Chemical Bonding

Lecture Notes: Introduction to Chemical Bonding

1. Introduction to Chemical Bonding

Every substance in nature is made up of atoms.
Atoms rarely exist independently (except noble gases).
Most atoms combine with others to form molecules or compounds.
This combination occurs through the formation of chemical bonds.

Definition:
A chemical bond is the attractive force which holds two or more atoms together in a stable chemical species.

2. Importance of Bonding

Explains the existence of molecules and compounds in nature.
Determines the stability and reactivity of substances.
Explains the physical state (solid, liquid, gas) of different substances.
Determines the structure and shape of molecules.
Governs physical properties like melting point, boiling point, hardness, solubility, electrical conductivity.
Provides a basis for chemical reactions (breaking old bonds and forming new ones).

Example:

Sodium (Na) reacts with Chlorine (Cl₂) → Ionic bond → Stable NaCl crystal.
Hydrogen (H₂) forms by covalent bond between two H atoms.

3. Octet Rule

Proposed by Lewis and Kossel (1916).
Atoms tend to achieve eight electrons in their outermost shell (octet), resembling the noble gas configuration.
This gives the atom maximum stability.

Key Points:

Atoms may lose, gain, or share electrons to complete their octet.
Explains formation of ionic and covalent bonds.

Examples:

Na (2,8,1) loses 1e⁻ → Na⁺ (2,8) = octet.
Cl (2,8,7) gains 1e⁻ → Cl⁻ (2,8,8) = octet.
H (1) + H (1) → H₂ (duet rule for H).

Limitations:

Does not explain stability of molecules with odd number of electrons (NO, NO₂).
Does not account for molecules with incomplete octets (BF₃).
Cannot explain expanded octets (>8 electrons, e.g., SF₆, PCl₅).

4. Cause of Bond Formation

Atoms form bonds to attain greater stability by achieving a lower energy state.

(a) Stability through Noble Gas Configuration

Atoms combine to complete their octet/duet.
Stable noble gases rarely form compounds because their outer shells are already complete.

(b) Lowering of Potential Energy

When atoms bond, the overall system’s potential energy decreases.
A bonded state is more stable than isolated atoms.

Example:

In H₂ molecule, the overlap of atomic orbitals results in a stable bond with bond energy of 436 kJ/mol.

(c) Electrostatic Forces

Attraction between positive nucleus of one atom and electrons of another leads to bond formation.
Balance of attractive and repulsive forces keeps atoms at an equilibrium bond length.

Summary

Chemical bonding explains why atoms combine.
Importance: Stability, structure, and properties of matter.
Octet rule: Atoms tend to have 8 valence electrons (noble gas configuration).
Cause of bonding: Atoms bond to achieve lower energy and greater stability.

Here are lecture notes on Types of Chemical Bonds for Class 11 – ISC Chemistry.

Lecture Notes: Types of Chemical Bonds

1. Ionic Bond

(a) Definition

An ionic bond (also called electrovalent bond) is formed by the complete transfer of electrons from one atom to another, resulting in the formation of oppositely charged ions held together by electrostatic forces.

(b) Formation by Electron Transfer

Metals (low ionization energy) lose electrons → form cations.
Non-metals (high electron affinity) gain electrons → form anions.
Strong electrostatic attraction between cation and anion = ionic bond.

Example:
Na (2,8,1) + Cl (2,8,7) → Na⁺ (2,8) + Cl⁻ (2,8,8) → NaCl

(c) Lattice Energy

Defined as the energy released when 1 mole of gaseous ions combine to form 1 mole of ionic solid.
Higher lattice energy → stronger ionic bond → higher melting point.
Depends on charge of ions and distance between them (Coulomb’s law).

Example: MgO has higher lattice energy than NaCl due to 2+ and 2− ions.

