Chemical Bonding And Molecular Structure Class 11 Notes Chemistry Chapter 4 - CBSE

Chapter : 4

What Are Chemical Bonding And Molecular Structure ?

Chemical Bond

The attractive force which holds various constituents (atoms, ions, etc.) together in different chemical species is called a chemical bond.

Kossel-lewis Approach To Chemical Bonding

Lewis postulated that atoms achieve the stable octet when they are linked by chemical bonds. G.N. Lewis, an American chemist introduced simple notations to represent valence electrons in an atom. These notations are called Lewis symbols. For example, the Lewis symbols for the elements of second period are as under: 


Significance Of Lewis Symbols

  • The number of dots around the symbol represents number of valence electrons which helps to calculate the common or group valence of the element.
  • Generally, the group valence of the elements is either equal to the number of dots in Lewis symbols or 8 minus the number of dots or valence electrons.

Octet Rule

According to Kössel and Lewis electronic theory of chemical bonding, atoms can combine either by transfer of valence electrons from one atom to another (gaining or losing) or by sharing of valence electrons in order to have an octet in their valence shells. This is known as octet rule.

Limitations Of The Octet Rule

  • The incomplete octet of the central atoms.
  • Odd-electron molecules.
  • The expanded Octet.

Other Drawbacks Of Octet Theory

  • Some noble gases, also combine with oxygen and fluorine to form a number of compounds like XeF2 , XeOF2 etc.
  • This theory does not account for the shape of the molecule.
  • It does not give any idea about the energy of the molecule and relative stability.

Modes Of Chemical Combination

  • Electrovalent bond or ionic bond
  • Covalent bond
  • Co-ordinate bond

Covalent Bond

The bond which is formed by the equal sharing of electrons between one or two atoms is called covalent bond. In these bonds, electrons are contributed by both.

  • When two atoms share one electron pair they are said to be joined by a single covalent bond.
  • When two atoms share two pairs of electrons, the covalent bond between them is called a double bond.
  • When combining atoms share three electron pairs, a triple bond is formed.

Ionic Or Electrovalent Bond

Ionic or Electrovalent bond is formed by the complete transfer of electrons from one atom to another. Generally, it is formed between metals and non-metals. We can say that it is the electrostatic force of attraction which holds the oppositely charged ions together. The compounds which is formed by ionic or electrovalent bond is known as electrovalent compounds.


It is the number of electrons lost or gained during the formation of an ionic bond or electrovalent bond.

Lewis Representation Of Simple Molecules (The Lewis Structures)

The Lewis dot Structure can be written through the following steps:

  • Calculate the total number of valence electrons of the combining atoms.
  • Each anion means addition of one electron and each cation means removal of one electron. This gives the total number of electrons to be distributed.
  • By knowing the chemical symbols of the combining atoms.
  • After placing shared pairs of electrons for single bond, the remaining electrons may account for either multiple bonds or as lone pairs. It is to be noted that octet of each atom should be completed.

Formal Charge

In a polyatomic molecule or ion, the formal charge may 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 is expressed as :

Formal Charge (F.C.) on an atom in a Lewis structure =[total number of valence electrons in the free atom] - [total number of non bonding (lone pair) electrons]-(1/2)[total number of bonding (shared, electrons]

Bond Parameters

Bond Length

It is defined as the equilibrium distance between the centres of the nuclei of the two bonded atoms. It is expressed in terms of A. Experimentally, it can be defined by X-ray diffraction or electron diffraction method.

  • The covalent radius is measured approximately as the radius of an atom’s core which is in contact with the
    core of an adjacent atom in a bonded situation
  • The Van der Waal’s radius represents the overall size of the atom which includes its valence shell in a non-bonded situation.

Bond Angle

It is defined as -the angle between the lines representing the orbitals containing the bonding – electrons.

It helps us in determining the shape and can be expressed in degree. Bond angle can be experimentally determined by spectroscopic methods.

Bond Enthalpy

Enthalpy is defined as the amount of energy required to break one mole of bonds of a particular type to separate them into gaseous atoms. It is also known as bond dissociation enthalpy or simple bond enthalpy. Unit of bond enthalpy = kJ mol-1 . It depends on the following two factors:

  • Greater the bond enthalpy, stronger is the bond. For e.g., the H—H bond enthalpy in hydrogen is 435.8 kJ
  • Greater the bond multiplicity, more will be the bond enthalpy. For e.g., bond enthalpy of C —C bond is 347 kJ mol-1 while that of C = C bond is 610 kJ mol-1.

In polyatomic molecules, the term mean or average bond enthalpy is used.

