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1.1 – 1.2: States of Matter

Overview

Matter exists in three primary states: solid, liquid, and gas. Each state has distinct
characteristics based on particle arrangement and energy.

Kinetic Particle Theory

This theory explains the properties of matter in terms of the movement of particles:

• Solids: Particles are tightly packed in a fixed arrangement. They vibrate but do
not move from their positions, giving solids a definite shape and volume.

• Liquids: Particles are close together but not in a fixed position. They can move
past one another, allowing liquids to flow and take the shape of their container while
maintaining a constant volume.

• Gases: Particles are far apart and move rapidly in all directions. Gases have
neither a fixed shape nor a fixed volume and will expand to fill their container.

Changes of State

• Melting: Solid to liquid. Particles gain energy and move more freely.

• Freezing: Liquid to solid. Particles lose energy and arrange into fixed positions.

• Boiling: Liquid to gas. Particles gain enough energy to overcome

intermolecular forces.

• Condensation: Gas to liquid. Particles lose energy and come closer together.

• Sublimation: Solid to gas without becoming liquid. Particles gain sufficient

energy to transition directly.

• Deposition: Gas to solid without becoming liquid. Particles lose energy

rapidly.
Diffusion

Diffusion is the movement of particles from an area of higher concentration to an area of

lower concentration. It occurs fastest in gases due to the large spaces between particles and
their high kinetic energy.

Visual Representation

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Example Questions

1. Multiple Choice: Which state of matter has particles that are closely packed in
a fixed position?

A) Solid

B) Liquid

C) Gas

D) Plasma

Answer: A) Solid

2. Short Answer: Explain why gases can be compressed more easily than solids
or liquids.

Answer: Gases have particles that are far apart with large spaces between them, allowing
them to be compressed more easily compared to solids and liquids where particles are
closely packed.

3. Diagram Labeling: Given a diagram showing particles in different

arrangements, identify which represents a solid, liquid, and gas.

Certainly! Let’s proceed with the next topic:

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2.1 – 2.7: Atoms, Elements, and Compounds

2.1: Elements, Compounds, and Mixtures

• Element: A pure substance consisting of only one type of atom. Examples

include oxygen (O₂), hydrogen (H₂), and iron (Fe).

• Compound: A substance formed when two or more elements chemically

combine in fixed proportions. For example, water (H₂O) is a compound of hydrogen and
oxygen.

• Mixture: A combination of two or more substances that are not chemically

bonded and can be separated by physical means. Air is a mixture of gases like nitrogen,
oxygen, and carbon dioxide.

2.2: Atomic Structure

• Protons: Positively charged particles found in the nucleus. Each has a relative
mass of 1.

• Neutrons: Neutral particles also located in the nucleus with a relative mass of
1.

• Electrons: Negatively charged particles orbiting the nucleus in energy levels or

shells. They have negligible mass.

• Atomic Number (Z): Number of protons in an atom’s nucleus.

• Mass Number (A): Total number of protons and neutrons in the nucleus.

2.3: Isotopes

• Isotopes: Atoms of the same element with the same number of protons but
different numbers of neutrons. For example, carbon-12 and carbon-14 are isotopes of
carbon.

2.4: Ions and Ionic Bonds

• Ion: An atom or group of atoms that has lost or gained electrons, resulting in
a charge.
 • Cation: Positively charged ion (loss of electrons).

• Anion: Negatively charged ion (gain of electrons).

• Ionic Bond: The electrostatic attraction between oppositely charged ions.

Common in compounds formed between metals and non-metals, such as sodium chloride
(NaCl).

2.5: Covalent Bonds

• Covalent Bond: A chemical bond formed when two atoms share one or more
pairs of electrons. Common in compounds formed between non-metals, such as water (H₂O)
and carbon dioxide (CO₂).

2.6: Giant Covalent Structures

• Diamond: Each carbon atom forms four covalent bonds in a tetrahedral

structure, resulting in a very hard substance.

• Graphite: Each carbon atom forms three covalent bonds, creating layers that
can slide over each other, making graphite a good lubricant.

2.7: Metallic Bonding

• Metallic Bond: The attraction between a lattice of positive metal ions and a
sea of delocalized electrons. This structure explains properties like electrical conductivity
and malleability in metals.

Visual Representation

Image Source: Save My Exams

Example Questions

1. Multiple Choice: Which subatomic particle has a relative mass of 1 and no

charge?

A) Proton

B) Neutron
C) Electron

D) Positron

Answer: B) Neutron

2. Short Answer: Define an isotope and give an example.

Answer: Isotopes are atoms of the same element with the same number of protons but
different numbers of neutrons. Example: Carbon-12 and Carbon-14.

3. Diagram Labeling: Given a diagram of an atom, label the protons, neutrons,

and electrons.

