7. Thomson’s Model
An atom has a spherical shape
(radius～10–10 m).
Positive charge is uniformly distributed
throughout the sphere.
Negatively charged electrons are
embedded in it like raisins in a
pudding.
8. Thomson’s Model
Mass of the atom is assumed to be
uniformly distributed all over the
atom.
Plum pudding Watermelon
9. Thomson’s Model
Electrons are embedded in an atom in such a way that
the most stable electrostatic arrangement is achieved.
Positively-
charged matter
Electrons
11. A stream of high energy 𝞪–particles was
directed at a thin gold foil (thickness ∼ 100 nm).
Rutherford’s Experiment
Radioactive
source
ZnS
screen
Gold foil
13. Observation Conclusion
Most 𝞪-particles passed
through the foil without
deflection
Presence of large
empty space
in the atom
Observation Conclusion
Few 𝞪-particles
were deflected by small
angles
Positive charge is
concentrated in a
very small region
Observation Conclusion
Very few 𝞪-particles
(∼1 of 20,000)
deflected at 180°
Small positively -
charged core at
the centre
16. α particles can be represented by :
B) He2+
C) He
A) Li+
D) None of the above
17. Imagine that in any atom about 50% of the space is occupied
by the atomic nucleus. If a silver foil is bombarded with α-
particles, majority of the α-particles would :
B) be absorbed by the nuclei
C) pass through the foil undeflected
A) be scattered
D) get converted into photons
18. Nucleus
An atom consists of
a small positively -
charged core at the
centre which
carries almost the
entire mass
of the atom.
It has negligible
volume as
compared to the
volume
of the atom.
19. Nucleus
Radius of the atom
Radius of the nucleus
～10–10 m
～10–15 m
Both protons and neutrons present in the
nucleus are collectively called nucleons.
20. Nucleus
R R0 (A)
=
R = Radius of nucleus of an element
A = Mass number of element
R0 = 1.11 x 10-15 m to 1.44 x 10-15 m
1
3
21. Extranuclear part
Nucleus is surrounded by
revolving electrons.
Electrons and
nucleus are held
together by
electrostatic forces
of attraction.
24. Rutherford’s Model: Drawbacks
It could not explain the line spectrum of the H atom.
It could not explain the electronic structure of the atom.
25. What is the ratio of specific charge of a proton to that of an
𝜶-particle?
B) 1 : 2
C) 1 : 4
A) 2 : 1
D) 4 : 1
26. How much time does light take to travel 1.0 m?
B) 3.3 microsecs
C) 3.3 nanosecs
A) 3.3 millisecs
D) 3.3 seconds
27. Smallest quantity of energy that can
be emitted or absorbed in the
form of EM radiation
Packet or bundle of energy
Quantum of radiation
Photon
Planck’s Quantum Theory
28. E h𝝂
= = h h = Planck’s constant
= 6.626 × 10-34 Js
Quantum Theory of Radiation
c
λ
E nh𝝂
=
n = Number of photons
= 0, 1, 2, 3, ….
_
30. An electric bulb marked as 60 watt emits light of wavelength
3000 Å. If 25% of the energy is emitted as light, what is the
number of photons emitted in one second?
B) 7.32 x 1020
C) 2.27 x 1019
A) 2.27 x 1020
D) 7.22 x 1019
31. 31
What is the equation E = h𝝂 indicate?
B) Photons are waves.
C) Photons are stream of particles.
A) Photons have both particle and wave nature.
D) No such inference can be drawn from the given
equation.
35. Which of the following statements is/are true in the context of
photoelectric effect?
A) The kinetic energy of ejected electron is independent of the
intensity of radiation.
B) It provided evidence for the quantum nature of light.
C) The number of photoelectrons ejected depends upon the
intensity of the incident radiation.
D) The kinetic energy of the emitted electrons depends on the
frequency of the incident radiation.
36. Electromagnetic radiation having λ = 310 Å is subjected to a
metal sheet whose work function = 12.8 eV. What will be the
velocity of the photoelectron having the maximum kinetic
energy?
B) 4.352 ✕ 106 m/s
C) 3.09 ✕ 106 m/s
A) No emission will occur
D) 8.72 ✕ 106 m/s
45. n = 1
n = 2
n = 3
n = 4
n = 5
n = ∞
13.6
eV
12.09
eV
-13.6 eV
-3.4 eV
-1.51 eV
-0.85 eV
-0.54 eV
0 eV
Energy Level Diagram for H atom
10.2
eV
3.4
eV
12.75
eV
13.06
eV
46. Rydberg’s Formula
RH Z2 x
1
λ
1, 2, 3, ...
n1
n1 + 1, n1 + 2, ...
n2
For any atom
1
n1
2
_ 1
n2
2
=
λ
For any atom
= 1
n2
2
1
n1
2
_
912
Z2
Å
47. Line Spectra is a characteristic of :
B) atoms
C) radicals
A) molecules
D) none of these
48. The wavelength of a spectral line for an electronic transition is
inversely proportional to :
B) the nuclear charge of the atom
C) the velocity of an electron undergoing transition
A) the number of electrons undergoing transition
D) the difference in the energy involved in the transition
49. de Broglie Wavelength (λ)
h
p
h
mv
λ = =
Momentum
of particle
p
Mass
of particle
m
Velocity
of particle
v
Planck’s
constant
h
p
0
λ
50. A beam of helium atoms move with a velocity of 2 × 104 m s–1.
Find the wavelength (in pm) of the particles constituting the
beam.
B) 3
C) 6
A) 4
D) 5
51. de Broglie’s Equation and Kinetic Energy
Multiplying both sides by m and rearranging
K.E. =
m2v2 2 K.E. × m
=
mv
mv2
1
2
√2 K.E. x m
=
52. de Broglie’s Equation and Kinetic Energy
h
λ =
Since, mv 2 K.E. x m
h
p
h
mv
λ
2 K.E. × m
de Broglie
equation
= √
√
= =
53. If the kinetic energy of a proton is increased to nine times, the
wavelength of the de Broglie wave associated with it would
become :
times
times
1
3
1
9
B) 9 times
C)
A) 3 times
D)
54. The mass of an electron is 9.1 × 10–31 kg. If its K.E. is
3.0 × 10–25 J, calculate its wavelength.
B) 0.89 × 10-7 m
C) 0.89 × 10-5 m
A) 0.89 × 10-6 m
D) 0.89 × 10-8 m
57. Heisenberg’s Uncertainty Principle
Δx . Δp
4𝜋
h
≥
Δx . m . Δv
Uncertainty in
position
Δx
Mass of
particle
m
Uncertainty in
momentum
Δp
Uncertainty in
velocity
Δv
≥
4𝜋
h
58. Principle of Optics
λ
Minimum error in the
position measurement (Δx) =
If a light (wavelength ‘λ’) is used to locate the
position of a particle, then
+
-
67. A golf ball has a mass of 40 g, and a speed of 45 m s-1. If the
speed can be measured within accuracy of 2 %, calculate the
uncertainty in the position.
69. Significance of Uncertainty Principle
2
Rules out the existence of
definite paths of electrons
3
Introduced concept of probability
of finding the electrons
1
Not an instrumental error, rather
conceptual error
72. The uncertainty in position of a particle of 25 g in space is 10-5 m.
What will be the uncertainty in velocity (m s-1) ?
B) 2.1 × 10-34
C) 0.5 × 10-34
A) 2.1 × 10-28
D) 5 × 10--24
73. The Heisenberg’s uncertainty principle is significant for :
B) a football
C) a jet aeroplane
A) a cricket ball
D) an electron
74. Quantum Numbers
Set of four numbers
required to define
an electron in an
atom completely
75. Quantum Numbers
Principal Quantum Number (n)
1
Azimuthal Quantum number (l)
2
Magnetic Quantum Number (ml )
3
Spin Quantum Number (s)
4
76. Principal Quantum Number (n)
Describes the size
of electron wave &
the total energy of
the electron
n = 1, 2, 3,...
Represented as K, L, M, N,...
77. Azimuthal Quantum Number (l )
2
Energy of the orbital in a
multielectron species (both n & l)
1
Designates the subshell to which
the electron belongs
Describes the 3-D
shape of the orbital
or the electron
cloud
85. Subshell Representations
n
1
l
0
2 0, 1
3 0, 1, 2
4 0, 1, 2, 3
Subshell
notation
1s
2s, 2p
3s, 3p, 3d
4s, 4p, 4d, 4f
Number of subshells in the nth shell n
86. Azimuthal Quantum Number (l )
Orbital angular
momentum (L)
l (l + 1) ħ
Subshell
s
Orbital angular
momentum
0
p 2 ħ
d 6 ħ
ħ =
h
2π
=
87. Magnetic Quantum Number (ml )
2 Describes the orientation of orbitals
3
Accounts for the splitting of lines of
atomic spectrum in magnetic field
1
Designates the orbital to which
the electron belongs
88. Magnetic Quantum Number (ml )
Subshell
s
Number of orbitals
1
p 3 (px, py, pz)
d 5 (dxy, dyz, dzx, dx - y ,dz )
2
2 2
f 7
2l + 1
Maximum number of
orbitals in a subshell =
2
92. Remember!
Maximum number of
electrons in a subshell
2 (2l + 1)
Subshell
l
s
0
Number of electrons 2
p
1
6
d
2
10
f
3
14
=
An orbital can accommodate maximum of 2 electrons.
93. Spin Quantum Number (s)
Spin of an electron
s = +
1
2
s =
1
2
Presence of two closely-spaced lines in
atomic spectrum
96. What is the total number of atomic orbitals in fourth energy
level of an atom?
B) 16
C) 32
A) 8
D) 4
97. Which of the following shell can contain maximum of 72
electrons?
B) 6th
C) 4th
A) 5th
D) 3rd
98. Which of the following statements concerning the four
quantum numbers is false?
B) l gives the shape of an orbital.
C) m gives the energy of the electron in the orbital.
A) n gives an idea of the size of an orbital.
D) s describes the spin of the electron.
99. Rules for Filling of Electrons in Orbitals
Rules
Aufbau principle
Hund’s rule of maximum
multiplicity
Pauli’s exclusion principle
101. Energies of Subshells of H-like Species
1s < 2s = 2p < 3s = 3p = 3d
< 4s = 4p = 4d = 4f < …
Order of
energy
Energy of single electron species depends only on the
Principal quantum number
104. Energy of Subshells of Multielectron Species
Different subshells have different energy
which depends on:
Principal quantum
number
Azimuthal quantum
number
105. ( n + l ) rule or Bohr-Bury’s Rule
Lower the
energy of subshell
Lower value of
(n + l )
106. Example
3d n + l 3+2 5
= = =
4s n + l 4+0 4
= = =
4s 4s
>
Value of n + l for :
4s will be filled before 3d
107. ( n + l ) rule
Subshell with lower ‘n’
value has lower energy
Two subshells with
same (n + l) value
108. Example
3d n + l 3 + 2 5
= = =
4p n + l 4 + 1 5
= = =
3d 4p
=
Value of n + l for :
3d will be filled first due to
the lower ‘n’ value
114. In a 3d subshell, all the five orbitals are degenerate. What does
it mean?
B) All the orbitals have the same shape.
C) All the orbitals have the same energy.
A) All the orbitals have the same orientation.
D) All the orbitals are unoccupied.
115. Degeneracy of the second excited state of H is y. Find the
value of y.
B) 9
C) 5
A) 3
D) 4
121. Hund’s Rule of Maximum Multiplicity
Deals with the filling of electrons
In the degenerate orbitals of
the same subshell
122. Hund’s Rule of Maximum Multiplicity
No electron pairing takes place in
the orbitals in a subshell
Until each orbital is occupied
by 1 electron with parallel spin
123. Hund’s Rule of Maximum Multiplicity
Hund’s rule is an empirical rule
Determines the lowest energy
arrangement of electrons
124. Why Maximum multiplicity?
Maximum spin of
an atom (S)
1
2
Spin Multiplicity (S.M.) 2S + 1
=
= x n
Spin Multiplicity (S.M.) Stability
128. sa pb dc ... Notation
Similar subshell represented for
different shells is differentiated by
Writing the principal
quantum number before the
respective subshell
129. sa pb dc ... Notation
2 p
5
Subshell ‘p’ belongs
to 2nd shell (n = 2)
Subshell ‘p’, l = 1 Number of
electrons = 5
F
19
9
Fluorine
1s2 2s2 2p5
130. Orbital Diagram Notation
1s2 2s2 2p5
Electron is Represented By
Upward arrow ( or ↿) Downward arrow ( or ⇂)
Represent all 4 Quantum numbers
⥮ ⥮ ⥮ ↿
⥮
F
145. Symmetry
Consequently, their shielding of one
another is relatively small
Electrons are more strongly
attracted by the nucleus
Have less energy
and more stability
146. Exchange Energy
Tends to exchange their positions
Energy released when two or more electrons
with the same spin in the degenerate orbitals
147. Exchange Energy
Number of exchanges that
can take place is maximum
When subshell is either half
filled or fully filled.
148. Exchange Energy
↿ ↿ ↿ ↿ ↿ ↿ ↿ ↿ ↿ ↿
↿ ↿ ↿ ↿ ↿ ↿ ↿ ↿ ↿ ↿
4 exchange by electron ‘a’ 3 exchange by electron ‘b’
2 exchange by electron ‘c’ 1 exchange by electron ‘d’
a b
c d
149. The orbital diagram in which both Hund’s rule and Aufbau
principle are violated is:
⥮ ⥮
↿
↿
⥮
⥮ ↿ ↿
⥮ ↿
↿↿ ↿ ↿
↿
2p 2p
2p 2p
2s
2s
2s 2s
150. Which rule is/are violated in the given electronic configurations?
⥮ ↿
⥮
↿ ↿
⥮ ↿
↿↿ ↿ ↿
↿
↿ ⥮ ⥮
2p
2s
2p
2s
2p
2s
2p
2s
Aufbau Pauli’s Hund’s
