Notes

Introduction to Chemical Bonding
Kossel-Lewis Theory & Octet Rule
Lewis Structures & Formal Charge
Ionic Bond
Covalent Bond & Bond Parameters
Resonance & Polarity
VSEPR Theory
Valence Bond Theory
Hybridisation (sp, sp², sp³, sp³d, sp³d²)
Molecular Orbital Theory
Hydrogen Bonding
Important Tables, Differences, and Exam Questions

Part 1: Introduction, Kossel-Lewis Theory & Octet Rule

1. Chemical Bond

Definition

A chemical bond is the attractive force that holds atoms or ions together to form molecules or compounds.

Why do atoms form chemical bonds?

Atoms combine to:

Achieve maximum stability.
Complete their outermost shell.
Lower their energy.

Stable electronic configuration = Noble gas configuration (Duplet or Octet).

2. Types of Chemical Bonds

Ionic (Electrovalent) Bond
Covalent Bond
Coordinate (Dative) Bond
Metallic Bond
Hydrogen Bond

3. Kossel-Lewis Theory (1916)

Main Idea

Atoms become stable by attaining the electronic configuration of the nearest noble gas.

This stability is achieved by:

Loss or gain of electrons → Ionic bond
Sharing of electrons → Covalent bond

Important Points

Noble gases are chemically stable because their outer shell is complete.
Metals usually lose electrons.
Non-metals usually gain or share electrons.

4. Valence Electrons

Valence electrons are the electrons present in the outermost shell of an atom.

They participate in chemical bonding.

Examples

Element	Valence Electrons
H	1
C	4
N	5
O	6
F	7
Na	1
Mg	2
Cl	7

5. Lewis Symbols (Electron Dot Symbols)

Lewis symbols represent the valence electrons of an atom using dots around its symbol.

Examples

H •
Na •
C ••••
N •••••
O ••••••
F •••••••
Ne ••••••••

Importance

Shows the number of valence electrons.
Helps predict bonding.
Used to draw Lewis structures.

6. Ionic (Electrovalent) Bond

Definition

An ionic bond is formed by the complete transfer of one or more electrons from one atom to another.

Formation Steps

Metal loses electron(s).
Non-metal gains electron(s).
Oppositely charged ions attract each other.

Example: Sodium Chloride (NaCl)

Na → Na⁺ + e⁻

Cl + e⁻ → Cl⁻

Na⁺ + Cl⁻ → NaCl

Characteristics of Ionic Compounds

Formed between metal and non-metal.
High melting and boiling points.
Conduct electricity in molten or aqueous state.
Soluble in water.
Hard and brittle.

7. Octet Rule

Definition

Atoms tend to gain, lose, or share electrons to obtain 8 electrons in their valence shell, similar to noble gases.

Hydrogen follows the duplet rule (2 electrons).

Examples

Na → loses 1 electron
Cl → gains 1 electron
CH₄ → shares electrons
H₂O → shares electrons

8. Limitations of Octet Rule

The octet rule is not applicable to all molecules.

(A) Incomplete Octet

Central atom has fewer than 8 electrons.

Examples:

BeH₂
BF₃
AlCl₃

(B) Odd Electron Molecules

Total number of electrons is odd.

Examples:

NO
NO₂

(C) Expanded Octet

Central atom has more than 8 electrons.

Examples:

PF₅
SF₆
H₂SO₄

9. Important Exam Points

Noble gases are stable due to complete outer shells.
Metals generally lose electrons.
Non-metals generally gain or share electrons.
Lewis symbols show only valence electrons.
Ionic bond is formed by electron transfer.
Covalent bond is formed by electron sharing.
Hydrogen follows the duplet rule.
Octet rule has three important exceptions:
- Incomplete octet
- Odd-electron molecules
- Expanded octet

Quick Revision

✔ Chemical bond → Force holding atoms together

✔ Stability → Complete outer shell

✔ Lewis symbols → Representation of valence electrons

✔ Ionic bond → Electron transfer

✔ Covalent bond → Electron sharing

✔ Octet Rule → 8 electrons in outer shell

✔ Duplet Rule → Hydrogen needs 2 electrons

✔ Exceptions → BeH₂, BF₃, NO, NO₂, PF₅, SF₆

Part 2: Covalent Bond, Lewis Structures & Formal Charge

10. Covalent Bond

Definition

A covalent bond is a chemical bond formed by the mutual sharing of electron pairs between two atoms.

It usually occurs between non-metal atoms.

Examples:

H₂
Cl₂
O₂
N₂
CH₄
H₂O

Formation of Covalent Bond

Atoms share electrons to complete their outer shell and achieve noble gas configuration.

Example: Chlorine molecule (Cl₂)

Electronic configuration of Cl:

Cl = 2, 8, 7

Each chlorine atom needs one electron to complete its octet.

Two chlorine atoms share one electron pair:

Cl — Cl

This forms a single covalent bond.

