1. Atoms are the basic building blocks of matter and consist of a small, dense nucleus surrounded by electrons. 2. Rutherford's gold foil experiment in 1911 showed that the atom has a small, dense nucleus containing positively charged protons and uncharged neutrons. 3. Niels Bohr proposed his model of the atom in 1913 in which electrons orbit the nucleus in fixed shells at specific energy levels, explaining atomic spectra.

1. ATOMIC STRUCTURE
Vinay Desai
M.Sc Radiation Physics
KIDWAI MEMORIAL INSTITUTE OF ONCOLOGY

2. • Every matter consists of basic entities called
elements.
• Each element is composed of smallest particles called
‘ATOM’.
• Atom- the name is derived from Greek language
Atomos means ‘Not to be cut’.
ATOM:

3. Atom and ‘DEMOCRITUS’
• DEMOCRITUS was a Greek
philosopher who began search
for description of matter more
than 2400 years ago (4th century
B.C.)
• Democritus was the person who
first suggested the existence of
ATOM & coined the name
ÁTOMOS’ means Not to be cut
or Indivisible.

4. Atomic structure
• Atom consists of positively charged
NUCLEUS at the centre and negatively
charged Electrons revolving around it.
• Radius of an atom -- 10-10 m.
• Radius of the nucleus -- 10-15 m.
• Nucleus consists of Protons and
Neutrons together called Nucleons.
• Most of the mass of an Atom is
possessed by Nucleus.

5. Representation of an Atom:
Example:
235U92 X-Uranium
A-235
Z-92

6. Theories of Atomic models:
John Dalton
• Matter is made of indivisible atoms,
they are indestructible.
• All atoms of a given (same) element
are identical in their physical and
chemical properties.
• Atoms of different elements differ in
their physical and chemical properties.
• Atoms of different elements combine
in simple whole-numbers ratios to form
Molecules
• Chemical reactions consist of the
combination separation or
rearrangement of atoms
• Limitations: It could not explain
•Why and how do atoms combine together to
form compound atoms (molecules)
•The nature of forces which hold atoms
together in compound atoms
•Why atoms cannot exist in free state and why
compound atoms can exist freely.

7. Theories of Atomic models:
J.J.THOMPSON
Plum pudding model(1904)
• Negative particles are evenly scattered
throughout an atom with a positively charged
mass of matter.
• Similar to that of chocolate chip icecream
• Later proved to be incorrect.
Limitations:
• It could not explain the result of scattering experiment explained by rhe
Rutherford experiment.
• It did not have any experimental evidence in its support .
Thermionic emission, photoelectric emission and ionization were explained on this basis.

8. Theories of Atomic models:
RUTHERFORD
Rutherford Gold foil experiment setup:
• If a thin foil metal is placed in the path of the beam, the image becomes diffuse.
• This due to the scattering of α- particles by the atoms of the foil.
• The particles scattered in various directions were counted by scintillation
counter
• It was found that although most of the particles scattered through angles of
the order of 10 or less
• But a small number say about 1 in every 10,000 scattered through 900 or even
1800
• The experiment is known as Rutherford's alpha particle scattering experiment.

9. Results of the Rutherford experiment
(a) The results that the metal foil
experiment would have yielded if the
plum pudding model had been
correct.
(b) Actual results

10.  The whole of the positive charge of atom must be concentrated in a very
small space
 Atom is mostly hollow inside
 Since α- particles are positively charged, the part of the atom deflecting
them must also be positive
 Most of the mass of the atom is concentrated in the nucleus
 In this model, the mass of the atom (leaving the mass of its electrons) and
its whole positive charge are concentrated at the centre of the atom in a
nucleus of radius 10-15 m
 The electrons are distributed around the nucleus in a hollow sphere of
radius 10-10 m
Conclusion of Rutherford experiment

11. Regarding stability of atom
• Electrons revolving around the nucleus have centripetal
acceleration
• According to electrodynamics, accelerated charged
particles radiate energy in the form of electromagnetic
waves
• Hence electromagnetic waves should be continuously
radiated by the revolving electrons
• Due to this continuous loss of energy of the electrons, the
radii of their orbits should be continuously decreasing and
ultimately the electron should fall into the nucleus
• Thus atom cannot remain stable
• Rutherford’s model also failed to explain the Line
spectrum.
Drawback’s of Rutherford’s Model

12. Niels David Bohr
Theories of Atomic models:
Bohr’s theory (Postulates)
1. Fixed circular orbits :The electrons move around the nucleus in
concentric circular orbits .
2. While revolving in stable orbits, the electrons do not radiate energy
in spite of their acceleration towards the centre of the orbit.
3. Each of the fixed orbits is associated with a definite amount of
energy called stationary energy.
The energy levels are numbered as 1, 2, 3, 4… or designated as
K,L,M, N …
4. Jumping of an electron from one energy level to the other (ground state and excited state) .
5. Principle of quantization of angular momentum of the moving electron an electron can move
only in that orbit in which the angular momentum of the electron around the nucleus is an
integral multiple of h/2π.
• NIELS DAVID BOHR A Danish physicist who developed Bohr model of
atomic structure, in which he introduced the theory of electrons orbiting
around the nucleus.

