NCERT Class 11 Chemistry Chapter 4 Notes Chemical Bonding and Molecular Structure - Download PDF

The NCERT Class 11 Chemistry chapter 4 Notes talks about Chemical Bonding and Molecular Structure. The main topics covered in Class 11 Chemistry chapter 4 notes are KÖssel-Lewis’s approach to chemical bonding, the octet rule, and its limitations, drawing Lewis structures of simple molecules, VSEPR theory, valence bond theory, molecular orbital theory of homonuclear diatomic molecules, and concepts of hydrogen bonding. This chapter on Chemical Bonding and Molecular Structure Class 11 notes also includes some FAQs to clarify some of the most frequent and important questions. The Chemical Bonding and Molecular Structure Class 11 notes pdf download will help students to easily get the cbse class 11 chemistry ch 4 notes.

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• NCERT Notes Class 11 Chemistry
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• NCERT Exemplar Class 11 Chemistry Chapter 4 Solutions Chemical Bonding and Molecular Structure

NCERT Class 11 Chemistry Chapter 4 Notes

Chemical Bonding and Molecular Structure

Chemical Bond: Chemical bond is formed when various constituents are held together in chemical species by force of attraction.

KÖSSEL-LEWIS APPROACH TO CHEMICAL BONDING

• In the periodic table, noble gases separate the highly electronegative halogens from the highly electropositive alkali metals.

• The formation of a negative ion from a halogen atom is with the gain of an electron and a positive ion from an alkali metal atom is with the loss of an electron.

• The negative and positive ions formed attain stable noble gas electronic configurations. The noble gases (except helium,) have a stable outer shell configuration of eight electrons,.

• Electrostatic attraction stabilizes the negative and positive ions.

Electrovalent bond: The bond is formed, as a result of the electrostatic attraction between the positive and negative ions.

Ionic bond: The bond is formed when an atom loses an electron and another atom gains an electron.

Electronic theory of chemical bonding: It states that atoms can combine either by sharing valence electrons or by transfer of valence electrons from one atom to another, so as to have an octet in their valence shells. This is known as the octet rule.

Covalent bond: It is a bond that is formed when two atoms share a pair of electrons between them. Molecules can have a single covalent bond, double bond, or even a triple bond.

Lewis Representation of Simple Molecules (Lewis Structures)

In terms of shared pairs of electrons and the octet rule, the Lewis dot structures depict bonding in molecules and ions. Electrons are represented by the dots. Such structures are called Lewis dot structures.

For example, lewis dot structure of water molecule:

Formal Charge

Lewis dot structures do not represent the real shapes of molecules. In polyatomic ions, the net charge is held by the ion as a whole, rather than by a single atom. However, possible to give each atom a formal charge. The formal charge of an atom is the difference between the number of valence electrons of an atom in an isolated or free state and the number of electrons assigned to that atom in the Lewis structure.

Limitations of Octet Rule

1. Sometimes number of electrons around the central atom is less than eight, hence resulting in an incomplete octet of the central atom.

2. In certain molecules, such as NO there is an odd number of electrons which results in an incomplete octet of the atoms.

3. Sometimes the number of electrons around the central atom is more than eight, hence resulting in an expanded octet of the central atom.

Lattice Enthalpy: The Lattice Enthalpy is an ionic solid is defined as the formation of one mole or change of ionic compound in its gaseous ions which keep all the other things standard. We can say that to check the strength of ionic compound lattice enthalpy is used.

Bond Parameters

Bond Length: It is the distance between the nuclei of two bonded atoms of a molecule. Or we can say that it is the sum of two bonded atoms covalent radii.

Bond Angle: It is the angle made in between two covalent bonds which are having the same origin from a similar atom.

Bond Enthalpy: It is the amount of energy needed to separate all the covalent bonds of a specific type. This exists in between two atoms in a gaseous state.

Bond Order: It is the number of bonds a molecule has between two atoms.

Resonance Structures

Resonance is a concept, wherein, whenever a single lewis structure cannot describe the structure of the molecules, then various possible canonical structures are used to accurately describe the molecule.

Polarity of Bonds

In heteronuclear bonds, the shared electron pair shift towards the more electronegative atom, resulting in polarity between the atoms, and the molecule is said to be a polar molecule.

Due to polarity, the molecule develops Dipole moment which is defined as the product of charge and distance of separation between the two bonded atoms.

VALENCE SHELL ELECTRON PAIR REPULSION (VSEPR Theory)

This theory predicts the shape of the covalent molecules.

