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Double Science (Chemistry) Retrospective Revision guide

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Last updated about 1 year ago
130 questions
Note from the author:
The iGCSE Edexcel Double Award Science (Chemistry) syllabus for students to RAG rate to aid with revision
1 Principles of chemistry
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2 Inorganic chemistry
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3 Physical chemistry
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4 Organic chemistry
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(a) States of matter

Question 1
1.

1.1
understand the three states of matter in terms of the arrangement, movement and
energy of the particles

Question 2
2.

1.2
understand the interconversions between the three states of matter in terms of:
• the names of the interconversions
• how they are achieved
• the changes in arrangement, movement and energy of the particles.

Question 3
3.

1.3
understand how the results of experiments involving the dilution of coloured solutions
and diffusion of gases can be explained

Question 4
4.

1.4
know what is meant by the terms:
• solvent
• solute
• solution
• saturated solution.

(b) Elements, compounds and mixtures

Question 5
5.

1.8
understand how to classify a substance as an element, compound or mixture

Question 6
6.

1.9
understand that a pure substance has a fixed melting and boiling point, but that a
mixture may melt or boil over a range of temperatures

Question 7
7.

1.10
describe these experimental techniques for the separation of mixtures:
• simple distillation
• fractional distillation
• filtration
• crystallisation
• paper chromatography.

Question 8
8.

1.11
understand how a chromatogram provides information about the composition of a
mixture

Question 9
9.

1.12
understand how to use the calculation of Rf values to identify the components of a
mixture

Question 10
10.

1.13
practical: investigate paper chromatography using inks/food colourings

(c) Atomic structure

Question 11
11.

1.14
know what is meant by the terms atom and molecule

Question 12
12.

1.15
know the structure of an atom in terms of the positions, relative masses and relative
charges of sub-atomic particles

Question 13
13.

1.16
know what is meant by the terms atomic number, mass number, isotopes and relative
atomic mass (Ar)

Question 14
14.

1.17
be able to calculate the relative atomic mass of an element (Ar) from isotopic
abundances

(d) The Periodic Table

Question 15
15.

1.18
understand how elements are arranged in the Periodic Table:
• in order of atomic number
• in groups and periods.

Question 16
16.

1.19
understand how to deduce the electronic configurations of the first 20 elements from
their positions in the Periodic Table

Question 17
17.

1.20
understand how to use electrical conductivity and the acid-base character of oxides to
classify elements as metals or non-metals

Question 18
18.

1.21
identify an element as a metal or a non-metal according to its position in the Periodic
Table

Question 19
19.

1.22
understand how the electronic configuration of a main group element is related to its
position in the Periodic Table

Question 20
20.

1.23
understand why elements in the same group of the Periodic Table have similar
chemical properties

Question 21
21.

1.24
understand why the noble gases (Group 0) do not readily react

(e) Chemical formulae, equations and calculations

Question 22
22.

1.25
write word equations and balanced chemical equations (including state symbols):
• for reactions studied in this specification
• for unfamiliar reactions where suitable information is provided

Question 23
23.

1.26
calculate relative formula masses (including relative molecular masses) (Mr) from
relative atomic masses (Ar)

Question 24
24.

1.27
know that the mole (mol) is the unit for the amount of a substance

Question 25
25.

1.28
understand how to carry out calculations involving amount of substance, relative
atomic mass (Ar) and relative formula mass (Mr)

Question 26
26.

1.29
calculate reacting masses using experimental data and chemical equations

Question 27
27.

1.30
calculate percentage yield

Question 28
28.

1.31
understand how the formulae of simple compounds can be obtained experimentally,
including metal oxides, water and salts containing water of crystallisation

Question 29
29.

1.32
know what is meant by the terms empirical formula and molecular formula

Question 30
30.

1.33
calculate empirical and molecular formulae from experimental data

Question 31
31.

1.36
practical: know how to determine the formula of a metal oxide by combustion
(e.g. magnesium oxide) or by reduction (e.g. copper(II) oxide)

(f) Ionic bonding

Question 32
32.

1.37
understand how ions are formed by electron loss or gain

Question 33
33.

1.38 know the charges of these ions:
• metals in Groups 1, 2 and 3
• non-metals in Groups 5, 6 and 7
• Ag+, Cu2+, Fe2+, Fe3+, Pb2+, Zn2+
• hydrogen (H+), hydroxide (OH–), ammonium (NH4+), carbonate (CO32–), nitrate
(NO3-), sulfate (SO42–).

Question 34
34.

1.39
write formulae for compounds formed between the ions listed above

Question 35
35.

