Year 12 Chemistry Module 6, Acid/Base Reactions ⏱ ~45 min Lesson 4 of 19 IQ1

Neutralisation in Everyday Life & Industry

In 1952, the London Great Smog killed an estimated 4,000 people in four days, coal-burning SO₂ dissolved in fog to form sulfuric acid aerosol. The UK government's 1956 Clean Air Act mandated industrial flue-gas scrubbing using Ca(OH)₂ to neutralise SO₂. The chemistry is exactly the same neutralisation reaction studied here: acid + base → salt + water, scaled to industrial volumes.

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Today's hook, In 1952, London's Great Smog killed 4,000 people in four days, SO₂ from coal fires dissolved in fog and formed sulfuric acid aerosol. The UK's 1956 Clean Air Act forced every industrial furnace to scrub SO₂ using Ca(OH)₂. Today: how does the same neutralisation chemistry that fixes your heartburn also prevent industrial disasters?

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Get oriented

Warm up first

Three quick questions from earlier lessons. Pulling old material back to mind before you learn something new makes the new material stick better, so this is not busywork.

Worksheets

Practise this lesson

Four printable worksheets that build from the foundations up to exam-style questions, start at whatever level suits you.

Build Foundations & guided practice Apply Application practice Master Mastery challenge Exam Exam-style questions

Think First

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.

Key Reaction Patterns, This Lesson

Acid + base → salt + water

e.g. 2NH₃ + H₂SO₄ → (NH₄)₂SO₄ | Ca(OH)₂ + SO₂ → CaSO₃ + H₂O | Mg(OH)₂ + 2HCl → MgCl₂ + 2H₂O

Acid + carbonate → salt + water + CO₂

e.g. CaCO₃ + 2HCl → CaCl₂ + H₂O + CO₂ | CaCO₃ + H₂SO₄ → CaSO₄ + H₂O + CO₂

Key industrial equations (memorise all):

Flue gas desulfurisation: Ca(OH)₂ + SO₂ → CaSO₃ + H₂O Ammonium sulfate fertiliser: 2NH₃ + H₂SO₄ → (NH₄)₂SO₄ Ammonium nitrate fertiliser: NH₃ + HNO₃ → NH₄NO₃ Antacid (carbonate): CaCO₃ + 2HCl → CaCl₂ + H₂O + CO₂ Antacid (hydroxide): Mg(OH)₂ + 2HCl → MgCl₂ + 2H₂O | Al(OH)₃ + 3HCl → AlCl₃ + 3H₂O

Know

Balanced equations for all three antacid reactions with HCl
The two fertiliser production equations (ammonium sulfate, ammonium nitrate)
The flue gas desulfurisation equation: Ca(OH)₂ + SO₂ → CaSO₃ + H₂O
Four application areas: antacids, soil correction, fertiliser, FGD, wastewater

Understand

Why 2 mol NH₃ is needed for H₂SO₄ but only 1 mol for HNO₃ (diprotic vs monoprotic)
Why CO₂ is produced by carbonate antacids but not by hydroxide antacids
How FGD prevents SO₂ from causing acid rain (acid + base at source)
Why pH control of industrial effluent has environmental importance

Can Do

Write and balance any industrial neutralisation equation from memory
Evaluate antacid choices for a patient using chemical criteria
Explain the environmental consequences of SO₂ and the chemistry of FGD
Write a high-mark extended response on pH control and environmental chemistry

Key Terms, scan these before reading

Antacid

A base (e.g., Mg(OH)₂, CaCO₃) that neutralises excess stomach acid (HCl); used in medicine.

Agricultural liming

Adding CaCO₃ or Ca(OH)₂ to acidic soil to raise pH and improve crop growth.

Diprotic acid

An acid with two ionisable protons (e.g., H₂SO₄, H₂CO₃); reacts in two stages.

Monoprotic acid

An acid with one ionisable proton (e.g., HCl, HNO₃, CH₃COOH).

Salt

An ionic compound formed from the cation of a base and the anion of an acid.

Wastewater treatment

Industrial neutralisation of acidic or alkaline effluent before discharge; often uses lime or sodium carbonate.

