This document discusses different types of chemical bonds including ionic bonding, covalent bonding, and metallic bonding. Ionic bonding involves the electrostatic attraction between oppositely charged ions when atoms gain or lose electrons. Covalent bonding occurs when atoms share pairs of electrons to gain stability. Metallic bonding results from the attraction between positively charged atomic nuclei and delocalized electrons in metals that act as the binding medium. The importance of chemical bonding is that it allows atoms to join together to form molecules and structures with unique physical and chemical properties essential for life.

1. CHEMICAL BONDING
BY: SEENAM IFTIKHAR

2. Chemical bond is an effect that causes certain
atoms to join together to form enduring
structures that have unique physical and
chemical properties.
DEFINITION:

3. EXAMPLE:

4. The importance of chemical bonding is nicely
illustrated by the structures of the two
compounds; ethanol and dimethyl-ether, both of
which have the simplest formula C2H6O.
The structural formulas reveal the very different
connectivity's of these two molecules whose
physical and chemistry properties are quite
different:
IMPORTANCE:

6. Ionic bonding
PRESENTED BY: MAZNA SALEEM

7. Ionic Bonding is a type of chemical bond that
involves the electrostatic attraction between
oppositely charged ions. These ions represent
atoms that have lost one or more electrons
(known as cations) and atoms that have
gained one or more electrons (known as
anions).
IONIC BONDING

8. EXAMPLE
• Sodium (2,8,1) has 1 electron more than a stable
noble gas structure (2,8). If it gave away that
electron it would become more stable.
• Chlorine (2,8,7) has 1 electron short of a stable
noble gas structure (2,8,8). If it could gain an
electron from somewhere it too would become
more stable.
Bonding in NaCl

9. HOW IT TRANSFERS???

10. EXAMPLE OF NaF...

11. EXAMPLE OF MgO...

12. BY: HIBA MUSHTAQ

13. INTRODUCTION
• Covalent bonding occurs when pairs of
electrons are shared by atoms. Atoms will
covalently bond with other atoms in order to
gain more stability, which is gained by forming
a full electron shell.
• Bond pair.
• Lone pair.

14. OCTET RULE
• The Octet Rule requires all atoms in a
molecule to have 8 valence electrons.

15. SINGLE BONDS
• A single bond is when two electrons--one pair
of electrons--are shared between two
atoms. It is depicted by a single line between
the two atoms.

16. EXAMPLE 1: HCl

17. DOUBLE BONDS
• A Double bond is when two atoms share two
pairs of electrons with each other. It is
depicted by two horizontal lines between two
atoms in a molecule.

18. EXAMPLE 2: Co2

19. TRIPLE BOND
• A Triple bond is when three pairs of electrons
are shared between two atoms in a molecule.
It is the least stable out of the three general
types of covalent bonds.

20. EXAMPLE 3: ACETYLENE

21. POLAR COVALENT BOND
• A Polar Covalent Bond is created when the
shared electrons between atoms are not
equally shared. This occurs when one atom
has a higher electronegativity than the atom it
is sharing with.
• Examples: Water, Sulfide, Ozone, etc.

22. EXAMPLES

23. NONPOLAR COVALENT BOND
• A Nonpolar Covalent Bond is created when
atoms share their electrons equally. This
usually occurs when two atoms have similar or
the same electron affinity.
• Examples of gas molecules that have a
nonpolar covalent bond: Hydrogen gas atom,
Nitrogen gas atoms, etc.

24. EXAMPLES

25. Co-Ordinate Covalent Bonding
Presented by Farah Deeba.BY: FARAH DEEBA

26. CO-ORDINATE COVALENT BOND
”A covalent bond in which both electrons come
from the same atom. Also known as dative
covalent bonding.”

