This file compiles all four uploaded Chapter 2 parts into one complete Ordinary Level HTML revision file. It also adds the missing syllabus points identified during checking.
Stable compounds, octet rule, ionic bonding and covalent bonding.
Chapter 2 Part 1: Chemical Bonding
1. Learning Objectives
Explain why atoms form chemical bonds.
Describe the stable electronic arrangement of noble gases.
State the octet rule in simple OL language.
Distinguish between ionic bonding and covalent bonding.
Explain how ions form by loss or gain of electrons.
Draw simple electron-transfer diagrams for ionic compounds such as sodium chloride and magnesium oxide.
Draw simple dot-and-cross diagrams for covalent molecules such as hydrogen, chlorine, water, methane and carbon dioxide.
Compare key properties of ionic and covalent substances.
Big Idea: Bonding is all about atoms becoming more stable by changing or sharing their outer electrons.
2. Why Do Atoms Form Bonds?
Most atoms are not completely stable on their own. They react with other atoms so that their outer shell becomes full, or more stable. This usually means achieving the same outer electron arrangement as a noble gas.
Atoms bond because stable outer shells are lower in energy.
Noble Gas Stability
Noble gases are very unreactive because their outer electron shells are already full. Other atoms often react in ways that give them a noble gas arrangement.
Noble Gas
Electron Configuration
Why Stable?
Helium
2
First shell is full.
Neon
2,8
Outer shell has 8 electrons.
Argon
2,8,8
Outer shell has 8 electrons.
Octet rule: atoms tend to gain, lose or share electrons to get 8 electrons in their outer shell.
Examiner Trap: Do not say every atom always needs 8 electrons. Hydrogen and helium are stable with 2 electrons in the first shell.
3. Ionic Bonding
Ionic bonding happens when electrons are transferred from one atom to another. This forms positive and negative ions. The oppositely charged ions attract each other strongly.
Metal atom loses electrons → positive ion
Non-metal atom gains electrons → negative ion
Example: Sodium Chloride, NaCl
Sodium has one electron in its outer shell. Chlorine has seven electrons in its outer shell. Sodium transfers one electron to chlorine.
Sodium loses one electron. Chlorine gains one electron. Na⁺ and Cl⁻ are formed.
Formation of ions:
Na → Na⁺ + e⁻
Cl + e⁻ → Cl⁻
The ionic bond is the strong electrostatic attraction between Na⁺ and Cl⁻.
Example: Magnesium Oxide, MgO
Magnesium has two outer electrons. Oxygen has six outer electrons. Magnesium transfers two electrons to oxygen.
Electron transfer:
Mg → Mg²⁺ + 2e⁻
O + 2e⁻ → O²⁻
Formula formed: MgO
Examiner Tip: In ionic bonding questions, always mention electron transfer, formation of oppositely charged ions, and strong attraction.
Common Mistake: Saying sodium chloride is made of molecules. It is not. Sodium chloride forms a giant ionic lattice.
4. Covalent Bonding
Covalent bonding happens when atoms share pairs of electrons. Covalent bonding usually occurs between non-metal atoms.
Covalent bond = shared pair of electrons
Single, Double and Triple Covalent Bonds
Bond Type
Meaning
Example
Single bond
One shared pair of electrons
H₂, Cl₂, CH₄
Double bond
Two shared pairs of electrons
O₂, CO₂
Triple bond
Three shared pairs of electrons
N₂
A covalent bond is a shared pair of electrons. More shared pairs make double or triple bonds.
Common Covalent Molecules at OL
Molecule
Formula
Bonding Description
Hydrogen
H₂
Two hydrogen atoms share one pair of electrons.
Chlorine
Cl₂
Two chlorine atoms share one pair of electrons.
Water
H₂O
Oxygen shares one pair with each hydrogen atom.
Methane
CH₄
Carbon shares one pair with each of four hydrogen atoms.
Carbon dioxide
CO₂
Carbon forms two double bonds, one with each oxygen atom.
Examiner Trap: In covalent bonding, electrons are shared, not transferred.
5. Ionic vs Covalent Bonding
Feature
Ionic Bonding
Covalent Bonding
Usually between
Metal + non-metal
Non-metal + non-metal
Electron behaviour
Electrons transferred
Electrons shared
Particles formed
Positive and negative ions
Molecules
Structure
Giant ionic lattice
Simple molecules or giant covalent structures
Melting/boiling points
Usually high
Usually low for simple molecules
Conductivity
Conduct when molten or dissolved in water
Usually do not conduct electricity
Examiner Tip: When comparing bonding, use the words transferred for ionic and shared for covalent. These are the magic words.
Syllabus Completion Additions: Chemical Compounds, Ionic Bonding and Covalent Bonding
This section adds the remaining Ordinary Level syllabus points that were missing or too brief in the original Chapter 2 parts, so that this compiled file matches the OL syllabus more closely.
1. Uses of Helium and Argon
Helium and argon are noble gases. They are very unreactive because they already have stable outer electron arrangements.
Gas
Why it is useful
Examples of use
Helium
Does not react easily and is very light.
Balloons, weather balloons, cooling systems, shielding gas in special welding.
