O level chemistry sec 3 exam revision notes

‘O’ Level Chemistry: Mid-Year Exam Revision Notes

We know Mid-Year Exams are just around the corner for most ‘O’ Level Chemistry Secondary 3 and 4 students.

Some of you may be feeling anxiety, or some might feel that while you study really well in class, you just aren’t getting the resources you need on the internet.

Tutopiya’s Team of Experts have already broken down the guide to Chemistry for ‘A’ Levels, and even helped your child prepare for the upcoming SA1 Primary School Examinations.

But we still have the GCE O Levels, which are crucial for your child to secure a top spot in a Junior College or Polytechnic of their choosing.

That’s why we’re breaking down systematically the top topics you’ll be tested for the O Level Chemistry Mid-Year Exam in 2019, especially if you are a secondary 3 or IP student.

Atomic structure

  • Isoelectronic
    • Atom, ions, groups of atoms having same total number of electrons
  • Isotopes
    • Atoms of the same element, having same number of protons, different number of neutrons
    • Finding Relative Atomic mass of isotopes
      • Add abundance of the first isotope multiplied by its atomic mass to abundance of the second isotope multiplied by its atomic mass and so on and so forth and then divide it by the number of isotopes
    • Electronic Structure
      • s,p,d,f,g orbital notation
        • Aufbau Principle
          • Electrons fill up orbitals of lowest energy first before proceeding to higher energy levels
        • Pauli exclusion Principle
          • An orbital can only hold max of 2 electrons of opposite spins
        • Hund’s rule of multiplicity
          • Electrons occupy different orbitals each first before pairing up in each orbital
        • 1s,2s,2p,3s,3p,3s,4s,3d……
        • Overlap in energies of certain layers
          • 4s and 3d, 3d is higher
          • In case of ions drawing electrons, will draw from 4s first before 3d even if it’s a higher orbital
        • s = 2, p =6, d=10, f= 14
        • Can write using noble gas core
          • 1s2,2s2,2p6,3s2 as [Ne] 3s2
        • Separation Techniques
          • Filtration
            • Separate mixture of insoluble solid and liquid
            • Mixture poured through filter paper
          • Sublimation
            • Change of a state of a substance from solid to gaseous without going through liquid state on heating
            • Mixture is heated gently, sublimate condenses on cooler sides of the funnel
          • Distillation
            • Fractional distillation: separate a mixture of miscible liquids using a fractionating column
              • Separates liquids in order of boiling points
                • Liquid with lowest boiling point distilled first
                • Liquid with highest boiling point distilled last
              • Chromatography
                • Separate and identify both coloured and colourless mixtures
                • Mixture must be dissolved in same solvent
                • Components travel at different rates over paper, depends on solubility in solvent. More soluble it is, the faster it moves
                • Rf value = Dist. Moved by substance/ Dist. Moved by solvent

Chemical Bonding

  • Dot and cross diagrams : Page 1
  • Ionic Bonding
    • Electrostatic force of attraction between two oppositely charged ions
    • Usually formed between non-metals and metals
    • High melting and boiling points
      • Due to the strong electrostatic forces of attraction between oppositely charged ions in the giant ionic crystalline lattice structure, a large amount of heat energy is needed to break these bonds
    • Ionic substances conduct electricity in aqueous and molten state but not in solid
      • In solid state, ions are held together by strong force of electrostatic attraction between oppositely charged ions
      • Can only vibrate around their own position
      • In liquid state, the strong forces of electrostatic attraction between oppositely charged ions are broken
      • Ions are free to move around
      • Electricity can be conducted by these mobile ions
    • Soluble in water
      • Water molecules can bond with both positive and negative ions, breaking up the lattice structure
    • Metallic Structures
      • Metallic structures consist of positive ions in a sea of delocalized valence electrons
      • Metallic ions and sea of delocalized valence electrons have strong force of electrostatic attraction between them
      • Can conduct electricity
        • Sea of delocalized valence electrons can move around freely and help to conduct electricity
      • Covalent Bonding
        • Two main types of Structure: Giant and Simple Molecular structures
        • Between two non-metals
        • Giant Molecular Structure
          • Covalently bonded: Millions of atoms joined by strong covalent bonds through the structure
          • Classic examples: Diamond, SiO2
          • High melting and boiling points\Hardness
            • Whole structure is held together by a network of strong covalent bonds between atoms
            • Large amount of energy is needed to break these bonds
          • Exceptions to normal GMS: Graphite
            • Layered structure, between layers of strong covalent bonds, there are weak Van Der Waal’s force
            • Van Der Waal’s force can be overcome easily and thus graphite is soft and easily broken
            • Can conduct electricity
              • There is a non-bonded valence electron on each carbon atom as only three covalent bonds are formed, thus these electrons can help to conduct electricity
            • Simple Molecular Structure
              • Strong covalent bonds between atoms within a molecule but weak intermolecular forces(Van Der Waal’s force or H-bonding) between molecules
                • Low Melting and Boiling point
                  • Due to weak intermolecular forces of attraction between molecules
                  • Little heat energy is needed to overcome these intermolecular forces of attraction
                • Non-conductor of electricity
                  • Neutral molecules that do not contain mobile ions or delocalized electrons
                • Exceptions of cases: Hydrogen Bonding
                  • N,O,F
                  • Hydrogen with N,O or F usually result in a covalently bonded atom with hydrogen bonding between molecules as well
                  • While they have Van Der Waal’s forces similar to normal Simple Molecular structure, they are also held together by hydrogen bonding and thus need more energy to overcome these forces

