CBSE Class 11 Chemistry Revision Notes Chapter 4 Chemical Bonding and Molecular Structure

Chemical Bonding and Molecular Structure explains why atoms combine and how bonds decide the shape, stability and polarity of molecules. In CBSE Class 11 Chemistry Chapter 4, students learn ionic bonds, covalent bonds, Lewis structures, VSEPR theory, hybridisation, molecular orbital theory and hydrogen bonding.

Atoms usually gain stability by combining with other atoms. This combination happens through a chemical bond, which holds atoms, ions or molecules together. Once students understand how bonds form, molecular shape, bond strength, boiling point, polarity and reactivity become easier to connect.

These CBSE class 11 chemistry revision notes chapter 4 cover the complete flow of the chapter in a quick revision format. The chapter starts with the Kossel-Lewis approach and octet rule, then moves to ionic bond, covalent bond, bond parameters, VSEPR theory, valence bond theory, hybridisation, molecular orbital theory and intermolecular forces.

Key Takeaways

• Chemical bond: The attractive force that holds atoms or ions together.
• Octet rule: Atoms combine to achieve stable noble gas configuration.
• VSEPR theory: Electron pair repulsions decide molecular shape.
• Hybridisation: Atomic orbitals mix to form equivalent hybrid orbitals.
• Molecular orbital theory: Atomic orbitals combine to form molecular orbitals.

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Chemical Bonding and Molecular Structure Class 11 Chemistry Notes: Chapter Overview

Chemical Bonding and Molecular Structure deals with one basic question: why do atoms combine? Most atoms become more stable when they lose, gain or share electrons. The type of bond formed depends on electron transfer, electron sharing, electronegativity difference and the stability gained by the atoms.

Class 11 Chemistry Chapter 4 Notes also connect bonding with molecular structure. The same chapter explains why water is bent, carbon dioxide is linear, methane is tetrahedral and ammonia is pyramidal. This makes the chapter important for both school exams and entrance-level Chemistry basics.

What Is Chemical Bonding?

Chemical bonding is the formation of a chemical link between atoms, ions or molecules. This link holds the combining species together and forms a stable chemical compound.

A chemical bond forms because the bonded system has lower energy than the separate atoms. Lower energy means higher stability.

Why Do Atoms Combine?

Atoms combine to become more stable. Many atoms achieve this by completing their outer shell, similar to noble gases.

Types of Chemical Bonds

Chemical bonds can form through electron transfer, electron sharing or attraction between molecules.

Kossel-Lewis Approach to Chemical Bonding

Kossel and Lewis explained chemical bonding in terms of electrons. They connected bonding with noble gas stability.

According to this approach, atoms combine by losing, gaining or sharing valence electrons. This helps them achieve a stable outer electronic configuration.

Lewis Symbols

Lewis symbols show valence electrons as dots around the chemical symbol of an element.

For example, sodium has one valence electron, so its Lewis symbol has one dot. Oxygen has six valence electrons, so its Lewis symbol has six dots.

Lewis symbols help students count the electrons available for bonding.

Octet Rule

The octet rule states that atoms tend to combine in such a way that each atom has eight electrons in its valence shell.

Hydrogen follows the duplet rule because its first shell is complete with two electrons.

Examples

Limitations of the Octet Rule

The octet rule is useful for many simple molecules, but it has limitations.

Ionic Bond

An ionic bond forms when one atom transfers electron or electrons to another atom. The atom that loses electrons becomes a cation. The atom that gains electrons becomes an anion.

The ionic bond is the electrostatic attraction between the positive and negative ions.

Example: Formation of NaCl

Na → Na⁺ + e⁻
Cl + e⁻ → Cl⁻
Na⁺ + Cl⁻ → NaCl

Electrovalency

Electrovalency is the number of electrons lost or gained by an atom during ionic bond formation.

In simple terms, electrovalency is equal to the charge on the ion.

Factors Governing the Formation of Ionic Bonds

Ionic bond formation depends on the ease of cation formation, anion formation and crystal lattice stability.

Ionic compounds form easily when a metal has low ionisation enthalpy and a non-metal has high negative electron gain enthalpy.

Lattice Enthalpy

Lattice enthalpy is the energy required to separate one mole of a solid ionic compound into its gaseous ions.

For example, lattice enthalpy of NaCl refers to the energy needed to separate solid NaCl into Na⁺(g) and Cl⁻(g).

Lattice enthalpy explains why ionic compounds are stable in the solid state.

Characteristics of Ionic Compounds

Ionic compounds have strong electrostatic forces between cations and anions.

Covalent Bond

A covalent bond forms when two atoms share one or more electron pairs. Covalent bonds usually form between non-metal atoms.

In a covalent bond, each atom contributes at least one electron to the shared pair.

Polar and Non-polar Covalent Bonds

A covalent bond may be polar or non-polar depending on electronegativity difference.

In a polar bond, the more electronegative atom attracts the shared pair more strongly.

Coordinate Bond

A coordinate bond is a covalent bond in which the shared electron pair is donated by one atom, but shared by both atoms.

Example: In NH₄⁺, nitrogen donates a lone pair to H⁺.

Lewis Dot Structure

Lewis dot structure represents the arrangement of valence electrons in a molecule or ion.

