GCSE
WJEC Eduqas GCSE in
CHEMISTRY
ACCREDITED BY OFQUAL
GUIDANCE FOR TEACHING
Teaching from 2016
This Ofqual regulated qualification is not available for
candidates in maintained schools and colleges in Wales.
Contents
1 – PURE SUBSTANCES AND MIXTURES
3
2 – PARTICLES AND ATOMIC STRUCTURE
10
3 – CHEMICAL FORMULAE, EQUATIONS AND AMOUNT OF SUBSTANCE
12
4 – THE PERIODIC TABLE AND PROPERTIES OF ELEMENTS
14
5 – BONDING, STRUCTURE AND PROPERTIES
23
6 – REACTIVITY SERIES AND EXTRACTION OF METALS
26
7 – CHEMISTRY OF ACIDS
37
8 – ENERGY CHANGES IN CHEMISTRY
47
9 – RATE OF CHEMICAL CHANGE AND DYNAMIC EQUILIBRIUM
52
10 – CARBON COMPOUNDS
66
11 – PRODUCTION, USE AND DISPOSAL OF IMPORTANT CHEMICALS AND
MATERIALS
69
12 – THE EARTH AND ITS ATMOSPHERE
76
2
1 – PURE SUBSTANCES AND MIXTURES
Spec Statement
Comment
(a)
explain what is meant by the
A pure substance (element or compound) contains only one
purity of a substance,
component. 'Pure' orange juice is not pure in the scientific
distinguishing between the
sense.
scientific and everyday use of the
term 'pure'
(b)
use melting point data to
distinguish pure from impure
substances
A pure substance melts at a fixed temperature. The presence
of impurities always lowers the melting point of a substance.
The melting point of a pure substance is sharp and impure
substances tend to melt over a small temperature range.
(c)
explain the differences between
elements, compounds and
mixtures
Elements are substances made up of only one type of atom.
Compounds are substances made of two or more different
types of atom that are chemically joined. They have
completely different properties to their constituent elements.
Mixtures consist of two or more substances not chemically
joined. The properties of the substances in a mixture remain
unchanged.
(d)
explain that many useful
materials are formulations of
mixtures, e.g. food and drink
products, medicines,
sunscreens, perfumes and paints
Many everyday products are mixtures in which each chemical
has a particular purpose. Formulations are made by mixing
the components in measured quantities to ensure that the
product has the required properties.
(e)
describe, explain and exemplify
the processes of filtration,
crystallisation, simple distillation
and fractional distillation
(f)
recall that chromatography
involves a stationary and a
mobile phase and that
separation depends on the
distribution between the phases
Candidates should know that different components travel
through the stationary phase at different rates because of
different distribution between phases. In the case of paper
chromatography this is due to differences in solubility; the
more soluble a component, the further it travels on the
chromatogram.
(g)
interpret chromatograms,
including measuring R f values
Candidates are expected to recall the expression used to
calculate R f values.
(h)
suggest chromatographic
methods for distinguishing pure
from impure substances
Candidates should know that gases can be analysed by gas
chromatography (where the mobile phase is a gas) and
liquids can be analysed by paper chromatography (where the
mobile phase is a liquid).
(i)
suggest suitable purification
techniques given information
about the substances involved
3
SPECIFIED PRACTICAL WORK
•
SP1A Determination of a melting point, e.g. for naphthalene (pure substance) or candle
wax (impure substance)
•
SP1B Separation of liquids by distillation, e.g. ethanol from water, and by paper
chromatography
4
Determination of a melting point, e.g. for naphthalene (pure substance) or
candle wax (impure substance)
Introduction
The melting point of a substance can be determined by taking temperature measurements
and drawing heating and cooling curves. By carrying out this experiment, you will be able to
determine a value for the melting point of naphthalene.
Apparatus
boiling tube
naphthalene powder
thermometer
400 cm3 beaker
Bunsen burner
tripod and gauze
heat resistant mat
stopwatch
clamp stand, clamp and boss
mineral wool
Diagram of Apparatus
5
Method
1. Fill a boiling tube with naphthalene powder to a depth of 4 cm and place a plug of
mineral wool into the top.
2. Clamp the boiling tube onto a clamp stand and immerse into the water in the beaker.
3. Heat the water in the beaker to 60 °C.
4. Continuously stir the naphthalene with a thermometer.
5. Record the temperature of the naphthalene every 30 seconds until the temperature
reaches 90 °C. Record the temperature when the change of state is observed.
Analysis
1. Plot a graph of time against temperature.
2. Determine a value for the melting point of naphthalene.
Risk Assessment
Hazard
Napthalene is
harmful
Risk
Vapour could be inhaled from the
boiling tube
Control measure
Use mineral wool to plug the top of the
boiling tube to reduce the vapour in the
lab and do not heat above 90 °C
Hot water can
scald
Bunsen burner
is hot
Hot water could splash when moving
apparatus
Flame could set light to / burn
individuals or equipment
Leave to cool before moving
Broken glass is
sharp
Thermometer could break when
using to stir and could cut the skin
Light on safety flame
Heat on medium / roaring flame
Use heatproof mat
Stir gently with the thermometer
Do not use a mercury thermometer
Teacher / Technician notes

Napthalene ‒ Refer to CLEAPSS Hazcard 63
Naphthalene must not be heated in an open boiling tube. The vapour level in the room can
be reduced by inserting mineral wool into the neck of the tube. Use a fresh plug each time.
It is essential to heat the naphthalene to 90 °C, so care must be taken when using high
temperatures.
6
Practical techniques covered
C1
Use of appropriate apparatus to make and record a range of measurements
accurately, including mass, time, temperature and volume of liquids and gases.
C2
Safe use of appropriate heating devices and techniques including use of a Bunsen
burner and a water bath or electric heater.
C6
Safe use and careful handling of gases, liquids and solids, including careful mixing of
reagents under controlled conditions, using appropriate apparatus to explore
chemical changes and/or products.
7
Separation of liquids by distillation, e.g. ethanol from water, and by paper
chromatography
Introduction
A mixture of liquids can be separated by distillation as each liquid will have a different boiling
point. In this experiment, you will separate ethanol from water. Ethanol has the lower boiling
point (78 ºC) and will therefore boil at a lower temperature than water. The vapour will travel
into the condenser where, in the cooler tube, it will condense back into a liquid and be
collected in a separate flask.
Apparatus
distillation apparatus
mixture of ethanol and water
anti-bumping granules
Bunsen burner
clamp stand, clamp and boss
thermometer
2 × watch glasses or evaporating basins
250 cm3 beaker
Diagram of Apparatus
8
Method
1. Add anti-bumping granules to the round bottomed flask (the distilling flask).
2. Add approximately 50 cm3 of the mixture of ethanol and water into the distilling flask.
3. Make sure there is a steady stream of water running through the condenser (from the
bottom to the top).
4. Heat the mixture, so that it boils gently.
5. Collect 5 cm3 of the liquid produced, stop heating and allow the apparatus to cool.
6. Test the liquid with a lighted splint.
Risk Assessment
Hazard
Ethanol is
flammable
Broken glass is
sharp
Risk
Large volumes of ethanol coming in
to contact with naked flame and
igniting during the distillation
Glassware could break when placing
in clamp stand and cause a risk of
cutting
Control measure
When not distilling, Bunsen burner
should be off
Make sure all parts of apparatus are
clamped securely
Care to not overtighten
Teacher / Technician notes
Reagents:

Ethanol ‒ Refer to CLEAPSS Hazcard 40A
A mixture of 80:20, water:ethanol should be used.
Care should be taken not to boil the mixture too vigorously as there is risk of it splashing into
the condenser and it may even crack the flask. Monitor the temperature in the neck and take
care that it doesn’t go higher than approximately 85 ºC.
The ethanol will evaporate due to its lower boiling point (78 ºC) then condense in the cold
condenser back to a liquid which is then collected in the beaker.
Practical techniques covered
C2
Safe use of appropriate heating devices and techniques including use of a Bunsen
burner and a water bath or electric heater.
C4
Safe use of a range of equipment to purify and/or separate chemical mixtures
including evaporation, filtration, crystallisation, chromatography and distillation.
9
2 – PARTICLES AND ATOMIC STRUCTURE
Spec Statement
Comment
(a)
recall and explain the main
features of the particle model in
terms of the states of matter and
changes of state, distinguishing
between physical and chemical
changes
(b)
use data to predict states of
substances under given
conditions
Candidates are expected to recall the melting point (0 °C)
and boiling point (100 °C) of water but data will be given for
all other substances.
(c)
explain the limitations of the
particle model in relation to
changes of state when
particles are represented by
inelastic spheres
This model provides no explanation as to why different
substances have different melting and boiling points.
This being the case, there must be some difference
between the particles of different substances.
(d)
describe how the particle model
does not explain why atoms of
some elements react with one
another
The explanation of why some elements react with others of
course requires the proton, neutron and electron model of
the atom and an understanding of electronic structure.
(e)
recall that experimental
observations suggest that atoms
are mostly empty space with
almost all the mass in a central
nucleus
Candidates should be familiar with the Geiger-Marsden
experiment and how this led to a new model of the atom.
Detailed recall of the experiment is not required.
(f)
describe the atom as a positively
charged nucleus surrounded by
negatively charged electrons,
with the nuclear radius much
smaller than the atomic radius
(g)
recall that the nucleus includes
protons and neutrons (except in
the case of 1H)
(h)
recall that atoms and small
molecules are typically around
10‒10 m or 0.1 nm in diameter
(i)
recall the relative charges and
approximate relative masses of
protons, neutrons and electrons
Charges need only be described in terms of positive (+1),
neutral (0) and negative (‒1). Protons and neutrons are
considered to have the same mass. That mass is given a
value of 1 atomic mass unit (amu). Electrons have a
negligible mass of approximately 2000 times less than that
of a proton/neutron.
10
(j)
explain why atoms as a whole
have no electrical charge
Candidates should be able to build upon this idea to
explain the charges found on simple ions e.g. Na+, Mg2+,
Cl– and O2–.
(k)
calculate numbers of protons,
neutrons and electrons in atoms
and ions, given atomic number
and mass number of isotopes
Candidates are expected to recall the definitions for atomic
number and mass number. They should use them to give
the numbers of protons, neutrons and electrons present in
any given atom/ion. Question papers in Chemistry will not
refer to proton number or nucleon number.
(l)
describe the electronic structure
of the first 20 elements
(m)
explain how the position of an
element in the Periodic Table is
related to the arrangement of
electrons in its atoms and hence
to its atomic number
Candidates should understand that an element's group
number corresponds to the number of electrons in the
outer shell of its atoms and that the period number is the
number of occupied electron shells.
(n)
describe what is meant by
isotopes and an element’s
relative atomic mass
Candidates should be able to describe the difference
between the atoms of different isotopes, in terms of the
numbers of neutrons present. They should be able to
calculate the relative atomic mass of elements with more
than one isotope.
(o)
explain that the arrangement
proposed by Mendeleev was
based on ‘atomic weights’; in
some cases the order was not
quite correct because different
isotopes have different masses
Early periodic tables placed the elements in strict order of
atomic weight. This resulted in gaps and some elements
being placed in the wrong group. Mendeleev overcame
some of the problems by leaving gaps for elements not yet
discovered. The discovery of isotopes made it possible to
explain why the order based on atomic weights was not
always correct.
11
3 – CHEMICAL FORMULAE, EQUATIONS AND AMOUNT OF SUBSTANCE
Spec Statement
Comment
(a)
use chemical symbols to write
the formulae of elements and
simple covalent and ionic
compounds
Recall of common formulae such as H 2 O and CO 2 is assumed.
(b)
deduce the empirical formula of a
compound from the relative
numbers of atoms present or
from a model or a diagram and
vice versa
(c)
recall and use the law of
conservation of mass
The law of conservation of mass states that no atoms are lost
or made during a chemical reaction so the total mass of the
products equals the total mass of the reactants. Candidates
will be required to use their understanding of this law to solve
numeric problems based on chemical equations.
(d)
use the names and symbols of
common elements and
compounds and the law of
conservation of mass to write
formulae and balanced chemical
equations and half equations
Only higher tier candidates will be required to write half
equations.
(e)
deduce the charge on ions of
elements in groups 1, 2, 3, 6 and
7
Assessment of this point will be in the context of ions not
appearing in the table given in the examination paper, e.g. Rb+,
Sr2+ or S2‒.
(f)
use the formulae of common ions
to deduce the formula of a
compound and write balanced
ionic equations
A table of formulae for common ions (including compound ions)
will be included in all examination papers. Candidates should
be able to apply their understanding in any context. Only
higher tier candidates will be required to write ionic
equations.
