Chemical Bonding and Molecular Structure: Theories, Chemical Bonding, and Bond Parameters

We have compiled all important concepts of Chemical Bonding class 11 Chemistry in the form concise revision PDF. These NCERT-based short notes are perfect for quick last-minute preparation before the Exam. It covers key concepts, chemical bonding theories, and bond parameters useful for CBSE exams, JEE Main, and NEET exams. You can download the Chemical Bonding and Molecular Structure PDF for free by clicking on the tab below.

📄 Download PDF

The Class 11 Chemistry chapter 4 covers fundamentally important topics. These topics are essential to understanding the bond formation and reaction mechanism to form new molecules. All the important topics and subtopics are listed below.

The process of forming of chemical bond to form a molecule by combining two or more atoms, ions, or molecules is called chemical bonding. Chemical bond is a force that keeps two or more atoms intact and forms a molecule of a chemical compound. In general, it is considered that atoms form chemical bonds to achieve a stable electronic configuration like their nearest noble gas. All important theories explaining the formation of chemical bonds in class 11 chemistry are discussed in this article.

For example, O₂ is a molecule that is formed by combining 2 oxygen atoms, and chemical bonds hold these two Oxygen atoms together. 

Applications of Chemical Bonding

There are many key theories developed to explain the chemical bonding and related aspects. First were Kössel and Lewis. In 1916  they proposed a logical explanation of chemical bonding through fulfilling valence electrons, which was based on the inertness of noble gases. The theory postulated that atoms achieve a stable configuration like noble gases when they are linked by chemical bonds. Afterwards, Valence bond theory and others provide a more detailed picture of the chemical bonding process. Below are four key theories discussed in this chapter.

• Kossel-Lewis Approach to Chemical Bonding
• Valence Shell Electron Pair Repulsion (VSEPR) Theory
• Valence Bond Theory (VBT)
• Molecular Orbital Theory (MOT)

All these concepts are discussed briefly on this page. You can also access complete information about these topics by visiting the related article, which is available below. 

The Kossel-Lewis approach is a combined term used for the theories proposed by Kossel and Lewis. Kossel's Theory explains the Ionic Bonding due to transfer of electrons between atoms. The Lewis Theory covers the formation of covalent bonding due to the sharing of electrons. Both, when combined called the Kossel-Lewis approach. You can read them in detail below

Lewis Theory of Chemical  Bonding

Gilbert Lewis explained understanding chemical bonds through sharing the valence electrons. This type of bond is known as a covalent bond. Important key concepts related to Lewis theory:

• Octet Rule: This rule states that all elements want to achieve stability by having eight electrons in their outermost shell to resemble inert gases. 

• Lewis Symbols: He used dots to represent the electrons in the outer shell.

• Lewis Dot Structure: These electron dot structures help in visualizing the sharing or transfer of electrons to form bonds. The lone pairs are represented as two dots on either side of the symbol of the element. other shared or transferred to form bond are represented through a line (dash).

• Lewis Theory Example: Both hydrogen atoms share their one electron with each other ot form a covalent bond between hydrogen atoms to form H₂.

  $$H\cdot +\cdot H\to H:H({\text{H}}_2)$$

Kossel Theory of Chemical  Bonding

Kossel theory is primarily focused formation of ionic bonds through the transfer of electrons especially between metals and non-metals. Important concepts related to Kossel theory are given below:

• Electrovalency: It is the number of electrons gained or lost to form the ionic bond. Atoms achieve a stable octet by gaining or losing electrons, forming ions..

• Electrostatic Attraction: There is cation (+ve) and anion (-ve) formation due to electron transfer. So, oppositely charged ions experience electrostatic attraction force, forming an ionic bond.

• Noble Gas Configuration: The atoms convert themselves by gaining or releasing electrons to achieve the electronic configuration of the nearest noble gas.