Types of Covalent Bonds

1. Single Bond

One pair of electrons is shared.
Represented by (—)

Examples:

H₂
Cl₂
CH₄

Bond order = 1

2. Double Bond

Two pairs of electrons are shared.
Represented by (=)

Examples:

O₂
CO₂
C₂H₄

Bond order = 2

3. Triple Bond

Three pairs of electrons are shared.
Represented by (≡)

Examples:

N₂
C₂H₂

Bond order = 3

11. Lewis Structures

Definition

Lewis structures are diagrams that show:

Arrangement of atoms
Shared electron pairs
Lone pairs of electrons

Rules for Drawing Lewis Structures

Step 1:

Calculate total valence electrons.

Example:
CH₄

Carbon = 4 electrons
4 Hydrogen = 4 electrons

Total = 8 electrons

Step 2:

Select central atom.

General rule:

Least electronegative atom is usually the central atom.
Hydrogen is never the central atom.

Examples:

Molecule	Central Atom
NH₃	N
H₂O	O
CO₂	C
BF₃	B

Step 3:

Join atoms using single bonds.

Each bond contains 2 electrons.

Step 4:

Complete octets of surrounding atoms.

Step 5:

Place remaining electrons on central atom.

If octet is incomplete, form multiple bonds.

Examples of Lewis Structures

1. Hydrogen Molecule (H₂)

H + H → H : H

One shared pair = Single bond

2. Oxygen Molecule (O₂)

Oxygen has 6 valence electrons.

Each oxygen shares two electrons.

O = O

Two shared pairs = Double bond

3. Nitrogen Molecule (N₂)

Nitrogen has 5 valence electrons.

Three electron pairs are shared.

N ≡ N

Three shared pairs = Triple bond

12. Formal Charge

Definition

Formal charge is the charge assigned to an atom in a molecule assuming equal sharing of bonding electrons.

Formula:

$\text{Formal Charge} = \text{Valence electrons} – \text{Non-bonding electrons} – \frac{1}{2}(\text{Bonding electrons})$ Formal Charge=Valence electrons−Non-bonding electrons−21(Bonding electrons)

Importance of Formal Charge

Helps select the most stable Lewis structure.
Helps compare different possible structures.
The structure with minimum formal charges is generally more stable.

Example: Oxygen in Ozone (O₃)

Ozone has two possible Lewis structures.

Formal charges help identify the more suitable structure.

13. Resonance

Definition

When a molecule cannot be represented by a single Lewis structure and requires two or more structures, the phenomenon is called resonance.

The actual molecule is a resonance hybrid.

Features of Resonance

Resonance structures have the same arrangement of atoms.
Only electron positions change.
The actual molecule is more stable than individual structures.
Resonance decreases energy and increases stability.

Examples of Resonance

1. Ozone (O₃)

Ozone has two equivalent structures.

Actual structure is a hybrid of both.

2. Carbonate Ion (CO₃²⁻)

Three resonance structures are possible.

All C–O bonds become equivalent.

3. Carbon Dioxide (CO₂)

CO₂ shows resonance due to different possible electron arrangements.

14. Bond Parameters

Bond parameters describe the characteristics of chemical bonds.

Main bond parameters:

Bond length
Bond angle
Bond enthalpy
Bond order

14.1 Bond Length

Definition

The equilibrium distance between the nuclei of two bonded atoms is called bond length.

Unit:

Picometre (pm)

Important Relation:

Higher bond order → Shorter bond length
Lower bond order → Longer bond length

Example:

C≡C < C=C < C–C

(Length increases)

14.2 Bond Angle

Definition

The angle between two bonds around the central atom is called bond angle.

Examples:

Molecule	Bond Angle
CH₄	109.5°
NH₃	107°
H₂O	104.5°
CO₂	180°

Bond angle determines molecular shape.

14.3 Bond Enthalpy

Definition

The energy required to break one mole of a particular bond in gaseous state is called bond enthalpy.

Unit:

kJ mol⁻¹

Important Point:

Higher bond enthalpy = Stronger bond

Example:

N≡N has very high bond enthalpy because it has a triple bond.

14.4 Bond Order

Definition

Bond order is the number of electron pairs shared between two atoms.

Examples:

Molecule	Bond Order
H₂	1
O₂	2
N₂	3
CO	3

Relationship:

Higher bond order:

Stronger bond
Higher bond enthalpy
Shorter bond length

Quick Revision

✔ Covalent bond = sharing of electrons
✔ Single bond = one shared pair
✔ Double bond = two shared pairs
✔ Triple bond = three shared pairs
✔ Lewis structure shows bonding and lone pairs
✔ Formal charge helps choose stable structures
✔ Resonance increases stability
✔ Higher bond order means stronger and shorter bond

Part 3: Ionic Bond, Lattice Enthalpy, Polarity of Bonds & Dipole Moment

15. Ionic (Electrovalent) Bond

Definition

An ionic bond is the electrostatic attraction between positively charged ions (cations) and negatively charged ions (anions).

