Chemical Bonding and Molecular Structure: Complete Guide & Formula Sheet with Free PDF Download (JEE & NEET)
Chemical Bonding — Competishun
Chemical Bonding and Molecular Structure: Complete Guide & Formula Sheet with Free PDF Download (JEE & NEET)
Chemical Bonding and Molecular Structure is one of the most fundamental and scoring chapters in Class 11 chemistry. It carries high weightage in JEE and NEET, with 3-4 questions appearing every year.
This chapter explains why atoms combine, how they form different types of bonds, and what determines the shape and properties of molecules. It covers the octet rule, ionic and covalent bonds, VSEPR theory, bond parameters, and molecular orbital theory.
This page gives you the complete guide to Chemical Bonding with all concepts explained in depth. You will find clear definitions, derivations, worked examples, and common mistakes to avoid. Download the free PDF below and keep it handy for quick revision before your JEE Main, JEE Advanced, or NEET exam.
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Download Free PDFWhat is Chemical Bonding?
Chemical Bonding is explained by the Kössel-Lewis approach, which is based on the octet rule. According to this rule, atoms combine to achieve a stable configuration of 8 electrons in their valence shell (or 2 for hydrogen and helium).[reference:1]
Ionic Bond
Formed by the complete transfer of electrons from a metal to a non-metal. Electrostatic attraction between oppositely charged ions.[reference:2]
Covalent Bond
Formed by the sharing of electrons between two non-metals. One pair of electrons forms a single bond.[reference:3]
Octet Rule and Lewis Structures
Lewis Symbols
Lewis symbols represent the valence electrons of an atom as dots around the chemical symbol. For example: Li·, ·Be·, ·B·, ·C·, ·N·, ·O·, ·F·, ·Ne·.[reference:5]
Lewis Structures — Step-by-Step
- Step 1: Count the total number of valence electrons. N = ΣVatoms + (-ve charge) − (+ve charge).[reference:6]
- Step 2: Place the least electronegative atom in the center (hydrogen and fluorine are never central).[reference:7]
- Step 3: Draw single bonds between atoms and fill the outer octets first.[reference:8]
- Step 4: If electrons are short, make multiple bonds. π = (octet demand − N) / 2.[reference:9][reference:10]
Formal Charge
Where V = valence electrons, L = lone pair electrons, B = bonding electrons.[reference:11]
The sum of formal charges must equal the overall charge of the molecule or ion.[reference:12]
Exceptions to the Octet Rule
| Type | Examples | Electrons |
|---|---|---|
| Incomplete octet | BeCl₂, BF₃, AlCl₃, BH₃ | 4 or 6 |
| Expanded octet | PCl₅, XeF₂, SF₆, IF₇ | 10, 12, or 14 |
| Odd-electron | NO, NO₂, ClO₂, O₂ | 11, 17, 19 |
| Period 2 elements have no d-orbitals, so they cannot expand their octet. For example, NCl₅ does not exist, but PCl₅ does.[reference:13] | ||
Ionic Bond — Formation and Properties
Conditions for Ionic Bond Formation
- Low ionization enthalpy of the metal (to remove electrons easily).[reference:15]
- High negative electron gain enthalpy of the non-metal (to accept electrons readily).[reference:16]
- High lattice enthalpy (to stabilize the ionic solid).[reference:17]
- Electronegativity difference (Δχ) > 1.7 (to ensure complete electron transfer).[reference:18]
Lattice Enthalpy
NaCl(s) → Na⁺(g) + Cl⁻(g) · U = +787 kJ mol⁻¹[reference:19]
| Factor | Trend |
|---|---|
| Charge | Doubling charge quadruples lattice enthalpy (z⁺z⁻).[reference:20] |
| Cation size | Smaller cation → higher |U|. LiF > NaF > KF > RbF > CsF.[reference:21] |
| Anion size | Smaller anion → higher |U|. LiF > LiCl > LiBr > LiI.[reference:22] |
| Charge beats size: Al₂O₃ > MgO > CaO > NaF > NaCl.[reference:23] | |
Properties of Ionic Compounds
- High melting and boiling points — due to strong electrostatic forces.[reference:24]
- Hard but brittle — shifting one layer aligns like charges, causing repulsion and fracture.[reference:25]
- Conduct electricity only in molten or aqueous state (when ions are free to move).[reference:26]
- Dissolve in polar solvents — hydration enthalpy compensates for lattice enthalpy.[reference:27]
- No isomerism — ionic bonds are non-directional.[reference:28]
Covalent Bond — Formation and Properties
Coordinate Bond (Dative Bond)
A coordinate bond is formed when one atom provides both electrons for sharing. The donor atom must have a lone pair, and the acceptor atom must have a vacant orbital.[reference:32]
In NH₄⁺, all four N–H bonds are identical — there is no distinction between the coordinate bond and the others.[reference:33]
Bond Parameters
| Parameter | Definition | Trend |
|---|---|---|
| Bond Length | Distance between nuclei of bonded atoms | Higher bond order → shorter bond length. H–F (92) < H–Cl (127) < H–Br (141) < H–I (161) pm.[reference:34] |
| Bond Enthalpy | Energy required to break one mole of bonds | Higher bond order → higher bond enthalpy. C–C (348) < C=C (614) < C≡C (839) kJ mol⁻¹.[reference:35] |
| Bond Order | Number of bonds between two atoms | O₃ (1.5), CO₃²⁻ (1.33), SO₃ (2.0), benzene (1.5).[reference:36] |
| Anomaly: F₂ (155) < Cl₂ (242) — F₂ has weak bond due to lone pair-lone pair repulsion.[reference:37] | ||
VSEPR Theory — Predicting Molecular Shapes
Key Rules of VSEPR Theory
- Lone pairs occupy more space than bond pairs, so they compress bond angles.
