The chemistry behind treating acid rain, manufacturing fertilisers that feed billions, and neutralising acidic factory waste before it enters waterways is all the same reaction you wrote in L02 — scaled up by orders of magnitude and with real consequences if it goes wrong.
Use the PDF for classwork, homework or revision. It includes key ideas, activities, questions, an extend task and success-criteria proof.
In 1952, a thick smog settled over London for five days. By the time it lifted, an estimated 4,000 people had died — and the death toll eventually climbed to 12,000. The smog was loaded with sulfur dioxide from coal-burning power stations. SO₂ dissolves in water to form sulfurous acid; when oxidised, it becomes sulfuric acid. The result was acid fog with a pH as low as 1.6 — more acidic than stomach acid — settling into the lungs of everyone who breathed it.
Modern coal-fired power stations now use a neutralisation reaction to strip SO₂ from exhaust gases before they reach the atmosphere.
Before you read on: What chemical do you think is used to neutralise acidic SO₂ gas in an industrial exhaust stack? Write the reaction you think would occur — including the products. You will return to this at the end of the lesson.
📚 Core Content
Before any industrial scale is considered, the most personal application of neutralisation chemistry is happening inside your own body right now — and understanding exactly which reactions antacids use, and why different formulations work differently, connects L02 directly to everyday chemistry.
The stomach maintains a pH of approximately 1.5–3.5 using hydrochloric acid (HCl) secreted by parietal cells. This acidity is essential for protein digestion and for activating the enzyme pepsin. When excess acid is produced — due to stress, diet, or medical conditions — the result is heartburn or indigestion. Antacids work by neutralising excess HCl using one of three common active ingredients, each with a distinct reaction:
| Active ingredient | Formula | Reaction type | Balanced equation | Gas produced? | Side effect |
|---|---|---|---|---|---|
| Calcium carbonate (Tums, Rennie) | CaCO₃ | Acid + carbonate | CaCO₃ + 2HCl → CaCl₂ + H₂O + CO₂ | CO₂ (burping) | Rebound acid secretion at high doses |
| Magnesium hydroxide (Mylanta, Milk of Magnesia) | Mg(OH)₂ | Acid + base | Mg(OH)₂ + 2HCl → MgCl₂ + 2H₂O | None | Mild laxative effect (Mg²⁺) |
| Aluminium hydroxide (Gaviscon) | Al(OH)₃ | Acid + base | Al(OH)₃ + 3HCl → AlCl₃ + 3H₂O | None | Constipating effect (Al³⁺) |
Two of the most economically significant applications of neutralisation outside industry operate at the scale of millions of hectares of farmland and billions of tubes of toothpaste — both rely on the same carbonate and hydroxide chemistry from L02.
Soil pH correction: Most agricultural soils have an optimal pH range of 6.0–7.5 for nutrient availability and microbial activity. Soils become acidic through natural leaching, acid rain, and nitrogen-based fertilisers (which produce H⁺ as a byproduct of nitrification). When soil pH falls below 5.5, aluminium and manganese ions become soluble and toxic to plant roots. Lime (CaCO₃) or agricultural lime (Ca(OH)₂, "slaked lime") is applied to acidic soil:
CaCO₃ + 2H⁺ → Ca²⁺ + H₂O + CO₂ (CaCO₃ in acidic soil)
Ca(OH)₂ + 2H⁺ → Ca²⁺ + 2H₂O (Ca(OH)₂ in acidic soil)
CaCO₃ acts more slowly (lower solubility) but is cheaper and safer to handle; Ca(OH)₂ acts faster but is caustic.
Toothpaste: Most toothpastes are mildly basic (pH 7–10) due to the presence of sodium fluoride, sodium bicarbonate (NaHCO₃), and calcium carbonate. The mildly basic pH neutralises acids produced by bacteria in the mouth (primarily lactic acid from fermentation of sugars), which dissolve tooth enamel (hydroxyapatite, Ca₅(PO₄)₃OH) below pH 5.5. Fluoride ions (F⁻) replace OH⁻ in hydroxyapatite, forming the harder, more acid-resistant fluorapatite (Ca₅(PO₄)₃F).
The Haber process produces ammonia — but ammonia itself cannot be directly applied to soil as a fertiliser in most conditions. The conversion of NH₃ into solid, stable, transportable fertiliser salts requires direct acid-base neutralisation reactions, and H₂SO₄ and HNO₃ are the two acids used at the largest scale globally.
