Theories of Covalent Bonding and Shape of Molecules

Bond Characteristics

Define the term ‘bond energy’

Bond energy is defined as the average amount of energy required to break all bonds of a particular type in one mole of the substance. It is a measure of the strength of the bond.

Key Information:

The SI unit for bond energy is kJ/mol (kilojoules per mole)
Some older literature may use kcal/mol (1 kcal/mol ≈ 4.184 kJ/mol)
Bond energy is always reported as a positive value (bond breaking is endothermic)

Note: Bond energy values are typically measured at standard conditions (298 K and 1 atm pressure).

Relate Bond Energy with Bond Strength

Bond energy is a direct measure of bond strength—the higher the bond energy, the stronger the bond, and the more energy is required to break it. Stronger bonds are typically shorter and more stable, while weaker bonds are longer and more easily broken.

Key Relationships:

Bond Energy & Bond Strength:
– Higher bond energy = Stronger bond (more stable, harder to break).
– Lower bond energy = Weaker bond (less stable, easier to break).
Bond Energy & Bond Length:
– Stronger bonds (higher bond energy) tend to be shorter (e.g., triple bonds are shorter and stronger than double or single bonds).
– Weaker bonds (lower bond energy) tend to be longer (e.g., single bonds are longer and weaker than double bonds).
Dipole Moment & Bond Energy:
– A stronger dipole (greater electronegativity difference between bonded atoms) generally leads to higher bond energy in polar covalent bonds.

Examples:

1. Covalent Bonds (Homonuclear Diatomics)

Bond	Bond Energy (kJ/mol)	Bond Strength	Bond Length (pm)
N≡N (Triple)	945	Very Strong	110
O=O (Double)	498	Moderate	121
F-F (Single)	158	Weak	142

N≡N has the highest bond energy, making it one of the strongest bonds in chemistry (hard to break, very stable).
F-F has the lowest bond energy, making it relatively weak (easily broken, reactive).

2. Polar Covalent Bonds (Dipole Effect)

Bond	Bond Energy (kJ/mol)	Electronegativity Difference (ΔEN)
H-F	567	1.78 (Highly Polar)
H-Cl	431	0.96 (Moderately Polar)
H-Br	366	0.76 (Less Polar)

H-F has the highest bond energy due to its strong dipole (large ΔEN).
H-Br has lower bond energy because it is less polar.

3. Ionic vs. Covalent Bonds

NaCl (Ionic Bond) – Very high bond energy (~787 kJ/mol) due to strong electrostatic attraction.
C-C (Covalent Bond) – Moderate bond energy (~346 kJ/mol), weaker than ionic but still strong.

Conclusion:

Strongest Bonds: Triple bonds (N≡N), ionic bonds (NaCl), and highly polar covalent bonds (H-F).
Weakest Bonds: Single bonds (F-F) and nonpolar bonds (C-H).

This relationship is crucial in predicting reactivity, stability, and physical properties of molecules.

Define the term “bond Length

The distance between the nuclei of two atoms forming a covalent bond is called bond length.

Key Points:

Bond length is typically measured in picometers (pm) or angstroms (Å)
1 Å = 100 pm = 10^-10 m
Bond length decreases with increasing bond order (single > double > triple)
Shorter bonds are generally stronger than longer bonds

Explain the ionic character of a covalent bond

A covalent bond involves the sharing of electrons between two atoms. However, if there is a significant difference in the electronegativities of the two atoms, the sharing of electrons is unequal. The more electronegative atom attracts the shared electron pair more strongly, resulting in a partial negative charge (δ−) on that atom and a partial positive charge (δ+) on the other atom. This unequal sharing of electrons introduces ionic character to the covalent bond.

Examples of Bonds with Ionic Character:

Bond	Electronegativity Difference	% Ionic Character
H-F	1.78	~43%
H-Cl	0.96	~20%
C-O	0.89	~18%

Key Factors Affecting Ionic Character:

Electronegativity difference between bonded atoms
Bond polarity (dipole moment)
Molecular geometry and symmetry

Predict the nature of bonding on the basis of electronegativity

The difference in electronegativity between two bonded atoms can be used to predict the nature of the bond.