Argon
Very unreactive, so it protects hot metals and filaments from reacting.
Electric light bulbs, welding, protective atmosphere in industry.
Examiner Tip: If asked why helium or argon is suitable, link the answer to chemical unreactivity.
2. Using the Octet Rule to Predict Formulas
At Ordinary Level, you should be able to use the octet rule and common ion charges to predict formulas for simple compounds.
Group 1 metal → usually +1 ion
Group 2 metal → usually +2 ion
Group 17 non-metal → usually −1 ion
Group 16 non-metal → usually −2 ion
Type of compound
Examples
How the formula is decided
Binary ionic compounds
NaCl, MgO, CaCl2, Al2O3
Total positive charge must balance total negative charge.
Hydroxides
NaOH, KOH, Ca(OH)2, Mg(OH)2, Al(OH)3
Hydroxide ion is OH−.
Carbonates
Na2CO3, CaCO3, MgCO3
Carbonate ion is CO32−.
Worked Example: Predict the formula of calcium chloride.
Calcium forms Ca2+ and chlorine forms Cl−.
Two chloride ions are needed to balance one calcium ion. Answer: CaCl2
Worked Example: Predict the formula of aluminium hydroxide.
Aluminium forms Al3+ and hydroxide is OH−.
Three hydroxide ions are needed. Answer: Al(OH)3
3. Minute Size of Ions and Molecules
Ions and molecules are extremely small particles. They are far too small to be seen directly with the naked eye.
An ion is a charged particle formed when an atom or group of atoms gains or loses electrons. A molecule is a neutral particle made when atoms are joined by covalent bonds.
Examiner Trap: A compound such as sodium chloride is made of ions, not molecules. A substance such as water is made of molecules.
4. Sodium Chloride Crystal Structure
Sodium chloride does not exist as separate NaCl molecules. It forms a giant ionic lattice. In this lattice, each sodium ion is surrounded by chloride ions and each chloride ion is surrounded by sodium ions.
The structure repeats in three dimensions to form a giant lattice, not separate NaCl molecules.
5. Ionic Materials in Everyday Life
Ionic material
Use
Why useful
Sodium chloride
Table salt / food preservation
Common stable ionic compound used in daily life.
Calcium carbonate
Building materials / antacid tablets
Hard, stable ionic compound.
Magnesium hydroxide
Antacid
Neutralises excess stomach acid.
Salt tablets / sports salts
Replacing salts lost in sweat
Provides important ions for the body.
Mandatory Experiment 2.1: Tests for Anions in Aqueous Solution
The Ordinary Level syllabus specifically requires the tests for chloride, carbonate, nitrate, and sulfate ions in aqueous solution. Learn the reagent, the observation, and the conclusion.
Anion
Test / Reagent
Positive Observation
Conclusion
Chloride, Cl−
Acidify with dilute nitric acid, then add silver nitrate solution.
White precipitate forms.
Chloride ion present.
Carbonate, CO32−
Add dilute acid, then test gas with limewater.
Fizzing / effervescence; limewater turns milky.
Carbon dioxide released, so carbonate present.
Nitrate, NO3−
Add sodium hydroxide solution and aluminium foil, then warm gently.
Ammonia gas produced; damp red litmus turns blue.
Nitrate ion present.
Sulfate, SO42−
Acidify with dilute hydrochloric acid, then add barium chloride solution.
White precipitate forms.
Sulfate ion present.
Carbonate test equation:
Ca(OH)2 + CO2 → CaCO3 + H2O
This is why limewater turns milky.
Nitrate test gas:
The gas formed is ammonia.
Damp red litmus paper turning blue confirms ammonia.
Good Practical Technique
Use clean test tubes to avoid contamination.
Add acids carefully and wear eye protection.
For chloride and sulfate tests, acidify first to remove interference from other ions.
For the nitrate test, warm gently rather than boiling strongly.
Examiner Tip: In practical questions, marks are often awarded separately for the reagent, the observation, and the conclusion. Learn all three.
6. Worked Examples
Example 1: Formation of chloride ion
Chlorine has electron configuration 2,8,7. It gains one electron to become Cl⁻ with configuration 2,8,8. This gives chlorine a stable outer shell.
Example 2: Formula of calcium chloride
Calcium forms Ca²⁺. Chlorine forms Cl⁻. Two chloride ions are needed to balance one calcium ion.
Formula = CaCl₂.
Example 3: Bonding in water
Oxygen has six outer electrons and needs two more to complete its outer shell. It shares one pair of electrons with each hydrogen atom, forming two covalent bonds.
7. Interactive Simulators / Virtual Labs
Simulator 1: PhET Build a Molecule
Use for: covalent molecules such as H₂O, CO₂, CH₄ and O₂.
Student task: Build at least four simple molecules and write their formulas and names.
Exam link: Helps students connect molecular formula with bonding and structure.
Suggested search: PhET Build a Molecule simulation.
Simulator 2: PhET Molecule Shapes Basics
Use for: introducing linear, planar and tetrahedral shapes before Part 2.
Student task: Add atoms around a central atom and observe how the shape changes.
Exam link: Prepares students for simple molecular geometry questions.
