Strand 1: Atomic Structure, Bonding & Related Properties
Sub-strand 1.1: Structure of Atoms
What is matter? It is anything that takes up space
Matter is made of tiny particles called? Atoms
Sketch and label the structure of the atom oxygen; mass #=16, 8p, neutron= 16-8=8
Define mass number? It is the sum of protons and neutrons also it is always bigger than the atomic number.
Define isotopes? Atoms of the same element that have the same atomic number but different mass number.
Define atomic number? It is equal to the number of protons and is equal to the number of electrons.
How to find neutron number? Mass# - proton#
Define electronic configuration? It describes the arrangement of the electrons in the energy levels. (2.8.8.18)
What are atoms? Are building blocks of all matter and have a central nucleus containing both neutron and proton.
What decides a type of atom? The protons
What is an element? A substance which contains only one type of atom
What is valence electron? The electrons in the outermost shell & they have the highest energy
Example; oxygen: Electron configuration; O (2.6)
Number of valence electron; 6
Lewis structure diagram /Electron dot diagram;
Why is an atom electrically neutral? Because the number of protons in an atom is equal to the number of electrons
Why is the elements in the periodic table unstable? It’s because they have incomplete energy shells except Helium, Neon and
Argon are stable.
Use the periodic Table to locate elements by their group and row position.
Describe the trend in electronic configuration across a row? Elements in the same period have the same number of energy shells
but they have different number of valence electrons, the number of valence electrons increase across a period.
Explain how the trend in electronic configuration across a row relates to atomic size? Electrons are added to the same energy
shell, the valence shell becomes populated with electrons at the same time more protons are added in the nuclei. The nuclei
become more positively charged. The electrostatic attraction between the valence electrons and the nuclei becomes stronger
pulling the energy shells inward hence the atom decrease in size.
Describe the electronic configuration down a group? Elements in the same group have the same number of valence electrons
but their valence electrons exist in different energy levels, energy level increases.
Explain how the trend in electronic configuration down a group relates to atomic size? As the number of energy shells increases,
the valence electrons are placed further away from the nucleus. The electrostatic attraction between the valence electrons
and the nucleus is weaker hence the atom increase in size.
Why elements are arranged in groups and periods?

Periods- represents the number of energy shells

Groups- represents the number of valence electrons
Group 1
Group 2
Group
Group
Group
Group
Group
Group
13
14
15
16
17
18
Period 1 1
2
H
Period 2
Period 3
Period 4
He
(1)
2
4
5
6
7
8
9
(2)
10
Li
Be
B
C
N
O
F
Ne
(2.1)
11
(2.2)
12
(2.3)
13
(2.4)
14
(2.5)
15
(2.6)
16
(2.7)
17
(2.8)
18
Na
Mg
Al
Si
P
S
Cl
Ar
(2.8.1)
19
(2.8.2)
20
(2.8.3)
(2.8.4)
(2.8.5)
(2.8.6)
(2.8.7)
(2.8.8)
K
Ca
(2.8.8.1)
(2.8.8.2)
Sub-strand 1.2: Chemical Bonding
What are the 2 types of chemical bonding? Ionic bonding and covalent bonding.
Draw the Lewis structure of a neutral atom? E.g. magnesium and oxygen
Ionic bonding. Define electrovalent or ionic bonding? Ionic bonding is the electrostatic force of attraction between the cation
and anion.
Show that metal atoms form positive ions (cations) when they donate electrons during ionic bonding. E.g. Potassium and
Magnesium
Show that non- metals form negative ions (anions) when they accept electrons during ionic bonding. E.g. Fluorine and Oxygen
Draw diagrams using Lewis structures to show the formation of ionic bonds between metals and non- metals.
E.g. magnesium and Oxygen
List the properties of ionic compounds:
 Ionic solids have very high melting points and boiling points, because a large amount of energy is needed to
disrupt the regular arrangement of ions in the solid
 Ionic solids do not conduct electricity, because the ions (charged particles) are not free to move within the
lattice.
 Molten (liquid state) ionic compounds and solutions of ionic compounds (i.e. dissolves in water) do conduct
electricity because ionic solutions are hard but brittle.
Why atoms form chemical bonds? In order for the atoms to be stable
Atoms can become stable by:

Giving electron to another atom

Taking electrons from another atom

Sharing electrons with another atom
What does chemical formula shows? The ratio of atoms in a compound
Covalent bonding. Define covalent bonding? It is the sharing of electrons and it occurs between non- metals.
