Unit 1 – Titration Calculations
Unit 1 - Jan 2009
Q1.
Magnesium oxide reacts with dilute hydrochloric acid according to the following
 equation.

How many moles of magnesium oxide, MgO, are required to neutralize 20 cm3 of
 0.50 mol dm−3 hydrochloric acid, HCl?
   A     0.0010
   B     0.0050
   C     0.010
   D     0.020

(Total for question = 1 mark)

Unit 1 - Jun 2009
Q2.
Calculate the mass of calcium hydroxide, Ca(OH)2, present in 100 cm3 of a 0.100 mol dm−3 solution.
[Assume the molar mass of Ca(OH)2 is 74.0 g mol−1.]
   A     0.570 g
   B     0.740 g
   C       1.85 g
   D       3.70 g
(Total for question = 1 mark)

Q3.
Sodium hydrogensulfate, NaHSO4, reacts with sodium hydroxide, NaOH, as shown below.

0.0100 mol of sodium hydrogensulfate is neutralized with dilute sodium hydroxide,
 concentration 0.200 mol dm−3.
Calculate the volume of sodium hydroxide required.
   A     20.0 cm3
   B     50.0 cm3
   C     100 cm3
   D     500 cm3
(Total for question = 1 mark)

Unit 1 - Jan 2010
Q4.
Magnesium chloride can be made by reacting solid magnesium carbonate, MgCO3, with
       dilute hydrochloric acid.
(a) Write an equation for the reaction, including state symbols.
(2)

(b) Give TWO observations you would make when the reaction is taking place.
(2)
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(c) In an experiment to make crystals of hydrated magnesium chloride, MgCl2.6H2O, magnesium carbonate was added to 25 cm3 of hydrochloric acid with concentration 2.0 mol dm-3. The molar mass of magnesium carbonate is 84.3 g mol-1.
(i) How many moles of acid are used in the reaction?
(1)
(ii) What mass of magnesium carbonate, in grams, reacts with this amount of acid?
(1)

(iii) Suggest why slightly more than this mass of magnesium carbonate is used in practice.
(1)
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 (iv) How would you separate the magnesium chloride solution from the reaction mixture in (iii)?
(1)
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(v) The magnesium chloride solution was left to crystallise. The crystals were separated and dried carefully. A sample of 3.75g of hydrated crystals, MgCl2.6H2O, which have molar mass 203.3 g mol-1, was obtained. Calculate the percentage yield of this reaction.
(2)

(vi) Give ONE reason why the yield of crystals is less than 100%, even when pure compounds are used in the preparation.
(1)
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 (d) Lattice energies can be measured using the Born-Haber cycle, or calculated from electrostatic theory. Lattice energies of magnesium chloride and magnesium iodide are shown below.

(i) What does this data indicate about the bonding in magnesium chloride?
(1)
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*(ii) Explain why there is a greater difference between the experimental (Born-Haber) and theoretical lattice energies for magnesium iodide, MgI2, compared with magnesium chloride.
(2)
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 (e) Blood plasma typically contains 20 parts per million (ppm) of magnesium, by mass.
(i) Calculate the mass of magnesium, in grams, present in 100 g of plasma.
(1)

(ii) Magnesium chloride can be used as a supplement in the diet to treat patients with low amounts of magnesium in the blood. Suggest ONE property which makes it more suitable for this purpose than magnesium carbonate.
(1)
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(Total for question = 16 marks)

Unit 1 - Jun 2010
Q5.
Which of these solutions does not contain the same total number of ions as the others?
   A     10.00 cm3 of 0.100 mol dm-3  NaCl(aq)
   B     20.00 cm3 of 0.050 mol dm-3  NaCl(aq)
   C     20.00 cm3 of 0.050 mol dm-3  MgCl2(aq)
   D     13.33 cm3 of 0.050 mol dm-3  MgCl2(aq)
(Total for question = 1 mark)

Unit 1 - Jan 2011
Q6.
(a) Coral reefs are produced by living organisms and predominantly made up of calcium carbonate. It has been suggested that coral reefs will be damaged by global warming because of the increased acidity of the oceans due to higher concentrations of carbon dioxide.
(i) Write a chemical equation to show how the presence of carbon dioxide in water results in the formation of carbonic acid. State symbols are not required
(1)

(ii) Write the ionic equation to show how acids react with carbonates. State symbols are not required.
(2)

(b) One method of determining the proportion of calcium carbonate in a coral is to dissolve a known mass of the coral in excess acid and measure the volume of carbon dioxide formed.
      In such an experiment, 1.13 g of coral was dissolved in 25 cm3 of hydrochloric acid (an excess) in a conical flask. When the reaction was complete, 224 cm3 of carbon dioxide had been collected over water using a 250 cm3 measuring cylinder.
(i) Draw a labelled diagram of the apparatus that could be used to carry out this experiment.
(2)