Antacids and stomach acid

Core Content

Cross-lesson links: The industrial neutralisation applications here extend the balanced equations from L02–L03 into real-world contexts. Acid rain chemistry (SO₂, NOₓ) connects to atmospheric buffer systems (L13). The agricultural liming reaction reappears in L18–L19 as a context for titration analysis. Polyprotic acids like H₂SO₄ are treated in detail in L04 and revisited in buffer calculations (L12–L13).

Antacids and Stomach Acid, Medical Neutralisation Chemistry

CaCO₃, Mg(OH)₂, Al(OH)₃ all neutralise HCl, but by different reaction types

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:

Scope note

We analyse antacids purely as acid-base chemistry (which base neutralises HCl, what products form, how much heat is released). This is not medical advice, choice of medication, dosing and treatment of persistent symptoms are decisions for a doctor or pharmacist, not conclusions you can draw from the neutralisation equation alone.

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⁺)

Must Know: All three antacid equations must be balanced, this is a standard short-answer exam task. Check each equation: CaCO₃ + 2HCl → CaCl₂ + H₂O + CO₂ (Ca=1✓, C=1✓, O=3✓, H=2✓, Cl=2✓). Mg(OH)₂ + 2HCl → MgCl₂ + 2H₂O (Mg=1✓, O=2✓, H=4✓, Cl=2✓). Al(OH)₃ + 3HCl → AlCl₃ + 3H₂O (Al=1✓, O=3✓, H=6✓, Cl=3✓).

Common Error: Students write CO₂ as a product for Mg(OH)₂ and Al(OH)₃ reactions with HCl. CO₂ is only produced when a carbonate or hydrogen carbonate reacts with an acid. Mg(OH)₂ and Al(OH)₃ are hydroxide bases, they react by the acid + base pattern and produce only salt + water. There is no carbonate ion present, so no CO₂ is possible.

Insight: Antacids provide temporary relief by raising stomach pH, but the stomach responds to increased pH by secreting more acid (a feedback mechanism). This "rebound effect" is particularly pronounced with CaCO₃. Proton pump inhibitors (PPIs, e.g. omeprazole) inhibit the enzyme that secretes HCl rather than neutralising it, they are more effective for chronic conditions but are not neutralisation chemistry.

Three antacid equations: CaCO₃ + 2HCl → CaCl₂ + H₂O + CO₂; Mg(OH)₂ + 2HCl → MgCl₂ + 2H₂O; Al(OH)₃ + 3HCl → AlCl₃ + 3H₂O, CO₂ is produced only from carbonate antacids; hydroxide antacids produce salt + water only; all three equations must be balanced for full marks.

Pause, copy the highlighted equations into your book before moving on.

Quick check, which antacid reaction produces no CO₂ and leaves no excess gas?

Soil pH correction and toothpaste

Soil pH Correction and Toothpaste, Agriculture and Dentistry

CaCO₃ and Ca(OH)₂ raise acidic soil pH · Toothpaste neutralises bacterial lactic acid

We just saw how antacids use carbonates and hydroxides to neutralise excess stomach HCl. That raises a question: Where else does neutralisation chemistry operate at large scale? This card answers it → the same reactions are used to correct acidic soil (liming) and to protect tooth enamel from bacterial acids.

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).

Acidic soil correction

Problem: Soil pH < 5.5; Al³⁺/Mn²⁺ toxicity

Neutralising agent: CaCO₃ or Ca(OH)₂

Reaction type: Acid + carbonate or acid + base

Target pH: Raise to 6.0–7.5

Toothpaste (acid protection)

Problem: Bacterial lactic acid dissolves enamel

Neutralising agent: NaHCO₃, CaCO₃, NaF

Reaction type: Neutralises lactic acid; F⁻ hardens enamel

Target pH: Maintain mouth pH > 5.5

Must Know: In HSC questions about soil pH correction, you must: (1) name the specific chemical (CaCO₃ or Ca(OH)₂), (2) write the balanced equation showing it reacting with soil acid, and (3) state the purpose (raising pH to optimal range for crop growth). All three components are required for full marks, naming the chemical alone is insufficient.