Ionic Bond

Electronegativity difference > 1.8
Electron transfer occurs
Example: NaCl (ΔEN = 2.23)

Polar Covalent

ΔEN between 0.4 – 1.8
Unequal electron sharing
Example: HCl (ΔEN = 0.96)

Nonpolar Covalent

ΔEN < 0.4
Equal electron sharing
Example: Cl₂ (ΔEN = 0)

Bond Length in Heteronuclear Molecules

In heteronuclear molecules (molecules with different types of atoms), a greater difference in electronegativity between the bonded atoms can lead to a shorter bond length. This is because the atom with higher electronegativity pulls the shared electron pair more strongly towards itself, effectively drawing the two nuclei closer together due to increased electrostatic attraction.

Key Points:

Greater EN difference → stronger electrostatic attraction → shorter bond length
This effect is most noticeable in polar covalent bonds
Bond length also depends on atomic size and bond order

Bond Length Comparison

H-F

ΔEN = 1.78

92 pm

H-Cl

ΔEN = 0.96

127 pm

H-Br

ΔEN = 0.76

141 pm

Note: As the electronegativity difference decreases, the bond length increases

Exemplify Dipole Moment

Dipole moment is a measure of the polarity of a molecule. It arises when there is a separation of positive and negative charges within the molecule due to differences in electronegativity. It is a vector quantity, having both magnitude and direction.

The magnitude of the dipole moment (μ) is given by:

μ = q × r

where:
q = magnitude of charge (C)
r = distance between charges (m)

Units of Dipole Moment

Debye (D) – Commonly used in chemistry
Coulomb-meter (Cm) – SI unit
Conversion: 1 D = 3.33564 × 10^-30 Cm

Example: HCl Molecule

Charge Separation

Chlorine is more electronegative than hydrogen, creating:

Positive end (δ⁺) on hydrogen
Negative end (δ⁻) on chlorine

Calculation

Given:

q = 1.602 × 10^-19 C
r = 1.275 × 10^-10 m

Step 1: Calculate in Cm

μ = (1.602 × 10-19 C) × (1.275 × 10-10 m)

= 2.04255 × 10-29 Cm

Step 2: Convert to Debye

μ = (2.04255 × 10-29 Cm) / (3.33564 × 10-30 Cm/D)

= 6.12 D

HCl Dipole Moment Representation

The arrow points toward the more electronegative atom

Predict Molecular Polarity from Molecular Shapes

While bond polarity depends on the electronegativity difference between atoms, molecular polarity depends on both the bond polarities and the overall shape or geometry of the molecule.

Nonpolar Molecules

Have polar bonds
Symmetrical molecular shape
Bond dipoles cancel out
No net dipole moment

Examples:

CCl₄

Tetrahedral

CO₂

Linear

Polar Molecules

Have polar bonds
Asymmetrical molecular shape
Bond dipoles do not cancel
Net dipole moment exists

Examples:

H₂O

Bent

NH₃

Trigonal pyramidal

Determining Molecular Polarity

Flowchart for determining molecular polarity

Dipole Moment in Molecules

CO₂ (Nonpolar)

Dipoles cancel → Nonpolar

H₂O (Polar)

Dipoles don’t cancel → Polar

Key Takeaways:

Molecular polarity depends on both bond polarity and molecular geometry
Symmetrical shapes often (but not always) lead to nonpolar molecules
Lone pairs frequently create asymmetry and polarity
The presence of a net dipole moment indicates polarity

Shape of Molecules using VSEPR Theory

Explain Valence Shell Electron Pair Repulsion (VSEPR) Theory

Valence Shell Electron Pair Repulsion (VSEPR) theory is a model used to predict and explain the shapes of molecules. The main postulate of this theory is that electron pairs in the valence shell of a central atom, whether bonding pairs or lone pairs, repel each other and tend to remain at maximum distance apart to minimize this repulsion. This arrangement of electron pairs determines the electron pair geometry, and the arrangement of only the bonding pairs determines the molecular geometry or shape.

Electron Pair Repulsion Strength

Strongest

Lone pair – Lone pair

Medium

Lone pair – Bonding pair

Weakest

Bonding pair – Bonding pair

A diagram showing lone pairs and bonding pairs

Draw Molecular Shapes Using VSEPR Theory

Step-by-Step Process:

Count valence electrons of the central atom
Add electrons contributed by the surrounding atoms
Adjust for ions (add for anions, subtract for cations)
Calculate electron pairs (total valence electrons ÷ 2)
Determine bonding pairs (number of atoms) and lone pairs (total pairs – bonding pairs)
Arrange electron pairs to minimize repulsion (linear, trigonal planar, tetrahedral, etc.)
Determine molecular shape based on bonding pair arrangement

