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UPDATED TO 2022-2024 SYLLABUS

CAIE AS LEVEL
CHEMISTRY
SUMMARIZED NOTES ON THE THEORY SYLLABUS
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CAIE AS LEVEL CHEMISTRY

e- lighter than P+ ∴ deflected at greater angle

1. Atomic Structure
1.4. Electronic Configuration
1.1. Subatomic Particle Electrons are arranged in energy levels called shells
Each shell is described by a principle quantum no. (P.Q)
Relative As the P.Q. increases, the energy of the shell increases
Subatomic Particle Relative Mass/ a.m.u
Charge Inside the shell, there are subshells: s, p, d and f
Protons (P) +1 1 Orbital: region in space where there is a maximum
Neutrons (n) 0 1 probability of finding an electron

Electrons (e-) -1 1/1840

1.2. Protons, Neutrons and Electrons

Mass concentrated within the centre, nucleus
An atom is electrically neutral; P+ = e- Each orbital can hold 2e-s in opposite directions
Atomic number or proton number (Z) = no. of protons When e-s are placed in a set of orbital of equal energy,
Atomic mass or nucleon no. (A) = no. of P + N they occupy them singly, and then pairing takes place
e-s placed in the opposite direction: both -ve charge & if
placed in the same direction, they’d repel. In the opposite
direction, they create a spin to reduce repulsion
Completely filled or half-filled (i.e. one e- in each orbital)
are more stable (reduced repulsion)

Examples of this Method are Shown Below:

Isoelectronic Ions: ions having the same no. of e-s

Isotopes are atoms of the same element with the same
proton number but different numbers of neutrons
Isotopes have similar chemical properties since they
have the same number of protons and electrons (so
chemical interactions are similar)
Isotopes have different physical properties since they
have different numbers of neutrons, causing them to
have different masses and, therefore, different
physical interactions. In certain cases (e.g., period 3 elements), an electron
would prefer the 4s orbital over 3d while filling up. For
1.3. Behaviour of a Beam of Subatomic example, 2 electrons in Titanium would be filled into the
4s orbital before 2 electrons are filled in the 3d orbital.
Particles This is because the 4s orbital is a more stable (lower)
energy level than the 3d orbital.
However, while losing electrons, the electrons from the 4s
orbital would be lost first, and then those from the 3d
orbital would be lost. This is because the 4s orbital is
outer than the 3d orbital.

1.5. Subshells

Protons: positively charged ∴ deflected to -ve pole s p d f

Neutrons: no charge ∴ not deflected Orbitals 1 3 5 7
Electrons: negatively charged ∴ deflected to +ve pole Max e-s 2 6 10 14

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Aufbau’s Principle: method of showing how atomic orbitals

are filled in a definite order to give the lowest energy
arrangement possible

Energy difference between 4s & 3d very small ∴ an e-

1.8. Factors Affecting Ionisation Energy
from 4s can be promoted to half-fill or full-fill 3d orbital to
Nuclear Charge
make the atom more stable
+ve charge due to protons in the nucleus
When filling, fill 4s before 3d and when removing, also
Greater nuclear charge means greater ionization
remove first from 4s
energy

1.6. Orbitals Shielding Effect

Inner shells of e-s repel outermost e-s, thus shielding
s orbitals are spherical, with the nucleus at the centre
them from the +ve nucleus. The more e- shells, the
p orbitals are dumbell-shaped greater is the shielding effect
Greater effect lower I.E because lesser attractive
force between nucleus & outer e-s
Atomic Radius
Distance from the centre of the nucleus to the
outermost orbit
As the number of electron shells increases, the atomic
radius increases
As the number of electrons in the outermost shell
increases, the atomic radius decreases as the
electrostatic attraction between the nucleus and outer
electrons increases
Greater radius lower I.E; a distance of outermost e- to
the nucleus is large ∴ less energy needed to remove
e-
Stable Configuration
High I.E needed to remove e-s from completely or half-
filled orbitals

A free radical is a species with one or more unpaired 1.9. General 1st Ionisation Energy
electrons. Trends
Note: ions and free radicals are different.
Down a group (decreases):
New shells added
1.7. Ionisation Energies (I.E) Attraction of nucleus to valence e-s decreases
Shielding effect increases
1st I.E: the energy needed to remove 1 mole of e-s from 1 Across a period (increases):
mole of a gaseous atom to form 1 mole of unipositive ions Shell number remains the same
Each Successive I.E is higher than the previous one Proton number increases
because as e-s are removed, protons > e-s ∴ the Effective nuclear charge increases
attraction between protons and remaining electrons Atomic radius decreases
increases
Successive I.Es have a large jump in their value when e-s 1.10. Trend in 1st I.E across 3rd Period
removed from the lower energy shell
Deduce group no. by checking when 1st big jump occurs

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As the positive charge on the cation increases, the

number of electrons lost increases, so the electrostatic
attraction between the nucleus and outer electrons
increases.

2. Atoms, Molecules and

Stoichiometry
2.1. Relative Mass
Atomic mass (Ar): weighted average mass of an atom
I.E of Al lower than Mg: e- removed in Al is from higher Molecular mass (Mr): mass of a molecule
energy 3p orbital which is further away from nucleus than Formula mass: mass of one formula unit of a compound
3s e- being removed from Mg. Nuclear attraction is less Isotopic mass: mass of a particular isotope of an element
for 3p than 3s ∴ I.E of Al is lower than Mg
Compared with 12C where one atom of 12C has mass of
-
I.E of S lower than P: e being removed in P is in a half exactly 12 units
filled, more stable 3p orbital whereas in S, the pairing of Unified atomic mass unit: u = 1.66 x 10-27kg
electrons in 3p results in increased repulsion ∴ less
energy need to remove an e- 2.2. The Mole
1.11. Ionic Radius Mole: amount of substance that has the same number of
particles (atoms, ions, molecules or electrons) as there
Ionic Radius: describes the size of an ion are atoms in exactly 12g of the carbon-12 isotope.
Avogadro’s constant: number of atoms, ions, molecules or
electrons in a mole = 6.02 × 10 23

2.3. Mass Spectra

Abundance of isotopes can be represented on a mass
spectra diagram
Positive Ion: smaller radius than original neutral atom
because shell no. decreases, the screening effect Peak Height
Relative Abundance = × 100%
Total Height
​

decreases, but the attraction of the nucleus increases.

Negative Ion: larger ionic radius than neutral atom
M ass × Relative Abundance
because e-s added while nuclear charge remains same Ar = ∑
100
​ ​

Groups 1 to 3 5 to 7
Ion Positive Negative

No. of shells
n−1 n
Across the period:

Proton no. and effective nuclear charge increases

Ionic radius decreases 2.4. Empirical and Molecular Formulae
Empirical Formula: gives the simplest ratio of different
Negative ions are always larger than positive ions in the atoms present in a molecule
same period as they have one more shell Molecular Formula: gives actual numbers of each type of
Ionic radius increases down the group since the number atom in a molecule
of electron shells increases. Molecular formula can be calculated using the Mr of a
As the negative charge on anion increases, the ionic compound and its empirical formula
radius increases since the number of electrons gained
increases such that the number of electrons exceeds the
M olecular F ormula = (Empirical Formula )n ​

number of protons.

l l

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Molecular Mass Dipole Moment: slight charges on atoms in a covalent

Where n =
Mass of Empirical Formula bond due to differences in electronegativity
​

The difference between the electronegativity of two

Atomic M ass × N o. of M oles atoms in a compound determines the overall dipole
% C omposition = × 100%
Molar Mass of Compound
​

moment and overall polarity of the compound.

A large difference in electronegativity will make the bond
more polar (more ionic in nature), but a small difference
in electronegativity will make the bond less polar (more
covalent in nature)

3.2. Ionic (Electrovalent) Bonding

An ionic bond is the electrostatic attraction between
oppositely charged ions.
Structure: giant ionic lattice, crystalline solids

2.5. Calculations Involving Mole

Concept
Mass Have high melting and boiling points
M oles =
Molar Mass
​

V olume of a Gas = M oles × 24

The formula applies to gases at r.t.p.
Unit of volume is dm3 and 1000cm3 = 1dm3
Moles
C oncentration = Coordination Number: number of oppositely charged ions
Volume
​

that surround a particular ion in an ionic solid

Concentration unit = mol dm−3 E.g: NaCl, MgCl2
Anhydrous: a compound in which all water molecules are
removed
Hydrated: a compound which has a number of water
3.3. Dot and Cross Diagrams
molecules associated with its crystalline structure
1. For anions, electrons are added to the atom (e.g.
Water of Crystallisation: these water molecules in a
hydrated compound are called water of crystallisation Chlorine forms Chloride Cl- ions.
2. For cations, electrons are removed from the atom
(e.g. Sodium forms Na+ ions.
3. Chemical Bonding 3. Atoms form the same number of bonds as the number
of electrons required to be added/removed for
3.1. Electronegativity and Bonding maximum stability. (e.g.: Chlorine will form 1 bond,
Sodium will form 1 bond, but Calcium Ca2+ will form 2
Electronegativity: the power of an atom to attract bonds
electrons to itself 4. The central atom may expand its octet to form more
The electronegativity depends on: bonds if necessary (e.g, S 8
The radius of atom (atomic size) inversely ∝ 5. In ionic compounds, electrons are shown to be
electronegativity completely transferred to each of the ions in the
Nuclear attraction directly ∝ electronegativity compound.
Electronegativity increases across a period because of
atomic radius ↓ and nuclear attraction ↑, so polarity ↑
Electronegativity decreases down a group because of
atomic radius ↑ and nuclear attraction ↓, so polarity ↓

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6. In covalent compounds, electrons are shown to be

shared between the atoms in the molecule.