ISO Electronic Molecules

Iso electronic molecules and ions have identical bond orders; for example, F2 and O2 2– have bond order 1. N2 ,CO and NO+ have bond order 3.

Resonance Structures

There are many molecules whose behaviour cannot be explained by a single-Lewis structure.

Resonance in the O3 molecule
(Structures I and II represent the two canonical forms while the Structure III is the resonance hybrid)

According to the concept of resonance, positions of nuclei, bonding and non-bonding pairs of electrons are taken as the canonical structure of the hybrid which describes the molecule accurately. For 03 , the two structures shown above are canonical structures and the III structure represents the structure of 03 more accurately. This is also called resonance hybrid.

Resonating Structure Of Phenol


Resonating Structure Of Nitrobenzene

Polarity Of Bonds

Non-Polar Covalent Bonds

When the atoms joined by covalent bond are the same like; H2 , 02 , Cl2 , the shared pair of electrons is equally attracted by two atoms and thus the shared electron pair is equidistant to both of them.

Polar Covalent Bonds

When covalent bonds formed between different atoms of different electronegativity, shared electron pair between two atoms gets displaced towards highly electronegative atoms.

Dipole Moment

It can be defined as the product of the magnitude of the charge and the distance between the centres of positive and negative charge. It is usually designated by a Greek letter ‘µ’. Mathematically, it is expressed as,

Dipole moment (µ) = charge (Q) × distance of separation (r) Dipole moment is usually expressed in Debye units (D). The conversion factor is 1 D = 3.33564 × 10–30 Cm Where, C is coulomb and m is meter.

The Valence Shell Electron Pair Repulsion (VSEPR) Theory

Sidgwick and Powell in 1940, proposed a simple theory based on repulsive character of electron pairs in the valence shell of the atoms. It was further developed by Nyholm and Gillespie (1957).

Main Postulates are the following

  • The exact shape of molecule depends upon the number of electron pairs (bonded or non bonded) around the
    central atoms.
  • The electron pairs have a tendency to repel each other since they exist around the central atom and the
    electron clouds are negatively charged.
  • Electron pairs try to take such position which can minimize the repulsion between them.
  • The valence shell is taken as a sphere with the electron pairs placed at maximum distance.
  • A multiple bond is treated as if it is a single electron pair and the electron pairs which constitute the bond
    as single pairs.

The repulsive interaction of electron pairs decrease in the order:

Lone pair (lp) – Lone pair (lp) > Lone pair (lp)– Bond pair (bp) > Bond pair (bp) – Bond pair (bp)

Applications Of Dipole Moment

  • For determining the polarity of the molecules.
  • In finding the shapes of the molecules. For example, the molecules with zero dipole moment will be linear or symmetrical. Those molecules which have unsymmetrical shapes will be either bent or angular. (e.g., NH3 with µ = 1.47 D).
  • In calculating the percentage ionic character of polar bonds.

Valence Bond Theory

Valence bond theory was introduced by Heitler and London (1927) and developed by Pauling and others. It is based on the concept of atomic orbitals and the electronic configuration of the atoms. According to this
theory, as these two atoms come closer new attractive and repulsive forces begin to operate.

Thus, attractive forces arise between:

  • The nucleus of one atom is attracted towards its own electron and the electron of the other and vice versa.
  • Repulsive forces arise between the electrons of two atoms and nuclei of two atoms. Attractive forces tend to
    bring the two atoms closer whereas repulsive forces tend to push them apart.

Orbital Overlap Concept

According to orbital overlap concept, covalent bond formed between atoms results in the overlap of orbitals belonging to the atoms having opposite spins of electrons.


The stability of a Molecular orbital depends upon the extent of the overlap of the atomic orbitals.

Types Of Orbital Overlap

Depending upon the type of overlapping, the covalent bonds are of two types,

Sigma (σ-bond)

Sigma bond is formed by the end to end (head-on) overlap of bonding orbitals along the internuclear axis. The axial overlap involving these orbitals is of three types:

  • s-s overlapping: In this case, there is overlap of two half-filled s-orbitals along the internuclear axis.
  • s-p overlapping: This type of overlapping occurs between half-filled s-orbitals of one atom and half-filled
    p-orbitals of another atom.
  • p-p overlapping: This type of overlapping takes place between half-filled p-orbitals of the two approaching atoms.

pi (π-bond)

π-bond is formed by the atomic orbitals when they overlap in such a way that their axes remain parallel to each other and perpendicular to the internuclear axis. The orbital formed is due to lateral overlapping or side wise overlapping.


