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.