Cover Page
Chemistry
Paper 1 (100 Marks): 30 MCQ, 40 marks Short –
Questions, 30 marks Structured Questions
1. Kinetic Particle Theory
2. Measurement and Experimental Techniques
3. Purification and Separation
4. Elements, Compounds, Mixtures
5. Atomic Structure
6. Ionic and Covalent Bonding
7. Metallic Bonding
8. Writing Chemical Equations
9. The Mole
10. Chemical Calculations
11. Acids and Bases
12. Salts
13. Oxidation and Reduction
14. The Periodic Table
15. The Atmosphere and Environment

Pure Chemistry SA2 Overall Revision Notes
Chapter 1
Kinetic Particle Theory
Kinetic Particle Theory States that all matter is made up of tiny particles that are in constant,
random motion
Solid Liquid Gas
Arrangement of the Close together in an Close together in a disorderly Far apart in a random
particles orderly arrangement arrangement arrangement.
Forces between particles Strong Strong None
Movement Vibrations about fixed Vibrations and movement Vibrations and movement
positions increase as throughout the liquid anywhere.
temperature increase.
Energy of the particles Less Energy More Energy
Density High (Particles close High (Particles close together) Low (Particles are far apart)
together)
Energy Changes Sublimation
Liquid  Gas
It occurs because particles at the surface of the
solid have enough energy to break away from
the solid and escape as a gas and particles of
the liquid are too weak to remain in that state.
Some examples of substances that sublime
include iodine and ammonium chloride.
Substances that sublime can change from
gaseous state to solid state through
condensation.
Dry ice is useful as it can keep food cold and
change into a gas without leaving liquid behind.
1. Between points A and B, the substance is completely solid. 1. Between points A and B, substance is completely
2. The substance starts melting at Point B. The temperature is liquid. It starts melting at point B which is the
called the melting point. freezing / melting point.
3. Between points B and C, the substance is melting and a 2. Between points B and C, the substance is freezing and
mixture of solid and liquid is present. Temperature is constant a mixture of liquid and solid is present. Temperature
as heat energy is taken to weaken and break the bonds. is constant as heat energy is given out to make bonds
4. The substance has completely melted at point C and became a between the particles.
liquid. 3. The substance has completely turned into a solid at C.
Diffusion
There are 2 definitions for diffusion:
Diffusion is a process in which particles travel from a region of higher concentration to a region of lower concentration.
Diffusion is a process whereby particles move freely to fill up any available space.
The 2 main factors affecting the rate of diffusion:
Relative Molecular Mass (the smaller the faster the speed)
Temperature (the higher the faster the speed of the molecules  Faster diffusion).

Pure Chemistry SA2 Overall Revision Notes
Chapter 2
Measurement and Experimental Techniques
Measurement SI unit and other units Instruments used
Time Seconds (s); minute (min) and Stopwatch and Stopclock
hour (h)
o
Temperature Kelvin (K); degree Celsius ( C) Mercury thermometer, clinical thermometer.
Mass Kilogram (kg), Grams (g) Beam Balance, electronic balance
Volume Cubic Metre (cm3) Beaker: Used to estimate the volume of a liquid.
Litre Measuring Cylinder: More accurate than beaker
Milimetre Burette: Measures up to nearest cm3. Accurately measures volume of
liquid to 0.1cm3
3
Pipette: Scale marked in 0.1cm divisions. Accurately measures fixed
volumes of liquids.
How a gas is collected depends on its physical properties, namely:
Solubility: How soluble the gas is in water.
Density: How dense the gas is compared to air.
Collecting Gases and Measuring Techniques
Method Solubility Density Example
Carbon Dioxide
More dense than air Hydrogen
Chloride
Gas is more soluble in water
Less dense than air Ammonia Gas
Gas is insoluble or slightly soluble Hydrogen
Gas is less dense
in water Oxygen
than air
Methane
A gas can be dried by passing it through a drying agent. Some examples of drying agents include:
Concentrated sulphuric acid
Quicklime (Calcium Oxide)
Fused calcium chloride
In the apparatus used, the tube introducing the gas is immersed in the acid whereas the exiting tube is not.
Concentrated sulphuric acid is used to dry most gases including chlorine and hydrogen chloride. However, it cannot be used
to dry ammonia as it reacts with ammonia.
Instead, Quicklime is used (placed above the moist ammonia).
Fused calcium chloride can also be used to dry most gases.
A gas syringe is used to measure the volume of a gas.
The gas is pushed in fully to expel any gas in the syringe.
As the gas from an external source enters the syringe, it pushes the plunger outwards.