Suggested search: PhET Molecule Shapes Basics.
Simulator 3: Ionic Bonding Interactive
Use for: electron transfer and ion formation.
Student task: Model sodium chloride and magnesium oxide, then identify which atom loses and which atom gains electrons.
Exam link: Reinforces ionic bonding diagrams and ion charges.
Suggested search: ionic bonding simulation electron transfer.
8. Examiner Secrets, Mistakes and Traps
Examiner Secret: Bonding questions are often marking-keyword questions. Use exact words: transfer, share, ions, opposite charges, strong attraction.
Common Mistake: Drawing ionic bonding as if electrons are shared. Ionic bonding involves transfer, not sharing.
Common Mistake: Forgetting the charge on ions. Sodium becomes Na⁺, chlorine becomes Cl⁻, magnesium becomes Mg²⁺, oxygen becomes O²⁻.
Examiner Trap: A covalent bond is not an attraction between ions. It is a shared pair of electrons.
Examiner Trap: Noble gases are stable because their outer shell is full, not because they have no electrons.
9. Exam Practice Questions
Q1. Explain why noble gases are generally unreactive. [2 marks]
Q2. Sodium reacts with chlorine to form sodium chloride. Explain the bonding in sodium chloride in terms of electron transfer. [5 marks]
Q3. Draw or describe the bonding in a molecule of water. [4 marks]
Q4. Compare ionic and covalent bonding using the words transferred and shared. [4 marks]
Q5. Explain why magnesium oxide has the formula MgO. [3 marks]
MCQs with Explanations
1. Ionic bonding involves: A. Sharing electrons B. Transferring electrons C. Losing protons D. Sharing neutrons
Answer: B. Ionic bonding occurs when electrons are transferred from one atom to another.
2. A covalent bond is: A. A shared pair of electrons B. A lost proton C. A neutron pair D. A positive ion only
Answer: A. Covalent bonds are shared pairs of electrons.
3. Which pair is most likely to form an ionic compound? A. Carbon and hydrogen B. Sodium and chlorine C. Oxygen and oxygen D. Nitrogen and hydrogen
Answer: B. Sodium is a metal and chlorine is a non-metal, so electron transfer occurs.
4. What ion does oxygen usually form? A. O⁺ B. O²⁺ C. O⁻ D. O²⁻
Answer: D. Oxygen gains two electrons to form O²⁻.
10. Mark Scheme
Q1. Noble gases have full outer electron shells [1]; therefore they are stable and do not easily gain, lose or share electrons [1].
Q2. Sodium has one outer electron [1]; chlorine has seven outer electrons [1]; sodium transfers one electron to chlorine [1]; Na⁺ and Cl⁻ ions form [1]; oppositely charged ions attract strongly [1].
Q3. Oxygen has six outer electrons [1]; each hydrogen has one electron [1]; oxygen shares one pair with each hydrogen [1]; two covalent bonds are formed / H₂O molecule formed [1].
Q4. Ionic bonding involves transfer of electrons [1]; covalent bonding involves sharing electrons [1]; ionic bonding forms ions [1]; covalent bonding forms molecules [1].
Q5. Magnesium loses two electrons to form Mg²⁺ [1]; oxygen gains two electrons to form O²⁻ [1]; charges balance in a 1:1 ratio, so formula is MgO [1].
11. Last-Minute Revision Sheet
Atoms bond to become more stable.
Noble gases are stable because their outer shells are full.
The octet rule means atoms tend to get 8 electrons in the outer shell.
Hydrogen and helium are stable with 2 electrons in the first shell.
Ionic bonding = electron transfer.
Metals form positive ions by losing electrons.
Non-metals form negative ions by gaining electrons.
I can explain how positive and negative ions form.
I can explain ionic bonding using electron transfer.
I can explain covalent bonding using shared pairs of electrons.
I can compare ionic and covalent bonding.
I can work out simple ionic formulas from charges.
I can answer OL exam questions using correct bonding keywords.
Part 2
Electronegativity, polarity, intermolecular forces and physical properties.
Chapter 2 Part 2: What This Part Covers
In Part 1, you studied why atoms bond and how ionic and covalent bonds form. In Part 2, we focus on why some covalent molecules have slightly charged ends and why some simple substances have surprisingly high boiling points.
Learning Objectives
Explain electronegativity as the attraction an atom has for shared electrons in a covalent bond.
Describe the trend in electronegativity across a period and down a group.
Use electronegativity difference to decide whether a bond is non-polar covalent, polar covalent or mainly ionic.
Explain the difference between polar and non-polar covalent bonds.
Describe intermolecular forces: van der Waals forces, dipole-dipole attractions and hydrogen bonding.
Link intermolecular forces to boiling point, melting point and solubility.
Examiner Tip: In bonding questions, always check whether the question asks about a bond inside a molecule or a force between molecules. They are not the same thing.
1. Electronegativity
Electronegativity is the ability of an atom in a molecule to attract the shared pair of electrons in a covalent bond.
Electronegativity = attraction for shared electrons
Electronegativity Trends
Fluorine is the most electronegative element. Metals usually have low electronegativity; non-metals usually have higher electronegativity.
