chemistry igcse: higher tier revision syllabus
TRANSCRIPT
Westminster School Chemistry Department
Chemistry IGCSE: Revision Notes Summer 2015
This booklet should be used by Westminster Students taking their IGCSE in
Chemistry in Summer 2015. Please do not use any older revision guides.
This booklet contains all the topics you need to know for your IGCSE Chemistry
paper, but not necessarily in sufficient detail. It must be used in conjunction with
your other resources (e.g. textbook, notebook, revision guide)
Exam Format
You will sit two Chemistry papers this summer.
Paper 1 – Two hours, tests most of the syllabus content, has a 2/3 weighting.
Paper 2 – One hour, tests all of the syllabus content, has a 1/3 weighting.
Both Papers will have questions testing your ‘investigative skills’. In the past
this has been assessed in a separate paper entitled ‘Written Alternative to
Coursework’
Edexcel 4CH0 IGCSE Specification
1. Principles of chemistry
a) States of matter
1.1 understand the arrangement, movement and energy of the particles in each of the three
states of matter: solid, liquid and gas;
Solid Particles closely packed & vibrate
Liquid Particles are still mainly touching, but gaps
have appeared. Particles are mobile
Gas Particles are far apart. Almost no forces of
attraction. Particles are mobile
1.2 describe how the interconversion of solids, liquids and gases are achieved and recall the
names used for these interconversions;
1.3 describe the changes in arrangement, movement and energy of particles during these
interconversions;
Melting Particles in solid vibrate faster and faster until the forces of attraction are no
longer strong enough to hold them together
Freezing Particles move more slowly until they are slow enough that forces of attraction
hold them in a solid
Boiling Particles move fast enough to break all forces of attraction between particles in a
liquid
Evaporation Fast particles near the surface have enough energy to break away and form a
gas. This does not happen at the boiling point
Sublimation Particles in solid vibrate faster and faster until they move fast enough to break
all forces of attraction
melting
Solid
Liquid Gas
freezing
condensation
boiling
sublimation
sublimation
b) Atoms
1.4 describe simple experiments leading to the idea of the small size of particles and their
movement including;
dilution of coloured solutions – e.g. dilution of potassium manganate(VII)
diffusion experiments – diffusion of bromine gas into air
1.5 understand the terms atom and molecule;
Atom Smallest particle of an element, made up of
protons, neutrons and electrons.
Molecule Two or more atoms chemically bonded to form
a discrete entity.
1.6 understand the differences between elements, compounds and mixtures;
Element A substance that cannot be split into anything
simpler by chemical means.
Compound Two or more elements chemically combined in
a fixed proportion. Cannot be separated by
physical means
Mixture Two or more substances mixed together in any
proportion (e.g. an alloy is a mixture not a
compound). Can be separated by techniques in
1.7 below
1.7 describe techniques for the separation of mixtures;
Simple Distillation Separates soluble solid & solvent [e.g.
brine, NaCl(aq)]
Fractional Distillation Separates mixtures of miscible liquids (e.g.
ethanol and water).
Filtration Separates insoluble solid and liquid (e.g.
sand and water)
Crystallisation Separates soluble solid and solvent [e.g.
CuSO4 (aq)]
Paper Chromatography Separates miscible liquids or solids
c) Atomic structure
1.8 recall that atoms consist of a central nucleus, composed of protons and neutrons,
surrounded by electrons, orbiting in shells;
1.9 recall the relative mass and relative charge of a proton, neutron and electron;
Particle Relative Mass Relative Charge
proton 1 +1
neutron 1 0
electron 1/1836
(do not say 0)
-1
1.10 understand the terms;
Atomic Number Number of protons
Mass Number Number of protons + number of neutrons
Isotope Atoms with the same number of protons but a
different number of neutrons
Relative atomic mass
(Ar)
The weighted average mass of all the isotopes
of an element relative to 1/12 of the mass of a
carbon-12 atom
1.11 calculate the relative atomic mass of an element from the relative abundances of its
isotopes. E.g.
Chlorine;
35Cl = 75% & 37Cl = 25%
5.3537100
2535
100
75
Ar
1.12 understand that the Periodic Table is an arrangement of elements in order of atomic
number
1.13 deduce the electronic configurations of the first twenty elements from their positions in
the Periodic Table
1.14 deduce the number of outer electrons in a main group element from its position in the
Periodic Table. E.g.