Molecular Shape Examples

SO₃ (Sulfur Trioxide)

6 bonding pairs (3 double bond units), 0 lone pairs
Trigonal planar geometry
Bond angles: 120°

PCl₃ (Phosphorus Trichloride)

4 bonding pairs, 1 lone pair
Trigonal pyramidal
Bond angles: 100°

H₂O (Water)

2 bonding pairs, 2 lone pairs
Bent/V-shaped
Bond angles: 104.5°

BeCl₂

2 bonding pairs, 0 lone pairs
Linear geometry
Bond angles: 180°

AlCl₃

3 bonding pairs, 0 lone pairs
Trigonal planar
Bond angles: 120°

SO₂

4 bonding pairs, 1 lone pair
Bent/V-shaped
Bond angles: 119°

VSEPR Geometry Summary

Electron Domains	Bonding Pairs	Lone Pairs	Geometry	Example
2	2	0	Linear	BeCl₂
3	3	0	Trigonal planar	BF₃
4	4	0	Tetrahedral	CH₄
4	3	1	Trigonal pyramidal	NH₃
4	2	2	Bent/V-shaped	H₂O

VBT, MOT and Hybridization

Explain Valence Bond Theory (VBT)

Valence Bond Theory (VBT) describes a covalent bond as the overlap of half-filled atomic orbitals in the valence shells of two atoms. When these orbitals overlap, a pair of electrons, one from each atom, occupies the overlapped region, resulting in bond formation. The greater the overlap between the atomic orbitals, the stronger the bond formed. VBT considers the molecule as a combination of individual atoms retaining their atomic orbitals, although they may overlap. VBT is concerned with both the formation of bonds and the shapes of molecules.

Key Features of VBT:

Focuses on atomic orbital overlap between bonding atoms
Electron pairs are localized between atoms
Accounts for both bond formation and molecular shapes
Atoms retain their atomic orbitals which may overlap during bonding

Sigma (σ) and Pi (π) Bonds

Sigma (σ) Bonds

Formed by head-on (axial) overlap of orbitals
Types: s-s, s-p, or p-p (end-to-end)
Electron density concentrated along the internuclear axis
All single bonds are σ bonds

Pi (π) Bonds

Formed by sideways (parallel) overlap of p orbitals
Electron density above and below the internuclear axis
Weaker than σ bonds due to less effective overlap
Multiple bonds: Double = 1σ+1π, Triple = 1σ+2π

Bond Type Examples

H₂ (H-H)

Single bond

1 σ bond

O₂ (O=O)

Double bond

1 σ + 1 π bond

N₂ (N≡N)

Triple bond

1 σ + 2 π bonds

Sigma vs. Pi Bonds Comparison

Property	Sigma (σ) Bond	Pi (π) Bond
Formation	Head-on orbital overlap	Sideways p-orbital overlap
Electron Density	Along the internuclear axis	Above/below internuclear axis
Strength	Stronger	Weaker
Rotation	Free rotation possible	Restricts rotation
Occurrence	All single bonds	Additional bonds in multiples

Key Takeaways:

VBT explains bonding through orbital overlap
Sigma bonds are stronger and allow free rotation
Pi bonds are weaker and restrict rotation
Multiple bonds combine both σ and π bonds

VBT, MOT, and Hybridization

3.3.1 Valence Bond Theory (VBT)

VBT describes covalent bonds as the overlap of half-filled atomic orbitals from two atoms. When orbitals overlap, electron pairs occupy the shared region, forming bonds. The bond strength increases with greater orbital overlap.

Sigma (σ) and Pi (π) Bonds

Sigma (σ) Bonds

Head-on (axial) orbital overlap
Types: s-s, s-p, p-p (end-to-end)
Electron density along the internuclear axis
All single bonds are σ bonds

Pi (π) Bonds

Sideways (parallel) p-orbital overlap
Electron density above/below the axis
Weaker than σ bonds
Double: 1σ+1π, Triple: 1σ+2π

Sigma and pi bond formation

Hybridization and Its Types

Hybridization is the mixing of atomic orbitals to form new hybrid orbitals with equivalent energy and shape. These hybrid orbitals are more effective in forming stable σ bonds.