(d) Properties of Ionic Compounds

Solid and crystalline at room temperature.
High melting and boiling points.
Conduct electricity in molten state or aqueous solution (ions free).
Generally soluble in polar solvents (e.g., water).
Brittle in nature (cleave along planes under stress).

2. Covalent Bond

(a) Definition

A covalent bond is formed by mutual sharing of electrons between two atoms, each contributing one or more electrons.

(b) Formation by Sharing of Electrons

Non-metallic atoms with high electronegativity share electrons.
Shared pair is counted for both atoms → achieving stable configuration.

Example: H₂ (each H shares 1e⁻ → duet), Cl₂ (each Cl shares 1e⁻ → octet).

(c) Types of Covalent Bonds

Single Bond – Sharing of one electron pair (2 electrons).
- Example: H₂, Cl₂, CH₄.
Double Bond – Sharing of two electron pairs (4 electrons).
- Example: O₂, CO₂.
Triple Bond – Sharing of three electron pairs (6 electrons).
- Example: N₂, C₂H₂.

(d) Lewis Structures

Diagrams showing valence electrons as dots and bonds as lines.
Help in visualizing shared pairs and lone pairs.

Examples:

H₂O:
H–O–H (two lone pairs on O)
CO₂:
O=C=O (each O double-bonded to C)

3. Coordinate (Dative) Bond

(a) Definition

A coordinate bond is a type of covalent bond in which both electrons in the shared pair are donated by one atom (donor), while the other atom (acceptor) contributes none.

(b) Donor–Acceptor Concept

Donor atom → has lone pair of electrons.
Acceptor atom → electron-deficient atom with vacant orbital.
After bond formation, it behaves like a normal covalent bond.

(c) Examples

Ammonium ion (NH₄⁺)
- NH₃ has a lone pair on N.
- H⁺ (no electron) accepts the lone pair.
- N → H coordinate bond forms NH₄⁺.
Sulphur dioxide (SO₂)
- S donates a lone pair to O forming a coordinate bond.
Boron trifluoride–ammonia complex (BF₃·NH₃)
- BF₃ has electron-deficient B atom.
- NH₃ donates lone pair from N.

Summary

Ionic bond: transfer of electrons → ions → lattice energy governs strength.
Covalent bond: sharing of electrons → single, double, triple bonds → represented by Lewis structures.
Coordinate bond: special covalent bond → donor provides both electrons.

Here are Lecture Notes on Theories of Bonding for Class 11 ISC Chemistry with structured explanations.

3. Theories of Bonding

A. Valence Bond Theory (VBT)

(1) Concept of Orbital Overlap

Proposed by Heitler and London.
A covalent bond is formed when two half-filled atomic orbitals overlap, and the electrons are paired.
Greater the overlap → stronger the bond.
Explains bond strength and bond length.

(2) Types of Overlap

Sigma (σ) Bond
- Formed by end-to-end (axial) overlap of orbitals.
- Stronger than π bond.
- Example: H–H in H₂, C–H in CH₄.
Pi (π) Bond
- Formed by lateral/sideways overlap of orbitals.
- Weaker than σ bond.
- Always formed in addition to a sigma bond (double or triple bonds).
- Example: O=O in O₂ (1σ + 1π), N≡N in N₂ (1σ + 2π).

B. Hybridization

(1) Definition

Hybridization is the mixing of atomic orbitals of nearly equal energy to form new equivalent hybrid orbitals with the same energy and shape.

(2) Types of Hybridization and Shapes (VSEPR Link)

Type	Orbital Mixing	Geometry	Example
sp	1 s + 1 p	Linear (180°)	BeCl₂, CO₂
sp²	1 s + 2 p	Trigonal planar (120°)	BF₃, C₂H₄
sp³	1 s + 3 p	Tetrahedral (109.5°)	CH₄, NH₃ (pyramidal), H₂O (bent)
sp³d	1 s + 3 p + 1 d	Trigonal bipyramidal (90°, 120°)	PCl₅
sp³d²	1 s + 3 p + 2 d	Octahedral (90°)	SF₆

VSEPR Theory (Valence Shell Electron Pair Repulsion) explains shapes by minimizing electron pair repulsions (bond pairs + lone pairs).
Example: NH₃ (sp³, but trigonal pyramidal due to 1 lone pair).