List the features of covalent bonding;
o Covalent bonds can from between ‘like atoms’ and also with ‘unlike atoms’.
o A covalent bond can be a single bond, double bond or triple bond.
Show using Lewis structures the formation of covalent bonds between non- metal.
List the properties of covalent compounds;
 Covalent bonding exists in solids, liquid and gasses at room temperature.
 Do not conduct electricity as they are made up of neutral particles.
 The melting and boiling points are lower than those of ionic solids
 Lower solubility in water than ionic solids.
Sub- strand 1.3; Chemical Reaction
Write the chemical formula for ionic compounds.
A) Iron ((𝚰𝚰𝚰) hydroxide
B) Aluminum oxide
Fe
OH
Al
O
3+
-1
Ans= 𝑭𝒆 (𝑶𝑯)𝟑
3+
2Ans= 𝑨𝒍𝟐 𝑶𝟑
Write the chemical formula for covalent compounds.
a) Nitrogen trioxide- 𝑵𝟐 𝑶𝟑
b) Dihydrogen sulphide- 𝑯𝟐 𝑺
Write simple word equations for decomposition reactions.
a) Sodium carbonate
Ans; Sodium carbonate
carbon dioxide + sodium oxide
b) Aluminum hydroxide
Ans; Aluminum hydroxide
water + aluminum oxide
Write simple word equations for combustion reactions.
a) Calcium
Ans; Calcium + oxygen
calcium oxide
b) Carbon tetrahydride
Ans; Carbon tetrahydride + oxygen
Carbon dioxide + oxygen
Write balanced molecular and ionic equations using symbols for the types of reactions mentioned above.
 Decomposition reaction of magnesium hydrogen carbonate balanced chemical equation;
Ans= 𝑴𝒈 (𝑯𝑪𝑶𝟑 )𝟐
𝑴𝒈𝑪𝑶𝟑 + 𝑪𝑶𝟐 + 𝑯𝟐 𝑶
 Combustion reaction of carbon balanced chemical equation;
Ans= 𝒄(𝒔) + 𝑶𝟐 (𝒈)
𝑪𝑶𝟐 (𝒈)
Write the activity series for cations [𝑁𝑎+ , 𝐶𝑎2+ , 𝑀𝑔2+ , 𝑍𝑛2+ , 𝐴𝑙 3+ , 𝐻 + , 𝐶𝑢2+ , 𝐴𝑔+ , 𝐴𝑢+ ] from the most
reactive to least reactive.
Write the activity series for anions [𝐶𝑙 − , 𝑂2− , 𝑆𝑂4 2− , 𝑁𝑂3 − , 𝐶𝑂3 2− ] from most reactive to least reactive.
Use the activity series to predict the reactivity of a particular compound.
Sub- strand 1.4; Inorganic Chemistry
Define inorganic compounds? Study of all compounds except organic compounds (compounds
containing C-H).
Describe the properties of inorganic compounds.
 High melting and boiling points
 Formed colored compounds
 Conduct electricity only in molten and aqueous state but not in solid state
 Slightly- highly soluble in water
 Form precipitates
Use the solubility rules to predict precipitation. The solubility rule is shown below:
Nitrates, 𝑁𝑂3−
All soluble
Aqueous/ spectator ions
Chlorides, 𝐶𝑙 −
All soluble except AgCl, PbCl
Sulphates, 𝑆𝑂42−
All soluble except 𝐵𝑎𝑆𝑂4 , 𝑃𝑏𝑆𝑂4, 𝐶𝑎𝑆𝑂4
Hydroxides, 𝑂𝐻 −
All insoluble except KOH, NaOH
Solid
Carbonates, 𝐶𝑂32− All insoluble except 𝐾2 𝐶𝑂3 , 𝑁𝑎2 𝐶𝑂3
Apply the solubility rules to test for the presence of the following ions: [𝐶𝑙 − , 𝑆𝑂4 2− , 𝐶𝑢2+ , 𝐹𝑒 2+ , 𝐶𝑂3 2− ]
Example: Sodium carbonate iron (||) sulphate
𝑁𝑎+ 𝐶𝑂32−
Ionic equation: only write for precipitates
2+
𝐹𝑒
√
𝐹𝑒 2+ (𝑎𝑞) + 𝐶𝑂32− 𝐹𝑒𝐶𝑂3 (𝑠) Iron (||) carbonate
𝑠𝑜42−
×
Example: Lead nitrate and sodium chloride
𝑃𝑏 2+ 𝑁𝑂3−
𝑃𝑏 2+ + 𝐶𝑙 −
𝑃𝑏𝐶𝑙2
×
𝑁𝑎+
𝐶𝑙 −
√
STRAND 2: QANTITATIVE CHEMISTRY & AQUEOUS CHEMISTRY
Sub-strand 2.1: Quantitative chemistry
Describe the quantity of matter.
