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AQA Additional Science – Unit C2
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Written by Peter Hill, BSc.
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GCSE Science Electronic Revision Guides
AQA Additional Science – Unit C2
Contents (click on a title)
Copyright notice – conditions of use
Topic 1: Atoms, Elements and Compounds
Topic 1: Ionic Bonding
Topic 1: Ionic Compounds
Topic 1: Covalent Bonding
Topic 2: Covalent Substances
Topic 2 Properties of Metals
Topic 2: Different Types of Polymer
C2 Topic 2: Nanoscience
Topics 1 & 2 Questions
Topics 1 & 2 Answers
Topic 3: Relative Atomic Mass and Formula Mass
Topic 3: Analysing Substances
Topic 3: Relative Formula Mass Calculations
Topic 3: Calculating Masses from Symbol Equations (H)
Topic 3: Percentage Yield
Topic 4: Measuring the Rate of Reactions
Topic 4: The Collision Theory
Topic 4: Catalysts
Topics 3 & 4 Questions
Topics 3 & 4 Answers
Topic 5: Energy Changes in Chemical Reactions
Topic 6: Acids and Bases
Topic 6: Acid Reactions and Salts
Topic 6: Making Salts
Topic 7: Electrolysis
Topic 7: Electrolysis of Aluminium Oxide and Sodium Chloride
Topics 5-7 Questions
Topics 5-7 Answers
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C2 Topic 1: Atoms, Elements and Compounds
Revision of atoms
All substances are made of atoms. A substance that is made of only one sort of atom is called
an element. Atoms are the smallest particles of an element that can take part in chemical
reactions. There are about 100 different elements.
Atoms have a very small central nucleus, which is made up of protons and neutrons and around
which there are electrons in shells (energy levels). The relative masses and electrical charges
are shown in the table below:
Name of particle
Relative mass
Relative charge
Proton
1
+1
Neutron
1
0
Electron
0.0005
-1
The structures of a hydrogen and carbon atom are shown as examples below:
CARBON ATOM
The outer shell contains
four electrons.
The inner shell contains
two electrons.
HYDROGEN ATOM
one electron
The nucleus consists
of one proton.
The nucleus consists
of six protons and
six neutrons.
= positively charged proton
= neutron (no charge)
= negatively charged electron
The atomic number is the number of protons that an atom contains. All atoms of a particular
element have the same number of protons. Atoms of different elements have different
numbers of protons.
Compounds
Compounds are formed when the atoms of two or more elements chemically combine. Chemical
bonds form between the atoms. The properties of the compound formed are completely
different to the elements it is composed of. In the reaction between hydrogen and oxygen for
example two hydrogen molecules react with one oxygen molecule to form the compound water:
H
H
H
H
H
H
+
O
O
O
O
H
H
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C2 Topic 1: Ionic Bonding
When elements react, their atoms join with other atoms to form compounds. This involves
giving, taking or sharing electrons in order to gain full outer shells. This gives the most stable
arrangement of electrons.
Ionic bonding
Compounds formed from metals and non-metals consist of ions. Metal atoms lose electrons to
form positively charged ions. Non-metal atoms gain electrons to form negatively charged ions.
Ions have the electronic structure of a noble gas i.e. they have full outer shells. Oppositely
charged ions are strongly attracted to each other and are held together by ionic bonds. The
diagram below shows the formation of a simple ionic compound sodium chloride:
Cl
Na
Na
A sodium atom gives its outer
electron to a chlorine atom.
Cl
Oppositely charged ions are formed
which then bond together strongly.
The group an element belongs to in the periodic table tells us the number of electrons in the
outer shell of its atoms. Elements in groups 1 and 2 of the periodic table only have 1 or 2
electrons in their outer shells so these form positive ions by losing their outer electrons.
Elements in groups 6 and 7 of the periodic table only need 1 or 2 electrons to fill up their outer
shells so these form negative ions by gaining extra electrons.
Structure and properties of ionic compounds
An ionic compound contains many ions held together by strong electrostatic forces of
attraction between oppositely charged ions. These forces act in all directions. Ionic compounds
form crystals because the ions are packed in a regular arrangement called a giant ionic lattice.
The diagram below shows the arrangement of sodium and chloride ions in a crystal of salt.
Ionic compounds show the following properties:

High melting and boiling points due to the strong forces
of attraction holding the ions together. It takes large
amounts of heat energy to break these bonds.

= sodium ion
= chloride ion
When melted or dissolved in water, ionic compounds
conduct electricity because the ions are free to move
and carry the current. This provides evidence that
these compounds are actually made up of ions.
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CONTENTS PAGE
C2 Topic 1: Ionic Compounds
The elements in Group 1 of the periodic table (the alkali metals) all react with non-metal
elements to form ionic compounds in which the metal ion has a single positive charge. This is
because they only have one electron in their outer shells to lose. Group 2 metals have two
electrons in their outer shells which they lose to form ions with a double positive charge.
The elements in Group 7 of the periodic table (the halogens) all react with the alkali metals to
form ionic compounds in which the halide ion has a single negative charge. This is because they
have seven electrons in their outer shells so they only need to gain one more. Group 6 elements
have six electrons in their outer shells so they gain two more to form ions with a double
negative charge.