Common Error: Students say toothpaste "neutralises acids completely" or "makes the mouth basic." The mechanism is more nuanced, toothpaste's mild basicity provides protection against acid attack, and fluoride ion replacement of hydroxide in hydroxyapatite provides structural protection. The neutralisation component reduces [H⁺] enough to prevent enamel dissolution (which begins below pH 5.5), not to make the mouth strongly basic.

Acidic soil (pH < 5.5 causes Al³⁺/Mn²⁺ toxicity) is corrected by CaCO₃ + 2H⁺ → Ca²⁺ + H₂O + CO₂ (slower, cheaper) or Ca(OH)₂ + 2H⁺ → Ca²⁺ + 2H₂O (faster, caustic), raising pH to 6.0–7.5; toothpaste (pH 7–10) neutralises bacterial lactic acid and F⁻ replaces OH⁻ in hydroxyapatite forming acid-resistant fluorapatite.

Add the highlighted point to your notes before the check below.

A farmer adds CaCO₃ to soil with pH 4.5. Why does the soil pH increase?

Deep dive, Extension Fertiliser production

Fertiliser Production, Acid-Base Chemistry at Industrial Scale

2NH₃ + H₂SO₄ → (NH₄)₂SO₄ | NH₃ + HNO₃ → NH₄NO₃

We just saw how neutralisation corrects acidic soil at the agricultural scale. That raises a question: Where does the nitrogen in those fertilisers come from, and how is NH₃ converted to a form plants can use? This card answers it → Haber-process NH₃ is neutralised with H₂SO₄ or HNO₃ to make stable solid fertiliser salts at industrial scale.

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, so 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, so it donates only one proton.

Ammonium sulfate

Formula: (NH₄)₂SO₄

Reaction: 2NH₃ + H₂SO₄ → (NH₄)₂SO₄

% N by mass: 21%

Consideration: Safe to handle; acidifies soil with long-term use

Ammonium nitrate

Formula: NH₄NO₃

Reaction: NH₃ + HNO₃ → NH₄NO₃

% N by mass: 34%

Consideration: Higher N content; fire/explosion risk; regulated

Urea (reference, not acid-base)

Formula: CO(NH₂)₂

Reaction: Industrial synthesis from NH₃ + CO₂

% N by mass: 46%

Consideration: Most concentrated; not a neutralisation product

Must Know: The equation 2NH₃ + H₂SO₄ → (NH₄)₂SO₄ is frequently tested, it connects Module 5 (Haber process, equilibrium) to Module 6 (acid-base neutralisation). Be able to write it from memory, balance it, and explain why 2 mol of NH₃ is required: H₂SO₄ is diprotic, it donates 2 protons, each requiring one NH₃ molecule to accept it.

Common Error: Students write NH₃ + H₂SO₄ → NH₄HSO₄ (ammonium hydrogen sulfate) instead of (NH₄)₂SO₄. This is the product of only the first ionisation of H₂SO₄. When equal or excess NH₃ is used (as in industrial production), both protons are neutralised and the product is (NH₄)₂SO₄. Always check the mole ratio before writing the product.

Insight: Long-term use of ammonium sulfate gradually acidifies soil because NH₄⁺ undergoes nitrification by soil bacteria: 2NH₄⁺ + 3O₂ → 2NO₂⁻ + 4H⁺ + 2H₂O, releasing H⁺ ions. This is why regular liming (soil pH correction with CaCO₃ or Ca(OH)₂) is necessary on farms using ammonium-based fertilisers, the acid-base chemistry of the fertiliser and the soil correction are directly linked.

Fertiliser production via Brønsted-Lowry acid-base neutralisation: 2NH₃ + H₂SO₄ → (NH₄)₂SO₄ (2 mol NH₃ because H₂SO₄ is diprotic); NH₃ + HNO₃ → NH₄NO₃ (1 mol NH₃ because HNO₃ is monoprotic), NH₃ is the proton acceptor (base) in both reactions.

Pause, write the highlighted equations into your book.

Why does 2NH₃ + H₂SO₄ → (NH₄)₂SO₄ require 2 moles of NH₃?

