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©2014 Marcus Ng ChemistryNotesforSecondarySchool O-LevelsPure Chemistry
2014 Chemistry Notes
For Secondary School
Pure Chemistry
O-Levels
By Marcus Ng

Chapter 1
Experimental Chemistry
1.1 Measurements
Physical Quantity SI Unit Apparatus Accuracy
Time Second (s)
Digital Stopwatch 0.01 s
Analog Stopwatch 0.1s
Temperature Kelvin (K)
Mercury Thermometer 0.01 K
Alcohol Thermometer 0.01 K
Mass Kilogram (kg)
Electronic Balance
Beam Balance
Length Meter (m)
Ruler 0.1 cm (1mm)
Vernier Calipers 0.01 cm (0.1mm)
Micrometer 0.001 cm (0.01mm)
Volume Cubic Meter (m3
)
Beaker
Measuring Cylinder 1 cm3
(1 ml)
Pipette* 0.1 cm3
(0.1 ml)
Burette 0.1 cm3
(0.1 ml)
How to read a Vernier Caliper reading & A Micrometer reading
Important Points to remember:
1. When measuring Time: Digital Stopwatch is more accurate than Analog Stopwatch
2. When measuring Length: Micrometer is more accurate than Vernier Calipers, which
are both more accurate than a Ruler
3. When measuring Volumes: Pipettes are only used for specific volumes
(10 cm3
, 25 cm3
or 50 cm3
)
4. When measuring Volumes: Burettes are the most accurate, followed by a Measuring
Cylinder and lastly a Beaker
5. When measuring Temperature: Maximum upper limit for Alcohol Thermometer is
351.15K (78 0
C)
6. When measuring Temperature: Mercury Thermometers are more expensive and toxic
then Alcohol Thermometers.

1.2 Separation Techniques
Technique Purpose
Filtration Used to separate an insoluble solid from a liquid
Crystallization Used to separate a soluble solid from a liquid
Distillation Used to separate a liquid from a soluble solid
Fractional Distillation Used to separate a liquid from a mixture of Miscible Liquids*
Separating Funnel Used to separate a liquid from a mixture of immiscible liquids*
Sublimation Used to separate a sublimable solid from a mixture of solids
Magnetic Attraction Used to separate a Solid that can be magnetized
*Miscible Liquids refer to liquids that can be completely mixed
Filtration
Crystallization

Distillation
Fractional Distillation

Separating Funnel
Sublimation

1.3 Collection of Gases
Properties of Gases
Gas Solubility Density
Ammonia Soluble in Water Less dense than air
Argon Insoluble in Water More dense than air
Carbon Dioxide Soluble in Water More dense than air
Carbon Monoxide Insoluble in Water Less dense than air
Chlorine Soluble in Water More dense than air
Helium Insoluble in Water Less dense than air
Hydrogen Insoluble in Water Less dense than air
Hydrogen Bromide Soluble in Water More dense than air
Hydrogen Chloride Soluble in Water More dense than air
Methane Insoluble in Water Less dense than air
Oxygen Insoluble in Water Similar density to air
Neon Insoluble in Water Less dense than air
Nitrogen Insoluble in Water Similar density to air
Nitrogen Dioxide Soluble in Water More dense than air
Sulphur Dioxide Soluble in Water More dense than air

1.4 Purity of a Substance
Important Points/Concepts to remember:
1. A Pure Substance melts and boils at a fixed and constant temperature
2. Impurities decreases the melting point of a substance
3. Impurities increases the boiling point of a substance
Methods to check the purity of a substance
1. Melting Point Determination
2. Boiling Point Determination
3. Chromatography

Chromatography
1. Chromatography can be used to determine the purity of a substance
2. Chromatography can be used to identify the substance
3. Chromatography can be used to separate components of a substance with different
solubilities in the same solvent and identify them.
4. There are 2 types of Paper Chromatography: Ascending and Descending
5. There can be 3 types of results (chromatograms that can be developed)
a. Only one spot is seen - A Pure Substance (one solute in sample)
b. More than one spot is seen - A Mixture (more than one solute in sample)
c. No spots are seen - No soluble* solute in mixture (*in the solvent used)
 Note: Some compounds are colourless and thus a locating agent need
to be used. An example of a locating agent is Ninhydrin, used for
locating amino acids.
6. Rf Value (Retention Factor) =
𝑑𝑖𝑠𝑡𝑎𝑛𝑐𝑒 𝑚𝑜𝑣𝑒𝑑 𝑏𝑦 𝑡ℎ𝑒 𝑠𝑢𝑏𝑠𝑡𝑎𝑛𝑐𝑒
𝑑𝑖𝑠𝑡𝑎𝑛𝑐𝑒 𝑏𝑒𝑡𝑤𝑒𝑒𝑛 𝑡ℎ𝑒 𝑠𝑡𝑎𝑟𝑡 & 𝑠𝑜𝑙𝑣𝑒𝑛𝑡 𝑙𝑖𝑛𝑒
solvent line
start line

1.5 Tests for Cations
Cation Symbol
Test
Add dilute sodium
hydroxide solution to a
solution of the substance.
Add dilute ammonia solution
to a solution of the
substance.
Flame Test
Aluminum Al3+
White precipitate
that dissolves in excess
sodium hydroxide
White precipitate
that is insoluble in excess
ammonia.
Ammonium NH4
+ Ammonia gas is
produced
No Chemical Reaction
Calcium Ca2+
White precipitate
that is insoluble in
excess sodium
hydroxide.
No Chemical Reaction
Red Flames are
produced
Copper (II) Cu2+
Pale Blue precipitate
sodium hydroxide
Pale blue precipitate
changing to deep blue
solution in excess
ammonia.
Green Flames are
produced
Iron(II) Fe2+
Pale green precipitate
excess sodium
hydroxide.
Pale green precipitate
ammonia.
Iron(III) Fe3+
Red-brown precipitate
excess sodium
hydroxide.
Red-brown precipitate
ammonia..
Lead(II) Pb2+
White precipitate
sodium hydroxide
White precipitate
ammonia.
Blue Flames are
produced
Magnesium Mg2+
White precipitate
excess sodium
hydroxide.
White precipitate
ammonia.
Zinc Zn2+
White precipitate
sodium hydroxide
White precipitate
ammonia

1.6 Tests for Anions
Anion Symbol Test Results
Chloride Cl-
Add aqueous silver nitrate solution to a
solution of substance
Or
Add acidified lead (II) nitrate solution to
a solution of substance
White precipitate that is soluble in
ammonia solution.
Bromide Br- Cream precipitate, that is slightly
soluble in ammonia solution.
Iodide I-
Pale yellow precipitate, that is
insoluble in ammonia solution.
Carbonate CO3
2- Add dilute hydrochloric acid to the
substance.
Carbon dioxide gas is produced
Nitrate NO3
-
Add dilute sodium hydroxide solution,
followed by Aluminum powder and
warm
Ammonia gas is produced
Sulphate SO4
2-
Acidify the solution of the substance
(Either HCl or HNO3)
Add solution of barium cations (BaCl or
BaNO3) to the solution.
White precipitate, does
not dissolve in excess dilute acid.
1.7 Tests for Gases
Gas Symbol Properties Litmus Test Splint Test Limewater test
Hydrogen H2
Colourless &
Odourless
Extinguish a
lighted splint
with a pop
sound
Oxygen O2
Colourless &
Odourless
Relights a
glowing splint
Carbon
Dioxide
CO2
Colourless &
Odourless
Turns moist Blue
litmus Red
Forms a White
precipitate
Chlorine Cl2
Greenish-yellow
with a Choking Smell
Turns moist Blue
litmus Red
and eventually
bleaches White
Ammonia NH3
Colourless with a
Pungent Smell
Turns moist Red
litmus Blue

Sulphur
Dioxide
SO2
Colourless with a
Choking Smell
Turns moist Blue
litmus Red
Chapter 2
Kinetic Particle Theory
2.1 Three States of Matter
Properties Solid Liquid Gas
Volume Fixed Volume Fixed Volume No Fixed Volume
Shape Fixed Shape No Fixed Shape No Fixed Shape
Compressibility Incompressible Negligibly compressible Very compressible
Packing
Particles are very
closely packed
Particles are closely
packed
Particles are very
far apart
Forces of
Attraction
Very strong forces
of attraction
between particles
Strong forces of
attraction between
particles
Very weak forces
of attraction
between particles
Motion
Particles vibrate
about a fixed
position
Particles can slide over
each other
Particles are in
random motion
Diagrammatic
Representation

2.2 Melting & Freezing
Important Explanation: Melting
During melting, the particles of a solid gain energy and vibrate until they overcome the
forces of attraction between the particles, moving faster and further apart. At this
point, there is no rise in temperature as all heat energy is used to separate the particles
at this point.
This is the melting point, at which the temperature remains constant until the whole solid
has melted into a liquid.
Important points to include:
1. Gaining/Losing Energy
2. Motion of particles
3. Forces of attraction
4. New motion of particles
5. No rise/drop in temperature as all heat energy is used to separate/combine the
particles
6. _____ Point where temperature remains constant
Graphical Representation

1. Between points a and b, the substance is in the Solid state
2. Between points b and c, the substance is in a mixture of Solid & Liquid states
3. Between points c and d, the substance is in the Liquid state
a
b c
d
Temperature/K
Time/min

2.3 Boiling & Condensation
During boiling, the particles of a liquid gain energy, sliding over each other until they
overcome the forces of attraction between the particles, moving faster, randomly and
very far apart. At this point, there is no rise in temperature as all heat energy is used to
separate the particles at this point.
This is the boiling point, at which the temperature remains constant until the whole liquid
has boiled into a gas.
Difference between Boiling & Evaporation
Boiling Evaporation
Occurs at boiling point Occurs at any temperature below boiling point
Occurs throughout the liquid Occurs only at the surface of the liquid
Bubbles observed No bubbles observed
Occurs quickly Occurs slowly
2.4 Sublimation
Example of Substances that sublimes are Carbon Dioxide, Naphthalene (Mothballs) and
Iodine
2.5 Diffusion
1. Diffusion of a substance is the movement of particles from a region of higher
concentration of a substance to a lower concentration of the same substance down
a concentration gradient.
 Examples include a teabag in a cup of water, or perfume in a room
2. The rate of diffusion is dependent on two factors: Temperature and Molecular size of
the particle.
 At a higher temperature, particles have more kinetic energy, resulting in them
moving faster. Hence they will be able to move at greater speeds from a
region of lower concentration to a region of higher concentration.
 Smaller particles are able to move faster, resulting in a faster rate of diffusion.

Chapter 3
Atomic Structure and Chemical Bonding
3.1 Atomic Structure (& Symbols to represent Atomic Structure)
Atoms are the basic building blocks of all matter.
Particle Relative Charge Relative Mass Location
Proton +1 1 Nucleus
Neutron 0 1 Nucleus
Electron -1 0.0005 Electronic Shells outside the Nucleus
1. Atomic Number (or Proton Number) → The number of Protons in an atom
 The number of Protons in an atom = The number of electrons
2. Mass number (or Nucleon Number) → The number of Protons + Neutrons
3. Atoms with same number of electrons and protons but different number of neutrons
= isotopes
𝑋 ← 𝑆𝑦𝑚𝑏𝑜𝑙𝑃𝑟𝑜𝑡𝑜𝑛 ( 𝐴𝑡𝑜𝑚𝑖𝑐) 𝑛𝑢𝑚𝑏𝑒𝑟 → 𝑏
𝑁𝑢𝑐𝑙𝑒𝑜𝑛 ( 𝑀𝑎𝑠𝑠) 𝑛𝑢𝑚𝑏𝑒𝑟 → 𝑎

Common Isotopes
Element Isotope
Proton
Number
Number of
Electrons
Nucleon
Number
Number of
Neutrons
Hydrogen
Protium 1 1 1 0
Deuterium 1 1 2 1
Tritium 1 1 3 2
Carbon
Carbon-12 6 6 12 6
Carbon-13 6 6 13 7
Carbon-14 6 6 14 8
Chlorine
Chlorine-35 17 17 35 18
Chlorine-37 17 17 37 20
3.2 Electron Arrangement
1. Electrons in an atom are arranged in energy shells
2. The arrangement of the electrons is call its electronic configuration
3. The first shell can hold a maximum of 2 electrons
4. The second and third shell can hold a maximum of 8 electrons
5. The outermost shell is called the valence shell. the electrons in this shell is called
valence electrons.
6. The shells & the no. of electrons can be represented by a dot-and-cross diagram.
Electronic Configuration & Dot-and-cross Diagram Examples
Element Electronic Configuration Dot-and-cross Diagram
Hydrogen 1
Carbon 2.4
Oxygen 2.6