Pure Chemistry SA2 Overall Revision Notes
Chapter 3
Purification and Separation
Chapter 3.1: Testing for Purity Chapter 3.2: Filtration
Separate a mixture of a solid and a liquid.
Pure Substance Impure Substance Upon filtration, the solid that remains on the
Fixed melting The impurities will lower the melting point. The filter paper is called the residue. The liquid
points greater the amount of impurities, the lower the that passes through the filter paper is called
melting point of the substance. the filtrate.
Cause melting to take place over a range of It should be done several times to ensure that
temperatures1. the entire solid is separated from the liquid.
Fixed boiling Impurities will increase the boiling point of a
points substance. The greater the amount of impurities,
the higher the boiling point of the substance.
1
If pressure acting on a liquid is increased, the
boiling point is raised. Likewise, if pressure is
decreased, boiling point will fall.
1 single spot seen Two or more spots seen on the chromatogram.
on the
chromatogram
Chapter 3.3: Crystallization Chapter 3.4: Evaporation to Dryness
Obtain a solid from the solution.
Solution is heated to make it 1. Easiest method to separate the soluble solid from the liquid.
saturated. 2. It is used to recover the solid particles from the mixture.
The saturated solution is allowed to 3. This process is done through evaporation.
cool. 4. As some compounds may decompose or break down on heating,
As it cools, crystals of solid appear. crystallization is thus recommended.
They are pure as impurities remain 5. In addition, when all the water is removed during evaporation, any soluble
dissolved in the solution. impurities present will be left on the crystals.
The crystals are filtered off and dried 6. This process is mainly used only if the solid has a high melting point or not
by pressing them on sheets of filter volatile (e.g. sea water to recover salt)
paper.
Chapter 3.5: Chromatography Chapter 3.6: Use of a Separating
Funnel
Used to separate and identify small amounts of substances that are
dissolved in solvents. Liquids that do not dissolve in each other are
Paper Chromatography can be used to separate and identify dyes in described as immiscible. (e.g. oil and water)
coloured ink.
Drops of the unknown ink and known dyes are placed side by side These liquids cannot mix with each other due
on a piece of filter paper. to their differences in density.
The paper is suspended into a solvent.
As the solvent moves up, the ink dissolves in the solvent and To separate immiscible liquids, a separating
move up the paper at different speeds and distances. funnel is used. The denser liquid is collected
The pattern of spots is called a chromatogram. first till it is all collected and the remainder is
A locating agent can be sprayed on the paper to make colourless collected into another beaker.
spots visible.
Every substance has an unique Rf value which can be calculated from Liquids like ethanol and water that mixes
Distance travelled by substance together completely to form a solution are
R f Value  miscible. A technique called fractional
Distance travelled by solvent
distillation can be used to separate them.
The Rf value of a substance does not change as long as
chromatography is carried out under the same conditions
(temperature, solvent, etc).