sp Hybridization

Mixing of one s + one p orbital
Forms two sp hybrid orbitals
Linear geometry (180° bond angle)
Example: BeCl₂, C₂H₂ (Acetylene)

sp² Hybridization

Mixing of one s + two p orbitals
Forms three sp² hybrid orbitals
Trigonal planar geometry (120° bond angle)
Example: BF₃, C₂H₄ (Ethylene)

sp³ Hybridization

Mixing of one s + three p orbitals
Forms four sp³ hybrid orbitals
Tetrahedral geometry (109.5° bond angle)
Examples:
- CH₄ (4 bonding pairs)
- NH₃ (3 bonding + 1 lone pair → trigonal pyramidal)
- H₂O (2 bonding + 2 lone pairs → bent)

Tetrahedral

trigonal pyramidal

bent

CH₄

4 bonding pairs

NH₃

3 bonding + 1 lone pair

H₂O

2 bonding + 2 lone pairs

sp³d (trigonal bipyramidal) Hybridization

Mixing of one s, three p and one d orbitals
Forms five sp³d hybrid orbitals
Trigonal bipyramidal geometry with bond angles of 120° between the three equatorial bonds and 90° between the axial and equatorial bonds.
d Orbital: In sp³d hybridization, the dz² d-orbital is typically involved. This is because its symmetry, which involves the z-axis, aligns well with the overlap required for bonding with the s and p orbitals in the hybridization process.
Examples:
- Phosphorus pentachloride (PCl₅) → trigonal bipyramidal.
- SF₄

Trigonal bipyramidal	Seesaw	T-Shaped	Linear

PCl₅ 5 bonding pairs	SF₄ 4 bonding + 1 lone pair	ClF₃ 3 bonding + 2 lone pairs	I₃^– 2 bonding + 3 lone pairs

sp³d² (octahedral) Hybridization

Mixing of one s, three p, and two d orbitals
Forms six sp³d² hybrid orbitals
Octahedral geometry with bond angles of 90° between atoms in an octahedral structure. These six sp³d² hybrid orbitals are arranged in a way that they point towards the six corners of an octahedron, which is a regular polyhedron with eight triangular faces.
d Orbitals: The d orbitals involved in sp³d² hybridization are typically the dx²-y² and dz² orbitals
Examples:
- SF₆ (sulfur hexafluoride)
- XeF₄ (xenon tetrafluoride)

Octahedral

Square pyramidal

Square planar

SF₆

6 bonding pairs

ClF₅

5 bonding + 1 lone pair

XeF₄

4 bonding + 2 lone pair

Molecular Shapes from Hybridization

Hybridization	Geometry	Bond Angle	Examples
sp	Linear	180°	BeCl₂, CO₂
sp²	Trigonal planar	120°	BF₃, C₂H₄
sp³	Tetrahedral	109.5°	CH₄
sp³ (3bp + 1lp)	Trigonal pyramidal	~107°	NH₃
sp³ (2bp + 2lp)	Bent/V-shaped	~104.5°	H₂O

Key Takeaways:

Hybridization explains molecular geometry and bond angles
More p orbitals in hybridization → more complex geometries
Lone pairs affect the final molecular shape (VSEPR)
Sigma bonds form first, and pi bonds add to multiple bonds

Molecular Orbital Theory (MOT)

Molecular Orbital Theory (MOT) describes how atomic orbitals combine to form molecular orbitals that belong to the entire molecule. These molecular orbitals are formed by the linear combination of atomic orbitals (LCAO).

Key Concepts of MOT:

Molecular Orbitals

Formed by combining atomic orbitals
Number of MOs = Number of AOs combined
Extend over the entire molecule

Orbital Types

Bonding (BMO): Lower energy, stabilizing
Antibonding (ABMO): Higher energy, destabilizing
Non-bonding: Neutral energy

Electron Filling

Electrons are filled into molecular orbitals according to rules similar to those for filling atomic orbitals:

Aufbau principle
Pauli exclusion principle
Hund’s rule

Formation of bonding (BMO) and antibonding molecular orbitals (ABMO) from atomic orbitals of Hydrogen atoms

MOT Applications for Homonuclear Diatomics

Using MOT, we can determine the electronic configuration of homonuclear diatomic molecules by filling electrons into the molecular orbitals in order of increasing energy.

**Orbital mixing in O₂ vs. N₂:** Due to the larger 2s-2p energy difference in O₂, s-p mixing is negligible, while in N₂, the closer energy levels permit significant interaction.

Molecular Orbital Energy Diagrams and s-p mixing

In the molecular orbital diagram of O₂, we assumed that the 2s orbitals interact only with other 2s orbitals. In contrast, the 2p orbitals interact only with other 2p orbitals, due to the significant energy difference between the 2s and 2p levels. However, this assumption does not hold for less electronegative atoms like N₂. In such elements, the energy gap between the 2s and 2p orbitals is smaller, allowing them to satisfy the energy condition required for orbital overlap.

