CBSE Class 11 Chemistry Revision Notes Chapter 4

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

Chemistry Class 11 Chapter 4 Notes will help students gain conceptual understanding and improve their exam performance. The concepts taught in Chapter 4 of Class 11 Chemistry are thoroughly summarized in these revision notes, which were prepared in accordance with the updated recent CBSE Syllabus.

Class 11 Chemistry Revision Notes for Chapter 4 – Chemical Bonding and Molecular Structure – Free Download

1. Introduction:

Atoms are generally incapable of free existence while groups of atoms of the same or different elements exist as one compound.  For example H2O2, H2O. A molecule is a collection of atoms that exist as a single species and share common traits. 

2. Chemical Bond 

The process of producing a chemical compound by forming a chemical link between two or more molecules, atoms, or ions is known as chemical bonding. The chemical link between such resulting molecules is called a chemical bond.

Types of Chemical Bonds Include:

• Ionic Bonds
• Covalent Bonds
• Hydrogen Bonds
• Polar Bonds 

Modes of Chemical Composition

An electrovalent bond, also known as an ionic bond, is created when one or more electrons are transferred from one atom to another. This process is known as electrovalency. This can happen in the following two ways:

• A covalent bond is formed when the two joining atoms provide equal amounts of shared electrons.
• A coordinate or a dative bond is formed when the shared electrons are contributed by only one of the atoms. 

3. Lewis Symbols 

In 1916, Kossel and Lewis successfully explained why atoms combine to form molecules based on the idea that noble gases have an electrical configuration. The electronic configurations of noble gas atoms are thought to be stable because they have little to no tendency to mix with atoms of any other element.

To represent an atom’s valence shell electrons, Lewis used simple symbols. The electrons in the outer shell are depicted as dots around the atom symbol.

Electrovalency

An element’s electrovalency is determined by how many electrons it gains or loses as an electrovalent connection forms. Chlorine and oxygen, for instance, have electro-valencies of one and two, respectively, after gaining one and two electrons. To put it simply, valency is equivalent to the charge of an ion.

Factors Governing the Formation of Ionic Bonds

• Ionisation Enthalpy (Ionisation Energy): The amount of energy needed to remove one electron from the outermost shell of an isolated atom in a gaseous state and convert it to a gaseous positive ion is called ionisation enthalpy.
• Electron Gain Enthalpy (Electron Affinity): The enthalpy change occurs as an extra electron is added to an isolated atom in the gaseous phase and it forms a gaseous negative ion called electron affinity.
• Lattice Enthalpy (Lattice Energy): Ionic compounds are created when positively and negatively charged ions come together.

Characteristics of Ionic Compounds 

Some of the characteristics of ionic compounds are: 

• Crystalline in Nature: Ionic compounds form crystalline structures, but their size is determined by the size of the anions and cations. As the number of anions equals the number of cations that combine to form ionic compounds, the crystalline structure of ionic compounds is uncharged.
• Solubility: Polar solvents like water are capable of dissolving ionic compounds. But in non-polar covalent solvents, they are less soluble.
• Since the ions of ionic compounds are bound together by strong electrostatic forces, they are stiff and brittle. Being held together by strong forces, breaking them requires a lot of force. 
• Boiling and melting points of ionic compounds are high.
• Ionic compounds are difficult to separate because powerful electrostatic forces keep them bound. As a result, to overcome such powerful forces between them, high melting and boiling temperatures are needed.

Refer to Chemistry Class 11 Chapter 4 Notes for further conceptual clarity. 

5. Covalent Bond

Polar or non-polar properties can be found in covalent bonds. Due to the fact that the more electronegative atom pulls the electron pair closer to itself, electrons in polar covalent chemical bonds are distributed unevenly. A polar molecule like water is an exemplar.

Conditions for Writing the Lewis Dot Structures 

The octet rule is used for writing electron dot structures of covalent compounds. According to this rule, all the atoms in a molecule have eight electrons in their valence shell, except for the hydrogen atom. This rule doesn’t apply to a hydrogen atom as the first shell of hydrogen is completed by two electrons to achieve the helium configuration.

Elements in group 17 like chlorine must share one electron; elements in group 15 must share three electrons; elements in group 16 like oxygen and sulphur must share two electrons, and so on until a stable octet is formed.

6. Formal Charge

A formal charge refers to a charge on an individual atom(s) in a polyatomic molecule. It is a deceptive charge linked to a single atom in the compound structure. Due to the flaws in the configuration of an atom involved in the synthesis of a compound, a formal charge is caused.

Note: We refer to it as deceptive because the real charge of a molecule or compound is dispersed throughout the structure of the compound.

Formal Charge Formula

We will now discuss the formal charge formula since we are familiar with the concept. Mathematically, the formal charge formula is as follows:

Formal Charge = Valence Electrons – 0.5 Bonding Electrons – Non-bonding Electrons

Significance of Formal Charge

Calculating formal charges aids in the selection of the most stable structure among the numerous Lewis structures. The most stable structure is the one with the least formal charges on the atoms.

9. Hybridisation

Hybridisation is the mixing of atomic orbitals of the same atom that have slightly different energies. This process causes energy redistribution and the creation of new orbitals that have the same energies and shapes. Hybrid orbitals are the new orbitals that are created as a result of this process.

Type of Hybridisation

Hybridisation can take many different forms depending on the orbitals involved, such as sp, sp2, and sp3. 

• Diagonal or sp hybridisation – The chemical bonding in compounds like alkanes that have triple bonds is explained by sp hybridisation. The 2s orbital is mixed with only one of the three p orbitals in this model, resulting in two p orbitals and two sp orbitals. 
• Diagonals of sp2 hybridisation: Other carbon compounds and molecules are explained by sp2 hybridisation. For example, the carbons in ethane have a double bond.