C. Molecular Orbital Theory (MOT)

(1) Basic Concept

Proposed by Mulliken and Hund.
Atomic orbitals of comparable energy combine to form molecular orbitals (MOs).
Two types of MOs:
1. Bonding orbital (lower energy, stable) → formed by constructive overlap.
2. Antibonding orbital (higher energy, unstable) → formed by destructive overlap.

(2) Energy-Level Diagram

For molecules up to O₂ (Z ≤ 8):
Order = σ(1s), σ*(1s), σ(2s), σ*(2s), σ(2px), [π(2py) = π(2pz)], [π*(2py) = π*(2pz)], σ*(2px)
For molecules after O₂ (Z > 8):
Order = σ(1s), σ*(1s), σ(2s), σ*(2s), [π(2py) = π(2pz)], σ(2px), [π*(2py) = π*(2pz)], σ*(2px)

(3) Bond Order

Bond Order=12(Nb−Na)\text{Bond Order} = \tfrac{1}{2} (N_b – N_a)

where, NbN_b = electrons in bonding orbitals, NaN_a = electrons in antibonding orbitals.

Bond order indicates stability and bond strength.
Higher bond order → stronger and shorter bond.

Examples:

H₂ → BO = 1 (stable)
He₂ → BO = 0 (unstable, molecule does not exist)
O₂ → BO = 2 (stable, paramagnetic due to unpaired electrons)
N₂ → BO = 3 (very strong, triple bond)

(4) Magnetic Properties

If all electrons are paired → diamagnetic (repelled by magnetic field).
If unpaired electrons present → paramagnetic (attracted to magnetic field).

Example:

O₂ is paramagnetic (2 unpaired electrons in π* orbitals).
N₂ is diamagnetic (all electrons paired).

Summary

VBT: Bond forms by overlap of half-filled orbitals (σ and π bonds).
Hybridization: Mixing of orbitals → specific geometries (explained by VSEPR).
MOT: Atomic orbitals combine → bonding and antibonding orbitals → explains bond order, stability, magnetism.

Here are Lecture Notes on Bond Parameters for Class 11 ISC Chemistry.

Bond Parameters

Bond parameters are measurable characteristics of chemical bonds. They help in understanding bond strength, stability, and geometry of molecules.

1. Bond Length

Definition: The average distance between the nuclei of two bonded atoms in a molecule.
Units: Picometre (pm), 1 pm = 10⁻¹² m.
Dependence:
- Type of bond: Single > Double > Triple (C–C: 154 pm, C=C: 134 pm, C≡C: 120 pm).
- Atomic size: Larger atoms → longer bond length.
- Bond order: Higher bond order → shorter bond.
Example: H–H bond length = 74 pm.

2. Bond Angle

Definition: The angle between two covalent bonds originating from the same atom.
Units: Degrees (°).
Dependence:
- Hybridization: sp (180°) > sp² (120°) > sp³ (109.5°).
- Lone pairs: Lone pairs repel more strongly → decrease bond angle.
- Multiple bonds: Greater electron density can slightly increase bond angle.
Examples:
- CH₄ = 109.5° (tetrahedral, sp³).
- H₂O = 104.5° (sp³, 2 lone pairs reduce angle).
- CO₂ = 180° (linear, sp).

3. Bond Enthalpy (Bond Energy)

Definition: The energy required to break one mole of a particular bond in gaseous state.
Units: kJ mol⁻¹.
Dependence:
- Higher bond order → higher bond enthalpy.
- Shorter bond length → stronger bond → more bond enthalpy.
Examples:
- H–H bond = 436 kJ/mol.
- C–H bond = 412 kJ/mol.
- C≡C bond enthalpy > C=C > C–C.