Element Group Electronic Configuration No. Outer Electrons
Carbon 4 2, 4 4
Phosphorus 5 2, 8, 5 5
Calcium 2 2, 8, 8, 2 2
d) Relative formula masses and molar volumes
1.15 calculate relative formula masses (Mr) from relative atomic masses (Ar). E.g.
Mr (CaCO3) = 40 + 12 + (3x16) = 100
1.16 & 1.17 understand the use of the term mole.
A dozen 12
A score 20
A mole 6.02 x 1023 (Avogadro’s Number)
1.18 carry out mole calculations using relative atomic mass (Ar) and relative formula mass (Mr).
e.g.
Mr
MassMoles
1.19 understand the term molar volume of a gas and use its values (24 dm3 and 24,000 cm3) at
room temperature and pressure (rtp = 25oC, 1 atm pressure) in calculations.
One mole of any gas occupies 24dm3 at rtp.
Moles
Volume24
e) Chemical formulae and chemical equations
1.20 & 1.21 write word equations and balanced chemical equations, with state symbols, to
represent the reactions studied in this specification.
1.21 use the state symbols (s), (l), (g) and (aq) in chemical equations to represent solids, liquids,
gases and aqueous solutions respectively. If you include state symbols and they are not
required you will not be penalised, so you might as well put them in. E.g.
CuCO3(s) + H2SO4(aq) → CuSO4(aq) + H2O(l) + CO2(g)
1.22 understand how the formulae of simple compounds can be obtained experimentally,
including metal oxides, water and salts containing water of crystallisation. E.g.
Copper Oxide: Measure mass before, heat in methane to reduce to copper, measure mass
after, find ratio of moles, determine empirical formula.
1.23 calculate empirical and molecular formulae from experimental data (see 1.22 above)
1.24 calculate reacting masses using experimental data and chemical equations. E.g.
Calculate the mass of carbon you would need to reduce 15.9g of copper (II) oxide.
1.25 calculate percentage yield. E.g.
In an aluminium smelting plant 45 tonnes of aluminium metal is produced from 100
tonnes of purified bauxite. What is the percentage yield?
1.26 carry out mole calculations using volumes and molar concentrations. E.g.
Calculate the volume of oxygen that is required for the complete combustion of 1dm3 of
octane vapor at rtp.
f) Ionic compounds
1.27 & 1.28 describe the formation of ions by the gain or loss of electrons and understand
oxidation as the loss of electrons and reduction as the gain of electrons. OILRIG.
Mg → Mg2+ + 2e- Oxidation
Cl2 + 2e- → 2Cl- Reduction
1.29 recall the charges of common ions in this specification. Ones to learn;
Sulphate SO42-
Carbonate CO32-
Nitrate NO3-
Hydroxide OH-
Ammonium NH4+
1.30 deduce the charge of an ion from the electronic configuration of the atom from which the
ion is formed. E.g.
Element Electronic
configuration
Ion
Calcium 2, 8, 8, 2 Ca2+
Oxygen 2, 6 O2-
Fluorine 2, 7 F- (fluoride)
1.31 explain, using dot and cross diagrams, the formation of ionic compounds by electron
transfer, limited to combinations of elements from Groups 1, 2, 3, and 5, 6, 7
1.32 understand ionic bonding as a strong electrostatic attraction between oppositely charged
ions.
1.33 understand that ionic compounds have high melting and boiling points because of
strong electrostatic forces between oppositely charged ions.
1.34 understand the relationship between ionic charge and the melting point and boiling point
of an ionic compound.
MgO has a higher mpt than NaCl (Mg2+ & O2- vs Na+ & Cl-)
1.35 describe an ionic crystal as a giant three-dimensional lattice structure held together by the
attraction between oppositely charged ions;
1.36 draw a simple diagram to represent the positions of the ions in a crystal of sodium chloride;
g) Covalent substances
1.37 & 1.38 describe the formation of a covalent bond by the sharing of a pair of electrons
between two atoms and understand covalent bonding as a strong attraction between the
bonding pair of electrons and the nuclei of the atoms involved in the bond.
