Chemistry Revision Notes
Contents
Section 1: Principles of Chemistry .............................................................................................................................. 4
States of Matter ..................................................................................................................................................... 4
Atoms ..................................................................................................................................................................... 4
Diffusion and Dilution ........................................................................................................................................ 5
Separation of Mixtures....................................................................................................................................... 5
Atomic Structure .................................................................................................................................................... 7
Relative Atomic Mass (Ar) .................................................................................................................................. 7
Electronic Structure............................................................................................................................................ 8
Relative Formula Masses and Molar Volumes of Gases ........................................................................................ 8
The Mole ............................................................................................................................................................ 8
Calculations ........................................................................................................................................................ 9
Chemical Formulae and Chemical Equations ......................................................................................................... 9
Empirical Formulae .......................................................................................................................................... 10
Molecular Formula ........................................................................................................................................... 11
Percentage Yield ............................................................................................................................................... 11
Ionic Compounds.................................................................................................................................................. 11
Covalent Bonding ................................................................................................................................................. 12
Metallic Bonding .................................................................................................................................................. 15
Electrolysis............................................................................................................................................................ 18
Conductivity ..................................................................................................................................................... 18
Ions ................................................................................................................................................................... 18
Products of Electrolysis .................................................................................................................................... 18
Electrolysis and Redox...................................................................................................................................... 18
Electrolysis calculations........................................................................................................................................ 20
The Chlor-Alkali Industry ...................................................................................................................................... 21
Section 2: Chemistry of the Elements ...................................................................................................................... 23
The Periodic Table ............................................................................................................................................ 23
Group I (The Alkali Metals)................................................................................................................................... 25
Group VII (The Halogens) ..................................................................................................................................... 25
1
Group 0 – The Noble Gases .................................................................................................................................. 26
Oxygen And Oxides .............................................................................................................................................. 27
Hydrogen and Water ............................................................................................................................................ 29
Reactivity Series ................................................................................................................................................... 30
Reduction and Oxidation...................................................................................................................................... 31
Test for Ions.......................................................................................................................................................... 31
Section 3 – Organic Chemistry ................................................................................................................................. 33
The Alkanes .......................................................................................................................................................... 33
Naming Hydrocarbons...................................................................................................................................... 33
Bromination...................................................................................................................................................... 33
Isomerism ......................................................................................................................................................... 34
Combustion ...................................................................................................................................................... 35
Alkenes ................................................................................................................................................................. 35
Addition Reactions ........................................................................................................................................... 35
Ethanol ................................................................................................................................................................. 36
Fermentation:................................................................................................................................................... 36
Addition of Steam to Ethene ............................................................................................................................ 36
Dehydration of Ethanol ........................................................................................................................................ 36
Section 4 – Physical Chemistry ................................................................................................................................. 38
Acids, Bases and Salts........................................................................................................................................... 38
Reactions of Acids ............................................................................................................................................ 38
Definitions ........................................................................................................................................................ 39
Salt Preparation.................................................................................................................................................... 39
Rules for Solubility:........................................................................................................................................... 39
Precipitation ..................................................................................................................................................... 39
Titration ............................................................................................................................................................ 40
Reaction of an Acid with an Insoluble Base ..................................................................................................... 40
Energetics ............................................................................................................................................................. 41
Measuring Enthalpy Changes ........................................................................................................................... 41
Calculating Enthalpy Changes .......................................................................................................................... 42
Rates of Reaction ................................................................................................................................................. 44
Factors that affect the rate of reaction ............................................................................................................ 44
Equilibria............................................................................................................................................................... 45
Le Chatelier’s Principle ..................................................................................................................................... 45
Section 5 – Chemistry in Society .............................................................................................................................. 46
2
Extraction of metals ............................................................................................................................................. 46
The Blast Furnace ............................................................................................................................................. 46
Aluminium Extraction ....................................................................................................................................... 47
Uses of Iron ...................................................................................................................................................... 47
Uses of Aluminium ........................................................................................................................................... 48
Crude Oil............................................................................................................................................................... 48
Crude Oil Processing......................................................................................................................................... 48
Cracking ............................................................................................................................................................ 48
Polymerisation ..................................................................................................................................................... 49
Addition ............................................................................................................................................................ 49
Drawing Polymers ............................................................................................................................................ 50
Condensation Polymerisation .......................................................................................................................... 51
Chemical Manufacture ......................................................................................................................................... 52
The Haber Process ............................................................................................................................................ 52
The Contact Process ......................................................................................................................................... 53
3
Section 1: Principles of Chemistry
States of Matter
STATE
DESCRIPTION
PROPERTIES AND EXPLANATIONS
Particles in close contact,
they are arranged in a
lattice. Vibrate around a
fixed point.
High density, cannot be compressed and has a fixed
shape.
Solid
As the solid is heated, the particles gain kinetic energy, and vibrate more rapidly.
When the melting point is reached, the particles have sufficient energy to
overcome the forces of attraction between them sufficiently and the particles will
move apart
LIQUID
In close contact in an
irregular fashion. Can move
around but cannot
separate.
They are less dense than solids, and still
incompressible. They fill the shape of their
container.
As temperature rises, the particles gain kinetic energy. When the boiling point is
reached, the particles have sufficient energy to overcome completely the forces of
attraction between them, and become separated.
GAS
Particles are spread far
apart, and in no regular
pattern. They are in rapid,
constant, random motion.
Low density and easily compressed. They expand to
fill their containers.
Atoms
An atom is the smallest particle of an element which can exist. An element is a substance containing one kind of
atom. A compound is two or more elements chemically combined in a fixed ratio as shown by its formula, where
as a molecule is two or more atoms joined together by chemical bonds. A mixture is two or more substances
combined.
4
Diffusion and Dilution
Some experiments can be done to prove the presence of tiny moving particles. If a small crystal of potassium
permanganate is placed in a beaker of water, it begins to dissolve, giving a purple solution. The purple colour
slowly spreads out from the crystal, as the particles (permanganate ions) move around randomly and spread out
through the water molecules. If the solution is then diluted further, the purple colour becomes paler, as the
permanganate ions become spread further apart. Since the colour caused by the particles can still be seen even
if only a tiny crystal is dissolved in a large volume of water, the crystal -must contain very many particles. These
particles must, then, be very small. The same effect can be seen if a drop of bromine is placed at the bottom of a
covered gas jar. The bromine evaporates, and the red vapour spreads out to fill the jar, as the bromine
molecules diffuse throughout the molecules in the air. This is much more rapid than diffusion of a coloured
solution, since the particles in a gas are much further apart, and move more rapidly.
Separation of Mixtures
To separate a compound into the elements from which it is made requires a chemical reaction; chemical bonds
must be broken, and this often requires a lot of energy. To separate the components of a mixture, it usually
requires a physical reaction. The appropriate method depends on the type of mixture.
Filtration
This is used to separate a solid from a liquid. The mixture is poured through a filter paper within a filter funnel.
Liquid (filtrate) passes through, whilst the solid (residue) remains in paper. This can be used to separate two
solids, if one is soluble and the other insoluble.
Crystallisation
When a solid is dissolved in water, it is possible to obtain the solid in the form of crystals. The solution is gently
heated in an evaporating basin, until about half of the water has evaporated. The remaining concentrated
solution is then left to cool, and the liquid to evaporate.
Simple Distillation
thermometer
water
out
Liebig
condenser
This is used to separate a liquid
solvent and a solute. The flask is
heated, causing the solvent to boil
off. The vapour rises, passes into the
condenser, and is cooled, causing it
to condense. The liquid will run ne
run off. The solute, having a much
higher boiling point, will remain in
the flask.
water
in
roundbottomed
flask
salt
solution
gauze
tripod
distilled
water
Bunsen
burner
5
Fractional Distillation
thermometer
water
out
Liebig
condenser
water
in
fractionating
column
a fraction
roundbottomed
flask
mixture of
liquids
gauze
This is used to separate a mixture
of liquids, based on their different
boiling points. The apparatus is
similar to simple distillation, but
also has a fractionating column
between the flask and the
condenser, giving a large surface
area. The substance with the
lowest boiling point will boil off
first. As the vapour rises through
the column, into the condenser, it
turns to a liquid and is collected.
The process repeats with a higher
temperature for another liquid,
and the next liquid is collected. The
process continues, until all of the
components in the mixture have
boiled, condensed, and been
collected.
tripod
Bunsen
burner
Chromatography
It is used to separate mixtures of coloured compounds which are soluble. A pencil line is drawn just above the
solvent line, and a small spot of each substance to be tested is placed on this line. The paper is then suspended
in a beaker with a solvent in it.