• Kossel Theory Example: The Sodium (Na), being a metal ato,m loses elctron to form cation (+ve) and Clorine gains that electron to complete its octet and form anion (-ve). Both ion form and ionic or electrovalent bond to form NaCl.

$$\text{Na}\to {\text{Na}}^{+}+e^{-};\text{Cl}+e^{-}\to {\text{Cl}}^{-}$$

$${\text{Na}}^{+}+{\text{Cl}}^{-}\to \text{NaCl}$$

Factors Affecting the Formation of Covalent and Ionic Bonds (as per Kossel-Lewis)

Here are several key factors which affect the formation of covalent and ionic Bonds. read below;

• Low Ionization Energy: Mostly Metals lose electrons to form cations. This happens because of the low ionization energy of Metal atoms.
• High Electron Affinity: Non-Metals, on the other hand, generally gain electrons to form an anion. They can do so due to their high electron affinity.
• Electronegativity Difference: A large electronegativity difference between the two atoms is the key for easy transfer of electrons and the formation of an ionic bond.

• Electronegativity Difference: There must be a small electronegativity difference between the two participating atoms for sharing the electrons to form a covalent bond.

• High Ionization Energy: When both atoms have high ionization energies, it allows atoms to just share the electrons, not transfer to form a covalent bond.  

• Mutual Attraction: The two atoms should be mutually attracted so that they can come close to shared electrons and overlap orbitals to form a covalent bond.

Read in detail about the Kossel-Lewis theory of chemical bonding by clicking on the link below:

Kossel-Lewis Approach to Chemical Bonding

Valence Shell Electron Pair Repulsion Theory, commonly known as VSEPR theory. This theory is primarily focused on predicting the shape (geometry) of molecules. The VSEPR Theory predicts the shape based on the repulsion between electron pairs around the central atom, so that the repulsion force can be minimized.  There are two types of Electron Pairs: Bonding pairs, which are shared between atoms, and  Lone pairs that are not involved in bonding. There are important key postulates in VSEPR Theory. 

• Electron Pair Repulsion: The bonding and lone electron pairs around the central atom repel each other and arrange themselves in a way so that the repulsion is minimized.
• Geometry Determination: The arrangement of electron pairs determines the geometry of the molecule. The shape of the molecule will be such that the repulsions between these pairs are minimum. 
• Repulsion Strength: The strength of repulsion follows the order:(Lone pair–lone pair > lone pair–bond pair > bond pair–bond pair).

Molecular Geometries based on Electron Pairs

Based on the arrangements of the electron pairs, the molecular structure types are given below:

The Valence Bond Theory (VBT) focuses on providing an explanation of the formation of covalent bonds through the overlapping atomic orbitals of two atoms. These atomic orbitals overlap and form covalent bond due to sharing the unpaired electrons. Based on overlapping orientation, there are two types of covalent bonds formed: σ (sigma), and π (pi). You can read the brief key properties below:

• A covalent bond forms due to the overlap of two atomic orbitals overlap, and their unpaired electrons are paired (Shared).
• The strength of the Covalent bond will depend on the extent of the overlapping
• In the overlapping process, the electron density increases in the region between nuclei.

Sigma (σ) and Pi (π) Bonds

Valence Bond Theory focuses on two types of bond formation based on eh overlapping of the orbitals. 

• Sigma (σ) Bond: When two orbitals are in head-on overlap along the internuclear axis, they form a Sigma bond.  These bonds are stronger and allow free rotation around the bond axis.

• Pi (π) Bond: When two orbitals (especially p orbitals) overlap sideways above and below the internuclear axis, they form a Pi bond. Pi bonds are weaker than sigma bonds, and there is no rotation around the bond axis in pi bonds..

Sigma Bond vs Pi Bond:

There are several types of overlap possible between two atoms. The process of two or more orbitals combining through overlapping is explained as Hybridization. Hybridization explains that atomic orbitals overlap and combine to form hybrid orbitals with specific shapes.

Valence Bond Theory (VBT)