Working out the formulae of ionic compounds
When metals react with non-metals the ratio of each ion in the compound formed depends on
the amount of charge each ion has. All of the negative charges in a compound must balance all
of the positive charges:
Example 1 – Sodium chloride
The sodium ion has a charge of +1 and the chloride ion has a charge of -1. Therefore sodium
chloride has a 1:1 ratio of sodium to chloride ions and the formula is NaCl:
Na
Cl
sodium atom
chlorine atom
Cl
Na
sodium chloride (NaCl)
Example 2 – Magnesium chloride
The magnesium ion has a charge of +2 and the chloride ion has a charge of -1. Therefore
sodium chloride has a 1:2 ratio of magnesium to chloride ions and the formula is MgCl2:
2
Cl
Mg
magnesium atom
Cl
Cl
Mg
Cl
magnesium chloride (MgCl2)
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CONTENTS PAGE
C2 Topic 1: Covalent Bonding
In covalent bonding non-metal atoms share pairs of electrons with each other so that they
gain full outer shells. Covalent bonds are strong. The diagrams below show covalent bonding in
molecules of elements. Only the outer shells of electrons need to be drawn. Three different
ways of representing each molecule are shown:
Hydrogen and chlorine atoms both need one
more electron to fill up their outer shells.
chlorine molecule (Cl2)
hydrogen molecule (H2)
H
H
H
H
H
H
Oxygen atoms need two more electrons to
fill up their outer shells so a double bond
containing two pairs of electrons is formed.
oxygen molecule (O2)
Cl
Cl
Cl
Cl
Cl
Cl
O
O
O
O
Compounds formed from non-metals only consist of molecules. The atoms are held together by
covalent bonds. The diagrams below show covalent bonding in simple molecules of compounds.
Hydrogen chloride forms when one hydrogen
atom bonds with one chlorine atom.
Water forms when two hydrogen
atoms bond with one oxygen atom.
shells.
water molecule (H2O)
hydrogen chloride molecule (HCl)
O
Cl
H
O
Cl
H
H
H
Cl
H
H
O
H
H
Ammonia forms when three hydrogen
atoms bond with one nitrogen atom.
ammonia molecule (NH3)
H
Methane forms when four hydrogen
atoms bond with one carbon atom.
methane molecule (CH4)
H
H
H
N
H
N
H
H
H
H
H
H
H
N
H
C
C
H
H
H
H
H
H
H
C
H
H
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CONTENTS PAGE
C2 Topic 2: Covalent Substances
Simple molecules
Substances that consist of simple molecules are gases, liquids or solids that have relatively low
melting and boiling points. They do not conduct electricity because the molecules do not have
an overall electric charge. Some examples of simple molecules are shown below:
O
O
Cl
Cl
O
C
O
O
H
Oxygen, chlorine and carbon dioxide are all gases at room temperature.
H
Water is a liquid at room temperature.
Higher Paper – Substances that consist of simple molecules have only weak forces of
attraction between the molecules (intermolecular forces). It is these forces that are
overcome, not the much stronger covalent bonds, when the substance melts or boils.
Giant structures
Atoms joined by covalent bonds can also form giant structures or macromolecules. Diamond and
graphite (forms of carbon) and silicon dioxide (silica) are examples of giant covalent structures
(lattices) of atoms. All of the atoms in these structures are linked to other atoms by strong
covalent bonds and so they have very high melting points.
In diamond each carbon atom forms four
covalent bonds with other carbon atoms. The
atoms are held together in a rigid, giant
covalent structure, so diamond is very hard.
In graphite each carbon atom bonds to
three others, forming layers. The layers
are free to slide over each other because
there are no covalent bonds between the
layers and so graphite is soft and slippery.
Explaining the properties of graphite (H)
There are weak intermolecular forces between the layers of graphite. These forces are so
weak that the layers are only held together loosely and they can slide over each other easily.
This makes graphite soft and slippery and explains how the graphite in a pencil can be rubbed
off onto paper.
In graphite one electron from each carbon atom is delocalised therefore graphite is similar to
metals. These delocalised electrons allow graphite to conduct heat and electricity. Graphite is
the only non-metal that has these properties.
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C2 Topic 2 Properties of Metals
Metals consist of giant structures of atoms arranged in a regular pattern. Metals can be bent
and shaped because the layers of atoms are able to slide over each other. Other elements can
be added to pure metals to form alloys. This makes alloys harder than pure metals. A model can
be used to explain how this happens:
FORCE
Pure metals have regular
arrangements of atoms in layers
that slide over each other easily if
forces are applied to them.
Different elements have different
sized atoms. Adding other
elements to the metal disrupts the
layers making it more difficult for
them to slide over each other.
Alloys are usually made from two or more different metals. Alloying metals changes their
properties and results in new materials which are more suited to their different uses. Most
metals in everyday use are alloys e.g. pure gold is too soft to make jewellery but it can be
hardened by adding metals such as zinc, copper and silver. Tin can is added to copper to form a
much harder alloy called bronze. Various forms of steel are made by adding small amounts of
carbon, chromium or nickel to pure iron.
Shape memory alloys
Nitinol is an alloy of nickel and titanium. It is an example of a shape memory alloy because if an
object made of nitinol is bent it returns to its original shape when heated. Nitinol is used in dental
braces. It can also be used in spectacle frames because if they become bent by e.g. being sat on
they can easily be reshaped by heating them.
Delocalised electrons conduct heat and electricity (H)
The electrons in the outer shells of metal atoms are delocalised (loose) and so free to move
through the whole structure. There is a structure of positive ions with electrons between the
ions holding them together by strong electrostatic attractions. This is what makes metals very
good conductors of heat and electricity.
= positively charged metal ion
= delocalised negatively charged electron
= direction of an electric current
through the metal
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C2 Topic 2: Different Types of Polymer
It is possible to produce a wide range of different polymers with properties that make each of
them suited to a particular use. The properties of polymers depend on what they are made
from and the conditions under which they are made. For example, low density (LD) and high
density (HD) polythene are produced using different catalysts and reaction conditions.
HD polythene has chains
lying close together.
LD polythene has side branches that
keep the chains further apart.
Thermosoftening polymers soften when heated. When they cool they harden into the new
shape. They can be reheated and remoulded over and over again. Thermosoftening polymers
consist of individual, tangled polymer chains.
Thermosetting polymers do not soften when heated and cannot be reshaped. They consist of
polymer chains with cross-links between them. Cross-links are chemical bonds holding the
polymer chains together.
