Covalent Bonding

• Covalent bonds — electron pair sharing between atoms (same or different non-metals).
• Lewis structures, octet rule, and exceptions (incomplete octet, expanded octet, odd electrons).
• Bond order, bond length, and bond enthalpy relationships.
• Polar covalent bonds and dipole moment direction (toward more electronegative atom).

Level 1 — Lewis structures and octet

Verbal: Atoms share electron pairs so each achieves noble gas configuration (usually octet). Single, double, triple bonds share 2, 4, 6 electrons respectively.

Symbolic: Bond order = (bonding − antibonding e⁻)/2; μ = q × d (dipole); FC = V − N − B/2.

Steps for Lewis: (1) Count valence e⁻ (2) Identify central atom (3) Connect with bonds (4) Complete octets (5) Minimise formal charges if needed.

Formal charge: FC = V − N − B/2 (V valence, N nonbonding, B bonding e⁻). Best structure: FC near 0, negative FC on electronegative atom.

Level 2 — Parameters and polarity

Bond enthalpy: Energy to break 1 mol bonds in gas — higher order → higher bond enthalpy (generally).

Electronegativity (Pauling): F highest (4.0). ΔEN > ~1.7 often ionic character significant; 0.4–1.7 polar covalent.

Dipole moment μ = q × d (debye). CO₂ μ = 0 (linear, vectors cancel); H₂O μ ≠ 0 (bent).

Resonance: Multiple Lewis structures — real structure is hybrid (e.g., O₃, CO₃²⁻, benzene preview).

NCERT spotlight — Resonance and formal charge

For O3, draw two equivalent Lewis structures with double bond on alternate oxygens. Real bond order is 1.5. Formal charge minimisation guides best structure.

VSEPR and polarity: CH4 nonpolar tetrahedral; NH3 pyramidal polar; H2O bent polar despite two O-H bonds.

Coordinate bond: One atom supplies both electrons in the shared pair — NH4+ formation from NH3 + H+ example in NCERT.

Draw Lewis structure of H₂SO₄ (simplified: S central, expanded octet). Identify polar bonds and net molecular dipole.

Step 1 — Valence: H 1×2 + S 6 + O 6×4 = 32 e⁻.
Step 2 — S central bonded to 4 O; two O also bonded to H (O−H).
Step 3 — S expanded octet with double bonds to two O (common NCERT representation).
Step 4 — S−O and O−H bonds polar (O more EN); tetrahedral-ish about S.
Step 5 — H₂SO₄ polar molecule — dissolves ionically in water (acid behaviour linked later).

Applications — bond energy and reaction enthalpy

H-H bond enthalpy 436 kJ/mol; H-Cl 431 kJ/mol. Reaction H2 + Cl2 -> 2HCl: bonds broken minus bonds formed approximates Delta H — Hess law connection. Bond order in N2 equals 3 (triple bond) explains high dissociation energy and inertness at room temperature compared to O2 bond order 2.

Deep dive — Lewis structures and molecular polarity

Draw Lewis for CO: triple bond with lone pairs satisfies octet on both C and O — bond order three, very strong. CO2 linear O=C=O — two double bonds, no dipole despite polar bonds. SO2 bent with lone pair on S — net dipole. Hypervalent molecules PCl5 (10 electrons around P), SF6 (12 around S) — expanded octet using vacant 3d orbitals in period 3+ elements per NCERT explanation. Formal charge calculation chooses best resonance structure: FC close to zero, negative FC on electronegative atom. Bond enthalpy average values: H-H 436, C-C 347, C=C 614, C≡C 839 kJ/mol — explain why saturated fats less reactive than unsaturated with double bonds. Coordinate covalent bond in NH4+: lone pair on N donated to H+ — product indistinguishable from normal covalent once formed. Metallic bond electron sea model preview — delocalised electrons explain conductivity malleability distinct from covalent network solids like diamond.

Review and practice drill

Review checklist: (1) Shared electron pairs. (2) Lewis structures and octet. (3) Formal charge minimisation. (4) Bond order length enthalpy trends. Practice: CO2 linear O=C=O — two double bonds, no dipole.

• Draw Lewis for NH₃ and state molecular geometry (preview VSEPR: pyramidal).
• Compare bond lengths C−C, C=C, C≡C.
• Why is HCl polar but Cl₂ non-polar?

Open the Practice tab for graded questions on Covalent.