Deep dive, Extension Flue gas desulfurisation

Flue Gas Desulfurisation, Preventing Acid Rain at the Source

Ca(OH)₂ + SO₂ → CaSO₃ + H₂O | Removes 90–99% of SO₂ before discharge

We just saw how fertiliser production generates NH₃-based salts via acid-base reactions. That raises a question: What happens when acidic gases like SO₂ escape into the atmosphere instead of being neutralised? This card answers it → SO₂ forms acid rain (pH 4–5) which corrodes infrastructure and kills aquatic life; FGD uses Ca(OH)₂ scrubbing to neutralise SO₂ at the source.

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:

Acidifies lakes and streams, killing fish and aquatic invertebrates
Leaches nutrients from soil, reducing forest productivity
Corrodes limestone and marble buildings: CaCO₃ + H₂SO₄ → CaSO₄ + H₂O + CO₂

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.

FGD with slaked lime

Reactants: Ca(OH)₂ + SO₂ (gas)

Balanced equation: Ca(OH)₂ + SO₂ → CaSO₃ + H₂O

Product: CaSO₃ (solid, collected)

Environmental benefit: Prevents SO₂ reaching atmosphere; reduces acid rain

FGD with limestone

Reactants: CaCO₃ + SO₂ (gas)

Balanced equation: CaCO₃ + SO₂ → CaSO₃ + CO₂

Product: CaSO₃ + CO₂

Environmental benefit: Cheaper reagent; CO₂ byproduct is a limitation

Acid rain damage to limestone

Reactants: CaCO₃ + H₂SO₄

Balanced equation: CaCO₃ + H₂SO₄ → CaSO₄ + H₂O + CO₂

Product: CaSO₄ (soluble, washes away)

Environmental impact: Explains building/statue corrosion

Must Know: The FGD equation Ca(OH)₂ + SO₂ → CaSO₃ + H₂O is a specific HSC-required equation. Note that SO₂ behaves as an acidic oxide, it reacts with a base (Ca(OH)₂) to form a salt (CaSO₃) and water. This follows the acid + base → salt + water pattern where the "acid" is the acidic oxide SO₂ rather than an aqueous acid. Treat SO₂ as the acid component when writing and balancing this equation.

Common Error: Students write SO₃ instead of SO₂ as the exhaust gas being neutralised. SO₂ is the primary combustion product of sulfur; SO₃ forms in smaller amounts by further oxidation. The FGD process primarily targets SO₂, writing SO₃ gives the wrong balanced product: Ca(OH)₂ + SO₃ → CaSO₄ + H₂O (a different reaction producing gypsum directly). The HSC-required equation involves SO₂.

A wet scrubber uses Ca(OH)₂ slurry to remove SO₂ from coal power station flue gas, producing clean gas and a gypsum by-product.

Flue gas desulfurisation converts acidic SO₂ to solid CaSO₃ before exhaust reaches the atmosphere

SO₂-rich flue gas meets a calcium hydroxide slurry that removes the sulfur dioxide.

In flue gas desulfurisation, acidic SO₂ is removed before exhaust is released, converting it into a collectable solid byproduct.

Flue gas desulfurisation: Ca(OH)₂(aq) + SO₂(g) → CaSO₃(s) + H₂O(l), SO₂ from coal combustion dissolves to form H₂SO₃/H₂SO₄ causing acid rain (pH 4–5) which corrodes CaCO₃ structures; FGD removes 90–99% SO₂ and produces CaSO₃ which can be oxidised to gypsum (CaSO₄) for commercial use.

Add the highlighted equation to your notes before the check below.

What is the correct primary equation for flue gas desulfurisation using slaked lime?

Deep dive, Extension. Two contrasting applications carry the core dot point, the antacid case and wastewater/soil treatment. Fertiliser production and flue gas desulfurisation are additional industrial cases, valuable context but not required for the core route.

Wastewater treatment

Wastewater Treatment, pH Control in Industrial Effluent

Acidic effluent + Ca(OH)₂ → salt + water · Legal requirement: pH 6.5–8.5 for discharge

We just saw that FGD prevents acidic SO₂ from reaching the atmosphere. That raises a question: What about acidic liquid waste that industrial processes generate? This card answers it → legal requirements mandate pH 6.5–8.5 for discharge, achieved by base neutralisation which also precipitates toxic heavy metal ions as insoluble hydroxides.

