CHEM REVISION
Unit 1: States of Matter
1.1 The Three States of Matter
Matter exists in three physical states: solid, liquid, and gas. These states differ in
particle arrangement, energy, movement, and forces of attraction.
Properties of Solids, Liquids, and Gases
Property Solid Liquid Gas
Shape Fixed Takes shape of No fixed shape
container
Volume Fixed Fixed Expands to fill
container
Particle Closely packed, Close but random Far apart, random
Arrangement regular pattern
Forces of Very strong Strong but weaker Very weak/negligible
Attraction than solids
Movement Particles vibrate in Particles slide past Particles move freely at
fixed positions each other high speeds
Compressibilit Not compressible Not compressible Highly compressible
y
Key Notes:
● Solids have a fixed shape and volume due to strong forces of attraction.
● Liquids have a fixed volume but no fixed shape. Particles move freely but
remain in contact.
● Gases have no fixed shape or volume. Particles move freely and can be
compressed easily.
1.2 The Kinetic Particle Theory
The kinetic particle theory explains the behavior of solids, liquids, and gases based on
particle energy and movement.
Main Ideas of the Kinetic Particle Theory
1. All matter is made of tiny particles.
2. Particles are in constant motion.
3. The amount of energy particles have determines their state.
Particle Movement in Different States
● Solids: Particles vibrate but stay in fixed positions.
● Liquids: Particles move past each other but remain close.
● Gases: Particles move randomly and rapidly in all directions.
Higher temperatures = more energy = faster movement!
1.3 Changes of State
Matter can change between states when energy is gained or lost.
Process Change Explanation
Melting Solid → Particles gain energy, vibrate more, and break free from
Liquid fixed positions.
Freezing Liquid → Particles lose energy, move slower, and form a fixed
Solid pattern.
Boiling Liquid → Particles gain energy, move faster, and escape as gas.
Gas
Condensati Gas → Particles lose energy, move closer, and form a liquid.
on Liquid
Sublimation Solid → Particles gain enough energy to skip the liquid phase
Gas (e.g., dry ice).
Key Points:
● Boiling happens at a fixed temperature throughout the liquid.
● Evaporation happens only at the surface and at any temperature.
● Condensation and freezing release energy, making particles move slower.
1.4 Effect of Temperature & Pressure on Gases
The behavior of gases is affected by temperature and pressure:
● Higher temperature → Particles move faster → Gas expands.
● Higher pressure → Particles move closer together → Gas compresses.
Example:
● A balloon expands when heated because gas particles move faster.
● A syringe compresses when pressure is applied, forcing gas particles closer.
1.5 Diffusion
Diffusion is the spreading of particles from a region of high concentration to low
concentration.
Factors Affecting Diffusion
1. Temperature: Higher temperatures = faster diffusion (more energy).
2. Particle Size: Lighter particles diffuse faster than heavier ones.
Example:
● Bromine gas in a jar spreads until evenly mixed.
● Potassium permanganate in water slowly turns the entire solution purple.
Unit 2: Atoms, Elements, and Compounds
2.1 Atomic Structure
Atoms: The Basic Unit of Matter
● Atoms are the smallest units of an element that retain the properties of that
element.
● Every atom consists of three subatomic particles:
1. Protons – Positive charge (+1), mass of 1 atomic unit, found in the
nucleus.
2. Neutrons – No charge (0), mass of 1 atomic unit, found in the nucleus.
3. Electrons – Negative charge (-1), negligible mass, move around the
nucleus in energy levels (shells).
Atomic Number and Mass Number
● Atomic Number (Z) = Number of protons in an atom.
● Mass Number (A) = Protons + Neutrons.
● Neutrons = Mass Number – Atomic Number.
Example:
● Carbon (C) has Atomic Number 6 → 6 protons, 6 electrons.
● Its Mass Number is 12 → Neutrons = 12 – 6 = 6.
Electron Arrangement (Configuration)
● Electrons orbit the nucleus in shells (energy levels).
● The first shell holds 2 electrons, the second and third shells hold 8
electrons.
● Electron configurations for first 20 elements:
○ Hydrogen (H) → 1
○ Oxygen (O) → 2,6
○ Sodium (Na) → 2,8,1
○ Calcium (Ca) → 2,8,8,2
2.2 Isotopes
Definition of Isotopes
● Isotopes are atoms of the same element with the same number of protons
but different numbers of neutrons.
● This means they have the same atomic number but different mass numbers.
