IGCSE Chemistry Topic 7: Acids, Bases and Salts

7.1 Acids & Bases 7.2 Oxides 7.3 Preparation of Salts

Hey Tara! Welcome to Topic 7 — Acids, Bases and Salts. This is one of the most predictable topics in the whole syllabus, which is brilliant news: the same three reaction patterns come up year after year, and once you can write them from memory you can answer half the questions on sight. The secret is that everything here traces back to a single particle — the hydrogen ion, H⁺. An acid releases it, a base mops it up, and whatever ions are left over pair up to make a salt. Along the way you will meet indicators (chemistry's traffic lights), the pH scale, the difference between a strong acid and a concentrated one (examiners adore this), and four different laboratory methods for making a salt — where the hard part is choosing the right one. Three subtopics, sixty practice questions, and a whole lot of colour changes. Let us get started!

7.1 The Characteristic Properties of Acids and Bases ▼

The Big Idea: It Is All About H⁺

Vinegar, lemon juice, car battery acid and your own stomach fluid look and taste nothing alike — yet chemically they behave in exactly the same way. Why? Because when any of them dissolves in water, the same particle is released: the hydrogen ion, H⁺. Everything an acid does, it does because of that ion.

Acid: a source of H⁺ ions in aqueous solution

An acid is a substance that produces hydrogen ions, H⁺(aq), when dissolved in water. The word "aqueous" matters. Dry hydrogen chloride gas is not acidic; dissolve it in water and it becomes hydrochloric acid.

Base: a substance that neutralises an acid

A base is a metal oxide, metal hydroxide, metal carbonate or ammonia — anything that reacts with an acid to form a salt. An alkali is a base that is soluble in water. It produces hydroxide ions, OH⁻(aq), in solution.

Memory Trick

"All alkalis are bases, but not all bases are alkalis." Copper(II) oxide is a base (it neutralises acid) but it will not dissolve in water, so it is not an alkali. Sodium hydroxide dissolves beautifully, so it is an alkali. Think of alkali as "the dissolved club" — a members-only subset of bases.

What H⁺ Actually Is

A hydrogen atom is one proton plus one electron. Remove the electron and all that is left is a bare proton. So H⁺ and "proton" mean exactly the same thing in acid–base chemistry. In real solutions the proton never floats free — it attaches to a water molecule to form the hydroxonium ion, H₃O⁺ — but at IGCSE you may write it simply as H⁺(aq).

Supplement

The Proton Transfer Definitions

Extended candidates need a sharper pair of definitions:

An acid is a proton (H⁺) donor.
A base is a proton (H⁺) acceptor.

This is a more powerful way of thinking, because it explains reactions where no water is present at all. When hydrogen chloride gas meets ammonia gas, a white smoke of ammonium chloride appears instantly:

HCl(g) + NH₃(g) → NH₄Cl(s)

HCl donates a proton ⇒ it is the acid. NH₃ accepts that proton to become NH₄⁺ ⇒ it is the base. No water, no solution, no pH — but still unmistakably an acid–base reaction.

Notice too that ammonia is a base without containing a single OH⁻ group. It is a base because of what it does (accepts a proton), not because of what it contains.

The Three Reactions of Acids

These three equations are the backbone of the entire topic. Learn them as word patterns first, then practise turning them into balanced symbol equations.

acid + metal → salt + hydrogen

Mg(s) + 2HCl(aq) → MgCl₂(aq) + H₂(g) Zn(s) + H₂SO₄(aq) → ZnSO₄(aq) + H₂(g) Observation: effervescence (fizzing); the metal gets smaller and disappears; the mixture warms up. Test for the gas: a lighted splint gives a squeaky pop.

Only metals above hydrogen in the reactivity series do this. Copper, silver and gold are below hydrogen, so they will not displace it — put a copper coin in dilute hydrochloric acid and absolutely nothing happens. That is a favourite exam trap.

acid + base → salt + water

With a metal oxide: CuO(s) + H₂SO₄(aq) → CuSO₄(aq) + H₂O(l) With a metal hydroxide: NaOH(aq) + HCl(aq) → NaCl(aq) + H₂O(l) Observation with CuO: the black solid disappears and the solution turns blue. With a solution, usually no visible change — only a temperature rise.

This is neutralisation. Because metal oxides and metal hydroxides are both bases, they follow the identical pattern: salt + water, every time, with no gas.

acid + carbonate → salt + water + carbon dioxide

CaCO₃(s) + 2HCl(aq) → CaCl₂(aq) + H₂O(l) + CO₂(g) Na₂CO₃(aq) + H₂SO₄(aq) → Na₂SO₄(aq) + H₂O(l) + CO₂(g) Observation: vigorous effervescence; the solid dissolves. Test for the gas: bubble it through limewater — it turns milky (a white precipitate of CaCO₃).

Hydrogencarbonates behave the same way: NaHCO₃ + HCl → NaCl + H₂O + CO₂. This is why baking soda fizzes in lemon juice.

The Fourth Reaction: Bases and Ammonium Salts

Working in the other direction, a base reacts with an ammonium salt to produce ammonia. This is the standard laboratory test for the ammonium ion:

NH₄Cl + NaOH → NaCl + H₂O + NH₃

Warm any ammonium salt with sodium hydroxide solution. Ammonia gas is given off: pungent smell, and it turns damp red litmus paper blue. Ammonia is the only common alkaline gas, so this test is conclusive.

Neutralisation at the Ionic Level

Strip away the spectator ions and every acid–alkali neutralisation is the same single reaction:

H⁺(aq) + OH⁻(aq) → H₂O(l)

This is the ionic equation for neutralisation. Learn it exactly, including the state symbols. The acidic ion and the alkaline ion destroy each other by forming water, which is neutral.

Take NaOH(aq) + HCl(aq) → NaCl(aq) + H₂O(l). Written out in full ions: Na⁺ + OH⁻ + H⁺ + Cl⁻ → Na⁺ + Cl⁻ + H₂O. The Na⁺ and Cl⁻ appear unchanged on both sides — they are spectator ions. Cancel them and only H⁺ + OH⁻ → H₂O survives.

Exam Tip

Neutralisation is always exothermic. Every acid–alkali mixture warms up, because forming the O–H bonds in water releases energy. If a question gives you a temperature-against-volume graph for a titration, the peak temperature marks the exact neutralisation point.

Indicators: Chemistry's Traffic Lights

An indicator is a dye that is one colour in acid and a different colour in alkali. You must know these three by heart. Note that only litmus has a distinct third colour for neutral solutions.

Indicator	In acid	In neutral solution	In alkali
Litmus	Red	Purple	Blue
Thymolphthalein	Colourless	Colourless	Blue
Methyl orange	Red	Orange	Yellow
Universal indicator	Red / orange / yellow	Green	Blue / indigo / violet
Phenolphthalein (useful extra)	Colourless	Colourless	Pink

Memory Trick

Litmus: "Blue to red, acid ahead." Methyl orange: the name tells you the middle colour is orange, so it must go red one way and yellow the other — and red is always the acid end, just like litmus. Thymolphthalein: starts with "thy" — think sky blue in alkali, invisible in acid.

Exam Tip

A single indicator tells you acid or alkali, not how acidic. Litmus turns red in pH 6 lemon-ish water and equally red in pH 1 battery acid. If a question asks you to compare the strength or acidity of two solutions, you must use universal indicator or a pH meter — naming litmus there loses the mark.

The pH Scale

The pH scale runs from 0 to 14 and measures how acidic or alkaline a solution is. It is really a measure of the concentration of H⁺ ions: the more H⁺ ions per dm³, the lower the pH.

pH below 7 — acidic. The lower the number, the more acidic.
pH exactly 7 — neutral (pure water, sodium chloride solution).
pH above 7 — alkaline. The higher the number, the more alkaline.

Universal indicator gives a continuous colour range: red through orange and yellow for acids, green at neutral, and blue through indigo to violet for alkalis.

pH range	Universal indicator colour	Description	Typical example
0–2	Red	Strongly acidic	Hydrochloric acid, stomach acid
3–4	Orange	Moderately acidic	Vinegar, orange juice
5–6	Yellow	Weakly acidic	Rainwater, black coffee, milk
7	Green	Neutral	Pure water, sodium chloride solution
8–9	Blue-green / blue	Weakly alkaline	Sodium hydrogencarbonate solution
10–12	Blue / indigo	Moderately alkaline	Limewater, milk of magnesia
13–14	Violet / purple	Strongly alkaline	Sodium hydroxide, oven cleaner

Exam Tip

Universal indicator is a mixture of several dyes, which is why it produces a range of colours rather than a single switch. Also remember the practical detail: you add a few drops to the solution, or dip a piece of UI paper in, then compare against a colour chart. Never write "I would look at the colour" without mentioning the chart.

Supplement

Strong and Weak Acids

This is the single most misunderstood idea in Topic 7, and examiners know it. The words strong and weak describe how completely an acid splits up into ions in water. They have nothing whatsoever to do with how much acid is in the bottle.

Strong acid: completely dissociated

HCl(aq) → H⁺(aq) + Cl⁻(aq) A single arrow. Every HCl molecule splits up, so the H⁺ concentration is as high as it can possibly be. Other strong acids: sulfuric acid (H₂SO₄) and nitric acid (HNO₃).

Weak acid: partially dissociated

CH₃COOH(aq) ⇌ H⁺(aq) + CH₃COO⁻(aq) A reversible arrow. Only a small proportion of the molecules split up at any moment — typically about 1 in 100 — so most stay whole. Other weak acids: carbonic acid (H₂CO₃), citric acid, and the sulfurous acid in acid rain.

Same concentration, same volume — but the strong acid has released every possible H⁺ ion while the weak acid keeps most of its hydrogen locked inside intact molecules.

Comparing 1.0 mol/dm³ Hydrochloric Acid with 1.0 mol/dm³ Ethanoic Acid

This exact comparison appears again and again. Both solutions have the same concentration, so any difference must come from the strength.

Property	Hydrochloric acid (strong)	Ethanoic acid (weak)	Reason
pH	About 0–1 (lower)	About 2–3 (higher)	More H⁺(aq) per dm³ in the strong acid
Rate of reaction with magnesium	Fast, vigorous fizzing	Slower, gentler fizzing	Higher H⁺ concentration ⇒ more frequent successful collisions
Electrical conductivity	Good	Poor	More mobile ions available to carry the charge
Volume of gas produced in total	The same	The same	Same moles of acid overall — the weak acid dissociates further as H⁺ is used up
Volume of alkali needed in a titration	The same	The same	Same moles of acid to be neutralised

Exam Tip — The Killer Distinction

Strong / weak = how completely the acid dissociates. It is a fixed property of the substance. You cannot make ethanoic acid strong.

Concentrated / dilute = how much acid is dissolved in a given volume of water. You control it with a measuring cylinder.

The two are independent. Concentrated ethanoic acid exists (glacial ethanoic acid, 17 mol/dm³) and it is still a weak acid. Dilute hydrochloric acid exists (0.01 mol/dm³) and it is still a strong acid. Never write "strong" when you mean "concentrated".

Memory Trick

Think of a football crowd. Strength = what fraction of the crowd is actually singing. Concentration = how many people are in the stadium. A packed stadium of quiet people (concentrated weak acid) can be quieter than a half-empty stadium of enthusiastic singers (dilute strong acid).

Weak Bases

The same logic works for bases. Sodium hydroxide is a strong base — it is fully dissociated into Na⁺ and OH⁻ ions, giving pH 14. Aqueous ammonia is a weak base — only a few molecules react with water to form NH₄⁺ and OH⁻, so a solution of the same concentration reaches only about pH 11.

NH₃(aq) + H₂O(l) ⇌ NH₄⁺(aq) + OH⁻(aq)

The reversible arrow tells you at a glance that ammonia is a weak base.

Worked Examples

Worked Example 1 Write balanced symbol equations, including state symbols, for: (a) zinc with dilute sulfuric acid, (b) magnesium oxide with dilute hydrochloric acid, (c) sodium carbonate solution with dilute nitric acid.