Thermosoftening polymer
Thermosetting polymer
Explaining the properties of thermosoftening and thermosetting polymers (H)
When thermosoftening polymers are heated the tangled polymer chains can uncoil and slide
past each other, making the polymer flexible.
In thermosetting polymers the cross-links between the polymer chains stop them sliding past
each other. The greater the strength of the forces (cross-links) the more energy is needed to
separate them and therefore the greater the strength and the higher the melting point of the
polymer.
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C2 Topic 2: Nanoscience
Nanoscience is becoming increasingly important in today’s world. It involves the use and control
of structures called nanoparticles that are very small (1 - 100 nanometres in size). A nanometre
(nm) is one billionth of a metre (0.000 000 001m).
Nanoparticles can occur naturally, for example in sea spray and volcanic ash. Humans can also
make them by accident, for example the combustion of fuels can produce very tiny particulates
(soot particles). Many nanoparticles are made by design using nanotechnology.
Nanoparticles of a material show different properties, compared to larger particles of the
same material. One of the reasons for this is the much larger surface area of the
nanoparticles compared to their volume i.e. they have a larger surface area to volume ratio.
Chemical reactions can happen more quickly if there is more exposed surface area.
Some uses of nanoparticles
Nanoparticles have the following uses:
Sunblock – Nanoparticles are being used in some sunblocks. They are
too small to reflect light but they block harmful UV rays. This makes
the sunblock transparent instead of the usual white colour.
Deodorants – Nanoparticles have also been used to make deodorants
that do not leave white marks on skin and clothes.
Sports equipment - Nanoparticles are added to plastics used for
sports equipment such as tennis rackets and golf clubs. This makes
the plastic stronger and more durable without increasing the weight.
Nanoscience may also lead to the development of:





new computers
new catalysts
new coatings
highly selective sensors that only detect a certain chemical
stronger and lighter construction materials.
Fullerenes (H)
Carbon can exist in a form called fullerenes. These are molecules shaped
like hollow balls or tubes (nanotubes). The carbon atoms are arranged in
hexagonal rings as they are in graphite. Different fullerenes contain
different numbers of carbon atoms.
Fullerene molecules
Fullerenes may be used in the future to deliver medicinal drugs into the
cells of the body. New lubricants and catalysts are also being developed
using fullerenes. Nanotubes are very strong and they conduct electricity.
They could be used for strengthening materials and in computer chips.
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C2 Topics 1 & 2 Questions
1. Describe the structure of an atom.
2. What does the atomic number of an element tell you?
3. What does the mass number of an element tell you?
4. What gives elements in the same group of the periodic table similar chemical properties?
5. Which group of elements in the periodic table have full outer shells of electrons?
6. Why do elements react to form compounds?
7. What type of compounds form when metals react with non-metals?
8. What happens to electrons when ionic bonds form?
9. What happens to electrons when covalent bonds form?
10. Describe the structure of an ionic compound such as sodium chloride (how the ions are
arranged).
11. Describe and explain two properties of ionic compounds.
12. If an aluminium ion is Al3+ and an oxide ion is O2- work out the formula of aluminium oxide.
13. Describe the physical properties of substances which consist of simple molecules.
14. Name two substances made of giant covalent structures.
15. Explain why graphite conducts electricity. (Higher paper)
16. How are shape memory alloys useful?
17. How do the properties of thermosoftening and thermosetting polymers differ?
18. Give two uses of nanoparticles.
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C2 Topics 1 & 2 Answers
1. Atoms have a small central nucleus, which is made up of protons and neutrons and around which there
are shells of electrons.
2. The atomic number of an element tells you the number of protons in the nucleus.
3. The mass number of an element tells you the sum of the protons and neutrons in the nucleus.
4. Elements in the same group of the periodic table have similar chemical properties because they have
the same number of electrons in their highest energy level (outer shell).
5. Group 0 of the periodic table have full outer shells of electrons.
6. When elements react, their atoms join with other atoms to form compounds. This involves giving,
taking or sharing electrons in order to gain full outer shells. This gives the most stable arrangement of
electrons.
7. Ionic compounds form when metals react with non-metals.
8. When ionic bonds form metal atoms lose electrons to form positively charged ions. Non-metal atoms
gain electrons to form negatively charged ions.
9. When covalent bonds form the atoms in the molecule share pairs of electrons.
10. An ionic compound contains many ions held together by strong electrostatic forces of attraction
between oppositely charged ions. These forces act in all directions. Ionic compounds form crystals
because the ions are packed in a regular arrangement called a giant ionic lattice.
11. Two properties of ionic compounds are:
- High melting and boiling points due to the strong forces of attraction holding the ions together. It
takes large amounts of heat energy to break these bonds.
- When melted or dissolved in water, ionic compounds conduct electricity because the ions are free to
move and carry the current.
12. Al2O3 because the positive and negative charges must balance. Two aluminium ions have a total
charge of 6+ and three oxygen ions have a total charge of 6-.
13. Substances which consist of simple molecules are gases, liquids or solids that have relatively low
melting and boiling points. They do not conduct electricity because the molecules have no overall electric
charge.
14. Diamond, graphite and silicon dioxide are substances made of giant covalent structures.
15. In graphite one electron from each carbon atom is delocalised therefore graphite is similar to
metals. These delocalised electrons allow graphite to conduct heat and electricity. (Higher paper)
16. If an object made from a shape memory alloy is bent it returns to its original shape when heated.
17. Thermosoftening polymers soften when heated. When they cool they harden into the new shape.
Thermosetting polymers do not soften when heated and cannot be reshaped.
18. Nanoparticles are used to make transparent sunblocks, deodorants that do not leave marks, and for
making plastics stronger in sports equipment.
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C2 Topic 3: Relative Atomic Mass and Formula Mass
The periodic table shows the atomic number and mass number for every element. The number
of protons in an atom of an element is its atomic number. The sum of the protons and neutrons
in an atom is its mass number e.g. for the element lithium:
The mass number is 7 because lithium
has 3 protons plus 4 neutrons.