- Multiple bonds are treated as single super-pairs — they repel more strongly.
- Order of repulsion: Lone pair-Lone pair > Lone pair-Bond pair > Bond pair-Bond pair.
The Angle Table — Every Bond Angle at a Glance
| Shape | Bond Angle | Examples |
|---|---|---|
| Linear | 180° | BeCl₂, CO₂, XeF₂, I₃⁻, NO₂⁺[reference:40] |
| Trigonal Planar | 120° | BF₃, SO₃, NO₃⁻, CO₃²⁻[reference:41] |
| Tetrahedral | 109.5° | CH₄, NH₄⁺, SO₄²⁻[reference:42] |
| Trigonal Pyramidal | 107° | NH₃[reference:43] |
| Bent | 104.5° | H₂O[reference:44] |
| Trigonal Bipyramidal | 90°/120° | PCl₅ |
| Octahedral | 90° | SF₆ |
| The bond angle table is one of the most frequently tested topics in JEE and NEET. Memorize these values.[reference:45] | ||
Factors Affecting Bond Angles
| Factor | Effect on Angle |
|---|---|
| Lone pairs on central atom | Lone pairs ↑ → Angle ↓. CH₄ (109.5°) > NH₃ (107°) > H₂O (104.5°)[reference:46] |
| Electronegativity of central atom | Central atom EN ↑ → Angle ↓. H₂O (104.5°) > H₂S (92°) > H₂Se (91°) > H₂Te (89.5°)[reference:47] |
| Electronegativity of surrounding atom | Surrounding EN ↑ → Angle ↓. Cl₂O (110.9°) > H₂O (104.5°) > F₂O (103.2°)[reference:48] |
| Multiple bonds / bulky groups | Bulkier groups ↑ → Angle ↑ (more repulsion) |
| Back bonding | Back bonding ↑ → Angle ↑[reference:49] |
| The reversal: NH₃ (107°) > NF₃ (102°), but PH₃ (93.5°) < PF₃ (97.8°). This is asked every year.[reference:50] | |
Molecular Orbital Theory (MOT)
Key Concepts
- Bonding Molecular Orbital (BMO): Lower energy, stabilized by in-phase combination of atomic orbitals.
- Antibonding Molecular Orbital (ABMO): Higher energy, destabilized by out-of-phase combination.
- Bond Order: BO = (NB − NA) / 2, where NB = electrons in bonding MOs, NA = electrons in antibonding MOs.
- Paramagnetism: Molecules with unpaired electrons (like O₂) are paramagnetic.
All Chemical Bonding Formulas at a Glance
| Formula | What It Means |
|---|---|
| FC = V − L − ½B | Formal charge calculation[reference:51] |
| π = (Octet Demand − N) / 2 | Number of π bonds in Lewis structure[reference:52] |
| |U| ∝ z⁺z⁻ / (r⁺ + r⁻) | Lattice enthalpy trend[reference:53] |
| ΔHsol = ΔHhyd − U | Enthalpy of solution[reference:54] |
| dA-B = rA + rB − 0.09|χA − χB| | Schomaker-Stevenson bond length[reference:55] |
| BO = (NB − NA) / 2 | Bond order from MOT |
| Memorise these formulas for Chemical Bonding. They are the key to scoring full marks in this chapter. | |
Common Mistakes in Chemical Bonding
- Forgetting the octet rule exceptions: Be, B, and elements from Period 3 can have incomplete or expanded octets. Period 2 elements cannot expand their octet.[reference:56]
- Confusing ionic and covalent bonds: Ionic bonds form between metals and non-metals (Δχ > 1.7). Covalent bonds form between non-metals (Δχ < 1.7).
- Misapplying VSEPR theory: Lone pairs repel more strongly than bond pairs and compress bond angles. This is a common source of errors.
- Forgetting that bond order and bond length are inversely related: Higher bond order → shorter bond length. This is a fundamental relationship in Chemical Bonding.
- Not checking formal charges: The sum of formal charges must equal the overall charge of the molecule or ion.[reference:57]
- Confusing bond enthalpy and bond dissociation enthalpy: Bond enthalpy is the average energy to break a bond in a molecule, while bond dissociation enthalpy is the energy to break a specific bond.
Why Chemical Bonding Matters for JEE and NEET
- High weightage: Chemical Bonding appears in 3-4 questions in every JEE Main, JEE Advanced, and NEET chemistry paper.
- Foundation for inorganic chemistry: Understanding Chemical Bonding helps you with coordination compounds, p-block chemistry, and solid-state chemistry.
- Direct scoring: Many questions are direct, especially VSEPR theory, bond angles, and bond order.
- Conceptual clarity: This chapter rewards students who understand the concepts rather than just memorizing formulas.
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