Chemical Bonding

🟢 Lite — Quick Review (1h–1d)

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Chemical Bonding — Key Facts for CUET Ionic bond: transfer of electrons (metal + non-metal); lattice energy $\propto \frac{z^+z^-}{r_0}$ Covalent bond: sharing of electrons (non-metal + non-metal); bond order = (bonding − antibonding electrons)/2 VSEPR: electron pairs repel; lone pairs compress bond angles; AX₂E has < 120°, AX₃ has = 120°, AX₄ has = 109.5° Hybridisation: sp³ (tetrahedral, 4bp), sp² (trigonal planar, 3bp), sp (linear, 2bp) Hydrogen bond: strong intermolecular force (20–40 kJ/mol); explain boiling points of NH₃, H₂O, HF ⚡ Exam tip: Formal charge = valence electrons − non-bonding electrons − (bonding electrons/2). Use to determine best Lewis structure

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🟡 Standard — Regular Study (2d–2mo)

For students who want genuine understanding of molecular bonding.

Chemical Bonding — CUET Chemistry Study Guide

Chemical bonds hold atoms together in molecules and compounds. The three primary bond types are ionic, covalent, and metallic. The type of bonding depends on electronegativity difference: large difference ($\Delta\chi > 1.7$) → ionic; small difference ($\Delta\chi < 0.4$) → nonpolar covalent; intermediate → polar covalent.

Ionic Bonding: Formed by complete transfer of electrons from metal to non-metal. Lattice energy (energy released when gaseous ions form solid lattice): $U = -\frac{N_A z^+z^- e^2}{4\pi\varepsilon_0 r_0}(1 - \frac{1}{n})$, where $n$ is Born exponent (8–12 depending on electronic configuration). Higher lattice energy means stronger ionic bond and higher melting point. NaCl (M.P. 801°C) vs KCl (M.P. 770°C) — smaller Na⁺ has higher lattice energy.

Covalent Bonding: Formed by sharing of electron pairs between atoms. Bond order determines bond strength and length: higher bond order = shorter, stronger bond. For O₂ (bond order 2): double bond, bond length 121 pm. For N₂ (bond order 3): triple bond, bond length 110 pm, very strong (945 kJ/mol).

Lewis Structures: Show all valence electrons. Octet rule (8 electrons per atom in compounds) applies to second-period elements. Exceptions: BH₃ (6 electrons), PCl₅ (10 electrons in P), SF₆ (12 electrons in S), XeF₄ (12 electrons in Xe), elements in period 3+ can expand octet.

VSEPR Theory: Valence Shell Electron Pair Repulsion theory predicts molecular geometry based on repulsions between electron pairs. Count total electron pairs (bonding + lone) around central atom:

• 2 bp, 0 lp → linear (180°)
• 3 bp, 0 lp → trigonal planar (120°)
• 2 bp, 1 lp → bent (~117°)
• 4 bp, 0 lp → tetrahedral (109.5°)
• 3 bp, 1 lp → trigonal pyramidal (~107°)
• 2 bp, 2 lp → bent (~104.5°)

Lone pairs repel more than bonding pairs, compressing bond angles. Water (2 bp + 2 lp) has bond angle 104.5° instead of 109.5°.

Hybridisation: Mixing of atomic orbitals to form new hybrid orbitals for bonding:

• sp³: four equivalent orbitals pointing to corners of tetrahedron (e.g., CH₄, NH₃, H₂O)
• sp²: three equivalent orbitals in plane at 120° (e.g., C₂H₄, BF₃)
• sp: two equivalent orbitals at 180° (e.g., C₂H₂, BeCl₂)

Example: Predict geometry of PCl₃ and SO₂. PCl₃: P has 5 valence e⁻ + 3×7 from Cl = 26 e⁻. Three bonding pairs + one lone pair → tetrahedral electron geometry, trigonal pyramidal molecular shape, ~107° bond angle. SO₂: S has 6 valence e⁻ + 2×6 from O = 18 e⁻. Two double bonds (count as 2 electron regions each) + one lone pair on S → trigonal planar electron geometry, bent molecular shape (~118°).

