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CHAPTER 1: Atomic Structure
An Atom is the smallest part of an element that can take part in a chemical
change. An atom is made up of protons, electrons and neutrons (except
Hydrogen).
A proton is a positively charged particle in the nucleus of an atom.
An electron is a negatively charged particle found in the orbitals outside the
nucleus of an atom.
A neutron is an uncharged particle in the nucleus of an atom.
An element is a substance containing one type of atom.
The mass of an atom is found in the nucleus as the protons and neutrons are
found there and have mass. The electrons form an electron cloud around the atom
that have a negligible (very small) mass.
Atomic Particle
Symbol
Relative Mass
Charge
1
electron
e
-1
1836
neutron
p
1
0
proton
n
1
1
NB:
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The charges of the atomic particles can be proved and confirmed by showing
their behaviour in magnetic fields (firing a beam of a charged atomic particle
past a positively and negatively charged pair of plates).
The proton beam bends toward the negative charged plate, the electron beam
bends toward the positively charged plate and the neutron beam continues in a
straight line.
The bends are caused by an repulsion of similar charges (positive and positive or
negative and negative) and attraction of opposite charges (positive and
negative). The neutron beam stays in a straight line because it has no charge.
The Atomic/Proton Number is the number of protons in the nucleus of an atom.
The Mass/Nucleon Number is the number of protons and neutrons in an atom.
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���� ������ = No. of Protons + No. of Neutrons
�������� = Mass Number − Proton Number
All atoms of the same element have the same number of protons in their
nucleus but the number of neutrons in the nucleus may differ.
Isotopes are atoms of the same element (same number of protons) with
different mass numbers (different number of electrons). Isotopes are also
represented with an isotopic symbol as seen below;
NB: Neutral atoms have the same number of
protons as they have electrons. But in
ions the atom looses or gains electrons to try and make their outermost shell stable.
When an element has a plus (+) on the right side of its symbol it has lost an electron
and when it has a minus (-) it has gained an electron
Cl = 17 protons = 17 electrons
Mg = 12 protons = 12 electrons
Cl− = 17 protons = 17 + 1
Mg2+ = 12 protons = 12 − 1
= 18 electroms
= 10 electrons
CHAPTER 2: Electrons in Atoms
 The electrons of an atom are arranged in energy levels outside the atoms.
 Energy Levels are the specific distance from the nucleus corresponding to the
energy of electrons. They are also known as quantum principle shells.
 shell 1: up to 2 electrons
 shell 2: up to 8 electrons
 shell 3: up to 18 electrons
 shell 4: up to 32 electrons.
 First Ionization Energy is the energy needed to remove one electron from each
atom in one mole of atoms (turns to ion from second onwards) in the gaseous
state to form one mole of gaseous 1+ ions.
NB: The definition of subsequent ionization energies would be the the level of
ionization to replace the words in red and the next level to replace the words in blue.
 The ionization energies can also be represented as chemical equations as follows;
−
First Ionization Energy: ��(�) → ��+
(�) + �
2+
−
Second Ionization Energy: ��+
(�) → ��(�) + �
Factors influencing the ionization energies of elements:
1. The size of the nuclear charge: ionization energy increases as the proton
number increases.
2. Distance of outer electrons from the nucleus: The further the outer electron
shell is from the nucleus, the lower the ionization energy.
3. Shielding effect of inner electrons: The ionization energy is lower as the
number of full electron shells between the outer electrons and the nucleus
increases.
NB: Successive ionization energies increase as more electrons are removed. Large
jumps in the IE reveal where electrons are being removed from the next principal
energy level. The number of electrons removed before the first large jump shows the
number of electrons in the outer shell.
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Sub-shells are regions of the principle quantum shells where electrons exist in
defined areas associated with particular amounts of energy.
 An Atomic Orbital is a region of space outside the nucleus that can be occupied
by a maximum of two electrons.
 Types of Orbitals
1. S Orbital: contains one orbitals (two electrons) and has a spherical shape
2. P Orbital: contains three orbitals (six electrons) and has an hour glass shape
3. D orbital: contains five orbitals (ten electrons) has a modified shape
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Electron Configuration is a way of representing the arrangement of electrons in
atoms by showing their principle quantum number (energy level), sub-shells and
number of electrons present.
NB: They are arranged in an unusual pattern. From 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 4s,
4p,….. except from Chromium and Copper who their 3d sub shell first.
NB: When filling the electrons each orbital must be filled with a single electron first
before a electrons with opposite spins fills the the empty space in the orbital.
 Patterns in Ionization Energy Across a Period:
1. General increase in ionization energy: The nuclear charge increases, the
distance between the nucleus and the outer electron remains reasonably
constant, The shielding by inner shells remains reasonably constant.
2. Rapid decrease in ionization energy between the last element in one period
and the first element in the next period: The distance between the nucleus
and the outer electron increases, the shielding by inner shells increases.
