Chemical bonding and molecular structure
Class 11 · Updated 2026-10-03. Published for practice.
Ionic vs covalent
Ionic bonding: transfer of electrons gives cations and anions; strong electrostatic attraction in a lattice. Typical for metal + non-metal with large electronegativity difference (NaCl, MgO).
Covalent bonding: atoms share electron pairs. Pure covalent is rare; most bonds are polar covalent with partial charges (δ+ and δ−). Fajan’s rules: small, highly charged cation polarises an anion and adds covalent character to an otherwise ionic bond.
Bond length decreases as bond order increases; bond enthalpy usually rises with stronger bonds. Know the trend, not every tabulated value.
Lewis structures and VSEPR
Draw Lewis structures: count valence electrons, place least electronegative atom central (usually), satisfy octets with bonds and lone pairs, check formal charges if needed.
VSEPR: electron pairs (bonding + lone) repel and spread out. Lone pairs occupy more space than bonding pairs, so they compress bond angles (H2O bent ~104.5°, not 109.5°).
Exam favourites: linear (BeCl2, CO2), trigonal planar (BF3), tetrahedral (CH4), trigonal pyramidal (NH3), bent (H2O), trigonal bipyramidal (PCl5), octahedral (SF6). Name the shape and give a rough bond angle.
Valence bond and hybridisation
Hybridisation is a bookkeeping model: mix atomic orbitals to match observed geometry. sp → linear; sp2 → trigonal planar; sp3 → tetrahedral; sp3d and sp3d2 for five and six electron-pair geometries.
Sigma (σ) bonds: head-on overlap along the bond axis. Pi (π) bonds: sideways overlap above and below a σ skeleton. Double bond = one σ + one π; triple = one σ + two π.
Resonance: when one Lewis structure is inadequate, write several contributing structures and delocalise π electrons (O3, CO32-, benzene). The real structure is a hybrid; bond lengths equalise between single and double character.
MO picture and hydrogen bonding
Molecular orbital theory (intro level): combine atomic orbitals to bonding and antibonding MOs. Bond order = (electrons in bonding MOs − electrons in antibonding MOs) / 2. O2 is paramagnetic because of unpaired electrons in π* orbitals — a fact Lewis structures mishandle.
Hydrogen bond: strong dipole–dipole attraction when H is bonded to N, O, or F and approaches another electronegative atom with a lone pair. Explains high boiling points of H2O, HF, and DNA base pairing.
Metallic bonding: delocalised electrons in a lattice of cations — conductivity, malleability. Van der Waals forces: weak attractions between molecules (London dispersion; dipole–dipole).