1.39 explain, using dot and cross diagrams, the formation of covalent compounds by electron
sharing for the following substances:
Hydrogen
Chlorine
Hydrogen Chloride
Water
Methane
Ammonia
Oxygen
Nitrogen
Carbon Dioxide
Ethane
Ethene
1.40 & 1.41 recall that substances with simple molecular structures are gases or liquids, or solids
with low melting points & explain this phenomenon in terms of relatively weak intermolecular
forces (this is the only time when this term is appropriate).
1.42 & 1.43 explain the high melting points of substances with giant covalent structures in terms
of the breaking of many strong covalent bonds and draw simple diagrams representing the
positions of the atoms in diamond and graphite.
1.44 explain how the uses of diamond and graphite depend on their structures, limited to
graphite as a lubricant and diamond in cutting
Graphite - weak forces between layers means they can slide over each other.
Diamond - rigid 3-D network of strong covalent bonds means it is the hardest naturally
occurring substance.
h) Metallic crystals
1.45 describe a metal as a giant structure of positive ions surrounded by a sea of delocalised
electrons;
1.46 explain the malleability and electrical conductivity of a metal in terms of its structure
and bonding;
Malleability Regular packing of positive metal ions
makes it simple for atoms to slide over
each other.
Conductivity Delocalised electrons are free to move
throughout the structure
i) Electrolysis
1.47 understand an electric current as a flow of electrons or ions (i.e mobile charged particles)
1.48 understand why covalent compounds do not conduct electricity
No mobile electrons or ions
Graphite is the exception. It has delocalised electrons between its hexagonal layers.
1.49 understand when ionic compounds conduct electricity;
Ionic compounds conduct when in solution – ions are mobile.
Ionic compounds conduct when molten – ions are mobile.
Ionic compounds do not conduct when solid – ions are not mobile.
1.50 describe simple experiments to distinguish between electrolytes and non-electrolytes. E.g.
Can the sample complete a circuit when two electrodes are put in?
1.51 recall that electrolysis (‘splitting up with electricity’) involves the formation of new
substances when ionic compounds conduct electricity. E.g.
Extraction of aluminium from bauxite (2Al2O3 → 4Al + 3O2).
1.52 describe simple experiments for the electrolysis, using inert electrodes, of molten salts
such as lead(II) bromide;
1.53 describe simple experiments for the electrolysis, using inert electrodes, of aqueous
solutions of sodium chloride, copper(II) sulphate and dilute sulphuric acid and
predict the products.
Anode = positive electrode, attracts negative ions
Cathode = negative electrode, attracts positive ions
Solution Anode Product Cathode Product
NaCl Cl2(g) H2(g)
CuSO4 O2(g) Cu(s)
H2SO4 O2(g) H2(g)
1.54 write ionic half-equations representing the reactions at the electrodes during electrolysis; e.g.
Cu2+(aq) + 2e- → Cu(s)
2Cl-(aq) → Cl2(g) + 2e-
2H2O(l) → O2(g) + 4H+(aq) + 4e-
1.55 recall that one faraday represents one mole of electrons (approx. 96,000 C) and calculate the
amounts of the products of the electrolysis of molten salts and aqueous solutions.
Q = IT
A solution of copper (II) sulphate is electrolysed for 15 minutes with 0.20 A. What mass
of copper is produced and on which electrode?
2. Chemistry of the elements
a) The Periodic Table
2.1 understand the terms group and period
Group = column
Period = row
2.2 recall the positions of metals and non-metals in the Periodic Table
Non-metals top right
Metals – the rest!
2.3 explain the classification of elements as metals or non-metals on the basis of their
electrical conductivity and the acid-base character of their oxides
Metals Conductors Basic oxides
Non-metals Non-conductors (except graphite) Acidic oxides
2.4 understand why elements in the same group of the Periodic Table have similar chemical
properties
Same number of electrons in outer shell.
2.5 recall the noble gases (Group 0) as a family of inert gases and explain their lack of
reactivity in terms of their electronic configurations.
Full outer shell of electrons.
b) The Group 1 elements - lithium, sodium and potassium – The Alkali Metals
2.6 describe the reactions of these elements with water and understand that the reactions
provide a basis for their recognition as a family of elements.