Mole concept- The mole measures the amount of a substance containing the same number of particles
as there are atoms in exactly 12 grams of carbon ( 𝟏𝟐𝟔𝑪), which is 𝟔. 𝟎 × 𝟏𝟎𝟐𝟑
Define the molar mass- is the mass of one mole of any chemical compound, (it’s also its atomic mass)
Mass- is the amount of matter in an object
Avogadro’s number- the number of elementary particles (which can be atoms, ions or molecules) in
one mole of a substance.
Calculate the number of moles and particles using Avogadro’s number.
Generally: Number of particles= moles × Avogadro’s number
Number of molecules= moles of molecules × Avogadro’s number
Total number of ions= moles of compound × number of ions in the compound × Avogadro’s number
Number of atoms in a compound= moles × number of atoms in the compound × Avogadro’s number
For example: find the number of;
a) Helium atoms in 0.5 moles of the gas- 0.5 × 1 × 6.02 × 1023 = 3.01 × 1023
b) H atoms in 0.65 moles of hydrogen gas- 0.65 × 2 × 6.02 × 1023 = 7.83 × 1023
#𝒎𝒐𝒍𝒆𝒄𝒖𝒍𝒆
)
#𝑵𝑨
Change the following quantities’ to the amount present in moles of the particles. (formula #n=
4.0×1023
c) 4.0 × 1023 𝑚𝑜𝑙𝑒𝑐𝑢𝑙𝑒𝑠 𝑜𝑓 𝐻2 0 − 6.02×1023 = 0.66 𝑚𝑜𝑙𝑒𝑠
Calculate moles using n=m/M (where m=mass of substance and M = molar mass of substance or atomic
mass of elements.)
Calculate the molar mass of:
d) Sodium Fluoride (NaF)e) M(𝐶6 𝐻12 06 )
𝑀(𝑁𝑎𝐹) = 𝑁𝑎 + 𝐹
=(12 × 6) + (1.01 × 12) + (16 × 6)
= 23 + 19
=72 + 12.12 + 96
= 42g/mol
=180.12g/mol
Calculate the moles of:
2.2
e) 2.2 g of NaCl- 58.5 = 0.03
3.4
f) 3.4 g of CaS- 288.8 = 0.01
For the molecule methane,𝐶𝐻4, find the:
f) Molar mass- 𝑀(𝐶𝐻4 ) = 𝐶𝑎 + 4𝐻
= 12 + 4 × 1.01
= 16.16g/mol
5.0
g) 5.0 g of Fe- 55.8 = 0.09
g) mass of 2 moles of methane- m= n × M
= 2 × 16.16
=32g
g) Number of methane molecules in 8g of the compound - #molecule= n of molecule × 𝑁𝐴
8÷ 16 = 0.5 𝑚𝑜𝑙𝑒𝑠
= 0.5 × 6.02 ×1023
=3.0 × 1023
𝑚 64
h) Amount of the compound in 64g of the compound- n=𝑀 16 = 4𝑚𝑜𝑙𝑒𝑠
i)
Mass of 3 mol of the compound- m = n × M 3 × 16.16= 48 g
Term 2
LESSON 4- STOICHIOMETRY
Stoichiometry- it is the quantitive relationship among reactants and products.
Determine the mole ratio of reactants and products in a balanced chemical reaction. Determine the
amounts (moles) of reactants and products using mole ratio in a balanced chemical reaction.
For example: propane (𝐶3 𝐻8 + 𝑂2 → 𝐻2 𝑂 + 𝐶𝑂2 )
a) Balance the above equation. 𝐶3 𝐻8 + 5𝑂2 → 4𝐻2 𝑂 + 3𝐶𝑂2
b) If 21g of propane is burned;
I.
II.
III.