Step 1: Identify the pattern for each

(a) acid + metal → salt + hydrogen. (b) acid + base (metal oxide) → salt + water. (c) acid + carbonate → salt + water + carbon dioxide. Get the pattern right and the products write themselves.

Step 2: Name the salt from the acid

Sulfuric acid gives a sulfate (SO₄²⁻), hydrochloric acid gives a chloride (Cl⁻), nitric acid gives a nitrate (NO₃⁻). Combine each with the metal ion, balancing the charges: Zn²⁺ + SO₄²⁻ → ZnSO₄; Mg²⁺ + 2Cl⁻ → MgCl₂; 2Na⁺ + NO₃⁻ → NaNO₃.

Step 3: Balance and add state symbols

(a) Zn(s) + H₂SO₄(aq) → ZnSO₄(aq) + H₂(g) — already balanced.
(b) MgO(s) + 2HCl(aq) → MgCl₂(aq) + H₂O(l) — two HCl are needed to supply two Cl.
(c) Na₂CO₃(aq) + 2HNO₃(aq) → 2NaNO₃(aq) + H₂O(l) + CO₂(g) — two nitric acids for two sodium ions.

(a) Zn(s) + H₂SO₄(aq) → ZnSO₄(aq) + H₂(g)
(b) MgO(s) + 2HCl(aq) → MgCl₂(aq) + H₂O(l)
(c) Na₂CO₃(aq) + 2HNO₃(aq) → 2NaNO₃(aq) + H₂O(l) + CO₂(g)

Worked Example 2 Rohan tests four colourless solutions W, X, Y and Z with universal indicator. W turns it green, X turns it red, Y turns it violet and Z turns it yellow. Rank them from most acidic to most alkaline, give an approximate pH for each, and predict what happens when a piece of magnesium ribbon is dropped into each.

Step 1: Convert colours into pH values

Red = strongly acidic, pH 0–2. Yellow = weakly acidic, pH 5–6. Green = neutral, pH 7. Violet = strongly alkaline, pH 13–14.

Step 2: Rank them

Most acidic to most alkaline: X (pH ~1) → Z (pH ~5) → W (pH 7) → Y (pH ~13).

Step 3: Predict the magnesium reactions

Magnesium reacts with acids to give hydrogen. In X there is rapid, vigorous effervescence and the ribbon dissolves quickly. In Z there is slow, gentle effervescence because the H⁺ concentration is far lower. In W (neutral) and Y (alkaline) there is essentially no reaction — magnesium does not react with alkalis under IGCSE conditions.

Order: X (most acidic, pH ~1) < Z (pH ~5) < W (pH 7) < Y (most alkaline, pH ~13). Magnesium fizzes vigorously in X, slowly in Z, and does not react in W or Y.

Worked Example 3 Two beakers each contain 50 cm³ of acid of concentration 0.10 mol/dm³. Beaker A holds hydrochloric acid, beaker B holds ethanoic acid. Identical 5 cm strips of magnesium are added to each. Describe and explain the differences and the similarities in what is observed.

Step 1: Establish what is the same

Same volume, same concentration, therefore the same number of moles of acid in each beaker (0.050 dm³ × 0.10 mol/dm³ = 0.005 mol). Both acids are monoprotic (one replaceable H per molecule).

Step 2: Establish what is different

Hydrochloric acid is strong: fully dissociated, so essentially all 0.005 mol of H⁺ is free in solution at once. Ethanoic acid is weak: only a small fraction of the molecules dissociate at any moment, so the concentration of free H⁺ is much lower.

Step 3: Predict the rate

A higher concentration of H⁺ gives more frequent successful collisions with the magnesium surface. So beaker A fizzes faster and more vigorously and the ribbon disappears sooner. Beaker B fizzes noticeably more slowly.

Step 4: Predict the total volume of gas — the mark most students lose

Assuming the magnesium is in excess of neither and the acid is the limiting reactant, the total volume of hydrogen is identical. As H⁺ is consumed in beaker B, the equilibrium CH₃COOH ⇌ H⁺ + CH₃COO⁻ shifts to the right to replace it, so eventually every acid molecule releases its proton. Weak does not mean incomplete — it means slower.

Beaker A (HCl) reacts faster with more vigorous effervescence because its H⁺ concentration is higher (fully dissociated). Beaker B (ethanoic) reacts more slowly. However, both eventually produce the same total volume of hydrogen, because both contain the same number of moles of acid and the weak acid dissociates further as the reaction proceeds.

Worked Example 4 A bottle is labelled "concentrated ethanoic acid, 17 mol/dm³". A student writes: "This must be a strong acid because it is so concentrated, so it will have a lower pH than 1 mol/dm³ hydrochloric acid." Identify every error and write a corrected statement.

Error 1: Confusing strength with concentration

"Concentrated" describes the number of moles of solute per dm³ of solution. "Strong" describes the degree of dissociation. Ethanoic acid is always a weak acid, at any concentration, because the –COOH group only partially ionises.

Error 2: Assuming the conclusion follows

Being concentrated does not make a substance strong. The two properties are independent variables.

Step 3: What about the pH claim?

Interestingly, the pH claim happens to be true — but for the wrong reason. 17 mol/dm³ ethanoic acid contains so many molecules that even a 1% dissociation gives a higher H⁺ concentration than 1 mol/dm³ HCl. The pH is low because of the huge concentration, not because of any strength. A correct answer must use the right reason to earn the mark.

Corrected: "Ethanoic acid is a weak acid at every concentration, because it is only partially dissociated. This bottle is concentrated, meaning a large number of moles per dm³. Its pH is low because of the very high concentration, not because it is strong."

Exam Tips for 7.1

1. Learn the three acid reactions as a chant: "metal → salt + hydrogen; base → salt + water; carbonate → salt + water + carbon dioxide." You will use them in Topic 7, Topic 9 and Topic 10.

2. Say "aqueous". The definition is "a source of H⁺ ions in aqueous solution". Leaving out those two words is a common lost mark, because dry HCl gas is not acidic.

3. Do not confuse base with alkali. Copper(II) oxide is a base but not an alkali. Use "alkali" only when the base is dissolved in water.

4. Colour changes need direction. Write "litmus turns from blue to red", not just "litmus is red". Examiners want the change, not the end state.

5. Strong ≠ concentrated. Cambridge sets at least one question a year on this. Strong = fully dissociated. Concentrated = lots of moles per dm³.

6. Rate differs, total amount does not. A weak acid reacts more slowly than a strong acid of the same concentration, but produces exactly the same total volume of gas and needs exactly the same volume of alkali to neutralise it.

7. Learn the ionic equation. H⁺(aq) + OH⁻(aq) → H₂O(l). It is worth a whole mark on its own and takes ten seconds to memorise.

8. Gas tests. Hydrogen → squeaky pop with a lighted splint. Carbon dioxide → limewater turns milky. Ammonia → damp red litmus turns blue. Say damp — dry litmus paper does not respond to a gas.

🌎 Apply It: Real-World Chemistry

Acid–base chemistry is quietly running your stomach, your farmland, your bathroom cleaner and the ancient marble of half the world's monuments.

A pharmacist in Mumbai recommends a chewable antacid tablet containing magnesium hydroxide and calcium carbonate to a customer with acid reflux. The customer notices the tablet makes him burp, and asks whether that means the tablet has gone off.

Explain the chemistry of the tablet, write two equations, and explain the burping.

▼

The Chemistry

Stomach acid is hydrochloric acid at about pH 1.5. Both ingredients are bases, so both neutralise it. Mg(OH)₂ + 2HCl → MgCl₂ + 2H₂O (base + acid → salt + water). CaCO₃ + 2HCl → CaCl₂ + H₂O + CO₂ (carbonate + acid → salt + water + carbon dioxide).

Why He Burps

The carbonate reaction releases carbon dioxide gas inside the stomach. There is nowhere for it to go except back up the oesophagus. Nothing has gone off — the burp is direct evidence that the neutralisation is working.

Chemistry Connection

Manufacturers choose weak or insoluble bases deliberately. A strong alkali such as sodium hydroxide would neutralise the acid brilliantly and then burn a hole in the stomach lining. Magnesium hydroxide is only slightly soluble, so it cannot push the pH much above about 9 — it self-limits. Chemistry chosen for safety, not just for effectiveness.

A tea planter in the Nilgiri hills tests his soil and finds a pH of 4.2. His agricultural officer tells him to spread powdered limestone (calcium carbonate) — but warns him not to use quicklime (calcium oxide) even though it works faster, and not to add fertiliser on the same day.

Explain the chemistry behind all three pieces of advice.

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Why Liming Works

Calcium carbonate is a base. It neutralises the excess H⁺ ions in the soil water: CaCO₃ + 2H⁺ → Ca²⁺ + H₂O + CO₂. The pH rises towards the 6–7 range where most crops absorb nutrients best. The fizzing you see is the carbon dioxide.

Why Not Quicklime

Calcium oxide is a far more reactive base and dissolves to give calcium hydroxide, a strong alkali. It can overshoot to pH 10 or higher, which locks up iron and manganese and scorches roots. Calcium carbonate is insoluble and reacts only as fast as acid is available, so it is self-limiting — the same design logic as the antacid tablet.

Why Not With Fertiliser

Most nitrogen fertilisers are ammonium salts. A base + ammonium salt reaction releases ammonia gas: NH₄⁺ + OH⁻ → NH₃ + H₂O. The expensive nitrogen literally floats away. Lime first, wait a few weeks, then fertilise.

Chemistry Connection

Three different syllabus reactions in one field: acid + carbonate, the meaning of a strong base, and base + ammonium salt → ammonia. Notice that the last one is the same reaction you use in the laboratory to test for the ammonium ion — the pungent smell over a freshly limed and fertilised field is a giant version of the test tube.

Conservators studying the marble of the Taj Mahal in Agra and the limestone of St Paul's Cathedral in London report the same damage: surfaces that were once sharply carved are now blurred and crumbling, and rainwater running off them is slightly acidic and contains dissolved calcium.

Explain the chemistry, and explain why a marble statue erodes faster than a granite one standing beside it.

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The Chemistry

Marble and limestone are both calcium carbonate. Sulfur dioxide and nitrogen oxides from burning fuels dissolve in rain to form sulfuric and nitric acids. This is the classic acid + carbonate reaction: CaCO₃ + H₂SO₄ → CaSO₄ + H₂O + CO₂. Even unpolluted rain is about pH 5.6 from dissolved CO₂ forming weak carbonic acid.

Why Marble and Not Granite

Granite is mostly silicates — quartz and feldspar — which are not carbonates and are not bases, so acids do not attack them. Only carbonate rock reacts. This is exactly why a fresh drop of dilute acid on an unknown rock, producing fizzing that turns limewater milky, identifies it as a carbonate.

Chemistry Connection

The calcium sulfate formed is slightly soluble and flakes away, taking the carved surface with it — so the damage is permanent, not a coating that can be cleaned off. Cambridge questions on acid rain almost always want two things: the named acid formed, and the equation with the carbonate.

A cleaning product manufacturer in Manchester sells two bathroom sprays. The limescale remover is labelled "contains ethanoic acid"; the oven cleaner is labelled "contains sodium hydroxide — corrosive". Both bottles carry a warning never to mix them, and the limescale remover is not recommended for marble worktops.

Explain the choice of chemical in each product and both warnings.

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Why Ethanoic Acid for Limescale

Limescale is calcium carbonate deposited from hard water. An acid dissolves it: CaCO₃ + 2CH₃COOH → (CH₃COO)₂Ca + H₂O + CO₂. A weak acid is chosen on purpose — it is strong enough to remove scale over a few minutes but gentle enough not to attack chrome, enamel or skin. Hydrochloric acid would work in seconds and then keep going.

Why Sodium Hydroxide for Ovens

Baked-on grease is fat. A strong alkali breaks fats down into soluble soaps — the same reaction used to make soap industrially. Only a strong alkali does this quickly, which is why the product is corrosive and needs gloves.