2Na(s) + 2H2O(l) → 2NaOH(aq) + H2(g)
Sodium moves randomly on surface of water.
Melts into a sphere.
Effervescence.
Sodium eventually disappears.
2.7 & 2.8 recall & explain the relative reactivities of the elements in Group 1
Cs > Rb > K > Na > Li
Reactivity increases down the group.
As we move down the group there is less electrostatic attraction between the outer
electron and the nucleus because it is further away. As a result the outer electron in easier
to remove.
.
c) The Group 7 elements - chlorine, bromine and iodine – The Halogens
2.9 & 2.10 recall the colours and physical states of the elements at room temperature and make
predictions about other halogens in the group.
Halogen Colour State
Fluorine Yellow-green Gas
Chlorine Green Gas
Bromine Brown Liquid
Iodine Grey Solid
Astatine Black Solid
2.11& 2.12 understand the difference between hydrogen chloride gas and hydrochloric acid and
explain, in terms of dissociation, why hydrogen chloride is acidic in water but not in
methylbenzene
Hydrogen chloride gas = HCl(g) – covalent molecule
Hydrochloric acid – HCl(aq) – dissociated in water [HCl(aq) → H+(aq) + Cl-(aq),
hence acidic].
Hydrogen chloride in methylbenzene is not dissociated hence not acidic.
2.13 recall the relative reactivities of the elements in Group 7.
F2 > Cl2 > Br2 > I2 > At2
Most reactive element at top of group (unlike alkali metals)
2.14 describe experiments to show that a more reactive halogen will displace a less reactive
halogen from a solution of one of its salts. E.g.
Cl2(aq) + 2KI(aq) → I2(aq) + 2KCl(aq)
Pale green solution → brown solution.
2.15 understand these displacement reactions as redox reactions. E.g.
Cl2(aq) + 2e- → 2Cl-(aq) Reduction
2I-(aq) → I2(aq) + 2e- Oxidation
d) Oxygen and oxides
2.16 recall the gases present in air and their approximate percentage by volume;
N2 78%
O2 21%
Ar 1%
CO2 0.04%
2.17 describe how experiments involving the reactions of elements such as copper, iron and
phosphorus with air can be used to determine the percentage by volume of oxygen in air. E.g.
2.18 describe the laboratory preparation of oxygen from hydrogen peroxide.
G = H2O2(aq)
H = MnO2(s) – manganese(IV) oxide catalyst
2H2O2(aq) → 2H2O(l) + O2(g)
2.19 describe the reactions with oxygen in air of magnesium, carbon and sulphur, and the
acidbase character of the oxides produced.
2Mg(s) + O2(g) → 2MgO(s)
Bright light, grey solid → white basic solid.
C(s) + O2(g) → CO2(g)
Acidic gas produced, CO produced if insufficient supply of oxygen.
S(s) + O2(g) → SO2(g)
Yellow solid → colourless acidic gas
S8(s) not used when writing balanced chemical equations.
2.20 describe the laboratory preparation of carbon dioxide from calcium carbonate and dilute
hydrochloric acid
CaCO3(s) + 2HCl(aq) → CaCl2(aq) + H2O(l) + CO2(g)
2.21 describe the formation of carbon dioxide from the thermal decomposition of metal
carbonates such as copper(II) carbonate;
CuCO3(s) → CuO(s) + CO2(g)
Green solid → Black solid + colourless gas
2.22 recall the properties of carbon dioxide, limited to its solubility and density;
Slightly soluble in water to produce carbonic acid [H2CO3(aq)].
Denser than air (collect by downward delivery).
2.23 explain the use of carbon dioxide in carbonating drinks and in fire extinguishers, in terms
of its solubility (fizzy drinks) and density (fire extinguishers).
2.24 recall the reactions of carbon dioxide and sulphur dioxide with water to produce acidic
Solutions;
CO2(g) + H2O(l) → H2CO3(aq) Carbonic acid
SO2(g) + H2O(l) → H2SO3(aq) Sulphurous acid
2.25 recall that sulphur dioxide and nitrogen oxides are pollutant gases which contribute to acid
rain, and describe the problems caused by acid rain
Damages trees
Makes lakes acidic – killing fish
Limestone buildings damaged
e) Hydrogen & water
2.26 describe the reactions of dilute hydrochloric and dilute sulphuric acids with magnesium,
aluminium, zinc and iron
Acid + Metal → Salt + Hydrogen
Effervescence observed, metal disappears.