Great! Here’s the next topic:

8.1 – 8.5: The Periodic Table

8.1: Periodic Table Structure

• The Periodic Table is an arrangement of elements in order of increasing

atomic number.

• Periods (horizontal rows): Show the number of electron shells.

• Groups (vertical columns): Contain elements with the same number of outer
electrons, hence similar chemical properties.

8.2: Group Properties

Group I – The Alkali Metals

• Very reactive, soft metals.

• Reactivity increases down the group.

 • React with water to form hydrogen gas and an alkaline solution:

2Na + 2H₂O → 2NaOH + H₂

Group VII – The Halogens

• Non-metals with colored vapors.

• Reactivity decreases down the group.

• Form salts with metals and displace less reactive halogens in solutions.

Cl₂ + 2KBr → 2KCl + Br₂

Group 0 – The Noble Gases

• Unreactive (inert) due to full outer electron shells.

• Used in light bulbs (argon), balloons (helium), and neon signs (neon).

8.3: Transition Elements

• Found in the center block (Groups 3–12).

• Have variable oxidation states.

• Form colored compounds.

• Often act as catalysts (e.g. Fe in the Haber process).

8.4: Metallic and Non-metallic Behavior

• Metals are found on the left and center; non-metals on the right.

• Metallic character increases down a group but decreases across a period.

8.5: Trends Across a Period

• Atomic radius decreases across a period (due to increased nuclear charge

pulling electrons closer).

• Ionization energy increases.

• Electronegativity increases.
 • Metallic to non-metallic character changes across a period.

Visual Representation

Image Source: Save My Exams

Example Questions

1. Multiple Choice: Which of the following elements is likely to form a colored

compound?

A) Sodium

B) Calcium

C) Iron

D) Oxygen

Answer: C) Iron

2. Short Answer: Why do noble gases not react easily?

Answer: Noble gases have full outer electron shells, making them stable and unreactive.

3. Displacement Reaction: Write the balanced equation for the reaction

between chlorine and potassium iodide.

Answer: Cl₂ + 2KI → 2KCl + I₂

Excellent! Here is the next detailed section:

12.1 – 12.4: Experimental Techniques and Chemical Analysis

(Including Collection of Gases and Drying Agents)

12.1: Measurement Techniques

• Measuring Mass: Use a digital/electronic balance for solids.

• Measuring Volume:

• Burette (accurate, variable volumes)

• Pipette (fixed volume, usually 25.0 cm³)

• Measuring Cylinder (less accurate)

• Measuring Temperature: Thermometer or digital probe.

12.2: Collection of Gases

Gases can be collected using three main methods depending on solubility and density:

Gas Collection Method Description Suitable for

Downward delivery Gas flows down into container due to being denser Chlorine,
hydrogen chloride

Upward delivery Gas escapes upward into inverted container Ammonia

Over water (displacement) Bubbles into a container filled with water Oxygen, hydrogen,
CO₂ (insoluble gases)

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Drying Gases

Different drying agents are used depending on the gas:

Drying Agent Used for Not Suitable for

Concentrated H₂SO₄ Drying acidic and neutral gases Not for alkaline gases like
ammonia

Quicklime (CaO) Drying ammonia Not for acidic gases

Anhydrous CaCl₂ General dryingNot for ammonia

12.3: Filtration, Crystallisation, and Distillation

• Filtration: Separates an insoluble solid from a liquid (e.g., sand from water).

• Crystallisation: Used to obtain a pure solid from a solution (e.g., copper(II)

sulfate crystals from solution).

• Simple Distillation: Separates a solvent from a solution (e.g., water from salt
solution).

• Fractional Distillation: Separates a mixture of liquids based on boiling points

(e.g., ethanol from water).

12.4: Paper Chromatography

Used to separate and identify substances in a mixture (e.g., dyes in ink).

• Rf value:

R_f = \frac{\text{Distance travelled by substance}}{\text{Distance travelled by solvent}}

 • A pure substance shows only one spot, whereas a mixture shows multiple
spots.

Visual Representations

Gas Collection Methods:

Paper Chromatography:

Example Questions

1. Multiple Choice: Which gas is best collected using upward delivery?

A) Oxygen

B) Ammonia

C) Carbon dioxide

D) Chlorine

Answer: B) Ammonia

2. Short Answer: Name a suitable drying agent for hydrogen gas.

Answer: Anhydrous calcium chloride (CaCl₂)

3. Practical Question: A student wants to separate a mixture of ethanol and

water. Which method should be used?

Answer: Fractional distillation

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Great! Here’s the detailed explanation for:

3.1 – 3.3: Stoichiometry

3.1: Chemical Formulae and Equations

• Molecular formula: Shows the actual number of atoms of each element (e.g.,
H₂O).

• Empirical formula: Simplest whole-number ratio of atoms (e.g., CH₂ for C₂H₄).

