The document discusses the nature of chemical bonds, defining them as the attractive forces between atoms. It explains why atoms combine, emphasizing the balance between attractive and repulsive forces, and outlines the octet rule for achieving stable electron configurations. The document categorizes chemical bonds into strong (ionic and covalent) and weak bonds, providing detailed characteristics and examples of each type of bond.

1. Lecturer
Department of Pharmacy
Northern University Bangladesh

2. References
1. Modern inorganic chemistry-R. D. Madan
2. Modern Inorganic Chemistry - S. Z. Haider
3. Essentials of Physical Chemistry -Bhal Tuli

3. A chemical bond is defined as the attractive force that holds
two or more atoms together in a molecule or an ion.
Why do atoms combine?
(Cause of chemical bonding)
By a close study of atoms and molecules it has been found
that atoms combine chemically for the following reasons:

4. Atoms consist of strongly positive nucleus and negative
electrons. When two atoms come closer to combine with
each other to form a bond between them, the attractive and
repulsive forces begin to operate between them. The
attractive forces are between the electrons of one atom and
the nucleus of other atom while the repulsive forces are
between the electrons or the nuclei of the two atoms.
When the two atoms approach each other these forces
counteract each other. The net result of these forces may be
either attraction or repulsion between the atoms.

5. On the other hand if the repulsive forces become dominant over
the attractive forces, the atoms do not combine and hence no
chemical bond is established between them.
Attractive force > repulsive force → formation of bond
For example in case of hydrogen atoms the net result is attraction
and two hydrogen atoms combine together to form H2 molecule.
On the other hand in case of He atom the net result is repulsion and
hence two He atoms do not combine together to form He2 molecule
If the attractive forces become
dominant over the repulsive forces
the net result is the attraction
between the atoms and hence they
combine together to form a chemic
al bond between them.

6. The atoms of noble gases do not normally react with other atoms to
form compounds. It is assumed that the outermost shell
configuration of the atoms of noble gases is a stable configuration
of 8 electrons which is known as octet. The two electrons in case of
He is also stable as octet which is known as doublet.

7. • The tendency of the atoms to have 8 electrons
in their outermost shell is known as octet rule
or rule of eight. Since helium atom has two
electrons, this rule is called doublet rule or
rule of two.
• The atoms possessing less than 4 electrons
generally tends to lose them while those
having more than 4 electrons in the outer most
shell tends to gain the electrons during the
chemical combination or bond formation to
attain stable configuration of the nearest
inert gas.

8. When two atoms combine together to form a bond there is an over
all decrease in the potential energy of the combining atoms, which
has greater stability.
The curve shown below represents the variation of potential
energy between the nuclei of the two atoms A and B which are
approaching closer to each other to form a bond between them.
The trend of the curve from right to left should be observed.
When two atoms A and B are far away, say at an infinite distance
from each other the attraction between them is zero. Hence there
is no possibility of formation of bond between them. This situation
has been represented by X.

9. • As the two atoms brought closer to each other i.e. as the
distance between the atoms is decreased, the attractive forces
between the nucleus and the electrons become more dominant
than the repulsive forces operating between the electrons of the
two atoms and hence energy of the system goes on decreasing
as shown by the downward trend of the curve. This decrease of
energy continues till a certain minimum value shown by Y in the
curve is obtained.
• Now if the atoms are brought still closer the repulsive forces
between the two nuclei at such small internuclear distance
becomes dominant and hence the energy of the system starts
increasing as shown by the upward trend of the curve.

10. The process by which the atoms of the elements rearrange their outer-most shell
electrons to get eight-electron outer-most shell configuration which is a stable
configuration takes place by the formation of chemical bond.
In general there are two main classes of bonds (or linkages) which hold the atoms
together in a molecule. These two main classes are subdivided into small groups.
Chemical Bonds
Strong Bonds Weak Bonds
1. Ionic or Electrovalent Bond
2. Covalent Bond
3. Co-ordinate or Deviate Bond
4. Metallic Bond
1. Hydrogen Bond
2. Van Deer Waals Interaction

11. • The chemical bond formed between two atoms by transfer of
one or more valence electrons from one atom to the other is
called ionic bond.
• This bond is also called electrovalent or polar bond.

