1
1
When potassium is added to water, it reacts vigorously and a coloured flame is seen.
The equation for the reaction is shown.
2K(s) + 2H2O(l) → 2KOH(aq) + H2(g)
3
(a) Calculate the volume, in cm , of hydrogen gas formed when 2.34 g of potassium is added to
excess water at room temperature and pressure.
Use the following steps.
(i)
Calculate the number of moles of potassium added.
= .......................................................mol [1]
(ii) Determine the number of moles of hydrogen gas formed.
= .......................................................mol [1]
(iii) Calculate the volume of hydrogen gas formed.
3
volume = ..........................................cm
[1]
[Total: 3]
2
Calcium nitrate crystals are hydrated and have the formula Ca(NO3)2•xH2O where x is a whole
number of molecules of water.
The student heats the crystals to remove the molecules of water.
2
Ca(NO3)2•xH2O(s) → Ca(NO3)2(s) + xH2O(g)
The student heats a sample of Ca(NO3)2•xH2O and forms 2.46 g of Ca(NO3)2 and 0.0600 moles
of H2O.
Determine the value of x. Use the following steps.
(a) Calculate the Mr of Ca(NO3)2.
Mr = ................................................ [1]
(b) Determine the number of moles of Ca(NO3)2 formed.
moles of Ca(NO3)2 formed = ........................ [1]
(c) Determine the value of x in Ca(NO3)2•xH2O.
x = .................................................. [1]
[Total: 3]
3
3
Calcium carbonate decomposes when heated.
calcium carbonate → calcium oxide + carbon dioxide
3
When 0.50 g of calcium carbonate decomposes, 120 cm of carbon dioxide gas is produced.
Calculate the volume of carbon dioxide gas produced when 0.10 g of calcium carbonate is used.
3
volume of carbon dioxide gas = .............................cm
[1]
[Total: 1]
4
A student investigates the reaction of small pieces of calcium carbonate with dilute hydrochloric
acid.
The hydrochloric acid is in excess.
CaCO3 + 2HCl → CaCl2 + CO2 + H2O
3
When 0.11 g of calcium carbonate is used, 25 cm of carbon dioxide gas is produced.
3
Calculate the mass of calcium carbonate needed to produce 100 cm of carbon dioxide gas.
mass of calcium carbonate = ......................g [1]
[Total: 1]
4
5
A student investigates the reaction of large pieces of copper(II) carbonate with dilute hydrochloric
acid. The hydrochloric acid is in excess.
CuCO3 + 2HCl → CuCl2 + CO2 + H2O
3
When 0.2 g of copper(II) carbonate is used, 38 cm of carbon dioxide gas is produced.
Calculate the volume of carbon dioxide gas produced when 0.50 g of copper(II) carbonate is used.
3
volume of carbon dioxide gas = ........................................ cm
[1]
[Total: 1]
6
Calcium carbonate decomposes when heated.
calcium carbonate → calcium oxide + carbon dioxide
3
When 0.50 g of calcium carbonate decomposes, 120 cm of carbon dioxide gas is produced.
Calculate the volume of carbon dioxide gas produced when 0.10 g of calcium carbonate is used.
3
volume of carbon dioxide gas = ................cm
[1]
[Total: 1]
7
When 4.8 g of magnesium reacts with excess oxygen, 8.0 g of magnesium oxide is formed.
Calculate the minimum mass of magnesium needed to produce 24.0 g of magnesium oxide.
minimum mass = ............................... g [1]
[Total: 1]
5
8
A student investigated the reaction of small pieces of zinc with dilute hydrochloric acid.
The hydrochloric acid was in excess.
Zn + 2HCl → ZnCl2 + H2
3
When 0.065 g of zinc is used, 24 cm of hydrogen gas is formed.
3
Calculate the mass of zinc needed to produce 96 cm of hydrogen gas.
mass of zinc = ................................g [1]
[Total: 1]
9
Lead is a metallic element in Group IV. One of the ores of lead is galena, which is an impure form
of lead(II) sulfide, PbS.
Lead also occurs in the ore cerussite, which contains lead(II) carbonate, PbCO3.
(a) Calculate the relative formula mass, Mr, of PbCO3.
Mr of PbCO3 = .............................................. [1]
6
(b) The Mr of PbS is 239.
