Chemical Bonding and Molecular Structure: Class 11 Chemistry Chapter 4 Notes

Key Terms and Definitions

Ionic Bond

The chemical bond formed through the complete transfer of one or more electrons from one atom to another, resulting in the formation of ions (cation and anion) held together by electrostatic forces.

Covalent Bond

A chemical bond formed by the mutual sharing of one or more pairs of electrons between two atoms, typically non-metals, to achieve a stable electron configuration.

Octet Rule

The tendency of atoms to prefer to have eight electrons in their valence shell. When atoms have fewer than eight electrons, they tend to react and form more stable compounds.

Bond Enthalpy

The amount of energy required to break one mole of a particular type of bond between two atoms in a gaseous state. It is a measure of bond strength.

VSEPR Theory

Valence Shell Electron Pair Repulsion Theory. It is a model used to predict the geometry of individual molecules from the number of electron pairs surrounding their central atoms.

Hybridization

The concept of mixing atomic orbitals of an atom with slightly different energies to form a new set of equivalent orbitals, called hybrid orbitals, which are more suitable for bonding.

Molecular Orbital Theory (MOT)

A theory that describes chemical bonding in terms of molecular orbitals, which result from the combination of atomic orbitals of the bonding atoms and are spread over the entire molecule.

Dipole Moment (μ)

A measure of the polarity of a chemical bond or molecule. It is the product of the magnitude of the charge and the distance of separation between the charges. It is a vector quantity.

Hydrogen Bond

A special type of dipole-dipole attraction between molecules, resulting from the attractive force between a hydrogen atom covalently bonded to a very electronegative atom (N, O, or F) and another nearby electronegative atom.

VSEPR Theory: Predicting Molecular Shape

The Valence Shell Electron Pair Repulsion (VSEPR) theory is a simple yet powerful tool for predicting the three-dimensional shape of molecules. Its core principle is that electron pairs in the valence shell of a central atom repel each other. To minimize this repulsion, these electron pairs orient themselves as far apart as possible in space. This arrangement determines the molecule's electron geometry. The theory considers both bonding pairs (involved in covalent bonds) and lone pairs (non-bonding electrons). A critical point is the hierarchy of repulsion: Lone Pair-Lone Pair (LP-LP) repulsion is the strongest, followed by Lone Pair-Bond Pair (LP-BP) repulsion, and finally, Bond Pair-Bond Pair (BP-BP) repulsion is the weakest. This is because lone pairs are localized on the central atom and occupy more space than bonding pairs, which are shared between two atoms. For example, in methane (CH₄), there are four bonding pairs and zero lone pairs around the central carbon. They arrange themselves in a tetrahedral geometry with bond angles of 109.5°. In ammonia (NH₃), there are three bonding pairs and one lone pair. The electron geometry is still tetrahedral, but the LP-BP repulsion is stronger than BP-BP repulsion, compressing the H-N-H bond angle to 107°. The final molecular shape (or geometry) only considers the positions of the atoms, so NH₃ is described as trigonal pyramidal. In water (H₂O) with two lone pairs, the repulsion is even greater, reducing the H-O-H angle to 104.5°, and its shape is described as bent or V-shaped.

Key Formulas & Concepts at a Glance

• Hybridization (H): A quick way to find hybridization is using the formula: H = ½ [V + M – C + A], where V = valence electrons of the central atom, M = no. of monovalent atoms, C = positive charge, A = negative charge.
• Bond Order (from MOT): Bond Order = ½ [N_b – N_a], where N_b is the number of electrons in bonding molecular orbitals and N_a is the number in antibonding molecular orbitals.
• Bond Order & Stability: A positive bond order (B.O. > 0) indicates a stable molecule. Higher bond order implies greater stability and higher bond dissociation enthalpy.
• Bond Order & Bond Length: Bond length is inversely proportional to bond order. A triple bond (B.O. = 3) is shorter than a double bond (B.O. = 2), which is shorter than a single bond (B.O. = 1).
• Dipole Moment & Polarity: A molecule with a net dipole moment (μ ≠ 0) is polar. Symmetrical molecules (like CO₂, CCl₄, BF₃) can have polar bonds but be non-polar overall because the individual bond dipoles cancel out.
• Fajan's Rules: Predict the degree of covalent character in an ionic bond. Covalent character is favored by: small cation, large anion, and high charge on either ion.
• Repulsion Order in VSEPR: Lone Pair–Lone Pair > Lone Pair–Bond Pair > Bond Pair–Bond Pair. This is key to predicting correct bond angles.
• Hydrogen Bonding: Affects physical properties like boiling point and solubility. Intermolecular H-bonding (between molecules, e.g., in H₂O) increases boiling point. Intramolecular H-bonding (within a molecule, e.g., in o-nitrophenol) decreases it.

