← Chemistry — Std 11
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Chemical Bonding

Ch. 5Std 11

Easy Overview

Why do atoms bond? Why is diamond hard but graphite soft — both pure carbon? Why does water boil at 100°C but H₂S at -60°C? Chemical bonding holds the answers. Atoms bond to achieve full outer shells (octet rule). They can transfer electrons (ionic), share (covalent), or pool (metallic). Ionic bonds form between metals and non-metals (electron transfer). Covalent bonds form between non-metals (electron sharing). VSEPR theory predicts molecular shapes. Hybridization explains how carbon forms sp³, sp², or sp bonds. Molecular Orbital Theory explains things Lewis can't — like O₂'s paramagnetism. Intermolecular forces (hydrogen bonding, dipole-dipole, London dispersion) determine boiling points, solubility, and even the structure of DNA and proteins.

Octet Rule and Its Limitations

Atoms tend to achieve 8 valence electrons (2 for H, He). Exceptions: incomplete octet (BeCl₂, BF₃), expanded octet (PCl₅, SF₆, IF₇ — period 3+ have d orbitals), odd-electron (NO, NO₂). Octet rule is a useful guideline, not universal.

Ionic Bond — Electron Transfer and Lattice Energy

Metal transfers e⁻ to non-metal, held by electrostatic attraction in crystal lattice. Lattice energy (U ∝ q₁q₂/r) determines strength: MgO > NaCl. Born-Haber cycle calculates U from known thermodynamic data. Properties: high MP, conducts when molten/dissolved, brittle (like charges repel when layers shift), water-soluble.

Covalent Bond — Electron Sharing

Single (σ), double (σ+π), triple (σ+2π). Bond length decreases, strength increases with order: C-C (154 pm, 348 kJ/mol), C=C (134, 614), C≡C (120, 839). Polarity from ΔEN: nonpolar (ΔEN≈0), polar covalent (0<ΔEN<1.7). Dipole moment μ = q×d (Debye). Covalent compounds: low MP, don't conduct.

Lewis Structures — Drawing Molecules

Count valence e⁻, place least EN atom center, connect with single bonds, distribute remaining e⁻ as lone pairs, form multiple bonds if octet incomplete. Formal charge = valence - (nonbonding + ½ bonding). Minimize formal charges; negative charge on more EN atoms. Limitations: no 3D, fails for delocalized systems.

VSEPR Theory — Predicting Molecular Shapes

Electron pairs repel → maximize separation. Steric number = bonding + lone pairs. SN2: linear (180°, BeCl₂). SN3: trigonal planar (120°, BF₃). SN4: tetrahedral (109.5°, CH₄); with 1 lone pair → trigonal pyramidal (NH₃, 107°); 2 lone pairs → bent (H₂O, 104.5°). SN5: trigonal bipyramidal (PCl₅). SN6: octahedral (SF₆). Lone pairs repel more than bonding pairs.

Hybridization — Mixing Orbitals for Better Bonding

sp³: 1s+3p → 4 tetrahedral (109.5°), 25% s character — alkanes. sp²: 1s+2p → 3 trigonal planar (120°), 33.3% s, one p for π — alkenes. sp: 1s+1p → 2 linear (180°), 50% s, two p for π — alkynes. More s character = shorter, stronger bonds. sp³d (PCl₅), sp³d² (SF₆).

Valence Bond Theory — Orbital Overlap

Bond forms when half-filled orbitals overlap with opposite spins. Greater overlap = stronger bond. σ bonds: end-to-end overlap (s-s, s-p, p-p along axis). π bonds: lateral p-orbital overlap above/below axis. σ stronger than π (more effective overlap). Double bond = σ+π, not twice as strong as single.

Molecular Orbital Theory — Delocalized Electrons

Atomic orbitals combine → bonding MO (lower energy) + antibonding MO (higher energy, *). Fill: Aufbau, Pauli, Hund. Bond order = (B - AB)/2. >0 = stable. O₂: bond order 2, 2 unpaired e⁻ in π* → paramagnetic (Lewis can't explain this). He₂: BO=0, doesn't exist. O₂⁺ > O₂ > O₂⁻ > O₂²⁻ in bond strength.

Polarity of Bonds and Molecules

Bond polarity from ΔEN. Dipole moment μ = q×d. Net μ = vector sum. Nonpolar: dipoles cancel (CO₂ linear, CCl₄ tetrahedral) or symmetric. Polar: H₂O (1.85 D). 'Like dissolves like': polar solvents dissolve polar/ionic; nonpolar dissolve nonpolar (fats, oils). Soaps: amphiphilic (nonpolar tail + polar head) → micelles.

Hydrogen Bonding — The Strong Intermolecular Force

H bonded to N/O/F attracted to lone pair on another N/O/F. Strength 5-30 kJ/mol — stronger than dipole-dipole (2-5) but weaker than covalent (150-800). Explains H₂O's high BP (100°C vs H₂S -60°C), ice floating (open hexagonal lattice), DNA base pairing. Intramolecular H-bonding in o-nitrophenol reduces water solubility.

Dipole-Dipole and London Dispersion Forces

Dipole-dipole: + end to - end of polar molecules. Dispersion (London): temporary dipoles from electron fluctuations. Exist in all molecules. Strength ↑ with molecular size (more e⁻ = more polarizable) and surface area (elongated > spherical). Explains trend: F₂, Cl₂ (gas), Br₂ (liquid), I₂ (solid).