Calculate the percentage of lead by mass in PbS.
percentage of lead by mass in PbS = .......... [1]
[Total: 2]
7
10 In an experiment, 1.61 g of Na2SO4•xH2O is heated until all the water is given off. The mass of
Na2SO4 remaining is 0.71 g.
[Mr: Na2SO4,142; H2O,18]
Determine the value of x using the following steps.
•
Calculate the number of moles of Na2SO4 remaining.
............................ mol
•
Calculate the mass of H2O given off.
................................ g
•
Calculate the number of moles of H2O given off.
............................ mol
•
Determine the value of x
x = ............................. [4]
[Total: 4]
8
11
Sodium hydrogencarbonate is found in baking powder.
When sodium hydrogencarbonate is heated it forms three products.
2NaHCO3 → Na2CO3 + H2O + CO2
Calculate the volume of carbon dioxide formed at room temperature and pressure when 12.6 g of
NaHCO3 is heated using the following steps:
•
determine the mass of one mole of NaHCO3
....................................... g
•
calculate the number of moles of NaHCO3
............................... moles
•
determine the number of moles of carbon dioxide formed
............................... moles
•
calculate the volume of carbon dioxide formed at room temperature and pressure.
................................... dm
3
[4]
[Total: 4]
9
12 On analysing hydrated iron(II) sulfate crystals, FeSO4•xH2O, a student found that one mole of the
crystals had a mass of 278 g.
Determine the value of x using the following steps:
•
calculate the mass of one mole of FeSO4
mass = ................................ g
•
calculate the mass of H2O present in one mole of FeSO4•xH2O
mass of H2O = ............................ g
•
determine the value of x.
x = .................................... [3]
[Total: 3]
13 Ammonia nitrate, NH4NO3, is a fertiliser.
Calculate the percentage by mass of nitrogen in ammonium nitrate.
...................................... % [2]
[Total: 2]
14 A sample of concentrated hydrobromic acid, HBr(aq), was electrolysed using platinum electrodes.
3
The concentration of the hydrobromic acid was 8.89 mol / dm .
10
3
(a) Calculate the concentration of the HBr(aq) in g / dm .
3
concentration of the HBr(aq) = ........................... g / dm
[1]
(b) Explain why concentrated HBr(aq) can conduct electricity.
...........................................................................................................................................
...........................................................................................................................................
...........................................................................................................................................
[2]
(c) Magnesium is not a suitable material from which to make the electrodes.
Explain why.
...........................................................................................................................................
...........................................................................................................................................
[1]
(d) Predict the product formed at the anode when concentrated HBr(aq) is electrolysed.
...........................................................................................................................................
[1]
(e) Write the ionic half-equation for the reaction occurring at the cathode.
...........................................................................................................................................
[2]
[Total: 7]
15 Nitrates such as ammonium nitrate are used as fertilisers.
The final stage in the production of ammonium nitrate is shown in the equation.
Ca(NO3)2 + 2NH3 + CO2 + H2O → 2NH4NO3 + CaCO3
Calculate the maximum mass of ammonium nitrate that can be produced from 820 g of calcium
nitrate, Ca(NO3)2, using the following steps.
The relative formula mass, Mr, of calcium nitrate, Ca(NO3)2, = 164.
11
(a) Calculate the number of moles of Ca(NO3)2 in 820 g.
........................... mol
[1]
........................... mol
[1]
Mr of NH4NO3 = ...........................
[1]
(b) Deduce the number of moles of NH4NO3 produced.
(c) Calculate the Mr of NH4NO3.
(d) Calculate the maximum mass of ammonium nitrate produced.
........................... g
[1]
[Total: 4]
16 Tin is a metallic element in Group IV. Its main ore is cassiterite which is an impure form of tin(IV)
oxide, SnO2.
Tin also occurs in stannite, Cu2FeSnS4.
12
(a) Calculate the relative formula mass, Mr, of Cu2FeSnS4.
Mr of Cu2FeSnS4 .....................................
[1]
(b) The Mr of SnO2 is 151.
Calculate the percentage of tin by mass in SnO2.
percentage of tin by mass in SnO2 .....................................
[1]
(c) The percentage of tin by mass in Cu2FeSnS4 is 27.6 %.
Use this information and your answer to (b) to suggest whether it would be better to extract
tin from SnO2 or Cu2FeSnS4.