Global food production depends on nitrogen fertilisers — nitrogen is the most commonly limiting nutrient for plant growth, and atmospheric N₂ is inaccessible to most plants without being converted to NH₃ first (the Haber process: N₂ + 3H₂ ⇌ 2NH₃). The NH₃ is then converted to two major solid fertiliser salts through straightforward acid-base neutralisation:
Ammonium sulfate (NH₄)₂SO₄: Produced by reacting ammonia with sulfuric acid.
2NH₃ + H₂SO₄ → (NH₄)₂SO₄
NH₃ is the Brønsted-Lowry base (proton acceptor); H₂SO₄ is the acid. Two moles of NH₃ are needed per mole of H₂SO₄ because sulfuric acid is diprotic — it donates two protons, each requiring one NH₃ to accept it.
Ammonium nitrate NH₄NO₃: Produced by reacting ammonia with nitric acid.
NH₃ + HNO₃ → NH₄NO₃
NH₃ acts as the base and HNO₃ as the acid. Only 1 mol NH₃ is needed because HNO₃ is monoprotic — it donates only one proton.
Acid rain is one of the most well-documented examples of uncontrolled acid-base chemistry causing large-scale environmental damage — and flue gas desulfurisation is the engineering solution that uses neutralisation chemistry at the scale of millions of tonnes of SO₂ per year.
Coal and oil contain sulfur as an impurity. When burned, sulfur is oxidised: S + O₂ → SO₂. Sulfur dioxide in the atmosphere dissolves in water droplets: SO₂ + H₂O → H₂SO₃, and is further oxidised: 2SO₂ + O₂ + 2H₂O → 2H₂SO₄. This acidic precipitation (acid rain, pH 4–5 or lower) causes significant damage:
Flue gas desulfurisation (FGD) — "scrubbing" — prevents SO₂ from reaching the atmosphere by reacting it with calcium hydroxide (Ca(OH)₂) in a wet scrubber before flue gas is released:
Ca(OH)₂(aq) + SO₂(g) → CaSO₃(s) + H₂O(l)
This is an acid + base reaction — SO₂ is acidic (it dissolves to form H₂SO₃) and Ca(OH)₂ is the base. The product CaSO₃ (calcium sulfite) is a solid that can be collected, or further oxidised to CaSO₄ (calcium sulfate, gypsum) which has commercial applications in wallboard and cement manufacture.
Every factory, mining operation, and chemical plant that produces acidic or alkaline liquid waste is legally required to neutralise that waste before discharge — and pH control of water systems is both an environmental imperative and an application of neutralisation chemistry at the largest scale most students will ever encounter.
Industrial processes generate enormous volumes of acidic or alkaline wastewater. Metal ore processing (acid mine drainage), electroplating (acidic pickling baths), and pharmaceutical manufacturing all produce effluent at pH values far outside the safe range for aquatic ecosystems (typically pH 6.5–8.5 for healthy freshwater systems).
Acidic industrial effluent is neutralised by adding a base — typically Ca(OH)₂ (cheap, readily available, effective), NaOH (more soluble, faster-acting but more expensive), or CaCO₃ (slowest but cheapest). A critical additional function: as pH rises during neutralisation, dissolved heavy metal ions such as Fe³⁺, Cu²⁺, Pb²⁺, and Zn²⁺ form insoluble hydroxide precipitates:
Fe³⁺(aq) + 3OH⁻(aq) → Fe(OH)₃(s)↓ (heavy metal precipitation)
| Effluent type | Source | Neutralising agent | Reaction | Additional benefit |
|---|---|---|---|---|
| Acidic (pH < 6) | Acid mine drainage, electroplating | Ca(OH)₂ or NaOH | Acid + base → salt + water | Precipitates heavy metal hydroxides for removal |
| Acidic (pH < 6) | Same as above | CaCO₃ | Acid + carbonate → salt + H₂O + CO₂ | Cheaper than Ca(OH)₂; slower |
| Alkaline (pH > 9) | Paper mills, cement works | CO₂ gas or dilute H₂SO₄ | CO₂ + H₂O → H₂CO₃; acid + base → salt + water | CO₂ is mild — avoids over-acidification risk |
The two fertiliser equations — 2NH₃ + H₂SO₄ → (NH₄)₂SO₄ and NH₃ + HNO₃ → NH₄NO₃ — represent reactions that collectively produce tens of millions of tonnes of fertiliser per year, feeding roughly 40% of the global population. The FGD equation Ca(OH)₂ + SO₂ → CaSO₃ + H₂O represents a reaction running 24 hours a day in hundreds of power stations worldwide, preventing millions of tonnes of SO₂ from becoming acid rain. These are not textbook abstractions — they are the largest-scale acid-base reactions on Earth.