HCl(aq) gives chloride salt of metal
H2SO4(aq) gives sulphate salt of metal.
2.27 describe the combustion of hydrogen
2H2(g) + O2(g) → 2H2O(l) Squeaky pop!
2.28 describe the use of anhydrous copper(II) sulphate in the chemical test for water
White → Blue
2.29 describe a physical test to show whether water is pure.
Boils at 100oC
f) Reactivity series
2.30 recall that metals can be arranged in a reactivity series based on the reactions of the metals
and their compounds: potassium, sodium, lithium, calcium, magnesium, aluminium, zinc,
iron, copper, silver and gold;
K > Na > Li > Ca > Mg > Al > Zn > C > Fe > H > Cu > Ag > Au
2.31 describe how reactions with water and dilute acids can be used to deduce the following
order of reactivity: potassium, sodium, lithium, calcium, magnesium, zinc, iron, and
copper
The more reactive the element the more vigorous the reaction.
Cu will not react with dilute acids as it is below H on the reactivity series.
2.32 deduce the position of a metal within the reactivity series using displacement reactions
between metals and their oxides, and between metals and their salts in aqueous solutions.
2.33 understand oxidation and reduction as the addition and removal of oxygen respectively
2.34 understand the terms: redox, oxidising agent and reducing agent
2Al(s) + Fe2O3(s) → 2Fe(l) + Al2O3(s) Thermite
Aluminium oxidised, iron reduced.
Aluminium reducing agent, iron(III) oxide oxidising agent.
Zn(s) + CuSO4(aq) → Cu(s) + ZnSO4(aq)
Zinc oxidised, copper reduced.
Zinc reducing agent, copper(II) sulphate oxidising agent.
2.35 recall the conditions under which iron rusts
Water and oxygen present
Reaction rate increases in presence of salt solution.
2.36 describe how the rusting of iron may be prevented by grease, oil, paint, plastic (barrier
protection) and galvanising (sacrificial protection).
2.37 understand the sacrificial protection of iron in terms of the reactivity series.
Zinc more reactive than iron, reacts preferentially.
g) Tests for ions and gases
2.38 describe simple tests for the cations;
Cation Test
Li+ Red flame
Na+ Strong yellow/orange flame
K+ Lilac flame
Ca2+ Brick red flame
NH4+ Add NaOH(aq), warm and test NH3(g) evolved with damp red litmus. Positive test will
go blue.
Cu2+ Add NaOH(aq), positive test = blue gelatinous ppt [Cu(OH)2(s)]
Fe2+ Add NaOH(aq), positive test = green ppt [Fe(OH)2(s)] (oxidises to Fe(OH)3(s) on
standing in air)
Fe3+ Add NaOH(aq), positive test = orange/brown ppt [Fe(OH)3(s)]
2.39 describe simple tests for the anions:
Anion Test
Cl- Add AgNO3(aq) & HNO3(aq), white ppt, Ag+(aq) + Cl-(aq) → AgCl(s)
Br- Add AgNO3(aq) & HNO3(aq), cream ppt, Ag+(aq) + Br-(aq) → AgBr(s)
I- Add AgNO3(aq) & HNO3(aq), yellow ppt, Ag+(aq) + I-(aq) → AgI(s)
SO42- Add BaCl2(aq) & HCl(aq), white ppt, Ba2+(aq) + SO4
2-(aq) → BaSO4(s)
CO32- Add HNO3(aq), or other acid, effervescence turns limewater cloudy
2.40 describe simple tests for the gases:
Gas Test
H2 Lit splint goes out with squeaky pop
O2 Relights glowing splint
CO2 Bubble through limewater. Goes cloudy.
NH3 Damp red litmus goes blue.
Cl2 Damp red or damp blue litmus bleached
3. Organic chemistry
3.1 explain the terms;
Homologous series Group of compounds with same general formula, similar chemical
properties & gradually changing physical properties.
Hydrocarbon Compound containing hydrogen & carbon only.
Saturated Compound with no C=C double bonds.