Direction
Trend
Reason at OL level
Across a period
Electronegativity increases
Atoms attract bonding electrons more strongly as nuclear charge increases.
Down a group
Electronegativity decreases
Outer electrons are further from the nucleus, so attraction is weaker.
Examiner Trap: Do not write that electronegativity means an atom attracts all electrons. It means attraction for the shared pair in a covalent bond.
2. Using Electronegativity to Predict Bond Type
The bigger the difference in electronegativity between two bonded atoms, the more unequal the sharing of electrons becomes.
Electronegativity Difference
Bond Type
Example
Explanation
Very small / zero
Non-polar covalent
Cl₂, H₂, O₂
Electrons are shared equally.
Moderate
Polar covalent
HCl, H₂O, NH₃
Electrons are shared unequally.
Large
Mainly ionic
NaCl, MgO
Electron transfer is more likely.
In HCl, chlorine attracts the shared electrons more strongly, so chlorine becomes slightly negative and hydrogen becomes slightly positive.
Worked Example: Is the H–Cl bond polar or non-polar? Chlorine is more electronegative than hydrogen, so the shared electrons are pulled closer to chlorine. The bond is polar covalent.
Common Mistake: Saying polar covalent bonds involve complete electron transfer. That describes ionic bonding. Polar covalent bonding still involves sharing, but the sharing is unequal.
3. Polar and Non-Polar Molecules
A molecule may contain polar bonds, but the whole molecule may still be non-polar if the bond polarities cancel because of the molecule's shape.
Molecule
Bond Polarity
Shape Effect
Overall Molecule
HCl
Polar H–Cl bond
Only one bond, so no cancellation
Polar
CO₂
Polar C=O bonds
Linear and symmetrical, so polarities cancel
Non-polar
H₂O
Polar O–H bonds
Bent shape, so polarities do not cancel
Polar
CH₄
C–H bonds have very small polarity
Tetrahedral and symmetrical
Non-polar
Shape matters. Symmetrical molecules can be non-polar even when they contain polar bonds.
Examiner Tip: For OL, know simple examples: HCl and H₂O are polar; CO₂ and CH₄ are non-polar.
4. Intermolecular Forces
Intermolecular forces are forces of attraction between molecules. They are much weaker than covalent bonds inside molecules, but they strongly affect physical properties.
Van der Waals Forces
Weak attractions between all molecules. They become stronger as molecules get bigger.
Dipole-Dipole Forces
Attractions between permanent positive and negative ends of polar molecules.
Hydrogen Bonding
A strong type of intermolecular attraction involving H bonded to N, O or F.
Bonding Inside vs Forces Between
Feature
Bond Inside a Molecule
Intermolecular Force Between Molecules
Example
Covalent bond in H₂O: O–H
Hydrogen bonding between H₂O molecules
Strength
Strong
Usually weaker
Broken during boiling?
No, water molecules remain H₂O
Yes, attractions between molecules are overcome
Hydrogen bonding explains why water has a higher boiling point than expected for such a small molecule.
Common Mistake: Saying boiling water breaks the O–H covalent bonds. Boiling only overcomes intermolecular forces between water molecules.
5. Effects on Physical Properties
The stronger the intermolecular forces, the more energy is needed to separate the molecules.
Property
Effect of Stronger Intermolecular Forces
Example
Boiling point
Higher boiling point
Water has a high boiling point due to hydrogen bonding.
Melting point
Usually higher melting point
Substances with stronger attractions are harder to separate.
Solubility
Polar substances often dissolve better in polar solvents
Many ionic and polar substances dissolve in water.
Volatility
Lower volatility
Liquids with stronger forces evaporate less easily.
Solubility Rule: Like Dissolves Like
Polar solvent: Water dissolves many polar and ionic substances because water molecules are polar.
Non-polar solvent: Non-polar solvents are better for dissolving non-polar substances such as oils.
Examiner Trap: Do not simply write “water dissolves everything.” It does not. Water is good at dissolving many ionic and polar substances.
Syllabus Completion Additions: Polarity, Solubility and Everyday Materials
This section completes the remaining practical and applied points from the OL syllabus section on electronegativity, polarity, and intermolecular forces.
1. Polarity Test for Liquids Using a Charged Plastic Rod
A simple classroom test for polarity is to bring a charged plastic rod close to a thin stream of liquid.
Liquid
Expected result
Meaning
Water
Stream bends towards the charged rod.
Water is polar.
Ethanol
May bend, but usually less strongly than water.
Ethanol is polar.
Hexane / paraffin-like liquids / oils
Little or no bending.
These are non-polar or weakly polar.
Examiner Trap: The charged rod does not prove ionic bonding. It is used here to show whether a covalent liquid is polar or non-polar.
2. Testing Solubility in Different Solvents
A useful practical comparison is to test substances in water and in a non-polar solvent such as hexane.
Substance
In water (polar solvent)
In hexane / non-polar solvent
What this shows
Sodium chloride
Dissolves
Does not dissolve
Ionic / polar substances usually dissolve in polar solvents.
Sugar
Dissolves
Does not dissolve
Polar molecules often dissolve in water.
Iodine
Very slight solubility
Dissolves more easily
Non-polar substances dissolve better in non-polar solvents.
