Chemical bonding can occur via ionic bonds or covalent bonds. Ionic bonds form when electrons are transferred from one atom to another, leaving cation and anion. Covalent bonds form when atoms share electrons via overlapping orbitals. The octet rule states that atoms seek to obtain eight electrons in their valence shell. Hybridization is the mixing of atomic orbitals to form new hybrid orbitals for bonding. Common hybridizations include sp, sp2, and sp3 which determine molecular geometry. Bond properties like order, length, energy are influenced by hybridization.

1. CHEMICAL BONDING
UNIT 3

2. Chemical Bond
It is defined as the attractive forces which hold the various chemical
constituents (atoms, ions, etc.) together in different chemical species.
Kossel-Lewis Approach to Chemical Bonding
According to this theory. atoms take part in the bond formation to
complete their octet or to acquire the electronic configuration of the
nearest inert gas atoms (Octet rule). This can be achieved by gaining,
losing or sharing the electrons.

3. Lewis Symbols
Valence electrons are reported by dots around the chemical symbol
of element, e.g.,
Ionic Bond
A chemical bond formed by complete transference of electrons from
one atom (metal) to another (non-metal) and hence, each atom
acquire the stable nearest noble gas configuration, is called ionic
bond or electrovalent bond, e.g., formation of sodium chloride

4. Factors affecting the formation of ionic bond
The factors are:
•Low ionization energy
•High electron affinity
•High lattice enthalpy
Low ionization energy: The metals with low ionization energy favor the
formation of ionic bond. As lower is the ionization energy more readily it will
lose electrons.
•High electron gain enthalpy: The non-metal participating should have
high electron gain enthalpy because more it will have attraction, for
upcoming electron more readily the bond will be formed.
•Lattice enthalpy: It is the amount of energy needed to break one mole of
bonds into its constituents, or the energy released when constituents
combine to form on 1 mole of a compound. More is the lattice energy, more
stable is the bond formed. All the compounds in which ionic bond is present
are called as ionic compounds.

5. General properties of ionic compounds:
1.Physical state: They form definite pattern that is crystal lattice. Crystal lattice
is 3D arrangement of cation and an anion.
For example, in NaCl crystal due to crystal formation they all are solids due to
strong bonding between constituents.
2.Melting and boiling point: They have high melting and boiling points
because of strong attraction between constituents.
3.Solubility: We know like dissolves like. So, polar compounds are soluble in
polar solvents. Now, ionic compounds have a charge that is they are polar.
Therefore, they will dissolve in polar solvents like water. So, all ionic compounds
are soluble in polar solvents and insoluble in organic solvents.
4.Electrical conductivity: It is due to free movement of free ions when ionic
compound is dissolved in water. When dissolved they break into ions and
conduct electricity.
5.Non directional in nature: When we are talking of directions in 3D structure,
we are talking about 3 coordinate .So, in NaCl or any other ionic compounds the
ion can take place in any direction .There direction is not fixed. Therefore, they
are non-directional in nature.

6. COVALENT BOND
They are formed by mutual sharing of electrons between combining atoms
Hydrogen chloride molecule Oxygen molecule
In Nitrogen molecule

7. Terms to know:
•Valence electrons: The electrons present in valence shell.
•Lone pair: The electrons that do not participate in bond formation.
•Bond pair: The electrons that participate in bond formation.
Lewis dot structures:
1.Chlorine molecule
Carbondioxide molecule
3. Methane

8. Coordinate bond
Coordinate bond is formed when shared pair of electrons comes only from one
atom. There is no mutual sharing of electrons. The one that donates electron is
called donor atom and other is called acceptor. The bond is also called dative
bond.

9. Lewis Representation of Simple Molecules (the Lewis
Structures)
The Lewis dot Structure can be written through the following steps:
(i) Calculate the total number of valence electrons of the combining
atoms.
(ii) 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.
(iii) By knowing the chemical symbols of the combining atoms.
(iv) 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.

12. Formal Charge
In polyatomic ions, the net charge is the charge on the ion as a whole
and not by particular atom. However, charges can be assigned to
individual atoms or ions. These are called formal charges.
It can be expressed as

13. Limitations of the Octet Rule
1.The incomplete octet of the central atoms:
In some covalent compounds central atom has less than eight electrons, i.e., it has an
incomplete octet. For example,
Li, Be and B have 1, 2, and 3 valence electrons only.

14. 2.Odd-electron molecules: There are certain molecules which have odd
number of electrons the octet rule is not applied for all the atoms.
The expanded Octet: In many compounds there are more than eight
valence electrons around the central atom. It is termed as expanded
octet. For Example,

15. Other Drawbacks of Octet Theory
(i) Some noble gases, also combine with oxygen and fluorine to form a
number of compounds like XeF2 , XeOF2 etc.
(ii) This theory does not account for the shape of the molecule.
(iii) It does not give any idea about the energy of The molecule and
relative stability.
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.