(g)
describe the physical states of
products and reactants using
state symbols (s, l, g and aq)
Candidates will be told to include state symbols in questions
where a specific mark is allocated for this skill.
(h)
calculate relative formula mass
of species separately and in
balanced chemical equations
Candidates should be able to calculate relative formula masses
using A r values.
12
(i)
use a balanced equation to
calculate masses of reactants
or products
Candidates should think of this as a progression from a
balanced symbol equation – and appreciate that
considering the masses of reactants and products is a
good opportunity to check that an equation is correctly
balanced. They will usually be given the balanced
equation in examination questions on this section.
(j)
calculate the empirical formula
of a compound from reacting
mass data
Candidates will not be expected to recall the methods used
to collect this type of data but they should show an
understanding of the principles involved when a
description is provided. They should be able to deal with
questions where the percentage composition of the
compound is given, as well as examples where actual
masses are provided. Candidates must show their
working in questions of this type and should be made
aware that data collected may possibly suggest a formula
different to that which they know to be correct, e.g.
incomplete reaction of magnesium with oxygen could
provide data that gives Mg 2 O as the formula for
magnesium oxide.
(k)
deduce the stoichiometry of an
equation from the masses of
reactants and products and
explain the effect of a limiting
quantity of a reactant
(l)
recall and use the definitions
of the Avogadro constant (in
standard form) and of the mole
Candidates should know that one mole of a substance
contains the same number of particles as one mole of any
other substance. The number of particles in a mole of a
given substance is the Avogadro constant. The value of
the Avogadro constant is 6.02 × 1023 per mole. There is no
requirement to recall this value.
(m)
explain how the mass of a
given substance is related to
the amount of that substance
in moles and vice versa
Candidates are expected to recall the relationship between
number of moles and mass in grams.
(n)
describe the relationship
between molar amounts of
gases and their volumes and
vice versa, and calculate the
volumes of gases involved in
reactions, using the molar gas
volume of 24 dm3 at room
temperature and pressure
Candidates should know that one mole of any gas
occupies the same volume under the same conditions of
temperature and pressure. This is 24 dm3 at room
temperature and pressure. Recall of this value is not
required.
13
4 – THE PERIODIC TABLE AND PROPERTIES OF ELEMENTS
Spec Statement
(a)
explain how the reactions of
elements are related to the
arrangement of electrons in
their atoms and hence to their
atomic number
(b)
recall the trends in melting
point/boiling point of elements
in Groups 1, 7 and 0
(c)
recall the reactions of Group 1
elements with Group 7
elements, with oxygen and with
water
Comment
Candidates should understand that elements with
the same number of electrons in their outer shell
undergo similar chemical reactions e.g. as seen in
Group 1 and Group 7.
Candidates are expected to recall observations
made during the reactions of lithium, sodium and
potassium in each case:
• Halogens – flame colours, white products
• Air/oxygen – tarnishing of freshly cut surface
• Burning in air/oxygen – flame colours, white
products
• Water – metals floating, movement on the
water surface and whether or not a ball is
formed, hissing sound, potassium only begins
to burn (lilac flame), lithium doesn't melt as it
reacts with water.
Observation includes sounds e.g. fizzing/hissing,
but ‘hydrogen formed’ is not an observation.
The tarnishing in air and reaction with water can be
easily demonstrated in the laboratory but burning
and reaction with halogens are best observed
through video clips.
In common with all specified reactions, candidates
should be able to name products and write word
and balanced symbol equations describing those
reactions.
Candidates should know that the elements of
Group 1 are known as the alkali metals.
(d)
interpret flame tests to identify
ions of Group 1 and other
metals [lithium, sodium,
potassium, barium, calcium and
copper(II)]
The flame colours expected are as follows:
• lithium – red (brick red will not be accepted)
• sodium – yellow-orange
• potassium – lilac
• barium – apple green (green will not be
accepted)
• calcium – brick red
• copper(II) – green
14
(e)
describe the advantages of
instrumental methods of
analysis, such as atomic
absorption spectroscopy
(sensitivity, accuracy and
speed)
(f)
interpret an instrumental result
given appropriate data in chart
or tabular form, when
accompanied by a reference
set in the same form
(g)
recall the reactions of Group 7
elements with Group 1
elements and with iron, and the
displacement reactions of
halogens
Candidates should recall the colours of chlorine,
bromine and iodine in their room temperature
states. Recall of the observations made during the
reactions with iron is not required but candidates
should know that it is the iron(III) salt formed in
each case.
Candidates should appreciate that displacement
reactions provide stronger evidence for the
decreasing reactivity down Group 7 than that
gained from the elements’ reactions with iron.
Factors such as the halogens’ different states at
room temperature can make it difficult to make a
fair comparison of their reactivities by observation
of their reactions with iron but they compete
directly against one another in displacement
reactions.
Candidates should know that solutions of halides
are colourless and that displacement of bromine
and iodine results in the formation of an orangebrown solution.
Candidates should know that the elements of
Group 7 are known as the halogens.
(h)
describe tests to identify
aqueous halide ions using
silver nitrate solution
Solutions of the halides produce different coloured
precipitates on addition of Ag+(aq):
• chloride – white precipitate
• bromide – cream precipitate
• iodide – yellow precipitate
Candidates should be able to name the insoluble
compounds formed and write word and balanced
symbol equations for the reactions.
Higher tier candidates should recognise
‘spectator ions’ which take no part in the
precipitation reaction and therefore be able to
write ionic equations for the reactions.
15
(i)
recall that Group 0 elements
are completely unreactive
Candidates should know that the elements of
Group 0 are known as the noble gases.
(j)
explain the reactivities (or
otherwise) of these elements in
terms of their electronic
structures and the desire to
attain/retain a full outer electron
shell
Candidates should understand that elements with
atoms containing full outer electron shells (Group
0) are unreactive and that other elements react in
order to try to attain the same state. They should
understand, for example, that the atoms of Group 1
metals lose one electron to do so, while those of
Group 7 elements gain one electron.
(k)
explain the trend in reactivities
of elements on descending
Group 1 and Group 7
Group 1 metals become more reactive down the
group. The increasing size of the atom/distance
from the positively charged nucleus makes it easier
for the outer electron to be lost.
Group 7 elements become less reactive down the
group. The increasing size of the atom/distance
from the positively charged nucleus results in a
smaller force attracting the additional electron.
(l)
predict properties from trends
within groups
(m)
predict possible reactions and
probable reactivity of elements
from their positions in the
Periodic Table
(n)
describe tests to identify
hydrogen, oxygen and chlorine
gases
When a lit splint is placed into a jar/tube containing
hydrogen gas, a squeaky pop is observed.
When a glowing splint is placed into a jar/tube
containing oxygen gas, it re-lights. (Please note
that reference to a flame glowing more brightly
is not acceptable as a test for oxygen gas.)
When damp blue litmus paper is placed in a
jar/tube containing chlorine gas, the paper first
turns red and is then bleached white.
(o)
interpret given data to identify
species from test results
(p)
describe metals and nonmetals and explain the
differences between them on
the basis of their characteristic
physical and chemical
properties
(q)
explain how the atomic
structure of metals and nonmetals relates to their position
in the Periodic Table
Candidates should recall the general physical
properties of metals and non-metals. They should
also know that metals form basic oxides while nonmetals form acidic oxides.
16
(r)
recall the general properties of
transition metals (melting point,
density, reactivity, formation of
coloured ions with different
charges and uses as catalysts)
as exemplified by titanium,
vanadium, iron and copper
Candidates should know that the transition metals
are found in the centre of the Periodic Table and
that they display the typical metallic properties of
high melting and boiling points, malleability, high
density, good electrical and thermal conductivity.
They are not very reactive and tend to react very
slowly in the natural environment. Transition
metals have a number of other properties. Many
are useful catalysts (e.g. iron in the manufacture of
ammonia, platinum in catalytic converters). They
can form more than one type of ion, e.g. Fe2+/Fe3+,
and their compounds are often coloured.
SPECIFIED PRACTICAL WORK
•
SP4 Identification of unknown substances using flame tests and chemical tests for ions
and gases
17
Identification of unknown substances using flame tests and chemical tests for
ions and gases
Introduction
Scientists need to identify the compounds that they are working with. To do this we use a
series of chemical tests that allow us to identify the different metal or non-metal ions that are
present in a compound.
These tests include:
•
•
•
•
Flame test
Test for carbonate ions
Test for Group 7 ions
Tests for gases
Flame test
Dip a damp wooden splint
into the solid sample being
tested. Put the sample into
the hottest part of a Bunsen
flame (air-hole open).
Result
Test for carbonate ions,
CO 3 2Add dilute hydrochloric acid.
Pipette the gas formed into
the limewater.
Result
Fizzes when acid is added
Test for Group 7 ions,
Cl-, Br- and IMake a solution by dissolving
the sample in water.
Add silver nitrate solution.
Result
Precipitate
colour
white
potassium, K+
Flame
colour
lilac
sodium, Na+
yellow
bromide,
Br-
cream/pale
yellow
calcium, Ca2+
brick red
iodide, I-
yellow
lithium, Li+
red
Ion
Gas formed turns limewater
milky
Ion
chloride,
Cl-
You will be provided with 5 solid compounds, labelled A, B, C, D and E.
You will use these tests to identify the five compounds you have been given.
18
Apparatus
5 × damp wooden splints
2 × wooden splints
Bunsen burner
heat proof mat
14 × test tubes
1 × dropping pipette
5 × spatulas
silver nitrate solution
dilute hydrochloric acid
limewater
small strip of magnesium
10 vol hydrogen peroxide
manganese oxide
Method ‒ Flame test
1.
2.
3.
4.
Take a damp wooden splint and dip it into sample A.
Hold the splint in the roaring (blue) Bunsen burner flame.
Record the flame colour obtained.
Repeat for each of the samples with a separate damp splint.
Method – Test for carbonate ions
1.
2.
3.
4.
Add one of the samples to a test tube.
Half fill a second tube with limewater.
Add hydrochloric acid to the sample and quickly attach the bung and side arm tube.
Record what happens to the limewater.
Diagram of Apparatus
19
Method – Test for Group 7 ions
1. Test each of the samples that did not give a positive result for the carbonate ion for
the presence of a Group 7 ion.
2. Add a small amount of the solid to a test tube.
3. Add de-ionised water to each solid to create a solution.
4. Add silver nitrate to the solution using a dropping pipette.
5. Record the colour of the precipitate formed.
Analysis
1. Use the reference tables to identify each of the unknown compounds.
Method ‒ Test for hydrogen
1.
2.
3.
4.
Half fill a test tube with hydrochloric acid.
Add a small strip of magnesium.
Place a lighted splint in the test tube.
Record your observations.
Method ‒ Test for oxygen
1.
2.
3.
4.
Quarter fill a test tube with hydrogen peroxide.
Add a spatula of manganese oxide.
Place a glowing splint in the test tube.
Record your observations.
20
Risk Assessment
Hazard
Hydrochloric
acid is an
irritant
Limewater
is corrosive
Silver nitrate is
toxic
Hot apparatus
can burn
Hydrogen
peroxide is
corrosive
Risk
Control measure
Hydrochloric acid could get onto the
skin when adding to test tube
Wash hands immediately if any
hydrochloric acid gets onto them / wear
laboratory gloves
Wear eye protection
Hydrochloric acid could get
transferred from the hands to the
eyes
Limewater could get onto the skin
when adding to test tube
Limewater could get transferred from
the hands to the eyes
Silver nitrate could get onto the skin
when adding to test tube
Silver nitrate could get transferred
from the hands to the eyes
Burns to skin when moving Bunsen
burner
Hydrogen peroxide could get onto
the skin when adding to test tube
Hydrogen peroxide could get
transferred from the hands to the
eyes
21
Wash hands immediately if any
limewater gets onto them / wear
laboratory gloves
Wear eye protection
Wash hands immediately if any silver
nitrate gets onto them / wear laboratory
gloves
Wear eye protection
Do not touch Bunsen burner until cool
Wash hands immediately if any
hydrogen peroxide gets onto them /
wear laboratory gloves
Wear eye protection
Teacher / Technician notes
In this experiment it is important that the splints are soaked in de-ionised water not tap
water.
Each splint should be no shorter than 10 cm.
An alternative to the damp splints is to use nichrome wires held in a bung or between tongs.
Reagents for flame tests:
Calcium chloride – Refer to CLEAPSS hazcard 19A
Lithium chloride – Refer to CLEAPSS hazcard 47B
Sodium carbonate – Refer to CLEAPSS hazcard 95A
Potassium iodide – Refer to CLEAPSS hazcard 47B
Potassium bromide – Refer to CLEAPSS hazcard 47B
Other reagents:
Hydrochloric acid – Refer to CLEAPSS hazcard 47A
Limewater – Refer to CLEAPSS hazcard 18
Silver nitrate – Refer to CLEAPSS hazcard 87
Hydrogen peroxide – Refer to CLEAPSS hazcard 50
Manganese oxide – Refer to CLEAPSS hazcard 60
Students should design their own table, but a suggested table format is shown below.