Both depend on the same principle: a base (Ca(OH)₂ or NH₃) neutralising an acid (SO₂ or H₂SO₄/HNO₃), producing a salt plus water (or other products). The chemistry is L02 — the scale is planetary.
"Mg(OH)₂ and Al(OH)₃ produce CO₂ when they react with HCl." CO₂ is only produced when a carbonate or hydrogen carbonate reacts with an acid. Hydroxide bases (Mg(OH)₂, Al(OH)₃) react by the acid + base pattern and produce only salt + water — no CO₂ is possible because there is no carbonate ion in the reactants.
"1 mol NH₃ is needed to react with H₂SO₄." H₂SO₄ is diprotic — it donates 2 protons, requiring 2 mol of NH₃ (one per proton). The correct equation is 2NH₃ + H₂SO₄ → (NH₄)₂SO₄. Writing NH₃ + H₂SO₄ → NH₄HSO₄ gives only the first neutralisation step (incomplete neutralisation with excess H₂SO₄).
"SO₃ is the gas neutralised by FGD." SO₂ is the primary combustion product of sulfur in coal and oil. The HSC-required FGD equation uses SO₂: Ca(OH)₂ + SO₂ → CaSO₃ + H₂O. SO₃ is a secondary product from further oxidation; writing SO₃ gives a different product (CaSO₄, gypsum) and will not match marking guidelines for the primary FGD reaction.
"Naming the neutralising agent is sufficient for wastewater treatment questions." HSC marking requires three components: (1) the neutralising agent, (2) the balanced equation, and (3) the environmental consequence of pH control — specifically the impact on aquatic ecosystem pH tolerance. Naming Ca(OH)₂ alone earns at most one of the marks available.
📝 Worked Examples
Problem: Write balanced molecular equations for the following industrial neutralisation reactions: (a) the production of ammonium nitrate fertiliser from ammonia and nitric acid; (b) the neutralisation of sulfur dioxide in flue gas using calcium hydroxide; (c) the neutralisation of hydrochloric acid in acidic industrial effluent using calcium carbonate.
Ammonium nitrate production: Reaction type: acid + base. NH₃ is the base (proton acceptor); HNO₃ is the acid (monoprotic — donates 1 H⁺). Product: NH₄⁺ + NO₃⁻ → NH₄NO₃. Unbalanced: NH₃ + HNO₃ → NH₄NO₃.
Balance check: N = 1 + 1 = 2 left; 1 + 1 = 2 right ✓. H = 3 + 1 = 4 left; 4 right ✓. O = 3 left; 3 right ✓. Already balanced.
FGD: Reaction type: acidic oxide + base → salt + water. SO₂ acts as the acid component; Ca(OH)₂ is the base. Product: CaSO₃ (calcium sulfite) + H₂O. Unbalanced: Ca(OH)₂ + SO₂ → CaSO₃ + H₂O.
Balance check: Ca = 1 ✓, S = 1 ✓, O = 2 + 2 = 4 left; 3 + 1 = 4 right ✓, H = 2 ✓. Already balanced.
Effluent neutralisation: Reaction type: acid + carbonate → salt + water + CO₂. Acid = HCl; carbonate = CaCO₃. Salt formed: Ca²⁺ + 2Cl⁻ → CaCl₂. Unbalanced: HCl + CaCO₃ → CaCl₂ + H₂O + CO₂.
Need 2 HCl to provide 2 Cl⁻ for CaCl₂. Balanced: 2HCl + CaCO₃ → CaCl₂ + H₂O + CO₂. Check: Ca=1✓, Cl=2✓, H=2✓, C=1✓, O = 3 left; 1 + 2 = 3 right ✓.
Answers: (a) NH₃ + HNO₃ → NH₄NO₃ (b) Ca(OH)₂ + SO₂ → CaSO₃ + H₂O (c) 2HCl + CaCO₃ → CaCl₂ + H₂O + CO₂
Problem: A patient with kidney disease cannot excrete large amounts of calcium or magnesium ions. They experience chronic acid reflux. Their pharmacist needs to recommend an antacid that: (i) effectively neutralises HCl; (ii) does not introduce significant Ca²⁺ or Mg²⁺ ions; (iii) does not cause excessive CO₂ gas production. Evaluate each antacid (CaCO₃, Mg(OH)₂, Al(OH)₃) against these criteria and justify a recommendation.