4. Bond Order

Definition: The number of chemical bonds between two atoms.
Formula (MOT): Bond Order=12(Nb−Na)\text{Bond Order} = \tfrac{1}{2}(N_b – N_a) where NbN_b = number of bonding electrons, NaN_a = number of antibonding electrons.
Dependence:
- Higher bond order → shorter, stronger bonds.
Examples:
- H₂ → BO = 1 (single bond).
- O₂ → BO = 2 (double bond).
- N₂ → BO = 3 (triple bond).
- He₂ → BO = 0 (molecule does not exist).

5. Polarity of Bonds

(a) Electronegativity Difference

If two atoms have different electronegativities → bond becomes polar.
Greater the difference → greater polarity.

(c) Examples:

HCl → polar (μ = 1.08 D).
CO₂ → non-polar overall (though each C=O bond is polar, the dipoles cancel out due to linear geometry).
H₂O → polar molecule (net dipole due to bent shape).

Summary

Bond length: Distance between nuclei (shorter = stronger).
Bond angle: Determines molecular shape (influenced by hybridization & lone pairs).
Bond enthalpy: Energy to break bond (greater for multiple bonds).
Bond order: Number of bonds (1, 2, 3 → increasing strength).
Bond polarity: Arises from electronegativity difference; measured by dipole moment.

Here are Lecture Notes on Polarity and Forces for Class 11 ISC Chemistry.

5. Polarity and Forces

1. Polar vs Nonpolar Covalent Bonds

Non-Polar Covalent Bond
- Formed when two atoms of same/similar electronegativity share electrons equally.
- No charge separation.
- Example: H₂, Cl₂, O₂, CH₄.
Polar Covalent Bond
- Formed when two atoms of different electronegativities share electrons unequally.
- Partial charges (δ⁺, δ⁻) develop → dipole moment.
- Example: HCl (Hδ⁺ – Clδ⁻), H₂O.

Key Point:

The greater the electronegativity difference, the more polar the bond.

2. Intermolecular Forces (Van der Waals’ Forces)

Intermolecular forces are weak attractive forces between molecules.
They determine physical properties like melting point, boiling point, solubility.

Types:

Dipole–dipole interaction
London dispersion forces
Hydrogen bonding

3. Dipole–Dipole Interaction

Present between polar molecules.
Positive end of one molecule attracts negative end of another.
Stronger than London dispersion forces, but weaker than covalent/ionic bonds.
Example: HCl molecules attract each other through dipole–dipole forces.

4. London Dispersion Forces (Instantaneous Dipole–Induced Dipole Forces)

Present between all molecules (polar and nonpolar).
Caused by instantaneous dipoles due to movement of electrons.
Strength increases with:
- Number of electrons (molar mass).
- Surface area of molecules.
Example: Noble gases (He, Ne, Ar) and nonpolar molecules (I₂, CH₄) held together only by London forces.

5. Hydrogen Bonding

Special, stronger type of dipole–dipole attraction.
Occurs when H is covalently bonded to highly electronegative atom (N, O, F) and attracted to lone pair of another electronegative atom.
Stronger than dipole–dipole, but weaker than covalent/ionic bond.

Types:

Intermolecular Hydrogen Bonding
- Between molecules.
- Example: H₂O (responsible for high boiling point), HF, NH₃.
Intramolecular Hydrogen Bonding
- Within the same molecule.
- Example: o-nitrophenol (–OH group H bonds with NO₂ group oxygen).

Consequences of H-Bonding:

High boiling/melting points (e.g., water, HF).
Anomalous expansion of water (ice is less dense).
Increased solubility of compounds with –OH, –NH₂ groups.