Chapter 3.7: Simple and Fractional Distillation
Simple Distillation Fractional Distillation
Diagram
Similarities in Apparatus Purpose / Things to note
Apparatus Boiling Chips Smoothen boiling and prevent bubbling.
Thermometer The bulb of the thermometer must be at the opening of the condenser to measure the
temperature of the vapour as it leaves the distilling flask.
Condenser It should be slanted to allow the gas to touch the walls of the condenser easily and transfer
heat to the cold water, allowing the vapour to cool down. If it is horizontal, the chances of the
vapour condensing are limited.
Cold running water should also be allowed to run from the bottom of the condenser and leave
from the top.
Water in / The “water in” is placed at the end to distribute the water to ensure that the solvent is
Water out entirely condensed to get back the solvent and to prevent the vapour from escaping.
1
Beaker If the distillate collected is volatile , place it into a bigger container filled with ice.
Fractioning Only in Fractional Distillation. Used to allow the liquid with higher boiling point to condense
Column and fall back whereas the liquid with lower boiling point will rise up.
Uses Recover a solvent from a non-volatile source. Separate mixtures whose boiling points differ less
Separate mixtures of liquids where boiling points differ at than 25°. Used in industries to obtain nitrogen, argon
least 25°C. and oxygen, or used to separate mixtures of liquids
like crude oil.
Steps taken 1. Solution boils; water vaporizes and enters the 4. Solution boils; solution vaporizes and enters the
condenser. condenser.
2. Water vapour is cooled in the condenser and 5. The liquid with lower boiling point will become
changes back into pure water which is collected in vapour and rise up. The liquid with higher boiling
the beaker. point will condense in the fractioning column.
3. The salt solution becomes more concentrated and 6. When all the liquid that has the lower boiling
eventually becomes salt if distillation is allowed to point condenses, the temperature will rise
carry on. quickly and allow the other liquid to rise over.
Temperature
100 120
recorded 100
Temperature
80
Temperature
80
60 60
40 40
20
20
0
0
Time
Time
1
The distillate collected can become gas easily with a very low boiling point (can be below room temperature).

Pure Chemistry SA2 Overall Revision Notes
Chapter 4
Elements, Compounds, Mixtures
Chapter 3.1: Elements Chapter 3.2: Compounds
1. An element is a substance that cannot be broken down into 1. A compound is a substance containing 2
simpler substance by chemical methods. or more elements chemically joined
2. The 3 most abundant element: Oxygen, Silicon, Aluminum together.
3. Can be classified by state (2 liquids – mercury and bromine) 2. It has different properties from its
or as metals, non-metals and metalloids. constituent elements.
a. Metals (except graphite) conduct electricity 3. Made up of molecules or ions
b. Non – metals include noble gases 4. A compound is made up of different
c. Metalloids conduct electricity when impure / heated. E.g. elements chemically combined in a fixed
Silicon. ratio.
4. Elements are made up of atoms or molecules. 5. An ion is an atom or a group of atoms
a. An atom is the smallest unit of an element. with an electrical charge. Many
b. A molecule is a group of 2 or more atoms chemically compounds which consist of ions, known
joined together (represented by a chemical formula) as ionic compounds and most are solids.
Chapter 3.3: Mixtures
1. A mixture consists of 2 or more substances not chemically combined together.
2. It may contain different elements (brass – alloy), different compounds (sea water, milk) or both elements and
compounds (air).
3. Differences between a pure compound and mixture:
Pure Compound Mixture
Has a fixed percentage by mass of each element. Has a variable composition by mass of each element.
Has a chemical formula. Does not have a chemical formula.
Cannot be separated into its elements by physical Can be easily be separated into its components by
means. physical means.
Has a fixed melting point and a fixed boiling point. Has a variable melting and boiling point.

Pure Chemistry SA2 Overall Revision Notes
Chapter 5
Atomic Structure
Chapter 5.1: Nuclear Model of an Atom Chapter 5.2: Isotopes
Particle Symbol Relative Mass Charge 1. Isotopes are atoms of the same element with
Proton p 1 +1 different number of neutrons.
Neutron n 1 0 2. Most elements consist of a mixture of isotopes.
Electron e 1/1840 -1 3. Isotopes have slightly different physical properties
but have identical chemical properties as they
1. Proton Number: Number of protons in an atom. have the same number and arrangement of
2. Nucleon Number: Number of protons and neutrons electrons.
in an atom. 4. The relative atomic mass of can be calculated by
3. Number of electrons = Number of protons. adding the atomic mass with percentage of each
Therefore atoms are neutral (no charge) isotope.
4. An atom can be described in symbol form. 5. Isotopes can be used for heart pacemakers.
Chapter 5.3: Arranging Electrons in Atoms
1. The furthest shell from the nucleus is called the outermost shell. The electrons in this shell are known as the
valence electrons.
2. The valence electrons are used to form chemical bonds.
3. Metals have few valence electrons whereas non-metals have many.