It is formed by transfer of electrons from one atom to another.

Formation of Ionic Bond

Conditions required:

Low ionisation enthalpy
- Atom should easily lose electrons.
- Usually metals.
High negative electron gain enthalpy
- Atom should easily accept electrons.
- Usually non-metals.
High lattice enthalpy
- Provides stability to ionic crystals.

Example: Formation of NaCl

Step 1: Formation of cation

Na → Na⁺ + e⁻

(Sodium loses one electron)

Step 2: Formation of anion

Cl + e⁻ → Cl⁻

(Chlorine gains one electron)

Step 3: Formation of ionic compound

Na⁺ + Cl⁻ → NaCl

The attraction between Na⁺ and Cl⁻ forms an ionic bond.

16. Characteristics of Ionic Compounds

1. Physical State

Usually crystalline solids.

2. Melting and Boiling Points

High due to strong electrostatic attraction.

3. Solubility

Generally soluble in polar solvents like water.

4. Electrical Conductivity

Conduct electricity in:
- Molten state
- Aqueous solution
Do not conduct in solid state because ions cannot move freely.

5. Hardness

Hard and brittle due to strong ionic forces.

17. Lattice Enthalpy

Definition

The energy required to separate one mole of an ionic solid into gaseous ions is called lattice enthalpy.

Example:

NaCl(s) → Na⁺(g) + Cl⁻(g)

Lattice enthalpy of NaCl:

+788 kJ mol⁻¹

Factors Affecting Lattice Enthalpy

1. Size of Ions

Smaller ions → Stronger attraction → Higher lattice enthalpy

Example:

LiF > CsF

2. Charge on Ions

Higher charge → Stronger attraction

Example:

MgO > NaCl

(Mg²⁺ and O²⁻ have higher charges)

3. Arrangement of Ions

Better packing of ions increases lattice stability.

18. Polarity of Bonds

Non-polar Covalent Bond

When two identical atoms share electrons equally.

Examples:

H₂
O₂
Cl₂
N₂

No charge separation occurs.

Polar Covalent Bond

When electrons are shared unequally due to difference in electronegativity.

Examples:

HCl
HF
H₂O

One atom becomes slightly negative (δ⁻) and the other slightly positive (δ⁺).

19. Electronegativity and Bond Polarity

Greater difference in electronegativity:

→ More polar bond

Example:

HF is more polar than HCl because fluorine is more electronegative.

20. Dipole Moment

Definition

Dipole moment measures the polarity of a bond or molecule.

It is represented by: $\mu = q \times r$ μ=q×r

Where:

μ = Dipole moment
q = Charge
r = Distance between charges

Unit

Debye (D)

Conversion:

1 D = 3.336 × 10⁻³⁰ C m

Factors Affecting Dipole Moment

1. Difference in Electronegativity

Greater difference → Higher dipole moment

2. Molecular Shape

Shape determines whether bond dipoles cancel or add.

Examples of Dipole Moment

1. H₂O

Bent shape
Dipole moments do not cancel
Polar molecule

Dipole moment = 1.85 D

2. CO₂

Linear shape
Two C=O dipoles cancel each other

Dipole moment = 0

3. BF₃

Trigonal planar
Dipoles cancel

Dipole moment = 0

4. NH₃

Pyramidal shape
Polar molecule

Dipole moment ≠ 0

21. Fajan’s Rules (Covalent Character in Ionic Bonds)

According to Fajan’s rules, ionic compounds can show covalent character.

Factors increasing covalent character:

1. Small Cation Size

Small cations have greater polarising power.

Example:

AlCl₃ is more covalent than NaCl.

2. Large Anion Size

Large anions are easily distorted.

Example:

LiI is more covalent than LiF.

3. Higher Charge on Cation

Higher charge increases covalent nature.

Example:

MgCl₂ is more covalent than NaCl.

Quick Revision

✔ Ionic bond = attraction between oppositely charged ions
✔ Metal loses electrons → Cation
✔ Non-metal gains electrons → Anion
✔ Lattice enthalpy decides stability of ionic compounds
✔ Smaller ions have higher lattice enthalpy
✔ Equal sharing → Non-polar bond
✔ Unequal sharing → Polar bond
✔ Dipole moment shows polarity
✔ Molecular shape decides overall polarity
✔ CO₂ and BF₃ are non-polar due to cancellation of dipoles

Part 4: VSEPR Theory and Molecular Shapes

22. Valence Shell Electron Pair Repulsion (VSEPR) Theory

Definition

VSEPR theory explains the shape of molecules based on repulsion between electron pairs present in the valence shell of the central atom.

