Chemical Bonding and Molecular Structure - Comprehensive Revision Notes

🎯 Overview

This chapter explores how atoms combine to form molecules through various types of chemical bonds. Understanding chemical bonding is fundamental to predicting molecular structure, properties, and reactivity patterns in chemistry.

📚 Core Concepts

1. Why do Atoms Bond?

Octet Rule and Stability

🔋 Driving Forces for Bond Formation:
• Tendency to achieve noble gas configuration
• Lowering of potential energy
• Increased stability
• Completion of valence shell (octet rule)

⚡ Energy Considerations:
• Bond formation is exothermic
• Energy is released when bonds form
• More stable configuration = lower energy
• Bond energy measures bond strength

📊 Octet Rule Exceptions:
• H, He, Li, Be: Duet rule (2 electrons)
• B: Often electron-deficient (6 electrons)
• P, S: Can expand octet (10, 12 electrons)
• Radicals: Odd number of electrons

2. Types of Chemical Bonds

Ionic Bonding

⚡ Ionic Bond Characteristics:
• Transfer of electrons from metal to non-metal
• Electrostatic attraction between ions
• Formation of cations and anions
• High lattice energy
• Usually between metals (Groups 1, 2) and non-metals (Groups 16, 17)

📊 Properties of Ionic Compounds:
• High melting and boiling points
• Soluble in polar solvents (water)
• Conduct electricity in molten/aqueous state
• Crystalline structure
• Hard and brittle

🔋 Factors Favoring Ionic Bonding:
• Large electronegativity difference (>1.7)
• Low ionization energy (metal)
• High electron affinity (non-metal)
• Small size of ions

Covalent Bonding

🤝 Covalent Bond Characteristics:
• Sharing of electron pairs
• Equal or unequal sharing based on electronegativity
• Directional nature
• Usually between non-metals

📊 Types of Covalent Bonds:
• Single bond: σ bond only
• Double bond: σ + π bond
• Triple bond: σ + 2π bonds

⚡ Factors Affecting Covalent Bonding:
• Similar electronegativity (<1.7 difference)
• High ionization energies
• Orbital overlap
• Bond energy and length

🔬 Covalent Bond Properties:
• Lower melting points than ionic
• Poor electrical conductivity
• Variable solubility
• Molecular structure
• Directional properties

Coordinate/Dative Bonding

🎯 Coordinate Bond Characteristics:
• One-sided electron pair donation
• Lewis acid-base interaction
• Both electrons from same atom
• Arrow notation (→)
• Common in complex ions

📊 Examples:
• NH₃ → BF₃ (Ammonia borane)
• CO → Fe (Metal carbonyls)
• H₂O → H⁺ (Hydronium ion)

⚡ Properties:
• Similar to covalent bonds once formed
• Important in coordination chemistry
• Common in acid-base chemistry

3. Lewis Structures and Formal Charge

Drawing Lewis Structures

📝 Step-by-Step Method:
1. Count total valence electrons
2. Identify central atom (least electronegative)
3. Draw skeletal structure
4. Distribute remaining electrons
5. Form multiple bonds if needed
6. Check octet rule satisfaction

🎯 Central Atom Selection Rules:
• Least electronegative (except H)
• Can expand octet (period 3+)
• Usually the atom present once
• Avoids symmetrical arrangement issues

✅ Octet Rule Verification:
• Main group elements: 8 electrons (except H, He)
• Expanded octet possible for period 3+
• Electron-deficient: B, Be
• Odd electron species: radicals

Formal Charge Calculation

📊 Formal Charge Formula:
FC = (Valence electrons) - (Non-bonding electrons) - ½(Bonding electrons)

🎯 Formal Charge Rules:
• Most stable structure has minimal formal charges
• Negative formal charges on electronegative atoms
• Positive formal charges on electropositive atoms
• Adjacent atoms should not have same sign charges
• Formal charges should be close to zero

📝 Example: CO₂
O: FC = 6 - 4 - ½(4) = 0
C: FC = 4 - 0 - ½(8) = 0
O: FC = 6 - 4 - ½(4) = 0

✅ Best structure: All formal charges = 0

4. VSEPR Theory

Basic Principles

📐 VSEPR Theory Fundamentals:
• Valence Shell Electron Pair Repulsion
• Electron pairs arrange to minimize repulsion
• Repulsion order: LP-LP > LP-BP > BP-BP
• Geometry determined by electron pair arrangement
• Molecular shape from atom positions