3.5. Covalent Bonding

Covalent Bond: the bond formed by the sharing of pairs of
electrons between the nuclei of two atoms.
Bonding Electrons: e-s involved in bond formation
7. Use a legend of different shapes (e.g. solid circle, Non-bonding electrons or lone pair: pair of valence e-s
unfilled circle, triangle, cross etc.) to show which that are not involved in bond formation
electrons come from which atom in the same way as Covalent compounds are made of molecules which are
colours have been used in the above diagrams. held together by weak intermolecular forces
Elements in period 3 can expand their octet, including the
3.4. Metallic Bonding compounds sulfur dioxide, SO2, phosphorus
pentachloride, PCl5, and sulfur hexafluoride, SF6.
Strong electrostatic forces of attraction between metal They have low melting and boiling points.
cations and delocalized mobile electrons
Structure: lattice of +ve ions surrounded by ­mobile e-s 3.6. Coordinate (Dative Covalent)
Strength of metallic bond increases with:
Bonding
Increasing positive charge on the ions in the lattice
Decreasing size of metal ions in the lattice A coordinate bond is a covalent bond where both
Increasing number of mobile e-s per atom electrons in the bond come from the same atom
Conditions:
An atom should have a lone pair of electrons
An atom should be in need of a pair of electrons
Donor: the atom that supplies the pair of electrons
Acceptor: the atom that accepts the pair of electrons
Coordinate bond is represented by an “→ ” drawn from
the atom donating to towards the atom accepting
Formation of Ammonium ion NH +
4: ​

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Note: Elements in Period 3 can expand their octet by making

use of the energetically accessible but lower lying d-subshell
for bonding. This means that some elements of period 3 can
bond with more than 4 electrons at once. (e.g: Sulfur,
Phosphorus etc)

3.8. Shapes of Molecules

The shape and bond angles of molecules depend on:
Formation of AlCl3 dimer ( Al2 Cl 6 ): The number of pairs of electrons around the central
atom
Whether these pairs are lone pairs or bonded pairs
Valence shell electrons are arranged in pairs to minimize
repulsion between themselves
Order of Repulsion Strength (VSEPR Theory):

Above 750oC, exists as vapor & covalent molecule AlCl3

As vapor cools, exists as dimer Al2Cl6 2 pairs of e’s
Bond angle as AlCl3 = 120o 2 bonded, 0 lone pair
Bond angle as Al2Cl6 = 109.5o Linear
180O
E.g. CO2
3.7. Orbital Overlap
For a covalent bond to form, atomic orbitals containing
unpaired valence electrons must overlap each other
S – S (Sigma Σ) 3 pairs of e’s
3 bonded, 0 lone pair
Trigonal planar
120O
E.g. BF3

S – P (Sigma Σ)
4 pairs of e’s
4 bonded, 0 lone pair
Tetrahedral
109.5O
E.g. CH4

P – P (Sigma Σ)

3 bonded, 1 lone pair

Pyramidal
107O
E.g. NH3
P – P (Pi π )

2 bonded, 2 lone pair

Angular
104.5O
E.g. H2O
Sigma bond has greater overlap ∴ σ > π
Pi bond cannot exist without a Sigma bond.

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5 pairs of e’s
5 bonded, 0 lone pair
Trigonal Bipyramid
90O and 120O
E.g. PF5

6 pairs of e’s
6 bonded, 0 lone pair
Octahedral
90O
E.g. SF6
sp 3 Hybridisation

When 1 s orbital combines with 3 p orbital.

They form linear molecules with 109.25 o angles, e.g.
3.9. Hybridisation SiC l4 , and all hydrocarbons containing 4 bonds like
​

C 2 H4 ( C − C ).
​ ​

Definition Has three times the property of p orbital properties than s

orbital.
Mix atomic orbitals like s and p to form new hybrid
orbitals like sp 2 , sp 3 , and sp .
The resultant orbital has different energy levels, shapes,
and properties.

sp Hybridisation

When 1 s orbital combines with 1 p orbital.

They form linear molecules with 180 o angles, e.g. C O2 ​

and all hydrocarbons containing 2 bonds like C 2 H2

​ ​

(Carbon-Carbon triple bond).

Has an equal amount of s and p orbital properties.

3.10. Bonds
sp 2 Hybridisation Bond Energy: the energy needed to break one mole of a
given bond in one mole of gaseous molecules
When 1 s orbital combines with 2 p orbital. Bond Length: distance between the centres of two nuclei
They form linear molecules with 120 o angles, e.g. BH3 ​

of two adjacent atoms

and all hydrocarbons containing 3 bonds like C 2 H4 (
​ ​

Double bonds are shorter than single bonds because

C = C ). double bonds have a greater negative charge density
Has twice the property of p orbital properties than s between the two atomic nuclei hence greater attraction
orbital. The bond length depends on the radii of the two bonded
atoms; the larger the radius, longer the bond length
The strength of the bond depends on the length of the
bond

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3.11. Hydrogen Bonding

The strongest type of intermolecular force in covalent Permanent Dipole-Dipole Forces
bonds
For hydrogen bonding to occur, we need: Weak forces present between polar molecules
A molecule having a H atom bonded to F, O or N Molecules are always attracted to charged rods, whether
Molecule having F, O or N atom with lone pair of e-s +ve or –ve because molecules have +ve and –ve charges

3.14. Summary

Note: Important properties of water (high melting/boiling

points, high surface tension) are due to the strong hydrogen
bonds present between water molecules.

3.12. Polar and Non-Polar

Polar Covalent Bonds

Bonds with slight ionic character

The bond formed with atoms of different electronegativity
Bonding e-s attracted more towards atom with greater
electronegativity ∴ unequal sharing of electrons ∴
molecule develops slight charges = Polar Molecule
Polar molecules have dipoles, electric charges of equal 4. States of Matter
magnitude and opposite sign
The greater the difference in electronegativity of the two
4.1. Basic Assumptions of Kinetic
bonded atoms, the greater is the ionic character
Theory
Non-Polar Covalent Bonds
Ideal Gas: a gas whose volume varies in proportion to
The bond formed between:
temperature and inverse to pressure.
Identical Atoms: the electronegativity of both atoms is
Noble gases such as helium and neon approach ideal
the same so a pair of electrons shared equally
behaviour because of their low intermolecular forces.
Symmetrical Polyatomic Molecules: dipoles of bond
expert equal & opposite effects, hence cancel charge Ideal Gas Laws:
Non-polar molecules have no overall charge
Gas molecules move rapidly and randomly
The distance between gas molecules is greater than the
3.13. Intermolecular Forces
diameter of molecules ∴ volume is negligible
Intermolecular Forces: weak forces present between two No forces of attraction/repulsion between molecules
covalent molecules All collisions between particles are elastic EK conserved ​

Temperature of gas related to average EK of molecules ​

Induced Dipole (Van Der Waals’ Forces) Conditions at which gases behave ideally:
High temperature
Very weak forces present between non-polar molecules Low pressure
Due to constant motion of e-s, at an instant, a non-polar
molecule develops poles due to distortion of electron Limitations of Ideal Gas Laws:
density giving rise to instantaneous dipole, which can
Real gases do not obey kinetic theory in two ways:
induce a dipole in the adjacent molecules
There is not zero attraction between molecules
Van der Waals forces increase with:
We cannot ignore the volume of molecules
increasing the number of contact points between
themselves
molecules; point where molecules come close
together Deviations visible at low temp. and high pressure
increasing number of electrons (+ protons) in
molecule Molecules are close to each other

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The volume of molecules is not negligible relative to the Particles continue to break away from the surface but are
container trapped in space above the liquid.
VDW forces present, pulling molecules to each other As gaseous particles collide, some of them hit the surface
Pressure is lower than expected from ideal gas of the liquid again and become trapped there.
The effective volume is less than expected from the ideal An equilibrium is set up in which the number of particles
gas leaving the surface is balanced by the number rejoining it.
Liquid water molecules ⇌ Vapor water molecules
4.2. General Gas Equations A fixed number of gaseous particles will be in the space
above the liquid in this equilibrium.
P V = nRT Vapour Pressure: pressure exerted by a vapour in
M ass × RT equilibrium with a liquid.
Mr = Vapour pressure increases as:
PV
​ ​