The solvent soaks up the filter paper, and
dissolves the coloured substances in each
sample, carrying them up the paper with it.
Different substances are carried different
distances. This can show how many
different components are present, and is
often used to compare several inks to see
if one matches an original sample used for
comparison.
lid
filter
paper
pencil
line
solvent
6
Atomic Structure
An atom is the smallest particle of an element which can exist. It is possible to split an atom into smaller
particles, but these will no longer display the properties of a particular chemical element.
RELATIVE MASS
RELATIVE CHARGE
LOCATION
Proton
1
+1
nucleus
Neutron
1
0
nucleus
Electron
Negligible (1/2000)
-1
orbiting nucleus
The mass of the atom is found in the nucleus (protons and neutrons). Electrons have a relatively insignificant
mass, and orbit the nucleus. The atomic number is the number of protons in one atom of a particular element.
This number determines which element a particular atom is. The atomic number is shown at the bottom-left of
the element’s symbol. An atom has no overall charge. It has an equal number of protons (positively charged)
and electrons (negatively charged). The mass number is the total number of protons and neutrons in one atom.
This tells us the mass of the atom, since the mass of the electrons is insignificant. The mass number is written at
the top-left of the symbol. Isotopes are atoms of the same element, with the same number of protons and
electrons, but a different numbers of neutrons.
Mass number =15
e.g. a nitrogen-15 atom:
Atomic number = 7
Relative Atomic Mass (Ar)
The relative atomic mass is the average mass of one atom, relative to 1/12 of the mass of a carbon-12 atom.
This can be calculated, if the isotopes and their abundances are known:
Chlorine consists of two isotopes, chlorine-35, and chlorine-37. 75% of chlorine atoms are Cl-35, and 25% are Cl37.
R.A.M.
= (75/100) x 35 + (25/100) x 37
= 35.5
The relative atomic mass is shown on the Periodic Table above the symbol of each element.
7
Electronic Structure
The electrons in an atom orbit the nucleus at certain fixed distances – shells. Each shell can hold a certain
maximum number of electrons – the pattern becomes quite complex beyond the element calcium, but for the
first 20 elements it is relatively simple:
Shell
1
2
3
4
Maximum number of electrons
2
8
8
(2)
The electrons fill up the shells, beginning with the 1st shell (closest to the nucleus), and moving on to the next
shell out when one is full. The number of outer electrons in the outer shell of an atom is equal to its group
number on the periodic table. Because all chemical reactions are the result of changes in the outer shell
electrons of the reacting atoms, atoms in the same group, having the same number of electrons in their outer
shell, will react in a similar way.
Relative Formula Masses and Molar Volumes of Gases
The relative formula mass of a compound can be found by adding up all of the individual atoms in the formula.
The percentage of an element in a compound in term of its mass is important, and can also find the formula of
an unknown compound.
What is the percentage by mass of oxygen in sodium nitrate, NaNO3?
Relative formula mass
Total relative mass of oxygen
= 23 + 14 + 3 x 16
= 3 x16
= 85
= 48
Oxygen therefore accounts for 48/85 of the total mass, so:
48
85
Percentage of oxygen by mass =
x 100%
= 56.5% (3 s.f.)
The element in question may appear in more than one place in the formula.
e.g.
What is the percentage by mass of oxygen in hydrated copper(II) sulphate, CuSO4.5H2O?
Relative formula mass
= 64 + 32 + 4 x 16 + 5 x (2 x 1 + 16)
Total relative mass of oxygen
= 9 x 16 (there are 9 O atoms in total)
Percentage by mass
= 144/250 x 100%
= 250
= 144
= 57.6%
The Mole
The mole is an amount of a substance. It is equal to 6.02x1023, which is also known as Avogadro ’s number. 2
moles of AlCl3 contains 2 moles of Al3+ ions, but 6 moles of Cl- ions. Moles also equal the amount of a substance.
Moles = mass/relative formula mass.
8
Calculations
Reacting Masses
e.g. What mass of oxygen is needed to burn 3.00kg of propane, C3H8?
Mr of propane = 3 x 12 + 8 x 1 = 44
1. Calculate the number of moles of
propane given in the question.
2. Write a balanced equation for the
chemical reaction.
3. Use the ratio given by the balancing
numbers in the equation to find the
number of moles of O2 required to
react with the C3H8.
4. Convert the moles of oxygen back
into a mass.
moles 
mass
relative formulamass
C3H8 + 5O2

3000g
44g
= 68.2 mol of C3H8
3CO2 + 4H2O
1 mol of C3H8 : 5 moles O2
68.2 x 5 = 341
so 68.2 mol C3H8 : 341 mol O2
mass = moles x Mr
= 341 x 32
= 10,900g
Mass of O2 = 10.9 kg (3 s.f.)
Gas Volumes
At room temperature and pressure one mole of a gas has the same volume 24dm3.
volume of gas (dm 3 )
volume of gas (cm3 )
24
24000
Moles of gas =
or
Solution Calculations
A concentration of 1mol/dm3 means that 1 mol of the substance is dissolved in each dm3 of water.
Concentration (mol / g dm-3) =
amount (mol)/gram s(g)
volume (dm 3 )
Chemical Formulae and Chemical Equations
The formula of a compound tells us the number of each type of atom in a molecule. The formula of sulphuric
acid is H2SO4. This tells us that there are two moles of hydrogen, one mole of sulphur and four moles of oxygen.
For ionic compounds, the formula can be worked out by making sure the total charge equals zero. The valency
9
can also be used. The valency represents the total numbers of covalent bonds an atom can form, or the charge
on an ion. Valencies can be worked out by using the periodic table:
Group
Valency
I
1
II
2
III
3
IV
4
V
3
VI
2
VII
1
O
0
The valencies of some elements in some compounds are given in the name in roman numerals, in brackets. For
example: Iron (III) Chloride. The valencies of some compounds and elements must be learned:
Name
Formula Valency
Hydrogen
H2+
1
Zinc
Zn
2
Nitrate
NO3
1
Hydroxide
OH
1
Carbonate
CO3
2
Sulphate
SO4
2
Ammonium (forms positive ion)
NH4
1
Diatomic elements are hydrogen, nitrogen, oxygen and the halogens.
Empirical Formulae
The empirical formula of a compound is the simplest whole number ratio of atoms in a compound.
e.g.
2.88g of magnesium is heated in nitrogen, and forms 4.00g of magnesium nitride. Find the
empirical formula of magnesium nitride.
Mass of nitrogen = 4.00 - 2.88 = 1.12g
Write the mass, or percentage by
mass, of each element.
Mg
N
2.88g
1.12g
Divide each mass (or percentage) by
the relative atomic mass of the
element, to convert to moles.
2.88
24
1.12
14
=0.120mol
=0.0800mol
Simplify this mole ratio by dividing
each number by the smallest.
0.120
 1.50
0.0800
0.0800
 1.00
0.0800
If this does not give whole number,
multiply each by an appropriate figure
– in this case, the values are doubled.
1.50 x 2 = 3.00
1.00 x 2 = 2.00
Having arrived at a whole-number ratio, state the ratio as a formula:
Empirical formula = Mg3N2
10
Molecular Formula
This is the actual number of each type of atom in one molecule of a molecular substance. This must be a whole
number multiple of the empirical formula, which means the relative molecular mass must be a whole number
multiple of the empirical formula mass.
A compound has the empirical formula CH2O, and a relative molecular mass of 120. What is its molecular
formula?
Relative mass of CH2O = 12 + 2 x 1 + 16 = 30
120 / 30 = 4, so there must be 4 lots of CH2O in the molecule to give the mass of 120
Therefore molecular formula must be C4H8O4
Percentage Yield
Percentage yield = actual yield/theoretical maximum yield x 100%
The percentage yield may not be 100% for several reasons. The reaction may not be complete, there may be
other side reactions occurring, or the product cannot be fully separated.
Ionic Compounds
There are three different types of bonding which hold together the atoms in substances which are ionic,
covalent and metallic. These occur because of the redistributing of electrons. Ionic bonding occurs in
compounds of a metal with a non-metal, as well as ammonium compounds. When dissolved in water, acids also
form ions. The atoms become stable from a full outer shell of electrons. The gaining/losing of electrons cause
these to become ions.
e.g. calcium chloride
CaCl2
1.
Calcium atoms have 2 electrons in their outer
shell.
Chlorine atoms have 7 electrons in their outer
shells.
2.
Electrons are transferred from the metal,
calcium, to the non-metal, chlorine.