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

Must Know: HSC questions on wastewater treatment often ask three things: (1) name the neutralising agent, (2) write a balanced equation, AND (3) explain the environmental importance of pH control. All three parts are required. The environmental importance must explicitly state the effect on aquatic ecosystems: "acidic effluent lowers the pH of the receiving waterway, reducing dissolved oxygen availability and making conditions uninhabitable for fish and aquatic invertebrates."

Common Error: Students answer wastewater treatment questions by only naming the neutralising agent without addressing the environmental consequence. Naming Ca(OH)₂ earns one mark; explaining that uncontrolled acidic discharge would lower waterway pH, dissolve carbonate sediments, reduce biodiversity, and introduce toxic heavy metal ions earns the remaining marks. The environmental reasoning is typically worth more marks than the chemical equation in HSC marking guidelines.

Insight: Acid mine drainage (AMD) is one of the most persistent environmental challenges in Australian mining. When sulfide ore minerals (pyrite, FeS₂) are exposed to air and water during mining, they oxidise spontaneously: 4FeS₂ + 15O₂ + 14H₂O → 4Fe(OH)₃ + 8H₂SO₄. The H₂SO₄ produced can generate AMD with pH as low as 2. The orange-red staining of rocks near mine sites is Fe(OH)₃ precipitate, visible evidence of AMD neutralisation chemistry occurring naturally.

Acidic industrial effluent must be neutralised to pH 6.5–8.5 before discharge using Ca(OH)₂, NaOH, or CaCO₃; as pH rises, dissolved heavy metal ions precipitate: Fe³⁺ + 3OH⁻ → Fe(OH)₃(s)↓, HSC answers must include: (1) neutralising agent, (2) balanced equation, (3) environmental consequence for the aquatic ecosystem.

Pause, copy the highlighted principle into your book.

A mining company adds Ca(OH)₂ to acidic wastewater containing Fe³⁺ ions. What TWO things happen?

Modelling neutralisation reactions

Modelling Neutralisation Reactions, Two Representations of the Same Chemistry

Particle diagram (H₃O⁺ + OH⁻ → 2H₂O) and symbolic ionic equation, two media for one reaction

We just saw five applications of neutralisation, written only as molecular equations. That raises a question: a molecular equation like HCl + NaOH → NaCl + H₂O is one way to represent what happens, but is it the only way to show the actual chemistry occurring between the particles? This card answers it → the same neutralisation can be modelled with a particle diagram showing H₃O⁺ and OH⁻ combining, and with a symbolic ionic equation, two different media for the identical reaction.

Every neutralisation reaction can be represented in more than one way, and NESA specifically requires you to be able to model neutralisation using a variety of representations, not just a single memorised molecular equation. Two of the most useful representations are a particle-level diagram and a symbolic ionic equation, and comparing them side by side shows what each medium reveals and what each medium hides.

Representation 1, the particle diagram: A particle diagram shows individual ions and molecules before and after the reaction, rather than formula units. When hydrochloric acid solution (containing H₃O⁺ and Cl⁻ ions) is mixed with sodium hydroxide solution (containing Na⁺ and OH⁻ ions), the diagram shows H₃O⁺ and OH⁻ particles colliding and combining to form neutral H₂O particles, while the Na⁺ and Cl⁻ particles remain unchanged, drifting freely in solution both before and after.

H₃O⁺ and OH⁻ combine to form water, while Na⁺ and Cl⁻ remain as spectator ions.

Particle-level model of neutralisation: only H₃O⁺ and OH⁻ react to form H₂O; Na⁺ and Cl⁻ are spectator ions, unchanged on both sides of the reaction.

Representation 2, the symbolic ionic equation: The same reaction can be written symbolically at three levels of detail. The full molecular equation, HCl(aq) + NaOH(aq) → NaCl(aq) + H₂O(l), shows every formula unit but hides which ions actually react. The full ionic equation, H₃O⁺(aq) + Cl⁻(aq) + Na⁺(aq) + OH⁻(aq) → Na⁺(aq) + Cl⁻(aq) + 2H₂O(l), lists every ion present. The net ionic equation, H₃O⁺(aq) + OH⁻(aq) → 2H₂O(l), cancels the spectator ions (Na⁺ and Cl⁻ appear unchanged on both sides) and shows only the species that actually react.