Cooking oil
Does not dissolve
Mixes with non-polar solvents
Non-polar materials do not mix well with water.
Like dissolves like: polar dissolves polar, non-polar dissolves non-polar.
3. Everyday Examples of Polar and Non-Polar Materials
Polar materials
Water
Ethanol
Ammonia solution
Sugar solution
Non-polar materials
Cooking oil
Wax
Petrol / hydrocarbons
Grease
Why this matters
Oil and water do not mix.
Grease is removed better by detergents.
Petrol is non-polar, so it dissolves many hydrocarbon substances.
Student task: Build simple molecules and observe how electronegativity and shape affect overall polarity.
Compare a non-polar molecule with a polar molecule.
Move electronegativity values and observe the bond dipole arrow.
Explain why symmetrical molecules may be non-polar.
Exam link: Helps with questions asking why HCl is polar but CO₂ is non-polar.
Simulator 2: Intermolecular Forces and States
Suggested source: PhET States of Matter simulation or similar molecular motion simulator.
Student task: Heat and cool particles and observe that changes of state involve separating particles, not changing the identity of the molecules.
Increase temperature and describe particle motion.
Explain why stronger attractions need more energy to overcome.
Connect the observation to boiling point.
Simulator 3: Solubility Investigation
Suggested source: ChemCollective or LabXchange solubility/polarity virtual activities.
Student task: Compare dissolving of polar and non-polar substances in different solvents.
Key conclusion: Like dissolves like.
7. Worked Examples
Example 1: Identify bond polarity
Is the Cl–Cl bond polar?
Both atoms are chlorine, so they attract the shared electrons equally. The bond is non-polar covalent.
Example 2: Explain high boiling point
Water has a higher boiling point than expected because water molecules form hydrogen bonds with each other. Extra energy is needed to overcome these intermolecular forces.
Example 3: Solubility
Why does sodium chloride dissolve in water?
Water is polar. The partial charges on water molecules attract the positive sodium ions and negative chloride ions, helping them separate from the crystal.
8. Examiner Secrets, Mistakes and Traps
Examiner Secret: When explaining polarity, mention both electronegativity difference and unequal sharing of electrons.
Examiner Secret: When explaining boiling point, use the phrase “energy is needed to overcome intermolecular forces.”
Common Mistake: Calling hydrogen bonding a covalent bond. Hydrogen bonding is an intermolecular force, not a covalent bond.
Common Mistake: Saying CO₂ is polar just because it has polar bonds. CO₂ is linear and symmetrical, so the bond polarities cancel.
Examiner Trap: Dipole-dipole attractions happen between polar molecules, not between ions in an ionic crystal.
9. Exam Practice Questions
Q1. Define electronegativity. [2 marks]
Q2. Describe the trend in electronegativity across a period in the periodic table. [2 marks]
Q3. Explain why the H–Cl bond is polar covalent. [3 marks]
Q4. Explain why CO₂ is a non-polar molecule even though its C=O bonds are polar. [4 marks]
Q5. Name three types of intermolecular forces. [3 marks]
Q6. Explain why water has a relatively high boiling point. [4 marks]
MCQs with Explanations
1. Electronegativity is the ability of an atom to attract: A. neutrons B. shared electrons C. protons D. ions only
Answer: B. Electronegativity means attraction for the shared pair of electrons in a covalent bond.
2. Which molecule is polar? A. CO₂ B. CH₄ C. HCl D. Cl₂
Answer: C. HCl has an unequal sharing of electrons and one slightly positive end and one slightly negative end.
3. Which is usually the strongest intermolecular force listed below? A. Hydrogen bonding B. Van der Waals forces C. Metallic bonding D. Ionic bonding
Answer: A. Among the intermolecular forces listed, hydrogen bonding is usually strongest. Metallic and ionic bonding are not intermolecular forces.
10. Mark Scheme
Q1. Electronegativity is the ability of an atom [1] to attract the shared pair of electrons in a covalent bond [1].
Q2. Electronegativity increases across a period [1] because atoms attract bonding electrons more strongly / nuclear attraction increases [1].
Q3. Chlorine is more electronegative than hydrogen [1]; the shared electrons are attracted more towards chlorine [1]; this gives H a slight positive charge and Cl a slight negative charge [1].
Q4. Each C=O bond is polar [1]; CO₂ is linear [1]; the molecule is symmetrical [1]; the bond polarities cancel, so the molecule is non-polar [1].
Q5. Van der Waals forces [1]; dipole-dipole attractions [1]; hydrogen bonding [1].
Q6. Water molecules are polar [1]; hydrogen bonding occurs between water molecules [1]; these forces are relatively strong [1]; more energy is needed to overcome them, giving water a high boiling point [1].
11. Last-Minute Revision Sheet
Electronegativity means attraction for shared electrons in a covalent bond.
Electronegativity increases across a period and decreases down a group.
Non-polar covalent bond = equal sharing of electrons.
Polar covalent bond = unequal sharing of electrons.
Polar molecules have a slightly positive end and a slightly negative end.
CO₂ is non-polar because it is linear and symmetrical.