(ii) Suggest how you would mix the acid and the coral to ensure that no carbon dioxide escaped from the apparatus.
(1)
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 (iii) Calculate the number of moles of carbon dioxide collected in the experiment.
                    [The molar volume of any gas is 24 000 cm3 mol−1 at room temperature and pressure.]
(1)

(iv) Complete the equation below for the reaction between calcium carbonate and hydrochloric acid by inserting the missing state symbols.
(1)
CaCO3(..............) + 2HCl(..............) → CaCl2(..............) + H2O(l) + CO2(..............)
(v) Calculate the mass of 1 mol of calcium carbonate.
      [Assume relative atomic masses: Ca = 40, C = 12, O = 16.]
(1)

(vi) Use your data and the equation in (iv) to calculate the mass of calcium carbonate in the sample and the percentage by mass of calcium carbonate in the coral.
                    Give your final answer to three significant figures.
(2)

(vii) When this experiment is repeated, the results are inconsistent. Suggest a reason for this other than errors in the procedure, measurements or calculations.
(1)
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(Total for question = 12 marks)

Unit 1 - Jun 2011
Q7.
0.400 g of magnesium ribbon reacted with exactly 22.2 cm3 of hydrochloric acid of concentration 1.50 mol dm−3.
400 cm3 of hydrogen gas was formed, the volume being measured at room temperature
 and pressure.
In the calculations that follow, use the following molar masses:
Mg = 24.0 g mol−1
Cl = 35.5 g mol−1
(a) Calculate the amount (in moles) of magnesium used.
(1)

(b) Calculate the amount (in moles) of hydrochloric acid used.
(1)

(c) Calculate the amount (in moles) of hydrogen produced.
      [Molar volume of any gas at room temperature and pressure = 24 000 cm3 mol−1]
(1)

(d) Show that the calculated amounts of magnesium, hydrochloric acid and hydrogen are consistent with the following equation for the reaction

(1)
(e) Calculate the maximum mass of magnesium chloride that would be formed in this reaction. Give your answer to three significant figures.
(3)
(Total for Question = 7 marks)
Q8.
The European Union has set a limit (with effect from January 2010) of 3.13 ppm for the proportion of the toxic gas carbon monoxide in the air that we breathe. This is equivalent to
   A    3.13%
   B    0.0313%
   C    0.000313%
   D    0.00000313%
(Total for Question = 1 mark)
Unit 1 - Jun 2012
Q9.
A solution contains 66 ppm of a solute. The mass of the solute dissolved in 1 kg of this solution is
   A   66 g
   B   0.66 g
   C   0.066 g
   D   0.000066 g
(Total for question = 1 mark)
Unit 1 - Jan 2013
Q10.
This question is about the preparation of the alum, potassium aluminium sulfate,
 KAl(SO4)212H2O.  It is a double salt consisting of potassium ions, aluminium ions and sulfate ions, and water of crystallization.
(a)  The first step of the preparation involves adding an excess of aluminium foil to 10 cm3 of 2 mol dm−3 potassium hydroxide to form potassium aluminate.
The equation for this reaction is

(i)   Write a balanced ionic equation for this reaction.
(1)

(ii)   Calculate the number of moles of potassium hydroxide used.
(1)

(iii)   Hence state the number of moles of aluminium that react with the potassium
         hydroxide.
(1)

(iv)   Use your answer to (iii) to calculate the mass of aluminium that reacts with the
         potassium hydroxide. Use the Periodic Table as a source of data.
(1)

(v)   Calculate the total mass of aluminium added to the potassium hydroxide if a
        10% excess of aluminium is required.
(1)

(vi)   Identify two hazards in this first step of the preparation.
(2)
Hazard 1
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Hazard 2
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(b)  The second step of the reaction is the addition of a slight excess of 1 mol dm−3 sulfuric acid.
(i)   Balance the following equation for the reaction
(1)

(ii)   Calculate the volume of the 1 mol dm−3 sulfuric acid that reacts with the potassium aluminate.
(1)

(iii)   State how you would show that the acid had been added in excess.
(2)
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*(iv)   State and explain the steps necessary to obtain pure, dry crystals from the mixture.
(4)
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(v)   Suggest the colour of the crystals.
(1)
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(vi)   Suggest the formula of another metal ion which could form an alum, in combination with potassium and sulfate ions.
(1)
(Total for question = 17 marks)

Unit 1 - Jun 2013
Q11.
How many moles of ions are present in 20 cm3 of 0.050 mol dm−3 calcium chloride solution, CaCl2(aq)?