12. Let us consider two atoms A & B; atom A has one electron in its
valence-shell while atom B has seven electrons. Thus A has one
electron in excess and B has one electron short than the stable
octet. Therefore A transfers an electron to B and converts into
positive ion (cation), A+. On the other hand, B accepts an electron
from A and converts into negative ion (anion), B–.
In this transaction both the atoms acquire a stable electronic
configuration of their nearest inert element. The resulting cation
and anion are held together by electrostatic force of attraction
which is called ionic bond or electrovalent bond.

13. 1) Number of Valence electron: The atom which is
converted into cation should posses 1, 2 or 3 valence
electrons while the atom which is converted into anion
should have 5, 6 or 7 valence electrons.
The element of group IA, IIA and IIIA satisfy this condition
for the cation and those of group VA, VIA and VIIA satisfy
this condition for anion.
2) The ionization energy of the metal atom should be low.
3) Electron affinity of the nonmetal should be high.

14. 4) The lattice energy of the ionic compound formed should high:
The energy released when one gram mole a crystal is formed from its
gaseous atoms is called the lattice energy of the crystal. Thus:
A+(g) + B-(g) A+B-(crystal) + Energy released (Lattice energy)
1 mole 1 mole 1 mole
Higher the value of the lattice energy of a crystal, the greater is the ease of
its formation i.e. greater will be the stability of the crystal.
5) Electronegativity difference of the two atoms forming the ionic bond
should be high:
In fact a difference of 2 or more is essential for the formation of ionic bond.
For example, since the electronegativity difference between Na and Cl is 2.1
(Na=0.9, Cl=3.0) Na and Cl will form an ionic bond in NaCl molecule.

15. 1. Physical state: Ionic compounds are crystalline solids at
room temperature.
2. Electrical conductivity: Ionic compounds do not conduct
electricity when they are in the solid state. The ionic solids conduct
electricity when they are water solution or in the molten state.
3. They are quite hard, have low volatility and high melting and
boiling points.
4. Solubility in polar and nonpolar solvent: Ionic solids are freely
soluble in polar solvents like H2O, liquid ammonia etc.
On the other hand ionic solids are insoluble or slightly soluble in
nonpolar solvents (organic solvents) such as C6H6, CCl4 etc.

16. 5. Crystal structure: Ionic solids do not exist as individual neutral
independent molecules rather they exist as three dimensional
solid aggregates which have definite geometric shape.
6. Highly brittle: Ionic solids are highly brittle, i.e. if a little external
force is applied on ionic crystals, they are easily broken. This
property is called brittleness.
7. High density: The electrostatic force of attraction existing
between the cation and anion in an ionic crystal bring these ions
very close to each other. This decreases the volume of crystal and
as a consequence this ionic crystal has high density.
8. They do not exhibit isomerism: Ionic bond involving
electrostatic lines of force between opposite ions is non-rigid and
non-directional. The ionic compounds are, therefore, incapable of
exhibiting stereoisomerism.

17. The chemical bond between two atoms in which the electrons
(in pair) are shared by both the participating atoms is called
covalent bond.

18. Hydrogen atom has one valence electron and oxygen atom
(2,6) has six valence electrons, and achieves the stable octet
by sharing two electrons with two hydrogen atoms (one with
each H atom.
Thus hydrogen atoms acquire 2
electrons and oxygen acquires 8
electrons in their respective outermost
shells and the shared electron pair
contributes a covalent bond between
hydrogen atoms and oxygen.
Some Other Examples: H2, N2, O2 ,
NH3 ,BF3, CH4 molecules

19. The covalent bond formed by the sharing of one, two or three
electron pairs between the participating atoms is called single,
double and triple covalent bond respectively.
Double and triple covalent bonds are called multiple covalent
bonds.
Type Representation No. of Electron Pairs Example
Single Covalent bond Single Dash (–) 1 pair=(1×2)=2 electrons
H–H, Cl–Cl
H–Cl,
H–O–H
Double Covalent bond Double Dash (=) 2 pairs=(2×2)=4 electrons O=O
Triple Covalent bond Triple Dash (≡) 3 pairs=(3×2)=6 electrons N≡N