The Two Warnings

Mixing them causes a vigorous exothermic neutralisation (H⁺ + OH⁻ → H₂O) that can spit hot liquid, and it wastes both products by converting them into a useless salt solution. And marble is calcium carbonate, so the limescale remover would not tell the difference between the scale and the worktop — it would etch both.

Chemistry Connection

Notice the design principle repeating across all these scenarios: where you need control, use a weak acid or base; where you need speed and can manage the hazard, use a strong one. The strong/weak distinction is not academic trivia — it is a product-safety decision.

A swimmer at a public pool in Singapore complains of stinging eyes. The pool technician measures the water at pH 8.4, well above the target range of 7.2–7.6, and adds a measured dose of sodium hydrogensulfate. An hour later the pH reads 7.4. A colleague asks why they do not just use hydrochloric acid, which is cheaper.

Explain what the technician has done chemically, and evaluate the colleague's suggestion.

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The Chemistry

The pool has drifted alkaline. Sodium hydrogensulfate dissolves to release H⁺ ions, which neutralise the excess OH⁻ ions: H⁺ + OH⁻ → H₂O. Removing OH⁻ lowers the pH back into the comfortable range. Human tears sit at about pH 7.4, which is exactly why that target is chosen.

Evaluating the Alternative

Hydrochloric acid would do the same job chemically and is cheaper, but it is a strong acid supplied as a corrosive liquid. Transporting, storing and pouring it around a public pool is far more hazardous, and a single overdose could crash the pH into the acidic range within minutes. Sodium hydrogensulfate is a solid that dissolves gradually, making it easier to dose accurately and much safer to handle.

Chemistry Connection

This is a perfect illustration of why the pH scale is logarithmic. Moving from 8.4 to 7.4 means increasing the H⁺ concentration by a factor of ten — a much bigger chemical change than the small numerical difference suggests. Exam questions often ask "how many times more acidic is pH 3 than pH 6?" The answer is 10 × 10 × 10 = 1000 times.

Practice Questions: 7.1

20 multiple choice questions. Click an option to check your answer.

Your Score 0 / 20

Question 1

Which statement correctly defines an acid?

A A substance that contains hydrogen

B A substance that produces hydrogen ions in aqueous solution

C A substance that produces hydroxide ions in aqueous solution

D A substance with a sour taste that dissolves metals

The definition requires H⁺ ions in aqueous solution. Option A is wrong because methane and ammonia contain hydrogen but are not acids. Option C defines an alkali. Option D describes properties, not a definition.

Question 2

Which of these substances is a base but not an alkali?

A Sodium hydroxide

B Potassium hydroxide

C Copper(II) oxide

D Aqueous ammonia

An alkali is a base that dissolves in water. Copper(II) oxide neutralises acids (so it is a base) but is insoluble, so it is not an alkali. The other three all dissolve to give OH⁻(aq).

Question 3

Magnesium is added to dilute sulfuric acid. What are the products?

A Magnesium sulfate + water

B Magnesium sulfate + hydrogen

C Magnesium sulfate + water + carbon dioxide

D Magnesium sulfide + hydrogen

Acid + metal → salt + hydrogen. Sulfuric acid always gives a sulfate (SO₄²⁻), never a sulfide (S²⁻) — a very common careless slip.

Question 4

A gas produced by adding acid to a solid turns limewater milky. The solid contained which ion?

A Sulfate

B Carbonate

C Nitrate

D Ammonium

Limewater turning milky is the standard test for carbon dioxide, and CO₂ is released when an acid reacts with a carbonate or hydrogencarbonate.

Question 5

What is the ionic equation for the neutralisation of any acid by any alkali?

A H₂ + O → H₂O

B H⁺(aq) + OH⁻(aq) → H₂O(l)

C H⁺(aq) + O²⁻(aq) → H₂O(l)

D Na⁺(aq) + Cl⁻(aq) → NaCl(aq)

The hydrogen ion from the acid and the hydroxide ion from the alkali combine to form water. Everything else is a spectator ion. Option D is the spectator ion pair, which is exactly what you cancel out.

Question 6

Litmus paper is placed in a solution of pH 12. What colour is observed?

A Red

B Purple

C Blue

D Yellow

pH 12 is alkaline and litmus is blue in alkali. Yellow would be methyl orange in alkali; purple is litmus in a neutral solution.

Question 7

Methyl orange is added to dilute hydrochloric acid. What is seen?

A A red colour

B A yellow colour

C It stays colourless

D A blue colour

Methyl orange is red in acid, orange at neutral and yellow in alkali. Colourless-in-acid describes thymolphthalein and phenolphthalein.

Question 8

Thymolphthalein is added to aqueous sodium hydroxide. What is the colour?

A Colourless

B Pink

C Blue

D Red

Thymolphthalein is colourless in acid and neutral solutions and blue in alkali. Pink in alkali is phenolphthalein — do not mix the two names up.

Question 9

Universal indicator turns green in a solution. Which statement must be true?

A The solution is pure water

B The solution has a pH of about 7

C The solution contains no ions

D The solution will not conduct electricity

Green means neutral, pH 7. But a neutral solution need not be pure water — sodium chloride solution is pH 7, full of ions, and conducts electricity well.

Question 10

How many times more concentrated in H⁺ ions is a solution of pH 2 compared with one of pH 5?

A 3 times

B 30 times

C 100 times

D 1000 times

Each unit on the pH scale is a ten-fold change. Three units means 10 × 10 × 10 = 1000 times. The pH scale is logarithmic, not linear.

Question 11

Which pair of statements about hydrochloric acid and ethanoic acid of the same concentration is correct?

A HCl has the lower pH and the higher electrical conductivity

B HCl has the higher pH and the higher electrical conductivity

C HCl has the lower pH and the lower electrical conductivity

D Both have the same pH and conductivity

HCl is fully dissociated, so it has more free ions: more H⁺ means a lower pH, and more mobile ions overall means better conductivity.

Question 12

Excess magnesium is added to 50 cm³ of 1.0 mol/dm³ hydrochloric acid and separately to 50 cm³ of 1.0 mol/dm³ ethanoic acid. Which statement about the total volume of hydrogen is correct?

A The hydrochloric acid produces more, because it is stronger

B The ethanoic acid produces more, because it has more hydrogen atoms

C Both produce the same volume, because both contain the same moles of acid

D The ethanoic acid produces none, because it is too weak to react

This is the classic trap. Weak means slower, not less. As H⁺ is used up, the equilibrium CH₃COOH ⇌ H⁺ + CH₃COO⁻ shifts right, so the weak acid eventually delivers every proton. Same moles of acid ⇒ same volume of hydrogen.

Question 13

Which statement about a concentrated solution of a weak acid is correct?

A It cannot exist — weak acids can only be dilute

B It contains many moles of acid per dm³, but only a small proportion are dissociated

C It is the same thing as a dilute solution of a strong acid

D Concentrating a weak acid turns it into a strong acid

Strength and concentration are independent. Glacial ethanoic acid is about 17 mol/dm³ — extremely concentrated — and still a weak acid, because the proportion of molecules that ionise is unchanged.

Question 14

In the reaction HCl(g) + NH₃(g) → NH₄Cl(s), which statement is correct?

A Neither substance is acting as an acid because there is no water

B HCl is the proton donor and NH₃ is the proton acceptor

C NH₃ is the proton donor and HCl is the proton acceptor

D Both substances donate protons

The proton-transfer definitions work without water. HCl gives away H⁺ (acid, proton donor); ammonia accepts it to become NH₄⁺ (base, proton acceptor). The product is the white smoke of ammonium chloride.

Question 15

An unknown white solid is warmed with sodium hydroxide solution. A pungent gas is given off which turns damp red litmus paper blue. The solid contained which ion?

A Carbonate

B Ammonium

C Nitrate

D Chloride

Base + ammonium salt → ammonia. Ammonia is the only common alkaline gas, so turning damp red litmus blue identifies it — and therefore identifies the ammonium ion in the original solid.

Question 16

A copper coin is placed in dilute hydrochloric acid. What is observed?

A Rapid effervescence and the coin dissolves

B Slow effervescence and the solution turns blue

C No visible reaction

D A white precipitate forms

Copper is below hydrogen in the reactivity series, so it cannot displace hydrogen from an acid. Only metals above hydrogen give the acid + metal reaction.

Question 17

Which observation shows that neutralisation is exothermic?

A The indicator changes colour

B The temperature of the mixture rises

C A gas is given off

D The solution becomes colourless

Exothermic means energy is released to the surroundings, so the temperature rises. In a titration, the temperature peaks exactly at the neutralisation point.

Question 18

Which row correctly matches an acid to the gas produced when it reacts with a named substance?

A HCl + copper(II) oxide → hydrogen

B HNO₃ + calcium carbonate → carbon dioxide

C H₂SO₄ + sodium hydroxide → hydrogen

D HCl + zinc → carbon dioxide

Only carbonates give carbon dioxide, and only metals above hydrogen give hydrogen. Acid + oxide and acid + hydroxide give salt and water only — no gas at all.

Question 19

Which of these is the best method to compare the acidity of two different acid solutions?

A Add blue litmus paper to each

B Add methyl orange to each

C Measure the pH of each with a pH meter

D Add thymolphthalein to each

Single indicators only tell you acid or alkali — litmus is equally red at pH 6 and pH 1. To compare degrees of acidity you need universal indicator or, better, a pH meter, which gives a numerical value.

Question 20

Aqueous ammonia of concentration 0.1 mol/dm³ has a pH of about 11, while 0.1 mol/dm³ sodium hydroxide has a pH of about 13. The best explanation is:

A Ammonia is more dilute than the sodium hydroxide

B Ammonia is a weak base, so it is only partially dissociated and produces fewer OH⁻ ions

C Ammonia contains no hydroxide group so it cannot be a base at all

D Sodium hydroxide contains a metal, which always makes solutions more alkaline

Both have the same concentration, so the difference must come from strength. NH₃ + H₂O ⇌ NH₄⁺ + OH⁻ is reversible and lies to the left, so fewer OH⁻ ions are released — a lower pH than the fully dissociated NaOH.

7.2 Oxides ▼

The Big Idea: Sort Oxides by What They React With

An oxide is a compound of an element with oxygen. There are hundreds of them, but for the exam you only need to sort them into four boxes. The sorting rule is beautifully simple: what does it react with?

Type of oxide	Reacts with acids?	Reacts with alkalis?	Typically formed by
Basic	Yes ✓	No ✗	Metals
Acidic	No ✗	Yes ✓	Non-metals
Amphoteric	Yes ✓	Yes ✓	A few metals near the metal/non-metal border
Neutral	No ✗	No ✗	A handful of non-metals

Memory Trick

Read the table diagonally. Metal oxides are basic; non-metal oxides are acidic. That single sentence handles the vast majority of examples. Then learn the two short exception lists: amphoteric (ZnO and Al₂O₃) and neutral (CO, N₂O, NO, H₂O). Everything else falls into place.

Basic Oxides

A basic oxide is a metal oxide that reacts with an acid to form a salt and water. It is a base, so it follows the neutralisation pattern exactly.

basic oxide + acid → salt + water

CuO(s) + 2HNO₃(aq) → Cu(NO₃)₂(aq) + H₂O(l) — black solid dissolves, blue solution forms MgO(s) + H₂SO₄(aq) → MgSO₄(aq) + H₂O(l) CaO(s) + 2HCl(aq) → CaCl₂(aq) + H₂O(l)

Common basic oxides: Na₂O, K₂O, CaO, MgO, CuO, FeO, Fe₂O₃, PbO (though lead oxides are also amphoteric in practice — at IGCSE stick with ZnO and Al₂O₃ as the amphoteric examples).

The oxides of the most reactive metals — sodium, potassium and calcium — also dissolve in water to give alkaline solutions:

CaO(s) + H₂O(l) → Ca(OH)₂(aq)

Calcium oxide (quicklime) reacts vigorously with water — the mixture steams and gets hot. The product, calcium hydroxide, is limewater. pH about 12. Soluble basic oxides therefore produce alkalis. Insoluble ones (CuO, Fe₂O₃) are still basic but do not change the pH of water.

