1. 1
   A student heats a small sample of copper carbonate, CuCO3, inside a sealed flask fitted with a pressure sensor, so that no gas can enter or escape, to test whether mass is conserved during its thermal decomposition: CuCO3 -> CuO + CO2. The flask contains 12.4 g of copper carbonate at the start (Mr: CuCO3 = 124, CuO = 80, CO2 = 44).

   (a)

   How does the total mass measured inside the sealed flask after heating compare with the total mass before heating?

   [1 mark]

   • AIt increases
   • BIt decreases
   • CIt stays the same
   • DIt cannot be determined

   (b)

   If the flask were opened immediately after heating, what would happen to the reading on a balance holding the flask and its contents?

   [1 mark]

   • AIt would stay the same
   • BIt would decrease, as carbon dioxide gas escapes
   • CIt would increase, as air enters
   • DIt cannot be predicted

   (c)

   Calculate the mass of carbon dioxide gas produced when the 12.4 g sample of copper carbonate decomposes completely.

   [2 marks]

   Total for question 1: 4 marks

2. 2
   A quality-control chemist compares two compounds used in toothpaste: sodium fluoride, NaF, and calcium carbonate, CaCO3 (Ar: Na = 23, F = 19, Ca = 40, C = 12, O = 16).

   (a)

   What is the relative formula mass of sodium fluoride, NaF?

   [1 mark]

   • A42
   • B61
   • C23
   • D19

   (b)

   What is the relative formula mass of calcium carbonate, CaCO3?

   [1 mark]

   • A116
   • B88
   • C60
   • D100

   (c)

   Calculate the percentage by mass of calcium in calcium carbonate.

   [2 marks]

   Total for question 2: 4 marks

3. 3
   A chemist analyses an unknown colourless liquid and finds it is made of 52.2% carbon, 13.0% hydrogen and 34.8% oxygen by mass (Ar: C = 12, H = 1, O = 16).

   (a)

   Calculate the empirical formula of the liquid, showing your working.

   [3 marks]

   (b)

   The relative formula mass of the liquid is found separately to be 46. Determine its molecular formula, and explain why the empirical and molecular formulae are the same in this case.

   [4 marks]

   Total for question 3: 7 marks

4. 4
   An industrial chemist is scaling up the production of calcium oxide by heating calcium carbonate: CaCO3 -> CaO + CO2. In one batch, 500 kg of calcium carbonate is heated until it decomposes completely (Ar: Ca = 40, C = 12, O = 16; assume the molar gas volume is 24 dm3/mol).

   (a)

   Calculate the mass of calcium oxide produced when the 500 kg of calcium carbonate decomposes completely, showing the number of moles involved.

   [6 marks]

   (b)

   State the mass of carbon dioxide gas released during this decomposition, explaining your reasoning using the law of conservation of mass, then calculate the number of moles of carbon dioxide this represents and the volume this gas occupies at room temperature and pressure.

   [6 marks]

   Total for question 4: 12 marks

5. 5
   A student weighs out 9.2 g of ethanol, C2H5OH (Mr = 46), to use in an experiment testing its combustion (Avogadro's constant = 6.02 x 10^23 per mole).

   (a)

   How many moles of ethanol are in the 9.2 g sample?

   [1 mark]

   • A2.0 mol
   • B0.02 mol
   • C4.6 mol
   • D0.2 mol

   (b)

   Approximately how many molecules of ethanol does the 9.2 g sample contain?

   [1 mark]

   • A1.2 x 10^23
   • B6.02 x 10^23
   • C3.0 x 10^22
   • D1.2 x 10^24

   (c)

   The student now wants exactly 0.5 mol of ethanol instead. Calculate the mass of ethanol this represents.

   [2 marks]

   Total for question 5: 4 marks

6. 6
   A technician dissolves 8.0 g of sodium hydroxide, NaOH (Mr = 40), in water and makes the total volume up to 2 dm3 of solution.

   (a)

   How many moles of sodium hydroxide were dissolved?

   [1 mark]

   • A2.0 mol
   • B0.2 mol
   • C0.02 mol
   • D20 mol

   (b)

   What is the concentration of this solution, in mol/dm3?

   [1 mark]

   • A4.0 mol/dm3
   • B0.4 mol/dm3
   • C0.1 mol/dm3
   • D1.0 mol/dm3

   (c)

   Calculate the concentration of this solution in g/dm3.

   [2 marks]

   Total for question 6: 4 marks

7. 7
   A brewery's quality lab titrates a 20.0 cm3 sample of dilute sulfuric acid used to clean equipment against a standard potassium hydroxide solution of concentration 0.150 mol/dm3, to check its concentration before disposal: H2SO4 + 2KOH -> K2SO4 + 2H2O. Exactly 24.0 cm3 of the potassium hydroxide solution is needed to neutralise the acid.

   (a)

   Calculate the number of moles of potassium hydroxide used in the titration.

   [3 marks]

   (b)

   Using the balanced equation and your answer to part (a), calculate the concentration of the sulfuric acid, in mol/dm3.

   [4 marks]

   Total for question 7: 7 marks

8. 8
   A laboratory checks the concentration of a batch of dilute nitric acid. A 25.0 cm3 sample of the acid is titrated against a standard sodium hydroxide solution of concentration 0.200 mol/dm3: HNO3 + NaOH -> NaNO3 + H2O. Exactly 15.0 cm3 of the sodium hydroxide solution is needed for neutralisation. A separate 25.0 cm3 sample of the same nitric acid is then reacted with excess calcium carbonate: CaCO3 + 2HNO3 -> Ca(NO3)2 + H2O + CO2 (Mr: NaNO3 = 85; assume the molar gas volume is 24 dm3/mol).

   (a)

   Calculate the concentration of the nitric acid, in mol/dm3, using the titration results, and calculate the mass of solid sodium nitrate that could be obtained by evaporating the neutralised solution.

   [6 marks]

   (b)

   A 25.0 cm3 sample of the same nitric acid (concentration 0.12 mol/dm3) is reacted with excess calcium carbonate. Calculate the number of moles of nitric acid in this sample, the number of moles of carbon dioxide gas produced, and the volume of this gas at room temperature and pressure.

   [6 marks]

   Total for question 8: 12 marks

End of questions