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1. Chapter#3
The Atomic Structure

2. Discovery of Fundamental Particles of
Atom
• Proton
• Electron
• Neutron

3. The Greek Period
In 2000 BC, The Greek Philosopher ‘Democritos’
gave a philosophy that all the substances are
composed of very small particles called
‘Atomos’ means indivisible in Greek.
The Greek rejected this concept and continued
to believe in Aristotle’s concept of 4 basic
elements i.e; fire, earth, water and earth.

4. Dalton’s Atomic Theory
John Dalton was an English scientist who proposed
this theory in 1808.
Main Points of Dalton’s Atomic Theory:
i) Matter is composed of indivisible particles
called atom.
ii) Identical atoms are known as elements.
iii) Elements are combined chemically to form
compound.
iv) Chemical reaction is only the rearrangement of
atoms and atoms themselves are not changed.

5. Discovery of Cathode Rays/Electrons
Discharge Tube Experiment
• Heinrich Giessler was a German Scientist who
invented the discharge tube to pass electric
current through air at low pressure.
• W. Crookes was the scientist who discovered
cathode rays through Discharge tube.
• J. J. Thomson discovered the properties of
cathode rays and concluded that these rays
are actually material particles.
• Stoney named them Electrons.

7. Properties of Cathode Rays
• Cathode rays travel in straight line so they cast sharp shadows of objects.
• These rays evolve from cathode and moves to anode.
• They can pass through thin foil.
• They carry negative charge.
• They can rotate a light paddle wheel hence they can exert mechanical
pressure and they have kinetic energy.
• Cathode rays consist of material particles known as electrons.
• When cathode rays strike with a meta they produce X-rays.
• Cathode rays does not depend upon the gas filled in the discharge tube.
• The e/m ratio of cathode rays remain same for all the gases.

8. Discovery of Protons
Discharge Tube Experiment
Protons/Anode Rays/Positive Rays/Canal Rays
Were discovered by Goldstein by using a
discharge tube with perforated cathode.

9. Properties of Anode Rays
• Anode rays travel in straight line so they cast sharp
shadows of objects.
• These rays evolve from anode and moves to cathode.
• They cannot pass through thin foil.
• They carry positive charge.
• Anode rays consist of material particles known as
protons and it is 1836 times heavier than electrons.
• Anode rays depend upon the gas filled in the discharge
tube.
• The e/m ratio of cathode rays is different for all the
gases.
So, Heavier atoms have lower e/m values.

10. NOTE
• After the discovery of electron and proton,
scientists gave their atomic models before the
discovery of neutron.
• But for the sake of rhythm we will study about
the discovery of radioactivity and neutron.

11. Radioactivity
The spontaneous disintegration of the unstable nucleus of an
atom in the form of invisible radiation is called radioactivity.
There are two types of radioactivity:
i) Natural radioactivity (Elements after Pb 82)
Discovered by Henry Bacquerel.
i) Artifical Radioactivity (Bombardment of particles)
Discovered by James Chadwick for discovery of neutron.
Radioactivity is a nuclear reaction:
92U238  90Th234 + 2He4

12. • Henry Bacquerel discovered the phenomena of
radioactivity by the emission of beta particles
through pitch blende.
• Madam Curie & Perrie Curie discovered that the
emission of particles is due to Radium.
• Rutherford discovered that there are three types
of radioactive rays.
• James Chadwick discovered artificial radioactivity
through which he discovered neutron.

13. Types of Radioactive Rays

14. Properties of Radioactive Rays
Property α - Particles  - Particles  - rays
Nature Helium nucleus electrons Electromagnetic
radiations
Mass 4 amu 9.11x10-31 kg 0
Charge Positive Negative No charge
Velocity < light ≈ light = light
Penetration Power Least (1 – 2 cm in
air)
More than alpha
(1 – 2 m in air)
Most
(15 – 20 cm in lead)
Ionisation Power highest medium least

15. Radioactive Decay

16. Discovery of Neutron
Artificial Radioactivity
James Chadwich (1934)
4Be9 + 2He4  6C12 + 0n1
neutrons can penetrate more than gamma rays.

17. Summary
Particle Discoverd by Method Mass in kg Charge in C
Electron W. Crookes Discharge Tube 9.11x10-31 kg -1.6x10-19 C
Proton Goldstein Discharge Tube 1.67x10-27 kg 1.6x10-19 C
Neutron James
Chadwick
Aritificial
Radioactivity
1.76x10-27 kg 0

18. Thomson’s Plum Pudding Model

19. Rutherford’s Atomic Model
(Nuclear Model of Atom)
Rutherford bombarded alpha particles from
polonium metal into the thin gold foil.
Gold was selected because of bigger atoms and
extremely malleable.
A photographic plate was also present behind
the gold foil.
Through this experiment he concluded that:

21. Defects in Rutherford’s Atomic Model
• If electron is emitting energy continuously
then it must stop afterwards and falls into the
nucleus making a spiral path.
• If energy is emitted by atom in the continuous
manner then we should get a continuous
spectrum but in contrast to it we get a line
spectra.

22. Spectra (band of light)
Plural (spectrum)
The dispersion of light when it is passed through
a prism is called spectrum.
There are two types of spectra:
1) Line Spectra
2) Continuous Spectra

23. Continuous Specta
(Formed by white light)
Red is the least deviated light having the highest wavelength (7000 Å), lowest frequency and lowest energy.
Violet is the most deviated light having the lowest wavelength (4000 Å), highest frequency and highest energy.