It was proposed by:

Sidgwick and Powell (1940)
Modified by Nyholm and Gillespie (1957)

Main Postulates of VSEPR Theory

1. Electron Pair Repulsion

Electron pairs around the central atom repel each other because electrons have negative charge.

2. Maximum Separation

Electron pairs arrange themselves as far apart as possible to reduce repulsion.

3. Shape of Molecule

Molecular shape depends on:
- Number of bond pairs
- Number of lone pairs

4. Multiple Bonds

A double or triple bond is treated as one electron pair.

5. Lone Pair Effect

Lone pairs occupy more space than bond pairs and cause greater repulsion.

Order of Repulsion

The repulsion decreases in the order:

Lone pair – Lone pair > Lone pair – Bond pair > Bond pair – Bond pair

Reason:

Lone pair electrons are attracted by only one nucleus, so they spread more.

23. Factors Affecting Molecular Shape

The shape depends on:

Number of electron pairs around central atom
Lone pairs present
Bond pair arrangement
Size of surrounding atoms

24. Shapes of Molecules Without Lone Pair

(A) Linear Shape

General Formula:

AB₂

Electron pairs:

2 bond pairs

Bond Angle:

180°

Examples:

BeCl₂
CO₂

Structure:

B — A — B

(B) Trigonal Planar Shape

General Formula:

AB₃

Electron pairs:

3 bond pairs

Bond Angle:

120°

Example:

BF₃

Shape:

Flat triangular arrangement

(C) Tetrahedral Shape

General Formula:

AB₄

Electron pairs:

4 bond pairs

Bond Angle:

109.5°

Example:

CH₄

(D) Trigonal Bipyramidal Shape

General Formula:

AB₅

Electron pairs:

5 bond pairs

Bond Angle:

90° and 120°

Example:

PCl₅

(E) Octahedral Shape

General Formula:

AB₆

Electron pairs:

6 bond pairs

Bond Angle:

90°

Example:

SF₆

25. Effect of Lone Pair on Molecular Shape

Lone pairs:

Occupy more space.
Increase repulsion.
Decrease bond angles.

Molecules Having Lone Pairs

1. Ammonia (NH₃)

Central Atom:

Nitrogen

Electron pairs:

3 bond pairs
1 lone pair

Shape:

Trigonal pyramidal

Bond Angle:

107°

Reason:

Lone pair-bond pair repulsion reduces angle from 109.5°.

2. Water (H₂O)

Central Atom:

Oxygen

Electron pairs:

2 bond pairs
2 lone pairs

Shape:

Bent / Angular

Bond Angle:

104.5°

Reason:

Two lone pairs cause strong repulsion.

3. Ammonium Ion (NH₄⁺)

Electron pairs:

4 bond pairs

Shape:

Tetrahedral

Bond Angle:

109.5°

4. Methane (CH₄)

Electron pairs:

4 bond pairs

Shape:

Tetrahedral

Bond Angle:

109.5°

Reason:

All electron pairs are equivalent.

26. Important Molecular Shapes Table

Molecule	Electron Pair	Lone Pair	Shape	Bond Angle
BeCl₂	2	0	Linear	180°
BF₃	3	0	Trigonal planar	120°
CH₄	4	0	Tetrahedral	109.5°
NH₃	4	1	Trigonal pyramidal	107°
H₂O	4	2	Bent	104.5°
PCl₅	5	0	Trigonal bipyramidal	90°,120°
SF₆	6	0	Octahedral	90°

27. Important VSEPR Examples

CO₂

Central atom: Carbon
Two double bonds
No lone pair
Shape: Linear
Dipole moment: Zero

BF₃

Central atom: Boron
Three bond pairs
No lone pair
Shape: Trigonal planar
Non-polar

NH₃

Nitrogen has one lone pair
Shape: Pyramidal
Polar molecule

H₂O

Oxygen has two lone pairs
Shape: Bent
Polar molecule

Quick Revision

✔ VSEPR predicts molecular shape
✔ Electron pairs arrange to minimize repulsion
✔ Lone pair repulsion is maximum
✔ Lone pairs reduce bond angles
✔ CH₄ → Tetrahedral
✔ NH₃ → Trigonal pyramidal
✔ H₂O → Bent
✔ BF₃ → Trigonal planar
✔ CO₂ → Linear
✔ SF₆ → Octahedral

Part 5: Valence Bond Theory, Orbital Overlap & Sigma (σ) and Pi (π) Bonds

28. Valence Bond Theory (VBT)

Definition

Valence Bond Theory explains the formation of covalent bonds by the overlapping of atomic orbitals containing electrons with opposite spins.

It was proposed by:

Heitler and London (1927)
Developed by Pauling and others

Main Ideas of Valence Bond Theory

1. Formation of Covalent Bond

A covalent bond forms when two half-filled atomic orbitals overlap.
The electrons in these orbitals have opposite spins.

Example:

H + H → H₂

Each hydrogen atom has one electron in 1s orbital.