7
Li
Lithium
The atomic number is 3 because
lithium has 3 protons.
3
To find the number of neutrons in
an atom subtract the atomic number
from the mass number.
Atoms of the same element can have different numbers of neutrons and are called isotopes of
that element. The different isotopes of an element have different masses which is why the
relative atomic mass of an element (Ar) is not always a whole number.
Relative Atomic Mass (Ar) (Higher paper)
The relative atomic mass of an element (Ar) is an average value for the isotopes of the
element. It compares the mass of atoms of the element with an atom of carbon-12 (C12) which
is the most common isotope of the element carbon.
The relative atomic mass is often just the same as the mass number, but if an element has
isotopes then the relative atomic mass is given as a decimal (the average of its isotopes).
Magnesium atoms for example have a relative atomic mass of 24 which shows that they are
exactly twice as heavy as a carbon-12 atom. Chlorine has a relative atomic mass of 35.5 because
it has different isotopes.
Relative Formula Mass (Mr)
The relative formula mass of a compound (Mr) is the sum of the relative atomic masses of the
atoms in the numbers shown in the formula. For example the Mr of magnesium chloride is:
MgCl2
1 magnesium atom = 24
2 chlorine atoms = 2 x 35.5 = 71
Therefore Mr = 24 +71 = 95
The mole
The relative formula mass of a substance in grams is known as one mole of that substance.
Therefore one mole of magnesium chloride has a mass of 95g.
To find out how many moles there are in a certain mass of a substance you divide its mass by
its relative formula mass. For example if you have 142.5g of magnesium chloride:
Number of moles = 142.5g ÷ 95 = 1.5 moles
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C2 Topic 3: Analysing Substances
Paper chromatography
Chemical analysis can be used to identify additives in foods.
Artificial colours can be detected and identified by paper
chromatography. A food colouring might be pure or contain a
mixture of dyes.
A small spot of the colouring is placed on a pencil line near the
bottom of a strip of chromatography paper. The paper is then
placed into a beaker containing a suitable solvent, e.g. water or
ethanol. The solvent must be kept below the level of the pencil
line. The solvent soaks up the paper carrying the dyes with it.
Gas chromatography
Instrumental methods of detecting and identifying chemicals are accurate, sensitive and rapid and
are particularly useful when the amount of a sample is very small. Gas chromatography is an
example of an instrumental method. It allows the separation of a mixture of compounds.
1. Different substances, carried by a gas, travel through a column packed with a solid material
at different speeds, so that they become separated.
2. The number of peaks on the output of a gas chromatograph shows the number of compounds
present. The position of the peaks indicates the retention time. This is the time taken for a
substance to travel through the column and it can be used to help identify the substance.
3. The output from the gas chromatography column can be linked to a mass spectrometer. This
can be used to very quickly and accurately detect and identify very small quantities of
substances leaving the end of the column.
Carrier gas
Detector
Column
Response
Sample injected
retention time
Time (min)
(Higher paper)
The mass spectrometer can also give the relative
molecular mass of each of the substances separated in
the column. The spectrometer produces a graph and the
molecular mass is given by the molecular ion peak.
Percentage
Oven
molecular mass
Relative molecular mass
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C2 Topic 3: Relative Formula Mass Calculations
The relative formula mass of a compound (Mr) is the sum of the relative atomic masses of the
atoms in the numbers shown in the formula. For example the Mr of magnesium chloride is:
MgCl2
1 magnesium atom = 24
2 chlorine atoms = 2 x 35.5 = 71
Therefore Mr = 24 +71 = 95
Calculating the percentage composition by mass of elements in a compound
From the formula of a compound it is possible to calculate the percentage by mass of the
elements it is made up of:
Percentage mass of element
=
Ar x no. atoms of the element
x
100
Mr of compound
Example: Find the percentage mass of oxygen in calcium carbonate (CaCO3).
Answer:
Ar of calcium = 40, Ar of carbon = 12, Ar of oxygen = 16
Mr of calcium carbonate = 40 + 12 + (3 x 16) = 100
Percentage mass of oxygen
=
3 x 16
x
100
=
48%
100
Calculating the empirical formula from reacting masses (H)
The empirical formula of a compound is the ratio of its elements in its simplest form. It is
possible to calculate this from the actual experimental masses of each element that react
together to form the compound. For each element its experimental mass is divided by its A r to
give the correct ratio. The ratio is then converted to its simplest form.
Example: In an experiment it was found that 12g of magnesium reacted completely with 8g of
oxygen. Find the empirical formula of magnesium oxide.
Answer:
Ar of magnesium = 24, Ar of oxygen = 16
Empirical formula of magnesium oxide
=
12 ÷ 24
:
8 ÷ 16
0.5
0.5
1
1
Therefore the simplest ratio of magnesium atoms to oxygen atoms is 1:1 and the empirical
formula of magnesium oxide is MgO.
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C2 Topic 3: Calculating Masses from Symbol Equations (H)
The masses of reactants and products can be calculated from balanced symbol equations. You
should be able to calculate the masses of individual products from a given mass of a reactant
and the balanced symbol equation.
Method
1. Write down the balanced symbol equation for the reaction.
2. Work out the relative formula masses (Mr) of the reactant and product. Remember to
multiply the Mr by the balancing number in front of each substance in the equation.
3. For the substance you are given the mass of, divide this mass by its Mr
4. Multiply the result from step 3 by the Mr of the substance you need to find the mass of.
The three examples below follow these four simple steps to calculate the mass of a product
each time:
Example 1 – What mass of carbon dioxide is formed when 24g of carbon is burned in air?
1.
C
+
O2
2.
12
3.
24 ÷ 12 = 2
4.
2 X 44 = 88g of CO2
CO2
44
Example 2 – What mass of magnesium oxide is formed when 12g of magnesium is burned in air?
1.