Evaluate CaCO₃: Reaction: CaCO₃ + 2HCl → CaCl₂ + H₂O + CO₂. Produces Ca²⁺ ions — fails criterion (ii) for a patient unable to excrete Ca²⁺. Also produces CO₂ gas — fails criterion (iii). Not suitable.
Evaluate Mg(OH)₂: Reaction: Mg(OH)₂ + 2HCl → MgCl₂ + 2H₂O. Produces Mg²⁺ ions — fails criterion (ii). Does not produce CO₂ — satisfies criterion (iii). Partially suitable but ion load is a concern. Not optimal.
Evaluate Al(OH)₃: Reaction: Al(OH)₃ + 3HCl → AlCl₃ + 3H₂O. Produces Al³⁺ ions — not Ca²⁺ or Mg²⁺, so criterion (ii) is satisfied for this patient's specific condition. Does not produce CO₂ — satisfies criterion (iii). Effectively neutralises HCl — satisfies criterion (i).
Recommendation: Al(OH)₃ best satisfies all three stated criteria. Note: high-dose Al³⁺ has its own long-term concerns in renal patients — this would need clinical consideration beyond the chemical criteria given here.
Answer: Al(OH)₃ — satisfies all three criteria: neutralises HCl effectively (Al(OH)₃ + 3HCl → AlCl₃ + 3H₂O), introduces Al³⁺ rather than Ca²⁺ or Mg²⁺, and produces no CO₂ gas. Clinical caveat: long-term Al³⁺ load in renal patients requires separate clinical consideration.
Problem: A coal-fired power station burns coal with 3% sulfur content and releases flue gas containing approximately 2000 ppm SO₂. The station has no flue gas desulfurisation system. Describe the chemistry by which SO₂ causes environmental damage to a nearby freshwater lake, and evaluate the effectiveness of installing a Ca(OH)₂ scrubbing system as a solution. Include balanced equations for all relevant reactions.
How SO₂ causes acid rain and lake acidification:
Sulfur in coal is oxidised during combustion: S + O₂ → SO₂.
SO₂ released into atmosphere dissolves in water droplets: SO₂(g) + H₂O(l) → H₂SO₃(aq).
In the atmosphere, SO₂ is further oxidised: 2SO₂ + O₂ + 2H₂O → 2H₂SO₄ (sulfuric acid).
Acid rain (pH 4–5) falls into the lake. As lake pH drops: (1) carbonate buffering is overwhelmed (HCO₃⁻ + H⁺ → H₂CO₃ → H₂O + CO₂); (2) aluminium ions become soluble and toxic to fish gills (Al(OH)₃(s) + 3H⁺ → Al³⁺ + 3H₂O); (3) aquatic invertebrates die; (4) fish populations collapse below pH 5.
Effectiveness of Ca(OH)₂ scrubbing: Installing a wet scrubber with Ca(OH)₂ slurry in the exhaust stack neutralises SO₂ before it is released:
Ca(OH)₂(aq) + SO₂(g) → CaSO₃(s) + H₂O(l)
CaSO₃ precipitates as a solid and is collected. If further oxidised: CaSO₃ + ½O₂ → CaSO₄ (gypsum — commercial value in construction). Modern FGD systems achieve 90–99% SO₂ removal, reducing flue gas SO₂ from 2000 ppm to 20–200 ppm — within regulatory limits.
Evaluation of effectiveness:
Chemical effectiveness: High — the neutralisation reaction is rapid, complete, and produces a manageable solid byproduct.
Economic consideration: Ca(OH)₂ is cheap and widely available; CaSO₄ byproduct can be sold as synthetic gypsum, partially offsetting operating costs.
Limitations: Does not address CO₂ emissions from combustion (which cause ocean acidification via a separate mechanism); requires continuous supply of Ca(OH)₂ and solid waste management; does not remediate the already-acidified lake (which would require direct liming of the lake).
Overall assessment: FGD with Ca(OH)₂ is highly effective at preventing further SO₂-related acidification by neutralising SO₂ at the source. It represents the most direct and chemically sound solution to the problem described.