20. 1. High ionization energy: Atoms which have high value of ionization energy
are incapable of losing electrons to form cations. Thus these elements form
covalent bond rather than ionic bond.
2. Equal electron affinities: For covalent bonding the two atoms must have
almost equal attraction for electrons.
3. Number of valence electron: each of the two atoms should have 5, 6 or 7
valence electrons (H atom has only one electron) so that both the atoms achieve
the octet by sharing 3, 2 or 1 electron pair. The nonmetals of VA, VIA and VIIA
groups respectively satisfy this condition.
4. Equal electronegativity: both atoms should have equal electronegativity so
that the transfer of electron(s) from one atom to another may not take place.
5. High nuclear charge and small internuclear distance: High charge on the
bonding nuclei and smaller internuclear distance favor the formation of covalent
bond.

21. 1. Physical state: Most of the covalent compounds are present as gases or
liquids of low boiling points under the normal condition of temperature and
pressure due to weak forces.
2. Covalent compounds don't conduct electricity in water: Electricity is
conducted in water from the movement of ions from one place to the other. These
ions are the charge carriers which allow water to conduct electricity. Since there
are no ions in a covalent compound, they don't conduct electricity in water.
3. Lower melting and boiling points: Except giant molecules other covalent
solids have relatively lower melting and boiling points than the ionic solids. Low
boiling points are due to the fact that the attractive forces between the covalent
molecules are weak van der Waal’s forces.
4. Solubility in polar and nonpolar solvents: Covalent solids are insoluble in
polar solvents like H2O but are readily soluble in nonpolar solvents like C6H6, CCl4
etc. Their solubility in nonpolar solvents is based on the principle ‘like dissolves
like’.

22. 5. Crystal structure: Crystals of solid covalent compounds are of two types:
Those in which every atom is bonded with other atoms by covalent bonds result
ing in the formation of giant molecule. Example: Diamond, silicon carbide (SiC),
aluminium nitride (AlN) etc.
Those which consist of separate layers. The covalent compounds containing
separate layers are said to have layer lattice structure. Example of such
compounds is: CdI2, CdCl2, BN, graphite etc.
6. Neither too hard nor too brittle: Covalent compound are neither hard nor
brittle. Rather they are soft and waxy, since they usually consist of separate
molecules.
7. Isomerism: Since covalent bonds are rigid and directional, they can give rise to
different arrangement of atoms in space. So a single molecular formula of a
covalent compound may represent a number of different compounds with different
properties. Thus covalent compounds can show isomerism.

23. Property
Ionic Compounds
(Electrovalent compounds)
Covalent Compounds
1. Physical State Crystalline solids at room temperature. Gases or liquids, solids(high molecular weight).
2. Crystal Structure
They consist of 3D solid aggregates
(ionic solids).
Two types (a) giant molecules (e.g. diamond )
and (b) separate layers (e.g. graphite).
3. Hardness and
brittleness
Hard and brittle. Soft and waxy.
4. Nature of reactions
They undergo ionic reactions (in solution)
which are fast and instantaneous.
They undergo molecular reactions
(in solution) which are slow.
5. Melting and boiling
points
Usually have high melting and boiling point. Covalent solids have low melting and boiling
points (exception-giant molecules).
6. Solubility
Freely soluble in polar solvents and insoluble
in or slightly soluble in nonpolar solvents.
Insoluble in polar solvents and readily soluble in
non-polar solvents.
7. Electrical conductivity
Conduct electricity in fused state or in
solution.
Poor conductors.
8. Formation
By the transfer of electrons from metal to a
non-metal.
By the sharing of electrons pair (s) between
non-metal atoms.
9. Isomerism Rare. Frequent.
10. Rigidity Non rigid and non-directional. Rigid and directional.

24. Electronegativity is a measure of the tendency of an atom to attract
a bonding pair of electrons.
The Pauling scale is the most commonly used. Fluorine (the most
electronegative element) is assigned a value of 4.0, and values
range down to caesium and francium which are the least
electronegative at 0.7.

