CH#3 Theories of Covalent Bonding Class 11

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  1. Lewis structures show electron sharing but do not explain molecular shape or bond energy.
  2. Covalent bonding is explained by Valence Bond Theory (VBT) and Molecular Orbital Theory (MOT).
  3. Valence Bond Theory explains covalent bonds by overlap of atomic orbitals.
  4. A covalent bond forms by overlap of half-filled orbitals.
  5. Greater overlap = stronger covalent bond.
  6. Sigma (σ) bond is formed by head-on overlap of orbitals.
  7. Pi (π) bond is formed by side-by-side overlap of parallel p-orbitals.
  8. Sigma bonds are stronger than pi bonds.
  9. H₂ molecule forms by s–s overlap, producing a σ bond.
  10. HF molecule forms by s–p overlap, producing a σ bond.
  11. A double bond consists of one σ and one π bond.
  12. A triple bond consists of one σ and two π bonds.
  13. Valence Bond Theory fails to explain oxygen’s paramagnetism.
  14. Molecular Orbital Theory treats electrons as delocalized over the molecule.
  15. Bonding molecular orbitals have lower energy than atomic orbitals.
  16. Antibonding molecular orbitals have higher energy than atomic orbitals.
  17. Bond order = (bonding − antibonding electrons) / 2.
  18. H₂ molecule has bond order 1 and is diamagnetic.
  19. O₂ molecule has bond order 2 and is paramagnetic.
  20. VSEPR theory predicts shape by minimizing electron pair repulsion.
  21. Lone pair–lone pair repulsion is strongest.
  22. Multiple bonds are treated as one electron pair in VSEPR theory.
  23. AX₂ molecules like CO₂ are linear (180°).
  24. AX₃ molecules like BF₃ are trigonal planar (120°).
  25. AX₄ molecules like CH₄ are tetrahedral (109.5°).
  26. Hybridization explains actual shape and valency of molecules.
  27. sp³ hybridization produces four equivalent tetrahedral orbitals.
  28. Methane (CH₄) shows sp³ hybridization and tetrahedral geometry.
  29. Ammonia (NH₃) is sp³ hybridized with pyramidal shape.
  30. Water (H₂O) is sp³ hybridized with bent geometry.
  31. sp² hybridization results in trigonal planar geometry.
  32. Ethene (C₂H₄) has one σ and one π bond between carbons.
  33. sp hybridization produces linear geometry (180°).
  34. Ethyne (C₂H₂) contains one σ and two π bonds.
  35. Bond energy is the energy required to break one mole of a bond.
  36. Higher bond energy indicates a stronger bond.
  37. Multiple bonds have higher bond energy than single bonds.
  38. Smaller atoms form shorter and stronger bonds.
  39. Bond length is the distance between bonded nuclei.
  40. Shorter bond length = stronger bond.
  41. Polar covalent bonds arise from electronegativity difference.
  42. Greater electronegativity difference means greater ionic character.
  43. Dipole moment measures bond or molecular polarity.
  44. Molecular polarity depends on bond polarity and shape.
  45. CO₂ is non-polar due to linear shape and dipole cancellation.
  46. H₂O is polar due to bent shape and net dipole moment.
  47. Polar substances dissolve in polar solvents.
  48. Non-polar substances dissolve in non-polar solvents.
  49. Stronger bonds require more energy to break.
  50. A reaction occurs when energy released exceeds energy absorbed