   A    0.0050
   B    0.0030
   C    0.0020
   D    0.0010

(Total for Question = 1 mark)

Q12.
Phosphorus(V) chloride, PCl5, reacts with water according to the equation

PCl5(s) + 4H2O(l) → H3PO4(aq) + 5HCl(aq)
If 1.04 g of phosphorus pentachloride (molar mass = 208 g mol−1) is reacted completely with water and the solution made up to 1 dm3, the concentration of the hydrochloric acid in mol dm−3 is
   A    0.001
   B    0.005
   C    0.025
   D    0.250

(Total for Question = 1 mark)

Unit 1 - Jun 2014
Q13.
Nickel is an element in the d-block of the Periodic Table.
(a)  Complete the electronic configuration of a nickel atom using the s, p, d notation.
(1)
1s2 .................................................................................................................... 

(b)  A sample of nickel is made up of three isotopes. The percentage abundances are shown in the table below.

Calculate the relative atomic mass of nickel. Give your answer to two decimal places.
(2)

(c)  Nickel reacts with carbon monoxide, CO, to give the compound nickel carbonyl, Ni(CO)4.
Ni(s) + 4CO(g) → Ni(CO)4(g)
(i)  Calculate the volume of carbon monoxide, in dm3, measured at room temperature and pressure, that is required to react completely with 5.87 g of nickel.
[Relative atomic mass: Ni = 58.7 
Molar volume of a gas = 24 dm3 mol−1 at room temperature and pressure.]
(3)

(ii)  Calculate the number of carbon monoxide molecules present in the volume of gas you have calculated in (c)(i). 
[The Avogadro constant, L = 6.02 × 1023 mol−1]
(1)

(d)  Nickel(II) nitrate, Ni(NO3)2, can be made by several different methods.
Method 1
Nickel(II) oxide, NiO, was reacted with dilute nitric acid according to the equation

NiO(s) + 2HNO3(aq) → Ni(NO3)2(aq) + H2O(l)
(i)  Calculate the volume of 2.00 mol dm−3 dilute nitric acid, in cm3, that was required to exactly neutralize 1.494 g of nickel(II) oxide.
Use the relative atomic masses: Ni = 58.7, O = 16.0
(3)

Method 2
A volume of 25.0 cm3 of 2.00 mol dm−3 nitric acid, HNO3, was transferred to a beaker. Solid nickel(II) carbonate, NiCO3, was added until it was in excess.
(ii)  Why was excess nickel(II) carbonate used?
(1)
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(iii)  Why must the beaker be much larger than the volume of acid used?
(1)
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(iv)  Write a balanced equation for the reaction between nickel(II) carbonate and dilute nitric acid, including state symbols.
(2)

*(v)  For Method 2, describe the practical steps that you would take to obtain pure dry crystals of hydrated nickel(II) nitrate, Ni(NO3)2.6H2O, from a mixture of nickel(II) nitrate solution and unreacted solid nickel(II) carbonate.
(4)
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(Total for question = 18 marks)
Q14.
In an experiment to make crystals of hydrated copper(II) nitrate, a sample of 5.60 g of copper(II) oxide was added to 50 cm3 of 2.50 mol dm−3 nitric acid. The following reaction occurred.
CuO(s) + 2HNO3(aq) → Cu(NO3)2(aq) + H2O(l)
(a)  Calculate the number of moles of each reactant present, and use this to show that the copper(II) oxide was in excess.
The molar mass of copper(II) oxide, CuO, is 79.5 g mol−1
(3)
Moles of copper(II) oxide added

Moles of nitric acid used

The copper(II) oxide is in excess because
 ............................................................................................................................................. 
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 (b)  The copper(II) nitrate solution was heated gently to concentrate it, and then left to crystallize. The mass of hydrated copper(II) nitrate crystals, Cu(NO3)2.6H2O, obtained was 12.52 g.
Calculate the percentage yield.
The molar mass of Cu(NO3)2.6H2O is 295.6 g mol−1.
(3)

(c)  Give one reason why the percentage yield is less than 100%, even though the nitric acid was completely reacted.
(1)
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*(d)  (i)  The nitrate ion, NO3−, contains both covalent and dative covalent bonds.
What is the difference between these types of bond?
(2)
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(ii)  Complete the dot and cross diagram to show the bonding in the nitrate ion. 
Only the outer electron shells for each atom need to be shown.
Represent the nitrogen electrons with crosses (×), and oxygen electrons with dots, (•). The symbol * on the diagram represents the extra electron giving the ion its charge.
(3)

  
(Total for question = 12 marks)