Silicon 2.8.4

Easy Reference Table (Please try to not rely on this.learn to derive these information from a periodic table)
Symbol Element Proton Number Electronic Configuration
𝐻1
1
Hydrogen 1 1
𝐻𝑒2
4
Helium 2 2
𝐿𝑖3
7
Lithium 3 2.1
𝐵𝑒4
9
Beryllium 4 2.2
𝐵5
11
Boron 5 2.3
𝐶6
12
Carbon 6 2.4
𝑁7
14
Nitrogen 7 2.5
𝑂8
16
Oxygen 8 2.6
𝐹9
19
Fluorine 9 2.7
𝑁𝑒10
20
Neon 10 2.8
𝑁𝑎11
23
Sodium 11 2.8.1
𝑀𝑔12
24
Magnesium 12 2.8.2
𝐴𝑙13
27
Aluminum 13 2.8.3
𝑆𝑖14
28
Silicon 14 2.8.4
𝑃15
31
Phosphorus 15 2.8.5
𝑆16
32
Sulphur 16 2.8.6
𝐶𝑙17
35
or 𝐶𝑙17
37 Chlorine 17 2.8.7
𝐴𝑟18
40
Argon 18 2.8.8
𝐾19
39
Potassium 19 2.8.8.1
𝐶𝑎20
40
Calcium 20 2.8.8.2
3.3 Formation of Ions
1. Atoms are generally naturally unstable. (With the exception of Group 0/ Group 8
elements)
2. This is due to the lack of a stable octet (or duplet) structure, with fully filled shells.
3. Atoms can form ions by gaining or losing valence electrons, in their attempt to obtain
a stable octet (or duplet) structure, with fully filled shells.
4. Metals usually lose electrons, forming Positive Ions, also known as Cations.
5. Non-Metals usually gain electrons, forming Negative Ions, also known as Anions.
Example
 Lithium can lose a valence electron to form Li+, a positive ion (cation) with a fully
filled valence shell of 2 electron, with an electronic configuration of 2.
 Magnesium can lose two valence electrons to form Mg2+, a positive ion (cation) with
a fully filled valence shell of 8 electron, with an electronic configuration of 2.8
 Oxygen can gain two valence electrons to form O2-, a negative ion (anion) with a
fully filled valence shell of 8 electron, with an electronic configuration of 2.8

3.4 Ionic Bonding
1. Ionic Bonds are formed when metallic atoms give their valence electrons to non-
metallic atoms. This enables both the metallic and the non-metallic atoms to achieve
a stable octet (or duplet) structure, with fully filled valence shells.
2. These ions formed are oppositely charged, and attract each other through strong
electrostatic forces of attraction, thus forming the ionic bond.
Examples
Na + Cl → NaCl
Sodium (Na) can lose a valence electron to Chlorine (Cl), forming NaCl, with the
positively charged Na+ ion, and the negatively charged Cl- ion.
Mg + O → MgO
Magnesium (Mg) can lose two valence electrons to Oxygen (O), forming MgO, with the
positively charged Mg2+ ion, and the negatively charged O2- ion.
Mg + 2 Cl → MgCl2
Magnesium (Mg) can lose two valence electrons, one to each Chlorine (Cl), forming
MgCl2, with the positively charged Mg2+ ion, and 2 negatively charged Cl- ions.

3.5 Structure of Ionic Compounds
1. A solid ionic compound has a giant lattice structure with alternating positively and
negatively charged ions.
2. The ions are held in fixed positions by strong electrostatic forces of attraction.
3.6 Physical Properties of Ionic Compounds
Physical Properties of Ionic
Compounds
Explanation in terms of their structure and
bonding
Ionic compounds are usually
crystalline solids at room
temperature
The ions are arranged in a highly regular fashion,
with strong electrostatic forces of attraction (ionic
bonds) between the ions.
Ionic compounds have high
melting and boiling points
The electrostatic forces of attraction between the
oppositely charged ions is very strong and extends
over the entire crystalline structure. Large amounts
of energy is required to separate the ions
Ionic compounds cannot conduct
electricity when solid, but do so in
molten or in aqueous form.
In the solid structure, the ions are held in fixed
positions. When molten or in aqueous solution, the
ions are mobile, so a flow of charge is possible.
Most ionic compounds are water
soluble, but insoluble in organic
solvents.
Water molecules are polar, and are attracted to the
charged ions in the ionic compound. This helps to
pull the crystalline structure as the solid dissolves.

3.7 Covalent Bonding
1. Covalent bonding occurs when the electrons are shared, so as to achieve a stable
octet/duplet structure.
2. Each pairs of shared electrons forms one covalent bond.
3. Covalent bonding occurs mainly between non-metals
Examples
H + H → H2
Two hydrogen atoms can share an electron each to form a covalent H-H bond, giving
both atoms a stable duplet structure.
O + O → O2
Two oxygen atoms can share two electrons each to form two covalent O-O bonds,
giving both atoms a stable octet structure.
H + H + O → H2O
Each hydrogen atom shares one electron with the oxygen atoms, forming 2 O-H
covalent bonds.

3.8 Physical Properties of Simple Covalent Compounds
Physical Properties of Simple
Covalent Compounds
Explanation in terms of their structure and
bonding
Simple Covalent Compounds
have low boiling and melting
points.
The inter-molecular forces of attraction are very
weak, hence very little energy is required to break
the forces apart.
Simple Covalent Compounds
cannot conduct electricity in any
state.
There are no mobile ions or electrons in simple
covalent compounds in any states.
Simple Covalent Compounds are
soluble in organic solvents but
not in water.
Simple Covalent Compounds have generally non-
polar molecules, and thus would be unable to
dissolve in a solvent like water with strong hydrogen
bonding, but would be soluble in a organic solvent
like ethanol, petrol or trichloromethane.
3.9 Metallic Bonding
1. Metals can give away their valance electrons to form ions stable octet/duplet
structures.
2. These valence electrons then form a “sea of electrons” which can move freely
among the metal cations, holding them in place in a metallic lattice structure.
3.10 Physical Properties of Metals
Physical Properties of
Simple Covalent
Compounds
Explanation in terms of their structure and bonding
Metals have high boiling
and melting points.
The electrostatic forces of attraction between the cations
and the delocalized electrons are very strong, and
would high temperatures to overcome them.
Metals can conduct
electricity and heat in any
state.
Metals can conduct heat and electricity in any state due
to the presence of the free electrons available to carry
electrical and thermal energy.

Metals are soft and
malleable.
Due to the sea of electrons, the cations can easily slide
over each other.
3.11 Giant Molecular Structure/ Macromolecules
1. Macromolecules contain billions of atoms per molecules, covalently bonded to each
other.
2. Examples of Macromolecules include diamond, graphite, silicon, silicon dioxide and
polymers (to be covered in greater detail in organic chemistry).
3.12 Diamond and Graphite
1. Diamonds and Graphite are macromolecules with giant structures.
2. Diamonds and Graphite are allotropes of carbon (Allotropes are different structural
arrangement of the same element).
Diamond Graphite
Structure
Every carbon atom is part of a
giant molecule.
The carbon atoms forms layers,
which are held together by weak
van der waal’s forces of
attractions.
Each layer is a flat (planar)
macromolecule.
Each carbon atom is bonded to
four other carbon atoms in a
tetrahedral arrangement.
Each carbon atom is connected to
only 3 other carbon atoms, with
one free electron per atom.
Properties
Very high boiling and melting point. Very high boiling and melting point.
Very poor conductor of electricity. Good conductor of electricity of
heat, due to the free electrons.
Very hard – the hardest substance
in the world. (Hardness: physically
resistant to change when force is
applied).
Very soft, due to weak Van Der
Waal’s forces between layers;
layers can easily slide over each
other.
Others
Silicon and Silicon Dioxide have
similar structures and similar
properties.

3.13 Elements, Compounds and Mixtures.
1. Elements are made up of only one kind of atoms, and can be found directly in the
periodic table.
2. Elements cannot be further broken down by physical or chemical means (at least for
O-levels syllabus).
3. Compounds are made of two or more different kinds atoms chemically combined in
a fixed proportion.
 E.g. Hydrochloric acid comprises of hydrogen atoms and chlorine atoms in a
1:1 fixed proportion only.
4. A mixture is made up of two or more elements and/or compounds physically
combined. They can be physically separated by physical means and do not have a
fixed proportion.
 E.g. Saltwater can be 50% salt 50% water, or 40% salt 60% water or 30% salt
70% water …
Mixture Compounds
Composition No fixed composition/ proportion.
The percentage of one
element/compound to another in a
mixture can vary.
Fixed composition/ proportion.
The percentage of one element to
another in a particular compound
is always the same.
Properties No set of properties of its own. It
exhibits a combination of the
properties of the constituent
components.
It has its own set of properties.
Melting Point
& Boiling Point
No fixed M.P. or B.P. A fixed M.P. and B.P.
Preparation No chemical reaction has to
occur.
A chemical reaction has to occur.
Separation Can be separated into its
components by physical means
Can only be separated into its
components by chemical means

Chapter 4
Stoichiometry & Mole Concept
4.1 Chemical Formulae
 Number of Atoms/Ions are denoted by subscript.
 Charge of ions are denoted by superscript.
 Metals and/or positive cations are placed first in the chemical formula.
 Brackets are used for repeated clusters of atoms (like anions)
o E.g. Magnesium Nitrate = Mg(NO3)2
 Mono- is a prefix for indicating only 1 of a particular atom/ cluster of atoms
o Carbon Monoxide = CO
 Di- is a prefix indicating 2 of a particular atoms/cluster of atoms
o Carbon Dioxide = CO2
 Tri- is a prefix indicating 3 of a particular atoms/cluster of atoms
o Dinitrogen Trioxide = N2O3
 Prefixes are only used for covalent molecules. For ionic compounds, the formulae have to be
deduced from the valency of the components.
 For transition metals, the valency is indicated in brackets
o Iron (II) has a valency of 2, Iron (III) has a valency of 3.
Common Ions
Name Formulae Name Formulae
Ammonium NH3+ Nitrate NO3-
Carbonate CO2 2- Nitrite NO2-
Chromate (VI) CrO42- Oxide O2-
Dichromate (VI) Cr2O72- Phosphate PO43-
Ethanoate CH3CO2- Sulphate / Sulfate SO42-
Hydrogencarbonate HCO2- Sulphite / Sulfite SO32-
Hydroxide OH- Sulphide / Sulfide S2-
4.2 Balancing Equations (With state symbols)
 A chemical equation is used to shows information in a chemical reaction.
o What chemicals are used (Reactants).
o What chemicals are created (Products).
o What states they are in: (s), (l), (g) or (aq)
 The reactants are always on the left, and the products on the right.
 Ensure the left side of the equation equals the right side.
Example 1:
Fe(s) + 2 C5H6(g) → Fe(C5H5)2(s) + H2(g)
 There is 1 Iron atom on both sides.
 There are 10 Carbon atoms on both sides.
 There are 12 Hydrogen atoms on both sides.
Example 2:
CH3CH2OH + CH3CO2H ⇌ CH3CO2CH2CH3 + H2O
 There are 4 Carbon atoms on both sides.
 There are 10 Hydrogen atoms on both sides.
 There are 4 Oxygen atoms on both sides.
 *The catalyst HCl is not included in the equation as it appears in the same form on both sides.