Sample
Flame test
observation
Carbonate test
observation
Group 7 test
observation
Name of
compound
A
B
C
D
E
Practical techniques covered
C6
Safe use and careful handling of gases, liquids and solids, including careful mixing of
reagents under controlled conditions, using appropriate apparatus to explore
chemical changes and/or products.
C8
Use of appropriate qualitative reagents and techniques to analyse and identify
unknown samples or products including gas tests, flame tests, precipitation reactions,
and the determination of concentrations of strong acids and strong alkalis.
22
5 – BONDING, STRUCTURE AND PROPERTIES
Spec Statement
Comment
(a)
describe and compare the nature
and arrangement of chemical
bonds in ionic compounds,
simple molecules, giant covalent
structures, polymers and metals
(b)
explain ionic bonding in terms of
electrostatic forces and the
transfer of electrons
(c)
construct dot and cross diagrams
to show ionic bonding in simple
ionic substances
It should be emphasised to candidates that the dot/cross
notation should be used to ensure that it is completely
clear which electrons have been transferred in forming ions
and that no electrons should appear to be in two places at
once.
(d)
explain the physical properties of
ionic compounds in terms of their
lattice structure
Candidates should use this model to explain why ionic
compounds have high melting/boiling points, are soluble in
water and conduct electricity when dissolved or in molten
form.
(e)
explain covalent bonding in
terms of the sharing of electrons
(f)
construct dot and cross diagrams
to show covalent bonding in
simple molecules
Use of the dot/cross notation to show from which atom a
given electron has come should again be emphasised.
(g)
explain the physical properties of
simple covalent substances in
terms of intermolecular bonding
Candidates should use this model to explain why simple
molecular substances have low melting/boiling points.
They should also explain why simple covalent substances
do not conduct electricity, even in molten form.
(h)
explain metallic bonding in terms
of electrostatic forces between
the ‘sea’ of electrons/lattice of
positive ions
(i)
explain the physical properties of
metals in terms of the above
model
Candidates should use this model to explain why, in
general, metals have high melting/boiling points, are good
conductors of heat and electricity and are malleable and
ductile. Higher tier candidates should be able to
explain the trend in melting/boiling point from sodium
to magnesium to aluminium, in terms of the numbers
of electrons lost by each atom and the charges on the
ions.
23
(j)
describe the limitations of the
different representations and
models of bonding, including dot
and cross diagrams, ball and
stick models and two and three
dimensional representations
(k)
recall that carbon atoms can
form four covalent bonds
(l)
explain that the huge number of
natural and synthetic organic
compounds we use today occur
due to the ability of carbon to
form families of similar
compounds, chains and rings
Different models are useful in explaining different ideas but
none of them capture all the important details. The
strengths and weaknesses of each model should be
considered, e.g. a dot and cross diagram shows exactly
which atom in a molecule has contributed each electron to
a bond but it does not show the shape of the molecule; a
ball and stick model shows the molecular shape but little
about the nature of the bond.
(m)
explain the properties of
diamond, graphite, fullerenes
and graphene in terms of their
structure and bonding
Candidates should recognise each of these as giant
structures containing covalent bonds. Candidates should
know that the very high melting points of diamond and
graphite are a result of the strong covalent bonding
present. Their differing hardness, brittleness, lubricating
and conducting properties are a result of each carbon atom
in diamond being strongly bonded to four others whilst
each one in graphite forms only three strong bonds.
Candidates should explain these differences in terms of the
graphite carbon atoms’ fourth ‘delocalised’ or ‘free’
electron. Candidates should know that fullerenes are cage
structures made entirely of carbon atoms.
Buckminsterfullerene is the most widely-known fullerene.
Its molecules are spherical and contain 60 carbon atoms.
Research into the use of fullerenes as drug delivery
systems in the body, in lubricants and as catalysts is ongoing. Graphene has been shown to be the strongest
material ever tested and also the best electrical conductor
but although claims have been made that it will transform
technology in the future, there are as yet no commercially
available 'graphene products'.
(n)
use ideas about energy transfers
and the relative strength of
chemical bonds and
intermolecular forces to explain
the different temperatures at
which changes of state occur
The lower the amount of energy required to break bonds/
overcome forces between particles, the lower the
temperature required to cause a change of state.
24
(o)
recognise that individual atoms
do not have the same properties
as bulk materials as
demonstrated by the different
properties of diamond, graphite,
fullerenes and graphene, which
all contain carbon atoms only,
and by nano-scale silver particles
exhibiting properties not seen in
bulk silver
(p)
recall the multiplying factors milli(10‒3), micro- (10‒6) and nano(10‒9)
(q)
compare nano-scale dimensions
(in the range 1-100 nm) to typical
dimensions of atoms and
molecules
One glucose molecule (C 6 H 12 O 6 ) is around 1 nm in
diameter therefore most nano-scale materials have particle
size in the order of a few hundred atoms.
(r)
describe the surface area to
volume relationship for differentsized particles and describe how
this affects properties
Candidates should be able to calculate surface area and
volume for cubes and use given formulae to do so for
spheres.
Nano-scale materials have properties different from those
for the same materials in bulk because of their high surface
area-to-volume ratio.
(s)
describe how the uses of nanoscale particles of silver and
titanium dioxide are related to
their properties
Candidates should know that nano-sized silver particles
are antibacterial, antiviral and antifungal and that they are
used in plasters, antiseptic sprays, refrigerator linings,
socks, deodorant sprays and so on. Nano-sized titanium
dioxide particles are used in some sun screens as they
absorb and reflect UV light but are also transparent so
more appealing to consumers. Self-cleaning glass is
coated with nano-scale titanium dioxide particles. These
catalyse the breakdown of dirt in the presence of UV light
and also cause water to spread out in a thin film, rather
than forming droplets on the surface. The combined effort
of sunshine and rainwater cleans the windows!
(t)
explain the possible risks
associated with the use of nanoscale particles of silver and
titanium dioxide, and of potential
future developments in
nanoscience
Candidates should appreciate that nanomaterials currently
used have been tested to ensure that they cause no
damage to individuals or the environment, but that their
long-term effects are as yet unknown. Some people have
expressed concern that nano-scale silver (deodorants) and
titanium dioxide (sun screens) are applied to the skin and
can therefore be easily absorbed into the body. While it
has been shown that these uses are safe in the short term,
there is no certainty that exposure over many years will not
result in problems.
25
6 – REACTIVITY SERIES AND EXTRACTION OF METALS
Spec Statement
Comment
(a)
explain how the reactivity of
metals with water or dilute acids
is related to the tendency of the
metal to form its positive ion
(b)
investigate the relative
reactivities of metals by
displacement (e.g. iron nail in
copper(II) chloride solution) and
competition reactions (e.g.
thermit reaction)
(c)
deduce an order of reactivity of
metals based on experimental
results
(d)
explain that the method used to
extract a metal from its ore is
linked to its position within the
reactivity series in relation to
carbon
Gold and silver are examples of metals that are found native.
Candidates should know that the most reactive metals are
extracted by electrolysis while those towards the middle of the
reactivity series can be chemically reduced. They may be
required to use information such as, “X is more/less reactive
than carbon…” to suggest a method of extraction for any
metal. Candidates should have an awareness of the
approximate position of common metals (and carbon and
hydrogen) in the reactivity series but detailed recall is not
required.
(e)
explain reduction and oxidation
in terms of loss or gain of
oxygen, identifying which species
are oxidised and which are
reduced e.g. during thermit
reaction and in the blast furnace
Candidates should be able to recognise loss or gain of
oxygen in any given reaction. They should be precise in their
descriptions e.g. iron(III) oxide – not iron – is reduced in the
blast furnace.
(f)
explain reduction and
oxidation in terms of gain or
loss of electrons, identifying
which species are oxidised
and which are reduced e.g.
during displacement reactions
and electrolysis
Candidates should be able to recognise gain or loss of
electrons in any given reaction e.g. Pb2+ ions are reduced
during the electrolysis of lead(II) bromide because they
gain electrons to form Pb atoms; Br‒ ions are oxidised
because they lose electrons. Defining reduction and
oxidation in terms of electrons is useful when reactions
do not involve oxygen.
The emphasis is on the understanding of the processes and
not on recall of colours of elements, compounds or solutions,
although information of this nature these may be given in a
question.
26
(g)
explain the principles of
extraction of iron from iron ore in
the blast furnace, including
reduction by carbon monoxide
and the acid/base reaction that
forms slag
Candidates are expected to name each of the raw materials
that are added to the furnace and to explain why they are
needed:
• Iron ore – source of iron
• Coke – as a fuel and to produce carbon monoxide for the
reduction
• Limestone – to remove impurities (slag formation when
limestone breaks down and reacts with sand from the
rocks)
• Hot air – provides oxygen so that coke can burn
Candidates should be able to write word and balanced symbol
equations for the combustion of carbon, reduction of iron(III)
oxide by carbon monoxide, decomposition of calcium
carbonate and the neutralisation reaction between calcium
oxide and silicon dioxide.
(h)
describe electrolysis of molten
ionic compounds, e.g. lead(II)
bromide, in terms of the ions
present and reactions at the
electrodes
(i)
recall that metals (or hydrogen)
are formed at the cathode and
non-metals are formed at the
anode in electrolysis using inert
electrodes
(j)
predict the products of
electrolysis of binary ionic
compounds in the molten state
(k)
explain why and how electrolysis
is used to extract reactive metals
from their ores
Candidates should know that for electrolysis to proceed,
compounds must be melted to release their ions. They
should explain electrolysis in terms of positive ions moving
towards the cathode where they gain electrons forming metal
atoms, and negative ions moving towards the anode where
they lose electrons forming molecules of the non-metal.
Higher tier candidates should be able to write half
equations for the processes taking place at the
electrodes.
Reactive metals like sodium and aluminium are extracted by
electrolysis. Electrolysis is required because carbon cannot
displace metals higher than it in the reactivity series. This
process uses vast amounts of electricity.
27
(l)
explain the principles of
extraction of aluminium from
aluminium ore (bauxite),
including the use of cryolite
Candidates should know that aluminium oxide (from bauxite)
dissolves in molten cryolite at a temperature much lower than
its melting point, therefore saving energy. Candidates should
know that aluminium ions travel to the cathode and that they
gain electrons and form aluminium atoms, whilst oxide ions
travel to the anode and lose electrons forming oxygen gas.
They should be able to write a balanced equation for the
overall reaction taking place.
Higher tier candidates should be able to write half
equations for the processes occurring at the cathode and
the anode.
Al3+ + 3e– → Al
2O2– → O 2 + 4e–
Candidates at both tiers should know that the oxygen formed
reacts with the carbon anodes, forming carbon dioxide gas
and requiring these to be replaced frequently.
(m)
evaluate the methods of
bacterial metal extraction and
phytoextraction
Higher tier candidates should be familiar with the
following methods of copper extraction but detailed recall
is not required.
Copper ores are becoming scarce and new ways of
extracting copper from low-grade ores include bacterial
extraction (bioleaching) and phytoextraction.
Bioleaching
• bacteria absorb copper compounds to form a solution
called a leachate
Phytomining
• plants absorb copper compounds through their roots
whilst growing
• plants are cropped
• cropped plants are burned to produce ash
• ash mixed with water to form a solution containing
copper compounds
Copper can be obtained from both solutions by either
displacement or electrolysis.
Both methods avoid the usual disadvantages associated
with mining and traditional extraction methods.
Advantages include: low cost, less environmental impact
than traditional methods
Disadvantages include: extremely slow process, toxic
chemicals formed
28
(n)
describe electrolysis of water in
terms of the ions present and
reactions at the electrodes
Candidates should know that hydrogen ions travel to the
cathode and that they gain electrons and form hydrogen gas,
whilst hydroxide ions travel to the anode and lose electrons
forming oxygen gas. They should be able to explain why the
volume of hydrogen formed is twice that of oxygen. They
should be able to write a balanced equation for the overall
reaction.
Higher tier candidates should be able to write a half
equation to show the reaction taking place at the cathode
and to balance the equation (atoms and charges) for the
reaction taking place at the anode. They are not required
to recall this equation.