P1 V1 P2 V2
=
​ ​ ​ ​

​ ​

T1 ​ T2 ​

Quantity Unit Conversion

4.4. Crystallisation
Pressure Pascal 1KPa = 1000Pa
Anhydrous
Volume m3 1m3 = 1000dm3 = 1x106cm3
Temperature OK OC Salt containing no water molecules.
+ 273
Hydrous
Standard Conditions: 101KPa and 298 K
Mols of One Gas Salt containing water molecules.
Mole Fraction = Total Mols of Gases ​

Partial Pressure of a Gas = M ole F raction × Water of Crystallisation

T otal P ressure
The number of water molecules present in the Hydrous salt.
4.3. Liquid State For example, C uSO4 ⋅ 5(H2 O) has 5 water of crystalisation.
​ ​

Particles touching but may have gaps 4.5. Recycling

Have EK slide past each other in random motion
​

Enthalpy of Fusion: heat energy required to change 1 Finite resource: resource which doesn't get replaced at
mole of solid into a liquid at its melting point the same rate that it is used up.
Heating a Solid (melting): Examples of finite resources: copper, aluminium, glass
Energy transferred makes solid particles vibrate Advantage of Recycling: ○ Saves energy ○ Reduces
faster environmental issues ○ Conserves ore supplies ○ Less
Forces of attraction weaken & solid changes to liquid wastage ○ Cheaper than extracting
Enthalpy of Vaporisation: heat energy required to change
1 mole of liquid into a gas at its boiling point
Heating a Liquid (vaporisation):
4.6. States of Matter and Properties
Energy transferred makes liquid particles move faster
Forces of attraction weaken
Solubility
The highest energy particles escape first
For solids, generally, solubility increases with increasing
Liquid starts to evaporate – temp. Below b.p.
temperature as the increase in temperature facilitates
Forces weaken further – particles move faster &
the overcoming of intermolecular bonds, making it easier
Spread
for the solid to dissolve.
Liquid boils – temp. At b.p.
For gases, generally, solubility decreases with increasing
The evaporation of a liquid in a closed container
temperature as the pressure of the gas increases
(pressure only affects the solubility of gases)

Conductivity

Solids (metals) are generally the best conductors of

electricity, while gases are the worst.
Constant evaporation from the surface Solids (metals) are generally better conductors of heat
than liquids, while liquids are better thermal conductors

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CAIE AS LEVEL CHEMISTRY

than gases. This is because of the proximity of molecules Layers kept together by weak Van der Waal’s forces
in solids, allowing heat to be transferred rapidly through High m.p./b.p. - strong covalent bonds throughout
vibrations of neighbouring molecules. Soft – forces between layers are weak
Conducts electricity - has delocalized electrons
4.7. Solid State
Ionic lattice

Silicon(IV) Oxide:
Each Si is bonded to 4 oxygen atoms, but each oxygen
is bonded to 2 Si atoms
Sand is largely SiO2
Similar properties to diamond

Metallic lattice
Hydrogen-Bonded Lattice:

In ice form, water molecules slow down and come closer

together
Due to polarity, molecules form hydrogen bonds between
lone pairs of oxygen & δ + charge of hydrogens
Each water molecule has 2 H-bonds
Simple molecular They arrange themselves into an open crystalline,
hexagonal structure
Due to large spaces, ice is less dense than water

Macromolecular Lattice:

Diamond: Effect of Hydrogen Bonding on Physical Properties:

High m.p./b.p. - each carbon forms four covalent Relatively high m.p./b.p.: many strong H-bonds
bonds High viscosity: hydrogen bonding reduces the ability of
Hard-tetrahedral structure water molecules to slide over each other
High surface tension: hydrogen bonds in water exert a
Doesn’t conduct heat or electricity – no free e-
Used for cutting as is the strongest known substance downward force on the surface of the liquid
and has sharp edges Ice is less dense than water: larger spaces between
molecules in a hexagonal structure

Simple Molecular Lattice:

Iodine:
Dark grey crystalline solid; vaporizes into purple gas
m.p./b.p. are slightly higher than room temp
Slightly soluble in water; dissolves in organic solvents
Graphite:
Diatomic molecules formed due to covalent bonds
Three strong (sp2) covalent bonds between individual atoms
Fourth e- in p orbital ∴ forms a pi bond, forming a Molecules have weak Van der Waals forces of
cloud of delocalised electrons above and below the attraction between them
planes

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Fullerenes:
Buckminsterfullerenes(C 60)
C atoms in pentagonal and hexagonal rings
Spherical
C 60 molecules held together by Van der Waals
forces
Can conduct heat and electricity
Very strong and tough
Insoluble in water
Low m.p./b.p. Endothermic Reactions

Nanotubes
C atoms in hexagonal rings only
Cylindrical
The structure is rod-like due to continuing rings
Conducts heat and electricity
Very strong and tough
Insoluble in water
High m.p./b.p.

5.3. Enthalpy Change Definitions

Standard Conditions (this syllabus assumes that these are
298K and 101 kPa) shown by ⦵

5. Chemical Energetics Standard Enthalpy Change of:

1. Reaction ΔH
5.1. Energy Change in Reactions 1. Enthalpy changes when 1 mole of element or
compound is completely reacted under
Exothermic Reactions Endothermic Reactions standard conditions in their standard states.
Energy given out Energy taken in 2. Combustion ΔHC ​

1. Enthalpy changes when 1 mole of element or

Surrounding warmer Surrounding cooler
compound is completely combusted under
Bond making Bond breaking standard conditions in their standard states.
ΔH positive 3. Formation ΔHF
ΔH negative ​

EReactants > EProducts 1. Enthalpy changes when 1 mole of a compound

EReactants < EProducts is formed from its elements under standard
conditions in their standard state.
Energy Change is measured in 1mol dm−3 4. Neutralization ΔHn ​

1. Enthalpy change when 1 mole of H + and

5.2. Reaction Pathway Diagrams OH − combine to form 1 mole of H2 O under ​

standard conditions in their standard states

Exothermic Reactions
5.4. Bond Energy

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Energy needed to break a specific covalent bond

5.8. Worked Examples - Hess Law
Also how much energy is released when a bond forms

5.5. Calculating Enthalpy Changes

Q = −mcΔT
Q
ΔH =
no. of moles
​

When substance dissolved in water use c & m of water

ΔT is a change in temp.: add –ve or +ve to show rise/fall

5.6. Hess’s Law

The total 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 are
the same.
Reason to Use Hess’s Law:
Standard conditions hard to maintain (e.g. Exo/endo)
Elements don’t always react directly

5.7. Calculation of Enthalpy Change, ΔH

Reaction from Formation

Formation from Combustion

Hydration from Anhydrous Salt

Reaction from Bond Energies

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Oxygen as peroxide; oxidation number = -1

1st group elements & hydrogen; oxidation number = +1
H with highly reactive metal; oxidation number = -1

Following these rules, all other atoms in a covalent bond

must balance out the charge

6.2. Redox Processes

Reaction where both oxidation and reduction occur
This can be shown with changes in oxidation numbers of
elements from the product side to the reactant side.
E.g. C a2 S i + 6H2 O → 2C a(OH)2 + SiO2 + 4H2
​ ​ ​ ​ ​

C a = +4 → +2 ⟹ gain of negative charge ∴ a

reduction
Si = −4 → + 4 ⟹ loss of negative charge ∴
oxidation
Disproportionation Reaction: When the same species
undergo oxidation and reduction in a reaction.

Oxidising Agent: A substance which helps to oxidise

another species by accepting electrons from it and itself
gets reduced in the process
Reducing Agent: A substance which helps to reduce
another species by donating electrons to it and itself gets
oxidised in the process

Source: Alt Academy

6. Electrochemistry
6.1. Calculating Oxidation Numbers
Ionic Molecules: group number = valence electrons
Covalent Molecules:
Rules: 6.3. Balancing Equations
Atoms in a diatomic molecule; oxidation number =
0 Equation: HI + HN O3 → I 2 + N2 O3 + H2 O
​ ​ ​ ​ ​

Oxygen in a compound; oxidation number = -2 Half ionic: 2I − − 2e− → I 2 N +5 + 2e− → N +3

​

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For every 2 iodines, there will be 1 nitrogen

7.5. Manufacture of Sulfuric Acid
Thus first put in correct ratio for iodine and nitrogen then
balance hydrogens and oxygens
Balanced: 4HI + 2HN O3 → 2I 2 + N2 O3 + 3H 2 O ​ ​ ​ ​ ​

7. Equilibria
7.1. Introduction
Reversible reaction: a reaction in which products can be
changed back to reactants by reversing the conditions
Dynamic Equilibrium: the state of a reversible reaction
carried out in a closed container where the rates of
forward and backward reactions are equal and constant