27. The atoms which donates the electron to form
it is called as the donor atom, while the atom
which accepts the pair of electron for bonding is
called as the acceptor atoms.
EXPLAINATION

28. EXAMPLES OF CO-ORDINATE BOND
• Lewis acid base reaction is an excellent example of the
co-ordinate covalent bond. For example the bond-
H3N: → BF3
-is a coordinate bond. Here nitrogen acts a donor atom.
The lone pair of electron in the nitrogen is donated to
the vacant p orbital of the boron. Here ammonia is
Lewis base and BF3 is Lewis acid.

29. EXAMPLES OF CO-ORDINATE
BOND
Boron hydride-ammonia complex.

30. Here, the nitrogen atom becomes the donor. The hydrogen
atom becomes the acceptor.
The linkage between N and H atoms is called coordinate
bond. It is represented by an arrow →.
AMMONIUM ION CO-ORDINATE BOND:

31. PROPERTIES
• Has all the characteristic of the covalent bond.
• Have low boiling and melting point.
• There are no columbic forces of attraction.
• Does not conduct electricity in the liquid or in
the dissolved state.
• Compounds are that much soluble in water.
• As strong as other covalent bonds.

32. Metallic bond
BY: UFAQ HAMID

33. METALLIC BONDING
Metallic bonding is the electrostatic attraction between the
positively charged atomic nuclei of metal atoms and the delocalized
electrons in the metal.
PROPERTIES OF METALS
1.Metals are shiny.
2.Metals conduct electricity because electrons are free to move.
3.Metals conduct heat because the positive nuclei are packed
closely together and can easily transfer the heat.
4.Metals have a high melting point because the bonds are strong
and a high density because of the tight packing of the nuclei.

34. The outermost electrons for most metals are only loosely bound to their nuclei
because of their relative remoteness from their positively charged cores. All valence
electrons of a given metal combine to form a "sea" of electrons that move freely
between the atom cores. The positively charged cores are held together by these
negatively charged electrons. In other words, the free electrons act as the bond
between the positively charged ions. Metallic bonds are non directional. As a
consequence, the bonds do not break when a metal is deformed. This is one of the
reasons for the high ductility of metals.
The schematic representation of metallic bonding is show below. The valence
electrons become dissociated with their atomic core and form an electron "sea" that
acts as the binding medium between the positively charged ions.
Examples for materials having metallic bonds are most metals such as Cu, Al, Au, Ag
etc. Transition metals (Fe, Ni etc.) form mixed bonds that are comprised of covalent
bonds (involving their 3d-electrons) and metallic bonds. This is one of the reasons why
they are less ductile than Cu, Ag and Au.
METALLIC BOND EXAMPLES

35. Ionic Bond Covalent Bond Metallic Bond
The transfer of electrons
between two atoms
having different electro
negativities forms this
bond.
This bond is formed by the
mutual sharing of
electrons between same
or different elements .
This bond is formed due to
the attraction between
kernels and the mobile
electrons in a metal
lattice.
This is a strong bond due
to electrostatic force of
attraction.
This is also a fairly strong
bond because the electron
pair is strongly attracted
by two nuclei.
This is a weak bond due to
the simultaneous
attraction of the electrons
by a large number of
kernels
This is a non-directional
bond.
This is a directional bond.
This is a non-directional
bond.
This bond makes
substances hard and
brittle.
This bond makes
substances hard and
incompressible.
This bond make
substances malleable
and ductile.
COMPARISON OF IONIC BOND COVALENT BOND
AND METALLIC BOND

36. BY: SHABANA RAEES

37. INTRA MOLECULAR FORCES
The intramolecular force is the sum of all the
forces holding a molecule or compound together.
These forces are stronger than intermolecular
forces. Intramolecular forces are categorized into
three types.
1. Metallic bonding
2. Ionic bonding
3. Covalent bonding

38. METALLIC BONDING:
This type of bonding is formed when metal
atoms share delocalized electrons.
IONIC BONDING:
This type of bonding is formed when metal atom
donate one or more electrons to non metal
atoms. An electrostatic attraction builds
between the two ions, which is known as ionic
bonding.
COVALENT BONDING:
Covalent bonding involves sharing a pair of
electrons. By sharing a pair of electrons both
atoms gain a stable electron configuration state.

