Overview
What you’ll learn
- State the solubility rules for common salts in water and use them to predict precipitates.
- Choose a preparation method for a given salt — excess insoluble solid, titration or precipitation — and describe each with equations.
- Describe how to obtain pure crystals from a solution (filtration, evaporation to the point of crystallisation, washing and drying) and carry out water-of-crystallisation calculations.
- Recall and apply tests for the cations Al³⁺, NH₄⁺, Ca²⁺, Cu²⁺, Fe²⁺, Fe³⁺ and Zn²⁺ using aqueous sodium hydroxide and aqueous ammonia.
- Recall and apply tests for carbonate, chloride, iodide, nitrate and sulfate ions.
- Recall and apply tests for hydrogen, oxygen, carbon dioxide, ammonia, chlorine, sulfur dioxide and water vapour.
- Deduce the identity of an unknown salt by combining observations from several tests.
In 30 seconds
- 1Solubility families: all sodium, potassium and ammonium salts dissolve; all nitrates dissolve; chlorides dissolve except AgCl and PbCl₂; sulfates dissolve except BaSO₄, PbSO₄ and CaSO₄; most carbonates are insoluble (Group 1 and ammonium excepted).
- 2Choosing the method: soluble salt from an insoluble base/carbonate/metal → excess-solid method; soluble salt of Na, K or NH₄ → titration; insoluble salt → precipitation from two soluble solutions.
- 3Add the insoluble reactant in excess so every trace of acid reacts, then filter it off before evaporating the filtrate to the point of crystallisation.
- 4Titration is needed for sodium/potassium/ammonium salts because no solid is left over to filter — volumes must be measured exactly.
- 5Precipitation pairs two soluble solutions whose ions swap to form the insoluble salt: Ag⁺(aq) + Cl⁻(aq) → AgCl(s), then filter, wash with distilled water and dry.
- 6Water of crystallisation is water built into the lattice in a fixed ratio: CuSO₄·5H₂O (blue) loses it on heating to leave white CuSO₄; anhydrous copper(II) sulfate turning blue is the standard test for water.
- 7Cation tests: Cu²⁺ blue ppt, Fe²⁺ green ppt (browns on standing), Fe³⁺ red-brown ppt — none dissolve in excess; Ca²⁺, Al³⁺ and Zn²⁺ give white ppts, but only the Al³⁺ and Zn²⁺ ones dissolve in excess NaOH; only Zn²⁺ also dissolves in excess NH₃.
- 8Key gas tests: limewater milky = CO₂; acidified KMnO₄ decolourised = SO₂; bleached damp litmus = Cl₂; red litmus turned blue = NH₃; popping splint = H₂; relighting splint = O₂; NH₄⁺ + warm NaOH gives that ammonia.
Notes · 1 of 6
Which salts dissolve?
Five family rules cover almost every salt you will meet; the exceptions are examiners favourites.
Before you can prepare or identify a salt you must predict its solubility. The rules hang on the ions, not the source: a nitrate group drags any partner into solution, while a carbonate group usually drags it out.
| Family | Soluble? | Exceptions to memorise |
|---|---|---|
| Sodium, potassium, ammonium salts | Always soluble | None — this trio never lets you down |
| Nitrates | All soluble | None |
| Ethanoates | All soluble | None |
| Chlorides | Soluble | Silver chloride (AgCl), lead(II) chloride (PbCl₂) |
| Sulfates | Soluble | Barium sulfate, lead(II) sulfate, calcium sulfate |
| Carbonates | Mostly insoluble | Soluble ones: Group 1 (Na₂CO₃, K₂CO₃) and (NH₄)₂CO₃ |
| Hydroxides | Mostly insoluble | Soluble ones: NaOH, KOH (Ca(OH)₂ sparingly soluble) |
Notes · 2 of 6
Preparing soluble salts
Excess solid, or a measured titration — the solubility of your starting materials decides.
Route 1 — the excess-solid method (insoluble base, carbonate or metal): warm the dilute acid, stir in the insoluble solid in excess so all the acid is used up, filter off the leftover solid, then evaporate the filtrate to the point of crystallisation and let it cool so crystals grow. Example: copper(II) sulfate from copper(II) oxide and dilute sulfuric acid, CuO + H₂SO₄ → CuSO₄ + H₂O.