Examples of Isotopes
Element Isotope Proton Neutron Mass
s s Number
Carbon Carbon-12 6 6 12
Carbon Carbon-14 6 8 14
Hydroge Hydrogen-1 1 0 1
n
Hydroge Deuterium 1 1 2
n (Hydrogen-2)
Hydroge Tritium (Hydrogen-3) 1 2 3
n
Key Points:
● Isotopes of the same element have the same chemical properties because
they have the same electron arrangement.
● Some isotopes are radioactive (e.g., Carbon-14 used in carbon dating).
Uses of Isotopes
1. Carbon-14 → Used for carbon dating (determining the age of ancient fossils).
2. Cobalt-60 → Used in cancer treatment (radiotherapy).
3. Uranium-235 → Used in nuclear reactors for energy production.
2.3 Elements, Compounds, and Mixtures
Elements
● Pure substances made of only one type of atom.
● Cannot be broken down into simpler substances.
● Found on the Periodic Table.
Examples:
● Oxygen (O₂), Hydrogen (H₂), Iron (Fe), Carbon (C).
Compounds
● Two or more elements chemically bonded together in a fixed ratio.
● Have different properties than the elements that make them.
● Cannot be separated physically (only by chemical methods).
Examples:
● Water (H₂O) – Made of hydrogen and oxygen, but has different properties.
● Carbon Dioxide (CO₂) – Made of carbon and oxygen, completely different from
both.
Mixtures
● Two or more substances physically combined (not chemically bonded).
● Can be separated by physical methods (e.g., filtration, distillation).
Examples:
● Air (a mixture of nitrogen, oxygen, CO₂, etc.).
● Seawater (a mixture of water, salt, minerals).
2.4 Differences Between Elements, Compounds, and Mixtures
Feature Element Compound Mixture
Compositio Only one type of Two or more elements Two or more substances
n atom chemically bonded physically mixed
Separation Cannot be Can only be separated by Can be separated by
broken down chemical methods physical methods
Properties Same as the Different from the original Retains properties of its
original element elements components
Example Oxygen (O₂), Water (H₂O), Carbon Air, Sand + Iron filings
Gold (Au) Dioxide (CO₂)
Unit 3: Stoichiometry
3.1 The Mole Concept
Definition of a Mole
● A mole is the amount of substance that contains 6.022 × 10²³ particles
(atoms, molecules, or ions).
● This number is called Avogadro’s constant.
Key Formulae
1. Moles = Mass (g) ÷ Molar Mass (Mr)n=MassMrn=MrMass
2. Number of Particles = Moles × Avogadro’s Number Particles=n×(6.022×1023)
3. Moles of Gas = Volume (dm³) ÷ Molar Gas Volume (24 dm³ at
RTP)n=Volume24
Example Calculations:
● Find the number of moles in 18g of water (H₂O).
○ Mr of H₂O = (2 × 1) + 16 = 18
○ Moles = 18 ÷ 18 = 1 mole
● Find the number of oxygen molecules in 2 moles of O₂.
○ Number of molecules = 2 × (6.022 × 10²³) = 1.204 × 10²⁴ molecules
3.2 Empirical and Molecular Formulas
Empirical Formula
● The simplest whole-number ratio of atoms in a compound.
Example Calculation:
A compound contains 40% sulfur and 60% oxygen. Find its empirical formula.
1. Convert mass to moles:
○ S: 40 ÷ 32 = 1.25
○ O: 60 ÷ 16 = 3.75
2. Divide by the smallest number:
○ S: 1.25 ÷ 1.25 = 1
○ O: 3.75 ÷ 1.25 = 3
3. Empirical formula = SO₃
Molecular Formula
● The actual number of atoms in a compound.
● Formula: Molecular Formula = (Empirical Formula) × Factor
● Factor = Molecular Mass ÷ Empirical Mass
Example Calculation:
● Empirical formula CH₂, Mr = 14
● Given molecular Mr = 28
● Factor = 28 ÷ 14 = 2
● Molecular formula = C₂H₄
3.3 Balancing Chemical Equations
● A balanced equation must follow the Law of Conservation of Mass.
● The number of atoms of each element must be the same on both sides.
Example:
Unbalanced:
H2+O2→H2O
Balanced:
2H2+O2→2H2O
3.4 Reacting Mass Calculations
Steps to Calculate Reacting Masses:
1. Write a balanced equation.
2. Find the moles of the given substance.
3. Use the mole ratio from the equation.
4. Find the mass of the required substance using: Mass = Moles × Mr.
Example Calculation:
How much iron(III) oxide (Fe₂O₃) is needed to produce 112g of iron (Fe)?