Explain your answer.
...........................................................................................................................................
...........................................................................................................................................
[1]
[Total: 3]
3
3
17 25 cm of a gaseous hydrocarbon, CxHy, were burnt in 150 cm of oxygen. This was an excess of
oxygen.
3
3
After cooling, the volume of the gases remaining was 100 cm . This consisted of 75 cm of
3
carbon dioxide and 25 cm of unreacted oxygen. The water that was produced in the reaction was
liquid.
All volumes were measured at the same temperature and pressure.
(a) What is meant by an excess of oxygen?
...........................................................................................................................................
[1]
13
(b) What was the volume of oxygen that reacted with the hydrocarbon?
3
........................... cm
[1]
(c) Complete the table to show the smallest whole number ratio of volumes.
volume of
hydrocarbon
reacted
smallest whole
number ratio of
volumes
:
volume of
oxygen
reacted
:
:
volume of
carbon
dioxide
produced
:
[1]
(d) Use your answer to (c) to balance the chemical equation. Deduce the formula of the
hydrocarbon.
CxHy(g) + ........O2(g) → ........CO2(g) + ........H2O(l)
formula of the hydrocarbon = .....................................
[2]
[Total: 5]
18 A student prepares magnesium sulfate crystals, MgSO4, by adding excess magnesium to dilute
sulfuric acid.
Magnesium sulfate crystals have the formula, MgSO4•xH2O, where x is a whole number of molecules
of water.
The student heats the crystals to remove the molecules of water.
MgSO4•xH2O(s) → MgSO4(s) + xH2O(g)
The student heats a sample of MgSO4•xH2O and finds it has lost 0.140 moles of H2O and has
2.40 g of MgSO4 remaining.
14
Determine the value of x .
Use the following steps.
(a) Calculate the Mr of MgSO4.
Mr = .................................. [1]
(b) Determine the number of moles of MgSO4 formed.
moles of MgSO4 formed = ........................... [1]
(c) Determine the value of x in MgSO4•xH2O.
x = .................................... [1]
[Total: 3]
19 Sodium reacts vigorously with water to form aqueous sodium hydroxide, NaOH, which is a strong
base.
The equation for the reaction is shown.
2Na(s) + 2H2O(l) → 2NaOH(aq) + H2(g)
3
Calculate the concentration of NaOH(aq) formed, in g / dm , when 0.345 g of sodium is added to
3
50.0 cm of distilled water. Assume there is no change in volume.
Use the following steps.
15
(a) Calculate the number of moles of Na added.
= ............................... mol [1]
(b) Determine the number of moles of NaOH formed.
= ............................... mol [1]
3
(c) Calculate the concentration of NaOH in mol / dm .
3
[1]
3
[2]
concentration of NaOH = .......................... mol / dm
3
(d) Determine the Mr of NaOH and calculate the concentration of NaOH in g / dm .
concentration of NaOH = .............................. g / dm
[Total: 5]
20 Ester Z has the empirical formula C3H6O and a relative molecular mass of 116.
Calculate the molecular formula of ester Z.
molecular formula = .......................... [1]
16
[Total: 1]
21 A student determines the concentration of a solution of dilute sulfuric acid, H2SO4, by titration with
aqueous sodium hydroxide, NaOH.
The equation for the reaction is shown.
H2SO4 + 2NaOH → Na2SO4 + 2H2O
Calculate the concentration of H2SO4 using the following steps.
3
3
• Calculate the number of moles in 25.0 cm of 0.200 mol / dm NaOH.
................................... mol
• Determine the number of moles of H2SO4 that react with the NaOH.
................................... mol
• Calculate the concentration of H2SO4.
3
.......................... mol / dm
[3]
[Total: 3]
3
22 A 25.0 cm sample of limewater is placed in a conical flask.
The concentration of Ca(OH)2 in the limewater is determined by titration with dilute hydrochloric
acid, HCl.
(a) Name the item of apparatus used to measure the volume of acid in this titration.
...........................................................................................................................................
[1]
17
(b) State the type of reaction which takes place.
...........................................................................................................................................
[1]
(c) As well as limewater and dilute hydrochloric acid, state what other type of substance must be
added to the conical flask.
...........................................................................................................................................
[1]
(d) The equation for the reaction is shown.