🎯 Key Assumptions:
• Electron pairs repel each other
• Repulsion depends on distance
• Lone pairs occupy more space
• Multiple bonds counted as single region
• Central atom determines geometry

📊 Electron Domain Geometry:
• 2 domains: Linear (180°)
• 3 domains: Trigonal planar (120°)
• 4 domains: Tetrahedral (109.5°)
• 5 domains: Trigonal bipyramidal (120°, 90°)
• 6 domains: Octahedral (90°)

Common Molecular Geometries

📐 2 Electron Domains:
• Linear: 2 bonding pairs, 0 lone pairs
• Bond angle: 180°
• Examples: CO₂, BeCl₂

📐 3 Electron Domains:
• Trigonal planar: 3 BP, 0 LP, 120°
  Examples: BF₃, CO₃²⁻
• Bent: 2 BP, 1 LP, <120°
  Examples: SO₂, O₃

📐 4 Electron Domains:
• Tetrahedral: 4 BP, 0 LP, 109.5°
  Examples: CH₄, NH₄⁺
• Trigonal pyramidal: 3 BP, 1 LP, <109.5°
  Examples: NH₃, PCl₃
• Bent: 2 BP, 2 LP, <<109.5°
  Examples: H₂O, H₂S

📐 5 Electron Domains:
• Trigonal bipyramidal: 5 BP, 0 LP
  Examples: PCl₅
• Seesaw: 4 BP, 1 LP
  Examples: SF₄
• T-shaped: 3 BP, 2 LP
  Examples: ClF₃
• Linear: 2 BP, 3 LP
  Examples: XeF₂

📐 6 Electron Domains:
• Octahedral: 6 BP, 0 LP, 90°
  Examples: SF₆
• Square pyramidal: 5 BP, 1 LP
  Examples: BrF₅
• Square planar: 4 BP, 2 LP
  Examples: XeF₄

5. Valence Bond Theory

Basic Principles

🤝 Valence Bond Theory Fundamentals:
• Atomic orbitals overlap to form bonds
• Overlapping orbitals must have proper symmetry
• Unpaired electrons participate in bonding
• Bond strength depends on overlap extent
• Directional nature of covalent bonds

📊 Orbital Overlap Types:
• σ bond: Head-on overlap, cylindrical symmetry
• π bond: Side-on overlap, nodal plane
• δ bond: Four-lobed overlap (for transition metals)

⚡ Hybridization:
• Mixing of atomic orbitals
• Formation of equivalent hybrid orbitals
• Explains molecular geometry
• Number of hybrids = number of atomic orbitals mixed

Types of Hybridization

📐 sp Hybridization:
• 1 s + 1 p → 2 sp orbitals
• Linear geometry (180°)
• Examples: BeCl₂, CO₂, C₂H₂

📐 sp² Hybridization:
• 1 s + 2 p → 3 sp² orbitals
• Trigonal planar (120°)
• Examples: BF₃, CO₃²⁻, C₂H₄

📐 sp³ Hybridization:
• 1 s + 3 p → 4 sp³ orbitals
• Tetrahedral (109.5°)
• Examples: CH₄, NH₃, H₂O

📐 sp³d Hybridization:
• 1 s + 3 p + 1 d → 5 sp³d orbitals
• Trigonal bipyramidal
• Examples: PCl₅, SF₄

📐 sp³d² Hybridization:
• 1 s + 3 p + 2 d → 6 sp³d² orbitals
• Octahedral (90°)
• Examples: SF₆, XeF₄

🎯 Hybridization Determination:
1. Draw Lewis structure
2. Count electron domains around central atom
3. Determine hybridization
4. Predict molecular geometry

6. Molecular Orbital Theory

Basic Principles

🔬 Molecular Orbital Theory Fundamentals:
• Atomic orbitals combine to form molecular orbitals
• MOs belong to entire molecule
• Electrons occupy MOs following rules
• Bond order determines stability
• Explains magnetic properties

📊 MO Formation Rules:
• Number of MOs = Number of AOs combined
• Energy similar to AOs with same symmetry
• Bonding MOs lower energy than AOs
• Antibonding MOs higher energy than AOs
• Non-bonding MOs similar energy to AOs