Ionic Equations

  • Ionic equations are chemical equations that have the non-participating ions removed
  • Key points
    • Do not remove solid/liquid/gas reactants/products
    • Remove aqueous reactants/products or parts of them that do not end up as a solid, liquid or gas in the final equation
    • Add charges to aqueous reactants/products that have had their counterparts removed
    • HCl (aq) +NaOH (aq) -> NaCl (aq) + H2O (l)
    • HCl and NaOH are aqueous. However, H and O become H20, a product that is a liquid, thus they are not removed but Cl and Na are.
    • Ionic Equation : H+(aq)  + OH (aq) -> H20 (l)

Acids, Bases and Salts

  • Acids
    • Substance that produces H+ ions as sole positive ion in water
    • Weak acids do not disassociate fully, strong acids disassociate fully
    • Turns blue litmus paper red
    • <7 on pH scale
    • Reactions with other substances
      • React with metals
        • Products are hydrogen gas and salt
      • React with carbonates
        • Products are carbon dioxide, salt and water
      • React with metal hydroxides/oxides (Neutralisation)
        • Products are salt and water
      • Bases
        • Substance that produces OH- ions as sole negative ion in water
        • Alkali = soluble base
        • Turns red litmus paper blue
        • >7 on pH scale
        • Reactions with other substances
          • React with acids (Neutralisation)
            • Products are salt and water
          • React with ammonium salts
            • Products are salt, water and ammonia gas
          • Alkali reacts with metal salt
            • Products are metal hydroxide and salt
          • Oxides
            • Metallic oxides
              • Basic Oxides
                • Reacts with acid to produce salt and water
                • CaO, MgO
              • Amphoteric Oxides
                • Reacts with both acids and bases
                • ZnO, Al2O3, PbO
              • Non-metallic oxides
                • Acidic Oxides
                  • Reacts with alkalis to produce salt and water
                  • NO2, CO2, SO2
                • Neutral Oxides
                  • Does not react with either alkali or water
                  • CO, NO
                • Solubility Rules
                  • Nitrates
                    • All are soluble
                  • Sulfates
                    • All soluble except Ba, Ca, Pb
                  • Chlorides
                    • All soluble except Ag, Pb
                  • Carbonates
                    • All insoluble except Group 1 Metal and ammonium carbonates
                  • Hydroxides and Oxides
                    • All insoluble except Group 1 Metal and some Group 2 Metals
                  • Reactivity of metals
                    • Most reactive : K, Na, Ca, Mg ,Al
                    • Most non-reactive : Gold, Platinum, Silver , Copper
                  • Preparation of Salts 