7.2. Le Chatelier’s Principle

When a chemical system in dynamic equilibrium is
disturbed (conditions changed) it tends to respond in such
a way so as to oppose the change and a new equilibrium
SO3 not dissolved directly into water because reaction
is set up
explosive and causes H2SO4 to vaporize
Forward reaction exothermic ∴ temp. not too high so the
rate of backward reaction doesn’t increase & is not too
low so particles have EK for collision and catalyst works
​

Since the reaction is highly exothermic, gases must be

cooled
7.3. Equilibrium Constants No impurities; otherwise, the catalyst will be poisoned
Atmospheric pressure is enough because equilibrium
Equilibrium constant expressed in terms of concentration already favours the product side
[Product]mols
KC = ​

[Reactant]mols
​

Only liquids and gases 7.6. Brønsted–Lowry Theory of Acids

Equilibrium constant expressed in terms of partial
pressure and Bases
p(Product)mols
KP = ​

p(Reactant)mols
​

Brønsted-Lowry Theory:
Only gases
An acid is a proton (H+) donor
Large value of K C /K P ​ ​ ⇒ equi. towards products side A bases is a proton (H+) acceptor
Amphoteric: substances that can act like bases or acids
Smaller value of K C /K P ⇒ equi. towards reactants
​ ​
Strong acid: an acid that dissociates completely in solution
side (e.g., HCl)
K C /K P changes only with changes in temperature
​ ​ Weak acid: an acid that dissociates partially in solution
The amount of reactants that disappear will always (e.g., ethanoic acid)
appear in the products in the same ratio as present in a Strong base: a base that dissociates completely in
balanced equation solution (e.g., NaOH)
Weak base: a base that dissociates partially in solution
7.4. Manufacture of Ammonia (e.g., ammonia)
Strong acids/bases react more vigorously than weak
acids/bases.
Strong acids have lower pH values than weak acids.
Strong bases have higher pH values than weak bases.
When an acid reacts with a base, salt & water are formed.
The pH changes in this neutralisation reaction can be

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graphed as shown in the image below: \n

8.5. Effect of Temperature

Number of collisions and chance of success will increase

8.6. Effect of Catalyst

Catalyst: a substance that increases the rate of reaction
but remains chemically unchanged at the end
It does not alter the chemical composition of substances
and only lowers the activation energy
It provides a new route or mechanism to follow for
8. Reaction Kinetics reactants that requires less energy

8.1. Introduction
Rate of a Reaction: change in concentration of reactants
or products per unit of time
Activation Energy: minimum energy colliding particles
must possess for a successful collision to take place
Catalysis: acceleration of a chemical reaction by a catalyst

8.2. Effect of Concentration Changes Curve unchanged; only activation energy changes
Homogeneous catalysts: reactant and catalyst are in the
Increasing conc. of reactants increases rate of reaction: same physical state
more particles per unit volume, collision rate between Heterogeneous catalysts: reactant and catalyst are in
reacting particles increases, ∴ rate of successful collision different physical states
increases, resulting in increased rate of reaction. Enzymes: a protein molecule that is a biological catalyst.
Most are specific to a substrate & function as lock-key
8.3. Effect of Pressure Changes
Increasing the pressure of the system increases the rate 9. The Periodic Table:
of reaction: more particles per unit volume, collision rate
between reacting particles increases, and rate of
Chemical Periodicity
successful collision increases, resulting in an increased
rate of reaction. 9.1. Introduction
8.4. Maxwell-Boltzmann Theory Group 1 2 3 4
Element Sodium Magnesium Aluminium Silicon
Explains effect of temp. & catalyst on rate of reaction
Character Metal Metalloid
Based on distribution of energy among reacting
Structure Giant metallic lattice Macromolecular
molecules under different conditions
Metallic bond between cations Covalent bonds
Bonding
and delocalized e- between atoms

Diagram

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9.2. Reaction of Elements with Oxygen With Produces

SO2(g) NaOH NaHSO3(aq)
Formulae Reaction Structure Oxid. Nature SO2(g) Excess NaOH Na2SO3(aq) + H2O
Burns yellow SO3(g) NaOH NaHSO4(aq)
Na Na2O(s) +1 Basic
flame SO3(g) Excess NaOH Na2SO4(aq) + H2O
Burns Bright Giant ionic
Mg MgO(s) +2 Basic
White Flame lattice
Burns Bright
9.6. Reactions of Elements with
Al Al2O3(s) +3 Amphoteric
White Flame Chlorine
Giant
Si SiO2(s) Coating +4 W. acidic
covalent Formula Structure Oxid. Nature
P2O3(s) Burns yellow +3 Na NaCl­(s) +1 Neutral
P S. Acidic Giant ionic
P2O5(s) flame Simple +5 Mg MgCl2(s) +2 Neutral
SO2(g) Burns blue molecular +4 Al AlCl3(s) +3 Acidic
S S. acidic
SO3(g) flame +6 Si SiCl4(l) +4 S. Acidic
PCl3(l) Simple molecular +3
P 2 O5 empirical formula of P 4 O10 P S. Acidic
PCl5(l) +5
​ ​ ​ ​

9.3. Reaction of Na & Mg with Water

9.7. Reactions of Chloride with Water
Na & 2Na(s) + 2H2O(l) Very fast, floats, forms
Water 2NaOH(aq) + H2(g) ball & dissolves
Mg & Mg (s) + 2H2O(l)
Very slow
Water Mg(OH) 2(aq) + H2(g)
Mg & Mg (s) + H2O(g) MgO(s) +
Very fast
Steam H2(g) Sodium chloride simply dissolves in water. Water is polar
∴ positive Na+ attracted to OH- while Cl- attracted to H+
9.4. Reaction of Oxides with Water MgCl2 slightly acidic because Mg ion has smaller radius &
higher charge ∴ attraction to water is so strong that H2O
Reaction Oxid. Nature loses a proton and solution becomes slightly acidic
Na2O(s) Na2O(s) + H2O(l) 2NaOH(aq) +1 S. Alkaline
MgO(s) MgO(s) + H2O(l) Mg(OH) 2(aq) +2 W. Alkaline
Al2O3(s) NO REACTION
SiO2(s)
P2O3(s) P2O3(s) + 3H2O(l) 2H3PO3(aq) +3
S. Acidic
P2O5(s) P2O5(s) + 3H2O(l) 2H3PO4(aq) +5
SO2(g) SO2(g) + H2O(l) H2SO3(aq) +4
S. Acidic 9.8. Atomic Radius
SO3(g) SO3(g) + H2O(l) H2SO4(aq) +6

9.5. Acid-Base Reactions

Aluminum oxide is amphoteric ∴ reacts with the acid and
base

Al2O3 + H2SO4 → Al2(SO4) 3 + H2O | Al2O3 + NaOH →NaAlO2

+ H2 O P+ in nucleus increases so nuclear charge increases
There are more e-, but increase in shielding is negligible
Silicon dioxide is acidic: SiO2 + NaOH (hot & conc.)
because each extra e- enters same principal energy level
Na2SiO3
Sulphur dioxide and trioxide are strongly acidic ∴ force of attraction between nucleus & e- increases ...
So atomic radius decreases.
With Produces

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of protons increases
9.9. Ionic Radius
9.13. First Ionisation Energy

Ionic radius decreases across a period however, since

non-metals gain electrons, they have one more shell than
metals therefore they always have a larger radius than
metal ions Generally, increases as no. of protons increases
Decrease Mg → Al: more distant and less effective
9.10. Melting Point nuclear charge on 3p orbital
Decrease P → S: in S, one electron paired ∴ causing
repulsion and easier to lose an electron

10. Group 2
10.1. Introduction

Na Al m.p. increases because delocalized e- per atom

increases making metallic bond stronger
Si has highest m.p. due to giant covalent structure
The larger the molecule size, the stronger the Van Der
Waals forces ∴ S 8 > P4 > Cl2 > Ar

9.11. Electrical Conductivity

m.p./b.p. decreases down group: atoms/ions get larger,

the distance between nuclei & e-s increases ∴ bonds
weaker
m.p./b.p. higher in gp. 2 than 1: 2e-s per atom donated
into delocalized system ∴ metallic bonding stronger
density increases down group: mass of atoms increases
faster than their size (volume) as atomic no. increases
Na < Mg < Al because no. of delocalized electrons which
can carry charge increases
10.2. Reaction of Group 2 Metals with
Silicon is a semi-conductor
Non-metals – covalent ∴ no charge Oxygen

9.12. Electronegativity
All group 2 metals tarnish in air, forming oxide coatings
Burn vigorously in oxygen-forming white solids

10.3. Reactions with Water

Metals: M (s) ​ + H2 O(l) → M (OH)2(aq) + H2(g)
​ ​ ​ ​

Metal Oxides: MO (s) + H2 O (l) → M(OH)2(aq)

​ ​ ​ ​

Solubility of MO (Metal Group 2 Oxide) and M(OH) 2(Metal

Increases across period because the bonded e- are in the
Group 2 Hydroxides) increases down group
same energy level but are attracted more strongly as no.

