it is related to structure of atoms.

1. Atomic Structure
By
Dr. Parul Jain
Lecturer, Engineering Chemistry
S. V. Polytechnic College, Bhopal

2. Objects
• Discovery of electron
• Discovery of Proton
• Discovery of neutron
• Rutherford’s Nuclear Model
• Bohr’s Model
• Orbital and shape of S,P, d orbital
• Bohr-Burry scheme
• Radioactivity

3. Dalton's Atomic Theory
• All matter is composed of minute indivisible
particles called atoms.
• Atoms cannot be created or destroyed.
• The indivisibility of an atom was later found
to be incorrect. Experiments showed that the
atom was composed of several smaller
particles like electron, proton , neutron etc.

4. Discovery of electron
• Electron was discovered by J. J. Thomson in 1897 when he
was studying the properties of cathode ray.
• Thomson won Nobel Prize in 1906 for discovering the
electron.
• In the experiment, J. J. Thomson had taken a discharge
tube. He applied a high electrical voltage around 10,000
volts between two electrodes and air at low pressure
around 0.001 mm Hg. He observed a glow behind the
anode . It means some rays was coming out from the
cathode to anode. These rays are called cathode rays .

5. Properties of cathode rays
1. They travel in straight lines.
2. They have mass and velocity means kinetic energy.
2.They deflect in magnetic field as well as in electric field
means they are consist of charged particles.
3. Applying electric field in the path of cathode ray deflects
the ray towards positively charged plate. Hence cathode ray
consists of negatively charged particles called electron.
4.They are independent of the material composition of the
cathode.
Mass=9.1 x 10-28
g,Charge= -1.6 x 10-19
coulomb
A diagram of J.J. Thomson's cathode ray tube. Image from Openstax, CC BY 4.0.

6. Discovery of Proton
In 1886, Goldstein had done experiments to prove the
existence of anode rays. In the experiment, he had taken a
discharge tube with a perforated cathode and air at low
pressure around 0.001 mm Hg. At high voltage around 10,000
volts, he observed a faint red glow behind the cathode.

7. Properties of Anode rays
1. They travel in straight lines.
2. They have mass and velocity means kinetic energy.
3.They deflect in magnetic field as well as in electric
field means they are consist of charged particles.
4. Applying electric field in the path of anode ray
deflects the ray towards negatively charged plate.
Hence anode ray consists of positively charged
particles called proton.
4.They are dependent of the material composition of
the cathode.

8. Proton (P)
• Mass=1.67 x 10-24 g
• Charge=+1.6 x 10-19 coulomb
Atomic number = no. of protons
= no. of electrons
• But scientists soon realized that mass of the
nucleus is approximately twice than the
number of proton.

9. Discovery of neutron
• In 1932, James Chadwick bombarded
beryllium with alpha particles and found an
electrically neutral particle. This new particle
had similar mass to proton and was charge
less. This particle was called neutron.
• 4Be9 + 2He4 (α) 6C12 + 0n1
• Mass=1.67 x 10-24 g
• Charge=0

10. Atomic Structure
What we have learned
Atoms are composed of
electrons
protons
neutrons
Atomic number = no. of protons
= no. of electrons
Atomic mass = no. of protons + no. of neutron
Protons and neutrons are located in the nucleus. Electrons
are found in orbitals surrounding the nucleus.

11. Subatomic Particles
Particle
Mass
(g)
Charge
(Coulombs)
Charge
(units)
Electron (e-
) 9.1 x 10-28
-1.6 x 10-19
-1
Proton (p) 1.673x10-24
+1.6 x 10-19
+1
Neutron (n) 1.675x10-24
0 0
mass p = mass n = 1840 x mass e-

12. Assignment
Q. Find out the no. of Electrons, Protons and
Neutrons in following atoms:
• 1H1 e-= 1, P=1, n=0
• 11Na23 e-= 11, P=11, n=12
• 15P31 e-= 15, P=15, n=16
• 17Cl35 e-= 17, P=17, n=18

13. Rutherford’s Gold Foil Experiment
• In 1911, Rutherford discovered nucleus in his famous gold foil
experiment.
• In this experiment, Rutherford bombarded a beam of alpha
particles on an ultrathin gold foil and then detected the
scattered alpha particles in zinc sulfide (ZnS) screen.

14. Observations
• Most of the particles go straight through the foil without any
deflection.
• Some of the alpha particles deflect at small angle.
• Very few even bounce back (1 or 2 in 20,000).

15. Conclusion
• Most of the part of an atom is empty.
• Atom consists of a small, positively charged part called
nucleus . This explained why a very small fraction of the
α particles were deflected.
• Nucleus of an atom is very small as compared to the total
size and contains most of the atom's mass that’s why very
few α particles were come back
• The number of negatively charged electrons dispersed
outside the nucleus is same as number of positively
charge in the nucleus. It explains the overall electrical
neutrality of an atom.