When two 1s orbitals overlap, a H–H bond is formed.

2. Energy Change During Bond Formation

When two atoms approach each other:

Attractive forces increase.
Energy decreases.
At a particular distance, maximum stability is achieved.

This distance is called bond length.

3. Bond Strength

Bond strength depends on the extent of orbital overlap.

Greater overlap → Stronger bond

29. Orbital Overlap Concept

Definition

The partial interpenetration of atomic orbitals during bond formation is called orbital overlap.

The overlapping orbitals form a covalent bond.

Conditions for Effective Overlap

For strong bond formation:

Orbitals should have:
- Similar energy
- Proper orientation
- Maximum overlap
Electrons should have opposite spins.

Types of Orbital Overlap

Orbital overlap is mainly of two types:

Sigma (σ) overlap
Pi (π) overlap

30. Sigma (σ) Bond

Definition

A sigma bond is formed by head-on overlap of atomic orbitals along the internuclear axis.

It is the strongest covalent bond.

Types of Sigma Overlap

1. s-s Overlap

Overlap between two s orbitals.

Example:

H₂ molecule

H(1s) + H(1s) → σ bond

2. s-p Overlap

Overlap between s orbital of one atom and p orbital of another atom.

Example:

HCl molecule

3. p-p Overlap

Overlap between two p orbitals along the axis.

Example:

F₂ molecule

Characteristics of Sigma Bond

Formed by head-on overlap.
Stronger than pi bond.
Electron density lies between two nuclei.
Allows free rotation around bond axis.
Every single bond is a sigma bond.

31. Pi (π) Bond

Definition

A pi bond is formed by sidewise overlap of parallel p orbitals.

Formation of Pi Bond

Two p orbitals overlap sideways.
Electron clouds are present above and below the plane of atoms.

Characteristics of Pi Bond

Weaker than sigma bond.
Formed after sigma bond formation.
Restricts rotation.
Present in multiple bonds.

Difference Between Sigma and Pi Bond

Sigma (σ) Bond	Pi (π) Bond
Formed by head-on overlap	Formed by sidewise overlap
Stronger bond	Weaker bond
Electron density along axis	Electron density above and below axis
Allows rotation	Restricts rotation
Present in all single bonds	Present only in multiple bonds

32. Multiple Bonds

Double Bond

Contains:

1 Sigma + 1 Pi bond

Examples:

O₂
C₂H₄

Triple Bond

Contains:

1 Sigma + 2 Pi bonds

Examples:

N₂
C₂H₂

33. Directional Nature of Covalent Bonds

Covalent bonds are directional because:

Atomic orbitals overlap in specific directions.
The orientation of orbitals decides molecular shape.

Examples:

Molecule	Shape
CH₄	Tetrahedral
NH₃	Pyramidal
H₂O	Bent

34. Limitations of Simple Orbital Overlap

Simple overlap theory cannot explain:

Exact shapes of molecules.
Bond angles.
Equivalent bonds in molecules like CH₄.

Example:

Carbon has:

One s orbital
Three p orbitals

Simple overlap predicts different C–H bonds, but experimentally all four C–H bonds in methane are identical.

This problem is explained by hybridisation.

Quick Revision

✔ VBT explains covalent bond formation
✔ Covalent bond forms by orbital overlap
✔ Greater overlap = stronger bond
✔ σ bond = head-on overlap
✔ π bond = sidewise overlap
✔ σ bond is stronger than π bond
✔ Double bond = 1σ + 1π
✔ Triple bond = 1σ + 2π
✔ Orbital orientation decides molecular shape

Part 6: Hybridisation and Molecular Geometry

35. Hybridisation

Definition

Hybridisation is the process of mixing atomic orbitals of nearly equal energy to form a new set of equivalent orbitals called hybrid orbitals.

The concept was introduced by Linus Pauling.

Main Features of Hybridisation

Number of hybrid orbitals formed is equal to the number of atomic orbitals mixed.
Hybrid orbitals have:
- Same energy
- Same shape
Hybrid orbitals are more effective in bond formation.
Hybrid orbitals arrange themselves to reduce electron pair repulsion.
Hybridisation determines molecular geometry.

Conditions for Hybridisation

Only valence shell orbitals participate.
Orbitals should have similar energy.
Promotion of electrons is not always necessary.
Filled orbitals can also participate in some cases.

36. Types of Hybridisation

Main types:

sp hybridisation
sp² hybridisation
sp³ hybridisation
sp³d hybridisation
sp³d² hybridisation

37. sp Hybridisation

Definition

Mixing of:

One s orbital + One p orbital

forms:

Two sp hybrid orbitals

Characteristics

Percentage of s-character = 50%
Percentage of p-character = 50%
Geometry = Linear
Bond angle = 180°

Example: BeCl₂

Formation:

Be has one s and one p orbital hybridisation.
Two sp orbitals are formed.
Each overlaps with Cl orbital.