4.3 Ionic Equations
 An ionic equation only shows the ions involved in the reaction
o An ion is involved in the reaction if its charge changes during the reaction.
o Ions uninvolved are called spectator ions
o Note: Insoluble compounds should not be broken up into its component ions
 Ensure the total charge on the left side of the equation equals the total charge on the right side.
Example 1:
Chemical Equation: CuCO3 (s) + 2 HCl(aq) → CuCl2(aq) + CO2(g) + H2O(l)
Ionic Equation: CuCO3 (s) + 2 H+
(aq) → Cu2+
(aq) + CO2(g) + H2O(l)
 Charge on the left side 2 x (+1) = +2
 Charge on the right side +2 = Charge on the left side
Example 2:
Chemical Equation: NaOH (aq) + HCl(aq) → NaCl(aq) + H2O(l)
Ionic Equation: OH-
(aq) + H+
(aq) → H2O(l)
 Charge on the left side - 1 +1 = 0
 Charge on the right side 0 = Charge on the left side
4.4 Relative Atomic Mass (Element) & Relative Molecular Mass
The relative atomic mass (Ar) of an element is the average mass of one atom of an element compared to
1
12
of the mass of a carbon-12 atom.
 Relative Atomic Mass (Ar) may sometimes have the same values as the mass number, but they are
conceptually DIFFERENT from each other.
o Mass number refer to the number of protons and neutrons in an atom. They can differ betweens
isotopes of the same elements. Atoms of different elements can have the same mass number.
o Relative Atomic Mass (Ar) refers to the AVERAGE mass of atoms of a particular element in
accordance with isotopic composition.
 Relative Atomic Mass (Ar) has no units.
The relative molecular mass (Mr) of a substance is the average mass of one molecule of the substance
compared with
1
12
of the mass of a carbon-12 atom.
4.5 % by Mass of an Element in a Compound
% by Mass of an Element in a Compound =
No.of Atoms x Relative Atomic Mass (Ar)of the Element
relative molecular mass (Mr) of the Compound
x 100%
Example 1:
% by Mass of Oxygen in Carbon Dioxide (CO2) =
No.of Atoms of Oxygen x (Ar)of the Oxygen
(Mr) of Carbon Dioxide
x 100%
=
2 x 16
12+16 x 2
x 100%
= 72.73% (2 decimal points)

4.6 Mole Concept
 A mole is the number of particles which contains the same number of atoms in a 12.0g sample of
carbon-12.
o This number is 6.02 x 1023. This number is also known as Avogadro's number.
 Mole can be abbreviated as mol.
4.7 Molar Mass and Molar Volume
 The mass (in grams) of 1 mole of a substance, is called its molar mass.
o The molar mass of a substance is equal to its relative atomic mass or relative molecular mass.
 E.g. The relative atomic mass (Ar) of Helium is 4.0. The molar mass of Helium is 4.0 g. The
mass of 1 mole of Helium atoms is 4.0 g. The mass of 6.02 x 1023 Helium atoms is 4.0 g.
 E.g. The relative molecular mass (Mr) of Carbon Dioxide is 44.0. The molar mass of Carbon
Dioxide is 44.0 g. The mass of 6.02 x 1023 Carbon Dioxide molecules is 44.0 g.
 The volume occupied by 1 mole of a Gas, is called the molar volume.
o The molar volume of ALL gases at r.t.p. is 24.0 dm3 = 24 000.0 cm3
 r.t.p refers to Room Temperature and Pressure
 r.t.p: Temperature = 25oC and Pressure = 1 atm
o The molar volume of ALL gases at s.t.p. is 22.4 dm3 = 22 400.0 cm3
 r.t.p refers to Standard Temperature and Pressure
 s.t.p: Temperature = 0oC and Pressure = 1 atm
o E.g. 1 mole of Chlorine gas at r.t.p. has a volume of 24.0 dm3. 1 mole of Bromine gas at r.t.p. also
has a volume of 24.0 dm3, despite having a larger atom than Chlorine.
Mole =
𝑀𝑎𝑠𝑠
𝑀𝑜𝑙𝑎𝑟 𝑀𝑎𝑠𝑠
 The number of moles present in a sample =
𝑀𝑎𝑠𝑠 𝑜𝑓 𝑡ℎ𝑒 𝑠𝑎𝑚𝑝𝑙𝑒
𝑚𝑜𝑙𝑎𝑟 𝑚𝑎𝑠𝑠 𝑜𝑓 𝑡ℎ𝑒 𝑠𝑎𝑚𝑝𝑙𝑒
Example 1: Calculate the number of moles in a 142.0g sample of Chlorine gas.
Molar Mass of Chlorine gas (Cl2) = 2 x 35.5g
= 71.0g
Number of moles of Cl2 present =
=
142 .0𝑔
71.0𝑔
= 2 mols
Example 2: Calculate the number of moles in a 100.0g sample of NaCl.
Molar Mass of NaCl = 23.0g + 35.5g
= 58.5 g
Number of moles of Cl2 present =
=
100 .0𝑔
58.5𝑔
= 1.71 mols (3.s.f)

4.8 Molar Solutions (Concentration)
 The concentration of a solution refers to the amount of solute in 1 dm3 of solution.
o 1 000 cm3 = 1 dm3
 Concentration can be presented in 2 ways: Concentration or Molarity
Concentration in (g dm-3
)=
𝑀𝑎𝑠𝑠 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑒 (𝑔𝑟𝑎𝑚𝑠)
𝑉𝑜𝑙𝑢𝑚𝑒 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛 (𝑑𝑚3)
Molarity (mol dm-3
or M) =
𝑁𝑢𝑚𝑏𝑒𝑟 𝑜𝑓 𝑚𝑜𝑙𝑒𝑠 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑒 (𝑚𝑜𝑙𝑠)
Molarity (mol dm-3
or M) =
Concentration in (g dm−3)
𝑀𝑜𝑙𝑎𝑟 𝑀𝑎𝑠𝑠 𝑜𝑓 𝑆𝑜𝑙𝑢𝑡𝑒 (𝑔𝑟𝑎𝑚)
Example 1: A 100 cm3 solution of HCl contains 1g of HCl.
Concentration of Solution =
𝑀𝑎𝑠𝑠 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑒 ( 𝑔𝑟𝑎𝑚𝑠)
𝑉𝑜𝑙𝑢𝑚𝑒 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛 ( 𝑑𝑚3)
=
1𝑔
100 𝑐𝑚3
= 0.1 g dm3
Number of Moles of HCl =
𝑚𝑎𝑠𝑠
𝑀𝑟
=
1 𝑔
36.5
= 0.0274 mols (3.s.f)
Molarity of Solution =
𝑁𝑢𝑚𝑏𝑒𝑟 𝑜𝑓 𝑚𝑜𝑙𝑒𝑠 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑒 (𝑚𝑜𝑙𝑠)
=
0.0274 𝑚𝑜𝑙𝑠
0.1 𝑑𝑚3
= 0.00274 mol dm-3
Molarity of Solution =
Concentration in (g dm−3)
𝑀𝑜𝑙𝑎𝑟 𝑀𝑎𝑠𝑠 𝑜𝑓 𝑆𝑜𝑙𝑢𝑡𝑒 (𝑔𝑟𝑎𝑚)
=
0.1g dm−3
36 .5
= 0.00274 mol dm-3
 The concentration of a solution changes w hen diluted
M1V1 = M2V2
M1 = Original Molarity
V1 = Original Volume
M2 = New Molarity
V2 = New Volume
Example 1: a 10 cm3 sample of a 1M HCl solution is diluted to 50cm3
M1V1 = M2V2
(1 M) x (10 cm3)= M2 x (50cm3)
M2 =
1 𝑀 𝑥 10 𝑐𝑚3
50 𝑐𝑚3
= 0.20 M

4.9 Empirical Formulae
 The empirical formulae shows the simplest integer ratio of the different types of atoms in a
compound.
o The empirical formulae for Benzene (C6H6) is CH
o The empirical formulae for Butane (C4H10) is C2H5
 The empirical formulae may be determined using the following table if the mass of each individual
constituent elements are given:
Element X Element Y
Mass of each Individual Element
Molar Mass
Number of Moles
Smallest Mole
Divide by the smallest Mole
Ratio
Example 1: A Sample of Iron Sulphide contains 5.373g of iron and 4.627g of sulphur.
Iron Sulphur
Mass of each Individual Element 5.373g 4.627g
Molar Mass 56.0 32.0
Number of Moles
5.373𝑔
56.0
= 0.0959 mol
4.627𝑔
32
= 0.145 mol
Smallest Mole 0.0959 mol 0.0959 mol
Divide by the smallest Mole
0.0959
0.0959
= 1
0.145
0.0959
= 1.51
Ratio 2 3
Empirical Formulae = Fe2S3
Example 2: A compound contained (by mass) 23.3% Magnesium, 30.7% Sulphur and 46.0% Oxygen.
Magnesium Sulphur Oxygen
Mass of a 100g Sample 23.3g 30.7 g 46.0 g
Molar Mass 24.0 32.0 16.0
Moles
23.3
24.0
= 0.97 mol
30.7
32
= 0.96 mol
46.0
16
= 2.88 mol
Smallest Mole 0.96 mol 0.96 mol 0.96 mol
Divide
0.97
0.96
= 1
0.96
0.96
= 1
2.88
0.96
= 3
Ratio 1 1 3
Empirical Formulae = MgSO3

4.10 Molecular Formulae
 The molecular formulae shows the actual number of atoms of each element in each molecule of a
compound
 Molecular Formulae = n x Empirical Formulae
o n =
Molecular Mass
Empirical Mass
Example 1: A Compound contained (by mass) 26.67% Carbon, 2.22% Hydrogen and 71.11% Oxygen.
One mole of the compound has a mass of 90.1g.
Carbon Hydrogen Oxygen
Mass of a 100g Sample 26.67 g 2.22 g 71.11 g
Molar Mass 12.0 1.0 16.0
Moles
26.67
12.0
= 2.22 mol
2.22
1.0
= 2.22 mol
71.11
16.0
= 4.44 mol
Smallest Mole 2.22 mol 2.22 mol 2.22 mol
Divide
2.22
2.22
= 1
2.22
2.22
= 1
4.44
2.22
= 2
Ratio 1 1 2
Empirical Formulae = CHO2
Empirical Mass = 12.0g + 1.0g + (2 x 16.0)g
= 45.0g
n =
Molecular Mass
Empirical Mass
=
90.1 𝑔
45.0 𝑔
= 2
Molecular Formulae = 2 x CHO2
= C2H2O4
Example 2: The empirical formulae of a compound is C2H4O. Its relative molecular mass is 88
Empirical Formulae = C2H4O
Relative mass = (2 x 12.0) + (4 x 1.0) + 16.0
= 44.0
n
Molecular Mass
Empirical Mass
=
88.0 𝑔
44.0 𝑔
= 2
Molecular Formulae = 2 x C2H4O
= C4H8O2

4.11 Theoretical Product Yield
 The theoretical product yield of a chemical reaction can be calculated from the mass (or volume for
gases) of the reactants, using a balanced equation
Example 1: 4.0g of Methane is completely burnt in excess oxygen to yield Carbon Dioxide and Water.
CH4 (g) + 2 O2 (g) = CO2 (g) + 2 H2O(l)
Number of moles of Methane =
𝑀𝑎𝑠𝑠 𝑜𝑓 𝑀𝑒𝑡ℎ𝑎𝑛𝑒
𝑀𝑜𝑙𝑎𝑟 𝑀𝑎𝑠𝑠 𝑜𝑓 𝑀𝑒𝑡ℎ𝑎𝑛𝑒
=
4.0 𝑔
12+(4 𝑥 1.0)
=
4.0 𝑔
16 .0
= 0.25 mol
Mole Ratio
CH4 : CO2
1 : 1
Number of moles of CO2 to be produced = Number of moles of Methane
= 0.25 mol
Mass of CO2 produced = Moles x Molar Mass
= 0.25 mols x [12.0g + (16.0g x 2)]
= 0.25 mols x 44.0g
= 11.0 g
Volume of CO2 produced at r.t.p = Moles x 24.0 dm3
= 0.25 x 24.0 dm3
= 6.0 dm3
Mole Ratio
CH4 : H2O
1 : 2
Number of moles of H2O to be produced = 2 x Number of moles of Methane
= 2 x 0.25 mol
= 0.5 mol
Mass of H2O produced = Moles x Molar Mass
= 0.5 mols x [(1.0g x 2) + 16.0g]
= 0.25 mols x 18.0g
= 9.0 g

4.12 Limiting Reagent
 A limiting reagent is a reactant that causes a reaction to stop once it's completely consumed. It limits
the amount of products to be formed.
 The limiting reagent can be identified by comparing the number of moles of each reactants with the
mole ratio of the reactants in the chemical equation.
Example 1: 5.6g of iron is burnt in 6.4g of sulphur to form iron (III) sulphide.
2Fe(s) + 3S(s) = Fe2S3 (s)
Number of moles of Iron =
𝑀𝑎𝑠𝑠 𝑜𝑓 𝐼𝑟𝑜𝑛
𝑀𝑜𝑙𝑎𝑟 𝑀𝑎𝑠𝑠 𝑜𝑓 𝐼𝑟𝑜𝑛
=
5.6 𝑔
56
= 0.1 mol
Number of moles of Sulphur =
𝑀𝑎𝑠𝑠 𝑜𝑓 𝑆𝑢𝑙𝑝ℎ𝑢𝑟
𝑀𝑜𝑙𝑎𝑟 𝑀𝑎𝑠𝑠 𝑜𝑓 𝑆𝑢𝑙𝑝ℎ𝑢𝑟
=
6.4 𝑔
32
= 0.2 mol
Mole Ratio
Fe : S
2 : 3
0.1 : 0.15 < 0.2
The limiting reagent is Iron
Number of moles of Fe2S3 produced =
1
2
x Number of moles of Iron
=
1
2
x 0.1 mol
= 0.05 mols
Mass of Fe2S3 produced = Moles x Molar Mass
= 0.05 mols x [(56.0g x 2) + (32.0g x 3)]
= 0.05 mols x 208.0g
= 10.4 g
Example 2: 0.05 moles of Zinc is added to 0.075 moles of HCl. Zinc (II) Chloride and H2 gas is produced.
Zn(s) + 2HCl(aq) = ZnCl2(aq) + H2(g)
Number of moles of Zinc = 0.05 mol
Number of moles of HCl = 0.075 mol
Mole Ratio
Zn : HCl
1 : 2
0.05 : 0.10 > 0.075
The limiting reagent is HCl
Number of moles of ZnCl2 produced =
1
2
x Number of moles of HCl
=
1
2
x 0.075 mol
= 0.0375 mols
Mass of ZnCl2 produced = Moles x Molar Mass
= 0.0375 mols x [65.0g + (32.0g x 3)]
= 0.0375 mols x 129.0g
= 4.84 g (3 significant figures)