11
3. Draw the final result, placing the ions in square brackets, and remembering to show the charge on each
ion. Each ion must have a full outer shell of electrons:
Calcium has lost 2 electrons, so has become a
2+ ion. The chlorine atoms have gained one
electron each, and have a charge of 1-. These
are now attracted to each other, and held
together by strong electrostatic attractions
between ions with opposite charges.
Metals ions always form positive ions. Non-metals form negative ions - except for hydrogen (H+) and ammonium
(NH4+) ions. The ions are generally arranged in an ionic lattice – a giant structure, placing positive ions next to
negative, to maximise the total attraction, such as in the sodium chloride lattice:
Na+ ion
Cl- ion
They have high melting and boiling points, because of their giant structure, with strong electrostatic attractions
between positive and negative ions throughout the entire structure. They are also brittle. When being bent the
similarly charged ions touch and repel, shattering the object. The size of the charge on an ion is equal to its
valency. Magnesium is in Group II, so has 2 outer electrons; it will therefore lose these when it reacts, and forms
a Mg2+ ion. Oxygen is in Group VI, and has 6 outer electrons; it therefore needs to gain 2 electrons to fill the
shell, and will form an O2- ion. The size of the charge on the ions affects the properties of the ionic compound.
For example, the melting point of magnesium oxide, MgO, is much higher than that of sodium chloride, NaCl.
MgO consists of ions with two units of charge – Mg2+ and O2- - which therefore attract each other much more
strongly than the singly charged Na+ and Cl- ions in NaCl, so much more heat energy is required to separate
them.
Covalent Bonding
A covalent bond is a shared pair of electrons. The outer shells overlap as the atoms share pairs of electrons, so
that both atoms can achieve a full outer shell. The bond holds the atoms together, because the positively
charged protons in the nuclei of the two atoms are both electrostatically attracted to the negatively charged
electron pair in the bond.
Covalent bonding can also be represented by a dot-and-cross diagram
e.g. oxygen
Draw correct number of electrons on outer shell:
12
Oxygen is in group VI, so there are six
electrons in the outer shell of each atom.
Draw a diagram of the molecule representing each bond as a
line between the atoms. One line represents one covalent
bond (a pair of electrons). In the dot-and-cross diagram, this is
represented as one dot and one cross:
Hydrogen
H
Chlorine
Cl
Nitrogen
N
N
Hydrogen
chloride
H
Cl
Carbon dioxide
O
C
O
H
N
H
H
Cl
Ammonia
H
13
Water
H
O
H
H
Ethene
H
C
C
H
H
H
Methane
H
C
H
H
Ethane
H
H
H
C
C
H
H
H
Most covalently bonded compounds have a simple molecular structure, meaning they have low melting and
boiling points due to the weak intermolecular forces which are easily broken. The covalent bonds are not broken
when it melts or boils.
Some form giant structures, in which each atom is covalently bonded to several others, with this pattern
repeating indefinitely to form a single, giant macromolecule, of unlimited size. This type of substance is best
illustrated using two allotropes of carbon. Allotropes are different structural forms of the same element.
Diamond
Graphite
14
Giant lattice, with each carbon atom forming four
strong covalent bonds to four other carbon atoms.
Layer lattice, with each carbon atom forming three
covalent bonds to three other carbon atoms, giving
hexagonal layers of atoms.
High sublimation temperature – strong covalent
bonds need to be broken, hence more energy is
required. The result is single carbon atoms with no
bonds, between them, meaning it is sublimated.
High sublimation temperature – strong covalent
bonds need to be broken, hence more energy is
required. The result is single carbon atoms with no
bonds, between them, meaning it is sublimated.
Uses: coating saw blades and drill bits – its extreme
hardness allows it to cut through any substance.
Uses: in lubricating oils, and as pencil ‘lead’ – the
weak layers allow it to slide and be transferred to
items such as paper.
Ionic compounds have higher melting points than covalent compounds as the ions are held together in a giant
structure by strong electrostatic forces. Covalent compounds have a simple molecular structure – although the
covalent bonds holding the hydrogen and oxygen atoms together are very strong, these are not broken on
melting, only the weak intermolecular forces are broken.
Metallic Bonding
This is found in metals and alloys. Each metal atom loses its outer shell electrons, becoming a positive ion. These
positive metal ions are closely-packed in a lattice. The outer shell electrons are delocalised and they are free to
move throughout the entire metal. It is the electrostatic attraction between the positive metal ions in the
lattice, and the ‘cloud’ of delocalised negative electrons which holds the metal together.
15
e-
e-
ee-
ee-
e-
Regular lattice of positive metal ions.
e-
e-
ee-
e-
e-
e-
e-
e-
e-
ee-
e-
e-
e-
e-
Delocalised ‘cloud’ of electrons, free to
move between the ions.
e-
Property
Explanation
High melting and boiling point,
high tensile strength
Ionic compounds have a giant structure, with strong electrostatic attraction
between positive metal ions and delocalised electrons, requiring large
amounts of energy to overcome it.
Malleable
The layers of metal ions can slide easily over each other. This can happen
without disrupting the metallic bonding.
Electrical conductors
There is a sea of delocalised electrons, free to carry the current.
16
Type Of
Substance
Formed
From
Metallic
Metals
Ionic
Metals and
non-metals.
Covalent
Molecular
Giant Covalent
Non-metals
Non-metals
Structure
Giant
Giant
Simple
Giant
Bonding
Description
Properties
Examples
Metallic
Lattice of positive metal ions held in a
cloud of delocalised electrons.
High melting and
boiling points;
electrical conductors
Gold, copper,
steel, silver
Ionic
Lattice of alternatingly charged ions,
held by electrostatic force.
High melting and
boiling points
Sodium chloride,
magnesium oxide
all salts.
Covalent
Help together by a shared pair of
electrons, with weak intermolecular
bonds.
Low melting and
boiling points
Water, ammonia,
diatomic
elements.
Giant three-dimensional tetrahedral
structure with no free electrons..
Sublimes at very high
temperatures; hard:
electrical insulator
Diamond
Layered hexagonal structure with
some free electrons.
Sublimes at very high
temperatures; soft;
electrical conductor
Graphite
Covalent
17
Electrolysis
Conductivity
There are two types of electric conductors: metallic conductors and electrolytes. In metallic
conductors, the electrons are delocalised, and carry the charge. These are solid metals, liquid metals,
and graphite. Electrolytes conduct electricity because of free moving ions. If an ionic substance melts
or is dissolved, it is an electrolyte. Covalently bonded acids – which are dissolved in water – are also
electrolytes, due to the disassociated ions. Electrolytes have a higher resistance than metallic
conductors, and electrolytes are decomposed by the passage of an electric current. This
decomposition is called electrolysis.
Ions
There are two ions present in the electrolyte, anions (positive) and cations (negative.) The anion is
attracted to the anode (negative electrode) and the cation is attracted to the cathode (positive
electrode). There are three main rules used to find the charge of an ion: metal ions, hydrogen ions,
and ammonium ions are ALWAYS positive. Non-metal ions are always negative. The size of the
charge is equal to the valency of the element.
Products of Electrolysis
The simplest examples of electrolysis involve a molten binary ionic substance. When this is
electrolysed, it breaks down into the two electrons from which it is made. The metal (cation) will
form at the cathode, and the non-metal (anion) will form at the anode.
If the compound is aqueous, the H+ and OH- ions complicate things. These may be discharged as
hydrogen and oxygen. When there are two ions of the same type involved in electrolysis, their
reactivity is the main factor. The less reactive element will be discharged, as its compound is less
stable.
The results of electrolysing an ionic compound in aqueous solution can be predicted by using the
rules below:


Hydrogen will be discharged at the cathode unless copper, silver or gold ions are present – in
which case they will be discharged.
Oxygen will be discharged at the anode unless chlorine, bromine or iodine are present (in
high concentrations), in which case they will be discharged as the halogen.
The electrodes are usually made from platinum or graphite, as they are unreactive.
Electrolysis and Redox
All electrolysis reactions are redox reactions. Oxidation is the loss of electrons, reduction is the gain
of electrons (OILRIG). Reduction will take place at the cathode, oxidation at the anode. The redox
reactions can be represented by two half equations, which show electrons being lost or gained,
represented with the symbol e-. When the two half equations are combined, the electrons must
cancel out, giving an ordinary equation for the whole reaction. Some examples:
Molten Zinc Chloride
As it is molten, the zinc chloride is split into molten zinc and chlorine gas. The positive Zn2+ ions are
attracted to the cathode, where they gain two electrons, and are reduced to zinc metal:
18
Cathode
Anode
Dilute Sulphuric Acid
The electrolysis of any acid produces hydrogen and oxygen. Electrolysis of compounds which
produce gases is done by a Hoffman voltameter, which collects the gases. The electrolysis occurs at
platinum electrodes, and the gas bubbles up into inverted burettes. The volume of hydrogen is
double the amount of oxygen, as the formula H2O suggests. The water is electrolysed, the acid is just
used as an electrolyte. The hydrogen ions go to the cathode and are reduced, and then pair up as
diatomic molecules:
Concentrated Aqueous Sodium Chloride
There are no copper, silver, or gold ions, so hydrogen will be produced at the cathode. The high
concentration of chloride ions mean chlorine will be produced as the anode and discharged as
chlorine gas.