16. 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. It can be expressed in degree. Bond
angle can be experimentally determined by spectroscopic methods.
• Bond Enthalpy
It is defined as the amount of energy required to break one mole of bonds of a
particular type to separate them into gaseous atoms.
Bond Enthalpy is also known as bond dissociation enthalpy or simple bond
enthalpy. Unit of bond enthalpy = kJ mol-1
Greater the bond enthalpy, stronger is the bond. For e.g., the H—H bond
enthalpy in hydrogen is 435.8 kJ mol-1.
The magnitude of bond enthalpy is also related to bond multiplicity. 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.

17. Bond Order
According to Lewis, in a covalent bond, the bond order is given by the
number of bonds between two atoms in a molecule. For example,
Bond order of H2 (H —H) =1
Bond order of 02 (O = O) =2
Bond order of N2 (N = N) =3
Isoelectronic 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. With the increase in bond order, bond enthalpy increases and
bond length decreases. For example,

18. Resonance Structures
There are many molecules whose behaviour cannot be explained by a single-Lew
is structure, Tor example, Lewis structure of Ozone represented as follows:
Thus, according to the concept of resonance,
whenever a single Lewis structure cannot
explain all the properties of the molecule, the
molecule is then supposed to have many
structures with similar energy. Positions of
nuclei, bonding and nonbonding 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.

19. Polarity of Bonds
Polar and Non-Polar Covalent 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.
Alternatively, we can say that it lies exactly in the centre of the bonding atoms. As
a result, no poles are developed and the bond is called as non-polar covalent
bond. The corresponding molecules are known as non-polar molecules.
For Example,

20. Polar bond: When covalent bonds formed between different atoms
of different electronegativity, shared electron pair between two atoms
gets displaced towards highly electronegative atoms.
For Example, in HCl molecule, since electronegativity of chlorine is
high as compared to hydrogen thus, electron pair is displaced more
towards chlorine atom, thus chlorine will acquire a partial negative
charge (δ–) and hydrogen atom have a partial positive charge (δ+)
with the magnitude of charge same as on chlorination. Such covalent
bond is called polar covalent bond.

21. Dipole Moment
Due to polarity, polar molecules are also known as dipole molecules and they possess dipole
moment. Dipole moment is defined as the product of magnitude of the positive or negative
charge and the distance between the charges.

22. 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:
(i) The exact shape of molecule depends upon the number of electron pairs
(bonded or non bonded) around the central atoms.
(ii) The electron pairs have a tendency to repel each other since they exist
around the central atom and the electron clouds are negatively charged.
(iii) Electron pairs try to take such position which can minimize the rupulsion
between them.
(iv) The valence shell is taken as a sphere with the electron pairs placed at
maximum distance.
(v) A multiple bond is treated as if it is a single electron pair and the electron
pairs which constitute the bond as single pairs.

25. Valence Bond Theory of Covalent Bond
According to this theory, a covalent bond is formed by the overlapping of two half-
filled atomic orbitals having electrons with opposite spins. It is based on wave nature
of electron.
Depending upon the type of overlapping, the covalent bonds are of two types, known as
sigma (σ ) and pi (π) bonds.
1. Sigma Bond (σ bond) Sigma bond is formed by the end to end (head-on)
overlap of bonding orbitals along the internuclear axis.
The following result in the formation of σ bond.
(i) s-s overlapping
(ii) S-p overlapping
(iii) P-p head to head overlapping (axial)

26. s-s overlapping: In this case, there is overlap of two half-filled s-orbitals along
the internuclear axis as shown below:
s-p overlapping: This type of overlapping occurs between half-filled
s-orbitals of one atom and half filled p-orbitals of another atoms.

27. p-p overlapping: This type of overlapping takes place between half filled p-orbitals of
the two approaching atoms.

28. Examples

29. 2. Pi Bond (π bond)
It is formed by the sidewise or lateral overlapping between p-
atomic orbitals [pop side by side or lateral overlapping]
π bond is a weaker bond than σ bond.
Strength of Sigma and pf Bonds
Sigma bond (σ bond) is formed by the axial overlapping of the atomic
orbitals while the π-bond is formed by side wise overlapping. Since
axial overlapping is greater as compared to side wise. Thus, the sigma
bond is said to be stronger bond in comparison to a π-bond.
Distinction between sigma and n bonds

31. Distinction between sigma and pi bonds

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

33. Conditions necessary for hybridisation:
(i) Orbitals of valence shell take part in the hybridisation.
(ii) Orbitals involved in hybridisation should have almost equal
energy.
(iii) Promotion of electron is not necessary condition prior to
hybridisation.
(iv) In some cases filled orbitals of valence shell also take part in
hybridisation.

34. Types of Hybridisation:
(i) 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.

35. (ii) 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°. Example. A few compounds in which sp2 hybridisation takes place are
BF3, BH3, BCl3 carbon compounds containing double bond etc.

38. 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.

39. The angle between sp3 hybrid orbitals
is 109.5°.
A compound in which
sp3 hybridisation occurs is, (CH4). The
structures of NH2 and H20 molecules
can also be explained with the help of
sp3 hybridisation.