Exception: There is a slight decrease in ionization energy between Beryllium and
Boron Magnesium and Aluminium. The first electron to be removed for B and Al
is in p orbital while that for Be and Mg is in s orbital 9(further from nucleus).
Exception: there is a slight decrease in ionization energy between Nitrogen and
Oxygen/ Phosphorus and Sulphur. This is due to the spin-pair repulsion. The
electron removed from the nitrogen / phosphorus is from an orbital that
contains an unpaired electron. The electron removed from the oxygen / sulphur
is from the orbital that contains a pair of electrons. The extra repulsion between
the pair of electrons in this orbital results in less energy being needed to remove
an electron.
 Patterns in Ionization Energy Down a Group:
1. The first ionization energy decreases as you go down a group in the Periodic
Table: The distance between the nucleus and the outer electron increases,
the shielding by complete inner shells increases.
 Spin Pair Repulsion is when a pair of electrons in the same orbital repel each
other because they have the charge.
 Free Radical is when species of atoms that have one or more unpaired set of
electrons.
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CHAPTER 3: Atoms, Molecules and Stoichiometry
Unified Atomic Mass Unit is one twelfth of the mass of a carbon atom.
Relative Atomic Mass is the weighted average of the mass of atoms in a given
sample of an element compared to the unified atomic mass unit.
Weighted Average Mass of Atom
�������� ������ ���� (�� ) =
Unified Atomic Mass Unit
������ ���� ������ ���� ������� =
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Isotopic Abundances × Percentage Abundance
100
Relative Isotopic Mass is the is the mass of a particular atom of an isotope
compared to the unified atomic mass unit.
Relative Molecular Mass is the weighted average mass of a molecule in a given
sample of that molecule compared to the unified atomic mass unit.
Relative Formula Mass is the weighted average of one formula unit compared
to the value of the unified mass unit
�������/����� ���s = No. of Ions Present × Atomic /Molar Mass of the Atom Present
A Hydrated Compound is a compound which contains a definite number of
moles of water in their structure (water of crystallization).
 An Anhydrous Compound is a compound containing no definite number of
moles of water in their structure (water of crystallization).
 The Avogadro constant is the number of specified particles (atoms, ions,
molecules or electrons) in a mole of those particles to numerical value a 6.02 x
1023.
��. �� ��������� = No. of Moles × Avogadro Constant
 A Mole is the amount of substance which contains 6.02 x 1023 specified particles
atoms, molecule ions or electrons.
Mass (g)
��. �� ����� (���) =
Molar Mass (g mol−1 )
Atomic Mass of Element
���������� ����������� �� �� ������� =
Formula Mass of Compound
Actual Yeild (given in question)
���������� ����� =
× 100
Theoratical Yeild (mass using stiocheometry)
 The Empirical Formula is the simplest whole number ratio of the elements
present in one molecule or formula unit of the compound.
NB: To calculate the empirical formula of an element you divide the masses given of
each element by the elements atomic mass, then you divide all by the lowest division
done above to obtain the lowest whole number ratio.
 The Molecular Formula is the formula that shows the number and type of each
atom in a molecule.
NB: To calculate the molecular formula divide the molecules relative molecular mass
by its formula mass to get what it was divided by. Multiply the answer by the
number of atoms in each atom to get how many atoms are actually in the molecule.
 When deducing the formula for a molecule you switch around their oxidization
numbers and reduce it to the simplest whole number form (in order to make the
overall charge of the molecule neutral).
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Molecular Formulas to Learn:
Ammonium
Carbonate
Hydrocarbonate
Hydroxide
Nitrate
Phosphate
Sulphate
Zinc
Silver
Oxidization Numbers
NH4 2−
CO3 2−
HCO3 −
OH−
NO3 −
PO4 3−
SO4 2−
Zn2+
Ag+
To balance a chemical reaction take note of the amount of each element are at
both side of the equation. Multiply the molecules in order to try to make similar
elements equal to one another.
NB: When writing a chemical reaction remember to write the state symbols solid (s),
liquid (l), gas (g) and aqueous (aq).
 Concentration is the amount of solute (in moles) dissolved in a stated volume of
solution
No. of Moles (mol. )
������������� (��� ��−� ) =
Volume of Solution (dm3 )
 Molar Gas Volume is the volume occupied by one mole of any gas at room
temperature and pressure (s.t.p.). One mole of gas occupies 24.0 dm³ at r.t.p.
Volume of Substance (dm3 )
��. �� ����� (���. ) =
Molar Gas Volume (dm3 )
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Topic 4: Chemical Bonding
Intramolecular Forces are the strong forces between atoms
Ionic Bonding is the electrostatic attraction between opposite charged ions
(anions and cations). An anion is a negatively charged ion while a cation is a
positively charge ion.
 In ionic bonding one atom (usually a metal) loses one or more electrons
to form a positive ion (cation). The electrons lost by the atom are received
by another atom (usually a non-metal) that gains the electrons to become
a negative ion (anion)
 The purpose of all bonds is to make all the reacting elements stable by
making their outermost electron shell either full or empty.
