Chemical Bonding and Molecular Structure

Kossel – Lewis Approach to Chemical Bonding 

Chemical bond is an attraction force between atoms of a molecule. 

In 1916 Kossel and Lewis succeeded in explaining the chemical bonding in terms of electrons. 

Octet Rule – Atoms of different elements try to attain electronic configuration like noble gas atoms or to complete their octet by chemical bonding. In other words, atoms of all main group elements tend to bond in such a way that each atom has 8 electrons in its valence shell so that the atoms will attain electronic configuration like noble gases. Thus, by chemical bonding atoms get stability like noble gases. 

Lewis postulated that atoms achieve the stable octet when they are linked by chemical bonds. 

Lewis Symbols – In the formation of a molecule, only the valence electrons (electrons of outermost shell of an atom) take part in chemical bonding. An American chemist G. N. Lewis introduced simple notations to represent valence electrons in an atom. These notations are known as Lewis symbols. For example, carbon has 4 electrons in its outermost shell. So, its Lewis symbol will be as follows –

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Electrovalent Bond – The bond formed, as a result of electrostatic attraction between the positive and negative ion is known as electrovalent bond. It is also called ionic bond. 

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The number of electrons lost or gained during the formation of an electrovalent linkage is termed as the electrovalency of the element.

Factors Which Affect The Formation of Ionic Bond – 

  • Ionization Enthalpy - Ionisation enthaly of any element is the amount of energy required to remove an electron from the outermost shell of an isolated atom in gaseous phase so as to convert it into a gaseous positive ion. It is also known as ionization energy. Lesser the ionisation enthalpy, easier will be the removal of an electron, that is formation of a positive ion and hence greater the chances of formation of an ionic bond. 

  • Electron Gain Enthalpy - Electron affinity or Electron gain enthalpy of an element is the enthalpy change that takes place when an extra electron is added to an isolated atom in the gaseous phase to form a gaseous negative ion. 

Higher is the electron affinity, more is the energy released and more stable will be the negative ion produced. Consequently, the probability of formation of ionic bond will increase. 

  • Lattice Energy - The energy released when the requisite number of gaseous positive and negative ions combine to form one mole of the ionic compound is called lattice enthalpy. 

Characteristics of Ionic Compounds –

  • Ionic compounds usually exist in the solid state. 

  • In ionic compounds ions get arranged in regular patterns. So, they have a definite crystal structure. For example, NaCl has octahedral crystal structure. 

  • Ionic compounds usually possess high melting and boiling points. This is because ions are held together by strong electrostatic forces in ionic compounds. 

  • Ionic compounds are usually soluble in polar solvents such as water.

  • Solutions of ionic compounds are good conductors of electricity. They are good conductors of electricity in their molten state as well. 

  • Ionic compounds take part in various chemical reactions. 

Covalent Bond – The bond formed by the sharing of electrons between atoms, is called covalent bond. This term and the idea was introduced by Langmuir in 1919.

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The dots in the above structure represent the valence electrons (Lewis symbols). Such structures are called Lewis dot structure. Two atoms may have single, double or triple covalent bond. 

Formal Charge – Formal charge of an atom in a polyatomic molecule can be defined as the difference between the number of valence electrons of that atom in an isolated or free state and the number of electrons assigned to that atom in the Lewis structure. It helps to select the most stable structure, that is, the one with least energy out of the different possible Lewis structures.

Formal charge on an atom in a Lewis structure = [total number of valence electrons in the free atom] – [total number of nonbonding electrons] - 1/2 [total number of shared electrons]

Example – Formal charge on atoms in carbonate ion. 

Lewis structure of CO32- ion –

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Formal charge on C atom = 4 – 0 - ½ (8) = 0

Formal charge on double bonded O atom = 6 – 4 - ½(4) = 0

Formal charge on single bonded O atom = 6 – 6 -  ½(2) = -1

Bond Parameters 

Bond Length – Bond length is defined as the equilibrium distance between the nuclei of two bonded atoms in a molecule. 

Factors Affecting Bond Length –

  1. Size of Atoms - The bond length increases with increase in the size of the atoms. For example, bond lengths of H–X are in the order -

HI > HBr > HCl > HF 

  1. Multiplicity of Bond - The bond length decreases with the multiplicity of the bond. Thus, bond length of carbon-carbon bonds is in the order -

C  ≡ C < C = C < C – C 

  1. Type of Hybridisation - As an s - orbital is smaller in size, greater the s-character, shorter is the hybrid orbital and hence shorter is the bond length. For example, 

Bond lengths - sp3 C–H > sp2 C–H > sp C–H 

s-character - (25%) (33%) (50%)

Bond Angle – Bond angle is defined as the angle between the orbitals containing bonding electron pairs around the central atom in a molecule or complex ion. It can be determined experimentally by spectroscopic methods. It is expressed ion degree. 

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Bond Enthalpy – Bond enthalpy is defined as the amount of energy required to break one mole of bonds of a particular type between two atoms in a gaseous state. 