13. I. No explanation for the spectra of multi electron systems:
Eg: He, Li
I. No explanation of fine spectrum of atoms:
III. No explanation for Zeeman and Stark effect : effect of
electric and magnetic fields on the spectral atoms.
•When a magnetic field is applied on an atom, its usually
observed spectral lines split. This effect is known as Zeeman’s
effect
•Spectral lines also get split in the presence of electric field. This
effect is known as Stark effect.
Limitations of Bohr’s Postulates

14.  Isotopes - Elements having same atomic number (protons) , but different
mass numbers (nucleons).
eg: 12C, 13C and 14C
are three isotopes of the element carbon with mass numbers 12, 13 and 14
respectively. The atomic number of all carbon isotopes is 6.
 Isobars - Elements having same mass number , but different atomic
numbers.
eg: 40S , 40Cl , 40Ar , 40K, and 40Ca
are isobars containing 40 nucleons; however, they contain varying atomic
number.
Classification of atoms:

15.  Isomers - Molecules with the same molecular formula but
different chemical structures. That is, isomers contain the
same number of atoms of each element, but have different
arrangements of their atoms in space.
eg: 131mXe54 is an isomer of 131 Xe54
m stands for meta-stable state
isomers represents identical atoms but they differ in nuclear energy states.
 Isotones -Nuclides having very same neutron number N ,
but different proton number Z.
eg:boron-12 Carbon-13 nuclei
both contain 7 neutrons

16. NUCLEAR STABILTY:
The Strong Force is exerted by
anything with mass to attract
other masses together and works
within a very short distance.
 Neutrons has no charge, but
have the strong force to bring
other nucleons together.
As a general rule, a nucleus will
need a neutron/proton ratio of 3:2
(or 1.5:1) in order to stay together.
This rule is more precise for larger
nuclei.
Of all known isotopes of natural
elements (about 1500), only 250
of them are stable

17.  .All of these stable isotopes have an atomic number in between 1 and 83.
 The mass of a nucleus will be less than the mass of all of the protons and
neutrons making it up. The difference is called the mass defect, which is
converted into energy if the nucleus is broken up.
 The amount of energy that keeps a nucleons together is called the Binding
Energy. This amount of energy is higher for nuclei that are stable than it
would be for unstable nuclei. (Joules)

18. Binding energy can be calculated by the formula E=mc2
(Einstein,s principle of equivalence of mass & energy relation)
Where,
c= speed of light ,m=mass & E = Energy
Eg: If a mass of 1 kg is converted into energy,
m=1kg speed of light is 3x108 m/s
E=1 kg x (3x108 m/s)2
E= 9 x 1016 J
Also,
Mass of electron at rest in terms of energy equivalent is given by,
m=9.1x10-31kg and speed of light is 3x108 m/s
E = 9.1x10-31 x (3x108)2
E=8.19x10-14 J
E=0.511 MeV

19.  Masses of atoms & atomic particles are
conveniently given in terms of amu.
 An amu is defined as 1/12 of the mass of the
12C6 Carbon atom.
1 amu= 1.66x10-27 kg
Mass expressed in terms of amu is known as
atomic mass or atomic weight.
Atomic mass and Energy Units

20. Charge and Mass of sub-atomic particles
Particle Charge of particle Mass of the particle
Proton 1.602176 x 10-19 Coulomb 1.00727 amu
Neutron Electrically neutral 1.00866 amu
Electron 1.602176 x 10-19 Coulomb 0.00548 amu
• Number of Protons = Number of Neutrons
• 1 amu = 1/12th of mass of an carbon atom.
• 1 amu = 1.66 x10-27 kg

21.  Avagadro's law- Every gram atomic weight of
a substance contains same number of atoms,
the number is referred as Avagadro's number
 Value of Avagadro's number
6.0221x1023 atoms per gram atomic weight
Gram atomic weight

22.  Basic unit of energy is joule (J)
 1 Joule is the work done when force of 1 Newton is
acting through a distance of 1 meter.
 Another energy unit in Nuclear physics is electron volt
(eV)
 1 eV is defined as the Kinetic energy acquired by an
electron in passing through a potential difference of
1V.
1eV=1Vx1.602x10-19 C
1eV = 1.602 x 10-19 J
Product of Potential difference and Charge of the electron
Atomic energy unit - Joule

23. DISTRIBUTION OF ORBITAL
ELECTRONS:
Postulates of Bohrs theory :
a) Electrons can exist in only those orbits for which angular
momentum of electron is an integral multiple of h/2Π
(h= plancks constant 6.62x10-34 )
b) No energy is gained or lost while an electron remains in any
one of the permissible orbits.
According to the model proposed by Niels
Bohr in 1913
Electrons revolve around the nucleus in
specific orbits.
They are prevented from leaving the
atom by the necessary centripetal force of
attraction between the positively charged
nucleus and negatively charged electron.
Bohr’s Atomic model

24. Arrangement of electrons in orbitals
Innermost orbit is called as K-shell.
Followed by orbital’s called L-shell, M-shell and N-shell.
The maximum no. of electrons in an orbital is given by the formula 2n2.
Eg: 1) Hydrogen atom has 1 electron in K-shell
2)Helium atom has 2 electrons in K-shell
3)Oxygen atom has 8 electrons (2 in K-shell, 6 in L-shell)

25. NUCLEAR FORCES:
There are four types of forces in nature
1)Strong nuclear force
2)Electromagnetic force
3)Weak nuclear force
4)Gravitational force
Strong nuclear force -Short range force(~10-15 m )

26. Energy level diagram of a particle in a nucleus
A)Particle with no
charge
B)Particle with positive
charge
U(r) - Potential energy
r - distance from centre of nucleus
B- barrier height
R- Nuclear radius

27.  Energy level diagram of decay of 60Co27 nucleus.
Nuclear energy levels:
60Co27 nucleus firstly emits β- particle with emission of photons
60Co27
60 Ni28 + e- + ν
Due to Nuclear transformation an neutron disintegrates in to a proton,an electron
& a neutrino

28. Thank you…
Vinay Desai
M.Sc Radiation Physics
Radiation Physics Department
KIDWAI MEMORIAL INSTITUTE OF ONCOLOGY
Bengaluru
E-mail:- vinaydesaimsc@gmail.com