Summary answer: SO₂ → H₂SO₃ → H₂SO₄ via atmospheric reactions; acid rain (pH 4–5) acidifies lake; carbonate buffering overwhelmed; Al³⁺ mobilised and toxic; ecosystem collapses. FGD: Ca(OH)₂ + SO₂ → CaSO₃ + H₂O; 90–99% SO₂ removal; CaSO₄ byproduct has commercial value. Evaluation: highly effective at source; does not address CO₂ or remediate existing lake acidification; economically viable with byproduct offsetting costs.
🧪 Activities
Look back at what you wrote in the Think First section. What has changed? What did you get right? What surprised you?
🔢 Multiple Choice
1. A power station installs a flue gas desulfurisation system using Ca(OH)₂. Which equation correctly represents the primary reaction occurring in the scrubber?
2. A farmer applies calcium carbonate to acidic soil with pH 4.5 to raise it to pH 6.5 for wheat cultivation. Which explanation correctly describes why the soil pH increases?
3. A student compares two fertilisers: ammonium sulfate (NH₄)₂SO₄ (produced by 2NH₃ + H₂SO₄ → (NH₄)₂SO₄) and ammonium nitrate NH₄NO₃ (produced by NH₃ + HNO₃ → NH₄NO₃). Which statement correctly compares the two production reactions?
4. A patient takes an antacid for acid reflux. Which of the following antacid reactions produces CO₂ gas as a product?
A patient takes an antacid for acid reflux. Select the option that antacid reactions produces CO₂ gas as a product?
5. A mining company discharges acidic wastewater containing dissolved heavy metal ions (Fe³⁺, Cu²⁺) into a treatment pond where Ca(OH)₂ is added. Which pair of statements correctly describes what happens during treatment?
✍️ Short Answer
6. Write balanced molecular equations for the reactions of all three common antacid active ingredients with hydrochloric acid: (a) calcium carbonate, (b) magnesium hydroxide, (c) aluminium hydroxide. For each, identify the reaction type and state whether CO₂ gas is produced. Explain why two of the three antacids do not produce CO₂. 4 MARKS
7. A fertiliser company uses ammonia (NH₃) produced by the Haber process to manufacture two products: ammonium sulfate and ammonium nitrate. (a) Write balanced equations for the production of each fertiliser. (b) Using Brønsted-Lowry theory, explain why 2 mol of NH₃ are needed to neutralise H₂SO₄ but only 1 mol is needed to neutralise HNO₃. (c) A student argues that H₂SO₄ is a stronger acid than HNO₃, and that is why more NH₃ is needed. Is this reasoning correct? Explain. 5 MARKS
8. Extended Response — The 1952 London Smog: A coal-fired power station releases 2000 ppm SO₂ in its flue gas with no treatment system. (a) Describe, with balanced equations, how SO₂ in the atmosphere leads to acid rain and the chemical damage it causes to a nearby limestone statue and a freshwater lake. (4 marks) (b) Evaluate the effectiveness of installing a Ca(OH)₂ scrubber as a solution. Include the balanced equation, the removal efficiency, any byproducts and their uses, and identify one limitation of this approach. (3 marks) 7 MARKS
Go back to your Think First predictions at the top of this lesson.
1. FGD: Ca(OH)₂ + SO₂ → CaSO₃ + H₂O. Acid component = SO₂ (acidic oxide — dissolves to form H₂SO₃); Base = Ca(OH)₂. CaSO₃ is the primary product because the FGD reaction involves SO₂ (not SO₃). CaSO₄ would only form if CaSO₃ were further oxidised (CaSO₃ + ½O₂ → CaSO₄) — this is a separate step, not the primary scrubber reaction.
2. Fertiliser production: 2NH₃ + H₂SO₄ → (NH₄)₂SO₄. H₂SO₄ is diprotic — it has two ionisable H⁺ (from both ionisations). Each H⁺ must be accepted by one NH₃ molecule (NH₃ is a Brønsted-Lowry base — proton acceptor), so 2 mol NH₃ needed per mole H₂SO₄. NH₃ + HNO₃ → NH₄NO₃. HNO₃ is monoprotic — only one ionisable H⁺, requiring only 1 mol NH₃.