1.40
draw dot-and-cross diagrams to show the formation of ionic compounds by electron
transfer, limited to combinations of elements from Groups 1, 2, 3 and 5, 6, 7
only outer electrons need be shown

Question 36
36.

1.41
understand ionic bonding in terms of electrostatic attractions

Question 37
37.

1.42
understand why compounds with giant ionic lattices have high melting and boiling points

Question 38
38.

1.43
know that ionic compounds do not conduct electricity when solid, but do conduct
electricity when molten and in aqueous solution

(g) Covalent bonding

Question 39
39.

1.44
know that a covalent bond is formed between atoms by the sharing of a pair of
electrons

Question 40
40.

1.45
understand covalent bonds in terms of electrostatic attractions

Question 41
41.

1.46
understand how to use dot-and-cross diagrams to represent covalent bonds in:
• diatomic molecules, including hydrogen, oxygen, nitrogen, halogens and hydrogen
halides
• inorganic molecules including water, ammonia and carbon dioxide
• organic molecules containing up to two carbon atoms, including methane, ethane,
ethene and those containing halogen atoms.

Question 42
42.

1.47
explain why substances with a simple molecular structures are gases or liquids, or
solids with low melting and boiling points
the term intermolecular forces of attraction can be used to represent all forces
between molecules

Question 43
43.

1.48
explain why the melting and boiling points of substances with simple molecular
structures increase, in general, with increasing relative molecular mass

Question 44
44.

1.49
explain why substances with giant covalent structures are solids with high melting and
boiling points

Question 45
45.

1.50
explain how the structures of diamond, graphite and C60 fullerene influence their
physical properties, including electrical conductivity and hardness

Question 46
46.

1.51
know that covalent compounds do not usually conduct electricity

(a) Group 1 (alkali metals) – lithium, sodium and potassium

Question 47
47.

2.1
understand how the similarities in the reactions of these elements with water provide
evidence for their recognition as a family of elements

Question 48
48.

2.2
understand how the differences between the reactions of these elements with air and
water provide evidence for the trend in reactivity in Group 1

Question 49
49.

2.3
use knowledge of trends in Group 1 to predict the properties of other alkali metals

(b) Group 7 (halogens) – chlorine, bromine and iodine

Question 50
50.

2.5
know the colours, physical states (at room temperature) and trends in physical
properties of these elements

Question 51
51.

2.6
use knowledge of trends in Group 7 to predict the properties of other halogens

Question 52
52.

2.7
understand how displacement reactions involving halogens and halides provide
evidence for the trend in reactivity in Group 7

(c) Gases in the atmosphere

Question 53
53.

2.9
know the approximate percentages by volume of the four most abundant gases in dry
air

Question 54
54.

2.10
understand how to determine the percentage by volume of oxygen in air using
experiments involving the reactions of metals (e.g. iron) and non-metals
(e.g. phosphorus) with air

Question 55
55.

2.11
describe the combustion of elements in oxygen, including magnesium, hydrogen and
sulfur

Question 56
56.

2.12
describe the formation of carbon dioxide from the thermal decomposition of metal
carbonates, including copper(II) carbonate

Question 57
57.

2.13
know that carbon dioxide is a greenhouse gas and that increasing amounts in the
atmosphere may contribute to climate change

Question 58
58.

2.14
practical: determine the approximate percentage by volume of oxygen in air using a
metal or a non-metal

(d) Reactivity series

Question 59
59.

2.15
understand how metals can be arranged in a reactivity series based on their reactions
with:
• water
• dilute hydrochloric or sulfuric acid.

Question 60
60.

2.16
understand how metals can be arranged in a reactivity series based on their
displacement reactions between:
• metals and metal oxides
• metals and aqueous solutions of metal salts.

Question 61
61.

2.17
know the order of reactivity of these metals: potassium, sodium, lithium, calcium,
magnesium, aluminium, zinc, iron, copper, silver, gold

Question 62
62.

2.18
know the conditions under which iron rusts

Question 63
63.

2.19
understand how the rusting of iron may be prevented by:
• barrier methods
• galvanising
• sacrificial protection.

Question 64
64.

2.20
the terms:
• oxidation
• reduction
• redox
• oxidising agent
• reducing agent
in terms of gain or loss of oxygen and loss or gain of electrons.

Question 65
65.

2.21
practical: investigate reactions between dilute hydrochloric and sulfuric acids and
metals (e.g. magnesium, zinc and iron)

(e) Acids, alkalis and titrations

Question 66
66.