🎯 Electron Filling Rules:
• Aufbau principle: Fill lowest energy first
• Pauli exclusion: Max 2 electrons per MO
• Hund's rule: Maximum unpaired electrons in degenerate MOs

MO Diagrams for Diatomic Molecules

⚡ Homonuclear Diatomic Molecules:

H₂, He₂, N₂, O₂, F₂ (Period 2):
• σ(1s), σ*(1s), σ(2s), σ*(2s), π(2p), σ(2p), π*(2p), σ*(2p)

📊 Bond Order Calculation:
Bond Order = ½(No. of electrons in bonding MOs - No. of electrons in antibonding MOs)

🎯 Stability Criteria:
• Bond Order > 0: Stable molecule
• Bond Order = 0: Unstable molecule
• Higher Bond Order = More stable
• Bond Order ≈ Bond Length (inverse relationship)

📊 Example: O₂ MO Configuration
- Total electrons: 12
- Bond order: ½(10 - 6) = 2
- Magnetic: Paramagnetic (2 unpaired electrons)
- Bond length: Intermediate

7. Hydrogen Bonding

Types and Characteristics

🤝 Hydrogen Bonding:
• H attached to highly electronegative atom (N, O, F)
• Interaction with lone pair on another electronegative atom
• Directional and relatively strong
• Important in biological systems

📊 Types of Hydrogen Bonds:
• Intermolecular: Between different molecules
  Examples: H₂O, NH₃, HF
• Intramolecular: Within same molecule
  Examples: o-nitrophenol, salicylic acid

⚡ Hydrogen Bond Properties:
• Strength: 5-30 kJ/mol (weaker than covalent, stronger than dipole)
• Directional: Linear arrangement preferred
• Affects boiling points and solubilities
• Important for DNA structure and protein folding

🔬 Factors Affecting Strength:
• Electronegativity of atoms involved
• Distance between atoms
• Solvent effects
• Temperature and pressure

📊 Advanced Concepts

1. Molecular Polarity

Determining Molecular Polarity

⚡ Polarity Determination:
• Depends on bond polarity and molecular geometry
• Vector sum of bond dipoles
• Symmetric molecules may be non-polar despite polar bonds

📊 Steps to Determine Polarity:
1. Identify electronegativity differences
2. Determine bond dipoles
3. Consider molecular geometry
4. Calculate vector sum of bond dipoles

🎯 Common Examples:
• Non-polar: CO₂ (linear, symmetric), CH₄ (tetrahedral)
• Polar: H₂O (bent, asymmetric), NH₃ (trigonal pyramidal)

⚡ Factors Affecting Polarity:
• Electronegativity difference
• Molecular geometry
• Lone pair presence
• Symmetry of molecule

2. Resonance

Resonance Structures

🔄 Resonance Concept:
• Multiple Lewis structures for same molecule
• Delocalization of electrons
• Actual structure is resonance hybrid
• Contributing structures differ only in electron position

📊 Resonance Rules:
• Octet rule must be satisfied
• Formal charges should be minimized
• Negative charges on electronegative atoms
• Positive charges on electropositive atoms
• Proper bond lengths

🎯 Resonance Examples:
• O₃ (ozone): 2 resonance structures
• NO₃⁻ (nitrate): 3 resonance structures
• Benzene (C₆H₆): 6 resonance structures (Kekulé forms)

⚡ Resonance Effects:
• Stabilization of molecules
• Equalization of bond lengths
• Increased stability
• Affects reactivity patterns

3. Bond Parameters

Bond Length, Energy, and Order

📏 Bond Length:
• Distance between nuclei of bonded atoms
• Inversely related to bond order
• Affected by atomic size
• Multiple bonds shorter than single bonds

⚡ Bond Energy:
• Energy required to break bond
• Directly related to bond order
• Multiple bonds stronger than single bonds
• Affected by atomic size and electronegativity

📊 Bond Order:
• Number of electron pairs shared
• Single bond: BO = 1
• Double bond: BO = 2
• Triple bond: BO = 3
• Fractional bond orders in resonance

🔗 Relationships:
• Higher bond order = shorter bond length
• Higher bond order = higher bond energy
• Larger atoms = longer bonds
• Higher electronegativity = shorter bonds