Summary

Polar bonds: unequal electron sharing; nonpolar bonds: equal sharing.
Dipole–dipole forces: attraction between polar molecules.
London forces: weakest, arise from temporary dipoles, present in all molecules.
Hydrogen bonding: strong dipole–dipole force with H attached to N, O, or F; explains special properties of water and biological molecules (DNA base pairing).

Here are Lecture Notes on Theories of Shapes of Molecules (VSEPR Theory) for Class 11 ISC Chemistry.

Theories of Shapes of Molecules

1. Valence Shell Electron Pair Repulsion (VSEPR) Theory

(a) Basic Concept

Proposed by Sidgwick and Powell.
Shape of a molecule depends on repulsion between electron pairs (bond pairs and lone pairs) around the central atom.
Electron pairs arrange themselves as far apart as possible to minimize repulsion → determines molecular geometry.

(b) Key Postulates

Both bond pairs (bp) and lone pairs (lp) of electrons repel each other.
The repulsive strength order:
lp–lp > lp–bp > bp–bp.
Lone pairs distort bond angles more strongly.
Multiple bonds are treated as a single region of electron density, but exert slightly more repulsion.

2. AXn Notation (VSEPR Formula)

General form: AXₙEₘ
- A = central atom
- X = number of atoms bonded to central atom
- n = number of bond pairs
- E = number of lone pairs on central atom
- m = number of lone pairs

Example:

CH₄ → AX₄ (tetrahedral, no lone pairs).
NH₃ → AX₃E (pyramidal, 1 lone pair).
H₂O → AX₂E₂ (bent, 2 lone pairs).

3. Common Molecular Geometries

Geometry	AXE Formula	Bond Angle(s)	Examples
Linear	AX₂	180°	BeCl₂, CO₂
Trigonal Planar	AX₃	120°	BF₃, SO₃
Tetrahedral	AX₄	109.5°	CH₄, CCl₄
Trigonal Bipyramidal	AX₅	120° (equatorial), 90° (axial)	PCl₅
Octahedral	AX₆	90°	SF₆
Bent (Angular)	AX₂E / AX₂E₂	<120° (SO₂), <109.5° (H₂O)	H₂O, SO₂
Trigonal Pyramidal	AX₃E	~107°	NH₃, PCl₃

4. Notes on Shapes

Linear → central atom has 2 bond pairs, no lone pairs (CO₂).
Trigonal planar → 3 bond pairs, planar (BF₃).
Tetrahedral → 4 bond pairs (CH₄); perfect symmetry.
Pyramidal → 3 bond pairs + 1 lone pair (NH₃); bond angle reduced to ~107°.
Bent → 2 bond pairs + lone pairs (H₂O); angle reduced to ~104.5°.
Trigonal bipyramidal → 5 bond pairs, expanded octet (PCl₅).
Octahedral → 6 bond pairs, highly symmetrical (SF₆).

Summary

VSEPR theory predicts shapes of molecules based on repulsion between electron pairs.
AXₙEₘ notation is used to classify molecules.
Common geometries: linear, trigonal planar, tetrahedral, trigonal bipyramidal, octahedral, bent, pyramidal.
Lone pairs cause distortion from ideal bond angles.

Here are Lecture Notes on Special Topics in Chemical Bonding for Class 11 ISC Chemistry.

Special Topics in Chemical Bonding

1. Resonance Structures

(a) Concept

Some molecules cannot be represented by a single Lewis structure.
Instead, two or more structures are drawn (resonating structures) with same arrangement of atoms but different electron distribution.
The actual molecule is a resonance hybrid, which is more stable than any individual structure.

(b) Conditions for Resonance

Same arrangement of atoms.
Only distribution of electrons (π or lone pairs) differs.
Resonance structures must have similar energies.

(c) Examples

O₃ (ozone):
O=O–O ↔ O–O=O
CO₃²⁻ (carbonate ion):
Three equivalent structures with delocalized π electrons.
Benzene (C₆H₆):
Two equivalent Kekulé structures, true structure is delocalized ring.

