OxfordAQA International GCSE Chemistry: The properties of acids and bases, and preparation of salts (9202)

This study guide teaches acids, bases and the preparation of salts for OxfordAQA International GCSE Chemistry (9202). It follows the OxfordAQA International GCSE Chemistry (9202) specification, Version 6.2, for exams May/June 2018 onwards, and covers every statement in sections 3.5.1 (statements a to g) and 3.5.2 (statements a to c). There are no tiers in this qualification, so the whole unit is for every candidate. The specification says any of its content may be assessed on Paper 1 and on Paper 2.

Quick links: course hub · printable checklist · condensed notes on acids and salts · original acids and salts questions · free diagnostics

What this unit covers

Statement	Skills you need
3.5.1 a	Say that metal oxides and hydroxides are bases, and that soluble hydroxides are alkalis
3.5.1 b	Describe neutralisation: acid + base gives a salt
3.5.1 c	Name the salt from any acid and base or alkali (chlorides, nitrates, sulfates)
3.5.1 d	Explain that ammonia solution is alkaline and is used to make ammonium salts
3.5.1 e	Describe how limewater reacts with carbon dioxide to give calcium carbonate
3.5.1 f	Link H⁺(aq) to acidity and OH⁻(aq) to alkalinity; use the pH scale 0 to 14; use universal indicator
3.5.1 g	Write the neutralisation equation H⁺(aq) + OH⁻(aq) → H₂O(l)
3.5.2 a	Make soluble salts from acids with metals, insoluble bases or alkalis; suggest a method for a named salt
3.5.2 b	Crystallise a salt solution to get the solid salt
3.5.2 c	Make insoluble salts by precipitation; name the reagents; explain how precipitation cleans water

None of the required practicals in section 6.2 of the specification sits in 3.5.1 or 3.5.2. The titration practical belongs to 3.6.4 (see the quantitative chemistry guide), but the alkali method below uses the same apparatus, so skills overlap.

Bases, alkalis and neutralisation (3.5.1 a to c)

A base is a metal oxide or a metal hydroxide. Copper(II) oxide, iron(III) oxide and magnesium hydroxide are all bases. An alkali is a base that dissolves in water, such as sodium, lithium or calcium hydroxide. Most metal oxides and hydroxides are insoluble, so most bases are not alkalis.

When an acid reacts with a base, a salt and water form. This is a neutralisation reaction.

acid + base → salt + water

The name of the salt has two parts.

• The first part comes from the metal in the base or alkali (magnesium, copper, lithium…).
• The second part comes from the acid:

Acid	Formula	Salts are called
Hydrochloric acid	HCl	chlorides
Nitric acid	HNO₃	nitrates
Sulfuric acid	H₂SO₄	sulfates

Worked example 1: naming salts and writing equations

Name the salt and write a balanced equation for (a) iron(III) oxide with hydrochloric acid, (b) magnesium hydroxide with hydrochloric acid.

(a) The metal is iron(III); the acid gives a chloride, so the salt is iron(III) chloride, FeCl₃. Iron(III) is 3+ and chloride is 1−, so three chlorides balance one iron.

Fe₂O₃ + 6HCl → 2FeCl₃ + 3H₂O

Check: 2 Fe, 3 O, 6 H and 6 Cl on each side.

(b) The salt is magnesium chloride, MgCl₂.

Mg(OH)₂ + 2HCl → MgCl₂ + 2H₂O

Check: 1 Mg, 2 O, 4 H and 2 Cl on each side.

Ammonia and limewater (3.5.1 d and e)

Ammonia is a gas that dissolves well in water. Some of the dissolved ammonia reacts with water to release hydroxide ions, so the solution is alkaline:

NH₃(aq) + H₂O(l) ⇌ NH₄⁺(aq) + OH⁻(aq)

Ammonia reacts with acids to make ammonium salts. No water forms in these reactions, because each ammonia molecule simply takes a hydrogen ion from the acid to become NH₄⁺:

NH₃ + HNO₃ → NH₄NO₃ (ammonium nitrate)

2NH₃ + H₂SO₄ → (NH₄)₂SO₄ (ammonium sulfate)

Both are used as fertilisers. The ammonium ion, NH₄⁺, takes the place of a metal ion in the salt.