H₂O is polar because it is bent and has polar O–H bonds.
Intermolecular forces act between molecules.
Hydrogen bonding is a strong intermolecular force involving H bonded to N, O or F.
Stronger intermolecular forces usually mean a higher boiling point.
Like dissolves like: polar substances often dissolve in polar solvents.
12. Self-Assessment Checklist
I can define electronegativity.
I can describe the trend in electronegativity across a period and down a group.
I can identify polar and non-polar covalent bonds.
I can explain why HCl is polar.
I can explain why CO₂ is non-polar.
I can name the three main intermolecular forces at OL.
I can explain hydrogen bonding in water.
I can link intermolecular forces to boiling point and solubility.
Part 3
Shapes of molecules and simple OL geometry.
1. Learning Objectives
Explain that electron pairs around a central atom repel each other.
Use simple VSEPR ideas to predict the shapes of common molecules.
Describe linear, trigonal planar and tetrahedral molecular shapes.
State the bond angles for linear, trigonal planar and tetrahedral molecules.
Recognise examples: BeCl₂, BF₃ and CH₄.
Use molecular shape to help explain some properties of simple covalent molecules.
OL Focus: At Ordinary Level, keep the explanation simple: electron pairs spread out as far apart as possible because they repel each other.
2. VSEPR Theory - The Big Idea
VSEPR stands for Valence Shell Electron Pair Repulsion. It is a model used to predict the shape of molecules.
Electron pairs repel and arrange themselves as far apart as possible.
The shape of a molecule depends mainly on the number of electron pairs around the central atom. In OL Chemistry, the three main shapes you must know are:
Number of bonding regions around central atom
Shape
Bond Angle
Example
2
Linear
180°
BeCl₂
3
Trigonal planar
120°
BF₃
4
Tetrahedral
109.5°
CH₄
Common Mistake: Students often draw molecules flat on paper and then assume all molecules are flat. CH₄ is not flat; it is three-dimensional tetrahedral.
3. Linear Molecules
A molecule is linear when there are two bonding regions around the central atom. These two bonding regions move as far apart as possible, giving a bond angle of 180°.
Two bonding regions around the central atom produce a straight-line molecule.
Worked Example: BeCl₂ has beryllium in the centre and two chlorine atoms bonded to it. Two bonding regions repel equally, so the molecule is linear with a 180° bond angle.
Examiner Trap: Do not call BeCl₂ “bent”. With two bonding regions and no lone pair treatment at OL, it is linear.
4. Trigonal Planar Molecules
A molecule is trigonal planar when there are three bonding regions around the central atom. The atoms lie in one flat plane and the bond angles are 120°.
Three bonding regions spread out into a flat triangular arrangement.
Worked Example: BF₃ has boron in the centre with three fluorine atoms bonded to it. The three bonding regions spread out evenly, forming a trigonal planar shape with 120° angles.
Common Mistake: “Trigonal planar” means flat triangular, not pyramidal.
5. Tetrahedral Molecules
A molecule is tetrahedral when there are four bonding regions around the central atom. The bonds point towards the corners of a tetrahedron. The bond angle is about 109.5°.
Methane, CH₄, is the classic tetrahedral molecule at OL.
Worked Example: In CH₄, carbon forms four single covalent bonds with four hydrogen atoms. Four bonding regions repel into a tetrahedral shape with bond angles of about 109.5°.
Examiner Trap: CH₄ is not square planar. Four bonds around carbon form a tetrahedral 3D shape.
6. Interactive Simulators and Virtual Tasks
Molecular Shapes Simulator
Suggested simulator: PhET - Molecule Shapes
Student Task: Build molecules with 2, 3 and 4 bonding regions. Record the shape and bond angle each time.
Exam Link: This directly supports questions asking students to identify linear, trigonal planar and tetrahedral shapes.
Build-a-Molecule Simulator
Suggested simulator: PhET - Build a Molecule
Student Task: Build simple covalent molecules such as CH₄ and compare formula, structure and shape.
Exam Link: Helps students connect molecular formula with displayed structure.
3D Molecule Viewer Task
Suggested simulator: Any classroom-approved 3D molecule viewer.
Student Task: Rotate CH₄ and observe why it cannot be represented perfectly by a flat drawing.
Exam Link: Prevents the common mistake of drawing tetrahedral molecules as flat squares.
7. Examiner Secrets, Mistakes and Traps
Examiner Secret: When asked for shape, give both the shape and the bond angle if possible. This often earns full marks faster.
Examiner Tip: Learn the three OL examples as a set: BeCl₂ = linear, BF₃ = trigonal planar, CH₄ = tetrahedral.
Common Mistake: Writing “triangular” instead of “trigonal planar”. Use the correct chemistry term.
Common Mistake: Giving 90° for CH₄. The tetrahedral angle is 109.5°, not 90°.
Examiner Trap: A displayed formula may look flat, but the real molecule may be 3D.
8. Worked Examples
Example 1: State the shape and bond angle of BeCl₂. Answer: Linear, 180°.
Example 2: BF₃ has three bonding regions around boron. Predict its shape. Answer: Trigonal planar. The bond angle is 120°.