 It can be represented by the use of a dot and cross diagrams as seen in the
examples below:
Covalent Bonding is the electrostatic attraction between the nuclei of two
atoms and a shared pair of electrons.
 A lone pair of electrons are a pair of electrons in the outermost shell of an
atom that are not involved in the bonding.
 There are 4 types of covalent bonds:
1.
Single Covalent Bond (X-Y): when two lone electrons/ one lone pair (one lone
electron from each atom) are shared by the atoms to make their outermost
electron shell stable.
2.
Double Covalent Bond (X=Y): when four lone electrons/two lone pairs(two lone
electron from each atom) are shared by the atoms to make their outermost
electron shell stable.
3.
Triple Covalent Bond (X≡Y): when six lone electrons/three lone pairs (three lone
electron from each atom) are shared by the atoms to make their outermost
electron shell stable.
4.
Co-ordinate/Dative Covalent Bond : when one atom provides both the
electrons needed for a simple covalent bond is bonded to a second atom with
an unfilled p-orbital (requires 2 electrons). It is represented by an arrow pointing
in the direction of the atom with an unfilled p orbital.
NB: There are exceptions to this rule with some specific elements in periods 15, 16
and 17 where they are either electron deficient (do not have a complete shell) or
have an extended octet (more electrons than needed in their outer shell. This is due
to to them using their unfilled p and d orbitals to store the excess electrons.
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Bond Energy and Bond Length
 Bond energy is the energy required to break one mole of a particular
covalent bond in its gaseous stat.
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 Bond length is the distance between the nuclei of two covalently bonded
atoms.
 A double bond has a smaller bond length than a single bond (although
larger than a single bond) due to a greater force of attraction between the
electrons because of a larger nucleic pull.
 Hence, the greater the bond length the smaller the bond energy between
the bonded atoms.
 Due to the bond of double bonds requiring more energy to break down
this then reduces its reactivity (bonds need to be broken to form new
ones). Because of this the longer the bond length the greater the bonds
reactivity.
Shape of Molecules (VSEPR Diagrams)
Sigma (σ)and Pi (π) Bonds
 Covalent bonds are formed when atomic orbital overlap each other. When
s orbitals overlap each other it forms a sigma (σ) bond and when p orbitals
overlap each other it forms a pi bond (π).
 At this level there are three compounds you need to know
1. Ethane
2. Ethene
Metallic Bonding is the electrostatic attraction between positive ions an
delocalised electrons.
 The electrons in metals loose their outer shell electrons to sea of
delocalised electrons and become positive ions.
 Delocalised electrons are electrons that are not associated to a particular
atom or bond.
 Properties of Metals/Metallic Bonding:
1. High Melting and Boiling Points: because it requires a lot of energy to
weaken the strong attraction between the positive ion and the
delocalised electrons.
2. Conduct Electricity: current flows through metals because their
electrons can move around (delocalised) unlike covalent and ionic
bonds.
3. Conduct Heat: their electrons are not bound to any atom or bond.
 Electronegativity is the power of a particular atom that is covalently bonded to
another atom to attract the bonding pair of electrons towards itself.
NB: Electronegativity increases across the period and down the group. Flourine has
the highest electronegativity.
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Factors Influencing Electronegativity:
1. The higher the nuclear charge (number of protons) the higher the
electronegativity. This is because of the strength of the pull of the electrons
to nucleus is increased.
2. An increase in atomic radius (distance of electron from nucleus) causes a
reduction in electronegativity. Because the electrons are further away from
the nucleus.
3. An increase in number of inner complete electron shells (shielding) causes a
decrease in electronegativity. Each shell reduces the nucleus force of
attraction to the valent electrons.
NB: If the electronegativity difference (difference between the electronegativity of
the atoms) is greater than 2 it is ionic but if it is lower than 1 it is covalent.
 Polar Bonds are bonds where the electron pair in the bond is drawn towards the
atom with the larger electronegativity making one end of the molecular slightly
positive compared to the other.
 Molecules could either be non polar (the electronegativities of the atoms are
the same and hence the electronegativity difference is 0) or polar
(electronegativity difference is not 0).
 The atom of a polar bond that is less less electronegative ha a partial delta
positive (δ+) charge while the atom with higher electronegativity has a
partial delta negative charge. The bond is represented with an arrow type
from the delta positive to the delta negative (δ-).
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NB: An increase in polarity leads to an increase in reactivity.
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Intermolecular Bonds are the weak bonds between molecules (they are much
weaker than intramolecular bonds). There are 2 types;
1. Instantaneous dipole - Induced dipole (id-id) Forces: These are the bonds
between non polar molecules. The electrons in atoms are in constant
movement and when the electrons are gathered in one area then one side
of the atom is more negatively charged than the other side. In id-id forces
the more negative side of one atom forms a temporary dipole with the
less negative side of another atom. The more the electrons in this
temporary bond the stronger the bond and the larger the energy required
to break the bond.