19
Aqueous Copper (II) Sulphate – Graphite Electrodes
The Cu2+ ions are present in the solution, so they are attracted to the cathode and are reduced to
Copper atoms. A salmon pink coating forms on the graphite cathode. No halide ions are present, so
oxygen is produced at the anode, as the OH- ions are oxidised.
Electrolysis calculations
It is possible to calculate the amount of a substance produced in electrolysis. The quantities depend
on the total number of electrons supplied to the ions, and the charge of the ions. The charge on one
mole of electrons is one faraday, and is equal to 96,500 Coulombs.
1 mol e-  1 F = 96500 C
Moles of electrons = charge (C) / 96500
The charge which has passed through a circuit can be found using:
Q=Ixt
Where
e.g.
Q = charge (Coulombs);
I = current (Amps);
t = time (seconds)
A current of 0.5A is passed through an aqueous solution of copper(II) sulphate for 2
hours. What mass of copper metal is deposited on the cathode?
Q=Ixt
= 0.5A x 7200s
= 3600C
Moles of electrons = charge / 96500
2+
Cu
(aq)
+ 2e
= 0.0373 mol e--
Cu(s)
2 moles of electrons gives 1 mole of Cu
moles of Cu = moles of electrons / 2
= 0.0373 / 2
= 0.0187 mol Cu
Mass = moles x molar mass
= 0.0187 x 63.5
= 1.19g
Mass of copper = 1.19 g
e.g.
For how long must a current of 0.1A be passed through dilute sulphuric acid in order
to produce 240cm3 of oxygen gas?
Moles of gas = volume / 24000
--
4OH (aq)
= 240/24000
2H2O(l) + O2(g) + 4e
= 0.010 mol O2
--
4 moles of electrons give 1 mole of O2
Moles of electrons = moles of O2 x 4
= 0.010 x 4
= 0.040 mol e-
Charge = moles of e- x 96500
= 0.040 x 96500
= 3860 C
20
t=Q/I
= 3860 / 0.1
= 38600 s
time = 643 minutes (3 s.f.)
The Chlor-Alkali Industry
The electrolysis of concentrated sodium chloride solution (brine) produces three important
chemicals, chlorine, hydrogen, and sodium hydroxide. The method of manufacture is called the
chlor-alkali industry. This is carried out in a diaphragm (membrane) cell:
chlorin
e
hydrogen
concentrated
salt solution in
sodium
hydroxide and
dilute salt
nickel cathode
titanium anode
diaphrag
m
The electrodes are in separate compartments, partitioned by a permeable diaphragm. At the anode
chlorine ions are oxidised, forming chlorine gas, collected from the top:
2Cl
--
Cl2(g) + 2e
(aq)
--
At the cathode, hydrogen ions are reduced, to form hydrogen gas, which is also collected at
the top of the cell:
+
2H (aq) + 2e
--
H2(g)
The H+ ions are formed when the water disassociates, and they are removed by the
electrolysis. Due to Le Chatelier’s principle, more water is then disassociated:
H2O(l)
+
--
H (aq) + OH (aq) .
The OH- ions are not involved in the electrolysis, and accumulate in the cathode
compartment. The Na+ ions are attracted to the cathode, but are not removed by
electrolysis, so they remain in solution. The solution in the cathode compartment is now
enriched in sodium and hydroxide ions, aka sodium hydroxide solution.
The diaphragm stops the hydroxide ions from diffusing back into the anode cell where they
would react with the chlorine., as the anode cell has a higher level of solution, meaning the
21
flow is from anode – cathode. Some sodium chloride remains in the solution, so the solution
is heated till the sodium chloride crystallises out, so it can be removed.
Sodium hydroxide is used:




To purify bauxite to make alumina, so aluminium can be extracted.
To make soap
To break down wood when making paper
To manufacture chemicals
Chlorine is used to make bleach, hydrochloric acid, PVC, and to sterilise water.
Hydrogen is used to manufacture ammonia and margarine, and as an alternative power
source.
22
Section 2: Chemistry of the Elements
The Periodic Table
The elements in the Periodic Table are arranged in order of increasing atomic number. The number
of shells is given by the Period of an element. The number of electrons in the outer shell is given by
the Group. The elements in Group 0 all have full outer shells. The elements can be classified as
either metals or non-metals largely on the basis of their electrical conductivity. Metals conduct
electricity well, whereas non-metals (except graphite) are insulators.
The vast majority of elements are metals. The metals are found to the left of the Periodic Table, and
the non-metals towards the right. Some elements close to the line, example: silicon, display
properties between those of metals and non-metals, are classed as semi-metals. The metals tend to
form positive ions, as they lose their outer shell electrons. The non-metals generally gain electrons,
and form negative ions, or bond covalently with other non-metals.
A further distinction between the metals and non-metals is in the chemical behaviour of their
oxides. Metal oxides are basic – they will react with acids to form salts, and some of them dissolve in
water to give alkaline solutions:
Non-metal oxides are acidic – they react with alkalis to form salts, and dissolve in water to give acidic
solution. Important examples are carbon dioxide, which is dissolved under pressure in fizzy drinks,
and sulphur dioxide, which dissolves in rainwater to make sulphuric acid, causing acid rain.
23
24
Group I (The Alkali Metals)
Properties
They have a low density and lithium, potassium and sodium float. They are shiny when freshly cut,
but rapidly tarnish. They are soft and have low melting points. They are silvery except for Caesium
which is pale gold, and they are very reactive and are stored under oil or in argon.
Trends
On descending the group, the alkali metals become more reactive, softer, and denser. The outer
electrons become further from the nucleus, meaning the attraction becomes weaker, making it
easier to remove the outer electrons, making it more reactive.
Reactions
With oxygen, to form oxides:
4Na(s) + O2(g)
2Na2O(s)
With water, to form hydroxides (hence the name, the alkali metals) and hydrogen:
2K(s) + 2H2O(l)
2KOH(aq) + H2(g)
Lithium reacts vigorously, fizzing around on the surface of the water, and appearing to dissolve, as it
forms soluble lithium hydroxide.
Sodium reactions are the same as Lithium, but slightly more reactive, and forms soluble sodium
hydroxide. With more heat produced, and sodium’s lower melting point, the sodium becomes
molten, and forms a ball of liquid metal.
Potassium reacts violently, fizzing around very rapidly in a molten ball, and appearing to dissolve as
it forms soluble potassium hydroxide. It also burns with a purple flame.
Group VII (The Halogens)
These are reactive diatomic non-metals, which give off poisonous fumes.
Fluorine, F2
Pale yellow gas, very toxic and extremely reactive.
Chlorine, Cl2
Pale green gas, dense and toxic.
Bromine, Br2
Dense, dark red liquid, gives off red-brown vapour. Toxic and corrosive.
Iodine, I2
Dark grey solid, sublimes to give a purple vapour. Forms a brown solution in
water, and a purple solution in hexane.
25
Trends
On descending the group, they become darker, and have a higher melting/boiling point. They
become less reactive as the distance of the outer shell from the nucleus increases, making it harder
to attract a new electron.
Reactions
The halogens burn vigorously when heated with the alkali metals, to form white crystalline halide
salts. These are ionic compounds.
2Na(s) + Cl2(g)
2NaCl(s)
Any halogen can be displaced from its compound (halide) by using a more reactive halogen. Such
reactions are done in aqueous solution.
1. Chlorine + sodium bromide
Word equation:
Chlorine + sodium bromide → sodium chloride + bromine
Cl2(aq) + 2NaBr(aq)
Formula equation:
2NaCl(aq) + Br2(aq)
The colourless chlorine water reacts with the colourless sodium bromide solution, producing a
solution which is orange due to the formation of aqueous bromine.
2. Chlorine + sodium iodide
Word equation:
Chlorine + sodium iodide → sodium chloride + iodine
Cl2(aq) + 2NaI(aq)
Formula equation:
2NaCl(aq) + I2(aq)
The colourless chlorine water reacts with the colourless sodium iodide solution, producing a solution
which is brown due to the formation of aqueous iodine.