Acidic Oxides

An acidic oxide is a non-metal oxide that reacts with a base or alkali to form a salt and water. Many of them also dissolve in water to give acidic solutions.

acidic oxide + alkali → salt + water

CO₂(g) + 2NaOH(aq) → Na₂CO₃(aq) + H₂O(l) SO₂(g) + 2NaOH(aq) → Na₂SO₃(aq) + H₂O(l) Note the pattern: still salt + water, exactly like ordinary neutralisation.

Dissolved in water, acidic oxides form acids:

Acidic oxide	Reaction with water	Acid formed	Where it matters
CO₂ (carbon dioxide)	CO₂ + H₂O ⇌ H₂CO₃	Carbonic acid (weak)	Fizzy drinks; natural rain at pH 5.6; ocean acidification
SO₂ (sulfur dioxide)	SO₂ + H₂O → H₂SO₃	Sulfurous acid	Acid rain from burning coal and sulfur-rich fuels
SO₃ (sulfur trioxide)	SO₃ + H₂O → H₂SO₄	Sulfuric acid (strong)	The final step of the Contact process
NO₂ (nitrogen dioxide)	Dissolves to give a mixture	Nitric and nitrous acids	Acid rain from vehicle engines and lightning
P₄O₁₀ (phosphorus(V) oxide)	P₄O₁₀ + 6H₂O → 4H₃PO₄	Phosphoric acid	Fertiliser and cola manufacture

Exam Tip

Do not confuse "acidic oxide" with "acid". Sulfur dioxide is a gas, not an acid — it becomes an acid only when it meets water. The examiner wants the phrase "it dissolves in water to form an acidic solution", or "it reacts with alkalis to form a salt and water".

Supplement

Amphoteric Oxides

An amphoteric oxide reacts with both acids and alkalis (bases) to form a salt and water. It sits on the fence. The two you must know are:

Zinc oxide, ZnO
Aluminium oxide, Al₂O₃

Zinc oxide behaving as a base

ZnO(s) + 2HCl(aq) → ZnCl₂(aq) + H₂O(l) Here it neutralises an acid — ordinary basic behaviour. The white solid dissolves to give a colourless solution.

Zinc oxide behaving as an acid

ZnO(s) + 2NaOH(aq) → Na₂ZnO₂(aq) + H₂O(l) The salt formed is sodium zincate. The white solid dissolves in hot concentrated alkali — behaviour no ordinary metal oxide shows.

Aluminium oxide does both too

With acid: Al₂O₃(s) + 6HCl(aq) → 2AlCl₃(aq) + 3H₂O(l) With alkali: Al₂O₃(s) + 2NaOH(aq) → 2NaAlO₂(aq) + H₂O(l), forming sodium aluminate

Two tests, four outcomes. This decision tree is the whole of 7.2 in one picture — and it is exactly how an examiner will set an "identify the oxide" question.

Exam Tip

The word amphoteric is worth a mark on its own, but only if you back it up. Full-credit phrasing: "Zinc oxide is amphoteric because it reacts with acids AND with alkalis, forming a salt and water in each case." Writing "it is both an acid and a base" is imprecise — ZnO is not an acid, it behaves as an acidic oxide towards alkalis.

Neutral Oxides

A neutral oxide reacts with neither acids nor bases, and dissolving it in water does not change the pH. The list is short, so just learn it:

Carbon monoxide, CO — the toxic gas from incomplete combustion
Dinitrogen oxide, N₂O — "laughing gas", used as an anaesthetic
Nitrogen monoxide, NO — formed in car engines and by lightning
Water, H₂O — technically hydrogen oxide, and famously neutral at pH 7

Memory Trick

Notice that CO is neutral but CO₂ is acidic, and NO is neutral but NO₂ is acidic. The extra oxygen makes all the difference. Remember it as: "one oxygen sits on the fence, two oxygens turn sour." Examiners love offering CO as a distractor in an acidic-oxide question.

The Periodic Trend Across a Period

Because metals are on the left of the Periodic Table and non-metals on the right, the character of the oxides changes steadily as you cross a period:

Element (Period 3)	Oxide	Metal or non-metal?	Classification
Sodium	Na₂O	Metal	Basic (dissolves to give a strong alkali)
Magnesium	MgO	Metal	Basic (slightly soluble, weakly alkaline)
Aluminium	Al₂O₃	Metal (borderline)	Amphoteric
Silicon	SiO₂	Non-metal (borderline)	Acidic (but insoluble in water)
Phosphorus	P₄O₁₀	Non-metal	Acidic
Sulfur	SO₂, SO₃	Non-metal	Acidic

The trend is basic → amphoteric → acidic from left to right, with the amphoteric oxide sitting exactly at the metal/non-metal borderline. That is not a coincidence — it is the whole reason aluminium oxide behaves the way it does.

Worked Examples

Worked Example 1 Classify each of these oxides as acidic, basic, amphoteric or neutral, giving your reasoning: Na₂O, SO₂, ZnO, CO, Fe₂O₃.

Step 1: Check the exception lists first

ZnO is one of the two amphoteric oxides you must know. CO is on the neutral list. Deal with the exceptions before applying the general rule, or you will misclassify them.

Step 2: Apply metal / non-metal to the rest

Sodium is a metal ⇒ Na₂O is basic, and it is soluble so it gives an alkaline solution. Sulfur is a non-metal ⇒ SO₂ is acidic, dissolving to give sulfurous acid. Iron is a metal ⇒ Fe₂O₃ is basic, though insoluble in water.

Step 3: Justify the exceptions

ZnO is amphoteric because it reacts with acids (ZnO + 2HCl → ZnCl₂ + H₂O) and with alkalis (ZnO + 2NaOH → Na₂ZnO₂ + H₂O). CO is neutral because it reacts with neither.

Na₂O = basic; SO₂ = acidic; ZnO = amphoteric; CO = neutral; Fe₂O₃ = basic.

Worked Example 2 A student has a white powder that could be magnesium oxide or zinc oxide. Describe a laboratory procedure that would distinguish them, including expected observations and equations.

Step 1: Identify the property that differs

Both are white solids and both are basic, so both dissolve in dilute acid — that test is useless on its own. The difference is that zinc oxide is amphoteric, so it also reacts with alkalis, whereas magnesium oxide does not.

Step 2: Design the test

Place a spatula of each powder into a separate test tube. Add a few cm³ of hot concentrated aqueous sodium hydroxide to each and warm gently.

Step 3: State the observations

The zinc oxide dissolves to give a colourless solution: ZnO + 2NaOH → Na₂ZnO₂ + H₂O. The magnesium oxide remains undissolved as a white solid, because MgO is purely basic and does not react with alkalis.

Step 4: Add a confirmatory test

Repeat with dilute hydrochloric acid: both dissolve, confirming both are basic oxides and that the alkali result was genuinely due to amphoteric character rather than the powder simply being soluble.

Warm each powder with aqueous sodium hydroxide. The one that dissolves is zinc oxide (amphoteric, forming sodium zincate); the one that stays as a white solid is magnesium oxide (basic only).

Worked Example 3 Sulfur dioxide released from a smelter in Zambia is removed from the waste gases by passing them through a spray of calcium hydroxide solution. (a) Explain why this works in terms of oxide classification. (b) Write an equation. (c) Explain why the same treatment would not remove carbon monoxide.

Part (a): The classification

Sulfur is a non-metal, so SO₂ is an acidic oxide. Calcium hydroxide is an alkali. Acidic oxide + alkali → salt + water, so the SO₂ is chemically absorbed rather than simply blown through.

Part (b): The equation

SO₂(g) + Ca(OH)₂(aq) → CaSO₃(s) + H₂O(l). The salt is calcium sulfite, which is only slightly soluble and settles out as a solid that can be collected. (In practice it is oxidised further to calcium sulfate and sold as gypsum for plasterboard.)

Part (c): Why CO is not removed

Carbon monoxide is a neutral oxide. It reacts with neither acids nor alkalis, so an alkaline spray has no chemical grip on it at all. CO must instead be removed by oxidising it to CO₂ using a catalyst.

(a) SO₂ is an acidic oxide and Ca(OH)₂ is an alkali, so they neutralise each other. (b) SO₂ + Ca(OH)₂ → CaSO₃ + H₂O. (c) CO is a neutral oxide, so it does not react with alkalis; it must be catalytically oxidised to CO₂ instead.

Exam Tips for 7.2

1. Start with metal or non-metal. Look up the element in the Periodic Table. Metal → basic; non-metal → acidic. That answers most questions in five seconds.

2. Learn the two exception lists cold. Amphoteric: ZnO and Al₂O₃. Neutral: CO, NO, N₂O, H₂O. Everything else follows the general rule.

3. Amphoteric needs BOTH halves. A definition that mentions only acids or only alkalis scores zero. Say "reacts with acids and with bases/alkalis, forming a salt and water in each case".

4. Basic does not mean soluble. CuO and Fe₂O₃ are basic oxides that will not dissolve in water at all. Only the soluble ones (Na₂O, K₂O, CaO) form alkalis.

5. Know the salt names for the amphoteric reactions. ZnO with NaOH gives sodium zincate; Al₂O₃ with NaOH gives sodium aluminate. These names appear on mark schemes.

6. Watch the CO / CO₂ and NO / NO₂ pairs. One is neutral, the other acidic. This is the single most common distractor in 7.2 multiple choice.

🌎 Apply It: Real-World Chemistry

Whether an oxide is acidic, basic or amphoteric decides how power stations are cleaned, why aeroplanes do not dissolve, and what happens to a coral reef.

A coal-fired power station in Uttar Pradesh installs flue-gas desulfurisation. Waste gases are sprayed with a slurry of powdered limestone and water, and the plant then sells the solid by-product to a plasterboard factory. Sulfur dioxide emissions fall by 95%, but carbon dioxide emissions are unchanged.

Explain the chemistry of the clean-up, and explain why CO₂ is not removed by the same process even though it is also an acidic oxide.

▼

The Chemistry

Sulfur dioxide is an acidic oxide; calcium carbonate is a base. The acidic oxide is neutralised: SO₂ + CaCO₃ → CaSO₃ + CO₂. Blowing air through oxidises the calcium sulfite to calcium sulfate, CaSO₄·2H₂O — which is gypsum, exactly the material plasterboard is made from.

Why CO₂ Escapes

Carbon dioxide is a much weaker acidic oxide than sulfur dioxide — carbonic acid is a weak acid, whereas sulfurous acid is considerably stronger. In a limestone slurry the CO₂ is not held; worse, the reaction above actually releases extra CO₂. Cleaning up sulfur dioxide has a small carbon cost.

Chemistry Connection

This is a lovely example of "acidic" being a spectrum rather than a switch. Both oxides are acidic, but only the more strongly acidic one is captured. It also links straight back to 7.1: the strength of the acid formed determines what the oxide can do.

Aluminium window frames on a house in Kolkata look identical after twenty monsoons, while the steel gate beside them has rusted through twice. Yet the reactivity series places aluminium above iron. A chemistry teacher also warns never to clean aluminium cookware with strong oven cleaner.

Resolve the apparent contradiction, and explain the oven-cleaner warning.

▼

The Protective Layer

Aluminium reacts instantly with oxygen to form a thin, hard, impermeable layer of aluminium oxide that seals the metal underneath. Rust, by contrast, is flaky and porous, so oxygen and water reach fresh iron beneath it and corrosion continues all the way through.

The Oven Cleaner Warning

Aluminium oxide is amphoteric. Oven cleaner is concentrated sodium hydroxide, and the protective oxide dissolves in it: Al₂O₃ + 2NaOH → 2NaAlO₂ + H₂O. Strip the oxide away and the exposed aluminium metal is attacked by the alkali too — the pan pits, darkens and can be ruined in minutes.