Ca(OH)2 + 2HCl → CaCl2 + 2H2O
3
3
3
20.0 cm of 0.0500 mol / dm HCll reacts with the 25.0 cm of Ca(OH)2.
3
Determine the concentration of Ca(OH)2 in g / dm . Use the following steps.
(i)
3
3
Calculate the number of moles in 20.0 cm of 0.0500 mol / dm HCl.
.............................................mol [1]
3
(ii) Determine the number of moles of Ca(OH)2 in 25.0 cm of the limewater.
.............................................mol [1]
3
(iii) Calculate the concentration of Ca(OH)2 in mol / dm .
3
....................................mol / dm
[1]
18
3
(iv) Determine the concentration of Ca(OH)2 in g / dm .
3
........................................g / dm
[2]
[Total: 8]
3
23 Determine the volume of CO2 gas given off when excess MgCO3 is added to 25.0 cm of
3
0.400 mol / dm HCl at room temperature and pressure.
MgCO3 + 2HCl → MgCl2 + H2O + CO2
Use the following steps.
•
3
3
Calculate the number of moles of HCl in 25.0 cm of 0.400 mol / dm of acid.
................................... mol
•
Determine the number of moles of CO2 gas given off.
................................... mol
•
3
Calculate the volume of CO2 gas given off in cm .
................................... cm
3
[3]
[Total: 3]
24 When solid copper(II) nitrate is heated copper(II) oxide, nitrogen dioxide and oxygen are formed.
2Cu(NO3)2 → 2CuO + 4NO2 + O2
19
Calculate the volume of nitrogen dioxide formed at room temperature and pressure when 4.7 g of
Cu(NO3)2 is heated.
Use the following steps:
(a) Calculate the mass of one mole of Cu(NO3)2.
............................... g [1]
(b) Calculate the number of moles of Cu(NO3)2 used.
........................ moles [1]
(c) Determine the number of moles of nitrogen dioxide formed.
........................ moles [1]
(d) Calculate the volume of nitrogen dioxide formed at room temperature and pressure.
3
........................... dm
[1]
[Total: 4]
25 The Avogadro constant is the number of particles in 1 mole.
23
The numerical value of the Avogadro constant is 6.02 x 10 .
20
(a) Calculate the number of molecules in 22.0 g of carbon dioxide, CO2.
.......................................molecules [2]
3
(b) Calculate the number of molecules in 6.00 dm of carbon dioxide gas at room temperature
and pressure.
.......................................molecules [1]
[Total: 3]
26 The equation for the reaction between powdered zinc carbonate and dilute nitric acid is shown.
(a) ZnCO3............ + 2HNO3............ → Zn(NO3)2............ + H2O............ + CO2............
Complete the equation by adding state symbols.
[2]
3
(b) A student found that 2.5 g of zinc carbonate required 20 cm of dilute nitric acid to react
completely.
Calculate the concentration of dilute nitric acid using the following steps:
(i)
Calculate the mass of 1 mole of ZnCO3.
.......................g [1]
(ii) Calculate the number of moles of ZnCO3 reacting.
...............moles [1]
21
(iii) Determine the number of moles of HNO3 reacting.
...............moles [1]
(iv) Calculate the concentration of HNO3.
3
..........mol / dm
[1]
[Total: 6]
22
27 Some chlorides are hydrated.
When hydrated barium chloride crystals, BaCl2•xH2O, are heated they give off water.
BaCl2•xH2O(s) → BaCl2(s) + xH2O(g)
A student carries out an experiment to determine the value of x in BaCl2•xH2O.
In an experiment, 4.88 g of BaCl2•xH2O is heated until all the water is given off. The mass of BaCl2
remaining is 4.16 g.
[Mr: BaCl2, 208; H2O, 18]
Determine the value of x using the following steps.
•
Calculate the number of moles of BaCl2 remaining.
................................... mol
•
Calculate the mass of H2O given off.
....................................... g
•
Calculate the number of moles of H2O given off.
................................... mol
•
Determine the value of x.
x = .................................... [4]
[Total: 4]
28 Ammonium nitrate, NH4NO3, is a compound of nitrogen.
Calculate the percentage by mass of nitrogen in ammonium nitrate.
percentage by mass of nitrogen = ................................ [2]
[Total: 2]
23
3
29 A sample of sulfuric acid, H2SO4, has a concentration of 0.75 mol / dm .
3
Calculate the concentration of sulfuric acid in g / dm .