Route 2 — the titration method (soluble alkali or carbonate of Na, K or NH₄): there is no insoluble solid to leave in excess, so the exact reacting volume is found with an indicator, then repeated without the indicator, and the solution evaporated to give crystals. Example: NaOH + HNO₃ → NaNO₃ + H₂O.
| Target salt | Reagent pair | Method |
|---|---|---|
| Zinc sulfate | ZnO (insoluble) + dilute H₂SO₄ | Excess solid, filter, crystallise |
| Copper(II) chloride | CuO (insoluble) + dilute HCl | Excess solid, filter, crystallise |
| Sodium sulfate | NaOH (soluble) + dilute H₂SO₄ | Titration, then evaporate |
| Ammonium nitrate | NH₃(aq) + dilute HNO₃ | Titration, then evaporate |
| Barium sulfate (insoluble) | BaCl₂(aq) + Na₂SO₄(aq) | Precipitation, filter, wash, dry |
Evaporate only to the point of crystallisation (small crystals form on a glass rod placed in a cooling drop), then leave the rest to cool slowly. Boiling to dryness spits hot solution, traps impurities and can decompose the salt.
Notes · 3 of 6
Preparing insoluble salts by precipitation
Two soluble solutions swap ions; the insoluble salt falls out and is filtered, washed and dried.
An insoluble salt cannot be grown from an acid-and-solid reaction, because you could never separate it from the leftover solid. Instead, make it fall out of solution: pick two soluble solutions whose ions can swap to give your target salt plus a soluble bystander. Silver chloride: AgNO₃(aq) + NaCl(aq) → AgCl(s) + NaNO₃(aq). Lead(II) iodide, the golden shower of school chemistry: Pb(NO₃)₂(aq) + 2KI(aq) → PbI₂(s) + 2KNO₃(aq).
The collected precipitate still carries a coating of the soluble leftover salt, so it is washed with distilled water on the filter paper and then dried — pressed between filter papers or left in a warm oven. The ionic equation strips the story to its core: Ag⁺(aq) + Cl⁻(aq) → AgCl(s).
Notes · 4 of 6
Water of crystallisation
Some crystals lock water molecules into their lattice — and the ratio can be calculated.
A hydrated crystal contains water molecules woven into its lattice in a fixed whole-number ratio — the water of crystallisation. Copper(II) sulfate crystallises as CuSO₄·5H₂O, blue crystals that heat to white anhydrous CuSO₄: CuSO₄·5H₂O(s) → CuSO₄(s) + 5H₂O(g). Add water back and the blue colour returns — which is why anhydrous copper(II) sulfate is the classic test for water. (Cobalt(II) chloride paper, blue to pink, does the same job.)
Yield bookkeeping belongs here too: the anhydrous salt has a smaller Mr than the hydrate, so quote which one you mean. A preparation that isolates 6.89 g of ZnSO₄·7H₂O when the theory allows 8.61 g ran at 6.89 ÷ 8.61 = 80% yield — the usual losses are crystals left in the beaker, solution lost in transfer, and the filtering and washing steps.
Notes · 5 of 6
Identifying cations: NaOH and ammonia
Two reagents, seven cations — the precipitate colours and their behaviour in excess do all the work.
Add a few drops of aqueous sodium hydroxide to the unknown solution: metal hydroxides that are insoluble crash out as coloured precipitates. Then add more NaOH — some white precipitates dissolve in excess (they are amphoteric, a hint of Chapter 6). A second run with aqueous ammonia splits the look-alikes.
| Cation | With NaOH(aq), a few drops | In excess NaOH | In excess NH₃(aq) |
|---|---|---|---|
| Cu²⁺ | Blue precipitate | Insoluble | Insoluble (deep blue solution) |
| Fe²⁺ | Green precipitate, browns on standing | Insoluble | Insoluble |
| Fe³⁺ | Red-brown precipitate | Insoluble | Insoluble |
| Ca²⁺ | White precipitate | Insoluble | No precipitate |
| Al³⁺ | White precipitate | Soluble (colourless) | Insoluble |
| Zn²⁺ | White precipitate | Soluble (colourless) | Soluble (colourless) |
| NH₄⁺ | No precipitate; on warming, pungent gas turns damp red litmus blue | — | — |
The precipitates themselves are hydroxides, so ionic equations are quick: Cu²⁺(aq) + 2OH⁻(aq) → Cu(OH)₂(s); Fe³⁺(aq) + 3OH⁻(aq) → Fe(OH)₃(s). The iron(II) green precipitate browns from the surface because air oxidises Fe(OH)₂ to Fe(OH)₃ — a favourite one-mark observation.