3. Bromine + sodium iodide
Word equation:
Formula equation:
Bromine + sodium iodide → sodium bromide + iodine
Br2(aq) + 2NaI(aq)
2NaBr(aq) + I2(aq)
The orange bromine water reacts with the colourless sodium iodide solution, producing a solution
which is brown due to the formation of aqueous iodine.
Group 0 – The Noble Gases
The gases of group 0 (helium, neon, argon, krypton, xenon, radon) are all colourless gases, and are
extremely unreactive – they do not form compounds with other elements and so are stable.
26
Oxygen And Oxides
Oxygen is the second-most abundant gases in the atmosphere:
GAS
ABUNDANCE
Nitrogen
78%
Oxygen
21%
Argon
0.9%
Carbon dioxide
0.04%
Water vapour
Variable
To find the percentage of oxygen in air, attach two gas syringes together with a glass tube, which
contain copper filings. Fill one gas syringe, empty the other, and heat the tube. The metal will be
oxidised, and the volume of air (at room temperature) is now 21% less.
copper
turnings
gas syringe
glass tube
Preparation Of Oxygen
Oxygen can be prepared by the decomposition of hydrogen peroxide. A manganese dioxide catalyst
is needed:
2H2O2(aq)
M nO2
2H2O(l) + O2(g)
27
hydrogen
peroxide
oxygen
gas
manganese
dioxide
beehive
shelf
Formation Of Oxides
Magnesium burns vigorously in air, with a brilliant white flame, to form magnesium oxide, a white
powder. Magnesium oxide is a metal oxide, and so is basic. It is slightly soluble in water, giving a
slightly alkaline solution of magnesium hydroxide (pH 10). It will react with acids, to form a salt and
water:
2Mg(s) + O2(g)
MgO(s) + H2O(l)
2MgO(s)
Mg(OH)2(aq)
MgO(s) + 2HCl(aq)
MgCl2(aq) + H2O(l)
Carbon burns steadily if heated in air, to form colourless carbon dioxide gas. If the supply of oxygen
is limited, some toxic carbon monoxide is also produced. Carbon dioxide is a non-metal oxide, and so
is acidic. It is slightly soluble in water, giving a weakly acidic solution of carbonic acid (pH 6).
C(s) + O2(g)
CO2(g) + H2O(l)
CO2(g)
H2CO3(aq)
Sulphur is a yellow solid, which burns in air with a bright blue flame, to form white fumes of sulphur
dioxide. It dissolves readily in water to form an acidic solution of sulphurous acid, H2SO3.
S(s) + O2(g)
SO2(g) + H2O(l)
SO2(g)
H2SO3(aq)
28
Carbon Dioxide
To prepare carbon dioxide add hydrochloric acid to marble chips (calcium carbonate).
CaCO3(s) + 2HCl(aq)
CaCl2(aq) + H2O(l)
Carbon dioxide is also formed in the thermal decomposition of metal carbonates. Green copper (II)
carbonate becomes black copper (I) oxide.
CuCO3(s)
heat
CuO(s) + CO2(g)
Carbon dioxide is used in fire extinguishers – it is unreactive, and denser than air, so gathers around
the fire, depriving it of oxygen. It is especially useful for electrical fires, when it is dangerous to use
water. Carbon dioxide is also dissolved, under pressure, in fizzy drinks. When the bottle is opened,
the pressure is released, and the carbon dioxide bubbles out of solution.
Acid Rain
Sulphur dioxide is formed when coal is burned in power stations. This dissolves in rainwater, forming
acid rain. This damages trees, kills fish in rivers and lakes, and damages limestone buildings. Similar
pollutants include nitrogen oxides (NO, NO2), formed when nitrogen in the air reacts with oxygen in
hot car engines. It can dissolve in rain water to form nitrous and nitric acids, which also contribute to
acid rain.
Hydrogen and Water
Hydrogen is the least dense but most abundant element in the universe.
Test for Hydrogen
To test for the presence of hydrogen gas, a sample of the gas is collected in a test tube, and a lit
splint is applied. Hydrogen ignites with a ‘squeaky pop’ and a blue flame will be seen.
Combustion of Hydrogen
When hydrogen burns it forms water (hydrogen oxide). Because the flame is hot, this will form as
water vapour.
2H2(g) + O2(g)
H2O(g)
Test for Water
White anhydrous copper (II) sulphate turns blue in the presence water and pure water boils at
exactly 100C.
29
Reactivity Series
A more reactive metal will displace a less reactive metal from a compound, and this usually occurs in
solution, but can occur in solids, if the metal is heated.
Copper metal + magnesium chloride solution: no reaction – copper is less reactive than magnesium.
Magnesium metal + copper sulphate solution: Magnesium is more reactive than copper, so can
displace it. A pink coating of copper forms on the surface of the magnesium, and the blue copper
sulphate solution slowly turns colourless, as it is converted to magnesium sulphate:
Mg(s) + CuSO4(aq)
MgSO4(aq) + Cu(s)
Displacement reactions can be used to establish a reactivity series for common metals:
METAL
REACTION WITH ACID
REACTION WITH WATER
Potassium
Violent - hydrogen produced is ignited
Sodium
Very vigorous – shoots around on the
surface of the water
Dangerously violent
Vigorous – fizzes around on the surface
of the water
Lithium
Calcium
Fizzes vigorously, forming hydrogen gas
Fizzes rapidly, forming hydrogen gas
Magnesium
Fizzes rapidly, forming hydrogen gas
Reacts very slowly, forming hydrogen gas
30
Zinc
Fizzes steadily, forming hydrogen gas
Reacts very slowly, forming hydrogen gas
Iron
Fizzes very slowly, forming hydrogen gas
Rusts slowly, but only if oxygen is
present
Copper
No reaction
No reaction
Reduction and Oxidation
Any reaction which involves oxygen being transferred between two other elements is a redox
reaction. A redox reaction is one in which both oxidation and reduction occurs.
Oxidation is the gain of oxygen and the loss of electrons.
Reduction is the loss of oxygen and the gain of electrons.
A substance which provides oxygen and takes electrons is the oxidising agent.
A substance which removes oxygen and provides electrons is the reducing agent.
Rusting
Rusting is the oxidation of iron, to form hydrated iron(III) oxide.
Only iron (and alloys of iron, such as steel) can rust. Other metals corrode. Rusting requires oxygen
and water. Rusting can be prevented by placing grease or oil on the iron which repels water. This is
applied regularly and is messy, normally for machinery. Paint and plastic coatings form a protective
layer, but once scratched the metal below starts to rust. Galvanising involves coating the iron in a
layer of zinc. If scratched the zinc acts sacrificially
Sacrificial protection involves placing a more reactive metal (such as magnesium or zinc) in contact
with the iron object. Because this sacrificial metal is more reactive, it will corrode in preference to
the iron – it is ‘sacrificed’ to protect the iron from rusting. This is used mainly on ships. Blocks of zinc
are bolted to the hull at regular intervals. These slowly oxidise, and protect the ship from rusting.
They must be regularly replaced, when they become corroded.
Test for Ions
+
Li
Na+
K+
Ca2+
NH4+
Cu2+
Fe2+
Fe3+
Cations:
Flame test – add conc. HCl to the
compound, dip wire loop in the
paste and hold in a Bunsen
burner blue flame.
Add sodium hydroxide and warm.
Add sodium hydroxide solution
Red colour
Persistent orange colour
Lilac colour
Brick red colour
Ammonia gas is produced which has a pungent smell and turns red litmus
blue.
Pale blue precipitate of Cu(OH)2
Dirty green precipitate of Fe(OH)2
Rusty brown precipitate of Fe(OH)3
31
-
Cl
BrISO42CO32-
Anions:
Add nitric acid and then silver (I) nitrate solution.
Add some hydrochloric acid and then barium chloride solution.
Add hydrochloric acid and bubble gas through limewater.
Gases:
Ammonia NH3
Carbon Dioxide CO2
Chlorine Cl2
Hydrogen H2
Oxygen O2
Water H2O
Damp red litmus paper
Bubble through limewater
Expose to damp blue litmus
Collect gas in a test tube and apply a lit splint
Collect gas in a test tube, apply a glowing splint
Pass through anhydrous copper sulphate crystals, or
test with anhydrous cobalt chloride paper
32
White precipitate of silver chloride.
Cream precipitate of silver iodide.
Yellow precipitate of silver bromide.
White precipitate of Barium Sulphate.
Carbon Dioxide gas is produced which
turns lime water cloudy.
Turns damp red litmus blue.
It turns cloudy.
The paper is bleached white.
Squeaky pop and blue flame.
Splint relights.
Copper sulphate – white to blue, cobalt
chloride paper – blue to pink.