3
.............................. g / dm
[2]
[Total: 2]
30 When hydrated magnesium sulfate crystals, MgSO4•xH2O, are heated they give off water.
MgSO4•xH2O(s) → MgSO4(s) + xH2O(g)
A student carries out an experiment to determine the value of x in MgSO4•xH2O.
Step 1 Hydrated magnesium sulfate crystals were weighed.
Step 2 Hydrated magnesium sulfate crystals were heated.
Step 3 The remaining solid was weighed.
(a) Describe how the student can ensure that all the water is given off.
...........................................................................................................................................
...........................................................................................................................................
...........................................................................................................................................
[2]
24
(b) In an experiment, all the water was removed from 1.23 g of MgSO4•xH2O. The mass of MgSO4
remaining was 0.60 g.
Mr: MgSO4 = 120; Mr: H2O = 18
Determine the value of x using the following steps.
•
Calculate the number of moles of MgSO4 remaining.
moles of MgSO4 = ........................
•
Caluculate the mass of H2O given off.
mass of H2O = ......................... g
•
Calculate the moles of H2O given off.
moles of H2O = .............................
•
Determine the value of x.
x = .................................. [4]
[Total: 6]
31 A student wanted to find the concentration of some dilute sulfuric acid by titration. The student
3
3
3
found that 25.0 cm of 0.0400 mol / dm NaOH(aq) reacted exactly with 20.0 cm of H2SO4(aq).
(a) Name a suitable indicator to use in this titration.
...........................................................................................................................................
[1]
25
3
(b) Calculate the concentration of the H2SO4(aq) in mol / dm using the following steps.
3
• Calculate the number of moles of NaOH in 25.0 cm .
moles = ..................................
3
• Deduce the number of moles of H2SO4 that reacted with the 25.0 cm of NaOH(aq).
moles = ..................................
3
• Calculate the concentration of H2SO4(aq) in mol / dm .
3
[3]
3
[2]
concentration = ......................... mol / dm
3
3
(c) Calculate the concentration of the 0.0400 mol / dm NaOH(aq) in g / dm .
concentration = ............................ g / dm
[Total: 6]
32 The equation for the decomposition of hydrogen peroxide is shown.
2H2O2(aq) → 2H2O(I) + O2(g)
3
3
25.0 cm of aqueous hydrogen peroxide forms 48.0 cm of oxygen at room temperature and pressure
(r.t.p.).
26
(a) Calculate the concentration of aqueous hydrogen peroxide at the start of the experiment using
the following steps.
• Calculate the number of moles of oxygen formed.
........................................ mol
• Deduce the number of moles of hydrogen peroxide that decomposed.
........................................ mol
3
• Calculate the concentration of hydrogen peroxide in mol / dm .
3
[3]
...........................................................................................................................................
[2]
.................................mol / dm
(b) Oxygen can also be produced by the decomposition of potassium chlorate(V), KClO3.
The only products of this decomposition are potassium chloride and oxygen.
Write a chemical equation for this decomposition.
[Total: 5]
27
33 PCl5 reacts with lithium fluoride, LiF, to form LiPF6.
PCl5 + 6LiF → LiPF6 + 5LiCl
Calculate the mass of LiF needed to form 3.04 g of LiPF6 using the following steps.
• Calculate the number of moles of LiPF6 formed.
[Mr: LiPF6, 152]
number of moles = .........................................
• Deduce the number of moles of LiF needed.
number of moles = .........................................
• Calculate the mass of LiF needed.
mass = ................................ g [3]
[Total: 3]
34 Dilute hydrochloric acid, HCl(aq), reacts with aqueous sodium carbonate, Na2CO3(aq).
The chemical equation for the reaction is shown.
2HCl + Na2CO3 → 2NaCl + CO2 + H2O
3
A 25.0 cm portion of Na2CO3(aq) was placed in a conical flask with a few drops of a suitable
3
indicator. It was titrated against HCl(aq) of concentration 0.180 mol / dm .
3
20.0 cm of HCl(aq) was required to reach the end-point.
3
Calculate the concentration of the Na2CO3(aq), in mol / dm , using the following steps.
28
(a) Calculate the number of moles of HCl used in the titration.