Sigma (σ) vs. Pi (π) Bonds

Worked Examples

• {"item":"Predicting Hybridization and Shape of SF₄","body":"1. Find Hybridization: Central atom is S (Sulfur). \n - Valence electrons (V) = 6\n - Monovalent atoms (M) = 4 (Fluorine)\n - Charge (C/A) = 0\n - H = ½ [6 + 4 – 0 + 0] = 5. \n - H=5 corresponds to sp³d hybridization.\n2. Determine Geometry and Shape: \n - 5 electron pairs give a trigonal bipyramidal electron geometry.\n - Sulfur forms 4 bonds with F, so there are 4 bond pairs.\n - No. of lone pairs = 5 (total pairs) - 4 (bond pairs) = 1 lone pair.\n - In trigonal bipyramidal geometry, the lone pair occupies an equatorial position to minimize repulsion.\n - The resulting molecular shape is a See-Saw."}
• {"item":"Calculating Bond Order of O₂","body":"1. Total Electrons: Oxygen atom has 8 electrons, so O₂ molecule has 16 electrons.\n2. MOT Configuration for O₂ (16e⁻): \n (σ1s)² (σ1s)² (σ2s)² (σ2s)² (σ2pz)² (π2px)² (π2py)² (π2px)¹ (π2py)¹\n3. Count Bonding and Antibonding Electrons:\n - N_b (Bonding) = 2+2+2+2+2 = 10\n - N_a (Antibonding) = 2+2+1+1 = 6\n4. Calculate Bond Order:\n - B.O. = ½ [N_b – N_a] = ½ [10 – 6] = ½ [4] = 2.\n - The bond order of 2 corresponds to the double bond in O=O. The presence of unpaired electrons in π* orbitals explains the paramagnetic nature of O₂."}

CBSE Exam Alert

A common trap for students is confusing 'electron geometry' with 'molecular shape'. Remember, electron geometry describes the arrangement of ALL electron pairs (bonding and lone pairs), while molecular shape describes the arrangement of ATOMS ONLY. For example, in XeF₄: there are 4 bonding pairs and 2 lone pairs, making a total of 6 electron pairs. The electron geometry is octahedral. However, since the two lone pairs are not 'seen' in the final structure, the molecular shape is square planar. Always start by finding the total number of electron pairs to determine the electron geometry, then place the lone pairs to minimize repulsion, and finally, describe the shape based on the atomic positions. This distinction is crucial for getting full marks.

Quick Revision Check

• Q: Why is the H-O-H bond angle in water (104.5°) less than the tetrahedral angle (109.5°)? A: Due to the stronger repulsion from the two lone pairs on the oxygen atom compared to the bond pairs. This LP-LP and LP-BP repulsion pushes the O-H bonds closer together.
• Q: Which molecule has a zero dipole moment: NH₃ or BF₃? A: BF₃. It has a symmetrical trigonal planar geometry where the three polar B-F bond dipoles cancel each other out. NH₃ has a trigonal pyramidal shape and a net dipole moment.
• Q: Arrange the following in increasing order of bond strength: O₂, O₂⁺, O₂⁻. A: The bond orders are O₂ (2.0), O₂⁺ (2.5), O₂⁻ (1.5). Since bond strength is proportional to bond order, the order is: O₂⁻ < O₂ < O₂⁺.
• Q: What is the hybridization of the carbon atoms in ethyne (C₂H₂)? A: Each carbon atom is triple-bonded to the other carbon and single-bonded to a hydrogen. This corresponds to sp hybridization.

Frequently Asked Questions