24. Line Spectra / Atomic Spectra
(Emitted by atoms in discharge tube)

25. Planck’s Quantum Theory
Proposed by Max Planck
Quanta = packets of energy
• Energy is released from atom in discontinuous
manner in the form of packets of energy
called quanta. Quanta of light is called photon.
• The energy of emitted light is directly
proportional to its frequency.
E=hυ
Here h=Planck’s constant (6.625x10-34 JS)
(6.625x10-27 erg. S)
This theory proves that energy is quantized.

26. Bohr’s Atomic Model
(Solar System Model of Atom)
• Based on Max Planck’s Quantum Theory.
• Neils Bohr removed the defects of Rutherfor’s atomic model.
Main Points:
i) As soon as atom revolves in its specified energy level it neither
radiates not absorbs energy.
ii) When it jumps from lower energy level to higher energy level it
absorbs energy and when it jumps back to lower energy level it
absorbs energy.
iii) The energy is absorbed and released in the form of quanta.
iv) The change in energy can be calculated by:
ΔE=E2-E1=hυ
v) Only those orbits are possible where the angular momentum of
electron is the integral multiple of h/2π
mvr = nh/2 π

28. Radius of Hydrogen Atom
By Neils Bohr

4
2
2
2
k
mZe
h
n
r 
r1=ao=0.529 Å
ao is known as Bohr’s radius for hydrogen.
rn=n2 x ao

29. Energy of an electron
1
2
3
.
1313 

 kJmol
n
En

30. Wave number (υ)
The number of waves per unit length is called wave
number.









 2
2
2
1
1
1
n
n
RH

Here, RH is known as Rydberg constant for hydrogen. The value of Rydberg constant
is 109678 cm-1
n1 is the lower orbit.
n2 is the higher orbit.

31. Hydrogen Spectrum
• The following series of spectral lines are obtained by
hydrogen gas.
Lyman Series (when electron jumps to 1st orbit)
UV region
Balmer Series (when electron jumps to 2nd orbit)
Visible region
Paschen Series (when electron jumps to 3rdorbit)
Infrared region
Bracket Series (when electron jumps to 4th orbit)
Far infrared region
Brackett Series (when electron jumps to 5thorbit)
Far infrared region

33. Defects in Bohr’s Atomic Model
• Applicable only for the atoms with an electron
in the last shell.

34. X-Ray
Roentgen Rays

35. X-Ray and Atomic Number
• Atomic Number discovered by Moseley.

36. Heisenberg’s Uncertainty Principle
• Given by Werner Heisenberg in 1925.
• It rejected Bohr’s circular orbit concept.
“It is impossible to determine the exact position
and momentum of an electron simultaneously
in an atom.”
mvr≈nh/2π

38. Quantum Mechanical Model
Wave Mechanical Model
(Ervine Schrodinger)

39. Orbit/Shell/Energy Level/Stationary State Orbital
(i) The fixed path in which electron
revolves.
(i) The region around nucleus where the
possibility of finding an electron is
maximum.
(ii) It is circular in shape. (ii) It exists in different shapes.
(iii) It is discovered by Bohr. (iii) It is discovered by Schrodinger.
(iv) An orbit contains 2 or more electrons. (iv) An orbital contains maximum two
electrons.
(v) The number of electrons can be
calculated by 2n2
(vi) The number of electrons can be
calculated by 2(2l+1)

40. Types of Orbitals
• There are 4 types of orbitals.
(i) s – orbital (s=sharp lines)
(ii) p – orbital (p=principal lines)
(iii) d – orbital (d=diffused lines)
(iv) f – orbital (fundamental lines)

41. s – sub shell
s - orbital
• It is spherical in shape.
• It can accommodate maximum 2 electrons.
• There is no nodal plane in s-orbital.
• It is non-directional.

42. p – subshell
contains px py pz orbitals
• It is dumb-bell in shape.
• It contains 6 electrons in 3 different orbitals
having different orientations.
• It is directional. The electron charge density is
maximum at the ends.
• The nodal plane lies between the point of
intersection of two nodes.

44. d – sub shell
dxy dyz dzx dx2-y2 dz2 orbitals
• It is double dumb-bell in shape.
• There are 5 orbitals in d – subshell.
• d – sub shell contains max 10 electrons.

45. f – sub shell
• They have complex saucer like shape.
• They contain maximum 14 electrons in 7
orientations.

46. Quantum Numbers
• It tells about the address of electron in an
atom.
• There are four types of quantum numbers.
(i) Principal Quantum number (n)  orbit
(ii) Azimuthal / subsidiary quantum number (l)  sub-shell
(iii) Magnetic quantum number (m)  orbital
(iii) Spin quantum number (s)  electron

47. Principal Quantum Number
• It tells about the number of orbit or the size of
orbital.
• It is denoted by ‘n’.
• Its value is in natural number.
• n=1 for K-shell, n=2 for L-shell etc.

49. Azimuthal Quantum Number
• It tells about the shape of orbital.
• Its value is in whole number (0, 1, 2, 3)
• The value can be calculated by the formula
l=0 to n-1

50. Magnetic Quantum Number
• It tells about the orientation of orbital.
• Its value is in integer.
• For any value of l we have 2l+1 values of m.
• For any value of l:
m=-l  0  +l

51. Spin Quantum Number
• It tells about the spin of electron in its own
axis.
• Its value is either +1/2 or -1/2

52. Electronic Configuration

53. Rules of Electronic Configuration
• There are four rules of electronic
configuration.
(i) Pauli’s exclusion principle
(ii) n+l rule
(iii) Aufbau Principle
(iv) Hund’s rule of maximum multiplicity

54. Definitions
• Atomic Radius
• Ionization Potential
• Electron Affinity
• Electronegativity