39. WHY IS INTRAMOLECULAR FORCES
GREATER THAN INTERMOLECULAR
FORCES?
Intramolecular forces generally involve covalent bonds, wherein
electrons are shared, and atoms are relatively close to each other.
Proximity of the atoms and the electron sharing make the forces
strong.
Intermolecular forces generally involve weak electrostatic
interaction, and atoms are farther from each other, rendering the
forces weaker than the intramolecular variety.
In general, the force between two objects is proportional to the
inverse of the square of the distance between them, or Force ~=
1/(distance*distance). The smaller the distance between the
objects, the GREATER the force between them.

40. WHAT IS THE DIFFERENCE BETWEEN INTERMOLECULAR
AND INTRA-MOLECULAR FORCES?
• Intermolecular forces are formed between molecules and, intra-
molecular forces are formed within the molecule.
• Intra-molecular forces are much stronger compared to
intermolecular forces.
• Covalent, ionic, and metallic bonding’s are types of intra-molecular
forces. Dipole-dipole, dipole-induced dipole, dispersion forces,
hydrogen bonding are some of the examples for intermolecular
forces.

41. BY: AREEBA SALEEM

42. The set of attractive and repulsive forces that
occur between the molecules as a result of the
polarity of the molecules.
INTERMOLECULAR FORCES

43. When two or more atoms are joined by chemical
bonds they form a molecule, electrons travel up to
the new molecule and are concentrated in the most
electronegativity atom area, the electronegativity is
defined as property that have the atoms or
molecules to attract electrons. The concentration of
electrons in a defined area of the molecule creates a
negative charge, while the absence of electrons
creates a positive charge.

44. Intermolecular forces acting between the
molecules are classified as:
•Permanent dipoles
•Induced dipoles
•Dispersed dipole.
•Hydrogen bonds
Within the 4 groups described above, the most
important forces are the top 3, also known as
Van der Waals forces.
CLASSIFICATION OF
INTERMOLECULAR FORCES

45. PERMANENT DIPOLES
This type of bonding occurs when two molecules
have positive and negative charges, they are polar
molecules that have polarity, electrostatically
attracting and forming the bond.

46. INDUCED DIPOLES
This type of bonding occurs when a non polar
molecule redistributes the concentration of
electrons (it has the ability to polarize) when it
approach of a polar molecule, so that it creates
a bond between two molecules.
In this case the polar molecule induces the
creation of the non-polar molecule in a polar
molecule.

47. DISPERSED DIPOLE
This latter case the binding occurs between polar
molecules but can be polarized, and when the
latter occurs attract each other creating a
molecular bond.
The binding energies generated by the
intermolecular forces are smaller than the
energies generated in the chemical bonds, but
exist in greater numbers compare with the number
of chemical bonds.

48. Intermolecular Forces: Chemical bonds:
•Intermolecular forces depend
on the temperature
•They are weaker than chemical
bonds, order of 100 times lower
•The bond distance is at the
level of microns
•Unions are not directed.
•Chemical bonds do not
depend so much about the
temperature.
They are stronger than
intermolecular forces
•The bond distance is very
small, in terms of Angstroms
•Unions are directed
DIFFERENCE BETWEEN INTERMOLECULAR
FORCES AND CHEMICAL BONDS

49. HYDROGEN BONDING
PRESENTED BY: ZAINAB SOHAIL

50. A Hydrogen Bonding is a dipole-dipole
attractive force that exist between two polar
molecules, containing a hydrogen atom
covalently bonded to an atom of F, O or N.
This attraction between positive hydrogen
and negative oxygen or flourine, is called
Hydrogen Bond.
HYDROGEN BONDING