Section 3 – Organic Chemistry
The Alkanes
The alkanes are a homologous series of hydrocarbons. This means they have the same general
formula, same chemical properties and follow a trend in physical properties. The general formula of
alkanes is CnH2n+2.
Naming Hydrocarbons
When naming an organic molecule, look for the longest continuous chain of carbon atoms, which
determines the root:
Number of carbon atoms Root
1
METH2
ETH3
PROP4
BUT5
PENT6
HEX7
HEPT8
OCTThe root is also given a suffix to identify the series to which the compound belongs, and all alkanes
have names ending in –ane.
Bromination
Alkanes will react with bromine if they exposed to ultraviolet light. This is a substitution reaction,
where one hydrogen atom is replaced with one bromine atom. Hydrogen Bromide is also produced:
CH4 + Br2  CH3Br + HBr
33
Isomerism
Isomers are molecules with the same molecular formula, but different structural formulae.
The formulae shown above are displayed formulae – they show every bond and every atom. One
line connecting the atoms represents one covalent bond.
34
Combustion
Alkanes are not very reactive, and so they are mostly used as fuels. When a hydrocarbon is
completely burned, it forms carbon dioxide and water.
C3H8 + 5O2
3CO2 + 4H2O
If insufficient oxygen is available, incomplete combustion occurs, and this produces the poisonous
carbon monoxide:
2C3H8 + 7O2
6CO + 8H2O
If there is even less oxygen, carbon is produced, leading to a sooty, yellow flame.
Alkenes
The alkenes are another homologous series of hydrocarbons, with the general formula CnH2n.
Addition Reactions
Alkanes are saturated hydrocarbons, as they only contain single bonds. This means they are
unreactive, and are used as fuels. Alkenes are unsaturated, as they contain double bonds. They are
35
more reactive as the double bond can be broken, leading to an addition reaction. Bromination is an
addition reaction:
H
H
C
C
H
Br
+
Br
Br
H
H
H
C
C
H
H
Br
This can be used to distinguish between alkanes and alkenes. When bromine water is added to an
alkene it rapidly decolourises, but when added to an alkane the solution stays orange.
Ethanol
Ethanol is part of the alcohols homologous series:
OH H
H
C C
H
H H
There are two main methods for its production:
Fermentation:
This is the conversion of glucose to ethanol and carbon dioxide which is done by anaerobic
respiration of yeast. A valve must be used to allow the carbon dioxide to escape, without allowing air
to enter the vessel. This is done at 40 degrees, which is the optimum temperature for the zymase
enzyme. This produces alcohol of 15% concentration, and is the only method used to produce
alcoholic drinks, as the flavour of the fruit juices fermented provides the flavour of the drink. The
alcohol is then distilled to produce more concentrated drinks, such as vodka. This is a batch process,
and so takes much longer.
C6H12O6(aq)
zy mase
2C2H5OH(aq) + 2CO2(g)
Addition of Steam to Ethene
This is carried out at a high temperature and pressure (300C, 60-70 atm), and is passed over a
phosphoric acid catalyst. This method is a continuous process, and is used in oil rich countries. This
produces purer alcohol.
C2H4 + H2O  C2H5OH
Dehydration of Ethanol
The above reaction (addition of steam to ethane) can be reversed. This is an elimination reaction.
Some mineral wool is soaked in ethanol and placed at the bottom of a horizontal boiling tube. Some
aluminium oxide is placed in the middle of the tube (catalyst) and a bung is inserted. The catalyst
36
and the ethanol are heated, and the vaporised ethanol passes over the catalyst, and breaks down to
form ethene and steam. The ethene can be collected over water.
C2H5OH  C2H4 + H2O
37
Section 4 – Physical Chemistry
Acids, Bases and Salts
Indicators are substances which determine whether a solution is acidic or alkaline.
Indicator
Colour in Acid
Colour in Alkali
Litmus
Red
Blue
Phenolphthalein
Colourless
Pink
Methyl Orange
Red
Yellow
There are two types of litmus, red litmus which turns blue in alkalis, and blue litmus which turns red
in acids.
Universal indicator shows how acidic or alkaline something is, and it goes from red (strong acid) to
green (neutral) to purple (strong alkali). Acidity is measured using the pH scale, which goes from 1
(acidic) to 7 (neutral) to 14 (alkaline).A pH probe can also be used.
Reactions of Acids
When acids react, they form salts:
ACID
FORMULA
Hydrochloric acid
Sulphuric acid
Nitric acid
SALT FORMED
HCl (aq)
H2SO4 (aq)
HNO3 (aq)
chloride salts
sulphate salts
nitrate salts
(Cl, valency 1)
(SO4, valency 2)
(NO3, valency 1)
Acid + Metal → Salt + Hydrogen
Metals which are below hydrogen in the reactivity series do not react with dilute acids.
Magnesium – reacts rapidly with dilute acids to make colourless solutions of magnesium
salts.
Zinc – reacts readily to make colourless salts.
Aluminium – does not react due to impervious aluminium oxide coating.
Iron – reacts slowly to form pale solutions of Iron (II) chloride or Iron (III) sulphate.
Acid + Metal Oxide → Salt + Water
Acid + Metal Carbonate → Salt + Water + Carbon Dioxide
Acid + Metal Hydroxide → Salt + Water
ACID +
METAL
PRODUCTS
SALT + HYDROGEN
38
METAL OXIDE
METAL CARBONATE
METAL HYDROXIDE
SALT + WATER
SALT + CARBON DIOXIDE + WATER
SALT + WATER
Definitions
Acids are substances which dissociate in water and form an H+(aq) ion. Acids are proton donors. Acids
only behave like acids when dissolved in water. Hydrogen chloride dissolves in water to form
hydrochloric acid, in methylbenzene is does not dissociate, and so does not behave like an acid.
Bases are substances which react with acids and form a salt and water. This includes metal oxides,
metal carbonates and metal hydroxides.
Alkalis are bases which dissolve in water to form the OH-(aq) ion. Most alkalis are metal
hydroxides, but ammonia is also an alkali.
NH4+(aq) + OH--(aq)
NH3(aq) + H2O(l)
Salts are formed when the hydrogen in the acid is replaced by a different positive ion –
usually a metal ion.
Neutralisation is the reaction of an acid with a base. A salt and water is always produced.
Salt Preparation
Rules for Solubility:
 All sodium, potassium and ammonium compounds are soluble.
 All chlorides except silver chloride are soluble.
 All sulphates are soluble, except for barium and calcium sulphate.
 All hydroxides are insoluble, except sodium, potassium and ammonium.
 All nitrates are soluble.
Precipitation
This is used to prepare insoluble salts. Two solutions are mixed, each one containing one of the
necessary ions. On mixing, the ions combine and form the salt, which precipitates. This is filtered off,
rinsed and left to dry. For example, to make Barium Sulphate, two solutions are needed, one of
Barium Nitrate, one of Sodium Sulphate (both are soluble). The solutions are mixed, the precipitate
filtered, rinsed and left to dry.
Equation:
Ba(NO3)2(aq) + Na2SO4(aq)
Ionic equation:
Ba2+(aq) + SO42--(aq)
BaSO4(s) + 2NaNO3(aq)
BaSO4(s)
Note that the ionic equation for the formation of any salt follows this simple pattern, of two
aqueous ions combining to make the precipitate, e.g.:
Ag+(aq) + Cl --(aq)
AgCl(s)
39
Titration
This is used to prepare soluble salts from an acid and a soluble base. It is normally used to prepare
sodium, potassium and ammonium salts. A known volume of acid is measured into a conical flask,
using a pipette, and some indicator is added. The alkali is placed in a burette and slowly added into
the indicator shows the solution is now neutral. The amount of alkali is noted, and the experiment is
repeated, until concordant titres are found. The titre is then added again, but this time no indicator
is added. The solution is then boiled till it is saturated, and then slowly heated till it crystallises.
Reaction of an Acid with an Insoluble Base
This method is used to prepare soluble salts, from an insoluble base. The acid is placed in a beaker
and warmed with a Bunsen burner. The insoluble metal oxide or carbonate is added and stirred. It is
added until it stops disappearing, and is in excess – the reaction is complete. It can now be filtered
off, leaving the pure salt solution behind. The solution is now gently warmed until it is saturated, and
then it is left to crystallise.
t e at
o e
e.g. barium sulphate
e o tion
o t e
a ro riate
meta nitrate
an o i m
a t.
e.g. potassium chloride
e o tion
o t e
a ro riate
a i an
a ai a
or
O
e.g. use BA
nitrate + sodium
LP AT
solutions
S
it a
o i m
ota i m or
ammoni m
at
S
e.g. use
D C L C
acid +
P TA
hydroxide
O
ti a o e
at
t not a
o i m
ota i m or
ammoni m
at
e.g. nickel
nitrate
40
et e
a ro riate
a i an
meta
ar onate or
o i e
e.g. use T C
acid + C L
carbonate or
oxide
Energetics
All chemical reactions result in a conversion of energy. Chemical reactions which release heat energy
are exothermic, and those which take in heat energy are endothermic. Due to the conservation of
energy, exothermic reactions result in the chemical energy of the products decreasing, and
endothermic reactions have the chemical energy of the products increasing.