4.13 Percentage Yield
 The actual yield of a reaction is usually lesser than the theoretical yield. Thus the percentage yield
can be used to determine the effectiveness of the process.
Percentage yield =
𝐴𝑐𝑡𝑢𝑎𝑙 𝑌𝑖𝑒𝑙𝑑
𝑇ℎ𝑒𝑜𝑟𝑒𝑡𝑖𝑐𝑎𝑙 𝑌𝑖𝑒𝑙𝑑
x 100%
Example 1: 50.0 cm3 of a 0.105 mol dm-3 aqueous calcium chloride solution is treated with 2.0g of Silver
Nitrate to form Silver Chloride. 1.45g of Silver Chloride is recorded
CaCl2(aq) + 2AgNO3(aq) = 2 AgCl(aq) + Ca(NO3)2 (aq)
Number of moles of CaCl2 = 50.0 cm3 x 0.105 mol dm-3
= 0.050 dm3 x 0.105 mol dm-3
= 0.00525 mol
Number of moles of AgNO3 =
𝑀𝑎𝑠𝑠 𝑜𝑓 AgNO3
𝑀𝑜𝑙𝑎𝑟 𝑀𝑎𝑠𝑠 𝑜𝑓 AgNO3
=
2.0 𝑔
108 .0𝑔+14.0𝑔 +(16.0𝑔 𝑥 3)
=
2.0 𝑔
170 .0 𝑔
= 0.0118 mol (3.s.f)
Mole Ratio
CaCl2 : AgNO3
1 : 2
0.00525 : 0.0105 < 0.0118
The limiting reagent is CaCl2
Number of moles of AgCl produced = 2 x Number of moles of CaCl2
= 2 x 0.00525 mols
= 0.0105 mols
Theoretical Mass of AgCl produced = Moles x Molar Mass
= 0.0105 mols x [108.0g + 35.5g]
= 0.0105 mols x 143.5g
= 1.51 g (3.s.f)
Actual mass of AgCl produced = 1.45 g
Percentage Yield =
1.45 𝑔
1.51 𝑔
x 100%
= 96.03% (2 Decimal Places)

4.14 Percentage Purity
 Percentage purity indicates the amount of pure substance present in a sample of chemical
substance.
Percentage Purity =
𝑀𝑎𝑠𝑠 𝑜𝑓 𝑃𝑢𝑟𝑒 𝑆𝑢𝑏𝑠𝑡𝑎𝑛𝑐𝑒 𝑃𝑟𝑒𝑠𝑒𝑛𝑡
𝑀𝑎𝑠𝑠 𝑜𝑓 𝑆𝑎𝑚𝑝𝑙𝑒
x 100%
Example 1: 4.35g of impure manganese (IV) oxide reacted with 48 cm3 of 1.0 mol dm3 of HCl.
MnO2(s) + 4HCl(aq) = MnCl2(aq) + 2 H2O(l) + Cl2 (g)
Number of moles of HCl = 48.0 cm3 x 1.0 mol dm-3
= 0.048 dm3 x 1.0 mol dm-3
= 0.048 mol
Mole Ratio
MnO2 : HCl
1 : 4
Number of moles of MnO2 present =
1
4
x Number of moles of HCl
=
1
4
x 0.048 mols
= 0.012 mols
Mass of Pure MnO2 present = 0.012 mols x Molar Mass of MnO2
= 0.012 mols x [55.0g + (16.0g x 2)]
= 0.012 mols x 87.0g
= 1.044g
Percentage Purity of MnO2 =
1.044 𝑔
4.35 𝑔
x 100%
= 24.00%

Chapter 5
Energy and Chemical Reactions
5.1 Chemical Energy
 All chemical substances store chemical energy
o This energy can be converted into heat, light, electrical or sound energies
 A reaction that gives out heat to the surrounding is an exothermic reaction
 A reaction that takes in heat from the surrounding is an endothermic reaction
 ∆H represents the change in heat energy of the reaction.
o It is the difference between the energy content of the products and the reactants.
o Exothermic Reactions have a negative ∆H
o Endothermic Reactions have a positive ∆H
5.2 Bond Energies
 Bond Energies measures the strength of a covalent bond.
 When two atoms are joined together by a chemical bond, heat energy is released.
o Hence, bond forming is exothermic
 When a chemical bond is broken, heat energy is consumed.
o Hence, bond breaking is endothermic
 The amount of energy consumed in breaking a chemical bond is known as the bond energy
o The same amount of energy is produced when the same bond is formed
∆H (Heat of reaction) = Total Heat Energy Absorbed - Total Heat Energy Released
Covalent Bond Bond Energy / kJmol-1
H - H 436
Cl - Cl 242
C - C 348
C - H 412
O - H 463
Cl - H 431
N - H 388
O = O 496
C = O 743
N ≡ N 945
C = C 838
Example 1: Combustion is an Exothermic Reaction
CH4 + 2O2 → CO2 + 2H2O
Covalent Bonds in reactants = (4 x C - H bonds) + (2 x O = O bonds)
Sum of Bond Energies in Reactants (Er) = (4 x 412kJ) + (2 x 496kJ)
= 2240 kJ
Covalent Bonds in Products = (2 x C = O bonds) + (4 x O - H bonds)
Sum of Bond Energies in Product (Ep) = (2 x 743 kJ) + (4 x 463 kJ)
= 3338 kJ
Overall Heat of Reaction (∆H) = Total Heat Energy Absorbed - Total Heat Energy Released
= 2240 kJ - 3338 kJ
= -1098 kJ

5.3 Energy Profile Diagrams
5.4 Collision Theory
 Collision Theory states that a chemical reaction occurs when reactant particles collide with each
other. However, Not all collisions will result in the formation of products.
 A collision is effective only when the reactant particles have enough energy to overcome the
activation energy of the reaction, as well as having to collide in the proper orientation.
 Therefore, the speed of the reaction will depend on the number of effective collisions between the
reactants.
5.5 Speed of reaction
 The speed of reaction refers to how fast reactants are used up or how fast products are formed
 Speed of reaction =
𝐶ℎ𝑎𝑛𝑔𝑒 𝑖𝑛 𝑎𝑚𝑜𝑢𝑛𝑡 𝑜𝑓 𝑟𝑒𝑎𝑐𝑡𝑎𝑛𝑡 𝑜𝑟 𝑝𝑟𝑜𝑑𝑢𝑐𝑡𝑠
𝑡𝑖𝑚𝑒
 Factors that affect the speed of reactions include:
o Concentration of chemicals involved
 An increase in concentration results in a higher speed of reaction
 More particles in a given volume results in an increase in frequency of effective collisions
o Temperature
 An increase in temperature results in a higher speed of reaction
 At higher temperatures, the particles have more kinetic energy and thus move faster, this
leads to an increase frequency of effective collisions
 At higher temperatures, more particles have the necessary energy to overcome the
activation energy needed for an effective collision.
o Pressure
 Changes in pressures only affects reactions where gases are involved.
 An increase in pressure results in a higher speed of reaction
 The same number of particles in a smaller volume results in an increase in frequency of
effective collisions
o Particle Size
 A decrease in particle size results in a higher speed of reaction
 Breaking up the particles results in greater total surface area, which in turn results in more
particles being able to collide per unit time
o Catalysts
 Three experimental methods to determine the speed of a reaction include:
o Measuring time for reaction to be completed
o Measuring quantity of products formed over a fixed time interval
 By volume of gas produced
 By mass of product
o Measuring quantity of reactants left over a fixed time interval
 By mass of Reactants
 Titration For acidic or basic reactants

5.6 Catalysts
 A catalyst is a substance that changes the rate of reaction, but itself is chemically unchanged at the
end of the reaction.
 A catalyst changes the rate of reaction by lowering the activation energy of a reaction
o Because the catalysed activation energy is lower than the uncatalysed activation energy, the
reaction will take place more quickly.
 Different reaction will require different type of catalysts.
o Each catalyst is usually specific to a particular reaction
 Catalysts are usually transition metals, or transition metal compounds
 Enzymes are an example of biological catalysts found in living cells, used to speed up the breaking
down of giant molecules such as proteins or starch.
 A catalyst does not change the amount of products obtained
 A catalyst does not change the ∆H of a reaction.
5.7 Redox Reactions
 Redox is reaction where reduction and oxidation occurs simultaneously
 Reduction occurs when a substance
o Gains Hydrogen
o Or Loses Oxygen
o Or Gains Electrons
o Or Decrease in oxidation state
 A Reducing Agent causes another substance to be reduced
o Hence a substance being reduced is an oxidising agent
 Oxidation occurs when a substance
o Loses Hydrogen
o Or Gains Oxygen
o Or Loses Electrons
o Or Increase in oxidation state
 An Oxidizing Agent causes another substance to be oxidized
o Hence a substance being oxidized is a Reducing agent
OIL RIG

Oxidation is Lose, Reduction is Gain
Example 1:
H2 (g) + CuO (s) → Cu (s) + H2O (l)
 The hydrogen gas gains oxygen, and thus is oxidized. (Undergoes oxidation)
o The hydrogen gas is a Reducing agent
 The copper oxide loses oxygen, and thus is reduced. (Undergoes reduction)
o The copper oxide is an Oxidizing agent
Example 2:
4CO (g) + Fe2O3 (s) → 3Fe (s) + 4CO2 (g)
 The carbon monoxide gas gains oxygen, and thus is oxidized. (Undergoes oxidation)
o The CO gas is a Reducing agent
 The iron (III) oxide loses oxygen, and thus is reduced. (Undergoes reduction)
o The Iron (III) oxide is an Oxidizing agent
Example 3:
4H2S (g) + Cl2 (g) → 2HCl (g) + S (s)
 The chlorine gas gains hydrogen, and thus is reduced. (Undergoes reduction)
o The chlorine gas is a oxidizing agent
 The hydrogen sulphide loses hydrogen, and thus is oxidized. (Undergoes oxidation)
o The hydrogen sulphide is an reducing agent
Example 4:
FeSO4 (aq) + Zn (s) → ZnSO4 (aq) + Fe (s)
 The iron gains electrons (From oxidation state +2 to 0), and thus is reduced. (Undergoes reduction)
o The iron is a oxidizing agent
 The zinc loses electrons (From oxidation state 0 to +2), and thus is oxidized. (Undergoes oxidation)
o The zinc is an reducing agent
Example 5:
2 Na (s) + Cl2 (g) → 2NaCl (s)
 The chlorine gas gains electrons (From oxidation state 0 to -1), and is reduced. (Undergoes
reduction)
o The chlorine gas is a oxidizing agent
 The sodium loses electrons (From oxidation state 0 to +1), and is oxidized. (Undergoes oxidation)
o The sodium is an reducing agent
5.8 Test for Oxidizing and Reducing Agents
Type Agent Half Equation Colour Change
Oxdizing
KMnO4
MnO4-(aq) + 8H+(aq) + 5e- → Mn2+ (aq) + 4H2O(l)
Purple to
colourless
Acidified Potassium
Manganate (VII)
Oxdizing
K2Cr2O7
Cr2O72-(aq) + 14 H+(aq) + 5 e- → 2Cr3+ (aq) + 7H2O(l) Orange to Green
Acidified
Oxdizing
Cl2
Cl2(g) + 2 e- → 2Cl-(aq)
Greenish Yellow
to ColourlessChlorine
Reducing
KI
I-(aq) → I2(aq) + 2 e- Colourless to
Reddish BrownAqueous