Refer to Chemistry Class 11 Chapter 4 Notes for further explanation. 

10. Valence Bond Theory: 

Types of Covalent Bond       

1. Sigma (σ) Bond:

Head-on positive overlap (same phase) of atomic orbitals along the internuclear axis leads to sigma bonds. They are the strongest covalent bonds because the participating orbitals directly overlap. Sigma bonds are typically found in all single bonds. The following atomic orbital combinations can be used to make sigma bonds: 

• s-s Overlapping – In s-s overlapping, one “s” orbital from each participating atom undergoes head-on overlapping along the internuclear axis. One s orbital must be half-filled before it can overlap with another.
• sp Overlapping – Here, one half-filled s orbital undergoes head-on overlapping with one half-filled p orbital along the internuclear axis. Such overlapping can be observed in ammonia. 
• p-p overlapping – Here, along the internuclear axis, one half-filled p orbital from each participating atom undergoes head-on overlapping.

2. The Pi (π) Bond:

pi bonds are created by the perpendicular positive side-by-side overlap (same phase) of atomic orbitals with the internuclear axis. Atomic orbital axis is parallel to one another, but in this type of covalent bonding, the overlap is perpendicular to the internuclear axis.

pi bonds are weaker than sigma bonds because of the significantly lower degree of overlapping. The strength of a sigma bond is always greater than that of a pi bond, though. 

Difference Between sigma and pi Bonds

Bond Parameters  

• Bond Length: Bond length is the equilibrium distance between the centres of the nuclei of two bonded atoms.
• Bond Energy: Also called bond dissociation enthalpy, it is the amount of energy required to break one mole of bonds to separate them into gaseous atoms.
• Bond Angle: It is the angle between the orbitals containing the bonding electrons i.e., the lines representing the directions of the bonds. 
• Bond Order: The number of bonds between two atoms is known as bond order.

11. Molecular Orbital Theory

The connected atoms’ orbitals lose their distinctive characteristics as they get closer to one another and combine to form larger orbitals. These bigger orbitals formed are referred to as molecular orbitals.

Important Features of M.O.T 

Some of the important features of M.O.T. are as follows: 

• Molecular orbitals exist around the nuclei of molecules, just like atomic orbitals exist around an atom’s nucleus.
• The molecular orbitals are different from the atomic orbitals they are created from. 
• The molecular orbitals have varying energy levels.
• The form of the atomic orbitals from which molecular orbitals are formed determines the shape of the molecular orbitals.

Conditions for Atomic Orbitals to Form M.O

• The combining A.O. must be of comparable energy.
• The combining atomic orbitals should overlap to a large extent as the greater the overlap, the more stable molecule is formed.

12. Metallic Bond 

A metallic bond is a simultaneous attraction between kernels and the mobile electrons that hold the kernel together. This model is called the electron sea model.

13. Hydrogen Bonding

It is the process of hydrogen atoms interacting with one another to form hydrogen bonds, a particular class of attractive intermolecular forces. This type of hydrogen bond is joined to the hydrogen atom by one highly electronegative atom and another highly electronegative atom that is nearby.

Due to the dipole-dipole interaction between a hydrogen atom that is bound to a highly electronegative atom and another highly electronegative atom that is proximate to the hydrogen atom, a specific class of attractive intermolecular interactions is created. This precisely refers to hydrogen bonding.

Conditions for Hydrogen Bonding

• The molecule should contain a highly electronegative atom that is linked to the hydrogen atom. 
• The smaller the size of the electronegative atom, the greater the electrostatic attraction. Hence, the size of such an atom must be small.

Types of Hydrogen Bonding

1. Intermolecular Hydrogen Bonding – 

Intermolecular Hydrogen Bonding takes place between different molecules of the different or same compounds.

1. Intramolecular Hydrogen Bonding – It takes place within a molecule itself.

14. Van Der Waals Forces

Weak intermolecular forces that are dependent on the distance between atoms or molecules are known as Van Der Waals Forces. It arises from the interactions between uncharged molecules or atoms.

Types of Vander Waals Force

• Ion Dipole Attraction – It is between an ion such as Na+ and a polar molecule such as Hydrochloric acid (HCl).
• Dipole-Dipole Attraction – It is in between two polar molecules such as Hydrogen Fluoride and Hydrochloric acid (HCl).
• Ion Induced Dipole Attraction – A neutral molecule is induced by an ion in this case.
• Dipole-Induced dipole attraction: A neutral molecule is induced as a dipole by another dipole in this case. 
• Induced Dipole – Also called London dispersion force, it is induced between two nonpolar molecules.

15. Dipole Moment

A dipole moment expresses the polar character or polarity of the molecules. It is defined by the product of charge and the distance of separation of atoms in a chemical bond. Refer to Chapter 4 Chemistry Class 11 Notes for further explanation. 

Unit of Dipole Moment

Its unit in the CGS system is debye (D).

Applications of Dipole Moment

• For predicting the nature of the molecules – Molecules with zero dipole moment are non-polar. On the contrary, molecules that have a dipole moment are polar. 
• For determining the shape of the molecule – A molecule with a dipole moment will not be symmetrical and may be bent or angular, while a molecule with a zero dipole moment will be linear and symmetrical.

Solved example 

Q1. Which of these has the largest dipole moment?

1. CF4 
2. CH3F 

iii. CH3OH 

1. CH4 
2. CO2 

Ans. (i) and (iv)

The dipole moment of CH4 and CF4 is zero as they both have tetrahedral structures and are symmetrical. The dipole moment of CO2 is also zero as it is linear. Out of the remaining two i.e., CH3OH and CH3F, since F is more electronegative than O, CH3F will have a high dipole moment.