Class 11 Chemistry Notes Chapter 4 (Chemical bonding and molecular structure) – Chemistry Part-I Book

Alright class, let's begin our detailed study of Chapter 4: Chemical Bonding and Molecular Structure. This is a cornerstone chapter, absolutely vital for understanding the properties and reactions of substances, and frequently tested in government exams. Pay close attention!

Chapter 4: Chemical Bonding and Molecular Structure - Detailed Notes

1. Introduction: Why do Atoms Combine?

• Tendency to Attain Stability: Atoms combine to achieve a more stable electronic configuration, usually resembling that of the nearest noble gas (ns²np⁶ configuration - Octet Rule).
• Lowering of Energy: When atoms bond, the potential energy of the system decreases, leading to stability.
• Kossel-Lewis Approach:
  • Lewis Symbols: Valence electrons are represented as dots around the element symbol. (e.g., :Ṅ:, :Ö:, ⋅Na)
  • Octet Rule: Atoms tend to gain, lose, or share electrons to achieve 8 electrons in their valence shell. (Exceptions exist, especially for H, Li, Be which aim for a duplet, and elements from Period 3 onwards).

2. Types of Chemical Bonds

A. Ionic (Electrovalent) Bond

• Formation: Formed by the complete transfer of one or more electrons from a highly electropositive atom (metal) to a highly electronegative atom (non-metal). This results in the formation of oppositely charged ions (cations and anions) held together by strong electrostatic forces (Coulombic attraction).
  • Example: Na (2,8,1) + Cl (2,8,7) → Na⁺ (2,8) + Cl⁻ (2,8,8) → NaCl
• Factors Favouring Ionic Bond Formation:
  • Low Ionization Enthalpy (IE): Easily removal of electron(s) from the metal atom.
  • High Negative Electron Gain Enthalpy (EGE): Easy acceptance of electron(s) by the non-metal atom.
  • High Lattice Enthalpy: Large amount of energy released when ions pack together to form the crystal lattice. This is the driving force for stability.
• Properties of Ionic Compounds:
  • Generally crystalline solids at room temperature.
  • High melting and boiling points (due to strong electrostatic forces).
  • Soluble in polar solvents (like water) but insoluble in non-polar solvents (like benzene, CCl₄).
  • Conduct electricity in molten state or aqueous solution (due to mobile ions), but not in solid state.
  • Ionic bonds are non-directional.

B. Covalent Bond

• Formation: Formed by the mutual sharing of one or more electron pairs between combining atoms, usually between non-metals or atoms with similar electronegativity.
  • Single Bond: Sharing of one electron pair (e.g., H₂, Cl₂, CH₄). Represented by a single line (-).
  • Double Bond: Sharing of two electron pairs (e.g., O₂, CO₂). Represented by a double line (=).
  • Triple Bond: Sharing of three electron pairs (e.g., N₂, C₂H₂). Represented by a triple line (≡).
• Lewis Structures: Diagrams representing covalent bonding by showing shared pairs (as lines) and lone pairs (as dots) around each atom, ensuring octets (or duplets for H) are completed.
• Formal Charge (FC): A hypothetical charge assigned to an atom in a molecule, assuming equal sharing of electrons in a covalent bond.
  • FC = [Total valence e⁻ in free atom] - [Total non-bonding (lone pair) e⁻] - ½ [Total bonding (shared) e⁻]
  • Helps in selecting the most stable Lewis structure (structures with lowest FCs, negative FC on more electronegative atom are preferred).
• Limitations of the Octet Rule:
  • Incomplete Octet: Central atom has < 8 valence electrons (e.g., BeCl₂, BF₃).
  • Expanded Octet: Central atom has > 8 valence electrons (possible for elements in Period 3 onwards due to available d-orbitals, e.g., PCl₅, SF₆, IF₇).
  • Odd-Electron Molecules: Molecules with an odd number of total valence electrons cannot satisfy the octet rule for all atoms (e.g., NO, NO₂).