Unsaturated Compound with at least one C=C double bond.
General formula Formula to represent a homologous series, e.g. CnH2n+2
Isomer Compound with the same molecular formula, but different
structural formula.
a) Alkanes
3.2 recall that alkanes have the general formula CnH2n+2
3.3 draw displayed formulae for alkanes with up to five carbon atoms in a molecule, and name
the straight-chain isomers
Alkane Molecular Formula Displayed formula
Methane CH4
Ethane C2H6
Propane C3H8
Butane C4H10
Pentane C5H12
3.4 recall the products of the complete and incomplete combustion of alkanes
Complete combustion = CO2(g) + H2O(l)
Incomplete combustion = CO(g) [poisonous], C(s)[soot], H2O(l)
3.5 recall the reaction of methane with bromine to form bromomethane in the presence of UV
light
CH4(g) + Br2(l) → CH3Br(g) + HBr(g)
Brown colour disappears
b) Alkenes
3.6 recall that alkenes have the general formula CnH2n
3.7 draw displayed formulae for alkenes with up to four carbon atoms in a molecule, and
name the straight-chain isomers.
Alkene Molecular Formula Displayed formula
Ethene C2H4
Propene C3H6
Butene C4H8
3.8 describe the addition reaction of alkenes with bromine, including the decolourising of
bromine water as a test for alkenes.
Br2(aq) + C2H4(g) → C2H4Br2(l)
Brown → colourless (bromine water decolourises)
c) Ethanol
3.9 describe the manufacture of ethanol by passing ethene and steam over a phosphoric acid
catalyst at a temperature of about 300°C and a pressure of about 60–70 atm.
CH2CH2(g) + H2O(g) → CH3CH2OH(l)
3.10 describe the manufacture of ethanol by the fermentation of sugars, for example
glucose, at a temperature of about 30°C
C6H12O6(aq) → 2C2H5OH(aq) + 2CO2(g)
Anaerobic conditions (no oxygen)
T = 30oC
Yeast (or other biological catalyst / enzyme)
Dissolved in water
Maximum purity is 15% as yeast is killed at a greater %.
Distillation gives 96% pure ethanol.
3.11 evaluate the factors relevant to the choice of method used in the manufacture of
ethanol, for example the relative availability of sugar cane and crude oil
Fermentation Hydration of ethene
Resources Renewable (sugar beet, sugar
cane, maize etc.). Best for
countries with available land
for growing crops.
Non renewable (from crude
oil). Best for countries with
access to crude oil.
Type of process Batch Continuous
Rate Slow (several days) Fast
Purity of product 15% max (needs distilling) 100%
Conditions Gentle temp and no high
pressure
High temp and pressure,
needing a high input of
energy, hence costly.
3.12 describe the dehydration of ethanol to ethene, using aluminium oxide catalyst.
C2H5OH(g) → C2H4(g) + H2O(l)
4. Physical chemistry
a) acids & alkalis
4.1 describe the use of the following indicators;
Indicator In Acid In Alkali
Litmus Red Blue
Phenolphthalein Colourless Pink
Methyl Orange Red Yellow
4.2 & 4.3 understand how the pH scale, from 0–14, can be used to classify solutions as strongly
acidic, weakly acidic, neutral, weakly alkaline or strongly alkaline and describe the use of
universal indicator to measure the approximate pH value of a solution.
4.4 define acids as sources of hydrogen ions, H+, and alkalis as sources of hydroxide ions,
OH¯
HCl(aq) → H+(aq) + Cl-(aq)
NaOH(aq) → Na+(aq) + OH-(aq)
4.5 predict the products of reactions between dilute hydrochloric, nitric and sulphuric acids; and
metals, metal oxides and metal carbonates (excluding the reactions between nitric acid and
metals)
Acid + metal → salt + hydrogen
Acid + metal oxide (base) → salt + water
Acid + metal hydroxide (alkali) → salt + water
Acid + metal carbonate → salt + water + carbon dioxide
4.6 recall the general rules for predicting the solubility of salts in water
all common sodium, potassium and ammonium salts are soluble (SPA rule)
all nitrates are soluble
common chlorides are soluble, except silver chloride (e.g. test for Cl-)
common sulphates are soluble, except those of barium and calcium (e.g. test for SO42-)
common carbonates are insoluble, except those of sodium, potassium and ammonium
4.7 describe how to prepare soluble salts from acids e.g.