Shape:

Linear

Bond angle:

180°

Other Examples:

CO₂
C₂H₂

38. sp² Hybridisation

Definition

Mixing of:

One s orbital + Two p orbitals

forms:

Three sp² hybrid orbitals

Characteristics

Percentage of s-character = 33.3%
Percentage of p-character = 66.7%
Geometry = Trigonal planar
Bond angle = 120°

Example: BF₃

Formation:

Boron uses one s and two p orbitals.
Three sp² orbitals are formed.
Each overlaps with fluorine orbitals.

Shape:

Trigonal planar

Bond angle:

120°

Other Examples:

C₂H₄ (Ethene)
SO₃

39. sp³ Hybridisation

Definition

Mixing of:

One s orbital + Three p orbitals

forms:

Four sp³ hybrid orbitals

Characteristics

Percentage of s-character = 25%
Percentage of p-character = 75%
Geometry = Tetrahedral
Bond angle = 109.5°

Example 1: Methane (CH₄)

Carbon configuration:

Ground state:
1s² 2s² 2p²

Excited state:
1s² 2s¹ 2p³

One s + three p orbitals mix:

→ Four sp³ orbitals

Each overlaps with hydrogen 1s orbital.

Shape:

Tetrahedral

Bond angle:

109.5°

Example 2: Ammonia (NH₃)

Nitrogen has:

3 bond pairs
1 lone pair

Hybridisation:
sp³

Shape:
Trigonal pyramidal

Bond angle:
107°

Reason:

Lone pair repulsion decreases bond angle.

Example 3: Water (H₂O)

Oxygen has:

2 bond pairs
2 lone pairs

Hybridisation:
sp³

Shape:
Bent / Angular

Bond angle:
104.5°

Reason:

Two lone pairs create strong repulsion.

40. sp³d Hybridisation

Definition

Mixing of:

One s + Three p + One d orbital

forms:

Five sp³d hybrid orbitals

Geometry:

Trigonal bipyramidal

Bond angles:

120°
90°

Example: PCl₅

Phosphorus forms five hybrid orbitals.

Five chlorine atoms form five P–Cl bonds.

Shape:
Trigonal bipyramidal

41. sp³d² Hybridisation

Definition

Mixing of:

One s + Three p + Two d orbitals

forms:

Six sp³d² hybrid orbitals

Geometry:

Octahedral

Bond angle:

90°

Example: SF₆

Sulphur forms six hybrid orbitals.

Six fluorine atoms attach to sulphur.

Shape:
Octahedral

42. Summary Table of Hybridisation

Hybridisation	Orbitals Mixed	Number of Hybrid Orbitals	Shape	Bond Angle	Example
sp	s+p	2	Linear	180°	BeCl₂
sp²	s+p+p	3	Trigonal planar	120°	BF₃
sp³	s+p+p+p	4	Tetrahedral	109.5°	CH₄
sp³d	s+p+p+p+d	5	Trigonal bipyramidal	90°,120°	PCl₅
sp³d²	s+p+p+p+d+d	6	Octahedral	90°	SF₆

43. Hybridisation and Bond Formation in Hydrocarbons

Ethane (C₂H₆)

Carbon is sp³ hybridised.
C–C bond:
sp³–sp³ overlap
All bonds are sigma bonds.

Ethene (C₂H₄)

Carbon is sp² hybridised.
Contains:
- One sigma bond
- One pi bond

C=C bond = σ + π

Ethyne (C₂H₂)

Carbon is sp hybridised.
Contains:
- One sigma bond
- Two pi bonds

C≡C bond = σ + 2π

Quick Revision

✔ Hybridisation explains molecular shapes
✔ sp → Linear → 180°
✔ sp² → Trigonal planar → 120°
✔ sp³ → Tetrahedral → 109.5°
✔ sp³d → Trigonal bipyramidal
✔ sp³d² → Octahedral
✔ CH₄ → sp³
✔ BF₃ → sp²
✔ BeCl₂ → sp
✔ PCl₅ → sp³d
✔ SF₆ → sp³d²

Part 7: Molecular Orbital Theory (MOT), Bond Order & Magnetic Properties

44. Molecular Orbital Theory (MOT)

Definition

Molecular Orbital Theory explains the formation of molecules by the combination of atomic orbitals to form molecular orbitals.

It was developed by:

Hund
Mulliken
Lenard-Jones

Basic Principles of MOT

1. Formation of Molecular Orbitals

When atomic orbitals combine, they form new orbitals called molecular orbitals.

These orbitals belong to the entire molecule.

2. Types of Molecular Orbitals

Two types of molecular orbitals are formed:

(A) Bonding Molecular Orbital (BMO)

Formed by constructive combination of atomic orbitals.
Has lower energy than atomic orbitals.
Increases stability of molecule.