51. • Hydrogen has no inner shell electron and is
very small in size, the positive charge density
developed is high
• The nucleus of hydrogen atom is exposed to
attraction by nearby electron cloud, a lone
pair electrons on the electronegative atom

52. EXAMPLES OF HYDROGEN
BONDING
In Proteins:

53. EFFECTS OF HYDROGEN
BONDING
On Boiling point:
bp +78 °C bp -45°C bp -25°C

54. On Water-solubility:
.………O ̶ H………O ̶ H………O ̶ H……………
ǀ ǀ ↙ ǀ
H H R-H H

55. BY: AISHA AMJAD

56. DEFINITION:
A dipole moment ​is a measurement of
the separation of two oppositely
charged .
The magnitude is equal to the charge
multiplied by the distance between the
charges and the direction is from
negative charge to positive charge.
Dipole moments are a vector quantity.
Dipole moments are measured in the
SI units Debye which is equal to the of
coulomb meters (C m).
(1 Debye equals 3.34 x 10-30 coulomb-
meters)

57. DIPOLE-DIPOLE ATTRACTION
• Attractive forces that exist
between molecules that
have permanent dipoles.
• These exist in any polar
substance.
• Weaker than Ion-Dipole
force
• Increased polarity, stronger
dipole-dipole attraction
H Cl
+ -
H Cl
+ -
dipole
dipole
Attractive force

58. The dipole–dipole interactions in (a) crystalline
CH3CN and (b) liquid CH3CN.
Dipole moments arise from differences in
electronegativity. The larger the difference in
electronegativity, the larger the dipole moment.
EXAMPLE

59. POLARITY OF THE DIPOLE
MOLECULE:
The dipole moment is a measure of the polarity of
the molecule.
POLAR MOLECULE.
An example of polar molecule is water ,ammonia
etc.
For molecules of
approximately equal mass
and size, the strength of
intermolecular attractions
increases with increasing
polarity.
NON POLAR MOLECULE.
An example of non polar molecule is carbon dioxide
and carbon tetrachloride etc.

60. VAN DER WAALS FORCES
BY: HAZIQA IFFRIN

61. • Vander Waals forces are also known as London forces.
• They are weak interactions caused by momentary changes in
electron density in a molecule.
• They are the only attractive forces present in nonpolar compounds.
Even though CH4 has no
net dipole, at any one
instant its electron density
may not be completely
symmetrical, resulting in a
temporary dipole. This can
induce a temporary dipole
in another molecule. The
weak interaction of these
temporary dipoles
constitutes van der Waals
forces.

62. • All compounds exhibit van der Waals forces.
• The surface area of a molecule determines the strength
of the van der Waals interactions between molecules.
The larger the surface area, the larger the attractive
force between two molecules, and the stronger the
intermolecular forces.
Surface area and van der Waals forces

63. • Van der Waals forces are also affected by polarizability.
• Polarizability is a measure of how the electron cloud
around an atom responds to changes in its electronic
environment.
Larger atoms, like iodine,
which have more loosely
held valence electrons,
are more polarizable
than smaller atoms like
fluorine, which have
more tightly held
electrons. Thus, two F2
molecules have little
attractive force between
them since the electrons
are tightly held and
temporary dipoles are
difficult to induce.

64. BY: ASRA FATIMA

65. APPLICATIONS OF CHEMICAL
BONDING
•Today, chemical bonding is understood as the
joining of atoms through electromagnetic force.
• Not all chemical bonds are created equal: some
are weak, and some very strong, a difference that
depends primarily on the interactions of electrons
between atoms.

66. • Almost everything a person sees or touches in
daily life—the air we breathe, the food we
eat, the clothes we wear, and so on—is the
result of a chemical bond, or, more accurately,
many chemical bonds.
• The strength of chemical bonds varies
considerably; there are "strong bonds" such
as covalent or ionic bonds and "weak bonds"
such as Dipole-dipole interaction, the London
dispersion force and hydrogen bonding.