Limewater is a solution of calcium hydroxide in water. It is an alkali. Carbon dioxide reacts with it to give calcium carbonate, which is insoluble, so the clear solution turns cloudy (milky) white:

Ca(OH)₂(aq) + CO₂(g) → CaCO₃(s) + H₂O(l)

This is why limewater is the test for carbon dioxide.

Hydrogen ions, hydroxide ions and pH (3.5.1 f and g)

In water, acids release hydrogen ions, H⁺(aq). These make a solution acidic. Alkalis release hydroxide ions, OH⁻(aq). These make a solution alkaline.

The pH scale runs from 0 to 14 and measures how acidic or alkaline a solution is.

pH	Meaning
0 to just below 7	acidic; the lower the number, the more acidic
7	neutral (for example, pure water)
just above 7 to 14	alkaline; the higher the number, the more alkaline

Using universal indicator

Universal indicator is a mixture of indicators that changes through a range of colours as pH changes. To measure the approximate pH of a solution:

1. Put a small sample of the solution in a test tube or on a spotting tile.
2. Add a few drops of universal indicator solution (or touch a strip of universal indicator paper with a glass rod dipped in the solution).
3. Compare the colour with the colour chart supplied with the indicator and read off the pH.

Typical colours run from red (very acidic, around pH 0 to 2) through orange and yellow (weakly acidic), green (neutral, pH 7), blue (weakly alkaline) to purple (very alkaline, around pH 12 to 14). Shades vary between brands, so always use the chart. The result is approximate, not a precise value.

The neutralisation equation

Whatever the acid and alkali, the reaction that actually happens is between hydrogen ions and hydroxide ions:

H⁺(aq) + OH⁻(aq) → H₂O(l)

The metal ion and the negative ion from the acid stay in solution unchanged and become the salt when the water is removed. As alkali is added to acid, H⁺ ions are used up and the pH rises towards 7; with alkali in excess, the pH goes above 7.

Making soluble salts (3.5.2 a and b)

There are three routes. All end by crystallising the salt solution.

Route 1: acid + metal

metal + acid → salt + hydrogen

Not every metal is suitable. Very reactive metals such as potassium, sodium and lithium react too violently with acids to be safe. Metals below hydrogen in the reactivity series, such as copper and silver, are not reactive enough to give hydrogen with dilute acids. Moderately reactive metals such as magnesium, zinc and iron work well. Add the metal in excess until bubbling stops, then filter off the unreacted metal. For the reactivity series, see the metals study guide.

Route 2: acid + insoluble base

Add the insoluble base (a metal oxide or hydroxide) to the acid until no more will react. Leftover solid shows all the acid has been used. Warming speeds this up. Filter off the excess; the filtrate is the salt solution.

Route 3: acid + alkali

Both reactants are solutions, so an excess cannot be filtered off. An indicator shows when they have exactly reacted.

1. Measure a known volume of alkali into a conical flask with a pipette and add a few drops of indicator.
2. Run acid in from a burette while swirling the flask. Stop at the first permanent colour change and note the burette reading.
3. Repeat with the same volumes but no indicator, so the salt is not contaminated with it (or remove it with activated charcoal).

Crystallising the salt

1. Heat the salt solution in an evaporating basin to boil off some of the water.
2. Stop when the solution is saturated: crystals appear on a cold glass rod dipped in and taken out.
3. Leave the solution to cool slowly; crystals form as the salt becomes less soluble.
4. Filter off the crystals, rinse with a little cold distilled water, then pat them dry with filter paper or leave them somewhere warm.

Do not heat to dryness: many salts decompose or lose water of crystallisation.