2. Permanent dipole - Permanent dipole (pd-pd) Forces: These are bonds
between polar molecules. Because their bonds have both a positive and
negative side already they form much stronger bonds with the other polar
atom. More energy is needed to break pd-pd bonds than that of id-id
forces.
Hydrogen Bonding: is a type of pd-pd intermolecular force and the strongest
of all the intermolecular bonds. It happens when a hydrogen bond covalently
bonded to a nitrogen, oxygen or fluorine (three most electronegative atoms) is
bonded to one of the three most electronegative atoms with an unfilled lone
pair of electrons. Hydrogen bonds increase the boiling points of molecules due
to large amount of energy required to break its bonds.
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The Peculiar Nature of Water:
1. High Boiling Point (due to the hydrogen bond occurring in the bonds
between water molecules)
2. High Surface Tension and Viscosity (hydrogen bond reduces the ability of
molecules to slide across one another as the atoms and electrons are not
as free to move)
3. Ice is less dense than water
Properties of the Types of Bonds:
1. Ionic Bonds: They are solid at room temperature due to the strong
electrostatic forces between their atom. They have a high boiling and
melting point for the same reason. Ionic bonds are soluble in water
because they are polar and do not conduct electricity because their
electrons are not free to move (involved in bonds)
2. Metallic Bonds: They are solid at room temperature due to the strong
electrostatic forces between their atom. They have a high boiling and
melting point for the same reason. Metallic bonds are not soluble in water
because their bonds are too strong to be disrupted by the polarity of
water. They conduct electricity because their electrons are free to move
(delocalised)
3. Covalent Bonds: They are liquid or gas at room temperature because the
intermolecular forces between their molecules are relatively week. Due to
this they have low boiling and melting points. Polar covalent bonds are
soluble in water but non polar covalent bonds are insoluble. Covalent
bonds do not conduct electricity as their electrons are not free to move
around (involved in bonding)
TOPIC 5: States of Matter
 The States of Matter and The Properties of Their Particles:
1. Gases: They have no fixed shape or volume. Their particles are far apart so
they can be compressed. Their particles are randomly arranged. Their
particles can move freely from place to place.
2. Liquids: They take the shape and volume of their container. Their particles
are close together so they can be slightly compressed. They are arranged
randomly and have a limited degree of movement in all directions.
3. Solids: They have a fixed shape and volume. Their particles are touching each
other so they cannot be compressed. They have a regular arrangement and
do not move position (vibrate on a spot).
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The Kinetic Theory of Gases: This is a theory that states that the particles in
gases are in constant movement. Because of this we can assume:
 There is no intermolecular forces between gases (hence can move freely)
 The temperature of gases is related to the average kinetic energy of its
molecules.
 All collisions between gas particles are elastic/no energy is lost when they
collide (because their kinetic energy does not reduce)
 The distance between gas particles are large (because they have space to
move around)
An Ideal Gas is one whose volume varies in proportion to the temperature and
in inverse to pressure. They fit into the kinetic theory of gases.
A Real Gas is one that does not fit into the kinetic theory of gases and does not
obey the ideal gas law. This happens especially at low temperatures and high
pressure.
The General Gas Equation
�� = ��� or �� =
�
�
�T
Where;
 P is Pressure in Pascal (Pa)
 V is the volume in m3 (1m3 = 1000dm3)
 n is the number of moles which can be calculated by dividing the mass of the
atom by its molar mass. It is measured in mol.
 R is gas constant with a value of 8.31JK-1mol-1
 T is the temperature in Kelvin/K (1K = 273°C)
 Changes in State of Matter:
 When solids are heated kinetic energy of the particles increase till it is high
enough to break the force of attraction between its particles and begin to
move around just like liquids do. This is called melting.
 When liquids are cooled they lose kinetic energy and do not move around
as freely. This causes an increased force of attraction between its
molecules till they are so strong the particles can only vibrate about a
point like solids do. This is called freezing.
 When liquids are heated the kinetic particles of its particles increase till it
is high enough to break the force of attraction between its particles and
separate completely and move freely just like gases (vapour). This is called
vaporization.
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 When gases (vapour)are cooled their kinetic energy is reduced and the
force of attraction between their particles continues to increase till they
are no longer to move freely and have limitations to their movement like
liquids. This is called condensation.
 As seen the changes in state of matter are reversible as the changes are
not permanent and be changed from one to another by heating and
cooling.
NB: An increase in temperature leads to an increase in kinetic energy that can
then be used to break bonds between atoms and molecules.
 In a closed system (gases) cannot escape the system and hence would be
trapped at the top of the container. The gases (vapour) at the top of the
contains become close to one another. After a while the vapour with
lower kinetic energies that cannot overcome the attractive forces of the
other vapours will be converted back to liquid.
Types of Structures Found in Elements and Compounds:
1. Giant Ionic Structures: They form a crystalline lattice (a 3D repeating
arrangements of its atoms). Ionic Lattices are hard as they have strong
electrostatic forces holding them together which also makes them have
high boiling and melting points. They are brittle because when their
layers are displaced like charged atoms may come together which would
cause the system to split up (like charges repel). They are soluble in water
because its molecules are made out of polar bonds and can only conduct
electricity in its molten or aqueous (dissolved) state (when its electrons
can move around).