Step 1: Write the balanced equation:
Fe2O3+3C→2Fe+3CO
Step 2: Find moles of Fe:
Moles=11256=2 moles
Step 3: Use the mole ratio (Fe₂O₃ : Fe = 1:2)
● Fe₂O₃ required = 2 ÷ 2 = 1 mole
Step 4: Find mass of Fe₂O₃:
● Mr of Fe₂O₃ = (2 × 56) + (3 × 16) = 160
● Mass = 1 × 160 = 160g
Answer: 160g of Fe₂O₃ is required.
3.5 Limiting Reactants
● The limiting reactant is the substance that is completely used up first in a
reaction.
● The excess reactant is left over.
Example Calculation:
Given: 8g of hydrogen reacts with 32g of oxygen to form water. Identify the limiting
reactant.
1. Balanced equation:
2H2+O2→2H2O
2. Find moles:
○ H₂: 8 ÷ 2 = 4 moles
○ O₂: 32 ÷ 32 = 1 mole
3. Use mole ratio (H₂ : O₂ = 2:1)
○ Oxygen is limiting (needs 2 moles of H₂ but we have 4, so H₂ is in
excess).
Answer: Oxygen (O₂) is the limiting reactant.
3.6 Gas Volume Calculations
1 mole of any gas at room temperature and pressure (RTP) occupies 24 dm³.
Formula:
Gas Volume (dm³)=Moles×24
Example Calculation:
Find the volume of CO₂ produced when 2 moles of CaCO₃ decompose:
CaCO3→CaO+CO2
● Mole ratio: CaCO₃ : CO₂ = 1:1
● Moles of CO₂ = 2 moles
● Volume of CO₂ = 2 × 24 = 48 dm³
Answer: 48 dm³ of CO₂ is produced.
Unit 5: Chemical Energetics
5.1 Exothermic and Endothermic Reactions
Definition of Energy Changes in Reactions
● Chemical reactions involve energy changes because bonds are broken and
formed.
● Bond breaking requires energy (endothermic).
● Bond formation releases energy (exothermic).
Exothermic Reactions
● Release heat energy to the surroundings.
● Temperature of surroundings increases.
● More energy is released when new bonds form than is absorbed to break bonds.
● ΔH (enthalpy change) is negative (-ΔH).
Examples of Exothermic Reactions:
1. Combustion
(burning fuels, e.g., methane burning in oxygen)
CH4+2O2→CO2+2H2O+heat
2. Respiration (energy release from glucose)
C6H12O6+6O2→6CO2+6H2O+energy
Endothermic Reactions
● Absorb heat energy from the surroundings.
● Temperature of surroundings decreases.
● More energy is required to break bonds than is released when new bonds form.
● ΔH (enthalpy change) is positive (+ΔH).
Examples of Endothermic Reactions:
1. Photosynthesis
(plants absorb light energy)
6CO2+6H2O+light energy→C6H12O6+6O2
2. Boiling and Melting (phase changes absorb heat)
5.2 Energy Level Diagrams
● Show the energy changes in reactions.
● Exothermic Reactions: Products have lower energy than reactants.
● Endothermic Reactions: Products have higher energy than reactants.
Example of an Exothermic Energy Profile:
Reactants→Products+Energy ReleasedReactants→Products+Energy Released
● Energy goes down in the diagram.
● ΔH is negative (-ΔH).
Example of an Endothermic Energy Profile:
Reactants+Energy Absorbed→ProductsReactants+Energy Absorbed→Products
● Energy goes up in the diagram.
● ΔH is positive (+ΔH).
5.3 Activation Energy (Ea)
● The minimum energy required for a reaction to start.
● Catalysts lower activation energy, speeding up reactions.
Example:
● Combustion of methane requires a spark to start.
● Enzymes act as biological catalysts to speed up reactions in cells.
Unit 8: The Periodic Table
8.1 Arrangement of Elements in the Periodic Table
Structure of the Periodic Table
● The Periodic Table arranges elements in order of increasing atomic number
(number of protons).
● Rows = Periods → Show the number of electron shells an element has.
● Columns = Groups → Show the number of electrons in the outer shell.
Example:
● Sodium (Na) is in Period 3, Group 1 → It has 3 electron shells and 1 outer
electron.
● Chlorine (Cl) is in Period 3, Group 7 → It has 3 electron shells and 7 outer
electrons.
Metals vs. Non-Metals in the Periodic Table
● Metals are found on the left and center of the table.
● Non-metals are found on the right.