3. Bond Parameters

• Bond Length: The average equilibrium distance between the nuclei of two bonded atoms in a molecule.
  • Factors: Decreases with increasing bond order (triple < double < single), decreases with increasing electronegativity difference, increases with atomic size.
• Bond Angle: The angle between the orbitals containing bonding electron pairs around the central atom in a molecule. Helps determine the shape.
• Bond Enthalpy: The energy required to break one mole of bonds of a particular type between two atoms in a gaseous state.
  • Factors: Increases with bond order, decreases with bond length. Represents bond strength.
• Bond Order: The number of bonds between two atoms in a molecule (1 for single, 2 for double, 3 for triple). In MOT, Bond Order = ½ (Nb - Na). Higher bond order implies greater stability and shorter bond length.

4. Polarity of Covalent Bonds & Dipole Moment

• Non-polar Covalent Bond: Formed between identical atoms (e.g., H₂, Cl₂) or atoms with negligible electronegativity difference. Shared electron pair is equally attracted.
• Polar Covalent Bond: Formed between atoms with significant electronegativity difference (e.g., HCl, H₂O). Shared electron pair is attracted more towards the more electronegative atom, creating partial positive (δ+) and partial negative (δ-) charges.
• Dipole Moment (µ): A measure of the polarity of a chemical bond or molecule. It's a vector quantity.
  • µ = q × d (where q = magnitude of charge separation, d = distance between charges).
  • Unit: Debye (D). 1 D = 3.33564 × 10⁻³⁰ C m.
  • Molecular Dipole Moment: The vector sum of individual bond dipoles.
    • Symmetrical molecules (like CO₂, BF₃, CCl₄, CH₄) have zero net dipole moment (µ=0) even if bonds are polar, because bond dipoles cancel out.
    • Unsymmetrical molecules (like H₂O, NH₃, HCl, CHCl₃) have a net dipole moment (µ≠0).
  • Dipole moment helps predict molecular geometry and the degree of polarity.
• Percentage Ionic Character: Can be estimated from the observed dipole moment and the calculated dipole moment assuming 100% ionic character.

5. Theories of Covalent Bonding & Molecular Shapes

A. VSEPR Theory (Valence Shell Electron Pair Repulsion Theory)

• Premise: The shape of a molecule is determined by repulsions between electron pairs (both bonding and lone pairs) in the valence shell of the central atom. These pairs arrange themselves to minimize repulsion and maximize distance.

• Order of Repulsion: Lone pair-Lone pair (lp-lp) > Lone pair-Bonding pair (lp-bp) > Bonding pair-Bonding pair (bp-bp).

• Predicting Shapes:

  1. Draw the Lewis structure.
  2. Count total electron pairs (bonding pairs + lone pairs) around the central atom.
  3. Determine the electron geometry based on the total number of electron pairs (Linear, Trigonal Planar, Tetrahedral, Trigonal Bipyramidal, Octahedral).
  4. Determine the molecular shape (actual arrangement of atoms) considering the effect of lone pairs on geometry. Lone pairs occupy more space and cause greater repulsion, distorting bond angles.

  Total e⁻ pairs	Bonding Pairs	Lone Pairs	Electron Geometry	Molecular Shape	Example	Approx. Bond Angle
  2	2	0	Linear	Linear	BeCl₂, CO₂	180°
  3	3	0	Trigonal Planar	Trigonal Planar	BF₃, SO₃	120°
  3	2	1	Trigonal Planar	Bent / V-shape	SO₂, O₃	<120°
  4	4	0	Tetrahedral	Tetrahedral	CH₄, SiCl₄	109.5°
  4	3	1	Tetrahedral	Trigonal Pyramidal	NH₃, PCl₃	<109.5° (~107°)
  4	2	2	Tetrahedral	Bent / V-shape	H₂O, H₂S	<109.5° (~104.5°)
  5	5	0	Trigonal Bipyramidal	Trigonal Bipyramidal	PCl₅	90°, 120°
  5	4	1	Trigonal Bipyramidal	See-Saw	SF₄	<90°, <120°
  5	3	2	Trigonal Bipyramidal	T-shape	ClF₃	~90°
  5	2	3	Trigonal Bipyramidal	Linear	XeF₂, I₃⁻	180°
  6	6	0	Octahedral	Octahedral	SF₆	90°
  6	5	1	Octahedral	Square Pyramidal	BrF₅, XeOF₄	<90°
  6	4	2	Octahedral	Square Planar	XeF₄, [ICl₄]⁻	90°