2H+ + 2e– → H 2
2OH– → O 2 + 2H+ + 4e‒
(o)
describe competing reactions in
the electrolysis of aqueous
solutions, e.g. copper(II) chloride,
sodium chloride and sulfuric acid,
in terms of the different species
present
Candidates should know that there are H+ and OH‒ ions
present in an aqueous solution as well as the ions from the
dissolved salt. They should know that metals lower in the
reactivity series than hydrogen are formed at the cathode. In
the case of copper(II) chloride solution, the products are
copper metal and chlorine gas. They should be able to write
a balanced equation for the overall reaction.
Higher tier candidates should be able to write half
equations for the processes occurring at the cathode and
the anode e.g.
Cu2+ + 2e– → Cu
2Cl– → Cl 2 + 2e–
Candidates at both tiers should know that when the dissolved
salt contains ions of metals higher in the reactivity series than
hydrogen, it is hydrogen gas rather than the metal that forms
at the cathode. Electrolysis of sodium chloride solution
therefore gives hydrogen gas and chlorine gas.
Electrolysis of sulfuric acid gives hydrogen gas and oxygen
gas. In this case candidates are not expected to explain the
reaction taking place at the anode.
Higher tier candidates should again be able to write half
equations showing the formation of hydrogen and
chlorine.
(p)
recall the properties of
aluminium, copper, iron and
titanium
(q)
explain how the properties of
metals are related to their uses
and select appropriate metals
given details of the usage
required
29
(r)
describe tests to identify
aqueous copper(II), iron(II) and
iron(III) ions
Candidates should know the colours of each of the hydroxide
precipitates formed. They should be able to write word and
balanced symbol equations for the reactions.
Higher tier candidates should be able to write ionic
equations for the reactions e.g.
Cu2+(aq) + 2OH‒(aq) → Cu(OH) 2 (s)
SPECIFIED PRACTICAL WORK
•
SP6A Determination of relative reactivities of metals through displacement reactions
•
SP6B Investigation into electrolysis of aqueous solutions and electroplating
30
Determination of relative reactivities of metals through displacement reactions
Introduction
Some metals are more reactive than others. In this experiment, a piece of metal is added to
a solution of a compound of another metal. A more reactive metal displaces a less reactive
metal from its compound. By carrying out this experiment, you will be investigating the
competition reactions of metals and produce a reactivity series of the metals.
Apparatus
dimple tray
100 cm3 beaker
4 × dropping pipettes
5 cm3 of each of the following at 0.1 mol/ dm3
 zinc sulfate
 magnesium sulfate
 copper(II) sulfate
 iron(II) sulfate
Approximately 1 cm length/square sample of the following metals.
 zinc
 magnesium
 copper
 iron
Diagram of Apparatus
31
Method
1. Using a dropping pipette, put a little zinc sulfate in four of the depressions of the
dropping tile. Do this for each solution in turn. Do not overfill dimples.
2. Put a piece of metal in each of the solutions, using the apparatus diagram as a guide.
3. Observe and record the changes in the solutions or metal samples.
Analysis
1.
Use your results to construct a reactivity series for the metals used. Write equations for
any reactions that occurred.
Risk Assessment
Hazard
Salt solutions
are harmful
Risk
Whilst dispensing the solutions they
can be squirted into eyes or if spilt
onto hands, solutions can be
transferred to eyes
Control measure
Wear eye protection
Wash hands when solutions spilt on to
hands
Teacher / Technician notes
Reagents
 Zinc sulfate ‒ Refer to CLEAPSS hazcard 108
 Magnesium sulfate ‒ Refer to CLEAPSS hazcard 59B
 Copper(II) sulfate ‒ Refer to CLEAPSS hazcard 27B
 Iron(II) sulfate ‒ Refer to CLEAPSS hazcard 38
 Zinc foil ‒ Refer to CLEAPSS hazcard 107
 Magnesium ribbon ‒ Refer to CLEAPSS hazcard 59A
 Copper foil ‒ Refer to CLEAPSS hazcard 26
 Iron ‒ Refer to CLEAPSS hazcard 38
Solutions may be dispensed in small beakers to each group of students or in small dropper
bottles.
Students may need two dimple trays per group, if trays do not contain 16 dimples.
Metals should be approximately 1 cm lengths/squares of ribbon or foil cleaned with an emery
cloth and as similar in size as possible.
Students will need to record which metals react with the solutions. A table may be useful.
Use a ✓ to show reactivity and a ✗ to show no reaction. The metals with the most ticks are
the most reactive.
32
Students should design their own table, but a suggested table format is shown below.
Zinc
Magnesium
Copper
Iron
Zinc sulfate
Magnesium sulfate
Copper(II) sulfate
Iron(II) sulfate
You can point out to students that there is no need to carry out the zinc/zinc sulfate,
magnesium/magnesium sulfate reactions, etc or allow them to decide for themselves if these
reactions are likely to lead to a positive result.
Remind students that they are looking for metal displacement, some solutions are slightly
acidic so bubbles of hydrogen can be seen. Explain that this doesn’t count as displacement.
Students may need to be given guidance of the sort of observations they may expect to see.
It may be best to get the class to tell you what they think the order of reactivity is while they
still have the evidence in front of them, so that discrepancies can be resolved.
There are many ways of carrying out this series of reactions. The one described here uses a
dimple tray, but it can be adapted with test tubes. The advantages of the dimple tray are the
small amounts of chemical involved and the way the results are displayed.
Practical techniques covered
C5
Making and recording of appropriate observations during chemical reactions
including changes in temperature and the measurement of rates of reaction by a
variety of methods such as production of gas and colour change.
33
Investigation into electrolysis of aqueous solutions and electroplating
Introduction
In this experiment you will carry out the electrolysis of copper(II) sulfate solution and link
your findings to industrial copper purification and copper plating.
Apparatus
250 cm3 beaker
2 × graphite electrodes (about 5 mm diameter)
clamp stand, boss and clamp
12 V d.c. power supply
leads and crocodile clips
200 cm3 copper(II) sulfate, about 0.5 mol/dm3
Diagram of Apparatus
Method
1.
2.
3.
4.
Measure 200 cm3 of copper(II) sulfate into the beaker.
Set up the apparatus as in the diagram.
Switch on the power supply.
After 2 minutes record any observations seen at the electrodes.
34
Risk Assessment
Hazard
Copper(II) sulfate
is harmful
Risk
Copper(II) sulfate splashed onto
hands whilst pouring could be
transferred to eyes
Control measure
Wear eye protection
Wash hands if copper(II) sulfate spilt on
them
Teacher / Technician notes
•
Copper(II) sulfate solution ‒ Refer to CLEAPSS hazcard 26
There are several ways of securing the graphite electrodes. Using a clamp stand and clamp
is probably the most convenient. They can also be fixed on to a small strip of wood or
cardboard resting on the top of the beaker.
A lamp can be included in the circuit to indicate that there is a flow of current.
As an extension to the basic experiment, strips of copper can be used in place of the
graphite rods.
After setting up the cell as shown students can observe changes to each of the electrodes.
They should see a deposit of copper forming on the cathode. This will often be powdery and
uneven. It can be explained that, if the current used is much lower, then the solid coating is
shiny, impermeable and very difficult to rub off; this process forms the basis of electroplating.
Bubbles of gas (oxygen) are formed at the anode.
Cathode
Cu2+(aq) + 2e-
Anode
2H 2 O(l)
Cu(s)
O 2 (g) + 4H+(aq) + 4e-
With copper electrodes, the copper anode dissolves. The reaction is the reverse of the
cathode reactions.
With graphite electrodes, the oxygen usually reacts with the anode to form CO 2 .
The results can lead to a discussion about electroplating and the electrolytic purification of
copper. It is useful to allow students to copperplate metal objects supplied by the school and
previously tested for their suitability. Personal items should not be used. In many cases, an
alternative redox reaction often takes place before any current is actually passed.
35
After doing the electrolysis as described above, the electrodes can be interchanged.
Students can then see the copper disappearing from the surface of the copper-coated
anode.
Cu(s)
Cu2+(aq) + 2e-
This leads to a discussion as to why, during electrolysis, the:
-
anode consists of an unrefined sample of the metal
cathode is made of pure copper or a support metal such as stainless steel.
Practical techniques covered
C3
Use of appropriate apparatus and techniques for conducting and monitoring chemical
reactions, including appropriate reagents and/or techniques for the measurement of
pH in different situations.
C5
Making and recording of appropriate observations during chemical reactions
including changes in temperature and the measurement of rates of reaction by a
variety of methods such as production of gas and colour change.
C6
Safe use and careful handling of gases, liquids and solids, including careful mixing of
reagents under controlled conditions, using appropriate apparatus to explore
chemical changes and/or products.
C7
Use of appropriate apparatus and techniques to draw, set up and use
electrochemical cells for separation and production of elements and compounds.
36
7 – CHEMISTRY OF ACIDS
Spec Statement
Comment
(a)
recall that acids react with
some metals and with bases
(including alkalis) and
carbonates
When an acid reacts with a metal, a solution of the
metal salt and hydrogen gas are produced. Metal
oxides and metal hydroxides are known as bases
and an alkali is a soluble base.
(b)
write equations predicting
products from given reactants
(c)
describe a test to identify
carbon dioxide gas
When carbon dioxide gas is passed
through limewater the solution turns milky. (Please
note that reference to extinguishing a lit splint or
flame is not acceptable as a test for carbon dioxide
gas.)
(d)
describe a test to identify
carbonate ions using dilute acid
Effervescence (fizzing) is observed when an acid
reacts with a carbonate. Note that ‘carbon dioxide
formed’ is not an observation.
(e)
recall that acids form hydrogen
ions when they dissolve in
water and solutions of alkalis
contain hydroxide ions
(f)
recall that acidity and alkalinity
are measured by pH and how
to measure pH using pH
indicator chart and digitally
Acids and alkalis can be classified as being either
strong or weak. Universal indicator and the pH
scale are used to for this purpose. Candidates
should recall associated colours, approximate pH
values and acid/alkali strength e.g. orange > pH
~3/4 > weak acid.
(g)
describe neutralisation as acid
reacting with base to form a salt
plus water (or with carbonate to
form a salt plus water and
carbon dioxide)
The reactions of acids with bases always produce
a metal salt and water and acids and carbonates
produce carbon dioxide gas in addition to a salt
and water. Neutralisation reactions are exothermic
and effervescence (fizzing) is observed when an
acid reacts with a carbonate.
(h)
prepare crystals of soluble salts
from insoluble bases and
carbonates
Candidates should know the method used to
prepare crystals of soluble salts from the reaction
of acids with insoluble bases and carbonates:
• excess base/carbonate to use up all acid;
• filtration to remove excess base;
• evaporation of water to form crystals.
They should know that small crystals can be
formed quickly by heating to evaporate until about
⅓ of the solution remains and leaving to cool.
Allowing the filtered solution to evaporate slowly
over a period of days results in the formation of
larger crystals.
37
(i)
use a titration method to
prepare crystals of soluble salts
and to determine relative
concentrations of strong acids
and strong alkalis
Candidates should know the method used to
prepare crystals of soluble salts from the reaction
of acids with alkalis:
• indicator and fixed volume of acid/alkali in flask;
• exact volume of alkali/acid needed for
neutralisation is measured and recorded;
• same fixed volume of acid/alkali in clean flask
and exact volume of alkali/acid needed for
neutralisation is added but with no indicator;
• evaporation of water to form crystals.
All candidates should be able to compare relative
concentrations of acid/alkali on the basis that if, for
example, 25cm3 of NaOH requires 30cm3 of HCl to
neutralise it, the alkali must be of higher
concentration than the acid.
(j)
describe test to identify
aqueous sulfate ions using
barium chloride solution
A white precipitate of barium sulfate forms when
barium chloride solution is added to a solution
containing sulfate ions. Higher tier candidates
should be able to write an ionic equation for the
reaction.
(k)
recognise that aqueous
neutralisation reactions can be
generalised to hydrogen ions
reacting with hydroxide ions to
form water
H+ + OH‒ → H 2 O
Candidates at both tiers should know that
neutralisation reactions can be summarised by this
ionic equation.
(l)
use and explain the terms
dilute and concentrated
(amount of substance) and
weak and strong (degree of
ionisation) in relation to
acids
Candidates should know that any acid (or any
solution) can be dilute or concentrated whilst
any given acid is either strong or weak, e.g.
hydrochloric acid is a strong acid (pH 1) and
ethanoic acid is a weak acid (pH 3). They
should understand that a strong acid is fully
dissociated whilst a weak acid is only partly
dissociated.
(m)
describe the observed
differences between reactions
of strong acids and weak acids
Candidates should know that weak acids, such as
ethanoic acid, react with metals, bases (including
alkalis) and carbonates in the same way as strong
acids but that the reactions occur more slowly and
are less exothermic. They should know that
ethanoic acid forms salts called ethanoates, e.g.
sodium ethanoate is formed when it reacts with
sodium hydroxide.
