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

Chemical bonding and molecular structure explain how atoms combine to form compounds and how the arrangement of atoms affects a substance's properties. Have you ever wondered why the shape of a molecule influences its behavior in chemical reactions? The way atoms bond together—whether through ionic, covalent, or metallic bonds—determines the strength, flexibility, and reactivity of the resulting compound. Chemical Bonding explains how different atoms combine to form molecules Apart from this Chapter 4 explains about properties of Chemical bonds, types of bonds, their nature, formation, and the molecular structure of different compounds.

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This Story also Contains

1. NCERT Class 11 Chemistry Chapter 4 Notes
2. KÖSSEL-LEWIS APPROACH TO CHEMICAL BONDING
3. Lewis Representation of Simple Molecules (Lewis Structures)
4. Bond Parameters
5. Valence Bond Theory
6. Hybridization
7. MOLECULAR ORBITAL THEORY
8. Hydrogen Bonding
9. NCERT Class 11 Notes Chapter-Wise
10. Subject-Wise NCERT Exemplar Solutions
11. Subject-Wise NCERT Solutions

This chapter explains several basic concepts of Chemistry that help students to develop a strong base on the subject which will help them not just in Board exams but also in competitive exams like JEE, and NEET. Some of the main topics covered in Class 11 Chemistry chapter 4 notes are KÖssel-Lewis’s approach to chemical bonding, the octet rule, and Limitations Of The Octet Rule, drawing Lewis Electron Dot Structures, Vsepr Theory, Valence Bond Theory, Molecular Orbital Theory of homonuclear diatomic molecules, and concepts of Hydrogen Bonding.

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NCERT Class 11 Chemistry Chapter 4 Notes

Chemical Bonding and Molecular Structure

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

Type of intermolecular forces

1. Ion - dipole

2. Ion-induced dipole or debye forces

3. Dispersion forces or London forces

4. Dipole - dipole force or keesom force

5. Hydrogen bonding

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,$n{s}^2$) have a stable outer shell configuration of eight electrons,$n{s}^{2}n{p}^6$.

• 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, a 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.

The number of electrons present in the outermost shell are known as valence electrons. For example, the electronic configuration of sodium (Na) is 2, 8, 1, thus, sodium has one valence electron. According to the long form of the periodic table, in the case of representative elements, the group number is equal to the number of valence electrons. The valence electrons in atoms are shown in terms of Lewis symbols. To write Lewis symbol for an element, we write down its symbol surrounded by a number of dots or crosses equal to the number of valence electrons. Paired and unpaired valence electrons are also indicated.

For example, lewis dot structure of the 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.

Thus, we calculate the formal charge as follows:

formal charge = valence shell electrons − lone pair electrons − 1/2 bonding electrons

Limitations of the 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 shifts 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.

Dipole moment(µ):

IT is vector sum of all the individual bond moments

Mathematically $\mu =q\times d$

$q=$ magnitude of charge

$d=$ distance between the centres of opposite charge

It is denoted by a small arrow with a tail on the positive centre and head pointing towards negative center

Fazan’s Rule and Covalent Character in Ionic Bond :

The covalent character in ionic bond is determined by Fazan’s rule. It simply says that no ionic bond is completely ionic, there is always some covalent character in ionic bond. When a cation approaches an anion, then the electron cloud of the anion is distorted and shifted towards the cation, this distortion is known as the polarisation of the anion.

The ability of the cation to distort the anion is known as polarising power and the ability of the anion to get distorted is known as polarisability.

The covalent character in ionic bond depends on the following factors:

• Size of the cation: Smaller the size of the cation, larger will be its polarisability.
• Size of the anion: Larger the size of the anion, larger will be its polarisability.
• Charge on cation and anion: More the charge on cation more will be its polarising power. Further, more the charge on anion, more will be its polarisability.