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Alkalinity of solution increases down the group

11.2. Oxidising Ability
Solubility of MSO4 decreases down group
M(CO3) don’t react with water Halogens have high electron affinity (they gain electrons
easily) hence, they are good oxidising agents
10.4. Reaction with Acids Oxidising ability decreases down the group because
electron affinity decreases as atomic size increases.
M (s) + Acid(aq) → Salt + Hydrogen
​ ​

11.3. Reactions of Halide Ions

MO (s) + Acid(aq) → Salt + Water
​ ​

X2(g) + H2(g) 2HX(g)

M(OH)x(s) + Acid(aq) → Salt + Water
​ ​

Product Reaction Description

MCO 3(s) + Acid(aq) → Salt + Water + Carbon Dioxide
​ ​
HF Reacts explosively in all conditions
HCl Reacts explosively in sunlight
10.5. Thermal Decomposition of Group HBr Reacts slowly on heating
2 Metals HI Forms an equilibrium mixture on heating

MCO 3(s) M O (s) + CO 2(g)

​ ​ ​
The thermal stability of halogen hydrides decreases down
the group because:
2M(NO 3 )2(s) 2MO (s) +4NO 2(g) + O 2(g)
​ ​ ​ ​ ​

The size of halogen atom increases

Nuclear attraction decreases
NO2: thick brown, acidic and soluble gas The H – X bond becomes longer and weaker
Thermal stability increases down the group ∴ Thus, less energy is needed to break the bond
decomposition becomes more difficult. Bond energies decrease down the group

10.6. Uses of Group 2 Metals 11.4. (Sub) Halide ions and aq. Silver
Calcium compounds:
Ions

Ag +(aq) + X-(aq) AgX(s)

Halide With Silver With dilute aq. With conc. aq.
Ion Nitrate ammonia ammonia
Cl‑ White ppt. ppt. dissolves

Br- Cream ppt. X ppt. dissolves

I- Yellow ppt. X X
Calcium oxide (lime): basic oxide used to neutralize acidic
soil and used as a drying agent for drying ammonia
Calcium carbonate (limestone): used as building material dissolves, it forms a complex ion:
(cement, concrete) etc., for extraction of iron, glass AgX(s) + 2NH3(aq) → [Ag(NH3) 2]+(aq) + X-
industry, neutralize soil or chemical waste
The complex ion formed is called Diamine Silver(I) ion

11. Group 17 [H3 N :⟶ Ag ←: N H3 ]+

​ ​

11.5. (Sub) Halide Ions and Aqueous

11.1. Trends in Colour and Volatility
Sulfuric Acid
Fluorine Yellow
Metal Halide + Conc. H2SO4(aq) Hydrogen Halide
Yellow- Gas
Chlorine
Green
m.p. & b.p. Volatility The colour Conc. H2SO4(aq) is an oxidising agent (except for chloride
Orange- and fluorides as it is not strong enough)
Bromine Liquid increases decreases gets darker
Brown This reaction is used for the preparation of hydrogen
down the down the down the
Grey, group ↓ group ↓ group ↓ halides
Iodine Black,
Solid Chlorine NaCl(s) + H2SO4(aq) HCl(g) + NaHSO4(aq)
Violet
Astatine Black Bromine NaBr(s) + H2SO4(aq) HBr(g) + NaHSO4(aq)

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HBr(g) + H2SO4(aq) Br2(g) + SO2(g) + H2O(l)

12.1. Lack of Reactivity of Nitrogen
NaI(s) + H2SO4(aq) HI(g) + NaHSO4(aq)
HI(g) + H2SO4(aq) I2(g) + SO2(g) + H2O(l) Nitrogen molecule has three strong covalent bonds.
Iodine The bond is very strong and requires high energy to split
HI(g) + H2SO4(aq) I2(g) + H2S (g) + H2O(l)
the two nitrogen atoms of a molecule.
6HI(g) + H2SO4(aq) 3I2(g) + S (s) + 4H2O(l) It reacts only under extreme temperature or pressure or
in the presence of a catalyst.
11.6. The Reactions of Chlorine with
Aqueous Sodium Hydroxide 12.2. Ammonium

Disproportionation: a reaction in which the same Lone pair of e-s of nitrogen forms a coordinate bond with
substance is oxidized and reduced simultaneously, the H+ ion
producing two different products
Formation: NH3(g) + H+ NH4+
When chlorine reacts with a solution of cold aqueous
Shape: tetrahedral
sodium hydroxide, the disproportionation goes to lower
Bond angle: 109.5o
oxidation states
Bond length: equal lengths
C l2 + 2N aOH → N aCl + N aClO + H2 O
​ ​

With a hot solution, the oxidation state of chlorine goes up

Displacement of ammonia from its salts:
to +V
warm
Any Ammonium Salt + Any Base Ammonia
3C l2 + 6N aOH → N aClO3 + 5N aCl + 3H2 O
​

​ ​ ​

This happens as the chlorate is formed by 12.3. Uses of Ammonia & its
disproportionation of hypochlorite and hypochlorous acid
Compounds
ClO− + 2HClO → ClO3− + 2HCl ​

Used in the production of nitric acid

Higher temperatures promote the formation of hypochlorous Used in the production of inorganic fertilizers
acid through the hydrolysis of hypochlorite and, therefore, Used in the production of nylon
speed up the reaction Used in the production of explosives

11.7. Important Uses of Halogens and 12.4. Oxides of Nitrogen

Halogen Compounds
N2(g) + O2(g)→ 2NO(g) or N2(g) + ½O2(g)→ NO2(g)

Fluorine: Naturally: during lightning, EA provided for N2 to react

To make chlorofluorocarbon (CFCs)
Man-made: in car engine, high temp. and pressure
As fluoride in toothpaste
Catalytic convertors: exhaust gases passed through
To make polytetrafluoroethylene (PTFE) – non-sticking
catalytic convertors containing a catalyst (platinum/
coating in pots and pans
palladium/nickel) helping to reduce oxides to nitrogen.
Bromine and Iodine: manufacture of photographic films
Catalytic role in oxidation of sulphur dioxide:
Chlorine:
In bleaches
To make PVC and chlorofluorocarbon (CFCs)
As solvents
Use of chlorine in water purification:
The oxidising power of chlorine is used in the
treatment of water to kill bacteria

Cl2(aq) + H2O(l)→ HCl(aq) + HClO(aq)

HClO(aq)→ HCl(aq) + O
12.5. Pollution

This disproportionation reaction produces reactive Acid Rain: SO3 + H2O→ H2SO4
oxygen atoms which kill bacteria 2NO2 + H2O → HNO3 + HNO2 or 4NO2 + 2H2O + O2 →
4HNO3

12. Nitrogen and Sulfur

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Damages trees & plants, kills fish and other river life,
buildings, statues and metal structures

Combustion Pollutants:

Nitrogen oxide (NO): formed by reaction of N2 and O2 in sp3 sp2 sp

the engine, forms acid rain and respiratory problems
All orbitals mix 2s, 2px, 2py mix 2s and 2px mix
Atmospheric oxides of nitrogen (NO & NO2) can react with
unburned hydrocarbons to form peroxyacetyl nitrate 3 sp2 orbitals 2 sp orbitals
4 sp3 orbitals
(PAN) which is a component of photochemical smog 1 pure p orbital 2 pure p orbitals
Carbon monoxide (CO): source: incomplete combustion of Ratio of characteristics s : p
hydrocarbon fuel, toxic effect on haemoglobin
1:3
1:2 1:1
12.6. Food Preservation
SO2 is used by itself or as a sulphite to preserve food

SO2 + H2O → H2SO3(aq)

SO2 & sulphites inhibit growth of bacteria, yeasts, etc. &

13.3. Classes of Compound
are reducing agents, so reduce rate of oxidation of food.
Used to prevent spoilage of dried fruit, dehydrated
Organic Family Suffix Example
vegetables and fruit juices.

