Class 11 Chemistry Notes Chapter 4 (Chapter 4) – Examplar Problems (English) Book

Alright class, let's dive deep into Chapter 4, 'Chemical Bonding and Molecular Structure'. This chapter is absolutely fundamental, not just for your Class 11 understanding, but it forms the bedrock for much of chemistry you'll encounter later, especially for competitive government exams. Pay close attention, as understanding bonding is key to understanding the properties and reactions of substances.

Chapter 4: Chemical Bonding and Molecular Structure - Detailed Notes

1. Introduction: Why do Atoms Combine?

• Goal: Atoms combine to achieve a more stable electronic configuration, usually resembling that of the nearest noble gas (ns²np⁶ configuration - the Octet Rule).
• Stability: Formation of a chemical bond leads to a decrease in the potential energy of the system, resulting in stability.
• Kossel-Lewis Approach: Explained bond formation based on achieving stable noble gas configurations.
  • Kossel: Explained ionic bonding via electron transfer.
  • Lewis: Explained covalent bonding via electron sharing, introducing Lewis Symbols (element symbol surrounded by valence electrons as dots).

2. Types of Chemical Bonds

• Ionic (Electrovalent) Bond:

  • Formation: Complete transfer of one or more valence electrons from a highly electropositive atom (metal) to a highly electronegative atom (non-metal). Results in the formation of oppositely charged ions (cation and anion) held by electrostatic attraction.
  • Example: NaCl (Na → Na⁺ + e⁻ ; Cl + e⁻ → Cl⁻ ; Na⁺ + Cl⁻ → NaCl)
  • Factors Favoring Formation:
    • Low Ionization Enthalpy (IE) of the metal.
    • High Electron Gain Enthalpy (EGE) (more negative) of the non-metal.
    • High Lattice Enthalpy (Energy released when one mole of ionic compound is formed from its gaseous ions).
  • Lattice Enthalpy: A measure of the strength of the ionic bond. Can be estimated using the Born-Haber Cycle.
  • Properties of Ionic Compounds: Generally crystalline solids, high melting and boiling points, soluble in polar solvents (like water), insulators in solid state but conductors in molten state or aqueous solution.

• Covalent Bond:

  • Formation: Mutual sharing of one or more pairs of electrons between two atoms (usually non-metals) to achieve stable octets (or duplets for H, Li, Be).
  • Types: Single (sharing 1 pair, e.g., H₂), Double (sharing 2 pairs, e.g., O₂), Triple (sharing 3 pairs, e.g., N₂).
  • Lewis Structures: Diagrams representing shared pairs (bond pairs, shown as lines) and unshared valence electrons (lone pairs, shown as dots) around atoms in a molecule or ion.
  • Formal Charge (FC): A hypothetical charge assigned to an atom in a molecule, assuming equal sharing of bonded electrons.
    • FC = [Total valence e⁻ in free atom] - [Total non-bonding (lone pair) e⁻] - ½ [Total bonding (shared) e⁻]
    • Helps select the most stable Lewis structure (lowest FCs, negative FC on more electronegative atom).
  • Limitations of Octet Rule:
    • Incomplete Octet: Central atom has < 8 valence e⁻ (e.g., LiCl, BeH₂, BF₃).
    • Odd-Electron Molecules: Molecules with an odd number of total valence e⁻ (e.g., NO, NO₂).
    • Expanded Octet: Central atom has > 8 valence e⁻ (possible for elements in period 3 and beyond due to available d-orbitals, e.g., PCl₅, SF₆, IF₇).

• Coordinate (Dative) Bond:

  • A special type of covalent bond where the shared pair of electrons is contributed by only one of the bonded atoms (donor), while the other atom (acceptor) only shares it. Represented by an arrow (→) from donor to acceptor.
  • Example: Formation of NH₄⁺ (N in NH₃ donates lone pair to H⁺), BF₃ ← NH₃.