2Mg
+
O2
2.
2 x 24 = 48
3.
12 ÷ 48 = 0.25
4.
0.25 X 80 = 20g of MgO
2MgO
2 x 40 = 80
Example 3 – What mass of copper is formed when 200g of copper oxide is heated with carbon?
1.
2CuO
+
C
2.
2 x 79.5 = 159
3.
200 ÷ 159 = 1.258
4.
1.258 X 127 = 160g of Cu
2Cu
+
CO2
2 x 63.5 = 127
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C2 Topic 3: Percentage Yield
The amount of a product obtained from a chemical reaction is known as the yield and it can be
calculated from the balanced symbol equation for the reaction.
The predicted yield is sometimes called the theoretical yield. When the actual yield of a
product is compared with the maximum theoretical yield as a percentage, it is called the
percentage yield:
percentage yield
=
actual yield
x
100
predicted yield
Even though no atoms are gained or lost in a chemical reaction, it is not always possible to
obtain the calculated amount (predicted yield) of a product for three main reasons:
1. The reaction may not go to completion because it is reversible:
In some chemical reactions, the products of the reaction can react to produce the starting
reactants. Such reactions are called reversible reactions and are represented as:
A + B
C + D
For example:
ammonium chloride
ammonia + hydrogen chloride
In a reversible reaction the reactants will never be completely converted into products because
some of the products are always reacting to form the reactants again.
2. Some of the product may be lost when it is separated from the reaction mixture:
Various techniques may be used to separate the product from the reaction mixture. Some of
the product may be left on the inside surfaces of equipment, or if it is filtered some will be
left in the filter paper.
3. Some of the reactants may react in ways different from the expected reaction.
The reactants may undergo other unwanted reactions that do not produce any product.
Therefore there will be less of the reactants left to produce the product.
Sustainable development and percentage yield
We can never achieve 100% yield from chemical reactions but the nearer we can get to this the
less energy and raw materials are wasted. It is important for sustainable development that we
have the most efficient production processes possible.
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C2 Topic 4: Measuring the Rate of Reactions
The rates of chemical reactions vary from very fast, explosive reactions to very slow
reactions. The rate of a chemical reaction can be measured by how quickly the reactants are
used up or how quickly products are formed:
Rate of reaction
=
amount of reactant used OR product formed
time
Four factors can increase the rate of a reaction:




increasing the temperature
increasing the concentration
increasing the surface area of a solid
adding a catalyst.
Measuring the rate of a reaction
Syringe to collect CO2 gas.
The rate of a chemical reaction can be measured using
the reaction between hydrochloric acid and marble chips
(calcium carbonate). The reaction needs to be timed while
recording results. The rate of reaction can be measured
by either collecting and measuring the volume of carbon
dioxide gas produced or by leaving the flask open on a
sensitive balance and recording the loss in mass due to
the escape of carbon dioxide.
hydrochloric acid
marble chips
The temperature and concentration of the acid can be changed to investigate the effect on
reaction rate. The marble chips can be broken into smaller pieces to investigate the effect of
surface area. Many small chips have a greater total surface area than fewer large chips.
Another method of following the rate of a reaction is by observing the formation or loss of a
colour or precipitate as a reaction proceeds.
Effect of changing conditions on the rate of reaction
The rate of reaction decreases when the concentration
of hydrochloric acid is reduced, the temperature is
lowered or larger marble chips are used. The blue curve
on the graph shows this.
higher rate of reaction
Amount of product
The rate of reaction increases when the concentration of
hydrochloric acid is increased, the temperature is raised
or the marble chips are broken up. The red curve on the
graph opposite shows this. The reaction occurs most
quickly at the beginning and then slows down as the
reactants are used up until the curve levels out when no
further product is being produced.
lower rate of reaction
Time
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C2 Topic 4: The Collision Theory
The rate of chemical reactions is affected by temperature, concentration, surface area and
catalysts. The collision theory explains how this happens. It all depends on how often the
reactant particles collide and how much energy they have to break bonds and react.
Temperature
Chemical reactions can only occur when reacting
particles collide with each other and with enough
no reaction
reaction
At low temperatures
there are low energy
collisions.
At high temperatures
there are high energy
collisions.
energy. The minimum amount of energy particles must
have to react is called the activation energy.
Increasing the temperature increases the speed of
the reacting particles so that they collide more
frequently and more energetically. This increases the
rate of reaction.
Concentration
Increasing the concentration of reactants in solutions
brings the particles closer together. This increases
the frequency of collisions and so increases the rate of
reaction.
Increasing the pressure of reacting gases brings the
particles closer together. This also increases the
frequency of collisions and so increases the rate of
low concentration
high concentration
small surface area
large surface area
reaction.
Surface area
If one of the reactants is a solid its total surface area
can be increased by breaking it up into smaller pieces.
The increased surface area increases the frequency of
collisions and so increases the rate of reaction.
Catalysts
Catalysts increase the rate of chemical reactions but
they are not changed or used up during the reaction.
They work by lowering the activation energy for the
reaction. Different reactions need different catalysts.
Reactants react more easily on the catalyst surface.
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C2 Topic 4: Catalysts
A catalyst speeds up the rate of a reaction but is not used up by the reaction. Catalysts work
by lowering the energy needed by the reactants to react. Catalysts are important in
increasing the rates of chemical reactions used in industrial processes.
Investigating catalysts
Hydrogen peroxide is a chemical that can be used to
investigate the action of a number of catalysts. It
Syringe to collect oxygen gas.
naturally breaks down very slowly to release oxygen gas
but adding a suitable catalyst greatly increases its rate
of breakdown.
The rate of reaction can be measured in the same way as
hydrogen peroxide
the reaction with hydrochloric acid and marble chips by
catalyst
timing the reaction while collecting and measuring the
volume of oxygen gas produced.
The effects of various catalysts such as manganese oxide, potato and liver can be observed to
find out which is the most effective.