Example 3: Explain why methane is tetrahedral. Answer: Carbon has four bonding regions around it. These repel and spread out as far apart as possible, giving a tetrahedral shape.
9. Exam Practice Questions
Q1. What does VSEPR theory help chemists to predict? [2 marks]
Q2. State the shape and bond angle of BeCl₂. [2 marks]
Q3. BF₃ is described as trigonal planar. Explain what this means. [3 marks]
Q4. Draw a simple diagram of methane, CH₄, and state its bond angle. [3 marks]
Q5. Explain why electron pairs around a central atom do not stay close together. [2 marks]
MCQs with Explanations
1. What is the shape of BeCl₂? A. Tetrahedral B. Linear C. Trigonal planar D. Pyramidal
Answer: B. BeCl₂ has two bonding regions around the central atom, giving a linear shape.
2. What is the bond angle in BF₃? A. 90° B. 109.5° C. 120° D. 180°
Answer: C. BF₃ is trigonal planar, so the bond angle is 120°.
3. Which molecule is tetrahedral? A. BeCl₂ B. BF₃ C. CH₄ D. CO₂ only
Answer: C. Methane, CH₄, has four bonding regions around carbon and is tetrahedral.
10. Mark Scheme
Q1. It predicts molecular shape [1] based on repulsion between electron pairs around the central atom [1].
Q2. Shape = linear [1]; bond angle = 180° [1].
Q3. Three bonding regions around the central atom [1]; all atoms lie in one flat plane [1]; bond angles are 120° [1].
Q4. Correct central carbon with four hydrogens [1]; tetrahedral arrangement shown or stated [1]; bond angle 109.5° [1].
Q5. Electron pairs have negative charge [1]; they repel each other and move as far apart as possible [1].
11. Last-Minute Revision Sheet
VSEPR means electron pairs repel each other.
Molecular shapes depend on bonding regions around the central atom.
2 bonding regions = linear = 180°.
3 bonding regions = trigonal planar = 120°.
4 bonding regions = tetrahedral = 109.5°.
BeCl₂ is linear.
BF₃ is trigonal planar.
CH₄ is tetrahedral.
Displayed formulas can make 3D molecules look flat, so be careful.
12. Self-Assessment Checklist
I can explain the basic idea of VSEPR theory.
I can identify a linear molecule and state its bond angle.
I can identify a trigonal planar molecule and state its bond angle.
I can identify a tetrahedral molecule and state its bond angle.
I can remember BeCl₂, BF₃ and CH₄ as key OL examples.
I can avoid confusing flat drawings with real 3D molecular shapes.
Part 4
Final review, summary and exam practice.
1. Learning Objectives
Compare ionic, covalent and metallic bonding.
Connect bonding type with physical properties.
Explain melting point, boiling point and conductivity using bonding.
Review simple molecular shapes at OL level.
Answer mixed exam-style questions on bonding and molecular geometry.
2. Bonding Big Picture
Chemical bonding explains why atoms join together and why substances have different properties.
Bonding controls structure, and structure controls properties.
First identify the atoms involved, then choose the bonding type.
3. Ionic Bonding Final Review
Ionic bonding happens when electrons are transferred from one atom to another, forming oppositely charged ions that attract.
Feature
Ionic Compounds
Particles
Positive and negative ions.
Structure
Giant ionic lattice.
Melting/boiling point
High, because strong attractions need lots of energy to break.
Conductivity
Conducts when molten or dissolved, not when solid.
Solubility
Many ionic compounds dissolve in water.
Examiner Tip: Ionic solids do not conduct because ions are fixed in place. Molten or dissolved ionic compounds conduct because ions can move.
4. Covalent Bonding Final Review
Covalent bonding happens when non-metal atoms share pairs of electrons.
Type
Example
Properties
Simple molecular covalent
Water, methane, carbon dioxide
Usually low melting/boiling point and poor electrical conductivity.
Giant covalent
Diamond, graphite, silicon dioxide
Very high melting point because many strong covalent bonds must be broken.
Examiner Trap: In simple molecular substances, melting usually breaks intermolecular forces, not the covalent bonds inside molecules.
5. Metallic Bonding Final Review
Metallic bonding is the attraction between positive metal ions and delocalised electrons.
Property
Explanation
Conduct electricity
Delocalised electrons can move and carry charge.
Malleable
Layers of metal ions can slide while bonding remains.
High melting point
Strong attraction between metal ions and electrons.
Shiny
Delocalised electrons interact with light.
Common Mistake: Saying metals conduct because ions move. In solid metals, electrons move, not metal ions.
6. Bonding and Properties Summary
Substance Type
Melting Point
Electrical Conductivity
Reason
Ionic solid
High
No
Ions fixed in lattice.
Molten ionic compound
High
Yes
Ions are free to move.
Simple covalent molecule
Low
No
No mobile charged particles.
Metal
Usually high
Yes
Delocalised electrons move.
Graphite
Very high
Yes
Delocalised electrons between carbon layers.
Diamond
Very high
No
No free electrons; all electrons used in bonds.
7. Molecular Geometry Final Review
Molecular shape depends on electron pairs around the central atom. Electron pairs repel and spread out as far as possible.
