2. Introduction/Historical background
Early Hindu (Maharishi Kapila, Maharshi Kanada) and Greek Philosophers (Leucippus, Democritus)
were the originators of the concept of atoms.
Greek word atomos (a = absence, tomos = cut)
Dalton’s Atomic Theory (1808)
It states that all matters are made up of small, indivisible particles known as ‘atoms’.
Postulates:
1. All matter consists of tiny indivisible particles called atoms.
2. All atoms of a specific element are identical in all respects. However, atoms of different elements
exhibit different properties and vary in mass and size.
3. Atoms can neither be created nor destroyed. (indestructible, indivisible)
4. Atoms of different elements can combine with each other in fixed whole-number ratios in order to
form compounds.
5. Atoms can be rearranged, combined, or separated in chemical reactions.
ATOMIC STRUCTURE
3. Limitations:
Dalton’s Atomic Theory
1. It does not account for subatomic particles.
2. It does not account for isotopes and isobars.
3. Elements need not combine in simple whole-number ratios to form
compounds.
4. The theory does not account for allotropes.
Merits:
1. The law of multiple proportions, the law of conservation of mass, and the law of
constant proportions are not violated by Dalton’s atomic theory.
2. The theory provides a basis to differentiate between elements (atoms) and
compounds (molecules).
4. Discovery of the electron
J.J. Thomson
Discovery of the proton
E. Goldstein
Discovery of neutron
James Chadwick (1932)
Late discovery due to electrically neutral nature
7. Atomic models
1. JJ Thomson’s Atomic model
2. Rutherford’s Atomic model
3. Bohr’s Atomic model
4. Bohr-Sommerfeld atomic model (Modification of Bohr’s atomic
model)
5. Wave mechanical and quantum mechanical model of the atom:
Quantum theory
8. Only few properties of atoms can be understand from this model thus, this theory
was soon rejected.
9. Rutherford’s Atomic model
(α-ray scattering experiment)
This experiment led to the discovery of nucleus so it is called nuclear
model of atom..
10. α-particles are shot out from radio active elements with
very high speed.
They are the nucleus of helium atoms. It has a charge of
+2 and mass 4 amu.
15. 1. Inability to explain the stability of atom:
According to this model, electron revolves around the positively
charged nucleus in a circular path.
As the electron revolves in a circular orbit, it is constantly subjected to
centripetal acceleration and radiates energy continuously ( as per
Maxwell's EM theory) . As a result due to this continuous loss of
energy, the electrons should follow spiral path towards the nucleus
and fall into it. Hence atoms must collapse, but they are stable.
Limitations
Radiation of energy by moving electron should give continuous atomic spectra but this doesn’t happen
as atoms give discontinuous line spectra which represent radiation of different frequencies.
16.
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19.
20. Thomson’s atomic model and Rutherford’s atomic model failed to answer any questions related to
the energy of an atom and its stability.
In the year 1913, Niels Bohr proposed an atomic structure model, describing an atom as a small,
positively charged nucleus surrounded by electrons that travel in circular orbits around the
positively charged nucleus as planets around the sun in our solar system, with attraction provided
by electrostatic forces, popularly known as Bohr’s atomic model.
Bohr theory Applicable to modified Rutherford’s atomic model by explaining that electrons move in
fixed orbitals (shells) and not anywhere in between and he also explained that each orbit (shell) has
a fixed energy level.
Rutherford basically explained the nucleus of an atom and Bohr modified that model into electrons
and their energy levels.
Bohr’s atomic model
21. 1. Stationary orbit
In an atom, electrons (negatively charged) revolve around the positively charged
nucleus in a in a fixed circular path or orbits or shells or energy level termed
“stationary orbit”.
Postulates of Bohr’s atomic model
22. Only those orbits are possible for which angular momentum of the electron is equal to the
integral multiple of h/2π.
Mathematically it can be expressed as
Where n = integer ( 1 , 2, 3, …… denotes no. of shell or orbit)
h = plank’s constant (6.624 × 10-27 erg-seconds)
m = mass of electron
v= velocity of electron
2. Quantization of angular momentum
23. 3. Quantization of energy.
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.
When an electron revolves around the nucleus in a fixed orbit, the electrons neither emit or
nor absorbs the radiation, i.e. energy of an electron is quantized.
24. An electron emits or absorbs energy when it jumps from one orbit or energy level to another.
The energy absorbed or emitted is equal to the difference between the energies of the two energy
levels (E1, E2) and is determined by Plank’s equation.
ΔE = hv
Where, ΔE = E2-E1
ΔE = energy absorbed or emitted
h= Plank’s constant
v= frequency of electromagnetic radiation emitted or absorbed
25. 4.Origin of atomic spectra.
(b)
The electrons in an atom move from a lower energy level to a higher energy level by gaining the
required energy and an electron moves from a higher energy level to lower energy level by losing
energy.
The released energy corresponds to the radiation of different wavelength.
26. Applications of Bohr’s Model of an Atom
• To explain the stability of atom.
• To explain the line spectra of hydrogen.
• To calculate the energy and velocity of electron.
• To calculate the radius of atomic orbital.
27. Limitations of Bohr’s Model of an Atom
• Bohr’s theory is applicable only to hydrogen-like species containing one electron
only e.g. Li2+. Bohr’s theory is applicable to hydrogen like atoms (single electron
system). Li2+ and H-atom consists of only one electron. He, He2+ consist of 2, 0
electrons respectively
• 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.
• It could not explain the spectra obtained from larger atoms.