• Balanced chemical equations: Show conservation of mass (equal atoms on

both sides).

Example:

\text{CH}_4 + 2\text{O}_2 \rightarrow \text{CO}_2 + 2\text{H}_2\text{O}

3.2: The Mole Concept

• Mole: The amount of substance containing 6.022 \times 10^{23} particles

(Avogadro’s number).

• Molar mass (M): The mass of 1 mole of a substance (in g/mol).

\text{Moles} = \frac{\text{Mass}}{\text{Molar Mass (M)}}

\text{Mass} = \text{Moles} \times \text{Molar Mass}

Example:

How many moles in 36 g of water (H₂O)?

Molar mass = 2(1) + 16 = 18 g/mol

\frac{36}{18} = 2 \text{ moles}

3.3: Reacting Masses and Volume Calculations

Reacting Mass Calculations:

1. Write a balanced equation.

2. Calculate the number of moles of a known substance.

3. Use mole ratios to find moles of the unknown.

4. Convert moles to mass or volume.

Example:

What mass of carbon dioxide is formed when 12 g of carbon burns in excess oxygen?

\text{C} + \text{O}_2 \rightarrow \text{CO}_2

Moles of C = \frac{12}{12} = 1 mole

Mole ratio C : CO₂ = 1:1

Mass of CO₂ = 1 mole × 44 g/mol = 44 g

Molar Gas Volume

At room temperature and pressure (r.t.p), 1 mole of gas occupies 24 dm³.

\text{Volume (dm³)} = \text{Moles} \times 24

Example:

How many dm³ of hydrogen are formed from 2 moles?

2 \times 24 = 48 \text{ dm³}

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Visual Representation

Empirical Formula Calculation

Example Questions

1. Calculation: Find the empirical formula of a compound with 12 g of carbon

and 4 g of hydrogen.

Moles of C = \frac{12}{12} = 1

Moles of H = \frac{4}{1} = 4

Simplest ratio: 1:4 → CH₄

2. Short Answer: Define the mole.

Answer: The mole is the amount of substance containing 6.022 \times 10^{23} particles.

3. Gas Volume: Calculate the volume of 0.5 moles of oxygen gas at r.t.p.

Answer: 0.5 \times 24 = 12 \text{ dm³}

Perfect! Let’s cover:

7.1 – 7.2: Acids, Bases, and Salts

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7.1: Properties of Acids and Bases

Acids:

• Release H⁺ ions in solution.

• pH less than 7.

• Turn blue litmus red.

• React with metals, bases, and carbonates.

General Reactions:

• Acid + Metal → Salt + Hydrogen

(e.g. HCl + Zn → ZnCl₂ + H₂)

• Acid + Base → Salt + Water

(Neutralisation)

(e.g. HCl + NaOH → NaCl + H₂O)

• Acid + Carbonate → Salt + Water + CO₂

(e.g. H₂SO₄ + CaCO₃ → CaSO₄ + H₂O + CO₂)

Bases:

• Accept H⁺ ions or release OH⁻ ions.

• pH greater than 7.

• Turn red litmus blue.

• Insoluble bases are called metal oxides/hydroxides.

• Alkalis are soluble bases (e.g., NaOH, KOH).

7.2: Preparation of Salts

Types of Salts

• Soluble salts: Made using acids + metals, bases, or carbonates.

• Insoluble salts: Made using precipitation.

Methods of Salt Preparation

Salt Type Method Example

Soluble (from acid + base) Excess base method CuSO₄ from CuO + H₂SO₄

Soluble (from acid + alkali) Titration NaCl from NaOH + HCl

Insoluble Salt Precipitation PbSO₄ from Pb(NO₃)₂ + H₂SO₄

Titration Steps

1. Add a known volume of acid to a conical flask.

2. Add a few drops of indicator (e.g., phenolphthalein).

3. Fill a burette with alkali.

4. Add alkali dropwise until the indicator changes color.

5. Record volume used and repeat for accurate results.

6. Repeat without indicator to obtain pure salt.

Indicators and pH

Indicator Acid Color Alkali Color

Litmus Red Blue

Methyl Orange Red Yellow

Phenolphthalein Colorless Pink

Solubility Rules (Common Salts)

Salt Type Solubility

All sodium, potassium, ammonium salts Soluble

All nitrates Soluble

Most chlorides Soluble (except Ag⁺, Pb²⁺)

Most sulfates Soluble (except Ba²⁺, Pb²⁺, Ca²⁺)

Most carbonates Insoluble (except Na⁺, K⁺, NH₄⁺)

Visual Representation

Titration Apparatus:

Example Questions

1. Short Answer: What is the name of the salt formed when hydrochloric acid
reacts with sodium hydroxide?
Answer: Sodium chloride (NaCl)

2. Multiple Choice: Which of the following salts is insoluble in water?

A) NaCl

B) KNO₃

C) PbSO₄

D) NH₄Cl

Answer: C) PbSO₄

3. Reaction Writing: Write a balanced chemical equation for the reaction

between copper(II) oxide and sulfuric acid.

