Chemical Bonding | IISER Smart Prep

1. Core Concept

Chemical bonding is the physical process responsible for the attractive interactions between atoms and molecules, which confers stability to diatomic and polyatomic chemical compounds. Atoms bond to achieve a lower potential energy and a stable octet (noble gas configuration).

2. Kössel-Lewis Approach

Kössel and Lewis provided the first logical explanation for valence: atoms bond to complete their valence shell octet.

Octet Rule: Atoms gain, lose, or share electrons to reach 8 electrons in their outer shell.
Lewis Symbols: Represented by valence electrons as dots around the element symbol.

Formal Charge (F.C.) = V - L - ½S

V = Valence e^-, L = Lone pair e^-, S = Shared/Bonded e^-. Helps identify the most stable Lewis structure (lowest net F.C.).

Octet Rule Limitations:

Incomplete Octet: LiCl, BeH₂, BF₃ (central atom < 8e^-).
Expanded Octet: PCl₅, SF₆, H₂SO₄ (central atom > 8e^- due to d-orbitals).
Odd-electron molecules: NO, NO₂.

3. Ionic or Electrovalent Bond

Formed by complete transfer of electrons from a metal (low IE) to a non-metal (high Δ_egH). Resulting ions are held by electrostatic force.

Factors favoring Ionic Bond:

Low Ionization Enthalpy of cation.
High magnitude of Electron Gain Enthalpy of anion.
High Lattice Enthalpy: The energy released when 1 mole of ionic crystal is formed from gaseous ions. (Lattice Energy ∝ Charge / Radius).

Fajan's Rule (Covalent character in Ionic bonds):

No bond is 100% ionic. Covalent character increases when:

Small cation (high polarizing power).
Large anion (high polarisability).
High charges on ions.
Cations with pseudo-noble gas configuration (ns²np⁶nd¹⁰, e.g., Cu⁺).

4. Bond Parameters

Parameter	Definition / Dependency
Bond Length	Equilibrium distance between nuclei. Decreases with higher B.O.
Bond Angle	Angle between orbitals containing bonding pairs.
Bond Enthalpy	Energy to break 1 mole of bonds. Bond Strength ∝ B.Enthalpy.
Bond Order	Number of bonds between atoms. B.O. = 1 (single), 2 (double), 3 (triple).

Resonance:

When a single Lewis structure can't explain all properties (e.g., O₃, CO₃^2-). Resonance hybrids have intermediate bond lengths and lower energy.

μ = Q × r

Dipole Moment: Vector quantity measured in Debye (D). μ = 0 for symmetric molecules (e.g., BF₃, CCl₄); μ ≠ 0 for polar molecules (e.g., H₂O, NF₃).

5. VSEPR Theory

Predicts geometry based on repulsion: LP-LP > LP-BP > BP-BP.

Type	BP	LP	Shape	Example
AX₂	2	0	Linear	BeCl₂
AX₃	3	0	Trigonal Planar	BF₃
AX₂E	2	1	Bent / V-shape	SO₂, O₃
AX₄	4	0	Tetrahedral	CH₄
AX₃E	3	1	Trigonal Pyramidal	NH₃
AX₂E₂	2	2	Bent / V-shape	H₂O
AX₅	5	0	Trigonal Bipyramidal	PCl₅
AX₄E	4	1	See-saw	SF₄
AX₃E₂	3	2	T-shape	ClF₃
AX₆	6	0	Octahedral	SF₆
AX₅E	5	1	Square Pyramidal	BrF₅
AX₄E₂	4	2	Square Planar	XeF₄

6. Valence Bond Theory & Hybridisation

VBT: Bond is formed by overlap of half-filled atomic orbitals. Focuses on directional nature.

Sigma (σ) bond: Head-on overlap (strong).
Pi (π) bond: Sideways overlap (weak; only possible if σ is already present).

Hybridisation Types & Geometry:

sp: Linear (180°). BeCl₂, C₂H₂.
sp²: Trigonal Planar (120°). BF₃, C₂H₄.
sp³: Tetrahedral (109.5°). CH₄, NH₃, H₂O.
sp³d: Trigonal Bipyramidal. PCl₅.
sp³d²: Octahedral. SF₆.
sp³d³: Pentagonal Bipyramidal. IF₇.

7. Molecular Orbital Theory (MOT)

Atomic orbitals combine to form Molecular Orbitals (MO). Number of MO = Number of AO.

Bonding MO: Lower energy, higher stability. (σ, π).
Antibonding MO: Higher energy, lower stability. (σ*, π*).

Bond Order (B.O.) = ½(N_b - N_a)

If B.O. > 0, species is stable. Stability ∝ B.O. Magnetic nature: Unpaired e^- → Paramagnetic; All paired → Diamagnetic.

Energy Sequence (MUST KNOW):

Up to N₂ (Z ≤ 14): σ1s < σ*1s < σ2s < σ*2s < (π2p_x = π2p_y) < σ2p_z < (π*2p_x = π*2p_y) < σ*2p_z.
O₂ and F₂ (Z > 14): σ1s < σ*1s < σ2s < σ*2s < σ2p_z < (π2p_x = π2p_y) < (π*2p_x = π*2p_y) < σ*2p_z.

Note: The position of σ2p_z changes due to s-p mixing in lighter elements.

8. Hydrogen Bonding

Special dipole-dipole interaction between H (bonded to F, O, N) and another electronegative atom.

Intermolecular: Between different molecules. Increases B.P. (e.g., H₂O, HF, alcohols).
Intramolecular: Within the same molecule. Reduces B.P., increases volatility. (e.g., o-nitrophenol, Salicylaldehyde).

Density Anomaly: Ice is less dense than water because H-bonding forms a cage-like structure in solid state, increasing volume.

9. Common Mistakes

Dipole Moment vectors: Don't forget that in NF₃, lone pair dipole opposes F dipoles, making it less polar than NH₃ where they reinforce.
See-saw positioning: In sp³d hybridisation with 1 LP (SF₄), the lone pair always occupies the equatorial position to minimize repulsion.
Magnetic nature of O₂: Standard Lewis octet theory fails; MOT correctly predicts it as paramagnetic (2 unpaired e^- in π* orbitals).

10. IAT Exam Focus Points

Bond Order Comparison: Arrange O₂, O₂⁺, O₂^-, O₂^2-. B.O.: 2, 2.5, 1.5, 1. Stability: O₂⁺ > O₂ > O₂^- > O₂^2-.
Isoelectronic Species Shapes: Species like CO₂, BeCl₂, N₃^- are all linear. CH₄, BH₄^-, NH₄⁺ are all tetrahedral.
Born-Haber Cycle: Understand the summation of energies (Sublimation + IE + Dissociation + EA + Lattice = Heat of Formation).
Bond Angles in H₂O and NH₃: 104.5° and 107°. Lower than 109.5° due to LP repulsion.

11. Practice Mock Test

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