𝑚
21𝑔
Calculate the amount (moles) of propane burnt 𝑛 = 𝑀 = 44.08𝑔/𝑚 = 0.476 𝑚𝑜𝑙𝑒𝑠
𝑀(𝐶3 𝐻8 ) = 3 × 12.0 + 8 × 1.01 = 44.08
Determine the mole ratio of propane to water and determine the amount (moles) of
water produced using the mole ratio: 𝐶3 𝐻8 : 4𝐻2 𝑂
x =1.91 moles
1: 4
0.48. . ∶ 𝑋
Calculate the mass of water produced. 𝑚 = 𝑛 × 𝑀
1.91 × 18.02 = 34𝑔
𝑀(𝐻2 𝑂) = 2 × 1.01 + 16.0 = 18.02
LESSON 5, 6 &7; PERCENTAGE COMPOSITION, EMPIRICAL FORMULA AND MOLECULAR FORMULAR
Percentage composition- is a measure of the mass of the different elements in a compound expressed
𝑚𝑎𝑠𝑠 𝑜𝑓 𝑠𝑎𝑚𝑝𝑙𝑒
as a percentage of the total mass. 𝑃𝑒𝑟𝑐𝑒𝑛𝑡 𝑐𝑜𝑚𝑝𝑜𝑠𝑖𝑡𝑖𝑜𝑛 = 𝑡𝑜𝑡𝑎𝑙 𝑚𝑎𝑠𝑠 𝑜𝑓 𝑐𝑜𝑚𝑝𝑜𝑢𝑛𝑑 × 100
E.g. 𝑀(𝑀𝑔𝑂) = 24.3 + 16.0 = 40.3
𝑀𝑔 →
24.3
× 100 = 60.3%
40.3
𝑂 → 39.7%
Determine the empirical and molecular formulae of compounds. Example:
Determine the empirical formula and the molecular formula of a drug which has the percentage
composition of: 60.60% C, 7.07% H, and 32.30% O. the molecular mass of the drug is 198g/mol.
𝑚
𝑛=𝑀
60.60
𝐶 = 12.0 = 5.05
C : H : 0
5.05 7 2.02
:
:
2.02 2.02 2.02
𝟐. 𝟓 ∶ 𝟑. 𝟓 ∶ 𝟏
7.07
𝐻 = 1.01 = 7
32.30
𝑂 = 16.0 = 2.02
2(𝐶2.5 𝐻3.5 𝑂1 )
𝑬𝑭 = 𝑪𝟓 𝑯𝟕 𝑶𝟐
𝑴𝒐𝒍𝒆𝒄𝒖𝒍𝒂𝒓 𝒇𝒐𝒓𝒎𝒖𝒍𝒂 = 𝒙(𝑬𝒎𝒑𝒆𝒓𝒊𝒄𝒂𝒍 𝒇𝒐𝒓𝒎𝒖𝒍𝒂)
𝒙=
𝑴(𝒎𝒐𝒍𝒆𝒄𝒖𝒍𝒂𝒓 𝒇𝒐𝒓𝒎𝒖𝒍𝒂)
𝑴(𝒆𝒎𝒑𝒆𝒓𝒊𝒄𝒂𝒍 𝒇𝒐𝒓𝒎𝒖𝒍𝒂)
= 𝟏𝟐. 𝟎 × 𝟓 + 𝟏. 𝟎𝟏 × 𝟕 + 𝟏𝟔. 𝟎 × 𝟐
= 𝟔𝟎 + 𝟔. 𝟎𝟔 + 𝟑𝟐
𝟏𝟗𝟖𝒈/𝒎𝒐𝒍
=
= 𝟐(𝑪𝟓 𝑯𝟕 𝑶𝟐 ) 𝑴𝑭 = 𝑪𝟏𝟎 𝑯𝟏𝟒 𝑶𝟐
𝟗𝟖. 𝟎𝟔𝒈/𝒎𝒐𝒍
Sub-strand 2.2: Aqueous chemistry
 Lesson 1- structure of water
Explain the high melting and boiling point of water in relation to the types of bonds: hydrogen bonding
is the strongest intermolecular force, and it needs more energy to break up the bond.
Describe the solubility of substances in water, including ionic and molecular substances: e.g.
Solubility of NaCl in water- force of attraction between water molecules and sodium chloride has to be
stronger, than the bonds between negative and positive ion to water.