Property Metals Non-Metals
Conductivity Good conductors of Poor conductors
heat/electricity
Appearance Shiny Dull
Malleability Can be hammered into shape Brittle
Melting & Boiling High Low
Points
State at Room Mostly solids Many gases or
Temperature liquids
8.2 Group Trends in the Periodic Table
Group 1: The Alkali Metals (Lithium, Sodium, Potassium, etc.)
● 1 electron in the outer shell → Very reactive!
● React with water to form alkalis and hydrogen gas.
● Reactivity increases down the group because:
1. Outer electron is further from the nucleus.
2. Weaker attraction to the nucleus → Easier to lose.
Example Reaction:
2Na+2H2O→2NaOH+H2
(Sodium + Water → Sodium Hydroxide + Hydrogen)
Group 7: The Halogens (Fluorine, Chlorine, Bromine, Iodine, etc.)
● 7 electrons in the outer shell → Very reactive!
● React with metals to form salts (e.g., NaCl, KBr).
● Reactivity decreases down the group because:
1. Outer electrons are further from the nucleus.
2. Harder to attract another electron.
Example Displacement Reaction:
Cl2+2KBr→2KCl+Br2
(Chlorine displaces Bromine because it is more reactive.)
Group 8: The Noble Gases (Helium, Neon, Argon, etc.)
● Full outer electron shell → Very stable and unreactive.
● Used in balloons (Helium), neon lights (Neon), and welding (Argon).
8.3 The Transition Metals
1. Found in the middle block of the Periodic Table.
2. Properties:
High melting & boiling points
Form colored compounds (e.g., Copper sulfate is blue)
Act as catalysts in reactions
Example:
● Iron (Fe) is used in the Haber Process as a catalyst to make ammonia.
Group 8: The Noble Gases (Helium, Neon, Argon, Krypton, Xenon, Radon)
Properties of Noble Gases
● Full outer shell of electrons → Very stable and unreactive (inert).
● Exist as monatomic gases (single atoms, not molecules).
● Low boiling points (increase down the group).
● Density increases down the group.
Why Are Noble Gases Unreactive?
● Full outer electron shell → No need to gain, lose, or share electrons.
● No tendency to form bonds → Exist as single atoms (monatomic gases).
Key Trends in Group 8
● Boiling points increase down the group → More electrons = Stronger forces.
● Density increases down the group → Atoms get larger and heavier.
Unit 12: Experimental Techniques & Chemical Analysis
12.1 Separation Techniques
Mixtures can be separated using physical methods because they are not chemically
bonded.
Separation Used For Example
Method
Filtration Separating an insoluble solid from a liquid Sand and water
Crystallization Separating a soluble solid from a solution Salt from seawater
by evaporating the solvent
Simple Separating a solvent from a solution Purifying water
Distillation from saltwater
Fractional Separating liquids with different boiling Separating ethanol
Distillation points from water
Chromatograph Separating dyes in ink or food colorings Identifying food
y additives
Filtration
● Used when a solid does not dissolve in a liquid.
● The liquid passes through filter paper (filtrate), and the solid remains
(residue).
Example:
● Sand and water can be separated using filtration.
● The water (filtrate) passes through, while the sand (residue) stays on the
paper.
Crystallization
● Used to obtain a pure soluble solid from a solution.
● The solvent is evaporated, leaving behind solid crystals.
Example:
● Salt from seawater is obtained by crystallization.
Distillation (Simple & Fractional)
● Simple Distillation: Separates a liquid from a solution by boiling and
condensing it.
● Fractional Distillation: Separates two or more liquids with different boiling
points using a fractionating column.
Examples:
● Simple Distillation → Purifying water from saltwater.
● Fractional Distillation → Separating ethanol (78°C) from water (100°C).
Chromatography
● Used to separate substances in a mixture (e.g., ink, dyes, amino acids).
● Components move at different speeds, forming separate spots.
Steps in Paper Chromatography:
1. Draw a pencil baseline on filter paper.
2. Place a spot of the sample on the baseline.
3. Dip the paper in solvent (e.g., water or ethanol).
4. Solvent moves up, carrying dyes at different rates.
5. Calculate the Rf value for each dye:
Rf=Distance moved by substance/Distance moved by solvent
Uses:
● Identifying dyes in food.
● Checking the purity of drugs.
12.4 Purity & Melting/Boiling Points
● A pure substance has a sharp, fixed melting and boiling point.
● Impurities lower melting points and raise boiling points.
Example:
● Pure water boils at 100°C and melts at 0°C.
● Saltwater has a lower freezing point and higher boiling point than pure
water.