3. Mining wastewater: Neutralising agent: Ca(OH)₂ (or NaOH or CaCO₃). Acid neutralisation (e.g. if acid is H₂SO₄): H₂SO₄ + Ca(OH)₂ → CaSO₄ + 2H₂O. Fe³⁺ precipitation: Fe³⁺(aq) + 3OH⁻(aq) → Fe(OH)₃(s)↓. Environmental importance: pH 3.2 is strongly acidic — far below the safe range (6.5–8.5) for aquatic life. Direct discharge would lower receiving waterway pH, dissolve carbonate sediments, mobilise toxic metals, and kill aquatic organisms from bacteria to fish.
1. NH₃ + H₂SO₄ → NH₄HSO₄: This equation is not wrong in isolation — it represents neutralisation of only the first H⁺ from H₂SO₄ (forming ammonium hydrogen sulfate). It would be correct if only 1 mol NH₃ was used with 1 mol H₂SO₄. Industrially, excess NH₃ is used to neutralise both protons: correct equation is 2NH₃ + H₂SO₄ → (NH₄)₂SO₄. Error in industrial context: only 1 mol NH₃ used when H₂SO₄ requires 2.
2. Ca(OH)₂ + SO₃ → CaSO₄ + H₂O: The error in context is using SO₃ rather than SO₂. SO₂ is the primary sulfur combustion product and the target of FGD. SO₃ forms in smaller amounts by further oxidation and would produce CaSO₄ (gypsum) directly — a different reaction. Correct primary FGD equation: Ca(OH)₂ + SO₂ → CaSO₃ + H₂O.
3. Mg(OH)₂ + 2HCl → MgCl₂ + H₂O + CO₂: Error — CO₂ cannot be a product. CO₂ is only produced when a carbonate or hydrogen carbonate reacts with an acid. Mg(OH)₂ is a hydroxide base — the reaction is acid + base → salt + water only. Correct equation: Mg(OH)₂ + 2HCl → MgCl₂ + 2H₂O.
4. "CaCO₃ releases OH⁻ ions": Incorrect — CaCO₃ does not contain OH⁻ and does not release OH⁻ when it dissolves. The correct mechanism: CaCO₃ + 2H⁺ → Ca²⁺ + H₂O + CO₂. The CO₃²⁻ ions react with H⁺ ions in the acid soil water (consuming H⁺). Reducing [H⁺] raises pH — without any OH⁻ being produced.
5. Two problems: (1) "Just raise to pH 7" is insufficient — the legal safe discharge range is 6.5–8.5, but even at pH 7, dissolved heavy metal ions (Fe³⁺, Cu²⁺) may still be present at toxic concentrations. Metal hydroxide precipitation requires pH to be raised high enough for each specific metal to precipitate (e.g. Fe(OH)₃ precipitates above pH ~3, but Cu(OH)₂ requires higher pH). Neutralisation alone does not remove these ions — filtration of the precipitated hydroxides is also required. (2) Describing neutralisation as "making it safe" oversimplifies — the wastewater may still contain toxic dissolved species, suspended solids, and organic compounds beyond the acid that require separate treatment.
1. B — The FGD reaction with Ca(OH)₂ targets SO₂ (the primary sulfur combustion product) and produces calcium sulfite (CaSO₃) and water. SO₂ acts as the acidic component (acidic oxide); Ca(OH)₂ is the base. Option A uses SO₃ (wrong gas) producing CaSO₄ (different product). Option C uses CaCO₃ not Ca(OH)₂. Option D gives a non-existent compound Ca(SO₂)₂ with incorrect stoichiometry.
2. B — CaCO₃ reacts with H⁺ in acidic soil by the acid + carbonate reaction: CaCO₃ + 2H⁺ → Ca²⁺ + H₂O + CO₂. This consumes H⁺ ions, reducing [H⁺] in soil water and raising pH. Option A is wrong — Ca²⁺ does not neutralise H⁺; it is the CO₃²⁻ that reacts. Option C is wrong — CaCO₃ is not a hydroxide and releases no OH⁻. Option D misrepresents the mechanism — CaCO₃ consumes H⁺ directly (it does not just buffer without pH change).
3. B — H₂SO₄ is diprotic (can donate 2 H⁺); HNO₃ is monoprotic (donates 1 H⁺). The number of moles of NH₃ required equals the number of protons the acid can donate — 2 for H₂SO₄, 1 for HNO₃. Option A incorrectly states HNO₃ is diprotic. Option C incorrectly links acid strength (not relevant here — both are strong acids) to mole ratio. Option D incorrectly states NH₃ accepts 2 protons — NH₃ accepts only 1 proton to form NH₄⁺.