Uses of catalysts in industry
Catalysts are important in speeding up chemical reactions in many industrial processes.
Catalysts save on energy costs by speeding up industrial processes and lowering the
temperature needed for reactions to proceed. Saving energy also helps sustainable
development.
Catalytic converters in vehicles reduce the amount of waste gases that come out of the
exhaust pipe. The platinum catalyst is spread over a honeycomb structure to greatly increase
the surface area. This increases the rate of catalysed reactions taking place. Catalytic
converters work best at high temperatures when the car engine is fully warmed up.
Disadvantages of catalysts

Catalysts can be very expensive.

Different reactions need different catalysts therefore a number of catalysts are
usually needed for the different stages in a manufacturing process.

Catalysts often need to be removed from the product and cleaned before they can be
used again.
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C2 Topics 3 & 4 Questions
1. Explain why the relative atomic mass of an element is not always a whole number.
2. Explain how the relative atomic mass of an element is calculated. (Higher paper)
3. How is the relative formula mass of a compound (Mr) calculated?
4. What quantity is one mole of a substance?
5. Calculate the relative formula mass of sodium oxide (Na2O)
(Ar of sodium = 23,
Ar of oxygen = 16)
6. Find the percentage mass of chlorine in magnesium chloride (MgCl2).
(Ar of magnesium = 24,
Ar of chlorine = 35.5)
7. In an experiment it was found that 10g of calcium reacted completely with 4g of
oxygen. Find the empirical formula of calcium oxide. (Higher paper)
(Ar of calcium = 40,
Ar of oxygen = 16)
8. What is meant by the term ‘percentage yield’ of a reaction?
9. Give three reasons why it is often not possible to obtain all of the calculated predicted yield
of a product.
10. What simple technique can be used to separate and identify dyes in a food colouring?
11. In gas chromatography what does the number of peaks on the chromatograph indicate?
12. Use the particle theory to explain why increasing the temperature increases the rate of a
reaction.
13. Use the particle theory to explain why increasing the concentration of reactants in
solutions increases the rate of a reaction.
14. Which substances increase the rate of chemical reactions without being changed
themselves?
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C2 Topics 3 & 4 Answers
1. Atoms of the same element can have different numbers of neutrons and are called isotopes of that
element. The different isotopes of an element have different masses which is why the relative atomic
mass of an element is not always a whole number.
2. The relative atomic mass of an element is an average value for the isotopes of the element. It
compares the mass of atoms of the element with an atom of carbon-12 (C12) which is the most
common isotope of the element carbon. (Higher paper)
3. The relative formula mass of a compound (Mr) is the sum of the relative atomic masses of the atoms
in the numbers shown in the formula.
4. One mole of a substance is the relative formula mass in grams.
5. The relative formula mass of sodium oxide (Na2O):
2 sodium atoms = 2 x 23 = 46,
1 oxygen atom = 16,
Therefore M r = 46 +16 = 62
6. The percentage mass of chlorine in magnesium chloride (MgCl2):
Mr of magnesium chloride = 24 + (2 x 35.5) = 95
Percentage mass of chlorine = (2 x 35.5) x 100 = 75%
95
7. Empirical formula of calcium oxide
=
10 ÷ 40
0.25
1
:
4 ÷ 16
0.25
1
Therefore the simplest ratio of calcium atoms to oxygen atoms is 1:1 and the empirical formula of
calcium oxide is CaO.
8. The percentage yield of a reaction is the actual yield of a product compared with the maximum
theoretical yield as a percentage.
9. Reasons why it is often not possible to obtain all of the calculated predicted yield of a product are:
- The reaction may not go to completion because it is reversible.
- Some of the product may be lost when it is separated from the reaction mixture.
- The reactants may undergo other unwanted reactions that do not produce any product.
10. Paper chromatography can be used to separate and identify dyes in a food colouring.
11. In gas chromatography the number of peaks on the chromatograph indicates the number of
compounds present.
12. Increasing the temperature increases the speed of the reacting particles so that they collide more
frequently and more energetically. This increases the rate of reaction.
13. Increasing the concentration of reactants in solutions brings the particles closer together. This
increases the frequency of collisions and so increases the rate of reaction.
14. Catalysts increase the rate of chemical reactions without being changed themselves.
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C2 Topic 5: Energy Changes in Chemical Reactions
When chemical reactions occur, energy is transferred to or from the surroundings.
Exothermic reactions
An exothermic reaction is one that transfers energy to the surroundings i.e. it involves the
release of energy usually in the form of heat, but sometimes light or sound as well. Examples of
exothermic reactions include:



combustion (burning)
many oxidation reactions
neutralisation (acid/alkali reactions).
Exothermic reactions have many everyday uses. For example self-heating cans which make
hot drinks have chemicals in their bases that react exothermically. Hand warmers use the
oxidation of iron oxide reaction to generate heat. As the disposable packet is opened, oxygen
reacts with iron, in the presence of a salt solution catalyst.
Endothermic reactions
An endothermic reaction is one that takes in energy from the surroundings i.e. energy must be
supplied to the reaction (usually in the form of heat). A fall in temperature indicates an
endothermic reaction is taking place as heat energy is taken in. Endothermic reactions include
thermal decompositions (break down of chemicals using heat). Some sports injury cool packs
are based upon endothermic reactions.
Measuring temperature changes in reactions
A rise in temperature indicates an
exothermic reaction is taking place as heat
energy is given out.
A fall in temperature indicates an
endothermic reaction is taking place as heat
energy is taken in.
Thermometer records
temperature change.
lid
Polystyrene cup
insulates contents.
reacting chemicals
Reversible reactions
If a reversible reaction is exothermic in one direction, it is endothermic in the opposite
direction. The same amount of energy is transferred in each case. For example:
endothermic
hydrated copper sulfate
anhydrous copper sulfate + water
exothermic
(BLUE)
(WHITE)
If blue hydrated copper sulfate crystals are heated they break down into white anhydrous
copper sulphate and water. If water is then added to the anhydrous copper sulphate heat
energy is released as it converts back to the hydrated form.