Molecule
Shape
Approximate Bond Angle
Reason
CH₄
Tetrahedral
109.5°
Four bonding pairs around carbon.
NH₃
Pyramidal
About 107°
Three bonding pairs and one lone pair.
H₂O
V-shaped / bent
About 104.5°
Two bonding pairs and two lone pairs.
CO₂
Linear
180°
Two regions of electron density around carbon.
Lone pairs repel more strongly than bonding pairs, reducing bond angles.
8. Interactive Simulators and Virtual Labs
Simulator 1: PhET – Molecule Shapes
Use for: Exploring molecular geometry and lone pairs.
Example 1: Sodium chloride has a high melting point because strong electrostatic attractions between oppositely charged ions require a lot of energy to overcome.
Example 2: Molten sodium chloride conducts electricity because ions are free to move and carry charge.
Example 3: Water is V-shaped because oxygen has two bonding pairs and two lone pairs of electrons.
Example 4: Metals conduct electricity because delocalised electrons are free to move through the structure.
10. Examiner Secrets, Mistakes and Traps
Examiner Secret: Property questions need a structure explanation, not just a property statement.
Examiner Tip: For conductivity, always identify the mobile charged particle: ions or electrons.
Common Mistake: Saying covalent molecules conduct because they contain electrons. Electrons in covalent bonds are not free to move.
Common Mistake: Saying ionic solids conduct electricity. Solid ionic compounds do not conduct because ions are fixed.
Examiner Trap: Graphite conducts, diamond does not, even though both are made of carbon.
11. Final Chapter 2 Exam Practice
Q1. Define ionic bonding. [2 marks]
Q2. Explain why sodium chloride has a high melting point. [3 marks]
Q3. Explain why molten sodium chloride conducts electricity but solid sodium chloride does not. [4 marks]
Q4. Define covalent bonding. [2 marks]
Q5. Explain why simple covalent substances often have low boiling points. [3 marks]
Q7. Give the shape and bond angle of carbon dioxide. [2 marks]
Q8. Explain why water is V-shaped. [3 marks]
MCQs with Explanations
1. Ionic bonding involves: A. sharing electrons B. transfer of electrons C. delocalised electrons only D. no electrons
Answer: B. Ionic bonding involves electron transfer and attraction between ions.
2. Metals conduct electricity because they contain: A. mobile ions B. delocalised electrons C. water molecules D. covalent molecules
Answer: B. Delocalised electrons carry charge.
3. CO₂ is: A. V-shaped B. pyramidal C. tetrahedral D. linear
Answer: D. Carbon dioxide is linear with a 180° bond angle.
4. Diamond does not conduct electricity because: A. it has no carbon B. all electrons are used in bonds C. it is a liquid D. it has free ions
Answer: B. Diamond has no delocalised electrons.
12. Last-Minute Revision Sheet
Ionic bonding = electron transfer.
Covalent bonding = sharing electron pairs.
Metallic bonding = attraction between positive metal ions and delocalised electrons.
Ionic compounds have high melting points.
Ionic compounds conduct when molten or dissolved, not when solid.
Simple covalent substances usually have low melting and boiling points.
Metals conduct because delocalised electrons move.
Diamond is giant covalent and does not conduct.
Graphite conducts because it has delocalised electrons.
CH₄ is tetrahedral, NH₃ pyramidal, H₂O V-shaped, CO₂ linear.
Lone pairs reduce bond angles because they repel more strongly.
13. Self-Assessment Checklist
I can compare ionic, covalent and metallic bonding.
I can explain melting point using bonding and structure.
I can explain electrical conductivity using mobile charged particles.
I can describe simple molecular shapes.
I can compare diamond and graphite.
I can answer mixed Chapter 2 exam questions confidently.
14. Mark Scheme
Q1. Transfer of electrons [1]; attraction between oppositely charged ions [1].
Q2. Giant ionic lattice [1]; strong electrostatic attractions between ions [1]; lots of energy needed to overcome attractions [1].
Q3. Solid ions are fixed in place [1] so cannot carry charge [1]; molten ions are free to move [1] and carry charge [1].
Q4. Sharing [1] of pairs of electrons [1].
Q5. Simple molecules have weak intermolecular forces [1]; little energy needed to overcome them [1]; so low boiling points [1].
Q6. Metals have delocalised electrons [1] that move and carry charge [1].
Q7. Linear [1]; 180° [1].
Q8. Oxygen has two bonding pairs and two lone pairs [1]; lone pairs repel more strongly [1]; giving a V-shaped/bent molecule [1].
Chapter 2 Syllabus Completion Check
This compiled version now includes the core Ordinary Level Chapter 2 content:
2.1 Chemical Compounds
2.2 Ionic Bonding
2.3 Covalent Bonding
2.4 Electronegativity
2.5 Shapes of Molecules and Intermolecular Forces
Mandatory Experiment 2.1: tests for chloride, carbonate, nitrate and sulfate ions
Applied points: helium and argon uses, ionic materials in daily life, polarity testing, and solvent-solubility testing
Use this file as your complete Chapter 2 master version. It compiles Parts 1–4 and includes the missing Ordinary Level syllabus points that were added during quality checking.