3. Bond Parameters

• Bond Length: Average equilibrium distance between the nuclei of two bonded atoms. Factors: Size of atoms (larger atoms → longer bond), Bond Multiplicity (Single > Double > Triple). Measured in Å or pm.
• Bond Angle: Angle between orbitals containing bonding electron pairs around the central atom. Influenced by hybridization and repulsion between electron pairs (VSEPR).
• Bond Enthalpy: Energy required to break one mole of a specific type of bond between atoms in the gaseous state. Factors: Size of atoms (smaller atoms → stronger bond), Bond Multiplicity (Triple > Double > Single), Lone pair repulsion (can weaken bonds). Unit: kJ mol⁻¹. Higher bond enthalpy indicates a stronger bond.
• Bond Order: Number of chemical bonds between a pair of atoms. (e.g., H-H: 1, O=O: 2, N≡N: 3). For resonance structures or MOT, it can be fractional. Higher bond order implies shorter bond length and higher bond enthalpy.

4. Polarity of Covalent Bonds & Dipole Moment

• Electronegativity (EN): Tendency of an atom in a chemical bond to attract the shared pair of electrons towards itself. F > O > N > Cl > Br > I > S > C > H.
• Polar Covalent Bond: Formed between two dissimilar atoms with different electronegativities. The shared pair is shifted towards the more electronegative atom, creating partial positive (δ+) and partial negative (δ-) charges. Example: H-Cl (Cl is δ-, H is δ+).
• Non-polar Covalent Bond: Formed between identical atoms (e.g., H₂, Cl₂) or atoms with nearly identical EN, where electron pair is shared equally.
• Dipole Moment (µ): Measure of the polarity of a chemical bond or molecule. It's the product of the magnitude of charge (q) and the distance of separation (d) between the charges. µ = q × d.
  • Vector quantity, represented by an arrow pointing from positive to negative end (→).
  • Unit: Debye (D). 1 D = 3.33564 × 10⁻³⁰ C m.
  • Molecular Dipole Moment: Vector sum of individual bond dipoles.
    • Diatomic Molecules: µ depends on the EN difference.
    • Polyatomic Molecules: Depends on both bond polarity and molecular geometry. Symmetrical molecules (like BF₃, CCl₄, CO₂) can have polar bonds but zero net dipole moment due to cancellation. Asymmetrical molecules (like H₂O, NH₃) have a net dipole moment.
• Percentage Ionic Character: Can be estimated from the observed dipole moment and the theoretical dipole moment assuming 100% ionic character. % Ionic Character = (Observed µ / Calculated µ for 100% ionic) × 100.
• Fajan's Rules: Predict the degree of covalent character in an ionic bond. Covalent character is favored by:
  • Small cation size.
  • Large anion size.
  • High charge on cation and/or anion.
  • Cations with pseudo noble gas configuration (e.g., Cu⁺) have higher polarizing power than those with noble gas configuration (e.g., Na⁺).

5. Theories of Covalent Bonding and Molecular Shape

• Valence Shell Electron Pair Repulsion (VSEPR) Theory:

  • Premise: Electron pairs (both bonding and lone pairs) in the valence shell of a central atom repel each other and arrange themselves in space to minimize repulsion and maximize distance.
  • Order of Repulsion: Lone pair-Lone pair (lp-lp) > Lone pair-Bond pair (lp-bp) > Bond pair-Bond pair (bp-bp).
  • Predicting Geometry:
    1. Draw the Lewis structure.
    2. Count total electron pairs (bond pairs + lone pairs) around the central atom.
    3. Arrange electron pairs to minimize repulsion (gives electron geometry).
    4. Determine the molecular shape based only on the positions of the atoms (ignoring lone pairs for the final shape name).
  • Examples:

    Electron Pairs	Arrangement (Electron Geometry)	Lone Pairs	Shape (Molecular Geometry)	Example	Hybridization
    2	Linear	0	Linear	BeCl₂, CO₂	sp
    3	Trigonal Planar	0	Trigonal Planar	BF₃, SO₃	sp²
    3	Trigonal Planar	1	Bent / V-shape	SO₂	sp²
    4	Tetrahedral	0	Tetrahedral	CH₄, SiCl₄	sp³
    4	Tetrahedral	1	Trigonal Pyramidal	NH₃, PCl₃	sp³
    4	Tetrahedral	2	Bent / V-shape	H₂O, H₂S	sp³
    5	Trigonal Bipyramidal	0	Trigonal Bipyramidal	PCl₅	sp³d
    5	Trigonal Bipyramidal	1	See-Saw	SF₄	sp³d
    5	Trigonal Bipyramidal	2	T-shaped	ClF₃	sp³d
    5	Trigonal Bipyramidal	3	Linear	XeF₂, I₃⁻	sp³d
    6	Octahedral	0	Octahedral	SF₆	sp³d²
    6	Octahedral	1	Square Pyramidal	BrF₅	sp³d²
    6	Octahedral	2	Square Planar	XeF₄	sp³d²