Represented as:

σ, π

(B) Antibonding Molecular Orbital (ABMO)

Formed by destructive combination of atomic orbitals.
Has higher energy than atomic orbitals.
Decreases stability.

Represented as:

σ*, π*

(* indicates antibonding orbital)

45. Conditions for Combination of Atomic Orbitals

Atomic orbitals combine effectively when:

1. Similar Energy

Orbitals having nearly equal energy combine easily.

Example:
2p–2p overlap

2. Proper Orientation

Orbitals must overlap in the correct direction.

3. Maximum Overlap

Greater overlap produces stronger molecular orbitals.

46. Types of Molecular Orbital Formation

1. Sigma (σ) Molecular Orbital

Formed by:

Head-on overlap

Examples:

s-s overlap
s-p overlap
p-p overlap

2. Pi (π) Molecular Orbital

Formed by:

Sidewise overlap of p orbitals

47. Electronic Configuration of Molecular Orbitals

Electrons fill molecular orbitals according to:

1. Aufbau Principle

Electrons occupy lower energy orbitals first.

2. Pauli Exclusion Principle

Maximum two electrons can occupy one orbital with opposite spins.

3. Hund’s Rule

Electrons occupy degenerate orbitals singly before pairing.

48. Energy Order of Molecular Orbitals

For molecules up to nitrogen (B₂, C₂, N₂):

Energy order:

σ2s < σ2s < π2pₓ = π2pᵧ < σ2p𝓏 < π2pₓ = π2pᵧ < σ2p𝓏

For oxygen and fluorine:

Energy order:

σ2s < σ2s < σ2p𝓏 < π2pₓ = π2pᵧ < π2pₓ = π2pᵧ < σ2p𝓏

49. Bond Order

Definition

Bond order is the number of bonds present between two atoms.

According to MOT: $\text{Bond Order}=\frac{N_b-N_a}{2}$ Bond Order=2Nb−Na

Where:

Nb = Number of electrons in bonding orbitals
Na = Number of electrons in antibonding orbitals

Importance of Bond Order

Higher bond order means:

✔ Stronger bond
✔ Shorter bond length
✔ Greater stability

Examples of Bond Order

H₂ Molecule

Electronic configuration:

(σ1s)²

Bond order:

= (2-0)/2

= 1

Therefore H₂ is stable.

He₂ Molecule

Electronic configuration:

(σ1s)²(σ*1s)²

Bond order:

= (2-2)/2

= 0

Therefore He₂ does not exist.

50. Magnetic Properties of Molecules

Molecules show two types of magnetic behaviour:

1. Diamagnetic Substances

Definition

Substances containing only paired electrons are called diamagnetic.

Properties:

Weakly repelled by magnetic field.

Examples:

H₂
N₂
CO

2. Paramagnetic Substances

Definition

Substances containing unpaired electrons are called paramagnetic.

Properties:

Attracted by magnetic field.

Examples:

O₂
NO
B₂

51. Oxygen Molecule (O₂) and Paramagnetism

According to MOT:

O₂ has two unpaired electrons in antibonding π* orbitals.

Therefore:

Oxygen is paramagnetic.
MOT successfully explains magnetic nature.

52. Bond Order and Stability Table

Molecule	Bond Order	Stability
H₂	1	Stable
He₂	0	Does not exist
N₂	3	Very stable
O₂	2	Stable
O₂⁺	2.5	More stable
O₂⁻	1.5	Less stable

53. Comparison of VBT and MOT

Valence Bond Theory	Molecular Orbital Theory
Bond formed by orbital overlap	Atomic orbitals combine to form molecular orbitals
Electrons belong to particular bonds	Electrons are spread over entire molecule
Cannot explain O₂ paramagnetism	Explains O₂ paramagnetism
Gives localized bonds	Gives delocalized orbitals

Quick Revision

✔ MOT explains molecular stability and magnetic properties
✔ Bonding orbital lowers energy
✔ Antibonding orbital increases energy
✔ Bond order = (Bonding electrons − Antibonding electrons)/2
✔ Higher bond order = stronger bond
✔ Unpaired electrons cause paramagnetism
✔ Paired electrons cause diamagnetism
✔ O₂ is paramagnetic due to two unpaired electrons
✔ He₂ is unstable because bond order is zero

Part 8: Hydrogen Bonding, Important Differences & Final Revision Sheet

54. Hydrogen Bonding

Definition

A hydrogen bond is an attractive force between hydrogen atom attached to a highly electronegative atom (F, O, N) and another electronegative atom having a lone pair of electrons.

Hydrogen bonding is weaker than covalent bonds but stronger than ordinary intermolecular forces.

Conditions Required for Hydrogen Bonding

Hydrogen must be directly attached to a highly electronegative atom:
- Fluorine (F)
- Oxygen (O)
- Nitrogen (N)
The electronegative atom should have a lone pair of electrons.