MO diagram for O₂, F₂

The typical energy level diagram for second-period homonuclear diatomic molecules involves the following order of increasing energy for the molecular orbitals formed from 2s and 2p atomic orbitals:

σ2s < σ*2s < σ2p < π2p = π2p < π*2p = π*2p < σ*2p

MO diagram for O₂

For N₂, B₂, and C₂

For diatomic molecules (like N $_{2}$ $B_{2}$ , and $C_{2}$ ), the order of $σ_{2 p z}$ and $π_{2 p x}, π_{2 p y}$ is reversed: In these elements, the energies of the 2s and 2p orbitals are closer in energy and meet the energy criterion for overlap.

σ2s < σ*2s < π2p = π2p < σ2p < π*2p = π*2p < σ*2p

MO diagram for N₂

Key Calculations

Bond Order

Bond order is defined as half the difference between the number of electrons in bonding molecular orbitals ( $N_{b}$ ) and the number of electrons in antibonding molecular orbitals ( $N_{a}$ ):

BO = (N_b – N_a) / 2

Where:
N_b = bonding electrons
N_a = antibonding electrons

Magnetic Properties

Paramagnetic: Unpaired electrons
Diamagnetic: All electrons paired

Example Molecules

Molecule	Configuration	Bond Order	Magnetism	Stability
He₂	(σ_1s)²(σ*_1s)²	0	Diamagnetic	Unstable
N₂	(σ_2s)²(σ*_2s)²(π_2p)⁴(σ_2p)²	3	Diamagnetic	Very stable
O₂	(σ_2s)²(σ_2s)²(σ_2p)²(π_2p)⁴(π_2p)²	2	Paramagnetic	Stable
O₂²⁺	(σ_2s)²(σ*_2s)²(σ_2p)²(π_2p)⁴	3	Diamagnetic	Very stable
O₂^2-	(σ_2s)²(σ_2s)²(σ_2p)²(π_2p)⁴(π_2p)⁴	1	Diamagnetic	Less stable

Key Takeaways:

MOT explains bonding through delocalized molecular orbitals
Bond order predicts bond strength and length
Unpaired electrons cause paramagnetism
Different energy level orders exist for light vs heavy diatomic molecules

Compare VBT and MOT

Valence Bond Theory (VBT) and Molecular Orbital Theory (MOT) are two fundamental approaches to understanding chemical bonding, each with distinct perspectives and applications.

Feature	Valence Bond Theory (VBT)	Molecular Orbital Theory (MOT)
Molecular View	Molecule as collection of atoms with localized bonds	Molecule as nuclei surrounded by delocalized molecular orbitals
Bond Formation	Overlap of atomic orbitals (retains atomic character)	Combination of atomic orbitals forming new molecular orbitals
Electron Description	Electrons localized between bonding atoms	Electrons delocalized over entire molecule
Magnetic Properties	Cannot explain paramagnetism of O₂	Explains O₂ paramagnetism via unpaired electrons
Excited States	Limited ability to describe excited states	Effectively describes excited states and spectroscopy
Resonance	Requires resonance structures for delocalization	Naturally accounts for electron delocalization
Bond Order	Qualitative (based on bond count)	Quantitative: BO = (N_b – N_a)/2 (N_b = bonding e^–, N_a = antibonding e^–)

When to Use VBT

Simple molecules with localized bonds
Predicting molecular geometry
Understanding hybridization
Organic chemistry applications

When to Use MOT

Molecules with delocalized electrons
Predicting magnetic properties
Understanding excited states
Diatomic and polyatomic molecules

Key Takeaways:

VBT emphasizes localized bonds while MOT emphasizes delocalized orbitals
MOT better explains magnetic properties and excited states
VBT is more intuitive for simple molecules
MOT provides quantitative bond order calculations

Effect of Bonding on Physical and Chemical Properties

Solubility of Ionic and Covalent Compounds

“Like dissolves like” principle:

Substances with similar bonding types and polarity tend to be mutually soluble.