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🔴 Extended — Deep Study (3mo+)

Comprehensive coverage for students on a longer study timeline.

Chemical Bonding — Complete CUET Chemistry Notes

Molecular Orbital Theory: More advanced than Lewis/VSEPR, MOT treats electrons as delocalised over the entire molecule. Orbitals combine to form bonding (lower energy) and antibonding (higher energy) molecular orbitals. Electrons fill in order of increasing energy (M.O. diagram).

For O₂: σ1s², σ1s², σ2s², σ2s², σ2pz², π2px² = π2py², π2px¹ = π2py¹ (triplet diradical, explains paramagnetism). Bond order = (8 bonding − 6 antibonding)/2 = 1.

Bond Length Trends: In a period, bond length decreases with increasing bond order (C–C > C=C > C≡C). Down a group, bond length increases (C–H < Si–H < Ge–H). Resonance structures delocalise electrons, making actual bonds intermediate in length (benzene: all C–C bonds = 139 pm, between single 154 pm and double 134 pm).

Fajans’ Rule: For ionic compounds with similarly sized ions, greater charge on cation → more covalent character. Larger cation + smaller anion → more covalent character. AgCl is covalent (large Ag⁺, polarising), while NaCl is ionic.

Dipole Moment: $\mu = Q \times r$. For polar covalent bonds, dipole moment measures separation of charge. Vector sum of individual bond dipoles gives molecular dipole. CO₂ (linear, O=C=O): bond dipoles cancel, μ = 0. H₂O (bent): bond dipoles don’t cancel, μ = 1.85 D. BF₃ (trigonal planar): μ = 0.

Hydrogen Bonding: Special dipole-dipole interaction when H is bonded to highly electronegative F, O, or N. Strength 20–40 kJ/mol (much stronger than van der Waals ~1 kJ/mol). Explains anomalous properties: ice density < water (hydrogen bonds create open cage structure in solid), high boiling points of NH₃, H₂O, HF compared to analogues (PH₃, H₂S, HCl).

Intermolecular Forces Hierarchy:

1. Ionic bonding (strongest, 400–4000 kJ/mol)
2. Covalent bonding (200–1100 kJ/mol)
3. Hydrogen bonding (10–40 kJ/mol)
4. Dipole-dipole (5–20 kJ/mol)
5. London dispersion (0.5–5 kJ/mol) — present in all molecules

Valence Bond Theory vs MOT: VBT: localised bonds formed by overlap of atomic orbitals (hybrid orbitals). MOT: delocalised electrons in molecular orbitals. VBT explains geometry well; MOT explains bond order, magnetism, and resonance better.

BORN-HABER CYCLE: Calculate lattice energy indirectly using Hess’s law. For NaCl: ΔHf° = ΔHat° + ½ΔHd° + ΔAi° + ΔEae° + ΔHl°. This shows that ionic compounds form because lattice energy released is greater than energy invested.

CUET Exam Patterns (2022–2024):

• VSEPR geometry predictions are most frequently tested (1–2 marks)
• Hybridisation and shapes (sp, sp², sp³) appear every year
• Formal charge calculations tested in 2023
• Hydrogen bonding examples (water, ammonia, HF) are common MCQs
• Exceptions to octet rule (PCl₅, SF₆) occasionally tested
• Common mistakes: forgetting lone pair repulsion compresses angles; misidentifying hybridisation in molecules with resonance

⚡ Key insight: In VSEPR, count the regions of high electron density — each bond (single, double, triple) counts as one region, each lone pair counts as one region. The central atom’s hybridisation = number of regions. For formal charge minimisation, put negative charge on more electronegative atoms and positive charge on less electronegative atoms.

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