.....................................mol
[1]
3
(b) Calculate the number of moles of Na2CO3 contained in the 25.0 cm portion of Na2CO3(aq).
.....................................mol
[1]
.....................................
[1]
3
(c) Calculate the concentration of the Na2CO3(aq) in mol / dm .
[Total: 3]
29
35 Ammonia reacts with oxygen as shown.
4NH3(g) + 5O2(g) → 4NO(g) + 6H2O(g)
3
3
Calculate the volume of oxygen at room temperature and pressure, in dm , that reacts with 4.80 dm
of ammonia.
3
volume = .....................................dm
[3]
[Total: 3]
36 Some car airbags contain sodium azide.
When a car airbag is used the sodium azide, NaN3, decomposes.
The products are nitrogen and sodium.
The equation for the decomposition of sodium azide is shown.
2NaN3(s) → 2Na(l) + 3N2(g)
3
Calculate the mass, in g, of sodium azide needed to produce 144 dm of nitrogen using the following
steps.
3
(a) Calculate the number of moles in 144 dm of N2 measured at room temperature and pressure.
moles of N2 = .....................................mol
[1]
(b) Determine the number of moles of NaN3 needed to produce this number of moles of N2.
moles of NaN3 = .....................................mol
[1]
30
(c) Calculate the relative formula mass, Mr, of NaN3.
Mr = .....................................
[1]
3
(d) Calculate the mass of NaN3 needed to produce 144 dm of N2.
.....................................g
[1]
[Total: 4]
37 Titanium is extracted from an ore called rutile. Rutile is an impure form of titanium(IV) oxide, TiO2.
Calculate the volume of chlorine gas, Cl2(g), at room temperature and pressure, that reacts
completely with 400 g of TiO2(s) using the following steps.
TiO2(s) + 2Cl2(g) + 2C(s) → TiCl4(g) + 2CO(g)
(a) Calculate the relative formula mass, Mr, of TiO2.
Mr of TiO2 = ...........................
[1]
........................... mol
[1]
(b) Calculate the number of moles in 400 g of TiO2.
31
(c) Determine the number of moles of Cl2 that react with 400 g of TiO2.
moles of Cl2 = ........................... mol
[1]
(d) Calculate the volume of Cl2 that reacts with 400 g of TiO2.
3
volume of Cl2 = ........................... dm
[1]
[Total: 4]
38 Iron(II) sulfate decomposes when heated strongly.
2FeSO4(s) → Fe2O3(s) + SO2(g) + SO3(g)
15.20 g of FeSO4(s) was heated and formed 4.80 g of Fe2O3(s).
[Mr, FeSO4 = 152; Mr, Fe2O3 = 160]
Calculate the percentage yield for this reaction.
........................... %
[3]
[Total: 3]
39 A teacher heated 18.8 g of copper(II) nitrate.
The equation for the reaction is shown.
2Cu(NO3)2 → O2 + 4NO2 + 2CuO
32
(a) Calculate the number of moles of copper(II) nitrate present in the 18.8 g.
.....................................mol
[2]
(b) Calculate the maximum number of moles of oxygen that can be made by heating 18.8 g of
copper(II) nitrate.
.....................................mol
[1]
3
(c) Calculate the maximum volume of oxygen at room temperature and pressure, in cm , that can
be made by heating 18.8 g of copper(II) nitrate.
3
.....................................cm
[1]
[Total: 4]
40 Calcium chlorate(V), Ca(ClO3)2, is made by reacting calcium hydroxide with chlorine gas.
6Ca(OH)2 + 6Cl2 → Ca(ClO3)2 + 5CaCl2 + 6H2O
3
8.88 g of calcium hydroxide and 7200 cm of chlorine gas are mixed together.
(a) How many moles is 8.88 g of calcium hydroxide?
..................................... mol
[2]
..................................... mol
[1]
3
(b) How many moles of chlorine gas is 7200 cm ?
33
(c) What is the maximum number of moles of calcium chlorate(V) that can be made from 8.88 g
3
of calcium hydroxide and 7200 cm of chlorine gas?
..................................... mol
[1]
(d) What is the maximum mass of calcium chlorate(V) that can be made from 8.88 g of
3
calcium hydroxide and 7200 cm of chlorine gas?
..................................... g
[2]
[Total: 6]