Pure Chemistry SA2 Overall Revision Notes
Chapter 6 – 8
Ionic, Covalent, Metallic Bonding and Ionic Equations
Chapter 6.1: Ionic Bonding Chapter 6.2: Covalent Bonding
8. Ionic bonds are formed in compounds of metals 1. Covalent bonds are formed between atoms of non-
and non-metals. metals.
9. They are formed when electrons are transferred 2. 2 atoms form a covalent bond by sharing electrons
from one atom to another, forming ions. from their valence shell.
10. These ions are held together by electrostatic 3. Some atoms form 2 covalent bonds with another atom
attraction, called ionic bonding. The larger the known as double bonds. In such a bond, 4 electrons
difference of charges of the 2 ions, the stronger are shared.
the bond. 4. The covalent bonds between the atoms in the bromine
11. First, the metallic atom loses electrons to molecules are strong, however there are only weak
become a positive ion. Then the non-metallic forces in between the molecules.
atom gains electrons to become a negative ion. 5. Covalent compounds have a simple or giant molecular
The 2 ions are then held together by ionic bonds. structure.
12. Ionic compounds consist of large numbers of 6. Simple molecular structure: Have low melting and
ions in a giant ionic structure. boiling points due to weak van der vaals forces
13. Ionic compounds have high melting and boiling between the molecules. They do not conduct
points as the ionic bonds are strong. electricity as there are no free ions/electrons and are
14. They only conduct electricity in the liquid state usually insoluble in water.
as the ions can move and carry the current. 7. Giant covalent structure: Consist of huge number of
15. They are often soluble in water. atoms joined together by covalent bonding. Have high
16. Examples are Sodium Chloride and Magnesium melting / boiling points as they are joined together by
Chloride. strong covalent bonds. Examples are
diamond/graphite (carbon) and silicon oxide.
Chapter 6.3: Noble Gases Chapter 7.1: Metallic Bonding
1. Noble Gases like oxygen, helium 1. A metal consists of an orderly arrangement of positive metal ions
and argon are unreactive. surrounded by a sea of electrons which are free to move about.
2. They do not form chemical bonds 2. The metallic bond is the force of attraction between the positive
with other atoms as they have a metal ions and negative-charged electrons.
stable electronic configuration. 3. Metals have a giant lattice structure.
3. Atoms of other elements form 4. Solid metals are good electrical conductors as they have free-
chemical bonds so that they attain moving electrons.
the electronic configuration of a 5. Metals are malleable as the layers of atoms can easily slide over
noble gas. each other.
Chapter 8: Writing Ionic Equations AgNO3
1. Steps in writing an ionic equation:
Write a balanced chemical equation of the reaction including the state symbols.
Identify ionic compounds that are soluble in water. Rewrite the chemical equation Ag+ NO3-
in terms of ions.
Note: Only break down those which are aqueous state into positive and
negative. For example:
H2SO4
Cancel out the spectator ions.
Write out the ionic equation.
2. We need ionic equations as a reaction may have a lot of ions but only 2 ions react.
3. Acids are covalent but form ions in water. 2H+ (SO4)2-

Pure Chemistry SA2 Overall Revision Notes
Chapter 9 – 10
The Mole and Chemical Calculations
Chapter 9.1: The Mole Concept
Chapter 9.2: Calculating Empirical and Chapter 10.1: Calculations from Chemical
Molecular Formula Equations
A balanced chemical equation gives us the mole ratio
There are several steps in finding the empirical of the reactants and products involved in a chemical
formula of a compound: reaction.
1. Find Percentage, Mass or Relative Atomic
Mass Therefore, there are 3 steps to calculating the mass of
2. Find no. of moles (Mass, Percentage or Ar a substance reacted or produced:
/ Molar Mass) 1. Convert whatever mass is given of a substance
3. Divide by smallest ratio into the number of mole.
Notes: 2. Compare the mole ratio from equation.
If the relative molecular mass of a 3. Convert the number of mole to the mass of the
compound is known, the molecular substance you want to find.
formula can be found from the empirical
formula. Calculating volume of reacting gases from chemical
Leave all answers in 3 sig. fig. equations:
When divided by smallest ratio if it has a Since 1 mole of any gas occupies 24 dm3 at r.t.p.,
decimal of 0.5, multiply by 2. If it is 0.33 or the volume of gas is proportional to the number
0.67, multiply by 3. If not, round up or of moles of the gas.
down. Note: 1dm3 = 1000cm3
Chapter 10.2: Limiting Reactants Chapter 10.3: Calculation on Concentration of
Solutions
A balanced chemical equation is used to The concentration of a solution gives the amount of
calculate the exact amounts of reactants used solute in 1dm3 of solution. It can be expressed in
up and products formed using its molar ratio. g/cm3 or mol/dm3 (molar concentration).
The reactant that is completely used up is
known as the limiting reactant. It determines
or limits the amount of product formed.
Once the limiting reactant is used up, the
reaction stops.
Always use the limiting reactant to
calculate the product.
Chapter 10.4: Acid-Base Titration Calculations Chapter 10.5: Percentage Yield and Percentage
Purity
(Conc. of A)(Vol of A) No. of mol of A Actual Yield
 Percentage Yield =  100%
(Conc. of B)(Vol of B) No. of mol of B Theroetical Yield
Mass of pure substance
Percentage Purity =  100%
Mass of sample