Worked example 2: suggesting a method for a named soluble salt

Suggest how to make crystals of copper(II) nitrate, and explain why you chose that method.

Step 1: choose the acid. A nitrate needs nitric acid.

Step 2: choose the copper compound. Copper is below hydrogen, so route 1 is out (with nitric acid, copper metal gives toxic nitrogen oxides, not hydrogen). Copper(II) oxide is an insoluble base, so use route 2.

CuO + 2HNO₃ → Cu(NO₃)₂ + H₂O

Step 3: write the method. Warm about 40 cm³ of dilute nitric acid. Add black copper(II) oxide a spatula at a time, stirring, until some solid remains; the solution turns blue. Filter, then crystallise the blue filtrate.

Step 4: check the excess. If 2.50 g of copper(II) oxide was added and 0.65 g of dry solid recovered, 2.50 − 0.65 = 1.85 g reacted. Leftover solid is the evidence that the acid was used up.

Worked example 3: choosing between routes

Which route would you use for (a) magnesium sulfate, (b) lithium chloride?

(a) Magnesium is moderately reactive: add excess magnesium ribbon to dilute sulfuric acid, filter and crystallise. Mg + H₂SO₄ → MgSO₄ + H₂

(b) Lithium is too reactive for route 1, and lithium hydroxide is an alkali: use route 3 with hydrochloric acid. LiOH + HCl → LiCl + H₂O

Making insoluble salts by precipitation (3.5.2 c)

An insoluble salt is made by mixing two solutions: one containing its positive ion and one containing its negative ion. The salt forms at once as a precipitate.

To choose the two solutions, pick soluble compounds:

• All sodium, potassium and ammonium compounds are soluble.
• All nitrates are soluble.

So a nitrate usually supplies the metal ion, and a sodium or potassium compound usually supplies the negative ion.

Worked example 4: naming the reagents for an insoluble salt

Name two solutions you would mix to make (a) silver bromide, (b) calcium carbonate, and write the equations.

(a) Silver ions from silver nitrate solution; bromide ions from sodium bromide solution. A cream precipitate forms.

AgNO₃(aq) + NaBr(aq) → AgBr(s) + NaNO₃(aq)

(b) Calcium ions from calcium chloride solution; carbonate ions from sodium carbonate solution. A white precipitate forms.

CaCl₂(aq) + Na₂CO₃(aq) → CaCO₃(s) + 2NaCl(aq)

Method

1. Mix the two solutions in a beaker and stir.
2. Filter to collect the precipitate on the filter paper.
3. Wash the solid with distilled water to remove the soluble salt left with it.
4. Dry the solid in a warm oven or between filter papers.

Precipitation in water treatment

Precipitation takes ions out of solution, so it is used to clean water.

• Drinking water: in hard water areas, calcium ions can be removed by precipitating them as calcium carbonate.
• Effluent (waste water): dissolved metal ions such as copper(II) or zinc can be removed by adding an alkali so that insoluble metal hydroxides precipitate; phosphate can be removed as an insoluble phosphate. The solids are then settled or filtered out.

Common errors

• Calling every base an alkali: copper(II) oxide does not dissolve, so it is not an alkali.
• Including water as a product when ammonia reacts with an acid.
• Saying “limewater goes cloudy” without naming calcium carbonate as the product.
• Choosing route 1 for copper (too unreactive) or sodium (too reactive).
• Forgetting to repeat the indicator method without indicator.
• For an insoluble salt, naming a solid as a reagent instead of two solutions of soluble compounds.

Where next

• Recap the whole unit quickly with the acids, bases and salts revision notes.
• Test yourself with the acids, bases and salts practice questions.
• Track progress on the course hub and the printable checklist, or take a free 10-minute diagnostic.

Official syllabus

OxfordAQA International GCSE Chemistry (9202) specification, Version 6.2, for exams May/June 2018 onwards, published by OxfordAQA. This guide covers section 3.5.1 The properties of acids and bases and section 3.5.2 Preparation of salts.