Potassium Iodide
Reducing
FeSO4
Fe2+(aq) → Fe3+(aq) + e- Greenish to
Yellow
Aqueous Iron (II)
Sulphate
Chapter 6
Acid, Bases and Salts
6.1 pH
 pH measures the concentration of H+ ions in a solution
 The pH scale ranges from 0 to 14
o Acids have a pH value of less than 7
 The lower the pH, the stronger the acid
 Sulphuric Acid in Car Batteries (pH 1) is a much stronger acid than lemon juice (pH 2)
o Bases and Alkalis have a pH value of greater than 7
 The higher the pH, the stronger the base/ alkali
 Sodium Hydroxide in bleach (pH 13) is a much stronger base than ammonia in fertilizer (pH 11)
o Neutral solutions (like water) have a pH of exactly 7
6.2 pH Indicators
 The most accurate way of measuring pH is using a pH meter
 An approximate way of measuring pH is using an indicator
 An indicator can tell the pH by with colour changes
Universal Indicator
0 1 2 3 4 5 6 7 8 9 10 11 12 13 14
Red Orange Yellow Green Green-Blue Blue Violet
Litmus
0 1 2 3 4 5 6 7 8 9 10 11 12 13 14
Red Blue
Phenolphthalein
0 1 2 3 4 5 6 7 8 9 10 11 12 13 14
Colourless Pink - Purple
Methyl Orange
0 1 2 3 4 5 6 7 8 9 10 11 12 13 14
Red Yellow
Screened Methyl Orange
0 1 2 3 4 5 6 7 8 9 10 11 12 13 14
Red Grey Green
Bromothymol Blue
0 1 2 3 4 5 6 7 8 9 10 11 12 13 14
Yellow Blue

6.3 Acids
Acids are substances that produces H+ ions in Water.
Properties of Acids
 Acids are sour
o Lemon's sourness is from citric acid
o Vinegar's sourness is from ethanoic acid
 Acids have pH of less than 7
o Acids change Blue Litmus paper Red
 Organic acids are acids that contain the -COOH group
o Examples of organic acids are Ethanoic Acid and Citric Acid
 Mineral acids are acids that are not organic
o Mineral acids are much stronger acids than organic acids
o Examples of mineral acids are Hydrochloric Acid and Nitric Acid
6.4 Reactions of Acids
Reaction of Acids with Bases
 Acids will react with bases to form an inorganic Salt and water only
 This reaction is called Neutralization
HX(aq) + ZOH(aq) → ZX(aq) + H2O(l)
Example 1
HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l)
Example 2
H2SO4(aq) + Ca(OH)2(aq) → CaSO4(aq) + 2 H2O(l)
Reaction of Acids with Metals
 Acids will react with Metal to form an inorganic Salt and Hydrogen Gas only
 This is due to displacement (to be covered under the Reactivity Series in the topic of Metals)
2 HX(aq) + 2 M(s) → 2 MX(aq) + H2(g)
Example 1
2 HCl(aq) + 2 Na(s) → 2 NaCl(aq) + H2(g)
Example 2
H2SO4(aq) + Mg(s) → MgSO4(aq) + H2(g)
Reaction of Acids with Carbonates
 Acids will react with Carbonates to form an inorganic salt, carbon dioxide and water.
2 HX(aq) + ZCO3(aq) → ZX2(aq) + CO2(g) + H2O(l)
Example 1
2 HCl(aq) + Na2CO3(s) → 2 NaCl(aq) + CO2(g) + H2O(l)
Example 2
H2SO4(aq) + CaCO3(s) → CaSO4(s) + CO2(g) + H2O(l)

Common Acids
Acids Formulae Cation Anion
Hydrochloric Acid HCl H+ Cl-
Sulphuric Acid H2SO4 H+ SO42-
Nitric Acid HNO3 H+ NO3-
Ethanoic Acid CH3COOH H+ CH3COO-
Phosphoric Acid H3PO4 H+ PO43-
Hydrofluoric Acid HF H+ F-
Hydrobromic Acid HBr H+ Br-
Hydroiodic Acid HI H+ I-
Carbonic Acid H2CO3 H+ CO32-
6.5 Bases and Alkalis
Bases are substances that react with an Acid to form a Salt and Water Only
Properties of Bases
 Edible bases taste bitter
 Alkalis feel slippery
 Bases have pH of more than 7
o Bases change Red Litmus paper Blue
 Bases are usually Metal oxides or Metal hydroxides
 Soluble Bases are called alkalis
o Group I hydroxides are readily soluble
o Group II hydroxides are sparingly soluble
o Group III or Transition Metal hydroxides are generally insoluble
6.6 Reactions of Bases
Reaction of Bases with Ammonium Salts
 Bases will react with Ammonium Salts to form an inorganic salt, ammonia gas and water
NH4X(aq) + ZOH(aq) → ZX(aq) + NH3(g) + H2O(l)
Example 1
NH4Cl(aq) + NaOH(aq) → NaCl(aq) + NH3(g) + H2O(l)
Example 2
(NH4)2SO4(aq) + Ca(OH)2(aq) → CaSO4(aq) + NH3(g) + 2 H2O(l)
Precipitation of Insoluble hydroxides
 Alkalis are used to precipitate out insoluble hydroxides from solutions of their salt
Example 1
2 NaOHaq) + CuSO4(s) → Na2SO4(aq) + Cu(OH)2(s) (Blue precipitate)
Example 2
2 NaOHaq) + MgCl2(s) → 2 NaCl(aq) + Mg(OH)2(s) (White precipitate)

Common Bases
Bases Formulae Cation Anion
Sodium Hydroxide NaOH Na+ OH-
Potassium Hydroxide KOH K+ OH-
Ammonium Hydroxide NH4OH NH4+ OH-
Calcium Hydroxide Ca(OH)2 Ca2+ OH-
Magnesium Hydroxide Mg(OH)2 Mg2+ OH-
Barium Hydroxide Ba(OH)2 Ba2+ OH-
Aluminum Hydroxide Al(OH)3 Al3+ OH-
Zinc Hydroxide Zn(OH)2 Zn2+ OH-
6.7 Oxides
 Oxides are formed when elements burn in Oxygen
 There are 4 types of Oxides: Acidic, Basic, Amphoteric and Neutral
 Non-Metallic oxides are acidic
o They have similar properties as acids, as well as undergo similar reactions as acids
o They form acids when dissolved in water
 Carbon Dioxide dissolves in water to form carbonic acid (H2CO3)
 Metallic oxides are basic
o They have similar properties as bases, as well as undergo similar reactions as bases
 Some Metallic oxides are amphoteric
o They show both acidic and basic properties
 They can neutralize both acids and bases.
 Aluminum oxide can react with hydrochloric acid to form aluminum chloride and water
 Aluminum oxide can react with sodium hydroxide to form sodium aluminate and water
o Some examples are Aluminum Oxide (Al2O3), Zinc Oxide (ZnO) and Lead (II)
 Some Non-Metallic oxides are neutral (Pure)
o They show neither acidic nor basic properties
o Some examples are Dihydrogen Oxide, Carbon Monoxide (CO) and Nitrogen Oxide (NO)
6.8 Solubility of Various Salts
Soluble Salts Insoluble Salts
All Nitrates
All Lead (II) Salts except Lead (II) NitrateAll Chlorides except Lead (II) Chloride
All Sulphate except Lead (II) Sulphate
All Group I and Ammonium Salts
All Carbonates except for Group I and
Ammonium Carbonates
All Sulphides (s2-) except for Group I and
Ammonium Sulphides
All hydroxides except for Group I and
Ammonium hydroxides
All oxides except for Group I and
Ammonium Oxides
Also Insoluble: Barium Sulphate, Calcium Sulphate and Silver Chloride
Sparingly soluble: Group II Hydroxides and Oxides

6.9 Preparation of Salts
 The method to prepare a salt depends on its solubility
6.10 Precipitation method
 The precipitation method to prepare an Insoluble salt
Step 1: Choosing the reactants. (They must be soluble)
Step 2: Mix the reactants.
Step 3: Wait for the insoluble salt to precipitate out. Stirring or heating may speed the reaction along.
Step 4: Filter out the insoluble salt.
Salt
Insoluble
Precipitation
Method
Soluble
Group I or
Ammonium
Salt
Titration
NOT Group I
or Ammonium
Salt
Crystalization
React Metal
with Acid
React Metal
Carbonate
with Acid
React Metal
Oxide with
Acid

6.11 Crystallization method
 The crystallization method is used to prepare a soluble salt that does not contain Group I or
Ammonium ions.
Step 1: Choosing the reactants. (Use the acid containing the anion, and the metal/ oxide/ carbonate.)
Step 2: Mix the reactants. Let the metal/ metal oxide/ metal carbonate be in excess.
Step 3: Wait for the reaction to complete. Stirring or heating may speed the reaction along.
Step 4: Filter out the excess metal/ metal oxide/ metal carbonate
Step 5: Heat the filtrate till saturated (when a thin layer of crystals are formed.
Step 6: Leave the filtrate to cool for more crystals to form.
Step 7: Filter out the crystals and dry
6.12 Titration method
 The Titration method is used to prepare a soluble salt that contains Group I or Ammonium ions.
 This method is based on the neutralization reaction.
Step 1: Choosing the reactants. (Use the acid containing the anion & the hydroxide containing the cation.)
Step 2: Pipette out 25 cm3 of one of the reactants into a conical flask. Add a few drops of indicator
Step 3: Add the other reactant into a burette
Step 4: Add the reactant in the burette into the conical flask drop by drop. Stop when the colour changes.
Step 5: Heat the filtrate till saturated (when a thin layer of crystals are formed.
Step 6: Leave the filtrate to cool for more crystals to form.
Step 7: Filter out the crystals and dry

Chapter 7
Periodic Table
7.1 Periodic Trends (An Overview)
 The Periodic Table is an arrangement of elements with an increasing number of Protons
o Number of valence electrons increases across a period (Left to right) from 1 to 8.
 The Periodic Table is arranged in vertical groups and horizontal periods.
o Elements of the same period have the same number of electron shells.
o Elements of the same group have the same number of valence electrons.
o Elements of the same group have similar chemical properties & form compounds with similar
chemical formulae.
 A zig-zag diagonal line divides the metals and non metals.
o Elements near the line are called metalloids & have characteristics of both metals & non-metals.
7.2 Electronegativity
 Electronegativity refers to the ability to gain electrons
 Less electronegative elements have a greater metallic character.
 Elements become more electronegative across the period (Left to right)
o Elements become less metallic across the period
 Elements become less electronegative down a group (Top to Bottom)
o Elements become more metallic down a group
 Chlorine is the most electronegative element, while Francium is the least Electronegative.
Name of Various Groups of Elements
Note: Only Group I (Alkali Metals), Group VII (Halogens), Group VIII (Noble Gases) and Transition Metals are examinable.

7.3 Group I Metals: Alkali Metals
 Group I Elements reacts with water to form Alkalis and hydrogen gas (Thus the name Alkali Metals).
2 Na(s) + 2 H2O(l) → 2 NaOH(aq) + H2(g)
 Group I Elements reacts with air to form Basic Oxides.
o These basic oxides dissolve in water to form alkalis
4 Na(s) + O2(g) → 2Na2O(s)
 Hot Group I metals react with chlorine gas to form Metal Chlorides Salts
2 Na(s) + Cl2(g) → 2NaCl(s)
 Group I Metals are very reactive (the most reactive in the periodic table)
o They easily react with cold water and air, and thus have to be stored in oil or vacuum.
o The reactivity of Group I metals increases down the group.
 Lithium reacts violently with water
 Sodium reacts very violently with water
 Potassium reacts explosively with water.
 Group I Metals are strong reducing agents
 Physical properties of Group I Metals include:
o They easily react with cold water and air, and thus have to be stored in oil or vacuum.
o They have low densities.
 The densities increases down the group.
o They have low melting points.
 The melting point increases down the group.
o They are shiny and silvery solids.
o They are very soft and can be easily cut with a knife or razor.
o They are good conductors of electricity and heat.