Chemistry Connection

Amphoteric behaviour is usually taught as an abstract classification. Here it has an immediate practical consequence: aluminium is superbly resistant to acids and weather but has one specific vulnerability — strong alkali. That is why aluminium is never used for tanks holding caustic soda.

Marine biologists monitoring the Great Barrier Reef report that surface seawater pH has fallen from about 8.2 to 8.1 since the industrial revolution, and that coral skeletons are becoming measurably thinner. A newspaper describes this as "the oceans turning acidic".

Explain the chemistry, evaluate the newspaper's wording, and explain why such a small pH change matters.

▼

The Chemistry

Carbon dioxide is an acidic oxide. It dissolves in seawater and reacts: CO₂ + H₂O ⇌ H₂CO₃, releasing H⁺ ions. Those extra H⁺ ions attack the calcium carbonate of coral skeletons and shells: CaCO₃ + 2H⁺ → Ca²⁺ + H₂O + CO₂ — the acid + carbonate reaction from 7.1, running slowly across an entire ocean.

Evaluating the Wording

The newspaper is imprecise. At pH 8.1 the ocean is still alkaline; it has not become acidic. The correct description is "acidification" — a shift towards the acidic end of the scale. This is a genuine exam distinction: moving from pH 9 to pH 8 is acidification, but the solution remains alkaline throughout.

Why 0.1 Matters

The pH scale is logarithmic. A drop of 0.1 corresponds to roughly a 26% increase in H⁺ concentration — a large chemical change hiding behind a small-looking number.

Chemistry Connection

Three syllabus points converge here: CO₂ as an acidic oxide, carbonic acid as a weak acid, and acid + carbonate as a reaction. Cambridge has set questions asking students to explain reef damage using nothing more than these three ideas.

A dermatologist in Australia recommends a "mineral" sunscreen containing 20% zinc oxide, and a paediatrician recommends a zinc oxide cream for nappy rash. Both note the compound is essentially inert on skin at pH 5.5, but a laboratory technician handling the same powder must avoid contact with strong alkalis.

Explain how one compound can be both a skin-safe barrier and an amphoteric reagent.

▼

Why It Is Safe on Skin

Zinc oxide is insoluble in water and skin sits at a mildly acidic pH of about 5.5 — far too weakly acidic to dissolve it at any appreciable rate. So it sits on the surface as a physical barrier, scattering ultraviolet light and keeping moisture off inflamed skin. Its insolubility is precisely why it is safe.

Why the Laboratory Warning

With concentrated acids or concentrated hot alkalis it reacts readily, showing its amphoteric character: ZnO + 2HCl → ZnCl₂ + H₂O and ZnO + 2NaOH → Na₂ZnO₂ + H₂O. Reactivity depends on conditions, not just on identity.

Chemistry Connection

Being amphoteric does not mean "reacts with everything all the time". It means the reaction is possible in both directions given suitable conditions. Exam answers should say "reacts with acids and with alkalis" — the conditions matter in the laboratory but the classification is about capability.

Air-quality monitors along a busy road in Delhi record high levels of nitrogen monoxide (NO) close to the traffic, but higher levels of nitrogen dioxide (NO₂) a few hundred metres downwind. Rainwater collected downwind has a pH of 4.3, while rain collected in a rural area 200 km away measures pH 5.6.

Explain the change from NO to NO₂, the difference in the two rainwater pH values, and why rural rain is not pH 7.

▼

NO Becomes NO₂

The intense heat inside an engine makes nitrogen and oxygen from the air combine: N₂ + O₂ → 2NO. Nitrogen monoxide is a neutral oxide, so at the roadside it does nothing to rainwater. In the open air it is slowly oxidised: 2NO + O₂ → 2NO₂. That takes time, which is why the NO₂ peak appears downwind.

Why the Downwind Rain Is Acidic

Nitrogen dioxide is an acidic oxide. It dissolves in rain to form nitric acid, a strong acid, driving the pH down to 4.3. The neutral NO could never have done this — the extra oxygen atom changes the entire chemistry.

Why Rural Rain Is pH 5.6, Not 7

All rain dissolves atmospheric carbon dioxide, which is an acidic oxide: CO₂ + H₂O ⇌ H₂CO₃. Carbonic acid is weak, so the pH settles at about 5.6. This is natural, not pollution. Genuine "acid rain" means a pH significantly below 5.6.

Chemistry Connection

This scenario is built entirely on the neutral-versus-acidic distinction for the NO/NO₂ pair. If you can explain why one is harmless to rainwater and the other is not, you have shown the examiner you understand oxide classification rather than just memorising lists.

Practice Questions: 7.2

20 multiple choice questions. Click an option to check your answer.

Your Score 0 / 20

Question 1

Which type of element usually forms a basic oxide?

A A metal

B A non-metal

C A noble gas

D A halogen

Metal oxides are basic; non-metal oxides are acidic. Noble gases do not normally form oxides at all.

Question 2

Which oxide is amphoteric?

A MgO

B SO₂

C Al₂O₃

D CO

Aluminium oxide reacts with both acids and alkalis. MgO is basic, SO₂ is acidic, and CO is neutral.

Question 3

Which oxide is neutral?

A NO₂

B N₂O

C CO₂

D SO₃

Dinitrogen oxide (laughing gas) is neutral, along with CO, NO and H₂O. NO₂, CO₂ and SO₃ are all acidic oxides.

Question 4

An oxide dissolves in both dilute hydrochloric acid and hot aqueous sodium hydroxide. It is:

A Acidic

B Basic

C Amphoteric

D Neutral

Reacting with both an acid and an alkali is the definition of amphoteric. The two examples you must know are ZnO and Al₂O₃.

Question 5

Carbon dioxide is bubbled through aqueous sodium hydroxide. What are the products?

A Sodium carbonate + water

B Sodium carbonate + hydrogen

C Carbon + sodium oxide

D No reaction occurs

CO₂ is an acidic oxide, so acidic oxide + alkali → salt + water: CO₂ + 2NaOH → Na₂CO₃ + H₂O. No gas is produced.

Question 6

Which oxide dissolves in water to produce a solution with a pH of about 13?

A CuO

B Na₂O

C SO₂

D ZnO

Sodium oxide is a soluble basic oxide: Na₂O + H₂O → 2NaOH, a strong alkali. CuO and ZnO are insoluble in water, and SO₂ gives an acidic solution.

Question 7

Which statement about copper(II) oxide is correct?

A It is an alkali because it is a metal oxide

B It is a basic oxide but is insoluble in water

C It is amphoteric

D It is neutral because it does not dissolve

CuO neutralises acids, so it is a base — but it is insoluble, so it is not an alkali. Insolubility does not make an oxide neutral; neutral means it reacts with neither acids nor alkalis.

Question 8

Sulfur trioxide dissolves in water. Which acid is formed?

A Sulfurous acid, H₂SO₃

B Sulfuric acid, H₂SO₄

C Hydrogen sulfide, H₂S

D Carbonic acid, H₂CO₃

SO₃ + H₂O → H₂SO₄. Count the oxygens: SO₂ gives sulfurous acid (H₂SO₃), SO₃ gives sulfuric acid (H₂SO₄).

Question 9

Zinc oxide reacts with hot aqueous sodium hydroxide. What is the salt formed called?

A Sodium zinc oxide

B Sodium zincate

C Zinc hydroxide

D Zinc sodate

ZnO + 2NaOH → Na₂ZnO₂ + H₂O. The salt is sodium zincate. The equivalent product from aluminium oxide is sodium aluminate.

Question 10

Across Period 3 from sodium to sulfur, the oxides change in character as:

A acidic → amphoteric → basic

B basic → amphoteric → acidic

C neutral → basic → acidic

D acidic throughout

Metals on the left give basic oxides; non-metals on the right give acidic oxides; aluminium oxide sits amphoterically at the borderline between them.

Question 11

Which pair contains one neutral oxide and one acidic oxide?

A CO and CO₂

B MgO and CaO

C ZnO and Al₂O₃

D SO₂ and SO₃

Carbon monoxide is neutral while carbon dioxide is acidic. B is two basic oxides, C is two amphoteric oxides and D is two acidic oxides.

Question 12

Which equation shows an oxide behaving as an acidic oxide?

A CuO + H₂SO₄ → CuSO₄ + H₂O

B SO₂ + 2NaOH → Na₂SO₃ + H₂O

C CaO + H₂O → Ca(OH)₂

D 2Mg + O₂ → 2MgO

An acidic oxide is defined by its reaction with a base or alkali. Only B shows an oxide neutralising an alkali. A and C show basic behaviour; D is just a synthesis reaction.

Question 13

A student adds an unknown white oxide to dilute nitric acid; it dissolves. She then adds a fresh sample to hot concentrated sodium hydroxide; nothing happens. The oxide is:

A Acidic

B Basic

C Amphoteric

D Neutral

Reacts with acid but not with alkali ⇒ purely basic. Magnesium oxide would behave exactly like this; zinc oxide would have dissolved in both.

Question 14

Why does unpolluted rainwater have a pH of about 5.6 rather than 7?

A Sulfur dioxide from volcanoes dissolves in it

B Carbon dioxide dissolves in it to form weak carbonic acid

C Nitrogen monoxide dissolves in it

D Water is naturally slightly acidic

CO₂ is always present in the air and is an acidic oxide: CO₂ + H₂O ⇌ H₂CO₃. NO is neutral so it would have no effect, and pure water is exactly pH 7.

Question 15

Which of these would not react with dilute hydrochloric acid?

A Zinc oxide

B Calcium oxide

C Carbon monoxide

D Iron(III) oxide

CO is a neutral oxide, so it reacts with neither acids nor alkalis. The other three are all metal oxides that behave as bases towards acid.

Question 16

Which statement best explains why aluminium cookware should not be cleaned with strong alkali?

A Aluminium is above hydrogen in the reactivity series

B The protective aluminium oxide layer is amphoteric and dissolves in alkali

C Aluminium oxide is a neutral oxide

D Alkalis always dissolve metals

Al₂O₃ + 2NaOH → 2NaAlO₂ + H₂O removes the protective barrier, exposing the metal. Alkalis certainly do not dissolve all metals — iron is unaffected.

Question 17

Sulfur dioxide in flue gases is removed by spraying with calcium hydroxide. Which classification explains why this works?

A SO₂ is a basic oxide and Ca(OH)₂ is an acid

B SO₂ is an acidic oxide and Ca(OH)₂ is an alkali

C SO₂ is amphoteric

D Both are neutral oxides

Sulfur is a non-metal, so SO₂ is acidic. Calcium hydroxide is an alkali, so the two neutralise: SO₂ + Ca(OH)₂ → CaSO₃ + H₂O.

Question 18

Which oxide would turn damp universal indicator paper red when the gas is passed over it?

A N₂O

B CO

C NO₂

D H₂

Nitrogen dioxide is acidic and dissolves in the damp paper to give an acidic solution. N₂O and CO are neutral, and hydrogen is not an oxide at all.

Question 19

Which statement about aluminium oxide is false?

A It reacts with hydrochloric acid to form aluminium chloride

B It reacts with sodium hydroxide to form sodium aluminate

C It dissolves readily in water to form an alkaline solution

D It forms a protective layer on aluminium metal

Aluminium oxide is essentially insoluble in water — that insolubility is exactly what makes it a good protective barrier. Amphoteric does not imply soluble.

Question 20

A gas X is a neutral oxide. Which of these is a valid conclusion about X?

A X must be harmless to humans

B X must be insoluble in water

C X will not react with either acids or alkalis

D X must be a metal oxide

Neutral is a purely chemical classification about acid–base behaviour. Carbon monoxide is neutral and lethally toxic, so "neutral" certainly does not mean "safe".

7.3 Preparation of Salts ▼

What Is a Salt?

A salt is the compound formed when the hydrogen ion of an acid is replaced by a metal ion (or by the ammonium ion, NH₄⁺). Every salt therefore has two parts: a positive ion from the base and a negative ion from the acid.

salt = metal (or ammonium) ion + acid's negative ion

NaCl comes from Na⁺ (from sodium hydroxide) and Cl⁻ (from hydrochloric acid). (NH₄)₂SO₄ comes from NH₄⁺ (from ammonia) and SO₄²⁻ (from sulfuric acid).