Answer: CuO + H₂SO₄ → CuSO₄ + H₂O

Awesome! Let’s go through:

10.1 – 10.2: Chemistry of the Environment

10.1: Air Composition and Pollutants

Composition of Clean, Dry Air (by volume):

• Nitrogen (N₂): 78%

• Oxygen (O₂): 21%

• Argon & other noble gases: 0.9%

• Carbon dioxide (CO₂): 0.04%

Air Pollutants and Their Sources:

Pollutant Source Effects Removal/Prevention

Carbon monoxide (CO) Incomplete combustion of fuels Toxic – binds to

haemoglobin Catalytic converters in cars

Sulfur dioxide (SO₂) Burning fossil fuels with sulfur Causes acid rain Use low-
sulfur fuels; scrubbers

Oxides of nitrogen (NO, NO₂) Car engines Acid rain; respiratory issues Catalytic
converters

Lead compounds Leaded petrol (now banned in many countries) Brain damage Use
unleaded petrol

Acid Rain

• Cause: Emission of SO₂ and NOₓ gases that dissolve in water vapour.

\text{SO}_2 + H_2O \rightarrow H_2SO_3

\text{2NO}_2 + H_2O \rightarrow HNO_2 + HNO_3

• Effects:

• Damages buildings (especially limestone).

• Harms aquatic life and soil fertility.

• Damages forests.

10.2: Greenhouse Effect and Global Warming

Greenhouse Gases:

• Carbon dioxide (CO₂)

• Methane (CH₄)

• Water vapour

Cause of Global Warming:

• Greenhouse gases trap infrared radiation, raising Earth’s temperature.

Sources:

• CO₂: Burning fossil fuels, deforestation.

• CH₄: Cattle farming, landfills, rice paddies.

Consequences:

• Rising sea levels.

• Climate change and extreme weather.

• Melting ice caps.

Solutions:

• Use renewable energy.

• Reduce deforestation.

• Improve energy efficiency.

Visual Representation

Air Pollution and Effects Diagram:

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Example Questions

1. Short Answer: What are two harmful effects of sulfur dioxide in the
atmosphere?

Answer: Causes acid rain; irritates the respiratory system.

2. Multiple Choice: Which gas is responsible for the formation of acid rain?

A) Nitrogen

B) Carbon dioxide

C) Sulfur dioxide

D) Oxygen

Answer: C) Sulfur dioxide

3. Explain: How does carbon monoxide affect the human body?

Answer: It binds to haemoglobin in red blood cells, reducing the oxygen-carrying capacity of
the blood, which can be fatal.

Great! Here’s the final topic:

12.2: Experimental Techniques – Separation and Purification

This section focuses on separating mixtures and purifying substances based on physical
properties such as solubility, boiling point, and particle size.
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Main Separation Techniques

Technique Purpose Example

Filtration Separate insoluble solid from a liquid Sand from water

Crystallisation Obtain a pure solid from a solution Copper(II) sulfate crystals from solution

Simple Distillation Separate a solvent from a solution Water from salt water

Fractional Distillation Separate liquids with different boiling pointsEthanol and water

Chromatography Separate components of a mixture (usually dyes) Ink or food dyes

Decanting Pouring off liquid from a settled solidOil and water (rough method)

Centrifugation Rapid separation of solid from liquid via spinning Blood components

Purification Techniques

• Crystallisation:

1. Heat solution to evaporate some water.

2. Leave to cool and allow crystals to form.

3. Filter and dry the crystals.

• Distillation:

• In simple distillation, the liquid is boiled, evaporated, and then condensed.

• In fractional distillation, a fractionating column is used to separate liquids

with different boiling points.

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Uses of Fractional Distillation

• Petroleum Refining: Crude oil → Petrol, diesel, etc.

• Ethanol Purification: Fermented sugar solution → Ethanol

Paper Chromatography Recap

• Stationary phase: Paper

• Mobile phase: Solvent (e.g. water or ethanol)

• Rf value:

R_f = \frac{\text{Distance moved by substance}}{\text{Distance moved by solvent front}}

• Pure substances = one spot, mixtures = multiple spots

Visual Representation

Separation Techniques Summary Diagram:

Example Questions

1. Short Answer: Which technique would you use to separate ethanol from
water?

Answer: Fractional distillation

2. Multiple Choice: Which method is used to obtain pure salt from a salt
solution?
A) Filtration

B) Distillation

C) Crystallisation

D) Chromatography

Answer: C) Crystallisation

3. Explain: Why is the solvent line drawn in pencil during paper

chromatography?

Answer: Because ink might dissolve in the solvent and interfere with results.