B. Valence Bond Theory (VBT)

• Premise: A covalent bond is formed by the overlap of half-filled atomic orbitals of valence shells of the combining atoms. The overlapping orbitals must have electrons with opposite spins.
• Types of Overlap & Bonds:
  • Sigma (σ) Bond: Formed by head-on (axial) overlap of orbitals (s-s, s-p, p-p axial). Stronger bond. Allows free rotation around the bond axis. All single bonds are σ bonds.
  • Pi (π) Bond: Formed by sideways (lateral) overlap of p-orbitals (p-p lateral). Weaker bond. Restricts rotation. Occurs in multiple bonds (double bond = 1σ + 1π; triple bond = 1σ + 2π).
• Hybridization: The concept of mixing atomic orbitals of slightly different energies to form a new set of equivalent orbitals (hybrid orbitals) with the same energy, shape, and size.
  • Need: To explain observed bond angles and equivalent bond lengths in molecules like CH₄, BF₃ etc., which cannot be explained by simple orbital overlap.
  • Types of Hybridization:
    • sp: Mixing one s + one p → two sp orbitals. Linear geometry (180°). (e.g., BeCl₂, C₂H₂)
    • sp²: Mixing one s + two p → three sp² orbitals. Trigonal planar geometry (120°). (e.g., BF₃, C₂H₄)
    • sp³: Mixing one s + three p → four sp³ orbitals. Tetrahedral geometry (109.5°). (e.g., CH₄, NH₃, H₂O)
    • sp³d: Mixing one s + three p + one d → five sp³d orbitals. Trigonal bipyramidal geometry (90°, 120°). (e.g., PCl₅, SF₄)
    • sp³d²: Mixing one s + three p + two d → six sp³d² orbitals. Octahedral geometry (90°). (e.g., SF₆, XeF₄)
  • Predicting Hybridization: Use the VSEPR method: Hybridization = Number of σ bonds + Number of lone pairs around the central atom. (2=sp, 3=sp², 4=sp³, 5=sp³d, 6=sp³d²)

C. Molecular Orbital Theory (MOT)

• Premise: Atomic orbitals combine to form an equal number of molecular orbitals (MOs) which belong to the molecule as a whole. Electrons fill these MOs according to Aufbau principle, Pauli exclusion principle, and Hund's rule.
• Linear Combination of Atomic Orbitals (LCAO): Atomic orbitals combine constructively (forming Bonding Molecular Orbitals - BMOs) and destructively (forming Anti-bonding Molecular Orbitals - ABMOs).
  • BMOs: Lower energy, higher stability, increased electron density between nuclei. (σ, π)
  • ABMOs: Higher energy, lower stability, node between nuclei. (σ*, π*)
• Energy Level Diagrams: Show the relative energies of MOs. The order of filling for diatomic molecules of Period 2 elements:
  • 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 (Note: σ2pz is lower in energy than π2px/π2py)
• Bond Order (BO): BO = ½ [Number of electrons in BMOs (Nb) - Number of electrons in ABMOs (Na)]
  • BO > 0: Molecule is stable and exists.
  • BO = 0: Molecule is unstable and does not exist (e.g., He₂, Ne₂).
  • Higher BO indicates greater stability and shorter bond length.
• Magnetic Properties:
  • Paramagnetic: Attracted by magnetic field. Have one or more unpaired electrons in MOs (e.g., O₂, B₂).
  • Diamagnetic: Repelled by magnetic field. All electrons in MOs are paired (e.g., N₂, F₂, C₂).
• MOT successfully explains: Paramagnetism of O₂, existence of species like He₂⁺, relative stabilities and bond lengths.