7.4 Group VII: Halogens
 Group VII Halogens form diatomic molecules with a single covalent bond
 Trends of Group VII Halogens include:
o The melting and boiling point increases down the group.
 Fluorine and Chlorine are gaseous at room temperature
 Bromine is liquid at room temperature
 Iodine and Astatine are solids at room temperature
o The colours of the Halogens get darker down the group.
 Fluorine is pale yellow
 Chlorine is yellowish green
 Bromine is reddish brown
 Iodine and Astatine are black
o The Halogens get less reactive down the group
 Fluorine is the most reactive, and astatine is the least reactive
 Group VII Halogens are strong oxidizing agents
 Halogens undergo displacement reactions.
o A halogen in a salt can be displaced by a more reactive halogen
 For example; Fluorine is the most reactive, and astatine is the least reactive, hence the
astatine in an astatine salt can be displaced by fluorine gas
2 NaAt(s) + F2(g) → 2 NaF(s) + At2(s)
 Group VIII Halogens reacts with water to form Acids or Acidic Solutions
Cl2 (g) + 2 H2O(l) → HCl(aq) + HOCl(aq)
 Group VII Halogens react with metals to form ionic halides (Salts)
2 Na(s) + Cl2(g) → 2NaCl(s)
 Group VII Halogens react with bases and alkalis to form ionic halides (Salts) due to its acidic nature.
2 NaOH(s) + Cl2(g) → NaOCl(s) + NaCl(s) + H2O(l)
 Physical properties of Group VII Halogens include:
o They have low boiling and melting points.
o They do not conduct heat or electricity in any state.
o They are sparingly soluble in water
o They are soluble in organic solvent like CCl4.
 Some uses of Halogens
o Fluoride is used in toothpaste to prevent tooth decay
o Chlorine and Iodine is used to kill bacteria
o Iodine is needed by the human body for proper thyroid gland function

7.5 Group VIII: Noble Gases
 Group VIII Noble Gases are chemically unreactive due to their stable octet structure (With fully filled
valence shells)
 Noble Gases exists as monoatomic gases at room temperature.
o Noble gases have very low boiling and melting points
 Some uses of Noble Gases
o Argon is used to fill light bulbs
o Neon is used to fill coloured glowing tubes
o Helium is used to fill weather balloons
7.6 Transition Metals
 Properties of Transition Metals include:
o They are strong hard metals with high boiling and melting points.
o They have high density
o They form coloured compounds
 Iron (II) oxide is green, while Iron (III) Oxide (Rust) is reddish brown.
o They form ions with variable charges
 Iron can form Fe2+ and Fe3+ ions.
o They are used as catalyst
o They are used to make alloys
 Steel is an alloy comprising of Iron and Carbon
Summary (Trends)

Chapter 8
Metals
8.1 Physical Properties of Metals
 High Boiling and Melting point.
o General exception to these are Mercury and Group I and Group II Metals
 Good Conductor of Heat and Electricity
o Due to sea of delocalized electrons
 Malleable (Ability of being flattened) & Ductile (Ability to be pulled into wires)
o Due to metallic bonding, in which the layers of atoms can easily slide over each other.
8.2 Alloys
 An alloy is a mixture of metal with another element
o This second element may be both either a metal or a non metal
 Pure metals are usually too soft to be used.
 Alloys strengthens metals to be used by disrupting the orderly arrangement of the metal atoms with
foreign atoms of different sizes.
 Some metals, like iron, oxidize or rusts easily.
o Hence alloys of these metals may be used in place of the metals due to their resistance to
oxidization or corrosion.
Examples:
Alloys Constituent Elements Uses
Bronze Copper and Tin Trophies
Brass Copper and Zinc Musical Instruments and Electrical plug pins
Pewter Tin, Antimony and Copper Dinnerware like plates and teapots
Industrial Steel Iron and Carbon Scaffoldings
Stainless Steel Iron, Chromium and Nickel Cutlery and surgical instruments
Chromium Steel Iron and Chromium Ball Bearings
High Speed Steel Tungsten and Vanadium High Speed Drills
8.3 Reactivity Series
 Metals differ greatly in their chemical reactivity
 Very reactive metals are unstable as a metal, but form very stable compounds
o These metals are not found uncombined in nature
 Less reactive metals are more stable as a metal.
o These metals can be found uncombined in nature
Most Reactive
Potassium Sodium Calcium Magnesium Aluminum Carbon Zinc Iron
K Na Ca Mg Al C Zn Fe
Potato Salad Can Make A Cunning Zebra Itchy
Tin Lead Hydrogen Copper Mercury Silver Gold
Sn Pl H Cu Hg Ag Au
These Large Helicopters Can Make Some Giddy
Least Reactive

8.4 Chemical Reactions of Metal
 All metals undergo displacement reactions
o A more reactive metal is able to displace a less reactive metal from its compounds
o E.g. Displacement of the less reactive copper by the more reactive zinc
Zn(s) + CuSO4 (aq) → Cu(s) + ZnSO4 (aq)
o E.g. Displacement of the less reactive lead by the more reactive magnesium
Mg(s) + PbO (s) → Pb(s) + MgO (s)
 Most Metals can react with water to produce hydrogen gas and either hydroxides or metal oxides.
o E.g. Reaction of Sodium with cold water
2Na(s) + 2H2O(l) → 2NaOH(aq) + H2 (g)
o E.g. Reaction of Aluminum with steam
2Al(s) + 3H2O(l) → Al2O3 (aq) + 3H2 (g)
 Most Metals can react with acids to produce salts and hydrogen gas
o E.g. Reaction of Calcium with Hydrochloric Acid
Ca(s) + 2HCl (aq) → CaCl2 (aq) H2 (g)
Metals Reaction with Water Reaction with Acids
Potassium K Reacts vigorouslywith cold water to
produce hydroxides and hydrogen gas
Reacts explosively with acids to produce
salts and hydrogen gasSodium Na
Calcium Ca
Reacts slowlywith cold water to produce
hydroxides and hydrogen gas
Reacts vigorouslywith acids to produce
salts and hydrogen gas
Magnesium Mg
Reacts with Steam to produce metal oxides
and hydrogen gas
Reacts rapidlywith warm acids to
produce salts and hydrogen gas
Aluminum Al
Reacts slowlywith hot acids to produce
salts and hydrogen gas
Zinc Zn
Iron Fe
Tin Sn
Does notreact to water or steam
Lead Pb
Copper Cu
Does notreact with acids
Mercury Hg
Silver Ag
Gold Au

8.5 Chemical Reactions of Metal Compounds
 The oxides of the less reactive metals can be reduced by carbon to produce the metal & CO2 only.
o E.g. Reduction of Zinc oxide
C(s) + 2ZnO(s) → 2Zn(s) + CO2 (g)
 The oxides of the less reactive metals can be reduced by hydrogen to produce the metal & H2O only.
o E.g. Reduction of Lead oxide
PbO(s) + H2(g) → Pb (s) + H2O (l)
 Some of the carbonates of the less reactive metals can be decomposed upon heating.
o E.g. Decomposition of Copper Carbonate
CuCO3(s) → CuO (s) + CO2 (g)
o E.g. Decomposition of Silver Carbonate
2Ag2CO3(s) → 4Ag (s) + O2 (g) + 2CO2 (g)
Metals
Reduction of Oxides
by Carbon
Reduction of Oxides
by Hydrogen
Decomposition of Metal
Carbonates
Potassium K
Does notreduce Does notreduce
Does notdecompose
Sodium Na
Calcium Ca
Metal carbonate decomposes
upon heating into metal oxides
and carbon dioxide gas
Magnesium Mg
Aluminum Al
Zinc Zn
Metal Oxides get reduced
by carbon to form metal
and carbon dioxide only
Metal Oxides get reduced
by hydrogen to form
metal and water only
Iron Fe
Tin Sn
Lead Pb
Copper Cu
Mercury Hg Carbonate decomposes upon
heating into metal,O2 and CO2
gas
Silver Ag
Gold Au
8.6 Extraction of Metals
 Metals high up in the reactivity series do not exist in the free state, but can be found in the form of
metal ores or metal salts
o These ores are normally in the form of oxides, sulphides or carbonates.
o Some common ores include
Metal Ore Compound in Ore
Aluminum Bauxite Aluminum Oxide
Iron Haematite Iron (III) Oxide
Zinc Zinc Blende Zinc Sulphide
 Very reactive metals form very stable compounds, hence they can only be extracted by electrolysis.
o All metals above carbon on the reactivity series can only be extracted by electrolysis.
 Less reactive metals can be extracted through reduction by carbon, hydrogen or more reactive metals
o All metals below carbon on the reactivity series can only be extracted through reduction
 Metals with low reactivity can be found usually in the free state, or can be extracted through heating.
o This refers to metals like mercury, silver and gold.

8.7 Extraction of Iron
 Iron is extracted from haematite (Fe2O3) through reduction in a blast furnace.
1. Haematite (Iron (III) Oxide), Limestone (Calcium Carbonate) and Coke (Carbon) are fed into the
blast furnace
2. Hot air is fed into the bottom of the furnace
3. The Coke burns in the hot air to form carbon dioxide.
C (S) + O2 (g) → CO2 (g)
4. The carbon dioxide is further reduced to carbon monoxide
CO2 (g) + C (S) → 2CO (g)
5. The carbon monoxide reduces the haematite to iron
CO (g) + Fe2O3 (s) → 2Fe (s) + CO2 (g)
6. The molten iron form is filled with sand particles, which can be removed using the limestone
7. The Limestone is decomposed by heat to produce calcium oxide and carbon dioxide.
CaCO3 (g) → CaO (s) + CO2 (g)
8. The calcium oxide reacts with the sand to form slag (calcium silicate)
CaO (g) + SiO2 (s) → CaSiO3 (s)
9. The slag is less dense than iron. Hence it floats on the molten iron & is removed from the top tap
10. The molten iron is removed from the bottom tap

8.8 Recycling of metals
 Metal ores resources are finite and limited. Hence it is important to recycle metals
 Some advantages of recycling include
o Recycling saves energy required to extract metals from ores
o Recycling reduce emission of greenhouse gases like carbon dioxide, produced in the extraction of
metals like iron from ores
o Recycling preserves scarce non-renewable raw materials
o Recycling reduces environmental air pollution and water pollution.
o Recycling reduced the amount of land required for the disposal of metals through landfills
o It is cheaper to recycle some metals like aluminum, than to extract them from the earth's crust
 Some disadvantages of recycling include
o Recycling is a time consuming process
o Recycling takes up a high amount of effort and human resources.
8.9 Rusting of Iron
 In the presence of water and oxygen, Iron rusts
 Some methods of preventing rust include:
o Coating the iron with a substance to prevent air and water from coming into contact with the metal
surface. These substances include
 Paint
 Oil or grease
 Electroplating the iron with a less reactive metal like tin or copper
o Sacrificial protection.

Chapter 9
Electrolysis
9.1 Electrochemical Cells
 There are 2 main types of Electrochemical Cells:
o Electrolytic Cells
 An Electrolytic cell has a battery or a direct current (D.C) supply to facilitate a chemical
reaction.
o Galvanic Cells (Also Known as Simple Cells)
 A simple cell produces an electrical charge
 It does not have a battery or a D.C Supply.
 It usually has a voltmeter to measure the voltage produced
9.2 Introduction to Electrolysis: Set-up and Apparatus
 Electrolysis is the process of decomposing an ionic compound into its constituent element through the
conduction of electricity through the ionic compound in the molten or aqueous state.
 The electric current facilitates Redox reactions at the electrodes
 Components of an Electrolytic cell:
o A Battery or D.C supply
 The positive end is drawn as the longer line
o 2 Electrodes (one positive and one negative)
 Terminals through which electrons enter or leave the electrolyte
 The electrode attached to the positive end of the battery is positive and vice versa.
 Some electrodes may be inert (is not involved in the electrolysis)
 Graphite (Carbon) and platinum are inert electrodes
o Electrolyte
 A Molten compound OR an aqueous solution which conducts electricity.
 Aqueous solutions of Acids, Bases or Salts can be used as the electrolyte.
 Molten Salts can be used as the electrolyte.
 Strong conductors of electricity are classified as strong electrolytes
 Examples include aqueous HCl, H2SO4, NaOH, KOH, NaCl and KI
 Weak conductors of electricity are classified as weak electrolytes
 Examples include weak acids and bases like Ethanoic Acid and Limewater
 Non-conductors of electricity are known as non-electrolytes
 Examples include Organic compounds like CCl4 or Sugar, and pure water
 The negative ions (Anions) are attracted to the positive electrode (Anode)
o The anions are oxidized at the anode
 The positive ions (Cations) are attracted to the negative electrode (Cathode)
o The cations are reduced at the cathode.
 Note: Electric charges travels from the positive end to the negative end, across the battery.