2. Giant Metallic Structures: Metals also form a lattice. These lattices are
strong and have a high melting and boiling point due to the forces
between the positive ion and the sea of delocalised electrons. They are
not brittle as the electrons in metals act in all directions so the bonds are
reformable (hence they are malleable and ductile). They conduct
electricity as their electrons are free to move but they are not soluble in
water because of how strong their bonds are. Alloys (mixture of two or
more metals or a metal and non-metal) makes metals stronger than their
original metal forms as the different sized atoms make it hard for the
atoms to slide over each other.
3. Simple Molecular Structures: They form lattices that are weak and have
low boiling and melting points due to the weak intermolecular forces
between their molecules. They cannot conduct electricity (their electrons
are not free to move) and are insoluble (covalent bonds are non-polar).
4. Giant Molecular/Covalent Molecules: They are lattices with high boiling
and melting points due to the strong bonds between their molecules and
atoms. There are 3 types;
a. Graphites: In graphite the carbons are arranged in planar layers. The
carbons in each row are joined by covalent bonds with delocalised
electron clouds below and above the row. Each row is joined by
weak intermolecular id-id forces. They are soft as their rows can
easily be displaced due to the weak forces between them and they
conduct electricity because of their delocalised electrons.
b. Diamonds: They are made out of a lattice of carbon bonds in a
tetrahedral arrangement. They are hard as there are only covalent
bonds between their atoms and molecules. They do not conduct
electricity because all their electrons are not free to move (fixed to
their bonds).
c. Silicon( IV) Oxide/Si2O: They are made out of a silicon and oxygen
lattice with similar properties to diamonds (hard and conduct
electricity)
5.
Fullerines:They are made out of nano-particles of carbon. There are 3
types needed to be known:
a. Buckminsterfullerene: They are made out of hexagonal and
pentagonal and hexagonal carbon rings that join up to create a
soccer ball like structure.They have are soft and have low
sublimation points (they turn straight to ga) as there are weak
intermolecular forces between their molecules. It is a poor
conductor of electricity and soluble in some solvents as not all its
electrons are used inn the bonding (most of them are).
b. Nanotubes: They are carbon molecules that bond together o form a
cylindrical like shape. They have high electric conductivity on their
surface as some of their electrons are delocalised. They also have
high tensile strength and high boiling and melting point as their
molecules are covalently bonded.
c. Graphene: It is a single isolated layer of graphite. It is the most
chemically reactive allotrope of carbon. It conducts electricity better
than graphite.
NB: Allotropes are structures made out of the same element. Graphite,
diamonds and all the fullerines are allotropes of carbon.
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TOPIC 6: Enthalpy Changes
An Exothermic Reaction is a reaction where heat energy is released (bond
forming) during the reaction. The value of ΔH is negative.
An Endothermic Reaction is a reaction where heat is absorbed (bond breaking)
during the reaction. The value of ΔH is positive.
Enthalpy Change is the total heat energy transferred during a chemical reaction.
�������� ������(��) = Enthalpy of Product(ΔHP) − Enthalpy of Reactant(ΔHR)
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A Reaction Pathway Diagram is a graph like diagram that represents the
enthalpy of the reactant (left) and enthalpy of the product (right) in a chemical
reaction showing the enthalpy change value and the activation energy.
Reaction Pathway Diagram for Exothermic Reaction:

Reaction Pathway Diagram for Endothermic Reaction:

Activation energy is the minimum energy that colliding particles must possess to
break bonds to start a chemical reaction. It is always positive (as it is bond breaking)
 Standard Conditions: in order to properly compare enthalpies we must use the
same conditions for everything. Thee standard conditions is a pressure of
101kPa and a temperature of 298K. Standard conditions are represented by
⦵ or in enthalpy change by ΔH⦵ .
 Varieties of Enthalpy Change:
A. Standard Enthalpy Change of Reaction(ΔHR⦵ ) is the enthalpy change
when the amounts of reactants shown in the stoichiometric equation
react to give products under standard conditions.
There are 3 sub-divisions;
i. Standard Enthalpy Change of Formation(ΔHF⦵ ) is the enthalpy change when
one mole of a compound is formed from its elements under standard conditions.
It can be endothermic or exothermic.
1
2Fe(s) + 1 O2(g) → Fe2 O3(s)
2
ii. Standard Enthalpy Change of Combustion(ΔH⦵ ) is the enthalpy change when
one mole of a substance is burnt in excess oxygen under standard conditions. It
is always exothermic.
CH4(g) + 2O(g) → CO2(g) + 2H2 O(l)
iii. Standard Enthalpy Change of Neutralization(ΔHNE⦵ ) is the enthalpy change
when one mole of water is formed by the reaction of an acid with an alkali
under standard conditions. It is always exothermic.