13. An Introduction to AS
Level Organic Chemistry -ane Methane

13.1. Introduction
Alkane
Organic Chemistry: study of hydrocarbons and their
derivatives
Carbon can form a variety of compounds because:
Carbon is tetravalent
-ene Ethene
Carbon-carbon bonds can be single, double or triple
Atoms can be arranged in chains, branches and rings
Homologous series: a series of compounds of similar
structures In which: Alkene
contain the same functional group
all share the same general formula
the formula of homologue differs from neighbour by
CH2
similar chemical properties -yne Ethyne
gradual change in physical properties as Mr increases
Functional group: an atom or group of atoms in an organic
molecule that determines the characteristic reactions of a
homologous series. Alkyne
Alkyl group: a reactive group which is alkane minus 1 H

13.2. Hybridisation
Hybridisation: mixing up of different atomic orbitals -ol methanol
resulting in new orbitals of equal energy.
Carbon’s Electron Configuration:
Alcohol

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Organic Family Suffix Example Organic Family Suffix Example

methoxy-
-oxy- -amine Methyl-amine
methane

Ether Amine

Chloro-
halo- -ane nitrile Ethyl nitrile
ethane

Haloalkane Nitrile

-benzene
also have
Nitro-
other
benzene
-al Methanal naming
Arene/Phenyl/Benzene methods

Aldehyde

methyl-
amine
methanamine

-one Propanone
Amide

Ketones 13.4. Types of Formulae

Displayed formula (Hexane)

Methanoic
-oic Acid
Acid

Carboxylic Acid

Structural formula (Hexane)

CH3-CH2CH2CH2CH2-CH3 or CH3(CH2) 4CH3
Skeletal formula (Hexane)
Methyl-
-ate
ethanoate

Molecular formula (Hexane)

Ester C 6H14

13.5. IUPAC Nomenclature

Select the longest chain as the main chain
Other carbon chains as substituent alkyl groups

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Give the lowest number C in the main chain to a Nucleophilic substitution (halogenoalkanes)
substituent Free radical substitution (alkanes)
If different alkyl groups are present on identical positions, Elimination reaction: more than one product formed,
give simpler alkyl smaller number small molecule removed from reactant (alcohols and
Two or more alkyl groups present, order alphabetically halogenoalkanes)
If the same substituent is repeated use the di, tri, or tetra Hydrolysis reaction: breaking down of molecule by water,
prefix sped up by acid or alkali (esters and alkenes)
If the ring of carbon is present, use the prefix “cyclo”
Write the position of the double bond in alkene, e.g. but-1-
13.9. Oxidation and Reduction
ene
Oxidation: addition of oxygen or removal of hydrogen
13.6. Breaking of Covalent Bonds Reduction: addition of hydrogen or removal of oxygen

Homolytic Fission: 13.10. Shapes of Ethane and Ethene

Two atoms sharing e- pair of similar electronegativity
Ethane: sp3 bonds, all sigma bonds
When bond breaks, each atom takes one e- from pair of
electrons forming free radicals
Free radicals: electrically neutral atoms or groups of
atoms with unpaired electrons are very reactive
Free radical reaction catalysed by heat or light

Heterolytic Fission:

Two atoms sharing e- pair are of different

electronegativity
When the bond breaks, one of the bonded atoms takes Ethene: Planar Shape, H – C – H bond = 120o
-
both bonding e s
Results in the formation of +ve and –ve ions
If +ve charged on C, it is called carbocation or carbonium
If –ve charged on C, it is called carbanion

Note: homolytic fission requires less energy than heterolytic

13.7. Types of Reagents

Nucleophilic reagent (nucleophile): donator of pair of e-

Benzene
Must have lone pair of e-s
Attack centre of +ve charge (positive pole)
Reaction with nucleophile called nucleophilic reactions
Examples: CH-, Cl-, NH3, H2O, CN-

Electrophilic reagent (electrophile): acceptor of pair of e-

13.11. Isomerism
+ve ions or e- deficient molecules
Existence of two or more compounds with the same
Attack regions of high e- density molecular formula but different structural formula
Examples: Br+, CH3+, AlCl3

13.8. Types of Reaction

Addition reaction: single product formed
Electrophilic addition (alkenes)
Nucleophilic addition (carbonyl compounds)
Substitution reaction: two products formed
Note:

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Straight chain alkanes have higher b.p. than branched

Branching makes molecule more spherical reduces
contact points VDW forces decreases

13.12. Chain Isomers

Isomers have different carbon chain length
Same chemical properties but slightly different physical
Example:
13.16. Optical Isomers
Arises from different arrangement of atoms or groups in
3D space resulting in two isomers
Have effect on polarised light
Chiral carbon: a carbon having 4 single bonds and 4
different atoms or groups
Isomers non-super-imposable images of each other
Have same physical and chemical properties
13.13. Position Isomers No. of optical isomers in a molecule containing n chiral
carbons = 2n
Isomers differ in position of substituent atoms or group or
the functional group
Same chemical properties but slightly different physical
Example: But-1-ene

14. Hydrocarbons
Example: But-2-ene
14.1. Properties
Generally unreactive:

All C–C bonds single; alkanes = saturated hydrocarbons

13.14. Functional Isomers Non-polar ∴ no center of charge to act as either
nucleophile or electrophile ∴ cannot attract polar
Isomers have different functional groups, belong to reagents like acids, bases, metals or oxidizing agents
different homologous series
Have different physical and chemical properties Physical properties:

Ratio of C : H Functional Gps. Example The volatility of the alkanes decreases and m.p/b.p
increases as number of carbon atoms increases
1:3 Alcohol & Ether C 2H­6O
Reason: increasing Van der Waals forces
1:2 Aldehyde & Ketone C 3 H6 O
1:2 14.2. Combustion
Carboxylic acid & Ester C 3H6O2
Must have O2
Used as fuel because they burn in oxygen to given out
13.15. Geometric (cis/trans) Isomers large amounts of energy
Alkanes kinetically stable in presence of O2; combustion
Shown only by alkenes occurs when necessary amount of Ea supplied
Arises due to restriction of double bond Reaction occurs only in gas phase
Only possible when each carbon has 2 different groups Complete: carbon dioxide + water
cis-trans isomers have different b.p. Incomplete: carbon monoxide + carbon (soot) + water
cis isomers have higher dipole General Equation of Hydrocarbon Combustion:
trans isomer of symmetrical alkene has zero dipole y y
C xHy + (x + ) O2 → xCO 2 + H2 O
4 2
​ ​ ​ ​ ​

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14.3. Substitution 14.5. Cracking

Alkanes react with halogens: Cl2 and Br2 Breaking of large less useful alkanes into useful, more
energy value smaller products using heat & catalyst
Example: Chlorination of Methane Products:
Reaction smaller alkanes and alkenes or
Reagent Condition
Type Mechanism smaller alkenes and hydrogen gas
Cl2(g) UV light Substitution Free Radical Thermal cracking: high temp. & pressure
Catalytic cracking: high temp. & catalyst

14.6. Hydrocarbons as Fuels

Source of alkanes: crude oil
Steady change in b.p. of alkanes allows crude oil to be
separated by fractional distillation
Catalytic conversion of CO and NOx:
2NO2 + 4CO → N2 + 4CO2
2NO + 2CO → N2 + 2CO2

14.7. Alkenes
Unsaturated hydrocarbons
Reactants are Halogens and Alkane. Contain at least one C=C double bond
Involves 3 steps: initiation, propagation, and termination. General formula: C nH2n (like cycloalkanes)
requires the action of UV light or Heat.
Source of alkenes:
Cracking alkanes
14.4. Initiation Dehydration of alcohols
More reactive than alkanes due to presence of double
Breakdown of Chlorine into radicals. bond; pi electrons loosely and more susceptible to attacks
Uses the action of UV light. by e- deficient groups like electrophiles
Starts the reaction. Alkenes combust completely carbon dioxide + water
Creates radicals. Give energy but not used as fuels; have other uses
C l2(g) → 2C l(g)
o ​ ​

14.8. Electrophilic Addition Mechanism

Propagation
Chlorine radical reacts with the alkyl/alkane.
Helps the reaction to propagate (chain reaction).
Maintains balance of radicals.
C H4 + C lo → C H3o + HCl
​ ​

C H3o + C l2 → C H3 C l + C lo
​ ​ ​

Termination
Electrophile forms by heterolytic fission
Chlorine radical and alkyl radical reacts. Electrophile attacks double bond
Stops propagation of reaction.
Pair of e-s from double bond migrate to electrophile and
Reduces the number of radicals.
π bond breaks
C H3o + C lo → C H3 C l
Carbocation formed which attacks the nucleophile
​ ​

2C lo → C l2 ​

C H3o + C H3o → C H3 C H3
14.9. Carbocations
​ ​ ​ ​

Hint:

C H3o and C lo represents radicals which is neutral atom

​

and an unpaired electron.