(n)
recall that as hydrogen ion
concentration increases by a
factor of ten the pH value of a
solution decreases by one
For example a solution with a hydrogen ion
concentration of 0.01 mol / dm3 has a pH of 2. A
0.1 mol / dm3 solution of hydrogen ions is 10
times as concentrated and has a pH of 1.
38
(o)
describe neutrality and
relative acidity and alkalinity
in terms of the effect of the
concentration of hydrogen
ions on the numerical value
of pH (whole numbers only)
For example if the hydrogen ion concentration
of a solution of pH 2 is reduced by a factor of
10 then its pH will be 3. If it reduced by a
further factor of 10 it will be 4 and so on.
(p)
explain how the
mass/number of moles of a
solute and the volume of the
solution is related to the
concentration of the solution
Candidates are required to recall the
relationship between mass/number of moles,
volume and concentration and be able to use it
appropriately.
(q)
explain how the
concentration of a solution in
mol / dm3 is related to the
mass of the solute and the
volume of the solution
Candidates should be able to interconvert
mass in grams and number of moles and
therefore express concentrations in g / dm3 as
well as mol / dm3.
(r)
explain the relationship
between the volume of a
solution of known
concentration of a substance
and the volume or
concentration of another
substance that react
completely together
Acid-base titrations should be used to
determine the exact volumes of acid/alkali
required to react completely. Titrations should
be repeated to establish the most accurate
values possible and the data should be used
alongside the relevant chemical equation to
calculate the concentration of the unknown
solution. Candidates should be familiar with
calculations relating to reactions with 1:1 and
other mole ratios.
SPECIFIED PRACTICAL WORK
•
SP7A Preparation of crystals of a soluble salt from an insoluble base or carbonate
•
SP7B Titration of a strong acid against a strong base using an indicator
39
Preparation of crystals of a soluble salt from an insoluble base or carbonate
Introduction
In this experiment you will make crystals of copper(II) sulfate. This can be done using either
copper(II) carbonate or copper(II) oxide.
copper(II)
carbonate
CuCO 3 (s)
+
sulfuric
acid
+ H 2 SO 4 (aq)
copper(II) + water + carbon
sulfate
dioxide
CuSO 4 (aq) + H 2 O(l) + CO 2 (g)
copper(II)
oxide
CuO(s)
+
copper(II)
sulfate
CuSO 4 (aq)
+
sulfuric
acid
H 2 SO 4 (aq)
Apparatus
100 cm3 beaker
stirring rod
filter funnel and paper
evaporating basin
50 cm3 measuring cylinder
0.5 mol/dm3 H 2 SO 4
copper(II) oxide or copper(II) carbonate
spatula
indicator paper
Access to:
electronic balance ± 0.1 g
Diagram of Apparatus
40
+
water
+
H 2 O(l)
Method
1. Measure 50 cm3 of sulfuric acid and pour into the beaker.
2. Measure approximately 4 g copper (II) oxide or 5 g copper(II) carbonate. (This does not
need to be precise as the solid will be in excess.)
3. Add the solid to the acid and stir thoroughly.
4. To ensure all the acid has reacted, touch the glass rod onto a piece of indicator paper.
If it is acidic continue stirring.
5. If the solution is neutral, pour the mixture into the filtration apparatus above the
evaporating basin.
6. Allow to evaporate for several days until dry.
Risk Assessment
Hazard
Sulfuric acid is
corrosive
Copper(II) sulfate is
harmful
Risk
Risk of splashing into eyes
whilst stirring
Risk of splashing into eyes
whilst stirring
Control measure
Take care whilst stirring and wear eye
protection
Take care whilst stirring and wear eye
protection
Hot tripod and
evaporating basin
can burn
Risk of burning hands when
touching hot tripod / basin
Leave apparatus to cool before moving
Teacher / Technician notes
Reagents:
•
•
•
•
Copper(II) oxide ‒ Refer to CLEAPSS hazcard 26
Copper(II) carbonate ‒ Refer to CLEAPSS hazcard 26
Sulfuric acid ‒ Refer to CLEAPSS hazcard 98A
Copper(II) sulfate solution ‒ Refer to CLEAPSS hazcard 26
50 cm3 of copper(II) sulfate solution requires medium to large evaporating basins. Quantities
can be reduced to suit available equipment. However it is vital that the solid is always in
excess.
It can be emphasised that the reason for adding the insoluble base in excess is to ensure all
of the acid has reacted and that a pure sample of the salt can thus be obtained.
At method point 6, it is possible to heat the evaporating basin to reduce the volume of
copper(II) sulfate solution by approximately a third using a Bunsen burner. This will reduce
the time needed to reach dryness.
There is also scope for extension work – the mass of the base added could be weighed
accurately and recorded. The mass of excess could then be obtained and thus the number
of moles of copper(II) sulfate produced could be calculated.
41
Practical techniques covered
C2
Safe use of appropriate heating devices and techniques including use of a Bunsen
burner and a water bath or electric heater.
C4
Safe use of a range of equipment to purify and/or separate chemical mixtures
including evaporation, filtration, crystallisation, chromatography and distillation.
42
Titration of a strong acid against a strong base using an indicator
Introduction
In this experiment sodium hydroxide is neutralised with hydrochloric acid to produce the
soluble salt, sodium chloride in solution. An indicator is used to show when neutralisation
has occurred. The solution could then be concentrated and crystallised to produce sodium
chloride crystals.
Apparatus
burette
measuring cylinder
100 cm3 conical flask
small filter funnel
white paper
dilute sodium hydroxide
dilute hydrochloric acid
indicator
clamp stand, boss and clamp or burette stand
Diagram of Apparatus
43
Method
1. Use the small funnel to fill the burette with acid. Run a little acid out into a waste
beaker to fill the part of the burette that is below the tap. Record the starting volume of
acid in the burette.
2. Accurately measure 25 cm3 of sodium hydroxide solution into a conical flask.
3. Add 2 drops of indicator.
4. Add 0.1 cm3 of acid at a time, swirl the flask after each acid addition. Keep adding acid
until the indicator changes colour. Record the final volume of acid in the burette.
5. Repeat steps 1-4 twice more.
Analysis
1. Calculate the volume of acid that was needed to neutralise the alkali in each repeat.
2. Calculate the mean volume of dilute hydrochloric acid needed to neutralise 25 cm3 of
the sodium hydroxide solution.
3. What do your results tell you about the concentration of the alkali?
Risk Assessment
Hazard
Hydrochloric
acid and
sodium
hydroxide are
corrosive
Risk
Hydrochloric acid or sodium
hydroxide spilling onto hands when
filling burette or measuring volume of
liquids
Control measure
Wear gloves
Wash hands immediately after contact
with solutions
Hydrochloric acid or sodium
hydroxide splashing into eyes when
filling burette
Wear goggles
Burette and
pipette made
from glass
which is brittle
and is sharp if
broken
Burette breaking when clamping
giving danger of cuts
Take care when clamping burette not
to overtighten
Pipette breaking when being
handled giving danger of cuts
Take care when using pipette
44
Teacher / Technician notes
Reagents:
•
•
Hydrochloric acid – Refer to CLEAPSS hazcard 47A
Sodium hydroxide – Refer to CLEAPSS hazcard 31
Sodium hydroxide and hydrochloric acid solutions do not need to be made up to a high
degree of accuracy, but should be reasonably close to the same concentration and less than
0.5 mol/dm3.
Burette stands and clamps are designed to prevent crushing of the burette by overtightening, which may happen if standard jaw clamps are used.
A white tile can be used to go under the titration flask, instead of white paper.
Students need training in using burettes correctly, including how to clamp them securely and
fill them safely. You should consider demonstrating burette technique, and give students the
opportunity to practise this. Students do not need the acid volume to start on zero in the
burette, but must ensure that the reading is not above zero.
In this experiment, a pipette is not essential and a measuring cylinder is acceptable.
However, a pipette and filler could be used to increase accuracy if desired.
There is an opportunity here with more able students to do quantitative measurements,
leading to calculations, but the primary aim is to introduce students to the titration technique
to produce a neutral solution.
Indicators you can use include screened methyl orange (green in alkali, violet in acid) and
phenolphthalein (pink in alkali, colourless in acid).
Students should design their own table, but a suggested table format is shown below.
1
Trial
2
Final volume of
acid in burette
(cm3)
Initial volume
of acid in
burette (cm3)
Titre (volume
added) (cm3)
45
3
Mean
Practical techniques covered
C1
Use of appropriate apparatus to make and record a range of measurements
accurately, including mass, time, temperature and volume of liquids and gases.
C3
Use of appropriate apparatus and techniques for conducting and monitoring chemical
reactions, including appropriate reagents and/or techniques for the measurement of
pH in different situations.
C6
Safe use and careful handling of gases, liquids and solids, including careful mixing of
reagents under controlled conditions, using appropriate apparatus to explore
chemical changes and/or products.
46
8 – ENERGY CHANGES IN CHEMISTRY
Spec Statement
Comment
(a)
distinguish between
endothermic and exothermic
reactions on the basis of the
temperature change of the
surroundings
The emphasis here should be on interpretation of
experimental data to identify exothermic and endothermic
reactions. There is no requirement for candidates to recall
examples of endothermic reactions but they do need to
know that combustion and neutralisation reactions are
exothermic.
(b)
draw and label a reaction profile
for an exothermic and an
endothermic reaction, identifying
activation energy
(c)
explain activation energy as the
energy needed for a reaction to
occur
Candidates should know that this is the minimum amount of
energy required to start a reaction.
(d)
calculate energy changes in a
chemical reaction by
considering bond making and
bond breaking energies
Candidates should be able to calculate the total amount
of energy required to break bonds and the total amount
released in forming bonds during a given reaction, and
use those values to find the overall energy change for
the reaction. They should explain that a reaction is
exothermic because more energy is released in forming
bonds than is required to break bonds, rather than by
stating simply that the overall energy change has a
negative value. They should be able to apply their
understanding to more complex questions, e.g. where
the given data is used to calculate a bond energy value.
(e)
recall that a chemical cell
produces a potential difference
until the reactants are used up
A simple cell can be made by externally connecting two
different metals immersed solutions of their own ions. Cell
potential differences vary depending on the relative
positions of the metals in the reactivity series. Cells stop
producing electrical energy when one of the reactants has
been used up.
(f)
evaluate the advantages and
disadvantages of
hydrogen/oxygen and other fuel
cells for given uses
Candidates should know that hydrogen is used in fuel cells
that are now being utilised to power cars. It is expected
that candidates are able to discuss the advantages and
disadvantages of its use as a fuel.
•
•
Advantages – produced from water therefore
renewable; water is the only product of its combustion
so burning hydrogen does not contribute towards global
warming or acid rain.
Disadvantages – requires large amounts of electricity to
produce hydrogen from water by electrolysis (how is
this generated?); storage requires bulky and heavy
pressurised containers; hydrogen is potentially
hazardous as it forms an explosive mixture with air.
47
SPECIFIED PRACTICAL WORK
•
SP8 Determination of the amount of energy released by a fuel
48
Determination of the amount of energy released by a fuel
Introduction
Fuels react with oxygen when they burn, releasing energy. You will burn four different
alcohols and compare the energy they give off.
alcohol + oxygen
carbon dioxide + water
Apparatus
clamp stand, clamp and boss
250 cm3 conical flask
100 cm3 measuring cylinder
thermometer
Access to:
electronic balance ± 0.01 g
4 × spirit burners containing methanol, ethanol, propanol, butanol
Diagram of Apparatus
49
Method
1. Measure 100 cm3 of water into the conical flask.
2. Clamp the flask at a suitable height so the spirit burner can be placed below it (as
shown in the diagram - make sure that the thermometer does not touch the bottom of
the flask).
3. Record the temperature of the water.
4. Record the mass of the spirit burner (including the lid and alcohol).
5. Place the spirit burner under the conical flask and light it.
6. Allow the burner to heat the water until the temperature rises by about 40 °C. Record
the temperature of the water.
