Chemical Bonding and Molecular Structure Class 11 Chapter Notes

Essential Chapter Glossary

Octet Rule

The chemical rule stating that atoms tend to combine in such a way that they each have eight electrons in their valence shells, giving them the same electronic configuration as a noble gas.

Formal Charge

The hypothetical charge assigned to an individual atom in a molecule, calculated as: [Total valence e- in free atom] - [Total non-bonding e-] - 0.5 * [Total bonding e-].

Dipole Moment (μ)

The measure of net polarity in a molecule. It is the product of the magnitude of charge (q) and the distance (d) separating the charges: μ = q × d. Unit: Debye (D).

Hybridization

The theoretical process of mixing atomic orbitals of slightly different energies of the same atom to produce a new set of equivalent orbitals known as hybrid orbitals.

Resonance

A phenomenon where a single Lewis structure cannot accurately represent a molecule; instead, the molecule is described as a hybrid of two or more contributing resonance structures.

Hydrogen Bonding

A strong dipole-dipole attractive force that occurs when hydrogen is covalently bonded to highly electronegative atoms like Fluorine, Oxygen, or Nitrogen (F, O, N).

VSEPR Theory and Molecular Geometries

The Valence Shell Electron Pair Repulsion (VSEPR) Theory provides a reliable method to predict the 3D shapes of covalent molecules. The fundamental postulate is that valence shell electron pairs (both bonding and non-bonding/lone pairs) surrounding a central atom repel each other and stay as far apart as possible to minimize repulsion and maximize stability.

Crucially, the repulsive forces follow a strict hierarchical order: Lone Pair - Lone Pair (lp-lp) > Lone Pair - Bond Pair (lp-bp) > Bond Pair - Bond Pair (bp-bp). When lone pairs are present, they distort the regular geometry by squeezing the adjacent bond angles. For example, while methane ($CH_4$) has a perfect tetrahedral shape with a bond angle of $109.5^\circ$ due to having 4 bond pairs, ammonia ($NH_3$) with 3 bond pairs and 1 lone pair is distorted to a trigonal pyramidal geometry with a reduced angle of $107^\circ$. Similarly, water ($H_2O$) with 2 bond pairs and 2 lone pairs is distorted to a bent/V-shaped geometry with an angle of $104.5^\circ$.

Valence Bond Theory (VBT) vs. Molecular Orbital Theory (MOT)

Step-by-Step Method to Determine Hybridization and Shape

1. Calculate the Steric Number (H) — Use the formula: H = 0.5 * [V + M - C + A], where V = Valence electrons of central atom, M = Number of monovalent atoms surrounding it, C = Charge on cation, and A = Charge on anion.
2. Determine Hybridization State — Correlate H to orbital set: If H = 2 (sp), H = 3 (sp2), H = 4 (sp3), H = 5 (sp3d), H = 6 (sp3d2).
3. Find Lone Pairs (lp) — Calculate lone pairs: lp = Steric Number (H) - Number of bonded atoms (bp).
4. Determine the Shape / Geometry — If lp = 0, molecular shape matches the regular geometry. If lp > 0, arrange lone pairs in positions minimizing lp-lp repulsion (e.g., equatorial positions in trigonal bipyramidal) and describe the final visible atomic framework shape.

Must-Remember Formulas and Rules

• Fajan's Rules for Covalent Character in Ionic Bonds: Covalent character is favored by a small cation size, high cation charge, and large anion size.
• Bond Order Formula in MOT: Bond Order = 0.5 * (Nb - Na), where Nb is the number of bonding electrons and Na is the number of antibonding electrons.
• Bond Stability: A higher Bond Order means greater bond stability and shorter bond length.
• Dipole Moment of Symmetrical Molecules: Highly symmetrical molecules like BF3, CCl4, and CO2 have a net dipole moment of zero (μ = 0) because individual bond dipoles cancel out.
• Oxygen (O2) Paramagnetism: The MOT electronic configuration reveals two unpaired electrons in the π2px and π2py orbitals, proving paramagnetic behavior.
• Exceptions to the Octet Rule: Incomplete octet (e.g., BeCl2, BF3), expanded octet (e.g., PCl5, SF6), and odd-electron molecules (e.g., NO, NO2).
• Types of Hydrogen Bonding: Intermolecular (between different molecules, e.g., H2O, HF) and Intramolecular (within the same molecule, e.g., o-nitrophenol).