Resonance — When One Structure Isn't Enough

Actual structure = hybrid of multiple canonical structures. Resonance energy = extra stability. Rules: same connectivity, same # of e⁻, more bonds = more stable, minimal charge separation, -ve charge on EN atoms. Benzene: two Kekulé forms, all C-C = 139 pm. CO₃²⁻: three equivalent forms. RCOO⁻: charge delocalized over both O. Shown with ↔.

Fajan's Rules — Ionic vs Covalent Character

Smaller cation + larger anion = more covalent character. Higher charges = more covalent. LiCl > NaCl (smaller cation). LiI > LiF (larger anion). AlCl₃ > NaCl (higher charge). AgCl insoluble (covalent from d¹⁰ configuration) vs NaCl soluble. AgF soluble (F⁻ too small) vs AgI insoluble (I⁻ highly polarizable).

Coordinate (Dative) Covalent Bond

Both e⁻ from same atom (donor) to acceptor with empty orbital. Once formed = indistinguishable from regular covalent. Examples: NH₄⁺ (NH₃→H⁺), H₃O⁺ (H₂O→H⁺), F₃B:NH₃, CO→metal (hemoglobin binding), metal complexes (NH₃, H₂O, CN⁻ → metal ion). Shown with → from donor to acceptor.

Metallic Bond — The Electron Sea Model

Positive metal ions in regular lattice surrounded by delocalized valence electrons ('sea'). Explains: electrical conductivity (mobile e⁻), thermal conductivity, malleability (layers slide, e⁻ sea adjusts), luster (e⁻ reflect light). Strength ∝ # delocalized e⁻ per atom: transition metals > alkali metals.

Sigma and Pi Bonds

σ: end-to-end overlap along axis, e⁻ density between nuclei — strong. Every single bond = σ. π: lateral p-orbital overlap above/below axis — weaker. Double = σ+π. Triple = σ+2π (π bonds perpendicular). π bonds restrict rotation → cis-trans isomerism. Conjugated π systems → delocalization.

Bond Parameters — Length, Energy, Angle, Order

Bond length: sp³C-C (154 pm) > sp²C=C (134) > spC≡C (120). Order: single (1) < double (2) < triple (3). Energy: C-C (348), C=C (614), C≡C (839 kJ/mol). π bond (~265) weaker than σ (~348). Angle: VSEPR/ hybridization. Shorter = stronger. Higher order = shorter and stronger.

Comparison of Bonding Types

ΔEN spectrum: >1.7 ionic (NaCl), 0.4-1.7 polar covalent (H₂O), <0.4 nonpolar (H₂). Ionic: high MP, conducts molten, brittle, water-soluble. Covalent molecular: low MP, doesn't conduct, variable solubility. Covalent network (diamond): extremely high MP, hardest. Metallic: conducts, malleable, lustrous.

Key Points

  • •Octet: 8 valence e⁻ (2 for H). Exceptions: incomplete, expanded, odd-electron
  • •Ionic: e⁻ transfer, lattice energy ∝ q₁qâ‚‚/r, Born-Haber cycle
  • •Covalent: sharing; single (σ) < double (σ+Ï€) < triple (σ+2Ï€)
  • •Lewis: minimize formal charges; negative on more EN atoms
  • •VSEPR: e⁻ pairs repel; lone pairs repel more than bonding pairs
  • •Shapes: linear 2, trigonal planar 3, tetrahedral 4, trigonal bipyramidal 5, octahedral 6
  • •Hybridization: sp³ (109.5°, 25% s), sp² (120°, 33.3% s), sp (180°, 50% s)
  • •VBT: σ (end-to-end, strong) and Ï€ (lateral, weaker)
  • •MOT: BO = (B-AB)/2; Oâ‚‚ paramagnetic (2 unpaired e⁻ in Ï€*)
  • •H-bonding: H-N/O/F; explains Hâ‚‚O high BP, ice density, DNA
  • •London dispersion: temporary dipoles; ↑ with size and surface area
  • •Resonance: hybrid of structures; extra stability = resonance energy
  • •Fajan's: small/high charge cation + large anion = more covalent
  • •Coordinate bond: donor→acceptor (NH₄⁺, H₃O⁺, metal complexes)
  • •Metallic: electron sea; explains conductivity, malleability, luster
  • •Î”EN > 1.7 ionic, 0.4-1.7 polar covalent, <0.4 nonpolar covalent

Practice Questions

  • Draw Lewis structures for Hâ‚‚O, NH₃, COâ‚‚, SO₄²⁻. Calculate formal charges.
  • Explain Hâ‚‚ formation using MOT. Calculate BO. Explain Oâ‚‚ paramagnetism.
  • Predict shapes of CHâ‚„, Câ‚‚Hâ‚„, Câ‚‚Hâ‚‚ using hybridization.
  • Why does ice float on water? Explain with H-bonding.
  • Predict shapes of NH₃, Hâ‚‚O, SF₆, PClâ‚… using VSEPR.
  • Explain resonance: benzene, CO₃²⁻, ozone.
  • Apply Fajan's: why LiCl more covalent than NaCl? Why AgCl insoluble but NaCl soluble?
  • Differentiate σ and Ï€ bonds. How does Ï€ bond cause cis-trans isomerism?
  • Explain metallic bonding with electron sea model.