Factors Affecting Bond Energy –

  1. Size of the Atoms - Greater the size of the atoms, greater is the bond length and less is the bond dissociation enthalpy, i.e., less is the bond strength. 

  2. Multiplicity of Bonds - For the bond between the same two atoms, greater is the multiplicity of the bond, greater is the bond dissociation enthalpy. This is firstly because atoms come closer and secondly, the number of bonds to be broken is more. For example, bond dissociation enthalpies of H2, O2 and N2 are in the order - H–H < O = O < N

  3. Number of Lone Pairs of Electrons Present - Greater the number of lone pairs of electrons present on the bonded atoms greater is the repulsion between the atoms and hence less is the bond dissociation enthalpy. 

Bond Order – In the Lewis description of covalent bond, the bond order is given by the number of bonds between the two atoms in a molecule. 

Isoelectronic molecules and ions have identical bond orders. Bond order of F2 and O22- is 1 and they are isoelectronic. 

With increase in bond order, bond enthalpy increases and bond length decreases. Thus, we can write -

Bond order ∝ bond enthalpy ∝ \[\frac{1}{Bond \; length}\]

Dipole Moment – Dipole moment can be defined as the product of the magnitude of the charge and the distance between the centres of the positive and negative is denoted by the Greek letter 'μ' . Mathematically, it can be expressed as follows –

Dipole moment = charge x distance of separation 

μ = Q × d

It is expressed in Debye units (D). 1D = 3.33564 10-30 C m (C = coulomb and m = meter)

It is a vector quantity. It is depicted by a small arrow with a tail on positive center and head pointing towards the negative center. For example, dipole moment of HCl is represented as follows –

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Significance of dipole moment - The molecules having zero moment are non-polar molecules and those having μnet ≠ 0  are polar in nature. 

The value of dipole moment can be used for determining the amount of ionic character in a bond. The formula for determining percentage of ionic character is given below –

Percentage of ionic character = \[\frac{Experimental\; value\; of \;dipole moment}{Theoretical \;value \;of \;dipole \;moment }\] ×100

The Valence Shell Electron Pair Repulsion (VSEPR) Theory

Valence shell electron pair repulsion theory explains the shapes of molecules. It is based on the repulsive interactions between electron pairs in the valence shell of the atoms. This theory was given by Sidgwick and Powell in 1940 and was further improved by Nyholm and Gillespie in 1957. Main postulates of VSEPR theory are as follows –

  • The shape of the molecule depends upon the number of valence shell electron pairs either bonded or non-bonded around the central atom. 

  • Pairs of electrons in the valence shell repel one another since their electron clouds are negatively charged. 

  • These pairs of electrons tend to occupy such positions in space that minimize repulsion and thus maximize distance between them. 

  • The valence shell is taken as a sphere with the electron pairs localizing on the spherical surface at maximum distance from one another. 

  • A multiple bond is treated as if it is a single electron pair and the two or three electron pairs of a multiple bond are treated as a single super pair. 

  • Where two or more resonance structures can represent a molecule, the VSEPR model is applicable to any such structure. 

Decreasing Order of The Repulsive Interaction of Electron Pairs –

Lone pair – lone pair > lone pair – bond pair > Bond pair – Bond pair


lp – lp > lp – bp > bp – bp 

Valence Bond Theory 

Valence bond theory was given by Heitler and London in 1927 and further developed by Pauling and others. It is based on electronic configuration of elements, the overlap criteria of atomic orbitals, the hybridization of atomic orbitals and the principle of variation and superposition. In class XI chemistry, valence bond theory is discussed at a basic level in qualitative and non – mathematical terms only. 

  • Attractive forces arise between – 

  • nucleus of one atom and its own electron. 

  • nucleus of one atom and electron of another atom.

  • Repulsive forces arise between –

  • Electrons of the two atoms 

  • Nuclei of the two atoms

If the magnitude of attraction force is more than repulsive force, then the formation of bond takes place and potential energy decreases. Ultimately, net force of attraction balances the force of repulsion and molecules acquire minimum energy and become stable.  

Sigma Bond – The strongest covalent bond which is formed by the head on overlapping of atomic orbitals is called sigma bond. It is denoted by . We find the sigma bond in alkanes, alkenes, alkynes. It is formed by s-s overlapping, s-p overlapping and p-p overlapping. Formation of sigma bond is given below between the orbitals- 

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Pi Bond – The covalent bond which is formed by lateral overlapping of the half-filled atomic orbitals of atoms is called pi bond. It is denoted by . We find pi bonds in alkenes and alkynes. Formation of pi bond is given below between the two orbitals – 

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The concept of hybridization was introduced by Pauling. Atomic orbitals combine to form a new set of equivalent orbitals known as hybrid orbitals. The phenomenon is known as hybridization. These hybrid orbitals take part in bond formation. 

The Main Features of Hybridization Are As Follows –

  • Number of hybrid orbitals = Number of atomic orbitals that take part in hybridization 

  • The hybridized orbitals are always equivalent in energy and shape. 

  • The hybrid orbitals are more effective in forming stable bonds than the pure atomic orbitals. 