2.28
describe the use of litmus, phenolphthalein and methyl orange to distinguish between
acidic and alkaline solutions

Question 67
67.

2.29
understand how to use the pH scale, from 0–14, can be used to classify solutions as
strongly acidic (0–3), weakly acidic (4–6), neutral (7), weakly alkaline (8–10) and
strongly alkaline (11–14)

Question 68
68.

2.30
describe the use of universal indicator to measure the approximate pH value of an
aqueous solution

Question 69
69.

2.31
know that acids in aqueous solution are a source of hydrogen ions and alkalis in a
aqueous solution are a source of hydroxide ions

Question 70
70.

2.32
know that alkalis can neutralise acids

f) Acids, bases and salt preparations

Question 71
71.

2.34
know the general rules for predicting the solubility of ionic compounds in water:
• common sodium, potassium and ammonium compounds are soluble
• all nitrates are soluble
• common chlorides are soluble, except those of silver and lead(II)
• common sulfates are soluble, except for those of barium, calcium and lead(II)
• common carbonates are insoluble, except for those of sodium, potassium and
ammonium
• common hydroxides are insoluble except for those of sodium, potassium and
calcium (calcium hydroxide is slightly soluble).

Question 72
72.

2.35
understand acids and bases in terms of proton transfer

Question 73
73.

2.36
understand that an acid is a proton donor and a base is a proton acceptor

Question 74
74.

2.37
describe the reactions of hydrochloric acid, sulfuric acid and nitric acid with metals,
bases and metal carbonates (excluding the reactions between nitric acid and metals)
to form salts

Question 75
75.

2.38
know that metal oxides, metal hydroxides and ammonia can act as bases, and that
alkalis are bases that are soluble in water

Question 76
76.

2.39
describe an experiment to prepare a pure, dry sample of a soluble salt, starting from
an insoluble reactant

Question 77
77.

2.42
practical: prepare a sample of pure, dry hydrated copper(II) sulfate crystals starting
from copper(II) oxide

(g) Chemical tests

Question 78
78.

2.44
describe tests for these gases:
• hydrogen
• oxygen
• carbon dioxide
• ammonia
• chlorine.

Question 79
79.

2.45
describe how to carry out a flame test

Question 80
80.

2.46
know the colours formed in flame tests for these cations:
• Li+ is red
• Na+ is yellow
• K+ is lilac
• Ca2+ is orange-red
• Cu2+ is blue-green.

Question 81
81.

2.47
describe tests for these cations:
• NH4+ using sodium hydroxide solution and identifying the gas evolved
• Cu2+, Fe2+ and Fe3+ using sodium hydroxide solution.

Question 82
82.

2.48
describe tests for these anions:
• Cl–, Br– and I– using acidified silver nitrate solution
• SO42– using acidified barium chloride solution
• CO32– using hydrochloric acid and identifying the gas evolved.

Question 83
83.

2.49
describe a test for the presence of water using anhydrous copper(II) sulfate

Question 84
84.

2.50
describe a physical test to show whether a sample of water is pure

(a) Energetics

Question 85
85.

3.1
know that chemical reactions in which heat energy is given out are described as
exothermic, and those in which heat energy is taken in are described as endothermic

Question 86
86.

3.2
describe simple calorimetry experiments for reactions such as combustion,
displacement, dissolving and neutralisation

Question 87
87.

3.3
calculate the heat energy change from a measured temperature change using the
expression Q = mcΔT

Question 88
88.

3.4
calculate the molar enthalpy change (ΔH) from the heat energy change, Q

Question 89
89.

3.8
practical: investigate temperature changes accompanying some of the following types
of change:
• salts dissolving in water
• neutralisation reactions
• displacement reactions
• combustion reactions.

(b) Rates of reaction

Question 90
90.

3.9
describe experiments to investigate the effects of changes in surface area of a solid,
concentration of a solution, temperature and the use of a catalyst on the rate of a
reaction

Question 91
91.

3.10
describe the effects of changes in surface area of a solid, concentration of a solution,
pressure of a gas, temperature and the use of a catalyst on the rate of a reaction

Question 92
92.

3.11
explain the effects of changes in surface area of a solid, concentration of a solution,
pressure of a gas and temperature on the rate of a reaction in terms of particle
collision theory

Question 93
93.

3.12
know that a catalyst is a substance that increases the rate of a reaction, but is
chemically unchanged at the end of the reaction

Question 94
94.

3.13
know that a catalyst works by providing an alternative pathway with lower activation
energy

Question 95
95.

3.15
practical: investigate the effect of changing the surface area of marble chips and of
changing the concentration of hydrochloric acid on the rate of reaction between
marble chips and dilute hydrochloric acid

Question 96
96.