25. In order to explain how a covalent bond is formed Heitler and London
in 1927 put forward a theory which is called Valence Bond Theory
(VBT).
The main postulates of Valence bond theory are as follows:
• A covalent bond is formed by the overlapping of two half-filled
valence atomic orbitals of two different atoms.
• The overlap of two atomic orbitals gives rise to a single bond
orbital which is a localised orbital and is occupied by both the
electrons having opposite spins.
• The electrons in the overlapping orbitals get paired and confined
between the nuclei of two atoms. The equilibrium distance at which
the two atomic nuclei are now held is called the Bond length.

26. • The electron density between two bonded atoms
increases due to overlapping. This confers stability to
the molecule. Greater the extent of overlapping,
stronger is the bond formed.
• The amount of energy given off or released per mole at
the time of overlapping of atomic orbitals to form a
bond is termed as Bond Energy or Stabilisation
Energy
Thus it is clear that, an atomic orbital of the outermost
shell of an atom containing one electron has a tendency to
overlap with another atomic orbital of another atom
containing one electron of opposite spin. This type of
overlap gives rise to the formation of a bond which is
called covalent bond.

27. A covalent bond is of two types depending on the type of
overlapping between the two atoms
(1) Sigma (σ) bond (2) Pi (π) Bond
A covalent bond which is formed between two atoms by the overlap
of their half-filled atomic orbitals along the line joining the nuclei of
both the atoms (i.e. along the nuclear axis, bond axis or molecular
axis, as it is called) is called a σ-bond.
In other words, when there is end to end overlapping of half-filled
atomic orbitals along the inter-nuclear axis, the bond resulted is
called sigma (σ) bond. This type of overlapping between the atomic
orbitals is also called “head-on” overlapping or “axial” overlapping.

28. (i) s-s overlap: In this type of overlap, half-filled s-orbital of one
atom overlaps with the half-filleds-orbital of the other atom.
e.g. The formation of hydrogen molecule from two H-atoms.
(ii) s-p overlap: In this type of overlap, s-orbital of one atom
overlaps with the half-filled p-orbital of the other atom.
e.g Formation of hydrogen chloride molecule from H- and Cl-atoms.
(iii) p-p overlap: Here singly-filled p-orbital of one atom overlaps
with the half-filled p-orbital of the other atom on internuclear axis
atoms (nuclear axis, bond axis or molecular axis).
e.g. The formation of Cl2 from two Cl-atoms.

29. A covalent bond which is formed between two atoms by the overlap of
their singly-filled p-orbitals along a line perpendicular to their nuclear
axis (side-to-side overlap) is called a π bond.
Example: Formation of O2 , N2 molecule
This type of overlapping takes place perpendicular to the inter-nuclear
axis as shown below:
p-p side-to-side overlap: This is also called side-wise or side-way or
lateral overlap. Here the overlap of two p-orbitals takes place along a
line perpendicular to the molecular axis. Hence the name lateral
overlap. This means that for this type of overlap the two p-orbitals must
be held parallel to each other, i.e. the orbital aids should be parallel.

30. Which one is strong bond. Sigma (σ) or Pi (π) ?

31. may precisely be
defined as the
phenomenon of
mixing up (or merging)
of orbitals of an atom
of nearly equal
energy, giving rise to
entirely new orbitals
equal in number to
the mixing orbitals
and having same
energy contents and
identical shapes.
Hybrid Orbital Geometry of Molecule Examples
sp Linear BeCl2, HgCl2, HC≡CH
sp2
Triangular Planar BBr3, BF3, H2C= CH2,C6H6
sp3
Tetrahedral CH4, SiCl4, H3C–CH3
sp2
d Square planar [Cu(NH3)4]+ 2
, Ni(CN)4
2–
, XeF4
sp3
d Trigonal bi-pyramidal PCl5
sp3
d2
Octahedral SF6, Co(NH3)6
+ 3
Fig. Possible hybridization schemes
Definition: Hybridization

32. The process by which an s and a p orbital mix and form two new
equivalent hybrid orbitals is called sp hybridization. This hybrid
orbital lie along a line and are , therefore, often referred to as
Linear hybrid orbitals.
The Electronic Configuration of carbon in the
ground state is 1s2 2s2 2p2. If energy is supplied to raise one of the
2s electrons to a higher energy level to fill the vacant 2p orbital,
the electronic configuration of carbon in the excited state is
1s2 2s1 2px
1 2py
1 2pz
1. Two sub-orbitals (i.e. 2s1 and 2px
1) of C
atom in the excited state mix together and redistribute their energy
to give a set of two new equivalent orbitals which are known as sp
hybrid orbitals.