Hybridisation is the process of intermixing of the orbitals of slightly different energies so as to redistribute their energies resulting in the formation of new set of orbitals of equivalent energies and shape.

Salient Features of Hybridisation

  • Orbitals with almost equal energy take part in the hybridisation.
  • Number of hybrid orbitals produced is equal to the number of atomic orbitals mixed,
  • Geometry of a covalent molecule can be indicated by the type of hybridisation.
  • The hybrid orbitals are more effective in forming stable bonds than the pure atomic orbitals.

Conditions necessary for hybridisation

  • Orbitals of valence shell take part in the hybridisation.
  • Orbitals involved in hybridisation should have almost equal energy.
  • Promotion of electron is not necessary condition prior to hybridisation.
  • In some cases filled orbitals of valence shell also take part in hybridisation.

Types of Hybridisation

  • sp hybridisation: When one s and one p-orbital hybridise to form two equivalent orbitals, the orbital is
    known as sp hybrid orbital, and the type of hybridisation is called sp hybridisation. Each of the hybrid orbitals formed has 50% s-character and 50%, p-character. This type of hybridisation is also known as diagonal hybridisation.
  • sp2 hybridisation: In this type, one s and two p-orbitals hybridise to form three equivalent sp2 hybridised
    orbitals. All the three hybrid orbitals remain in the same plane making an angle of 120°.
  • sp3 hybridisation: In this type, one s and three p-orbitals in the valence shell of an atom get hybridised to form four equivalent hybrid orbitals. There is 25% s-character and 75% p-character in each sp3 hybrid orbital.
    The four sp3 orbitals are directed towards four corners of the tetrahedron.

 Linear Combination Of Atomic Orbitals (Lcao)

The formation of molecular orbitals can be explained by the linear combination of atomic orbitals. Combination takes place either by addition or by subtraction of wave function as shown along side. The molecular orbital formed by addition of atomic orbitals is called bonding molecular orbital while molecular orbital formed by subtraction of atomic orbitals is called antibonding molecular orbital.


Conditions for the combination of atomic orbitals

  • The combining atomic orbitals must have almost equal energy.
  • The combining atomic orbitals must have same symmetry about the molecular axis.
  • The combining atomic orbitals must overlap to the maximum extent.

Types Of Molecular Orbitals

  • Sigma (σ) Molecular Orbitals: They are symmetrical around the bond-axis.
  • pi (π) Molecular Orbitals: They are not symmetrical, because of the presence of positive lobes above and
    negative lobes below the molecular plane.

Electronic Configuration And Molecular Behaviour

The distribution of electrons among various molecular orbitals is called electronic configuration of the molecule.

Stability Of Molecules

If Nb = Number of electrons occupying bonding orbitals.

Na = Number of electrons occupying antibonding orbitals.


(i) If Nb > Na molecule will be stable.

(ii) If Nb > Na molecule will be unstable.

Bond Order

Bond order is defined as half of the difference between the number of electrons present in bonding and antibonding molecular orbitals.

Bond order (B.O.) = 1/2 [Nb -Na]

The bond order may be a whole number, a fraction or even zero. It may also be positive or negative.

  • Nature of the bond: Integral bond order value for single double and triple bond will be 1, 2 and 3
  • Bond-Length: Bond order is inversely proportional to bond-length. Thus, greater the bond order, smaller will be the bond-length.
  • Magnetic Nature: If all the molecular orbitals have paired electrons, the substance is diamagnetic. If one or more molecular orbitals have unpaired electrons, it is paramagnetic e.g., O2 molecule.

Hydrogen Bonding

When highly electronegative elements like nitrogen, oxygen, fluorine are attached to hydrogen to form covalent bond, the electrons of the covalent bond are shifted towards the more electronegative atom. Thus, partial positive charge develops on hydrogen atom which forms a bond with the other electronegative atom. This bond is known as hydrogen bond and it is weaker than the covalent bond. For example, in HF molecule, hydrogen bond exists between hydrogen atom of one molecule and fluorine atom of another molecule. It can be depicted as :


Types of Hydrogen Bonds

  • Intermolecular hydrogen bond: It is formed between two different molecules of the same or different compounds. For Example, in HF molecules, water molecules etc.
  • Intramolecular hydrogen bond: In this type, hydrogen atom is in between the two highly electronegative F, N, O atoms present within the same molecule. For example, in o-nitrophenol, the hydrogen is in between the two oxygen atoms.

H - - - F > H - - - O > H - - - N

10 kcal/mol > 7 kcal/mol > 2.0 kcal/mol