Solubility of sucrose, 𝐶6 𝐻22 𝑂6 in water- sucrose is a polar molecule, it will be soluble in water because
they have like polarities.
Use solution vocabulary and solubility rule to communicate and predict the solubility of specific
substances
Intramolecular force- force between the atoms in a molecule (e.g. covalent or ironic bonding)
Intermolecular force- force of attraction between the molecules (e.g. hydrogen bonding)
 Lesson 2- concentration of solutions
Describe the concentration of a solution in gram/liter (g/l) and moles/liter (mol/l).
The concentration of a solution is called its molarity. Molarity can be expressed as a measure of the
amount of solute dissolved in a given volume of solvent (mols/L) or measure of the mass of dissolved
solute in a given volume (g/L).
Calculate the concentration of solutions;
1. 𝑐𝑜𝑛𝑐𝑒𝑛𝑡𝑟𝑎𝑡𝑖𝑜𝑛 =
𝑚𝑎𝑠𝑠 𝑜𝑓 𝑑𝑖𝑠𝑠𝑜𝑙𝑣𝑒𝑑 𝑠𝑜𝑙𝑢𝑡𝑒 (𝑚𝑜𝑙𝑠)
𝑛
= 𝑣 = 𝑚𝑜𝑙𝑠/𝐿 𝑜𝑟 𝑚𝑜𝑙𝑠𝐿−1
𝑣𝑜𝑙𝑢𝑚𝑒 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛 (𝐿)
2. 𝑐𝑜𝑛𝑐𝑒𝑛𝑡𝑟𝑎𝑡𝑖𝑜𝑛 =
𝑚𝑎𝑠𝑠 𝑜𝑓 𝑑𝑖𝑠𝑠𝑜𝑙𝑣𝑒𝑑 𝑠𝑜𝑙𝑢𝑡𝑒 (𝑔)
𝑚
= 𝑣 = 𝑔/𝐿 𝑜𝑟 𝑔𝐿−1
𝑣𝑜𝑙𝑢𝑚𝑒 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛 (𝐿)
Examples;
solution Mass of NaCl Volume of water (L) Concentration (c=m/v)
Concentration (c=n/v)
1
10g
25ml
0.025L
400g/l
6.84mols/l
2
15g
50ml
0.05L
300g/l
5.13mol/s
NOTE: change the masses into moles using the formula (n=m÷M)
 Lesson 3- PH of solutions
Describe acidic and basic solutions;
Acidic solutions
Basic solutions
 pH less than 7
 pH greater than 7
 taste sour
 taste bitter
 turn blue litmus paper to red
 turn red litmus paper to blue
Compare the levels of acidity or alkalinity of solutions using pH values.
STRAND 3- PHYSICAL CHEMISTRY
Sub- strand 3.1: Energy Changes
Describe the law of conservation of energy; it states that energy cannot be created nor destroyed, but
is always conserved.
Define potential energy & give examples- is stored energy and the energy of position (gravitational) e.g.
seeds, battery & coal
Define kinetic energy & give examples- moving objects doing work e.g. sound, heat &
Electricity.
List examples of kinetic energy- wind & hydropower
Relate the transfer of energy within a chemical system causes it to do work;
Define work- is to move or displace matter.
Sub- strand 3.2: Chemistry of The Atmosphere
Describe the composition of the atmosphere (percentages of different gases);
 78 % Nitrogen gas, 𝑁2
 21 % Oxygen gas, 𝑂2
 0.9 % Argon, Ar
 0.03 % Carbon dioxide gas, 𝐶𝑂2
Name the major greenhouse gases produced by human activities and natural events;
Water vapor (𝐻2 𝑂), carbon dioxide (𝐶𝑂2 ), methane (𝐶𝐻4) & nitrous oxide (NO).
Describe the protective ability of ozone (𝒐𝟑 ) - the ozone layer absorbs the sun’s harmful high energy
ultra- violet (UV) radiation and protects the earth.
Write a word equation for the combustion of carbon in the atmosphere;
Carbon + oxygen
carbon dioxide + water
Give examples of sources of carbon from human activities; burning of fossil fuels, agriculture &
deforestation.
Write a word equation for the reaction of sulphur dioxide from volcanic activity with water in the
atmosphere; sulfur dioxide (g) + water (l)
sulfurous acid (g)
Describe the effects of sulphurous and sulphuric acid in the form of “acid rain” on the environment.
