Chemical Bonding - Class 11 Chemistry

Chemical Bonding

Overview: Ionic and Covalent bonds, VSEPR theory, Hybridization, and Molecular Orbital Theory (MOT).

1. Types of Bonds

Ionic Bond: Complete transfer of electrons (Electrovalent). Favorable when IE low (metal) and EA high (non-metal). High Lattice Energy.
Covalent Bond: Sharing of electrons. Directional.
Coordinate Bond: Shared pair contributed by only one atom (Donor-Acceptor).

% Ionic = 16(ΔX) + 3.5(ΔX)² (Hannay Smith)

2. VSEPR Theory

Valence Shell Electron Pair Repulsion theory predicts geometry.

Repulsion order: lp-lp > lp-bp > bp-bp.

sp (BeCl2): Linear (180°).
sp2 (BF3): Trigonal Planar (120°).
sp3 (CH4): Tetrahedral (109.5°).
sp3d (PCl5): Trigonal Bipyramidal.
sp3d2 (SF6): Octahedral.

3. Hybridization

Mixing of atomic orbitals to form new equivalent hybrid orbitals.

Sigma (σ) bond: Head-on overlap. Stronger.
Pi (π) bond: Sideways overlap. Weaker.

H = 1/2 [V + M - C + A]

V=Valance e-, M=Monovalent atoms, C=Cation charge, A=Anion charge.

4. Molecular Orbital Theory (MOT)

Linear Combination of Atomic Orbitals (LCAO).

Bonding MO (σ, π): Lower energy, stable.
Antibonding MO (σ*, π*): Higher energy, unstable.

Bond Order = 1/2 (N_b - N_a)

Positive BO: Stable. Zero/Neg BO: Unstable.

5. Other Concepts

Dipole Moment (μ)

μ = q × d (Debye D)

Measure of polarity. Vector quantity.

Fajan's Rule

Covalent character in ionic bond increases with:

Small cation size.
Large anion size.
High charge on cation/anion.
Pseudo noble gas config of cation (e.g. Cu+).

Numericals - Chemical Bonding

Numericals

Octet Rule

Q1. Draw Lewis structure of CO. Evaluate Formal Charge on C.

:C ≡ O:

FC = V - L - 1/2 S

FC on C = 4 - 2 - 1/2(6) = 4 - 2 - 3 = -1.

(Note: C has negative charge in CO).

Dipole Moment

Q2. Resultant dipole moment of water (μ=1.85D, angle=104.5).

Vector sum of two O-H bond moments.

μ_R = 2 x cos(θ/2).

Net μ is non-zero (Polar).

Hybridization

Q3. Determine hybridization of XeF4.

H = 1/2 (V + M - C + A)

V=8 (Xe), M=4 (F).

H = 1/2 (8 + 4) = 6.

Hybridization = sp3d2.

Structure: Square Planar (2 lone pairs).

Bond Order

Q4. Calculate Bond Order of O2+.

O2 has 16e. O2+ has 15e.

Config: σ1s2...σ2pz2 π2px2=π2py2 π*2px1.

Nb = 10, Na = 5.

BO = 0.5 (10 - 5) = 2.5.

Paramagnetic (1 unpaired electron).

Shape

Q5. Predict shape of ClF3.

V=7, M=3. H = 1/2 (10) = 5 (sp3d).

Positions: 5. Atoms: 3. Lone pairs: 2.

T-shaped (Lone pairs at equatorial).

Fajan's Rule

Q6. Which is more covalent: LiCl or KCl?

Li+ is smaller than K+.

Smaller cation polarizes anion more.

LiCl is more covalent.

Resonance

Q7. Bond order of Ozone (O3).

Two resonating structures. One double one single.

Total bonds = 3. Resonating positions = 2.

BO = 3/2 = 1.5.

Sigma and Pi Bonds

Q8. Count σ and π bonds in C2H2 (Acetylene).

H-C ≡ C-H

Two C-H sigma bonds. One C-C sigma bond. Total σ = 3.

Two pi bonds in triple bond. Total π = 2.

Hydrogen Bonding

Q9. Why is H2O liquid but H2S gas?

O is small and highly electronegative, forms intermolecular H-bonds.

S is large, H-bonding negligible.

Association of molecules in H2O raises boiling point.

Formal Charge

Q10. Formal charge on central O in O3?

Central O has 1 double bond, 1 single bond, 1 lone pair.

FC = 6 - 2 - 1/2(6) = 6 - 2 - 3 = +1.

Formulas & Facts - Bonding

Equations & Formulas

Concept	Formula
Formal Charge	V - L - 1/2 S
Hybridization (H)	1/2 (V + M - C + A)
Bond Order (MO)	1/2 (Nb - Na)
Percent Ionic Char	16ΔX + 3.5ΔX²
Dipole Moment	μ = q × d
Bond Length	sp3 > sp2 > sp
Bond Enthalpy	Triple > Double > Single
H-Bond Strength	F-H...F > O-H...O > N-H...N

50 NEET Facts

Key points for Chemical Bonding.

