📝 Chapter Notes & Revision

Chemical Bonding and Molecular Structure

🏫 MP BoardClass 11Chemistry

📐 Formula & Cheat Sheet (English)

MP Board Class 11 Chemistry

Chapter 4: Chemical Bonding and Molecular Structure

Quick Revision Notes & Formula Sheet

1. Kössel-Lewis Approach & Octet Rule (अष्टक नियम)

  • Octet Rule: Atoms combine either by transfer of valence electrons (gaining or losing) or by sharing of valence electrons in order to attain a stable octet (8 electrons in their valence shell).
  • Exceptions to Octet Rule:
    • Incomplete octet of central atom: $LiCl$, $BeH_2$, $BF_3$ (Central atom has $< 8$ electrons).
    • Odd-electron molecules: $NO$, $NO_2$ (Contain unpaired electrons).
    • Expanded octet: $PF_5$, $SF_6$, $H_2SO_4$ (Central atom has $> 8$ electrons due to presence of vacant d-orbitals).

2. Formal Charge (फॉर्मल आवेश)

Formal charge helps in selecting the lowest energy (most stable) structure among different Lewis structures.

$$\text{Formal Charge (FC)} = V - L - \frac{1}{2} S$$

Where:

  • $V$ = Total number of valence electrons in free atom
  • $L$ = Total number of non-bonding electrons (lone pair electrons)
  • $S$ = Total number of bonding (shared) electrons

3. Ionic or Electrovalent Bond

Formed by complete transfer of electrons from one atom to another.

  • Factors favoring Ionic Bond formation:
    1. Low Ionization Enthalpy of cation-forming metal.
    2. High negative Electron Gain Enthalpy of anion-forming non-metal.
    3. High Lattice Enthalpy ($\Delta_{lattice}H$) of formed ionic compound.

4. Bond Parameters

  1. Bond Length: Equilibrium distance between nuclei of two bonded atoms.
  2. Bond Angle: Angle between orbitals containing bonding electron pairs around central atom.
  3. Bond Enthalpy: Amount of energy required to break one mole of bonds of a particular type between two atoms in a gaseous state.
    • $\text{Bond Enthalpy} \propto \text{Bond Order} \propto \frac{1}{\text{Bond Length}}$
  4. Bond Order (आबंध कोटि): Number of bonds between two atoms in a molecule.
    • Isoelectronic molecules/ions have same bond order (e.g., $F_2$ and $O_2^{2-}$ have Bond Order = 1).

5. Dipole Moment ($\mu$) & Fajan's Rule

Dipole Moment (द्विध्रुव आघूर्ण)

Measures the polarity of a covalent bond. $$\mu = q \times d$$

Where:

  • $q$ = Magnitude of charge
  • $d$ = Distance of separation
  • Unit: Debye ($D$). ($1 \text{ D} = 3.3356 \times 10^{-30} \text{ C}\cdot\text{m}$)
  • Net Dipole Moment ($\mu_{net}$) for polyatomic molecules:
    • Non-polar molecules (Symmetrical): $\mu_{net} = 0$ (e.g., $CO_2, BF_3, CCl_4$)
    • Polar molecules (Asymmetrical): $\mu_{net} \neq 0$ (e.g., $H_2O, NH_3$)
    • Note: Dipole moment of $NH_3 > NF_3$ because orbital dipole of lone pair in $NF_3$ opposes the resultant dipole of $N-F$ bonds.

Fajan's Rules (Covalent character in Ionic Bonds)

  • Covalent Character is maximum when:
    1. Size of Cation is small.
    2. Size of Anion is large.
    3. Charge on Cation or Anion is high.
    4. Pseudo-noble gas configuration of cation ($ns^2 np^6 nd^{10}$).

6. VSEPR Theory (Valence Shell Electron Pair Repulsion)

Predicts the 3D shape of molecules based on electron pair repulsions.