Enthalpy is the chemical energy change, and is given the symbol Δ . Δ is always given in terms of
the chemicals, and not the surroundings. This means exothermic reactions have negative enthalpy
changes, and endothermic reactions have positive ones.
Measuring Enthalpy Changes
It is possible to measure the enthalpy change by using a reaction to heat or cool a known mass of
water. The enthalpy change can be measured by using the formula:
Δ = m c ΔT
Where: Δ = energy supplied by water (joules), m = mass of water (grams), c = specific heat capacity
of water (4.2 J/g/C), and ΔT = the change in temperature of the water (C). Since an increase in the
temperature of the water means a decrease in the energy of the chemicals, to find the enthalpy
change of the reaction, use:
ΔH = - m c ΔT
41
If the reaction occurs in solution, the mass of the solution is used.
Enthalpy change is commonly given per mole, and the molar enthalpy change is given in kilojoules
per mole.
e.g.
100g of water were placed in a copper calorimeter above a fuel burner containing hexane,
C6H14. Burning the hexane caused the temperature of the water to rise from 18 to 44. The mass of
the burner decreased from 98.30g to 97.87g. What is the enthalpy of combustion of 1 mole of
hexane?
Formula mass of hexane
Temperature rise
= 44 – 18
= 6 x 12 + 14 x 1
= 86
= 26C
Mass of hexane burned = 98.30 – 97.87 = 0.43
So
Moles of hexane burned
= mass / molar mass
Energy supplied to water
= m c ΔT
Enthalpy change
= - m c ΔT
Enthalpy change per mol
ΔH
= 0.43 / 86
= 0.005 mol
= 100g x 4.2 J/g/C x 26C
= -10920 J
= -10920J / 0.005 mol
= -2184000 J/mol
= -2184 kJ/mol
Calculating Enthalpy Changes
Making bonds releases energy and breaking them requires energy.
42
= 10920 J
To calculate the enthalpy change for any reaction:
1. Write a balanced equation for the reaction.
2. Find the total bond energy of every bond in the reactant molecules, remembering to take
into account the balancing numbers (it may help to draw out the molecules).
3. Find the total bond energy of every bond in the product molecules, remembering to take
into account the balancing numbers.
4. The overall enthalpy change is given by:
ΔH = bond energy of reactants – bond energy of products
What is the enthalpy change when one mole of methane is burned?
e.g.
CH4 + 2O2
CO2 + 2H2O
H
H
C
O
H
+
O
O
O
O
C
+
O
O
H
H
O
H
H
BOND
BOND DISSOCIATION ENTHALPY
(kJ/mol)
C
H
412
O
O
496
Total bond energy for reactants = 4 (C  H) + 2 (O = O)
H
C
O
743
O
H
463
= 4 x 412 + 2 x 496
= 2640 kJ/mol
Total bond energy for products
= 2 x (C = O) + 4 x (O  H) = 2 x 743 + 4 x 463
= 3338 kJ/mol
Reactant – product energies
= 2640 – 3338
= -698 kJ/mol
ΔH = -698 kJ/mol
The negative value shows that the reaction is exothermic.
43
In some cases, the actual change is different to the theoretical one, and if the obtained value
is lower, there is either some heat loss, or incomplete combustion.
Rates of Reaction
The rate of reaction is the change in amount of reactant (or product) per unit time.
The rate of reaction may be monitored in several ways:





The volume of gas produced in a gas syringe
The loss of mass due to gas escape
pH change
temperature change
colour change
To measure the above, it is crucial to have a unit time.
Factors that affect the rate of reaction
The four factors which affect the rate of a chemical reaction:




concentration/pressure of liquid/gaseous reactants
surface area of solid reactants
temperature
catalysts
For a reaction to take place the reactant particles need to collide with energy greater than the activation
energy. Increasing the first two factors increases the rate at which they collide, so there are more frequent
collisions. Increasing the third factor increases the success rate of the collisions (as each particle has energy
closer to the activation energy). The addition of the catalyst lowers the activation energy of the reaction. Thus
increasing all four factors ensures there are more frequent, more successful collisions. The catalyst increases
the rate of a chemical reaction without being used up in the overall reaction. The catalyst provides an
alternative pathway for a reaction, with a lower activation energy. Catalysts are specific – certain ones catalyse
certain reactions, but not others. The catalyst is not a reactant, and should not be written as part of the
equation. It is written above the arrow.
Measuring the rate
44
Equilibria
Most chemical reactions will proceed to completion once started, but others are reversible, and can
occur in both the backward and forwards directions. These are dented by the
arrow.
These reactions do not reach completion, they reach an equilibrium, where the rates of the
forwards and the rate of the backwards reactions are the same. Chemical equilibria are
dynamic equilibria in that the change is constantly occurring, but as both reactions are
occurring at the same rate, the amount of products and reactants remains constant.
Le Chatelier’s Principle
This principle is a way of predicting the outcome of changing an equilibrium mixture, and it states:
‘When a change of conditions is imposed on a system in equilibrium, the position of equilibrium will
shift in a direction so as to oppose the change in conditions.’
Change Imposed
Increase in temperature
Decrease in temperature
Increase in pressure
Decrease in pressure
Adding a reactant
Removing a reactant
Using a catalyst
Effect on Equilibrium
Moves to the endothermic reaction
Moves to the exothermic reaction
Moves to the side with fewer moles of gas
Moves to the side with more moles of gas
Moves to the opposite side
Moves towards the side involving that reactant
No effect – but increases the rate of both.
45
Section 5 – Chemistry in Society
Extraction of metals
Metals are extracted from ores. Ores are rocks which contain a sufficient amount of a peculiar
metal.
The extraction of metal can be divided into three categories:



Least reactive – may be found unreacted. They still need to be removed from their ores, but
not chemically broken down. This includes gold and platinum.
Moderately reactive metals – they are below carbon in the reactivity series. Extracted in a
displacement reaction with carbon. This includes lead, copper and iron.
Metals which are more reactive than carbon need to be extracted by electrolysis. This
includes aluminium.
The Blast Furnace
Three things are added to the top of the blast furnace:
Iron ore – often haematite, Fe2O3
Coke – mainly consists of carbon. Made by purifying coal.
Limestone – calcium carbonate, CaCO3
Hot air is fed through the bottom. The coke burns in this air:
This reaction is exothermic, and heats up the furnace. The carbon
dioxide then reacts with more coke to form carbon monoxide:
The carbon monoxide reacts with the haematite, to form molten
iron and carbon dioxide.
The calcium carbonate thermally decomposes to form calcium oxide
and carbon dioxide:
The Calcium oxide reacts with the silicon oxide (an impurity in
haematite) to make calcium silicate, which forms slag, floating
above the molten iron. Slag is used for road surfacing and making
cement:
In the process, the iron oxide is reduced, and the carbon monoxide is oxidised.
46
Aluminium Extraction
The main ore of aluminium is bauxite, which consists of alumina (aluminium oxide). Aluminium is
above Carbon in the reactivity series, so it must be displaced with electrolysis.
Aluminium Oxide has a high melting point, and is insoluble in water, and so is dissolved in molten
cryolite, so it can be electrolysed. The cryolite is at 900 degrees, whereas molten aluminium oxide
would be at 2000 degrees. This means using cryolite is cheaper and easier. The cathode is the lining
of the cell, and the anodes are the large blocks dipped in the electrolyte. The electrodes are made of
graphite. The molten aluminium, after forming on the cathode, sinks to the bottom, and is tapped
off. The hot oxygen reacts with the electrodes, forming carbon dioxide, and so they often need
replacing. This process is expensive due to the electricity needed.
The aluminium oxide consists of aluminium ions (Al3+) and oxide ions (O2-). The aluminium ions
(cations) are attracted to the cathode, where they are reduced to aluminium atoms:
3+
Al + 3e
--
Al
The oxide ions (anions) are attracted to the anodes, where they are oxidised to oxygen gas:
2--
2O
O2 + 4e
--
Uses of Iron
It is very cheap and abundant, and therefore used in lots of things:


Pig iron which is straight from the blast furnace can be moulded. If it is remelted and
remoulded, it is cast iron and is very impure (4% carbon). This is very hard and brittle, and
used in manhole covers and guttering.