Representation	What it shows	What it hides	Best used for
Particle diagram	Individual H₃O⁺/OH⁻ combining; spectator ions visibly unchanged	Exact stoichiometric ratios for large numbers of moles	Building conceptual understanding of what "neutralisation" means at the particle level
Full molecular equation	Overall reactants and products, easy to balance	Which species are actually reacting versus spectating	Stoichiometry calculations, mass and mole relationships
Net ionic equation	Only the species that actually react, H₃O⁺(aq) + OH⁻(aq) → 2H₂O(l)	The identity of the salt formed	Explaining why acid + base always gives the same net reaction regardless of the specific acid or base

Must Know: Every strong acid + strong base neutralisation has the identical net ionic equation, H₃O⁺(aq) + OH⁻(aq) → 2H₂O(l), regardless of which specific acid or base is used. The molecular equation changes (different salt forms each time), but the particle diagram and the net ionic equation of a strong acid/strong base pair are always the same, because the spectator ions are the only thing that differs.

Common Error: Students treat the molecular equation as the only valid representation and cannot explain what is physically happening between the particles. When asked to "model" a neutralisation reaction, a molecular equation alone is not a model of the particle-level process, a particle diagram or a net ionic equation is required to show that only H₃O⁺ and OH⁻ actually react while Na⁺ and Cl⁻ (or any other spectator pair) take no part in the reaction.

Neutralisation can be modelled with at least two different representations: a particle diagram (H₃O⁺ and OH⁻ combine to form H₂O; spectator ions like Na⁺ and Cl⁻ are drawn unchanged before and after) and a symbolic ionic equation (full molecular: HCl + NaOH → NaCl + H₂O; net ionic: H₃O⁺(aq) + OH⁻(aq) → 2H₂O(l), spectators cancelled). Every strong acid/strong base pair shares the identical net ionic equation and particle-level picture, only the spectator ions differ.

Pause, copy the highlighted point into your book before moving on.

A particle diagram of HCl(aq) + NaOH(aq) shows H₃O⁺ and OH⁻ combining to form H₂O, while Na⁺ and Cl⁻ are drawn unchanged on both sides. What does this diagram demonstrate that the molecular equation HCl + NaOH → NaCl + H₂O does not show directly?

Consolidate and apply

Real-World Anchor

The Chemistry That Feeds and Protects

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.

Common Misconceptions, L04

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"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.

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"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).

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"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.

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"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

Worked Example 1, Straightforward · Band 3

Writing and Balancing Industrial Equations

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

Worked Example 2, Intermediate · Band 4–5

Evaluating Antacid Chemistry for a Specific Patient

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.

Worked Example 3, Hard · Band 6

Extended Response, Environmental Evaluation of pH Control (7 marks)

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.

✍️ Fill in the Blanks +4 XP

Complete these statements about neutralisation applications:

Antacids treat excess stomach acid by neutralising HCl with a . Agricultural lime (CaCO₃) raises soil pH by reacting with acidic soil components in a reaction. Acid rain (pH < 5.6) forms when SO₂ reacts with to produce H₂SO₃.

Complete the Learn phase to unlock Practice.

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Practice questions

Activities

Industrial Equations: Match, Balance and Evaluate

For each industrial context below, identify the reaction type, write and balance the equation, and answer the analysis question. Show all working.

A power station burns high-sulfur coal. Write the balanced equation for the FGD reaction using Ca(OH)₂ to neutralise SO₂. Identify the acid component and the base component. Explain why CaSO₃ rather than CaSO₄ is the primary product in this reaction.
A fertiliser factory receives a tank of H₂SO₄ and a supply of NH₃ from the Haber process. Write the balanced equation for the fertiliser produced. Explain, using Brønsted-Lowry theory, why 2 mol of NH₃ are required per mole of H₂SO₄. If the factory also produces ammonium nitrate, write that equation and explain why only 1 mol of NH₃ is needed.
A mining company discharges acidic wastewater containing Fe³⁺ ions at pH 3.2. Name one suitable neutralising agent and write the equation for neutralisation of the acid component. Write a second equation showing what happens to Fe³⁺ ions as pH rises. Explain why preventing the discharge of this wastewater at pH 3.2 is environmentally critical.

