⚛️ Introduction to Chemical Bonding
What is a Chemical Bond?
A chemical bond is the force that holds together two or more atoms, molecules or ions. The properties of a substance depend on the type of the chemical bond between its atoms.
The term chemical bond includes ionic, covalent, dative, metallic bonds, as well as intermolecular forces (van der Waals forces). However, being weak enough, van der Waals forces are usually not termed as pure chemical bonds.
🌟 Key Insight: NaCl has ionic bond and is solid, but water is a covalent compound and liquid. This shows how bond type affects physical properties!
Historical Perspective
Lewis concept of bonding gives a simple explanation of the formation of all types of bonds. According to this theory, atoms make bonds to complete their outermost shells to have noble gas like configuration. This is mostly attained through the formation of an octet in the valence shell.
🔗 Types of Chemical Bonds
Ionic Bond
According to the Lewis theory, the ionic bond is formed by the complete transfer of electrons from an atom with low ionization energy to another atom with high electron affinity.
Formation of NaCl:
1. Na atom (2,8,1) loses one electron to form Na⁺ (2,8)
2. Cl atom (2,8,7) gains one electron to form Cl⁻ (2,8,8)
3. Oppositely charged ions are held together by strong electrostatic forces
Covalent Bond
A covalent bond is formed by the mutual sharing of electrons between two atoms. While sharing electrons, each atom completes its valence shell and attains the nearest noble gas configuration.
| Bond Type | Description | Example |
|---|---|---|
| Non-polar Covalent | Equal sharing of electrons between identical atoms | H₂, Cl₂, O₂ |
| Polar Covalent | Unequal sharing due to electronegativity difference | HCl, H₂O, NH₃ |
Dative/Coordinate Covalent Bond
A dative bond is formed between two atoms when the shared pair of electrons is donated by one of the bonding atoms. This bond is formed by overlapping of one completely filled orbital with one empty orbital.
Examples:
1. H₃O⁺ (Hydronium ion)
2. CO (Carbon monoxide)
3. O₃ (Ozone)
4. NH₄⁺ (Ammonium ion)
⚡ Electronegativity & Bond Polarity
Electronegativity Scale
The difference in electronegativity of two bonded atoms provides an approximate measure of the bond polarity and an indication of the type of bond.
| ΔEN Range | Bond Type | Examples |
|---|---|---|
| 0 – 0.4 | Non-polar Covalent | H-H, Cl-Cl |
| 0.4 – 1.8 | Polar Covalent | H-Cl, H₂O |
| ＞1.8 | Ionic | NaCl, KF |
Dipole Moment
A molecule with δ⁺ charge on one part and δ⁻ charge on the other part is called a dipole and such a molecule is said to have a dipole moment. It is a quantitative measurement of the polarity of a bond or a molecule.
💡 Important: CO has a dipole moment (0.12 D) but CO₂ does not. This is because CO₂ is linear and symmetrical, while CO has an asymmetric electron distribution.
🔺 VSEPR Theory
Basic Postulates
Valence Shell Electron Pair Repulsion (VSEPR) model describes the shapes of molecules based on the electron pairs that surround the central atom. The main postulates are:
1. Both lone pairs and bond pairs participate in determining molecular geometry
2. Electron pairs arrange themselves to maximize distance and minimize repulsion
3. Lone pairs occupy more space than bond pairs
4. Repulsion order: lp-lp ＞ lp-bp ＞ bp-bp
Common Molecular Shapes
| Electron Pairs | Lone Pairs | Shape | Examples | Bond Angle |
|---|---|---|---|---|
| 2 | 0 | Linear | BeCl₂, CO₂ | 180° |
| 3 | 0 | Trigonal Planar | BF₃, SO₃ | 120° |
| 4 | 0 | Tetrahedral | CH₄, CCl₄ | 109.5° |
| 4 | 1 | Trigonal Pyramidal | NH₃, H₃O⁺ | ＜109.5° |
| 4 | 2 | Bent/V-shaped | H₂O, OF₂ | ＜109.5° |
| 5 | 0 | Trigonal Bipyramidal | PCl₅, I₃⁻ | 120° & 90° |
| 6 | 0 | Octahedral | SF₆, XeF₄ | 90° |
🧬 Atomic Orbital Hybridization
What is Hybridization?
A process in which atomic orbitals of slightly different energies and shapes are mixed together to form a new set of equivalent orbitals of same energy and same shape is called hybridization.
Types of Hybridization
| Type | Orbitals Mixed | Geometry | Bond Angle | Examples |
|---|---|---|---|---|
| sp | 1s + 1p | Linear | 180° | BeCl₂, CO₂ |
| sp² | 1s + 2p | Trigonal Planar | 120° | BF₃, SO₃ |
| sp³ | 1s + 3p | Tetrahedral | 109.5° | CH₄, NH₃, H₂O |
| sp³d | 1s + 3p + 1d | Trigonal Bipyramidal | 120° & 90° | PCl₅ |
| sp³d² | 1s + 3p + 2d | Octahedral | 90° | SF₆ |
Determining Hybridization
To determine hybridization: Count the number of atoms bonded to central atom + number of lone pairs on central atom.
Examples:
CH₄: 4 bonds + 0 lone pairs = 4 → sp³
NH₃: 3 bonds + 1 lone pair = 4 → sp³
H₂O: 2 bonds + 2 lone pairs = 4 → sp³
BF₃: 3 bonds + 0 lone pairs = 3 → sp²
BeCl₂: 2 bonds + 0 lone pairs = 2 → sp
🌌 Molecular Orbital Theory (MOT)
Key Postulates
1. Atomic orbitals overlap to form molecular orbitals characteristic of the whole molecule
2. Two atomic orbitals form two molecular orbitals: bonding (lower energy) and antibonding (higher energy)
3. Bond order = ½ (No. of bonding electrons – No. of antibonding electrons)
4. MOT successfully explains paramagnetism of O₂
MO Diagrams Comparison
| Molecule | Bond Order | Magnetic Property | Bond Type | Stability |
|---|---|---|---|---|
| H₂ | 1 | Diamagnetic | Single | Stable |
| He₂ | 0 | Diamagnetic | None | Not formed |
| N₂ | 3 | Diamagnetic | Triple | Very stable |
| O₂ | 2 | Paramagnetic | Double | Stable |
| F₂ | 1 | Diamagnetic | Single | Less stable |
🎯 Critical Insight: MOT explains why O₂ is paramagnetic (has unpaired electrons) while N₂ is diamagnetic. This was a major success of MOT over VBT.
🚀 Study Strategies for Chemical Bonding
Master Bond Type Identification
Learn to calculate electronegativity difference (ΔEN) to predict bond type: ΔEN ＜ 0.4 = non-polar covalent; 0.4-1.8 = polar covalent; ＞1.8 = ionic.
VSEPR Shape Prediction
Practice counting electron pairs (bond pairs + lone pairs) to predict molecular shapes. Remember: lone pairs cause bond angle compression.
Hybridization Patterns
Memorize the correlation: 2 electron groups → sp; 3 groups → sp²; 4 groups → sp³; 5 groups → sp³d; 6 groups → sp³d².
MOT Applications
Focus on calculating bond orders and predicting magnetic properties. Remember: Bond order = ½(bonding e⁻ – antibonding e⁻).