1) The presentation summarizes atomic structure and chemical bonding models including fundamental particles, Rutherford and Bohr atomic models, quantum numbers, and types of chemical bonds. 2) It discusses limitations of earlier atomic models and introduces concepts like electron energy levels, photon energy, and spectral lines. 3) Various types of chemical bonds are defined including ionic, covalent, and metallic bonds. Molecular orbital theory of covalent bonding is also introduced.

1. Welcome
To
Our Presentation

2. Presented to
Our Respectable Teacher
Md Monir Hossain
Lecturer
Department of Pharmacy
Mawlana Bhashani Science and Technology University
Tangail 1902, Bangladesh

3. Course Title : Inorganic Pharmacy I
Course Code : PHAR-1103
Presentation on:
Atomic Structure & Chemical Bond

4. Presented by – Group -3
1.Toma Khanam (PHA-19015)
2.Md Mehedi Hasan Sagor (PHA-19016)
3.Mim Akter Brishti (PHA-19020)
4.Sabbir Ahmed (PHA-19021)
5.Md Hafizur Rahman (PHA-19022)

5. Contents :
 Fundamental particles
 Theories of atomic structure
 Rutherford & Bohr atom model & its
limitation
 Quantum numbers
 Pauli’s exclusion principle
 Origin of spectral line
 Electronic concept of valency
 Different types of chemical bond
 Molecular orbital theory of co-valent
bonding

6. Fundamental Particle:
The fundamental particle represent the matter that
can not be subdivided into smaller and simpler
particles.
Scientists thought that they had found finally
fundamental particles when John Dalton discovered
atom in 1803 and he declared that atom could not be
divided into smaller , simpler particle .For almost 100
years after Dalton discovered atoms they were
accepted as the fundamental particles of matter

7. But starting the late 1890s with the discovery of
electrons particles smaller and simpler than
atoms where identified. Within a few decades
,protons and Neutrons were discovered .
Ultimately ,hundreds of sub atomic particles were
found .
The key characteristic of fundamental particle is
that they have no internal structure.

9. Types of fundamental particles :
There are two types of fundamental particles
1. Matter particles some of which combine to
produce the word about us.
2. Force particles one of which the photon is
responsible for electromagnetic radiation.

10. These are classified in the standard model of
particle physics, which theroses how the basic
building of matter interact ,governed by
fundamental forces .
Matter particles are fermions, while force
particles are boson.

11. Theory of Atomic Structure
• John Dalton Model
• Thomson's Atomic Model
• Rutherford's Nuclear Atomic Model
• Bohr's Atomic Model
Many atomic model was invented in 18th century. Those are :-
Here, we will describe about last two models

12. Rutherford Atomic Model
Ernest
Rutherford
The Rutherford model, also
known as planetary model is a
model which tried to describe
an atom devised by Ernest
Rutherford.

13. 1. Atom contains a heavy and positively
charged part at its center. This central
part of the atom is called nucleus.
Rutherford proposed, in 1911, what is now
called the Rutherford model of the atom.
He put forward these postulates that sum up
most of the model:

14. 2. The volume occupied by the nucleus is only a
minute fraction of the total volume of the atom, i.e.
the size of the nucleus is very small as compared to
that of the whole atom.
Although the nucleus is small in size, it is heavy
due to the presence of all protons and neutrons
in it. The mass of the electrons is negligible.
The nucleus has positive charge.

15. 3. Rutherford's model of atom bears a close
resemblance with the solar system in which the
massive sun plays the role of the massive nucleus and
the planets play the role of the revolving electrons.

16. Limitation of Rutherford atomic model:
Rutherford model was unable to explain the
stability of an atom. According to Rutherford
postulates, electrons revolve at a very high speed
around a nucleus of an atom in a fixed orbit.
However ,Maxwell explained accelerated charged
particles release electromagnetic radiation .
Therefore ,electrons revolving around the nucleus
will release electromagnetic radiation

17. • The electromagnetic radiation will have
energy from the electronic motion as a
result of which the orbits will gradually
shrink . Finally the orbits will shrink and
collapse in the nucleus and atom.
According to the calculations , if Maxwell’s
explanation is followed Rutherford’s model
will collapse with 10-8 seconds
• Therefore ,Rutherford atomic model was
not following Maxwell’s theory and it was
unable to explain an atoms stability

18. 3. Rutherford’s theory was incomplete
because it did not mention anything
about the arrangement of electrons in
the orbital. This was one of the major
drawbacks of Rutherford atomic
model.