Thus covalent character for chlorides follows this order:

NaCl < MgCl₂ < AlCl₃

In this case, the charge on the cation increases, thus its polarising power also increases.

Further, for cation size, the covalent character follows the below order:

LiCl > NaCl > KCl > CsCl

In this case, as the size of the cation increases, its polarising power decreases.

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.

Hybridization

The features of hybridization are:

1. The number of hybrid orbitals is equal to the number of 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 a 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.

$s{p}^2$ hybridization: This hybridization involves the mixing of one and two p-orbitals to form three equivalent $s{p}^2$ hybridized orbitals.

$s{p}^3$ hybridization: This hybridization involves the mixing of one s-orbital and three p-orbitals of the valence shell to form four $s{p}^3$ hybrid orbitals.

$ds{p}^2$hybridization: This hybridization involves the mixing of one d-orbital, one s-orbital, and two p-orbitals of the valence shell to form four $ds{p}^2$hybrid orbitals.

$d^{2}s{p}^3$ hybridization: This hybridization involves the mixing of two d-orbitals, one s-orbital, and three p-orbitals of the valence shell to form six $d^{2}s{p}^3$ hybrid orbitals.

$s{p}^{3}d$hybridisation: This hybridization involves the mixing of one s-orbital, three p-orbitals, and one d-orbital to form five $s{p}^{3}d$hybrid orbitals.

$s{p}^3{d}^2$hybridisation: This hybridization involves the mixing of one s-orbital, three p-orbitals, and two d-orbitals to form six $s{p}^3{d}^2$hybrid orbitals.

MOLECULAR ORBITAL THEORY

1. As the electrons of atoms are present in the various atomic orbitals, in 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 that have 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 a bonding molecular orbital and the other is an antibonding molecular orbital.

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

6. In an atom, an atomic orbital gives the 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 following 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

Hydrogen bonds are strong forces which occurs when a hydrogen atom bonded to an electronegative atom approaches a nearby electronegative atom such as O, N, F, etc.. Greater the electronegativity of the atom will result in an increase in hydrogen-bond strength. The hydrogen bond is stronger intermolecular force, but it is weaker than a covalent or an ionic bond. Hydrogen bonds are responsible for holding together DNA, proteins, and other macromolecules.

Formation of Hydrogen Bond

A hydrogen bond is an electromagnetic attraction that occurs between a partially positively charged hydrogen atom attached to a highly electronegative atom and another nearby electronegative atom. A hydrogen bond is a type of dipole-dipole interaction; it is not a true chemical bond. This hydrogen bond attraction can occur between the different molecules (intermolecularly) or within different parts of a single molecule (intramolecularly).

Types of Hydrogen Bonding

There are two types of hydrogen bonding, i.e:

• Intermolecular Hydrogen Bonding: Intermolecular hydrogen bonding occurs when the H-atom of one molecule and an electronegative atom of another molecule are close to each other. For example, hydrogen bond between the molecules of hydrogen fluoride. Intermolecular hydrogen bonding results into association of molecules. Thus, it increases the melting point, boiling point, solubility, etc.

• Intramolecular Hydrogen Bonding: Intramolecular hydrogen bonding occurs the hydrogen atom and an electronegative atom of the same molecule. Intramolecular hydrogen bonding results in the cyclization of the molecules and prevents their association. Thus, the properties of these compounds like melting point, boiling point, etc. are usually low. For example, intramolecular hydrogen bonding is present in molecules such as o-nitrophenol, o-nitrobenzoic acid, etc.

Significance of NCERT Class 11 Chemistry Chapter 4 Notes

Chapter 4 chemistry class 11 notes are helpful to revise the chapter and to get an idea about the main topics covered in the chapter. Also, these NCERT Class 11 Chemistry chapter 4 notes are useful for competitive exams like VITEEE, BITSAT, JEE Main, NEET, etc. The chemistry class 11 chapter 4 notes PDF download can be very helpful in reading through the paper.

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