Naming the Salt from the Acid

The acid decides the second word of the salt's name. Learn this table — you will use it in every single salt-preparation question.

Acid	Formula	Negative ion supplied	Salt family	Example
Hydrochloric acid	HCl	Cl⁻	chloride	Sodium chloride, NaCl
Sulfuric acid	H₂SO₄	SO₄²⁻	sulfate	Copper(II) sulfate, CuSO₄
Nitric acid	HNO₃	NO₃⁻	nitrate	Calcium nitrate, Ca(NO₃)₂
Phosphoric acid	H₃PO₄	PO₄³⁻	phosphate	Sodium phosphate, Na₃PO₄
Ethanoic acid	CH₃COOH	CH₃COO⁻	ethanoate	Sodium ethanoate, CH₃COONa
Carbonic acid	H₂CO₃	CO₃²⁻	carbonate	Sodium carbonate, Na₂CO₃

Memory Trick

"The metal comes first, the acid comes last." Read the question backwards from the acid: hydrochloric → chloride; sulfuric → sulfate; nitric → nitrate. And watch the trap: sulfuric acid gives sulfate (SO₄²⁻), never sulfide (S²⁻) — sulfides come from hydrogen sulfide, which is not on your syllabus.

Solubility Rules — The Table That Decides Everything

Before you can choose a preparation method you must know whether your target salt is soluble or insoluble in water. There is no way around learning this table.

Compound type	Soluble	Insoluble
Sodium, potassium and ammonium salts	ALL soluble — no exceptions	—
Nitrates	ALL soluble — no exceptions	—
Chlorides	Most are soluble	Silver chloride, lead(II) chloride
Sulfates	Most are soluble	Barium sulfate, lead(II) sulfate; calcium sulfate is only slightly soluble
Carbonates	Only sodium, potassium and ammonium carbonate	All the rest are insoluble
Hydroxides	Only sodium, potassium and ammonium hydroxide	All the rest are insoluble; calcium hydroxide is slightly soluble

Memory Trick — Two Halves

Top half: "Nice People Are Never Silly" — Nitrates, Potassium, Ammonium, Na (sodium) are always Soluble.

Bottom half: carbonates and hydroxides are insoluble by default — the only ones that dissolve are the sodium, potassium and ammonium versions from the top half. Then remember just five insoluble oddities: silver chloride, lead chloride, barium sulfate, lead sulfate and (slightly) calcium sulfate.

Exam Tip

Do not over-generalise. "All carbonates are insoluble" is wrong — sodium carbonate (washing soda) dissolves happily, and that is exactly why it works in precipitation reactions. Equally, "all sulfates are soluble" is wrong — barium sulfate is famously insoluble, which is why it is used as the "barium meal" in medical X-rays. Examiners set questions on precisely these two over-generalisations.

Choosing the Right Method

There are three methods, and the whole skill lies in choosing correctly. Ask two questions, in this order:

Is the salt I want soluble or insoluble? If insoluble → precipitation. Stop here.
If it is soluble, are my starting materials soluble too? If one reactant is an insoluble solid → excess solid method. If both reactants are solutions → titration.

Two questions, three routes. Sketch this decision tree in the margin of your exam paper before you answer any salt-preparation question.

Method 1: Acid + Excess Insoluble Solid

Use this when you want a soluble salt and one of your reactants is an insoluble solid — a metal, an insoluble base (metal oxide or hydroxide) or an insoluble carbonate.

Example: preparing copper(II) sulfate crystals from copper(II) oxide and dilute sulfuric acid.

Warm about 25 cm³ of dilute sulfuric acid in a beaker. Warming speeds the reaction up; do not boil.
Add black copper(II) oxide a spatula at a time, stirring, until no more dissolves and some solid remains at the bottom. This is your signal that the acid has been completely used up — the solid is now in excess.
Filter the mixture to remove the excess copper(II) oxide. The filtrate is a blue solution of copper(II) sulfate.
Evaporate the filtrate in an evaporating basin until about half the water has gone and a hot saturated solution remains — the "point of crystallisation". Test by dipping a glass rod in; crystals forming on the rod means it is ready.
Cool the solution slowly to allow crystals to grow. Slow cooling gives larger, better-formed crystals.
Filter off the crystals and dry them between two pieces of filter paper, or in a warm oven at low temperature.

CuO(s) + H₂SO₄(aq) → CuSO₄(aq) + H₂O(l)

The same procedure works with a metal (Zn + H₂SO₄), an insoluble hydroxide, or an insoluble carbonate (CuCO₃ + H₂SO₄), which also gives off CO₂.

Exam Tip — The Three Marks Everyone Loses

(1) Say why you add excess: "to make sure all the acid has reacted". If any acid were left, it would contaminate your crystals.

(2) Say why you filter: "to remove the unreacted excess solid". Do not just write "filter".

(3) Do NOT evaporate to dryness. Heating a hydrated salt too far drives off the water of crystallisation and you get an anhydrous powder instead of crystals. Write "evaporate until the point of crystallisation, then leave to cool and crystallise".

Method 2: Titration (Both Reactants Soluble)

Use this when you want a soluble salt and both reactants are solutions — typically an acid plus a soluble alkali (NaOH, KOH, aqueous ammonia) or a soluble carbonate (Na₂CO₃).

Here the excess-solid trick is impossible: if you add too much sodium hydroxide it simply dissolves and you cannot see it. So you must measure the exact volume needed — that is what a titration is for.

Use a pipette to transfer exactly 25.0 cm³ of the alkali into a conical flask.
Add 2–3 drops of a suitable indicator (methyl orange or thymolphthalein).
Fill a burette with the acid and run it in, swirling, until the indicator changes colour permanently. Record the titre.
Repeat until you obtain concordant results (within 0.10 cm³ of each other).
Now repeat the titration exactly using the same volumes but with no indicator added. The product is a pure salt solution rather than one contaminated with dye.
Evaporate to the point of crystallisation, cool, filter and dry as before.

NaOH(aq) + HCl(aq) → NaCl(aq) + H₂O(l)

You cannot make sodium chloride by the excess-solid method, because every sodium compound is soluble — there would be nothing to filter.

The burette delivers acid drop by drop into a known volume of alkali. The white tile behind the flask makes the first permanent colour change unmistakable.

Method 3: Precipitation (For Insoluble Salts)

Use this when the salt you want is insoluble. The trick is to pick two soluble compounds that between them supply the two ions you need, mix their solutions, and let the insoluble salt drop out.

Example: preparing barium sulfate. You need Ba²⁺ and SO₄²⁻. Choose a soluble barium compound (barium chloride or barium nitrate) and a soluble sulfate (sodium sulfate or dilute sulfuric acid).

BaCl₂(aq) + Na₂SO₄(aq) → BaSO₄(s) + 2NaCl(aq)

Ionic equation: Ba²⁺(aq) + SO₄²⁻(aq) → BaSO₄(s) The Na⁺ and Cl⁻ are spectator ions and stay dissolved. Observation: an immediate thick white precipitate.

The procedure is short:

Mix the two solutions and stir.
Filter to collect the precipitate as the residue.
Wash the residue with distilled water while it is still in the filter paper — this rinses away the soluble spectator-ion salt that would otherwise contaminate it.
Dry the solid in a warm oven or between filter papers.

Exam Tip

The washing step is worth a mark on its own and is the one candidates forget. Say "wash with distilled water to remove the soluble sodium chloride" — name the impurity if you can. And note there is no evaporation and no crystallisation in this method: the product is already a solid.

Insoluble salt wanted	Soluble reactant 1	Soluble reactant 2	Appearance of precipitate
Silver chloride, AgCl	Silver nitrate	Sodium chloride	White (darkens in light)
Barium sulfate, BaSO₄	Barium chloride	Sodium sulfate / dilute H₂SO₄	White
Lead(II) iodide, PbI₂	Lead(II) nitrate	Potassium iodide	Bright yellow
Calcium carbonate, CaCO₃	Calcium chloride	Sodium carbonate	White
Copper(II) hydroxide, Cu(OH)₂	Copper(II) sulfate	Sodium hydroxide	Pale blue

Memory Trick

To pick your two reactants, use the nitrate and sodium rule: nitrates are always soluble and sodium salts are always soluble. So take your positive ion as a nitrate and your negative ion as a sodium salt, and the two starting solutions are guaranteed to dissolve. Want lead(II) sulfate? Use lead(II) nitrate plus sodium sulfate. It works every time.

Water of Crystallisation

Many salts crystallise with water molecules built into the crystal lattice. This is called water of crystallisation, and a salt containing it is hydrated. A salt with the water removed is anhydrous.

CuSO₄·5H₂O(s) ⇌ CuSO₄(s) + 5H₂O(l)

Left: hydrated copper(II) sulfate — blue crystals. Right: anhydrous copper(II) sulfate — white powder. Heating drives the water off (endothermic); adding water back turns it blue again (exothermic). This is the standard test for water.

Hydrated salt	Formula	Common name	Colour change on heating
Copper(II) sulfate-5-water	CuSO₄·5H₂O	Blue vitriol	Blue → white
Cobalt(II) chloride-6-water	CoCl₂·6H₂O	—	Pink → blue
Sodium carbonate-10-water	Na₂CO₃·10H₂O	Washing soda	Colourless crystals → white powder
Calcium sulfate-2-water	CaSO₄·2H₂O	Gypsum	→ plaster of Paris
Magnesium sulfate-7-water	MgSO₄·7H₂O	Epsom salts	Colourless crystals → white powder

Exam Tip

Water of crystallisation explains the "do not evaporate to dryness" rule. If you keep heating, you drive off the water that is chemically part of the crystal and destroy the crystal structure. You end up with an anhydrous powder, not the blue crystals the question asked for. This single point appears on mark schemes for salt preparation year after year.

Worked Examples

Worked Example 1 Choose the correct preparation method, name the starting materials and write a balanced equation for each of: (a) zinc sulfate, (b) potassium nitrate, (c) lead(II) iodide.

(a) Zinc sulfate — check solubility first

Most sulfates are soluble, and zinc sulfate is one of them ⇒ soluble salt. Zinc is a metal (insoluble in water) and zinc oxide/carbonate are insoluble too, so I have an insoluble reactant available ⇒ Method 1, excess solid. Add excess zinc oxide to warm dilute sulfuric acid, filter off the excess, evaporate to the point of crystallisation, cool, filter and dry.
ZnO(s) + H₂SO₄(aq) → ZnSO₄(aq) + H₂O(l)

(b) Potassium nitrate — the trap

All potassium salts and all nitrates are soluble ⇒ soluble salt. But every potassium compound is also soluble, so there is no insoluble solid to add in excess and nothing to filter off. That rules out Method 1 ⇒ Method 2, titration. Titrate 25.0 cm³ of potassium hydroxide against dilute nitric acid with methyl orange, record the titre, repeat without indicator, then evaporate and crystallise.
KOH(aq) + HNO₃(aq) → KNO₃(aq) + H₂O(l)

Lead(II) iodide is insoluble ⇒ Method 3, precipitation. I need Pb²⁺ and I⁻ from two soluble compounds: use lead(II) nitrate (all nitrates soluble) and potassium iodide (all potassium salts soluble). Mix, filter, wash the residue with distilled water, dry.
Pb(NO₃)₂(aq) + 2KI(aq) → PbI₂(s) + 2KNO₃(aq)
Ionic: Pb²⁺(aq) + 2I⁻(aq) → PbI₂(s). A bright yellow precipitate forms.

(a) Excess solid method: ZnO + H₂SO₄ → ZnSO₄ + H₂O. (b) Titration: KOH + HNO₃ → KNO₃ + H₂O. (c) Precipitation: Pb(NO₃)₂ + 2KI → PbI₂ + 2KNO₃.

Worked Example 2 Ananya wants to prepare pure, dry crystals of magnesium sulfate (Epsom salts) starting from magnesium carbonate powder and dilute sulfuric acid. Write a full method, explaining the purpose of each step, and give one safety precaution.