7. Extinguish the flame carefully and record the mass of the burner.
8. Repeat steps 1-7 with each of the other alcohols.
Analysis
1. Calculate the temperature rise for each fuel.
2. Calculate the mass of each alcohol burnt.
3. Calculate the energy released for each alcohol using the following equation.
Energy released from alcohol per gram (J) =
mass of water (g)×temperature increase (℃)×4.2
mass of alcohol (g)
Risk Assessment
Hazard
Methanol is harmful
and highly
flammable
Ethanol is highly
flammable
Risk
May set light to / burn
individuals or equipment Vapour
can cause irreversible damage
May set light to / burn
individuals or equipment
Propanol is highly
flammable and an
irritant
Butanol is highly
flammable and
harmful if swallowed
May set light to / burn
individuals or equipment
Vapour may irritate respiratory
system and may irritate skin if
spilt
50
Control measure
Work in a well ventilated lab
Wear eye protection and ensure work
station is clear
Work in a well ventilated lab
Wear eye protection and ensure work
station is clear
Work in a well ventilated lab
Wear eye protection and ensure work
station is clear
Work in a well ventilated lab
Wear eye protection and ensure work
station is clear
Rinse immediately if spilt on skin
Teacher / Technician notes
Methanol ‒ Refer to CLEAPSS hazcard 40B
Ethanol ‒ Refer to CLEAPSS hazcard 40A
Propanol ‒ Refer to CLEAPSS hazcard 84A
Butanol ‒ Refer to CLEAPSS hazcard 84B
Pentanol should not be used as a fume cupboard is needed ‒ Refer to CLEAPSS hazcard
84C.
Spirit burners should not be used for more than one alcohol. Make sure that the wick fits
tightly in the holder and the holder sits tightly in the container.
Students should not fill or refill spirit burners.
An extension activity could be to plot a graph of the number of carbon atoms in the alcohol
against the energy released per gram.
No repeats are planned in this experiment, but can be carried out if time allows.
Alternatively, groups can compare results to discuss reproducibility.
Students should design their own table, but a suggested table format is shown below.
Alcohol
Initial
mass
of
burner
(g)
Final
mass
of
burner
(g)
Change
in mass
of
burner
(g)
Initial
temperature
(ºC)
Final
temperature
(ºC)
Energy
Temperature
released
increase
per gram
(ºC)
(J)
Practical techniques covered
C1
Use of appropriate apparatus to make and record a range of measurements
accurately, including mass, time, temperature and volume of liquids and gases.
C5
Making and recording of appropriate observations during chemical reactions
including changes in temperature and the measurement of rates of reaction by a
variety of methods such as production of gas and colour change.
C6
Safe use and careful handling of gases, liquids and solids, including careful mixing of
reagents under controlled conditions, using appropriate apparatus to explore
chemical changes and/or products.
51
9 – RATE OF CHEMICAL CHANGE AND DYNAMIC EQUILIBRIUM
Spec Statement
Comment
(a)
suggest practical methods for
determining the rate of a given
reaction – from gas collection,
loss of mass and precipitation
(including using data-logging
apparatus)
Candidates should recognise that a rate measures
a change over a given time. They should be
familiar with gas collection and mass loss methods
of studying the rates of reactions such as acids and
metals/carbonates, as well as the precipitation
reaction of dilute hydrochloric acid and sodium
thiosulfate.
(b)
explain any observed changes
in mass in non-enclosed
systems during a chemical
reaction using the particle
model
(c)
interpret rate of reaction graphs
Candidates should be able to draw a tangent to a
curve and use this to calculate rate at a given
point.
(d)
describe the effect of changes
in temperature, concentration
(pressure) and surface area on
rate of reaction
The rate of reaction is increased by increasing
temperature, concentration (pressure) and surface
area. Candidates should appreciate that
decreasing solid particle size increases surface
area.
(e)
explain the effects on rates of
reaction of changes in
temperature and concentration
(pressure) in terms of
frequency and energy of
collision between particles
Candidates should understand that particles of
reactants must collide in order for a reaction to
occur and that these collisions must have energy
greater than the activation energy to be
'successful'. The greater the number of successful
collisions in a given time, the faster the
reaction/higher the rate.
(f)
explain the effects on rates of
reaction of changes in the size
of the pieces of a reacting solid
in terms of surface area to
volume ratio
(g)
describe the characteristics of
catalysts and their effect on
rates of reaction
A catalyst is a substance that increases the rate of
a reaction while remaining chemically unchanged.
(h)
identify catalysts in reactions
Candidates are not expected to recall the names of
specific catalysts other than those named in other
parts of the specification, e.g. iron in the Haber
process. They should know that a catalyst does
not appear as a reactant in a chemical equation.
52
(i)
explain catalytic action in terms
of activation energy
Catalysts increase the rate of a reaction by
lowering the minimum energy required for
successful collisions.
(j)
recall that enzymes act as
catalysts in biological systems
Candidates should understand what is meant by an
enzyme’s optimum temperature and that enzymes
are denatured at high temperature e.g. over about
60°C. No explanation of how enzymes work, i.e.
the lock and key idea, is expected.
(k)
recall that some reactions may
be reversed by altering reaction
conditions
Candidates should know that a reversible reaction
is one that can go in either direction. Certain
conditions may favour the forward reaction whilst
others favour the backward reaction. The ⇌
symbol is used to represent a reversible reaction.
(l)
recall that dynamic equilibrium
occurs when the rates of
forward and reverse reactions
are equal
When a reversible reaction occurs in a closed
system, dynamic equilibrium is reached when the
forward and backward reactions occur at exactly
the same rate. Once equilibrium is reached, the
concentrations of the reactants and products
remain constant (but not necessarily equal).
(m)
predict the effect of changing
reaction conditions
(concentration, temperature
and pressure) on equilibrium
position and suggest
appropriate conditions to
produce a particular product
Candidates should recall Le Chatelier's
principle and be able to apply it to any given
example.
SPECIFIED PRACTICAL WORK
•
SP9A Investigation into the effect of one factor on the rate of a reaction using a gas
collection method
•
SP9B Investigation into the effect of one factor on the rate of the reaction between dilute
hydrochloric acid and sodium thiosulfate
•
SP9C Investigation into the effect of various catalysts on the decomposition of hydrogen
peroxide
53
Investigation into the effect of one factor on the rate of reaction using a gas
collection method
Introduction
Magnesium reacts with dilute hydrochloric acid to produce hydrogen. The equation for the
reaction is as follows:
magnesium
Mg(s)
+ hydrochloric
acid
+
magnesium
chloride
2HCl(aq)
MgCl 2 (aq)
+
+
hydrogen
H 2 (g)
The rate at which the hydrogen gas is produced can be used to determine the rate of the
reaction.
In this experiment you will study the effect of changing the concentration of the hydrochloric
acid on the rate of the reaction.
Apparatus
250 cm3 conical flask
single-holed rubber bung
delivery tube to fit conical flask
trough or plastic washing-up bowl
100 cm3 measuring cylinder
250 cm3 measuring cylinder
clamp stand, boss and clamp
stopwatch
magnesium ribbon in 3 cm lengths
1 mol/dm3 hydrochloric acid
54
Diagram of Apparatus
Method
1. Set up the apparatus as shown in the diagram.
2. Measure 20 cm3 of 1 mol/dm3 hydrochloric acid using the 25 cm3 measuring cylinder.
Pour the acid into the 250 cm3 conical flask.
3. Fill the other measuring cylinder with water, make sure that it stays filled with water
when you turn it upside down and clamp above the trough.
4. Add a 3 cm strip of magnesium ribbon to the flask, put the bung into the flask and start
the stopwatch.
5. Record the volume of hydrogen gas given off every ten seconds. Continue timing until
no more gas appears to be given off.
6. Repeat steps 2-5 using 10 cm3 of the hydrochloric acid and 10 cm3 of water to make the
total volume used 20 cm3.
Analysis
1. Plot a graph of volume of hydrogen gas (y-axis) against time (x-axis), for both
concentrations of hydrochloric acid and label the lines appropriately.
Risk Assessment
Hazard
Hydrochloric
acid is an
irritant
Risk
Control measure
Hydrochloric acid could get onto the
skin when adding to measuring
cylinder
Hydrochloric acid could get
transferred from the hands to the
eyes
Wash hands immediately if any
hydrochloric acid gets onto them / wear
laboratory gloves
55
Wear eye protection
Teacher / Technician notes
The magnesium ribbon should be clean and free from obvious corrosion or oxidation. Clean
if necessary by rubbing lengths of the ribbon with an emery cloth to remove the layer of
oxide. To ensure that most of the magnesium surface is under the surface of the acid, it
should be folded into a zigzag shape.
The bungs in the flasks need to be rubber. Corks are too porous and will leak. The tube
through the bung should be a short section of glass, and then a flexible rubber tube can be
connected. These can be pre-prepared before the reaction so all the student has to do is
push the bung into the flask.
Gas syringes can be used instead of troughs of water and measuring cylinders. Syringes
should not be allowed to become wet, or the plungers will stick inside the barrels. The
apparatus set up for this procedure is shown in the diagram below:
Reagents:
•
•
Hydrochloric acid – Refer to CLEAPSS hazcard 47A
Magnesium ribbon – Refer to CLEAPSS hazcard 59A
A 3 cm length of magnesium ribbon has a mass of 0.04 g and should yield 40 cm3 of
hydrogen gas when reacted with this excess of acid.
If a graph of volume (y-axis) against time (x-axis) is drawn, the slope of the graph is steepest
at the beginning. This shows that the reaction is fastest at the start. As the magnesium is
used up, the rate falls. This can be seen on the graph, as the slope becomes less steep and
then levels out when the reaction has stopped (when no more gas is produced).
No repeats have been included in the method, but students can compare results with other
groups to make judgements on reproducibility.
56
Practical techniques covered
C1
Use of appropriate apparatus to make and record a range of measurements
accurately, including mass, time, temperature and volume of liquids and gases.
C3
Use of appropriate apparatus and techniques for conducting and monitoring chemical
reactions, including appropriate reagents and/or techniques for the measurement of
pH in different situations.
C5
Making and recording of appropriate observations during chemical reactions
including changes in temperature and the measurement of rates of reaction by a
variety of methods such as production of gas and colour change.
C6
Safe use and careful handling of gases, liquids and solids, including careful mixing of
reagents under controlled conditions, using appropriate apparatus to explore
chemical changes and/or products.
57
Investigation into the effect of one factor on the rate of reaction between dilute
hydrochloric acid and sodium thiosulfate
Introduction
Sodium thiosulfate reacts with hydrochloric acid to form a solid precipitate of sulfur. The
formation of this precipitate makes the solution become cloudy, and so the rate at which this
cloudiness appears can be used as a way to measure the rate of the reaction. The equation
for this reaction is as follows:
sodium
thiosulfate
+ hydrochloric
acid
Na 2 S 2 O 3 (aq) + 2HCl(aq)
sodium
chloride
+ water
+ sulfur + sulfur
dioxide
2NaCl(aq)
+ H 2 O(l) + SO 2 (g) + S(s)
The rate at which this precipitate forms can be changed by changing the conditions under
which the reaction is carried out.
In this experiment you will study the effect of changing the temperature of the sodium
thiosulfate solution.
Apparatus
10 cm3 measuring cylinder
25 cm3 measuring cylinder
250 cm3 conical flask
white paper with cross marked on it
stopwatch
1 mol/dm3 hydrochloric acid
thermometer
Access to:
40 g/dm3 sodium thiosulfate solution at 5 °C
40 g/dm3 sodium thiosulfate solution in a waterbath at 60 °C
58
Diagram of Apparatus
Method
1. Draw a cross on a square of white paper.
2. Measure 25 cm3 of hot sodium thiosulfate using the 25 cm3 measuring cylinder and
pour into the conical flask. Record the temperature of the solution.
3. Using the 10 cm3 measuring cylinder, measure out 5 cm3 of the hydrochloric acid.
4. Place the conical flask onto the cross and add the hydrochloric acid. Swirl the flask to
mix the contents and at the same time start the stopwatch.
5. Look down at the cross from above the mixture.
6. Stop the stopwatch as soon as the cross disappears.
7. Record the time taken for the cross to disappear.
9. Repeat steps 2 to 7 for different temperatures of sodium thiosulfate, made according to
the table below.
Volume of sodium thiosulfate solution
at 60 °C (cm3)
25
20
15
10
5
0
59
Volume of sodium thiosulfate solution
at 5 °C (cm3)
0
5
10
15
20
25
Analysis
1. Plot a graph of the temperature of sodium thiosulfate against the time taken for the
cross to disappear.
Risk Assessment
Hazard
Risk
Hydrochloric acid is an irritant
Damage/irritation to skin
There may be transfer from
the hands to the eyes
causing irritation
Damage/irritation to skin
There may be transfer from
the hands to the eyes
causing irritation
Inhalation of gas may cause
damage/irritation to the lungs
Burns or scalds if the hot
sodium thiosulfate is knocked
over
Sodium thiosulfate is an
irritant
Sulphur dioxide gas
produced is an irritant
Hot water can scald/burn
Control measure
Wash skin immediately if
contact made with
hydrochloric acid
Wear safety goggles
Wash skin immediately if
contact made with
sodium thiosulfate
Wear safety goggles
Carry out in a well ventilated
space
Keep maximum temperature
to 60 oC
Teacher / Technician notes
The crosses on the paper can be pre-prepared and laminated.