HCl(aq) + NaOH(aq) → NaCl(aq) + H2 O(l)
 Measuring Enthalpy Change:
 You can measure the enthalpy of a liquid using a calorimeter. By using a known
amount of volume (mass) of a substance we use a calorimeter to measure the
temperature change and use the below formula to measure for enthalpy
change.
�� = ����
 Where:
 ΔH is the enthalpy change in Joules (J)
 m is the mass of the liquid in g (1cm3 volume = 1g mass)
 C is the specific heat capacity of liquids that is 4.18 J g -1 °C -1
 ΔT is the temperature change (final - initial) in °C
Specific heat capacity is the energy needed to rise the temperature of 1g of
substance by 1°C.
If there is a rise in temperature the reaction is exothermic and ΔH is negative but
if there is a fall in temperature the reaction is endothermic and ΔH is positive.
 Hess’s Law states that the enthalpy change in a chemical reaction is
independent of the route by which the chemical reaction takes place as long as
the initial and final conditions and states of reactants and products are the same
for each route.

Enthalpy Cycle for Formation Reaction
Hence, ΔH1 + ΔHr (indirect route) = ΔH2 (direct route)
 Enthalpy Cycle for Combustion Reactions
Hence, ΔH1 (direct route) = ΔHr + ΔH2 (indirect route)
 Ways of Calculating Bond energies for Enthalpy Cycles:
1. When given the relevant enthalpies of formation
2. When given the average bond energies
Using the average bond energies or relevant enthalpy of formation is not as
precise. This is because different bonds have different on energies dependent on the
environment (climate) that it is done in the solution (liquid) that it is performed in.
 The Exact Bond Energy is the energy needed to break a specific covalent bond in
its gaseous state in specific environments
 The Average Bond Energy is the average bond energy needed to break a specific
covalent bond in its gaseous state in various environments.




CHAPTER 7: Redox Reactions
Oxidization is the loss of electrons from an atom, molecule or ion. There are 3
situation where oxidization occurs:
1. The addition of oxygen by an atom, molecule or ion.
2. The removal of hydrogen by an atom, molecule or ion.
3. The increase in oxidization number by an atom, molecule or ion.
Reduction is the gain of electrons from an atom, molecule or ion. There are 3
situation where oxidization occurs:
1. The removal of oxygen by an atom, molecule or ion.
2. The addition of hydrogen by an atom, molecule or ion.
3. The decrease in oxidization number by an atom, molecule or ion.
A Redox Reaction is a reaction where both oxidization and reduction take place
at the same time.
Half Equations and Balancing of Half Equations:
 A Half Equation show oxidization or reduction in isolation from the main
equation. For Example;
���� ��������: 2Na(s) + 2Cl2(g) → 2NaCl(s)
����������� ���� ��������: Na → Na+ + e−
��������� ���� ��������: Cl2 + 2e− → 2Cl−
 When balancing this equation you have make sure the amount of
electrons in each are the same. Hence here we times the oxidization
equation by 2
2Na + Cl2 → 2Na+ + 2Cl−
 Oxidization Numbers: There are rules concerning the oxidation number of
elements;
 The oxidization number of uncombined elements is zero
 Group 1 elements have an oxidization of +1
 Group 2 elements have an oxidization of +2
 Fluorine is always -1
 Hydrogen is always +1 (except in metal hydrides where it is -1)
 Oxygen is always -2 (unless when paired with fluorine that is more
electronegative where it is +2)
 The oxidization number of an ion is equivalent to its charge
 The sum of the oxidation number of a compound (not an ion) is zero
 In a compound the more electronegative atom is the negative oxidization
number.
For oxidization numbers that we do know that are pared with oxidization
numbers that we know, we can make an algebraic expression that allows the
sum of the known and unknown to equate to zero (in a compound) or the charge
of the ion.
 Oxidizing and Reducing Agents:
 An Oxidizing Agent is an atom, compound or ion that increases the
oxidization number of another atom or ion. In the process it goes through
reduction.
 A Reducing Agent is an atom, compound or ion that decreases the
oxidization number of another atom or ion. In the process it goes through
oxidization.
 Spectator Ions are atoms or ions that their oxidization numbers do not
change throughout the reaction
 Naming Compound With Regards to their Oxidization Numbers:
 The oxidization number is written in brackets in for elements with multiple
oxidization numbers for instance;
The oxidization no. of N in N2O is +1. So this compound is nitrogen(I) oxide.
The oxidization no. of N in NO is +2. So this compound is nitrogen(II) oxide.
The oxidization no. of N in NO2 is +4. So this compound is nitrogen(IV) oxide
 An ion containing oxygen and one other element has the suffix ‘-ate’ (except
for hydroxide, OH-)
 An acid (hydrogen + an oxide) containing oxygen has the suffix ‘-ic’

CHAPTER 8: Equilibria
A reversible reaction is a reaction in which the products can be changed back to
the reactants by reversing the conditions (heating and cooling). They are made
out of forward and backward reactions .