16. The Nuclear Model of the atom
Based on his experimental results, Rutherford made the
following assumptions about the structure of the atom:
• Most of the part of an atom is empty and it consists of a very
small, positively charged part called nucleus.
• The size of nucleus (10-13 cm )is very small in comparison to
the size of atom(10-8 cm) and it contains most of the atom's
mass. It means all the protons are situated in nucleus.
• Negatively charged electrons are revolved around the nucleus
in empty space of atom. Since atom is neutral, number of
electrons are equal to the number of protons.
Image of Rutherford atom from Wikimedia Commons, CC-
BY-SA-3.0.

17. Defects of Rutherford’s model
1 According to Maxwell’s electromagnetic theory, charge
particle when accelerated, must emit energy. So when
negatively charged electron accelerated around the nucleus, it
must radiate energy and due to continuous loss of energy
orbit of electron must decrease continuously. Consequently
electron will fall into the nucleus. This shows that atom is
unstable. But in actual practice Atom is most stable.
2. If the electrons emit energy continuously, they should form
continuous spectrum .But actually line spectrum is obtained.

18. Bohr’s Model
The Bohr model of an atom was based upon Planck's
quantum theory of radiation.
E = hν
where: h is Planck's constant ,h=6.626x 10-34 Js
ν is frequency
Postulates
• In an atom, electrons revolve around the nucleus in a
closed circular path called orbit or shell.
• Each orbit or shell has a fixed energy so it is also called
energy levels.
• The energy levels are represented by an integer n (n=1, 2,
3…) known as the principle quantum number. 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.

19. Bohr’s Model
• When electrons in an atom move from a lower energy level to
a higher energy level they require energy and when electron
moves from a higher energy level to lower energy level they
lose energy.
• How much energy they will absorb or emit is determined by
the energy difference of those energy levels. According to the
Planck’s relation:
Δ E = E 2 − E 1
= h ν2 -h ν1
=h (ν 2- ν1 )=hΔν
• The angular momentum L(m v r ), of the orbiting electron is
quantized such that
m v r = n h/2π where n = 1, 2, 3, ...

20. Merits and Demerits
Merits
• Bohr’s model can explain the stability of atom.
• It was successful to explain the spectrum of hydrogen atom
and other hydrogen-like atoms and ions.
Demerits
• It could not explain the spectra obtained from multi electrons
atoms.
• It could not explain fine structure and hyperfine structure of H
atom
• Bohr’s model of an atom failed to explain the Zeeman Effect
(effect of magnetic field on the spectra of atoms).
• It also failed to explain the Stark effect (effect of electric field
on the spectra of atoms).
• It violates the Heisenberg Uncertainty Principle.

21. ORBITAL
ORBITAL
It is a region in three dimensional space around the
nucleus where the probability of finding an electron
is maximum.
• It represents three dimensional motion of electron
around nucleus.
• The maximum no. of electrons in an orbital is 2.
• Orbitals have different shapes.

22. Difference between orbit and orbital
Orbit
• It is well-defined circular path
followed by revolving electron
around nucleus.
• It represents two dimensional
motion of electron around
nucleus.
• The maximum no. of electrons
in an orbit is 2n2.
• It is circular in shape.
• They are non-directional in
character
Orbital
• It is a region of space around
the nucleus where the
probability of finding an
electron is maximum.
• It represents three
dimensional motion of
electron around nucleus.
• The maximum no. of electrons
in an orbital is 2.
• Orbitals have different shapes.
• They are directional in
character

23. S orbital
• S orbital is spherically symmetric around the
nucleus of the atom.
• Non-directional

24. P-orbital
• Dumb-bell in shape
• Directional

25. d- orbital
• Double dumb-bell shape
• The five d orbitals are called
Non axial d orbitals-dxy, dyz, dzx
•

26. d- orbital
• Axial d orbitals- dx² - Y², dz²

27. Bohr- Burry Scheme
1. Maximum number of electrons that can be accommodated
in a shell is given by 2n2 ,where n = shell number
n = 1, Maximum number of electrons = 2 x (1) 2 = 2
n = 2, Maximum number of electrons = 2x(2) 2 =2 x 4 = 8
n = 3, Maximum number of electrons= 2x(3) 2= 2 x 9 = 18
n = 4, Maximum number of electrons = 2x(4) 2= 2x16 = 32
2. The maximum number of electrons that can be
accommodated in the outer most orbit is 8 and in
penultimate shell is 18.
3. A new orbit begins when an orbit contains 8 electrons.
4. The outermost orbit cannot have more than 2 electrons and
penultimate shell more than 8 electrons, unless the inner
shells are completely filled by 2n2 rule .

29. Assignment
• Q. Write the electronic configuration of
following atoms:
• 16S and 53I
• 16S- 2,8,6
• 53I- 2,8,18,18,7

30. Thanks