions.
Mistake: Confusing ammonia (NH₃) with the ammonium ion (NH₄⁺).
One is a covalent gas, the other an ion in salts.
Instead: NH₃ is the alkaline gas; NH₄⁺ is what it becomes in acids and fertiliser salts.
Mistake: Assuming the recycled unreacted gases increase the yield per pass.
Each pass still converts only a fraction.
Instead: Recycling raises the OVERALL conversion — the same per-pass yield, repeated many times.
Check yourself
Quick quiz
20 questions with instant explanations. About 5 minutes.
Question 1 of 3
- Question 1: current question
- Question 2: not answered
- Question 3: not answered
Recall reversibility
The symbol ⇌ in an equation shows that the reaction
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Practise
Topical questions
Exam-style practice in three tiers, from recall to data questions.
- 4Foundation
- 5Standard
- 3Challenging
12 questions worth 47 marks, each with a model answer, mark scheme and examiner tips.
Practise topical questionsExam practice
Practice paper
A 30-mark paper set out like a school paper, with a separate mark scheme. Give yourself 40 min.
Question paper
Chapter 5 Practice Paper
30 marks · 40 min · 4 pages
5 multiple-choice questions and structured questions with answer lines. Chapter 5 (reversible reactions, dynamic equilibrium, Haber process, NPK fertilisers).
Mark scheme
PROChapter 5 Practice Paper — mark scheme
Answers · Marking points · 3 pages
Model answers with a mark for each creditworthy point — for marking at home or in class.
Where this chapter is tested
Each school sets its own scope — check your assessment outline.
Questions
Frequently asked questions
What does Sec 4 Chemistry Chapter 5 cover?
It covers sub-topic 3.7 of O-Level Pure Chemistry 6092: reversible reactions (hydrated salts as the classic example), dynamic equilibrium in closed systems, the Haber process with its temperature/pressure/catalyst compromises, ammonia’s conversion to nitric acid, and NPK fertilisers including eutrophication.
Why does the Haber process use 450 °C when lower temperatures give more ammonia?
Because the forward reaction is exothermic, cooling raises the yield — but the rate collapses. 450 °C is the compromise that produces useful ammonia per hour; recycling the unreacted gases recovers most of the yield loss.
Does the iron catalyst increase the amount of ammonia made?
No. It speeds the forward and reverse reactions equally, so equilibrium is reached sooner but the equilibrium position — the eventual yield — is unchanged. Temperature and pressure move the yield; the catalyst only buys time.
How is ammonium sulfate prepared in the lab?
By titration, since both reactants are soluble: neutralise ammonia solution against sulfuric acid with an indicator, repeat without the indicator, then evaporate to crystallisation. 2NH₃ + H₂SO₄ → (NH₄)₂SO₄.
When is this chapter usually tested in school?
Ammonia and equilibrium close out the reaction chemistry of Sec 4 and appear in the prelim and O-Level papers, often as a conditions-compromise question fused with rate reasoning. Scope varies by school, so check your school’s assessment outline.
Up next
Chapter 6: The Periodic Table