33. One of the 'sp' hybrid orbitals overlaps
with the 'sp' hybrid orbital of the other
carbon atom to form a sp-sp sigma
bond. The remaining two sp hybrid
orbitals of the two carbon atoms
overlap with the s orbitals of two
hydrogen atoms to form two sigma
(sp-s) bonds.
The unhybridized 2py
1 and 2pz
1 orbitals
of two carbon atoms overlap laterally to
form two pi bonds. The combination of
one (sp-sp) sigma bond and two (p-p) pi
bonds gives rise to the triple bond of
acetylene.

34. Formation of ethyne
molecule. (a) Shows
scheme of overlaps;
(b) shows the bonds
σ bonds being
indicated by straight
lines; and (c) shows
ball-and-stick model
of ethyne (acetylene)
Thus acetylene
molecule contains
one σ and a two π
bonds between the
two carbons and
each carbon is
linked with one
H-atom through σ b
onds.

35. Theory: According to MOT, put forward by Hund and Muflikan all
the atomic orbitals (AOs) of the atoms participating in the formation
of the molecule approach nearer to each other and get mixed up to
give an equivalent number of new orbitals that now belong to the
molecule as a whole. These new orbitals are called molecular
orbitals (MOs).
According to this method, the molecular orbitals are formed by the
linear combination (addition or subtraction) of atomic orbitals of the
constituent atoms of the molecule.
Explanation: Let us consider the simplest case of H2 molecule
consisting of two hydrogen atoms represented by HA and HB.

36. The atomic orbitals of these atoms are represented by the wave
functions ψA and ψB. When these atoms approach each other there
come two possibilities.
(1) Molecular orbital is formed by the addition of wave functions of
atomic orbitals. It can be represented by ψ(MO) = ψA + ψB ………..(i)
The M.O. formed is called bonding molecular orbital. It lowers the
energy and brings about the stability in the system.
(2) Molecular orbital is formed by the subtraction of wave functions
of atomic orbitals. It can be represented by ψ*(MO) = ψA – ψB …….(ii)
The MO formed is called antibonding molecular orbital. This type
of MO corresponds to higher energy state. It has net disruptive
effect. That is why this MO is termed as antibonding molecular
orbital, distinguished by attaching an asterisk (*).

37. A covalent bond which is formed by mutual sharing of two electrons
both of which are provided entirely by one of the linked atoms or
ions is called co-ordinate bond.
Co-ordinate bond is also sometimes referred to as co-ordinate co
valent bond or dative bond. The compounds containing a
coordinate bond are called coordinate compounds. The pair of
shared electrons is called lone pair. The atom which furnishes the
electron pair is called donor or ligand while the other atom which
accepts the electron pair is called acceptor. A coordinate bond is
represented by an arrow which points away from the donor to the
acceptor.
e.g. H2O2, O3, SO2, SO3 Molecules

38. Definition: The particular type of bonding which holds the metal
atoms together in a metal crystal is called metallic bonding.
THE ELECTRON SEA MODEL : A metal is regarded as a group of
positive metal ions packed together as closely as possible in a
regular geometric pattern and immersed in a sea of electrons (called
electron-pool or electron-gas or electron-cloud) which move about
freely (mobile or delocalised electrons) in the vacant valence orbitals
.The attractive force that binds the metal ions to the mobile electrons
is called metallic bond and the force of attraction between the metal
ions and the mobile electrons holds the atoms together.

39. A hydrogen bond is the attractive force between the hydrogen attac
hed to an electronegative atom of one molecule and an electronega
tive atom of a different molecule.
In order to understand the concept
of hydrogen bond let us consider
a molecule, say, AH in which H
atom is linked with a strongly
electronegative but very small
atom X (X may be N, O or F) by a
normal covalent bond.