An alternative method can also be followed using the method set out on CLEAPSS card
C195. It reduces the volume of reactants used so enabling more sets of equipment to be
created.
Reagents
•
•
Hydrochloric acid – Refer to CLEAPSS hazcard 47A
Sodium thiosulfate – Refer to CLEAPSS hazcard 95A
No repeats have been included in the method, but reproducibility can be checked by
comparing results with other groups. As temperatures will vary across groups, the whole
class data could be plotted onto one graph.
More able candidates could calculate and plot the rate of the reaction using
1
.
time (s)
Students should design their own table, but a suggested table format is shown below.
Time taken for cross to disappear (s)
Recorded temperature (°C)
60
Practical techniques covered
C1
Use of appropriate apparatus to make and record a range of measurements
accurately, including mass, time, temperature and volume of liquids and gases.
C3
Use of appropriate apparatus and techniques for conducting and monitoring chemical
reactions, including appropriate reagents and/or techniques for the measurement of
pH in different situations.
C5
Making and recording of appropriate observations during chemical reactions
including changes in temperature and the measurement of rates of reaction by a
variety of methods such as production of gas and colour change.
C6
Safe use and careful handling of gases, liquids and solids, including careful mixing of
reagents under controlled conditions, using appropriate apparatus to explore
chemical changes and/or products.
61
Investigation into the effect of various catalysts on the decomposition of
hydrogen peroxide
Introduction
Hydrogen peroxide naturally decomposes to release oxygen. The equation for this reaction
is as follows:
hydrogen peroxide
2H 2 O 2 (aq)
water
2H 2 O(l)
+
+
oxygen
O 2 (g)
The rate at which this decomposition occurs is very slow. However the presence of a
catalyst will increase the rate at which the decomposition takes place. In this experiment you
will study how different catalysts affect the rate of decomposition.
The rate at which the oxygen gas is produced can be used to determine the rate of the
reaction. By changing the catalyst used, the rate at which the oxygen gas is produced can
be changed.
Apparatus
250 cm3 conical flask
single-holed rubber bung
delivery tube to fit to conical flask
trough or washing up bowl
2 × 100 cm3 measuring cylinder
stopwatch
clamp stand, boss and clamp
Access to:
10 vol hydrogen peroxide
approximately 0.5 g samples of various catalysts
62
Diagram of Apparatus
Method
1. Set up the apparatus as shown in the diagram.
2. Measure 50 cm3 of hydrogen peroxide using a measuring cylinder and place it in the
conical flask.
3. Fill the other measuring cylinder with water and make sure that it stays filled with water
when you turn it upside down.
4. Connect the bung and delivery tube and place it under the measuring cylinder.
5. Add 0.5 g of a catalyst to the flask, put the bung back into the flask and start the
stopwatch.
6. Record the volume of gas given off every 10 seconds. Continue timing until no more
oxygen appears to be given off.
7. Repeat steps 2-6 for another two catalysts.
Method
1. Compare the results for the three catalysts and reach a conclusion.
Risk Assessment
Hazard
Hydrogen
peroxide is
corrosive
Risk
Hydrogen peroxide could get onto
the skin when adding to test tube
Hydrogen peroxide could get
transferred from the hands to the
eyes
63
Control measure
Wash hands immediately if any
hydrogen peroxide gets onto them /
wear laboratory gloves
Wear eye protection
Teacher / Technician notes
Reagents:
•
•
•
•
•
Hydrogen peroxide – Refer to CLEAPSS hazcard 50
Manganese(IV) oxide – Refer to CLEAPSS hazcard 60
Copper(II) oxide – Refer to CLEAPSS hazcard 26
Zinc oxide – Refer to CLEAPSS hazcard 108B
Iron – Refer to CLEAPSS hazcard 55A
Students should be given the following catalysts to choose from – manganese(IV) oxide, iron
oxide, liver, potato, iron, copper(II) oxide, yeast, zinc oxide.
In this experiment the students measure the volume of oxygen gas produced at 10 second
time intervals. It is important that the concentration of the hydrogen peroxide is no greater
than 10 vol, as this is safe for the students to use and creates a sufficient volume of gas to
be recorded.
It would save time if the masses of each catalyst were pre-weighed. Try to ensure with the
liver and potato options, that they are freshly prepared.
The bungs in the flasks need to be rubber. Corks are too porous and will leak. The tube
through the bung should be a short section of glass, and then a flexible rubber tube can be
connected. These can be pre-prepared before the experiment so all the student has to do is
push the bung into the flask.
No repeats are expected for this experiment and students can compare reproducibility by
comparing their results with those of other groups.
Gas syringes can be used instead of troughs of water and measuring cylinders.
Syringes should not be allowed to become wet, or the plungers will stick inside the barrels.
The apparatus set up for this procedure is shown in the diagram below.
64
Once the data from each experiment has been collected the students can construct a graph
of volume of oxygen gas (y-axis) against time in seconds (x-axis) for each of the catalysts on
the same axes.
Practical techniques covered
C1
Use of appropriate apparatus to make and record a range of measurements
accurately, including mass, time, temperature and volume of liquids and gases.
C3
Use of appropriate apparatus and techniques for conducting and monitoring chemical
reactions, including appropriate reagents and/or techniques for the measurement of
pH in different situations.
C5
Making and recording of appropriate observations during chemical reactions
including changes in temperature and the measurement of rates of reaction by a
variety of methods such as production of gas and colour change.
C6
Safe use and careful handling of gases, liquids and solids, including careful mixing of
reagents under controlled conditions, using appropriate apparatus to explore
chemical changes and/or products.
65
10 – CARBON COMPOUNDS
Spec Statement
Comment
(a)
recall that crude oil is a main
source of hydrocarbons and is a
feedstock for the petrochemical
industry
Candidates should know that crude oil is a complex
mixture of hydrocarbons that was formed over millions of
years from the remains of simple marine organisms.
(b)
describe and explain the
separation of crude oil by
fractional distillation
Candidates should know that crude oil is boiled/vaporised
before it enters the fractionating column and that the
hydrocarbons present condense at different heights in the
column. The lower the boiling point, the higher in the
column a compound is collected. They should know that
fractions are mixtures containing hydrocarbon compounds
that have similar boiling points and that these have similar
chain lengths. They are not expected to recall details such
as the range of chain lengths present in the constituent
hydrocarbons of different fractions but they should know
the uses of the following fractions: petroleum gases;
gasoline/petrol; naphtha; kerosene; diesel; lubricating oil;
fuel oil; bitumen.
Candidates should know that the longer the chain lengths
of the hydrocarbons present in a fraction, the higher its
boiling point range. They should also recall the effect of
increasing chain length on different fractions’ colour
(colourless – yellow – brown), viscosity, ease of ignition
and cleanliness of burn.
(c)
describe the fractions as largely
a mixture of compounds of
general formula C n H 2n+2 which
are members of the alkane
homologous series
(d)
describe the production of
materials that are more useful by
cracking
The cracking process involves heating fractions obtained
from crude oil to a high temperature in the presence of a
catalyst. This causes the hydrocarbon molecules present
to decompose forming smaller molecules, including an
alkene. There is greater demand for the smaller
hydrocarbons, and alkenes such as ethene, are the
starting material for the production of many plastics.
(e)
explain how modern life is
crucially dependent upon
hydrocarbons and recognise that
crude oil is a finite resource
Candidates should appreciate that hydrocarbon fuels are
vital for travel and electricity generation and that alkenes
produced by cracking are the basis for production of most
plastics. Without crude oil our lifestyles would be
unrecognisable. As crude oil reserves are used up,
governments will eventually have to decide between
burning the remaining oil and using it for other purposes.
66
(f)
recognise functional groups and
identify members of the same
homologous series
(g)
name and draw the structural
formulae, using fully displayed
formulae, of the first four
members of the straight chain
alkanes, alkenes, alcohols and
carboxylic acids
These should include the following compounds:
but-1-ene and but-2-ene
propan-1-ol, propan-2-ol, butan-1-ol and butan-2-ol
(h)
predict the formulae and
structures of products of
reactions (combustion, addition
across a double bond and
oxidation of alcohols to
carboxylic acids) of the first four
and other given members of
these homologous series
Combustion – alkanes, alkenes and alcohols
Addition across double bond – hydrogen and bromine; not
H—Br
Oxidation – 1° alcohols only
(i)
recall that it is the generality of
reactions of functional groups
that determine the reactions of
organic compounds
(j)
recall the basic principles of
addition polymerisation by
reference to the functional group
in the monomer and the
repeating units in the polymer
(k)
deduce the structure of an
addition polymer from a simple
alkene monomer and vice versa
(l)
explain the basic principles of
condensation polymerisation
by reference to the functional
groups of the monomers, the
minimum number of functional
groups within a monomer, the
number of repeating units in
the polymer, and simultaneous
formation of a small molecule
Candidates should know that in addition polymerisation
reactions many small unsaturated molecules (monomers)
join together to form very large saturated molecules
(polymers). They should understand that the reactivity of
the monomer arises from the presence of its double bond,
and that as polymerisation happens one of the bonds
breaks to allow the molecule to join to another. They
should be able to draw the structural formulae of the
ethene, propene, vinylchloride and tetrafluoroethene
monomers and describe the reactions forming their
respective polymers in the form of an equation using ‘n’
monomer molecules.
Candidates should explain that condensation
polymerisation involves monomers with two functional
groups. (These are usually two different monomers
but a condensation polymer can form from one
monomer with two different functional groups.) When
these monomers react they join together forming a
small molecule such as H 2 O. Polyesters and
polyamides are examples of condensation polymers.
Recall of the names and structures of reacting
monomers is not required.
67
(m)
recall that DNA is a polymer
made from four different
monomers called nucleotides
and that other important
naturally-occurring polymers are
based on sugars and amino
acids
Candidates should know that DNA is a polymer. The
monomer units of DNA are nucleotides, and the polymer is
known as a 'polynucleotide'. There are four different types
of nucleotides found in DNA. They are given one letter
abbreviations – A, G, C and T. DNA is a double-stranded
macromolecule where two polynucleotide chains are held
together by weak forces.
Other naturally-occurring polymers include:
• starch – polymers made of sugars
• proteins – polymers made of amino acids
68
11 – PRODUCTION, USE AND DISPOSAL OF IMPORTANT CHEMICALS AND
MATERIALS
Spec Statement
(a)
explain the importance of the
Haber process in agricultural
production
Comment
Candidates should understand that a great proportion of
the ammonia made world-wide is used to produce
nitrogenous fertilisers in order to maximise crop growth and
availability of food for the increasing global population.
Candidates should know that nitrogen gas for the Haber
process is obtained from the air and that hydrogen is
usually made from methane. They should recall the basic
reaction conditions used in the process – temperature in
the range 350-450°C, pressure in the range 150-200 atm
and iron catalyst. They should be able to write word and
balanced symbol equations for the reaction taking place.
(b)
recall the importance of nitrogen,
phosphorus and potassium
compounds in agricultural
production and the potential
drawbacks of over use
Nitrogen – to make proteins for strong stems and healthy
leaves
Phosphorus – helps roots grow and fruit ripen
Potassium – helps protect against disease and frost
damage
Candidates should recall the problems associated with
fertilisers being washed into waterways and be able to give
a basic explanation of eutrophication.
(c)
describe the industrial production
of fertilisers as several integrated
processes using a variety of raw
materials
Ammonium nitrate and ammonium sulfate are both used in
fertilisers. They are made by neutralising ammonia with
nitric acid and sulfuric acid respectively. Detailed recall of
these processes is not required. Fertilisers for specific
uses are formulated by mixing different amounts of the
various compounds required to attain the appropriate NPK
values.
(d)
compare the industrial production Candidates should be able to describe how crystals of
of fertilisers with laboratory
ammonium nitrate and ammonium sulfate can be made in
syntheses of the same products
the laboratory. Detailed recall of the industrial production
processes is not required.
(e)
describe tests to identify
ammonia gas and ammonium
salts
When damp red litmus paper is placed into a jar/tube
containing ammonia gas, it turns blue. Candidates are not
required to know the hydrogen chloride gas test. (Please
note that the pungent smell of ammonia gas is not a
chemical test.)
Ammonium salts can be identified by adding sodium
hydroxide solution and warming gently to release ammonia
gas (which turns damp red litmus paper blue).