Worked Calculations & Examples

• {"title":"Example 1: Calculating Formal Charge in Ozone (O3)","description":"Let's label the three oxygen atoms in O3 as central (O1), single-bonded (O2), and double-bonded (O3).\n- Central O1: Valence = 6, Non-bonding = 2, Bonding = 6 (3 bonds). Formal Charge = 6 - 2 - 0.5(6) = +1.\n- Double-bonded O3: Valence = 6, Non-bonding = 4, Bonding = 4. Formal Charge = 6 - 4 - 0.5(4) = 0.\n- Single-bonded O2: Valence = 6, Non-bonding = 6, Bonding = 2. Formal Charge = 6 - 6 - 0.5(2) = -1.\nResult: The overall sum is 0, matching the neutral molecule's charge.*"}
• {"title":"Example 2: Determining Bond Order of O2, O2+, and O2-","description":"Using MOT configuration for Oxygen (16 electrons):\n- O2 (16 e-): Bond Order = 0.5 (10 - 6) = 2.0 (Paramagnetic)\n- O2+ (15 e-): Bond Order = 0.5 (10 - 5) = 2.5 (More stable than O2)\n- O2- (17 e-): Bond Order = 0.5 (10 - 7) = 1.5 (Less stable than O2)\nStability Order: O2+ > O2 > O2-*"}

Exam Traps & Board Cues

1. Shape vs. Geometry Confusion: In VSEPR, 'Geometry' refers to the arrangement of electron pairs (including lone pairs), while 'Shape' or 'Molecular Structure' refers only to the position of atoms. For example, $NH_3$ has a tetrahedral geometry but a trigonal pyramidal shape.

2. MOT Configuration Switch: Remember that for elements with atomic number $Z \le 7$ (like $B_2$, $C_2$, $N_2$), the $\pi_{2px}$ and $\pi_{2py}$ orbitals are lower in energy than the $\sigma_{2pz}$ orbital. For $Z > 7$ (like $O_2$, $F_2$), the $\sigma_{2pz}$ orbital is lower in energy than the $\pi_{2px}$ and $\pi_{2py}$ orbitals. Don't mix up this ordering during bond order calculations!

Quick Revision Check

• Why is liquid water (H2O) more dense than solid ice? In ice, intermolecular hydrogen bonding creates a highly ordered, open cage-like tetrahedral structure with larger volume/gaps. Upon melting, these hydrogen bonds partially collapse, allowing water molecules to pack more closely, increasing density.
• Which of the two, NH3 or NF3, has a higher dipole moment and why? NH3 has a higher dipole moment. In NH3, the orbital dipole due to the lone pair and the N-H bond dipoles point in the same direction, reinforcing each other. In NF3, the highly electronegative F atoms pull density away, causing N-F bond dipoles to oppose the lone-pair dipole, partially canceling it.
• What is the hybridization of the phosphorus atom in PCl5, and why are axial bonds longer than equatorial bonds? The hybridization is sp3d (trigonal bipyramidal). The axial bonds suffer greater repulsion from the equatorial bond pairs (repulsion at 90 degrees) than the equatorial bonds do, causing the axial bonds to stretch and become longer and weaker.
• Why does He2 molecule not exist according to Molecular Orbital Theory? He2 would have 4 electrons. Its MOT configuration is σ1s² σ1s². The Bond Order = 0.5 (Nb - Na) = 0.5 * (2 - 2) = 0. Since the bond order is zero, the molecule is unstable and cannot exist.

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