40. Since the electronegative atom, X attracts the electron pair towards
itself, thus will have a lone pair of electrons. This leaves H atom wit
h a large partial positive charge. Consequently AH molecule will be
have as a dipole which is represented as:
Evidently the dipole has X as its negative end and H as its positive
end.
Now if another molecule like XH (same molecule) or YH (different
molecule) (X and Y are strongly electronegative atoms) which also
forms a dipole X - or Y - H respectively is brought near A - dipole, t
hese two dipoles will be attracted towards each other by electrostat
ic force of attraction which is represented by a dotted or dashed lin
e and is called hydrogen bond or hydrogen bonding.

41. Inter-molecular hydrogen bond: This type of hydrogen bond
occurs between two or more molecules of the same or different
compound. e.g.: NH3, H2O etc.
Intra-molecular hydrogen bond: This type of hydrogen bond is
formed between a hydrogen atom and an electronegative atom
present in the same molecules. O
C
H
O
H
o-hydroxy benzaldehyde

42. 1. Physical state of water: Without hydrogen bonding, H2O would
have existed as a gas like H2S. In that case no life would have been
possible.
2. Applications in biological investigations: Hydrogen
bonding also exists in molecules of living system like various
tissues, organs, skin in animals. The molecular conformation of
protein, nucleic acid etc. are determined by hydrogen bond.

43. Protein: Hydrogen bonding plays an important role in determining
the three-dimensional structures adopted by proteins. In these
macromolecules, bonding between parts of the same
macromolecule cause it to fold into a specific shape, which helps
determine the molecule's physiological or biochemical role.
The breakdown of protein into amino acid is due to the weak hydrogen bond.
The NH and CO vertically adjacent to each other in the helix are linked
together by H- bond.
DNA: The double helical structure of DNA, for example, is due
largely to hydrogen bonding between the base pairs, which link one
complementary strand to the other and enable replication.
Drug action: Hydrogen bonding increases the activity of drug.
For example, (–) ephedrine is more active than (+) ephedrine
since (–) ephedrine forms hydrogen bonding with biological receptor

45. Van der Waals forces are very short lived intermolecular attractive
forces which are believed to exist between all kinds of atoms,
molecules and ions when they are sufficiently close to each other.
1. Dipole-dipole interaction (Keesom Forces)
2. Ion-dipole interaction
3. Dipole-induced dipole interaction (Debye Forces)
4.Instantaneous dipole-induced dipole interaction
(London Forces)

46. 1. Dipole-dipole interaction (Keesom Forces): These forces are
found in polar molecules having permanent polarity in them. When
polar molecules are brought nearer to each other, they orient
themselves in such a way that the positive end of the one dipole
attracts the negative end of another dipole and vice versa. This is the
strongest of all other types of van der Waals forces.
2. Ion-dipole interaction
When NaCl is put in H2O, it dissolves in it, since the negative ends of
water molecule dipoles aggregate around Na+ ions and the positive
ends around Cl- ions.

47. 3. Dipole-induced dipole interaction (Debye Forces)
This type of force is found in a mixture containing polar and nonpolar molecules. When
a nonpolar molecule is brought near to a polar molecule, the positive end of the polar
molecule attracts the mobile electrons of the non-polar molecule and thus polarity is
induced in nonpolar molecule. Now both the molecules become dipoles and
hence the positive end of the polar molecule attracts the displaced electron cloud of no
npolar molecule.
4. Instantaneous dipole-induced dipole interaction (London Forces)
A nonpolar atom or molecule may be visualized as a positive center surrounded by
asymmetrical negative electron cloud. But as the electron cloud oscillates, the electron
cloud becomes denser on one side of the molecule than on the other side. The
displacement of electron cloud creates an instantaneous dipole temporarily. Thus the
non-polar molecule becomes temporarily polar. Now when the self-polarized molecule
is brought near to a non-polar molecule, it polarizes the neighboring molecule by
disturbing its electronic distribution and thus an induced dipole is created. These
temporarily polar molecules are mutually attracted by weak attractive forces.