• Valence Bond Theory (VBT):

  • Premise: A covalent bond forms by the overlap of half-filled atomic orbitals of valence shells of the two atoms. The electrons in the overlapping orbitals get paired.
  • Types of Overlap:
    • Sigma (σ) Bond: Formed by head-on (axial) overlap of orbitals (s-s, s-p, p-p). Stronger bond. Allows free rotation. First bond formed between two atoms is always sigma.
    • Pi (π) Bond: Formed by sideways (lateral) overlap of p-orbitals (p-p). Weaker bond. Restricts rotation. Formed in addition to a sigma bond in multiple bonds (double, triple).
  • Hybridization: Concept introduced by Pauling to explain observed shapes and bond angles. It's the mixing of atomic orbitals of slightly different energies on the same atom to form a new set of equivalent energy hybrid orbitals with specific directional characteristics.
    • Types: sp (Linear), sp² (Trigonal Planar), sp³ (Tetrahedral), sp³d (Trigonal Bipyramidal), sp³d² (Octahedral).
    • Predicting Hybridization: Number of hybrid orbitals = (Number of sigma bonds) + (Number of lone pairs) around the central atom. Or use formula: H = ½ [V + M - C + A], where V=valence e⁻ of central atom, M=no. of monovalent atoms attached, C=charge on cation, A=charge on anion.

• Molecular Orbital Theory (MOT):

  • Premise: Atomic orbitals combine to form an equal number of molecular orbitals (MOs) spread over the entire molecule. Electrons fill these MOs according to Aufbau, Pauli, and Hund's rules.
  • Linear Combination of Atomic Orbitals (LCAO): Method used to form MOs.
  • Types of MOs:
    • Bonding Molecular Orbitals (BMOs): Lower energy, increased electron density between nuclei, stabilizes the molecule (e.g., σ, π).
    • Antibonding Molecular Orbitals (ABMOs): Higher energy, node between nuclei, destabilizes the molecule (e.g., σ*, π*).
  • Energy Level Diagrams: Different for molecules up to N₂ (due to s-p mixing) and for O₂, F₂, Ne₂.
    • For B₂, C₂, N₂: σ1s < σ1s < σ2s < σ2s < (π2px = π2py) < σ2pz < (π2px = π2py) < σ*2pz
    • For O₂, F₂, Ne₂: σ1s < σ1s < σ2s < σ2s < σ2pz < (π2px = π2py) < (π2px = π2py) < σ*2pz
  • Applications:
    • Bond Order (B.O.): B.O. = ½ [Number of electrons in BMOs (Nb) - Number of electrons in ABMOs (Na)].
      • B.O. > 0: Molecule exists and is stable.
      • B.O. = 0: Molecule does not exist or is unstable (e.g., He₂, Ne₂).
      • Higher B.O. → Greater stability, Shorter bond length, Higher bond dissociation energy.
    • Magnetic Nature:
      • Paramagnetic: Contains unpaired electrons in MOs (attracted by magnetic field). Example: O₂, B₂.
      • Diamagnetic: All electrons in MOs are paired (repelled by magnetic field). Example: N₂, F₂, C₂.
    • Explains properties VBT cannot (e.g., paramagnetism of O₂, existence of ions like H₂⁺, He₂⁺).