Add insoluble metal carbonate to acid in excess to ensure all acid has reacted.
Filter excess metal carbonate.
Evaporate half water and leave to crystallise.
Filter and dry crystals
4.8 describe how to prepare insoluble salts using precipitation reactions
Add reactants together in equal proportions
Filter solid product
Wash and dry solid product
4.9 describe how to carry out acid-alkali titrations, to prepare a soluble salt of sodium, potassium
or ammonium.
Titration!
Add metal hydroxide to conical flask using 25.0cm3 pipette.
Add a few drops of named indicator (e.g. methyl orange)
Add acid to burette.
Add acid to metal hydroxide solution, drop-wise near endpoint. State colour change (e.g.
orange → red)
Repeat titration without indicator.
Evaporate half water and leave to crystallise.
Filter and dry crystals.
b) energetics
4.10 recall that chemical reactions in which heat energy is given out are described as
exothermic and those in which heat energy is taken in are endothermic.
4.11 describe simple calorimetry experiments for reactions, such as combustion, displacement,
dissolving and neutralisation in which heat energy changes can be calculated from
measured temperature changes.
E = mcT
c is the specific heat capacity of the water you are heating (4.18 J K-1 g-1)
m is the mass of the water you are heating.
4.12 calculate molar enthalpy change from heat energy change
molar enthalpy change = heat energy change ÷ moles
4.13 understand the use of ΔH to represent molar enthalpy change for exothermic and
endothermic reactions
exothermic – H negative
endothermic – H positive
4.14 & 4.19 represent exothermic and endothermic reactions on a simple energy level diagram,
along with activation energy;
4.15 & 4.16 recall that the breaking of bonds is endothermic and that the making of bonds is
exothermic and use average bond energies to calculate the enthalpy change during a simple
chemical reaction
Breaking bonds requires energy
Making bonds releases energy
c) Rates of reaction
4.17 describe experiments to investigate the effects of changes in surface area of a solid,
concentration of solutions, temperature and the use of a catalyst on the rate of a reaction
4.18 & 4.20 & 4.21 describe and explain the effects of changes in surface area of a solid,
concentration of solutions,pressure of gases, temperature and the use of a catalyst on the rate of a
reaction.
Factor Effect
Surface area of solid Increasing surface area (smaller particles)
increases rate, due to higher frequency of
collisions between reactant particles.
Concentration Increasing concentration speeds up reaction
due to greater number of particles per unit
volume, means higher frequency of collisions
between reactant particles.
Pressure of gases Increasing pressure speeds up reaction due to
greater number of particles per unit volume,
means higher frequency of collisions between
reactant particles.
Temperature Increasing temperature increases collision
frequency as particles move faster, but also
increases proportion of successful collisions as
collisions have more energy.
Catalyst Lowers activation energy by providing an
alternative reaction pathway.
d) Equilibria
4.22 recall that some reactions are reversible and are indicated by the symbol ⇌ in equations.
4.23 describe reversible reactions such as the dehydration of hydrated copper(II) sulphate and the
effect of heat on ammonium chloride.
CuSO4.5H2O(s) ⇌ CuSO4(s) + 5H2O(l)
Blue → White
NH4Cl(s) ⇌ NH3(g) + HCl(g)
White solid → pungent gas → white solid reforms near top of test tube
4.24 understand the concept of dynamic equilibrium.
Forwards and reverse reaction occurring at same rate
4.25 predict the effects of changing the pressure and temperature on the equilibrium position in
reversible reactions.
Equilibrium responds to counter any change made.
Increasing pressure favours side with fewer moles.
Increasing concentration of reactants drives the equilibrium to the right, the
product side.
Increasing the temperature drives the equilibrium in the endothermic direction.
5. Chemistry in Society
a) Extraction and uses of metals
5.1 explain how the methods of extraction of the metals in this section are related to their
positions in the reactivity series.