 Effects of acid rain: acidification of bodies of water, damage of vegetation and damage to
building materials (limestone and marble)
Discuss the effects of carbon dioxide dissolving in seawater on coral.
 Water + carbon dioxide
carbonic acid
The ocean absorbs 30 % of the 𝐶𝑂2 released into the atmosphere. When the levels of 𝐶𝑂2 in the
atmosphere increase due to the greenhouse effect, then the amount absorbed by the ocean also
increases. When 𝐶𝑂2 dissolves in seawater, it increases the acidity of the seawater. In other words, the
pH level of the seawater decreases. This is referred to as ocean acidification. The increase in acidity level
stops the buildup of 𝐶𝑎𝐶𝑂3 in the seawater. Coral polyps draw on this 𝐶𝑎𝐶𝑂3 to build their skeleton
which become coral. As a result, the growing of coral is reduced and leads to the destruction of coral
reefs.
Sub-strand 3.3- Rate of Reaction
Lesson 1& 2: collision theory
Describe the collision theory- the collision theory states that the reacting particles must collide into
each other and the rate of a chemical reaction is proportional to the number of collision between
reactant particles.
Define activation energy or energy hill- it is the minimum amount of energy needed to start a reaction
Explain the effect of temperature on the rate of collision of particles; increase in temperature
increases the rate of chemical reactions. This is because when temperature increases, the reacting
particles gain energy and move faster, resulting in greater chances of effective collision between them
to form products.
Explain the effect of pressure or change in volume on the rate of collision of particles; if the reactants
are gases, at higher pressure, there will be more particles per unit volume. Therefore more particles will
collide per second and the rate of reaction will increase.
Explain the effect of concentration on the rate of collision of particles; increase in concentration of the
reactants increases the rate of chemical reactions. This is because the more the reacting particles, the
more the collisions occurring between them, thus higher chances of effective collision.
Explain the effect of surface area on the rate of collision of particles; increase in surface area of the
reacting particles increases the rate of chemical reactions. This is because when the surface area is
increased, more particles will be exposed to collide amongst each other at a given time and create more
effective collision.
Apply collision theory to real life examples; a puddle of water dries up faster on a hot day
Factor- temperature; when the temperature is high, the water particles gain more kinetic energy. The
water particles have a greater chance of collision and therefore the rate of reaction is faster.
Sub- Strand 3.4- Chemical Equilibrium System
Define and give examples of open and closed systemsClosed system- can exchange energy only between the system and surroundings but not matter. E.g.
Open system- can exchange matter and energy between the system and surroundings. E.g. boiling
water in a saucepan without lid.
Define reversible reactions- reactions that can go in either direction
Demonstrate examples of reversible reactions-
Define dynamic equilibrium- it is when the concentration of the reactants and products become
constant, the rate of the forward and backward reaction will become equal.
Write simple reversible equations for evaporation and condensation-𝐻2 𝑂 (𝐿) ↔ 𝐻2 𝑂 (𝑔)
Write simple reversible equations for the dichromate- chromate system𝐶𝑟2 𝑂7 2− (𝑎𝑞) + 𝑂𝐻 − (𝑎𝑞) ↔ 2𝐶𝑟𝑂4 2− (𝑎𝑞) + 2𝐻 + (𝑎𝑞)
Dichromate (orange)
chromate (yellow)
Write simple reversible equations for the anhydrous copper sulphate- hydrous sulphate system𝐶𝑢𝑆𝑂4. 7𝐻2 𝑂 + Heat ↔ 𝐶𝑢𝑆𝑂4 + 7𝐻2 𝑂
Relate the reversibility of reactions to closed systems; reversible reactions that happen in closed
system eventually reach equilibrium. This is because none of the reactants or products can escape. The
reactants react to give the products. When the products build up, they react and give back the
reactants. As the reaction continues, the rate of the forward reaction decreases while the rate of the
backward increases, then they will become constant and will become equal.