19. Bohrs Atomic model
In atomic physics, Bohr
model or Bohr diagram,
presented by Niels Bohr in 1913,
is a system consisting of a small,
dense nucleus surrounded by
revolving electrons —similar to
the structure of the Solar System.
Niels Bohr

20. Niels Bohr proposed, in 1913, what is now called
the Bohr model of the atom. He put forward these
three postulates that sum up most of the model:
1. The energy levels are represented by an integer
(n=1, 2, 3…) known as the quantum number. This
range of quantum number starts from nucleus side
with n=1 having the lowest energy level. The orbits
n=1, 2, 3, 4… are assigned as K, L, M, N…. shells
and when an electron attains the lowest energy
level it is said to be in the ground state.

21. 2. The stationary orbits are attained at
distances for which the angular
momentum of the revolving electron is
an integral multiple of the
reduced Planck's constant: mevr=nh/2π
Bohr’s Model of an Atom
where n = 1, 2, 3, ... is called
the principal quantum number,
and ħ = h/2π. The lowest value of n is 1;
this gives a smallest possible orbital
radius of 0.0529 nm known as the Bohr
radius. Once an electron is in this lowest
orbit, it can get no closer to the proton.

22. 3. Electrons can only gain and lose
energy by jumping from one allowed
orbit to another, absorbing or
emitting electromagnetic radiation
with a frequency ν determined by
the energy difference of the levels
according to the Planck relation:
ΔE=E2-E1 =hv
where h is Planck's constant.

23. Limitations of Bohr’s Model of an Atom
1. Bohr’s model of an atom failed to explain the
Zeeman Effect (effect of magnetic field on
the spectra of atoms).
2. It also failed to explain the Stark effect
(effect of electric field on the spectra of
atoms).
3. It violates the Heisenberg Uncertainty
Principle.
4. It could not explain the spectra obtained
from larger atoms.

24. Quantum numbers
 Quantum numbers are a set of values that
describes the state of an electron including its
distance from the nucleus, the orientation and
type of orbital where it is likely to be found, and
its spin.
 What is Quantum Numbers ?

25. The Four Electronic Quantum Numbers
 Four quantum numbers can describe an electron
in an atom completely. These quantum numbers
are:
1. Principal quantum number (n)
2. Azimuthal quantum number (ℓ)
3. Magnetic quantum number (m)
4. Spin quantum number (s)

26. 1. The principal quantum number (n) describes
the electron shell, or energy level, of an electron.
The value of n ranges from 1 to the shell
containing the outermost electron of that atom,
that is
N=1,2,3…..
For example, in caesium (Cs), the
outermost valence electron is in the shell
with energy level 6, so an electron in
caesium can have an n value from 1 to 6

27. 2. The azimuthal quantum number (ℓ) (also
known as the angular quantum
number or orbital quantum number)
describes the subshell, and gives the
magnitude of the orbital angular
momentum through the relation
L2=ħ2 ℓ (ℓ + 1)
3. The magnetic quantum number (mℓ) describes
th specific orbital (or "cloud") within that
subshell, and yields the projection of the
orbital angular momentum along a specified
axis:
Lz = mℓ ħ

28. 4. The spin projection quantum
number (ms) describes the spin
(intrinsic angular momentum) of the electron
within that orbital, and gives the projection of
the spin angular momentum S along the
specified axis

29. Pauli exclusion principle
In 1925 Wolfgang Pauli put forward a principle
which controls the assignment of values to the
quantum numbers of an electron in an orbital.
Pauli exclusion principle can be stated as
“It is impossible for two electrons residing in the
same orbital of a given poly-electron atom (same
atom) to have the same values of all the four
quantum numbers.”

30. Illustration of the principle
In order to illustrate the principle let us consider
helium atom which has two electrons (Atomic
number = 2) in is orbital for which a = 1, 1 = 0 and in
= 0. The four quantum numbers for the two electrons
in is orbital are as follows

31. Photon energy
 Photon energy is the energy carried by a single
photon.
 light comes in discrete packets called photons
and the energy of each photon is set by its
color or wavelength
 From Einstein, we known that the photon
energy is inversely proportional to its
wavelength

32. Electron orbits around nucleus
img

34. From quantum mechanics, only certain orbits are allowed. Each orbits has
a specific energy.

35. How atoms emit light
 The emitted photon has an energy which is exactly
the energy difference between the orbits that the
electron had before and after.
 Because only certain energies are allowed for the
electron orbits, only certain energies of photons
can be produced. We call these the spectral lines
of hydrogen.