9.3 Electrolysis of Molten Electrolytes
 A molten ionic compound is used as the electrolyte
 Inert electrodes are used
 The metallic cations are reduced at the cathode to form metal atoms
 The non-metallic anions are oxidized at the anode to form a simple covalent molecule
 Note: Heat may be necessary to keep the ionic compound molten
Example 1: Electrolysis of Molten Sodium Chloride (NaCl)
At the Carbon Cathode (Negative Electrode), Sodium metal is discharged
Na+(l) + e- → Na(s)
At the Carbon Anode (Positive Electrode), Chlorine gas is discharged
2Cl-(l) → Cl2(g) + 2e-
Overall Reaction
2NaCl(l) → Cl2(g) + 2Na(s)
Example 2: Electrolysis of Molten Magnesium Oxide (MgO)
At the Carbon Cathode (Negative Electrode), Magnesium metal is discharged
Mg2+(l) + 2e- → Mg(s)
At the Carbon Anode (Positive Electrode), Oxygen gas is discharged.
2O2-(l) → O2(g) + 4e-
Overall Reaction
2MgO(l) → O2(g) + 2Mg(s)
9.4 Uses of Electrolysis of Molten Electrolytes
 Electrolysis of molten electrolytes can be used to extract reactive metals from their ores
o Metals more reactive than zinc is extracted by electrolysis
Example 1: Extraction of Aluminium from Bauxite
At the Carbon Cathode (Negative Electrode), Aluminium metal is discharged
Al3+(l) + 3e- → Al(l)
At the Carbon Anode (Positive Electrode), Oxygen gas is discharged.
2O2-(l) → O2(g) + 4e-
Overall Reaction
2Al2O3(l) → 3 O2(g) + 4Al(l)

9.5 Preferential Discharge
 In the electrolysis of aqueous solutions, there are several cations or anions present.
 When there are several cations or anions, the ions that are preferentially discharged are dependent
on the following few criteria:
o The position of the ion on the electrochemical series
 An ion lower on the electrochemical series may be discharged first.
o Concentration of Ion
 An ion of higher concentration may be discharged first
Electrochemical Series: Cations
Ease of Discharge
In Concentrated Solutions In Dilute Solutions
Cations Product Cations Product
Difficult K+, H+ H2 K+, H+ H2
Na+, H+ H2 Na+, H+ H2
Ca2+, H+ H2 Ca2+, H+ H2
Mg2+, H+ H2 Mg2+, H+ H2
Al3+, H+ H2 Al3+, H+ H2
Zn2+, H+ H2 Zn2+, H+ H2
Ni2+, H+ Ni Ni2+, H+ H2
Pb2+, H+ Pb Pb2+, H+ H2
H+ H2 H+ H2
Cu2+, H+ Cu Cu2+, H+ Cu
Easy Ag+, H+ Ag Ag+, H+ Ag
Electrochemical Series: Anions
Ease of Discharge
In Concentrated Solutions In Dilute Solutions
Anions Product Anions Product
Difficult SO42-, OH- O2 (+ H2O) SO42-, OH- O2 (+ H2O)
NO32-, OH- O2 (+ H2O) NO32-, OH- O2 (+ H2O)
Cl-, OH-
Cl2 Cl-, OH- O2 (+ H2O)
Br-, OH-
Br2 Br-, OH- O2 (+ H2O)
I-, OH-
I2 I-, OH- O2 (+ H2O)
Easy OH- O2 (+ H2O) OH- O2 (+ H2O)
Example 1: Electrolysis of Dilute Aqueous Sodium Chloride (NaCl)
At the Carbon Cathode (Negative Electrode),
 The Na+ and H+ ions are attracted to the cathode
 The H+ ions are preferentially discharged (Hydrogen gas is discharged)
2H+(aq) + 2e- → H2(s)
At the Carbon Anode (Positive Electrode),
 The Cl- and OH- ions are attracted to the cathode
 The OH- ions are preferentially discharged (Oxygen gas is discharged)
4OH-(aq) → O2(g) + 2H2O(g) + 4e-
Overall Reaction
2H2O(l) → 2H2(g) + 2O2(g)

Example 2: Electrolysis of Concentrated Aqueous Sodium Chloride (NaCl)
 The Na+ and H+ ions are attracted to the cathode
2H+(aq) + 2e- → H2(g)
 The Cl- ions are preferentially discharged (Chlorine gas is discharged)
2Cl-(aq) → Cl2(g) + 2e-
Overall Reaction
2H+(aq) + 2Cl-(aq) → H2(g) + Cl2(g)
↓
2H+(aq) + 2OH-(aq) + 2Na+(aq) + 2Cl-(aq) → H2(g) + Cl2(g) + 2Na+(aq)+ 2OH-(aq)
↓
2NaCl(aq) + 2H2O(l) → 2NaOH(aq) + H2(g) + Cl2(g)
Example 3: Electrolysis of Concentrated Aqueous Copper Chloride (CuCl2)
 The Cu2+ and H+ ions are attracted to the cathode
 The Cu2+ ions are preferentially discharged (Copper Metal is discharged)
Cu2+(aq) + 2e- → Cu(s)
 The Cl- ions are preferentially discharged (Chlorine gas is discharged)
2Cl-(aq) → Cl2(g) + 2e-
Overall Reaction
Cu2+(aq) + 2Cl-(aq) → Cu(s) + Cl2(g)
Observations: The bluish-green colour of the solution is discharged, and pink deposits are formed on the
cathode
9.6 Uses of Electrolysis of Aqueous Compounds
 The electrolysis of aqueous solutions can be used to manufacture various chemical compounds
Example 1: Electrolysis of Concentrated Aqueous Sodium Chloride (NaCl)
At the Carbon Cathode (Negative Electrode), Hydrogen gas is discharged
At the Carbon Anode (Positive Electrode), Chlorine gas is discharged
Overall Reaction
2NaCl(aq) + 2H2O(l) → 2NaOH(aq) + H2(g) + Cl2(g)
Sodium hydroxide and chlorine gas is produced from the Electrolysis of Concentrated Aqueous Sodium
Chloride

9.7 Reactive Electrodes
 If reactive (active) electrodes are used, the anode could be oxidized in place of the anions in the
electrolyte
o Note: the cation in the electrolyte would usually contain the cationic form of the reactive anode
o The cation would then be reduced at the cathode
Example 1: Electrolysis ofAqueous Copper Sulphate (CuSO4) with copper electrodes
At the Copper Anode (Positive Electrode),
 The SO4
2-
and OH-
ions are attracted to the cathode, but neither is discharged as the copper atoms in the
electrode loses electrons more readilythan either anions
 Cu atoms loses 2 electrons to form Cu2+
ions.
 The copper anode becomes smaller over time,and eventually completelydissolves in the electrolyte.
Cu(s) → Cu2+
(aq) + 2e-
At the Copper Cathode (Negative Electrode),
 The Cu2+
and H+
ions are attracted to the cathode
 The Cu2+
ions are preferentiallydischarged (Copper metal is deposited on the cathode)
 The copper cathode becomes larger over time
Cu2+
(aq) + 2e-
→ Cu(s)
Overall Reaction:There is no change in the electrolyte.
9.8 Uses of Reactive Electrodes: Electroplating
 Electroplating can be used to prevent rust or corrosion of metals by covering the metal with a layer of
another metal (that is usually corrosion resistant.
o The item to be plated is to be the cathode
 Galvanizing is a form of electroplating where an iron or steel item is covered in a layer of zinc
 In the set-up for electroplating:
 The item must be able to conduct electricity
o The metal used to coat the cathode is to be the anode
 The electrolyte is to be an aqueous salt of the anode.
Example 1: Electroplating an aluminium statuette with copper
Electrode:Copper anode and Aluminium (statuette) cathode
Electrolyte: Aqueous Copper Sulphate
At the Copper Anode (Positive Electrode),
 The copper atoms loses 2 electrons to form Cu2+
ions.
Cu(s) → Cu2+
(aq) + 2e-
At the Aluminium (Statuette) Cathode (Negative Electrode),
 The Cu2+
and H+
 The Cu2+
 The Aluminium (statuette) cathode becomes coated in a layer of copper
Cu2+
(aq) + 2e-
→ Cu(s)

9.9 Galvanic Cell (Simple Cells)
 A Galvanic Cell produces an electrical charge by converting chemical energy to electrical energy
 A galvanic cell is also known as a simple cell
 Components of a galvanic cell:
o NO Battery or D.C supply
o 2 Electrodes of different metals
 The more reactive metal will oxidize, and be the negative electrode (the anode)
 The less reactive metal will become the positive electrode (the cathode)
o An Electrolyte solution
 Sometimes, 2 different electrolyte solutions can be used, connected by a salt bridge
o A Voltmeter is often included to measure the voltage produced by the galvanic cell
 The voltage produced depends on the positions of the two metal electrodes in the reactivity
series
 The further apart the metals are on the metal reactivity series, the greater the voltage
produced
Example 1: A Galvanic Cell with Copper and Zinc Electrodes,with Sulphuric Acid as the electrolyte
 As zinc is more reactive than copper the zinc will lose 2 electrons and be oxidized.
o Hence the zinc electrode becomes the negative electrode (the anode)
o Therefore the copper electrode is the positive electrode (the cathode)
At the Zinc Anode (Positive Electrode),
 The SO4
2-
and OH-
ions are attracted to the cathode, but neither is discharged as the zinc atoms in the electrode
loses electrons more readilythan either anions
 The zinc atoms loses 2 electrons to form Zn2+ ions.
Zn(s) → Zn2+
(aq) + 2e-
 The Zn2+
and H+
2H+(aq) + 2e- → H2(g)

Example 2: A Galvanic Cell with Copper and Zinc Electrodes, with Zinc sulphate and Copper sulphate as the
electrolyte solutions,connected by a saltbridge
 As zinc is more reactive than copper the zinc will lose 2 electrons and be oxidized.
o Hence the zinc electrode becomes the negative electrode (the anode)
o Therefore the copper electrode is the positive electrode (the cathode)
At the Zinc Anode (Positive Electrode),
 The SO4
2-
and OH-
ions are attracted to the cathode, but neither is discharged as the zinc atoms in the electrode
loses electrons more readilythan either anions
 The zinc atoms loses 2 electrons to form Zn2+
ions.
Zn(s) → Zn2+
(aq) + 2e-
 The Cu2+
and H+
 The Cu2+
Cu2+
(aq) + 2e-
→ Cu(s)
Overall Reaction
Zn(s) + Cu2+(aq) → Zn2+
(aq) + Cu(s)

Chapter 10
Organic Chemistry
10.1 Introduction to Organic Chemistry
 Organic Chemistry is the study of Carbon based compounds
o Except for Carbon Monoxide, Carbon Dioxide and metal Carbonates
 Most organic compounds also have hydrogen, and some also have oxygen.
o Organic compounds with carbon and hydrogen atoms are called hydrocarbons
 Important reminder: Carbon Atoms only can form 4 bonds around each one
 In organic chemistry, there are many compounds with similar chemical properties, and have a general
formula
o This family of compounds is known as a homologeous series
o All compounds in a homologus series typically have a common functional group, and differ by a -
CH2 unit.
o Compounds in a homologeous series have similar chemical properties but different physical
properties like boiling and melting points
 4 Main homologeous series to be taught include
o Alkanes
 Has no functional group
 Has the suffix -ane
 Has the general formula of CnH2n + 2
o Alkenes
 has a double bond between 2 carbon atoms
 Has the C=C functional group
 Has the suffix -ene
 Has the general formula of CnH2n
o Alcohols
 Has the -OH functional group
 Has the suffix -anol
 Has the general formula of CnH2n + 1OH
o Carboxylic Acid
 Has the -COOH functional group
 Has the suffix -anoic Acid
 Has the general formula of CnH2nO2 or Cn-1H2n-1COOH
No. Of Carbons
1 2 3 4
Meth- Eth- Prop- But-
Functional
Groups
Alkanes
CnH2n + 2
-ane
Methane Ethane Propane Butane
CH4 C2H6 C3H8 C4H10
Alkenes
CnH2n
-ene
Ethene Propene Butene
C2H4 C3H6 C4H8
Alcohol
CnH2n + 1OH
-anol
Methanol Ethanol Propanol Butanol
CH3OH C2H5OH C3H7OH C4H9OH
Carboxylic
Acid
CnH2nO2
- anoic
Acid
Methanoic Acid
Ethanoic
Acid
Propanoic
Acid
Butanoic Acid
HCOOH CH3COOH C2H5COOH C3H7COOH

10.2 Petroleum and Crude Oil
 Crude oil is a mixture of many thousands of different hydrocarbons with different properties.
o To make crude oil useful, batches of similar compounds with similar properties need to be sorted
and separated by fractional distillation.
o In fractional distillation, the crude oil is heated to make it vaporise. The vapour is then
cooled.
o Different fractions of the oil are collected at different temperatures.
 The larger hydrocarbons are not as useful as the smaller hydrocarbon
o Cracking is a process that can be used to break larger hydrocarbons into smaller ones
o Cracking is done by passing the vaporised hydrocarbon over a solid catalyst
o Cracking produces alkenes