1. Octet Rule exceptions Incomplete octet (LiCl, BeH2, BCl3). Expanded octet (PF5, SF6). Odd electron (NO).

2. Ionic Bond Conditions Low IE of metal, High EA of non-metal, High Lattice Enthalpy.

3. Lattice Enthalpy Energy released when 1 mole of crystal forms from gaseous ions. Depends on charge product (q1q2) and 1/r.

4. Born Haber Cycle Method to calculate Lattice Enthalpy using Hess's Law.

5. Solubility of Ionic Soluble in polar solvents (Hydration Energy > Lattice Energy).

6. Covalent Bond Sharing of electrons. Formed between non-metals.

7. Sigma Bond Axial overlap (s-s, s-p, p-p). Stronger. Free rotation possible.

8. Pi Bond Sideways overlap (p-p). Weaker. Restricted rotation. Always formed after sigma.

9. Hybridization Concept to explain equivalency of bonds (e.g., 4 C-H bonds in CH4).

10. sp Hybridization Linear, 180 deg. (BeCl2, C2H2). 50% s-character.

11. sp2 Hybridization Trigonal Planar, 120 deg. (BF3, C2H4). 33% s-character.

12. sp3 Hybridization Tetrahedral, 109.5 deg. (CH4). 25% s-character.

13. Bond Angle & Lone Pair Lone pairs repel more. CH4(109.5) > NH3(107) > H2O(104.5).

14. Drago's Rule No hybridization in PH3, AsH3, SbH3. Bond angles ~90 deg. Pure p-orbitals used.

15. Bent's Rule More electronegative substituent prefers hybrid orbital with less s-character (axial position in PCl5).

16. Axial vs Equatorial In sp3d (PCl5), Axial bonds are longer and weaker than Equatorial bonds.

17. SF6 sp3d2. Octahedral. All bonds equal length.

18. IF7 sp3d3. Pentagonal Bipyramidal.

19. XeF2 sp3d. 3 Lone pairs equatorial. Linear shape.

20. XeF4 sp3d2. 2 Lone pairs axial. Square Planar shape.

21. Molecular Orbital Theory Explains paramagnetism of O2 (2 unpaired e- in antibonding pi*).

22. Bond Order (BO) Directly prop to Stability and Bond Energy. Inversely prop to Bond Length.

23. BO of O2 series O2+ (2.5) > O2 (2.0) > O2- (1.5) > O2-- (1.0).

24. BO of N2 series N2 (3.0) > N2+ (2.5) = N2- (2.5).

25. Homo vs Hetero atomic LCAO is effective when atomic orbitals have comparable energies.

26. Hydrogen Bonding Dipole-dipole attraction with H bonded to F, O, N.

27. Intramolecular H-bond Within same molecule (o-nitrophenol). Lowers boiling point, prevents association.

28. Intermolecular H-bond Between molecules (water, alcohol). Increases boiling point, viscosity.

29. Ice Density Ice (cage-like structure) is less dense than water. Max density at 4C.

30. Dipole Moment (CO2) Zero (Linear, vectors cancel).

31. Dipole Moment (BF3) Zero (Trigonal planar, vectors cancel).

32. Dipole Moment (NH3 vs NF3) NH3 > NF3. In NH3, orbital moment and bond moments add up. In NF3, they oppose.

33. Fajan's Rule Covalency High charge, small cation, large anion -> High Polarization -> High Covalent character.

34. Melting Point trend (Fajan) SnCl2 (ionic) > SnCl4 (covalent).

35. Color (Fajan) AgCl (white) -> AgBr (pale yellow) -> AgI (yellow). Polarization causes color (charge transfer).

36. Resonance Delocalization of pi electrons. Increases stability.

37. Bond Length in Benzene All C-C bonds equal (1.39 A), intermediate between single (1.54) and double (1.34).

38. Formal Charge Bookkeeping of electrons. Structure with lowest formal charges is most stable.

39. Banana Bond 3c-2e bond in Diborane (B2H6). Tau bond.

40. Back Bonding BF3 acts as Lewis acid. Sideways overlap of filled F p-orbital to empty B p-orbital. p-pi-p-pi back bond.

41. Coordinate Bond in NH4+ N gives lone pair to H+. Structure is tetrahedral.

42. Metallic Bond Sea of electrons model. Electrical conductivity, Malleability.

43. Van der Waals Forces Weak. Dispersion (London) forces, Dipole-Dipole, Dipole-Induced Dipole.

44. Strength of forces Ionic > Covalent > Metallic > H-bond > Dipole-Dipole > London.

45. Boiling point of Noble gases Increases down group due to increasing London forces (size).

46. KHF2 Exists as K+ and [HF2]-. Very strong Symmetric H-bond in [F-H-F]-.

47. Solubility of Sulfate BeSO4 is soluble (High Hydration). BaSO4 is insoluble (Small Hydration vs Lattice).

48. Thermal Stability of Carbonates Increases down group. Li2CO3 decomposes easily. K2CO3 is stable.

49. Bond Angle in Ether C-O-C angle > 109.5 due to steric repulsion of alkyl groups.

50. NO2 Odd electron molecule. Paramagnetic. Dimerizes to N2O4 (Diamagnetic) at low temp.

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