Pure Chemistry SA2 Overall Revision Notes
Chapter 11 – 12
Acids, Bases and Salts
Chapter 11.1: Acids Chapter 11.2: Bases
6. An acid is a compound which produces hydrogen 1. Bases are metal oxides and hydroxides.
ions when dissolve in water. 2. A base is a substance which reacts with an acid to
7. A strong acid is one that is completely ionized in give a salt and water only known as neutralization.
water (HCl) whereas a weak acid is only partially 3. Base + Acid  Salt + Water
ionized in water. 4. Base + Ammonium Salts  Salt + Water + Ammonia
8. Acids have a sour taste, turn blue litmus red. 5. Soluble bases are also known as alkalis.
9. Acid + Metal  Salt + Hydrogen
10. Acid + Base  Salt + Water
11. Acid + Carbonate  Salt + Water + Carbon Dioxide
Chapter 11.3: Alkalis Chapter 11.4: Neutralization
12. Alkalis are bases which are soluble in water. 16. The reaction between an acid and a base is called
13. Alkalis produce hydroxide ions when dissolved in neutralization.
water. 17. In such a reaction, the hydrogen ions from the acid
14. Alkalis are slippery, turn blue litmus red. react with the hydroxide ions of the alkali.
15. Alkali reactions are similar to base reactions. 18. Neutralization reactions are exothermic.
Chapter 11.5: Oxides Chapter 11.6: Indications and pH
23. Oxides are compounds of oxygen with another element 19. Indicators are substance that turn to
(Copper oxide, sulfur dioxide). different colours in acidic and alkali
24. There are 4 types of oxides: solutions.
a. Acidic Oxides: Oxides of non – metals, dissolve in 20. The pH of a solution is a number that shows
water to give acids. React with bases to give a salt and how acidic or alkaline a solution is. A neutral
water. solution like ethanol has a pH of 7, an acidic
b. Amphoteric Oxides: Oxides of metals react with both solution like hydrochloric acid has pH of 2
acids and alkalis to form salt and water. (Zinc Oxides / and an alkaline solution like sodium
Aluminum Oxide) hydroxide has a pH of 14.
c. Basic Oxides: Oxides of metals, some dissolve in water 21. The lesser the pH, the more acidic, the
to give alkalis. React with acids to give a salt and higher the pH , the more alkaline the
water. substance is.
d. Neutral Oxides: Oxides of non – metals, do not react 22. Farmers neutralize excess acidity in the soil
with acids or bases. (Carbon Oxide, Nitrogen Oxide, by adding calcium hydroxide (slaked lime)
Water) to the soil.
Chapter 12.1: Salts Chapter 12.2: Preparation of Salts
23. Insoluble salts are prepared using the
16. A salt is obtained from an acid when the hydrogen ion precipitation method (mix 2 soluble salts and
of an acid is replaced by a metal or ammonium ion. obtain insoluble salt through filtration)
17. Solubility Table: 24. Soluble salts are prepared using acid + excess
Soluble Salts Insoluble Salts metal/insoluble metal oxide/carbonate (if
Carbonates SPA (Sodium, Potassium, All the rest reacting with insoluble substance) or titration
/ Hydroxides Ammonium) (acid + alkali reaction).
Chlorides All the rest Lead/Silver 25. After getting the salt solution, saturated it and
Nitrates All None then let it crystallize to get the salt.
Sulfates All the rest Barium,
Calcium, Lead