������� ��������: CuSO4 ∙ 5H2 O → CuSO4 + 5H2 O
��������� ��������: CuSO4 + 5H2 O → CuSO4 ∙ 5H2O
NB: Reversible reactions
are represented by two half equation arrows ⇌
 An Equilibrium Reaction is a reaction that does not go to completion and in
which the reactants and products are present in fixed concentration ratios.
 Characteristics of Equilibrium reactions:
1. It is dynamic (this means that the reactants are being converted to products
at the same rate the products are being converted back to the reactant).
2. The concentration of reactants and products remain constant at equilibrium.
3. It requires a closed system (a system in which matter is not gained or lost).
 Changing Equilibrium Position:
La Chatelier’s Principle states that If one or more factors that affect a dynamic
equilibrium is changed, the position (refers to the relative amount of product or
reactants present) of equilibrium moves to minimize this change.
A. Changes in Concentration:
 When the concentration of reactants increases the equilibrium shifts
towards the right to reduce the effect. The shift towards the right
makes more products to be formed until equilibrium is restored.
 When the concentration of product increases the equilibrium shifts
towards the left to reduce the effect. The shift towards the left makes
more reactants to be formed until equilibrium is restored.
NB: When substances that are not initially present are added to the mixture the
equilibrium remains untouched because it will equally affect both sides.
B. Changes in Pressure:
 When pressure is increased the the equilibrium shifts in the direction of
the lower number of moles (lower concentration). This is done so that a
lower number of the increase substance is formed.
 When pressure is decreased the the equilibrium shifts in the direction
of the higher number of moles (higher concentration). This is done so
that a higher number of the lowered substance is formed.
NB: Changes in pressure only affects reactions where gases are the reactants and
products (solids and liquids are not compressible)
NB: A lowered pressure means a lowered concentration and an increased pressure
means a lowered concentration.
C. Changes in Temperature:
 In an endothermic reaction (ΔH is positive) an increase in temperature
causes the equilibrium to shift towards the product and a decrease in
temperature causes the equilibrium to shift towards the reactant.
 In an exothermic reaction (ΔH is negative) an increase in temperature
causes the equilibrium to shift towards the reactant and a decrease in
temperature causes the equilibrium to shift towards the product.
catalyst is a substance that increases the rate of reaction. Catalysts reduce the
time for the reaction to occur by reducing the temperature change at which they
occur at. They do not affect the position of equilibrium.
 Equilibrium Expressions
An equilibrium expression is a simple relationship that links the equilibrium constant
in terms of concentration (KC) to the equilibrium concentrations and the equilibrium
constant in terms of pressure (KP) the equilibrium partial pressure.
A. In the equation; �� + �� (���������) ⇌ �� + �� (�������)
NB: A
NB: Solids
are ignored in the KC because their concentrations remain constant.
NB: The units for concentration is mol dm-3. Concentration is represented by square
brackets ([]). So to find out the units for the KC can be calculating by dividing the
concentration units of the product by the reactant.
NB: The initial concentration of the reaction is always equal to the concentration at
equilibrium.
 Factors Involving Change in KC
1. Concentration and Pressure: The value of KC does not change when the
concentration and the pressure of the reaction is altered. This is
because the ratio of the concentration of the product and reactant will
remain the same after the equilibrium is shifted to restore the
equilibrium.
2. Temperature: When the temperature changes the value of KC also
changes in the same order as discussed earlier.
B. Partial Pressure is the pressure exerted by a particular gas in a mixture.
PA
������� �������� =
PA + PB + PC
����� �������� (������ ) = PA + PB + PC + . . . .
In the equation; � + �� (���������) ⇌ �� (�������)
���
�� =
�A × ��
B
Where p is the partial pressure of each substances
NB: The unit of pressure is Pascals (Pa). The unit of KP is the units of each individual
pressures of the product of the individual pressures of the reactant.
 The Mole Fraction the number of moles of a particular gas in a mixture divided
by the total number of moles of all the gases in the mixture.
No. of moles in a particular gas
���� �������� =
Total No. of moles of all the gases in the mixture
������� �������� = Mole Fraction × Total Pressure


A Neutralization Reaction is a reaction of an acid with a base (alkali) to form a
salt and water.
An acid is a proton (H+) ion donor. They have the ability to neutralize a base.
Acids contain hydrogen atoms (they release hydrogen ions when dissolved in
water).
A base (alkali) is a proton (H+ ion) acceptor. The have the ability to neutralize an
acid. Bases (alkali) contain oxides (O) or hydroxides ions (OH-). They accept the
hydrogen ions released by the acid.
NB: An alkali is a base that is soluble in water

NB: Water
is an amphoteric. This means it can act as both as an acid and base.
 Types of Acids and Bases:
Degree of Dissociation is the extent to which a molecule of an acid ionise in a
solvent to produce H+ or a base to ionise in a solvent and produce OH- ions in a
solvent
A. Strong Acids: A strong acid is an acid which dissociates completely in a
solution (produces many H+ ions in solvent). They have a very low pH as
there is a high concentration of hydroxium ions in the solution. They have a
pH between 0 and 3.