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C2 Topic 6: Acids and Bases
An acid is any substance with a pH number of less than 7 and a base is any substance with a pH
number of more than 7. Metal oxides and hydroxides are bases. Soluble hydroxides are called
alkalis.
Universal indicator and the pH scale
Hydrogen ions, H
+
(aq),
make solutions acidic and hydroxide ions OH
-
(aq),
make solutions alkaline.
The pH scale is a measure of the acidity or alkalinity of a solution.
Universal indicator contains a mixture of indicators which change to a number of different
colours depending on the acid or alkali being tested. Universal indicator can be used in a
solution or paper form. The colour change gives an estimate of the pH number of the substance
being tested.
pH
0
1
2
3
5
4
6
7
NEUTRAL
WEAK ACIDS
STRONG ACIDS
Acids becoming stronger
8
9
10 11 12 13 14
WEAK ALKALIS
STRONG ALKALIS
Alkalis becoming stronger
Neutralisation reactions
Acids and bases react together to form a salt and water. Salts are neutral (pH 7) so this is
called a neutralisation reaction.
ACID
+
BASE
SALT
+
WATER
In neutralisation reactions hydrogen ions react with hydroxide ions to produce water. This
reaction can be represented by the equation:
H
+
(aq)
+
OH
-
(aq)
H2O(l)
State symbols
(s) = SOLID,
(l) = LIQUID,
(g) = GAS,
(aq) = AQUEOUS (dissolved in water)
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C2 Topic 6: Acid Reactions and Salts
Acids can undergo a number of reactions to produce different salts:
1. METAL + ACID
SALT + HYDROGEN
e.g. magnesium + sulfuric acid
Mg(s)
+
magnesium sulfate + hydrogen
H2SO4(aq)
MgSO4(aq)
+
H2(g)
2. METAL OXIDE + ACID
SALT + WATER
e.g. copper oxide + sulfuric acid
copper sulfate + water
CuO(s)
+
H2SO4(aq)
CuSO4(aq)
3. METAL HYDROXIDE + ACID
SALT + WATER
e.g. sodium hydroxide + hydrochloric acid
NaOH(aq)
+
+ H2O(l)
sodium chloride + water
HCl(aq)
NaCl(aq)
+ H2O(l)
Ammonium salts
Ammonia dissolves in water to produce an alkaline solution. It is used to produce ammonium
salts such as ammonium nitrate which are very important as fertilisers.
e.g. ammonia + nitric acid
NH3(aq) +
HNO3(aq)
ammonium nitrate
NH4NO3(aq)
Naming salts
The first part of the name of a salt comes from the metal, or the metal in the oxide or
hydroxide. The second part of the name comes from the acid.
Hydrochloric acid produces chloride salts. Nitric acid produces nitrate salts. Sulfuric acid
produces sulfate salts.
The table below gives examples of acids, and their salts:
Acid
Reactant
Salt
Hydrochloric acid
Magnesium
Magnesium chloride
Hydrochloric acid
Sodium hydroxide
Sodium chloride
Hydrochloric acid
Zinc oxide
Zinc chloride
Sulfuric acid
Copper oxide
Copper sulfate
Sulfuric acid
Calcium hydroxide
Calcium sulfate
Sulfuric acid
Ammonia
Ammonium sulfate
Nitric acid
Zinc
Zinc nitrate
Nitric acid
Iron oxide
Iron nitrate
Nitric acid
Sodium hydroxide
Sodium nitrate
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C2 Topic 6: Making Salts
Making soluble salts
Soluble salts can be made by reacting acids with:



metals – not all metals are suitable; some are too reactive and others are not reactive
enough. For example an excess of solid magnesium metal can be added to hydrochloric
acid. When all of the acid has reacted a pure solution of magnesium chloride can be
separated from the excess magnesium metal.
insoluble bases – the base is added to the acid until no more will react and the excess
solid is filtered off. For example copper oxide can be reacted with sulphuric acid to form
copper sulphate salt. The excess copper oxide is filtered off to leave a pure solution of
copper sulphate.
alkalis – the correct amount of alkali can be added to an acid to neutralise it and form a
salt solution. An indicator can be used to show when the acid and alkali have completely
reacted. For example an exact amount of sodium hydroxide solution can be added to
hydrochloric acid to form a pure solution of sodium chloride salt.
If a salt solution is left to evaporate solid crystals of the salt can be produced. The salt
solution can also be heated gently to evaporate the water off more quickly.
Making insoluble salts
Insoluble salts can be made by mixing appropriate solutions of ions so that a precipitate of the
insoluble salt is formed. For example if the colourless solutions of lead nitrate and potassium
iodide are mixed together in a flask a golden yellow precipitate of solid lead iodide is produced.
Pb(NO3)2(aq)
The two solutions
are mixed.
+
2KI(aq)
PbI2(s)
+ 2KNO3(aq)
The reaction forms
a yellow precipitate.
The salt can then be separated from the solution by filtering. It is then washed and dried to
produce a solid yellow powder.
Uses of precipitation reactions
Precipitation reactions can be used to remove unwanted poisonous ions from solutions, for
example in treating water for drinking or in treating effluent (sewage).
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C2 Topic 7: Electrolysis
When an ionic substance is melted or dissolved in water, the ions are free to move about within
the liquid or solution. Passing an electric current through ionic substances that are molten or in
solution breaks them down into elements. This process is called electrolysis and the substance
that is broken down is called the electrolyte.
During electrolysis, positively charged ions move to
the negative electrode, and negatively charged ions
move to the positive electrode.
At the negative electrode, positively charged ions gain
electrons (reduction) and at the positive electrode,
negatively charged ions lose electrons (oxidation).
d.c. supply
positive
electrode
negative
electrode
If there is a mixture of ions, the products formed
depend on the reactivity of the elements involved.
