Sec 3 Chemistry Salts & Qualitative Analysis topical questions

State whether each salt is soluble or insoluble in water.

1. (a)Sodium carbonate[1]
2. (b)Silver chloride[1]
3. (c)Calcium sulfate[1]
4. (d)Ammonium nitrate[1]

Name the gas identified by each test.

1. (a)Relights a glowing splint[1]
2. (b)Turns limewater milky[1]
3. (c)Bleaches damp litmus paper white[1]

State the colour of the precipitate formed when aqueous sodium hydroxide is added to solutions of each cation.

1. (a)Cu²⁺[1]
2. (b)Fe³⁺[1]
3. (c)Fe²⁺[1]
4. (d)Ca²⁺[1]

Name the reagent (and one observation) for testing each solution.

1. (a)A solution to show it contains sulfate ions[1]
2. (b)A solution to show it contains iodide ions[1]
3. (c)A solution to show it contains carbonate ions[1]
4. (d)A solution to show it contains nitrate ions[1]

Copy-and-learn colours: aqueous sodium hydroxide is added to three cation solutions.

1. (a)State the colour of each precipitate.[3]
2. (b)State what happens to the iron(II) precipitate after standing in air.[1]

For each salt, state whether it is soluble in water and name a preparation method that suits it.

1. (a)Barium sulfate[1]
2. (b)Sodium sulfate[1]
3. (c)Copper(II) sulfate[1]

A student prepares zinc sulfate crystals by warming dilute sulfuric acid and stirring in zinc oxide little by little until some solid remains. The mixture is filtered and the filtrate evaporated to the point of crystallisation and cooled. Explain (a) why the zinc oxide is added until no more dissolves, (b) why the mixture is filtered, and (c) why the filtrate is not evaporated to complete dryness.

Aqueous silver nitrate is mixed with aqueous sodium chloride.

1. (a)Write the full chemical equation for the reaction.[2]
2. (b)Write the ionic equation and explain why the sodium and nitrate ions do not appear in it.[2]

25.0 cm³ of aqueous sodium hydroxide is neutralised by exactly 20.0 cm³ of 0.0500 mol/dm³ dilute sulfuric acid (H₂SO₄ + 2NaOH → Na₂SO₄ + 2H₂O).

1. (a)Calculate the moles of sulfuric acid used.[2]
2. (b)Calculate the concentration of the sodium hydroxide in mol/dm³.[2]
3. (c)Explain why titration, not the excess-solid method, is used for this salt.[1]

4.92 g of hydrated magnesium sulfate crystals, MgSO₄·xH₂O, are heated to constant mass. The anhydrous MgSO₄ left weighs 2.40 g. (Mr: MgSO₄ = 120, H₂O = 18.)

1. (a)Calculate the moles of anhydrous MgSO₄.[2]
2. (b)Calculate the moles of water driven off.[1]
3. (c)Deduce x in MgSO₄·xH₂O.[2]

Solution P gives a white precipitate with a few drops of aqueous sodium hydroxide. The precipitate dissolves when more sodium hydroxide is added. A second sample of P with aqueous ammonia gives a white precipitate that does not dissolve in excess.

1. (a)Name the three cations that could give a white precipitate with NaOH.[1]
2. (b)Use the ammonia result to identify the cation in P, explaining your reasoning.[2]
3. (c)Write the ionic equation for the formation of the white precipitate.[2]

Write ionic equations, with state symbols, for two precipitation reactions.

1. (a)Silver nitrate solution + potassium bromide solution (cream precipitate).[2]
2. (b)Lead(II) nitrate solution + potassium iodide solution (yellow precipitate).[2]

20.0 cm³ of dilute nitric acid neutralises 25.0 cm³ of 0.100 mol/dm³ potassium hydroxide exactly (HNO₃ + KOH → KNO₃ + H₂O).

1. (a)Calculate the concentration of the nitric acid in mol/dm³.[3]
2. (b)Name the salt present in the neutral solution.[1]

Solution M is blue in colour. Adding aqueous sodium hydroxide gives a pale blue precipitate that does not dissolve in excess NaOH. Acidifying another sample with dilute nitric acid and adding barium nitrate solution gives a white precipitate.

1. (a)Identify the cation and the anion in M.[2]
2. (b)Name M and write the ionic equation for the precipitate formed with sodium hydroxide.[3]

A white solid Z is analysed. It dissolves in water. Warming a sample with aqueous sodium hydroxide releases a pungent gas that turns damp red litmus blue. Acidifying a second sample with dilute nitric acid and adding barium nitrate solution gives a white precipitate.

1. (a)Identify the gas released with sodium hydroxide, name the ion it proves, and write the ionic equation for its formation.[3]
2. (b)Name the white precipitate and the ion it proves.[2]
3. (c)Write the formula of Z.[1]

Excess zinc oxide is reacted with dilute sulfuric acid and the solution crystallised. The planned charge of 2.43 g of ZnO (Mr = 81) reacts completely. The actual crop of zinc sulfate-7-water crystals, ZnSO₄·7H₂O (Mr = 287), is 6.89 g.

1. (a)Calculate the moles of ZnO used.[2]
2. (b)Calculate the theoretical mass of ZnSO₄·7H₂O expected.[2]
3. (c)Calculate the percentage yield.[2]

You are to prepare a pure, dry sample of barium sulfate from common laboratory reagents.

1. (a)Choose two suitable reagent solutions and explain why both must be soluble.[2]
2. (b)Outline the steps after mixing, up to the dry solid.[2]
3. (c)Write the ionic equation for the reaction.[2]

A white solid W dissolves in water to give a colourless solution. Adding aqueous sodium hydroxide gives a white precipitate that dissolves in excess NaOH; aqueous ammonia gives a white precipitate that does not dissolve in excess. Adding dilute hydrochloric acid to a fresh sample of the solution gives a white precipitate. Warming a further sample with NaOH and aluminium foil releases a pungent gas that turns damp red litmus blue.

1. (a)Deduce the cation in W, explaining how the chloride test separates the two remaining candidates.[3]
2. (b)Deduce the anion in W, naming the test used.[2]
3. (c)Name W.[1]