🎯 Problem-Solving Strategies

1. Lewis Structure Drawing

Step-by-Step Approach

📝 Systematic Method:
1. Count total valence electrons
2. Identify central atom
3. Draw skeletal structure
4. Complete octets of outer atoms
5. Complete octet of central atom
6. Form multiple bonds if needed
7. Calculate formal charges
8. Choose best structure

✅ Verification Checklist:
• All atoms have complete octets (except H)
• Total valence electrons used correctly
• Formal charges minimized
• Most electronegative atoms have negative charges
• Structure is chemically reasonable

2. Hybridization Determination

Quick Method

🎯 Hybridization Rules:
• Count sigma bonds + lone pairs around central atom
• 2: sp hybridization
• 3: sp² hybridization
• 4: sp³ hybridization
• 5: sp³d hybridization
• 6: sp³d² hybridization

📊 Example Determination:
NH₃:
• N has 3 sigma bonds + 1 lone pair = 4
• sp³ hybridization
• Trigonal pyramidal geometry

CO₂:
• C has 2 sigma bonds + 0 lone pairs = 2
• sp hybridization
• Linear geometry

3. VSEPR Geometry Prediction

Systematic Approach

📐 Geometry Prediction Steps:
1. Draw Lewis structure
2. Count electron domains around central atom
3. Determine electron domain geometry
4. Identify number of lone pairs
5. Predict molecular shape
6. Estimate bond angles

🎯 Bond Angle Approximations:
• Lone pairs repel more than bonding pairs
• Bond angles decrease with more lone pairs
• Multiple bonds repel more than single bonds
• Use VSEPR to predict approximate angles

📊 Common Angle Ranges:
• 180°: Linear molecules
• 120°: Trigonal planar
• 109.5°: Tetrahedral
• 90°: Octahedral and related

📈 Common Mistakes and How to Avoid Them

1. Lewis Structure Errors

Common Problems

❌ Frequent Mistakes:
• Incorrect counting of valence electrons
• Wrong choice of central atom
• Incomplete octets
• Incorrect formal charge calculations
• Ignoring resonance structures

✅ Correct Approach:
• Double-check electron counting
• Follow central atom selection rules
• Ensure octet completion
• Calculate formal charges systematically
• Consider all reasonable resonance structures

2. Hybridization Errors

Misconceptions

❌ Common Errors:
• Using electron domains instead of sigma bonds + lone pairs
• Confusing geometry with hybridization
• Not considering lone pairs in hybridization
• Wrong hybridization for expanded octet

✅ Correct Approach:
• Count sigma bonds + lone pairs
• Remember hybridization explains geometry
• Include lone pairs in domain count
• Use d orbitals for expanded octet cases

3. VSEPR Application Errors

Geometry Prediction Mistakes

❌ Common Problems:
• Ignoring lone pair repulsion effects
• Wrong bond angle predictions
• Confusing electron domain geometry with molecular shape
• Not considering molecular symmetry

✅ Correct Approach:
• Always consider lone pair repulsion
• Adjust angles based on lone pairs
• Distinguish electron domain from molecular geometry
• Consider overall molecular symmetry

🔗 Integration with Other Topics

1. Connection to Periodic Properties

Trends in Bonding

📊 Periodic Trends:
• Electronegativity increases across period
• Atomic size decreases across period
• Ionization energy increases across period
• These affect bond type and strength

🎯 Bond Type Predictions:
• Metal + non-metal: Ionic
• Non-metal + non-metal: Covalent
• Similar electronegativity: Covalent
• Large electronegativity difference: Ionic

2. Connection to Thermodynamics

Bond Energies and Reactions

⚡ Bond Energy in Reactions:
• ΔH = Σ(bonds broken) - Σ(bonds formed)
• Bond breaking requires energy (endothermic)
• Bond formation releases energy (exothermic)
• Used to calculate reaction enthalpy

📊 Applications:
• Predict reaction spontaneity
• Calculate activation energies
• Understand reaction mechanisms
• Compare reaction pathways

3. Connection to Spectroscopy

Molecular Structure Determination

🔬 Spectroscopic Methods:
• IR spectroscopy: Bond vibrations
• NMR: Molecular structure and bonding
• X-ray crystallography: Exact molecular geometry
• UV-Vis: Electronic transitions