37. Spectral lines
A spectral line is a dark or bright line in an otherwise
uniform and continuous spectrum, resulting from
emission or absorption of light in a narrow frequency
range, compared with the nearby
frequencies. Spectral lines are often used to identify
atoms and molecules.
Each element (hydrogen, helium, neon, mercury, iron,
…) has its own particular set of energy levels and its
own set of spectral lines.

38. There are two types of spectral lines in the
visible part of the electromagnetic spectrum:
1. Emission lines – these appear as discrete
coloured lines, often on a black background,
and correspond to specific wavelengths of
light emitted by an object.
Types of Spectral lines

39. 2. Absorption lines – these appear as dark bands,
often superimposed on a coloured continuum, and are
the result of specific wavelengths being absorbed
along the line-of-sight

40. Spectral lines of
hydrogen
The length of each
arrow determines the
energy and therefore
the wavelength of the
photon emitted

41. Uses of spectral lines
Because each element has it own unique
pattern of spectral lines, the spectral lines from
stars can be used to determine the
composition.
The spectral lines can be used to determine
the relative number of atoms of each
elements.

42. Chemical Bond
What is Chemical Bond ?
A chemical is defined as the
attractive force that holds two or
more atoms together in a molecule
or an ion

43. Different types of chemical
bond
We have the following four types of bonds which
hold the atoms together in a molecules
Weak bondsStrong bonds
Ionic
Bonds
Co-valent
Bond
Co-ordinate
Bond
Metalic bond
Hydrogen
bond
Van-der waals bond

44. Ionic bond
What is Ionic Bond ?
The chemical bond formed between two atoms
by the 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.

45. Example of ionic compound :
1. MgO molecules.
2. CaCl2 molecules

46. Covalent Bond
The chemical bond between two atoms in which the electrons (in
pairs) are shared by both the participating atoms is called covalent
bond.

47. Examples of Co-valent bond
• H2 molecule is composed of two H atoms, each having one
valence electron. Each contributes an electron to the
shared pair and both atoms acquire stable helium
configuration
• In HF molecule H atom attains a doublet while F atom
achieves an octet of electrons.

48. In chemistry, Molecular orbital (MO) theory is a
method for describing the electronic structure of
molecules using quantum mechanics . Electrons are
not assigned to individual bonds between atoms, but
are treated as moving under the influence of
the nuclei in the whole molecule.
Molecular Orbital Theory

49. Valence bond theory :
According to this theory a covalent bond is formed
between the two atoms by the overlap of half filled
valence atomic orbitals of each atom containing
one unpaired electron. Valence bond theory
considers that the overlapping atomic orbitals of
the participating atoms form a chemical bond.
Because of the overlapping, it is
most probable that electrons should be in the bond
region

50. The two types of overlapping orbitals are sigma and
pi. Sigma bonds occur when the orbitals of two shared
electrons overlap head-to-head. Pi bonds occur when two
orbitals overlap when they are parallel.

51. A coordinate covalent bond, also known as a dative
bond or coordinate bond is a kind of 2-center, 2-
electron covalent bond in which the
two electrons derive from the same atom. The
bonding of metal ions to ligands involves this kind of
interaction. This type of interaction is central
to Lewis acid-base theory.
Co-ordinate Covalent bond
What is Co-ordinate Co-valent Bond ?

52. Example of Co-ordinate Co-valent bond

53. Metallic bond
Metallic bonding is a type of chemical bonding that
rises from the electrostatic attractive force
between conduction electrons (in the form of an
electron cloud of delocalized electrons) and
positively charged metal ions. It may be described
as the sharing of free electrons among
a structure of positively charged ions (cations).

54. Example of Metalic
Bond :

55. Dipole-dipole forces are attractive forces
between the positive end of one polar
molecule and the negative end of another
polar molecule.
Dipole-dipole force
What is Dipole-dipole force ?

56. The figures show two arrangements of polar iodine
monochloride (ICl) molecules that give rise to dipole-
dipole attractions.

57. Dipole-dipole forces have strengths that
range from 5 kJ to 20 kJ per mole. They are
much weaker than ionic or covalent bonds
and have a significant effect only when the
molecules involved are close together
(touching or almost touching).
Properties of Dipole-dipole forces

58. Reference :
1. Modern Inorganic chemistry – RD Madan
2. Introduction to Inorganic Pharmacy -S.Z Haidar
3. https://www.britannica.com/
4. http://Wikipedia.org
5. http://khanacademy.org

59. Thank
You