5.2 describe and explain the extraction of aluminium from purified aluminium oxide by
electrolysis, including;
the use of molten cryolite as a solvent and to decrease the required operating temperature
the need to replace the positive electrodes [graphite reacts with O2(g) to form CO2(g)]
the cost of the electricity as a major factor
5.3 write ionic half-equations for the reactions at the electrodes in aluminium extraction
Al3+ + 3e- → Al
2O2- → O2 + 4e-
5.4 describe and explain the main reactions involved in the extraction of iron from iron ore
(haematite), using coke, limestone and air in a blast furnace
C + O2 → CO2 generates heat
CO2 + C → 2CO
Fe2O3 + 3CO → 2Fe + 3CO2
CaCO3 → CaO + CO2
CaO + SiO2 → CaSiO3 removes impurities
5.5 explain the uses of aluminium and iron, in terms of their properties. E.g.
Al – plane fuselage – low density
Fe (steel) – bridges – high tensile strength
b) Natural oil and gas
5.6 recall that crude oil is a mixture of hydrocarbons
5.7 & 5.8 describe how the industrial process of fractional distillation separates crude oil into
fractions and recall the names and uses of the main fractions obtained from crude oil: refinery
gases, gasoline, kerosene, diesel, fuel oil and bitumen;
5.9 describe the trend in boiling point and viscosity of the main fractions
Heavier fractions – greater boiling points, greater viscosity
5.10 recall that incomplete combustion of fuels may produce carbon monoxide and explain that
carbon monoxide is poisonous because it reduces the capacity of the blood to carry oxygen.
5.11 recall that, in car engines, the temperature reached is high enough to allow nitrogen and
oxygen from air to react, forming nitrogen oxides which contribute to acid rain.
5.12 & 5.13 recall that fractional distillation of crude oil produces more long-chain hydrocarbons
than can be used directly and fewer short-chain hydrocarbons than required and 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°C. e. g.
C18H38 → 2C6H12 + C6H14
c) Synthetic polymers
5.14 recall that an addition polymer is formed by joining up many small molecules called
monomers.
5.15 draw the repeat unit of addition polymers, including poly(ethene), poly(propene) and
poly(chloroethene).
5.16 deduce the structure of a monomer from the repeat unit of an addition polymer.
5.17 recall that nylon is a condensation polymer.
5.18 understand that the formation of a condensation polymer is accompanied by the
release of a small molecule such as water or hydrogen chloride.
5.19 recall the types of monomers used in the manufacture of nylon (a polyamide).
5.20 draw the structure of nylon in block diagram format.
d) The manufacture of some important chemicals
5.21 recall that nitrogen from air, and hydrogen from natural gas or the cracking of hydrocarbons,
are used in the manufacture of ammonia.
N2(g) + 3H2(g) ⇌ 2NH3(g)
5.22 describe the manufacture of ammonia by the Haber process, including the essential
Conditions;
temperature of about 450°C
a pressure of about 200 atmospheres
an iron catalyst
5.23 understand how the cooling of the reaction mixture liquifies the ammonia produced and
allows the unused hydrogen and nitrogen to be recirculated.
5.24 recall the use of ammonia in the manufacture of nitric acid and fertilisers.
5.25 recall the raw materials used in the manufacture of sulphuric acid;
Sulphur from volcanoes (Poland/USA)
Air
Water
5.26 describe the manufacture of sulphuric acid by the contact process, including the essential
conditions;
S(l) + O2(g) → SO2(g)
2SO2(g) + O2(g) ⇌ 2SO3(g)
SO3(g) + H2SO4(98%) → H2S2O7(l) Oleum (fuming sulphuric acid)
H2S2O7(l) + H2O(l) → 2H2SO4(l)
a temperature of about 450 °C
a pressure of about 2 atmospheres
a vanadium(V) oxide catalyst
5.27 recall the use of sulphuric acid in the manufacture of detergents, fertilisers and paints
5.28 describe the manufacture of sodium hydroxide and chlorine by the electrolysis of
concentrated sodium chloride solution (brine) in a diaphragm cell;
5.29 write ionic half-equations for the reactions at the electrodes in the diaphragm cell.
2Cl- → Cl2 + 2e- Anode
2H+ + 2e- → H2 Cathode
5.30 recall important uses of sodium hydroxide, including the manufacture of bleach, paper and
soap; and of chlorine, including sterilising water supplies and in the manufacture of bleach and
hydrochloric acid.