3.16
practical: investigate the effect of different solids on the catalytic decomposition of
hydrogen peroxide solution

(c) Reversible reactions and equilibria

Question 97
97.

3.17
know that some reactions are reversible and this is indicated by the symbol ⇌ in
equations

Question 98
98.

3.18
describe reversible reactions such as the dehydration of hydrated copper(II) sulfate
and the effect of heat on ammonium chloride

(a) Introduction

Question 99
99.

4.1
know that a hydrocarbon is a compound of hydrogen and carbon only

Question 100
100.

4.2
understand how to represent organic molecules using empirical formulae, molecular
formulae, general formulae, structural formulae and displayed formulae

Question 101
101.

4.3
know what is meant by the terms homologous series, functional group and isomerism

Question 102
102.

4.4
understand how to name compounds relevant to this specification using the rules of
International Union of Pure and Applied Chemistry (IUPAC) nomenclature
students will be expected to name compounds containing up to six carbon atoms

Question 103
103.

4.5
understand how to write the possible structural and displayed formulae of an organic
molecule given its molecular formula

Question 104
104.

4.6
understand how to classify reactions of organic compounds as substitution, addition
and combustion
knowledge of reaction mechanisms is not required

(b) Crude oil

Question 105
105.

4.7
know that crude oil is a mixture of hydrocarbons

Question 106
106.

4.8
describe how the industrial process of fractional distillation separates crude oil into
fractions

Question 107
107.

4.9
know the names and uses of the main fractions obtained from crude oil:
refinery gases, gasoline, kerosene, diesel, fuel oil and bitumen

Question 108
108.

4.10
know the trend in colour, boiling point and viscosity of the main fractions

Question 109
109.

4.11
know that a fuel is a substance that, when burned, releases heat energy

Question 110
110.

4.12
know the possible products of complete and incomplete combustion of hydrocarbons
with oxygen in the air

Question 111
111.

4.13
understand why carbon monoxide is poisonous, in terms of its effect on the capacity
of blood to transport oxygen
references to haemoglobin are not required

Question 112
112.

4.14
know that, in car engines, the temperature reached is high enough to allow nitrogen
and oxygen from air to react, forming oxides of nitrogen

Question 113
113.

4.15
explain how the combustion of some impurities in hydrocarbon fuels results in the
formation of sulfur dioxide

Question 114
114.

4.16
understand how sulfur dioxide and oxides of nitrogen contribute to acid rain

Question 115
115.

4.17
describe how long-chain alkanes are converted to alkenes and shorter-chain alkanes
by catalytic cracking (using silica or alumina as the catalyst and a temperature in the
range of 600–700 ºC)

Question 116
116.

4.18
explain why cracking is necessary, in terms of the balance between supply and
demand for different fractions

(c) Alkanes

Question 117
117.

4.19
know the general formula for alkanes

Question 118
118.

4.20
explain why alkanes are classified as saturated hydrocarbons

Question 119
119.

4.21
understand how to draw the structural and displayed formulae for alkanes with up to
five carbon atoms in the molecule, and to name the unbranched-chain isomers

Question 120
120.

4.22
describe the reactions of alkanes with halogens in the presence of ultraviolet
radiation, limited to mono-substitution
knowledge of reaction mechanisms is not required

(d) Alkenes

Question 121
121.

4.23
know that alkenes contain the functional group >C=C<

Question 122
122.

4.24
know the general formula for alkenes

Question 123
123.

4.25
explain why alkenes are classified as unsaturated hydrocarbons

Question 124
124.

4.26
understand how to draw the structural and displayed formulae for alkenes with up to
four carbon atoms in the molecule, and name the unbranched-chain isomers
knowledge of cis/trans or E/Z notation is not required

Question 125
125.

4.27
describe the reactions of alkenes with bromine to produce dibromoalkanes

Question 126
126.

4.28
describe how bromine water can be used to distinguish between an alkane and an
alkene

(e) Synthetic polymers

Question 127
127.

4.44
know that an addition polymer is formed by joining up many small molecules called
monomers

Question 128
128.

4.45
understand how to draw the repeat unit of an addition polymer, including
poly(ethene), poly(propene), poly(chloroethene) and (poly)tetrafluoroethene

Question 129
129.

4.46
understand how to deduce the structure of a monomer from the repeat unit of an
addition polymer and vice versa

Question 130
130.

4.47
explain problems in the disposal of addition polymers, including:
• their inertness and inability to biodegrade
• the production of toxic gases when they are burned.