STRAND 4: OXIDATION AND REDUCTION
Sub-strand 4.1; oxidation and reduction, lesson 1& 2- Definition of Oxidation-Reduction Reaction
Define oxidation in terms of hydrogen, oxygen, electronsDefine reduction in terms of hydrogen, oxygen, electronsOxidation
Reduction
Oxygen
 Gain of oxygen
 Loss of oxygen
2𝑀𝑔 + 𝑂2 → 2𝑀𝑔𝑂
𝐶𝑢𝑂 + 𝐻2 → 𝐶𝑢 + 𝐻2 𝑂
Mg is oxidized to MgO by gaining oxygen
Hydrogen

Loss of hydrogen
𝐻2 𝑆 + 𝐶𝑙2 → 2𝐻𝐶𝑙 + 𝑆
𝐻2 𝑂 is oxidized to S by losing H atom
Electrons (OIL RIG)

Loss of electrons
𝑀𝑔 → 𝑀𝑔2+ + 2è
Mg is oxidized to 𝑀𝑔2+ by loss of electrons
Oxidation number

Increase in oxidation number
CuO is reduced to Cu by losing oxygen

Gain of hydrogen
𝐻2 𝑆 + 𝐶𝑙2 → 2𝐻𝐶𝑙 + 𝑆
𝐶𝑙2 is reduced to HCl by gaining electrons

Gain of electrons
𝑂2 + 4è → 2𝑂2−
𝑂2 is reduced to 𝑂2− by gaining electrons

Decrease in oxidation number
Determine the oxidation state of each atom in a given element, molecule or ion;Type equation here.
2𝑀𝑔(𝑠) + 𝑂2(𝑔) → 2𝑀𝑔𝑂(𝑠)
𝐶𝑙2 = 0, 𝐶𝑎2+ = +2, 𝐻2 𝑂 = 𝐻 = +1
0
0
+2 -2
O= -2
Explain oxidation and reduction reactions in terms of changes in oxidation states (number);
Mg is oxidized to 𝑀𝑔2+ since it’s ON has increased from 0 to +2 in 2MgO, while oxygen is reduced since
it’s ON has decrease from 0 to -2.
Lesson 3- Oxidants and Reductants
Define the term oxidant or oxidizing agent- oxidants (oxidizing agents) are substances that enhances
the oxidation reaction to take place but itself gets reduced
Determine the oxidant in a chemical reactionDetermine the reductant in a chemical reactionIdentify the common oxidizing agents (oxygen, chlorine, metals with dilute acids, hydrogen peroxide,
permanganate, and dichromate)
Name
Oxidizing Agents
Appearances & reduction equation
2−
Oxygen (𝑂2 )
(colorless) 𝑂2 → 𝑂 (color depends on the oxide form)
Chlorine (𝐶𝑙2 )
(greenish yellow) 𝐶𝑙2 → 2𝐶𝑙 − (colorless)
Iodine (𝐼2 ), bromine (𝐵𝑟2 )
(brown) 𝐼2 → 2𝐼 − 𝐵𝑟2 → 2𝐵𝑟 −
(colorless)
−
−
2−
Permanganate ion (𝑀𝑛𝑂 4 ) (purple) 𝑀𝑛𝑂 4 → 𝑀𝑛 (colorless)
Dichromate ion (𝐶𝑟2 𝑂2− 7)
(orange) 𝐶𝑟2 𝑂2− 7 → 𝐶𝑟 3− (green)
Hydrogen peroxide (𝐻2 𝑂2)
(colorless) 𝐻2 𝑂2 → 𝐻2 𝑂 (colorless)
Dilute acids (𝐻 − )
(colorless) 𝐻 − → 𝐻2 (colorless)
Define the term reductant or reducing agent- Reductants (reducing agents) are substances that
enhances the reduction reaction to take place but itself gets oxidized.
Identify the common reducing agents (metals e.g. zinc, magnesium and iron; carbon, sulfur dioxide,
carbon monoxide)
Reducing Agents
Metal e.g. Zn, Mg & Fe (Colorless) 𝑍𝑛 → 𝑍𝑛2+ , 𝑀𝑔 → 𝑀𝑔2+ (colorless) 𝐹𝑒 → 𝐹𝑒 2+(pale green)
Carbon
(black) 𝐶 → 𝐶𝑂2(colorless)
Sulfur dioxide
(colorless) 𝑆𝑂2 → 𝑆𝑂2− 4 (colorless)
Ferrous ion
(pale green) 𝐹𝑒 2− → 𝐹𝑒 3−(brown)
Carbon monoxide (CO) (colorless) 𝐶𝑂 → 𝐶𝑂2 (colorless)
Lesson 2- Balancing Redox reactions
Balance simple redox equations of the type;