Acids, bases & salts Banker

1. (a) Which salt is insoluble in water?A sodium carbonateB barium sulfateC copper(II) nitrateD potassium chloride [1]

A student makes crystals of hydrated copper(II) sulfate by adding excess copper(II) oxide to warm dilute sulfuric acid.

1. (a) Write the word equation for the reaction. [1]
2. (b) Explain why copper(II) oxide is added in excess. [1]
3. (c) Describe how the excess copper(II) oxide is removed. [1]
4. (d) Describe how crystals of hydrated copper(II) sulfate are then obtained from the filtrate. [3]

A student adds zinc carbonate powder, a little at a time, to dilute hydrochloric acid to make zinc chloride.

1. (a) Write the balanced equation for the reaction. [2]
2. (b) Describe what is seen as the zinc carbonate is added. [1]
3. (c) State how the student knows that the zinc carbonate is in excess. [1]
4. (d) Suggest why the mixture is filtered before crystallising. [1]

State whether each salt is soluble or insoluble in water.

1. (a) sodium carbonate [1]
2. (b) calcium carbonate [1]
3. (c) silver chloride [1]
4. (d) ammonium sulfate [1]
5. (e) lead(II) nitrate [1]
6. (f) calcium sulfate [1]

A student makes zinc sulfate crystals by adding excess granulated zinc to dilute sulfuric acid in a beaker.

1. (a) Describe what the student sees when the zinc is added to the acid. [1]
2. (b) Describe a test to identify the gas produced. [1]
3. (c) State how the student can tell that the reaction has finished. [1]
4. (d) Write the balanced equation for the reaction. [2]
5. (e) Explain why the solution should not be heated to dryness to obtain the crystals. [1]

Barium sulfate is an insoluble salt used in medical X-ray imaging.

1. (a) Name two aqueous solutions that could be mixed to make barium sulfate. [2]
2. (b) Write the ionic equation, with state symbols, for the formation of barium sulfate. [2]
3. (c) Describe how a pure, dry sample of barium sulfate is obtained from the mixture. [3]
4. (d) Name the method of preparation used. [1]

Potassium nitrate is used in fertilisers. It can be made from aqueous potassium hydroxide and dilute nitric acid.

1. (a) Explain why potassium nitrate cannot be made by adding an excess of a solid to the acid and filtering. [2]
2. (b) Describe how a titration is used to make a pure solution of potassium nitrate. [4]
3. (c) Write the equation for the reaction. [1]

A student heats hydrated magnesium sulfate, MgSO₄·xH₂O, in a crucible, cools and reweighs it, and repeats until the mass is constant.Mass of empty crucible = 21.40 gMass of crucible + hydrated salt = 26.32 gMass of crucible + salt after heating to constant mass = 23.80 g

1. (a) Calculate the mass of hydrated salt used and the mass of water lost. [2]
2. (b) Calculate the percentage by mass of water in the hydrated salt. Give your answer to 3 significant figures. [1]
3. (c) Use Mr MgSO₄ = 120 and Mr H₂O = 18 to find x. [2]
4. (d) Explain why the student heated the crucible to constant mass. [1]

1. (a) Which method is most suitable for preparing a sample of silver chloride?A titration of an acid with an alkaliB adding excess metal to an acid and filteringC mixing two aqueous solutions to form a precipitateD adding excess carbonate to an acid and filtering [1]

Some salts form crystals that contain water.

1. (a) State the meaning of the term hydrated, as used for a salt. [1]
2. (b) Write the equation for the reaction when water is added to anhydrous cobalt(II) chloride to form CoCl₂·6H₂O. Use the correct arrow. [2]
3. (c) State the colour change seen when water is added to anhydrous cobalt(II) chloride. [1]
4. (d) Explain how this change can be used as a test for water. [1]

2.50 g of hydrated copper(II) sulfate, CuSO₄·xH₂O, is heated until no more water is lost. 1.60 g of anhydrous copper(II) sulfate remains. Mr: CuSO₄ = 160, H₂O = 18.

1. (a) Calculate the number of moles of anhydrous copper(II) sulfate. [1]
2. (b) Calculate the number of moles of water lost. [1]
3. (c) Deduce the value of x. [1]
4. (d) Describe the colour change on heating. [1]

A student reacts 4.0 g of copper(II) oxide with excess dilute sulfuric acid and crystallises hydrated copper(II) sulfate.CuO + H₂SO₄ → CuSO₄ + H₂O, then CuSO₄ + 5H₂O → CuSO₄·5H₂OMr: CuO = 80, CuSO₄·5H₂O = 250.

1. (a) Calculate the maximum mass of CuSO₄·5H₂O that can be made. [2]
2. (b) The student obtains 9.8 g of crystals. Calculate the percentage yield. [1]
3. (c) Suggest one reason why the yield is less than 100%. [1]

1. (a) When 4.30 g of hydrated calcium sulfate, CaSO₄·xH₂O, is heated to constant mass, 3.40 g of anhydrous CaSO₄ remains. Mr: CaSO₄ = 136, H₂O = 18. What is the value of x?A 1B 2C 5D 10 [1]

25.0 cm³ of 0.200 mol/dm³ aqueous sodium hydroxide is exactly neutralised by 20.0 cm³ of dilute sulfuric acid. Mr: Na₂SO₄ = 142.