Step 1: Confirm the method

Magnesium sulfate is soluble; magnesium carbonate is insoluble. So this is Method 1, excess insoluble solid. The reaction is acid + carbonate → salt + water + carbon dioxide: MgCO₃(s) + H₂SO₄(aq) → MgSO₄(aq) + H₂O(l) + CO₂(g).

Step 2: React with excess

Warm about 25 cm³ of dilute sulfuric acid gently in a beaker. Add magnesium carbonate a spatula at a time, stirring. Effervescence occurs as CO₂ is released. Keep adding until the fizzing stops and solid remains undissolved — the visual signal that all the acid has reacted. Purpose: to ensure no acid is left to contaminate the crystals.

Step 3: Filter

Filter the mixture through filter paper in a funnel. Purpose: to remove the excess unreacted magnesium carbonate. The filtrate is a colourless solution of magnesium sulfate.

Step 4: Crystallise

Transfer the filtrate to an evaporating basin and heat over a water bath until roughly half the water has evaporated. Test with a cold glass rod — crystals appearing on it means the point of crystallisation is reached. Purpose: to concentrate the solution so it becomes saturated. Do not evaporate to dryness, or the water of crystallisation will be driven off and you will get anhydrous powder instead of MgSO₄·7H₂O crystals.

Step 5: Cool, filter, dry

Leave the hot saturated solution to cool slowly. Solubility falls as the temperature drops, so crystals grow. Slow cooling gives larger, purer crystals. Filter them off and dry between two pieces of filter paper (or in a warm oven at low temperature).

Safety

Wear eye protection throughout — dilute sulfuric acid is an irritant and the hot saturated solution can spit during evaporation. Use tongs for hot apparatus.

Warm the acid, add excess MgCO₃ until fizzing stops and solid remains, filter off the excess, evaporate the filtrate to the point of crystallisation, cool slowly to crystallise, then filter and dry between filter papers. Never evaporate to dryness. Wear eye protection.

Worked Example 3 In a titration, 25.0 cm³ of 0.100 mol/dm³ sodium hydroxide solution required exactly 20.0 cm³ of dilute sulfuric acid for neutralisation. Calculate the concentration of the sulfuric acid in mol/dm³ and in g/dm³. (Aᵣ: H = 1, O = 16, S = 32)

Step 1: Write the balanced equation

2NaOH(aq) + H₂SO₄(aq) → Na₂SO₄(aq) + 2H₂O(l). The mole ratio is 2 NaOH : 1 H₂SO₄ because sulfuric acid is diprotic — each molecule supplies two H⁺.

Step 2: Convert cm³ to dm³

Divide by 1000. Alkali: 25.0 ÷ 1000 = 0.0250 dm³. Acid: 20.0 ÷ 1000 = 0.0200 dm³. Forgetting this step is the single most common titration error.

Step 3: Moles of the known substance

moles = concentration × volume = 0.100 × 0.0250 = 2.50 × 10⁻³ mol NaOH.

Step 4: Use the ratio

2 mol NaOH reacts with 1 mol H₂SO₄, so divide by 2: moles of H₂SO₄ = 2.50 × 10⁻³ ÷ 2 = 1.25 × 10⁻³ mol.

Step 5: Concentration in mol/dm³

concentration = moles ÷ volume = 1.25 × 10⁻³ ÷ 0.0200 = 0.0625 mol/dm³.

Step 6: Convert to g/dm³

Mᵣ of H₂SO₄ = (2 × 1) + 32 + (4 × 16) = 2 + 32 + 64 = 98. concentration in g/dm³ = 0.0625 × 98 = 6.125 g/dm³, which is 6.13 g/dm³ to 3 significant figures.

The sulfuric acid concentration is 0.0625 mol/dm³, equivalent to 6.13 g/dm³.

Worked Example 4 A student titrates 25.0 cm³ of hydrochloric acid of unknown concentration against 0.200 mol/dm³ sodium hydroxide. Her titres are 22.40, 21.50, 21.60 and 21.55 cm³. Calculate the concentration of the acid, and explain how you treated the results.

Step 1: Handle the results correctly

The first titre (22.40) is a rough titre and is well outside the others — it is discarded. The remaining three, 21.50, 21.60 and 21.55, are concordant (all within 0.10 cm³ of one another). Never simply average all four.

Step 2: Mean titre

(21.50 + 21.60 + 21.55) ÷ 3 = 64.65 ÷ 3 = 21.55 cm³ = 0.02155 dm³.

Step 3: Moles of NaOH

moles = 0.200 × 0.02155 = 4.310 × 10⁻³ mol.

Step 4: Apply the mole ratio

NaOH + HCl → NaCl + H₂O, a 1 : 1 ratio. So moles of HCl = 4.310 × 10⁻³ mol. (Compare with the previous example — using 2 : 1 here would halve the answer and lose the marks.)

Step 5: Concentration of the acid

Volume of acid = 25.0 cm³ = 0.0250 dm³. concentration = 4.310 × 10⁻³ ÷ 0.0250 = 0.1724 mol/dm³.

Discarding the rough titre and averaging the three concordant results gives a mean titre of 21.55 cm³. The hydrochloric acid concentration is 0.172 mol/dm³ (3 s.f.).

Worked Example 5 5.00 g of hydrated magnesium sulfate crystals, MgSO₄·xH₂O, is heated until the mass is constant. 2.44 g of anhydrous magnesium sulfate remains. Calculate x. (Aᵣ: H = 1, O = 16, Mg = 24, S = 32)

Step 1: Find the mass of water driven off

mass of water = 5.00 − 2.44 = 2.56 g. This water was the water of crystallisation held inside the lattice.

Step 2: Moles of anhydrous salt

Mᵣ of MgSO₄ = 24 + 32 + (4 × 16) = 120. moles = 2.44 ÷ 120 = 0.02033 mol.

Step 3: Moles of water

Mᵣ of H₂O = 18. moles = 2.56 ÷ 18 = 0.1422 mol.

Step 4: Find the simplest ratio

Divide both by the smaller number: MgSO₄ = 0.02033 ÷ 0.02033 = 1. H₂O = 0.1422 ÷ 0.02033 = 6.99 ≈ 7.

Step 5: Sanity check

x must be a whole number, and 6.99 rounds cleanly to 7. The formula MgSO₄·7H₂O is Epsom salts — a familiar compound, which confirms the answer. "Heated to constant mass" matters: it proves all the water has gone.

x = 7, so the formula of the hydrated salt is MgSO₄·7H₂O.

Exam Tips for 7.3

1. Always start with solubility. Write "soluble" or "insoluble" next to the target salt before choosing anything. Insoluble → precipitation, always.

2. Sodium, potassium and ammonium salts can never be made by the excess-solid method. Every compound of those three metals is soluble, so nothing would remain to filter. They must be made by titration.

3. Two filtrations in Method 1. One to remove the excess solid, one at the end to collect the crystals. Questions often ask "why is the mixture filtered?" — check which filtration they mean.

4. No filtration at all in Method 2. Both reactants are solutions, so there is nothing to filter out. Writing "filter off the excess alkali" in a titration question is an instant giveaway.

5. Repeat the titration without indicator. Otherwise your salt crystals contain dye. This is a standard mark.

6. Wash the precipitate. In Method 3, always add "wash with distilled water" to remove the soluble salt formed alongside it.

7. Never evaporate to dryness. Say "evaporate to the point of crystallisation, then leave to cool". Boiling dry destroys the water of crystallisation.

8. In titration sums, do the ratio check. HCl and HNO₃ are 1 : 1 with NaOH; H₂SO₄ is 1 : 2. Write the balanced equation first, every single time.

9. cm³ ÷ 1000 = dm³. Write the conversion down as a separate line so you cannot forget it under pressure.

10. Discard the rough titre. Average only concordant results (within 0.10 cm³).

🌎 Apply It: Real-World Chemistry

Salt making is not a school exercise — it is how fertiliser, medical imaging agents, photographic film and half a chemical industry actually get produced.

A fertiliser plant near Kanpur produces ammonium sulfate by reacting ammonia gas with sulfuric acid. The plant chemist insists the ratio must be controlled to within a fraction of a percent, because neither excess ammonia nor excess acid can be tolerated in the final product.

Which laboratory method is this the industrial version of, and why is precise control unavoidable here but not when making copper(II) sulfate?

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The Method

Both reactants are soluble — ammonia dissolves readily and sulfuric acid is a solution — and the product, ammonium sulfate, is soluble too. So this is the industrial equivalent of a titration: 2NH₃ + H₂SO₄ → (NH₄)₂SO₄.

Why Precision Is Unavoidable

There is no insoluble solid anywhere in the system, so there is nothing to filter. Any excess ammonia stays dissolved in the product; any excess sulfuric acid does too. The only way to get a pure salt is to add exactly the right amount — which is precisely why titration exists as a method.

The Contrast

For copper(II) sulfate you can be gloriously careless: throw in far too much copper(II) oxide and simply filter the surplus out. The insolubility of the reactant does the purification for you. Ammonium salts do not offer that luxury.

Chemistry Connection

This is the whole logic of method selection in one factory. The question "is one reactant insoluble?" is not about chemistry trivia — it is about whether you can purify by filtration or must purify by measurement. Every salt-preparation exam question is testing that judgement.

A radiographer at a hospital in Leeds gives a patient a "barium meal" — a thick white suspension of barium sulfate — before an X-ray of the digestive tract. Barium ions are extremely toxic. The patient asks why she is being asked to swallow a poison.

Explain why the procedure is safe, how the compound is made, and why barium chloride could never be used instead.

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Why It Is Safe

Barium sulfate is one of the classic insoluble sulfates. It does not dissolve in water, and crucially it does not dissolve in stomach acid either. Because it never releases free Ba²⁺ ions, it cannot be absorbed into the bloodstream — it passes straight through and out. The toxicity of barium belongs to the ion, not to the element locked in an insoluble lattice.

How It Is Made

By precipitation, since the target salt is insoluble: BaCl₂(aq) + Na₂SO₄(aq) → BaSO₄(s) + 2NaCl(aq). Ionic equation: Ba²⁺(aq) + SO₄²⁻(aq) → BaSO₄(s). The precipitate is filtered and then washed thoroughly with distilled water — a step that in this application is genuinely life-critical, because any residual soluble barium chloride clinging to the crystals would be absorbed and could kill.

Why Not Barium Chloride

Barium chloride is soluble. Swallowing it would release Ba²⁺ ions directly into the body. It is a laboratory reagent, not a medicine.

Chemistry Connection

Next time you write "wash the precipitate with distilled water to remove soluble impurities", remember this is not a box-ticking mark scheme phrase — in a hospital it is the difference between a diagnostic image and a fatality. Barium sulfate also blocks X-rays because barium has a high proton number, which is why it shows up white on the film.

In the salt pans of the Rann of Kutch in Gujarat, seawater is channelled into shallow basins and left to evaporate under the sun. Workers harvest common salt from the earlier pans, but the last, most concentrated pans yield a bitter mixture that is sold separately for making Epsom salts and plaster.

Explain, in terms of solubility and crystallisation, why the salts separate in that order and why the final pans are the bitter ones.

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The Chemistry of Fractional Crystallisation

Seawater is a mixed solution of many salts. As water evaporates, each salt becomes more concentrated. A salt begins to crystallise once its own solution becomes saturated. The least soluble salts saturate first, so calcium sulfate (gypsum) comes out early, then sodium chloride, and finally the very soluble magnesium salts.

Why the Last Pans Are Bitter

Magnesium sulfate and magnesium chloride are extremely soluble, so they stay dissolved long after the sodium chloride has crystallised out. They concentrate in the final brine, called bittern, and taste bitter. MgSO₄·7H₂O crystallised from it is Epsom salts — and note the seven waters of crystallisation built into the lattice.

