🧪 AS Level Chemistry – Chapter 3: Chemical
Bonding (9701)
🗓️ Date: 15 July 2025
🔹 3.1 Electronegativity and Bonding
🔹 What is Electronegativity?
Electronegativity is the ability of an atom to attract the shared electrons in a covalent bond towards
itself. - It is a relative scale, developed by Linus Pauling. - More electronegative atoms pull electrons closer
to themselves.
🔹 What affects electronegativity?
1. Nuclear charge: More protons = stronger attraction.
2. Atomic radius: Smaller atoms = stronger pull on electrons.
3. Electron shielding: More inner shells = weaker attraction of outer electrons.
🔹 Trends in Electronegativity:
Direction
Trend
Reason
Across a Period
Increases →
More protons, smaller atomic size
Down a Group
Decreases ↓
More shielding and larger atomic radius
🔹 Bond Type Based on Electronegativity Difference:
• 0 to ~0.4: Non-polar covalent
• ~0.4 to 1.7: Polar covalent
• >1.7: Ionic bond
📘 Note: Pauling values are usually provided in exams.
🔹 3.2 Ionic Bonding
⚡ What is Ionic Bonding?
• Ionic bonding is the electrostatic force of attraction between positively and negatively charged
ions.
• It usually forms between metals (lose electrons) and non-metals (gain electrons).
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🧪 Examples:
• NaCl: Na gives 1 e⁻ to Cl → Na⁺ and Cl⁻
• MgO: Mg gives 2 e⁻ to O → Mg²⁺ and O²⁻
• CaF₂: Ca²⁺ and 2F⁻
🔹 Properties of Ionic Compounds:
• High melting & boiling point: Strong ionic lattice.
• Conducts electricity: Only when molten or aqueous (ions are free to move).
• Soluble in water: Because water is polar and stabilises ions.
🔹 3.3 Metallic Bonding
🔩 What is Metallic Bonding?
• Electrostatic attraction between positive metal ions and a sea of delocalised electrons.
• Electrons are free to move = good conductors.
🧪 Example: Copper (Cu)
🔹 Properties:
• Good conductors: Free electrons carry charge.
• High melting points: Strong attractions.
• Malleable & ductile: Layers of atoms can slide.
🔁 Electrons are mobile = metals are flexible yet strong.
🔹 3.4 Covalent and Coordinate Bonding
🔗 What is a Covalent Bond?
• A shared pair of electrons between two atoms.
• Attraction is between shared electrons and both nuclei.
🧪 Covalent Molecules:
Molecule
Description
H₂
Single covalent bond (1 shared pair)
O₂
Double bond (2 shared pairs)
N₂
Triple bond (3 shared pairs)
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Molecule
Description
Cl₂
Single bond between two Cl atoms
HCl
Polar covalent (different ENs)
CO₂
2 double bonds (O=C=O), linear
NH₃
3 bonds + 1 lone pair (pyramidal)
CH₄
4 bonds, tetrahedral, sp³
C₂H₆
Single bonds only
C₂H₄
Double bond (contains π bond)
🔸 Octet Rule:
• Atoms want 8 electrons in their outer shell for stability.
• Exception: Period 3 elements like P, S can hold more than 8 electrons.
🔸 Expanded Octet (Period 3 Elements):
Molecule
Explanation
SO₂
S has 10 electrons (double bonds)
PCl₅
5 bonds = 10 electrons
SF₆
6 bonds = 12 electrons
➕ Coordinate (Dative Covalent) Bond:
• One atom donates both electrons in a shared pair.
• Represented with an arrow (→).
🧪 Examples:
• NH₄⁺: NH₃ donates lone pair to H⁺
• Al₂Cl₆: Cl donates lone pair to Al atom
🔹 Orbital Overlap: σ and π Bonds
🔸 Sigma (σ) Bonds:
• Head-on overlap of orbitals.
• Strongest type of covalent bond.
• Found in all covalent bonds (single/double/triple).
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🔸 Pi (π) Bonds:
• Sideways overlap of unhybridised p-orbitals.
• Only in double and triple bonds.
Molecule
Bond Type
H₂
1σ
C₂H₆
All σ
C₂H₄
1σ+1π
N₂
1σ+2π
HCN
1σ+2π
🌐 Hybridisation:
• Orbitals mix to form new hybrid orbitals for bonding.
Hybrid Type
Shape
Bond Angle
Example
sp
Linear
180°
BeCl₂
sp²
Trigonal planar
120°
BF₃
sp³
Tetrahedral
109.5°
CH₄
🔹 3.5 Shapes of Molecules (VSEPR Theory)
🧠 VSEPR = Valence Shell Electron Pair Repulsion
Electron pairs (bonding or lone) repel each other and arrange to minimise repulsion.
🔢 How to Predict Shape:
1. Count bonding and lone pairs on central atom.
2. Apply geometry:
Molecule
Shape
Bond Angle
Reason
BF₃
Trigonal planar
120°
3 bonding pairs
CO₂
Linear
180°
2 bonding pairs
CH₄
Tetrahedral
109.5°
4 bonding pairs
NH₃
Pyramidal
107°
3 bonding + 1 lone pair
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Molecule
Shape
Bond Angle
Reason
H₂O
Bent (non-linear)
104.5°
2 bonding + 2 lone pairs
PF₅
Trigonal bipyramidal
120° & 90°
5 bonding pairs
SF₆
Octahedral
90°
6 bonding pairs
📌 Lone pairs repel more than bonding pairs → smaller bond angles.
🔹 3.6 Intermolecular Forces & Bond Properties
1️⃣ Hydrogen Bonding (H-bonding):
Occurs when H is bonded to a highly electronegative atom (N, O, or F). - H-F, H-O, H-N
🧪 Examples:
• H₂O – between O and H of different molecules
• NH₃ – N and H
🔹 Properties of Water Due to H-bonding:
• High BP/MP: More energy needed to break H-bonds.
• High surface tension: Strong cohesive forces.
• Ice is less dense than water: Open lattice structure held by H-bonds.
2️⃣ Bond Polarity and Dipoles:
• Caused by difference in electronegativity.
• More EN atom pulls electron pair closer → partial charges (δ⁺, δ⁻).
• Dipole = molecule with partial positive and negative sides.
🧪 Examples:
• HCl → Polar covalent (Cl is more EN)
• CH₄ → Non-polar (symmetrical)
3️⃣ Van der Waals’ Forces (All IMFs except actual bonds):
Types:
1. id-id (London dispersion): Temporary dipoles in non-polar molecules. Weakest.
2. pd-pd: Between permanent dipoles. Stronger.
3. Hydrogen bonding: Special type of pd-pd when H is bonded to N, O, or F.
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Strength (Weak → Strong):
London < pd-pd < H-bond < Ionic/Covalent/Metallic
🔹 3.7 Dot-and-Cross Diagrams
🧊 What are They?
• Show how electrons are shared or transferred using dots & crosses for each atom.
🧪 What to Represent:
1. Ionic Compounds: Full transfer of electrons, brackets, charges.
2. NaCl, MgO, CaF₂
3. Covalent Compounds: Shared pairs.
4. H₂, O₂, N₂, HCl, CH₄, NH₃, CO₂
5. Dative Bonding:
6. NH₄⁺ → Arrow from N to H⁺
7. Al₂Cl₆ → Lone pair on Cl to Al
💡 You may show expanded octets (SF₆, PCl₅) with extra shared pairs.
✅ You are now fully covered for Chapter 3 – Chemical Bonding! If you reread and understand these
notes + practice MCQs and structured Qs, you’ll be ready for your test.
Would you like me to quiz you or create past paper-style questions now?
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