Pure Chemistry SA2 Overall Revision Notes
Chapter 13
Redox (Oxidation and Reduction)
Chapter 13.1: Oxidation Chapter 13.2: Reduction
2. The term oxidation represents a substance that is 1. The term reduction represents a substance that is
oxidized if it: oxidized if it:
Gain Oxygen Loses Oxygen
Loses hydrogen Gains hydrogen
Loses electrons Gains electrons
Increases its oxidation state after a reaction. Decreases its oxidation state after a reaction.
Chapter 13.3: Oxidation State
3. Oxidation – Increase in Oxidation State
The oxidation state is the charge an atom of an element would have if it existed as an ion in a compound
even if it is covalently bonded.
To work out the oxidation state of an atom, the following rules are applied:
 The oxidation state of a free element is zero.
 The oxidation state of a simple ion is the same as the charge on the ion.
 The oxidation states of the atoms present in the formula of a compound add up to zero.
 The total of the oxidation states of the atoms in a polyatomic ion is equal to the charge on the ion.
To determine the oxidation state of atoms in a compound, find the charges of each atom in the compound.
That is its oxidation state or number.
If the oxidation state of a substance increases, the substance has been oxidized.
Chapter 13.4: Redox Reactions Chapter 13.5: Oxidizing and Reducing Agents
6. Oxidation and Reduction usually occur together in 4. An oxidizing agent is a substance that produces the
reactions known as redox reactions. oxidation of another substance. Example: Potassium
7. Many redox reactions involve either electron Manganate / Potassium Dichromate / Chlorine.
transfer or change in oxidation state. 5. A reducing agent is a substance that produces the
8. Not all reactions are redox reactions, for example reduction of another substance. Example: Potassium
neutralization reactions and decomposition of Iodide, Carbon monoxide, hydrogen.
carbonates with heat are not.

Pure Chemistry SA2 Overall Revision Notes
Chapter 16
The Periodic Table
Chapter 16.1: Features of the Periodic Chapter 16.2: Patterns in the Periodic Table
Table
6. From left to right across a period, the proton number
1. A horizontal row of elements is called a increases, number of valence electrons increase, gradual
Period. change from metals to non-metals.
2. A vertical row of elements is called a Group. 7. Elements in the same group have the same number of
Group I: Alkali Metals, Group VII: Halogens, valence electrons, form ions with similar charges, form
Group 0: Noble Gases compounds with similar formulae, and have similar
3. Elements to the left of the zig-zag line are physical and chemical properties.
metals; elements to the right are non-metals. 8. When going down a group, the proton number, number
4. Period number: Number of electron shells. of electron shells and relative atomic mass increases.
5. Group number: Number of valence electrons. The properties also change.
Chapter 16.3: Group I (Alkali Metals) Chapter 16.4: Group VII (Halogens)
13. Elements are soft, silvery solids, low densities and 9. Elements are non – metals with low boiling
melting points. points. The consists of diatomic molecules.
14. They react vigorously with cold water to form an alkaline 10. They react with metals to form salts.
solution of metal hydroxide and hydrogen gas. 11. They have 7 valence electrons and gain
15. Elements burn brightly when heated in chlorine gas electrons with a negative charge of 1 in
forming metal chloride. compounds.
16. Have 1 valence electron. 12. When going down the group, the melting
17. Forms ions with a positive charge of 1. and boiling points increase, become darker
18. When going down a group, the melting points decrease in colour and become less reactive.
and reactions become more reactive.
Chapter 16.5: Group 0 Elements (Noble Gas)
19. The noble gases are all unreactive gases as they have full electron shells.
20. They are all monatomic gases.
21. Some important uses of noble gases:
a. Helium is used in balloons as it has a low density.
b. Argon and neon are used in light bulbs as they are unreactive.
c. Argon is used in extraction and welding of metals as it is unreactive and protect the hot metal from
reaction with oxygen in the air.