2. Overview of presentation
• Bohr’s atomic model
• Postulates of Bohr’s atomic model
• Successes of Bohr’s model
• Calculations based on Bohr’s model
• Limitations of Bohr’s Model
• Bibliography
3.
4. Bohr’s Atomic Model
• An atom is made up of three particles,
electrons, protons and neutrons. Electrons
have a negative charge and protons have a
positive charge whereas neutrons have no
charge. They are neutral. Due to the
presence of equal number of negative
electrons and positive protons, the atom as
a whole is electrically neutral.
5. • The protons and electrons are located in a
small nucleus at the centre of the atom. Due
to the presence of protons, the nucleus is
positively charged.
• The electrons revolve rapidly around the
nucleus in fixed circular paths called energy
levels or shells. The 'energy levels' or 'shells'
or 'orbits' are represented in two ways:
either by the numbers 1, 2, 3, 4, 5 and 6 or
by letters K, L, M, N, O and P. The energy
levels are counted from centre outwards.
6. • Each energy level is associated with a fixed
amount of energy. The shell nearest to the
nucleus has minimum energy and the shell
farthest from the nucleus has maximum
energy.
• There is no change in the energy of
electrons as long as they keep revolving
with the same energy level. But, when an
electron jumps from a lower energy level to
a higher one, some energy is absorbed while
some energy is emitted.
7. • When an electron jumps from a higher
energy level to a lower one, the amount of
energy absorbed or emitted is given by the
difference of energies associated with the
two levels. Thus, if an electron jumps from
orbit 1 (energy E1) to orbit 2 (energy E2), the
change in energy is given by E2 - E1.
8. • The energy change is accompanied by
absorption of radiation energy of E = E2 E1 =
h where, h is a constant called 'Planck's
constant' and is the frequency of radiation
absorbed or emitted. The value of h is 6.626
x 10-34
J-s
. The absorption and emission of
light due to electron jumps are measured by
use of spectrometers.
9. Postulates of Bohr’s Atomic
Model
• Electrons revolve round the nucleus with
definite velocities in concentric circular
orbits situated at definite distances from the
nucleus. The energy of an electron in a
certain orbit remains constant. As long as it
remains in that orbit, it neither emits nor
absorbs energy. These are termed
stationary states or main energy states.
10. • Bohr proposed that the angular momentum
of an electron is quantized. Thus, the
motion of an electron is restricted to those
orbits where its angular momentum is an
integral multiple of h/2π, where h is
Planck’s constant.
• Thus we have the relationship mvr = nh/2π,
where m is mass of electron, v is velocity of
electron of said orbit, r is radius of that
orbit, n is a simple integer.
11. • The stationary states or allowed energy
levels are only those where n = 1, 2, 3, ……
This is called Bohr quantum condition.
• The energy of an electron changes only
when it moves from one orbit to another.
An electronic transition from an inner orbit
to outer orbit involves absorption of energy.
Similarly, when an electron jumps from an
outer orbit to inner orbit it releases energy,
which is equal to the difference between the
two energy levels.
12.
13. • The energy thus released in the form of a
radiation of a certain frequency appears in
the form a line in the atomic spectrum. If
the energy of an electron in the outer orbit
(n2) is E2 and energy of electron in the inner
orbit (n1) is E1 then E2 - E1 = ΔE = hν.
• The value of n could be small integers 1, 2, 3
and these correspond to the first, second,
third, and so on. Quantum states are shells
for the electron; n is termed as principal
quantum number.
14. • Based on the Bohr theory ,Bohr calculated
the radii of the various orbits and the
energies associated with the electrons
present in those shells.
15. Successes of Bohr’s Model
• When electron jumps from lower energy
level to higher energy level, it absorbs
relevant amount of energy and this results
in the absorption spectrum.
16. • When an electron drops to higher level
from lower level, it emits some amount of
energy and emission spectrum is observed.
• Since there is only one electron in hydrogen
atom, there should be one line in hydrogen
spectrum. But in Bohr theory, there are
infinite number of orbits, so more than one
line is observed in spectrum.
17. Calculations based on Bohr’s
Model
• Radius of nth orbit
According to bohr model, the attraction
force between electron and nucleus is
balanced by centrifugal force of electron
which is due to motion of electron and tend
to take electron away from nucleus.
18.
19.
20.
21.
22. • Since the value of Z is constant for an atom,
r n α n2 , so radius increases with increasing
the value of n.
• If the value of n is constant , rn α 1/Z
• Hence, radius of a particular orbit
decreases with increasing the atomic
number.
23.
24.
25. • Energy of electron in nth
orbit
According to Bohr atomic model, the
maximum energy value of electron at
infinite is zero because of negligible
attraction force between electron and
nucleus at infinite distance.
Hence, as electron comes closer to
nucleus, the energy becomes negative.
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30.
31. Limitations of Bohr’s Model
• Bohr model could not explain those atoms
which have more than one electron like
lithium, helium. This model was applicable
only for those atoms which have one
electron.
• Bohr theory explained only spherical orbits.
There was no explanation for elliptical
orbits.
32. • This model failed to explain Zeeman Effect
and stark effect.
• Bohr model could not explain the
uncertainty principle of Heisenberg.
33. • Bohr model was not related with
classification and periodicity of elements.
• By using Bohr atomic model, one can’t
explain the intensity of spectrum line.
34. • Bohr model could not explain the wave
nature of electron. It explained only particle
nature of electron.