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This leads to the rupture of the double bond

Two compounds are formed
Products formed depend on alkene

Markovnikov’s principle: an electrophile adds to an

unsymmetrical alkene so that the most stable carbocation
is formed as an intermediate
Hydrogen binds to carbon with more hydrogens

Inductive effect of alkyl groups:

Alkyl groups donate e-­­ to the ring
Producing a positive inductive effect
A larger alkyl group has a weaker inductive effect

14.10. Addition Reactions 14.12. Polymerisation

Hydrogenation (Alkene + H2 Alkane) Repeated addition of 1000s of alkene molecules
Reagent: H2(g) (monomer) to each other forming a macromolecule
Condition: Polyethene:
Catalyst: Nickel LDPE: cling wrap
Temp.: 100oC HDPE: water pipes, wire insulation
Press.: 2 atm.
Use: convert liquid oils to saturated solid fats
Halogenation (Alkene + X2 Dihaloalkane)
Reagent: Halogen(aq) Polychloroethene (PVC):
Condition: r.t.p./dark Water pipes
Halogenation (Alkene + Hydrohalogen Halogenoalkane) Insulation of wires
Reagent: Hydrohalogen(g)
Condition: r.t.p.
Hydration (Alkene + H2O(g) Alcohol)
Reagent: steam General conditions: high pressure,
Condition: high temperature and catalyst
Catalyst: H3­PO4 – phosphoric acid Disadvantages:
Non-biodegradable: does not break down so increases
Temp.: 300oC
the amount of volume needed for landfill sites
Press.: 70atm
Combustion produces harmful gases which contribute
to global warming e.g. SO2, CO­2 and HCl from PVCs
14.11. Oxidation of Alkenes Disposal of Polymers:
Recycle existing plastic
Both oxidation and addition to double bond involved
Make polymers biodegradable by adding starch units
KMnO4 changes from purple/pink to colourless

With Cold Dilute Acidified KMnO4/H+ 15. Halogen Compounds

Diol is formed
15.1. Types of Halogenoalkanes

Primary 1o (S N 2)

With Hot Concentrated Acidified KMnO4/H+

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Bimolecular – two molecules involved in 1st step

Primary and secondary halogenoalkanes

15.4. Nucleophilic Substitution Reaction

Secondary 2o (S N 2 and S N 1)
Hydrolysis (R – X + OH- R – OH + X-)
Reagent: strong alkali; NaOH(aq) or KOH(aq)
Condition: heat/reflux
Fluoroalkanes are not hydrolysed because the C – F
bond is too strong
Ease of hydrolysis increases: Primary < Secondary <
Tertiary
Tertiary halogenoalkanes can be hydrolysed without
Tertiary 3o (S N 1)
alkali
Note: if any Cl- or Br- ions present in NaOH(aq), these
ions will interfere with reaction
Nitrile (cyanide) (R – X + CN- RCN + X-)
Reagent: KCN or NaCN in ethanol
Condition:
­
Solvent: Ethanol
Heat/Reflux
Reaction forms a C – C bond; therefore number of
15.2. Strength of C – Hal Bond Carbon increases; name has one more carbon
Primary Amines (R – X + NH3 RNH2(l) + HX(g))
Polar Nature Bond Energy Reactivity Reagent: Ammonia (NH3)
Fluoro Condition: Ammonia in alcohol under pressure in a
Chloro sealed container
Decrease Decrease Increases Note: If excess concentration of ammonia used, HX
Bromo
reacts with it forming NH4X
Iodo
Electronegativity Bond length increases, bond
decreases energy decreases, lower EA so
15.5. Reflux
down group more reactive
Many organic reactions proceed slowly
Heating done under reflux to prevent volatile organic
15.3. Nucleophilic Substitution solvents to evaporate
Mechanism similar to simple distillation
Mechanism
The C-X bond is a polar bond and has partial charges due
to the high electronegativity of halogen.
The δ+ carbocation is easily susceptible to attack by a
nucleophile

S N 1 Mechanism:

15.6. Elimination Reaction

R – X + OH- Alkene + X- + H2O

Unimolecular – only one molecule involved in 1st step Mechanism:
Secondary and Tertiary halogenoalkanes

S N 2 Mechanism:

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Reagent: ethanolic NaOH or KOH

Conditions: temp. 60oC, reflux
Secondary 2o
OH- acts as a proton acceptor; it accepts the H+ loss from
the halogenoalkanes during elimination
Elimination becomes progressively easier

Primary < Secondary < Tertiary

Note: the carbon atom adjacent to carbon with halide

must have at least one hydrogen attached to it.
Tertiary 3o

15.7. Uses of Halogenoalkanes

CFCs are inert and can be liquefied easily: Strength of C –
X bond is very high, hence do not decompose easily and
are not flammable.
Uses:
As propellants in aerosol cans
Source of Alcohols:
As solvents in dry-cleaning
Hydration of alkenes
As refrigerant for freezers and fridges
Fermentation
Fire extinguishers, insecticides and pesticides

16.2. Properties
15.8. CFCs Effect on Ozone Layer
Physical Properties:
Destroys the ozone layer
CFCs escape the atmosphere and, because of their Colourless liquids at r.t.p
inertness, remain without further reaction until they reach b.p. and density increases with increasing C atoms and
the stratosphere and ozone layer. also with increasing OH groups
In the stratosphere, high energy U.V causes the Cl atom to
split CFC molecule forming Cl⋅, which reacts with ozone
This is a catalytic cycle where one Cl⋅ can react with many Boiling Point:
O3 thus causing destruction of ozone layer: b.p. decreases→

Cl⋅ + O3(g) ⋅OCl(g) + O2(g)

Because:
⋅OCl(g) + O(g) Cl⋅ + O2(g)
Can react and breakdown another O3 molecule

Note: the alternative is using HCFCs (replace Cl with H or b.p. of alcohols > alkenes as they have hydrogen bonds
more F atoms) as they break down more easily and do not
release Cl → less effect on the ozone layer
Solubility of Alcohols in Water:

Smaller alcohols mix completely with water since strong

16. Hydroxy Compounds hydrogen bonds occur between alcohol and water
As hydrocarbon nature increases (i.e. more C-C… bonds),
16.1. Types of Alcohols the non-polar character outweighs the ability of the OH to
form hydrogen bonds and ∴ solubility decreases
Small alcohols (e.g. ethanol) are good solvents for both
Primary 1o
polar and non-polar compounds as they have polar and
non-polar components

16.3. Reaction with Sodium

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R – OH + Na(l) RO- Na+ + ½ H2(g) Type of Reaction: Nucleophilic Substitution

Type of reaction: acid-base R – OH R – X

Reagent used: liquid sodium metal Forming
Producing: Reactions Condition
Reactivity of alcohols decreases with increasing chain Reagent
lengths of hydrocarbon Conc. HCl
Reaction less vigorous than that of Na and water which Zn +
RCl(l) +
shows water is a stronger acid than alcohol Heat/Reflux
H2 O
SOCl2
16.4. Reaction with Carboxylic Acids RCl(l) +
SO2(g) +
HCl(g)
Alkyl r.t.p
PCl5 RCl(l)
Chlorides
+
POCl3(aq)
R– + HCl(g)
OH+
PCl3 RCl(l)
+
Heat/Reflux
H3PO3(aq)
Reagent Condition Type of Reaction
+ HCl(g)
Heat-reflux
R-COOH Esterification NaBr + HBr(g)
Conc. H2SO4
H2SO4(aq) RBr(l) +
HBr Alkyl H2 O
as fats, oils & flavours in fruits r.t.p
P + Br2 – Bromides PBr3(g)
Used in food flavourings and perfumes and as solvents warm RBr(l) +
PBr3 H3PO3(aq)
16.5. Hydrolysis of Esters PI3(g) RI(l)
P + I2 – Alkyl
+ r.t.p
warm PI3 Iodide
H3PO3(aq)

16.8. Oxidation of Alcohols

Reagent: Oxidising agents

Reagent Type of Reaction

16.6. Dehydration of Alcohols Acidified K2Cr­2O7 Acidified KMnO4

Oxidation
Orange to Green Pink to Colourless
Alcohol(l) Alkene + H2O(l)
Condition: png)
Conc. H2SO4 OR
Tertiary alcohols not oxidised because no hydrogens
H3PO4 at 180oC OR attached to carbon with OH group so oxidising agent
Al2O3 at 300oC colour does not change
Type of reaction: Elimination
16.9. Tests for Alcohols
Mechanism:
Result with:
Reagent
Primary Secondary Tertiary
Na metal Bubble of H2 Gas

Adjacent carbon to carbon with OH must have at least one K2Cr­2O4 /H+ Green X

hydrogen (tertiary cannot undergo dehydration) KMnO4/H+ Colourless X

16.7. Halogenation

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17. Carbonyl Compounds

17.1. Introduction
Boiling Point:
Note: HCN is a poor nucleophile with few CN- Ions, the
reaction is slow. To increase CN- conc.:

Make HCN react in the presence of alkali

Solubility: HCN + OH- H2O + CN-

Smaller carbonyl compounds: completely soluble as they Addition of KCN and dilute H2SO4 can provide HCN and
form hydrogen bonds with water molecules; are good
solvents for polar & non-polar solutes more CN- ions
Larger carbonyl compounds: polar nature decreases, and
non-polar nature increases, ability to form hydrogen 17.4. Reduction of Carbonyl Compounds
bonds decreases
Type of Reaction: nucleophilic addition (H- ions)
17.2. Production of Carbonyl Reducing agents:
NaBH4 – sodium tetrahydrioborate
Compounds LiAlH4 – lithium aluminium hydride
H2/Pt or Ni
Oxidation of Primary Alcohols (Aldehydes)
Aldehydes ⟹ 1o Alcohols
Reactants: acidified K 2 C r 2 O7 /K M nO3 .
​ ​ ​ ​
R-CHO + 2[H] RCH2OH
Distillation to separate Aldehyde from other compounds. Ketones ⟹ 2o Alcohols
Aldehyde can easily be oxidized into Carboxylic acid, so it R-CO-R + 2[H] R-CH(OH)-R
needs to be distilled as soon as it forms.
C H3 C H2 OH + [O] => C H3 C HO + H2 O.
​ ​ ​ ​

17.5. Testing Carbonyl Compounds

Oxidation of Secondary Alcohols (Ketones) 2,4,-dinitrophenylhydrazine:

Reactants: acidified K 2 C r 2 O7 /K M nO3 .

