Chemical Bonding — JEE Main & Advanced Notes

Explain structure and properties through ionic/covalent bonding, VSEPR, hybridisation, resonance, dipole moment and molecular orbital theory.

Priority: Must Do. Unit: Inorganic Chemistry. Level: Moderate.

How the uploaded material was used: Mapped from bonding, VSEPR, hybridisation, resonance and molecular orbital theory practice material. The final student-facing notes and questions are original, rewritten and copyright-safe.

These are the ideas that decide most correct answers in Chemical Bonding.

• Lewis structures — the foundation of bonding analysis: Drawing a correct Lewis structure is mandatory before any other bonding analysis. Procedure: (1) Count total valence electrons for all atoms. (2) Connect atoms with single bonds using the least electronegative atom as central. (3) Complete octets on outer atoms first using remaining electrons. (4) If the central atom is short of an octet, form double or triple bonds by shifting lone pairs from outer atoms. Formal charge = valence electrons minus nonbonding electrons minus half the bonding electrons. The best Lewis structure minimises formal charges and places negative formal charge on more electronegative atoms.
• VSEPR theory — geometry from electron pair repulsion: Valence Shell Electron Pair Repulsion theory predicts molecular geometry by minimising repulsion among all electron pairs around the central atom. Count the steric number SN = sigma bonds plus lone pairs. SN=2: linear (180°). SN=3: trigonal planar with 0 lone pairs (120°) or bent with 1 lone pair (about 117°). SN=4: tetrahedral (0 LP, 109.5°), trigonal pyramidal (1 LP, about 107°), bent (2 LP, about 104.5°). SN=5: trigonal bipyramidal — lone pairs prefer equatorial positions. SN=6: octahedral — lone pairs prefer trans positions. Lone pairs repel more strongly than bonding pairs because they are not shared between two nuclei.
• Hybridisation — linking geometry to orbitals: sp hybridisation (SN=2) gives 2 equivalent orbitals at 180° and leaves 2 unhybridised p orbitals for pi bonds. sp2 (SN=3) gives 3 orbitals at 120° and leaves 1 unhybridised p orbital for 1 pi bond. sp3 (SN=4) gives 4 orbitals at 109.5°. sp3d (SN=5) gives trigonal bipyramidal. sp3d2 (SN=6) gives octahedral geometry. Always draw the Lewis structure first, then count SN, then assign hybridisation. Number of pi bonds equals number of unhybridised p orbitals being used.
• Resonance — delocalisation and stabilisation: When a single Lewis structure cannot represent the bonding adequately (O3, SO3, benzene, NO3-), multiple resonance structures are drawn. The real molecule is the resonance hybrid — bonds are intermediate between the extremes. Resonance lowers the energy of the molecule (resonance stabilisation energy). Rules: all resonance structures must have identical atom positions (only electrons move); the structure with more bonds, less charge separation, and negative charge on the more electronegative atom contributes more to the hybrid.
• Molecular orbital theory for diatomics: In MOT, atomic orbitals combine to form bonding MOs (lower in energy, stabilising) and antibonding MOs (higher in energy, destabilising, marked with asterisk). Fill MOs using Aufbau, Pauli, and Hund rules — exactly like atomic orbital filling. Bond order = (bonding electrons minus antibonding electrons) divided by 2. Bond order of 0 means the molecule does not exist (like He2). N2 has bond order 3 — the most stable diatomic. O2 has bond order 2 and two unpaired electrons in pi*2p → O2 is paramagnetic, which Lewis structure cannot predict. MO energy order for B2 to N2: sigma1s, sigma*1s, sigma2s, sigma*2s, pi2p (=pi2p), sigma2p, pi*2p (=pi*2p), sigma*2p. For O2 and F2, sigma2p moves below pi2p.
• Ionic bonding, lattice energy and Fajans rules: Ionic bonding forms when the electronegativity difference exceeds about 1.7 (Pauling scale). Lattice energy (energy released forming 1 mol of crystal from gaseous ions) is larger for higher ionic charges and smaller ionic radii — consistent with Coulomb attraction. The Born-Haber cycle uses Hess's law to calculate lattice energy indirectly from measurable quantities. Fajans rules: a small, highly-charged cation strongly polarises the electron cloud of a large anion, pulling electron density toward itself and introducing covalent character. AlCl3 is partially covalent; NaCl is essentially fully ionic.
• Hydrogen bonding — conditions and consequences: H-bonding occurs when hydrogen is covalently bonded to N, O, or F (the three electronegative, small atoms with lone pairs), and the partially positive H attracts a lone pair on a nearby electronegative atom. Intermolecular H-bonding raises boiling point, melting point, and viscosity — this explains the anomalously high boiling points of H2O (100°C), HF (19.5°C), and NH3 (-33°C) compared to other hydrides of the same group. Intramolecular H-bonding (within the same molecule, e.g., ortho-nitrophenol, maleic acid) results in a lower boiling point than the para or trans isomer where only intermolecular H-bonding is possible.
• Dipole moment and molecular polarity: A bond is polar if the bonded atoms have different electronegativities. Whether the molecule is polar depends on the vector sum of all individual bond dipoles — and this depends critically on the molecular geometry. CO2: linear, two equal and opposite C=O dipoles cancel perfectly → non-polar despite having polar bonds. H2O: bent geometry, dipoles add up → polar (mu = 1.85 D). BF3: trigonal planar → non-polar. NF3: trigonal pyramidal → polar (lone pair contributes). CCl4: tetrahedral → non-polar. CHCl3: one H instead of Cl breaks the symmetry → polar.