Ionic Compounds

Soluble in polar solvents (e.g., water)
Dissolution involves ion-dipole interactions
Water molecules hydrate ions, overcoming lattice energy

Hydration of NaCl in water

Polar Covalent Compounds

Soluble in polar solvents
Dissolution involves dipole-dipole interactions
May form hydrogen bonds with water

Alcohol dissolving in water

Nonpolar Covalent Compounds

Soluble in nonpolar solvents
Dissolution involves London dispersion forces
Generally insoluble in polar solvents

Oil is insoluble in water

Chemical Properties of Ionic and Covalent Compounds

Ionic Compounds

Conduct electricity when molten or dissolved
Fast ionic reactions (e.g., precipitation)
High melting/boiling points due to strong lattice energy
Brittle crystal structure

Covalent Compounds

Non-conductive (unless ionized in solution)
Slower molecular reactions (e.g., combustion)
Lower melting/boiling points due to weaker intermolecular forces
Variable physical states at room temperature

Examples:

NaCl (ionic):

MP: 801°C
Conducts when molten

CH₄ (covalent):

MP: -182°C
Non-conductive

Directional vs. Non-directional Bonds

Covalent Bonds (Directional)

Result from orbital overlap in specific directions
Create definite molecular shapes (VSEPR)
Bond strength depends on overlap orientation
Example: CH₄ tetrahedral structure

Tetrahedral structure (CH₄)

Ionic Bonds (Non-directional)

Result from electrostatic attraction in all directions
Form crystal lattices based on ion packing
Strength depends on charge and distance
Example: NaCl cubic crystal structure

Key Takeaways:

Bond type determines solubility through “like dissolves like”
Ionic compounds have distinct properties from covalent
Covalent bonds are directional, ionic bonds are non-directional
Physical properties reflect bond strength and interactions

Chemistry

Bond Characteristics

Define the term ‘bond energy’

Key Information:

Relate Bond Energy with Bond Strength

Key Relationships:

Examples:

1. Covalent Bonds (Homonuclear Diatomics)

2. Polar Covalent Bonds (Dipole Effect)

3. Ionic vs. Covalent Bonds

Conclusion:

Define the term “bond Length

Key Points:

Explain the ionic character of a covalent bond

Examples of Bonds with Ionic Character:

Key Factors Affecting Ionic Character:

Predict the nature of bonding on the basis of electronegativity

Ionic Bond

Polar Covalent

Nonpolar Covalent

Bond Length in Heteronuclear Molecules

Key Points:

Bond Length Comparison

Exemplify Dipole Moment

Units of Dipole Moment

Example: HCl Molecule

Charge Separation

Calculation

Step 1: Calculate in Cm

Step 2: Convert to Debye

HCl Dipole Moment Representation

Predict Molecular Polarity from Molecular Shapes

Nonpolar Molecules

Examples:

Polar Molecules

Examples:

Determining Molecular Polarity

Dipole Moment in Molecules

CO2 (Nonpolar)

H2O (Polar)

Key Takeaways:

Shape of Molecules using VSEPR Theory

Explain Valence Shell Electron Pair Repulsion (VSEPR) Theory

Electron Pair Repulsion Strength

Draw Molecular Shapes Using VSEPR Theory

Step-by-Step Process:

Molecular Shape Examples

SO3 (Sulfur Trioxide)

PCl3 (Phosphorus Trichloride)

H2O (Water)

BeCl2

AlCl3

SO2

VSEPR Geometry Summary

VBT, MOT and Hybridization

Explain Valence Bond Theory (VBT)

Key Features of VBT:

Sigma (σ) and Pi (π) Bonds

Sigma (σ) Bonds

Pi (π) Bonds

Bond Type Examples

H2 (H-H)

O2 (O=O)

N2 (N≡N)

Sigma vs. Pi Bonds Comparison

Key Takeaways:

VBT, MOT, and Hybridization

3.3.1 Valence Bond Theory (VBT)

Sigma (σ) and Pi (π) Bonds

Sigma (σ) Bonds

Pi (π) Bonds

Hybridization and Its Types

sp Hybridization

sp² Hybridization

sp³ Hybridization

sp³d (trigonal bipyramidal) Hybridization

sp³d² (octahedral) Hybridization

Molecular Shapes from Hybridization

Key Takeaways:

Molecular Orbital Theory (MOT)

CO₂ (Nonpolar)

H₂O (Polar)

SO₃ (Sulfur Trioxide)

PCl₃ (Phosphorus Trichloride)

H₂O (Water)

BeCl₂

AlCl₃

SO₂

H₂ (H-H)

O₂ (O=O)

N₂ (N≡N)

MO diagram for O₂, F₂

For N₂, B₂, and C₂