4. A — CO₂ is produced only when a carbonate reacts with an acid (acid + carbonate → salt + water + CO₂). CaCO₃ contains carbonate — so CaCO₃ + 2HCl → CaCl₂ + H₂O + CO₂ produces CO₂. Mg(OH)₂ and Al(OH)₃ are hydroxide bases (no carbonate) — they produce only salt + water. NaOH is a hydroxide base — salt + water only.
5. C — Ca(OH)₂ reacts with H⁺ in the acidic wastewater to raise pH (neutralisation). As pH rises, [OH⁻] increases. Heavy metal ions form insoluble hydroxide precipitates: Fe³⁺ + 3OH⁻ → Fe(OH)₃↓ (and similarly for Cu²⁺). These precipitates can be filtered out. Option A is wrong — heavy metals do precipitate as pH rises. Option B is wrong — Ca(OH)₂ raises pH (reduces acidity). Option D is wrong — Ca(OH)₂ does not react directly with metal ions as stated; the mechanism is pH-dependent precipitation via OH⁻.
Q6 (4 marks): (a) CaCO₃ + 2HCl → CaCl₂ + H₂O + CO₂. Type: acid + carbonate. CO₂ produced — yes [1]. (b) Mg(OH)₂ + 2HCl → MgCl₂ + 2H₂O. Type: acid + base. CO₂ produced — no [1]. (c) Al(OH)₃ + 3HCl → AlCl₃ + 3H₂O. Type: acid + base. CO₂ produced — no [1]. Explanation: CO₂ is only produced when a carbonate or hydrogen carbonate reacts with an acid — a carbonate ion (CO₃²⁻) accepts H⁺ to form H₂CO₃, which decomposes to CO₂ + H₂O. Mg(OH)₂ and Al(OH)₃ contain no carbonate ion; they react by the acid + base pattern producing only salt and water — no CO₂ is possible [1].
Q7 (5 marks): (a) 2NH₃ + H₂SO₄ → (NH₄)₂SO₄; NH₃ + HNO₃ → NH₄NO₃ [1 each = 2 marks]. (b) Brønsted-Lowry: NH₃ acts as a proton acceptor (base). H₂SO₄ is diprotic — it contains two ionisable hydrogen atoms and can donate two protons (H⁺). Each H⁺ must be accepted by one NH₃ molecule to form NH₄⁺. Therefore 2 mol NH₃ is required per mole of H₂SO₄. HNO₃ is monoprotic — it contains only one ionisable hydrogen and donates only one proton, requiring only 1 mol NH₃ [2]. (c) The student's reasoning is incorrect. Acid strength (strong vs weak) describes the degree of ionisation — the fraction of molecules that donate protons to water. The mole ratio of NH₃ required depends only on the number of protons that can be donated per molecule (proticity), not on the acid's strength. Both H₂SO₄ and HNO₃ are strong acids — fully ionised — but H₂SO₄ requires 2 mol NH₃ because it is diprotic (2 ionisable H⁺), not because it is stronger [1].
Q8 (7 marks): (a) Combustion: S + O₂ → SO₂ [½]. Atmospheric chemistry: SO₂ + H₂O → H₂SO₃; 2SO₂ + O₂ + 2H₂O → 2H₂SO₄. Acid rain (pH 4–5) produced [1]. Limestone statue: CaCO₃ + H₂SO₄ → CaSO₄ + H₂O + CO₂. CaSO₄ is slightly soluble and washes away, causing surface erosion and loss of detail [1]. Freshwater lake: pH drops; carbonate buffering overwhelmed; Al(OH)₃(s) + 3H⁺ → Al³⁺ + 3H₂O — Al³⁺ is toxic to fish gills; aquatic invertebrates and fish populations collapse [1–2 marks depending on detail]. (b) FGD equation: Ca(OH)₂ + SO₂ → CaSO₃ + H₂O [1]. Removal efficiency: modern FGD removes 90–99% of SO₂, reducing 2000 ppm to ≤20–200 ppm — within regulatory limits [1]. Byproduct: CaSO₃ can be oxidised to CaSO₄ (gypsum), used commercially in wallboard and cement manufacture, partially offsetting operating costs [½]. Limitation: FGD does not remove CO₂ (which causes ocean acidification and climate change through a separate mechanism); it also does not remediate the already-acidified lake — separate liming of the lake would be required [½].
pH, pOH and the Water Dissociation Constant
Tick when you've finished all activities and checked your answers.