1. (a) Write the equation for the reaction. [1]
2. (b) Calculate the concentration of the sulfuric acid in mol/dm³. [3]
3. (c) Calculate the mass of sodium sulfate in the neutral solution. [2]

Washing soda crystals are sold in markets as a cleaning agent.

1. (a) Plan an experiment to find the value of x in washing soda, Na₂CO₃·xH₂O, a hydrated salt. Include how you would use your results. You may use common laboratory apparatus. [6]

1. (a) What is the colour of methyl orange in aqueous sodium hydroxide?A redB orangeC yellowD blue [1]

Name the products formed when dilute hydrochloric acid reacts with each substance. For (a) and (b) also give one observation.

1. (a) magnesium [2]
2. (b) calcium carbonate [2]
3. (c) sodium hydroxide solution [1]
4. (d) State the name given to the reaction in (c). [1]

A student tests household vinegar with different indicators.

1. (a) State the colour of litmus, methyl orange and thymolphthalein in vinegar (an acid). [3]
2. (b) A student dips red litmus paper into vinegar and sees no change. Explain why this does not show that vinegar is an acid. [1]
3. (c) Name an indicator that would give the approximate pH of the vinegar. [1]

Water can act as an acid or as a base.

1. (a) Define an acid and a base in terms of proton transfer. [2]
2. (b) NH₃ + H₂O ⇌ NH₄⁺ + OH⁻Identify the acid and the base in the forward reaction. Explain your answer. [2]
3. (c) HCl + NH₃ → NH₄ClWhich substance acts as a base in this reaction? [1]

A mixture of solid ammonium chloride and calcium hydroxide is warmed.

1. (a) Name the gas produced. [1]
2. (b) Complete the equation: 2NH₄Cl + Ca(OH)₂ → CaCl₂ + ___ + ___ [2]
3. (c) Describe the test for this gas. [1]
4. (d) Explain why farmers should not spread calcium hydroxide on soil at the same time as an ammonium fertiliser. [2]

A farmer in the Swat valley tests his soil and finds pH 4.5, which is too acidic for the crops he wants to grow.

1. (a) Name a substance that the farmer could add to raise the pH of the soil. [1]
2. (b) Write the ionic equation for the neutralisation of an acid by an alkali. [1]
3. (c) Explain, in terms of ions, why the pH of the soil increases. [2]
4. (d) Suggest why the farmer should not add too much of the substance. [1]

1. (a) Which oxide is amphoteric?A zinc oxideB calcium oxideC sulfur dioxideD carbon monoxide [1]

A teacher burns magnesium and sulfur separately in oxygen. Each product is shaken with water containing universal indicator.

1. (a) State the colour of universal indicator in the solution from magnesium oxide, and suggest its pH. [2]
2. (b) State the colour of universal indicator in the solution from sulfur dioxide, and suggest its pH. [2]
3. (c) State the general rule linking the type of element to the type of oxide. [1]

Classify each oxide as acidic, basic, amphoteric or neutral.

1. (a) magnesium oxide [1]
2. (b) sulfur dioxide [1]
3. (c) aluminium oxide [1]
4. (d) carbon monoxide [1]
5. (e) phosphorus(V) oxide [1]

Aluminium oxide is described as an amphoteric oxide.

1. (a) Define the term amphoteric. [1]
2. (b) Write the balanced equation for the reaction of aluminium oxide with hydrochloric acid. [2]
3. (c) Describe what you would do to show that aluminium oxide also behaves as an acidic oxide, and what you would see. [2]
4. (d) Explain why the oxide of a metal such as magnesium does not dissolve in aqueous sodium hydroxide. [1]

1. (a) Oxide W does not react with dilute hydrochloric acid or with aqueous sodium hydroxide. What is W?A carbon monoxideB carbon dioxideC zinc oxideD copper(II) oxide [1]

Universal indicator shows the pH of a solution.

1. (a) Match each substance to its likely pH: stomach acid, lemon juice, pure water, oven cleaner. pH values: 1, 3, 7, 13. [2]
2. (b) Name the ion present in excess in all alkaline solutions. [1]
3. (c) Describe how universal indicator paper is used to find the pH of a solution. [2]

A student has 0.10 mol/dm³ hydrochloric acid and 0.10 mol/dm³ ethanoic acid.

1. (a) Suggest which acid has the higher pH. [1]
2. (b) Explain the difference in pH in terms of ions. [2]
3. (c) Write the equation for the dissociation of ethanoic acid in water. [2]
4. (d) Magnesium ribbon is added to each acid. State how the rate of reaction compares. [1]

1. (a) Which solution has the lowest pH?A 0.1 mol/dm³ hydrochloric acidB 0.1 mol/dm³ ethanoic acidC 0.01 mol/dm³ hydrochloric acidD 0.1 mol/dm³ sodium hydroxide [1]

Three acids, P, Q and R, all at 0.50 mol/dm³ and in excess, each react with 0.12 g of magnesium ribbon.Acid: P, Q, RpH: 0.3, 2.5, 0.3Volume of hydrogen after 60 s / cm³: 92, 21, 95

1. (a) Which acids are strong? Explain your choice using the data. [2]
2. (b) Explain, in terms of particles, why acid Q gives less gas in 60 s. [2]
3. (c) All the magnesium eventually reacts in each acid. Explain why the final volume of hydrogen is the same in each case. [2]