Steel which has 1.5% carbon is called high-carbon steel, and is very hard but brittle. It is used
in drill bits.
47



Mild steel contains 0.25% iron, and is stronger and harder. This is used in car bodies, girders
and ships.
Pure iron, also called wrought iron, is used as decorative railings, as it is soft and malleable.
If the iron is alloyed with chromium and nickel, stainless steel is formed. This protects the
iron from rust and corrosion. It is used in cutlery, kitchen sinks, and in gardening tools.
Uses of Aluminium
Aluminium is not very strong, and is used as an alloy, due to its low density, corrosion resistance and
as it is a good conductor. It is used in aircraft (low density, corrosion resistance, strength when
alloyed), saucepans (low density, corrosion resistance, high thermal conductivity); high-voltage
power cables (low density; corrosion resistance; high electrical conductivity – the cables have a steel
core to increase strength).
Crude Oil
Crude Oil Processing
Crude Oil is a mixture of hydrocarbons. These are simple covalently bonded molecules, with weak
intermolecular forces.
Fractional distillation is used to separate crude oil into useful components (fractions). The more
carbon atoms the molecule contains, the stronger the intermolecular forces and the higher its
boiling point.
The crude oil is heated to around 350°C in a furnace, and the compounds in the oil enter the bottom
of the fractional distillation tower as a mixture of gases. They rise up the tower, and as they do so,
they cool down. When a molecule reaches a temperature below its boiling point, it condenses, turns
to a liquid and falls down the tower, to be collected in a series of trays. The compounds with a
similar boiling point condense in the same area, and are removed as part of the same fraction. The
compounds with the lowest boiling points reach the top of the tower without condensing, and are
removed as part of the refinery gases section.
There are many types of hydrocarbon – straight chain, branched chain, and ring molecules – but the
carbon always forms four bonds, and the hydrogen one. Crude oil consists of mostly alkanes.
Cracking
If alkane vapour is passed over a heated catalyst, it thermally decomposes. The alkane is broken
down into two smaller molecules, an alkane and an alkene, or an alkene and hydrogen. If an alkane
and an alkene are produced, the alkene is usually smaller. The hydrogen can be burned as a fuel to
power the refinery. The catalyst is a mixture of silicon dioxide and aluminium oxide, and this is done
as 500 degrees centigrade. Cracking is very useful, as it breaks down the longer, less useful alkanes
found in crude oil such as diesel and naptha into shorter alkanes like petrol, which is in high demand,
and into alkenes, which are more reactive than alkanes, and are can be used to make plastics.
C8H18
C8H18
catalyst
HEAT
C6H14 + C2H4
C8H16 + H2
48
Polymerisation
Addition
Alkenes can add to other alkenes, and the monomers join to form polymers.
49
H
H
H
H
H
H
C
C
C
C
C
C
H
H
H
H
H
H
H
H
H
H
H
H
C
C
C
C
C
C
H
H
H
H
H
H
High temperature and pressure
Catalyst
The chain above shows three repeat units, with one repeat unit being a length of polymer
made from a single monomer.
Polymers:
Name
Monomer
Repeat Unit
Notes, Uses, and
Properties
Also called polythene,
and is used to make
plastic bags, carrier
bottles and packaging.
Poly(ethene)
Ethene
Poly(propene)
Propene
Commonly called
polypropylene, it is
used to make ropes,
crates, and many
other items.
Poly(chloroethene)
Chloroethene
Commonly called PVC,
it is the strong rigid
material used to make
doors, window frames
and drainpipes.
Drawing Polymers
To draw a polymer, draw out a monomer so that the C=C bond is horizontal, and all the other groups
are vertical. Then break the C=C double bond, and draw the new bonds at the side, then add the
brackets.
50
H3C
CH3
C
C
redraw
H
CH3
CH3
CH3
C
C
H
CH3
polymerise
2-methylbut-2-ene
CH3
CH3
C
C
H
CH3
poly(2-methylbut-2-ene)
To find the monomer from which a polymer is made, isolate two adjacent carbon atoms, and then
replace the C=C double bond.
CH3
H
CH3
H
CH3
H
C
C
C
C
C
C
Cl
H
Cl
H
Cl
H
is made from
CH3
H
C
C
Cl
H
The name of a polymer is always poly followed by the name of monomer in brackets e.g.
ethene goes to poly(ethene)
Condensation Polymerisation
A condensation reaction can be used to join many monomers to form a polymer, and this reaction
differs from addition reactions as different monomers are used.
If one monomer contains an –OH group, and the other contains an –H, these combine to form water.
The remainders of the monomers join to form the condensation polymer.
Polyamides
One of the most important condensation polymers is nylon, which is used in ropes, carpets and clothing. Nylon
is a polyamide, which is made from a carboxylic acid (-COOH) and an amine (-NH2). Each of the monomers is
double ended, which lets them join together alternately by forming an amide link (-CONH-) to make the
polymer. This can be seen below, with the central portion of the monomers being represented by a block:
O
O
C
HO
H
C
H
N
OH
H
N
O
O
H
H
C
C
N
N
H
+
O
H
H
Polyesters
Polyesters are condensation polymers, which are like polyamides, but are formed from a carboxylic
acid (-COOH), and an alcohol (-OH), and both monomers are double ended. This results in the
formation of an ester link (-COO-), with water being produced as a side product.
51
O
O
C
C
HO
O
OH
O
H
O
O
C
C
O
O
H
+
O
H
H
Polyesters such as terylene are used to manufacture material for clothing.
O
HOOC
C6H4
COOH
+
HO
C2H4
OH
C
O
C6 H4
+
C
O
C2H4
O
H
H
Chemical Manufacture
The Haber Process
This is the combination of hydrogen and nitrogen to make ammonia.
N2(g) + 3H2(g)
2NH3(g)
The hydrogen is obtained by reacted methane with steam, or by cracking hydrocarbons:
CH4(g) + H2O(g)
C10H22
heat
cataly st
Ni
CO(g) + 3H2(g)
C10H20 + H2
The nitrogen is obtained from the air.
The hydrogen and nitrogen are fed into the reaction vessel in the ratio of 3:1 respectively. The
reaction is reversible:
N2(g) + 3H2(g)
2NH3(g)
ΔH = -92kJ/mol
Temperature:
450°C
Pressure:
200 atmospheres
Catalyst:
Finely divided iron
The forward reaction is exothermic, so the temperature is low enough to give a reasonable yield, but high
enough to produce that yield at an acceptable rate. The iron catalyst increases the rate of the reactions.
Increasing the pressure increases the rate of the forward reaction, and also increases the rate, but a too high
pressure is dangerous and expensive, meaning a pressure of 200 atmospheres is used. The yield of ammonia is
15%, and the mixtures of gases which leave the vessel are cooled, until the ammonia condenses, and is
52
O
removed as a liquid. The hydrogen and oxygen are recycled into the reaction vessel. The ammonia produced is
used to manufacture fertilisers, and to make nitric acid.
Nitrogen from fractional
distillation of liquid air
450C
200 atm
iron catalyst
H2 and N2
Cooling
NH3
Liquid
NH3
Unreacted gases
Hydrogen from natural gas
and steam
The Contact Process
This is the manufacture of Sulphuric Acid.
Firstly, sulphur is burned in oxygen to make sulphur dioxide.
S(s) + O2(g)
SO2(g)
The sulphur dioxide is then reacted with more oxygen to produce sulphur trioxide:
2SO2(g) + O2(g)
2SO3(g)
This is a reversible reaction, and the forward reaction is exothermic, so the temperature is a
compromise between a higher yield and a fast rate of reaction. Increasing the pressure
would increase the yield and increase the rate, but the yield is already very high at one
atmosphere, so the pressure is kept low. The catalyst increases the rate of reaction.
Temperature:
450C
Pressure:
1-2 atm
Catalyst:
Vanadium(V) oxide – V2O5
The sulphur trioxide will dissolve in water to form sulphuric acid, but this is incredibly
exothermic, and will vaporise the solution. Instead, the sulphur trioxide is dissolved in
concentrated sulphuric acid, to give oleum. Water is then added to the oleum, producing
concentrated sulphuric acid.
H2SO4(l) + SO3(g)
H2S2O7(l)
H2S2O7(l) + H2O(l)
2H2SO4(l)
Sulphuric acid is used in the manufacture of fertilisers, manufacture of detergents, and
manufacture of paints.
53