📊 Structure-Property Relationships:
• Bond length affects vibrational frequency
• Molecular symmetry affects spectral patterns
• Bond polarity affects electronic transitions
• Geometry determines spectral selection rules

📊 Previous Year Questions Analysis

JEE Main & Advanced Pattern (2009-2024)

Question Distribution

📊 Topic-wise Distribution:
• Lewis structures and formal charge: 20%
• VSEPR theory and geometry: 30%
• Hybridization and orbital theory: 25%
• Molecular orbital theory: 15%
• Hydrogen bonding and polarity: 10%

📈 Difficulty Level:
• Easy: 35% (Basic concepts and direct applications)
• Medium: 45% (Multi-step problems and reasoning)
• Hard: 20% (Complex molecules and integrated concepts)

🎯 Frequently Asked Concepts:
• VSEPR geometry prediction
• Hybridization determination
• Lewis structure drawing
• MO theory for diatomic molecules
• Polarity and dipole moments

Common Question Types

📝 Type 1: Structure Determination
• Draw Lewis structures for given molecules
• Determine formal charges
• Identify best resonance structure
• Predict molecular geometry

📈 Type 2: Geometry and Hybridization
• Determine hybridization from structure
• Predict bond angles
• Explain molecular shapes
• Compare geometries of related molecules

🔄 Type 3: MO Theory Applications
• Draw MO diagrams for diatomic molecules
• Calculate bond orders
• Predict magnetic properties
• Compare stabilities

📊 Type 4: Advanced Concepts
• Explain hydrogen bonding patterns
• Determine molecular polarity
• Analyze resonance effects
• Apply bonding concepts to reactions

🎯 Type 5: Integrated Problems
• Combine bonding with thermodynamics
• Relate structure to properties
• Connect to periodic trends
• Apply to real-world molecules

🎯 Quick Reference Summary

Essential Formulas and Concepts

📊 Formal Charge:
FC = (Valence electrons) - (Non-bonding electrons) - ½(Bonding electrons)

📐 Bond Order (MO Theory):
BO = ½(No. of electrons in bonding MOs - No. of electrons in antibonding MOs)

🎯 Hybridization Rules:
• Electron domains = σ bonds + lone pairs
• 2 domains: sp
• 3 domains: sp²
• 4 domains: sp³
• 5 domains: sp³d
• 6 domains: sp³d²

📏 Bond Length Trends:
• BO ↑ → Bond Length ↓
• Atomic Size ↑ → Bond Length ↑
• Multiple bonds < Single bonds

Key Points to Remember

✅ Always check octet rule satisfaction
✅ Calculate formal charges for resonance structures
✅ Use VSEPR for geometry prediction
✅ Consider lone pair repulsion effects
✅ MO theory explains magnetic properties
✅ Hydrogen bonding requires H-N, H-O, or H-F
✅ Resonance stabilizes molecules

🔗 Practice Problems

Basic Level

1. Draw Lewis structure for CO₃²⁻ and determine formal charges
2. Predict geometry and hybridization of PCl₃
3. Determine if H₂O is polar or non-polar

Medium Level

1. Draw MO diagram for N₂ and calculate bond order
2. Compare bond angles in NH₃, H₂O, and CH₄
3. Determine hybridization of central atom in SF₄

Advanced Level

1. Explain why O₂ is paramagnetic using MO theory
2. Draw resonance structures for NO₃⁻ and determine best structure
3. Predict geometry and explain polarity of XeF₄

📱 Digital Resources

Interactive Learning Tools

• Lewis Structure Builder: Build and verify structures
• VSEPR Simulator: Visualize molecular geometry
• MO Diagram Builder: Create MO diagrams
• Hybridization Calculator: Determine hybridization

Video Resources

• Concept Videos: Detailed concept explanations
• Problem Solving Sessions: Step-by-step solutions
• 3D Visualization: Interactive molecular models

🎯 Master Chemical Bonding with comprehensive notes, visual aids, and extensive practice!

This complete revision guide will help you excel in JEE questions related to chemical bonding and molecular structure! 🚀

Chapter: Chemical Bonding and Molecular Structure | Class: 11 | Subject: Chemistry | Last Updated: October 2024