(d) Importance

Explains stability (resonance energy).
Explains equal bond lengths in molecules like CO₃²⁻ and benzene.

2. Partial Ionic Character of Covalent Bonds (Fajan’s Rule)

(a) Concept

Even in covalent bonds, electron sharing may be unequal, giving the bond partial ionic character.
Ionic compounds may also have covalent character if the cation distorts the anion’s electron cloud.

(b) Fajan’s Rule (Polarizing Power and Polarizability)

Cation factors (polarizing power):
- Smaller size → higher polarizing power.
- Higher charge → higher polarizing power.
Anion factors (polarizability):
- Larger size → easier to distort → more covalent character.
- Higher charge → more polarizable.
Covalent character increases when:
- Small, highly charged cation + large, highly charged anion.

(c) Examples

LiI more covalent than NaI (Li⁺ smaller, more polarizing).
AlCl₃ shows covalent character though formed from metal + nonmetal.

3. Metallic Bonding

(a) Electron Sea Model

Metals consist of a lattice of positive metal ions surrounded by a “sea” of delocalized valence electrons.
Metallic bond = electrostatic attraction between cations and delocalized electrons.
Explains:
- Electrical conductivity (free electrons move under electric field).
- Malleability and ductility (layers of metal ions slide while bond remains intact).
- Metallic lustre (electrons absorb/emit light).

(b) Band Theory of Metals

Extends the electron sea model using quantum mechanics.
Atomic orbitals overlap to form bands:
- Valence band: highest range of energies filled with electrons.
- Conduction band: higher energy levels where electrons can move freely.
Metal: conduction band overlaps with valence band → free electron flow.
Insulator: large energy gap (band gap) between valence and conduction bands.
Semiconductor: small band gap → conductivity increases with temperature or doping.

Summary

Resonance: actual molecule = hybrid of multiple structures; stabilizes molecules like CO₃²⁻ and benzene.
Partial ionic character (Fajan’s rule): cation polarizes anion → bond has covalent nature; depends on charge and size of ions.
Metallic bonding: explained by electron sea model (delocalized electrons) and band theory (electronic band structures); explains conductivity, ductility, and metallic properties.

Video Lectures

Lewis Structure

Concept Map

Lecture Notes: Lewis Structure

1. Introduction

Lewis structure (also called electron dot structure) represents the arrangement of valence electrons around atoms in a molecule.
Developed by Gilbert N. Lewis (1916).
Shows bonding pairs (shared electrons) and lone pairs (non-bonded electrons).

2. Basic Rules

Valence electrons: Count total valence electrons of all atoms.
Central atom: Usually the least electronegative atom (except H, which is always terminal).
Bonds first: Place single bonds between central atom and surrounding atoms.
Complete octets: Distribute remaining electrons to satisfy octet (duet for H).
Check electron count: Ensure total electrons used = total valence electrons.
Multiple bonds: If atoms lack octet, form double or triple bonds.

3. Key Concepts

Bonding pair: Shared electrons (represented by a line “—”).
Best Lewis structure → minimal formal charges.
Lone pair: Non-shared electrons (represented by dots).

Formal Charge (FC)

4. Stepwise Example: CO₂

Count valence electrons: C (4) + 2 × O (6) = 16.
Skeleton: O—C—O.
Place bonds: 2 bonds (4 e⁻ used).
Remaining = 12 e⁻ → place on O atoms.
Central C incomplete → make double bonds: O=C=O.
Check: All atoms have octet, no formal charges.

5. Exceptions to Octet Rule

Incomplete octet: e.g., BeCl₂, BF₃.
Expanded octet: Atoms with d-orbitals (e.g., PCl₅, SF₆).
Odd-electron species: e.g., NO, NO₂.

6. Importance

Predicts molecular geometry (basis of VSEPR theory).
Helps in determining bond order, polarity, and reactivity.
Useful in acid-base chemistry, coordination complexes, and resonance structures.