Chemical Bonding — JEE Main & Advanced Notes

Count the valence electrons of all atoms, add one per negative charge and subtract one per positive charge. Connect the atoms, distribute electrons, then check formal charges and the total electron count. Hydrogen needs a duet; an octet is a useful main-group guide, with electron-deficient and odd-electron exceptions.

Formal charge = valence electrons of the free atom − nonbonding electrons − half the bonding electrons. The sum of formal charges must equal the species' charge. Formal charge is bookkeeping within a Lewis structure, not the same as oxidation state or measured partial charge.

A single bond contains one σ bond; a double contains one σ and one π; a triple contains one σ and two π. A multiple bond still counts as one region around a central atom in a VSEPR shape calculation.

Count bonding regions and lone pairs around the central atom. The electron-domain arrangement includes both; molecular shape describes the positions of atoms. There is no fixed two-degree decrease per lone pair. Bond angles depend on the atoms, bonds and electron distribution.

• AX₂: linear; ideal angle 180°.
• AX₃: trigonal planar; ideal angle 120°. AX₂E is bent.
• AX₄: tetrahedral; ideal angle about 109.5°. AX₃E is trigonal pyramidal; AX₂E₂ is bent.
• Five domains: trigonal bipyramidal arrangement. AX₄E is seesaw; AX₃E₂ is T-shaped; AX₂E₃ is linear.
• Six domains: octahedral arrangement. AX₅E is square pyramidal; AX₄E₂ is square planar.

In the five-domain model, lone pairs prefer equatorial positions to reduce 90° interactions. With two lone pairs in an octahedral arrangement, opposite positions give the square-planar case. X counts bonded atoms and E counts central lone pairs.

Resonance contributors keep the same atom connectivity and differ in electron placement. The molecule does not jump back and forth between drawings. Equivalent contributors can explain equal bond lengths, such as the three C–O bonds in carbonate.

For simple main-group localised bonding, two, three and four electron domains commonly correspond to sp, sp² and sp³ descriptions. These are bonding models. The familiar sp³d and sp³d² labels for hypervalent examples are exam bookkeeping conventions; they do not establish substantial d-orbital participation in the real bond.

To assess molecular polarity, combine geometry with the directions and magnitudes of bond dipoles. Equal surrounding atoms in a symmetric geometry can give cancellation. A lone pair is a prompt to inspect shape, not a universal proof of polarity: XeF₄ has lone pairs but its square-planar bond dipoles cancel.

Use an energy ordering appropriate to the species, then apply Pauli and Hund. In the standard second-period homonuclear diagrams, π2p lies below σ2p for B₂, C₂ and N₂; σ2p lies below π2p for O₂ and F₂. Include antibonding orbitals, marked with an asterisk, when filling.

Bond order = (bonding electrons − antibonding electrons)/2. Unpaired electrons produce paramagnetism. For O₂, the two highest electrons occupy separate π* orbitals; adding or removing electrons changes both bond order and the pairing.

Compare bond lengths using bond order within a related series, not as an absolute law across unrelated elements. A basic MO diagram giving zero bond order predicts no ordinary net covalent bond in that model; it is not proof that weakly bound species can never exist.

Ionic and covalent character form a continuum. In a simple electrostatic comparison of similar lattices, larger charge products and smaller interionic separation increase the magnitude of attraction. Do not compare lattice energies using charge alone when structure and size differ.

Fajans' reasoning asks whether a cation strongly distorts an anion's electron cloud: small, highly charged cations and large, easily polarised anions favour greater covalent character.

Common school-level hydrogen-bond donors contain H bonded to N, O or F; an acceptor needs an available electron-rich site. Hydrogen bonding can occur within a molecule or between molecules. Intermolecular association can raise boiling points, but neither melting-point order nor an exact boiling point follows from counting hydrogen bonds alone.

Revision sequence: electron count → Lewis structure → domains and shape → polarity → MO filling where needed. Revisit atomic structure and orbital filling, then browse the Chemistry chapter directory.

Further reading: molecular shape and polarity and molecular orbital theory. These notes and worked problems were independently written for LearnAtMyPlace.