  • Order of Repulsion: $$\text{Lone Pair - Lone Pair (LP-LP)} > \text{Lone Pair - Bond Pair (LP-BP)} > \text{Bond Pair - Bond Pair (BP-BP)}$$

Geometry vs. Shape Table:

Electron PairsBond PairsLone PairsHybridizationGeometryShapeExamples
220$sp$LinearLinear$BeCl_2, CO_2$
330$sp^2$Trigonal PlanarTrigonal Planar$BF_3, BCl_3$
321$sp^2$Trigonal PlanarBent / V-shape$SO_2, PbCl_2$
440$sp^3$TetrahedralTetrahedral$CH_4, NH_4^+$
431$sp^3$TetrahedralTrigonal Pyramidal$NH_3, H_3O^+$
422$sp^3$TetrahedralBent / V-shape$H_2O$
550$sp^3d$Trigonal BipyramidalTrigonal Bipyramidal$PCl_5$
541$sp^3d$Trigonal BipyramidalSee-Saw$SF_4$
532$sp^3d$Trigonal BipyramidalT-shaped$ClF_3$
523$sp^3d$Trigonal BipyramidalLinear$XeF_2$
660$sp^3d^2$OctahedralOctahedral$SF_6$
651$sp^3d^2$OctahedralSquare Pyramidal$BrF_5$
642$sp^3d^2$OctahedralSquare Planar$XeF_4$

7. Valence Bond Theory (VBT) & Hybridization

Types of Covalent Bonds:

  1. Sigma ($\sigma$) Bond: Formed by end-to-end (axial) overlap of atomic orbitals. Stronger bond.
  2. Pi ($\pi$) Bond: Formed by lateral (sideways) overlap of atomic orbitals. Weaker bond.
    • Single bond = $1\sigma$
    • Double bond = $1\sigma + 1\pi$
    • Triple bond = $1\sigma + 2\pi$

Shortcut Formula to find Hybridization Number ($H$):

$$H = \frac{1}{2} \left[ V + M - C + A \right]$$

Where:

  • $V$ = Valence electrons on central atom

  • $M$ = Number of monovalent atoms attached (H, F, Cl, Br, I)

  • $C$ = Positive charge on cation

  • $A$ = Negative charge on anion

  • Hybridization Determination:

    • $H = 2 \rightarrow sp$
    • $H = 3 \rightarrow sp^2$
    • $H = 4 \rightarrow sp^3$
    • $H = 5 \rightarrow sp^3d$
    • $H = 6 \rightarrow sp^3d^2$
    • $H = 7 \rightarrow sp^3d^3$

8. Molecular Orbital Theory (MOT)

1. Energy Level Diagrams:

  • For molecules with $\le 14$ electrons ($B_2, C_2, N_2$): $$\sigma 1s < \sigma^* 1s < \sigma 2s < \sigma^* 2s < (\pi 2p_x = \pi 2p_y) < \sigma 2p_z < (\pi^* 2p_x = \pi^* 2p_y) < \sigma^* 2p_z$$

  • For molecules with $> 14$ electrons ($O_2, F_2, Ne_2$): $$\sigma 1s < \sigma^* 1s < \sigma 2s < \sigma^* 2s < \sigma 2p_z < (\pi 2p_x = \pi 2p_y) < (\pi^* 2p_x = \pi^* 2p_y) < \sigma^* 2p_z$$

2. Key Formulas in MOT:

  • Bond Order (BO): $$\text{Bond Order} = \frac{1}{2} (N_b - N_a)$$ (where $N_b$ = Bonding electrons, $N_a$ = Anti-bonding electrons)

  • Stability Criteria:

    • If $N_b > N_a \rightarrow$ Molecule is Stable ($\text{BO} > 0$)
    • If $N_b \le N_a \rightarrow$ Molecule is Unstable ($\text{BO} \le 0$)

  • Magnetic Character:

    • If all MOs have paired electrons $\rightarrow$ Diamagnetic (प्रतिचुंबकीय)
    • If one or more MOs have unpaired electrons $\rightarrow$ Paramagnetic (अनुचुंबकीय) (e.g., $O_2, B_2$ are paramagnetic)

9. Hydrogen Bonding (हाइड्रोजन आबंध)

Formed when $H$ atom is covalently bonded to a highly electronegative atom ($F, O, N$).

  1. Intermolecular H-Bonding: Formed between two different molecules of same or different compounds.
    • Examples: $H_2O, HF, NH_3$, alcohol in water.
    • Effect: Increases Boiling Point and Viscosity.
  2. Intramolecular H-Bonding: Formed within the same molecule.
    • Examples: o-Nitrophenol, Salicylaldehyde.
    • Effect: Lowers Boiling Point compared to Intermolecular counterpart (e.g., o-nitrophenol has lower B.P. than p-nitrophenol).