Qualitative Analysis

  • Test for cations:
    • NaOH
      • White precipitate produced
        • Insoluble in excess
          • Ca
        • Soluble in excess
          • Al, Pb, Zn
        • Green precipitate produced
          • Fe 2+  (insoluble in excess)
        • Reddish-brown precipitate produced
          • Fe 3+ (insoluble in excess)
        • Blue precipitate produced
          • Cu 2+ (insoluble in excess)
        • NH3
          • No precipitate
            • Ca
          • White precipitate produced
            • Soluble in excess
              • Zn
            • Insoluble in excess
              • Al, Pb
            • Fe 2+, Fe 3+ and Cu 2+ are similar to reactions in NaOH
          • Note: To identify between Al and Pb, react both with chloride ion. Al with chloride ion produces white ppt while Pb with chloride ion produces colourless solution
        • Test for anions
          • Carbonate
            • Add dilute acid: Carbon dioxide, water and salt is produced
          • Sulfate
            • Acidify with dilute nitric acid, add aqueous barium nitrate
            • White ppt formed
          • Chloride and Iodine
            • Acidify with dilute nitric acid , add aqueous silver nitrate/lead (II) nitrate)
            • Chloride : White ppt , Iodine : Yellow ppt
          • Nitrate
            • Add aqueous NaOh and aluminium foil/Devarda’s alloy and then warm gently
            • Effervescene of colourless and pungent gas which turns moist red litmus paper blue

 Test for gases

  • Hydrogen
    • Lighted splint (extinguished with pop sound)
  • Oxygen
    • Glowing splint (rekindles the splint)
  • Carbon Dioxide
    • Limewater (white ppt is formed)
  • Sulfur dioxide
    • Add acidified aqueous potassium dichromate (VI)
    • Effervescence of colourless and pungent gas which turn moist blue litmus paper red
    • Orange acidified potassium dichromate (VI) turns green
  • Chlorine
    • Observation: Effervescence of greenish yellow and pungent gas which turns moist blue litmus paper red and then bleaches it
  • Ammonia
    • Observation: Effervescence of colourless and pungent gas that turns moist red litmus paper blue

Chemical Calculation

  • 1 mol = 6.02 x 10^23
  • No of moles = Mass(in g)/Molar mass
    • Note: When facing diatomic molecules, need to multiply >.>
  • For finding Empirical Formula and Molecular Formula
  • Lithium forms a compound with composition 8.00% lithium, 36.8% sulfur and 55.2% oxygen. (a) Find the empirical formula of this compound. (b) Relative molecular mass of the compound is 174. Find the molecular formula of the compound.
Mass [in 100g]836.855.2
Molar mass73216
Mols8/7 = 1.1436.8/32 = 1.1555.2/16 = 3.45
Mole ratio113


Empirical formula = LiSO3

Molecular formula = n x empirical formula

n = 174 / [7 + 32+ 16×3]

n = 2     Therefore, molecular formula is (LiSO3) x 2 = Li2S2O6

  • Molar gas volume = 24 dm3
  • Molarity
    • Concentration in mol/dm3 x molar mass = concentration in mass/dm3

Chemical Periodicity

  • Metals reactivity increase down the group
    • Atomic size of metals increases, atoms can lose valence electron easily to form positive charge ions
  • Non-metals reactivity decreases down the group
  • Electronegativity increase àand up the group so F would be most electronegative
  • As Mr increases, Van Der Waal’s force gets stronger and stronger, thus covalent compounds further on the table would have higher boiling and melting points
  • Group I – Alkali Metals
    • Properties
      • Shiny, silvery, metallic solid (metal)
      • Soft
      • Melting point decreases down group
      • Li, Na and K have low density, can float in water
      • Reactivity increases down table
    • Group VII – Halogens
      • Appearances
ElementsChemical FormulaColourState
ChlorineCl2Greenish yellowGas
BromineBr2Reddish BrownLiquid


  • Properties
    • Exists as diatomic molecules
    • Simple molecular structure with weak Van Der Waals’ force existing between molecules
    • Reactivity decreases down table
    • Displacement reactions
      • More reactive halogen displaces less reactive halogen from compound
      • Cl2(g) + 2 NaBr (aq)  -> 2 NaCl (aq) + Br2  (l)
    • Group VIII – Noble Gases
      • Unreactive monoatomic gases
      • Very stable electronic structure
      • All have low melting and boiling point- Van Der Waal’s force