40. In NH3 , the valence shell (outer) electronic configuration of nitrogen in the
ground state is
sp3 hybridization in NH3 molecule

41. In NH3 molecule three unpaired electrons in the sp3 hybrid orbitals and a
lone pair of electrons is present in the fourth one. These three hybrid orbitals
overlap with 1s orbitals of hydrogen atoms to form three N–H sigma bonds.
We know that the force of repulsion between a lone pair and a bond pair is
more than the force of repulsion between two bond pairs of electrons. The
molecule thus gets distorted and the bond angle is reduced to 107° from
109.5°. The geometry of such a molecule will be pyramidal

42. In case of H2O molecule, the four oxygen orbitals (one 2s and three 2p) undergo sp3 hybridisation forming
four sp3 hybrid orbitals out of which two contain one electron each and the other two contain a pair of
electrons. These four sp3 hybrid orbitals acquire a tetrahedral geometry, with two corners occupied by
hydrogen atoms while the other two by the lone pairs. The bond angle in this case is reduced to 104.5°
from 109.5° and the molecule thus acquires a V-shape or angular geometry
sp3 hybridization in H2O molecule

43. Therefore in ethane C–C bond length is 154 pm and each C–H
bond length is 109 pm

44. sp2 Hybridisation in C2H4 :

45. In the formation of ethene molecule, one of the sp2 hybrid orbitals of carbon atom
overlaps axially with sp2 hybridised orbital of another carbon atom to form C–C sigma
bond. While the other two sp2 hybrid orbitals of each carbon atom are used for
making sp2–s sigma bond with two hydrogen atoms. The unhybridised orbital (2px
or 2py ) of one carbon atom overlaps sidewise with the similar orbital of the other
carbon atom to form weak π bond, which consists of two equal electron clouds
distributed above and below the plane of carbon and hydrogen atoms. Thus, in
ethene molecule, the carboncarbon bond consists of one sp2–sp2 sigma bond and
one pi (π ) bond between p orbitals which are not used in the hybridisation and are
perpendicular to the plane of molecule; the bond length 134 pm. The C–H bond is
sp2–s sigma with bond length 108 pm. The H– C–H bond angle is 117.6° while the
H–C–C angle is 121°. The formation of sigma and pi bonds in ethene

46. sp Hybridisation in C2H2 : In the formation of ethyne molecule, both the carbon atoms
undergo sp-hybridisation having two unhybridised orbital i.e., 2py and 2px.
One sp hybrid orbital of one carbon atom overlaps axially with sp hybrid orbital of the
other carbon atom to form C–C sigma bond, while the other hybridised orbital of each
carbon atom overlaps axially with the half filled s orbital of hydrogen atoms forming s
bonds. Each of the two unhybridised p orbitals of both the carbon atoms overlaps sidewise to
form two p bonds between the carbon atoms.
So the triple bond between the two carbon atoms is made up of one sigma and two pi
bonds as shown in Fig.

47. (i) Formation of PCl5 (sp3d hybridisation):
The ground state and the excited state outer electronic configurations of phosphorus (Z=15)
are represented below.

49. (ii) Formation of SF6 (sp3d2 hybridisation):
In SF6 the central sulphur atom has the ground state outer electronic configuration
3s23p4. In the exited state the available six orbitals i.e., one s, three p and two d are singly
occupied by electrons. These orbitals hybridise to form six new sp3d2 hybrid orbitals, which
are projected towards the six corners of a regular octahedron in SF6. These six sp3d2
hybrid orbitals overlap with singly occupied
orbitals of fluorine atoms to form six S–F sigma
bonds. Thus SF6 molecule has a regular
octahedral geometry as shown in Fig

50. Octahedral geometry of SF6 molecule

51. MOLECULAR ORBITAL THEORY
The salient features of this theory are :
(i) The electrons in a molecule are present in the various molecular orbitals as the electrons of atoms are
present in the various atomic orbitals.
(ii) The atomic orbitals of comparable energies and proper symmetry combine to form molecular orbitals.
(iii) While an electron in an atomic orbital is influenced by one nucleus, in a molecular orbital it is influenced by
two or more nuclei depending upon the number of atoms in the molecule. Thus
an atomic orbital is monocentric while a molecular orbital is polycentric.
(iv) The number of molecular orbital formed is equal to the number of combining atomic orbitals. When two
atomic orbitals combine, two molecular orbitals are formed. One is known as bonding molecular orbital while
the other is called antibonding molecular orbital.
(v) The bonding molecular orbital has lower energy and hence greater stability than the corresponding
antibonding molecular orbital.
(vi) Just as the electron probability distribution around a nucleus in an atom is given by an atomic orbital, the
electron probability distribution around a group of nuclei in a molecule is given by a molecular orbital.
(vii) The molecular orbitals like atomic orbitals are filled in accordance with the aufbau principle obeying the
Pauli’s exclusion principle and the Hund’s rule

52. Energy Level Diagram for Molecular Orbitals
The bond order of H2 molecule can be
calculated as given below: