1. STRUCTURE OF ATOM
Rutherford’s nuclear model of atom
Particle nature of electromagnetic radiation
(Plank quantum theory, Photoelectric effect)
Atomic spectrum (line spectrum of hydrogen)
Bohr model of atom for hydrogen
Dual behaviour of matter
Heisenberg’s uncertainty principle
Orbitals and quantum numbers
Filling of orbitals in atom
Stability of completely filled and half filled subshells
2. Rutherford’s Nuclear Model of Atom
Rutherford proposed an atom model based on his α–
particle scattering experiment.
He bombarded a very thin gold foil with α–particles.
The Experiment: A stream
of high energy α–particles
from a radioactive source
was directed at a thin gold
foil. The thin gold foil was
surrounded by a circular
fluorescent zinc sulphide
screen. Whenever α–
particles struck the screen,
a tiny flash of light was
produced at that point.
3. Observations Conclusions
Most of the α– particles passed through
the gold foil without any deviation.
most space in the atom is empty.
A small fraction of the α–particles was
deflected by small angles.
positive charge of the atom is
concentrated in a very small volume at
the centre called nucleus.
A very few α– particles (∼1 in 20,000)
bounced back, that is, were deflected
by nearly 180°.
The volume occupied by the nucleus is
negligibly small
4. The nuclear model (Planetary model) of atom
• All the positive charge and most of the mass of the atom
are concentrated in an extremely small region called
nucleus.
• Electrons are revolving round the nucleus with a very high
speed in circular paths called orbits.
• Electrons and the nucleus are held together by
electrostatic forces of attraction.
5. Drawbacks or Limitations of Rutherford’s atom
model
Rutherford’s model cannot explain the stability of the
atom.
He could not explain the electronic structure of atom.
6. Particle Nature of Electromagnetic Radiation:
Planck’s Quantum Theory
The important postulates of quantum theory proposed by Max
Planck are:
1. Atoms and molecules could emit or absorb energy not in a
continuous manner, but discontinuously in small packets of energy
called quanta or photons.
2. The energy (E ) of each quantum of radiation is proportional to its
frequency (ν).
It is expressed by the equation, E = hν Where ‘h’ is known as
Planck’s constant and its value is 6.626×10–34 J s.
7. Photoelectric effect
It is the phenomenon of ejection of electrons by certain metals
when light of suitable frequency incident on them.
The electrons ejected are called
photoelectrons.
8. The electrons are ejected from the metal surface as soon as the
beam of light strikes the surface. i.e., there is no time lag between
the striking of light beam and the ejection of electrons from the
metal surface.
The number of electrons ejected is proportional to the intensity or
brightness of light.
For each metal, there is a minimum frequency known as threshold
frequency [ν0], below which photoelectric effect is not observed.
The kinetic energy of the ejected electrons is directly proportional to
the frequency of the incident light.
10. Greater the energy possessed by the photon, greater will be
transfer of energy to the electron and greater the kinetic energy
of the ejected electron.
A more intense beam of light contains larger number of photons,
so the number of electrons ejected is also larger.
11. Line Spectrum of Hydrogen
When an electric discharge is passed through gaseous hydrogen,
the H2 molecules dissociate and the energetically excited hydrogen
atoms produced emit electromagnetic radiation of discrete
frequencies.
The hydrogen spectrum consists of several series of lines named
after their discoverers. The first five series of lines are Lyman,
Balmer, Paschen, Brackett and Pfund series.
12. Johannes Rydberg proposed an equation for finding the wave number of
the different lines in Hydrogen spectrum. The expression is:
n1- final energy level
n2 – initial energy level
13. The electron in the hydrogen atom can move around the nucleus in
circular paths of fixed radius and energy. These paths are called
orbits or stationary states or allowed energy states. These energy
levels are numbered as 1,2,3 etc or as K, L, M, N, etc. These
numbers are known as Principal quantum numbers.
BOHR’S MODEL FOR HYDROGEN ATOM
14. The energy of an electron in an orbit does not change
with time. However, when an electron absorbs energy, it
will move away from the nucleus (i.e. to a higher energy
level) and when it loses energy, it will move towards the
nucleus (i.e. to a lower energy level).
15. The radius of orbits can be given by the equation:
where a0 = 52.9 pm.
Thus the radius of the first stationary state is 52.9 pm
(called the Bohr radius). As n increases, the value of r will
increase.
16. The frequency of radiation absorbed or emitted when transition
occurs between two stationary states that differ in energy by ΔE, is
given by:
18. Dual Behaviour of Matter – de Broglie’s equation
de Broglie proposed that like radiation, matter also show both particle
and wave nature. This is known as dual behaviour of matter. i.e.
electrons should have momentum as well as wavelength. He gave the
following relation between wavelength (λ) and momentum (p) for
material particles.
19. What will be the wavelength of a ball of mass 1kg moving with a velocity
10m/s.
20.
21. Heisenberg’s Uncertainty Principle
“it is impossible to determine simultaneously, the exact position and
momentum (or velocity) of a moving microscopic particle like
electron”.
Where Δx is the
uncertainty in position
and Δp is the
uncertainty in
momentum and Δv is
the uncertainty in
velocity of the particle.
22. If the error in position of an electron is 1.1A0, find error in measurement of
momentum
23.
24. Quantum mechanical model of atom/ Wave
mechanical model of atom
The branch of science that
takes into account the dual
behavior of matter is called
wave mechanics
This model describes
electron as a 3
dimensional wave in the
electric field of positively
charges nucleus
The equation for such a
wave is derived by
Schrodinger.
25. The solution of Schrodinger equation provides a set of numbers,
known as quantum number, which describes energies of electrons
in atoms, information about the shape and orientation of orbitals
These are principal quantum number, azimuthal quantum number
and magnetic quantum number
n, l, m, s
Except to these, there is an additional quantum number called
spin quantum number which give information about the spin of
electrons in the orbitals
26. Principal quantum number (n) : size of orbital
It is a positive integer with value n= 1,2,3….
It identifies the shell, if n=1, K shell; n=2, L shell; n=3, M shell……
The value of energy of electron increases with increasing value of n
27. Azimuthal / angular momentum quantum number (l):
subshell
It gives idea about the shape of orbital. Value: 0 n-1
Azimuthal q.no
value(l)
Name of orbital shape
0 s Spherical
1 p Dumb bell
2 d Double dumb bell
3 f 8 fold
28. Magnetic quantum number (m) : orientation of orbital
Value of m varies from –l to +l
l
value
Name of
orbital
m values Number
of orbitals
Orientation of orbitals
0 s 0 1
1 p -1,0,+1 3
2 d -2,-1,0,+1,+2 5
3 f -3,-2,-1, 0, +1,
+2,+3
7
Number of orbitals in a given subshell + 2l+1
29. n l (0-> n-1) Name of
orbital
m value Orbital
notation
Representation
1 0 s 0 1s
2 0
1
s
p
0
-1,0,+1
2s
2p
3 0
1
2
s
p
d
0
-1,0,+1
-2,-1,0,1,2
3s
3p
3d
4 0
1
2
3
s
p
d
f
0
-1,0,1
-2,-1,0,1,2
-3,-2,-1,0,1,2,3
4s
4p
4d
4f
30. How many subshells and orbitals are present in n = 3?
n l m number of orbitals
3 0 0 1
1 -1,0,1 3
2 -2,-1,0,1,2 5
Number of orbitals= 9
Short cut = n2
31. Using s, p, d, f notations, describe the orbital with the following quantum
numbers
n=2 l=1
n=4, l=0
n=5,l=3
n=3, l=2
Find the number of electrons for which n=2, m=0, s=+1/2