​ ​ ​ ​

Distillation to separate Ketone from other compounds.

Ketones cannot be further oxidized.
C H3 C HOHC H3 + [O] => C H3 C OC H3 + H2 O.
​ ​ ​ ​ ​

17.3. Nucleophilic Addition with HCN It is a nucleophilic addition & condensation/elimination.

Forms: red/orange ppt.
The m.p. of the ppt. can be used to identify individual
Reagent Condition Type of Reaction
aldehydes and ketones
HCN HCN w/alkali or HCN w/KCN Nucleophilic Addition
Tests Given only by Aldehydes:
Since HCN was added, the carbon chain increases
The product formed is hydroxynitrile or cyanohydrine Tollen’s Reagent
Aldehydes are more susceptible to nucleophilic attacks
than ketones Solution of Ag(N O3 )2 + N H3 excess
​ ​ ​

More minor carbonyl compounds are more reactive Aldehyde + Tollens agent Silver Mirror
Product has a chiral carbon ∴ exhibits optical isomerism Ag + reduced to Ag , and –C HO oxidized to the acid

Mechanism: 2Ag + + RC HO2 => Ag + RCOOH − + H +

Fehling’s Solution
C uSO4 in ammonia solution
​

Aldehyde + Fehling’s Solution Red ppt.

C u 2 + reduced to C u (I) oxide and –C HO oxidized
​ ​

to acid.

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2C u 2 + +RC HO => 2C u + > +RCOOH −

​
Primary Halogenoalkanes
Ammonia
Heat under Pressure
18. Carboxylic Acids and Ethanol (Catalyst)

Derivatives Reaction:
C H3 C H2 B r + N H3 => C H3 C H2 N H2 + HBr
​ ​ ​ ​ ​ ​

18.1. Introduction
19.3. Formation of Nitriles
Weak acids don’t dissociate completely
Forms hydrogen bonds: Reactants and Conditions
High melting and boiling points
High solubility of smaller carboxylic acids Halogenoalkanes
Forms hydrogen-bonded dimers when pure vapour, liquid K CN
or solid & when dissolved in non-polar organic solvents Heat under Reflux
Ethanol (Catalyst)

Reaction:
C H3 C H2 B r + K CN => C H3 C H2 C N + K Br
​ ​ ​ ​

18.2. Formation of Carboxylic Acids 19.4. Formation of Hydroxynitriles

From alcohols: complete oxidation of primary alcohols Reactants and Conditions
From aldehydes: oxidation of aldehydes
From nitriles: acid/base hydrolysis of a nitrile Aldehydes or Ketones
HCN
Heat under Reflux
KCN (Catalyst)

Reaction:
C H3 C H2 C HO + HCN =>
​ ​

C H3 C H2 C (OH)(CN )H
18.3. Formation of Salts
​ ​

C H3 C H2 CO + HCN =>
​ ​

C H3 C H2 C (OH)(CN )C H3
​ ​ ​

Heterolytic fission of the hydroxyl bond (-OH)

Salts called carboxylates
19.5. Hydrolysis of Nitriles

Reactants and Conditions

Nitriles
Heat Under Reflux
Acid/Base
19. Nitrogen Compounds
Acid Hydrolysis

19.1. Introduction Products:

Types: Salt
Carboxylic Acid
Amines ( −N H2 ) ​

Amides ( −CO − N H2 ) {not included in AS}

​
Reaction:
Nitriles ( −CN ) C H3 C H2 C N + HCl => C H3 C H2 C O2 H + N H4 C l
​ ​ ​ ​ ​ ​

Base Hydrolysis
19.2. Formation of Primary Amines
Products:
Reactants and Conditions
Base

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Salt of Carboxylic Acid LDPE HDPE

Hard plastic made with non-
Reaction: Highly branched chains
polar polymer
C H3 C H2 CN + N aOH => C H3 C H2 C O2 N a + N H3
​ ​ ​ ​ ​ ​

Low melting point Unbranched chains

Chains cannot pack closely
20. Polymerisation due to random branching, High melting point
leaving gaps
20.1. Addition Polymerisation Low density
Closely packed chains, fewer
gaps
Monomers are alkenes; the polymer formed is polyalkene High density
Alkene joins to itself, and no molecule is lost
E.g. nA → −A − A − A − A− or (−A−) ­ n ​

20.2. Environmental Impact of

The molecular mass of the polymer is multiple monomers
Poly(alkenes)
Physical Properties of Polymers:
Non-biodegradable
Only C and H present = only VDW forces present
They can not be decomposed by the action of bacteria
Presence of Cl or F = dipoles = permanent dipole
Dangerous to Decompose
interaction
Produce harmful gases into the air when burned
Longer chains with fewer side chains have stronger VDW
It takes hundreds of years to turn into decomposable
forces
material
Results in large amounts of landfills or Garbage
Patches at Oceans

21. Organic Synthesis

polyethene polytetrafluoroethene
21.1. Introduction
Branch of Organic Chemistry.
Used to construct new chemicals artificially from
petroleum or other organic compounds.
For example, make propanoic acid from ethane.
Involves multiple steps and reactants.
It can have more than one route for a specific product.

21.2. Things to Remember

Reactants and products
Reaction Mechanisms
Electrophilic Addition
Nucleophilic Substitution
Free Radical Substitution
Nucleophilic Addition
Conditions for reactions like
Organic Compounds like ethanol
Heat under Reflux
Pressure/Closed Container
Catalyst like N aBH4 , acids ​

Poly Vinyl Chloride (poly-chloroethene)

LDPE HDPE
Soft plastic made with non-
Using Ziegler catalyst
polar polymer

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Chemical substances are identified by their mass-to-

charge ratio.
The substance is ionized and given charge.
The charged ions pass through magnetic and electric
fields to calculate the mass-to-charge ratio of each ion
present.

22. Analytical Techniques The abundance of the ions is used to identify different
elements and functional groups of a compound.
22.1. Infra-red Spectroscopy
22.4. Isotopic Abundance of Elements
This is when a sample being analysed is irradiated with
electromagnetic waves in the infra-red region of the Identifies the relative abundances of different isotopes in
electromagnetic spectrum. a sample.
The machine used is a spectrophotometer, and it detects
the intensity of wavelengths of infra-red that pass through
the sample
The energy absorbed corresponds to changes in the
vibration of bonds, leading to the bond being stretched,
bend and twist
At a specific frequency, the resonance frequency, the
largest vibrations are obtained
Each type of vibration will absorb characteristic
wavelengths of infra-red radiation
We can hence identify the presence (or absence) of
different functional groups from the absorbance pattern
The 2 left most peaks represent the 2 isotopes of Chlorine
on an infra-red spectrum
i.e., Cl-35 and Cl-37.
The right-most 3 peaks represent the Chlorine molecules
with different arrangements with Cl-35 and Cl-37.
The ratio between the Cl-35 and Cl-37 peaks represents
the relative abundance of the isotopes.

22.5. Organic Fragment Pattern

Alkyls like C H3 - form Fragment patterns on Mass
​

22.2. Monitoring Air Pollution Spectrometer

We can use the Molecular ion peak ( M + ) as identified ​

IR spectroscopy identifies particular bonds in a molecule, here

and so each pollutant will show a different pattern of
absorption – this allows the identification of the pollution
It is also possible to measure the concentration of each
pollutant with the different amounts of absorption

22.3. Mass Spectrometry

Introduction

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Here, 72 is the Molecular ion peak. For 2 Chlorine:

It is the total mass of the organic molecule.
Used to identify Carbon atoms present in the Compound
using formula.

100 ∗ M +1
n (N umber of C arbon atoms) =
​

1.1 ∗ M +
​

M +1 Represents the abundance of peak to the right of

​

molecular ion peak.

It is smaller than the M + peak and caused by the C-13
​

isotope.
1 in every 100 Carbon is C-13, so it has a ratio of M + ​

depending on the number of carbon atoms.

Different atoms and functional groups create several
There are M +2 and M +4 peaks alongside M + peak.
peaks that the table can identify.
​ ​ ​

They have a difference of 2 m/z.

M +4 , M +2 , M + are in the ratio 9:6:1.
22.6. Halogen Identification
​ ​ ​

If the conditions are met, there are 2 Chlorine atoms.

Uses M +2 peak to identify Chlorine and Bromine atoms

​

For Bromine:
in the organic compound.

For 1 Chlorine:

There are M +2 peaks alongside M + peak.

​ ​

They have a difference of 2 m/z.

M +2 , M + have almost same height.
​ ​

M +2 is in 3:1 to M +
​ ​
If the conditions are met, there is a Bromine atom.
Both have a difference of 2 m/z due to the isotopes of
Chlorine.
If the conditions are met, there is a Chlorine atom.

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Chemistry

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