Notes · 6 of 6
Identifying anions and gases
Acidified silver nitrate and barium nitrate unmask anions; eight gas tests finish the toolkit.
| Anion | Test | Positive result |
|---|---|---|
| Carbonate CO₃²⁻ | Add dilute acid | Effervescence; gas turns limewater milky |
| Chloride Cl⁻ | Acidify with dilute HNO₃, add AgNO₃(aq) | White precipitate (AgCl) |
| Iodide I⁻ | Acidify with dilute HNO₃, add AgNO₃(aq) | Yellow precipitate (AgI) |
| Nitrate NO₃⁻ | Add NaOH(aq) and aluminium foil, warm | Pungent gas turns damp red litmus blue (ammonia) |
| Sulfate SO₄²⁻ | Acidify with dilute HNO₃, add Ba(NO₃)₂(aq) | White precipitate (BaSO₄) |
| Gas | Test and result |
|---|---|
| Hydrogen H₂ | Lighted splint → pop |
| Oxygen O₂ | Glowing splint → relights |
| Carbon dioxide CO₂ | Bubble through limewater → milky |
| Ammonia NH₃ | Damp red litmus → blue; pungent |
| Chlorine Cl₂ | Damp litmus → bleached white |
| Sulfur dioxide SO₂ | Acidified potassium manganate(VII) → purple to colourless; sharp smell |
| Water vapour | Anhydrous copper(II) sulfate white → blue (or cobalt paper blue → pink) |
The acidifying step before silver nitrate or barium nitrate matters: unacidified tests are fooled by carbonate ions, which also give white precipitates (Ag₂CO₃, BaCO₃). Dilute nitric acid destroys carbonates first, leaving chloride, iodide and sulfate to answer for themselves.
Revise
Key terms
The words examiners expect you to use precisely.
- Soluble salt
- A salt that dissolves in water, e.g. all sodium, potassium, ammonium salts and all nitrates.
- Insoluble salt
- A salt that does not dissolve appreciably in water, e.g. AgCl, PbCl₂, BaSO₄, CaCO₃.
- Precipitate
- The insoluble solid that forms when two solutions react: Ba²⁺(aq) + SO₄²⁻(aq) → BaSO₄(s).
- Excess-solid method
- Salt preparation stirring an insoluble base, metal or carbonate in excess into acid, filtering, then crystallising the filtrate.
- Titration
- Measured reaction of an acid with an alkali to find the exact reacting volumes — the route to pure Na, K and NH₄ salts.
- Point of crystallisation
- The stage where a hot solution holds just enough dissolved salt for crystals to appear on cooling.
Show 6 more key termsShow fewer key terms
- Water of crystallisation
- Water molecules incorporated into a crystal lattice in a fixed ratio, written as a dot formula such as CuSO₄·5H₂O.
- Hydrated / anhydrous
- Hydrated: containing water of crystallisation. Anhydrous: with all lattice water removed, e.g. white CuSO₄.
- Qualitative analysis
- Identifying which ions or gases are present using characteristic test-tube reactions, rather than measuring amounts.
- Limewater
- Calcium hydroxide solution; turns milky with CO₂ because insoluble CaCO₃ forms.
- Acidified (test)
- Dilute nitric acid added first so carbonate ions cannot fake a positive silver nitrate or barium nitrate result.
- Filtrate / residue
- The filtrate is the liquid passing through the filter paper; the residue is the solid left on it.
Revise
Common mistakes
Where students most often lose marks in this chapter — and what to do instead.
Mistake: Treating every chloride and sulfate as soluble.
The rules carry named exceptions.
Instead: AgCl and PbCl₂ are insoluble chlorides; BaSO₄, PbSO₄ and CaSO₄ are insoluble sulfates.
Mistake: Choosing titration for copper(II) sulfate or excess solid for sodium nitrate.
The method is fixed by the solubility of the starting materials.
Instead: Insoluble solid + acid → excess-solid method; Na/K/NH₄ salts → titration; insoluble product → precipitation.
Show 10 more mistakesShow fewer mistakes
Mistake: Boiling the salt solution to complete dryness to get crystals.
The spitting solution traps impurities and may decompose the salt.
Instead: Evaporate to the point of crystallisation, then let the concentrated solution cool slowly.
Mistake: Adding the solid reactant in exactly stoichiometric amount.
You cannot measure the exact end-point by eye, so acid survives.
Instead: Add in excess, then filter the surplus away — guaranteeing no acid remains.
Mistake: Skipping the acidification before silver nitrate or barium nitrate.