  • Gases can be used to act as inert atmosphere
    • Argon and helium
      • During welding, if aluminium is welded in air, hot metal will catch fire and burn in oxygen
      • Argon and helium provides a unreactive atmosphere to prevent this
    • Transition Metals
      • Strong hard metals with High Density
      • High Melting and Boiling point – Metallic bonds
      • Variable oxidation state
      • Forms coloured ion
        • Iron (II) – Pale green , Iron (III) – Pale yellow, Copper (II) – Blue
      • Use as catalysts
      • Exceptions
        • Scandium, Zinc(Most likely one to be tested) , Silver
        • Only one oxidation state : Zn2+
        • Form white compounds in solid states
        • Coloured compounds in aqueous states

Atmosphere and Environment

  • Air: 79% Nitrogen, 20% Oxygen, 1% Noble gas, 0.03% Carbon Dioxide, 0.5% Water Vapour
  • Air pollution: Chemicals in air with high enough concentrations that it that harms living organisms
    • Pollutants : Harmful substances found in environment
      • Carbon monoxide
        • Source: Forest fires, incomplete combustion of fuel in cars
        • Reacts with haemoglobin, forming stable carboxyhaemoglobin which prevents blood from transporting oxygen
        • Paralyzes brain activity, headaches, fatigue, impaired judgement
        • Colourless, tasteless, odourless
      • Sulfur dioxide
        • Source: Volcanic eruptions, burning of fossil fuels with sulfur as impurity
        • Poisonous choking gas: Irritates eyes, attacks lungs , causing breathing difficulties, leading to bronchitis
        • Forms acid rain
          • Sulfur dioxide dissolves in water, forming sulfurous acid, H2SO3
          • Sulfurous acid oxidizes to form sulphuric acid (sulphuric acid rain)
          • Corrodes metal and limestone structures
          • Poor health and stunted growth in fish
          • Absorption of needed nutrients by plants affected, replaced by toxic ions ,killing plants
        • Oxides of Nitrogen
          • Source: Car exhaust fumes (high temperature in engine causes nitrogen and oxygen to react) and lightning (through heat released by lightning)
          • Nitrogen Dioxide: red-brown toxic gas, unpleasant pungent odour
          • Causes eye irritation, damage lung tissues and blood vessels
          • Forms acid rain
          • 4NO2 + 2H20 + 02à 4HNO3
        • Lead
          • Lead accumulates in body, causing damage to the brain, liver, kidneys, central nervous system
          • Symptoms: Loss of appetite, vomiting, convulsions
        • Methane
          • Source: Bacterial decay of vegetation, fires, mining, decaying animal dung, rubbish in landfills
          • Colourless, odourless gas
          • Under strong sunlight, reacts with nitrogen dioxide forming photochemical smog
          • Greenhouse gas, causes global warming
        • Unburnt hydrocarbons
          • Source: Incomplete combustion of petrol in car engines
          • Component in photochemical smog
        • Photochemical smog
          • Mixture of pollutants: dust, nitrogen oxides, ozone, unreacted hydrocarbons, peroxyacyl nitrates (PAN)
          • Brownish haze, painful eyes, reduced visibility
          • Causes headache, eye , nose, throat irritation, impaired lung function, coughing and wheezing
        • Ozone
          • Combines with unburnt hydrocarbons to form PAN, which causes tearing of eyes
            • Dangerous of asthmatic patients
            • Damage rubber in car tires and fabrics
            • Damage plants
          • Ozone layer depletion
            • Ozone layer acts as giant sunscreen to protect Earth’s surface from harmful UV radiation
            • CFCs: aerosol, refrigerators, cleaning solvents
            • CFC molecules rise into upper atmosphere, forming chlorine atoms, chlorine atoms react with ozone molecules
            • Ozone layer is destructed, UV rays are let through, causes skin cancer, eye cateracts , severely damages plant growth
          • Catalytic converter
            • Contain catalysts: platinum and rhodium
            • Carbon monoxide reacts with nitrogen oxide to form nitrogen and carbon dioxide
            • Unburnt hydrocarbons oxidized to carbon dioxide and water
            • Found in cars mostly to reduce pollution caused
          • Ways to minimize acid rain
            • Treat soil with calcium carbonate/calcium oxide to neutralize excess acidity
            • Flue gas desulfurization
              • Sulfur dioxide is removed from flue/waste gases by reacting with aqueous suspension of calcium carbonate, forming calcium sulphite
                • CaCO3 + SO2 àCaSO3 + CO2
              • Calcium sulphite oxidized to form calcium sulphate

O Level Chemistry Tuition

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