  • These hybrid orbitals are directed in space in some preferred direction to have minimum repulsion between electron pairs. Thus, forms a stable arrangement. 

  • The type of hybridization indicates the geometry of the molecules. 

Conditions for Hybridization -

  • The orbitals present in the valence shell of the atom are hybridized. 

  • The orbitals undergoing hybridization should have almost equal energy. 

  • Promotion of electrons is not necessary before hybridization. 

  • It is not necessary that only half - filled orbitals can take part in hybridization. Even filled orbitals of the valence shell take part in hybridization. 

Types of Hybridization – Hybridization takes place by involvement of various orbitals such as s, p, and d orbitals. It has following types -

sp Hybridization – one s and one p orbitals take part in hybridization and each sp hybrid orbital has 50% s - character and 50% p – character. Two sp- hybrid orbitals are formed. 

sp2 Hybridization – one s and two p orbitals participate in this hybridization and form three hybrid orbitals. Each of these hybrid orbitals show 33% s character and 67% p character. 

sp3 Hybridization – one s and three p orbitals participate in this hybridization and form four hybrid orbitals. Each of these hybrid orbitals show 25% s character and 75% p character. 

dsp2 Hybridization – one - s, two p and one d orbitals participate in this hybridization and form four hybrid orbitals. 

sp3d Hybridization – one s, three p and one d orbitals participate in this hybridization and form five hybrid orbitals. 

sp3d2 Hybridization – one - s, three p and two d orbitals participate in this hybridization and form three hybrid orbitals. 

d2sp3 Hybridization – one - s, three p and two d orbitals participate in this hybridization and form three hybrid orbitals. 

S. No. 

Atomic Orbitals used 

Hybrid Orbitals Formed 

Bond Angle 





Two sp orbitals 





s, p, p

Three sp2 orbitals 


Planar trigonal 



s, p, p, p

Four sp3 orbitals 





d, s, p, p

Four dsp2 orbitals 


Square planar 



s, p, p, p, d

Five sp3d orbitals 

90° and 120°

Trigonal bipyramidal 



s, p, p, p, d, d

Six sp3d2 orbitals 

90° and 120°

Square pyramidal 



d, s, p, p, p

Five dsp3 orbitals 


T – shaped 


Molecular Orbital Theory 

Molecular orbital theory is another approach to explain chemical bonding in molecules. It was given by Mulliken and Hund in 1932. The molecular orbital theory considers the entire molecule as a unit with all the electrons moving under the influence of all the nuclei present in the molecule. Salient features of molecular orbital theory are as follows –

  1. Like an Atomic orbital which is around the nucleus of an atom there are molecular orbital which are around the nuclei of a molecule. 

  2. The molecular orbitals are entirely different from the atomic orbitals from which they are formed. Atomic orbitals fuse together and form molecular orbitals. Conditions for atomic orbitals to form molecular orbitals –

  • The combining atomic orbitals should be of comparable energy.

  • The combining atomic orbitals must overlap to a large extent. Greater the overlap, stable the molecule form. 

  1. The valence electrons of the constituent atoms are considered to be moving under the influence of nuclei of participating atoms in the molecular orbital.

  2. The molecular orbitals possess different energy levels like atomic orbitals in an isolated atom. 

  3. The shape of molecular orbitals is dependent upon the shapes of atomic orbitals from which they are formed.

  4. Molecular orbitals are arranged in order of increasing energy just like atomic orbitals. 

  5. The number of molecular orbitals formed is equal to the number of atomic orbitals combining in bond formation. 

  6. Like atomic orbitals, the filling of electrons in molecular orbitals is governed by the three principles - Aufbau principle, Hund’s rule and Pauli’s exclusion principle.

Hydrogen Bonding 

Hydrogen bonding can be defined as the attraction force which binds the hydrogen atom of one molecule with the electronegative atom of another molecule. It is also called hydrogen bridge. It is a very weak bond. 

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Conditions For Hydrogen Bonding – The molecule must contain a highly electronegative atom such as F, Cl, Br etc. and the size of the electronegative atom should be small. 

Types of Hydrogen Bonding – It Is of Two Types –

  • Intermolecular Hydrogen Bonding - When hydrogen bonding takes place between different molecules of the same or different compounds, it is called intermolecular hydrogen bonding. Example – water alcohol. 

  • Intramolecular Hydrogen Bonding - The hydrogen bonding which takes place within a molecule itself. The bond is formed between the Hatom of one group with the more electronegative atom of the other group. Example – Hydrogen bonding in an o-nitrophenol molecule. 

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This ends our coverage on the topic “Chemical bonding and chemical structure”. We hope you enjoyed learning and were able to grasp the concepts. You can get separate articles as well on various subtopics of this article such as ‘Bond angle’, ‘Molecular orbital theory’ etc. on Vedantu website. We hope after reading this article you will be able to solve problems based on the topic. If you are looking for solutions of NCERT Textbook problems based on this topic, then log on to Vedantu website or download Vedantu Learning App. By doing so, you will be able to access free PDFs of NCERT Solutions as well as Revision notes, Mock Tests and much more.