Electrolysis of molten lead bromide
Molten lead bromide can be separated into lead metal and bromine gas using electrolysis:
positive
electrode
Negatively charged bromide
ions are attracted towards
the positive electrode. Here
they are oxidised to bromine
atoms by losing electrons.
negative
electrode
lead ion
bromide ion
Positively charged lead ions
are attracted towards the
negative electrode. Here
they are reduced to lead
atoms by gaining electrons.
Electroplating
Electrolysis is used to electroplate objects. It is often used to coat a cheap metal with a more
expensive one such as silver. The negative electrode is the object that you want to
electroplate. The positive electrode is the metal that you want to coat the object with.
The electrolyte must contain ions of the plating metal e.g. silver nitrate solution to plate with
silver or copper sulphate solution to plate with copper.
Half equations for reactions at electrodes (H)
Reactions at electrodes can be represented by half equations, for example:
At the positive electrode: 2Br– ➞ Br2 + 2e–
or 2Br– - 2e– ➞ Br2 + 2e–
At the negative electrode: Pb2+ + 2e– ➞ Pb
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C2 Topic 7: Electrolysis of Aluminium Oxide and Sodium Chloride
Extraction of aluminium by electrolysis
Aluminium is found as an ore called bauxite which is mainly composed of aluminium oxide.
Aluminium is manufactured by the electrolysis of a molten mixture of aluminium oxide and
cryolite. Cryolite is a chemical that brings the melting point of aluminium oxide down from
about 2000°C to around 900°C. This makes the process much easier to carry out and reduces
the production costs.
The electrodes are made of graphite which is a form of carbon. Aluminium forms at the
negative electrode and oxygen at the positive electrode. The positive electrode is constantly
wearing away because it reacts with the oxygen to produce carbon dioxide.
solid crust
carbon positive electrode
molten mixture of
bauxite and cryolite
carbon negative electrode
molten aluminium
Large scale manufacture of chlorine and its uses
The electrolysis of sodium chloride solution produces hydrogen and chlorine. Sodium hydroxide
solution is also produced. Chlorine is produced on an industrial scale by the electrolysis of sea
water. Chlorine is a toxic gas which makes its manufacture hazardous. Chlorine has three main
uses:



the manufacture of bleach
the manufacture of plastics
it is added to water supplies and swimming pools to kill bacteria.
Sodium hydroxide is a strong alkali with many uses including the production of cleaning agents
and soap.
positive
electrode
Cl2
H2
negative
electrode
hydrogen ion
chloride ion
hydrogen molecule
chlorine molecule
Sodium hydroxide solution
forms during the process.
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C2 Topics 5-7 Questions
1. Explain what is meant by the term ‘exothermic reaction’.
2. Give two examples of exothermic reactions.
3. Explain what is meant by the term ‘endothermic reaction’.
4. What indicates that an endothermic reaction is taking place?
5. Which ions make solutions acidic?
6. Which ions make solutions alkaline?
7. What does universal indicator contain?
8. Complete the following equations:

ACID + BASE =

METAL + ACID =
9. What important use does ammonium nitrate have?
10. Name the salt produced by reacting hydrochloric acid with magnesium.
11. Name the salt produced by reacting sulphuric acid with copper oxide.
12. Name the salt produced by reacting nitric acid with zinc.
13. How can insoluble salts be made?
14. Which electrode do positively charged ions move towards during electrolysis?
15. Explain what happens to the ions at each electrode during electrolysis.
16. What is molten lead bromide separated into during electrolysis?
17. What must the electrolyte contain in the process of electroplating?
18. How is aluminium extracted from its ore?
19. Name three products from the electrolysis of sodium chloride solution.
20. What are the three main uses of chlorine?
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C2 Topics 5-7 Answers
1. An exothermic reaction is one that transfers energy to the surroundings i.e. it involves the release of
energy usually in the form of heat, but sometimes light or sound as well.
2. Examples of exothermic reactions include:
- combustion (burning)
- many oxidation reactions
- neutralisation (acid/alkali reactions).
3. An endothermic reaction is one that takes in energy from the surroundings i.e. energy must be
supplied to the reaction (usually in the form of heat).
4. A fall in temperature indicates an endothermic reaction is taking place as heat energy is taken in.
5. Hydrogen ions, H+(aq), make solutions acidic.
6. Hydroxide ions OH-(aq), make solutions alkaline.
7. Universal indicator contains a mixture of indicators which change to a number of different colours
depending on the acid or alkali being tested.
8. ACID + BASE = SALT + WATER
METAL + ACID = SALT + HYDROGEN
9. Ammonium nitrate is a very important fertiliser.
10. Magnesium chloride is the salt produced by reacting hydrochloric acid with magnesium.
11. Copper sulphate is the salt produced by reacting sulphuric acid with copper oxide.
12. Zinc nitrate is the salt produced by reacting nitric acid with zinc.
13. Insoluble salts can be made by mixing appropriate solutions of ions so that a precipitate of the
insoluble salt is formed.
14. Positively charged ions move towards the negative electrode during electrolysis.
15. During electrolysis at the negative electrode, positively charged ions gain electrons (reduction) and
at the positive electrode, negatively charged ions lose electrons (oxidation).
16. Molten lead bromide is separated into lead metal and bromine gas during electrolysis.
17. In the process of electroplating the electrolyte must contain ions of the plating metal e.g. silver
nitrate solution to plate with silver or copper sulphate solution to plate with copper.
18. Aluminium is extracted from its ore by the electrolysis of a molten mixture of aluminium oxide and
cryolite.
19. The electrolysis of sodium chloride solution produces hydrogen, chlorine and sodium hydroxide
solution.
20. The three main uses of chlorine are the manufacture of bleach, the manufacture of plastics and it is
added to water supplies and swimming pools to kill bacteria.
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