Types of Hydrogen Bonding

Hydrogen bonding is of two types:

Intermolecular hydrogen bonding
Intramolecular hydrogen bonding

55. Intermolecular Hydrogen Bonding

Definition

Hydrogen bonding occurring between different molecules is called intermolecular hydrogen bonding.

Examples:

1. Water (H₂O)

Water molecules form hydrogen bonds with each other.
This increases boiling point.

2. Hydrogen Fluoride (HF)

HF molecules are connected through hydrogen bonding.

3. Ammonia (NH₃)

Nitrogen forms hydrogen bonds due to lone pair.

Effects of Intermolecular Hydrogen Bonding

1. Higher Boiling Point

More energy is required to separate molecules.

Example:

H₂O has a higher boiling point than H₂S.

2. Increased Viscosity

Molecules stick together more strongly.

3. Greater Solubility

Compounds capable of hydrogen bonding dissolve better in water.

56. Intramolecular Hydrogen Bonding

Definition

Hydrogen bonding that occurs within the same molecule is called intramolecular hydrogen bonding.

Example:

o-nitrophenol

Hydrogen bonding occurs between:

–OH group
–NO₂ group

Effects of Intramolecular Hydrogen Bonding

Decreases intermolecular attraction.
Lowers boiling point.
Increases volatility.
Reduces water solubility.

57. Hydrogen Bonding in Water

Water shows strong hydrogen bonding because:

Oxygen is highly electronegative.
Oxygen has two lone pairs.
Each water molecule can form multiple hydrogen bonds.

This gives water:

High boiling point
High surface tension
Liquid state at room temperature

58. Important Differences

Ionic Bond vs Covalent Bond

Ionic Bond	Covalent Bond
Formed by transfer of electrons	Formed by sharing of electrons
Usually between metal and non-metal	Usually between non-metals
Strong electrostatic force	Directional bond
High melting point	Lower melting point generally
Conducts electricity in molten state	Usually poor conductor

Covalent Bond vs Coordinate Bond

Covalent Bond	Coordinate Bond
Both atoms contribute electrons	One atom contributes both electrons
Example: H₂	Example: NH₄⁺
Normal sharing	Donation of lone pair

Sigma Bond vs Pi Bond

Sigma (σ) Bond	Pi (π) Bond
Head-on overlap	Sidewise overlap
Stronger	Weaker
Allows rotation	Restricts rotation
First bond formed	Formed after sigma bond
Present in single bonds	Present in multiple bonds

Bond Order vs Bond Length

Higher Bond Order	Lower Bond Order
Stronger bond	Weaker bond
Shorter bond length	Longer bond length
Higher bond energy	Lower bond energy

Example:

C≡C > C=C > C–C

(Bond strength decreases)

59. Important Molecular Shapes Summary

Molecule	Hybridisation	Shape	Angle
BeCl₂	sp	Linear	180°
BF₃	sp²	Trigonal planar	120°
CH₄	sp³	Tetrahedral	109.5°
NH₃	sp³	Pyramidal	107°
H₂O	sp³	Bent	104.5°
PCl₅	sp³d	Trigonal bipyramidal	90°,120°
SF₆	sp³d²	Octahedral	90°

60. Important Formulas

Dipole Moment

$\mu = q \times r$ μ=q×r

Where:

μ = Dipole moment
q = Charge
r = Distance between charges

Formal Charge

$FC = V – L – \frac{B}{2}$ FC=V−L−2B

Where:

V = Valence electrons
L = Lone pair electrons
B = Bonding electrons

Bond Order (MOT)

$BO=\frac{N_b-N_a}{2}$ BO=2Nb−Na

Where:

Nb = Bonding electrons
Na = Antibonding electrons

61. Important Exam Points

Octet Rule Exceptions

Incomplete Octet:

BeCl₂
BF₃

Odd Electron Molecules:

NO
NO₂

Expanded Octet:

PF₅
SF₆

Important Examples to Remember

Linear:

CO₂, BeCl₂

Trigonal Planar:

BF₃

Tetrahedral:

CH₄

Pyramidal:

NH₃

Bent:

H₂O

Trigonal Bipyramidal:

PCl₅

Octahedral:

SF₆

Chapter Final One-Page Revision

✅ Chemical bond holds atoms together.
✅ Ionic bond forms by electron transfer.
✅ Covalent bond forms by electron sharing.
✅ Lewis structures show bonding and lone pairs.
✅ Octet rule explains stability of atoms.
✅ VSEPR theory predicts molecular shapes.
✅ Lone pair repulsion > bond pair repulsion.
✅ VBT explains bond formation by orbital overlap.
✅ σ bond is stronger than π bond.
✅ Hybridisation explains shapes of molecules.
✅ MOT explains stability and magnetic properties.
✅ O₂ is paramagnetic due to unpaired electrons.
✅ Hydrogen bonding affects boiling point and physical properties.