Chemical Bonding and Molecular Structure

• Kössel-Lewis Approach & Octet Rule

  • Octet Rule: Atoms lose, gain, or share electrons to attain 8 valence electrons
  • Lewis Symbols: Representations showing valence electrons as dots
  • Limitations of Octet Rule
    • Incomplete Octet: e.g., $\text{LiCl}$, $\text{BF}_3$, $\text{BeH}_2$
    • Odd-Electron Molecules: e.g., $\text{NO}$, $\text{NO}_2$
    • Expanded Octet: e.g., $\text{PCl}_5$, $\text{SF}_6$, $\text{H}_2\text{SO}_4$

• Types of Chemical Bonds

  • Ionic or Electrovalent Bond
    • Transfer of electrons from metal to non-metal
    • Favored by: Low Ionization Enthalpy, high negative Electron Gain Enthalpy, high Lattice Enthalpy
  • Covalent Bond
    • Mutual sharing of electron pairs
    • Can be single, double, or triple bonds
  • Coordinate (Dative) Bond
    • Shared pair of electrons donated by a single atom

• Bond Parameters

  • Bond Length: Equilibrium distance between nuclei of two bonded atoms
  • Bond Angle: Angle between orbitals containing bonding electron pairs
  • Bond Enthalpy: Energy required to break one mole of bonds
  • Bond Order: Number of bonds between two atoms
  • Resonance: Blending of multiple Lewis structures to represent a single molecule (e.g., $\text{O}_3$)
  • Polarity of Bonds
    • Dipole Moment ($\mu$): $\mu = \text{Charge } (Q) \times \text{Distance } (r)$
    • Non-polar ($\mu = 0$) vs. Polar ($\mu \neq 0$)
    • Fajan's Rules: Partial covalent character in ionic bonds (favored by small cation, large anion, high charges)

• VSEPR Theory

  • Core Concept: Geometry depends on repulsion between valence shell electron pairs
  • Repulsion Order: Lone Pair - Lone Pair > Lone Pair - Bond Pair > Bond Pair - Bond Pair
  • Molecular Geometries
    • Linear: e.g., $\text{BeCl}_2$
    • Trigonal Planar: e.g., $\text{BF}_3$
    • Tetrahedral: e.g., $\text{CH}_4$
    • Trigonal Pyramidal: e.g., $\text{NH}_3$ (1 Lone Pair)
    • Bent / V-Shaped: e.g., $\text{H}_2\text{O}$ (2 Lone Pairs)

• Valence Bond Theory (VBT)

  • Orbital Overlap Concept: Bond forms by overlap of half-filled atomic orbitals
  • Types of Overlap & Bonds
    • Sigma ($\sigma$) Bond: Axial/Head-on overlap (Stronger)
    • Pi ($\pi$) Bond: Lateral/Sidewise overlap (Weaker)

• Hybridization

  • Concept: Mixing of atomic orbitals to form equivalent hybrid orbitals
  • Types and Geometries
    • $sp$: Linear ($180^\circ$), e.g., $\text{BeCl}_2$, $\text{C}_2\text{H}_2$
    • $sp^2$: Trigonal Planar ($120^\circ$), e.g., $\text{BF}_3$, $\text{C}_2\text{H}_4$
    • $sp^3$: Tetrahedral ($109.5^\circ$), e.g., $\text{CH}_4$, $\text{NH}_3$, $\text{H}_2\text{O}$
    • $sp^3d$: Trigonal Bipyramidal, e.g., $\text{PCl}_5$
    • $sp^3d^2$: Octahedral, e.g., $\text{SF}_6$

• Molecular Orbital Theory (MOT)

  • Linear Combination of Atomic Orbitals (LCAO)
  • Types of Molecular Orbitals
    • Bonding MO (BMO): Lower energy, higher stability
    • Antibonding MO (ABMO): Higher energy, lower stability
  • Key Formulas & Concepts
    • Bond Order: $\text{BO} = \frac{1}{2}(N_b - N_a)$
    • Stability: Stable if $N_b > N_a$
  • Magnetic Behavior
    • Paramagnetic: Presence of unpaired electrons
    • Diamagnetic: All electrons paired

• Hydrogen Bonding

  • Condition: H atom bonded to highly electronegative atoms ($\text{F}$, $\text{O}$, $\text{N}$)
  • Types
    • Intermolecular H-Bond: Between different molecules (e.g., $\text{HF}$, $\text{H}_2\text{O}$)
    • Intramolecular H-Bond: Within the same molecule (e.g., o-nitrophenol)