10.3 Isomerism
 Organic compounds with same chemical formula but different structural formula are known as
isomers
Example 1: Isomers of Butane Chemical Formula: C4H10
Example 2: Isomers of Butanol Chemical Formula: C4H9OH
Example 3: Propane Chemical Formula: C3H8
(EXAMPLE OF WHAT'S NOT A SET OF ISOMERS)
ALL 3 are the same structure

10.4 Alkanes
 Alkanes are organic compounds with only Carbon and Hydrogen atoms with NO Functional groups
 Alkanes have the general formula of CnH2n + 2
 Alkanes are generally unreactive.
Methane Ethane Propane Butane
HCH3 CH3CH3 CH3CH2CH3
CH3CH2CH2CH3
Or
CH3CHCH3
CH3
10.4a Combustion
 Alkanes undergo Combustion
o Alkanes burn in Oxygen to form Carbon Dioxide and Water Vapour
Example 1: Combustion of Methane, CH4
CH4 + 2O2 → CO2 + 2H2O
Example 2: Combustion of Butane, C4H10
2C4H10 + 13O2 → 8CO2 + 10H2O
10.4b Substitution Reaction
 In the presence of light, Alkanes undergo Substitution reaction with halogens
o Observation: The coloured halogens will decolourise
Example 1: Substitution Reaction of Methane, CH4 with chlorine gas, Cl2
CH4 + Cl2 → CH3Cl + HCl
Example 2: Substitution Reaction of Butane, C4H10 with bromine, Br2
C4H10 + Br2 → C4H9Br + HBr
Note: The substitution reaction could proceed further
C4H10 + Br2 → C4H9Br + HBr
C4H9Br + Br2 → C4H8Br2 + HBr
C4H8Br2 + Br2 → C4H7Br3 + HBr

10.5 Alkenes
 Alkenes are organic compounds with a double bond between 2 carbon atoms
 Alkanes have the general formula of CnH2n
 Alkenes are unsaturated organic compounds (Has 1 or more double bonds)
o Alkanes are saturated organic compounds (Has no double bonds)
Ethene Propene Butene
H2C=CH2 H2C=CHCH3
H2C=CHCH2CH3
Or
H2C=CCH3
CH3
Or
CH3CH=CHCH3
10.5a Combustion
 Alkenes undergo Combustion
o Alkenes burn in Oxygen to form Carbon Dioxide and Water Vapour
Example 1: Combustion of Butene, C4H8
C4H8 + 6O2 → 4CO2 + 4H2O
10.5b Substitution Reaction
 In the presence of light, Alkenes can undergo Substitution reaction with halogens
10.5c Addition Reaction with Halogens
 Alkenes can undergo addition reactions with halogens in the absence of light
Example 1: Addition of aqueous bromine, Br2 to Ethene, C2H4
C2H4 + Br2 → C2H4Br2
Note: This can be used as a test to differentiate Alkanes from Alkenes. Alkenes can decolourise bromine
in the absence of light, while Alkanes cannot decolourise bromine in the absence of light
Example 2: Addition of aqueous bromine, Br2 to Butene, C4H8
C4H8 + Br2 → C4H8Br2
Example 3: Addition of chlorine gas, Cl2 to Butene, C4H8
C4H8 + Cl2 → C4H8Cl2

10.5d Addition Reaction with Hydrogen Gas
 Alkenes can undergo addition reactions with Hydrogen Gas
o This process is known as hydrogenation
 Hydrogenation is used to change vegetable oil into margarine
 Hydrogenation is used to change Alkenes to Alkanes
 Reaction Conditions: 200 oC with Nickel Catalyst
Example 1: Addition of Hydrogen Gas, H2 to Ethene, C2H4
Temperature: 200 oC
Pressure: 1 atm
Catalyst Used: Nickel Catalyst
C2H4 + H2 → C2H6
10.5e Addition Reaction with Water Vapour
 Alkenes can undergo addition reactions with Water vapour
o This process is known as Hydration
 Hydrogenation is used to change Alkenes to Alcohol
 Reaction Conditions: 300 oC, 60 atm with Phosphoric Acid Catalyst
Example 1: Addition of Water Vapor, H2O to Ethene, C2H4
Temperature: 300 oC
Pressure: 60 atm
Catalyst Used: Phosphoric Acid (H3PO4)
C2H4 + H2O → C2H5OH
10.5f Addition Polymerisation
 Alkenes can undergo addition reactions with itself to form long chains of polymers
o This process is known as Polymerisation
Example 1: Polymerisation of Ethene to form polyethene
nC2H4 → - [C2H4]n-

 Polyethene can be used to make plastic items like plastic bags, plastic bottles, etc...
10.6 Alcohol
 Alcohols are organic compounds with a -OH functional group
 Alcohols have the general formula of CnH2n+1OH
 Like Alkanes and Alkenes, Alcohol can undergo combustion and substitution reactions.
 Alcohol can be prepared through the hydration of alkenes (see 10.5e), or through the fermentation
(see 10.6a)
 Alcohols can be used as an organic solvent, or as fuel.
Methanol Ethanol Propanol Butanol
CH3OH CH3CH2OH
CH3CH2CH2OH
Or
CH3CHCH3
OH
4 Isomers
Check 10.3 Example 2
10.6a Fermentation
 Alcohol can also be prepared through the fermentation
o Enzymes are added to break down the glucose in sugar or starch to produce ethanol and carbon
dioxide.
o The fermented mixture is then filtered and the alcohol is extracted through fractional distillation.
o Note: Fermentation must be carried out in an oxygen free environment to prevent the alcohol from
turning into carboxylic acids.
Temperature: 37 oC
Pressure: 1 atm
Catalyst Used: Enzymes in Yeast
C6H12O6 → 2C2H5OH + 2CO2
10.6b Oxidation Reaction
 Alcohols oxidize to form carboxylic acids and water
o For the oxidation, Acidified Potassium Dichromate (VI), Acidified Potassium Manganate (VII), or
even atmospheric oxygen can be used.
Example 1: Oxidation of Ethanol using Acidified Potassium Dichromate (VI)
Conditions: r.t.p
Observations: Orange Acidified Potassium Dichromate (VI) changes to green
C2H5OH + Cr2O7
2- + 10H+ → CH3COOH + 2Cr3+ + 6H2O
Example 2: Oxidation of Propanol using Acidified Potassium Manganate (VII)
Conditions: r.t.p
Observations: Purple Acidified Potassium Manganate (VII) decolourises
C3H7OH + MnO4
- → C2H5COOH + MnO2 + H2O
Example 3: Oxidation of Butanol using Atmospheric Oxygen
C4H9OH + O2 → C3H7COOH + H2O
10.6c Condensation Reaction: Esterification

 See 10.7b

10.7 Carboxylic Acid
 Carboxylic Acids are organic compounds with a -COOH functional group
 Carboxylic Acids have the general formula of CnH2nO2 or Cn-1H2n-1COOH
 Carboxylic Acids are weak acids
o Weak acids only partially dissociate in water to form H+ ions
o As a (weak) acid, carboxylic acids also undergo all acid reactions
Methanoic Acid Ethanoic Acid Propanoic Acid Butanoic Acid
HCOOH CH3COOH CH3CH2COOH
CH3CH2CH2COOH
Or
CH3CHCH3
COOH
10.7a Acid Reactions
 Carboxylic Acids can undergo Neutralization
o Carboxylic Acids reacts with bases to form salt and water only
 Carboxylic Acids can undergo reactions with metals
o Carboxylic Acids reacts with metals to form salt and hydrogen gas only
 Carboxylic Acids can undergo reactions with metal carbonates
o Carboxylic Acids reacts with metal carbonates to form salt, water and carbon dioxide
10.7b Condensation Reaction: Esterification
 Esterification is a condensation reaction
 Alcohols and carboxylic acids react to form Esters
o Esters have the functional group -OCO-
 Ester are named after the alcohol and the carboxylic acid used to prepare the ester
o The first term follows the alcohol used
 Methanol used = methyl
 Ethanol used = ethyl
o The second term follows the carboxylic acid used
 Methanoic Acid used = methonate
 Ethanoic Acid used = ethanoate
o Hence: E.g. Ethanol and butanoic acid produces Ethyl Butanoate
 Esters are sweet-smelling compounds
 Reaction Conditions: Sulphuric Acid Catalyst
Example 1: Preparation of Methyl Ethanoate
Reagents: Methanol and Ethanoic Acid
Conditions: Sulphuric Acid as Catalyst
Observations: A sweet smell is produced
CH3OH + CH3COOH → CH3OCOCH3 + H2O
Example 2: Preparation of Methyl Propanoate
Reagents: Methanol and Propanoic Acid
Conditions: Sulphuric Acid as Catalyst
Observations: A sweet smell is produced
CH3OH + C2H5COOH → CH3OCOCH2CH3 + H2O

10.8 Macromolecules
 Macromolecules are giant molecules that are formed when thousands of smaller units of identical
molecules or atoms are joined together
o This smaller units are called monomers
o The process of joining these monomers into a macromolecule is called polymerisation
 There are 2 types of macromolecules
o Natural Macromolecules
o Synthetic (man-made) Polymers
 Synthetic polymers can be made through addition polymerisation (addition polymers) or
condensation polymerisation (Condensation Polymers)
10.8a Addition Polymers
 Addition polymers are made from unsaturated monomers through an addition reactions
 Alkenes undergo addition polymerisation to form polyalkenes
10.8b Condensation Polymers
 Condensation polymers are made from condensation reactions
o A small molecule like water is produced through condensation reaction
 Monomers used in condensation reactions have 2 functional group on both ends
o A Diol has 2 -OH functional group on either sides
o A Dicarboxylic Acid has 2 -COOH functional group on either sides
o A Diammine has 2 -NH2 functional group on either sides
Diol Dicarboxylic Acid Diammine
HO OH HOOC COOH H2N NH2
= -[CH2]n-
 2 condensation reactions that can be used to produce condensation polymers are
o Esterification (Between diols and dicarboxylic acids)
 The -OCO- functional group between the alcohol and the acid is called the ester linkage
 Terrylene is a polymer with ester linkages, made from diols and dicarboxylic acids.
o Amide Condensation (Between diamine and dicarboxylic acids)
 The -CONH- functional group between the acid and the amine is called the amide linkage
 Nylon is a polymer with amide linkages, made from diamines and dicarboxylic acids.

Chapter 11
The Environment
11.1 Air
 Air comprises of
o ≈ 78% Nitrogen Gas (N2)
o ≈ 21% Oxygen Gas (O2)
o ≈ 1% Argon Gas (Ar)
o Very small amounts of Carbon Dioxide and other rare gases.
11.2 Carbon Cycle & the Greenhouse Effect
 The carbon cycle shows how carbon is circulated around the world
o All living creatures, plants, animals and humans, release carbon dioxide as part of respiration
o All living creatures, plants, animals and humans, also release carbon dioxide through decay and
decomposition
o Plants consume carbon dioxide during photosynthesis
o Animals and Humans consume carbon in the form of food (be it in the form of both plants or other
animals)
o Animals and Humans release carbon in the form of methane
 These two gases, Carbon Dioxide and Methane, are major contributors to the green house effect.
 Greenhouse gases, such as methane and carbon dioxide, are responsible for the green house effect,
which traps heat in our earth's atmosphere.
 The green house effect is essential to sustaining life as the earth would otherwise be too cold to
survive.
 However, too much greenhouse gas leads to global warming, which results in
o Melting of polar caps
o Rising sea levels causing floods in low lying land
o Changing weather patterns such as increase in rainfall in some areas, and possibly causing
floods
o Changing weather patterns such as decrease in rainfall in some areas, resulting in an increase in
number of deserts, as well as possible famine due to crop distruptions.
 Some causes of the increase of Greenhouse gases are
o Increase in use of fossil fuel
o Deforestation
o Decay of vegetation due to deforestation
o increased farming of rice fields
11.3 Carbon Monoxide
 Besides carbon dioxide, the burning of fossil fuels also produces Carbon Monoxide, especially when
there is insufficient oxygen.
o The major source of carbon monoxide is from the burning of petrol in vehicles
 Carbon Monoxide is harmful as it is a poisonous gas
o It binds with the haemoglobin in our blood and prevents it from carrying the oxygen that the body
needs.
 The release of Carbon Monoxide can be reduced by the use of catalytic converters in vehicles.
o The catalytic converter converts the carbon monoxide to carbon dioxide.
11.4 ChloroFluoroCarbons (CFCs)
 The earth is protected by a layer of ozone which absorbs dangerous Ultra-Violet rays from the sun.
o The UV rays would otherwise cause severe damage to vegetations, as well as higher risk of skin
cancer.
 The ozone layer is constantly being destroyed by CFCs used in aerosols, refrigerators and cleaning
solvents

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