Level: ISC Chemistry Class 11
Author: Dr E. Ramanathan PhD
Topic Tree
Here is a structured Topic Tree for Chemical Bonding – ISC Class 11 Chemistry:
Chemical Bonding – Topic Tree
1. Introduction to Chemical Bonding
- Importance of bonding
- Octet rule
- Cause of bond formation (stability, lower energy)
2. Types of Chemical Bonds
- Ionic Bond
- Formation by electron transfer
- Lattice energy
- Properties
- Covalent Bond
- Formation by sharing of electrons
- Types: single, double, triple
- Lewis structures
- Coordinate (Dative) Bond
- Donor-acceptor concept
- Examples
3. Theories of Bonding
- Valence Bond Theory (VBT)
- Concept of orbital overlap
- Types of overlap: σ (sigma), π (pi)
- Hybridization
- Definition
- Types: sp, sp², sp³, sp³d, sp³d²
- Shapes of molecules (VSEPR theory connection)
- Molecular Orbital Theory (MOT)
- Formation of bonding and antibonding orbitals
- Energy-level diagrams
- Bond order and magnetic properties
4. Bond Parameters
- Bond length
- Bond angle
- Bond enthalpy
- Bond order
- Polarity of bonds (dipole moment, electronegativity difference)
5. Polarity and Forces
- Polar vs Nonpolar covalent bonds
- Intermolecular forces
- Dipole–dipole interaction
- London dispersion forces
- Hydrogen bonding
6. Theories of Shapes of Molecules
- VSEPR Theory
- AXn notation
- Common molecular geometries:
- Linear, trigonal planar, tetrahedral, trigonal bipyramidal, octahedral, bent, pyramidal
7. Special Topics
- Resonance structures
- Partial ionic character of covalent bonds (Fajan’s rule)
- Metallic bonding (electron sea model, band theory)
Important Terms and Defintions
Chemical Bonding Glossary (45 Terms)
| Term | Definition | Example / Hint |
|---|---|---|
| 1. Chemical Bond | Attractive force holding atoms together in a stable arrangement. | H₂, NaCl |
| 2. Octet Rule | Atoms tend to have 8 electrons in outer shell for stability. | Ne, Cl⁻ |
| 3. Duet Rule | H & He stable with 2 valence electrons. | He, H₂ |
| 4. Electronegativity | Ability of atom to attract shared electron pair. | F > Cl > Br |
| 5. Ionization Energy | Energy needed to remove an electron from gaseous atom. | Na → Na⁺ |
| 6. Electron Affinity | Energy released when electron is added to gaseous atom. | Cl + e⁻ → Cl⁻ |
| 7. Ionic Bond | Bond formed by transfer of electrons. | NaCl, MgO |
| 8. Lattice Energy | Energy released when gaseous ions form 1 mole of ionic solid. | Na⁺ + Cl⁻ → NaCl |
| 9. Covalent Bond | Bond formed by sharing of electron pairs. | H₂O, CH₄ |
| 10. Bond Pair | Pair of electrons shared between two atoms. | H–O–H |
| 11. Lone Pair | Pair of valence electrons not shared in bonding. | O in H₂O |
| 12. Coordinate Bond | Both electrons in shared pair from one atom. | NH₄⁺ |
| 13. Resonance | Delocalization of electrons represented by multiple structures. | CO₃²⁻, Benzene |
| 14. Bond Length | Distance between nuclei of two bonded atoms. | H–H = 74 pm |
| 15. Bond Angle | Angle between two covalent bonds at same atom. | H₂O = 104.5° |
| 16. Bond Order | Number of bonds between two atoms. | O₂ = 2, N₂ = 3 |
| 17. Bond Enthalpy | Energy required to break one mole of bonds in gaseous state. | H–H = 436 kJ/mol |
| 18. Sigma (σ) Bond | Formed by head-on orbital overlap. | H–H, C–H |
| 19. Pi (π) Bond | Formed by lateral overlap of orbitals. | C=C in ethene |
| 20. Valence Bond Theory (VBT) | Bonding explained by overlap of atomic orbitals. | H₂ molecule |
| 21. Hybridization | Mixing of atomic orbitals to form new orbitals. | sp³ → CH₄ |
| 22. VSEPR Theory | Electron pairs arrange to minimize repulsion → predict shape. | NH₃ = pyramidal |
| 23. Molecular Orbital (MO) | Orbitals formed by linear combination of atomic orbitals. | H₂ MO diagram |
| 24. Bonding MO | Constructive overlap → stabilizing orbital. | σ(1s) in H₂ |
| 25. Antibonding MO | Destructive overlap → destabilizing orbital. | σ*(1s) |
| 26. Bond Polarity | Unequal sharing of electrons due to electronegativity difference. | H–Cl |
| 27. Dipole Moment | Measure of bond polarity (charge × distance). | H₂O = 1.84 D |
| 28. Fajan’s Rule | Predicts covalent character in ionic compounds. | LiI more covalent than NaI |
| 29. Hydrogen Bond | Attraction between H (bonded to N, O, F) & another electronegative atom. | H₂O, HF |
| 30. Metallic Bond | Attraction between metal cations & sea of delocalized electrons. | Cu, Fe |
| 31. Isoelectronic Species | Species with same number of electrons but different nuclei. | O²⁻, F⁻, Ne |
| 32. Formal Charge | Hypothetical charge if bonded electrons are equally shared. | O in NO₃⁻ |
| 33. Polar Covalent Bond | Unequal sharing of electrons. | H–Cl |
| 34. Non-Polar Covalent Bond | Equal sharing of electrons. | H₂, Cl₂ |
| 35. Electrovalency | Number of electrons lost/gained in ionic bonding. | Na⁺ = 1 |
| 36. Covalency | Number of electron pairs shared in covalent bonding. | C in CH₄ = 4 |
| 37. Hybrid Orbitals | New equivalent orbitals after hybridization. | sp² → BF₃ |
| 38. Interstitial Compounds | Small atoms occupy interstitial positions in metals. | Fe₃C, TiH₂ |
| 39. Crystal Field Theory (CFT) | Explains splitting of d-orbitals in metal complexes. | [Cu(H₂O)₆]²⁺ |
| 40. Ligand | Ion/molecule donating electron pair to central atom. | NH₃, H₂O |
| 41. Coordination Number | Number of ligands attached to central atom. | [Co(NH₃)₆]³⁺ → 6 |
| 42. Bond Energy | Average energy required to break a bond. | C–C = 348 kJ/mol |
| 43. Delocalized Electrons | Electrons not confined to a bond; spread across atoms. | Benzene π-electrons |
| 44. Band Theory | Metallic bonding explained by overlapping orbitals forming bands. | Valence vs Conduction band |
| 45. Crystal Lattice | Regular 3D arrangement of ions/atoms/molecules in a solid. | NaCl structure |