Steps to Draw Lewis Dot Structures

The least electronegative atom usually occupies the central position. Hydrogen is always terminal.

Formal Charge

Formal charge is the charge assigned to an atom in a Lewis structure. It helps compare possible Lewis structures and choose the most stable one.

Formal Charge = Valence Electrons - Non-bonding Electrons - 1/2 Bonding Electrons

Formal charge is a bookkeeping tool. It does not always represent the actual charge on the atom.

Bond Parameters

Bond parameters describe the strength, size and arrangement of bonds in a molecule.

Bond Length

Bond length is the equilibrium distance between the nuclei of two bonded atoms.

As bond order increases, bond length decreases.

Bond Enthalpy

Bond enthalpy is the amount of energy required to break one mole of bonds in gaseous molecules.

For the same pair of atoms, a triple bond has higher bond enthalpy than a double bond, and a double bond has higher bond enthalpy than a single bond.

Bond Angle

Bond angle is the angle between the orbitals containing bonding electron pairs around the central atom.

Bond angle helps predict molecular shape.

Bond Order

Bond order is the number of bonds between two atoms.

Higher bond order usually means shorter and stronger bond.

VSEPR Theory

VSEPR theory stands for Valence Shell Electron Pair Repulsion theory. It predicts molecular shape using repulsions between electron pairs around the central atom.

The main idea is simple: electron pairs arrange themselves as far apart as possible to minimise repulsion.

This is why molecules with lone pairs, such as NH₃ and H₂O, have smaller bond angles than methane.

Shapes of Molecules Using VSEPR Theory

Valence Bond Theory

Valence bond theory explains covalent bond formation through overlap of half-filled atomic orbitals.

A bond forms when two half-filled orbitals overlap, and the overlapping region contains paired electrons with opposite spins.

Sigma Bond

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

Sigma bonds are stronger than pi bonds because the extent of overlap is higher.

Every single covalent bond is a sigma bond.

Pi Bond

A pi bond forms by sideways overlap of unhybridised p-orbitals. The overlap happens above and below the internuclear axis.

A pi bond is weaker than a sigma bond because the extent of overlap is lower.

Difference Between Sigma Bond and Pi Bond

Hybridisation

Hybridisation is the mixing of atomic orbitals of similar energy on the same atom to form new equivalent orbitals.

These new orbitals are called hybrid orbitals. Hybridisation explains molecular shapes and bond angles more clearly.

sp Hybridisation

sp hybridisation happens when one s orbital and one p orbital mix to form two sp hybrid orbitals.

sp² Hybridisation

sp² hybridisation happens when one s orbital and two p orbitals mix to form three sp² hybrid orbitals.

sp³ Hybridisation

sp³ hybridisation happens when one s orbital and three p orbitals mix to form four sp³ hybrid orbitals.

In NH₃ and H₂O, the central atom is also sp³ hybridised, but lone pairs change the observed shape.

Molecular Orbital Theory

Molecular orbital theory explains bonding through molecular orbitals formed by the combination of atomic orbitals.

When atomic orbitals combine, they form two types of molecular orbitals: bonding molecular orbital and antibonding molecular orbital.

Electrons fill molecular orbitals according to Aufbau principle, Pauli’s exclusion principle and Hund’s rule.

Conditions for Formation of Molecular Orbitals

The number of molecular orbitals formed is equal to the number of atomic orbitals combined.

Bond Order in Molecular Orbital Theory

In molecular orbital theory, bond order is calculated using the number of electrons in bonding and antibonding molecular orbitals.

Bond Order = 1/2 × (Number of electrons in bonding orbitals - Number of electrons in antibonding orbitals)

Magnetic Nature of Molecules

Molecular orbital theory helps explain magnetic behaviour.

O₂ is paramagnetic because it contains unpaired electrons in antibonding molecular orbitals.

Metallic Bond

Metallic bond is the attraction between metal kernels and mobile electrons.

In metals, valence electrons are delocalised and move freely. This is called the electron sea model.

Hydrogen Bonding

Hydrogen bonding is an attractive force between a hydrogen atom bonded to a highly electronegative atom and another electronegative atom nearby.

Hydrogen bonding is usually seen when hydrogen is bonded to fluorine, oxygen or nitrogen.

Types of Hydrogen Bonding

Hydrogen bonding affects boiling point, solubility, viscosity and molecular association.

Van Der Waals Forces

Van Der Waals forces are weak intermolecular forces between atoms or molecules.

They are weaker than ionic and covalent bonds but important in explaining physical properties.

Dipole Moment

Dipole moment measures the polarity of a bond or molecule. It is the product of charge and distance between the centres of positive and negative charges.

Dipole Moment = Charge × Distance

The SI unit is coulomb metre, but the common unit used in Chemistry is debye (D).

Molecular Polarity

Molecular polarity depends on bond polarity and molecular shape.

A molecule with polar bonds may have zero dipole moment if the bond dipoles cancel out due to symmetry.

Chemical Bonding and Molecular Structure Class 11 Chemistry Chapter 4 CBSE Notes: Quick Revision Tables

Important Concepts

Bond Parameters Summary

Shape and Hybridisation Summary

Key Terms from CBSE Class 11 Chemistry Revision Notes Chapter 4

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