B. Weak Acid: A weak acid is an acid which dissociates partially in a solution
(produces few H+ ions in solvent). They have a lowered pH than strong acids
(still below 7) as the concentration of hydroxium ions is lower. They have a
pH between between 4 and 6.
C. Strong Bases: A strong base is a base which dissociates completely in a
solution (produces many OH- in solvent). They have a very low pH as there is
a high concentration of hydroxide ions in the solution. They have a pH
between 11 and 14.
D. Weak Bases: A weak base is a base which dissociates partially in a solution
(produces few OH- in solvent). They have an increased pH that that of strong
bases (still above 7) as the concentration of hydroxide ions is lower. They
have a pH between 8 and 10.
NB: The lower the pH value the higher the number hydroxium (H3O+) in the solution
and the lower lower the number of hydroxium (H3O+) in the solution. A low pH
(below 7) is acidic and a high pH is alkali (above 7). A neutral substance (like salt) is 7.
Distinguishing Between a Strong Acid and a Weak Base:
1. pH values: a strong acid has a lower pH value than a weak acid. This is
because the concentration of the H+ ions in the solution of a strong acid
would be higher. pH indicators measure the amount of hydroxium ions in a
solution (the more H3O+ ions the lower the pH).
2. Electrical Conductivity: a strong acid has a greater electrical conductivity
than those of weak acids of the same concentration. This is because the
concentration of hydrogen ions (and other ions) is greater in strong acids.
3. Reactivity with Metals: When a piece of magnesium ribbon is added to a
strong acid, a steady stream of hydrogen bubbles is observed. When we
add a piece of magnesium ribbon to a weak acid of the same concentration,
only a few bubbles are observed. This is because the concentration of
hydrogen ions is greater in strong acids.
 Indicators and Acid Base Titrations:
An Acid-Base Indicator is a compound that has two different ranges of colours
depending on the pH of the solution in which it is placed. It changes colour over a
narrow range of pH values.
Name of dye
Colour at pH Range
Colour at
lower pH
higher pH
Methyl Orange
Red
3.2 - 4.4
Yellow
Methyl Red
Red
4.2 - 6.3
Yellow
Bromothymol
Yellow
6.0 - 7.6
Blue
Phenolphthalein
Colourless
8.2 - 10
Pink/Violet
Universal pH indicator
Red/Orange
0-14
Blue/Purple
The Universal pH Indicator Scale

CHAPTER 9: Rate of Reaction
 Rate of Reaction is the change in the amount or concentration of a particular
reactant or product per unit time. It is measured standardly in mol dm-3 s-1.
Change in amount of reactant or product
���� �� �������� =
time
NB: The rate of reaction can be calculated by taking the gradient of the tangent of the
curve of concentration over time. Calculating the rate of reaction over shorter
intervals is more accurate as the rate decreases over time.
 The Collision Theory states that in order for particles to react when they collide
they must have sufficient energy and collide in their correct orientation.

Types of Collisions:
1. Ineffective Collision: It is where the particles collide without sufficient
kinetic energy and not in their correct orientation. The reaction does not
occur.
2. Effective Collision: It is where the particles collide with enough kinetic
energy to react in the proper orientation and a chemical reaction occurs.
According to the collision theory a reaction will speed up if the frequency of
collisions increase and the amount of particles with energy greater than the
activation energy increases.
 Activation Energy is defined as the minimum energy required for colliding
particles to react.
 Effects Affecting the Rate of Reaction:
 Concentration: The greater the concentration of the reactants the greater
the rate of reaction. This is because the random motion of more particles in
the solution result in a more frequent collisions between reacting particles.
NB:
 Temperature: An increase in the temperature leads to the average kinetic
energy of the particles increases. Particles will move around more quickly at
a higher temperature and hence have more frequent collisions. An increase
in temperature leads to an increased rate of reaction because the particles
move around quicker (more frequent collisions) and the amount of
successful collisions increases (number of particles with enough energy
increases).

A Catalyst is a substance that increases the rate of a chemical reaction but is
unchanged at the end of the reaction.
Types of Catalysts
1. Homogeneous Catalysts: When the reactant and the catalyst are in the
same phase (state of matter).
2. Heterogeneous Catalysts: When the reactant and the catalyst are in
different phases (most heterogeneous catalysts are solids that catalyse
gaseous/liquid reactants).
NB: Heterogeneous catalysts are usually faster than homogeneous reactions.


Catalysis is the increase in rate of a chemical reaction brought about by the
addition of particular substances that are not used up in the reaction.

Boltzmann’s Distribution is a graph showing the number of molecules with a
particular kinetic energy plotted against kinetic energy. They they change when
there are differences in concentration and temperature but not when a catalyst
is added. The shaded area refers to the amount of molecules with energy
greater than the activation energy.