Below are the postulates of this theory:

1. The number of valence shell electron pairs decides the shape of a molecule around the central atom.

2. Pairs of electrons in the valence shell repel one another because of the negatively charged electron clouds.

3. The pairs of electrons occupy such positions in space where there is minimum repulsion so it maximise the distance between them.

4. The valence shell is assumed as a sphere with the electron pairs present on the spherical surface at a maximum distance from one another.

5. Multiple bonds are considered as a single electron pair and the two/three electron pairs of multiple bonds are taken as a single super pair.

6. When two or more resonance structures represent a molecule, the VSEPR theory is applied to any such structure.

The repulsive interaction of electron pairs is given as:

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

Valence Bond Theory

When two atoms approach each other, at a large distance there is no force, however, as the molecule starts approaching closer they experience a force of attraction, and after a distance when the two atoms come close enough they experience repulsive force.

So, when two atoms approach each other and form a bond, energy is released and this energy is called bond enthalpy.Types of Overlapping

Sigma(σ) bond: Also called head-on overlap or axial overlap, this type of covalent bond is formed when the bonding orbitals overlap end to end along the internuclear axis.

• s-s overlapping: The type of overlapping can be seen in half-filled s-orbitals of one atom and half-filled s-orbitals of another atom.

• s-p overlapping: The type of overlapping can be seen in half-filled s-orbitals of one atom and half-filled p-orbitals of another atom.

• p–p overlapping: The type of overlapping can be seen in p-orbitals of the two approaching atoms.

pi(π ) bond: The type of overlapping can be seen in a covalent bond that is formed when the bonding orbitals overlap sidewise, such that their axes are parallel to each other but perpendicular to the internuclear axis.

Hybridisation

The features of hybridization are:

1. The number of hybrid orbitals is equal in number to the atomic orbitals that hybridize.

2. The hybridized orbitals are always the same in energy and shape.

3. The hybrid orbitals form more stable bonds than pure atomic orbitals.

4. The hybrid orbitals are directed in minimum repulsion space giving a stable arrangement.

Types of Hybridisation

sp hybridization: This hybridization involves the mixing of one s and one p orbital to form equivalent sp hybrid orbitals.

hybridization: This hybridization involves mixing of ones and two p-orbitals to form three equivalent hybridized orbitals.

hybridization: This hybridisation involves mixing of one s-orbital and three p-orbitals of the valence shell to form four hybrid orbital.

hybridization: This hybridization involves mixing of one d-orbital, one s-orbital, and two p-orbitals of the valence shell to form four hybrid orbital.

hybridization: This hybridization involves mixing of two d-orbital, one s-orbital and three p-orbitals of the valence shell to form six hybrid orbital.

hybridisation: This hybridization involves mixing of one s-orbital, three p-orbitals, and one d-orbital to form five hybrid orbital.

hybridisation: This hybridisation involves mixing of one s-orbital, three p-orbitals, and two d-orbitals to form six hybrid orbitals.

MOLECULAR ORBITAL THEORY

1. As the electrons of atoms are present in the various atomic orbitals, the same way the electrons in a molecule are present in the various molecular orbitals.

2. The molecular orbitals are formed by the combination of atomic orbitals which is having comparable energies and proper symmetry.

3. A nucleus influences an electron in an atomic orbital, however, in a molecular orbital, it is influenced by two or more nuclei. Therefore, an atomic orbital is monocentric and a molecular orbital is polycentric.

4. The number of molecular orbitals is equal to the number of combining atomic orbitals. Two atomic orbitals combine, to form two molecular orbitals, where one is bonding molecular orbital and the other is antibonding molecular orbital.

5. The bonding molecular orbital has lower energy and the antibonding molecular orbital has more energy, hence is less stable.

6. In an atom, atomic orbital gives electron probability distribution around a nucleus, and in a molecule molecular orbital, gives the electron probability distribution around a group of nuclei.

7. The molecular orbitals are also filled in accordance with the followings rules and principles the Aufbau principle, Pauli’s exclusion principle, and Hund’s rule.

Molecular orbitals are formed by Linear Combination of Atomic Orbitals (LCAO) where atomic orbitals have the same energy or nearly the same, same symmetry and overlap is maximum.

Hydrogen Bonding

A hydrogen bond is formed between an H atom and an electronegative atom due to the force of attraction, where H has a partial positive charge and the electronegative atom has a partial negative charge.

There are two types of Hydrogen bonds:

Intermolecular hydrogen bond: A hydrogen bond is formed between an H atom and an electronegative atom of two different molecules. For instance, H-bond in HF molecule, alcohol or water molecules, etc.

Intramolecular hydrogen bond: A hydrogen bond is formed between an H atom and an electronegative atom of the same molecule. For instance, H-bond in o-nitrophenol.

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