Chemistry Connection

This is your laboratory crystallisation step scaled up to several square kilometres, and it shows why the instruction "evaporate to the point of crystallisation and cool slowly" matters. Slow crystallisation lets one salt come out at a time in large, pure crystals; rapid evaporation to dryness would leave you with a fine, contaminated mixture of everything at once.

A water technician in Birmingham tests a river sample suspected of contamination from a metal-plating works. She adds a few drops of dilute nitric acid followed by barium nitrate solution to one portion, and dilute nitric acid followed by silver nitrate solution to another. The first gives a white precipitate; the second gives a white precipitate that darkens on standing in sunlight.

Identify the two ions present, explain the role of the nitric acid, and explain why nitrate salts are used as the test reagents.

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Identifying the Ions

A white precipitate with acidified barium nitrate indicates sulfate ions: Ba²⁺(aq) + SO₄²⁻(aq) → BaSO₄(s). A white precipitate with acidified silver nitrate that darkens in light indicates chloride ions: Ag⁺(aq) + Cl⁻(aq) → AgCl(s). Silver chloride decomposes photochemically to silver metal, which is exactly the chemistry behind old photographic film.

Why Add Nitric Acid First

To remove carbonate ions, which would otherwise form white precipitates of barium carbonate or silver carbonate and give a false positive. The acid converts any carbonate to CO₂ gas, which bubbles away. This is a genuine mark on the mark scheme and is frequently forgotten.

Why Nitrates Specifically

All nitrates are soluble, so barium nitrate and silver nitrate are guaranteed to be in solution and to supply free Ba²⁺ and Ag⁺ ions. If you tried barium sulfate or silver chloride as the reagent you would be adding an insoluble solid — useless. And crucially, the nitrate ion cannot itself form a precipitate with anything, so it can never confuse the result.

Chemistry Connection

Every qualitative test for an anion is really a precipitation reaction run backwards: instead of choosing reactants to make a known insoluble salt, you use a known insoluble salt to reveal an unknown ion. The solubility table you learned for salt preparation is the same table that powers the whole of qualitative analysis.

A quality-control chemist at a vinegar bottling plant in Modena, Italy, must certify that each batch contains 5.00% ethanoic acid by mass. She pipettes 25.0 cm³ of diluted vinegar into a conical flask, adds phenolphthalein, and titrates against standardised 0.100 mol/dm³ sodium hydroxide. Her supervisor asks why she does not simply use a pH meter, which would be faster.

Explain what the titration measures that a pH reading cannot, and why phenolphthalein rather than methyl orange is chosen.

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What the Titration Measures

A pH meter measures the concentration of free H⁺ ions at that instant. Ethanoic acid is a weak acid, so only a small fraction of its molecules are dissociated at any moment — the pH reading would drastically underestimate how much acid is actually present. A titration measures the total moles of acid, because as the alkali removes H⁺, the equilibrium CH₃COOH ⇌ H⁺ + CH₃COO⁻ shifts right until every molecule has been neutralised.

Why Phenolphthalein

A weak acid neutralised by a strong base gives a salt (sodium ethanoate) whose solution is slightly alkaline, so the equivalence point sits at about pH 9 rather than pH 7. Phenolphthalein changes colour in the pH 8–10 range, matching that end point. Methyl orange changes at about pH 3–4 and would signal the end point far too early, giving a low result.

Chemistry Connection

This is the sharpest possible demonstration that strength and amount are different quantities. Weak acid, low H⁺ concentration, high pH — and yet exactly the same number of moles of alkali needed as an equally concentrated strong acid would require. If you understand this scenario, you will never again write "a weak acid needs less alkali to neutralise it".

Practice Questions: 7.3

20 multiple choice questions. Click an option to check your answer.

Your Score 0 / 20

Question 1

Which salt is produced when magnesium reacts with dilute nitric acid?

A Magnesium nitride

B Magnesium nitrate

C Magnesium sulfate

D Magnesium chloride

Nitric acid always gives nitrates (NO₃⁻). A nitride contains N³⁻ and comes from nitrogen itself, not from nitric acid — a favourite distractor.

Question 2

Which of these salts is insoluble in water?

A Sodium carbonate

B Ammonium sulfate

C Barium sulfate

D Potassium nitrate

Most sulfates are soluble but barium sulfate and lead(II) sulfate are not. All sodium, potassium and ammonium salts, and all nitrates, are soluble without exception.

Question 3

Which method should be used to prepare pure dry crystals of copper(II) sulfate from copper(II) oxide?

A Add excess copper(II) oxide to warm dilute sulfuric acid, filter, then crystallise

B Titrate copper(II) oxide against sulfuric acid using methyl orange

C Mix copper(II) nitrate and sodium sulfate solutions and filter the precipitate

D Electrolyse copper(II) sulfate solution

The target salt is soluble and one reactant (CuO) is an insoluble solid, so use the excess-solid method. Titration is impossible with an insoluble solid, and copper(II) sulfate is soluble so precipitation would not work.

Question 4

Why is excess solid added in the excess-solid method?

A To speed up the reaction

B To ensure all the acid is used up, so none contaminates the crystals

C To increase the yield of the salt above 100%

D To make the solution more concentrated

Leftover acid would remain in the filtrate and be evaporated into the crystals. Excess solid guarantees the acid is the limiting reactant, and the surplus solid is simply filtered off.

Question 5

Which salt cannot be made by adding an excess of an insoluble solid to an acid?

A Zinc chloride

B Copper(II) nitrate

C Sodium chloride

D Magnesium sulfate

Every sodium compound is soluble, so there is no insoluble sodium solid to add in excess and nothing to filter off. Sodium, potassium and ammonium salts must be made by titration.

Question 6

Which pair of solutions would be mixed to prepare lead(II) sulfate by precipitation?

A Lead(II) sulfate and sodium nitrate

B Lead(II) nitrate and sodium sulfate

C Lead(II) chloride and barium sulfate

D Lead(II) carbonate and sulfuric acid

Both starting materials must be soluble. Nitrates are always soluble and sodium salts are always soluble, so lead(II) nitrate + sodium sulfate is the reliable choice. Option C uses two insoluble compounds.

Question 7

After collecting a precipitate by filtration, why is it washed with distilled water?

A To dissolve any remaining precipitate

B To remove the soluble salt formed alongside it

C To neutralise the precipitate

D To make the precipitate crystallise

The spectator ions form a soluble salt (for example sodium nitrate) that clings to the residue. Distilled water rinses it away without dissolving the insoluble product.

Question 8

In a titration preparation of sodium chloride, why is the experiment repeated with no indicator added?

A Because the indicator would react with the acid

B Because the indicator would contaminate the salt crystals

C Because the indicator changes the volume needed

D Because the indicator prevents crystallisation

The indicator is a coloured dye. Evaporating a solution containing it would give stained crystals, so the exact titre is found first, then repeated cleanly.

Question 9

Why should a salt solution not be evaporated to dryness when preparing crystals?

A The salt would decompose into its elements

B The water of crystallisation would be driven off, giving a powder instead of crystals

C The salt would dissolve again

D The yield would exceed 100%

Hydrated salts contain water built into the lattice. Overheating removes it, converting CuSO₄·5H₂O into anhydrous white CuSO₄ powder rather than blue crystals.

Question 10

25.0 cm³ of 0.100 mol/dm³ NaOH is exactly neutralised by 25.0 cm³ of HCl. What is the concentration of the HCl?

A 0.0500 mol/dm³

B 0.100 mol/dm³

C 0.200 mol/dm³

D 0.400 mol/dm³

NaOH + HCl is a 1 : 1 ratio. Equal volumes and a 1 : 1 ratio mean equal concentrations. Option A would apply if the acid were diprotic.

Question 11

25.0 cm³ of 0.200 mol/dm³ NaOH is exactly neutralised by 25.0 cm³ of sulfuric acid. What is the concentration of the acid?

A 0.100 mol/dm³

B 0.200 mol/dm³

C 0.400 mol/dm³

D 0.0500 mol/dm³

2NaOH + H₂SO₄ → Na₂SO₄ + 2H₂O, a 2 : 1 ratio. Moles NaOH = 0.005; moles H₂SO₄ = 0.0025; concentration = 0.0025 ÷ 0.025 = 0.100 mol/dm³. Always write the equation before dividing.

Question 12

A student records titres of 24.90, 23.20, 23.15 and 23.25 cm³. What mean titre should be used?

A 23.63 cm³ (all four averaged)

B 23.20 cm³ (the three concordant results averaged)

C 24.90 cm³ (the largest value)

D 23.15 cm³ (the smallest value)

The first titre is a rough one and is discarded. (23.20 + 23.15 + 23.25) ÷ 3 = 23.20 cm³. Concordant means within 0.10 cm³ of each other.

Question 13

Which piece of apparatus delivers a fixed, accurately known volume such as 25.0 cm³?

A Burette

B Pipette

C Measuring cylinder

D Conical flask

A pipette delivers one fixed volume very accurately. A burette delivers a variable volume that you read off; a measuring cylinder is far less accurate than either.

Question 14

2.50 g of hydrated copper(II) sulfate is heated to constant mass, leaving 1.60 g of anhydrous salt. Which statement is correct? (Mᵣ CuSO₄ = 160, H₂O = 18)

A 0.90 g of water was lost and x = 3

B 0.90 g of water was lost and x = 5

C 1.60 g of water was lost and x = 5

D No water was lost

Water lost = 2.50 − 1.60 = 0.90 g ⇒ 0.90 ÷ 18 = 0.050 mol. CuSO₄ = 1.60 ÷ 160 = 0.010 mol. Ratio = 1 : 5, so the formula is CuSO₄·5H₂O.

Question 15

Which statement about solubility is correct?

A All carbonates are insoluble

B All sulfates are soluble

C All nitrates are soluble

D All chlorides are soluble

Nitrates are the one family with no exceptions at all. Sodium, potassium and ammonium carbonates are soluble; barium and lead sulfates are insoluble; silver and lead chlorides are insoluble.

Question 16

What is the correct ionic equation for the precipitation of silver chloride?

A AgNO₃(aq) + NaCl(aq) → AgCl(s) + NaNO₃(aq)

B Ag⁺(aq) + Cl⁻(aq) → AgCl(s)

C Ag(s) + Cl₂(g) → AgCl(s)

D Na⁺(aq) + NO₃⁻(aq) → NaNO₃(aq)

An ionic equation shows only the ions that actually change. Option A is the full equation, not the ionic one; option D shows only the spectator ions, which are precisely what you cancel.

Question 17

Which salt would be prepared by titration rather than by the excess-solid method?

A Copper(II) chloride

B Ammonium nitrate

C Zinc sulfate

D Iron(II) sulfate

All ammonium compounds are soluble, so aqueous ammonia and nitric acid are both solutions with nothing to filter. The other three all have an insoluble oxide, carbonate or metal available.

Question 18

In the excess-solid method, how do you know when enough solid has been added?

A The indicator changes colour

B Solid stops dissolving and remains at the bottom of the beaker

C The solution turns colourless

D The mixture starts to boil

Undissolved solid is the visual signal that all the acid has been used up. That is exactly why no indicator is needed in this method — the excess solid does the job.

Question 19

Which observation is made when dilute hydrochloric acid is added to excess solid copper(II) carbonate?

A No reaction, because copper is below hydrogen in the reactivity series

B Effervescence, and a blue-green solution forms

C A white precipitate forms

D Hydrogen gas is released

CuCO₃ + 2HCl → CuCl₂ + H₂O + CO₂. Acids react with all carbonates regardless of the metal's position in the reactivity series — that rule applies only to acid + metal.

Question 20

A student needs pure dry crystals of calcium carbonate. Which sequence is correct?

A Titrate calcium hydroxide against carbonic acid, then evaporate

B Add excess calcium carbonate to hydrochloric acid, filter, evaporate

C Mix calcium chloride and sodium carbonate solutions, filter, wash, dry

D Evaporate limewater to dryness

Calcium carbonate is insoluble, so precipitation is the only route. Both starting solutions are soluble (all chlorides except silver and lead, and all sodium salts). Option B would produce calcium chloride, not carbonate.