6. Hydrogen Bonding

• Definition: An attractive force between a hydrogen atom covalently bonded to a highly electronegative atom (F, O, or N) and another nearby highly electronegative atom (F, O, or N). It's a special type of dipole-dipole interaction, stronger than van der Waals forces but weaker than covalent/ionic bonds. Represented by dotted lines (---).
• Conditions: H must be bonded to F, O, or N; the other electronegative atom must be small and highly electronegative (usually F, O, or N).
• Types:
  • Intermolecular H-bonding: Occurs between different molecules (e.g., HF, H₂O, NH₃, alcohols, carboxylic acids). Leads to association of molecules.
  • Intramolecular H-bonding: Occurs within the same molecule (e.g., o-nitrophenol, salicylaldehyde). Leads to cyclization.
• Consequences:
  • Abnormally high boiling points and melting points (e.g., H₂O > H₂S).
  • Solubility (alcohols/sugars in water).
  • Lower density of ice compared to water.
  • Structure and stability of proteins and nucleic acids (DNA).
  • Intramolecular H-bonding decreases boiling point and solubility compared to the para isomer (e.g., o-nitrophenol vs p-nitrophenol).

Multiple Choice Questions (MCQs)

1. Which of the following molecules has the highest dipole moment?
   (a) CO₂
   (b) CH₄
   (c) NH₃
   (d) NF₃

2. According to VSEPR theory, the shape of the SF₄ molecule is:
   (a) Tetrahedral
   (b) Square planar
   (c) See-saw
   (d) Trigonal bipyramidal

3. What is the hybridization of the central atom in PCl₅?
   (a) sp²
   (b) sp³
   (c) sp³d
   (d) sp³d²

4. Which of the following species is paramagnetic according to Molecular Orbital Theory?
   (a) N₂
   (b) O₂²⁻ (Peroxide ion)
   (c) O₂
   (d) F₂

5. The correct order of bond order for the species O₂, O₂⁺, O₂⁻, O₂²⁻ is:
   (a) O₂⁺ > O₂ > O₂⁻ > O₂²⁻
   (b) O₂ > O₂⁺ > O₂⁻ > O₂²⁻
   (c) O₂²⁻ > O₂⁻ > O₂ > O₂⁺
   (d) O₂⁻ > O₂²⁻ > O₂⁺ > O₂

6. In which of the following pairs does the first species have a higher boiling point than the second due to hydrogen bonding?
   (a) H₂S, H₂O
   (b) PH₃, NH₃
   (c) HCl, HF
   (d) CH₄, SiH₄

7. Which factor does not favor the formation of an ionic bond?
   (a) High lattice enthalpy
   (b) Low ionization enthalpy of metal
   (c) High electron gain enthalpy of non-metal
   (d) High ionization enthalpy of metal

8. The formal charge on the central oxygen atom in the ozone (O₃) molecule (considering one double and one single bond) is:
   (a) 0
   (b) +1
   (c) -1
   (d) +2

9. Which molecule contains both sigma (σ) and pi (π) bonds?
   (a) CH₄
   (b) H₂O
   (c) C₂H₄ (Ethene)
   (d) NH₃

10. According to Fajan's rules, which compound has the most covalent character?
    (a) NaCl
    (b) MgCl₂
    (c) AlCl₃
    (d) LiCl

Answer Key for MCQs:

1. (c) NH₃ (Pyramidal shape, lone pair adds to bond dipoles) [NF₃ has lower µ due to opposing dipoles]
2. (c) See-saw (Based on trigonal bipyramidal electron geometry with one lone pair in equatorial position)
3. (c) sp³d (5 electron pairs around P)
4. (c) O₂ (Has two unpaired electrons in π* MOs)
5. (a) O₂⁺ (BO=2.5) > O₂ (BO=2.0) > O₂⁻ (BO=1.5) > O₂²⁻ (BO=1.0)
6. (c) HF (Strong H-bonding) > HCl (No H-bonding) [Note: H₂O > H₂S, NH₃ > PH₃ are also correct examples, but the question asks for the pair where the first species is higher]
7. (d) High ionization enthalpy of metal (Makes cation formation difficult)
8. (b) +1 (Structure: O=O⁺-O⁻. Central O: 6 - 2 - ½(6) = +1)
9. (c) C₂H₄ (Contains C=C double bond, which consists of one σ and one π bond)
10. (c) AlCl₃ (Al³⁺ has the highest charge and smallest size among the cations listed, leading to maximum polarization of Cl⁻)

Study these notes thoroughly. Focus on understanding the concepts behind VSEPR, Hybridization, and MOT, as direct application questions are very common. Good luck with your preparation!