Check Your Understanding

Multiple Choice

+5 XP

A fresh set drawn from this lesson's question bank, with feedback shown immediately. +5 XP per correct · +25 XP all correct

✍️ Short Answer

Extended Questions

ApplyBand 3

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

UnderstandBand 4

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

EvaluateBand 6

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

✅ Comprehensive Answers

Activity 1, Industrial Equations: Match, Balance and Evaluate

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.

Multiple Choice Answers

Flue gas desulfurisation. The FGD reaction with Ca(OH)₂ targets SO₂ (the primary sulfur combustion product) and produces calcium sulfite (CaSO₃) and water: Ca(OH)₂ + SO₂ → CaSO₃ + H₂O. SO₂ acts as the acidic oxide; Ca(OH)₂ is the base. Using SO₃ (wrong gas) would give CaSO₄; CaCO₃ is a different reagent; and Ca(SO₂)₂ is not a real compound.

Liming acidic soil. CaCO₃ reacts with H⁺ in acidic soil by the acid + carbonate reaction: CaCO₃ + 2H⁺ → Ca²⁺ + H₂O + CO₂. This consumes H⁺, reducing [H⁺] in soil water and raising pH. It is the CO₃²⁻ that reacts, not Ca²⁺; CaCO₃ is not a hydroxide and releases no OH⁻.

Diprotic vs monoprotic fertiliser acids. H₂SO₄ is diprotic (donates 2 H⁺); HNO₃ is monoprotic (donates 1 H⁺). The moles of NH₃ required equals the number of protons the acid can donate, 2 for H₂SO₄ and 1 for HNO₃. The ratio is set by proticity, not by acid strength, and NH₃ accepts only one proton to form NH₄⁺.

Which antacid releases CO₂. CO₂ is produced only when a carbonate reacts with an acid (acid + carbonate → salt + water + CO₂): CaCO₃ + 2HCl → CaCl₂ + H₂O + CO₂. Mg(OH)₂, Al(OH)₃ and NaOH are hydroxide bases with no carbonate, so they produce only salt + water.

Heavy-metal precipitation in wastewater. Ca(OH)₂ reacts with H⁺ in the acidic wastewater to raise pH (neutralisation). As pH rises, [OH⁻] increases and heavy metal ions form insoluble hydroxide precipitates: Fe³⁺ + 3OH⁻ → Fe(OH)₃↓, which can be filtered out. The mechanism is pH-dependent precipitation via OH⁻, not direct reaction with the metal ions.

Short Answer Model Answers

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].

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Retrieve and reflect

Check what actually stuck

Rapid Review

Which two antacids react with HCl without releasing CO₂, and why?

Why does 2NH₃ + H₂SO₄ → (NH₄)₂SO₄ need 2 mol of NH₃, but NH₃ + HNO₃ → NH₄NO₃ needs only 1?

Write the flue gas desulfurisation equation and name the acid and base components.

When Ca(OH)₂ is added to acidic mining wastewater, what happens to dissolved Fe³⁺?

What is the net ionic equation shared by every strong acid + strong base neutralisation?

Revisit Your Initial Thinking

Look back at what you wrote in the Think First section. Recall the 1952 London Great Smog: Ca(OH)₂ + SO₂ → CaSO₃ + H₂O is the reaction that the UK's 1956 Clean Air Act mandated in every industrial furnace. What has changed in your understanding? What did you get right? What surprised you?

You predicted which chemical is used to neutralise SO₂ in flue gas. Was it Ca(OH)₂? Did you predict the correct product (CaSO₃, not CaSO₄)?
Did you correctly identify SO₂ as the acidic component and Ca(OH)₂ as the base? Could you now explain this as an acid + base → salt + water reaction where SO₂ acts as an acidic oxide?
Could you now extend your answer to include: (1) the full chemistry of how SO₂ becomes acid rain if untreated, and (2) the commercial value of the CaSO₃/CaSO₄ byproduct?

I have completed this lesson and reviewed my answers.

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