69
(f)
interpret graphs of reaction
conditions versus rate
(g)
explain the trade-off between
rate of production of a desired
product and position of
equilibrium in some
industrially important
processes
(h)
explain how the commercially
used conditions for an
industrial process are related
to the availability and cost of
raw materials and energy
supplies, control of
equilibrium position and rate
(i)
describe the social and
environmental impact of
decisions made in siting
chemical plants
(j)
calculate the percentage yield of
a reaction product from the
actual yield of a reaction
This would involve a simple percentage calculation where
theoretical yield is also given.
(k)
calculate the theoretical yield of
a product from a given amount of
reactant
This is more challenging and involves using the balanced
symbol equation to calculate the theoretical mass of
product formed.
(l)
define the atom economy of a
reaction
Candidates should know that the atom economy is a
measure of the amount of starting materials that end up as
useful products. A higher atom economy is likely to favour
sustainability and lower costs.
(m)
calculate the atom economy of a
reaction to form a desired
product from the balanced
equation
Candidates are required to recall an expression used to
calculate atom economy.
(n)
explain why a particular
reaction pathway is chosen to
produce a specified product
given appropriate data such as
atom economy (if not
calculated), yield, rate,
equilibrium position and
usefulness of by-products
No detailed recall is required here. Questions will
involve weighing up the various factors listed in order
to evaluate the sustainability of different reaction
pathways.
This requires an understanding of reaction yield as
well as reaction rate. In reversible reactions it may be
better to use conditions that give a low yield at a good
rate than conditions which give a high yield but at a
low rate. Candidates should consider these ideas in
the context of the conditions used in the Haber
process. Explanations should include application of
Le Chatelier's principle.
70
(o)
describe the conditions which
cause corrosion and the process
of corrosion, and explain how
mitigation is achieved by creating
a physical barrier to oxygen and
water and by sacrificial protection
Candidates should recall that corrosion is the breaking
down of a material, especially a metal, by chemical
reactions with substances in the environment. They should
know that the rusting of iron is an example of corrosion and
that both air and water are necessary for iron to rust. They
should be able to explain corrosion prevention by applying
a barrier (e.g. grease, paint, electroplating) and by
sacrificial protection (coating with a more reactive metal
such as zinc to galvanise iron, magnesium blocks attached
to steel ships).
(p)
describe the composition of
some important alloys in relation
to their properties and uses
Candidates should recall the compositions, properties and
uses of stainless steel, duralumin, nitinol and brass.
(q)
compare quantitatively the
physical properties of glass and
clay ceramics, polymers,
composites and metals
Detailed recall is not required here.
(r)
explain how the properties of
materials are related to their
uses and select appropriate
materials given details of the
usage required
(s)
describe the basic principles in
carrying out a life-cycle
assessment of a material or
product
Candidates should know that a life-cycle assessment
(LCA) is carried out to assess the environmental impact of
products at each of the following stages:
• extraction and processing raw materials
• manufacture and packaging
• use, repair and maintenance during its lifetime
• disposal or recycling at the end of its useful life
(t)
interpret data from a life-cycle
assessment of a material or
product
Candidates should be able to carry out simple LCA
comparisons using numerical data (e.g. energy, water
consumption, waste) for a given material or product, e.g.
paper versus plastic shopping bags.
(u)
describe a process where a
material or product is recycled for
a different use, and explain why
this is viable
Aluminium, steel and glass can be recycled indefinitely
without losing quality so it is possible, for example, for
aluminium from drinks cans to be made into aircraft parts
and bicycle frames. Glass not suitable for re-melting can
be used as aggregate in concrete or ground down to make
building sand. Plastic bottles can be melted down and
made into any number of products, from garden furniture to
fleece jackets.
71
(v)
evaluate factors that affect
decisions on recycling
A number of factors have to be considered before the
recycling of a material becomes commonplace. Economic
viability is probably the most important one. No private
company will recycle a product if the costs are greater than
the return. This explains why some types of plastic are not
recycled. On the other hand, recycling waste plastic
reduces the amount of waste either ending up in landfill or
being burned. Equally important are the benefits of
conserving crude oil reserves and, because recycling uses
less energy than production, reducing fossil fuel use.
Recycling metals conserves the raw materials and uses
much less energy. Recycling aluminium requires
approximately 5% of the energy used to extract the metal
from bauxite. The reduction in energy for recycling means
that less electricity is needed and so there are smaller
associated greenhouse emissions.
SPECIFIED PRACTICAL WORK
•
SP11 Determination of the percentage of water in a hydrated salt, e.g. copper(II) sulfate
72
Determination of the percentage of water in a hydrated salt, e.g. copper (II)
sulfate
Introduction
Copper(II) sulfate bonds to molecules of water when crystallised. This gives a recognisable
blue colour. Upon strong heating, the water is evaporated leaving white, anhydrous
copper(II) sulfate. By carrying out this experiment, you will be able to calculate the
percentage of water in copper(II) sulfate crystals.
Apparatus
crucible
Bunsen burner
copper(II) sulfate (CuSO 4 .xH 2 O)
spatula
Access to:
electronic balance ± 0.01 g
Diagram of Apparatus
73
Method
1.
2.
3.
4.
Record the mass of an empty crucible.
Add 6 spatulas of copper(II) sulfate, record the new mass.
Heat on a tripod with a roaring flame for several minutes.
Once the solid has completely changed from blue to white allow to cool then record the
final mass of anhydrous copper(II) sulfate.
Analysis
1. Calculate the mass of hydrated copper(II) sulfate used.
2. Calculate the mass of water lost.
3. Calculate the percentage of water in this copper(II) sulfate.
Risk Assessment
Hazard
Copper(II) sulfate
is harmful
Risk
Risk of splashing into eyes whilst
stirring
Control measure
Take care whilst stirring and wear eye
protection
Hot apparatus can
burn
Burning hands when moving hot
equipment
Allow sufficient cooling time before
removing crucible from tripod
Grip tripod at the base at all times
Teacher / Technician notes
Copper(II) sulfate ‒ Refer to CLEAPSS hazcard 27C
Six spatulas of copper(II) sulfate are required to detect the loss of mass – roughly 30-40 %.
It is worth demonstrating at the end of the practical that the procedure can be reversed by
adding deionised water to the anhydrous copper(II) sulfate. The dehydrated copper(II)
sulfate can be re-used by hydrating it in this way - CARE - this is very exothermic.
More able students could be asked to calculate the number of water molecules per molecule
of copper(II) sulfate.
Students should design their own table, but a suggested table format is shown below.
Mass of crucible (g)
Mass of crucible + hydrated copper(II) sulfate (g)
Mass of crucible + anhydrous copper(II) sulfate (g)
Mass of hydrated copper(II) sulfate (g)
Mass of water lost (g)
Percentage of water
74
Practical techniques covered
C1
Use of appropriate apparatus to make and record a range of measurements
accurately, including mass, time, temperature and volume of liquids and gases.
C2
Safe use of appropriate heating devices and techniques including use of a Bunsen
burner and a water bath or electric heater.
C6
Safe use and careful handling of gases, liquids and solids, including careful mixing of
reagents under controlled conditions, using appropriate apparatus to explore
chemical changes and/or products.
75
12 – THE EARTH AND ITS ATMOSPHERE
Spec Statement
Comment
(a)
interpret evidence for how it is
thought the atmosphere was
originally formed
Candidates should know that several theories have been
suggested to account for the formation of the Earth’s early
atmosphere, but many scientists agree that it is most likely
to have formed from gases expelled by volcanoes. Carbon
dioxide, water vapour and ammonia make up the greatest
proportion of volcanic gases.
(b)
describe how it is thought that
the an oxygen-rich atmosphere
developed over geological time
Candidates should know that the surface of the Earth cooled
over time and that water vapour present in the early
atmosphere condensed forming the oceans. They should
appreciate that this happened quickly, in geological terms,
and that other changes took far longer. The percentage of
carbon dioxide has decreased to a fraction of one percent as
a result of a number of processes, the most important being
photosynthesis. Photosynthesis began as green plants
evolved, using up carbon dioxide and releasing oxygen into
the atmosphere for the first time. The evolution of marine
animals followed over hundreds of millions of years and
much carbon dioxide was locked into limestone and chalk
formed from their shells. More still was locked into fossil
fuels formed many millions of years ago from the remains of
simple marine organisms (crude oil and natural gas) and
larger land plants (coal). Ammonia decomposed on reaction
with oxygen forming nitrogen, which became the most
abundant gas in the atmosphere. These changes occurred
over billions of years.
(c)
recall the approximate
composition of the present day
atmosphere
nitrogen 78%
oxygen 21%
argon (+ other noble gases) 0.9%
carbon dioxide 0.04%
(d)
describe the greenhouse effect
in terms of the interaction of
radiation with the Earth’s
atmosphere
Greenhouse gases (water vapour, carbon dioxide and
methane) in the atmosphere maintain temperatures on Earth
within a range that supports life. They allow short
wavelength radiation to pass through the atmosphere to the
Earth’s surface but absorb the outgoing long wavelength
radiation coming from the Earth's surface causing an
increase in atmospheric temperature. The greenhouse
effect is an entirely natural phenomenon.
(e)
explain global warming in terms
of an ‘enhanced greenhouse
effect’
Global warming is the result of an ‘enhanced greenhouse
effect’. It is the impact on the climate of the additional heat
retained due to the ever-increasing amounts of carbon
dioxide (and other greenhouse gases) released into the
Earth’s atmosphere since the industrial revolution.
(f)
evaluate the evidence for manmade causes of climate change,
Candidates should appreciate that the overall trend over the
past 200 years shows a correlation between carbon dioxide
76
(g)
including the correlation between
change in atmospheric carbon
dioxide concentration and the
consumption of fossil fuels, and
describe the uncertainties in the
evidence base
concentration and atmospheric temperature, although there
are anomalies over short time periods. They should be
aware that sunspot activity has been suggested as an
alternative cause for the observed temperature changes and
that this has been ruled out by the majority of scientists on
the basis of the available evidence.
describe the potential effects of
increased levels of carbon
dioxide and methane on the
Earth's climate and how these
may be mitigated, including
consideration of scale, risk and
environmental implications
Increased levels of carbon dioxide could cause:
• Climate change e.g. hotter summers in some parts of the
world (droughts) and increased rainfall (flooding) in
others
• Higher rate of melting of ice caps, polar sea ice, glaciers
• Rising sea levels
• Changes in food production capacity of some regions
• Impact on wildlife
Increases in carbon dioxide levels can be slowed and
eventually reversed by:
• Increasing use of alternative energy sources
• Improved energy conservation
• Carbon capture and storage
• Carbon off-setting e.g. tree planting
(h)
describe the major sources of
carbon monoxide, sulfur dioxide,
oxides of nitrogen and
particulates in the atmosphere
and explain the problems caused
by increased amounts of these
substances
Candidates should recall the main sources and effects of
atmospheric pollution.
• Carbon monoxide – incomplete burning of fossil fuels;
poisonous
• Sulfur dioxide – burning fuels containing sulfur; affects
the respiratory system, causes acid rain (see below)
• Nitrogen oxides – burning fossil fuels; affects the
respiratory system, causes acid rain
• Particulates (smoke and soot) – burning fossil fuels;
affects the respiratory system, reduces visibility,
damages the appearance of buildings
They should know that fossil fuels contain many impurities
including many sulfur-containing compounds and that these
produce sulfur dioxide on burning. Sulfur dioxide forms a
solution of sulfuric acid on contact with water in the
atmosphere and this falls as acid rain. Candidates should
know that ‘clean’ rain is weakly acidic (pH ~5.5) and that
acid rain has a pH in the range of about 2-4. They should
know that acid rain lowers the pH of lakes etc., damaging
aquatic life and damaging forests and vegetation. They
should know that it damages buildings (particularly those
made of limestone [calcium carbonate]) and increases the
rate of corrosion of metal structures such as bridges and
statues.
77
(i)
describe the principal methods
for increasing the availability of
potable water in terms of the
separation techniques used,
including ease of treatment of
waste water, ground water and
salt water
Candidates should know potable water is water that is safe
to drink, rather than 'pure' water.
Candidates should know the stages in the treatment of the
public water supply using sedimentation, filtration and
chlorination.
• Sedimentation – in reservoirs/tanks, larger solid particles
settle under gravity
• Filtration – through layers of sand and gravel, removes
smaller insoluble particles
• Chlorination – chlorine added to kill bacteria, prevents
disease/makes it safe to drink
Candidates should know that the simplest method for
desalination of sea water is distillation. This involves boiling
sea water which uses large amounts of costly energy,
preventing it from being a viable process in many parts of
the world. Candidates should be aware that other methods
are also used, e.g. the use of membrane systems, but they
are not required to know any details of such methods. They
should be able to discuss the potential of desalination as a
source of drinking water in different parts of the world in
terms of proximity to the sea, availability of ‘cheap’ energy
and a country’s wealth.
78