Carbonates also give white precipitates and cause false positives.
Instead: Add dilute nitric acid first; it removes CO₃²⁻ as CO₂ so only Cl⁻/I⁻/SO₄²⁻ respond.
Mistake: Mixing up iron(II) and iron(III) precipitate colours.
The colours anchor the whole identification.
Instead: Fe²⁺ → green precipitate that browns on standing; Fe³⁺ → red-brown from the start.
Mistake: Claiming calcium ions from any white precipitate.
Ca²⁺, Al³⁺ and Zn²⁺ all give white precipitates with NaOH.
Instead: Test behaviour in excess NaOH and in excess NH₃ to split the three.
Mistake: Concluding Zn²⁺ whenever a precipitate dissolves in excess NaOH.
Al³⁺ behaves identically with NaOH.
Instead: Only Zn²⁺ also dissolves in excess aqueous ammonia — that is the discriminator.
Mistake: Saying only CO₂ turns limewater milky, so any milky result is CO₂.
SO₂ also clouds limewater.
Instead: Confirm with acidified potassium manganate(VII): SO₂ decolourises it, CO₂ does not.
Mistake: Testing for the nitrate ion with silver nitrate.
Silver nitrate catches chloride and iodide, not nitrate.
Instead: Nitrates need NaOH plus aluminium foil, warmed — ammonia released proves NO₃⁻.
Mistake: Washing a precipitate with tap water.
Tap water carries dissolved ions that contaminate the sample.
Instead: Wash with distilled water, then dry between filter papers.
Mistake: Drying crystals in a hot oven to remove every trace of water.
Heat strips water of crystallisation and spoils the hydrate.
Instead: Press between filter papers or use a desiccator; leave the lattice water in place.
Check yourself
Quick quiz
20 questions with instant explanations. About 5 minutes.
Question 1 of 3
- Question 1: current question
- Question 2: not answered
- Question 3: not answered
Apply solubility rules
Which salt is insoluble in water?
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Practise
Topical questions
Exam-style practice in three tiers, from recall to data questions.
- 4Foundation
- 5Standard
- 3Challenging
12 questions worth 56 marks, each with a model answer, mark scheme and examiner tips.
Practise topical questionsExam practice
Practice paper
A 30-mark paper set out like a school paper, with a separate mark scheme. Give yourself 40 min.
Question paper
Chapter 7 Practice Paper
30 marks · 40 min · 4 pages
5 multiple-choice questions and structured questions with answer lines. Chapter 7 (solubility rules, salt preparation, water of crystallisation, cation/anion/gas tests).
Mark scheme
PROChapter 7 Practice Paper — mark scheme
Answers · Marking points · 3 pages
Model answers with a mark for each creditworthy point — for marking at home or in class.
Where this chapter is tested
Each school sets its own scope — check your assessment outline.
Questions
Frequently asked questions
What does Sec 3 Chemistry Chapter 7 cover?
It covers sub-topic 3.2 of O-Level Pure Chemistry 6092 — solubility rules, choosing and describing salt-preparation methods, water of crystallisation, purity and yield — plus sub-topic 1.3, the full test tables for cations (NaOH and ammonia), anions and gases.
How do I quickly decide which salt-preparation method to use?
Look at solubility twice: if the target salt is soluble, use acid + excess insoluble solid (or titration when both reactants are soluble, as for sodium, potassium and ammonium salts). If the target salt is insoluble, precipitate it by mixing two soluble solutions, then filter, wash and dry.
Why is zinc the only cation whose precipitate dissolves in both excess NaOH and excess ammonia?
Zinc hydroxide is amphoteric (it reacts with the strong base in excess NaOH) and also forms soluble complex ammines with excess ammonia. Aluminium hydroxide shares the amphoteric half but not the ammonia half, and calcium hydroxide does neither in ammonia — which is exactly how the three white precipitates are told apart.
Why must the nitric acid be added before silver nitrate or barium nitrate tests?
Carbonate ions would otherwise give white precipitates too (silver carbonate, barium carbonate) and fake a positive result. Dilute nitric acid destroys the carbonate first, releasing CO₂, so only chloride, iodide or sulfate can precipitate afterwards.
When is this chapter usually tested in school?
Salts and qualitative analysis close out Sec 3, so expect heavy weighting in the end-of-year examination — and qualitative analysis deductions reappear in every O-Level Paper 2. Scope varies by school, so check your school’s assessment outline.
Up next · in preparation
Sec 4 Chapter 1: Oxidation and Reduction