• Formal charge = valence electrons minus nonbonding electrons minus half of bonding electrons
• Steric number (SN) = sigma bonds to central atom + lone pairs on central atom
• VSEPR shapes: SN=2 linear; SN=3 trigonal planar or bent; SN=4 tetrahedral/trigonal pyramidal/bent; SN=5 trigonal bipyramidal; SN=6 octahedral
• Bond order (MOT) = (electrons in bonding MOs minus electrons in antibonding MOs) divided by 2
• Higher bond order implies shorter bond length, stronger bond, and higher bond dissociation energy
• Dipole moment: mu = q x d (charge times bond length); unit is Debye (D)
• Net molecular dipole = vector sum of all bond dipoles; geometry determines if they cancel
• Lattice energy proportional to Z+Z-/r; higher charge and smaller ions give larger lattice energy
• Fajans rules: covalent character increases with smaller cation, larger anion, higher cation charge
• Bond angles reduce from ideal by approximately 2 degrees per lone pair on central atom

Derivation / logic hint: Do not plug values blindly. Start from conservation of mass/charge, equilibrium definition, energy balance, electron movement, structure-property relation, or stability of the product/intermediate.

Most negative marks in this chapter come from condition errors, not lack of memory.

• Confusing electron geometry with molecular shape: SN=4 with 2 lone pairs has tetrahedral electron geometry but a bent molecular shape. The molecular shape describes only the positions of atoms, not lone pairs.
• Ignoring lone pairs in VSEPR: lone pairs repel more strongly than bonding pairs since they are not shared between two nuclei. Always count all lone pairs on the central atom when determining SN and predicting shape.
• Applying the octet rule to all atoms: BF3 has only 6 electrons on B (electron-deficient, Lewis acid). SF6 and PCl5 have expanded octets (d orbital involvement). H needs only 2 electrons. Recognise the period — 3rd period and beyond can expand the octet.
• Assigning hybridisation before drawing the Lewis structure: the Lewis structure determines how many lone pairs are on the central atom. Without it, SN is wrong and hybridisation is wrong. The sequence must always be: Lewis structure → SN → hybridisation.
• Concluding polarity from bond polarity alone: CO2 has polar C=O bonds but is non-polar overall (linear geometry, dipoles cancel). Shape is essential for the polarity determination. Do not skip the geometry step.

For JEE Main, prioritise direct formula use, NCERT-aligned facts, named-reaction recognition, trend comparison and quick elimination. Target 60–90 seconds per question.

• Lewis structures and formal charge
• Ionic bonding and lattice energy
• Fajans rules and covalent character
• VSEPR theory and molecular shapes

For JEE Advanced, combine ideas. Expect assertion-reason, integer, multiple-correct, paragraph-style and hidden-condition problems. Before finalising, ask which assumption the question is testing.

• Fajans rules and covalent character
• VSEPR theory and molecular shapes
• Hybridisation (sp sp2 sp3 sp3d sp3d2)
• Resonance and delocalisation
• Molecular orbital theory (MOT)
• Dipole moment and molecular polarity
• Hydrogen bonding

Use this block in the final 24–48 hours before a mock.

• Lewis structure: count total valence electrons, form single bonds, complete outer octets, then form multiple bonds if central atom is short. Compute formal charges.
• VSEPR: SN = sigma bonds + lone pairs on central atom. Use SN to determine electron geometry, then remove lone pairs to get molecular shape.
• MOT: fill MOs in energy order. BO = (bonding minus antibonding)/2. Unpaired electrons in final configuration means paramagnetic (O2 is paramagnetic). N2 has BO=3.
• Resonance: identical atom positions, only electrons move. More bonds and less charge separation means a better contributing structure.
• H-bonding conditions: H bonded to N, O, or F. Intermolecular H-bonding raises BP. Intramolecular H-bonding lowers BP compared to the intermolecular isomer.