6. Hydrogen Bonding

• Definition: A special type of dipole-dipole interaction occurring when hydrogen is bonded to a highly electronegative atom (F, O, N) and is attracted to another nearby electronegative atom. It's weaker than covalent/ionic bonds but stronger than van der Waals forces. Represented by a dotted line (---).
  • Example: Hδ+---Xδ- (where X = F, O, N)
• Types:
  • Intermolecular H-bonding: Occurs between different molecules (e.g., H₂O, HF, NH₃, alcohols). Leads to association of molecules, resulting in high boiling points, high viscosity, high solubility in water.
  • Intramolecular H-bonding: Occurs within the same molecule (e.g., o-nitrophenol, salicylaldehyde). Leads to chelation (ring formation), decreased boiling point (less association), and decreased water solubility compared to isomers with intermolecular H-bonding (like p-nitrophenol).
• Consequences: Anomalous properties of water (high BP, density maximum at 4°C, structure of ice), structure of proteins and DNA, etc.

Practice MCQs for Government Exams

1. Which of the following molecules has the highest dipole moment?
   (a) CO₂
   (b) CH₄
   (c) NH₃
   (d) BF₃

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

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

4. Which of the following species is paramagnetic according to Molecular Orbital Theory?
   (a) N₂
   (b) O₂
   (c) F₂
   (d) C₂²⁻

5. What is the bond order of N₂⁺?
   (a) 2
   (b) 2.5
   (c) 3
   (d) 1.5

6. Which type of bond is formed by the sideways overlap of p-orbitals?
   (a) Sigma (σ) bond
   (b) Pi (π) bond
   (c) Ionic bond
   (d) Hydrogen bond

7. Which of the following exhibits intramolecular hydrogen bonding?
   (a) Water (H₂O)
   (b) Ethanol (C₂H₅OH)
   (c) o-Nitrophenol
   (d) p-Nitrophenol

8. The correct order of bond angles in H₂O, NH₃, and CH₄ is:
   (a) CH₄ > NH₃ > H₂O
   (b) NH₃ > CH₄ > H₂O
   (c) H₂O > NH₃ > CH₄
   (d) CH₄ > H₂O > NH₃

9. Formation of an ionic bond is favoured by:
   (a) High ionization enthalpy of metal & high electron gain enthalpy of non-metal
   (b) Low ionization enthalpy of metal & high negative electron gain enthalpy of non-metal
   (c) Low ionization enthalpy of metal & low electron gain enthalpy of non-metal
   (d) High ionization enthalpy of metal & low negative electron gain enthalpy of non-metal

10. In which of the following molecules does the central atom have an expanded octet?
    (a) BF₃
    (b) NH₃
    (c) H₂O
    (d) SF₆

Answer Key:

1. (c) NH₃ (Unsymmetrical pyramidal shape with lone pair)
2. (c) See-saw (Based on trigonal bipyramidal geometry with one lone pair in equatorial position)
3. (c) sp³d (5 electron pairs around P)
4. (b) O₂ (Has two unpaired electrons in π* MOs)
5. (b) 2.5 (N₂ has BO=3; N₂⁺ loses one electron from BMO, BO = ½(10-5) = 2.5)
6. (b) Pi (π) bond
7. (c) o-Nitrophenol (H of -OH group forms H-bond with O of adjacent -NO₂ group)
8. (a) CH₄ (109.5°) > NH₃ (~107°) > H₂O (~104.5°) (Due to increasing lone pair repulsion)
9. (b) Low ionization enthalpy of metal & high negative electron gain enthalpy of non-metal (Also needs high lattice enthalpy)
10. (d) SF₆ (Sulphur has 12 electrons in its valence shell)

Make sure you understand the underlying concepts behind each point and MCQ answer. Revise VSEPR shapes, hybridization, and MOT diagrams thoroughly. Good luck with your preparation!