Chemical Bonding and Molecular Structure

🟢 Lite — Quick Review (1h–1d)

Rapid summary for last-minute revision before your exam.

Chemical Bonding — Key Facts

Chemical bonds hold atoms together in compounds. The three primary bond types are ionic, covalent, and metallic. Bond formation releases energy (exothermic) — breaking bonds requires energy.

Ionic Bonding:

• Transfer of electrons from metal to non-metal
• Example: Na (1s²2s²2p⁶3s¹) → Na⁺ (2s²2p⁶) + e⁻
• Cl (1s²2s²2p⁶3s²3p⁵) + e⁻ → Cl⁻ (1s²2s²2p⁶3s²3p⁶)
• Ionic compounds: high melting/boiling point, conduct electricity in molten state, soluble in polar solvents

Covalent Bonding:

• Sharing of electron pairs between non-metals
• Single bond (σ): 2 electrons shared
• Double bond (σ + π): 4 electrons shared
• Triple bond (σ + 2π): 6 electrons shared
• Polar covalent: unequal sharing due to electronegativity difference

Electronegativity (Pauling Scale):

• F = 4.0 (most electronegative)
• Trend: increases across period, decreases down group
• Difference > 1.7: ionic bond
• Difference 0.4–1.7: polar covalent
• Difference < 0.4: non-polar covalent

⚡ ECAT Exam Tip: Use the Fajan rule — cations with high charge and small size cause distortion of the electron cloud (polarisation), making ionic compounds more covalent.

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🟡 Standard — Regular Study (2d–2mo)

For students who want genuine understanding…

Lewis Structures (Electron Dot Structures):

Steps to draw Lewis structures:

1. Count total valence electrons
2. Identify central atom (usually least electronegative, can form multiple bonds)
3. Connect atoms with single bonds
4. Complete octets of outer atoms
5. Use lone pairs on central atom to form double/triple bonds if needed

Examples:

CO₂ (Carbon dioxide):

• Total valence e⁻ = 4 + 6 + 6 = 16
• O=C=O (double bonds to satisfy carbon’s octet)

NH₃ (Ammonia):

• Total valence e⁻ = 5 + 3(1) = 8
• Central N with 3 bonding pairs and 1 lone pair
• Trigonal pyramidal geometry

VSEPR Theory (Valence Shell Electron Pair Repulsion):

Steric Number	Electron Pair Geometry	Molecular Shape
2	Linear	Linear
3	Trigonal planar	Trigonal planar, bent
4	Tetrahedral	Tetrahedral, trigonal pyramidal, bent
5	Trigonal bipyramidal	See-saw, T-shaped, linear
6	Octahedral	Octahedral, square pyramidal, square planar

Hybridisation:

Hybridisation	Geometry	Bond Angle	Example
sp	Linear	180°	BeCl₂, C₂H₂
sp²	Trigonal planar	120°	BF₃, C₂H₄
sp³	Tetrahedral	109.5°	CH₄, NH₃
sp³d	Trigonal bipyramidal	90°, 120°	PCl₅
sp³d²	Octahedral	90°	SF₆

⚡ ECAT Exam Tip: For molecules with resonance (like O₃, NO₃⁻), all resonance structures are valid but the actual structure is a hybrid. Formal charge calculation: FC = Valence e⁻ - (Non-bonding e⁻ + ½ Bonding e⁻).

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🔴 Extended — Deep Study (3mo+)

Comprehensive coverage for students on a longer study timeline.

Molecular Orbital Theory (MOT):

Antibonding orbitals have a node between nuclei (higher energy); bonding orbitals have no node (lower energy).

Bond Order: $$Bond\ Order = \frac{N_b - N_a}{2}$$

where N_b = electrons in bonding orbitals, N_a = electrons in antibonding orbitals.

MOT for O₂:

• σ1s, σ1s, σ2s, σ2s, σ2pz, π2px = π2py, π2px = π2py, σ*2pz
• Total 16 electrons
• Bond order = (10 - 6)/2 = 2
• O₂ is paramagnetic (2 unpaired electrons in π* orbitals)

Hydrogen Bonding:

A special dipole-dipole attraction when H is bonded to highly electronegative F, O, or N.

Conditions:

1. H bonded to F, O, or N
2. Electronegative atom with lone pair
3. Partial positive H attracted to lone pair

Examples:

• Water (H₂O): explains high boiling point (100°C), ice being less dense than liquid water
• Ammonia (NH₃): boiling point -33°C (higher than PH₃ which has only van der Waals)
• HF: boiling point 19.5°C
• DNA double helix: A-T and G-C base pairing

Dipole Moment (μ):

$$\mu = q \times d$$

Unit: Debye (D); 1 D = 3.336 × 10⁻³⁰ C·m

For CO₂: linear, μ = 0 D (vectors cancel) For H₂O: bent, μ = 1.85 D (vectors don’t cancel)

Van der Waals Forces:

1. London dispersion forces: Instantaneous dipole-induced dipole (present in all molecules)

   • Strength increases with molecular mass and surface area
   • Only intermolecular force in non-polar molecules

2. Debye forces: Polar molecule induces dipole in non-polar molecule

3. Keesom forces: Dipole-dipole interactions (in polar molecules)

Solubility Rules:

“Like dissolves like” — polar solvents dissolve polar/ionic solutes; non-polar solvents dissolve non-polar solutes.

Lattice Energy (Born-Haber Cycle):

Lattice energy = energy released when gaseous ions form one mole of solid ionic compound.

$$U \propto \frac{Z^+ Z^-}{r_0}$$

where Z = ion charge, r₀ = internuclear distance.

⚡ ECAT 2024 Analysis: MOT questions on bond order and paramagnetism appear frequently. Fajan rule and dipole moment calculations are also tested. Remember: He₂ has bond order 0 (doesn’t exist); Ne₂ has bond order 4 but is very unstable.

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